Full text of "General Chemistry"

See other formats

GENERAL CHEMISTRY

Walter E. Miller

Associate Professor

Joseph A. Babor

Professor Emeritus

Department of Chemistry
The City College of The City University of New York

SCIENTIFIC BQOK AGENCY

22 RAJA WOODMUNT STREET,

CALCUTTA- I.

Indian Edition first published 1968 by Scientific
Book Agency, Calcutta-1, India

Original English language edition publish'd by Wnu
C. Brown Company Publishers, Dubuque, Iowa,
52001. U.S.A.

Library of Congress Catalog Number* 65-21888

All rights reserved. No part of this book may be
reproduced in any form or by any process without
permission in writing from the copyright owner

This book has been published with the assistance
of the Joint Indian-American Standaid Works
Programme

Sales area of Indian Edition : India, Southeastern
Asian Countries, The Philippines, Continent of
Africa

Price in India Rupees Eight and Fifty Paise

Printed by Photo-offset at S. Antool & Co. Private
Ltd. 91, Acharya Prafulla Chandra Road, Calcutta-9

Preface

This -text is traditional in its approach and in its viewpoint. It is a revision
of the well-known Babor textbooks in general chemistry used by thousands
of students for over a quarter of a century to incorporate the current concepts
of chemical theory. The book presents the major fundamental concepts and
such descriptive chemistry a beginning student may be expected to know
and which he should be able to grasp at the end of one year. The book
does not attempt to include every facet of modern chemical theory. The
authors have not catered to popular appeal. They have not included only
that which is intellectually interesting or theoretically delightful. Merely be-
cause a subject is dull it has not been omitted. What a student should know
and what he can understand meaningfully have been the sole criteria for
inclusion in the subject matter.

To a great extent the historical approach is relied upon to give the stu-
dent an insight to the romance of chemistry, the development of chemical
concepts, and the operation of the scientific method. Whereas some may
think that the laws of chemical combination and the development of atomic
weights are outmoded and belong to an earlier generation, the authors
believe that^ these subjects offer excellent examples of chemical reasoning from
experimental data. Hence much emphasis is given throughout the book to
experimental data to show the foundations upon which theory is built. Many
students come to look upon the laboratory as a place where work is done by
an inferior breed and do not quite realize that, despite its marvelous theoreti-
cal advances, chemistry is still primarily an experimental science. If there
is any theme to the textbook it is that chemistry is an experimental science
and that theory is based upon experimental data!

The bulk of the theoretical subject matter is in the first twenty-five chapters
which may well be used for an independent course in the principles of
chemistry. Special attention has been given to the subjects of chemical

Preface

equilibrium and kinetics, and a brief chapter on qualitative analysis has
been included. In many cases both ionic and molecular equations are given
side by side. Though the student is advised that only ions ( and not molecules )
exist in 'Certain compounds, molecular equations are given illustratively to
avoid the viewpoint that reagent bottles of ions alone exist on the laboratory
shelf. In the solution of illustrative problems particular attention has been
given to dimensions and significant figures.

Essential descriptive chemistry has not been minimized though much of
the descriptive chemistry of elements and compounds has been compressed
into tabular form. For the nonchemistry major in particular, there will prob-
ably be no course other than general chemistry wherein a basic knowledge
of industrial practice will be obtained to enable him to function in later life
as an intelligent “citizen chemist."

In this revision we have aimed at the same general student level as
in previous editions. Though there are sections which require some knowledge
of mathematics, the authors believe that little material will be found in this
textbook which cannot be mastered by the average college freshman if he
makes an honest and conscientious effort.

We would like to acknowledge our deep gratitude to Professor Emil
Margolis for his valuable suggestions and to Elizabeth Zawadsky for her
secretarial assistance.

Walter E. Miller
Joseph A. Babor

New York. N. Y.

April, 1965

Contents

Chapter Page

1. Introduction: The Philosophy of Science 1

1. The Scientific Method. 2. Scientific Explanation. 3. Matter and Mass.
4. Properties and Conditions. 5. Elements, Compounds, and Mixtures.
6. Abundance of the Elements. 7. Units of Measurement. 8. Temperature and
Heat. 9. Pressure.

2. The States of Matter: Cases 13

1. The Gaseous State. 2. Behavior of Gases. 3. Change in Volume of a Gas
with Pressure: Boyle's Law. 4. Change in Volume of a Gas with Tempera-
ture: Charles' Law; Absolute Temperature. 5. Simultaneous Correction for
Temperature and Pressure. 6. The Ideal Gas Law. 7. Dalton’s Law of Partial
Pressures. 8. Graham’s Law of Diffusion. 9. The Kinetic Molecular Theory.

10. Explanation of the Gas Laws by the Kinetic-Molecular Theory. 11. De-
viations from the Gas Laws. 12. Critical Temperature and Pressure.

3. The States of Matter: Solids and Liquids 33

L The Solid State. 2. The Unit Cell. 3. Melting and Phase Changes.

4. Crystal Binding. 5. Surface Tension. 6. Viscosity. 7. Vapor Pressure.

8. Changes of State. 9. Diffusion in Liquids* and Solids.

4. Chemical Change 49

1. A Chemical Change — The Rusting of Iron. 2. Law of Conservation of
Mass. 3. Law of Definite -Proportions. 4. Law of Multiple Proportions.

5. Dalton’s Atomic Theory. 6. Explanation of the Laws of Chemical Change.

5. Atomic and Molecular Weights 55

1. Gay-Lussac’s Law of Combining Volumes. 2. Avogadro’s Hypothesis.

3. Molecular Weights of Gases. 4. The Avogadro Number and the Mole.

5. Weight and Volume Relationship of Gases. 6. Determination of Atomic
Weights.

viii

Content's

Chapter

6. Chemical Calculations

1. Symbols and Formulas 2 Equations 3. IVi ventage Composition from a
Formula. 4 The Formula from the Percentage Composition 1 Chemical
Equations. 6 Weight— Weight Pioblems 7 Weight— Volume Problems.
8. Volume— Volume Problems. 9 Valence. 10. Equivalence

7, Energy Changes in Chemical Reaction 81

X. Conservation of Energy. 2 Measurement oi the Enow ot a Che meal
Change. 3. Thcrmoehenutal Equations 4 Hess’ Law ot Constant Heat
Summation 5. Heat of Formation.

8. The Kinetics of Chemical Reactions • • • * 93

1. Factors Affecting the Rate of Chemical Reaction 2 Reaction Rate
Theory. 3 Classification of Reactions 4. First Orclei Reactions

9. Chemical Equilibrium * 107

1. Reversible Reactions 2. Chemical Equilibrium 3 Factors Affecting the
Point of Equilibrium. 4. The Concentration Factor 5. The Equilibrium
Equation. 6. Heterogeneous Equilibrium. 7. The Pressure Factor. S. ’flit 1
Temperature Factor 9. Le Chateliei’s Principle.

10. The Periodic Law 125

1. Periodicity and Early Attempts at Classification. 2. T he Periodic Table.

3. Uses of the Periodic Table, t. Theoretical Meaning oi the Penodu Table.

11. Atomic Structure—I: The Nucleus 135

I. Early Concept of the Atom. 2. Discovery of the Electron. 3 Discovery
of X-rays. 4. Electromagnetic Radiation. 5. Discover} of Radioactivity.

6. Types of Rays from Radioactive Material. 7. The Proton and the Neu-
tron. 8. The Nuclear Atom. 9. Atomic Numbers. 10 Nuclear Structure.

II. Mass Spectrograph. 12. Isotopes. 13. Atomic Weight Scales.

12. Atomic Structure— II: Electronic Structure 150

I. The Neutral Atom. 2. Bohr Theory 3. Quantum Mechanics. 4. Quantum
Numbers. 5. Pauli Exclusion Principle. 6. Summary. 7 Electron Configura-
tions of the Atoms. 8. Electron Configuration and Atomic Properties.

9. Electron Configuration and the Periodic Table 10. Origin of Spectra.

II. The Spectroscope.

13. The Chemical Bond 169

1. The Octet Rule. 2. The Ionic Bond. 3. The Covalent Bond. 4. The Co-
ordinate Covalent Bond. 5. Oxidation Number. 6. Partial Ionic Character
of Covalent Bonds. 7. Polarity of Covalent Molecules. 8. The Hydrogen
Bond. 9. Atomic Size. 10. Reaction Mechanism for Bond Formation, the
Bom-Haber Cycle. 11. Resonance. 12. Potential Energy and the Molecule.

13. The Nature of the Covalent Bond. 14. Properties of the Ionic and
Covalent Bonds.

14. Oxygen

1. Preparation of Oxygen. 2. Properties of Oxygen. 3. Uses of Oxygen.

Contents ix

Chapter Page

15. Hydrogen 200

l. Preparation of Hydrogen. 2. Properties of Hydrogen. 3. Isotopes, of
Hydrogen. 4. Order of Activity of the Metals.

16. Water and Hydrogen Peroxide 206

1. The Structure of Water. 2 Properties of Water. 3. Water as a Solvent.

4. Hydrates. 5. Physical States of Water 6. Water Vapor. 7. Purification
of Water. 8. Structure of Hydrogen Peroxide. 9. Preparation of Hydrogen
Peroxide. 10. Properties of Hydrogen Peroxide.

17. Solutions 219

1. Solution, Mixture, and Compound 2. Types of Solutions. 3. Nomencla-
ture. 4. Concentration Units 5. Kinetic-Molecular View of Solution. 6. Sat-
urated Solutions. 7. Supersaturated Solutions. 8. Solubility. 9. Law of Par-
tition

18. Properties of Solutions 235

I. Raoul t’s Law. 2. Deliquescence. 3. Boiling and Freezing Points of Solu-
tions. 4. Molecular Weights of Substances in Solution. 5. Osmotic Pressure.

6. Eutectic Mixtures. 7 Ideal Solutions. 8. Fractional Distillation. 9. Con-
stant Boiling Mixtures.

19. Solutions of Electrolytes 249

1. Electrical Conductance. 2. Conductance and Concentration. 3. Deviations
hom Raoult’s Law. 4. Chemical Reactions of Electrolytes 5. The Arrhenius
Theory. 6. Degree of Ionization. 7. Strong Electrolytes.

20. Ionic Equilibria— I: Acids and Bases 260

1. Ionization of Weak Electrohtes. 2. Strong Electrolytes. 3. Acids and
Bases. 4. Polyprotic Acids. 5. Ionization of Water. 6. The pH Scale. 7. In-
dicators. 8 Neutralization. 9. Acidimetry and Alkalimetry: Titration

21. Ionic Equilibria— II: Common Ion Effect & Hydrolysis 277

1 Common Ion Effect. 2. Hydrolysis. 3. Buffer Solutions.

22. Ionic Equilibria— III: Precipitates and Their Dissolution 284

1. Solubility Product. 2. Significance of the Solubility Product, 3. Calcula-
tion of the Value of the K sp . 4. Calculation of Solubility from the K S p.

5. Reactions Between Ions. 6. The Dissolution of Precipitates. 7. Complex
Ions. 8 Amphoteric Hydroxides.

23. Oxidation and Reduction 296

I. Ion-Electron Method of Balancing Equations. 2 Partial Ionic Equations.

3. Oxidation-Number Method of Balancing Redox Equations. 4. Equiva-
lent Weight of an Oxidizing or Reducing Agent. 5. Oxidation-Reduction
Problems.

Contents

Chapter 1 ACE

24. Electrochemistry I: The Voltaic Cell ...... 306

1. The Voltaic Cell. 2. Operation of the Zn-CL Voltaic Cell 3. Nomencla-
ture Applicable to a Voltaic Cell. 4. The Darnell Voltaic Coll 5. Measure-
ment of the Electromotive Force of a Voltaic Cell. 6 Single Electrode
Potentials. 7. Meaning of the Electromotive Series. 8 Calculation of the
Electromotive Force of a Voltaic Cell. 9. Electromotive Force and Equili-
brium Constant. 10 Electrode Potential and Concentration. 11. Concentration
Cells. 12. The Dry Cell

25. Electrochemistry II: The Electrolytic Cell 223

1. Electrolysis of Molten Sodium Chloride, NaCl. 2 Electrolysis of Aqueous
Sodium Chloride. 3. Electrorefming. 4 The Storage Cell 5 Electrical Units.

6. Quantitative Electrochemistry.

26. The Halogen Elements 334

1. General Properties of the Halogens. 2. Oxidation States of the Halogen
Elements. 3. Chemical Nomenclature. 4. Preparation of Flourine. 5. Proper-
ties of Flourine. 6. Preparation of Chlorine. 7. Properties of Chlorine. 8. Uses
of Chlorine 9. Preparation of Bromine. 10. Properties of Bromine 11. Uses
of Bromine. 12. Preparation of Iodine. 13. Properties of Iodine. 14. Uses of
Iodine. 15. Interhalogen Compounds.

27. The Hydrogen Halides 349

1. Methods of Preparation. 2. General Properties of the Hydrogen Halides.

3. Properties and Uses of Hydrogen Fluoride. 4 Properties and Uses of
Hydrogen Chloride. 5. Properties and Uses of Hydrogen Bromide and
Hydrogen Iodide. 6. Oxides of Chlorine. 7. Oxygen Acids of the Halogens
and Their Salts. 8. The Hypohalous Acids. 9. The Halous Acids. 10. The
Halic Acids 11. The Perhalic Acids.

28. The Elements of Group VIB: The Sulfur Family 358

1. General Properties of the Sulfur Family. 2. Occurrence of Sulfur.

3. Physical Properties. 4. Allotropes of Sulfur. 5. The Extraction of Sulfur.

6. Properties of Sulfur. 7. Uses of Sulfur. 8. Preparation of H^S. 9. Proper-
ties of K 2 S. 10. Polysulfides. 11 Preparation of S0 2 . 12. Properties of SO.,.

13. Uses of SO*. 14. Preparation of SO,. 15. Properties of SO, r 16. Prepara-
tion of H 2 SO r 17. Properties of H 2 S0 4 . 18. Uses of H 2 S0 4 . 19. Thiosulfates.

20. Selenium, Tellurium, and Polonium.

29. The Elements of Group VB: Nitrogen 375

1. General Properties of the Nitrogen Family. 2. Preparation of Nitrogen.

3. Properties of Nitrogen. 4. Uses of Nitrogen. 5. Preparation of NH a .

6. Properties of NH ;r 7. Liquid Ammonia. 8. Uses of NH 3 . 9. Other Com-
pounds of Nitrogen and Hydrogen. 10. Oxidation State +1. 11, Oxidation
State -j-2. 12. Oxidation State -f-3. 13. Oxidation State 4-4. 14. Oxidation
State H-5. 15. Preparation of HNO r 16. Properties of HNO,, 17, The Nitro-
gen Cycle.

30. The Elements of Group VB: Phosphorus, Arsenic, Antimony, and

Bismuth 390

1. Preparation of Phosphorus. 2. The Phosphorus Molecule. 3. Allotropes
of Phosphorus. 4. Properties of phosphorus. 5. Uses of Phosphorus. 6. Phos - 0
phine. 7. Halides of Phosphorus. 8. Oxides of Phosphorus. 9. Acids of
Phosphorus. 10. Orthophosphoric Acid and Phosphates. 11. Pyrophosphoric
Acid and Pyrophosphates. 12. Metaphosphoric Acid and Metaphosphates.

13. Phosphorus Acid and Phosphites. 14. Hypophosphorus Acid and Hypo-
phosphites. 15. Preparation of the Elements. 16. Properties and Uses, 17. Hy-
4rogen Compounds. 18. Halogen Compounds. 19. Oxides and Acids.

20. Sulfur Compounds.

Contents

Chapter Page

31. The Elements of Group IVB: Carbon 404

I. General Properties^ 2 Allotropes of Carbon. 3. Properties of Carbon

4. Carbon Dioxide. 5. Preparation of Carbon Dioxide. 6. Properties of
Carbon Dioxide 7. Carbonates and Hydrogen Carbonates. 8. Uses of Carbon
Dioxide. 9. Carbon Monoxide 10. Preparation of Carbon Monoxide.

II. Properties of Carbon Monoxide. 12. Other Compounds of Carbon.

32. Organic Chemistry I: The Hydrocarbons 418

1. The Hybrid Tetrahedral Bond. 2 Classes of Organic Compounds.

3, The Alkane Series of Hydrocarbons. 4. Isomerism of Carbon Compounds.

5. Nomenclature. 6. Reactions of the Alkane Hydrocarbons. 7. Structure of
the Alkane Hydrocarbons. 8. The Alkene Series of Hydrocarbons. 9. Struc-
ture of Ethylene 10. Reactions of the Alkene Hydrocarbons. 11. Diene
Hydrocarbons. 12. The Alkyne Series of Hydrocarbons. 13. Cyclic Aliphatic
Hydrocarbons. 14 Aromatic Compounds. 15. Reactions of the Aromatic
Hydrocarbons. 16. Natural Sources of Organic Compounds.

33. Organic Chemistry II: Derivatives of the Hydrocarbons 440

1. Halogen Derivatives. 2. Alcohols. 3. Phenols. 4. Ethers. 5. Aldehydes
and Ketones. 6 Acids. 7 Esters. 8. Amines and Amides. 9. Amino Acids.

10. Nitro Compounds. 11. Summary. 12. Optical Activity.

34. Organic Chemistry III: Polymers and Biochemistry 453

1. Natural and Synthetic Rubber. 2. Polymers and Polymerization. 3. Food
and Related Compounds. 4. Carbohydrates. 5. Photosynthesis. 6. Poly-
saccharides. 7. Paper 8. Fats. 9. Proteins. 10. Vitamins. 11. Hormones.
12. Enzymes.

35. The Elements of Group IVB: Silicon 467

1. Preparation of Silicon. 2. Properties and Uses of Silicon. 3. Silicon
Hydrides. 4. Silicon Halides. 5 Silicon Dioxide. 6. Silicic Acids. 7. Silicates.

8. Glass. 9 Varieties of Glass. 10. Clay and Ceramics. 11. Silicones.

36- Colloid Chemistry 477

1. Colloidal Dimensions. 2. Types of Colloidal Dispersions. 3. Preparation
of Colloidal Dispersions. 4. Filtration and Dialysis. 5. Optical Properties.

6. Colligative Properties. 7. Adsorption Properties 8. Electrical Properties.

9. Emulsions, Detergents, and Soaps. 10. The Realm of Colloids.

37. Metals 489

1. Nature of the Metallic State. 2. Occurrence of Metals in Nature. 3. Metal-
lurgy. 4. Special Metallurgical Processes. 5. Alloys. 6. Chemical Properties
of Metals. 7, Complex Ions. 8, Nomenclature of Complex Ions.

38. Elements of Group IA; The Alkali Metals 603

1. General Properties of the Alkali Metals. 2. Nomenclature. 3. Occurrence
of the Alkali Metals. 4. Preparation of the Alkali Metals. 5. Uses of the
Alkali Metals. 6. Reactions of the Alkali Metals. 7. Compounds of the
Alkali Metals.

30. The Elements of Group IIA: The Alkaline Earth Metals 514

1. General Properties of the Alkaline Earth Elements. 2. Occurence of the
Alkaline Earth Elements. 3. Nomenclature of the Alkaline Earth Elements.
4. Preparation and Uses of the Alkaline Earth Elements. 5. Reactions of the
Alkaline Earth Elements. 6. Compounds of the Alkaline Earth Elements.

7. Hardness of Water. 8. Softening of Hard Water. 9. Ion Exchange.

Contents

xii

Chapter Page

40. Transition Elements— I: General Properties 529

1. General Properties of Transition Elements 2. The Elements oi Croup
IIlX 3. The Elements of Group IVA. 4. The Elements of Cioup VA

41. Transition Elements— II; Iron, Cobalt, Nickel and the Platinum Metals.. 543

1. General Properties of the Iron Family. 2. Metallurgy oi Iron. 3. Wrought
Iron. 4 Steel 5 Heat Treatment oi Steel. 6. Propel ties oi lion. 7 Com-
pounds oi Iron. 8 Complex Ions. 9. Cobalt. 10. Nickel 11 The Platinum
Metals. 12 Ruthenium and Osmium 13. Rhodium and Iridium 14. Pal-
ladium and Platinum.

42. Transition Elements— III: The Elements of Groups VIA and VILA ....562

1 The Elements of Group VIA. 2. Chromium. 3. Oxidation States of Chio-
mium. 4. Molybdenum and Tungsten. 5. The Elements of Group VIZA
6 Manganese. 7. Oxidation States of Manganese 8 Technetium and
Rhenium.

43. Transition Elements— IV: The Elements of Group IB 574

1 General Properties oi the Copper Family. 2. Metalluigs oi Copper
3. Uses of Copper 4. Properties of Copper. 5 Compounds oi Copper.

6. Preparation of Silver. 7 Compounds of Silver 8. Photography. 9 Gold.

44. Transition Elements— V: The Elements of Group II B 585

1 General Propel ties of the Zinc Family. 2. Metallurgy of Zinc. 3. Proper-
ties of Zinc. 4. Compounds of Zinc 5 Metallurgy of Mercury 6. Properties
of Mercury. 7. Compounds of Mercun . 8. Physiological Action of Mercury.

45. The Elements of Group IIIB: The Aluminum Family 593

1. General Properties of the Aluminum Family 2. Preparation oi Aluminum
3. Use of Aluminum 4 Properties of Aluminum 5. Compounds oi Mum-
mum.

46. The Elements of Group IVB: Germanium, Tin, and Lead 604

1. Transistors 2 Allotropie Forms of Tin 3. Properties of Tin 4 Properties
of Lead.

47. The Elements of Group O: The Noble Gases 612

1. General Properties of the Noble Gases. 2. Discovery of the Noble Gases.

3. Properties of the Noble Gases. 4. Cryogenics 5 Chemical Properties
of the Noble Gases.

48. Nuclear Chemistry I 620

I The Group Displacement Rule. 2. Kinetics of Radioactive Transformation.

3. Detection of Radioactivity. 4. Properties of the Nucleus 5. Artificial
Nuclear Reactions. 6. The Acceleration of Charged Particles.

49. Nuclear Chemistry II 038

1. Nuclear Fission. 2. The Cham Reaction. 3. Separation of Isotopes.

4. The Nuclear Reactor. 5. The Transuranium Elements. 6. Nuclear Fusion.

7. Applications of Nuclear Energy.

50. Analysis .?.649

1. Instrumental Analysis. 2. Chemical Analysis 3. Schemes of Analysis.

4. Analytical Procedure. 5. Anion Analysis.

Appendix 659

Index 677

Introduction:
The Philosophy of Science

Broadly defined, science is classified knowledge, an accumulated body of
facts based upon observation and then collated in an attempt to achieve
regularity among them. These facts are concerned with natural phenomena
and are classified into specific branches of knowledge so that general laws
can be formulated, experimentally tested, and verified, thereby increasing
man's understanding of the world about him. Science is the search for order.

The term “science,” as used here, is synonymous with the older terms,
“natural philosophy” and “natural science,” and embraces many fields, among
them astronomy, biology, chemistry, geology, and' physics. Chemistry is one
of the oldest of sciences, its origin being contemporary with that of astronomy.
To the ancient Babylonians and Greeks, astronomy included the problem
of the structure of matter. The Greek word “chemeia” fiist appears in
fourth century writing and originally was used to designate the art of metal
working. Later the Arabs added the prefix “al” and “alchemy” came to signify
the arts of chemistry in general. Modern chemistry, however, dates from the
late eighteenth century when the performance of quantitative experiments
brought maturity to the science.

As early as 450 b.c. Empedocles tried to explain the existence of all
varieties of matter in terms of four primary elements— earth, water, air, and
fire. Combinations of these in varying proportions were believed to produce
all the known species of matter. Leucippus and Democritus were the first
to conceive the atomic nature of matter. Antedating the very heart of the
science of chemistry some 2000 years before the dawn of the Atomic Age,
Democritus proposed, about 400 b.c., that matter consisted of indestructible
moving atoms, qualitatively alike but differing in mass, size, and shape.
Later Plato and Aristotle discarded the atomic concept and their great in-
fluence upon subsequent thought hindered scientific advance for centuries.

The early Greeks observed various evident facts and framed them into
theories such as the geometry of Euclid. Greek philosophy, however, was
based upon unveiling the mysteries of nature by discussion and logic alone.
Experimental work to determine factual data was considered unnecessary or

Introduction. The> Philosophy of Scwru'e

inferior to the obtaining of general propositions by logic. This belief in
logic alone held sway for 2000 years— till the Renaissance and the beginning
of modem history.

But man does not come into this world bringing a priori knowledge with
him. Knowledge comes from facts and facts come from experimental observa-
tion alone. Facts are discovered, not invented or revealed, and chemistry is
primarily an experimental science to wrest from nature data concerning the
behavior of matter in chemical reactions. Facts, by themsehes, undigested
and unclassified, do not lead to fundamental knowledge. Facts have value
and meaning only when they are correlated, systematized, and “explained/’
To accomplish this, two kinds of logic, inductive and deductive, are employed.
Induction is the process of correlating a large number of separate but related
facts so that a general rule can be inferred from them, whereas deduction
is the reverse process whereby a general rule, once established, can be used
to predict new facts. During the early growth of a science, scientific knowl-
edge first obtained is generally qualitative and descriptive but as the science
develops and becomes more mature, general laws are discovered. Such laws
are best described in the precise and unambiguous language of mathematics,

1. The Scientific Method. Distinction must be made between “science"
and the scientific method.” Science is a field of study, just as are history
and economics. Contrary to widely held public opinion, science is not a
universal panacea and does not include all phases of human endeavor. There
are perhaps some spheres in which science is an intruder. Science is built
upon certain fundamental postulates and axioms which preclude its universal
application.

The scientific method, however, is a technique of investigation and is
applicable in every field of physical endeavor. The proper use of the
scientific method involves several distinct operations. The first is the experi-
mental observation and collection of pertinent facts or data; this requires
controlled measurements in the laboratory. The second step is the collation
and grouping of the data followed by a study of their possible mutual rela-
tionships. Thereafter it may be possible for the scientist to make a general
statement which covers the relationships observed among the experimental
a a and perhaps to formulate a quantitative law of nature. Such scientific
generalizations or laws are by no means to be considered inviolate. Further
research may bring to light new facts at variance with a scientific law. The
law then either is modified to include the newer facts or some entirely new
hypothesis may be proposed and the old law discarded. Even so the accuracy
o a scientific law is not absolute. Its conclusions are limited by any
assumptions made in deriving the law, a point which is too often forgotten

SalWwTf* U , S l° f . a laW ' , N ° kw is valid outsii3e the domain orig-
noi J f for * by lts , initial postulates. Furthermore, from the view-

point oh modern physics, there is reason to believe that absolute accuracy
of measurement is an unrealizaWe state of affairs. * X

™„? eSpite th i C ma ™ l0US achieve ments of science, particularly in the present

istrv Man” kn °w ed f e iS StiU fragmentary in many fields including P chem-

ToToft !l C ! aWS n ° W in , v ° gUe are rec °S ni2ed as being incomplete,
ten the student acquires the impression that so much is blown and

Introduction: The Philosophy of Science

so little remains to be discovered. There is more that is new and waiting
to be discovered than is now found between the covers of textbooks. It is
to this end that scientific research is directed— to reduce the vast area of
the unknown!

2. Scientific Explanation. From the ultimate point of view, scientific
explanations describe only how or the manner in which certain phenomena
operate; they do not tell us why events take place in one particular way in
preference to another. In the last analysis, too, scientists did not create the
universe— they merely describe it and its operation.

In formulating a scientific explanation, the scientist uses his imagination
and tries to construct some plausible state of affairs, a so-called model , that
would account for an observed phenomenon. Imagination is not at all incom-
patible with the scientific method; in fact it is essential to it. Many great
discoveries have been due to accidents happening to the prepared mind.
The great scientist is one equipped with a creative imagination— one who,
through deep insight, can hypothesize systems or laws that can interpret
observed phenomena. Indeed many of our most familiar concepts are truly
philosophical constructs— energy, temperature, and force, among others— in-
ventions of the human mind without real existence in the sense of a material
object, but fruitful in that their use brings quantitative order to the vast
quantity of data which nature yields.

Yet a scientist does not accept an invented state of affairs as valid if it
is merely in accord with the immediate facts that he is trying to explain. In
addition the imagined system must also serve as a basis for the prediction
of new facts or phenomena that can be observed or tested experimentally.
When these predictions have been verified by further experiment then the
assumptions are held to be valid. The assumed state of affairs is called an
hypothesis . To many students it appears that a scientific hypothesis is
fabricated from thin air and thereafter it seems to be a remarkable circum-
stance that experimental data adjust themselves to be in agreement with
the hypothesis. This is putting the cart before the horse and nothing could
be further from the truth. A scientific hypothesis originates from experimental
data and to be operationally useful must be capable of verification by ex-
perimental techniques. The hypothesis that life exists on other systems
similar to our solar system cannot yet be experimentally verified in view
of our present state of space travel.

An hypothesis may be of such generality and profound significance that,
after repeatedly proving its experimental validity over many years, it may
be dignified through common usage by calling it a law . A scientific law is
a statement, either qualitative or mathematical, of a cause-and-effect rela-
tionship. It summarizes a multitude of separate observations and expresses
what is believed to be the invariable consequence of a given set of condi-
tions and operations. The statement that “iron will rust if exposed to moist
air” rs a scientific law, as are Newton's mathematical laws of motion. To
explain laws, theories 1 are proposed. A theory proposes a model, or mental
construct, from which the laws of nature can be derived. A theory not only

x The term “theory” is used here in a strictly technical sense, distinct from any
meanings associated with it in ordinary language.

Introduction The Philosophy of hn onto

permits the deduction of the laws of nature lor which it was devised but
also enables the prediction of new' laws. Both terms, theory and law, have
equal validity and deserve equal respect. 1 ' A theory is concerned with objects
which are not directly tangible or material, whereas a law can be verified
directly with macroscopic, material objects. Thus Dalton’s Atomic Theory
deals with the mental construct of atoms which individually are not ap-
parent, while Maxwell’s Electromagnetic Theory of Radiation deals with
hypothetical waves and vectors. Newton’s Laws of Motion can be demon-
strated directly with billiard balls or bullets, objects readily at hand.

John Dalton conducted numerous experiments on the behavior of gases.
In endeavoring to interpret his experimental results he speculated about the
ultimate constitution of matter. Influenced by the writings oi Democritus
and of Isaac Newton, Dalton proposed the theory that all matter is composed
of discrete, indivisible particles called atoms and that chemical compounds
are formed by the union of atoms of different elements in simple numerical
proportions. Dalton’s Atomic Theory offered a theoretical basis for the
stoichiometric laws of chemical combination (Chapter 6). The theory did
not attempt to explain why chemical reactions* took place. Moreover, we now
believe that Dalton was wrong in his idea that the atom is indivisible and
is the ultimate unit of matter. Yet the formulation of the theory brought
order to the chemical thought of the time, and recent discoveries have caused
the theory to be extended rather than discarded. Similarly, New ton’s Laws
of Motion, so dominant in the study of mechanics, have been shown to be
a limiting case of the more general Theory of Relativity.

At this point some words should be added concerning the philosophy of
science insofar as it deals with the nature of truth and reality. At first
the idea that the scientist is not particularly concerned with absolute truth
as the philosopher sees it may seem peculiar. A scientist is faced with
factual experimental data. His objective is to explain these data. In so doing
he may propound an hypothesis. If the hypothesis proves successful in the
manner previously described, it is called “true.” New data obtained perhaps
from newly discovered experimental techniques, however, may invalidate
the hypothesis. A new proposal, if successful, is then purported to be the
“truth.” The older hypothesis is now considered either limited in scope
or false and incorrect. The history of science is replete with cases wherein
one theory supplanted another yet each in its turn was considered the
“truth.” Science has been called a collection of pointer readings. It deals
with appearances, of which the events are linked by theoretical concepts
attempting to bring order to the whole. Perhaps the succession of theories is a
closer and closer approximation to the “truth.” Logic would dictate that there
is, for a given phenomenon, one underlying reality or truth. Perhaps in time
we may learn the fundamental truth but science as now constituted is prag-
matic. What theory works best, and if many work equally well then the
simplest, is considered the “truth.” Who can say what relationship the con-
ceptual framework of the scientist bears to the reality behind it?

Furthermore, to ultimate reality modern physics has brought an inherent
haziness, p articularly in atomic and nuclear problems so that in the last

2 TJ e view that a hypothesis passes through successive stages nf nomenclature —
hypothesis, theory, and law — as it gains scientific stature, is erroneous.

Introduction The Philosophy of Science

analysis there may be no single definite answer. In such cases the best
answer that we may be able to give will be in terms of probability only.
For the macroscopic problems of everyday life, however, this divergence
from a single, unique truth-value is of such minute magnitude that it need
not concern us.

3. Matter and Mass. The chemist deals with matter and its transforma-
tions from one species to another. Though we are all familiar with “matter”
a fundamental definition of this term is difficult. In physical science there
are times when there is no alternative but to define arbitrarily certain
basic terms which serve as a foundation upon which the edifice of science
can be reared. Such a term is “matter.” If an object has mass and occupies
space it is said to contain “matter.” Any material object has matter; a total
vacuum does not.

“Mass” is another fundamental quantity which defies basic definition.
Mass is a measure of the inertia of a body. Its concept stems from Newton’s
Laws of Motion. The first law that a body at rest remains at rest and a
body in motion continues to move in a straight line, when acted upon by
no unbalanced force, implies the property of inertia to all objects; the second
law states that an unbalanced force accelerates a body, and that the con-
stant of proportionality between the force and the acceleration produced
is the mass of the body. The larger the mass of an object the greater is its
inertia. The distinction between “mass” and “weight” should nevertheless
be clearly understood. “Mass” refers to the quantity of matter in an object
while the “weight” of an object is measured by the gravitational attraction
of the earth for it. In outer space beyond the gravitational pull of the earth
and other astral bodies, an object would be weightless yet its mass would
be identical with that on the surface of the earth. Mass is measured in units
of grams or pounds and should not be confused with bulk, which refers
to the space occupied by a body or its physical dimensions.

4. Properties and Conditions. The scientist distinguishes between species
of matter by their properties. Just as each person has certain inherent char-
acteristic properties such as the color of his eyes or hair which distinguish
him as an individual, chemical species can be identified one from th& other
by certain properties which are specific to a given species. Among these
properties are color, density, melting point, boiling point, and solubility.
These specific properties are sometimes called intensive properties. It can
be said safely that there are not two different substances in the universe
which have completely identical properties. Indeed if such were the case
they could not be distinguished from each other and would be considered
to be the same substance.

Chemical analysis is essentially a detective story in which we try to match
the properties of an unidentified sample of matter with properties previously
recorded for known substances. In the table following are listed some com-
mon substances and some of their properties. With the aid of such a table,
an extraterrestrial visitor would have little trouble in distinguishing among
these species. Certain properties, such as shape, volume, or v/eight, possessed
by a particular sample of material are not characteristic of the species but only
of that particular sample. Because these properties depend upon the quantity

Introduction The Philosophy of Science

Table I -A

Properties of Common Substances

Property

Salt

Silver

Copper

Wafer

Physical State at 20 °C

solid

liquid

Color

colorless

yellow

white

red

eoloiless

Taste

‘sal tv*

none

Density g/cm 3

2.17

19.32

10. 19

8 961

1.00

Melting Point, °C

801

1063

961

1083

0 00

Boiling Point, °C

1413

2966

2212

2595

100

Solubility in Water at

20 °C. gram /I it ei

360

insoluble

of material present and are not primarily dependent upon the nature ot
the material they are sometimes called extensive properties. Two samples
of copper, having different dimensions, will both he recognized as copper
because their specific properties are identical. It may be noted that the
ratio of two extensive properties, mass to \olume, is itself a specific property,
density.

A sample of matter can also have other properties* which are impressed
upon it by an external source. These are more properly called the conditions
under which the sample exists. Two important conditions are temperature
and pressure.

The properties listed in Table i-A are some of a general group called
physical properties. The breaking in two of a piece of chalk, the conversion
of solid ice to liquid water and then to gaseous steam all result in changing
solely the physical properties of the substance concerned without altering
its chemical nature. Ice, liquid water, and steam are all the same chemical
species. Changes wherein the chemical nature of a material is unaltered
are termed physical changes.

Matter also has chemical properties. These involve the transformation from
one species of matter to another, a process known as chemical change. Thus,
if carbon is burned in oxygen, a chemical change results; and a new chemical
entity, carbon dioxide, is produced. Because new species of matter result
from a chemical reaction the properties of the initial substances, or reactants,
disappear and in their stead appear the properties of the products of the
chemical reaction. All the properties of carbon dioxide, both physical and
chemical, are different from those of both carbon and oxygen

This then is the primary business of chemistry— to study chemical trans-
formations and all their accompanying details, the conditions which enable
us to control the reaction, and through theory, to gain greater insight con-
cerning tpe mechanism and reasons for chemical reaction!

5. Elements, Compounds, an<i Mixtures. A material consisting of bur one
species of matter is a pure substance. Pure substances may be elements
or compounds. A pure substance that cannot be decomposed bv ordinary
chemical means mto simpler substances, nor so formed by the combination
of Other substances, is defined to be an element. The combination of elements

Introduction : The Philosophy of Science

results in new substances called compounds which, in turn, can be decom-
posed into their constituent elements.

These definitions obviously are empirical. Later an element will be
defined more precisely as a substance composed solely of one atomic
species. It is within the realm of possibility that a substance now considered
an element will be decomposed into simpler substances. In that case the
simpler substances would be defined as elements During the early history
of chemistry such discoveries were frequent as new techniques of chemical
decomposition were developed but that such a discovery should be made
in the present light of chemical knowledge is extremely remote.

Examples of elements are hydrogen, oxygen, carbon, silver, gold, and
chlorine; of compounds, water, salt, sugar, and sulfuric acid. Hydrogen and
oxygen can be combined to form water but they cannot be decomposed into
simpler chemical substances. Conversely the compound water can be broken
down into the elements from which it was formed. When uncombined, ele-
ments are said to be in the elemental or free state. One hundred and three
elements are known at the present time. The number of compounds is far
greater; over a half a million compounds are known. A list of the known
elements is on the inside front cover of this book.

Most of the commonly seen objects about us are not pure substances but
are mixtures in that they contain two or more substances intermingled. In
the rock granite three distinct substances can be distinguished: glasslike
fragments of quartz, dark flakes of mica, and pink portions of feldspar.
Each component of a mixture retains its original specific properties so that
mixtures, unlike compounds, can be separated into their components no
matter how finely subdivided the mixture may be. In a mixture of sand
and salt, the salt is still soluble in water while the sand is not. Separation
can be effected merely by adding water, which dissolves the salt, and then
pouring off the liquid. The waiter can subsequently be removed from the
salt by distillation. From a mixture of iron and copper, the iron can be
separated because it alone, of the two, is magnetic and a magnet will pick
out the iron and leave the copper. When a mixture contains only a very
small quantity of a substance or substances, it is customary to refer to it
as an impure substance and the small quantities are considered to be
impurities. When a chemist uses terms such as iron, water, or salt, he
implies the pure substance and not just any sample of that material.

6. Abundance of the Elements. The elements are not equally distributed
in nature. Table 1-B lists the relative abundance by weight of the elements
in the earth’s crust. Many elements are familiar to us, for example, copper,
silver, lead, and carbon, but such familiarity has little relation to the
element’s abundance. Indeed, the second most abundant element, silicon, is
quite unfamiliar in the elemental state. The relative rarity of most of the
elements is emphasized by the fact that the ten elements in Table 1-B con-
stitute over 98% of the earth’s crustal weight. The remaining elements may
be considered almost as impurities, though exceedingly important ones. Thus
carbon, the element which is a constituent of most chemical compounds, a
major source of the earth’s fuel, and the heart of all living matter, does
not appear among these most abundant elements. Oxygen, which occurs in

Introduction, The Philosophy of Si unite

the atmosphere, in water, and in many rocks, alone comprises about halt
the weight of the earth, while oxygen and silicon make up more than
three-quarters of its weight.

Table 1-B

Composition of the Earth's Crust by Weight

Element

Per cent in solid
portion of
earth* s crust

Per cent m ocean

Per rent in shell
including the
atmosphere

Oxygen

47.33

85.78

50.02

Silicon

27.74

25.80

Aluminum

7.85

7.30

Iron

4.50

4.18

Calcium

3.47

0.05

3.22

Sodium

2.46

1.14

2.36

Potassium

2.46

0.04

2.28

Magnesium

2.24

0.14

2.08

Hydrogen

0.22

10.67

0.95

Titanium

0.46 i

0.43

Other Elements

1.27

2.18

1.38

There is evidence to indicate that the composition of the earth varies
from the crust to the center. For example the average density of the
earth's crust is about 2.8 grams per cubic centimeter (g/cm :i ) whereas the
average density of the entire planet has been estimated as approximately
5.5 g/cm 3 . Further seismic and magnetic data lead to the belief that the
core of the earth, 4000 miles in diameter, consists mainly of iron (90.8%)
with a small fraction of nickel (8.6%) and has a density of about 9.6 g/cm*\
Based on these estimates calculations place iron as the most abundant
element of the entire planet, about 35.4%, with oxygen second, about 27.8% .

7. Units of Measurement. We have said repeatedly that chemistry is a
quantitative science, that it measures phenomena and obtains numbers re-
lating to their magnitudes. This makes necessary the establishment of sets
of units in terms of which data can be obtained. Experimental variables,
differing in kind, necessitate different types of units. Measurements of
length, be it the dimension of an object or the wavelength of light, require
different units from those required in the measurement of temperature.

Based upon the three fundamental quantities of physical science, dimension,
mass , and time, two major systems of units developed. They are the
Metric System and the English System. The Metric System is an outgrowth
of the French Revolution and was adopted in that country as the sys-
tem of weights and measures in 1793. Though also legally adopted by
the United States in 1866, the Nfetric System has not come into general
use in this country; however, it is used almost exclusively in scientific work.
The Metric System is a decimal system. A basic unit is defined arhitrarilv
prefixes attached thereto indicate tenfold multiples and submultiples
o* this unit. These prefixes and their values are listed in Appendix II.

Introduction • The Philosophy of Science

The unit of length in the metric system is the meter, abbreviated simply
as m. The centimeter (cm) and millimeter (mm) are frequently used
subdivisions of the meter. Abbreviations of units generally are not followed
by periods, thus mm and not mm. is the way the abbreviation is written.
Originally the meter was thought to be one ten millionth of the length of the
earth’s meridian passing through Paris, but this was found to be in error,
and the meter was established as an arbitrary length between two lines
etched on a platinum-iridium bar when the bar is at 0°C. This bar was
kept in the vaults of the International Bureau of Weights and Measures at
Sevres, France, and duplicate standards were made available to other
countries. On October 14, 1960, the Eleventh General Conference of Weights
and Measures adopted a new international standard of length based on the
orange-red line in the spectrum of the krypton isotope 86. The new definition
of the meter as 1,650,763.73 wavelengths of this orange-red line replaces
the platinum-iridium bar as standard. The new definition relates the meter
to a constant of nature which is independently reproducible, universally
available, and indestructible. Units of area are dimensionally the square
of linear units, as square centimeters (cm 2 ), and units of volume are the
cube of linear units, as cubic centimeters (cm 3 ), sometimes also abbreviated
as cc.

The primary unit of mass is the kilogram (kg), of which the prototype
is a platinum-iridium cylinder also kept at Sevres. For the range of masses
ordinarily used in the laboratory, the chemist finds it more convenient to
use subdivisions, the gram (g) and the milligram (mg). The standard of
volumetric capacity is the liter ( 1 ) , which is the volume occupied by one
kilogram of water at 4°C, the temperature at which water has its maximum
density, and under normal atmospheric pressure. This statement fixes the
maximum density of water as one gram per milliliter at 4°C. For measure-
ment of fluids, the liter and milliliter (ml) are commonly used units in
chemistry. One milliliter is not exactly equal to one cubic centimeter, one
milliliter being equal to 1.00027 cm 3 . The milliliter is a unit based on mass
whereas the cubic centimeter is the volume of a cube one centimeter on
edge. Originally it was intended that the two should be equal and only
in extremely accurate work does the difference between them assume im-
portance.

The unit of time is the second (sec). Formerly defined as a specific
part of a year, the second was redefined in 1964 also on the basis of a wave-
length of light. The frequency of a spectral line emitted by the cesium-133
isotope' is specified as 9,192,631,770 cycles per second.

Because of its more commonly used units the Metric System is also
called the centimeter.-gram-second (cgs) system. In the English System
the basic units of length, mass, and time are the foot, pound, and second,
respectively, and hence this system is also designated the “fps” system.
Appendix II shows the relation between units of the two systems.

8. Temperature and -Heat. Temperature is a measure of the intensity
or concentration of heat energy in a material object. Heat is a form of
energy whereas temperature is an index of the direction in which heat
will flow. Spontaneously heat will always flow from an object at high

Intwiim turn The Phth'sophif >>t u m r

temperature to one of lowei tempera tuie. A common misconception is that
the temperature is a measure of the total quantity of heat within a body.
This is incorrect. The Atlantic Ocean, being so \ast a quautiiv of matter,
has a large amount of heat energy within it yet its temperature, or intensity
of heat, is quite low because its heat is distiibuted tluoughout so laige a
mass.

Tempeiature is measured by a thermometer of which the most common
is the mercury-in-glass thermometer. The temperature scale most commonly
used in scientific work is that invented by the Swedish astronomer, Anders
Celsius, in 1742. On this scale the freezing point of water is designated arbi-
trarily as zero and the boiling point of water is taken to be 1(X), hot li
measured at normal atmospheric pressure. The interval between these points
is divided into 100 equal parts known as degrees Celsius, abbicviated C.
This designation has not, however, come into geneial use by scientists in
America who refer to tempei attires on this scale as degteos Centjgiade,
also °C. In practice it is difficult to establish a temperature scale defined
by more than one fixed point. In 1960 the Eleventh Geneial Confeience on
Weights and Measures adopted the “International Practical Scale*’ wherein
the triple point of water (page 213) is defined as 0.01 degrees. The interval
between the triple point and the boiling point of water is thus 99.99 degiees
and hence is not strictly centigrade any longei.

Another scale of thermometry commonly us<'d in English speaking
countries is the Fahrenheit scale, proposed in 1721 b\ the German physicist.
Daniel Gabriel Fahrenheit. On this scale the freezing point of water is
32°F and the boiling point is 212°F. The interval of ISO degrees on the
Fahrenheit scale between the freezing and boiling points of water cor-
responds to an interval of 100 degrees centigrade: thus each centigrade
degree equals 9/5 Fahrenheit degree. To convert theimometrr readings
from one scale to another, these relationships hold.

°F = ( C C) -I- 32 and °C = — ( F - 321
o 9

All thermometers are based upon the variation of some property of
matter with a change in temperature. The mercury-in-glass thermometer
depends upon a change in dimension, that is, the change in volume of
liquid mercury with temperature. Expansion or contraction takes place
within a capillary tube which accentuates the observation of the volume
change in one linear dimension. The limits of the mercury thermometer an*
the freezing point and the boiling point of mercury, approximate!) from
-38°C to 350°C, For temperatures beyond these limits other fluids or instru-
ments are used. At 1ow t temperatures a fluid such as alcohol or hydrogen
gas can be used. Temperatures from -200 r C to over 100Q C C can be measured
by the resistance thermometer, '"which depends upon the variation with
temperature of the electrical resistance of a metal, usually platinum, and
the thermocouple which depends upon the change in electrical potential
(voltage) with temperature produced at the junction of two dissimilar
metals such as iron and copper. The optical pyrometer, based upon the

Introduction: The Philosophy oj Science

change in color and the intensity of radiation emitted by an incandescent
body, can determine temperatures in excess of 3500° C.

9. Pressure. Pressure is defined as force per unit area. A force of 10
pounds operating over an area of 2 squaie inches produces a pressure of
5 pounds per square inch. In that it is not a measure of a total effect
but rather of an intensity, pressure is akin to temperature.

To measure pressure a barometer is used. In its simplest form the
barometer consists of a glass tube about one meter in length and closed
at one end. It is filled with mercury and then the open end, temporarily
held closed with a thumb, is inverted into a rcservoii of mercury, after which
the thumb is removed. The level of the mercury in the tube falls somewhat,
leaving in the upper end a vacuum known as Torricellis vacuum. (Figure
1.1) This instrument was invented in 1643 by Evangelista Torricelli, an
Italian mathematician and secretary to Galileo. Since the column of mercury
is held up by the pressure of the air acting upon the surface of the mercury
in the reservoir outside the tube, this pressure can be measured in teims
of the height of the mercury column. Atmospheric pressure is not constant
but may vary from day to day. Standard atmospheric pressure, or a pressure
of one atmosphere, has been defined equal to a height of 76.0 cm of mercury
at 0°C and sea level. One atmosphere corresponds to a pressure of
1.013 X 10' J dyne/'enr or 14.7 lb/in 2 . Atmospheric pressure, which depends
upon the weight of air over a unit area of the earth’s surface, varies con-
siderably with height above sea level. At an altitude of 10,000 feet, atmospheric
pressure is about 400 mm whereas at a height of 10 miles the pressure
is about 40 mm.

Figure 1 . 1 . A Mercury Barometer.

The column of mercur> in the vertical tube is
sustained by the atmospheric pressure exerted
upon the surface of the mercury in the cup, C.
The height, h, of the column of mercury is a
measure of the atmospheric pressure.

Introduction The Philosophy of Scu tv c

QUESTIONS

1. Define and differentiate among each of the following, science, the scientific
method, and scientific explanation.

2. Distinguish among hypothesis, theory, and law of nature. Which is most
likely to undergo modification? Explain.

3. List examples of theories and laws with which \ou art* familial. What
experimental facts does each try to explain?

4. Devise an experiment by which it is desired to determine how atmospheric
pressure varies with height above sea lex el. What data and controls would
be necessary?

5. The volume of a substance varies with temperature as shown by the data
below:

Volume, liters

7.5

10.0

12.5

15.0

I 20.0

Tempeiature, °C

100

200

300

400

| 600

Derive an equation which represents these data. Use your equation to calculate
the volume at 500 °C.

6. Distinguish between chemical and physical propeities. List chemical and
physical properties of (a) iron (b) ox \ gen (c) wood. (Note you can look up
the properties of the first two m this book or a chemical handbook.)

7. What properties could be used to distinguish between (u N water ami gasoline
(b) salt and sugar (c) aluminum and mercury?

8. Distinguish among element, compound, and mixture, giving examples of each.

9. What experimental evidence would indicate whether the following substances
were pure substances or mixtures- (a) paper (b) ink <c) solder (d) salt (e) 14
karat gold (f) ice?

10. Which of the following are physical changes and which are chemical changes:
(a) corrosion of iron (b) compression of air (c) burning of gasoline (d) forma-
tion of snow (e) addition of liquid silver to liquid gold (f) dissolving sugar
m water (g) explosion of dynamite.

11 .

12 .

13.

14 .

What are the international standards of length, mass, and time p
A Hquid has a density of 0.80 g/ml. (a) What will be the volume of 200 grams
of this liquid? (b) What weight of this liquid could be held by a container
, meters lon S by 50 centimeters wide by 12 centimeters high? (e) What
volume of water will have the same weight as 500 milliliters of this liquid?

(a) Convert 20 degrees Centigrade to degrees Fahrenheit; 40*C to °F

(b) Convert 96 6«F to °C; 212°F to °C (c) At what temperature will
Centigrade and Fahrenheit thermometers read the same value?

The pressure, p of a height, h, of a uniform fluid ol density, d, is given
y p dg know that one atmosphere pressure corresponds to 1.013 v 10 n
dyne/cm^ and also to 14.7 lb/in*. The density of mercury is 13.6 g/ml and
e dyne equals one g cm/sec 2 ; the acceleration of gravity, g, is 981 cm/sec 2 .

The States of Matter:

Cases

There are three states of matter: solid, liquid, ana gas. Any substance
may exist in any of the three states depending upon the surrounding con-
ditions of temperature and pressure. Water thus may have the form of
solid ice, liquid water, or gaseous steam. The three states of matter can
be defined empirically as follows. A solid has a definite shape and a definite
volume, both of which are maintained quite independently of the shape
of the container in which the solid may be placed. A liquid has a definite
volume but its shape will conform to that of its container. A gas has neither
a definite volume nor a definite shape inherently but will completely fill and
take the shape of its containing vessel. The basis of these empirical proper-
ties will become evident when the structure of these states is considered
from a fundamental molecular and kinetic viewpoint. In this chapter we
shall discuss the properties and structure of the gaseous state, the simplest
of the three, and we shall defer to the next chapter the study of the solid
and liquid states.

1. The Gaseous State. No matter what the volume in which a given
quantity of gas is placed, it spreads uniformly throughout the entire space.
In other words, a gas is infinitely expandable. Another property characteristic
of the gaseous state is a high degree of compressibility. When a given
quantity of a gas is subjected to an increase in pressure, its volume is
diminished considerably more than is an equal volume of solid or liquid
subjected to the same pressure. Gases have relatively low densities as
compared with solids and liquids. Under ordinary conditions one cubic
centimeter of water weighs approximately one gram whereas one cubic
centimeter of gaseous water weighs about 0.0008 gram under the same
conditions.

These properties of gases create difficulties in the handling of gases
experimentally. Unlike a solid, a gas cannot be held directly in the hand.
Gases, and liquids also, must be kept in a container, but unlike liquids,
gases must be held in a closed container for obviously they would escape
if the container were open.

J ill' nf \Littri (last

As we shall see, weights of substances are of prune nnpu: lance >n chemical
calculations. Weights can be obtained by weighing diiectb upon a chemical
balance or scale. Because the densities of gases aie so low, the mntainci loi
even a small weight of gas is relatively laige, peihaps o\en laigei than
the balance itself. Further the container generally weighs far umic than the
gas it contains so that a small erioi in weighing the eontainei may be gi eater
than the weight of the gas itself. One solution to tins problem is in the fact
that the weight of a gas is proportional to its \olume, and \o!usnes oi gases
are fairly easy to measure. If the density of a gas is known, a measurement
of the volume of a gas enables its weight to be deteunined In calculation,
since the weight equals the density times the volume. \lso, inasmuch as
gaseous volumes are large they can be measured with accuracy. In contrast,
however, with the solid and liquid states the volumes of gase> vary ap-
preciably with temperature and pressure. To specify solely the volume of
a gas conveys little information unless simultaneously the temperature and
pressure are stated.

A convenient technique for the laboratory prepaiation and collection
of gases is shown in Figure 2.1. The gas is prepared In chemical reaction
in a generator and then collected by downward displacement of a liquid
in which the gas is not soluble and with which it does not react The col-
lecting vessel, called a eudiometer , is usually transpaient and is calibrated
to indicate the volume at various ley r els The eudiometer is fust f illed com-
pletely with a liquid and is then inverted in a jar containing the same
liquid. The gas is collected by allowing it to bubble into the open end of
the eudiometer, thus displacing the confining liquid dowmvuui into the jar.
Before reading the volume of the gas collected, the eudiometer is laised or
lowered so that the level of the liquid inside is the same as that outside of

Ficrure 2 . 1 . Laboratory Preparation of a Gas.

The State* of Matter: Gases

the jar. Figure 2.2. Under these conditions the pressure of the gas collected
within the eudiometer is equal to the pressure exerted upon the surface
of the liquid in the jar, namely, atmospheric pressure. The latter can then
be read separately on a barometer. The temperature of the gas is obtained
by allowing sufficient time to elapse so that the gas comes to the same
temperature as the room and then reading a thermometer held in the air
near the eudiometer. In this way the volume of a gas at a specific tempera-
ture and pressure can be obtained.

The pressure under which the gas, G, is confined in the eudiometer can be varied
by raising or lowering the eudiometer m the liquid, L, in the jar. In A, the
pressure of the gas is less than atmospheric pressure. Atmospheric pressure equals
the sum of the gas pressure plus the pies sure due to the height of the column of
liquid, hj. so that the pressure of the gas equals atmospheric pressure minus the
pressure corresponding to h t . In B, the pressure of the gas is greater than atmospheric
pressure by a pressure corresponding to the pressure of the columns of liquid, h r
In C, the pressure of the gas equals atinosphtne pressure.

Figure 2.2. The Pressure of a Gas.

2. Behavior of gases. How does the volume of a gas vary with pressure
and temperature? Inasmuch as both \ ariables affect the volume we must resort
to a common mathematical technique. Where a phenomenon depends upon
a number of independent variables, the problem is dissected by keeping all
but one of the variables constant, mentally or experimentally as required,
and then observing how the phenomenon under consideration varies with
the one remaining variable. This is then done in turn for each of the
other variables. Unlike volumes, the weight of a given sample of gas is
unchanged by variations in temperature and pressure. A specific volume

The States of Matter Gases

of a gas such as one liter, at different temperatures and pressures, will,
however, have different weights.

3. Change in Volume of a Gas with Pressure: Boyle s Law, Let us
consider' first how the volume of a gas varies with prossuie it the tempeiature
is kept constant. It is common knowledge that when we push down the
piston upon the gas (air) in a bicycle pump, thereby mo easing the pres-
sure, the gas contracts or diminishes in volume. In 3662, Robert Boyle,
a British physicist, expressed this behavior of all gases as follows. Al con-
stant temperature the volume occupied by the same sample of a gas is
inversely proportional to the pressure. This statement is known as Boyles
Law. If the pressure on a gas is doubled, the volume thus becomes half as
great. If the pressure on the gas is reduced to one-fifth its value, the volume
increases fivefold.

Boyle's Law can be expressed mathematically as,
y 1 (Voltime is proportional to the reciprocal of the pres-

P sure, or inversely proportional to the prossuie l 1

v. * 1

or V = K "p

(1) or PV = k, where k is a constant depending upon the quantity of

gas and the temperature

Thus for a given sample of a gas at constant temperature, the product
of the pressure and the corresponding volume, such as Pi and V,, P^. and V..,
etc., will be equal to the same constant value.

(2) P, V 1 - P, V,

In scientific work, subscripts are used to refer to diffeicnt conditions,
sometimes also called states, of the substance under investigation. Where a
change occurs, the subscript Y refers to the initial state, and the subscript Y
refers to the final state.

The data in Table 2- A, obtained for four grams of helium gas at 0 c C,
illustrate Boyle’s Law and the constancy of the PV product.

Table 2-A
Boyles Law

Pressure (atmosphere)

Volume (liter)

PV ( atmosphere-liter )

1.0852

20.65

22.41

1.0020

22,37

22.41

0.8067

27.78

22.41

0.6847

32.73

22,41

0.5387

41.61

22.41

0.3550

63.10

22,41

0.1937

115.65

22.41

1 Where a proportional relationship exists, the proportionality symbol oc , can be
replaced by equality and a constant, symbol k. Thus “8” is proportional to “2," or 8 « 2;
therefore 8 = k X 2. The value of the constant is obviously “4."

The States of Matter: Gases

The data of Table 2.1 are shown graphically in Figure 2.3.

01 020304050607 0 809 10 11 12
Pressure m atmospheres

Figure 2.3. Relation Between Pres-
sure and Volume of a Gas.

At a constant temperature, the volume
of a given sample of a gas is in-
versely proportional to the pressure.
The reciprocal relationship shown by
the curve is a hyperbola of the type

Helium is an exceptional example. For most gases at ordinary tempera-
tures, the PV product is almost but not quite constant. Later we shall see
that Boyle's Law is an approximation holding best at high temperatures
and at low pressures.

Within its limitations, Boyle's Law can be used to calculate what the new
volume will be when a gas is subjected to a change in pressure. From
Equation 2

(3) V, = V, -|i

Equation 3 states that the final volume, VL, will equal the original volume,
V„ multiplied by the ratio of the pressures. It is a simple matter to decide
logically, other than by memorization of the equation, which pressure is
in the numerator and which is in the denominator of the ratio. If the new
pressure is greater than the old, the new volume must be smaller and the
pressure ratio must be a fraction less than unity, and the greater pressure
therefore must be in the denominator. If the new pressure is smaller than
the old, the reverse holds.

Example 1: A quantity of gas occupies 76.8 ml at a pressure of 772 mm of
mercury. What will be the volume of the gas at 760 mm?

Solution: The final pressure, 760 mm, is smaller than the initial pressure,
772 mm. This change in pressure will result in an increase in the volume of the
gas. The fraction by which the original volume must be multiplied is therefore
greater than unity, that is, the numerator is greater than the denominator. The
new volume is:

V., = V, = 76.8 ml X = 78.0 ml

P 2 760 mm

7’ /if States of Matter • 0<wcs

Example 2: At what pressuie will a quantity of a gas, which fjccispies 100 ml
at a pressure of 720 mm, oceupv a volume oi 84.0 ml

Solution. The mathematical expression for Ernie's Law, equation 2, may he
solved for ? 2 thus

Since the volume is to be deci eased, the pressure must be mei eased. Hence
the fraction

V, , . 100 ml

— r must be greater than unity, or — 7 - — .

V 2 v 84 0 m3

and P 2 . = 720 mm

100 mi
84.0 ml

85 7 mm

NOTE: In wot king problems it is advisable to use the units of the vniiables
in addition to their numeucal values. Units max be multiplied and divided in the
same manner as numbeis.

^ ,, Jn , 48 cm-

Thus 6 cm \ 8 cm = 48 cm- and ~7 b cm

S cm

Units are especially useful in conversion from one set to anothei.

0.50 ft x L2

ft 10 mile

44 x

sec 5280 ft

min

mile

~TT

Units which are- apparently unrelated max
6.0 liter X 1.2 atmospheie -■
1

7.2 liter atmosphere X — -

30 degree

also be multiplied and divided.
- 7.2 liter atmospheie

liter atmosphere
degree

Students are sometimes confused by the fact that the ; dual compul-
sion of a gas, as in a bicycle pump, does cause its temperature to rise
whereas Boyle’s Law is based upon the premise that the temperature re-
mains constant. For Boyle’s Law to hold when gases are compressed or
expanded, the gas is brought to its initial temperature before the final
volume is measured,

4. Change in Volume of a Gas with Temperature: Charles’ Law; Absolute
Temperature. Now let us consider how the volume of a given sample* of
gas varies with a change in temperature if the pressure is kept constant.
It is common experience that^ if a gas is heated its volume increases
whereas if it is cooled its volume decreases. The behavior of gases due to
changes in temperature was first studied by Jacques Alexandre Cesar Charles,
a French physicist. In 1802 the French chemist, Louis Joseph Gay-Lussac
showed that, at constant pressure, all gases increase in volume by the same

The States of Matter: Gases

fractional amount when heated over the same temperature range. For each
degree centigrade rise in temperature a gas increases in volume by 1/273
of its volume at zero degrees centigrade, that is, 1/273 of the volume
that the gas would have occupied at 0°C. The value, 1/273, is the coefficient
of thermal expansion of all gases, and has units of deg -1 . In the rare case
of having just 273 ml of a gas at 0°C initially, the volume would be 274 ml
at 1°C, 275 ml at 2°C, 272 ml at -1°C, 271 ml at -2°C, and so on. This
leads to the remarkable conclusion that at ~273°C the volume of a gas will
shrink to zero, a most unusual annihilation of matter, but of course this
does not occur. All gases liquefy before -273°C is reached and so the laws
of gaseous behavior no longer apply.

At constant pressure, the volume of a gas thus depends upon, or is a
function of, the temperature; V = f(t). Let us assume that we have a
certain volume of a gas, V 0 , at 0°C. Then the volume, V t , at any higher
temperature, t, will be^he initial volume plus the expansion in volume due
to the rise in temperature (contraction for a decrease in

V t = V u + expansion

The expansion in volume per degree rise in temperature, is

temperature).

V„ and for a

change in temperature, t, above 0°C, the expansion is

273

V„ t.

(4) Thus V t

v -+ ik v ' A

= ~ V

• [>+£]- -

[273 + t]
273

273

[273 + t]

Thus the volume of a gas, V t , is linearly proportional to (273 + t).

A new scale of temperature may now be defined which has its zero
273 degrees below the zero on the centigrade scale. This new temperature
scale is called the Absolute Temperature scale, (symbol °A), or the Kelvin
Temperature scale, (symbol °K), for its inventor, Lord Kelvin (William
Thomson). The Kelvin scale and the Centigrade .scale are related by
T ( °K) = t (°C) + 273; in order to convert from degrees centigrade to
degrees Kelvin, 273 is added to the centigrade temperature. In scientific
work, the symbol ‘T always represents degrees Kelvin, while V represents
degrees centigrade. The magnitude of the degree, and hence an interval ,
is identical on both scales.

Equation 4 may now be written

( 5 )

Vt =-^2-T

* 273

or V t = k T where k is a constant whose value

depends upon the quantity of gas and the pressure.

We have thus derived that, at constant pressure, the volume of a given
sample of gas is proportional to the Kelvin or absolute temperature. This state-
ment is known as Charles’ Law and is illustrated graphically in Figure 2.4.

The States of Matter: Gasfis

50 100 150 200 250 SUO 150 400 tSO
Temperature, k

Figure 2.4,

Relation Between Temperature and
Volume of a Gas.

The straight linos indicate* that the
volume of a gas is a linear function
of the Absolute temperature Lino A
i.s for a gas whose volume is 273 ml
at 273'' K (O’ C). Lino B is for a gas
whose volume is UK) ml at 273 ^K.
Both lines “point” to the* origin,
but both art* dotted to indicate that
liquefaction occurs at low tempera-
tures and the gas laws are no longer
applicable

Since — = k, for the same sample of a gas at constant pressure but for

different temperatures, T x and T»,

V, V., T .

(6) = Ai. and V, = V, ii

1 1 A 2 1 1

The new volume of a gas, V.,, after a temperature change at constant
pressure, equals the original volume, V,, multiplied by the ratio of the
absolute temperatures. As with Boyle’s Law, it is an easy matter to decide
which temperature is in the numerator and which is in the denominator.
If the new temperature is greater than the old, the new volume must be
larger and the value of the temperature ratio must be greater than unity.
With a decrease in temperature, the reverse holds.

Example 3: A sample of air occupies a volume of 300 ml at 25 °C. If the
pressure is maintained constant, what will be the volume of the same sample of
air at ~5°C?

Solution: The first step is to change the centigrade temperature to Absolute
temperatures; 25°C = 298 °K, and -o°C ~ 268°K. Will the new volume be
greater or less than 300 ml? The new volume will be less because the temperature
decreases. The fraction by which the original volume is multiplied is therefore
less than unity; the numerator is less than the denominator, that is, 268/298.
The new volume will be:

V 2

= 300 ml x

268 deg
298 deg

= 270 ml

5, Simultaneous Correction for Temperature and Pressure* In deriving
the equations of gaseous behavior, one variable, either temperature or pres-
sure, was held constant. In the event that both temperature and pressure

The States of Matter; Gases

vary simultaneously, calculation is made first for either the change in
temperature or in pressure, and then for the remaining variable. Thus Boyle s
and Charles' Laws may be combined where there is a change in both
temperature and pressure.

(7) V 2 = Vl X |l X

Example 4: The observed volume of a quantity of a gas at 10°C and 750 mm
pressure is 300 ml. Calculate the volume of the gas at 0°C and 760 mm pressure.

Solution: The final volume can be obtained by direct substitution of the
proper values in Equation 7. This equation also states that the final volume is
the initial volume multiplied by a pressure factor and a temperature factor. Since
the pressure is increased, the volume will decrease due to the pressure change,
and the pressure factor is less than unity, or 750/760. Since the temperature is
decreased, the volume will decrease due to the temperature change, and the
temperature factor is less than unity, or 273/283. The final volume will be:

V 2

300 ml x

750 mm
760 mm

273 deg
283 mm

~ 286 ml

Because volumes of gases change with temperature and pressure, so called
“standard conditions'* have been arbitrarily defined. These are 0°C and
760 mm pressure, sometimes abbreviated STP (standard temperature and
pressure). Often it is necessary to calculate what the volume of a gas
would be at standard conditions. If Example 4 had been worded, “Find
the volume at STP,” the solution would be unchanged from the foregoing.

6. The Ideal Gas Law. Since the volume of a gas is inversely proportional
to the pressure and is directly proportional to the absolute temperature, the
gas laws can be combined elegantly into a single equation

V = k -T or PV = JcT

where k is a constant whose magnitude depends upon the quantity of the
gas and the units chosen for the pressure, volume, and temperature. If the
pressure is in atmospheres, the volume in liters, the temperature in degrees
absolute, and the quantity of gas is one mole, a chemical unit of quantity
of matter comprising 6.023 X 10 23 molecules, then the constant k has the
value of 0,08206 liter atm/deg mole. This is known as the molar gas constant
and is denoted by the special symbol “R.” Hence for one mole,

(8) PV = RT
and for any number of moles, n,

(9) PV = nRT

Equation 9 is known as the Ideal Gas Law. Any one of the variables in it
can be calculated if the other three are known. The concept of the mole is
discussed in detail in Chapter 5.

The Stutt'% <)f Matter, Oases

Example 5: Calculate the volume oi one molt* oi a gas at a tempai atme of
2T^C and a pressure of TOO rnrn

, t . . _ nKT

Solution * Equation 9 ls solved foi whence \ ~ ~—

The tempeiatuse must he in degiees absolute, and il the \aluo oi 0 08208
1 ami/deg mole is used lor R, then the piessuie must be m atmosphufs and the
volume will be found m liters.

1.00 mole x 0.08206

V =■

litei atm
deg molt*

300 deg

TOO mm

20 7 liters

760

atm

If the problem were given for inoie than one mule, for u moles, then the
answer would be (n s 26 7) liter.

7. Daltons Law of Partial Pressures. The gas laws apply not only to
individual gases but also to mixtures ot gases. When two or more gases that
do not react chemically with each other are mixed, each gas behaves in-
dependently as if it alone were in the containing vessel. Each gas makes
a contribution to the total pressure, its contribution being that pressure
which it would exert if it were alone in the container. The total pressure
is additive and is the sum of the pressures, now called partial pressures, of
the individual gases This behavior of mixed gases is generalized by Dalton’s
Law of Partial Pressures: Each component of a gaseous mixture exerts a
partial pressure which is the pressure it would produce if it existed alone
in the same volume, and the total pressure of the mixture is the sum of
the partial pressures of all the components.

Thus, if a quantity of a gas. A, in a 2 liter vessel has a pressure of
0.5 atm and if another gas, B, in a 2 liter vessel has a pressure of 0.75 atm,
then both gases placed in a single 2 liter container will produce a total
pressure of 1.25 atm, the partial pressure of the gas A therein being 0.5 atm
and that of gas B being 0.75 atm, provided the temperature remains constant
in all cases. Air is a gaseous mixture composed of approximately 20% oxygen
and 80% nitrogen by volume. At a total pressure of 760 mm the partial
pressure of the oxygen is 20% of the total pressure or 1/5 x 760 mm ~~ 152 mm;
the partial pressure of the nitrogen is 8071 of the total pressure or 4/5 X
760 = 608 mm. The gas laws can be employed for mixtures of gases using
the total pressure of the mixture or to the individual gases using their
partial pressures.

Application of Dalton's Law is common when gases are collected over
water, as in a eudiometer tube. Some of the water evaporates into the
space occupied by the gas and intermingles with it, and thus we do not have
a single pure gas but a mixtuie of two gases, the collected gas and water
vapor. The total pressure of the gas sample is therefore the sum of the
partial pressures of the collected gas and the water vapor. The pressure
exerted by water vapor depends upon the temperature alone and is given
for various temperatures in Appendix III. To find the pressure exerted

The States of Matter: Gases

by the gas a\one, that is, its partial pressure, the vapor pressure of water
at the temperatme of the sample is subtracted from the total pressure
of the mixture. This value would also be the pressure of the dry gas in the
same volume and at the same temperature.

Example 6. 500 ml of hydrogen gas are collected over water at 40 °C and at
765 mm pressure. What will be the volume of dry hydrogen at standard conditions?

Solution: The hydiogen collected is not pure hydrogen but a mixture of
hydrogen and water vapor. The vapor pressure of water at 40 °C is 55 mm so
that the partial pressure of the hydrogen in the mixture is the total pressure
minus the vapor pressure of water, i.e, (765 mm - 55 mm) or 710 mm. Equation 7
can be used to find the volume of dry hydrogen at STP.

= 500 ml x

273 deg
313 deg

710 mm
760 mm

= 407 ml

The pressure ratio is 710/760 because the change in pressure of the dry hydrogen
is from 710 mm to 760 mm, resulting in a decrease in volume.

8. Graham’s Law of Diffusion. When a gas jet is opened the odor of gas
soon can be detected in the room, even when the air is presumably at rest.
The gas obviously penetrates the air and distributes itself throughout the
room. The same phenomenon is observed readily when a colored gas such
as bromine is released in a glass jar. The bromine spreads uniformly through-
out the entire space in the vessel. This behavior of gases and vapors is called
diffusion . This process of diffusion occurs even when two gases are separated
by a porous partition such as unglazed porcelain.

All gases do not diffuse at the same rate. At a given temperature and
pressure, a light gas, that is, one of low density, diffuses more rapidly than
a heavy one, a gas of high density, in 1832, Thomas Graham showed that
the rates of diffusion of gases are inversely proportional to the square roots
of their densities. This statement is known as Grahams Law of Diffusion,
If r 1 and n. represent the rates of diffusion of two gases whose densities
are di and dj, respectively, Graham's Law may be written

Example 7 . At STP, one liter of oxvgen weighs 1.429 g and one liter of
hydrogen weighs 0.0899 g. Calculate the relative lates of diffusion of the two
gases.

Solution' From Equation 10

^hydrogen
t OAygeii

dpxygen

dhydrogen

1.429 g/i
0.0899 g/1

=. 3.99

Oxygen is about 16 times as dense as h\drogen, and hence diffuses only about
1/4 as rapidly as hydrogen.

A technique based upon gaseous diffusion was used to separate partially
the isotopes of uranium for the “atomic” bomb (Chapter 49) during World
War II, and today still is the principal source of the Uranium 2515 isotope.

//«' States uf Matter Gust's

This isotope, present to the extent oi but 0.7 r c m nutuial maimim, is
separated from the Uranium-- 18 isotope bv successi\e diiiusions of the
gaseous compound, uranium hexafluoride.

9. The Kinetic Molecular Theory, Inasmuch as (ill gases, irrespective
of their chemical nature, conform to the gas laws with good accuracy it
would appear probable that all gases have a similar mechanical stiueture.
The model postulated for the observed behavior oi gases is called the Kinetic-
Molecular Theory. As its name implies there uie two basic postulates-
kinetic and molecular.

The main assumptions of this theory are:

I. Molecular: Gases are composed of minute, discrete particles called mole-
cules. The molecules of any one gas are all alike.- The molecules art* separated
by average distances large in comparison to their own diameteis, except
at high pressures. This permits the molecules to be considered mathematically
as points. Between the molecules is empty space. This assumption accounts
for the high degree of compressibility of gases. An inctoase m pressure
merely forces the molecules closer together and the space between them
diminishes. The molecule itself is assumed to be incompressible. It is further
assumed that there are no attractive forces between tin* molecules,

II. Kinetic* The .gaseous molecules are in rapid motion. They move in
straight lines in all directions unless they collide with other molecules or die
walls cf the container, whereupon they rebound and continue to move off
in a straight line. The molecular collisions. are deemed to be elastic, that is,
the stfm of die kinetic energies of the colliding molecules before and after
collision is unchanged. The average distance travelled by the molecules be-
tween collisions is called the mean free path. This molecular motion accounts
for the fact that a gas will fill any inclosed space and, in conjunction with
the large spaces between the molecules, accounts for the phenomenon of
diffusion.

Pressure is due to the number of collisions which the molecules make
with the walls of a container per unit area and per unit time. It is the
machine gun-like bombardment of the molecular bullets upon the container
wall that gives rise to pressure.

The velocities of the molecules and their kinetic energies depend upon
the temperature. For any object, or a molecule, of mass W and velocity
V the kinetic energy is given by l k m u\ The higher the temperature the
greater the molecular velocity and hence the kinetic energy. At a given
temperature, the kinetic energies of the same quantities of two different
gases are equal. At absolute zero die kinetic energy of a molecule becomes
zero. At a given temperature all molecules of a gas do not have the same
value of velocity and energy. Rather there is a distribution of velocities
and energies, as shown in Figure 2.5. Only a small fraction of the molecules
have extreme values of velocity; most have an intermediate value. For
each temperature there is a mean value of velocity characteristic of that
temperature. When the temperature is raised, the distribution curve shifts
to the region of higher velocities so that the number of molecules having

2 Except for the minor perturbations imposed by isotopes (Chapter II )

The States of Matter: Gases

these higher velocities is increased. The mean velocity of a molecule of
hydrogen at STP is about 1.85 kilometers per second; for oxygen the velocity
is one-fourth that of hydrogen, or about 0.46 km/sec.

10. Explanation of the Gas Laws by the Kinetic-Molecular Theory.
Boyles Law : If, for example, the volume of a gas is decreased, it is the
free space or distance between the molecules that is diminished. If the tempera-
ture remains constant, the velocity of the molecules will also be unchanged
and consequently a greater number of collisions will occur with the walls
of the container, resulting in an increased pressure. The pressure will
increase in proportion as the volume is decreased, as described by Boyle’s Law.

Number

Molecules

The distribution of velocities for an assemblage of gaseous molecules is shown
for two temperatures. The curves are not symmetrical bell curves. The ordinates,
A and B, represent the arithmetic mean velocities for the lower and higher tempera-
tures, respectively. Few molecules have extreme velocities. About half the mole-
cules have speeds within 30% of the mean velocity. Only about one m 50,000 has
a velocity greater than four times the mean value. Since kinetic energy depends
upon velocity, the energy could have been plotted as the abscissa without altering
the shape of the curves. The total energy of the gas would then be proportional
to the area under a given curve. For one mole of an ideal gas, the energy
3

is equal to RT.

Figure 2.5. Distribution of Molecular Velocities and Energies.

Charles 9 Law : A rise in temperature results in an increased velocity of
the molecules. If the volume is maintained constant, the frequency of mole-
cular collision with the walls of the container will increase and the pressure
will increase. This is a corollary to Charles’ Law: that if the volume
of a gas is constant, the pressure is proportional to the absolute temperature.
If the pressure is to be kept constant, however, the volume of the container
must be expanded. The molecules, with their increased velocities at higher
temperature, have to move a greater distance before colliding with the

i'lu \ of si utter Gases

container wall and so the frequency oi collision pa unit an m and the
pressure will be unchanged

Daltons Law: Because oi the relatively great distances between gaseous
molecules and because of the absence ot intermoiecului atti action, molecules
of different gases act independently of each other because ot tins inde-
pendent behavior, each gas produces a pn-ssuie a*> d it war alone in the
container, and the total pressure of a mixture is the sum oi the individual
or independent pressures.

Grahams Late : For a given numbei of gas molecules the total kinetic
energy is independent of the nature of the gas and depends upon the
temperature only. For the same number of molecules ot two diifeient
gases, 1 and 2, at the same temperature the kinetic < names me equal so

that i mi Mi 2 = 4- ni,u/. The ratio of then velocities is --

2 2 " u. \ m x

Hence the relative velocities of the molecules are rnasrk piupoitionul to
the square roots of then respective masses Smei Ua* speed *»* (effusion
of a gas is dependent upon the velocity of the individual mnlecidm and since
the density of a gas is proportional to the mass oi its molecules, if follows
that the relative rates of diffusion of gases me mveiselv piopoitmual to
the square roots of their densities.

11. Deviations from the Gas Laws. The laws of gaseous behavior, namely
Boyle's Law, Charles’ Law, Dalton's Law, and Graham's Law, are not
obeyed exactly. In other woids, experimental determination is not in per-
fect agreement with theoretical calculation. At ordinary tempi latures and
pressures, such deviations usually are not greater than a few pel cent and
calculated volumes serve satisfactorily as approximations. The concept of
a gas that would obey the gas laws exactly is. however, of value to the
scientist. Such a gas is known as an Ideal Gas in distinction to actual or
real gases. A gas having the attributes listed in the Kinetic-Molecular Theory*,
namely, relatively large distances between the molecules and no attractive
forces between them, would be an ideal gas. No such gas actually ovists.
The equation PV = nRT, a composite of the gas laws, holds exactly only
for ideal gases and hence is known as the Ideal Gas Law. Real gases are
said to be more or less ideal by the degree to which they approximate the
gas laws. Generally real gases approach ideal behavior at high temperatures
and at low pressures.

The deviation from ideal behavior of a real gas at high pressures be-
comes evident when we consider that during the compression of a gas only
the free space between the molecules is diminished and not the volume
of the molecules themselves. At ordinary conditions the volume occupied
by the molecules is only about 0.0014 of the total gaseous volume but at a
pressure of one or two hundred, atmospheres the volume of the molecules
is an appreciable fraction of the total volume. Hence when additional
pressure is applied to a gas already under great pressure, that is, when the
molecules of a gas are already close together, the contraction is less than
Boyles Law would indicate. Since the volume does not diminish in pro-

The States of Matter: Gases

portion to the increase in pressure, the PV product is larger than that
predicted by Boyle’s Law, and the gas is less compressible than an ideal gas.

In an ideal gas the molecules are assumed to have no attractive forces
and hence no tendency to cohere when they collide. If gases were ideal,
this lack of attraction infers that the liquid state would not exist. By suf-
ficient cooling and compression, however, it has been possible to liquefy
all gases. This indicates that there me attractive forces which become im-
portant especially at low temperatures where the kinetic energies of the
molecules are low. Because of the attractive forces the compressibility ot
a real gas at low temperature is greater than that of an ideal gas. As a gas is
compressed the attractive forces draw the molecules even closer together
than would be the case if no such forces existed. In effect the attractive forces
act like an additional pressure to produce a smaller volume than that pre-
dicted by the gas law's. These attractive forces between molecules are the
so-called Van der Waals forces. They are electrical in nature and will be
considered in Chapter 13.

In 1879 J. D. Van der Waals, a Dutch physicist, proposed an equation
which added two teims to the Ideal Gas Law and thereby took into account
the attractive forces between molecules and the volume occupied by the
molecules themselves. The Van der Waals equation for one mole of a gas is

(H) (p + ^)(v-b) = RT

Both “a" and “b” are constants whose values are specific for each gas; “a” ex-
presses the attractive forces between molecules and “b” is the volume occupied
by the molecules. Where the attractive forces and the volume of the molecules
are negligible, Van der Waals’ equation approaches the Ideal Gas Law
as a limit.

The deviations of gases from ideal behavior are shown in Figure 2.6,
where the PV product is plotted against P at constant temperature for a
hypothetical real gas.

Pressure
X Volume
PV

Fig-aie 2.6 L

Variation in PV with P at
constant temperature.

Line A represents the behavior
of an ideal gas. At constant
temperature the product PV is
constant. Line B is the behavior
of a real gas. At very low pres-
sures, it follows line A indicat-
ing ideal behavior. At moderate
pressures attractive forces pre-
dominate, the gas is more com-
pressible than an ideal gas, and
the PV product is below that of
an ideal gas. At high pressures,
the volume of the molecules
themselves make the gas less
compressible than an ideal gas,
and the PV product is higher
than that ot an ideal gas.

Pressure, P

Table 2-B

Physical Properties

The States of Matter. Gases

Density of
liquid at
critical
temperature
and pressure
g/ml

0.031

0.430

0,311

0.460

0.573

0.235

0.45

0 52

0.102

Critical

pressure

atmospheres

12.8

49.7

33.5

72.8

111 5

71.7

1 i . i

45.8

217.7

Critical

temperature

°C

q oq ^ p no cj iq

c> co o r}5 oi cd r- oi

-jo

1 i 1 1 1

Density

STP

g n

0.0899

1.429 |

1.251

1.250

1.977

3.214

0.771

1.978

2.927

0.717

999.9

Melting
point at

1 atmosphere
°G

^ O c,-j co

^ ^ ^ (g tw ^ S?

<£ 22 2? £: 4^ co oi jo **

oi ©1 03 ^ r-l *~i r~4

> * * l CO 1 1 1 1 1

00*0

Boiling
point at

1 atmosphere
°C

-252.8

-182.8

-195.8

-190

-78.5 (sublimes)

- 34.6

- 33.35

- 88.49

- 10.0

-161.5

100.00

Substance

1. Hydrogen

2. Oxygen

3. Nitrogen

4. Carbon monoxide

5. Carbon dioxide

6. Chlorine

7. Ammonia

8. Nitrous oxide

9. Sulfur dioxide

10. Methane
*

11. Water

The States of Matter. Gases

12. Critical Temperature and Pressure. All known gases have been
liquefied by the application of suitably low temperatures and sufficiently
high pressures For each gas, however, there is a temperature above which,
no matter how high the applied pressure may be, liquefaction will not
occur. Above this temperature the molecules are moving too rapidly and
their kinetic energies are too great for the attractive forces between the mole-
cules to cause them to cohere. This maximum temperature at which a gas
can be liquefied is known as the critical temperature , and the pressure
just required to cause liquefaction at the critical temperature is called the
critical pressure. The “boundary line” between a liquid and its vapor above
it is the meniscus of the liquid. If a liquid were heated in a sealed transparent
tube, at the critical temperature the meniscus would disappear and there
would be no distinction between the liquid and gaseous states. The critical
temperature is not to be confused with the boiling point, which is the
temperature at which liquefaction will occur under a pressure of one
atmosphere. If the temperature of a gas is lowered, the kinetic energy of
its molecules becomes less, and hence less pressure is required to liquefy
the gas until at its boiling point one atmosphere is sufficient. In Table 2-B,
critical and other data are given for several substances; all are gases under
ordinary conditions except for water which is included for comparison. In-
spection of the table indicates that hydrogen cannot be a liquid at room
temperature but chlorine, which is a gas at room temperature, could be
liquefied by the application of sufficient pressure.

QUESTIONS

In doing problems, the student should not “automate” an equation by merely
plugging in the proper values for the variables. No equation should be used
without an understanding of its theory and limitations! In addition, the proper
units should be inserted in each equation used to solve a problem. Mathematical
operations are simplified by the use of exponentials and logarithms. The student
is advised to read the section on Mathematical Operations in the appendix. It
is suggested that attention be given to the section entitled “Significant Figures” to
avoid carrying out unnecessary multiplications ancj divisions.

1. What are the three states of matter? What are some of the important physical
properties of each state which aid in identifying a substance?

2. State and illustrate the following laws: Boyle’s, Charles’, Dalton’s, and
Graham’s. Write the equation for each.

3. What are the main assumptions of the kinetic-molecular theory? Show how
the theory explains the various gas laws.

4. On the basis of the kinetic-molecular theory, explain how the pressure
of a gas should vary with temperature if the volume is kept constant.

5. Define and illustrate the following: critical temperature, critical pressure,
boiling point. What would be the theoretical critical temperature and boiling
point of an ideal gas? Explain.

6. Explain the statement that “all molecules of a gas do not have the same
energy at the same temperature.”

7. What is meant by the average velocity of the molecules of a gas? On what
conditions does this velocity depend?

T/ic Sta!<‘\ of KUt'tf! ( asrs

8. How is the critical tempei ature related to the attractive iou-es between
gaseous molecules? In which gas would the attractive toices he gieater,
nitrogen or chlorine?

9. In converting to or from STP, 273 degrees and 760 mm an* alwavs »»*, op-
posite sides of the tempeiature and pressure latios, one m ti.e 'mnu-i,, *.*r
and one in the denominator. Win should this be so?

10. Explain the deviations of real gases from ideal behiiMoi What an* the causes
of these deviations? What is indicated bv the lowest point of cmv« B su
Figure 2.6?

11. Which of the substances in Table 2-B could be liquefied at -1 00 ( J \t one
atmosphere pressure, which are liquid at —100 G‘ J Which uc solid' 0

12. Under what conditions of temperature and piessme would the \ an dei W a ils’
equation approach the Ideal Gas Law?

13. (a) A quantity of gas occupies So liteis at 750 mm piessme at 21 ( \\ hut
volume will it occupy at standaid conditions' 0 (b) \t 800 mm and IS t J

Ahs. 7S.0 htejs, ON 1 hit is

14. A tank is tested to withstand a piessure of 12 atnmspheies. it is lilletl with
a gas at 0°C to a pressuie of 7 atmospheo s Will it hr sale to subjn f the
filled tank to a temperature of 300° C?

A»s \n

15. An elastic balloon is filled at 20 °C and one atmosphere to a volume of 1000
cu ft with hydrogen gas. When iclcased it rose to a height where the tempera-
ture was -30°C and the balloon expanded to 30(H) cu it What is the
atmospheric pressure at the new height* 5

Ans\ 0 27 atm

16. A certain quantity of gas was collected over vvatei at 21 t" and 756 nun
pressure and measured 246 ml. What would he the volume oi the third
gas at 40° C and 780 mm pressure?

Arts: 247 ml

17. A sample of air is collected over water at 25 °C and a total pi ensure of
759 mm. What is the partial pressure of oxygen m the sample*?

Ans, 146 8 mm

18. A quantity of a gas A, exerts a pressure of 250 mm in a volume of 200 ml.
Another gas, B, exerts a pressure of 140 mm in a volume of 75 ml U both
gases are placed in a 150 ml container, what will be the total pressure
and the partial pressure of each gas r ^

19. (a) What will be the pressure, in mm, of one mole of a gas m a 2 50 liter
volume at 20°C? (b) What will be the pressme of 3.5 moles of this gas umh’i
the same conditions?

20. A quantity of 40 ml of a gas, initially at 100 °C and 1.0 atmosphere pievnue,
is cooled to 10°C and compressed into a volume of 10 ml. What is the
new pressure of the gas?

21. A quantity of 20 ml of a gas, initially at 10°C and 400 mm pressure, is
compressed to a pressure of 600 mm. What must the final temperature he
for the volume to remain unchanged?

22. If two liters of hydrogen diffuse through a small opening in two hours,
how long will it take one liter of each of the following gases to diffuse
through the same opening under exactly the same conditions: oxygen, nitrogen

The States of Matter: Gases

carbon dioxide? The densities of the gases, in g/1, are: hydrogen, 0.0899;
o\\ gen, 1.43; nitrogen, 1.25, carbon dioxide, 1.96.

Ans : 4.0 hrs; 3.7 hrs; 4.7 hrs

23. At STP a one hundred liter container weighs 234.85 g when filled with
hydrogen. What will it weigh when filled with oxygen?

24. Calculate the density of oxygen at 127°C and 950 mm pressure.

25. The weight of 350 ml of a gas at 700 mm pressure and at 40°C is 0.825
grams. Calculate the density of the gas in g/1 at STP.

26. Two liters of gas A and three hteis of gas B, each initially at two atmospheres
pressure, are mixed. The volume of the mixture is then made ‘one liter by
changing the pressure. What is the partial pressuie of each gas in the final
mixtuic? Assume the temperature remains constant throughout the process.

27. (a) Using Van dei Waals’ equation, calculate the pressuie exerted by one mole
of carbon dioxide gas at 27° C and 10 0 atm pressure. For carbon dioxide,
a = 3.6 atm literVmole- and b = 0.043 iitei/mole. (b) Calculate the pressure
assuming carbon dioxide is an ideal gas.

APPENDIX

Kinetic Equation for Gases

The Ideal Gas Law can be derived from the fundamental postulates of the
Kinetic-Molecular Theory (page 24). Assume an assemblage of n gaseous mole-
cules in a cubic container of side L; each molecule has a mass, m, and a velocity, u
(Figure 2.7). Let us derive an expression for the pressure on the right-hand face
of the contamei . The pressure, P, which is due to the collisions of the molecules with

the walls of the container is the force, F, per unit area, (i, of a container wall:

P — F/a. Force is the change of momentum with time,
and momentum is the product of mass and velocity, or

mu. For the collision of a single molecule, such as A,

with the container wall, the momentum before collision
is -f-mii and after collision is -mi/, The magnitude of the
momentum is unchanged by collision since collisions are
assumed to be elastic (and the container immovable) but
velocity is a vectorial quantity and its direction must be
indicated by the •+■ and — signs. Hence the change in
momentum upon the collision of a single molecule is
-Tmu - [~mu) == -f2mu momentum change/ collision.

To obtain the force, this quantity must be multiplied by
the number of collisions per unit time. Assume that an
unimpeded molecule travels a distance, 2L, back and forth across the container
between collisions, then

Figure 2.7

distance

time

distance

collision

The force exerted bv one molecule is

u collision

2L time

„ momentum change

2 mu — X

collision

u collision
2L time

mu 2

The Sidfc? n/ Vtiffer Gases

If the assumption is made that one-third of the molecules are moving m each of
the X, Y, and Z directions, then n/3 molecules will strike the iight-hand wall of

ri m ti-
the container. The total force is equal to ~ —

Since P = F/a and the area of the wall is L 2 , then

_ nmu~ v 1 timu-

F “ ~3lT ' L 2 ~ 3L 3

But L 3 equals the volume, V, of the container, so that P
1

PV = -j— nrmr . This is the kinetic equation for gases.

Since the kinetic energy of the ti molecules is mnu J , then
PV = 2/3 E where E is the kinetic eiu*ig>.

The kinetic energy is proportional to the absolute tempoiatwo, T, hence
PV = kT where k is a constant dependent upon the quantity of gas
For one mole of a gas, k is equal to the molar gas constant, H, ami P\ - - KT;
for n moles, PV = nRT.

The States of Matter:
Solids and Liquids

The liquid and solid states are sometimes called the condensed states of
matter. Here the attractive forces between the molecules, which are at a
minimum in the gaseous state, have become of sufficient relative importance
because of reduced temperature, increased pressure, or a combination of both,
to cause condensation.

The distances between the particles in the liquid and solid states are much
smaller than in the gaseous state. In the solid state such distances are of the
order of I0' 8 cm, or 1 A (A = Angstrom unit; 1 A =1 x 10" 8 cm). Distances
between particles in the liquid state are generally slightly larger, but in the
gaseous state these distances are a thousand or more times greater. This can
be inferred from a comparison of the volumes occupied by equal weights of
the same substance in the different states. Thus, at its normal melting point,
the density of solid mercury is 14.2 g/cm 3 , that of liquid mercury is 13.7
g/ cm 3 , while the density of mercury vapor at its normal boiling point is 0.0039
g/cm 3 . Indeed the particles in the condensed states may be considered to be
so closely packed that they are almost touching. 1 Compared with gases, both
solids and liquids are quite incompressible. At 20° C, an increase of pressure
of one atmosphere on one liter of liquid water will reduce its volume by
approximately 5 x 10" 5 liter. Similarly the variation of volume with tempera-
ture is small. A one degree increase in temperature will increase a one liter
volume of iron by 3.5 x 10“ 5 liter.

Because the forces between the particles in the liquid and solid states
are large and because these forces differ in magnitude and in nature from
substance to substance, there is no single general equation which describes
the behavior of the condensed states as does the Ideal Gas Law for the
gaseous state, where the intermolecular forces were so small they could be
assumed to be zero.

^“Touching” is an anthropomorphic word. Many scientific concepts, so precisely de-
scribed by mathematics, are limited by semantics. In Chapter 13, we shall see that all
particles are surrounded by electrical fields. The interaction of these fields might be
construed as “touching.”

The State.-, of Mutter Solid.- and Lit/, nth

L The Solid State. Most solids are crystalline. Crystals are substances
which have definite geometnc external loims. They have plane laces which
meet to form straight line edges and have angles between the faces which
are characteristic of the substance. This most evident chaiaetenstic of
a crystalline solid, its symmetry of geometric form, is due to regularities
in the positions of corresponding faces and edges. Because of these legu-
larities, a crystal has axes and planes of symmetiy. When rotated once about
an axis of symmetry, a crystal will present a like appearance, two, tluce.
four, or six times. The axis is then said to be an axis of two-, three-, hnu~, or
sixfold symmetry. The axis through the centers of two opposite faces of a
cube thus is an axis of fourfold symmetry A plane of symmetiy divides a
crystal into two similar halves which have the relation of minor images to
each other.

On this basis crystal fonns can be divided into six systems. The* classifi-
cation is based upon the relative lengths of the axes of symmetiy and the
number of pairs of axes that are at right angles to each other, ("ailing the
axes a, b , and c (and cl in the hexagonal system where there uie lour pint-
cipal axes), the six systems are summarized in Table 3-A. It will !><> seen
that crystals may vary from the highly symmetric cubic to the completely

Table 3-A

The Crystal Systems

System

Axes

Angles

hxamplr

Cubic

a = h = c

all axes at right angles

Bock salt

Tetragonal

a — b ^ c

all axes at right angles

White tin

Orthoihombic

a ^ b ^ c

all axes at light angles

Rhombic snlfui

Monoclinic

CL h zfc c

only one paii at right
angles

Calcite

Triclinic

a ^ b ^ c

no a\e$ at right angles

Potassium dichmmate

Hexagonal

a = b = c 3 * d

thiee equal axes in one
plane at 60° to each
other. The fourth
axis, d, at 90 to the
plane of the othei
three

(haphite

general triclinic system. With the exception of those in the cubic system,
crystals are anisotropic, that is, many of their properties are not the same
in all directions. Thus the heat conductivity and optical properties* such
as the speed with which light is transmitted, vary with the direction in
which these properties are measured in the crystal.

Only rarely do we find a perfectly formed crystal. Due to interferences
in their growth, crystals are usually distorted to some degree. Perfect crystals
can, however, be grown by suspending a tiny seed crystal in a saturated
solution of the substance to be crystallized and then allowing the solution
to evaporate slowly. If crystals grow on the bottom of a container, some

The States of Matter. Solids and Liquids

Simple

One paiticle is centered
at each corner of a cube.

Cubic System

Body centered

One particle is centered
at each corner and one
is centered in the body
of the cube.

Face centered

One particle is centered
at each corner and one
is centered in each plane
face of the cube.

Tetragonal System

Monoclinic System

Body

Simple centered

End face

Simple centered.

Simple

Orthorhombic System

Body

centered

End face
centered

Triclinic System Rhombohedral System

Hexagonal System

a, b, and c represent the axes of the unit cell; (X , /3 : and 7 represent the angles
between the axes.

Figure 3.1 The Fourteen Unit Cells.

Vhv S tat( \ t>/ Miitt* ’ Soluis unu Liquid?

faces grow more rapidly than others, producing distorted and nonsym-
metric shapes.

2. The Unit Cell. When a crystal is shattoied h\ a hammer blow the
fragments, though much reduced in si/e, tend to hav e similai plane faces
and interfacial angles as did the original crystal. If we visualize a progres-
sive subdivision of a macroscopic crystal to smaller and smaller pi optu turns,
ultimately we would come to the smallest entity which still shows the same
crystalline form and has the same physical pioperties as does the macro-
scopic crystal. This smallest subdivision is called the unit cell Kuither
subdivisions of the unit cell would produce entities winch do not have the
physical properties of the ordinary crystal. A macroscopic crystal thus can
be considered as being built up of unit cells end-to-end m three dimensions,
much like a three-dimensional brickwork. The unit cell itself is composed
of particles arrayed in a specific pattern and held together In electrical
forces. These particles are very close but at fixed distances from each other,
their centers approximately 10~ H cm apart. The lines joining these centers
intersect at angles characteristic of a given substance. Such an ordered array
of particles is the basis of a crystal's symmetry and is our fundamental
model of crystalline structure.

A crystal has the same type of symmetry as dot's the unit eel! and the
crystal system thus depends upon the unit cell Unit cells winch art' simple
cubes therefore form crystals having cubic xynunctn. There ;ue fourteen
kinds of unit cells. These are illustrated in Figure 3.1. The illustrations in
Figure 3.1 were drawn to show primarily the shape of the unit cells. The
particles are not to scale. A more realistic view, with the particles drawn
to scale, is given in Figure 3.2 for a simple cubic crystal, iieie the proximity
and close packing of the particles are indicative of their “touching.”

The specific arrangement of the particles in a unit cell is determined by
a technique known as X-ray diffraction analysis. The particles uie spaced
at such distances that they can act as a diffraction grating for radiation of

The lines between particles denote a unit cell. The center points of the particle's
throughout the crystal establish a geometric array known as the crystal lattice.

Figure 3.2. Arrangement of Particles in a Simple Cubic Crystal.

The States of Matter: Solids and Liquids

X-ray wavelengths, much as lines drawn on a metal or glass plate and spaced
about 1/4000 cm apart can diffract visible light. In this technique a beam of
monochromatic (single wavelength) X-rays is directed against the surface
of a crystal. The “reflected” beam will have various intensities depending
upon the angle of incidence upon the, face of the crystal. The angle at
which maximum intensity of the reflected beam occurs can be determined
by rotating the crystal. Under these conditions, substitution in the general
equation for optical diffraction, enables one to calculate d, the distance'

where n = an integer, the order of the diffracted beam
n A = 2 d sin 0 ^ — the wavelength of the X rays

6 = the angle for maximum intensity of reflection

between the planes of the particles within the crystal. From quite ordinary
geometric viewpoints it is then’ possible to postulate the arrangement of
particles within a unit cell, the dimensions of the cell, and thereby the macro-
scopic crystalline structure.

3. Melting and Phase Changes. Actually the particles in a crystal are not
at rest but vibrate about mean central positions at which, for many mathe-
matical purposes such as the concept of the unit cell, the particles may be
assumed to be fixed. The magnitude of the vibration is determined by the
nature of the particles, the electrical forces between them, and the tem-
perature. When the temperature of a solid is raised, the particles vibrate
with increased energy but still about the same mean positions so that the
over-all shape of the crystal is maintained. At a certain temperature some
particles achieve sufficient energy to overcome the attractive forces which
bind them to fixed positions and move off independently. Thereby the solid
loses its definite form and the result is the liquid state. The temperature at
which this occurs is the melting point of the solid. For a pure substance
the melting point is a single definite value. In the process of melting, a
solid absorbs a definite quantity of heat energy per unit weight. This is
called the heat of fusion. During the melting of ice, 80 calories are required
to melt one gram; the heat of fusion of ice is thus 80 cal/g. Though heat is
supplied during the melting process, the temperature remains constant.
Work must be done and hence energy must be supplied to separate and
to pull away each particle in the crystal lattice against the electrical
forces attracting it to its neighbors. The energy supplied by the heat of
fusion serves this purpose. Just such an amount of heat is absorbed as is
required to overcome the attractive forces without causing an increase in
temperature.

Many substances can exist in more than one crystalline form. This
phenomenon is known as polymorphism and, in the case of elements, the
various forms are called allotropes. In most cases of polymorphism one
form f^n change to another as it is heated above or cooled below a tem-
perature that is definite for the particular substance. This temperature is
called the transition temperature or transition point of the substance. Thus
sulfur crystals formed at room temperature have the rhombic form. On
being heated above 96 °C, the transition temperature, they change slowly
to another crystalline form, the monoclinic. During such a change a definite

7 he States of Mattel Solids and Ltt/utds

Quantity of energy is absorbed for each gram ot sulfur turns ormed. Obvi-
omlv there is no essential difference between melting and tnm.sihon to
another form in the solid state. Such changes, and it wdi be seen that
boiling also belongs to this category ot change, aie hi until ^ <1, amirs of
phase A phase is a homogeneous legion ot matter and ditters m meaning
from the term “state.” Though there are only three' states of mattei there
can be any number of phases. A single state ot mattei pci force exists as a
single phase but oil and water, which do not mix, exist as two phases m the
liquid state. The distinguishing mark of a phase is a hounduiv imo about
it which separates it from another region ot matter. Examples aie the
meniscus between a liquid and the air abo\ e it, and the line ot demarcation
between oil and water. Many solid alloys are composed ot more than
one phase. Thus a change of state also involves a change of phase but
the reverse is not necessarily true. Both phase changes and changes of state
occur at characteristic constant temperatures and imolve the abruption
or emission of a definite quantity of heat energy per gram.

4. Crystal Binding. On the basis of the particles composing the crystal,
crystals can be classified as atomic, molecular, or ionic In an atomic er>stal
like iron, it is iron atoms which make up the unit cell and wlm h occupy
the points of the crystal lattice. In solid iodine it is iodine molecules, each
composed of two iodine atoms, whereas .with ordinal \ salt, it is ions of
sodium and chlorine which compose the unit cell. An ion is an atom or a
group of atoms which have acquired an electric charge, hi all cases the
binding between the particles is electrical in nature. The bond, or electrical
attraction, between ions is strong and, as a result, ionic crystals generally
have high melting points. The bonding between molecules is due to the
so-called Van der Waals’ forces. These are identical with the attractive

forces which cause deviations in the gas laws and cause condensation of a
vapor to the liquid state. They are relatively weak and so molecular crystals
are soft, have low melting points, and may even sublime, that is, pass
directly from the solid state to the vapor state without forming liquid,
Iodine and some organic solids are of this type, eg., naphthalene, the
major constituent of moth balls.

Certain rare solids show no evidence of regular structure and are known
as amorphous solids. They may be formed by the sudden condensation
of a vapor upon a cold surface. Matter vaporized from the hot filament
of an incandescent bulb and deposited upon the relatively cold surface
of the glass envelope is a structureless solid of this nature. Time is re*
quired for such condensed particles to orient themselves into fixed posi-
tions if a crystalline pattern is to result. Amorphous solids can often be
changed into crystalline form by holding them for some time at an elevated
temperature (annealing).

Another apparent class of solids are vitreous or glass-like materials.
Examples are ordinary glass and tar at a low temperature. These are not
truly solids but are supercooled liquids of such large viscosity that they
appear to be solids. Being liquids they have no crystalline structure. When
heated they do not melt at a definite temperature but soften gradually
and become more fluid. They are isotropic and, under great pressure, will

The States of Matter: Solids and Liquids

flow slowly. With the passage of time, the particles of a vitreous substance
may very slowly orient themselves into a crystalline pattern.

The Liquid State

At the melting point, particles or groups of particles, of the solid
attain energies sufficient to break away from the main body of the crystal.
The attractive forces between the solid particles are insufficient, relative
to the increased kinetic energy which the particles then have, to maintain
fixed positions and hence a definite shape. These forces are still strong
enough, however, to keep the particles close together in a condensed state,
namely, the liquid state. In the liquid state there is no definite orientation
among the particles composing the liquid. Just above the melting point
there may be some clusters of particles which maintain some orientation
inherited from the solid state but any semblance of organization diminishes
as the temperature, and hence the kinetic energy of the particles, is increased
until finally no trace of organization remains and each particle moves at
random. It is this variation in the degree of orientation which make the
liquid state the most difficult to treat theoretically. The complete random-
ness of the gaseous state and the perfect order of the solid state, lying at
the two extremes of structural organization, admit of simpler theoretical
treatment.

5. Surface Tension. Unlike ga'ses, liquids (and solids) have surfaces. The
mere existence of a surface gives rise to a special kind of properties in con-
trast to bulk properties. Thus molecules, not on the surface but within the
body of a liquid, are surrounded on all sides by other liquid molecules.
On such an interior molecule there will be no net force in any direction,
when averaged over a period of time. A molecule in the surface of a liquid,
however, solely because of its position, will have a net force acting upon
it in toward the body of the liquid, as shown in Figure 3.3. In a freely

Figure 3.3.

Forces Acting on Molecules
in a Liquid.

“A” represents a surface molecule
and ‘ d a molecule jn the body of
the liquid. The other circles repre-
sent neighboring molecules.

There are no molecules above a sur-
face molecule to counterbalance the
attractive forces due to molecules be-
low it.

On surface molecules there is a re-
sultant force acting m toward* the
body of the liquid.

An interior molecule is surrounded
on all sides by other molecules and
there is no net force acting upon it.

The State* of Matte ? Solids and Liquids

falling droplet of liquid, this net force toward the center causes the drop
to assume a spherical shape. For a given inass of a substance, a sphere
has the minimum surface area or ratio of surface to volume.

One of the properties of liquids ansing from the imbalance of force
upon surface molecules is surface tensum . The surface of a liquid appears
to be under a tension and acts as though it were a stretched membrane,
Water has a relatively high surface tension, strong enough to support a
steel needle or the weight of certain insects. The surface tension of a
liquid can be decreased by dissolving therein a surface active material
such as soap or a detergent. This will also increase the wetting power of
the liquid. It is surface tension which causes certain liquids to rise in a
capillary tube. Indeed, surface tension is measuied quuntitatnelv by the
height of the rise of a liquid of known density in a capillaiv tube or known
diameter. Surface tension measurements are of great importance in bio-
logical work. The technique of measuring surface tension by determining
the force required to pull a platinum wire ring through a liquid surface
was first proposed by the biologist, Lecornte DuXouy, An increase m tem-
perature decreases the surface tension and theoretically the surface tension
should become zero when the critical temperature is reached. No surface
is then present and liquid and gas merge without any apparent meniscus
between them. Solids may also be construed to base sm face tensions, which
are difficult to measure and are very large, hundreds of times greater than
those of liquids. (Table 3-B)

Table 3-B

Surface Tension and Viscosity

Substance

Melting
Point ,
°C

Boiling
Point ,

°C

Surface Tension ,
dxjne/cm
at 20'C at 50 C

Viscosity,
etnftpatse
! at 20 C : at 50°C

Water

0.0

100.0

72.8 67.9

i 005

0.656

Benzene

5.5

80.1

2S.9 23.0

0 63 }

0.498

Ethyl Alcohol

-117.3

78.5

22.3 19 8

1 .200

0.834

Diethyl Ether

-116.3

34.6

17.0

0.233

Mercury

- 38.9

356.9

484 470

1 ,53

1.39

Glycerol

290

850

Air

0.0181

0.0195

Wax

4 x 10"

6. Viscosity. Viscosity concerns the bulk flow of a fluid material Gases
have relatively low viscosities, 'while solids can be said to have extremely
high viscosities. The wide range in viscosities in liquids is due to intermole*
cular forces which vary from liquid to liquid and about which not too much
is known. Viscosity can be visualized as a frictional drag as one layer of a
fluid flows over a neighboring layer. Molecules “jumping 1 ' from one layer

The States of Matter: Solids and Liquids

to another give rise to a viscous drag. An increase in temperature increases
the viscosity of a gas but generally decreases that of a liquid. In rare cases,
notably that of liquid sulfur, the viscosity increases with temperature to a
maximum and then falls. Viscosity can be measured by a variety of means
of which the most common is the rate of flow through a cylindrical tube.
Motor oils are classified according to their viscosities. An SAE 20 (Society
of Automotive Engineers) oil refers to the length of time it takes a fixed
volume of oil to pass through a calibrated opening. The higher the number,
that is, the longer the time of flow, the greater is the viscosity. In cold
weather we might use the less viscous SAE 10 oil in an engine while dur-
ing the summer the more viscous SAE 30 oil would be required.

7. Vapor Pressure. Many of the viewpoints of the Kinetic-Molecular
Theory developed for gases are also applicable to the liquid state. Within
the body of a liquid, the liquid molecules, though very close to each other,
move about in all directions, colliding with and rebounding from each other,
the walls of the containing vessel, and the surface of the liquid. Because
of the inward force acting on surface molecules, the surface acts as an
energy barrier, in effect repelling molecules back into the body of the liquid.
But, as with gases, there exists a distribution of kinetic energies among the
liquid molecules (page 24). Of those that come to the surface, some possess
sufficient kinetic energy to break through the surface barrier, thereby be-
coming gaseous molecules, more properly vapor phase molecules, in the
space above the liquid. This breaking away from the liquid, which occurs
only at the liquid surface, is called evaporation.

The rate of evaporation, that, is, the number of molecules which escape
from the liquid surface in a given time, depends on A) the area of the liquid
surface, B) the temperature, which determines the fraction, of the liquid
molecules which have sufficient energy to escape, arid C) the extent to
which the gaseous molecules move away from the vicinity of the liquid
surface. Continued evaporation will, of course, cause complete conversion
of the liquid into vapor and, in time, the liquid will dry up. (Figure 3.4a.)

Let us now consider a closed container so that the resulting gaseous
molecules cannot escape into free space but are confined to a relatively
small volume above the surface of the liquid (Figure 3.4b). Under these
circumstances, it is entirely possible that some of the gaseous molecules, in
their own haphazard meanderings, will plunge back into the body of the
liquid, or condense. When the liquid is first placed in the container and
the container is covered, there are only a few gaseous molecules in the
space above the liquid, so that the rate of return to the liquid is smaller
than the rate at which molecules escape from the liquid. In the course of
time, the number of gaseous molecules increases and hence the rate of re-
turn to the liquid also increases. Finally a condition is reached in which
the number of molecules leaving the liquid in any given period of time equals
the number returning in that same interval, hat is, the rate of evaporation
equals the rate of condensation. There will then exist a state of equilibrium,
and the concentration of molecules in the gaseous state will remain un-
changed. If, under these conditions of equilibrium, the pressure of the vapor
were to be measured, the value obtained would be called the vapor pressure

Solids

d liquids

the Stati's of MiJtir

of the liquid (Figure 3,4c). The vapor pressurr is not jmt am \alut* of the
pressure of the vapor but is the value which obtains onlv umK*r conditions
of equilibrium between the liquid and the \ apoi *

From an open container, a In a closed container, before W hen e«|iii!ibrmin is estab-

liquid will evaporate com- equilibrium is reached, the hshed the manometer read-

pletely and zero pressure is manometer reading will mg Iwcomes Miustaut The

read on the manometer. continue to rise. pressure lead on it equals

the vapoi piessnrr* nf the
liquid

The small circles represent molecules and the small arrows then dtmtions of mo-
tion. The heavy arrows indicate the relative Kites of evaporation and toiulensa-

tion. The height, h, which is the difference m levels of the liquid iinmnni in

the two arms of the manometer, is a measure of the pre^suie of the >\sten» to

which the manometer is connected, m this case, the pressure of tin* sapor

Figure 3.4. Vapor Pressure and Equilibrium.

It should be emphasized that, under the conditions of equilibrium, there
is no cessation of activity but both processes, vaporization and condensation
—opposing in a sense— proceed at the same rate, and hence there is no further
change in measurable properties. This concept of dynamtr rquihhrturtu as
distinguished from a static equilibrium, is of extreme importance in physical
science and will be remarked time and time again. To express an equilibrium
such as that between liquid water and water vapor, we write

Water (liquid) Water (vapor)

where the arrow to the right indicates the process of vaporization of
liquid water and the arrow to the left <— indicates the process of conden-
sation. The double arrow thus represents the dual processes of vaporiza-
tion and condensation proceeding simultaneously in opposite directions.

The average kinetic energy of liquid molecules also varies with tem-
perature. With an increase in temperature, molecular velocities increase
and so also the number of molecules which have the energy required to
break through the surface. Hence the rate of evaporation increases/ result-
ing in a higher concentration of gaseous molecules and a higher vapor
pressure at the higher temperature. The variation in vapor pressure with
temperature for four liquids-water, ethyl alcohol, diethvl ether, and mer-
cury-is shown graphically in Figure 3.5, and data for the vapor pressure
of liquid water as a function of temperature are given in Appendix III.

The States of Matter: Solids and Liquids

Liquids which have appreciable vapor pressures are said to be volatile.
Those that have extremely low vapor pressures, such as mercury, are often
called nonvolatile. At any given temperature, the liquid with the higher
vapor pressure is the more volatile; for example, ether is more volatile
than water.

If air, or any gas, contains a vapor at a pressure equal to the latters vapor
pressure, it is said to be saturated with that vapor. Thus, on a day when
the temperature is 25 r C and the vapor pressure of the water in the atmos-
phere is 23.8 mm, the air is saturated with water vapor. If we attempted
to inject more water vapor into the air and thus to increase the water
vapor pressure at the same temperature, 25 °C, condensation of the excess
water vapor above that required to maintain the vapor pressure of 23.8 mm
would occur. The pressure of the vapor cannot exceed the equilibrium or
maximum possible vapor pressure.

The vapor pressure of a volatile liquid increases with a rise in tempera-
ture and, at a temperature at which its vapor pressure equals atmospheric
pressure, or the external pressure impressed upon it, the liquid boils. This
temperature is the boiling point. The normal boiling point is the tempera-
ture at which a liquid has a vapor pressure equal to normal atmospheric
pressure, or 760 mm. In Figure 3.5 a horizontal line has been drawn at a
pressure of 760 mm. The temperature at which this line intersects the
vapor pressure curve of a given liquid is the normal boiling point of that
liquid. For water the normal boiling point is 100°C. Unlike the critical
temperature which is an invariant property of a substance dependent upon
its molecular properties only, the nature of the boiling point is relatively
trivial in that it is simply the temperature at which the vapor pressure of a
liquid equals the ambient pressure. Only then can bubbles of vapor survive
in the liquid and produce the mechanical action of boiling. Indeed the
curves in Figure 3.5 also represent the boiling points of the various liquids
under different pressures. By reducing the pressure upon it to less than
760 mm, water can be made to boil at a temperature below 100 °C. On
the top of Mont Blanc, 15,779 feet above sea* level, where the average baro-
metric pressure is about 420 mm, water boils at a temperature slightly over
84° C. On the other hand, if the external pressure on water is increased
above 760 mm, the boiling point exceeds 100° C. In steam boilers where the
pressure is 11,659 mm, water boils at 200 °C. The home pressure cooker
operates on this principle to produce temperatures above the normal boiling
point of water. In organic chemistry, in particular, many liquid substances
decompose if heated in an attempt to boil them at atmospheric pressure.
Such a liquid can be distilled under reduced pressure. Thereby its boiling
point can be reduced to a temperature at which decomposition will not
occur and the liquid can be purified by 4 Vacuum distillation.”

A volatile liquid gives off vapor at all temperatures. The difference
between evaporation and boiling is that evaporation takes place at the sur-
face only. Boiling, more correctly termed ebullition, is the formation of
water vapor in the form of bubbles throughout the entire liquid. These
bubbles rise to the surface and burst.

During boiling, a change of phase from the liquid to the vapor state,
energy must be supplied to vaporize the liquid. The quantity of heat energy

The States of Matter Soltth and Lujuids

required to vaporize one gram of a liquid at its boiling point is called the
heat of vaporization. For water at 100" C, the heat of \upon/ution is 540
calories per gram. The heat of vaporization is not constant but varies with
temperature, decreasing with an increase in temperature and becoming zero
at die critical temperature.

The vapor pressures of different liquids differ markedly even at the same tem-
perature. A vertical line gives the vapor pressures of the four liquids at a specific
temperature. Thus, at 25°C the vapor pressures of water, alcohol, and ether are
23.8 mm, 59.0 mm, and 537 mm, respectively; the vapor pressure of mercury',
too small to be read from the graph, is 0.002 mm.

Figure 3.5. Vapor Pressure of Liquids as a Function of Temperature.

Solids may also have vapor pressures. The process whereby a solid
gives off vapor molecules directly without the intermediate formation of
liquid is known as sublimation. Generally the vapor pressures of solids are
low but iodine and naphthalene are examples of solids which have appreci-
able vapor pressures at ordinary temperatures. As with liquids, the vapor
pressure of a solid increases with temperature. At 0°C ice has a vapor
pressure of 4.6 mm; at — 5°C the vapor pressure of ice is 3.0 mm ( Appendix
IV). Even if the surrounding temperature never rose above the melting
point so that liquid was never formed, ice would evaporate and "dry up,”
given sufficient time. The melting point can now be defined more pre-
cisely as the temperature at which solid and liquid have the same value of
vapor pressure (Figure 3.6). At this temperature both phases will be in
equilibrium and can exist alongside each other indefinitely. This equilibrium
can be represented as

Solid Liquid

The States of Matter: Solids and Liquids

At the melting point there is a dynamic equilibrium and two processes are
going on at the same rate-the melting of solid to form liquid and the freez-
ing of liquid to form solid.

Vapor

Pressure

Line AB represents the variation of the vapor pressure of a solid with temperature;
line BC represents the variation of the vapor pressure of a liquid with temperature,
and is analogous to the curves of Figure 3.5. The temperature, D, corresponding
to the intersection, B, of the two vapor pressure curves is the freezing point. Tem-
perature E, at which the liquid has a vapor pressure of one atmosphere, is the
boiling point.

Figure 3.6. The Relation of Freezing Point and Boiling Point to Vapor Pressure.

Many liquids can be supercooled or superheated, that is, cooled or
heated beyond temperatures at which equilibrium transition to solid or
vapor, respectively, should have occurred but did not. Very pure water has
been cooled many degrees below 0°C without the formation of ice. If the
tendency to crystallize is small, usually due to a high viscosity, a liquid can
be supercooled to a rather low temperature. It is this phenomenon which
accounts for the formation of vitreous or glassy substances. Superheating
or supercooling of a liquid is common but superheating of a solid above
its equilibrium melting point is practically unknown. A supercooled sub-
stance is not at equilibrium. The introduction of even the slightest quantity
of that phase which is stable under the supercooled conditions enables
equilibrium to be established, and the supercooled phase is transformed
immediately to the stable form.

8. Changes of State. For a pure substance, bqtl^fpelting and boiling
(and their counterparts— freezing and condensation ’ £he boiling point)
take place at constant temperatures. Addition o* ~ & at temperatures be-

in/ S'.,**

* M J

fluids

«.»ld

tween those at which a change oi 'date occurs Minch uu teases *h*. kinetic
energies of the particles concerned and hcike the bmpnntme nt the sub-
stance without causing a change nt vtate H .t M*ini is heated no change of

state occurs till the melting point is reached \t this fempnutun tun states,

the solid and the liquid, can coexist and transition imm ih« solid to the
liquid takes place at a constant temperature Kvni tVueh heat is being
added (the heat of fusion 1 just so long as some solid is pi«*srw the tem-
perature will not rise but will remain constant Not until all the solid

has been converted to liquid, does the teinpeiahne aeam begin to use.
With a continued increase 1 in tempeiature, the liquid naitules have ever-
increasing kinetic energies and an increasing number gam suHm lent energy
to escape the surface barrier and therein create luuhri values of vapor
pressure. At the boiling point the transformation inmi liquid to vapor takes
place rapidly, with ebullition. Here, too, the tempt latino icnuins constant
so long as some liquid remains. Continued heating above tin boding point
results solely in a greater kinetic eneigv oi the gaseous mulct ules No
further change of state can occur. The heating process desiphed is illus-
trated graphically in Figure 3,7, such a graph of temperature against time
is known as a heating curve

Time

The graph represents the heating ot a substance from below its melting point to
above its boiling point. Along line AB the solid is being heated, only one state,
the solid, is present. At point B, the solid begins to melt; the temperature cm re-
sponding to point B is the melting point. From B to C. two states, the solid ami
liquid, are present, and the temperature remains constant. At point ( ' all the solid
has been converted to liquid and then the ternperatuie rises along the line CD.
Only the liquid state is present from C to D. Point D is the hod mg point The
temperature remains constant from D to E as the liquid bmls, along line DK two
states, liquid and gas (vapor) are present. At point E all the liquid has been trans-
formed to gas and from this point on only the gaseous state exists. The length
of the segments in the heating curve depends upon the quantity of material being
heated and the rate of heating. The reverse process of cooling a substance from
me gaseous to the solid state would give a graph which is the mirror image of
Figure 3.7 and which is known as a cooling curve .

Figure as Heating Curve for a Pure Substance.

The States of Matter; Solids and Liquids

9. Diffusion in Liquids and Solids. Diffusion which takes place rela-
tively rapidly in gaseous substances proceeds very slowly in liquids. The
proximity of liquid particles and the absence of any appreciable free space
between them makes for a small free- path of travel. It may take years for
a dye placed at the bottom of a vessel containing a liquid to diffuse uni-
formly throughout the liquid. This is why mechanical mixing is so essential
for uniform dispersion in liquid systems. Even in. solids some diffusion is
observed. If solid gold and lead are placed in contact, in time each metal
will be found to have migrated into the other to a slight extent. No general
equation is known for the diffusion of solids or liquids.

QUESTIONS

1. What are the three classes of solids? What are the characteristics of each
class?

2. Define and illustrate die following: polymorphism; allotrope; transition tem-
perature; melting point, boiling point.

3. Define crystal lattice and unit cell. In what manner are the two related?

4. Is it possible for two allotropic forms of a substance to exist together? Explain.

5. The coefficient of thermal expansion of solid iron is 3.5 X 10' 5 deg 1 ; the
coefficient of compressibility of liquid water is 5 x 10* 5 atm -1 . Compare these
values with the corresponding values for an ideal gas.

6. In terms of the kinetic-molecular theory discuss the following: vapor pressure;
evaporation; melting; freezing; evaporation; condensation, sublimation.

7. Compare the relative binding forces in a crystal which melts and in one
which sublimes.

8. For a volatile substance the melting point is the temperature at which solid
and liquid have the same value of vapor pressure. Can only solids which have
vapor pressures melt?

9. Discuss and illustrate the relationship between ( a ) vapor pressure and boiling
point; (b) vapor pressure and melting point.

10. What would be the boiling point of ether and ethyl alcohol under 300 mm
pressure?

11. What is surface tension? Why do droplets of a liquid assume spherical shape?
What is the surface tension of a liquid at its critical temperature? Why
should surface tension decrease with temperature?

12. What is viscosity? Why should the viscosity of a liquid decrease with tem-
perature?

13. Using the concept of equilibrium explain (a) why ice will melt in an ice-water
mixture if the temperature is raised above 0°C; (b) why liquid water will boil.

14. In Figure 3.6 what would be the meaning of extrapolating the solid and
liquid vapor pressure curves beyond point B?

15. Wbat would happen to supercooled water at -1°C if a crystal if ice were
added to it at that constant temperature? Expl$ ; n your answer

16. On one set of axes draw a heating curve and a celling curve fo r Jiyl alcohol.
Label all significant points. (Use Table 3-B)

17. What is the vapor pressure of water atop Mont Blanc, where the temperature
is 5°C?

fhc Stains of Matter Solids and Liquids

18. Distinguish between “state” and “phase.” Can there he moie than one gase-
ous phase? Give an example of two phases in the solid state.

19. Explain why the temperature remains constant timing a change ot state or
of phase.

20. One liter of oxygen is collected over water at 27 C and a pressure 4 (atmos-
pheric) of 786 mm. What would be the volume of the drv gas at STP?

21. (a) When 1000 g of water freeze at 0°C how mam calones ate liberated?
(b) If the specific heat of iron is 0.115 calone/gram degree, to what temjkTa-
ture could 10 g of iron be raised from an initial temperatuie of 0 C y

Chemical Change

To the alchemist chemical change whereby one substance is transformed
into another must have seemed magical and mysterious beyond comprehension.
Even to the modern scientist it is not without its sense of fascination and
wonder. Many chemical changes such as the explosion of gunpowder proceed
in an overtly striking fashion with the emission of heat and light; others proceed
about us without our consciously being aware of them, such as the rusting
of iron or the conversion of the food we eat into the energy to keep us going
about our daily activities. Because of its relative simplicity we shall consider
specifically the rusting of iron in some detail, though the conclusions we shall
derive will be applicable to all chemical changes.

1. A Chemical Change— The Rusting of Iron. In the elemental state, iron
has the characteristic properties which we ordinarily associate with metals.
For example, it has a metallic luster and it is a good conductor of heat, among
other physical properties which the reader can readily bring to mind. If iron
is exposed to the atmosphere, however, it crumbles away in the form of a
red powder, which has its own distinct set of properties quite different from
those of the original iron. This type of change is known as a chemical change
or a chemical reaction. In rusting, iron combines chemically with the oxygen
of the atmosphere to form a new chemical species, the red iron oxide. The
reacting substances in a chemical reaction are known technically as the
reactants and the substances formed as the products. A chemical equation can
be written to represent the rusting of iron as follows:

Iron plus Oxygen yield Iron Oxide
or Iron + Oxygen Iron Oxide

The iron oxide in this equation is the red oxide of iron known as rust. It
will be seen later that there are three possmle binary compounds of iron and
oxygen, or iron oxides. Whereas the rusting of iron proceeds spontaneously, time is
an important factor in all chemical changes. Rates of chemical reactions, the
study of which is called kinetics (Chapter 8) are dependent upon a number of
factors and vary widely. Under ordinary conditions the rusting of iron proceeds
rather slowlv.

< ‘It, 'n:< ( '?W!iRt-

Since the earliest days of worded lustc »i> mam thcmica! nvulmn s have
been known and utilized for practical puipos^ hut wluh ehenuxUy re-
mained a non-quantitative science, little thoou*tie.d j»iu»uv, uas made to-
wards their understanding. Indeed, modem ehenmin. wl.uh dates iiom the
development of quantitatixe technique, is reiutweK umuvi ha\i:ig its be-
ginnings in the late eighteenth eentuiy. It was dining those stars that
meaningful experimental data concerning chemical ion turns wen fust ob-
tained and generalizations formulated which led to Daltons \tonue Theory.

2. Law of Conservation of Mass. The first of these genrtuh/ntions was
the recognition that, in a chemical change the total mass of reactants and
products remains unchanged in a closed system The last pin use must he
emphasized because, in the event that a gas, geneialh invisible, is one
of the products and escapes the reaction container, tin* mass of material
that the experimenter would weigh at the completion of a chemical icaction
would be less than the muss initial!) Before it was understood that gases
had weight, the escape of a gas from an open container during a chemical
reaction led to results which weie interpieted eiioneoush In the lusting
of iron, the iron oxide (rust') formed weighs more than did the initial iron
because of the latter s combination with the gas, owgen H the mass of
the oxygen is included, however, the sum of the masses <4 all chemical
species before and after the reaction is a constant \alue Ibis experimental
fact is now known as the Law of Comm tit ion rt Mass Though today the
content of this statement appears trite and all but apparent to even the
uninitiated in science, there is really no a priori reason whv mass should be
conserved in a chemical reaction, or conversely, then* is no reason why
mass should not be gained or lost in a chemical change, so we should not
be disconcerted at the inability of the alchemists to arme at this principle.

3. Law of Definite Proportions. If we were to measure caret ullv the

weights of iron and oxygen which actually combine to form iron oxide,
or conversely to analyze the iron oxide produced for the weights of iron
and oxygen composing it, we would find that in as many samples of iron
oxide as we took for analysis, the ratio of the weight of iron to the weight
of oxygen with which it is combined is a constant value In other words,
the iron oxide has a definite composition by weight. Thus Hi. 6 grams
of iron will always combine completely with exactly 48.0 grams of oxy-
gen to produce exactly 159.6 grams of iron oxide, and theiefme the
fraction of iron in iron oxide is equal to 111.6/(111.6 48.0 1 0.700,

or 70.0%. Similarly the oxygen percentage in iron oxide must be 30.0%.
It should not be inferred that the two numbers 1IL7 and 48. 0 arc the only
weights of iron and oxygen which will combine. Am other arlntrun weights
of iron and oxygen initially chosen would have combined in the same ratio
of 111.6 to 48.0, still resulting, ; u- an iron oxide of the same composition as
the foregoing. Thus 111.6/4^0 or 232 grams of iron would combine with
1.00 grams of oxygen t< produce 3.32 grams of iron oxide. If either the
iron or the oxygen is present in excess of the 111.6 to 48.0 ratio, that excess
would have remained uncombined and the ratio by weight of the iron and
oxygen which did combine would be in the proportion of 111.6 to 48.0.

Chemical Change

These experimental facts can be summarized as tlW Law of
Proportions , which states that when two (or more) substances combine to
form a specific compound they combine in a fixed and definite ratio by weight.
This leads to the conclusion that a pure compound, independent of its source,
has a definite composition by weight. 1

4. Law of Multiple Proportions. Experimental investigation resulted in
another unusual fact. The same two elements may form two or more com-
pounds. Indeed, iron and oxygen form three different binary compounds,
each having its own separate properties. For any one of these compounds
the weight ratio is fixed (Definite Proportions) but these ratios will be
different in the several compounds. Let us consider two of the iron-oxygen
compounds for which the weight ratios are given below.

Iron-Oxygen Weight Ratio

Iron (III) Oxide (ferric oxide; rust) 111.6 : 48.00 or 2.325 : 1

Iron (II) Oxide (ferrous oxide) 167.4 : 48.00 or 3.488 : 1

The weights of iron in the two compounds which combine with unit
weight (such as one gram) of oxygen aie in the ratio of 3.488 to 2.325,
which is exactly three to two within the limits of experimental error. If
the weight of oxygen chosen as index had been some number other than
one, the weights of iron would have changed proportionately but their ratio
would still be three to two. Similarly if the iron weight had been fixed,
the ratio of the oxygen weights in the two compounds would have been
two (in Compound B) to three (in Compound A).

That a simple whole number ratio should so result is striking. Investiga-
tion of other groups of compounds has led to similar results. Hydrogen and
oxygen thus combine to form two different compounds, water and hydrogen
peroxide; carbon and oxygen form two common oxides 2 ; mercury and chlorine
form two chlorides; and there are five distinct compounds of nitrogen and
oxygen, all of which indicate this relationship as shown in Table 4-A.

These facts can be summarized in a general statement known as the
Law of Multiple Proportions : If two elements unite to form more than
one compound, a small whole number ratio exists between the weights of
one of the elements which combine with a fixed weight of the other element.

5. Dalton's Atomic Theory. The Laws of Conservation of Mass, Definite
Proportions, and Multiple Proportions summarized the experimental facts
concerning chemical change which were known to chemists in the first years
of the nineteenth century. Nothing of a theoretical nature has been discussed
so far, but underlying such broad generalizations there must exist some
orderly state of affairs which gives rise to the simple relations noted. In
the history and development of science such a status of knowledge arises
again and again and invites a theory concerning fundamental phenomena.

1 In Chapter 32 we shall see that two elements can form different compounds having
identical weight compositions. This effect is due to different arrangements of the atoms
within the molecules, a phenomenon known as isomerism.

2 In addition to carbon monoxide and carbon dioxide, carbon forms two other oxides:
carbon suboxide and nvellitic anhydride.

Chrtnu t J Change

Table 4- A

Laws of Mujliiflk Phoi*ohiio\s

Substance

Composition

By Weight

Hydrogen

(ht/gni "

I. Water

1.008

8 00

Hydrogen Peroxide

l.OOS

10 00 1 2 • S . 00 )

Carbon

( hifgen

II. Carbon Monoxide

12.00

10 00

Carbon Dioxide

12.00

02 00 1 2 ■ 10 . 00 )

Chlorine !

Mercury

III. Mercuric Chloride

35.45

100 3

Mercurous Chloiide

33.45

200.6 \2 * 100 . 3 )

Sitiogen

Oxygen

IV. Nitrous Oxide

14.01

S 00

Nitric Oxide

14.01

16 00 (2 * 8 00 )

Nitrogen Trioxide

14.01

2 1 OO (3 * 8 . 00 )

Nitrogen Tetroxide

14.01

32 00 M • 8 . 00 )

Nitrogen Pentoxide

14.01

40.00 (5 < 8 . 00 )

The man who proposed such a theory concerning the fundamental nature
of chemical species was John Dalton, an English schoolmaster. In oidrr to
explain chemical behavior, in 1808 Dalton proposed simple 1 postulates which
are known collectively as Dalton’s Atomic Theory. Stated from a modern
viewpoint they are:

A) Elements (and compounds) are composed ultimately of minute, discrete,
indivisible particles called atoms.

B) All atoms of the same element are identical in all pioperties, physical
and chemical. These properties arc different from those of atoms of
other elements.

C) Atoms are the units of chemical change.

D) Compounds are formed by the union of integral numbers of atoms of
different elements in simple numerical proportions as 1:1. 1:2, 1.3, 2:3, etc.
To Dalton chemical reaction was merely the interchange of atoms of

the reacting substances. From this viewpoint an element can be defined
as a substance composed of atoms of only one type whereas in a compound
different kinds of atoms are combined.

In Chapter 11 it will be seen that, in the light of modern mvlv.it theory,
Dalton s postulates are not rigorously true but insofar as the fundamental view-
pomt is concerned, Daltons Atomic Theory of an indivisible billiard ball-hke atom
H V V j. a ?. 7 ^ or calculations in chemical stoichiometrv . Dalton assumed
the indivisibility of atoms and a single definite mass for atoms of the same
element. He, of course, made no^ inference concerning the internal structure of
themselves. Actually atoms can decompose to simpler particles (radio-
' 1 ^ /. k a P^ r 48) and all atoms of the same element do not have identical

asses (iso opes) (Chapter II). In ordinary chemical change, however, atomic
ElT ” ltlon *£ es not occur ; Though individual masses of atoms of the same
n may iffer, m even the minutest quantity of matter undergoing chemical

Chemical Change

change the number of atoms involved is enormous and they have been so
thoroughly mixed that any collection of them will have statistically the same
average atomic weight. It is with this average atomic weight that the chemist
is concerned. Though our ideas concerning the nature of atoms, particularly their
internal structure, have become more sophisticated, our faith in the basic truth
of the atomic theory, at least insofar as chemical change is concerned, is not
diminished.

6. Explanation of the Laws of Chemical Change. Conservation of Mass :
In a chemical change, though the atoms interchange partners, their nature
and number before and after the reaction remains unchanged. Inasmuch
as the atoms are indivisible and the mass of each atom is invariant, the
total mass before and after chemical reaction must be constant. 3 Thus the
conservation of mass is more truly the conservation of atoms. Later we shall
have to satisfy this requirement by an arithmetic operation known as balancing
a chemical equation, that is, to make the numbers of each kind of atom
on each side of a chemical equation (before and after reaction) the same.

Definite Proportions: Because a given compound consists of specific num-
bers of atoms each of which has a definite weight, it must inevitably have
a definite composition by weight. The iron oxide, rust, is a chemical com-
pound which always is composed of two atoms of iron combined with three
atoms of oxygen. Inasmuch as an atom of iron has a definite weight, as does
the oxygen atom its own fixed weight, the composition by weight of iron
oxide will always be the same, that is, the relative weight of iron to oxygen
in rust will always be equal to the ratio of the weight of two iron atoms
to that of three oxygen atoms.

Multiple Proportions : Where two or more compounds are formed by the
same elements, fixing the weight of one of the elements establishes a definite
number of atoms of that element. In each compound with that number of
atoms, an integral number of atoms of the other element must combine. This
integral number will be different for the several compounds, but the ratio of
these integral numbers of atoms will itself be a small whole number.

Thus, in the two compounds of hydrogen and oxygen, water and hydro-
gen peroxide, the number of atoms of oxygen combining with two atoms
of hydrogen (a fixed weight) is one in the case of water and two in the
case of hydrogen peroxide. Their ratio, and so also the ratio of their
weights, must be two to one.

An insight of the structures of the two molecules will clarify how such
a situation is possible. In the structures following, H represents a hydrogen
atom, the symbol O, an oxygen atom, and a dash the bond between two atoms.

H-O-H H-O-O-H

water hydrogen peroxide

3 Because mass and energy are interconvertible according to Einstein's equation, E = me 2
(page 630), strictly speaking mass is not conserved in a chemical reaction where energy is
liberated or absorbed. The quantity of mass corresponding to the energy involved in an
ordinary chemical reaction is so slight, however, that the deviation from conservation of
mass is beyond the range of detection of present analytical balances. For all practical
purposes mass can still be considered to be conserved.

Chemical Change

In the two iron oxide compounds, the number of atoms of oxygen combin-
ing with two atoms of iron is two in one compound ami three in the other.

QUESTIONS

1. State and illustrate the Law of Conservation ol Mass the Law of Definite
Proportions in die light of Dalton's Atomic Theoiv

2. List the postulates of Dalton’s Atomic' Theorv .

3. Explain the Laws of Conseivation of Mass, Definite Piopoiimns. and Multiple
Proportions in the light of Dalton’s Atomic Theoiv.

4. Explain the Law of Definite Proportions using as an example the compound
A 2 B. The atomic weight ol A is 23 and the atomic weight of B is 32

5. Show how the following compounds dlustiate the Law of Multiple Piopor-
tions: AIL and AJB.

8. Draw reasonable structural formulas for the two oxides ol non which show
\vh> these compounds illustute the Law of Multiple Pi opm turns

7. The percent by weight of iron in iron oxide trusD is 69.0% < ,i ' What weight

of iron is in 316 grams of iron oxide'* {b) What weight ol oxsgen will com-
bine with 100 grams of iron to form non oxidt ‘ J

8. In the two iron oxides on page 51, what i> the ratio ot die vu.’hts ol owgen

that will combine with 55.85 grams of iron '*

9. Chlorine and oxygen form two different compounds. In compound A the
percent of chlorine by weight is 81.6% and in compound B the peicent by
weight of chlorine is 59.7%. Show how these data dlustiate the Law of
Multiple Proportions.

10. There aie two oxides of carbon. In one giam of one ovule there is .V7 giam

of carbon and in one gram of the other oxide there is 3/11 gram of carbon.

Show how these quantities illustrate the Law of Multiple Propoi turns

11. When 30 g of a compound C are decomposed, 10 g of an element A and
20 g of an element B are formed. If 15 g of A are mixed with 50 g of B
what will be the weight of C formed and what will be the total weight
of the final mixture?

12. Ammonia contains 17.76% hydrogen and 82.24% nitiogen. When 3.77 g
of hydrogen react completely with 26.23 g of nitrogen, 30 (K) g of hydra/.ine
are formed. Show how these data illustrate the Law of Multiple Proportions.

Atomic and Molecular Weights

To make further use of the atomic theory, we must know the relative
weights of the atoms. It is simple to determine experimentally the weights
of elements which will combine with each other. If not only the weights of
two elements but also the number of atoms of one element which combines
with one atom of the other element were known, the relative weights of
the atoms of the elements could then be determined. Hydrogen and oxygen
combine to form water in the ratio of 1.008 to 8.000 by weight. If a water
molecule contains one atom of hydrogen and one atom of oxygen, then
these weights also express the relative weights of the two kinds of atoms.
If a molecule of water contains two atoms of hydrogen and one atom of
oxygen, however, the relative weights of the hydrogen and oxygen atoms
will be 0.504 and 8.000, or 1.008 and 16.000. The combining weight of an
element can be defined as that weight of an element which will combine
with 8.000 parts by weight of oxygen or 1.008 parts by weight of hydrogen. 1
The combining weight is also called the equivalent weight of an element
(Page 77). The combining weights do not represent the relative weights
of the atoms of the elements, but there is a relationship between the com-
bining weight and the atomic weight.

1. Gay-Lussac’s Law of Combining Volumes. The problem of finding
the relative weights of atoms remained unsolved for fifty-two years after
the formulation of Dalton’s Atomic Theory. In 1808 the first clue to the
solution of the problem was proposed by a French professor of chemistry,
Joseph Louis Gay-Lussac. Gay-Lussac discovered that when gaseous sub-
stances interact their volumes (both reactants and products, if gases)
are in the ratio of small whole numbers. This is known as the Law of
Combining Volumes; it is assumed that the temperature and pressure at
which the gaseous volumes are measured are the same before and after
reaction.

lAs we shall see later, the combining weight is more accurately that weight which
combines with three parts by weight of the carbon-twelve isotope.

W. nt tu

< ‘tit! fttltir \Y rights,

Two volumes of hydrogen gas thus will
gas to produce two volumes of watei as
above 10OC. The relathe xolutnes can be

1 Vol

1 Vol.

1 Vol

Hydrogen O \ ygt ’»

ieact with one \ohune oi oxsgen
steam, when the t*‘iiipeiatuic is
pictuiod as billow*?

— [1 ~va. | [_i y< >1 j

Strain

Whatever the actual volumes of the gases mu\ he, the mtio of the volumes
of hydrogen to oxygen to steam is 2 to 1 to 2 V gum \oiume of hydrogen
thus will combine with half its volume of o\\ gen and produce a \olume of
steam equal to that of the initial hydiogen Foi example 71 ml of hvdrngen
will react completely with 37 ml of oxygen to pioduce 71 ml of steam. If
100 ml of hydrogen were used, the excess Indrogen, 20 ml, would icmain
unreactcd. It should be noted that there is no law* of comenahon of volume.

The 2:1:2 ratio above is not universal. Each individual gaseous reaction
has its own specific integral volume ratio. YWm*u chloiinc and hydiogen
gases combine to form hydrogen chloride gas, the volume i elation is 1 volume
of chlorine to 1 volume of hydrogen to 2 volumes of hxdrogen chloride gas.

1 Vol.

-f-

1 Vol. !

1 Vol. 1 I

... ......... M 1

1 Vol. j

Chlorine

Hydrogen

Hydrogi n '

Chi unde

When nitrogen and hydrogen react to form ammonia gas, the volume mtio
is 1 to 3 to 2.

1 Vol.

1 +

1 Vol.

-*j 1 Vol. | | 1 Vol.

Nitrogen

II ydrogen

Amm*in\a

2. Avogadro’s Hypothesis. To explain this striking relationship between
the volumes of gases, Amadeo Avogadro, an Italian professor of phvsies in
1811 proposed the hypothesis that equal volumes of gases under identical
conditions of temperature and pressure contain the same numbei of molecules.
In effect, this hypothesis sets up a proportionality between the volume of
any gaseous substance and the number of molecules therein. Under the
same conditions of temperature and piessure, two liters of hydrogen would
contain twice the number of molecules as would one liter of oxygen and
a number equal to that in two liters of steam. Since the ratio of combining
volumes is identical with the ratio of combining molecules and siftee the
latter ratio is a small whole number so also is the ratio of the combining
volumes.

Several times we have stated that oxygen molecules wen* diatomic and
we now have sufficient information to prove logically that this is so. Since
two volumes of hydrogen react with one volume of oxygen to produce two
volumes of steam, it can be inferred from Gav-Lussac’s Law and Avogadro’s
hypothesis that two molecules of hydrogen react with one molecule of oxy-
gen to produce two molecules of steam. Concerning the numbers of atoms
which constitute each molecule we have no information but, at the very
least, each molecule of water produced must contain at least one oxygen
atom. Let us assume that this is the case. Then, in the two water mole-
cules produced there is a total of two oxygen atoms, one for each mole-
cule. These two oxygen atoms could have come only from the one oxy-
gen molecule which reacted. Hence this one molecule must contain two
oxygen atoms and the oxygen molecule is diatomic. More critically this proof

Atomic and Molecular Weights

indicates only that the number of oxygen atoms in an oxygen molecule must
be even, and the minimum number is two. It is theoretically possible that
a water molecule contains more than one oxygen atom; if the number were
two, then the oxygen molecule would contain four atoms. With the reaction
just considered it cannot be shown that the hydrogen molecule is diatomic
but similar reasoning for the reaction between hydrogen and chlorine would
prove that the hydrogen molecule cannot have an odd number of atoms.
From the reaction of nitrogen and hydrogen to produce ammonia, it can
be shown that the nitrogen molecule cannot be monatomic.

3. Molecular Weights of Gases. A single molecule of a substance is too
small to be isolated and weighed directly but if we were to weigh equal
volumes of different gases under the same conditions of temperature and
pressure, we would obtain the weights' of equal numbers of molecules of
these gases. The ratio of these weights would also be the ratio of the weights
of single molecules of these gases. For example, for two gases A and B,

Weight of 1000 molecules of gas A 1000 X weight of one molecule of gas A
Weight of 1000 molecules of gas B ~~ 1000 X weight of one molecule of gas B

_ Weight of one molecule of gas A
Weight of one molecule of gas B

If a weight is assigned to one molecule the relative weights of other
molecules are thereby fixed. These relative weights of the molecules are
their molecular weights.

From the early nineteenth century to 1961, the element oxygen was
chosen as standard, and to the oxygen atom was arbitrarily assigned a relative
atomic weight of 16.0000. Thereby the molecular weight of oxygen was
32.0000. For reasons that will become apparent later, the present standard
of atomic weights is the weight of the carbon-twelve (C 12 ) isotope. The C 12
isotope is a species of carbon atom and to it has been assigned the weight
of 12 atomic mass units. The atomic weights listed in Table 5- A are weights
of the atoms relative to this standard. Being a relative weight the atomic
weight of an element is truly a dimensionless number. If we define the atomic
mass unit (amu) as one-twelfth the mass of an atom of C t2 , however,
then the atomic weights can be said to have this unit. Thus, the atomic
weight of hydrogen is 1.00797 amu and the atomic weight of mercury is
200.59 amu. In practical calculations we may use any system of weight units
with the atomic weights. Chemists generally use grams, and when the
atomic weight of an elemen is taken in grams it is called the gram-atomic
weight (Page 59).

On the present C 12 scale of atomic weights the atomic weight of oxygen
is 15.9994 and thereby the molecular weight of oxygen is 31.9998. Thus to
determine the molecular weight of a gaseous substance A, equal volumes
and hence equal numbers of molecules *)f the gas and of oxygen are
weighed at the same conditions of temperature and pressure. The ratio
of the weights of these equal volumes is equal to the ratio of the molecular
weights of the two gases. Thus

weight of a given volume of gas A molecular weight of A

weight of the same volume of oxygen 31.9998

Atomic and Molcenlur Weigf

Table 5-A

Table of Aiomic; Wfighis

(Based on Carbon- 12)

Element

Symbol

Atomic

dumber

Atomic

Weight

Element

Symbol

Xtimoe

\'umhi’r

I \h rmic
j Weight

Actinium

(227)

Meieun

! 200 59

Aluminum

26.9815

Moh bdonnm

1 95 9 \

Americium

(243)

NVodv in mm

1 1 H 24

Antimony

121.75

Neon

_:u is, i

Argon

39.948

Neptunium

\l J

i '

Arsenic

74.9216

Nickel

; 58 7!

Astatine

(210)

Niobium

I 92 906

Barium

137.34

Nitrogen

! 1 0067

Berkelium

(249)

Nobel mm

103

f 35 i )

Beryllium

9.0122

Osmium

1 90,2

Bismuth

208.980

0\\ gen

1 5 999 \

Boron

10.811

Palladium

106 1

Bromine

79.909

rhosphoius

803)738

Cadmium

112.40

Platinum

197 09

Calcium

40.08

Plutonium

9 1

! 3 !3 i

Californium

(219)

Polonium

s i

,210 1

Carbon

12.01115

Potassium

59 102

Cerium

140 12

Piaseothmuim

1 to 907

Cesium

132 905

Pxomcthiuni

tint

Chlorine

35.453

Protactinium

(251 )

Chromium

51.996

Radium

226 05

Cobalt

58.9332

Radon

, 222 »

Copper

63.54

Rhenium

ISO 2

Curium

(245)

Rhodium

103 905

Dyspiosium

162.50

Rubidium

85 47

Einsteinium

£s

(253)

| Ruthenium

4 1 i

| 101 07

Erbium

167.26

Samarium

j Sm

! 1 50.35

Europium

151.96

Scandium

' 41.956

Fermium

100

(254)

Selenium

St*

! 34

78.96

Fluorine

IS 9984

Silicon

! 14

28.086

Francium

(223)

Silver

j 47

107 870

Gadolinium

157.25

Sodium

22 9898

Gallium

! 69.72

Strontium

: os

ST. 62

Germanium

! 72.59

Sulfur

52 0G t

Gold

: 196.967

Tantalum

ISO 948

Hafnium

178.49

Technetium

(99'.

Helium

4.0026

Tellurium

127.60

Holmium

164.930

Terbium

158.924

Hydrogen

1.00797

Thallium

204 37

Indium

114.82

Thorium

232.038

Iodine

126.9044

Thulium

Trn

168.934

Iridium

192.2

Tin

118.69

Iron

55.847

Titanium

47.90

Krypton

83.80

Tungsten

183 83

Lanthanum

138.91

Uranium

238.03

Lawrencium

103

(257)

Vanadium

50,942

Lead

207.19

Xenon

131.30

Lithium

6.939

Ytterbium

173,04*

Lutetium

174.9T

Yttrium

88.905

Magnesium

24.312

Zinc

65.37

Manganese

54.9380

Zirconium

91.22

Mendelevium

101

(256)

The values enclosed in parentheses are those of the mass number of the most
stable isotope.

Atomic and Molecular Weights

It can be determined experimentally that the gram-molecular weight of
oxygen gas, 31.9998 grams, occupies 22.414 liters at standard conditions
(STP). Hence the weight in grams of 22.414 liters of any gas at STP is
the gram-molecular weight of that gas. The volume, 22.414 liters, is known
as the gram-molecular volume (GMV). Actually this value pertains to the
gram-molecular weight of an ideal gas at STP; for real gases at STP, ex-
cluding hydrogen, the value is slightly lower.

4. The Avogadro Number and the Mole. The number of molecules which
is in the volume at STP occupied by the gram-molecular weight of a gas,
and in the gram-molecular weight of any substance (at any conditions)
is 6.023 X 10 23 , or 602,300,000,000,000,000,000,000. This number is known as
the Avogadro Number.

The concept of the mole is one of the most important in chemistry.
Like the word “dozen,” the mole is a number— the Avogadro Number. Thus
we can speak of a mole of molecules, or the Avogadro number of molecules;
a mole of atoms, the Avogadro number of atoms; or of a mole of electrons,
the Avogadro number of electrons. The quantity of matter containing an
Avogadro Number of entities is a mole of that species. The weight of the
Avogadro number of molecules is the molecular weight in grams, or the
gram-molecular weight (GMW); the term gram-molecule is also used to
indicate a collection of the Avogadro number of molecules, and hence its
weight is also the gram-molecular weight. Similarly, a mole of atoms, or the
Avogadro number of atoms, is called the gram-atom and its weight is
the gram-atomic weight (GAW). Thus the gram-molecular weight of oxygen,
31.9998 grams, contains the Avogadro number of oxygen molecules, and
the gram-atomic weight of carbon 12 , 12.0000 grams, contains the Avogadro
number of carbon atoms.

Thus the mole also represents the weight in grams of the Avogadro,.
number of entities. The number of moles represented by a weight in grams
of a substance is given by

where n = the number of moles
( 1 ) n = — g = the weight in grams

M M = the gram-molecular weight (for molecular species)

or the gram-atomic weight (for atomic species)

Example 1: How many moles are represented by 100 grams of water? The
gram-molecular weight of water is 18.0 gram/mole.

Solution:

JL = 100 g

M 18.0 g/mole

— 5.55 moles

The volume occupied by a mole, or gram-molecular weight of a substance,
is known as the molar volume. For a gas at STP, this is 22.414 liters.

A clear distinction must be made between a molecule and an atom.
Whereas the atom is the smallest chemical entity involved in chemical
change, it is not necessarily the smallest unit of matter which can exist in-
dependently under ordinary conditions in nature. Atoms frequently are not

\ttmih* i:iu! Mol* cuhn Wrti'ht'i

individually stable, and two or more combine to innn a unit which can
exist stably. This smallest chemical unit capable of independent existence
is called the molecule. Most gaseous molecules, suc h as In drogi n, oxygen,
chlorine, and nitiogen, are diatomic, that is, two atoms combine to form
one molecule. Some, such as neon and argon, are monatomic. Hu* gum-
molecular weight of hydrogen, 2.016 grams, contains 6.023 . 1 0 J '* molecules
of hydrogen. Each molecule contains two atoms so that one mole* of hydro-
gen contains (2 X 6.023 X 10~ :t ) atoms of hydrogen One gram (1 00S g)
of hydrogen, the gram-atomic weight, contains 6.023 * 10 ‘ atoms of hydro-
gen. The actual weight of a single atom U> r molecule' can be calculated
by dividing the gram atomic weight (oi the gram-moleeulai weight) by
the Avogadro Number. Thus the actual weight of a single hvdiogen atom is

1.008

gram

gram-atom

1.673 X Id-'" 4

6.023 X 10 23

atom

gram-atom

gram

ate mi

At the root of the accuracy of the quantitative chomh al calculations to
follow is the extreme magnitude of the Avogadio Numbei \\Y shall see
later (Chapter 8) that chemical reaction is basically statistical on a mole-
cular scale. As such the “exactness** of our calculations depends upon the
number of molecules with which we deal. If the Avogadro Numbei were
1000, there might be no science of chemistry!

5. Weight and Volume Relationship of Gases. Many problems con-
cerning the weight and volume relationships of gases hmge upon the fact
that, at STP, the gram-molecular weight (GMW) of a gas will occupy the
gram-molecular volume (GMV) of 22.414 liters.

In the following examples, all the gases are assumed to be ideal and
calculations are made to a precision of only three significant figures,

Example 1: Calculate the weight of one liter (the density) of hvdiogen chloride
gas at STP. The gram-molecular weight of hydrogen chloride is 36 5 grant 'molts

Solution: Since, at STP, the gram-molecular weight of hvdrogen chloride,
36.5 g/mole, occupies 22.4 liters (per mole), the weight of one htei is

36.5 gram/mole
22.4 liter/mole

1.62 gram/ liter

Example 2: Calculate the molecular weight of hydrogen chloride if the gas
has a density of 1.62 grams per liter at STP.

Solution: The weight of 22.4 liters of a gas at STP is its gram-molecular
weight. Hence

1.62 gram/liter x 22.4 liter/mole = 36.5 gram/mole

The gram-molecular weight is 36.5 g/mole and the molecular weight is 36.5,
Note that it is only for gases that the volume occupied by one mole i$ 22.4 liters.
No special value pertains to solids or to liquids. For liquid water, whose density

Atomic and Molecular Weights

at STP is LOO g/ml and whose gram -molecular weight is 18.0 g/mole, the molar
volume is

= 18.0 ml/mole

1.00 g/ml

If the weight of a given volume of a gas at conditions other than standard is
to be calculated, the first step is to find the volume the sample of gas would
occupy at STP. Then the GMW-GMV relation at STP is used.

Example 3: Calculate the weight of 15.0 liters of carbon dioxide gas measured
at 100° C and 756 mm pressure. The molecular weight of carbon dioxide is 44.0.

Solution: The volume that this sample of carbon dioxide would occupy at STP is

15.0 liter x

273 deg
373 deg

756 mm
760 mm

10.9 liters at STP

Since the GMW of carbon dioxide is 44.0 g/mole and since this weight occupies

22.4 liters at STP, the weight occupies 22.4 liters at STP, the weight of 10.9
liters at STP is

44.0 gram/mole

22.4 liter/mole

x 10.9 liter = 21.5

Example 4: Calculate the molecular weight of a gas, 15.0 liters of which
measured at 100°C and 756 mm pressure, weigh 21.5 grams.

Solution. As in Example 3, the gaseous volume is first converted to STP. This
volume is 10.9 liters. The GMW is that weight which occupies 22.4 liters at STP.
Thus

22.4 liter/ mole x
10.9 liter

21.5 g = 44.0 g/mole

The Ideal Gas Law permits a simple, direct solution to all problems of
the type illustrated. Since PV = nRT and n = g/M, then

(2) PV = JL RT

v M

If any four of the five variables, P, V, g, M, and T, in Equation 2 are known,
the fifth can be calculated.

Example 5: Calculate the molecular weight required in Example 4 using the
Ideal Gas Law.

Solution: Equation 2 can be transformed to solve for the molecular weight, M

21.5 g x 0.0821

liter atm
mole deg

x 373 deg

756 mm
760 mm/ atm

x 15.0 liter

= 44.0 g/mole

6. Determination of Atomic Weights. If we wish to find the atomic
weight of an element today we can look it up in Table 5-A and if we
know the formula of a compound its molecular weight can be obtained merel}
by summing up the proper values of atomic weights. The implications con-

V tomn i n<i tilui \\ ( i gilts

tained in Avogadro’s hypothesis wore oveilooked until 1S5S In that year
the Italian professor of chemistrv, Stanislao Cannizzuio by a tour do force
of chemical logic, showed how Axogadro’s Hypothesis could ho used to \ icld
the first exact determination of atomic weights.

Molecules are composed of integral numbeis oi atoms In one molecular
weight of a compound there thus would be integial multiples 0 } the atomic
weights of the atoms making up the compound Chemical auaKsis of a com-
pound can give the percent by weight of each atomic species m the compound.
Multiplying the molecular weight of the compound by the expciimentally
determined weight percent of the elements theicin will give tin- weights of
the elements in one molecular weight of the compound Such a weight must
represent some integral number of the atomic weights of the elements The
weight will be equal to the gram-atomic weight if only one atom ot the
element is present in the molecule, or it will he a multiple oi the gram-
atomic weight if more than one atom of the element is present m one
molecule.

For a specific element whose atomic weight is to be determined, if a
large number of compounds of the element arc taken, it is vers likek that
in at least one of these compounds chosen there will be onK one atom
of the element in a molecule. Such a compound will yield the smallest
value for the weight of the element in question, and Canni//aio seasoned
that this smallest weight of the element found in the molecular weights of
& large number of compounds of that element would represent the weight
of one atom, or its gram-atomic weight. Indeed other weights of the element
in the molecular weights of the series of compounds chosen should bo in-
tegral multiples of this smallest weight, corresponding to two or innic atoms
of the element per molecule.

Cannizzaro chose gaseous or volatile compounds of the element whose
atomic weight was to be determined because the molecular weights of
such compounds could readily be determined from their weight of 22.4*
liters at STP. In Table 5-B data are given for the determination of the
atomic weight of the element chlorine by Cannizzaro’s method.

Many elements, particularly the metals, form no or very few volatile
compounds so that the method of Cannizzaro cannot be used to determine
their atomic weights. It is of historical interest that, in 1819, two French
scientists, Pierre Louis Dulong and Alexis Therese Petit observed that the
product of the gram-atomic weight (g/GAW) and the specific heat (calorie/
gram degree) for solid elements is approximately 6.2 calorie/CAW degree.
The specific heat of a substance is the quantity of heat in calories required
to raise the temperature of one gram of a substance one degree centigrade.
Though Dulong and Petit’s rule is not exact it can be used to estimate
the approximate atomic weight of an element from a knowledge of its
specific heat. The combining weight of an element with eight grams of
oxygen can be determined with great precision. This weight is an "integral
multiple or submultiple of the atomic weight of the element but without
a knowledge of the relative number of atoms of the element and of oxygon
in one molecule of the compound the atomic weight cannot be determined.
Calculation of an approximate atomic weight by Dulong and Petit’s rule

Atomic and Molecular Weights

serves as a pointer to indicate which multiple of the combining weight is
the exact atomic weight.

Table 5-B

Determination of Atomic Weight by Cannizzaro’s Method

Substance

Molecular

Weight

111

Percent

Chlorine

By Weight

Weight of
Chlorine In
One GMW
of the
Compound

Formula
of the
Compound

Chlorine

71.0

100.0

Cl,

Methyl Chloride

50.5

35.5 !

CH,C1

Chloroform

119.5

89.1

CHC1 3

Caibon Tetrachloride

154.0

92.0

142.0

CC1 4

Hydrogen Chloride

36.5

97.5

35.5

HC1

Ethylene Chloride

99.0

71.5

■DEB.

c,h 4 cl

Phosphorous Chloride

137.5

79.9

■M

pci.

Column I gives the name of the chlorine compounds. The molecular weights
m Column II were obtained from the weight of 22.4 liters of the compound at
STP. The data in Column III were obtained by chemical analysis of the compounds.
Multiplying the figures in Columns II and III gives the values for Column IV.
In this column the weights of chlorine must correspond to a whole number of
atomic weights. The least weight of chlorine is 35.5 and this is taken to be the
atomic weight of chlorine. Elemental chlorine gas itself is thus diatomic. Column
V gives the molecular formula of the compound as it is known today. In the
event that chance had caused omission of all compounds containing but one
atom of chlorine (a possibility which diminishes with the number of compounds
chosen), and if the value 35.5 corresponded to two atoms of chlorine so that

35.5

the atomic weight was truly — — - or 17.8, then we should have found in Column

IV multiples of 17.8 corresponding to three or five atoms of chlorine, namely
53.4 and 89.0. The fact that these values are not found, however, is further veri-
fication that 35.5 corresponds to the weight of one atom of chlorine.

Example 6: One gram of oxygen combines with 3.750 g of a metallic element,
X, whose specific heat is 0.010 calorie/gram degree. For the element X calculate
(a) the combining weight (b) the accurate atomic weight.

Solution : The combining weight of an element is that weight which reacts
with eight parts by weight of oxygen. For X this is

3.750 g of X
1.000 g of oxygen

x 8.000 g o‘f oxygen = 30.00 g of X

This value, 30.00, is a multiple or submultiple of the atomic weight, e.g. 15.00,
20.00, 30.00, 60.00, 90.00, etc. If one molecule of the compound of X and
oxygen has a 1:1 ratio of atoms of X to atoms of oxygen then the atomic weight

Mdiiiu and Stoh'iulur Weights

of X is 60.00; if this ratio is 4:1 the atomic weight of X is MOD Which of
the possible values is the correct atomic weight will be indicated b> a calculation
of an approximate value by Dulong and Petits tide That multipit 1 of the com-
bining weight closest to the value so calculated is the accurate atomic weight of X.

Atomic weight =

6.2 cal/CAW deg

0.0 10 cal/g deg

=- 62

c;a\\

The approximate atomic weight is 62. The multiple ot the combining weight closest
to this value is 60.00 and hence the accurate atomic weight of \ is 60 00.

QUESTIONS

1. State Gay-Lussac's Law of Combining Volumes. How does Avogadio’s hypoth-
esis explain this law?

2. Describe an experiment which would enable you to determine (.0 the mole-
cular weight of a gas and (b) the molecular weight of a liquid which boils
at 50°C.

3. Prove logically that nitrogen cannot be a monatomic molecule.

4. Outline Cannizzaro's method for obtaining atomic weights.

5. ' Distinguish between mole, molecular weight, gram-molecular weight, gram-

molecular volume, Avogadro number.

6. The population of the earth is about 3 * lQ n people. The population of a mole
is 6 X 10 2a entities. Subtract (3 x 10 u ) from (6 •

7. What is the weight of an Avogadro number of (a) tin atoms (b) CO, molecules;
(c) NaCl formula units?

8. A liter of gas, measured at STP, weighs 2.324 g. What is (a) tin 1 molecular
weight of the gas (b) the gram-molecular weight (c) the weight of oih* mole?

Ans; (a) 52.1 (b) 52.1 g tc) 52 1 g

9. Calculate the density of NO gas at (a) STP (b) 27 °C and 2.0 atm.

Ans; (a) 2.05 g/l (b) 3.73 g/1

10. What is the molecular weight of a gaseous compound, of which 0.684 g oc-
cupies 131 ml at 85°C and 757 mm pressure? Am; 154

11. A volume of 605 ml of a gas at 100° C and 762 mm pressure weighs 0.920 g.

Calculate the molecular weight of the gas. Am; 46.5

12. What volume will 10.0 g of ammonia, NH 3 , occupy at 150 C and 500 mm

pressure? * Am; 31.0 liters

13. The molecular weight of carbon dioxide, 0O 2 , is 44.0. Calculate the weight
of five liters at CO a at (a) STP (b) 50 °C and 380 mm pressure.

Ans; (a) 9.82 g

14. The density of a gaseous element is five times that of oxygen under the

same conditions. If its molecules are triatomic what is the atomic weight of
the element? Am: 53

15. For the gaseous reaction, 2 NO + 0 2 — ► 2 N0 2 , what volume of Q 2 must

react to produce 3.0 liters of NO, if all gases are measured at (a) STP (b) 20
and 750 mm? Am; (a) 1.5 liters

16. For the gaseous reaction, C S H 8 + 5 O z 3 C0 2 -f 4 H 2 C, what volume

of 0 2 measured at 20 °C and 750 mm pressure will be required to react
with 5.0 liter of propane, C 3 H 8 , measured at STP? Ans; 27 liters

Atomic and Molecular Weights

17. Calculate the number of molecules in a drop of water which weighs 0.040 g.

A ns: 1.3 X 10 21

18. Calculate the number of molecules in (a) 1 1.0 g of CCX (b) 5.6 liters of C0 2
at STP (c) 5.6 liters of CO_. at 127°C and 800 mm pressure.

19. Calculate the weight of a single molecule of C0 2 .

20. The pressure in a 200 ml evacuated container is 2.0 x 10’ u mm at 20°C.
How many molecules lemain in the container? How would the nature of the
gas in the container affect your answer?

21. What is the equivalent weight of a metal which combines with oxygen in

the ratio of 10.8 g of the metal to 9.6 g of oxygen? Ans\ 9.0

22. A mass of 80.0 g of a pure. solid metal is heated to 100°C and then quickly

immersed in 400 g of water at 20.0° C. The temperature of the water rises
to 21.0°C. Calculate (a) the specific heat of the metal and (b) its approximate
atomic weight. Arts: (a) 0.0633 cal/ g deg (b) 98

23. A metallic element has a specific heat of 0.030 cal/g deg. If 6.552 g of the
element combine with oxygen to form 8.152 g of an oxide, what is the most
accurate value for the atomic weight of the element that can be calculated
from these data?

24. A mass of 69.667 g of a solid metallic element combines with 8.000 g of

oxygen. The specific heat of the element is 0.0305 cal/g deg. What is the
accurate atomic weight of the element? Ans: 209.0

25. In Example 6 (Page 63) what ratio of atoms of X to atoms of oxygen would
result in an atomic weight of X equal to (a) 30.00 (b) 120.0.

26. The data in the following table pertains to compounds of the elem X.
(a) Complete the last column in the table.

Compound

Molecular Weight

Percent of X

Number of Atoms of
X in one molecule of
the compound

82.3

63.7

97.7

53.6

(b) What is the approximate atomic weight of X? (c) If 0.2300 g of an oxide is
found to contain 0.1600 g of oxygen, what is the accurate atomic weight of X?

27. For 19.2 g of sulfur dioxide gas, S0 2 , calculate (a) the number of moles of
S0 2 (b) the number of molecules of S0 2 in one liter at STP (c) the total
number of atoms in 19.2 g of S0 2 (d) the weight of one S0 2 molecule.

28. Atoms of elements A, B, and C combine to form a compound in the atomic
ratio of 1:6:2 respectively. The atomic weights of A, B, and C are 64, 4,
and 16. Calculate the maximum weight of the compound formed by the reaction
of 1.28 g of A, 3.01 x 10 23 atoms of B, and 0.04 gram-atom of C.

Chemical Calculations

1. Symbols and Formulas. To recoid information concerning t
change, chemists have developed a simple notation in whuh each element
is represented by a standard symbol. The alchemists used ustionoimt al Injures
derived from a belief in die “sympathy” of the metals and the planets, Dalton
used circular figures. (Figure 6.1)

6 c o •

Gold Silver llydiogen Lai bon

tor the Sun lor the Moon

Alchemical Symbols Dalions S^mboi-s

Figure 6.1. Early Chemical Symbols.

Modem chemical symbols, as proposed in 1811 by the Swedish chemist,
Jons Jakob Berzelius, are merely the first letter of the element's name. Thus
H stands for hydrogen and C for carbon. To distinguish between elements
whose names start with the same letter, a second letter is adde d to the symbol
of one of the elements. The symbol for calcium therefore is La and that
for chlorine is Cl. Only the first letter is capitalized. Some of the symbols
are derived from the Latin names of the elements, particularly hu those
elements which were known in antiquity, for example. An for gold (aurum)
and Ag for silver (argentum). The symbols of the elements are listed in
Table 5-A and are recognized by international agreement.

More than merely the chemical species is meant. when a chemical symbol
is written. The symbol “H” not only denotes one atom, of hydrogen but
also specifies a given weight of # hydrogen, namely the atomic weight. 1.008
amu. Inasmuch as a molecule of a substance consists of a definite number
of atoms, the substance can be represented by a formula containing the
symbols of the elements. The formula therefore represents one mole-
cule of hydrogen, which is composed of two atoms of hydrogen. The mole-

Chemical Calculations

cular weight of hydrogen would be the weight of two hydrogen atoms, or
2.016 amu. Subscripts in a formula indicate the internal constitution of a
molecule by specifying the numbers of each kind of atom within the mole-
cule. Thus Fe^Oi is the formula for an iron oxide (rust), each molecule
of which consists of two atoms of iron and three atoms of oxygen. One
molecule of this iron oxide has a molecular weight of (2 X 55.8) H- (3 -b 16.0),
or 159.6 anm. Strictly speaking, the sum of the atomic weights of the ele-
ments m a formula is the formula iceight. The simple entity represented by
the formula may not exist as such in nature The molecule may actually exist
as associated formula units each having a weight which is an integral multiple
of the formula weight and the term molecular weight should be applied to
the weight of that configuration. Also, as we shall see later, many “mole-
cules” for which apparently molecular formulas can be written, as in the
case of sodium chloride, NaCl, do not exist as molecular entities but are dis-
sociated into units known as ions. For such substances we refer to the for-
mula weight, and the formula weight in grams, or the gram-formula weight
(GFW), would represent the weight of the Avogadro number of formula
units, or one mole of formula units.

An integer written before a formula specifies a number of molecules of
a substance. Thus 2 Fe^Oj represents two molecules of iron oxide (rust)
and a total weight of (2 X 159.6), or 319.2 amu. Such a number preceding
a formula multiplies everything which follows it, whereas a subscript mul-
tiplies only the symbol (or symbols inclosed m parentheses) which precede
it. In the formula Ca 3 (P0 4 ).> the subscript “4” refers only to the oxygen
atom before it, whereas the subscript “2” multiplies the entire (P0 4 ) group.

2. Equations. A chemical change such as the rusting of iron can be
written m the form of a chemical equation with the use of chemical sym-
bols. A chemical equation is written so that the formula(s) of the reacting
substances, or reactants, appear on the left, and the products of the reaction
on the right of an arrow. The arrow or equal sign may be interpreted as
“react to form.” Thus for the rusting of iron,

Fe + Ch -> Fe,0<

Inspection of this “equation” indicates an apparent discrepancy. We start
with two oxygen atoms, as one molecule, CL, and somehow produce three
oxygen atoms in the iron oxide, Fe 2 0.<. Two oxygen atoms weigh 32.00
amu and three would weigh 48.00 amu. The equation, as it now stands,
apparently contravenes the Law of Conservation of Mass. To satisfy this
law we place coefficients before each reactant and product, as necessary,
to make the total numbers of each atom on the left and right of the arrow
equal. This process is known as balancing the equation.

To balance the oxygen atoms we would have first

Fe + 3 O 2 — ^ 2 Fe 2 0,

but. the iron atoms must also be balanced so that the completely balanced
equation would be 4 Fe + 3 0 2 2 Fe 2 0 3 .

It would be incorrect to write 4 Fe + 6 0-^2 Fe>0 :i because oxygen does
not occur in nature as a monatomic molecule but only as the diatomic molecule,

gg Chi mu a? Ctdruhticm

Oo. The equation must xeprcsent the nutuie* ot tin* substance 4 * as th(*\ tniK exist;
when we use oxygen in an experiment, it comes as diatomic inoi«*v ulus, O Iron
does exist as individual atoms. Smnlurh it would be imvucit to balance the fore-
going equation by writing Ftc O, — Fe.O iuvau.se this also tnmpeis with the
chemical nature of the iron and oxygen molecules Natme takes no cogm/ance
of our paper mechanics and the meie writing ot torinuLis tor subsumes does not
cause them to exist. The number “1” is not written in balancing equations, where
no number appears before a formula, the value " 1” is undeistood.

Algebraic Method for Balancing Equations

In addition to the trial and eiroi method ot balancing equations sxstematic
methods exist, some of which we shall have occasion to examine m detail d hupter
22). One technique is purely algcbiaie m that it sets up simultaneous equations
based upon the fact that the numbers of each species ot atom on both sides of
the balanced equation must be equal. The funly complicated osutiou between
potassium permanganate, KMnO t , and hvdrochlonc acid, HCI, < an be* used as an
illustration. The unbalanced equation foi the leaction i,\

* KMnO, + b HCi -> c KCl 4 d MnCl, - v Cl, ” fHO

where a coefficient from a to / has been inserted befon each molecular toimula.
Balancing the equation requires determination of the piopei values of these co-
efficients. Because the number of potassium atoms {K atoms) on each side of the

equation must be equal, therefore a = c.

Similarly for the Mn atoms a ~ d

for the O atoms 4 a /

for the H atoms b ~ 2f

for the Cl atoms b = c 2d -f 2e

We now arbitrarily assign the value of “1” to any of these coefficients If a is

chosen to be “1,” then c = 1, d = 1, f = 4, b = 8, and e — 5/2.

The balanced equation would then be

KMnO, + 8 HCI -> KC1 + MnCL + - CL + 4 H.O

Because only integral numbers of molecules react, the properly balanced equation
would be double the foregoing or

2 KMn0 4 + 16 HCI -> 2 KCI + 2 MnCL, + 5 Cl, +■ 8 H.O

3. Percentage Composition from a Formula. Since both the qualitative
and quantitative composition of a compound are represented hy a chemical
formula, the. percentage by weight of each of the elements in a compound
can be calculated from its formula. In effect analytical computations can he
made in that we can calculate the results that would be obtained by a
laboratory analysis of the pure substance represented by the formula.

Example r 1: The formula of ethane is C,H tl . What is the percentage composi-
tion by weight of this compound?

Solution: The formula indicates that there are two atoms of carbon and six
atoms of hydrogen in one molecule of ethane. Thus there are two atomic weights
of carbon and six atomic weights of hydrogen in one molecular weight of ethane.
The total weight of the elements in the formula, or the formula weight (molecular

Chemical Calculations

weight) of ethane is (2 x 12.0) + (6 x 1.0), or 30.0. The percentage of each of the
elements is obtained by dividing the weight of each element in one formula weight
by the formula weight of the compound.

2 x 19 0

Carbon = — = 0.800 = 80.0%

uU,U

Hydrogen = 6 * J - = 0.200 = 20.0%

Since the total percentage of all the constituents in a compound must add up to
100% the percentage of the carbon or of the hydrogen could have been obtained
by difference if the other were known, but it is preferable to work out both per-
centages independently as a check.

4. The Formula from the Percentage Composition. Conversely when the
proportions by weight of the elements in a compound are known, the formula
of the compound can be calculated. To write a formula we must know the
number of atoms of each element in one molecule of the compound, or the
number of atomic weights of each element in one molecular weight. If the
percentage of each element is divided by its respective atomic weight, the
quotients obtained are in the ratio of the number of atoms of each element
present in one molecule of the compound. An illustration will make this clear.

Example 2: A compound contains 80.0% of carbon and 20.0% of hydrogen.
What is the formula of the compound?

Solution. The formula is C x H y where x and y are integral numbers. Dividing
each percentage by the atomic weight of the respective element we obtain quotients
which stand to each other in the ratio of the number of atoms of carbon to atoms
of hydrogen in the formula.

80.0 g carbon ^ 1.00 g-atom carbon __ ^ g-atom carbon

100.0 g compound 12.0 g carbon * g compound

20.0 g hydrogen x 1.00 g-atom hydrogen __ ^ g-atom hydrogen

100.0 g compound 1.01 g hydrogen g compound

The ratio of these numbers, or the ratio of the numbers of gram-atoms of
carbon and hydrogen in the compound, is the same as the ratio of atoms. Because
the numbers of atoms in a formula must be integral it would be incorrect to write
for the formula C 0 066 H o 108 . However the ratio of these numbers, 0.066 to 0.198,
or 1 to 3, gives the smallest whole number ratio of carbon atoms to hydrogen
atoms. For every atom of carbon there are three atoms of hydrogen. In other
words the simplest formula is CH*.

The problem could have been stated in terms of other analytical data such as
giving the actual weights of the elements in a specific sample of the compound;
thus 18.0 g of carbon combine with 4.5 g of hydrogen to form 22.5 g of the
compound ethane. Dividing these weights by the respective atomic weights would
yield th « same 1 to 3 ratio.

The simplest formula, or the “least common denominator” formula, is known
as the empirical formula . For the foregoing example this is CH*; however this
formula is not necessarily the molecular formula. The molecular formula can be CH 3
or a multiple of the empirical formula, for example, C 2 H 6 , C 3 H 9 , C 4 H 12 , etc., be-
cause in each of these formulas the ratio of carbon to hydrogen atoms is also 1

to 3, and so each oi these lomtulas w«mM \:*'id the '• urn u i d 1 * pen outage of
caibon to hydrogen, In oulei to deteniun* the tnii-t m-’liiuK , l u i*-u. i *«*i the
compound, the moleeulai weight of the t umpui mi mi^! i>< hnouj, i hi molecular
weights of the possible compounds listed .ue .e tol’ow » t I I ^ 1 > l \ [ H« 30,0;
CjH,, = 45 0; C,fltj 600, It \u* hn<iw hu iii*4.eu** that 2 1 *» i*te^ uf ethane,
a ‘gas, at STP weigh appioMmateh 5u m tmv 9*< u th« d.mtdr mulOple of the
empiiic.il formula is the mokvulai tnimula. n.uneb < H

The following problem is given m oidei to shun the tabulation of the
molecular formula when additional data air mnn, so that tlx molecular
weight can also be calculated.

Example 3. A substance contains 37. S3 oi i.nben, h V> oi hvtliugm. and
55.9% of chlorine. When vaporized, 2 59 g of tin substaw e njn 624 ml

at 10()°C and 775 mm piessuu*. What is the muleeuhu tunuuU ut the com-
pound? What is the accurate molecular weight <4 the mmpnnuil*'

Solution: The relative amounts oi caibon, hvdtogen. and thimine m the com-
pound are 37.8, 6.3, and 55.9 respect i\ eh . Dividing ea« h In the atomic weight

of the respective element, we obtain y—j* 5.15 hn i.nbun, — — 6 $ tui huho-

55.9

gen, and = 1.575 for chlorine.

w.O

The ratio of the number of caibon atoms to that ui hv dmgeu and that of
chlorine is 3.15 : 6.3 : 1.575. How to e\pu*ss tins latm m simple whole num-
bers may not be apparent on inspection, but it each I, Mm is divided bv the

smallest (1.575), we obtain: =: 2 for carbon, ~~r 4 tot hvdogen, \ 1

1.575 1 o*.> l.o.o

for chlorine. The simplest or empmea! tonuula is <d,H,C| and its tounuU
weight is 63.46. The molecular formula is either the empmcal inutitihi or some
multiple of it, and the molecular weight of the compound is the tot inula weight
of the empirical formula or some multiple of it

The molecular weight of the compound can be calculated from the data given
in the statement of the problem. The volume occupied bv 2 59 g <4 the sub-
stance at STP is

273 deg K 775 mm

V (at STP) = 624 ml *

373 deg K 760 mm
The gram-molecular weight is the weight of 22.4 liters at STP, Thu
2.59 g _ GMW

166 ml oi 0. 1 16 liter

0.466 liter .22.4 liter/mole

; hence the GMW =- 124.5 g mole

The GMW could also have been calculated directly ftom the Ideal Gas Law.

The approximate molecular weight of the compound is 124.5. The actual mole-
cular weight is that multiple of the formula weight of the empirical foimula
which is nearest to the value of the approximate molecular weight. In this case,
it is 2 x 63.46, or 126.92. Thus the molecular formula is C t HsCL and the exact
molecular weight is 126.92.

5. Chemical Equations. From balanced chemical equations, calculations
can be made concerning the weight and volume relationships existing among
the reactants and products of chemical change. Such chemical problems are
quite simple and should properly be called chemical arithmetic. They in-

Chemical Calculations

volve nothing more than the idea of simple proportion and are no more
difficult than the problem, “If three apples cost seven cents, how many
apples can be bought for thirty-five cents?” Somehow students find this
difficult when the units are chemical and not fruit. Thus “if three grams
of oxygen react with seven grams of iron, how many grams of oxygen will
react with thirty-five grams of iron?” Whereas the economic market gives
the rate of exchange between apples and money, for oxygen and iron this
proportion is given by the balanced chemical equation.

A chemical equation represents not only the nature of the reacting sub-
stances and the products formed but also the relative weights of these
substances since each is associated with a definite relative weight, namely,
its formula weight.

For the rusting of iron, 4 Fe + 3 0 2 2 Fe 2 0 3 , the balanced equation

states that 4 atoms of iron react with 3 molecules of oxygen to form 2 mole-
cules of iron oxide. Any common multiple of these numbers also will react
completely. Thus 400 atoms of iron will react with 300 molecules of oxygen
to form 200 molecules of iron oxide, and 4 times the Avogadro number of
iron atoms will react with 3 times the Avogadro number of oxygen mole-
cules to form 2 times the Avogadro number of iron oxide molecules.
Hence 4 atomic weights of iron (4 X 55.8 = 223.2 amu) will react with
3 molecular weights of oxygen (3 X 32.0 = 96.0 amu) to produce 2 mole-
cular weights of iron oxide (2 X 159.8 = 319.2 amu).

6. Weight— Weight Problems. The following examples are typical of the
calculation of the weight of one substance that will react with, or be pro-
duced from, the weight of another substance.

Example 4: In the process of rusting what weight of oxygen will react with

35.0 grams or iron?

Solution: Method I

The balanced equation is 4 Fe + 3 0 2 — » 2 Fe 2 0 3

The solution of this problem hinges upon the fact that four atomic weights of
Fe react with three molecular weights of 0 2 .

First we find the number of gram-atomic weights of Fe represented by the

35.0 g of Fe. The number of gram-molecular weights of O , that will react with
this quantity 7 of Fe will be three-fourths as many. Then to find the weight in
grams of 0 2 that this number of, gram-molecular weights of O a represents, we
multiply by the GMW of 0 2 .

Since 55.8 grams is the gram-atomic weight of Fe, 35.0 grams of Fe equals

35.0 g of Fe

55.8

g of Fe
GAW of Fe

gram-atomic weights of Fe.

Since 3 molecular weights of 0 2 react with 4 atomic weights of Fe,

35.0 g of Fe ^ 3 GMW of 0 2 is the number of gram-molecular

pr , o gof Fe x 4 GAW of Fe weights of O s that react.

55 8 GAW of Fe

( 'hcmit iil ( 'alt uhitttms

But each gram-molecular weight of O, wtuglis *>2 0 gi«un s . *md luuut* the weight of
0 2 is

35.0 g of Fe

55.8

g of Fe

3 GMW of O.

4 GAW of Fe

.52 0

g nl O
CMW nl O

15 1 gs ains 0,

GAW of Fe

The weight of the Fe.,0, piodueed tan also 1 h* ink-, dated in a smulai maimer.

35.0 g of Fe

5« 8 JLSlIS

GAW of Fe
= 50.1 grams Fe 2 O a

2 CMW oi Fe () U oi Fe O

4 GAW of ¥c

GMW of Fe. ()

Method II

The balanced equation gives the weight i alios ul uMi-t.mts ami pmducts, as
calculated in paragraph 5. These \ nines of the weight ratio aie listed undei the
respective substances in the equation following, x is tin* weight ot oxvgen to he
calculated that reacts with 35.0 giains of iion.

35.0 x

4 Fe 3 O. 2 Fe.O

223 2 96.0 319.2

Since 223.2 grains of iron react with 96.0 glams of owgen, the weight of owgeu,
x, can be determined by the following proportion:

35.0 g of Fe _ x g of 0 :

223.2 g of Fe ~ 96.0 g of O,

Essentially Methods I and II are the same.

and x — 15,1 giams of t>.

Example 5: What weight of hydrogen will be produced bv the reaction of

10.0 g of zinc with excess hydrochloric acid, HC1?

Solution: The balanced equation is Zn 4 2 HC1 — ■* ZnCl 2 -t 1G

10.0 g of Zn v/ 1 GMW of H 2 ^ p g of H,

Ae ^ rt of Zn 1 GAW of Zn * ‘ # ~ GMW of IL

8 CAW of Zn
= 0.309 g of hydrogen

In many cases it is unnecessary to know the completely balanced equa-
tion in order to solve a stoichiometric problem. For example, the weight of
phosphoric acid, H 3 P0 4 , that can be made from elemental phosphorus is
limited by the fact that there is one P atom in H 3 P0 4 . The weight relation-
ship between P and H 3 P0 4 is such that one P atom can form only one mole-
cule of H 3 PO 4 , and hence the weight ratio of P to H 3 PD 4 is 31 to 98 Similar-
ly, in the formation of Fe 2 0 3 from Fe it will take two atoms of Fe to form
one molecule of Fe 2 O a and the weight ratio of Fe to Fe a () 3 is (2 X 55.8)
to 159.8.

7. Weight— Volume Problems, Often it is necessary to calculate the
volume of a gas that will react or be produced in a chemical reaction when
a given weight of a substance is used. Since one gram-molecular weight of
a gas occupies 22.4 liters at STP, this volume can be substituted for each
gram-molecular weight of the gas in the calculation.

Chemical Calculations

Example 6: When 35.0 grams of uon rust (a) what volume of oxygen at STP
will react? (b) what volume of oxygen measuied at 25 °C and 770 mm pressure
will react?

Solution: Method I

(a) 35.0 g of Fe v 3 GMW of (X ^ 1 liter of CX

ft* a g of Fe 4 GAW of Fe * * 4 GMW of O,

GAW of Fe
= 10.5 liter CX at STP

(b) If the gaseous volume is to be calculated at conditions other than STP, the
volume at STP can be converted to the desired conditions, or the foregoing ex-
pression can be "extended.”

298 deg K 760 mm

273 deg K 770 mm

Method II

10.5 liter x

= 11.3 liter CX

(a) In place of each GMW of oxygen, the GMV of 22.4 liters can be written.
Then x would be the volume of oxygen at STP that will react with 35.0 grams
of iron.

Hence

35.0 g of Fe
223.2 g of Fe

35.0 x

4 Fe x 3 0.. 2 Fe.O,

223.2 (3 X 22.4)

x liter of 0 2 at STP
67.2 liter of 0 2 at STP

and x = 10.5 liter CX at STP

(b) The volume of oxygen at STP can be converted to the desired conditions as
in Method I.

To find the weight of a substance required to produce a given volume
of a gas, the procedure is reversed. First we calculate the volume of the gas
at STP, and die number of moles this represents, then the number of moles
of the reactant concerned, and finally its weight.

Example 7: What weight of calcium carbonate, CaCO s , will react with excess
hydrochloric acid, HC1, to produce 10.0 liters of carbon dioxide, C0 2 , at 25 °C
and 770 mm pressure?

Solution : The equation for the reaction is

CaC0 3 + 2 HC1 CaCl 2 + CCX + H 2 0.
At STP this quantity of CO a would occupy

10.0 liter x

273 deg K
298 deg K

x 770 mm
760 mm

9.28 liter C0 2 at STP

This volume of C0 2 is — = 0.414 mole CO a

22.4 liter/mole

From the balanced equation we see that one moje (formula weight) of
CaCCX yields one mole of C0 2 ; therefore 0.414 mole of CaCCX will be re-
quired. Since the molecular weight of CaC0 3 is 100, the weight of CaCO s is

0.414 mole X 100 01 ~ y = 41.4 grams CaC0 3

CrMW

( /< 'tLitum

8. Volume— Volume Problems. \\ lien* two m inun* <mm*s mis- *m ul\ ed u\ a
chemical reaction we can calculate the volume of out u hah wih nstct

with, or be produced from, a given \oluine ol anothei gas '1 tie solution to
such problems is provided b\ C»a\ -Lussuc s 1 .aw ot ( omhtnmg Volumes,
namely, that there is a small integral iatio between the solium s el gases m
a chemical reaction. This latio is the same as the latm nt *h< numbers of
molecules which react or arc produced and lienee is given b> the eort fluents
before each gaseous substance tn the balanced ehemn a! equation I he iatio
of these coefficients is identical with tin* ratio ni the gas* nus volumes This
relation holds provided the gases concerned aie at tin same temperature
and pressure; if they are not, then their volumes must be mm cited to some
common conditions.

Example 8: The equation im the burning o! pmpams < lb, is
C, 5 H s 5 (), — 3 TO * Ml <>

(a) What volume of oxygen is lequned to react with 3 1 ht< n ^ < f pmpane J >h What
volume of carbon dioxide will be pioduced 4 M It “> 1 liters ni pmpane .or burned
in 40.0 liters of oxygen, what \olume of each snhstam e will u ab Ml gases
are measured at the same conditions.

Solution * (a) From the equation the ratio <>t the number uf moles of piopam*
and oxygen is 1 to 5. The ratio ol the volumes ol piopane ami nwgrn whnh re-
act is the same. Since the volume o( ovgeu icqnned is hve times the volume
of propane, the volume of ovvgen is 5 3 I liteis 27 0 btuv

(b) The iatio of the volumes of propane ami c at bon dioxide is l to v The volume
of carbon dioxide is therefore three times the volume of piopane, oi 3 • ,V| liters
— 16.2 liters;

oxygen in excess of the 27.0 liteis will remain umeaeted, Ilnur t 10 0 27 Ub or

13.0 liters of oxygen will he left ovei. If 40.0 liters of eaeh gas wen* mixed the
propane would be in excess. With 40.0 liters of oxygen, onk SO liteis oi propane
would react and 32,0 liters of propane would lemam

9. Valence. If we consider the equations foi the react ions of the metals,
sodium, zinc, and aluminum, with hydrochloric acid,

2 Na -b 2 HC1 — 2 XaCl 4. U.

Zn + 2 HC1 -* ZnCl, * H,

2 A1 + 6 HC1 -* 2 AlCl , + 3 H.

we note that the ratio of the number of atoms of hydrogen liberated by the
reaction of one atom of the metal is 1:1 in the case of sodium, 2:1 in the

case of zinc, and 3:1 in the case of aluminum. Further, if we compare the

formulas of the chlorine compounds, sodium chloride, XaCl, calcium chloride,
CaCl 2 , and aluminum chloride, A1CI*, we see that elements can have dif-
ferent combining capacities in that one sodium atom combines with one
chlorine atom, one calcium atom- with two chlorine atoms, and one aluminum
atom with three chlorine atoms.

The chemical combining capacity of an element is called its valence.
Again whenever a quantitative scale of values is to he established a standard
must be chosen.. The standard of combining capacity chosen is the combin-
ing capacity of the hydrogen atom, and thus also of the chlorine atom

Chemical Calculations

inasmuch as one hydrogen atom combines with one chlorine atom, as in
hydrogen chloride, HC1, so that hydrogen and chlorine must have equal
valences. The valence of an element thus is defined as the number of atoms
of hydrogen, or of chlorine, with which one atom of thg element will com-
bine or displace in chemical reaction. The valence of sodium, calcium, and
aluminum would be one (univalent), two (bivalent), and three (trivalent),
respectively. It should be noted that the valence of an element is a pure
number.

In many cases a knowledge of the molecular formula permits one to de-
duce the valence of an element. From the formulas H 2 0 and CH 4 it is
evident that the valence of oxygen is two and that of carbon is four. In
some cases groups of atoms, or radicals, undergo chemical reaction as a unit
Each radical, as an entity, has its own valence. That the sulfate radical,
(SOj), has a valence of two is readily verified from the formula for sul-
furic acid, H_>SOi.

A fundamental understanding of valence and the mechanism of chemical
reaction must be based on a knowledge of atomic structure. In Chapter 13
we shall see that valence can also be defined as the number of electrons
lost or gained by an atom (electrovalence) or the number of pairs of elec-
trons shared by an atom (covalence).

All elements and radicals do not combine with each other. In Table 6-A
are listed the valences of common elements and radicals. For reasons which
will become apparent later, elements and radicals in one column can com-
bine with elements and radicals in the other column but elements within a
given column do not generally combine with each other.

When combination to form a stable molecule does occur, the total valence
or combining capacity of the combining elements or radicals must be equal.
A knowledge of valences thereby enables us to predict the formulas of com-
pounds. For example, what is the formula of calcium phosphate? The
formula will consist of two parts, the calcium and the phosphate. Table 6-A
shows that calcium, Ca, has a valence of 2, and the phosphate radical,
(PCh), has a valence of 3. Since only integral numbers of atoms and radi-
cals may be taken, though it is algebraically true that IV 2 calcium atoms
would have a valence of 3 and so balance the valence of a phosphate radical,
three calcium atoms with a total valence of 6 are required to combine with
two phosphate radicals also having a valence of 6. The formula for calcium
phosphate is therefore Ca, { (P0 4 );>.

Tacitly we have assumed that an element has but one characteristic val-
ence, but under different conditions some elements exhibit more than one
combining capacity. That this must be the case is evident from the Law of
Multiple Proportions. In many of our illustrative problems we have used
the iron oxide, Fe,0„ wherein the iron has a valence of 3. Iron also ex-
hibits ft valence of 2, as in FeO. The term “iron oxide” is not sufficiently
specific. Where a metallic element has more than one valence, modern
nomenclature indicates the valence 1 by writing a Roman numeral in paren-
thesis after the name of the element. Thus Fe 2 0» is written iron(III)

'More properly the oxidation number (page 175) and not the valence is specified by the
Roman numeral; however the two, oxidation number and valence, are closely related.

Table fcf-A

Common Valences of Some Eleven is and K a me a us

Element

Symbol VaU

Illustrative Element

Symbol

Val-

Illustrative

ence

Compound or

ence

Compound

Radical

Ammonium

(NH 4 )

NH.C1

Acetate

iC'H,COO!

~T~

CH,C OOH

Copper(I); cuprous

CuCi

Arsen ite

(AsO.)

KAsO.

Copper (II); cupric

CuCl,

Hydrogen

ihc:o,)

NaHCO,

Hydrogen

HC1

| carbonate

HBr

Mercury(I);

HgCl

| Bromide

mercurous

B ruinate

(BrO,)

NaBiOj

Mercury(II);

HgCl.

Chloride

HC1

mercuric

Chlorate

(CIO,)

KC10 3

Potassium

KOI

Cyanide

(CN)

KCN

Silver

KgCl

Fluoride

Sodium

NaCl

Hydroxide

Iodide

Nitrate

Nitrite

Dihydrogen

(OH)

(NO.)

(NO,,)

(H..PO.)

NaOH

HNO a

NaNO..

n\iH 7 po 4

Barium

Cadmium

Calcium

Cobalt

BaCl.

CdCI 2

CaCl,

CoClj

phosphate

Iron(Il); ferrous

FeCl.

Permanganate

i MnO.)

KMnO*

Iron(III); ferric

Fed-,

Carbonate

(CO,)

Na/A

Lead

PbCl. i

Chromate

(CrO.)

K.CrO,

Magnesium

MgCL

Oxide

H,0

Manganese(II);

MnCI.

Peroxide

(0..)

manganous

Sulfate

(SO.)

H.SO*

Manganese(IV);

MnCI,

Sulfide

h 3 $

manganic

Monohydrogen

(HPO.)

N.i,HP0 4

Nickel

Strontium

Tin(H); stannous
Tin(IV); stannic

NiCl.
SrCl.,
SnCL
SnCl. |

Phosphate

Arsenate

Phosphate

(AsO.)

(PO.)

h,po 4

. . 5

Carbon

Silicon

CO,

SiO.

Aluminum

Antimony

A1C1,

SbCl 3

Bismuth

BiClj

Chromium

CiCl,

Oxide and FeO is iron (II) Oxide. An older nomenclature uses the suffix
“ ic ” for compounds of the higher valence and the suffix * ous” to indicate
compounds with the lower valence. In this system F<uO ;} is ferric oxide
and FeO is ferrous oxide. Some elements have many valences. Nitrogen
has five but, on the other hand, elements such as sodium, potassium, cal-
cium, aluminum, hydrogen, and oxygen practically never show more than
one valence.

The fact that an element may have more than one valence may appear
to make the concept of valence of little value. This is not so. If an element
has two valences it forms two distinct groups of compounds, for example,
iron (III) and iron (II) compounds. The compounds of a given group are
similar to each other but differ markedly from compounds of the other

Chemical Calculations

group. Thus iron (III) compounds, such as FeCl, and Fe 2 (S0 4 ) 3 , have
similar properties which are different from those of the iron (II) com-
pounds, FeCl 2 and FeSO t . It is as though iron in the iron (III) condition
were an entirely different element from iron in the iron (II) condition.
Each is easily recognized and distinguished from the other in laboratory
practice.

10. Equivalence. In chemical reaction the numbers of molecules of the
reacting substances generally are not equal. Indeed such a situation is rare.

It is true of the reaction Zn + H 2 S0 4 — > ZnS0 4 -j- H 2
but not of Zn + 2 HC1 ZnCl 2 + H 2

What- is true of both these reactions is that the combining capacities of the
reactants are equal. Thus the zinc atom has a valence of 2. In terms of
valence it is equivalent to two hydrogen atoms and in both reactions one
zinc atom did displace two hydrogen atoms. Only one molecule of H 2 S0 4
is required to furnish two hydrogen atoms but two molecules of HC1 are
necessary.

The hydrogen atom is taken as the unit of chemical equivalency and the
chemical equivalent is defined as that quantity of a substance which will
react with or contain one atomic weight of hydrogen— or one atomic weight
of chlorine since both hydrogen and chlorine are univalent and therefore
equivalent to each other. Thus one atomic weight of zinc contains two
chemical equivalents, one molecular weight of sulfuric acid two chemical
equivalents, and one molecular weight of hydrochloric acid one chemical
equivalent. Two molecular weights of hydrocloric acid would contain two
chemical equivalents. It is evident that chemical reactions take place in such
manner that the numbers of chemical equivalents of the reactants are equal.
The concept of valence has enabled the chemist to reduce chemical re-
actions to the same standard of exchange, namely, the chemical equivalent.
From this viewpoint, valence may be defined as the number of equivalents
in one molecular weight of a substance (or in one atomic weight in the
case of a monatomic molecule).

The equivalent weight , or the weight of one chemical equivalent, is
defined as that weight of a substance which, will react with or contain one
atomic weight of hydrogen (1.008) or of chlorine (35.45). Thus the equiva-
lent weight of zinc is its atomic weight divided by two, since one atomic
weight of zinc will react with two atomic weights of hydrogen and hence
one-half the atomic weight of zinc is equivalent to one atomic weight of
hydrogen. Similarly the equivalent weight of H 2 S0 4 is one-half its molecular
weight while for HC1 the equivalent weight is equal to the molecular weight.
The atomic weight and the molecular weight are integral multiples of the
equivalent weight. It is now apparent why the weight known as the com-
bining weight, or the equivalent weight ( page 55 ) , was defined as that
weight which would combine with 8.000 parts by weight q£ oxygen.

Specifically for an atomic reactant,

( 1 )

equivalent weight

atomic weight
atomic valence

For a molecule composed oi two ladicals, since the total valence <»i each
radical is the same.

moleculai weight <01 humnli weight /

( 2 ) equivalent weight - to tal valencr

As with atomic and molecular weights, equivalent weights 1 u\ r no units.
However a weight in grams equal to the equivalent weight ot a substance
is known as the gram -equivalent weight (CEW ), 01 siinph tin* gram-cquii -
alent.

Example: Calculate the equivalent weights ol Ah AhOIF , A! (SO,
Solution

For Ah equivalent weight

atomic weight 27 H

atomic valence ]

For A1(OH) h , equivalent weight

moleculai weight
valence (ol one Al atom*

For A1 2 (S0 4 ) 3 , equivalent weight

moleculai weight
valence (of two Al atoms »

57 0

QUESTIONS

1. What is the complete significance of the symbol What is the signifi-

cance of the formula, H 2 S0 4 ? What information can be deduced horn this
formula?

2. Why is it incorrect to write the formula for water as* II </ Would 11,0.
be correct?

3. What information is expressed in a chemical equation? Why must chemical
equations be balanced?

4. Define valence; equivalent weight; chemical equivalent. What is the relation-
ship between atomic weight and equivalent weight?

5. Balance the following equations:

(a) Fe 2 0 3 4 H, -» Fe + H,0

(b) Al + HC1 A1C1, + H..

(d) CaC0 3 4 HC1 -» CaCL 4 H.,0 + CO..

(e) NaOH 4 HC1 NaCf 4 H.,0

(f) . Cu 4 HNO a Cu(NO a )o 4- NO., + H.,0

(g) KC10 S KC1 + O.,

(h) NH 3 4 Oo -> NO + H.,0

(i) Ca 3 P 2 + HoO Ca(OH), 4 PH,

(j) KMn0 4 4 H 2 S0 4 4 H 2 S -» K 2 S0 4 4 MnSO, - li.O - S

(k) I 2 4 HN0 3 HIO, 4- NO 4 H 2 0

6. Calculate the formula weights (molecular weights) ot the following com-
pounds: ZftS; Ca 3 p 2 ; Co(OH).,; AL(SO,) 3 ; C,.H,0; C.H ,OH; CH,COOH;

C„H 22 0 ii.

7. Calculate the percent composition of each of the following compounds: (a)
Na 2 S0 4 (b) Na 2 \S0 3 (c) K 2 CrO t (d) K 2 Cr a 0 7 (e) C a H fc O (f) O.H^O,.

Amo (a) Na = 32.4%: S = 22.5%; O = 45.1%.

Chemical Calculations

8. Calculate the empirical formulas iroi
positions are given below.

(a)

Na = 29.11%
S = 40.51%
O = 30.38%

(b)

C = 62.02%
H = 10.42%
O = 27.56%

the compounds whose weight com-
ic) (d)

11 = 3.70% Ca = 65.93

C = 44.44% P = 34.07

N = 51.85%

Am: (a) Na 2 S 2 0..

9. Calculate the weight ot coppei in 500 g of each of the following compounds:

(a) CuCl 2 ; (b) Cu(NO,). ; (c) CuSo 4 • 5H,0; (d) CuoFe(CN),.

Atw: (a) 235 g; (b) 169 g; (c) 127 g; (d) 93 g.

10. What weight ot Fe 2 O ti will result horn the rusting of 35.0 g of iron?

11. What weight of potassium chlorate is lequired to produce 200 g of oxygen

by the following reaction. 2 KCIO, — » 2 KC1 4- 3 0 2 ? Ans: 510 g.

12. (a) What volume ot C0 2 , measured at 20°C and 720 mm pressure, will be
produced by the leaction of 530 g of sodium carbonate, Na,,CO„ accoiding
to the following reaction. Na 2 C0 3 -f 2 HC1 2 NaCl +■ H 2 0 4- C0 2

(b) What weight of NaCl will be simultaneously produced?

13. To produce a volume of 100 liters of H._> at 67 °C and 680 mm pressure
what weight of Mg will be lequired bv the reaction of (a) Mg 4- HoSG 4
MgS0 4 4- H, (b) Mg + 2 HC1 -» MgCl> + H 2 ?

14. How many molecules of H 2 are liberated by the reaction of 1.0 g of Mg
with excess H.,S0 4 ?

* 4 Ans: 2.5 x 10

15. For the oxidation of ammonia, 4 NH 3 4- 5 0 2 4 NO 4- 6 H s O,

if 20 liters of NH-, and 0 2 each are mixed and reacted, calculate (a) the volume
of the final mixture (b) the number of moles of NO produced. All substances
are gases.

16. For the reaction of 15.0 liters of H 2 , taken at STP, according to the equation:
Ny 4- 3 H 2 — ► 2 NH.j, calculate (a) the number of moles of NH 3 pro-
duced (b) the weight of NH 3 produced (c) the volume of NH 3 produced at
STP (d) the volume of N 2 required at STP.

17. What weight of H 2 S0 4 can be produced from five tons of sulfur?

Ans: 15.3 tons.

18. To produce 100 g of an artificial sapphire, A1 2 0 3 , what weight of A1 is
required?

19. From 53 g of sodium carbonate, Na 2 C0 3 , what weight of (a) NaCl and (b) C0 2
can be produced?

20. The boiling point of benzene, C 6 H 6 , is 80°C. What volume would be oc-
cupied by 3.90 g of benzene vapor at 100 °C and 760 mm pressure?

21. What is the weight of 86.6 liters of H 2 S gas at 200 °C and 400 mm pressure?

22. A gaseous compound contains 0.84% hydrogen, 10.05% carbon, and 89.10%
chlorine. At 150°C and 760 mm pressure the density of the gas is 3.43 g/1.
What is the molecular formula of the compound?

Ans; CHC1 3 .

23. A gaseous hydrocarbon contains carbon and hydrogen in the ratio of 4 : 1
by weight. One liter of the gas at STP weighs 1.356 g. What is the mole-
cular formula of the gas?

Ans: CoH 6 .

Chemual Calculations

24. The atomic weight of a solid metallic element is 45 0 \nal\siv
6.90 g of an oxide of the element contain 4 50 g of the elements
is the valence of the element? (b) What is* the formula of the o\uh

25. The equivalent weight of magnesium is 12.1. (a) What weight of

be produced by the action of 60.5 g of Mg with excess HCP do
weight of CL, would 12.1 g of Mg combine? An a.

dmus that
What

(a) 3.
H, would
M ith what
(a * 5.0 g.

Energy Changes In
Chemical Reaction

Chemical and physical changes invariably are accompanied by energy
changes. Many occur with the liberation of heat. Thus the burning of coal,
a chemical change involving the union of solid carbon and gaseous oxygen
from the air to produce gaseous carbon dioxide, gives off heat. Reactions
or changes in which energy is liberated are called exothermic. The term in-
cludes changes in which the energy may be liberated in forms other than
heat, such as radiant or electrical energy. Other reactions occur with the
absorption of energy, that is, heat or some other form of energy is consumed
when the process occurs. When ice melts to liquid water, heat is absorbed
from the surrounding air; the temperature of the air would be lowered if no
other source of heat were present. Reactions or changes which absorb energy
are called endothermic .

1. Conservation of Energy. Careful study of energy changes and common
experience have established the fact that energy can be neither created nor
destroyed. In any physical or chemical process energy is conserved. In a
closed system, one which is isolated from its surroundings, the sum total of
energy remains constant, though the energy may be converted from one
kind to another. This generalization is known as the Law of Conservation
of Energy, or the First Law of Thermodynamics, the science which deals
with the transformations of energy into its various forms. Heat thus can
be converted into mechanical energy in a steam engine and mechanical
energy can be converted into electrical energy by means of a generator.
The Law of Conservation of Energy constitutes an energy balance sheet
for such processes. Though the idea of conservation of energy was accepted
vaguely by 1840, the first statement of the general principle was made by
a German physician, Julius Robert Mayer, in 1S42. The concept first occurred
to him when he noted that the venous blood of sailors in the tropics was a
bright red due to the lesser oxygen requirement there to maintain normal
body temperature. Later James Prescott Joule, an Eriglish physicist who had
studied under John Dalton, quantitatively measured the conversion of work
into heat in several different ways and gained final acceptance of the

principle. Though the Law of Conservation of Eneigv has nevei been
proved mathematically to be true, it is accepted from everyday experience
as a cardinal principle of physical science. The l mted States Patent Office
will reject a priori any patent application which pi opuses to conti a\ene this
law, e.g., a perpetual motion machine.

More accurately the law should be known as the I’onseivatnm of Mass-
Energy because it holds strictly only tor systems where the mass is unchanged.
Since mass and energy aie irterconvei tible it is possible tor one In appear
or to disappear at the expense of the othei but the sum total of both mass
and energy will remain constant. In ordinarv chemical and phvsieal changes,
however, any change in mass is so slight as to be negligible so the Law of
Conservation of Energy does apply in pi action! woik.

The physicist defines energy as the capacity tor doing work. Energy
may be either kinetic or potential. Kinetic energy is energy due to motion,
like the energy of a moving projectile or running water Potential energy
is stored or latent energy, like the energy an object has bv vutuc ot its
position. Potential energy and kinetic energy are transformable from one to
the other. A rock resting on a precipice has potential eneigy, if it falls over
the edge its potential energy becomes kinetic energy. During its fall, its
potential energy at any point has decreased by an amount just equal to its
gain in kinetic energy. Raising the rock to its original position would restore
to it its initial value of potential energy. The potential energy inherent in
chemical substances, and which can he released in chemical reactions, is
called chemical energy. Dynamite has potential chemical eneigy. The explo-
sion of dynamite may produce sound energy, radiant energy, mechanical
energy in that the atmospheric gases are expanded, and perhaps oven elec-
trical energy with the proper machine to harness and transform chemical
to electrical energy. The Law of Conservation of Energy states that the de-
crease in chemical energy due to the reaction of the dynamite is equal to
the sum of the several forms of energy released by the chemical change.
The change of one form of energy into another is not always 1 (K) percent
efficient. For example, when mechanical energy is converted into electrical
energy, some of the mechanical energy is changed into heat energy because
of friction. This constitutes a loss or degradation of energy in that it is not
available for doing work, but the energy does not disappear. The First Law
states that the mechanical energy initially available equals the sum of the
electrical energy and the heat energy produced. Another example of an
incomplete energy conversion is the change of electrical energy mto radiant
energy in an incandescent lamp; the resistance of the filament converts some
of the electrical energy into heat, an engineering loss if the most efficient
conversion of electrical energy into heat is desired, but the energy balance
required by the Law of Conservation of Energy is maintained. A prime
function of engineers is to devise means of making such energy transforma-
tions into the desired form more efficient.

During the last few decades the concept of energy has become the dom-
inant viewpoint in physical science. Common usage has given to the term
energy an aura of familiarity, yet energy is truly a philosophic and mystic
concept. Even the definition of energy does not describe it as something

Energy Changes in Chemical Reactions

tangible but only as a capacity. We cannot hold a bottle of energy per se in
our hands to be shown about like an ordinary material object, yet the con-
cept is as real as is the page before you.

Whether a process will occur, or a chemical reaction proceed, can be
interpreted from the viewpoint of energy. Suppose you have a ball in your
hand and let it go. It falls. Why? To say that gravitational attraction pulls
the ball down does not tell us much fundamentally. Obviously the ball at a
higher level possesses more potential energy than does the ball on the ground,
so that it is more pertinent to say that the ball, in falling, went from a
higher energy level to a lower one. It is a matter of common experience
that water, of its own accord, does not flow uphill and that heat does not
spontaneously flow from* a lower temperature to a higher temperature.
Water can be made to flow uphill and heat from a colder to a warmer ob-
ject but it requires the intervention of a machine, a pump, or a refrigerator
which will expend energy and do work. The principle common to the two na-
tural events is that, in any natural process, that is, spontaneously, systems go
from higher to lower energy levels. The condition of lowest energy corresponds
to the one of greatest stability, and is the state of equilibrium. Hence a
corollary statement would be that all systems tend to attain the state of
equilibrium.

This viewpoint can be extended to chemical reactions. The direction in
which a chemical reaction will proceed, if at all, is but another facet of
the general principle that systems tend to attain the lowest possible energy
available to them. If we conceive that the reactants and the products of a
chemical reaction exist at energy levels (which have no relation to gravita-
tional levels ) , whether a reaction will proceed spontaneously in one direction
or the other depends upon whether the “reactants” (which are written on the
left side of a chemical equation only by convention) or the “products” exist
at a lower level of energy. A chemical reaction will proceed spontaneously
to form products only if the reactants have a higher energy than do the
products. In other words, chemical reactions proceed in order to attain a
state of equilibrium and the driving force behind the reaction is the differ-
ence in energy between the reactants and the equilibrium state.

The particular energy which governs the direction of chemical change is
called the Free Energy (symbol F) by the chemists. It is closely related to
the heat energy liberated or absorbed in a reaction but is not equal to it.
Though the heat energy is not an exact criterion of the spontaneity of a
reaction, in general, reactions which are highly exothermic proceed spontan-
eously and rapidly whereas endothermic reactions do not unless some external
energy source is applied.

These concepts are the basis of the Second Law of Thermodynamics. It is
the Second Law which makes heat energy proportional to the absolute
temperature, and which sets an upper limit to the efficiency of a thermo-
dynamic process or a heat engine. In its mathematical treatment, the Second
Law introduces the concept of entropy, (symbol S). In any process or
chemical reaction involving a heat transfer, the change in entropy, AS is
given by

( 1 )

where

q is the quantity of tiansfeuecl at
the absolute temperature, T

A is a symbol commonk used to denote
a change in a variable

Consider as a model an idealized, hictionless engine operating on a
fluid in a cyclic process, a series of steps through which the s\stem is hi ought
back to its initial state. A quantity of heat, <p, is taken in at a higher
temperature, T c , and an amount of heat, q,, is dkchurgvd with the exhaust
fluid at a lower temperature, T,. The work, w, clone In this ideal engine
equals the quantity of heat “used* by the engine and is equal to the dif-
ference, (q 2 - qi). The engine efficiency is given W

( 2 )

Efficiency =

work done
heat taken in

T,-T

The value of the efficiency calculated by this equation is tin* maximum
possible efficiency attainable by any actual heat engine operating between
the temperatures, T> and T^ An actual engine, e.g., a steam engine or gasoline
internal combustion engine, can only approach this maximum value in ef-
ficiency and its efficiency will always be less because of eneigv losses
through incomplete utilization of the available heat energx and iuction.
Theoretically, 100 percent efficiency can be achieved only by discharging
the exhaust fluid at absolute zero. This would mean that all the heat
energy had been absorbed from a given sample of fluid.

The entropy change, AS, is also a criterion of the spontaneity of a process
and so is related to the change in free energy, AF. For a reaction occurring
at constant pressure and at constant temperature,

(3) AF ~ AH - TAS where AH is the heat energy emitted or absorbed

in the reaction

In any spontaneous or naturally occurring process, the free energy change,
AF, is negative but the entropy change, AS, is always positive, that is, the
absolute value of the entropy of the products is always greater than that of
the reactants.

The concept of entropy has taken on many philosophical implications.
It can be related by statistical mechanics to the degree of order or disorder
of the entities comprising a chemical system. In the throwing of a pair of
dice, a “seven* is more probable than a “two” because there are more ways
in which a “seven” can be made up than a “two.” The greater the number
of ways in which a system can be made up, tire larger is its entropy. This
relationship is given by the Third Law of Thermodynamics,

(4) S = k loge W where S = the absolute value of the entropy of a system

k = the Boltzmann constant, L38 X lC ltJ erg/
molecule degree 1

W = the number of ways a system can be made
up

1 The Boltzmann constant is the gas constant per molecule, H/N.

Energy Changes in Chemical Reactions

Thus the configuration “seven” has a greater entropy than that of a “two.”
In an assemblage of perfectly ordered entities, where by definition perfect
order admits of but one way of arrangement, the absolute value of the
entropy of such a configuration is zero. A pure crystalline solid at absolute
zero is perfectly ordered and has an entropy of zero. When melting occurs
the random motion of liquid molecules has a greater degree of disorder than
existed in the solid. There are more ways in which the liquid particles can
be arranged to make up a given liquid state than was the case with the solid
so that the entropy of the liquid is greater. As the liquid is heated further
above its melting point the randomness of its molecular motion increases
and so, too, its entropy. Because the gaseous state has the greatest degree
of disorder, the entropy of a gas is greater than that of the liquid from which
it was formed.

From a statistical viewpoint the state of equilibrium, or the state of
maximum stability, is one which can be made up in more ways than any
other state. Hence the equilibrium state corresponds to one of maximum
entropy. Since all systems tend to attain equilibrium, the entropy increases
in any spontaneous process, physical or chemical. The magnitude of the
entropy change is a criterion of the tendency of the process to proceed. So,
too, a mixture of two gases does not separate spontaneously into its pure
components. The random mixing has a greater degree of disorder and a
greater entropy than the sum of the entropies of the separate gases.

2. Measurement of the Energy of a Chemical Change. The energy of a
chemical reaction is determined experimentally from the amount of heat
given off or absorbed during the interaction of definite quantities of the re-
acting substances. The unit of heat is the calorie (cal), which is defined
as the amount of heat required to raise the temperature of one gram of
water one degree centigrade, specifically from 14.5° C to 15.5°C. Frequently
the kilocalorie (kcal) or kilogram calorie is used; it is equal to 1000 calories.
The heat unit used by engineers is the British Thermal Unit (B.T.U.). It
is the amount of heat required to raise the temperature of one pound of
water one degree Fahrenheit. One B.T.U. is equal to 252 calories. Other
energy units can be used instead of calories. Thus 1.00 calorie = 4.18 X 10 7
erg = 3.09 foot pound = 1.16 X 10 -6 kilowatt hour.

To measure the energy liberated or absorbed as heat in a chemical
change, the reaction to be investigated is carried out in an apparatus called a
calorimeter. Known weights of the interacting substances are allowed to
react in an insulated vessel so that all the heat produced by the reaction
is absorbed by a known weight of water. The temperature of the water
is taken before and after the reaction, and from the increase in the tempera-
ture of the water the number of calories is calculated. The quantity of heat,
q, given up or absorbed by any substance due to a change in its temperature
is given by

(5) q = C X ni X t where C = the specific heat of the substance (see

following )

m = the mass of the substance
t = the temperature change of the sub-
stance

The specific heat of a substance is defined as tin* numlwi ol calories
required to raise the temperature of our .main nt a Mibst amv one degree
centigrade. It follows from the definition ol the euloiie that the specific
heat of water is one calorie pci giam pei degree l he specific heat of a
substance can be calculated from Equation 1 b\ introducing a known
quantity of heat, q, into a given mass, m, and meusmmg the temperature
rise, t, thereby produced. Values ol specific heat van with the temperature
and also depend upon whether they are measured urnlci conditions of con-
stant pressure or of constant volume In detcnnmmg specific heats
under constant pressure conditions, the volume of a substance mucuses
when heated and thus work is done against the external pressure.
A portion of the heat input is used to do tins woik, leaving a lesser
amount (than under constant volume conditions whcic no external work
is done) to increase the kinetic eneigv of the molecules and so laise
the temperature. Under constant piessure conditions, then, a gieutct quantity
of heat is required to raise the temperature of a given mass ot a substance
than under constant volume conditions, and so the specific heat at constant
pressure (C p ) is always greater than the specific heat at constant volume
(CJ. Appendix V lists the specific heats of some common substances.

3. Thermochemical Equations. Chemical equations can be used, to express
not only the stoichiometric relationships ol the reacting substances hut also
the energy change of a reaction. Thus

(6) 2 H, + O, 2 11,0 i- 136,634 cal

means that two moles of hydrogen combine with one mole of oxygen to
form two moles of water with the liberation of 136.634 calories The heat
specified with an equation pertains to a reaction as written. Koi the i ruction
h 2 + y 2 cc -* h,o the heat liberated would be 68,317 ealones because
this equation refers to the reaction of only half the quantities ol materials
as was used in the first equation, that is, one mole ol hvdmgen will
react with half a mole of oxygen (16 g or 11.2 liters at STT) to form one
mole of water and liberate 68,317 calories. Tin* value of 68,317 calories
may be called either the heat of combustion of hydrogen or the heat of
formation of water, depending upon ones viewpoint.

In addition, the quantity of heat evolved or absorbed m a given reaction
depends upon variables other than the chemical nature of the leactants:
the physical states of the substances, the temperature of the reaction, and
the manner in which the reaction is carried out Thus whethei the water
formed in the reaction above is liquid or gaseous will affect the magnitude
of the heat liberated by an amount equal to the heat of vaporization <4 water.
The states of the substances in a thermochemical equation are indicated
by placing after each formula the pertinent designation; (.0 tor solid, (/)
for liquid, and (g) for gas or vapor. More exactly then

<7) 2 H«(g) + Oqg) 2 H,0(7) ~r 136634 cat

In some cases it is even necessary to specify which solid allotropc is the
reactant involved because different heats of reaction result from the reaction

Energy Changes m Chemical Reactions

of the different allotropic forms. Thus for the two forms of sulfur, rhombic
and monoclinic,

S (rhombic) + 0 2 (g) S0 2 (g) + 70,960 cal
S ( monoclinic ) + 0 2 (g) S0 2 (g) + 71,030 cal

Chemical reactions can be carried out under conditions of constant pres-
sure or of constant volume, at the option of the experimenter. The former
is the more common state of affairs. An experiment carried out in the open
on the laboratory bench is done under a constant pressure of one atmosphere.
If a reaction is carried out in a sealed container so that the volume of the
reacting system is fixed, it is done under conditions of constant volume.
The energy of a chemical reaction differs somewhat under these two
conditions. To specify critically one or the other, the chemist uses the
symbol AH for heats of constant pressure processes and AE for the heats
of constant volume processes; AH and AE are thermodynamic variables known
as the enthalpy and the internal energy , respectively. The symbols AH and
AE are written separately after a thermochemical equation. By convention,
for an exothermic reaction since the chemical system loses heat, both
AH and AE are given negative signs; for an endothermic reaction since the
system gains heat, these variables are given positive signs.

Since the energy of a chemical reaction tends to change the temperature
of the reacting system, values of heats of reaction are given on the assump-
tion that the temperature of the products after reaction is brought back to
that of the initial reactants so that the reaction can be assumed to be iso-
thermal This temperature is generally specified in degrees Kelvin as a
subscript to AH or AE; thus AH 298 or AE 29 3 for a reaction which occurs at
25° C. Thus a complete thermochemical equation for the formation of water
by the reaction of hydrogen and oxygen at 25 °C and at constant pressure
would be

(8) 2 H 2 (g) + 0 2 (g) -* 2 H 2 0(!) AH„ a = -136,634 cal

Under constant pressure conditions the volume of a reacting system may
change and pressure-volume work, or mechanical work, will be done against
the external pressure. This mechanical work is given by P(AV), where P
is the constant external pressure and AV the change in volume of the system.
Where only liquids and solids are involved, the change in volume, AV, is
practically zero and no mechanical work is done. Where gases are involved,
and there is a change in their number of moles, P(AV) may be appreciable.
Since some heat energy is required to do this mechanical work in a constant
pressure process, unlike a constant volume process where AV = 0, AH and
AE will not be equal. The difference between the two is the mechanical
work done in the constant pressure process, or P(AV).

(9) AH = AE + P(AV)

If the process is also carried out at constant temperature (isothermal),
P(AV) = BT(An) and

(10) AH = AE + RT(An) whore K sc? the molar g is constant, its value

in calorics is 1 9 ST cal mole deg
T the absolute iemperatuu\ °K
An ~ the numbei of moles of gaseous
products miuiis the niimKr of
moles of titLsrom icactauts.

For Equation 8 , An = [(()) - (2 ■* D], or -3. Foi reactions wheie only
liquids and solids are involved, An = 0, and lienee Ali and AH aie equal,
These symbols, AH and AE, can also be used for the heats involved i n
changes of state. To indicate the endothermic heat nf vuponzution for the
boiling of water at 10 O°C, 9720 cal/mole, and the exothermic heat of fusion
for the freezing of water at 0 °C, 1436 cal/mole, we may wntc

H 2 0 (l) H 2 0(g) AH ltl =* . 9720 cal

H 2 0 (l) -+ H s O(s) AH . 7 1 =-. - 1 136 cal

A thermochemical equation enables one to calculate how much heat
would be evolved or absorbed by the reaction of a given quantity of a *sub-
stance.

Example 1: How much heat will be produced when 3“ 0 inarm of non rust,
according to the following thermochemical equation?

4 Fe(s) + 3 0 2 (g) -> 2 Fe/3, AH,„, ~ -381 t heal

Solution: The equation states that 381,400 calories are given off when four
gram-atomic weights of iron rust. Therefore the rusting of 33.0 g of He rust, the
amount of heat produced will be

35.0 g

55.8

GAW

^ 381.4 keal
X 4 GAW

= 59.6 keal

The attractive forces which hold together the atoms within a molecule
constitute the chemical bond. To decompose a mole of hydrogen molecules
into hydrogen atoms, H 2 -» 2 H, 104 kilocalories a re required. This value,
104 kcal/mole, is said to be the H — H bond energy in the IF molecule. A
given type of bond between two atoms in a molecule has a specific bond
energy. In a reaction chemical bonds are broken and others are formed,
To break a bond an input of energy is required. In the formation of a
bond, energy is released. The net energy of a chemical reaction is the differ-
ence in the energies required to break the various bonds (since more than
one bond may be broken) and the energies released in the formation of
new bonds. In some cases it is possible to estimate heats of reaction from
a knowledge of the bond energies of the reactants and the products.

4. Hess’ Law of Constant Heat Summation. A direct consequence of
the Law of Conservation of Energy is the generalization discovered in 1840
by a German professor of chemistry, Germain Henri Hess, who thereby
became virtually the founder of thermochemistry. Hess pointed out that the
heat evolved in a chemical process is the same whether it takes place in
one or in several steps. An equivalent statement would be that the quantity
of energy liberated or absorbed during a chemical or physical change
depends only upon the initial and final substances and their conditions, and
not on the intermediate steps or the path of a process. Mathematically the

Energy Changes in Chemical Reactions

energy of a system is said to be a point function and its value depends only
upon the present state of a system and is independent of its past history. 2
Hess' Law is merely an alternative statement of the Law of Conservation of
Energy.

These statements permit chemical equations to be handled in accordance
with the rules of elementary algebra. They can be added, subtracted, reversed,
multiplied by a number, and the heat of the reaction will be affected accord-
ingly.

The validity of Hess' Law can be illustrated by the equality of the
heats of formation of barium oxide, BaO, when prepared in two different ways.

Direct preparation:

(a) Ba(s) + % 0 2 (g) BaO(s) AH = -124,400 cal

Indirect preparation:

( b ) Ba(s) -f 0*(g) BslOj(s) AH = -141,600 cal

(d) Ba (*) + %O t (g) BaO(«) AH = -124,400 cal

Equations b and c are treated as simultaneous algebraic equations; b and
c are added and common terms are cancelled. The result is equation d which
is identical with equation a for the direct method.

On paper, if not on the laboratory bench, a chemical reaction can be
reversed; the heat of reaction will be unchanged in magnitude but its sign
will change. Thus, for the decomposition of two moles of water, the reverse
of equation 8:

2 H z O(Z) 2 H 2 ( g ) + 0 2 (g) AH* = +136,634 cal

It is obvious that the heat of vaporization and the heat of condensation
are equal, but opposite in sign; so likewise are the heat of fusion and the
heat of solidification.

It is possible to determine indirectly the heat of a reaction, even though
it is impossible to carry out the reaction and measure the heat experimentally,
provided that data for other pertinent reactions are available. For example,
it is impossible to measure the heat evolved when carbon burns solely to
carbon monoxide, CO. But the heats evolved when carbon burns to carbon
dioxide, C0 2 , and also when CO bums to C0 2 , can be determined experi-
mentally with good precision. The thermochemical equations are:

(e) C (graphite) + 0 2 (g) — » C0 2 (g) • AH — -94,052 cal

(f) CO (g) + y 2 0*(g) C0 2 (g) AH = -67,636 cal

If f is subtracted from e, then

(g) C (graphite) + % O z (g) CO(g) AH = -26,416 cal

2 If this were not so, conservation of energy would not exist. If a certain number
of calories were liberated in going from a State A to a State B, and if a lesser amount
were required to return the system by another “path” to the initial State A, then in
a repetitive or cyclic process, an operator would accrue an excess of energy which could
be used for other purposes. Very simply the operator would be getting something for
nothing! Unfortunately the same amount of energy is required to* return to .the initial
conditions of State A as was evolved in going from A to B!

Depending upon the type of the chemical icaetion, heats nf icactions
may masquerade under many names, if the reaction is a combustion, the
heat of reaction is called the heat of combustion. Other heat terms used
are: heat of formation , vxhich refers to the heat liberated or absorbed when
a unit amount (gram or mole) of the substance is burned from its elements,
heat of solution y which is the heat liberated or absorbed when a unit amount
of substance dissolves in a specified amount of sobent. heat ol dilution)
heat of neutralization, heat of hydration , etc.

5. Heat of Formation. The heat of formation of a compound is defined
as the amount of heat in a reaction in which one mole of the compound
is formed from its elements, all substances being m then >i tin lard states.
For thermochemical purposes the standard state is defined as one atmosphere
pressure and 25°C. The elements, in the stable hum in which exist at
the standard state, are arbitiarily assigned a \alue of zero loi their boats of
formation. By definition then, the heat of formation of liquid wab*r is
-68.3 kcal/mole; that for CCb is -94.1 kcal/mole and fm (Ois -2b 4 le al mole.
The heats of formation of a great number of substance* are known, a few
are given in Table 7-A.

If the heats of formation of all the substances in a chemical inaction are
known, the heat of the reaction can be calculated. The heat of inaction is
the sum of the heats of formation of the product* nanus the sum of the
heats of formation of the reactants. Data such as m Table 7 A cun be used
taking proper consideration of the number of moles of each substance in
the balanced chemical equation.

Table 7-A

Standard Heats

of Formation, AH f , kcal/mole a

i 25' C & 1 atm

h 2 o(1)

- 68.32

N'aCl(s)

- 9S.28

HjO(g)

- 57.80

NaOH(s)

-!02 0

HF(g)

- 64.2

H..so*(n

-198.0

HCl(g)

- 22.06

CaO(s)

- 151.0

HBr(g)

- 8.66

CaCO,(»)

-288 -1

HI(g)

4- 6.20

Fc,0,(«)

-190,7

H 2 S(g)

- 4.82

Al,0,(.s)

-399 i

NH 3 (g)

- 11.04

CH,(g)

- 17.89

CO(g)

- 26.42

C,H,(g)

- 20.24

co 2 (g)

- 94.05

C,H,(g)

- 24.82

C t H s {g)

4 54.19

C 2 H 4 (g)

4 12,58

4 11.72

NOTE: A negative sign for AH f means that the heat of formation is exothermic,
a positive sign represents an endothermic heat of formation.

Example 2; .Calculate the heat of combustion of methane, CH* at 25*0. The
•reaction is

GH*(g) + 2 0,.(g) -* C0 2 (g) + 2 H a O (f)

Solution The sum of the heats of formation of the products is:

AH,

for one mole of CO^g)

= - 94.05

AH f

for two moles of is (2 x -68.31)

= -136.62

AH f

for the products

= -230.67

AH f

for one mole of CH^g)

= - 17.89

AH,

for 0,(g)

= 0.00

AH,

for the reactants

= - 17.89

The heat of the le act ion, or more accurately, the enthalpy change, AH 298 ,
= (-230.67) - (-17.89) = -212.78 kcal

QUESTIONS

1. Define and illustrate with specific examples: (a) exothermic reaction (b) endo-
thermic reaction (c) heat of reaction (d) heat of formation (e) heat of solution
(f) molar heat of condensation (g) enthalpy change (h) entropy change.

2. Using the data in Appendix V, derive a relationship for the metallic elements
which would lead to the detei ruination of the approximate atomic weight.

3. What dictates whether or not a process will occur? Discuss the criteria of
spontaneity.

4. State the Law of Conservation of Energy. Illustrate, using as examples, a
mechanical process and a chemical process. List other properties of matter
which are conserved.

5. State and illustrate Hess’ Law. What broader generalization includes Hess’ Law?

6. What are the uses of thermochemical equations?

7. In writing thermochemical equations, why is it important that the physical
state of the reacting substances be indicated?

8. Write an equation which represents the heat of formation of sulfuric acid,

h,,so 4 .

9. Under what conditions of temperature should a heat engine operate for maxi-
mum efficiency?

10. Explain why the entropy increases during the process of melting.

11. Which has a higher value of entropy, mercury liquid or mercury vapor at its
normal boiling point? Explain.

12. Calculate the number of calories in (a) one liter-atmosphere (b) one BTU
( c ) one watt-hour.

13. How much heat is required to raise 100 g of iron from (a) 100°C to 150°C

(b) 20° C to 70°C? An* (a) 530 cal

14. A piece of metal weighing 50.0 g and at a temperature of 140° C was

placed into one liter of water at 25.00 °C. The temperature of the water
rose to 25.55 °C. Assume the density of water is 1.00 g/ml and calculate the
approximate atomic weight of the metal. Ans: 62

15. 120 g of a metal whose specific heat is 0.025 cal/g deg are heated to 100°C
and then added to 40.0 g of water initially at 20.0 °C. What will be the final
temperature of the metal-water mixture?

16. Calculate the heat, in calories and in BTU, required to change twenty grams

of ice at 0°C to steam at 100°C. Ans: 14,380 cal; 57 BTU

17. When one pound of fuel oil is burned, 18,000 B'H aie pmduoed What
weight of oil is required to heat 1,000 gallons ni wain imm U) F to 190 U F?

An \ 66 6 lb

18. Calculate the enthalpy change foi the ti.mstoi nudum from ihoaibie snlhu to
monoclinic sulfur.

19. The heat combustion of one gram-atom of caibnn (diamond) is DIB heal,
and that for carbon (graphite) is 94 8 keal Calculate the beat oi the miction

C (graphite) — * C (diamond)

20. Calculate the heat of sublimation at 0 c C hu H , 0 {\) — » > IFOuF, if at 0 C.

H 2 0(s) -v H.O(I) AH = 4-1,4 40 eal

H a O(l) -4 H,0(g) AH =. A 10,370 cal

21. Calculate AH for the reaction. 2 Al 4- Fe ,0 , AFC). t- 2 Fo

22. For the reaction 2 H 2 0(g) — ► 2 H 2 (g) 4- CK(g) AH - -115 6 keal;

calculate AE.

23. At 25°C the heat of formation, AH f in heal/ mole, is -68 3 fm IFOfJ) and
-57.8 for H 2 0(g). Calculate the heat of \apon/ahnn tor HO at 25 (

24. (a) Calculate the heat of combustion of ueetslene, ( IF, at 25 (\

2 C 2 H 2 (g) A 5 0..(g) — 4 CO .(«) -t 2 FI .()(!'»

Given: 2 C(graphite) A H.,(g) CTF(g) AH - <-5h2 heal

C( graphite) A 0 2 (g) C0 2 (g) AH v. -9 J, I heal

2 H 2 (g) A 0 2 (g) 2 H.O<n AH = -6S.8 heal

(b) How much heat would be evolved when t\sent\ grams of <* IF burn?

Ansa (a) -310.6 heal mule C .IF

25. (a) Calculate the heat of combustion of propane, C il M at 25 C, to CO, and
H 2 0 (b) Assume that air initially at 25 C is the source of tin* uwgcn, and
using the specific heats in Appendix V, calculate the tempeiatme pi educed in
burning one mole of C { H H .

26. Calculate the heat of formation of butane, C,H to , at 25 C m hcal/mole, given
that 2 C,H 10 (g) A 13 0 2 (g) -> 4 C0 2 (g) A 5 H.OiO AH = -688 0 keal

Am.v. AH f - —29 S hcal/mole

27. From the data in Table 7-A calculate the heat of h}drugenution of acetslene,
C 2 H 2 . The reaction is C 2 H 2 (g) A 2 H 2 (g) — * CJi, ( (gF

28. Calculate the enthalpy change, AH, and the entropy change, AS, for (a) the
boiling of one mole of water at 100°C; (b) the condensation of one mole of
steam at 100 °C.

29. From the Ideal Gas Law, PV = nRT, derive P(Av) = RT(An) foi a gaseous
reaction at constant pressure and constant temperature.

The Kinetics of
Chemical Reactions

Chemical kinetics is the study of the rates and the mechanisms of chemical
reactions. A balanced chemical equation merely states the initial reactants,
the final products, and their over-all stoichiometric relationships. It gives no
indication that a number of intermediate steps or a succession of elementary
reactions may be involved in the transformation of reactants into ultimate
products. Chemical kinetics is concerned with the intermediate steps or
mechanism of a reaction, the rates of these steps, and the factors such as
concentration, temperature, and catalyst which affect the rate. Some chemical
reactions seem to take place instantaneously while others proceed at an
extremely slow rate. The rate of a chemical reaction is measured in terms of
the concentration change per unit time of some reactant or product. Which
substance is chosen is a matter of experimental convenience since the reactants
and products are linked by the stoichiometric relationships given by the
balanced chemical equation. Experimentally the rate can be determined by
measuring the change in some property of the reaction system. Thus, if one
of the substances, reactant or product, is colored, the rate can be measured
in terms of the change in color intensity of the system. Or if one of the
substances is a gas, its change in volume or pressure will indicate the rate
of reaction.

1. Factors Affecting the Rate of Chemical Reaction. Many factors in-
fluence the rate of a chemical reaction. In general, the rate of reaction depends
upon:

(A) The Nature of the Reacting Materials

Everyday observation attests to the fact that iron reacts with the oxygen
in the air whereas gold will not react under similar conditions. It is apparent,
though sometimes forgotten, that the rate of a chemical reaction depends
upon the nature of the reacting substances.

(B) State of Subdivision

Substances must be in contact to react. Reactions involving a solid
and another substance can take place only at the surface of the solid. The

I fu' Kiiu tu \ >0 ('Sumu'ul fU actum

material within the body of a solid cannot contact an rxtcmal reactant and
so cannot react with it. If the solid is ground to ^ fine poxxdt a, tin* surface
area is increased, more molecules can then ir.ut m a gn »*n tunc and the
reaction will proceed at a faster rate. The finer the state of subdivision the
more rapidly will a reaction occur Indeed eetfuin fineh pexxdeied solid
substances react so rapidly that they constitute an explosive ha/nd. for
example, dust explosions, and the surging flame pmdneed In a magicians
throwing lycopodium powder into a small flame. Out* lenson w h\ a chemist
frequently prefers to carry out chemical icactions m solution is that in
solution the ultimate limit of subdivision i> readied because the leactants
are subdivided to molecular dimensions.

The rate of a reaction is proportional to the comentiution ol eaeh of the
reacting substances. This statement is known as the Law ul Mass \et:nn and
was first proposed in 1853 by the Norwegian chemists, Cato Maximilian Culd-
berg and Peter Waage. For a gaseous substance, concentration max be re-
placed by pressure since the concentration of a gaseous substance xaries
direct!} as its pressure. By concentration is meant the amount j)ci unit i olume
and not the total amount of a substance. This can be illustrated h> the re-
action between zinc, Zn, and sulfuric acid, 1 1 .SO

Zn + H_.SC), -v ZnSO, - H

To similar pieces of Zn m two containers add 50 ml and 500 ml »<*spectivcly
of the same dilute H_.SO,. The total amounts of fCSO s ait' m the ratio of
1 to 10 but the concentiations of FLSO, arc the same. It xx ill bo ohsoixed that
the reaction rate in both cases is the same. If, hovvoxer, to similar pieces
of Zn in two containers we add the same volume of the dilute HSOj and
then dilute one with water, the action of the diluted acid upon the Zn is
slowed. The same amount of acid is present in both cases but the concentra-
tions are different. Any unit of mass or volume can be used but chemists
commonly express concentrations in moles per liter.

(D) Temperature

The speed of a chemical reaction increases with a rixt in temperature.
In general , a rise of 10°C roughly doubles the velocity of a chemical inaction.
A reaction that has a definite rate at 20°C will proceed about txxicr as fast
at 30°C, about four times as fast at 40°C, and about 8,000 times as fast at
150°C. Reactions whose rates at ordinary • temperatures air negligible may
become extremely rapid at high temperatures. The rusting of iron, the oxida-
tion of carbon or coal, the coagulation of albumin m the preparation of a
hard boiled egg all are reactions which are accelerated at elevated tempera-
tures. Some reactions require extremely high temperatures to give any evi-
dence of chemical change. Water does not begin to dissociate mto hydrogen
and oxygen until heated to about 1000° C.

(E) Catalysis

Frequently the addition of an "‘apparently inert” substance has tin* effect
of increasing, and sometimes of decreasing, the rate of a chemical reaction.

The Kinetics of Chemical Heart ion

The addition o t a minute quantity of solid manganese dioxide, Mn0 2) to
heated potassium chlorate, KC10 { , markedly accelerates the production of
oxygen It appears that the MnOj does not enter the reaction because the net
reaction, 2 KC10< — » KC1 -j- 3 O j5 is unaltered and the MnO a can be recovered
with no loss in its mass. Such a substance which changes the velocity of a
chemical reaction without itself undergoing a change is termed a catalyst
and the phenomenon is known as catalysis. Thus the action between HC1 and
0 2 is quickened by the presence of cupric chloride, CuCL; the presence of an
iodide causes hydrogen peroxide, R^Ch, to decompose rapidly into H 2 0 and
Oo, the reaction between Hj and Ch, too slow to be observed at ordinary
temperatures, takes place rapidly in the presence of platinum.

Catalysts do not initiate reactions which would not normally proceed.
They merely change the rales of chemical reactions Each catalyst is specific
for a given reaction and no single substance will catalyze all reactions though
some, such as platinum metal, will catalyze many reactions. Often a reaction
is catalyzed by one of its products. Such action is known as autocatalysis .
The reaction between potassium permanganate, KMnO*, and oxalic acid,
H.CjOi, in acid solution speeds up as the reaction progresses, even though
the concentrations of the original reactants decrease:

2 KMnO, + 5 + 3 H,SO s K 2 SO, + 2 MnS0 4 + 10 C0 2 + 8 H 2 G.

The phenomenon is due to the presence of manganese (II) ion, a product of the
reaction. If a manganese (II) compound is added to the original mixture, the
reaction will be speeded up immediately.

2. Reaction Rate Theory. From a fundamental molecular point of view,
let us now develop a theory concerning rates of reaction. In this discussion
we shall consider homogeneous reactions, or reactions in which the reactants
and products are all in one phase, though the concepts we shall develop are
also applicable to heterogeneous reactions in which more than one phase is
present. Also we shall use the term molecule in a general sense to mean any
fundamental particle of a reactant.

( A ) Collision

For simplicity let us consider a homogeneous gaseous reaction between
two molecules. There arc three prerequisites before chemical reaction will
take place. First the reacting molecules must collide. As applied to molecules
collision is a nebulous term. A physical collision between two rigid, billiard
ball-like objects is readily visualized but molecules are not rigid billiard
bails. In Chapter 12 we shall see that a molecule can be considered as
an electric field of a definite geometric pattern, and collisions between
molecules can then be construed as an approach to such proximity that there
is appreciable interaction between these electric fields.

The number of collisions which molecules make with each other in the
gaseous state can be calculated. This number is enormous so that, if every
collision resulted in a fruitful reaction, all chemical reactions would be well
nigh instantaneous. This is certainly not the case, however, since many re-
actions require appreciable time. The reaction between hydrogen and oxygen
proceeds very slowly even at 250° C, though the number of collisions between
these molecules is of the order of 10 28 per second in each cubic centimeter

1 he KinetU'* o/ ChetHtful lUactum

of gas. Consequently collision alone is not suffic ient prerequisite for chemical
reaction to occur and there must he another factor which has a hem mg upon
the rate of reaction. However, increasing the concentration of am reactant will
increase the number of collisions between reacting molecules and hence the
rate of reaction, as indicated by the Law of Mass Action.

(B) Energy of Activation

It is believed that colliding molecules do not react unless thev possess a
minimum amount of energy, called the energy of actuation { symbol AE„h
Let us consider the elementary reaction between a hvdrogeu molec ule and a
chlorine atom.

(1) H, + Cl HC1 4- H

Though neither a chlorine atom nor a hydrogen atom can luxe piolonged
independent existence they may exist for very brief periods as ft or radieah
and participate in the actua 1 mechanism of a reaction. For this reaction to
happen, a chlorine atom must collide with a hydrogen molecule and break
the bond which joins the two hydrogen atoms together A certain amount
of energy, equal at least to the energy of the hydrogen-hydrogen bond, will
be required to rupture the hydrogen molecule. A chlonne atom, having an
energy less than this value, will be unable to break the bond and miction
will not take place. Hence, in addition to collision a second preiequisite lor
chemical reaction is that the colliding molecules must possess a minimum
energy, the energy of activation. The value of this energy is different for
each specific reaction.

That there is an attractive force or bond between the atoms of a hvdrogeu mole-
cule is evident from the existence of the hydrogen molecule as a stable entity.
The energy which holds the two hydrogen atoms together is known as the hand
energy. For the H 2 molecule the bond energy is 104 kcal niokv that is, 104 kcal
are necessary to dissociate one mole of hydrogen molecules into hvdrogeu atoms,
For the HC1 molecule the bond energy is 103 kcal/ mole. From the viewpoint of
bond energies. Equation 1 could be conceived as the rupture of an H - H bond foI%
lowed by the formation of an H - Cl bond.

Even though the products of a reaction have a lower energy than the
reactants, it may be necessary first to activate the reactants to a still higher
energy level before reaction can occur. It is as if an energy wall or barrier
existed between the reactants and the products and the reactants must first be
elevated over its peak.

The course of an hypothetical reaction is represented in Figure 8.1, where
the energy of the reacting system is plotted on the vertical axis against an
arbitrary reaction coordinate which indicates the extent to which the reaction
has proceeded from reactants to products. Point R represents the energy of
the initial state, or reactants, and point P, the energy of the final state, or
products. If the path of the reaction were directly from R to P as indicated
by the line joining the two points in Figure 8.1a, there would be zero activa-
tion energy and the reaction would be instantaneous. Such a reaction would
be analogous to the release of a ball which then drops directly to the floor, the
ball held in the hand being equivalent to the initial state of the process and

The Kinetics of Chemical Reaction

the ball on the floor to the final state. If there is a sheet of plywood in the
path of the ball to the floor, however, the ball does not fall directly to the
floor. The intervening wood acts as an energy barrier and the ball can reach
the floor only if it is given sufficient energy to overcome this barrier, that is
to break through the wood. In the same way it is conceived that chemical
reactions which are not instantaneous do not proceed directly from R to P,
but that a potential energy barrier intervenes and the reactants must first
be given an increment of energy, the energy of activation, to enable them
to pass over this barrier in order to form the products of the reaction. The
value of the energy of activation is equal to the difference in energy between
the height of the potential energy barrier and the initial energy of the re-
actants (Figure 8.1b). Two reaction paths are drawn in Figure 8.1b, both
starting from the same initial state, R. Path RR'P represents an exothermic re-
action while path R R" P represents an endothermic reaction. The energy of
activation has no direct bearing on whether a reaction is exothermic or endo-
thermic. If P is below R, the reaction is exothermic; if P is above R, the reaction
is endothermic.

R represents the initial energy level of the reactants; P represents the energy
level of the products. A reaction with zero energy of activation is shown in a.
In b the full line represents an exothermic reaction and the broken line an endo-
thermic reaction. In both cases, the activation energy, AE a , is the same. R' and R"
represent the energy of the activated complex.

Figure 8.1. Energy of Activation.

Let us return to the reaction of Equation 1 and assume that a chlorine

atom with the necessary energy of activation to react approaches a hydrogen

molecule. The bond between the hydrogen atoms is not rigid so that as

the chlorine atom comes closer the hydrogen atoms separate somewhat.

When the chlorine atom has penetrated to a certain distance the three atoms

form a complex configuration in which there are tentative bonds to each other,

thus H . . • H

• •

/ he h,u< f;. -i

1 h Ki action

This configuration is called the tit'iitttii tt <‘t vipU v ano cm i espon* Is to points
R r and R" in Figiae 8.1b. Thenviftei the chlmmo alum lamj brm bond
with one of the hydiogen atoms and go oli as an IK'i mnlivim l mini other
circumstances the ehlmmv atom may mbourni so wbuh (,hi 'Ur *wo hydro-
gen atoms will come togethei again as a H_ molecule ami no u*not mu will
have taken place.

approach of a Cl atom II ... H
(2) H 2 +C1 — * \

no reaction; return to Cl
initial leactants activated
complex

leuelioii to pu tdods

> na - h

It we consider a reaction, suclr as the iever.se of Equation 1,

H + HC1 — » HL ”p Cl, we can set 1 from Figuie S Ih that the emigy of
activation for such a reaction, if exothenmc, is given bv A Kb oi the diiference
in energy between points P and R' \ P and R" foj an endothciium teartiou).
It is apparent, theiefoie, that the difference m values between the encodes
of activation foi the reverse and the foiward reactions is equal U> the net
energy absorbed or emitted m the chemical reaction

(3) AK (leaction ) “ AE., - AF/,

We have already noted that there is a distribution ot eueign s in an
assemblage of gaseous molecules at a given tempeiatuie i page 2 A Let us
arbitrarily set a value for the energy of activation, designated AK,,, and
represented by the vertical line in Figme 8 2 Then o»!\ those molecules
which ha\e energies equal to or greatei than this value, that is to the
right of this line, can react; those with energies lower lhan this value cannot
react. The relative areas under the curve to the right and to the left of
the AE., line give the ratio of the number of molecules which oui and which
cannot react, respectively. Obviously at temperature T, only a relatively small
number of molecules have sufficient energy to enable them to icuct If
the temperature is raised the distribution curve shifts as shown In the
broken line in Figure 8.2. For a given reaction, the value of tin 1 activation
energy is independent of the temperature so that at the highei temperature,
the number of molecules having the minimum energy icqmted to react
is greater and the reaction proceeds more rapidly. For a ten degree 1 rise in
temperature, the total energy of a mole of gaseous particles, as given by
3/2 RT, increases only about 3 percent (10 parts in 300), The number of
molecules having energies equal to or greater than AE, is approximately
doubled, however, as indicated by the relative areas under the two curves
to the right of the AE tt in**. For this reason the rate of reaction is roughly
doubled by a ten degree rise in temperature.

A catalyst provides a path different from that of the uncatalyzed reaction
through which the reaction may proceed to the same final products. In this
new path the energy of activation is lower than in the uncatalyzed reaction.
A door is § catalyst to exit from an otherwise sealed room. With a catalyst
present, at a given temperature a larger fraction of the reactant molecules

The Kinetics of Chemical Reaction

Figure 8.2.

Activation Energy and
Temperature.

T 3 is a temperature higher than
T The number of molecules having
an energy equal to or greater than
AE a at temperature T is given by
the crosshatched area under the T
curve to the right of AE a . At
temperature T x , this number is
given by the dotted area under
the T 1 curve. If T x is ten degrees
higher than T, the ratio of these
areas is about two to one

would have the lower energy of activation and the reaction would pioceed
more rapidly. For a catalyzed reaction the points R' and R" in Figure 8.1 b
would be lower. A possible mechanism is that the catalyst actually does
take part in the chemical reaction by foirmng an intermediate compound
from which it is ultimately regenerated. If the reaction, A -|- B AB, occurs
slowly because of a high activation energy, addition of a catalyst, C, alters
the mechanism or path of the reaction to one of lower actiyation energy
so that it proceeds more rapidly.

(4) A + C — > AC (AC is an intermediate compound)

(5) AC + B — > AB + C (regeneration of the catalyst)

The catalyst is regenerated to continue its catalytic activity so that only
very small quantities are needed, and it is ultimately recoverable.

Even though molecules collide with an energy equal to or greater than
the energy of activation, reaction may not occur. Since molecules do have
definite geometric configurations, a further requisite for chemical reaction is
that the collision must occur at a position where reaction can take place or be
initiated. Whereas it is recognized that such geometric considerations are
a factor, the structure of each molecule is specific, and perhaps quite com-
plex, so that each case must be considered individually. Hence this steric
factor will be recognized here but its further consideration will be left for
more advanced textbooks.

3. Classification of Reactions.

(A) Molecularity

A chemical reaction may be complex in that it consists of- a number of
elementary processes, each of which takes place in a single step. Reactions
may be classified as unimolecular, bimolecular, trimolecular , etc., according
to the number of molecules which reaet in a single elementary' step. The
reaction of individual molecules is a unimolecular reaction. An example is
the decomposition of nitrogen pentoxide, N 2 0 5 .

2 N.O. 4 N0 2 + 0 2

100

I In A.i/Jt tu n r/ii ini. f.'/ action

Another unimolecular reaction is the disintegration oi ludiuaotnr elements
(Chapter 48). Reactions in which two molecules collide and then ieact
are called bimolecular reactions. Equation l is an example. Othei examples are
the reaction between hydrogen and iodine \apoi to proJuu- huirogen iodide,
H 2 ■+ I, — > 2 HI, and the reaction between sodium ii> dioxide and ethyl
acetate in water solution to produce sodium acetate and etlwl alcohol,
NaOH + CH,COOC,H- — CH COONa * CILOII. Occasional a re-
action is effected by the simultaneous collision of three molecules Such a
reaction is trimolecular. Triinolecuku reactions are laie since the simultan-
eous collision of more than two molecules is impiohabh* from a statistical
viewpoint. An example of a trimolecular reaction is 2 NO * O 2 NOj,
for which the actual mechanism is NO -r NO * O. 2 NO . A bah
lanced chemical reaction gives no indication of its moleeuiantv This de-
pends upon the actual reaction mechanism. A reaction. 2 V - B — > AjB,
is not necessarily trimolecular. Rather it may be a seiies of umnmlecular
and bimolecular steps, for example,

A r- B -> AB
AB - A — A,R

The decomposition of N\Or, m Equation 3 ma> appeal to he bimolecular but
consists of a numbei of steps, the first of which is ummoleeulai .

N 2 0 3 NO, + N’O,

NO, + NO i( — NO + NO. 4- NO,

NO + NOj — » 2 NO,

Reactions which proceed one after the other so that a product of one
reaction is a reactant in a succeeding reaction are known as consecutive
reactions. Thus A — > B —» C indicates consecutive reactions in which the
reactant; A, first forms B which, in turn, reacts to form the final product, C.
If any one of the steps in consecutive reactions is much slower than the others,
the rate of the over-all reaction is governed by the rate of this slowest step.
Sometimes a molecule can react in more than one way and thereby form
different products. Such alternative or parallel reactions can be represented

* A <c

Here the reaction of A yields the products B and C by two

separate mechanisms. In parallel reactions the most rapid step by which
A is used up governs the over-all reaction rate. The product of a reaction in
turn may react to re-form the initial reactant. In such opposing reactions ,
as A B and B A, we shall see in the next chapter that the reactant,
A, is not completely used up but rather a state of equilibrium develops
in which both A and B are present.

(B) Order of a Reaction

A separate means for the classification of reaction rates is based upon the
mathematical expression for the rate of a reaction. This can be deduced
from the Law of Mass Action. For a general reaction,

( 6 )

A + B + . . . products

The Kinetics of Chemical Reaction

101

the rate is propoitional to the concentration of A times the concentration
of B, each concentration being raised to some power. For gases, pressures
can be used in place of concentrations.

The rate of reaction depends upon the probability of molecular collision.
The probability for the successive occurrence of sepaiate events is the product
of their individual probabilities. For this reason it is the product of the con-
centrations of the i cactants and not some other function such as their sum, which
appears in the rate equation.

If the rate of the reaction is denoted by the symbol, v, ( velocity ) and the
concentration of a reactant by the symbol [A] or [B], where the bracket [ ]
represents the concentiation of whatever chemical species is written withm
it, we may write

v oc [A]" 1 [B] w ( oc — 4 is proportional to*)

and thus (7) v — k [A]"' [B] w

The constant, k, is called the reaction rate constant and the exponents,
m and n y are numbers which are not necessarily integral. The order of a
reaction is defined as the sum of these exponents, m and n, the reaction
is said to be mt h order in A and nth order in B. The value of the reaction
rate constant is different for each reaction, for a given reaction, k is indepen-
dent of concentration but depends upon the temperature. The values of the
exponents of the rate equation are determined experimentally. We must
remember that the experimenter has only a macroscopic viewpoint and
not a microscopic insight to the actual reaction mechanism. The reaction
rate can be observed only through the macroscopic change in the concen-
tration of some reactant or product. The exponents, m and n are so chosen
to fit a rate equation to the actual experimental data. Particularly because of
the complications of consecutive, parallel, and opposing reactions, the values
of the exponents generally will not be whole numbers. Thus, two experiments
may be carried out at the same temperature, in one of which the concen-
tration of one reactant, such as A, is double that in the other experiment.
In the experiment with the higher concentration of A, the rate will be greater.
If, for example, the reaction rate for the higher concentration of A is eight
times that for the lower concentration of A, then m must be three, since
8 ~ 2 m . Similarly by changing the concentration of B in a definite ratio,
not necessarily twofold while holding that of A constant in two separate
experiments, and measuring the relative rates of reaction, the value of n
can be determined.

Many unimolecular reactions are also first order, many bimolecular re-
actions second order, and trimolecular reactions third order. But it is not
a necessary condition that the molecularity of a reaction should equal its
order. Indeed such a coincidence is fortuitous. The monomolecular decomposi-
tion of NX3r, is also first order since its experimentally determined rate of
decomposition is given by

(8) v = k [N 2 O fl ]

Doubling the concentration of N 2 O r , would double its rate of reaction. The
bimolecular reaction. H 2 + I 2 -» 2 HI, is second order. Its rate is

(9) v — k [H 2 ][I,]

102

( f'> miral Reaction

IL> K.7i <r:.

The reaction is first order with roypeit u> 11 and also iirst mder m I 2 ,
but second order over-all Doubling the com ontiatum of eithei IL or p
would double the rate oi the reaction, doubling the i ouceniiahun of both
reactants simultaneously would increase the inaction rate fourfold. The
reaction 2 NO -j- Q, — * 2 NO ; is third order, suit** the reaction rate is ex-
pressed by v : k [NO] [NO] [(>. 1 oi

(10) v -- k [NO] J [0,1

A bimolecular action may be tiist older d one of the leaetauts is present
in so large an excess that its concentration may he considered practically
unchanged as the reaction proceeds For example, ethvl acetate reacts with
water to form acetic acid and ethyl alcohol.

CH..COOCJHL r HO H — CM COOH - C H OH

If a small amount of ethyl acetate is dissolved in a large excess oi water the
change in the concentration of the water is practical!) ml as the reaction
proceeds. The reaction rate will depend on the concentration oi the ethyl
acetate only, and will he first order.

An example of a reaction which has a fractional reaction order is that
between ethylene and iodine, C,H* • I ■ — * C Hd,* hi its early stages the
reaction rate is 5/2 order since it follows the rate equation

( 11 ) V =r k [C.H.l [Ll 3/a

4. First Order Reactions. As the course of a reaction proceeds, the con-
centrations of the reacting substances tlecrea.se exponentially with time, as
shown in Figure 8.3. Mathematically, a first order reaction is one in which
a fixed fraction of a substance reacts in a given interval of time. Let us
assume a first order reaction in which 0.20 of a reactant is used up in the
first hour. Then in each successive hourly interval, 20 percent of what was
present at the beginning of that hour would react. ‘Hits is illustrated by
the data in Table 8-A where, for simplicity only, the initial concentration of
the reactant is taken as one mole per liter.

Table 8-A

Concentration

The Kinetics of Chemical Reaction

103

Figure 8.3.

Change in Concentration During
Chemical Reaction.

Starting from the initial concentra-
tion, c 0 , at zero time, the concentra-
tion of a reactant still present de-
creases exponentially with time. The
time for the concentration to de-

crease to half its initial value, ,
is the half life period, t , as shown
by the dashed lines.

The equation giving the exponential decrease in concentration with time
for a first order reaction is

(12) c = c 0 e~ kt (Page 106)

where c is the concentration still present at a time, t, and c 0 is the initial
concentration; k is the reaction rate constant and e is the base of natural
logarithms.

From Equation 12 it can be shown that
(id) logio c = - t + A

(14) and log 1(

2.303

where A is a constant
k

logi,

2.303

Equation 13 indicates that it is the logarithm of the reactant concentration
which is a linear function of time. If the logarithms of these concentration
values are plotted against time, a straight line graph will be obtained. The
value of the reaction rate constant can then be determined from the slope

1 ^

of this line, which is equal to —

Z.oUo

From experimental data as in Table 8-A the value of k can be calculated
by substitution in Equation 14, and once k is known it is possible to calculate
the concentration of a reactant still present at any time, t.

Example 1: Using the data for the first order reaction in Table 8-A for the
concentration at two hours, calculate (a) the reaction rate constant, k; (b) the
concentration of the reactant present at seven hours.

Solution : (a) Substituting in Equation 14, at t = 2.00 hr, c = 0.640 mole/liter,

. 1.00 mole/liter k ,

l0g 0.64 mole/liter ~ 2.303 2 '°° **

, 2.303 , 1.00 mole/liter

k = „ x log — , • = 0.223 hr 1

(b)

log

2.00 hr
1.00 mole/liter

0.64 mole/liter
0.223 hr* 1

2.303

c = 0.210 mole/liter

x 7.00 hr

104

Fhr kit.

r * 1 ' wu'al Reaction

The chemist finds it convenient to calculate the half life period, t^, the
length of time at which just half the initial concentt itmn will icmain, and
so also when half has reacted, namely, when c l\„ as slum it m Figure 8.3,

Substitution of c = V& c u in Equation U givrs foi the half life period

0 693

(15) t % - —

0.693

In Example 1, the half life period would he * {i ^ OU! Some re-

flection will make it apparent that for a first oulei icuchon it takes twice
as long for three-quarters of a substance to react as it does for one-half
to react.

We have seen that the rate of a reaction mei eases with a rise in tempera-
ture and is enhanced by a low value of the actuation eueigv for the re-
action. An empirical equation relating the reaction rate constant, k, to tempera-
ture and activation energy was first proposed by Svante \rrhenius. It is
similar in form to the equation for a first ordei leactnm late.

Al* H

( 16 ) k = se~~

where s is a constant known as the “frequency factor,' " T is the absolute
temperature, and AE a is the energy of activation.

Many important scientific phenomena follow the mathematics of a first
order reaction. Among these are the rate of decay of radioactive elements,
the decrease in the intensity of radiation as it passes through an absorbing
medium, biologic growth, and compound interest if the exponent in Equation 12
is positive. For chemical reactions other than first ordei, analogous rate
equations can be derived. With rare exceptions all are exponential.

QUESTIONS

1. List the factors which affect the rate of a chemical reaction. Explain each
from a kinetic point of view.

2. State and explain the Law of Mass Action, What is a reaction rate constant?

3. What is meant by the energy of activation? Draw a potential energy diagram
illustrating this energy. For what reason should there exist an energy of
activation?

4. Draw a potential energy diagram illustrating energy of activation for an
endothermic reaction.

5. Distinguish between the energy of activation and the heat of reaction.

6. How does a catalyst operate to speed up a reaction? Draw the path of a
catalyzed reaction in Figure 8.1b.

7. Explain how a rise in temperature increases the rate of a chemical reaction,

8. Explain why the reaction rate as shown in Figure 8.3 should continuously
decrease.

9. What prerequisites must be met before two molecules, A and B, can react?

10. Define free radical, activated complex, autocatalvsis, mechanism of a reaction,
half life period.

The Kinetics of Chemical Reaction

105

11. If the energy of activation toi a chemical reaction were zero, what percent
of collisions would be fruitful? Explain.

12. Distinguish between “moleculanty” and “order” of a reaction.

13. Write the rate law expressions and state the ordei for the following single

step reactions: (a) A + 2 B — * AB 2 (b) A + l k B., AB

14. For a gaseous leaction between A and B, the rate determining step is
2 A + B C. (a) Write the rate law for the reaction (b) How would the re-
action rate be affected ( 1 ) if the concentration of B alone is doubled (2) if both
the concentrations of A and B are doubled (3) if the concentration of B
is doubled and, of A is halved (c) How could the concentrations of A and B
be varied so that the reaction rate would remain unchanged?

15. Given the following kinetic data for a gaseous reaction which occurs in a
single step. 2 A 4- B — * A_,B AH = -20 keal.

Experiment

No.

Initial Concentration
of A, mole/ liter

Initial Concentration
of B , mole /liter

Concentration
of A 2 B , mole/ liter,
formed

after one hour

1.0

2.0

1.0

3.0

■ ■ ■

3.0

1.0

4.0

1.0

IHuH

(a) Write the rate expression for the reaction (b) Calculate the value of the
reaction rate constant, k. What are its units? (c) What would be the con-
centration of AoB after one hour if the initial concentrations of A and B
were each 2.0 mole/ liter? (d) How would the value of k be affected by
an increase in temperature?

16. For the data in Table 8- A, calculate the concentration of the reactant remaining
at nine hours and at twelve hours.

17. Draw a graph of the logarithm of concentration against time for the data
in Table 8-A, and calculate the reaction rate constant from the slope of
the graph.

18. The decay of the radiation activity of a radioactive element is first order,
namely, A = A u e~ kt , where A is the activity at any time, t, starting from an
initial activity, A 0 , and k is the decay (reaction rate) constant. If the radiation
activity decreases 25 percent m one year calculate (a) the value of k (b) the
half life of the radioactive element.

APPENDIX

First Order Reaction Rate Equation

If dc represents the change in the concentration of a reactant during the

dc .

time interval, dt, the rate of a chemical reaction is -7- For a first order reaction
* dt

the rate is proportional to the concentration, c. Thus
dc

(1 ) "t- = - kc, where k is the reaction rate constant.

The Kiru'tH'* of Chemical Reaction

The negative sign is necessan because tic is a decrease and is mheiently a
negative quantity. Tiansposing Equation 1,

By the integral calculus,

(3) In c = -k t + A', where A' is a constant oi infegiation
Converting from natural logarithms to logarithms with the base 10,

2.303 !og lt> c = -kt -f A
k

(4) log c = - Tn ' nQ * w ^ ere A' — A '2.303 { Equation 13, page 103)

Substituting the conditions in Equation 2 that, at zero tune when the reaction
begins, the concentration is the initial concentration, c\„ and at anv subsequent
time, t, the concentration is c.

In c - In c„ = -k(t - 0)

(5) or

In — = -k t

(Equation 12. page 103)

Equation 4, in terms of logarithms to the base 10, is

) log — = - 2.3(13 (

Multiplying Equation 6 hv -l gives

(Equation 14, page 103)

The half life period, t is obtained by substituting e ~~ M» e„ in Equation 7.

bg i rt= ioos 11 *

_ 2.303 log 2 _ 0.(593

(9)

0.693

(Equation 15, page 104)

Chemical Equilibrium

1. Reversible Reactions. A reversible reaction is one in which the
products formed can, in turn, react to re-form the original substances. We
have already intimated that such reactions exist by writing A —> B and
B — » A for opposing reactions in our study of kinetics. Rather than write
two separate equations, the chemist indicates the reversibility of a chemical
reaction by writing a pair of oppositely directed arrows, thus A B. Not
all reactions are reversible in practice. Though on paper it is no great
effort to write a double arrow, nature does not always obey our paper
mechanics. The reaction, 2 KC10 3 -» 2 KC1 -f 3 0 2) is irreversible in that
KC1 and 0 2 will not combine directly to form KC10 3 . An irreversible re-
action proceeds in one direction only.

The reaction between red hot iron and steam is reversible. If steam is
passed over heated iron, magnetic oxide of iron, Fe 3 0 4 (not to be confused
with rust, Fe 2 0 3 ) and hydrogen are formed.

(1) 3 Fe(s) + 4 HoO(g) — Fe 3 0 4 (s) + 4 H,(g)

If this reaction is carried out in an open container, a product of the re-
action, hydrogen gas, will escape and the reverse reaction between it and
Fe :i Oi will not be possible. The reaction is then said to goto completion, that
is, the reactants will be consumed completely in the formation of products
and, in time, there will be present solely the products of the reaction. By their
very nature, all irreversible reactions go to completion. However, if the
iron-steam reaction is carried out in a sealed container so that the products
of the reaction cannot escape, then the reverse reaction can take place.
What will be apparent experimentally to the macroscopic observer (you
and me) is that, in time, the chemical reaction seems to stop before all the
iron is converted to Fe 3 0 4 . The reacting system seems to have reached a
state of equilibrium and thereafter no further change is apparent. On the
molecular level, however, there are two reactions proceeding at the same
time as indicated by the double arrow, each undoing the effect of the other
so that the reaction cannot go to completion. Both reactants and products

107

then will always be present jn equilibrium with t\uh othui. Thw air in a
state of balanced dynamic actixitv ami not in a statu* state uf iest or in-
activity. A chemical reaction thus will always be incomplete when the re-
verse reaction also takes place under the saint' cnmhtions

2. Chemical Equilibrium. Let us try to pictuu* the kmetu s uf a rever-
sible reaction. The Deacon process for making chlonue fioin Indrogen
chloride is reversible.

(2) 4 HClU) + O^S) ^ 2 CLu* t 2HO(g»

When HC1 and Ch, in the ratio of 4 to 1 by \uhum\ and hence bv moles,
are heated to about 450“ C, they react slowly to produce Cl. and ll .O. In
the presence of a catalyst, CuCL, the reaction is speeded up but does not
go to completion. About 20 percent of the HC1 and () rein iins unchanged.
Consider the forward reaction between HC1 and ().. It onlv HC1 and 0>,
are mixed initially, it is evident that at first no reaction other than that be-
tween these two is possible. This reaction will start at a maximum rate
because the concentrations of the reacting substances are at then highest
values. How r ever, as the reaction proceeds the concent! ations of these sub-
stances decrease and hence the rate of the reaction between them decreases.
At the start of the reaction there are no Cl- and ILO molecules but they
will increase in number, and in concentration, as the action proceeds. The
rate of the reverse reaction will increase similarly. We haxe a system m
which one reaction, starting at a maximum, decreases m rate progressively,
and the reverse reaction, starting at zero rate, continually im reaxes in late.
Finally the two rates must become equal at some point m tune. When this
condition is reached all four substances are being formed just as fast as
they are being consumed, and no further changes in concentration take place
No change occurs in the rates of the reactions and the two opposing re-
actions go on at equal rates. There is no cessation of either reaction; the
molecules are constantly colliding with each other and undergoing change.
There is present a dynamically balanced system, not a statically balanced
one. All reversible reactions must lead to this condition of apparently sus-
pended action, a condition of chemical equilibrium .

The relationships between the rates of the two opposing reactions in a
chemical equilibrium are shown diagramaticaily in Figure 9.1 At the start
the forward reaction proceeds at a high rate, A, but its rate falls off with
time. The rate of the reverse reaction is zero at the start but increases
with time. Eventually the two rates become equal at point B. From B
to C the two reactions proceed at the same rate; the system is at equilibrium.
At equilibrium there are two opposing reactions going on at the same time
with the same rates.

The fact that an equilibrium exists does not mean that a reaction has
progressed just 50 percent to completion and that then there is no further
change. Indeed this is seldom the case. The extent to which a reaction has
proceeded in the direction of the products, or “to the right/* when equi-
librium is reached is known as the point of equilibrium . Corresponding to
any degree of completion, the point of equilibrium can be any value be-
tween 0 and 100 percent, A low value means that the reaction has gone

Chemical Equilibrium

109

to the right only to a small extent before equilibrium was reached, a high
value indicates that, at equilibrium, the reaction has gone to the right to a
large extent. For an irreversible reaction the point of equilibrium is 100
percent to the right. For the Deacon process at 450° C, the point of equi-
librium is about 80 percent towards completion, that is, about 80 percent
of the initial concentrations of HC1 and 0 2 have reacted to form products.
Sometimes the chemist indicates the fact that an equilibrium is displaced
preferentially in one direction by an unequal length of the double arrows.
To show that the products are favored in the equilibrium mixture the
Deacon process would be written as:

(3) 4 HC1 + 0 2 * 2 HoO + Cl 2

Figure 9.1. Diagram Showing the Rates of Reactions in the System
4 HC1 + 0 2 ^ 2 HP + 2 Cl 2

The forward reaction starts at a maximum rate A, which decreases as the reaction
proceeds. The reverse reaction starts at zero rate, but increases as the reaction
proceeds. Eventually the two rates become equal at B and the system is at
equilibrium.

3. Factors Affecting the Point of Equilibrium. The relative concentra-
tions of the interacting substances at equilibrium, or the point of equilibrium,
is determined by the relative reaction rates of the forward and reverse re-
actions. The attainment of equilibrium can be viewed as an adjustment of
these concentrations until the rates of the opposing reactions are equal. By
altering the conditions under which a reaction is taking place, the forward
and reverse reaction rates may be changed unequally and a new point of
equilibrium will be established. Such a change in a chemical system, re-

suiting in a new point of equilibrium, is known as a displacement of the
equilibrium. Many industrial reactions come to equilibrium and arc 1 thereby
incomplete. By altering the conditions on tin* s\stem *n equilibrium, the
chemist and the engineer attempt to displace tin* point of equilibrium as
much as possible in the direction ot the desired pioduct.

Of the factors that affect the rate* of reaction, onlv changes m concen-
tration (or pressure for gases ? and tempeiatme ultei the point ot equi-
librium. The state of subdivision may ha\e a small imlnect effect and is
usually neglected. A catalyst has no effect upon the point of equilibrium
since it affects both the forward and the recerse n ac tions identically Since
a catalyst reduces the activation energy for both the foiwaid and the reverse
reactions by a like amount, both reactions are speeded up equally. Equi-
librium is reached in less time but the point of equihbnum is the same as
when no catalyst is used.

4. The Concentration Factor, To illustsatc tin* effect ot change's m con-
centration upon the point of equilibrium, let us consider again the equilibrium

(4) 4 HCl(g) + Oqg) ^ 2 Cl itC * 2 H

When equilibrium is attained the two opposing reactions proceed at the
same rate and we have a mixture m which all four substances are present,
each in a concentration invariant with time. If any of the concent lations
are thereafter changed, the point of equilibrium will be displace!

Let us assume that this equilibrium mixture is m a closed vessel of
constant volume so that any change in the quantity of one of the reacting
substances constitutes a change in the concentration "of that react, mt, Further
to be certain that any observable effect will be due solely to a concentra-
tion change, let us maintain the equilibrium mixture at a constant temper-
ature, an experimental state of affairs easily accomplished by placing the
vessel in a thermostat or surrounding it with heating or cooling coils as
necessary. Now into the equilibrium mixture let us inject additional HC1
and thereby increase its concentration. The increase in the concentration
of the HC1 will result in a greater number of collisions between the HC1
and the 0 2 molecules. The rate of the forward reaction will therefore in-
crease. The concentrations of the Cl. and the H.O (products) are not
immediately affected by the addition of the HC1 but presently the in-
creased rate of reaction between the HC1 and O. results in an increased
concentration or yield of the Cl 2 and HoO, which now in turn undergo the
reverse reaction at a faster rate. As the newly injected reactant, HCi, is
used up the rate of the forward reaction diminishes from its new high level,
while die reverse reaction increases in rate as the concentration of products
increases. Under the new conditions there will again be reached in time
a new point at which the forward and reverse reaction rates are equal and
a state of equilibrium will again result. The length of time it takes for a
new equilibrium to be established depends upon each individual reaction
and will vary from the almost instantaneous to the appreciable. The net
result of increasing the concentration of the HCI is an increased concen-
tration of the products Cl 2 and H 2 0, that is, the point of equilibrium has
been displaced to the right. Also at the new state of equilibrium the HCI

Chemical Equilibrium

111

concentration is greater than it was initially merely because it was the
substance added while the concentration of 0> is diminished because it is
used up in forming the higher concentrations of Cl> and H 2 0.

Increasing the concentration of the 0 2 alone would produce a result
similar to that described for the HC1 addition. On the other hand, if the
concentration of either the Cl> or H 2 0, or both, were increased, the point of
equilibrium would be displaced to the left, resulting in higher concentra-
tions of HC1 and 0 2 . In general an increase in the concentration of any of
the reacting substances displaces the point of equilibrium in the direction
of its products. The student might now consider how the point of equi-
librium would be affected if any one of the substances in an equilibrium
mixture were removed, e.g., if the Cl 2 concentration were reduced. It is
evident that a reversible reaction can be made to go to completion by the
removal of one of the products as fast as it is formed.

5. The Equilibrium Equation. We may be forgiven if we remind the

reader that the qualitative aspect of a theory, to be truly of value, must lend
itself to quantitative interpretation. With the mathematical background
of reaction rate theory already developed from the Law of Mass Action, this
can be done quite simply for the concept of chemical equilibrium.

For the general homogeneous gaseous reaction

(5) A(g) + B(g) C(g) + D(g)

let us assume that both the forward and the reverse reactions are uncom-

plicated single-step second order processes. Then the rates of the forward and
reverse reactions are given by

Vf = k f [A] [B] and v r - k r [C] [D]

where the subscripts f and r characterize the forward and the reverse
actions respectively. At equilibrium the rates of the opposing reactions
equal, that is, v f = v r . Therefore,

kf [A] [B] = k r [C] [D]

By separating the reaction rate constants from the concentrations
equation can be rewritten in the form

kf [C] [D]
k r “ [A] [B]

The ratio of two constants is itself a third constant, so that
K is called the equilibrium constant and we may write for it

( 6 )

_ fC] [D]
[A] [B]

Even if the opposing reactions were not single step processes the same
expression for the equilibrium constant, Equation 6, would have been ob-
tained because any kinetic deviations would appear identically in both the
numerator and the denominator and thus cancel each other out. In other

words the state of equilibuum is lmlepoiuDni ot am mrrhamxm by which
it is attained.

More generally for the 1 faction

( 7 )

a A(g) -i- h B{ £4 ) r C { g) -p «/ D<

the expression for the equilibrium constant is

[C]‘ |D]"

(8)

[A]* [Bp

In general, to write the equilibinun constant expression h»i a chemical
reaction, we write a fraction wherein the numciatoi contains the product
of the concentrations of the chemical products and the denominator the
product of the concentrations of the reactants, each cnmrntiation being
raised to a power which is equal to the coefficient pieceding the sub-
stance in the balanced equation for the reaction. Thus for the Deacon process

(»)

[HjOp icy-

[hci] 4 toy

Again where gases are concerned, instead of molui
sures may be used in writing an equilibrium constant
Thus

concent! ations, pres-
expression.

( 10 )

K„

PVl,

The subscript "p” to the equilibrium constant indicates that K is m terms
of pressure. The value of K,, is not equal numerically or dimensionally to
the value of the equilibrium constant written in terms of concentration, now
designated for distinction as K e . The two equilibrium constants are related by

(11) K p = K«.(RT)** n where R -= the molar gas constant,

0.0821 liter atm/mole cleg
T ~ the absolute temperature*

An *-r the number of moles of gaseous prod-
ucts minus the number of moles of gas-
eous reactants; for the Deacon process,
An = -I (page 88)

The significance of the equilibrium constant is this. If we were to mix
any arbitrary concentrations of chemical species such as A, B, C, and 1) of
Equation 7 these concentrations would adjust themselves hv the chemical
reactions involved so that the values of the concentrations at equilibrium,
when inserted into the equilibrium constant expression, would always equal
the same value— the value of the equilibrium constant K. This holds true
whether we start solely with the reactants A and B, or with the products
C and D, or with any mixture of these.

For a giver* equilibrium system the value of the equilibrium constant is
characteristic; it is independent of the relative concentrations and depends
solely upon the temperature. That the value of the equilibrium constant

Chemical Equilibrium

113

remains constant at a given temperature is evident from the experimental
data given in Table 9-A for the leversible reaction,

C0 2 (g) + H,(g) H 2 0(g) + CO(g),

at 986 °C. In this case the expression for the equilibrium constant is:

[H 2 0] [CO]

(12)

K =

[CO,.] [H.,]

Table 9-A

Initial Mixture ,
moles/liter

Final Mixture ,
moles /liter

(at 986°C)

CO,

h 2

CO = h 2 o°

10.1

89.9

0.70

80.38

9.46

1.59

30.1

69.9

7.18

46.82

23.00

1.58

49.1

51.9

2152

22.82

27.83

1.57

60.9

39.1

34.67

12.77

26.28

1.56

70.3

29.7

47.66

6.76

22.79

1.61

*Note. The concentrations of CO and H 2 0 are equal inasmuch as their initial con-
centrations were zero and they are produced in a 1 :1 molar ratio.

The concentrations of the reacting substances in Table 9-A appear to be
totally unrelated, especially since the initial concentrations of C0 2 and H 2
were chosen arbitrarily and are not necessarily in the 1 to 1. molar ratio
in which they react. Yet when the final equilibrium concentrations are sub-
stituted in the equilibrium constant expression the values obtained for the
equilibrium constant are the same, after making due allowance for ex-
perimental error.

An equilibrium constant, K, applies to an equation in the “specific form
in which it is written. Thus for an equation written in the reverse sense
(backwards) the value of the equilibrium constant is the reciprocal of that
for the forward equation. If K 0 is the equilibrium constant for the reaction,
A + B C -f D, for the reaction, C + D ^ A I B, the equilibrium
constant has the value 1/K t . For the reaction, 2 A -f 2 B 2 C + 2 D,
the constant would be Ki 2 .

In writing an equilibrium constant expression, inasmuch as the con-
vention is to write the concentrations of chemical products in the numerator
and the concentrations of reactants in the denominator, a large numerical
value of the constant indicates that the chemical products are present in
concentrations large in comparison to the concentrations of the remaining
reactants, and hence the chemical reaction has proceeded to the right to
a large extent at equilibrium. Similarly, a small value of the equilibrium
constant indicates a chemical reaction which proceeds to the right but
slightly. The magnitude of the equilibrium constant is thus a measure of
"chemical affinity,” or the tendency of a reaction to proceed. From a ther-
modynamic point of view, the point of equilibrium is related to the differ-

114

v <iuthbriwn

( 'h mi, al

ence in free energy A F, of the u*actunt> and products m a chemical re-
action. The greater the decrease m free rnfi^ m going tiom reactants to
products, the more will the point of equilibrium he to the light m the
direction of the formation of the products, and so the larger will he the
value of the equilibrium constant, inasmuch as both the iur rnergv change
and the equilibrium constant are measures oi the extent to which a re-
action proceeds, the)' are related.

(13) - AF RT log, h - RT log. Q

where A F = the free energy change of a chemical reaction
K = the equilibrium constant of the reaction

Q — an expression similar m form to that for K but containing
the actual initial concentrations [m pressures) of the reactants
and the final concentrations of the products. It differs from K
in that the latter is calculated by inserting into its expression the
\aiues of concentration which exist at equilibrium
R = the molar gas constant
T = the absolute temperature

In the event that all reactants and products are in their standard states,
e.g., for a gas, a pressure of one atmosphere, then Q corrals one. and the
final term vanishes, leaving

(14) - AF° ■= RT log, K

where the superscript indicates that all substances are in their standard
states, and AF“ is known as the standard free energy change. It should be
noted that the more negative the value of AF, that is, the greater the de-
crease in free energy, the larger is the value of K.

The equilibrium constant is of fundamental importance. Later on we
shall see that other useful tools of the chemical trade, such as the ioniza-
tion constant and the solubility product constant, are nothing other than
applications of the basic equilibrium constant despite their difference in
names.

Example 1: The dissociation of phosphorus pentachloride, POL,, vapor at high
temperature into phosphorus trichloride, PCb, and chlorine, CL, results in the fol-
lowing equilibrium.

PCl 5 {g) PCl,(g) + CL(g)

If 1.00 mole of PC1 5 alone is placed initially in a volume of 2.00 liters at 25<PC
it is found that at equilibrium the concentrations of the PC1-, PCi ; , and CL are*
respectively, 0.375 mole/liter, 0.125 mole/liter, and 0.125 mole/liter. Calculate the
equilibrium constant for the reaction at 250°C.

Solution: The equilibrium constant expression is:

_ [PCl 3 j [Cl 2 j

[pcy

Chemical Equilibrium

115

Substituting the values of the concentiations at equilibrium , we obtain

__ 0.125 mole/liter x 0.125 mole/liter

0.375 mole/ liter

and K = 0.0416 mole/liter

Note that the units of K will depend upon the form of each equilibrium constant
expression.

Example 2: Calculate the concentration of Cl 2 at 250° C in an equilibrium mix-
ture in which the concentration of PCI, is 0.500 mole/liter and the concentration
of PCI., is 0.200 mole/liter.

Solution: Now that the value of the equilibrium constant is known, we can
solve for the unknown concentration of CL by substituting the concentrations of the
other components of the equilibrium mixture into the expression for K.

Thus

0.0416 mole/liter =

0.200 mole/liter X [Cl a ]
0.500 mole/liter

Solving for [CL], we obtain

[CL] =• 0.104 mole/liter

Example 3: When 1.00 mole of PC1 3 is placed initially in a volume of 10.0 liters
at 250° C, what will be the concentration of CL> at equilibrium?

Solution : This problem differs from the conditions in Example 1 in that the 1.00
mole of PC1 5 is placed not in 2.00 liters but in a different volilme, 10.0 liters. The
value of K will be unchanged but the fraction of the 1.00 mole which reacts, or
the degree to which the reaction goes to the right, will be different.

Since we do not know how much of the PC1 5 will react, let x represent the
number of moles of PCL, that react to establish equilibrium. Therefore the number
of moles of PCL, still remaining at equilibrium, which is equal to the initial number
of moles less the number which react, is (1.00 - x).

The balanced chemical equation, PC1 5 — > PC1 3 + Cl 2 , tells us the relative
number of moles which react and are produced. Thus one mole of PC1 5 reacts to
produce one mole of PC1 3 and one mole of CL; more generally, the number of
moles of PC1 5 which reacts produces an equal number of moles of PC1 3 and also
of CL. From the reaction of x moles of PC1 5 there will be produced x moles of
PC1 3 and x moles of CL.

The concentration in mole/liter of each substance at equilibrium is

PC1 5

1.00 -

10.0

PCL,

~ 10.0

Cl 2

__ X

“ 10.0

- mole/liter
mole/liter

mole/liter

116

( lu imcat htjtiHibrium

Substituting in the expiession for K

0,0416 mole,/ liter =

mole \ v / x
liter / \ 10.00

1 .00 - x mole
i() 0 liter

mole ^
Titer /

and solving for x, x = 0.469 mole

The concentration of Cl 2 is

0.469 mole
W "liter

or 0.0169 mole/hter

Note that this value is also equal to the concentration of PCl„ while the con-
centration of PC1 & remaining at equilibrium is ~~ ~ ~ > or 0.0531 mole/liter.

1(1 0 utci

The degree of dissociation of the PCI*, usually represented In the symbol a,
is defined as the fraction of one mole of POL, which reacts. In this problem,
a is also numerically equal to r, or

0.469 mole _

a = -r — ■ — r- = 0.469, OK 46.9% dissociation

1.00 mole

Example 4: If 2.50 moles of PC1 5 at 250° G are placed in a 10.0 liter volume
which already contains 0.400 moles of CL, calculate: U) the concentiation of PCI,
at equilibrium (b) the degree of dissociation of POL, (e) the total pressuie of the
equilibrium mixture.

Solution: The student would do well to formulate the solution of equilibrium
problems in the following manner:

PCI, PCI,

Initial number of moles 2.50 0

Let x equal the number of moles of PCI., whic h reacts.
Number of moles at equilibrium 2.50 — x x

Total number of moles (2.50 - x) 4- x -f (0 lOO : x)

0.400

0.400 4 x
2.90 4 x

Concentration at equilibrium

K c - 0.0416 =

and x = 0.732 mole

2.50 - x x

10.0 10.0

(•

v / 0,400 + x\
\ 10.0 /
2.50 - x \

( 10 . 0 ) )

(WOO -*■ I
10 0

(a) The concentration of PC1 3 equals = 0.0732 mole/litcr

Chemical Equilibrium

0.732 mole
2.50 mole

117

(b) The degree of dissociation, equals

= 0.293, or 29.3%

(c) On the assumption that all the gases behave ideally, the Ideal Gas Law can
be applied, using the total number of moles present in the equilibrium mixture,
(2.90 + x) or 3.63 moles, to calculate the total pressure.

nRT 3.63 mole x 0.0821 liter atm/mole deg x 523 deg K

P “ ~ V “ 10.0 liter

= 15.6 atm

The partial pressures of the individual gases could be calculated similarly by
substituting the number of moles of each into the Ideal Gas Law.

Example 5. If the concentration of Cl 2 in Example 4 were doubled by the
addition of Cl 2 alone (1.132 moles) at 250 °C, what would be the value of the
equilibrium constant?

Solution. The value of the equilibrium constant is unchanged by the addition
of CL, or any of the reacting substances, if the temperature remains constant.
From a superficial arithmetic viewpoint it might seem that doubling the concen-
tration of CL would double K but this is not so. The increased concentration of
CL will increase the rate of the reverse reaction and more PC1 5 will be formed
at the expense of PC1 3 and CL, until the two opposing reaction rates are again
equal. During this time the concentrations of PC1 3 and of Cl 2 will decrease and
the concentration of PCL, will increase until, at the new point of equilibrium, they
will be such that when substituted into the expression for K, the quotient will
still equal 0.0416.

Problems concerning more complicated reactions in which the ratio of the re-
acting molecules is not 1 to 1 are treated in the same manner.

Example 6: At 527 °C, 2.00 moles of N a and 0.500 mole of H 2 are introduced
to a 4.00 liter container. The reaction

N 2 (g) + 3 H 2 (g) 2 NHs(g)

takes place and at equilibrium the total pressure is 40.72 atmospheres. Calculate:
(a) the number of moles of NH< present at equilibrium (b) the partial pressure
of NH< at equilibrium (c) K c .

. , . N 2 H a NH,

ton. num ber of moles 2.00 0.500 0

Let x equal the number of moles of N 2 which react. (Note below)

Number of moles at equilibrium 2.00 — x 0.500 — 3x 2x

Total number of moles (2.00 - x ) 4- (0500 — 3x) 4- 2x = 2.50 — 2x

. 2.00 - x 0.500 - 3x 2x

Concentration at equilibrium — 4*00 "TOO*

Note: What x represents is at the discretion of the person solving the prob-
lem. x could have been made equal to the number of moles of H* which reacts

118

< intuit til /’( fuihhrium

or the number or
moles ox the othei
tiples of v This

moles of Mi winch aie piotburd but tim. the inmibeis of
substances would have been fiiUtmnai. and not internal, mul-
ls mourn ement but ue\ eithele^s also tviieit

ta) Assuming that the* Ideal Gas Law is applicable. nKi

40.72 atm - 4.00 htei --= (2.50-20 mole ' O.OS21 htei atm mole chat S00 deg K

and solving tor i, i 0 010 mole.

The number of moles of NH, is 2a, oi 0.020 mole

(b) The paitial pressure of NIL can also be ealenlated hom the Ideal <»as Law.
n R T (2.x) R T 0.020 • 0.0821 * so ° .

6. Heterogeneous Equilibrium. Equilibrium systems mu\ be homogene-
ous or heterogeneous depending upon the number of pluses involved in
the equilibrium. Where the reacting substances present in an equilibrium
system are all in one phase the equilibrium ns said to he homogeneous.
When more than one phase is involved the equilibrium is heterogeneous.
So far we have confined our study only to homogeneous equilibria though
the concepts of equilibria are applicable to both homogeneous and hetero-
geneous equilibria. Contrary to what vve might expect, equilibrium constant
expressions for heterogeneous equilibria are simpler than those for homo-
geneous equilibria. In a heterogeneous equilibrium involving gases and
solids, the concentrations of the solids are assumed to be constant. Only
the surface molecules of a solid can react and their number per square centi-
meter of surface, or the surface concentration, is a fixed value for a given
solid substance. Independent of the hulk of the solid, the reacting concen-
tration of a solid is constant. Hence the concentrations of solids, where
gases are also present, do not appear in the equilibrium constant cxpiession;
the expression for K is written in terms of the gaseous substances only in
the manner we have already indicated.

For the heterogeneous decomposition of calcium carbonate, CaCO,,

(15) CaCO,(.s) CaO(s) + COqg)

we might write an equilibrium constant, K'
whence

[CaOl [CO,

K' -T a - CO ] [CO.

[CuCOJ

[CaO]

Since the concentrations of the solids CaCO and 'CaO are constant, the
term on the left, K' , is also constant. The value of this com-

Chemical Equilibrium

119

posite term is taken to be the equilibrium constant, K, for the reaction. Thus
(16) K = [C0 2 ] and, in terms of pressures, K p — ? C 02

Example 7 . At 800 °C the dissociation pressure of CaCO s is 167 mm. Cal-
culate the value of the equilibrium constant, K p .

Solution: Besides being constant, the vapor pressures of most solids are negli-
gibly small in comparison to gaseous pressures. Hence the entire pressure of 167mm
is due to the C0 2 formed.

Therefore K p — P 0Oi = 167 mm.

If the pressure above the equilibrium mixture is reduced to a value below
167 mm at 800° C more CaC0 3 will decompose to keep up the pressure of 167 mm.
If additional CO._, is injected into the equilibrium mixture, tending to raise the
pressure above 167 mm, some C0 2 will combine with CaO and form CaC0 3 ,
thereby .reducing the C0 2 pressure to 167 mm.

7. The Pressure Factor. Changing the pressure upon a system at equi-
librium will affect the point of equilibrium insofar as the pressure change
affects the concentrations of the reacting substances. In general when the
pressure upon an object is increased, it tends to shrink or decrease in volume.
Chemical systems behave similarly and, under increased pressure, also shrink
if possible. When the pressure upon a chemical system is increased, the
point of equilibrium shifts in the direction which will result in a smaller
volume for the system. For equilibria involving gases this shift will be in
the direction of the smaller number of gaseous molecules. Conversely a
reduction in pressure displaces the point of equilibrium in the direction of
the larger number of gaseous molecules. Note that we are speaking here
of a change in pressure upon the entire system, and not of a change in
pressure of an individual reactant. The latter type of change is a change
in the concentration of a single reactant only and can be treated as such.
Also it is assumed that the pressure change occurs while the temperature
remains constant.

For the equilibrium PC1 5 PC1 3 + Cl 2 , an increase in pressure dis-
places the point of equilibrium to the left, thereby decreasing the degree
of dissociation of PCI*. The higher the pressure, the lower the degree of
dissociation. A decrease in pressure shifts the point of equilibrium to the
right. The calculations of Examples 1 and 3 illustrate this behavior. In the
larger volume of Example 3, 10.0 liters, produced by a reduction in the
external pressure upon the system, the number of moles of Cl 2 is 0.469 as
compared with the 0.125 moles of Cl 2 in the smaller volume of 1.00 liter
in Example 1. The greater degree of dissociation of PC1 5 in Example 3 in-
dicates that a reduction in pressure shifted the point of equilibrium to the
right. Where the number of moles of gaseous products and reactants are
equal, as in the reaction H 2 (g) + Mg) ^ 2 HI(g), a change of pres-
sure has no effect upon the point of equilibrium.

Since the densities of solids and liquids are much greater than that of
gases, their volumes can be considered negligible in comparison with that
of gases. In a heterogeneous equilibrium involving gases, only the number

120

ChcJtut al l tjiu librium

of gaseous moles are counted m determining which side of a chemical equa-
tion presents the smaller volume, hoi the reaction,

2 C{s) -r O.(g) ^ 2 CO (g).

an increase in pressure would displace the equilibrium to the left since there
are two moles of gas on the right and only one on the leit oi the equation.
Just as with concentration changes, changes m pressure affect onlv the point
of equilibrium and not the value of the equilibrium constant

8. The Temperature Factor. For a given icaction the equilihnuin con-
stant has a definite value at a definite temperature This value varies with
temperature because the rates of the forward and leverse reactions are af-
fected unequally by the same temperature change The forward and reverse
reactions have different activation energies 1 and a giv en tcmperatuie use
always speeds up to a greater extent that reaction with the higher activa-
tion energy. Thus the ratio of the reaction rate constants, k 5 /kj, will change
and so also the equilibrium constant, K, will change with temperature. In
the Deacon process an increase in tcmperatuie decreases the yield of chlorine.
The increase in temperature increases the rate of the reverst 1 reaction to a
greater extent than it does the rate of the forward reaction, and the point
of equilibrium is displaced to the left. Conversely a decrease m tempera-
ture shifts the point of equilibrium to the right. Howes ei, in the system
PC1 5 PCh + Cl 2 , an increase in temperature favors the forward miction
and the equilibrium is displaced to the right. The increase n\ temperature in-
creases the speed of the decomposition of PCI-, more than it does the speed
of the recombination of the PCL with CL.

The effect of temperature upon the point of equilibrium was generalized
by the Dutch chemist, Jacobus Henricus Van’t Hoff as follows. When the
temperature of a system in equilibrium is raised, the equilibrium point is
displaced m the direction which absorbs heat. The converse of this state-
ment is also true. For example, the forward reaction of the Deacon process
is exothermic while the reverse reaction is endothermic.

(17) 4 HCl(g) + 0,(g)^± 2 CL (g) + 2 H 2 0(g) AH ^ -28,000 cal

When we increase the temperature of this system, the equilibrium is dis-
placed to the left, in the direction of the action which absorbs, or uses up,
heat. Thus a greater percentage yield of chlorine will be obtained by low-
ering the temperature.

Note that, though a greater percent of the reactants is converted in Oh at
lower temperatures, time is always a factor in chemical reaction. At higher tem-
peratures, the more rapid rat.es ^of reaction, albeit at a lower percentage vield,
may produce a greater quantity of Cl 2 in a working clay. Industrially, an optimum
temperature may have to be chosen as a compromise between a low percent y ield
and a higher reaction rate which will give the maximum quantity of the desired
product in a given time.

The decomposition of PC1 5 is an endothermic reaction:

(18) PCl.(g) ^ PCI 3 (g) + Cl 2 (g) AH ~ -1-30,000 cal

1 Theoretically it is possible that the forward and reverse reactions have identical
activation energies, but such a situation would he extremely rare.

Chemical Equilibrium

121

The forward reaction absorbs heat, and hence an increase in temperature
favors this action and the equilibrium is displaced to the right. A decrease
in temperature would shift the equilibrium to the left; a lesser percent
of PCl«j would decompose.

It may be convenient to visualize heat energy as one of the reactants in
a chemical equation. For Equation (18) the endothermic heat of ‘reaction
can be written as an integral part of the equation, thus:

PC Mg) + 30,000 cal PCl :i (g) + Cl 2 (g)

We have seen that increasing the concentration of a chemical species in an
equilibrium system displaces the point of equilibrium in the opposite direction.
Similarly, increasing the concentration of heat energy, or raising the tempera-
ture, will drive the reaction the direction opposite to the side of the equation
on which the heat is written as a positive quantity. For the PCh equilibrium
system, an increased concentration of PC1 5 shifts the point of equilibrium to
the right; so too, will an increase in the concentration of heat energy, or a rise
in temperature. For the exothermic reaction between C 2 H 4 and H 2 at 900°C,
written as C 2 H» -|- H 2 CoH* -f 24, (XX) cal, an increase in temperature of
the system at equilibrium would displace the point of equilibrium to the
left. Thus the conclusions derived for concentration changes could apply
equally well to heat energy or temperature changes.

Van’t Hoff’s generalization also applies to physical equilibria in saturated
solutions. If a solid, dissolving in a nearly saturated solution, absorbs heat,
then heating a mixture of the solid and a saturated solution of the solid
should cause more of the solid to dissolve as the temperature rises.

Solid -f heat Solution

The rise in temperature would favor the endothermic action— the solution
of the solid. Most solids are more soluble at higher temperatures than at
lower temperatures. (See solubility curves, page 230) When warm satu-
rated solutions are cooled, some of the solid precipitates. Such precipitation
is accompanied by an evolution of heat. There are, however, some sub-
stances which precipitate from saturated solutions with absorption of heat.
Such solids should be less soluble at elevated temperatures, because as the
temperature is raised, that action takes place which absorbs heat— the pre-
cipitation of the solid. This is the case with anhydrous sodium sulfate.

9. Le Chatelier’s Principle. The generalization of Van’t Hoff and the
effects of concentration and pressure changes upon systems at equilibrium
are particular cases of the principle first stated *in 1884 by the French
chemist, Henri Louis Le Chatelier. Le Chatelier s principle states: If some
stress (change in temperature, concentration, or pressure) is brought to
bear upon a system in equilibrium* the point of equilibrium will shift in the
direction that tends to undo or minimize the stress. In other words, systems
in equilibrium readjust in such a manner as to restpre the original condi-
tions and to minimize the effect of the external change imposed on the system.

Let us apply Le Chatelier s principle to an equilibrium system for which
the forward reaction is exothermic, as given below.

(19) A (g) + 3 B(g) 2 C(g) AH = - X caP

2 A practical reaction of this type is N 2 -j- 3 H a ;z± 2 NH 3 aH = —22 keal.

122

* ' . m>

i I tjuihiirium

A) Concentration. Suppose foi rvnmph that tlu* mhh * -uti utiou ot A is
increased. The point of equilibrium will shift m that dut«hou wlmh will
tend to undo this increased concentration at V nann K to th** light This
results m the formation 'of more (- and a icdiution m tlv uu leased con-
centration of A Note that the ongmal conditions .no not upinilmed and
that the shift m the equilibrium onh tend s to testmt fho nmyiu! condi-
tions. In general an increase in the eoncentintum ot un\ suKt.uuo \n the
equilibrium system displaces the point of equilibrium towanis tin opposite
side of the chemical equation. A i eduction m tin* * oik enh ation of any
species would displace the equilibrium m the direction of that sprues Thus
a reaction can be made to go to completion b\ the complete removal of
one of the products.

B) Pressure: If the pressure upon a gaseous s\stem m equthhimm is in-
creased, the point of equilibrium shifts in the direction which tends to
minimize the applied pressure, namely, in the direction wlmh will pro-
duce a decrease in volume*. In Equation 19, four \olumes of aaseous react-
ants produce two volumes of gaseous products, m tin torw.ud reaction. If
the pressure upon the system is increased the resulting shift in the equi-
librium will take place so as to reduce the increased pressure, that is, in
the direction of the smaller volume. \ shift m the other direction would
increase the number of moles and further increase the pressure Thus
the displacement will he to the right and the proportion of products. C in

' this case, will increase. Similarly, if the pressure is reduced tin' equilibrium
will be displaced to the left since the tendency will be to maintain the
original higher pressure,

C) Temperature . An increase in temperature favors the reaction which tends
to reduce the temperature, that is, the reaction which proceeds with the
absorption of heat. Increasing the temperature of tin* equilibnum mixture
in Equation 19 would favor the reverse icaction. The point of equilibrium
would be displaced to the lei t with the formation of more \ and B. A
reduction in temperature would favor the heat-producing reaction with a
resultant increased comersion of A and 13 into C These conclusions are
in accord with Van’t Hoffs generalization.

Le Chatelier’s principle can be applied also to physical systems at equi-
librium. If a mixture of solid ice and liquid water in equilibrium at 0 c C
is heated, the ice melts because in so doing it tends to ketp the temperature
down by absorbing heat, namely, its heat of fusion. Also, if the ice-water
mixture is compressed at constant temperature, the ice melts because liquid
water has a greater density than does ice and so would occupy a smaller
volume. In melting the ice tends to reduce the applied increased pressure.

The principle of Le Chatelier is one of the most comprehensive of sci-
entific generalizations with implications not merely in the realm ot chemistry
but in ail fields of scientific thought. It is a general statement of the physi-
cal Law of Inertia— that systems at equilibrium resist change and tend to
maintain the status quo. Thereby it implies many physical laws, among
them Newtons Laws of Motion, Lenzs Law of Electromagnetic Induction,
and even the economic Law of Supply and Demand, wherein the price
of a commodity may be construed as the point of equilibrium.

Chemical Kiiuihbiium

123

QUESTIONS

1. Define unci illustrate what is meant by (a) reversible reaction (b) chemical
equilihiiiun

2. Distinguish between the point of equilibrium and the equilibrium constant.

3. Which of the factois that affect the rate of a reaction also affect (a) the point
of equilibimm (b) the equilibrium constant?

4. Explain the effect of a catalyst upon the point of equilibrium.

5. Write the expression for the equilibrium constant for the following.

(a) 2 SO, 4- O, 2 SO, (b) SO, + % O, SO,
(c)3At2B^2C + D AH = 4- 2 kcal

6. State and illustiate (a) Van't Hoff's principle (b) Le Chatelier’s principle.

7. Define and illustrate (a) homogeneous equilibiium (b) heterogeneous equi-
librium.

8. Why is it unnecessary to know the volume of the reaction mixture in cal-
culating the equilibrium constant for the reaction of Table 9-A?

9. What will be the effect of increased tempeiature upon the point of equilibrium
and the value of the equilibrium constant upon the following gaseous systems
at equilibrium?

(a) H, 4- Cl. ^z± 2HC1 AH =-44 kcal (b) N 2 + CL 2NO AH = f43 kcal.

10. What will be the effect of increased pressure upon the point of equilibrium
and the value of the equilibiium constant upon the following gaseous systems
at equilibrium?

(a) 2 SO, 4- O, 2 SO, (b) N, 4- 0 2 ^± 2 NO

11. The equation for the production of ammonia by the Haber Process is

N,(g) 4- 3 H,(g) z=± 2 NHJg) AH = -22 kcal.

(a) State the conditions of temperature and pressure under which the highest
yield of NH, would be produced (b) industrial practice is to operate at elevated
temperatures where the percent of NIL, produced is relatively low. Why is
this done?

12. (a) Show that for an equilibrium in which the numbers of molecules on both
sides of the balanced equation are equal, changes in total pressure have no
effect on the point of equilibrium (b) what should be the nature of a chemical
equilibrium so that temperature change has no effect on the point of equi-
librium?

13. Prove that when two chemical equations are added, for which the equi-
librium constants are Kh and K^, the equilibrium constant, K, for the result-
ant reaction is K, ’X K 2 .

14. For the equilibrium at 500°C, H,(g) 4- Io(g) 2 HI(g), the equilibrium
constant, K p , is 45.0 (a) What are the units of K p for this reaction? (b) Calculate
K p for 2 HI H, 4- I 2 (c) calculate K p for Vz H 2 4- % I 2 — HI.

15. The value of the equilibrium constant, K p , is 0,042 at 250°C for the equi-
librium, PCI, PCI, 4- CL. Calculate the value of K p for

(a) PCI, 4- CL, ;=± PCI, (b) 2 PCI, 2 PCI, 4- 2 Cl 2 .

16. For the gaseous equilibrium, 3 A 4- B 2 C 4- D, it is found that 0.5 mole
of A remains at equilibrium after starting with two moles of A and B
eaeh in a two liter vessel. Calculate (a) the concentration of A, B, C, and D,
present at equilibrium (b) K c .

124

Chi mu kI equilibrium

17. Two solids, A and B, are m equilibrium, Apv> z=± B (s' If an increase in
pressure causes the complete transformation of A into B, whuh soli d has the
larger density? Explain your answer

18. For N 8 O 4 0?; 2 NO./gb the decree of dissociation is 20 percent at 25°C,

If 0.10 mole of N 2 0 4 is introduced mto one litei vessel calculate pri K c

(b) the total pressure of the equilibrium mixture at 25 C

19. For the gaseous equilibrium. CO 4 - Cl, ±± COOT, AH r= 12 keal. How
would the value of the equilibrium constant and the point of equilibrium be
affected by (a) heating the reaction mi.xtuie ^b) introduction of a catalyst

20. In illustrative Example 6 calculate (a) the percent comeismn of NC into \ T H {
(b) the partial pressure of H 2 at equilibrium.

21. When 2.00 moles of I 2 and 8.00 moles of H, are heated at 500 C in a one
liter container, 3.00 moles of HI are formed at equilibrium Calculate (a! the
number of moles of I 2 and H 2 present at equilibrium (bi the concentration of
I 2 and H 2 (c) K c (d) the total pressure (e) the partial pressure of the H 2 .

22/ For the reaction H 2 4- C0 2 H.O + CO, K p = 1.60 at l,00<rC. Calculate
the number of moles of CO present at equilibrium if (j) two moles of H and
CO each are mixed initially (b) four moles of each me mixed? ( c) If K ?1 is 2.00
at 1,100°C, is the forward reaction exothermic or endothermic* 1

23. For the gaseous reaction at 400°C, N 2 4- 3 H ; 2 NH,, il 1.0 mole N 2
and 2.0 mole H 2 are mixed initially in a two liter container, and if 0 1 mole
of NH 3 is present at equilibrium, calculate (a) K c (b) the total pressure of the
equilibrium mixture.

24. For the reaction, N 2 4- 0 2 2 NO, K p = 2.25 - 10* 4 3t 2,000 C If air

(80 percent N 2 and 20 percent 0 2 ) is heated to 2,000 O, what will be the
percent NO in the equilibrium mixture?

25. Solid ammonium hydiosulfide, NH 4 HS, dissociates as follows;

NH<HS(s) NH s (g) + H.S/g)

When an excess of NH 4 H$ is placed in a previously evacuated container, the
total pressure at equilibrium is 400 mm. (a) Write the expression for the equi-
librium constant (b) calculate the value of the equilibrium constant. HINT:
The pressure of the NH 4 HS is constant and negligibly small, in comparison
to that of the gases present.

26. For the dissociation at 800°C, CaCO/s) CaO ($) -f C0 2 (g) K„ = 167mm.
What is the minimum number of moles of CaCO k that must be placed
in a two liter container at 800 °C in older to establish the foregoing equi-
librium?

The Periodic Law

At the present time there are over one hundred known chemical elements
with the possibility that the techniques of nuclear science may discover,
or rather synthesize, additional elements (Chapter 48). To the ancients only
a dozen or so elements were known and these were among the most inactive
ones, such as copper and gold. These were either found free in nature or
their compounds could be readily decomposed into the free elements. With
the passage of centuries and with the discovery and isolation of more ele-
ments and their compounds, the mass of factual chemical information in-
creased. If one had to study in detail the properties of each substance a
simple book of chemistry would indeed be a veritable encyclopedia. Particu-
larly since the establishment of Dalton’s Atomic Theory, chemists have
looked for some relation or unifying principle concerning the properties of
the elements and their compounds. Fortunately there is such a relationship
so that all. the elements can be classified into a few main groups. Within a
given group the properties of the elements are generally quite similar,
showing only a qualitative gradation in character between successive mem-
bers. This enables us to limit our initial detailed study of chemistry to repre-
sentative members of each such group with the knowledge that other mem-
bers of the group will have properties quite like those of the element studied
in detail.

1. Periodicity and Early Attempts at Classification. If the elements are
listed in the order of increasing atomic weight, starting with hydrogen as
the first, it is found that every element is not completely different in
character from every other element but that the properties of the elements
recur or repeat themselves in cycles. This recurrence in properties results in
families of elements which have similar properties. The ordinal number
of an element in such a tabulation is called the atomic number and is an
important factor in atomic structure (Chapter 11).

Let us list the first seventeen elements in ascending order of atomic weight.
Atomic Number 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17

Element H He Li Be B C N O FNeNaMgAl Si P S Cl

Atomic Weight 1.0 4.0 6.9 9.0 10.8 12 14 16 19 20 23 24 27 28 31 32 35.5

Omitting hydrogen from consideration for the moment, the elements from
helium to fluorine show properties quite distinct from eaeli other Won, how-
ever, an unreactive gas, has properties \ei> much like these of helium, Sodium
the element after neon, has properties similar to those of lithium, the element
after helium, and so on with magnesium and hen Ilium, aluminum and boron,
silicon and carbon, phosphorus and nitrogen, silicon and nxygen, and chlorine
and fluorine. We might now predict that the next element Hoi chlorine

should be like neon and indeed it is. \igon is also an m*Tt gas like neon.

Despite the fact that argon has a larger atomic weight • 3*) than potassium
(39.1), its properties fit in with those of neon and helium, whereas those
of potassium are like sodium and lithium For these two elements, argon and
potassium, to take their proper positions the older of atomic weight must
be inverted. The atomic weight thus cannot bo the controlling factor in
determining the properties of an element, in the next chapter we shall see
that this determining factor is the atomic number

If the elements are rearranged in accordance with this recurrence of
properties the following table is obtained

He Li Be B C X O F

Nc Nrt Mg Al Si r s c:i

Ar K

This periodicity, or recurrence, in properties can lx* illustrated by the
valence of the elements with respect to hydrogen and chlorine, which ascends
from zero to four and then reverts to zero. The valence* towards oxygen,
which increases from zero to seven, is also recurrent. The following table of
compounds, in which the number in parentheses represents the valence of the
first element in the compound’s formula, will make this clear
Valence to H and Cl

He(0) LiCl(l) BeCL(2) BC1.,(3) CC1 4 (4); CH*(4) NH-(3) OH ,(2) FH( 1 )
Valence to Oxygen

He(0) Li 2 0 ( I ) BeO(2) B,0,(3) CO, (4) N,0,(5) F,0(D

Valence to H and Cl

Ne(0) NaCl(l) MgCl,(2) A1C1..(3) Sia i (4);SiH 4 (4)PH l (3) SfF(2^ CIH(l)
Valence to Oxygen

Ne(0) Na,0(l) MgO(2) Al a 0 8 (3) SiO a (4) P 2 0. ( (5)S0 i (fn CbO ; (7>

Valence to H and Cl

Ar(0) KC1(1)

Valence to Oxvgen
Ar(0) KoO(I)

The periodicity in chemical properties is not limited to valence. Physical
properties also show this recurrence. Indeed there arc very' few properties
which are not periodic, for example, radioactivity. An excellent illustration
of periodicity is the variation in atomic volume. In Figure 10.1 advantage
is taken of our present knowledge concerning all the elements, and atomic
volume, or atomic weight/density, is plotted against the atomic number of
the elements. It is apparent that the atomic volume rises to a peak and
then drops to a minimum cyclically. This recurring variation in properties

12S

I'hi’ V* nadir Law

can be generalized m a statement known as the JVmixIk 1 <aw tin * pioperties
of the elements aie periodic functions of the atomic mnnbei 1 his may be
the grandest generalization in all chemisti>.

2. The Periodic Table. The major credit tor the discovery of the Periodic
Law is given to a Russian chemist, Dmitri Mendelejeff In 1869, Memhdejeff
published an arrangement of the elements then known to exist, some 56,
with the periodic concept so clearly and completely developed that his name
has been associated with this classification, although quite independently
Lothar Meyer, a German chemist, armed at the 1 same conclusions the follow-
ing year. Earlier, in 1863, the English chemist, John A. R. New lands had
noted that “the eighth element, starting from a given one, is a kind of
repetition of the first, like the eighth note of an oetaxe in music." 'Hus system
failed when carried further to include the higher elements, and was ndiculed
at the time of its presentation, although yeais later New lands was honored
for the basic truth of his discovci>\

Mendelejeff’s table, brought up to date in xiew of oui present knowledge
of a greater number of elements, is shown in Figuie 10 2, A modern ex*
panded version is given in Figure 10.3 and on the inside back cuxer. The
horizontal rows of the Periodic Table are called periods oi series The verti-
cal rows or columns are called groups. It is in the critical gioups that w r e
find elements of similar, properties, sometimes called families of elements.
The first period consists of two elements, hydrogen { No. 1 ) and helium
(No. 2), which is a noble ( unreactive ) gas as are all elements in Group 0.
The next sixteen elements form two periods of eight each horn lithium
(No. 3) to neon (No. 10), and from sodium (No. Ill to argon (No. IS),
Period four begins with potassium (No. 19), a metal similar to lithium and
sodium, and ends with the noble gas, krypton (No. 36), a series of eighteen
elements. This period is called a long period in distinction to Periods Two
and Three, which are short, having but eight elements each. In this long series
the ten additional elements from scandium (No. 21) to zinc (No. 30) are
called transition elements. Period Five, from rubidium (No. 37) through
xenon (No. 54), is also a long series of eighteen elements containing ten
transition elements, from yttrium (No. 39) through cadmium (No. 48).
Period Six comprising the elements from cesium (No. 55) through radon
(No. 86) is the longest series with a total of thirty-two elements, and in-
cludes transition elements from lanthanum (No. 57) through mercury (No. 80).
Within this series is a special family of fourteen elements from cerium ( No. 58)
through lutecium (No. 71), all of which are placed in one cubicle of the
Periodic Table along with lanthanum and are therefore called the lanthanide
series . An older name for this special family is the rare earth series. The
seventh and final period starts with francium (No. 87) and extends to the
element of highest atomic number known, lawrencium (No. 103). Within
this period is the actinide series which starts with actinium (No. 89). It
is possible that elements of atomic number higher than 103 may be syn-
thesized by the techniques of nuclear chemistry. In this event they would
merely be added on to the present terminus of the Periodic Table.

Vertical groups are numbered from zero through eight, and with the
exception of Groups 0 and VIII, each is divided into two subgroups, A and B.

129

Figure 10.2. Mendelejeff's Periodic Arrangement oi the Elements.
[Modified and brought up to date]

VIA VII A VIII

Figure 10.3. Periodic Classification of the Elements.

The Periodic Law

131

To emphasize this division into subgroups in the original Mendelejeff table
(Figure 10.2) elements in the A subgroup are placed to the left and elements
in the B subgroup to the right. In the extended form (Figure 10.3) sub-
groups are assigned separate columns. The elements in these A and B
families of the same group frequently have little more in common than their
valences. Thus all the elements of Group I exhibit a valence of one in their
compounds but the chemical and physical properties of the elements and
compounds of the alkali metal family. Group IA, are quite unlike those of
the elements and compounds of the copper family, Group IB. Group 0 has
but one family while Group VIII consists of three sets of three elements, one
set in each of the long periods.

It has been mentioned that valence is a periodic property. One character-
istic valence of each element is the number of the group in which it appears.
If the symbol R is used to represent an element, general type formulas for
compounds formed with hydrogen, oxygen, or chlorine are given at the head
of each column of the Table in Figure 10.2. Thus in Group II A, the
formulas RC1 2 and RO indicate that beryllium has a valence of two and will
form the compounds BeCh and BeO. Similar behavior is to be expected from
the other elements in Group II A. In Group V B, arsenic will form the
compounds, AsH ri and As 2 0 5 . In Group VIII the platinum metals, osmium,
iridium, and platinum, do show the rare valence of eight, as in the compound,
0s0 4 . Approximately half the elements exhibit more than one valence but
only the maximum valence is represented by the type formula.

Elements which are generally considered to be metals are to the left
of the zig-zag line in Figure 10.3, whereas the nonmetallic elements are
to the right of this line. Though the elements within a family bear strong
resemblances to each other, there is a gradation in their properties as
we proceed down the group. Thus the boiling points of the elements in
Group 0 increase as the atomic number increases. In the A subgroups, as
a rule, the elements at the top are the least active chemically, the activity
increasing with an increase in atomic number. Thus in Group IA, the alkali
metals, the activity increases as we go from lithium to cesium. However, in
Group VII B, the halogen family, chemical activity decreases with increasing
atomic number; iodine is the least active and fluorine the most active element
of this subgroup.

3. Uses of the Periodic Table. The Periodic Table is of service to
chemists in several ways. First, it furnishes a comprehensive classification
of the elements, enabling us to correlate our knowledge of the properties of
the individual elements and their compounds. We shall use this concept
of the Periodic Table constantly in the succceeding chapters.

Second, with the aid of the Periodic Table the existence of new elements
and their properties has been predicted. Mendelejeff himself made re-
markable predictions concerning, at that time, as yet undiscovered elements.
When he proposed his arrangement, the elements scandium (No. 21), gallium
(No. 31), and germanium (No. 32) were unknown. In order to maintain
the regularity noted in the first part of his table and to present elements
with like properties as a family in a vertical column, it was Mendelejeffs
genius to leave vacant spaces for elements which he said were still un-

132

The Pt nodic Law

discovered. From the position of the vacant spaces Mendelejeff was able to
predict the atomic weights and other properties of the unknown elements.
The discovery of the three elements within twenty \ours of his piediction
and the accuracy with which Mendelejeff foretold their pioportie> served
to justify the validity of the periodic relationship How closely one of
Mendelejeffs predictions, his hypothetical Ekasihcon, camt‘ to the truth
can be appreciated fiom the following table.

Table 10- A

Ekasilicon (E$)

(Predicted in 1871 by Mendelejeff)

Atomic weight, 72
Specific gravity, 5.5
Specific heat, 0,072 cui/g deg
Element will be dirty grav and on cal-
cination will give a white powder of
EsCh.

Element will decompose steam with dif-
ficulty.

Acids will have a slight action, alkalies
no pronounced action.

The action of sodium on Es0 2 or on
EsK 2 F 6 will give the element.

The oxide Es0 2 will be a refractory and
have a sp. gr. 4.7. It will be less
basic than Ti0 2 and Sn0 2 , but greater
than Si0 2 .

The chloride EsC 1 4 will be a liquid with
a boiling point under 100° C and a
sp. gr. 1.9 at 0°C.

t U'rniajuum ( Ge *

(Discovered in 1888 by \\ inkier)

Atomic w eight, 72 6

Specific gnivitv, 5.36

Specific heat, 0 U76 cal. g deg

Coloi of element is gr.n ish white and
on ignition \iekls a white ovule KleO.k

The element docs not decompose liquid
water.

The element is not attached b\ HGl, but
is bv aqua legia Solutions of KOU
have no action, but the element is
oxidized bv fused alkalies.

Gciinanmm is made bv reduction of
Ge ().. with carbon, or of Gofv J\, with
sodium.

The oxide GeCT is a refiactun anti has
a sp. gi. 4.703. It is a vers weak base.

Geimanium chlonde (GeClj) boils at
86 5 C C and has a sp. gr 1.88 at 20 U C.

Mendelejeff also predicted the existence of an element between molybdenum
(No. 42) and ruthenium (No. 44) and it was not until 1947 that the element,
technetium (No. 43), which fit into the vacant space, was prepared by the
techniques of nuclear chemistry.

Finally the periodic system also helped to decide the correct values
for the atomic weights of some of the elements. The failure of some elements
to fall into the group to which they evidently belonged led to the precise
determination of their atomic weights. On the basis of such reasoning erron-
eotis atomic weights assigned to beryllium, cesium, and uranium were
corrected.

Though no elements from atomic numbers 1 to 103 remain to be discovered
we can still use the periodic system much in the same way today. Thus few

The Periodic Law

133

of us will ever see or experiment with the element strontium (No. 38).
With some knowledge of the properties of calcium and barium, and assum-
ing a gradation in properties along the group, we can, however, predict
quite well the nature and magnitude of strontium’s properties.

4. Theoretical Meaning of the Periodic Table. Many interesting questions
are posed by the Periodic Table.

A) Why indeed should there be the phenomenon of periodicity and a
Periodic Table at all?

B) Why should the periods have 2, 8, 8, 18, 18, and 32 elements?

C) Why should the members of a given vertical subgroup have similar
properties? Why should there be a gradation of properties within a family
of elements?

D) Why should there be two subgroups, A and B, within a vertical group?

E) Why are the elements of Group 0 almost completely inert?

F) What is the source of the difference between a metal and a nonmetal?
Why are these two classes not intermingled throughout the table? Why
are they located in different regions of the table?

G) Why do the lanthanide and actinide series of elements fit into a single
cubicle of the Periodic Table?

Furthermore the position of hydrogen in the table is not without fault.
On the basis of its physical properties hydrogen is a nonmetal whereas
some of its chemical properties would classify it in Group I A with the
alkali metals. Also the table emphasizes but one valence of each element
whereas other valences are often more important than the one indicated by
the group in which the element is placed. For example, the important
valence of copper is 2, not 1. To bring some elements into the proper groups
as indicated by their properties, the order of increasing atomic weights has
been reversed in four instances. Argon precedes potassium, cobalt precedes
nickel; tellurium precedes iodine; and thorium precedes protoactinium. It
must be re-emphasized that, though atomic weights and atomic numbers
follow each other closely, the properties of an element are a function solely
6f the atomic number. The subject matter of some of these questions was
at one time called the “defects” of the Periodic Table but logical answers
are now known on the basis of the structures of the individual atoms. Hence
it is to this topic of atomic structure that we must next turn our attention.

QUESTIONS

1. List some properties of elements which are periodic and some which are not.

2. What are some of the advantages of the extended form of the Periodic Table
(Fig. 10.3) over Mendelejeffs version (Fig. 10.2)?

3. Discuss the uses of the Periodic Table.

4. On the basis of the Periodic Table write formulas for the compounds formed
between (a) silicon and chlorine (b) antimony and sulfur (c) oxygen and
germanium (d) astatine and hydrogen (e) lithium and hydrogen.

134

Thv Periodic Law

5. The densities, m g/cm 3 , of the Group I A metals are Li ~ 0.334, Na ~ 0 . 971 ,
K = 0.862; Rb = 1.532; Cs = 1 . 90 . Calculate the atomic volumes of these
metals and decide whether density oi atomic volume shows a greutei regularity.

6. Draw a graph of the boiling points of the Group 0 dements against atomic
number (Table 47-A).

7. Draw a graph of the melting points of the elements, atomic numbers 1 through
20, against atomic number.

8. Explain how each of the following properties is i elated to the position of an
element in the Periodic Table, (a) valence (b) atomic weight (c) dectuc con-
ductivity; (d) formulas of the chloride compounds.

9. Predict the properties of (a) strontium, atomic numbex 38 (b) iodine, atomic
number 53.

Atomic Structure— 1
The Nucleus

So far we have directed our study mainly to the macroscopic properties
of matter directly observable to the experimenter. The fundamental units
of chemical reaction are atoms. For an insight into the nature of chemical
reaction and the chemical bond it is essential that we have an under-
standing of the structure of atoms. The development of the modern theory of
atomic structure affords an excellent example of the growth of scientific
knowledge— of how knowledge grows by standing upon the shoulders of
previous knowledge, and of how many apparently unrelated disciplines ulti-
mately yield a single unifying theory. Probably there is no subject that holds
greater fascination and is more fundamental to a knowledge of chemistry
than the subject of atomic structure.

1. Early Concept of the Atom. Up to the end of the nineteenth century
Dalton's hypothesis, that atoms were indivisible and that all atoms of a given
element were identical in every respect, served well in the development
of chemical knowledge. The hypothesis, strengthened by a growing belief
that the transmutation of elements was only an alchemist's dream, eventually
led to the concept, generally held but not expressly stated, that atoms were
rather like billiard balls and that the atoms of one element were made of
a kind of matter different from that of any other clement. What this matter
was, no one knew! The discovery in 1897 that identical particles called
electrons could be obtained from atoms of different elements, together with
the experimental observation that atoms of one kind of element did change
into atoms of a different element through radioactive emission, showed that
the indivisible billiard ball concept was not in keeping with the experimental
facts.

2. Discovery of the Electron. The first important steps in the develop-
ment of the modern theory of atomic structure were the discovery of the
electron and the determination of its properties. It has been known since about
1853 that gases under very low pressure will conduct electricity when a
high voltage is impressed across the gas. The apparatus for studying this

136

Atomic SfJiiif'in — 1 I'he Xticleus

phenomenon is known as a vacuum discharge tube, 01 Crookes tube, after
Sir William Crookes who did much of the early experimental work m this held.
In principle the vacuum discharge tube is quite similar to the colored electric
signs used in advertising, e.g., the neon tube. It consists of a gas filled tube
into which are sealed two metal plates, or electrodes, usually at opposite
ends. If the tube is evacuated till the gas pressure is extremely low, between
0.01 mm and 0.001 mm of mercury, and if a high electiie potential of
several thousand volts is impressed across the electrodes, an electric dis-
charge occurs (Figure 11.1 a). Streaks of white light apparently emerge from
the negative electrode, or cathode, and move in a straight line towards the

C.lthotlr

A large potential difference is applied to the electrodes of a tube
which contains a gas at low pressure. A current consisting of negative
particles passes through^ the tube from cathode to anode and the
tube glows. The signs and - below and above the tube in (b) are
the poles of the external electric field. The cathode rays move in a
circular path and are deflected to the positive pole.

Figure 11.1. Vacuum Discharge Tube.

Atomic Structure — I: The Nucleus

137

positive electrode, or anode. Where they strike the glass a glow or fluorescence
appears. These streaks or beams of light are called cathode rays and consist
of small particles of matter moving at high velocities. When an electric
or magnetic field is placed around such a beam of particles in a Crookes
tube, the beam is deflected from its straight line path (Figure 11.1 b).
Since the beam is deflected in an electric field, the particles composing the
beam must themselves be electrically charged. Further since the particles
are attracted to the positive pole of the electric field they must be negatively
charged, since opposite charges attract and like charges repel each other.
The particles were called electrons and are the fundamental units of negative
electricity. By observing the deflections of cathode rays and the curvatures
of their paths in magnetic and electric fields of known strengths, the English
physicist. Sir Joseph J. Thomson, was able to determine in 1897 the velocities
of the particles and the ratio of their charge, e, to their mass, m. The value
of this ratio, e/m, is 5.273 X 10 17 e.s.u./g (electrostatic units per gram). This
value of e/m was independent of the metal used as the cathode and of the
nature of the gas in the tube, indicating that the nature of the particles
emitted from various metallic cathodes was the same.

The actual charge of the electron, e, was determined in 1917 by the
American physicist, Robert A. Millikan in his celebrated oil drop experiment.
The presently accepted value is 4.803 X 10~ 10 e.s.u. It is thus possible to
calculate the mass of the electron, m, by dividing e by e/m.

m e iectron

4.803 X 10~ 10 e.s.u.

5.273 X 10 17 ±i±-
g

= 9.108 X 10- 28 g

Since the mass of a hydrogen atom is 1.673 X 10~ 24 gram, the electron has
a mass of about 1/1837 of that of the hydrogen atom. On the atomic weight
scale, an electron has a mass of 0.000548 a.m.u.

3. Discovery of X-rays. In 1895 the German physicist, Wilhelm Roentgen,
studied the greenish glow inside the vacuum discharge tube and discovered
a new type of radiation emanating from the tube. These rays caused certain
minerals to fluoresce and had the ability to penetrate light-tight paper and
to fog a photographic plate wrapped in the paper. The rays originate from
the metal which serves as the target upon which the stream of cathode rays,
or electrons, impinge (Figure 11.2).

A high voltage is impressed across the
terminals of the tube, causing a stream of
electrons to flow toward the target. The
kinetic energy of the electrons striking the
target give rise to X-rays. Modern X-ray
tubes use filaments heated by electricity
as the source of electrons.

Figure 11.2. Early Type of X-ray Tube.

138

^tnutun

■l the Sudan s

The kinetic energ> of the electrons is absorbed b\ the atoms of the target
element and is then emitted as ladiation b> a process to be consulted in
the next chapter. Not knowing what caused these Tit vs and for want of a
better name, Roentgen called them X-rays. The raws arc* not deflected by
magnetic oi electric fields and hence are uncharged but arc* decreased in
intensity by passage through dense objects such as metals 01 hones and
so cast shadows of such materials upon a fluorescent screen 01 photoguphic
film.

4, Electromagnetic Radiation. X-ra\s are radiation of the* same type as
visible light in that both belong to the famih of alrctroma^nctu radiation .
The term electromagnetic arises from the fact that the Scottish physicist
James Clerk Maxwell developed a successful mathematical theory of light
in 1864, in which light was considered to be propagated by varving electric
and magnetic fields, hence an electromagnetic wave*. Foi any wave motion,
the product of die wavelength, X, and the frequency, v. or waxes per second,
equals th.e velocity of the wave.

( 1 )

X cm X v

sec

\eloeit\ of the wave

It is the velocity, 2.998 X 10 10 cm/sec, which characterizes the family of
wave motion known as electromagnetic radiation What differentiates one
kind of radiation from another within a, given famih are the values of X
and v. The difference between X-rays and visible light is that X-rays have
much higher frequencies, approximately 3 X 10 ls sec” 1 as compared with
4 X 10 14 sec' 1 for visible light, and correspondingly lower wavelengths.
For both X-rays and visible light, however, the product of X and v equals
2.998 X 10 10 cm/sec.

The several types of electromagnetic radiation are listed in Table 11-A
and are drawn as a spectrum in Figure 11.3. The distinction between the
types is based upon different characteristics exhibited by waves of different
frequency. Thus X-rays and radio waves pass through material objects which
visible light cannot penetrate. The ranges in wavelength of the radiations
are not separated by sharp dividing lines and may overlap. Any particular
type of radiation can be further subdivided. Thus radio waves include long
waves, broadcast frequencies, short waves, television frequencies, radar, and
microwaves. Each of these subregions has its own distinct properties.

Table ii-A

Electromagnetic Radiations

Tt;pe of Radiation Wavelength (cm) Frequency (sec 1 )

Power Waves Infinite to 3 X lO 6 zero to 1 x I0 4

Radio Waves 3 x 10° to 3 x 1(H 1 x 10 4 to I x ]0 U

Infrared 3 x 1(H to 7.5 X 10~ 5 I X I0 U to 4 x 10 14

Visible 7.5 x 10-» to 4 x 10-* 4 x ID 14 to 7.5 x 10 1

Ultraviolet 4 x 10~ 5 to 1 x 1 <H 7.5 x 10 14 to 3 x I0*»

X-rays 1 x 10-« to I X IO' 10 3 X 10 16 to 3 x 10 20

Gamma Rays 1 x l()-® to 1 X 10- 12 3 X 10*» to 3 * ID 22

Cosmic Rays 1 x 10“ u to zero 3 x 10 21 to infinite

Atomic Structure — h The Nucleus

139

Frequency (see* 1 )

3 x 10 25

3 x 10 21 3 x 10 17

3 x 10 13

3 x 10®

3 x 10 5

1 1

Cosmic

Rays

1 L

r i i

Gamma v Ultra 3

Rays X - rays Violet H
>

1 1 1 1

i i

-§j Infrared

i i

— t r

Short

Wave

Radio

i i

1 1

Long

Wave

Radio

__] 1

10- 15 1(H 3 10- 11 10:® KH 10-« 1CH 1CH 10 10* 10 5 10 7

Wavelength (cm)

Note the wide range of radio waves and the extremely narrow range of visible
light. The shortest waves are cosmic lays of extremely high frequency and energy.

Figure 11.3. Electromagnetic Spectrum.

Another unit frequently used to characterize radiation is the wave number,

A , where X is in centimeters; the wave number has units cm -1 . It is
X

the number of wavelengths in one centimeter and so is proportional to the
frequency.

Electromagnetic radiation is also energy. The energy, E, of a given
type of radiation is directly proportional to its frequency and is given by
the simple expression

(2) E = hv

where h is a universal constant known as the Planck constant, equal to
6.62 X 10~ 27 erg sec. The energy of an X-ray whose frequency is 5.00 X 10 17
sec -1 is

E = 6.62 X 10- 27 erg sec X 5.00 X 10 17 sec- 1 = 3.31 X 10' 9 erg

Thus X-rays, being of higher frequency than visible light, are more energetic;
cosmic rays are the most energetic electromagnetic radiation known. Where
extremely high energies are concerned, a unit of energy frequently used
is the electron volt , ( symbol eV ) . One electron volt is the amount of energy
an electron acquires in passing through a one volt difference of potential,
for example, in the discharge tube of Figure 11.1, and equals 1.602 X 10 -12
erg/electron, or 23,061 cal/mole. Since all forms of energy are interconvertible,

(3) E = Vz mv 2 =. hv = eV

5. Discovery of Radioactivity. Every new discovery stimulates scientific
investigation. One year after the discovery of X-rays, the French physicist
Henri Becquerel, while investigating a supposed relation between fluorescence
and X-rays, placed various fluorescent materials on photographic plates that
were protected from exposure to visible light by being wrapped in heavy
black paper. Visible light cannot penetrate such a wrapping but X-rays or a
similarly energetic radiation can, so that the photographic plate would become
fogged or show a blackening after development in the same manner as a
photographic negative. Most of the substances so treated by Becquerel did

140

Afnmtc r ’<< turr — l /7n* Ytidetis

not affect the plate but compounds of uranium did Vpparentlv these com-
pounds emitted some kind of invisible radiation haunt* properties similar to
those of X-rays. This spontaneous emission of a pcMietiatim* mcikiiinn was called
radioactivity (radiation activity'

Thereafter, Pierre Curie and his wife, Maria SUodouska Ohio, decided
to undertake a study of the phenomenon described In Hecquerel The de-
tails of their labor make a scientific saga and have been deseiibrd m many
popular books but it is sufficient to sa> here that they discovered other
elements which are also radioactive, among them radium Though radium
has a radiation activity perhaps a million times that of uranium, its con-
centration in the mineral pitchblende from which both elements aie extracted
is less than one-millionth of one percent, or one ounce per bundled tons of
pitchblende. Little wonder it is then that the Curies had to process, by most
tedious chemical procedures, literally tons ot pitchblende ore befoie they
ultimately extracted 0.1 gram of a relatively pure radium compound.

6. Types of Rays from Radioactive Material. Further m\ estimation con-
cerning the nature of the mysterious radiation being spontaneously emitted
by naturally radioactive substances such as indium and uranium indicated
that such radiations were really composed of three different types of rays.
When passed through a strong electric or magnetic field, an original single
beam of radiation is split into three components. Figure i 1.4 shows a sample
of radioactive material in a lead container with a small opening As the

M.F. is the magnetic field created by a
magnet whose poles are at right angles
to the path of the rays, and perpendicu-
lar to the plane of the paper. A diagram
similar to this was submitted by Marie
Curie in her doctorate thesis in 1903.

Figure 11.4.

Rays from Radioactive
Material.

Most of tin* lays from the
radioactive substance are
stopped by the lead block.
Only those that are in a
narrow beam lease the block
and enter the magnetic field,
M F One portion of the
beam is deflected m a direc-
tion that indicates it is com-
posed of positively charged
particles {alpha). A second
portion of the beam is un-
deflected {gamma). A third
portion (beta) is deflected
very sharply, and, unlike
the first, or alpha , beam is
spread out.

Atomic Structure — b The Nucleus

141

radiation emerges it passes between the poles of an electromagnet, one
pole above and one below the plane of the paper. The radiation is thereby
split into three beams known as alpha rays , beta rays , and gamma rays.

The alpha rays proved to be composed of doubly positively charged
particles, each of which has a mass approximately four times that of a
hydrogen atom. These particles were shown to be charged ions of the element
helium, He 2 + The beta rays turned out to be negative electrons and the
gamma rays, radiation similar to X-rays but of higher energy (Table 11-B).
Alpha and beta particles are ejected by radioactive material at high velocities.
The initial velocities of alpha particles are of the order of 10,000 to 20,000
miles per second; those of beta particles about 100,000 miles per second,
while gamma rays travel with the velocity of light, 186,000 miles per second.
The penetrating power of alpha particles is slight. They are able to pene-
trate 3-4 cm of air, only a few sheets of paper, and not more than 0.002 cm
of aluminum. The penetrating power of beta rays is about 100 times greater,
whereas gamma rays are highly penetrating and can pass through several
centimeters of lead. The ionizing power of these radiations is in reverse
order to their penetrating power.

Besides the photographic plate there are several other methods for the
detection of radioactivity. Among these are the electroscope, the scintillation
counter, the cloud chamber, the bubble chamber, the ionization chamber, and
the Geiger-Muller counter. In the chapters on Nuclear Chemistry these will
be considered in some detail. It is the emission of particles and rays by
radioactive atoms that makes it possible to follow them through chemical
reactions even though their concentrations are very low. The significance of
radioactivity, insofar as atomic structure is concerned, is that a concept such
as Dalton's indivisible “billiard ball” type of atom is no longer tenable. If
alpha and beta particles are ejected from atoms it is obvious that the atom
must be composed of entities smaller than the atom itself.

7. The Proton and the Neutron. The discovery of negatively charged
rays (electrons) in the cathode ray tube suggested to scientists the possible
existence of positively charged rays emanating from the anode of such a
tube. In 1886 the German physicist Eugene Goldstein, using a Crookes
tube in which the cathode was perforated, discovered such rays passing
through the cathode away from the positively charged anode (Figure 11.5).

Figure 11.5 .
Positive-Ray Tube.

The positive rays formed
in a Crookes tube move
towards the cathode and
can be detected after they
pass through perforations
in the cathode.

Cathode rays Positive rays

142

Uonur ^fnuturi — I rlti \uctvus

These lays are also deflected m magnetic and electnc fields hut in a direction
opposite to the deflection of electrons in the same fields, indicating that they
are streams of positively charged particles, and hence wcte called positive lavs.
Unlike the cathode rays, the natuie of the positive lavs was dependent
upon the gas used to fill the Ciookes tube The largest value of o/in in
deflection experiments was obtained when the tube was filled with hvdiogen.
Assuming that the charge, e, on the pai tides is the same in dilfeient gases,
the mass, m, of the positive iay particles in hvdiogen is less than foi any
other gas. This lightest positive particle is called the proton. The pioton is
the unit of positive charge and has a mass, 1 0072S amu. almost equal to
that of the hydrogen atom.

Another particle of fundamental impoitance m atomic* structure, the
neutron, was discovered in 1932 by the English phvsicist, James Chadwick.
During the bombardment of light elements, notably beryllium, by alpha
particles from a naturally radioactive element, particles aic produced which
are unaffected by electric or magnetic fields. These pa? tides are uncharged,
hence the name neutron, and have a mass onlv slightly gi eater than the
proton, 1.00867 a.m.u. Because it is uncharged the neittion pinch ices no
ionization in its path but has a high penetrating power.

8. The Nuclear Atom. The fundamental structural components of all
atoms are protons, neutrons, and electrons. The mass and charge of these
particles, and other fundamental particles not chrectlv involved m atomic
structure, are given in Table 11-B Knowing the component parts of atoms,
however, does not tell us their structures. The first to piopo.se a definite
structure of atoms was Sir J. J. Thomson in 1904. lie assumed a uniform
positive sphere of electricity in which were imbedded elections, arranged
in such a way as to conform with the periodic ariangement of elements in
the Mendelejeff table. In 1911 a most decisive experiment in the theory of
atomic structure was carried out by the British physicist, Ernest Ruther-
ford. It had been shown earlier by H. Geiger and E. Marsden that high
speed alpha particles could pass through thin metallic foils, though some
were deflected from their course or “scattered/* Because of its relative-
ly large mass and kinetic energy as compared with an electron in an
atom, an alpha particle would not he deflected by a close approach to such
electrons. It is to be expected, however, that alpha particles, themselves
positively charged, would be deflected if they passed near the relatively
heavy positive charges (protons) within the atoms comprising the foil.
Rutherford measured these deflections of the alpha particles from their
initial course. Most of the alpha particles went through the foil undeflected,
an appreciable number were deflected through a small angle, a lesser
number through a larger angle, and a few, about 1 in 8,000, were bent
back in their paths 90° or more from their initial course (Figure 11.6).

These results disproved Thomson’s hypothesis because positive charges
uniformly distributed within an atom would not give the distribution of
alpha particles at the various angles that Rutherford obtained. Rutherford
therefore proposed that the positive charge within an atom was concentrated
into a very small volume, the nucleus of the atom, with the remainder of

Atomic Structure — I: The Nucleus

143

A thin beam of alpha particles, obtained by letting the alpha particles from
a radium source pass through a pinhole m a metallic plate, is directed at a
gold foil. Most of the particles pass through the foil undeflected but many
are deflected through various angles from their original path. A very few are
deflected through an angle greater than 90°

Figure 11.6. Scattering of Alpha Particles By a Gold Foil.

the atom being mainly empty space. Thus most of the alpha particles would
encounter only empty space as they passed through atoms of the foil and
would be undetected from their straight line paths. A small number would
come close enough to the nucleus to be repelled slightly and so be deflected
through a small angle. A still smaller number would pass closer to the
nucleus and be turned through a larger angle, and only a very few would
make “head-on” collisions with the nucleus and be reflected back upon their
original courses (Figure 11.7).

The closer the approach of an alpha particle to the positively charged nucleus,
the greater is the angle through which it is deflected.

Figure 11.7. Deflection of Alpha Particles by the Nucleus of an Atom.

Thus the atom has an exceedingly small, positively charged nucleus
which is surrounded by a complement of electrons that are also very small.
Rutherford calculated that the nucleus of an atom has a diameter of the
order of 10~ 12 cm or about one ten-thousandth of the diameter of an atom.

144

Atomic Strm tun ■ — / 7*/ic \uc!ru$

Tabic 11-B

FuNDAM*.\r\L PUVIK LVS

Particle

Mass

gram

a m u

Charge

Electron

9.108 • 10-'

0 0005 48

-1°

Proton

1.672 1CH ‘

1.00728

Neutron

1 675 < 10-’ 4

] 00867

Positron

9 1 OS v 10-‘ s

0 0005 IS

+ 1

Neutrino

less than that of election, 0

Meson, fx

207 times that of electron

4-1, 0,-1

Meson, tt

273 times that of

electron !

Meson, k

966 times that of

election '

Antiproton

1.672 * 10- 24

1.0072S

-1

Alpha particle* °

6.642 v 10-2*

4 0015

•unit charge is 4.803 X 10 ~ l ° eMl 00 included for otni.wrisou onK

9. Atomic Numbers. In 1914 the English phvsicist, H. G }. Moseley,
determined the magnitude of the positive charge on the nucleus of an atom
by obtaining the X-ray spectra emitted by different metallic elements. A
spectrum consists of the characteristic radiation emitted by a substance and
when photographed or viewed through a spectroscope, appears as a series
of lines or bands arranged in order of wavelength. A spectrum is obtained
when energy is absorbed by a substance and then re-emitted as radiation.
When passed through a prism or a diffraction grating the radiation is dis-
persed into its component frequencies. These are characteristic of an element
and can be used for analytical purposes. Much of our fundamental theoretical
knowledge has come from spectrographic data.

Because the distances between the atoms of solid elements, approximately
10~ 8 cm, are of the same order of magnitude as the wavelengths of X-rays,
a crystal of an element can be used as a grating for the diffraction of an
X-ray beam into its component frequencies as spectral lines. Moseley con-
structed X-ray tubes with anodes of different metals. Moving with large
kinetic energies the electrons from the cathode, upon striking the anode,
excited the atoms therein into emitting their characteristic spectra which
were then recorded upon a photographic plate. Spectra of the different
elements showed the same pattern but corresponding lines for the different

1 Mn

E 1

| Co

Figure 11.8. Moseley's X-ray
Spectra.

The wavelengths of the corres-
ponding lines increase progressive-
ly with increase in the atomic
weight of the element.

Wavelength

Atomic Structure — 1; The Nucleus

145

elements were shifted slightly in frequency. In Figure 11.8 just two, for
simplicity, of the many lines in a spectrum are shown for the elements from
titanium to copper. These lines vary in wavelength progressively with the
atomic weight of the element used as anode, or more accurately with the
numerical position of the element in the Periodic Table. When the square
roots of the frequencies of the corresponding lines in each spectrum are
plotted against the value of such position (H — 1, He = 2; Li = 3 . . .
Fe = 26), a linear relationship is disclosed (Figure 11.9).

Atomic

number

VHP requency X ]0“ 8

The square roots of the frequencies of the corresponding lines in the X-ray
spectra of the elements are directly proportional to atomic numbers of the*
elements.

Figure 11.9. Illustration of Moseley's Law.

Moseley interpreted his findings as follows; “We have here a proof that
there is in the atom a fundamental quantity which increases by regular steps
as we pass from one element to the next. This quantity can only be the charge
on the central positive nucleus.” The atomic number is thus identified as
being equal to the number of unit positive charges on the atomic nucleus.

10. Nuclear Structure. The internal structure of the nucleus is still some-
what of a mystery. At one time it was thought that the nucleus contained
protons and electrons but it was difficult to explain the stability of such a
configuration when closely packed. Some thirty more or less stable “elemen-
tary particles” have so far been identified, but for our present purposes it
will be sufficient to consider that the nucleus of an atom is composed of pro-
tons and neutrons only, held together by forces not yet definitively known.
These essential components of nuclei, protons and neutrons, are often referred
to by the general term nucleon.

The number of protons in the nucleus is equal to the atomic number and
is also equal to the charge, Z, of the nucleus. This number of protons alone is in-
sufficient to account for the weight of the nucleus (except in the case of hydro-
gen) so that additional particles which contribute mass but not charge to the
nucleus, the neutrons, must be present therein. The sum of the numbers of

146

.Wuwur bttmtuj' — I l hi 1 Xuclcui

protons and neutrons m an atomic nucleus is a wholr number, defined as
the mass number, A. Hence, the number ot neutrons is equal to the mass
number of the atom, A, less the atomic number Z. In sumnurs :

A = Mass number, a whole number which is the sum of the protons
and neutrons

Z — N umbei of protons, the atomic numbei, the positive ehaige of
the nucleus

A - Z = Number of neutrons

The composition of the nuclei of atoms of lithium, iron, and uranium are
shown in Figure 11.10.

Lithium nucleus lion nucleus Uranium nucleus

A = 7; Z = 3 A = 56, Z ~ 26 A 238; Z = 92

Figure 11.10 . Atomic Nuclei.

For elements ot low atomic weight, the numbers of protons and neutrons
are approximately equal. With increasing atomic weight, the ratio of neutrons
to protons increases to a maximum value of about 1.5.

11. Mass Spectrograph. It was stated earlier that electric and magnetic
deflection experiments can determine the ratio of charge to mass of charged
particles. Such analysis of positive rays has enabled the measuiement of
atomic masses to an accuracy of one part in a hundred thousand. In passing
through an electric or magnetic field the charged particles, or ions, of a
positive ray are deflected in proportion to their charges, masses, and velocities.
Experimentally, positive ions produced by ionizing a gas or a metal vapor
with an electron gun, are accelerated to a uniform velocity and collimated
by passage through a series of slits. The resulting beam of positive ions is
then passed through a magnetic field. The ions traverse a circular path, and
for those ions whose charges are the same, the curvature of the path will
depend solely upon the mass of the particle. Where the ions then strike a
photographic plate, this apparatus produces lines which resemble a spectrum
and is therefore known as a mass spectrograph. Each line represents particles
of a specific mass and the relative intensities of the lines give the relative
amounts of the several species present in the positive ray. The photographic
plate can be calibrated through the use of elements of known atomic mass
as standards. One type of mass spectrograph is shown in Figure 11.11.
If the photographic plate is replaced by a metallic plate connected to an
electrometer or other device for measuring the ion current, and if the ap-
paratus is modified to measure the current due to ions of a specific mass,
it is known as a mass spectrometer

Atomic Structure — I The Nucleus

147

12. Isotopes. Since the masses of both the proton and the neutron are
very close to unity, and since the contribution of any electrons to the mass
of an atom is almost negligible, atomic weights should be very close to

Figure 11.11. A Mass Spectrograph.

X = Ion source

S a and S 2 = Collimating slits

P = Photographic plate

Mj and M 2 are positions on the plate at which
ions of different mass strike The mass of the
ions striking at M, i^ less than the mass of
the ions at M 2 .

whole numbers. Indeed nearly half the elements do have atomic weights
which differ by no more than 0.1 from integral values. Some, though, such
as neon (atomic weight = 20.2) and chlorine (atomic weight = 35.5) have
atomic weights which differ appreciably from whole numbers. When chemi-
cally pure neon is examined in the mass spectrograph, two lines are obtained
on the photographic plate, one for an. atomic mass of 20.00 and the other
for a mass of 22.00. No line is obtained for an atomic mass of 20.2. A measure-
ment of the relative densities, or photographic blackness, of the two lines
show the particles of mass 20.00 to be present in natural neon in the ratio
of 9 to 1 as compared with the particles of mass 22.00. Evidently neon
consists of two kinds of atoms, one with a mass of 20.00 to the extent of
90 percent and the other with a mass of 22.00 to the extent of 10 percent.
The weighted average of these two gives an atomic weight of 20.2 to neon
as found in nature.

These kinds of atoms, which differ solely in mass but have the same
atomic number, are called isotopes. Alternatively, the macs number can be
defined as the whole number nearest the weight of an isotopic species,
whereas the atomic weight is the weighted average of the masses of the
isotopes found in nature. Inasmuch as the atomic numbers of the isotopes
of an element are the same, the difference between the two kinds of atoms
is due to a different number of neutrons in the nuclei of the isotopes. In
the case of neon, the lighter isotope has 10 neutrons and the heavier one
12 neutrons. Most of the elements, especially those of even atomic number,
have isotopes. The element tin has as many as ten. Because of the relatively
large difference in their masses, a specially important pair of isotopes are
those of hydrogen. Atoms ot ordinary hydrogen have a single proton as the
nucleus and hence a mass of one. An isotope of mass two is known as
deuterium or heavy hydrogen and has a nucleus, the deuteron, which con-

148

Atoniu ^t:u< t'i;<

■/ / In \ucleus

sisis of one proton and one neution The irlutixt* abundance oi the two
isotopes is 99.985 percent for that of ones one and OOlo percent for
that of mass two. A third isotope, tiitium, of ma^ thiee. does not oc-
cur, naturally but is a pioduet of mielem tiansnmtatinn and is itself
radioactive. Because the isotopes of an element have almost identical chemical
properties their separation is difficult. Such separations au* considered in
Chapter 49. The existence of atoms haxmg the same mass numbers hut
different atomic numbers is also possible. Such atoms are called iwhars and
are different chemical species from each other.

To specify the particular isotope of an element when wilting foinmlas or
equations, a symbolism has been adopted m which the mass number is
written as a superscript after the chemical symbol, and the atomic number
as a subscript before the symbol. For example

50 Sn 118 and ^Sn 1 - 2 aie isotopes
and soSn 1 - and , 2 Te 122 are isnbuis

Actually it is redundant to write both the chemical symbol and the atomic
number subscript, since a given element can have but one atomic number,
so that on occasion the subscript is omitted.

13. Atomic Weight Scales. Prior to the adoption in 1901 of the atomic
weight scale based upon the mass of the carbon- 12 isotope, thrie were two
scales of atomic weights. The chemical atomic weight scale was based on
the arbitrary assignment of the value sixteen to natuially occurring oxygen.
This was done long before the notion of isotopes was conceived. However,
mass spectrographic analysis indicated that oxygen contained three isotopes:
99.76% of 0 1H , 0.04% of O 17 , and 0.20 % of () ls . Physicists therefore adopted a
scale which was based upon the assignment of the value sixteen to only
the most abundant isotope of oxygen, On this scale the atomic weight of
natural oxygen, the weighted mean of the three isotopes, was 16.0044
(0.9976 X 16.0000) + (0.0004 X 17.0045) (0.0020 X 18.0037) - 16.0044.

Atomic weights on the physical scale were thus 16.0044/16.0000, or 1.000275
times greater than those on the chemical atomic weight scale. This dis-
crepancy has been resolved by the adoption of a single unified scale of
atomic weights based upon the C 12 isotope, whose weight has been assigned
the integral value twelve.

QUESTIONS

1. Discuss the discovery of the electron. List the properties of the electron.

2. What is meant by electromagnetic radiation?

3. How were X-rays discovered? What is the relation between X-rays and other
electromagnetic radiations?

4. Define nucleon, proton, and neutron. List the properties of protons and
neutrons.

5. What was the contribution of radioactivity' to the concept of atomic structure?

6. Identify the types of rays given off by naturally radioactive substances.

Atomic Structure — T The Nucleus

149

7. Draw a diagram showing how a single beam emitted by a naturally radioactive
substance would be dispersed in an electric field.

8. Discuss the experimental evidence and the reasoning which led to the concept
of the nucleus.

9. What was the contribution of Moseley to the modern concept of the atom?

10. What are positive rays and how can they be produced?

31. What is a mass spectrograph? How does it separate particles of different
masses?

12. Define isotope. What is the structural reason for the existence of isotopes?

13. Distinguish between (a) isotope and isobar (b) mass number and atomic weight.

14. What is the nuclear structure of each of the following atoms: (a) vanadium,
aaV B1 (b) bromine, 35 Br 80 (c) tin, 50 Sn 118 (d) lead, 82 Pb 206 ?

15. What are the nuclear structures of the isotopes of (a) calcium, aoCa 40 and 20 Ca 44
(b) zmc, 30 Zn 64 and 30 Zn 6c (c) mercury, 80 Hg 2O ° and 80 Hg 201 ?

16. Calculate the atomic weights of the elements below for which the isotopic
compositions are given.

Element

Mass Number

Isotopic Weight

Abundance

Copper

Sulfur

31.972

95.06%

32.972-

0.74%

33.969

4.20%

17. How were quantitative stoichiometric calculations affected by the adoption
of the atomic weight scale based upon C 12 ?

18. Calculate the velocity and kinetic energy, in ergs and electron volts, of an
electron equivalent to radiation having an energy of 1 X 1(H erg.

19. Calculate the energy, in ergs and in electron volts of (a) an alpha particle having
a velocity of 10,000 mile/ sec (b) a beta particle having a velocity of 100,000
mile/ sec.

20. What is the wavelength of radiation whose energy is one electron volt? What

type of ladiation would this be? Ans : 1.24 x 10~ 4 cm

Atomic Structure— II
Electronic Structure

1. The Neutral Atom. We have seen that the atomic nucleus is posi-
tively charged, the magnitude of such charge being equal to the atomic
number, Z. Inasmuch as a complete atom is electrically uncharged or neu-
tral, there must also be in the atom negathe charges in such number as
to equal the positive charge of the nucleus. The unit of negative charge is
the electron. The number of electrons in an atom is thus also equal to the
atomic number, Z. For example, the element oxygen whose atomic number
is eight has eight positive charges, or protons, in the nucleus and eight
electrons external to the nucleus. Table 12- A summarizes the structures
of the first ten elements of the Periodic Table.

Table 12-A

Structure of Atoms of the First Ten Elements

Element

Atomic

number

Mem

number

Protons m
nucleus

! Neutrons in
nucleus

f . -

Electrons

A-Z

Hydrogen

Helium

Lithium

Beryllium

Boron

Carbon

Nitrogen

7 i

Oxygen

8 ;

Fluorine

Neon

2. Bohr Theory. A knowledge of the total number of electrons within
an atom, however, gives no information concerning any order which the
electrons may have. Much of the theory concerning the arrangement of

Atomic Structure — II: Electronic Structure

151

the electrons had its origin and development during the years 1913 to 1926.
The subject is by no means closed for much that is new is being contributed
to the theory of atomic structure even today.

In the early 1900’s there were suggestions that the electrons were dis-
posed about the nucleus much like a planetary system. Whereas such an
array appears logically plausible it is contrary to the principles of electro-
magnetic theory. Classical electrodynamics can prove that a charged particle
which is accelerated must radiate energy. An electron, even though mov-
ing at constant speed in a circular orbit is being accelerated since its direc-
tion is constantly changing, and so should emit energy. With a lesser energy,
the radius of the electron’s orbit will decrease and the electron should ulti-
mately spiral into the nucleus. Thus classical theory dictates that a planetary
configuration should not be stable. Nevertheless, in 1913, the Danish physi-
cist, Niels Bohr, at the time one of Rutherford’s research students, proposed
that, contrary to classical theory, the electron could exist in a stationary state
of constant energy indefinitely without emitting energy. Bohr visualized
that the electron could move in circular orbits or shells about the nucleus
and in each orbit the electron had a definite and fixed value of energy. The
crux of the Bohr theory is that electrons could not assume any arbitrary
value of energy but that only definite values of energy could be had by
electrons, that is, the electron could exist only in discrete energy levels.
This meant that the electron could be only in specific orbits or shells.
This concept of definite energy levels, as opposed to a continuum of energy
values which a system might have, is the heart of modern quantum theory.
In quantum theory, energy, like matter, has been made discontinuous. Each
atom has its own specific scheme of electronic energy levels in which an
electron may exist. An electron cannot have a value of energy intermediate
between these energy levels. Its energy is said to be quantized.

The symbol n is used to designate a specific energy level or orbit; n is
called the principal quantum number and may have only integral values, the
lowest being one. Thus, n = 1 refers to the lowest electronic energy level
of an atom, or the orbit nearest the nucleus; n = 2 refers to the second
lowest energy level, etc. In the Bohr theory the energy of an electron
varies with the value of n only. In an alternate nomenclature, the lowest
energy level is called the K level (or orbit or shell); the second, L; the
third, M; the fourth, N, and so on alphabetically. These letters vtere chosen
because it will be seen presently that the Bohr theory was capable of ex-
plaining the origin of spectra and' in X-ray spectra a nomenclature had
already been developed wherein certain series of lines were called the K
series, the L series, etc. In Figure 12.1a there is shown a -schematic energy
level diagram, while Figure 12.1b is the corresponding orbitlike viewpoint.
By assuming that the angular momentum of an electron in its orbit was also
quantized, and by equating the force of attraction between the nucleus
of a hydrogen atom and the orbital electron with the centrifugal force on
the electron, Bohr was able to calculate the energy of the electron in its
orbit and also the radius of the orbit (page 166). The radius of the first
Bohr orbit in hydrogen is 0.529 X 10* 8 cm, or 0.529 A. (A = Angstrom
unit; 1 Angstrom unit = I X 10* 8 cm.)

"'»•!< ^tincture

152

A tomu Sfiw r?.j, — i l l

■ f\ n -- 1

M — K n 3

1. e a :

E n i

(a)

Schematic Energy Level Diagram
E,, K V. , c.ik! K, ate the fivrtl oneigv
levoK c jir'.'tpoii'lnnr to principal quantum
numhei . 1, 2, A, and A, «nd tlit* K, L, M,
,>nd N shcllv lopectiwlv

,b)

OrDitiike Representation ot an Atom
(J) t m*ijN the nut it'ii'. The fixed
entiev l«*\ih » m he MMiali/etl as \ uvular
mbits LaTi th;*. was modified to mciuclc
elhptkid «uiu».

Figure 12.1

3. Quantum Mechanics. To be able to pmdict the hituic path that a
material object will take it is necessary to know at a giwu time its position
and its momentum, that is, the product of its mass and velocity. A basic
principle of classical mechanics is that it is theoretically possible to measure
simultaneously both the position and the momentum of an object with in-
finite exactness, and any inability to do so piacticallv was asenbed to the
limitations of the measuring instrument. In H126, the German physicist
Werner Heisenberg showed that this was not the ease but that then* is a
limit to the accuracy of measurement. The very act n? measurement dis-
turbs the system being measured and introduces an uncertainty m the
final result. The uncertainty is extremely small, of the order of Planck's
constant. When we deal with macroscopic objects the error in our measure-
ments due to Heisenberg’s Uncertainty Principle is negligible. But in deal-
ing with subatomic particles whose properties are of the same order of
magnitude as Planck’s constant, the crroi in measurement can be relatively
great, Heisenbergs Uncertainty Principle can be expie.ssed mathematically
by

(1) Ax * Ap ^ h

The minimum er*G r in the measurement of the position, x, of an object
multiplied by the minimum error in the measurement of its momentum, p,
can not be less than Planck's constant, h. Thus both the position and mo-
mentum cannot simultaneously be measured with absolute exactness. Indeed
a greater exactness in the measurement of one is offset by an increased
uncertainty in the measurement of the other. Because of Heisenbergs Un-
certainty Principle an orbital electron cannot be treated as a minute particle

Atomic Structure — 11 Elect) owe Structure

153

with an exactly known position and momentum at a given time. In dealing
with subatomic particles the principles of classical mechanics fail.

Earlier, in 1923, the French physicist Louis de Broglie had suggested
that a particle of matter could be associated with a “matter wave” whose
wavelength, X, was given by

^ ^ __ h _ _h_ where p is the momentum of a particle

mv p having a mass m and a velocity v.

Three years later the German physicist Erwin Schrodinger developed the
mathematical theory known as quantum mechanics or wave mechanics
which united the particle and the wave aspects of matter. The mathematics
is complex, involving the solution of second order differential equations and
will not be considered here but the conclusions of the theory are pertinent.
In quantum mechanics the electrons are not restricted to fixed orbits. The
fixed orbits of the Bohr theory are replaced by probabilities of finding the
electron at any given position. Whatever the electron may be and whatever
it may be doing, it is visualized as creating a diffuse cloud-like electric
field whose intensity varies with distance from the nucleus. An electron in
a K shell creates a three dimensional spherical cloud of electricity. The
intensity of the electric field at any point, or the electron cloud density,
is a measure of finding the electron at that point. In theory an electron
has a finite probability of being at any point from the nucleus to infinity.
However, maximum intensity of the electric field calculated for the hydro-

Distance from the nucleus, A
(a)

The intensity of the electron field, or the
probability of finding the hydrogen electron
at a given point, as calculated By quantum
mechanics shows a maximum at 0.529 A> the
radius of the first Bohr orbit.

(b)

The electron cloud density of
the electron in a hydrogen atom,
or of an electron in the K shell of
any atom, is spherically symmetri-
cal. The nucleus is at the core of
the sphere.

Figure 12.2. Electron Cloud Density.

154

\ioinii t'timturi — It litttumu Sow turc

gen electron occurs at the same distance when 1 Bohi phu cd the electron
orbit, 0.529 A, as shown in Figure 12 2. In this sense the Bohr theon and
Schrodingers quantum mechanics are in agi cement, in mam cases, paitieu-
larly in the drawing of diagrams, we shall use the p.uticie new point ol the
electron, but only for simplicity of visualization, alwavs realizing that such
a picture is not reality

4. Quantum Numbers, A solution of Schrodingers equation is known
as a wave function and is itself a mathematical equation. The \va\e func-
tion for an electron is called an orbital An orbital gi\ es the spatial distri-
bution of the electron cloud density. Alternative!) it can he considered a
state of the electron. To specify the state of an electron m a given oibital,
four quantum numbers are necessary. These are represented In the sym-
bols n s l m , and s.

n = the principal quantum number. It can take on onl\ intcgial values,
the lowest being 1. Thus a value of n - 2 would refer to the second lowest
principal electronic energy level.

I ~ the azimuthal quantum number. It may have integral \ulm*s from
zero up to a maximum of ( n — 1). If n = 1, / can he onlv 0; d n 2 , the
possible values of / are 0 and 1; if n ~ 3, / can he 0, 1, oi 2. Electrons
are designated as s, p, cl, or f electrons, a nomenc lature which also origi-
nated in spectroscopy, according to whether the value of ! is 0, 1, 2, or 3,
respectively. For example, a p electron would be one for which 1 1. For

a given value of n , the energy of a given state increases with an increase
in the value of l. Thus in increasing order of energy the orbitals are v, p,
d, and f.

m = the magnetic quantum number. It may hin t* integral values rang-
ing from +/ to —l. Thus, if l = 2, the possible values of m art* f 2, + 1» 0,
— 1, and —2. In all, a total of (2 1 *f 1) values of m are possible.

s = the spin quantum number. This is distinct from and not to be
confused with the symbol s which indicates a value of / -- 0. The spin
quantum number may have one of two values only, ! l 2 or h».

In the Bohr theory, the principal quantum number n represented the
orbit in which the electron revolved and l described the' angular momentum
of the electron in its orbit. The energy of the electron was specified by its
values of n and l. The quantum number m was introduced to account for
the splitting of spectral lines in a magnetic field and was supposed to
indicate different orientations of the orbit in space. In quantum mechanics
the quantum numbers n, l, and m are merely integral coefficients in a
mathematical equation; the value of l determines the shape of an orbitals
electron cloud (page 161) and m its spatial orientation in a magnetic field.
The spin quantum number, s , was introduced to explain the existence of
certain pairs of spectral lines. The electron behaves as if it were spinning
about its own axis in one direction or the other, denoted by a ^ or -
sign. By specifying its four quantum numbers the quantum state of an
electron is fully described.

5. Pauli Exclusion Principle. In 1925 the Austrian physicist, Wolfgang
Pauli, proposed the Exclusion Principle that no two electrons in a given
atom or molecule can have all four quantum numbers the same, that is,

Atomic Structure— ll. Electronic Structure

155

two electrons cannot exist in the same quantum state. Though this principle
has not yet been proved theoretically it accounts for the arrangement of
the Periodic Table.

With the aid of the Pauli Exclusion Principle and the rules defining the
quantum numbers let us calculate the number of possible quantum states
available to electrons. Where n = 1, l must be 0, m must be 0, but the
spin quantum number, $, can be 4 % or — Only two quantum states
are possible when n = 1 as tabulated below.

n l m s

State I 10 0 4"%

State II 10 0-%

Since the Pauli Exclusion Principle assigns a single electron to each state,

the maximum number of electrons possible, where n = 1, is two. In other
words, the first principal energy level (or the first orbit or the K shell)
can hold no more than two electrons. Since l = 0 for both of these elec-
trons, both are s electrons.

When n = 2, l may be 0 or 1. Again if l = 0, m must be 0 and s can
be either -f % or — It is evident that the maximum number of electron
states where l = 0 is two and is independent of the value of the principal
quantum number, n. If l = 1, m can be +1, 0, or —1. For each of these
m values, there can be two s values, 4-% and — Thus for l = 1, there
can be a total of six electron states, or six p electrons. This maximum
value of six p electrons for n = 1 is also independent of the value of n.
All told, for n = 2, eight states are possible, as shown below, and hence
eight electrons complete the second orbit, or the L shell of electrons.

+% / s electrons

-%j (1 = 0)

+ 1

+% \

+ 1

—Vi I

+% V p electrons

-%( (1=1)

-1

—Vi j

Where n = 3, l may be 0, 1, or 2. The cases for l = 0 and l = 1 have
already been considered. When Z = 2, the possible m values are 42, 4-1,
0, —1, and — 2. With two spin values for each value of m, the number
of states possible when l = 2 is ten, or there can be a maximum of ten d
electrons. The total number of states for n — 3 is (10 4- 6 4- 2) or 18.

Where n ;= 4, an analogous calculation results in 32 for the total number
of possible states. This sequence of 2, 8, 18, and 32 for the number of elec-
tron states is the reason for the sequence of 2, 8, 18, and 32 for the
number of elements in successive periods of the Periodic Table.

V.'.m*

Slrurtvrr

6 Summary. in iffeet tin- values wt / subdivide the ;»iiuu:m! quantum
level designated by n into subshells winch an* the orbit. ih Vn\ mbiial can
contain no more than a pair of elections 1 best pail'. of dev. irons have the
same values of m /, and m but opposite values of the spin quaiPum number,
s. For any given value oi tie 4 pmu ipul quantum nnmbt i ». thru cam he only
one s orbital and thus only two v elections in that quantum state Sumiarly
there can be no more* than three 4 p orbitals oi six p eh i linos no mote than
five d orbitals, or ten d electrons; and no more than seven I orbitals, or
fourteen / electrons. The maximum number of different lands of orbitals,
$, p, d , or /, associated with a principal quantum number u. is equal to
the value of n The maximum nunihet of all orbitals e equ.il to *r and
since each orbital can contain two elections the mavuumn mi ml mm of elec-
trons in a principal quantum level is 2/r. Thus a complete h shell [n 1)

consist*: of but one s orbital with a pair of electrons The l she!! [n 2)

can contain two kinds of orbitals, v and p, with two and six ele< fours, re-
spectively, for a total of eight. The \f shell (n - 3 can have p, and d

orbitals containing 2, 6, and 10 electrons, respectively, or < ighteen electrons
in all. The N shell (n = 4) can contain foui kinds of nilnt*ds v p. tL and /,
with 2, 6, 10, and 14 electrons, respectively for a total of ihufv two electrons.

A symbolism has been devised to indicate the number and kmd of elec-
trons which may be in the various energy states oi an atom 'Hie principal
quantum number is written first and tin » the letter, ,v, p, J, or /, to specify
the value of /. A superscript is added to denote the actual number o* elec-
trons in the specified state, Tims 2 p i refers to four elections with values
of n — 2 and Z = L or four p electrons in the L shell. This symbolism
uniquely defines the energy state of the four electrons. The distribution
of the orbital electrons is summarized in Table 12-B.

Table 12-B

Distribution of Electrons

Shell

Principal
Quantum
Number, n

Azimuthal
Quantum
Number , l

Orbitals
( Subshells )

Maximum
Number of
Electrons

Complete
( hbital

1.5-

*1 g

2.5*

6 i 8

1 2

2 j

as 2

3 p

6 > 18

3p«

10 )

3d 10

2 \

4s s

6 32

10 ( 32

4pC

4 d

4 d U)

14 )

4 fu

Atomic Structure — II Elccttonic Strut turn

157

7. Electron Configurations of the Atoms. The simplest atom, the hy-
drogen atom, has a single electron in the lowest possible energy level, the
Is. Starting from this point, the process of constructing the electronic con-
figurations of the other atoms involves the insertion of additional electrons,
as we proceed across the Pei iodic Table from atom to atom into the un-
occupied quantum state of the lowest energy. In this way the orbitals are
completed and with the filling of all the orbitals of a principal quantum level,
a new shell must be started. In the order of their increasing energy and hence
the order in which they will be filled, the orbitals are:

Is, 2s, 2p , 3s, 3 p, 4s, 3d, 4 p, 5 s. Ad, 5p, 6s, 4f, 5 d, 6p, 5/, 6 d

This progression is shown in Figure 12.3a together with a simple mnemonic
for it in Figure 12.3b.

The sequence in which the orbitals are filled
gives rise to other limitations for the electronic
configurations of atoms. These are:

1. a. An outermost or highest energy level can
contain no more than eight electrons (except
for the K shell whose maximum is two).

b. A next-to-the-outeimost shell can contain
no more than eighteen electrons

2. An outermost level can contain no more than
two electrons if the next-to-the-outermost shell
has fewer than its allowable maximum of 2 n 2
electrons. Based on current spectrographic and
chemical data, the electronic configurations of
the elements are given in Table 12-C.

6s-

6(1

5s-

5 d

4 p

3 8

3 P

2s —

2 P

(a)

The orbitals are filled in order
along the diagonals from right to
left.

(b)

Figure 12.3. Energy Level Diagram of the Orbitals.

158

Atomic Structure— II- Electronic Structure

Table 12-C

Electron Structures of Single Atoms of the Elements

tructurc * — II Electronic Structure

159

Table 12-G (continued)

K> lo to K> IO to K) to K) to ro lo to to l<j to h-

160

Structure.

v/< r.r.c tti fur. — // Ehctn-nic

The first element, hydrogen { / I*, has but one olet Iron m the low-

est principal quantum level, the k she!!. The electron is an .v electron and
the electronic configura? ion ior the hydrogen atom would he represented
by the symbol lx 1 . The 1 » c \t civim-nt, helium { 7 2 1 has two electrons

in the K shell and both are.x ehetrons- its electronic coni iguruuun is Is-.
With tlies e two electrons the K shell b tilled. Helium is an nmn gas and
completes the first period of the Periodic Table.

With the element after helium a new principal quantum level, the next
lowest in energy or the L shell, must he started. Urns lithium (/ - 3),

with three electrons, has its K shell filled with two Is electrons and the
third electron is in the second slit'll as a 'Is electron. Lithiums configura-
tion is Is* 2 2s 1 . In beryllium ('A 1 ) the ,v orbital of the second shell is

completed and the structure is Lv~ 2w. Xote that the addition o i the super-
scripts gives the total number of electrons in an atom and so also its
atomic number, Z. Since the 2s orbital is completed in beryllium the next
element, boron (Z 5) must start a 2/> orbital, l.v 2 2.v‘ J 2/>\ There can be
three 2p orbitals so that a total ol six such electrons can be accommodated.
This takes us through the element neon (Z 10) which has a complete L
shell and a configuration, lx 2 2s 2 2p n . Like helium neon is an inert gas
and also the terminal element of a series in the Periodic 'fable.

Thereafter the filling of the 3.v and the 3 p proceeds regularly from
sodium (Is 2 2s 2 2 /; <J 3s l ) to the inert gas argon (lx 2 2v 2 2/f 3s 2 3 n n ), the
element which completes the third series of the Periodic Table. Since the
3d orbitals are higher in energy than is the Tv orbital, the latter is filled
first in potassium and calcium, (lx 2 2s 2 2/d 3x 2 3/d 4x 2 ) . The five* 3d orbitals
can hold ten electrons and these are then filled from scandium { i.v 2 2s 2 2 p n
3s 2 3 p e 3d 1 4s 2 ) to zinc (Is* 2s 2 2 /d 3v 2 3/d 3d 1 " Tv 2 ). Thereafter the 4p
orbitals are occupied through the inert gas krypton ( lx 2 2v 2 2/d fix 2 3/d 3d 10
4-s* 2 4/d) and the end of the fourth period of the Periodic Table.

Upon occasion this regular progression must be modified, as in the
cases of chromium and copper, to conform with experimental evidence,
primarily spectroscopic. Thus chromium is ( lx 2 2x 2 2/d 3x 2 3/d 3 d : * 4.x 1 ) rather

than (lx 2 3d 4 4x 2 ). Apparently the half completed 3d orbital is a more

stable configuration. But in general the filling of the orbitals of lowest
energy and the principal energy levels proceeds in an orderly manner to
the last element of the Periodic Table. In some cases, particularly the
lanthanide and the actinide series, the electron configurations are not
known with complete certainty. Some of these elements have been pre-
pared only in microscopic quantities, insufficient for definitive experimental
work.

8. Electron Configuration and Atomic Properties. An electron within
an orbital creates an electron field with a definite shape and spatial orien-
tation. The field of an x electron is spherical, that is, the electric charge
density produced by the electron has the form of a spherical cloud with
a maximum intensity at a fixed distance from the central nucleus. Such
an electron field has been drawn in Figure 12.2b. The electron fields of p
electrons have a pattern which is dumbbell-shaped. Indeed the three p
orbitals are spatially oriented at right angles to each other in the same

Atomic Structure — 11: Elect! onic Structure

161

manner a s are the x, y, and z axes of rectangular coordinates. The p or-
bitals are therefore distinguished from one another as the p v , p y , and p 2
orbitals. These are shown in Figure 12.4 The fields of d and f electrons
have lobes which are also spatially oriented but in a more complex fashion
than are the p orbitals. We shall see in Chapter 13 that the spatial orienta-
tion of orbitals is a major factor in determining the geometric shape of a
molecule and the orientation of the chemical bonds comprising it.

Figure 12 A. Spatial Orientation of p Orbitals.

Where more than one orbital of the same type is available in a prin-
cipal quantum level, e.g., the three p orbitals in the L shell, the incoming
electrons distribute themselves among the several orbitals, one in each, so
as to retain the same, or parallel, values of spin, all + or all — . In boron,
the single p electron can be in the ;k, the p >n or the p z orbital. But in
nitrogen, (Is 2 2 s 2 2 p 3 ), each of the three possible 2 p orbitals has one elec-
tron in it, each electron having the same spin value. Of the 2 p 3 electrons,
one electron is in the p x , the p y , and the p z orbitals so that the electron
configuration of nitrogen could be written more accurately as (Is 2 2s 2 2 p* 1
2p y l 2 Pz 1 ). Such an electronic configuration wherein a set of orbitals is
half filled with one electron in each seems to have special stability.

Pairing in the 2p orbitals does not begin until oxygen (Is 2 2s 2 2p 4 ), in
which one of the 2p orbitals is complete. Two electrons can exist in, or
pair off in, the same orbital only if they have opposite spins, + and — .
Electrons of the same spin will repel each other and cannot exist in the
same orbital since they would have the same values for all four quantum
numbers. Spin orientation and the pairing of electrons can be represented
as shown in Figure 12.5.

A spinning electron acts as a small magnet. Atoms with unpaired elec-
trons have a net magnetic field and interact with an external magnetic
field. They are paramagnetic in that they are pulled from a weak into a
strong magnetic field. The greater the number of unpaired electrons, the
more intense is the magnetic property of an atom so that magnetic measure-
ments are an important tool in the clarification of atomic structure. If all
the electrons of an atom are paired the net electron spin is zero and so also

162

Atomic Structure ■ — 11. Electronic Structure

Nitrogen

2 P

®(D©

Ox> gen

® © 0

Neon

® ® ®

Each circle represents an orbital. An arrow Within a circle indicates occupancy
of the orbital by an electron, and the two directions of the auows, up or down,
indicate the two possibilities of the spin quantum number. Xu the nitrogen atom,
both the Is and 2s orbitals have paired electrons, the p electrons, one m each p
orbital, have parallel spins. In the oxygen atom, one v orbital is complete and
contains a pair of electrons, opposite m spin. In neon, an inert gas, all electrons
are paired.

Figure 12.5. Electron Spin Orientation.

is their magnetic field. Such atoms are diamagnetic and are repelled by an
external magnetic field. To determine whether a substance is diamagnetic
or paramagnetic, it can be suspended from a balance beam so that it is
half inside and half outside the space between the poles of a strong electro-
magnet. When the electromagnet is activated, a diamagnetic substance is
repelled by the field while a paramagnetic substance is pulled into the
field. The force required to restore the balance to its initial point would be
a quantitative measure of the magnetic susceptibility of the substance.

9. Electron Configuration and the Periodic Table. It is now apparent
that the arrangement of the Periodic Table, deduced by Mendelejeff on the
basis of the physical and chemical properties of the elements, is rooted in
their electronic structures. The noble gases in Group 0 are elements whose
s and p orbitals of the outermost energy level are complete, ns 2 np H
except for helium where a single s orbital completes the K shell. The periods
of 2, 8, 18, 18, and 32 elements involve the filling of orbitals from one inert
gas structure to the next.

Similarity in the chemical behavior of different elements is due to a
similarity in their electron structures. The alkali metals of Group I A all
have a single unpaired s electron, in their outermost energy levels. The
halogen elements of Group VII B have seven electrons in their outermost
shells or a deficiency of one electron to complete the p orbitals of those
shells, ns- np\ The chemical properties of an element are determined pri-
marily by the electron configuration of the outermost shell but the structure
of the next-to-the-outermost shell does have a marked influence. The separa-
tion into A and B subdivisions of a group is due to a difference between the
configurations of their next-to-outermost shells. In a B subgroup, the next-
to-the-outermost shell has a total of eighteen electrons, n$r np b nd u \ In an
A subgroup, the next-to-the-outermost shell has unoccupied or incomplete d
orbitals; the total number of electrons varies from 8 to 17. Thus, in Group

Atomic Structure — II: Elect: onic Structure

163

I A, the next-to-the-outermost shell has 8 electrons, while in Group I B, this
shell has completely filled d orbitals, resulting in an additional 10 electrons,
or 18 in all.

In the fourth period, the first transition series of ten elements from
scandium to zinc is due to the building up of the 3d orbitals from 8 to 18
electrons. Similarly the transition series from yttrium to cadmium in the fifth
period results from the filling of the 4 d orbitals though this does not pro-
ceed so regularly as did the buildup of the 3d orbitals. In the fifth period,
one might expect that a third transition series of ten elements would start
with lanthanum, but instead of the 5d orbitals the 4 f orbitals are first oc-
cupied. Inasmuch as these electrons enter a principal energy level, the N
shell, which is two removed from the outermost since some orbitals in the
O and P shells are already occupied, the elements in which the 4 f orbitals are
being filled have remarkably similar chemical properties. Hence they are
quite correctly placed in a single cubicle of the Periodic Table, along with
lanthanum. This is the lanthanide series of elements, sometimes called the
rare earth series. With the conclusion of the lanthanide series at lutecium,
there appears a transition series to mercury during which the 5 d orbitals
fill up. Presumably the series starting with actinium, the actinide series, is
similar to the lanthanide series in that the 5/ orbitals are being filled. This
series is concluded with the element of highest atomic number, lawrencium.

10. Origin of Spectra. Though an electron must exist in a definite energy
level it can be raised or “excited” to a higher energy level by the absorption
of energy. Since the energy levels have fixed values, only a definite amount
of energy can be absorbed by the electron in the process of excitation. This
is equal to the difference in energy between the initial electronic energy
level and the level to which the electron has been excited. Normally the
electron does not remain long in the excited level, 10 8 second or less. It
returns spontaneously to the lower energy level and thereby emits the definite
quantity of energy which it previously absorbed, as radiation of a specific
frequency, that is, a spectral line. The frequency, v , of the spectral line is
related to the energy difference, E, by

(1) E = hv

Some possible electronic transitions between energy levels in the sodium
atom are shown in Figure 12.6. These are shown for the 3s electron only.
The transition from a lower to a higher energy level, for example the ex-
citation 3s -» 4p, involves the absorption of energy. The reverse transition,
4 p — > 3s, results in the emission of energy as a spectral line. The frequency
of this spectral line is equal to difference in energy between the 4p and 3s
levels divided by Plancks constant.

Electrons in complete energy levels, such as the K and L shells in the
sodium atom, are relatively strongly bound in the atom and extremely high
energies are required to eject them or to excite them to higher energy levels.

164

Atomic Structure — U. Electronic Structure

These high energies give rise to X-ray spectra. Spectra in the visible and ultra-
violet regions are produced by the excitation of the outermost electrons
which are the most loosely bound in an atom. For sodium this is the 3 $
electron. Ordinarily this electron is in the “ground” or lowest energy level
it can attain, the 3 s level. But there are available higher energy levels to
which this electron can be excited with moderate energies. In theory a spec-
tral line could result, after absorption of energy, from a transition from
any higher to any lower energy level. Not all transitions are permissible,
however; certain selection rules limit transitions for both ahsorption and
emission to s < — > p\ p d; d < — > /.

Because each atom has its own specific set of energy levels, each atom
emits a characteristic spectrum which can serve to identify or “to finger-
print” the atom in analytical chemistry. Some atomic spectra in the visible
region are shown in Figure 12.7.

11. The Spectroscope. To obtain a spectrum from a source of radiation
an instrument called a spectroscope is used to resolve or to split up this
radiation into its component frequencies. The essentials of a spectroscopic
instrument are a source of radiation, a dispersion medium, and a recording
device. In increasing order of energy, the common sources into which the

* Cornplrtr
1 OrhiUi*

- - J

fhe electronic configuration of sodium is Is 2 2 p* 3s x . Transitions are

shown .for the 3$ electron only. Electrons in the complete K and L shells are more
strongly hound by the nucleus than the &s electron. The spacing between the
energy levels is not to scale but is in agreement with the relative energies of the
orbitals, i.e., the order in which the orbitals are filled first.

Figure 12.6. Energy Level Diagram of the Sodium Atom.

TimlmilnnliilnlnliiliilihnllmilmiliiiiTiiiilin tlinilinillmllmiliiiiliiiiiriinlmiliHilm

Spectra in die visible region for some atoms are shown. The solar spectrum or
white light & a continuous spectrum with frequencies from the red to the violet.
Special lines, invisible to the eye, may also be present in the infrared and uitra>
violet

Figure 12 . 7 . Spectra of Atom*.

Atomic Structure — II: Electronic Structure

165

substance under investigation is inserted are the flame of a Bunsen burner,
the direct current arc, and the high voltage spark. The electrons of the sub-
stance are excited by the energy of the source. The dispersion medium,
which is either a prism or a diffraction grating, resolves the emitted radia-
tion into separate frequencies which are then recorded or observed as in-
dividual lines by the eye, or through the magnitude of a meter reading,
or on a photographic plate. The instrument derives its name from the
method of observing the spectral lines. Observation by the eye is done by
a spectroscope ; recording with a meter is done, by a spectrometer or a
spectrophotometer; and where a photographic record is obtained, the in-
strument is a spectrograph. Basically all are the same. Lenses and a slit
are necessary to collimate and to focus the light beam and the spectral
lines, which are images of the slit. The arrangement of a prism spectro-
graph is shown in Figure 12.8.

A narrow beam of light from the source, X, is selected by the slit, S. It is
collimated by the lens, L„ into parallel rays which enter the prism, P. In passing
through the prism the radiation is dispersed into its Component wavelengths which
are then focussed by lens, L 2 , upon the photographic plate, C. The spectral lines
on the plate are images of the slit. Light of long wavelength is bent least in
passing through the prism.

Figure 12.8. The Prism Spectrograph.

QUESTIONS

1. Discuss the Bohr concept of the atom. How does it differ from the Dalton
model?

2. In what respects is the Bohr theory incorrect?

3. What is meant by the term "energy level”?

4. What were the contributions of de Broglie and of Schrodinger to atomic
structure?

5. List the' quantum numbers and their characteristics.

6. What is the Pauli Exclusion Principle and how does it bear upon atomic
structure?

7. Define "orbital.”

jgg Atomic Structure -—/I: FAectronic Structure

8. What is meant by the pairing of electrons? Do all atoms of odd atomic
number have unpaired electrons? Do atoms of even atomic number have
completely paired electrons?

9. What is the maximum number of electrons in (a) a principal quantum level
(b) an orbital (c) in the s\ p % d> and f orbitals of an atom?

10. What is meant by the symbols 3s“; 4d'; 5 /*?

11. What is the atomic number of the following elements whose electronic
structures are: Is* 2s* 2p‘ ; 3s 1 ; la* 2s* 2 &v* 3p« 3rf« 4 a-?

12. What element dr elements has in its electronic configuration these electrons:
4p»; 4 f ; 3p<; 5f ;

13. Why is the rubidium atom Is® . . . 4 a- 4p* ! 5s* rather than Is* . , . 4s* 4p° 4d*?

14. Which of the following are not possible? Give a reason for your answer.

Ip*; 3s 3 ; 4P; 2p K ; 5 7p J ?

15. Define paramagnetism and diamagnetism. What is the cause of each?

16. Which of the following atoms will be paramagnetic? B, V, Zn, Fb, La?

17. What is meant by the spatial orientation of orbitals? Give an example,

18. From the viewpoint of atomic structure, what is (a) a transition series (b) the
lanthanide series?

19. What is the origin of (a) visible spectra (b) X-ray spectra (c) gamma rays?

20. Design a spectroscope. Can a spectroscope be used for the ultraviolet?
Explain.

21. What part{s) of the aljom largely account for (a) mass of the atom (b) density
of the atom (c) similarity of chemical properties in a group of the Periodic
Table?

22. From the viewpoint of electronic configuration distinguish between the A and
B subgroups of the Periodic Table.

23. How is the fact that the fourth period of the Periodic Table has 18 elements
related to electron configurations?

24. Calculate the deBroglie wave length of (a) an electron moving with a velocity
of 1 X 10* cm/sec (b) of a 10 gram bullet moving at the same speed.

25. The difference in energy between two electronic energy levels is 10.0 electron
volts. Calculate the frequency and wavelength of a spectral line which would
be emitted by an electronic transition between these levels. In what region of
the electromagnetic spectrum is this line?

26. How might techniques and the meaning of measurement be affected in our
macroscopic world if Planck's constant were one, and not 6.5 x 1G~ 27 erg see?

APPENDIX

Bohr Derivation for the hydrogen atom. The angular momentum of an electron
of mass m moving in a circular orbit of radius r with a velocity v is equal to mvr.

Bohr assumed that this angular momentum was quantized in units of r — . Only

2 IT

those orbits were possible whose angular momentum was an integral multiple,
n * °* m > & is the principal quantum number.

Atomic Structure — II: Electronic Structure

167

( 1 )

mvr = n

_h_
27 r

The electron is acted upon by two opposing forces in its orbit, the centrifugal

mv-s

force, , and the Coulombic force of attraction between the electron and the

positively charged nucleus. For a nucleus of charge, Ze, the attractive force is

Ze • e Ze 2

“7“

given by Coulomb's Law as
Equating the two forces,

(2)

and solving for r.

mv z

Ze 2

r 2

(3)

r =

Ze 2

mv 2

(4)

Substituting the value of v obtained from equation (1) into equation (3),

n 2 h 2

1 4 V s m e 2 Z

For the hydrogen atom, Z = 1 and the smallest orbit is that for which n = 1.
Substituting these and the values of h, 7 r, m, and e in Equation 4 gives a value for
r of 0.529 X 10“ 8 cm.

The energy levels for the hydrogen atom can also be calculated. The total

Ze 2

energy, E, is the sum of the kinetic energy, Vi mv 2 , and the potential energy,

Note: The potential energy is the integral of the attractive force

jst * = _ »

J r 2 r

Ekin 4" Epot

mv 2

Ze 2 Ze 2

r 2r

Ze 2

(5)

Ze 2
2 r

Substituting the value of r from Equation (4)

E =

27T 2 m e 4 Z 2
h 2 n 2

For a given atom, only the quantum number, n, in Equation 6 is variable.
Since n can take on only integral values, the energy, E, is also quantized.

Radiation of a definite frequency is absorbed or emitted by the transition of
an electron from one energy level, n l9 to another, n 2 . Collecting all the terms

168

Atomic Structur^-ll: Electronic Structure

which are constant (for a given atom), Equation 6 can be written for the energy
difference between two such levels.

( 7 )

-n t

E„, —

27r*m e 4 Z 2
h 2

"JL

ns 1

Since E n2 - E 0l = hy and v = c A, the frequency (or wavelength) of the
corresponding spectral line can be calculated.

The Chemical Bond

The attractive forces which hold together the atoms within a molecule
constitute the chemical bond. Our model so far developed for atomic structure,
indicates that the chemical bond must be electrical in character and somehow
related to the electron structures of atoms. We shall see that chemical
bonds are formed through the transfer or the sharing of electrons from the
outer shells and orbitals of atoms. In chemical reactions bonds are broken
and reformed in order to attain more stable electronic configurations.

1. The Octet Rule. The stability of the electron configuration of an atom
can be measured in terms of the energy, defined as the first ionization po-
tential, which is required to remove completely an electron frpm a neutral
atom. The greater the value of an ionization potential, the lower is the
value of energy which an electron configuration itself possesses, and thus
the more stable it is. The energy required to remove an electron from an
already singly ionized atom is called the second ionization potential, and so
on. In Figure 13.1, the first ionization potentials of the atoms are plotted
against atomic number.

Note that the ionization potentials of the elements are periodic functions of the
atomic numbers of the elements.

Figure 13 J. Ionization Potential* of the Element*.

170

The Chemical Bond

The ionization potentials of the noble gases constitute the peaks of
the graph and so are greater than those of the elements that immediately
precede or follow them in the Periodic Table. Thus the electron configura-
tion of the neon atom is more stable than that of the fluorine atom or the
sodium atom, and argon is more stable than chlorine or potassium. Chemical
evidence of the unusual stability of the noble gases is the almost complete
unreactivity of these elements. Only the larger noble gas atoms react and
even these form stable binary compounds only with fluorine, e.g., XeF 4 .
Again since all systems, electron configurations included, tend to attain
maximum stability, that is, the lowest possible energy level available to
them, it is reasonable to assume as a starting point that atoms react insofar
as possible to acquire a noble gas electron configuration, an outermost shell
with an octet of electrons, ns 2 np ft .

2. The Ionic Bond. Let us consider the reaction between a sodium atom
and a chlorine atom. The electron configuration of sodium is Is 2 2$r 2p 6 3s 1 .
By losing the single 3s electron, the sodium atom would expose as its outer-
most shell a complete L shell containing eight electrons, 2s 2 and thereby
have a neon-like configuration. On paper it is also possible for a sodium
atom to gain seven electrons in its M shell and to form thereby an octet,
but this never happens. Energy calculations for this process or merely logical
consideration of the two alternatives makes the latter extremely improbable.
Similarly, a chlorine atom (Is 2 2s 2 2p $ 3 s 2 3 p®) could gain one electron
and thereby have a total of eight electrons in its outermost or M shell,
3 s 2 3p\ an argon-like configuration. The transfer of the 3s electron from a
sodium atom to a chlorine atom would result in the formation of an octet
for each atom. This electron transfer constitutes the chemical reaction be-
tween the two atoms. Using an orbitlike viewpoint for the electron con-
figuration of an atom, the reaction between a Na 2S atom and a Cl*® atom can
be represented as follows:

/ / /

■7 8 2

V V \

K L M
Na M atom

M K L

Cl** atom

In this transfer of an electron, the atomic nuclei and the completed electron shells,
insofar as the number of electrons is concerned, remain unchanged. This is char-
acteristic of chemical reaction.

Figure 13.2. The Transfer of the 3s Electron from a Na Atom to a Cl Atom.

What is the result of the transfer of an electron from the sodium atom
to the chlorine atom? The chlorine nucleus, unchanged, is charged +19, but
outside the nucleus there are now a total of 20 electrons, resulting in a
net electrical charge of -1 on the chlorine “atom” This entity is called the
chloride (not chlorine!) ion, an ion being an atom or a group of atoms
which is electrically charged. Similarly die sodium ion, produced when

The Chemical Bond

171

the sodium atom loses an electron, is charged +1. An ion is indicated by
writing its chemical nature and the magnitude of its charge at the upper
right of the chemical symbol, thus, the calcium ion, Ca 2 + or Ca++ and the
sulfate ion, S0 4 2 - or S0 4 **~.

The reaction between the sodium atom and the chlorine atom can be
represented more simply as

• •

(1) Na o + . Cl J -» Na+ + Cl-

• •

where the symbols Na o and . Cl : , sometimes called Lewis symbols for

the American chemist, G. N. Lewis, stand for the kernels of the respective
atoms, that is, all but the external electron shells. The symbols o and
represent electrons in the outermost shells of the Na and Cl atoms respec-
tively, and Na+ and Cl** the ions formed by the reaction. No inference is to
be drawn, here or later, from the use of different symbols tor electrons;
an electron is an electron, independent of its source. The difference in
symbols is useful merely to indicate the different origins of the electrons.
An alternate method of representing the electron transfer is

(2) Na - e Na+ or Na Na+ + e

Cl + # -> Cl- Cl ^ Cl~ - e

where e represents the electron.

Whereas elemental chlorine exists in nature only as diatomic molecules, Cl 2 ,
and not as individual atoms, each chlorine atom of the molecule would react
with a sodium atom in the manner indicated. For simplicity in illustrating the
mechanism of the reaction only one chlorine atom was shown. The balanced equa-
tion for the reaction as it occurs in nature would be 2 Na -f Cl 2 — > 2 Na + + 2 Cb

The opposite charges on the sodium and chloride ions result in an
electrostatic attraction, or bond, between them. This type of bond is called
the ionic bond and the resultant compound is said to be ionic. Though
generally written in the molecular form as NaCl, sodium chloride, even in
the solid state, is actually composed of Na + and Ch ions, and more properly
should be written as Na + Cl“. In summary, an ionic bond is formed when
one or more electrons is transferred completely from one atom to another,
thereby producing charged particles, or ions, which hold each other^ by
electrostatic attraction. Ionic bonds are formed mpst readily by those elements
whose neutral electron configurations differ from an octet by one, two,
or at most, three electrons. The tendency to do so is also roughly in that
order. Metallic elements at the left of die Periodic Table have one, two,
or three electrons in excess of an octet and so lose electrons to form positive
ions whereas nonmetallic elements on the right lack but a few electrons
'towards an octet and so gain electrons to form positive ions. Inasmuch as
it is the very same electrons which are lost by one atom that are gained
by the other, reacting atom, it is apparent that, in chemical reactions in-
volving the formation of ionic bonds, the number of electrons lost is equal

172

The Chemical Bond

to the number gained. The numbers of atoms that react must adjust them-
selves to this requirement. Thus sodium combines with sulfur to form the
ionic compound, sodium sulfide, Na^S. The reaction is 2 Na + S — ► Na 2 S
and the electron mechanism is

Na*

(3) S : -> 2 Na + + S 2 ~

Na® **

The loss of electrons is defined as oxidation ; it results in a more posi-
tively charged (algebraically) entity. The gain of electrons is called reduction;
it results in a more negatively charged entity. The charge on the ion is called
its electrovalence and the ionic bond is thereby also known as the electro -
valent bond . For sodium, the electrovalence is +1, for chlorine -1, and
for sulfur -2. Electrovalence, which is specific to ions, is one kind of valence
and is distinguished by placing the sign of the ionic charge before the valence
number. In many cases the absolute value of the electrovalence equals
numerically the valence as given in Table 6-A. Some atoms can have more
than one value of electrovalence. Under different conditions certain atoms
can lose or gain different numbers of electrons. An atom of iron (Is 2 2s 2
2 p* 3s 2 3 p 6 3 dE® 4s 2 ) can lose either two electrons, the 4s electrons and
form an Fe 2 + ion, or three electrons, both 4s electrons and one 3d electron*
In the latter case, the Fe s + ion would be formed, with five 3 d electrons,
one in each of the five d orbitals.

3. The Covalent Bond. The ionic bond results in charged particles. The
marked differences in properties (Section 14) between certain substances
indicates that some are not ionic and that a different kind of chemical bond
must exist. In addition to the transfer of electrons an electron octet can be
formed by the sharing of electrons. Let us consider the combination of
two chlorine atoms to form a chlorine molecule, Cl 2 . In the chlorine mole-
cule there are no ions. With two identical atoms there is obviously no
greater tendency on the part of one to gain or the other to lose electrons.
In the chlorine molecule, one electron from the outermost shell of each
chlorine atom is shared between them. This results in an octet of electrons,
and a noble gas configuration in the outermost shell of each, atom since
the shared electrons are counted in each atom's total Thus,

(4) SCI* + *ci; -» S Cl S Cl s (or Cl*)

** >« * • •«

Such a bond wherein electrons are shared by the members of the bond is
called a covalent bond, A single covalent bond consists of a pair of electrons,
opposite in spin. The electrons of a covalent bond are in the same orbital
and hence must have opposite spins to produce a stable bond. If the electrons
had parallel spins they would repel each other and no bond would be
formed. By far the great majority of chemical compounds have covalent
bonds.

In covalent bonding the valence of an atom, called its covalence, is
equal to the total number of covalent bonds or electron pairs linking it to
other atoms in the molecule. In the Cl* molecule, where a single pair of

The Chemical Bond

173

electrons is shared, the chlorine atoms are said to have a covalence of
one. A common symbol for the covalent bond is a dash, thus — , and the

• • 0 9

Cl 2 molecule might then be represented as : Cl — Cl % . Such a formula in

• • 09

which the bonds are represented by lines between the bonded atoms is called
a structural formula or graphic formula. For simplicity the nonbonding elec-
trons may be omitted, and the chlorine molecule written simply as Cl - Cl.
More than one pair of electrons may be shared between two atoms. The
sharing of four electrons would be a double covalent bond, symbol — , and
of six electrons a triple covalent bond, symbol = Such bonds are especially
reactive. Most carbon compounds containing double or triple bonds be-
tween carbon atoms are unsaturated in that they can react to add atoms to the
carbon atoms joined by such bonds, thereby forming a saturated compound
having only single covalent bonds between the bonded atoms.

Typical examples of covalent bonding are:

Molecular

Electronic

Structural

Substance

Formula

Hydrogen

h 2

h:h

H - H

Water

h 2 o

h:o:h

o o

H — O — H

:ci:

Carbon Tetrachloride

CC1 4

:ci:c:cis

Cl - C - Cl

•• #fl ••

:ci:

• •

Ammonia

nh 3

h:n:h

O 0

H — N - H

Carbon Dioxide

co 2

’b::cj:o*

Acetylene

c 2 h 2

h:c»c:h

III

In the examples given, the carbon atom has a covalence of four; nitrogen
three; oxygen two; and hydrogen one. Unlike electrovalence, covalence is
never given a + or sign; under ordinary conditions the covalent bond does
not disrupt to yield Ions. Within an ion composed of more than one atom
the atoms can be bound by covalent bonds. In the nitrate ion, NOr, the
nitrogen and oxygen atoms are joined by covalent bonds though, the entire
ion can be linked by an ionic bond to another positive ion, such as Na +
4. The Coordinate Covalent Bond* It is not necessary that the covalent
bond be formed from an electron contributed by each member of the bond.
One member may contribute both electrons. Inspection of the NH 3 structure
reveals that the nitrogen atom has two unshared electrons not directly involved
in bonding. These can be used to bond covalently the NH 3 molecule to

174

The Chemical Bond

some other atom, ion, or molecule which might require two electrons. The
hydrogen ion, H + , which is merely a proton, can thus be bonded to a NH*
molecule to form the ammonium ion, NH* + . The positive charge of the
proton is then considered to be distributed over the entire NH 4 + ion since
all the hydrogen atoms are chemically equivalent,

(5) H:N:H + H +

Such a covalent bond, where formation is due to the contribution of both
bonding electrons by one member of the bond, is sometimes called a coordinate
covalent bond. Only in its method of formation does a coordinate covalent
bond differ from an ordinary covalent bond. Once formed both types of
covalent bonds are alike in properties. The atom or ion which supplies
the pair of electrons is called the donor or ligand. The atom receiving
the electron pair is the acceptor ; Not every pair of electrons is available
for coordinate covalent bonding. Coordinate covalent bonds are most fre-
quently formed by those atoms whose normal covalence is already filled.

Other examples of coordinate covalent bonds are the sulfate ion, S0 4 2 ~
and the blue, tetrahydrated copper ion, Cu(H 2 0) 4 :i ' l “ (Figure 13.3).

h :n: h
» «

o o

:o”

: 6 : s i o :

00 •» oo

%Ot

The sulfate ion has a charge of 2 Two elec-
trons, shown as £ have been gained from an-
other atom(s) to "which it is linked by an ionic
bond.

In the Cu(H 2 0) 4 2+ ion, four molecules of
water are symmetrically linked to the cupric
ion, Cu 2 +, by coordinate covalent bonds. The
oxygen atoms are the donor atoms and the
Cu 2 + ion is the acceptor.

Figure 23.3. Coordinate Covalent Bonding.

The octet rule is not universally applicable. For some covalent compounds
octet formation is not possible. In boron trifluoride, BF S , the fluorine atoms
obey the octet rule but the boron atom is “electron deficient” in that it has
only six electrons in its external shell.

:F:

:F :b: f:

F — Tl _ F

The Chemical Bond

175

However BF S can form an “addition compound” with NH 3 . The compound
is BF 3 — NH :) , and in it two electrons from the nitrogen atom produce an
octet around the boron atom by coordinate covalent binding.

:f: h

• • • o X •

:f :b Inih or

• • • o 0 X

:F: H

• •

F H

I I

f-b-n-h

Since the K shell can accommodate a maximum of eight electrons, 2 s 2 2p°,
elements in the second period of the Periodic Table can form no more
than four covalent bonds. Thus the compound nitrogen trichloride, NCI,,
exists but nitrogen pentachloride, NC1 5 , a molecule which would require
five covalent bonds between the nitrogen atom and the chlorine atoms, has
not been prepared. However the element phosphorus, in the same family
as nitrogen but in the third period, forms the compound, phosphorus
pentachloride, PCl r >. In this case the octet can be expanded to ten
electrons because 3 d orbitals, in addition to 3s and 3p orbitals, are
available for bonding. Where the difference in energy between or-
bitals is not too great, electrons may be “promoted” to the higher
level during chemical reaction. Promotion of one 3 p electron to a 3 d level
in phosphorus enables the formation of two additional covalent bonds, or
five in all. For the nitrogen atom, such promotion is energetically im-
probable since the 3d level is much higher in energy than is. the 2 p level.
Other molecules analogous to PCI* are sulfur hexafluoride, SF 6 , and iodine
heptafluoride, IF 7 . Particularly with elements of high atomic number where
a variety of orbitals may be used in bonding, the octet rule is not followed
strictly.

5. Oxidation Number. It is possible to assign a number to each atom
in a molecular formula on the basis that the molecule as a whole is electrically
neutral and that each atom contributes an algebraic portion to that neutrality.
Such a number is known as the oxidation number or oxidation state. In its
most common compounds the hydrogen atom is arbitrarily assigned an oxi-
dation number of +1. 1 From this starting point, the oxidation numbers of
other atoms can be determined. The formula, HC1, indicates that the oxidation
number of the chlorine atom must be -1 for the HC1 molecule to have a
net charge of zero, and if each hydrogen atom is +1 in H 2 0, then the oxygen
atom must have an oxidation number of -2. In nitric acid, HNO a , if H is
+1, and each O is -2, then N must be +5. This is evident from the algebraic
equation: 4-1 -f (N) + [3 X (-2)] = 0. The oxidation number of an
atom in the elemental state is zero. Within a compound, the oxidation number
of an atom is a positive or a negative integer but never zero, A given atom
may evidence different oxidation numbers in different compounds. In addi-

176

The Chemical Bond

tion to +5, the nitrogen atom has oxidation numbers of +4 (NO s ), +3 (NaCh),
+2 (NO), +1 (Ns»0), zero (N 2 ), and -3 (NH 3 ). Such a range is unusual;
most atoms show but one or two oxidation numbers.

Though superficially alike, the oxidation number of an atom is not to
be confused with its valence. The latter is a pure number denoting a com-
bining capacity, whereas the oxidation number is an algebraic value. In
many cases oxidation numbers bear some relation to the electrovalence of
the ion concerned but frequently they are fictitious in that the atoms or the
ions to which they refer do not actually exist. In nitric acid there are no
N 6+ ions; the stable entity is the nitrate ion, (N0 3 )~.

6. Partial Ionic Character of Covalent Bonds. So far we have discussed
the chemical bond in terms of two extreme cases, the ionic bond and the
non-ionic covalent bond. However chemical bonds can partake somewhat
of the characteristics of each type of bonding and so have an intermediate
character. Just as atoms have different tendencies to lose electrons, so too
they vary in their attractions for electrons. Thus the 'electrons in a covalent
bond may not be equally shared by the members of the bond. An atom
having a large affinity for electrons will tend to attract to itself the electrons
of a covalent bond Though this attraction is insufficient to produce charged
ions, the electrons will be closer to the atom with the greater electron at-
traction and such an atom will be relatively more negative than the other
member of the bond. Thus a covalent bond can have a partial ionic character.

Electron affinity is measured in terms of the energy which is released
when a neutral atom gains an electron. This is difficult to determine experi-
mentally. Related to the electron affinity, but more qualitatively useful, is
the electronegativity concept based on bond energies and proposed by
the American chemist, Linus Pauling. Based upon an arbitrary scale, the ele-
ments have been assigned electronegativity values which are measures of
the attractions of their atoms for electrons in a covalent bond. Electronega-
tivities are listed in Table 13-A both in order of magnitude and in their rela-
tion to the Periodic Table. The nonmetallic elements which generally have a
tendency to gain electrons have high values of electronegativity, and con-
versely the metals have low electronegativities. Fluorine, the most electro-
negative element, has a value of 4.0 whereas cesium, with little tendency to
gain electrons and indeed a great tendency to lose electrons, has a value of
0.7. The metals in general have values less than 1.7.

A covalent bond formed between atoms of the same electronegativity
would result in the bonding electrons being equally shared between them
an<f the resultant bond would be purely covalent. Such is the case with the
diatomic molecules, H 2> 0 2 , N 2 , and Cl*. The greater the difference in
electronegativities between two elements forming a covalent bond the greater
is the ionic character of the bond. From the values given in Table 13-A
we could predict that carbon disulfide, CS*, would have very little ionic
character whereas hydrogen chloride, HC1, should have appreciable ionic
character. Indeed the electronegativity difference may be so great as to
render a bond wholly ionic. As a general rule, a difference of 1,7 corresponds
to approximately 50% ionic character, as shown.

The Chemical Bond

177

Difference in Electronegativity;

0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8 2.0 2.2 2.4 2.6 2.8 3.0
Percent Ionic Character:

1 4 9 15 22 30 39 47 55 63 70 76 82 86 89

The strength of a covalent bond is also related to the electronegativity dif-
ference; the larger this difference the greater is the bond strength. Thus
ammonia, NH 3 , is a more stable compound, that is, less easily decomposed

Table 13- A

Electronegativity

Values

2.1

4.0

1.8

3.5

1.8

' N

3.0

1.8

1.0

1.5

2.0

2.5

3.0

3.5

4.0

3.0

1.7

2.8

1.5

0.9

1.2

1.5

1.8

2.1

2.5

3.0

2.5

1.5

2.5

1.2

2.5

1.0

0.8

1.0

1.8

2.0

2.4

2.8

2.4

1.0

2.1

1.0

0.8

1.0

1.8

1.9

2.1

2.5

,2.1

0.9

2.1

0.9

2.0

0.8

0.7

0.9

1.8

1.9

2.0

0.8

1.9

0.7

No values of electronegativity have been assigned

1.8

to the Noble Gases

into its elements, than is phosphine, PH 3 . Whereas electronegativity is a use-
ful concept it is not the only factor which governs the character of a covalent
bond so that conclusions drawn solely from it must be treated with caution.

7. Polarity of Covalent Molecules. Where the same atoms combine to
form a covalent bond, such as in the Cl 2 molecule, neither atom has a greater
tendency to attract electrons and the electron pair of the covalent bond will
be “midway” between the chlorine atoms. In an HC1 molecule the electron
pair will be closer to the Cl atom than to the H atom since the former
has a greater electronegativity. There are no ions in the molecule but due
to the unequal distribution of the electron pair, the Cl atom, is negative rela-
tive to the H atom. The Cl atom is said to have a partial negative charge
and the H atom a partial positive charge. Viewed as a separation of two
centers of electrical charge, the HC1 molecule is said to be polar or a dipole
by analogy to a magnet with its two poles some distance apart. The Cl 2
molecule is thereby nonpolar (Figure 13.4).

With more complex molecules the determination of polarity is not so
simple. Just as it is possible to determine the center of gravity of a physical
object or the geometric center of population for the United States, however,
given sufficient data a point in the molecule could be calculated at which

178

The Chemical Bond

\ X 00

v Cl 1 Cl 0

X \ 00

H :ci :

« 0

8+ 8-

(± 3

Cl 2

HCl

Nonpolar; covalent elec-
tron pair equally shared
between the Cl atoms.

Polar; covalent electron pair

Schematic representation

closer to the Cl atom. The H

of the HC1 molecule; a

atom is positive relative to the
CL atom. The symbol 5 repre-
sents the partial charge.

dipole.

Figure 13.4. Polar and Nonpolar Molecules.

we might consider the net contribution of all the electrons in a molecule
to be localized, namely, the center of negative charge. This could also
be done for the positive charges, or atomic nuclei, of a molecule. Those
molecules in which the centers of negativity and positivity coincide are non-
polar. Where they do not coincide the molecule is polar and the greater
the separation of the negative and the positive centers, the greater is the
polarity of the molecule. Another example of a nonpolar molecule is carbon
tetrachloride, CCl*. Even though the electrons in any single C — Cl bond
are closer to the G1 atom than to the C atom, the tetrahedral symmetry
of the CC1 4 molecule makes for a net polarity of zero for the entire mole-
cule. Because of its asymmetry the trichloromethane molecule, CHC1 S , is
polar (Figure 13,5).

Polarity is measured in terms of the molecular dipole moment , fi. This is
the product of the magnitude of the charge, e, of the molecular dipole and the
distance, d, which separates them: fi = e X d. In an electric field the + and -
ends of a polar molecule tend to orient themselves with the field, as shown

X X

CIS

Cl*, c

X •

X X

-+■ XX , « X X

—

CIS

C 1 - + C + -C 1 sen c :ci s

Cl - 4- c + - Cl

1 XX XX

T x.

sen

—

Cl “

CC1 4 molecule
nonpolar

CHCI S molecule
polar

Figure 13.5. Polarity and Molecular Symmetry.

+ and - signs indicate the relative charges of the bonded atoms
* T 5 electrons; x = chlorine electrons; * = hydrogen electron

Carbon is less electronegative than chlorine but more electronegative than hydro-
gen. No ions are present in either CCl 4 or CHC1 S , but CHC1 3 is polar.

The Chemical Bond

179

in Figure 13.6a. The dipole moment can be measured by determining the
ratio of the capacities of an electrical capacitor with air as the medium be-
tween the capacitor plates, that is, with air as the dielectric, and with the
substance whose dipole moment is to be determined as the dielectric med-
ium. The dipole moment is a useful tool in determining the geometry of .mole-
cular structure. From the magnitude of the dipole moment one can decide
whether a molecule has a symmetric or an asymmetric structure. A molecule
with zero dipole moment is nonpolar and symmetric. The water molecule, H 2 0,
could possibly be a linear symmetric molecule or an asymmetric one but
since water has an appreciable dipole moment the triangular structure in
Figure 13.6b is preferred. On the other hand, despite the difference in
electronegativities between a carbon and an oxygen atom, the carbon dioxide
molecule, CQ 2 , has zero dipole moment and hence a linear symmetric structure.

CUD
GEZ3
C E3

(±_

c±z

4-*

< (±Z

“b

Figure 13.6 a. Orientation of a
Dipol,© in an Electric Field.

A and B are electrically charged
metallic plates in a polar liquid. The
pokr molecules, indicated as dipoles,
align themselves in the direction of
the field.

H + -O- + H

+ H

A linear symmetric structure would have An asymmetric structure,

no dipole moment.

Figure 13.6 b. The Water Molecule.

8. The Hydrogen Bond. The normal valence of a hydrogen atom is
one but in some cases it can act as a bond between two atoms. For example,
in the polar water molecule, the oxygen atom is so electronegative that it
attracts to itself the covalent electron pair, leaving the hydrogen atom as
essentially a proton. This relatively positive hydrogen atom will be attracted
to the oxygen atom of a neighboring water molecule and thereby act as a
bridge to bind two water molecules together. Such a bond is known as a
hydrogen bond or a protonic bridge (Figure 13.7).

H H The -j- and - signs indicate relative charge.

O o H “ O The hydrogen bond is the H atom between the two O

H — . -f- — atoms.

Figure 13.7. The Hydrogen Bond.

180

The Chemical Bond

The unusually high boiling point of a compound with so low a molecular
weight as H 2 6 (100°C) compared with the boiling points of analogous hydro-
gen compounds of the Group VIB elements, H 2 S (-62 C C), H 2 Se (-42°C), and
H 2 Te (-4°C), is believed due to the association of water molecules by
hydrogen bonding. To vaporize water the hydrogen bonds must be broken;
the energy required to do this results in a high boiling point. Only the most
electronegative elements, F, O, and N, can be linked by hydrogen bonds.
The HF 2 “ ion, in H 2 F 2 and in KHF 2> is held together by a hydrogen bond
between the two fluorine atoms, [F — H — F]~.

9, Atomic Size, A useful generalization can be drawn between the size
of an atom and the magnitude of its ionization potential and electron affinity.
Coulomb's Law of electrostatic attraction states that the force of attraction,
F, between two charges of magnitude q 2 and q 2 is proportional to the product
of the charges and inversely proportional to the dielectric constant, €, of
die medium between the charges and the square of the distance, d, between
them.

it — q*

F “ T3F

Hence distance has a more pronounced effect upon the attractive force be-
tween charged particles than does the magnitude of the charge. Doubling
either charge doubles the force of attraction but doubling the distance de-
creases the force fourfold.

Coulomb's Law can be applied to the attractive force which exists be-
tween the nucleus and the valence electrons of an atom. Within a given
atom, electrons closest to the nucleus will be more strongly held and hence
more difficult to remove For example, in the sodium atom, an atom with
complete K and L shells, more energy is required to eject an electron
from the K shell than from the L shell. Electrons in the K shell are closer
to the nucleus than those in the L shell and so are bound more strongly.
Within a given principal quantum shell electrons are most easily removed
from the orbitals of highest energy, that is, in inverse order to that in
which they filled the orbitals. In the L shell which contains both $ and p
electrons, the p electrons are more easily removed.

A second factor in the ease of removing electrons is that electrons in an
inner shell screen outer electrons from the full magnitude of the nuclear
charge. Thus electrons in the K shell of the sodium atom are attracted by
the full +11 charge of the sodium nucleus but the effective nuclear charge
for attraction of electrons in the L shell is less than +11 because of the
screening effect of the K electrons. The effective nuclear charge for the
sodium atom has been estimated to be +2,2. For both these reasons, the
distance factor and the screening effect, the ionization potential of an electron
in an inner shell is greater than that of one in an outer shell On the other
hand, if it were possible, an electron would be more readily attracted into
an inner shell vacancy than into one in an outer shell.

Figure 13,1 showed that the ionization potential was a periodic function.
A line connecting the low points of that graph would give the variation in
ionization potential for the alkali metals, Li, Na, K, Rb, and Gs. These

The Chemical Bond

181

atoms have a single s electron in their outermost shells, ns 1 . Below are
given the values of the first and second ionization potentials and the radii
of the alkali metal atoms.

Element

Atomic Radius, A

1.55

1.90

2.35

2.48

2.67

1st Ionization Potential, eV

5.39

5.14

4.34

■ 4.18

3.89

2nd Ionization Potential, eV

75.3

47.1

31.7

27.4

23.4

The values of the ionization potentials decrease with an increase in atomic
size. Potassium has a lower first ionization potential than sodium mainly
because its 4s electron is farther removed from the nucleus than is the
corresponding 3s electron of sodium. For a given atom the value of its
second ionization potential is always greater than that of its first ionization
potential. The second electron is removed, not from a neutral atom as was
the first electron, but irom an ion which is already charged +1 and hence
has a greater tendency to hold on to its remaining electrons. In the case
of the alkali metals there is a large difference between the values of the
first and the second ionization potentials because the second electron must
be removed from a complete octet. Complete orbitals and shells show unusual
stability and large amounts of energy are required to disrupt them.

For the nonmetallic halogen group of elements, F, Cl, Br, and I, where
the tendency in chemical reaction is to gain electrons, there is a similar
variation in ionization potential with atomic size, as shown below.

Element

Atomic Radius, A

0.72

0.99

1.14

1.33

1st Ionization Potential, eV

17.42

13.01

11.84

10.44

2nd Ionization Potential, eV

34.8

23.7

19.1

19.4

Electron Affinity, eV

3.63

3.78

3.54

3.24

Electronegativity

4.0

3.0

2.8

2.5

Here, too, the magnitude of the first ionization potential indicates that an
electron can be removed most readily from the largest atom, I, and least
readily from the smallest, F. One the other hand we might expect that an
electron would be most readily gained by the smallest atom, F, and least
readily by I. Except for chlorines' peak value of electron affinity this is
generally true. Even so fluorine does have the highest value of electro-
negativity.

Within any horizontal period of the Periodic Table, the trend of the ion-
ization potential is to increase as we go from left to right. Where electrons
enter the same principal shell, the simultaneous increase in nuclear charge
and the resultant greater nuclear attraction for electrons tend to make the
atoms progressively smaller till a noble gas configuration is attained. In
Figure 13.1 the peaks in the jagged lines in passing from Li to Ne, and
from Na to Ar, show that the relationship between the ionization potential
and atomic size is followed generally but not too strictly.

When a neutral atom loses an electron, the size of the resultant ion is
less than that of the parent atom. With the alkali metal atoms an entire

182

The Chemical Bond

quantum shell disappears. In any case, where more than one electron is
present in the shell, the decreased repulsion between a fewer number of
electrons in the shell results in a contraction and a smaller size for the ion*
Conversely, the gaining of an electron by an atom produces a negatively
charged ion whose size is greater than that of the initial atom. This is due
to an increased repulsion between a greater number of electrons, particularly
since the additional electron goes not into a new shell, but into one already
well populated with electrons. These points are illustrated by the following
data.

Element

Atomic Radius, A

1.55

1.90

2.35

2.48

2.67

0.72

0.99

1.14

1.33

Ionic Radius, A

0.60

0.95

1.33

1.48

1.69

1.36

1.81

1.95

2.16

Though the computation for the energy of a chemical reaction involves
more than just the loss or gain of electrons the following general concept
is useful qualitatively. In chemical reactions, metals tend to lose electrons
and nonmetals tend to gain electrons. This tendency is a measure of the
activity of a substance. The larger the size of a metallic atom the greater
wilhbe its tendency to lose electrons and the more active it will be. Con-
versely, the smaller a nonmetallic atom the greater will be its chemical
activity. These rules hold well for a family of elements or a vertical group
in the periodic table, but they must be used with caution in comparing
elements in different groups.

10. Reaction Mechanism for Bond Formation; the Bom-Habcr Cycle.
Equation 1 for the formation of sodium chloride is an oversimplification
of the reaction mechanism between sodium metal and chlorine gas. Actually
it represents the reaction between isolated atoms of sodium vapor and
chlorine gas to form separate ion-pairs, Na + and Cl\ At ordinary tempera-
tures, sodium is a solid consisting of closely bound atoms, chlorine is a gas
composed of diatomic molecules, and sodium chloride is a three dimensional
lattice of Na+ and Cl~ ions.

The reaction between sodium metal and chlorine gas at ordinary tempera-
tures may be visualized as consisting of the following series of steps:

(1) the sublimation from solid sodium of atoms into the gaseous
state. This requires an input of energy, the sublimation energy, equal to
26 fccal/g-atom.

Na(#X Na(g) AE = +26 kcal

(2) the dissociation of chlorine molecules, Cl 2 , into chlorine atoms, Cl
The energy of dissociation is also endothermic and is 58 kcal per mole of €1*.

Vi Cl.(g) Cl(g) AE = +29 kcal

(3) the ionization of a sodium atom to form a sodium ion, Na 4 * This
requires an energy equal to the ionization potential of sodium, 5.14 eV or
118 kcal/g-atom.

Na(g) Na + {g) +* e

AE s= +118 kcal

The Chemical Bond

(4) the formation of a chloride ion, Cl", when the chlorine atom gains
the electron lost by the sodium atom. This process emits energy equal to
the electron affinity of chlorine, 3.78 eV or 87 kcal/g-atom.

Cl(«) + e -> Cl -(g) AE = -87 kcal

(5) the combination of the ions, Na + and Ch, to form gaseous ion-pairs,
Na^Cl", followed by the union of the separate ion-pairs to form an ionic
crystal lattice. In this process energy known as the lattice energy is evolved.
The lattice energy is defined as the energy released when gaseous ions,
initially at infinite distance from each other, come together to form an ionic
solid. The lattice energy can be calculated from Coulomb's Law; for sodium
chloride it is -185 kcal/mole.

Na + (g) + Cl-(g) Na+Cl-te) I
Na+Cl-te) Na+Ch(j?) )

AE = -185 kcal

The net energy for the over-all reaction is the algebraic sum of the
energies involved in the separate steps (Hess' Law) or

(4-26 + 29 4* - 87 - 185) = -99 kcal.

Thus

Na(*) + Cl,(g) Na+Cl~(s) AE = -99 kcal

This procedure, based on the Law of Conservation of Energy, whereby
a thermodynamic equation is broken down into a series of elementary steps,
is known as the Born-Haber Cycle.

In the event the temperature of the reaction were above the boiling
point of sodium metal, the sublimation step would be omitted. If the reaction
were, carried out above the boiling point of sodium chloride the formation
of the crystal lattice from the gaseous ion-pairs would not be applicable;
only that portion of the lattice energy due to the pairing of the gaseous ions
would be emitted. If a reaction is carried out in solution, steps for the
solvation of the ions, that is, combination with molecules of the solvent,
would have to be added.

11. Resonance. Up to now we have been writing the formula of a covalent
substance in terms of a single molecular structure. Thus for C0 2 we
wrote 0=0=0. However two additional structures, 0=C— O and
O — C = 0, have comparable energies and are plausible. Which one, then,
is the correct structure or is there a correct one? 2 We have no way of know-
ing. If the equations of quantum mechanics are solved on the basis that
each of the three structures -makes a partial contribution, not necessarily
equal, to the character of carbon dioxide, however, we get better agreement
with experimental fact for the value of bond energy and interatomic distance
than if a single structure alone were used. We say that the structure of
carbon dioxide “resonates" between the several forms, known as canonical

^Footnote: To help one appreciate the fact that 0=C — O and O — C=0 are
truly two distinct formulas and not identical, assume a COs molecule contains two
different oxygen isotopes, an O 10 atom and an O 18 atom. There are thus two possible
positions for the triple bond, between the C and O 18 and between the C and O 18 .

184

The Chemical Bond

forms . Such a situation where more than one structure, approximately equiva-
lent in energy, can be written for a molecule is called resonance. The term is
perhaps an unfortunate one in that it gives the impression that the several
different forms have a real, independent existence and that the molecule is
sometimes in one form and sometimes in another, or that an equilibrium
exists among the several structures. None of these is true in factl A molecule
has its own specific structure which may not be definitively known so that
it is clarifying at times to speak of distinct canonical forms which, however,
are merely mental constructs. A weighted mean of the properties of these
canonical forms would equal the properties of the molecule represented by
them. A compound with partial ionic character might be said to resonate
between covalent and ionic structures, each making a contribution to the
properties of the molecule. Thus for the HC1 molecule, two canonical struc-
tures are postulated, one purely covalent and one purely ionic, and the
actual molecular structure is conceived as a resotumce hybrid between the
two extremes (Figure 13.8).

H 5 CIS H+JCIJ-

«• •*

Figure 13.8. Resonance Forms of the HC1 Molecule.

Resonating structures differ only in the position of electrons and must have
the same numbers of paired and unpaired electrons. The concept of resonance
is very frequently used to describe molecular structure. A common example
is the benzene molecule, C 8 H 6 , for which two resonance forms are shown
(Figure 13.9).

^C'

“C - H

,C - H

Figure 13,$, The Benzene Molecule.

The energy calculated for a molecular system based upon resonating
structures is lower than the energy of any individual structure. In other
words, the stability of a resonance hybrid is greater than any of its canonical
forms. Generally the greater the number of resonance forms (of equivalent
energy) which can be written for a molecule the greater is the stability of
that molecule. As a familiar analogy, again we can resort to the game of
dice for clarification of this statement. A “seven” is more stable, that is,
more probable, than a “two” because the number of structures or ways in

The Chemical Bond

185

which it can be made up is greater than the number which can make up a
two.

The concept of resonance can be interpreted literally only in t§rms of
the integral calculus of which it is a direct consequence. The integratioh of
a differential equation yields many solutions. A linear combination, addition
or subtraction of these solutions, is itself a solution to the differential equa-
tion. For a molecular system, a combination of solutions to the differential
equations of quantum mechanics gives a lower energy for the system than
does any of the individual solutions. The separate solutions correspond to
the separate canonical forms of the resonance hybrid and the lowering of the
energy is the stabilizing energy due to resonance. The canonical forms have no
more real existence than do the mathematical equations representing them!

12. Potential Energy and the Molecule. The simplest covalent bond is
that holding two hydrogen atoms together in the hydrogen molecule, H J H.
Let us consider the interaction of the Is electrons of two hydrogen atoms,
initially a relatively large distance apart, as they approach each other to
form a hydrogen molecule. In Figure 13.10 the potential energy of the
two hydrogen atoms is plotted against interatomic distance. If the electrons
in the two hydrogen atoms have parallel spins, the atoms will repel each

Two hydrogen atoms having parallel
electron spins will repel each other and
not form a H 2 molecule. The negative
slope of the curve indicates a force of
repulsion between the hydrogen atoms.

(b)

Two hydrogen atoms having opposite
electron spins? can form a stable bond and
a H 2 molecule. The positive slope of the
curve from A to B represents the force
of attraction between the hydrogen atoms
separated by large distances. The nega-
tive slope from A to C represents a force
of repulsion when the atoms come closer
together than d 4 , the equilibrium dis-
jtance about which the atoms in the mole-
cule oscillate.

Figure 13.10. Potential Energy Curves for the Approach of
Two Hydrogen Atom*.

1S6

The 0 h 'Rond

other and a stable bond will not be formed. The negative slope of the curve
in Figure 13.10 a represents a repulsive force which increases as the atoms
come closer together. Two atoms with a total energy, E x would approach
to a minimum distance, d u and then fly apart without forming a bond.

If the electrons have opposite spins a stable bond can be formed. As the
atoms approach each other, attractive forces come into play as indicated
by the positive slope of the portion of the curve from point B to point A
in Figure 13.10 b. The potential energy decreases to a minimum point, A,
of maximum stability. Thereafter an approach of the hydrogen atoms to
closer distances gives rise to repulsive forces. Tire energy minimum in the
curve indicates the formation of a stable bond between the two hydrogen
atoms, that is, a hydrogen molecule. If the molecule has a total energy, E*
the hydrogen atoms will oscillate between a distance of closest approach,
d*, and of farthest separation, d 3 . The distance, d 4 , corresponding to the
minimum in the potential energy curve, is the equilibrium distance about
which the atoms oscillate. It is the value listed in tables as the interatomic
distance. For the H> molecule, this is 0.743 A. An ordinate between E 2 and
the potential energy curve, such as ef > represents the difference between the
total energy and potential energy and is thus the kinetic energy of the
hydrogen atoms. Where E 2 intersects the potential energy curve, at g and h „
the kinetic energy is zero and the energy of the system is solely potential.
At these points the oscillating atoms reverse their direction of motion. Much
the same viewpoint can be applied to any diatomic molecule and to the
forces of attraction and repulsion which produce the stable bond between
ions in an ionic bond.

13. The Nature of the Covalent Bond. The explanation of the nature of
the covalent bond has been one of the major achievements of quan-
tum mechanics. A covalent bond is produced when orbitals overlap.
The electron clouds of the orbitals of different atoms merge to produce an
increased density of electronic charge between the bonded atoms. This
concentration of charge between the atoms constitutes the bond and is what
we have represented very simply as a covalent electron pair. The increased
charge density between the atoms indicates a large probability of finding
the electron pair between the bonded atoms but these electrons have a finite
probability of being anywhere in the molecule. It is not possible to say that
either of .the bonding electrons belongs to this or that atom; they belong
to the molecule as a whole. The energy released as the atoms converge
in the process of overlapping is the bond energy. The greater the extent
of the overlap the stronger is the covalent bond. From this viewpoint,
Figure 13.11 shows different representations of the convergence of the
Is orbitals of two hydrogen atoms to form a hydrogen molecule.

The combination of oxygen with hydrogen to form water involves the
overlap of the Is orbitals of the two hydrogen atoms with the two incomplete

orbitals of an oxygen atom (Is 2 2s 2 2p y 2p„). We have seen in
Figure 12.4 that p orbitals are directed at right angles to each other. The
bonds which they form should be similarly oriented in space and we might
predict that the angle between the O— H bonds in water should be 90* as
shown in Figure 13.12. Actually the angle between the O—H bonds is 105°.

The Chemical Bond

187

The spherically symmetrical Is orbitals, shown as electron clouds, overlap to form the
covalent bond. The increase in charge density between the atoms in the H 2 molecule
constitutes the bond. The atomic nuclei are indicated by the dark centers.

The electron field intensity is plotted in a manner similar to that of contours on
a map. The formation of a bond is indicated by the fact that some contours link
both atoms in the H 2 molecule.

A vertical view of the electron field intensity of (b) is shown.

Figure 13,11 Formation of the Hydrogen Molecule by Overlap of
Atomic Orbitals.

A partial explanation is due to the polarity of the bond. The more electro-
negative oxygen atom draws to itself the bonding electron pair, and the
mutual repulsion of the relatively positive hydrogen atoms increases the bond
angle to greater than 90°. With the related molecules, H 2 S and H 2 Se, be-
cause of the larger size and lesser electronegativity of the central S or Se
atom, the bond angles are 92° and 90°, respectively. In covalent bond
formation, however, we have so far considered only the overlap of “pure”
orbitals. Within an atom orbitals may combine to form a new hybrid orbital L
Thus an s and p orbital of an atom may combine to form a hybrid sp
orbital which will have its own character and spatial orientation. These hybrid
orbitals can overlap other orbitals to form covalent bonds. The bond angle of
105° for H a O results primarily from a hybrid bond with about 20% s char-
acter and 80% p character. The subject of hybrid bonds will be taken up
in some detail during the study of carbon and its compounds (Chapter 32),
The distance between the centers of the nuclei of two covalently bonded
atoms is known as the bond distance . Half the bond distance between two

188

The Chemical Bond

similar atoms is the covalent bond radius of the atom. The C—C bond
distance is 1.54 A and hence the covalent bond radius of a carbon atom is
0.77 A. The value of a covalent bond radius depends upon whether the
atoms are linked by a single, double, or triple bond, the values decreasing in
that order. The C=C bond radius is 0.67 A, and that for C=C is 0.60
A. When added, covalent bond radii give a good approximation in many
cases to the actual bond length. Since the covalent bond radius of a hydrogen
atom is 0.30 A the C— H bond length in methane, CH«, can be calculated
to be 1.07 A. The greater stability of resonant structures, however, gives
them a shorter bond length than is calculated by this simple addition.

p orbital

s orbital

(a) General overlap of an s and a p orbital. The shaded area is the region

of overlap.

(b) Reaction of an oxygen atom with two hydrogen atoms. For the oxygen
atom, only the unfilled 2 p y and 2p % orbitals are shown. The filled 2 p orbital
is perpendicular to the plane of this page. The H — O — H angle shown is 00°

Figure 13.12 . Formation of a Water Molecule.

14. Properties of the Ionic and Covalent Bonds. The purpose of all theory
is to enable one to predict events on a macroscopic scale. The theory of the
chemical bond, whether ionic or covalent, must be related to experimentally
observable criteria. In the tabulation of properties which follows, the extreme
cases of a purely ionic structure and a purely covalent structure are compared.
We have seen that covalent substances can have a partial ionic character and
this will be reflected in their properties in that such covalent compounds
will have somewhat ionic characteristics. Indeed a series of compounds could
be presented wherein there would be a gradation in properties from the purely
covalent to the completely ionic.

The Chemical Bond

189

(A) The nature of the bond : Ionic compounds consist of distinct electrical-
ly charged particles, the ions, bound by Coulombic electrostatic attraction.
Covalent compounds consist of discrete molecules within which the bonding
between atoms is due to the sharing of electron pairs. It is this difference
which underlies the divergence in properties between ionic and covalent
substances.

Sodium chloride is an ionic solid. In a crystal of sodium chloride, the
Na + and Cl" ions are arranged in a regular cubic array. Each Na+ ion is
surrounded by six equidistant Cl" ions, and each Cl“ ion is surrounded
similarly by six Na+ ions. The crystal lattice is held together by the electro-
static attraction of neighboring Na + and Cl“ ions (Figure 13.13). In the
sense that there is no one specific Na+ ion which “belongs” to a specific Cl"
ion, there are no individual molecules of NaCl, though of course the numbers
of Na + and Cl" ions must be equal. The entire crystal of sodium chloride
might be visualized as a single giant molecule. In the vapor state, however,
individual molecules of NaCl do exist.

Iodine is a molecular solid. In a crystal of iodine, I 2 , an iodine atom
is linked to one other iodine atom by a covalent bond. No ions are present.
Since one iodine atom is joined to only , one other iodine atom the two can
be distinguished as a separate entity, the I 2 molecule (Figure 13.14). Be-

Figure 13.1 3. Sodium Chloride, Figure 13.14. Iodine,

rm Ionic Crystal. a Molecular Crystal.

tween the molecules there are Van der Waals forces similar to the attractive
forces between molecules of a gas. Crystals held together by Van der Waals
forces are sometimes called molecular crystals-^ Van der Waals’ forces are
also electrical in nature. Molecules vibrate and rotate so that the electrical
fields about molecules are not static. Actually the electron cloud patterns
we have been drawing represent average values over a period of time. At
any instant the electric field within a molecule may be oscillating so that
it is more concentrated in one region of the molecule than in another. The
mere proximity of two such molecules, which are in effect oscillating dipoles,

190

The Chemical Bond

results in an attraction which holds the separate molecules together. These
attractive forces are the Van der Wauls forces. They are of significant
magnitude only when molecules are very close as they are in a solid or
liquid or highly compressed gas. Boiling involves the separation of individual
molecules against Van der Waals attraction so that the boiling point is a
measure of these forces. The Van der Waals forces depend upon the number
of electrons and the surface area of a molecule. In turn these are both
roughly proportional to the molecular weight of a substance so that we
should expect to find that the boiling point increases with an increase in
molecular weight. Such a relationship is evidenced by the boiling points of
the inert gases.

Substance Helium Neon Argon Krypton Xenon Radon

Boiling Point, °C -269 -246 -186 -153 -107 -61

The ionic bond and the covalent bond are approximately equal in strength,
that is, the energy required to disrupt them is about equal. Compared with
these interatomic forces the Van der Waals forces are weak. It is relatively
easy to separate one molecule from another held by Van der Waals forces.
Thus molecular crystals are generally soft. Examples of molecular crystals
are solid carbon dioxide and naphthalene.

(B) Melting point and boiling point : The melting and boiling points
of covalent compounds are generally lower than those of ionic compounds.
The following data for typical ionic and covalent compounds illustrate this
point.

Ionic Compounds

Melting Point, °C

BoSing Point, *C

Sodium Chloride, NaCl

801

1413

Tin (II) Chloride, SnCl 2

246

623

Aluminum Oxide, Al 2 O a

Covalent Compounds

2050

2250

Carbon Tetrachloride, CCl*

-23

Tin (IV) Chloride, SnCl 4

-33

114

Aluminum Iodide, All 3

191

360

To produce a change in state of an ionic compound the strong electrostatic
forces between ions must be overcome. With a covalent substance only the
weak Van der Waals' forces between molecules must be overcome to separate
one molecule from another so that the melting points and boiling points
of these compounds will be relatively low. The covalent bond within a mole*
cule is not broken in a change of state. For similar reasons ionic substances
have low volatilities whereas covalent compounds have relatively high vapor
pressures. Indeed, iodine molecules sublime, that is, pass directly from the
solid to the vapor state without intermediate liquefaction.

(C) Electrical conductivity : A molten ionic substance or an aqueous
solution of one will conduct electricity, A covalent substance is a non*
conductor. Sodium chloride fused or dissolved in water is a conductor, but
carbon tetrachloride is not,

(P) Solubility: Ionic substances, are generally soluble in polar solvents,
that is, solvents whose molecules are polar, and insoluble in nonpolar solvents.

The Chemical Bond

191

For covalent compounds the converse is true. In water, a polar solvent,
NaCl is quite soluble, I* is slightly soluble, and CC1 4 not at all soluble. In
the nonpolar solvent, benzene, NaCl, is insoluble but I 2 and CC1 4 are readily
soluble.

(E) Rates of chemical reaction : With ionic compounds the rate of a
chemical reaction is generally very rapid. The reaction between solutions
of NaCl and AgN0 3 , both ionic compounds, is well nigh instantaneous.

Na+Ch + Ag+NCV Na+N0 3 - -f Ag+ Cl*

Reactions between covalent substances proceed more slowly and may take
hours or even days before the reaction is complete. The conversion of
sugar, C*H 12 0«, to ethyl alcohol, C 2 H 5 OH,

C fl H 32 0« 2 C 2 H 5 OH+ 2 C0 2

and the synthesis of ammonia from its elements

N* 4* 3 H 2 -> 2 NH 3

are examples of reactions between covalent compounds which take place
slowly.

QUESTIONS

1. Define ionization potential. What does the second ionization potential of an
element measure? Why is the second ionization potential greater than the
first?

2. How does the first ionization potential vary with the position of an element
in the periodic table?

3. Distinguish between eleirtrovaUent and covalent bonds. Give examples of each.

4. Illustrate the electron transfer between the following which react to form an
ionic bond: K and Br; Ca and S; Mg and Br; A1 and O.

5. Distinguish between oxidation number and valence, using A1C1 S as an example.

6. Calculate the oxidation numbers of the following atoms: (a) S in H 2 S (b) S
in H 2 $0 4 (c) Ga in CaSO* (d) A1 in A1 2 (S0 4 ) 8 (e) K in KN0 3 (f) Mn in
KMnb 4 (g) Cr in K 2 Cr 2 0 7 (h) N in NH*C1 (i) the N atoms in NH 4 NT0 3 .

7. Define and illustrate the coordinate covalent bond.

8. Draw electron configurations for H a S, CaCl 2 , PC1 3 , SiBr 4 , Na 2 0, H 2 0 2 , N 2 H 4T
CH a O, CH s -CH 2 C1. Indicate whether each substance is ionic or covalent.

9. Draw electron configurations for the ions: N0 3 ~, C10 3 *, C10 4 _ P0 4 8 ' 0 2 2 \

IQ. In the S0 4 2 * ion, indicate which bonds are ionic, covalent, and coordinate

covalent.

11. Why does the compound Na s AlF 6 exist but Na 3 BF 6 does not? Why does
PCI* exist but not N€l s ?

12. From the viewpoint of electron configurations, is (CH 3 ) 3 N a stable molecule?

13. Would it be possible to form a coordinate covalent bond between NH a and an
oxygen atom?

14. Define oxidation and reduction in terms of electrons. Write an equation for
a reduction and an equation for an oxidation.

The Chemical Bond

192

15. What is meant by the electronegativity of an atom? How is it related to
ionic and covalent bonds? Distinguish between electronegativity and electron
affinity.

18. Define (a) polarity (b) polar molecule (c) dipole (d) partial ionic character.

17. Define and illustrate what is meant by a hydrogen bond. Between what
atoms are hydrogen bonds formed?

18. What is the concept of resonance? Write resonance structures for CO*, N,Q,
C # H # , SO a , Oj(ozotie). and the NO;,- ion.

19. What is the relation between (a) atomic radius and ionization potential (b) atomic
radius and electronegativity? illustrate these relationships with typical groups
of the Periodic Table.

20. What is meant by effective nuclear charge?

21. Draw and explain the potential energy diagram for a stable diatomic molecule.

22. What is meant by the “overlapping of orbitals”? Illustrate this for the water
molecule,

23. What orbitals are used in the formation of the NH ; , molecule? Predict the
spatial orientation of the N-H bonds.

24. From the viewpoint of electron configuration explain why the iron atom ex-
hibits two oxidation states, +2 and +3. Why is a Fe 4 1 ion unlikely?

25. Discuss the effect of ionic and covalent bonding on the following properties
(a) electrical conductivity (b) melting and boiling points (c) solubility in polar
and nonpolar solvents (d) rate of chemical reaction (c) hardness.

26. The heats of atomization for one mole of H, and one mole of 0, are both
endothermic; +104.2 kcal and +118.3 kcal, respectively. The heat of forma-
tion of one mole of gaseous water is -57.8 kcal. Using Hess' .Law, calculate
the O-H bond energy.

27. Calculate the lattice energy of potassium chloride, KCi, given the reaction

K(g) + % Cl,(g) -» K+Cl-(s) AE = -105 kcal

28. Elements A, B, and C have atomic numbers Z, Z +2, and Z + 3, respectively.
B is a noble gas. (a) Which element has the highest value of ionization poten-
tial? (b) Which element has the highest value of electronegativity? (c) What is
the formula of a compound formed between A and C? (d) What type of bond-
ing would this compound probably have? (e) What type of bonding would a
compound between A and the element just to its left in the Periodic Table
probably have? (f) What type of bonding would a compound between A and
the element just below, it in the Periodic Table probably have? (g) Would the
'Compound in (f) be polar or nonpolar?

Oxygen

Oxygen is the most abundant element on earth by weight, comprising
about 47% of the earth's crust. By weight the atmosphere contains 21%
oxygen and pure water contains 89% oxygen. Most common minerals such
as limestone, granite, feldspar, and sandstone contain high percentages of
oxygen, as do all plants and animals in their tissues. Oxygen was discovered
that is, first prepared in the pure state as elemental oxygen, by Carl Wilhelm
Scheele, a Swedish apothecary, sometime before 1773. Publication of this
work was delayed and did not appear until 1777. Meanwhile, in 1774, Joseph

Table 14-A

Properties of Oxygen

Symbol

Molecular formula

Atomic Number

Atomic Weight

15.9994

Electron Configuration

is 2 2*2 2p*

Isotopes (Mass Number

16 (99.91)

Abundance in Earth's

and Percent)

17 ( 0.01)

Crust, wt %

47.3

18 ( 0.08)

Physical State at STP

colorless gas

Ionization Potential, 1st,

13.61

Density at STP, g/1

1.43

Density of liquid, g/ml

1.14

Electron Affinity, eV

1.47

Melting Point, °C

—218.4

Electronegativity

3.5

Boiling Point, °C

-183.0

Oxidation Numbers

-2, -1

Critical Temperature, °C

118.8

Covalent Radius, A

0.74

Critical Pressure, atm

49.7

Ionic Radius, A

1.40 (-2)

Solubility in water at STP,
ml/1 ; 48.9

Heat of Atomization,

kcal/g-atom : 59.2

M is ceU aneotts; Caseous oxygen is colorless, odorless, and tasteless; liquid and
solid oxygen are pale blue. The O a molecule is paramagnetic.

194

Oxygen

Priestley, an English clergyman and amateur scientist, independently pre-
pared oxygen by heating mercury(II) oxide, HgO. Priestley is usually credited
with the discovery of oxygen. Scheel es priority as the discoverer of this
element was established in 1892, more than 100 years later, mainly on the
basis of the excellent laboratory notes that he kept and which, fortunately,
had been preserved.

1. Preparation of Oxygen, (A) From air: Since' atmospheric oxygen is
chemically uncombined in the mixture of gases we call “air,” it can be sepa-
rated from the other components of the atmosphere, primarily nitrogen
and argon, by simple mechanical means without resorting to a chemical
process. The process used is the distillation of liquid air. Liquid air is pro-
duced in a machine whose operation is described ir* Figure 14.1,

Figure 14.1. Liquid- Air Machine.

Air purified from dust, carbon dioxide,
and water vapor is compressed to about 200
atmospheres by pump A. The resultant
heat of compression is absorbed in the
cooler B. The cold, high pressure air en-
ters the liquefier C via the inner tube of
a concentric pair of tubes. At tbe end of
the inner tube is an orifice D through
which the air expands to a lower pressure
in chamber E. This expansion results in a
cooling effect. The cold air is led back
through the outer tube of the spiral F to
precool the succeeding incoming high pres-
sure air. On expansion this preoooleo air
becomes still lower in temperature until
ultimately expansion results in sufficient
cooling to reduce the temperature of the
atmospheric gases below their critical tem-
peratures andT ultimately their bolHng points,
whereupon drops of liquid air form and
collect in E. The liquid air k then dktlled.

Since nitrogen and argon have boiling points lower than oxygon they
come off first when liquid air is distilled, leaving a residue of pale bk«
liquid oxygen, commonly referred to as LOX. The liquid oxygen ir vapor-
ized and stored in steel cylinders at pressures about 2,000 lb/in*. Ordinary
temperatures are above the critical temperatures of these substances so that
neither liquid air nor liquid oxygen can be kept indefinitely at these tem-
peratures. They can be kept as liquids for appreciable periods of time in
a special double-walled container, or Dewar flask. The inside walls of the
container are silvered and the space between the walls is evacuated to
minimize the transfer of heat from the surroundings to the contents of the
flask. In time the liquid air boils off. The common thermos bottle is similar
to the Dewar flask.

(B) From compounds of oxygen : It would appear that water should
be a simple source of oxygen by decomposing it into its elements hydrogen
and oxygen, thus; 2 H 2 0 2 H 2 + 0 2 . Water is, however, a very stable

compound. Heating it to 2,00Q°C causes only about 2% decomposition. Yet
the decomposition of water can be accomplished at ordinary temperatures

Oxygen

195

by electrical energy. When a direct current is passed through water con-
taining a little sulfuric acid or sodium hydroxide, since pure water alone
is a poor conductor of electricity, the water is decomposed into hydrogen
and oxygen according to the foregoing equation. Oxygen gas is formed at
the positive electrode and hydrogen gas at the negative electrode. The
process is called electrolysis . The reaction which takes place at each elec-
trode is:

(1) 4 H 2 0 + 4 e 2 H 2 (g) + 2(OH)— (at the negative electrode)

(2) 2 H 2 0 — 4 e — ► 0 2 (g) + 4 H+ (at the positive electrode)

The action of the electric current is discussed in Chapter 25. In this or
in any electrolysis the ordinary laws of stoichiometry hold and the products
of the electrolysis of water are the same as would be formed by its thermal
decomposition. Thus the ratio by volume of hydrogen to oxygen produced
is 2 to 1 and the weight ratio is 1 to 8. Figure 14.2 shows a convenient
apparatus devised by A. W. Hofmann for collecting the products as they
leave the electrodes.

Figure 14.2 . Electrolysis of Water.

During the passage of a direct electric 'current sup-
plied by the battery, B, oxygen is liberated at the
positive electrode, while hydrogen is liberated at the
negative electrode. The ratio by volume of hydrogen
to oxygen is 2 : 1. The apparatus is designed to prevent
the mixing of the hydrogen and oxygen. The gases rise
to the surface of the confining liquid (water) and are
collected in the separate arms of the apparatus.

Other oxygen compounds are not so stable as water and can be more
easily decomposed by heating. Certain oxides (a binary compound with
oxygen) and also more complex compounds of oxygen can be completely
decomposed whereas others yield but a part of their oxygen content.

(3) 2 HgO -> 2 Hg + 0 2 (heating of mercury (II) oxide)

(4) 2 BaO 2 BaO + O g (heating of barium peroxide)

(5) 2 KClOg-* 2 KCI + 3 0 2 (heating of potassium chlorate)

The thermal decomposition of potassium chlorate, KC10. H , is a common
laboratory preparation for oxygen. This reaction is catalyzed by the addi-
tion of manganese dioxide, MnO z .

196 Oxygm

Oxygen is also generated when water is dropped slowly on solid sodium
peroxide, Na 2 0 2 .

(6) 2 Na,O s + 2 H 2 Q O- + 4 NaOH

A laboratory preparation is one which can be carried out on a small scale
with not too complex equipment to produce small quantities of a substance for
experimentation (Figure 14.3), Cost is not a primary consideration. In a com-
mercial preparation, economics is a major consideration and apparatus which can
handle large flow rates are usual. Thus the liquefaction of air would be a com-
mercial, and not a laboratory, preparation of oxygen. The electrolysis of water
lends itself to laboratory preparation and could be a commercial preparation of
oxygen only where electricity is cheap.

Figure 14.3.

Laboratory Preparation of Oxygen.

A mixture of KCIO, and a little MnO a is
heated in a test tube. Because oxygen is not
appreciably soluble in water it is collected in
a bottle, initially filled with water, by displace-
ment of the water In a pneumatic trough. The
displacement of a liquid in which a gas is
insoluble is a standard technique for collect-
ing a pure sample of a gas.

2, Properties of Oxygen. (A) Physical properties ; The physical proper-
ties of oxygen are given in Table 14-A.

(B) Chemical properties: It might be expected that the 0 2 molecule,
in accordance with a valence of two for the oxygen atom, would have a

double covalent bond, 5Q:sOS , but its paramagnetic property indicates

•» •*

a structure containing unpaired electrons, :0 : O: .

At ordinary temperatures, oxygen is quite inactive but at higher tem-
peratures it combines with almost all the elements to form oxides, with
the evolution of much energy.

(7) C + O* -+ CO* (above 500°C)

(8) 4 Fe + 3 O* 2 Fe*O s (slowly at low temperature, but rapidly at

high temperature)

(9) 2 F* + O s 2 F*0 (occurs even with LOX, because of high

reactivity of fluorine)

Compounds of elements that react with oxygen will also react in such
a manner that each element of the compound will form an oxide*

Oxygen

197

(10) 2 C 2 H 2 + 5 O a — > 4 C0 2 + 2 H 2 0 (burning of acetylene, C 2 H 2 ;

the oxyacetylene torch )

(11) 2 ZnS + 3 0 2 -> 2 ZnO + 2 S0 2 ( roasting of an ore, zinc sul-

fide, ZnS )

(12) C 2 H 5 OH + 3 0 2 -» 2 C0 2 + 3 H 2 0 ( oxidation of ethyl alcohol,

C 2 H 5 OH)

The oxides of nonmetallic elements, e.g, S0 2 and N 2 Oj;, when dissolved
in water, yield acidic solutions and so are called acid anhydrides , whereas
the oxides of metallic elements, e.g., CaO and Na 2 0, are basic anhydrides
in that they form basic aqueous solutions. Acids and bases are considered
in detail in Chapter 20.

(13) SO a + H 2 0 H 2 S0 3 (sulfurous acid)

(14) CaO + H 2 0 -» Ca(OH) 2 (calcium hydroxide; a base)

3. Uses of Oxygen. Apart from its being necessary to sustain terrestrial

life, atmospheric oxygen is used in many industrial processes. The metal-
lurgical roasting of ores, the oxidation of sulfur in the preparation of sul-
furic acid, the ammonia oxidation process in the manufacture of nitric acid,
the oxidation reactions in the several furnaces used in the production of
iron and steel all involve reactions with oxygen. So do the combustion re-
actions of gas and oil which are still the prime source of industrial energy
and the oxidation of food to furnish body energy.

A fuel system consists of a combustible material and an oxidizer. The
ordinary internal combustion engine and all atmospheric jet engines rely
upon the most common oxidizer, atmospheric oxygen, to react with the
fuel used, e.g., gasoline. To operate in outer space where there is no at-
mosphere, a rocket engine, which is a type of thermal heat engine, must
incorporate within its mechanism both the fuel and the oxidizer. Liquid
oxygen is a commonly used oxidizer for fuels such as gasoline, alcohol,
hydrogen, and hydrazine. Combustion is not limited to the burning of a
substance in oxygen or in air. Other gases will support combustion; thus
hydrogen will burn in fluorine. For combustion to occur three requirements
must be met. First, there must be a combustible substance, that is, a fuel.
Carbon dioxide, C0 2 , will not bum because the carbon is already completely
oxidized. Second, there must be an oxidizer and last, the combustible
substance must be raised to its kindling temperature, a function which
may be performed by a spark or a match. The visible evidence of most
combustions, namely a flame, is merely the burning of one gas in another.
Wood bums with a flame because the heat distills volatile substances from
it; nonvolatile charcoal glows and bums without flame.

Since the rate of a chemical reaction depends upon the concentrations
of the reacting substances, a reaction with pure oxygen will occur more
rapidly and produce a higher temperature than with atmospheric oxygen.
In addition when a substance burns in air, a part of the heat produced’
is used to warm the atmospheric nitrogen so the temperature does not rise

198

°«««*

as high as when pure oxygen is used. Being highly exothermic the re-
actions of hydrogen and of acetylene with pure oxygen are applied as
the oxyhydrogen and oxyacetylene torches, respectively. Temperatures of
3,000°C, sufficient for the welding of steel or the cutting of iron plates
several inches thick, can be attained. A Bunsen burner, using natural gas
which is primarily methane, CH t , and air can attain temperatures of about
1,500° C.

Ozone

The substance ozone is composed only of oxygen atoms and has the
formula 0 3 . Such an existence of different forms of an element is called
allotropy . Ordinary oxygen and ozone are said to be allotropcs. Allotropy
also results when an element exists in two or more crystalline forms, e.g.,
rhombic and monoclinic sulfur, graphite and diamond.

Table 14-B _

Properties of Ozone

Molecular formula

o :t

Molecular weight

; 48.00

Melting point. °C

—250° C

Boiling point,

: — 112°C

Critical temperature, C

- 5°C

Density at STP, g/i

: 2.14

Solubility in water at STP,

500

Density of liquid, g/ml

: 1.71

ml/1

Miscellaneous: Caseous O* is colorless and has a pungent odor; liquid O a is blue.

Ozone is prepared by passing oxygen or air through an electric discharge.
The reaction is endothermic and a mixture containing about 5% of O a is
produced.

(15) 3 O* 2 0 3 AH = +68 kcal

In the stratosphere some O s is formed by the action of ultraviolet light.

Since the O s molecule is not paramagnetic, all of its electrons must be
paired. The structure of the molecule is best explained on the basis of two
resonance covalent forms (Figure 14.4).

An O. molecule contains two covalent bonds, one of which is a double covalent
bond. The double bond resonates between the two 0-0 bonds to produce the
two forms. In both the bond angle is 127° and the bond distance is 1.26 A.

Figure 14.4. Resonance Forme of the Oasone Molecule.

Ozone is unstable and slowly reverts to oxygen; liquid ozone may de-
compose explosively. Ozone is a more potent oxidizing agent than oxygen;
of the common substances only fluorine is a stronger oxidizing agent. Be-
cause it can oxidize many colored organic compounds to colorless products

Oxygen

199

ozone acts as a bleaching agent for oils, waxes, flour, and ivory. A fuel
system of ozone and cyanogen, C 2 N 2 , gives a higher combustion tempera-
ture than any other rocket propellant system. Since microorganisms are
effectively destroyed by ozone, it is used to sterilize drinking water.

QUESTIONS

1. Draw the nuclear structures and the electron configurations of the three
isotopes of oxygen.

2. What principles are involved in the production of oxygen by the liquefac-
tion of air?

3. Draw electron configurations which illustrate why the oxygen molecule is
paramagnetic and the ozone molecule is not.

4. Distinguish between isotope and allotrope. What are the isotopes and allo-
tropes of oxygen?

5. Write balanced chemical equations for the preparation of oxygen.

6. Write balanced chemical equations for the complete combustion of the fol-

lowing: (a) cadmium sulfide, CdS (b) propane, C 3 H 8 (c) hydrogen sulfide,
H 2 S (d) benzene, C*H 6 (e) phosphine, PH 3 (f) octane, C S H IS .

7. Define and illustrate the following terms: (a) oxide (b) acid anhydride (c) oxida-
tion (d) combustion (e) flame.

8. Using the Periodic Table, write formulas for the oxides of silicon, radium,
scandium, rubidium, antimony, sulfur, and chlorine.

9. Why should ozone be a more potent oxidizing agent than oxygen?

10. Fifty grams of KC10 3 are completely decomposed into KC1 and 0 2 . The

0 2 is collected over water at 20 °C and a total pressure of 759 mm. Cal-

culate (a) the volume of the 0 2 collected (b) the weight of the 0 2 .

Ans: (a) 15.2 liter

11. Which requires more oxygen for complete combustion: 10 g of ethane, C 2 H e ,
or 10 g of ethyl alcohol, C 2 H 5 OH?

12. What weight of prosphoric acid, H s P0 4 , can be formed by the combustion
of 25 g of phosphorus followed by combination of the product with water?

Ans.: 79 g

13. What volume of oxygen would be produced at STP by the addition of excess

water to 52 g of Na 2 0 2 ? Ans: 7.5 liter

14. (a) What volume of oxygen would react with 100 liters of' methane, CH 4 , both

gases being measured at 20 °C and one atmosphere pressure? (b) What will
be the volume of the resulting products? Ans: (a) 200 liters

15. (a) What weight of oxygen would be produced by the electrolysis of 49.5 g

of water? (b) What weight of hydrogen will be produced simultaneously?

Ans: (a) 44.0 g

16. "What weight of Ca(OH) 2 can be formed from the CaO produced by the
heating of 7.0 g of calcium in air?

17. A one gram sample of KC10 3 is heated and 150 ml of oxygen are collected
at STP. What percent of the KC10 S was decomposed?

18. If the one gram sample of KClOs in problem 17 contained an impurity
(sand) and if it were heated tiH no more .oxygen was evolved, and if the
volume of oxygen collected at STP were 150 ml, what is the percent of
KCIO3 in the initial sample?

Hydrogen

Though only a trace of elemental hydrogen is found in the earths at-
mosphere, less than one part in 200,000, hydrogen is widely distributed. As
combined hydrogen it constitutes 11% by weight of water and is a con-
stituent of almost all organic compounds and most acids and bases. On
the sun hydrogen is very abundant; it is believed that the source of solar
energy is the nuclear conversion of hydrogen into helium (page 645).
Hydrogen was first prepared in 1766 by the English physicist. Sir Henry
Cavendish. It was so named by the French chemist, Antoine Lavoisier,
from the fact that it burned to form water.

L Preparation of Hydrogen. (A) From water: Since elemental hydrogen
does not occur in nature it must be prepared by liberation from its com-
pounds. Water is the most abundant and cheapest source of hydrogen. We
have already noted that water is a compound very stable to thermal decom-
position, and that both hydrogen and oxygen can be prepared by the elec-
trolysis of water which has been made conducting by the addition of a
small quantity of sulfuric acid or sodium hydroxide. Electrolytic H a , 99.9%
pure, is the purest form of commercial hydrogen.

1) Reaction of metals with water: The very active alkali metals and
alkaline earth metals displace hydrogen from water. With the alkali metals
the reaction is violent, even explosive. Other less active metals require
an increased temperature for this reaction to take place.

(1) Na + H s O

-» NaOH

H s

( takes place at room tempera-
ture )

(2) Mg + H.O

-*■ MgO

{with boiling water)

(3) 3 Fe + 4 H.O

^ Fe 3 0 4

+ 4H,

(with red hot iron and super-
heated steam)

2) The water gas reaction : When steam is passed over red hot coke
at 600°C a mixture of hydrogen and carbon monoxide, called water gas,
is formed.

(4)

G + H 2 0 H 2 + CO

Hydrogen

201

Table I5-A

Properties

of Hydrogen

Symbol

Molecular Formula

Atomic Number

Atomic Weight

1.00797

Electron Configuration

Isotopes (Mass Number

1 (99.985)

Abundance in Earth's Crust, wt % : 0.88

and Percent)

2 ( 0.015)

Physical State at STP

colorless gas

3 (trace )

Density at STP, g/1

0.0899

Ionization Potential, 1st,

Density of liquid, g/ml

0.071

13.60

Melting Point, °C

-259.2

Electron Affinity, eV

0.70

Boiling Point, °C

-252.7

Electronegativity

2.1

Critical Temperature, °C

-234.5

Oxidation Numbers

+ 1, -1

Critical Pressure, atm

12.8

Covalent Radius, A

0.37

Solubility in water at STP,

Ionic Radius, A

2.08 (-1)

ml/l

21.5

Heat of Atomization,

kcal/g-atom

: 52.1

Miscellaneous: Gaseous hydrogen is colorless, odorless, and tasteless. It is the

lightest substance known.

Water gas is a commercial fuel* The H 2 is difficult to separate from the
CO but when the mixture is further reacted with steam, using a suitable
catalyst, additional H 2 is produced and the CO is converted to C0 2 . This
is readily removed from the H 2 by dissolving it in cold water or in a solu-
tion of alkali.

(5) CO + H a O H 2 + C0 2

(B) From acid : In an acid the concentration of hydrogen ion, H+ is
greater than in water. Since the rate of a reaction depends upon the con-
centration of the reactants, many metals that react slowly or not at all with
water react readily at room temperature to displace hydrogen from acids.

(6) Zn + 2 HC1 H 2 + ZnCl 2 (a common laboratory preparation)

(7) Fe + H 2 S0 4 -> H 2 + FeS0 4

Certain inactive metals such as copper, silver, mercury, and gold react
neither with water nor acid to produce free H 2 .

(C) From alkali : Aluminum reacts with concentrated alkali as sodium hy-~
droxide, NaOH, to produce H 2 and the aluminate ion, Al(OH) 4 ~ Zinc and
silicon undergo similar reactions.

(8) 2 A1 + 2 NaOH + 6 H 2 0 ->3H 2 + 2 NaAl(OH) 4 (sodium aluminate)

(D) From hydrides : Compounds of active metals and hydrogen, the hydrides,
liberate H 2 when treated with water. With calcium hydride, CaH 2 , the
reaction is:

(9)

CaH 2 + 2H 2 0 2 H 2 + Ca<OH) 2

202

Hydrogen

2. Properties of Hydrogen. (A) Physical properties : Table 15-A.

(B) Chemical properties : At ordinary temperatures hydrogen is relatively
inert but, under the proper conditions, it will combine directly with most
elements.

1) Reaction with nonmetals : At room temperature a mixture of H 2 and
0 2 does not react but, if ignited and if the concentration of each reactant is
within certain limits, the reaction is explosive. So too, mixtures of H 2 and
air may be explosive. Pure hydrogen does not explode but will bum quietly
at its air interface. Hydrogen also combines explosively with fluorine, even
at liquid hydrogen temperatures, and chlorine, but less readily with bro-
mine, iodine, sulfur, and nitrogen.

(10) H* + X 2 2 HX (X 2 represents a halogen element; F 2 ,C1 2 ,

(11) H 2 + S — * H 2 S (at elevated temperatures)

(12) 3 H 2 + N 2 2 NH S (Haber Process for ammonia, NH»)

Except for the reaction with I 2 all these reactions are exothermic.

2) Reaction with metals: Upon heating to a few hundred degrees many
metals combine with H 2 to form hydrides . The hydrides of the very active
elements, as LiH, NaH, and CaH 2 , are electrovalent compounds wherein the
hydrogen atom has gained an electron to form the unusual hydrogen nega-
tive ion, with an oxidation number of -1, H".

(13) £ Na + H 2 2 Na + H (hydrogen and liquid sodium at 400°C)

The fused metal hydrides conduct electricity. In the electrolysis of molten
calcium hydride, CaH a , for example, the H~ ion would migrate to the
positive electrode, lose an electron and be liberated as free H 2 . The metal
hydrides are strong reducing agents. UH can reduce CO a to carbon and,
because of its low density, finds use in the removal of CO, from the atmosphere
of a space ship.

3) Reaction with compounds : Because of its tendency to combine with
oxygen, hydrogen reacts with many metal oxides and reduces them ^ to
the free metal.

(14) CuO + H* Cu + H a O

(15) Fe*0< + 4 H 2 3 Fe + 4 H t O

With unsaturated organic compounds, H* is added directly to produce
saturated compounds, a reaction known as hydrogenation. This is illustrated
by the reaction of ethylene, C a H*, and H a , to form ethane, C a H«.

(16) H H

H H

Hydrogen

The hydrogenation of unsaturated liquid oils is used commercially to con-
vert liquid fats, such as vegetable oils, into solid fats (Crisco, Spry). Under
suitable conditions, H 2 will react with CO to form a variety of compounds
ranging from methyl alcohol, CH 3 OH, to long chain carbon compounds.

(17) CO + 2 H 2 CH 3 OH

This reaction is the basis of the catalytic Fischer-Tropsch process whereby
Germany produced much of its gasoline during World War II. Coal can
also be hydrogenated to produce liquid hydrocarbons.

3. Isotopes of Hydrogen. Besides ordinary hydrogen of mass , number
one, 1 H 1 , two isotopes are known. One of these, having a mass number of
two, has been given the name deuterium , and a symbol, D or a H 2 ; the other
with a mass number of three, a H 3 , is called tritium. The nucleus of the deu-
terium atom is the deuteron and water formed from deuterium, D 2 0, is
known as heavy water. In ordinary water, about one* atom in every 4,500
hydrogen atoms is deuterium* tritium is radioactive and occurs in natural
hydrogen only to the extent of 1 part in 10 lT

In general the properties of isotopes are almost identical, particularly for
isotopes of high mass number. Because of the relatively large percentage
difference in mass between the nuclei of the hydrogen isotopes, however,
an appreciable difference in their properties exists. Some of the physical prop-
erties of H 2 and D 2 , and of ordinary water, H 2 0, and heavy water, D 2 Q, are
compared in Table 15-B.

Table 15-B

Physical Properties of H 2 , D 2 , H 2 0, and D a O

Property

h 2

d 2 h 2 o

d 2 o

Molecular weight

2:016

4.028 13.016

Melting point, °0

-259.2

-254.5 0.00

3.80

Boiling Point, °C

-252.7

—249.5 100.00

101.42

Density at 20° C

0.0835 g/1

0.1673 g/1 0.998 g/ml

1.106 g/ml

Solubility of NaCl, g/1000 g

359

305

Deuterium was discovered in 1932 by H. C. Urey, F. G. Brickwedde,
and G. M. Murphy. The production of heavy water, that is, the separation
of D 2 0 from ordinary water, is discussed in Chapter 49.

4. Order of Activity of the Metals. The different abilities of metals to
displace hydrogen from water and from acids would indicate that the metals
can be arranged in an order of chemical activity. This order is given in
Table 15-C. In addition to the displacement of hydrogen this order of ac-
tivity includes various types of reactions involving metals, such as the re-
duction of metal oxides by hydrogen or by carbon, and the decomposition
of such oxides by heat.

The activity series is useful in predicting the direction in which certain
chemical reactions will proceed since the relative activities of the elements
in this series applies not only to the displacement of hydrogen but also to

204

Hydrogen

Table 15-C

Order of Activity of the Metals

React with cold water; react vio- ^

lently with acids

' LS ^

l K 1

t Ba j
I Sr
Ca \
i Na I

Oxides not reduced by hydrogen
' or carbon monoxide

React with steam and acids, as

HC1 and H 2 $0 4 to liberate hy-
drogen |

f Mg

i Al 1
Mnj
1 Zn '

jFei
Ni '
Sn

i Pb j

| Oxides reduced by carbon at high
| temperature

H (

React with oxidizing acids as j

HN0 3 and H 2 S0 4 but do not j
liberate hydrogen \

[»

! Ag
k Hg 1

| 1

Oxides reduced by hydrogen or
carbon monoxide at high tem-
| peratures

Dissolve only in aqua regia \

(1 HCl : 3 HNO :i ) j

: Pt j

: Au /

f !

| Oxides decomposed by heat

the displacement of any element in the series from its compounds. In gen*
eral a higher element, that is, a more active element, will displace any ele-
ment below it from compounds of the lower element. If iron metal is
placed in a solution of copper sulfate, CuS0 4 , the copper is displaced by
the iron; copper metal is set free and iron goes into solution as FeSCh.

(18) Fe + CuSO, -> Cu + FeSO,

The reverse reaction, the displacement of iron from FeSO« by Cu, will not
take place spontaneously. Silver, which is below copper in the activity
series, does not displace copper from its compounds, nor will copper dis-
place hydrogen from acids to give free H».

The oxides of the metals vary widely in their stability. The oxides of the
least active elements, such as Ag*0 and HgO, are easily decomposed by
the heat of the Bunsen burner, whereas oxides of the very active elements,
such as CaO and MgO, are not decomposed even at the high temperature
of an electric furnace, 3,000°C. This relative stability is also related to die
activity of the metals. Active metals form stable compounds that are dif-
ficult to decompose and conversely, compounds of inactive metals are rela-
tively unstable and easily decomposed. The activity series is also known
as the Electromotive Series. Quantitative data relating chemical activity to
electric potentials can be obtained by electrochemical measurements as dis-
cussed in Chapter 24.

Hydrogen

QUESTIONS

1. Draw the nuclear structures and the electron configurations of the three iso-
topes of hydrogen.

2. Distinguish between ordinary water and heavy water.

3. Write balanced chemical equations for the preparation of hydrogen by the
reaction of (a) a metal and water (b) a nonmetal and water (c) a metal and
an acid (d) a metal hydride and water (e) electrolysis.

4. Write separate equations for a laboratory preparation and a commercial
preparation of hydrogen.

5. What is water gas? How can hydrogen be separated from water gas?

6. Explain and illustrate the process of hydrogenation.

7. List the chief industrial uses of hydrogen.

8. Why does hydrogen react more readily with nonmetallic elements than with
metallic elements?

9. How does the hydrogen in an alkali metal hydride differ from the hydrogen
in a binary compound with a nonmetallic element, such as HC1. What ex-
perimental evidence verifies your statement?

10. Which of the following metals would liberate H 2 from (a) cold water (b) acid:
Mn, Na, Sn, Cu, Ba, Al, Ag? Which oxides of these metals can be decomposed
by the heat of a Bunsen burner?

11. Indicate which of the following reactions will occur and complete the equation.

(a) Cu +AgNO s (b) Cu + PbCl 2 (c) H 2 -f ZnO (d) Al + Fe 2 O s (e) Ni + FeS0 4
(f) Sn 4- MnS0 4 .

12. (a) A balloon contains 50 liters of H 2 . One liter of air at STP weighs 1.293 g.
If the balloon material weighs 25 g, what additional weight can be attached
to the balloon to just prevent it from rising at STP? (b) If the balloon material
were flexible, and if the temperature rose to 50 °C while the pressure re-
mained at 760 mm, what additional weight could be attached to the balloon?

Ans: (a) 35 g

13. To reduce 3.98 g of CuO with H 2 calculate (a) the number of moles of H 2
required (b) the number of molecules of H 2 (c) the weight of H 2 in grams
(d) the volume of H 2 at 27 °C and 600 mm pressure. Ans; (a) 0.050 mole

14. (a) What weight of H 2 can be obtained from the reaction of 8.0 g of potassium
with water? (b) What volume will the H 2 occupy at STP? (c) How many
molecules of H 2 were formed?

Ans: (a) 0.21 g (b) 2.3 1 (c) 6.15 x 10 22 molecules

15. What weight of Mg metal is required to liberate 2.0 x 10 24 molecules of H 2

by (a) reaction with water (b) reaction with HC1? Ans: (a) 80 g

16. What volume of water gas, measured at 25 °C and 800 mm pressure, can be
produced from 1,0 kg of coal which contains 94% carbon? Ans: 3640 liters

17. At 500° C , Kp is 5.5 for the gaseous reaction CO -4- H 2 0 — > H 2 + C0 2 . If
a mixture of one mole of CO and two moles of H 2 0 are mixed in a 6.0 liter con-
tainer at 500°C, at equilibrium what will be (a) the mole percent of H 2

(b) the concentration of H 2 in mole/liter?

Water and
Hydrogen Peroxide

Water is the most abundant and widely distributed chemical on earth.
It occurs as solid ice and snow, most commonly as the liquid which covers
almost three-quarters of the earths surface, and as water vapor in the at-
mosphere; as much as 50,000 tons of water vapor can be present in the air
over one square mile of the earths surface. Water is present in all living
matter and in food; it comprises almost 65% of the human body. The ancients
considered it to be one of their four “elements,” earth, air, fire, and water,
from which all substances could be formed. Not until 1781 was water first
recognized as a compound by the English chemist, Sir Henry Cavendish,
who prepared it by the combustion of hydrogen in air.

1. The Structure of Water The water molecule is a covalent com-
pound resulting from the combination, or overlapping, of the two half-
filled 2 p orbitals of an oxygen atom with the Is orbitals of two hydrogen
atoms. The H— O—H angle is 105° Because of the large electronega-
tivity of the oxygen atom, water is a polar molecule with the oxygen atom
as the relatively negative end of a dipole. (Figure 16.1)

(a) Electronic representation of the H a O molecule; the electrons are closer to
the O atoms than to the H atoms.

(b) Geometric representation of the electron fields. The central line indicates
the polar nature of the H z O molecule; the O atom is negative relative to the H
atoms.

Figure 26 J. The Water Molecule.

Water and Hydrogen Peroxide

207

In the vapor state only individual H 2 0 molecules exist. In the liquid
state the water dipoles orient themselves with their positive ends, (hydro-
gen) attracted to the negative ends (oxygen) of their neighbors so that
water molecules are linked to each other by hydrogen bonds between oxygen
atoms of adjacent water molecules. The degree of association is not fixed
but varies with the temperature. It is greatest just above the freezing poirit
and least at the boiling point. The association of water molecules is the
cause of the low vapor pressure, high heat of vaporization, and the high
boiling point and melting point of water relative to analogous compounds
of Group VIB elements. H 2 S, H 2 Se, and H 2 Te (Figure 16.2).

Figure 16,2. Boiling points of the hydro-
gen compounds of the Group VIB elements.

Water has the highest boiling point of the
hydrogen compounds of the Group VIB ele-
ments. The boiling points of H 2 S, H 2 Se, and
H 2 Te follow a normal trend.

Molecular Formula :

Melting Point, °C :

Boiling Point, °C :

Critical Temperature, °C:

Critical Pressure, atm

Density at 4°C, g/ml :

Table 16-A

Properties of Water

h 2 o

0.00

100.0 (at
760 mm)
374

2X7.7

1.000

Molecular Weight ; 18.015

Vapor Pressure at 20 °C, mm : 17.4

Heat of Fusion at 0°C, cal/g : 79.7

Heat of Vaporization at 100 °C,
cal/g ; 540

Specific Heat at 15 °C,

cal/g deg *• 1*^0

Cubic Coefficient of Expansion
at 20° C, deg- 1 : 2.07 x 1(H

Miscellaneous: Pure water is odorless, colorless, and tasteless. It is a poor conductor
of electricity. By weight, the ratio of hydrogen to oxygen is 1:8; two volumes
of H 2 gas combine with one volume of 0 2 gas to form water. The apparent
coincidence that the physical properties of water are peculiarly integral is
due to the fact that water is chosen to be the standard for many physical
properties and is arbitrarily assigned a value of one (density, specific heat;
or zero and one hundred (melting point and boiling point).

In the solid state a crystal of ice can be considered to be a giant molecule
wherein the individual H— O— H molecules are linked tetrahedrally to
each other. X-ray studies of ice show that the oxygen atom of each H 2 <J

208

Water and Hydrogen Peroxide

molecule is bound to four hydrogen atoms; each hydrogen atom is held
between two oxygen atoms. The tetrahedra are joined to form the macro-
scopic ice crystal as shown in Figure 16.3. The open structure of the tetra-
hedron makes the density of ice less than that of the more closely packed
liquid water molecules at 0°C.

2. Properties of Water. (A) Physical properties : Table I6-A. At 4°C
water possesses its maximum density of one gram per milliliter. At tem-
peratures both above and below 4°C, water expands and its density de-
creases as shown in Table 16-B.

Table 16-B

Density of Water

Temperature, °C

Density , g/ml

Temperature , °C

Density, g/ml

0 ice

0.917

1.00000

0 water (I)

0.99987

0.99999

0.99993

0.99997

0.99973

0.99999

0.99823

When water changes into ice at 0°C, it undergoes considerable expansion
so that ice floats on liquid water. This expansion upon freezing is an un-
usual phenomenon; for nearly all other substances the density of the solid
is greater than that of the liquid so that they contract upon changing from
the liquid to the solid state. Water has a higher heat capacity than any
other substance except hydrogen. Thus large bodies of water, such as lakes

fca a crystal of ice each oxygen atom is attached to four hydrogen atoms which
it shares with other oxygen atoms. The four hydrogen atoms surrounding the oxygen
atom town a tetrahedron, as shown in a, in that each hydrogen atom would occupy
the c omer of a regum four-sided figure with an oxygen atom at its center. The
tetrahedra are not independent but are joined through the hydrogen atoms, one be-
tween two oxygen atoms, to form an ice crystal, a portion of which is shown in b.

Figure 2&& The Structure of Ice.

Water and Hydrogen Peroxide

209

and oceans, change in temperature more slowly than the rocks and soil that
make up the land, and tend to regulate the temperature of the air either
by the absorption of large quantities of heat as the water warms up during
daylight hours or by liberating heat during the cooler night hours.

(B) Chemical properties: Many chemical properties of water have al-
ready been inferred in discussing the properties ~of hydrogen and of oxygen
such as thermal stability, the reactions of water with metals, nonmetals, metal-
lic oxides, and nonmetallic oxides. 1 * * * * *

Furthermore, many reactions do not take place unless there is present
at least a minute quantity of water. Perfectly dry hydrogen and chlorine
do not react unless a trace of moisture is present. Substances like sulfur
and phosphorus, which bum vigorously in ordinary (moist) oxygen, com-
bine slowly, even at elevated temperatures, with oxygen that has been
carefully dried. Iron rusts in air only when moisture is present; in dry air
no action occurs. Iron also rusts in water when air (oxygen) is present,
but if the water is freed of all dissolved oxygen, iron will not rust. The
decay of organic matter, which is essentially an oxidation process that is
promoted by bacteria, takes place only when the material is wet or moist.
In all these reactions the water can be considered to be a catalyst whose
presence is necessary for reaction to occur.

The subject of ionization in solution will be taken up in detail later
(Chapter 19). For the present we might note that pure water, and also
the water in an aqueous solution, dissociates into hydrogen ions, H+, and
hydroxide ions, OH”, to a very slight extent. Between water and its ions
there is an equilibrium,

(1) H 2 0 ^=± H+ + OH*"

In one liter of pure water, the concentration of H+ and that of OH" are
each only 1 X 10' 7 mole/liter. Hence water is a poor conductor of electricity
because the important factor in determining the conductivity of such a sub-
stance is the concentration of the ions therein. In pure water, the [H + ]
and the [OH“] are necessarily equal so that water is neutral, that is, neither
acidic nor basic* Actually, the hydrogen ion, H+, which is truly a proton
only, cannot have independent existence in the presence of other water
molecules. It combines with a water molecule to form H+(H 2 0) or H s O+
through the formation of a coordinate covalent bond to the oxygen atom
of a water molecule. The H 3 0+ ion is called the hydronium ion.

3. Water as a Solvent. Water is an excellent though not a universal
solvent. The polar nature of water makes it, in general, a good solvent for
polar compounds and ionic substances in particular. Solution of an ionic
compound such as Na+Ch involves the combination of the individual ions,
Na+ and Cl", with water molecules. The negative ends of the water mole-
cules, considered as dipoles, attract the positive ions while the positive ends

1 The student should not forget that, where a chemical reaction such as A + B — »

C + D is cited as an illustration of a chemical property of A, it is also a chemical

property of B, and a means of preparation of C and or D. Also for reactions which are ,

reversible, ana not all are in practice, such a reverse reaction would also be a chemical

property of C and of D. This is of great advantage in minimizing the quantity of memoriza-

tion for examination purposes.

210

Water and Hydrogen Peroxide

attract the negative ions. The result is a clustering of water dipoles around
each ion, resulting in the solution of the salt. The dissolved ion is said to
be hydrated , or solvated , by a number of water molecules, x, the number
depending on the concentration of the solution and the temperature. Thus
the sodium ion would more truly be [^(HaO)*]* . In addition, the
dielectric constant of a medium affects the force of attraction between
charged particles in that medium according to Coulomb's Law (page
180). The larger the value of the dielectric constant, the smaller is
the attractive force. The dielectric constant of water is 80, an unusually
high value, so that the attractive force between ions in water, e.g., Na+
and Cl" is reduced to such an extent that the ions have little tendency to
attract each other. Thus they separate and move about the solution as in-

Hydration increases the size of
the ions; the distance between the
ions is also increased.

ofalo

(b)

(c)

The forces between the
ions are reduced by
the dielectric.

When the «dt dissolves in water, there is a strong ion-dipole interaction. The
negative ends of the water dipoles are attracted to the positive ions, the positive
ends of the dipoles to the negative ions. The extent to which the attraction pro-
ceeds depends upon such factors as the radius of the ion, the charge on the ion,
and the dipole moment of die solvent

Figure ISA. Solution of Solid NoCl (after J. H, Hildebrand).

Wafer and Hydrogen Peroxide

211

dependent entities. These concepts are illustrated in Figure 16.4. It should
not be inferred that all ionic substances are soluble in water. Many, such
as CaC0 3 , and indeed most naturally occurring minerals, are insoluble in
water.

Many substances dissolve in water because of specific chemical reactions
with water. Pure liquid hydrogen chloride, HC1, is a polar covalent com-
pound. It dissolves in water because of the chemical reaction.

(2) HC1 + H s O H 3 0+ + Cl-

Similarly, sulfur dioxide, S0 2 , and phosphorus trichloride, PC1 3 , dissolve
in water.

(3) S0 2 + H 2 0 H 2 S0 3

(4) PC1 3 + 3 H 2 0 H 3 P0 3 + 3 HC1

For most nonpolar substances, e.g., H 2 , gasoline, and paraffin, water is gen-
erally a poor solvent though here also exceptions exist. Alcohol and sugar
are soluble in water because of hydrogen bonds between the oxygen atoms
of these molecules and those of water. In essence, again we have a clustering
of water molecules around the dissolved molecule. Solution of a substance
in water thus involves an interaction with water molecules. The general
theory of the nature of liquid solution is complex but, in general, substances
will dissolve in each other if their molecular structures are analogous. Thus
polar substances generally dissolve in polar solvents, and nonpolar sub-
stances in nonpolar solvents.

4. Hydrates. Many substances incorporate water in their molecular
formulas. If white anhydrous (without water) copper sulfate, CuS0 4 , is
dissolved in water, a blue solution results. On slow evaporation of the
water, blue crystals deposit. When analyzed these show the following com-
position by weight: CuSO*, 63.9% and H s O, 36.1%, corresponding to the
formula CuSO* 5 H 2 0.

63.9

Formula weight of CuS0 4

159.6

36.1

36.1 ~

Formula weight of H 2 0

18.0

When the blue crystals are heated, water is given off and white anhydrous
CuS0 4 is left. This reversible reaction may be expressed by the equation

(5) CuSO* + 5 H 2 0 ^ CuSO* • 5 H 2 0

Such a definite chemical compound with water as CuS0 4 5 H 2 0 is known,
as a hydrate , and the water therein is called the water of hydration . A dot
is placed between the formula of the anhydrous substance and the water
of hydration to indicate the chemical bond between them. When the mole-
cular or formula weight of a hydrate is calculated, the weight of the water
molecules must be included. The fact that a substance can form more than
one hydrate, e.g., CuSO* * 5 H 2 0; CuS0 4 * 3 H 2 0; CuS0 4 * H 2 0 is an ex-
cellent illustration of the Law of Multiple Proportions.

212

Water and Hydrogen Peroxide

In a crystalline salt the water of hydration is generally covalently linked
to the metallic ion, though certain negative ions, e.g., [SO*] 2- , can also be
hydrated. Thus CuSO. -5 H a (> is more truly

r 2 +

H s O OHj

yPl [ SO< ( H “ 0 )] 2 "

h 2 o \>h 3

The blue solution of a copper salt owes its color to the blue [Cu(H 2 0) 4 ] 2+
ion. Other examples of hydrates are

MgCi* * 6 H a O, or [Mg(H 2 0) 8 “* 2 Cl ], and Cl 2 * 6 H 2 0.

Because they can lose water to their surroundings, hydrates have vapor
pressures. As with pure water, the vapor pressure of a hydrate varies with
the temperature but is always less than that of pure water at a given tem-
perature. If the pressure of the water vapor in the atmosphere is less than
that of a hydrate, the hydrate will give up water and form a lower hydrate
or an anhydrous salt. The loss of water from a crystal of a hydrate is usually
accompanied by the crumbling of the crystal to a powder, a phenomenon
called efflorescence . Water may also be absorbed from the atmosphere.
A substance capable of doing this is hygroscopic and can be used as a
drying agent. Anhydrous CaCl 2 absorbs water from gases because it forms
hydrates having low vapor pressures; however, hydrate formation is not
essential in order that a substance act as a drying agent. Phosphorus(V)
oxide, P 4 O 10 , and concentrated sulfuric add are very strong drying agents
or dessicants.

5. Physical States of Water. Figure 16,5 consolidates many of the facts
concerning the physical states of H 2 0. With line OA we are already familiar.
It is the vapor pressure curve for liquid water as a function of temperature
(page 44), Any point on this curve gives the pressure and the temperature
at which the two states, liquid and vapor, are in equilibrium with each
other. We should recall that “to be in equilibrium” means to be able to
exist side by side indefinitely. This is indicated in the diagram by the
symbol L V, Line OA is also the boiling point curve for water in
that it represents temperatures at which water will boil under external
pressures. Similarly, line OB is the sublimation curve for H a O. It gives the
vapor pressure of ice as a function of temperature and represents the equi-
librium between solid and vapor, indicated by the symbol V. Line OC
is the melting point {freezing point) curve. It gives the pressures and
temperatures at which liquid and solid may be in equilibrium, indicated by
the symbol S L. The fact that this curve slopes to the left indicates
that the freezing point of water decreases with increasing pressure.

Suppose we were to consider the species, H a O, at a certain temperature
and pressure and plot the point representing these conditions in Figure 16.5
as point P. Since this point lies in the area bounded by lines OC and OA*
it would indicate that, under the conditions of point P, species H s O is in

Water and Hydrogen Peroxide

213

the liquid state. Indeed any point within this area would represent liquid
water and that area has been so labelled. Similarly any point, such as O
in die area to the left of lines OB and OC, would represent H 2 0 in the
solid state. Also any point such as R to the right of lines OB and OA,
but to the left of the vertical dashed line, would represent water vapor’
Within any of these areas only one physical state is stable. The lines 2 repre-
sent conditions where two states may exist in equilibrium; on OA, liquid
and vapor; on OC, solid and liquid; and on OB, solid and vapor. The point
O is known as the triple point because under its conditions three states,
solid, liquid, and vapor, can exist at equilibrium.

If we were to draw a horizontal line from Q through P to R, such a line
would represent the heating of water at constant pressure. Such a process
has already been discussed in Chapter 3. Starting at Q where the water is

Temperature °C

Figure 16.5 . Phase Diagram for Water.

solid, heating would merely warm the solid ice until the temperature reaches
that corresponding to point M, the intersection of the horizontal line QPR
and the melting point curve, OC. Here the solid melts at constant tem-
perature. In Figure 16.5 the time variable is not a factor as it was in Figure
3.7; while melting takes place, the system remains at point M. After all the
ice has melted, the temperature rises as the water is heated but the water
remains as liquid until point T is reached. Here the water boils off as
vapor at a constant temperature, the boiling point, T. Further heating
above point T would merely serve to increase the temperature of the vapor.

2 Mathematically a line may be considered to be an indefinite number of points end
to end.

214

Water and Hydrogen Peroxide

It would be illuminating for the student to draw other horizontal and
vertical lines and to visualize the processes which they represent* Inasmuch
as liquid water cannot exist above its critical temperature, line OA ter-
minates at 374° C. If vapor is heated above this temperature no visible
transformation, or change of state, occurs so that only a broken line, AD,
has been drawn to separate the “gas” from the "vapor” state. This line
clarifies the distinction between a “gas” and a “vapor.” A vapor is the
“gaseous” substance below its critical temperature whereas a true “gas”
exists only above its critical temperature.

It should be pointed out that Figure 16.5 was not derived theoretically
but represents only experimentally determined fact. Once obtained it can
be used to predict the course of experimental changes of state. Figure
16.5 is properly entitled the phase diagram for H*0. In place of “state”
the term “phase” would be more accurate (page 38). At pressures higher
than those shown in Figure 16.5, ice can exist in several solid phases. Dif-
ferent crystalline arrangements of the H*0 molecules give rise to different
physical properties.

6. Water Vapor, The concentration of the water vapor in the air is vari-
able. Over large bodies of water the air tends to become saturated with
water vapor, and the partial pressure of the water vapor in the air ap-
proaches the maximum value (vapor pressure) at the prevailing tempera-
ture. In order to express the relative amount of moisture in the air, the term
humidity is employed. When the air is saturated with water vapor— that is,
when the partial pressure of the water vapor equals the vapor pressure
of water at that temperature— the humidity is 100 percent. If the partial
pressure is only half the vapor pressure at the given temperature, the relative
humidity , or degree of saturation, is 50 percent. The average humidity is
about 66 percent— that is, the air is normally about two-thirds saturated
with water vapor. At 20° C the vapor pressure of water is 17.4 mm. If
the partial pressure of the water vapor in the air at 20°C is 10.52 mm, the

humidity is

10,52
17.4 *

or 60.5 percent. If the temperature is raised to 23°C

without changing the water content (vapor pressure of water at 23°C is

20.9 mm), the relative humidity becomes

10.52
20.9 *

or 50.3 percent. If the

same air is cooled to 12°C, at which the vapor pressure is 10.52 mm,
the humidity becomes 100 percent, and the air is saturated. If the temp-
erature falls below that at which a given sample of air has a humidity
of' 100 percent (below its dew point) the excess water separates out as
dew, or at temperatures below 0°C as ice, in the form of snow or frost,
The relative humidity depends, therefore, upon two factors, the partial
pressure of water vapor in the air and the temperature.

7. Purification of Water. In this chapter there is no section on the chemi-
cal preparation of water. Water is so abundant in nature that chemical
preparation is unnecessary. If need be, water could be prepared by the

Water and Hydrogen Peroxide

215

burning of hydrogen or any hydrogen containing compound in, oxygen or
in air.

Water as found in nature is not pure. It may contain suspended solids
or dissolved gases and solids. The commercial problem is to prepare water
of sufficient purity from available natural water sources. Filtration and
sedimentation will remove suspended matter, but soluble impurities can
be removed only by distillation and chemical treatment. For municipal
purposes, water is purified by precipitating therein aluminum hydroxide,
Al(OH) 3 , a gelatinous precipitate. As this precipitate settles slowly, it
traps and carries down suspended matter and bacteria. The water is then
filtered through beds of sand, and chlorine is added to kill any remaining
microorganisms.

Ground water commonly contains dissolved salts of calcium, iron, and
magnesium and is known as hard water. These dissolved salts in no way
make the water unfit for drinking. They do affect the cleansing action of
soap and produce a stonelike deposit, known as boiler scale, within steam
boiler tubes. The chemical treatment io remove such impurities from water
is considered in Chapter 39. Sea water contains about 4 percent of dis-
solved solids, mainly NaCl. Though it offers an almost inexhaustible sup-
ply of water, the expense of purifying sea water, including any transpor-
tation costs to inland areas, is at present prohibitive. In some areas of the
world and in this country, water is a critical natural resource and much
research is being carried out on techniques to prepare industrially usable
water.

In the laboratory pure water is prepared by distillation. This is an ex-
pensive procedure and is .uneconomical commercially. Dissolved gases are
evolved during the early stages of distillation; the later fractions are free
of such impurity. If necessary the water may be redistilled. Because of
the solvent action of water on the, alkaline constituents of glass, glass ves-
sels are unsatisfactory for storing chemically pure water. Fused silica can
be used. The solubility of atmospheric gases, particularly carbon dioxide,
makes it difficult to keep chemically pure water for any length of time.

Hydrogen Peroxide

Hydrogen peroxide, H 2 0 2 , is considered in conjunction with water be-
cause it, too, is composed of only hydrogen* and oxygen. The properties
of hydrogen peroxide are unique, however, and bear little resemblance to
those of water.

Table 16-C

Properties of Hydrogen Peroxide

Molecular Formula

: H 2 O s

Molecular Weight

s 34.016

Melting Point, °C

: -0.89

Boiling Point, °C

: 151

Density at 20° C, g/ml

’ : 1.46

Miscellaneous: Pure hydrogen peroxide is a colorless, syrupy liquid, soluble in
all proportions iti water, alcohol, and ether. When concentrated it is
subject to violent decomposition.

216 Wafer and Hydrogen Peroxide

8. Structure of Hydrogen Peroxide. The structure of the hydrogen per-
oxide molecule is H

••

The oxygen atoms are linked directly to each other with a single covalent
bond and the 0—0 bond length is 1.49 A; the H— O— O angle is about
110° and the O— H bonds are not in the same plane, the "interplanax”
angle being about 110° also. The peroxide ion. O r , has a valence of
two and each oxygen atom therein has an oxidation number of —1, A
dioxide, such as Mn0 2 , bears a superficial resemblance to, but differs struc-
turally from, a peroxide. In a dioxide the oxygen atoms are not bonded
to each other and each oxygen atom has an oxidation number of —2, just
as in H 2 0.

9. Preparation of Hydrogen Peroxide, Hydrogen peroxide can be pre-
pared by the reaction of a salt of hydrogen peroxide, such as sodium peroxide,
Na 2 0 2 , or barium peroxide, Ba0 2 , with an acid at low temperature e.g. 0°C.

(5) Na 2 0 2 + 2 HC1 2 NaCl + H 2 G 2

(6) Ba0 2 + H 2 SO< -» RaS0 4 + H>O a

This method produces an approximately 3% aqueous solution of H 2 O a , use-
ful as an antiseptic.

Solutions of higher concentrations are prepared by the electrolysis of
50% H a SO* at high current densities. Hydrogen is produced at the cathode,
and persulfuric acid, H 2 S 2 O s at the anode. TTie H a S 2 0* reacts with water
forming H 2 0 2 .

(7) H 2 S 2 Os + 2 H 2 0 -> 2 H 2 SO* + H a O a

The H 2 S0 4 is returned to the electrolytic cell so that the net reaction is

(S) 2 H*0 H 2 0 2 + H*

Solutions containing up to 30% H 2 O a are thus produced. Higher concen-
trations, up to 85% H 2 0 2 , can be produced by distillation under reduced
pressure. The 85% solution has been used as the oxidizer in rocket propel-
lants. More dilute solutions, 6%, are used for bleaching organic matter,

10. Properties of Hydrogen Peroxide. ( A) Physical properties: Table* 18-B.
(B) Chemical properties ; Hydrogen peroxide is unstable both to heat and
to light.

(9) 2 H 2 0 2 2 H*0 + 0 2 AH = -46,200 cal

Dilute solutions keep fairly well especially if there is present a trace of
acid or a stabilizer, as acetanilide, which acts as a negative catalyst for to
decomposition process. Solid particles, such as dust, salts of heavy metals,
and alkalies tend to increase the rate of decomposition of H 2 0 2 . To pre-

Water and Hydrogen Peroxide

217

serve the solution it is stored in brown bottles to keep out the more ener-
getic rays of sunlight.

In aqueous solution, H 2 0 2 is a very weak acid.

(10) Ho0 2 2 H+ + O a 2 -

With certain compounds, H 2 0 2 reacts as an oxidizing agent.

(11) 2 HI + H 2 0 2 I 2 + 2 H 2 0 (Iodide ion, I - , is oxi-

dized to iodine, I 2 )

(12) H 2 S0 3 + H 2 0 2 — > H 2 S0 4 + H 2 0 (oxidation of sulfurous

acid to sulfuric acid)

(13) PbS + 4 H 2 0 2 — > PbS0 4 + 4 H 2 0 (oxidation of sulfide ion

to sulfate ion)

Paintings with white lead as the base which have darkened in time due
to the formation of black PbS can be restored to their original tint by wash-
ing with H 2 0 2 , which forms the white PbS0 4 . The antiseptic and bleach-
ing properties of H 2 0 2 are due to its oxidizing action.

Although usually considered an oxidizing agent, H 2 0 2 can function as a
reducing agent when reacted with a more powerful oxidizing agent than it is.

(14) Ag s O 4- H 2 0 2 -*» 2 Ag + H 2 0 + 0 2 (reduction of silver oxide)

(15) 2KMn0 4 + 8 H 2 S0 4 + 5 H 2 0 2 -> 2MnS0 4 +

2 KHSO, + 8 H 2 0 + 50 2
(reduction of permanganate ion, Mn0 4 “ in acid solution)

When H 2 0 2 . acts as a reducing agent, oxygen is formed; as an oxidizing
agent water is formed.

QUESTIONS

1. For what reasons has water been selected as a standard for many physical
properties?

2. Explain why ice does not form on the surface of a lake.

3. Why does pressure cause ice to melt?

4. 'Which will produce the severest bum: (a) water at 100°C (b) steam at 100°C
(c) hot air from an oven at 100 °C? Explain briefly.

5. Compare the structures of a water and a hydrogen peroxide molecule.

6. Illustrate the formation of a hydronium ion from a proton and a water mole-
cule.

7. Explain the dissolution of an ionic compound in water. For what type of com-
pounds is water a good solvent?

8. How are the dielectric constant of a solvent and the polarity of a solute re-
lated to solubility?

9. Define and illustrate the terms: hydrate; hydride; water of hydration; hyroxide;
anhydrous; dessicant; efflorescence; hygroscopic.

3. From Figure 16.5 select a set of conditions where water exists solely as (a) solid
(b) liquid (c) vapor {d) gas. What is the condition of H z O along line OA?
along OB?

218

Water and Hydrogen Peroxide

11. In Figure 16.5 draw a straight line which represents the transformation of
(a) solid to vapor (b) liquid to solid (c) vapor to solid (d) liquid to solid to vapor,

12. Using Figure 16.5 describe what would happen if the pressure on liquid
water, initially at 200°C and 100 atm pressure, were reduced at constant
temperature.

13. Explain why the temperature of a mixture of ice and water at 0°C does not
change upon heating or cooling until one of the phases disappears.

14. What impurities are found in natural water? How can pure water be prepared?

15. Define relative humidity. On what factors does it depend?

16. The partial pressure of water vapor in a sample of air at 25°C is 12.0 mm.

Calculate the relative humidity. Am: 50.7 %

17. If a sample of air at 24 °C must be cooled to 16°C to deposit dew, what is the

relative humidity of the original sample? Ans: 61.3%

18. What weight of water is present in 150 liters of air at 20 °C (a) saturated with
water vapor (b) if the relative humidity is 60% ?

19. Write balanced chemical equations for three modes of behavior of water.

20. Write balanced chemical equations for (a) the preparation of H 2 O s and its be-
havior as (b) an oxidizing agent and as (c) a reducing agent,

21. (a) What weight of H 2 0 2 can be prepared from 50 g of Ba0 2 ? (b) If the H a 0 a

obtained is decomposed completely what volume of 0 2 , collected over water,
will be produced at 25 °C and 778 mm? Ans: (a) 10 g

22. Assuming complete decomposition of equal weights of H 2 D and H 2 0 2 , which
would give greater yields of (a) oxygen (b) hydrogen?

23. The element barium forms both a dioxide and a peroxide. Both have the
formula Ba0 2 . Draw the structural formula of each compound.

24. What volume of ice at 0°C would be obtained from condensing and freezing

5.00 liters of steam measured at 100°C and 760 mm? Ans : 3.15 cm 8

25. How much heat is required to convert 20 g of ice at 0°C to steam at 100°C?

Ans: 14,380 cal

26. How much heat is given off when 1,000 g of steam at 100°C is condensed

and the water formed is cooled to 20°C? Ans: 820 cal

27. If 60 ml H 2 and 45 ml O a are exploded at 25 °C what will be the composition

of the final mixture? Ans: 15 ml 0,

28. Calculate the formula of aluminum sulfate hydrate, 9,54 g of which lost

4.61 g of water upon complete dehydration. Ans ; Al 2 (S0 4 ) s * 18 H,0

29. What weight of water is required to convert 50 g of CuSO* • 3 H s O to
CuSO* . 5 H 8 0?

30* What will be the weight of water evolved and the weight of anhydrous salt
produced from the heating of one mole of the hydrates {&) BaCI* * H s 0
(b) Na*B 4 0 7 • 10 H a O?

Solutions

If a few crystals of purple potassium permanganate are added to water,
the solid substance dissolves and disperses as evidenced by the spread of the
purple color throughout the liquid. Although the solid has a greater density
than the water there is no tendency for the dissolved solid to settle even
after long standing. The colored liquid appears to be perfectly homogeneous,
and it is believed that the dissolved solid is subdivided and dispersed as
particles of molecular dimensions, that is, as molecules, atoms ,or ions. Such
a homogeneous “molecular” mixture is called a solution. The subject of solu-
tions is of great theoretical and practical importance. Many biological and
other natural processes are directly concerned with solutions, while in the
laboratory and in industry most chemical reactions are carried out in solution.

1. Solution, Mixture, and Compound. A careful distinction should be
made among solutions, mixtures, and compounds. In a mechanical mixture,
no matter how finely the mixture may be ground, it is possible to detect
nonhomogeneity because the properties of the component substances are
unchanged. Thus a finely ground mixture of iron and sulfur can be com-
pletely separated on the basis of the magnetic properties of the iron or the
solubility of sulfur in carbon disulfide. A solution differs from an ordinary
mixture in that the subdivision of the “mixed” substances in the solution is down
to molecular dimensions. This is not to infer that the components of a solu-
tion cannot be separated by the proper physical means. Distillation of a
solution of salt and water can produce pure salt and pure water. In a com-
pound the simple substances from which it was formed have lost their identities
during the chemical change producing the compound and can no longer
be identified as such. Thus in carbon dioxide we find none of the character-
istics of the elements carbon and oxygen. In the »formation of a solution,
the properties of a substance are not greatly altered; the change is often
mainly a physical one involving dispersion into the molecular state. Further-
more, it is possible to vary the composition of a solution within rather wide
limits whereas the composition of a pure compound is fixed. The properties
of a solution, such as vapor pressure, boiling point, and freezing point, vary

220

Sofuthnt

with the proportions of the substances making up the solution, but com-
pounds have fixed and definite physical properties. If an insoluble sub-
stance is mixed with water and allowed to stand, the substance will settle out.
Such mixtures are not homogeneous and are called suspensions. Unite
solutions they can be separated into their components by filtration.

2. Types of Solutions. Solutions may exist in the gaseous, liquid, or solid
states. Prior to making a solution, the components thereof may be in a state
different from that of the final solution and from this viewpoint nine classes
of solutions can be distinguished (Table 17-A).

Table 17-A

Solutions

Example ■

A) Gaseous

1. Gas in a gas

Any mixture of gases, as air

Solutions

2. Liquid in a gas

Water in air

3. Solid in a gas

Naphthalene in air

B) Liquid

1. Gas in a Liquid

Carbon dioxide in water

Solutions

2. Liquid in a liquid

Alcohol in water

3. Solid in a Liquid

Sugar in water

C) Solid

1. Gas in a solid

Hydrogen in palladium

Solutions

2. Liquid in a solid

Mercury in copper

3. Solid in a solid

Gold in silver

Since solutions in the liquid state are of greatest importance to the chemist,
we shall limit our discussion in this chapter to such solutions.

3. Nomenclature. The terms solute and solvent refer to the components
of a solution. These are terms primarily for convenience of expression and
they bear no intrinsic theoretical significance. The solute refers to the
dissolved substance and the solvent to the dispersion medium. When a
solution is formed from two gases it is theoretically meaningless to speak
of one gas as being the solute and dissolving in the other as solvent since
the two gases interpenetrate and mutually dissolve in each other. This is
true of all solutions! Common usage is such that the substance which is
present in smaller proportion is called the solute and the other the solvent
When a solid is dissolved in a liquid, the solid usually is termed the solute and
the liquid the solvent because of a “common sense” feeling that the solid is
dispersed throughout the liquid. In the case of a 10% solution of alcohol in
water, the alcohol is the solute and the water is the solvent, but if we were
to add pure alcohol to this solution till the alcohol concentration exceeded
50%, are we to infer that the alcohol and the water have changed roles?
When the solution contains a relatively small amount of solute compared
to the amount of solvent, it is a dilute solution. If the relative amount of
solute is large, the solution is said to be concentrated. It is obvious that
concentrated solutions are possible only when the solute is very soluble.

4. Concentration Units. Hie concentration of a solution can be expressed
by any of the following:

Solutions

221

A. Weight fraction (weight percent; percent by weight)

B. Volume fraction (volume percent; percent by volume)

C. Molarity

D. Normality

E. Molality

F. Mole fraction (mole percent)

In the discussion that follows we- shall consider only the most common
type of solution— an aqueous binary solution. A binary solution is one that
contains only a single solute and a single solvent. An aqueous solution is
one in which water is the solvent. Whatever the makeup of the solution, how-
ever, whether nonaqueous, ternary, etc., the general principles that follow
are applicable.

A. Weight Fraction

Weight fraction (of solute) = — SQ - ut ~

Total weight of solution

_ Weight of solute

Weight of solute + Weight of solvent
Multiplying the Weight Fraction by 100 gives the weight percent or percent
by weight

Example 1; 15.0 g of Na 2 S0 4 are dissolved in 100 g of H 2 0. Calculate the
percent of Na 2 SC) 4 by weight.

Solution: Weight Fraction of Na 2 S0 4 =

15.0 g Na 2 S0 4

15.0 g Na 2 S0 4 + 100 g H z O

0.130

g Na 2 SQ 4
g solution

or 13.0% Na 2 S0 4

The percent by weight of the water could be similarly calculated or it could be
obtained by difference since, in a binary solution, the sum of the percents of
solute and of solvent must total 100.

If the quantity of either the solute or of the solvent is given in units of
volume, its weight must be calculated from a knowledge of its density.

Example 2: 15.0 g of Na 2 S0 4 are dissolved in 100 ml of H s O at 25 °C. The
density of water at 25° C is 0.997 g/ml. Calculate the percent of Na 2 S0 4 by weight.

Solution: The weight of the H a O = 100 ml x 0.997 g/ml = 99.7 g

Weight Fractionof _ 15.0 g Na 2 SQ 4 = g Na,SQ 4

Na 2 S0 4 - lg Q g Na2S04 + 99.7 g H 2 0 °’ g soln*

The results of Examples 1 and 2 are almost the same numerically only because
the density of the solvent, H a O, is close to unity. With another solvent whose
density differed appreciably from unit value the two answers would be quite
different.

If the density of the solution is known the weight of solute in a given
volume of solution can be calculated.

x The abbreviation “soln” will be used for “solution.”

222

Solutions

Example 3: What weight of Na 2 S0 4 is present in 100 ml of a 10.0% by weight
aqueous solution, the density of which is 1.08 g/ml?

Solution: The total weight of the solution (solute 4* solvent) is
100 ml X 1.08 g/ml = 108 g
10.0% of this weight is the weight of the Na 2 S0 4 .

0.100 g Na.SO,

X 108 g soln

10.8 g Na,S0 4

97.2 g H 2 0.

1.00 g soln

The weight of water present would be 108 g — 10.8 g
B, Volume Fraction

This method of expressing concentration is rarely used by the chemist
but is sometimes used industrially where both solute and solvent are liquids

or gases.

Volume Fraction —

Volume of solute
Total volume of solution

Unlike weights, the volumes of liquids and gases arc not necessarily additive
(page 243). If 50 ml of alcohol are mixed with 50 mi of water the total
volume is not 100 ml. There is a small contraction in volume and the solution
volume is 97.5 ml. In some cases solution results in an expansion of volume.

Example 4: When 50.0 ml of alcohol are mixed with 50.0 ml of ILO, the volume
of solution is 97.5 ml. Calculate the percent of alcohol by volume.

Solution:

Volume Fraction of alcohol =

50. Q ml alcohol
97.5 ml soln

ml alcohol
ml soln

Variations of the foregoing two methods of expressing concentrations appear
in the literature; thus the concentration may be given as a weight per unit volume

of solution, ^ ^ te -. In Example 3 the concentration of Na 2 $0 4 is also

. volume of solution

0.108 g/ml, or 108 g/3. Since percent is parts per hundred, on this basis the
concentration of Na 2 S0 4 is 10.8%, that is, 10.8 units of weight per 100 units of
solution volume. Strictly speaking, this is incorrect since percent is a unitless number
and is properly given only when the concentration ratio, numerator and denomina-
tor have identical units.

Unless the concentration units are clearly specified in the systems so far
considered, error may arise because of ambiguity in a loose statement of the
concentration. This is one reason why the chemist prefers to use the following
chemical units of concentration.

C. Molarity (Mohr solutions)

A one molar solution contains one mole (GMW) of solute in one liter of
solution. Any similar ratio, such as one-half mole in one-half liter of solution, is
also a one molar solution. Concentration in terms of molar units, or molarity,
is thus the number of moles of solute per liter of solution. Molarity is speci-
fied by a number signifying the magnitude of the concentration followed
by the letter M, thus 0.1M for a one-tenth molar solution. For example, 0.5
mole of solute in one liter of solution would be 0.5M, and 0.12 mole of solute
in 0.50 liter of solution would be 0.24M.

Solutions

223

To prepare a 1M solution of any solute, one mole of the solute is weighed
out and dissolved in sufficient solvent to make one liter of solution. In order
to measure this final solution volume accurately, a special container known
as a volumetric flask is employed (Figure 17.1). The flask has a long narrow
netk on which is etched a thin line. When a liquid is placed in the flask so
that the lowest portion of the meniscus is tangent to the line, the flask then
contains precisely the volume of liquid at the temperature indicated on the
flask. The weighed quantity of solute is placed in the flask, sufficient water is
added to dissolve the solute completely, and additional water is then added
to bring the total volume up to the etched line. Volumetric flasks of different
capacities are available for preparing various volumes of solutions of known
concentrations. Solutions of accurately known concentration are known as
standard solutions. In making up such a solution the temperature must be
specified inasmuch as volumes vary with temperature and hence solutions
based on molarity would have slightly different concentrations at different
temperatures. In preparing a 1 M solution it is incorrect to dissolve one mole
of solute in one liter of solvent. The final volume of the resulting solution
would not be equal to one liter.

For the system of molarity an equation relating the weight of solute, the
solution volume, and the molarity can be derived from the definition of a
molar solution.

( 1 )

Molarity =

number of moles of solute
volume of solution in liters

(2) Number of moles of solute, n = M X V

(3) Number of moles of solute, n =

GMW

224

Solutions

(4)

(5)

GMW

M X V

w - M X V X GMW

where w = weight of solute in grams
Af =r. molarity of solution
V - volume of solution in liters
GMW -• gram-molecular weight of
solute

If any three of the factors in Equation 5 are known the fourth can be
calculated.

For ionic substances such as NaCl which exist as ions even in the solid
state it is not accurate to speak of the “molecular weight."’ Further, in solution
molecules of a solute may associate or dissociate. In these cases it is preferable
to speak of the formula weight of a substance or the sum of the atomic
weights of all the atoms in a formula. One mole would contain the Avogadro
number of formula units. Hence the one molar solution as defined in the
foregoing is sometimes referred to as a one formal solution, in that it con-
tains one formula weight in grams, or the gram formula weight (GFW), of
solute per liter of solution. For the system of formality , the symbol F is used,
e.g., 0.5F, and an equation analogous to Equation 5 is applicable.

(6) w = F X V X GFW where w = weight of solute in grams

F = formality of the solution
V = volume of solution in liters
GFW = gram formula weight of solute

Example 5: What weight of Na 2 S0 4 should be dissolved in water to make
200 ml of a 0.250M solution? The molecular weight (formula weight) of Na 2 S0 4
is 142.

Solution:

w = 0.250 X 0.200 liter x 142 — = 7.10 g

liter mole

Example 6: What is the molarity of a solution if T.10 g of Na 2 SG 4 are dissolved
in 200 ml of solution?

Solution:

M = Yr — 7 — r = °- 250 mole/liter

V X GMW 0,200 liter X 142 g/mole

Example 7: What is the molarity of the solution in Example 3?

Solution:

M = (Hr)

Equation 2 can be used for dilution problems. Diluting a solution does
not change the weight of solute therein so the number of moles, or (M X V),
remains constant. Therefore.

Solutions

225

(7) Mx X Vx = M 2 X V 2

(before dilution) (after dilution)

Example 8: How much water should be added to 50 ml of a 0.25M solution
of Na 2 S0 4 to make its concentration 0.10M?

Solution: By Equation 7, 50 ml x 0.25M = V 2 X 0.10M

V 2 = 125 ml

Since the final volume is to be 125 ml, 75 ml of water should be added to the
original 50 ml. In the use of Equation 7 the volumes need only to be expressed
in the same units, not necessarily in liters.

D. Normality ( Normal solutions)

A one normal solution contains one gram equivalent weight (GEW) of
solute in one liter of solution. In terms of normality, concentration is the num-
ber of gram-equivalents per liter of solution and is indicated by the symbol
N, for example, 0.35 N. This system of concentration is applicable only to
those compounds where the concept of equivalent weight is meaningful,
primarily acids, bases, and salts.

For calculations involving normal solutions, equations analogous to those
for molar solutions can be derived.

(8) Number of gram-equivalent weights of solute = N X V

(9) w = N X V X GEW where w = weight of solute in grams.

N = normality of solution
V = volume of solution in liters
GEW = gram-equivalent weight of
solute

Example 9: What is the normality of a solution if 7.10 g of Na 2 S0 4 are
dissolved in 200 ml of solution?

Solution

GEW =

For Na 2 S0 4 ,
GMW

142 gram/mole

'V x GEW

2 gram-equivalent/mole
7.10 g

= 71.0

gram

gram-equivalent

= 0.500N

0.200 1 X 71.0 g/g-equiv \ liter /

Comparison of the answers in Examples 6 and 9 indicates that, with the
same weight of Na 2 S0 4 in a given volume of solution, the normality is just twice
the molarity. This must be so because the relation between molarity and normality
is the same as that between the GMW and GEW.

Since a gram-molecular weight always contains one or more gram-equivalent
weights, the dissolving of one GMW corresponds to the solution of one or
more GEW. For the same solution, then, the normality is always equal to
or an integral multiple of the molarity.

The student may well ask why the system of normality has been devised.
We have noted that, in a chemical reaction, the number of moles of reactants
are not necessarily equal but the number of their chemical equivalents are
always equal. Hence for two reactants in solution the values of (N X V)
must be equal.

226

Solutions

(10) Ni X Vi = N s X V 2 where the subscripts 1 and 2 refer to the

different reactants.

Example 10: What volume of 0.20\ T HOI will react with 1,25 liters of
0.602V NaOH?

Solution :

_ jVi x Vi

0.60iV x 1.25 liters

= 3.75 liters HC1

As with Equation 7 the volumes in Equation 10 need be only in the same
units; Equation 10 can also be used for problems of dilution where normality
is concerned.

The gram-molecular weight and the gram-equivalent weight are rela-
tively large quantities to use in laboratory experiments. To express the quan-
tities of solutes in small volumes of solutions, smaller units of mass are
more convenient. In one milliliter of a molar solution there is present one-
thousandth of a mole, or a millimole ; in one milliliter of a normal solution
the quantity of solute is one-thousandth of an equivalent, or a milliequiva -
lent. One millimole per liter would be 0.001 M; one milliequi valent per
liter would be 0.001 N. Millimoles per milliliter and millieqmvalents per liter
are numerically equal to molarity and normality, respectively.

A not-too-apparent advantage of solutions in general, and of standard
molar and normal solutions in particular, is this. If we wish a certain weight
of solid we have to weigh it out, a relatively tedious procedure. The solid
is thereafter generally dissolved in a suitable solvent if a chemical reaction
is to be carried out. However, if we have a standard solution of that
solid, all we need do to obtain a given weight of the solute, already dis-
solved, is-to take a volume of that solution as calculated by Equation 5 or 9.
Taking a volume is much simpler than weighing.

E. Molality (Molal solutions)

A one molal solution contains one mole (GMW) of solute in one thousand
grams of solvent. The symbol for molality is m, for example, 0.25m. Since
the molality of a solution is the number of moles of solute per one thousand
grams of solvent, then

( 11 )
m =

w 2 1000

(GMW) 2 x Wx

where w 2 = weight of solute in grams
Wx = weight of solvent in grams
(GMW ) 2 = gram-molecular weight of solute

General usage Is that the subscripts "1*" and “2” indicate the solvent and
solute, respectively.

Solutions

227

F. Mole Fraction

The mole fraction, symbol X, of a substance is the number of moles of
that substance divided by the total number of moles of all species present
in the solution. In a given solution a mole fraction can be calculated for
each substance present but the sum of these fractions must add up to one.
Since it is a ratio of like quantities, mole fraction has no units.

For a single solute in a single solvent, the mole Fraction of Solute, X 2 , is

X 2 =

moles of solute

( 12 )

total moles moles of solute + moles of solvent

n 2 where n 2 = moles of solute
n x + n 2 n x = moles of solvent

X, =

The mole fraction of the solvent, X 1 would be

n x

n x + n 2

> or (1-X 2 )

Since n

GMW

(13) w 2

X, = (GMW)z

Wi w 2

(GMW)i (GMW ) 2

where

w 2 = weight of solute in grams
Wi = weight of solvent in grams

(GMW) 2 = gram molecular weight of solute
(GMW)i = gram molecular weight of solvent

Example 12; What is the mole fraction of the solute in Example 3?

Solution:

10.8 g

142 g/mole

97.2 g 10.8 g

18.0 g/mole 142 g/mole

= 0.0139

This means that 1.39 moles of every 100 moles of solution is the solute,
Na 2 S0 4 . The mole fraction of the solvent, H 2 0, could be similarly calculated
to be 0.9861.

Unlike molar and normal solutions, mol 1 concentrations and mole frac-
tions are not on the basis of solution volume. In the next chapter we shall
see that certain properties of solutions depend upon the relative numbers
of molecules of solute and solvent. This makes necessary a unit of concen-
tration on a weight or mole basis, such as molality or mole fraction. Such
concentrations are temperature independent whereas the volumes of liquids,
and hence the number of molecules in a specified volume of a liquid, will
vary with the temperature.

228

Solutions

The information given in Example 3 has been used for many of the
illustrative examples. This infers, and correctly so, that the several units of
concentration are convertible one to the other if given the proper data. In
addition to all the foregoing units, the chemist sometimes uses parts per
thousand, and for very dilute solutions, parts per million (ppm). Since
percent is parts per hundred, multiplication of percent by 10 will give
parts per thousand; multiplication by 10,000 will give parts per million.

5. Kinetic-Molecular View of Solution. When a crystal of a soluble
substance is placed in a suitable solvent it gradually becomes smaller and
may ultimately disappear. Molecules, atoms, or ions of the solute, through
the m echanism of solution, leave the surface of the solid. Because of their
kinetic energy and collisions with the solvent molecules they are dispersed
throughout the solution. The viewpoint is essentially the same as that of
the kinetic molecular theory of gases. The intermingling of solute and solvent,
or their diffusion, is much slower in liquid solutions than in gaseous solutions
but eventually the solution becomes uniform throughout. In solid solutions
diffusion is practically nonexistent but does occur to a slight extent. If two
pure metals, A and B, are placed in contact, after a long time some atoms
of A will have crossed the interface and be found in B while some B atoms
will be found in A.

6. Saturated Solutions. In dissolving, solute molecules leave the solid
surface at a rate dependent upon the temperature. On the other hand, dis-
solved molecules may return to the solid state by condensing upon the
solid surface; the rate of this process depends upon the solution concentra-
tion. If the rate of solution is always greater than the rate of recrystallization
then the solid will dissolve completely. If a situation is attained where both
rates are equal, however, the quantity of solid then present will remain in-
definitely. There will then be a dynamic equilibrium between the solid un-
dissolved solute and the dissolved solute in solution.

(14) Solute (solid) Solute (dissolved)

A solution which is in equilibrium with excess solute (or which could be in
such equilibrium) is called a saturated solution (Figure 17.2). The con-

figure 17.2. A Saturated Solution.

Only surface molecules can leave the
solid. An equilibrium exists between
excess solid solute and dissolved solute
so that the rates of solution and con-
densation are equal.

Solutions

229

siderations applied to the equilibrium in a saturated solution are quite analo-
gous to those for the phenomenon of vapor pressure.

That the processes of solution and deposition continue after a saturated
solution is formed can be proved by a simple experiment. Suspend a broken
or irregular crystal in a saturated solution of the same material and maintain
a constant temperature. The suspended fragment will neither lose nor gain
weight. It will, however, form a more perfect crystal since solution will
occur at certain faces and edges and deposition at others.

A solution containing a lesser quantity of dissolved solute than a saturated
solution of the same solute is an unsaturated solution. The concept of satura-
tion has nothing to do with whether a solution is dilute or concentrated.
The solubility of calcium carbonate, CaC0 3 , is so slight that a saturated
solution is extremely dilute. Saturated solutions of very soluble substances,
however, are also concentrated.

7. Supersaturated Solutions. It would be incorrect to define a saturated
solution as “one which has dissolved in it the most it could dissolve.” In
general, solids are more soluble at higher temperatures. If a solution saturated
at one temperature is heated with an excess of solute of this type, the
excess solute will dissolve. If the solution is then cooled, the excess solute
usually deposits from the solution, always maintaining equilibrium between
the solid and the solution. However, in some cases the solid does not
recrystallize upon cooling. Such a solution thus has dissolved in it more than
a saturated solution, and is said to be supersaturated. In a supersaturated solu-
tion no dynamic equilibrium exists because no solid is present. For the
excess solute to precipitate some solid material 2 must be present to act as
a nucleus upon which crystallization can take place. Indeed, if a minute
crystal of the solute is added to a supersaturated solution, the system proceeds
toward a state of equilibrium and crystallization of the excess solute com-
mences immediately, the added crystal acting as a nucleus for the formation
of the solid. The solution remaining will then be saturated.

8. Solubility. Factors other than temperature also affect the rate of
solution. Like evaporation, solution takes place at a surface so that the
rate of solution can be increased by increasing the surface of the solute
exposed to the solvent. Finely divided material dissolves more rapidly than
the same weight of solid in the form of one large crystal. For this reason
chemists grind* a solute before dissolving it. Stirring of a solute with the
solvent results in a constant removal of the saturated solution produced at
the surface of the solid, thereby exposing it to fresh or unsaturated solvent.
The same effect can be accomplished by suspending the solute in the solvent
(Figure 17.3).

The solubility of a substance is that weight of it which is dissolved in a
saturated solution. For statistical purposes, the number of grams of a sub-
stance required to saturate 100 grams, or sometimes 100 ml, of a solvent at
a given temperature is usually defined as the solubility of that substance
at the cited temperature.

2 Not every solid will produce the effect described but a crystal of the solute
always will.

230

Solutions

Figure /7,3. DiaioluUon of a Suspended Solid
As the solid dissolves in the solvent, the solution thus
formed is denser than the sohvnt, and hence flows down
through the solvent, This behavior sets up countercurrent
flows: the dense solution flows downward; the less dense
solvent flow’s upward and comes in contact with the
solid, which dissolves to form a more dense solution,
This process continues until the solid is dissolved or the
solution becomes saturated. In either event the solution
eventually is of uniform density.

Changes in temperature always affect the solubility of a solute. For solid
solutes the solubility generally increases with a rise in temperature whereas
the solubility of gases decreases with temperature. A convenient way of
representing solubility data is by means of solubility curves, where solu-
bility is plotted against temperature (Figure 17.4).

Figure 17.4.

Solubility Curves.
The curves show the varia-
tion in solubility of various
solutes with change in temp-
erature. Except for gases,
the solubility usually in-
creases with increased temp-
erature. On a weight basis
the solubility of a gas is so
much less than that for a
solid solute that a different
solubility scale is used; it is
shown on the right. A break
in the solubility curve indi-
cates a chemical change at
the temperature of the
break. At points A and B,
known as transition points,
different hydrates of calcium
chloride are formed.

A saturated solution in equilibrium with excess solute behaves in accord
with Le Chatelier’s principle. Most solid solutes absorb heat during the
solution process. This involves a change from the solid to the liquid state,
an endothermic process analogous to melting.

(15) Solute (solid) Solute (dissolved) AH = + calories

A rise in temperature will shift the point of equilibrium in the direction
which absorbs heat, that is, to the right. Hence more solid will dissolve and

Solutions

231

the solubility is proportional to the temperature. In some cases the solution
of a solute is followed by an exothermic chemical reaction between solute
and solvent molecules. The heat evolved in this chemical change may exceed
the endothermic heat of solution. In such cases the temperature will rise
as the solute dissolves, even though the process of solution itself is endo-
thermic. This is the case with NaOH. However if NaOH dissolves in an
almost saturated solution, the process is completely endothermic.

When a gas dissolves in a liquid, the process involves a change from the
gaseous to the liquid state, an exothermic process akin to condensation.
Thus, the fact that the solubility of gases is inversely proportional to the
temperature is also in accord with Le Chatelier’s principle. When a glass
of cold water stands in a warm room, bubbles of air form on the inner
surface of the container. Air is more soluble in water at a lower temperature,
and as the water warms up it becomes saturated with the air. Thereafter
the air “precipitates” out.

Only in the case of gaseous solutes does pressure have any appreciable
effect on solubility. It was found in 1803 by William Henry that the escap-
ing tendency, or the partial pressure of a gaseous solute, is proportional to*
its mole fraction in solution; conversely the solubility of a gas is proportional
to the pressure at which it is supplied. Henry’s Law can be written

(16) P 2 = k X 2 where P 2 = the pressure of the gas

k = a constant

X 2 — the mole fraction of the gaseous solute

With a mixture of gases the solubility of each component is propor-
tional to its own partial pressure.

9. Law of Partition. When two different liquids, such as alcohol and
water, are mixed they may dissolve, one in the other, in any proportion.
Such pairs of liquids are said to miscible in all proportions. Other pairs
may dissolve in one another only to a limited extent, forming two separate
layers in which the less dense floats upon the more dense liquid. Such
pairs are said to be immiscible . If ether and water are mixed together and
allowed to stand, two layers form (Figure 17.5). The upper layer consists

Ether

Water

Ether

Figure 17.5* Partially Immiscible Liquids.

When ether and water are shaken together to form a mixture
and the mixture is allowed to stand for a while, the ether
and water separate into two layers. However, a small amount
of water dissolves in the ether and a small amount of ether
dissolves in the water.

232

Solutions

of ether saturated with water and the lower layer of water saturated with
ether.

When a solute is added to two immiscible liquids, and the system is
thoroughly shaken and then allowed to stand until the two liquids separate,
the solute distributes itself between the two solvents. At equilibrium the
ratio of the concentrations of the solute in each of the two solvents is equal
to the ratio of the solubilities of the solute in each solvent alone. This state-
ment is known as the Law of Partition , Iodine, I s is about 200 times more
soluble in ether than in water. If ether is added to a solution of iodine in
water and the mixture is well shaken, the concentration of iodine in the
ether layer will be 200 times the concentration of iodine in the water layer.
If equal volumes of ether and a water solution of iodine are taken the amount

of iodine remaining in the water layer will be —r of the original amount and

aU\/

the balance of the iodine will be in the ether layer. A condition of equilibrium
exists between the solute, I a , in both solvents.

( 17 )

I 2 (in water) 1% (in ether)

The ratio of the concentrations of iodine in the two solvents at equilibrium is

(18)

[hi (in ether) 200

[I 2 ] (in water) 1

The constant, K, known as the distribution coefficient , is truly an equilibrium
constant for Equation 17.

By means of successive treatments with an immiscible solvent in which
iodine is very soluble it is possible to remove practically all of the iodine
from an aqueous solution. This process is called extraction . Two extractions
with equal volumes of ether will reduce the concentration of iodine in a

water solution to about — - —
40,000

of its original concentration. The extraction

of essential oils and drugs from plants and the extraction of silver from molten
lead are hut two of the industrial applications of the Law of Partition.

QUESTIONS

1. Compare the characteristics of solutions, mixtures, and compounds.

2. What is meant by the terms "solute” and "solvent”?

3. What factors affect the rate of solution of a solid solute?

4. How can the composition of a solution be expressed? Define, each unit.

5. From the definition of normality, derive w = N x V x GEW.

6. For what purpose and how is a volumetric flask used?

7. How will a change in temperature affect the concentrations of a 1 M and a
Iro solution? Explain.

8. Define and illustrate (a) saturated solution (b) supersaturated solution (c) con-
centrated solution (d) standard solution (e) solubility (f) solubility curve.

Solutions

233

9. In what ways are the phenomena of vapor pressure and saturated solutions
alike?

10. From the viewpoint of equilibrium explain why solubility varies with tempera-
ture.

11. If the heat of solution of a solute is endothermic, what will be the effect of
an increase in temperature on its solubility? Explain.

12. If the solubility of a substance decreases with temperature, is its solution
process exothermic or endothermic? Explain.

13. Describe what occurs when a crystal of solute is introduced into a super-
saturated solution of the solute.

14. Why are nonpolar solutes likely to be insoluble in water?

15. Explain why the mixture of gases recovered from boiling water which has
been exposed to air has a different composition, from atmospheric air.

16. What weight of MgS0 4 is required to make up 300 ml of (a) a 0.25M solution

(b) a 0.25N solution? Ans: (a) 9.0 g (b) 4.5 g

17. What volume of 0.45N solution can be made from 56.6 g of NH 4 C1?

Ans : 2.35 liters

18. In order to prepare 250 ml of a 1.0QN solution, what weights of the follow-
ing are required (a) Na 2 $0 4 (b) Na 2 SO 4 *10H 2 O (c) A1 2 (S0 4 ) 3 ?

Ans: (a) 17.8 g

19. Calculate the normality and the molarity of a solution which contains 4.00 g

MgBr 2 in 250 ml of solution. Ans: 0.221 N

20. How much water should be added to 54 ml of a 0.15 N aqueous solution
of AlClj, to make the concentration (a) 0.10N (b) 0.01M?

Ans: (a) 27 ml (b) 216 ml

21. What weight of sugar, C a2 H 22 0 11 , should be dissolved in 250 g of water to

make a 0.500m solution? Am: t 42.8 g

22. A liter of copper sulfate solution contains 60.9 g of CuS0 4 . The density of
the solution at 18°C is 1.09 g/ml. Calculate the (a) molarity (b) normality

Am: (a) 0.381M (b) 0.762 N (c) 5.76% (d) 0.370m (e) 0.0068

23. Commercial sulfuric acid is an aqueous solution containing 95.7% H 2 S0 4 by
weight. The density of the solution at 20° C is 1.84 g/ml. Calculate the

(a) molarity (b) normality (c) molality (d) mole fraction.

24. What weight in grams of a sulfuric acid solution containing 58% H 2 S0 4 is

required to make one liter of 0.50N solution? Ans : 42.3 g

25. Calculate the mole fraction of a solute in a one molal solution in which benzene,
C 6 H 6 , is the solvent.

26. From Figure 17.4 determine what weight KNO a will crystallize if 200 g of

water, saturated with the salt at 70°C, are cooled to 20°C. Ans: 212 g

27. From Figure 17.4 determine what weight of water is needed to dissolve

100 g of KBr at 50°C. Ansi* 125 g

28. What volume of 0.50IV NaOH will neutralize 20 ml of (a) 0.50 N HC1

(b) 0.50N H 2 S0 4 (c) 0.50M H 2 S0 4 ? Ans : (a) 20 ml (b) 20 ml (c) 40 ml

29. What volume of 0.24N HC1 will react with 45.6 ml of 0.54 N NaOH?

Ans : 103 ml

30. What weight of NaOH can be neutralized by 50 ml of an acid which is 1.4N?

Ans: 2.8 g

234

SoJtrtionj

31. How many gram-equivalents of Ca(OH) s and of Ca*+ ion are present in

4.0 liters of O.oOiV solution? Ans: 2.0 g-equiv

32. An aqueous solution contains 0.005 g of iodine in one liter. Calculate the
concentration of iodine that would remain after shaking with (a) 100 ml of
ether (b) two successive 50 ml volumes of ether. What conclusions can you
draw concerning the relative efficiency of extraction?

33. If one millimole of nitrogen dissolves in a certain volume of water at 20°C

under a pressure of 760 mm, what weight of nitrogen will dissolve under a
pressure of 1140 mm? Ans : 0.042 g

Properties of Solutions

The physical properties of solutions differ from those of the pure solvent.
The differences in properties can be divided into two classes:

A) Those changes which depend upon both the nature and the concen-
tration of the solute, e.g., density and electrical conductivity. For example,
when sodium chloride is dissolved in water, there is a decrease in the total
volume (the total volume of the solution is less than the volume of the salt
plus the volume of the solvent), whereas when ammonium chloride is dis-
solved there is an expansion in the total volume. In solutions of sugar and
water, there is almost no change in volume.

B) Those changes which depend upon only the concentration of the
solute. In this class the changes in properties vary only with the number
of solute particles in a given quantity of solvent and are independent of their
nature, whether they be molecules, ions, or atoms. Such properties are
known as coUigative properties and include vapor pressure, boiling point,
freezing point, and osmotic pressure. It is with these properties that this
chapter is concerned.

1. Raoult’s Law. A volatile substance, such as water, alcohol, or ben-
zene, produces a vapor pressure. A solution of a nonvolatile solute, such as
sugar, in a volatile solvent also has a vapor pressure, but its vapor pressure
is always less than that of the pure solvent at the same temperature, and
the larger the concentration of the solution, the greater is the depression of
the vapor pressure. This is shown in Figure 18.1 where the vapor pressures
of pure solvent and solutions of a nonvolatile solute in that same solvent
are plotted as a function of temperature.

In 1881 the French physicist, F. M. Raoult proposed a quantitative rela-
tionship between the lowering of the vapor pressure of a solution and its
concentration. In dilute solutions of a nonvolatile solute in a volatile sol-
vent, the depression of the vapor pressure is proportional to the number of
moles of the solute in a given weight of the solvent. Thus the vapor pres-
sure' lowering depends upon the molality of a solution or its mole fraction.
Solutions containing the same numbers of moles, even of different nonvola-

Propert

Figure 18.L

Vapoj: Pressures of Pure
Solvent and Solutions.
Curve A is for pure solvent;
it is identical with those in
Figure 3.5. Both Curves B
and C are for solutions. At a
given temperature their vapor
pressures are less than that of
the pure solvent; C is for a
solution whose concentration
is greater than that in B.

tile solutes, in the same weight of a given solvent have the same vapor
pressures.

The molecular hypothesis serves to explain Raoult s law. We have seen
that the vapor pressure is dependent upon the number of molecules leaving
the surface of a liquid per unit time. In a pure liquid, the molecules at the
surface are all alike, while in a solution some of the molecules at the sur-
face, the solute molecules, are not volatile. The concentration of the solvent
molecules at the surface of the solution is less than that in the surface of the
pure solvent. Hence the rate at which they leave the surface of the solution
will be less than that for the pure solvent. For example, in a solution where
the number of solute molecules is one fifth of the total, that is, the mole
fraction of solute, X<>, is one fifth, the rate at which solvent molecules evapo-
rate will be reduced by one fifth. Both the numbers of molecules leaving
and returning to the surface of the liquid per unit time are equal. How-
ever, the number leaving the surface has been reduced by one fifth by the
presence of the solute, and the number returning is also one fifth less than
that in the case of the pure solvent. Hence, die vapor pressure of the pure
solvent has been reduced by one fifth and the vapor pressure of the solution
is thereby four fifths that of the pure solvent,

Raoults Law can be stated quantitatively that the relative lowering of the
vapor pressure is equal to the mole fraction of the solute. If p<> is the vapor
pressure of the pure solvent and p the lower vapor pressure of the solution,

then p 0 — p is the actual lowering of die vapor pressure, and — ~ JB. is the

relative lowering. This is equal to the mole fraction of the solute, X*. Hence

(1) P<l ~ P = X,

Since Xj = (1 — X*), from Equation 1 can be derived

(2) p = X, p„

that is, the vapor pressure of the solution equals the mole fraction of the
solvent times die vapor pressure of the pure solvent. Thus if the relative

Properties of Solutions

237

lowering of the vapor pressure is 20%, the vapor pressure of the solution
is 80% of the pure solvent vapor pressure.

When water is the solvent, organic substances as sugar or urea yield
solutions whose vapor pressures are in accord with Raoults Law whereas
solutions of most inorganic solutes, such as acids, bases, and salts, have
vapor pressures that are lower than those of organic solutes at the same
molal concentration. This increased depression of the vapor pressure, or
apparent deviation from Raoult’s Law, is explained by the dissociation of
inorganic solutes into ions resulting in a greater concentration of solute
particles. Raoult’s Law is strictly applicable only to dilute solutions of
nonvolatile, non-dissociating solutes.

Example 1: What is the vapor pressure at 30°C of a solution containing 50 g
of sugar, C 12 RL a 0 n , in 500 g of water? The vapor pressure of H a O at 30°C is
31.5 mm.

Solution: (a) By Equation 2, the mole fraction of the solvent, water, is

v -

A j — i

“HiiO n C ia Il 22 Oii

500 g

__ 18.0 g/mole

1 ~ 500 g ~ 50-0 g

18.0 g/mole 342 g/mole

p = X,p„ = 0.995 x 31.5 mm = 31.3 mm
(b) By Equation 1, X, = (1 - X*) = (1

=S 0.995

- 0.995)

= 0.005

Since

Solving,

p = 31.3 mm

then

31.5 mm — p
31.5 mm

0.005

2. Deliquescence. A solution of a very soluble substance may have a
vapor pressure that is less than the pressure of the water vapor in the
atmosphere surrounding it. Such a solution will not evaporate. On the
contrary, the solution will take up moisture from the air until its vapor
pressure equals the partial pressure of the water vapor in the air. Equi-
librium will then exist. All solids tend to adsorb (or condense on their
surfaces) small quantities of water from the air. This adsorbed moisture
forms a film of saturated solution. If the solid is very soluble, this film of
solution will have a low vapor pressure. The solution adsorbs more water
from the air, in an effort to reach equilibrium; more of the solid dissolves
as the solution is diluted by the adsorbed moisture. This process goes on
until the entire bulk of the material has dissolved in the water extracted
from the air. The process stops when the vapor pressure of the solution
equals the partial pressure of water vapor in the air. This behavior of a
substance, which in fact is due to the low vapor pressure of a solution, is
called deliquescence . Deliquescence is not the “opposite” of efflorescence
which concerns only hydrates.

288

Properties of Solution

3. Boiling and Freezing Points of Solutions. As a consequence of the
depression of the vapor pressure of a solution, the boiling point and the
freezing point of a solution differ from those of the pure solvent. Since
the boiling point of a liquid is the temperature at which its vapor pressure
is one atmosphere and since the addition of a nonvolatile solute lowers the
vapor pressure, it follows that such addition of a nonvolatile solute will
raise the boiling point, because the solution must’ be heated to a tempera-
ture above the boiling point of the pure solvent in order that the vapor
pressure of the solution will equal one atmosphere.

The addition of one mole of a nonvolatile solute to 1000 gm of water,
a one molal solution, raises the boiling point 0.52 °G, that is, the solution
will boil at 100.52°C. The elevation in boiling point, 0.52'C, is called
the molal boiling point elevation constant for water. The symbol for this

constant is K b and its units are

deg

mole/kg solvent

or deg kg/mole.

Each

solvent has its own characteristic value of this constant.

The boiling point elevation of a solution other than one molal is pro-
portional to its molality; a 2m aqueous solution boils at 101.04°C. Hence

(3) AT b = K b m where AT b = boiling point elevation, degrees

K b =** molal boiling point elevation constant,
deg kg/mole

m = molality of the solution, mole/kg

Substituting the definition of molality (Equation 11, page 226) into
Equation 3 gives

(4) AT*

K x 2222

b (GMW) 2 w,

Tabfe ia-A

Molal Freezing-Point and Boiling-Point Constants

Solvent

Freezing
point of
! pure solvent
°C

Freezing

1 point
constant, K f
deg kg/mote

Boiling
point of
pure solvent
°C

Boiling

point

constant, K b
deg kg/mole

Acetic Add, HC 3 H 3 O s

16.7

3.90

118.5

3.07

Aniline, C»H S NH*

-6.2

5.87

184,4

3.52

Benzene, C«H S

5.48

5.12

80.15

2.53

Carbon tetrachloride, CC1 4

-22.9

3.18

76.8

5.03

Chloroform, CHC1„

-63.5

4.67

61.3

3.86

Diethyl Ether, (C 2 H 5 ) a O

-117.0

1.79

34.6

2.02

Phenol, C„H 8 OH

7.3

181.2

3.56

Water, H s O

0.00

1.86

100.00

0.52

In Figure 18.2, line OA represents the vapor pressure of pure liquid
water as a function of temperature. This line is identical with line OA of
Figure 16.5. Line BA t gives the vapor pressure of a one molal aqueous

Properties of Solutions

Figure 18.2.

Effect of Solute on the Va-
por Pressure Freezing
Point, and Boiling Point
of Water.

The solid line curve, BOA,
applies to pure water which
freezes at 0°C (Point O) and
boils at 100° C (Point A). The
broken line curve applies to
a one molal solution; its
freezing point is given by B
and its boiling point by A r

solution of a nonvolatile, nonionizing solute. Because at any given tempera-
ture the solution vapor pressure is lower than that of pure water, the vapor
pressure of the solution at 100°C, point D, is less than one atmosphere.
The solution will not boil until it is heated to a temperature indicated by
point Ai, where its vapor pressure equals atmospheric pressure. The tem-
perature at point A x is 100.52° C, For solutions other than one molal there
would be vapor pressure curves more or less parallel to the lm curve lying
above or below it, depending upon the concentration but always below that
of the pure water curve. In Figure 18,2 line OB represents the vapor pres-
sure of ice, as it did in Figure 16.5. The freezing point of pure water is
point O, the temperature at which the solid ice and the liquid water have
the same vapor pressures and so can coexist in equilibrium with each other
indefinitely. The freezing point of the lm solution, point B, is the tem-
perature at which the vapor pressure curves for the solid ice and the solu-
tion intersect. This temperature is lower than the freeezing point of the pure
water.

Quite analogous considerations are applicable to the freezing point de-
pression of a solution as were employed for the boiling point elevation,
and similar equations can be developed. With water as solvent, a one
molal solution freezes at — 1.86° C, and the molal freezing point constant,
K f , for water is thus 1.86 deg kg/mole. As with K b , the values of K f are
different for different solvents. Table 18- A lists values of these constants
for various solvents. Thus the freezing point depression of a solution de-
ls molality.

K t m where AT* = freezing point depression, degrees

Kf = molal freezing point depression
constant, deg kg/mole
m — molality of the solution, mole/kg

w 2 v 1000
' (GMW)s w,

pends upon i

(5) AT, =

(6) AT, =

240

Properties of Solutions

Example 2: Calculate (a) the boiling point and (b) the freezing point of the
sugar solution in Example 1.

Solution: (a) by Equation 4,
deg kg

AT h = K h m = 0.52

mole

50 g

342 g/mote

1000 g/k r

500 g

The boiling point of the solution is 100.00^0 + 0.15°C = 100.15°C
(b) By Equation 6,

0.15 deg

AT, = K f m ~ 1.86

deg kg

mole

50
342 g,

g/mole

1000 g/kgl
500 g J

0.54 deg

The freezing point of the solution is 0,(XPC - C.54°C = -0.54°C

The molality of the solution could have been calculated separately, m = 0,29,
and its value employed to multiply K h and K, in order to calculate AT h and AT f .
Note that a given solution gives the same proportionate elevation in boiling point
as it does depression in freezing point.

4. Molecular Weights of Substances in Solution. There are many com-
pounds whose molecular weights cannot be determined by weighing a
known volume of the substance as a gas at certain conditions. Some com-
pounds are not readily volatilized while others are decomposed by heating.
By measuring the boiling point elevation or the freezing point depression
of solutions of such substances, however, their molecular weights can be
determined by the use of Equation 4 or Equation 6.

A) The boiling point method : To determine the molecular weight of a
substance by the boiling point method, a known weight of the substance
is dissolved in a definite weight of a solvent. The boiling point of the solu-
tion is then determined. Knowing the boiling point of the pure solvent aid
the value of K h for the solvent, the unknown molecular weight can be cal-
culated by substitution in Equation 4.

Example 3: A solution of 0.500 g of a solute of unknown molecular weight
in 40.0 g of benzene boils at 80.40°C. Calculate the molecular weight of the solute.

Boiling point of benzene = 80.15°C; K b for benzene = 2.53 deg kg/mok

Solution: AT b = 80.40°C - 80.I5°C ~ 0.25°C.

Substituting in Equation 4,

Solving for the molecular weight of the solute; (GMW) t = 126 g/mole.

bfote that the very small difference in temperature for the boiling point eleva-
tion necessitates precise measurements. To obtain a molecular weight to three
significant figures, thermometers which can be read to 0.001°C must be used.

g )The freezing point method : The freezing point of a solution contain-
ing a known weight of a solute in a known weight of solvent is determined.
The calculations for this method are similar to those for the boiling point
method except that Equation 6 is used.

Properties of Solutions

241

Example 4; The freezing point of a solution containing 0.300 g of a solute in
20.0 g of benzene is 4.20°C. Calculate the molecular weight of the solute.

Freezing point of benzene = 5.48°C; K £ for benzene = 5.12 deg kg/mole.

Solution : AT t = 5.48°C - 4.20°C = 1.28°C.

Substituting in Equation 6,

ij » ^ = 512 d< * [s *

Solving for the molecular weight of the solute; (GMW) 2 = 60.0 g/mole.

The depression of the vapor pressure itself, 'which is proportional to
the molality, could also be used to calculate the molecular weight (see
Equation 1 and Example 1). However, boiling point and freezing point
measurements, especially the latter, are much more precise so they are
used almost exclusively for molecular weight determinations of solutes.

5. Osmotic Pressure, A semipermeable membrane is one which permits
the passage of certain molecules through it but prevents the passage of
others. Such a membrane may offer no obstacle to the free circulation of
solvent molecules but will resist the passage of solute molecules. Various
substances may serve as semipermeable membranes, for example, cellophane,
parchment paper, the walls of living cells, and animal membranes. A very
efficient membrane is a film of copper (II) hexacyanoferratel(II) [cupric
ferrocyanide], Cu 2 Fe(CN) 6 , deposited in the pores of a clay cup.

When a solution and a pure solvent such as water are separated by
a semipermeable membrane there is a net flow of water through the mem-
brane from the pure solvent into the solution. This phenomenon is known as
osmosis . The water will continue to flow unless a pressure sufficient to
prevent it is exerted upon the solution or until the solution rises and builds
up a sufficient back hydrostatic pressure. The pressure required to pre-
vent the flow of solvent from the pure solvent through a semipermeable
membrane into a solution is the osmotic pressure of the solution. In Fig-
ure 18.3 a semipermeable membrane separates pure water from an aque-
ous solution of sugar. Water can pass through the membrane readily in
either direction, but sugar is restricted to the solution side of the mem-
brane. Since the concentration of water is greater in pure water than it is
in the solution, the rate at which water molecules pass from the pure water
to the solution is greater than the rate at which they pass from the sugar
solution to the pure water. There is thus a net flow of water from the
pure water to the solution.

Like the other colligative properties, the osmotic pressure of a solution
is independent of the nature of the solute molecules but is proportional to
the molality of the solution and its temperature. Solutions of different
substances having the same molal concentrations exert the same osmotic
pressure at the same temperature. Osmosis will occur not only from a
pure solvent to a solution but also from a solution of lower concentration
to one. of higher concentration. Equilibrium is established when the con-
centrations of the two solutions are equal.

242

I'mnvtws Of Solutions

f l

Figme J8.3.

The Meaauremeat of
Osmotic Pressure.

Water flows from the pare
water through the semiper-
rovahle membrane into the
sugar solution, increasing its
volume and thus causing a
rise in the liquid level in the
stem of the apparatus. The
height, h, is a measure of tie
osmotic pressure of the solu-
tion. At equilibrium, the
level in the stem remains
constant and the rate of flow
of water across the membrane
in both directions is equal.
Inasmuch as water passing
through the membrane dilutes
the sugar solution, a better
technique is to determine
what external pressure must
be applied to the surface, S,
of the sugar solution to pre-
vent osmosis from taking
place. Such a pressure is the
more correct osmotic pressure.

A solution containing one mole of a solute dissolved in 1000 grams of
water produces an osmotic pressure of 22.4 atmospheres at 0°C. The de-
pendence of osmotic pressure upon concentration and temperature follows
an equation that is strikingly similar to the Ideal Gas Law:

(7) ?rV = nRT where tt =s osmotic pressure

V = volume of solution
n = number of moles of solute
T = absolute temperature
R = molar gas constant; 0.0825 liter atm/mole
deg

From the determination of the osmotic pressure of a solution of known
concentration at a given temperature this equation can be used to deter-
mine the molecular weight of a solute. The inherently high values of osmotic
pressure make this technique particularly useful for solutes of high molecular
weight. The molecular weight of hemoglobin was determined by this
means and found to be 68,000. Perhaps it should be re-emphasized that
the discussion of solutions to this point has dealt only with dilute solu-
tions of nonvolatile and non-dissociating solutes.

6. Eutectic Mixtures, In the discussion of the freezing point depression
on page 239, it might be questioned why the same vapor pressure curve
for the solid form of water was used to determine the freezing points of
both the pure solvent and the solution. When a dilute solution freezes the
solid that first crystallizes is pure solvent. For example, if a dilute aque-
ous sugar solution freezes the initial solid frozen out is pure ice. When

Properties of Solutions

243

pure solvent freezes out of a solution, the concentration of the solute in-
creases and the freezing point of the remaining solution is decreased.
Further cooling causes more solvent to crystallize, and as the process con-
tinues the freezing point becomes progressively lower as the concentration
of the solute increases. The solution may eventually become saturated with
the solute and at this point further cooling causes both solvent and solute
to crystallize out together in the same proportion as they are in solution.
There is then no further change in concentration of the solution and
hence no change in freezing point. A solution of such concentration, in
which both solute and solvent freeze out in a definite ratio at a constant
temperature, is called a eutectic mixture; its composition is known as the
eutectic composition and its freezing point is the eutectic temperature. For
a solution of sodium chloride the eutectic temperature is — 21 °C and the
eutectic composition is 23.3% NaCl; for calcium chloride solutions the eutec-
tic is at — 55 °C and a composition of 29.8% CaCl 2 .

Practical use is made of the fact that a solute lowers the freezing point
of a solvent. Snow (m.p. = 0°C) will melt if NaCl is added to it; a 23.3%
NaCl— H 2 0 mixture will remain liquid even if the temperature falls to — 21 °C.
To prevent the water in an automobile radiator from freezing, various sub-
stances are added in concentrations depending upon the temperature an-
ticipated. Such solutions are known as antifreeze solutions; common anti-
freeze solutes are methyl alcohol, CH 3 OH; ethylene glycol, C 2 H 4 (OH) 2 ; and
glycerol, C 3 H 5 (OH) 3 .

7. Ideal Solutions. When two miscible volatile components form a liquid
solution, the vapor pressure of the product is usually different from the sum
of the vapor pressures of the two components of the mixture. Each par-
ticular mixture of the two^ components has its own specific boiling point and
gives off a vapor which is in equilibrium with the mixture. The vapor
pressure might be accurately calculated if each component exerted a partial
vapor pressure which was proportional to the mole fraction of that com-
ponent ( Raoult s Law), but very few solutions of volatile components act
strictly in accord with Raoult’s Law. A solution which obeys Raoulfs Law
is known as an ideal solution .

Deviations from Raoult s Law are due to factors that involve attractive
and repulsive forces between molecules in the solution. When benzene,
C 6 H 6 , and toluene, C 7 H 8 , are mixed to form a solution, Raoult s Law holds.
The reason for this ideal behavior lies in the similarity of molecular struc-
ture and fields of force between the two components; the attraction be-
tween a molecule of benzene and between a molecule of toluene is about the
same as the attraction between two molecules of the same substance. This
concept leads to the empirical definition that an ideal solution of two liquids
is one in which the heat of solution is zero and also one in which the
solution volume is additive, that is, the solution volume equals the sum of
the volumes of the separate components.

. 8. Fractional Distillation. For solutions of volatile liquid components
whose individual vapor pressures are not the same, the composition of the
vapor is different from the composition of the liquid with which it is in

244

Properties of Solutions

equilibrium. The processes occurring when such a solution is boiled may
be visualized from the curves in Figure 18.1, in which the boiling points
are plotted against the composition for a hypothetical mixture of volatile
components X (b.p. = 60°C ) and Y (b.p. — 120 C ). We will assume that
these substances are miscible in all proportions. The lower curve gives the
boiling points of all possible liquid mixtures of X and Y. The upper cum
(broken line) shows the composition of the vapor in equilibrium with the
boiling liquid. Thus a horizontal line drawn at any temperature gives the
compositions of liquid and vapor which are in equilibrium at that tempera-
ture. Such a line indicates that the vapor coming off is richer in the more
volatile component than is the liquid from which it originates. Suppose
that we distill a mixture containing one mole of X and one mole of Y,
The mixture will boil at 80 5 C (point A on the lower curve) and the distil-
late (condensed vapor which was in equilibrium with the boiling liquid)
coming off at that temperature will have a composition given by point B on the
upper curve, or 0.2 mole of Y to 08 mole of X. The composition at point
B is richer in the more volatile component, 0.8 X, than is the liquid from
which it was distilled, 0.5 X. If the distillate is now subjected to a second
distillation, it will boil at 65°C, point B', and yield a distillate correspond-
ing to point C on the upper curve, or 0.0 1 mole of Y to 0.96 mole of X.
Thus two successive distillations, for this particular system, serve to en-
rich a 50 mole percent X solution to a 96 mole percent X solution. This
enrichment of X would occur in the distillate while the liquid remaining
would become correspondingly more concentrated in Y. By repeated dis-
tillation of the separata fractions, practically complete separation of the
original mixture can be accomplished. This process is termed fractional
distillation .

By employing fractionating columns, devices in which the vapors are
condensed and redistilled automatically, separation of a mixture of two
or more volatile components may be realized in one continuous distillation
(Figure 18.5). Fractionating columns are employed in the separation of
crude oil into its components such as gasoline and kerosene, and in the
distillation of commercial alcohol and liquors.

9, Constant Boiling Mixtures, Theoretically, the components of a solu-
tion of X and Y of Figures 18.4 could have been separated completely into
pure X and pure Y. Certain liquid mixtures have distillation curves which
exhibit a maximum boiling point, as drawn in Figure 18,6. Distillation of
such a mixture will yield only one pure component and a mixture that
boils at a constant temperature. At this constant temperature the distillate
coming off is a mixture which has the same composition as the liquid being
boiled. Continued boiling of such a mixture does not change the composi-
tion of the liquid Or the boiling point, both of which thereafter remain
constant until the last drop of liquid evaporates. Such a mixture is called a
constant boiltog mixture or azeotropic mixture . Thus solutions of substances
forming constant boiling mixtures cannot be completely separated by frac-
tional distillation. They can be separated into one pure component and a
mixture of that composition which corresponds to the maximum or mini-
mum boiling point.

Properties of Solutions

245

120
110

100

° 90

2 80

1 70

| 60
| 50

40
30
20
10
0

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0

Composition, Mole Fraction Y

Figure 18.4. Distillation Diagram.

Figure 18.5.

Schematic Diagram of a Bubble-cap
Fractionating' Column.

The liquid to be distilled, of composi-
tion A in Figure 18.4, is introduced into
the bottom of the apparatus, which may
be heated bv a flame or an electric
heater. Distilled vapor of composition
B rises to the first plate. As it bubbles
its way through "bubble-caps' and a
layer of previously condensed liquid,
this vapor also condenses. The liber-
ated heat of condensation serves to
vaporize some of the previously con-
densed liquid. The new vapor has the
composition C. A third distillation from
the second plate would produce further
enrichment, and so on up through the
column. The tubes, T, return the over-
flow condensate to a lower plate. The
vapor leaving the top plate is most en-
riched and goes to the condenser.

Temperature

246

Properties of Solutions

A. Maximum Boiling Point Distillation
Diagram.

An example is an HG1 — H .O mixture,
for which the azeotrope is at 110*0 ami
20% HCI. Mixtures richer than this in
HCl can be separated into pure HC! and
a 20% HCl mixture, the azeotrope. Mix-
tures less concentrated in HCI than the
azeotrope can be separated into pure
water and the same 20% HCl azeotropic
mixture.

B. Minimum Boiling Point Distillation
Diagram.

An example is an alcohol-benzene mix-
ture, for which the azeotrope is at 68°C
and 45% alcohol. Mixtures containing
more than 45% alcohol can be separated
into pure alcohol and the azeotropic mix-
ture whereas mixtures containing less than
45% alcohol can be separated into pure
benzene and the azeotropic mixture.

Figure J8.6. Distillation Curves Exhibiting Maximum
and Minimum Boiling Points.

QUESTIONS

L What properties of solutions are dependent solely upon the number of particles?

2. State Raoult’s Law. Write two alternate equations for Raoult’s Law,

3. Starting with p = X 1 p t> , derive — ~ =*

Po*

4. Draw a graph of vapor pressure against mole fraction of solvent for a solution
which obeys Kaoult’s Law.

5. Two beakers* one containing pure water and the other an aqueous solution of
sugar, are placed under an inverted belf jar. Describe and explain what will
occur,

6. Explain deliquescence. Why will only very soluble substances deliquesce under
ordinary conditions?

7. Define osmosis; semipermeable membrane.

8. If a container made of a semipermeable membrane is filled with a 0.5 m solu-
tion of sugar in water and suspended in a 0.1m sugar solution, what action
will occur?

Properties of Solutions

247

9. Explain why a solution of a nonvolatile solute has a boiling point and a freez-
ing point different from that of the pure solvent.

10. Define (a) molal freezing point depression constant (b) molal boiling point
elevation constant.

11. (a) For the same weight of solute, which would be more effective in lowering
the freezing point of water: sugar, C 12 H 22 O n ; glycerol, C 3 H 5 (OH) 3 ; ethylene
glycol, C 2 H 4 (OH) 2 ? (b) Which would have the highest osmotic pressure?

12. What experimental data must be obtained to determine the molecular weight
of a solute from (a) vapor pressure measurements (b) freezing point measure-
ments?

13. The freezing point of whole milk is — 0.56 & C. How could one determine
whether a sample of milk had been watered to increase its volume?

14. Define and illustrate a eutectic mixture. What is the reason for the existence
of a eutectic? What is the lowest temperature obtainable by the addition of
NaCl to ice?

15. Define ideal solution. How do ordinary solutions differ from ideal solutions?

16. Describe and explain graphically fractional distillation. From Figure 18.4 de-
termine (a) die boiling point of a liquid mixture containing 4 moles of Y and
1 mole of X (b) the composition of the vapor in equilibrium with this mixture
(c) the boiling point of the condensed vapor.

17. Describe the operation of a fractionating column. Using Figure 18.4, how
many distillations are required to produce a liquid of 0.9 mole fraction X from
a liquid of 0.1 mole fraction X? If the vapor and liquid curves were “closer”
together, how would this affect the number of distillations required?

18. Define and illustrate a constant boiling mixture.

19. Why is it impossible to separate, by distillation only, a mixture of two sub-
stances that forms a constant boiling mixture? Into what could the components
of a mixture which exhibits a minimum boiling point be separated?

20. The vapor pressure of water at 30 °C is 31.5 mm. Calculate the vapor pres-
sure of each of the following solutions at the same temperature.

(a) 1.0 mole sugar, 4- 50 moles water Arts: (a) 30.9 mm

(b) 100 g sugar 4- 60 moles water (b) 31.4 mm

21. For a solution of 30.7 g of glycerol, C 3 H 5 (OH) 3 , in 400 g of water, calculate

(a) the boiling point (b) the freezing point.

Ans: (a) 100.43°C (b) -1.54°C

22. What weight of naphthalene, C 10 H 8 , must be dissolved in 500 g of benzene,

C 6 H G to produce a solution that freezes at 0.0°C? Ans : 67.5 g

23. What weight of ethylene glycol, C 2 H 4 (OH) 2 , should be dissolved in 20 liters
of water to produce a solution which freezes at — 9.3°C?

24. Calculate the freezing point of each of' the following solutions:

(a) 1.0 mole solute + 40 moles water Ans: (a) — 2.58°C

(b) 0.82 g sugar, C 12 H 22 O n , 4- 12.0 g water (b) -0.37 c C

25. Calculate the boiling point of each of the following solutions:

(a) 85.5 g sugar C 12 H 22 O n 4- 500 g water Ans: (a) 100.26°C

(b) 0.2 mole solute + 6.0 moles carbon tetrachloride, CC1 4

26. Calculate the osmotic pressure of the solution in Example 25 (a) at 20 °C. The

density of the solution is 1.07 g/ml. Ans; 10.9 atm

248

Properties of Solutions

27. The freezing point of pure aniline, (J tt H,NH a , is -6.20°C. A solution con-
taining 1.28 g of the solute naphthalene, C,„H 6 , in 100 g of aniline freezes at
-6.79°C. What is the molal freezing point constant for aniline?

28. What is the molecular weight of a nonvolatile compound, if 10.0 g of it dis-
solved in 100 g of water produces a solution that boils at 100.26°C?

Arts: 200

29. Calculate the molecular weight of the solutes in the following solutions;

(a) 98 g of solute in 45 moles of water; freezing point = -1.52°C

(b) 12.5 g of solute in 100 g of benzene; freezing point = 2.92*C

(d) 50 g of solute in 500 g of water; vapor pressure at 30 °C = 31.3 mm

Ans: (a) 148 (b) 250

Solutions of Electrolytes

The previous chapter dealt with solutes that did not dissociate. Certain
solutes, however, such as acids, bases, and salts, dissociate into electrically
charged ions. How this dissociation affects the properties of solutions is
the subject of this chapter.

1. Electrical Conductance. Pure water does not conduct an electric
current but certain substances, when dissolved in water or even when fused
in the absence of water, conduct electricity. Solutions of other substances,
such as ordinary sugar, do not allow the passage of an electric current. A
simple apparatus for testing a solution for electrical conductance is shown
in Figure 19.1. It consists of a source of electric potential (battery), a means
of detecting the flow of an electric current (lamp or meter), and two elec-
trodes immersed in the solution to be tested, all connected in series.

(a) Pictorial Diagram (b) Schematic Diagram

B = Battery; L = Lamp (a meter could be used in place of the lamp f or_greater
sensitivity): S = the solution whose conductance is to be tested; E and E — Jiiec-
trodes (generally graphite or a metal which will itself not react with the solution;
the electrode. E, connected to the positive terminal of the battery is the positive
electrode (4-) and that comiected to the negative terminal, E', i* the negative
electrode (-).

Figure 19.1, Apparatus for Testing Electric Conductance of Solution*.

250

Solutions of Electrolytes

If the electrodes are immersed in a conducting solution, the lamp will
glow. The brightness of the lamp, or the magnitude of the meter reading,
is an indication of the conducting ability, or conductance, of the solution.
If the solution is a nonconductor, or only slightly conducting, the lamp will
not glow.

A substance which forms a conducting solution when dissolved in water
is known as an electrolyte and the passage of an electric current through
such a solution is electrolysis. A good conductor of electricity is a strong
electrolyte whereas a poor conductor is a weak electrolyte. During electrolysis
chemical reactions take place at the electrodes. When a solution of hydrogen
chloride is electrolyzed, hydrogen gas is produced at the negative electrode
and chlorine gas at the positive electrode. The quantitative relationships,
known as Faraday’s Laws, between the amount of electricity and the resultant
electrolytic products are taken up in the chapter on Electrochemistry (Chap-
ter 25).

2. Conductance and Concentration. The resistance of any conductor, R,
is directly proportional to its length, L, and inversely proportional to its
cross-sectional area, A.

> . , ^ L where r the specific resistance of the conductor; it

(1) R = r —

A is the resistance' of a cube 1 cm on edge

The reciprocal of the resistance, — , is the conductance of a substance

and the reciprocal of the specific resistance, _ , is its specific conduct-

ance, k . The unit of resistance is tire olun and the unit of conductance is
the reciprocal ohm , ohnr 1 or mho* which indicates that at least some scientists
have a sense of humor. Values of specific resistance for some materials are
given in Table 19-A.

Table '19-A

Electrical Resistance

Substance

! Temperature , °C

Specific Resistance, ohm cm

Copper

Um x

Mercury

95.77 x 10 -«

Potassium chloride, I M

8.93

0 . 001 M

681.0

Acetic acid, 1 M

758

0.001 M

24400

Water

2.5 x 10 7

To compare the conductances of different solutions the molar conductance
and the equivalent conductance are used. The molar conductance of a solu-
tion is defined as the conductance of a solution containing one mole of solute.
Because for different electrolytes one molecule forms different numbers of
ions, the equivalent conductance is a better measure of the relative con-
ductances of solutions. The equivalent conductance is defined as the conduct-
ance of a solution containing one gram-equivalent of a compound. This defini-

Solutions of Electrolytes

251

tion specifies no concentration since the gram-equivalent of solute may be
dissolved in any volume of solvent. The equivalent conductance thus may be
determined over a range of concentrations.

The equivalent conductance, a, is related to the specific conductance,
k, by

/£\ A _ y ^ 1000 * where V e = volume of solution in ml

^ * e 2 s/ which contains one gram-

equivalent of solute

N = normality of the solution

The equivalent conductance is found to be not independent of the solu-
tion concentration. For strong electrolytes, such as potassium chloride, it
increases as the solution is diluted and finally reaches a maximum value,
designated a,,. For weak electrolytes, such as acetic acid, the equivalent
conductance increases progressively with dilution; it does not reach a maxi-
mum within the range of . experimental determination. Data for the conduct-
ances of potassium chloride and acetic acid are given in Table 19-B, and the
variation of equivalent conductance with concentration is shown in Figure 19.2.

Table 19-B

Variation of Equivalent Conductance

Potassium Chloride

Acetic Acid

Concentration ,
g-equiv/ liter, N

Volume per

Specific

Equivalent

Specific

Equivalent

g-equiv, ml

Conductance

1.0

0.1

Bi§!M

0.1119

0.01289

111.9

128.9

0.00052

5.2

0.001413

141.3

0.000163

16.3

10«

0.0001469

146.9

0.0000492

49.2

0.0001

0.0000

-10 7

0.00001489

148.9

149.9°

390,6*

* =, calculated values.

Dilution, liter/gram-equivalent
(decreasing concentration -»)

With increasing dilution the equivalent conductance increases, For a strong
electrolyte it reaches a maximum value, a 0 * (Dilution is the reciprocal or con-
centration.)

Figure 19.2. Relationship between Equivalent Conductance and
Concentration.

252

Solutions of Electrolytes

3. Deviations from Raoult’s Law. An aqueous solution of a nonvolatile
nonelectrolyte has a vapor pressure depression, a boiling point elevation, a
freezing point depression, and an osmotic pressure in agreement with- Raoult’s
Law. An electrolyte dissolved in water yields a solution for which the fore-
going properties are greater than those prescribed by Raoult’s Law. For
example, a lm solution of sugar would freeze at -1.86°C in agreement with
Raoult’s Law but a 1m aqueous solution of sodium chloride, an electrolyte,
freezes at a temperature below -1.86°C. Such solutions of electrolytes are
said to deviate from Raoult’s Law. Moreover, the percentage deviation in-
creases as the solutions are made more dilute. A measure of the deviation
from Raoult’s Law can be expressed by a value, i, which is the ratio of the
experimentally observed change in a colligative property to that predicted
by Raoult’s Law, thus:

i — ATf (experimentally determined)

AT, (calculated from Raoult’s Law)

Similar ratios of boiling point elevations, AT*,, vapor pressure depressions,
Ap, and osmotic pressures, v, are also valid for i. The values of I at various
concentrations for three electrolytes are given in Table 19-C and these data
are plotted in Figure 19.3.

Table 19-C

The Value of i at Various Concentrations

Sdute

i * Concentration in motes per 1000 g water

0.5

0.05

0.01

0.005

KCl

1.80

1.89

1.94

1.96

k,so 4

2.32

2.57

2.80

2.86

K s Fe(CN),

2.45

3.02

3.60

3.68

4. Chemical Reactions of Electrolytes. Chemical reactions between electro-
lytes in water solution differ markedly both in rate and in type from re-
actions between non electrolytes, and also from reactions of electrolytes
in the dry state. Reactions between electrolytes are generally very rapid,
sometimes almost instantaneous, whereas, those between nonelectrolytes are
generally much slower. If a solution of copper nitrate is added to a solution
of ammonium carbonate, both electrolytes, there is an instantaneous precipita-
tion of copper carbonate.

Cu(NO»)* + (NH 4 ) s C 0 3 -* CuCOj(s) + 2 NH 4 NO s

When dry copper nitrate and dry ammonium nitrate are mixed no apparent
reaction takes place. If the mixture is heated the products are entirely dif-
ferent from those produced in solution. Reactions between organic com-
pounds, which are generally nonelectrolytes, may take many hours to go
to completion.

Electrolytes appear to be composed of two, rarely more, kinds of atoms
or radicals, each of which exhibits its own characteristic properties in solu-

Solutions of Electrolytes

253

Figure 1S.3. Variation of i with Concentration.

tion. For example, all copper compounds that act as electrolytes in solution,
as CuS0 4 , CuCl 2 , etc., undergo the same reaction as does Cu(N0 3 ) 2 . Solu-
tions of these copper compounds have in common many chemical properties
which are characteristic of the copper radical and are not affected by the
nature of the other radical with which the copper is combined. Similarly
solutions of sulfates, as CuS0 4 , Na 2 S0 4 , H 2 S0 4 , etc., all have common proper-
ties characteristic of the sulfate radical. These properties are quite indepen-
dent of the nature of the other radical present in the compound. Thus the
CuS0 4 molecule is visualized as having properties which are due to the
independent behavior of the copper radical and of the sulfate radical.

5. The Arrhenius Theory. In order to explain these facts concerning solu-
tions of electrolytes, the Swedish chemist Svante Arrhenius in his doctoral
thesis in 1887 advanced the following postulates, which collectively are known
as the Arrhenius Theory of Ionization.

A) Electrolytes when dissolved in water or some other solvents dis-
sociate partially into smaller units, known as ions.

B) The ions are electrically charged. The charge on any ion is equal to
its oxidation number.

C) An equilibrium exists between the ions and undissociated molecules.

Thus, Molecules ^ Ions e.g., H 2 S ^ H + + (HS)~

D) Ions act independently of one another and the undissociated molecules;
they have their own specific properties.

In the light of subsequent developments the Arrhenius theory is known
to be inadequate. The concepts are valid for solutions of weak electrolytes

254

Snluthm* tff Electrolyte

but are incorrect for .solutions oi strong rlectrolytt's. At the time that
Arrhenius presented his theory the electron was unknown and the idea that
charged particles could exist in an electrically neutral solution was ridiculed
by most chemists. Nevertheless the theory explained much of the experimental
data concerning solutions and in 1903 Arrhenius received the Nobel prize in
chemistry.

Conduction takes place in a solution of an electrolyte because electrically
charged particles, the ions, are present. 1‘nder the influence of an electric
field, the potential applied to the electrodes, the ions are attracted to the
oppositely charged electrodes. This migration of the ions to the electrodes
constitutes the electric current in tin* solution. At the electrodes the ions
undergo an electrode reaction, through the gain or loss of electrons, whereby
their electrical charges are neutralized, and new chemical products are
formed. In the process of ionization the number of negative charges and
positive charges formed are equal Thus despite the fact that electrical charges
are present, the solution is electrically neutral

The conductance of any solution, that is, the rate at which electrical
charge flows through it, depends upon three factors:

1. the concentration of the ions in the solution

2. the charge on the ions

3. the mobility or speed of migration of the ions

A strong electrolyte is one which is highly ionized, that is, it produces
a large ionic concentration, whereas a weak electrolyte yields few ions. When
a solution is diluted, Arrhenius assumed that the mobility of the ions, and
the ionic charges, remained unchanged. He ascribed the increase in equivalent
conductance upon dilution to an increase in the number of ions. In the
equilibrium that exists between the undissociated molecules and their ions,
dilution shifts the point of equilibrium in the direction of ion formation,
thereby increasing the ionic concentration. The maximum value of conduct-
ance, attained at high dilution corresponds to complete, or 100%, dis-
sociation of the molecules into ions. Further dilution beyond this point
cannot increase the number of ions and so the conductance remains constant.
Conversely concentration of an electrolytic solution would decrease the de-
gree of ionization and its conductance would be lower. ’

Ionization results in a greater total number of particles of all kinds, if we in-
clude both the ions and the undissociated molecules still present. Since the
magnitudes of the vapor pressure depression, the boiling point elevation,
the freezing point depression, and the osmotic pressure depend only upon
the number of particles present, whatever their nature, the values of these
properties will be greater than predicted by Raoult’s Law, which assumes
no dissociation of the solute. Furthermore as the solution is diluted, the
degree of ionization increases and the deviation from Raoult’s Law increases*
For a compound such as potassium chloride which dissociates into two ions,
the value of i should approach 2 as a limiting value. Theoretically a Itn
solution of KC1, if completely ionized, would yield a 2m solution of ions.
For K 2 S0 4 and K s Fe(CN) 6 , i should approach values of 3 and 4, respectively.

K 2 SO*~»2K++ ($0 4 ) 2 -
K*Fe(CN) e 3 K+ + [Fe<CN) e J*~

Solutions of Electrolytes

255

The properties of an electrolyte are truly the properties of the ionic
species it produces. Copper sulfate exhibits properties of the ions it forms,
namely, the Cu 2 + and (SO*) 2 ” ions. Inasmuch as the ions have their own
specific properties and since they react independently of one another and
of the undissociated molecules, all copper compounds, which in solution
yield Cu 2 + ions, will show similar properties on that account. Also all
compounds producing (SO*) 2 " ions will give evidence of those properties
specific to sulfates. This viewpoint is fundamental in qualitative analysis .

6. Degree of Ionization. Though any colligative property can be used,
the degree of ionization of a solute can be computed most readily from
freezing point or conductance measurements.

A) The freezing point method: For a 1.00m aqueous solution of a non-
ionizing solute, the freezing point depression is 1.86° C. The freezing point
of a solution of an electrolyte is proportional to the total concentration
(molality) of all species of particles present, including ions and undissociated
molecules. This molality will be in proportion to the actual freezing point
depression of the solution as 1.00 is to 1,86.

As an example, let us calculate the degree of ionization of a 0.1m solution
of acetic acid. Acetic acid, H(C 2 H 3 0 2 ), is a weak electrolyte and ionizes
into hydrogen ions (truly hydronium ions) and acetate ions.

H(C 2 H s 0 2 ) ^ H+ + (C 2 H 3 0 2 )-

An equilibrium exists between the undissociated H(C 2 H 3 0 2 ) molecules and
the H+ and (C 2 H 3 0 2 ) - ions, all three being present at equilibrium. When
equilibrium is established, the acetic acid molecules are ionizing into hydro-
gen ions and acetate ions at a rate equal to that at which the ions are
recombining into acetic acid molecules.

The ordinary principles of chemical equilibrium are applicable. To
calculate the total number of moles of all species present at equilibrium,
let a denote the unknown degree of ionization, that is, the fraction of the
initial number of acetic acid moles which ionizes; a is a pure number whose
value ranges from 0 to 1. In the 0.1m solution of acetic acid, the number of
moles of acetic acid which ionizes is thus 0.1 a . Since ionization is a chemical
reaction, its equation states that, for each mole of acetic acid which ionizes,
there will be produced one mole of hydrogen ion and one mole of acetate
ion. If 0.1 a mole of H(C 2 H 3 0 2 ) ionizes there will be produced 0.1 a mole
of H + ion and 0.1 a mole of (C 2 H 3 0 2 )“ ion. The number of moles of un-
ionized, molecular H(C 2 H s 0 2 ) remaining will be the original number less
that which ionizes, or (0.1 - 0.1 a). These statements are summarized below.

H(C 2 H 3 0 2 )

H+

(C 2 H 8 0 2 )-

Initial concentration 0.1

Let a equal the degree of ionization

Number of moles which re- 0.1 a

0.1 a

act or are produced

Concentration at equilibrium (0.1 - 0.1 a)

0.1 a

Total number of moles [ (0.1 - 0.1 a) + 0.1

a + 0.1 a]

= ( 0.1 + 0-1 a )

256

Solutions of Electrolytes

By experiment, the freezing point of the 0 ,1m acetic solution is found
to be -0.19°C. Substituting in AT* = k*m (Chapter 18, Equation 5)

0.19 deg = 1.86 X (0.1 + 0.1 a)

mole kg

and solving, ct = 0.022, that is, in a 0.1 molal solution acetic acid is 2.2% ionized,
B) Conductance measurements : In the Arrhenius theory, the maximum
or limiting value of equivalent conductance, a 0 , corresponds to 100% ion-
ization of the solute. The ratio of the actual equivalent conductance, a, of
a given solution to its limiting value is the degree of ionization.

If a solution at a certain concentration has a value of equivalent conductance
half its maximum value, the solute is 50% ionized at that concentration.

Example h At 25°C the equivalent conductance of 0.001m acetic acid is
49.5 oh nr 1 cm 2 equiv- 4 . The maximum value of equivalent conductance, a 0 > is
390 ohnr 1 cm 2 equiv 4 . Calculate the degree of ionization of 0.001 m acetic acid,

Solution:

— A — 49.5 ohm* 1 cm s equiv*
a ~~ A 0 ~~ 390 ohm 4 cm 2 equiv 4 ~~

For strong electrolytes, values of *\<» can be obtained by graphical means
as indicated in Figure 19.2. Better, a graph of equivalent conductance
against the square root of concentration gives a straight line which can be
extrapolated to zero concentration to give a value of a«. For weak electro-
lytes the rapid change in equivalent conductance with dilution does not
admit such extrapolation. The conductance of an electrolyte is, however,
the sum of the conductances of its ions. At extreme dilution, known to the
chemist as infinite dilution , the equivalent conductance, a 0 , is

a 0 = L + In where L = the ionic conductance of the
+ - + positive ion

= the ionic conductance of the
negative ion

This statement is known as Kohlrausch’s law of the independent migration
of ions. Hence a 0 for acetic add would be the sum of the ionic conductances
of the hydrogen ion and the acetate ion.

Ao, H(C*H*0 2 ) = lot H+ 4- f<b (C 2 H$0 2 )“"

For each of the strong electrolytes, NaCI, HC1, and Na(C 2 H 3 Ot), values
of A 0 can be obtained graphically; these are, respectively, 126, 425, and 91,
at 25°C. Arithmetic then enables the calculation of a*, for acetic acid.
a <hh(c*h*o 2 ) ~ Ao,hci + Ao, Ntci ^ 425 + 91-126=390

Values of ionic conductivities, 1©, for some common ions are given to
Table 19-D.

Solutions of Electrolytes

257

Table 19-D

Equivalent Ionic Conductances at Infinite Dilution, 1 0 , at 25 °C,

ohm- 1 cm 2 equiv* 1

H+

349.8

OH-

198.0

Li+

38.7

Cl-

76.3

Na+

50.1

Br~

78.4

K+

73.5

76.8

% Ca 2 +

59.5

NO a -

71.4

% Mg 2 +

53.1

% SO, 2 -

79.8

7. Strong Electrolytes, Subsequent discoveries, primarily that of X-rays
in 1895 and tKeir application to the determination of the structures of crys ta l-
line solids, have proved the Arrhenius theory inadequate. The theory is
valid for dilute solutions of weak electrolytes but fails to predict the experi-
mental results for moderately concentrated solutions. For solutions of strong
electrolytes the Arrhenius theory is invalid.

Let us consider sodium chloride as a typical strong electrolyte. We have
seen that molecules of NaCl do not exist. In the solid state there are no NaCl
entities as such, but only sodium ions, Na+, and chloride ions. Cl” so that
solid sodium chloride is already 100% ionized. Consequently there can be
no equilibrium between undissociated molecules' and ions in solution. In
solution the Na+ and Cl” ions merely dissociate. The term “dissociation” is
more proper than the term "ionization” when speaking of a substance which
is already ionized in the solid state. Dissociation refers merely to the separa-
tion of ions whereas ionization infers the formation of ions from neutral
molecules.

Nevertheless the three criteria listed on page 254 which determine the
conductance of a solution are still valid. To explain the increase in conduct-
ance of a strong electrolyte upon dilution, P. Debye and E. Huckel in 1923
developed a quantitative theory which is the basis of the modem treat-
ment of strong electrolytes. They assumed that strong electrolytes are 100%
ionized and that it is the mobility of the ions which varies with concentra-
tion. In solution each ion is believed to be surrounded, through electrostatic
attraction, by an "atmosphere” of oppositely charged ions. This atmosphere
reduces the mobility of an ion; for example, it retards its motion towards an
electrode during electrolysis. Thus a positive Na+ ion would have about it an
atmosphere of negatively charged Cl” ions while the Cl" ions would be
surrounded by an atmosphere of positive Na+ ions, as shown in Figure 16.4.
The negative Cl” ions around a Na+ ion would act as a "drag” and decrease
the latter s mobility. Thus the ions do not exhibit their maximum freedom
of movement. As the sodium chloride solution is diluted, the average dis-
tance between the Na+ and Cl” ions increases and the hindering effect of
an oppositely charged atmosphere is correspondingly reduced. Ionic mobili-
ties consequently increase with dilution. Ultimately dilution results in such
a distance between the ions that the effect of an ionic atmophere is reduced
to negligible proportions and the ions exhibit their maximum mobility and
conductance, a condition which corresponds to 100% ionization in the
Arrhenius theory.

258

Solutions of Electrolytes

Because an oppositely charged ionic atmosphere prevents an ion from
exerting its full individual effect, a solution of an electrolyte has an apparent,
or effective, concentration which differs from its nominal concentration. The
effective concentration is known as the activity, for which the symbol is a ,
and the ratio of the effective concentration to the nominal concentration,
c, is called the activity coefficient, {.

/ =

The activity coefficient decreases with an increase in concentration; for solu-
tions of different electrolytes at the same concentration the activity coefficient
is inversely proportional to the charges on the ions of the electrolyte. Activi-
ties can be determined by measurements of vapor pressure, freezing point,
and electromotive force but a description of such measurements is beyond
the scope of this book.

QUESTIONS

1. Distinguish between electrolytes and noneleetrolytes. Give examples of each,

2. Describe an experiment which would enable one to determine whether
benzene was an electrolyte.

3. How is the conductance of a solution measured?

4. Define resistance, conductance, specific conductance, molar conductance,
equivalent conductance, infinite dilution,

5. From their definitions, derive the units of specific conductance and equivalent
conductance.

6. What is meant by a “deviation by an electrolyte from Ruoult's Law”? What
is the cause of such deviation?

7. list the factors on which the conductance of an electrolyte depends,

8. By what methods can the degree of ionization of an electrolyte be determined?

9. Explain the statement, “the value of i approaches 3 as a limiting value for
potassium sulfate, K 2 $0 4 .”

10. What is the salient feature of the Debye- Huckel theory for strong electrolytes?
In what respect does this theory differ from the Arrhenius theory?

11 . How is the increase in conduction with dilution for a strong electrolyte ex-
plained by the (a) Arrhenius theory and (b) Debye-Huckel theory?

12. Define (a) activity (b) activity coefficient. How are they related?

13. Why should the activity coefficient vary inversely with (a) concentration and
(b) ionic charge?

14. What reasons can you suggest for the order in which the ionic conductances
of the alkali metal ions increases, namely, Li+, Na+, K*?

15. Which solution has a lower vapor pressure: 171 g of sugar, C, 2 H 22 O n , in a
given quantity of water, or 55.5 g of CaCl 2 in the same quantity of water?

18. Assuming complete dissociation what ratio of weights of NaCI and CaCI*
should be dissolved in water to give solutions having (a) the same vapor pres-
sure and (b) the same freezing point?

Solutions of Electrolytes

259

17. Theoretically what is the maximum possible boiling point of aqueous solutions
of (a) 0.2m NaCl and (b) 0.5m Ca(N0 3 ) 2 ?

18. The freezing point of a 0.0100m solution of acetic acid is -0.0195°C. Cal-
culate the degree of ionization of the acetic acid. Compare your answer with
that on page 256 and explain any difference in values.

19. For the acetic acid solution of Problem 18 calculate the boiling point.

20. At 25 °C the equivalent conductance of a 0.0040m aqueous solution of acetic
acid is 25.6 ohm -1 cm 2 equiv -1 . Calculate (a) the degree of ionization of the
acetic acid (b) the freezing point of the solution.

21. The compound, MA, is a weak electrolyte, ionizing MA M + + A". Calculate
the freezing point of a 0.050m aqueous solution of MA which is 10% ionized,

22 The specific conductance of a 0.0050N solution of a weak electrolyte, MA, is
7.0 x 10“ r * ohm -1 cm" 1 . The equivalent ionic conductances of the M+ and A-
ions are 150 and 50 ohm - * 1 cm 2 equiv~\ respectively. Calculate the degree of
ionization of MA. Ans: 7.0%

Ionic Equilibria— I
Acids and Bases

Inasmuch as most chemical reactions and analytical procedures in inor-
ganic chemistry involve ionic substances, the principles of ionic equilibria
are of fundamental importance. To the equilibrium that exists between
molecules and ions in a solution of a weak electrolyte, the concepts of
chemical equilibrium are applicable. These include LeChatelier's principle
and the expression of an equilibrium constant In the sense that this chapter
merely applies the principles of chemical equilibrium to ionic equilibria in
solution, little that is basically new is introduced.

1. Ionization of Weak Electrolytes. A) Weak acid: Acetic acid is a
typical weak acid. We have written its ionization as

H(C*H>O t ) H+ + (C,H»0,)-

For brevity the chemist sometimes uses the symbols HAc for acetic add and
At r for fee acetate ion. Using these symbols,

(I) HAc ^H+ + Ac~

In aqueous solution fee hydrogen ion, H+, which is simply a proton, can-
not exfat independently. The diameter of a proton, about 10' 1 * cm, is far
less than that of a simple ion, about 10* cm. Its large ratio of charge to
size, in conjunction wife fee presence of polar water molecules wherein
fee oxygen atoms contain unshared electron pairs, causes the proton to
unite wife a water molecule to form the hydronium ion, H s O+.

(2) H + + H,0 -* H s O+ or H + 4- :A - H -* I] H - A — HJ

More properly fee ionization of acetic add should be written

( 3 ) HAc + H*0 H.O+ + Ac-

Ionic Equilibria— I: Acids and Bases

261

For this equilibrium an expression for an equilibrium constant can be
written in a manner identical to that described on page 112.

[HsO+3 [Act]
[HAc] [H a O]

Since the concentration of water remains practically constant, it is custom-
ary to carry the factor, [H 2 0], over to the right side of the equation and
to include it in the value of the equilibrium constant.

( 5 )

[H.O+] [Ac-]
[HAc]

K x X [H 2 0] = K

K, not Ki, is defined as the equilibrium constant; for an ionization reaction,
the equilibrium constant is called the ionization constant.

If Equation 1 were used for the ionization of acetic acid, a similar ex-
pression would have been obtained except for the replacement of H s O+
by H+.

[H+] [Ad _ v
(6) [HAc] K

Though the free proton never occurs in solution, for simplicity the
chemist frequently writes ionization constant expressions in terms of the
hydrogen ion, H+, rather than the hydronium ion, H 3 0+. Whenever the
symbol H+ is used for an ionization in^ aqueous solution, however, it is to
be understood that the actual species present is the H 3 0+ ion.

Example 1: At 25 °C, by conductance measurements, it is found that a 0.10M
solution of acetic acid is 1.3% ionized. Calculate (a) the hydrogen ion concentra-
tion, and (b) the ionization constant.

Solution ; (a) Since 1.3% of the HAc undergoes the ionization reaction, the
number of moles of HAc that ionizes is (0.10 x 0.013) = 0.0013 mole. The [H+]
is therefore 0.0013 mole/liter.

(b) Since the H+ and the Ac- ions are produced in a 1 to 1 ratio, the [Ac-]
is also 0,0013 mole/liter. The concentration of undissociated HAc molecules re-
maining is (0.10 — 0.0013) = 0.0987 mole/liter.

0,0013 x 0.0013
0.0987

1.8 X 10- 5

Because the degree of ionization is small, in the calculation of HAc the value of
0.0013 could be neglected with respect to 0.10. If the value of t).10 were used
for the HAc, little more than a one percent error in the calculated value of K
would result. Particularly in this problem, the relative precision of the values
0.10 and 0.0013 warrant discarding the 0.0013.

Once the value of K is known it can be used for other equilibrium
calculations but only at the same temperature.

hmic Equilibria —7: AcidEsr and Bases

Example 2: For a 0,50 M solution of acetic acid at 25°C, calculate (a) the
hydrogen ion concentration and (b) the degree of ionization.

Solution:

Initial concentration, mole/iiter

HAc H 4 Ac-

0.50 0 0

Let x equal the moles per liter ot HAc that ionize. The numbers of moles
per liter of H+ and of Ac- ions that are formed are also equal to x.

Concentration at equilibrium, mole/ liter 0,50 —

v _ [H + ] [Ac-] _ (*)(*)

“ [HAc] 0.50 - x

X X

1.8 x H)~s

Solving for x (again the approximation may be made that the value of x in the
denominator may be neglected in comparison to 0.50; otherwise a quadratic equa-
tion must be solved),

x = 3.0 x HH

(a) The concentration of H t is 0,0030 mole/iiter; the [Ac*] is the same and
the [HAc] is 0.497 mole/liter.

(b) The degree of ionization, a, is — —

U.50

0.0030

0.50

= 0.006, or 0,6% ionized.

The fact that a weak electrolyte is only slightly ionized indicates that it
has a relatively great tendency to remain as undissociated molecules. Whereas
HAc has little tendency to ionize into H*+ and Ac~ ions, the reverse tendency
for H + and Ac~ ions to combine into molecular HAc is correspondingly
great. Thus for the reverse of Equation 1, H f F Ac* HAc, the ex-
pression for the equilibrium constant will be the reciprocal of Equation 6
and the value of the equilibrium constant for the combination of H+ and

Ac* ions will be

1.8 X 10*

or 5.5 X 10*, a large number.

B) Weak base: Ammonia, NH S , is a weak base. It reacts with water
to form ammonium ion, NH*+, and hydroxide ion, OH~

(7) NH« + H a O NH«+ + OH-

The expression for the ionization constant of ammonia is

/n V _ [NH<+1 [OHi

[NHs]

Again the concentration of water, [H a O], does not appear in the ionization
constant expression because it has been included in the value of K,
Appendix VI gives data concerning the ionization of a number of elec-
trolytes. The ionization constant for an acid is sometimes designated at K $ ,
and for a base as K b .

Certain salts, which are covalently bonded, are also slightly ionized.
There is not complete separation of the salt into ions when it is dissolved

Ionic Equilibria — 1: Acids and Bases

263

in water as is the ease with ionic salts. An example is mercury (II) chloride,
HgCI 2 .

HgCl, HgCl+ + Ch

2. Strong Electrolytes. Strong electrolytes are assumed to be complete-
ly ionized in solution. Since only ions are present and no undissociated mole-
cules in a solution of a strong electrolyte, there can be no equilibrium and the
discussion in the preceding sections is inapplicable. The writing of an
ionization constant expression for a strong electrolyte is without meaning.
Some typical strong electrolytes are hydrochloric acid, HC1; nitric acid,
HNO,; bases such as sodium hydroxide, NaOH, and calcium hydroxide,
Ca(OH) 2 ; and salts such as sodium chloride, NaCl, and sodium acetate,
Na(C 2 H,0 2 ).

3. Acids and Bases. Acids and bases have been variously defined
throughout the history of chemistry. The earliest definitions were purely
experimental. Thus acids turned blue litmus red, reacted with metals to
give H 2 , and neutralized bases; bases turned red litmus blue, neutralized
acids, and had a soapy feel. In the Arrhenius theory an acid was defined
as a compound which, in water solution, ionized to produce hydrogen ions;
and a base was a compound which, similarly, produced hydroxide ions. A
salt was a compound whose direct ionization yielded neither H+ ions nor OH"
ions. Thus HC1 was an acid, NaOH a base, and NaCl a salt. The reaction
of an acid and a base was a neutralization reaction which formed a salt and
water. These definitions have been found neither sufficient nor accurate.
For example, hydrogen chloride, HC1, dissolved in the nonaqueous solvent,
benzene, C 6 H 6 , gives no evidence of the presence of ions yet it affects in-
dicators and combines with ammonia to form ammonium chloride, a typical
acid-base reaction. So, too, gaseous HC1 combines with gaseous NH 3 . In
liquid ammonia, sodium amide, NaNH a , is a base. Liquid ammonia ionizes,
NH 3 H+ + NH 2 "; the amide ion, NHr, is a strong base akin to the
OH" ion where water is the solvent.

More general definitions of acids and bases were proposed independ-
ently in 1923 by the Danish chemist, J. N. Bronsted, and the British chemist,
J. M. Lowry: an acid is any substance that can transfer a proton to another
substance whereas a base is any substance that can combine with a proton.
Briefly, an acid is a proton-donor and a base is a proton-acceptor. The
transfer of a prot6n is known as protolysis , and a reaction involving proton
transfer is a protolytic reaction.

From this point of view, in the ionization of acetic acid,

HAc + H 2 0 ;=± HaO + + Ac-

molecular acetic acid, HAc, is an acid because it donates a proton to the
H 2 0 molecule; the latter, in accepting the .proton, thereby acts as a base;
If we now consider the reverse reaction, the hydronium ion, H a O+, is an
acid and the acetate ion, Ac~ is a base. Thus an acid and a base yield an
acid and a base.

(9) HAc 4- H 2 0 H 3 0+ + Ac-

acid x base 2 acid 2 basej

264

Ionic Equilibria — 1 : Acids and Bases

For the ionization of ammonia,

(10) NH< r 11,0 ^ OII-

basei acids acidj base-

An add and a base, related to each other as are the HAc and the Acr
ion, and the NH.» and the NH^ ion, so that one can be formed from the
other by the loss or gain of a proton, are known as a conjugate acid-base
pair. In general,

Acid ^=± Proton 4~ Base

In this system, an acid or a base may be any species, a neutral molecule
or an ion. Even hydrated metal ions, as A!(H .O) n :i 1 , can give off protons
and act as acids. The equations below illustrate these points; the same
subscripts indicate a conjugate acid-base pair.

Acid,

Base,

Acid,

Base,

(11)

HCI

H,0

H;,0 +

ci-

(12)

IhSO.

H a O

h 3 o+

4 -

HSOr

(13)

HSOr

4 *

H a O

"T— "

H n O+

-J.

SQ/-

(14)

A!(H 5 0) b *+

4 "

H a O

Ai( H a O) 5 (OH) 2 +

(15)

HC1

NH,

NH,+

Cl~ ( gaseous reaction)

(16)

HAc

NHj

NH.+

Ac - (in liquid ammonia)

According to the Brdmted- Lowry concept, the strength of an acid is
measured by its tendency to lose protons and the strength of a base by
its tendency to gain protons. It follows that the stronger an acid is, the
weaker is its conjugate base; the converse is also true. Because HAc is a
weak acid, the Ac* ion is a strong base. In an acid-base equilibrium, ioniza-
tion will proceed in that direction which forms the weaker acid and the
weaker base. Since acetic acid is only slightly ionized Equation 9 proceeds
to the right only to a small extent and we could predict that the H 3 0+ ion
is a stronger add than is HAc, and the Ac* ion is a stronger base than
H 2 0. Since HC1 is a stronger acid than the H 3 0 + ion, it is highly ionized
and H 2 0 is a stronger base than the Cl“ ion. In the following table, a num-
ber of acids are listed in the order of decreasing strength; their conjugate
bases in the second column are therefore in the order of increasing strength.

Any add in Table 19-B will react appreciably with any base below its
conjugate base in the table. Thus HNO a will react with HS“ ion whereas
NH*+ will not to any appreciable extent. The stronger the add and the
stronger the base the greater will be the extent of reaction between them.

Whether a substance acts as an add or a base may not be evident from
its formula but depends upon the reaction it undergoes. In Equation 9
water is a proton-acceptor (base; whereas in Equation 10 water acts as
a proton-donor (add). Thus the water molecule can act both as an add
and as a base, dependent upon its mode of behavior. When confronted with
a stronger add than itself, the water molecule acts as a base; against a

Ionic Equilibria— I: Acids and Bases

265

Table I9-B

Strengths of Acids and Bases

Conjugate Acid

Conjugate Base

hcio 4

Perchloric acid

cio 4 -

Perchlorate ion

HCI

Hydrogen chloride

Cl-

Chloride ion

h 2 so 4

Sulfuric acid

hso 4 -

Hydrogen sulfate ion

HNOj

Nitric acid

NO-

Nitrate ion

H 3 0+

Hydronium ion

h 2 o

Water

HS<V

Hydrogen sulfate ion

so, 2 -

Sulfate ion

h s po 4

Phosphoric acid

h 2 po 4 -

Dihydrogen phosphate
ion

Acetate ion

CHhCOOH

Acetic acid

CH.COO-

A1(H 2 0) 6 *+

Hexaaquoaluminum (ill) AI(H 2 0) 5 (0H)-+
ion

Hydroxopentaaquo-
aluminum (III) ion

h 2 s

Hydrogen sulfide

HS-

Hydrosulfide ion

NH 4 +

Ammonium ion

NH {

Ammonia

HCN

Hydrogen cyanide

CN-

Cyanide ion

h 2 o

Water

OH-

Hydroxide ion

c 2 h 5 oh

Ethyl alcohol

C a H s O-

Ethoxide ion

nh 3

Ammonia

nh 2 -

Amide ion

h 2

Hydrogen

Hydride ion ^ r

weaker acid the water molecule acts as an acid, and the weaker acid behaves
as a base if chemically possible. Such a substance which can act as either
an acid or a base is called amphiprotic. Inspection of Equations 12 and 13
indicates that the hydrogen sulfate ion, HS0 4 “, is also amphiprotic. On
the other hand, because they have but one mode of reaction, the hydroni-
um ion, H 3 0+, can act only as an acid, and the sulfate ion, S0 4 2 ”, only
as a base.

In aqueous solution the strong acids, HC10 4 , HC1, H 2 S0 4 , and HN0 3 ,
all appear to have the same acidity. All react with water almost completely
to form the same acid, H :{ 0+, which is the strongest acid that can exist in
aqueous solution. Thus equivalent concentrations of these acids show the
same maximum acidity because they contain about the same concentra-
tion of the same acid. This equalizing of acid strength in aqueous solution
is known as the levelling effect. Similarly the strong bases seem to have
the same basic strength because OH” ion is the strongest base that can
exist in water solution and bases stronger than OH” ion are levelled to the
strength of that ion. In liquid ammonia as solvent, the strongest acid that
can exist is the ammonium ion NH 4 +, and the strongest base the amide ion,
NH 2 ~-. According to Bronsted theory, the ionization of ammonia would be

(17) NH 8 + NH». NH 4 + + NH*”

In solvents other than water, such as glacial acetic acid, the difference in
acid strength between the strong acids becomes evident because it is more
difficult to- force an acetic acid molecule to accept a proton than it is with
water. In acetic acid the ionization of perchloric acid is;

(18) HCI0 4 + CH 3 COOH -» CH 3 COOH 2 -h + C10 4 ”

For a substance to act as a Bronsted acid, it must contain a proton which
it can donate. In many reactions protons are not transferred yet the re

266

itmir Ktfmlihm — 1. Adds and Bases

actions have the characteristics of an ueid-huse reaction. The most general
definition of acids and basts was proposed in 1923 by the American chemist,
G. N. Lewis. According to Lewis, a base is any species that has an un-
shared pair of electrons which it can furnish to form a coordinate covalent
bond, and an acid is a substance that accepts the electron pair. Thus a
base is an electron-pair donor and an acid is an electron-pair acceptor.
In the chemical reaction between NIL and BF, (page 175) NH* is a
Lewis base and BF { is a Lewis acid.

If F

| |

(19) H - NS + B - H

I I

In the combination of sulfur trioxide, SO<, and the oxide ion, O 2 ", of a
metal oxide such as CaO, the SO* is a Lewis acid and the O- ion a Lewis
base.

( 20 )

l O

H F
N* : B

H - N *. B - F

I i

II F

The proton is a Lewis acid because it can accept an electron pair to
complete its valence shell while the OH" ion, NIL, and SO., act as Lewis
acids because they have unshared electron pairs they can donate. The
Lewis concept is the most fundamental and includes the Arrhenius and
Bronsted definitions. Substances such as BF*, difficult to classify within
the framework of the Arrhenius or Bronsted theories, are included within
the Lewis definitions. An examination of every Bronsted acid-base reaction
will show that the base donates an electron pair to a proton; thus every
Bronsted base is also a base in the Lewis sense.

The Lewis concept based on electron pairs makes the field of acid-base
chemistry a subdivision of covalent bond formation. For a substance to act
as a Lewis acid, it needs merely to have an unoccupied orbital in a valence
shell Even a simple metal ion, such as AP + , can act as a Lewis add
by forming coordinate covalent bonds to the oxygen atoms of H 2 G molecules,
as in the hydrated ion, [Al(H 2 0)«] a * f . Molecules which are electron de-
ficient, as BF 3> or which can expand the octet by utilizing vacant d orbitals,
as SiF # , also can function as Lewis acids. The strength of a Lewis add
depends upon its tendency to attract electrons. For simple ions, this is
directly proportional to the charge of the ion and inversely proportional to
the size of the ion. On both these counts the Al s + ion is a stronger Lewis
add than Mg 2 *.

Ionic Equilibria— I: Acids and Bases

267

No theory of acids and bases is inherently the “correct” one. Chemistry
is a descriptive science and the several theories are best applicable in differ-
ent situations. For aqueous solutions the Bronsted theory is perhaps most
applicable. For simplicity, though, we shall hereafter write equations for
ionization in aqueous solution in terms of the hydrogen ion, H+, though
recognizing that the ionic species actually present is the hydronium ion,
H 3 0+.

4. Polyprotie Acids/ An acid such as HC1 which yields but one pro-
ton upon dissociation is called a monoprotic acid, and sometimes a mono-
basic acid because it combines with OH~ ion in a 1 to 1 mole ratio.
Molecules of some acids can dissociate more than one proton. Thus H 2 S0 4
can dissociate two protons per molecule and H 3 P0 4 three protons per mole-
cule. Such acids are known as polyprotic acids ( polybasic acids); H 2 S0 4
is diprotic ( dibasic ) and H 3 P0 4 is triprotic ( tribasic ).

Polyprotic acids ionize in steps. The primary ionization is always greater
in extent than the secondary, and so on. This is so because the first proton
is released from a neutral molecule, but the second from an ion which is
negatively charged and so tends to hold protons with a greater attractive
force than did the original molecule. Thus hydrogen sulfide, H 2 S, (more
properly hydrosulfuric acid) ionizes in water to give a proton and the hydro-
sulfide ion, HS~.

(21) H 2 S H+ + HS-

The hydrosulfide ion ionizes further to yield another proton and the sulfide
ion, $ 2 "

(22) HS‘ H+ + S 2 *

Since the secondary ionization of a polyprotic acid is not so extensive as
the primary ionization, H 2 S is a stronger acid than is the HS~ ion. The HS~
ion is amphiprotic; it is also a base, but a weaker base than the S 2 ~ ion.

Each ionization of a polyprotic acid can be represented by an equilibrium
constant expression and an ionization constant. For the two successive ion-
izations of H 2 S, we can write

( 23) m m = Kl = Li x io-

[H*S]

(24) [H+] [S'*-] _ K _ 10 x 10 -i4

^ ' [HS]

Equations 23 and 24 can be combined by ordinary multiplication

[H+] [HS-] v [H+] [S*-]
[H 2 S] [HS-]

= K x X K* = (1.1 X 10-) (1.0 X 10 ^ 4 )

[H+3« [S*-] __ R =
[H,S]

(25)

1.1 X Id 21

Ionic Equilibria — l; Acids and Bases

The value of K in Equation 25 would correspond to a chemical equation
obtained by the addition of Equations 21 and 22, that is,

(26) H-S 2 H+ f S'-- K = 1.1 X 10* 1

5. Ionization of Water. The fact that water does exhibit a small, though
measurable, electrical conductivity indicates that it is ionized slightly. The
equation for its ionization is

(27) H*0 H* t GH-

(27a) or H 3 0 f H-O H.O* - OH-

acidi base.. acid... base t

In Equation 27a water molecules act both as an acid and a base. At 25°C
equilibrium is attained when about 0.0000002% of the water has been ionized.
The expression for the ionization constant of water is

[H+3 [OHi _
[H.O]

or the equivalent

[H,0+] (OH-] „

[h 3 o] (h 3 o)

K,.

Since the concentration of undissociated H 3 0 molecules is practically con-
stant, the value of [H s O] may be included in the ionization constant for
water, K w , thus:

(28) [H+] [OH-J =K, X [H.O] = K*

or [H,0+] [OHi = K t X [H.O]-’ =K W

At 25°C, the value of this constant, K„, also called the ion product of water,
is 1 X 10-*"*, when concentrations are expressed in moles per liter.

In pure water or in any dilute aqueous solution the equilibrium expressed
in Equation 27 (or 27a) is always present and the product of the concen-
trations of the hydrogen and hydroxide ions must always equal 1 X I0" M .
In pure water, the concentration of hydrogen ion equals that of hydroxide
ion, that is, [H+] = [OH*]. Hence

(29) [H+] = [OH-] =^Tx~W r = 1 X lit’ mole/liter

The degree of ionization, and so the concentration of hydrogen and hy-
droxide ion and the ion product of water, increase with a rise in tempera-
ture.

Because die ionization constant for water remains constant even when
electrolytes are dissolved in it, it follows that, as die concentration of hy-
drogen ion increases, the concentration of hydroxide ion decreases propor-
tionally. Conversely, as the concentration of the hydroxide ion increases,
the hydrogen ion concentration decreases. For any aqueous solution,

[H+]

1 X IQ 44
[OHI

1 X 10 44
[H+]

(30)

and

[OH-] *

Ionic Equilibria— I: Acids and Bases

Thus, in a solution in which the concentration of hydrogen ion is one mole
per liter the concentration of hydroxide ion must be 1 X 10 44 mole per
liter. When the molar concentration of OH- ion is 0.01, the concentration
of H+ ion is

[H+]

1 X 10 44

0.01

= 1 X 10* 12 mole/liter

In aqueous solutions of both acids and bases, both hydrogen ion and hy-
droxide ion are present. An acid solution is one in which the [H+] is

greater than the [OH“], that is, greater than 1 X 10- 7 mole/liter. A basic

solution is one in which the [OH“] exceeds the [H+]. Where the [H+] is

equal to the [OH"], the solution is neutral.

6. The pH Scale. The hydrogen ion concentration is a measure of
the acidity, or the basicity, of an aqueous solution. In many branches of
chemistry we deal with solutions wherein the [H+] is very small, e.g., in
neutral or basic solutions. The handling of such concentration values would
be mathematically cumbersome. Chemists have found it convenient to use
a measure of acidity known as the pH scale, first proposed by S. P. Sorensen
in 1909. The pH is defined as

(31)

pH = log

[H+]

— log [H+] where [H+l is in mole/liter.

The following examples will show what is meant by pH, and how it is
calculated

(a) If the [H+] is 1 X lO 7 mole/liter, pH = -log (1 X 1(H) = -(-7) = 7

(b) If the [H+] is 0.0001 or 1 X 1CH, pH = -log (1 X 1(H) = 4

The reverse calculation of the [H+] from a knowledge of the pH can
be done as follows:

(d) If the pH is 5.80 what is the [H+]?

pH = —log [H+]; hence 5.80 = —log [H + ] = log

[H+l

the antilog of 5.80 is 6.31 X 10 5

6.31 X 10 5 == — -~r

[H+]

[H+l = \ = J_ X 10* B = 0.16 X 10- 5 = 1.6 X 10-®M

6.31 X 10 5 6.31

Pure water or a neutral aqueous solution has a pH of 7, An acid solution
is one having a pH less than 7, and a basic solution is one with a pH greater
than 7. The lower the pH value, the greater the acidity.

270

ionic Eijuihhrui — I: Acids and Bases

Example 2: Calculate the pH of a 0.020AI solution of NH
Solution: The equation for the ionization equilibrium of Nil ^ is
NH, +H..O ^ NII 4 + OH-

and the ionization constant expression is

[NH,-] [Oil-] _ t .
k

l.H x 10*

NH, NH 4 + OH"

Initial concentration, mole/iiter 0.020 0 0

Let x equal the moles per liter of NH , that undergo ionization. The numbers
-jd£ moles per liter of NH 4 -f and of OH“ ions that are fonned are also equal
to x. The concentration of NH< is {0.020 — x), but neglecting the slight de-
crease in its concentration due to the ionization, |NH ; ] - 0.020 mole/ liter.

Concentration at equilibrium 0.020 x x

_ <*> w

0.020

1.8 x kh

x ss 6.0 x 10 *

Thus the hydroxide ion concentration, [OH ], is 6.0 a 10'* mole/liter (not the[H+])

[H+]

1 * 10 - 1 *

[OH-]

pH = - log [H 4 *] =s - log (1.67 x

1 < I0~*f
6.0 x 10 4

io-*o - -

« 1.67 x 10-n
(0.22 -11) = 10.78

Akin to the pH as a measure of the strength of an acid or a base is the pK.
The pK is related to the ionization constant, K, in a manner analogous to
that by which the pH is related to the [H + j; pK ss — log K. The pK of
water is 14. A pOH unit can be similarly defined; pOH — log [OH - ].
For water or an aqueous solution,, the relation holds that pH f pOH = 14.

7. Indicators, A chemical indicator is a substance which exhibits dif-
ferent colors at different values of pH, The indicator thymol blue is yellow
in a solution whose pH is 6.0 or less and is blue if the pH is greater than
9.6. At values of pH between 8.0 and 9.6 there is a gradual change in color
from yellow to blue.

A substance which is an indicator is either a weak add or a weak base.
The molecular formula of an indicator is generally complex, most being
organic dyes. For one which is a weak acid, as is thymol blue, we might
represent the formula as H(In), a substance which can ionize slightly in
aqueous solution to produce a proton and a residual “indicator’" ion, In -
A simplified representation of the equilibrium which exists is

(32)

H(In) ;=± H+ + hr

The H(In) molecule and the ion, In - , are differently colored. An equilibrium
constant expression for the indicator may be written

( 33 )

* = [H+] [In - ]
[H(In)]

Ionic Equilibria — I: Acids and Bases

271

The color of & solution to which the indicator has been added will depend
upon the ratio of the concentrations of the H(In) and the In' ion, that is,

[ ~ H(in yj For a given indicator this ra tio. in turn, will depend upon the

hydrogen ion concentration. According to LeChatelier’s principle, in an
acidic solution the high [H + ] displaces the equilibrium of Equation 32 to
the left. The [H(In)] will be relatively large and the [In - ] correspondingly
small. The solution will thereby evidence the color of the H(In) form.
In a basic solution, the removal of the hydrogen ion by the base will .dis-
place the equilibrium to the right and the color of the In - ion will be ap-
parent. Similar considerations apply to a weak base type of indicator whose
formula might be represented as In (OH).

The color change of an indicator is not generally at a pH of 7. Each
indicator changes hue at a fairly definite value of pH, that is, over a range
of one or two pH units, which is characteristic of the indicator and its equi-
librium constant. This pH value may be anywhere along the pH scale.
An indicator exhibits an intermediate color when [H(In)] equals [hr],

r-r -»

that is, — ~ — =-~ = 1. At this point the concentration of H+ in the solution
[H(In)J

is equal to the equilibrium constant for the indicator, or the pH pK.
Figure 20.1 gives the colors of some common indicators and the pH values
at which their color changes occur.

CH+J = 10-1 10* s 10-5 ifc-7 10-9 10-11 10-1»

pHts O 1 l 3 4 5 6 7 8 9 10 11 12 13 14 Indicator P K

! Yellow ; 7

more j

Neutral ! 1 more •

} to Green*

» to Blue >

“? fc Violet

(Water) j j basic |

Methyl Violet 1.0

Thymol Blue 1.5; 8.9

Congo Red 4.0

Methyl Orange 3,7

Brom Cresol Green 4 6

Methyl Red 5.1

Litmus ca. 6 5

i l Orange ,

| I to Yellow

j j

lYellow

« ” !v.Uow j |

1 ! Blue

| Bh “ i

B i“Wj

to KedJ

! i | i !

■

Red

Rea ! i

.3W!

: 1 1 • 1
| Yellow ! !

Yellow

{Vellnw
■ to Cceen

1 | 1 I I'-
ll \ 1 i

J J Green 1 ;

i *

i I J j 1

y I ! I Yellow ;

I 1 Red

** _L 1

* { i i J

Purpje ! 1 | {

to Blue » ! ! BIue»

I"" 1

Red >

Rod- { | i *

“tn^w ! ! Yellow

Neutral Red 7.4

Phenol ph thalem 9.4

Thymolphthalein 9.8

— T 1 ' 1 -

i t I Colorless { J i

i Colorless ! | to^Red i \ Re

{ j Colorless {

l Colorless i to ^lue *

Blue

If the pH of a solution were varied slowly, the color change of ,an indicator would
take place over an interval ef 1 or 2 pH units. By the pH at which an indicator
changes color is meant the mid-point of the color change interval. If the pH is
changed rapidly over this interval only the two extreme colors of the indicator
would be observed.

Figure 2Q.L Color Change and pH Range of Indicators.

272

ionic Equilibria — /; Acids and Bam

The pH of a solution can he determined to within a few tenths of a
unit by comparison with solutions of known pH to which the proper indi-
cator has been added. In this comparative method the indicator used should
be one whose color change occupies a pH interval which includes the pH
of the unknown solution. In using an indicator only a small quantity is
required. For solution volumes of about 50 ml a couple of drops of indi-
cator are sufficient. The pH of a solution may be accurately determined
by an electrochemical measurement of the potential of u cell in which the
solution of unknown pH is the electrolyte, (Chapter 24)

8. Neutralization. When an acid, as HCl, reacts with a base, as NaOH,
to form a salt, NaCl, and water, the molecular equation can be written

(34) HCl + NaOH - NaCl + H a O

However, the HCl, NaOH, and NaCl are all strong electrolytes, completely
dissociated, so that they exist solely as ions and the equation should be
written, more properly,

(35) H+ -t Or + Na+ + OH- -* Na * f Or + H a O

inasmuch as the Na + and Cb ions appear on both sides of the equation,
they truly do not take part in the reaction. They are merely “spectator
ions’* which must be present to maintain electric neutrality in solution
but they can be omitted in writing the equation. This leaves

(38) H+ + OH-~* H a O

The sole reaction, when HCl is added to NaOH, is the combination of H+
and OH* ions to form water. Such a reaction of the ion of an acid and
the OH- ion of a base to form water is known as neutralization. In aqueous
solution neutralization is a protolytic reaction forming water. In the neu-
tralization of any strong acid by any strong base, water is the only product
and hence the same amount of heat is evolved, 13,400 calories per mole
of water formed, independent of the molecular formulas of the acid and
the base. When a weak acid or a weak base is neutralized, the value of
13,400 is modified by the amount of heat required to ionize the add or the
base as the ions are used up in the formation erf water*

9. Addimetry and Alkalimetry : Titration. In a neutralization reaction,
as in all chemical reactions, the numbers of chemical equivalents of the
reacting substances are equal For the reaction between HCl and NaOH,
the number of equivalents of HCl equals the number of equivalents of
NaOH. In a solution the number of gram-equivalents of the solute equals
the normality of the solution times its volume in liters.

(37) Number of gram-equivalents = N X V

It follows, therefore, for a neutralization reaction, the product (N X V)
for the add equals the (N X V) for the base, Thus

(38) W. X V* « N b X V*

Ionic Equilibria — I: Acids and Bases

273

where the subscripts a and b refer to acid and base, respectively. In Equa-
tion 38, since V appears on both sides the volumes need not be in liters.
However, V* and V;, must be expressed in the same units of volume, e.g., ml.

If any three factors of Equation 38 are known, the fourth can be cal-
culated. The equation can thus be used analytically to determine the un-
known concentration (normality) of an acidic or basic solution. The ex-
perimental procedure is known as titration. To determine the unknown
normality of a solution of HC1, a definite volume of it is taken and an
indicator such as phenolphthalein is added. In an acid solution this indi-
cator is. eolorless. To the HC1 solution measured quantities of a solution

A definite volume of a solution (HC1), whose normality is to be determined is
placed in a beaker. This volume may be obtained from a buret or a pipet, both
of which are volumetric apparatus designed to deliver known volumes of a liquid.
The buret has a stopcock at its lower end to shut off the flow of liquid, A few
drops of indicator are added to the solution in the beaker. While this solution
is continually stirred, a standard solution of the other reagent (NaOH) is added
slowly from a buret to the point where the indicator shows a permanent change
in color. This is the end point of the titration. The volume of this solution is
recorded and the normality of the unknown solution can be calculated by Equa-
tion 38.

Figure 202. The Titration Process.

of NaOH of known normality are added slowly from a buret, as indicated
in Figure 20.2. So long as insufficient NaOH has been added to react with
all the HC1, that is, while excess HC1 is still present, the pH will remain
fairly constant and the phenolphthalein will be colorless. When one drop

274

Ionic Equilibria — 7; Acids and Bases

of excess NaOH has been added that is, one drop of NaOH more than is
required to react with the HG1, the pH will rise sharply dm* to the excess
NaOH and the phenolphthalein will exhibit the pink color of a basic solu-
tion. This point, or the volume of NaOH at which the indicator color
change occurs, is called the end point of the titration.

The normality of the acid can now be calculated since we know the
normality and the volume of the NaOH required to neutralize a known
volume of HC1,

Example 3; In a titration, 100 ml of an HC1 solution require 37.5 ml of
3.00 N NaOH for neutralization. Calculate (a) the normality of the acid and (b) the
weight of HCI in the 100 ml.

Solution:

(a) *\ r no, *
N nn =
(b) w =

100 ml ss 3.00.V

3.0QN x 37.5 ml
100 ml

*V x V x GEW

37.5 ml

U25N

vv = 1.125 ff'-g 1 -- x 0.100 liter x 36.5 S-~=41.1 g

liter g-eqmv a

In a similar manner the normality of a base can be determined by titrating
a definite volume of an unknown base with an acid of known normality.
The selection of the proper indicator to be used in a titration depends
upon the relative strengths of the acid and the base, that is, the pH of
the solution at the end point of the titration. For example, in the titration
of NaOH with acetic acid, when equivalent amounts of acid and base have
been added, the result is a water solution of sodium acetate. Sodium acetate
is a salt which is completely ionized into the weak-acid Na 4 * ion and the
strong-base CHsCOO* ion. By hydrolysis (Chapter 21), a solution of sodi-
um acetate has a pH of about 8,5 and an indicator that changes at this pH
should be used. In practice any indicator with a pH range between 5 and
9 can be employed when one of the solutions is either a strong base or a
strong acid because at the end point one additional drop of the strong
acid or base will cause a large change in pH. Phenolphthalein is com-
monly used in such cases because of its sharp color change,

QUESTIONS

1. Write equilibrium constant expressions for the ionization of (a) the weak
hydrocyanic acid* HCN (b) the HSO-f ion (e) a weak base, M(OH) (d) the
secondary ionization of H a S.

2. Define acid and base according to (a) the Arrhenius theory (b) the Bronsted-
Lowry theory.

3, In what respects is the Arrhenius theory inadequate?

4, Illustrate what is meant by a Bronsted conjugate acid and base.

5. What governs the direction in which a Brdnsted acid-base reaction will take
place?

6, Explain the statement, "the stronger an acid, the weaker its conjugate base.*

Ionic Equilibria — 1: Acids and Bases

275

7. Write equations illustrating how the following can act as a Bronsted acid
and as a Bronsted base: (a) a molecule (b) an ion (c) a hydrated metal ion.

8. Define and illustrate what is meant by an amphiprotic substance.

9. Define the term “polvprotic acid/* Write the stepwise ionization of the
triprotic acid, H 3 P0 4 . Label all Bronsted acids and bases therein.

10. Why is the primary ionization of a polyprotic acid always greater in extent
than the secondary ionization?

11. Arrange the following in increasing order of acidity: PO* 3 -; H 2 P0 4 “; HP0 4 2 -.

12. Arrange the following in increasing order of basicity: (a) H 2 S : HS~; S 2 ~; H 2 S0 4 ;

(b) Cl”; OH"; H a O; H s O+.

13. Which is the stronger acid in each pair: (a) NH 4 + and NH S (b) CN” and OH”?

14. What is the Lewis concept of acid and base? Why is it the most general
of acid-base definitions?

15. Explain why the following can act as Lewis acids: (a) Zn 2 + ion (b) A1C1 3

16. Predict whether the following can act as Lewis acids or bases: (a) BeF 2
(b) Br- ion (c) H 3 C - NH 2 (d) Ag+ ion.

17. Derive the expression for the ion product of water. Prove that for water
pH 4- pOH = 14

18. What is an indicator? Why does the color of an indicator change with a
change in pH?

19. Describe the procedure loiown as titration. How is the proper indicator selected
for a titration?

20. Write an equation for the neutralization of a strong acid and a strong base.
What conclusion can you draw in the event a heat of neutralization is not
13,400 cal/mole?

21. At 25 °C a O.OlOAf solution of monoprotic formic acid, HCOOH, is 4.2%
ionized. Calculate the ionization constant of the' acid.

22. For a 0.010M solution of HCN at 25 °C calculate (a) the hydrogen ion con-
centration (b) the hydroxide ion concentration.

23. Calculate the pH corresponding to each of the following hydrogen ion con-
centrations: (a) 1 x 10-* (b) 2 X 10” 3 (c) 5 x 1(H (d) 3.6 x 10” 10 .

Ans: (a) 3.0 (b) 2.7 (c) 2.3* (d) 9.4

24. Convert the following pH values to their hydrogen ion concentrations: (a) 6.0

(b) 3.4 (c) 6.8 (d) 8.9. Which of -these are basic solutions?

25. Calculate the pH of the following solutions (make any necessary assumptions):
(a) 0.01M HC1 (b) 0.01M NaOH (c) 0.01M acetic acid (d) 0.01M ammonia.

26. Assume a weak acid, HX, which ionizes HX H+ 4- X”, and for which
K = 9.0 X 10” 7 . For a 0.40M solution of HX calculate (a) concentration of
H+ (b) concentration of X - (c) concentration of undissociated HX (d) pH of the
solution (e) concentration of OH” (f) pK value of the acid.

27. For a 0.10M solution of H 2 S calculate the concentrations of (a) H+ (b) HS~

(c) S 2 ". Assume that the H+ comes solely from the primary ionization and that
the secondary ionization does not decrease the HS _ concentration.

28. Calculate the pH of a O.OIOM solution of H s P0 4 . Ans: 2.25

29. Using Figure 20.1 estimate the approximate pH of the following solutions:
(a) a solution is red with congo fed and red with neutral red (b) a solution is
yellow with thymol blue and red with methyl orange (c) a solution is yellow
with neutral red and colorless with phenolphthalein.

276

tome Equilibria— -I: Acids and Bases

30. If 75.0 ml of 0.10N HC1 are added to 50 mi of O.WN NaOH, what is the

pH of the resulting solution? Ans: 1.70

31. To what volume of water should 50 ml of O.HhV HC1 be added to malfA -
solution of pH = 2.50?

32. For the addition of 200 ml of 0.10N NaOH to 100 ml of 0.10N HC1, draw
a graph of pH against volume of NaOH.

33. The pH of a 0.010N solution of KCN is about 11.5. If you were to titrate
Q.OIOiV HCN with 0.0874N KOH, what indicator would you select? Ex-
plain your choice.

34. To neutralize 28.5 ml of a sulfuric acid solution, 35.5 ml of G.25N NaOH
are required. Calculate the normality of the acid solution. How would your
answer be affected if the acid were HCl instead of H 2 S0 4 ? Ans: 0.31N

35. (a) Calculate the normality of an acid solution, of which 36.4 ml neutralize
24.6 ml of a solution containing 56.2 g NaOH per liter (b) if the acid is
sulfuric acid, what weight of H 2 SO, is in the 36.4 ml?

Ans : (a) 0.97 N (b) 1.73 g

36. What volume of 0.25M H 2 S0 4 is required to neutralize 10.0 g NaOH?

Ans: 500 ml

37. If 50 ml of HCl of pH — 1.0 and 50 ml of HCl of pH as 5.0 are mixed, what
will be the pH of the mixture? Assume the volumes are additive.

Ionic Equilibria— II
Common Ion Effect
& Hydrolysis

1. Common Ion Effect. In a solution of acetic acid, the concentration of
hydrogen ion produced in the equilibrium, HAc H + + Ac", results in a
pH of about 3. When the salt, sodium acetate, is added the pH of the
solution rises to about 5, a hundredfold decrease in the hydrogen ion con-
centration. The salt, sodium acetate, is completely dissociated into sodium
and acetate ions, so that its addition adds Na+ ions and Ac‘ ions to the
solution of HAc. The Na + ions have negligible effect but the increased
concentration of Ac' ions, in accord with Le Chateliers principle, displaces
the point of equilibrium to the left in the direction of molecular HAc. This
results in a decreased H+ concentration and a consequent rise in the pH.
The displacement of an ionic equilibrium by the addition, from an external
source, of an ion which is already present in the equilibrium mixture is
known as the common ion effect . The lowering of the acidity of a solution
of a weak acid, HAc, by the addition of a salt, NaAc, containing the same
negative ion, Ac~, as the weak acid is an example of the common ion effect. This
action is sometimes referred to as the repression of the ionization of the
weak electrolyte, in this case the HAc. Thus the hydrogen ion concentration
can be altered by the addition of a sr’t containing no H+ itself. Conversely,
if HC1 is added to the original HAc solution, the concentration of Ac r ion
would be decreased.

The operation of the common ion effect can be illustrated by an example
involving acetic acid and sodium acetate.

Example I; (a) Calculate the pH of a O.lOAf solution of acetic acid, which
is 1.34% ionized at 25° C; (b) Calculate the pH of a 0.10 M solution of acetic acid,
to one liter of which has been added 0.10 mole of solid sodium acetate.

Solution: (a) Since the 0.10M HAc solution is 1.34 ionized, then

[H+] = 0.00134 mole/liter
[Act] = 0.00134 mole/liter
[HAc] =5 0.09866 mole/liter

The pH of this solution = —log (0,00134) = 2.87

278

Ionic Equilibria — II: Common Ion Effect and Hydrolysis

(b) Assuming the volume remains constant at one liter upon the addition of
the sodium acetate, the added NaAc represents an additional concentration of
Act ion, equal to 0.10 mole/liter. If no displacement of the HAc equilibrium
occurred, the concentration of Ac* ion would be 0.10134 mole/liter. However,
due to the displacement of the HAc equilibrium to the left, x moles of H+ and
x moles of Ac* combine to form .* moles of molecular HAc. When equilibrium
is re-established the solution will contain the following concentrations:

[H+] = (0.00134 - x) mole/liter
[Ac-] = (0.10134 - x) mole/liter
[HAc] = (0.09886 + x) mole/liter

Substituting these values in the ionization constant expression for HAo:

y - fH+j [A Cl _
K [HAc] “

(0.00134 - x) (0.10134 - x)
(0.09866 + x)

= 1.75 x 10-5

Solving, x = 0.00132 mole/liter

The equilibrium concentrations are therefore:

[H+] = (0.00134 - 0.00132) = 0.00002 mole/liter

[Ac-] = (0.10134 - 0.00132) = 0.10002 mole/liter

[HAc] = (0.09886 + 0.00132) as 0.09998 mole/liter

The ionization of the acetic acid has been repressed because the [H+] has
been reduced, by the addition of sodium acetate, from 0.00134 mole/liter to
0.00002 mole/liter. The degree of ionization of the solution containing both 0.10M
HAc and O.10M NaAc is 0.02% and the pH of the solution is 4.7.

Again the mathematics could have been simplified if we had negelected the
value of x, which we found to be 0.00132, in computing the values of [HAc] and
[Ac*]. Particularly in this case of the common ion effect where the ionization of
the acetic acid is even less than it is in water alone, we may assume - that the
[HAc] in the equilibrium mixture equals its initial concentration independent of
any decrease that occurs through ionization, i.e., 0,10 mole/liter. Similarly the
[Ac-] may be assumed to come solely from the added salt, NaAc, and to be 0.10
mole/liter. With these assumptions, the calculation of the final [H+3 is simplified.
rti.fi ^ a |a

Thus - — = 1.8 x jo-8 ( w fth the lesser precision, the value

of 1.8 x 10-® may be used for K)
and [H+J = 1.8 X 10-* mole/liter (compare with the more accurate value

of 2 x 10-8 mole/liter}

The common ion effect is hot observed with strong acids, nor with weak
adds if a high concentration of other ions is present. The pH of a HO
solution is not raised appreciably, if at all, by the addition of NaCl. The
common ion effect also applies to solutions of weak bases. In a solution of
ammonia, the hydroxide ion is due to

NH» + H,0 NH.+ + OH-

asd its concentration can be reduced by the addition of a salt, such as
NH.C3, which yields the NH«+ km.

An important application of the common ion effect is the control of the
sulfide ion concentration, [S**], in the hydrogen sulfide equilibrium,

H*S ±5 2 H+ + S*-

Ionic Equilibria — II: Common Ion Effect and Hydrolysis

279

In a saturated solution at ordinary temperatures and pressures the concen-
tration of H 2 S is 0.10 mole/liter. Substituting this value for the [H 2 S] in

we obtain

[H+]» [S»]
[H 2 S]

= 1.1 X 10- 21

(1) [H+] 2 * [S 2 -] = 1.1 X 10* 22

Thus in a saturated solution of H 2 S, the ion product is 1.1 X 10* 22 . The
concentration of the S 2 ' ion can be controlled by varying the concentration
of the H+ ion also present in the solution. Addition of a strong acid to a
saturated solution of H 2 S will drive the H 2 S equilibrium to the left and the
[S 2 "] will decrease correspondingly. If a base is added, H+ ion will be re-
moved, the equilibrium will shift to the right, and the [S 2 -] will increase.
This manipulation of the [S 2 “] is the basis of the classical method of separa-
tion of metal ions in qualitative analysis by precipitation of their sulfides
(Chapter 50).

Example 2: Calculate the concentration of sulfide ion in a saturated solution
of (a) H 2 S containing 0.3M H+ ion; (b) H 2 S containing 6.0 x 10“ 4 M OH- ion.

Solution: (a) For a saturated solution of H 2 S:

[H+] 2 [S 2 -] = 1.1 X IO -22

(b)

[s 2 -]

[H+]

LI X 10-22
(0.3) 2

= 1.2 X 10- 21 M

[OH-]

1 x 10- 14
6.0 x I (H

= 1.67 x 10-uAf

[S 2 -]

1.1 X 10-22

(1.67 x 10-n)2

= 0.40M

2. Hydrolysis. An aqueous solution of sodium chloride is neutral and has
a pH of 7. However the solutions of other apparently neutral salts have values
of pH other than 7. An aqueous solution of sodium acetate is slightly basic
whereas one of ammonium chloride is slightly acidic. This is due to a reaction
between the ions of the salt and water, known as hydrolysis, literally "cleavage
by water.”

In an aqueous solution of sodium acetate, there are present Na+ and Ac
ions from die NaAc, and H+ and OH“ ions from the slight ionization of
water. The question arises,. is it possible for any of these ions to combine?
Certainly there can be no combination between the Na+ and H+ ions, nor
between the Ac and OH' ions, because they have like charges. Nor can we
expect combination between the Na+ and OH" ions, for the reason that
sodium hydroxide is a strong electrolyte, completely dissociated into ions.
However the combination between H + and Ac ions will take place to an
appreciable extent because acetic acid is a weak acid.

The species which are present and the reactions which, take place in an
aqueous solution of sodium acetate can be formulated as follows:

280

Ionic Equilibria — //: Common loti Effect and Hydrolysis

or simply

(2)

Na+

Ac~

h 3 o

OH-

H+

HAc

Ac +

H a O

HAc +

GH-

base-

acid,

base.

Sodium ions, Na+, are omitted from Equation 2 since they do not enter
the reaction and are merely spectator ions. As a result of the combination
of the H+ and Ac ions, the hydrogen ion concentration is decreased and
more water ionizes to maintain its ionic equilibrium. The concentration
of the OH' ion increases in such quantity required by the ion product,
[H+] [OH'] = 10 w , and the solution is basic. Essentially the hydrolysis of
sodium acetate is the reaction of its negative ion with water. Other salts
which exhibit similar behavior are sodium cyanide, NaCN; sodium amide,
NaNH 3 ; and sodium carbonate, Na 3 C0 3 .

The hydrolysis of ammonium chloride, NH.CI, yields an acidic solution
because of the increased concentration of hydrogen ions produced.

(3) NH«+ + HjO NH, -f H a O+

acidj base., base.. acid 3

Hydrated metal ions also undergo hydrolysis. A solution of ZnCl 2 is acidic
because of the hydrolysis of the zinc ion, Zn(H»0)* a+ .

(4) Zn(H 3 0)«*+ + H a O Zn(H 3 0),(OH) + -f H s O+

A salt of both a strong acid and a strong base, e.g., NaCl, does not hydrolyze
because little or no reaction takes place between the ions of the salt and
water.

Hydrolysis is the reverse of neutralization. It is the reaction between
the ion of a salt and water to produce a weak acid and OH" ion, or a weak
base and H+ ion. The hydrolysis of Ac ion is the reverse of the neutralization
of HAc by an ionic hydroxide, e.g., NaOH, while the hydrolysis of NH 4 +
is the reverse of the neutralization df NH, by a strong acid, e.g., HC1. The
extent to which a hydrolysis reaction proceeds, or the degree of hydrolysis,
depends upon the weakness of the molecular acid or base formed. Since
hydrolysis results in an ionic equilibrium, a hydrolysis constant can be cal-
culated and quantitative calculations can be done in the same manner as
were those in die previous chapter.

As an example, the hydrolysis of the acetate ion is typical of that for any
negative ion. For the two simultaneous equilibria in an aqueous solution of
Ac ion, we may write the simplified ionic equation:

H*0 ;=± H+ -f OH-
Ac- -f H+ HAc

for which K w = [H+] [OH-]

for which

[HAc]
[Ac] [H+]

Khac

Ionic Equilibria — II: Common Ion Effect and Hydrolysis

281

Solving both expressions for [H+] and equating the values.

whence

( 5 )

[H+] =

[OH*]

_ [HAc]
[Ac]

'•HAc

[HAc] [OH']
[Ac]

K w

‘■HAc

= Kh

Kh is the hydrolysis constant; it is an equilibrium constant for the hydrolysis
reaction. Equation 2. In general, the hydrolysis constant is equal to the ratio
of K w to the ionization constant for the weak electrolyte produced in the
hydrolysis reaction. For the hydrolysis of NH 4 +, Equation 3, there can
be derived

( 6 )

[NH 3 ] [H 3 0+] _ v _ K w

[NH 4 +] h K N h,

Generally the ionization constant of the weak electrolyte is larger than K w ,
so that KJ, is small, and so too the extent of hydrolysis is small.

Example 3: For a 0.50 M sodium acetate solution, calculate (a)' the hydroxide
ion concentration; (b) the percentage hydrolysis; (c) the hydrogen ion concentra-
tion; (d) the pH.

Solution: (a) The hydrolysis reaction is Acr + H 2 0 HAc+ OH - , for which

K h

K„

Kw _ 1 X 10- 14

K H ac 1-8 x 10-*

[HAc] [OH-]

5.6 x 10-M

= 5.6 x 10-10

Ac- HAo OH-

Initial concentration, mole/liter 0.50 0 0

Let the [OH*] formed equal x mole/liter; then [HAc] is also x mole/liter and
the [Ac*] is (0.5 - r) mole/liter.

Concentration at equilibrium, mole/ liter 0.50-x x x

Substituting these values in the expression for K h ,

(ffi L = 5.6 x 10-i«

0.5 - x

Again since the value of K h is small, the extent of hydrolysis will be slight; the
value of x will be small in comparison to 0.50 and can be omitted from the
denominator. The resulting expression can be more easily solved for x.

0.50

5= 5.6 x lO-io

and x = 1.7 x 10*5

The concentration of the OH~ ion is 1.7 x 10* 5 mole/liter.

(b) the percentage hydrolysis or the degree of hydrolysis is the number of
moles of Ac* ion which hydrolyze divided by the initial number of moles of Ac*

282

Ionic Equilibria — H: Common Ion Effect and Hydrolysis

ion. Since the number of moles of OH- ion produced is equal to the number of moles
of Ac- ion which hydrolyze, then

degree of hydrolysis

[QHj

Initial [Aei

L 7 x IQ* mole/ liter
0.5 mole/ liter

- 3.4 X 10-5

= 3.4 x 10-*%

(d) pH = -log [H+i = -log <6.0 x lO-i") = 9.22

3. Buffer Solutions. A small amount of acid or base added to pure water
will change the pH markedly. If a drop of concentrated HC1 were added
to a liter of water the pH would change from 7.0 to approximately 4.0, a
one thousandfold change in the {H + J. For this reason it is difficult to pre-
serve the pH of an aqueous solution. If stored in a glass container, enough
alkali may be dissolved from the glass to raise the pH of a solution. If a
solution is exposed to air, small quantities of CO a dissolve and lower the pH
due to the formation of carbonic acid.

For many purposes it is desirable to prepare a solution which maintains
a constant pH. For example, a solution containing both acetic acid and
sodium acetate has a pH of about 5, and if a small amount of either add or
base were added the pH would change only slightly. This behavior can be
explained as follows. The acetic acid equilibrium is: HAc + Ac'.

If acid is added to this solution, most of the hydrogen ions of the additional
acid unite with the large reserve of acetate ion available from the sodium
acetate to form molecular acetic acid. Thus there is little increase in the
concentration of the hydrogen ion and the pH remains practically constant.
If a base is added to the solution, the basic ions such as OH' combine with
the hydrogen ions present but simultaneously more of the acetic acid ionizes
to maintain the concentration of the hydrogen ion at nearly its original value.
Such a solution, which maintains a relatively constant pH even upon addi-
tion of add or base, is known as a buffer solution.

A buffer solution can be prepared from (A) a weak acid and a salt of
the negative ion of the acid, e.g., HAc and NaAc, (B) a weak base and a
salt of the positive ion of the base, e.g., NH» (or NH.OH) and NH«C1, (C)
a mixture of salts, e.g., Na 3 P0 4 and Na*HPO«. A buffer solution does not
have unlimited capacity to maintain a constant pH. If large amounts of
add or base are added, the concentrations of the buffer mixture may be
insuffident to combine with the additional add or base and so the pH will
change. For each buffer mixture there is a pH range over which it ads
most effectively. The effective pH interval for buffered HAc — NaAc mixtures
is between 3.7 and 5.7. The actual pH of such a buffer mixture depends on
the relative concentrations of die HAc and the Ac ion. Its value can be cal-
culated as described in Example 1. Indeed the solution of 0.10M HAc and
0.10M NaAc is a buffer solution having a pH of 4.7. The pH of Wood is
maintained at about 7.4 prindpally by the action of the buffer mixture,
HjCOs (or CO* and H t O) and the bicarbonate ion, HCO,\

Ionic Equilibria — II: Common Ion Effect and Hydrolysis

283

QUESTIONS

1. What is meant by the common ion effect? Illustrate its operation.

2. The pH of a 0.01 M solution of HCN is 5.7. How will the pH be affected

by the addition of (a) NaCl (b) NaCN (c) HC1?

3. Define and illustrate hydrolysis. What is meant by a hydrolysis constant?

4. Formulate the hydrolysis of (a) NaCN (b) Na 2 S (c) NaNO a . Would the solu-
tions of these salts be acidic, neutral, or basic?

5. When acetic aqid is “neutralized” with sodium hydroxide the solution at
the end point has a pH of abt>ut 9. Account for the basicity of the solution.

6. Which solution is more basic, O.IM NaCN or 0.1M NaHCO s ? Explain.

7. Derive an expression for the hydrolysis constant, K h , for the NH 4 +* ion.

8. To one liter of water, 0.1 mole of HC1 and 0.1 mole of KCN are added. As-
suming the total volume remains one liter, what species are present and in
what concentrations?

9. Define and illustrate what is meant by a buffer mixture. How does such a
mixture stabilize the pH of a solution?

10. A solution of 0.010M HC1 is saturated with H 2 S. Calculate the concentration
of S 2 * ion present.

11. What is the concentration of OH* ion and the pH of a 0.050M solution of
NH S which also contains 0.10 mole of NH 4 C1 per liter?

Ans: 9.0 x 10-0; 8.9

12. (a) Calculate the pH of a solution prepared by mixing 100 ml of 0.020M

HAc and 100 ml of 0.020M NaAc. (b) If the resulting solution were diluted
tenfold what- would the pH be? Am: (a) 4.7

13. Calculate the percent hydrolysis and the pH of (a) 0.10M NH 4 C1 (b) 0.01M

Na 2 CO s . Ans: (a) 0.0074%; 5.1

14. For the hydrolysis of the Zn(H 2 0) 4 2 + ion, K h = 2.3 x 10 10 . Calculate the

pH of a 0.010M solution of ZnCl 2 . Am: 5.83

15. Estimate the pH of a buffer solution made by dissolving 3.0 millimoles of
NaCN and 4.0 millimoles of HCN in 200 ml of solution.

16. What should be the ratio of concentrations of HAc and NaAc to make a
buffer solution having a pH of 4.00?

17. What substances, and in what concentrations, would you use to prepare a
buffer solution of pH equal to 8.50?

Ionic Equilibria— ITT
Precipitates and
Their Dissolution

X. Solubility Product, A saturated solution has been defined as one
in which there is an equilibrium between excess undissolved solute and
dissolved solute. In a saturated solution of a solid ionic solute there is an
equilibrium between excess solid solute and ions of the solute in solution.
In such a solution we shall assume that there are no molecules but that
only ions are present. This assumption is justified on two counts: first,
we shall consider only ionic solutes, substances which are already com-
pletely ionized in the solid state and for which the process of solution merely
serves to separate the ions, and second, we shall deal only with solutes
which are very slightly soluble so that their saturated solutions are ex-
tremely dilute. In such extreme dilution any dissolved solute can reason-
ably be assumed to be completely dissociated.

In a saturated solution of silver chloride, AgCl, a sparingly soluble ionic
salt, the following equilibrium exists, as shown in Figure 22.1.

(1) AgCl(s) Ag+ + Clr

For this equilibrium we can write an equilibrium constant expression,

( 2 )

k = FAg^i tea-]

1 [AgCl(s)]

Inasmuch as a solid can react only upon its surface, and for any given solid,
the number of molecules or ions per' square centimeter of surface area is a
fixed value, the "concentration” of a solid is constant. Thus the [AgCl(s)]
is constant and by cross-multiplying in Equation 2 its value can be included
in Ki, whence

(3) Ki[AgCI(s)] = K sp = [Ag+] [Cl“]

Such a value of the equilibrium constant, K, p , is known as a solubility
product constant ,

Ionic Equilibria — 111: Precipitates and Their Dissolution

285

Figure 22.1. Equilibrium in a Saturated
Solution of AgCL

AgCl is a slightly soluble, ionic solid. In a
saturated solution of AgCl the rate at which
the ions, Ag+ and Cl”, leave the solid equals
the rate at which they crystallize upon the
solid AgCl surface.

The solubility product constant may be defined as the product of the
ionic concentrations, in moles per liter, each concentration being raised to a
power which is equal to the number of moles of the ion produced by the
dissociation of one mole of the salt Thus for silver sulfide, Ag 2 S, which
dissociates

(4) Ag 2 S(s) 2 Ag+ + S 2 -; K 8p - [A g^] 2 [S 2 -]

(5) For Fe(OH) s ($) Fe 3 + + 3 OH“; K sp = [Fe*+] [QH“] 3

2. Significance of the Solubility Product. For an ionic solute, the value
of Kg P is a number which represents the ion product of the concentrations
of the ions which can exist in a saturated solution . For AgCl, the value of
the K ap at 25°C is 1.56 X 10" 10 . If we attempt to mix in a solution a concen-
tration of silver ion, Ag+, and a concentration of chloride ion, Cl", such
that the product of these concentrations exceeds the value of the K 8 p, then
AgCl wifi precipitate-in such amount that the product of the remaining
concentrations of Ag+ and Cl“ ions in the saturated solutions will just equal
the value of the K, p . A solution of AgCl in which the ion product of Ag+
and Cl" ions is less than the value of the Ka P is unsaturated and no precipi-
tate of AgCl will form.

In a saturated solution of AgCl alone in water, the [Ag+] equals the
[Cl*] because the dissociation of AgCl yields Ag+ ion and Cl" ion in a 1
to 1 ratio. In such a solution, if the [Ag+] equals x mole/liter, then the
[Ch] also equals x mole/liter, and by Equation 3,

1.56 X 10- 10 = [Ag+] [Ch] = (%)(*) = x 2
and x = 1.25 X 10~ 5

that is, the [Ag+] = [C1-] = 1.25 X 10' 5 mole/liter

It is not a necessary condition, however, that the two ions be present
in equal concentration. For example, if the [Ag+] is 1.56 X 10* 4 mole/liter,
then the [Cl*] which can exist in equilibrium with it is 1.00 X 10* 6 mole/liter,
because only then will the product of the two ionic concentrations equal the
value of the K, p , or 1.56 X 10' 10

286

Ionic Equilibria— III: Precipitates and Their Dissolution

In theory, the K ap is closely related to the ion product of water, K w .
Both are specific applications of the equilibrium constant. If we attempt
to mix in a solution such concentrations of H+ and of OH" so that their
ion product exceeds K w , then molecular water will “precipitate ” Further-
more the [H+] and [OH*] concentrations need not be equal, except in
pure water.

Example 1: At 25°C equal volumes of 0.0010M Ag+ ion, e.g., O.OOlOAf
AgNOg, and 0.0010M Cb ion, e.g., 0.0010M NaCl, are mixed.

(a) Will a precipitate of AgCl form?

(b) If a precipitate forms, how much AgCl will precipitate?

AgCl with 0.00050M Ag+ ion?

Solution: (a) If no precipitate occurs, both the Ag+ and the Cl" concentra-
tions would be reduced to on8-half their initial values merely because the mixing
of equal volumes doubles the volume of the solution in which the initial quantities
of Ag+ and Cl' axe dissolved. Thus, assuming no precipitation of AgCl, the ionic
concentrations would be 0.0005M Ag+ and 0.0005M Cb. The corresponding
ion product would be

(5 X 10-4) (5 x 10-4) = 2.5 x 10- 7

Since this number exceeds the value of the K gp , 1.56 X HH°, precipitation of
AgCl will occur.

(b) Let x equal the moles per liter of Ag+ which precipitate; the number of
moles per liter of Cl* which must precipitate is also x, and the number of moles
of AgCl which precipitates is x. The concentration of Ag+ remaining in the
saturated solution is (0.0005 — x), and the Ch concentration is (0.0005 — x). The
product of these concentrations must equal the K gp .

(0.0005 - x) (0.0005 - x) = 1.56 x MH«

Solving, x = 0.000475,

The [Ag+] and the [Cb] remaining in solution is 0.000125 mole/liter and
the number of moles of AgCl which precipitate is 0.000475 moles.

K sp = 1.56 x xo-io = [Agf] [Cb] = (5 x 1(H) [Cb]

1.56 x 10- W

[Cb] =; — g ~ — = 3 * 1 x 10 ~ 7 mole/ liter

Note that the [Ag+ ] Is not equal to the [Cb] in this case. The result indicates
that, if the [Cb] is greater than 3.1 X 10“ 7 M, then AgCl will precipitate. If the
[Cl*] is less than 3.1 x 10“ 7 M, no precipitate will form.

A common analytical procedure is the selective precipitation of metal
sulfides by hydrogen sulfide. On page 279 we saw that the S 2_ ion concen-
tration in a saturated solution of H 2 S could be controlled by fixing the H+
ion concentration. Metal sulfides are generally insoluble but their solubili-
ties vary over a wide range; the K sp for copper (II) sulfide, CuS, is 3.6 X 1(H 2
whereas that for manganese (II) sulfide, MnS, is 1.4 X 1(H 5 . A given S 2 -
ion concentration may be sufficient to precipitate the very insoluble metal

Ionic Equilibria — 111: Precipitates and Their Dissolution

287

sulfides but not to precipitate the relatively soluble ones. Thereby their
separation can be accomplished.

Example 2: For CuS the K sp is 3.6 x 10-* 2 ; for MnS the K sp is 1.4 x 10- 1B . A
solution containing 0.00010M Cu 2 + ion, 0.10M Mn 2 + ion, and 0.0010M HC1
is saturated with H 2 S. Will CuS and MnS precipitate?

Solution: In 0.0010M HC1, the [H+] is 1.0 X 10 - 3 M, and the [S 2 "] is

[s 2 -]

1.1 X 10-22

[H+P

1.] x 10-22
(1.0 X HH)2

1.1 x 10- 16 M

The ion product of CuS is [Cu 2 +] [S 2 -].

This equals (1.0 x 10-4) (u x 10-ie) - 1.1 x 10* 20 .

This number exceeds the K sp for CuS, 3.6 X 10' 42 , so that CuS will precipitate.
The ion product of MnS is [Mn 2 +] [S 2 ”].

This equals (1.0 x 10*) (1.1 x 10*«) = 1.1 x 10* 7 .

This number is less than the K sp for MnS, 1.4 X 10“ 15 , so that MnS will not
precipitate.

Note that the concentration of Mn 2 + ion is 1,000 times greater than that of
the Cu 2 +, yet MnS does not precipitate whereas CuS does.

To cause precipitation of the MnS the S 2_ concentration must be at least

1.4 X 10*5

0.10

or 1.4 X 10* 4 M.

MnS can be precipitated by making the solution sufficiently basic. The [H+]
must be decreased to a value such that the [S 2 “] in equilibrium with it is at least
1.4 X IQ* 4 M.

3. Calculation of the Value of the K gp . The value of the solubility
product constant can be determined from the solubility of the compound
concerned. Note that the term “solubility" means the concentration in a
saturated solution. The following examples illustrate how the value of the
K sp is calculated.

Example 3: At 25 °C the solubility of silver chloride is 0.00180 grams per liter.
Calculate the K sp for AgCl at 25 °C.

Solution: For AgCl, K gp = [Ag+] [Cl“].

In this expression, the concentrations of Ag+ and Cl* ions must be in moles
per liter. Since the GMW of AgCl is 143.3 g/mole, the molar solubility of AgCl is
0.00180 g/Iiter

143.3 g/mole

1.25 X 10-5 mole/liter.

Because the dissociation of one mole of AgCl produces one mole of Ag+ ion
and one mole of Cl“ ion,. the [Ag+] and [Cl*] will also be 1.25 x 10* 5 mole/liter.
Hence at 25 °C,

K sp = (1.25 x 10-5) (i.25 x 10-5) _ 1.56 x 10“ 10
It is evident that, if the solubility of AgCl varies with temperature, the value

288

Ionic Equilibria — III: Precipitates and Their Dissolution

Example 4: At 25°C, the solubility of silver chromate, Ag 2 Cr0 4 , is 1.3 X 10-4
mole/liter. Calculate the K gp .

Solution: The dissociation of silver chromate is: Ag 2 Cr0 4 (s) 2 Ag+ -f Cr0 4 2 -
for which K gp = [Ag+] 2 [Cr0 4 2 -].

Because the dissociation of one mole of Ag 2 Cr0 4 yields two moles of Ag+ ion
and one mole of Cr0 4 2 ~ ion, the Cr0 4 2 - ion is related to the Ag 2 Cr0 4 in a 1 to 1
ratio and its concentration will equal the molar solubility of the Ag 2 Cr0 4 . The
[Ag+] will be twice that of the Cr0 4 2 * ion, or twice the molar solubility of the
Ag 2 Cr0 4 .

Hence [Cr0 4 2 -] = 1.3 X 10" 4 mole/liter and [Ag+] = 2(1.3 x 10-*) —
2.6 x 10*4 mole/liter.

The K sp = (2.6 x 10* 4 ) 2 (1.3 x 10- 4 ) = 8.8 x 1(H 2 .

Appendix VII lists the values of the solubility, product constants for a
number bf compounds. Solubility product theory, namely, that there is
such a value as the K ap which remains constant, is applicable only to very
slightly soluble ionic solutes. For this reason, all K 8p values are very small.
The theory does not apply to “soluble” compounds such as NaCl, nor even
to slightly soluble substances if a high concentration of other substances
is present. Though K sp values are temperature-dependent, the values given
in Appendix VII are applicable at ordinary temperatures.

4. Calculation of Solubility from the K 8t> . Conversely a knowledge of
the K sp value for a solute enables one to calculate the solubility of the
solute. On page 285, from the*K gp for AgCl we determined the concen-
trations of the Ag+ and C1‘ ions that were present in a saturated solution
of AgCl in pure water. Because of the 1 to 1 relationship between the
AgCl and the Ag+, the molar concentration of the Ag+ ion equals the
molar concentration of the AgCl which dissolved to produce the saturated
solution, that is, the solubility of the AgCl. Hence the solubility of AgCl
= [Ag*] = 1.25 X 10~ 5 mole/liter = (1.25 X 10” 5 ) mole/liter X
143.3 g/mole = 0.00180 g/liter. For Ag2Cr0 4 , the solubility equals the con-
centration of the Cr0 4 2 ’' ion, or one-half the concentration of the Ag* ion ;
in the saturated solution.

Example 5; At 20° C, the solubility product constant of magnesium hydroxide,
Mg(OH) 2 , is 3.2 X 10* 11 . Calculate the solubility of Mg(OH) 2 .

Solution: The dissociation of magnesium hydroxide is:

Mg(OH) 2 (*) Mg 2 * + 2 OH”.

Let x equal the moles per liter of Mg(OH) 2 that dissolve. The dissociation
equation states that x moles of Mg(OH) 2 yield x moles of Mg 2 + ion and 2x moles
of OH‘ ion. Hence

[Mg 2 *] = x mole/liter

[OH’ ] = 2x mole/liter

The K, p = [Mg*+] [OH’] 2 = (x)(2s) 2 = 4 = 3.2 x 10-“

Solving, x = 2.0 x 10*4.

At 20°C, the solubility of Mg(OH) 2 is 2.0 x 10* 4 mole/liter. In a saturated
solution of Mg(OH)* the [Mg 2 *-] is 2.0 X 1(H mole/Hter and the [OH~] is
4.0 X 10- 4 mole/liter.

Ionic Equilibria — 111: Precipitates and Their Dissolution

In solving this problem, any OH*' ion contributed by the ionization of water
has been neglected. In pure water, the [Off-] is 1 x 10* 7 mole/liter and this
value can be disregarded in comparison to 4.0 x 10* 4 mole/liter. Further, the
presence of the OH* ion from the Mg(OH) 2 will repress the ionization of water
by the common ion effect so that its OH- ion contribution will be even less than
1 X lO* mole/liter.

Let us now consider the case where a solute is dissolved, not in pure
water, but in a solution which already contains an appreciable concentra-
tion of an ion common to one of the ions in the solute.

Example 6: Calculate the solubility of Mg(OH) 2 in 0.010M NaOH at 20°C.

Solution: Before starting the problem we might predict the qualitative effect
of the NaOH. From Le Chatelier s principle we should expect that the relatively
high OH - ion concentration, Q.OlOiVf from the completely dissociated NaOH, will
shift the magnesium hydroxide equilibrium in the direction of Mg(OH) 2 . This
would reduce the [Mg 2 +]. The solubility of Mg(OH) 2 , which is equal to the
[Mg 2 +], should thus be less than that in pure water.

As in Example 5, let x equal the moles per liter of Mg(OH) 2 that dissolve. By
the dissociation of Mg(OH) 2 there are formed x mole/liter of Mg 2 + ion and 2x
mole/liter of OH* ion.

The ionic concentrations are:

[Mg 2 +] = x mole/liter.

[OH’ ] = (2x + 0.01) mole/liter, inasmuch as the O.OlOAf NaOH

initially present yields a OH* of 0.010 mole/liter to which the 2x mole/liter
from the dissociation of Mg(OH) 2 must be added.

K sp = [Mg 2 +] [OH**] 2 = (x)(2x + 0.010) 2 ='3.2 x HH 1

Solving, x = 3.2 x 10* 7 .

Solving for x is simplified by neglecting the small value of 2x, which turns out to
be 6.4 X 1Q“ 7 , in comparison with 0.010. The equation to be solved is thereby:

(x)(0.010) 2 = 3.2 X 10*1,

Thus, in the O.OIOM solution of NaOH, the [Mg 2 +] is 3.2 x 10- 7 mole/liter
and the [OH-] is 0.010 mole/liter. The O.OIOM OH“ ion concentration reduces
the solubility of Mg(OH) 2 from 2.0 x 10* 4 moles per liter to 3.2 x 10* 7 moles
per liter, approximately a thousandfold.

The reduction in solubility illustrated by Example 6 is another example
of the common ion effect. A practical application of this is to ensure the
“completeness” of precipitation in chemical procedures. Example 6 indi-
cates that if we wish to precipitate the maximum quantity of Mg 2 + ion as
the hydroxide, Mg(OH) 2 , an excess of OH' ion should be added beyond
that quantity required to initiate precipitation. This would leave in solu-
tion only a very small concentration of Mg 2 + ion, which may be beyond the
limit of detection by ordinary analytical methods. Truly complete precipi-
tation can never be accomplished if an equilibrium exists in solutiob. Some
Mg 2 +, albeit an insignificant amount, must still be present in the saturated
solution to satisfy the K sp for Mg(OH) 2 , For most quantitative purposes.

290

Ionic Equilibria— *111: Precipitates and Their Dissolution

if 99.9% or more of a substance has been precipitated the chemist con-
siders it to be ‘completely” precipitated.

5. Reactions Between Ions. Till now we have considered various aspects
of equilibrium in ionic solutions. Perhaps we should digress for a bit to
see how ionic reactions could be made to proceed to “completion.” Let us
consider what reaction might occur if dilute solutions of potassium chloride,
KC1, and sodium nitrate, NaNO s , were mixed. One might be tempted to
say that the following reaction would take place,

KC1 + NaNO s -» KNO s + NaCl

but this would be incorrect. Both KC1 and NaN0 3 are strong electrolytes,
completely dissociated so that the initial mixture contains the four ions, K+,
C1-, Na+, and N0 3 ‘, in addition to the slight concentrations of H+ and OH"
from the ionization of water. Except for the formation of water, none of these
ions has any tendency to combine with any other and to form a molecular com-
pound. The sole result of mixing the KC1 and NaN0 3 solutions is to pro-
duce a mixture of ions. No chemical reaction, in the ordinary sense, takes
place.

Any chemical reaction will go to completion if one or more of the products
is removed from the equilibrium mixture. Thereby the reverse reaction
cannot take place and the forward reaction, being the only one, proceeds
to completion. A reaction between ions will go to completion if one of
the products is such that ions are removed from the equilibrium mixture.
Such will be the case if one of the products is (A) an un-ionized or wealdy
ionized molecule, (B) a gas, or (C) a precipitate. The following three equa-
tions illustrate these points.

(6) Na+ + OH- + H+ + Or -> Na* + Or + H 2 0

(7) 2 Na+ + C0 3 2 - + 2 H* + 2 Or 2 Na+ + 2 Ch + H a O + C0 2 (g)

(8) Ag+ + NOa- + H* + Ch-* H* + NO s - + AgCl(s)

In all three cases ions are removed from solutions, either physically as a
gas which escapes the reaction vessel or chemically by being locked up in
an unionized molecule or in a precipitate.

6. The Dissolution, of Precipitates. Precipitation and the subsequent
redissolving of a precipitate are among the most common procedures in
analytical separations. By the addition of OH“ to a solution of Mg 2 *,
Mg(OH) 2 can be precipitated. If a solution of HCi is added to this precipitate
of Mg(OH) 2 , the precipitate dissolves. This dissolution of the precipitate
must mean that the ionic concentrations of the Mg 2 * and OH" now present
are such that their ion product, [Mg 2 *] [OH"] 2 , no longer exceeds the
value of the K sp for Mg(OH) 2 . The addition of the HCI therefore must
have reduced either or both the [Mg 2 *] or [OH"]. In this case it is the
OH“. Addition of a high concentration of H+, from HCI or any other strong
acid, to the equilibrium existing in a saturated solution of Mg(OH) 2 , causes
a reaction of the H* and the OH" ions to form water. This can be formu-
lated as follows;

Ionic Equilibria — III: Precipitates and Their Dissolution

291

( 9 )

Mg(OH) 2 (s)

Mg 2 + + OH-

H+

h 2 o

The removal of the OH* ion causes the Mg(OH) 2 equilibrium to shift to
the right, in the direction that will produce more OH’ ion and incidentally
more Mg 2 + ion. The OH* ions are removed as rapidly as they are pro-
duced, however, by the formation of weakly ionized water. If sufficient
H+ ion is added, die Mg(OH) 2 equilibrium is displaced to such an extent
that all the Mg(OH) 2 dissolves. Even though the [Mg 2 +] increases, the
OH* is reduced to such value that the ion product, [Mg 2 +] [OH*] 2 , falls
below the value of the K sp , In all cases of the dissolution of a precipitate,
the ion product is reduced to a value below that of the K gp for the pre-
cipitate by the addition of a reagent, which decreases the concentration of
one or more of the ions involved in the equilibrium.

The dissolution of Mg(OH) 2 can be explained from another point of
view. Equation 9 indicates that there is a competition for the OH* ions
between the Mg 2+ ions and the H+ ions. The H+ ion is a stronger acid
than is the Mg 2 + ion, that is, it has a stronger affinity for OH* ion, and so
wins out. The K sp for Mg(OH) 2 , 3.2 X 10" n , measures the tendency of
Mg (OH) 2 to yield OH’ ions; the K w for water, I X 10* 14 , measures the
tendency of H 2 0 to yield OH" ions. In Example 5 we found that, in a solu-
tion of Mg(OH) 2 , the [OH~] is 4.0 X 10" 4 M; in pure water it is 1 X 10* 7 M.
Hence water has a lesser tendency to yield OH* ion than does a satu-
rated solution of Mg(OH) 2 . Conversely, H+ ions, particularly in the high
concentration that exists in an acid solution, have a greater tendency to
combine with OH’ ions and form H 2 0 than do Mg 2 + ions to form Mg(OH) 2 .
This combination of the ions, H+ and OH’, displaces the over-all equilibrium
towards the formation of water with the result that the Mg(OH) 2 dissolves.

7. Complex Ions, The addition of ammonia, NH 3 , to a precipitate of
AgCl dissolves the precipitate. In this case Ag+ ion is removed from the
saturated solution by the formation of a complex ion, Ag(NH 3 ) 2 +. This
ion is called the silver-ammonia complex ion, or the diammine silver ion;
the term “ammine” refers to the combined ammonia molecules. Though per-
haps similar in appearance, the Ag+ ion and the Ag(NH 3 ) 2 + ion are quite
distinct entities.

The equilibria involved in the addition of NH 3 to solid AgCl are:

(10) AgCl(s) Ag+ + Ct

2 NH 3

Ag(NH s ) 2 +

292

Ionic Equilibria — III: Precipitates and Their Dissolution

If sufficient NH S is added, the Ag+ ion concentration is reduced by the
formation of die Ag(NH 3 ) 2 + ion to such an extent that the ion product,
[Ag+] [Ct], is less than the K sp for AgCl and so the AgCl dissolves.

For die equilibrium between a complex ion and its component parts an
equilibrium constant expression can be written. For the dissociation of. the
Ag(NH 3 ) 2 + complex ion,

(11) Ag(NH 3 ) 2 + Ag+ 4 2 NH a

and Ka = where K* = the dissociation constant for the

[Ag(NH 3 ) 2 + ] complex ion

K* is a measure of the stability of the silver ammonia complex ion, Ag(NH 3 ) 2 +,
that is, its tendency to dissociate into Ag+ ion and NH 3 . The value of K*
for the Ag(NH 3 ) 2 + ion is 6.8 X 10~ 8 ; the small value is evidence of the
stability of the ion.

A complex ion may be defined as an ion composed of more than one
atomic species. Under this broad definition the sulfate ion, S0 4 2 “, could be
classed as a complex ion. More precisely the term complex ion is applied
to the case where a metallic ion is combined with a neutral molecule or
another ion; such a molecule or ion is called a ligand . Examples of complex
ion formation are:

(12) Cu 2 + 4* 4 NH 3 ^ Cu(NH 3 ) 4 2 + (the copper-ammonia complex ion) 1

(1§) Ag+ 4 2 CN- Ag(CN) 2 ~ (the silver-cyanide complex ion)

In aqueous solution all simple metal ions are believed to exist as
complex ions containing water molecules; for example,

(14) Cu 2 + 44H 2 0^ Cu(H 2 0) 4 2+

Complex ion formation is quite common but some metal ions, notably
those of the alkali metals and the alkaline earth metals, form no complex
ions other than through combination with H 2 0 molecules in water solu-
tion, Equilibria for the more common complex ions and their dissociation
constants are tabulated in Appendix VIII

In a complex ion, the number of ligands to which the metal ion is bonded
is called the coordination number of the metal ion. In both Ag(NH 3 ) 2 + and
Ag(CN) 2 “ the coordination number of the silver ion is two. In Fe(CN) 8 3 ‘,
the coordination number of the iron is six. Generally, but not infallibly,
the coordination number of a metal ion is twice its oxidation number.
Most of the complex ions listed in Appendix VIII follow this rule. The oxida-
tion number of Ag+ is one and it combines with two ligands. The resultant
charge of a complex ion equals the sum of the charges of its components.
For the silver-ammonia complex ion, the net charge is 41; 41 from the
Ag+ ion plus 0 from the two neutral NH 3 molecules. The Ag+ ion also

See page 501 for a more specific nomenclature for complex ions.

Ionic Equilibria — III: Precipitates and Their Dissolution

293

forms a complex ion with two thiosulfate ions, S 2 0 3 2 i the net charge on
this complex ion is +1 + (2 X -2) = -3, and the formula of the complex ion
is Ag(S 2 0 3 )2 8 “ The structure of complex ions and the nature of their bond-
ing is discussed more fully in the chapter on metals (Chapter 37).

8. Amphoteric Hydroxides. If a small amount of NaOH solution is
added to a solution of Zn 2 + ion, the K 8p of zinc hydroxide is exceeded and a
precipitate of Zn(OH) 2 is formed.

(15) Zn*+ + 2 OH" Zn(OH) 2 (s)

Addition of a strong add to the precipitate of Zn(OH) 2 produces the ex-
pected dissolution, in a manner similar to the dissolution of Mg(OH)^ dis-
cussed in Section 6.

(16) Zn(OH 2 (s) +2H+-> Zn*+ + 2 H 2 0

If we add to the precipitate of Zn(OH) 2 additional NaOH beyond that
amount required to produce its precipitation we should expect a more com-
plete precipitation, in view of the discussion on page 289. However the
addition of excess NaOH gives rise to an unexpected result in that the
Zn(OH) 2 redissolves. This dissolution in excess NaOH is explained by the
formation of a soluble complex ion, the zincate ion, Zn(OH) 4 2_

(17) Zn(OH) 2 (s) + 2 OH'-» Zn(OH) 4 2 -

Such a substance, as Zn(OH) 2 , which dissolves in both strong acid and
in strong base is said to be amphoteric.

In effect, an amphoteric hydroxide can act both as an acid or as a
base. Towards an acid stronger than the amphoteric hydroxide, it acts as
a base; with a base stronger than the amphoteric hydroxide, it acts as
an acid.

(180 Zn(OH)* 2 -^ 2 OH* 4- Zn(OH) 2 (s) + 2 H+ Zn 2 + + 2 H 2 0

acting as an acid to the addition of OH* acting as a base to the addition of H+

The equilibria can be displaced in either direction by the addition of acid
or base.

The equations above present an oversimplified picture of amphoterism
inasmuch as, in aqueous solution the zinc ion is hydrated and exists as
Zn(H 2 0) 4 2 + rather than as the simple Zn 2 + ion. Upon the addition of NaOH,
the H 2 0 molecules are progressively replaced by OH" ions; the Zn(H 2 0) 4 a +
ion acts as an acid in that protons donated by its water molecules com-
bine with the added OH” ions, thereby leaving OH* ligands in the complex
ion. If the Zn(H 2 0) 4 2+ ion is represented structurally as

h 2 o

2 +

294

Ionic Equilibria — III: Precipitates and Their Dissolution

the first stage in the successive replacement of H z O molecules by the ad-
dition of OH - ion is

(19)

HoO

OH 2

h 2 o

/ Zll \

oh 2

2 +

+ OH-

HO OH 2 *

H 2 C> OH 2

+ h 2 0

Further addition of OH“ ion yields Zn(0H) 2 (H 2 O) 2 , Zn(0H) 3 (H 2 0)~
and ultimately Zn(OH) 4 2 "

The equilibria are summarized in the single equation following.

(20)

OH- OH-

Zn(H 2 0) 4 2 + Zn(H 2 0) s (OH)+ Zn(H 2 0) 2 (0H) 2 (s)

H+ H+

OH- OH-

Zn(H 2 0)(0H) 3 Zn(OH) 4 a ”
H+ H+

Note that the addition of two equivalents of OH' ion to the Zn(H 2 0) 4 2 +
ion precipitates the neutral species, Zn(H 2 0) 2 (0H) 2 or dimply Zn(OH) 2 .
Upon addition of acid co Zn(OH) 4 2 ~ the equilibria are displaced to the left
and the reverse of the reaction with OH - takes place.

Other amphoteric hydroxides and their corresponding complex ions are:

Al(OH) 3 , aluminum hydroxide; Al(OH) 4 “ aluminate ion

Cr(OH) 3 , chromium hydroxide; Cr( OH) 4 “, chromite ion

Pb(OH)*, lead hydroxide; Pb(OH) 4 2 “ plumbite ion

Sn(OH) 2 , tin hydroxide; Sn(OH) 4 2 ", stannite ion

QUESTIONS

1. Define and illustrate the term “solubility product.” What is the significance
4 of a K* p ?

2. Write solubility product expressions for (a) CaC0 3 (b) Cu(OH) 2 (c) Ag 2 S
(d) Al(OH) s .

3. For the compounds whose solubilities are given following, calculate the cor-
responding values of K bp : (a) Agl ; 1.2 x 10-3 mo le/liter (b) Mg(OH) 2 ;

1.3 x 10-4 m ole/liter (c) Ba(Cr0 4 ) ; 2.5 x 10-3 gram/liter ; (d) Ag 2 Cr0 4 ;

4.3 x 10‘ 2 gram/liter.

4. From the values of K sp given below, calculate the solubilities of the
compounds: (a) AgBr ; 7.7 X 1(H» (b) PbCl 2 ; 1.7 X 1(H (c) Ag 2 S ; 1.6 X KM*

5. List and illustrate the circumstances under which an ionic reaction will go
to completion.

6. What condition must prevail in order that a precipitate (a) be formed (b) go
into solution?

Ionic Equilibria — III: Precipitates and Their Dissolution 295

7. If the solubility of a compound, AB, increased 10% over a given temperature
interval, how would the value of its K gp be affected?

8. Define and illustrate (a) complex ion (b) ligand (c) coordination number.

9. Formulate a reaction illustrating the dissolution of. a precipitate through the
formation of a complex ion.

10. In terms of solubility product theory, explain the solution of the following
compounds: (a) CaC0 3 in HC1 (b) AgCl in NH 3 (c) Zn(OH) 2 in H 2 S0 4
(d) Zn(OH) 2 in NaOH.

11. Define and illustrate amphoteric hydroxide . How would you prove experi-
mentally that Be(OH) 2 is amphoteric?

12. Explain why the presence of an appreciable quantity of NH 4 C1 prevents the
precipitation of Mg(OH) 2 when NH 3 is added to a solution of Mg 2 + ion,

13. What is the maximum concentration of Pb 2 + ion that can exist in solution
with a Cl* ion concentration of 0.020 mole/liter?

14. A solution contains 0.010M Mg 2 + ion. What is the minimum concentration
of OH" ion required to initiate precipitation of Mg(OH) 2 ?

15. To a solution containing 1 X 10’ 5 M Ag+ ion there is added 1 x 10^M Br~
ion. (a) Will a precipitate of AgBr form (b) if so, what will be the concen-
tration of Ag+ remaining in solution?

16. Calculate the concentration of NH 4 + ion required to prevent precipitation of
Mg(OH) 2 from a solution containing 0.0020M Mg 2 + and 0.010 M NH 3 .

17. A solution containing 0.00020M Cd 2 + ion and 0.10M HC1 is saturated with
H 2 S. (a) Will cadmium sulfide, CdS, precipitate (b) what is the minimum
concentration of HC1 required to prevent precipitation of CdS?

18. The dissociation constant, K d , for the Ag(NH 3 )+ 2 ion is 6.8 x 10^. In a
solution containing 0.10M Ag(NH 3 ) 2 + and 0.0026M NH 3 (a) what is the con-
centration of Ag+ ion (b) what must be the concentration of added Cl" ion
to precipitate AgCl?

19. To one liter of solution are added 0.0030 moles of Ag+ ion and 0.0050
moles of Cl" ion. Neglecting dilution Ja) will AgCl precipitate (b) what con-
centration of NH 3 in this solution would prevent formation of a precipitate?

20. (a) In comparing two compounds, does a lower value of Kg P mean a smaller
solubility? (b) Which is the more soluble compound, Al(OH) 3 or FeS? For
Al(OH) 3 , Kgp = 1.9 x IQ-**; for FeS, K sp = 3.7 x 1<H*.

Oxidation and Reduction

The term oxidation originated prior to our present understanding of atomic
structure when it was defined as a chemical combination with oxygen.
Conversely, reduction was the removal of oxygen or, what was effectively
the same thing, combination with hydrogen. Obviously such definitions are
limited to reactions with oxygen or hydrogen such as C + 0 2 — > C0 2 and
Fe 3 0 4 -f- 4 H 2 -» 3 Fe + 4 H 2 0, but would not include an equation analogous
to the oxidation of carbon such as C + ^ Cl 2 — > CC1 4 . If the reactions of
carbon with oxygen and with chlorine are compared we see that, in both
cases, the carbon atom increased in oxidation number from 0 to 4, whereas
the oxygen and chlorine both decreased in oxidation number, the oxygen
from 0 to -2 and the chlorine from 0 to -1.

More generally, then, oxidation is defined as the loss of electrons by a
chemical entity, and reduction as the gain of electrons. Any chemical re-
action that involves either of these processes must also involve the other
for if , in a reaction some chemical species (atom, molecule, or ion) loses
electrons, some other must gain these very same electrons. Neither an oxida-
tion nor a reduction can occur alone. Both must take place concurrently
and such chemical reactions are therefore called oxidation-reduction reactions
or, in short, redox reactions . If atoms, molecules, or ions lose or gain electrons
in oxidation-reduction reactions, a change in oxidation number must be in-
volved. Algebraically, the loss of electrons leaves a more positively charged
species while the gain of electrons produces a less positively charged en-
tity. Thus oxidation results in an increase in oxidation number and reduction
in a decrease in oxidation number. (Figure 23.1).

Based on the fact that the numbers of electrons lost and gained in redox
reactions must be equal, special techniques for balancing such equations
can be developed. The one we shall consider first in some detail is the
Ion-Electron Method because it is *true to life” in that it deals with the
chemical species that actually exist and with the actual ionic reactions which
take place in redox reactions. It is with these reactions that we shall he
concerned when we study electrochemical theory in the

Oxidation and Reduction

297

1. Ion-Electron Method of Balancing Equations. Let us consider a set
of rules which will make the balancing of redox equations relatively simple.
We shall then prbceed to balance one or two since the best, and perhaps the
only way, to learn the method is to apply it.

+ 6

4 - 5
+ 4
+ 3

Step 1: Obviously we must have an equation to balance,
that is, we must know the "skeleton” equation with all .the
reactants and products so that the problem in balancing is
solely to determine the proper coefficients before each substance.
Let us balance the equation for the reaction between copper
and concentrated nitric acid.

(1) Cu + HNO s Cu(N0 3 ) 2 + N0 2 + H 2 0

+ 2
+ 1

- 1

- 2

- 3

- 4

- 5

- 6

Figure 23 J.
Oxidation
and

Reduction

Step 2: "Parse” the equation for those chemical entities—
atoms, molecules, or ions— having independent existence which
themselves, or within which, some atom undergoes a change in
oxidation number. Then write what we shall call the partial
ionic equations.

In this equation there are three possible reactants— copper,
Cu, hydrogen ion, H+, and nitrate ion, N0 3 “ It is apparent that
the copper atom changed in oxidation number inasmuch as it
is uncombined on the left, and so has an oxidation number of
zero; whereas on the right, as copper ion in Cu(N0 3 ) 2 , it must
have an oxidation number other than zero since it is a part of
a molecule which, as a whole, is electrically neutral, In Cu(N0 3 ) 2
the oxidation number of the copper ion is +2 because the
nitrate ion, N0 3 “ is -1. It is true that the beginner’s lament,
"But how should we know that?” can be answered only with
the reply, "Experience!” but if one observes that the hydrogen
ion is -+• l in HN0 3 then the N0 3 ~ ion alone must be +1.

For copper we can now write a partial ionic equation in which the oxida-
tion numbers are indicated as superscripts.

(2) Cu° -> Cu 2+

This reaction of the copper is obviously an oxidation. There must also
be another partial ionic equation representing a reduction. Almost by the
process of elimination, inasmuch as hydrogen remains unchanged in oxida-
tion number, +1 i* 1 HN0 3 and in H 2 0, it must be the N0 3 ~ ion which has
undergone some change. As part of the Cu(N0 3 ) 2 produced no change oc-
curred to the N0 3 - ion, but in the formation of the N0 2 molecule there was
a change in the oxidation number of a nitrogen atom from +5 to +4.
Granting that we must know, if only by expedience, that N0 2 is a gas and,

29 $ Oxidation and Reduction

ionic equation is NO a ' NO 2 0 . We now have the two partial ionic
equations:

(3) Cu° -» Cu 2 +

(4) NOr N0 2 °

Step 3: Balance the partial ionic equations atomically. Equation 3 is
already balanced since there is but one copper atom on each side. Equation 4
is not balanced; the nitrogen atoms are but there are three oxygen atoms
on the left and only two on the right. We might be tempted to add an
oxygen atom, O, or perhaps Vz0 2 to the right, e.g., NO s * NO a + O, but
this would be incorrect since it is seen from the initial skeleton equation
that no free oxygen atoms or molecules are produced. Oxygen atoms, how-
ever, are a product of the reaction in the form of water and so such oxygen
atoms as are needed to balance the partial ionic equations atomically are ob-
tained in the form of water molecules, one H 2 0 molecule per each oxygen
atom required. Thus we have NG 3 ~-* N0 2 + H a O. Now, however, with the
addition of a H 2 0 molecule, we have introduced another problem in that
the equation is unbalanced with respect to hydrogen atoms, two on the
right and none on the left. Again perhaps we might be tempted to write
H 2 + NO a “ — ► NOo + H 2 0, but this would be improper because no free
hydrogen gas, H 2 , is involved in the reaction. Hydrogen atoms do appear
on the left as part of the acid, HN0 3 , and so the number of hydrogens we
need are obtained in the form of hydrogen ions, H+. Since two are required
the correct partial ionic equation is: 2 H + + N0 3 “ N0 2 + H z O. This

method of obtaining hydrogen and oxygen atoms for balancing partial ionic
equations is very common, particularly in acid solution. In basic or neutral
media, a less frequent situation, these atoms could be acquired in the form
of hydroxide ions, OH“ if such were involved in the equation.

We have now

(7) Cu° Cu 2 +

(8) 2H+ + N0 3 --* N0 2 + H 2 0

Step 4i Balance the partial ionic equations electrically since the net
electric charges on the left and right of each equation, as it now stands,
are unequal. In equation 7, for example, the neutral Cu° is not electrically
balanced by the charged Cu 2 + ion. The medium of exchange in redox re-
actions is the electron and the charges on both sides of the equation are
equalized by the gain or loss of the proper number of electrons, each being
a unit of negative charge for which we will use the symbol e. For Cu° to
become the more positively charged Cu 2 + two electrons must be lost:
Cu° - 2 e Cu 2 +. In purely an algebraic sense we could have balanced
this equation by the gain of two positive charges but this would be im-
proper. The unit of positive charge is the proton and protons exist only
as constituents of atomic nuclei. We have seen these are not involved in
ordinary chemical reactions; only the valence electrons are so involved. Thus
the second partial ionic equation is completed bv ad diner rmn t n

Oxidation and Reduction

299

the left side: 2 H+ + N0 3 ~ 1 e — » NO a -)- H 2 0, since the net charge on

the left of 2 H+ and N0 3 ” is +1, while that of the products, both neutral
molecules, is zero. The two complete partial ionic equations are:

(9) Cu° - 2 e Cu 2 + (oxidation)

(10) 2H+ + NO~ + 1 e N0 2 + H 2 0 (reduction)

By definition Equation 9 is an oxidation and Equation 10 is a reduction.

Step 5: Equalize the numbers of electrons lost and gained. These must
be equal because it is the very same electrons lost in the oxidation of the Cu°
that are gained in the reduction of the N0 3 ~ ion. This step is done by multi-
plying the entire Equation 10 by two. Hence

(11) Cu° - 2 e Cu 2+

(12) 4H+ + 2 NOr + 2 e 2 NO, + 2 H 2 0

Step 6: Add the partial ionic equations, omitting the electrons lost and
gained which cancel because they are equal. Thereby

(13) Cu° + 4 H+ + 2 N0 3 - Cu 2 + + 2 N0 2 + 2 H 2 0

This ionic equation is, in fact, the actual chemical reaction which takes place
when copper and nitric acid react. The ionic equation affords a deeper
insight to what truly happens than does the molecular equation. Observe
that it is not HN0 3 molecules which react but rather H+ and NO a ~ ions,
and these not in equal numbers.

Step 7: Insert the coefficients of the ionic equation before their cor-
responding entities in the original molecular equation. Occasionally a mini-
mum of adjustment may be necessary. Thus a “1” would appear before the
Cu and a “4” before the HNO s for the 4 H+. The "2 NOr” required are
thereby already taken care of by the "4 HNOs/' With a "2” before both
the N0 2 and the H 2 0 the molecular equation is balanced.

(14) Cu + 4 HNO s -> Cu(N0 3 ) 2 + 2 N0 2 + 2 H z O

The apparent discrepancy between the 4 H+ and the 2 NO s ~ in the ionic
equation, though both must be linked as HN0 3 in the molecular equation,
is due to the fact that this method considers only those species which
actually react in the oxidation-reduction process. To offset the electric charge
of the Cu 2-1- in the Cu(N0 3 ) 2 , two NO s “ ions are present merely as spectator
ions to preserve electric neutrality in the solution. They do not enter the
chemical reaction but they must be included in the molecular equation
which is necessary for stoichiometric calculations.

An oxidizing agent is a substance which causes an oxidation process to
take place. In this reaction the Cu is oxidized; since the N0 3 ” ion (or the
HNO a ) brought this about, it is the N0 3 " ion (or the HN0 3 ) which is the
oxidizing agent. Note that the oxidizing agent is itself reduced. Similarly, the
reducing agent > that is, the substance causing reduction to take place, is
the Cu which is itself oxidized.

300

Oxidation and Reduction

Consider now the reaction between potassium permanganate, KMn0 4 ,
and hydrochloric acid, HCL

Step 1: The skeleton equation is:

KMn0 4 + HC1 -» KC1 + MnCl 2 + Cl 2 + H a O

Step 2 : The partial ionic equations are:

Cl” Cl 2
Mn0 4 ” -* Mn 2 +

Step 3: Balance the partial ionic equations atomically:

2 Cl” -> Cl 2

MnOr + 8 H+ Mn 2 + + 4 H 2 0

Step 4: Balance the partial ionic equations electrically:

2 Cl" - 2 e -> Cl 2 (oxidation)

Mn0 4 ” + 8 H + + 5 e -* Mn 2 + -f 4 H 2 0 (reduction)

Step 5: Equalize the numbers of electrons lost and gained; multiply the
first equation by “5” and the second by “2 ”

10 Cl" -10 e 5 Cl 2

2 MnOr + 16 H+ + 10 e -* 2 Mn 2 + + 5 Cl 2 + 8 H 2 Q

Step 6: Add the partial ionic equations:

2 MnOr + 16 H+ + 10 Cl--* 2 Mn 2 + + 5 Cl 2 + 8 H 2 0

Step 7: Insert coefficients into the molecular equation:

2 KMnO, + 16 HC1 -* 2 KC1 + 2 MnCl 2 + 5 Cl 3 + 8 H 2 0

In this case a bit of trimming in the form of a “2” before the KC1 is required.
All other coefficients are obtained from the balanced ionic equation (Step 6)
and the “2 KMnO/’ made necessary the “2 KC 1” to balance the K + ions,
which were spectator ions. The MnOr ion (or the KMn0 4 ) is the oxidizing
agent and the Ch ion (or the HC1) is the reducing agent.

2, Partial Ionic Equations. Many partial ionic equations are of sufficient
importance to warrant being tabulated.

(15) NO,; + 2 H+ + 1 e -* NO a + H 2 0 (in concentrated HNO,)

(16) NO,- + 4 H+ + 3 e -> NO +2 H,0 (in dilute HNO,)

(17) MnO.” + 8 H + -r 5 e — > Mn 2 + + 4 H 2 0 (in acid solution)

(18) MnOr + 2 H 2 0 + 3 e -* Mn0 2 + 4 OH* (in neutral or basic

solution)

(19) CtzO? 2 * -f- 14 H + + 6 e -* 2 Cr 8 ++ 7 H 2 0 (in acid solution)

The reduction of a simple metal ion, M +n . to its elemental state, M, and
the reduction of a nonmetal, X,., to its simple ionic state, X" tt , may be
represented ]oy the following equations, where n is the number of electrons
transferred in the reaction.

Oxidation and Reduction

301

(20) M +n + n e — ■> M° e.g., Cu 2 + -j- 2 e —> Cu°

(21) X + n e -» X- n e.g., S + 2 e -» S 2 ‘

Though the foregoing partial ionic equations have been written as re-
ductions, they can also be used in the reverse sense, as oxidations, for the
balancing of equations by the ion-electron method. With equal correctness
we could have written, Jfor example,

(22) N0 2 + H 2 0-le-> NO," + 2 H+

as an oxidation. The skeleton equation will determine in which direction
a partial ionic equation should be written. This implies that these reactions
can possibly go in either direction, that is, that they are reversible and that
the principles of chemical equilibrium can be applied. In many cases this
is so. The direction in which an ionic reaction actually does proceed and
the manner of determining this direction will be considered in the next
chapter.

3. Oxidation-Number Method of Balancing ReHox Equations. Another
technique for balancing redox equations is the Oxidation-Number Method .
This procedure is similar to the Ion-Electron Method except that we use
only the oxidation numbers of the individual elements within the ions or
molecules that are involved in the electron transfers. In many cases this
method uses fictitious .entities because it is the entire ion or molecule, and
not the component .atoms, which has independent existence and undergoes
the electron change.

Let us balance the reaction between Cu and HNO a (Equation 1) by
the Oxidation-Number Method. The skeleton equation is rewritten below
but now with the oxidation number of each element above its symbol.

0 +1 .5-2 +2+5-2 +4-2 +1-2

Cu + HNO, -+ Cu (N0 3 ) 2 + no 2 + h 2 o

It is evident that the copper and nitrogen atoms have changed in oxidation
number. The partial equations, including the electron transfers, are

Cu° 2 e -* Cu 2 + (oxidation)

N 5 + + 1 e -> N 4 + (reduction)

Thereafter the procedure is much the same as with the Ion-Electron Method.
Each partial ionic equation is multiplied by the proper integer so that
their addition will cancel out the electron transfers. The resulting coefficients
are then placed before the proper molecules in the skeleton molecular
equation.

1 X [Cu° - 2^ Cu 2 +]

2 X [N 5 + + 1 e N 4 + ]

Addition of these equations gives

Cn« 4- 2 NM- Cn2+ 4- 2N 4 +

302

Oxidation and Reduction

Only one of these coefficients has to be changed for the molecular equation.
Since there are four nitrogen atoms on the right side, the coefficient “2”
before the HNO s must be increased to “4.” Thereafter only a minimum of
additional balancing is required— a “2” for the H 2 0, and thus

Cu + 4 HNOa Cu(N 0 3 ) 2 + 2 N0 2 + 2 H z O

In both methods of balancing the numbers of electrons lost and gained axe
identical Since it is not the method of balancing an equation which deter-
mines experimental fact this is necessarily so,

4, Equivalent Weight of an Oxidizing or Reducing Agent. The equivalent
weight of a substance has been defined as that weight which would react
with or produce one atomic weight of hydrogen. From that aspect the
standard of equivalency, in other than redox reactions, is the hydrogen atom.
In oxidation-reduction reactions, the electron is the medium of exchange,
and since one electron is equivalent to one hydrogen ion, the standard of
equivalency is the electron. The equivalent weight of an oxidizing or re-
ducing agent is its formula weight divided by the number of electrons lost
or gained by one formula weight, as given by the balanced ionic equations.
Equivalent weights can be calculated in terms of atoms, molecules, or ions,
whatever be the reacting species. In the examples already considered, the
equivalent weights are:

Cu =

63.5

5 HN0 3 =

36.5

KMn0 4 =

158.0

The equivalent weight of a substance is dependent upon the specific
reaction it undergoes. For dilute HNO a reacting with Cu to produce NO
(instead of NO a ), the equivalent weight would be 1/3 of the formula
weight; for KMn0 4 reacting in basic solution the equivalent weight would
be 1/3 of its formula weight (see Section 2).

5. Oxidation-Reduction Problems. In every case of an oxidation-reduc-
tion reaction, since the number of electrons gained and lost are equal,
the number of equivalents of the actual reacting species , i.e., of the oxidizing
and reducing agents, must be equal. This will also be true for the molecular
formulas in the balanced molecular equation if the atoms exhibiting a change
in oxidation number appear in only a single product. In the reaction between
Cu and HN0 3 , all four N0 3 ~ ions are not reduced. Two are reduced to NO
but two are spectator ions which undergo no reaction but which must
be included in the molecular equation to maintain over-all neutrality by
balancing the +2 charge of the Cu 2+ ion. In the reaction, HBr + H 2 S0 4

Br 2 + SO. -f-, H 2 0, the Br atpm of the reducing agent, Br~ ion, and
the S atom of the oxidizing -agent, H 2 S0 4 , each form only one product.

Example 1: What volume of 2.00 N H 2 S0 4 solution will react with 16,2
grams of HBr?

Solution : 2 HBr + H 2 S0 4 Br 2 + S0 2 +2 H 2 0
% Br- - 2 e Br 2

S0 4 2-‘+ 4.H+ 4- 2 e S0 2 + 2 H z O

Oxidation and Reduction

303

The balanced molecular equation need not be known but the partial ionic equations
and the electron changes must be known so that the number of equivalents per mole
for each reactant can be calculated.

Because two formula weights of HBr lose two electrons, the equivalent weight

of HBr =

formula weight
electron change

80.9

The number ot gram-equivalents of HBr = — = 0.200

80.9 g/ g-equiv

For a solution, the number of gram-equivalents = V (liter) X ]V (normality)

The number of gram-equivalents of HN0 3 = V x 2.00

The numbers of gram-equivalents of HBr and H 2 $0 4 are equal,

0.200 = V x 2.00
V = 0.100 liter

Example 2: What volume of 2.002V H 2 S0 4 solution will react with 60 ml of
0.50N HBr solution?

Solution: We have already considered this type of problem, the reaction be-
tween two solutions of known concentration, in our discussion of titration.

Vh 2 so 4 x N H2S04 = V HBr X N HBr
Vn 2 so 4 x 2.00 N = 60 ml x 0.50 N

V H 2 S0 4 = IS

Example 3: What volume of 0.2 5JV KMn0 4 in acid solution (H 2 S0 4 ) will
react with 5.6 liters of H 2 S gas measured at STP?

Solution : The skeleton equation is:

KMn0 4 4 H 2 S 4 - H 2 S0 4 S 4 MnS0 4 + K 2 S0 4 4 H a O
and the partial ionic equations are:

Mn0 4 - 4 8 H+ 4 5 e -> Mn+ 2 4 4 H 2 0
H 2 S -2 S 4 2H +

Since H 2 S is a gas one mole occupies 22.4 liters at STP.

The number of moles of H 2 S

5.6 liters
22,4 liters/mole

Because the electron change for one mole of H 2 S is two, each mole contains
two gram-equivalents.

The number of gram-equivalents of H 2 S = 2

g-equiv x 5.6 liter
mole 22.4 liter/mole

The number of gram-equivalents of KMn0 4 = V x 0.25

g-equiv

liter

These numbers of gram-equivalents are equal, so

g-equiv

mole

X 5.6

liter/mole

= V x 0.25

g-equiv

liter

V = 2.0 liters

304

Oxidation and Reduction

QUESTIONS

1. Define and illustrate (a) oxidation (b) reduction (c) oxidizing agent (d) re-
ducing agent.

2. Define the equivalent weight of an oxidizing or reducing agent. How does
this differ from that for a reactant that is not oxidized or reduced?

3 . From Section 2, calculate the equivalent weight of (a) concentrated HN0 3
acting as an oxidizing agent (b) dilute HN0 3 acting as an oxidizing agent

4. Balance the following oxidation-reduction equations. In each case label the
oxidation and the reduction and name the oxidizing agent and the reducing
agent.

(a) Cu 4 HN0 3 Cu(N0 8 ) 2 4- NO 4H 2 0

(b) KMnO* 4 KC1 4 H 2 S0 4 K 2 S0 4 4 MnSO, 4 Cl, + H z O

(d) K 2 Cr 2 0 T 4 HCl -» Cl 2 4 KC1 4- CrCl 3 4 H z O

(e) Zn 4- HNO s Zn(N0 8 ) 2 + NH 4 NO s 4 H z O

(f) I 2 + HNO a HI0 3 4- NO, + H 2 0

(g) KMn0 4 4- FeSO* 4 H 2 S0 4 -» KHS0 4 4 Fe 2 (S0 4 ) 3 4 MnS0 4 4 H 2 0

(h) P 4 fiNO s 4 H 2 0 -* H 3 P0 4 4 NO

(i) Crl 3 4 Cl 2 4 KOH -> K 2 Cr0 4 4 KI0 4 4 KC1 4 H a O

(j) CrCla 4 NaC10 s 4 NaOH NaCr0 4 4 NaCl 4 H 2 0

5. Balance the following equations by the Ion-Electron Method. In each case
label the reduction and the reducing agent.

(a) Br 4 SO* 2 - 4 H4 ^ Br 2 4“ S0 2

(b) I, 4 S-A 2 - I- 4 SA 2 -

(d) Mn0 4 4 SO a 2 “ Mn 2 4 4 - SO* 2- (in acid solution)

(e) Fe 4 Fe 8 4 — » Fe 2 4

6 . Balance the following equations by the Oxidation*Number Method. In each
case label the oxidation and the oxidizing agent.

(a) CuO 4 NH a -» Cu 4 N 2 4 H 2 0

(b) NH 3 4 0 2 NO 4 H z O

(d) FeS 2 4 0 2 — > Fe 2 O s 4 S0 2

(e) H a S0 3 4 H 2 0 2 -> H 2 S0 4 4 h^O

7. What weight of KMnO* is required to prepare 1.50 liters of 0.100N solution

to be used as an oxidizing agent in acid solution? Ans: 4.74 g

8 . Whkt volume of 0:10N K 2 Cr 2 0 7 i£ required to oxidize the Fe 2 4 ion in 50 ml
of 0.050N FeS0 4 in acid solution?

9. What volume of 0.10M KMnO* is required to liberate the iodine from

50 ml of 0.10N HI solution? Ans: 10 ml

10. In Problem 4f, what weight of iodine will react with 100 ml of 0.50N

hno 3 ?

Oxidation and Reduction

305

11. In Problem 4c, (a) how many equivalents of HI will react with 1.50 equiva-
lents of H 2 S0 4 (b) how many gram-equivalents of HI will react with 29.4 g
of H 2 S0 4 (c) what volume of 9.20 N H 2 S0 4 will react with 6.4 g of HI?

12. For the reaction: H 2 S 4- HN0 3 — > S + NO -f H 2 0, show that the equivalent
, weights of H 2 S and HNO s have the same ratio as the weights which react

as given by the balanced equation.

13. For the reaction in Problem 12, what volume of H„S gas, taken at STP, will
react with (a) 2.1 g of HN0 3 (b) 150 ml of 0.20 N HN0 3 ?

Electrochemistry I
The Voltaic Cell

The movement of an electrically charged particle, whatever its nature,
constitutes an electric current. The nature of the charged particle varies,
depending upon whether the electrical conduction is metallic, ionic, or
gaseous. In metals, an electric current is the movement of electrons; in solu-
tion and in molten salts, it is the migration of ions, both positive and nega-
tive; whereas in gases both ions and electrons move.

The study of electrochemistry is concerned with ionic conduction in solu-
tion and in fused salts. There are two aspects to electrochemistry:

A) the conversion of chemical energy to electrical energy, that is, the
generation of an electric potential (voltage), the source of which is a
chemical reaction. Devices which accomplish this are variously called voltaic
celh or primary cells or batteries.

B) the conversion of electrical energy to chemical energy, that is, the
application of an electric potential to an electrolyte, either a solution or a
fused salt, resulting in a flow of current which produces chemical reactions
and chemical products at the electrodes. This process is known as electrolysis
and the apparatus in which it is accomplished is an electrolytic cell. All
electrolytic cells are secondary cells.

1. The Voltaic Cell. Let us consider the reaction of zinc, Zn, and chlorine,
Cl 2 ,

(1) Zn Cl 2 — > ZnCl 2

This is an oxidation-reduction reaction for which the partial ionic equations are

(2) Zn - 2 e Zn 2 +

(3) Cla 4" 2 e — > 2 Cl“

The reaction is the transfer of two electrons from a Zn atom to a Cl 2 mole-
cule, resulting in the formation of a Zn 2 + ion and two Cl™ ions, the mixture
of which constitutes zinc chloride, ZnCl 2 , If a stick of Zn is placed directly
into chlorine water, a solution of Cl 2 in H 2 0, this reaction takes place.

Electrochemistry I: The Voltaic Cell

307

Electrons are transferred at the point of contact between a Zn atom and a
Cl 2 molecule. Under these circumstances any energy evolved in the re-
action is released as thermal energy or heat.

In order to obtain the energy of the reaction as electrical energy the re-
actants must be separated to prevent the direct transfer of electrons. The
chemical reaction is unchanged. Electrons still go from the Zn to the Cl 2
but they are provided with a different path. This is a wire or any metallic
conductor, the external circuit, which connects the Zn to the Cl 2 , as shown
in Figure 24.1. Both the Zn rod and the Cl 2 "dip” into water or some solu-
tion through which the movement of ions completes the electrical circuit.
For practical purposes the Zn rod may dip into a solution of an inactive
electrolyte, such as NaCl, while the Cl 2 may be dissolved in a similar solu-
tion. A porous partition separates the cell into two compartments and pre-
vents direct mixing of the reactants, Zn and Cl 2 . Ions, however, such as
Cl", Zn 2- K and Na + , can pass through the partition readily. Since the Cl 2

Meter

Cell
Cl in

NaCl Solution

The Zn rod is immersed in a NaCl
solution. The Cl 2 is dissolved in a
NaCl solution in which a platinum
wire, Pt, is suspended. A porous par-
tition separates the solutions. A wire
from the Zn to the Pt completes the
circuit.

Figure 24 A. A Voltaic Cell.

itself is a nonconductor unlike the metallic Zn, a rod or a wire of an inert
substance, such as platinum or carbon, which will not react with the Cl 2
or the electrolyte but which will conduct an electron current, is placed
therein. This wire is called an electrode and, being an electrical conductor,
its function is to carry electrons from the external circuit directly to the
Cl 2 . Because the Zn rod is a solid metal it can serve the dual function as its
own electrode to deliver electrons to the external circuit.

That a current is produced is easily shown by connecting the electrodes
of the cell to a galvanometer or a voltmeter. The circuit is then complete
as shown in Figure 24.1 and the needle of the instrument will deflect as
an indication of the current flow. Any resistor such as an electric light bulb,
connected between the Zn and the Pt, would also serve to complete the
circuit. As the current* flows, Zn dissolves to form Zn 2 + ions (Equation 2)
while at the Pt electrode Cl 2 gains electrons to form Cl" ions (Equation 3).

When the Zn and the Cl 2 are placed in separate containers a salt bridge is
used instead of the porous diaphragm. A salt bridge is an inverted U-tube

308

Electrochemistry I: The Voltaic Cell

containing a solution of an electrolyte, such as NaCl, and it, too, permits
the passage of ions but effectively prevents mixing of the reactants (Fig-
ure 24.2).

Meter

Cl. in

NaCl Solution

The salt bridge serves
the same purpose as a
porous partition, to keep
the two solutions from
mixing mechanically. Ions
can readily pass through
the salt bridge, which
contains an electrolyte
such as NaCl.

Figure 24.2. Voltaic Cell with a Salt Bridge.

2. Operation of the Zn— Cl 2 Voltaic Cell. Because of its electron con-
figuration a zinc atom has a tendency to lose its two 4s electrons. On the
other hand a chlorine atom must gain a single electron to complete its valence
shell. If a rod of Zn is placed in water the Zn atoms tend to form Zn 2 +
ions (Equation 2). The, Zn 2 + ions are soluble in water. The electrons remain
on the Zn rod since free electrons do not exist in solution; only ions or mole-
cules normally exist in aqueous solution. Because of the Zn 2 + ions present,
the water solution would be positive with respect to the Zn rod or, con-
versely, the Zn rod with excess electrons upon it would be negative with
respect to the solution. Since the separation of electrical charge creates a
difference in potential there is a potential difference between the Zn rod
and the solution (Figure 24.3). This potential difference prevents the forma-
tion of no more than an insignificant amount of Zn 2 + ions, and thus also

The Zn rod sends Zn 2 + ions into solution causing the solution to
acquire a positive charge while the electrons lost by the Zn re-
main on the rod, imparting to it a negative charge.

Figure 24.3. The Zinc Half-Cell,

Electrochemistry I: The Voltaic CeU

309

the concurrent production of electrons. For two reasons the chemical re-
action thus does not proceed to any noticeable extent nor does the cell
therefore operate. Electrons must be removed from the Zn rod and the Zn 2 +
ions in solution must be removed from the neighborhood of the Zn rod or,
more accurately, the positive electric field in the solution must be neutralized.

If a wire is connected from the Zn rod to a Pt wire dipping into the
solution of Cl 2 , electrons can move through the wire from the Zn to the Cl 2
where electrons could be consumed by the oxidation of the Cl 2 (Equation 3).
But here, too, the gaining of electrons would proceed only to an insignificant
extent unless the Cl - ions formed were removed from the vicinity of the
Pt electrode and the negative electric field neutralized. Thus, for the cell to
operate more than momentarily, Cl” ions must migrate out of their compartment
through the porous partition or the salt bridge in the direction of the Zn
rod while, simultaneously, Zn 2 + ions move from their compartment toward
the Pt electrode. For the movement of one Zn 2 + ion there is the concurrent
migration of two Cl” ions so that there is produced in time an equivalent
mixture of Zn 2 + and Cl” ions, or ZnCl 2 . Under these circumstances the cell
will operate to deliver Current. Because of the equivalent migration of ions,
all portions of the solution have equivalent quantities of positive and negative
ions and hence are electrically neutral.

In the compartment containing the Zn rod in a solution of NaCl, the
movement of two Na + ions out of this compartment would be equivalent
to the migration of one Zn 2 + ion. For practical purposes pure water is not
used as an electrolyte. In pure water the ionic concentration is so slight
that the internal electrical resistance of the cell would be so high as to
reduce any current to a very small value.

It will bear repetition that a voltaic cell is merely a device which utilizes
an oxidation-reduction chemical reaction so that the energy of the reaction
is liberated as electrical energy rather than as heat energy. The magnitude
of the electrical energy produced is identical to that of the heat energy which
would result if the chemical reactants were mixed directly. A cell will
operate just so long as the chemical reaction on which it is based can
proceed. It will cease functioning when either one of its reactants is exhausted
or when a state of equilibrium is reached between the concentrations of
reactants and products of the chemical reaction so that the reaction no longer
has a tendency to proceed.

3. Nomenclature Applicable to a Voltaic Cell. The zinc rod, being the
source of the negatively charged electrons, is the negative (-) electrode.
Because the Cl 2 creates a deficiency of electrons at the Pt electrode, and
also because in the -external circuit electrons move from the Zn to the
Pt, the Pt electrode is positive ( + ). If the complete cell were sealed into
a container with two outside terminals, one connected internally to the Zn
electrode and the other to the Pt electrode, the Zn terminal would be marked
(-) and the Pt terminal ( + )*

By definition the electrode at which oxidation occurs is the anode and
that at which reduction takes place is the cathode . The definitions of anode
and cathode are independent of the polarity of the electrodes. In the Zn — Cl 2
cell, the Zn electrode is the anode and the Pt electrode is the cathode. In a

310

Electrochemistry I: The Voltaic Cell

voltaic cell the anode is (-) and the cathode is ( + ). Ions which migrate
toward die anode are called anions; ions which travel toward the cathode
are cations. Thus the Cl' ion is the anion and the Zn-+ ion is the cation. In
all cases the anion is negative and the cation is positive.

In the external circuit the electric current is the flow of electrons from the
negative electrode to the positive electrode. By long-standing convention for
electrical circuits, however, it has been defined that “current” flows from
die positive to the negative electrode. In the solution the flow of electricity
or the current is the migration of ions, both positive and negative, in op-
posite directions, as shown in Figure 24.4.

- — Conventional Ctinent

The electrons given up by the Zn (-)
electrode travel through the external
circuit to the Pt (~j-) electrode. By
convention, “current” is said to flow
from the (+) to the (-) electrode. The
circuit is completed by the flow of
anions and cations in solution, in op-
posite direction and through the
porous partition.

Figure 24.4. Current Flow in a Voltaic Cell

The potential or, preferably, the electromotive force (symbol emf) of a
voltaic cell is the driving force or pressure on the electrons and ions. It is the
sum of two tendencies— that of a Zn atom to lose electrons and that of a
Cl 2 molecule to gain electrons. The quantitative determination of electromotive
force is taken up later in the chapter.

4. The Daniell Voltaic Cell. In theory any oxidation-reduction reaction
can be used as a basis for a voltaic cell but, for practical reasons, some lend
themselves more readily to commercial application. The displacement of copper
by the more active element, zinc, is the basis of the Daniell voltaic cell.

(4) Zn + Cu 2 + Zn 2 + + Cu

A possible arrangement for such a cell is illustrated in Figure 24.5a, The
reactants, Zn and Cu 2 + ion, the latter in a solution of CuS0 4 , are separated.
Any inert electrode can be used as the cathode, but inasmuch as Cu is
deposited thereon, a Cu electrode generally is used. Since Zn 2 + ions are
produced at the anode, the initial solution in that compartment is usually
a solution of a zinc salt, such as ZnS0 4 . A porous partition separates the
ZnS0 4 solution from the concentrated solution of CuSO*. In the commercial
cell no porous partition is used. Instead advantage is taken of the difference
in densities of the two solutions ZnS0 4 and CuS0 4 ; the dilute ZnS0 4 solution

Electrochemistry I: The Voltaic CeU

311

is floated on the saturated CuS0 4 solution, which is maintained saturated by
the presence of excess CuS0 4 • 5 H 2 0 crystals in the bottom of the cell.
Whatever the physical arrangement of the cell may be, the chemical char-
acteristics are unchanged.

(a) A schematic representation of a voltaic cell utilizing the reaction:

Zn -j- — > Zn 2 *i" -f- Cu

Zinc

ZnS0 4 solution
CuS0 4 solution
Copper

CuS0 4 crystals

(b) The Zn - Cu 2 + cell arranged as a single compartment voltaic cell, known as
the Daniell or Gravity Cell, A Zn “crowfoot” is suspended in dilute ZnS0 4 solution
that is less dense than, and hence floats on, a saturated CuS0 4 solution. To keep
the latter saturated, crystals of CuS0 4 * 5 H a O are placed at the bottom of the jar.
The Cu electrode consists of strips of sheet metal fastened together. An insulated
wire is connected to the Cu electrode and reaches up through the solutions to form
the (-}-) terminal of the cell; the Zn is the (-) electrode.

Figure 24.5. The Daniell or Gravity Cell.

5. Measurement of the Electromotive Force of a Voltaic Cell. The electro-
motive force of a cell can be measured by a voltmeter but this will not give
too accurate a result because the current drawn for the voltmeter causes
changes within the cell which affect the voltage to be measured. To measure
accurately the electromotive force generated by the Zn — Cl 2 cell, as a
typical example, the meter in the circuit of Figure 24.1 is removed and the

312

Electrochemistry I: The Voltaic CeU

electrodes of the cell are connected to an external source of potential whose
magnitude can be varied by the experimeter and which is so connected
that its potential opposes that of the cell whose electromotive force is to be
measured. The apparatus so used to apply an external potential is called a
potentiometer . A simple type of potentiometer is shown in Figure 24.6. If
the potential opposing the cell is increased gradually the current as read on
the galvanometer will decrease, become zero, and then reverse in direction.
The applied electric potential from the external source which causes the cur-
rent flow to become zero and hence just balances the unknown ceil potential
is the electromotive force of the cell.

When the applied potential is less than the cell potential, the cell acts as
a voltaic cell to deliver current to the external circuit. When the applied
potential is less than the cell potential, however, current flows in the op-
posite direction and the electrode reactions are the reverse of what they are
when the cell acts as a voltaic cell. Under such conditions the cell is being
electrolyzed. Zinc will plate out on the zinc electrode, Zn 2 + + 2 e -» Zn,
and chloride ion will be converted to chlorine at the platinum electrode,
2 Cl* - 2 e -» Cl 2 .

6. Single Electrode Potentials. A cell may be considered as being com-
posed of two half cells, each consisting of an electrode and the solution in
which it is suspended. In one half cell the oxidation takes place and in the
other the reduction. The potential between an electrode and its solution
is called the single electrode potential or half-cell potential. Such a potential
is a measure of the tendency of the electrode reaction to take place.

Though it is a simple procedure to measure the total potential of a com-
plete voltaic cell with a potentiometer, no satisfactory technique for the
measurement of the absolute value of a half-cell potential alone has yet
been contrived. To get around this difficulty an arbitrary value of single
electrode potential is assigned as a standard to a particular half-cell and
then all other single electrode potentials are measured relative to that
standard. Such a "modus operandi” is a common procedure in scientific
work. To a half-cell consisting of hydrogen gas at one atmosphere pressure
bubbling through a one molar solution of hydrogen ion around an inert
platinum electrode there is arbitrarily assigned a single electrode potential
of zero. The reaction of this standard hydrogen electrode as part of a voltaic
cell is

(5) Hj 2 H+ + 2 e, if it acts as the anode

(6) 2 H+ + 2 e -» H 2 , if it acts as the cathode*

To indicate the equilibrium existing between H 2 and H+, the two equations
can be combined.

(7) H a 2 H + + 2 e

Whether the hydrogen electrode acts as the anode or the cathode in a
cell depends upon the relative tendency of the other half-cell reaction to
gain or lose electrons. The assignment of the value zero to the hydrogen
electrode does not mean that the tendency of reactions 5 and 6 to proceed

Electrochemistry 1: The Voltaic Cell

313

(a) (b)

(a) A schematic diagram of the use of a potentiometer in measuring the electro-
motive force of a cell. A porous partition or a salt bridge can be used.

(b) A pictorial diagram of a slide wire potentiometer.

The potentiometer consists of a battery, A, which can deliver a potential greater
than that of the voltaic cell, X, to be measured. Battery A is in series with a high
resistance slide wire, MN, along which a sliding contact, P, can be moved. One end
of the slide wire, M, and the contact, P, are connected to the electrodes of the
cell, X. In one of the leads to cell X there is a galvanometer, G, to detect current
flow to or from X. The polarities of A and X must be so connected that their poten-
tials oppose each other. The sliding contact is moved along the slide wire to a
point, P, where the galvanometer reads zero current. At this point the electro-
motive force of X is equal and opposite to the potential delivered by the. poten-
tiometer circuit. If the slide wire has a uniform resistance throughout its length, the
electromotive force of X is

Length of MP
Length of MN

Potential of A.

Figure 24S. Measurement of Electromotive Force.

is zero. Rather the value ^zero” is to be considered in its algebraic sense,
merely as a number greater than -1 but less than +1.

To determine then the single electrode potential of an electrode in a
solution, a complete cell is set up consisting of the unknown as one half-cell
and the standard hydrogen electrode as the other half cell. The potential
of this cell, as read on a potentiometer, is ascribed completely to the unknown
half-cell since the standard hydrogen electrode has been assigned a value
of zero. Because an electrode potential varies with the ionic concentration
in which the electrode is immersed (Section 10), it is customary to use
solutions in which the ionic concentrations are one molar. Values of single
electrode potentials so obtained are known as standard single electrode
potentials (symbol E°). As an example, to determine the standard single
electrode potential of Zn in 1 M Zn 2 + ion the cell illustrated in Figure 24.7
may be used. The cell reaction is Zn 2 H+ — > Zn 2 + + H 2 . The cell poten-

314

Electrochemistry h The Voltaic Cell

tial measured on the potentiometer is 0.76 volt, and this potential of 0.76
volt is ascribed to the half -cell reaction, Zn Zn 2 +(1M) + 2 e.

A convention has been devised to represent the design of a voltaic cell.
The cell of Figure 24,7 is thereby represented as

Zn | Zn 2 + ( 1M ) || H+(l M) | H 2 (g, 1 atm), Pt

The reduced species of the half-cell reactions, and the electrodes, are
generally written at the extremes of the formulation while the oxidized species
are written in the middle. The half-cell in which the oxidation reaction occurs,
and hence the negative electrode, is written on the left, the reduction on
the right. A single vertical line indicates the separation of an electrode
from its solution, or one solution from another; a double line represents a
salt bridge. The symbols above to the right of the 1 1 stand for the standard
hydrogen electrode.

If the ZnJZn 2 + half-cell is replaced by a Cu|Cu 2+ half-cell, the poten-
tiometer reads 0.34 volt but the polarity of the* leads from the voltaic cell
to the potentiometer must be reversed. Where before current flowed through
the external circuit from the zinc to the hydrogen now current flows from
the hydrogen to the copper. Hydrogen has a greater tendency to lose elec-
trons than does copper so the cell reaction which takes place is H 2 + Cu
2 H+ + Cu. This cell reaction could be represented by

Pt, H*(g, 1 atm) | H+(1M) || Cu 2 +(1M) | Cu

By this procedure Table 24-A was developed. The table lists the standard
single electrode potentials, E°, of elements and ions in order of decreasing
potential, the term "standard” referring to the fact that all ionic concen-
trations are one molar. The electrode reactions in Table 24-A are written
as oxidations whereby electrons are produced. The corresponding potentials,
known as oxidation potentials , apply to the electrode reactions as written
in the second column! To distinguish whether an electrode reaction is a
more or less potent electron producer than hydrogen, positive ( + ) and
negative (-) signs are placed before its value of electrode potential. Those
reactions with higher oxidation potentials, and hence a greater tendency to
yield electrons than the reaction H 2 — > 2 H + + 2 e, are arbitrarily assigned
positive ( + ) values of potential and those with lower oxidation potentials
than hydrogen have negative (-) potentials. Not only the reactions of ele-
ments, but any oxidation or reduction reaction, by employing an inert elec-
trode such as platinum immersed in 1 M solutions of the ions involved,
can be used as a half-cell and its standard electrode potential determined
against the standard hydrogen electrode. The oxidation potentials of some
such important half-cell reactions are also included in Table 24-A. This
table is also known as the Electromotive Series.

7. Meaning erf the Electromotive Series. The tendency of an electrode re-
action to proceed is measured by the magnitude of its potential. The more
positive this potential the greater is the tendency of its reaction to occur. The
activity of an element is measured by its tendency to enter the ionic state.
A metal tends to lose electrons so that the higher its oxidation potential as

Electrochemistry I: The Voltaic Cell

315

Half-cell

The zinc half-cell and the standard hydrogen electrode constitute a complete voltaic
cell. In each half-cell the ionic concentrations are one molar. The cell potential is
read on the potentiometer, P. The measured potential is the standard single electrode .
potential of zinc.

Figure 24.7. Measurement of the Single Electrode Potential of Zinc.

read directly from Table 24-A the more active is the metal and the more
potent a reducing agent it is. Thus the metals listed in Table 24-A are in
order of decreasing chemical activity and the Electromotive Series is the
quantitative verification of the order of activity given for the metals in
Table 15-C. On the other hand, nonmetals tend to gain electrons and act
as oxidizing agents. Thus the lower its position in the electromotive series
the better an oxidizing agent is a nonmetallic element. Elemental fluorine, F 2 ,
is the most powerful oxidizing agent known.

Table 24-A indicates that the tendency of Zn to lose electrons and form
Zn 2 + ions, in a 1 M solution of Zn 2 + ion, is measured by a value of potential
equal to +0.76 volt. For the reverse reaction, the gaining of electrons by
Zn 2 + ion to form Zn metal, Zn 2 + + 2 e — > Zn, the absolute magnitude of
the potential is unchanged but the algebraic sign is reversed and the
potential is -0.76 volt. The negative value of the potential denotes that Zn 2 +
ion has little tendency to gain electrons; Zn 2 + ion would be a relatively poor
oxidizing agent to the extent that Zn metal is a relatively good reducing
agent.

The potential of -2.85 volts indicates that fluoride ion, F”, has a very small
tendency to act as a reducing agent and yield electrons to form fluorine, F 2 .
For the reaction, F 2 + 2e 2 F~, which is the reverse of that in the table,
however, the potential is +2.85 volts. The tendency for F 2 to gain electrons

S16

Electrochemistry I; The Voltaic Cell

Table 24-A

Standard Electrode Potentials at 25°C, Volts

Electrode

Electrode Reaction

Potential

u i

La+

Li Li+ 4- e

4-3.04

K i'

K+

Kz=± K+ + e

4-2.92

Ba 2+

Ba ^=± Ba 2 + 4- 2 e

4*2.90

1 Ca 2 +

Ca Ca 2+ + 2 e

4-2.87

Na+

Na Na+ 4- e

4-2.71

| Mg 2 +

Mg :=± Mg 2+ + 2e

4*2.37

A1 |

Al*+

A1 r± Al 8 + 4* 3 e

4*1.66

1 Mn 2 +

Mn =± Mn 2 * 4- 2 e

4-1.18

Zn 2 +

Zn :=± Zn 2 + + 2 e

4:0.76

Fe*+

Fe Fe 2 * + 2e

4-0.44

1 Cd 2 +

Cd ;=± Cd 2 * + 2 e

4-0.40

Co*+

Co Co 2 * + 2 e

4-0.29

Ni*+

Ni Ni 2 * + 2 e

4-0.22

Sn 2 +

Sn Z± Sn 2 * + 2 e

4-0.13

Pb 2 +

Pb Pb 2 * + 2e

4-0.12

H+

H 2 ^±2H+ + 2e

0.00

Sn 2 '

1 Sn 4 +

Sn 2 * Sn 4 * -f 2 e

-0.15

! Cu®+

Cu Cu 2 * + 2 e

-0.34

o 2

1 OH-

4 OH“ 0 2 + 2 H a O + 4 e

-0.40

1 Cu+

Cu Cu* 4 e

-0.51

I. 1

I-

2I^I 2 + 2e

-0.54

Fe 2+ 1 Fe*+

Fe 2 * s± Fe 8 * + e

-0.74

1 H&H-

2 Hg =± Hg 2 2 * + 2 e

-0.79

1 Ag+

Ag Ag* 4- e

-0.80

| Hg^f

Hg Hg 2 * + 2 e

-0.86

H gj

2 + | Hg 2 !

Hg 2 2 * ^ 2 Hg 2 * + 2e

-0.92

Br,

1 Br-

2 Br- ;zi Br 2 4* 2 e

-1.06

Cr»+ 1 Cr-A 2 *

2 Cr 3 + 4- 7 H 2 0 ;=± Cr 2 0 7 2 - 4 14 H+ + 6 e

-1.10

o 2

1 h 2 o

2 H 2 0 0 2 + 4 H+ + 4 e

-1.23

ci,

1 ci-

2 Cl- Cl 2 + 2 e

-1.36

1 Au*+

Au ;=± Au 8 + 4- 3 e

-1.50

Mn*+ I MnOr

Mn 2 * 4- 4 H 2 0 ^ MnCL- 4- 8 H+ + 5 e

-1.52

F* 1

1 F-

2 F- F 2 4- 2 e

-2.87

is large and hence F 2 is a strong oxidizing agent— the most potent, in fact,,
because no other substance in the table has a greater potential to gain
electrons. Permanganate ion, MnO*- in acid solution is also a strong oxidizing
agent because the reaction, Mn0 4 ~ + 8 H+ + 5 e Mn 2 + 4 H 2 0, has
a potential of +1.52 volts. In general, the reduced state of an element can
reduce the oxidized state of any element lower than it in the electromotive
series. Thus Zn can reduce Cu 2+ ion, Zn + Cu 2+ — > Zn 2 + Cu, but the
reverse reaction will not take place spontaneously inasmuch as Cu cannot
reduce Zn 2 + ion. In the construction of a voltaic cell the element or electrode
reaction higher in the electromotive series will be the source of electrons
and consequently the negative electrode, even though the value of its
single electrode potential has a more positive value.

Electrochemistry I: The 'Voltaic Cell

317

8. Calculation of the Electromotive Force of a Voltaic Cell. Table 24-A
enables the calculation of the potential of a voltaic cell from a knowledge
of the oxidation-reduction reaction taking place therein, provided that all
ionic concentrations are 1 M.

Example 1: Calculate the potential of a voltaic cell based upon the reaction:

Zn + Cu 2 +(1 M) Zn 2 +(1M) + Cu

Solution: First we write the partial ionic equations for the reaction in a manner
identical to that used in balancing redox equations by the ion-electron method.
This is an important feature of the ion-electron method— that the partial equations
used to balance redox equations are identical to the half -cell reactions which occur
at the anode and cathode of a voltaic cell. For the reaction given:

Zn — > Zn 2 + 4- 2 e
Cu 2+ + 2 e — > Cu

For each of the half-cell reactions the potential can be obtained from Table 24-A.
For the first equation, the oxidation of Zn, the potential is read directly from the
table as +0.76 volt. The second equation, the reduction of Cu 2 + ion, does not
appear in the table but its reverse reaction, for which the potential is -0.34 volt,
does. Thus for the reaction Cu 2+ + 2 e — » Cu, the potential is +0.34 volt.

Just as the addition of the two partial ionic equations gives the cell reaction,
so the addition of the corresponding potentials gives the cell potential.

Zn Zn 2 + + 2 e E<> = +0.76 volt

Cu 2 + + 2 e -» Cu E° = +0.34 volt

Zn + Cu 2 + -» Zn 2 + + Cu E° celI = +1.10 volt

The principle we have used is to add the single electrode potentials ot tne half
cell reactions that actually take place in the voltaic cell- The same result could
also have been obtained by taking the algebraic difference of the two single
electrode potentials as given in Table 24-A. Thus

E°, ;ell = +0.76 volt - (-0.34 volt) = +1.10 volt

This follows because Table 24-A gives only oxidation potentials and in every redox
reaction there must be both an oxidation and a reduction. For the reduction the
potential is opposite in sign to that in the table.

Sometimes a half-cell reaction is written as a multiple of the equation as given
in Table 24-A. However the potential, which measures a tendency to proceed, is
unchanged. For both Ag — » Ag + + e and 2 Ag — » 2 Ag+ + 2 e, the potential
is —0.80 volt. The second equation merely indicates that twice as much Ag reacts
but the potential or the tendency of a Ag atom to lose an electron is identical
in both cases. For a voltaic cell based on the reaction written as
Zn + 2 Ag+ Zn 2 + + 2 Ag,

or any multiple or submultiple thereof, the potential, E°, is + 1.56 volts.

9. Electromotive Force and Equilibrium Constant Whether or not a
given oxidation-reduction reaction will take place spontaneously and the
direction in which it will proceed can be inferred from the algebraic sign
of the potential for the corresponding voltaic cell. A reaction will be spontan-
eous in hat direction for which the calculated potential is positive ( + );
if this potential is negative (-) it is the reverse reaction which proceeds

318

Electrochemistry I: The Voltaic Cell

spontaneously. Thus for the reaction Cu + Zn 2 + Cu 2+ + Zn, the cal-
culated potential is -1.10 volt. The reaction would not proceed as written
from left to right. In Example 1 we saw that it is the reverse reaction which
occurs with a tendency measured by a potential equal to +1.10 volt.

Since the magnitude of a cell potential is a measure of the tendency of
the corresponding oxidation-reduction reaction to proceed, it must be related
to the reaction equilibrium constant which expresses the same tendency. The
equilibrium constant, K, for the cell reaction can be calculated from its
standard potential, E°, at 25°C by the following equation first derived in
1889 by the German physicist, W. H. Nemst.

(8) E°

= *I]nK
n3

where n = the number of electrons transferred

in the complete balanced cell reaction
R = the molar gas constant
T = the absolute temperature
2F = the Faraday, a constant (page 329)

Substituting the values, R = 1.987 cal/mole deg; T = 298°K (25°C);
£F = 23,060 cal/volt equiv, and converting from natural logarithms to
logarithms based on the value 10, the value of

2.303 RT

0.0591 volt equiv/mole.

(9)

Hence at 25°C

E°

0.0591

log K

The Nemst equation thus unites two important fields of chemistry-
chemical equilibrium and electrochemistry. By the measurement or the cal-
culation of the standard electrode potential, E°, for an oxidation-reduction
reaction the equilibrium constant, K, for that reaction can be calculated.

Example 2: Calculate the equilibrium constant at 25°C for the reaction of
Example 1.

Solution: Substituting in Equation 8 the value of E°, found to be +1.10 volt
in Example 1 and n = 2.

+ 1.10

0.0591

log K

K = 1.7 x 10^7

This value of K applies to the equilibrium constant expression,

T „ [Zn 2+ ]

[Cu 2 +]

For the reverse reaction, E° is —1.10 volt, and the value of its equilibrium constant
is the reciprocal of the K value found above, or

1.7 X 1037

10. Electrode Potential and Concentration* The potentials in Table 24-A
involve solutions whose ionic concentrations are 1 M and gases whose pressures

Electrochemistry 1: The Voltaic Cell

319

are one atmosphere. If these concentrations and pressures are other than
unit value the resultant potential will differ somewhat from the standard
potential. Le Chatelier s principle indicates that the actual oxidation potential
of a metal would be greater than the standard potential if the ionic con-
centration were less than 1 M and smaller if the ionic concentration exceeds 1 M.

In general, for any half-cell reaction the potential varies with the con-
centrations of the products and the reactants of the reaction, and at 25°C is
given by the equation

E = E° -

0.0591

[products]

[reactants]

Here "products” and "reactants” refer to the direction of the half-cell re-
action as dictated by the direction in which the total cell reaction is written.
For the half-cell reaction, Fe 2 + — » Fe 3+ + e ,

(for which E° = +0.74 volt)

E = E° - log 5 Fe8+] (for which E° = -0.74 volt)

2 [Fe 2+ ]

For the reaction written, Fe 3 + + e Fe 2+ ,

E = E° - log I Fe2 ll - (for which E° = +0.74 volt)

2 6 [Fe 3 *]

For the same concentration ratio of the ions the magnitude of the potential, E
(and of E°) will be the same but opposite in sign.

In a manner similar to the writing of an equilibrium constant expression,
where a coefficient other than one appears before a species in a half-cell
equation, the concentration of that species must be raised to a power equal

to the coefficient preceding it. For the reaction, 2 Cl” -* Cl 2 + 2 e

0.0591

log ISd
[Cl-] !

Where a solid, or a gas at one atmosphere pressure, is part of a half-cell
reaction, their "concentrations” are assumed to be one and so do not appear
in the expression for the potential. For Zn($) Zn 2+ + 2 e

0.0591

log [Zn 2+ ]

and for H 2 (g, 1 atm) 2 H + + 2 e

E = E 0 - log [H+] :

log [H+] 2

When the concentrations of the products and of the reactants of the half-cell
reaction are unit value, or even when their ratio is one, the potential equals
the standard electrode potential as given in Table 24-A. For a half-cell
reaction in which n = 1, a tenfold change in concentration changes the poten-
tial from E° by 0.0591 volt and where n= 2 by volt.

320

Electrochemistry I: The Voltaic Cell
single electrode potential at 25°C for the reaction

Example 3: What is the
Zn Zn 2 +(0.01M) + 2 e?

Solution:

E = E°

E = +0.76 -

0.0591

log [Zn 2 +]

log (0.01) = +0.82 volt

The potential, or the tendency of Zn to go into solution as Zn 2 + ion, is greater
in a concentration of 0.01M Zn 2+ than in 1 M Zn 2 +.

Example 4: Calculate the potential of a voltaic cell based upon the reaction
Zn + Cu 2 +(0.1M) -» Zn 2 +(0.5M) + Cu

Solution: The potentials for the half-cell reactions, according to the equation
as written, are calculated and then added tc obtain the cell potential. The half-cell
reactions are:

Zn -> Zn 2 +(0.5M) + 2 e
Cu 2 +(0.1M) + 2 e Cu
and the corresponding single electrode potentials are:

E = E° - log [Zn 2 +] = +0.76 - log 0.5 = +0.77 volt

E =, E°

0.0591

log

[Cu 2 +]

= +0.34 -

0.0591

log

-i— = +0.31 volt

0.1

The cell potential = +0.77 volt + 0.31 volt = 1.08 volt

Note that the cell potential could have been calculated in a single step by an
equation which combines both single electrode potential equations.

u , xr'fi 0.0591 1

E = E° cell — log

[Zn 2 +]

[Cu 2 +]

+ 1.10

0.0591 . 0.5

T - 108 oT

+ 1.08 volt

11. Concentration Cells. Because the potential between an electrode and
a solution of its ions varies with the concentration of these ions, a cell,
constructed by placing electrodes of the same metal in contact with solu-
tions of its ions of different concentrations, will produce a difference of
potential. Such voltaic cells are known as concentration ceUs . For any con-
centration cell, E° is zero; the driving force, or potential, is due solely to
a difference in concentration and if all concentrations were unit value
there would be no tendency for the cell reaction to proceed. An example
of a concentration cell is: Zn | Zn 2+ (0.1M) II Zn 2+ (1M) I Zn. The net
cell reaction is Zn 2+ (1M) -» Zn 2 +(0.1M); that is, the concentrated Zn 2 " 1 "
ion solution becomes more dilute and the more dilute solution increases in
concentration. The potential of this cell, which can be calculated by Equa-
tion 10, is +0.0295 volt.

12, The Dry Cell. The so-called dm cell (Figure 24,8) consists of a cylin-
drical zinc container which serves as the negative electrode and a central

Electrochemistry J: The Voltaic Cell

321

carbon rod, the positive electrode. The electrolyte is a pasty wet mixture
consisting of ammonium chloride, NH 4 C1, manganese dioxide, Mn0 2 , and
zinc chloride, ZnCl 2 . A layer of porous paper separates the zinc from the
electrolyte. The cell is sealed with pitch to prevent evaporation; if the cell
were truly dry, ionic migration could not take place. At the negative electrode
the Zn is oxidfized to Zn 2 +. The electrons thereby produced travel thr ou gh
the external circuit to the positive carbon electrode. Here ammonium ion,
NH 4 +, gains an electron and is reduced; NH 4 + + e -» NH S + H. The
Mn0 2 oxidizes the liberated hydrogen which would otherwise accumulate
as a film of gas on the carbon electrode and decrease the flow of current.
This phenomenon is known as gas polarization and the Mn0 2 , which prevents
this action, is a depolarizer . The NH 3 unites with the Zn 2 + ions formed to
produce a soluble complex ion, Zn(NH 3 ) 4 2+ . Thereby no NH S gas is evolved
nor does the concentration of Zn 2 + increase. The potential of the dry cell
is approximately 1.53 volts.

Figure 24.8. The Dry Cell.

QUESTIONS

1. Compare the nature of the conducting particles in electrical conduction
through (a) metals (b) aqueous solutions (c) molten salts (d) gases.

2. Define and illustrate the following terms: (a) voltaic cell (b) single electrode
potential (c) standard electrode potential (d) cell potential (e) half-cell re-
action (f) anode (g) cathode (h) negative electrode (i) positive electrode.

3. Draw diagrams of voltaic cells based on the following reactions. In each
case label the anode, cathode, negative electrode, positive electrode, direction
of electron flow in the external circuit, direction of ionic flow in solution, and
indicate the reaction at each electrode, (a) Zn 4* HC1 (b) Fe 2+ + Br 2
(c) HC1 + KMn0 4 (d) Cu + Ag+.

4. Calculate the potentials of the cells in Problem 3, assuming unit concentration
for each species.

5. What is the purpose of a porous partition or a salt bridge in a voltaic cell?

Electrochemistry I; The Voltaic Cell

6. Ions migrate to oppositely charged electrodes in a voltaic cell. How is
electric neutrality maintained throughout the solution?

7. How is the potential of a voltaic cell measured? Draw a schematic diagram
of the electric circuit used.

8. (a) How is a single electrode potential measured (b) what is the meaning
of a. single electrode potential (c) how is a single electrode potential affected
by concentration?

9. In what respects would cells based on the following reactions differ:

(a) Zn 4* 2 Ag4 Zn 2+ + 2 Ag (b) % Zn 4 Ag+- Vz Zn 2 + 4 Ag?

10. (a) Describe the Daniell cell (b) how would its voltage be affected if the
concentration of the CuS0 4 solution were less than saturated?

11. (a) Describe the dry cell (b) what is the function of each ingredient?

12. (a) What is meant by a concentration cell (b) why is E° for a concentration
cell equal to zero (c) what is the value of the equilibrium constant for a
concentration cell reaction?

13. Write the cell reaction and calculate E° for the following:

(a) Cu | Cu 2 + || Ag+ | Ag (b) Fe | FoH- || Zn 2 4 | Zn

14. Calculate the single electrode potentials for oxidations represented by the
following half-cells: (a) Cd | Cd 2 +(1M) (b) Cd | Cd 2 +(0.02M)

15. Will reaction take place between the following pairs? Justify your answer,
(a) Fe and Fe*4 (b) Cu*4 and Zn 2 4 (c) I 2 and Cl- (d) Co 2 + and Ni
(e) Na + and Cl" (f) Mg and I- (g) H 2 and Sn 4 4 (h) Sn 2 + and Cr 2 0 7 2 -
(in acid solution),

16. Calculate the equilibrium constants for the following reactions:

(a) Sn + Pb 2 + -> Sn 2 + + Pb (b) Zn + 2 H+ Zn 2 + + H 2
(c) Cd + Fe 2+ Cd 2 4 + Fe (d) Cl 2 + 2 Br- -+ 2 Cl- + Br 2
(e) Vz Cl 2 + Bir -V Or -f % Br 2 (f ) 2 Ag + Cu 2 + -» 2 Ag4 + Cu

17. Calculate the potentials of voltaic cells based on the reactions:

(»> Sn + Pb 2 4(lAf) Sn 2 +(1M) + Pb

(b) Zn + 2 H4(1M) Zn 2 +(0.1M) + H 2 (g, 1 atm)

(d) 2 Ag 4- Cu 2 +(0JM) -» 2 Ag+(0,1M) 4- Cu

18. For the reaction Sn 4* Pb 2 4(lM) — > Sn 2+ (xM) + Pb, what should be the
concentration of Sn 2 4 which will just prevent the reaction from proceeding
to the right? Note: calculate the [Sn 2+ ] which will give a cell potential of zero.

19. Given the equilibrium: Sn 2 4 -}- Fe 3 + Fe 2 4 -j- Sn 4 4 ? 'when one mole of
Sn 2 + and one mole of Fe 3 4 are added to one liter of water what will be
the concentration of Fe 2 4 present at equilibrium? Note: calculate the equili-
brium constant for the reaction and then treat the problem as one in chemical
equilibrium; assume one liter of water and one liter of solution are synonomous.

20. For a voltaic cell utilizing the reaction, Ag4 -f Fe 2 4 Ag 4- Fe a 4, how
much will the potential change if (a) the concentration of Fe 3 + ion is in-
creased tenfold (b) the quantity of Ag is doubled?

Electrochemistry II
Hie Electrolytic Cell

When two chemically inert electrodes are connected to a source of
direct current, such as a voltaic cell, each electrode takes on the polarity of
the terminal to which it is connected. The electrode connected to the nega-
tive terminal becomes negative because an excess of electrons is driven
upon it by the potential of the direct current source; the electrode connected
to the positive terminal becomes positive because electrons are removed
from it, leaving an excess of positive charge. If the two electrodes are im-
mersed in an electrolyte, and if the potential is sufficiently high, electrolysis
takes place and chemical reactions occur at the electrodes. An ion, atom,
or molecule gains electrons at the negative electrode (reduction) and simul-
taneously some substance loses electrons at the positive electrode (oxidation).
To decompose a compound electrolytically, the potential applied must be
at least equal to, usually greater, than the potential that the products of
the electrolysis would generate if set up as a voltaic cell.

As an illustration of a typical electrolytic process, the electrolysis of a
hypothetical electrolyte, M+X”, is shown in Figure 25.1. Attracted by the
field of the negative electrode, positive ions, M+, migrate to it where they
gain electrons and are reduced*. M+ + e M°. By definition this electrode
is the cathode . In the case of electrolytic cells the cathode is the negative
electrode; positive ions attracted to it are cations. Simultaneously negative
ions, X”, migrate to the positive electrode where they lose electrons and are
oxidized: X” -e X°. This positive electrode, where oxidation takes place,
is the anode and the negative ions are anions.

For the loss of each electron at the anode there is a corresponding gain
of one electron from the cathode. Thus if an X - ion gives an electron to the
anode, in the external circuit there is effectively a flow of one electron
from the anode through the source of potential to the cathode, where an
M+ ion gains an electron. There exists a “conservation of electrons” in that
the number of electrons lost in the exidation process must equal the number
gained in the reduction. With ions of unequal charge, such as M 2 + and X®“,
the numbers of positive and negative ions which react at the electrodes'

324

Electrochemistry 11: The Electrolytic Cell

+ e — ►
reduction
at

cathode

B is a source of direct current, the potential of which is sufficient to remove
electrons from electrode A and to deliver them to electrode C; thereby these
electrodes become positive (-f ) and negative (-), respectively. Positive ions, M+,
are attracted to and migrate to the negative electrode, C, where they remove
electrons and form elemental M. Negative ions, X", move toward the positive
electrode where they lose electrons and form elemental X.

Figure 25.1. Electrolysis of M + X“

so adjust themselves to maintain this equality, that is, three positive to two
negative ions.

The movement of electrons in the external circuit and the migration of
ions, in both directions, within the electrolyte constitute the flow of current
in the complete electric circuit. In Figure 25.1 the electrodes are assumed
to be chemically inert in that they serve merely as conductors of electrons
to and from the electrolyte. In some electrolytic cells the electrode substance
itself may undergo oxidation or reduction, e.g., in the electrorefining of
copper, die anode reaction is the oxidation of a Cu electrode to Cu 2 +.

1. Electrolysis of Molten Sodium Chloride, NaCl. Solid NaCl consists
of Na+ and Cl~ ions. If the NaCl is molten the ions are free to move.
Figure 25.1 can be used to illustrate the electrolysis of molten NaCl, if the
M+ ion represents a Na+ ion and the X~ ion a Cl- ion. The Na+ ions migrate
to the negative electrode and are reduced: Na + -be—* Na°; at the anode,
Cl ions are oxidized: Cl“ — e —> Cl°. Since Cl° atoms can not exist inde-
pendently two combine to form a stable Cl 2 molecule and so the oxidation of

i iOD is „ v ^ xinen Meetly as: 2 Cl~ - 2 e Cl 2 . For the “conservation of
electrons, inasmuch as two electrons are produced by the oxidation of
two Cl- ions, there must occur simultaneously: 2 Na+ + 2 e 2 Na°. The
products of the electrolysis of molten NaCl are sodium metal, Na, at the
cathode and chlorine gas, Cl 2 , at the anode. The Downs cell for production
o metallic sodium and chlorine is an industrial application of this electro-
lytic reaction.

2, Electrolysis of Aqueous Sodium Chloride. When a moderately con-
centrated aqueous solution of NaCl is electrolyzed, H 2 gas is evolved at

Electrochemistry II: The Electrolytic Cell

325

the cathode, Cl 2 gas at the anode, while OH - ion is produced in solution.
How can we account for these products, particularly for the fact that H 2
and not Na is produced at the cathode?

In molten NaCl only one reaction was possible at each electrode— the
reduction of Na+ to Na° at the cathode and the oxidation of Cl" to Cl 2 at
the anode. In an aqueous solution of NaCl there exist: Na + ions, Cl" ions,
H 2 0 molecules, and also H+ ions and OH" ions (both approximately I0~ T
Molar) from the slight ionization of water. Hence there is more than one
species of cation, namely Na + and H+, which in theory, can be reduced
at the cathode. The possible electrode reactions are:

Na+ + e Na° and 2 H+ + 2 e -* H 2

Which of these electrode reactions will take place? The standard electrode
potentials listed in Table 24-A enable us to make an educated guess. We
have already seen that the magnitude of its potential is a measure of the
tendency of an electrode reaction to proceed. Because Na is higher than H 2
in the table, Na has a greater tendency to form Na + ion -than does H 2 to
form H+ ion. Conversely H+ ion is more easily reduced to the elemental
state than is Na+ ion. For metals the magnitude of the oxidation potential
is a measure of the tendency to enter or to remain in the ionic state. This
tendency decreases as we down the table. On the other hand, the lower the
position of a metal in the table, the more readily are its ions reduced, that
is, the more readily does the reverse of the reactions as given in Table 24-A
proceed.

A quantitative calculation, verifying this conclusion, can be made quite
simply. Let us assume a 1M solution of NaCl in water; to this concentration
the values of the oxidation potentials in Table 24-A apply strictly. Hence

Na - e Na+(1M) E° = +2.71 volts

If the concentration of the H + ion were 1M its oxidation potential
would be 0.00 volt but where its concentration is 10~ 7 M the potential is
0.42 volt.

H 2 - 2 e 2 H+(10- 7 M) E = +0.42 volt

During electrolysis the reduction reaction at the cathode is the reverse of
the foregoing reactions and so, too, would be the signs preceding the values
of the potentials.

Na+(lAf) + e .Na E = -2.71 volts

2 H+( 10" 7 Af ) + 2 e -> H 2 E = -0.42 volt

Since -0.42 volt is algebraically greater than -2.71 volt the reduction of H+
ion proceeds in preference to the reduction of Na+ ,ion. Indeed, no Na+ ion
is ever discharged in the electrolysis of aqueous NaCl. 1 As the H+ ion is
reduced at the cathode, water molecules ionize to maintain the H+ ion con-

1 Where a mercurv cathode is used, Na metal can be discharged because it dissolves
in the mercury and its activity is thereby reduced.

326

Electrochemistry II: The Electrolytic CeU

centration. The equations for the electrode reduction of H+ ion and for
the ionization of H 2 O can be combined:

2H++2e^H,

2 H 2 Q 2 H + + 2 OH“

2 H 2 0 + 2 e -> H 2 + 2 OH-

and the last equation is generally written as the cathode reaction. Thus
the electrode reaction of H 2 0 molecules, or the discharge of H+ ions, leaves
in solution an excess of OH" ions. These, with the Na + ions, which are un-
affected by the electrolysis, constitute a solution of sodium hydroxide, NaOH.

An analogous situation exists with the anions present, Cl' and OH'. Two
anodic oxidations may be possible:

2 Cl- Cl 2 + 2 e and 4 OH“ 0 2 + 4 H+ + 4 e

The second reaction can also be written, by the addition of four H + ions to
each side of the equation and the formation of H 2 0 on the left side, as
2 H 2 0 0 2 -f- 4 H+ + 4 e. The potential for the first of the two electrode

reactions is the greater so that Cl" ions are oxidized to Cl 2 ; the OH“ ions
(or H 2 0) do not undergo anodic reaction.

Thus the electrode reactions, and the net reaction, for the electrolysis
of an aqueous solution of NaCl are:

cathode: 2 H 2 0 + 2 e H 2 + 2 OH"

anode: 2 Cl' -» Cl 2 + 2 e

net reaction: 2 H 2 0 + 2 Cl“ H 2 + Cl 2 + 2 OH~

No electrons appear in the net reaction which is a complete oxidation-
reduction chemical reaction. The electrolysis of aqueous NaCl is an im-
portant industrial process for the production of H 2 , Cl 2 , and NaOH.

In the electrolysis of a dilute aqueous solution electrode reactions occur
which are independent of the nature of the dissolved electrolyte. These are:

cathode: 2 H 2 0 + 2 e -> H 2 + 2 OH'

anode: 2 H 2 0 -> 0 2 + 4 H+ + 4 e

On the assumption that the H + ions and the OH - ions produced are per-
mitted to mix and combine to form H 2 0, the net reaction is obtained by
doubling the cathode reaction and adding:

net reaction: 2 H 2 0 -» 2 H 2 + 0 2

Identical electrode reactions would take place in the electrolysis of dilute
aqueous solutions of other salts such as sodium, sulfate, Na 2 S0 4 , where
neither the Na+ ion nor the S0 4 2- ion undergoes electrode reaction.

3. Electrorefining. The selective discharge of ions based on their re-
duction potentials is the basis of electrorefining and of analytic separations.
Where two ions are present in solution, the applied potential can be so regu-

Electrochemistry II: The Electrolytic Cell

#27

lated as to discharge and plate out only that ion which more readily under-
goes electrode reaction. For example, in a solution containing ,copper ion,
Cu 2+ , and zinc ion, Zn 2 +, practically all of the Cu 2 + ion can be plated out
at a potential which is insufficient to plate out the Zn 2 + ion.

4. The Storage Cell. When a voltaic cell operates to deliver electric
current a chemical reaction occurs and chemical energy is transformed into
electrical energy. The cell is spent or discharged when the chemical reaction
is complete or an equilibrium between reactants and products within the
cell is reached. If an opposite current is passed through the spent cell from
an external source of D.C. potential, that is, if the cell is now electrolyzed,
chemical reactions occur which are the reverse of those which took place
when the cell acted as a voltaic cell. The initial reactants are regenerated and
the cell may again be operated as a voltaic cell for the production of electric-
ity. In this way electrolysis stores up electrical energy as chemical energy
which subsequently is converted into electrical energy during the discharge
of the cell. Such a reversible cell is called a storage cell or accumulator.

The most common type of storage cell is the lead storage cell. The
electrodes consist of lead gratings. The openings in the gratings are filled
with spongy lead, Pb, in one plate and with lead dioxide, Pb0 2 , in the
other. The electrodes are alternately arranged and separated by insulators
of wood or perforated rubber. The electrolyte is a solution of sulfuric acid,
H 2 S0 4 , about 3.7 Molar (Figure 25.2). In this condition the cell is in the
charged state and can operate as a voltaic cell with a potential of about
2.04 volts. During discharge the electrode reactions are:

cathode: PbO 2 (s) + 4H+ + S0 4 2 - + 2 e -> PbS0 4 (s) + 2 H 2 0

E° = 1.685 volts

anode: Pb(s) + S0 4 2 ~ -» PbS0 4 (s-j + 2 e E° = 0.355 volt

In both electrode reactions the final product is solid lead sulfate, PbSO*,
which adheres to the electrodes. The net chemical reaction of the cell is:
Pb(s).+ Pb0 2 (s) + 4H++2 SO, 2 - PbS0 4 (s) + 2 H 2 0 E° = 2.04 volts
As the cell discharges the removal of H+ and S0 4 2 " ions decreases the con-
centration and hence the density of the H 2 S0 4 solution. The degree of
charge or of discharge of the cell can be determined by measuring this
density. In a fully charged cell the density of the electrolyte is about
1.30 g/ml; in a discharged cell the density drops to about 1.15 g/ml. In
order to charge the cell a current from an external source is passed through
it in a direction opposite to that of the discharge current. The resulting
electrode reactions are the reverse of those just given and Pb, Pb0 2 , and
H 2 S0 4 are reformed. Since the potential of a single cell is approximately
2.04 volts, three lead storage cells in series produce six volts, and six cells
give twelve volts.

Another storage cell of importance is the Edison cell, in which the
electrolyte is a solution of potassium hydroxide, KOH, and the electrodes
are iron, Fe, and hydrated nickel (IV) oxide, Ni0 2 . The electrode reactions
upon discharge and the net reaction are:

328

Electrochemistry II: The Electrolytic Cell

cathode: NiO 2 (s) + 2 H 2 0 + 2 e Ni(OH) 2 ($) + 2 OH"

E° = +0.49 volt

anode: Fe($) + 2 OH“ Fe(OH) 2 ($) + 2 e E° = +0.88 volt

net reaction: Ni0 2 ($) + Fe($) + 2 H 2 0 — > Ni(OH) 2 ($) + Fe(OH) 2 ($)

' E° = 1.37 volts

Advantages of the Edison cell are that no corrosive acid is involved and
the output voltage remains constant because only solid electrode materials
are used up while the concentration of the electrolyte, OH" ion, remains
constant.

5. Electrical Units. The coulomb (symbol Q) is a measure of the quantity
of electricity which flows in a circuit in a specified interval of time. A
coulomb represents the flow of a definite number of electrons; 96,500 coulombs
correspond to the Avogadro number of electrons, 6.02 X 10 23 electrons.

The ampere (symbol I) is a measure of the rate of flow of electrical
charge. For example it corresponds to the number of electrons, or ions, per
second. One ampere is a flow of current equal to one coulomb per second. Thus

(1) 1 = 9 - and Q = It where I = the current in amperes

* Q — the quantity in coulombs

t = the time in seconds

The volt (symbol E) measures the electromotive force or the electrical
pressure driving the electrons through a circuit.

The ohm (symbol R) measures the resistance offered to the flow of current.
Current, voltage, and resistance are related by Ohm's Law:

(2) I (ampere) = ; E = 1 R and R = —

R (ohm) I

The joule is a unit of both energy and work. Electrical energy in
coulombs is the product of volts and coulombs.

(3) Energy = E Q = E I t

The watt is a unit of power. Power is the rate of doing work or the
rate of dissipation of energy. One watt is defined as a rate equal to one
joule per second; in electrical units one watt equals one volt times one ampere.

(4) Power = ; Power (watts) = E (volts) X I (amperes)

Time

6. Quantitative Electrochemistry. In 1833 Michael Faraday found that,
in electrolysis, the quantity (weight) of material undergoing chemical change
at an electrode is proportional to:

A) the quantity of electricity (coulombs) that pass through an electrolyte

B) the equivalent weight of the substance undergoing the electrode re-
ackiotv

Electrochemistry II: The Electrolytic Cell

Specifically the passing of 96,500 coulombs through a cell causes one gram-
equivalent of chemical reactfbn at each electrode. This value of 96,500
coulombs is a unit known as the Faraday .

In Figure 25.3 a number of electrolytic cells are connected in series. If
the value of the current and the duration of the electrolysis are so adjusted
that 96,500 coulombs pass through the circuit and hence through each cell,
the weights of the electrochemical products are:

When the cell is fully charged, the negative electrode consists of Pb, the positive
electrode of PbOa> and the electrolyte is H^SO^ Upon discharge, VbSO* is
formed on both electrodes according to the reactions shown above. Charging the
cell requires an external source of D.C. potential greater than that produced by
the cell upon discharge. During’ charging electrolysis occurs within the cell and
the electrodes and electrolyte are regenerated to their initial states.

Figure 25,2 . The Lead Storage Cell.

330

Electrochemistry II; The Electrolytic Cell

hydrogen:

1.01

grams

(2 H+ + 2 H 2 ; in cell 1)

copper:

63.6

grams

(Cu 2+ + 2 e Cu; in cell 2)

iron:

55.8

grams

(Fe 2-1 " + 2 e — » Fe; in cell 3)

iron:

55.8

grams

(Fe 3+ + 3 e Fe; in cell 4)

oxygen:

16.0

grams

(2 HjO 0 2 + 4 H+ + 4 e; in cell 2)

chlorine:

35.5

grams

(2 Cl- -4 Cl 2 + 2 e; in cells 1, 3, and 4)

A quantity of electricity other than one Faraday would yield a proportionate
quantity of electrochemical products* Thus 0.5 Faraday or 48,250 coulombs

63 6

would plate out 0.5 X — grams of copper.

Faradays Laws represent one of the most exact experimental facts in
scientific work* This is because 96,500 coulombs constitute one mole, or the
Avogadro number, of electrons and an electrode reaction is essentially a
counting process— a doling out or a removal of a whole number of electrons
at each electrode. To each H + ion that is reduced at a cathode, one electron
is given; to each Cu 2 + ion, two electrons; to each Fe 2+ ion, two electrons;
and to each Fe 3 + ion, three electrons. With a definite number of electrons,
or coulombs, available at an electrode only a definite number of ions
(or atoms or molecules) can be oxidized or reduced at that electrode.
Since one gram-equivalent of chemical reaction at an electrode represents
the transfer of one mole of electrons, or 96,500 coulombs, the weights of the

* j LZ * _

__ _|_

□

1 J

□

HC1

CuSCL

FeCh

The electrolytic cells are connected in series so that the same quantity of electricity
passes through each cell. One Faraday causes one gram-equivalent of chemical
change at each electrode. In cell 1, 1.008 g of H 2 wul be formed at the cathode
and 35.5 g of Cl 2 at the anode; in cell 2, 31.8 g of Cu at the cathode and 8.0 g
of 0 2 at the anode. In cells 3 and 4, 35.5 g of Cl 2 will be liberated whereas
27.9 g of Fe will be deposited in cell 3, but only 18.6 g of Fe in cell 4.

Figure 25.3. Faraday's Law.

Electrochemistry II: The Electrolytic Cell

331

substances undergoing electrode reaction are proportional to their equivalent
weights.

This relationship is the basis for the definition of the equivalent weight
of an oxidizing agent or a reducing agent. In an oxidation-reduction reaction
one formula weight of a substance gains or loses a specific number of elec-
trons. That weight which corresponds to the transfer of one mole of electrons
is the gram-equivalent weight.

Faraday’s Laws can be summarized in a simple equation based on the
relation that one Faraday, or 96,500 coulombs, corresponds to one gram-
equivalent of electrode reaction.

Actual number of coulombs
^ passing through a cell

96,500 coulombs

Actual weight in grams of substance
undergoing electrode reaction

One gram-equivalent weight (GEW)

(6)

I X t
96,500

where

I = the current in amperes
t = the time in seconds

to = the weight in grams of the substance
undergoing electrode reaction

(7) w =

I X t
96,500

A == the formula weight of the substance
undergoing electrode reaction

n = the electron change in the electrode
reaction

It should be noted that neither the voltage nor any concentration is directly
concerned in the statement of Faradays Laws.

Faraday’s Laws apply not merely to electrolytic cells but also to voltaic
cells. The weight of a substance undergoing electrode reaction in a voltaic
cell is also proportional to its equivalent weight and to the number of coulombs
which the voltaic cell delivers, so that equations 5, 6, and 7 are applicable.

Example 1: A current of 1.50 amperes at a potential of 6.0 volts is passed
through 4.0 liters of 1.2 M CuS0 4 for 2.00 hours. Calculate the weight of each
electrolytic product.

Solution: The electrode reactions are:

cathode: Cu 2 + + 2e — > Cu
anode: 2 H 2 0 0 2 + 4 H+ + 4 e

The weight of Cu produced is

__ 1,50 amp x 2,00 hr x 3600 sec/hr x 1 coulomb/ amp see 63.6 g/g-atom
96,500 coulomb/g-equiv 2 g-equiv/g-atom

w = 3.56 grams of Cu.

The weight of O z produced is

_ 1.50 amp x 2.00 hr x 3600 sec/hr x 1 coulomb/amp sec v 32.0 g/mole

~ 96,500 coulomb/g-equiv 4 g-equiv/mole

to = 0.896 gram of O a .

332

Electrochemistry 11: The Electrolytic Cell

If the temperature and pressure were given, the volume of 0 2 under these
conditions could be calculated.

H+ ions are also produced; in conjunction with the S0 4 2 - ions present they
constitute a solution of H 2 S0 4 . The weight of the H+ ions and of the H 2 S0 4
produced can also be calculated in a manner identical to the foregoing. The ratio,

X X , is the number of Faradays and so also the tiumber of gram-
96,500

equivalents of H+ ion and of H 2 S0 4 produced. The value of this ratio is 0.112
Faraday; hence the number of gram-equivalents of H 2 S0 4 produced is 0.112.
Uniformly dissolved in the 5.0 liters of solution, this would make a 0.028 Normal
solution of H 2 S0 4 .

Example 2: How much Zn undergoes electrode reaction in a dry cell which
delivers 0.100 ampere for 90 minutes?

Solution: The weight of Zn is

— Q-1QQ amp x §0 min x 60 sec/min x 1 coulomb/ amp sec 65.2 g/g-atom
96,500 coulomb/g-equiv 2 g-equiv/ g-atom

w = 0.183 gram of Zn.

QUESTIONS

1. (a) In what characteristics are voltaic cells and electrolytic cells similar
(b) in what do they differ?

2. In electrolysis what are the polarities of anode and cathode? Explain.

3. Draw a schematic diagram of the apparatus you would use to plate out lead,
Pb, from a solution of lead sulfate, PbS0 4 .

4. What are the electrode reactions during the electrolysis of aqueous solutions
of (a) HC1 (b) H 2 S0 4 (c) Na 2 S0 4 (d) CuS0 4 using copper electrodes
(e) A1C1 3 using platinum electrodes?

5. A solution is composed of 1 M MgCl 2 and 1 M K 2 S0 4 . What electrode re-
actions will occur during electrolysis of the solution using platinum electrodes?

6. A solution contains equal concentrations of Ni 2 + and Cu 2+ ions, (a) What
conditions are necessary to plate out each metal separately? (b) draw a
schematic diagram of the apparatus to be used.

7. What should be the characteristics of a voltaic cell in order that it can be
regenerated by electrolysis?

8. (a) Describe the composition of a lead storage cell (b) what are the electrode
reactions on charge and on discharge (c) what is the minimum voltage re-
quired for recharging (d) why can the degree of charge be determined by
the density of the electrolyte?

9. Define (a) ampere (b) coulomb (c) Faraday (d) volt (e) watt (f) joule.

10. (a) State Faraday’s Laws (b) why are the laws so exact an experimental

fact (c) why is the voltage not involved in the direct application of Faraday’s
Laws?

Electrochemistry 11: The Electrolytic CeU

333

11. A voltage of 3.0 volts is impressed across platinum electrodes dipping into
molten NaCl. Will electrolysis occur? Explain.

12. An aqueous solution contains lAf CuCl 2 and 1M ZnCl 2 . (a) At what voltage
will the Cu 2 + begin to plate out (b) at what voltage will the Zn 2 + begin
to plate out (c) what potential is required to plate out 99.99% of the
Cu 2+ (d) how much Zn 2 *^ will be plated out at the potential in part c?

13. (a) A solution contains equal concentrations of Zn 2 + and Cd 2 + ions. Is it
possible to separate the ions by electrodeposition? (b) would it be possible
if the [Zn 2 +] = 1 M and the [Cd 2 +] = 10~ 6 M?

14. A 300 watt lamp is operated at 120 volts for two hours. Calculate (a) the
current through the lamp (b) the electrical energy in kilowatt-hours, joules,
and calories (c) the weight of water that could be heated from 20°C to 50°C
by this amount of energy.

15. How long will it take a current of four amperes at six volts to deposit 10.0 g
of copper from a solution of CuS0 4 ?

16. (a) What weight of zinc will dissolve in a Daniell cell if the cell runs a
motor that draws 0.50 ampere for 90 minutes (b) what weight of copper
will be deposited simultaneously?

17. When 12.0 g of cadmium are plated out from a CdS0 4 solution at a potential
of 1.5 volts, calculate (a) the work done (b) the power dissipated if the
cell is operated for one hour.

18. In two hours, 3.2 g of Zn were oxidized in a dry cell, (a) How many
coulombs were produced (b) what was the average current?

19. The following record is taken from an experimenters notes of an electrolysis.

Solution electrolyzed: 4.0 liters 0.8 ON CuS0 4 Electrodes: Pt

Voltage : 2.0 volts Current : 2.5 amperes

Temperature : 20° C Pressure : 768 mm

(a) what weight of copper was deposited (b) what volume of oxygen,
collected over water, was produced?

20. During an electrolysis of molten NaCl, 4.6 g of sodium were obtained, (a) How
many moles of electrons were transferred (b) what weight of Cl 2 was produced?

The Halogen Elements

In Group VIIB of the Periodic System there are five elements that con-
stitute the family known as the halogens (Greek, halos : the sea, salt producer).
These are fluorine, chlorine, bromine, iodine, and astatine. So minute a quan-
tity of the short-lived radioactive element astatine has been isolated that little
is experimentally known of its properties and so it will be omitted from
the following discussion even though the Periodic Law enables us to infer
what its properties are. The similarity in electron configurations of the
halogens is shown by the following table:

Element

Atomic

Number

4s 4p 4 d 4/

6s 6p

Fluorine

Chlorine

2 5

Bromine

2 6

2 5

Iodine

2 6

fill 1 #

2 5

Astatine

2 6

2 6 10

MHB

2 6 10

2 5

1. General Properties of the Halogens. The neutral halogen atoms have
outermost shells containing seven electrons, ns 2 * * np 5 . Being one electron short
of an octet the halogen atoms gain one electron from metallic atoms to form
the halide ions, F~ Cl“, Bit, I", and At r. The halide ions have a noble gas
electron configuration, n s 2 np 6 . If X represents a neutral halogen atom then

(1) X + e X- and X 2 + 2 e 2 X~

With atoms of similar electronegativity the octet can also be attained by the
formation of a covalent bond. In the gaseous state the halogen molecules
are diatomic; a covalent bond joins the two halogen atoms.

(2) SX* + ♦ x: :x:x: (X-X or x 2 )

•« •• to ••

The halogen elements are all quite reactive, the order of activity being

fluorine, chlorine, bromine, and iodine. The order of activity decreases with

The Halogen Elements

335

increasing atomic size. The fluorine atom is the smallest and attracts an
electron more strongly than the other halogen atoms, whereas the iodine
atom is the largest and the least reactive halogen. On the other hand, the
fluoride ion is the most stable and the iodide ion the least stable of the
group. An electron, once gained by a fluorine atom, is the most difficult
to remove. The fluopne atom is, in fact, the most powerful of oxidizing
agents in that it has the greatest ability to remove electrons from other
substances. Hence the fluoride ion is the most difficult to oxidize to the
free atom; the iodide ion is the easiest to oxidize and so iodides are better
reducing agents than the other halide ions. Some of the more important
properties of the halogen elements are listed in Table 26-A.

Table 26-A

Properties of the Halogen Elements

Property

Fluorine

Chlorine

Bromine

Iodine

Astatine

Symbol

Molecular Formula

Cl*

Br 2

Atomic Number

Atomic Weight

18.9984

35.453

79.909

126.9044

(210)

Isotopes (mass numbers

19 (100)

35 (75.4)

79 (50.6)

127 ( 100)

and per cents)

37 ( 24.6)

SI (49.4)

Abundance in Earth s

0.07

0.031

1.6 X 10-4

3 X 10-6

Crust, %

Physical State at STP

gas

liquid

solid

Color of Vapor

pale yellow

green yellow

red brown

violet black

Odor

sharp, irri-

intensely

tating

irritating

Density at STP

1.70 g/1

3.21 g/1

3.12 g/ml

4.93 g/ml

Melting Point, °C

-219.6

-101.0

-7.2

113.7

(302)

Boiling Point, °C

-188.2

- 34.7'

58.0

183

Critical Temperature, °C

-129

144

311

553

Critical Pressure, atm

102

Solubility in water at

decomposes

0.090 (g)

0.210 ( l )

0.00133 (s)

20 °C, mole/ liter

Heat of Fusion, kcal/mole

0.061

0.77

1.26

1.87

Heat of Vaporization,

kcal/mole

0.78

2.44

3.59

4.99

Heat of Atomization at

28.61

26,71

25.48

25°C, kcal/g-atom

Ionization Potential, 1st, eV

17.42

13.01

11.84

10.44

Electron Affinity, eV

3.63

3.78

3.54

3.20

Electronegativity

4.0

3.0

2.8

2.5

Covalent Radius, A

0.64

0.97

1.14

1.33

Ionic Radius, A (-1)

1.36

1.87

1.95

2.16

(4-7)

0.26

0.39

0.50

Oxidation States

-1

4T,4-3,+5,-f7

4-1, 4-5

4-1, -f 5, +7

Oxidation Potential, volt

-2.87

-1.360

-1,065

-0.535

(-0.2)

for 2 X" — 2 e — * X 2

336

The Halogen Elements

2. Oxidation States of the Halogen Elements. In addition to an oxidation
state of -1 the halogens form compounds with oxidation states from +1
to +7. Table 26-B lists representative compounds for each oxidation state of
the halogen elements; where there is a blank place no compound with the
corresponding oxidation state is known. Only in the -1 state is the halogen
bond ionic; all other oxidation states involve covalent halogen bonds.

Table 26-B

Oxidation

States of the

Halogen Elements

Oxidation

State

Fluorine

Chlorine

Bromine

Iodine

-1

F- ion

HF; NaF

Cl - ion

HCI; NaCl

Br- ion
HBr; NaBr

1“ ion

HI; Nal

f 2

Cl 2

Br 2

I 2

f 2 o

C1 2 0

HCIO

NaCIO

Br 2 0

HBrO

NaBrO

HIO

NalO

f 2 o 2 (FO)

HC10 2

NaC10 2

C10 2

BrO,

i 2 o, (I0 2 )

HC10 S

NaC10 s

HBrO s

NaBr0 3

HIOa

NaIO s

Cl 2 O s (CIO,)

C1 2 0 7

HC10 4

NaCIO,

HIO,

H,IO s (HIO, • 2H 2 0)
NalO,

3. Chemical Nomenclature. Table 26-B offers an opportunity to survey
a system of nomenclature which the chemist applies to a series of oxygen-
containing adds and their salts. Consider the series of compounds HC1,
HCIO, HC10 2 , HClOa, and HC10 4 . All are acids whose formulas differ
only in the number of oxygen atoms per molecule. The oxidation state of
the chlorine atom also increases with the oxygen content of the molecule.
Based on this difference, the chemical nomenclature uses certain prefixes
and suffixes to distinguish between the compounds.

The prefix for the binary acid, containing only two elements, is “hydro
which implies that no oxygen is present in the acid molecule; thus HC1 is
hydrochloric acid. The suffix “ide” refers to a binary compound between
a metallic atom and a nonmetallic atom; NaCl is sodium chloride. Where

The Halogen Elements

337

the nonmetallic element is oxygen, the compound is an oxide; CaO is
calcium oxide.

HC10 2 is called chlorous acid, and HC10 3 is chloric acid; the suffix “ous”
meaning lower than and the suffix “ic” meaning higher than, both referring
to the relative number of oxygen atoms in the two molecules. HCIO, having
less oxygen than chlorous acid (HC10 2 ), is called %pochlorous acid, the
prefix “hypo” also meaning lower than. HC10 4 , with more oxygen than
chloric acid (HC10 3 ), is called perchloric acid, the prefix “per” indicating
higher than .

For the salt corresponding to the “ous” acid, the “ous” suffix is replaced
by “ite.” NaCIO, the sodium salt of hypochlorous acid, is sodium hypo-
chlorite. Similarly the corresponding salt of an oxygen-containing “ic” acid
is an “ate” salt. This system of nomenclature is summarized below.

Acids

State

Salts

hydro..

. ic

HC1

hydrochloric acid

-1

ide

NaCl

sodium chloride

hypo

ous

HCIO

hypochlorous acid

+ 1

hypo

.. ite

NaCIO

sodium

hypochlorite

ous

hcio 2

chlorous acid

ite

NaC10 2

sodium chlorite

HC10 S

chloric acid

ate

NaClOs

sodium chlorate

per.

hcio 4

perchloric acid

per

ate

NaCIO*

sodium

perchlorate

The names of oxides are also based on the number of oxygen atoms, in
the molecule; thus C1 2 0, chlorine monoxide; C10 2 , chlorine dioxide; and
C1 2 0 7 , chlorine heptoxide.

Fluorine

4. Preparation of Fluorine. Fluorine does not occur free in nature. Com-
bined fluorine is fairly widely distributed in rocks and minerals such as fluor-
spar or fluorite, CaF 2 ; cryolite , Na s AlF«; and fluor-apatite, CaF 2 * 3 Ca 3 (P0 4 ) 2 ,
It is present in bones (0.02-0.65%), in the enamel of the teeth (0.33-0.59%) and
in sea water (2 mg/1). Fluorspar serves as the source of practically all fluorine
compounds. The name fluorine is derived from the fact that fluorspar crystals,
after exposure to bright sunlight, glow or “fluoresce” in the dark.

The preparation of elemental fluorine involves the oxidation of the
fluoride ion.

(3) 2 F“ -2 e F 2

Because fluorine itself i§ the most potent oxidizing agent, all attempts to
liberate free fluorine from fluoride salts by chemical oxidizing agents fail.

338

The Halogen Elements

The element is prepared by the electrolysis of a solution of anhydrous hydro-
gen fluoride, HF, in potassium hydrogen fluoride, KHF 2 .

(4) 2HF -* H*(g) + F 2 (g)

One type of apparatus for the production of fluorine is the electrolytic cell
shown in Figure 26.1. It consists of a copper container with a graphite anode
and a graphite cathode. The V shape of the cell effectively forms a two com-
partment cell so that the H 2 produced at the cathode and the F 2 produced
at the anode do not mix. The electrolytic mixture, 40% HF and 60% KHF 2 ,
is heated to about 95 °C since the melting point of the mixture is about 72°C.
Although attacked by fluorine, copper becomes coated with a thin impervious
coat of copper fluoride, CuF 2 , which resists further action. Other metals,
such as platinum, silver, nickel, Monel, and special steels, behave similarly.

5. Properties of Fluorine. Some of the important physical properties
of fluorine are listed in Table 26-A. Fluorine is the most reactive element. It

reacts directly with all other elements except nitrogen, oxygen, chlorine,
and the noble gases, helium, neon, and argon. However, compounds are
formed with all the elements except for the three noble gases mentioned.
Fluorine displaces the other halogens and most of the other nonmetallic
elements from their compounds with hydrogen and the metals; it reacts
vigorously with water.

(5) F, + 2 MX-* 2 MF + X 2 (where M represents a metallic element

and X represents a nonmetallic element)

(6) 2 F 2 + 2 H 2 0 ~* 4 HF + O a (F 2 0 may also be formed)

The Halogen Elements

339

A jet of steam bums in an atmosphere of fluorine like illuminating gas
Wood, glass, and asbestos ignite and bum in fluorine. In some liquid
rocket propellants fluorine is used as the oxidizing agent. With liquid hydro-
gen as the fuel the reaction is highly exothermic.

(7) H 2 + F 2 2 HF AH = —128 kcal

Other liquid propellant mixtures are hydrazine, N 2 H 4 , and F 2 : and ammonia
NH 3 , and F 2 .

The hydrogen atoms of hydrocarbons, compounds composed only of car-
bon and hydrogen atoms (Chapter 32), can be replaced by fluorire atoms.
Such compounds are called fluorocarbons and are known commercially as
Freon. The simplest is CF 4 , carbon tetrafluoride. Fluorocarbon compounds
are extremely inert and are resistant to the action of acids, bases, and

oxidizing agents. A polymerized tetrafluorethylene

F F
-C-C-

I I

F F

is the plastic.

Teflon, an electric insulator which has the property of not being wetted by
any liquid. Fluorocarbons also find use as aerosol propellants, solvents, fire-
extinguishing agents, and dielectrics.

Sodium fluoride, NaF, and the mineral cryolite, Na 3 AlF 6 , in powdered
form are used as insecticides. Calcium fluoride, CaF 2 , is an important con-
stituent of teeth and bones. Small amounts of fluorides, less than one part
per million, in drinking water are believed to be beneficial in preventing
tooth decay; however, more than three parts per million may cause abnormal
growth of teeth and mottling of the enamel' Cryolite is used as the solvent
in the metallurgy of aluminum by the Hall electrolytic process (page 597).
Uranium hexafluoride, UF 6 , is the only volatile compound of uranium and
has been used in the separation of the isotopes of uranium by gaseous
diffusion ( page 641 ) .

Chlorine

6. Preparation of Chlorine (Greek, ckloros : light green). The most abundant
compound of chlorine is sodium chloride, NaCl. Sea water contains about
3.8% NaCl, 0.3% MgCl 2 , and 0.08% KC1. The Dead Sea and Great Salt Lake
(Utah) contain about 18% NaCl. The evaporation of inland seas has led
to the formation of large deposits of salts; at Stassfurt (Germany) the deposit
is over 1000 feet thick, in several strata.

The ultimate source of chlorine is NaCl. Some methods of preparation
depend upon the use of HC1 which can be prepared by the reaction of
NaCl with concentrated H 2 S0 4 .

(8) NaCl 4- H 2 S0 4 HCl(g) + NaHSO, (at low temperatures)

(9) 2 NaCl + H 2 S0 4 2 HCl(g) + Na,>S0 4 (above 500°C)

The hydrogen chloride, HC1, comes off as a gas and is dissolved in water
to produce hydrochloric acid.

340

The Halogen Elements

Oxidation of chloride ions, Ch, to chlorine molecules, Cl 2 , can be ac-
complished by (A) chemical oxidizing agents and (B) electrolysis of aqueous
solutions of chlorides.

(A) By chemical oxidizing agents: Chlorine can be displaced from chlor-
ides by fluorine, as in Equation 5. Under ordinary conditions hydrogen
chloride and oxygen react very slowly to produce chlorine and water.

(10) 4 HC1 + 0 2 ^ 2 Cl, + 2 H 2 0

When a mixture of HC1* and O s is passed through a chamber containing
a catalyst such as a copper salt, and heated to 400°C, however, the reaction
|s speeded up and proceeds to the extent of about 80%. This is the Deacon
Process, formerly used extensively as a commercial source of chlorine but
now replaced by more economical methods.

A method frequently used in the laboratory involves the reaction of
manganese dioxide, Mn0 2 , an effective and cheap oxidizing agent, with
concentrated hydrochloric acid.

(11) 4 HC1 + Mn0 2 Cl 2 + MnCl 2 4 - 2 H 2 0

The formation of HC1 and its subsequent oxidation can be accomplished in
one operation by treating a mixture of solid NaCl and Mn0 2 with concen-
trated H 2 S0 4 .

(12) 2 NaCl + MnO, + H 2 S0 4 Cl, + 2 NaHSO, + MnSO, + 2 H a O

Other oxidizing agents such as potassium permanganate, KMn0 4 , potassium
dichromate, K 2 Cr 2 0 7 , and lead dioxide, Pb0 2 , may be substituted for Mn0 2
in the reaction with HC1.

(12) 2 KMnG 4 + 16 HC1 5 Cl 2 + 2 KC1 + 2 MnCl 2 + 8 H 2 0

(13) K 2 Cr 2 0 7 + 14 HC1 3 Cl 2 + 2 KC1 + 2 CrCl s + 7 H a O

(14) Pb0 2 + 4 HC1 Cl 2 + PbCl 2 + 2 H a O

(B) By electrolysis of aqueous solutions of Cl" ion: Almost all com-
mercial chlorine is prepared by the electrolysis of aqueous NaCl. The net
electrolytic reaction is

(15) 2 Na+ + 2 Cl" + H 2 0 -* Cl 2 + H 2 + 2 Na+ + 2 OH~

The H 2 and the NaOH, both of which are valuable by-products, can react
with the chlorine if they are allowed to mix. In order to operate the electrolytic
process economically these reactions must be prevented.

(16) H 2 + Cl 2 2 HC1

(17) 2 NaOH + Cl 2 NaCl + NaCIO + H z O

All electrolytic cells for producing chlorine from brine fall into two
classes: (1) diaphragm cells and (2) mercury cells. In a diaphragm cell
the anode and[ cathode compartments are separated by a porous diaphragm
which permits^the migration of ions but prevents the physical mixing of

The Halogen Elements

841

At the anode, chloride ion is oxidized by the' removal of an electron; Cl- - e — > CI°,
Atoms of chlorine combine to form diatomic molecules, Cl 2 . At the cathode,
hydrogen ions are reduced to free H 2 ; 2 H+ -f- 2 e — > H 2 . The sodium ions migrate
to the negative electrode and form a solution of NaOH with the OH- ions which are
produced in high concentration by the displacement of the water ionization equili-
brium caused by the removal of the H+ ions.

figure 26.2. A Typical Chlorine Diaphragm Cell

the electrolytic products. Brine is fed into the anode side of the diaphragm
so that there is always a slow percolation of electrolyte throvjgh it to the
cathode eomparent. Cl 2 is drawn off from the space above the anode, H 2
from the top of the cathode compartment, while a mixture of NaOH and
some excess brine is removed from a layer formed at the bottom of the
cathode compartment. The principle of the diaphragm cell is shown in
Figure 26.2. The NaOH is purified by fractional crystallization. When the
caustic solution is concentrated by partial evaporation, the NaCl, being less
soluble, crystallizes and is removed. The most common type of diaphragm
cell is the Hooker Type S Sell. The anode consists of a number of graphite
blades in parallel while the cathode is a series of steel gauze "fingers” upon
which a felt of asbestos has been drawn by suction; this sheathing of asbestos
is the diaphragm. The mercury cell is an elongated, slightly inclined trough
with a gas-tight cover in which are imbedded graphite anodes. During elec-
trolysis, sodium metal is plated out in a flowing stream of mercury which
acts as the cathode. The formation of an amalgam by the sodium substantially

342

The Halogen Elements

prevents it from reacting with either the water of the brine solution or the
chlorine which is removed from a pipe in the top of the cell. The' amalgam
is subsequently reacted with water in another cell to produce NaOH and H 2 .
In the United States diaphragm cells produce about 75% of all electrolytic
chlorine but mercury cells are becoming increasingly important. The produc-
tion of electrolytic chlorine alone uses over 2% of the national electric power,
7. Properties of Chlorine. Physical properties: These are given in Table
26-A.

Chemical properties: Chlorine combines directly with all metals and with
all of the nonmetals except fluorine, oxygen, nitrogen, carbon, and the noble
gases. By indirect methods, however, chlorine compounds of all the elements
except the noble gases can be prepared. With both metals and nonmetals
the direct reaction with chlorine forms chlorides. The metal chlorides are
ionic whereas the chlorine compounds of the nonmetals are covalent.

(18) 2 Na + Cl, NaCl (2 Na+ + 2 Or)

(19) Cu + Cl* CuCL (Cu*+ + 2 Cl~)

(20) 2 P 4- 3 Cl 2 — > 2 PC1 3 (phosphorus trichloride; with excess Cl 2 ,

phosphorus pentachloride, PC1 S , is formed)

(21) H 2 4- Cl 2 2 HC1 (no perceptible reaction occurs in the dark;
in direct sunlight or ultra-violet light the reaction takes place explosively)

Chlorine also reacts with compounds containing hydrogen.

(22) H 2 S 4- Cl 2 -> 2 HC1 4- S

(23) C„H 6 4- 3 Cl 2 6 HC1 4“ 6 C (hydrocarbons bum in Cl 2 )

In some cases, with methane, CH 4 , for example, chlorine not only combines
with the hydrogen atoms of the compound but also replaces the hydrogen
removed.

(24) CH 4 4- Cl 2 HC1 4“ CH S C1 (methyl chloride)

(25) Ch 3 Cl 4- Cl 2 — * HC1 + CH 2 C1 2 (methylene chloride)

(26) CH 2 C1 2 4- Cl 2 — » HC1 4- CHC1 3 (chloroform)

(27) CHCls + Cl 2 — » HC1 4- CC1 4 (carbon tetrachloride)

In the carbon compounds above both the hydrogen and the chlorine atoms
are covalently bonded to the central carbon, as shown below for Equation 24.
H H

(28) H:GH +:C1:C1:-*H:C1: + h:c:ci:

•* •f *• •• •• • •

H H

When chlorine dissolves in water it reacts with the solvent. The resulting
solution is known as “chlorine water.”

( 29 ) Cl 2 + H.O ^ HC1 + HOC1

The Halogen Elements

343

(30) :C1 Cl: + H:o:H ^ h+ + Cl- + H:*o:ci:

• • •• *• •« •«

The equilibrium constant for this reaction is 5 X 1(H. The low value indicates
that the reaction goes to the right only to a small extent. Thus a solution
of chlorine water contains a mixture of H+ and Cl' ions and molecular
hypochlorous acid in equilibrium with dissolved molecular chlorine and
water. HOC1 is unstable and decomposes slowly, particularly on exposure
to sunlight or a catalyst. For this reason chlorine water is stored in brown
glass bottles..

(31) 2 HOC1 2 H++ 2 Cl~ + O a

As this reaction proceeds the chlorine water equilibrium is displaced to the
right so that, ultimately, the solution of chlorine water contains * only HC1
and some dissolved O a .

Chlorine reacts with substances which yield a high hydroxide ion con-
centration such as NaOH or Ca(OH) 2 .

(32) Cl 2 + 2 OH- Cl- + OCh + H 2 0

With bromide ion, Br~ and iodide ion, I”, chlorine reacts to displace the less
active halogen elements. A distinctive test for Cl' ion is the formation of
a white precipitate of AgCl with silver ion, Ag+.

8. Uses of Chlorine. Chlorine is one of the most important industrial chemi-
cals, current annual production in the United States being about 5 million tons,
60% of the world s output. It is readily liquified by pressure and is stored and
shipped in steel tanks. Large quantities are used in bleaching agents. The
bleaching effect of Cl 2 depends .upon the formation of HOC1 which is a strong
oxidizing agent Dry chlorine does not bleach. Bleaching action is due to HOC1,
produced in moist chlorine by the reaction of Cl 2 and H 2 0, which oxidizes the
coloring substances to colorless compounds. Chlorine is too destructive to
be used on fibers of animal origin (wool, feathers, and silk); even fibers of
vegetable origin (cotton and wood pulp) are weakened. About 65 per cent
of the total chlorine produced is used as a bleaching agent in the manu-
facture of paper pulp, and 20 per cent in the textile industry.

Chlorine is used extensively to sterilize drinking water and water in
swimming pools. All pathogenic microorganisms can be destroyed by chlor-
ine in concentrations as low as 0.3 to 2 parts per million of water.
Among the many uses of chlorine are the manufacture of dyes, synthetic
rubber, DDT (diphlorodiphenyltrichloroethane), drugs, explosives, and im-
portant chlorides of sulfur, carbon, titanium, silicon, phosphorus, arsenic,
aluminum, tin, and antimony.

Chlorine gas is poisonous. It was the first poison gas used in World War I.
Since then more toxic chemical agents have been developed, most of which
are chlorine compounds of organic molecules. So far as is known no poison
gas was used in World War II though all the major powers involved had
ample supplies of chemical agents. Among the more important chemical
warfare agents are phosgene, COCl 2 , a lung irritant like chlorine; dichlor-
diethylsulfide (mustard gas), CIGH 2 -CH 2 -S-CH 2 -CH 2 CI, which causes skin

344

The Halogen Elements

burns and blisters; chloracetophenone, C 6 H 5 -CO-CH 2 Cl, a solid aerosol which
causes nasal irritation and sneezing. The more recent extremely toxic nerve
gases are organic phosphates containing a halogen atom; these inhibit the
mechanism of nerve transmission.

Bromine

9. Preparation of Bromine (Greek, bromos : bad smell). Bromine, Br 2 ,
can be prepared from bromide ion, Br~, by reactions analogous to those
used* for the preparation of chlorine.

(33) 2 Br- + Cl 2 2 Cl- + Br 2

(34) 2 Br + Mn0 2 + 4 H+ Br 2 + Mn*+ + 2 H a O

A large portion of the bromine produced in this country is obtained by
electrolysis of the liquors remaining after the separation of NaCl from the
brines pumped from salt wells. On evaporation of these brines NaCl crystal-
lizes out and the more soluble bromide remains in solution. Since bromide
ion is more easily oxidized to the neutral atom than chloride ion, it is
possible to operate the electrolytic cell at a sufficiently, low voltage to liberate
the bromjne but not the chlorine. The bromine liberated at the anode is
removed by a stream of air, from which it can be separated by condensation.

Sea water contains 70 parts per million (0.0070%) of bromine as Br~
ion. To extract the bromine the sea water is acidified with H 2 S0 4 and then
treated with Cl 2 . The free bromine is removed by blowing air through the
solution and then absorbing it in a solution of sodium carbonate, Na 2 C0 3 .

(35) 3 Br 2 + 3 Na 2 CO s 5 NaBr + NaBrO s + 3 CO,

The bromine is subsequently recovered from solution by treatment with H 2 S0 4
and steaming out the free Br 2 vapor.

(36) 5NaBr + NaBrO s + 3 H 2 S0 4 3 Br 2 + 3 Na 2 S0 4 +3 H 2 0

More than 35,000,000 gallons of sea water are treated daily at one plant
on the coast of North Carolina.

10. Properties of Bromine, Physical properties : See Table 26-A.
Chemical properties : The chemical properties of bromine resemble closely

those of chlorine. The bromine atom is larger than the chlorine atom and
so gains an electron less readily than does the chlorine atom. Bromine is
thus less active than chlorine and its reactions proceed less vigorously, as
evidenced by smaller heats of reaction. Reactions similar to those of chlorine
could be written for bromine with hydrogen, active metals and nonmetals,
hydrocarbons, and water. Bromine vapor also produces serious inflammation
of the respiratory tract and if the liquid comes in contact with the flesh,
severe bums are produced which are very slow to heal.

11. Uses of Bromine. The major industrial use of bromine, current pro-
TOction being about 40,000 tons annually, is in the manufacture of tetraethyl
le^d, (C2H5. f 4 Pb, a component of ethyl gasoline. The process involves the
formation * of ethyl bromide, C 2 H s Br, by the reaction of Br 2 and ethane,

The Halogen Elements

345

C a He, followed by the reaction of C 2 H 3 Br with an alloy of sodium and lead
(written NaPb).

(37) C a H« + Br 2 C 2 H 5 Br + HBr

(38) 4 C 2 H 5 Br + 4 NaPb (C 2 H 5 ) 4 Pb + 4 NaBr + 3 Pb

Ethylene dibromide, C 2 H 4 Br 2 , is added to the gasoline along with the tetra-
ethyl lead so that lead bromide, PbBr 2 , instead of lead oxide, PbO, is formed
when the gasoline bums. The more volatile PbBr 2 is blown out with the
exhaust gases whereas PbO would remain and foul the engine cylinders.

Certain metal bromides, such as those of sodium, potassium, calcium,
and strontium, are used in medicine as nerve sedatives. Silver bromide, AgBr,
is employed in the manufacture of photographic films and papers.

lodttte

12. Preparation of Iodine (Greek, iodes : violetlike). Iodine, I 2 can be
prepared by any of the reactions used for the preparation of chlorine and
bromine. The main source of commercial iodine is sodium iodate, NaI0 3 ,
which occurs in the beds of Chile saltpeter, MaN0 3 . After crystallization of
NaNO s from a solution of Chile saltpeter, the mother liquor is treated with
sodium hydrogen sulfite, NaHSO s , and free iodine is liberated. Upon evapora-
tion of the reaction mixture to dryness, the iodine is recovered and purified
by sublimation.

(39) 2 I0 3 - + 2 HSCV + 3 S0 8 2 - -» I 2 + SO, 2 " H 2 0

13. Properties of Iodine. Physical properties: See Table 26-A. At ordinary
temperatures, iodine is a volatile solid so that, in a closed container, the
characteristic violet color of the vapor is evident. If heated, solid I 2 sub-
limes.

Iodine is very slightly soluble in water, imparting a perceptibly brown
color to the solution. It is quite soluble in alcohol, in ether, and in aqueous
solutions of iodides, forming brown solutions. In chloroform, carbon disulfide,
and many hydrocarbons, iodine forms violet 4 solutions resembling the color
of the vapor. In the violet solutions the iodine is present in molecular
state, I 2 , while in the brown solutions the iodine appears to combine with
the solvent. The dissolving of I 2 in aqueous solutions of 1“ ion is due to the
formation of a tri-iodide ion, I 3 “, which has a red-brown color.

(40) I 2 + I- I 3 - (I, + KI KIs)

Chemical properties : Excluding astatine, iodine has the largest atomic
radius of the halogens. It has the least tendency to gain electrons and is the
weakest oxidizing agent (best reducing agent) and conversely, the I" is
most readily oxidized to I 2 . Iodine is the most metallic halogen and, indeed,
shows a slight electrical conductivity.

When minute quantities of iodine, as little as 0.001 mg, are added to
a starch emulsion, an intense blue-black color develops. This is a specific
test for free I 2 . Starch-iodide paper, which is paper impregnated with starch
and KI, is used in the laboratory as a test for the other halogens and oxidizing
agents as ozone and hydrogen peroxide. Exposure of the moist paper to

346

The Halogen Elements

oxidizing agent releases free I 2 which, with the starch present, turns the
paper blue-black.

In analytical chemistry the reaction of I 2 .and sodium thiosulfate, Na 2 S 2 0 3 ,
is important for the quantitative determination of oxidizing agents.

(41) I 2 + Na 2 S 2 0 3 -» 2 Nal + Na 2 S 4 0 6 (sodium tetrathionate)

An excess of a solution of I" ion is added to the oxidizing agent. The I 2
produced is then titrated with a standard solution of Na 2 S 2 0 3 ; starch may be
used as the indicator since the products are colorless and the end point
.will be indicated by the disappearance of the blue color. The determination
of the quantity of I 2 produced enables the calculation of the initial quantity
of the oxidizing agent. Such a procedure is common in analysis. Where a
substance can not be titrated directly an excess of reagent is added; either
the remaining reagent or a product of the reaction with the reagent is then
titrated.

14. Uses of Iodine. Iodine is best known for its use as an antiseptic in
the form of "tincture of iodine” which is a solution of the element in alcohol.
Iodine is used to produce iodoform, CHI 3 which is also used as an antiseptic.
The manufacture of many organic compounds, such as aniline dyes, requires
the use of iodine. Silver iodide, Agl, is used in photography in the same
manner as AgCl and AgBr.

Iodine is essential to the well-being of man. The thyroid gland contains
a compound of iodine called thyroxin ; lack of the proper amount of iodine
in the diet causes a deficiency of this compound that results in a condition
called goiter. Injection of iodothyrine, obtained from the thyroid of the sheep,
into the human gland often serves as a cure. In some communities where
goiter is prevalent, iodides are added to the drinking water and to table salt
to prevent its development.

15. Interhalogen Compounds. The halogens react with each other to form
interhalogen compounds. In these an atom of the larger halogen element
combines with an odd number of atoms of the smaller halogen element. The
interhalogen compounds are covalent molecules. Liquid IC1 and IC1 3 conduct
electricity indicating that their bonds have some ionic character. The inter-
halogen compounds are unstable and are decomposed by water and by heat.
The compounds are listed below according to type.

ab 3

ab 5

ab 7

GIF

cif 3

BrF

BfF 3

BrF,

IF*

IF,

BrCl

ICl

ICl,

TBr

IBr,

The Halogen Elements

347

QUESTIONS

1. Correlate the relative activities of the halogen elements with their electron
configurations.

2. Predict the values for astatine in the blank spaces of Table 26-A.

3 (a) List the halogens in increasing order of oxidizing strength, (b) What

experimental evidence verifies your list (c) which halogen is the best reducing
agent?

4. Compare the halogens with respect to (a) atomic size (b) ionization potential
(c) electron affinity (d) electronegativity.

5. How are the elemental halogens prepared from the halides in general? Write
the ionic equation.

6. Describe the production of fluorine. Why is it impossible to liberate fluorine
from an aqueous solution of a fluoride?

7. What are the chief sources of the halogen elements? Write equations for-
the preparation of each halogen element.

8. Which has a higher percentage of fluorine, calcium fluoride or cryolite?

9. Describe the operation of the electrolytic cells for making chlorine. Write the
electrode equations.

10. Write equations for laboratory methods of preparing the halogen elements.

11. Write equations to show how chlorine reacts with the following: (a) Mg and
Fe (b) P 4 and S 8 (c) CH 4 and C 2 H 4 (d) H 2 0 and KOH.

12. List the oxidation states of chlorine and a compound which represents each
state.

13. Give the formula and the name of the acids of chlorine, and of the potassium
salt of each acid.

14. A water solution contains both Cl - and I" ions. How could you verify the
presence of each ion?

15. Write equations for the manufacture of tetraethyl lead. Why is ethylene
dibromide added to gasoline containing tetraethyl lead?

16. Why is iodine more soluble in a solution of KI than in H 2 0?

17. Which interhalogen compound of type AB on page 346 is most polar? Explain.
Draw the electron dot structure for C1F and for C1F*.

18. (a) What weight of chlorine can- be obtained from 20 g of CaCl 2 (b) what
volume will this chlorine occupy at 70° C and 686 mm pressure? Ans : (a) 12.8 g

19. To produce 25 liters of Cl 2 at STP by reaction of HC1, what weight is re-
quired of (a) KMn0 4 <b) K 2 Cr 2 0 7 ? Ans: (a) 70.5 g

20. What volume of Cl 2 , measured at 20 °C and 800 mm pressure, is required
to displace all the bromine from 490 ml of a Br~ ion solution?

21. Sea water contains 70 parts per million of bromine by weight, Assuming
100% efficiency, what weight of sea water must be treated to obtain 140 lb
of Br 2 ?

22. Concentrated hydrochloric acid, density 1.20 g/ml and containing 40% HC1
by weight, is reacted with Mn0 2 to produce CL. If 100 ml of the acid are
consumed, (a) what weight of Mn0 2 reacted (b) what volume of Cl 2 at
STP was produced?

348

The Halogen Elements

23. (a) What quantity of electricity is required to produce 89.6 liters of Cl
at STP in a diaphragm cell (b) What weight of NaOH is produced simultan-
eously?

24. Balance the following equations by an oxidation-reduction method. In each
case, name the reducing agent and calculate its equivalent weight.

(a) KBr + Mn0 2 + H 2 S0 4 KHS0 4 + MnS0 4 + Br 2 + H*0

(b) Cl 2 + NaOH NaCl + NaCIO + H 2 0

(d) Br 2 4* Na 2 C0 3 — > NaBr H- NaBrO a -f C0 2

25. What pressure will 5.0 grams of gaseous Freon, CC1 2 F 2 , exert at 20 Q C in a
250 ml container?

The Hydrogen Halides

The halogen elements form not only simple hydro-acids but also oxygen
acids (oxyacids) in which the oxidation state of the halogen ranges from
-1 to +7. The properties of the hydrohalogen acids are given in Table 27-A.

Table 27-A

OF THE

Hydrohalogen Acids

Formula

HCl

HBr

Molecular Weight

20.0

36,5

80.9

128

Physical State at STP

liquid

gas

Color of Vapor

Odor

Density at STP, g/1

all the vapors are colorless
all have sharp, irritating odofrS
1.63 3.61

5.71

Melting Point, °C

-83.1

-114.2

-86.9

-50.8

Boiling Point, °C

19.9

-84.9

-66.7

-35.4

Critical Temperature, °C

151

Critical Pressure, atm

Solubility in water at

20° C, g/100 g solution

Constant Boiling Mixture
(a) at 760 mm, °C

120

110

126

127

(b) weight % of acid

35.4

20.2

47.8

57.5

1.14

1.10

1.49

1.70

Dielectric Constant of

Liquid

Heat of Formation, AH,
kcal/mole

-64.0

-22.0

-13.5

+0.8

Percent Dissociation at

1000°C

1 X 10- 18

0.144

350

The Hydrogen Halides

1. Methods of Preparation. (A) The direct combination of the halogen
element with hydrogen:

( 1 ) jj 2 + X 2 2 HX where the symbol X refers to a halogen atom

(B) The double displacement, or metathesis, of a metal halide, MX, and
a strong acid:

(2) MX + H + HX + M+ (CaF 2 + H 2 S0 4 -*2 HF(g) + CaS0 4 )

When H 2 S0 4 is used only HF and HC1 can be prepared pure by this method
because HBr and HI react with concentrated H 2 S0 4 , an oxidizing agent
sufficiently potent to oxidize Br~ and T ions to the free elements.

(3) 2 HBr + H 2 S0 4 Br 2 + SO, + 2 H z O

(4) 8 HI + H 2 S0 4 4 I, + H 2 S + 4 H.O

Instead of H 2 S0 4 some acid which does not react with HBr or HI and which
is nonvolatile under the conditions of the experiment must be employed.
Phosphoric acid, H 3 P0 4 , is a nonoxidizing and nonvolatile acid which meets
these requirements.

(5) KBr + H 3 P0 4 HBr(g) + KH 2 P0 4

(6) PX« + 3 H 2 0 -> 3 HX(g) + HJPO,j

(PBr» + 3 H.O 3HBr(g) + H 3 P0 3 )

The phosphorous acid, H 3 P0 3 , is nonvolatile and remains in solution. Phos-
phorus halides can be prepared by direct combination of phosphorus and
the halogen element.

(D) Reaction with compounds of hydrogen

(7) H 2 S + X 2 2 HX + S

(8) H 2 0 + X 2 HX + HOX (with all halogens except F 2 )

(9) CH 4 + X 2 HX + CH S X (with hydrocarbons)

2. General Properties of the Hydrogen Halides. The names hydrogen
fluoride, hydrogen chloride, hydrogen bromide, and hydrogen iodide refer
to the compounds as gases; the aqueous solutions are called hydrofluoric
acid, hydrochloric acid, hydrobromic acid, and hydriodic acid.

The hydrogen halides are colorless gases having sharp pungent odors.
They are extremely soluble in water and fume in moist air. At 0°C, 507
volumes of HC1 will dissolve in one volume of water; commercial concen-
trated hydrochloric acid is 12M or 37% HC1 by weight. HBr and HI are
even more soluble than HC1. All of the adds form constant boiling mixtures,
azeotropic mixtures in which neither the boiling point nor the composition
changes during distillation (Table 27-A).

, « *

The bond between the hydrogen and the halogen atoms is covalent, .

The Hydrogen Halides

351

If heated sufficiently the hydrogen halides decompose into their elements.
As expected HI is the least stable; its heat of formation is lowest and less
energy is required to decompose it. The dry compounds are not very reactive.
The addition of a small amount of water, with the consequent production
of H s O+ and X- ions, greatly increases the activity: HX+ H 2 O H a O+ + X~.
Except for HF, the aqueous solutions are strong acids and exhibit the typical
chemical properties pertaining to acids and also of the respective halide
ions. HF is a weak acid with an ionization constant of 7 X 10 -4 .

3. Properties and Uses of Hydrogen Fluoride, Hydrogen fluoride mole-
cules are not monomolecular but are associated due to hydrogen bonding
between two fluorine atoms, forming (HF) U . In the liquid state at low
temperatures, n may be as high as 6 and in the gaseous state as low as 2.
This tendency to associate is the result of the extreme electronegativity of the
fluorine atom; electrons are withdrawn from the hydrogen atom, leaving
essentially the positive proton to act as the hydrogen bond. The other hydro-
gen halides show no tendency to associate. The association of HF accounts
for those properties, such as boiling point, whose values are “out of line”
with those of the other halides. It also accounts for the formation of the

«• ••

hydrogen fluoride ion, HF a “ or [IFl H :F:]~, if it is assumed that (HF) a
ionizes as H(HF 2 ).

Hydrogen fluoride is distinguished by its action on silicon compounds.
With quartz, SiO*, and glass (a mixture of silicates), the following reactions
occur:

(10) Si0 2 + 4 HF S«F 4 (g) + 2 H z O

(11) CaSi0 3 + 6 HF SiF 4 (g) + CaF* + 3 H 2 0

Hydrogen fluoride is used in the chemical analysis of silicate material and
for etching glass. In etching, the glass is covered with a thin film of paraffin
and the design to be produced is scratched through the film with a stylus.
The surface is then exposed to the action of the gas or of a solution of
the add. The paraffin coating protects the unexposed glass. Burets, ther-
mometers, and other graduated glassware are engraved in this manner;
electric light bulbs are frosted by etching with HF. The fumes of hydro-
fluoric acid are intensely poisonous; in contact with the skin the concentrated
solution produces burns and sores which heal with difficulty.

4. Properties and Uses of Hydrogen Chloride. Dry hydrogen chloride
reacts with the more active metals to form the metal chloride and hydrogen.
The reactions of hydrochloric acid written below are typical.

(12) Zn + 2 HC1 ZnCl 2 + H 2 (with metals more active

than hydrogen)

(13) Fe 2 0 3 + 6 HC1 2 FeCl s + 3 H 2 Q (the dissolving of metal

oxides)

(14) NaOH + HC1 NaCl + H 2 0 (neutralization of bases)

352

The Hydrogen Halides

(15) Na 2 CO s + 2 HC1-* 2 NaCl + C0 2 + H 2 0 (reaction with carbonates)

(16) AgNOs + HC1 -» AgCl + HNO s (precipitation of insoluble

chlorides)

The first four reactions are characteristic of any strong acid; the reaction
with AgN0 3 is specific to the Cl" ion and Ag + ion. Silver chloride, AgCl,
is a white insoluble substance and the reaction serves as an analytical test
for both the Cl' and Ag+ ions. In general all metal chlorides are soluble
with the exception of those of silver, lead, and mercury (I); AgCl, PbCk,
and Hg 2 Cl 2 .

Next to sulfuric acid, hydrochloric acid is the most widely used acid in
industry. It is used in making pickling baths for removing the oxide coating
from iron before plating or galvanizing, for the production of glucose from
starch, in the preparation of glue from animal tissue, and in the manu-
facture of dyes, drugs, and many metal chlorides. Hydrochloric acid plays
an important part in the digestion of certain foods in the stomach. Pepsin,
which is present in the gastric juice, does not function in the digestive process
unless HC1 is also present. In the gastric juice HC1 is produced from NaCl
that is taken into the body with food.

5. Properties and Uses of Hydrogen Bromide and Hydrogen Iodide. The
properties of HBr and of HI are very much like those of HC1. Both form
strong acids but neither has any extensive uses not satisfied by the cheaper
HC1. Both AgBr and Agl are insoluble yellow compounds useful as specific
tests for the Br and 1“ ions in analytical chemistry,

6. Oxides of Chlorine. The important oxides of chlorine are Cl 2 0, C10 2 ,
and C1 2 0 7 (Table 26-B). Their electronic structures are shown below; each
oxygen-chlorine bond is a single covalent bond.

•• ••

:ci:o:ci; :o:ci:o: :o:Ci:o:Gl:q:

**:o :**:o :**

•• ••

Chlorine Monoxide Chlorine Dioxide Chlorine Heptoxide

Chlorine monoxide, C1 2 0, is an unstable yellow-brown gas. It is soluble
in water and is the anhydride of HOC1. It can be prepared by distillation of
HOC1 under reduced pressure or by passing Cl 2 over dry mercury (II)
oxide, HgO.

(17) £ Cl 2 + HgO C1 2 0 + HgCl 2

The gas decomposes explosively when heated even slightly but with proper
care it can be liquified at about 3°C; its freezing point is -116°C.

Chlorine dioxide, C10 2 , is a red-yellow gas, liquifying at — 10°C and freezing
at -59°C. It is formed whenever chloric acid, HC10 3 , decomposes.

(18) KCIQs + H 2 S0 4 HClOa + KHSO,

(I®)- 4 HClOa 4 C10 2 + O a + 2 H z O

The Hydrogen Halides

353

The gas is very explosive, decomposing into Cl 2 and O s even if warmed
slightly. Because of the violence attending this reaction , it is dangerous to
add concentrated H 2 S0 4 , or any strong acid , to a dry chlorate . C10 2 can
be prepared safely by die reaction of KC10 3 , and the weak oxalic acid,
H 2 C 2 0 4 ; by the electrolysis of a solution of sodium chlorite, NaCIO; and
by the action of Cl 2 on moistened NaC10 2 .

(20) 2 KC10 3 + 2 H 2 C 2 0 4 2 C10 2 + K 2 C 2 0 4 +2 C0 2 + 2 H a O

(21) 2 NaC10 2 + H a O 2 C10 2 + H 2 + 2 Na+ + 2 OH’

(22) Cl 2 + 2 NaC10 2 2 C10 2 + 2 NaCl

C10 2 is one of the rare molecules containing an odd number of electrons;
it is paramagnetic and several resonance forms probably exist. With C10 2
as starting material all compounds of chlorine can be prepared. When
diluted with air, C10 2 is safe to handle and is used commercially to bleach
cellulose products and for sterilizing water.

Chlorine pentoxide, C1 2 0 7 , is a colorless, oily liquid that boils at 82°C.
It, too, is explosive when heated above its boiling point or when struck
roughly. It is prepared by removing the elements of water from perchloric
acid, HC10 4 , by means of a strong dehydrating agent such as phosphorus(V)
oxide, P 4 O 10 .

(23) 4 HC10 4 + P*O 10 2 Cl 2 0 7l + 4 HPO*

The other halogen oxides are of limited importance.

7. Oxygen Acids of the Halogens and Their Salts. Only for chlorine
is there known a complete series of oxygen acids and their salts with oxida-
tion states of +1, +3, +5, and +7. The electron dot formulas for the oxygen
acids of chlorine are shown below; the structure of HC1 is included
for comparison. Even though for pedagogical reasons the formulas are some-
times written differently, e.g., HOC1 vs. HCIO, it should be noted that in all

•• • *

:0: :o:

H:ci:

h:o:ci:

h:o:ci:o:

HKCBCKOi

H:0:cl:o:

:o:

*•

Hydrogen

chloride

Hypochlorous

acid

Chlorous

acid

Chloric

add

Perchloric

add

cases the hydrogen atom is bonded directly to an oxygen atom. This H— O bond
is a covalent bond with appreciable ionic character; the Cl— O bond is essential-
ly covalent. In salts the hydrogen atom is replaced by a metallic atom and
the resulting bond is ionic.

The strength of the oxygen acids as adds increases with the number of
oxygen atoms in the molecule and hence also with the oxidation state of the
chlorine atom. Thus HC10 4 is a stronger acid than HCIO. The ionization
constant of HCIO is 3 X 10" 8 and that for HC10 2 is 1 X 10~ 2 . HC10 a is a
strong acid and HC10 4 is perhaps the strongest of all acids. With increasing
positive oxidation state of the chlorine atom, the electrons binding the

354

The Hydrogen Halides

hydrogen atom are drawn away from it. The H— O bond becomes relatively
more ionic and the tendency to yield the proton increases. This relationship
between acid strength and oxygen content is applicable to a series of oxygen
acids of any element; thus H 2 S0 4 is a stronger acid than H 2 S0 3 . The thermal
stability of the oxygen acids and their salts also increases with increasing
number .of oxygen atoms. On the other hand, the strength of an oxygen
acid as an oxidizing agent depends inversely upon the number of oxygen
atoms in the molecule; HCIO is a stronger oxidizing agent than HCIO*.

8. The Hypohalous Acids. Except for fluorine the reaction of the
halogen elements with water is

(24) X 2 + H 2 0 H+ + X- + HOX

The extent of the reaction is small for Cl 2 , less for Br 2 , and still less for I 2 .
If a base such as NaOH is used, at low temperatures the salts NaX and
NaOX are formed and the reaction proceeds to a greater extent.

(25a) X 2 + 2 OH” X- + OX” + H 2 0

(25b) Cl 2 + 2 NaOH NaCl + NaOCl + H*0

A solution containing NaCl and NaOCl, truly a mixture of Na + , Cl”, and
OCl~ ions, is prepared by electrolyzing a solution of NaCl under conditions
which allow the products formed at the electrodes, Cl a and NaOH, to interact.

The hypohalous acids have been prepared only in aqueous solution. They
are very weak acids, the order of decreasing acidity being HOC1, HOBr, HOI;
HOI is actually slightly basic. This gradation in acidity is in agreement with
the increasing atomic size and decreasing electronegativity of the halogen
elements. These acids are strong oxidizing agents; the oxidizing strength
follows the same gradation. The hypohalous acids are unstable to heat and
to light.

(26) 3 HOX ^2H+ + 2X' + HXO s (upon heating)

(27) 2 HOX 2 H+ + 2 X' + 0 2 (with sunlight or a catalyst)

Bleaching powder, CaOCl 2 , commercially called "chloride of lime” is pre-
pared by passing Cl 2 over slaked lime, Ca(OH) 2 .

(28) Cl 2 + Ca(OH) 2 CaOCl 2 + H 2 0

The formula for bleaching powder indicates that it is a chloride-hypochlorite
/Cl

compound: Ca^ When treated with an acid, HOC1 is liberated.

^OCl

High test hypochlorite, H.T.H., is calcium hypochlorite, Ca(OCl) 2 .

9. The Halous Acids. Chlorous acid, HC10 2 , is the only halogen acid
known of the type HX0 2 . When C10 2 is dissolved in water both HClO*
and HC10 3 are formed; dissolved in NaOH the sodium salts, NaC10 2 and
NapO s , are produced.

(29) 2 C10 2 + H 2 0 HC10 2 ■+■ HClOs

The Hydrogen Halides

355

10. The Halle Acids. Halic acids, HX0 3 , can be prepared by the reaction
of a halate salt and an acid. A solution of chloric acid, HC10 8 , is obtained
by the addition of dilute H 2 S0 4 to a solution of barium chlorate, Ba(G10 3 ) 2 .
The precipitation of BaSG 4 , which can be removed by filtration, drives the
reaction to the right.

(30) Ba 2+ + 2 C1CV + 2 H+ + S0 4 2 - -» 2 H+ + 2 C10 8 " + BaS0 4 (s)

Iodic acid, HIO s , is usually made by the oxidation of I 2 with concen-
trated HNO a .

(31) I 2 + 10 HNOa 2 HIOs + 10 NO* + 4 H 2 0

HC10 S and HBrO a are known only in aqueous solution but the more stable
HI0 3 can be isolated as white anhydrous crystals. HC10 3 can be concen-
trated to about 40% by evaporation below 40° C. If heated in an attempt
to concentrate it further it decomposes violently into C10 2 . It is an active
oxidizing agent but cannot be prepared for such use due to its instability.
The same results can be obtained by using an acidified solution of KC10 8
as an oxidizing agent. The heating of HI0 3 produces its anhydride, I 2 O s ,
a white solid stable up to 300 °C, above which it decomposes into its elements.

Potassium chlorate, KC10 3 „ is the most important salt of chloric acid.
It is prepared by passing Cl 2 into a hot concentrated solution of KOH.

(32a) 3 Cl 2 + 6 OH- C1CV + 5 Cl- + 3 H a O

(32b) 3 Cl 2 + 6 KOH KCIO, + 5 KC1 + 3 H 2 0

The electrolysis of a hot solution of KC1 with provision for the mixing of
the electrolytic products, Cl 2 and KOH, is a more economical method of
producing KC10 3 by the same chemical reactions. Upon partial evaporation
of the fi*al solution the KC10 3 , which is less soluble than KC1, crystallizes.
When heated, KC10 8 decomposes in two ways.

(33) 4 KCIO* -» 3 KC10 4 + KC1 (at low temperature in the absence of

a catalyst)

(34) 2 KCIO3 2 KC1 + 3 0 2 (with a catalyst as Mn0 2 )

KCIO3, or its acidified solution, is used as an oxidizing agent both in the
laboratory and industrially for the manufacture of matches and fireworks.
Mixtures of KC10 3 , by intent or by accident, with carbon, sulfur, sugar, or
other readily oxidizable substances, are dangerously explosive.

11. The Perhalic Acids. Perchloric acid, HC10 4 is prepared by treating
a perchlorate with an acid.

(35) KCIO, + H a S0 4 -» HC10 4 + KHS0 4

HC10 4 is a colorless liquid and is the only halogen oxygen acid which can
be obtained in the pure state. Heated above 90°C, it also decomposes
explosively but by distillation of an aqueous solution under reduced pressure
the pure substance can be prepared. The potassium salt, KC10 4 , is pro-
duced by heating KC10 3 without a catalyst at low temperature. It can also

356

The Hydrogen Halid®

be. formed by prolonged electrolysis of a KC1 solution; this results in the
anodic oxidation of the C10 s " first formed to CIO*

The perchlorates are oxidizing agents but less vigorous than the chlorates.
KCIO* is used as an oxidizing agent in rocket propellants. With few ex-
ceptions perchlorate salts are highly soluble. KCIO* is sparingly soluble
and is used in analytical procedures for the detection and quantitative
determination of potassium. Anhydrous Mg(C10 4 ) 2 and Ba(C10 4 ) 2 absorb
water so avidly that they are employed as drying agents, almost equal in
efficiency to P 2 O s .

Periodic acid, HIO*, can be prepared by the reaction of I 2 and HCIO*.
(36) I 2 + 2 HCIO* 2 HIO* + Cl 2

Evaporation of the resulting solution yields a white deliquescent solid,
H 5 IOe, another form of periodic acid. Because of its larger size the iodine
atom can coordinate with six oxygen atoms. Attempts to prepare HBrO* or
its salts have so far been unsuccessful.

QUESTIONS

1. From each of the following sets, select the substance which best fits the re-
quirement specified: Explain your choice.

(1) Biggest atom Cl, I, Br, F

(2) Smallest ionization potential Cl, I, Br, F

(3) Best reducing agent Cl", I", Br% F~

(4) Strongest acid HCIO.,, HOC1, HC!O a , HCIO*

(5) Most stable compound CaF s , CaBr 2 , CaCl 2 , Cal 2

(6) Weakest acid HC1, HBr, HI, HF

(7) Largest electron affinity Cl, I, Br, F

(8) Smallest relative electronegativity Cl, I, Br, F

(9) Highest degree of hydrogen bonding HC1, HBr, HI, HF

. (10) Paramagnetic C10 4 ~, C1 2 0, C10 2 , GO<f

2. State the names and formulas of the oxygen acids of chlorine and of their
sodium salts.

3. Write the formulas of the following compounds: (a) potassium hypobromite
(b) calcium bromate (c) silver chloride (d) silver chlorate (e) magnesium
perchlorate (f) sodium iodate.

4. Write equations for the preparation of the hydrogen halides.

5. Why cannot H 2 S0 4 be used in preparing .HBr and HI from their salts?

6. Write equations few: the preparation from Cl 2 of (a) potassium chlorate and
(b) calcium perchlorate.

7. Draw the electron dot structures of (a) H 2 F 2 (b) C10 2 (c) C10 2 " ion (d)
HC10 3 . Which of these might have resonance hybrid forms?

8. Hydrofluoric acid etches glass, yet is said to be a weak acid. Explain this
statement.

9. Will an aqueous solution of sodium chlorite be acidic, neutral, or basic?
Explain.

10. In the reaction: I 2 + 2 HC10 4 — > Cl 2 -4- 2 HIO*, explain why iodine ^dis-
places” chlorine?

The Hydrogen Halides

357

11. Balance the following equations by an oxidation-reduction method. In each
case calculate the equivalent weight of the oxidizing agent.

(a) HOC1 + H 2 0 +1* -» HI0 3 + HC1

(b) I 2 + Q, + H 2 0 HIO s + HC1

12. What weight of KBr is required to make LOO kg of AgBr?

13. If 2000 liters of Cl 2 , measured at STP, react with hot KOH solution, (a) what
weight of KC10 3 is produced and (b) what weight of KOH reacted?

14. What weight of iodic acid can be* made from 200 g of iodine?

15. What volume of 0.10N HBrcan be prepared by the hydrolysis of 20gqf PBr 3 ?

16. Ten grams of PC1 3 are hydrolyzed and the HC1 produced is dissolved in 250 g
of water. The solution of HC1 is used to generate C0 2 by reaction with CaCO s .
Assuming no losses what volume would be the C0 2 thus formed occupy at STP?

The Elements of Group VIB
The Sulfur Family

The elements of Group VIB of the Periodic System are oxygen, sulfur,
selenium, tellurium, and polonium. Oxygen has already been considered
separately. Being the first member of the group, it has the smallest atomic
size and the highest value of electronegativity so that in many respects it
differs markedly from the other elements of the group. The remaining
elements are called the Sulfur Family. The electron configurations of these
elements are given below; other properties are listed in Table 28-A.

Atomic

Element

Number

6s 6p

Oxygen

Sulfur

2 4

Selenium

2 6 10

2 4

Tellurium

2 6 K)

2 6 10

2 4

Polonium

2 6

2 6 10

2 6 10 14

2 6 10

2 4

1. General Properties of the S rlfur Family. All the elements have six
valence electrons, ns 1 2 * np 4 ; two of the p orbitals contain unpaired electrons. The
elements are nonmetals but as the size of the atbm increases from oxygen to
polonium, electrons are retained less strongly and some metallic characteristics
appear in the larger elements of the group. Thus sulfur is a yellow solid and an
excellent electric insulator; tellurium has a silvery metallic appearance, and
both selenium and tellurium show a small electrical conductivity. Oxygen
alone of the group is a gas at room temperature; the other elements are all
solids. All exist in allotropic forms.

The principal oxidation states of the Group VIB elements are -2, +*4,
and +6. To attain an electron configuration of eight electrons in the outer-
most shell the elements react to gain two electrons, thereby exhibiting an

oxidation state of —2. In the oxidation states of +4 and 4- 6 only covalent bonds

are formed by the Group VIB atoms. As with the halogens the tendency to

The Elements of Group VIB: The Sulfur Family

359

Table 28- A

Properties of the Group VIB Elements

Property

Oxygen

Sulfur

Selenium

Tellurium

Polonium

Symbol

Molecular Formula

o 2

S R

Se 8

Te s

Atomic Number

Atomic Weight

15.9994

32.064

78.96

127.60

210

Isotopes (mass numbers
and percents)

16(99.91)

32(95.02)

74 ( 0.96)

120 ( 0.09)

17( 0.01)

3S( 0.75)

76 ( 9.12)

122 ( 2.46)

18( 0.08)

34( 4.21)

77 ( 7.50)

123 ( 0.87)

36( 0.02)

78 (23.61)

124 ( 4.61)

80 (49.96

125 ( 6.99)

■

82 ( 8.85)

126 (18.70)
128 (31.79)
130 (34.49)

Abundance in Earth’s

47.3

0.052

9 X 10-«

1.8 X 10-7

3 X 10 -14

Crust, %

Physical State at STP

colorless

yellow

gray or red

gray

gas

solid

solid i

solid

Allotropes

ozone

rhombic

hexagonal (gray)

hexagonal

monoclinic

mcnoclinic (red)

amorphous

Melting Point, °C

-218.4

114.5 (r)
119.3 (m)

217.4 (h)

449.8

Boiling Point, °C

-183.0

444.6

684.8

1,390

Density at STP, g/ml

2.07 (r)
1.96 (m)

4.80 (h)

4.50 (m)

6.24

Heat of Fusion, kcal/mole

350

1600

4270

Heat of Vaporization,

814

2520

4350

kcal/mole

Ionization Potential, 1st, eV

13.61

10.36

9.75

9.01

8.4

Electronegativity

3.5

2.5

2.4

2.1

Covalent Radius, A

0.74

1.04

1.17

1.37

1.53

Ionic Radius, A (-2)

1.40

1.84

1.98

2.21

(+6)

0.29

0.42

0.56

Oxidation States

-2, -1

-2

+1, +2

-f-l, 4-4, 4-6

4-4, 4-6

+4, +6

Oxidation Potential, volt for
H 2 X X2- + 2H+ + 2e

-1.23

4-0.40

4-0.72

Key to symbols: (r) = rhombic; (m) =monoclinic; (h) — hexagonal.

gain electrons and so to act as oxidizing agents decreases with increasing
atomic size. Hydrogen telluride, H 2 Te, in which the tellurium atom has an
oxidation state of -2, is a good reducing agent. Its oxidation potential of
-^•0.72 volt indicates that Te 2- is almost as good a reducing agent as is zinc
metal.

The atomic size of an element in the sulfur family is larger than that
of its adjacent halogen, one greater in atomic number. Hence these elements
are less electronegative than the halogens; tellurium is no more electronegative
than hydrogen. In combination with more electronegative elements, the
Group VIB elements show positive oxidation states. Oxygen, being second
only to fluorine in electronegativity, shows positive oxidation states only

360

The Elements of Group VIB: The Sulfur Family

in combination with fluorine: +2 in 0F 2 (F 2 0) and +1 in 02F 2 (F 2 0 2 ). For
the other members of the group the oxidation states of +4 and +6 are
most common. In the halogen family, Group VIIB, wherein an element
has seven electrons in its valence shell, the principal oxidation states have
odd values: -1, +1, +3, +5, and +7. In the sulfur family the most stable
oxidation states are even: +2, +4, and +6. The difference of two units
or a multiple thereof is due to the greater stability of existence and the
interaction of electrons in pairs. In Table 28-B the oxidation states of the
.Group VIB elements are listed together with typical compounds for each state.

Table 28-B

Oxidation States of the Sulfur Family

Oxidation

State

Sulfur

Selenium

Tellurium

-2

O 2 "ion; Na g O
H 2 0

S 2 " ion; Na 2 S

h 2 s

Te 2 ~ ion; Na.,Te
H 2 Te

-1

Na 2 0 2

0 2

S(S 8 )

Se(Se 8 )

•Te(Te 8 )

o 2 f 2

SO(u)

TeO

S*0 3 (u)

S0 2 ; HjSOj,

SO :i -~ ion; Na,SO :J

Se0 2 ; H 2 SeO a
SeO s 2 ~ ion; Na 2 SeO a

TeO a ; H 2 Te0 3

SO a ; H 2 S0 4

S0 4 2 ~ ion; Na 2 S0 4

Se0 3 (u); H 2 Se0 4
Se0 4 2 - ion; Na 2 Se0 4

Note: the symbol (u) represents an unstable species.

Sulfur

2. Occurrence of Sulfur. Sulfur has been known from the beginning of
history. It is mentioned in the Bible and in early classical records. In the
free state sulfur is found in the volcanic regions of Sicily, Japan, and Mexico;
extensive subterranean deposits occur in Louisiana and in Texas and off
their shores. In the combined state sulfur is more widely distributed as
sulfides and sulfates. Many important metallic ores are sulfides; among them
are galena, PbS, cinnabar , HgS, zinc blende , ZnS, and iron pyrites , FeS 2 . An
important sulfate is gypsum, CaS0 4 • 2 H 2 0.

3. Physical Properties (See Table 28-A). Sulfur is a brittle, pale-yellow
solid without taste or odor. It is insoluble in water but soluble in carbon di-
sulfide, CS 2 , from which it may be crystallized at room temperature in the
rhombic form (Figure 28,1 a). When sulfur is heated some unusual physical
transformations take place. Ordinary sulfur melts at 114.5°C, forming just
above the melting point a thin straw-colored liquid. On further heating the
Bqoid becomes darker in color and more viscous. At about 200° G the

The Elements of Group, VIB: The Sulfur Family

361

(b)

(a) A crystal of rhombic sulfur, the stable variety at temperatures below 95.5°C.

(b) Crystals of monoclinic sulfur, stable between 95.5 °C and 119.3°C.

Figure 28.1. The Solid Forms of Sulfur.

liquid sulfur is dark brown and so viscous that it shows little tendency to
flow even from an inverted container. With continued heating the liquid
becomes more fluid again and boils at 444.6° C. When liquid sulfur, just
above its melting point, is cooled slowly, long transparent needle-like
crystals of sulfur are formed. This variety of solid sulfur is known as mono-
clinic sulfur (Figure 28.1 b). If liquid sulfur is heated almost to the boil-
ing point and the resulting dark brown liquid is cooled rapidly by pour-
ing into cold water, a dark rubbery amorphous mass known as plastic
sulfur is produced. Plastic sulfur is insoluble in CS 2 .

4. Allotropes of Sulfur. At one atmosphere pressure and below 95.5° C
the variety of solid sulfur which is stable is rhombic sulfur; above 95.5° C
the stable form is monoclinic sulfur. At the transition temperature, 95.5°C,
both forms can be in equilibrium. Pure monoclinic sulfur melts at 119.3°C.
If rhombic sulfur is kept at a temperature between 95.5° C and 119.3 °C,
ultimately the rhombic variety will be transformed completely into the
monoclinic form but the change is very slow, involving, as it does, a solid
state transition. For this reason monoclinic sulfur is more readily obtained
when liquid sulfur is cooled from temperatures just above its melting point;
the first transformation is to the monoclinic form. If monoclinic sulfur is
cooled below 95.5°C it then changes slowly to the rhombic form. If rhombic
sulfur is heated rapidly to temperatures above the transition temperature,
95.5° C, there is insufficient time to establish equilibrium with the mono-
clinic form. Thus the melting point observed is that of a superheated rhombic
sulfur, 114.5°C, rather than that of the monoclinic form, 119.3°C.

Both rhombic and monoclinic sulfur, and the yellow liquid just above
the melting point, are composed of Sg molecular units. Different geometric
orientation of these molecules within the crystal give rise to the rhombic and
monoclinic forms. In the S 8 molecule, X-ray analysis indicates that the sulfur
atoms are arranged in a puckered eight-membered ring with a single covalent
bond between each sulfur atom (Figure 28.2 a).

Usually the viscosity of a liquid decreases as the temperature rises. This
is true for liquids whose molecules are unchanged during the heating process.
The increase in the viscosity of liquid sulfur with increased temperature is
ascribed to the fupture of the S 8 rings to form chains of sulfur atoms

362

The Elements of Group VIB: The Sulfur Family

(a) The eight sulfur atoms are arranged in a staggered octagonal ring with four
of the atoms in one plane and the other four in a parallel plane.

(b) The electron structure of the S 8 molecule; the sulfur atoms are joined by single
covalent bonds.

Figure 28.2 . Structure of the Sulfur Molecule, S 8 .

(Figure 28.2 c). Entanglement of these chains during liquid flow results
is an increased resistance to flow and hence an increased viscosity. Further-
more, the end atoms of the chains have a single unpaired electron so that
terminal atoms of two chains can form a covalent bond and join to produce
longer chains; *S— S— S u — S— S<\ At temperatures above 200°C the increased
kinetic energy of the molecules reduces the size of these chains, resulting
in a corresponding decrease in viscosity. The brown color of liquid sulfur
is ascribed to the unpaired electrons at the end of the sulfur chains; these
can absorb wavelengths of visible radiation so as to appear brown. Plastic
sulfur is a supercooled liquid; it is amorphous in that it contains no fixed
arrangement of the sulfur atoms. The elasticity of plastic sulfur is due to the
uncoiling of the chains under tension and their recoiling when the tension
is removed. At temperatures below 95.5° C, plastic sulfur reverts slowly to
crystalline rhombic sulfur.

Just above the boiling point sulfur vapor consists of S 8 molecules. At
higher temperatures, S« molecules are produced and at 1000° C dissociation
into S 2 molecules is complete. The series of changes that sulfur undergoes
when heated may be summarized as follows;

95.5°C 119.3°C 444.6°C to 1000°C

Sa(rhombic) S 8 (monoclmic) S 8 (l) ^ S 8 (g)^±S 2 (g)

5. The Extraction of Sulfur. Until the present century most of the
world's supply of sulfur was obtained from Sicily, There the sulfur bearing
rock is stacked on a sloping floor and ignited. Part of the sulfur bums as
fuel; the remainder melts and collects at the bottom of the kiln.

Currently about 80% of the world's Sulfur is obtained from the deposits
in Louisiana and Texas, employing a process devised by Herman Frasch,
a petroleum engineer. The process is not unlike that for the extraction of
petroleum from oil wells. Since the sulfur deposit lies in a stratum about 900
feet below the surface, the Frasch process consists of making a boring
thretigh the bverlying material and driving four concentric pipes, the largest
inches in diameter, down into the deposit as shown in Figure 28.3.

The Elements of Group VIB: The Sulfur Family

363

Figure 28,3. The Frasch Process for Extracting Suliur.

Through the two outer pipes superheated water, heated under pressure to
170°C, is pumped down into the deposit, melting the sulfur which collects m
a pool at the bottom of the well. When sufficient time has elapsed for the
melting of a quantity of sulfur hot air is forced down the innermost pipe.
The air forms a froth with the molten sulfur which is forced up to the surface
through the remaining pipe. The liquid sulfur is run into enormous wooden
storage bins where it solidifies. The nature of the process, essentially a melt-
ing and a recrystallization of the sulfur which leaves the impurities un-
disturbed, yields a product about 99.5% pure, sufficiently pure for most
commercial purposes. Such a process of recrystallization is a standard
technique for the purification of chemical samples in the laboratory. Com-
mercially, sulfur can be obtained in cylindrical sticks known as “roll sulfur”
or as a fine powder obtained by the condensation of sulfur vapor and
called “flowers of sulfur.”

364

The Elements of Group V1B: The Sulfur Family

6. Properties of Sulfur. Physical properties : These are listed in Table 28-A.
Chemical properties: Sulfur is a reactive element, though less so than
oxygen. It combines directly with most of the elements, especially when
heated; with all metals except gold, platinum, and palladium to form sulfides
and with all nonmetals except nitrogen, iodine, and the noble gases. Typical
reactions are given below. 1

(1) Fe + S FeS (vigorous reactions; evolve enough heat to make
the masses incandescent; a mixture of Zn and S

(2) Zn + S — > ZnS Is sometimes used as a rocket propellant by amateur
experimenters; it should be handled with caution)

(3) S + 0 2 SO* (heated S bums in O* or air)

(4) S + H 2 -» H 2 S (a slow and incomplete reaction)

(5) 2 S + C -» CS 2 (S vapor over a bed of glowing coke)

7. Uses of Sulfur. Most of the annual output of sulfur, over five million
tons, is used in the manufacture of sulfur compounds, particularly sulfuric
acid, H2SO4. Small quantities are used for the vulcanization of rubber and
as an ingredient of insecticides and gunpowder.

HYDROGEN SULFIDE
Table 28-C

Properties of Hydrogen Sulfide

Molecular Formula :

h 2 s

Molecular Weight

34.08

Melting Point, °C :

- 85.6

Boiling Point, °C

-61.8

Critical Temperature, °C

100.4

Critical Pressure, atm

88.9

Density at STP, g/1 :

1.52

Solubility in water at
STP, volume ratio

52:1 (0.23M)

Heat of Fusion, kcal/mole :
Heat of Formation, kcal/mole:

0.568
- 4.80

Heat of Vaporization,

kcal/mole

••

4.46

* ••

Molecular Structure: The H 2 S molecule is covalent: I S— H; the H-S— H angle is 92°.

Miscellaneous: H 2 S is found in sulfur spring water and in volcanic gases. The
gas is produced whenever organic matter containing sulfur decomposes in the
absence of air At STP, H 2 S is a colorless gas with an odor of rotten eggs; it is
poisonous.

8. Preparation of H 2 S. At ordinary temperatures hydrogen and sulfur
do not combine, but as the temperature is increased reaction begins and
ahnost complete union occurs at 310°C. Unlike that of chlorine or of oxygen
with hydrogen the direct combination of hydrogen and sulfur takes place

?1 In writing equations in which sulfur appears, for the sake of simplicity we shall use
fee symbol S for- fee sulfur molecule, though we recognize its true natures to be S 8 .

The Elements of Group VI-B: The Sulfur Family

365

slowly. Hydrogen sulfide is most readily prepared by the action of a dilute
acid on a metal sulfide.

(8) S*- + 2 H+ H 2 S(g) (FeS+ 2 HC1 H 2 S(g) + FeCl 2 )

9. Properties of H 2 S. Physical properties : These are given in Table 28-C.
Chemical properties : (A) The relatively small heat of formation of H 2 S
indicates that it is not a very stable molecule.lf heated, it decomposes into
its elements, a result which is in accord with the fact that the preparation
by direct combination does not proceed readilv

(B) When ignited, H 2 S bums in air.

(7) 2 H 2 S + 3 0 2 2 SO z + 2 H 2 0 (in excess air or oxygen)

(8) 2 H 2 S + 0 2 —» 2 S + 2 H 2 0 (in a limited supply of air)

The combustion of any sulfide in excess oxygen yields S0 2 and oxides
of the other elements present in the original compound.

(9) Pb + H 2 S PbS + H 2

Copper and silver tarnish in air containing a trace of H 2 S due to the formation
of a black sulfide film.

(D) In the presence of moisture H 2 S reacts with S0 2 to liberate sulfur.
The sulfur deposits in volcanic regions are thought to be due to the reaction
of the gases, H 2 S and S0 2 , both of which are present in volcanic vapors.

(10) 2 H 2 S + S0 2 3 S + 2 H a O

(E) A solution of H 2 S is acidic and is properly called hydrosulfuric
acid. Most chemists, however, refer to the solution also as “hydrogen sulfide/’
Because H 2 S is a diprotic acid it ionizes in two stages.

(11) H 2 S H+ + HS 1 K x = 1,1 X 10- 7

(12) HS- H+ + S 2 " K 2 = 1.0 X 10~ 15

In common with all diprotic acids H 2 S gives rise to hydrogen salts and to-ao**-
mal salts, e.g., NaHS and Na 2 S, These are formed by the reaction of H 2 S a!nd
NaOH. The first reaction, while H 2 S is present in excess,, is

(13) OH- + H+ + HS- -» HS- + H a O (NaOH + H 2 S NaHS + H a O)
Upon further addition of NaOH,

(14) OH- + H+ + S 2 - S 2 - + k a O (NaOH + NaHS Na 2 S +

Aqueous solutions of both NaHS and Na 2 S are strongly basic due to hydrolysis.

(F) H 2 S is a reducing agent: The S 2- ion loses two electrons to another
species, reducing it, and is' itself oxidized to free sulfur.

(15) 2 Fe 8 + + H 2 S 2 Fe 2+ + 2 H + + S (reduction of iron(III) to iron(H;)

(16) Br a + H 2 S -> 2 HBr + S (reduction of halogen elements)

366

T he Elements of Group V1B: The Sulfur Family

(G) Sulfides of the metals show an extreme yariation in solubility.
For this reason selective precipitation of metal sulfides by H 2 S is the basis
of the classical scheme of qualitative analysis (Chapter 50).

10. Polysulfides. Sulfur dissolves in a solution of a soluble metal sulfide,
combining with the S 3 - ion to form compounds of higher sulfur content
called polysulfides. With K 2 S, the polysulfides K 2 S 2 , K 2 S 3 , K 2 S 4 , and K 2 S 5 can
be formed. The polysulfide ions may be represented as S„ 2_ , where n is
2 to 5. The sulfur atoms are linked by single covalent bonds; the disulfide ion,
S 2 2 -, is similar in structure to the peroxide ion, 0 2 3 '.

[ :s-s: f [ :s-s-s-s-*s: t

«• »» •• •• *• •• ••

SULFUR DIOXIDE

Table 28-D

Properties of Sulfur Dioxide

Molecular Formula

so 2

—

Molecular Weight

: 64.06

Melting Point, °C

: - 72.7

Boiling Point, °C

: -10.0

Critical Temperature, °C

: 157

Critical Pressure, atm

: 78

Density at STP, g/liter

: 2.86

Solubility in water at
STP, volume ratio

: 80:1 (3.5M)

Heat of Vaporization,
kcal/mole

: 6.07

Heat of Formation,
kcal/mole

: -70.9

Molecular Structure: Two equivalent electron dot structures are possible for the
S0 2 molecule; the structure of S0 2 is a resonance hybrid of the two forms. The
O— S—O angle is 120°.

•• •*

•o;

% ••

Miscellaneous : S0 2 is found in volcanic gases and in the atmosphere near indus-
trial centers where it is formed during the combustion of coal containing sulfur
compounds. At STP, S0 2 is a colorless gas with a characteristic pungent odor.
Because of the polar nature of the S0 2 molecule, liquid S0 2 is a solvent for
many substances.

11. Preparation of S0 2 . Sulfur dioxide can be prepared by

(A) The burning of sulfur or sulfur compounds

(17) S + 0 2 S0 2 ( g) AH = -70.9 kcal

(18) 4 FeS 2 + 11 0 2 -» 8 S0 2 ( g) + 2 Fe 2 0 3 (roasting a metal sulfide ore)

(B) The reaction of an acid with a sulfite or a hydrogen sulfite salt

(19) SO* 2 - + 2H+-4 S0 2 (g) + H 2 0 (Na 2 S0 3 + 2 HC1 S0 2 + H a O +

2 NaCl)

(20) HSQ a ~ + H+ S0 2 (g) + H 2 0 (NaHSQ 3 + IlCl S0 2 + H 2 Q + NaCl)

The Elements of Group VIB: The Sulfur Family

367

(21) Cu+ 2 H 2 S0 4 S0 2 (g) + CuS0 4 + 2 H a O

(22) C + 2 H 2 S0 4 2 S0 2 (g) + CO, + 2 H a O

12. Properties of SO*. Physical properties : These are listed in Table 28-D.
Chemical Properties : (A) At high temperatures and in the presence of a
suitable catalyst, S0 2 is oxidized to SO.,.

(23) 2 S0 2 •+■ O a — » 2 SO,

(B) SO* dissolves readily in and combines with water to form sulfurous
acid, H 2 S0 3 .

(24) SO, + H a O H 2 SO.,

The HgSOg molecule exists only in aqueous solution. If such a solution is
boiled Equation 24 is reversed and S0 2 gas is expelled. The structure of

the H 2 S0 3 molecule is H—O— S— O— H

•• | ••

:o:

• •

(25) H,SO, H+ + HSCV K, = 1.3 X 10~ 2

(26) HS0 3 - H+ + S0 3 2 " K 2 = 6.2 X 1 <h 8

Normal and hydrogen sulfite salts are possible. The reaction of NaOH with
H 2 S0 3 , to produce NaHS0 3 and Na 2 S0 3 , is analogous to that occurring
with H 2 S. The hydrolysis of Na 2 S0 3 results in a basic solution because of
the marked proton affinity of the SO s 2 ~ ion. A solution of NaHSO a is acidic
because the HS0 3 ~ ion has a greater tendency to yield protons than to com-
bine with them.

(D) H 2 S0 3 and its salts in acid solution are reducing agents. A variety
of oxidizing agents such as 0 2 , Cl 2 , H 2 0 2 , and KMnO* oxidize H 2 S0 3 to H 2 S0 4 ,
and S0 3 2- ion to SO* 2- ion.

(27) 2 H 2 S0 3 + 0 2 2 H 2 S0 4 (a sulfite solution in contact with air

always contains some sulfate ion)

(28) 5 SO* 2 - + 2 MnOr + 6H+->5 S0 4 2 - + 2 Mn*+ + 3 H 2 0 (oxidation

by KMnO*)

(E) Both normal and hydrogen sulfite salts decompose on heating.

(29) 4 Na 2 SO s Na 2 S + 3 Na 2 S0 4 (auto-oxidation-reduction)

(30) 2 NaHSOs Na 2 SO s + S0 2 + H 2 0

(F) In alkaline solution the sulfite ion combines with sulfur to form
the thiosulfate ion, SjOa 2 ”

The Elements of Group VI B: The Sulfur Family

(31) Na 2 S0 3 + S Na 2 S 2 0 3 .

13. Uses of S0 2 . Large quantities of sulfur dioxide are used as an inter-
mediate in the production of H 2 S0 4 , H 2 S0 3 , and various sulfites. Because
of the ease with which it can be liquified, by a pressure of three atmospheres
at 20 °C, sulfur dioxide is used as a refrigerant to some extent. Since it com-
bines directly with many organic coloring substances to form colorless com-
pounds, H 2 S0 3 is employed as a bleaching agent, especially for such materials
as wool, silk, straw, and other fabrics that might be destroyed by chlorine.
Calcium hydrogen sulfite, Ca(HS0 3 ) 2 , is used in the manufacture of paper
pulp; it dissolves lignin, a substance that holds together the fibers of cellulose
in die wood.

SULFUR TRIOXIDE
Table 28-E

Properties of Sulfur Trioxide

Molecular Formula

so 3

Molecular Weight

: 80.066

Melting Point, °C

16.8

Boiling Point, °C

: 44.8

Density at STP, g/ml

2.75

Heat of Formation,*
kcal/mole

: 22.2

Heat of Fusion, kcal/mole

: 1.92

Heat of Vaporization,
kcal/mole

: 9.48

Molecular Structure: The S0 8 molecule is a resonance hybrid of three equivalent
forms. The atoms lie in the same plane forming an equilateral triangle, with the
sulfur atom at the center and the oxygen atoms at the comers. All three O— S— O
angles are 120°.

JO!

* •

:o:

Miscellaneous: Solid SO a exists in at least two different forms. When pure SO s ,
a colorless liquid at ordinary temperatures, is frozen, a glassy transparent solid
is formed. When a trace of moisture is added to liquid S0 8 it changes to a mass
of white, needle-shaped crystals that resemble asbestos; this form sublimes without
melting.

14. Preparation of SO s . Sulfur trioxide is made by the direct union of
S0 2 and

(32) 2 SO* + 0 2 ;=± 2 SO s AH =? -44.4 kcal

The reaction is reversible. At 400° C equilibrium is reached when there is
almost complete conversion to SO a but the reaction takes place very slowly.
If the temperature is raised to increase the rate of reaction, the percent con-
version to S0 8 is reduced since the reaction is exothermic. In die presence
of a catalyst, such as finely divided platinum or vanadium pentoxide, V 2 0&,

The Elements of Group VIB : The Sulfur Family

the reaction rate is increased (without affecting the point of equilibrium)
and almost complete conversion to SO s is obtained at 400°C.

15. Properties of S0 3 . Physical properties : These are given in Table 28-E.
Chemical properties : Sulfur trioxide is the anhydride of H 2 SO*. When ex-
posed to moist air it fumes, forming minute droplets of a solution of H 2 SO*.
It also combines with 100% H 2 S0 4 to form oleum or pyrosulfuric acid, H 2 S 2 0 7 .
This, too, fumes in air and is known as fuming sulfuric acid.

(33) S0 3 + H 2 0 — » H 2 SO 4 AH = —40.5 kcal

(34) S0 8 + H 2 S 0 4 H 2 S 2 0 7

SULFURIC ACID
Table 28-F

Properties of Sulfuric Acid

Molecular Formula

h 2 so 4

Molecular Weight

: 98.08

Melting Point, °C

: 10.49

Boiling Point, °C

: decomposes

Density at 15 °C

: 1.84

Constant Boiling Mixture

Heat of Formation

: -189.8

a) at 760 mm, °C

317

kcal/mole

b) weight % of acid

98.54

Heat of Fusion, kcal/mole

0.24

Heat of Vaporization,
kcal/mole

12.0

Molecular Structure: The oxygen atoms are tetrahedrally arranged about the
central sulfur atom.

Miscellaneous: Pure, anhydrous (100%) H 2 S0 4 is a colorless, odorless, oily liquid.
If heated to boiling it decomposes into S0 3 and H 2 0, with the S0 3 coming off in
greater proportion. At 317°C an azeotropic mixture containing 98.54% H 2 S0 4
results. Commercial concentrated sulfuric acid, known as oil of vitriol, contains
about 98% acid. H z S0 4 mixes with water in all proportions, with the generation
of a great deal of heat. For this reason, when preparing a more dilute solution
of the acid, the acid is always added to water, slowly and with stirring, to distribute
the heat of solution throughout the volume of water. If the reverse is done, adding
water to the acid, the less dense water floats on the acid and the heat liberated
at the interface may produce steam and subsequent spattering of the hot acid, a
potentially dangerous event. In contact with the skin, H 2 SQ 4 produces severe bums.

16. Preparation of H 2 S0 4 . Sulfuric acid is prepared industrially by two
methods: (A) the contact process and (B) the lead chamber process . The
contact process derives its name from the fact that the oxidation of S0 2 to S0 3
takes place in contact with the surface of a catalyst. Usually sulfur is the
source of the S0 2 used in the process. Oxidation of the S0 2 is accomplished

370

The Elements of Group VIB: The Sulfur Family

by passing it, mixed with air, over a Pt or V 2 0 5 catalyst at 400-450 °C. The
S0 3 formed is then dissolved in “water.” However if the S0 3 is passed
directly into pure water, a fog consisting of droplets of H 2 S0 4 solution is
produced. This technical difficulty is overcome by absorbing the SO s in
98% H2SO4 solution, in which it is readily soluble. The final solution is
maintained at a concentration of 97-99% by a regulated influx of water. The
contact process can be represented by the following series of equations.

(35)

s + o 2 -* so 2

(burning of sulfur)

(36)

2 S0 2 + 0 2 -> 2 SO a

(catalytic oxidation of SO z )

(37)

S0 3 - 1- h 2 so* ^ h 2 s 2 o 7

(solution in sulfuric acid)

(38)

H 2 S 2 0 7 + H 2 0 -> 2 H 2 SO*

(dilution with water)

More than half of the H 2 S0 4 produced in this country is made by the contact
process; all newly constructed H 2 S0 4 plants use this process.

The chemistry of the lead chamber process differs from that of the con-
tact process primarily in the manner of oxidation of SC 2 to S0 3 . In the
chamber process this oxidation is effected in large lead-lined chambers
through the reaction of S0 2 and oxides of nitrogen, NO and N0 2 . These
'are obtained by the thermal decomposition of nitric acid, HN0 3 ; the lead
is protected by the formation of a layer of lead sulfate, PbS0 4

The actual chemical reactions for the oxidation of S0 2 in the chamber
process are not definitively understood. In the presence of air and water,
S0 2 reacts with the oxides of nitrogen to produce nitrosylsulfuric acid,
SO z (OH) (N0 2 ). With excess water the nitrosylsulfuric acid hydrolyzes to
form H 2 S0 4 .

(39) 2 S0 2 + NO + NO z + H 2 0 + O, 2 S0 2 (.0H) (NO,)

(40) 2 S0 2 (0H)(N0 2 ) +H 2 0^2 H 2 S0 4 + NO + NO a

What gives credence to this mechanism is that the compound, nitrosylsulfuric
acid, can be isolated in the laboratory as a white solid; however, in plant
operation these “chamber crystals” are not permitted to form. The oxides
of nitrogen are regenerated and used over again. In effect, they function
as a catalyst and upon their recovery depends the economic efficiency of
the process. “Chamber acid” is about 60-70% H 2 S0 4 . It can be concentrated
by evaporation, a costly procedure. It is cheaper, but more impure than
“contact acid,” but is suitable for many purposes where extreme purity is
not required.

17. Properties of H 2 SO** Physical properties: These are listed in Table 28-F.
Chemical properties : (A) Acid properties: Sulfuric acid shows the usual
properties of a diprotic acid, ionizing in two stages and forming normal sul-
fate and hydrogen sulfate salts, e.g., Na 2 S0 4 and NaHS0 4 . Sulfuric acid is
a strong acid; in dilute solution it is a somewhat weaker acid than HC1
and HN0 3 , but stronger than H 2 S0 3 .

(41) H 2 S0 4 H + + HS0 4 ~ (virtually complete reaction)

(42) HSO<f H+ + SO* 2 - K 2 = 1.2 X 10*

The Elements of Group VIS : The Sulfur Family

371

An aqueous solution of Na 2 S0 4 is essentially neutral. The degree of hydrolysis
of this salt is very small because the S0 4 2 ~ ion has but a slight tendency to
combine with a proton. However, a solution of NaHSO* is strongly acidic
due to the appreciable ionization of the HS0 4 ~ ion.

(B) Oxidizing properties: Hot concentrated sulfuric acid is an oxidizing
agent. Reactions with Cu, C, HBr, and HI have already been cited. Metals
more active than hydrogen also are oxidized by cold, dilute H 2 S0 4 in its
ordinary behavior as an acid due to the displacement of hydrogen by the
more active metal.

(C) Dehydrating action: Concentrated sulfuric acid combines vigorously
with water to form a series of hydrates, of which the best known is the
monohydrate, H 2 S0 4 * H s O. The heat of solution for the anhydrous acid,
19 kcal/mole, is an energy greater than that produced by many “chemical
reactions.” This action with water is so pronounced that H 2 S0 4 frequently
removes hydrogen and oxygen from compounds, particularly from those
which contain these elements in the two to one ratio that exists in water.
Thus paper and wood, which are largely cellulose, (C 6 H 10 0 5 )n, and sugars
such as sucrose, Ci 2 H 22 Oii, are charred by concentrated H 2 S0 4 and carbon
is set free.

(43) C^HagOn + 11 H 2 S0 4 12 C + 11 H 2 SG 4 • H 2 0

This dehydrating action of H 2 S0 4 is used to dry gases, which do not react
with the acid, by bubbling them through the acid, and to eliminate water
in many chemical reactions, such as nitration in the manufacture of dyes
and explosives.

(D) Thermal instability: When heated to about 150° C pure anhydrous
H 2 S0 4 decomposes into S0 3 and H 2 0. The escaping vapors are richer in
S0 3 than in H a O. The remaining solution begins to boil at about 290°C,
finally forming an azeotropic mixture which boils at 317° C and has a com-
position of 98.54% H 2 S0 4 .

(E) Reactions based on low volatility: Because of the high “boiling
point” and low volatility of H 2 S0 4 , it is employed in the preparation of
more volatile acids, such as HC1 and HNO a .

(44) NaCl + H 2 S0 4 -* HC1 + NaHSO,

(45) NaNOa + H 2 S0 4 HNO* + NaHSO.

Theoretically, these reactions are reversible, but when concentrated solu-
tions of H 2 S0 4 are used and the mixtures are warmed the more volatile
acids are driven off and the reactions proceed to the right

18. Uses of H 2 S0 4 * Sulfuric acid is the most widely used manufactured
compound in industry; fifteen million tons, calculated as 100% acid, are
produced annually. Major industrial uses are:

(A) the production of fertilizers, namely, ammonium sulfate and soluble
phosphate fertilizers (B) the refining of petroleum and the removal of im-
purities from crude oil, gasoline, and kerosene (C) the pickling of steel
whereby rust and scale are dissolved from a steel surface by dipping into

372

The Elements of Group VI B: The Sulfur Family

a bath of H 2 S0 4 solution prior to coating with zinc, tin, or enamel (D) as an
electrolyte in the lead storage battery and in the electrometallurgy of certain
metals (E) the manufacture of a great variety of chemicals: paints, pigments,
textiles, plastics, coal tar products such as drugs, dyes, and disinfectants,
metal sulfates, and other acids.

19. Thiosulfates. When a solution of sodium sulfite, Na 2 S0 3 , is boiled
with free sulfur, sodium thiosulfate, Na 2 S 2 O s , is formed. In this reaction the
SO s 2 - ion acts as a Lewis acid and the S atom as a Lewis base.

(46) Na 2 S0 3 + S -r> Na 2 S 2 0 3

The thiosulfate ion, S 2 0 3 2 ”, is analogous to the sulfate ion, SO* 2 ”, from the
viewpoint that an oxygen atom of the sulfate ion has been replaced by a
sulfur atom; indeed the prefix thio means “sulfur.” The sulfur atoms in the
S 2 (V~ ion; the central sulfur atom has an oxidation state of +6 whereas
the coordinated sulfur atom is ~2.

The S 2 0 3 2 " ion forms a complex ion with Ag+ ion. For this reason, sodium
thiosulfate, Na 2 S 2 0 3 , incorrectly called “hypo,” is used in “fixing” photo-
graphic film and paper by dissolving any remaining silver halide on the
film which has not been reduced by the developer.

(47) AgBr(s) + 2 S 2 0 3 2 ” Ag(S 2 0 3 ) 2 3 ” + Br~

Because it contains a sulfur atom in an oxidation state of -% the S 2 0 3 2 " ion
is readily oxidized and acts as a reducing agent. The quantitative determination
of I 2 is based on the reduction by a standard solution of Na 2 S 2 0 3 .

(48) I 2 + 2 S 2 0 3 2 ~ 2 1“ + S^ 2 - ( tetrathionate ion)

Stronger oxidizing agents such as* Cl 2 convert the S 2 0» 2 “ ion to SO* 2 ” ion.
Thiosulfuric add, H 2 S 2 0 3 , is unstable; the S a 0 3 2 " ion decomposes upon the
addition of acid.

(49) S 2 <V” + 2 H+ (H 2 S 2 0 3 ) H 2 S0 3 + S

20. Selenium, Tellurium, and Polonium. The element selenium (Greek,
selene : the moon), occurs as the free element associated with sulfur and
in the combined state as a selenide in various sulfide ores. Selenium has
several allotropic forms. Gray or “metallic” selenium has a hexagonal, rhombo-
hedral structure and is the stable variety at room temperature. Red selenium
exists in two different monoclinic forms. If the red variety is melted and
allowed to solidify, it turns into the gray form, Red selenium is soluble
in carbon disulfide, CS 2 , whereas gray selenium is not. Black and red
amorphous forms are also known.

The Elements of Group VIB: The Sulfur Family

373

"Metallic" selenium conducts electricity poorly but when exposed to
light, its conductivity increases greatly and in proportion to the intensity
of the radiation. For this reason it finds use as the selenium photoelectric
cell, a device which converts radiant energy into electrical energy, and.
which lends itself to many measuring and control operations. Because the
junction of selenium and a metal conducts electricity preferentially in one
direction, the selenium rectifier is used to convert alternating current to
direct current.

Selenium has its main use in the glass industry. Ordinary glass has a
green tint due to the presence of a light iron impurity but when a small
amount of selenium is added to the melt, the resulting glass is colorless.
When larger amounts of selenium are added to glass a ruby red glass is
obtained, which is used in automobile tail-lights and traffic signals.

The element tellurium (Latin, telluris: of the earth) occurs mainly as
a telluride in ores of lead, gold, and silver. Both selenium and tellurium
are obtained as by-products from the sludge that collects at the anode
during the electrolytic refining of metals such as copper, nickel, and lead.
The stable form of tellurium is hexagonal; it is isomorphous with gray selenium.
Few uses for tellurium have been developed. It is alloyed with lead to make
the latter more corrosion resistant, and with copper to make it more
machinable.

Polonium, discovered and named by Mme. Marie Curie for her native
country, Poland, is an extremely rare and radioactive element. It is one of
the products in the series of elements produced by the radioactive decay
of uranium and so is found in uranium bearing materials such as pitch-
blende. It is an alpha emitter like radium and its half-life is 140 days. Little
is known of its chemical, properties.

Chemical properties: The chemical properties of selenium and tellurium
are similar to those of sulfur. Representative compounds of these elements
are listed in Table 28-B. Like H 2 S0 4 , selenic acid, H 2 Se0 4 , is a strong acid
but telluric acid is weak. Telluric acid also differs in that* its formula is
H 6 TeO e rather than H 2 Te0 4 . As was the case with iodine, the large size
of the tellurium atom enables it to coordinate with six, rather than with
four, oxygen atoms. Selenous acid, H 2 Se0 3 , and tellurous acid, H 2 Te0 3 ,
have ionization constants of 2 X l(k 2 and 2 X 1Q“ 3 > respectively.

The generalization concerning the strength of the halogen acids as
oxidizing and reducing agents is valid also for the Group VIB elements.
The acids H 2 Se0 3 and H 2 Te0 3 are better oxidizing agents than is H 2 SO a
but are not so potent as H 2 Se0 4 or H 0 TeO 6 . These last two are more powerful
oxidizing agents than H 2 S0 4 . Of the group, H 2 Te, H 2 Se, and H 2 S, the best
reducing agent is H 2 Te and the weakest is H 2 S. The last three compounds
are poisonous gases.

QUESTIONS

1. Discuss the eftect of atomic size on the properties of S, Se, and Te.

2. How is sulfur obtained commercially?

3. Discuss the nature of the sulfur molecule.

374

The Elements of Group VIB .* The Sulfur Family

4. Describe the allotropie forms of sulfur. Is the conversion from rhombic to
monoclinic sulfur exothermic or endothermic? Explain.

5. Explain why the viscosity of liquid sulfur increases with temperature.

6. Why are the principal oxidation states of sulfur even-numbered?

7. Write equations for the preparation of (a) H 2 S (b) a metal sulfide (c) sodium
polysulfide (d) sodium hydrogen sulfide.

8. Write equations illustrating the reducing action of H 2 S.

9. Draw electron dot structures for the resonance hybrids of S0 2 and S0 3 ,

10. Does the thiosulfate ion have resonance hybrids? Explain.

11. Write equations for the preparation of (a) S0 2 (b).H 2 S0 3 (c) NaHS0 3 .

12. Draw electron dot structures for (a) sulfurous acid (b) sulfuric acid (c) pyro-
sulfuric acid (d) tetrathionic acid (e) thiosulfate ion.

13. Balance by the ion-electron method: (a) KMn0 4 + H 2 S0 3 (b) KMn0 4 + H 2 S.

14. Write balanced chemical equations for the preparation of sulfuric acid from
its elements and water.

15. List the relative strengths of the acids in which the Group VIB element
has an oxidation state of +6. Explain your choice.

16. Both H 2 S0 3 and H 2 S0 4 are diprotic acids. Write equations for their stepwise
ionizations. Which is a stronger base, HSO s - ion or HS0 4 “ ion?

17. Write equations illustrating chemical properties of H 2 S0 4 .

18. Given three white compounds, K^O#, K 2 S0 4 , and K 2 S 2 0 3 , how would you
proceed to distinguish one from the other?

19. What volume of H 2 S, measured at 60 °C and 740 mm pressure, can be
obtained by treating 60 g of zinc sulfide with excess acid? Ansi 17.3 liter

20. What weight of sulfur will be produced when ten liters of S0 2 and fifteen

liters of H 2 S, both measured at 27 °C and 760 mm pressure, are mixed in
the presence of moisture? Ans: 29.2 g

21. What volume of S0 2 , taken at 40 °C and 570 mm pressure, is produced
by the action of excess sulfuric acid on sodium hydrogen sulfite?

22. What volume of 0.40N HCI is necessary to react with excess Na 2 SO s to pro-
duce 50 g of NaHSO s ?

23. What is the molarity of commercial sulfuric acid which has a density of 1.84
g/ml and contains 97% H 2 S0 4 by weight?

24. A ton of pyrites, FeS 2 , 95% pure, is burned and the S0 2 produced is
converted into sulfuric acid. Calculate the weight of H 2 S0 4 produced.

25. Brass is an alloy of zinc and copper. A 0.100 g sample of brass, on being
dissolved in acid and then saturated with H 2 S, yields 75.8 mg of copper(II)
sulfide. What is the percent of copper in the alloy?

The Elements of Group VB

Nitrogen

In Group VB of the Periodic System there are five elements, nitrogen,
phosphorus, arsenic, antimony, and bismuth, known collectively as the Nitrogen
Family. The electron configurations of these elements are shown below; other
properties of these elements are listed in Table 29-A.

Element

Atomic

Number

2$ 2 p

3s 3p 3d

4s 4p 4d 4f

5$ 5 p 5 d

6 s 6p

Nitrogen

2 3

Phosphorus

2 6

2 3

Arsenic

2 6

2 6 10

2 3

Antimony

2 6

2 6 10

2 6 10 14

2 3

Bismuth

2 6

2 6 10

2 6 10 14

2 6 10

2 3

1. General Properties of the Nitrogen Family. The Group VB elements
have five valence electrons, ns 1 2 np 3 ; the three p electrons are not paired
but are distributed, one in each of the three p orbitals. The relationships which
have been observed previously concerning the gradation of properties within
a group are again evident with the elements of Group VB. Thus as the
atomic size increases there is a transition from the nonmetallic nitrogen
through the quasi-metallic arsenic and antimony to the metallic bismuth.
Phosphorus is, perhaps, the most typical element of the group; by virtue of
its small size, . nitrogen shows some anomalous characteristics.

None of the Group VB elements are especially reactive, being much less
so than the adjacent elements of the sulfur family and the halogens. Of the
three successive elements, phosphorus, sulfur, and chlorine, the phosphorus

atom has the smallest nuclear charge and is the largest in size. The phosphorus
atom is three electrons short of. a noble gas electron configuration, the sulfur
atom two electrons, whereas the chlorine atom needs but one electron to
complete an octet. All these factors work to reduce the tendency of the

376

The Elements of Group VB; Nitrogen

Table 29-A

Properties of the Group VB Elements

Property

Nitrogen

Phosphorus

Arsenic

Antimony

Bismuth

Symbol

Molecular Formula

As*

Sb 4

Atomic Number

•83

Atomic Weight

14.0067

30.9738

74.9216

121.75

208.980

Isotopes (mass numbers

14 (99.63)

31 (100)

75 (100)

121(57.25)

209(100)

and percents)

15 ( 0.37)

123(42.75)

Abundance in Earth’s Crust, %

0.03

0.11

10-«

10-"

Physical State at STP

colorless

solid

gray-white

gas

solid

Allotropes

white;

gray;

red

yellow

Melting Point, °C

-210.1

44.1(w)

814

630(w)

271

597 (r)

(at 36 atm)

Boiling Point, °C

“195.8

280 (w)

431 (r)(s)

610(s)

1440

1560

Critical Temperature, °C

-147.1

Critical Pressure, atm

33.5

Density at STP, g/ml

0.00125

1.82(w)

5.72(g)

6.58(g)

9,80

2.34(r) 1

3.9 (y)

5 3 (y)

Heat of Fusion, kcal/mole

0.086

0.15

6.62

4.74

2.6

Heat of Vaporization,

0.67

2.97(w)

34.5(s)

16.2

36.2

kcal/mole

7.2 (r)(s)

Heat of Atomization at 25 °C,

112 6

79.8

69.4

62.7

47.5

kcal/g-atom

Ionization Potential, 1st, eV

14.45

10.55

9.81

8.64

8.30

Electronegativity

3.0

2.1

2.0 1

1.8 '1

1.7

Covalent Radius, A

0.74

1.06

1.21

1.41 i

1.52

Ionic Radius, A (+5)

0.11

0.34

0.47

0.62 1

0.74

Oxidation States

all states

“2, -3

-3

-3(u)

“3 toH-5

+ 1. +3, +4, +5

+3, +5

4*3, 45

43, 45

Key to symbols: (w) = white; (r) = red; (g) = gray; (y) = yellow; (s) = sublimes;
(u) = unstable

Group VB elements, relative to comparable elements of Groups VIB and
VIIB, to attract electrons into the valence shell.

Of the Group VB elements, only the two smallest atoms, nitrogen and
phosphorous, are capable of forming stable negative ions of oxidation state
-3, and nitrogen reacts more readily in this respect than does phosphorus.
Bismuth forms no stable compounds in which it shows a negative oxidation
state. Except for the N 8 ” and P 3 ~ ions, atoms of the Group VB elements
attain a noble gas electron configuration only by forming covalent bonds.
Ionic compounds, such as NaN0 3 , are common but in such cases the
Group VB element is part of an ion within which it is covalently bonded.
The principal oxidation states are -3, +3, and +5, values which are one
less than for the Sulfur Family. Thus all the elements form hydrogen com-
pounds of the type MH 3 and oxides of the types M 2 0 3 and M 2 O s . Nitrogen
is uncommon in that it exhibits all oxidation states from -3 to +5 (Table 29-B)

The Elements of Group V/B.* Nitrogen

377

The acid character of the oxides decreases as we go down the group
from nitrogen to bismuth, and the metallic nature of the element increases.
Thus N 2 O 3 and P 2 O 3 are acidic, As 2 0 3 and Sb 2 0 3 are amphoteric, whereas
Bi 2 O s is basic. All the oxides of type M 2 0 5 are acidic, but Bi 2 0 5 is least so.
Except for nitrogen, all the Group VB elements are solid at ordinary tempera-
tures. Only nitrogen and bismuth do not show allotropic modifications.

NITROGEN

In view of the great variety of nitrogen compounds, and of both their
industrial and biological importance, we shall consider nitrogen and its
compounds separately. Nitrogen is not too abundant even though, as elemen-
tal N 2 , it constitutes 78% of the atmosphere by volume. In the combined
state nitrogen is found principally as nitrates, KN0 3 and NaN0 3 . The
latter, Chilean saltpeter, was the chief source of nitrogen compounds until
chemical methods were devised, in the early part of the twentieth century,
for converting the N 2 of the air into nitrogen compounds. Such procedures
are known as the “fixing of nitrogen / 7 As a constituent of proteins, nitrogen
is an essential part of all living matter.

Table 29-B

Oxidation States of the Nitrogen Family

Oxidation

State

Nitrogen

Phosphorus

Arsenic

Antimony

Bismuth

-3

N 2 3 “ ion

SbH,

BiH a (u)

NH,; NH*+ ion

ESSE2BE51

■■HI

-2

NJL

-1

l'if 1 fw 1 WBBBI

■HHHHHI

HHIHl

HHHHHI

■

p(pj

■ ■

BamW

mSSmKtm

■ ■

■

4-3

n 2 o 8

P 4 0 6

As 4 0 6

Sb 4 O s

Bi 4 O s

hno 2

h 3 po,

H3ASO3

H 3 Sb0 3

N0 2 - ion

As 0 3 3 - ion

Sb 8 + ion

Bi 8 4 ion

SbO+ ion

BiO+ ion

NCls

PC1 3

AsC1 3

SbCl 3

BiCl 3

p 2 o 4

EISsnIH

Bi 2 0 4

n 2 o 5

As 4 O 10

Sb 4 Oio

Bi 2 U 5

hno 3

h 3 po 4

H 3 As0 4

HSb(OH),

N 0 3 - ion

PO 4 3 - ion

♦

PCI,

AsF s

SbCl 5

*No pentahalides of N and Bi are known; AsC 1 3 is unknown;
(u) = unstable

2, Preparation of Nitrogen. Nitrogen can be obtained from the air simply

by removing the oxygen. This leaves the other components of the atmosphere,

e.g., the noble gases, as a small impurity. Industrially nitrogen is obtained

378

The Elements of Group VB; Nitrogen

by the fractional distillation of liquid air. Chemically pure nitrogen can be
prepared by the decomposition of nitrogen compounds.

(1) NH 4 + + N0 2 " N 2 + H 2 0 (warm saturated solutions of NH+ ion and

N<V ion are mixed; ammonium nitrite is
unstable)

(2) 2 NH 3 + 3 CuO N 2 + 3 Cu + 3 H 2 0 (NH 3 gas over red hot CuO)

3. Properties of Nitrogen. Physical properties : These are given in Table
29-A. At STP nitrogen is a colorless, odorless, and tasteless gas having a
solubility in water about half that of oxygen, 23 ml/liter.

Chemical properties: At ordinary temperatures nitrogen is relatively un-
reactive. Indeed it serves to dilute the more reactive oxygen of the atmosphere.
Lavoisier named the element azote , or lifeless, a term by which it is still known
in France. The nitrogen molecule is diatomic and nonpolar, with a triple
covalent bond between the nitrogen atoms.

:N = N:

The N = N bond is extremely stable. Its heat of dissociation, 225.2 kcal/mole,
is larger than that for any diatomic molecule. It is this high dissociation
energy which accounts for the unreactivity of the N 2 molecule and con-
versely for the tendency of many nitrogenous compounds to decompose into
the more stable N 2 . At high temperatures, nitrogen combines directly with
hydrogen and oxygen, also with more active metals to form ionic nitrides,
such as Li 3 N, Ca s N 2 , and AIN. Less active metals and some nonmetals form
covalent nitrides such as P 3 N 5 , BN, Si 3 N 4 , and NI S , but these are best pre-
pared indirectly by the reaction of ammonia and the element concerned.

4. Uses of Nitrogen. Large quantities of nitrogen are used in various
processes for “fixing atmospheric nitrogen,” and the production thereby of
nitrogen compounds, notably NH 3 . The major use of nitrogen compounds
is as a fertilizer to supply plant life with its necessary nitrogen. Some uses
of N 2 are based on its inactivity; where an inert atmosphere is required for a
chemical process or for storage of an otherwise reactive substance, N 2 may
be used.

AMMONIA
TahU 29-C

Properties of Ammonia

Molecular Formula

NH S

Molecular Weight

17.02

Melting Point, °C

-77.7 -

Boiling Point, °C

-33.4

Critical Temperature, °C

132.4

Critical Pressure, atm

111.5

Density at STP, g/1

0.771

Density of Liquid, g/ml

0.677

Heat of Fusion, kcal/mole

1.84

Heat of Vaporization,
kcal/mole

5.88

Heat of Formation,
kcal/mole

; 11.0

Solubility in water at STP . :
volume ratio

1160:1

The Elements of Group VB: Nitrogen

379

Molecular Structure: The NH* molecule is covalent. It has a tetrahedral shape
with the N atom at the apex and the H atoms forming an equilateral base.
.The N-H bond distances are 1.016 A and the height of the pyramid is 0.360 A.
The angle between the N-H bonds is the tetrahedral angle of 109°28'. The nitrogen
atom vibrates between two equally stable positions, one above and one below
the plane of the hydrogen atoms; this “inversion” frequency of vibration is
2.387013 x 10 10 per second and is the basis of the “ammonia atomic clock.”

Side View Top View

Miscellaneous: NH 3 is produced in nature by the action of bacteria upon organic
matter in the soil. At STP, NH 3 is a colorless gas with a distinctive pungent
odor. Diluted in a large quantity of air it is not poisonous but is irritating to the
mucous membranes.

5. Preparation of NH 3 . Ammonia was known in alchemical times. Lavoisier
prepared it first by heating ammonium chloride, NH 4 C1 ("sal ammoniac,”
a name derived perhaps from the Egyptian god, Ra Ammon), and calcium
hydroxide, Ca(OH) 2 . Ammonia can be prepared by the following reactions:

(A) Heating an ammonium salt, NH^, with a strong base, OH~:

(1) NH 4 + 4- OH" — » NH 3 + H 2 0 (NH 4 Cl + NaOH NH a + H 2 0 +NaCl)

(B) Hydrolysis of nitrides of active metals:

(2) Mg 3 N 2 + 6 H 2 0 2 NH 3 + 3 Mg(OH) 2

(3) N/+ 3H 2 2 NH 3 AH = -22.0 kcal

Operating at a pressure of 200 atmospheres and at 500 °C with a catalyst
of iron mixed with potassium aluminate, a 15% conversion to NH 3 is ob-
tained. After passing over the catalyst the gases are cooled and the NH 3 is
recovered by liquifaction under pressure. Unreacted N 2 and H 2 are then
recycled through the catalyst chamber. The similar French Claude process
achieves a yield of 40% NH 3 by operating af a higher pressure of 1000
atmospheres.

(D) As a by-product in the production of coke: Soft coal contains about
1% combined nitrogen. When such coal is heated in the absence of air
(destructive distillation) about one-fifth of such nitrogen is evolved as NH S .

(E) By the Cyanamide Process: Lime, CaO, and coke, C, are heated in
an electric furnace to form calcium carbide, CaC 2 . Nitrogen, obtained from

The Elements of Group VB; Nitrogen

liquid air, is passed over CaC 2 at 1000°C, producing calcium cyanamide,
CaCN 2 , which is then treated with high pressure steam to yield NH 3 .

(4) CaO + 3 C CaC 2 + CO

(5) CaC 2 + N 2 CaCN 2 + C

(6) CaCN 2 + 3 H 2 0 -h> 2 NH S + CaC0 3

The Cyanamide Process employs inexpensive raw materials but requires
large amounts of energy so that it has been displaced by the more economical
Haber Process.

6. Properties of NH 3 . Physical properties : These are listed in Table 29-C.
Chemical properties : (A) The nitrogen atom of the NH S molecule has an
unshared pair of electrons. This electron pair can form a coordinate covalent
bond tc an atom or ion whose valence shell is incomplete. Thus NH 3 can
combine with a proton:

(7) NH 3 + H+ NH 4 + (NH s + H 2 0 NH 4 + + OH"; K = J.8 X 1(H)

Thus NH 3 is. a base and an aqueous solution gives a basic reaction. The
aqueous solution is also called ammonium hydroxide because ‘ammonia water”
was once thought to contain the definite compound NH 4 OH, which then
ionized slightly to yield NH* + and OH- ions: NH a + H a O NH 4 OH
NH 4 + 4- OH". It is likely that between the NH 3 and H a O molecules there
is a hydrogen bond, and the oscillation of the proton between the nitrogen
and oxygen atoms makes the distinction between (NH 3 + H 2 0) and NH 4 OIi
arbitrary. Perhaps it is most correct to say that a solution of ammonium
hydroxide consists primarily of hydrated ammonia molecules which ionize
slightly into NH*+ and OH~ ions. The formation of a coordinate covalent
bond is also the mechanism of the reaction between NH 3 and BF 3 and of
the combination of NH 3 with many simple metallic ions, such as Cu 2+ to
form complex ions.

(B) The NH 3 molecule is analogous to the H s O molecule. If H 2 0 can
be represented as H(OH) then NH 3 can be represented as H(NH 2 ). The
NH 2 - radical is known as the amide group and can be substituted for other
negative groups. Such a substitution is termed ammonolysis because of its
similarity to the hydrolysis reaction of water.

(8) HgCl 2 + 2 NH 3 -» HgNH 2 Cl + NH 4 + + Ch

(9) Na + NH S NaNH 2 + H 2 (analogous to the reaction with H 2 0)

( 10) 3 Mg + 2 NH 3 -* Mg s N 2 + 3 H 2 (at high temperatures to form nitrides)

(D) The oxides of less active metals are reduced by NH S .

(11) 3 CuO + 2 NH a -*■ N 2 + 3 Cu + 3 H 2 0 (at elevated temperatures)

(E) Ammonia burns in oxygen and in air.

(12) 4 NH a + 3 0 2 2 N 2 + 8 H s O (in pure oxygen)

The Elements of Group VB: Nitrogen

381

(13) 4 NH 3 + 5 0 2 4 NO + 6 H 2 0 (in oxygen-enriched air over a

Pt gauze catalyst at 700° C)

The last reaction is the Ostwald Process, or the ammonia-oxidation process,
and is one of the steps in the industrial production of nitric acid, HN0 3 .

7. Liquid Ammonia. Mention has been made of the chemical similarity
between the NH 3 and H 2 0 molecules. The extreme solubility of NH 3 in
water is due to the highly polar nature of both the NH 3 and the H a O
molecules, and their combination to form hydrated ammonia molecules.

Liquid NH 3 is an excellent solvent. Though not quite so good a solvent
for electrolytes as H a O because of its lower dielectric constant (22 vs 81 for
H 2 0), it is a better solvent for many covalent compounds. Like H 2 0,
liquid NH 3 is a poor conductor of electricity and is only slightly ionized.

(14) 2 NH 3 NH^ + NH 2 -

By comparison with the ionization of water, theNH 4 + and NK 2 " ions in a
liquid ammonia system, are analogous to the H 3 0+ and OH" in a liquid
water system. Thus solutions of NH 4 + salts in liquid NH 3 exhibit acid
characteristics; they affect indicators and react with active metals to liberate
hydrogen.

(15) Zn + 2 NH 4 + -» Zn 2 + 4- 2 NH 3 4-H 2

Unlike H 2 0, liquid NH 3 dissolves alkali metals and alkaline earth metals
without their reacting to liberate hydrogen. Such dilute solutions are bright
blue and are excellent conductors of electricity; these solutions are fairly
stable, decomposing slowly to form amides such as sodium amide, NaNH 2 .

8. Uses of NH> Almost all industrial ammonia is used in the production
of other nitrogen compounds, especially HN0 3 and NH 4 + salts, for which,
indeed, it is the keystone chemical. Because it is readily liquified and be-
cause of its high heat of vaporization, exceeded only by that of water, NH 3
is used as a refrigerating liquid. At 20 °C a pressure of only 8.4 atmospheres
suffices to liquify NH ;1 vapor. Since the heat of fusion of water is 79 cal/g,
the vaporization of one gram of NH 3 theoretically can abstract sufficient
heat to freeze four grams of water at 0°C. A dilute solution of NH 3 , ammonia
water, finds use as a household cleanser because of its slight basicity.

9. Other Compounds of Nitrogen and Hydrogen. Hydrazine, N 2 H 4 , is
formed by the oxidation of NH 3 with NaClO. Its relation to NH a is
analogous to that of H 2 0 2 to H 2 0. Hydrazine is a colorless toxic liquid with a
density of 1.01 g/ml; its melting point is 1°C and its boiling point is 113°C. It
is soluble in polar solvents and is. less basic than NH 3 .

(16) N 2 H 4 4- H s O N 2 H 5 + 4- OH- K= 8.5 X 10" 7

Hydrazine is far more reactive than NH 3 . With liquid 0 2 , F 2 , or H 2 0 2 as
the oxidizer, hydrazine is a high-performance fuel as a rocket propellant.
The reactions are highly exothermic; with H 2 0 2 it is self-igniting.

(17) N 2 H 4 + 2H 2 0->N 2 + 4 H 2 0

382

The Elements of Group VB: Nitrogen

A derivative, dimethyl hydrazine, H 2 N~N(CH 3 ) 2 , is used as the fuel in
military missiles. The — N=N— bond is the basis of the important group
of "azo” cpmpounds in organic chemistry.

When N 2 H 4 is reduced with nitrous acid, HN0 2 , a colorless liquid,
hy dr azoic acid, HN 3 , is formed. Hydr azoic acid is a weak acid; its salts are
called azides and contain the azide ion, N 3 ". Both the acid and its salts are
unstable; lead azide, Pb(N 3 ) 2 > is a detonator for military explosives. Hydro-
xylamine, NH 2 OH, is prepared by the reduction of nitric oxide, NO. It is a
white solid whose aqueous solution is weakly basic; it is unstable and a
powerful reducing agent.

The electron dot structures for these three compounds are shown below.

H H

h:n':n: h

h:n ::n::n:

h:n:o:h

Hydrazine

Hydrazoic Acid

Hydroxylamine

OXIDES AND OXYGEN ACIDS OF NITROGEN

10. Oxidation State +1. Nitrous oxide, N 2 0, and hyponitrous add,
H 2 N 2 0 2 , are compounds in which the nitrogen atoms exhibit an oxidation
state of +1. Nitrous oxide is prepared by heating ammonium nitrate, NH 4 N0 3 .
Care must be taken not to heat too strongly or an explosion may result.

(18) NH 4 N0 3 N 2 0 + 2 H 2 0

At STP, N 2 0 is a colorless gas with a faint odor and a sweetish taste. It melts
at -102.4°C and bods at -89.5°C. The N 2 0 molecule is linear and has two
resonance forms.

: N = N 8 : :N = N-o:

• •

When inhaled in small doses, N a O induces a condition of exhilaration often
accompanied by boisterous laughter; on this account the gas is often called
‘laughing gas.” Larger doses of N 2 0 mixed with oxygen are used as a dental
anesthetic. The solubility of N 2 0 in fats is the basis of "self-whipped” cream.
The N 2 0 is dissolved under pressure in a cartridge containing cream; upon
release of the pressure, tiny bubbles. of the gas produce a foam or whipped
cream. When heated above 550°C, N a O decomposes into its elements* For
this reason a glowing splint bursts into flame when immersed in the gas,
much as it does in pure oxygen. Although the formula of N 2 0 is the
anhydride of hyponitrous acid, H 2 N 2 0 2 , the gas does not combine with
water though it is soluble to the extent of 60 volumes per 100 volumes of
water. Hyponitrous acid is a white unstable solid; it decomposes upon
standing into N a O and H 2 0. Its structural formula is

H-o: :o-H

X N = N /

The Elements of Group V B: Nitrogen

383

11. Oxidation State +2. In nitric oxide, NO, the nitrogen atom exhibits
an oxidation state of +2. The NO molecule is one of the few stable molecules
with an odd number of electrons, and two resonance forms can be written for it.

sN = o: :n = o:

Though the two forms appear to indicate that an unpaired electron divides
its time equally between the nitrogen and oxygen atoms, it is more likely
that this electron is distributed over the entire molecule. Thus a tendency
to dimerize and to form «(NO) 2 is decreased and N 2 0 2 is found only to a
small extent in the liquid state. Because of the unpaired electron, the NO
molecule is paramagnetic.

Nitric oxide can be prepared by:

(A) the reaction of less active metals, such as Cu and Ag, with dilute HNO»

(19) 3 Cu + 8 H+-+ 2 NOr -» 3 Cu*+ + 2 NO + 4 H z O

The NO produced by this reaction is not pure inasmuch as other oxides of
nitrogen are simultaneously produced, though NO is the chief product.

(B) the Ostwald Process or the catalytic oxidation of NH 3

(20) 4 NH 3 + 5 0 2 -> 4 NO + 6 H a O

(21) n 2 + 0 2 -» 2 NO AH = +43.5 kcal

This highly endothermic reaction was the basis of the now obsolete Arc
Process, whereby a mixture of N 2 and 0 2 , passed through an electric arc,
gave a low yield, perhaps 5%, of NO. Nitric oxide is formed during lightning
discharges; it has been estimated that such natural phenomena fix 40 mil-
lion tons of nitrogen annually.

Nitric oxide is a colorless gas at STP; its melting point is -163.6°C and
its boiling point is -151,8°C. It combines with oxygen spontaneously; the
reaction is readily apparent since the brown gas, nitrogen dioxide, N0 2 , is
the product.

(22) 2 NO + 0 2 2 NO, AH = -27.8 kcal

With metal ions such as Fe 2 +, Cu 2 +, and Co 2 +, NO forms complex ions.

(23) Fe 2 * + NO Fe(NO) 2 + (a brown complex ion used as the brown

ring test for nitrates and nitrites)

12. Oxidation State +3. Dinitrogen trioxide, N 2 O a , nitrous acid, HN0 2 ,
and nitrites, NOjf, are compounds in which the nitrogen atom has an oxida-
tion state of +3. When an equimolecular mixture of NO and NO z is cooled,
condensation to a blue liquid takes place at -21 °C. This liquid is dinitrogen
trioxide, N 2 0 3 . It is the true anhydride of nitrous, acid; upon warming it
decomposes into NO and N0 2 . The structure of N 2 O s is

•° X

• ^ / \

N- N

;o.

•• •*

384

The Elements of Group VB: Nitrogen

When an acid is added to a cold dilute solution of a nitrite salt, N0 2 ~,
a pale blue solution containing nitrous acid, HN0 2 , is obtained.

(24) NCV + H+ HN0 2 (KN0 2 + H 2 S0 4 -> HN0 2 + KHSO,)

Nitrous acid is a weak acid. It is unstable and is known only in solution;
upon warming it decomposes.

(25) 3 HN0 2 HNQ 3 4 - 2 NO + H 2 0

The salts of HN0 2 , the nitrites, are stable compounds, of which the most
important is NaN0 2 . It is prepared by heating sodium nitrate, NaN0 3 , to a
high temperature, or by the reduction of NaN0 3 with Pb, Fe, or C.

(26) 2 NaNOa -» 2 NaN0 2 + O s

(27) NaN0 3 + Pb -> NaNO, + PbO

The structures of HN0 2 and two resonance forms of the N0 2 ~ ion are
shown below.

• • ••

N-O-H

II "

:o:

Nitrous acid Nitrite ion

1 N-O:

II ••

*• m »

N=0

• •

:0:

• •

Since the nitrogen atom has an intermediate oxidation state of -3, both HN0 2
and the N0 2 ~ ion can act either as an oxidizing agent or as a reducing agent.

13. Oxidation State +4. Nitrogen dioxide, N0 2 , and its dimer, dinitrogen
tetroxide, N 2 0 4 , exist in equilibrium with each other.

(28) 2 NO* N 2 Q 4

Both substances are gases; N0 2 is brown and N 2 0 4 is colorless. Due to the
presence of some N0 2 the mixture ordinarily appears brown. Because the
formation of N 2 0 4 from N0 2 is an exothermic reaction, the color of the
mixture becomes lighter when cooled due to the shifting of the equilibruim
to the right. When heated the brown color is intensified as the percent of
N0 2 increases.

The N0 2 molecule contains an unpaired electron and three resonance
forms are possible wherein this lone electron shares its time equally
among the three atoms of the N0 2 molecule. Because of this unpaired
electron, N0 2 is colored and paramagnetic. In the formation of an N 2 0 4
molecule, two molecules of NO a share their unpaired electrons; N 2 0 4 is
thereby colorless and not paramagnetic. The molecular structures of N0 2
and N 2 0 4 are

• N

* *

/°- :

:n:

.p:

Nitrogen dioxide

Dinitrogen tetroxide

The Elements of Group VB: Nitrogen

385

Nitrogen dioxide can be prepared by (A) the oxidation of NO by 0 2 ;
(B) the reaction of less active metals with concentrated nitric acid; (C) the
thermal decomposition of nitrates of heavy metals.

(29) 2 NO + 0 2 2 N0 2

(30) Cu + 4 H+ + 2 N0 3 - -* CirM- + 2 N0 2 + 2 H 2 0

(31) 2 Pb(NO s ) 2 2 PbO 4- 4 NO, + 0 2
Nitrogen dioxide reacts with water to form nitric acid.

(32) 3 NO, + H 2 0 2 HN0 3 + NO

14. Oxidation State +5. Compounds containing nitrogen in the +5
oxidation state are dinitrogen pentoxide, N 2 0 5 , nitric acid, HNO a> and nitrates,
N0 3 ". Dinitrogen pentoxide is prepared by the action of a powerful dehy-
drating agent such as phosphorus ( V ) oxide, P*O 10 , upon HN0 3 .

(33) 4 HN0 3 + P 4 O 10 2 N 2 O s + 4 HPO,

Dinitrogen pentoxide is a white solid and is unstable above 30° C, decom-
posing into N0 2 and 0 2 . It combines with water to form HN0 3 . The
structure of N 2 O s is

NITRIC ACID
Table 29-D

Properties of Nitric Acid

Molecular Formula

: HNO a

Molecular Weight

: 63.016

Melting Point,

: -41.8

Boiling Point, °C

: 86

Density at 20 °C, g/ml

1.50

Constant Boiling Mixture

a) at 760 mm, °C

b) weight % of acid

: 121
: 68.4

Heat of Fusion, kcal/mole

Heat of Formation,
kcal/mole

: 0.60

: -41.4

Heat of Vaporization,
kcal/mole

: 7.2

Molecular Structure; Two

resonance forms are possible.

•

H-O-N^ ’ H-O-N^*

.. ••

* i •_

• • *

Miscellaneous: Pure anhydrous HNO s is a colorless liquid which fumes in moist
air. It is miscible in all proportions with H 2 0. Commercial HNO a is yellow due to

386

The Elements of Group VB; Nitrogen

some decomposition into NO a which remains dissolved. Large amounts of N0 2
can dissolve in HNQ 3 to give a red-brown solution called fuming nitric acid, which
is more reactive than ordinary concentrated HNO ;i . Dilute HN0 3 stains the skin
yellow due to its reacti on with proteins; concentrated HNQ 3 causes painful bum s.

15. Preparation of HNO s . Nitric acid can be prepared by:

(A) The reaction of a nitrate salt with the less volatile sulfuric acid

(34)

NaN0 3 + H 2 S0 4 -*■ HNOs +

khso 4

(B)

A series of steps involving the fixation

of atinospheric nitrogen

(35)

N 2 + 3H,->2 NH 3

(Haber Process)

(36)

4 NH 3 + 5 0 2 -» 4 NO 4- 6 H.O

(Ostwald Process)

(37)

2 NO + 0 2 2 N0 2

(oxidation of NO)

(38)

3 N0 2 + H 2 0 -* 2 HNO s + NO

(solution of NO z )

An alternate route is through the preparation of NO by the Arc Process.
The NO produced in the last step is recycled to the reaction chamber for
further oxidation to N0 2 .

16. Properties of HN0 3 . Physical properties : These are listed in Table 29-D.
Chemical properties : Nitric acid is a strong acid, a powerful oxidizing
agent, and is unstable at elevated temperatures.

(A) Instability: Upon heating or exposure to light, nitric acid decomposes.

(39) 4 HNO s -> 4 NO, + 0 2 + 2 H z O

(B) Oxidizing properties: Nitric acid, or the nitrate ion in acid solution,
is an oxidizing agent which undergoes a variety of reactions depending upon
the concentration of the HN0 3 and the nature of the reducing agent and
its concentration. We have seen that the reactions of dilute and concentrated
HN0 3 with Cu differ in that NO is produced with dilute acid and N0 2 with
concentrated acid. Below are listed the possible half-reactions of HN0 3 in
producing the several oxidation states of nitrogen, together with their oxida-

tion potentials.

Half-reaction E° (volt)

NO, + H a O N0 3 - + 2 H+ + e -0.80

HN0 2 + H 2 0 NOr + 3 H+ + 2 e -0.94

NO +2 H 2 0 -» NOr + 4HH 3 e -0.96

N 2 0 + 5 H 2 0 -» 2 N0 3 - + 10 H+ + S e -1.12

N 2 4-6 H z O -* 2 N0 3 - + 12 H+ + 10 e -1.25

(NH 2 OH)H+ + 2 H 2 0 N0 3 ~ + 8 H+ + 6 e -0.73

N 2 H 5 + 4 - 6 H 2 0 -* 2 NO/ 4- 17 H+ 4 - 14 e -0.83

NH4 4 - 3 H 2 0 -» N0 3 - 4 - 10 H+ 4 - 8 e -0.88

Inasmuch as the potentials of these half-reactions are negative, N0 3 " ion in
acid solution is a better oxidizing agent than H+ ion alone. The oxidizing
property of a “nonoxidizing” acid such as HC1 is due to the reaction,

The Elements of Group VB: Nitrogen

387

2 H+ 4* 2 e — » H 2 , for which E° is 0.00 volt. Hence HC1 cannot oxidize
metals such as copper and silver which are below hydrogen in the electro-
motive series. However, HN0 3 can oxidize all but the most inactive metals,
notably gold and platinum. Water is always produced by the oxidizing re-
action of HNO s ; free H 2 is never formed even with the active metals.
Representative reactions of HNO s are given below.

1. Reactions with less active metals such as Pb, Ag, and Cu

(40) Pb + 4 H+ 2 N(V Pb 2 + + 2 N0 2 + 2 H 2 0 (with concentrated

HNO.)

(41) 3 Ag + 4 H+ + NCV Ag + + NO + 2 H 2 0 (with dilute HNO s )

2. Reactions with more active metals such as Zn and Fe

(42) 4 Zn+ 10 H+ + 2 NO s ~ 4 Zn 2 + +N.O+5 H 2 0 (with dilute HNO*)

(43) 5 Zn + 12 H+ + 2 NO s - -> 5 Zn 2 + + N 2 + 6 H 2 0 (with very dilute

hno 3 )

(44) 4Zn + 10 H+ + N(V-*4Zn 2 + + NH 4 + + 3 H a O (with very dilute

hno 3 )

3. Reactions with nonmetals such as P, I 2 and S

(45) 3P + 5H+ + 5N0 3 - + 2H 2 0^3H S P0 4 + 5.N0 (with concen-

trated hno 3 )

(40) 3 I 2 + 4 H+ + 10 N0 3 - 6 IO s - + 10 NO+ 2 H 2 0

(47) S + 2H+ + 2 NCV -* 2H+ + SO* 2 - + 2 NO

4. Reactions with inorganic compounds such as H 2 S and S0 2

(48) 3 H 2 S + 2 H+ + 2 NO s ~ ->3S + 2NO + 4 H.O

(49) 3 S0 2 + 2 H+ + 2 NO a - 3 S0 4 2 - + 2 NO

5. Reactions with organic compounds such as toluene, C 6 H 5 GHi, and gly-
cerol, C 3 H 5 (OH) 3

(50) C 6 H 5 CH 3 + 3 HN0 3 ^3H 2 0 + C 6 H 2 (N0 2 ) 3 CH 3 (trinitrotoluene)

(51) C 3 H 5 (OH) 3 + 3 HN0 3 3 H 2 0 + C 3 H 5 (N0 3 ) 8 (nitroglycerine)
Both trinitrotoluene and nitroglycerine are violently explosive; nitroglycerine
is an oily liquid which may be absorbed in wood pulp to form the less
sensitive dynamite. With cellulose, HNO a reacts similarly to form nitro-
cellulose; wool or turpentine, moistened with fuming HNO a , bursts into
flame.

6. Aqua regia: Neither HC1 nor HN0 3 alone can dissolve elemental gold
or platinum. A mixture of one part of HNO s with three parts of HC1, known
as aqua regia (royal water), oxidizes these metals due to the formation of
the complex ions, AuC 1 4 " and PtCl 6 2 '

(52) Au + 4H+ + NCV + 4 Or -> AuClr + NO + 2 H z O

(53) 3 Pt + 16 H+ + 4 N0 3 - + 18 Cl- -► 3 PtCl 6 2 - + 4 NO + 8 H s O

388

The Elements of Group VB: Nitrogen

(C) Nitrates: Salts of nitric acid, prepared by the reaction of HN0 3
with metals, metal oxides or hydroxides, are known as nitrates. The nitrate
ion, NCV, is planar and has three resonance structures.

*•

:o:

All nitrate salts decompose upon heating; the thermal decomposition of
NH4NO3 and of NaNOs has already been mentioned.

17. The Nitrogen Cycle. All living organisms require nitrogen in some
combined form but neither plants nor animals can assimilate the elemental
nitrogen of the atmosphere. Nitrogen compounds are taken in by plants
and converted by biochemical processes to proteins. Animals obtain their
necessary nitrogen by feeding upon plants or other animals.

The removal of nitrogen from the soil is balanced by certain processes
which replenish the supply of nitrogen compounds in the soil. These processes
involve:

(A) the action of nitrogen-fixing bacteria which live in nodules on the roots
of leguminous plants such as peas, beans, clover, and alfalfa

(B) the decay of plant and animal matter

(C) the formation of NO from atmospheric N 2 by lightning discharges, and its
subsequent oxidation and reaction with water

(D) the deliberate addition of nitrogenous fertilizers to the soil.

This circulation of nitrogen in nature constitutes an equilibrium between
the various users and nroducers of nitrogen compounds and is known as
the Nitrogen Cycle.

QUESTIONS

1. Predict the properties of the elements of Group VB on the basis of their
electron configurations and atomic dimensions.

2. What factors reduce the tendency of the Group VB elements to attract elec-
trons relative to comparable elements of Group VI BP

3. List the oxidation states of nitrogen. Illustrate each with a specific compound.

4. Describe the preparation of elemental nitrogen. What method could prepare
chemically pure nitrogen? What uses does nitrogen have?

5. Write equations for the production of ammonia from the following substances:
(a) (NH 4 ) 2 S0 4 (b) Mg 3 N 2 (c) concentrated ammonium hydroxide solution.
How can ammonia be dried?

6. Describe the following processes for the manufacture of ammonia: (a) Haber
^b) Cyan amide.

7. State the experimental conditions which favor the formation of ammonia
in the Haber Process.

8. What oroperties of ammonia make it suitable for use as a refrigerant?

The Elements of Group VB: Nitrogen

389

10. Sketch the shape of the ammonium ion, using the drawing of the ammonia
molecule as a model.

11. Explain why an aqueous solution of ammonia gives a basic reaction.

12. What is ammonolysis? Write reactions of liquid ammonia that are analogous
to those of liquid water.

13. From the viewpoint of molecular structure explain why oxides of nitrogen
are acidic.

14. Write equations for the heating of (a) ammonium nitrate (b) sodium nitrate
(c) nitric acid.

15. Describe the methods for the production of nitric acid from its elements and
water; state the conditions and write all equations. Whv is concentrated nitric
acid usually yellow?

16. Write equations for the reaction between concentrated nitric acid and (a) silver
(b) lead (c) phosphorus (d) iodine (e) sulfur.

17. Write equations for the reaction between dilute nitric acid and (a) silver
(b) copper (c) zinc (d) hydrogen sulfide (e) iron (II) sulfate.

J8. How could you determine whether (a) a sample of N 2 0 contained NO (b) a
sample of NO contained N 2 ?

19. What volume of N z O, measured at STP, can be obtained from 80.0 g of

NH 4 N0 3 ? Ans; 22.4 liters

20. What volume of 0.50M HN0 3 can be produced from the acid obtained bv

the oxidation of 170 g of NH S ? Ans\ 20 liters

21. What volume of concentrated HNO s (density 1.4 g/ml; 68% HNO a ) is re-
quired to neutralize 250 g of limestone (assume the limestone is 100% CaC0 3 )?

Ans; 33 ml

22. What volume of 1.5M HNO s is required to oxidize 4.0 g of Fe 2+ to Fe 3+ ?

Ans : 16 ml

23. What is the pH of one liter of 0.10M NH 3 to which has been added 0.30

mole of solid NH 4 C1? Ans; 8.8

24. Assuming 100% yield, what weight of NH 3 can be produced from nine tons

of carbon by the cyanamide process? Ans; 8.5 tons

25. A refrigerating engine uses NH 3 as its operating fluid. Assuming 100% ef-
ficiency of heat transfer, to freeze 100 g of H 2 0 (a) what weight of NH 3
liquid must be vaporized and (b) what volume would this weight of NH 3
vapor occupy at its normal boiling point, assuming that the vapor behaves
as an ideal gas?

The Elements of Group VB
Phosphorus, Arsenic,
Antimony, and Bismuth

Phosphorus

Phosphorus (Greek, phosphoros : light hearer) was discovered in 1669 by
the German alchemist, Brandt. While searching for the philosopher’s stone,
a substance that was thought capable of changing baser metals into gold,
he distilled a mixture of sand and evaporated urine and obtained a substance
which had the property of glowing in the dark. Much later Lavoisier proved
this substance to be the element phosphorus.

Because of its reactivity with oxygen, phosphorus does not occur free
in nature but is widely distributed in the form of phosphates. The most
abundant and the most important is the mineral apatite , Ca 3 (P0 4 ) 2 * CaCl 2
(or CaF 2 ), which makes up a large part of the phosphate rock deposits in
the southeastern states. Phosphorus is an essential component of both plant
and animal nutrition. The animal body contains complex organic phosphorus
compounds in the nerves, muscles, and brain; calcium phosphate constitutes
about 30% of the bones and teeth.

1. Preparation of Phosphorus. Phosphorus is now made by heating na-
tural calcium phosphate, Ca 3 (P0 4 ) 2 , with sand, Si0 2 , and coke, C, in an
electric furnace at a temperature of 1450° C (Figure 30.1). The over-all
reaction is

(1) 2 Ca 3 (P0 4 ) 2 + 6 Si0 2 + 10 C P 4 + 6 CaSiO, + 10 CO

This reaction may be viewed as taking place in two stages; first, a displace-
ment of phosphorus (V) oxide, P 4 O 10 , by the Si0 2 followed by a reduction
of the P 4 O 10 by C.

(2) 2 Ca 3 (P0 4 ) 2 + 6 SiO, 6 CaSi0 3 + P 4 O 10

(3) P 4 O 10 4- 10 C P* + 10 CO

At the temperature of the furnace, the phosphorus vapor distills off and

The Elements of Group VB: Phosphorus , Arsenic , Antimony, and Bismuth

391

CaSiOs collects as a liquid slag at the bottom of the furnace; it is drawn
off at intervals and discarded.

Calcium Phosphate,
Sand and Coke

Figure 30 A. An Electric Furnace for the Production of Phosphorus.

• •

(a) (b)

W The electron structure of the P 4 molecule. The phosphorus atoms are joined
by single covalent bonds; each phosphorus atom has one pair of unshared electrons.

(b) The tetrahedral arrangement of the phosphorus atoms.

Figure 30.2. The Phosphorus Molecule, P 4 .

2. The Phosphorus Molecule, Vapor density measurements indicate that
the phosphorus molecule has the formula P*. The molecular structure is

392

The Elements of Group VB: Phosphorus , Arsenic , Antimony , ami Bismuth

tetrahedron. Each phosphorus atom is joined by a single covalent bond to
each of the other three atoms, and each has one pair of unshared electrons.
The sharing of three electron pairs and the unshared pair enable each
phosphorus atom to have eight electrons in its valence shell (Figure 30.2).
If phosphorus vapor is heated above 800° C, the P 4 molecule dissociates into
P 2 molecule. Presumably the P 2 molecule has a structure similar to that of
the N 2 molecule, with a triple bond between the phosphorus atoms.

3. Allotropes of Phosphorus. Solid phosphorus exists in several allotropic
modifications: white phosphorus, red phosphorus, violet phosphorus, and
black phosphorus. Only the white and red forms are of general interest.

(A) White phosphorus. White phosphorus is a translucent, wax-like solid,
obtained by the condensation of phosphorus vapor. Exposed to light the
white variety changes slowly to the more stable red form. A superficial coating
of red phosphorus is produced, giving the mass a yellowish appearance and
hence this form is sometimes called yellow phosphorus. When exposed to
moist air, white phosphorus glows in the dark due to slow oxidation and
liberation of some of the reaction energy as green light. Such a production
of light by a chemical reaction is known as chemiluminescence . The crystal
lattice of white phosphorus is composed of P* molecules held together by
weak Van der Waals forces.

White phosphorus ignites spontaneously in air at about 40°C. Because
of this flammability it must be stored under water to prevent oxidation and
spontaneous combustion. It is almost insoluble in water but is extremely
soluble in carbon disulfide, and less so in alcohol, ether, liquid NH 3 , and
liquid S0 2 . Even at ordinary temperatures white phosphorus has an ap-
preciable vapor pressure. Both the solid and the vapor are poisonous. When
breathed continuously, small amounts of the fumes cause necrosis, or rotting
of the bones of the jaw and nose; a dose of 0.15 gram of the solid can
be fatal. White phosphorus should never be handled with the bare fingers;
forceps or tongs should be used. Burns produced by phosphorus are painful
and slow to heal.

(B) Red phosphorus. Red phosphorus is obtained by heating the white
variety to about 250° G in the absence of air; the change is catalyzed by a
trace of iodine.

(4) Phosphorus (white) -» Phosphorus (red) AH = -17.6 kcal

Because the transformation of white to red phosphorus is exothermic, the
red form is the more stable. Red phosphorus is made up of minute crystals
in which P 4 tetrahedra are joined together in long chains through phosphorus-
phosphorus bonds. These bonds extend throughout the entire crystal which
is, in effect, a giant molecule. This structural difference between red and
white phosphorus causes a wide disparity in their properties.

Red phosphorus has a greater density. Up to about 300° C it has little
vapor pressure and it does not ignite till about 250° C. Heated at atniospheric
pressure it vaporizes without melting; only under great pressure does it
melt at about 600° C. Red phosphorus is insoluble in those solvents which
dissolve the white variety, It does not phosphoresce and it is not poisonous.

The Elements of Group V-B: Phosphorus, Arsenic , Antimony, and Bismuth

393

(C) Other Forms of Phosphorus. Red phosphorus dissolves in molten
lead; if the mass is allowed to solidify and if the lead is dissolved, crystals
of violet phosphorus remain. The properties of red and violet phosphorus
are essentially the same. It is believed the two are chemically identical and
differ only in crystal size. Metallic or black phosphorus is obtained by
heating white phosphorus at 200°C under high pressure, about 10,000
atmospheres.

4. Properties of Phosphorus. Physical properties : These are given in
Table 29-A.

Chemical properties : Phosphorus is an active element, much more so than
nitrogen; it combines directly with oxygen, sulfur, the halogens, and many
metallic elements. White phosphorus is more reactive than the red form
though the products of their chemical reactions are identical. Specific re-
actions will be given as the compounds of phosphorus are considered.

5. Uses of Phosphorus. The major use of phosphorus is in the preparation
of phosphate salts to be used as fertilizers. Other uses are the manufacture
of matches, phosphor bronzes, and as a rat poison. Because of the poisonous
nature of phosphorus vapor the use of elemental phosphorus in matches has
been legally prohibited in all countries. At present two types of matches are
in general use, the ‘strike anywhere” match and the safety match. The former
contains in its head a mixture of an oxidizing agent (KNO a ), a combustible
substance (sulfur or rosin), and a glue binder. Though combustible this
mixture does not ignite when rubbed, The mixture is therefore tipped with
a compound of phosphorus (P 4 S 3 )> and an oxidizing agent Pb0 2 ). The
two parts of the match head are usually colored differently. Friction produces
sufficient heat to ignite the P 4 S 3 , which in turn ignites the match head
and then the wood or paper. In the safety match the head contains no
phosphorus but is a mixture of an oxidizing agent (KC10 3 ) and a com-
bustible substance (Sb 2 S. t ). A side of the match box or match book cover is
coated with a mixture of red phosphorus and powdered glass. Friction vapor-
izes a trace of the red phosphorus which ignites and sets fire to the head of
the match.

Phosphorus bums in oxygen with a brilliant white flame, forming clouds
of white, solid phosphorus(V) oxide, P4O10. In moist air such a cloud gives
a dense, opaque fog of minute droplets of phosphoric acid which remains
suspended in the atmosphere. For this reason, the burning of white phos-
phorus has been used in warfare to produce a screening smoke. Also because
of its combustibility, phosphorus has found application in incendiary bombs
and in tracer bullets.

6. Phosphine. Unlike ammonia, phosphine, PH 3 , cannot be prepared by
direct union of its elements. It can be made by (A) boiling white phos-
phorus in a concentrated solution of KOH or NaOH and (B) the hydrolysis
of phosphides of active metals.

(5) P 4 + 3 OH- + 3 H 2 0 PH, + 3 H 2 P0 2 “ (hypophosphite ion)

(6) Ca 8 P 2 + 6 H a O 2 PH 3 + 3 Ca(QH)*

394

The Elements of Group VB: Phosphorus , Arsenic , Antimony , and! Bismuth

At STP phosphine is a colorless, poisonous gas with an offensive odor.
Its melting point is -132° C and its boiling point is -86° C. The structure
of the phosphine molecule is analogous to that of ammonia. Phosphine
is less stable than ammonia and is readily decomposed by heat into its
elements. Unlike ammonia, phosphine is only slightly soluble in water, 11
volumes per one of water at STP, and yields no basic compound analogous
to NH 4 OH. With the hydrogen halides, phosphine combines to form phos-
phonium compounds resembling in composition the ammonium salts, PH 4 C1,
PH 4 Br, and PHJ. Of these, phosphonium iodide, PH 4 I, a colorless crystalline
solid, is the most stable. The phosphonium compounds do not yield the
phosphonium ion, PH 4 +, in solution but are decomposed by water into PH 3
and acid.

7. Halides of Phosphorus. The halogens react readily with phosphorus
to form halides, of which the most common are trihalides, PX 3 , and penta-
halides, PX5. The reactions of phosphorus with chlorine are typical.

(7) P 4 + 6 Cl 2 4 PC1 3 (phosphorus (III) chloride; when excess P is

used)

(8) P 4 + 10 Cl 2 4 PCI5 (phosphorus(V) chloride; when excess Cl 2 is

used)

The pentahalides can also be formed by heating the trihalide with the
free halogen.

(9) PC1 3 + Cl 2 PCI*

Of the types PX 3 and PX 5 , only PI 5 is not known; mixed halides, such as
PFC1 2 and PF a Cl 2 , have also been prepared, as has also the unusual P 2 I 4 .
The structures of some halogen compounds of phosphorus are shown in
Figure 30.3.

PC1 3 PC1 5 POCl a

A tetrahedral molecule with A bipyramidal molecule; A tetrahedral molecule with
an atom at each apex. three chlorine atoms form an the phosphorus atom at the

equilateral triangle around center of the tetrahedral
the central phosphorus atom. body.

Figure 30.3. Structures of PC1 3 , PC1 5 , cmd POCH a .

The Elements of Group V-B ; Phosphorus , Arsenic , Antimony, and Bismuth

395

In PC1 8 the phosphorus atom forms five covalent bonds so that ten
electrons are involved in its bonding. The availability of 3 d orbitals, in ad-
dition to its 3s and 3p orbitals, enables the phosphorus atom to expand its
valence shell beyond eight electrons and to form compounds of which
nitrogen is incapable. Thus NC1 5 is not known; d orbitals are unavailable
to the nitrogen atom. In the solid state, or in solvents of high dielectric con-
stant, PC1 5 has an ionic structure, [PC1 4 ]+ [PC1 6 ]-. In the PC1 6 - ion, the
phosphorus atom uses twelve electrons in bonding.

The phosphorus halides are completely and irreversibly hydrolyzed by
water; this hydrolysis has previously been cited as a* preparation of the
hydrogen halides.

(10) PC1<, + 3 H 2 0 — » 3 HCI + H 3 P0 3 (phosphorous acid)

(11) PC1 5 + H 2 0 — > 2 HCI + POCI3 ( phosphorus ( V) oxychloride;

partial hydrolysis in a limited
amount of cold water)

(12) PC1 5 + 4 H 2 0 -» 5 HCI + H3PO4 ( orthophosphoric acid; complete

hydrolysis in excess water)

8. Oxides of Phosphorus. Phosphorus forms the oxides, P 4 0 6 , P4O10,
and PgOjg. The first two, corresponding to oxidation states of +3 and +5,
respectively, are of most importance; P 8 0 16 can be considered to be a compound
of the other two. The compounds P 4 0 6 and P 4 O 10 are often written as their
empirical formulas P 2 0. s and P 2 0 8 , and on this account are referred to as
phosphorus trioxide and phosphorus pentoxide, respectively. However, vapor
density measurements indicate that the true molecular formulas are P 4 O e
and P 4 O 10 . This is not surprising in view of the fact that elemental phosphorus
is tetra-atomic so that formation of the oxides merely involves coordination
of the phosphorus atoms with six or ten oxygen atoms as shown in Figure 30.4.

Both oxides are crystalline white solids. They react with water to form
acids. Phosphorus (III) oxide dissolves slowly in water to form phosphorous
acid, H s PO«.

(13) P 4 Or, + 6 H.O 4 HJPO.

Phosphorus ( V) oxide forms a number of acids, the specific one formed
depending upon the amount of water used. The reaction of P 4 O 10 with
water is violent; the compound has a great affinity for water and is one
of the most powerful dehydrating agents known.

(14) P 4 O 10 + 2 H 2 0 -*• 4 HPOa (metaphosphoric acid)

(15) P 4 0 1() + 4 H 2 0 2 H 4 P 2 0 7 (pyrophosphoric acid)

(16) P.O,. + 6 H 2 0 -» 4 H 3 P0 4 (orthophosphoric acid)

No oxidation-reduction is involved in any of these reactions; the oxidation
state of the phosphorus in the acid is the same as that in the oxide from
which it was produced. The relation between these phosphoric acids is
merely the degree of hydration. Where the same anhydride forms more

396

The Elements of Group VB: Phosphorus , Arsenic , Antimony, and Bismuth

P4O

O = Oxygen
• ss Phosphorus

The phosphorus atoms are located at the comers of a regular tetrahedron; oxygen
atoms are located at positions in between; in P 4 O 10 a phosphorus atom and its four
mounding oxygen atoms also form a tetrahedral structure.

Figure 30.4 . Structures of P 4 O s and P 4 O 10 .

than one acid, differing only in the degree of hydration, the most hydrated
acid is called the ortho acid, and the least hydrated acid, the meta acid.

In practice the hydration of P 4 O 10 cannot be stopped at the intermediate
stage of pyrophosphoric acid. The latter is made by the heating and dehydra-
tion of orthophosphorie acid at about 250°C; heating to higher temperatures
causes further dehydration and the formation of metaphosphoric add.

(17) 2 H 3 P0 4 HJP 2 O 7 + H 2 0 (heating at 250°C)

(18) H 3 P0 4 HPO s + H 2 0 (heating above 400°C)

9. Acids of Phosphorus. The more important acids of phosphorus are the
three phosphoric acids in which the oxidation state of the phosphorus is +5,
phosphor ous acid, H 3 P 0 3 , wherein the oxidation state is + 3 , and hypo-
phosphorous add, H 8 P0 2 , in which the phosphorus is +1, In the pure
state at STP, these adds are all colorless, deliquescent solids; all are soluble
in water. The structures of the acids are drawn below.

10. Orthophosphorie Add and Phosphates. The most important of the
acids of phosphorus is orthophosphorie add, H 3 P0 4 , commonly referred to
simply as “phosphoric acid." Pure H 3 F0 4 can be prepared by the oxidation
of P 4 to P 4 O 10 , followed by solution of the oxide in excess water. Com-
mercially, an 85% solution of the add is made by treating pulverized Ca 3 (P0 4 ) 2
with concentrated H 2 S0 4 , separating the CaS 0 4 by filtration, and then

The Elements of Group VB: Phosphorus, Arsenic, Antimony, and Bismuth

397

:A:

.. | ••

H-O-P-O-H

• • | • •
: 0 :

H h

I I

:o: :o:

•• | •• ..

H-O-P-O-P-O-H

•• j •• . .

:0: to:

Orthophosphoric Add Pyrophosphoric Add
H

H-O-P-O-H

•• | ••

: 0 :

• *

Phosphorous Acid

•• I ••

H — O— P=0

Metaphosphoric Acid
H

•• | ••

: o-p-o-h

•• | •*

Hypophosphorous Acid

evaporating. The resulting concentrated solution of H 8 P0 4 is a viscous liquid
commonly known as “syrupy phosphoric acid”

Phosphoric acid is moderately strong. It is triprotic and ionizes in three
stages, the degree of ionization decreasing with each successive stage.

(19) H 3 P0 4 H+ + H 2 P0 4 “ (primary ionization; K* = 7.5 X 10”®)

(20) H 2 P0 4 ” -» H+ + HPO^ (secondary ionization; K 2 = 6.2 X 10” 8 )

(21) HP0 4 2 ” -> H+ + P0 4 ®- (tertiary ionization; K 3 ~ 1.2 X 1(H 2 )

Corresponding to the three stages of ionization, H 3 P0 4 forms three series
of salts known as phosphates. Examples are sodium dihydrogen phosphate,
NaH 2 P0 4 ; disodium hydrogen phosphate, Na 2 HP0 4 ; and trisodium phosphate,
Na B P0 4 . These phosphates can be prepared by adding a calculated quantity
—one, two, or three equivalents— of NaOH to H 3 P0 4 .

An aqueous solution of NaH 2 P0 4 is slightly acid. The H 2 P0 4 ” ion is
amphiprotiq and can act both as an acid and a base.

as base as acid

(22) H 3 P0 4 + OH” 4- H a O + H 2 POr + H a O H s O+ + HP0 4 2 ‘

Because the reaction with water to form H 3 0+ ions has the greater tendency
to proceed, the solution is acidic. A solution of Na 2 HP0 4 is basic because the
HP0 4 2 - ion is a better proton acceptor than donor. Its predominant reaction is

(23) HP0 4 2 - + H 2 0 -> H 2 P0 4 ” + OH-
Solutions of Na 3 P0 4 are strongly basic due to hydrolysis.

(24) P0 4 8 ' + H 2 0 HP0 4 2 ” +OH-

The solubility of these salts decreases in the order: NaH 2 P0 4 > Na 2 HP0 4 ,
and Na 3 P0 4 . On the first two of these salts, the effect of heat is to remove
the elements of water, leaving NaP0 3 in the first case and Na 4 P 2 0 7 in the
Second; Na s P0 4 is unaffected by heat.

398

The Elements of Group VB: Phosphorus , Arsenic , Antimony , ancf Bismuth

Phosphate salts are used mainly as fertilizers. Plant life requires phos-
phorus in the form of soluble orthophosphate compounds. Phosphate rock,
Ca 3 (P0 4 ) 2 , is not sufficiently soluble to serve as a source of phosphorus
for plants. A more soluble and suitable phosphate is calcium dihydrogen
phosphate, Ca(H 2 P0 4 ) 2 ; it is prepared by the reaction of phosphate rock
with H 2 S0 4 . More than 10 million tons of phosphate rock are converted
annually into fertilizer. A mixture of Ca(H 2 P0 4 ) 2 and gypsum, CaSCV 2H 2 0,
forms a dry mass, sold as “superphosphate of lime” Other uses of the
phosphate salts are: NaH 2 P0 4 as an ingredient of baking powders and in
die treatment of boiler water (hard water) to prevent scale formation;
Na 2 HP0 4 in the 4 weighting” of silk; and Na 3 P0 4 as a detergent as a water
softener, and in boiler water treatment.

11. Pyrophosphoric Acid and Pyrophosphates. Pyrophosphoric acid,
H 4 P 2 0 7 , ik prepared by the moderate heating and dehydration of H 3 P0 4 ,
two molecules of which lose one molecule of water between them. Dissolved
in water, H 4 P 2 0 7 slowly hydrolyzes back to the ortho-acid. Though H 4 P 2 0 7
is a tetraprotic acid only two series of salts have been isolated. These are
the dihydrogen pyrophosphates, such as Na 2 H 2 P 2 0 7 , and the normal pyro-
phosphates, as Na 4 P 2 0 7 . Because it forms soluble complexes with Ca 2 + and
Mg 2 + ions, Na 4 P 2 0 7 is used as a water softener; such complex formation re-
duces the concentration* of Ca 2 + and Mg 2+ ions in solution to such a value
that no precipitate is formed with either soap or carbonate ion, C0 3 2 "
(Chapter 39).

12. Metaphosphoric Acid and Metaphosphates. Metaphosphoric acid is
formed by the strong heating of ortho- or pyrophosphoric acid, or by
the addition of a small amount of water to P 4 O l0 - It is a viscous, glassy
substance which also reacts slowly with water to form H 3 P0 4 . Metaphosphoric
acid has the empirical formula, HPO s . These units polymerize, forming rings
and long chains in which a phosphorus atom is joined tetrahedrally to four
oxygen atoms and the latter form bridges between the phosphorus atoms, so
that the true molecular formula is (HP0 3 ) n . The chain type of polymer is
shown below.

H H H H H

i A A A A

iiii

n>/N/N/- N/ jjN}/ fso/

o o o o o

The salts of metaphosphoric acid are similarly polymerized, e.g., (NaP0 3 ) n -
The sodium salt can he prepared by heating NaH 2 PO*, Na 2 H 2 P 2 0 7 , or
NaNH 4 HP0 4 (microcosmic salt). Trisodium metaphosphate, Na 3 P 3 0 9 or
(NaP0 3 ) ? , is obtained by heating NaH 2 P0 4 at 550° C for several hours.
If the heating. is done at 70Q°C a colorless liquid is formed. When this liquid
is cooled rapidly there is produced a glasslike, water soluble material known
as “sodium hexametaphosphate” because it was once believed to have the

The Elements of Group VB: Phosphorus , Arsenic, Antimony, and Bismuth

399

formula (NaP0 3 ) 6 . This substance also forms a stable complex with Ca 2 +
ion and so is used as a water softener sold under the trade name of “Calgon”

13. Phosphorous Acid and Phosphites. Phosphorous acid, H 3 PO s , is pre-
pared by the hydrolysis of phosphorous (III) halides. Only two of the
hydrogen atoms in the molecule are ionizable. The nonionizable hydro-
gen differs from the other two in that it is bonded directly to the phos-
phorus atom. Hence the formula of phosphorus acid would, be more proper-
ly written as H 2 (HP0 3 ). The acid is moderately strong. When heated
it decomposes into PH 3 and H 3 P0 4 .

X2S) H 3 PO* H+ + H 2 P0 3 - Ki = 1.6 X 10-*'

(26) H 2 P0 3 -~> H+ + HPCV- K 2 = 7.0 X 10~ 7

Two types of salts, known as phosphites, are possible, e.g., NaH 2 P0 3 and
Na 2 HP0 3 . Since they are readily oxidized to phosphates, phosphorous acid
and phosphites are good reducing agents. They reduce the halogens, Ag+ ion
to Ag, and SCV" ion to S.

14. Hypophosphorous Acid and Hypophosphites. Hypophosphorous acid,
H 3 P0 2 , is a weak monoprotic acid. Only the hydrogen atom linked to the
oxygen atom is ionizable. The formula of the acid is better written as
H(H 2 P0 2 ). The hypophosphite ion, H 2 P0 2? is a product of the reaction
between P 4 and KOH. The acid is a strong reducing agent.

ARSENIC, ANTIMONY, AND BISMUTH

The elements, arsenic, antimony, and bismuth, occur to a small extent
in the free state. The important ores are sulfides and oxides:

As 2 S 2 , realgar * Sb 2 S 3 , stibnite Bi 2 S 3 , bismuthinite

As 2 S 3 , orpiment Bi 2 O s , bismite

FeAsS , arsenopyrite
As 2 O s , claudetite

15. Preparation of the Elements. The elements are produced by reduction
of their oxides with carbon; sulfide ores are first roasted to the oxide. The
reactions for arsenic are typical

(27) 2 As 2 S« + 9 0 2 As 4 0 6 + 6 S0 2 (roasting of the sulfide)

(28) As 4 O e + 6 C -» As 4 + 6 CO (reduction of the oxide)

Special processes for arsenic and antimony are:

(29) FeAsS As 4 + FeS (thermal decomposition of arsenopyrite)

(30) Sb 2 S» + 3 Fe -» 2 Sb + FeS (reduction with iron)

16. Properties and Uses. The physical properties of these elements are
given in Table 29-A. Both arsenic and antimony exist in two crystalline
forms, a stable gray metallic form and a metastable yellow form. Bismuth
exists only as a gray metal, with a pinkish tinge. Ordinary arsenic is a steel
gray crystalline' solid, metallic in appearance. It is brittle and is a poor

400

The Elements of Group V B: Phosphorus , Arsenic , Antimony , and Bismuth

conductor of electricity, though a good conductor of heat. Heated to 610°C,
arsenic sublimes; if the vapors of arsenic are cooled rapidly, the yellow,
less stable allotrope is formed. Analogous to white phosphorus, the yellow
form of arsenic has the molecular formula, As*, is very volatile, and is
soluble in CS 2 . It transforms rapidly into the more stable metallic variety.

Arsenic is relatively inert at ordinary temperatures, but when heated
in air it bums with a bluish flame, producing white clouds of the solid
oxide, As 4 O 10 . It does not displace hydrogen from acids, but with nitric
acid and other powerful oxidizing agents, it is oxidized in the same way as
phosphorus and forms arsenic acid, H 3 As0 4 . It unites with the halogens, with
sulfur, and with many metals to form arsenides.

Arsenic alloys with many metals. The addition of 0.5 percent of arsenic
to lead lowers the melting point of the lead and produces an alloy harder
than pure lead. Such an alloy is used for making shot that is more nearly
spherical than shot made with other materials. Only about 100 tons of ele-
mental arsenic are used annually in the United States,

Antimony is very brittle, a poor conductor of heat and electricity, and is
much less volatile than metallic arsenic. Yellow antimony changes so rapidly
to the metallic form that it can be kept only at very low temperatures. The
reactions of antimony are similar to those of arsenic. The principal- use of
antimony is in the making of alloys such as type metal, britannia, and
babbitt. Much antimony is used in the manufacture of plates for lead storage
batteries. Most of the world’s supply of antimony comes from China; the
United States uses about 20,000 tons annually.

Bismuth is used in making alloys of low melting point, much of which
melt below the boiling point of water, e.g.. Wood’s metal, which melts at
71 °C. Such alloys find use in automatic sprinkler systems and as electric fuses.

17. Hydrogen Compounds. The compounds, arsine, AsH 3 , stibine, SbH 8 ,
and bismuthine, BiH 3 , are colorless, extremely poisonous gases with char-
acteristic odors. They can be prepared by

(A) the hydrolysis of metal arsenides, stibnides, or bismuthides

(31) Zn 3 As 2 + 6 H 2 0 -* 2AsH s + 3 Zn(OH) 2

(B) the reduction of active metals by a halide in acid solution

(32) AsC 1 3 + 3 Zn + 3 HC1 AsH 3 + 3 ZnCl 2

To produce BiH 3 by this reaction the more active reducing agent. Mg,
must be used. The hydrogen compounds are unstable, decomposing tinder low
heat into their constituent elements. AsH 3 decomposes readily, SbH s ex-
plosively, and BiH s is so unstable that little is known of its properties.

The thermal instability of AsH 3 and SbH 3 , coupled with the reductive
method of preparation, is the basis of the sensitive Marsh test for the detec-
tion of these elements. Pure Zn and HC1 are added to the sample to be
tested for As or Sb. Any AsH 3 or SbH s produced is passed through a tube
wherein it is heated and decomposed. The presence of As or Sb is indicated
by the formation of a black deposit or mirror, which is due to condensation
of the As or Sb vapor in the cooler portion of the tube beyond the point

The Elements of Group VB: Phosphorus , Arsenic , Antimony , Bismuth

401

of heating. An arsenic mirror is soluble in NaCIO; an antimony mirror is
not. Quantities of arsenic as low as 0.0001 g can be detected so that deliberate
arsenic poisoning, in contrast with medieval times, is now out of vogue.

18. Halogen Compounds. Arsenic, antimony, and bismuth form trihalide
compounds of the type MX 3 with all the halogens. The only pentahalides
which are known are AsF*, SbF 3 , and SbCl lV The halogen compounds are
prepared by the direct union of the elements. All the trihalides are hydrolyzed
by water.

(33) AsCL + 3 H 2 0 3 HC1 + H 3 As0 3 (arsenious acid)

(34) SbCl« + H 2 0 2 HC1 + SbOCl (antimony (III) oxychloride)

(35) BiCl 3 + H 2 0 2 HC1 + BiOCl (bismuth(III) oxychloride)

Unlike the hydrolysis of PCb the above reactions are all reversible so that
SbOCl and BiOCl, both white solids, dissolve upon the addition of concen-
trated HC1. The SbO+ and BiO+ ions are known as the antimonyl and
bismuthyl ions, .respectively, and their formation is another indication of
the relatively greater metallic character of antimony and bismuth.

19. Oxides and Acids. The common oxides of arsenic, antimony, and
bismuth are tabulated below. Of these the most important are those in which
the Group VB element has an oxidation state of +3.

Oxidation state +3

As 4 0 6

Sb 4 O e

Bi 2 0*

Oxidation state +4

Sb 2 0 4

Bi 2 0 4

Oxidation state +5

AS 4 O 10

Sb 4 Oio

Bi 2 0«5

When arsenic or an arsenic compound is burned in air, arsenic (III) oxide,
As 4 0 6 , is produced. It is a white solid, known commercially as white arsenic
or simply as “arsenic.” Of all arsenic compounds it is the most important,
about 60,000 tons being produced annually. The compound is poisonous
and most of its production goes into the manufacture of weed killers and
insecticides. Considerable quantities are also used in the manufacture of
glass to insure a colorless product. Arsenic (III) oxide is slightly soluble in
water, forming the weak arsenious acid, H 3 AsG 3 . It reacts readily with strong
bases to yield arsenites, e.g., Na 3 As0 3 and NaAs0 2 . Even with strong acids
there is some reaction so that the oxide is somewhat amphoteric. Antimony (III)
oxide is definitely amphoteric and bismuth (III) oxide is a weak base.

Arsenic(V) oxide cannot be prepared by oxidation of the (III) oxide.
When As 4 0$ is treated with concentrated HNO a , white crystals of ortho-
arsenic acid, H s As0 4 , are formed.

(36) As 4 0 6 + 8 HN0 3 + 2 H 2 0 -> 4 H 3 As0 4 + 8 N0 2

Heating H 3 As0 4 results in its dehydration to As 4 O 10 . Although the acids
themselves are not known, salts of pyroarsenic acid, H 4 As 2 0 7 , and metaarsenic
acid, HAs 0 3 , have been prepared. Salts containing arsenic(V) are taown
as arsenates. Both arsenites and arsenates are toxic. The oxides of antimony
and bismuth are similar to those of arsenic. In addition, the oxides Sb 2 0 4

402

The Elements of Group V B : Phosphorus , Arsenic , Antimony , and Bismuth

and Bi 2 0 4 are obtained when either the (III) oxide or the (V) oxide is
heated in air. These oxides may be considered as salts, Sb(Sb0 4 ) and
Bi(Bi0 4 ). Free antimonie acid and bismuthic acid are unknown though their
salts have been prepared, e.g,, NaSb(OH) 6 and NaBiO s . The formula .of
antimonie acid is presumably HSb(OH) e > in which the antimony atom has
a coordination number of 6.

20. Sulfur Compounds. The following sulfides are known:

Oxidation state +2

As 2 S 2

Oxidation state +3

As 2 S 3

Sb 2 S 3

Bi 2 S 3

Oxidation state +4

Sb 2 S 4

Oxidation state +5

As 2 S 5

SbsSs

The (III) sulfides can be prepared by direct combination with sulfur, or
by the reaction of H 2 S with a solution containing a (III) compound of the
Group VB element.

(37) 2 H 3 As0 3 + 3H 2 S-^ As 2 S 3 + 6 H 2 0

Both As 2 S 3 and Sb 2 S 3 are soluble in alkali metal sulfides, such as Na 2 S, by
forming complex ions; Bi 2 S 3 does not undergo this reaction. The reaction
is reversed and the sulfide reprecipitated by addition of acid.

(38) As 2 S 3 + 3 S -» 2 AsS 3 3 ~ (thioarsenite ion)

(39) 2 AsSs 8 - + 6H+^ As 2 S 3 + 3 H 2 S

Arsenic(V) sulfide can be prepared by fusing As 2 S 3 with S, or by passing
H 2 S through a solution of an arsenate. In Na 2 S solution, As 2 S 5 dissolves
by forming the thioarsenate ion, AsS 4 3 ~

QUESTIONS

1. Compare phosphorus with the other elements of Group VB. Write the formulas
of analogous compounds.

2. Discuss the allotropes of phosphorus. Why should one allotrope be more
active chemically than another?

3. Describe the production of phosphorus, writing pertinent equations.

4. Write equations for the following reactions; (a) P -j- HNO a (b) P+ KOH
(c) P + I 2 (d) P + Mg (e) PBr a + H 2 0 (f) P 4 O 10 + H 2 Q (g) P 4 0 6 + H a O
(h) Ca 3 P 2 + H 2 0 (i) H s P 0 3 + Ag+.

5. Write equations for the preparation of PH 3 . Compare the properties of PH 3
and NH a .

6. Write equations f6r the preparation of PC1 3 and PC1 5 . What is the nature
of the bond between the phosphorus atom and (a) the Cl atom in PC1 5
(b) the Ca atom in Ca 8 P 2 ?

7. Account for the fact that PF 5 is stable but PI 5 has not been prepared.

8. (a) Which is the more acidic oxide (1) N 2 0 6 or N 2 O a (2) N 2 0 5 or P 4 O 30
(b) which is the stronger acid (I) H 3 PC 4 or H 3 PO a (2) H 3 P0 4 or H 3 As0 4 ?
Explain briefly.

The Elements of Group VB; Phosphorus , Arsenic , Antimony , and Bismuth

403

9. What experimental evidence is there for the use of the formulas P 4 0 6 and
P 4 O 10 instead of the simpler empirical formulas?

10. Starting with elemental phosphorus, write equations for the preparation
of (a) H 3 P0 4 (b) Na 2 H 2 P0 4 (c) NaPO s .

11. Tabulate the phosphoric acids and their sodium salts. Name each.

12. Write equations for the successive stages of ionization of H 3 P0 4 . Explain
why a solution of NaH 2 P0 4 gives a slightly acidic reaction and Na 2 HP0 4 a
slightly basic reaction.

13. Calculate the hydrolysis constant, K h , for the reaction in Equation 23.

14. How are the elements arsenic, antimony, and bismuth prepared from their
ores?

15. Write electron dot structures for hypophosphorous acid, phosphorous acid, and
pyrophosphoric acid. Explain why hypophosphorous acid is monoprotic.

16. How do BiOCl and NaOCl differ in structure?

17. Write equations for the heating of (a) H 3 P0 4 (b) NaH 2 P0 4 (c) NaNH 4 HP0 4
(d) Na 3 P0 4 .

18. What weight of phosphorus can be obtained from 40 kg of phosphate rock which

contains 75% Ca 3 (P0 4 ) 2 ? Ans: 6.0 kg

19. Twenty liters of PH 3 taken at 25 °C and 700 mm are burned completely. The

combustion products are dissolved in water to make one liter of solution.
What is the normality of the acid formed? Ans: 2.3 N

20. What weight of “superphosphate of lime” can be made from 3.0 tons of
Ca s (P0 4 ) 2 ?

21. Calculate the concentrations of HoP0 4 ~, HP0 4 2 ~, and P0 4 a ~ ions in 0.10M

h 3 po 4 .

The Elements of Croup IVB

Carbon

The elements of Group IVB, carbon, silicon, germanium, tin, and lead
stand midway in the Periodic System. The electron configurations of these
elements are listed below; other properties are given in Table 31-A.

Element

Atomic
Number •

1 K
Is

2s 2 p

3s 3 p 3 d

4s 4 p 4 d 4 f

5s 5 p 5 d

6s 6p

Carbon

2 2

Silicon

2 6

2 2

Germanium

2 6

2 6 10

2 2

Tin

2 6

2 6 10

2 2

Lead

2 6

2 6 10

2 6 10 14

2 6 10

2 2

1. General Properties. All the elements have four valence electrons,
ns 2 ,, np 2 . In theory, the elements may either gain or lose four electrons to
attain a noble gas configuration, but the large values of energy that would be
involved in such a transfer preclude this possibility and no compounds
exist in which there are simple quadruply charged ions. The vast majority
of compounds are formed through covalent bonding.

Perhaps no group in the Periodic System is more striking in its gradation
from the nonmetallic carbon to the metallic lead. The normal variation in
properties with increase in atomic number is again evident in this group.
These trends are shown below.

Increasing
metallic
character
of the
elements

Decreasing

co 2

Decreasing

ch 4

acidity

SiOj

stability

SiH 4

(increasing

Ge0 2

of the

GeH,

Sn]

basicity)

Sn0 2 }

hydrogen

SnH 4

of the
oxides

PbO*f

compounds

PbH 4

Because of the marked change in properties from nonmetallic to metallic
within thctgroup, it is expedient to discuss the nonmetals, carbon and silicon.

The Elements of Group TVS: Carbon

405

Table 3 1-A

Properties of the Elements of Group IVB

Property

Carbon

Silicon

Germanium

Tin

Lead

Symbol

msm

Atomic Number

Atomic Weight

12.011

28.086

72.59

118.69

207.19

Isotopes (mass numbers

12(98.89)

28(92.18)

70(20.45)

112( 0.95)

202( 0.0004)

and percent)

13( 1.11)

29( 4.71)

72(27.41)

1 14( 0.65)

204( 1.37)

3(K 3.11)

73( 7.77)

115( 0.34)

200(26.25)

74(36 f 58)

116(14.24)

207(20.82)

76( 7.79)

117( 7.57)
118(24.01)
119( 8.58)
120(32.97)
122( 4.71)
124( 5.98)

208(51.56)

Abundance in Earth's Crust, %

0.032

25.7

0.0007

0.004

0.0016

Physical State at STP

solid

gray

solid

gray

solid

Allotropes

diamond;

gray

gray;

graphite;

“amorphous

white

“amorphous”

black

Melting Point, °C

3500(s)

1410

937

232

327

Boiling Point, °C

4830 ?

2680

2830

2687

1717

Density at STP, g/cm ; *

3.51(d)

tsss

5.32

5.75(g)

11,35

2.22(gr)
1.8-2. 0(a)

7.30(w)

Heat of Fusion, keal/mole

11.7

7.6

1.72

1.14

Heat of Vaporization,

170

105

79.9

69.4

42.9

keal/mole

Ionization Potential, 1st, eV

11.26

8.15

8,13

7.33

7,42

Electronegativity

2.5

1.8

Covalent Radius, A

0.77

1.11

1.22

1.41

1.47

Ionic Radius, A (4-4)

0.41

0.53

0.73

0.84

Oxidation States

Oxidation Potential, volt

-4 to +4

-4, 4>4

-4, +2, +4

4-2, 4-4

4-2, +4

4-0.13

for M — > M a + -jh 2 e

0.0

4-0.14

Key to symbols: (s) = sublimes; (d) = diamond; (gr) = graphite; (g) = gray; (w) = white
(a) ~ amorphous

separately and to defer the study of the metals, germanium, tin, and lead,
to later chapters where the metallic elements are considered.

Carbon

Although carbon (Latin, carbo : coal) constitutes but 0.032% of the earths
crust, over 500,000 compounds, or more than 90% of all known compounds, con-
tain carbon. Because of this vast number of carbon compounds a separate
branch of chemistry, organic chemistry , is devoted solely to their study. In na-
ture, carbon occurs both as the free element and in the combined state. Elemen-
tal carbon is found in two allotropic forms, diamond and graphite. Many
chemists believe there is a third form, amorphous or nonciystalline carbon,
e.g., coal, charcoal, lamp black, and soot, but it is more likely that this

406

The Element 4 of Group TVB: Carbon

form is composed of microscopic crystals of graphite. In compounds, carbon
is found combined with hydrogen in natural gas and petroleum as hydro-
carbons; with oxygen as carbon dioxide; and in vast deposits of metal
carbonates such as limestone, CaCO s , and magnesite, MgC0 8 . All living
matter, both plant and animal, is composed of carbon-containing compounds.

2. Allotropes of Carbon. (A) Diamond. Natural diamonds are found in a
number of localities but the richest deposits are located in South Africa and
Brazil. Diamonds occur in a hard blue clay and resemble rough pebbles.
For ornamental purposes the natural stone is cut by grinding- new faces
with diamond dust in definite patterns. The Cullinan diamond, found in
South Africa in 1905, was the largest known specimen, weighing 3106 carats 1 ;
it was subsequently cut into nine principal gems which now form part of
the British Crown Jewels. Diamonds are usually tinted a faint yellow but
may be colored red, green, blue, or black due to slight impurities; in some
cases the color of the diamond enhances its value. Imperfect or black
diamonds, known as borts or carbonados, have no value as gems but
are used in tools for cutting and drilling, and for polishing other
stones. The brilliance of a diamond is due to its high refractive index, 2.4;
that of glass is about 1.5. In passing through a diamond, light is reflected by
the many facets into which the gem is cut and is also diffracted into its
component wavelengths to give the diamond its sparkle and “fire ”

The structure of the diamond crystal is tetrahedral. Each atom is joined by
a single covalent bond to each of four other carbon atoms which are
situated at the comers of a tetrahedron. This structure is continuous so that
the crystal is truly a single macromolecule (Figure 31.1 a). The carbon-
carbon distance is 1.54 A. Because of its compact structure, diamond has a
density, 3.5 g/cm 3 , which is relatively high for an element so low in atomic
weight. Because its electrons are tightly held in covalent bonds and are
not free to move, diamond is a nonconductor of electricity, though a good
conductor of heat. The melting point of diamond is probably higher than
that of any other element. Diamond is brittle but is the hardest substance
known in that it will scratch, but cannot be scratched by, any other substance.

At ordinary temperatures diamond is chemically inert. Heated in air,
the product of combustion is pure carbon dioxide, with perhaps a trace of
ash. If diamond is heated in the absence of air, a slow transformation will
take place to graphite, the more stable form.

(1) C (diamond) — > C (graphite) AH — ^450 cal

(B) Graphite. Probably no two allotropic forms differ from each other
so strikingly as do diamond and graphite. Graphite is a soft solid with a
grayish-black metallic luster and a slippery, greasy feel. It conducts electricity
fairly well, its conductivity being about 1/1000 that of a metal.

Graphite crystallizes in hexagonal plates. Its carbon atoms are arrayed
in planes or layers composed of hexagonal rings (Figure 31.1 b). Each carbon
atom is joined by a single covalent bond to two other carbon atoms and
by a double bond to a third carbon atom (Figure 31,2). The bonding

carat is 0.20 gram.

The Elements of Group IV B; Carbon

407

(a) Diamond

(b) Graphite

The carbon atoms in diamond are tetrdhedrally arrayed and their arrangement in
graphite is hexagonal.

Figure 31.1. Crystal Structures of Diamond and Graphite.

I I I
->" c / c '

' C ^ c / C ^c /C ^c^

Figure 31.2 . Graphite
A top view of the carbon atoms in a
portion of a graphite layer is shown.
Hexagons in adjacent layers are not
directly over each other but are dis-
placed; hexagons in alternate layers
are in line.

electrons are not so strongly held as in diamond and the positions of the
double bonds are not fixed so that resonance forms are possible. The
mobility of its electrons, characterized by the shifting double bonds, gives
graphite its electric conductivity and metallic appearance. The forces bond-
ing the carbon atoms within a hexagon are strong, this carbon-carbon distance
being 1.42 A. But the layers of hexagons are separated by a greater distance,
3.4 A, so that carbon atoms in different planes are attracted to each other
by weaker forces. The weakness of the interplanar forces accounts for the
soft, flaky character of graphite; also because one plane can slide readily over
another, graphite is a good lubricant. On the other hand, because of the
strong covalent forces operating between the carbon atoms within a hexagon,
graphite also has a high melting point.

408

The Elements of Group IV B: Carbon

The heat of combustion of graphite is less than that of diamond. Of the
two allotropic forms, graphite has the lesser internal energy and is the
more stable modification (Equation 1).

(2) C (graphite) + 0 2 C0 2 AH = -94,030 cal

(3) C (diamond) + 0 2 — > C0 2 AH = —94,480 cal

The problem of converting graphite into diamond is one which, for
obvious reasons, has intrigued many chemists. The problem is simply one
of puckering the planar hexagonal rings of graphite into a tetrahedral
configuration— and, lo, we have diamondsl In nature this is accomplished
under the great pressures and temperatures which exist below the surface
of the earth. Very small diamonds were reported to have been fabricated
in the laboratory by the French chemist, Henri Moissan, in 1893. Molten
iron dissolves carbon readily. Upon sudden chilling of a saturated solution
of carbon in iron, some dissolved carbon separates out as a fine powder,
consisting in part of true diamonds. The largest of these had a diameter
of 0,5 mm, too small to have any value as a gem. In 1955, the General
Electric Company disclosed the successful production of artificial diamonds,
about 1.7 mm in diameter, by subjecting graphite in an electric furnace to
temperatures of 2800°C at a pressure of 1,500,000 pounds per square inch.
A metal catalyst-chromium, iron, tantalum, and others— must be used to
initiate the growth of the diamond crystals.

Natural graphite is widely distributed. Large deposits occur in Ceylon
and Siberia, and to a lesser extent in the United States. Synthetic graphite
can be made from coal by the Acheson process. Anthracite coal or coke, with
some iron oxide as a catalyst, is heated in an electric furnace at 3500°C.
Using an alternating current of 40,000 amperes at 200 volts the conversion
of coal into graphite is complete in about 24 hours. Graphite is used in the
manufacture of electrodes because of its electric conductivity,* in crucibles
for high temperature work because of its chemical inactivity, as a lubricant
directly or in an oil or water suspension, and mixed with clay in 'lead”
pencils because of its softness and color. The higher the proportion of
graphite in the pencil, the softer is the “lead.” At one time it was thought
that graphite contained lead and it was therefore called black lead or
plumbago .

(C) Amorphous carbon. Other forms of carbon include coal, coke, wood
and- bone charcoal, and lampblack. These are amorphous forms of carbon.
Coal is fossil vegetable matter. Slow decomposition resulted in a loss of
hydrogen as hydrocarbons, and oxygen as carbon dioxide, with the residue
becoming progressively richer in free carbon. During the stages of meta-
morphosis, or “carbonization,” the carbon content increased from about 40%
in the original wood (cellulose) to about 60% in peat, 70% in brown coal
or lignite, 78% in bituminous coal, and about 90% in anthracite coal. The
fuel value of these products is dependent upon the carbon content.

Coke is almost pure carbon. It is produced by the destructive distillation,
or heating in the absence of air, of bituminous coal. During the distillation,
several volatile products, such as ammonia, illuminating gas, benzene, and

The Elements of Group TVB: Carbon

409

coal tar, are given off. The residue, composed primarily of free carbon, is
known as coke.

Wood charcoal is made by the destructive distillation of wood. Valuable
by-products include acetic acid, methyl alcohol (wood alcohol), and ace-
tone. Wood charcoal is highly porous and has a vast internal surface area.
One cubic centimeter of wood charcoal may have a total surface area as
large as 10 million square centimeters. Upon this surface, charcoal adsorbs
many times its volume of gases, actually vapors, such as NH 3 , H 2 S, and Cl 2 .
Hence charcoal is used in the canister of a gas mask as an adsorbent for
toxic gases. The adsorptive ability of charcoal may be increased by treat-
ing it with high temperature steam. This breaks down the pores into a
greater surface area and removes any initially adsorbed gases. Charcoal so
treated is called “activated charcoal.”

Animal charcoal or boneblack is the residue obtained from the destruc-
tive distillation of bones. It consists of approximately 10% carbon and is
especially useful for removing coloring matter in the refining of sugar.

Lampblack or carbon blade is obtained when natural gas and other
carbon compounds are burned in a limited supply of air. The resulting black
soot is one of the purest varieties of amorphous carbon, containing 98.6%
carbon and 1.4% hydrogen. It is used in the compounding of rubber
for tires to make a tougher product, and also in printers' and india ink,
shoe polish, and black paint.

3. Properties of Carbon. Physical properties : These are listed in Table 31-A.
Chemical properties : At ordinary temperatures, carbon is chemically un-
reactive. Above 500° C, carbon combines directly with oxygen to form
carbon monoxide, CO, and carbon dioxide, C0 2 ; with sulfur to form carbon
disulfide, CS 2 , and with many metals and nonmetals to give carbides.

Table 31-B — _

Properties of Carbon Dioxide and Carbon Monoxide

Property

Carbon Dioxide

Carbon Monoxide

Molecular Formula

C0 2

Molecular Weight

44.01

28.01

Physical State at STP

colorless gas

Melting Point, °C

-56 (at 5.3 atm)

-200

Boiling Point, °C

sublimes at -78.5

-190

Critical Temperature, °C

31.1

-140

Critical Pressure, atm

72.8

35.0

Density at STP, g/1

1.96

1.25

Density of liquid,' g/ml

0.95

0.T93

3:1

Solubility in water at STP,
volume ratio

170:1

4. Carbon Dioxide. Carbon dioxide is present in the atmosphere,
0.03-0.04% by volume. The respiration of animals and the combustion of
carbonaceous fuels are similar in that 'both processes release C0 2 to the
atmosphere while removing 0 2 . Conversely, green plants take in C0 2 and
liberate 0 2 during the process of photosynthesis. In photosynthesis, CU 2 ana

410

The Elements of Group IVB: Carbon

H 2 0 react, utilizing solar energy and with chlorophyll as a catalyst, to form
carbohydrates such as cellulose, starch, and sugar, and free 0 2 . Through
this carbon (and oxygen) cycle, a material balance results which keeps the
atmospheric C0 2 content fairly constant. Carbon dioxide is not poisonous:
the harmful effects of “C0 2 poisoning” are more likely the result of a de-
ficiency of 0 2 . Indeed, C0 2 acts as a stimulant to the lungs to initiate
breathing.

Because of its high critical temperature, carbon dioxide can be liquified
at room temperature. It is stored as a liquid under pressure in steel con-
tainers. If the pressure is released, evaporation of some of the liquid may
cause a portion of the remainder to be cooled below its freezing point and
solid C0 2 is produced as a snowy mass. The solid has a vapor pressure of
760 mm at -78.5*C and a pressure of 59 atm at 20°C. Hence, at atmospheric
pressure the solid evaporates without melting and is used as a cooling medium
popularly called “dry ice” because it leaves no liquid residue as does
“H 2 0 ice ”

The C0 2 molecule is linear. The experimentally determined carbon- oxygen
distance of 1,15 A is smaller than the value calculated from theory if only
the first of the structures below existed, so that the C0 2 molecule is best
portrayed as a resonance hybrid of three structures.

• O —O—O * • o=c — O ■ * O — 0—0 *

• * ••

5. Preparation of Carbon Dioxide. (A) The complete combustion of
elemental carbon or of compounds of carbon in excess oxygen or air produces
carbon dioxide and oxides of the other elements combined with the carbon.
All such combustions are exothermic.

(4)

C 0 2 — . COj

AH = - 94 kcal

(5)

CH, + 2 0 2 C0 2 + 2 H 2 0

AH = -212 kcal

(6)

C 2 H 5 OH + 3 0 2 2 C0 2 + 3 H s O

AH = -327 kcal

(B) The reaction of a carbonate, C0 8 2 ~, or a hydrogen carbonate HC0 3 ",
with a strong acid, H + , is a laboratory method of preparing C0 2

(7) C0 3 2 - + 2 H+ C0 2 + H 2 0 (CaCOs + 2 HC1-*C0 2 + H a O + CaCl 2 )

(8) HCO s - + H+ C0 2 + H 2 0 (NaHCOs + HC1 C0 2 + H s O + NaCl)

(9) CaC0 3 -> C0 2 + CaO (except NaHCOs and K 2 CO s )

(10) 2 NaHCOs C0 2 + Na 2 COs

(D) The fermentation of sugar (glucose) for the production of alcohol
is a commercial source of C0 2 .

(U) C*Kl 2 Ob 2 C0 2 + 2 C 2 H 5 OH (ethyl alcohol)

6, Properties of Carbon Dioxide. Physical properties : These are given
in Table STB.

The Elements of Group IVB: Carbon

411

Chemical properties : Carbon dioxide is a very stable compound; at 2000°C
it is only 1.8% dissociated into its elements. It is reduced by red-hot carbon
to CO, and by active metals at high temperature to free carbon.

(12) C0 2 -j- C — > 2 CO

(13) C0 2 + 2 M g 2 C + 2 MgO (with burning Mg; also K, Na, and Zn)

Since C0 2 is formed by the complete oxidation of carbon, it does not bum,
nor does it support combustion. As little as 2.5% C0 2 in air will extinguish
a burning candle. Carbon dioxide dissolves in water to form carbonic acid,
H 2 C0 3 . The acid has never been isolated and is known only in aqueous solu-
tion. Its structure is presumably

0=C

O-H

T>— H

*•

Carbonic acid is a weak diprotic acid.

(14) H 2 C0 3 ^ H+ + HCCV

(15) HCO a - H+ + C0 3 2 -

Ki =. 4.3 X 10- 7
K 2 = 4.7 X 10- 11

As written in Equation 14 the primary ionization infers that all the dis-
solved C0 2 has reacted with H 2 0 to form H 2 CO s . Only a small fraction
of the CO a , however, is in the acid form so that the primary ionization is
more properly written as

(16) C0 2 + HoO -» H+ + HCO s - K x = 4.3 X Kb 7

The situation is analogous to that of NH 3 and NH±OH. The equilibria in an
aqueous solution of C0 2 can be summarized by

(16) C0 2 + H 2 0 H 2 CO s H+ + HC 0 3 t H+ + COa 2 *

7. Carbonates and Hydrogen Carbonates. Two series of salts can be
formed by H 2 C0 3 , namely, the normal carbonates, e.g., Na 2 COa, and the
hydrogen carbonates or bicarbonates, e.g., NaHC0 3 . When CC 2 reacts with
excess NaOH, Na 2 C0 3 is formed; when an excess of CO g is used, NaHCO s is
produced.

(17) CO a + 2 OH“ CO s 2 - + H 2 0 (C0 2 + 2 NaOH -» Na 2 C0 3 + H 2 0)

(18) C0 2 + OH- HCOr (CO, + NaOH NaHCO s )

Also, if HC0 3 “ ion further reacts With OH" ion, the CO3 2 " ion is formed,
whereas reaction of the CO,! 2 - ion with C0 2 yields the HCOs“ ion.

(19) HCOa” + OH--> C0 3 2 - + HoO (NaHCOs + NaOH Na 2 C0 3 + H 2 0)

(20) CO s 2 - + C0 2 + H 2 0 2 HCOr (Na a CO» + CO, + H 2 0 2 NaHCOs)

Both the carbonates and the hydrogen carbonate salts of the alkali metals
are soluble in water, the carbonates being the more soluble. Other metal
carbonates, however, are insoluble whereas their hydrogen carbonates are
generally more soluble. The calcium compounds offer an interesting example:
CaC0 3 is insoluble but Ca(HCO ; *) 2 is soluble. When C0 2 is passed into a

412

The Elements of Group 7VB; Carbon

solution of Ca(OH) 2 there is first precipitated CaC0 3 which then redissolves
as more C0 2 is added due to the formation of Ca(HC0 3 ) 2 .

The weakness of the secondary ionization of H 2 C0 3 indicates that the
CO3 2 - ion readily combines with protons. Hence the hydrolysis of C0 3 2 " ion
yields a strongly basic solution; the pH of 1 M Na 2 C0 3 is approximately 12.

(21) CO3 2 - + H 2 0 -» HCXV + OH-

Solutions of HCO3- ion are also basic by hydrolysis; the pH of 1M NaHC0 3
is about 8.3. The HC0 3 ' ion can act both as an acid and as a base, but
its combination with protons proceeds to a greater extent than does its
production of protons.

(22) H 2 C0 3 + OH- H2O + HCO3- + H 2 0 ;=± H 3 0+ + CO s 2 -

Because the HC0 3 ~ can both accept and yield protons, two such ions can
interact.

(23) HCOr + HCO3- H 2 C0 3 + C0 3 2 - K = 1.1 X Kb 4

It is due to this equilibrium that C0 3 2 ~ ion is converted to HC0 3 “ ion by
the addition of CO z . The carbonic acid-hydrogen carbonate equilibria are im-
portant factors in maintaining the pH of the blood.

The CO s 2- ion has a planar structure with the three oxygen atoms in
triangular symmetry about the central carbon atom. For the carbon to have
four covalent bonds there must be a double bond between it and one of
the oxygen atoms. The X-ray analysis of CaC0 3 , however, indicates that
the three carbon-oxygen bonds are identical. One pair of shared electrons
is not located between a specific C— O bond but belongs to the CO s 2 “ ion
as a whole, giving some partial double bond character to each bond, so
that the structure of the ion is best depicted as a resonance hybrid of
three forms.

••

:o-c

dm 1

8. Uses of Carbon Dioxide. The principal use of carbon dioxide is in
the production of carbonated beverages. The solubility of C0 2 in H 2 0 is
increased by pressure “Soda water” is an aqueous solution of C0 2 under
a pressure of 3 to 4 atmospheres; under reduced pressure the C0 2 escapes
to produce the familiar bubbles. Because C0 2 does not support combustion
it finds use as a fire extinguisher. One such type is shown in Figure 31.3.
Another type of fire extinguisher is a tank filled with liquid C0 2 under
pressure; when its valve is opened, a stream of gaseous and solid C0 2 is
emitted through a funnel shaped nozzle. This extinguisher leaves no
liquid residue which might damage any materials not affected by fire.

9. Carbon Monoxide, Carbon monoxide is also a colorless, odorless, and
tasteless gas but, unlike C0 2 , it is poisonous. When inhaled, it combines with
the hemoglobin of the red blood cells to form .a relatively stable compound.

"'O'.

:o-c:

The Elements of Group TVB: Carbon

413

carboxy-hemoglobin . The hemoglobin is thereby rendered unable to perform
its normal function, namely, to combine with oxygen, to transport it through-
out the body, and to release it where required. The oxygen supply to the
body becomes inadequate and death occurs when about one-third of the
hemoglobin has entered into such combination. One volume of carbon
monoxide in 10,000 of air produces symptoms of poisoning while one in
800 can produce death in thirty minutes. Because of its lack of odor, carbon
monoxide is an insidious poison. Due to incomplete combustion the exhaust
gases from an internal combustion engine contains 1 to 8% CO. This concen-
tration is sufficient to make the air of a small closed garage deadly in
just a few minutes. Canaries are especially susceptible to carbon monoxide
poisoning and are often used in mines to indicate dangerous concentrations
of the gas, although instrumental techniques have been developed to deter-
mine the concentration of carbon monoxide in the air.

In its physical properties carbon monoxide bears a striking resemblance
to nitrogen. Both gases have the same molecular weight and the same
number of valence electrons. This might imply a triple bond structure for
CO but its electronic structure is best formulated as a resonance hybrid of
three structures, one of which contains a triple bond like N 2 .

:C=o: :C=o: :C-0:

••

10. Preparation of Carbon Monoxide. Carbon monoxide can be prepared
in several ways.

(A) The combustion of carbon or its compounds in a limited supply
of oxygen

(24) 2 C + 0 2 2 CO AH = -52,8 kcal

If an is used, a combustible mixture of CO (30-40%) and N 2 is obtained.
This mixture is known as producer gas and is used as an industrial fuel.
About 500° C, the combustion of carbon, even in a limited supply of air,
yields almost pure CO a . But at 1000°C, the product of such a reaction with
excess carbon is almost solely CO.

(B) The reduction of carbon dioxide by carbon at high temperatures

(25) C0 2 + C 2 CO

When a bed of coal burns in a furnace, C0 2 is formed where the entering
air first reacts with the carbon. As the C0 2 rises through the upper layer
of red-hot coal it is reduced to CO. These reactions are significant in the
production of iron in the blast furnace.

(26) h 2 0 + C CO + Ha AH = -29.0 kcal

The mixture of CO and H 2 , called water gas, contains about 50% CO and is
also used extensively as a fuel. Water gas is also a source of hydrogen.

(D) The decomposition of oxalic acid, (COOH) 2 , or formic acid, HCOOH,
in the presence of sulfuric acid as a dehydrating agent

414

The Elements of Group IVB: Carbon

(27) (COOH) 2 CO + C0 2 + H 2 0 (a laboratory preparation)

(28) HOOOH CO + H 2 0

Although CO appears to be anhydride of HCOOH, Equation 28 is not
reversible.

Figure 31.3. A Type of C0 2 Fire Extinguisher.
H 2 S0 4 and a solution of NaHCO s are in separate
compartments. When the extinguisher is inverted
the two liquids mix. The C0 2 generated develops
an internal pressure which forces a stream of
liquid out of the nozzle in the direction of the fire.

11. Properties of Carbon Monoxide. Physical properties ; These are given
in Table 31-B.

Chemical properties: Since the oxidation state of carbon in carbon mon-
oxide is +2, carbon monoxide reacts to form the more stable +4 state.

(A) Carbon monoxide burns in oxygen or air to form carbon dioxide.
The energy liberated in this reaction is the basis of the fuel value of CO.

(29) 2 CO + 0 2 -> C0 2 AH = -135.2 kcal

‘(B) Carbon monoxide combines with chlorine in the presence of sun-
light or a charcoal catalyst.

(30) CO + Cl 2 — » COCl 2 (carbonyl chloride or phosgene)

Phosgene (which contains no phosphorus) is a poisonous gas and was used
as a chemical warfare agent in World War I.

(C) Carbon monoxide is a reducing agent; its reduction of metal oxides
at high temperatures is an important step in many metallurgical operations.

(31) Fe 2 0* + 3 CO 2 Fe + 3 CO,

(D) Carbon monoxide combines directly with many metals, particularly
with those those ol Group VIII of the Periodic System, forming compounds

The Elements of Group IVB: Carbon

415

called metal carbonyls. Examples are nickei carbonyl, Ni(CO) 4 , and iron
carbonyl, Fe(CO) 5 .

12. Other Compounds of Carbon. Carbon disulfide, CS 2 , is made by
the direct combination of carbon and sulfur in an electric furnace.

(32) C + 2 S CS 2 AH = +22 kcal

Carbon disulfide is a colorless, highly volatile and flammable liquid, which
freezes at — 108°C and boils at 46°C. It has an unusually high index of
refraction (1.63). Because its heat of formation is endothermic, CS 2 is
unstable and undergoes slow decomposition, upon standing; it turns yellow
and its not unpleasant odor becomes disagreeable due to the formation
of other sulfur compounds. Though insoluble in water, CS 2 is an excellent
solvent for nonmetallic elements such as S 8 , Se 8 , P 4 , Br 2 , and I 2 , and for
many organic substances such as waxes, fats, oils, resins, and rubber. Carbon
disulfide is used as an insecticide and in the manufacture of rayon, cello-
phane, and carbon tetrachloride.

Carbon tetrachloride, CC1 4 , is made by passing chlorine into carbon
disulfide using iodine as a catalyst.

(33) * CS 2 + 3 Cl* CC1 4 + S 2 CL

The CC1 4 is separated from the higher boiling S 2 C1 2 by fractional distillation.
Carbon tetrachloride is a colorless, volatile liquid, freezing at -23 °C and
boiling at 78°C. It is a good solvent for grease and fat and consequently
is used in dry cleaning. Because CC1 4 is noncombustible and because its
vapor is five times as dense as air, it is used as a fire extinguisher. It is
most effective in fighting oil fires for the liquid mixes with the oil instead
of sinking to the bottom as would water. When using a CC1 4 extinguisher,
care must be taken not to inhale any of the vapors; the vapor of CCl 4 itself
is slightly toxic, and COCl 2 is often formed when CC1 4 is heated in air to
a high temperature.

At very high temperatures, carbon combines with many metals and
nonmetals to form carbides. The preparations of calcium carbide, CaC 2 ;
and of silicon carbide, SiC, by heating a mixture of the oxide and coke in
an electric furnace, are typical.

(34) CaO + 3 C CaC 2 + CO (at about 3500°C)

(35) Si0 2 + 3 C — > SiC + 2 CO

Most carbides form splendidly shaped and colored crystals; SiC crystallizes
in purple or black, iridescent, hexagonal plates. The carbides are brittle but
very hard; SiC is almost as hard as diamond and is sold under the trade
name of “Carborundum ” It, and the carbides of boron, B 4 C, and tungsten,
WC, are used as abrasives. The carbides of the more active metals react with
water to form hydrocarbons; other carbides are extremely unreactive. Com-
mercially the most important carbide is CaC 2 . It reacts with water to yield
acetylene, C 2 H 2 , a gaseous compound used as a fuel and also as an inter-
mediate in the synthesis of many organic, compounds.

(36) ' CaC 2 + 2 H 2 0 C 2 H 2 + Ca(OH) 2

416

The Elements of Group 1VB • Carbon

QUESTIONS

1. For the elements of Group 1VB, account for the gradation of (a) metallic
character (b) acidity or basicity of the oxides (c) stability of the hydrogen
compounds.

2. Compare the properties of the allotropic forms of carbon. Account for the
high melting points of both graphite and diamond. What properties might be
predicted from their crystal structures?

3. Why does graphite exhibit resonance forms, whereas diamond does not?

4. How can synthetic diamonds be made? How is graphite manufactured? What
are the uses of diamond and graphite?

5. Distinguish among peat, lignite, anthracite coal, bituminous coal, boneblack,
lampblack, black lead, coke, charcoal.

6. Why does solid C0 2 sublime, and not melt, at one atmosphere pressure?

7. How are the following made (a) calcium carbide (b) silicon carbide (c) carbon
disulfide (d) carbon tetrachloride? List properties and uses of each.

8. How is carbon dioxide, gas and solid, produced commerically? State their
uses. How would you make liquid carbon dioxide?

9. Draw resonance forms of (a) carbon dioxide (b) the carbonate ion (c) car-
bon monoxide.

10. Explain why any strong acid reacts with any carbonate to form carbon dioxide.

11. Why are aqueous solutions of the alkali carbonates and hydrogen carbonates
basic?

12. Formulate the various equilibria that exist in carbonated water.

13. Write equations for the reactions which occur when carbon dioxide is
bubbled into a solution of lime water until the solution again becomes clear.

14. Why does soda water remain quiescent in a closed bottle but bubble when the
bottle is opened?

15. What principles are involved in extinguishing a fire? Explain the operation
of the “bicarbonate-acid” extinguisher and the “pyrene” extinguisher. How
could burning magnesium be extinguished?

16. For carbon dioxide and for carbon monoxide compare the (a) chemical prop-
erties (b) physiological properties.

17. What is producer gas? water gas? How is each made? How could you
separate the components of water gas? Why might producer gas be preferred
as a fuel to coal?

18. Complete and balance the following equations; where applicable, write ionic
equations also.

C 3 H s (propane) 4- 0 2
(C 2 H 5 ) 2 0 (dimethyl ether) 4- 0 2

khco 3 + h 2 so 4

NaHCO s + H 3 P0 4
C0 2 + KOH

CaC 2 + H a O

CS 2 4- 0 2

FeO 4- CO

CO 4* Cl 2

Ni 4* CO

The Elements of Group IV B: Carbon

417

19. What weight of NaHC0 3 is needed to produce 37 liters of dry C0 2 at

50 °C and 760 mm when treated with excess acid? Am : 123 g

20. A gaseous fuel has the following composition by volume: 30% H 2 , 40% CO,
25% N 2 and 5% C0 2 . What volume of air, containing 21% 0 2 by volume,
is required for complete combustion of 5000 cubic feet of the fuel gas?

Ans ; 8333 ft®

21. What volume of CO z , measured at 20° C and 780 mm, is required to convert

10.0 g of Na 2 C0 3 to NaHCO s ? Am: 2.25 liters

22. What volume of air at STP containing a normal amount of C0 2 must be passed
through a solution of Ba(OH) 2 to precipitate 5.0 g of BaC0 3 ? Am: 1900 liters

23. What volume of CO, measured at STP, must be burned to yield 200 kcal?

Ans: 33.2 liters

24. From information in this chapter, calculate AH for Equation 25.

25. Calculate the equilibrium constant for the following reaction:

HCO s - + H 2 0 H 2 C0 8 + OH-

Organic Chemistry I
The Hydrocarbons

The term organic chemistry dates back to the late eighteenth century
when chemists, still under the influence of alchemy and its philosophy,
thought that certain compounds of carbon and hydrogen, and others derived
from them, could be produced only in nature by living organisms, and that
it was impossible to make them from inorganic or mineral substances. But
in 1828 the German chemist, Friedrich Wohler, converted an inorganic salt,
ammonium cyanate, NH 4 OCN, a compound not produced by either plant
or animal, into the compound urea, CO(NH 2 ) 2 . Until then urea, which Js
found in the urine and is formed by the oxidation of nitrogenous material
during the digestive processes in animals, had been considered capable of
being produced only by living organisms. Wohler’s synthesis, however, pro-
duced an organic compound from an inorganic substance that was not as-
sociated in any way with the life processes of either plants or animals. Though
urea and ammonium cyanate contain the same number of carbon, oxygen,
nitrogen, and hydrogen atoms in a molecule, they are unlike in that they
have different arrangements of these atoms within the molecule.

The transformation of ammonium isocyanate into urea can be expressed by:

H H

(1) [NH 4 ] + [0-C=N]‘ -» H-N-C-N-H

This chemical change is an example of the type known as internal re--
arrangement, and again illustrates the importance of structure in determining
the nature of a molecule. Internal rearrangements are quite common in
organic chemistry. They involve no change in the number or kind of atoms
within a molecule; the reaction consists solely of an alteration in the dis-
position of the atoms in a molecule. Wohler’s discovery sounded the death
knell of the vital foxce theory, which held that comnounds of carbon could

Organic Chemistry I; The Hydrocarbons

419

be formed only through the influence of the living or vital force of a bio-
logical organism. Organic chemistry is now defined as the study of the
compounds of carbon. Inorganic chemistry deals with the compounds of
all elements other than carbon. 1 Since Wohlers time, many thousands of
organic compounds, including many not found in nature, have been syn-
thesized in the laboratory. More than half a million organic compounds
are known, far more than the total of inorganic compounds (about 50,000),
and the number possible is theoretically without limit.

1. The Hybrid Tetrahedral Bond. The carbon atom has four valence
electrons; by sharing electrons to form four covalent bonds it attains a noble
gas configuration. The electron configuration of an individual carbon atom
is Is 2 , 2s 2 , 2 px 1 , 2p y 1 , 2p z °. It would appear that, in compound formation,
the carbon atom should form bonds solely with its two unpaired p electrons
or perhaps use its 2 s 2 electrons in addition to make possible a total of four
bonds. In the latter case, bonds formed with the s electrons would differ
in properties from those formed with the p electrons. Experimentally this
is found to be not so. The four covalent bonds that a carbon atom forms
are identical in all respects. Thus methane, CH 4 , is a symmetrical tetrahedral
molecule, with a carbon atom at the center of the tetrahedral body and a
hydrogen atom at each of the four comers of the tetrahedron (Figure 32.1 a).
The bond distances (1.09 A), the bond angles (109.5°), and the bond
energies (101 kcal/mole) are all the same.

The equivalence of die four carbon bonds is due first to the promotion,
during reaction, of one 2s electron to the unoccupied 2p orbital, giving
the configuration Is 2 , 2s 1 , 2 p x \ 2 p y \ 2 p z l . These separate orbitals then com-
bine, or hybridize, to form four composite but equivalent orbitals, each
of which is known as a sp* hybrid orbital , because it was formed through
the combination of one s and three p orbitals. In effect the electron clouds
of the separate s and p orbitals interact to form a new pattern of electron
field intensity, the sp 8 orbital. Bonds formed by hybrid orbitals (through
overlap with other orbitals) are stronger than those formed by the pure
orbitals involved in the formation of the hybrid orbital and they, too, have
spatial orientation. The sp s bonds are directed towards the comers of a
regular tetrahedron.

The nomenclature sp s represents an orbital which has 1/4 s character and
3/4 p character (where no superscript- is written the number 1 is presumed).
The sum of the superscripts indicates the number of electrons involved
in bonding and also the total number of such bonds that can be formed;
thus only four sp s bonds can be formed.

In diamond (page 406) the carbon-carbon bonds are sp 3 bonds. In the case
of graphite (page 406), the bonds between carbon atoms within a hexagon are
formed by the overlap of sp 2 hybrid orbitals. Such orbitals lie in the same plane
and are directed at an angle of 120° from each other; for this reason the carbon
atoms in graphite form coplanar hexagons. Only three of the four carbon
electrons are used in forming the sp 2 orbitals; one p orbital in each carbon

1 Also included in inorganic chemistry are most of the simple carbon compounds studied

in the last chapter, such as CO, C0 2 , the metal carbonates, etc.

420

Organic Chemistry I; The Hydrocarbons

atom remains. Its electron field is at right angles to the hexagonal plane,
one lobe projecting above and one below this plane. A small overlapping
between these p orbitals acts as the weak bonding force between adjacent
planes. Other types of hybrid orbitals will be taken up as they become
necessary.

2. Glasses of Organic Compounds. The principal reason for the great
multiplicity of carbon compounds is the ability of the carbon atom to combine
with itself, that is, to form covalent bonds not only to atoms of other
elements but also to other atoms of carbon. Thus the carbon atom can
combine with four hydrogen atoms, as in methane, CH 4 , or the carbon

(a)

Tetrahedral
structure of CH 4

Electron cloud
density of an
sp 3 orbital

(b)

Orientation of
sp 3 orbitals

Is orbital
of hydrogen

sp 3 orbital
of carbon

Formation of CH 4 by the over-
lap of the carbon sp s orbitals
with the hydrogen Is orbitals.

Figure 32.1. Orbitals of the Carbon Atom.

atom can share its electrons with three hydrogen atoms and another carbon
atom, which in turn can be bound to other carbon atoms, as shown below.

H H

1 |

H H H H

14 11

H-C-H

1 1

H-C-C-H

i i

111

h-c-c-c-c-h

1 1

H H

u u

Methane, CH*

Ethane, C 2 H«

Butane, CJHio

No other atom has so marked an ability to combine with others of its own
species as does the carbon atom; only a few, notably silicon, boron, and
nitrogen, have even a limited ability to do so.

Ite carbon-to-carbon linkages are not limited to “straight chains” (see
Section 7). Branched chains and closed chains or cyclic structures can be
formed as shown by the following skeleton diagrams. For the sake of clarity
and brevity in depicting the carbon chain these diagrams are incomplete;
in a complete structural formula each carbon atom must have a total of four
valence bonds.

Organic Chemistry I: The Hydrocarbons

421

-C-C-C-C-C-

-C-C-C-C-
1 |

c-c-c

1 1
c c

C-C-A

“Straight” chain

Branched chain

Closed chain

or ring

On the basis of their structure organic compounds are divided into two
broad classes, aliphatic compounds and aromatic compounds. The aliphatic
class includes those in which the carbon atoms form open chain compounds
and certain cyclic compounds whose structures are related to the open chain
compounds. The aromatic compounds are benzene and those which resemble
benzene in their chemical behavior. The original meanings of the words
aliphatic (fatty) and aromatic (fragrant) are now without significance. We
shall take up the aliphatic compounds first.

Aliphatic Compounds

3. The Alkane Series of Hydrocarbons. The hydrocarbons, or compounds
of carbon and hydrogen only, are the simplest organic compounds. Methane,
CH 4 , is the first member of a series of compounds known as the alkane
series or the paraffin series of hydrocarbons. In Table 32-A are listed some
properties of a few members of this group.

Table 32- A

Aucanes

Formula

Name

State

at20°C

Melting
Faint , °C

Boiling

Point, °C

Density*

g/ml

Number oj
Isomers

ch 4

Methane

gas

-183

-162

0.415

c 2 h„

Ethane

gas

-172

-88

0.561

C a H s

Propane

gas

-187

-42

0.585

C,H 10

n-Butane

gas

-138

0.60

c,h 12

n-Pentane

liquid

-130

C„H U

n-Hexane

liquid

-95

0.659

C,H le

n-Heptane

liquid

"-90

0.884

c„h 18

n-Octane

liquid

-57

126

0.703

c*h 20

^-Nonane

liquid

-54

151

0.718

^10^22

tt-Decane

liquid

-30

174

0.730

^18^88

/i-Octadecane

solid

308

0.785

60523

c 20 h 42

/i-Eicosane

solid

0.778

! 366319

Isobutane

gas

-159

-12

0.60

c 5 h 12

Isopentane

liquid

-160

0.620

C 5 Hj 2

Neopentane

gas

-17

9.5

0.610

♦Tie density is given at 20 °C except for those substances which are gases; for these
the density is that of the liquid at its boiling point.

422

Organic Chemistry I ; The Hydrocarbons

A general formula, expressing the composition of the members of the
alkane series, is C„H 2w . 2 , in which n is an integer. The composition of any
one member of this series differs from that of the preceding or the following
member by one carbon and two hydrogen atoms, or CH 2 . Such a series
of compounds showing a constant differential in formula between successive
members is known as a homologous series. Table 32-A is evidence for the
generalization that the members of a homologous series show a regular
progression in properties. Except for the first few members of the series
the boiling point increases 20 to 30 degrees for each carbon atom added
to the hydrocarbon chain; the melting point and density also increase but
not quite so regularly.

All the bonds in an alkane molecule are single covalent bonds. The
carbon-carbon bond energy is 80.5 kcal/mole and the carbon-hydrogen bond
energy is 101 kcal/mole. Because the carbon and hydrogen atoms do not
differ much in electronegativity and because of their symmetry the alkane
molecules are nonpolar. They are insoluble in water but generally dissolve
in other organic solvents; all are less dense than water. The alkanes occur
in petroleum and its products, such as natural gasoline, kerosene, lubricating
and fuel oils, and paraffin.

4. Isomerism of Carbon Compounds. The use of structural formulas in
organic chemistry is almost compulsory for the molecular formula alone is
insufficient to distinguish between certain organic compounds. For example
there are three different compounds which have the same molecular formula,
C 5 Hi 2 . Though the molecule of each contains five carbon and twelve hydrogen
atoms, nevertheless each has different chemical and physical properties. The
difference between the molecules of these compounds is in the arrangement
of the carbon and hydrogen atoms in the molecules as shown by the following
structural formulas.

H H H

H H H H H

H H H H

H '6' K

1 t 1 1 1

Mil

H-C-C-C-C-H

\ 1 /
h-c-c-c-h

III

H H H H H

H A H H

/ 1 \

W ,C V \h

/|\

/i\

H H H

Pentane

Isopentane

Neopentane

The existence of two or more compounds having the same molecular
formula but possessing different properties because of different molecular
structures is called isomerism ; the compounds are said to be isomers of each
other. The greater the number of carbon atoms in a molecule, the larger is the
number of isomers possible. This is indicated in the last column of Table 32-A;
for a compound having the formula C 4 oH 8 2 the number of isomers theoretically
possible is 69,491,178,805,831.

Organic Chemistry I: The Hydrocarbons

423

5. Nomenclature. Even with the smaller numbers of compounds in
inorganic chemistry it is not uncommon to find that the same compound
might be known by different names. With so vast a number of organic
compounds and the complication introduced by isomerism, some system of
distinction among them is imperative. Accordingly there are several systems
of nomenclature for deriving the name of a compound from its formula.

(A) The system or rules for naming the hydrocarbons, adopted by the
Council of the International Union of Pure and Applied Chemistry (IUPAC),
is in part as follows: 2

(l)The name of a compound is derived from the longest continuous carbon
chain in the compound. The prefix of such name specifies the number of
carbon atoms in the chain. For compounds containing up to four carbon
atoms in the chain, the prefixes are historical and have become established
through long usage; beginning with a five carbon atom chain, Greek or
Latin numerical prefixes are used.

Prefix

Number of Carbon atoms

Prefix

Number of Carbon atoms

meth-

one

hex-

six

eth-

two

hept-

seven

prop-

three

oct-

eight

but-

four

non-

nine

pent-

five

dec-

ten

For the alkane series, the suffix-ane is added to these prefixes. In other
series of compounds we shall encounter, different suffixes are used.

(2) A branched chain hydrocarbon is regarded as a derivative of the
hydrocarbon having the longest continuous chain in the compound.

(3) The carbon atoms in the longest chain are numbered in sequence,
1, 2, 3, 4, etc.

(4) The branched or substituted groups are named, giving the location
of the group along the chain by indicating the number of the carbon atom
to which the substituted group is attached. In this connection, the initial
numbering of the carbon atoms is started from the end of the chain which
will result in the smallest numbers for the substituted groups in the name of
the compound. The name of a branched group is also based on the foregoing
prefixes in indicating its carbon chain length, but the suffix -yl is added.
This suffix denotes a group or “radical," not a complete compound. Such
groups derived from the alkane series are called alkyl radicals and the chemist
often represents them by the symbol R in structural formulas. Thus an alkyl
radical is an alkane hydrocarbon less one hydrogen atom so that the radical
has one unsatisfied bond or an unpaired electron; the CH 3 group is the
methyl radical, the C 2 H 5 group the ethyl , etc. Groups obtained by dropping
a hydrogen atom from an aromatic hydrocarbon are known as aryl groups.

2 The complete rules for naming organic compounds are tabulated in the “Handbook
of Chemistry and Physics” (Chemical Rubber Publishing Co.), a useful reference book
with which the student would be wise to become familiar.

424

Organic Chemistry I: The Hydrocarbons

As an example of this system of nomenclature let us name the alkane
with the following carbon skeleton:

A c

C-C-C-C-C-C-C-C
1234567 8

This hydrocarbon contains 12 carbon atoms but its longest continuous chain
(without backtracking) has only 8 carbon atoms so that the compound is
considered to be an octane. The carbon atoms of this chain have been
numbered from the left since this will result in smaller numbers for the
attached groups. An ethyl radical (two carbon chain) is attached to the third
carbon atom and two methyl groups (one carbon chain) to the fifth carbon.
Accordingly, the name of the compound is 3-ethyl-5,5-dimethyloctane . The
numeral 5 is repeated to indicate that there are two substituted groups on
the fifth carbon atom; the prefix di denotes that they are both the same,
in this case, methyl groups. By this system the three isomers of pentane
on page 422 would be named pentane, 2-methylbutane, and 2,2 dimethyl-
propane. This method of nomenclature is also the basis for naming com-
pounds other than the alkane hydrocarbons; additional modifications will be
introduced as they are required in our study of organic chemistry.

(B) Another system of nomenclature distinguishes between normal and
iso- compounds. Normal compounds are those in which all the carbon
atoms are in a continuous chain. Thus C— C—C—C— C is a normal hydro-
carbon, normal pentane, written n- pentane; the prefix n- indicates a normal
compound. Such distinction is unnecessary in the IUPAC system because
there can be only one normal alkane. "Pentane” means a continuous chain
configuration with the formula C 5 H 12 ; any isomer would have a different
root name. The iso- prefix refers to a branched chain compound, this prefix
is applied only to the isomer which has a methyl radical attached to the
carbon atom one removed from the end of the carbon chain. Of the two
isomers of n- pentane only one can be called isopentane; neopentane is
neither a normal nor an iso- compound and hence cannot be named by
this system.

(C) A carbon atom may also be classified on the basis of the number of
other carbon atoms to which it is attached. A primary carbon atom is one
bonded to only one other carbon atom; a Secondary carbon atom to two
other carbon atoms; and a tertiary carbon atom to three other carbon atoms.
In isopentane, carbon atoms 1 and 4 at the ends of the chain are primary
carbon atoms; carbon atom 8 is a secondary carbon atom; and carbon atom 2
is a tertiary carbon atom.

6. Reactions of the Alkane Hydrocarbons. The alkane or paraffin ("having
little affinity”) hydrocarbons are characterized by their chemical inertness.

Organic Chemistry I: The Hydrocarbons

425

They are unaffected by boiling with concentrated sulfuric acid or strong
bases, and are but slightly acted on by oxidizing agents such as concentrated
nitric acid.

(A) Reactions with the halogens (halogenation). Chlorine and bromine
react with the alkanes, replacing part or all of the hydrogen atoms by
an equal number of halogen atoms. Such a reaction is termed substitution.
Fluorine reacts explosively but iodine does not react directly with the
alkanes. The four hydrogen atoms of methane can thus be replaced successively
by chlorine atoms to form the compounds, CH 3 C1, CH 2 C1 2 , CHCla,
and CC1 4 . Experimentally the reaction is difficult to control or to stop at
any particular stage of substitution so that a mixture of products results.
It is generally characteristic of organic reactions that no single reaction
occurs; parallel reactions take place simultaneously and no one product is
formed to the exclusion of others. With longer carbon chains substitution
by a halogen can result in isomers depending upon the positions taken by
the halogen atom(s). Thus several isomers of a compound can be formed
and it is frequently difficult to separate them.

(B) Reaction with oxygen. The complete combustion of a hydrocarbon
with oxygen yields carbon dioxide and water. Because a large amount of
heat energy is evolved in such a reaction the hydrocarbons are used as
fuels. The combustions of CH 4 and C S H 18 are typical.

(2) CH 4 + 2 O z C0 2 + 2 H 2 0 AH = -213kcal

(3) 2 C 8 H 18 + 25 0 2 -> 16 C0 2 + 18 H 2 0 AH = -2604 kcal

(C) Pyrolysis (“cleavage by heat”). When the alkane hydrocarbons are
heated in the absence of air to a high temperature (400-800°C), pyrolysis
occurs. A number of different reactions can take place depending upon the
hydrocarbon itself, temperature, pressure, and the presence of a catalyst.
Generally the longer hydrocarbon chains are broken into smaller ones and
hydrogen is one of the products of the reaction. In the pyrolysis of propane
at 700° C two reactions occur to about an equal extent.

(4) C 3 H 8 CH 4 + C 2 H 4 (ethylene)

(5) C 3 H 8 -» H 2 — {- C 3 H 8 (propylene)

Ethylene and propylene belong to a series of hydrocarbons known as the
alkenes.

In the petroleum industry the pyrolysis of long chain alkanes is known
as “cracking.” This breaks down the larger molecules of the kerosene and
gas oil fractions of petroleum into more volatile shorter chain compounds,
such as octane, which can be used as gasoline.

7. Structure of the Alkane Hydrocarbons. The structure of the CH 4 mole-
cule is that of a regular tetrahedron. Whereas the structures of the normal
hydrocarbons from propane on are written in the text with the carbbn atoms
in a straight line, in reality these atoms form a zigzag structure as demanded
by the continuance of the tetrahedral pattern throughout the molecule. The
carbon-carbon distance of 1.54 A is the same as that in the diamond

426

Organic Chemistry I* The Hydrocarbons

Figure 32.2. The Propane Molecule,

Only the carbon atom skeleton and one hydrogen
atom are shown to illustrate bond distances and the
bond angles of the molecule.

CH CjJK* C 3 H 8

Figure 32.3. Structures of Alkane Molecules.

CH 4 c 2 h 6

Figure 32.4. Molecular Models.

End View End View

(a) (b)

(a) Eclipsed Confirmation (b) Staggered Conformation.

Fiaure 32.5. Confnrrn/rHnrn* nf th*> Ffh/m« Mrtlamila

Organic Chemistry 1: The Hydrocarbons

427

structure and the bonding orbitals contributed by the carbon atoms are
all of the sp 3 type (Figure 32.2). The drawings in Figure 32.3 emphasize
bond angles, but, in reality, the proximity of the bonded atoms is such that
the molecular models in Figure 32.4 are more realistic. Molecules do not
have rigid conformations as the illustrations in Figures 32.3 and 32.4 might
imply. In the alkanes there can be free rotation about the single bond “axis”
joining two carbon atoms. Two aspects of this rotation for the ethane
molecule are shown in Figure 32.5. Such different arrangements of atoms
within a molecule which can be achieved without the breaking of bonds
are called conformations. In the eclipsed conformation the hydrogen atoms
are crowded closer together than they are in the staggered conformation.
There arises a repulsion, known as nonbonding interaction, which makes
the potential energy of the eclipsed conformation larger than that of the
staggered conformation; for ethane this is about 3 kcal/mole. With larger
molecules the possible conformations often have a bearing upon their
chemical properties but these subtle differences in structure are better left
to courses in organic chemistry.

8. The Alkene Series of Hydrocarbons. Another homologous series of
hydrocarbons has the general formula C n H 2n - It is known variously as the
alkene series, the olefin series, or the ethylene series from its first member,
ethylene, C 2 H 4 . The compounds in this series contain two hydrogen atoms
less than the corresponding member of the alkane series. Hypothetically
we may consider them to be formed from an alkane by the removal of
two hydrogen atoms from adjacent carbon atoms. This removal of a hydro-
gen atom would leave an unpaired electron on each carbon atom. The
two unpaired electrons couple to form an additional covalent bond between
the carbon atoms; the result is a double covalent bond between the carbon
atoms. This double bond is characteristic of alkene molecules and denotes
that two pairs of electrons are shared by two adjacent carbon atoms.

H H
•# •«

H-C-C-H

I I

H H
Ethane

minus

— two H — »
atoms

hJJ-h _

Ethylene

The schematic representation above for the preparation of an alkene hydro-
carbon is purely for illustrative purposes and cannot be realized in practice.
Ethylene can be prepared by eliminating the elements of water from a
molecule of ethyl alcohol, C 2 H 5 OH, using concentrated sulfuric acid as the
dehydrating agent.

H H

H-C-C-O-H

Ethyl alcohol

h 2 so 4

H H

H— C=C— H + H 2 0

Ethylene

428

Organic Chemistry l; The Hydrocarbons

The main industrial source of alkene hydrocarbons, however, is the cracking
of long chain alkane compounds found in petroleum.

Alkene hydrocarbons are named after the corresponding alkanes by
replacing the ending ~ane by -ene. The location of the double bond is
specified by a numerical prefix which indicates the number of the carbon
atom preceding the double bond, as shown below.

C—C = C-C-C C~C = C-C-C

2-pentene 3-methyl-2-pentene

The properties of some members of the alkene series are listed in Table 32-B.

Table 32-B

Alkenes

Formula

Name

State

at 20° C

Melting

Point, °C

Boiling

Point, °C

Density *
g/ml-

c 2 h 4

Ethylene (ethene)

gas

-169

-102

0.566

C,H e

Propylene (propene)

gas

-185

-48

c,h 8

1-Butene

gas

-130

- 6.5

0.668

c 5 h 10

1-Pentene

liquid

0.643

c«h 12

1-Hexene

liquid

-138

63.5

0.675

c 7 h 14

1-Heptene

liquid

-119

0.698

c„h 16

1-Octene

liquid

-104

122.5

0.716

^0^18

1-Nonene

liquid

- 81

146

0.731

C l0 H 20

1-Decehe

liquid

- 66

171

0.743

c,h 8

m-2 -Butene

gas

-139

0.667

C.H S

trans- 2-Butene

gas

-106

0.649

*The density is given at 20 °C except for those substances which are gases; for
these the density is that of the liquid at its boiling point.

With the alkene hydrocarbons two new types of isomerism arise. The
first depends merely upon the position of the double bond in the molecule.
Thus C 4 H 8 is the molecular formula for both 1-Butene and 2-Butene.

c= C-C-C C-C=C-C

1-Butene 2-Butene

Secondly, unlike singly bonded carbon atoms, thfcse joined by a double bond
cannot rotate freely about the axis between them. The double bond hinders
such rotation so that the molecule at that site has a fixed geometric configura-
tion. If we consider 2-Butene, two structures are possible depending upon
whether the end methyl groups are on the same side or on opposite sides of the
double bond axis. The two compounds have different properties (Table 32-B)
and are known as geometric isomers. They are distinguished by the prefixes
ci$- (Latin: on this side ) and trans- (Latin: across) in accordance with
the relative positions of the methyl groups; in the cis- compound the methyl

Grgan'c Chemistry I: The Hydrocarbons

429

groups are on the same side of the double bond and in the tram- compound
they are on opposite sides.

H-C-CHa H-C-CHa

H II

H-C-CHa HaC-C-H

ds-2-Butene trans-2-Butene

9. Structure of Ethylene. It might seem that the double bond between
the carbon atoms of an allcene molecule should be more stable than the
corresponding single bond of an alkane, but this is not so. The members of the
alkene series are more reactive than those of the alkane series. The C— C
single bond energy is about 80 kcal/mole whereas that of the C=C double
bond is about 100 kcal/mole, less than double the single bond energy. A
possible inference is that the second bond is less strong than the first and
this is, in fact, found to be the case.

In the ethylene molecule, orbitals of the carbon atoms overlap with the
s orbitals of the hydrogen atoms to form the C— H bonds and also with each
other to form a double bond. From each carbon atom there must be four
bonds. In three of these bonds (two carbon-hydrogen bonds and one carbon-
carbon bond), a carbon atom uses hybrid sp 2 orbitals. These are oriented to
give a planar trigonal structure having a bond angle of 120°. (Figure 32.5 , a
and b). Because of a similarity in electron cloud distribution to the bond
formed by the overlap of two pure s- orbitals these bonds are called cr (sigma)
bonds. So far we have accounted for three electrons of each carbon atom. One
p electron still remains of the original three 2 p electrons (one having been
an s electron which was promoted to a p level). The electron cloud of this
remaining p electron consists of two lobes, one above and one below the
plane of the sp 2 orbitals. (Figure 32.6 c). The cloud of the p electron of
one carbon atom overlaps with that of its counterpart p electron of the other
carbon atom to form the second bond of the double bond. This second bond
is perpendicular to the C 2 H 4 plane and is different in kind from those formed
with the sp 2 orbitals. (Figure 32.7). It is called a n (pi) bond and is
weaker than a cr bond. Hence the double bond in an alkene molecule is
truly composed of two kinds of bonds, a relatively strong cr bond with a bond
energy of about 60 kcal/mole and a weaker rr bond with an energy of about
40 kcal/mole. Whereas the total bond strength of a C=C double bond,
about 100 kcal/mole, is greater than that of a single C— C bond formed by
the overlap of two sp 2 orbitals, about 80 kcal/mole in ethane, the weaker tt
bond of the ethylene double bond is more easily broken ' than is the single
C— C bond of ethane and hence ethylene is more reactive. The greater total
bond energy of the double bond results in a smaller C=C distance of 1.34 A
in ethylene compared with the C— C distance of 1.54 A in ethane.

10. Reactions of the Alkene Hydrocarbons. The alkene hydrocarbons are
known as S unsaturated compounds because they contain less hydrogen than
their alkane counterparts and can undergo addition reactions at the site of
the double bond to form compounds which contain only single bonds between
the carbon atoms and are thereby saturated .

430

Organic Chemistry 1: The Hydrocarbons

(a) (b) (c)

(a) Trigorial orientation of sp 2 bonds (b) Incomplete structure of C 2 H 4 molecule
showing o- bonds only (c) Incomplete structure of C 2 H 4 molecule showing orbitals
of the lone p electrons, one for each carbon atom.

Figure 32.6. Formation of the 7 r Bond in the C a H 4 Molecule.

(a) (b)

(a) The r bond between the carbon atoms is due to the overlap of the lone p
electrons; it is perpendicular to the C 2 H 4 plane (b) Dimensions of the C 2 H 4 mole-
cule; — = <r bond; = w bond.

Figure 32.7. The C 2 H 4 Molecule.

(A) Reaction of an alkene with hydrogen yields an alkane

(6) C2H4 + H 2 -> C 2 H 6 (a catalyst, Ni or Pt, is required)

(B) Water can be added to form an alcohol in the presence of H+ ion
which acts as a catalyst

(7) C 2 H 4 + HOH C*H 5 OH (ethyl alcohol)

(C) A halogen element, such as Br 2 , can be added directly. In this case
there is no formation of HBr as in the substitution reaction with an alkane.

(8) C*H 4 + Br a C a H*Br 2 (1,2-dibromomethane)

(D) A special type of addition reaction is one by which two alkene
molecules add to each other or polymerize. The polymerization of ethylene
into 1-butene can be represented by

H H

| |

H H H

1 1 1

-O-

"ft's -

=C— C— C— H

I |

H-~— _

A 0

H H

Organic Chemistry I: The Hydrocarbons

431

Apart from its uses as an intermediate in the manufacture of organic
chemicals, ethylene gas has some unusual properties. Mixed with oxygen
it is a quick acting surgical anesthetic. A small percentage of ethylene in
the air artificially hastens the ripening of citrus fruits and also shortens the
dormant period through which potatoes and other tubers pass before they
start to grow.

11. Diene Hydrocarbons. Hydrocarbon molecules can contain more than
one double bond. Alkene hydrocarbons containing two double bonds are
known as dienes or diolefins . They have the type formula C w H 2n . 2 and are
named after the corresponding alkane by replacing the suffix - ane by - diene .
Numerals indicate the position of the double bonds; thus, 1,3-butadiene is
CH 2 =CH~CH=CH 2 . The chemical properties of a diene depend upon the
arrangement of its double bonds. The arrangement in 1,3-butadiene wherein
the double bonds are separated by two carbon atoms joined by a single bond
is known as a conjugated double bond. Besides ordinary addition across a
double bond between adjacent carbon atoms (1,2 addition), conjugated
dienes undergo 1,4 addition to the carbon atoms at the extremes of the
conjugated system; a double bond is then formed between the central
carbon atoms of the system. Such reactions are of importance in polymeriza-
tion to form long chain molecules, e.g., rubber.

12. The Alkyne Series of Hydrocarbons. A third homologous series of
hydrocarbons, the alkyne series, has the general formula, C n H 2n - 2 * These
molecules contain a triple bond between adjacent carbon atoms. The nomen-
clature of this series uses the suffix -yne and the position of the triple bond
is designated by a numerical prefix. The triple bond is composed of one
relatively strong sp hybrid bond formed by the hybridization of one s and
one p orbital, and two weaker bonds formed through the overlap of the
two remaining p electrons on each carbon atom. Hence alkyne compounds
are even more reactive than those in the alkene series.

Acetylene (ethyne), C 2 H 2 , is the first and most important member of
the alkyne series, some of which are listed in Table 32-C.

Table 32-C _____

Alkynes

Formula

Name

State

at 20° C

Melting
Point , °C

Boiling
Point , °C

Density *
g/ml

c 2 H,

Acetylene (ethyne)

gas

- 82

sublimes-

0.621

c,h 4

Propyne

gas

-101

- 23

0.678

c<h 6

1-Butyne

gas

-122

8.6

0.688

c,h 8

1-Pentyne

liquid

- 98

0.695

1-Hexyne

liquid

-124

0.719

C*H S

2-Butyne

liquid

- 24

0.694

C a H I( ,

2-Hexyne

liquid

- 92

0.730

c 6 h 10

3-Hexyne

liquid

- 51

0.725

*The density is given at 20 °C except for those substances which are gases; for
these the density is that of the liquid at its boiling point.

432

Organic Chemistry I: The Hydrocarbons

Because sp orbitals are oriented at 180° from each other, that is, they
lie in a straight line, the acetylene molecule is linear and has the folloVring
structure:

H— C^C—H where — = cr bond

r\ = 7T bond

Acetylene is produced by the reaction of calcium carbide, CaC 2 , and water.
(9) CaC 2 + 2 H 2 0 C 2 H 2 + Ca(OH) 2

Inasmuch as CaC 2 can be obtained by the reaction of CaO and C, acetylene
can be synthesized from abundant and cheap natural materials. In turn, it
is used in the synthesis of a large number of organic compounds. Acetylene
is a colorless poisonous gas which burns with a yellow smoky flame. It
forms addition compounds in much the same way as do the members of
the alkene series. Acetylene is unstable and, when compressed, may decom-
pose violently into its elements. To avoid explosions it is dissolved in acetone
under pressure and in this form is used for illuminating purposes. Mixed
with oxygen in the oxy-acetylene torch, it produces an intensely hot flame
with temperatures up to 3500°C ( oxy-hydrogen flame gives 2800°C) which
is used in the welding and cutting of metals.

13. Cyclic Aliphatic Hydrocarbons. Some aliphatic hydrocarbon mole-
cules, found principally in petroleum, have a cyclic structure. The prefix
cyclo - is used to name these compounds and their reactions are generally
the same as their open chain analogs. Examples are

CH 2

h s c ^ch 2
h 2 c — ch 2

Cydopentane

H 2 C 5 2CH

\ 4 3

H 2 C CH CH 3

3-Methylcyclopentene

14. Aromatic- Compounds. Benzene, CJ e H 6 , is the first member of a
homologous series of hydrocarbons whose general formula is C n H 2 n 0 . Other
members of this series are toluene, C 6 H 5 CH 3 , and three isomeric xylenes,
C 6 H 4 (CH 3 )2 (Table 32-D). The structure of the benzene molecule is such

h-c^ C ^c-h-~ 4-

H — C ^ ^C— H

11 s

II 1

h-c^ c /C-h —f.

H— C C — H

(a)

<b)

Figure 32,8. Two Resonance Forms of the Benzene Molecule.

Organic Chemistry I: The Hydrocarbons

433

The hexagonal shaped electron clouds are due to the overlap
of the electron clouds of lone p electrons, one from each
carbon atom, above and below the plane of the benzene
ring.

Figure 32.9 . The Benzene Molecule.

ii ra iv

I represents a resonance hybrid of II and HI; IV is commonly used because of its
simplicity but confusion may arise since it is also used to designate cyclohexane.

Figure 32 JO. Symbols Used for the Benzene Nucleus,

434

Organic Chemistry I: The Hydrocarbons

that six carbon atoms form a planar, symmetrical, hexagonal ring, with alter-
nate double and single bonds between the carbon atoms and with a hydrogen
atom attached to each carbon atom. This structure was first proposed by
the German chemist, F. A. Kekul4, in 1865. Two equivalent structures can
be written depending upon the location of the double bonds so that the
benzene molecule is considered to be a resonance hybrid of the two forms.
(Figure 32.8)

The chemical behavior of the double bonds in benzene is different
from that of aliphatic double bonds. For example, the benzene molecule
undergoes substitution for the hydrogen atoms rather than addition across
the double bond. The drawings in Figure 32.8 are not meant to imply that
the benzene molecule has the structure depicted in (a) half the time and
that of (b) the other half of the time, nor is there an equilibrium between
the two forms due to some oscillation of electrons between adjacent carbon
atoms. The structures depicted in Figure 32.8 have no real existence but
are useful and can be considered as pictorial extrapolations of theory. The
benzene molecule has its own unique properties. Despite the pictorial
alternation of the double bonds, the six carbon-carbon bonds are equidistant
and equivalent in all respects. Each such bond behaves as if it were a
hybrid with partial single and double bond character. The carbon-carbon
bond length is 1.39 A, intermediate between that of an aliphatic single bond
and double bond. One electron from each carbon atom, or six in all, is believed
to be delocalized in that it belongs not to a specific carbon-carbon bond
but serves to bind the benzene ring as a unique entity. The six equivalent
carbon-carbon bonds in benzene are c r bonds from the overlap of sp 2 orbitals
so that the benzene molecule is planar with an angle of 120° C between
the C— C bonds. The formation of these bonds leaves six lone p electrons,
one from each carbon atom, which produce electron clouds above and be-
low the plane of the hexagon. In a manner quite analogous to that of the
ethylene molecule, the overlap of the electron clouds of adjacent p electrons
along the hexagon yields two continuous hexagonal shaped it electron clouds
as shown in Figure 32.9. It is characteristic of all aromatic compounds that
they have at least one benzene ring in their structure. In writing reactions
involving aromatic molecules, rather than write out the benzene structure in
detail, the chemist uses a symbol to designate a benzene ring or "nucleus”
common symbols are shown in Figure 32.10.

Benzene, toluene, and the xylenes are colorless liquids at room tempera-
ture. They have an aromatic odor, are highly flammable, and are ex-
cellent solvents for other organic compounds such as grease, resins, and
rubber. Two other aromatic hydrocarbons of importance are naphthalene,
CioHs, and anthracene, Cj 4 H 10 (Figure 32.11). They are related to benzene
but are composed of more than one benzene ring; two carbon atoms are
common in adjacent rings. Naphthalene is commonly seen as moth balls
or flakes.

15. Reactions of the Aromatic Hydrocarbons. The benzene nucleus is
unusually stable. As a resonance hybrid, it has a lower energy, and hence
a greater stability, than that predictable from any of its separate resonance

Organic Chemistry I: The Hydrocarbons

435

forms. It is this stabilizing resonance energy, about 36 kcal/mole, that gives
rise to the set of properties known as aromatic properties. Benzene behaves
chemically like a saturated hydrocarbon, though, with a catalyst, it can be
made to add hydrogen and to form cyclohexane. The main mode of reaction

Table 32-D

Aromatic Hydrocarbons

Formula

Name

State

at 20°C

Melting
Point, °C

Soiling

Point, e C

Density
at 20°C, g/mt

c 6 h 6

Benzene

liquid

5.5

0.879

c 7 h,

Toluene

liquid

- 95

111

0.866

c 8 h 10

O-Xylene

liquid

- 25

144

^8^10

m -Xylene

liquid

- 48

139

0.864

C 8 H 10

p-Xylene

liquid

138

0.861

c 10 h 8

Naphthalene

solid

218

1.15

C U H 10

Anthracene

solid

217

354

1.25

of benzene is to replace one or more of its hydrogen atoms by some
substituent. With chlorine and a catalyst, chlorobenzene is formed.

(10) C 6 H 6 + Cl 2 C 6 H 5 C1 + HC1

+ Cl 2

4- HC1

Because all the carbon atoms are equivalent there can be but one chloro-
benzene. The replacement of two hydrogen atoms by two chlorine atoms
results in a dichlorobenzene, of which there can be three isomers depending
upon the positions of the chlorine atoms. Such isomers are distinguished
by the prefixes ortho (abbreviated: o-), meta (m-), or para (p-); ortho, if the
two substituted groups are on adjacent carbon atoms; meta, if one carbon atom
intervenes; and para, if two carbon atoms intervene.

Cl Cl Cl

ortho-dichlorobenzene meta-dichlorobenzene para-dichlorobenzene

Groups, such as the methyl radical, CHs, can also be substituted for the
hydrogen atoms of a benzene ring.

436

Organic Chemistry I: The Hydrocarbons

toluene oxylene m-xylene p-xylene

C7H8 C s Hio C 8 Hio CgHjo

If more than two groups are substituted on the benzene ring, the com-
pounds are named by indicating the positions of the substituents by a
numerical prefix. Thus the structure of l-chloro-3,5-dimethyl-2,4,6-trinitro-
benzene is:

no 2

Cyclic compounds exist in which the ring contains atoms of other
elements in addition to carbon. Such compounds are classed as heterocyclic
compounds . Examples of heterocyclic compounds are:

H— Cr ^C-H

II II

. — C C — H

1! 1

H C N

[— C (L-H

Pyridine

Thiazole

16. Natural Sources of Organic Compounds. The most abundant source
of hydrocarbons is petroleum though many organic compounds are obtained
from the distillation of coal tar. Since World War II a major source of
aromatic compounds has been the hydroforming process which converts
alkanes and cycloalkanes into aromatic compounds. The passage of methyl-
cyclohexane over an oxide catalyst (A1 2 0 3 and Mo 2 0 3 ) at high temperature
and pressure (55G°C and 300

Organic Chemistry I: The Hydrocarbons

437

Crude petroleum ("rock oil”) is an oily liquid that varies in color from
place to place. Used as a medicament by the Indians and the early Western
settlers, it was first found as an oil that collected on the surface of some ponds in
Pennsylvania. Petroleum is not a single compound but is a complex mixture of
all sorts of hydrocarbons from the simple CH 4 to compounds containing
30 or 40 carbon atoms, and consists primarily of alkanes and cycloalkanes
known in the industry as naphthenes. What were once the complicated
organic molecules of living plants and animals have been transformed by
millions of years of decay and geologic conditions into the hydrocarbons
of petroleum. Petroleum deposits are located in oil-bearing sand under folds
in the overlying rock. A well is drilled through the rock till the oil-bearing
stratum is reached. Often the oil is under pressure and gushes up but more
frequently it is pumped to the surface.

Many products are obtained from petroleum by fractional distillation
since the various hydrocarbons therein have different boiling points. There
is a rough correspondence between molecular weight and the boiling point
of a petroleum constituent. If petroleum is heated the temperature of the
boiling mixture rises steadily as the more volatile components are removed,
and by changing the receiver at definite temperatures a rough separation
into several different fractions is obtained. Each fraction is still a mixture of
hydrocarbons. Some of the products obtained from fractional distillation in
an oil refinery are listed in Table 32-E.

The fraction containing the lubricating oils is redistilled, using superheated
steam, and is divided into several portions— light, medium, and heavy motor
oils. When the oils are chilled, solid paraffin, improperly called “wax,” separates
out and is removed. Such oils are called “paraffin base”; others which leave
a residue of pitch and tar are known as “asphalt base” oils. Today the
principal use of the volatile fractions of petroleum is as a fuel. The demand
for gasoline has led to the cracking process, a method of converting some

Table 32-E

Petroleum Products

Product

Carbon

Number

Distillation
Range , °C

Uses

Gas

c i — C 4

less than 20 °C

Heating fuel; making carbon black
and casing head gasoline

Petroleum ether

c 5 — C 6

20—60

Solvent

Ligroin (naphtha)

C e — C 7

60—100

Solvent; dry cleaning

Gasoline

C 5 — C ls

40—220

Fuel for internal combustion engines

Kerosene

C 12 — C 16

200—320

Fuel for jet engines; illumination

Gas oil (fuel oil)

c 15 — c 18

275—375

Fuel for diesel engines; furnace oil;
“cracking” material

Lubricating oil

C 16 — C 20

over 350

Lubrication

Greases; petroleum

c 20 up

over 350

Lubrication; ointments

Paraffin ("wax”)

c 25 up

solid

M.Pt 50-55

Candles; waterproofing

Pitch; tar

Polycyclic

structures

solid residue

Waterproofing; artificial asphalt

Petroleum coke

Polycyclic

solid residue

Fuel; carbon electrodes

438

Organic Chemistry I: The Hydrocarbons

of the larger molecules with higher boiling points that are within the fuel
oil range to smaller ones with lower boiling points that are within the
gasoline range. When the higher fractions are heated with a catalyst to
about 500°C under pressure, the large molecules are decomposed or "cracked”
into smaller ones. This process has more than doubled the yield of gasoline
from petroleum at the expense of the higher boiling oil. Simultaneously,
alkenes are formed which can be used as the raw materials for the production
of other aliphatic compounds.

In an internal combustion engine a mixture of gasoline vapor and air
is burned rapidly and smoothly. This reaction is highly exothermic and the
expansion of the hot gaseous products, C0 2 and H 2 0, moves the engine
piston. The greater the compression ratio, that is, the ratio of the volumes
in the engine cylinder with the piston at both ends of its stroke, the
greater is the efficiency of the engine. If too high a compression ratio is
employed, however, the mixture bums too rapidly or detonates. The result
is a violent jar against the piston, the engine "‘knocks,” and its efficiency
decreases. Some gasoline mixtures can be used at higher compression
ratios than others without knocking. The tendency to knock is related to
the structure of the hydrocarbon molecule used as fuel and is measured
by the octane number. To compare the anti-knock properties of gasoline two
pure components of gasoline, 2,2,4-trimethylpentane (an isomer of octane,
C 8 Hi 8 , called ""isooctane”) and normal heptane, C 7 Hi 6 , are taken as standards.
Isooctane has good antiknock properties and is arbitrarily assigned an octane
number of 100 while normal heptane, which has poor antiknock properties,
is given an octane number of 0. The percent of octane in a synthetic mixture
of normal heptane and the octane having the same antiknock property as
the gasoline in question is the octane number rating of the gasoline.
Branched chain alkanes, cycloalkanes, and unsaturated hydrocarbons have
good, antiknock properties and are sometimes added to straight run gasoline.
This permits a higher compression ratio in engine design and a greater
efficiency of operation. Some modern fuels have better antiknock properties
than isooctane and hence an octane number greater than 100. The octane
number of gasoline can also be increased by the addition of substances that
act as negative catalysts for the burning rate of hydrocarbons. Lead tetraethyl,
Pb(C 2 H 5 )4 is most commonly used in the United States.

In the production of coke, the destructive distillation of coke yields a
black tarry liquid, rich in aromatic hydrocarbons and known as coal tar.
Its fractional distillation, in a manner quite analogous to that of petroleum,
yields a variety of aromatic compounds, from which can be synthesized dyes,
drugs, plastics, and a host of other compounds.

QUESTIONS

1. Why are molecular formulas insufficient, and structural formulas necessary,
in die study of organic chemistry?

2. Account for the equivalence of the four carbon bonds in carbon compounds

Organic Chemistry I: The Hydrocarbons

439

3. Ilustrate what is meant by each of the following terms: (a) straight chain
hydrocarbon (b) branched chain hydrocarbon (c) homologous series (d) cyclic
structure (e) isomerism.

4. What are the general formulas for the following series of hydrocarbons:
(a) alkane (b) alkene (c) alkyne (d) benzene?

5. To what homologous series might these compounds belong? Draw the structure
of each: (a) C 4 H 8 (b) C 6 H 6 (c) C 5 H 12 (d) C 6 H 10 (e) C 5 H 8 (f) C 7 H 8 .

6. (a) Draw the structural formulas of the five isomers of hexane and n#me
each according to the IUPAC system (b) Draw the isomers of heptane and
name each.

7. Draw a graph of boiling point against the number of carbon atoms for
the compounds of the alkane series.

8. What is the structural difference between saturated and unsaturated hydro-
carbons? What chemical reaction will distinguish between these types?

9. Discuss the carbon-carbon single bond in terms of its bonding orbitals. Draw
the structural formula for ethane and label the types of bonds therein.

10 (a) What is the nature of the double bond in ethylene? (b) Explain the

reactivity of the double bond in terms of its bond character.

11. What is the spatial arrangement of the carbon atoms in butane?

12. What is meant by the term “geometric isomer”? What is the significance of
the prefixes cis and trans? Why are they inapplicable to butane? Which would
have a higher polarity, cis- or trans- butene?

13. What are the names of the hydrocarbons represented by the following?

(b) C-C= C-C-C-C-C

A A A

14. Write structural formulas for (a) 1-pentene (b) 2,2-dimethyIpropane (c) 2,
3-dimethyl-4-ethyl-3-octene (d) 2-methyl- 1,3-butadiene (e) 1,4-dimethylben-
zene (p-toluene) (f) methylcyclobutane (g) isohexane (h) tetramethylmethane.

15. What is a diene hydrocarbon? What are conjugate double bonds? What is
meant by 1,4 addition to butadiene?

16. Write reactions which illustrate polymerization of (a) ethylene and (b)
butadiene.

17. Discuss the structure of the benzene molecule including (a) the significance
of its resonance forms (b) the nature of its double bonds (c) its stability.

18. Illustrate what is meant by a heterocyclic compound,

19. Define the terms: (a) 'pyrolysis (b) cracking (c) compression ratio of an
engine (d) octane number.

20. List some important products obtained in the distillation of petroleum.

21. Combustion of 2.8 g of a hydrocarbon formed 8.8 g of C0 2 and 3.6 g of

H a O. At 27 °C and 600 mm pressure, 5.6 g of the compound occupied 3.1

liters. What is the molecular formula of the hydrocarbon?

22. Construct a table, listing the orbitals used in bonding in the following com-

pounds and their spatial orientations: (a) CH 4 (b) C 2 H ft (c) C 2 H 4 (d) C 2 H 2
(el C»H* (f) 2-methvlnrooene.

(a) c— C-C-C-C— C-C

A A A
A

Organic Chemistry II
Derivatives of the Hydrocarbons

The carbon-carbon bond, both in aliphatic and aromatic compounds, is
extremely stable and, in the great majority of organic reactions, such bonds
are not broken. At least theoretically, therefore, all other organic compounds
can be considered as being derived from the hydrocarbons through the re-
placement of one or more hydrogen atoms by another kind of atom or group
of atoms (radical). In this way it is possible to visualize a great number
of carbon compounds that have as a structural skeleton the same chain
or ring of carbon atoms found in the hydrocarbons, and it is customary to
regard such compounds as derivatives of the hydrocarbons. The substituted
group is known" as a functional group because it is principally this group
which determines the properties and behavior of the compound. It is con-
venient to classify hydrocarbon derivatives according to the nature of the
functional group.

1. Halogen Derivatives. The substitution of a halogen atom for a hydro-
gen atom of a hydrocarbon molecule yields a halogen derivative of that
hydrocarbon. Methyl chloride, CH 3 C 1 , and chlorobenzene, C 6 H fi Cl, both
formed by the direct reaction of chlorine with methane, CH 4 , and benzene,
C e He, respectively, are halogen derivatives of those compounds. For sim-
plicity in designating an aliphatic derivative, the symbol *R* is used to
represent the alkyl nucleus to which the functional group is attached. For
aromatic compounds the specific ring symbols are generally used, but some-
times also the symbol ‘Ar’ for the aromatic nucleus. Thus the general
formula for a chlorine derivative of an aliphatic hydrocarbon is written

R— Cl or R:C1; that for an aromatic hydrocarbon would be (Oj” ri ox Ar— Cl.

The most common reaction of alkyl halides is substitution for its halogen
atom. With NaOH the product is an alcohol; with NH S an amine is formed.

(1) R~X + NaOH NaX + R-OH (an alcohol)

CH3CI + NaOH -» NaCl + CH 3 OH (methyl alcohol)

CH3CI + NHa — > HC 1 + CH3NH2 (methylamine)

Organic Chemistry II: Derivatives of the Hydrocarbons

441

2. Alcohols. For aliphatic compounds, the —OH or hydroxyl group is
characteristic of the class of compounds known as alcohols. In naming the
alcohols the suffix - 61 is used and, where necessary, the position of the
OH group in the parent carbon chain is specified by a numerical prefix.
The structures of some common alcohols are shown below and their properties
given in Table 33-A.

H-C-OH

Methanol
Methyl alcohol

H H

i i

H-C-C-OH

I 1

H H

Ethanol
Ethyl alcohol

H H H

i i I

H-C-C-C-OH

1-Propanol
n-Propyl alcohol

H H H

i i i

H-C-C-C-H

H OH H

2-Propanol
Isopropyl alcohol

Table 33-A

Alcohols

Formula

Melting
Point , °C

Boiling
Point , °C

Density

gfml

Solubility
g/lOOg H 2 0

CH 3 OH

Methanol

liquid

- 97

64.5

0.793

CH 3 CH 2 OH

Ethanol

liquid

-115

78.3

0.789

ch 3 ch 2 ch 2 oh

l -Propanol

liquid

-126

0.804

CH 3 (CH 2 ) 2 CH 2 OH

1 -Butanol

liquid

- 90

118

0.810

7.9

CH 3 (CH 2 ) 3 CH 2 OH

1-Pentanol

liquid

- 78

138

0.817

2.3

ch 3 chohch 3

2-Propanol

liquid

- 86

0.789

In general alcohols with carbon numbers up to twelve are liquids; higher alcohols
are solids. Alcohols with the prefix 1- are normal alcohols in that the OH group
is attached to a terminal carbon atom. In 2-propanoJ the OH group is attached to
the second carbon atom in the carbon chain; this compound is also a secondary
alcohol in that the OH group is bonded to a secondary carbon atom.

The -OH group of an alcohol is quite different in character from that
of an inorganic hydroxide. The C— OH bond in an alcohol is covalent
and does not ionize to give hydroxide ion. However the OH group is a
very polar group so that, unlike the parent hydrocarbon, the lower alcohols
are soluble in water. The solubility decreases as the length of the carbon
chain increases and as the influence of the hydrocarbon portion of the
molecule becomes relatively more important. On the other hand, alcohols
make versatile solvents. The hydrocarbon, or nonpolar portion of the mole-
cule makes it a good solvent for molecular solutes while the polar nature of
the molecule, especially in the lower alcohols, enables it to dissolve some
ionic solutes. Alcohols are commonly used as solvents for carrying out organic
reactions.

The relatively hi gh boiling point of an alcohol such as CHsOH (64.5 C)
compared with those of CH 4 (— 162°C) and CHsCl (— 24°C), is attributed to

442

Organic Chemistry II: Derivatives of the Hydrocarbons

association of alcohol molecules by hydrogen bonds which join the oxygen
atoms of two or more molecules.

R-O • • • H-0

k k

Alcohols can be prepared by the hydrolysis of alkyl halides (Equation 1 )
or by the addition of the elements of water to an alkene.

(2) C 2 H 4 + H 2 0 C 2 H 5 OH (with acid as a catalyst)

Until 1920 methanol, commonly called wood alcohol, was produced by the
destructive distillation of wood. Today almost all methanol is produced
synthetically by the catalyzed reaction of CO and H 2 .

(3) CO -f 2 H 2 CH s OH (at 400°C and 300 atm; ZnO catalyst)

Ethanol, the common variety of alcohol or grain alcohol, is prepared prin-
cipally by the hydration of ethylene and by the fermentation of sugar
obtained from molasses. With glucose, C 6 Hi 2 0 6 , the fermentation reaction is

(4) C e H 12 0 6 2 C 2 H 5 OH + 2 CO.

Distillation of the resulting solution, about 10% ethanol, yields a constant
boiling mixture of 95% ethanol. Of all organic chemicals, ethanol ranks first
both in quantity and in value of production. From it, as a keystone, almost
all the other types of aliphatic compounds can be synthesized. It is used in
the manufacture of ether, acetaldehyde, acetic acid, and many other chemi-
cals. Both methanol and ethanol are employed as fuels and as industrial
solvents.

Methanol is poisonous if taken internally or even if breathed for a pro-
longed period. Ethanol is classed medically as a hypnotic. It is the alcohol
of intoxicating beverages; in large quantities it, too, is poisonous. To render
industrial ethanol unfit to drink, it is denatured by the addition of a non-
potable substance such as methanol or gasoline. Such denaturants are
difficult to separate from the ethanol.

Alcohols can have more than one —OH group, one on each of two or
more carbon atoms. Such compounds with more than one functional group
are called polyfunctional compounds. The structural formulas of ethylene
glycol, C 2 H 4 (OH) 2 , and glycerol (glycerin), CaH 5 (OH) s are

H H

H H H

h-c-A-h

1 1

H-C-A-A-H

OH OH

1 1 1

OH OH OH

Ethylene glycol

1 , 2 -ethanediol

Glycerol

1,2,3-propanetriol

Ethylene glycol is prepared from ethylene and is used as an antifreeze for
automobile radiators (“Prestone”). Glycerol is a bv-oroduct of the hvdrolvsis

Organic Chemistry 11: Derivatives of the Hydrocarbons

443

of fats and oils in the manufacture of soap. It is also used as an antifreeze,
and in the production of nitroglycerin and cosmetic formulations.

3. Phenols. Aromatic compounds in which the —OH group is attached
directly to the ring structure are not alcohols. They are phenols and differ
markedly in properties from alcohols. The simplest and most important phenol
is hydroxybenzene, C 6 H G OH; it is named phenol and is commonly known
as carbolic acid. The methyl phenols are called cresols (Figure 33.1), A

1,4-Benzenedioi
Hydroquinone

Figure 33 J. Phenols.

compound such as benzyl alcohol, C 6 H 5 CH 2 OH, in which the OH group
is not bonded directly to the benzene ring but is on a substituted aliphatic
chain, is properly an alcohol and not a phenol Though phenol and the
cresols are components of coal tar, most of their supply is synthesized from
benzene and toluene. In one such process, benzene is chlorinated to chloro-
benzene which is then reacted with aqueous NaOH; the Cl atom is thereby
replaced by an OH group to yield phenol.

Phenols in general are weak acids with ionization constants about 1(H 0 ;
phenol will react with a strong base, such as NaOH, to form a salt, sodium
phenoxide, C 8 H 5 ONa+. Pure phenol is a colorless, crystalline solid, moderate-
ly soluble in water. It is poisonous and corrosive to tissue; a dilute solution
is used as an antiseptic. It is also used in the preparation of dyes, drugs, and
resins, e.g., Bakelite. Hydroquinone (1,4-benzenediol) is a mild reducing
agent and is therefore used widely as a photographic developer.

4. Ethers. The elimination of a single molecule of water between two
molecules of an alcohol yields an ether.

(5) R-io H + H! 0 -R' R-O-R' + H z O

The general formula for an ether is R—O— R', wherein the radicals, R and R',
may be the same or different, aliphatic or aromatic. The simplest ether is
dimethyl ether, CH 3 OCH 3 , which is an isomer of ethanol, C 2 H 5 OH. Their
properties are quite different, however. Diethyl ether, C 2 H 5 OC 2 H 5 , the
common “ether,” is the most important of the ethers. It is a colorless, volatile,
highly flammable liquid, and is used as an anesthetic. Ethers are chemically
unreactive. Unlike alcohols, they have no hydrogen atoms bonded directly
to oxygen atoms and so cannot form hydrogen bonds. As a consequence

CH 3

m-Cresol

444 Organic Chemistry 11; Derivatives of the Hydrocarbons

they do not associate in the liquid state and so are highly volatile and have
low boiling points.

5. Aldehydes and Ketones. Aldehydes have the type formula R— d]r=0
R

and ketones the formula R— C— O. The groups, R and R', may be the same,
different, aliphatic, or aromatic. Both aldehydes and ketoues contain the
C=0, or carbonyl , group. It is this group which primarily determines their
properties. Because the aldehyde group, — CHO, can be bonded to only one
other atom it must be at the “end” of a molecule whereas the carbonyl group
of a ketone is “within” the carbon chain. Aldehyde compounds are specified
by the suffix -at and ketones by the suffix -one. The structures of some
aldehydes are shown in Figure 33.2 and their properties are listed in Table 33-B.

H— C=0
Formaldehyde

h 8 c1ch s

Acetone

Figure 33.2 . Aldehydes and Ketones.

Table 33-B

Aldehydes and Ketones

Formula

Name

State
at 20°C

Melting

Point , °C

Boiling
Point, °C

Density *
g/ml

Solubility
g/lOOg Hfi

HCHO

Formaldehyde

gas

0.815

very sol.

CH s CHO

Acetaldehyde

liquid

0.783

C e H„CHO

Benzaldehyde

liquid

1.050

0.33

CH s COCH 8

Acetone

liquid

- 94

0.792

C e H s COCH s

Acetophenone

solid

202

1.026

ins.

C 8 H 5 COC„H 5

Benzophenone

solid

306

1.1

ins.

♦The density is given at 20 °C except for those substances which are gases; for these
the density is that of the liquid at its boiling point.

Aldehydes can be prepared by the oxidation (or dehydrogenation) of
primary alcohols. The simplest aldehyde, formaldehyde, HCHO, is made
by the catalytic oxidation of methanol at about 55Q°C.

H 3 C-C:=0

Acetaldehyde

Benzaldehyde

Acetophenone

Benzophenone

Organic Chemistry II: Derivatives of the Hydrocarbons

445

( 6 )

H-C-O JH + [O] H— C=0 + H a O

2 CH 3 OH + O z 2 HCHO + 2 H 2 0

Formaldehyde is a gas with a sharp odor. A 37% aqueous solution, known
commercially as "formalin,” is used as an embalming fluid to preserve ana-
tomical specimens, and as a fumigant. Formaldehyde has a tendency to
polymerize and to form the solid paraformaldehyde, (CH 2 0) n , from which,
however, it can be regenerated upon heating. Bakelite, a synthetic resin, is
produced by the reaction of formaldehyde and phenol. Acetaldehyde is
synthesized by the mild oxidation of ethanol, or by the addition of water
to acetylene, C 2 H 2 . It is used chiefly in the production of acetic acid,
CH3COOH.

Ketones are prepared by the oxidation of secondary alcohols. The oxida-
tion of 2-propanol yields acetone (dimethyl ketone), the simplest and most
important of the ketones.

(7) CH 3 -CHOH-CH 3 + [O] CH 3 -CO-CH 3 + h 2 o

Acetone is also prepared by the Weitzmann process for the fermentation of
starch, a process developed during World War I by Chaim Weitzmann,
later the firsc president of Israel. Acetone is an excellent solvent, used so
industrially for varnishes, plastics, cellulose derivatives, and smokeless powder.

The typical reaction of aldehydes and ketones is addition across the
C=0 double bond; aldehydes undergo such reaction moie readily than do
ketones. By the addition of H 2 , both aldehydes and ketones can be readily
reduced to alcohols, a reaction which is the reverse of their preparation.
Ketones are not easily oxidized but aldehydes are oxidized to organic acids
even by mild oxidizing agents. Hence aldehydes act as reducing agents;
they reduce Ag(NH 3 ) 2 + ion to silver metal, and Fehlings solution (a solution
of a complex copper (II) tartrate ion) to red insoluble copper (I) oxide, Cu 2 0.

6. Acids. The carboxyl , or — COOH, group is characteristic of the type
of organic compounds known as carboxylic acids, whose general formula is
O

R~C— OH. The carboxyl group must be a terminal group. The structures
of some important carboxylic acids are drawn in Figure 33.3 and their
properties are given in Table 33-C.

—OH

O O

Formic

Acid

Acetic

Acid

Oxalic

Acid

Benzoic

Acid

Figure 33.3. Carboxylic Adda.

446

Organic Chemistry II: Derivatives of the Hydrocarbons

Table 33-C

Carboxylic Acids

Formula

Name

State

at20°C

Melting
Point, °C ’

Boiling
Point, °C

Density

g/ml

Solubility
S/I00g Hfi

HCOOH

Formic

liquid

8.4

100.5

1.220

CH 3 COOH

Acetic

liquid

16.6

118

1.040

HOOC-COOH

Oxalic

solid

189

1.653

9.5

CH s (CH 2 ) 18 COOH

Stearic

solid

290*

0.847

insol.

C e H 0 COOH

Benzoic

solid

122

249

1.260

0.18

o-C 8 H 4 (COOH) 2

Phthalic

solid

231

■

0.70

o-HOC 8 H 4 COOH

Salicylic

solid

159

0.22

*At 100 mm pressure

The carboxylic acids can be produced by the oxidation of the aldehydes.
O O

(7) CH 3 -C-H + [O] CH 3 -C-OH

An acid may be considered the “final organic stage” in the oxidation of a
hydrocarbon, the first stage of which was an alcohol. The successive stages
of oxidation from ethane to acetic acid are:

(8) C 2 H 6 C 2 H 5 OH CH3CHO CH3COOH

Further oxidation of the acid by strong oxidizing agents yields C0 2 and H 2 0 ,
Formic acid is specially synthesized by the reaction of CO and aqueous
NaOH at 200°C and 100 lb/in 2 pressure. This yields the sodium salt, sodium
formate, HCO<>Na + Upon acidification the acid is formed.

H+

(9) CO + NaOH HCOO-Na+ * HCOOH

Acetic acid is obtained by the oxidation of acetaldehyde and by the destruc-
tive distillation of wood. It is the most common and most important organic
acid. Vinegar is a dilute solution, about 4% acetic acid, while glacial acetic
acid is the pure acid, so called because it freezes to ice-like crystals.

Carboxylic adds are polar molecules and, like alcohols, associate by
hydrogen bonding. The dimeric form of acetic acid is

O-H-O

H S C-CT ^C-CHs

O-H-O'"'

Carboxylic acids are weak; we are already familiar with the behavior of
acetic add, the classic example of a weak acid. The acid is monoprotic, the
only ionizable hydrogen atom being the one in the carboxyl group. With
strong bases, such as NaOH, salts are formed.

(10) CH3GOOH + NaOH -» H 2 0 + CH s COO~Na +

Organic Chemistry II: Derivatives of the Hydrocarbons

447

The principal source of the long chain acids, e.g., stearic acid, is the hydro-
lysis of fats; hence such acids are known as fatty acids .

Acids can contain more than one carboxyl group, e.g., oxalic acid. They
are polyprotic and, as in inorganic chemistry, the primary ionization is the
strongest. Their reactions are similar in kind to those of the monocarboxylic
acids in that one or more of the ionizable hydrogen atoms can be replaced.
Many acids, known as hydroxy-acids , contain both carboxyl and hydroxyl
groups, e.g., lactic acid. Examples of these typ* s of acids are shown in Figure
33.4. Oxalic acid occurs in plants, such as rhubarb and sorrel, and in tart

ch 3

H 2 C-COOH

h 2 c-cooh

ho-c-cooh

H-C-OH

ho-c-cooh

COOH

ho-c-cooh

h 2 c-cooh

Lactic Acid

Malic Acid

£[

Tartaric Acid

Citric Acid

Figure 33.4. Hydroxy Acids.

fruits as a calcium or potassium salt. Lactic acid is formed in sour milk by
the fermentation of lactose (milk sugar), C^H^On, a sugar much like cane
sugar. Malic acid is found free or as salts in fruit, such as apples, pears, and
cherries. Tartaric acid is widely distributed in fruits; it occurs as the salt,
potassium hydrogen tartrate, KHC*H 4 O e , in grapes. Citric acid occurs in
citrus fruits, such as lemons and oranges.

Another group of organic acids is the sulfonic acids. The sulfonic group,
— S0 2 0H, is strongly acidic. Benzene sulfonic acid, C 6 H 5 S0 2 0H, is an example;
it is prepared by the reaction of concentrated H 2 S0 4 directly with benzene.

7. Esters, The product of the reaction between an acid and an alcohol is
an ester . The reaction, called esterification , is slow and reversible. It is usually
carried out at high temperature and in the presence of an acid, which acts
as a catalyst.

( 11 )

r_C-|OH_+ Hj-Q-R'

R-

,-C-O-R'

+ H a O

Oxygen isotope tracer experiments have proved that, in esterification, it is
the acid molecule which contributes the OH group to the formation of water.
The general formula of an ester is RCOOR', so that an ester differs from an
acid in that the ionizable hydrogen atom of the carboxyl group is replaced
by a radical, usually alkyl. As a consequence, esters do not ionize. In naming
an ester, the radical originating from the alcohol is stated first. Thus, methyl
anptatp is an ester obtained from methyl alcohol and acetic acid.

448

Organic Chemistry II: Derivatives of the Hydrocarbons

The esterification of acetic acid and ethanol is a classic example of an
equilibrium in solution, and to which Le Chatelier s principle is applicable.

(12) CHgCOOH + C 2 H 5 OH CHsCOOCaH* + H a O

Since the total number of inoles is unchanged in the reaction, an equilibrium
constant, K, can be written in terms of the number of moles, n, of each
substance,

( 13 ) K = ggg- X ° w ^ - er - = 4.0 (at 2S°C)

Uacid X alcohol

The reaction of an ester with a strong base, such as NaOH or KOH, is
called saponification. The ester hydrolyzes to yield the alcohol and the sodium
salt of the acid from which the ester was originally constituted.

The low molecular weight esters are liquids which have pleasant, flower-
like or fruit-like odors, and are used in perfumes and artificial flavorings.
Isoamyl acetate, CHsCOOCsHu, has the odor of bananas; methyl butyrate,
C 8 H 7 COOCH 3 , that of pineapples; and octyl acetate, CH 3 COOC 8 H i7 , that
of oranges; methyl salicylate, C 6 H 4 (OH)COOCH 3 , is Oil of wintergreen.
The esterification of the phenolic OH group of salicylic acid with acetic
acid yields acetylsalicylic acid, CH 3 COOC 6 H*COOH, or aspirin. Waxes
are esters of high molecular weight alcohols and acids; beeswax is
CisHsiCOOCaiHag. Esters make excellent solvents; ethyl acetate and butyl
acetate are used extensively as lacquer solvents.

8. Amines and Amides. An amine may be considered as a derivative of
ammonia, NH 3 , in which one or more organic radicals, aliphatic or aromatic,
has been substituted for the hydrogen atoms. There can be primary, second-
ary, or tertiary amines according to whether the nitrogen atom is bonded to
one, two, or three carbon atoms, respectively. The corresponding formulas
are RNH 2 , R 2 NH, and R S N. Quaternary ammonium salts, such as R*NX, can
also be prepared. To denote an amine, either the suffix amine or the prefix
amino is used. Aromatic amines are named as derivatives of aniline, C e H 5 NH 2 ,
the simplest aromatic amine. Some simple amines are shown in Figure 33.5.

H S C . CH 3

H 3 C-NH 2 ^NH I

H S C H 3 C-N-CH 3

Methylamine Dimethylamine Trimethylamine

Figure 33.5, Amines.

The chemical behavior of an amine is very similar to that of NH S . As
in NH 3 , the nitrogen atom of an amine has an unshared pair of electrons.
Hence the amines are basic and form salts with strong acids analogous to
ammonium, NH* + , salts.

(14) CH3NH2 -f- HC1 CH 3 NH 3 +C1“ (methylammonium chloride)

< 22 >- nh -

Aniline

Organic Chemistry II: Derivatives of the Hydrocarbons

449

The alkylamines have fish-like odors. The aromatic amines are toxic and
can be absorbed through the skin. The amines find use in the synthesis of
dyes, medicinals, and photographic developers.

An amide has the general formula, R— C— NH 2 , and may be considered

the product of the reaction between a carboxylic acid and NH 3 , with the
splitting off of water. An amide can be formed by heating an ammonium
salt of an organic acid to high temperature. The nomenclature uses the suffix
amide ; CH 3 CONH 2 is acetamide, and C 6 H 5 CONH 2 is benzamide. The amides
are neutral; when heated with acid or base, the amides are hydrolyzed to
the component carboxylic acid and NH 3 , or their salts.

9. Amino Acids. An aminoacid is a compound which contains both
carboxyl and amino groups. The — NH 2 group can be attached to any carbon
atom other than that of the — COOH group. If it is bonded to the adjacent
carbon atom the acid is known as an alpha, a, aminoacid. Successive posi-
tions of the — NH 2 group along the carbon chain with respect to the —COOH
group are designated by the corresponding successive letters of the Greek
alphabet; beta, /3, gamma, y, etc. The simplest aminoacid is glycine (amino-
acetic acid), NH 2 CH 2 COOH. The aminoacids are colorless, high melting,
crystalline solids which are generally insoluble in nonpolar solvents but are
soluble in water. They have high dipole moments. Such properties are in-
dicative of a salt-like structure so that aminoacids are believed to exist as
dipolar molecules ( “zwitter ion") due to the shift of a proton from the
carboxyl to the amino group. Thus the structure of glycine is NH 3 + CH 2 COO
The NH 3 + group acts as an acid and the COO" group as a base so that
an aminoacid is amphoteric and can react with both acid and base. The
ionization constants for aminoacids, however, are very low; for glycine,
K a is 1.6 X 10~ 10 and K b is 2.5 X 10 -12 . Aminoacids are obtained by the
hydrolysis of proteins and are of great biochemical importance in that they
are the units from which proteins are built.

10. Nitro Compounds. When hydrocarbons are treated with HN0 3 one
or more nitro groups, — N0 2 , can be substituted for hydrogen atoms to
produce nitro derivatives having the general formula R-N0 2 . Thus C 2 H 5 N0 2
is nitroethane and C 6 H 5 N0 2 is nitrobenzene. The nitration of phenol yields
picric acid, 2,4,6-trinitrophenol, and the nitration of toluene gives 2,4,6-trini-
trotoluene (TNT), both of which are explosives. The aromatic nitro com-
pounds are rather more important than the aliphatic ones.

11. Summary The various types of hydrocarbon derivatives are listed
in Table 33-D.

12. Optical Activity. The electromagnetic theory of radiation postulates
that light is a wave motion in which electric and magnetic fields vibrate,
or vary in intensity, at right angles to each other and to the direction in
which the light travels. In ordinary light these vibrations can be in any
of the infinite number of planes which might be drawn through a line repre-
senting the direction of propagation of the light, Tlane-polarized light is light
whose vibrations are restricted to only one plane. The light transmitted after

450

Organic Chemistry II: Derivatives of the Hydrocarbons

Table 33-D

Hydrocarbon Derivatives

Functional

Group

Formula

Compound

Example

Name

-OH

R-OH

Alcohol

ch 3 ~oh

Methyl alcohol

—OH

Ar-OH

Phenol

C 9 H s -OH

Phenol

kXh

CH3-H-H

-C-H

Aldehyde

Acetaldehyde

R— C— R'

Ketone

CHo—C— CH.

Acetone

-C-OH

CH 3 -C-OH

R-C-OH

Acid

Acetic acid

-O-

R— O— R'

Ether

CH 3 — 0—C 2 H 5

Methylethyl

ether

Xo-

R-C-OR'

Ester

ch 3 -c-o-c 2 h 5

Ethyl acetate

-nh 2

r-nh 2

Amine

CH.-Nhk

Methylamine

r-c-nh 2

ch 3 -c— nh 2

— C— NHo

Amide

Acetamide

-no 2

r~no 2

Nitro

ch 3 -no 2

Nitromethane

passing through certain substances, such as calcite, a form of calcium car-
bonate, CaC0 3 , is plane-polarized. A substance which is optically active
rotates the plane of polarized light, that is, when light vibrating in one plane
only passes through such a substance it emerges vibrating in a different plane
which is at some angle to the initial plane. If an optically active substance
rotates the plane of polarized light clockwise, or to the right, it is said to
be dextrorotatory; if the rotation is counterclockwise (to the left) the sub-
stance is levorotatory. A schematic illustration of these concepts is given
in Figure 33,6.

All types of organic compounds can be optically active. A compound is
optically active if its structure and the structure of its mirror image are not
superimposable. In Figure 33.7 the structures of 2-chlorobutane and its mirror
image are drawn. Molecular models would most easily indicate that the two
structures are not superimposable. However, without removing a molecular
structure from the plane of the paper, no rotation, or twisting, or sliding
will enable one structure to be superimposed upon and match the other
structure. They bear the same relation to each other as do the right and left
hands. Thus there are two types of 2-chlorobutane molecules which differ
because of the spatial orientation of their atoms and hence are isomers. Such

Organic Chemistry II: Derivatives of the Hydrocarbons

451

(a) (b) (c)

(a) Ordinary light vibrates in an infinite number of planes; for light travelling
perpendicular to the plane of the page, the arrows represent vibrations in
the plane of the page.

(b) Plane-polarized light vibrates in only one plane, shown as AB.

(c) AB represents the plane of polarized light prior to its passing through an
optically active substance; after such passage the plane of the polarized light
is CD. The plane has been rotated through an angle a.

Figure 33.6. Optical Rotation.

isomers which are mirror images of each other are known as enantiomers .
They have identical physical properties except for the direction of rotation
of the plane of polarized light, and they have identical chemical properties
except toward other optically active substances.

A carbon atom attached to four different groups is known as an asymmetric
carbon atom. In 2-chlorobutane the second carbon atom of the butane chain
is asymmetric; the four different groups to which it is bonded are H, CH 3 , Cl,
and C 2 H 5 . Most compounds containing an asymmetric carbon atom are op-
tically active so that the presence of such a carbon atom is a good indication
of optical activity. Optical activity is an important tool in determining the
configurations, or arrangements of atoms, in organic compounds but a detailed
study of this subject is better left to courses in organic chemistry.

The two structures are not superimposable; M represents a mirror.

Figure 33.7 . Structure of 2-Chlorobutane.

QUESTIONS

1. What is the significance of the term “functional group”?

2. What is the functional group in each of the following types of organic com-
pounds: alcohol; ether; aldehyde; ketone; carboxylic acid; sulfonic acid; amine;
nitro compound; ester; phenol; alkyl chloride; amide?

452 Organic Chemistry lh Derivatives of the Hydrocarbons

3. Draw the structure of each of the following; (a) l-chloro-2-butanol (b) 3-amino-
butanoic acid (c) propanone (d) ethanal (e) p-aminobenzoic acid (f) 2-phenyl-
2-propanol (g) m-nitrophenol (h) propyl acetate (i) benzamide (j) /3-hydroxy-
pentanoic acid (k) l s 2-ethanediol.

4. What are the products obtained by the progressive oxidation of a hydro-
carbon? Illustrate these products starting with propane.

5. Write structural formulas to illustrate the reaction of 2-pentene with water.
Name the product (s) formed.

6. Write the structural formulas and name the isomers of alcohols having the
formula, C 4 H 9 OH.

7. Ethanol and dimethyl ether are isomers, yet one associates and the other
does not. Explain.

8. What types of organic compounds can associate by hydrogen bonding? Draw
a probable structure for two such associated molecules.

9. Draw structural formulas to illustrate the following reactions: (a) the forma-
tion of an ether (b) methyl bromide and NaOH (c) benzoic acid and
ethanol (d) acetic acid and NH 3 (e) hydrolysis of propyl acetate (f) oxidation
of a secondary alcohol (g) reduction with hydrogen of propanal (h) 2-chloro-
propane and NH 3 .

10. What products might be formed in the dehydration of a mixture of ethanol
and propanol? How could they be separated?

11. What is the difference between an aldehyde and a ketone? Which would
be produced by the oxidation of a secondary alcohol?

12. Why must aldehyde and carboxyl groups be terminal groups?

13. The molecular formula of acetic acid is C 2 H 4 0 2 . Why is it monoprotic?

14. Which of the following would be most soluble in water: propane, 1-propanol,
1-hexanol? Explain your choice.

15. Using structural formulas write equations for the 1 following conversions:
(a) 2-propanol to propylene (b) 1-butene to 2-bromobutane (c) ethylene
to acetamide (d) ethanol to ethyl acetate.

16. Starting with coke, C, and limestone, CaCO s , outline a series of reactions
which would produce acetic acid.

17. Outline a synthesis of methyl salicylate (oil of wintergreen), starting with
phenol.

18. How could an impurity of palmitic acid be removed from gasoline?

19. What theoretical weight of acetic acid could be obtained from the oxidation
of 100 g of ethanol?

20. Two moles of ethanol are mixed with one mole of acetic acid in a four liter
container at 25 °C. The equilibrium constant for the reaction at 25° C is 4.0.
How many moles of ester will be present at equilibrium?

21. What is meant by optical activity? What property must a compound have
for it to be optically active? Draw diagrams which illustrate this property.

22. Which of the following are optically active? In each case label an asymmetric
carbon atom: (a) ethanol (b) 2-methyl-l-butanol (c) 3-methylhexane (d) pro-
pionic acid (e) alanine, CH 3 -CHNH 2 — COOH (f) glycerol (g) glyceraldehyde,
CH 2 OH-CHOH~CHO (h) lactic acid.

23. list properties in which enantiomers differ and properties in which they are
alike.

Organic Chemistry III
Polymers and Biochemistry

1. Natural and Synthetic Rubber. Natural rubber is a hydrocarbon ob-
tained from latex, the milky fluid from the inner bark of .certain tropical
trees, e.g., hevea brasiliensis. When latex is heated, or acidified with acetic
acid, the suspended hydrocarbons coagulate and can be removed from the
residual liquid. This product is the raw rubber of commerce. It is tacky and
sticky, soft when hot, and hard and brittle when cold. If stretched it does
not return to its original dimensions. These undesirable characteristics are
minimized by heating the raw rubber with sulfur to about 130° C, a process
known as vulcanization and discovered in 1839 by Charles Goodyear. During
vulcanization the rubber may be compounded or blended with other sub-
stances such as lampblack (carbon black) or zinc oxide. These substances
improve the wearing qualities of the rubber and make it tough and durable,
a quality necessary for automobile tires.

Upon destructive distillation raw rubber yields the hydrocarbon, C 5 H 8 ,
isoprene, or 2-methyl-l, 3, -butadiene. The isoprene molecule has a conjugated
double bond structure. Such molecules can polymerize by combining with
themselves end to end, and the rubber “molecule” may be considered to be
a long chain polymer of isoprene.

(1) CH S H CHa H GH 3 H CH 3 H

CH 2 =C— C=CK 2 — CH 2 — C=C— CH 2 — CH 2 — (i=C— CH 2 — CH 2 — C=(L~ CH 2 —

isoprene three polymerized isoprene units of rubber

A rubber ‘molecule” may contain approximately 2,000 to 6,000 isoprene
units and have a corresponding molecular weight of 130,000 to 400,000.
In a given sample of rubber or of any polymer, however, the molecules are
not all of the same length and molecular weight.

The long, threadlike rubber molecules are intertwined with each other
in a random manner and have a somewhat coiled structure. When rubber

454

Organic Chemistry III: Polymers and Biochemistry

is stretched the molecules are straightened out but the interlaced struc-
ture prevents their moving past one another. When the stretching force
is removed, the rubber molecules assume their original configurations. Even
after polymerization of the isoprene units, the rubber molecules still con-
tain double bonds. During vulcanization it is to the double bonds of
adjacent rubber molecules that sulfur atoms add, cross-linking or “stitching”
them together to form a three dimensional network. The degree of cross-
linking determines the ability of the rubber molecule to coil and uncoil
and hence controls the rigidity of the rubber product. Most soft rubber
products, such as rubber bands and tires, contain 5-10% sulfur by weight,
whereas hard rubber objects such as combs and bowling balls, may con-
tain as much as 40% sulfur. Exposed to oxygen, especially in sunlight, the

rubber molecule is oxidized at its double bonds, splits into smaller frag-

ments, and the characteristic properties of rubber disappear.

The manifold uses of rubber scarcely need enumeration. More than
one million tons are used annually in the United States. It is a strategic
material, and the isolation of the United States during World War II
from its sources of natural rubber was almost disastrous. Rubberlike ma-
terials, so called “synthetic rubber,” can be synthesized by the polymeriza-

tion of molecules similar to isoprene. Chloroprene, 2-chloro- 1,3-butadiene,
polymerizes in a manner analogous to that of isoprene to give the synthetic
rubber, neoprene. Neoprene is more expensive than natural rubber, but it
stretches better, has greater abrasion resistance, and is more resistant to
hydrocarbon solvents.

2. Polymers and Polymerization. When passed through a hot tube at
600°C, acetylene, C 2 H 2 , polymerizes to form benzene, C„H 0 . However, the
use of the term high polymer or simply polymer is restricted to a large
molecule built up by the joining together of a great many small molecules
in a systematic manner. The small unit molecule is referred to as a monomer
and the joining process is polymerization. Common examples of polymers
are rubber, plastics, resins, synthetic fibers, proteins, and cellulose.

There are two methods of forming polymers: addition polymerization
and condensation polymerization . In addition polymerization, the monomer
molecules add together; the empirical formula of the polymer is the same
as that of the monomer. Ethylene, C 2 H*, polymerizes by addition to form
polyethylene, which is used in making plastic containers, thin packaging
film, insulation, and tubing.

HH HHHHHHHH

M 1 1 1 1 1 1 m

(2) C=C > -C-C-C-C-C-C-C-C-

II I I II I I I I

HH HHHHHHHH

Propylene and derivatives of ethylene polymerize similarly; among the latter
are the important vinyl polymers (the vinyl group is C=C).

Organic Chemistry III: Polymers and Biochemistry

455

CH 2 =CHC1

CH 2 =CH(CN)

CH 2 =CH(C 6 H 5 )

cf 2 =cf 2

vinyl chloride; polymerizes to polyvinyl chloride used in
enamels, phonograph records, and the making of Koroseal.

vinyl cyanide (acrylonitrile) ; polymerizes to Orion used to make
a synthetic textile fiber.

vinyl benzene (styrene); polymerizes to polystyrene,
tetrafluorethylene; polymerizes to Teflon.

A copolymer is a mixed polymer in which the chain consists of more
than one kind of monomeric unit. Thus butadiene and styrene copolymerize
to form the synthetic rubber, Bum S; over 85% of the synthetic rubber
produced in World War II was Buna S.

( 3 )

H H H H

H H

i i

+ C = C »

H H H H H H

<U=LL.i- L

h i /*\ i

butadiene

styrene

a portion of Buna S

The monomers in copolymerization need not be in a one to one ratio;' in
Bum S the ratio of butadiene to styrene is six to one.

In condensation polymerization, the reacting molecules split out a simple
molecule, usually water, upon combining to form the polymer. The formation
of Bakelite , the condensation product of phenol and formaldehyde, is an
example. The hardness of Bakelite is due to the formation of a three-dimen-
sional cross-linked structure.

a portion of Bakelite

Nylon is formed by the reaction of adipic acid, HOOC(CH 2 )*COOH, and
1,6-diaminohexane, H 2 fr(CH 2 ) e NH 2 ; both ends of both molecules have func-
tional groups. A portion of the Nylon molecule is shown below.

HO OH HO OH

CH 2 ) J-N- ( CH 2 ) 6 — N— (L(CH 2 ) 4 — C— N—

The process of polymerization can be carried out by heat and pressure
alone, but the addition of a small amount of a substance known as an initiator

456

Organic Chemistry III: Polymers and Biochemistry

enables the process to proceed more readily. Peroxides are commonly used
as initiators. The peroxide molecule, e.g., R—O : O-R, is decomposed into
a fragment, R—O * , an incomplete molecule, or free radical, with an un-
paired electron. This adds itself ( to a molecule of the monomer, which in
turn becomes a free radical. This then combines with another molecule of
the monomer, etc. Propagation of this chain results in a molecule of the
high polymer. The polymer chain may be terminated at any stage of its
growth by the combination of two free radicals to form a stable molecule
of polymer. The fact that such combination occurs at random accounts for
the different lengths of polymer molecules in a given sample.

Plastics are polymers which can be molded or cast into a .desired form,
If the plastic can be softened upon heating and remolded, it is said to be
thermoplastic. Such plastics are generally composed of long chain linear
molecules and are soluble in some solvent. Polymers which form fibers such
as Nylon and Dacron are of this type. A solution of such a polymer may
be extruded into a precipitating medium to form a thread of the polymer.
Stretching the thread as it is formed aligns the polymer molecules in one
direction and imparts strength to the fiber. Thermosetting plastics are in-
soluble and cannot be remolded, once formed; their molecules form a strong,
highly cross-linked, three dimensional network.

3. Food and Related Compounds. Although plants feed on very simple
substances such as; carbon dioxide, water, and soluble nitrate and nitrite
salts, only complex compounds will nourish animal life. Certain compounds
necessary to the well-being of man are incapable of being synthesized by his
metabolic processes but can be manufactured by plants. To a large extent
animal life is dependent upon the organic compounds synthesized by plants.
That branch of chemistry which deals with the chemistry of life processes,
biochemistry , is perhaps, by its very nature, the most important of all. Bio-
chemistry includes medicinal chemistry, metabolic processes, and perhaps
the essence of life itself, but such subjects are beyond the scope of this book.

There are five classes of foodstuffs, the materials of which foods are
composed, which are essential in the diet of man. These are carbohydrates ,
fats , proteins , vitamins , and mineral matter. All but the last are organic
chemicals. Though water and oxygen are necessary to all living processes
they are not ordinarily classified as foodstuffs.

4. Carbohydrates. The carbohydrates are a group of compounds com-
posed of carbon, hydrogen, and oxygen, in which the ratio of hydrogen atoms
to oxygen atoms is the same two to one ratio that exists in water, namely,
C w H 2 a?0<r. No water molecules are present as such in carbohydrates though
the formula is sometimes written as C,,(H 2 0)*. The most important carbo-
hydrates are sugars, starches, cellulose, and related compounds; the suffix -ose
commonly denotes a carbohydrate. Some important carbohydrates are:

Glucose (dextrose, grape sugar) C 6 H 12 0 6 Lactose (milk sugar) C^H 2 20 n ‘H 2 0
Fructose (levulose, fruit sugar) C 6 H 12 Q 6 Maltose (malt sugar) C 12 H 2 20n'H 2 0
Sucrose (ordinary sugar) C 12 H 22 O u Starch (C 6 U 1Q 0^) x

Cellulose (C 6 H 10 O 5 ) y

Organic Chemistry lib Polymers and Biochemistry

457

Carbohydrates are either polyhydroxy aldehydes, polyhydroxy ketones,
or their condensation products which can be hydrolyzed to these. There are
three main classes of carbohydrates. If a carbohydrate cannot be hydrolyzed
to simpler compounds it is a monosaccharide; if it can be hydrolyzed into
two monosaccharide molecules it is a disaccharide ; and if hydrolysis yields
many monosaccharide molecules it is a polysaccharide. Monosaccharides can
be further classified; one containing an aldehyde group is an aldose and
one containing a ketone group' is a ketose. A carbohydrate is also known as a
triose, tetrose, pentose, hexose, etc., depending upon the number of carbon
atoms in the molecule. As shown in Figure 34.1, open chain and cyclic structures
have been postulated for carbohydrate molecules.

H— C=0

H-C-OH
HO— C— H

H-C-OH

h 2 c-oh

Glucose: an aldohexose Fructose: a 2-ketohexose

Figure 34.1. Structures of Glucose and Fructose.

The monosaccharides are the simplest carbohydrates. Most are pentoses
and hexoses, and the most important are glucose and fructose, both of which
have the formula C e Hia0 6 and hence are isomers of each other. Glucose
is the unit of which starch, cellulose, and glycogen (animal starch) are com-
posed and probably mere glucose units exist in nature than any other organic
group. It is found in fruits usually mixed with fructose; ripe grapes may
contain up to 30% sucrose. In human blood it is present in small quantity,
about 0.1%, and in normal urine in negligible quantity. In persons suffering
from diabetes the glucose content of urine may be 3-10%. Glucose can be
prepared by boiling starch, a polysaccharide, with dilute acid; the H + ions
act as a catalyst for the hydrolysis of the starch.

(5) (C 6 H 10 O 5 )„ + n H 2 0-+ nC fl H 12 0 6

starch glucose

or (CqHioO{i)» ^ Dextrins 1 — > C 12 H 22 O 11 CeHiaOe

starch maltose glucose

1 See page 460.

45S

Organic Chemistry HI: Polymers and Biochemistry

Complete hydrolysis of corn starch yields glucose; partial hydroylsis yields a
mixture known as corn syrup. Not so sweet as cane sugar or sucrose, it is
equally valuable as a food, and considerably cheaper. It is used as a sweeten-
ing agent in the manufacture of candy and preserving fruits. Fructose occurs
in many fruits and in honey; it is the sweetest of all sugars. Glucose, fructose,
and starch can be converted into* ethanol by fermentation.

A disaccharide can be considered as the condensation product of two
molecules of a monosaccharide through the elimination of a molecule of water.

(6) 2 CfiH^Oe — ^ G12H22O31 H2O

The most important of the disaccharides is sucrose (Figure 34.2).

ch 2 oh

H-C

CHOH

CHOH O

CHOH

H-C 1

CH 2 OH

[

H-C

CHOH

-C-O

CH 2 OH

Glucose unit Fructose unit

Figure 34.2. The Sucrose Molecule.

Upon hydrolysis one molecule of sucrose yields one molecule of glucose and
one molecule of fructose. The hydrolysis of sucrose is known as inversion
and the resulting mixture of glucose and fructose is called invert sugar.

All monosaccharides and most disaccharides are reducing sugars. They
will reduce Fehlings solution, precipitating red Cu 2 0. This reaction is a
quantitative test for these sugars. Sucrose is an exception and is not a re-
ducing sugar inasmuch as it contains no aldehyde or ketone group.

Sucrose occurs in sugar cane (15-20%), beets (12-15%), maple sap (2%),
coffee, and honey. The extraction of sucrose from sugar cane involves a
number of operations. The juices of the plant are obtained by crushing the
raw material in a series of roller mills and leaching with the minimum
amount of water. About 95% of the sugar is thus extracted; the remaining
cellulosic material, called bagasse , is used in the manufacture of the insulating
material, "Celotex” A suspension of Ca(OH) 2 in water is then added to the
extract to neutralize the free acids in the sugar cane sap and to precipitate

Organic Chemistry -III: Polymers and Biochemistry

459

certain impurities. The filtrate from this step contains about 10% sucrose.
Under reduced pressure it is evaporated until the water content is about
8% and then allowed to crystallize. The mother liquor, separated by centrifuga-
tion, is molasses and is used to make alcohol. The sugar thus obtained, about
95% sucrose, is brown in color and impure. Refined sugar, about 99% pure,
is obtained by redissolving, decolorizing the solution with charcoal, and
recryslallizing.

Lactose occurs in the milk of mammals (3-5%). It is composed of a unit
of glucose and one of galactose, an isomer of glucose. Lactose is obtained
as a by-product in the manufacture of cheese. Maltose is present in beer
and in com syrup, and is composed of two units of glucose. Both lactose
and maltose reduce Fehling’s solution.

5. Photosynthesis. The term photosynthesis refers to a process whereby
green plants convert C0 2 and HoO into carbohydrates, utilizing the radiant
energy of the stm and in the presence of chlorophyll as a catalyst. The net
reaction is

(7) n C0 2 + n H 2 0 -^[C(H 2 0)]„ +n0 2

The reaction is highly endothermic, 112 kcal being required for the reaction
of one mole of CO a . The initial step in photosynthesis is the absorption of
a quantum of radiant energy by the chlorophyll molecule, a complex organic
molecule containing magnesium ion, C r ,5H 72 N 4 0 5 Mg. 1 Chlorophyll is green
because it absorbs the red and orange wavelengths of sunlight permitting
green light to be transmitted or reflected. The absorbed radiant energy
is tranformed into chemical energy by a complex mechanism, the complete
understanding of which has so far defied the most brilliant chemists.
Photosynthesis is thus a mechanism whereby solar energy is stored as chemical
energy in the form of a carbohydrate molecule. During the combustion
of a carbohydrate to C0 2 and H 2 0, as in the burning of the cellulose of wood
or in the digestion of food, the reverse of Equation 7 takes place and the
chemical energy stored in the carbohydrate molecule is released. Through
the process of photosynthesis, carbohydrates form the ultimate source of
energy for plant and animal life. It has been estimated that, every second,
five million calories of energy are stored as carbohydrates by plants. Photo-
synthesis is but one of a large class of photochemical reactions which are
either initiated or maintained by means of radiation.

6. Polysaccharides. Polysaccharides are polymers composed of mono-
saccharide units linked through oxygen atoms to form chains of various
lengths. The two most important polysaccharides, starch and cellulose, have
the formula (C«H 10 O 5 ) fl . For starch the value of n is about 300 to 400
whereas for cellulose n is over 1500. The molecular weight” of cellulose
molecule may be from 250,000 to more than 1,000,000. The structure of
starch is shown in Figure 34.3.

Chlorophyll has been synthesized in the laboratory by Professor Robert B, Woodward
of Harvard University and independently by a group of chemists in Germany.

460

Organic Chemistry 111: Polymers and Biochemistry

Starch is found abundantly in grain, tubers, and fruit. It forms the
reserve food supply of seeds and plants, and is the cheapest and most abund-
ant of human foods. Starch occurs as granules which are insoluble in cold

ch 2 oh

\?A

— O . C

\5 i ?/*

°\F

HoV*

C—

ffi-O-

vV \° H

~c / ! x c—

J/OH

1 1

• H

OH ! H

ch 2 oh

Figure 34.3. The Structure of Starch.

water. When treated with hot water the starch granules burst their surround-
ing membranes and form a colloidal dispersion known as “starch solution.”
A sensitive test for starch is the intense blue color produced when a dilute
solution of iodine is added to it The complete hydrolysis of starch (and also
cellulose) yields the monosaccharide, glucose. Controlled hydrolysis of starch
yields, in turn, dextrins, maltose, and ultimately glucose. Corn syrup is a
mixture of these. Dextrin can also be prepared by heating starch, previously
moistened with HC1, to about 200 °C. Dextrins are gum-like substances,
soluble in water, which are used as adhesives, particularly on postage stamps
and envelopes. Besides its direct consumption as a food starch has many indus-
trial uses. Hundreds of modifications, based on a difference in source and pre-
paration (grinding and heat treatment) are used as adhesives, in the sizing of
cotton and paper and in the laundry industry to give strength and stiffness
to the final product.

Glycogen, or animal starch, is stored in the liver and muscles as a reserve
food supply. Muscle glycogen breaks down into lactic acid, the oxidation
of which provides the energy for muscular activity.

The most abundant of all carbohydrates is cellulose, so named because
it* forms the walls of plant cells and makes up the bulk of the framework
and woody structure of a plant. The long chain cellulose molecules lie side by
side, held together by hydrogen bonding through the many OH groups present,
to form rope-like structures. Bundles of these molecules form the elements
of plant fibers. In some plants, notably cotton and flax, the cellulose fibers
are sufficiently long so that they can be mechanically twisted, spun into
thread, and woven into textiles. Woody plants consist of relatively short
cellulose fibers; paper is a mat of these fibers interlaced. Cellulose is a
stable compound, resistant to the action of most simple solvents and chemical
reagents. Because of this property it can be used as filter paper, which

Organic Chemistry III: Polymers and Biochemistry

461

is practically pure cellulose, while because of its hydroxyl groups cellulose
can react as does an alcohol to form esters. For a great many industries
cellulose is the raw material.

7. Paper. The best varieties of paper are made of rags, preferably linen
while the cheaper grades are made from wood pulp. The wood is cut into
chips and heated with a solution of calcium hydrogen sulfite, Ca(HS0 3 ) 2 .
This dissolves the gummy material known as lignin , which, together with
cellulose, makes up the solid part of the wood. The pulpy material is then
washed, beaten to reduce it to shreds, and bleached with chlorine. The
mass, consisting of a suspension of cellulose fibers in water, is spread on
screens, drained, pressed, and dried, forming a sheet of interlaced fibers of
cellulose. During the process other substances are usually added to impart
special properties to the finished product. Thus glue or gelatine (size)
prevents ink from running; clay and barium sulfate (fillers) give body to
the paper and make possible the production of a smooth surface by running
the sheet between the rollers of the paper machine. Colored fillers are
added to the pulp when special tints are required.

When cellulose, e.g., cotton, is treated with a mixture of HN0 3 and
H 2 S0 4 , the ester, cellulose nitrate, is formed. The degree of nitration, which
depends upon the concentration of the acids and the duration of the process,
determines the properties and uses of the end product. Completely nitrated
cellulose has three nitrate groups per glucose unit, C 6 Ht 0 2 (0N0 2 )3, and
contains 14.7% nitrogen. Guncotton , often called cellulose trinitrate, is almost
completely nitrated cellulose. Lower nitrates of cellulose containing 2-3 nitrate
groups per glucose unit (about 11% nitrogen), are collectively called pyroxylin.
Collodion is a solution of pyroxylin in a mixture of alcohol and ether, which
upon evaporation leaves a transparent film. Celluloid is a mixture of pyroxylin
and camphor; with a little alcohol it forms a plastic mass which can be
molded into any desired shape and will then harden into permanent form.
Pyroxylin varnishes and lacquers are made by dissolving pyroxylin in various
solvents, such as methyl alcohol, acetone, and amyl acetate.

Cellulose acetate is an ester made by treating cellulose with the anhydride
of acetic acid. It is used in making safety photographic film because it
does not burn rapidly, nor form noxious gases as does cellulose nitrate upon
combustion. It is also used in making textile fibers, as Celanese and Acetate .

Cellulose, either as wood pulp or cotton, reacts with NaOH and CS 2
to form a viscous yellow liquid, cellulose xanthate or viscose . This liquid is
soluble in water, but reprecipitates cellulose when the solution is acidified
with H2SO4. If the solution of viscose is extruded through minute holes into
an acid solution, cellulose is regenerated as solid filaments. These filaments
are then twisted together to form a thread that is used in making fabrics.
This is the viscose process for the manufacture of Rayon. Cellophane is
made in the same manner, but the viscose is forced through a narrow slit
so that the cellulose is precipitated in the form of thin sheets instead of
as threads.

8. Fats. Fats are esters of the trihydroxy alcohol, glycerol, C 3 H 5 (OH) 3 ,
and straight chain carboxylic acids.

462

Organic Chemistry III: Polymers and Biochemistry

r_0-O-CH 2

R— c_o— ch 2

R_C-0-CH 2

The acids from which the fats are derived are the so-called 'Tatty acids”
With few exceptions these are straight chain carboxylic acids containing
an even number of carbon atoms. Some of the more important fatty acids
are: lauric acid, CH 3 (CH 2 ) 10 COOH; palmitic acid, CH 3 (CH 2 K*COOH; stearic
acid, CH 3 (CH,) 16 COOH; oleic acid, cis-CH 3 (CH 2 ) 7 CH=:CH(CH 2 ) 7 COOH.

Naturally occurring fats are mixtures of several different esters of gly-
cerol. Fats derived mainly from saturated acids are usually solid whereas
those from unsaturated acids are liquid fats, generally called oils. Hydro-
genation of the double bonds in the unsaturated liquid fats “hardens” them
into solid fats. In this way, cheap fats such as cottonseed oil and soy bean
oil, which were formerly discarded as waste products, are converted into
solids having the consistency of butter, e.g., oleomargarine, Crisco, and Spry,

The hydrolysis of a fat or oil by boiling it with a strong base such as
NaOH or KOH yields glycerol and the sodium or potassium salts of the
fatty acids of the fat or oil. The process is known as saponification and the
sodium or potassium salts are soaps. With glyceryl stearate, or stearin,
(C 17 H 35 COO) 3 C 3 H 5 , the saponification reaction is

(8) (C 17 H 35 COO) 3 C 3 H 5 + 3 NaOH -» 3 C 17 H 35 COONa + C 3 H 5 (OH) 3

glyceryl stearate sodium stearate glycerol

Glycerol is always a by-product of soap manufacture; indeed most glycerol
is produced by this process.

Soap is a specific term— a soap is the salt of a high molecular weight
fatty acid. Soap does not include the entire class but is merely one member
of the larger class of detergents. To refer to any detergent as a soap is akin
to speaking of a “rubber cork.” Ordinary soap is a sodium salt which is soluble
in water but salts of metals other than the alkali metals form insoluble soaps,
e.g. 4 calcium stearate. The cleansing action of soap is discussed in Chapter 36
and the action of soaps with hard water in Chapter 39. Most commercial
detergents are sodium salts of monoalkyl sulfates having the formula,
R O SO 2 — ON a, in which the alkyl radicals vary from C 7 Hi 5 to C 18 H 37 .

9. Proteins. Proteins are essential to life processes and are probably
the most important of all substances found in plant and animal life. In
animals, proteins make up the structural framework much as cellulose does

Organic Chemistry III: Polymers and Biochemistry

463

in plants. Their chief function, however, is to build and to repair body
tissue since only proteins have the capability of reproduction and growth.
There are many thousands of different proteins, perhaps as many as 50,000
in the human body, each having a structure which enables it to accomplish
a specific function. Muscle, nerves, hair, and skin are composed principally
of proteins.

All proteins are nitrogenous substances containing about 16% nitrogen
in addition to carbon, hydrogen, and oxygen. Certain proteins contain small
amounts of other elements such as sulfur (egg albumin), phosphorus
(casein), and iron (hemoglobin). Determination of the molecular weights
of proteins yields values ranging from about 12,000 for insulin to many mil-
lions for tobacco mosaic virus. When proteins are hydrolyzed by heating
them with aqueous solutions of acids or bases, they break down into amino
acids. These aminoacids constitute the building blocks of protein mole-
cules. Protein molecules are polymers formed by the splitting out of a
molecule of water in the reaction of the — COOH group of one amino acid
with the — NH 2 group of another.

HHO HHO HHO HHOHHOHHO

II II I I II r - I 1 II I I II I I II I I II

(9) H— N— C— C-j-OH H-j-N— C— C^OH H-l-N— C— C—OH -* ~N-C-C-N-C-G-N-C-C—

i A A A k A

O H

The grouping — C— N— is known as a peptide linkage and hence a protein
molecule is a polypeptide.

Twenty-six different aminoacids have been isolated from the hydrolysis
of proteins. Though a single protein may yield several aminoacids, no one
protein gives all twenty-six. Because protein molecules can differ in the
number, the kind, and the arrangement of the aminoacids, the number of
possible protein structures is statistically very large.

The geometric shape of the protein molecule is the subject of intensive
research. X-ray analysis indicates that, structurally, proteins can be classi-
fied broadly either as fibrous proteins or as globular proteins. In the mole-
cules of fibrous proteins the polypeptide chains have a coiled spring and
helical configuration and hydrogen bonds between different parts of the
same chain keep the helix in its coiled position. Several chains may lie
parallel to each other and be intertwined to form a rope-like structure.
Fibrous proteins are insoluble in water. Fibroin . , in silk, and keratin , in hair,
skin, wool, and nails, are proteins of this type. The globular proteins have
a spherical shape. Segments oi the helical chain are folded back upon each
other to give a random arrangement. The globular proteins are soluble in
aqueous solutions. They are concerned primarily with the regulation of life
processes, a function that requires solubility in body fluids. Examples of
globular proteins are insulin and hemoglobin.

464

Organic Chemistry III: Polymers and Biochemistry

It may be that it is not the proteins themselves which are essential to
the diet but the aminoacids derived from their hydrolysis. These amino-
acids are carried by the blood to various sites in the body where they can
be used as building blocks for the manufacture of specific protein molecules
required by the body. Of the twenty-six aminoacids obtained from pro-
teins only nine are essential to man in that these nine cannot be synthesized
by man’s metabolic processes, whereas the others, the so-called nonessential
aminoacids, can be manufactured by mans biochemistry. Several of the
essential aminoacids can be synthesized in the laboratory.

Except for glycine all aminoacids contain at least one asymmetric carbon
atom and all are optically active. It is a curious fact that only the levoamino
acids are found in nature; the dextro enantiomorphs can be obtained only
by synthesis. The chemistry of life on this planet is built on a system of
levoamino acids. The dextro forms cannot be assimilated by man and, since
their reactions towards other optically active substances differ from those
of the levo forms, may even be poisonous.

10. Vitamins. The vitamins are accessory food materials which are neces-
sary for the nutrition of man. The term vitamins was coined in recognition
of the fact that their absence from the diet resulted in certain specific diseases.
Vitamins are specific chemicals and it was only a few years ago that the
techniques of organic chemistry were able to determine the exact chemical
nature of some vitamins. Thus vitamin A, lack of which results in a malfunc-
tion of the eyes, has the structure

H 3 C CHs CH* CHs

i i

C-CH=CH-C=CH-CH=CH-C=CH-CH 2 OH

X /

ch 2

c-ch s

Certain inorganic elements are also necessary for good health. Among
the common elements found in tissue are Na, K, Ca, Fe, Mg, S. P, and the
halogens but certain rarer elements such as Cu, Mn, Zn, and Co are also
required in minute quantities. In many cases the specific function of these
trace elements is unknown.

11. Hormones. In addition to vitamins, the body requires hormones in
order to function properly. Although vitamins are obtained in the food,
hormones are synthesized within certain organs of the body— the .endocrine
or “ductless” glands— and pass directly into the blood. The exact number of
hormones produced in the body is not known, but several have been identified
and their functions ascertained. Thyroxin , CisHuCXNI^ is produced by the
thyroid gland. A lack of this hormone causes goiter and cretinism, a disease
producing deformity of the body and impairment of the function of the
brain. Adrenalin , [C 6 H 3 ( OH ) 2 CH ( CH 2 NHCH 3 ) OH] , is produced by the
adrenal glands and acts as a stimulant of the heart and lungs. Insulin, a pro-
tein in chemical composition, is produced in the pancreas and aids in the
metabolism of sugar. The lack of sufficient quantity of this hormone causes

Organic Chemistry III: Polymers and Biochemistry

465

the disease known as diabetes , in which the glucose in the blood and the
urine increase abnormally.

12. Enzymes. The combustion of sugar to carbon dioxide and water can
be readily carried out on the laboratory bench but a high temperature is re-
quired to initiate and to maintain the reaction. In the human body the
oxidation of a sugar such as glucose proceeds quite rapidly at a temperature
of about 37 °C due to the presence of organic catalysts or enzymes. Enzymes
are proteins which act as catalysts for the metabolic reactions taking place
in a living organism. It has been estimated that there are over 20,000 different
enzymes in the human body, each a highly specific catalyst for a particular
biochemical reaction. In the stomach the enzyme pepsin catalyzes the
hydrolysis of proteins into aminoacids whereas lipase catalyzes the decomposi-
tion of fats. Though starch and cellulose are identical in that they are con-
structed from the same monosaccharide, glucose, enzymes in the human
body are able to catalyze the hydrolysis of starch but not of cellulose. More
practically, human beings can digest starch but not cellulose.

An enzyme molecule contains a special group, or coenzyme , attached to
the protein portion of the molecule, or apoenzyme ; the substance whose
reaction is catalyzed by an enzyme is known as the substrate. Neither the
separate coenzyme or apoenzyme exhibits catalytic activity. A factor in
the specificity of an enzyme for a substrate is the geometric shape of their
molecules, and it is believed that they must fit together in “lock and key"
fashion. It is for this reason, perhaps, that the human body can tolerate
only levoamino acids. Many enzymes have been isolated in crystalline form
but the molecular structure of no enzyme has so far been determined, nor
has the definite mechanism of any enzymes operation. This remains one
of the most important biochemical problems.

QUESTIONS

1. Define and illustrate the following: monomer; polymer; copolymer; isomer.

2. Distinguish between addition polymerization and condensation polymerization.
Write a reaction which illustrates each method,

3. What method of polymerization is involved in the formation of (a) rubber
from isoprene and (b) protein from aminoacids?

4. Draw th^ structural formula for two units of a polymer formed from
NH 2 (CH 3 ) 3 NH 2 and HOOC(CH 2 ) 2 COOH.

5. What is meant by a three dimensional cross-linked polymeric structure?
What effect does cross-linking have on the properties of a polymer?

6. How can peroxides act as initiators of polymerization?

7. What is the difference in structure and in properties of thermosetting and
thermoplastic polymers?

8. From what monomer could the polymer, Saran , — CH 2 CC1 2 CH 2 CC1 2 — , be made?

9. Which of the following reactions could theoretically lead to polymer formation
(a) ethanol and acetic acid (b) ethylene glycol and acetic acid (c) ethylene
glycol and oxalic acid (d) glycerol and phthalic acid?

10. Distinguish among carbohydrate; cellulose; starch.

466

Organic Chemistry III: Polymers and Biochemistry

11. Define and give an example of each of the following: monosaccharide; dis-
accharide; polysaccharide; aldose; ketose; hexose.

12. Write reactions for (a) the hydrolysis of sucrose (b) the hydrolysis of a fat
(c) the condensation of three molecules of an aminoacid to form a polypeptide.

13. What is the net chemical reaction in photosynthesis? What is the function
of chlorophyll? What is the source of the reaction energy?

14. Would a solution of sodium stearate be acidic, basic, or neutral? Explain,

15. What is the chemical composition of a protein? What are the two broad
classes of proteins? To which would an enzyme belong?

16. A favorite building block in nature is the isoprene molecule. Draw lines on
the structure of the Vitamin A molecule which subdivide it into isoprene
units.

17. Distinguish among enzyme, hormone, and vitamin.

The Elements of Group IVB

Silicon

After oxygen, silicon (Latin, silexi flint) is the most abundant element
comprising about 25.8% of the earths crust. Excluding the seas and the
atmosphere, the solid part of the earth's crust is composed primarily of silica
Si 0 2 , and silicates, compounds which contain a — O— Si— O— structure. Metallic
minerals and ores, other than those containing silicon, constitute but a small
part of the earth s crust. In fact silicon compounds play a role in the mineral
world similar to that of carbon in the plant and animal world. Most common
rocks and clay, except for limestone, CaC 0 3 , and dolomite, CaC0 3 * MgCO ;i ,
contain silicon; sand is largely SiO z . Silica , Si0 2> is found as quartz, flint, and
in the semiprecious gems, amethyst, opal, and onyx; any color in these is due
to impurities. The compositions of some common minerals are given in
Table 35-A.

Table 35-A

Common Mineral Silicates

Feldspar

KAlSi s 0 8

Mica

KAl 3 Si 3 O 10 (OH) j

Kaolin (clay)

A1 2 S1 2 0 5 (0H ) 4

Zeolite

Na 2 Al 2 Si 3 O 10 * H 2 0

Garnet

Fe 3 Al 2 ( Si 0 4 ) 3

Asbestos

Mg,CaSi*0 12

Zircon

ZrSi0 4

Talc

MfeSi 4 O 10 (OH ) 2

Even in the organic world small amounts of silicon are found in many
plants. Oat hulls and the hard, lustrous portions of straw and bamboo
contain as much as 40 percent Si0 2 , while the skeletons of some marine
organisms, sponges and diatoms, are mainly Si0 2 . The latter can be found
in beds hundreds of feet thick known as diatomaceous earth.

1. Preparation of Silicon. Elemental silicon exists in two allotropic forms;
a gray, crystalline, semimetallic variety in which the silicon atoms have a

468

Elements of Group IVB; Silicon

tetrahedral arrangement, and a brown, amorphous form. Gray silicon is brittle
but hard. It can scratch glass and has a hardness of 7 on the Mohs scale,
in which diamond, the hardest known substance, has a hardness value of 10.
Silicon can be prepared by the following methods:

(A) Reduction of Si0 2 by a strong reducing agent such as Mg or A1

(1) SiO a + 2 Mg -» Si + 2 MgO

Silicon produced in this way is a brown amorphous powder. Some magnesium
silicide, Mg s Si, is produced simultaneously by a side reaction.

(2) Si0 2 + 4Mg-4 Mg 2 Si + 2 MgO

(B) Reduction of Si0 2 by carbon in an electric furnace

(3) SiO a + 2 C Si + 2 CO

This method produces the gray ‘crystalline variety of silicon. The process is
similar to that employed for making silicon carbide, SiC, except that smaller
quantities of carbon are used so that pure silicon is produced in greater
proportion. If Fe 2 0 3 is added to the charge an industrially important alloy
of iron and silicon, called ferrosilicon, is obtained.

2. Properties and Uses of Silicon. Silicon is the second element of Group
IVB. Its properties are listed in Table 31-A. At ordinary temperatures silicon
is unreactive. It is unaffected by air, water, or acids excluding HF, but does
combine with F 2 at low temperatures and with the other halogens and 0 2 at
higher temperatures to form the respective binary compounds, SiX 4 and Si0 2 .
With many metals at high temperatures, silicon combines to form silicides
such as Mg 2 Si. Although it does not react with acids, silicon reacts with cold
NaOH or KOH, an indication of its nonmetallic character.

(4) Si + 2 OH~ + 2 H 2 0 Si(V~ + 2 H 2

Despite its abundance the properties of silicon are such that it has few
industrial uses. Its major use is in the steel industry in the form of ferrosilicon
as a deoxidizer and as an alloying agent to modify the properties of steels.
Duriron is an acid resisting alloy containing 14 percent Si. Pure silicon is* a poor
conductor of electricity but its conductivity increases with a rise in tempera-
ture, a behavior contrary to that of metals. Such substances, which include
boron and germanium, are called semi-conductors and their main use is
in transistors and electric rectifiers.

3. Silicon Hydrides. Compounds of silicon and hydrogen are produced
by the action of acids on metal silicides; the reaction is analogous to that
of acids on metal carbides.

(5) Mg 2 Si + 4 HC1 SiH 4 + 2 MgCl 2

The hydrides of silicon form a series of compounds known as silanes,
similar in formula and covalent structure to the paraffin hydrocarbons. Their
general formula is Si n H 2n _ 2 , but they are much fewer in number than the
hydrocarbons, the highest member having n equal to 8. Silane, SiH 4 , is a

Elements of Group 1VB: Silicon

469

colorless gas with a disagreeable odor. Disilane, Si 2 H 6 , is also a gas but
trisilane, Si 3 H 8 , is a liquid, as are the higher members of the series. The
silanes are less stable than the hydrocarbons, igniting spontaneously in air
and decomposing slowly when in contact with water.

(6) SiH 4 + 2 0 2 -> Si0 2 + 2 H 2 0

(7) SiH 4 + 4 H 2 0 — > 4 H 2 + H 4 Si0 4 ( orthosilicic acid)

4. Silicon Halides. The silicon halides can be prepared by the direct
combination of the elements but more convenient preparations for silicon
tetrafluoride, SiF 4 , and silicon tetrachloride, SiCl 4 , are:

(A) The action of HF on Si0 2 or a silicate

(8) Si0 2 + 4 HF -» SiF 4 + 2 H a O

(B) Passing Cl 2 over a mixture of Si0 2 and C at high temperature

(9) Si0 2 + 2 C + 2 Cl 2 -> SiCl 4 + 2 CO

At room temperature SiF 4 is a colorless gas and SiCl 4 a colorless liquid. Both
react vigorously with water.

(10) 3 SiF 4 +.4 H 2 0 — > H 4 Si0 4 -I- 2 H 2 SiF 6 (fluosilicic acid)

(11) SiCl 4 + 4 H.O H 4 Si0 4 + 4 HC1

Fluosilicic acid, H 2 SiF 6 , is a strong acid, soluble in water and stable only
in solution. If heated it decomposes into HF and SiF 4 . Orthosilicic acid,
HjSi0 4 , is insoluble in water.

5. Silicon Dioxide. Commonly known as silica , silicon dioxide, Si0 2 , oc-
curs naturally as quartz, a colorless, hard, and brittle solid. Pure quartz
melts at about 1600 °C to form a viscous liquid. Upon cooling, the liquid
does not readily recrystallize but supercools to form a glassy substance known
as "quartz glass.” This supercooled liquid can be shaped into various forms
such as beakers, flasks, and tubes. Vessels made of quartz have many
advantages. The material is less soluble than glass and has a much smaller
coefficient of thermal expansion, permitting it to undergo sudden changes
in temperature without breaking. A quartz tube can be heated to red heat
and plunged into water without fracture whereas under the same conditions
ordinary glass would shatter. Unlike glass, which is opaque to ultraviolet
radiation, quartz is transparent to both visible and ultraviolet light. Hence
it is used in optical apparatus where- ultraviolet transmission is a requisite,
for example, as the envelope of mercury vapor lamps used as a source of
ultraviolet rays.

It is interesting to compare the structures of Si0 2 and C0 2 . In solid
C0 2 the structural unit is the C0 2 molecule. These separate C0 2 molecules
are held together by weak intermolecular, Van der Waals forces. Relatively
little energy is required to separate one molecule from another and so solid
C0 2 is highly volatile and has a low melting point. In contrast, the unit of
structure in Si0 2 is not an Si0 2 molecule. In a crystal' of Si0 2 , the silicon
and oxygen atoms are arranged in tetrahedral, fashion; a silicon atom is

cated at the center and an oxygen atom at each of the four corners of a
gular tetrahedron. These SiO* tetrahedra are joined together through the
aring of oxygen atoms at the corners of two adjacent tetrahedra. This
Si„0_ S i- network forms a three dimensional structure which extends
roughout the entire crystal, which may therefore be regarded as a macro-
olecule. The network consists of strings of tetrahedra, which may be strung
ie by side or may be linked in spiral fashion, giving rise to left-handed
right-handed quartz, which rotate the plane of polarized light to the
ft or to the right, respectively. (Figures 35.1 and 35.2) Since the bonds
; tween the silicon and oxygen atoms are single covalent bonds, relatively
tong and not easily broken, Si0 2 has a high melting point and is a hard
ibstance.

(a) Right-handed < b > Left-handed

Figure 35 . 1 . Crystals of Natural Quartz.

Silicon atom
Oxygen atom

° ,

' i» \

si \

/Si O

-i ;>o

O I n O 1 o

-O' -n

Si y

O x Si"

\i| \

Si \

1 !

Si 0

- 4>:0

o x

■•O' N

si"

° h 7

i i

! v O
O

' I x o

1 ^
\ Si/

/Si O
o I N o , o
O

X Si X

1/ IN

Si'

(b) Parallel strings of tetrahedra
Figure 35.2 . Structure of Quartz.

6. Silicic Acids. Though it is insoluble and does not react with water,
Si0 2 may be regarded as the anhydride of many silicic acids which can be
represented by the formula x Si0 2 • y H 2 0. Orthosilicic acid, H 4 Si0 4 , is pre-
pared by treating an aqueous solution of sodium silicate, Na 2 Si0 3 , with HC1.

(12) Na 2 Si0 3 + 2 HC1 + H*0 H 4 Si0 4 + 2 NaCl

If the acid is added to the solution of sodium silicate, a colloidal dispersion and
then a jelly-like precipitate of silicic acid is formed within a few seconds. If the
silicate is poured into the acid, however, no precipitation occurs; the silicic acid
remains in colloidal dispersion.

By successive dehydration of fI 4 Si0 4 many other silicic acids may theoretically
be formed. Although all are hypothetical they are frequently mentioned in
the chemical literature. Salts of these acids exist, however, and many are
found in nature as minerals.

(13)

-H 2 0

H 4 Si0 4 ► H 2 SiO„ (metasilicic acid)

-H 2 0

* Si0 2

Incompletely dehydrated H 4 Si0 4 , containing about 6% water, is called silica
gel. Because it has a porous structure with a large surface area, silica gel
is employed as a drying agent, as a decolorizing agent, and as an adsorbing
agent ror organic solvents and Valuable vapors such as gasoline.

Series of polysilicic acids are hypothetically possible. Disilicic acids and
trisilicic acids can be considered as the products of the dehydration of two
and three molecules of H 4 Si0 4 , respectively.

(14)2H 4 Si0 4 H 6 Si 2 O v — > H 4 Si 2 0g — > H 2 Si 2 0 5 — > 2 Si0 2

(15) 3 H 4 Si0 4 H 10 Si 3 O 13 — ^ HsSiaOio—^ H 3 Si 3 0£>

H 4 Si 3 0* — » H 2 Si s 0 7 -»3 Si0 2

7. Silicates. The silicates may be regarded as the salts of the various silicic
acids and can be divided into two groups: the soluble silicates and the
insoluble mineral silicates. Only the silicates of the alkali metals are soluble
in water. Of these, sodium (meta)silicate. Na 2 Si0 3 , is the most important.
It is made by fusing sand with Na 2 C0 3 or by boiling sand with NaOH.

(16) Si0 2 “j" Na 2 C0 3 — ^ Na 2 Si0 3 -j- 0O 2

(17) Si0 2 + 2 NaOH Na 2 Si0 3 + H 2 0

Pure Na 2 Si0 3 is a glassy solid but it is sold commercially as a concentrated
aqueous solution, a viscous, syrupy liquid known as water ghss. The solution
is used for fireproofing timber and textiles, for preserving eggs, and, because

it gives a basic reaction by hydrolysis, as a filler in cheap soaps and washing

powders. The formula Na 2 Si0 3 is an oversimplification and a more correct
one would be Na 2 0 * x Si0 2 ; a solution of water glass probably contains
molecular species such as Na 2 Si 3 0 7 , etc.

All the mineral silicates are characterized by a tetrahedral structure
wherein a central silicon atom is joined by single covalent bonds to four
oxygen atoms. The structure of a silicate crystal is a soft of fabric of these
Si0 4 tetrahedra. The tetrahedra may exist as independent units or they

472

Elements of Group IVB; Silicon

may be joined by sharing one or more of the oxygen atoms at the corners
of the tetrahedra so that oxygen bridges are formed between silicon atoms.
Different structures result depending upon whether no, one, two, three, or
four oxygen atoms of an Si0 4 tetrahedron are bonded to the silicon atoms
of neighboring tetrahedra. It is this difference in the mode of linking the
tetrahedra which results in the different silicon-oxygen atom ratios in the
various silicate formulas.

The cause of this behavior is the fact that the total number of valence
electrons available to one silicon atom and four oxygen atoms, 28, is in-
sufficient by 4 to complete an octet around each atom, assuming that only
single covalent bonds are formed. The sharing of oxygen atoms reduces this
insufficiency; the remaining electrons which are required are supplied by
the metal atoms of the silicate compound. Where no oxygen atoms are shared
between SiO* tetrahedra there results the simplest of all silicates, the ortho-
silicates, in which there are individual SiO.* 4 ' ions. (Figure 35.3) If two tetra-

An extended silicate chain

(d)

Figure 35.3. Structures of Silicate Ions.

hedra are joined at one corner by the sharing of a single oxygen atom,
the unit of the silicate structure becomes the Si 2 0 7 ,J ~ ion. Three tetrahedra
can share oxygeii atoms and form a ring structure as in the Si 3 0 9 6 - ion.
This ion and the analogous ring structures of the Si 4 Oi 2 8 “ and Si 6 Oi 8 12 ~ ions
all have the empirical formula of the simple metasilicate ion, Si 2 0 3 2 ~, so
that their formulas are sometimes written ( Si 2 0 H ) n ~ 2n . All these intricate
ions are discrete and are held together within the crystal through electro-
static attraction between them and the positive metal ions of the silicate

Elements of Group IV B: Silicon

473

compound. The nature of the binding force here is the same as that in a
simple crystal of NaCl.

In a ring structure each Si0 4 tetrahedron shares two of its oxygen atoms,
one atom with each of the two neighboring tetrahedra. A similar bonding
can exist in an open ring or extended chain structure wherein each tetra-
hedron, except the two end ones, also forms two oxygen bridges. Such an
extended chain often results in a fibrous structure; asbestos has* this type of
structure. In the case of three oxygen bridges per silicon atom, a two di-
mensional sheetlike structure, as in talc or in muscovite mica, is produced.
With four oxygen bridge atoms, a three dimensional framework results;
examples of this are feldspar, the zeolite rocks, and indeed silica itself.

Frequently some of the silicon atoms of the three dimensional network of
silicate anions are replaced by aluminum atoms. Since aluminum is trivalent
and silicon is tetravalent, an additional univalent ion enters the lattice
structure to maintain electric neutrality. The formula for feldspar, KAlSi a 0 9 ,
is an example.

Because the covalent bonds holding the anions together are strong, the
silicate minerals have high melting points, and because of an open and

porous structure relatively low densities. It is interesting to conjecture

that, because of these properties, as the liquid magma of the earth cooled
the silicates rose to the surface and crystallized early to form the crust
upon which we live.

8. Glass. The earliest examples of man-made glass are glazed beads
found in Egypt and believed to have been made over 14,000 years ago. The
first glass containers, jars for cosmetics, were made in Egypt about 2,000 b.c.
Glass is a mixture of silicates containing an excess of Si0 2 . In the sense that

it does not have the ordered internal structure characteristic of crystals,

glass is not a solid. It is a supercooled liquid of extreme viscosity; it does not
have a definite melting point but softens gradually upon heating.

Ordinary glass is a mixture of sodium and calcium silicates having an
approximate composition of Na a O; CaO; 6 Si0 2 . It is made by fusing together
sodium carbonate (soda ash), Na 2 CO d , limestone, CaC0 3 , and sand, Si0 2 ,
at temperatures over 1,500 °C.

(18a) Na 2 CO,i + SiO, Na 2 SiO d + C0 2

(18b) GaCOa -j“ Si0 2 — > CaSiOy -j- C0 2

When all the CO a is expelled a clear melt results. This clear, viscous material
can then be poured into molds or stamped with dies to produce pressed
glass ware. Bottles, beakers, and similar objects are made by taking up a
sufficient quantity of the molten glass on an iron pipe, inserting it into a
mold, and blowing till the outline of the mold is filled. Automatic machines
can produce bottles at the rate of 250 a minute. To make window or plate
glass, liquid glass is poured out upon a flat table and rolled into a sheet
which is then ground flat and polished on both sides. Glass that has been
allowed to cool rapidly is brittle and likely to fracture on receiving a scratch
or a slight shock. Glassware is therefore annealed by heating it for some
time at a temperature just below its softening point, and then cooling at a

474

Thicmcnts of Group TVB .* Silicon

slow rate. Laminated safety glass is made by cementing together two sheets
of glass, between which is placed a layer of transparent plastic. Upon impact
the flexible center layer holds any broken fragments in place.

9. Varieties of Glass, Soda-lime glass ( Si0 2 , 75%; Na s O, 15%; CaO, 8%;
Al 2 0 3 , 2%) is the ordinary glass of commerce and is called soft glass because
it softens at a relatively low temperature and is easily worked into differ-
ent shapes. Soda-lime glass is slowly attacked by chemical agents and even
by water. It is used for window panes, bottles, glass bricks, and glass
fibers. Hard glass, of higher melting point, is made by substituting K 2 C0 3
for the Na 2 C0 3 and is used for making chemical apparatus which is to with-
stand higher temperatures.

Borosilicate glass (Si0 2 , 80%; Na 2 0, 4%; K 2 0, 0.6%; CaO, 0.4%; B 2 0 3 ,
12%; A1 2 0 3 , 3%), of which Pyrex is an example, is used for chemical glass-
ware and baking dishes. It resists the attack of chemical reagents and has
a small coefficient of thermal expansion.

Flint glass or lead glass (Si0 2 , 45.5%; Na 2 0, 3.5%; K 2 0, 4%; CaO, 3%;
PbO, 44%) is used for optical work and decorative purposes (“cut” glass)
because of its high index of refraction.

Glass is colored by fusing oxides of metals which yield colored silicates
into the liquid mass. Iron and chromium produce a green glass; cobalt and
copper, blue; and manganese, violet or amethyst. Ruby glass contains finely
divided gold or selenium. Milk glass is made by adding CaF 2 , Sn0 2 , or
Ca 3 (P0 4 ) 2 .

10. Clay and Ceramics. The term clay is applied to a mixture of sub-
stances produced by the weathering of silicious rocks containing feldspar,
KAlSi 3 O s . Through the action of water and carbon dioxide, the potasssium
is slowly removed, together with a portion of the silica, and a hydrated
aluminum silicate is formed. When the rock undergoing this process is pure
feldspar the result is a white plastic mineral known as kaolin or china clay ,
Al 2 Si 2 0 5 (OH)4. Most pottery and bricks are made from clay. A mixture
of clay and water is molded into a desired shape, dried, and heated or “fired”
at high temperature in a kiln. Complex and little understood reactions take
place during which the mixture acquires a rigid structure. Clay usually
contains an impurity of Fe 2 0 3 which imparts a red tint to the final product.
Earthenware such as tile, brick, stoneware jugs, etc., is made from impure
clay while china and porcelain are made from pure clay, free from iron.
A glaze can be produced by spreading a paste of finely ground feldspar
and silica upon the surface of the china or porcelain and firing again at a
higher temperature. Colored decorations are obtained by applying the oxides
of suitable metals to the glaze and firing for a third time, whereupon the
oxides form colored silicates.

11. Silicones. In recent years chemists have synthesized a class of com-
pounds called silicones . These contain polymeric — Si— O— Si— chains with
organic radicals attached to the silicon atoms. As with organic molecules,
the silicones have linear., cyclic, or cross-linked structures. The structure
of a linear silicone polymer is

Elements of Group TVB: Silicon

475

R R R

R— Si— O— Si -O-Si-R

i i i

B E R

The structure is analogous to a metasilicate chain, in which two oxygen
atoms have been replaced by organic radicals, R. In the simplest silicone,
R is a methyl radical.

The silicones can be prepared by the hydrolysis of organosilicon chloride
compounds of the type, R 2 SiCl 2 , followed by polymerization through the
elimination of H 2 0.

(19) R 2 SiCl 2 + 2 H 2 0 -> R 2 Si(OH) 2 + 2 HC1

R R R

(£0) HO— SiiOH-H | O— Si-[OH— H‘ O-Si-OH -»

I I i

R R R

I ! I

HO-Si-O-Si-O-Si-OH + 2 H 2 0

i I I

R R R

Physically, the silicones may be oils, greases, or resins, depending upon

their structural complexity. A linear molecule with about 10 silicon atoms
in the chain would be an cil while a cross linked structure would yield a resin-
ous polymer. The silicones have unusual chemical and physical properties. They
are chemically inert, stable at high temperatures, and water repellent. The
oils have an extremely small temperature coefficient of viscosity and are
useful as lubricants under large variations in temperature. Because of their
good dielectric properties, silicone resins find use as electric insulators.

QUESTIONS

1. Write equations for the preparation of elemental silicon. In what allotropic
forms does silicon exist? Though silicon is the second most abundant element
what properties mitigate its extensive commercial use?

2. How are silanes prepared? Compare the properties of silanes and alkane
hydrocarbons. Why should silicon not form as many hydrogen compounds as
does carbon? How would one attempt to prepare the hypothetical "silanoF?

3. Compare the chemical and physical properties of CO s and Si0 2 . Account
for any similarities and differences.

4. What is the structure of silica? What other substances have a similar struc-
ture? Why can the silica crystal be considered a macromolecule? How do you
account for the fact that a quartz crystal is optically active?

Elements of Group 1VB: Silicon

476

5. Write reactions for (a) preparation of water glass (b) H 4 SiO, (c) dehydra-
tion of H 4 Si0 4 .

6. In the fluosilicate ion, SiF/-, six fluorine atoms are symmetrically spaced
about a central silicon atom. Draw the spatial and electron dot structures.
What orbitals are used in bonding?

7. The mineral beryl is Be 3 Al 2 Si 6 0 18 . The Si 6 0 is 12 - ion has a ring structure in
which there are six silicon atoms. Draw the structure of the ion.

8. Compare the structures of the following ions: Si 2 O r fi- ; P 2 0 j 4 ~; SjO^ - .

9. What is the type of structure that causes a fibrous type of mineral such as
asbestos?

10. Account for the (a) easy cleavage of mica and (b) the slipperiness of talc.

11. Write reactions for the manufacture of glass. How does the structure of glass
differ from that of quartz? In what properties do quartz and glass differ?
What characteristics does glass have that make it so cheap commercially?

12. Why does Pyrex glass withstand thermal shock whereas soft glass shatters?

13. Write the reaction for the preparation of a silicone. Draw the structure of the
linear polymer, (CaHj^jSijO.,, and discuss its probable properties.

14. Draw the probable structure of a cross-linked silicone resin. How would
the structure of a silicone oil differ from that of a solid resin?

Colloid Chemistry

Solutions have been distinguished from suspensions by the size of the
dispersed particles. In a solution the dissolved particles, which may be
atoms, molecules, or ions, are of molecular dimensions whereas in a suspension
the dispersed particles consist of very large aggregates of molecules. A mix-
ture containing dispersed particles intermediate in size between those of a
solution and a suspension is known as a colloidal dispersion , and the dis-
persed substance is called a colloid. Primarily because of their dimensions,
colloids take on special properties which are independent of the chemical
nature of the colloid.

The term colloid (Greek, kolla : glue) was first used in 1861 by the
Scottish chemist, Thomas Graham (Grahams Law of Diffusion of Gases).
Graham chose this name because of his observation that crystalline sub-
stances, which he called crystalloids, yielded true solutions while amorphous
substances, such as starch, gelatine, and glue, gave apparently homogeneous
solutions, sometimes containing a slight uniform turbidity, in which the dis-
persed particles showed but little tendency to diffuse or to settle out. This
distinction between crystalloids and colloids as different kinds of matter
has been dropped. Today the term colloid does not refer to a definite cate-
gory of matter, but rather to particles of any matter whose dimensions lie
within a certain range.

1. Colloidal Dimensions. Particles in colloidal dispersion are intermediate
in size between those in true solution and those in coarse suspension. How-
ever, there is neither a sharp separation in size nor a sharp change in proper-
ties between colloidal particles and particles of molecular dimensions on
one hand and particles of “macroscopic” size on the other hand. Hence the
range of size of colloidal particles has been arbitrarily defined. Particles
with at least one dimension in the range from 10“ 7 cm to 1(H cm are said
to be colloidal. 1 Inasmuch as molecular magnitudes are of the order of 1(H cm,

1 Alternative methods of expressing colloidal dimensions are:

Since 1m (micron) = 0.001 mm or 0.0001 cm, colloidal dimensions are 0.001m to 1 m
S ince 1 mM (millimicron) = 10- 7 cm, colloidal dimensions are 1 mu to 1000 m/i
Since 1 A (Angstrom unit) = 1CH cm, colloidal dimensions are 10 A to 10,000 A

478

Colloid Chemistry

colloidal particles are approximately 10 to 10,000 times larger than simple
molecules. Particles larger than 10' 2 * 4 * * cm are said to be in coarse suspension;
this size is about the limit of the resolving power of the best visible micro-
scope, using blue light. Particles smaller than 10~ 7 cm are considered to be
in true solution; this lower limit is near the limit of visibility of the ultra-
microscope.

The size of a particle alone offers no information as to its constitution.
A colloidal particle may be an aggregation, hundreds or thousands, of atoms
or molecules, or simply a single giant molecule. Many polymers have colloidal
dimensions. Indeed the definition of colloidal dimensions includes fibers
(one colloidal dimension) and films (two colloidal dimensions) as well as
three dimensional solids, liquid droplets, or gaseous bubbles. Provided the
particle falls within the size range associated with colloids it will exhibit
properties characteristic of colloidal dispersion. In Table 36- A are listed
data which may give a better perspective of colloidal dimensions.

Table 36-A

Size Range, A

C— C bond distance in ethane

1.54

Limit of resolution of the electron microscope

Smallest protein molecule

(Mol. Wt. = 10,000)

Hemoglobin molecule

(Mol. Wt. = 68,000)

Limit of resolution of the ultramicroscope

Gold particles in colloidal dispersion

20-200

Globulin molecule

(Mol. Wt. = 160,000)

30 x 30 x 150

Tobacco mosaic virus

(Mol. Wt. = 42 x 10*)

140 x 140 x 2500

Limit of resolution of visible light microscope

2500

Bacterial cells

10000

2. Types of Colloidal Dispersions. The terms colloidal state and colloidal
dispersion are sometimes used. Strictly speaking these are incorrect. There
are only three states of matter— gaseous, liquid, solid— while in solution par-
ticles of molecular dimension only are dissolved. The use of the term
colloidal dispersion is most proper. The substance that is dispersed as particles
of colloidal size is called the dispersed phase and the medium in which
these particles are dispersed is the dispersion medium .

Each of the three states of matter can be dispersed in colloidal sub-
division in media which may be gaseous, liquid, or solid. Within these broad
limits there are four important classes of colloidal dispersions, namely, sols 9
geb, aerosols , and emulsions. The most important of the colloidal systems
are the sob, in which a solid is dispersed in a liquid. Milk of magnesia is

a sol in which solid magnesium hydroxide, Mg(OH) 2 , is dispersed in colloidal

dimensions throughout water. A gel is a special type of colloid in which

the dispersed solid forms a three dimensional fibrous network in whose

inteistices the liquid dispersion medium is retained. Both the solid and

Colloid Chemistry '

479

liquid phases are continuous, giving the gel a jelly-like texture. When dilute
HC1 is added to a dilute solution of sodium silicate, a milky jelly-like
colloidal dispersion of hydrous silicon dioxide is formed slowly; the partial
dehydration of this product yields silica gel. An aerosol *s a colloid in which
a liquid or a solid is dispersed in a gas, e.g., a fog or a smoke. The spray
from a DDT bomb or from an .airplane spray tank is an aerosol while
cigarette smoke is an aerosol of cigarette ash in air. An emulsion is a
colloidal dispersion of one liquid in another. Milk consists of colloidal
droplets of butter fat in water. No example of a colloidal dispersion
of one gas in another gas is known; such “mixtures” form true solutions.

Colloidal dispersions may also be classified according to the physical state
of the dispersed phase and the dispersion medium, as shown in Table 36-B.

Table 36-B

Types of Colloidal Dispersions

Dispersed

phase

Dispersion

medium

Common name

Examples

Liquid

Gas

Liquid aerosol

Clouds, mists, fogs, sprays

Solid

Gas

Solid aerosol

Smoke, dust

Gas

Liquid

Foam

Foams, whipped cream

Liquid

Emulsion

Emulsions, mayonnaise, milk

Solid

Liquid

Sol

Starch suspension, paints,
printing inks

Gas

Solid

Pumice stone, white flowers,
gray hair, floating soap

Liquid

Solid

Gels, jellies, cheese,
opal (H 2 0 in SiO a )

Solid

Colored gems, some alloys

Further, depending upon the relative affinity of the dispersed phase for
the dispersion medium, colloidal dispersions may be divided into broad
classes. If the tendency to go into or tp remain in colloidal dispersion is slight,
the dispersed phase is said to be lyophobic (solvent fearing). A prefix may
be used to denote the nature of the dispersion medium; thus, hydrophobic ,
when the medium is water. Such colloids are easily precipitated by the addi-
tion of electrolytes, and if dried do not readily reform colloidal dispersions.
In general, lyophobic colloids are mainly aqueous dispersions of inorganic
substances, such as the metallic elements or metallic oxides, sulfides, etc.

If the affinity between the dispersed phase and the dispersion medium
is high, the colloidal substance is said to be lyophilic (solvent loving), or,
in the case of aqueous dispersions, hydrophilic. Dispersions of hydrophilic
colloids are more viscous than water itself and are more stable than hydro-
phobic colloids in that they are not readily precipitated by the addition
of electrolytes. They behave reversibly in that they can be separated from

480

Colloid Chemistry

the dispersion medium, by evaporation for example, and dried, and subse-
quently the dried material, upon being mixed with the dispersion medium,
will readily return to colloidal dispersion. Examples of hydrophilic colloids
are gelatin and starch. Upon the addition of water these substances spon-
taneously disperse to form colloidal systems. Such a direct disintegration
into particles of colloidal size, upon addition of the proper agent, is known
as peptization . Gums, resins, and shellacs behave similarly to form colloidal
dispersions in solvents other than water.

3. Preparation of Colloidal Dispersions. Inasmuch as colloidal particles
are intermediate in size between coarse particles and particles of molecular
dimension they can be prepared by two methods. We can start with coarse
particles and disintegrate them into particles of colloidal dimensions by
dispersion methods; these generally involve physical procedures. Or we
can start with particles of molecular dimensions and by condensation methods
cause them to grow into aggregates of colloidal size; such methods require
chemical procedures.

(A) Dispersion methods: Though it is difficult to attain colloidal sizes
by ordinary mechanical grinding alone, grinding in special mills known as
colloid mills, wherein a material is sheared between closely spaced discs
rotating at high speeds in opposite directions, reduces many substances to
colloidal dimensions. Such mills are used widely in industry for grinding
paint pigments and face powders. Spraying processes using atomizers that
give fine, uniform sprays can produce colloids. Powdered milk of colloidal
particle size is obtained by dehydrating milk spray. Homogenized milk is an
emulsion which is produced by forcing milk under pressure through orifices
in a metal plate in order to break up the fat globules. By causing an electric
arc between electrodes of a metal under the surface of a liquid, a colloidal
dispersion of a metal in a liquid can be prepared. The metal is vaporized by
the electric arc and its vapor then condenses to form particles of colloidal
size. The process is known as the Bredig arc method. The method is obviously
a combination of dispersion into particles of atomic size followed by their
condensation and growth into colloidal dimensions. The peptization of sub-
stances such as gelatin and glue may also be considered a process of dispersion.

(B) Condensation methods: Many chemical reactions yield insoluble
substances. A precipitate is formed if such a substance grows beyond colloidal
dimensions and forms aggregates of macroscopic size but if the growth
stops at colloidal dimensions a sol is produced. When H 2 S is added to a
solution of arsensic(III) oxide a colloidal dispersion of arsenic (III) sulfide
is formed.

(1) As 2 0 3 + 3 H 2 S As 2 S 3 + 3 H 2 0

When concentrated FeCl 3 solution is added to a large quantity of hot water,
the Fe(OH) s produced is colloidal.

(2) FeCl 3 + 3 H 2 0 Fe(OH) 3 + 3 HC1

Colloidally dispersed metals can be produced by the reduction of their ions.
Thus a gold(III) chloride, AuCl s , solution, when treated with a reducing

Colloid Chemistry

481

agent such as formaldehyde, will yield elemental gold in colloidal dispersion.
Since the wavelengths of light scattered by particles in colloidal dispersion
depend upon the size of these particles, gold sols can vary in color from
blue to ruby red. This technique is specially applicable to metals below
hydrogen in the electromotive series, for example, Pt, Hg, and Cu. For
molecular solutes, when a solution is mixed with a liquid in which the
substance is less soluble, the solute may separate out as aggregates of
colloidal dimensions. If a solution of sulfur or rosin in alcohol is poured
into water a colloidal dispersion is formed.

4. Filtration and Dialysis. Colloidal particles are small enough to pass
through ordinary filter paper. Indeed filter paper can serve as a rough but con-
venient device for separating colloidal particles from those in coarse suspension
since the diameter of the pores in ordinary filter paper is about Ifi. To separate
colloidal particles from those in true solution, such as inorganic electrolytes,
filters with much finer pores must be used. Such pores are available in parch-
ment, cellophane, and certain animal membranes, through which only particles
of molecular dimensions can pass. The technique of separation is known as
dialysis (Figure 36.1) and is frequently used to purify colloids. from dissolved

The colloidal dispersion to be purified is placed in a bag of parchment paper,
cellophane, or other special membrane. The bag is suspended in a beaker con-
taining pure solvent’ (the dispersion medium, as water, alcohol, etc.). The colloidal
particles cannot diffuse through the membrane but the materials in solution, mole-
cular or ionic, tend to establish equilibrium concentrations on both sides of the
membrane by diffusing through the membrane into the solvent in the beaker. Ulti-
mately this material would be distributed more or less in equal concentrations
within and outside the bag, but if the solvent in the beaker is continuously replaced
by fresh solvent the concentration of the dissolved material outside the membrane
is kept low. Because of the concentration gradient, there is a continuous flow of
the dissolved material from within the bag to the pure solvent. In this manner
the soluble impurities in the colloidal dispersion can be removed completely.

Figure 36.1 . Dialysis.

482

Colloid Chemistry

electrolytes and other impurities. The process is slow since the movement
of any ions through the membrane is solely by diffusion. Sometimes pressure
is applied to the colloidal dispersion, in which case the process is known
as ultra-filtration. In olectrodialysis , an electric potential is used to hasten
the movement of the ions across the membrane.

5. Optical Properties. In contrast with suspensions, but like solutions,
colloidal dispersions may be perfectly clear to the naked eye and when
viewed under the microscope. Indeed the upper limit of size of colloidal
particles was set at approximately the limit of visibility in the ordinary
microscope. To the naked eye, however, colloidal dispersions often appear
turbid, especially when examined at right angles to the bath of the beam
of light illuminating the colloid. This is the Tyndall effect; it is due to the
fact that colloidal particles are large enough to scatter light (Figure 36.2).

Source of
light.

True Solution

Colloidal Dispersion

A beam of light passing through a true solution is invisible. When the same beam
is passed through a colloidal dispersion it becomes visible due to the scattering
of the light by the colloidal particles.

Figure 36.2. The Tyndall Effect.

For the same reason a sunbeam is visible in a darkened room, the light being
scattered by the dust motes in the air.

For particles to be capable of scattering light or of deflecting light
from its straight line path, they must have dimensions approximately equal
to the wavelength of the light scattered. The wavelengths of visible light,
approximately 4 X 10“ 5 cm to 7.5 X 10~ 5 cm, fall within the range of dimen-
sions of colloidal particles so that many such particles are capable of scat-
tering visible light. In solution, particles of molecular dimensions are far too
small to scatter visible light; such a beam of light passes through a solution
unaffected and hence a solution appears to be perfectly clear.

In an ordinary microscope the observer s eye, the object to be viewed,
and the source of the light all lie in a straight line. A microscope so ar-
ranged that the line of sight is at right angles to the beam of incident light
is called an ultramicroscope (Figure 36.3). With such an instrument particles
too small to be seen through the ordinary microscope can be “observed”
Though colloidal particles themselves are invisible in the ultramicroscope,

Colloid Chemistry

483

Figure 36.3. The Ultramicroscope.

An intense beam of light at right angles to the
line of vision of the miscroscope is focused, by
means of a lens, to a sharp point in the
colloidal disperson on the microscope stage.
The microscope is then focused on the point
of light within the dispersion. Particles too
small to be detected by the ordinary micro-
scope can be detected by this means as irregu-
larly moving points of light. Using ultra-
violet instead of visible light and photograph-
ing the images, scientists have detected particles
only one hundred times larger than the hydro-
gen molecule.

because of its ability to scatter visible light, each colloidal particle appears
as a bright spot on a dark field-

Colloidal particles observed by the ultramicroscope appear to be in a
state of rapid, irregular, dancing motion called the Brownian movement
after the botanist, Robert Brown, who first observed it in 1827 in pollen
grains suspended in water (Figure 36.4). The motion is caused by the im-

Figure 36.4. Brownian Movement.

Irregular path of a single colloidal particle as ob-
served in the ultramicroscope. The movement is due
to the bombardment of colloidal particles by mole-
cules of the dispersion medium.

pact of molecules of the dispersion medium striking the suspended colloidal
particles. The smaller the size of the colloidal particles the more vigorous
is its Brownian motion. Brownian movement is additional evidence for the
kinetic theory of matter. It is a factor in the absence of settling' of colloidal
particles even though their density is greater than that of molecules of the
dispersion medium. A gold sol prepared by Michael Faraday in 1857 shows
no detectable settling and is still perfectly clear after more than 100 years.
Because of gravitational attraction, however, there is “settling” in that there
results a gradation in particle size with the smaller particles at the top and
the more massive ones at the bottom. The equilibrium distribution that is
ultimately attained is the result of two opposing forces, gravitational attrac-
tion and Brownian motion.

484

Colloid Chemistry

6. Colligative Properties. We have seen that the addition of a solute
to a solvent lowers the vapor pressure, lowers the freezing point, raises the
boiling point, and gives rise to the phenomenon of osmotic pressure. The
magnitude of these changes depends upon the number of dissolved particles
alone and not upon their nature. In colloidal dispersions the magnitude of
these changes is so small as to be almost negligible. A given weight Of sub-
stance dispersed as molecules yields a much larger number of particles than
does the same weight dispersed as colloidal particles. If we assume that the
average number of molecules, comprising one colloidal particle of As 2 S 3 ,
for example, is 1,000, a colloidal dispersion containing one formula weight
of As 2 S 3 in 1,000 grams of water will have a freezing point depression of
only 0.00186 °C, a value which is normally quite negligible and within the
error of measurement. Such small freezing point depressions give rise to
very high calculated molecular weights for colloidal particles.

7. Adsorption Properties. Because of their positions, surface molecules
have unsatisfied or “residual” binding forces by which they can adsorb
or condense foreign materials. The forces involved in adsorption are either
Van der Waals forces or the same electrical attractions that cause ordinary
chemical combination. The amount of adsorption depends on the extent
of the surface area available for adsorption. Activated carbon is an excellent
adsorber because its innumerable capillaries constitute an enormous surface
area. For a similar reason, the fine granular condition of silica gel makes
it a good adsorbent.

Some idea of the relative surface area which can be produced by sub-
dividing a given mass of material may be gained from the following illustration.
For a cube whose edge is 1 cm in length the surface area is 6 cm 2 . If this
cube is divided into cubes 1 mm on edge, or 1,000 cubes in all, the total
area of the surfaces of the smaller cubes thus formed is 60 mm 2 , a tenfold
increase. The total surface area is inversely proportional to the length of the
edge of the cube, so that if the process of subdivision is continued until
the particles are cubes I mp in length, the total surface area is 60,000,000 cm 2 ,
or about 1.5 acres.

In comparison with their volumes colloidal particles have large surface
areas and hence have great adsorptive capacities. Many colloidal properties
depend upon this attribute so that colloid chemistry is largely the chemistry
of surface effects. Because of their adsorptive powers many colloidal materials
are not "pure.” In many cases the deliberate addition of a hydrophilic colloid
stabilizes a hydrophobic colloid. The hydrophilic colloid is adsorbed as a
protective layer or “coating” and thereby prevents the hydrophobic particles
from uniting with their neighbors to form aggregates sufficiently large to
precipitate. In this manner gum arabic acts as a protective colloid to prevent
the precipitation of the coloring materials in inks.

8. Electrical Properties. Not merely other colloids or neutral molecules
but also ions can be adsorbed by colloidal particles. An important property
of colloidal particles is that they can be electrically charged. This can be
shown experimentally inasmuch as most aqueous colloidal dispersions, and
all the hydrophobic ones, conduct electricity. When an electric current is

Colloid Chemistry

485

passed through an AS 2 S 3 sol, the colloidal particles migrate to the anode,
evidence that the As 2 S 3 particles bear a negative charge. At the anode the
charge is neutralized and the colloidal dispersion coagulates as a precipitate
of As 2 S 3 . In other colloidal systems, for example, an Fe(OH) s sol, particles
move to the cathode, indicating that they are positively charged. The migra-
tion of colloidal particles in an electric field is known as electrophoresis.

In a given colloidal system all the colloidal particles have the same type of
charge. Most metal oxides and hydroxides have positive charges while the
free metals and most sulfides carry negative charges. The charge on a colloidal
particle is due to the adsorption on its surface of a particular ion which
is present in the dispersion medium. Preferential adsorption of positive ions
by a colloidal particle results in a positive charge on the particle whereas
adsorption of negative ions produces negatively charged colloidal particles.
Colloidal particles of Fe(OH ) 3 preferentially adsorb Fe 3+ ions and so be-
come positively charged whereas As 2 S 3 colloidal particles adsorb S 2 “ ions and
are negatively charged (Figure 36.5). In the absence of any electrolyte many

Ions are adsorbed on the surface of a colloidal particle. The entire colloidal dis-
person remains electrically neutral because there are present ions of charge opposite
to those adsorbed throughout the solution.

Figure 36.5 . Adsorption of Ions by Colloidcd Particles.

colloids adsorb H + and OH“ ions from water; most proteins adsorb these
ions but at different sites in the molecule.

The adsorption of ions is a major factor in stabilizing a colloidal dis-
persion. Inasmuch as the colloidal particles in a given system have the same
charge they repel each other ana hence the tendency to collide and to
coalesce is reduced. Neutralization of the adsorbed charges generally permits
the coagulation of the colloidal particles so that the addition of a sufficiently
high concentration of an electrolyte will precipitate a colloidal dispersion.
The addition of HC1 precipitates an As 2 S 3 sol; the increased concentration
of H + ions neutralizes the adsorbed S 2 “ ions through the formation of H 2 S.
Similarly .a Fe(OH ) 3 sol is precipitated by the addition of alkali. When they

486

Colloid Chemistry

are mixed, oppositely charged colloids such as As 2 S 3 and Fe(OH) s precipitate
each other. The flocculating power of an ion, or its ability to cause colloidal
precipitation, depends largely upon its charge. For equal molar concentra-
tions of K + , Ca 2+ , and Al 3+ ions the flocculating power is in the ratio of
1 : 80 : 1500.

The formation of deltas at the mouth of rivers is partially attributable
to the precipitation of colloids. River water contains clay and silt in colloidal
dispersion. The colloid is negatively charged and stabilized by the adsorp-
tion of hydroxide ions. When the river water reaches the salty sea water the
colloidal clay is precipitated by the neutralization of its negative charge by
Na 4 * and Mg 2+ ions of the sea water. Over a period of thousands of years
sufficient precipitate accumulates to form a delta.

Aerosol particles in smoke and dust are colloidal and often electrically
charged. A method developed by the American chemist, Frederick Cottrell,
neutralizes the charges on such particles and causes them to settle out. The
gases or fumes, carrying the colloidal particles in suspension, are passed be-
tween electrodes maintained at a high difference of potential. The charged
colloidal particles are attracted, to the oppositely charged electrodes where,
upon discharge, they are deposited as dust. Both positive and negative col-
loidal particles are precipitated. The process is employed not only to prevent
the escape of noxious industrial smokes into the atmosphere, blit also to
recover valuable products that would otherwise escape from the stacks of
kilns, smelters, and furnaces.

9. Emulsions, Detergents, and Soaps. Under ordinary circumstances oil
and water are immiscible. If an oil is shaken with water the oil can be
broken down into droplets of colloidal dimensions but upon standing, the
droplets coalesce to macroscopic dimensions. Separation of the liquids into
two separate layers occurs and a stable colloidal dispersion is not formed.
By the addition of certain substances, known as emulsifying agents , to such
immiscible liquid mixtures there can be formed an emulsion, a stable colloidal
dispersion of one liquid in another. Such substances act either by coating
the dispersed droplets with a protective film or by decreasing the surface
tension of one or both of the liquids so that their tendency to coalesce is
decreased. Materials which decrease the surface tension of liquids are
called wetting agents. Common emulsifying agents include detergents and
soaps, which operate to reduce surface tensions, and gelatin, gum arabic,
and other lyophilic colloids, which tend to form protective coatings or films.
Mayonnaise is an emulsion of olive oil in vinegar, stabilized by egg yolk as
the emulsifying agent; milk is an emulsion of butterfat in water with the
protein caSein serving as the emulsifying agent.

The cleansing action of detergents and soaps depends upon their ability
to act as emulsifying agents. Both detergents and soaps consist of molecules
which contain a polar end and a long nonpolar carbon chain end. Examples
of each are:

CH 8 (CH 2 ) 15 CH 2 COO- Na+
sodium stearate, a soap

CH3(CH 2 ) 10 CH 2 OSC>3- Na+
sodium lauryl sulfate, a detergent

Colloid Chemistry

487

The nonpolar end of these molecules is soluble in organic substances, such
as oils and greases, whereas the polar end is soluble in water. When an
oil is stirred with a soap solution, for example, the hydrocarbon ends of the
soap molecules dissolve in the oil droplets while the polar ends stick out into
the surrounding water. Such an orientation of the soap molecules is known
as a micelle (Figure 36.6). In effect, the oil droplet is surrounded by a

The polar end of the soap molecule dissolves in the water while the nonpolar, or
hydrocarbon portion, dissolves in the oil droplet, producing an oriented soap film
around the droplet.

Figure 36.6. A Micelle.

water soluble layer and, as such, can be washed away in the bulk of the
surrounding water. Repulsion of similar charges around the droplet prevents
them from coalescing and a stable emulsion is obtained. With solid dirt
particles the hydrocarbon ends of the soap or detergent molecules are
adsorbed to the surface of the dirt and the polar ends of these molecules
again form a water soluble, protective film around the dirt particle. Vigorous
agitation of the oil or dirt assists the washing process by breaking down these
materials to small, perhaps colloidal dimensions.

10. The Realm of Colloids. Colloidal dispersions are of the greatest im-
portance. The protoplasm of living cells is colloidal and hence biological
processes are concerned in some way with the properties of colloids. Organic
polymers can have dimensions which fall within the colloidal range. Thus
rubber, polystyrene, nylon, and proteins, to mention but a few examples,
exhibit colloidal properties. In the inorganic world, the photographic emulsion
and the process of ore flotation are colloidal. It would be difficult to over-
emphasize the importance of the field of colloids.

488

Colloid,' Chemistry

QUESTIONS

1. Distinguish between suspension, solution, and colloidal dispersion.

2. What is the range of size of colloidal particles? How was this range selected?

3. Into what class would spherical particles belong if their diameters are

(a) 6 X 10- 7 cm (b) 0.02/u (c) 2.0 A (d) 0.7 m/x (e) 7.0/x (f) 0.001 cm?

4. Define and illustrate the following terms: (a) dispersed phase (b) dis-
persion medium (c) adsorption (d) peptization (e) gel (f) jelly (g) sol
(h) emulsion (i) lyophobic (j) hydrophilic (k) micelle (1) electrophoresis.

5. Describe methods of preparing colloidal dispersions.

6. How could colloidal dispersion of the following be prepared: (a) gold

(b) an oil (c) Fe(OH) 3 (d) sulfur?

7. List properties in which colloidal dispersions and solutions differ.

8. Describe processes for the purification of colloidal dispersions.

9. A sol contains particles ranging in size from 0.2/x to 0.6/x, and an impurity
of NaOH. How could the two be separated?

10. What is Brownian motion? To what is it due?

11. Explain why colloidal dispersions do not settle out.

12. Explain why material in colloidal dispersion has only a negligible effect on the
freezing point of the dispersion medium.

13. Assuming Raoult's Law is obeyed calculate the relative freezing point de-
pressions in water of 1000 grams of (a) a solute whose molecular weight is
100 and (b) a colloid whose molecular weight is 100,000.

14. What is the Tyndall effect? What is the explanation for it? What is the principle
of the ultramicroscope?

15. To what property does charcoal owe its adsorptive power?

16. What is the source of the charge on a colloidal particle? How could you
demonstrate the presence and the nature of such charge?

17. What is the principle of operation of the Cottrell precipitator?

18. Why does colloidal clay in a river precipitate when it meets the ocean?

19. What is the function of an emulsifying agent? Give examples.

20. Explain the cleansing action of soaps and detergents.

21. Which of the following would be most effective, mole for mole, in coagulat-
ing a sol composed of positively charged particles: (a) NaCl (b) MgCl 2

22. Why does “French dressing/' composed of vinegar, salt, and olive oil give
an emulsion which is less durable than that of mayonnaise dressing?

23. Calculate the freezing point depression and the osmotic pressure of a one
percent aqueous solution of hemoglobin.

Metals

Though the term metal does not lend itself to a completely unambiguous
definition and though there are many exceptions to the criteria which follow,
nevertheless there are certain qualities which are generally applied to those
substances classified as metals. If a crystalline substance has the bright ap-
pearance known as metallic luster and is a good conductor of heat and
electricity, it is classified as a metal. Certain mechanical properties such as
high tensile strength, toughness, and hardness are also usually associated
with the metallic state (Table 37-A).

Table 37-A

Properties of Metals and Nonmetals

Metals

Nonmetals

Metallic luster

No metallic luster

Opaque

May be translucent or transparent

Good conductor of heat and electricity

Poor conductor; may be an insulator

Malleable and ductile

Brittle and nonductile

High density

Low density

Solid at ordinary temperature

May be solid, liquid, or gas

Not all metals, however, exhibit all the properties ascribed to the metallic
state while some nonmetals have properties which we consider to be metallic.
While we think of metals as having high melting points and high densities
some have relatively low melting points, for example, mercury (-39°C)
and gallium (30°C), whereas lithium, sodium, and potassium of the alkali
metals all have densities less then that of water. Indeed the alkali metals
are soft enough to be cut with a penknife. On the other hand, the non-
metals carbon and sili con have high melting points and exist in allotropic
forms which are among the hardest substances known.

490

Metals

1. Nature of the Metallic State. Many of the properties of metallic ele-
ments can be explained on the basis of their crystalline structures and their
electron configurations. X-ray diffraction studies of solid metals reveal that
the great majority these elements crystallizes in one of three arrangements
or lattices: the face-centered cubic, the body-centered cubic, and the close-
packed hexagonal structures.

In the face-centered cubic lattice one atom of the metal occupies each
comer of a cube and one atom is at the center cf each plane face of the
cube so that each atom is surrounded by twelve equidistant neighboring atoms.
This number of equidistant neighbors to which an atom is "joined" is known
as its coordination number . Examples of metals, crystallizing in this manner
are copper, silver, gold, calcium, aluminum, and lead.

In the body-centered cubic structure one atom occupies each corner of
a cube and one atom is at the center of the body of the cube. The co-
ordination number is eight so that this is a less closely packed structure
than is the face-centered cubic. The alkali metals and barium crystallize in
this fashion. The low density and the softness of the alkali metals is prob-
ably due to this structure.

The close-packed hexagonal is a structure in which the atoms occupy
the comers of a planar hexagon; planes containing these hexagons lie one
above the other. The coordination number is twelve. Beryllium, magnesium,
zinc, cadmium, and titanium crystallize in this structure.

The close-packing of metallic atoms accounts for the relatively high
densities of most metals. Nonmetallic elements have a maximum coordination
number of four, for example, carbon and silicon; other nonmetals have lower
coordination numbers. The ability of metals to deform, their malleability
and ductility, is also consistent with these structures. Deformation of a metal
causes a slippage of atoms along adjacent planes within the crystal. Atoms
then occupy positions previously held by their neighbors, but because all
metal atoms are identical, the attractive forces within a crystal are un-
changed. In an ionic crystal, slippage which shifts like charged ions into
positions one above the other results in repulsion between the ionic planes
and rupture of the crystal (Figure 37.1).

Applied

Force

Metal

Ionic
Crystal

The numbered circles represent metal atoms; circles containing -Hand - represent
ions of the charge indicated. If a shearing force is applied to a metallic crystal
and an ionic crystal resulting in a shift of one atomic oripnic diameter, intermetallic
forces are unchanged, but in an ionic crystal, repulsive forces are produced between
like ions in adjacent planes which cause disruption of the crystal.

Figure 37,1 Deformation of a Metal and an Ionic Crystal.

Metals

491

From the viewpoint of their electron configurations, metal atoms have
in most cases less than four electrons in their highest principal energy level
whereas, except for hydrogen, helium, and boron, nonmetal atoms have
four or more electrons in this level. Neither the covalent bond, the simple
ionic bond, nor Van der Waals forces between molecules offers an adequate
explanation of the bonding between metal atoms. Because of the high co-
ordination number of a metal atom in the crystalline state there is an in-
sufficient number of valence electrons for the atom to form covalent electron
pair bonds with all of its nearest neighbors. Thus a calcium atom, which has
but two valence electrons and which crystallizes in the face-centered cubic
system, cannot form covalent bonds to each of its twelve neighbors. In effect,
each calcium atom contributes 2/12, or 1/6, of an electron to each of its
neighbors; between two calcium atoms the equivalent of 2/6, or 1/3, of an
electron acts as the bond.

The ordinary type of ionic bonds cannot be formed in metals because
they would involve the transfer of electrons but, since all atoms are alike
in a given element, one atom does not have a greater attraction for elec-
trons than does another. Further, the generally low volatility, high melting
point, and hardness of metals are not in keeping with the postulation of
weak Van der Waals forces such as exist in a molecular crystal; also the
crystal structures of metals negate the existence of molecular units.

Hence a new type of bond, the so-called metallic bond , must be invoked.
It is believed that, in a metal, the atoms release their valence electrons so
that *a metal is truly a lattice of positive ions permeated by a cloudi of
electrons. The electrons are not localized as they would be in a specific
bond or even over a molecule but belong to the metallic crystal as a whole,
binding the entire crystal and holding the ions in fixed positions. The quan-
tum energy levels available to these electrons are very closely spaced so that
a small amount of energy serves to excite an electron to a higher level.
The opacity of metals is due to the fact that these free electrons can absorb
any frequency of light.

Because the cloud of electrons is mobile, the application of an external
potential causes a movement of these electrons, that is, a flow of electrical
current, from the negative to the positive terminal and thereby a metal is a
conductor of electricity. The electric conductivity of a metal is an electronic
conductance as distinguished from the ionic conductance of solutions of
electrolytes and of fused salts. Only the free electrons in the cloud move
in the metals, whereas ions move in solutions and fused electrolytes. Non-
metals, whose valence electrons are firmly bound in localized ionic or
covalent bonds, are nonconductors of electricity.

The electrical resistance of a metal increases with a rise in temperature.
This is attributed to an increased vibrational motion of the metallic atoms
which interferes with the flow of electrons. At temperatures close to absolute
zero, atomic oscillations are so diminished that metals become extremely good
conductors, or even superconductors. Indeed a current induced in a lead
ring at temperatures below 7°K will circulate for years without the input
of additional energy. The study of materials at extremely low temperatures

492

Metals

is known as cryogenics . Many unusual properties of matter first appear at
temperatures near absolute zero.

In the Periodic Table the metallic elements are on the left and the non-
metallic elements on the right. Certain elements, known as metalloids , are
intermediate in character and have both metallic and nonmetallic properties.
In this class are boron, silicon, germanium, arsenic, antimony, tellurium, and
astatine, A line drawn on the Periodic Table from boron to astatine will
delineate roughly the metallic from the nonmetallic elements. Of the 103
elements known at this writing, 79 may be classed as metals, 7 as metalloids,
and 17 as nonmetals.

2. Occurrence of Metals in Nature. A few metals are found uncombined
in nature. From the order of chemical activity of the metals it follows that
only those elements that are below hydrogen in the series can exist in
nature in the free or elemental state. For if an element above hydrogen is
exposed to water or to moist air, hydrogen will be displaced and the hy-
droxide, oxide, or carbonate of the element will be formed. Hence the
only metals found in the elemental condition are copper, mercury, silver,
gold, the platinum metals, and the members of the arsenic family. Because
of the solvent action of water (rain, streams, etc.), only insoluble compounds
of metals occur in nature. The soluble salts are found in sea water or as
underground deposits formed by the evaporation of prehistoric seas or inland
lakes. These insoluble naturally occurring compounds are mainly oxides,
hydroxides, sulfides, carbonates, sulfates, and silicates.

A mineral is any element or any inorganic compound found in nature.
An ore is a deposit of a mineral worthy of commercial development. The
concentration of a metal in an ore must be sufficiently high to make its
extraction profitable. Thus an ore with 10% iron is worthless while one
with 10% copper is worth working; commercial iron ores contain 35-50%
iron while ores with 0.002% gold are good ores. The chemical composition
of an ore is also a factor in determining its workability, or the ease with
which the metal can be extracted; a sulfide ore of iron containing 45% iron
would be rejected in favor of a lower percentage oxide ore. Important
classes of ores are listed in Table 37-B.

3. Metallurgy, The science and practice of extracting metals from their
ores is called metallurgy. Each metal presents an individual problem in
metallurgy, depending upon its chemical properties and the nature of the
ore from which it is to be extracted. The metallurgical operations employed
in extracting heavy metals such as iron, copper, lead, and zinc, include three
principal steps: A) preliminary treatment of the ore, B) reduction of the
ore, and C) refining the metal.

(A) Preliminary treatment: When taken from the ground most ores con-
tain large quantities of "foreign material such as rock, clay, sand, and lime-
stone. These impurities in the ore are known collectively as gangm. Some
high grade ores require no concentration but most contain so high a per-
centage of gangue that it is frequently worth while to apply some mechani-
cal process at the start which concentrates the ore by separating out the
great bulk of the gangue, thereby effecting a saving in bulk handling and

' Metals

493

Table 37 -B

Classes of Ores

Type of Ore

Examples

Native metals

Gold, silver, copper, platinum, mercury, arsenic, antimony, bis-
muth

Oxides

Iron, aluminum, copper, tin, manganese, chromium, and others

Sulfide

Zinc, cadmium, mercury, copper, silver, lead, cobalt, nickel,
arsenic, antimony, bismuth

Carbonate

Iron, copper, lead, zinc. The compounds CaC0 3> SrCO s , BaC0 3 ,
and MgCOg are used to prepare compounds of these metals.

Sulfate

CaS0 4 • 2 H 2 0, SrS0 4 , BaS0 4 , PbS0 4 .

Halide

NaCl, KC1, MgCl 2 * H 2 0; found in salt beds and in sea water.

Silicate

Though silicate minerals are common, they are generally un-
suitable as ores because of chemically difficult metallurgy. The
silicates of Be, Zn and Ni are important.

in subsequent fuel costs. A coarse and inefficient separation can sometimes
be effected by hand picking where labor costs are very cheap. However,
the material is usually finely ground and separation of the gangue from the
ore is accomplished by taking advantage of their specific gravities; Through
the use of purely physical methods, such as shaking or jiggling on an in-
clined table, the less dense particles of gangue can be washed away from the
rest of the mass. For a magnetic ore such as magnetite, Fe 3 0 4 , huge magnets
can be used to separate the ore from the crushed powder.

In some cases, for example, copper sulfide and zinc sulfide ores, ad-
vantage is taken of the preferential wetting of the ore by an oil. The ore
is ground to a very fine powder and then mixed with water containing about
one percent of an oil. Air is then bubbled through the suspension, producing
a heavy froth or foam on the surface. The metal sulfide is wetted by the
oil but the gangue is not, and the sulfide-oil mixture is carried to the sur-
face by air bubbles surrounded by films of oil. From the surface the
froth containing the sulfide can be skimmed off; the gangue remains in
the water under the surface-oil froth. By this flotation process it is possible
to concentrate over 90% of a sulfide ore into one-tenth of its original bulk,

Oxidfe ores are already suitable for reduction; other ores must be roasted
to form the metal oxide. Roasting consists of heating the finely divided
material in a current of air or of air-enriched furnace gases. Sulfide ores
are burned to oxide and carbonates form oxides through the loss of C0 2 .
Sulfur and arsenic are removed from the ore as volatile oxides and excess
water are driven off.

494

Metals

(1) 2MS + 30 2 ->2M0 + 2S0 3 (2 ZnS + 3 0 2 -*■ 2 ZnO + 2 S0 2 )

(2) MCO s -> MO + CO, (ZnCO, ZnO +C0 2 )

(B) Reduction: the next step is the reduction or smelting process. This
is carried out at high temperature in special furnaces where the ore is
mixed with a reducing agent. The most common reducing agent is carbon,
either in the form of coal or coke. The carbon itself or the carbon monoxide,
CO, produced by the incomplete combustion of the carbon, reduces the
oxide to the free metal.

(3) MO + C M + CO (CuO + C -> Cu + CO )

(4) MO + CO -> M + C0 2 (CuO + CO Cu + C0 2 )

Even after concentration most ores still contain some gangue. A substance,
known at the flux, is added during the reduction process to combine chemi-
cally with the residual gangue and to form a $hg. The latter is a substance
which is immiscible with the metal and has a low melting point. Where
the metal is liquid at the temperature of the furnace the slag, having a lesser
density, floats on the metal and protects it from oxidation. Because of this
layering the two can be tapped off separately. When the gangue is sand or
silicious material, a flux which will give a basic reaction such as limestone,
CaCOs, is uised. If the gangue were limestone, an acidic flux such as silica,
Si0 2 , would be used. In either case, fused glassy calcium silicate, CaSiO*,
is produced.

(5) Si0 2 —f- CaC0 3 — ^ C&Si0 3 -j— C0 2

Volatile metals such as zinc are driven off in the form of vapor during smelt-
ing and are recovered by condensation; in such cases no flux is required.

Some metals are not readily reduced from their oxides by carbon.
Though more expensive, other more powerful reducing agents such as alum-
inum, iron, or hydrogen must be used. By the use of aluminum, a number
of metals such as manganese and chromium are now obtained on a commercial
scale in a high state of purity. The process is known as aluminothermy or the
Goldschmidt process.

(6) Cr 2 0 3 ”4“ % Al — > 2 Or — j- A1 2 0 3

The sulfides of some metals, such as SbS and PbS, can be treated directly
with a reducing agent like iron.

(7) Sb 2 S 3 + 3 Fe 2 Sb + 3 FeS

The oxides of the very active metals generally are not reducible by chemi-
cal reducing agents and, if at all, only at extreme temperatures. Thus the
oxides of the alkali metals and of calcium, magnesium, and aluminum can
be reduced by carbon in an electric furnace but under these circumstances
the metal combines with carbon to form a carbide. Such metals are usually
prepared by the electrolysis of their fused salts, e.g., NaCl and CaCl 2 .
Aluminum is prepared industrially by the electrolysis of a solution of its
oxide, A1 2 0 3 , in fused cryolite, Na$AlF 6 , as the solvent.

Metals

495

(C) Refining: The crude metal produced by reduction is not 100% pure
but contains impurities such as small amounts of slag, some unreduced oxides
or sulfides, dissolved gases, and other metals, and hence must be refined or
purified if such impurities are objectionable. Low boiling point metals such
as zinc and mercury can be refined by distillation. Metals with low melting
points such as tin can be melted on an inclined table, thereby enabling the
low melting point metal to drain away from the higher melting solid im-
purities; the process is known as liquation .

However, electrolysis is the most common method of refining. The impure
metal serves as the anode and a fused salt or an aqueous solution of a salt
of the metal is the electrolyte. When the proper voltage is employed the
desired metal is deposited upon the cathode. Impurities remain either in
solution or as a sludge at the bottom of the electrolytic cell. Electrorefining
produces metals of the highest degree of purity; among the metals refined
electrolytically are copper, gold, zinc, lead, iron, aluminum, and chromium.

4. Special Metallurgical Processes. Inactive metals which occur in the
free state such as copper, silver, and gold can be obtained by liquation which
separates them from the gangue. These metals also dissolve in mercury to
form a solution from which the metal can be recovered subsequently by
distilling off the mercury.

Some metallic ores are treated with solutions of chemical reagents to bring
them alone into solution, a leaching process. The metal is then obtained by
precipitation with a reducing agent or by electrolysis. Thus gold and silver
form soluble complex ions with cyanide ion, CN~. For gold, the solution by
CN“ ion and reprecipitation by zinc are as follows:

(8) 4 Au + 8 CN- + 0 2 + 2 H 2 0 -» 4 Au(CN) 2 ~ + 4 OH'

(9) 2 Au(CN) 2 - + Zn 2 Au + Zn(CN)* 2 -

We are well aware that, in this nuclear age, the discoveries of science exert
a pronounced, if not decisive, effect upon political and economic affairs. For tech-
nological progress to govern the direction of historical events is not, however, a
recent trend.

The Stone Age existed when man had little technical knowledge to enable
him to make implements of materials other than stone or wood. Stone Age man
had no metallurgical capability.

Copper Oxide, CuO, being a compound of the relatively inactive metal, copper,
is easily reduced by carbon. Some ancestor, probably in Egypt as early as 2500 b.c.,
may have discovered the extraction of metallic Cu by heating a copper, ore with
a carbonaceous substance, perhaps wood. The mixing of copper with tin to produce
an alloy, bronze, having more desirable mechanical properties than the Cu alone,
resulted in the period known as the Bronze Age, wherein utensils and artifacts
of war were made of bronze The Trojan War, circa 1200 b.c., was fought with
bronze weapons.

For a very good chemical reason the Bronze Age preceded the Iron Age.
Iron is a more active element than copper so that iron oxide, Fe 2 0 3 , is not reduced
by carbon at the temperature of an ordinary fire. A higher temperature is required
and this can be obtained by using blasts of hot air in the furnace. This “blast

496

Metals

furnace” secret was zealously guarded by its discoverers, the Hittites. Iron, and
its alloy, steel, are much harder and tougher metals than are copper and bronze
so that the nation which could fabricate weapons of iron had a distinct military
advantage.

Only seven metals, Au, Ag, Cu, Hg, Fe, Sn, and Pb, were known to the ancients.

5. Alloys. A metallic product containing two or more metals, or 0ven
a metal and a nonmetal provided that the mixture has metallic properties, is
known as an alloy. Alloys are most commonly prepared by melting the two
metals together in the proper proportions and then allowing the melt to
solidify. In most cases the metals are completely miscible with each other
in the liquid state; lead and zinc are an exception. If one of the metals is
mercury the alloy is known as an amalgam. Alloys have certain desirable
properties, such as hardness or corrosion resistance, which the seperate metals
may not have. The magnitude of the properties of an alloy is not generally
the average of the properties of the individual metals; thus the melting
point of an alloy may be higher or lower than any of its components. Most
metallic materials of commerce are alloys, and not pure metals, e.g., steel
bronze, brass, 14 karat gold, etc. For industrial purposes alloys can be divided
into two classes— ferrous and nonferrous alloys. A list of the more important
alloys is given in Appendix IX. X-ray diffraction studies show that alloys
can be one of three different structural types: (A) an intimate mixture of
separate crystals of the metallic components; alloys of copper and lead are
of this type; (B) a solid solution in which the metals dissolve in each other;
silver-gold alloys are of this type; or ( C ) an intermetallic compound; examples
are the brasses, CuZn, CuZn 3 and Cu 5 Zn 8 , also Zn 3 Ag 2 and Cu 9 A1 4 . In such
interirfetallic compounds metals do not show their conventional valences.

6. Chemical Properties of Metals. Metals are electropositive. Most of the
common metallic compounds are salts and, unless the metal is part of a
negatively charged complex ion, the metal forms the positive ion. Because
of their relatively low attraction for their valence electrons, metal atoms
generally react as reducing agents. We have already seen that the oxidation
potential of a metal, as given by the Electromotive Series, is a measure of
its reducing ability. Some, such as the alkali metals, are strong reducing agents
and others, as gold and silver, are weak because of a small tendency to lose
electrons. Also the hydroxides of metals are generally bases while those of
nonmetals are acids; an aqueous solution of the metal oxide, CaO, is basic
whereas that of the nonmetal oxide, S0 3 , is acidic. In some measure the
strength of a base is determined by the size of the metal atom: The larger
this atom, the smaller is its attractive force for electrons, and the greater
is the tendency of a metal hydroxide to dissociate into OH" ions and metal
ions.

Many water solutions of metallic salts are acid due to hydrolysis. In
aqueous solution most metallic ions do not exist as simple ions but rather
as aquated complex ions with H 2 0 molecules as the ligand. Thus the zinc
ion is truly Zn(H 2 0) 4 2+ and can react with water to yield protons. Complex
ions are formed by many metals. The solution of silver chloride, AgCI, by
ammonia, NH 3j is due to the formation of the complex Ag(NH 3 ) 2 + ion.

Metals

497

Some of the heavy metals form part of a negative complex ion as in the
complex cyanides and chlorides, Ag(CN) 2 -, Fe(CN) 6 4 -, and Pt(Cl) 6 2 -.

7. Complex Ions. The formation of complex ions by metal ions* is a
common reaction. A complex ion has been defined as an ion consisting of
more than one atom but for the present we shall consider only complex ions
which consist of a metal ion combined with other negative ions or neutral
molecules. The groups attached to the metal ion are called ligands and their
number is the coordination number of the metal ion. Usually, but not in-
fallibly, the number of ligands or the coordination number, is twice the
charge of the metal ion; the net charge of the complex ion is the algebraic
sum of the charges of the species which go into making up the complex ion.

By comparison with the next higher noble gas configuration metal ions
are electron deficient. In other words they have orbitals which are unoccupied
or incompletely filled by electrons. Hence they combine with negative ions
or molecules which have unbound electron pairs that can form coordinate
covalent bonds to the metal ion. Most electronegative atoms, such as F, O,
N, Cl, and C, can so act as electron donors to the metal atom. Some complex
ions containing H 2 0 molecules are of the ion-dipole type wherein the positive
metal ion attracts the negative end, the oxygen atom, of the H 2 0 molecule.
In complexes containing the highly electronegative fluoride ion, F~, the bond
within the iqn i$ mainly ionic but the great majority of complex ions are
formed" through coordinate covalent bonding.

In the formation of these complex ions the coordinate bonds are hybrid
bonds, of which the most common are sp, sp 3 , dsp 2 , and d 2 sp B . The complex
ion has a spatial orientation corresponding to the hybrid bonds of which it
is composed. In the tetramminenickel(II) complex ion, Ni(NH 3 ) 4 2+ , the Ni 2+
ion has a coordination number of four and forms four tetrahedrally oriented
sp 3 bonds with four NH 3 molecules. The electron configuration of the neutral
Ni atom is Is 2 2s 2 2 p* 3 s 2 3 p e 3d 6 3d 1 3d 1 4 s 2 and that of the Ni 2+ ion identical
but for the loss of the 4s 2 electrons. The paramagnetic nature of these species
is in agreement with the two unpaired 3d electrons proposed in these con-
figurations. The electron configuration of the nickel atom in the complex
ion, Ni(NH 3 ) 4 2+ , may be represented by

Is 2 2s 2 2p 6 3s 3 3p* 3d 6 3d 1 3d 1 4s 2 4 p 2 4 p 2 4p 2

not involved in complex ion form four sp 3 bonding

bonding orbitals

This configuration has eight electrons more than does the simple Ni 2+ ion.
These eight electrons, one pair being donated by each of four nitrogen atoms
of the NH 3 molecules, form the four hybrid sp B bonds in the complex ion.
Note that six of these electrons entered 4p orbitals which were originally un-
occupied in the neutral Ni atom. Because the hybrid bonds are of the sp z
type the complex ion has a tetrahedral structure. The Ni atom is at the center
of the tetrahedral body and the NH 3 molecules are at the comers with the
nitrogen atoms coordinately bonded to the Ni atom (Figure 37.2). From

498

Metals

NH,

/ ' x

/ ' s v
/ Ni \ ^ — -)NE

/ - *v~ /

HsN«rC^ \

- O/

NHa

H H
H~N

H~N

h"h

X X

X N

N“H

N-H
H H

2 +

N f x X CN

5 Ni 1

1 X X !

NC- CN

(a) (b) (c)

(a) the tetrahedral shape of the Ni(NH 3 ) 4 2+ ion

(b) the coordinate bond is represented by an arrow from the atom which donates
the pair of electrons for the bond.

Figure 37.2. Structures of Nickel(II) Complex Ions.

the Lewis definitions of acid and base that a base is an electron-pair donor
and an acid an electron-pair acceptor, the NH 3 molecule is a base and the
Ni 2+ ion is an acid so that complex ion formation by coordinate bonding
can be considered as but another aspect of acid-base theory.

In the 4-coordinate tetracyanonickel(II) ion, Ni(CN)* 2+ , though the
total number of electrons is the same as in the Ni(NH 3 ) 4 2+ ion, the bonds are
of the hybrid dsp 2 type and the ion consequently has a square planar struc-
ture. Its electron configuration is

is 2 2s 2 2p 6 3s 2 3p 6 3d 2 3d 2 3d 2 3d 2 3d 2 4s 2 4p 2 4p 2 4p°

not involved in complex ion form four dsp 2 bond-
bonding ing orbitals

The absence of unpaired electrons is in accord with the fact that the
Ni(CN)/" ion is not paramagnetic.

An ion such as Cu 24 " can form two structurally similar types of complex
ions, one using 3d 4s 4 p 2 orbitals, whereby an inner d orbital, the 3d, is used,
or 4s 4 p 2 4d orbitals which use an outer d orbital, the 4d, the terms inner and
outer here referring to the d orbital employed in the bonding. Both con-
figurations yield square planar structures but can be distinguished because
the sp 2 d hybrid has the greater magnetic moment. The use of outer orbitals
results in bonds which are weaker and more polar than those formed from
inner orbitals. Transition elements which have unoccupied inner orbitals
form complex ions readily. Thousands of stable complexes are formed by
just these atoms alone: cobalt (III), chromium (III), platinum (II), and plat-
inum (IV). The regular, or nontransition elements, do not have incomplete
inner shells. Nevertheless an element such as aluminum can use outer 3d
orbitals in the formation of the complex Al(OH) 6 3 ~ ion.

For an element to have a coordination number greater than four, it must
use d orbitals since the combination of $ and p orbitals alone can form but

Metals

four bonds. Thus any complex ions which might be formed by elements
from hydrogen to fluorine can use only $ and p orbitals for a maximum
coordination number of four, e.g., BeF* 2 - and BFr, Elements higher in
atomic number than neon have available d orbitals of relatively low energy
which can be used in complex formation, and almost all such elements
form complex ions with a coordination number of six. For 6-coordinate ions
the most stable structure is the octahedral and the bonds are either inner
d?sp 8 or outer sp s d 2 hybrid orbitals.

In some cases, but not generally, complex ion formation results in a
noble gas configuration. Thus, the 6-coordinate complexes of cobalt(III)
such as Co(NH 3 )e 3+ have a total of 36 electrons and an electron configuration
similar to that of the noble gas, kryton. The cobalt atom is said to have an
effective atomic number (EAN) of 36. Other complex ions which have
noble gas configurations are 6-coordinate iron (II), 4-coordinate capper (I),
4-coordinate zinc(II), 8-coordinate molybdenum (IV), and _8-coordinate plat-
inum(IV).

In Table 37-C are listed the more important simple metal ions which
form complex ions and the manner of their bonding. Some ions appear under
more than one heading because a metal ion can form different types of
hybrid ions with different ligands.

Table 37-C

Hybrid Bonds in Metal Complex Ions

Ion Structure

Linear

Tetrahedral

Planar

Octahedral

Dodecahedral

Coordination Number

Hybrid bond

$p z

dsp 2

Inner

Outer

d*$p z

tPsp*

spPd 2

Metal ion

Cu+

Cu +

Ni 2+

Cr 2+

Al*+

Mo*+

Ag+

Be 2+

Cu 2+

Cr 8+

Ga 8+

Mo*+

Au +

Zn 2 +

p t 2 +

Mn 2 +

In*+

W 4 +

Cd 2 +

Pd 2 +

Fe 2+

Tis+

Hg 2 +

Fe 3+

Zn 2 +

Ni 2 +

Co 2+

Cd 2+

B 8 +

Co 8+

Hg 2 +

Ni 2 +

Ge 4+

Pt 4+

Inasmuch as the type of hybrid bond formed depends upon the electron
configuration of the metal atom it is not surprising that, under each type
of bond in Table 37-C there are ions of elements which exist as groups
in the Periodic Table. Thus, under sp s there are Be, Zn, Cd, and Hg; under
sp B d 2 are the elements of Group III. Within a given group of the Periodic
Table, complexing ability increases with increasing atomic number and,

500

Metals

where an element has more than one oxidation state, with increasing oxida-
tion number. Only even coordination numbers appear in Table 37-C; com-
plex ions with odd coordination numbers exist but are rare. Table 37-D lists
the common electron donor atoms and important groups in which they occur.

Table 37-D

Electron Donor Atoms in Complex Ion Formation

Donor Atom

Group

Name

Example

h 2 o

aquo

Cu(H 2 0) 4 2 +

Tetraaquocopper ( 11 )

OH-

hydroxo

Al(OH),*-

Hexahydroxoaluminum ( III)

NH S

ammine

Cu(NH 3 ) 4 2+

Tetramminecopper (II)

no 2 -

nitro

Co(N 0 2 ) 6 s -

Hexanitrocobalt ( III )

CN-

cyano

Cu(CN) 4 2 -

Tetracyanocopper ( II )

carbonyl

Ni(CO) 4

Tetracarbonylnickel (0)

S 2 -

thio

HgSjj 2-

Dithiomercury (II)

SCN-

thicyanato

Fe(SCN) 6 «-

Hexathiocyanatoiron ( III )

Halogen

F-, Cl-
Br-, I-

halo, e.g.,
chloro

CuCl 4 2 -

Tetrachlorocopper ( II )

It is possible to replace by chemical reaction some or all of the ligands
in a given complex ion by other ligands. It follows that a complex ion can
contain more than one type of ligand, as do [C0(NH 3 )5(H 2 0)] 3+ and
[CO(NH 3 ) 5 (Cl)] 2+ . The phenomena of hydrolysis and amphoterism may now
be interpreted as the interchange of the ligands, H a O and OH" in a complex
ion. Thus

(11) [Al(H 2 0)e] 8+ + OH- [A1(H 2 0) 5 (0H)] 2 + + H*0

(12) [Al(OH).]»- + H s O+ [A1(H 2 0)(0H 5 ] 2 - + h 2 o

Though it is beyond the scope of this book to consider the subject in any
detail it should be mentioned that complex ions can exist in isomeric forms,
With a square planar structure there can be geometric isomerism because of
different spatial arrangements of the ligands about the central metal -atom.
Thus two planar isomeric structures of M(a) z (b) 2 are

Metals

501

where a and b are ligands attached to the metal atom, M. This cis-trans
isomerism is but one of several types possible in complex ion§. The basic
theory of complex ions was proposed in 1903 by the Swiss chemist, Alfred
Werner; his postulates remain essentially correct today.

8. Nomenclature of Complex Ions. Like other ions a complex ion is part
of a molecule. In writing the formula of a substance containing a complex
ion, the complex ion is inclosed in brackets and the other ion which com-
pletes the molecule is written outside the brackets. When the complex ion
is written alone its net charge is shown outside the brackets. Thus tetrammine-
c0 PP er (^) chloride, [Cu(NH 3 ) 4 3Cl2, is made up of the tetramminecopper(II)
ion, [Cu(NH 3 ) 4 ] 2+ ion, and two chloride ions, 2CK

In naming a complex compound certain rules are observed.

(A) Negatively charged groups within the complex ion have names
ending in -o; neutral groups have no characteristic suffix. The number of
each kind of group is indicated by a prefix.

(B) Negative groups are specified before neutral groups in a complex
ion containing both.

(D) In complex anions the metal ends in -ate; in cations and in neutral
complexes the metal name is unchanged.

Examples of this system of nomenclature are:

[Ni(H 2 0) 6 ]S0 4

[Cr(NH 3 ) 4 Cl 2 ]Cl

K 4 [Fe(CN) 6 ]

K 2 [CuCl 4 ]

Hexaaquonickel ( II ) sulfate
Dichlorotetramminechromium ( III ) chloride
Potassium hexacyanoferrate(II)

Potassium tetrachlorocuprate(II)

QUESTIONS

1. Give a brief summary of the theory of the structure of metals. How does
this theory explain a metal's heat conductivity, electric conductivity, de-
formability, and opacity?

2. From the viewpoint of electron configuration what is a metal? Where are
metals located on the Periodic Table?

3. Explain why, in Group IV B, carbon is a nonmetal but lead is a metal.

4. What types of unit cells are commonly found in metals? What characteristics
does each have?

5. Which of the following metals occurs in the elemental state in nature (a) zinc
(b) copper (c) magnesium (d) mercury (e) lead?

6. List the chemical properties of metals in general. Illustrate each by a balanced
chemical equation,

7. Discuss the general procedures in metallurgy. Using ZnS as a typical ore write
chemical equations where applicable for each step.

502

Metals

8. Define (a) gangue (b) flux (c) slag^d) mineral (e) ore (f) flotation process
(g) alloy.

9, Why is aluminum not obtainable from its ore by reduction with carbon?
What substances could be used as reducing agents in metallurgy?

10. For what metals would electrometallurgical processes be necessary?

11. Which of the following would be basic and which acidic (a) P 4 0 1<} (b) BaO
(c) C10 2 (d) CdO (e) Si0 2 (f) Al(H a O) s ®+?

12. Iron (III) chloride is dissolved in water. Would the solution be acidic, basic,
or neutral? Explain your choice.

13. How is complex ion formation between a metal ion and a ligand related to
Lewis acid-base theory?

14. What hybrid bonds are important m complex ion formation and to what
spatial configurations do they give rise?

15. What is meant by the spatial isomerism of a complex ion?

16. Name the following complex ions (a) Ag(C'N) 2 " (b) Zn(NH 3 ) 4 2+
(c) Fe(H 2 0) 6 s+ (d) Ag(S 2 0 3 ) 2 »- (e) Al(H 2 0) 4 (OH) 2 + (f) Fe(CN) 6 *“

17. Write formulas for the following: (a) dichlorotetracyanoferrate(III) ion

(b) tetrachlorocuprate(II) ion (c) chlorotriamminenickel(II) sulfate (d) tri-
aquotriammineplatinum ( IV ) ion.

18. List some common ligands. What do they have in common?

19. Write electron dot formulas for the following (a) Ni(H 2 0) 4 2 + (b) Ag(NH 3 ) 2 +

20. What is the distinction between dsp 2 and sp 2 d orbitals?

21. Why do coordination numbers greater than four require the use of d orbitals?

22. Using Pt 2 + as the metal ion write formulas for a (a) coinplex cation
(b) complex anion (c) neutral complex (d) complex ion which could show
isomerism.

Elements of Group IA
The Alkali Metals

The elements of Group IA of the Periodic Table constitute the family
known as the alkali metals (Arabic, al-qili: ashes of saltwort). These elements
are lithium, sodium, potassium, rubidium, cesium, and francium. Of these,
sodium and potassium and their compounds are the most important and our
attention wifi be directed mainly to them. Little is experimentally known
about francium, an extremely rare and short-lived radioactive element, so
that it will not be considered here, though its properties could well be
deduced from those of the other members of the group. The electron con-
figurations of the alkali metals are listed below and their properties are
given in Table 38-A.

Element

Atomic

Number

2s 2 p

3s 3p 3d

4s 4p 4c? 4f

5s 5p 5 d

6s 6p

Lithium

Sodium

2 6

Potassium

2 6

Rubidium

2 6

2 6 10

2 6

Cesium

2 6

2 6 10

2 6

Francium

2 6

2 6 10

2 6 10 14

2 6 10

2 0

1. General Properties of the Alkali Metals. A neutral atom of an alkali
metal has an outermost shell which contains but one electron, ns 1 , one elec-
tron more than an atom of the noble gas preceding it in the Periodic Table.
In chemical reactions this electron is lost, leaving a stable octet as the outer-
most shell of the resulting ion.

(1) M M+ + e ; :NaS«-» :Nat+ + e

Almost without exception the alkali metals form electrovalent compounds
in which they are univalent and have an oxidation number of +1.

504

Elements of Group IA: The Alkali Metals

As expected, the ionization potentials of the alkali metals decrease with
an increase in atomic number, from Li to Cs, due to increasing atomic size
and a lesser attraction of the nucleus for the valence electron. Except for
lithium, the oxidation potentials follow a trend similar to that of the ioniza-
tion potentials. It is commonly the case that the first element in a vertical
group of Periodic Table shows anomalous behavior, primarily because

Table 38-A

Properties of the Alkali Metal Elements

Property

Lithium

Sodium

■a

Cesium j

Francium *

Symbol

Atomic Number

Atomic Weight

6.939

22.990

39.102

85.47

132.905

223

Isotopes (mass numbers
and percents)

6( 7.4)
7(92.6)

23(100)

39 (93.08)
40* ( 0.01)
41 ( 6.91)

85 (72.2)
87* (27.8)

133(100)

Abundance in Earth s
Crust, %

0.0067

2.85

2.60 |

0.03

0.0007

Physical State at STP

all are silver-white

solids

Density at STP, g/cm*

0.53

0.86

1.53

1.90

Melting point, °C

180

97.9

62.7

39.0

28.4

Boiling Point, °C

1331

883

776

696

680

Heat of Fusion, kcal/mole

0.72

0.62

0.55

0.50

Heat of Vaporization,
kcal/mole

32.2

21.3

18.5

16.5

15.8

Heat of Atomization,
kcal/mole

37.1

26,0

21.5

20.5

18.8

Ionization Potential, 1st, eV

5.39

5.14

4.34

4.18

3.89

Electronegativity

1.0

0.8

0.7

Atomic Radius, A

1.57

2.02

2.16

2.35

Ionic Radius, A (+1)

0.60

1.33

1.48

1.69

1.76

Oxidation States

Oxidation Potential, volt
for M M+ + e

+3.045

+2.714

+2.925

+2.923

Hydration Energy of M + ,
kcal/mole

123

represents radioactive species

Elements of Group IA: The Alkali ■ Metals

505

of its relatively small size. Indeed it is an empirical fact that die first elements
in Groups I A, II A, and II I A closely resemble the second elements in the
higher adjacent groups.

Though both are functions of energy, a clear distinction should be
made between the ionization potential and the oxidation potential The ioniza-
tion potential is a measure of the energy required to remove an electron from
an isolated gaseous atom. The oxidation potential refers to a process occurring
in aqueous solution. Such a process may be viewed as taking place in three
steps; first, the separation of a single metal atom from die bulk metal,
then the ionization of the metal atom, and finally the hydration of the metal
ion by molecules of the solvent. The oxidation potential is a measure of the
net energy involved in these three steps. These are written below for a
univalent metal atom, M.

(1)

M (bulk solid) -* M (single atom)

AH = +

(2)

M -» M + + e

(3)

M+ + n H 2 0 M(H 2 0) n +

AH _ -

For the small Li + ion the hydration energy is so exothermic as to result in
an oxidation potential greater than that of the other alkali metals.

The alkali metals are very soft and can be cut with a krife. When
freshly cut, they have a silvery white luster characteristic of all metals. Be-
cause of their high chemical reactivity, however, they tarnish rapidly in
air and so must be kept under an unreactive liquid such as kerosene to pre-
vent them from oxidizing. Cesium, indeed, ignites spontaneously in air or
iu water. The melting points of the alkali metals are low and decrease with
increasing atomic number; cesium melts at little above room temperature.
This quality stems from a relatively small attraction between neighboring
alkali metal atoms while the gradation in melting point is due to a decrease
of these attractive forces with increasing atomic size. The softness of these
metals is likewise a result of their weak interatomic forces.

The ease with which the valence electron can be removed is reflected
by the fact that the spectra of the alkali metals^ can be obtained by a low
energy source of excitation, the Bunsen flame. When placed in such a flame,
the elements and their compounds impart to it characteristic colors; lithium,
crimson; sodium, yellow; potassium, violet; rubidium, red; and cesium, blue.
A rapid qualitative test for these elements can be carried out by holding
in the flame a platinum wire which has been previously dipped into the
substance to be tested and noting the color of the flame. Quantitative results
can be obtained by the use of the more elaborate instrument, the flame
photometer. Flame, tests are extremely sensitive; as little as 10" 10 gram of
sodium will impart a detectable color to the flame. A related phenomenon
is the photoelectric effect, whereby radiation of sufficient energy causes the
ejection of an electron from an atom. With cesium and rubidium, the

506

Elements of Group 1 A: The Alkali Metals

alkali elements of lowest ionization potential, this can be accomplished with
the relatively low energy of visible light so that these elements find use
in photoelectric cells.

2. Nomenclature, The name lithium is derived from the Greek word
lithos, meaning stone. The metal was discovered in 1817 by Johan August
Arfvedson who named the element for his belief that its occurrence was
confined to the mineral kingdom. However, lithium is found in the ash of
some plants. Both sodium and potassium were discovered in 1807 by the
English chemist, Sir Humphrey Davy, by the electrolysis of their moist solid
hydroxides. The symbols Na and K are derived from the words natron and
kali. Natron is the Spanish term applied to the mineral alkali, Na 2 CO^,
obtained from the ashes of sea plants and from the mineral trona ,
Na 2 C0 3 • NaHC0 3 * 2 H 2 0. Kali is the German word for potash or the
vegetable alkali, K 2 CO s , obtained from the ash of land plants. The elements,
rubidium (Latin, rubidus: red) and cesium (Latin, caesius : blue-gray),
were discovered in 1860 during a spectroscopic examination of minerals
obtained from Durkheim mineral water by the pioneers of spectrum analysis,
the German scientists, B. W. Bunsen and G. R. Kirchoff. The elements are
named after the colors of their characteristic spectral lines. Francium, named
for France, was discovered in 1939 by MUe. Marguerite Perey in an investiga-
tion of anomalous beta ray activity of the radioactive element, actinium.

The common use of alkali metal compounds has given rise to historically
accepted names which might be noted here for the sake of clarity.

NaOH caustic soda

NaNOs Chile saltpeter KN0 3 saltpeter

NanCOs soda ash K 2 CO :r potash

Na 2 CO s * 10 H g O washing soda

NaHCOs bicarbonate of soda; baking soda

3. Occurrence of the Alkali Metals. Because of their chemical reactivity
the alkali metals do not occur in the elemental state. Compounds of the alkali
metals are widely distributed but, except for unusual natural events, the
simple compounds are not found on land because of their high solubilities.
Insoluble silicate rocks such as the feldspars, NaAlSi s Os and KAlSi 3 0 8 , to-
gether with CaAl 2 Si 2 0 8 , make up as much as 60% of the earth's crust. Through
weathering processes these silicates are converted into simpler soluble salts.
Sodium compounds eventually find their way to the sea but potassium
compounds, a major plant nutrient, are selectively absorbed by plant life
so that only a small concentration of potassium compounds is found in sea
water. There is no substitute for potassium compounds in plant nutrition.
Approximately 95% of the potassium compounds consumed in the United
States is used in agriculture, usually as mixed fertilizers. For this purpose
the most widely used compound is KC1, then K 2 S0 4 .

Under certain geologic conditions a body of sea water may become
isolated and, with a suitably arid climate, large deposits of salt can be

Elements of Group IA: The Alkali Metals

507

built up. In this way were formed the large underground beds of rock salt,
NaCl, the deposits of Chile saltpeter; NaNO s , of borax, Na 2 B 4 0 7 • 10 H 2 0,
in Death Valley, and the Stassfurt salt beds composed of KC1 and
KC1 • MgCl 2 ■ 6 H 2 0.

Compounds of lithium, rubidium, and cesium are widely but sparingly
distributed. The most important lithium minerals, in which the lithium con-
tent varies from 4-10%, are lepidolite , a complex hydrated fluosilicate,
K 2 Li 3 Al 4 Si 702 i( 0 H or F) 3 , spodwnene , LiAlSi 2 0 6 , and amblygonite, LiAlFP0 4 .
Rubidium and cesium usually occur together and associated with the other
alkali metals, as in lepidolite or feldspar. Though rubidium does not occur
as the major ingredient in any mineral its common occurrence as a minor
element in many minerals makes it more abundant than lead.

4. Preparation of the Alkali Metals. To prepare an alkali metal requires
the reduction of its +1 ion.

(4) M + e M°

Because of the large tendency of an alkali metal atom to lose its valence
electron the reverse process of reducing a metal ion to a neutral atom re-
quires much energy and cannot be accomplished by the usual chemical
reducing agents except under extreme conditions. Hence the alkali metals
are generally prepared by the electrolysis of their fused compounds. Elemen-
tal sodium is produced by the electrolysis of molten NaCl; the Downs
electrolytic cell is shown in Figure 38.1. The net equation for the electrolysis is'

(5) 2 Na+ + 2 Ch 2 Na(Z) + Cl 2 (g)

Formerly, sodium metal was produced by the electrolysis of molten NaOH.
Despite the advantage of a lower operating temperature since the melting
point of NaOH is 318°C this process is not used extensively today because
of the necessity for the prior conversion of NaCl to NaOH.

Lithium metal is obtained by the electrolysis of a fused mixture of LiCl
and another alkali halide such as KC1. Potassium, rubidium, and cesium are
usually not produced by electrolysis; potassium dissolves in molten KC1 or
KOH. These metals are best prepared by the reduction of their carbonates
or hydroxides at moderately high temperatures by Ca, Fe, or C. Such reduc-
tion by a “less active” element is due to the higher volatility of the alkali
metal; its removal from the reaction mixture drives the reaction to the
right. For example,

(6) Rb 2 CO* + Ca 2 Rb(g) + CaO + CO a

5. Uses of the Alkali Metals. Sodium is the only alkali metal which has
extensive use; about 125,000 tons are produced annually. Its most important
use is in the production of the antiknock compound, lead tetraethyl,
Pb(C 2 H 5 ) 4 . As a reducing agent, sodium finds application in organic chem-
istry for the reduction of animal and vegetable oils to fatty alcohols and in
the production of Ti, Zr, Cb, Be, and other difficult to reduce metals. Sodium
is also used to prepare borohydrides, ingredients of a class of high energy
fuels for turbojet and rocket engines. Sodium forms alloys with many metals

508

Elements of Group IA: The Alkali Metals

The Downs electrolytic cell has a circular anode of graphite A surrounded by an
annular iron or steel cathode C. A steel gauze diaphragm, S, is suspended between
the electrodes to separate the electrolytic products. The electrolyte is a molten
mixture of NaCl and Na 2 C0 3 or CaCl £ . The addition of a salt to the NaCl
(m.p. = 801 °C) yields a bath which melts about 600°C. The heat necessary to
keep the electrolyte molten is produced by the passage of the electric current.
Metallic Na, liberated at the cathode, floats on the molten electrolyte and over-
flows through the outlet pipe, P, into the storage compartment, R. Chlorine gas,
formed at the anode, is a by-product of the process. It is withdrawn through the
pipe, (); additional NaCl is added through the opening, B .

Figure 38.1.

The Downs Cell for the Production of Metallic Sodium by Electrolysis of
Fused Sodium Chloride.

but not with iron. Alloys of sodium with mercury (sc^um amalgam), and
with lead, and the more recent dispersions of finely di\ ided sodium particles
in a nonreactive fluid are convenient forms of reducing agent. Because liquid
sodium has excellent heat transfer characteristics it is used as a nuclear
reactor coolant and inside the hollow steel exhaust valves of aircraft engines.
Because of its bright yellow hue sodium vapor lamps are used for illumination.

Inasmuch as there are few uses for potassium which sodium does not
satisfy equally well and because of 'the greater difficulty and expense in
producing it, potassium metal is not used commercially. The principal use
of cesium is in photoelectric cells; other uses are in infrared telescopes such
as the "snooperscope” and as a "getter” in vacuum tubes. Rubidium has
similar applications. Both metals have possibilities as fuels for ion-drive
rocket engines in space flight. The metal vapor is ionized by passing it
through a heated platinum grid or an electric arc. By subsequent passage
through electric and magnetic fields the resulting ions can be accelerated to
300,000 miles per hour and then expelled to provide thrust for a space
vehicle.

Elements of Group l A: The Alkali Metals

509

6. Reactions of the Alkali Metals. The alkali metals are exceedingly re-
active. They combine directly with oxygen, hydrogen, sulfur, phosphorus, and
the halogen elements. They react with water, alcohol, and ammonia to dis-
place hydrogen, and with other metal compounds, e.g., A1C1 S , to displace
less active metals.

(A) Reaction with oxygen: Heating an alkali metal with 0 2 yields, in
the case of lithium only, the monoxide, Li 2 0; in the case of sodium, the per-
oxide; Na 2 0 2 , is formed; the others give “superoxides,” e.g., JC0 2 . The only
oxide of commercial importance is Na 2 0 2 ; the peroxides can be considered
to be salts of hydrogen peroxide, H 2 O a .

(B) Reaction with hydrogen: Heating an alkali metal with H 2 forms a
hydride, MH. The hydrides are white, crystalline, ionic compounds which
are unusual in that hydrogen is present as a negative ion, H~ During elec-
trolysis of a fused hydride, H 2 is produced at the anode. The hydrides are
decomposed by water.

(7) NaH + H 2 0 NaOH + H 2

Chemically the hydrides are good reducing agents, better than H 2 itself.

(8) 2 H- ^ H 2 + 2 e E° = +2.25 volts

(C) Reactions with water, alcohol, and ammonia: These reactions are
similar in that H 2 is displaced by the alkali metal. With H 2 0 the reaction
is energetic and the metal hydroxide is produced; with C 2 H 5 OH, sodium
ethoxide is formed; and in the presence of a catalyst the metals react with
dry NH S gas to form amides, M(NH 2 ).

(9) 2 M+ 2 H 2 0 ^ 2 MOH +H 2 (2 Na + 2 H 2 0 -» 2 NaOH + H 2 )

(10) 2 M + 2 ROH -» 2 ROM + H 2 (2 Na + 2 C 2 H 5 OH-> 2 C 2 H 5 ONa + H 2 )

(11) 2 M + 2 NH 3 2 MNH, + H 2 (2 Na + 2 NH a 2 NaNH 2 + H 2 )

In anhydrous liquid ammonia the metals dissolve, forming solutions of
intense blue color which conduct electricity. By evaporation of the ammonia
the metals can be recovered unchanged* The conductivity of the solution is
attributed to the presence of metal +1 ions and solvated electrons; these
electrons give rise to the blue color.

7. Compounds of the Alkali Metals. The alkali metals form typical metal
salts. Since these compounds are ionic they have properties which are a

• composite of those of the specific alkali metal ion and the anion which make
up the compound. The properties of the more important anions have al-
ready been considered.

The compounds of the alkali metals have far greater commercial im-
portance than do the metals themselves. With few exceptions, the salts are
readily soluble in water and are colorless except where the anion is colored.
The hydroxides form strongly basic, or alkaline , solutions. Whereas all other
hydroxides lose H 2 0 and carbonates lose C0 2 when heated below their
melting, points, the hydroxides and carbonates of the alkali metals can be
heated without decomposition.

510

Elements of Group IA: The Alkali Metals

Of all the alkali metal compounds, those of sodium are the most important
and it is to these that we shall direct our attention with the understanding
that, with few exceptions, the analogous compounds of the other alkali metals
have similar properties.

Sodium chloride , NaCl: The elements sodium and chlorine are essential
to human life, and NaCl most readily satisfies this requirement. References
to salt are found in the Bible and expressions such as “the salt of the earth’'
have become idiomatic. Salt has been used as currency and has been a
commodity for taxation. Directly or indirectly, NaCl is the raw material
from which almost all important sodium and chlorine compounds are made.
It need not be manufactured since it occurs naturally and abundantly as
rock salt and as brine. With an annual production of over 24 million tons
the salt industry ranks among the largest in the nonmetals field. The greatest
consumer of NaCl is the chemical industry itself, in the electrolysis of brine
for the production of Cl 2 and NaOH and in the Solvay process for the pro-
duction of NaHC0 3 and Na 2 C0 3 . Pure NaCl is not a hydrate nor does it
absorb water. Ordinary salt usually contains a little MgClo which is hygro-
scopic and is likely to cause the salt to cake.

Sodium hydroxide , NaOH: This is prepared by two methods, one chemical
and one electrochemical.

(A) The lime process, for which the reaction is

(12) Na 2 C0 3 + Ca(OH) 2 -> 2 NaOH + CaCO«(s)

The reaction is driven to the right because of the insolubility of CaCO H ,
which precipitates and can be filtered off to leave a solution of NaOH.

(B) The electrolysis of brine for which the net reaction is

(13) 2 Na+ + 2 Cl- + 2 H 2 0 -> 2 Na + + 2 OH" + H,(g) + Cl,(g)

The Hooker S cell, a type of diaphragm cell wherein an asbestos diaphragm
serves to separate the products of electrolysis, has been described. The
mercury cell, used to a large extent in countries other than the United
States, is an elongated trough IVz to 6 feet wide by 25 to 85 feet long, along
the bottom of which flows a shallow stream of mercury which acts as the
cathode. Above this is an assembly of graphite blocks, the anode, while
saturated brine is the electrolyte. The Na, liberated at the cathode, forms
a dilute amalgam, 0.1-0.2 percent Na, with the mercury. Thereby its activity
is reduced and it does not react with the aqueous electrolyte. The amalgam
flows from the brine cell into another cell, the soda cell or decomposer,
where it meets a countercurrent of water in which the Na reacts to form
NaOH and H 2 . The mercury is then recirculated through the brine cell.

Pure NaOH is a white, hygroscopic solid, very soluble in water with the
evolution of much heat. An aqueous solution is strongly basic and will
absorb C0 2 from the air,

(14) 2 OH- + C0 2 -» COs 2 - + H 2 0

Solutions of NaOH slowly attack glass so that glass containers become
etched and take on a cloudy appearance in time. In chemical technol-

Elements of Group IA: The Alkali Metals

511

ogy, NaOH is used where a strong base is needed, in the manufacture of
rayon and other textiles, in petroleum refining, and for saponification in the
production of soap. A commercial solution of NaOH is also known as lye.

Sodium carbonate, Na 2 C0 3 , and sodium hydrogen carbonate , NaHC0 3 :
The most important industrial source of these two compounds is the Solvay
or ammonia-soda process, discovered by the Belgian chemist, Ernest Solvay.
In this process, a solution of NaCl is saturated with NH* gas and then C0 2
is added under pressure. The formulation following shows the various ions
which are present in the equilibrium mixture and the reactions which take
place.

NH 3 + H 2 0 NH 4 + + OH*

C0 2 + 2 H 2 0 H 3 0+ + HCO3-

NaCl Cl* + Na+

i 4

2 H 2 0 NaHCOs(s)

The NaHC0 3 precipitates because it is but slightly soluble in the liquid of
the reaction mixture and so can be removed by filtration. By heating, it can
be converted to the carbonate, Na 2 C0 3 .

(16) 2 NaHCOs -> Na 2 C0 3 + H 2 0 + C0 2

The Solvay process is an excellent example of material economy in chemical
technology. The reactants are NaCl, CO a , and NH 3 , of which the C0 2 and
NH 3 are recovered and used over and over again. The C0 2 is obtained
initially from the decomposition of limestone, CaCO s , during the manufac-
ture of lime, CaO,

(17) CaC0 3 -> CaO + C0 2

and is recovered when the NaHC0 3 is heated to produce Na 2 C0 3 . The lime,
CaO, is used to recover the NH 3 by heating it with the ammonium chloride,
NH4CI, which remains in the mother liquor.

(18) CaO + 2 NH4CI 2 NH 3 + CaCl 2 H a O

The NHs is expensive and upon its recovery depends the financial success
of the process. In the over-all process the raw materials permanently con-
sumed are solely NaCl and CaC0 3 , both relatively cheap, and the only
by-product is CaCl 2 . The latter has few uses and is often discarded as waste.
Because it is not sufficiently insoluble, KHC0 3 , and hence K 2 C0 3 , cannot
be made by a similar procedure.

Both NaHC0 3 and Na 2 C0 3 can be produced electrolytically by passing
C0 2 into a solution of NaCl which is being electrolyzed. The . C0 2 reacts
with the NaOH formed during the electrolysis. With excess C0 2 , NaHC0 3
is formed and with a lesser quantity of C0 2 , Na 2 C0 3 is produced.

512

Elements of Group I A: The Alkali Metals

(19) OH" + C0 2 -> HC0 3 " (excess C0 2 )

(20) 2 OH" + C0 2 CO 3 2 - + H 2 0

The annual production of Na 2 C0 3 of over AVz million tons, almost all by
the Solvay process, rivals that of sulfuric acid and of chlorine in importance.
It is consumed mainly in manufacturing other chemicals, both inorganic
and organic. The largest single use is in the manufacture of glass. By
hydrolysis, an aqueous solution of Na 2 CO n is strongly basic and many of its
uses, such as in cleansers and water softeners, depend upon this property.
The chief use of NaHCO s is as an ingredient of baking powder where it acts
as a leavening agent through the formation of C0 2 which causes the baking
dough to rise. Baking powders contain NaHC0 3 and an acidic substance
such as sodium aluminum sulfate. When moistened, the hydrolysis of
A1(H 2 0) 6 3+ ions yields H + ions which react with HC0 3 " ions to form C0 2 .

Potassium salts are generally less hygroscopic than are the corresponding
sodium salts. Potassium nitrate, KNO s , is an ingredient of black gunpowder
together with charcoal and sulfur; it acts as the oxidizing agent for the
C and S. The action of gunpowder is essentially a rapid combustion. Once
ignited, gunpowder bums with the evolution of much heat and a large
volume of gas which can serve to propel a projectile. The reaction is by
no means a simple one; gaseous products include oxides of N, C, and S
and solid products are K 2 C0 3 and K 2 S0 4 . Black powder produces volumes
of heavy smoke; .organic nitro compounds, such as smokeless powder, form
no solid residue on explosion.

QUESTIONS

1. Account for the following properties of the alkali metals: (a) low first
ionization potentials (b) good conductors of heat and electricity (c) softness
(d) visible flame spectra (e) only one oxidation state (f) colorless ions
(g) few complex ions (h) hydroxides are basic.

2. For the element francium, fill in the blank spaces in Table 38-A and predict
its chemical properties.

3. How are the alkali metals protected from oxidation? What other elements
must be stored under a liquid?

4. How can the alkali metals sodium and potassium be prepared? Write per-
tinent equations.

5. Account for the fact that lithium has the lowest first ionization potential but
the highest oxidation potential of the alkali metals.

6 . Correlate the blue color of sodium vapor with the fact that two of its
prominent spectral lines are yellow.

7. Write formulas for the following: (a) rock salt (b) saltpeter (c) soda ash
(d) potash (e) caustic soda (f) baking soda (g) bicarbonate of soda (h) hypo.

8 . One of the most efficient carriers of hydrogen is lithium hydride. Explain this
statement.

9. What would you predict to be the relative stability of the alkali metal hydrides?

10. Balance the reaction of NaH with water by the ion-electron method.

Elements of Group IA: The Alkali Metals

513

11. Write balanced chemical equations for the reaction of sodium metal with

(a) H 2 (b) 0 2 (c) S (d) H 2 0 (e) CH 3 OH (f) NH, (g) MgCl 2 .

12. Account for the different products resulting from the reaction of the alkali
metals with oxygen.

13. In each case starting with NaCl, write equations for the preparation of
a) NaOH (b) NaHSO* (c) Na 2 S0 4 (d) NaNO s (e) NaH.

14. (a) Write equations for the production of Na 2 CO g by the Solvay process

(b) what economies are practiced to make the process a financial success?

(d) what are the uses of NaHCO s and of Na 2 C0 3 ?

15. Write equations for the reaction of NaOH and (a) C0 2 (b) Si0 2 .

16. A baking powder is composed of NaHC0 3 and Al 2 (SOJ 3 . Write ionic equa-
tions to illustrate its action.

17. What weight of lithium can be obtained theoretically from 5.00 kg of the

mineral lepidolite, LiAl(Si0 3 ) 2 ? Ans: 186 g

18. What weight of washing soda can be obtained from ten tons of NaCl by

the Solvay process, assuming 70% recovery? Ans : 17.1 tons

19. In the Solvay process (a) what weight of CaCl 2 is produced as a by-product
for every ton of NaHCO a formed, assuming 98% recovery of NH 3 (b) what
volume of C0 2 , measured at STP, is required to form 4.20 lb of NaHCO s ?

Ans: (a) 0.65 ton (b)*508 liters

20. What weight of NaHCO s is required to produce ten liters of C0 2 at 15 °C

and 770 mm by (a) heating the NaHC0 3 and (b) reacting NaHC0 3 with
H 2 S0 4 ? Ans: (a) 72 g (b) 36 g

21. What weight of NaOH can be made by treating 2.00 kg of Na 2 C0 3 with

slaked lime, Ca(OH) 2 ? Ans: 1.51 kg

22. Two liters of 1.0N NaCl are electrolyzed. If sufficient electricity is passed
through the solution to produce 0.1N NaOH, what is the concentration of
the remaining NaCl?

23. What weight of K 2 C0 3 is required to make 200 ml of 0.60 N solution?

The Elements of Group HA
The Alkaline Earth Metals

The elements beryllium, magnesium, calcium, strontium, barium, and
radium comprise Group IIA of the Periodic System and are known collectively
as the alkaline earth metals. The term earth was applied by the early chemists
to nonmetallic substances insoluble in water and unaffected by heat. Those
which gave alkaline reactions, e.g., lime, CaO, were designated the alkaline
earths. Formerly the term alkaline earth was applied only to the elements
calcium, strontium, and barium. Beryllium and magnesium have certain
properties which are different from those of the other three elements and
were classified under Group IIB with zinc, cadmium, and mercury. However]
the present deeper insight into electron configurations has resulted in the
classification of beryllium and magnesium as alkaline earth elements. The
electron configurations of these elements are shown below and their properties
are listed in Table 39-A.

Element

Atomic

Number

4s 4 p Ad 4f

5 s 5 p 5d

6s 6p

Beryllium

Magnesium

Calcium

2 6

Strontium

2 6

Barium

2 6

Radium

2 6

2 6 10

1. General Properties of the Alkaline Earth Elements. The elements of
Group IIA follow immediately after the alkali metals in the Periodic System.
As compared with an atom of the alkali metal element immediately preceding
it, an alkaline earth atom has one proton more in its nucleus and an addi-

The Elements of Group IIA: The Alkaline Earth Metals

515

tional electron in its outermost valence shell. If adjacent alkali and alkaline
earth elements, such as potassium and calcium, are compared, the greater
nuclear charge of the alkaline earth atom attracts the electron shells relatively
closer to the nucleus so that the alkaline earth atom has a smaller atomic
size and a smaller ionic size. Because their valence electrons are more
strongly bound, the alkaline earth elements are less active chemically and
are weaker reducing agents than the alkali metals. But in general, the
alkaline earth elements are typical metals with properties similar to those
encountered with the alkali metals.

The alkaline earth atoms have two valence electrons, ns 2 . They react to
lose these, thereby forming positive ions of the type, M 2+ ; the +2 oxidation
state is the only one exhibited. The high oxidation potentials of Ca, Sr, and
Ba, higher than that for Na, are attributable to the high heats of hydration,
of their ions.

Because of its relatively small size the first element of the group, beryllium,
shows distinctly anomalous properties; its properties closely resemble those
of aluminum in Group IIIA. Except for beryllium, the alkaline earth metals
are strong reducing agents, their compounds are electrovalent, and their
hydroxides are basic. Beryllium forms compounds with appreciable covalent
character and beryllium hydroxide, Be(OH) 2 , is amphoteric. The elements
have the usual silver-white metallic luster when freshly cut but tarnish
rapidly, except for beryllium and magnesium which form transparent oxide
coatings and so can be kept in air without protection. Harder than the
alkali metals, beryllium is hard enough to scratch glass and most other
metals whereas barium is but slightly harder than lead. As were the alkali
metal ions, the ions of the alkaline earth elements are also colorless; any
color in their compounds is to be attributed to the anion. The more active
of the elements, Ca, Sr, and Ba, also give characteristic flame tests: calcium,
orange-red; strontium, crimson; and barium, light green. In contrast to the
salts of the alkali metals, many of the common salts of Mg, Ca, Sr, and Ba
are insoluble in water. Their fluorides, hydroxides, carbonates, and sulfates
(except MgS0 4 ) are but slightly soluble. The solubilities of the carbonates
and sulfates decrease, whereas those of the hydroxides increase, with increas-
ing atomic number of the metal.

2. Occurrence of the Alkaline Earth Elements. None of the alkaline
earth elements occur in the elemental state. The important minerals are the
insoluble carbonates, sulfates, and silicates. Small quantities of beryllium are
found in over thirty minerals but its most important source is the silicate,
beryl , 3 BeO • A1 2 0 3 • 6 Si0 2 or Be 3 Al 2 Sis0 lf} , When colored green by chromium
compounds the mineral is known as emerald; aquamarine is a blue-green
form.

Magnesium is the eighth most abundant element in the earths litho-
sphere. Though there are more than sixty manesium-bearing minerals, the most
important commercial sources of the element are dolomite, CaC0 3 * MgCOs,*
magnesite , MgC0 3 , brucite, MgO ■ H 2 0, olivine, (Mg,Fe) 2 Si0 4 , and the dis-
solved Mg 2+ ion in sea water and in well brines. It has been estimated that
a cubic mile of sea water contains about 5,700,000 tons of Mg 2+ ion. Silicates

516

The Elements oi Group UA: The Alkaline Earth Metals

Table 39-A

Properties of the Alkaline Earth Elements

Property

Beryllium.

;

Magnesium

Calcium

Strontium

Barium

Radium *

Symbol

Atomic Number

Atomic Weight

24.312

40.08

87.62

137.34

| 226.0

Isotopes (mass numbers

9(100)

24(78.6)

40(96.92)

84( 0.56)

130( 0.10)

and percents)

25(10.1)

42( 0.64)

86( 9.86)

132( 0.10)

26(11.3)

43( 0.13)

87( 7.02)

134( 2.42)

44( 2.13)

88(82.56)

135( 6.59)

48( 0.003)

138( 7.81)

48( 0.18)

137(11.32)

138(71.66)

Abundance in Earth’s

| 0.00001

2.09

3.63

0.0002

0.05

1 X 10“12

Crust, %

Physical State at STP

all are silver-white

solids

Density at STP, g/cm 3

1.86

1.74

1.55

1 2.60 1

3.74

5.0

Melting Point, °C

1280

651

845

850

960

Boiling Point, °C

2480

1110

1440

1380

1640

1100

Heat of Fusion, kcal/mole

2.8

2.1

2.2

1.8

Heat of Vaporization,

70.4

30.8

35.8

33.2

36.1

kcal/mole

Heat of Atomization,

kcal/mole

Ionization Potential, eV, 1st

9,32

7.61

6.11

5.69 !

5.21

5.28

2nd

18.12

14.96

11.82

10.98 j

9.95

10.10

Electronegativity

1.5

1.2

1.0

0.9

Atomic Radius, A

0.89

1.37 j

1.74

1.91

1.98

2.28

Ionic Radius, A (+2)

0.31

0.65

0.99

1.10

1.29

1.52

Oxidation States

+2 i

Oxidation Potential, volt

+1.85

+2.37

+2.87

+2.89

+2.90

+2.92

for M — » M 2 + + 2 e
Hydration Energy of M 2 +,

500

460

395

355

305

kcal/mole

'■'represents radioactive species

of magnesium are common and are of economic importance, notably talc or
soapstone , H 2 Mg 3 (Si0 3 ).i, asbestos , Mg 3 Ca(Si0 3 ) 4s and meerschaum ,
Mg 2 Si s 0 8 * 2 H 2 0.

Calcium, the fifth most abundant element, occurs in a variety of forms
of CaC0 3 , itself the most common of all minerals not containing silicon.
Among these are limestone, marble, the cave formations of stalactites and
stalagmites eggshell, seashell, pearl, coral, and chalk, a marine formation
largely composed of shells of minute organisms, foraminifera , of which a
noteworthy deposit is the cliffs along the English Channel. In . the foregoing
examples crystalline character is almost lacking but there exist two varieties

The Elements of Group 1IA; The Alkaline Earth Metals

517

of natural crystalline CaC0 3 ; as calcite , in the form of Iceland spar, satinspar ,
and dogtooth spar, and as the rarer aragonite. Other calcium minerals are
the hydrated sulfate, gypsum , CaS0 4 * 2 H.,0, and the anhydrous form,
anhydrite, CaS0 4 ; fluorspar, CaF 2 ; and phosphate rock , Ca 3 (P0 4 ) 2 ; and
many complex silicates. Calcium compounds are present in sea water, from
which it is taken up by various marine animals and converted into CaC0 3
to form their outer protective shells. Plants require small amounts of calcium.
From the plant the calcium finds its way into the bodies of animals where,
as the phosphate, Ca 3 (P0 4 ) 2 , it forms the principal component of bones
and teeth.

The important strontium and barium minerals are the sulfates and the
carbonates; celestite, SrS0 4 , strontianite, SrCO«, barite, BaS0 4 , and witherite,
BaC0 3 . Most newly produced radium comes from pitchblende and radium-
bearing slimes from the Congo. Radium is one of the radioactive daughter
elements formed in the radioactive disintegration series of uranium and
so occurs in conjunction with uranium ores, about one part in three million.

3. Nomenclature of the Alkaline Earth Elements. Beryllium was dis-
covered in 1789 by the French chemist, Nicolas-Louis Vauquelin during an
analysis of emerald; its name is derived from the mineral beryl. Formerly
the element was called glucinum (Greek, glykys : sweet) because of the
sweet taste of its compounds. The elements magnesium, calcium, strontium,
and barium were first isolated by electrochemical means by Sir Humphrey
Davy in 1808, the year after his discovery of sodium and potassium. The
origin of the elements 7 names are: magnesium (Greek, magnesia lithos : the
Magnesian stone) for a mineral found in Magnesia, a district of Thessaly;
calcium (Latin, calx : lime) for the substance, lime, one of the most im-
portant structural materials today and in ancient times; strontium, after
Strontian, a village in Scotland, from which the mineral strontianite was
first recognized to contain a new “earth” different from lime and barite;
barium (Greek, barys : heavy) after the mineral barite, often called heavy
spar because of its high specific gravity, 4.3 to 4.6. Radium, as the bromide,
RaBr 2 , was first separated by Pierre Curie and his wife, Marie Sklodowska
Curie, from pitchblende in 1898. The pure element was isolated in 1911 by
Mme. Curie and Andre Debieme by electrolysis; it was named for its high
radiation intensity.

4. Preparation and Uses of the Alkaline Earth Elements. Preparation of
an alkaline earth element requires the reduction of its -f*2 ion, either by
electrolysis of a molten halide salt or by an appropriate chemical reducing
agent.

Beryllium metal is produced by the reduction of BeF a with magnesium,
usually in a high frequency induction electric furnace. It can also be made
by the electrolysis of fused BeCl 2 . Both beryllium metal and its compounds
are toxic, prolonged exposure causing the disease known as berylliosis, about
which little is known other than that it causes a kind of tumor that affects
the skin tod the lungs. Though beryllium is the metal of second lowest

518

The Elements of Group 1IA: The Alkaline Earth Metals

atomic number, it is the fifth lightest after Li, Na, K, and Rb. Its unusually
high melting point and strength-to-weight ratio have led to its use in
missiles and jet aircraft. In this connection its most celebrated use was as
the shield for the Project Mercury space capsule.

Until 1926, when the hardening effect of beryllium upon copper was
discovered it was little more than a laboratory curiosity but this discovery
has led to the establishment of the beryllium industry. By far its largest
use is as an alloying agent. A Be-Cu alloy containing 1.9% Be has a hardness
and a tensile strength equal to that of steel. Be-Cu alloys are widely used in
springs, electrical contacts, gears, and in miniature electrical components.
Be-Ni alloys closely resemble stainless steels. Pure beryllium is used as
windows in X-ray tubes; like other elements of low atomic number it is
transparent to X-rays. Beryllium absorbs fewer neutrons than any other
known structural material so that it is used as a moderator in nuclear reactors
to reduce the speed of fission neutrons and as a reflector material to reflect
neutrons back into a reactor core and thus reduce neutron leakage. About
80 tons of beryllium are produced annually.

In the United States the principal source of magnesium is raw sea water
which contains about 0.3% MgCl 2 . Almost pure calcium oxide, CaO, obtained
by roasting oyster shells, is used to precipitate the Mg~ + ions from sea water
as Mg (OH) 2 . After filtration, the Mg (OH) 2 is neutralized with HC1 to form
a solution of pure MgCL. This is evaporated and the anhydrous MgCl 2 is
melted and electrolyzed to produce magnesium ' metal of 99.8% purity. The
temperature of the electrolytic cell is above the melting point of the mag-
nesium, which rises to the top of the cell and is run off into molds. The Cl 2 ,
an electrolytic by-product, is converted into HC1 and used over again. This
process is known as the Dow process and the Dow Chemical Company’s 36,000
ton electrolytic plant at Freeport, Texas, is the main producer of magnesium
from sea water. This city is also the site of the nation’s first demonstration
plant for the conversion of sea water into fresh water; it uses a distillation
method and was placed into operation on June 21, 1960.

In the silicothermic process, calcined dolomite, which contains MgO and
CaO, is mixed with ferrosilicon, an alloy of Fe and Si, formed into briquets
and charged into retorts which are electrically heated to 1150°C and evacuated.
The Si reduces the MgO producing Mg vapor which condenses on a re-
movable condenser from which the metal is scraped off as 99.98% Mg.
The iron of the ferrosilicon merely acts as a carrier of the Si and takes no
part in the reaction; the CaO combines with the Si0 2 produced and remains
as a slag.

Magnesium is the lightest of the rigid and hence structurally useful metals.
It is used in lightweight metal construction and in making special alloys,
particularly with aluminum, c.g., Dowmetal and magnalium. With the in-
crease in production of magnesium since World War II to the present level
of 50,000 tons annually, the price of the metal has dropped from $3.50 a
pound in 1915 to less than $0.40 a. pound today. As a reducing agent, mag-
nesium is used to reduce TiCh in the production of titanium, a metal much
used in the steel industry for the removal of nitrogen from steel Smaller

The Elements of Group II A: The Alkaline Earth Metals

519

quantities of the elements Zr, Hf, U, and Be are similarly produced. Mag-
nesium anodes are used for electrolytic protection of iron and steel, par-
ticularly for ground pipes, water tanks, and hulls of ships. The vigorous re-
action of magnesium powder or ribbon with oxygen to produce an intense
light rich in ultraviolet radiation has led to its use in photography and in
signal flares. Because of the high temperature produced in this combustion,
sticks of magnesium have been employed as incendiary bombs in warfare.
Since burning magnesium reacts with H 2 0, CO.,, and sand (Si0 2 ), none
of these is effective in extinguishing it.

The elements calcium, strontium, barium, and radium are produced by
aluminothermy is a vacuum or by the electrolysis of a molten halide, usually
the chloride. The compounds of these elements are of far more commercial
importance than the elements themselves; of the group, calcium and its
compounds are the most important chemically. Since calcium will reduce
any oxide to its metal, its principal use is as a reducing agent. Industrially
the metals Cr, Ti, Zr, Th, and U are produced by calcium reduction of their
oxides. Calcium is also used to deoxidize iron, steel, and copper, to debis-
muthize lead, to remove nitrogen from argon, to desulfurize petroleum, to
dehydrate alcohol, and to reduce organic compounds. Strontium and barium
are used as getters to remove the last traces of gas in vacuum tubes. The
brilliant crimson color which strontium imparts to a flame is utilized in
many pyrotechnical applications, e.g., fireworks, signal flares, and tracer
bullets. The most important application of radium is in radiotherapy for
the treatment of cancer. Since radioactivity is a property of the nucleus,
compounds of radium are equally as radioactive as the element itself. Despite
much intensive research on the extraction of radium, the tedious but ingenious
method of fractional crystallization first used by the Curies is still used. The
one part in three million of radium in pitchblende is separated as RaBr 2
from which elemental radium can be prepared by electrolysis, if desired.

5. Reactions of the Alkaline Earth Elements. Except for beryllium, the
alkaline earth elements are typical metals and their reactions are similar to
those of the alkali metals, though less vigorous. They combine directly with
the nonmetals oxygen, hydrogen, the halogens, sulfur, nitrogen, and carbon.
With H 2 0 and NH«, H 2 is liberated and the hydroxide and the amide are
produced, respectively. The reaction of the elements with water offers a
good illustration of their relative activity. Beryllium reacts not at all even
with steam, magnesium reacts with boiling water, and the others react readily
with cold water.

6. Compounds of the Alkaline Earth Elements. Oxides : The oxides are
formed when the metals burn in air but, except for BeO which is obtained
from beryl, they are manufactured commercially by heating the carbonates;
BeCO a is unknown. Lime, CaO, is the most important of the oxides. It is
the second largest tonnage chemical in the United States, about 10 million
tons being produced annually. Its manufacture and use are almost as old
as recorded history; lime was used as plaster and mortar in building the
Egyptian pyramids 4,500 years ago. It is prepared commercially by the
calcination of limestone, CaCO s , at about 1,000 °C in large furnaces or kilns .

520

The Elements of Group HA; The Alkaline Earth Metals

(1) CaCO,(*) - CaO (-*) + C0 2 (g) AH ^ 21 kcal

The reaction is reversible. By forcing a draft of air through the lime kiln,
the C0 2 is swept out and its partial pressure in contact with the limestone
is kept at a* low value so that the reaction goes to completion. There are
about 7,000 uses for lime. By far their greatest number is chemical and in-
dustrial but the largest single use is in construction in the preparation of
mortar, plaster, and cement. Among the varied chemical uses are the
manufacture of steel, alkalies, calcium carbide, paper, glass, and bleaching
powder. Lime is used extensively in agriculture, primarily, for soil conditioning.

Magnesia , MgO, is the principal product of the magnesium industry.
If produced by the calcination of MgCO a at temperatures above 1,500°C
it is known as dead burnt magnesia. This does not react with water, has a
high melting point, and is used as a basic refractory for lining steel furnaces.
Mixed with asbestos (85% MgO), a light fluffy material is produced which
is employed as an insulator for covering steam pipes.

Hydroxides: BeO does not react with water; MgO reacts very slowly but
the other alkaline earth oxides react rapidly and -exothermically to form
hydroxides.

(2) MO + H 2 0 M(OH) 2 (CaO + H.O Ca(OH),. AH = -15.1 kcal

Both the molar heats of formation and the molar solubilities of the alkaline
earth hydroxides increase with increasing atomic number of the alkaline earth
element, as shown by the following data.

Be(OH) 2 Mg(OH) a Ca(OH) 2 Sr(OH) 2 Ba(OH),

Heat of Formation kcal/ mole 5.4 15.1 ~ 17;7 22.3

Solubility at 20° C, mole/ liter 2 X 10- 6 2 X 10- 4 0.022 0.066 0.23

The reaction of lime, CaO, and water is known as the slaking of lime and
the product calcium hydroxide, Ca(OH) 2 , is called slaked lime. The reaction
starts slowly but the heat liberated raises the temperature of the reaction
mixture to a point where it proceeds rapidly. Indeed the storing of unslaked
lime in a combustible container is a fire hazard. The solubility of Ca(OH) 2
in water is small. A clear solution is known as limewater while a suspension
of Ca(OH) 2 in water is called milk of lime. Lime that, is exposed to air slowly
absorbs H 2 0 and C0 2 ; such material is air-slaked lime .

Both CaO and Ca(OH) 2 react with atmospheric C0 2 to form CaC0 3 .

(3) Ca(OH) 2 + CO, CaCOa + H 2 0

It is this reaction which is the source of the usefulness of lime as a structural
material. Through it limestone, CaC0 3 , from which the CaO was made
originally, is regenerated albeit in a different physical shape. Mortar is made
by mixing slaked lime and sand in about a 1 : 3 ratio with water to form a
pasty mass. The hardening or setting of mortar consists, at first, of the
evaporation of water followed by the slow conversion over a period of years
of Ca(OR) 2 to CaC0 3 through reaction with atmospheric C0 2 .

The Elements of Group IIA: The Alkaline Earth Metals

521

Beryllium hydroxide, Be( OH ) 2 , is amphoteric, a behavior in accordance
with the small size of the Be 24 " ion. Hie other alkaline earth hydroxides are
basic. Unlike lime, MgO reacts slowly with water and dead burnt MgO
not at all. Since Mg(OH) 2 is only slightly soluble in water it is readily
precipitated by the reaction of solutions of Mg 2+ ion and a strong base,
OH” ion. Mg (OH) 2 is soluble in excess NH 4 + ion; it is not precipitated by
the addition of NH 3 to a solution of Mg 2 + ion when NH 4 + salts are present
in high concentration. The slight solubility of Mg(OH) 2 yields a concen-
tration of OH" ion in its saturated solution so small that it can be taken
internally without danger; suspensions of the material in water, sold under
the name of milk of magnesia , are used for reducing the acidity of the
stomach and as a mild purgative. Though Ca(OH) 2 is only slightly soluble
it is a strong base and is the cheapest of all “alkalies.”

Sulfates : Calcium sulfate occurs in two forms: anhydrite, CaS0 4 , and
gypsum (selenite; alabaster ), CaS0 4 * 2 H s O. Anhydrite has limited use, mainly
in the manufacture of ammonium sulfate, but over ten million tons of gypsum
are mined annually primarily for the production of plaster. In its early history,
dating back to Chinese legend, gypsum was used principally in die field of
art. Structurally, it has withstood the test of time in the pryramid of the
Pharaoh Cheops and in the masterpieces of Leonardo da Vinci.

When gypsum is heated to a temperature not above 125°C, it is dehydrated
partially by losing three-fourths of its water of hydration.

(4) 2[CaS0 4 • 2 H 2 0] -^(CaSO^H.,0 + 3 H 2 0

The product, a hemihydrate, is plaster of Paris or calcined gypsum. This
material has properties quite different from those of dead burnt gypsum
formed by prolonged calcining at temperatures above 200°C. When plaster
of Paris is mixed with water it forms a plastic mass which, through the
reversal of Equation 4, soon sets to a solid mass of interlaced crystals of
gypsum. During the setting process there is an expansion in volume so that,
if the reaction is carried out in a mold, all crevices are filled to produce
a sharp casting. Plaster of Paris is used for making plaster walls, casts of
statuary, surgical casts, pottery molds, etc. Casts can be rendered smooth
and nonporous by dipping them in paraffin. When mixed with wood pulp
and allowed to set in the form of sheets, plaster of Paris forms a material
much used in building construction as wall boards and partitions.

Magnesium sulfate crystallizes and is sold as the hydrate, MgS0 4 ■ 7 H 2 0,
known as Epsom salts . It is found abundantly in mineral springs (Epsom,
England) and in salt deposits, and is used medicinally as a purgative and for
the weighting of cotton goods. Barium sulfate is an important white pigment
in paint. It is a component of lithopone, a mixture of BaS0 4 and ZnS.
This pigment has excellent covering power, is not poisonous and does not
turn black as do lead paints upon exposure to H 2 S.

Carbonates : The normal* carbonates are among the most insoluble com-
pounds of the alkaline earth elements. Calcium carbonate, CaC0 3 , is readily
precipitated by carbonate, ion, CO a 2 ~. When C0 2 is passed into limewater,
Ca(OH) 2 , CaC0 3 is first precipitated (Equation 3) but an excess of the

522

The Elements of Group HA: The Alkaline Earth Metals

gas causes the precipitate to redissolve due to the formation of the soluble
calcium hydrogen carbonate, Ca(HCO a ) 2 .

(5) CaCOs -j- CO -2 Ht" hhO Ca~ 4 2 HCOv*

This reaction is reversible since heat will drive off C0 2 and cause the pre-
cipitation of CaCO, { . In substrata, limestone regions, the formation of natural
caves is due to the solvent action of water containing C0 2 on the limestone.
When water containing dissolved Ca(HCO :5 ) 2 reaches a cave, it may lose
its C0 2 due to a decreased pressure and deposit CaCO* in columns or in
icicle-like formations. The hanging columns are known as stalactites and the
formations of CaCO ;i on the floor of the cavern that separate from the drip-
pings are stalagmites.

7. Hardness of Water. Natural ground waters, ultimately the major
source of supply of industrial and home waters, are usually not pure. In
their course over and through the ground they are likely to dissolve various
mineral salts. Two of these are especially common in natural waters. One
is the slightly soluble CaS0 4 and the other is CaCO* which, though
insoluble in pure water, is soluble in water containing C0 2 through
the formation of Ca(HCO, ; ) a . Often FeCO a and MgCO a are similar-
ly dissolved. Water containing dissolved Ca 24 , Fe 24 , or Mg 24 salts is said
to be hard water, but by far the worst offender is the Ca 24 ion.

These dissolved salts in no way make the water unfit for drinking. Their
ill effects appear in two other functions of water, namely, in washing with
soap and in boiler water. With a typical soap such as sodium stearate,
C 17 H a5 COONa, Ca 24 ions, for example, react to precipitate the soap as an
insoluble calcium salt, calcium stearate, Ca(C 17 H a5 COO) 2 , that has no value
as a cleansing agent because of its insolubility.

(6) 2 C 17 H a5 COONa + Ca 24 -* Ca{C 17 H a ,COO) 2 (s) + 2 Na 4

If hard water is used for washing, the first soap entering solution is used
to precipitate Ca 24 ions, and only after this reaction has removed all the Ca 24
ions present can the soap lather and act as a cleanser. When water contain-
ing Ca(HC0 3 ) 2 is used in boilers, C0 2 is driven off by heating and CaCO*
precipitates in the boiler tubes as a stone-like layer called boiler scale . If
water containing CaS0 4 is evaporated in boilers, CaSO t also crystallizes as a
scale in the boiler tubes. Scale results in a loss of heat conductivity through
the metal of the boiler tube with the result that more fuel must be used
to obtain the same quantity of steam. Further, if the stone-like scale suddenly
cracks or develops a fissure due to overheating and water comes in contact
with hot iron, the sudden increase in pressure within the boiler or the re-
action of steam and red hot iron to produce H 2 can cause an explosion.

8. Softening of Hard Water. Because of these harmful effects, the Ca 24 ,
Fe 24 , or Mg 24 ions should be eliminated from water prior to its use in
washing or as a boiler feed. The removal of these ions from hard water is
known as the softening of hard water. Hardness in water may be classified
into two types: A) temporary , or carbonate hardness, and B) permanent , or
noncarbonate hardness. Temporary hard water contains Ca(HC0 3 ) 2 and is

The Elements of Group HA: The Alkaline Earth Metals

523

so called because it can be softened by boiling whereas permanent hard
water, such as one containing CaS0 4 , cannot he softened by boiling. In terms
of fuel, however, boiling is too expensive a method for the removal of tempor-
ary hardness.

The simplest way to soften either temporary or permanent hard water
is to remove the dissolved Ca 2+ ions as a precipitate, generally as insoluble
CaC0 3 . Since Ca(HCO,) 2 is an acid salt the addition of almost any base will
convert it to the insoluble normal salt, CaC0 3 , whose precipitation will
thereby remove the hardness from the water. In the household, NH 3 can be
added to water for this purpose. On a large scale, the cheapest base, Ca(OH) 2
is used.

(7) Ca 2+ + 2 HCOr + 2 NH, CaCO„(*) + 2 NH t + + CO/ 2 -

(8) Ca 2+ + 2 HCXV + Ca 2 + + 2 OH" -» 2 CaCO.(s) + 2 H.O

An excess of Ca(OH) 2 , beyond that required to neutralize the Ca(HC0 3 ) 2 >
is to be avoided; otherwise the water would still contain Ca 2 + ions and be
hard. Permanent hardness, usually caused by CaS0 4 , c&n be removed by the
addition of Na 2 C0 3 .

(9) Ca 2+ + S0 4 2 ’ + 2 Na+ + CO, 2 ’ CaCO { (s) + 2 Na+ + S0 4 2 ’

The use of Ca(OH) 2 , followed by Na 2 CO» for water softening is known as
the soda-lime treatment.

Besides its precipitation as CaCO,, other chemical reagents have been
used to remove Ca 2+ from solution. Borax, Na 2 B 4 0 7 • 10 H 2 0, and trisodium
phosphate, Na 3 P0 4 , are both basic by hydrolysis and also precipitate insoluble
calcium salts. With sodium hexametaphosphate, Na 8 P«O l8 , and with sodium
pyrophosphate, Na,P 2 O r , Ca 2 + forms complex ions, thereby removing the
simple Ca 2+ ion from solution.

9. Ion Exchange. Another technique for the softening of hard water is
ion exchange. Today the most commonly used materials in this process are
the ion exchange resins. An ion exchange resin is an organic polymer— a three
dimensional, insoluble hydrocarbon network throughout which are dispersed
ionic sites at which ions can be bound The most successful skeletal network
so, far synthesized is a copolymer of styrene, CeH->— CH=CH 2 , and divinyl-
benzene, CH=CHi>— C 6 H 4 — CH=CH 2 ; ionic groups can be introduced at fixed
positions in the network by a number of synthetic chemical procedures, e.g.,
sulfonation or reaction of the polymer with concentrated H 2 S0 4 . The skeletal
network is porous and the interstices fill with water (or solution in which
the resin is immersed) through which ions can move freely. However, ions
can be bound electrostatically at an ionic site and they can be replaced at
these sites by other ions of like charge.

There are two types of ion exchange resins. In one the insoluble matrix
contains negatively charged groups which bind positive ions and hence is
known as a cationic ion exchange resin because only cations or positive ions
are exchangeable at these ionic sites (Figure 39.1). Anions would not be
bound by a cationic resin but could move through the water permeating the

524

The Elements of Group II A: The Alkaline Earth Metals

Figure 39.1. A Cationic Ion Exchange Resin.
The insoluble polymeric matrix, shown as strands,
has negative groups permanently attached to it.
At these groups can be bound cations which can
be displaced by other cations. Within the matrix
is water or a solution through which the ions
can migrate. An anionic ion exchange resin has
fixed positive groups which can bind negative
ions.

matrix. The other type of resin is the anionic ion exchange resin in which
the matrix has positively charged sites at which negative ions or anions can
be bound and replaced or exchanged by other anions.

A cationic ion exchange resin may be formulated as R~ M + , where R~
represents an exchange site within the resin and M + a cation bound in the
resin. Electric neutrality throughout the resin must be maintained so that
a bivalent cation, M 2+ , would be bound to two exchange sites. Similarly, an
anionic ion exchange resin can be represented by R + X~, where X“ is a nega-
tive ion bound to a positively charged exchange site, R + , in the resin. A
cationic ion exchange resin in the sodium form, that is, a resin in which
the exchange sites are occupied by Na+ ions, can be written as R"Na + . If
a hard water containing Ca 2+ ions is passed over the resin, the Na + ions
will be replaced by the Ca 2+ ions.

(10) 2 R~Na+ +Ca 2 + (R-) 2 Ca 2+ + 2 Na+

Thus an exchange of Na + ions for Ca 2+ ions has been accomplished and the
Ca 2+ ions are now bound in the insoluble resin matrix. The Na + ions are
released into the effluent water but they are chemically harmless and have
no ill effects. By this exchange process the water has been softened.

An ion exchange resin has a limited capacity for exchange of ions. In
time all the exchange sites of the resin will be occupied by Ca 2+ ions. The
resin is then no longer capable of absorbing additional Ca 2+ ions and is
said to be exhausted. It can be regenerated, however, by passing over it a
concentrated solution of Na + ions, e.g., NaCjl. Through the reversal of Equa-
tion 10, the Ca 2+ ions are now displaced from the resin into the effluent
liquid and the resin is regenerated into the sodium form, R“Na + . Indeed
the reaction represented by Equation TO is an equilibrium and can be driven
in either direction depending upon the relative concentrations of the Na +
and Ca 2+ ions.

Ion exchange resins can be used not only to soften but also to demineralize
water. Let us consider a water containing dissolved NaCl. If first this water
is passed over a cationic ion exchange resin in the hydrogen form, prepared

The Elements of Group II A: The Alkaline Earth Metals

525

by treating the resin with a strong acid, the Na + ions will replace the H+
ions and be bound in the resin.

(11) R-H+ + Na+Cl- R-Na+ + H+Cl"

The effluent water thus contains the acid, HC1, instead of the salt, NaCl.
The acid solution is then passed through an anionic ion exchange resin in
the hydroxide form, R + OH% prepared by treating an anionic resin with a
strong base such as NaOH, and the Cl~ ions are bound therein while the
OH" ions of the resin are released.

(12) R+OH" + H+Cl- -» R+Cl- +H 2 0

The H + and OH" ions displaced from the resins combine to form H a O while
the Na + and Cl" ions are locked up in the cationic and anionic resins, re-
spectively. Instead of two separate resin treatments, a mixed bed of a
cationic and anionic resin is used commercially. It should be noted that this
procedure only demineralizes the water in that it removes solely the ionic
components. It does not completely purify a water as does distillation since
any nonionic species still remain in the demineralized water. Ion exchange
can be used for the removal of salts from sea water but, for such a high
concentration of electrolyte which exists in sea water, it is too expensive
and economically not feasible.

Whereas the prime use of ion exchange resins was originally in water
purification, many varied uses based upon the foregoing principles have been
developed. Among these are the demineralization of glycerol and other
organic chemicals, decalcification of blood, concentration of metal ions from
dilute waste waters, separation of the rare earth metals, color removal
from sugar, and catalysis by ion exchange resins in specific forms.

Certain rocks, such as the zeolites, which are hydrated sodium aluminum
silicates, e.g., analcite, Na 2 Al 2 Si 40 12 * 2 H 2 0, and synthetic aluminosilicates
also exhibit the property of ion exchange in that their Na + ions are replaceable
by other cations. Their capacity for exchange, however, in terms of gram-
equivalents of ions absorbed per gram of material is less than that of the
organic ion exchange resins. The phenomenon of ion exchange is not limited
to the organic resins and the aluminosilicates. Any insoluble ionic substance
in contact with ions in solution can undergo ion exchange. Plants receive
much of their nutrition from the soil through this process.

Ion exchange resins are available commercially in the form of spherical
beads which are generally employed packed into a column. Also available are
ion exchange materials in the form of thin sheets or membranes, cationic or
anionic. In contrast to the beads, these are not used for their capacity to
absorb ions but rather as barriers selective to the migration of ions since a
cationic ion exchange membrane will permit only cations to pass through
it while an anionic membrane will allow only the passage of anions. Set up as
partitions in a three compartment electrolytic cell they can be used to de-
mineralize water (Figure 39.2). A demonstration plant using this process
to prepare fresh water has been built by the Department of the Interior.

10. Cement and Concrete. Portland cement is produced by heating a
pulverized mixture containing approximately 75% limestone and 25% clay to

526

The Elements of Group IIA: The Alkaline Earth Metals

Anode

Cathode

The water to be demineralized, and containing an electrolyte such as Na 2 S0 4 ,
is placed ii\ the central compartment of a three compartment electrolytic cellf The
Na + ions migrate through the cationic ion exchange membrane to the cathode.
The electrode reaction at the cathode produces NaOH but the OH" ions are
trapped in this compartment because they cannot diffuse back through the cationic
membrane which is permeable only to positive ions. Similarly, SO/- ions leave
the central compartment and migrate towards the anode. Acid is produced by the
anode reaction but the H + ions are constrained from leaving this compartment
because they cannot pass through the anionic membrane. Thus the central com-
partment is demineralized and simultaneously high concentrations of acid and base
are produced in the electrode compartments. Multi-compartment cells can be
devised wherein demineralization and concentration occur in alternate compartments.

Figure 39.2 . An Electrolytic Cell for Demineralizing Water.

about 1,500° C. The essential ingredients are CaO, Si0 2) A1 2 0 3 , and Fe 2 0 3 .
Some natural rocks contain these elements in the suitable proportions. The
powdered mixture is “burned” in a slightly inclined rotary kiln through which
it is moved slowly by the rotation of the inclined cylinder. Kilns may be up
to 500 feet in length and 12 feet in diameter and constitute some of the largest
pieces of moving machinery in any industry.

In the lower third of the kiln, complex reactions occur whereby calcium
silicates and calcium aluminate are formed; the most important of these
are Ca-jSiCh, Ca 3 Si0 57 and Ca 3 (A10 3 ) 2 . At the high temperature of the kiln
the material is partially fused and sintered together in the form of hard
glassy balls called “clinker.” After cooling the clinker is mixed with 2-3%
gypsum, which acts as a retardant towards setting, and ground to a fine
powder. The resulting product is Portland cement. The name was given to
this class of cement, of which there are several kinds, by its inventor in
1824 because the concrete made from it resembled the building limestone

The Elements of Group 1IA: The Alkaline Earth Metals

527

from the Isle of Portland (England). In the United States, Portland cement
was first made in 1875 and over 300 million barrels are now produced
annually.

The setting of cement involves a number of complex hydration and hydroly-
sis reactions not all of which are known. The end products are calcium alum-
inosilicates which form a strong rigid network of crystals throughout the
mass. That the setting is not a simple reaction, but consists of several
reactions, is shown by the fact that the initial setting takes place in about
24 hours and is followed by a period of slow hardening which requires
about a month. Most Portland cements attain their xiltimate strength in one
year and half their ultimate strength in about seven days.

Portland cement alone has little utility; its use is in binding mineral
aggregate into concrete. There is no substitute for cement in this application.
Concrete is made by mixing cement, sand (fine aggregate) and crushed
stone or gravel (coarse aggregate); a common proportion is 1:2:4. Water is
added till the mixture has the proper consistency and thereafter the mixture
sets to a solid of high structural strength suitable for roads, walls, foundations,
and more recently, as shielding for nuclear reactors. Since it does not require
C0 2 from the air during the hardening process as does Ca(OH) 2 , concrete
will set under the surface of water and is so used in building dams, piers,
and docks.

QUESTIONS

1. Select three properties of the alkaline earth elements and describe how they
vary with the atomic number of these elements. Are there any anomalies?

2. Account for the hardness of beryllium and the softness of barium.

3. Which is the larger ion, Ca 2+ or K + (b) which is a stronger base, Be(OH) 2
or Ba(OH) 2 ? Explain briefly.

4. Would you expect the alkaline earth metal ions (a) to be colored and (b) to
form complex ions? Explain.

5. By what methods are the alkaline earth metals prepared? Compare these
methods with those for preparing the alkali metals.

6. Define, giving formulas where possible: (a) lime (b) slaked lime (c) milk
of lime (d) limewater (e) mortar (f) plaster of Paris (g) cement (h) con-
crete (i) ion exchange resin.

7. Starting in each case with CaCO„ write equations for the preparation of

(a) CaC, (b) Ca(HC0 3 ) 2 (c) CaS0 4 (d) CaOCl 2 (e) CaH 2 .

8. Write equations for the (a) formation of plaster of Paris from gypsum

(b) setting of plaster of Paris.

9. How is cement manufactured? What is concrete? How does the setting of
cement differ from the setting of mortar? Why can concrete set under -water?

10. Write equations to show why limewater becomes milky when CO a is first
bubbled through it but then clears upon further addition of GO*.

11. Define “hard water.” Distinguish between temporary and permanent hard
water. What are the ill effects of hard water?

12. List techniques for softening hard waters and write equations where applicable.

528

The Elements of Group 1IA: The Alkaline Earth Metah

13. A hard water contains 2,000 parts per million of CaS0 4 . What weight of

Na.>C0 3 * 10 H 2 0 is required to remove 90% of the hardness from one million
gallons of the water? Am: 16 tons

14. Discuss the ion exchange method for demineralization of water. Does this
method produce pure water? Write an equilibrium constant expression for
the exchange of Na + and Ca 2 + ions on an ion exchange resin.

15. Devise a five compartment ' electrolytic cell using ion exchange membranes
which will both demineralize and concentrate in alternate compartments.

16. At 25 °C, C’aC 2 0 4 is soluble to the extent of 0.0033 g in 600 ml of solution,

Calculate its solubility product. Am: 1.8 X 10-°

17. What weight of CaH 2 reacting with water is needed to fill a 500 liter tank

with hydrogen at 20 °C and 780 mm pressure? Am: 450 g

18. Explain why the equilibrium constant expression for the decomposition of
Ca0O 3 is: K = P C02 .

19. At 800° C, the value of the equilibrium constant for the decomposition of
CaC0 3 is 0.220 atm (a) Calculate the pressure of C0 2 produced 800°C
in a five liter container (b) if 10.0 g of CaC0 3 were placed in a five liter
container at 800 °C, how much would remain undecomposed?

Transition Elements— I
General Properties

The transition elements are those elements which occupy the “center” of
the Periodic Table between Groups III A and II B. There are four series of
transition elements. The first transition series is the group of ten elements
in Period 4 from scandium (Z = 21) to zinc (Z = 30). The second is in
Period 5 and also consists of ten elements from yttrium (Z — 39) to cadmium
(Z = 48). The third series is in Period 6, starts with lanthanum (Z = 57),
includes an inner transition series of fourteen elements from cerium ( Z = 58 )
to lutecium (Z = 71) called the lanthanide series, and ends with mercury
(Z = 80). The fourth transition series, in Period 7, begins with actinium
(Z = 89) and also includes an inner transition series of fourteen elements
called the actinide series, from thorium (Z — 90) to the last element cur-
rently known, lawrencium (Z = 103). Presumably element No. 104 would
then continue the normal transition series started with actinium and would
appear in the Periodic Table under hafnium, but presently actinium is the
only member of this fourth transition series. Except for the last series which
is incomplete, because the number of elements known runs out at No. 103,
each transition series starts with an element in Group IIIA and ends with
an element in Group IIB, as shown in Table 40-A.

The genesis of a transition series is its building up of the d or f orbitals
of an inner principal energy level while the outermost energy level i*emains
essentially unchanged. In a transition series, a d subshell is being completed
while in an inner transition series an f subshell is being filled. Since the d
orbitals can accommodate ten electrons, each of the three complete transition
series contains ten elements. In the first, the 3d orbitals are being filled, in
the second the 4d orbitals, and in the third the 5d orbitals." Completely oc-
cupied f orbitals contain fourteen electrons and so each inner transition series
has fourteen elements, the first due to the filling of 4 f orbitals and the
second due to the filling of 5 f orbitals.

A brief consideration of the “formation of the elements in Period 4, in
which the first transition series appears, will also afford a review of atomic

530

Transition Elements * — I: General Properties

Table 40-A

The Transition Series of Elements

VIIA

VIII

1st Transition Series

2nd Transition Series

3rd Transition Series

Au i
i

1st Inner Transition Series; Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Lanthanide Series

4th Transition Series

2nd Inner Transition

Series; Actinide Series Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lw

structure. Argon (Z = 18) is a noble gas with the electron configuration,
2-8-8, or more critically, Is 2 2s 2 2 p B 3s 2 3 In the next element, potassium
(Z = 19), the 19th electron starts the fourth principal quantum shell; even
though potassium has available 3d orbitals, the 4s orbital has a lower energy
than a 3d orbital and hence the configuration of potassium is 2-8-8- 1, or
Is 2 2 s 2 2 3s 2 3 p 6 4s 1 , rather than 2-8-9, which would require the occupying
of a 3d orbital. With the next element, calcium (Z ~ 20), the 4s 2 orbital
is filled.

With scandium (Z = 21), the 21st electron could possibly be a 4 p electron.
However, the 3d orbitals have an energy lower than do the 4p orbitals so

that they are occupied first. With scandium we have the beginning of the

first transition series as shown in Table 40-B. Since the five d orbitals can
accommodate ten electrons, the filling of the 3d orbitals from scandium to
zinc constitutes the first transition series. The filling of the orbitals is not al-
ways regular and in some cases the electron configurations are not definitively
known. For example, in the first transition series where the outermost
shell is 4s 2 , Sc has one 3d electron, Ti has two, V has three, but Cr, which
might be expected to have four 3d electrons, has five instead through the
promotion of one 4$ electron to the 3d level. A configuration whereby the
3d orbitals are half filled with one electron in each appears to be particularly
stable.

The next element, Mn, again has the configuration 4s 2 . From Fe on,

electrons are added to the 3d orbitals but, when Cu is reached, a process

similar to that which occurred with Cr again is evident. Instead of having
3d 9 4s 2 , Cu completes its d orbitals at the expense of a 4$ electron and has the
configuration, 3d 10 4s 1 . With the next element, Zn, both the 3d and 4s orbitals
are complete.

Transition Elements—!: General Properties

531

Table 40-B

The 1st Transition Series

Element

Atomic

Number

2 p

Table 40-C

The 1st Inner Transition Series

Ele-

ment

Atomic

Number

Sm •

10.

In the second and third transition series a similar progression is exhibited,
wherein the 4d and 5 d orbitals, respectively, are being occupied. By some
chemists the terminal elements of the three transition series, namely, Zn,
Cd, and Hg, are not classified as transition elements because their d and s
orbitals are complete. A transition element is thereby defined as one having
incomplete, but partially occupied, d orbitals; however, Gu, Ag, and Au,
elements with complete d orbitals but with only one electron in the outer-
most s orbital, are also included. Even though their properties may be at
variance with most of the transition elements, we shall also include the

532

Transition Elements ■ — I: General Properties

elements Zn, Cd, and Hg within the class of transition elements because thev
conform best to a regular viewpoint of electron configuration.

In the formation of an inner transition series, electrons enter f orbitals, as
indicated in Table 40-C. Since there can be seven such orbitals the first
inner transition series, the lanthanide series, from Ce to Lu in which the if
orbitals are being filled, numbers fourteen elements. Being of lower energy
these 4/ orbitals are occupied in preference to the 5 d orbitals. So, too, the
second inner transition series of fourteen elements, the actinide series, from
Th to Lw, is due to the filling of the 5/ orbitals.

It should be noted that electrons which enter d orbitals in the formation
of a transition series go into a next-to-outermost principal energy level,
while electrons which go into f orbitals in an inner transition series enter a
principal energy level third from the outermost. Transition elements also fit
into vertical groups of the Periodic Table, all of which are the B subgroups.
It is now evident that the difference between the A and B subgroups is that
the elements in the B subgroups have partially filled d orbitals in the case
of transition series elements, partially filled f orbitals in the case of inner
transition series elements, except for Zn, Cd, and Hg, wherein the d orbital
filling process is completed and for Cu, Ag, and Au, which complete their
cl orbitals at the expense of the outermost 6' orbital. Nontransition elements,
or Subgroup A elements, have in their next-to-outermost shells 8 or 18 elec-
trons; electrons added to such elements enter outermost shells. For the transi-
tion elements the first element in the subgroup is characteristic and gen-
erally the most important and it is to these that most of our attention will
be directed.

1. General Properties of Transition Elements. Because their atoms have
few electrons in their outermost shells, the transition elements are metals
and generally exhibit typically metallic characteristics. No nonmetal is a
transition element. More than half the known elements and most of the
metals are transition elements. Most are rare or occur widely dispersed in
trace amounts but some, such as Ti and Zr, are more abundant than the
moi’e familiar metals, Cu, Sn, and Pb, while a few, e.g., Cu and Fe, are of
major industrial importance. Unlike the elements of Groups IA and IIA, it
is the transition elements themselves, as metals, rather than their compounds,
which are important and more well-known. Many of the transition metals
were laboratory curiosities until the twentieth century. With the advance of
present day technology in electronics, nuclear energy, missiles, and space-
craft, most of these metals have acquired specific and essential uses.

Some of the transition metals have not been prepared in the completely
pure state, and in such cases their physical and chemical properties are
not definitively known. In general, the transition elements have high densities,
high melting points, and high boiling points. The values of these increase
from the first to the second to the third transition series so that Os, Ir, and
Pt aie the densest of all elements while W and Re have boiling points ex-
ceeded only by that of carbon. These properties arise from the relatively
small atomic sizes of the transition elements in conjunction with the formation
of strong interatomic covalent bonds.

Transition Elements — I: General Properties

533

Because electrons enter an inner shell and the number of electrons in
the outermost shells of the transition elements remains essentially constant,
there is considerable horizontal similarity in physical and chemical properties
among the elements of a given series. As an example, the data given in
Table 40-D indicate the slight change in atomic volume for the elements
in the first transition series and the slight variation in ionization potential;
similar trends are found in the other series.

Table 40-D

Atomic Radii and Ionization Potentials of the Fibst Transition Series

Element

Atomic Radius, A

1.439

1.324

1.168

1.165

1.157

1.149

1.173

1.249

Ionization

Potential, 1st, eV

6.358

6.818

6.743

6.76

7.432

7.896

7.862

7.633

7.724

9.931

In the case of an inner transition series the resemblance in properties is so
marked that die fourteen elements are placed within one cubicle of the
Periodic Table. Indeed, the chemical properties of the lanthanide series of
elements are so alike that they can be separated only with extreme difficulty.

The transition elements are less reactive than the Group IA and IIA ele-
ments, but their activities vary greatly. The elements of Group IIIA, Sc, Y,
and La, have high positive oxidation potentials and are extremely reactive
while those of Group IB, Cu, Ag, and Au and the platinum metals have
negative oxidation potentials and are quite unreactive.

In addition to using electrons in the outermost shell for compound
formation, a variable number of inner d electrons can also be used for this
purpose so that transition elements are generally characterized by multiple
oxidation states. Thus titanium has oxidation states of 4-2, +3, and +4.
In the +2 state only the 4s 2 electrons are used while in the 4-3 state one
3 d electron is also used and in the +4 state both 3d electrons are used.
Generally the maximum oxidation number equals the total number of elec-
trons available for bond formation and this, in turn, is equal to the number
of the subgroup in which the element is. The most common oxidation states
are +2 and -f 3. The +1 oxidation state, corresponding to the loss of a single
s electron, is rare except for the Group IB elements, Cu, Ag, and Au. Only
Sc, Y, and La in Group IIIA; Zr and Hf in Group IVA; Ag in Group IB; and
Zn and Cd in Group IIB show single oxidation states of +3, +4, +L and + 2 ,
respectively. In the lower oxidation states compounds are essentially ionic.
In the higher states, where a greater number of electrons are used in bond-
ing, compounds are covalent and many transition elements form anions, for
example, chromate, CrO 2 ~, dichromate, Cr 2 0--”, manganate, Mn0 4 2 \ per-
manganate, MnOr, vanadate, VOs*", titanate, Ti0 3 2 ', ferrocyanide, Fe(CN) 6 4 ~,
and ferricyanide, Fe ( CN ) 6 3- .

534

Transition Elements — l: General Properties

Not merely their classification but also most of the properties of the
transition elements and their compounds stem from the existence of unpaired
d electrons. Thus it is characteristic of these elements that they form
numerous complex ions. Incomplete or unoccupied inner d orbitals of a
transition element can be used to form complex ions with molecules or ions
which can donate electrons and complete the vacant orbitals through the
formation of coordinate covalent bonds. It has already been noted that the
transition element ions, Cr 34 , Co 34 , Pt 44 , and Pt 24 , form thousands of
complex ions.

Whereas the transition elements, except for Cu and Au, have the usual
gray metallic luster their compounds are generally colored. Only the ions
of the elements of Groups IIIA and IIB, and Cu 4 , Ag + , and Ti 44 , form color-
less compounds. Almost all inorganic colored substances are compounds of
transition elements; ions of elements at both extremes and at the top of
the Periodic Table are colorless. The color of transition compounds is due
to the excitation of an impaired d electron between closely spaced energy
levels. If the difference in energy between these levels corresponds to a
frequency in the visible spectrum, certain wavelengths will be selectively
absorbed from white light which passes through a compound of a transition
element so that its transmitted light appears colored. Thus the Fe 34 ion is
yellow because its absorbs blue radiation. The Sc 34 and Ti 44 ions are color-
less because they have no 3d electrons while the Cu 4, and Zn 24 ions are
colorless because they have completely paired 3d electrons. In different
oxidation states an element can exhibit different colors, e.g., the green Fe 2 h
'and the yellow Fe 34 . Closely related to color is the property of para-
magnetism. Again those elements and ions with unpaired electrons are
paramagnetic; iron is the outstanding example. Many of the transition ele-
ments act as effective catalysts for gaseous reactions, such as the Haber
process, but as yet there is no general theory relating atomic structure to
catalytic properties.

2. The Elements of Group IIIA. These elements never occur separately;
they are always found in the combined state associated with each other
and with other metallic elements. It is extremely difficult to isolate the
pure metals because of their marked similarity in properties. This is especially
true of the two inner transition series, particularly the lanthanide series. Isola-
tion of any of these elements requires complex and costly chemical processing
but each can be obtained ultimately by the electrolysis of a fused halide
or oxide. The properties of the elements are listed in Table 40-E.

Scandium , Yttrium , and Lanthanum : These are rare elements with no
commercial uses other than that for yttrium in microwave transmission where
a Y-Fe alloy shows low loss in transmitting microwave energy. The main
source of scandium is the mineral, thortveitite , (Sc,Y) 2 0* • 2 Si0 2 . The
principal yttrium and lanthanum minerals are gadolinite , a silicate of lanthan-
ide elements, and mormzite, a phosphate of these elements. Experimental
knowledge of these elements, particularly of scandium, is fragmentary. The
elements are reactive, having high oxidation potentials, over two volts,
sufficient to displace H 2 from cold water vigorously. Only one oxidation

536 Transition Elements — h General Properties

state, +3, is known. The ions are colorless and are not paramagnetic since
the loss of the 4s 2 electrons and the single 3 d electron leaves a configuration
with no unpaired electrons. Compounds of these elements can be prepared
by conventional methods; their chemistry resembles that of aluminum.

The Lanthanide Series: The lanthanide series of elements, frequently called
the "rare earths” because of their apparent scarcity at the time of their
discovery and because their oxides resembled the alkaline earth oxides, are
the elements of atomic number 57 to 71. Some of the rare earths are not
truly rare; the abundance of cerium, 0.001%, for example, is greater than that
of mercury and tin.

During the early development of the Periodic Table chemists could not
understand how these fourteen ""stepchildren” after lanthanum could be
fit into the table without upsetting its symmetry. With our present under-
standing of atomic structure we know that the lanthanide elements owe
their existence to a belated filling of the 4/ orbitals. No element of atomic
number lower than 57 has f electrons.

The mineral monazite contains about 50% rare earth compounds and is
the most common source of these elements. Because of the similarity in
their properties and because of their concomitant occurrence, separation of
the lanthanide elements is extremely difficult and they are used industrially
without separation, Misch metal is the general term applied to mixtures of
lanthanide metals. Separation of the lanthanide elements is possible, however,
though tedious, because there is a slight gradation in properties from
lanthanum to lutecium. The increase in nuclear charge gives rise to a slight
shrinkage in atomic volume known as the lanthanide contraction. The ac-
ompanying slight decrease in basicity and difference in solubility have en-
abled the separation of the lanthanide elements through repeated precipita-
tion and fractional crystallization. More recently the difference in strength
of adsorption upon a column of ion exchange resin followed by elution with
complexing agents has yielded excellent separations yielding 99.99% pure
metals. After loading a cation ion exchange resin with a mixture of lanthanide
ions, a reagent, such as citric acid, which forms complexes with these ions is
passed through the column. The heavier lanthanides are complexed more
strongly and move down the column more rapidly, resulting in a selective
elution. The metals are obtained ultimately by electrolysis, usually of the
chlorides, or by high temperature metallothermy. The elements are used as
alloying agents, in coloring glass, in lighter flints because they are pyrophoric,
and in the cores of arclight carbons to give a brilliance more closely ap-
proximating natural sunlight.

The lanthanide elements show the typical reactions of active metals. All
have an oxidation state of +3; Sm, Eu, and Yb are also bivalent, while Ce
is the only one of the group which has a +4 oxidation state. The compounds
of the lanthanides do not resemble tri valent elements so much as they do
the alkaline earths. Unlike Sc, Y, and La, most of the lanthanide ions are
colored and paramagnetic; Lu 3 +, Yb*+, and Gd 3+ salts are colorless because
these ions have complete or half complete 4 f orbitals. Cerium is the most
important element of the group. It can be separated from the other lanthanides
relatively easily because it alone forms a +4 oxidation state and can be

Transition Elements — I: General Properties

537

precipitated as Ce(QH) 4 . Cerium is also the easiest of these metals to reduce
to the elemental state by electrolysis.

The Actinide Series : This series of elements is analogous to the lanthanide
series in that the 5 f orbitals are being filled. Only the first three elements
after actinium, namely thorium, protoactinium, and uranium, are found in
nature; the others, the so-called transuranic elements of atomic number
greater than 92, have been synthesized in the laboratory by the techniques
of nuclear physics. Because all the actinide elements are radioactive and
because some of the transuranic elements have been prepared only in trace
amounts, some of their electron configurations and properties are not pre-
cisely known.

Unlike the lanthanides, the properties of the actinide elements are not
so similar in properties. Like Sc, Y, and La, actinium forms only a -f 3 oxida-
tion state. Only the +4 state of thorium is stable but most of the other
actinides form multiple oxidation states. These differences permit modes
of separation which make the actinide elements easier to separate than the
lanthanide elements. The numerous oxidation 'states arise from the relative
ease with which 5 f electrons are used in bonding, unlike the 4/ electrons
of the lanthanide elements. Presumably the 4 f electrons, being closer to
the nucleus, are held more strongly and so are not so readily available for
bond formation. The chemical properties of Ac, Th, and Pa resemble those
of the elements in Groups IIIA, IVA, and V A, respectively, but the properties
of uranium differ markedly from those of the Group VIA elements. For ex-
ample, tungsten is quite inert but uranium is very reactive.

The radioactive properties and uses of uranium, neptunium, and plu-
tonium in the field of nuclear energy are discussed in Chapter 49, but the
importance of uranium warrants a brief survey of its chemistry. In 1789 the
German chemist, M. H. Klaproth, recognized that the mineral pitchblende
contained a new element which he named uranium in honor of the planet
Uranus, discovered by Sir William Herschel in 1781. Not until 1842 was
uranium metal itself isolated by E. M. Peligot, and not until 1896 was the
radioactive property of uranium discovered by Henri Becquerel. The develop-
ment of nuclear energy has exalted uranium from an element of purely
academic interest to one of international importance. The first 'radioactive
element to be discovered, uranium is the fountainhead of all the transuranic
elements.

The principal ores of uranium are pitchblende , U s 0 8 , and camotite,
KaO * 2 UO s * V 2 0 5 * 3 H 2 0. In the United States the largest deposits are
in the Colorado Plateau. Though its occurrence, 0.0004% in the earth's crust,
classes uranium as a rare element it is more abundant than mercury or
gold. About 7 million tons of uranium ore are processed annually to yield
about 16,400 tons of U 3 O s valued at approximately $300,000,000. The metal
is obtained by leaching the ore with a strong acid such as H 2 SO* to form a
solution of uranyl ion, U0 2 2+ . This is followed by solvent extraction or by
treatment with an ion exchange resin to remove the uranium compounds.
Solvent extraction is a common technique for concentrating a substance based
upon its greater solubility in one of two immiscible liquids. In this case it
consists of mixing the aqueous uranyl solution with an immiscible solvent

538

Transition Elements — 1: General Properties

such as kerosene. In the kerosene is dissolved an organic compound, an
alkyl phosphate or amine, which reacts with the uranyl ion to form a com-
plex compound soluble in the kerosene but insoluble in water. From this
the uranium is precipitated by NH 3 or NaOH as U0 2 * 2 H 2 0. By reaction
with HF, this is converted to UF 4 and the metal is then obtained by re-
duction with Mg.

The melting point of uranium metal is about 1130°C and its boiling point
about 3900° C; its density is high, 19.05 g/cm 3 . Elemental uranium has an
oxidation potential to U 3 ~ of +1.85 volts, greater than that of aluminum,
and exhibits the typical reactions of an active metal. It combines with
atmospheric oxygen to form a film of protective oxide. The many oxidation
states of uranium together with typical compounds are listed in Table 40-F.
The most important states are the +4 and +6; the latter arises from the
use of the 7 sr, the Qd 1 , and the 5/ 3 electrons in bonding. In 1M acid solution,
the oxidation potentials for the several states are:

1.85 v 0.61 v -0.62 v -0.05 v
( 1 ) u » U 3+ > U 4+ » U0 2 + » U0 2 2+

Table 40-F

Oxidation States of Uranium

State

Formula

Description

an unstable state; few compounds are known

uf 3

solid compounds are stable; U +3 ion displaces H 2 from H 2 0

uo 2

a brown oxide, soluble in strong acid

(U 2 0 5 )

unstable and disproportionates into U 4+ and U0 2 2+

uo 3

i orange oxide

uf 6

the only volatile compound of uranium at ordinary tempera-
tures; used for the separation of uranium isotopes by gaseous
diffusion

UO a 2 +

a yellow ion; uranyl salts are the common commercial uranium
compounds; addition of a base precipitates insoluble metal di-
uranates, e.g., 1C 2 U 2 0 T . Uranyl salts are fluorescent; sodium di-
uranate, Na 2 U 2 0 7 * 6 H O, is used to make a yellow fluorescent
glass.

u 3 o 8

may be considered as U0 3 * U 2 0 5

3. The Elements of Group TV A The elements of Group IV A are titanium,
zirconium, and hafnium. Their electron configurations are as shown below.

Element

Atomic

Number

2s 2p

3s 3 p 3d

4s 4p 4 d 4 f

5s 5 p 5 d 5 f

6s Bp 6d

2 6

2 6 2

2 6

2 6 10

2 6 2

2 6

2 6 10

2 6 10 14

2 6 2

Transition Elements — I: General Properties

539

Titanium (Greek, Titanes : the Titans) is the most important element of
the group. Although it is the ninth most abundant element and the fourth
most abundant structural metal, only two minerals are of commercial im-
portance, ilmenite , FeTiO ;j , and rutile , Ti0 2 . The metal was discovered by
William Gregor in 1790, but because of its difficult metallurgy was not used
industrially till over a century later, and then only as an additive to iron
and steel. Since World War II there has been a sharp increase in titanium
production. The metal owes its recent importance to a combination of light-
ness, structural strength, and corrosion resistance. The strength-to-weight ratio
of titanium alloys exceeds that of any other metal. Titanium alloys are used
in airplane frames and power plants; each jet airplane uses up to 2500 pounds
of the metal. Much titanium is also used to coat welding rods.

At temperatures below 600 °C, titanium metal is more inert than aluminum
and resists corrosion in Sea water where stainless steels are attacked rapidly.
But at high temperatures, titanium is very reactive and consequently it is
difficult to produce the pure metal. Metallothermic reduction yields alloys
with the reducing metal and, at the high temperatures required for the re-
duction, the liquid titanium combines with all electronegative elements, with
all gases except the noble gases, and with all known refractory materials.
Comiriercially pure titanium is prepared principally by chlorinating Ti0 2 to
TiCl 4 followed by reduction of the tetrachloride with Na or Mg is an inert
gas atmosphere. Small quantities of chemically pure metal are produced by
the decomposition of Til* vapor on a hot tungsten wire.

The +4 oxidation state of titanium is due to the use of both 4 s and both
3 d electrons as bonding electrpns; -f 2 and +3 states are also known. In acid
solution, the oxidation potentials of titanium and its ions are:

1.63 v 0.37 v -0.lv

(2) Xi » Ti 2+ ► Ti s+ » TiO a +

Titanium (IV) oxide, TiO a , is the most important compound. It is resistant
to acid, alkali, and saline solutions, and because of its high opacity and
chemical inertness, is used as a white pigment in paints, plastics, rubber,
and other products. About 99% of all ilmenite goes into manufacturing Ti0 2
for pigments. Unlike lead-base paints, titanium pigments are not poisonous
nor are they discolored by H 2 S. Synthetic crystalline Ti0 2 is a semiprecious
gem. It has a greater refractive index and brilliance than diamond but it is
softer so that it is more easily scratched, a mode of distinction between the
two. When fused with other metal oxides or carbonates, Ti0 2 forms titanates,
e.g., CaTiO :i . The halide, TiCl 4 , is a colorless liquid which is used to make
smoke screens since its hydrolysis with atmospheric moisture yields a fine
suspension of TiO a .

Zirconium (Arabic, zargtin; gold color) is found as the silicate , ZrSi0 4 ,
and as the oxide, baddeleyite, Zr0 2 . The element was discovered in 1789 by
M. H. Klaproth. Because their chemical properties are almost identical the
small percentage of hafnium that occurs in conjunction with zirconium went
undetected until its discovery in 1923 through X-ray analysis by D. Coster
ar»rf n Wpvpw Prior tn that date all “mire” zirconium compounds contained

540

Transition Elements — I: General Properties

the concomitant hafnium as an impurity. Both elements are prepared by
reduction of their tetrachlorides in a manner similar to that for making
titanium metal. The method of separating the zirconium from the hafnium
is a complex operation involving solvent extraction.

Because zirconium metal has strength, corrosion resistance, and a low
absorption for neutrons it finds use in nuclear reactors as structural mem-
hers, as fuel cladding, and as a fuel moderator. For such nuclear purposes
the zirconium must be free of hafnium. In their natural mineral form, zircon
and baddeleyite, and also synthetic zirconia , Zr0 2 , are used as refractories
because of their high melting points, over 2500 °C.

Hafnium (Latin, Hajnia : Copenhagen) always occurs in conjunction with
zirconium minerals to the extent of 1-2% and is recovered as a by-product
of hafnium-free zirconium. Currently the only significant use of hafnium
is as a neutron absorber for controlling nuclear reactors. Its thermal neutron
absorption cross-section is 105 barns, compared to 5.8 for titanium and 0.18
for zirconium. Hafnium is resistant to corrosion at high temperatures, is
stable under the intense radiation within a nuclear reactor, and does not re-
quire cladding. Zirconium and hafnium form compounds only in the -f-4
oxidation state but otherwise their chemistry is similar to that of titanium.

4. The Element's of Group VA. The elements of Group VjA are vanadium,
niobium (formerly called columbium), and tantalum. Their electron con-
figurations are given below.

Element

6s 6p Qd

2 6

2 6 3

2 6

2 6 10

2 6 4

2 6

2 6 10

2 6 10 14

2 6 3

Though vanadium is widely distributed and its abundance, 0.017%, is
about equal to that of Cu, Ni, and Pb combined, it is not a readily available
element. It occurs mainly in conjunction with uranium ores of the Colorado
Plateau; two of the principal minerals are vana (Unite, Pb 4 (V0 4 ) 3 * PbCl, and
carnotite , from which vanadium is extracted largely as a by-product of
uranium production. In 1801 A. M. del Rio isolated a metal from Mexican
lead ore which he named “erythronium” but which he decided later was
merely a basic lead chromite. This decision was erroneously confirmed in
1805 by Collet-Discotils but in 1830 Nils G. Sefstrom rediscovered del Rio’s
erythronium as a new element, vanadium ( Vanadis , a Norse goddess).

Like the metals of Group IVA, elemental vanadium is difficult to prepare
because of its high melting point and its reactivity at elevated temperatures
with other metals, gases, and even ceramics. Little pure vanadium, however,
is used industrially; ferrovanadium , an alloy of Fe and V, is the single most
important vanadium product. It is prepared by reducing the mixed oxides,
V-0 5 and Fe 2 0 3 , by aluminothermy or in an efectric furnace. Most of the
ferrovanadium is used to make steel alloys. Added to steel in concentrations

Transition Elements — I: General Properties

541

as little as 0.01-0.25%, vanadium imparts strength, toughness, and resistance
to impact and abrasion; 0.25% vanadium in steel almost doubles its tensile
strength. In his Model T automobile, Henry Ford was the first to tecognize
the merits of vanadium steel.

The oxidation states of vanadium, +2, +3, +4, and +5, arise from the
use of its 4s electrons and from one to three of the 3 d electrons as valence
electrons. The most important compound of vanadium is vanadium(V) oxide,
V 2 0 5 , a red solid. It is prepared by the decomposition of ammonium, meta-
vanadate, NH 4 VO s , and is the starting material for the preparation of other
vanadium compounds.

(3) 2 NH 4 V0 3 -» V 2 O r > + 2 NH 3 + H a O

In strong acid solutions, V 2 0 5 dissolves to form the yellow vanadyl ion,
VO 2 *, while in solutions of alkali hydroxides, vanadates, VOs" and V0 4 3 ~, are
produced. V 2 O g is used widely as a catalyst in oxidation reactions, notably
the oxidation of S0 2 to SO a in the contact process for making H 2 S0 4 and in
the synthesis of many organic compounds.

Niobium and tantalum usually occur together, the principal minerals
being columbite, FeNb 2 0 6 , and tantalite , FeTa 2 O fi ; they were named for
Niobe, the daughter of Tantalus, a Greek mythological character, and for
Tantalus.

Niobium was discovered in 1801 by Charles Hatchett, an English chemist, in a
mineral in the British Museum. The mineral had been sent to England over a
century earlier by John Winthrop, the first Governor of the Massachusetts Bay
Colony, Hatchett named the element columbium and the mineral columbite. The
following year, the Swedish chemist, A. G. Ekeberg, isolated a new element
and named it tantalum , in allusion to the tantalizing difficulty he encountered in
dissolving its mineral. In 1844, it was found that Ekeberg’s tantalum was truly a
mixture of two elements, tantalum and a second one which was named niobium.
Further investigation proved that columbium and niobium were one and the same.
Though the dual nomenclature has persisted, the name niobium is the official one.

As w T as the case with zirconium and hafnium, separation of niobium and
tantalum is difficult because of a marked similarity in their chemical proper-
ties. Fusion of the ores with NaOH followed by leaching the soluble com-
ponents with water leaves a Nb-Ta residue which can be separated by solvent
extraction methods. Ultimately, elemental niobium is obtained by reduction
of Nb 2 0 5 with carbon while elemental tantalum is prepared mainly by elec-
trolysis of fused K 2 TaF 7 . Of the two metals, tantalum has the greater indus-
trial importance. It is added to certain steels and high temperature alloys
as ferrotantalum-niobium (40% Fe; 40% Nb; 20% Ta). The high melting point
of tantalum, greater than that of any metal except tungsten and rhenium, and
its extreme corrosion resistance, form the basis of its use in heat exchangers
and other industrial equipment. Though tantalum dissolves in concentrated
alkali solutions it is inert to the action of all common acids except HF and
hot concentrated H 2 S0 4 . Because of its inertness, tantalum is used for
surgical sutures, repairs, and prosthetics. Tantalum has many electronic uses.
Being an excellent thermionic electron emitter it is used in cathodes but its
largest consumption is in electrolytic capacitors. These pass electric current

542

Transition Elements — I; General Properties

in one direction only because of a thin tenacious film formed on the surface
of the tantalum by anodic oxidation. A new type of miniature electron valve
the cryotron , consists of a tiny strip of tantalum around which a fine niobium
wire is wrapped. The combination becomes superconductive at about 3°K
(in a bath of liquid helium), and may provide new advances in computer
design and in radar devices.

QUESTIONS

1. Locate the various transition series and inner transition series of elements
on the Periodic Table, Which element begins each series?

2. What is a transition series from the viewpoint of electron configuration? Why
are transition elements metals? Which of the transition series includes an
inner transition series? To what is an inner transition series due?

3. Why do elements of a given transition series show a marked similarity in
chemical and physical properties?

4. Account for the following properties of transition element ions: (a) they are
colored (b) they form numerous complex ions (c) they show many oxidation
states.

5. Write the electron configurations of the following ions: Co 2+ ; Cu + ; La 3+ ;
C’e 4+ ; Pr 3+ ; Sm 2+ ; Nd 3+ ; Fe 2+ ; Fe 3 +. Which of these ions should be
colored?

6. Account for the greater number of oxidation states shown by elements of
the actinide series as compared with elements of the lanthanide series,

7. Discuss the use of ion exchange resins in separating the lanthanide elements.

8. Calculate the percent difference in atomic volume (cm 3 /GAW) between
(a) lanthanum and lutecium and (b) cesium and barium. The density of
lutecium is 9.74 g/ cm 3 .

9. Which is the more basic oxide, La 2 0 3 or Lu 2 0 3 ? Explain.

10. Why are Ti 2 + and Ti 3+ compounds readily oxidized to the +4 state?

Transition Elements— II
Iron, Cobalt, Nickel
and the Platinum Metals

In each transition series there is a triad of elements: the first, the iron
family composed of iron, cobalt, and nickel; the second, the light platinum
trihd composed of ruthenium, rhodium, and palladium; and the third, the
heavy platinum triad composed of osmium, iridium, and platinum. In con-
trast to the normal vertical group relationships of the Periodic System, the
elements within each of these horizontal triads bear a greater resemblance in
properties to each other than they do to elements above or below in the
vertical sequence. In his original Periodic Table, Mendelejeff recognized this
and reserved an eighth column for these triads which constituted a separate
Group VIII.

1. General Properties of the Iron Family. For the iron triad the electron
configurations below and the properties listed in Table 41- A offer evidence
of the considerable similarity in physical and chemical properties of these
elements.

Element

Atomic

Number

mam

The electron configurations of iron, cobalt, and nickel differ only in the
number of 3d electrons, the outermost energy level in each case being 4s 2 .
Inasmuch as the 4s 2 electrons are readily lost these elements all evidence
an oxidation state of -^2. Under suitable oxidizing conditions an additional
electron can be removed from the 3 d orbital to form the +3 state. In the
case of iron, the third electron is easily removed so that the Fe 3+ ion has
a single unpaired electron in each of its 3 d orbitals, a configuration which

544

Transition Elements — II: Iron, Cobalt , Nickel and the Platinum Metals

we have seen has a special stability. From cobalt a third electron is not so
readily removed while the NF+ state is found only in Ni 2 0 3 and as a part
of a few complex ions.

Tabic 41 - A

Properties of Iron, Cobalt, and Nickel

Properttj

Iron

Cobalt

Nickel

Symbol

Atomic Number

Atomic Weight

55.847

58.933

58,71

Isotopes (mass numbers

54 ( 5.90)

59(100)

58(67.76)

and percent)

56(91.52)

57 ( 2.25)

58 ( 0.33)

60(26.17)

61 ( 1.25)

62 ( 3.66)

64 ( 1.16)

Abundance in Earth’s

5.0

0.001

0.02

Crust, %

Physical State at STP

gray

silver-white

solid

Density at STP, g/cm 3

7.86

8.83

8.90

Melting Point, °C

1535

1490

1450

Boiling Point, °C

2730

2890

2840

Heat of Fusion, kcal/mole

3.67

3,64

4.20

Heat of Atomization,

100

102

101

kcal/mole

Ionization Potential, eV, 1st

7.90

7.86

7.63

2nd

16.10

17.3

18.2

Electronegativity

1.8

Atomic Radius, A

1.165

1.157

1.149

Ionic Radius, A (4-2)

0.75

0.72

0.69

Oxidation States

4-2, +3

+2, +3

+2, +4

Coordination Numbers

6 (+2,4-3)

4,6 (+2)
6(+3)

4,6 (+2)

Oxidation Potential, volt

M M 2+ +2e

4-0.440

+0.277

+0.250

M 3 + + e

-0.771

-1.82

M 4- 2 OH~ -» M(OH) 2

4-0.887

+0.73

+0.72

+ 2 e

M(OH) 2 + (OH)--»

+0.56

-0.20

M(OH) 8 4- e

Transition Elements*— ~1I: Iron , Cobalt , Nickel and the Platinum Metals 545

In common with the other transition elements, iron, cobalt, and nickel
form colored compounds and numerous complex ions. Most of the complex
ions show octahedral coordination. Because of their unpaired electrons the
elements are paramagnetic but so strongly are they attracted to a magnetic
field that the special term ferromagnetic is applied to them. Not only are
the electrons oriented in a magnetic field but also the atoms tend to align
themselves in the same direction and to remain so oriented when the field
is removed.

Iron

Although elemental iron is not found in nature, the wide distribution of
iron compounds in rocks and soils and the relative ease with which these
compounds are reduced by carbon made possible its early discovery and
application to the manufacture of tools and weapons. Iron is one of the
seven metals known to the earliest ancients. The technique of smelting iron
is supposed to have been discovered by the Hittites about 1400 b.c., yet iron
did not assume true industrial importance till the eighteenth century.

Comprising about 5% of the earth’s crust, iron is the fourth most abundant
element and the second most abundant metal, after aluminum. In most
geologic theories the core of the earth is presumed to be liquid iron while,
alloyed with 3-8% nickel, elemental iron is the major constituent of most
meteorites. Iron ores are found throughout the world: in the United States,
principally in the Lake Superior region and in Alabama, and also in the
U.S.S.R., France, West Germany, the United Kingdom, Canada, and Venezuela.
The United States is the world’s principal producer of iron ore, mining about
90 million tons of ore annually or about one-fourth of the world’s supply.
Even so, this country currently imports about one-quarter of its total needs.

The chief ores of iron are the oxides and the carbonate. Among these are:

Fe20 3 Iron(III) oxide (ferric oxide); red hematite; red ore

Fe 2 0 3 * n H 2 0 Brown hematite; limonite
Fe 3 0 4 Magnetic oxide of iron; magnetite; lodestone

FeCOs Iron(II) carbonate; siderite

Although quite abundant, iron pyrites , FeS 2 , is not an ore of iron but is a
source of sulfur for the manufacture of H 2 S0 4 . Compounds of iron are also
found in living matter, for example, in the chlorophyll of plants and in the
hemoglobin of blood.

2. Metallurgy of Iron. Approximately 94% of the iron ore consumed in
the United States is reduced in blast furnaces, the remainder in open hearth
furnaces. The ore is first calcined to remove water, to decompose carbonates,
and to oxidize any sulfides and organic matter which may be present. Coke
is the reducing agent used. Ores containing lime or magnesia are mixed with
an acid flux such as sand or clay and conversely, ores- containing sand or
clay are mixed with limestone in order to form a fusible slag. The blast
furnace (Figure 41.1) is a tapered cylinder 100 feet high and 25 feet in diameter
and is lined with firebrick. A charge of ore, coke, and flux is introduced at
the top while a blast of hot air is admitted through the tuyeres near the

546

Transition Elenu

nts — II: Iron » Cobalt , .V

W rte Ptamum Metals

Figure 41 J. A Blast Furnace.

lineated m^FiJurT 4 T 2 T \t °‘j c ** rri ng in the blast furnace are de-

gressively JS to w m “dTldT"* °* ,de ° f ta "i “ ^
Upper part of the furnarv* k, oxlcie and then spongy iron metal in the

is decomposed into CaO ancfccT ~? noxide * In the same Te Z io * the CaC0 3
fore the LmatWof sla 2 CaSn* I ™ T * S reduCed be '

nace; otherwise a large portion of .l 56 ^ 115 ln th< ; middle region of the fur-
SiO a with FeO to form a E of nr?™ would be lost W union of the
middle of the furnace at it ^ °i iron ^^ metasilicate, FeSi0 3 . Just below the
furnace, at its widest part, the temperature exceeds the melting

Transition Elements — II: Iron, Cobalt, Nickel and the Platinum Metals

547

FesOa+CO-*!™* 00 ’ C
2FeO+COjj- — 600“C

/FeO+CO-*-

Fe + COsV— -750’C
f CaCOj— *■

CaO+COj\

CaO+SiOi V— -1000”C

CaSi0 3

(Coke*}- Solid Iron)

{Liquidkon Uquid f 1300°C

C+CCW2CO

Air

Slag \

j-==>

Molten Iron ”j

■ Hfc

— - 1500°C

Figure 41.2. Reactions in the Blast Furnace.

point of the iron (which melts at a temperature lower than pure iron be-
cause of its dissolved carbon) and of the slag. Both the iron and the slag
melt in this region and form two immiscible liquid layers at the bottom
of the furnace; the less dense slag layer collects over the molten iron and
protects it from reoxidation. Each is withdrawn periodically from openings
at the bottom of the furnace. Some of the slag is later used in making cement.

At the level of the tuyeres vigorous combustion of the coke takes place.
Due to the presence of excess air, C0 2 is first formed. But as it passes
upward through the overlying layers of white hot coke, the C0 2 unites
with the latter to form CO which, in fact, acts as the reducing agent in the
upper region of the furnace. The hot gases leaving the top of the furnace
contain about 25% CO and are readily combustible. They are binned in
special stoves to preheat the air which enters the furnace through the tuyeres.

The metal produced by the blast furnace is known as pig iron. It contains
2, 5-4.3% carbon and often as much silicon, with varying quantities of man-
ganese, sulfur, and phosphorus. A typical pig iron analysis might give this
composition: 4.0% C; 2.0% Si; 0.80% Mn; 0.15% P; and 0.05% S. The average
blast furnace will produce about 1,000 tons of pig iron daily and about half
that quantity of slag, Once lighted the furnace is not shut down until the
refractory lining has to be replaced, usually about every three years. During
this period the furnace will produce a million tons of iron and represents
a capital investment of about $15* million, excluding the cost of auxiliary
equipment such as coke ovens.

548

Transition Elements — 11: Iron , Cobalt, Sickel and the Platinum MetaU

When pig iron is remelted and cooled in molds it is known as cast iron.
This is very hard but too brittle to be used for parts of machinery that are
subjected to high stress. Cast iron expands upon solidifying and so is used
in casting variously shaped articles such as cooking ranges, stoves, pipes,
radiators, toys, and similar devices.

3. Wrought Iron. Wrought iron is made from pig iron by removing the
major portion of its impurities; it is the purest form of ii on made commer-
cially, containing no more than 0.2% carbon. It is made by two methods. In
the more modern method, molten iron that has been purified by the Bessemer
process (Section 4) is poured slowly into a liquid iron silicate slag. The
resulting spongy iron-slag mixture is cast into ingots which are then rolled
and hammered until the excess slag is squeezed out. In the older and more
laborious method, pig iron and scrap iron are heated upon a layer of iron
oxide in a reverberatory furnace. Carbon and sulfur are oxidized by the iron
oxide and escape as CCh and SO a . Phosphorus and silicon are similarly
oxidized and then combine with excess iron oxide to form a slag. Wrought
iron has a fibrous structure, partly due to the presence of thin films of slag
between layers of pure iron. It is extremely tough and resistant to shock
so it is used to make objects that are subjected to sudden and repeated stresses
such as anchors, chains, and bolts, and also for ornamental ironwork,

4. Steel. Steel is an alloy of iron and carbon, the carbon content being
from 0.04-1.7%. Other alloying elements, such as Mn, Si, Ni, Cr, W, V, Mo,
and Co may be added intentionally to impart desirable mechanical and
chemical properties not obtainable with carbon alone, and such steels are
known as alloy steels. The ingredient which most influences the properties
of iron or steel, however, is carbon. The specific effects of some elements
added to steel and the special uses of the resultant steels are listed in
Table 41-B.

In general, steel is made by removing the impurities from pig iron and
then adding the proper quantities of carbon and other required elements.
There are four steelmaking processes. These, with their relative capacities
of current total steel production, are:

A) the Bessemer Process; 2%

B) the Open-hearth Process; 85%

C) the Oxygen Converter Process; 3%

D) the Electric Furnace Process; 10%

Bessemer Process : Industrially the first used technique for large scale
production of steel was the Bessemer process (Figure 41.3). A Bessemer
converter is a pear-shaped vessel ten or more feet high, open at the top, and
lined with a refractory such as silica or magnesia which also acts as a flux.
The converter is mounted on trunnions so that it can be tilted to permit pour-
ing out the finished product. A charge of molten pig iron, 15-40 tons at a
time, is poured into the converter and sc blast of hot air is blown through
the molten metal from a number of small openings in the bottom. Except
for phosphorus and sulfur, carbon and some of the undesirable elements are

Transition Elements — II: Iron , Cobalt , Nickel and the Platinum Metals

549

Table 41-B

Alloy Steels

Name

Composition

Characteristic properties

Uses

Manganese

10-18% Mn

Extremely hard and resist-
ant to wear.

Grinding machinery,
safes, etc.

Chrome-

1-10% Cr

Great tensile strength, re-

Axles and other parts

Vanadium

0.15% V

sistance to stress and
torsion.

of automobiles.

Tungsten

10-20% W

Retains temper at high

High-speed cutting

3-8% Cr

temperatures.

tools.

Molybdenum

6-7% Mo

(same as tungsten)

Nickel

2-4% Ni

Resists corrosion, great
hardness and elasticity.

Drive shafts, gears,
cables, etc.

Invar

36% Ni

Practically no expansion.

Meter scales and
pendulum rods.

Nickel-

1-4% Ni

High tensile strength, great

Armor plate.

Chromium

0.5-2% Cr

hardness and elasticity.

;

Alnico

20% Ni
12% A1

5% Co

Highly magnetic.

Powerful permanent
magnets.

"18-8"

18% Cr

8% Ni

Resists corrosion.

Cutlery, instruments.

Figure 41 .3. A Bessemer Converter.

550

Transition Elements — 11: Iron, Cobalt, Nickel and the Platinum Metals

oxidized and escape as gases or form slag constituents. When the impurities
are burned out, the air is shut off and the required amounts of carbon and
other elements are added to form the steel. The converter is then tipped
on its trunnions and the molten steel poured into molds. The Bessemer process
is a rapid one, taking only 10-15 minutes to produce a batch of steel but it
does not permit close control of the finished product which is a low grade
inexpensive steel suitable, however, for steel frameworks.

Open-hearth Process : In this process, pig iron usually drawn directly from
the blast furnace without solidification, scrap iron, and iron ore are placed
onto the hearth of a gas fired furnace (Figure 41.4). Most open-hearth fur-
naces are of the reverberatory type with the hearth in the shape of a large,
oval dish and have a basic lining which is a thick layer of CaO or MgO.
The oxygen of the ore oxidizes the impurities; CO escapes while oxides of
silicon and phosphorus are absorbed by the basic lining. Since the process
of oxidation takes about eight hours, samples of the steel can be withdrawn
from the furnace at intervals for laboratory analysis. When the desired
composition is attained other alloying agents may be added and the steel
poured into molds. Efficient utilization of the heat from waste gases is
made possible by a special regenerative heating system. Through a system
of valves and flues, a checkerwork of brick at each end of the furnace can

be connected to either the incoming air supply or the stack for waste gases.

The spent hot gases pass down through the checkerwork and impart heat
to it; valves are then reversed and the incoming fuel gas and air move through
the hot chambers, are preheated, and meet and burn in the form of an

enormous blowpipe or blast lamp over the charge in the hearth. The hot

gases then leave through the set of chambers on the opx>osite side of the
furnace where that checkerwork is thereby preheated.

Oxygen Converter Process : The newest entry into the field of steelmaking
is the Basic Oxygen Converter process, frequently referred to as the Linz-
Donawitz process for the Austrian towns in which it originated. This is the
fastest process for the making of high quality steel; 100 ton melts can be

Figure 41.4. An Open-Hearth Furnace.

Transition Elements — II: Iron, Cobalt , Nickel and the Platinum. Metals

551

completed in 30 minutes. The process is especially adaptable to the usual
run of basic pig iron produced in the United States. A basic oxygen con-
verter is pear-shaped, similar to, but much larger than, a Bessemer con-
verter and without its tuyeres at the bottom. It is lined with magnesia fire-
brick and is held in trunnions which allow a complete 360° rotation of the
furnace. The converter is charged with molten pig iron (70%), scrap iron
(12-24%), and a flux. A water-cooled lance is lowered to a predetermined
position, about six feet, above the molten charge. From the tip of the lance,
oxygen gas is blown directly onto the surface of the melt at jet velocities of
1,500 miles per hour and at a rate of 13,000 cubic feet per minute, producing
temperatures of 2000-2500 °C. Under the jet of oxygen the metal loses its C,
Si, and Mn content through oxidation and its density increases to about
7.1 g/ml. The liquid pig iron surrounding the purified metal just formed in
the oxygen zone has a lower density, about 6.5 g/ml. Due to this difference
in density the refined metal sinks from the reaction zone to the bottom of
the converter and is replaced by impure metal.

Electric Furnace Process : Electric furnaces of the arc type and of the
induction type are used to manufacture virtually all kinds of steel from the
simple carbon steels to the high alloy steels. These furnaces utilize nearly
100% scrap iron and produce very fine steel free from slag and from blow-
holes due to occluded gases. The removal of undesirable elements is by
the same mode as in the basic-lined open-hearth furnace. Advantages
of the electric furnace are that the temperature can be controlled and also,
because higher temperatures are possible, special steels with high melting
points can be made.

Crucible steel is a special steel made by melting wrought iron with the
required amount of carbon and other elements in a graphite crucible. f>uch
steels are of high grade and are used for razors, files, and cutlery. The famed
blades of Damascus and of Toledo were made by the crucible process.

5. Heat Treatment of Steel. Not only does the composition of the steel
determine its properties but so also does its heat treatment. Heated to red
heat and allowed to cool slowly steel is comparatively soft but if chilled
suddenly by plunging into ice water it becomes intensely hard and brittle. By
cautious reheating of the chilled steel to 250-300 °C, the brittleness is removed
but the hardness remains. The degree of hardness is controlled by the tempera-
ture to which the steel is reheated during this process of tempering . Steels
having a carbon content of 0.5-1.5% permit tempering. Most of the carbon
in steel is present as iron carbide or cementite, Fe g C. When the hot steel
is chilled suddenly the cementite remains in a state of supersaturation and the
product is hard and brittle. On reheating the cementite crystallizes and a
mixture of hard crystals of the carbide and soft iron is produced, a material
that is less brittle and somewhat softer than the original steel,

6. Properties of Iron. Chemically pure iron can be prepared by reducing
pure Fe 2 O s with H 2 and heating the product for some time in a vacuum to
remove adsorbed hydrogen. It is not made on a commercial scale. The
physical properties of iron are given in Table 41-A. As indicated by its oxida-
tion potential of +0.44 volt iron is a fairly active metal. It displaces H 2 from

552

Transition Elements — II: Iron , Cobalt, Nickel and the Platinum Metals

strong acids and from steam at high temperatures. The metal combines di-
rectly with sulfur and with moist chlorine while finely divided iron ( iron filings
or steel wool) burns in oxygen or in air. When exposed to moist air or
when submerged in water containing dissolved oxygen, iron corrodes or
rusts, forming hydrated iron (III) oxide, Fe 2 0 :i * n H 2 0.

Although iron dissolves in both dilute and concentrated HN0 3 it does
not dissolve in fuming HN0 3 nor does it exhibit the normal properties ex-
pected of iron after such treatment. It is said to be passive and does not
react with acids nor, for example, displace Cu 2+ from solutions of its salts.
If the metal is scratched or struck, however, the passivity is destroyed and
the iron then exhibits its ordinary properties. Presumably the passivity is
due to the formation of a thin surface film or coating of iron oxide which
protects the underlying metal. Other metals such as Co, Ni, Mo, W, and Cr
can also be rendered passive by suitable treatment.

Corrosion is a, general term which refers to the undesirable conversion
of a free metal or an alloy to a compound, generally in a natural or in-
dustrial environment. The corrosion of iron is known as rusting, a process
of great economic importance since about 20% of the annual iron production
goes to replace that lost by rusting, a loss exceeding $3 billion. Rust is a
red-brown solid, a hydrated iron (III) oxide of varying composition best
represented by the formula Fe^Oa * n H 2 0. Rust formed on a metallic surface
does not adhere to it; the rust flakes off, exposing the underlying metal to
further corrosion. For iron to rust, both 0 2 and H 2 0 must be present. The
over-all reaction for the rusting of iron, an oxidation-reduction reaction, is

(1) 4 Fe + 3 0 2 + 2n HX> 2(Fe 2 O i * n H 2 0)

The reaction does not take place in one step. Despite intensive research, the
actual mechanism, or the steps in the rusting process, is not completely under-
stood. One possible mechanism is:

(2) Fe Fe 2+ + 2 e l oxidation of Fe to the +2 state by H + ions

(3) 2 H + + 2 e 2 H j ^ rom either acid or H 2 0 with the formation of

( free H radicals

(4) 4 H -f 0 2 -» 2 H 2 0 (oxidation of H to H 2 0)

(5) 4 Fe 24 * + 0 2 + (4+2n)H,0 -» 2(Fe«Oa * n H a O)+8 H+ (rust formation)

This mechanism is in agreement with the experimentally observed facts
that the presence of acid increases the rate of corrosion and that ordinary
iron and steel, which contain a number of substances of lower oxidation
potential than the iron, corrode more rapidly than does pure iron. Such 'im-
purities in the iron form with it a voltaic cell or an electric couple in which
iron is the anode. Rust itself forms a couple with pure iron so that once
it is produced corrosion proceeds more rapidly, a type of autocatalysis.
When iron is strained in any way, such as by bending or pressing, it becomes
more reactive and as a result corrosion is more rapid in the neighborhood
of punched holes than around drilled holes in structural steel.

Rustproofing is effected by coating iron with an impervious film, such as
oil, grease, paint, chromium, zinc, tin, etc., which, prevents the iron from

Transition Elements — II: Iron, Cobalt, Nickel and the Platinum, Metals

553

coming in contact with 0 2 or H 2 0, or both. By exposing the metal to steam
a thin, continuous, and adherent film of magnetic oxide of iron, Fe 3 0 4 ,
can be formed on its surface to protect it from rusting. Electroplating the
iron with zinc or dipping it into a bath of molten zinc will protect the iron
from rusting even though a hole is cut through the zinc down to the iron
base. Though both metals are thereby exposed to oxidizing conditions the
zinc, having the higher oxidation potential, would be oxidized in preference
to the iron. On the other hand, tin plating does not protect the underlying
iron in the event of a puncture. In such circumstances the iron is more
readily oxidized and, indeed, protects the tin from corrosion. The protection
of iron through the preferential oxidation of more active metals such as
zinc or cadmium is an example of cathodic protection. In a voltaic cell con-
taining iron and zinc electrodes, the iron would be the cathode and the
zinc the anode. Hence to protect iron pipes, or the steel hulls of ships, these
may be connected to more active metals known as sacrificial electrodes
because they are preferentially oxidized while the iron is unaffected.

7. Compounds of Iron; Iron(II) Compounds. Iron forms two series, of
compounds, iron (II) or ferrous , in which the oxidation state is +2 and
iron (III) or ferric , in which the oxidation state is +3. Iron (II) compounds
are generally green in color and are readily oxidized in the iron (III) state,
even upon exposure to air; iron (II) solutions can be kept in the +2 state
by having present some iron metal and acid. When iron reacts with a non-
oxidizing acid such as HC1 or dilute H 2 S0 4 , Fe 2+ is produced.

(6) Fe + 2 H+ -> Fe 2+ + H 2

Solutions of Fe 2+ ion are but slightly hydrolyzed since Fe(OH) 2 is a fairly
strong base. By the reaction of Fe 2+ ion and a strong base, Fe(OH) 2 is
precipitated as a white gelatinous material which turns brown due to its
oxidation to Fe(OH) 3 or a hydrated Fe 2 0 3 .

(7) Fe 2 -f- 2 OH" Fe(OH 2 (s)

(8) 4 Fe(OH) 2 + 0 2 + H 2 0 4 Fe(OH) 3

The most important iron (II) compound is the sulfate, FeS0 4 ■ 7 H 2 0,
known commercially as green vitriol or copperas and used as a mordant
in dyeing, as a disinfectant and weed killer, and in the manufacture of
black ink. It is formed as green monomclinic crystals by evaporating a
solution prepared by the reaction of iron and dilute H 2 S0 4 . A considerable
quantity of the salt is recovered from pickling baths into which sheet iron
was dipped in order to remove oxide scale prior to its being coated with zinc
or enamel. A stable iron(II) compound is the double salt, iron (II) ammonium
sulfate, (NH 4 ) 2 S0 4 • FeS0 4 * 6 H 2 0, known as Mohr’s salt. This compound
is not oxidized by air and so is used to prepare standard solutions of Fe 24 *
ion employed for oxidation-reduction reactions in analytical chemistry. Iron(II)
sulfide, FeS, is prepared by the direct union of iron and sulfur or as a black
precipitate from the reaction of Fe 2+ and S 2 ' ions in basic solution, a reaction
important in the qualitative detection of iron.

The slow oxidation of Fe 2+ ion to Fe 3+ ion by oxygen of the air is
utilized in the manufacture of writing ink. Iron (II) tannate, which is present

554

Transition Elements — 11: Iron , Cobalt, Nickel and the Platinum Metals

in inks as iron ( II ) sulfate and tannic acid, is a soluble colorless salt. When
the ink is spread on paper the Fe 2+ ion is slowly oxidized to Fe 3+ ion and
a black insoluble iron (III) tannate is formed. A blue or black dye is added
to the ink to make the writing visible initially. Fresh ink stains usually can
be washed out with water if applied at once, but after oxidation of the Fe 2 +
ion a reducing agent must be used to reduce the iron (III) compound to the
soluble iron (II) tannate.

Iron(lII) compounds : Iron (III) salts are usually yellow or red and their
solutions are quite acidic due to rather extensive hydrolysis of Fe(H 2 0) 6 3+ ion.

(9) Fe(H 2 0) 6 3 + + H.O -> Fe(H 2 0) 5 (0H) 2+ + H,0+ 6.3 X 10~ 3

The brown color in solutions of Fe 3+ salts is probably due to the hydrolytic
formation of iron (III) hydroxide, Fe(OH) 3 , which remains in colloidal dis-
persion. By the reaction of larger concentrations of Fe 3 + and OH~ ions,
Fe(OH) a is formed as a red-brown, gelatinous precipitate. Dehydration of
the Fe(OH) 3 yields Fe 2 0 3 , which is used as “rouge” and the pigment,
Venetian red. The most important iron (III) compound is the chloride, FeCl 3 ,
prepared by passing Cl 2 over heated iron or through FeCl 2 solution. It is
very soluble and forms several hydrates, of which the most important is
FeCls * 6 H 2 0, a yellow deliquescent solid. Solutions of Fe 3+ salts are
reduced by H 2 S or by S 2 " to the Fe 2+ state.

8. Complex Ions. Both oxidation states of iron form numerous complex
ions. Of special note are those formed with cyanide ion, CN", the hexa-
cyanoferrate(ll) or ferrocyanide ion, Fe(CN) 6 1 % and the hexacyanofer-
rate(III) or ferricyanide ion, Fe(CN) 6 3 ~.

Though similar in formula and in structure, these ions differ in oxidation state;
the o and the i of ferrocyanide and ferricyanide, respectively, indicate that they
are derived from the older terminology of ferrous and ferric.

The Fe(CN) 0 4 ~ ion is colored yellow while the Fe(CN) 6 3 ~ ion is red. These
ions are extremely stable, dissociating not at all into iron or cyanide ions and
hence can be considered as single entities just as, for example, the SO* 2 " ion.
Because of the lower oxidation state of the iron therein, the Fe(CN) 6 4 " ion
can act as a reducing agent and can be oxidized to Fe(CN) 6 3 ~ ion; this is
done industrially with Cl 2 . The most common salts of these ions are those
of sodium and potassium, e.g., K 4 Fe(CN) 6 and K 3 Fe(CN) 6 . With other metal
ions characteristic precipitates useful in analytical chemistry are formed. The
products of the reactions with Fe 2+ and Fe 3 ^ ions are of interest and are
tabulated below.

Fe 2+

K 2 Fe[Fe(CN) 6 ]
a white precipitate

FeK[Fe(CN) 6 ]
a blue precipitate;

Turnbulls blue

Fe 3 +

FeK[Fe(CN) 6 r

a blue precipitate;
Prussian blue

no precipitate;
brown solution

Reagent
K*Fe ( CN ) 6

K 3 Fe(CN)<j

Transition Elements — II: Iron , Cobalt , Nickel and the Platinum Metals

555

The composition of the precipitates varies with the conditions under which
the reaction is carried out. There is some suspicion that Prussian blue and
Turnbull's blue are identical compounds inasmuch as they have the same
molecular formula. Nevertheless the reactions indicated afford a means of
testing for Fe 2 + and Fe 3 + ions and for distinguishing between them.

Prussian blue is employed in making paints and laundry bluing whereas
Turnbull's blue is formed when blueprints are developed. Blueprint paper
is impregnated with a mixture of iron (III) salts, e.g., iron (III) ammonium
citrate, and K 3 Ke(CN) 6 . Upon exposure to strong light (under a photographic
negative or an ink drawing on transparent paper) part of the Fe 3 * salt is
reduced to the Fe 2 * state. The print is developed by immersing it in water;
Turnbull's blue is precipitated where reduction took place and the remainder
of the salts is washed away, leaving a positive print in white on blue.-

The Fe 3 * ion also combines with thiocyanate ion, SCN~ to form a series
of complex ions from Fe(SCN) 2 * to Fe (SCN)* 3 -. The latter has a deep red
color which is detectable at concentrations of 10' 5 M and hence is used as
a sensitive test for Fe 3 * The Fe 2-1- ion does not react with SCN" ion; any
red coloration given by it is due to Fe 3 * as an impurity.

9. Cobalt. The name of the metal is derived from the German word
Kobolcl, meaning evil spirit or goblin. It was applied by miners in Saxony,
about 1500 a.d., to arsenic-bearing cobalt ores that looked as if they were
copper-bearing but which evolved poisonous fumes on smelting and which
would not yield metals when treated by ordinary metallurgical methods.
Such characteristics were attributed to supernatural influences. The metal
was first isolated by Georg Brandt in 1735,

Cobalt is a relatively rare element, constituting but 0.001% of the earth's
crust. Its principal minerals are smalt it e, (CoNi)As 2 : cobaltite, CoAsS; carrol-
Ute, Co 2 CuS 4 ; and linnaeite, Co 3 S 4 , but seldom are any cobalt minerals found
in sufficient quantity to be mined for cobalt alone. Consequently, cobalt is
produced mainly as a by-product or a coproduct with other metals, principal-
ly Cu, Ni, Fe, and Pb, and cobalt ores are usually worked to obtain cobalt
salts rather than the free metal. Extraction of cobalt from its ores or con-
centrates requires various complex chemical processes to accommodate the
diverse characteristics of individual ores, mainly to separate the cobalt
from the other metals with which it occurs. The final step is either oxidation
of a cobalt oxide, usually in an electric furnace, or electrolysis of a cobalt
sulfate solution. Small quantities of the metal are also produced by alumino-
thermy. Most cobalt is used in manufacturing alloys such as high speed tool
steels, high temperature alloys, and permanent magnet alloys. Stellite , a non-
magnetic cobalt alloy (55% Co; 25% Cr; 15% W; 5% Mo) is hard and non-
corrosive and is used for high speed tools and surgical instruments. Alloys
containing up to 65% cobalt withstand operating temperatures of 1000°C and
are used as nozzle vanes and tail cone$ of jet engines. Cobalt is more ferro-
magnetic than iron; in the Alnico series of permanent magnets cobalt is
alloyed with varying proportions of Al, Ni, and Fe.

The properties of cobalt are given in Table 41-A. Pure cobalt metal is
silver-white with a reddish tinge. It dissolves slowly in acids and is made
passive by concentrated HNO g . Like iron, cobalt forms two series of com-

556

Transition Elements — 11: Iron , Cohalt , Nickel and the Platinum Metals

pounds, cobalt (II) or cobaltcus, and cobalt (III) or cobaltic but unlike iron
the Co 2+ state is stable. Any attempt to oxidize Co 2+ to the Co 3+ state
results in the precipitation of cobalt (III) hydroxide, Co(OH) 3 . The Co 3+ ion
is a sufficiently strong oxidizing agent to oxidize H 2 0 to 0 2 . Few simple
cobalt (III) salts are known, among them CoF 3 and Co 2 (S0 4 )3‘18 H 2 0,
and these decompose in aqueous solution. Both oxidation states of cobalt
form a tremendous number of complex ions. Of these the most stable are
derived from the Co 3+ state; examples are Co(CN) 6 3+ , [Co(NH 3 )5(H 2 0)] 3 ‘ l ' 3
[Co(NH 3 ) 5 Cl] 2 +, and Co(N0 2 ) 6 3 -.

In solution or in the hydrated state, cobalt (II) salts are red or pink due
to the Co(H 2 0) 6 2+ ion whereas the anhydrous salts are blue. Thus, gentle
heating of the pink salt, CoCl 2 * 6 H 2 0 (truly [Co(H 2 0) 6 ] Cl 2 ), turns it blue
but standing in a moist atmosphere restores the pink color. On account of
these changes the salt is used in sympathetic inks and in toy instruments
intended to forecast the weather. Cobalt (II) oxide, CoO, is black and dis-
solves in molten glass, imparting a blue color to it; it is used to produce
stained glass and a blue glaze on pottery,

10. Nickel. Toward the end of the seventeenth century the German term
kupfernickel , or "Old Nick's Copper/' was applied to ores resembling copper
ores from which no copper could be smelted but in its stead a hard white
metal which could not be fabricated into useful articles. In 1751, A. F. Cron-
stedt showed that these ores contained a new metal to which he gave the
name nickel. Though there are few large deposits of nickel ores, the element
constitutes about 0.02% of the earth's crust, an amount about twice as
much as that of Cu, Zn, and Pb combined. The principal ores are
pentlandite, NiS * FeS, and garnierite , a hydrous nickel magnesium silicate,
(Ni,Mg)SiG 3 * n H 2 0. The metallurgy of nickel is similar to that of cobalt
and involves many types of processes depending upon the characteristics of
the individual ores. Smelting in a blast furnace yields a matte of iron, copper,
and nickel sulfides. Through a recent advance in technology the nickel can
be obtained directly by electrolysis of the molten matte, or the matte can be
roasted to oxides and then reduced with carbon to produce a mixture of
metals from which the nickel can be extracted by the Mond process. In this
advantage is taken of the fact that nickel forms a compound with carbon
monoxide, Ni(CO) 4 . This is a colorless liquid which boils at 43°C but de-
composes with the deposition of elemental nickel when its vapor is heated to
200° C. Thus when CO is passed over the mixture of matte metals at about
80°C the volatile Ni(CO) 4 is formed and is subsequently decomposed at
200 °C to deposit nickel.

The largest use of nickel is in nonferrous alloys such as Monel metal
(67% Ni; 30% Cu; 3% Fe), in coinage, and in "nickel silver/' a Cu-Ni-Zn
alloy which resembles silver. A small percentage of nickel in steel improves
its strength and resistance to shock, as in armor plate, while certain types
of stainless steels, containing 6-22% Ni, are used widely in the chemical
and food processing industries. Much nickel is electroplated for decorative
and protective purposes while the finely divided element is used as a catalyst
for the hydrogenation of oils.

Transition Elements — II: Iron , Cobalt , Nickel and the Platinum Metals

557

The physical and chemical properties of nickel are similar to those of
cobalt. Though salts of Ni J+ ion are readily prepared no simple salts of Ni 3+
ion are known. The color of the hydrated Ni 2+ ion is green. The reaction of
Ni 2+ and OH~ ions precipitates the light green Ni(OH) 2 . When this is
treated with a strong oxidizing agent a black oxide is formed whose composi-
tion is not definitely known but which may be NiO a . This is itself a strong
oxidizing agent and is used as the cathode of the Edison storage battery. Many
complex ions are formed by nickel, most with a coordination number of
six such as the octahedral Ni(NH 3 ) 6 2+ but some with a coordination number
of four such as the tetrahedral Ni(NH 3 ) 4 2+ and Ni(CO) 4 and the planar
Ni(CN) 4 2 -.

11. The Platinum Metals. The second and third triads of Group VIII
of the Periodic Table are likewise transition elements analogous to Fe, Co,
and Ni, and collectively are known as the platinum metals. Ruthenium,
rhodium, and palladium are called the light platinum metals because their
densities are much less than those of osmium, iridium, and platinum, the
heavy platinum metals. The electron configurations for these triads of ele-
ments are given below and their properties are listed in Table 41 -C.

Element

Atomic

Number

2s 2p

3s 3p

4s 4p 4 d

5s 5p 5d 5 f

g 51 * 1

Ruthenium

Rhodium

Palladium

Osmium

6 6

Iridium

6 7

Platinum

6 9

The exploring Spaniards in Mexico and South America found a white
metal which resembled silver and named it “platina,” the Spanish word for
Tittle silver/’ To prevent its being used to adulterate gold die Spanish gov-
ernment prohibited the export of the metal from South America and ordered
a shipment to be thrown into the sea, at a location not precisely known. Today
all the platinum group metals, except palladium, command prices higher
than gold— and gold is sometimes used to adulterate platinum. Annual world
production of all six metals is about 35 tons, valued at about $35 million.
The Union of South Africa, Canada, and the U.S.S.R. are the leading pro-
ducers of these metals. Each metal has specific and important uses and Pt,
Pd, and Ir are among the strategic metals being stockpiled. Native platinum
contains all the platinum group metals; these were isolated during the study
of natural platinum during the early 1800’s. As indicated by their negative
oxidation potentials the elements have little chemical reactivity. They occur
mainly in the free state as placer deposits in the form of small grains or
scales in sand and gravel, mixed with each other or with the ores of other

558

Transition Elements — II: Iron, Cobalt , Nickel and the Platinum Metck

Table 41-C

Properties of the Platinum Metals

Property

Palladium

Osmium

Iridium

Platinum

Symbol

Atomic Number

Atomic Weight

101.07

102.905

106.4

190.2

192.2

195.09

Isotopes

96,98,

99,100,

101,102,

104

103

102,104,

105,106,

108,110

184,186,

187,188,

189,190,

192

191,

193

190,192,

194,195,

196,198

Abundance in Earth’s

Crust, %

10-n

10-7

10-io

5 x 10-’

Physical State at STP

gray

solid

white

solid

silver

solid

gray

solid

silver

solid

silver

solid

Density at STP, g/cm : *

12.80

12.42

12.0

22.7

22.65

21.45

Melting Point, °C

2400

1965

1555

2700

2450

1770

Boiling Point, °C

3700

3560

4400

4350

4010

Heat of Fusion, kcal/mole

6.1

5.2

4.0

7.0

■ 6.3

4.7

Heat of Atomization,
kcal/mole

144

130

160

150

135

Ionization Potential, eV, 1st

7.36

7.46

8.33

8.7

9.2

9.0

Atomic Radius, A

1.241

1.247

1.278

1.255

1.260

1.290

Oxidation States

2,3,4,

6,7,8

1,3,4

2,4

2,3,4,

6,8

1,2,3,

4,6

2, 3, 4, 6

Coordination Numbers

4(4-2)

6(+4)

6(4-4)

4(4-2)

6(4-4)

4(4-2)

6(4-3)

6(4-4)

4(4-2)

6(4-4)

Oxidation Potential

M M 2+ + 2 e

-0.45

-0.6

-0.987

-0.9

-1.1

-1.2

metals such as Au, Cu, Cr ? and Fe. Separation of the heavier platinum-bearing
particles from the less dense sand yields grains containing over 75% Pt.

12. Ruthenium and Osmium. Ruthenium (for Ruthenia, Ukraine) was
isolated in 1844 by Karl Klaus. Osmium (Greek, osme: to smell, because of
the odor of 0s0 4 ) was discovered by B. Tennant in 1803 in the residue left
after the solution of crude platinum in aqua regia. These metals, the first
in each triad and below iron in the Periodic Table, resemble the latter in
some respects. They are gray metals, not unlike iron, while the other platinum
metals are white and more like cobalt and nickel. Osmium has the greatest
density of any known substance. Both metals are hard and brittle and are
used in the making of hard alloys for fountain pen tips, phonograph needles,
and machine bearings. Heating the metal in air forms their tetroxides, RuO t
and 0s0 4 ; these are solids with low melting and boiling points. The vapor

Transition Elements — II: Iron, Cobalt, Nickel and the Platinum Metals

559

of 0s0 4 is poisonous. Although not an acid, 0s0 4 is known as osmic acid,
presumably because of the strong oxidizing power of osmium (VIII).

13. Rhodium and Iridium. Rhodium (Greek, radon : rose), named for the
rose-red color of its salts, was discovered in 1803 by W. H. Wollaston while
iridium (Greek, iris : rainbow), named for its green* red, and violet salts,
was isolated by Tennant, also in 1803. The metals are hard and are used
principally as alloying agents to improve the properties of platinum and
palladium. An alloy of 90% Pt and 10% Ir was chosen by the International
Committee on Weights and Measures to preserve the original standards of
length and weight. Electroplated rhodium is- used in reflectors, electronic
components, and in jewelry because of its high brilliance and corrosion re-
sistance. Radioisotope Ir 102 is used in radiography. Roth metals resist the
action of aqua regia but react with Cl 2 at high temperatures and form a
number of complex ions analogous to those of cobalt.

14. Palladium and Platinum. Palladium was named for the asteroid Pallas
which was discovered the same year, 1803, that Wollaston discovered the
metal. Of the entire group, platinum is the most important and the most
abundant element. It is a soft, silver-white, malleable and ductile metal which
can be welded at red heat. Because of its outstanding resistance to heat,
oxidation, and most chemical attack, platinum is used to make jewelry and
laboratory utensils. The metal is attacked by Cl 2 and dissolves in aqua
regia; it reacts with fused alkalies such as KOH to form platinates, Pt(OH) 6 2_ ,
but the alkali carbonates such as Na 2 C0 3 have no effect upon it. With lead
and tin, platinum readily forms fusible alloys while with sulfur, silicon, and
phosphorus it becomes very brittle so that care must be exercised in the
laboratory to avoid heating in platinum apparatus any compounds that yield
these elements in the free state. Probably half the platinum in the United
States is used as a catalyst for the manufacture of high octane gasoline and for
ammonia oxidation in the production of HN0 3 . When a little finely divided
platinum, known as platinum black, is introduced into a mixture of H 2 and
0 2 , the mixture explodes.

For both platinum and palladium the most important oxidation states are
+2 and +4; both elements also form a large number of complex ions. The
most important compounds of platinum are the chloroplatinates, PtCl 6 2- ;
the dissolving of the metal in aqua regia is due to the formation of this ion.

(10) 3 Pt + 16 H+ + 18 Cl- + 4 NOr -4 3 PtCW 2 - + 4 NO + 8 H 2 0

Evaporation of the resulting solution yields chloroplatinic acid, H 2 PtCl e * 6 H 2 0.

The chemistry of palladium and its industrial uses are similar to those of
platinum. Palladium is the only element of the group which reacts with HNO s
alone. Since finely divided palladium has the property of adsorbing several
hundred times its own volume of hydrogen gas, the metal makes, an effective
catalyst for hydrogenation reactions. The hydrogenated metal is a good re-
ducing agent; it will reduce Hg 2+ , Fe 3 ^*, and ions of metals having a lower
oxidation potential than hydrogen,

560

Transition Elements-IL Iron, Cohalt , Nickel and the Platinum Metals

QUESTIONS

1. Outline the steps in the metallurgy of iron starting with Fe 2 0 3 as ore. Write
equations for the reactions which, occur in the blast furnace. What is the pur-
pose of using limestone and coke in the smelting process? Why must all the
iron ore be reduced before the slag starts to form?

2. Distinguish among (a) pig iron (b) wrought iron (c) steel.

3. Describe the processes for the manufacture of steel.

4. Why is the lining of an open hearth furnace made of lime and magnesia?

5. What is meant by “passivity”? How can iron be made passive?

6. Propose a reasonable mechanism for the rusting of iron. Why does iron,
which has started to rust, rust more rapidly than at the start? Discuss methods
for the prevention of rusting of iron and steel.

7. Write formulas for (a) complex ions which illustrate the coordination num-
bers of iron (b) the cyanide complex ions of Fe 2+ and Fe 3 +.

8. Explain why the stability of the +2 state toward oxidation increases in the
order: iron, cobalt, nickel.

9. From the viewpoint of electron configuration why are the properties of iron,
cobalt, and nickel similar? Does each element form a stable series of com-
pounds in the^+2 and +3 states? What electrons are used in bonding in each
ionic state formed?

10. What electrons are used in bonding in the following complex ions Ni(NH 3 ) 4 2+ ;
Ni(NH s ) 6 2 +; Co(NH 3 ) 6 3 +; Co(CN)*H>

11. Which ion, Fe(CN) 6 3 - or Fe(CN) 6 4 ~, would be expected to have a magnetic
moment?

12: Write balanced equations for the following reactions: (a) the conversion of
Fe 2 O s to FeS0 4 (b) preparation of Fe 2 (S0 4 ) 3 from Fe(OH) 2 (c) oxidation
of FeS0 4 by air in the presence of H 2 S0 4 (d) reduction of FeCl 3 by H 2 S
in acid solution.

13. What chemical tests can be used to distinguish between (a) Fe 2 + and Fe 3+
(b) Co 2 + and Fe 3 + (c) Fe 2 + and Ni 2 +' (d) C'o 2 + and Ni 2 +?

14. Explain (a) the acidic reaction of a solution of FeCl 3 (b) the change in
color of a FeS0 4 solution from green to yellow on standing (c) the increased
rate of corrosion of a tin can cut through to its iron base.

15. The oxidation potential of oxygen in acid solution is -1.3 volt and in basic
solution -0.4 volt. Explain why air will oxidize Co 2 + in basic solution but not
in acid solution.

16. What industrial processes use platinum as a catalyst? List other uses of the
platinum metals.

17. What weight of hematite, Fe 2 0 3 , containing 30% silica and 10% water, is

required to produce one ton of iron? Ans: 2.33 tons

18. What weight of FeS0 4 will be oxidized by 400 ml of 1.0N KMn0 4 in acid

solution? Ansi 60.8 g

19. What volume of 0.10N KMn0 4 is required to oxidize 0.25 g of FeS0 4 in

acid solution? A ns: 33 ml

20. What volume of Cl 2 gas at 20°C and 770 mm is required to oxidize 2.0 liters

of 2. ON FeClj solution? Ans : 47.4 liters

21. What weight of FeCl 2 will be formed by the reduction of 250 g of FeCl s

in solution with elemental Fe? Ansi 292 g

Transition Elements — II: Iron , Cobalt , Nickel and the Platinum Metals

561

22. A sample of steel weighing 2.5 g is heated in oxygen till all the carbon in the

sample is converted to C0 2 . The gas is absorbed in 14.510 g of a KOH
solution. At the end of the experiment the KOH solution weighs 14.588 g.
What is the percent of carbon in the steel? Ans : 0.85%

23. (a) What weight of nickel is present in an electroplated nickel film on

one surface of a sheet of copper 1.0 meter long and 0.50 meter wide, if the
nickel coating is 0.007 mm thick? (b) What quantity of electricity was re-
quired to plate out this quantity of nickel? Ans : (a) 31.1 g

24. Calculate the concentration of Fe 3+ ion in the saturated solution formed

by the addition of a solution of FeCl 3 to an equal volume of 1.0 M NH 3 which
also contains 21.4 mg/ml of NH 4 C1. Ans: 6.6 x 10~ 25 mole/liter

Transition Elements— III
The Elements of
Groups VIA and VIIA

1* The Elements of Group VIA- Group VIA of the Periodic Table in-
cludes the elements chromium, molybdenum, and tungsten. The electron
configurations of these metals are given below and their properties are listed
in Table 42-A.

Element

Atomic

Number

2s 2 p

3s 3 p 3 d

4.9 4p 4 d 4f

5 s 5 p 5 d

p
6 s

Chromium

2 6

2 6 5

Molybdenum

2 6

2 6 10

2 6 5

Tungsten

2 6

2 6 10

2 6 10 14

2 6 4

2. Chromium. Chromium (Greek, chroma : color) was discovered in 1798
by the French chemist, L. N. Vauquelin, while analyzing a new mineral, the
“red-lead of Siberia"" or crocoisite , lead chromate, PbCr0 4 ; the element was
named for the varied colors of its compounds. Chromite , Fe(Cr0 2 )2, is the
sole commercial ore of chromium. The pure metal can be prepared by re-
ducing the oxide with aluminum. However; since the most important , use of
chromium is as a component of steel alloys it is unnecessary to prepare the
pure metal. An alloy of iron and chromium, ferrochrome , is produced by the
reduction of chromite or of a mixture of Fe 2 0 3 and'Cr 2 0 3 with carbon in
an electric furnace.

(1) Fe(Cr0 2 ) 2 + 4 C -» Fe + 2 Cr + 4 CO

From its high oxidation potential it can be inferred that chromium -is an
active metal. It dissolves in dilute HC1 and H 2 S0 4 , liberating H 2 and form-
ing Cr^ + ion. If chromium is dipped into concentrated HN0 3 , however, it

Transition Elements— -III: The Elements of Groups VIA and VII A

563

Table 42 - A

Properties of the Elements of Group VIA

Property

Chromium

M olybdenum

Tungsten*

Symbol

Atomic Number

Atomic Weight

51.996

95.94

183.85

Isotopes (mass numbers
and percent)

50( 4.31)
52(83.76)

53 ( 9.55)

54 ( 2.38)

92(15.05)

94 ( 9.35)
95(15.76)
96(16.56)

97 ( 9.60)
98(24.00)

100 ( 9.68)

ISO ( 0.13)
182(26.31)
183(14.28)
184(30.64)
186(28.64)

Abundance in Earth’s

Crust, %

0.037

10- 6

5 x 10-5

Physical State at STF

blue-white

solid

silver-white

solid

gray

solid

Density at STP, g/cm 3

7.14

10.2

19.26

Melting Point, °C

1890

2620

3370

Boiling Point, °C

2500

4800

5900

Heat of Fusion, kcal/mole

3.3

6.6

8.4

Heat of Atomization,
kcal/mole

80.5

156

202

Ionization Potential, eV, 1st

6.76

7.18

7.98

Electronegativity

1.7

2.1

2.3

Atomic Radius, A

1.17

1.29

1,23

Ionic Radius, A (+6)

0.52

0.82

0.88

Oxidation States

+2,4-3, +6

+2,+3,+4

+5,+6

+2, +3, +4
+5, +6

Oxidation Potential, volt
for M M 3+ + 3 e

+0.74

+0.2

+0.05

*Tungsten is also known as Wolfram.

becomes passive and does not react with HC1 or other reagents. As with
iron, this passivity is presumably due to an impervious oxide film so thin that it
is transparent. Other oxidizing agents and exposure to air for a considerable
time also produce this passive condition. Apart from its lustrous beauty, for
this reason chromium metal is used to plate steel and copper objects.

Chromium metal is electroplated from a bath of sodium dichromate,
Na 2 Cr 2 0 7 , and H 2 S0 4 , Though the metal is most commonly known in the

564

Transition Elements — III: The Elements of Groups VIA and VII A

form of chromium plate, about 0.0075 cm thick, on automobile trim and
household appliances, its main use is in the manufacture of steel alloys
to which, in concentrations of 12% or more, it imparts resistance to corrosion,
heat, and impact. Common stainless steel is approximately 64% Fe, 18% Cr,
and 8% Ni. Fifty percent of all chromite mined is used in the production
of alloys, 40% in the manufacture of chromite brick for refractory furnace
linings, while the remainder is used to manufacture chromium chemicals.

3. Oxidation States of Chromium (Table 42-B). Chromium shows three
oxidation states, +2, +3, and +6. Only in the +2 and +3 states does
chromium form the positive ion of salts while in the + 6 state, chromium
exists only as an acid anhydride or as a part of an oxy-anion. The general-
ization that the acid character of the oxides of an element increases with an
increase in its oxidation number is well illustrated by chromium. Chro-
mium (II) oxide, CrO, is basic, chromium (III) oxide, Cr 2 0 3 , is amphoteric,
while chromium (VI) oxide, CrO s , is acidic.

Chromium (II), or chromous compounds, such as CrCl 2 , can be pre-
pared by the reaction of elemental chromium and acid or by the reduction
of the Cr 3+ ion with an active reducing agent such as Zn or H 2 S0 3 . The
+2 state is unstable, in aqueous solution being readily oxidized to the +3
state and hence is rarely seen in the laboratory. Although solid chromium(II)
compounds have various colors, aqueous solutions have a blue color due
to the Gr(H 2 0) 6 2+ ion.

The chromium (III), or chromic state, is the most stable of the three.
Chromium (III) oxide, Cr 2 0 3 , the most common oxide of chromium, is pre-
pared by:

(2) (NH^CraOT--* Cr 2 0 3 -f N 2 + 4 H 2 0 (heating ammonium dichromate)

(3) Na 2 Cr 2 0 7 + S Cr 2 0 3 + Na 2 S0 4 (heating a dichromate with sulfur)

Cr 2 O a is an infusible, green solid, very slowly affected by acids, and as
chrome green is used as a pigment in paints. It has been already noted that
chromium (III) forms a large number of complex ions with the ligands,
NH 3 , H z O, CN-, halide ions, and others. In most of these the coordination
number of the chromium is six, the shape of the ion is octahedral, and the
bonding to the ligand is of the inner d 2 sp 3 type. Aqueous solutions of
chromium (III) are violet due to the presence of the Cr(H 2 0) 6 3+ ion. Ad-
dition of a high concentration of Cl" ion to such a solution yields a green
solution because of the displacement of H 2 0 molecules by Cl" ions and the
formation of the [Cr(H 2 0) 4 Cl 2 ] + ion. Solutions of chromium (III) salts are
acid due to hydrolysis.

(4) Cr(H 2 0) 6 3+ + H 2 0 -> [Cr ( H 2 0 ) 5 ( OH ) ] 2+ + H a O+

When NH 3 is added to a solution of a Cr 8 ~ salt, chromium (III) hvdroxide
forms as a pale blue, gelatinous precipitate.

(5) Cr(H 2 0) 6 3+ ’+ 3 OH- Cr(0H) 3 (H 2 0) 3 + H 2 0

The hydroxide, sometimes written Cr(OH) 3 * 3 H 2 0 or simply Cr(OH)a,
is amphoteric, dissolving in both strong acid and strong base.

Transition Elements— III: The Elements of Groups VIA and VllA

(6) Cr(0H) 3 (H 2 0) 8 + 3H+-» Cr(H 2 0) e »+

(7) Cr(0H) 3 (H 2 0) 3 4 OH" -» [Cr(0H) 4 (H 2 0) 2 ]~ + H 2 0

Dehydration of the [Cr(0H) 4 (H 2 0) 2 ]‘ ion ultimately yields the chromite,
Cr0 2 ‘ ion.

In the 4-6 state the most important compounds of chromium are the
chromates, Cr0 4 2 ", and the dichromates, Cr 2 0 7 2 ". Chromate compounds can
be prepared by the oxidation of the +3 state in an alkaline medium. The
technique of fusion with an alkali metal hydroxide or carbonate is a general
method for the preparation of soluble alkali metal salts of insoluble metal
acidic oxides.

(8) 2 Cr 2 0 3 + 4 K 2 CO s 4 3*0 2 — > 4 K 2 Cr0 4 4* 4 C0 2

The alkali metal chromates are soluble but those of the heavy metals, such
as PbCr0 4 , are not.

Table 42-B

Oxidation States of Chromium

Oxidation

State

Ionic

Species

Oxide

Hydroxide

Character

Derivative

Color

Cr 2 +

CrO

Cr(OH) a

basic

CrClj

blue

Cr»+

Cr 2 0 3

Cr(OH) 3

amphoteric

CrCl 3

green

Cr(OH) 4 “

(HCr0 2 )*

NaCr(OH) 4

green

Cr0 2 ~

NaCr0 2

green

Cr0 4 2 “

CrO*

(H 2 CrO.)

acidic

K 2 Cr0 4

yellow

Cr 2 0 7 2 "

H 2 Cr 2 0 7

K 2 Cr 2 0 7

orange

•Compounds whose formulas are inclosed in parentheses are hypothetical and have
not been prepared.

Chromic acid, H 2 Cr0 4 , which might be expected to be formed upon the
addition of an acid to a chromate salt, is not stable. Instead, a condensa-
tion occurs through the elimination of H 2 0 with the formation of the
dichromate ion, Cr 2 0 7 2 "; dichromic acid, H 2 Cr 2 07, however, exists only in
solution.

(9) 2 CrO, 2 - 4-2H + ^ Cr 2 0 7 2 ~ 4 H 2 0 K= 4.2 X 1W 4

Because the Cr0 4 2 " ion is yellow and the Cr 2 0 7 2- ion is orange a shift in
the equilibrium with a change in pH is apparent through a change in color
of the solution. Le Chatelier s principle is applicable so that in acid solution
the Cr 2 0 7 2 " ion predominates and the solution color is orange; in basic solu-
tion the color is yellow. No oxidation-reduction is involved in the chromate-
diehromate eq uilib rium. When a solution of a dichromate is added to one
containing ions of the heavy metals the normal chromate is precipitated, not

566

Transition Elements— III: The Elements of Groups VIA and VI lA

the dichromate. No hydrogen salts, such as KliCrO.,, exist, but only normal
chromate salts, e.g., K 2 Cr0.j, and dichromate salts, e.g., K 2 Cr 2 0 7 . The Cr 2 0 7 2 ~
ion is analogous to the S 2 0 7 2 - ion and, akin to sulfur, in high concentrations
of a strong acid, more highly condensed poly-anions can be formed, e.g.,
Cr 3 Oio 3 ~. When concentrated H 2 S0 4 is added to a dichromate or chromate, the
end product of ,uch condensation, chromium (VI) oxide, CrO.,, or chromic
anhydride, separates out as red needles. This is a powerful oxidizing agent
and, in concentrated H 2 S0 4 , is known as “cleaning solution” for glass appara-
tus because, of its ability to oxidize grease. Acidified solutions of Cr 2 0 7 2 -
are powerful oxidizing agents, capable of oxidizing all but the weakest
reducing agents.

(10) 2 Cr 8+ + 7 H 2 0 ^ Cr 2 O r a - + 14 H+ + 6e E° = -1.33 volt

Such solutions are used in titrations in volumetric analysis for the determina-
tion of iron by the oxidation of Fe 2+ to Fe 3 +

(11) 6 Fe 2 + + Cr 2 0 7 s - + 14 H+ ^ 6 Fe 3 + + 2 Cr 3 * -f 7 H a O

The oxidation potentials for chromium are summarized below.

(A) In acid solution

0.91 v 0.41 v -1.33 v

Cr *-Cr 2 + ►•Cr 1+ »^Cr 2 0 7 2 -

(B) In basic solution

0.4 v

Gr *-Cr(OH) 2

I 1.2 v -

1.1 v -0.13 v

*-Cr(OH) 4 - *-Cr0 4 2 -

4. Molybdenum and Tungsten. The mineral molybdenite , MoS 2? has an
appearance similar to that of graphite and, up to the middle of the eighteenth
century, was considered identical with it. In 1778 K. W. Scheele showed
that when molybdenite is treated with HNO a it leaves a white residue which
has acidic properties. He called the residue molybdic acid (Greek, motybdos :
lead) and proved that molybdenite was a sulfide of an element, molybdenum.
In 1781, Scheele also showed that the mineral, then called tungsten (Swedish,
heavy stone), and thought to be an ore of tin, contained an acid which
he named tungstic acid, and that the mineral wolframite, (Fe, Mn)W0 4 ,
contained the same acid. The elements were isolated later, molybdenum in
1790 by P. J. Hjelm and tungsten in 1793 by the d’Elhujar brothers in Spain.
Tungsten is also known as ux)lfram (German, wolf-soot), the name from
which its chemical symbol is derived.

Transition Elements — III: The Elements of Groups VIA and VII A 567

Both molybdenum and tungsten metals can be produced by the reduction
of their oxides, M0O3 and W 0 3 , with hydrogen or carbon in an electric
furnace at about 1000 °C. Though apparently simple chemically, the metal-
lurgy is complicated by the extremely high melting points of the metals so
that they are produced as powders. Tungsten has the highest melting point
of any metal but, by heating compressed powder to incipient fusion by
means of an electric current through it, coherent rods of the metal can
be formed. These rods can be hammered, rolled, or drawn into wires with
diameters as small as 0.01 mm. The tensile strength of tungsten wire exceeds
that of any other metallic substance. Inasmuch as the predominant use of
these metals is as an additive to steel they need not be prepared in the
pure state, but, like chromium, can be used as iron alloys, ferromolybdenum
and ferrotungsten , which can be prepared directly by metallothermic methods.
The metallurgical desirability of these metals is based upon the extreme
hardness and wear resistance of their alloys even at elevated temperatures;
hence they are employed in high speed machine tools which retain their
temper even at red heat. Because of its high melting point and extremely
low vapor pressure, tungsten is used extensively for filaments in electric
incandescent lamps and electron tubes. Although carbon filaments cannot
be heated above the temperature at which a yellow’ glow r is obtained without
vaporizing, tungsten has practically no vapor pressure even at white heat.
Since the light emitted by an incandescent object is proportional to its
temperature, it is possible to convert a greater .portion of the electrical energy
into light by using tungsten instead of carbon filaments.

5. The Elements of Group VIJA. The elements of this group are man-
ganese, technetium, and rhenium. The electron configurations are given
below and the properties of these elements, so far as they are known, are
listed in Table 42-C,

Element

Atomic

Number

2s 2 p

3 s 3 p 3 d

4s 4 p 4d 4f

5s 5 p 5 d

Manganese

2 6

2 6 5

Technetium

2 6

2 6 10

2 6 5

Rhenium
;

2 6

2 6 10

2 6 10 14

2 6 5

6. Manganese. A relatively abundant element, manganese ranks ninth
among the metals with an estimated abundance of 0*01%. Compounds of
manganese were known to the ancient Egyptians and Romans and were
employed for bleaching glass under the name of “magnes ” Until the eighteenth
century these were considered to be compounds of iron but in 1774, during
his investigation and discovery of chlorine, Scheele was the first to indicate
that the mineral pijrohmte , MnO a , contained a new element. It was named
manganese (Latin, magnes; magnet) to distinguish it from magnesium. At

568

Transition Elements— III: The Elements of Groups VIA and VI1A

Table 42-C

Properties of the Elements of Group VIIA

Property

Manganese

Technetium*

Rhenium

Symbol

Atomic Number

Atomic Weight

54.9381

186.2

Isotopes (mass numbers
and percent)

55(100)

185(37.07)

187(62.93)*

Abundance in Earth’s

Crust, %

0.1

none

10-*

Physical State at STP

gray-white

solid

Density at STP, g/cm 3

7.44

11.5

20.5

Melting Point, °C

1245

2125

3180

Boiling Point, °C

2040

5630

Heat of Fusion, keal/mole

3.5

5.5

7.9

Heat of Atomization,
kcal/mole

66.7

155

186

Ionization Potential, eV, 1st
2nd

7.43

15.7

7.87

Atomic Radius, A

1.17

1.28

Ionic Radius, A (+7)

0.47

Oxidation States

-h2,-f 3, +4
+6 , +7

+2,+4,+7

-1

4-3,+4,4-6,+7

Oxidation Potential, volt
for M -» M 2+ 4- 2 e

1.18

*A radioactive species.

least one hundred minerals contain manganese but those of greatest economic
importance are the oxides; the most important ore is the dioxide, pyrolustte,
Mn0 2 ; others are braunite, Mn 2 O a , and hausmannite , Mn 3 0 4 .

The pure element cannot be prepared by reduction with carbon inasmuch
as manganese combines with carbon at high temperatures to form the carbide,
Mn 8 C. Almost pure manganese can be obtained by reduction of one of its
oxides with aluminum while the pure element can be prepared by the
electrolysis of a solution of Mn 2+ ion using a mercury cathode. Manganese
is a gray-white metal tinged faintly pink, brittle but harder than iron. It is
an active metal; it tarnishes in moist air, displaces H 2 from strong acids, and
combines with C, N 2 , P, and S at high temperatures. The principal use of

Transition Elements — III: The Elements of Groups VIA and V1IA

the element is in the manufacture of steel, primarily to remove sulfur, the
main cause of brittleness in steel. For this it is unnecessary to use pure man-
ganese so that alloys of iron and manganese such as ferromanganese (70-
80% Mn ) and spiegeleisen ( 15-20% Mn ) , are more generally employed. These
cpn be made by reducing the mixed oxides of manganese and iron with carbon
in an ordinary blast furnace or in an electric furnace. Added in low concentra-
tions, manganese acts as a scavenger to remove sulfur and, to a lesser extent,
oxygen by reducing any iron oxide or iron sulfide present. The resulting
sulfides and oxides of manganese are but slightly soluble in molten iron
and are removed in the slag. In higher concentrations, up to about 14%, man-
ganese imparts special toughness and hardness to steel.

7. Oxidation States of Manganese. A summary of the oxidation states
of manganese is given in Table 42-D.

Table 42-D

Oxidation States of Manganese

Oxidation

State

Ionic

Species

Oxide

Hydroodde

Character

Derivative

Color

Mn 2+

MnO

Mn(OH) 2

Basic

MnClj

pink

Mn®+

Mn 2 O s

Mn(OH) s

Weakly basic

MnCl 3

violet

Mn*+

Mn0 2

Amphoteric

MnCl 4

brown

Mn0 3 2 -

(H 2 MnO s )*

CaMnOg

brown

MnO* 2-

(Mn0 3 )

(H 2 MnOJ

Weakly acidic

K 2 Mn0 4

green

MnOr

Mn 2 0 7

HMnO*
1

Strongly acidic

KMnO*

purple

♦Compounds whose formulas are inclosed in parentheses are hypothetical and have
not been prepared.

Although they are stable in solid compounds, not all of the oxidation
states are stable in solution particularly at elevated temperatures. Thus,
while MnCl 3 and K 2 MnO* are stable in dry air they disproportionate and
undergo oxidation-reduction in solution. Manganese (II) compounds can be
prepared by reduction of a higher oxide or by reaction of Mn0 2 and HC1;
MnCl* is unstable and decomposes into MnCl 2 and Cl 2 .

(12) MnO a + H 2 MnO + H a O

(13) MnO a + 4 HC1 -» MnCl 2 + Cl 2 + 2 H 2 0

Addition of NH S or an alkali metal hydroxide to a solution of Mn 2 + ion
yields a pinkish-white precipitate of Mn(OH) 2 , In the presence of a high
concentration of NH* + ion, this is soluble so that it will not be precipated
by NH 3 in such an environment. The hydrated Mn 2 + ion is one of the few
pink ions.

570 Transition Elements— III: The Elements of Groups VIA and VII A

Simple manganese (III) ions are unstable in aqueous solution and dis-
proportionate.

(14) 2 Mn 3 + + 2 H 2 0 Mn 2+ + MnO,. + 4H f

The +3 state is stable only in compounds which yield but a slight concen-
tratipn of Mn 3+ ion, for example, the insoluble Mn s O a and the slightly dis-
sociated complex ions, Mn(CN) 6 3 “ or MnCl-~".

The Mn 4+ ion is unknown in aqueous solution. Manganese(IV) oxide,
Mn0 2 is a difficultly soluble brown-black solid and is amphoteric. It may
be regarded as the anhydride of the hypothetical manganous acid, H g MnO s ,
a salt, of which, calcium manganite, CaMnO a , has been prepared. In acid
solution, Mn0 2 is a fairly strong oxidizing agent.

(15) Mn 2+ + 2H 2 0^ MnO,<*) + 4 H+ + 2 e E° = -1.23 volt

An important industrial chemical, the uses of MnO* are derived from its
oxidizing power. About 30,000 tons are used annually in the manufacture
of dry cells, in which its purpose is as an oxidant to prevent the formation
of free H 2 gas which would otherwise polarize the carbon cathode and
render the cell inoperable. It is also used in the chemical industry as an
oxidizing agent in the manufacture of hydroquinone, a photographic developer,
and of potassium permanganate. Its classical use to decolorize glass is due to
its reduction and the formation of a small concentration of red manganese (III)
ions in the melt. The red color is complementary to the green tint produced
by iron (II) impurities, resulting in a glass which is apparently colorless. In
current practice, selenium is now used almost universally as. a glass decolorizer.

When Mn0 2 is fused with K 2 C0 3 , or with KOH and an oxidizing agent
which can supply oxygen, there is formed bright green potassium manganate,
KsMnQ 4 .

(16) 2 Mn0 2 + 2 K 2 C0 3 + 0 2 -» 2 K 2 MnQ 4 + 2 CO*

In neutral or in acid solution, manganate ion, Mn0 4 2 “ disproportionates and
so can be preserved only in the presence of alkali. If the concentration
of the qjH- ion is decreased by the addition of even a weak acid, such as
acetic acid or carbonic acid, Mn0 2 and the permanganate ion, Mn0 4 “, are
produced.

(17) 3 MnO*^ + 4H+-> MnOr + 2 Mn0 2 + 2 H z O

For this reason, manganic acid, H 2 Mn0 4 , is unknown. Unlike the
Cr0 4 2- 1 Cr 2 0 7 2 “ equilibrium, the Mn0 4 2 ~ | Mn0 4 ~ equilibrium involves an
oxidation-reduction.

Potassium permanganate, KMn0 4 , is the most important of the per-
manganates. It has the deep purple color characteristic of the MnOr ion,
apparent even in extremely dilute solution. The salt is prepared commercially
by the oxidation- of K 2 Mn0 4 with Cl 2 in alkaline solution.

(18) 2 MnO* 2 ~ + Cl* 2 MnO,- + 2 Ck

Transition Elements — III: The Elements oj Groups VIA and VII A

571

The MnOr ion is a powerful oxidizing agent, particularly in acid solution,
but tbe nature of its oxidizing action differs in acid and in base. In acid, the
MnOr ion is reduced to the Mn- + state while in basic solution it is reduced
to tbe intermediate +4 state and precipitates as Mn0 2 .

(19) Mn 2+ 4 H 2 0 MnOr 8 H + -f— 5 e E° = -1.51 volt (in acid)

(20) Mn0 2 + 4 OH- MnOr + 2 H s O -f 3 e E° = -0.59 volt (in base)

So strong an oxidizing agent is the MnOr ion in acid solution that only the
weakest reducing agents fail to be oxidized by it. Conversely, only extremely
powerful oxidizing agents are capable of oxidizing Mn 2+ to MnOr. Examples
of the oxidizing action of MnOr ion are its oxidation of sulfurous acid,
H 2 SO s , and of sodium sulfite, Na 2 S0 3 , in basic solution.

(21) 2 MnOr + 5 H,SO : , .+ H,.0 -> 2 Mn 2 + + 5 SO, 2 - + 4H S 0+

(22) 2 MnOr + 3 S0 8 2 - -f H..O 2 MnO, + 5 SO, 2 - + 2. OH-

The oxidation of Fe a+ in acid solution by MnO,- is an important reaction
in analytical chemistry.

(23) 5 Fe-+ + MnOr + 8 H+ -* 5 Fe 3+ + Mn 2+ + 4 H 2 0

The oxidation potentials for the several states of manganese are sum-
marized below,

(A) In acid solution

1.18 v -1.51 v -0.95 v -2.26 -0.56

Mn ► Mn 2 + ► Mn 3 + ► MnO-, ► MnO , 2 " ► MnO, -

--1.23 v-

-1.70 v-

(B) In basic solution

1.57 v -0.1 v -0.2 -0.60 -0.56

Mn *-Mn (OH ) ►Mn ( OH ) , ►MnO, ►MnO , s ~ ►MnOr

-0.05-

--0,59-

In analytical chemistry, KMn0 4 is the most widely used oxidizing agent
for the quantitative analysis of substances having reducing properties. Among
such procedures are the direct .oxidation of simple ions such as Fe 2+ , Mn 2+ ,
Cu + , Ti® + , Mo 3 " 1 ", U 4+ , and the acids HN0 3 , H 2 S0 3 , H 2 S, and HgO* In
titrations the color of the MnOr ion acts as its own indicator for the com-
pleteness of reaction* For example, in the analysis of Fe 24- ion a standard
solution of KMnO* is added from a buret. While excess Fe 2+ is still present,
the MnOr is reduced to Mn 2+ and the purple color of the MnOr ion is

572

Transition Elements— III: The Elements of Groups VIA and VllA

not evident. When all the Fe 2 f has been oxidized to the Fe 3+ state, the next
drop of KMn0 4 solution imparts a pink color to the reaction mixture. This
is the end point of the titration; no special indicator need be added. The
relation, V x X Ni = V 2 X N 2 , is then applicable, or

Vp e 2+ X Nr e 2!+ = V MnO*“ X NMn 04 "'

Despite its many oxidation states, manganese forms few complex ions;
among these are Mn(CN) c 4 ', Mn(CN) 6 3 ~> and MnCl 4 2 ".

8. Technetium and Rhenium, Technetium (Greek, technetos: artificial)
is the first element prepared by man and it was named in recognition of
this achievement. The element is radioactive and occurs in nature, if at all,
only in trace amounts. It was first isolated in 1937 by C. Perrier and E.
Segrd in Italy from a sample of molybdenum which had been bombarded
by deuterons in the cyclotron at Berkeley, California. The element can be
prepared by reduction of TcO,f ion electrolytically or with iron or copper
but at present it has no- commercial value and is an academic curiosity.

Rhenium (Latin, Rhenus : the Rhine) is an extremely scarce element
with an occurrence of but one part per billion in the earth's crust. It was
discovered in 1925 by the German chemists, W. Noddack, Ida Tacke, and

O. Berg who, four years later, isolated one gram of the metal from 66 0 kg
of Norwegian molybdenite. Rhenium powder can be prepared by the electro-
lytic reduction of potassium peirhenate, KRe0 4 , or by hydrogen at about
1000°C. Of all the metals, rhenium has the second highest melting point
and is the fourth most dense. Some rhenium is used in thermocouples and
in thermionic filaments but most of the metal -*($700 per pound) is being
used for research and to develop commercial applications.

QUESTIONS

1. How are the metals, chromium, tungsten, and molybdenum, prepared?

2. In what respects does chromium resemble the elements of Groups VA and
VII A? Why should Cr 2+ and Cr 3+ resemble the corresponding oxidation
states of vanadium and manganese?

3. By what chemical means can an insoluble transition metal oxide, such as
Cr 2 0 3 , be brought into aqueous solution?

4. Account for the fact that aqueous solutions of chromium (III) ions exhibit
different colors.

5. Write complete equations for the following: (a) conversion of chromite to
K 2 Cr 2 0 7 (b) heating ammonium dichromate (c) solution of solid Cr(OH) s
in acid and in base (d) lead nitrate plus potassium chromate (e) hydrolysis
of Cr(H 2 0) e .

6. Balance by the ion-electron method: (a) oxidation of SO s 2 ~ by Cr 2 0 7 ^ in
acid solution (b) oxidation of iron (II) sulfate by Mn0 4 ~ in acid solution
(c) oxidation of Na 2 S0 3 by KMnO*.

7. Starting with MnQ 2 , write equations for the preparation of (a) KMn0 4
(b) KjjMnO* (c) MnCl 2 .

8. Illustrate what is meant by the term "disproportionate.*'

Transition Elements — 111: The Elements of Groups VIA and VUA

573

9. Which is the stronger oxidizing agent in acid solution, Cr 2 0; 2 * ion or MnO*~
ion? Explain briefly.

IQ. On the basis of their oxidation potentials, can Cr 2 + be formed from Cr
and Cr*+?

11. Why cannot the Mn0 4 2 ~ ion be prepared by reduction of the Mn0 4 ~ ion?

12. Why is the Mn0 4 2 ~ ion unstable in acid solution but stable in basic solution?

13. Will atmospheric oxygen oxidize Mn 2 + to Mn0 2 in acid solution?

14. From a violet aqueous solution of chromium (III) chloride, AgN0 3 precipitates
all of the “chlorine” present but from a green solution of chromium (III)
chloride only one third of the “chlorine” can be precipitated as AgCl. Ex-
plain, writing ionic and molecular formulas of pertinent species.

15. Why does not an alkali metal carbonate precipitate chromium carbonate upon
addition to a solution of C t H+ ions, but instead precipitates Cr(OH) a ; K sp for
Cr(OH) 3 is 6.7 X 10~ 31 .

16. Calculate the ratio of the concentrations of the Cr 2 0 7 2 ~ ion to the Cr0 4 2 ~ ion
in a solution whose pH is 7.0.

17. Calculate the standard potential of a voltaic cell wherein Fe 2H ~ is oxidized
by Cr 2 0* 2 -.

18. Why is KMn0 4 widely used in quantitative analysis?

19. A one gram sample of an iron alloy is dissolved in H 2 S0 4 , forming Fe 2+

ions. If 50.0 ml of 0.100N KMn0 4 are required to titrate the Fe 2 +, what
is the percent of iron in the alloy? Ans: 27.9%

20. What volume of nitrogen measured at 20 °C and 750 mm pressure can be

obtained by heating 15 g of ammonium dichromate? Ans: 1.4 liters

21. What quantity of K 2 Cr 2 0 7 can be prepared from one ton of chromite ore

containing 20% Cr 2 O a ? Ans : 775 lb

22. What weight of KMn0 4 is required to make 250 ml of 0.40N solution to be

used as an oxidizing agent in acid solution? Ans : 3.16 g

23. What volume of 0.10 N K 2 Cr 2 0 7 solution is required to oxidize 0.25 g of

Fe 2+ in acid solution? Ans: 45 ml

Transition Elements— IV
The Elements of Croup IB

The elements of Group IB include copper, silver, and gold, collectively
known as the Copper Family. Their electron configurations are shown below
and their properties are listed in Table 43- A.

Element

Atomic

Number

2$ 2 p

3s 3p 3 d

4s 4p Ad 4f

5s op Bd

Copper

2 6

2 6 10

Silver

2 6

2 6 10

Gold

2 6

2 6 10

1. General Properties of the Copper Family. With these elements, the
3 d, 4 d, and 5 d orbitals, which were in the process of being filled in the
three transition series, are complete. In the outermost principal quantum
level of the Group IB elements there is a single s electron, ns 1 , but despite
this apparent similarity in electron configuration to the elements of Group IA
there is a decided contrast in properties between the elements of the copper
family and the alkali metals.

The alkali metals are very soft, have very low densities, and melt at
rather low temperatures whereas Cu, Ag, and Au are much harder, have
much higher densities, and melt around 1000° C. Chemically, the alkali metals
are extremely reactive. They have but a single oxidation state, form only
ionic compounds, and no complex ions. In general their compounds are
soluble and their oxides are strongly basic. On the other hand, Cu, Ag, and
Au are relatively unreactive. They show more than one oxidation state, form
both ionic and covalent compounds, and numerous complex ions. Their com-
pounds are generally insoluble and only silver oxide is a strong base. In
the main this dissimilarity in properties can be traced to the smaller atomic
and ionic sizes of the Group IB elements and their having available d electrons

574

Transition Elements — IV: The Elements of Group IB

575

for bond formation. In the next to the outermost shell the Group I A elements
have eight electrons, none of which are d electrons, whereas the Group IB
elements have in the corresponding shell eighteen electrons, of which ten are
d electrons.

The elements of the copper family are known as the coinage metals.
Because of their rarity, especially pleasing luster, and chemical inactivity,
silver and gold in particular have been prized from prehistoric times as

Table 43-A

Properties of the Elements of Group IB

Property

Copper

Silver

Gold

Symbol

Atomic Number

Atomic Weight

63.54

107.870

196.967

Isotopes (mass numbers
and percent)

63(69.09)

64(30.91)

107(51.35)

108(48.65)

197(100)

Abundance in Earth s

0.0001

5 x 10 -s

io-»

Crust, %

Physical State at STP

red

solid

silver

solid

yellow

solid

Density at STP, g/cm 3

8.92

10.50

19.3

Melting Point, °C

1083

960.5

1063

Boiling Point, °C

2325

1950

2600

Heat of Fusion, kcai/mole

3.11

2.70

3.03

Heat of Vaporization,

72.8

61.0

77.5

kcal/mole

Heat of Atomization,

81.1

68.4

84.7

kcal/mole

Ionization Potential, eV, 1st

7.68

7.54

9.18

2nd

20.34

21.40

19.95

3rd

29.5

35.9

29.7

Electronegativity

1.9

2.4

Atomic Radius, A

1.27

1.44

1.40

Ionic Radius, A (+1)

0.96

1.13

1.37

Oxidation States

+l,+2

4*l,*f2

+1,4-3

Coordination Numbers

2,4 (+1)

4,6 (+2)

2(4-1)

4(4-2)

2,4 (+1)
4(+3)

Oxidation Potential, volt

-0.521

-0.799

-1.68

for M M+ -4-

576

Transition Elements — IV: The Elements of Group IB

ornaments and as media of exchange. Because they occur in the free state
in nature, gold was probably the first metal to be used by man, and copper
the second, but copper was the first metal to be fashioned into utensils and
instruments of warfare. The discovery of copper in the late Stone Age ushered
in the Copper or Bronze Age. This period of history had its greatest develop-
ment in Egypt and records have been found of the working of copper mines
about 3800 b.c. on the Sinai Peninsula by the Egyptian Pharaoh, Seneferu.
Silver ornaments and utensils, dating as early as 3000 b.c., have been unearthed
in Asia Minor. Silver was adopted by the ancient Romans and by the first
Continental Congress in the United States as their monetary standards. In
1792 tlie dollar was defined by Congress as 412.5 grains of silver, 0.900 fine
(90% pure) and, except for minor variations concerning alloying elements,
that definition has not been changed. Pure gold and silver are too soft to
be used directly in coinage and so for this purpose they are alloyed with a
small percentage of copper. The alloy has increased hardness and improved
wearing qualities. United States coinage silver contains 90% Ag and 10%
Cu, while British sterling is 92.5% Ag and 7.5% Cu. Jewelry silver contains
80% Ag and 20% Cu. In gold alloys, the amount of gold is usually expressed
in carats ; pure gold is 24 carat. British gold coinage is 22 carat and American
coinage is 21.6 carat. White gold, used in jewelry, is an alloy of Au and Ni.

Copper

Copper (Latin, cuprum : Cyprus, where it has been mined since 2500 b.c.)
is found in both the elemental and combined forms. Though widely dis-
tributed, native copper generally occurs in small deposits but masses weigh-
ing several tons have been found in the large native copper deposits of the
Lake Superior region. Numerous compounds of copper are found in nature
but the principal ores are the sulfides and the oxides. Among these are
chalcopyrites, CuFeS 2 ; bomite , Cu 3 FeS 4 ; chalcolite , Cu 2 S; cuprite , Cu 2 0;
malachite, CuCO s * Cu(OH) 2 ; and azurite, 2 CuCOs * Cu(OH) 2 .

2. Metallurgy of Copper, The metallurgy of copper varies with the
nature of the individual ore. When the ore contains native copper, it is
necessary merely to separate the metal by grinding the ore, washing the
metal free from the gangue, and melting the powdered copper with a suitable
flux. If the ore contains an oxide or carbonate which is soluble in acid, the
crushed ore is leached with dilute H 2 S0 4 and the metal is then obtained
from the solution by electrolysis.

Over 85% of current copper production, however, is derived from sulfide
ores. These usually contain less than 10% Cu, together with small quantities
of Ag, Au, the platinum metals, As, Sb, Se, and Te. The extraction of copper
from a sulfide ore involves several laborious steps, primarily because the
ore contains a large amount of FeS which must be converted to FeO and
then removed as an FeSiO s slag. Generally the steps in the processing of the
ore are:

(A) Concentration: The flotation process is now used almost exclusively.

(B) Roasting: Much of the sulfur is removed as SO*, later used for the

Transition Elements— TV: The Elements of Group IB

577

manufacture of H 2 S0 4 , as are other volatile components such as As and Sb;
some iron is converted to FeO.

(C) Smelting : This is carried out in a reverberatory furnace. At smelting
temperatures and in a neutral or reducing atmosphere, Cu 2 S is formed
while excess sulfur combines with iron to form FeS. In the liquid state,
Cu 2 S and FeS are miscible in all proportions; this mixture, whether liquid
or solid, is known as copper matte .

(D) Converting : The matte, with added sand, is transferred to a converter,
similiar to a Bessemer converter, and thin streams of air are blown through
the molten mass. The Cu 2 S thereby forms Cu metal.

(1) 2 Cu 2 S + S0 2 2 Cu 2 0 + 2 S0 2

(2) 2 Cu 2 0 + Cu 2 S -> 6 Cu + S0 2

The Cu is cast as large slabs and, as it solidifies, dissolved S0 2 is expelled,
giving the surface of the metal a blistered appearance and hence its name
“blister copper.”

(E) Refining : Blister copper is 99% pure, For the major use of copper
as an electrical conductor, blister copper is insufficiently pure. Even a small
amount of impurity greatly increases the electrical resistance of copper so
that further refining is necessary. Initial refining is done by blowing com-
pressed air through the remelted metal to oxidize any remaining impurities,
followed by stirring with poles of green wood. Reducing gases, such as
CO and H 2 , released by destructive distillation of the wood, reduce any Cu 2 0
formed. The “poled” Cu is cast into slabs about 3 feet square and about
3/4 inch thick and final refining is done by electrolysis. The slabs of par-
tially refined Cu are made the anodes in an electrolytic bath of CuS0 4
acidified with HUSO*, while the cathodes are thin sheets of pure Cu suspended
alternately between the anodes. During the passage of the electric current,
Cu 2 + ions from the electrolyte are discharged at the cathode so that it is
gradually built up of 100% Cu. At the same time an equivalent quantity of
Cu dissolves from the anode forming Cu 2+ so that, in effect, there is a
transfer of Cu from the impure anode to the pure cathode. Impurities in
the copper may be viewed as two types, those which are chemically more
active than Cu and those which are less active. By proper regulation of
the applied potential, metal impurities which are more active go into solution
and remain there as ions while the less active impurities do not dissolve
but are deposited as a sludge which accumulates at the bottom of the cell.
From this sludge are recovered Ag, Au, Pt, and Pd, sometimes in quantity
sufficient to pay for the refining process. Electrolytically refined Cu is
99.95% pure.

3. Uses of Copper. Second only to iron, copper is the most important
industrial metal. The United States is the most important copper-producing
country, producing over one million tons annually or more than 25% of the
world s production; next in order are Chile and the U.S.S.R. Many by-products
or co-products are also recovered during copper processing. Among these
are 40% of the 40 million pounds of molybdenum produced annually; 600,000

578

Transition Elements — TJ: The Elements of Group fg

pounds, or almost all, of the selenium produced; 30% of all gold and 28%
of all silver produced; and half a million tons of H 2 S0 4 .

Among the less expensive metals, copper is the best conductor of elec-
tricity so that approximately half of all the copper consumed is for electrical
applications. Virtually all underground and building wiring is made of copper,
though aluminum is now used for most overhead high voltage lines. Copper
is also used in photoengraving and in electrotyping. A wax or a. plaster cast
is made of each page of type and this is rubbed with graphite to make it a
conductor. The cast is then used as the cathode in an electrolytic bath of
CuSO*, the anode being a copper plate. Copper metal is deposited as a
film on the graphitized cast, reproducing every detail of the type or en-
graving. When the deposited film is sufficiently thick the copper plating
is stripped off and backed with lead. Because of its resistance to corrosion,
copper is also used for kettles, stills, plumbing goods, roofing, etc.

The remaining copper produced is used in the manufacture of alloys
which number over 1,000 different kinds but which can be classified into
four main types: brass, bronze, nickel-silver , and and cupro-nickeh Brass is
defined as any copper-base alloy containing zinc as the principal alloying
metal; it may contain smaller quantities of other elements such as Sn, Al,
Fe, Pb, Mn, and Ni. Ordinary bronze is an alloy of copper and tin but the
term bronze is seldom used alone so the simple Cu-Sn alloys have come
to be known as phosphor bronze because of the small residual phosphorus
content. The term bronze has been extended to copper-base alloys where
the principal alloying element is other than tin; thus aluminum bronze,
beryllium bronze, and silicon bronze. Formerly known as German silver,
nickel silver contains no silver but is an alloy of Cu, Ni, and Zn, Cupro-
nickels are essentially alloys of copper and nickel containing 10-30% Ni.

4. Properties of Copper, Copper is the only distinctly reddish-colored
metal. It is less active than hydrogen and so does not liberate H 2 from acids.
The oxidizing acids, HNO a and H 2 S0 4 , react with the metal to produce
NO or NO a and S0 2 , respectively, along with water.

(3) 3 Cu + 8 HNO s 3 Cu(N0 3 ) 2 + 2 NO + 4 H z O (with dilute HNO*;

concentrated acid
yields N0 2 )

(4) Cu + 2 H 2 S0 4 “> CuS0 4 + SO a + 2 H z O (with hot concen-

trated H 2 S0 4 )

At ordinary temperatures dry air has no effect on copper; at about 300°C,
CuO is formed and at I000 o C, Cu 2 0 is produced. In the presence of CO*
and atmospheric moisture, copper becomes covered with a green basic car-
bonate similar to malachite, which then acts as a protective coating.

(5) 2 Cu + 0 2 + C0 2 + H 2 0 -» Cu(OH) 2 • CuCO s

5. Compounds of Copper; Copper(l) Compounds . Copper forms two series
of compbunds, copper(I) or cuprous , in which the oxidation state is -}-l>
and copper (II) or cupric , in which the oxidation state is +2* In aqueous
solution the simple Cu + ion is unstable and disproportionates.

Transition Elements — IV; The Elements of Group IB

579

(6) 2 Cu+ Cu- + 4* Cu (2 CuCl CuCl 2 + Cu)

The +1 oxidation state is stable only in insoluble compounds such as the
copper(I) halides, Cu 2 0, Cu 2 S, or as a part of a complex ion, as CuCl 2 “,
Cu(CN) 3 ~, and Cu(NHj) 2 + . In solution the Cu + ion is colorless though solid
Cu 4 * compounds exhibit a variety of colors.

When copper(II) chloride, CuCl 2 , is boiled with Cu metal in the
presence of concentrated HCI the reverse of Equation 5 takes place and
CuCl is formed. This is soluble in concentrated HCI because of the forma-
tion of the dichlorocuprate complex ion, CuCl 2 ~.

(7) CuCl(s) -j" Cl” — > CuCl 2 “

If water is added to the solution, the Cl” ion concentration is decreased and
white CuCl precipitates.

A red solid, Cu a O can be prepared by adding NaOH to an acid solution
of CuCl. The simple hydroxide, Cu(OH) is unknown; when heated with basic
solutions of Cu 2+ salts, reducing agents precipitate Cu 2 0. This reaction is the
basis of the test for reducing sugars such as glucose.

Copper ( II) compounds ; The ordinary compounds of copper seen on the
laboratory bench are Cu 24 * compounds. The simple Cu 24 " ion is colorless
but aqueous solutions and hydrated salts are blue due to the presence
of the hydrated ions, Cu(H 2 0) t 2+ . Aqueous solutions are acid by hydrolysis.

(8) Cu{H 2 0) 4 2+ + H a O -» [Cu(H 2 0) 3 (0H)]+ + H 3 0+

Addition of OH - ion to Cu 2+ ion precipitates pale blue, gelatinous Cu(OH) 2 .
When heated, or when precipitated in hot solutions, the hydroxide loses
water to form black CuO. In excess NH ;i , Cu(OH) 2 dissolves by forming the
deep blue complex teiramminecopper(II) ion, Cu(NH 3 ) 4 2+ . This deep blue
color formed in NH# solution is characteristic and serves as a sensitive test
for Cu 24 *.

The most important copper salt is the sulfate. The anhydrous salt, CuSO*,
is white but, when crystallized from aqueous solution, blue crystals of the
pentahydrate, CuSO* * 5 H 2 0, sometimes known as blue vitriol , are formed.
If heated or if the pressure of water vapor with which it is in equilibrium
is reduced, the pentahydrate loses water in steps to form the trihydrate,
CuSO* • 3 H 2 0, the monohydrate, CuSO* * H a O, and ultimately the anhy-
drous salt Copper sulfate is used in the electrorefining of copper in the Daniell
cell, and as a fungicide. Copper is poisonous to the lower organisms so that
copper sulfate is added in small quantities, one or two parts per million, to
prevent the growth of. algae in reservoirs and swimming pools.

The halides, CuCI 2 and CuBr 2 , can be formed by direct combination of
the elements but Cul 2 does not exist. Concentrated aqueous solutions of
CuCl 2 are green due to the presence of a mixture of yellow CuCl* 2 ” and
blue Cu(H 2 0) 4 24 * ions. Dilution yields a blue solution due In the replace-
ment of ehloro groups by H 2 0 molecules. When a solution of I” ion is added
to one containing Cu 24 * ions, white CuJ is precipitated and I 2 is liberated.

(9) 2 Cu 24 * 4 - 4 f -> 2 Col(f) 4” I 2

580

Transition Elements — TV: The Elements of Group IB

This reaction is used for the quantitative determination of Cu 2+ ion. The
I 2 formed, which is proportional to the concentration of the Cu 2+ ion, is
determined by titration with a standard solution of Na 2 S 2 0 3 .

•Copper (II) sulfide, CuS, is extremely insoluble and can be precipitated
by H 2 S even in the presence of a high H+ ion concentration. It can be
dissolved, however, by the oxidizing action of HNOs.

(10) 3 CuS -J- 8 H"*" -)- 2 NOs" — * 3 Cu 2+ — j— 3 S — j— 2 NO -(- 4 H 2 0

In acid solution the oxidation potentials for copper are:

-0.52 v -0.15 v

Cu *-Cu+ Cu 2+

1 0.34 v 1

Silver and Gold

6. Preparation of Silver. The name silver is derived from the Anglo-Saxon
soelfor and the symbol Ag from the Latin argentum. The principal silver
ores are native silver, sulfides such as argentite, Ag 2 S, and horn silver, AgCL
The element is found associated with other metals, chiefly Cu, Pb, and Zn.
In the processing of its ores the chemical unreactivity of silver poses the
greatest problem because of the difficulty of getting it into solution. Ores of
native Ag and AgCl are leached with Hg, in which the Ag dissolves to form
an amalgam; AgCl reacts with Hg because Ag is displaced by the more
active Hg. The liquid amalgam is separated from the waste rock and then
distilled in iron retorts. The volatile Hg boils off, is condensed and reused,
while the Ag remains in the retort. Because of its inertness, Ag is oxidized
with difficulty but, in the presence of CN~ ion, air oxidation proceeds due
to the formation of the Ag(CN) 2 " complex ion. Air is blown through a
suspension of pulverized ore in NaCN solution for about two weeks, after
which the Ag is precipitated by the addition of Zn metal.

(11) 4 Ag + 8 CN“ + 0 2 + 2 H 2 0 4 Ag(CN) 2 ~ + 4 OH- (solution of Ag)

(12) 2 Ag(CN) 2 ~ + Zn 2 Ag + Zn(CN) 4 2 ’ (reprecipitation of Ag)

The recovery of Ag during the electrorefining of Cu has been noted.
Much of the Pb obtained from PbS ores also contains appreciable amounts
of Ag. Molten Zn and Pb are immiscible and Ag is about 3000 times more
soluble in Zn than it is in Pb. The “purified” Pb (containing Ag) is melted
and thoroughly mixed with a small quantity, 1-2%, of Zn. Most of the Ag
dissolves preferentially in the Zn and, when the mixing is stopped, the
less dense Zn floats to the top carrying with it the Ag. This Zn-Ag alloy
is skimmed off and the Zn is removed by distillation. This extraction of Ag is
known as the Parkes Process. Mexico is the chief silver producing country,
the United States second. Annual world production is about -8000 tons of
which the U.S. supplies about 16%.

Transition Elements — IV: The Elements of Group IB

581

Silver is the best known conductor of heat and electricity but its price
precludes its extensive use in place of copper or aluminum. Although silver
has been used for coinage for centuries and silver bullion now serves as a
base for paper currency, its principal use in recent years has been industrial.
Among industrial uses are the manufacture of photographic materials, sterling
tableware and electroplate, and silver solders and brazing alloys. Lesser uses
are for dental alloys, mirrors, and high efficiency Ag-Zn batteries for space
craft. Cheaper metals are often silverplated. The article to be plated is
made the cathode in a bath of KAg(CN) 2 while the anode is a bar of pure
Ag. The Ag(CN) 2 ~ ion yields a slight concentration of Ag + ion which plates
out on the cathode while simultaneously the Ag anode dissolves.

Cathode: Ag+ + e Ag (Ag(CN) 2 - Ag+ + 2 CN-)

Anode: Ag + 2 CN~ -» Ag(CN) 2 “ + e

Through the use of the complex Ag(CN) 2 “ ion, the Ag is deposited more
evenly and more slowly than from a solution of AgN0 3 . The film of deposited
silver has a flat appearance but assumes a brilliant luster when burnished.
Mirrors are silvered by reducing an ammoniacal solution of AgNO s with
a weak reducing agent such as formaldehyde or glucose.

7, Compounds of Silver, The commonly met oxidation state of silver is
+1* Under extreme oxidizing conditions compounds of silver in the +2
and +3 states can be formed, e.g., AgO and Ag 2 0 3 , but these are rarely
encountered. The Ag* 1 * ion is colorless. Like Cu(OH), the compound Ag(OH)
is unknown. When a base is added to a solution of Ag + ions, brown Ag 2 0
is precipitated. Though this-is only slightly soluble an aqueous suspension is dis-
tinctly basic. Heating readily decomposes Ag a O into its elements. In the labora-
tory the most common silver compound is the nitrate, AgNO ;j , a colorless very
soluble salt. It is prepared commercially by the reaction of Ag and HN0 3 ,
and is the starting material from which other silver compounds are made.
An aqueous solution is neutral, further indication that Ag 2 0 is a strong
base. Thin sticks of AgNO s , also known as lunar caustic , are used as an
antiseptic. Organic materials, such as the skin or cloth, reduce AgN0 3 to
Ag, which is deposited as a black stain.

Silver salts show a wide range of “insolubilities.” Of the common simple
compounds only AgNO s and AgF are appreciably soluble. The other silver
halides form as curdy precipitates when solutions of Ag + ion and a halide
ion are mixed; AgCl is white, AgBr pale yellow, and Agl yellow. The com-
pounds AgaO, AgCl, and AgBr, but neither Agl nor Ag 2 S, are soluble in
NH S due to the formation of the Ag(NH 3 ) 2 + ion but all dissolve in solutions
of thiosulfate ion, S 2 O s 2 ", by forming the Ag(S 2 0 3 ) 2 3 “ ion, and in solutions
of CN“ ion with the formation of the Ag ( CN ) 2 “ ion. It follows therefore
that the concentration of the Ag 4 * ion in equilibrium with the complex ion,
Ag(CN)jf, must be smaller than the concentration of Ag + ion in the saturated
solution of any of these insoluble silver salts.

8. Photography. Except for AgF the silver halides are sensitive to light
When exposed to light freshly precipitated halides turn purple and eventually
black because the light decomposes the silver halide into Ag and the halogen

582

Transition Elements — TV: The Elements of Group IB

element. Photographic film is made by coating celluloid with a dispersion
of colloidal AgBr in gelatin. When light having a certain minimum energy
is incident upon the film, it causes either the reduction of AgBr to Ag or
it causes an incipient reduction so that Ag + ions struck by light will' be
preferentially reduced by a mild reducing agent. The exact nature of this
process is not completely understood.

The steps in the complete photographic process are (A) exposure ; (B) de-
veloping; (C) fixing; and (D) printing.

(A) , Exposure: The film is exposed to the image of the object as formed
by the camera lens. By exposure is meant the total quantity of light which
strikes the film. It is the product of two factors: intensity of the light source
and the time of exposure. In a camera the intensity is controlled by the
iris diaphragm and the time by the shutter speed. The incident radiation
must have an energy sufficient to decompose the AgBr. Blue light, and
radiation of higher frequency, have the requisite energy; red light does
not but the photographic emulsion can be sensitized to red light by mixing
certain organic dyes in the emulsion. The exposure, usually for a fraction of
a second, produces no visible effect on the film.

(B) Developing : The exposed film is then placed in a bath containing
a developer. A developer is an alkaline solution of a weak organic reducing
agent, such as metol, pyrogallol, or hydroquinone, which affects the AgBr
rapidly only in those spots where reduction has already been started by
the action of the light during the exposure. The film soon begins to show
black spots of Ag metal in places that were most strongly acted on by
light. The spots increase in area and gradually build up a coherent picture
in which the bright parts of the object are represented by dark areas and
vice versa. On account of this reversal the developed film is called a negative.
Development must be carried out in total darkness for panchromatic film
or in a faint red light for film not sensitive to red.

(C) Fixing : When sufficiently developed to the proper density or black-
ness of image, the negative is removed from the developer. In the white
areas of the negative there remains unreduced AgBr which must be removed
to prevent the blackening of the entire film on later exposure to light. This
is accomplished by immersing the negative in the fixer, a solution of Na 2 S 2 0 3
whereupon the AgBr forms the soluble Ag(S 2 0 3 ) 2 3 ~ ion, leaving the white areas
as transparent celluloid. The negative is then washed thoroughly in water to
remove all but the Ag metal image and the gelatin.

(D) Printing: The finished picture or print is made by passing light,
generally from an incandescent lamp source, through the negative on photo-
graphic paper coated with an emulsion of AgBr or AgCl. This produces on
the paper a latent reversal of the negative image. The paper is then developed
and fixed in the same manner as was the negative and a positive print of
the original subject is thereby produced.

Color photography is more complex. The film consists of three emulsion
layers, each containing a dye which is sensitive to one of the three primary
colors. Upon exposure and development, images are produced in each of

Transition Elements — IV: The Elements of Group IB

583

the three layers. As a final step the Ag is removed and the superposition
of the dyes forms the colored image.

9. Gold. The name of this element is its Anglo-Saxon name and its
symbol is derived from the Latin, aurum : shining dawn. Though very rare,
gold is probably the oldest metal known to man. It is found chiefly as
native gold, scattered through gravel as pbcer gold, or disseminated in
veins of quartz as vein gold. In the combined state, gold occurs as the
tellurides, AuTe a and AuAgTe 2 . Gold is always found associated with silver,
and the metallurgies of the two elements are similar— amalgamation and
cyanide leaching. The gold can be refined by electrolysis or the silver can
be removed by treatment with HNO a . Acid treatment is impossible if the
gold content exceeds 25% and, in this event, a suitable amount of silver
is fused with the alloy to decrease the gold content below 25%, a procedure
known as quartation. Most of the world’s production is absorbed by gov-
ernments to provide stability for, paper currency. The main nonmonetary
uses of gold are in jewelry and the decorative arts. Gold leaf is used for
gilding and lettering, while baser metals are frequently plated with gold.
Gold in colloidal dispersion is easily prepared; the color of the finest ruby
glass is due to such gold.

Gold is the most malleable and ductile of all metals and can be rolled
into leaf about 0.00001 mm thick. The metal is very unreactive. Not one of
the common acids reacts with Au and even F 2 does not attack it at ordinary
temperatures. Gold does react with selenic acid, fLSeO*, and also dissolves
in aqua regia through the formation of the complex AuC 1 4 ~ ion.

(13) Au + 3 NO*" + 4 Cl- + 3 H+ AuCV + 3 NO, + 3 H s O

Gold has two oxidation states, gold(I) or aurous and gold(III) or auric .
Like the Cu + ion tine Au 4 * ion is unstable in aqueous solution, forming
Au 8+ and Au, but it does form stable complex ions such as Au(CN) 2 ". Two
oxides are known, Au a O and Au 2 0 3 , and the corresponding hydroxides,
Au(OH) and Au(OH) 3 . The lower hydroxide is a weak base but Au(OH) s
is a weak acid, capable of reacting with a strong base to form an aurate
salt, e.g., NaAuO*. Because of its unreactivity, the simple compounds of gold
are unstable to heat, being decomposed into their elements.

QUESTIONS

1. For what chemical reasons were the elements Cu, Ag, and Au known before
iron to ancient man?

2. Account for the similarities and the differences in properties between the
elements of Groups IB and 1A.

3. In Group. I A the ionization potential decreases with increasing atomic number
whereas the reverse is true of the Group IB elements. Explain this difference.

4. Outline the metallurgy of copper, writing chemical equations where applicable,

5. Blister copper contains impurities of Zn and Au. How are these removed?

6. Write equations for (a) preparation of CuCl from Cu (b) dropwise addi-
tion of NH a to CuS0 4 (c) addition of NH S to AgBr(s) (d) addition of

584

Transition Elements — IV: The Elements of Group IB

Na 2 S 2 0 3 to Agl(s) (e) addition of KCN to Cu(OH) 2 ($) (f) solution of
Au in aqua regia (g) CuCl(s) and excess HCI (h) Cu and concentrated HNO*.

7. A solution contains AgN0 3 and Cu(N0 3 ) 2 . How would you separate and
identify the presence of the metals? Write pertinent equations.

8. Copper(II) hydroxide dissolves in NH 3 , HCI, and KCN solutions; H 2 S pre-
cipitates CuS from the NH S and HCI solutions but not from the KCN solution.
Arrange the following solutions according to increasing Cu 2+ ion concentra-
tion and give reasons for your listing: solutions of CuCl 2 , Cu(NH 3 ) 4 2 + ion,
Cu(CN) 4 2 ’ ion, and saturated solutions of Cu(OH) 2 and CuS.

9. (a) Why are Cu 2+ compounds more stable than Cu + compounds? (b) why
is the stable state of silver 4-1 whereas that for copper is 4-2?

10. Outline the metallurgy of silver and of gold. Write chemical equations where
applicable.

11. What principle is involved in the Parkes process for obtaining silver?

12. Outline the steps in the complete photographic process. Write chemical
equations where applicable. What reagent could be substituted for Na^Oa
as fixing agent?

13. What would be the effect if an exposed film were immersed first in the
fixing solution and then in the developer?

14. Which film should be more sensitive to longer wavelengths of visible radiation,
one composed of AgCl or one of AgBr? Explain briefly,

15. What concentration of NH S is theoretically required. to dissolve 0.010 mole
of each of the silver halides in one liter of solution?

Ans: AgCl, 0.22M; AgBr, 3.0M; Agl, 68M

16. What weight of AgBr can be dissolved in one liter of 0.50M NH 3 ?

17. What is the concentration of Ag+ ion in a saturated solution of H 2 S con-
taining 0.3 M HCI?

18. What is the solubility of AgCN(s) in a solution of pH equal to 4.00?

19. A 12.6 g sample of a copper ore is brought into solution, reacted with
excess KI, and titrated with 25.5 ml of 0.400M Na 2 S 2 0 3 . Calculate the
percent of Cu in the ore.

20. Silver is electroplated on a flat metal object whose surface area is 50 cm 2 .
If a current of 0.10 ampere flows for 2.0 hours at 6.0 volts, what will be the
thickness of the deposit, assuming it to be of uniform thickness?

Transition Elements— V
The Elements of Group HB

Zinc, cadmium, and mercury comprise the elements of Group IIB; their
electron configurations are given below and their properties are listed in
Table 44-A.

Element

Atomic

Number

2s 2 p

3? 3 p 3 d

4s 4p 4d 4 f

5s 5 p 5d

Zinc

2 6

2 6 10

Cadmium

2 6

2 6 10

Mercury

2 6

2 6 10

1. Genera! Properties of the Zinc Family, In these elements, both the
s orbital in the outermost shell and the inner d orbitals are complete so that
these elements at the end of the three transition series are not truly transition
elements but start anew the progression of regular elements. Nevertheless
the presence of d electrons in the atoms of the Group IIB elements is conducive
to certain transition-like properties and there is a distinct difference in
properties between these elements and the Group IIA elements in a man-
ner akin to the difference in properties between the Group IA and IB
elements. Thus the Group IIB elements have lower melting points and
boiling points, and greater densities than the adjacent elements of Group
IIA. Also they are less reactive, have higher ionization potentials, their hy-
droxides are less basic, and they form a number of complex ions. Zinc and
cadmium show only one oxidation state, +2, but mercury has, in addition,
a +1 state. The ions of the Group IIB elements are colorless and non-
paramagnetic,

Zinc

The name of the element is of unknown origin, the German term zink
being first used by Paracelsus. Because of its reactivity, elemental zinc is not

586

Transition Elements— V : The Elements of Group 11$

Table 44-A

Properties of the Elements of Group IIB

Property

Zinc

Cadmium

Mercury

Symbol

Atomic Number

Atomic Weight

65.37

112.40

200.59

Isotopes (mass numbers

64(48.87)

106 ( 1.22)

196( 0.16)

and percent)

66(27.62)

67 ( 4.12)
68(18.70)
70( 0.69)

108 ( 0.89)
110(12.43)
111(12.86)
112(23.79)
113(12.34)
114(28.81)

116 ( 7.66)

198(10.02)

199(16.92)

200(23.10)

201(13.22)

202(29.74)

204 ( 6.84)

Abundance in Earth’s

0.00004

10~ 7

10-7

Crust, %

Physical State at STP

blue-white'

solid

silver

solid

silver

liquid

Melting Point, °C

419.4

320.9

-38.9

Boiling Point, °C

906

767

357

Heat of Fusion, kcal/mole

1.77

1.46

0.56

Heat of Atomization,

31.2

26.8

14.7

kcal/mole

Ionization Potential, eV, 1st

9.36

8.96

10.38

2nd

17.89

16.84

18.65

3rd

40.0

38.0

31.9

Electronegativity .

1.6

1.7

1.9

Atomic Radius, A

1.25

1.41

1.44

Ionic Radius, A (+2)

0.74

0.97

1.10

Oxidation States

4-2

+1.+2

Coordination Numbers
Oxidation potential, volt

2(+l)

4(+2)

for M — * M 2+ 4- 2 e

4-0.763

4-0.403

-0.854

Hydration Energy of M 2 +

492

436

441

kcal/mole

found in nature. The chief ore is sphalerite or zinc blende, ZnS. Except for
a unique deposit of zincite , ZnO, in New Jersey, other widespread and abun-
dant minerals such as ZnC0 3 and Zn 2 (0H) 2 Si0 3 resulting from the oxida-
tion of sphalerite, have little significance as ores.

Transition Elements— V: The Elements of Croup 1IB

587

2. Metallurgy of Zinc. Almost all zinc metal is smelted from ZnS. The
ore is first concentrated by flotation, then roasted to the oxide, and reduced
to metal with carbon. At the temperature of the reduction, 1200°C, well above
the boiling point of zinc, the metal is produced as a vapor which is con-
densed to a liquid and then cast into ingots known as slab zinc or spelter .
If the vapor is condensed directly to the solid it forms a fine powder or
zinc dust. Zinc is also produced by an electrolytic process. The ore, ZnS,
is roasted under such conditions as to produce ZnS0 4 and then leached with
dilute H 2 S0 4 . The resulting solution of ZnS0 4 is treated with Zn dust to
remove less active metal impurities such as Cu 2+ , Cd 2+ , and Pb 2+ ions, for
if these were not removed they would subsequently deposit during electrolysis
along with the Zn on the cathode. Impurities of metals more active than
Zn are not plated out because the potential applied is sufficient only to re-
duce Zn 2+ ion. Though the electrolytic solution is acidic, H + ion is not
reduced at the cathode because of the large overvoltage for the discharge
of H 2 gas on Zn. The zinc is deposited upon aluminum electrodes from
which it is stripped at intervals, melted, and cast. During electrolysis, H 2 SO*
is regenerated and later used for leaching more roasted ore. Electrolytic zinc
is about 99.95% pure. The residues from the distillation and leaching processes
contain economic quantities of Pb, Ag, and Au, and are shipped to a lead
smelter for further processing.

High purity zinc is also made by zone refining. This is a process in which
a short molten zone moves slowly but repeatedly through a bar of the
metal to be refined, redistributing the dissolved impurities through the
mechanism of fractional crystallization as the molten zone travels from
one end of the bar to the other. The molten zone contains a melting liquid-
solid interface and a freezing liquid-solid interface. At the melting inter-
face all solid materials become mobile. If an impurity is one which lowers
the melting point of the Zn, it concentrates in the liquid phase and moves
with it. If the impurity is one that raises the melting point of the Zn, it
tends to concentrate in the solid Zn and the liquid becomes depleted in
that impurity. This technique has produced Zn with Pb and Cd contents
of but one part per million of each. The annual world production of Zn,
about ZVa million tons, is exceeded only by that of steel, aluminum, and
copper. Canada is the largest producer, followed by the United States and
the U.S.S.R.

8. Properties of Zinc. Zinc is a blue-white crystalline metal, brittle at
ordinary temperatures but malleable between 120°C and I50°C so that it
can be rolled into sheets between heated rollers; it then retains its flexibility
when cold. From 200°C to 300°C, zinc again becomes brittle. Though zinc
is a reactive element, under ordinary atmospheric conditions it becomes
coated with a continuous and adherent film of a basic carbonate which
protects it from further action. Zinc will reduce steam at high temperatures
and displace H 2 from dilute acid, although when the Zn is pure, the action
is very slow due to the large overvoltage for the deposition of H 2 on Zn.
Zinc also dissolves in solutions of alkali metal hydroxides with the formation
of the complex ion, Zn(OH) 4 2 "

(1) Zn + 2 OH- + 2 H 2 0 Zn(OH) 4 2 - + H 2

588

Transition Elements— V : The Elements of Group IIB

To prevent its corrosion, millions of tons of steel are coated with zinc
by electroplating, galvanizing ? or sherardizing. In hot dip galvanizing, the
metal to be protected is first immersed into a flux, usually a mixture of
molten ZnCl 2 and NH 4 C1, and then into liquid zinc at about 455°C. Pure
zinc coats the metal as it is withdrawn. Sherardizing involves heating the
steel object with zinc dust in a rotating drum. Another important use of
zinc is in die casting. Such castings are produced by forcing molten zinc
alloys into steel dies at high temperature and pressure. Die casting enables
the production of intricate parts within close limits of precision and yields
extremely smooth surfaces that are easily plated. Automobile hardware, fuel
pumps, and hydraulic brakes are commonly zinc die castings. Much zinc
is also used in brass alloys, in dry cells, for chemical manufacture, and as
sacrificial anodes to protect boilers, pipelines, and ship hulls.

4. Compounds of Zinc. The only oxidation state of zinc is +2. In aqueous
solution and in hydrated salts, the zinc ion has the form, Zn(H 2 0) 4 2+ ; this
hydrolyzes to yield acidic solutions. The reaction of Zn 2 + and OH" ions
gives a white gelatinous precipitate of Zn(OH) 2 . The hydroxide is amphoteric,
dissolving in excess strong base through the formation of the Zn(OH)* 2 -
complex ion. Complex ions are also formed with NH 3 and CN" ion,
Zn(NH 3 ) 4 2+ and Zn(CN) 4 2 ~, respectively.

In industry the most important compound of zinc is ZnO. It is made
by burning Zn vapor or by roasting ZnS ores in an excess of air. The finely
divided oxide is caught in filter bags which allow the gases formed to pass
through. All rubber contains about 5% ZnO since the oxide aids the curing proc-
ess and reinforces the rubber bond. Under the name of zinc white , or Chinese
white , ZnO is used as a paint pigment. Though white lead has superior
covering power as a pigment, ZnO does not darken when exposed to H 2 S. Be-
cause of its mild basic and antiseptic properties, ZnO is an ingredient of
some ointments.

Among the common metals, ZnS is the only white sulfide. The value of
its solubility product, I X 10" 22 , is sufficiently small that ZnS can be pre-
cipitated in neutral and basic solutions but large enough so that, in an acid
solution such as 0.3 M HCI, the S 2 " ion concentration is insufficient for
precipitation. This solubility of ZnS in acid solution affords a means of
separating Zn 2+ ion from more insoluble metal sulfides such CuS, Ag 2 S, and
CdS. Because it can convert the energy of an incident electron beam into visible
light, impure ZnS is used as a phosphor in the manufacture of fluorescent
screens, e.g., in television picture tubes.

Cadmium

Ancient brass workers noted that a brown dust collected in the flues of
their foundries and Pliny applied the term cadmia (Greek; earth) to this
material. The element cadmium was discovered in 1.817 by the German
chemist, F. Strohmeyer, as an impurity in ZnC0 3 . No ore is mined for cadmium
alone; its compounds occur in small quantities, 0.17-1.4% in zinc ores and
the metal is obtained as a by-product of zinc processing. Being more volatile
than zinc, cadmium is in the first portion of the distilled metal to come
over in the smelting of zinc ores whereas, in the electrolytic process for zinc,

Transition Elements— V: The Elements of Group IIB

589

cadmium is one of the less active metals precipitated by the addition of
zinc dust prior to electrolysis. Annual production of cadmium in the United
States is 5,000 tons, about half the world’s output. Like zinc, cadmium is
used in electroplating iron and steel to prevent corrosion, in alkaline storage
batteries, and in low melting and antifriction bearing alloys. Because of its
unusually high absorbance for thermal neutrons, 2450 bams, the metal is
used to control and to shield nuclear reactors.

In appearance and in chemical properties, cadmium is very similar to
zinc. It has a lower oxidation potential, however, and is not so active. The
hydroxide, Cd(OH) 2 , is not amphoteric but is more soluble and more
basic than Zn(OH) 2 . The sulfide, CdS, is yellow and its K Bp of 1.4 X 10~ 28
is low enough so that it can be precipitated in acid solution; it is used as
a yellow pigment.

Mercury

Mercury is one of the seven metals known to the ancients and has been
found in Egyptian tombs dating back to 1500 b.c. Its name is derived from
the Latin Mercurius , the god of commerce and its symbol from the word
for liquid silver, hydrargyrum » Some native mercury is found occasionally
but the principal ore is the red sulfide, cinnabar , HgS.

5. Metallurgy of Mercury. To obtain the metal the ore, HgS, is simply
roasted in air.

(2) HgS + 0,> ^ Hg(g) + SO,

Mercury vapor distills from the furnace and is condensed. In this instance
roasting does not yield the metal oxide because the oxide of mercury is itself
decomposed by heat. Solid impurities in mercury can be removed by filtra-
tion through chamois while by distillation, dissolved impurities such as other
metals, are left behind. Mercury can also be purified by pouring it, in a
stream of minute drops, through dilute HNO a to dissolve more active metals.
Among the nonferrous metals mercury ranks tenth in quantity of output.
Italy and Spain dominate the worlds production, about 10,000 tons a year;
the metal is shipped in iron flasks.

6. Properties of Mercury. Mercury is the only metal that is liquid at
room temperature; cesium and gallium melt at 28.5°C and 29.8°C, respectively.
At ordinary temperatures, mercury has a low vapor pressure, 0.001 mm at
28°C and but 0.28 mm at 100°C. Because of the convenient range through
which the metal remains liquid, -38.9°C to 357°C, its uniform coefficient of
thermal expansion over this range, and its nonwetting of glass, mercury is
used in the manufacture of thermometers. Its high density and low vapor
pressure make it ideal for filling barometers and manometers. The vapor
of mercury is colorless and monatomic.

Except for iron and platinum all metals dissolve in or are wetted by mer-
cury to form amalgams. When the concentration of the dissolved metal is
low, amalgams are liquid but as the quantity of solute is increased an
amalgam becomes stringy and eventually solid. The chemical reactivity, of
an active metal such as sodium is decreased by amalgamation because its.
concentration is, in effect, decreased by the formation of a solution, namely,

590

Transition Elements— V: The Elements of Group HB

sodium in mercury. Sodiufn amalgam is used as a reducing agent and amal-
gams of tin, silver, and gold are used in dentistry.

Though one of the poorer metallic conductors, large quantities of mer-
cury are employed in electrical devices which utilize its liquid properties,
such as switches, rectifiers, etc. Mercury vapor does not conduct electricity
when cold but once an arc is struck it conducts readily with the emission
of radiation rich in ultraviolet. A major use of the metal is as the cathode
in the electrolytic preparation of Cl 2 and NaOH. A relatively new application
is frozen mercury patterns for precision casting; because of its smooth surface
and its low, uniform thermal expansion, mercury is superior for this purpose
to plastic or wax.

Chemically mercury is quite unreactive. At ordinary temperatures it does
not oxidize in air but does combine slowly with O a near its boiling point to
form HgO. Mercury is oxidized by HNO :i> both dilute and concentrated,
and by hot, concentrated H 2 S0 4 .

(3) 3 Hg + 8 H+ + 2 NO ; f -> 3 Hg 2+ + 2 NO + 4 H.O (dilute HN0 3 )

(4) Hg + 2 H 2 S0 4 -» Hg 2+ + S(V~ + SO, +2 H 2 0 (hot concentrated H 2 SO t )

7. Compounds of Mercury; Mercurij(I) Compounds. Mercury exhibits
two oxidation states, mercury (I) or mercurous , in which the oxidation state
is +1 and mercury (II) or mercuric , in which the oxidation state is +2. The
-fl ion is unusual in that it exists in the form of a double ion, Hg, 2+ , wherein
the two mercury atoms are covalently bonded to each other, thus + HgSHg + .
Experimental evidence for this structure is the nonparamagnetic character
of this ion. If the +1 ion were simply Hg + it would be j>aramagnetic because
of a single unpaired electron.

Mercury (I) compounds can be prepared by heating Hg 2 * 1 " compounds
with elemental Hg. The chemical behavior of the Hg,~ + ion is similar to
that of the Ag + ion. Except for Hg,F 2 , which decomposes in water, the
mercury (I) halides are insoluble, Hg 2 I 2 being the least soluble. The most
important of the mercury(I) compounds is Hg,Cl,. It can be sublimed
from a heated mixture of HgCL and Hg and is also obtained as a white
crystalline precipitate when solutions of Hg 2 2+ and Cl~ ions are mixed. Under
the name of calomel, Hg 2 Cl 2 is used in medicine as a purgative and anthel-
mintic. Unlike Ag + ion, Hg 2 2+ ion does not form complex ions with NH S ,
CN" or S 2 0 3 2 '. When NH 3 is added to Hg 2 Cl, an oxidation reduction reaction
takes place. The mixture turns black due to the formation of finely divided
free Hg, and white insoluble mercury (II) amidochloride, HgNH 2 Cl, is also
formed. This reaction serves as a qualitative test for Hg 2 2+ ion.

(5) Hg 2 Cl 2 + 2 NH* Hg(I) + HgNH,Cl(s) + NH 4 * + Cl~

The compounds, Hg 2 0 and Hg 2 S, are unstable, decomposing into the Hg 2+
compound and Hg.

Mercury (II) compounds : The +2 is the more common of the two
oxidation states of mercury. The mercury (II) halides are soluble in water,
the iodide, Hgl„ again being the least soluble. In the presence of ex*-

Transition Elements— V: The Elements of Group IIB

591

cess halide ion the solubility is increased because of the formation of
complex ions, e.g., Hgl 4 2 ". Because of such complex ion formation the halide
salts are poor conductors of electricity. The chloride, HgCl 2> can be prepared
by direct union of its elements or by sublimation from a mixture of HgS0 4
and NaCl. The iodide, Hgl 2 , is a scarlet salt which changes to yellow when
heated above 128°C.

When a small concentration of tin (II) chloride, SnCl 2 , is added to HgCl 2 ,
a white precipitate of Hg 2 Cl 2 is first formed. When excess SnCl 2 is used
the white precipitate turns gray and then black because of the formation
of free Hg.

(6) 2 Hg 2+ + Sn 2+ Hg 2 2+ + Sn 4+ (low concentration of Sn 2+ )

(7) Hg 2 2+ + Sn 2+ 2 Hg (l) + Sn 4 ^ (excess Sn 2 +)

Generally, any reducing agent capable of reducing Hg 2+ to Hg 2 2+ can also
reduce it to Hg metal.

No hydroxides of mercury exist. The addition of NaOH to a solution
of Hg 2+ ion precipitates a yellow variety of the oxide, HgO; when this
oxide is heated it changes to the more common red variety. Strongly heated,
HgO decomposes into its elements. By the reaction of H 2 S and Hg 2 + ion,
HgS is formed as a black precipitate; if this black sulfide is sublimed a red
form is obtained. Of the common metal sulfides, HgS is the most insoluble,
having a K 8I> of 1.6 X 10' 54 . The common acids, and even hot concentrated
HN0 3 , do not attack HgS but aqua regia dissolves it through the formation
of the complex HgCl* 2- ion. Mercury (II) fulminate, Hg(ONC) 2 , is prepared
by the reaction of Hg and HN0 3 in alcohol solution. It is a white solid that
explodes violently when struck and is used in making percussion caps.

The oxidation potentials for mercury in acid solution are:

-0.789 v -0.920 v

Hg ^ Hg* 2 + Hg 2+

f f

I 0.854 v 1

8. Physiological Action of Mercury. The soluble compounds of mercury
and mercury vapor are poisonous, and even prolonged contact of liquid
mercury with the skin should be avoided. To act as a poison a sub-
stance must be absorbed within the bc^y and interfere with its physiological
reactions. Also known as bichloride of mercury and corrosive sublimate ,
HgCl 2 is extremely poisonous and as little as 0.2 gram may be fatal. Taken
internally it combines with the proteins of the kidneys, thereby preventing
that organ from removing metabolic waste products. Antidotes for mercury
poisoning are egg white and milk, the proteins of which precipitate mercury(II)
compounds. Because of its insolubility, however, calomel is not poisonous;
its medicinal action depends upon its stimulation of the liver and otbeT
organs of secretion.

592

Transition Elements— V: The Elements of Group 11$

QUESTIONS

1. Account for the following: (a) the Zn 2+ ion is smaller than the Ca 2 + ion
(b) the oxidation potentials of the Group I IB elements decrease with increasing
atomic number (c) HgO is readily decomposed by heat but ZnO is not
(d) Zn, Cd, and Hg readily form complex ions but the elements of Group IIA
do not (e) mercury is the only Group II element to show more than one oxida-
tion state.

2. Write equations for the metallurgy of (a) zinc starting with ZnS as an ore
(b) mercury starting with HgS. How can each metal be refined?

3. Explain why the Zn 2+ and Hg 2 2+ ions are nonparamagnetic.

4. Why does impure Zn react more rapidly with acid than does pure Zn?

5. Write equations for the reaction of (a) Zn and NaOH (b) Zn(OH) g ($) and
NaOH (c) Zn(OH) 2 ($) and NH 3 (d) Hg 2 + and Hg (e) Hg 2 + and Sn 2 +
(f) "Hg 2 Cl 2 and NH S (g) ZnS and CN~.

6. Would the following reagents reduce Hg 2+ and to what state: Zn, Fe 2 +
I-*, and H 2 S?

7. How could you distinguish bhemically between (a) Zn 2+ and Cd 2 + (b) HgC^
and Hg 2 Cl 2 (c) Zn(OH) 2 and Cd<OH) 2 (d) ZnCl 2 and HgCl 2 ?

8. A solution contains a mixture of Zn 2+ , Cd 2+ , and Hg 2 + ions. How could
you confirm the presence of each?

9. Explain why a solution of zinc nitrate is acidic.

10. Calculate the H 3 0 + ion concentration in a solution of 0.01M ZnCl 2 . The
hydrolysis constant, K h , is 2.5 X 10” 10 .

11. What will be the concentration of Zn 2 + ion when 50 ml of 0.010M ZnCl 2
and 50 ml of 0.010M NH 3 are mixed?

12. Calculate the lowest value of the pH at which ZnS can be precipitated by H 2 S.

13. A solution which contains 0.0010M Zn 2 + and 0.0010M Cd 2+ in 0.30M
HC1 is saturated with H 2 S. Calculate the concentrations of Zn 2+ and Cd 2+
ions remaining in solution.

The Elements of Group HIB
The Aluminum Family

The family of elements comprising Group IIIB includes boron, aluminum,
gallium, indium, and thallium, and is known as the Aluminum Family for
its most important member. The electron configurations of these elements are
listed below and their properties are given in Table 45-A.

Element

Atomic
. Number

2s 2 p

3s 3p 3d

6s 6p

Boron

2 1

Aluminum

2 6

2 1

Gallium

2 6

2 6 10

2 1

Indium

2 6

2 6 10

2 1

Thallium

2 6

2 6 10

2 1

1. General properties of the Aluminum Family. The atoms of these
elements have three electrons in their outermost principal quantum shell,
ns 2 up 1 , All the elements have an oxidation state of +3 and, in addition, the
three largest members of the family show an oxidation state of +1 through
the loss of only the single outermost p electron. The* atomic and ionic radii
of atoms of the Group IIIB elements are less than those of the neighboring
Group IIA and IA atoms. Many characteristics of the Group IIIB elements
derive from the small size of their ions coupled with their high positive
charge since both factors increase the tendency of these atoms to hold elec-,
trons strongly.

The elements of the aluminum family are less reactive and have lower
oxidation potentials than do both the alkali metal elements and the alkaline
earth elements. Even so, in view of the high values of the third ionization
potentials, their reactivities and oxidation potentials are unexpectedly high,
probably because of the large heats of hydration of the small, highly charged

594

The Elements of Group II IB: The Aluminum Family

ions. The decrease in oxidation potential with increasing atomic number, a
trend at variance with that in Groups IA and IIA, is in accord with a
corresponding decrease in the ionic heats of hydration.

Simple (or hydrated) ions of the- type M 3+ where three electrons are
lost are unusual and covalent bonds are formed more frequently. Boron, in
fact, holds its valence electrons so strongly that it forms no B a+ ions but only
covalent bonds. The smallest element of the group, boron, is nonmetallic
but the other elements of the aluminum family have distinctly metallic
properties. The hydroxides of the Group IIIB elements are less basic than
those of the Group IA and IIA elements; indeed, the hydroxide of boron,
B(OH) s , is the weak acid, boric acid, and its formula is generally written
as H3BO3. The normal variation in hydroxide basicity is again evident since
Al(OH) 3 and Ga(OH) 3 are amphoteric, In(OH) s is slightly basic, whereas
Tl(OH) s is distinctly basic.

Table 45-A

Properties

OF THE

Elements

OF Grouf IIIB

Property

Boron

Aluminum

Gallium

;

Indium

Thallium

Symbol

Atomic Number

Atomic Weight

10.811

26.9815

69.72

114.82

204,37

Isotopes (mass numbers

10(18.98)

27(100)

69(60.2)

113( 4.16)

203(29.5)

and percent)

11(81.02)

71(39.8)

115(95.84)

205(70.5)

Abundance in Earths

0.0003

8.13

0.0015

1 X 10-5

0.0003

Crust, %

Physical State at STP

brown

silver

gray

silver

blue-white

solid

Density at STP, g/cm 3

2.34

2.70

5.91

7.31

11.85

Melting Point, °C

2150

660

29.8

156J2

303.6

Boiling Point, °C

2550

2050

1980

2047

1457

Heat of Fusion, kcal/mole

5.3

2.55

1.34

0.78

1.02

Heat of Vaporization,

129

70.2

61.2

54.1

38.7

kcal/mole

Heat of Atomization,

141

77.5

65.0

58.2

43.0

kcal/mole

Ionization Potential, eV, 1st

8.30

5.98

6.00

5.78

6.11

2nd

25.15

18.82

20.51

18.83

20.42

3rd

37.92

28.44

30.6

27,9

29.7

4th

259.3

120.0

63.8

57.8

50.5

Electronegativity

2.0

1.5

1.6

1,7 |

1.8

Atomic Radius, A

0.80

1.248

1.245

1.497 |

1.549

Ionic Radius, A (+3)

0.20

0.50

0.62

0.81

0.95

Oxidation Slates

•+•1*4-3

+-1,4-3

+1,4-3

Coordination Numbers

4,6

Oxidation Potential, volt

4-1.66

+-0.53

4-0,34

-0.72

for M — > M 3 + +- 3 e

The Elements of Group IIIB: The Aluminum Family

595

Boron

The term borax (Arabic, buraq) is mentioned in early Latin writings
but the element was not isolated till 1808 by the French chemists, J. L. Gay-
Lussac and L. J. Thenard. Free boron does not occur in nature. The desert
area of Southern California is the world's main source of boron mineral
of which the principal ones are borax , Na 2 B 4 0 7 * 10 H 2 0, and kemite ,
Na 2 B 4 0 7 * 4 H.O.

Elemental boron can be prepared as a brown amorphous powder by the
high temperature reduction of boron oxide, B 2 O a , with Mg; the MgO formed
is removed by solution in HC1. By the electrolysis of a fused mixture of potas-
sium fluoborate, KBF 4 , the oxide, B 2 0 3 , and KC1 at about 1000°C, 99%
pure boron can be obtained. If the amorphous form of boron is fused and
then cooled, a crystalline variety, variously described as shiny metallic or
colorless crystals, is produced. This form is brittle but hard as diamond and,
like Si and Ge, is a semiconductor. Elemental boron has minor uses as a
deoxidizer for nonferrous metals, a grain refiner for aluminum, and a thermal
neutron absorber.

The properties of boron are very similar to those of the element silicon
in Group IVB. Except for reacting with the most potent oxidizing agents, F 2
and HN0 3 , boron is quite inert at room temperature. At higher temperatures,
it combines with 0 2 , burning with a green flame, and with other non-
metallic elements. With many metals, boron reacts to form borides, e.g.,
MguBz, and with fused alkalis to form borates, e.g., K 3 BO a .

The compounds of boron are of more importance than the element itself.
The oxide, B 2 0», is a hard brittle glass which, when molten, can dissolve
metal oxides in a manner analogous to that of Si0 2 to form clear glasses
which may be colored depending on the metal oxide used. In water, B 2 0 3
dissolves to form boric acid, H 3 BO a . Boron carbide, CB 4 , formed by the
high temperature combination of the elements, is among the hardest sub-
stances known and is used as an abrasive.

In contrast to normal metal halides the boron halides are nonionic, covalent
molecules. At room temperature, BF 3 and BC1 3 are gases, BBr 3 is a viscous
liquid, and BI 3 is a crystalline solid. All are colorless and are completely
hydrolyzed by water. These molecules are electron deficient in that there
is an insufficient number of electrons to establish an octet around the boron
atom and consequently they can react as Lewis acids. Thus BF 3 forms ad-
dition compounds with NH 3 and with HF. With the latter, fluoboric acid,
HBF 4 , is the product.

( 1 )

h:f:

BlF -> H + +

« o

A strong acid, HBF 4 dissociates into a proton and the fluorborate ion, BF\ .
Though HBF 4 has not itself been isolated many of its salts are known. The
other halogen atoms do not form analogous complex ions with boron. By

596

The Elements of Group IlIB: The Aluminum Family

virtue of its ability to function as a Lewis acid, BF 3 is widely used as a
catalyst in organic syntheses, e.g., the Friedel-Crafts reaction for attaching
alkyl chains to benzene rings.

The most common acid of boron is orthoboric acid , H 30 3> usually re-
ferred to simply as boric acid . It is prepared by the reaction of borax and
a strong acid such as H 2 S0 4 , from which it crystallizes in thin white plates
when cooled.

(2) Na 2 B 4 0 7 + H 2 S0 4 + 5H 2 0-^4 H 3 B0 3 (s) + Na 2 SO*

The solubility of H 3 BO s in water increases markedly with temperature,
from about. 4% at 18°C to about 30% at 100°C, The acid is monobasic and
so weak, its ionization constant being 5.8 X that it hardly affects litmus
and can be used safely as an antiseptic eyewash. When heated above 100°C,
H 3 BO s loses water to form successively metaboric acid , HB0 2 , then tetrar
boric acid, H 2 B 4 0 7 , at about 140° C, and ultimately B 2 0 3 . Few orthoborate
salts are known. The most important salt of boric acid is borax or sodium
tetraborate, Na 2 B 4 0 7 * 10 H 2 0. Because H 3 BO.j is a weak acid, aqueous
solutions of borax are basic by hydrolysis.

(3) B 4 0 7 2 - + 7 H 2 0 4 H 3 B0 3 + 2 OH“

Because of its basicity and because it forms insoluble borates with calcium
and magnesium, borax is utilized for softening water and in cleanser mixtures.

When heated, borax fuses to form a glass also capable of dissolving metal
oxides. If its formula is written as 2 NaB0 2 * B 2 0 3 , it will be seen that the
excess B 2 0 3 present can combine with metal oxides. For this reason borax is
used as a flux in welding operations. A superficial oxide coating, which would
prevent fusion of the metal surfaces, is dissolved by the molten borax. This
action is also the basis of the laboratory borax bead tests since many oxides
impart characteristic colors to molten borax.

With hydrogen, boron forms a series of compounds, the boron hydrides
or boraneSy which bear some resemblance to the silicon hydrides and the
alkane hydrocarbons. The empirical BH 3 has not been prepared, the simplest
boron hydride being diborane , B 2 H 6 , a colorless gas, and the highest being
decaborane, B 10 H 14 , a solid. These belong to an experimentally incomplete
borane series, B n H n -f 4 , and can be formed by the reaction of Mg s B 2 and
acid. The boranes are electron deficient in that, according to the ordinary
rules of chemical valency, there appear to be top few electrons to satisfy
all the bonds required for the formation of a stable molecule. Thus diborane,
if its structure were analogous to that of ethane, C 2 H«, namely H 3 B~BKU
would require seven bonds or fourteen valence electrons but. only twelve
are available. For diborane the experimental data are best explained by a
structure containing hydrogen bridges:

•

9 9

•

• •

•

'H'

H *

* H

Ihe Elements of Group IIIB: The Aluminum Family

597

It is more likely, however, that the bonding .electrons are not localized but
that all contribute in some measure to binding the molecule as a whole.
The boron hydrides show varying degrees of stability. They are hydrolyzed
by water and are decomposed by oxidizing agents. They inflame spon-
taneously in moist air and diborane reacts explosively with Cl 2 . Because of
their higher heats of combustion than the hydrocarbons and their rapid
burning rates, the boron hydrides and some of their alkyl derivatives are
considered potential jet and rocket fuels; the heat of combustion of diborane
is about 18 kcal/gram.

Aluminum

The relative abundance of aluminum, 8.13% of the earth's crust, marks it
as the most abundant metal and the third most abundant element on earth
but its chemical reactivity precludes its being found in the free state in
nature. The most common minerals are the aluminosilicates, among which
are the feldspars, as orthochise, KAlSi 3 O s ; the micas, as muscovite ,
KH 2 Als(Si0 4 ) B ; and the clays, as kaolin, H 2 Al 2 (Si0 4 ) 2 * H 2 0. Because of the
extreme difficulty in extracting aluminum from these compounds the only
commercial source from which the metal is obtained is the hydrated oxide,
bauxite , AJ 2 0 3 * n H 2 0. Though the leading producers of aluminum are the
United States, Canada, and the U.S.S.R., the largest deposits of bauxite are in
Guinea, Australia, Jamaica, Ghana, and Surinam. Domestic deposits occur
mainly in Arkansas, Alabama, and Georgia. The anhydrous oxide, A1 2 0 3 , occurs
as corundum and as emery while the mineral cryolite, Na 3 AlF 6 , though not
an ore, is utilized in aluminum metallurgy. Aluminum is a constituent of
many precious stones. The ruby and sapphire, both natural and synthetic,
are A1 2 0 3 colored by a trace of other metal oxides while the emerald, garnet,
aquamarine, and lapis lazuli are aluminum silicates.

2. Preparation of Aluminum. Despite the common occurrence of aluminum
minerals the metal was .not readily obtainable till the late nineteenth century.
The growth of aluminum from a laboratory curiosity to a metal of commercial
importance second only to iron makes an interesting chapter in the world's
industrial history. The early Romans applied the term alumen to any sub-
stance with an astringent taste. In 1808 Sir Humphrey Davy recognized” that
the earth, alumina, was the oxide of an element which he named aluminum
but he was unable to isolate the element. The German chemist, Friedrich
Wohler, is generally credited with first isolating the pure metal in 1827 by
the reduction of A1C1 3 with the more active potassium, though an impure
sample had been prepared by Hans Christian Oersted two years earlier.
It was then evident that the metal had desirable characteristics which would
make it a useful material if it could be produced cheaply. In 1854 the
French chemist, E. H. Ste. Claire Deville substituted the cheaper sodium
for potassium as the reducing agent and the .price of aluminum dropped from
$160 in 1827 to $10.00 a pound in 1862. When H. Y. Castner, an American,
perfected the electrolytic process for making sodium, the price per pound
dropped to $4.00 in 1886. That same year, Charles M. Hall, based on work
started while a student at Oberlin College, and Paul Heroult, a French
metallurgist, independently discovered the technique of obtaining aluminum

598 The Elements of Group II IB: The Aluminum Family

metal throughihe electrolysis of alumina, Al.Oj, dissolved in cryolite, Na 3 AlF 6 ,
In 1885 less than 300 pounds of aluminum were produced in the United
States but ten years later, through the introduction of the electrolytic Hall-
Heroult process, production had increased to 900,000 pounds and the price
had fallen to $0.60 a pound. Today over two million tons of aluminum are
produced annually in the United States at a price of about $0.27 a pound
Because of its high melting point, 2050°C, pure A1 2 0 3 alone cannot be
fused readily nor electrolyzed. The oxide, however, dissolves in molten cryolite,
forming a solution which melts below 1000 °C. This solution of A1 2 0 3 in
Na 3 AlF 6 can be electrolyzed at 900-1000°C to produce 99.5% pure aluminum.
The apparatus is shown in Figure 45.1. The electrode reactions are:

(4) cathode: Al s+ + 3 e -» A1

(5) anode: 2 0 2 ~ - 4 e O s

The electrolytic cell is a carbon-lined steel shell, approximately 20' by 10" by 3'
deep.. The carbon lining acts as the cathode and one or more carbon blocks sus-
pended in the electrolyte serve as the anode. The electrolyte is Al 2 0 3 dissolved in
molten Na s AlF 6 as solvent. At the temperature of the bath the A1 is liquid; it
collects at the bottom of the cell and is drawn off through a tap. As the metal
is withdrawn, more A1 2 0 3 is added to the bath by breaking the frozen crust of
cryolite. The carbon anodes are consumed by reaction with the O a produced and
must he replaced frequently.

Figure 45 A. The Hall-Heroult Process For Mocking Aluminum.

Neither the Na + nor the F~ ions are discharged at the electrodes since the ions
of the less active elements, Al 3+ and O 2 ", require a lower potential and are
preferentially discharged. In this process the cryolite plays the role of solvent
similar to that played by water in aqueous solutions of electrolytes. Because
it is only the dissolved A1 2 0 3 that is consumed the cryolite solvent should
theoretically last indefinitely but, in practice, some is lost by vaporization
and must be replaced. Because the equivalent weight of aluminum is low
and because the electrical resistance of the bath is high, electrolytic produc-
tion of almqmum uses 3% of all domestic electric power generated; about

The Elements of Group 1IIB: The Aluminum Family

eight kilowatt-hours of electrical energy are required to produce one pound
of aluminum.

Because it is so active an element, aluminum cannot be purified readily
so it is necessary to purify the A! 2 0 3 prior to electrolysis. The chief impurities
in raw bauxite are Fe 2 0 3 and Si0 2 and these are removed by taking ad-
vantage of the amphoterism of A1 2 0 3 . In the Bayer purification, bauxite is
digested with concentrated NaOH solution at about 150°C. The Fe 2 O s is un-
affected, the SiO s is precipitated as an aluminum silicate, and the Al 2 O s is
brought into solution as the aluminate ion, Al(OH)r. The solution is filtered
and Al(OH) 3 is precipitated by blowing air containing C0 2 through the
solution to cool it and seeding the supersaturated solution formed. Calcining
the Al(OH) a then yields over 99% pure A1 2 0 3 .

The production of “super-purity” aluminum, over 99.99% Al, requires a
second electrolysis. In the Hoopes process, three immiscible molten layers
are used. The bottom, which also serves as the anode, consists of the “im-
pure” aluminum to which copper has been added to increase its density.
The intermediate layer is cryolite with added BaF 2 to give it a density inter-
mediate between that of impure aluminum and 100% aluminum. The top
layer, which acts as the cathode is pure aluminum. The mechanism of puri-
fication is quite analogous to that of the electrorefining of copper.

3. Uses of Aluminum. Industrially aluminum metal is more important
than its compounds. It finds many uses because of its low density, high
strength to weight ratio, high electric and thermal conductivity, resistance
to corrosion, low neutron absorption cross-section, high reflectivity for heat
and light, nontoxicity, and nonmagnetic and nonsparking properties.

The major uses of aluminum are in construction, for example in buildings
as roofing and siding; in aircraft and buses as their frames; in cars as pistons,
crankshafts, and trim; in the superstructure of ships, and in household ap-
pliances. Though aluminum is ductile and malleable the pure metal is difficult
to work on a lathe because its sticks to the cutting tools. Many alloys,
however, can be machined with ease. Duralumin (95 5% Al; 3 % Cu; 1% Mn;
0.5% Mg) is an alloy which is light but nearly as strong as steel. Thin sheets
of aluminum are often used for wrapping foods and the powdered metal,
suspended in oil, fs used in making “silver” paint. The 200-inch mirror in
the Mt. Palomer telescope is coated with aluminum. Aluminum is the lightest
of the common metals. Comparing wires of equal cross-section the electrical
conductivity of aluminum is about two-thirds that of copper but for wires
of equal length and weight its conductivity is twice that cf copper; thus
aluminum cable reinforced with a steel core has displaced copper for cross
country transmission lines,

4. Properties of Aluminum. Aluminum is a relatively active metal. When
freshly cut, it is silver-white but tarnishes in moist air due to superficial
oxidation. Further action is prevented by a firmly adhering, continuous film
of oxide that forms on the surface. For this reason, though one might
expect that so active an element should displace H 2 from boiling water,
aluminum can be used for cooking utensils. Though the metal displaces H 2
from acids and alkalies, HN0 3 renders it passive; in fact, HNO$ is shipped

600 The Elements of Group H1B: The Aluminum Family

in aluminum containers. At high temperatures, aluminum combines directly
with most nonmetals.

The high reactivity of aluminum is well illustrated by its reduction of
metal oxides. When a mixture of finely divided Al and Fe 2 0 3 is ignited by
a suitable fuse such as Mg ribbon, a highly exothermic reaction occurs.

(6) 2 Al + Fe 2 O s 2 Fe ( l ) + AUMZ) AH = -202 kcal

At the temperature produced by the reaction, 3000° C, both the Fe and the
Al 2 0 3 formed are liquid. Liquid iron produced by this reaction is used in
welding operations. Since the iron collects at the bottom of the reaction con-
tainer (a refractory material) in a layer separate from the A1 2 0 3 , it can be
tapped and run into a sand mold which surrounds the preheated ends of
the metal to be joined. In this manner repairs can be made to immovable
heavy steel machinery such as rails or crankshafts. The mixture of Al and
Fe 2 0 3 is known as thermite and was also used as military incendiary during
World War II. Analogous aluminothermic reactions are employed to pro-
duce samples of metals, e.g., chromium and manganese, whose oxides are
not readily reduced by carbon or where a carbon-free specimen of such
a metal is desired.

5. Compounds of Aluminum. Aluminum has but one oxidation state, -f‘3,
and its chemistry is quite regular. The chemistry of aqueous solutions of
aluminum salts can be interpreted in terms of the hydrated aluminum ion,
Al(H 2 0) 6 3 +. Though this formula has not been verified it is so written in
accordance with the formula of the salt, A1C1 3 * 6 H 2 0. Thus aqueous solu-
tions of aluminum salts are acid by hydrolysis. When NH« is added to
a solution of aluminum salt, a white, gelatinous precipitate of aluminum
hydroxide is formed.

(7) A1(H 2 0) 6 3 + + 3 OH- A1(0H) 3 (H 2 0) 3 ($) + 3 H 2 0

The hydroxide, usually written simply as Al(OH) 3 , is amphoteric and dissolves
in excess strong base with the. formation of the Al(OH) 4 “ ion, so that if
NaOH is added cautiously to an aluminum salt a precipitate of Al(OH) 3
first forms and then redissolves. Aluminum salts do not react with NH 3 to
form complex ions.

The precipitation of Al(OH) 3 is used for water purification. As the
freshly precipitated gelatinous material settles slowly it entraps finely sus-
pended material and bacteria in the water. Precipitated Al(OH) s also is
capable of adsorbing organic dyes and advantage is taken of this property
in dyeing textiles, notably cotton and linen. The hydroxide is first precipi-
tated in the fibers of the cloth by dipping the cloth successively into a solu-
tion of A1 2 (S0 4 ) 3 and a basic solution such as Na 2 C0 3 . The fabric is then
immersed in the dye; the dye, which alone does not adhere to the fiber,
is fixed in the fabric by the Al(OH) 3 . Other precipitated hydroxides, such
as those of iron, chromium, and tin, also exhibit this dye-adsorbing property
and are known as mordants .

When Al(OH) 3 is heated to dull redness, about 400°C, it is dehydrated
to A1 2 0 3 , a soft white powder which will reabsorb water and is soluble in

The Elements of Group IUB: The Aluminum Family

both acid and alkali. But if heated strongly the resulting Al 2 O s will undergo
none of these reactions; presumably oxygen bridges are formed between the
aluminum atoms to give an insoluble polymeric structure.

In addition to bctuxite , AlsOj is found in nature as corundum and
emery, substances nearly as hard as diamond and used as abrasives Ahmdum,
an artificial corundum made by heating bauxite to incipient fusion, about
2500°C, in an electric furnace is also used as an abrasive and as a refractory
material for crucibles and furnace linings. Synthetic rubies and sapphires are
manufactured by melting, in an oxyhydrogen flame, powdered A1 2 0 3 con-
taining sufficient quantity of a coloring oxide. The flame is allowed to im-
pinge upon the end of a fire-clay rod where the molten alumina forms a
carrot-shaped single crystal known as a bottle , usually % to 1 inch in diameter,
2 to 4 inches long, and weighing from 75 to 250 carats. The boule is then
heat-treated or split to relieve the inherent strain, after which it is cut
and polished into gems. Such gems are not imitations because they are
identical chemically with the natural products and can be distinguished
only by experts through the presence of characteristic minute air bubbles.

A clear distinction should be made between the properties of the pure
anhydrous aluminum halides and their aqueous solutions. Except for the
fluoride, the aluminum halides are nonionic, covalent compounds which
have relatively low melting and boiling points and are soluble in weakly
polar organic solvents such as ether and pyridine. These halides are white
deliquescent solids which have a high vapor pressure and sublime without
melting. Vapor density measurements and molecular weight determinations
in organic solvents indicate that the molecules are dimeric, e.g., A1 2 C1 6 ,
wherein the formation of chlorine bridges between the aluminum atoms
results in an octet of electrons around each aluminum atom.

* • •• #• •

.ci' ci ci;

^>1;^

Cl <C1'^ .Cl*.

With fluorine, the aluminum atom forms a series of complex ions ranging
from A1F 2 + to A1F« S -, the anion in cryolite. The tendency to form complex
ions with the larger halogen atoms is much less but the AlClr ion is known
and probably is the cause of the corrosion of aluminum in sea water.

Anhydrous aluminum chloride is an effective catalyst for many organic
reactions, notably the Friedel-Crafts reaction and the cracking of gasoline.
Depolymerization of the dimeric molecule yields the monomer, A1C1 8 , which
is an electron deficient molecule like BF 3 and can behave as a Lewis acid.
Thus AlCls forms double halides with halides of nonmetals and even forms
addition compounds with such unlikely molecules as N0 2 , H 2 S, PCls, SO a
and NH*.

With water the aluminum halides react vigorously with the evolution of
much heat Though the anhydrous salts are poor conductors of electricity,

(302 The Elements of Group MB: The Aluminum Family

their aqueous solutions are good conductors because of the ions produced
therein.

(8) A1 2 C1« + 9 H 2 0.-» 2 Al(H 2 0)e 34 + 8H+ + 12 Ch

Evaporation of these solutions yields hydrated salts, e.g., A1C1 { • 6 H 2 0; when
heated this salt does not lose water but decomposes into Al s O fl and HCI.

Potassium aluminum alum, KA1(S0 4 ) 2 *12 H y O, is the prototype of a
class of compounds known as alums. It is readily obtained by the evaporation .
of a solution containing equimolecular amounts of its component salts, K 2 S0 4
and Al 2 (SO*)»* Alums have the general formula, M(I)M(III) (SO») 2 * 12 H 2 0
where M(I) is a singly charged cation such as K 4 , Na 4 , NHj 4 , or Tl 4 and
M(I11) is a triply charged cation such as AF 4 , Cr* 4 , Fe : ’' 4 , or TP 4 . Hence
an alum need not contain aluminum. Most alums crystallize in the octahedral
form and are isomorphous , that is, they form crystals which are alike in shape
and belong to the same system. Thus, a transparent layer of ammonium alum,
(NH 4 )A1(S0 4 ) 2 ‘ 12 H a O, can be grown over a purple crystal of chrome
alum, KCr(S0 4 ) 2 * 12 H 2 0, with no alteration of the octahedral shape.
(Figure 45.2)

Gallium, Indium, and Thallium

All three of these elements were first discovered spectroscopically; gallium
in 1877 by the French chemist, P. de Boisbaudran who named it for the
Latin word for France, Gallia ; indium in 1863 by the German chemists, F.
Reich and T. Richter, and so named for its characteristic indigo-blue spectral
line; and thallium in 1861 by the English physicist, Sir William Crookes,
who named the element for the Greek thallos , a budding twig, because
of its characteristic green spectral line.

The metals are extremely rare; the earths crust contains about 15 grams
of gallium per ton while indium and thallium are even less abundant. No
known minerals contain these elements as a major constituent but rather
they occur sporadically associated with the ores of many other metals. Gallium
occurs in bauxite and is obtainable as a by-product of the Bayer process
while all three metals are found in the flue dust from the smelting of zinc
ores. The metals can be prepared ultimately by electrolysis or by reduction
with a more active metal, but since none of them have major industrial uses
production is small.

The metals are soft, much like lead, and gallium is unique because it
is liquid over a temperature range of almost 2000° C. Gallium is used in

The Elements of Group IIIB: The Aluminum Family

optical mirrors and for dental alloys with gold; as a liquid it is used to seal
glass joints and valves in vacuum equipment. Indium finds use in semiconduc-
tor devices and in special high speed bearing alloys. Thallium metal and its
compounds are toxic; the latter are used to exterminate rats and ants. The
reactions of these metals and their ions are quite similar to those of
aluminum; the increased basicity of their hydroxides has already been noted.

QUESTIONS

L Account for the following: (a) the heat of combustion of boron is .greater
than that of carbon (b) BF 3 is a Lewis acid (c) boron forms only covalent
compounds (d) the difference between the third and fourth ionization po-
tentials of boron is greater than the same difference for thallium (e) the
three largest elements of the aluminum family have more than one oxida-
tion state ({) fused aluminum chloride is a poor conductor of electricity but
an aqueous solution is a good conductor (g) nitric acid is shipped in aluminum
containers (h) alums are isomorphic (i) H 3 B0 3 is monobasic (j) feldspar
is not used as an ore of aluminum (k) borax is a soldering flux.

2. Describe the electrometallurgy of aluminum. What are the electrode reactions?
Why cannot aluminum be obtained by the electrolysis of an aqueous solution
of aluminum ions?

3. Why must bauxite be purified prior to the electrolysis of A1 2 0 8 ? In what
manner are impurities of Fe 2 0 3 and Si0 2 removed?

4. Write equations for the following: (a) hydrolysis of BF 3 (b) hydrolysis of
A1 2 (S0 4 ) 3 (c) hydrolysis of Na 2 B 4 0 7 (d) A1 and NaOH (e) preparation of
Cr metal by aluminothermy (f) preparation of boric acid from borax
(g) Al 2 (S0 4 ) s and NH 3 (h) the dropwise addition of NaOH to A1 2 (S0 4 ) 3 .

5. Why does not an aluminum carbonate precipitate when solutions of A1 2 (S0 4 ) 3
and Na 2 CO s are mixed?

6. How could you separate the components of a mixture of magnesium, zinc,
and aluminum hydroxides? Write pertinent equations.

7. Draw the structural formula for aluminum chloride. What evidence is there
for such a formula? Why can aluminum chloride act as a catalyst?

8. Write electron dot formulas for addition compounds of aluminum chloride
and (a) NH S (b) S0 2 .

9. What weight of H 3 BO s can be made from 100 grams of hydrated borax?

10. In the Hall-Heroult process what weight of A1 2 0 3 will be decomposed by a
current of 10.0 amperes flowing for 24 hours?

11. What weight of A1 2 (S0 4 ) 3 • 18 H a O is required to make 2.0 liters of 0.40N
solution?

The Elements of Croup IVB
Germanium, Tin, and Lead

In Chapters 31 and 35 the first two elements of Group IVB, carbon and
silicon, were studied and it was there noted that the metals of this group
would be deferred to a later chapter. These elements are germanium, tin, and
lead. Their electron configurations were given in Chapter 31 and their
properties in Table 31-A. None of these metals is found uncombined in
nature. They are quite rare, with even the commonplace tin and lead counted
among the least abundant elements, but fortunately their occurrence is in
concentrated and workable deposits. Tin and lead were known to ancient
man and their use predates the Christian era by 2000 to 4000 years.

Unlike carbon and silicon these elements have two oxidation states, +2
and +4. In the -f-2 state only the outermost np 2 electrons are used in bond
formation whereas in the +4 state all four outermost electrons, ri sr np 2 , are
used, a situation similar to that of Ga, In, and T1 in Group IIIB . With the
increase in atomic size from germanium to lead, the +2 state becomes rela-
tively more stable and of greater importance than the +4 state.

Germanium

It was from the mineral argyrodite , Ag 4 GeS 4 , that C. A. Winkler first
isolated the element, germanium (Latin, Germania : Germany). Germanium
satisfied the properties foretold by Mendelejeff for an element of atomic
number 32, which he called eka-silicon and had predicted would have proper-
ties similar to those of silicon.

The occurrence of germanium in the earth’s crust, but one part in 150,000,
marks it as a decidedly rare element. Sulfides of germanium are found in
minute quantities associated with the sulfide minerals of many other ele-
ments and in some eases the compounds appear to be complex thiogermanates,
as in argyrodite. The most abundant occurrence is in sphalerite, ZnS, in
quantities ranging from traces to tenths of one percent. The separation of
the germanium from the more abundant zinc is a complex chemical process
involving the roasting of the zinc ore to the oxide, then solution in HC1
from which volatile GeCl 4 can be distilled, followed by conversion of the

The Elements of Group JVB; Germanium , Tin , ancZ Lead

605

GeCl 4 to GeO s . The metal is then obtained by reduction of Ge0 2 in a cur-
rent of H 2 in an electric furnace at 650 °C till no more H a O is evolved* The
H 2 is then replaced by helium or nitrogen and the temperature is raised to
1000°C to melt the germanium powder produced. About 45,000 pounds of
germanium, valued at over $60 million, are consumed annually in this 1 country.

Except for the formation of two oxidation states the compounds of ger-
manium resemble those of silicon in most respects. It is the property of
germanium metal as a semiconductor, however, which provides its only
significant use in diqdes, transistors, arid in power rectifiers. For these
purposes it is necessary to have ultra-high purity germanium having impuri-
ties measured only in parts per billion. Less than one part per 100 billion
of some impurities has an observable effect on the behavior of germanium in
transistors. The extreme purity required can be achieved by zone refining.
A germanium ingot 12 to 20 inches long and about one inch in diameter is
passed through a series of induction coil heaters, each of which produces a
molten zone. While pure germanium crystallizes at the trailing edge of each
molten zone the impurities remain within the molten zone and are gradually
isolated toward one end as the bar moves through the heaters. After a num-
ber of passes the end of the ingot containing the impurities is cut off. So
produced, ultra-pure germanium is polycrystalline but for use as a semi-
conductor the metal must be in the form of a single crystal. Two methods of
single crystal growth are in use. In one, a seed crystal is dipped into molten
germanium and slowly withdrawn. By adjusting the heat input and the
heat loss, germanium will grow upon the seed until the mass has been
converted into a single crystal. The second method is a modification of zone
melting. A seed crystal is placed at one end of a carbon-coated silica crucible
filled with ultra-pure polycrystalline germanium. By carefully moving a
molten zone toward the seed, melting into it, and then slowly withdrawing
the molten zone the whole can be converted into a single crystal.

Like diamond, pure germanium crystallizes in a tetrahedral structure. In
this structure the four valence electrons of each germanium atom are held
strongly in covalent bonds to four neighboring germanium atoms. No free
electrons are available to conduct an electric current and a pure germanium
crystal should be totally nonconducting. Thermal energy disrupts some of
the covalent bonds, however, so some electrons are freed to conduct current.
With an increase in temperature, more electrons are set free so that the
conductivity of a semiconductor increases with a rise in temperature.

X. Transistors. If a trace of an impurity is introduced into a pure ger-
manium crystal its current carrying characteristics are changed radically.
If the impurity is an atom such as arsenic or antimony which has five elec-
trons in its outermost shell, replacement of a germanium atom in the crystal
lattic will introduce a fifth electron which will be in excess and which
can find no "place” or energy level available within the normal covalent
bonding structure. This electron is thus free to move throughout the crystal,
that is, to carry a current Relatively few impurity atoms are required to yield
an appreciable current when an external potential is applied across the
crystal. The procedure of adding an impurity to a pure germanium crystal

606

The Elements of Group IVB: Germanium, Tin, and Lead

is known as doping , and germanium doped by pentavalent donor atoms is
known as N-type germanium because its conductivity is due to negative
charge carriers, or electrons.

On the other hand, if the impurity added to a crystal of germanium is
an atom with only three valence electrons, such as gallium or indium, one
of the four covalent bonds of a germanium atom will be deficient one elec-
tron. The absence of an electron is referred to as a “hole” and it acts like
a positively charged particle when a voltage is applied across the germanium
crystal. To fill the vacancy an electron moves into the hole, thereby creat-
ing a new hole in the position, or energy level, vacated. In effect, the hole
moves in a direction opposite to that of the electron flow and behaves like
an actual mobile positive particle. Germanium doped by trivalent acceptor
atoms is called P-type germanium because conduction can be considered
due to the movement of positive charges, namely, the holes.

For electronic applications, a P-N junction diode is made by joining
P-type germanium to N-type germanium, whereas a junction triode consists
of two P-N junctions. There are two types of transistors, the P-N-P and the
N-P-N, but the electronic application of these devices, the diode as a recti-
fier and the triode as an amplifier or oscillator, is beyond the scope of
this text.

Tin

The name of this element is derived from the Anglo-Saxon tin and the
symbol from its Latin name, stannum. The only commercial ore of tin is
cassiterite, Sn0 2 . Deposits of economic importance occur in few areas,
principally in Malaya, Indonesia, and Bolivia; no important deposits occur
in the United States though half of the annual world production of 200,000
tons is consumed in this country.

Elemental tin is prepared by the reduction of cassiterite with carbon
in a reverberatory or blast furnace at 1200° C, using a suitable flux. Liquid tin
is drawn from the bottom of the furnace and cast in blocks. Refining is by
liquation. The tin is heated on an inclined hearth to a temperature slightly
above the melting point of pure metal but below the melting points of the
impurities. These, mainly compounds of iron and copper, remain as an
infusible dross while the greater part of the pure low melting tin flows
off. Tin can also be refined by electrolysis but, because of the small demand
for high purity metal, the process is rarely used. Considerable secondary tin
is recovered from tin-plated scrap metal. The annual world production of
tin, approximately 200,000 tons, ranks fifth among the nonferrous metals,
being exceeded only by aluminum, copper, lead, and zinc.

An unusual combination of properties— low melting point, resistance to
corrosion and fatigue,^ malleability, and ability to alloy with other metals—
accounts for the many uses for which tin is preferred. The largest use of
tin is as a protective coating for steel, copper, and other stronger but more
reactive metals, but mainly as tin plate for cans used in the preservation
of food. Tin plate is sheet steel coated with a thin, coherent film of tin
formed by dipping into a bath of molten tin or by electroplating. Tin has a
lower oxidation potential than iron, hov^ver, so if the tin coating is punctured

The Elements of Group TVS: Germanium Tin , and Lead

607

the more active iron will corrode rapidly. Another major use of tin is in
tin-lead alloys known as soft solders and for alloys such as bronze, gun metal,
babbitt, pewter, and type metal.

2. Allotropic Forms of Tin. Solid tin exists in three allotropic modifica-
tions as shown below.

13.2“ C 161 °C 232 °C

(1) Sn (alpha) Sn (beta) Sn (gamma) Liquid

gray tetrahedral white tetragonal rhombic

Ordinary tin is in the beta form, silver-white, ductile and malleable so
that it can be beaten into thin sheets of tin foil. When a bar of beta, tin
is bent, a low crackling sound known as '"tin cry” can be heard due to the
slippage of its crystals. When heated for some time above 161 °C, tin be-
comes brittle due to its conversion to the rhombic form or gamma tin.

Alpha tin is a gray friable powder whose crystal structure is tetrahedral
like diamond and germanium. Transition to this modification occurs below
13.2° C but even though the ambient temperature drops below this value,
conversion between solid phases is very slow so that powdery, gray tin is
not formed unless the low temperature is maintained for a long period of
time. The rate of the change increases as the temperature is lowered, attain-
ing a maximum value at -48°C, and appears to be catalyzed by the presence
of gray tin so that once the conversion has started it proceeds rapidly. Be-
cause the change spreads from a central point of origin, much like an in-
fection, the transformation is called “tin disease” or “tin pest.” Several
cases have been reported where ordinary tin, during an exceptionally cold
winter, changed to the powder variety. In 1851, Erdman observed that the
tin pipes of an organ had crumbled to a powder while in 1868 a shipment
of block tin stored in the custom house of St. Petersburg was found to have
changed in the same manner.

Similar allotropic conversions are shown by many metals, and divergent
data for their properties is often due to experimentation with metals in
partially transformed states, that is, metals containing different percentages
of their allotropic forms.

3. Properties of Tin. The metal is but slightly more reactive than hydro-
gen arid is not affected by moist air. From dilute acids, such as HC1 and
H 2 S0 4 , tin slowly displaces H 2 , being oxidized to the +2 state. In cold,
dilute HNOh the metal dissolves to form Sn(NO,) 2 but with concentrated
HN0 3 insoluble metastannic acid, H 2 Sn0 3 , or hydrous tin (IV) oxide,
Sn0 2 * n H 2 0, is produced. Tin also reacts with alkali hydroxides to form
st annite ion, Sn(OH) : r» and H 2 .

(2) 4 Sn+ 10 H+ + NO,- 4 Sn a+ + NH 4 + + 3H 2 0 (cold dilute HNO a )

(3) $n + 4H + + 4 NO,” H 2 SnO, + 4 N0 2 + H a O (concentrated HNO,)

(4) Sn 4- OH” 4- 2 H 2 0 Sn(OH),” 4- H 2

608

The Elements of Group IV B: Germanium , Tin, and Lead

Tin forms two series of compounds, tin (II) or stannous, in which the
oxidation state is +2, and tin (IV) or stannic, in which the oxidation state
is Simple Sn 4+ ions do not exist, however, because their compounds
•are covalent and in aqueous solution are highly hydrolyzed. Both oxidation
states constitute part of many complex ions, e.g., SnCl 4 2 ", Sn(OH) 3 ~, SnCl 6 2 -,
and Sn(OH) 6 2 ~. When tin is dissolved in a non-oxidizing acid, the Sn 2+ ion is
formed. This is readily oxidized to the +4 state and this, in turn, can be re-
duced to the +2 state even by weak reducing agents. Solutions of Sn 2+ are
preserved from oxidation by the addition of tin metal. Because of appreciable
hydrolysis, aqueous solutions of both Sn 2+ and Sn 4+ compounds are acidic
but the tendency to hydrolyze can be reduced by the presence of a strong
acid. The oxidation potentials for tin are summarized below.

0.14 v -0.15 v

In acid solution: Sn Sn 2+ Sn 4+

0.91 v 0.90 v

In basic solution: Sn Sn(OH) 4 2 ~' +» Sn(OH) 6 2 "

When tin is burned in an excess of air or oxygen white SnO z is produced.
The lower oxide, SnO, is a black powder which can be prepared by heating
Sn(OH) 2 . The two hydroxides of tin are formed as white precipitates when
a soluble base is added to solutions of Sn 2+ and Sn 4+ compounds. They
usually are formulated as Sn(OH) 2 and Sn(OH) 4 but are more likely
hydrous oxides, SnO • n H 2 0 and Sn0 2 * n H a O. Both hydroxides are am-
photeric, dissolving in excess base with the formation of the stannite ion,
Sn(OH) 3 ", and the stannate ion, Sn(OH) e 2 ”, respectively.

The chlorides of tin are its most commonly used compounds. The direct
combination of Sn and Cl 2 forms tin (IV) chloride, SnCl 4 . This is a colorless
liquid which fumes in moist air and is extensively hydrolyzed in aqueous
solution. The lower chloride, SnCl 2 , can be made by the reaction of Sn and
HC1 or by the reduction of SnCl 4 . Aqueous solutions of SnCl 2 become turbid
due to the hydrolytic formation of the slightly soluble basic chloride, Sn(OH)Cl.
A weak reducing agent, SnCl 2 can reduce Hg 2+ to Hg, and Fe 34 * to Fe 24

The sulfides of tin, brown SnS and yellow SnS 2 , can be precipitated from
dilute solution by the addition of H 2 S to solutions of Sn 2+ and Sn 4+ ions,
respectively. In high concentrations of sulfide ion, S 2 ~, e.g., the alkali metal
sulfides, SnS 2 dissolves with the formation of the thiostannate ion, SnS 3 2 ', a
behavior which distinguishes it from another yellow sulfide, CdS. Acidifica-
tion of the thiostannate solution reprecipitates SnS 2 and H 2 S is evolved. In
contrast, SnS is insoluble in S 2 " ion, e.g., (NH 4 ) 2 S but is soluble in ammonium
polysulfide, (NH 4 ) 2 S 2 , due to the oxidation of SnS to SnS 3 2,

Lead

The name of the element is Anglo-Saxon but the symbol, Pb, is obtained
from the Latin plumbum . Despite its inactivity, no elemental lead occurs in
nature, the common minerals being galena , PbS; cerussUe, PbCO B ; and

The Elements of Group IVB: Germanium, Tin, and Lead

609

mglesite , PbS0 4 . Of the three, galena is the most abundant in the United
States and lead is obtained almost exclusively from it. The ore is first roasted
to convert the PbS to PbO and then charged into a blast furnace with
coke and a suitable flux. Air is blown through the mass, burning the coke
to CO and producing a temperature of 1400°C. The CO reduces the PbO
to Pb, which is tapped as a liquid from the bottom of the furnace. Impuri-
ties of active metals in the lead are removed by reheating to just above the
melting point of lead whereupon oxides of these metals are formed and can
be skimmed off as a dross. In addition to base metal impurities the lead
often contains sufficient gold and silver to warrant their extraction by the
Parkes process. For most commercial purposes, however, lead so refined is
sufficiently pure but some lead, especially if bismuth is present above ac-
ceptable limits, is refined electrolytically by the Betts process. In this process,
the impure lead is made the anode and the electrolyte is a solution of lead
fluosilicate, PbSiF 6 .

Though lead is a rare metal, in tonnage produced it ranks fifth among
all metals, exceeded in quantity only by steel, copper, aluminum, and zinc.
The annual world smelter production is about 2Vz million tons, of which
the United States is the leading producer, supplying about 1/5 of this total.
Other important producers are the U.S.S.R., Australia, and Mexico. The use
of lead in storage batteries (32%) and in the manufacture of tetraethyl lead
(16%) consumes about half of all the lead produced. Other important uses
are for cable coverings, paint pigments, and in various alloys, particularly
with Sb and Sn. The uses in paint pigments and in tetraethyl lead are dis-
sipative but the greater part used for other purposes is recoverable and is
returned to commerce as secondary metal.

Because lead markedly attenuates nuclear and x-radiation, an important
though small use is for shielding purposes. In nuclear applications another
advantage is that the lead itself does not become radioactive when subjected
to such radiation. However, it must be free of impurities which can be trans-
formed into radioactive isotopes through radioactive bombardment. The lead
atom is the end product of all natural radioactive disintegration series
(Chapter 48). As found in nature, lead is a mixture of two isotopes, Pb 206
and Pb 208 so that lead obtained from uranium ores has a density of 11.27 g/cm 3
as compared with 11.35 g/cm 3 for natural lead.

4. Properties of Lead, Lead is the softest, most malleable, and the
densest of the base metals. It is relatively inert and has good corrosion re-
sistance. When freshly cut the metal has a lustrous surface which soon -dulls
upon exposure to air due to superficial oxidation. A coating of oxide is first
formed which is converted into a protective film of basic carbonate,
Pb 3 (0H) 2 (C0 3 )2. Lead reacts with HC1 very slowly and cold concentrated
H 2 S0 4 has little effect due to the formation of a surface coating of insoluble
PbSO*, but HNO a attacks lead readily.

In contact with hard water lead becomes covered with a protective film
of insoluble salts, such as the sulfate, basic carbonate, or phosphate. Be-
cause they contain no dissolved salts, distilled water and rain water do not
form such a coat but the metal is attacked bv water with dissolved oxvcen.

610

The Elements of Group IV B: Germanium , Tin, and Lead

forming Pb(OH) 2 which is noticeably soluble. Since soluble lead compounds
are poisonous, lead pipes can be used to convey drinking water only when
the water is somewhat hard.

Two oxidation states are shown by lead, the lead (II) or plumbous and
the lead (IV) or plumbic , but almost all of its important compounds are in
the +2 state. Though lead is the most metallic of the Group IVA elements,
its formation of plumbite and plumbate anions indicates some nonmetallic
character. Most. lead compounds are insoluble, the most common soluble
salts being the nitrate and the acetate. Though the Pb 2+ ion is less hydrolyzed
than the Sn 2+ ion, to prevent precipitation of a basic salt through hydrolysis
some acid must be present in a Pb 2+ solution.

Four oxides of lead are known: the yellow monoxide or litharge , PbO;
the brown oxide, Pb0 2 ; the trioxide, Pb 2 0 3 ; and the orange-red oxide known
as red lead or minium, Pb 3 0 4 . The last two, Pb 2 O s and Pb 3 0,, may be viewed
as compounds of lead in which the lead atoms are present in both the +2
and +4 states. When lead is heated in air, PbO is formed. This oxide is
soluble in acid and also in alkali due to the formation of the plumbite ion,
Pb(OH) 4 2 ~ or HPbOr; it is used in the manufacture of lead glass and in
glazing pottery. The most important lead (IV) compound is Pb0 2 , obtained
as a brown powder by the oxidation of plumbite ion with hypochlorite ion in
basic solution.

(5) Pb(OH) 3 - + OC1- -> Pb0 2 + Ch + OH~ + H 2 0

In acid solution, Pb0 2 is a strong oxidizing agent and is used as one of the
electrodes of the lead storage battery. In concentrated alkali, Pb0 2 dissolves
to give plumbates, e.g., Pb(OH) 6 2 ~ or Pb0 3 2 ". By heating PbO to a tempera-
ture not above 550 °C Pb 3 0 4 is formed but above this temperature the red
lead decomposes to reform PbO and 0 2 . Red lead is employed in making
flint glass, ceramics, and in red paint for ironwork.

The hydroxide, Pb(OH) 2 , precipitates upon the mixing of solutions of
Pb 2+ and OH" ions. It dissolves somewhat in water to yield a faintly basic
solution and it is amphoteric. Lead (IV) hydroxide, Pb(OH)*, does not exist
Lead (II) chloride, PbCl 2 , is a white solid that is insoluble in cold water
but readily soluble in hot water, a property which distinguishes it and en-
ables its separation from two other insoluble white chlorides, AgCl and
Hg 2 Cl 2 . A similar behavior is shown by PbBr 2 but the solubility of yellow
Pbl 2 is quite independent of temperature. In solutions containing an excess
of halide ion the solubility of the lead halides is increased due to the forma-
tion of complex ions such as PbCl 4 2 ".

When alkali metal carbonates are added to solutions of Pb 2+ ion, the
basic lead carbonate, Pb 3 (0H) 2 (C0 3 ) 2 , is formed. This compound is im-
portant as the paint pigment, white lead . It is prepared commercially by the
action of air, carbon dioxide, and acetic acid upon lead metal. White lead
has excellent covering power but suffers from the disadvantage of darkening
when exposed to traces of H 2 S in the atmosphere. Black PbS is the most
insoluble of all lead compounds but’ ean be dissolved in HNO*.

The Elements of Group IVjB: Germanium ) Tin , and Lead

611

The oxidation potentials of lead are tabulated below.

0.13 v -1.46 v

In acid solution: Pb Pb 2 + Pb 4+

0.54 v -0.28 v

In basic solution: Pb Pb(OH) 4 2 " Pb0 2

QUESTIONS

1. Outline the steps in the metallurgy of germanium. Write chemical equations
where applicable. How is ultra-high purity germanium obtained?

2. What principles are involved in the process of zone refining?

3. What is meant by “polycrystalline” germanium? How is it converted to a
single crystal?

4. Explain (a) why germanium is a conductor but diamond is not and (b) why
the electric conductivity of germanium increases but that of lead decreases
with a rise in temperature.

5. How are the electrical properties of germanium changed by the introduction
of a trace of (a) arsenic and (b) indium? Distinguish between N-type and
P-type germanium; by what means does each conduct electricity? Which type,
if any, would result for the introduction of (a) boron (b) bismuth (c) silicon
(d) beryllium?

6. Outline the metallurgy of tin. Write pertinent chemical equations. How is
tin refined? What is the main use of the metal?

7. What are the allotropic forms of tin? Which form exists at room temperature?
Will the tin on a “tin can” in the freezer compartment of a refrigerator change
to alpha tin?

8. Explain the effectiveness of tin plate as a corrosion preventative.

9. Will 0.00IM SnCL spontaneously disproportionate to Sn and SnCl 4 ?

10. Which is the more basic, Sn(OH) 2 or Sn(OH) 4 ? Explain briefly.

11. Why is the tendency of Sn 2+ to hydrolyze reduced by the addition of acid?

12. Write chemical equations for the metallurgy of lead. How does it differ from
that of tin? How can lead be purified? What is the major use of lead metal?

13. Write balanced chemical equations for the following: (a) Sn and NaOH
(b) Sn and dilute HN0 3 (c) SnCL and Hg 2 + (d) SnCl 2 and FeCl 3 (e) SnS
and S 2 ' (f) PbO and NaOH.

14. Why should not excess concentrated HC1 be used to precipitate PbCL?

15. An alloy contains Ag, Pb, Sn, and Cu. How could it be brought into solution
and by what chemical means could each ion then be detected and confirmed?

16. What is the oxidation state of lead in Pb 2 0 3 ?

17. What weight of white lead can be prepared from 100 grams of lead?

The Elements of Group O
The Noble Gases

The gases, helium, neon, argon, kryton, xenon, and radon, constitute
the elements of Group O and are known as the noble gases . Their electron
configurations are listed below and their properties are given in Table 47-A.

Element

Atomic

Number

3s 3 p 3d

4s 4 p 4d 4f

5s op 5 d

6s 6 p

Helium

Neon

2 6

Argon

2 6

Krypton

2 6

2 6 10

2 6

Xenon

2 6

2 6 10

2 6

Radon

2 6

2 6 10

2 6

1. General Properties of the Noble Gases. Each horizontal series of the
Periodic Table is terminated by a Group 0 element. In their outermost prin-

cipal energy levels the Group 0 elements have eight electrons, ns 1 2 np°, except
for helium in which this shell has a maximum capacity of two electrons, Ir.
As we have noted frequently such an octet configuration has a minimum
energy and unusual stability, and is characterized by an almost complete
chemical inertness. The molecules of the noble gas elements are monatomic.
The low boiling points of the Group 0 elements and the narrow ranges
over which their liquid states exist are also a reflection of the "com-
pleteness” of the electric fields surrounding these atoms and of the low
attractive forces between them. The atoms of the Group 0 elements have
no electron affinity and have high ionization potentials, higher than those
of the alkali metal or halogen elements. However, the first ionization po-
tentials of the larger elements, xenon and radon, are lower than those of chlorine
and nitrogen*

The Elements of Group O: The Noble Gases

613

Table 47-A

Properties of the Elements of Group 0

Property

Helium

Neon

Argon

Krypton

Xenon

Radon *

Symbol

Atomic Number

Atomic Weight

4.0026

20.183

39.948

83.80

131.30

222

Isotopes (mass numbers

1 3( 0.00013)

20(90.51)

36( 0.337)

78( 0,35)

124( 0.10)

220

and percent)

4(99.99987)

21( 0.28)

38( 0.063)

80( 2.27)

126( 0.10)

222

22( 9.21)

40(99.60)

82(11.56)

128( 1.91)

83(11.55)

129(26.24)

84(56.90)

130( 4.05)

86(17.37)

131(21.24)

132(28.93)

134(10.52)
136( 8.91)

Abundance in Earth s

5 X 10-4

0.0018

0.932

1 X 10-4

8 x 10-o

trace

Atmosphere, % by volume

Physical State at STP

all

are colorless gases

Density at STP, g/I

0.1785

1.784

1 3.708 1

5.85

9.73

Density of Liquid, g/ml**

0.126

1.40

3.06

4.40

Melting Point, °C

-27 1.9***

-248.7

-189.3

-157.0 i

-111.5

-71

Boiling Point, °C

-268.9

-245.9

-185.7

-152.9

-107.1

-62

Critical Temperature, °C

-267.9

-229

-122

-63

-17

104.5

Critical Pressure, atm

2.26

26.9

48.0

54.3

58.2

62.4

Ionization Potential eV, 1st

24.46

21.45

15.68

13.93

12.08

10.75

2nd

54.14

41.07

27.76

24.56

21.21

3rd

40.90

36.8

32.0

Atomic Radius, A

0.93

1.60

1.91

2.20

♦Radioactive element
♦♦At the boiling point
♦♦♦At 26 atmospheres

2, Discovery of the Noble Gases. The discovery of the noble gases is a
romance of chemical science. When he formulated his Periodic Table in
1869, Mendelejeff had no knowledge of these as yet undiscovered elements.
As early as 1785, Henry Cavendish observed that when he tried to combine
all the nitrogen in a given sample of air with additional oxygen, a small
amount of gas which would not combine remained. He estimated this residual
amount to be about 1/120 of the original air. This observation, foreshadowing
the discovery of the noble gases in the atmosphere, was disregarded by
scientists for over 100 years.

Argon: In 1894, the English physicist, John William Strutt, the third Lord
Rayleigh, observed that nitrogen prepared by removing all other known
gases from air was denser than pure nitrogen prepared chemically from am-
monia. The difference in densities, 1.2572 'g/1 for atmospheric nitrogen as
compared with 1.2506 g/1 for chemical nitrogen was significant in that it was
beyond the range of experimental error. In collaboration with the English
chemist. Sir William Ramsay, upon whom he called to hek) exnlain this

6X4

The Elements of Group O: The Noble Gases

baffling discrepancy, Rayleigh passed atmospheric nitrogen over hot mag-
nesium, which combines with N 2 to form a solid nitride, Mg 3 N 2 . Again there
was a small quantity of residual gas, 0.933% by volume of the original air,
which would not combine with any other element. Ramsay showed that
this gas emitted a spectrum distinct- from that of any other element and
hence was itself a new element. Because of its chemical inactivity the
element was named argon (Greek, argos: inactive). Actually the crude
argon obtained from air also contained small amounts of the gases neon,
kryton, and xenon but these were as yet unknown to the experimenters.

Helium : Through the use of the spectroscope the element helium was
observed first in the sun s chromosphere more than 25 years before it was
recognized on earth. During a total eclipse of the sun in 1868, P. J. C. Janssen
in India and J. N. Lockyer in England directed spectroscopes towards the
suns chromosphere and observed a yellow line at 5876 A which could not
be identified with any element then known on earth. To it Lockyer ascribed
a new element which he named helium (Greek, hellos : the sun). In 1889,
W. F. Hillebrand, an American chemist of the U. S. Geological Survey, ob-
served that certain uranites, minerals containing uranium, evolved a gas
when treated with acid. Hillebrand believed the gas to be nitrogen so it
remained for Sir William Ramsay to disclose definitively the presence of
terrestrial helium. After the discovery of argon in the atmosphere Ramsay
looked for other sources of the element. Hearing of Hillebrand’s experiments
he suspected that the gas might be argon and obtaining a sample of the
uranium bearing mineral, cleveite, from Hillebrand in 1894, repeated the
latter s experiments and subjected the purified gas to spectral analysis. Much
to Ramsay's surprise the spectrum was different from that of argon. When
the spectrum was compared with that of solar helium and found to be identi-
cal, terrestrial helium was thereby discovered. Later Ramsay showed that
traces of helium are present in the earth’s atmosphere.

Neon (Greek, neos : new); Krypton (Greek, kryptos: hidden); Xenon
(Greek, xenos: stranger): With the discovery of argon and helium came the
conviction that a new family of inert elements existed which would fit into
the Periodic Table between the halogen family and the alkali metals. With
the aid of his assistant, M. W. Travers, Ramsay continued to search for other
inert gases. In vain they sought to obtain these gases by heating rare
minerals; their efforts were then directed towards liquid air which, fortunately,
became available at that time. By careful fractional evaporation of liquid
air and by employing the spectroscope to detect minute quantities of sub-
stances, Ramsay and Travers isolated three new gaseous elements in 1898-
krypton first, then neon by fractionation of an argon residue and, finally,
xenon by fractionation of a krypton residue.

Radon : In 1900, two years after the discovery of radium by the Curies,
the German physicist, F. E. Dorn, showed that one of the radioactive dis-
integration products of radium is a gas. The heaviest gaseous element known,
it was first called ‘emanation of radium” by Rutherford but is now known
as radon.

The Elements of Group O: The Noble Gases

615

3. Properties of the Noble Gases. Except for helium and radon, the noble
gases are prepared by the fractional distillation of liquid air. Neon has a
relatively high electrical conductivity, about 75 times that of air, and is used
in the familiar orange-red electric signs. The passage of an electric current
through a tube filled with neon at low pressure ionizes the neon atoms
and results in the emission of its characteristic radiation. Neon lamps are
used as beacon lights because they cost less to operate than do incandescent
lamps and give better penetration through fog. The color of the emitted
light can be modified by mixing argon or mercury vapor with the neon
in tubes made of various colored glass. Argon is used to fill ordinary in-
candescent lamps at low pressure. The inert gas reduces the rate of evapora-
tion of the tungsten filament, thereby increasing the life of the lamp.
Because argon is a poor conductor of heat, poorer than nitrogen which was
formerly used for this purpose, less heat is lost and the lamp can be operated
at higher temperatures and efficiency. Krypton and xenon have no extensive
uses. An electric dicdiarge passed through a mixture of krypton and xenon
produces an intense light of short duration, about 1/50,000 second, and
such lamps have been employed for taking high speed photographic ex-
posures. Radon is radioactive and is utilized in radiotherapy for the treat-
ment of malignant growths. For this purpose it is collected from radium
salts and sealed in small tubes which are placed, near the tumor.

Helium is the most widely used of the noble gases. It occurs as a minor
constituent of natural gas, principally in two gas fields extending from
western Kansas through Oklahoma and into the Texas panhandle. Gas from
these fields normally contains 0.3- 1.0% helium and it is estimated that over
90% of the Free World’s helium is contained in these gas deposits. Because
of its strategic applications all helium production is under the control of the
federal government and, with the possible exception of the U.S.S.R., the
element is produced by no other country. Helium is extracted from natural
gas by a low temperature gas liquefaction process, developed at the U. S.
Bureau of Mines, which utilizes the extremely low boiling point of the ele-
ment; the product is 99.995% pure helium. About 500 million cubic feet of
the gas are produced annually at a cost of about one cent per cubic foot.
Helium is also a disintegration product of those radioactive nuclides which
yield alpha particles and is found, possibly in a state of solid solution, in all
minerals containing uranium and thorium. One gram of radium yields 0.11 ml
of helium per year.

Before World War II, helium was fused principally as a lifting gas for
lighter than air aircraft, Although its density is twice that of hydrogen the
lifting power of helium is 92% that of hydrogen. This follows from fine fact
that the lifting power of a gas whose density is less than that of air is
due to the difference between its weight and that of an equal volume of
air. Today helium is still used to lift scientific research balloons into the
upper atmosphere. Toward the end of World War II, helium shielded arc
welding was developed and now provides the major use of the gas. Helium
is used as an inert gaseous shield surrounding an electrode in arc welding
without a flux. Before this technique was perfected, some types of metals

616

The Elements of Group O; The Noble Gases

could not be joined directly to pieces of their own kind but only by being
welded to a different metal placed between. Helium also finds use as an
inert atmosphere during the heat treatment of various metals. Because of
its ability to diffuse through microscopic openings helium is used to detect
minute leaks in apparatus requiring perfect seals. Mixed with oxygen, helium
provides a breathing atmosphere for persons suffering from respiratory ail-
ments. The helium-oxygen mixture is much lighter than air and can flow
through restricted respiratory passages more readily and so less muscular
effort is required for breathing. Such mixtures are also used by divers and
caisson workers; when helium is substituted for atmospheric nitrogen, the
pathological condition known as the "bends” is inhibited since the absence
of nitrogen precludes the formation of nitrogen bubbles in the blood when
the pressure on the worker is decreased. No helium bubbles are formed
because of the relatively low solubility of helium in body fluids even at
high pressures; at 25 °C helium is 2 Y 2 times less soluble than nitrogen in water.

The largest single user of helium is the National Aeronautical and Space
Administration (NASA). Helium is used to force liquid fuels into the rocket
engines of space vehicles, to purge the engines of impurities, and to pres-
surize the thin steel envelope and empty fuel tanks of a missile in space.

Helium is an element with many extreme properties. It has a lower
solubility and a lower refractive index than any other gas. It diffuses more
rapidly, conducts heat better, transmits sound at higher velocity than any
gas except hydrogen, and conducts electricity better than any gas except
neon. Of all gases it is the most difficult to liquefy and, under its own vapor
pressure alone, remains liquid down to absolute zero. The solid can be formed
but only by die application of considerable pressure.

Between its critical pressure, 5.20°K and 2.18°K, liquid helium, specifically
the isotope He 4 , Shows no unusual properties other than a slight decrease
in viscosity with decreasing temperature. But at 2.18°K, a temperature known
as the lambda point, there occurs an abrupt transition to a new liquid phase
which has several unique properties. This phase is known as Helium II to
distinguish it from Helium I, the ordinary liquid phase stable above 2.18°K.
Unlike the usual type of phase change, that between the liquid helium phases
takes place without a latent heat or a change in density. Helium II is termed
a “superfluid As its temperature is lowered, its viscosity becomes vanishingly
small in that it can pass through the finest of .capillaries with little application
of pressure. Though an electrical insulator, Helium II is the best known con-
ductor of heat, its heat conductivity being 3 X 10 7 times greater than that
of Helium I, one degree higher in temperature and 2,000 times greater
than that of copper at room temperature. Above the lambda point, liquid
helium boils continually in, the normal manner but below this temperature
all bubbling ceases because heat is transported rapidly to the surface where
only evaporation takes place. The conduction of heat in Helium II appears
to take place in pulses analogous to the pressure wave of sound, a new
phenomenon to which has been given the name “second sound/' Another
curious characteristic of Helium II is its mobile film that apparently defies
gravity. Thus, a helium film flows up over the rim of an empty beaker

The Elements of Group O: The Noble Gases

617

partially immersed in a Helium II bath until the levels inside and outside
the beaker are equal If the beaker is then raised, the helium again climbs
out of the beaker down to the surrounding surface. Oddly, the He 3 isotope
shows none of these superfluid properties.

Despite the fact that X-ray studies indicate no crystalline array of the
atoms in Helium II, an explanation advanced for its peculiar behavior is that
its atoms are perfectly ordered and that Helium II has zero entropy even
though in the liquid state. At one or two degrees above absolute zero the
Helium II atoms are in their lowest possible energy levels where they have
no translational energy but solely vibrational energy, the so-called zero
point energy . This condition, coupled with the low attractive forces between
and the low density of helium atoms, gives rise to the superfluid.

4. Cryogenics. Temperatures in the neighborhood of absolute zero have
been achieved by expansion of cooled helium at high pressure much in the
same manner as the production of liquid air. By this technique temperatures
as low as 0.8° K have been attained and through the use of the magnetic
properties of the nucleus, a temperature as low as 0.00001°K has been
reached. The ability to produce temperatures within a few degrees of
absolute zero has given birth to the branch of science known as cryogenics,
the study of matter and its properties at very low temperatures. At these
extreme temperatures we can study matter shorn of thermal chaos and the
random motion of molecules. One of the most impressive of cryogenic phenom-
ena is superconductivity, discovered in 1911 by the Dutch physicist, H.
Kamerlingh-Onnes, who first liquefied helium. At temperatures near absolute
zero certain metals, notably those in the center of the Periodic Table, lose
all measurable resistance to the flow of electricity. Indeed, a current induced
in such a metal in the shape of a doughnut has persisted without any
noticeable loss in magnitude for years. The temperature at which a metal
becomes superconductive depends upon the individual metal; aluminum
attains superconductivity at 1.14°K while niobium does so at the relatively
high temperature of 9.22°K. Compounds can also become superconductors;
niobium nitride, NbN, does so at about 15°K.

Cryogenic applications are not lacking. The cryotron, a tiny strip of
tantalum around which a fine niobium wire is wrapped becomes a super-
conductor in a bath of liquid helium and can perform the functions of
electronic tubes. Molecular fragments, or free radicals, normally having a
life span of but a few thousandths of a second in chemical reactions, can
be trapped and studied more leisurely by freezing their motion in liquid
helium for an indefinite extension of their life spans. MASER, or microwave
amplification by stimulated emission of radiation, and other amplifiers plagued
by noise, operating in liquid helium baths gain about 1,000 times in amplifica-
tion over conventional operation.

5. Chemical Properties of the Noble Gases. Because of the stability of
the octet configuration it was believed that the Group 0 elements were
incapable of forming electrovalent or covalent compounds in the ordinary

618

The Elements of Group O: The Noble Gases

sense and few serious attempts were made to prepare any such compounds.
But in 1962, fluorides of xenon and radon were produced by ordinary
chemical techniques. The first true noble gas compound, Xe(PtF 8 ), was
prepared by the interaction of xenon and the powerful oxidizing agent,
platinum hexafluoride, PtF 6 . Thereafter a difluoride, XeF y , a tetrafluoride,
XeF 4 , and a hexafluoride, XeF 0 , were prepared by direct combination of
the elements. All are colorless, crystalline solids, stable at room temperature.
The tetrafluoride, XeF 4 , is made by heating one volume of xenon and five
volumes of fluorine in a closed container to 400° C and then quenching in
cold water. The compounds are soluble in HF but are hydrolyzed by water.
Hydrolysis of XeF 4 and XeF 6 forms Xe0 3> a white crystalline solid, which
is unstable and explodes when heated gently or even when rubbed. A radon
fluoride, RnF 2 , has been prepared by heating a radon-fluorine mixture to
400°C and also a krypton fluoride, KrF 4 , by passing an electric discharge
through a mixture of one volume of krypton and two volumes of fluorine in
a container cooled to liquid air temperatures. As might be expected, KrF*
is less stable thermally than XeF 4 ; at 20° C it decomposes into its elements
at a rate of about 10% per hour.

So far binary compounds of only the larger noble gas atoms have been
prepared. These compounds show even numbered oxidation states, as would
be expected from elements of Group 0. The nature of the bonding in the
noble gas compounds is presently a matter of extensive discussion. The chemi-
cal reactivity of xenon and radon may be due to the availability of unoccupied
4/ or 5 d orbitals, not much higher in energy than the completed outermost
shell. In the electric field created by the extremely electronegative fluorine
atom, an electron from each of two complete orbitals of the xenon atom
may be promoted to orbitals of higher energy. Thereby four half-filled
orbitals would be created and four covalent bonds could be formed, as in
XeF 4 . Attempts to prepare simple oxides and chlorides have so far failed but
the existence of XeOF 4 and XeOF M is indicated by mass spectrometry. The
larger inert gas atoms, argon, krypton, and xenon, also form hydrates analogous
to the hydrates of other gaseous elements, e.g., Kr * 6 HX). Conceivably, too,
the noble gas atoms could form covalent bonds with electron deficient
molecules. Because a sufficiently potent source of energy can remove electrons
from a noble gas atom, certain helides or compounds of helium with other
elements have been formed in an electron discharge tube. The species,
He 2 ~, HeH", and HeH 2 ~, have been so produced but these have only transi-
tory existence with life periods of the order of 10~ s second.

QUESTIONS

1. To what is the inactivity of the noble gases due? Do these elements have
unoccupied orbitals in any principal quantum levels?

2. Discuss briefly the occurrence and the discovery of the noble gases. For
what reasons were these elements discovered so much later than less abundant
elements?

3. Prove that helium has 92% of the lifting power of hydrogen.

4. List some uses of helium gas and liquid.

The Elements of Group 0: The Noble Gases

619

5. What are the properties of helium that make it useful for the detection of
leaks? In a test run, how would you detect any leakage of helium?

6. Account for the relatively high electrical conductivity of neon.

7. How are temperatures in the neighborhood of absolute zero produced? What
is the lowest temperature so far obtained?

8. Distinguish between Helium I and Helium II. Distinguish between He 4 and
He*. What unusual properties does helium show at 1°K?

9. Is there any limiting factor, other than absolute zero, to the lowest temperature
which might be attained?

10. Define (a) cryogenics (b) cryotron (c) superconductivity.

11. Why should the larger noble gas atoms react more readily with fluorine than
might neon or helium?

12. What orbitals could be involved in the formation of (a) XeF 2 (b) XeF 4
(c) RnF 4 (d) KrF 4 ?

L3. Would it be theoretically possible for argon to form an addition compound
with BF 3 ?

Nuclear Chemistry I

We have seen that the discovery of radioactivity was of prime importance
in formulating the theory of atomic and nuclear structure. All elements are
not naturally radioactive. The property of radioactivity, one of the few which
is not periodic, is restricted to the heaviest elements only, from lead to
lawrencium. From naturally radioactive species only three types of radiation
are emitted, namely, the alpha particle, the beta particle, and the gamma
ray. The properties of these, and of other subatomic particles we shall discuss
later, are given in Table 48- A.

1. The Group Displacement Rule. Before we inquire as to what change,
if any, occurs upon the emission of a particle or ray, it must be emphasized
that the phenomenon of radioactivity is a property of the nucleus of an atom.
All radioactive emanations come from the nucleus and any changes which
occur take place in the nucleus of the atom; subsequently the electron con-
figuration of the atom is modified to conform to the nuclear charge. In the
same manner as nonradioactive species, radioactive elements have chemical
properties which depend upon their electron configurations. But though a
radioactive element can combine chemically with different elements to
form different compounds, in these different chemical compounds the same
radioactive nucleus exhibits identical radioactive properties.

The nucleus, it will be recalled, is composed of protons and neutrons.
The former are present in number equal to the atomic number and so deter-
mine the chemical nature of the nuclide . The term nuclide is used for a
species of atom characterized by the composition of its nucleus, that is, the
numbers of protons and neutrons therein. Also, the entire mass of the atom is,
for all practical purposes, concentrated in the nucleus since the external elec-
trons have but negligible mass. Hence by the emission of an alpha particle
(charge -f*2; mass 4 amu) from a radioactive nucleus, there remains a nuclide
which is two less in positive charge or atomic number and four less in mass
than the parent nuclide. Such emission therefore results in the formation of a
new element! With the ejection of a beta particle (charge -1; mass 0 anrnb the

♦One atomic mass unit (airni) = 1.66 X 10~ 24 gram

♦♦One electron volt (eV) — 1.6021 X 10~ 12 erg/particle = 23061 cal/mole

622

'Nuclear Chemistry 1

resulting nuclide is one greater in atomic number but has practically the
same mass. The emission of a gamma ray only leaves the radioactive nuclide
unchanged. It is evident that the nature of the resultant nuclide can be ob-
tained by the application of ordinary algebra, that is, by taking the algebraic
difference between the numerical values of charge and mass of the parent
nuclide and the emitted particle, as indicated in Table 48-B. These algebraic
statements are known as the Group Displacement Rule. Similar algebraic
rules will hold later in determining the nature of the resulting nuclide upon
the emission of a nucleus of particles other than alpha and beta. With few
exceptions, a given radioactive nuclide emits only one particle at a time, either
alpha or beta, but gamma rays can be emitted in conjunction with alpha or
beta emission. The latter arise from the transition of the daughter element
from an excited to a lower nuclear energy level.

Table 48-B

The Group Displacement Rule

Alpha emission:

Beta emission:

Gamma emission:

Z X A 2 He 4 + z- 2 Y a ~*

Z X A -> ml efi + z+ iY A

Z X A hv + Z Y A

where X = the initial radio-
active nuclide
and Y = the resultant
nuclide

The radioactive decay of uranium, 02 U 238 , to lead offers a good illustration
of the Group Displacement Rule. The steps in this process are listed in
Table 48-C and are represented graphically in Figure 48.1. Uranium, an
alpha emitter, is the first in a series of radioactive elements which gives off
alpha and beta particles with the ultimate formation of lead. Two other
naturally radioactive series are known, the thorium and the actinium series,
so named for their initial members, but the end product of all naturally
radioactive series is a nonradioactive isotope of lead. A fourth radioactive
series, the neptunium series, composed entirely of man-made isotopic species
not found in nature, was discovered during the period of intense nuclear re-
search during World War II. The end product of this series is bismuth.
Because of the characteristics of alpha and beta particles, the mass numbers
of all members of a given series leave the same remainder when divided by
four. If m represents an integral value, the four radioactive series can then
be classified according to the mass numbers of their members:

thorium series
neptunium series
uranium series
actinium series

= 4 m + 0
= 4m-f 1
= 4m + 2
= 4 m + 3

Nuclear Chemistry I

623

Atomic number

Figure 48 J. The Decay Path of IP 8 to Pb 206 .

Though a given nuclide emits only one specific particle, a sample of uranium,
as an example, will be found to give off alpha, beta, and gamma emission
simultaneously because it contains a number of different radioactive nuclides
from its disintegration series.

2. Kinetics of Radioactive Transformation. Different radioactive elements
emit particles, or decay, at different rates. The rate of radioactive decay follows
the same mathematical relationships as does a first order chemical reaction
in that a given fraction of the atoms disintegrate in the same time interval
Thus the number of radioactive atoms still remaining at any time, anc

624

Nuclear Chemistry 1

hence the activity of a radioactive substance, decrease exponentially with
time as given by

(i)

N = No e~ Xt

where N is the number of radioaetive atoms still present at a time, t, from
a sample in which the initial number of atoms was N 0 ; X is a radioactive
decay constant, specific for each radioactive nuclide, and corresponds to the
reaction rate constant of a chemical reaction.

The half life, tj., of a radioactive nuclide, that is, the length of time for
the radioactive intensity and hence the mass of a given quantity of an element
to decrease to half the value present at the start of any arbitrarily designated
period is

( 2 )

0.693

The rate of disintegration, which is indicated by the value of either X or

ti , is an important characteristic property of a nuclide; values for those
*2

nuclides in the uranium series are given in Table 48-C. Half-life periods vary

Table 48-C

The Uranium Series

Element

Atomic

Number

Mass

Number

Particle

Emitted

Half-life period

Uranium

238

alpha

4.51 * 10® yr

Thorium

.90

234

, ' beta

24.1 day

Protoactinium

234

beta

1.18 min

Uranium

234

alpha

2.48 x 10* yr

Thorium

230

alpha

8.0 X 10 4 yr

Radium

226

alpha

1620 yr

Radon

222

alpha

3.82 day

Polonium i

218

alpha

3.05 min

Lead

214

beta

26.8 min

Bismuth

214

beta

19.7 min

Polonium

214

alpha

1.6 x 10- 4 sec

Lead

210

beta

19.4 yr

Bismuth

210

beta

5.0 day

Polonium

210

alpha

138 day

Lead

206

Stable

Nuclear Chemistry I

625

widely for radioactive elements, from billions of years to minute fractions
of a second. The shorter the half-life period, the more radioactive is the
element. The rate of radioactive decay is independent of any conditions
of temperature or pressure which man can impress upon radioactive elements.

If we consider only a single disintegration step, we might ask how long
it would take for a given quantity of a nuclide to be transformed completely
into its decomposition product. This question is mathematically unanswerable!
Why any specific radioactive nucleus should undergo radioactive emis-
sion at a certain instant in time is not yet known, and so it is impossible to
predict when a given nucleus will emit a particle. In such circumstances
when the laws governing the behavior of the individual are unknown, only
the methods of statistical analysis are applicable. When we deal with a
large number of radioactive nuclei, accurate predictions can be made con-
cerning the time when a given fraction of them will have undergone radio-
active transformation or decay, but again we can make no prediction con-
cerning which of the atoms, even if they could be labelled and identified
one from the other, will be in this fraction. Similar considerations apply to
actuarial statistics. The time of death of a specific individual, or even when
the last individual of a group will die, can hardly be predicted with accuracy]
But with a group of people it is possible to predict when a certain fraction
will have died, and the larger the group the more accurate is the prediction.
The financial well-being of life insurance companies attests to the practical
verification of such calculations. The accuracy of such actuarial calculations
is directly proportional to the number of individuals, and in radioactivity
to the number of atoms in the sample. Because of the extremely small masses
of atoms, even a minute quantity of material contains an enormous number of
atoms. A microgram of uranium contains more than 10 15 atoms, a number ap-
proximately one million times greater than the entire population of the earth.
Even if we were to calculate the length of time for 99.9999% of a gram of
uranium to undergo radioactive disintegration, at the end of this period the
remaining 0.0001% would still contain approximately 10 15 atoms of uranium.
Consequently calculations concerning the time when any given fraction, indeed
practically all, of a radioactive substance will have disintegrated can be made
with extreme accuracy. But not ALL since the last minute fraction of a
percent, as from 99.9999% to 100% must pass through the numerical region
where only a small number of atoms remain, and here the statistical laws
governing radioactivity fail because they do not hold for small numbers.
More generally, the equations for reaction kinetics, both chemical and
nuclear, are fundamentally statistical and are valid only when, and be-
cause, they deal with large numbers of particles.

The half-life period of radium is 1620 years. Thus, if we had one gram of
radium now, in 1620 years half a gram would remain. After the next 1620 years,
one-half of the quantity left after the first 1620 years would have decayed.,
so that one-quarter gram of radium would still be left. Such additional cal-
culations could be continued with accuracy to a point where the mass of the
residual radium contained so few atoms that the statistical laws on which
these calculations are based are no longer valid.

626

'Nuclear Chemistry 1

Example I: The half life of radon, Rn 2 --, is 3.82 days. Starting with a sample
of 0.800 g of pure radon, how long will it be before 0.500 g remains?

Solution-. From N = N„ e- * ‘ there can be derived for first order reaction
kinetics,

, N _ _X , , t 0-691

l0gl ° No ” " 2.303 1 d H A.

Since the number of radon atoms, N, is proportional to its mass, m, then

, m A .

l0gl ° m 0 “ " 2.303 *

The value of the decay constant, A, is found first.

A =
Then

0.691

0.691
3.82 days

= 0.181 day- 1

2.303 , m _ 2.303 , 0.500

“ A~ Sl ° "57““" 0.181 day" 1 ” gl ° 0.800

t = 2.59 days

3. Detection of Radioactivity. Many devices have been used to detect
radioactive emission. Among them are: A) the photographic image; B) the
scintillation counter; C) the electroscope; D) the cloud chamber; E) the
bubble chamber; F) the ionization chamber; and G) the Geiger-Miiller counter.

Except for the scintillation counter and perhaps the photographic process,
these techniques depend upon the formation of ionized species during the
passage of the radioactive emanation through matter. When an energetic,
rapidly moving, charged particle, such as an alpha or beta particle, passes
through a gas its kinetic energy may be sufficient to eject electrons from the
atoms or molecules with which it collides. The positive ions thereby produced
and the ejected electrons constitute gaseous ion-pairs which make the gas
through which the nuclear radiation has passed an electrical conductor. In
a given medium, the total number of ion-pairs formed depends primarily on
the energy of the moving particles. To form a single ion-pair requires the
expenditure of approximately 33.5 electron volts on the part of a moving
particle. Because of their relatively large mass and charge, alpha particles
produce the greatest number of ion-pairs per centimeter of path, about
50,000-100,000 in air under ordinary conditions. Under the same conditions,
beta particles have a lesser ionizing power, forming only a few hundred ion-
pairs, but their paths are much longer than those of alpha particles. On the
other hand, beta particles have a greater penetrating power than do alpha
particles which can be stopped by a few sheets of paper. Hence an alpha
emitter could be held in a glass bottle with impunity but, taken internally,
despite the small distance through which an alpha particle passes the result-
ing high degree of ionization produced in tissue is potentially very harmful.
Because neutrons are uncharged, they do not cause primary ionization
through collision in the manner of alpha and beta particles. Gamma rays,

Nuclear Chemistry I

627

akin to X-rays and ultraviolet radiation, may also have energies sufficient
to eject orbital electrons. In turn primary electrons, produced initially by any
of these incident particles or radiation, may have enough energy to cause
secondary ionization upon their collision with other molecules.

The oldest technique for the detection of radioactivity, in fact the manner
by which it was discovered, is the photographic image. When incident upon
a photographic emulsion, radioactive emanations initiate the same process
as does visible light upon a film or plate in a camera. Upon development,
a blackening occurs where the film was struck by a radioactively emitted
particle and the intensity of such blackening is proportional to the radio-
activity of the source. By using layers of stripped film, the continuous track
of an individual particle can be determined. Important radio-uses of photog-
raphy are in safety badges and in radioautographs, whereby a radioactive
object is placed directly upon film.

Certain substances emit visible and ultraviolet light when struck by
nuclear radiation, much as the coating on the inside face of a cathode ray
tube glows when electrons fall upon it. Two such common scintillators or
phosphors are zinc sulfide and naphthalene crystals. In his classic alpha
particle scattering experiments, the phenomenon of scintillation was first used
by Rutherford to count the numbers of such particles. Today instruments
known as scintillation counters transform the light produced in the scintilla-
tion process into an electric current by means of photoelectric cells and the
magnitude of such current is a measure of the activity of the radioactive
source.

In its oldest form, the electroscope consisted of two thin leaves of metal,
usually gold, suspended from a metallic rod which was insulated from its
surroundings. When a static electric charge was placed upon the rod, it dis-
tributed itself uniformly on the leaves which then separated at their free
ends due to the mutual repulsion of like charges. When a radioactive sub-
stance was brought near, the air surrounding the leaves was ionized. The
resulting ion-pairs discharged the leaves, causing them to collapse to their
uncharged state and the rate of fall of the leaves measured the intensity
of the radioactive source. In present-day electroscopes the gold leaves have
been replaced by quartz fibers rendered electrically conducting by a coating
of graphite, but the operating principle remains the same.

The cloud chamber ; invented by the Scottish physicist, C. T. R. Wilson,
consists of a piston and cylinder within which is dust-free air saturated
with water vapor. If the piston is pulled down rapidly, the resulting adiabatic
expansion cools the air in the cylinder. Under these conditions the air con-
tains more water vapor than is required to saturate it. But a further requisite
for “dew” to settle is the presence of an object to act as a center upon which
condensation of water vapor can take place. Without such “condensation
nuclei” the air will remain supercooled and will not condense. In normal
atmospheric air, dust particles act as such centers for the formation of a cloud
or mist. Wilson found that gaseous ions can also function as centers around
which water vapor will condense. Thus, if nuclear radiation passes through
the supercooled air of a cloud chamber, condensation occurs upon the trail
of ions produced by the radiation as a close array of fine droplets which

628

Nuclear Chemistry I

has the appearance of a track. This track is the path of an individual particle
and is the closest that man has come to the direct observance of an elementary
particle. Because of their high ionizing power, alpha particles produce heavy,
straight tracks; beta particles yield thin tracks, and gamma rays very faint
tracks. Cloud chamber tracks of alpha * and beta particles are shown in
Figure 48.2.

Alpha particle tracks
S is the radioactive source of the
alpha particles.

The branched path is due to the
collision of an alpha particle with
an atomic nucleus.

Beta particle tracks
The curvature of the tracks is due
to a magnetic field which has been
impressed around the chamber.

Figure 48.2, Tracks of Alpha and Beta Particles in a Cloud Chamber.

Invented by the American physicist, D. A. Glaser, the bubble chamber is
analogous in principle to the cloud chamber in that ions can also act as
centers for the formation of a track of small bubbles of vapor during the
passage of a particle through a superheated liquid. Liquid hydrogen is com-
monly used. Bubble chambers are far more sensitive than cloud chambers
and have been specially useful in the study of paths of high energy particles.

The ionization chamber is a partially evacuated tube containing two metal
electrodes across which a high difference of electric potential is placed. The
tube is filled with an ionizable gas at low pressure. The electrodes may
be parallel plates but usually the positive electrode is a wire running axially
through the center of the tube and the negative electrode is a thin sheet
of metal along the inner surface of the tube. A simplified schematic diagram
of an ionization chamber is drawn in Figure 48.3. Under normal conditions
the chamber is nonconducting but the passage of nuclear radiation through
the gas in the tube produces ions which are attracted to the oppositely charged
electrodes and are recorded as a current on a meter. The radioactive inten-
sity is indicated by the magnitude of this current.

Nuclear Chemistry I

T = chamber tube
C == cathode
A = anode
I = insulator
M = meter
W = thin window for
radiation entry

The meter indicates the magnitude of the ionization current. In practice, the
magnitude of this current is so small that a vacuum tube amplifier is used in place
of M and the amplified current is read on a special meter.

Figure 48.3. An Ionization Chamber.

The Geiger-Muller counter is similar in construction to the ionization
chamber but is operated with a much higher voltage across the electrodes.
At this higher potential the ion-pairs produced by the entry of a single
particle into the tube are accelerated to such an extent on their way to the
electrodes that their kinetic energies are sufficient to cause secondary ioniza-
tion by collision. This internal gas amplification, as high as ID 8 , produces
an avalanche of ions which results in a pulse of current. Individual particles
emitted by a radioactive source can thereby be counted as separate pulses.
The intensity of a radioactive source is measured by the number of pulses
in a given time, transformed instrumentally into flashes of a neon bulb, clicks
on a loud speaker, or recorded by a counting device. The Geiger-Muller
counter is the most commonly used instrument for the detection of nuclear
radiations, particularly beta and gamma.

4. Properties of the Nucleus. Scientific research has uncovered many em-
pirical facts concerning the atomic nucleus, but to date no theory of the
nucleus and its behavior has been universally acceptable. It might be said
that nuclear theory now occupies a status similar to that of atomic theory
prior to the advent of John Dalton. Dalton had available certain quantitative
facts concerning chemical reactions upon which he based his theory but it
remains to weld the facts concerning the nucleus into a coherent theory.

First the very stability of the nucleus must be explained. Classical theory
would indicate that the protons within the nucleus should fly apart due to
their mutual repulsion. Yet obviously they do not! Somehow, the neutrons
within the nucleus and, it is believed, particles called mesons intermediate in
mass between electrons and protons, exert a stabilizing influence.

The stability of a specific nucleus can be calculated in terms of the binding
energy which holds it together. If the mass of a given nucleus is calculated
by summing up the masses of its separate nucleons, the value thus obtained
is always greater than the actual nuclear mass. For example, the total weight
of the individual nucleons in a helium nucleus is (2xf*00728) + (2X1*00,867),
or 4.0319 amu, but the actual mass of a helium nucleus is observed to be
somewhat less, 4.0015 amu. This difference in mass is known as the mass

630

Nuclear Chemistry 1

defect of the nucleus. A mass defect is always found, that is, mass is always
lost in the formation of a stable nucleus from its component nucleons. In the
calculation of a mass defect, precise values of the masses of the proton
and the neutron must be used because the actual magnitude of the mass
defect, which is obtained by difference, is never very great.

This apparent discrepancy in mass, or the mass defect, which seems to
contradict the Law of Conservation of Mass, has been explained through
the equivalence of mass and energy proposed in 1905 by Albert Eisnstein
in his Theory of Relativity. Mass and energy are interconvertible according
to the relationship

(3) E = me P where E is the energy corresponding to a mass, m,

and c is the velocity of light

Nuclear Chemistry I

631

If m is expressed in grams, and c in cm/sec, the energy is given in ergs.
Inasmuch as the velocity of light, 2.99 X 10 10 cm/sec, is so huge a number,
even an extremely small quantity of mass corresponds to a vast amount of
energy. One gram of mass is equivalent to approximately 9 X 10 20 ergs,
or 2 X 10 in calories, or 2.5 X 10 (i kilowatt-hours of electrical energy. Be-
cause of the large values of energy involved in nuclear calculations the electron
volt is more conveniently used as a unit. One atomic mass unit is equivalent
to 931 million electron volts.

Thus the mass defect represents a liberation of energy in the process of
forming a stable nucleus from its components. This energy is the binding
energy of the nucleus, since its value would also be equal to the energy
required to break up a nucleus into its component nucleons. Because dif-
ferent nuclei contain different numbers of nucleons, a better relative measure
of nuclear stability is the binding energy per nucleon, calculated by dividing
the binding energy of a nuclide by its mass number. Thus the binding
energy of Ne-° is 160 Mev and that of Bi 209 is 1640 Mev, yet the binding
energies per nucleon are 8.0 Mev and 7.8 Mev, respectively.

If the binding energy per nucleon is plotted against mass number as in
Figure 48.4, it is seen that not all nuclei are equally stable. Maximum
stability occurs in the region of mass numbers centering about the element
iron. Possibly this has some bearing for the earths core consisting primarily
of iron. Minimum nuclear stability is shown by those elements at the two
extremes of the mass number range, that is, elements of both high and low
mass numbers are unstable with respect to elements of intermediate mass
number.

For all stable nuclides excluding those of hydrogen and helium, the
number of neutrons in a nuclide is equal to or greater than the number of
protons. In Figure 48.5 there is plotted for the naturally occurring stable
nuclides the ratio of the number of neutrons to protons, n/p, against the
atomic number, Z. For the lightest elements, excluding hydrogen, this ratio is
unity but slowly increases with atomic number to approximately 1.5. Apparent-
ly a greater ratio of neutrons is required to contribute stability to atoms of
high atomic number. For each element there is a range of the n/p ratio
within which a nuclide can be stable. Where this ratio is too high, as in
6 C 14 , such a nuclide will emit a beta particle. This has the effect of decreas-
ing the number of neutrons by one and of increasing the number of protons
by one, thereby decreasing the n/p ratio. Where the n/p ratio of a nuclide
is too low, as in 7 N K \ a positron is emitted.

(4) beta emission: r»C 14 — > 7 N 14 + ^e ()

(5) positron emission: 7 N 1 * 6 C 1:t +

The positron is a particle having the same mass as the negative electron;

their charges are equal in magnitude but opposite in sign. Its time of ex-

istence as a free positron is very short, approximately 10 ~ 9 second. When
it unites with a negative electron, both particles are annihilated and in
their place appears gamma radiation of equivalent energy. This evanescent
particle was discovered by the American physicist, C. D. Anderson, in

632

Nuclear Chemistry I

1932 through cloud chamber photographs taken during cosmic ray investiga-
tions, though its existence had been postulated earlier by the English
mathematical physicist, P. A. M. Dirac.

Nuclides may be classified also according to whether the numbers of
neutrons and protons within them are odd or even. Four possibilities exist,
as shown in Table 48-D, which also gives their distribution among the
naturally occurring stable nuclides.

Table 4S-D

Types of Nuclei

Type

Number of Protons

Number of Neutrons

Number of Species

Even

164

Even

Odd

III

Odd

Even

Odd

Apparently nature prefers a nuclear configuration where the numbers of
neutrons and protons are both even. Nature will tolerate the presence of
odd numbers of either neutrons or protons but dislikes the odd-odd type of
nucleus. Further the four such species are among the lightest elements;
they are X H 2 , a Li 6 , 5 B 10 , and 7 N 14 , and no stable odd-odd nuclide of mass
number greater than 14 is known. Especially stable are those nuclides which
contain a number of protons or of neutrons equal to the so-called ‘magic
numbers”: 2, 8, 20, 28, 50, 82, and 126. In particular, the elements having
20 or 50 protons, or 20, 50, or 82 neutrons, have the largest numbers of
stable isotopes.

To be acceptable, a theory of the nucleus must weld these apparently
unrelated properties into a coherent whole. At present two theoretical models
of the nucleus, the liquid drop model and the shell model , which emphasize
different aspects of the nucleus, are in use. The liquid drop model was pro-
posed by Niels Bohr in 1936, It postulates that nucleons have a random
arrangement and attract each other much as do the molecules in a drop
of liquid. Forces akin to surface tension tend to keep the nucleus close
to spherical shape. The liquid drop model obtains support from the fact
that nuclear density for all atoms is approximately constant, about
10 14 g/cm*; it has been useful in interpreting the mechanism of fission.
The nuclear shell model, for which much convincing evidence was pre-
sented by Mrs. Maria Goeppert-Mayer in 1948, postulates that nucleons
occupy energy levels analogous to electronic energy levels, The “magic
numbers” are thereby related to completed nuclear shells, while the greater
stability of nuclides with even numbers of protons or neutrons is due to a
pairing of nuclear spins.

Nuclear Chemistry * 1

633

5. Artificial Nuclear Reactions* Radioactivity is the spontaneous disinte-
gration of unstable nuclei, in the process of which new elements are formed.
In 1919 Rutherford discovered that nuclear transformations could be induced
in ordinarily stable nuclei by bombarding them with energetic particles.
Using alpha rays from naturally radioactive radon as projectiles incident
upon nitrogen gas, Rutherford was able to show that oxygen and protons
were produced. The equation representing this nuclear change is

(6) 7 N 14 + 2 He 4 -> 1 H 1 8 0 17

(nitrogen (alpha (proton) (oxygen atom,

atom) particle) isotope 17)

This conversion of nitrogen to oxygen was the first deliberate man-made
transmutation of elements, the realization of the alchemists* dream.

To classify and to simplify the writing of nuclear reactions, a notation
has been developed. The element being bombarded is written in the con-
ventional manner, such as N 14 , and this is followed by a parenthesis inside
of which is stated the symbol for the projectile, then a comma, and the
symbol for the ejected particle; the element formed by the nuclear reaction
is written immediately after the closed parenthesis. Thus the notation for the
reaction of Equation 6 would be N 14 (a, p)0 17 . The discovery of the neutron
in 1931 by the English physicist, James Chadwick, was made by the re-
action: Be 0 (a, n)C 12 , that is,

(7) 4 Be» + 2 He 4 c C 12 4- 0 n a

(beryllium (alpha (carbon (neutron)

atom) particle) atom)

Nuclear reactions may be classified on the basis of the projectile used
and the ejected particle. The reactions in Equations 6 and 7 would be
classified as (a, p) and (a, n), respectively. Other reactions of importance
are (n,p), (d,p), and (n,2n), for which examples are cited below.

(8)

20 Ca« + on 1 -» in K« + a HP

( n >p)

(9)

S2 Pb 207 + JH 2 2S Pb 208 + iH 1

(d,p)

(10)

22 Ti 48 + 0 n J -> 22 Ti 47 + 2 on 1

(n,2n)

Different

products can be formed by the reaction

of a given nuclide

and projectile since the energy of the projectile is an important factor in
determining the course of a reaction. Like chemical reactions, nuclear re-
actions may not proceed along a single path and several possible
transformations can occur at the same time, especially if the target
element has several isotopes. When 8 Li 6 is bombarded with deuterons, the
following reactions can occur:

;jLi 7 +

(d,p)

(d,n)

( 11 )

;Li 6 + X H 2

2 He 3 + 2 He 4 + 0 n l

634

Nuclear Chemistry I

The products of an induced nuclear reaction may themselves be un-
stable and disintegrate radioactively in a manner similar to that of the
naturally radioactive heavy elements. The first such production of artificial
radioactivity was reported in 1934 by Irene Joliot-Curie, the elder daughter
of Marie Curie, and her husband, Frederic Joliot. Since then hundreds of
man-initiated nuclear reactions have been achieved for the production of
radioactive nuclides; a few examples of artificially induced radioactivity
are given below. Many of the unstable nuclides synthesized by nuclear
bombardment had no previous existence on earth.

(12)

5 B 10

-r

.He*

7 N 1:t

on 1 ;

7 N 13

oC 13

+ie°

(13)

7 n>*

iH 1

8 0 1S

8 O ls

7 N 15

■Hie 0

(14)

»Co“ +

on 1

.jCo 80

, 7 Co«®

— ^

2 8 Ni 80

-ie°

(15)

T8 Pt 198 +

iH 2

T8 Pt 107

iH 1 ;

™Pt 107

70 Au 187 +

-ie°

6. The Acceleration of Charged Particles. Until 1931, alpha particles from
naturally radioactive sources were the sole particles experimentally avail-
able to induce nuclear reactions. In practice, alpha particles make rather
inefficient projectiles since a nuclear reaction can be induced only if the
projectile strikes the nucleus and, even though possessed of great energies,
less than one alpha particle in a million scores a direct hit upon a nuclear
target. Most of the energy of an alpha particle is dissipated in ionizing
the atoms of the substance through which it passes. Of greater importance
is the repulsion of the positively charged alpha particle by nuclei, which are
themselves positively charged. Such repulsion deflects the alpha particle
from striking nuclei in its path.

Today devices are available which can accelerate charged particles to
energies in the billions of electron volts. These energetic particles can then
be directed at a nuclear target and nuclear reaction can thereby be produced
at will in the laboratory. Among these devices are the Van de Graaff genera-
tor, the linear accelerator, and the cyclotron and its modifications such as
the synchrocyclotron, the betatron, and the synchrotron.

The Van de Graaff generator utilizes the electrostatic principles that a
sphere can accept any charge and that the discharge of electricity occurs most
readily from points. Positive charges from a high direct current source of
potential, about 5,000 to 20,000 volts, are directed from a pointed conductor
onto a moving belt, made of a nonconducting material such as silk, rayon,
or paper. The charges are carried by the belt to a spherical surface where a
potential as high as 8-12 Mev can be built up.

The linear accelerator consists of a series of separate hollow cylindrical
tubes, arranged along a common axis, with successive tubes at increasing
potentials. Opposite potentials are applied to adjacent tubes and the nature
of these potentials is alternated in phase with the passage of the charged
particles through the tubes. In moving through the tubes, die particles re-
ceive increments of energy as they pass through die gap from one tube
to the next and emerge with energies of several Mev.

Nuclear Chemistry I

635

The most important device for accelerating positive ions, however, is
the cyclotron, invented by Ernest O. Lawrence of the University of California,
for whom the element of atomic number 103 was named. As in the linear
accelerator, successive accelerations are used to impart energy to charged
particles. But instead of moving in a linear path, the particles move in a cir-
cular course through two hollow electrodes in the shape of D’s, and hence
termed Dees. An alternating potential synchronized to the movement of
the particles, in conjunction with a magnetic field at right angles to the plane
of the Dees, moves the particles in a circular path of ever increasing radius
from Dee to Dee at higher and higher energy. Ultimately the particles leave
the apparatus through a thin metallic window with energies as high as 700
Mev and impinge upon a nuclear target, as shown in Figure 48.6.

D and FF are the hollow semicircular Dees
C is the source of the charged particles

F is a deflecting plate at high negative potential which deflects the particles from
their spiral paths
W is the exit window
T is the target

The magnet is not shown. It is at right angles to the plane of the Dees, one pole
above and the other below the Dees. The magnet weighs about 15 tons and the pole
faces are about 15 feet in diameter.

Figure 48.6. The Cyclotron.

The synchrocyclotron, or frequency modulated cyclotron, uses a single
Dee and varies the frequency of the alternating potential to compensate for
the relativistic increase in mass of the rapidly moving particles. A particle
may attain a speed half that of light and hence its mass would be about
15% greater than it is at rest. The largest accelerator now in use is the
synchrotron at Brookhaven National Laboratory. It has a diameter of 84C

ggg Nuclear Chemistry l

feet and can accelerate protons to an energy of 33 Bev; a 70 Bev proton
accelerator is under construction in the Soviet Union.

Because the neutron is electrically neutral it cannot be accelerated through
the use of electrical machines. Neutrons are very effective projectiles, how-
ever, for inducing nuclear reactions since they are not repelled by the
nuclei of target elements, and collisions between them and target nuclei have
a high probability of occurrence. The only source of neutrons is through
nuclear reactions such as Be 9 (a,n)Be 8 ; H 2 (d,n)BP; and Li*(p,n)Be 7 . Gamma
rays can also be employed to initiate nuclear reactions, as in Be 9 (Y,n)Be 8 ,
but because gamma rays do not have high energies in comparison with
accelerated particles, their use is limited.

QUESTIONS

1. What is the Group Displacement Rule? Write specific nuclear reactions which
illustrate the rule.

2. List three chemical properties of radium. How is the free metal obtained?

3. What is meant by a radioactive disintegration series? Name four such series.
What is the end product of each? Account for the fact that the remainder,
after dividing the mass number of each member of a given series by four,
is the same value.

4. Draw a diagram, similar to Figure 48.1, starting with thorium-232, in which
the particles emitted in sequence are: alpha, beta, beta, alpha, alpha, alpha,
beta and gamma, beta, alpha, and beta. Indicate the name, atomic number,
and mass number of each member of the series.

5. Arrange alpha particles, beta particles, and gamma rays in order of (a) ion-
izing power and (b) penetrating power.

6. Show that the mass of a gamma ray is hv/c 2 . Calculate the mass of a gamma
ray having an energy of .1.0 Mev in amu and in grams.

7. Define the electron volt and state its equivalent values in ergs and in
calories. How much energy, in calories, is represented by a mole of protons
having an average energy of 3.0 Mev?

8. What is the kinetic order of radioactive decay? What is meant by the half-life
period ? According to mathematical theory, how long would it take for a
given sample of a radioactive substance to decay completely?

9. (a) For an element which has a half-life of 100 days, how long would it
take for 40% of an initial amount to disintegrate (b) what portion will
have decayed in 150 days?

10. From an initial quantity of 16 g of uranium how long would it be before
(a) 1.0 g remains (b) 0.75 g remains?

11. One gram of radium is placed in a sealed container, (a) How many radium
atoms -remain after one half-life period (b) after what length of time will
one billion radium atoms still remain (c) at the end of one year, what weight
of helium will be present in the container, assuming its source is solely
radium?

tfuchir Cht mtstry 1

637

I& A radioactive nuclide, V, mints %m alpha p.utulc tu form a stable nuclide, Y.
The halMtfe t»f X is UHHI \r*us, An auaKsis of a geologic sample of X anti Y
gives 0,0! 5' f of \ In vs eight, What ean lie said a Inn it the age of the sample?

13. List the devices einphncd to detect ludnuetivitv . State the basic principle
or reaction involved in each,

14. (a) Why are the tracks of an alpha particle in a cloud chamber denser
than those of a beta particle \ b) vsh> is a bubble chamber more sensitive
than a clmid ehamber?

15. Briefly discuss those charae tern tics of atomic nuclei which must be ex-
plained in a complete theory of the nucleus.

16. What is meant by the mass dvfi'tt of a nucleus? What explanation accounts
for it?

17. Calculate the mass defect of the cat bon- 1 2 isotope.

18. Calculate the mass of the proton, neutron, and the deuteron in Mev.

19. (a) The most sensitive chemical balances can weigh approximately to one
nmrogram, calculate the energy equivalent to this mass (b) calculate the
mass lost m a chroma! reaction which evolves 30 leal

20. What is meant by the “binding energy of the nucleus” and the “binding
energy |*er nucleon”? l>raw a graph of bind mg energy per nucleon against
mass number. Which are the most stable nuclei in nature?

21. Calculate the binding energy of (a) the deutmm (b) the C 1 - isotope (c) sort
isotope of Ke ai * which weighs 55.94826 attm

22. (a) How* is the ratio of nrutrons to protons m .« nuclide related to the
stability of the nuclide (bl what is the relation of odd and even numbers of
protons and neutrons m nuclides to their stabilities? List the odd-odd nuclides
known,

23. Briefly discuss the two theories promised as a model of the nucleus.

24. What is meant b> “artificial nuclear reaction? How can it lie induced?
Write an equation for such a reaction. What two particles are emitted by
most artificial radioactive elements?

25. Use the nuclear “notation” to express the reactions in Equations 12 to 15.

26. Write equations for the nuclear reactions represented by the following ex-
pressions, (a) B l0 (d,ti)C:'i (bl Na**(d,p)Na»* <e) Mg**(n,p)Na**

(d) Al«(«.n)P.

27. Why ore alpha particles more effective projectiles for inducing nuclear re-
actions in elements of low atomic number?

28. Briefly discuss the methods of accelerating charged particles. Can neutrons
be accelerated in these devices? Why are neutrons specially effective as pro-
jectiles?

Nuclear Chemistry II

1. Nuclear Fission. The artificial nuclear reactions we have so far con-
sidered have dealt with atoms of low atomic number and the nuclides formed
differed but slightly in atomic number from that of the target nucleus.
Such reactions might be viewed as ‘chipping” reactions wherein only a small
increment is either chipped off or added to a nucleus. In 1934, the Italian
physicist, Enrico Fermi, reported that when uranium was bombarded with
slow neutrons other elements were formed and in 1939, O. Hahn and F.
Strassmann, in Germany, identified one of the products as the element barium.
Since uranium and barium differ so widely in atomic number, the results
were bewildering and it remained for Lise Meitner and O. R. Frisch to
make the suggestion that the uranium nucleus, after neutron capture, split
into two nuclei of about equal size. Apparently scientists found it diffi-
cult to conclude that a nucleus could be split in two.

The splitting of a nucleus into two approximately equal fragments is
known as fission. To date fission has been observed only with the heaviest
elements, notably thorium, uranium, and plutonium but it has been brought
about in tantalum by 400 Mev alpha particles. The energy of an incident
particle must be sufficient to overcome a nuclear potential energy barrier
akin to the energy of activation in chemical reaction. Fission could be ex-
pected from the relation between binding energy and mass number (Fig-
ure 48.4), which indicates that the nuclei of heavy elements are unstable
relative to those of the medium-heavy elements. Some heavy nuclei undergo
spontaneous fission but at an extremely slow rate; the half life period for
the spontaneous fission of U 288 is almost 10 16 years and that of U 285 is
even greater.

Perhaps the most important fission reaction is that of U 235 by slow or
thermal neutrons of about 0.025 ev energy. The nuclear reaction is

/ , e -4oX 80 - 108

(1) aaU 285 + on 1 J + 1 to 3 on 1 ~f* energy

\ s^Y 12 ^ 58

638

Nuclear (:fa mi\try II

639

where X ami Y are the iixMou hagim-iit.v The equation militates th.it a given
nuclide can split in different ways st» that its fission produets consist of a wide
range of nuclides. Fission of does not generally yield two equal parts

felt rather a light fragment having a mass numher between <S0 and 108 and
a heavy fragment with a mass number between 125 and 153. Whatever
these fission fragments may be. boss ever, the sum of their atomic numbers
must add up to 92. and the sum of their mass numbers, including the neutrons
produced, to 236.

All fission reactions are accompanied hv the release of enormous amounts
of energy, the greater purt Wing thermal energy as the kinetic energy of the
fission fragments but a good projairtion consisting of radiation in the form
of gamma rays. The .source of this fission energy is a relatively large loss
in mass during the nuclear reaction, the sum of the masses of the products
is less than that of the reactants and tin's difference in mass appears as
energy, in the reaction

(2) f „n' - ..Sr”-"' +- „Xe w ■+■ 2 „n'

die loss in mass is 0.233 iimu, representing an energy emission of 208 Mev.
It is noteworthy that this huge release of energy was initiated by a neutron
having only 0.025 eV of energy. The complete fission of one pound of
uranium would liberate an energy equivalent to the burning of over two
million pounds of coal or the explosive force of 10,000 tons of TNT.

In most cases of fission, the products are themselves unstable and undergo
radioactive decay at various rates, they have a high n/p ratio and form a
short series of beta purtiele emitters till a stable nuclide configuration is
attained. In Equation 2, both ..Sr"* and ,,Xe n * are radioactive, decaying by
beta activity to the stable nuelides. „Mo M and , : I,a ,;w respectively.

2. The Chain Reaction. A striking feature of the neutron-induced fission
of U**‘ is its production of additional neutrons. The number of neutrons
emitted per atom of uranium undergoing fission was shown as being from
(me to three. The actual number of neutrons released, whether one, two,
or three, depends upon the specific mode of fission of a uranium atom but
die over-all average for a sample of IT J!U is 2.47. Over 99“# of these second-
ary neutrons are emitted very quickly after the absorption of the initial
incident neutron, prohably within 10 13 second. In turn, these neutrons cause
fission of additional uranium nuclei and thereby continue the fission process.
Thus a series of reactions, or a chain reaction, could theoretically proceed
until all the atoms in a sample of uranium had undergone fission. This is
the basic principle common to both the uranium nuclear bomb, incorrectly
termed the "atomic bomb," and the controlled nuclear reactor. Assuming
that the number of neutrons doubles in each generation and that there are
no neutron losses, after 80 generations there would have been produced 2"°,
or 10* 4 , neutrons, enough to cause the fission of one gram-atom of uranium.

In this discussion it has been assumed that all the neutrons produced by
fission continue to propagate further fission reactions so that, in principle,
the injection of a single neutron into a sample of uranium could trigger the
fission of ail the uranium atoms therein. In practice such is not the case.

640

Nuclear Chemistry II

Any effect which results in the loss of neutrons prior to their causing fission
will break the reaction chain and neither a bomb nor a controlled reactor
will be possible. Neutrons can be lost in two ways: A) by escape through
the surface of the uranium mass and B) through absorption in nonfission
capture processes. The minimum requirement for maintenance of a chain
is that after each fission event there will still remain, with due account for
neutron losses, at least one neutron available to cause the fission of another
nucleus. On this basis, there is defined a factor, k 9 which is the ratio of the
number of neutrons available for fission in any one generation to the num-
ber in the preceding generation. If k is one or more, a chain reaction is
possible; if its value is less than one, even ever so slightly, a chain reaction
cannot be maintained.

To reduce the number of neutrons lost by escape so that the value of
k is at least one, the size or the mass of the uranium must be greater than
a certain minimum value. This curious fact arises because the probability
of neutron escape depends upon the surface area of the uranium mass
whereas the probability of collision with a uranium nucleus depends upon the
mass, or volume, of uranium present. On the assumption that the uranium mass
has a spherical shape, as the mass of uranium is increased, its volume, which
is proportional to the cube of the spherical radius, increases to a greater
extent than does its surface, which varies as the square of the radius. Hence
an increase in size favors neutron collision rather than escape. There will
be a minimum size, defined as the critical size , which is required to reduce
the proportion of neutrons that escape so that a chain reaction can take place.
With a quantity of uranium less than the critical size, the rate of neutron
escape is too great to maintain a chain reaction. With sizes greater than the
critical, the chain reaction occurs spontaneously. Even the advent of a
stray neutron into a critical mass of uranium could initiate the chain re-
action. The value of the critical size is not a constant but depends upon
the purity of the uranium and its actual geometric shape which may be other
than spherical; the value can also be reduced by surrounding the uranium
mass with a material which will reflect escaping neutrons back into the
uranium mass.

Neutrons can also be lost through nonfission capture by impurities in
the uranium or by the U 238 isotope, Impurities can be removed by chemical
procedures which will yield almost 100% pure uranium. Natural uranium,
however, consists of a mixture of isotopes, 99.3% U 238 , only 0.7% U 235 (the
desirable isotope), and a trace of U 234 , so that purified uranium metal ob-
tained from natural sources contains these isotopes in the same proportions.
From this mixture of isotopes the U 235 must be separated. The separation
of U 235 from U 238 poses an enormous problem since it cannot be accomplished
by chemical means; the electron configurations of both isotopes are identical
and it is upon these configurations that chemical properties depend.

3, Separation of Isotopes. Most methods for the separation of isotopes
are based on a slight disparity in physical properties which arises from the
difference in mass. The techniques which have been employed for separat-

Nuclear Chemkt n/ I!

tog lP a from V Jis art*: A* gaseous diffusion; R ) thermal diffusion; C) eiec-
tromaguetic separation; ami D^ centrifugal separation.

In the f*mvow diffusion method , natural uranium is converts! to the
volatile compound, I'F*, hi reality a mixture of U- ,r ‘F„ and r- MH F„. Diffusion
of these compounds through a porous harrier follows Cruiuuns Law, and a
very slight enrichment of the U* 1: ‘ compound, proportional to the square
root of the compounds* masses, 352/349 or 1.0014, will he effected. In order to
produce 99% pure about 4,000 diffusion barriers in series are re-

quired. The first gaseous diffusion plant, with its acres of diffusion harriers
and its thousands of miles of piping went into full operation at Oak Ridge,
Tennessee, in 1945.

The thermal diffusion method is has<*d on the tendency of heavy molecules
of a gas to concentrate in the cooler region of a container which is heated
unequally. When a heated wire is stretched along the axis of a vertical
tube, the lighter molecules collect near the hot wire and the heavier ones
near the wider wall Convection currents then cause the lighter molecules
to rise and the heavier ones to fall resulting in a higher concentration of
the lighter molecules at the top ami the heavier ones a* the lmttom of the
tube. This method can also he used for pure liquids and for solutions. With
a tube iilicHit 30 meters long, and a temperature difference of 800°C,
almost complete separation of the chlorine isotopes, Cl :ts and CP 7 , can be
effected. A large scale thermal diffusion plant was first built at Oak Ridge
to 1944.

The principle of the electromagnetic method is similar to that of the
mass spectrograph, A stream of electrons traversing the vapor of UF* pro-
duces positive, singly charged uranium ions, U*. These are accelerated
by an electric potential and the ion lieam passes into a chamber where a
magnetic field bends it into a semicircular path. Ions having different masses
follow slightly different trajectories and are caught in separate collectors.
The electromagnetic method was the first to yield appreciable quantities of
U***. Though no lunge* user! for that purpose, it is still employed in the
separation of isotopes of over fifty stable elements.

Centrifugal separation depends upon the fact that, in a rapidly rotating
container, a force acts upon the molecule,* therein so that the heavier ones
concentrate at the outer side of the container while the lighter ones collect
near the axis of rotation. In principle, this method offers better promise
than gaseous diffusion lieeause it depends upon tin* absolute difference
between the masses of the isotopes rather than on the square root of the
ratio of the masses. To utilise this method, a pilot plant was constructed
but the project was abandoned because of the magnitude of the engineering
problems.

Isotopes have the same electron configurations and hence undergo the
same chemical reactions, but they do so at different rates. The activation
energy for a chemical reaction is such that a molecule containing the lighter
isotope generally reacts more rapidly. For isotopes of the heavy elements,
this difference in rate is negligible but for the lighter elements, and par-

(542 Nuclear Chemistry U

ticularly for hydrogen, the difference is appreciable. Deuterium has been
concentrated by the so-called exchange reaction ,

(3) H,(g) + D a O(Z) Do(g) + H 2 0(Z)

The rate of the reaction to the left is about three times as great as it is to
the right, that is, the equilibrium constant for the reaction is about 0,3.
If ordinary gaseous hydrogen, a mixture of its isotopes, is mixed with steam
and a catalyst to attain rapid equilibrium, the isotopic exchange yields
an enrichment of the heavier isotope in the steam, which is then condensed*
By successive operations using this technique, heavy water containing 99,5%
D 2 0 is prepared commercially for a few cents a gram. The hydrogen isotopes
have also been separated by fractional electrolysis. During the electrolysis
of water at high current densities, the lighter isotope is liberated preferen-
tially and so deuterium is concentrated in the water residue.

4. The Nuclear Reactor. In a nuclear bomb, all the successive individual
fission reactions and the accompanying energy release take place in so
short an interval of time, within 10 -c second, that a violent explosion is the
result. In a nuclear reactor, the over-all rate of fission is controlled so that
the energy is released at a relatively slow rate and explosion does not occur.
The principles underlying each device are the same and a given mass of
uranium undergoing fission in a bomb or in a reactor will yield the same
amount of energy.

To make a nuclear bomb, it is merely necessary to bring together the
critical size of uranium and inject neutrons into the mass. To regulate the
rate of fission in a reactor, rods of cadmium or of boron steel are inserted
throughout the uranium mass. Both cadmium and boron nuclei have large
cross sections for the capture of slow neutrons. The depth to which the
rods are inserted in the uranium mass controls the rate of capture of neutrons
and hence also the number left free to carry on the fission chain. To shut
down a reactor, the rods are inserted to considerable depth; to start up a
reactor, the rods are withdrawn till the required power output is attained.

The source of neutrons to initiate the primary fission in a reactor can
be any nuclear reaction by which neutrons are produced, e.g., Be d (a } n)C 12 .
The secondary neutrons emitted in the fission process are fast neutrons with
energies of 1-2 Mev. The probability of fission for U 335 , however, is greatest
with thermal neutrons having energies about 0.025 eV, Thermal neutrons
are those with energies approximately equal to the kinetic energies of gaseous
molecules; at 20°C, the thermal energy of a gas molecule, given by 3/2 kT
(where k is the Boltzmann constant), is about 0,038 eV. The high energy
neutrons produced by fission can be slowed down to thermal energies by
having them collide with lightweight nuclei such as those of carbon or
hydrogen. On the assumption that a neutron loses half its kinetic energy
in an elastic collision with such a nucleus, only about 18 collisions are
required to reduce the neutrons energy to thermal values. Graphite is
therefore mixed with the uranium in a regular geometric array in such
proportion that collisions of the initially fast neutrons with the carbon atoms
of the graphite slow them to thermal velocities, whereupon a subsequent

Nuclear Chemistry U

643

collision with a uranium nucleus results in fission. Because of its function
the graphite is known as a moderator. Heavy water is an excellent modera-
tor; ordinary water and paraffin are cheap but not completely satisfactory
as moderators because hydrogen has an appreciable tendency to absorb
neutrons. A United States patent on a neutron reactor as described in the
foregoing was granted to Enrico Fermi and Leo Szilard on May 17, 1955,
and the rights thereto were assigned to the Atomic Energy Commission. 1
A schematic diagram of the components of a reactor is shown in Figure 49.1.

Slow Neutron to Start the Keaction

The above reactions
are repeated.

Figure 4$J. A Schematic Diagram of a Fission Chain Enaction.

*U.S. Patent No. 2,708,656; s primer on nuclear fission, it may be purchased from
the Superintendent of Documents, Washington, D.C. for $0.25.

644

Nuclear Chemistry It

Industrial production of nuclear energy involves serious technological
problems. First the energy produced must be removed from the interior of
the reactor. This removal of heat is accomplished by employing a cooling
fluid which can be pumped through a piping network within the core of
the reactor. The heated coolant then flows through a heat exchanger, a unit
outside the reactor where the heat is absorbed by water in the production
of steam, after which the coolant returns to the reactor to repeat its cycle.
Heavy and ordinary water, liquid sodium, and hydrocarbon liquids and
gases have been used as coolants.

. Quite another problem is the safety of the operating personnel. Nuclear
radiation can be extremely dangerous to health and, as a consequence,
there has developed the field of health physics for the protection of person-
nel against the hazards of radiation. Obviously the reactor itself must be
shielded. For most reactors an eight-foot thickness of concrete forms a satis-
factory shield. Furthermore, a coolant such as sodium forms the radio-
active isotope Na 24 while in the reactor core and beta and gamma radiation
from it while outside the core is hazardous.

As fission proceeds, the U 235 fuel is consumed but in the ‘nuclear ashes”
are numerous by-products of the fission, a mixture of more than thirty ele-
ments from zinc to gadolinium, many of them radioactive. The disposal of
such radioactive “wastes” from reactors and processing plants is also a prob-
lem since such wastes cannot be unceremoniously discharged into soil or
stream for fear that they may ultimately appear in human food.

5. The Transuranium Elements. An atom of U 238 can capture a neutron
without undergoing fission and form the isotope U 239 . This isotope is unstable;
it is a beta ray emitter and has a half-life period of 23 minutes. The result of
this emission is an element of atomic number 93, named neptunium, the first
of the transuranie elements, so called because they are higher in atomic
number than uranium. In turn, Np 239 is also a radioactive beta emitter,
thereby producing an element of atomic number 94, plutonium. These trans-
formations are summarized below.

(1) 92 U 238 + on 1 * 2 U 239 -> ^ + * 3 Np 239 .,<?<> + JPu* 3 *

Relatively stable, Pu 239 is an alpha emitter with a half-life period of
24,400 years. By the nuclear 'reactions in Equation 1, Pu 289 can be produced
in a uranium reactor and can be readily separated from the U 238 matrix
because it differs chemically from uranium. The Pu 339 nuclide assumes
special importance because, like U 2 * 5 , it is fissonable by slow neutrons and
so- can be used in a bomb or in a controlled reactor. Further, a uranium
reactor can be designed to produce plutonium besides undergoing fission
so that the quantity of fissionable material therein remains unchanged or
even increases. Such a reactor, which produces power through fission and
simultaneously regenerates fissionable material, is known as a breeder reactor.

Plutonium is the first man-made element to be produced on a large
scale. At Hanford, Washington, a plant was constructed in 1943 for its
manufacture; a remarkable feature of this plant is that it was designed
from chemical data based upon experiments with only half a gram of

Nuckar Chenn*try It

645

plutonium- In addition to Np' '' and IV' 11 '. many other isotopes of these
elements are known, in all there are eleven isotopes of neptunium with
mass numbers iioin 231 to 21! ami sixteen isotopes of plutonium with mass
numbers from 232 to 24Ci and each has its own specific- activity and half-life
period. To date, additional trausur.inic elements through atomic number
103 have been prepared In nuclear bombardment of uranium, neptunium,
and plutonium with neutrons, alpha particles, and even carbon and nitrogen
tons. Furthermore, four elements whose atomic numbers are 43, 61, 85, and
87, undiscovered till the 19.30s but for which positions k*d been left vacant
fa the Periodic Table, were first made by nuclear technology. To sum up,
there base been synthesized fifteen elements which either did not exist
previously on earth or wen* present in sueh trace quantities that they passed
undetected. These elements are technetium, promethium, astatine, francium,
and the trunsur.mie elements of atomic number 93 to 103.

6. Nuclear Fusion. The elements of \ ery low- mass numbers are also
unstable with respect to nuclei of intermediate mass number and, in principle,
it is possible for nuclei of these light elements to combine and form more
stable single nuclides of heavier mass. Such nuclear reactions are known as
fusion reactions. For the lighter elements, Figure 48.4 indicates that the
binding energy per nucleon is u-lativeJy Imv but that there is a sharp
rise in the values uf these binding energies. Hence fusion reactions should
release tremendous energies, far greater than do fission reactions. It has
been suggested that the source of the sun's energy is the fusion of hydrogen
to helium. Tin- process consists of several steps but the over-all reaction is

(2) 4 ill’ 2 ..He* 4- 2 M e" -f- energy

The two positrons produced combine with the two electrons originally as-
sociated with the hydrogen atoms to yield additional energy. Similar fusion
reactions have Ih-cii effected on earth, notably

(3) ill* a- ,H* -* 3 He* (fusion of two deuterium nuclei)

(4) ,11' 4 ill-* -* t He* -f- nil 1 (fusion of tritium and deuterium)

In order to initiate a fusion reaction, the bombarding particles must have
energies corresponding to temperatures of millions of degrees and hence
such reactions are known as thermonuclear reactions. Such temperatures
can be achieved through a nuclear fusion reaction which can then act as a
trigger for a fusion reaction. This is the basis of the hydrogen fusion bomb
or the thermonuclear lxrmb. Fusion of deuterium, which comprises about
0.015% of the hydrogen atoms in water, offers the prospect of an almost
inexhaustible supply of energy. Furthermore, in such a “fusion reactor” there
would be no radioactive wastes, though shielding would be necessary.

At temperatures of millions of degrees, matter exists as a totally ionized
system of positive ions and electrons known as plasma. Obviously one prob-
lem in the utilization of fusion reactions for the commercial production of
energy is the development of a container for this plasma since any earthly
material would vaporize if subjected to plasma temperatures. Confinement

646

Nuclear Chemistry II

of plasma in ordinary containers, however, has been achieved through the
use of electric and magnetic fields which pinch the plasma in towards the
center and away from the walls of the container, a phenomenon known as
the pinch effect.

7. Applications of Nuclear Energy. The uses of nuclear energy may be
classified broadly into two categories, military and peaceful For the military,
nuclear bombs offer explosive effects with magnitudes which far exceed
those of chemical bombs. The first nuclear bomb was dropped on Hiro-
shima, Japan on August 6, 1945, and effectively put an end to World
War II. The bomb is one more example in a long series illustrating the pro*
found effect which scientific technology has had upon military, political, and
economic affairs.

The explosion of a nuclear bomb has several effects. The temperature
produced at its center is of the order of 10 million degrees. Heat radiating
from this source is sufficient to ignite primary fires several miles distant
from the explosion while secondary fires can result from damage to elec-
trical installations. The blast effect, including the positive pressure phase
and the more dangerous succeeding suction phase, is equivalent to the
explosion of thousands of tons of TNT, and complete destruction of ma-
terial can be produced many miles from the center of the explosion. The
neutron flux, and particularly the gamma radiation emitted, are extremely
penetrating and can cause human casualties. Death or radiation sickness may
result depending upon the magnitude of an individual's exposure. The char-
acteristic mushroom cloud of a nuclear burst carries with it large quantities
of radioactive material which, after being conveyed by atmospheric winds,
falls out ultimately and may constitute a radioactive hazard, immediately
and in the future through absorption into foodstuffs.

Controlled nuclear reactors are being used to supply industrial power and
in the power plants of ocean going vessels and submarines. The most con-
structive applications of nuclear energy, however, arise from the variety of
isotopes, both radioactive and stable, which have been made available for
experimental use in scientific work. Though isotopes of a given element ex-
hibit identical chemical behavior; the radioactive isotope or radioisotope can
be detected through its radioactivity while a nonradioactive isotope can be
identified by the mass spectrograph. Radioisotopes can serve as tracers
in following the course of a chemical or biological reaction. An element
whose normal isotopic composition in a compound has been deliberately
changed by an experimenter is said to be labelled or tagged . For example, if
to ordinary carbon dioxide there is added a small percentage of that com-
pound containing the radioactive C u isotope, the route of the C0 2 in respira-
tion and in photosynthesis can be followed.

Many scientific problems, which could not be solved by the ordinary
means available, have yielded to tracer techniques. Only a few of these
problems can be mentioned here but these may suffice to give an idea of
the diverse uses to which radioisotopes have been put. Radioactive phos-
phorus, P 32 , in the form of a phosphate salt has been used to determine

Nurlear C)um'\lrtj il

647

the relative intake ul the element from the .soil and from added fertilizer.
By labelling CO., or HO with ()'\ the question as to whether the oxygen
produced in photosynthesis came from the CO., or from the H s O was decided
in favor of the water. In medicine radioisotopes are used for diagnosis and
for therapy. The flow of Wood through the body can be followed by an
injection ol a solution of sodium chloride, labelled with Na 24 Malfunction-
ing of the thyroid gland nan be detected and subsequent treatment carried
out by the use of radioiodine. I U1 . since iodine* concentrates in this gland.

Since* the very early days of radioactivity, radium and its emanations
have been used iur the treatment of cancer because of their ability to in-
hibit the growth of cancerous cells. But radium is very expensive. Artificial
Co"", which is relatively cheap and has a high activity for the emission
of gamma rays, is now widely used as a substitute. It is made by exposing
pure Co*" to neutrons in a reactor and has a half-life period of 5.3 years.

Many strictly analytical chemistry techniques use radioisotopes; by their
use the solubility of sparingly soluble salts has been determined. In industry,
radioisotopes monitor the manufacture of metals, paper, plastics, and textiles.
Industrial applications also include the study of the catalysis of organic
reactions and of frictional wear. Neutron and gamma radiation is proving
effective for the sterilization and preservation of food. The Food and
Drug Administration has approved for human consumption canned bacon
sterilized with Co"" or Cs ,r, v; the FDA has also approved the disinfestation of
wheat and the sprout inhibition of potatoes with gamma radiation of energies
less than 2.3 Mev.

An unusual application of radioactivity is the dating of past events. The
age of the earth has been estimated from the ratio of uranium to lead found
in uranium minerals. Since Pb auft is the end product of the radioactive decay
series starting with U :;<s , a quantity of lead is always found associated with
uranium. On the assumption that all the lead was initially uranium, and
knowing the over-all decay rate from uranium to lead, a period of time
can be calculated for when all the lead was uranium. Allowance can be made
for any lead in a sample which is not of radioactive origin, and uranium-
lead ratios thus interpreted indicate that the oldest uranium minerals have
an age of 4.5 x 10* years. Similarly, the radioactive carbon istotope, C 14 ,
with a half -life of 5,570 years, can be used to date organic matter. A small
percentage of C 14 exists in COj molecules in the atmosphere, being formed
by the bombardment of atmospheric nitrogen with neutrons produced by
cosmic radiation, that is, N' 4 (n,p)C 14 . In time all living organic matter,
plant and animul, attains an equilibrium concentration of the C 14 isotope
in its carbonaceous content. After the death of the plant or animal, C 14 is
no longer taken in and what C 14 is then present decays without being
replenished. Measurement of its residual activity enables one to determine
the age of specimens of archaeologic and geologic interest, such as textiles
from Egyptian tombs and wood from the last Ice Age. Tritium, with a
half-life of 12.26 years, has been used to date events up to thirty years old,
such as how long moisture remains in the air between evaporation and sub-
sequent precipitation.

648

Nuclear Chemistry 11

QUESTIONS

1. Distinguish between fission and fusion. Write a nuclear reaction which
illustrates each.

2. What is the origin of the energy of (a) fission reactions (b) fusion reactions?

3. A neutron having an energy of 0.03 eV can initiate a nuclear fission reaction
which yields 208 Mev. Does this statement contravene the Law of Con-
servation of Energy? Explain.

4. Write the nuclear reaction for the fission of U 23r \ if one of the products is
Ba 140 and three neutrons are produced per atom of uranium,

5. What is meant by a chain nuclear reaction? Illustrate by means of a nuclear
equation.

6. How does a nuclear bomb differ from a nuclear reactor in principle?

7. How is the rate of a nuclear reactor controlled? What is the function of a
moderator? How is the heat of a reactor removed from its core and con-
verted to useful energy?

8. Discuss the method of slowing down neutrons of high energy to thermal
energies.

9. If a neutron loses half its energy upon elastic collision, how many collisions
are required to reduce its energy from 1 Mev to 0.1 eV?

10. Design and draw a diagram of a nuclear reactor. Label all materials.

11. List the hazards associated with the operation of a nuclear reactor.

12. What is meant by the term critical size? What factors determine its value?

13. Discuss the methods of separating isotopes. On what principle does each
depend? Why are chemical means of little value?

14. Which method of isotope separation would be most efficient for (a) chlorine
(b) hydrogen (c) uranium?

15. How could a mixture of U 285 , U 288 , and Pu 2 -™ be separated?

16. How are neptunium and plutonium produced? Write pertinent equations,

17. Write equations for the decay by beta activity of Xe 18ft to La 182 .

18. Why are fusion reactions known as thermonuclear reactions? How can fusion
reactions be initiated? How can they be contained?

19. List some applications of nuclear energy.

20. Define (a) plasma (b) radioisotope (c) labelled compound (d) breeder
reactor (e) thermal neutron.

Analysis

1. Instrumental Analysis, Chemical analysts, as defined by Webster, is
“the determination, which may or may not involve actual separation, of one
or more ingredients of a substance either as to kind or amount.” Analysis
to determine the kind or chemical nature of a substance is known as
qualitative analysis , whereas the determination of the amount, quantity, or
percentage of an ingredient is quantitative analysis. Classic quantitative
methods have been subdivided further as nr at', i metric and volumetric, on the
basis of whether the amount is resolved by weighing or through the use of
volumes of standard solutions, as in a titration.

in recent years the trend in analysis has been towards the use of instru-
mental techniques based upon the physical properties of the substances
being determined. In earlier pages we referred to certain instrumental tech-
niques hot did not so designate’ them at the time. Among these are the
mass spectrograph which separates and detects isotopes on the basis of their
masses, and the "optical" spectrograph which disperses the characteristic fre-
quencies of spectral lines emitted by excited atoms or molecules of a sub-
stance. Optical spectrographs may operate in the infrared, visible, ultra-
violet, or X-ray region, each of which furnishes different but pertinent in-
formation concerning the nature of a substance. Thus infrared spectra yield
knowledge of the vibrations ami rotations of molecules, information of im-
mense value in determining the liond distances and bond strengths of organic
molecules. The converse of emission, the selective absorption of optical wave-
lengths yields much the same information and is utilized in various spec
trophotometers.

Iq theory almost any physicul property can be used for analytical purposes.
The wide range of instrumental analysis is indicated in Table 50-A which
list* but a few erf the techniques in current use and the properties upon
which they are based. Today analysis might be more properly defined as the
characterization of a material, to include its composition, properties, and
mmlittos.

650

Table 50-A

Analysis

Instrumental Techniques

1 .

2 .

Technique

Property

Spectrography

Emission or absorption of characteristic wave-
lengths of radiation

Thermal Conductivity

Difference in the conduction of heat by gases

Nuclear Magnetic Resonance
Spectrography

Polarography

Transitions between the spin quantum levels of
an atomic nucleus in a magnetic field

Difference in electric potential required to re-
duce ions in solution at a microelectrode

6 .

8 .

Potentiometric Titration

Conductance Titration
Polarimetry

The difference in potential between a reference
electrode and an indicator electrode immersed
in a solution being titrated

The change in electrical conductance of an
ionic solution during titration

Rotation of the plane of vibration of polarized
light

Refractometry

The change in direction of the path of a beam
of light in passing from one medium to another
of different density

2. Chemical Analysis. Despite the rapid advance of instrumental tech-
niques, it will be worth our while to consider some of the procedures of
classical quantitative analysis wherein chemical reactions are used to deter-
mine the composition of a compound or mixture of compounds. The basic
problem in such qualitative analysis is to determine the various elements
which comprise a given unknown sample.

Let us pose for ourselves the problem of analyzing a solution which con-
tains a number of different inorganic cations and anions. The problem could
be resolved quite simply if there were a specific reagent, with a charac-
teristic reaction, for each ion. Unfortunately such is not the case* There arc
some reagents, usually organic substances, specific for certain ions but these
are few and require special conditions. Generally a reagent reacts with sev-
eral ions in a similar manner. For example, the addition of chloride ion,
Cl”, will indicate the presence in solution of silver ion, Ag + , through the
formation of a white precipitate of silver chloride, AgCL But lead(II) ion,
Pb 2 *, will also give a similar white precipitate with CI~ ion so that its
formation is not unambiguous indication of the presence of Ag + ion. Very
few tests can be made for a given ion in a solution without consideration
of the total composition of the solution and the possible presence therein
of an ion which may interfere with a given test and yield an erroneous
result. Therefore the general procedure that is followed is to separate the
sample into small groups of ions which exhibit similar chemical behavior
to one reagent* From these groups individual ions can then be isolated and

Analysis 051

confirmed, or it may he possible to test for the ion directly in the smaller
group.

The principles im’olved in the sep.uations of qualitative analysis are
not new hut have already been considered in some detail, Among these are
the solubility product, amphoterism, and complex ion formation so that
an inquiry into the procedures of qualitative analysis affords an opportunity
to review the principles of ionic equilibria and the chemical properties of
the more common elements.

In the event that a sample to be analyzed is a solid, the first step is to
dissolve it. Various solvents are tried, in each case noting whether any portion
of the sample dissolves inasmuch as it may not be a single substance. The
solvent always tried first is water, and then the common reagents, HC1,
NaOH, NH. ( , UNO.,, and aqua regia.

HC1 dissolves active metals ami many oxides, hydroxides, carbonates, and sulfides.
NaOH dissolves a few metals, e.g., A1 and Zn, and amphoteric oxides and
hydroxides.

NH ; , dissolves compounds which form soluble ammonia complexes.

HNO;, dissolves, in udditiun to those substances dissolved by HC1, all but the
noblest metals and muny relatively insoluble sulfides.

Aqua regia dissolves all metals and sulfides, in addition to those substances dis-
solved by HC1 and HNO, individually.

In some cases, notably with silicate substances, it is necessary to resort to high
temperature fusion with NaOH or Nn,(X) a , followed by solution in one of the
foregoing solvents.

To the experienced chemist, these simple solubility experiments yield
many clues concerning the nature of a sample. In analytical chemistry, one
acts the role of a detective and :..any seemingly innocent, sometimes over-
looked, observations and procedurt s can offer pertinent information. Ob-
viously, total solution of the sample in water negates the presence of sub-
stances insoluble in water, for example, AgCl and many oxides, sulfides,
and carbonates. Certain generalities concerning the solubilities of common
salts in water are worth noting.

A) Almost all Na + , K + * and NH, + salts are soluble.

B) Nitrates, nitrites, acetates, and chlorates are generally soluble.

C) All Cl , Br, and 1“ salts are soluble except those of Ag + , Pb a+ ,
and Hg, il+ .

D) The sulfates of Ba' i+ , Sr- + , and Fb- + arc; insoluble. Those of Ca a+ ,
Ag + , and Hg,* + are sparingly soluble, while all others are -oluble.

E) The hydroxides of Na+, K+, and NH, + are soluble. Those of Ba a +,
Sr* + , and Ca' J+ are slightly soluble. All others are soluble.

F) All carbonates, phosphates, chromates, and sulfites are insoluble ex-
cept those of Na + , K + , and NH, + . Many hydrogen carbonates and dihydro-
gen phosphates are soluble.

G) All sulfides are insoluble except those of Na + , K + , NH, + , Ca a +-
Ba*+, and AI*+.

652

Analysis

From the color of a solution valuable information can be inferred. A
colorless solution rules out the presence of the following colored cations
and anions, written below in their simplest form without water of hydration.

Cu 2+

blue

Cr0 4 2 '

yellow

Ni a+

green

Cr 2 0 T -~

orange

Co 2+

pink

MnOr

purple

Fe 3+

yellow brown

Cr 3+

violet

Caution should be observed since the absence of color may be due to an
insufficient concentration of an ion, particularly those of lighter hues, while
on the other hand a dark color can mask a more pale color or the mixture
can give a resultant hue which is not specifically recognizable.

Whether a solution is acid or basic to litmus also provides information
since many ions cannot remain dissolved in solutions which are basic while
others cannot exist in neutral solution. Thus ordinary concentrations of
Mg 2+ cannot remain in simple basic solution because Mg(OH) a would
precipitate and, similarly, Fe 2+ is insoluble in neutral solution due to the
precipitation of Fe(OH) 3 .

3. Schemes of Analysis. Many schemes of analysis are feasible. For the
purpose of discussion, let us assume that we have an aqueous solution
containing thirteen elements in ionic form: Na + , Mg 2+ , AF + , K*+, Ca 2+ ,
Cu 2 +, Zn 2 + Ag+, Sn*+ Ba 2 +, Hg 2 2 +, Fe 2 +, Pb 2 +, and, in addition, NH*+
ion. Since only these common ions are to be considered an abbreviated,
though representative, scheme of analysis will suffice. The student would
do well to consider possible modifications of the scheme presented or even
the "invention” of other schemes. The detection of anions will be deferred
to a later section while the analysis of gases is so specialized that we shall
not consider it at all.

The essential procedures of a systematic scheme of analysis involve selec-
tive precipitation due to a difference in solubility product and selective
solution through complex ion formation or amphoterism. Physical separa-
tion of a precipitate or of a residue is then accomplished by filtration or
by the more rapid centrifugation.

The fourteen ions listed above can be separated chemically into five
groups depending upon their common behavior to specific reagents. For
example, only three of the fourteen form insoluble chlorides, namely AgCl,
Hg 2 Cl 2 , and PbCl 2 . Chlorides of the other ions are soluble and thus remain
in > solution. The insoluble chlorides can be removed, free of the ions in
solution, by filtration or centrifugation. On such a basis, the fourteen ions
can be classified into the following five groups:

Group 1: Cations which form insoluble chlorides; Ag 4 *, Hg 2 2+ , Pb 2+

Group II; Those remaining cations whose sulfides are insoluble in
0.3 M HC1; Cu 2+ , Sn 4 *, Pb 2+ (Pb 2+ is not precipitated com-
pletely in Group I)

Amh/w

653

Group HI: Those* remaining cations whose hydroxides or sulfides pre-
cipitate from basic solution; Al :t + , Fe 2 *, Zn 2 *

Group IV: Cations which precipitate as carbonates in slightly basic
solution; Mg 2r , Ca- \ Ba 2+

Group V: Those ions, almost all of whose compounds are soluble in
water; Xa\ K r , Nil, f

In the scheme that follows, detailed laboratory instructions are omitted.
Only such procedure is given as to indicate the principles and techniques
whereby separations are accomplished and the results are confirmed. Under-
standing of the scheme will be enhanced if an equation is written for each
chemical reaction involved in the scheme,

4. Analytical Procedure. Group I; AgCi; Hg.Cl-; PbCL. To the original
solution, 3 M HCI is added. This precipitates the insoluble chlorides, AgCl,
Hg-.Gl-, and PbClj. Centrifuging the mixture separates the precipitated
chlorides from those cations whose chlorides are soluble and remain in
solution. The latter, in the supernatant liquid after centrifuging, is labelled
collectively as Groups II-V.

Since some of the supernatant liquid, containing the Group H-V ions, adheres to
the precipitate the solid is washer! with a small volume of the HCI reagent that
was used to cause the precipitation. Such jvashing of a precipitate with the
precipitating agent is a general procedure .

The three insoluble chlorides can now be separated one fro**, the other
and confirms! individually. Of the three, only PbCb is soluble in hot
water. Hence heating tin* precipitate with water dissolves the PbCl* and
leaves the AgCl am! Hg-Ch as a residue. After centrifuging the hot mixture,
the presence of Ph- * in the supernatant solution can be confirmed by
the addition of K*CrO # solution which precipitates the characteristic yellow
PhCrO*.

To the residue of AgCl and Hg a Ch, 6Af NH a is added and the mixture
centrifuged. The AgCl dissolves as the complex ion, Ag(HHn) a * and remains
in the supernatant liquid. A gray-black residue of Jig and HgNH a Cl con-
firms the presence of Hgr 4 . By making the solution of Ag(NH : ,) 8 * acidic
with 3 M IICI, the presence of Ag ■ is confirmed by the re-formation of a
white precipitate of AgCl.

Group II; CuS; SnSj; PbS. The solution labelled Groups II-IV is neu-
tralized with MH* and then sufficient 3 M HCI is added to make its con-
centration 0.3M. The concentration of hydrogen ion is now such that when
H*S is added, the equilibrium concentration of S 2 ' ion produced is so low
that it is sufficient to precipitate only the extremely insoluble metal sulfides.
These are CuS, SnS 8 , and PbS; other ions precipitated in this group, but not
included in this scheme, are Cd if , Hg 2 *» Sn 2 *, As 8 *, As®*, Sb 8 *, Sb 5 *, and
Bi 8 *, After centrifugation, the supernatant liquid from the precipitated sul-
fides is set aside as Groups III-V.

The student may wonder why Pb 2 ^ appears again in Group II when,
apparently, it was precipitated as PbCl a in Group I. This is because, in a

654

Analysis

saturated solution of PbCl 2 , a sufficiently high concentration of Pb 2+ still
remains to form a precipitate of PbS in Group II upon the addition of H 2 S.
Hence if Pb 2+ is detected in Group I it should also be found in Group II.
It is possible, however, that the original solution may contain too low a
concentration of Pb 2+ to precipitate as PbCl 2 in Group I but still sufficient
to be detected as PbS in Group II.

A solution can be saturated with H 2 $ either by bubbling the gas through
it or by the addition of thioacetamide, CH 3 CSNH 2 . The thioacet amide acts as a
source of H 2 S.

S O

it II

CH s -C-NH 2 + H s O+ +H a O -» CHy—C— OH + NH 4 + + H a S

The SnS 2 can be separated from the black residue of CuS, SnS 2 , and
PbS by the addition of concentrated NaOH (6M) since it alone of the
three sulfides is amphoteric and dissolves as Sn(OH )« 2 ~ or as Sn$ 3 2 ",
After centrifuging, acidification with HC1, and addition of H 2 S to the super-
natant liquid, a yellow precipitate of SnS 2 is produced.

Since the yellow precipitate may be elemental sulfur from the acidification of
polysulfide ion, S 2 2 “, which may be present, more certain confirmation of tin is
obtained by dissolving the precipitate in HCl and adding a reducing agent
such as elemental iron to form SnCL. If HgCl 2 is then added, a white precipitate
of Hg 2 Cl 2 , which may turn gray-black confirms the presence of Sn 4+ in the
original sample.

The remaining solid sulfides of Group II, CuS and PbS, are dissolved
in HN0 3 (3 Af). By the addition of dilute H 2 S0 4 (1 M), the lead is precipi-
tated as PbS0 4 , leaving a blue solution of Cu 2 + . Lead can be confirmed as
in Group I as PbCr0 4 , while the formation of a deep blue color upon the
addition of NH 3 to the Cu 2+ solution confirms its presence as Cu (NH 3 ) 4 2+ .

Group III: Al(OH) 3 ; FeS; ZnS. From the solution labelled Groups III-V,
aluminum, iron, and zinc can be precipitated from ammoniacal solution with
H 2 S either as hydroxides or as sulfides. When the solution is made basic
with NH S and H 2 S is added, AI(OH)*,, FeS, and ZnS precipitate. These
compounds are not so insoluble as the Group II sulfides and require a higher
concentration of S 2 ~ and a basic solution in order to precipitate. Other ions
that would also precipitate in Group III are Cr* + as Cr(OH) :h and Co 2+ ,
Ni 2+ , and Mn 2+ as sulfides. After centrifuging, the supernatant liquid
contains the ions of Groups IV-V.

The Group III precipitate is dissolved in dilute HCl (3M) and boiled
to drive out excess H 2 S. Addition of excess NaOH (6M) and the oxidizing
agent H 2 0 2 precipitates the iron as red-brown Fe(OH) 3 . Both aluminum
and zinc hydroxides are amphoteric and remain in solution as Al(OH)*~
and Zn(OH) 4 2 -, respectively. The Fe(OH) a is removed, dissolved in
HCl, and KSCN is added. The appearance of the blood-red color of
the complex ion, Fe(SCN) e 3 “, confirms the presence of Fa 8 * (Fe 2 * in
the original sample). The alkaline solution of Al(OH) 4 ~ and Zn(OH) 4 2 " is
acidified to form the simple ions, Al 3+ and Zn 2+ Addition of NH* {3M)

Analysis

655

then precipitates the gelatinous, white AI (OH)*, leaving the Zn(NH a ) 4 a +
ion in solution. Confirmation of the zinc is through the formation of a
white precipitate of ZnS upon addition of H,S to the supernatant solution.

Group IV: MgCO*; CaCO t ; BaCO.i. From the solution set aside as
Groups IV-V, magnesium, calcium, and barium can be precipitated as MgCO ;t ,
CaCOt, and RaGO, from ammoniaca! solution by (NH 4 ) s CO a . The carbon-
ates of the Group V ions are soluble and remain in the supernatant liquid
when the mixture is centrifuged.

If an excessive concentration of NH, 4 ion is present, Mg 24 is never completely
precipitated in this group. The concentration of the CO a *- ion can be reduced
by excess NH 4 4 to such a point that Mg* 4 is incompletely or not at all pre-
cipitated as MgCO v Therefore Mg 24 is frequently tested for in Group V.

The precipitated carbonates of Group IV are dissolved in dilute acetic
add, HCgHjOv* From this solution Ba- 4 is precipitated as yellow BaCrO,
by the addition of K-CrO,. The barium can be confirmed by the pale green
color of a flam? test which can be made upon the precipitate or its solution
in HC1.

A flame test involves the use of a platinum or chrome! wire generally mounted
in one end of a glass tube. Before using it for a flame test, the wire is cleaned
by dipping it in concentrated HC1 and then holding it briefly in the Bunsen flame.
The procedure is repeated until no color is imparted to the flame. The wire is
then dipped into the substance to be tested, which may be a solid paste or a
solution, and then again inserted into the flame. The color, if any, now imparted
to the flame, is characteristic and gives evidence of a specific metal. This color
is due to the emission of characteristic spectral frequencies. Flame tests can be
used for only those elements whose atoms have ionization potentials sufficiently
low that the energy of a Bunsen flame can excite them. The characteristic colors
imparted by some metals are;

Lithium: bright red Cesium: violet

Sodium: bright yellow Calcium: brick red

Potassium: violet Strontium: carmine red

Rubidium: violet Barium: pale green

The acetic acid solution of Ca 24 and Mg 24 is neutralized with NH*
Upon the addition of ammonium oxalate, (NH 4 )*C a 0 4 , to this solution, a
white precipitate of calcium oxalate, CaC a 0 4 , forms; MgCaO* is soluble.
Calcium can then be confirmed by the brick red color of a flame test. From
the ammonlacal solution. Mg 24 is precipitated as the highly crystalline, white
magnesium ammonium phosphate, MgNH»PO*, by the addition of Na 2 HPO,.

Group V: Na 4 , K 4 , NH* 4 ions. In the Group V solution, there remain
those ions whose chlorides, sulfides, hydroxides, and carbonates are all
soluble. The NH* 4 ion is included in this group because many of its chemical
reactions are similar to those of K 4 . The NH* 4 is always detected, however,
by an independent test upon the original solution because of the use of
NH a as a reagent throughout the scheme. To a small portion of the original
solution concentrated NaOH (8 M) is added. Upon warming, NH ;t vapor
is evolved and can be recognized by its characteristic pungent odor; also

656

Analysis

a strip of moist red litmus paper held in the path of the escaping vapor
turns blue.

The Na + and K + ions are tested for in the Group V solution independ-
ently by specific reagents without separation. The Na + is detected in
acetic acid solution through the formation of a yellow precipitate upon
the addition of zinc uranyl acetate, while K + is determined as a yellow
precipitate formed upon the addition of sodium cobaltinitrite. Since NH t **
gives an identical reaction it must be removed prior to the test for K 4 .
Flame tests can also be used for K* and Na+ but caution must be observed
since sodium compounds are present as an impurity in almost all sub-
stances and the bright yellow color which they impart to a flame often
masks the color formed by potassium. To detect K + , the flame is usually
viewed through a blue cobalt glass since the yellow light of the sodium
flame is absorbed by such a glass but the violet color emitted by potassium
is transmitted. The test for sodium is so sensitive, and this element is so widely
distributed, that only a lasting bright yellow flame is evidence of its presence
in a sample undergoing analysis.

Though obvious, it should be noted that if a reagent added initially to
precipitate an entire group yields no precipitate, it is unnecessary to go
through the procedures for the individual ions of that group and the analyst
may proceed directly to the next group. Thus, if the addition of H 2 S to the
Group II solution results in no precipitate, then no Cu 2 * 4 *, Sn 4+ , or Ph 2 ’ 1 '
is present and these ions need not be tested for individually; the analyst
should then go directly to the procedure for Group III.

5. Anion Analysis. Though systematic schemes of analysis for anions
have been developed, it is simpler in this case to make individual tests which
indicate the presence or absence of specific anions. Tests for anions are
usually made upon the dry salts or freshly prepared solutions since many
anions react with each other in aqueous solution. The few examples cited
below are representative of reactions used to test for anions.

When the anions, C0 3 2 ", S 2 ~, SO» 2 “, or NCV are treated with a strong
acid, such as H 2 S0 4j gases are formed. These are, respectively, CO s , H 2 S,
S0 2 , and NO. In addition to the observation of effervescence in each case,
the characteristic odors of H 2 S and SO> can be detected while NO forms
brown NO s upon mixing with the oxygen of the air. Many of the tests used
for confirmation of cations can also be used for the detection of anions, but
in the reverse sense. Whereas Cl- ion was used to precipitate Ag + ion in
Group I, ^conversely Cl“ can be identified through the formation of a white
precipitate of AgCl upon the addition of Ag + , Similarly, Bi r ion gives a
pale yellow precipitate of AgBr and I" ion a yellow precipitate of AgL Other
confirmatory precipitates used in cation analysis which can also be used
for the identification of anions are CaC 2 0 4 , BaCO a , MgNH 4 PG 4 , BaS0 4 PbS,
and PbCrO*.

A tuibjsh

657

QUESTIONS

1. List propeities of matter which lend themselves to techniques of instrumental
analysis.

2. Devise analytical instruments based on (a) heat conduction of gases (b) rela-
tive tendency to vaporise from an adsorbing material.

3. Outline in tabular form the scheme of analysis for the fourteen ions in Section 3.
Indicate the group leagent and the separations within a group. Write chemical
equations whete applicable.

4. Suppose HU! were not added to precipitate the ions of Group I but an
analyst proceeded directly with Group II. How would this affect the determina-
tion of the Group if ions?

5. (a) Calculate the concentration of Pb 2t in a saturated solution of PbCL.
(b) What concentration of S 2 ion is required to form a precipitate of PbS with
the l’b 2 * concentration of part (a)? (e) what is the maximum Pb 24 * con-
centration that would not be detected in Group I but would be found in
Group II?

6. Write chemical equations for the separation of the following pairs of ions:
(a) life 8- and fig 2 * (b) *V» • and AH* (c) Cu-’f and Pfe 24 * (d) Hg 2 +
and (»*' te) AH * and fti« ' .

7. A solution contains Pb 2 ‘ , F e* * , and Ba- r ions. Devise a scheme for the
isolation and confirmation of each ion present.

B. A solution contains the following cations: Hg a - M , Cu 2 *\ Zn 24 *, Ca 24 *, and
K * . Outline .i scheme of analysts for the detection and confirmation of each.

9. By what means could the following solids be brought into solution for
analytical purposes (a) zinc sulfide <b) mercury (II) sulfate (c) iron(III)
oxide tdl lead (ID carbonate (e) silver nitrate (f) copper (II) oxide?

10. Why does AHOH? . titul not ALS ; precipitate in Group III?

!L The chloride of a metal, M, is soluble. The sulfide has a value of 2 x 10 5 5
and the carbonate a value of 2 *• 10 *\ In which analytical group could
this metal he determined? Explain your choice,

12, In which groups could the following ions he determined: Cd 24 *, NF + ? Which
sulfide has the smaller solubility, CdS or NiS? Explain.

13. (a) Why is not Na.UO* used to precipitate the ions of Group IV rather
than (NHjKUO, |b) why is the test for NH t ' done on a separate portion
of the unknown?

14, Write chemical equations for the separation of the following groups of
anions: (a) CO, 2 , C.1 , and UrO, 2 (b) CO a ‘-“ and S 2 ' (c) Br- and S0 4 H

15. A mineral is suspected of being a mixture of galena, PbS, and limestone,
CnCQ v Outline a met lux! for the determination of each ion which might be
present.

Appendices

APPENDIX I
Mathematical Operations
Exponential Arithmetic:

Frequently very small numbers, such as 0,0000032 and 0.000000065, are
encountered in problems. Such numbers are difficult to handle, since it is
difficult to locate the decimal point when arithmetical operations are per-
formed. Some times numbers in chemical measurements are very large,
such as the number of molecules in a mole, 003,000,000,000,000,000,000,000.

In order to express such numbers, exponentials, or powers of 10, are used.
The exponent (power) represents the number of times 1 must be multiplied
by 10. Thus 10 = 1 X 10 = 10* , 1000 « 1 X 10 X 10 X 10 = 1 X 10»
as 10*. 100,00ti = 10* and so on. But what can be done with the small
numbers? Consider a series of numbers and their exponentials; 10,000 = 10*,
1000 ® 10", 100 as 10*. 10 as 10*. Each time we divide by 10, the exponent
becomes 1 less; hence, we may continue the table thus; I - 10°, 0.1 = 10-*,
0.01 as 10**, 0.001 ~ 10-*, and so on. A negative exponent, then, shows how
many times we divide 1 by 10.

A number, such as 15,000, can be separated into a power of 10 and an-
other number, as 15 X 1000 (or 15 X 10’) or 1.5 X 10,000 (1.5 X 10 4 ).
Likewise, 0.000015 — 1.5 X 0.00001 (or 1.5 X 10* 5 ), and so on. For num-
bers greater than 1, the exponent is positive and is equal to the number of
places the decimal point has been shifted to the left to form the coefficient.
Thus 15,000 = 15,000 X 10° = 1.5 X 10 4 , as shown above; the decimal
point has been shifted 4 places to the left to form the coefficient.

For numbers less than 1, the exponent is negative and is equal tp the
number of places the decimal point has been shifted to the right to form the
coefficient. Thus 0.000015 = 1.5 X 10**; the shift is 5 places to the right.
Whenever the decimal point is shifted one place to the right, it is equivalent

ApjH'iuticcb

to multiplying by 10 or adding 1 to the exponent; thus 0.000015 X 10
= 0.00015 = 1.5 X 10+ Whenever the decimal point is shifted one place
to the left, it is equivalent to dividing by 10 or subracting 1 from the ex-
ponent; thus 0.000015 ~ 10 = 0.0000015 =r 1.5 X 10 *

The multiplication, 100 X 100,000 — 10,000,000, written as exponentials,
is 10 2 X 10 5 = 10?. Note that 2 + 5 = 7; that is, in multiplication exponents
are added.

The division, 100 ~ 100,000 = 0.001, written as exponentials, is 10- -f* 10 r ‘
— 10~ 3 . Note that 2 — 5 = —3; that is, in division exponents are subtracted .

In the multiplication and division of large and small numbers, the num-
bers are changed to the exponential form. The coefficients and exponen-
tials are then operated on separately:

800,000 X 35,000,000 X 0.0007
15,000 X 0.021

8 X 10* X 3.5 X 10 7 X 7 X 10~ 4
1.5 X 10 4 X 2.1 X 10-®

x io r, + T -*- 4 + 2

1.5 X 2.1 x

V M« = fiOO. V IftT

Logarithms

The logarithm of a given number is the exponent of the power to which a
selected number, called the base, must be raised in order to produce the
given number. The base is usually 10; hence the logarithm of a number, is
the exponent of the power to which 10 must be raised in order to produce
that number. Thus, since 1000 = 10 3 , the logarithm of 1000 is 3. This is
written as log 1000 = 3. It should be noted that log 1 = 0, and that
log 10 (the base) = 1.

Furthermore, since 0.1 = ^ = 10~\ log 0.1 = -1. Similarly, it is ob-
vious that log .01 = -2, and log .001 = -3, and so on.

The logarithm of any number consisting of 1, preceded or followed by
zeros, is a whole number and can be obtained by finding the power to which
10 must be raised in order to produce the number. The logarithms of other
numbers cannot be obtained directly and must be found from a specially
computed table (the log table).

The logarithm of any number between 1 and 10 is a decimal, for it is more
than 0 and less than 1. The logarithm of any number between 10 and 100
is more than 1 and less than 2; therefore it is 1 plus a decimal. The loga-
rithm of any number between 100 and 1000 is 2 plus a decimal, and so on.
A logarithm, therefore, usually consists of two parts: a whole number, called
the characteristic , and a decimal, called the mantissa. The characteristic
may be positive or negative. When it is negative, the minus sign (-) is
written over the characteristic, that is, 1. The mantissa is positive in every
case. The logarithm of 2 is 1.3010 (+1+0.3010), but the logarithm of 0.2
is 1,3010 (-1+0.3010).

The characteristics of logarithms can always be determined by the fol-
lowing rules:

Appendices

661

Rule 1. For <i number "renter than I, the characteristic is one less than the
number of integral places in the number. ( By integral places is meant the
figures in a whole number, or in that part of a numlx'r that is to the left of
the decimal point.)

Rule 2. For a number wholly decimal, the characteristic is Motive and is
numerical/^ ow greater than the number of zeros immediately follcncing the
decimal point.

For example, according to rule 1. tin* ^characteristic of the logarithm of
874 is 2, because 674 is a whole number having three integral places, or
figures, and 2 is 1 less than 3, Similarly, the characteristic of log 67.4 is I
because there art* two integral places, anti the characteristic of log 6.74 is 0.

According to rule 2. the characteristic of log 0.674 is 1, since no zero im-
mediately follows the decimal point. Likewise, the characteristic of log
0.00674 is 3, since there are two zeros to the right of the decimal point

Finding the Logarithm of a Number

Appendix X is a four-place logarithm table. From the previous dis-
cussion the characteristic of the logarithm of a number can lx; readily ascer-
tained.

The mantissa is obtained from the table in the following way: Note that
the mantissa for the logarithms of 2, 20, 2000, or 20,000,000 is the same,
that is 0.3010, because in determining the logarithms of the above numbers,
you locate the numlter 20 in the first column in the table under No. and
then look under the column 0 and find the mantissa 3010. Determine the
characteristic and write the complete logarithms, that is, 0.3010; 1.3010;
3.3010; 7.3010. If the number is 201, the mantissa is found by locating the
number 20 in the first column and then moving over to the right under the
column headed 1 where you find the mantissa for 201, which is 3032. If
the number has four digits, as 4675, the logarithm is determined by locating
the number 46 in the first column; then moving over to the right under the
heading 7, we find 6093, which is the mantissa of 467.

But it is necessary to add to this an additional amount to account for
the digit 5 in the fourth place of the number. This is located by moving
over to the second set of columns on the right side of the table. Under the
heading 5 and on the same line as 46 in the first column we find the number
5, This is added to the mantissa of 467 to obtain the mantissa of 4675, or
8663 plus 5 equals 6698. The logarithm of 4675 is 3.6698.

Finding the Number Corresponding to a Logarithm

After on has mastered the method of finding the logarithm of a number
from the tables, it is not difficult to learn how to find the number corre-
sponding to a logarithm. Let us start with u definite problem, such as:
what is the number whose logarithm is 2.7195? The procedure is as follows:

Glance down the column headed 0 until yon find the first two figures of
the mantissa, tliat is, 71. Note that they appear beside number 52 as the
first two figures of 7160. Now run along the line to the right until you find
the next two figures, or the mantissa closest to the given one. Note that

Appendices

This corresponds to the number 524, but we have an additional amount
to account for, that is, the difference between the given mantissa and that
found in the table. The difference is 7195 minus 7193 or 2 in the fourth
place. Now glance along the line across the page until you find the num-
ber 2 in the second set of columns containing the additional quantities corre-
sponding to the fourth place in the number. Note that the figure 2 appears
rnider the headings 2 and 3. Here you have your choice of selecting either,
depending upon the fifth figure of the original number. Let us select the
first 2, under heading 2. The number at the head of the column must be
added to the number 524 as the fourth figure, thus: 5242. Since the
characteristic of the logarithm is 2, the number whose logarithm is 2.7195
is 524.2. The number 524.2 is the antilogarithm of 2.7195.

Appendix XI is a four-place antilogarithm table. Such a table is used
in the same way as logarithm tables, except that the first two digits of the
mantissas are read as decimals in the column labeled Log., the third digit
appears as the headings of the other columns, and the fourth as the head-
ings of the second set of columns on the right side of the table. The
antilogarithm appears within the table, just as the logarithm does in the
table of logarithms (Appendix XI). The characteristic determines the posi-
tion of the decimal point.

Multiplying and Dividing by Logarithms

The product of a factor to one power and the same factor to another
power is equal to the factor raised to the sum of the powers, that is,
10 4 X 10 2 = 10°. Since logarithms are the powers to which the number 10
must be raised to obtain the given number, when two numbers are multi-
plied, the logarithm of the product is equal to the sum of the logarithms of
the two numbers. Let us use a very simple illustration. The product of 4
and 3 is 12, Using logarithms, we have log 4 plus log 3 equals log 12.

log 3 = 0.4771
log 4 = 0.6021

sum = 1.0792 = log (3 X 4).

Consulting the logarithm tables, we find that the number corresponding
to logarithm 1.0792 is 12.

Rule for multiplying. To multiply two or more numbers by using loga-
rithms , add the logarithms of the numbers , and the sum will be the logarithm of
the product. Find the number corresponding to this logarithm, and the result
wiU be the answer.

In dividing a number raised to a power by the same number raised to an-
other power, the quotient is the number raised to a power which is the
difference between the two original powers, thus: 10 6 10 2 = 10 4 . Since

logarithms are the exponents, the logarithm of a quotient is the logarithm of
the dividend minus the logarithm of the divisor. Again let u$ use. a simple
illustration. Divide 84.12 by 12. Using logarithms, we have, log 84.12 minus
log 12 equals log quotient.

Appendices

log 84,12 r- 1.9249
log 12 1.0792

difference 0.8457 -- log (84.12 - 12).

The number corresponding to logarithm 0.8*157 is 7.01, which is the answer.

Rule for dividing. To divide one number by another by means of loga-
rithms , subtract the logarithm of the divisor from A he logarithm of the dividend,
and the difference will be the logarithm of the quotient. Find the number corre-
sponding to this logarithm , and the result will be the quotient.

For more complex arithmetical operations by logarithms, consult a mathe-
matics textbook or your instructor.

Significant Figures

A measurement of the weight of an object can be made with different
degrees of precision. The object can be weighed on a rough balance, and
it will be noted at a glance that it weighs about 86 g. We are certain that
its weight is between 85 and 87 g. The statement that the weight is 88 g
means that it lies between 86 4 1 and 88 - 1 g. We signify this by writ-
ing 88 ± l g. If we take pains to weigh the object on a better balance we
can state that it weighs 86.4 g or that its weight lies between 86.3 and 86.5 g
or better still that it weighs 86.4 ± 0.1 g. The weight 86.4 g indicates greater
precision than 86 g, though both are correct. By weighing very carefully
on a still more sensitive balance we find that the weight of the object is
86.37 g, or, to better indicate the weight found, we should write 86.37 ± 0.01 g.
The last measurement of the weight, that is, 86.37 g is of still higher precision
than the weighing which gave 86.4 g.

Note that the weight is expressed as the number of units of weight ( gram
or fraction of a gram) which must be put together to yield a weight equal
to that of the object being weighed. The number 86.37 contains number
characters or digits in fotir places. It is a number with four significant figures.
Each significant figure in its place denotes a number of grams, or of tens
of grams, or of tenths or hundredths of a gram. The weight when expressed
as 86 g has but two significant figures and is of lower precision. The precision
of a measurement is indicated by the number of significant figures. If the
weight were expressed as 0.08637 kg, or as 8.637 eg, it would still be a
measurement of four significant figures and of the same precision as 86.37 g.
The zeros in the number 0.08637 are not significant. They are used only to
locate the decimal point. A zero used to denote zero amount in a definite
place is, however, significant. For example, 5.43 cm denotes a value between
5.42 and 5.44 cm, but 5.430 denotes a value between 5.429 and 5.431 cm.
The zero is therefore significant. Some confusion arises from a value such as
86370 mg. As it is stated one does not know whether the zero is significant.
For this reason it is better to express numbers with the first digit in the
units place multiplied by a power of ten, 8.637 X 10* g.

Measurements of other properties, such as length, volume, temperature,
density, and the like, are also made with different degrees of precision.

664

Appendices

Such precision is indicated by the number of significant figures in the numeri-
cal expression.

In the solution of chemical problems, the following rules should apply in
the calculations:

Rule 1. A numerical value should have but one uncertain figure. Thus,
if a weight were known to be 86.4 dfc 0.1 g, it would be of no value to put a
zero or any other figure after the 4.

Rule 2. If less precision is required and figures are to be dropped from
a more precise value, increase the last retained figure by 1 if the digit fol-
lowing is 5 or greater. If it is less than 5, leave the last retained figure as
it is. Thus, if we know the weight of an object to a precision indicated by
86.37 g, but need its value to only three significant figures, we would assign
it a value of 86.4 g. If, however, we need the value with only two significant
figures, it would be taken as 86 g.

Rule 3. In adding or subtracting values, retain the number of significant
figures in the sum that appears in the numbei with the least significant
figures.

Rule 4. In multiplying or dividing, retain enough figures in the result
so that a difference of one in the last place yields a percentage error no
greater than that made by a difference of one in the last place in the least
precise value used in the computation.

665

Appendices

APPENDIX II

Prefixes of the

Metric System

kilo

1000

one thousand times the basic unit

hecto

100

one hundred times

deka

ten times

—

deci

0.1

one- tenth of the basic unit

centi

0.01

one-hundredth

milli

0.001

one-thousandth

micro

0.000001

one-millionth

Length

The Metric System:

1 meter (lm) = 10 decimeters (10 dm) = 100 centimeters (100 cm) = 1000
millimeters (1000 mm) = 39.37 in = 1,094 yd
1 kilometer (1 1cm) = 1000 meters = 0.6214 mile

Volume

1 stere = 1 cubic meter = 1000 cubic decimeters = 1.308 cu yd
1 cubic decimeter = 0.0353 ft* = 61.023 in 8 = 1000cm 8

Capacity

1 liter is the volume of pure water at 4°C and 760 mm pressure which weighs
1 kilogram. 1 liter = 1.000027 cu dm = 1000.027 cm 8
1 liter = 1000 milliliters =s 100 centiliters = 10 deciliters
1 liter = 33.31 fluid ounces = 1.056 liquid quarts
1 liter = 1.816 dry pints = 0.908 dry quarts

Weight

1 gram (g) = weight of 1 cm 8 of water at 4°C = 0.035 oz avdp.

1000 g ss 1 kilogram (kg) aas 2.205 lbs avdp.

1000 kilograms = 1 metric ton = 2205 lbs avdp.

1 2b avdp* ss 453.6 g
1 oz avdp. a= 28.35 g

666

Appendices

APPENDIX III

Vapor Pressure of Water

Temperature , °C

Vapor

Pressure , mm

Temperature, °C

Vapor

Pressure, mm

4.6

28.1

6.5

29.8

8.0

31.5

8.6

33.4

9.2

35.4

vflB

9.8

37.4

10.5

39.6

msm

11.2

41.8

iEM

11.9

55.0

12.7

92.2

■9

13.6

149.2

14.5

233.8

15.4

355.5

16.3

526.0

17.4

100

760.0

18.5

101

787.5

19.7

102

815.8

20.9

103

845.1

22.2

150

3581.

23.6

200

11588.

25.1

230

20925.

26.5

APPENDIX IV

Vapor Pressures of Ice and Liquid Water at Temperatures

Below 0°C

Temperature, °C

Vapor Pressure ; mm

Ice

Water

1.95

2.15

3.01

3.16

3.28

3.41

3.57

3.67

3.88

3.96

4.22

4.26

4.58

Appendices

667

APPENDIX V

Specific Heat

Substance

Temperature

°C

Gram-

Molecular

Weight*

Specific

Heat, C ••
cal/g deg

Molar

Specific Heat***
cal/mole deg

Argon

39.95

0.1253

5.004

Helium

-180

4.003

1.25

5.004

Oxygen

32.00

0.2178

6.970

Nitrogen

28.02

0.2477

6.941

Hydrogen

2.016

3.389

6.832

Carbon dioxide

44.01

0.1989

8.754

Sodium

22.99

0.285

6.555

Lead

207.2

0.0305

6.320

Uranium

238.0

0.0276

6.56

Silver

107.9

0.0558

5.93

Gold

197.0

0.0310

6.113

Iron

55.85

0.106

5.914

Water (ice)

-10

18.02

0.487

Water (liquid)

18.02

1.000

18.02

Water (steam)

110

18.02

0.481

'♦Cram-atomic weight in the case of monatomic molecules and metallic dements.
♦♦This is the specific heat under conditions of constant pressure (page 86),
‘♦♦Specific heat per gram-atom for monatomic molecules and metallic elements.

Appendices

APPENDIX VI

Ionization Constants at 25° C

Substance

Reaction

^ionization

Acetic acid, CH 3 COOH

CH 3 COOH -» H+ 4- CH 3 COO-

1.75 x 10-5

Benzoic acid, C 6 H 5 COOH

C 8 H 5 COOH H+ + C g H 5 COO

6,3 x 1 (H

Boric acid, HBO s

hbo 2 -* h+ + bo 2 -

6.4 x lO-io

Carbonic acid, H 2 CO s

H 2 C 0 8 H+ + HCO 3 -

HCCV H+ + CO 3 2 -

3.0 X 10-*

6.0 x 10-11

Hydrocyanic acid, HCN

i HCN H+ + CN-

4.0 X IO -10

Hydrofluoric acid, HF

HF — > H+ + F“

7.0 X 10-1

Hydrogen peroxide, H 2 0 2

H 2 0 2 -» H+ + HOG"*

2.4 x 10-1*

Hydrogen selenide, H 2 Se

H 2 SE H+ + HSe-

1.9 X 10-«

Hydrogen sulfide, H 2 S

H 2 S H+ + HS-
HS- H+ + S 2 -

1.1 x 10-T
1.0 x lO-i-i

Hydrogen telluride, H 2 Te

H 2 Te -> H+ + HTe-

2.4 x 10-3

Hypobromous acid, HOBr

HOBr -> H+ + OBr~

2.0 X 10-®

Hypochlorous acid, HOC1

HOC1 -* H+ + OC1-

5.6 x 10-8

Nitrous acid, HNO a

HN0 2 -» H+ + N0 2 -

4.5 x 10-1

Oxalic acid, H 2 C 2 0 4

H 2 C 2 0 4 -* H+ + HC 2 0 4 -
HC 2 0 4 - -> H+ + C 2 0 4 *-

3.8 X 10-*
6.2 x 10-5

Phenol, C G H 5 OH

C 8 H 5 OH H+ + C e H s O-

1.3 x 10-i«

Phosphoric acid, H 3 P0 4

H 3 P0 4 -+ H+ + H 2 P0 4 -

h 2 po 4 - -> H+ + hpo 4 *-

HP0 4 *- H+ + P0 4 s-

7.52 x 10-®
6.23 X 10- 8
3.6 x 10*i®

Phosphorous acid, H 3 PO a

h 3 po 3 -» h+ + h 2 po 3 -

H 2 PO s - H+ + HPO s 2-

7.0 x 10-®

2.0 X 10-5

Propionic acid, C 2 H 6 COOH

c 2 h 5 cooh -* h+ 4 - c 2 h s coo-

1.4 X 10-«

Sulfurous acid, H 2 SO s

h 2 so 3 -» H+ + HSO s -
HSO s - H+ + SO,*-

1.2 X 10-*
.5.0 X 10-«

Ammonia, NH 8

NH 8 + H 2 0 -»• NH 4 + + OH-

1.77 X 10*5

Aniline, C 6 H 5 NH 2

c„h s nh 2 + h 2 o -> C 8 H 5 NH 3 + + OH-

4.2 X 10-w

APPENDIX VII

SOLUBJULXTY PRODUCT CONSTANTS,

Substance

Aluminum hydroxide, Al(OH) a
Barium carbonate, BaC0 3
Barium chromate, BaCrO*

Barium fluoride, BaF,

Barium sulfate, BaS0 4
Calcium carbonate, CaCO s
Calcium fluoride, CaF 3
Calcium hydroxide, Ca(OH),
Calcium oxalate, CaC a 0 4
Cadmium hydroxide, Cd(OH) 2
Cadmium sulfide, CdS
Cobalt (II) hydroxide, Co(OH) a
Cobalt (III) hydroxide, Co(OH) s
Cobalt (III) sulfide, CoS
Copper (I) chloride, CuCl
Copper (I) bromide, CuBr
Copper (I) iodide, Cul
Copper (II) carbonate, CuCQ s
Copper (II) hydroxide, Cu(OH) 2
Copper (II) sulfide, CuS
Iron(II) carbonate, FeCO,

Iron (II 5 hydroxide, Fe<OH),

Iron (II) sulfide, FeS
Iron (HI) hydroxide, Fe(OH) s
Lead (II) carbonate, PbC0 3
Lead(II) chromate, PbCr0 4
Lead (II) chloride, PbCi a
Lead (II) bromide, FbBr a
Lead (II) iodide, Pbl*

Lead(II) hydroxide, Pb(OH) 2
Lead (II) sulfate, PbS0 4
Magnesium fluoride, MgF 2
Magnesium hydroxide, Mg(OH) a
Magnesium oxalate, MgC*G 4
Manganese (II) hydroxide, Mn(OH)
Manganese(II) sulfide, MnS
Nickel (II) carbonate, NiCo s
Nickel(II) hydroxide, Ni(OH)*
Nickel (It) sulfide, NiS
Silver chloride, AgCl
Silver bromide, AgBr
Stiver iodide, Agl
Silver chromate, Ag 2 Cr0 4
Silver sulfide, Ag*S
Strontium carbonate, SrCO*
Strontium sulfate, SrSO*

Tin(II) hydroxide, Sn(OH) 2
Tin (II) sulfide, SnS
.Zinc hydroxide, Zn(OH)*

Zinc sulfide, ZnS

K, p

1.9 x

8.1 x HH

2.4 x HHO

1.7 x 10-e

1.1 x KF°

8.7 x io-o

3.9 x 10-n

7.9 x l0-«
2.6 x 10-»

1.2 x 10- 1 *
7.1 x 10-^8

2.0 x lO-io

2.5 x KM*

3.0 X 10- 2 «

1.8 x 10-*

5.3 x IQ-*

5.0 x 10-12

1.4 x lO-io

5.5 x 10-20

3.5 x HH*

2.1 x lO-n

1.6 x 10- IS

3.7 x lO-io
4.0 x 10-88
3.3 x 10-1*

1.8 x lO-K

1.7 x l0- ft

6.4 x 10-«

8.7 x 10-o
2.6 x 10-18

1.8 X 10-8

6.5 x 10 -o
3.2 x lO-ii

8.6 x 10-8

4.5 X 10-i*
1.4 x 10-is

1.4 x 10- T

1.6 x lO-i*
LI x 10-**
1.8 x lO-io

7.7 x 10-18

1.5 x 10-18

8.8 x lO-i*

1.6 x 10 -*o
1.6 x ID- 9

2.8 x 10~ T

4.9 X 10-28
8.0 x 10-»
1.8 x 10-i*
l£ X ID - *®

670

APPENDIX Vm

Dissociation Constants, K d , Fob Complex Ions

Equilibrium

Ag(NH s ) 2 + Ag+ + 2 NH 3

Cd(NH 3 ) 4 2 + Cd 2 + + 4 NH S
Cu(NH 3 ) 2 + Cu+ + 2 NH 3

Cu(NH 3 ) < 2 + Cu 2 + + 4 NH 3

Co(NH s ) 3 2 + ^ Co 2 + + 6 NH 3

Co(NH 3 ) s *+ Co 3 + + 6NH,

Ni(NH 3 ) 4 2 + Ni 2 + + 4 NH 3

Ni(NH 3 ) 6 2 + Ni 2+ + 6 NH 3

Zn(NH 3 ) 4 2 + Zn »+■ + 4 NH 3

CuCI 2 - Cu+ + 2 Cl-
HgCl* 2 - Hg 2 + + 4 Cl-
Ag(CN) 2 - ;=± Ag+ + 2 CN-
Cd(CN) 4 2 - Cd 2 + + 4 CN-

Cu(CN) 2 - Cu+ + 2 CN-
Fe(CN) s 3 - Fe*+ + 6 CN-

Fe(CN) 6 <- Fe 2 + + 6 CN~

Zn(CN) 4 2- Zn 2 + + 4 CN"

Ag(S 2 O s ) 2 *- Ag+ + 2 S 2 0*

Fe(CNS) 6 3 - ;=± Fe®+ + 6 CNS-
Zn(OH) 4 2- Zn 2 + + 4 OH-

K d

6.8

1 ( H *

2.5

10-7

1.3

10 -“

4.6

10 -“

1.4

10- 5

2.2

10-34

4.8

10 -s

2.1

10 -s

1.0

10 -*

2.9

10-8

1.1

10-16

3.8

10-18

1.4

10-17

1.0

10 -“

1.0

10 -*“

1.0

10-37

1.3

10-17

4.2

10 -“

3.1

10 ~*

3.5

10-18

Appendices

671

Alloys

Composition

Admiralty metal

Aluminum bronze

Babbitt metal
Bel! metal
Brass (red)
(yellow)
Britannia metal
Constantan
Duralumin

Dowmetal D

Fusible metals
(Lipowitz)

(Bose)

(Woods)

Gun metal
Magnalium

Manganese bronze

Monel metal

Pewter
Solder (soft)
(medium)
(hard)

Speculum metal
Type metal

Cu 70%, Zn 29%, Sn 1%

Cu 90%, A1 10%

Sn 90%, Sb 7%, Cu 3%

Cu 78%, Sn 22%

Cu 90%, Zn 10%

Cu 67%, Zn 33%

Sn 90%, Sb 8%, Cu 2% j
Cu 60%, Ni 40%

Ai 95.5%, Cu 3%, Mn 1%,
Mg 0.5%

A! 8.5%, Mn 0.15%,

Cu 2.0%, Cd 1.0%,

Zn 0.57, , Mg 87.85%

Bi 50%, Ph 27%, Sn 13%,
Cd 30%

Bi 50%, Pb 27.1%,

Sn 22.9%

Bi 50%, Pb 25%,

Sn 12.5%, Cd 12.5%

Cu 90%, Sn 10%

Al 90%, Mg 10%

Cu 90%, Zn 5%, Sn 3%,
Mn 2%

Ni 72%, Cu 26.5%,

Fe 1.5%

Sn 75%, Pb 25%

Pb 67%, Sn 33%

Pb 50%, Sn 50%

Pb 33%, Sn 67%

Cu 67%, Sn 33%

Pb 82%, Sb 15%, Sn 3%

marine fittings; condenser
tubes for use with salt
water

hard, non-corrodible;
equipment exposed to cor-
rosive liquids
anti-friction linings
bells, gongs, etc,
gold paint, cheap jewelry
tubes, sheets, cartridges, etc.
cheap tableware
thermocouples
airplane and automobile
parts

light metal of high tensile
strength

fuse plugs in automatic
sprinkler systems

gears, castings, etc.
balance beams, light instru-
ments

propeller blades on ships

propeller blades, wire, sheet,
pipe* etc.; non-corrodible
cups, mugs, etc.
plumbers" solder

takes high polish, reflectors
casting type

672

APPENDIX X,

Logarithms

Appendices

No.

“IT

0 1

TTTT

tthtt

6 1

tttt

tit

"67

~~7~ ~8 “9

0000

3043 (

)086(

3128

)170|0212

0253

0294

0334

0374

12 I

17 21

25 I

29 33 37

0414

3453

3492 <

3531

356910607

0645

0682

0719

0755

8 11

19 23

26 30 34

0792

3828

3864

3899<

>93410969

1004

1038

1072

1106

7 10

14 17 21

24 28 31

1139

1173

1206

1239

1271|1303

1335

1367

1399

1430

6 10

13 16 19

23 26 29

1461

1492

1523

1553

158411614

1644

1673

1703

1732

9 I

12 15

18 |

21 24 27

1761

1790

1818

1847

1875

1903

1931

1959

1987

2014

11 14 17

20 22 25

2041

2068

2095

2122

2148

2175

2201

2227

2253

2279

11 13 16

18 21 24

2304

2330

2355

2380

2405

2430

2455

2480

2504

2529

10 12 15

17 20 22

2553

2577

2601

2625

2648

2672

2695

2718

2742

2765

9 12 14

16 19 21

2788

2810

2833

2856

2878

2900

2923

2945

2967

2989

9 11

13 |

16 18 20

3010

3032

3054

3075130961

3118

3139

3160

3181

3201

8 10 13

15 17 19

3222

3243

3263

3284

3304

3324

3345

3365

3385

3404

14 16 18

3424

3444

3464

3483

3502

3522

3541

3560

3579

3598

10 12

14 15 17

3617

3636

3655

3674

3692

3711

3729

3747 3766

3784

9 11

13 15 17

3802

3820

3838

3856

3874

3892

3909

3927 3945

3962

9 11

12 14 16

3979

3997

401414031

4048

4065

4082

4099

4116

4133

9 10

12 14 15

4150

4166

418314200

4216

4232

4249

4265

4281

4298

8 10

11 13 15

4314

4330

4346 4362

4378

4393

4409

4425

4440

4456

11 13 14

4472

4487

450214518

4533

4548

4564

4579

4594

4609

11 12 14

4624

4639

4654|4669

4683

4698

4713

4728

4742

4757

10 12 13

4771

4786

4800

481414829

4843

485714871

488614900!

10 11 13

4914

4928

4942

495514969

4983

499715011

5024

5038

10 11 12

5051

5065

5079

5092)5105

5119

513215145

5159

5172

9 11 12

5185

5198

5211

522415237

5250

526315276

5289

5302

9 10 12

5315

5328

5340

5353|5366

5378

539115403

5416

5428

9 10 11

544115453

5465

5478

5490

55021551415527

5539

5551

9 10 11

556315575

5587

5599

5611

5623

5635

5647

5658

5670

8 10 11

568215694

5705

5717

5729

5740

5752

5763

5775

5786

8 9 10

579815809

5821

5832

5843

5855

5866

5877

5888

5899

8 9 10

591115922

5933

5944

5955

5966

5977

5988

5999

6010

8 9 10

6021

6031

6042

6053

6064

6075

608516096

6107

6117

1 1

8 9 10

411

6128

6138

6149

6160

6170

6180

619116201

6212

6222

7 8 9

6232

6243

6253

6263

6274

6284

6294)6304

6314

6325

7 8 9

6335

6345

6355

6365

6375

6386

6395 6405

6415

6425

7 8 9

|6435

6444

6454

6464

6474

6484

{6493[650316513

6522

7 8 9

6532

6542

6561

6571

6599

16609

6618

7 8 9

6628

6637

6646

6656

6665

6675

16684

6693

6712

7 7 8

6721

6730

6739

6749

6758

6767

6776

6785

6794

6803

6 7 8

6812

6821

6830

6839

6848

6857

6866

6875

6884

6893

6 7 8

6911

\mn

lUWfi

MB,

ir SEES

\mz\

lEEra

6 7 8

699C

16998

7007

7016

7024

7033

7042

7050

7059

7067

6 7 8

707C

\7M

t7093

7103

7110

7118

7126

7135

7143

7152

8 7 8

716C

71CW

17177

7185

7193

7202

7210

7218

7226

6 7 7

7243

7251

725S

726’;

7275

7284

7292

7300

7308

7316

6 6 7

7324

7335

7340

► 7348

7356

7364

7372

7380

7388

7396

6 6 7

II 0

1 1

1 2

£3.

1 4

1 5

1 7

L_8_

9 1

17 8 9

Appendices

673

Logarithms

No.

I t

'‘Oil 2 ’ 3 | 4 ! 5 | 6 7 8 D i 1 2 3

5S'17404'7412'7419’7427'7435 ( 7443 7451 : 7459.7466'7474 . 1 2 Z
56:7482 7490 7497 7505 7513 7520 7528.7536 7543 7551 ! 1 2 2
57! 7559 7566 7574 75827589 7597 7604 7612 7619 7627.. 1 2 2
58)17634 78427649-7657 7664:76727679:7686:7694 77011! 1 1 2
59;|7709 7716 7723 77317738774577527760 7767j7774l| 112

4 5 6
! 3 4 5

3 4 5

3 4 4

7 8 9

1 5"6~7
15 6 7
i 5 6 7
5 6 7
|567

60! |7782 ! 7789!7796 7803
61 11785378607868 7875
62i'7924 7931 7938 7945
63'!79938000 8007,80141
64; 80628069 8075 8082

65 :'8129 8136 8142 8149
66!|8195'8202 8209 8215
67(8261 8267 8274 8280
68(18325 8331 8338,8344
69H8388839584018407

7810 7818 7825 7832 783917846 ! 1 1 2
788217889 7896790317910, 7917!| 1 1 2
7952 7959 79667973 7980;7987i 1 1 2
80218028 8035:804l!8048 l 8055i! 112
8089 8096;8102!8109.81ie:8122ij 1 1 2

3 4 4

3 3 4

15 6 6
5 6 6
5 6 8
5 5 6
5 5 6

8156 8162 8169~8n6!8I82:8189[ 1 1 2
8222)8228 8235824 1824818254* 1 1 2
8287 8293 8299:8306 83! 2.83 19> 1 1 2
835 1 8357 8363 8370|8376;8382; 1 1 1 2
8414;8420!8426;8432 8439 8445 | 1 1 2

3 3 4

2 3 4

5 5 6
5 5 6
5 5 6
4 5 6
4 5 6

70'!84518457 846384708476 84828488'8494 8500'8506: 1 1 2!
71 1!8513 8519 85258531 8537!8543;8549'8555!8561 8567li 1 12
72 85738579 85858591 :85978603!8609'8615, ! 8621 ’8627!; 112!
731186338639 8645 8651 8657 l 8663;866918675!8681t8686 i 112 1
74!!8692!869887048710!8716!8722j8727:873318739!8745!! 1 1 2 j

2 3 4

4 5 6
4 5 5
4 5 5
4 5 5
4 5 5

75, 8751 ;87568762'8768
76118808881488208825
77I18865 88718876 8882
78! 18921 8927 8932,8938
79! !8976!8982!898?,8993

80 !9O31»036 9042 9047"
81 1 !9085 ! 9090'9096!9101
82! [913819143 9149)9154
83119191 191969201 '9206
84! I9243924892 539258

8774}8779 , 8785|8791 1879718802; j l 1 2
88318837 8842 ! 8848 8854 8859 ! 1 l 2
8887 8893 8899 8904 89108915^ 1 1 2
8943!8949:8954|8980*8965'897i!j 112
8998l9004|9009i9015!9020!9025il 1 I 2

2 3 3

2 3 3 1
2 2 3|

4 5 5
4 5 5
4 4 5
4 4 5
4 4 5

9053 9058190631906919074
9106(9112:9117 912219128
91591918519170 917519180
9212:921719222 9227 9232
92631926919274)92799284

90791

9133;

9186)

92381

92891

112

1 1 2

2 3 3

4 4 5
4 4 5
4 4 5
4 4 5
4 4 5

85' 19294 !9299!9304!9309l9315!9320j932S!9330
86ll9345 l 9350!9355:9360 9365 , 9370*9375'9380
87H9395 , 9400!9405!9410{9415!9420l9425j9430
88|!9445!9450;9455i9460[9465;9469 9474 9479
89) 19494 *9499 9504 |9509|9513)9518l9523l9528

9335

9385

9435

9484

9533

9340

9360

9440

9489

9538

1 1 2

0 1 1
Oil

0 1 1

2 3 3

2 2 3

4 4 5
4 4 5
3 4 4
3 4 4
3 4 4

Hill

iliil

liiii

95621956619571 1957819581
960919614196191962419628
9657)9661 196681967 1)9675
9703)9708 971319717)9722
975Q)9754t9759|9763!9768

9586!

9633!

96801

97271

9773|

0 1 1
Oil
Oil

0 1 1
10 11

2 2 3)
2 2 3

2 2 3

3 4 4
) 3 4 4
3 4 4
3 4 4
| 3 4 4

95[ 19777 19782 9786 9791 19795 98001980519809
96 l !9823!982? 983219836,9841 9845 985019854
97 1 1 9868,9672. 98779881 19886 9890)9894(9899
9819912,99179921 !9926!9930 9934)9939)9943
991)9956 19961 l 9965i9969!9974|9978i9983!9987

9814

9859

9903

9948

9991

9818

9863

9908

9952

9996

0 1 1

2 2 3

2 2 3 1
2 2 3

2 2 3!
2 2 3]

3 4 4
3 4 4
3 4 4
3 4 4
3 3 4

II 0 | 1 | 2 ! .3 | 4 1 5 J 6J„ 7.| 8J 9.1JL 2 3J

LMLiJ

XJrrJ

674

Appendices

APPENDIX XI.

Antilogarithms

Log,

T& [ n 2~T~3 I 4~1 5~! er '7 •; 8' |~9"!r 'l" g~^"i ~i~T~e~r~ , ?~8'~8

.00

.01

.02

.03

.04

1000

1023

1047

1072

1096

1002

1026

1050

1074

1099

1005

1028

1052

1076

1102

1007

1030

1054

1079

1104

1009

1033

1057

1081

1107

1012

1035

1059

1084

1109

1014

1038

1062

1086

1112

1016

1040

1064

1089

1114

1019

1042

1067

1091

1117

1021

1045

1069

1094

1119

0 0 1

0 0 1
Oil

1 1 1

1 1 1
111
111

1 1 2

2 2 2
2 2 2
2 2 2
2 2 2
2 2 2

.05

.06

.07

.08

.09

1122

1148

1175

1202

1230

1125

1151

1178

1205

1233

1127

1153

1180

1208

1236

1130

1156

1183

1211

1239

1132

1159

1186

1213

1242

1135

1161

1189

1216

1245

1138

1164

1191

1219

1247

1140

1167

1194

1222

1250

1143

1169

1197

1225

1253

1146

1172

1199

1227

1256

Oil

112

2 2 2
2 2 2
2 2 2
2 2 3
2 2 3

.10

.11

.12

.13

.14

1259

1288

1318

1349

1380

1262

1291

1321

1352

1384

1265

1294

1324

1355

1387

1268

1297

1327

1358

1390

1271

1300

1330

1361

1393

1274

1303

1334

1?65

1396

1276

1306

1337

1368

1400

1279

1309

1340

1371

1403

1282

1312

1343

1374

1406

1285

1315

1346

1377

1409

Oil

0 1 1
Oil

112
12 2
12 2
12 2
12 2

2 2 3
2 2 3
2 2 3
2 3 3
2 3 3

.15

.16

.17

.18

.19

11413

11445

11479

1514

[1549

1416

1449

1483

1517

1552

1419

1452

1486

1521

1556

1422

1455

1489

1524

1560

1426

1459

1493

1528

1563

1429

1462

1496

1531

1567

1432

1466

1500

1535

1570

1435

1469

1503

1538

1574

1439

H472

1507

1542

1578

1442

1476

1510

1545

1581

0 1 1
Oil
Oil
Oil
Oil

12 2
12 2
12 2
12 2
12 2

2 3 3
2 3 3
2 3 3

2 3 3

3 3 3

.20

.21

.22

.23

.24

1585

1622

1660

1698

1738

1589

1626

1663

1702

1742

1592

1629

1667

1706

1746

1596

1633

1671

1710

1750

1600

1637

1675

1714

1754

1603

1641

1679

1718

1758

1607

1644

1683

1722

1762

1611

1648

1687

1726

1766

1614

1652

1690

1730

1770

1618

1656

1694

1734

1774

Oil

0 1 1

12 2

2 2 2

3 3 3
3 3 3
3 3 3
3 3 4
3 3 4

.25

.26

.27

.28

.29

1782

1824

1866

1910

1954

178611791

182811832

187111875

191411919

1959|1963

11795

1837

1879

1923!

1968

1811

1854

1897

1941

1986

1816

1858

1901

1945

1991

■

2 2 2

2 2 3

2 2 3 1
2 2 3|

3 3 4
3 3 4
3 3 4
3 4 4
3 4 4

.301

.311

.321

.331

. 34 ]

1995

2042

2089

2738

2188

2000120041200912014
2046|2051|2056|2061
20941209912104(2109
2143 21481215312158
219312198 2203 2208

201812023
2065 2070
211312118
21632168
221312218

2028

2075

2123

2173

2223

2032

12080

2128

2178

2228

2037

2084

2133

2183

2234

Oil

112

2 2 3

2 2 3 1
2 2 3

2 2 3!
2 3 3 1

3 4 4
3 4 4
3 4 4

3 4 4

4 4 5

.35

.36

.37

.38

.39

2239

2291

2344

2399

2455

2244

2296

2350

2404

2460

2249

2301

2355

2410

2466

2254

2307

2360

2415

2472

1225912265

23122317

236612371

24212427

247712483

2270

2323

2377

2432

2489

2275

2328

2382

2438

2495

228012286

233312339

238812393

244312449

2500|2506

112

2 3 3

4 4 5
4 4 5
4 4 5
4 4 5
4 5 5

.40

.41

.42

.43

.44

I2512I2518I2523I2529
2570(25761258212588
2630126361264212649
2692I2698I2704I2710
2754|2761 2767 2773

12535

2599

2655

2716

2780

2541

2600

2661

2723

2786

12547

2606

2667

2729

2793

2553(2559

261212618

267312679

273512742

2799|2805

2564

2624

2685

2748

2812

1 1 2
112
112
112
112

2 3 4

3 3 4

4 5 5
4 5 5
4 5 6
4 5 6
4 5 6

.451

.461

. 47 !

.48

.49

2818

2884

2951

13020

113090

28251283112838

28911289712904

29581296512972

30271303413041

3097(310513112

2844

2911

2979

3048

3119

2851

2917

2985

3055

3126

2858

2924

2992

3062

3133

2864

2931

2999

3069

3141

2871(2877

293812944

3006(3013

307613083

3148|3155

112

1 1 2
112
112

1 1 2

3 3 4

3 4 4

5 5 6
5 5 6
5 5 6
5 6 6
5 6 6

1.0 I 1 I 2/| 3(415(61718(9 II 1231456(789

675

Antilogarithms

Log.

" i: o 1 2 13 4 5 6 7 8 9 1 2 3 , 4 5 '6 ]' 7^8 9

.50' 3 1 62 3 1 70 3 1 77 3 184 3 192 31 99 3206 3214 3221 3228' 1 1 2 3 4 4 5 6 7

.51 >'3236 3243 3251 3258 3266 3273 3281 3289 3296'3304'i 122 345 56?

J2>i3311 3319 3327 3334 3342 3350 3357 3365 3373 3381 1! 122 345 567

.53 3388 3396 3404 3412 3420 3428 3436 3443 3451 3459" 1 22 345 667

.54 3467 3475 3483 349! 3499 3508 3516 3524 3532 >3540 | 122 345 667

755 3548 3556 3565 3573' 3581 '3589 3597 3606 3614 3622 ,T TT 2 13 4 5 6 7 7

.56 3631 3639 3648 3656 3664 3673 3681 3690 3698 3707'! 1231345 678

-57i 3915 3924 3733 3741 3750 3758 3767 3776 3784 3793 1 231345 678

.58 3802 38113819 3828 3837 3846 3855 3864 3873 3882" 1 23*445 678

.593890 3899 3908 3917 3926 3936 3945 3954 3963 3972 * 123(455 678

^0> 3981 3990 3999 4009 4018 4027 4036 4046 4055 4064'"' 12314 5~6|678
.61114074 4083 4093 4102 4111 4121 4130 4140 4150 4159" 123|456|789
^2!'41694178’4188 4198 4207 4217 4227 4236 4246 4256' 1 23!456!789
.63' 4266 4276 4285 4295 4305 4315 4325 4335 4345 4355 1 23 1 4 5 6 1 789
,64| [4365 4375*4385 4395 4406 4416 4426 4436 4446 4457 • 1 23)4561789

.65114467 4477 4487 4498 4508 4519 4529 4539 4550 4560'” 1 23|456>789
J&6II4571 4581 4592 4803 4613 4624 4634 4645 4656 4667 1 1 23 1 4 5 6 1 79 10

.67 > '4677 4688 4699 47 1 0 4721 4732 4742 4753 4764 4775!! 1231457 89 10

.66"4786'4797 4808 4819 4831 4842 4853 4864 4875 4887' 1 23(467 89 10

•69|!4898'4909 4920 4832 4943 4955 4966 4977 4989 5000' 1 1 231567 89 10

.70M50125023 5635'5047 5058 5070 50825093 51055117 ! 124>567I89 11
.71l>5129'5140'5152'5164 5176 5188>5200 5212 5224 5236" 1 2 415 6 7 8 10 11

.72[|5248I5260 1 5272 5284 5297 5309 5321 5333 5346 5358 > 1 2 4 I 5 6 7 9 10 11

.731I5370 , 5383;5395 , 5408 , 5420 543315445 5458’5470 5483" 1 3 4 15 6 8 9 10 11

.74I!5495'5508!552!'5534I5546 5559’5572 5585 : 5598 5610 : 1 3 4 | 5 6 8 9 10 12

.75 >5623 5636'5649'5662 5675 5689 5702'57i'5 5728 5741> 1 3 4 I 5 7 8 ! 9 10 12

.76>>5754 5768'5781'5794 5808 582r5834'5848 5851 5875" 1 3 4 15 7 8 9 11 12

.77II5888 I 5902>5916'5929 5943'5957!5970>5984'5998!6012!I 1 3 4 5 7 8 10 11 12

.78!!6026'6039'e053 6067 6081 >609516109 6124 6138!6152>1 1 3 4 6 7 8 10 11 13

.79C6166 618016194 6209 6223 6237>6252 6266!6281 >6295>| 1 3 4)6 7 9 10 11 13

.80 , l6310 l e324 6339%353W68 6383 : 6397 6412!6427;6442l! 1 3 4 6 7 9 10 12 13

.81!I6457 6471'6486 l 6501'6516i6531>6546 l 6561!6577'6592!! 2 3 5 6 8 9 11 12 14

.82l'6607!6622 6637'6653>6668>6683!6699 , 6714!6730!6745i! 2 3 5 6 8 9 11 12 14

.83> , 6761’6776‘6792'6808'6823(6839!6855!68?ll6887!6902>! 2 3 5 6 8 9 11 13 14

J84l!6918 , 6934>8950'6966!6982>6998!7015!7031!?047!7063>! 2 3 5 6 8 10 11 13 15

.85H7070 ! 709ei7H2i?129’7145'7161!7178 7194I7211!7228!l 2 3 5 7 8 10 12 13 15

.861(7244 >72611727817295 7311I7328I734517362I7379I7396I1 2 3 5 7 8 10 12 13 15

.87l!7413!?430!7447>7464!7482!7499175iei7534l7551!7588!! 2 3 5 7 8 10 12 14 16

J8>!7586>7603 I 7821>7838'7656 , 7674!7691I7709I7727I7745I| 2 4 5 7 9 11 12 14 18

.8?>!7762'7780 7798 7816 7834>7852 , 7870!7889!7907!7925[| 2 4 5 7 9 11 13 14 16

!901 >7943 7962 7980 7998 8017* 8035>S054 807218091 !8 1 1 0 f 2 4 6 7 9 11 | 13 15 17

.91 1'8128’8147>8166 8185 8204>8222>8241 >8260(827918299 2 4 6 8 9 11 I 13 15 17

.92: 8318 : 8337 8356'8375 8395 841418433 84531847218492! 2 4 6 8 10 12 | 14 15 17

.93118511(8531 '8551 85?0'8590!8610>8630'8650I8670!8690! 2 4 6 8 10 12 14 16 18

.94|!8710>87308750:87708790'8810i8831!8851'8872l8892| 2 4 6 8 10 12 ( 14 16 18

.95»891318933'8954'8974 T 8995 9bl6'9036>9057'9078:9099 2 4 6 8 10 12 15 17 19

.96(1912019141 | 9161>9183 < 9204'9226>9247I9268!9290I9311 2 4 6 8 11 13 15 17 19

.97ll9333>9354f9376 9397!9419>9441!9482(9484l9506l9528 2 4 7 9 11 13 15 17 20

.9SII9550I9572!9594 I 9616(9538!9661I9683>9705(97279750 2 4 7 9 11 13 16 18 20

.99!!9772!9795!9817>6840'9863:9886!9908!9931|0944{9977 2 5 7 9 11 14 16 18 20

J J j 1. | 2 | 3 | 4 1, 5..|,6. 1.7, 181 9 II 1 2314561789

Name Index

Acheson, E. G„ 408
Anderson, C, D., 831
Arrhenius, S., HM, 253
Avogadxo, A., 58, 59

Becquerel, H., 139, 53?
Berg, O., 572
Bohr, N„ 150, 166, 632
Bom, M„ 182
Boyle, R„ 16, 25
Brandt, G., 390
Briekwedde, F, G., 20-3
Bronsted, 263
Brown, H., 483
Burnen, R. W., 506

Cannizzaro, S., 62
Castner, H., 597
Cavendish, R, 613
Celsius, A., 10
Chadwick, )„ 142, 633
Charles, J., 18, 25
Collct-Disootils, R, 540
Coster, G., 539
Cottrell, F», 486
Coulomb, C, 180
Crookes, W., 136, 602
Curie, Marie, 140, 517, 634
Curie, Pierre, 140

Dalton, J., 4, 22, 26, 51
Davy. R, 506. 517, 597
de Boisbaudran, F„ 602
de Bzo^ie, L. f 158
Debterae, A., 517
Debye, P., 257
del Rio, A., 540
Democritus, 1

Dirac, P,» 632
Dorn, F., 614
Dulong, P„ 82
Dumas, j., 190
Du Nuoy, L., 40

Edison, T., 327
Einstein, A., 630
Ekeberg, A., 541

Fahrenheit, G., 10
Faraday, M., 328, 483
Fermi, E., 638, 643
Frasch, H,, 362
Frisch, O., 638

Gay-Lussac, L., 18, 55, 74, 595
Geiger, R, 142, 629
Glaser, D»> 628
Goeppert-Mayer, M., 632
Goldstein, E., 144
Goodyear, C,, 453
Graham, T., 23, 26, 477
Gregor, W., 539

Haber, F„ 182, 379
Hahn, O m 638
Hall, G, 597
Hatchett, C„ 541
Heisenberg, W., 152
Henry, W., 231
Heroult, P., 597
Herschel, W., 537
Hess, G., 88
Hevesy, G., 539
HiUebrand, W., 614
Hjelm, P., 586

678

Name Index

Hofmann, A., 195
Huckel, E., 257

Janssen, P., 614
Joliot, F., 634
Joliot-Curie, I., 634
Joule, J., 81

Kekute, R, 434

Kelvin, Lord (W. Thomson), 19

Kirchoff, G., 506

Klaproth, H., 537, 539

Klaus, K., 558

Kohlrausch, F., 256

Lavoisier, A., 200
Lawrence, E., 635
LeChatelier, H., 121, 260
Lewis, G., 171
Lockyer, J., 614
Lowry, J., 263

Marsden, E., 142
Mayer, R.> 81
Maxwell, J., 4, 138
Meitner, L., 638
Mendelejeff, D., 128, 613
Meyer, L., 128
Millikan, R., 137
Moseley, H., 144
Murphy, G., 203

Nemst, W., 318
Newlands, J., 128
Newton, I., 3, 122
Noddack, W., 572

Oersted, H., 597

Planck, 138, 152
Paracelsus, 585
Pauli, W., 154
Pauling, L., 176
Peligot, E., 537

Perey, M., 506
Perrier, C., 572
Petit, A., 62
Priestley, J., 193

Ramsay, W„ 613, 614
Raoult, F., 235
Reich, F., 602
Rayleigh, Lord, 613
Richter, T., 602
Roentgen, W., 137
Rutherford, E., 142, 633

Sainte-Claire Deville, E., 397
Scheeie, K., 193, 566, 567
Schrodinger, E., 153
Sefstrom, N., 540
Segre, E., 572
Solvav, E., 511
Strassman, F., 638
Strohmeyer, F., 588
Strutt, J., 613
Szilard, L., 643

Tacke, I., 572
Tennant, R., 558, 559
Thenard, L., 595
Thomson, J., 137, 142
Torricelli, E„ II
Travers, M., 614
Tyndall, J., 482

Urey, H., 203

Van der Waals, J., 27, 38, 189
van t Hoff, J., 120

Weitzmann, C„ 445
Werner, A., 501
Wilson, C., 627
Winkler, G., 604
Winthrop, J,, 541
Wohler, R, 418, 597
Wollaston, W., 559

Subject Index

The letter Y following a page number refers to a table on that page; the letter V refers
to a footnote.

absolute temperature, 19
absolute zero, 19
abundance of the elements, 7
acetaldehyde, 444

acetic acid, 445; equilibrium, 261; glacial,

440

acetone, 444

acetylene, 431

Acheson process, 408

acid, defined, 203, 268; organic, 444;

strength, 353, 373
acidimetry, 272

acids and bases, 260-274; Bronsted concept,
203; Lewis concept, 200
actinide series, 128, 529, 537; elements,
532

activated charcoal, 409
activated complex, 98
activation, 90
activity, in solution, 258
activity, order of, 203, 314
addition compound, 175
addition polymerization, 454
addition reactions, 429, 431, 446
to&enaUxt, 404
adsorption, 404, 485
aerosols, 478, 479t
alabaster, 5?'
albumin, 483
alcohols, 441
aldehydes, 444
akkee, 457

aliphatic, defined, 421; compounds, 421-432
alkadienes, 431
alkali metals, 503-512

- «■■■ -t.« - awH

Sa w ia i i i iwt r y,

alkaline earth metals, 514-527

alkane hydrocarbons, 421-427; table, 421;

reactions, 424; structure, 42$
atkene hydrocarbons, 427-431; table, 428;

reactions, 429; structure, 429
alkyl radical, 423

aikyne hydrocarbons, 431-432; table, 431
allotropic forms, 86, 361, 392, 406, 467, 607
alloy steels, 549t,

alloys, defined, 496; of copper, 578; of
iron, 549
a Ini co, 555

alpha particles, 140, 144t, 621 1
alumina, 597
aluininothermy, 600
aluminum, 593, 594t, 597
alums, 602
alundum, 601

amalgam, defined, 496; 508

amalgamation, 580, 583, 589

americium,

amides, 448

amines, 448

aminoacids, 449; in proteins, 463, 464
ammonia, 378; ionization of, 264, 380
liquid, 381, 509; oxidation process, 381
ammonifying bacteria, 388
ammonium hydroxide, 380
ammonolysis, 380
amorphous carbon, 408
amorphous silicon, 468
amorphous solids, 38
ampere, 328
amphiprotie, 285, 449
amphoteric hydroxide, 293
analysis, defined, 649; cation, 653; anion,
656

680

Subject Index

angiesite, 609
Angstrom unit, 33
anhydrite, 517, 521
aniline, 448

anion, 310, 323; analysis, 656
anion exchange resins, 524
anisotropic crystals, 34
annealing of glass, 473
anode, 137
anthracene, 434, 435t
anthracite, 408
antifreeze solution, 243
antiknock gasoline, 344, 438
antimony, 376, 399-402
apatite, 391
apoenzyme, 465
aquamarine, 515, 597
aqua regia, 204, 559, 591
aragonite, 517
arc process, 383
argentite, 580
argon, 612
argyrodite, 604

aromatic, defined, 421; compounds, 432-
436; reactions, 434

Arrhenius theory, assumptions of, 253;
limitations, 253

arsenic, 376, 399-402; Marsh’s test for, 400
arsenopyrite, 399
arsine, 400

artificial radioactivity, 663

asbestos, 467t, 473, 516

aspirin, 448

astatine, 334, 335t

asymmetric carbon atom, 451

atmospheric pressure, 11

atomic bomb (see nuclear bomb)

atomic heat, 62

atomic mass unit (amu), 57

atomic hypothesis, 4

atomic-molecular hypothesis, 51, 55-64

atomic number, 144

atomic radius, 180

atomic structure, 135-168

atomic theory, Dalton’s, 51

atomic transmutation, 620, 633

atomic volume, 126

atomic weight, defined, 57; 147; determina-
tion of, 61; scales, 148
average velocity of gas molecules, 25
Avogadro’s hypothesis, 58
Avogadro number, 59
azeotrope, 2 44
azurite, 576

bacteria, and nitrogen fixation, 388
bagasse, 458
Bakelite, 445, 455
baking powder, 512

baking soda, 506

balancing equations, 67; oxidation-reduc-
tion, 297
barite, 517
barium, 514, 516t
barometer, 11
base, defined, 263, 266
basic oxygen converter process, 550
battery, defined, 306; dry cell, 320; lead-
storage cell, 327; Edison cell, 327
bauxite, 597
Bayer process, 599
benzaldehyde, 444

benzene, 432, 435t; derivatives of, 434;

structure of, 432
benzoic acid, 445
benzyl alcohol, 444
beryl, 515, 517, 519
beryllium, 514, 516t
Bessemer process, 548, 577
beta particles , 140, 144t, 621t
Betts process, 609
himolecular reactions, 99
binary mixture, distillation, 243
binding energy, 629
biochemistry, 456
bismuth, 376, 399-402
bituminous coal, 408
black lead, 408
black phosphorus, 393
blast furnace, 545
bleaching action, 343
bleaching powder, 354
blister copper, 577
blueprints, 555
blue vitriol, 579

Bohr theory, 150; derivation for hydrogen
atom, 166
boiler scale, 522

boiling point, 43; of solutions, 238
boiling-point method, molecular weights,
240

bond, coordinate covalent, 173; covalent,
172, 186; single, 172; double, 173; triple,
173; ionic, 170; distance, 187; properties,
188

bond energy, 96
boranes, 596
borax, 523, 595
boric acid, 596
Bom-Haber cycle, 182
borohydrides, 507
boron, 593, 594t, 595
borosiiicate glass, 474
Boyles Law, 16, 25
brass, 578

Bredig arc method, 480
breeder reactor, 644
British thermal unit, 85

Subject Index

681

bromine, 344; oxygen acids of, 353

Brens ted concept of acids and bases, 263

bronze, 578; Age, 495, 576

Brownian movement, 483

brucite, 515

B.T.U,, 85

bubble chamber, 628

buffer solution, 282

Buna-S, 455

butane, 426, 42 It

cadmium, 586t, 588
calcite, 517
calcium, 514, 516t
calcium carbide, 432

calcium carbonate, 521; equilibrium, 118

calcium hydrogen carbonate, 522, 523

calcium hydrogen sulfite, 461

calcium oxide, 519, 526

calcium sulfate, 521

calculations, stoichiometric, 68-74

calomel, 590

calorie, 85

calorimeter, 85

Cannizzaro method, 62

carat, 406

carbides, 415, 568. 595
carbocyclic compounds, 432
carbohydrates, 450; digestion of, 465
carbon, 404415; allotxopes, 406; asym-
metric, 451; dating, 647
carbon compounds, inorganic, 409-415; or-
ganic, 418-465
carbonic acid, 411
carbon-fourteen, 631, 647
carbon tetrachloride, 173, 178
carbon-twelve standard, 57, 148
carbonyl group, 444
carboxyl group, 445
camotite, 537, 540
casein, 463
cassiterite, 606
cast iron, 548
catalysis, 94, 369, 534
catalyst, 95; organic, 465
cathode, 136, 310, 323
cathode rays, 137
cation, 310, 323
cation exchange resins, 523
caustic soda (see sodium hydroxide)
Celanese, 461
celestite, 517

cell, primary (see voltaic ceil)
cellophane, 461
ce&uJoid, 461

cellulose, 456, 459, 460; acetate, 461; ni-
trate, 461
CtobteoL 458

Celsius thermometer scale, 10
cement, 525
cementite, 551

centigrade temperature scale, 10

centrifugal separation, 641

ceramics, 474

ccrussite, 608

cesium, 503, 504t

chain reaction, 639

chaleopyrite, 576

chalcolite, 576

changes of state, 45

charcoal, animal, 409; wood, 409

charge of electron, 137, 621t

Charles' Law, 18, 25

chemical atomic weight scale, 57, 148

chemical bond, 169-191

chemical change, defined, 6; 49-54

chemical equilibrium, 107-122

chemical equivalent, 77; in titration, 272

chemical reaction, rate of, 93-104

chemical warfare, 343

chemiluminescence, 392

Chile saltpeter, 506, 507

China clay, 474

chloric acid, 355

chloride of lime, 354

chlorine, 334, 335t, 339; in chemical war-
fare, 343; Deacon process, 324, 340;
electrolysis of sodium chloride, 340
chloroform (trichloromethane), 178
chlorophyll, 459
chloroprene, 454
chlorous acid, 354
chromates, 565
chrome green, 563
chromite, 562
chromium, 562, 564t
cinnabar, 360,589
cis-compounds, 428
citric acid, 447
Claude process, 379
elaudetite, 399
clay and clay products, 474
cleaning solution, 566
clinker, 526
close-packing, 490
cloud chamber, 627
coagulation treatment, 600
coal, 408
coal tar, 43$
cobalt, 544t, 555
cobalt-sixty, 647
coenzyme, 465
coinage metals, 575, 576
coke, 408

collisions, molecular, 95
collodion, 481

Subject Index

color, 534: in flame tests, 655

combining volumes of gases, 55, 74

combining weight, 55, 77

combustion, 197

common ion effect, 277

complex ions, 291, 497, 534

composition by weight, 50, 68

compound, defined, 6

concentration cell, 320

concentration factor, in equilibrium, 110, 122

concentration units, ,220

concrete, 525

condensate, 243

condensation polymerization, 454
conductance, 249; equivalent, 250; molar,
250; specific, 251
conformations, 427
conjugate acid-base pair, 264
conjugated double bond, 431, 453
consecutive reaction, 100
conservation of energy, 53n, 81
conservation of mass, 50, 53, 630
constant-boiling mixtures, 244
contact process, 369
coordinate covalence, 173
coordination number, 292, 490, 497
copper, 574-580
copper matte, 577
copperas, 553
com syrup, 458, 460
corrosion, 552
corundum, 597
Cottrell precipiator, 486
coulomb, defined, 328
Coulomb’s Law, 180, 210
couples, electric, 552
covalent bond, 172, 186, 188
cracking, 425, 437
cresols, 444
critical pressure, 29
critical size, 640
critical temperature* 29
crocoisite, 562
Crookes tube, 136, 141
cross-linking, 454
crucible steel, 551
cryogenics, 492, 617
cryolite, 597, 598
cryotron, 542, 617
crystal binding, 38
crystal forms, 34

crystal lattice, 36; of metals, 490
crystal structure, principles, 36
cupric compounds (see copper)
cuprite, 578

cuprous compounds (see copper)
cyanamide process, 382
cyimide process, 580, 583
cyclic hydrocarbons* 432

cyclo-compounds, 432, 438
cyclohexane, 435
cyclotron, 635

Dacron, 456
Dalton’s Law, 22, 26
Daniell cell, 310
DDT, 343

Deacon process, 108, 324, 340
definite proportions, 50, 53
deliquescence, 237
deltas (river), 486

derivatives of hydrocarbons, 441-451, 450t
detection of radioactivity, 628
detergents, 462, 486
deuterium, 148, 203, 642, 645
deuteron, 203, 621t
deviations from the gas laws, 26
dew point, 214
dextrin, 457, 460
dextrorotatory, 450
diabetes, 457, 465
dialysis, 481
diamagnetic, 162
diamond, 406, 408, 419
diaphragm cell, 340
diatomaceous eaarth, 467
dichromate ion, 316t, 565
dielectric constant, 180, 210
diene hydrocarbons, 431
diffusion, 23, 27, 47; in separation of iso-
topes, 641

dinitrogen pentoxide, 385

dinitrogen trioxide, 385

diolefins, 431

dipole, 177; moment, 178

diprotic acids, 267

disaccharides, 457, 458

discharge potential, 323

displacement, 203, 350

displacement rule in radioactivity, 620

disproportionate, 569, 570

dissociation, 257

distillation, 7; fractional, 243; vacuum, 43
distribution coefficient, 232
distribution of molecular energies, 25, 41,
98

divinylbenzene, 523
dolomite, 515

doner atoms, 263, 266; in complex ions*
500t

double bond, 173, 427; in benzene, 434

double decomposition, 350

Dow process, 518

Downs process, 507

dry cell, 320

dry ice, 410

orbital, 497, 498. 499t

Suited IfH&x

683

orbital, 497, 499, 563
d+sjr* orbital, 499t
Dulong and Petit rule, 62
duralumin, 599
duriron, 468

dust, 627; explosions, 94

earth, age of, 647

earth*?; crust, composition of, 7

ebullition, 43

Edison storage cell, 327

effective atomic number (EAN), 499

effective nuclear charge, 180

efficiency, theimodynamic, 84

efflorescence, 212

ekasiliccn, 132

electrical energy, 82, 306; units, 328
electric ample, 552

electric furnace, for silicon, 468; for steel,
551

electrochemical cell, 306-332; voltaic cell,
308; electrolytic cell, 323
electrode, defined, 327
electrode potential, measurement of, 312;

standard, 313, 318t
electrodialysis, 482

electrolysis, 250, 306, 323; of sodium der-
ide, 324

electrolytes, conductance of, 249; solutions
of, 249-258

electrolytic cell, 323-332; and demineraliza-
tion, 525

electrolytic refining of metals, 326
electromagnetic radiation, 138, 449
electromagnetic separation of isotopes, 146,
641

electrometallurgy, 494
electromotive force, defined, 310; measure-
ment, 311 and equilibrium constant, 314
electromotive series, 314
electron, 135, 144t
electron cloud, 153'

electron configurations, summary, 156; de-
velopment of, 150; and atomic proper-
ties, 160

electronegativity, 178
electronic conductance, 306
electrophoresis, 485
electrorefining, 495, 577
electroscope, 626
electrotyping, 578
electrovalence, 172, 178
element, defined, 8
elements, abundance of, 7
emerald, 515, 597
emery, 597
emulsifying agent, 486
emulsion, 478, 479t, 488

enantiomers, 45 h
endothermic reaction, 81
end point, 274

energy, conservation of, 81; distribution,
25, 41, 98; interconvertibility, 81; in-
ternal, 87; of activation, 96
energy changes, in chemical reactions, 81;

measurement, 85
energy level, 151
energy units, 85
enthalpy, 87
entropy, 83
enzymes, 465
Epsom salts, 521
equations, chemical, 87
equilibrium, defined, 42, 44
equilibrium, chemical, 107-122; displace-
ment of, 109; heterogeneous, 118; homo-
geneous, 118; ionic , 260-294; Le Chate-
liers law; 121; point of, 108
equilibrium constant, 111; and EMF, 317;

weak electrolytes, 260; solubility, 284
equilibrium equation, derivation of. 111
equilibrium, factors, 110, 119, 120
equivalence, 77

equivalent weight, 55, 77; in oxidation-
reduction, 302
esters, 447
etching of glass, 351
ethane, 420, 42 It
ethanol, 427, 441
ethers, 443

ethyl alcohol (see ethanol)
ethylene, 427; structure of, 429
ethylene glycol, 442
ethyne (see acetylene)
eudiometer tube, 15
eutectic mixture, 242
evaporation, 41
exchange reaction, 642
exothermic reaction, defined, 81
extraction, 322; of silver, 580

Fahrenheit scale, 10

Faraday's laws, 328

faraday, defined, 329

fats, 461; hydrolysis of, 482

fatty acids, 447, 482

Fehlings solution, 445, 458, 459

feldspar, 7, 467t, 473, 506, 597

fermentation, 458

ferric compounds (see iron)

ferrochrome, 562

ferromagnetism, 545

ferrosilicon, 468

ferrous compounds (see iron)

fertilizers, 506

fibroin, 463

684

Subject index

filtration of colloids, 481
fire extinguishers, 412

first order reactions, 102; rate equation
(derivation), 105

fission, nuclear, 638; factor, k, 640

flame, defined, 197

flame tests, 505, 515, 655

flocculating power, 486

flotation process, 493, 576

flowers of sulfur, 363

fluorine, 335, 337; electrolytic cell, 338

fluorspar, 517

flux, 494, 545

foods, 456

foraminifera, 516

formaldehyde, 444

formal solution, 224

formic acid, 445

formula, defined, 66; determination of, 69;

graphic, 173
formula weight, 67

fractional distillation, 243; of petroleum,
437

fractional electrolysis, 642
francium, 503
francium, 503, 504t
Frasch process, 362
free energy, 83, 114
free radicals, 96

freezing point, defined, 37, 44; changes
in, 240, 252; eutectic, , 242
freezing point method, for molecular
weights, 240
freon, 339

Friedel-Crafts reaction, 596, 601

fructose, 456, 457

fuel, 437t

fuel system, 197

functional group, defined, 440; 450t
furnace, basic oxygen, 550; blast, 545; elec-
tric, 551; open hearth, 550
fusion, heat of, 37, 46; alkali, 541; nuclear,
638

galena, 360, 608
gallium, 602

galvanic cell (see voltaic cell)
galvanized iron, 588
gamma rays, 140, 640
gangue, 492
garnet, 467t, 597

gases, 13*29; determination of molecular
weight, 57; laboratory preparation, 14;
laws of, 16; liquefaction of, 29; measure*
ment of, 14; mixed, 22; solubility of, 231;
volume-weight relations, 60
gaseous products, 290; in thermochemistry,
87

gasoline, 344, 437

Gay Lussac's Law of Combining Volumes
55, 74

Geiger-Muller counter, 629
gels, 478

germanium, 405, 604

German silver, 578

glass, 473; decolorization, 373, 570

glucose, 456, 457

glycerol, 442, 461

glycine, 449, 464

glycogen, 457, 460

gold, 580

Goldschmidt process, 494
Graham's Law, 23, 26
gram-atom, 59
gram-atomic weight, 59
gram-equivalent weight, defined, 77
gram-formula weight, 67
gram-molecular volume, 59
gram-molecular weight, 59
granite, 7

graphic formula, 173

graphite, 406, 419

gravity cell, 310

green vitriol, 553

Group O elements, 612

Group IA elements, 503

Group IB elements, 574

Group IIA elements, 514

Group IIB elements, 585

Group IIIA elements, 534

Group IIIB elements, 593

Group IV A elements, 538

Group IVB elements, 404, 467, 604

Group VA elements, 540

Group VB elements, 375, 390

Group VIA elements, 562

Group VIB elements, 358

Group VIIA elements, 567

Group VIIB elements, 334

Group VIII elements, 543

group displacement rule, 620

guncotton, 461

gunpowder, 512

gypsum, 517, 526

Haber process, 12 ln„ 379
hafnium, 535t, 540
half-cell potential, 312
half-life period, 104, 624
halie acids, 355
Hall process, 597

halogen derivatives of the hydrocarbons, 441
halogens, 334-346; halogen acids, 350; oxy-
gen compounds of, 352-356; organic
derivatives, 440
halous adds, 354

hardness, 406, 489; Mohs scale, 488

Subject Index

685

haul water, 522
health physics, 644

heat, 9, 81; specific heat, 86; units of heat,
85

heat of formation, 90; fusion, 87, -16;

vaporization, 44
heavy water, 203, 642
Heisenbergs Uncertainty Principle, 152
helium, 612, 61 3t, 615; Helium I and II, 616
hematite, 545
hemoglobin, 412, 483
Henry's Law, 231
Hess, law of, 88
heterocyclic compounds, 436
homologous series, 422
Hooker Type S cell, 341
Hoopes process, 599
hormones, 464
humidity, 214

hybrid orbital, 187, 419; in complex ions,
497, 499t
hydrates, 211
hydrazine, 381
hydra zoic .acid, 382
hydrides, 202, 509

hydrocarbons, 418-439; aliphatic, 421; iuo«
matic, 432; cyclic, 432; derivatives, 441-
451, 450t; nomenclature, 423; unsatur-
ated, 427-432
hydroforming process, 436
hydrogem, 200-205; hydrogen bonds, 179,
351; isotopes, 203, 643
hydrogen bromide, 349, 352
hydrogen chloride, 349, 351
hydrogen electrode, 312
hydrogen fluoride, 349, 351
hydrogen iodide, 349, 352
hydrogen peroxide, 215
hydrogen sulfide, 364; equilibrium, 267,
278; in analysis, 653
hydrolysis, 279
hydroninm ion, 209
hydrophilic colloidal dispersions, 479
hydrophobic colloidal dispersions, 479
hydroquinone, 443

hydrosulfurie acid (see hydrogen sulfide)
hydroxy acids, 447
hydroxylamine, 382
droxyl radical, 442
hygroscopic, 212
hypochlorous acid, 342, 354
hypohalous acids, 354
hyponitrous acid, 382
hypophosphorous acid, 399
hypothesis, 3

ice, structure of, 207
Iceland spar, 517

ideal gas law, 21
ideal solution, 243
immiscible liquid, 231
indicators, 270,
indium, 602

inert gases (see noble gases)
initiator (in polymerization), 456
ink, 553

inner d orbitals, 498

instrumental analysis, 649

insulin, 463, 464

interhalogen compounds, 346

internal rearrangement, 418

inversion of ammonia, 379; of sugar, 458

invert sugar, 458

iodine, 335, 345; oxygen acids of, 853
ion-electron method of balancing equations,
297

ion exchange, 523, 537
ionic bond, 170

ionic compounds, dissolving of, 209, 283

ionic crystal, deformation of, 490

ionic conductance, 249; dissociation, 280;

equilibria, 260-294; radii, 182
ionization, 170, 280; chamber, 829; constant,
281; Arrhenius theory, 258; modern
theory, 257; potential, 169, 505
ions, complex, 291, 497, 534
iridium, 558t, 559

iron, 543-555; Iron Age, 495; pyrites, 360,
545

irreversible reaction, 107, 290
isobars, 148
iso-compounds, 424

isomerism, carbon compounds, 422, 428; in
complex ions, 500
isomorphism, 602
isooetane, 438
isoprene, 453

isotopes, 147; of hydrogen, 203; of lead,
609; separation of, 640
isotropic substances, 34

joule, 328

kaolin, 467 1, 474, 597
Kelvin scale, 19
keratin, 463
kerosene, 437t
ketones, 444
ketose, 457

kindling temperature, 197
kinetic equation for gases, derivation, 31
kinetic-molecular theory, 24; of solution,
228

kinetics, chemical, 93; radioactivity, 623*
knocking (of engines), 344, 438
Koroseal, 455

686

Subject Index

kupfemickel, 556
krypton, 612, 613t

lactic acid, 447
lactose, 447
lambda point, 616
lampblack, 409
lanthanide contraction, 536
lanthanide series, 128, 529, 532, 536; ele-
ments, 53 It
lanthanum, 534, 535t
lapis lazuli, 597
latex, 453
lattice energy, 183
laughing gas, 382

law, defined, 3; Boyle's, 16, 25; Charles,
18, 25; combining volumes, 55; con-
servation of energy, 81; conservation of
mass, 50; Coulomb's, 180, 210; Daltons,
22, 26; definite proportions, 50; Dulong
and Petit, 62; Faraday’s, 328; Gay-Lus-
sac’s, 55; Graham’s, 23, 26; Henry’s, 231;
Hess’, 88; Ideal gas, 21; Le Chatelier’s,
121, 230; Moseley’s, 144; multiple pro-
portions, 51; octaves (Newlands’), 128;
partition, 231; radioactive displacement,
620; Raoult’s, 235; Van’t Hoff’s, 120
leaching, 495, 580, 583
lead, 405, 604, 608
lead storage cell, 327
Le Chatelier’s Law, 121, 230, 448
lead chamber process, 369
lead dioxide, 327, 610
levelling effect, 265
levorotatory, 450

Lewis acid, defined, 266; 595, 601

ligand, 174, 292, 497

light quantum, 138, 152

lignin, 461

lignite, 408

lime, 519

limestone (see calcium carbonate), 519

limewater, 520, 521

lipase, 465

liquation, 495, 606

liquefaction of gases, 29, 194

liquid air, 194

liquid drop model of nucleus, 632
liquid state, 33, 39-47
litharge, 610

lithium, 503, 504t, 506, 507
lunar caustic, 581
lung-irritant gases, 343
lyophilic colloidal dispersions, 479
lyophobic colloidal dispersions, 4"

magnesia, 517, 520
magnesium, 514, 518t

magnetic separation, 7; of isotopes, 641

magnetic susceptibility, 162

magnetite, 545

malachite, 576

malic acid, 447

maltose, 459

manganese, 567

manganese dioxide, 570

Marsh’s lest for arsenic, 400

mass defect, 629

mass, defined, 5; conservation of, 50, 53,
630

mass-energy relationship, 630
mass number, 146
mass spectrograph, 146
match industry, 393

mathematical operations (see Appendix I)
matter, states of, 13
mayonnaise, 486

mechanism, 100; of corrosion, 552
meerschaum, 516

melting point, defined, 37, 44 (see freezing
point)

Mendelejeffs Periodic System, 128
mercury, 589-591
mercury cell, 340, 511
mesons, 144t, 829
meta acids, 396
meta compounds (organic), 435
metasilicic acid, 471
metal hydrides, 202, 509
metallic couples, 552
metallic bond, defined, 491
metallic elements and Periodic Table, 131,
492

metallic state, 489-501

metalloids, 492

metallurgy, general, 492

metals, 489-501; order of activity, 203

metaphosphoric acid, 398

methane, 420, 421t

methanol, 441, 444

methylamine, 448

Metric System, 8

mica, 7, 467t, 473, 597

micelle, 487

milk, 486; of lime, 520; of magnesia, 521

milliequivalent, 226

millimole, 228

mineral, defined, 492

mineral matter, 456, 464

mineral springs, 521

Misch metal, 538

mixture, defined, 7, 219; constant boiling,
244

moderator, 643
Mohr’s salt, 553

molal boiling point constant, 238
racial concentration, 226

Subject Index

687

molal freezing point constant, 269
molar conductance, 250
mole fraction, defined, 227
moleeularity (kinetic), 99
molecular motion, 24

molecular weights, 59; of gases, 57; of
solutes, 240
molecule, defined, 60
molar solution, defined, 222
mole, defined, 59
molybdenum, 564t, 566
monazite, 584, 586
Mond process, 556
monoclinic sulfur, 87, 861
monoprotic acids, 267
monosaccharide, 457
mordants, 558, 600
mortar, 520
Moseleys Law, 144
multiple proportions, law of, 51, 53
mustard gas, 348

naphthalene, 434, 435t
naphthenes, 437
neon, 612, 614
neoprene, 454
neptunium, 644
Nernst equation, 318
nerve gases, 344
neutralization, 272
neutrino, 621 1
neutron, 144, 144t, 621
neutron -proton ratio, 631
Newlands' Law of Octaves, 128
nickel, 556; ammonia complex ion, 497;
cyanide complex ion, 498; nickel-silver
alloy, 556

nitrogen, 375-388; 613
niobium, 535t, 541
nitrates, 365, 888
nitric acid, 385
nitric oxide, 333
nitrites, 384

nitro-compounds, organic, 449
nitrogen cycle, 388
nitrogen dioxide, 384

nitrogen family, other elements of, 390-402
nitrogen fixing bacteria, 388
nitrogen tetroxide, 384
nitrogen trioxide, 383
nitrous acid, 383
nitrous oxide, 382
noble gases, 170, 612-618
nomenclature, 336; of organic compounds,
423, 428, 431, 435, 436; of complex ions,
501

nonmetallic elements, 492
nonpolar molecule, 177
normal eomoound for came). 424

normal solution, defined, 225
nuclear atom, 142
nuclear bomb, 642, 646
nuclear chemistry, 620-648
nuclear fission, 638
nuclear fusion, 615
nuclear notation, 633
nuclear reactor, 642
nucleon, defined, 145

nucleus, bombardment of, 633; model of,
632

nuclide, defined, 620
nutrition, 456
nylon, 455

octane number, 438
octaves, law of (Newlands'), 128
octet rule, 169
ohm, 250, 328

oil, see petroleum; viscosity of, 41
olefin series, 427
oleomargarine, 462
olivine, 515

open hearth process, 550
opposing reactions, 100
optical activity, 449; of quartz, 469
orbital, defined, 154; energy levels, 157;

hybrid, 187, 419, 499t
order; and disorder, 84, 617; of a reaction,
100; of activity of metals, 203
ore, defined, 492; concentration of, 492
organic chemistry, 418-465
orlon, 455
orpiment, 399
ortho acid, 396
orthoclase, 597

ortho compounds (organic), 435
orthophosphorie acid, 396
osmium, 558
osmosis, 241

Ostwakl process, 381, 383
outer d orbitals, 498
oxalic acid, 446t, 447
oxidation, defined, 172, 296-303
oxidation number (state), 175
oxidation number method of balancing
equations, 301

oxidation potentials, 314, 505; Table, 316

oxidation-reduction, 306-323

oxidizing agent, 292

oxyacetylene torch, 198, 432

oxygen, 193-198; isotopes

oxygen converter process for steel, 550

oxyacetylene torch, 198, 432

ozone, 198

paints, 521, 539
palladium, 558t, 559

688

Subject Index

paper, 461

para compounds (organic), 435

paraffin hydrocarbons, 421

parallel reactions, 100

paramagnetism, 161, 534

Parkes process, 580, 609

partial ionic equations, 300

partial ionic character, 176

partial pressures, law of (Dalton’s), 22

partition, law of, 231

passivity, 552

Pauli Exclusion Principle, 154
peat, 408
pepsin, 352, 465
peptide, 463
peptization, 480
percentage composition, 68
perchloric acid, 353, 355
perhalic acids, 355

Periodic Law, 125-133; and electron con-
figuration, 162
permanganate ion, 316t, 570
peroxides (see hydrogen peroxide), 509
petroleum, 437
pH, 269

phase, 38; changes of, 38, 45
phase diagram for water, 212
phenols, 443
phosgene, 414
phosphate rock, 517
phosphates, 396
phosphine, 393
phosphorous acid, 395, 399
phosphorus, 375, 390-399
photoelectric cell, 373; effect, 505
photoengraving, 578
photographic flash powder, 519
photography, 581, 626
photosynthesis, 459, 647
physical atomic weight scale, 148
physical change, 6
pi (t) bond, 429, 432, 434
pig iron, 547
pigment, 521, 539
pinch effect, 646
pitchblende, 517, 537
pK, 270

placer mining, 583
Planck’s constant, 139, 152
plasma, 645
plaster of Paris, 521
plastics, defined, 456
plastic sulfur, 361
platinum, 558t, 559
platinum metals, 557, 558t
plumbago, 408
plumbic compounds (see lead)
plumbous compounds (see lead)
plutonium, 644

pOH, 270

polar molecules, 177

polarization, of cell, 321

polarized light, 449

poling, 577

polonium, 359, 372

polybasic acids, 267

polymer, defined, 454

polymerization, 430, 453, 454, 463

polypeptide, 463

polyprotic acids, 267

polysaccharides, 457, 459

polystyrene, 455

polysulfides, 366

Portland cement, 525

positive rays, 141

positron, 62 It, 631

potash, 506

potassium, 503, 504t, 506-509
potassium chlorate, 195, 355
potassium permanganate, 570; in analysis,
571

potentiometer, 312

pressure, defined, 11; of mixed gases, 22;
partial, 22

primary carbon atom, 424
primary cell, 306-321
producer gas, 413

projectiles, for bombarding nuclei, 633, 645

promotion, of electron, 175

propane, 42 It

propanol, 441

properties of matter, 5

propylene, 454

proteins, 462

protolysis, 263, 272

proton, 142, 144t, 621

Prussian blue, 554

pyrex glass, 474

pyridine, 436

pyrites, 360, 545

pyrolysis, 425

pyrophosphoric acid, 398

pyrosulfuric acid, 369

pyroxylin, 461

quantum mechanics, 152
quantum numbers, defined, 151, 154
quartation, 583

quartz, 7, 469; crystal lattice, 470

radioactive displacement rule, 820
radioactive disintegration series, 822
radioactivity, 139, 620; artificial, 663; de-
tection of, 626; kinetics of, 623
radioisotopes, 646
radium, 514, 518t

Subject Index

689

radon, 612, 614
Haoult’s Law, 235

rare earth elements {see lanthanide ele-
ments), 128

rate of chemical reaction, factors, 93
rayon, 461

rays, from radioactive materials, 140
reaction rate theory, 95; rate equations, U>1
reactions, kinetic classification of, 99
realgar, 399

rearrangement, internal, 418
red lead, 610
red phosphorus, 392
reducing agent, 299

reduction, 296-303; defined, 172; in metal-
lurgy, 494

refining, 826, 495; of copper, 577
refrigerating machines, 194, 381
relative humidity, 214
residual forces, 48*1

resistance, electrical, 328; of semiconductors,
468; of metals, 491
resonance, 183

reverberatory furnace, 548, 550

reversible reaction, 187

rhenium, 588t, 572

rhodium, 558t, 559

rhombic sulfur, 87, 361

roasting, 493

rock salt, 507

Roentgen rays (see X-rays)

rubber, 453

rubidium, 503, 504 1

ruby, 597, 601

rust, 49-51, 552

rust-proofing, 552

ruthenium, 558

rutile, 539

sal ammoniac, 379

salt bridge, 307

saltpeter, 506, 507

sapphire, 597, 601

saponification, 482

saturated compound, 173, 429

saturated solution, 228, 285

scandium, 160, 534, 535t

scheme of .nalysis, 652

science, d ined, 1; philosophy of, 1-5

scientific method, 2; scientific explanation, 3

scintillation counter, 626

secondary carbon atom, 424

secondary cell, 306, 823-332

second order reaction, 101

selenium, 359, 372

semi-conductors, 468

semipcrmeable membrane, 241, 481

separation of isotopes, 640

shell model of nucleus, 832
sherardized steel, 588
sigma (or) bond, 429, 432, 434
significant figures (see Appendix I)
silanes, 188

silica (see silicon dioxide), 467, 473

silica gel, 471

silicates, defined, 487; 471

silicic acids, 471

silieides, 408

silicon, 487-475

silicon carbide, 415, 488

silicon dioxide, 469

silicones, 474

silieothermie process, 518

silver, 580

single lx>nd, 172, 422

singh* electrode potentials, 312; Table, 318
single electrode potentials, 812, 316t
size ranges, 478t
slag, 494, 548
slaked lime, 520

smelting, 494; of copper, 577; of iron, 545
soap, defined, 482; cleansing action, 488;

and hard water, 522
soapstone (see tale)
soda ash (see sodium carbonate), 506
soda-lime glass, 474
soda-lime treatment of hard water, 523
sodium, 503-512
sodium carbonate, 506, 511
sodium chloride, 510; electrolysis of, 324;

structure, 173, 189
sodium hydrogen carbonate, 506, 51!
sodium hydroxide, 506, 510
sodium phosphates, 397; and hard water,
523

sol, 478, 479t

solids, classification of, 34

solid state, 33-39

solubility, 210, 229, 651; of mixed gases, 22
solubility product, 284-294
solute, defined, 220

solutions, 219-258; changes in boiling and
freezing point, 238; conductivity of, 249;
defined, 219; of electrolytes, 249-258; of
gases in liquids, 231; kinetic-molecular
view, 228; of liquids in liquids, 220,
222; osmotic pressure, 241; properties
of, 235-246; rate of, 229; types of, 220;
vapor pressure, 235; vs. mixture, 219
SoJvay process, 511
solvent, defined, 220; extraction, 537
sp orbital, 431, 497, 499t
$px orbital, 419, 429
sp* orbital, 419, 429, 497, 499t
$p%d orbital, 498
sp^d- orbital, 499
specific heat, defined, 86

Subject Index

spectra, origin of, 163; of alkali metals, 505

spectroscopy, 138, 163

spelter, 587

sphalerite, 586, 604

spiegeleisen, 569

spontaneous reaction, 83; and equilibrium,
113; and electromotive force, 317
stainless steel, 564
stalactites and stalagmites, 522
standard conditions (STP), defined, 21
standard solution, 223
stannic compounds (see tin)
stannous compounds (see tin)
starch, 456, 457, 459, 460
steel, defined, 548; manuafcture, 548-551
stellite, 555
stibnite, 399

stoichiometric calculations, 70

storage cell, lead, 327; Edison, 327

strength of acids, 263, 353, 373

strontium, 514, 516t

structural formula, 173

styrene, 455, 523

subatomic particles, 621t

sublimation, 44

substitution reaction, 425

substrate, 465

sucrose, 456, 458

sugars, 456

sulfate?, 370

sulfides, 364, 366

sulfites, 367

sulfonic acids, 447

sulfur, 87, 358

sulfur dioxide, 366

sulfur family, 358-373

sulfur trioxide, 368

sulfuric acid, 369

sulfurous acid, 367

superconductivity, 542, 617

supercooling, 45

supersaturation, 229

surface tension, 39

suspensions, 200, 478

supersaturation, 229

symbols, chemicals, 66; nuclear, 633

synchrocyclotron, 635

synchrotron, 635

talc, 467t, 473, 516

tantalum, 535t, 541

tar, coal, 438

tartaric acid, 447

technetium, 568t, 572

Teflon, 339, 455

tellurium, 359, 372

temperature, defined, 9; scales, 10

factor, fa equihbrium, 120

tempering, 551
tertiary carbon atom, 424
tetraethyl lead, 344, 438, 507
tetrahedral structure, 419,
tetrathionate, 372
thallium, 602
theory, defined, 3
thermal neutrons, 638, 642
thermite, 600

thermochemical equations, 86
thermodynamics, first law, 81; second law,
83; third law, 84
thermometric fluids, 589
thermometric scales, 10
thermonuclear reactions, 645
thermoplastic, 456
thermosetting, 456
thiosulfates, 372
thiazole, 436
third order reaction, 102
thorium, 822
thyroxin, 464
tin, 405, 604, 606
tin plate, 553
titanium, 535t, 539
titration, 272
TNT, 449

toluene, 432, 435t, 436
tracers, radioactive, 448, 646
trans-compounds, 428
transistors, 605

transition elements, general properties, 529-
534

transition points, 37
transition series, defined, 163, 529
transmutation, nuclear, 620, 633
transuranium elements, 537, 044
triehloromethane, structure, 178
trinitrotoluene, 449

triple bond, 173, 431; in nitrogen, 378; in
carbon monoxide, 413
triprotie acids, 267
tritium, 148, 203, 645, 647
tungsten, 564t, 566
Turnbull's blue, 555
Tyndall effect, 482

ultramicroscope, 482
Uncertainty Principle, 152
unimolecular reaction, 99, 101
unit cell, 36

units of measurement, 8, 477n
unsaturated compound, 173, 429
uranium, 537

uranium hexafluoride, 24, 641
uranium isotopes, 24, 638; separation, 641
uranium-lead ratio, 647
uranium Ore, 537

Subject Index

691

uranium series, decay path, 622
urea, 418

valence, 74; atomic structure and, 162;

electrovalence, 172; covalence, 172
Van de Graaff generator, 634
van der Waals equation, 27; forces, 38, 189,
392, 469, 484, 491
vanadium, 535 fc, 540
van’t Hoff, law of, 120
vapor pressure, 41
vinegar, 447
vinyl polymers, 454
viscose process, 461
viscosity, 40; of sulfur, 361
vital force theory, 418
vitamin A, structure, 464
vitamins, 464
vitreous substances, 38
volt, 328

voltaic cell, 306-321
volume fraction, 222
volumetric flask, 223
vulcanization, 453

washing soda (see sodium carbonate)
water, 206-215; as catalyst, 209; chemical
properties, 209; density of, 208; of hydra-
tion, 211; hardness in, 522; heavy, 203,
642, 643; of hydration, 211; impurities
in, 214; ionization of, 209, 268; physical
properties, 207, 212; as solvent, 209; struc-
ture of, 186, 206; vapor, 214

water gas, 413

water glass, 471

water purification, 215, 000

water softening, 522-525

watt, 328

waxes, 448

weak acids, 260

weak bases, 262

weight, equivalent, 55, 77, 302

weight fraction, defined, 221

white lead, 610

white phosphorus, 392

wolfram (see tungsten)

wrought iron, 548

X-rays, 138; crystal analysis, 36, 490; dis-
covery, 137; spectra, 144, 151; tube, 187
xenon, 612, 613t; compounds of, 618
xylene, 432, 435t, 436

yttrium, 534, 535t

zeolite, 467t, 525
zero point energy, 617
zinc, 529-531, 585-588
zinc blende, 360, 586
zinc-chlorine voltaic cell, 308
zincite, 586
zircon, 467t, 539
zirconium, 535t, 539
zone refining, 587, 605
zwitterion, 449

Periodic Classification of the Elements

£8

Xe
! 54

5 s

£s

!§B

tsg

t§8

i§S

•3

•

VII B

« cT

« PS

■

«§

VI B

** to

33 O

PC PC

■

O oc

«s

<*> CM

H vo

Po
i 84

VB j

rt JJ*

33 O
§ of

■

jo*-*

cnw

PQoS

IVB

SCO

PC PC

■

CJ co

1 Sn

-CcQ

0.00

III B

■

CQ lO

rjco

■

<5>

VIII

<3 Si

£§

PC'tr

VII A

£3

4*8

VIA

«*

■

£2

■

( IV A

* M

S «

■

*h4

. «•»

D 9,

PS PS

■

►*8

*— <

•ag

4®

■

<38

<2 go

PC oo

•M*

EC*-*

a<*>

s§S

£fe

■a

HU*

Period 1

»§

c3S

<2S

£S

■3$