PRINCIPLES 
OF 
CHEMISTRY 
HILDEBRAND Principles of Chemistry 
BY 
JOEL H. HILDEBRAND, Ph.D. 
Associate Professor of Chemistry in the 
University of California. 
(PRELIMINARY EDITION) 
Copyright, 1917, by Joel H. Hildebrand 
Lederer, Street & Zeus Co., Publishers, 
Berkeley, California. 
. 1917 CONTENTS 
CHAPTER I. 
Page 
KINDS OF MATTER 7 
The Domain of Chemistry-Properties Used to Recognize Substances- 
Heterogeneous Mixtures-Solutions-Compounds and Elements-Sym- 
bols-Classification of Elements. 
CHAPTER II. 
THE GAS LAWS AND THE KINETIC THEORY * 15 
Kinetic Theory-Relation between Quantity and Pressure-Relation be- 
tween Pressure and Volume-Relation between Pressure and Concentration 
-Relation between Pressure, Mass and Speed of the Molecules-Effect 
of Temperature-Absolute Zero of Temperature-Effect of changing both 
Temperature and Pressure-Partial Pressure of different Gases in Mix- 
tures-Avogadro's Rule. 
CHAPTER III. 
WEIGHT RELATIONS IN CHEMICAL REACTIONS. ATOMIC 
THEORY 24 
Law of Conservation of Mass*-Law of Simple Multiple Proportions- 
Atomic Theory-Atoms and Molecules, Formulas-Numerical Values 
of Atomic Weights - Interpretation of Formulas - Gram-Atoms and 
Gram-Molecules or Mois-Chemical Equations-Calculation of Weight 
Relations of Atomic Weights*-of Formulas. 
CHAPTER IV. 
VOLUME OF GASES. MOLECULAR WEIGHTS 31 
Choice of Formulas and Corresponding Atomic Weights-Changes in 
Volume or Pressure in Gas Reactions-Volume of One Mol of Any 
Gas-Calculation of Weight of a given Volume or vice-versa-Deter- 
mination of Relative Weights of a Gas and Air-Determination of 
Molecular Weight-Direct Relation between Volumes of Gases and 
Weights of other Substances in Reactions-Complete Interpretation 
of Chemical Equations-Molecular Weights of dissolved Substances. 
CHAPTER V. 
TYPES OF CHEMICAL COMPOUNDS 39 
Relation between Stability of Compounds and Metallic or Non-metallic 
Character of their Constituent Elements-Oxides and Sulfides-Halides 
-Relation of the Relative Positive and Negatiye Characters of Elements 
to the Reactions of their Compounds, (a) Replacement, (b) Metathesis, 
(c) Addition - Classification of Compounds - Acids - Bases - Salts 
-Anhydrides. 
CHAPTER VI. 
VALENCE. WRITING EQUATIONS ......T 49 
Assignment of Valence Numbers-Elements exhibiting more than one 
Valence-Use of Valence in classifying Compounds-Writing Formulas 
with the Aid of Valence-Writing Equations-Oxidation and Reduction. 
CHAPTER VIL 
CHEMICAL NOMENCLATURE 55 
Names of Elements-Most positive Element named first-Binary Com- 
pounds designated by Suffix -ide-Compounds whose positive Element 
may show different Valences-Oxy-acids and Salts-Sulfo-acids and Salts 
-Complex Halogen Acids and Salts-Acid and Basic Salts-Partially 
Dehydrated Acids and their Salts . 
CHAPTER VIII. 
QUANTITATIVE RELATIONS INVOLVING SOLUTIONS OF 
KNOWN CONCENTRATION .. 60 
Concentration in Terms of Mois per Liter-Titration*-Standardizing 
Solutions-The Equivalent: Normal Concentration-Summary-Ex- 
amples-Titration involving other Types of Reaction. 
CHAPTER IX. 
THERMO-CHEMISTRY 65 
The Calorie - Experimental Determination of Heats of Reaction - 
Numerical Representation of Heats of Reaction-Effect of the State of 
the Reacting Substances-Indirect Determination of Meats of Reaction 
-Fuel Value of Foods. ' CHAPTER* X. - Page 
BEHAVIOR OF SUBSTANCES DISSOLVED IN WATER. IONIC 
THEORY . 71 
Abnormally great Lowering of the Freezing Point-Independent Migra- 
tion of Ions in Electrolysis-Ions carry definite electric Charges-The 
Atom of Electricity-Properties of dilute aqueous Solutions of Strong 
Electrolytes are the Sum of independent Sets of Properties, hence 
independent Substances are present-Heats of Reaction in dilute 
Solution of Strong Electrolytes depend on reacting Ions only-Degree 
of Ionization of strong Electrolytes-Weak Acids and Bases are less 
ionized-Weak Salts-Ionization of weak polybasic Acids in Steps'- 
Ionization in other Solvents than Water-Ionization of fused Salts- 
Ionization of Water. 
CHAPTER XL 
THE SPEED OF CHEMICAL REACTIONS 81 
Application of the Kinetic Theory-Effect of Temperature-Effect of 
Concentration-Numerical Relation between Concentration and Speed 
-Effect of Stirring and of Contact Surface in Reactions occurring at 
the Boundaries between Phases-Catalysts-Enzymes. 
CHAPTER XII. 
CHEMICAL EQUILIBRIUM. THE EFFECT OF CONCENTRATION.... 87 
Reversibility of Chemical Reactions-It is possible to have all Sub- 
stances involved present together in Chemical Equilibrium-Effect on 
Equilibrium of changing Concentration-Quantitative expression of the 
Mass Law-Dissociation Constants of Electrolytes-Chemical Equations 
do not indicate Amounts present at Equilibrium. 
CHAPTER XIII. 
CHEMICAL EQUILIBRIUM. PROPERTIES WHICH MAY AFFECT 
CONCENTRATION 96 
Volatility-Solubility-Ionization of Water-Ionization of weak Acids, 
Bases and Salts-Complex Ions, (a) Ammonia, (b) Cyanide, (c) Halide, 
(d) Oxalate-Competition between the foregoing Factors, (a) Hydro- 
lysis, (b) The Solution of Hydroxides, (c) of Oxides, (d) of Ampho- 
teric Hydroxides, (e) of Sulfides in Hydrogen Ion, (f) .of Amphoteric 
Sulfides, (g) of Salts of other weak Acids, (h) by forming weak Salts, 
(i) by forming complex Ions, (j) by forming Ions of Amphoteric 
Hydroxides, (k) Systematizing the Reactions of a given Ion-Miscel- 
laneous Transformations. 
CHAPTER XIV. 
CHEMICAL EQUILIBRIUM. THE EFFECT OF PRESSURE AND 
TEMPERATURE 116 
Effect of changing total Pressure-Effect of changing Temperature- 
Simultaneous Consideration of all Factors governing Reactions, (a) 
"Contact Process," (b) Synthesis of Ammonia, (c) Synthesis of Nitric 
Oxide, (d) "Cracking" of Petroleum. 
CHAPTER XV. 
OXIDATION AND REDUCTION - - 123 
Writing Equations, (a) The Substances produced, (b) Assignment of 
Valence Numbers, (c) Change in Valences, (d) Balancing Valence 
Changes, (e) Final Balancing-Relative Oxidizing and Reducing Powers 
of Substances, (a) Metals and their Ions, (b) Solubility of Metals in 
Water, Acids and Alkalis, (c) General Table-Oxidizing Power and 
Speed of Oxidation-Substances acting both as oxidizing and reducing 
Agents-Solution of Insoluble Sulfides-Storage Batteries. 
CHAPTER XVI. 
THE PERIODIC SYSTEM -136 
Forms of Periodic Table-Density-Melting Point-Tensile Strength- 
Valence-Noble and Base Metals-Metals and Non-metals-Strength of 
. Acids and Bases-Amphoteric Sulfides-Stability of Hydrides-Compari- 
son of Oxides and Hydrides-Elements that form Complex Ions'- 
Resemblance of first member of Group to second of succeeding Group- 
Radioactive Elements-Prediction of unknown Elements. 
CHAPTER XVII. 
DISPERSED SYSTEMS •.... -146 
Brownian Movement-Optical Properties' of Colloidal Solutions- 
Absorption-Electrical Migration-Colloidal Particles-Relation between 
Charge and Stability, Coagulation-Emulsions-Jellies-Emulsion Col- 
loids-Protective Colloids-Preparation of Colloidal Solutions, (a) 
Dispersion Methods, (b) Condensation Methods. PREFACE. 
This book has been written primarily to serve as a reference book 
in the course in General Chemistry and Qualitative Analysis in the 
University of California. There are, it is true, many excellent texts 
in existence, but whenever an instructor uses a text written by someone 
else he faces the dilemma that a departure from the order and manner 
of presentation of the printed text results in the loss of much of its 
value to the student as a reference book, while a close adherence to 
it, on the other hand, interferes with the freedom and individuality, and 
hence with the enthusiasm and effectiveness of the teacher. 
The difficulty is not altogether avoided by writing one's own book, 
as probably every teacher is tempted to do, for a book tends to end the 
period of experimentation upon the class and to substitute therefor year 
after year of complacent stereotyped repetition. A wise old master of 
the teaching art doubtless had this in mind when he replied to the 
question as to why he did not write a text book by saying, "I would 
not be such a fool, for then I could not change my mind." 
Now there are two facts which seem to the writer to indicate a 
way to avoid these difficulties. The first is that the difficulties of the 
student are not concerned with the simple matters of fact so much as 
with general principles, methods and points of view. It is comparatively 
easy for him to learn the properties and uses of" substances from any 
text-book, lecture or laboratory course, but much harder for him to learn 
to write equations, to use them in calculations of weight and volume, 
to understand and apply such principles as the Ionic Theory or the Mass 
Law. Repetition, thought and practice are required in such cases. 
The second of these facts is that writers of text-books have been in 
the habit of introducing the general principles in the particular connec- 
tions which seem best to them at the time of writing, and it is usually 
unsatisfactory to change the order followed by the book. The difficulty 
is not so much that the material is not well presented, but rather, that 
a topic like the Mass Law, for example, is introduced in a chapter 
headed sulfur, and must remain there if the illustrations are to be of 
any value, whereas, of course, this principle no more belongs under 
sulfur than it does under any other element, and the instructor using 
the book may wish to introduce it at some totally different place. On 
the other hand, the purely descriptive material could be rearranged almost 
ad libitum. 
It seems to the author, accordingly, that the ideal text-book would 
consist of two volumes, one of General Principles, the other of Descrip- 
tive Chemistry, the latter a book of moderate dimensions on the plan 
of Roscoe and Schorlemmer's "Treatise on Chemistry" or Abegg and 
Auerbach's "Handbuch der anorganischen Chemie." Since, however, 
the descriptive material is available in so many existing books, the limit 
of the author's present ambition is a presentation of the Principles of 
Chemistry in such a way as to allow, in its use, the maximum flexibility 
consistent with the subject. It is hoped that the treatment of these 
principles in this volume is such that the various chapters and sections 
can be used in connection with almost any arrangement of descriptive 
material. Furthermore, although a certain logical order has been followed 
4 in the text, the effort has been to make the presentation of each topic 
such that considerable departure from the given order might be possible 
without loss of intelligibility. As a matter of fact, the order in which 
the topics are taken up in our own course does not fully correspond 
to that here used. Some parts of a chapter such as that on the Periodic 
System, for example, may well be introduced in a course much earlier 
than other parts, and the arrangement of paragraph headings throughout 
has been made with the end in view of facilitating such rearrangement. 
This separate treatment of Principles is, however, not only undertaken 
to allow flexibility but also to emphasize the Principles of Chemistry to 
a degree that seems justified by the present status of chemical science. 
The earlier history of all natural sciences is that of description and 
classification, and their adepts were necessarily men of well developed 
powers of observation and memory rather than of inference and deduc- 
tion. We still start the beginner in the historic way, asking him to 
commit to memory facts concerning the elements and their compounds. 
With the development of the science, however, there has arisen a con- 
siderable body of general principles which can serve as the basis for 
deduction, and the profound change in the character of chemical research 
during the last generation bears witness to the development of the later 
type of scientist, the one who seeks causes, and aims to develop means 
for predicting and controlling phenomena. In recognition of this change 
elementary texts have been "modernized" by the introduction of chapters 
on the Ionic Theory, Phase Rule, Chemical Equilibrium, etc., which are 
often intrusions foreign to the rest of the book, and so they never become 
part of the usable tools of the student. Although, of course, chemistry 
is yet far from being an exact science, we nevertheless have made far 
more progress in this direction than is recognized in the traditional plan 
of instruction, and progress in both pure science and industry is becoming 
constantly more dependent upon ability to predict new facts by the aid 
of general principles rather than upon a knowledge of existing facts. 
A remarkable plea for such a change in the teaching of chemistry 
is made by the late Professor H. LeChatelier in the preface to his book 
"Legons sur Carbone." The entire preface is worth reading by all teachers 
of chemistry. In it he says: "Pour comprendre la necessite de modifier 
profondement les methodes d'enseignment de la chimie, il suffit de com- 
parer un cours de chimie et un cours de physique .... Dans 1'enseign- 
ment de la physique on ne parle que des lois des phenomenes naturels. 
En chimie, au contraire, c'est une enumeration indefinie de petits faits 
particuliers: Formules de combinaisons, densites, couleurs, action de tel 
ou tel corps, recettes de preparation, etc L'enseignement de la 
chimie minerale est completement immobile depuis soixante-quinze ans." 
He goes on to contrast this orthodox method of presentation of chemistry 
with the chapter headings from the work of Lavoisier: "De la combination 
du calorique.-De la formation des fluides aeriformes.-Nomenclature des 
acides.-Quantite de chaleur degagee dans les differentes combustions.- 
De la fermentation putride.-Des seis neutres, etc.," all of which deal 
with general subjects instead of descriptions of substances. He admits 
of course, that it is not yet possible to consider chemistry an exact science 
whose presentation can be made solely on the basis of general principles, 
but holds, nevertheless, that the progress that has been made in this 
direction has been far from recognized in the traditional plan of presen- 
tation. 
It is not intended, by any means, that this volume should furnish the 
entire material for a course in general chemistry. It is designed distinctly 
5 as a reference book rather than a text book. Our own course includes 
much of the traditional descriptive material, and besides, for entrance to 
the University course we prescribe High School chemistry, which is 
usually largely descriptive in nature. Our aim is, however, to develop 
principles rather than simply to study elements and compounds descrip- 
tively. The laboratory manual for our course has been written with 
this end in view by Professor W. C. Bray and Dr. Ludwig Rosenstein. 
Considerable thought has been expended in devising the exercises at 
the end of each chapter. The material in the text does not consist of 
facts to be memorized so much as of tools, to learn to use which requires 
a great deal of drill. These exercises constitute, therefore, an important 
part of the book. 
For whatever merit the book may possess, the author desires to make 
very great acknowledgements to his colleagues. Our own course has 
been made to a large extent in the weekly conferences of instructors, and 
the freedom of criticism always enjoyed in these meetings has had a 
very large effect in determining the contents of this volume. Especial 
thanks are due for the valuable criticism given by Professor G. N. Lewis, 
Professor W. C. Bray, Dr. W. L. Argo and Dr. G. E. Gibson, who were 
kind enough to read the manuscript. 
Joel H. Hildebrand. 
Berkeley, Cal., 1917. 
6 CHAPTER I. 
KINDS OF MATTER. 
The Domain of Chemistry. Various aspects of the material things 
we see about us may attract our attention, such as the uses to which 
these things may be put, the beauty of form or color they may possess, 
or the materials out of which they are made. It is the last of these that 
is the concern of chemistry. Two vases may have nothing in common 
from an artistic standpoint, and a vase and a plate may serve entirely 
different purposes, and yet, from a chemical point of view, all three may 
be identical by reason of the material, porcelain, out of which they are 
made. Chemistry does not concern itself with articles or objects, as 
typified by such words as chair, pen, bottle, nail, wire, but rather with the 
substances or materials out of which these and other objects may be 
made, as typified by such words as wood, glass, iron, copper, clay, sugar. 
It is the task of the chemist to ascertain all of the qualities or properties 
of every kind of matter, so that the vast number of different kinds may 
be recognized and distinguished from each other. 
It is also the task of the chemist, to determine how these substances 
or materials may be obtained, or, when obtained, how they may be 
preserved. The materials out of which objects are made often do not 
in the free state, and the enormous quantities of it which are used for 
making such a multitude of things must be obtained from the various 
iron ores, by heating them with coke or charcoal in a blast furnace. It 
is of course very important to know how much coke to use with a given 
amount of ore, and howr much iron should be obtained from the ore, 
because the cost of iron is partly determined by these factors. The 
chemist, therefore, must be concerned not only with iron, its nature and 
properties, but also with the conditions necessary for getting it from its 
ores, and the quantities of the materials involved in the process. And 
then, finally, he must determine how to prevent loss of iron from rusting. 
The housewife, in making biscuits with the aid of baking soda and 
sour milk, finds that if too little soda is used the biscuits will not be 
"light," whereas if too much soda is used they will be yellow, and the 
taste will be impaired. She must therefore know, not only that when 
soda and sour milk are mixed they give a gas that will make the dough 
porous, but also the relative amounts necessary for the best results. 
In the processes just referred to it is obvious that new kinds of 
matter have been produced. The iron ore and the coke are quite evidently 
different stuff from iron. The sour milk loses its sour taste when mixed 
with sufficient soda, and a gas is evolved quite different from either of 
the previous materials. Any such change, in which the kind of matter is 
altered, is called a chemical change, or a chemical reaction. When 
iron is drawm into wire, or made into nails, it is still iron, and no chemical 
change has taken place, but when it rusts it is no longer iron but a brown 
powder of very different nature from iron. When marble is cut into 
various shapes it is still marble, whether in the form of a statute, a table 
top or a building block, possessing all of the characteristics by which we 
recognize it as marble. If, however, we heat it for a while to a high 
temperature it changes in appearance, losing weight, becoming the 
familiar substance we call quicklime. This again is a chemical reaction, 
7 since a new substance is produced. Another chemical reaction takes 
place if we allow water to come in contact with the quicklime. The 
color changes from light grey to white, heat is evolved, the mass swells 
and falls apart to a powder, or, if water is in excess, forms a paste. 
The resulting material is called slaked-lime, a substance with properties 
different from those of marble or quicklime. 
Often we direct our attention, in dealing with chemical reactions, 
not so much to the nature of the substances involved, or to their relative 
amounts, as to the energy used up or liberated by the process. When 
oil burns it unites with something from the air, as we must infer from 
the necessity of an air supply, and forms gaseous products, which ordi- 
narily do not attract our attention except for the necessity of providing 
for their escape. The most important thing is the energy made available 
by the reaction, which we may wish to use for light, heat or power. In 
a dry cell zinc is used up and electrical energy is obtained. The question 
as to what becomes of the zinc is overshadowed by the question as to 
how much electrical energy should be gotten from a definite amount of 
zinc. 
We see, then, that chemistry is concerned with substances, and their 
properties; with the changes or reactions whereby other substances are 
formed; with the conditions necessary for bringing about or preventing 
these reactions; and with the relative amounts of matter and energy 
involved. 
Some hint may be gotten from the illustrations given above of the 
immense importance of chemistry to the material welfare of mankind. 
Great service has been rendered by those who have applied it to the 
satisfaction of human needs. It must not be forgotten, however, that these 
needs are mental and spiritual as well as material, and that to enlarge 
man's mental horizon is quite as worthy an endeavor as it is to increase 
his physical comfort. To seek out and discover Nature's mysteries is 
not a pursuit that requires utilitarian justification. The world owes an 
incalculable debt to the explorers who have led the way to new fields of 
thought and endeavor. It is these men who pave the way not only for 
the miner, the railroad builder and the farmer, but, most important of all, 
for others who love Nature and whose spirits are enriched by her knowl- 
edge. To those having this pioneer spirit chemistry offers wide uncharted 
realms, to explore which is among the most fascinating of pursuits. 
Properties Used to Recognize Substances. Many substances have 
characteristics so marked that there is little likelihood of their being 
confused with other substances. The color of copper distinguishes it from 
other metals. The elacticity of rubber, the taste of sugar, the odor of 
ether, serve to distinguish these substances, even in the absence of other 
tests, from those encountered in ordinary life. We constantly apply 
such obvious tests as those of color, luster, hardness, odor and taste. 
We distinguish quite readily between solids, liquids and gases. Very 
often, however, we need to refine our methods of observation, or to rise 
less obvious properties for distinguishing substances. Instead of being 
content with saying that lead is "heavier" than iron, we find it desirable 
to express this "heaviness" or density numerically by determining the 
weight of a unit volume, usually the number of grams per cubic centi- 
meter. The refractive index if a transparent body can be determined 
with great exactness and is a most valuable means of identification. 
The coefficient of expansion, the conductivity for heat or electricity, the 
boiling point, the melting point, the heat of fusion or of vaporization, 
8 are all quantities having characteristic values for individual substances, 
and which may be accurately determined by appropriate methods. 
The chemical changes which substances may undergo offer important 
means of identification. Gold may be indistinguishable from certain 
samples of brass in appearance, but nitric acid at once differentiates them 
by dissolving the latter. Powdered talc and starch in a face powder 
might readily be differentiated by the fact that the latter will swell up 
and dissolve in boiling water, or that it will readily burn. 
Heterogeneous Mixtures. As we attempt to apply our tests to dis- 
tinguish substances, we notice that some materials give a rather ambiguous 
answer. What shall we consider to be the color, density or hardness of 
a piece of granite? On close examination we find it composed of several 
kinds of mineral, having different degrees of hardness, different colors 
and different properties in general. Although we might determine the 
density of a given piece of granite it would be folly to talk about the 
density of granite in general, because the constituents of granite are 
present in different proportion in different samples. Granite is obviously 
a mixture and not a pure substance, and its properties change abruptly 
from one small region to another. These regions within which the 
properties suffer no abrupt changes, in this case quartz, mica and feldspar, 
are called phases. A mixture containing more than one phase is called 
a heterogeneous mixture. Another example is muddy water. Here we 
cannot see the different phases quite so readily, but a microscope or a 
filter shows quite distinctly that two phases are present. Some alloys, like 
solder, appear quite uniform to the eye, as if they contained but one phase, 
whereas, by careful etching with acid and using a microscope with 
reflected light, it is evident that the alloy is heterogeneous, more than 
one phase being present. 
Evidently, then, the first step in our examination and classification 
of matter is to separate heterogeneous mixtures into their component 
homogeneous phases. This may be accomplished by purely mechanical 
means, which may vary with the nature of the separation. A suspended 
solid may be removed from a liquid by filtration. A fog (liquid particles) 
or a smoke (solid particles) may be removed from a gas by the same 
means. Two solids, may be picked apart, though such a process would 
often be impracticable. Where two kinds of solid particles have different 
densities it is sometimes possible to employ a liquid whose density is such 
that one solid will rise to the surface while the other stays on the bottom, 
as would be the case if a mixture of sand and sawdust were thrown into 
water. Two liquids which do not mix, like kerosene and water, are 
easily separated by skimming or some equivalent process. Where one 
is emulsified in the other, the highly dispersed globules of the one may 
often be brought together prior to a separation, as when butter is obtained 
from cream. 
Solutions. Having separated heterogeneous mixtures into single 
phases, our next step is the examination of these phases. We will find 
that a phase may consist of one substance, or of two or more substances, 
when it would be called a solution. For example, we may take sugar 
and water, each a pure substance, and make a solution, uniform throughout 
in properties. The sugar cannot be seen with a microscope or filtered 
out by a filter. Such a solution has the same appearance as a pure 
substance. It can be distinguished from a pure substance, however, by 
the fact that its properties depend on the relative amounts of the substances 
9 from which it is made. The properties of a pure substance, are always the 
same at the same temperature and pressure. Pure water always has the 
same density provided that the pressure and temperature at which the 
density is measured are kept the same. Ice always melts to water at 
the same temperature unless a different pressure is applied to it, and 
even then for every pressure there is a definite melting point. Similarly 
occur in a free state in nature. Iron, for example, is found but rarely 
the boiling point of water depends only upon the barometic pressure. 
The properties of solutions, on the other hand, depend on the composition 
as well as on the temperature and pressure. The properties of a sugar 
solution are not implied when the temperature and pressure are stated, 
as is the case with pure sugar or pure water. In order to predict the 
Fig. 1. Hardened Steel. 
Fig. 2. Wrought Iron. 
density or sweetness of such a solution, or the temperature at which 
it would freeze or boil, we would have to know its composition, i. e., how 
much sugar there is in a given amount of water. In general, then, 
the properties of a solution depend not only upon the temperature and 
pressure, hut upon the composition as well. 
The term solution must not be restricted to liquid solutions, although 
they furnish the most familiar examples. All gases are completely miscible 
with each other, forming but one phase, so that every mixture of gases is 
a solution. Alloys of silver and gold, no matter what the relative amounts 
Fig. 3. Grey Cast Iron. 
Fig. 4. Annealed Steel. 
10 of the two metals, contain but one kind of crystal, the properties of which, 
such as density, color, electrical conductivity, melting point, change con- 
tinuously with the composition. This is a solid solution, and may be 
contrasted with alloys of copper and silver, which contain two kinds of 
crystals through a certain range of composition. This could be shown, 
first, by the microscope, second, by the fact that alloys of these metals, 
though of different composition, would all begin to melt at the same 
temperature. This is true because a change in the proportion of the 
metals produces no change in the nature of the two crystals but simply 
in their relative amount, and the melting and boiling points of substances 
never depend upon how much of the substance is taken. 
Hardened steel is a solid solution, for though it contains carbon it 
shows but one kind of crystal under the microscope as shown in Fig. 1, 
just as does wrought iron, Fig. 2, which is nearly pure iron. Grey cast 
iron, however, contains graphite crystals, as shown in Fig. 3, and is 
obviously heterogeneous, as is annealed steel, Fig. 4, which consists of an 
intimate mixture of pure iron, of the kind shown in Fig. 2, with an iron 
carbide. 
Solutions and pure substances can be differentiated experimentally by 
changing the state of aggregation of part of the material, as by melting, 
crystallizing or boiling, whereupon one component usually changes in state 
more readily than the other, so that the new phase will be of different 
composition from the old and hence have different properties. Air can be 
shown to be a mixture of gases by liquefying it and allowing the liquid 
to evaporate slowly. During the process it will be found that the boiling 
point changes, that the color of the liquid residue changes, and that the 
air evaporating gradually changes in properties, as shown, for example, 
by the fact that the brightness with which a candle will burn in it increases 
gradually from the first to the last portions. 
If liquid air is distilled in a scientifically constructed still it is possible 
to separate it into two nearly pure constituents, just as water and alcohol 
can be separated. One of these constituents is found to be slightly 
lighter than air; it can be condensed to a colorless liquid boiling at 
-194° C; it is very inert chemically, reacting with but few other sub- 
stances. The name nitrogen is given to it. The other constituent is 
slightly heavier than air; it gives when condensed at low temperatures a 
blue liquid boiling at -182.5°C, and it reacts readily with many sub- 
stances. Ordinary combustible materials burn in it very vigorously. 
It is called oxygen. 
Compounds and Elements. Having found out how to distinguish 
between mixtures and pure substances, let us ask the further question: 
of what are pure substances made? This question may be answered by 
means of appropriate chemical reactions. We may determine what sub- 
stances may be made to unite to form the substance we are investigating, 
a process called synthesis, meaning putting together; or we may separate 
the substance into simpler constituent substances, a process called analysis, 
meaning taking apart. If molten lead is exposed to the air it becomes 
coated with a scum, and if this scum is constantly removed the lead 
entirely changes into another substance, called litharge, with a considerable 
increase in weight, indicating that something has been absorbed from the 
air. Of the two main constituents of the air only one, the oxygen, will 
transform the lead in this way. We would conclude from these facts 
that the familiar substance litharge is composed of lead and oxygen, or, 
11 as chemists would say, is a compound of lead and oxygen. This would 
be the synthesis of litharge. 
On the other hand, we might take some litharge, put it into a 
crucible, heat it to a red heat, pass over it some illuminating gas, and we 
would find that it changed into lead, with a loss in weight. Evidently 
litharge contains lead plus some other constituent which has been removed 
by the illuminating gas. Further experiments might be performed to 
prove this other constituent to be oxygen. This would be the analysis of 
litharge, and from it, likewise, we would conclude that litharge is a 
compound of lead and oxygen. 
Again, if we pass an electric current through water, using platinum 
or gold electrodes to introduce the current into the water, we find that 
gases are given off at the electrodes and that water is used up. If we 
collect the gas given off at the negative electrode we find that it is very 
light, soap bubbles filled with it rising like balloons, and that it will 
burn in air with a faintly luminous very hot flame. To this gas is 
given the name hydrogen. The other gas is slightly heavier than air, 
and will not burn in air, but burning substances thrust into it will continue 
to burn with greatly increased vigor and brilliancy. This gas we may 
identify as oxygen, which we have found to be also one of the constituents 
of air. From this analysis of water we would conclude that water is 
a compound of hydrogen and oxygen. The analysis might be made in 
many other ways; and we might also synthesize water by allowing hydro- 
gen and oxygen to combine, which they do very readily. 
Now it has never been possible to resolve lead, oxygen and hydrogen 
into simpler substances. These substances, and all others which have 
never by ordinary chemical reactions* been resolved into simpler sub- 
stances, are called elements. There are known at the present time about 
eighty-three elements. They can form almost endless combinations with 
each other; considerably over one hundred thousand are known and new 
ones are constantly being prepared. 
Table I contains a list of the chemical elements. Many of them are 
well known substances. Many are rare, occurring in but small quantities. 
Symbols. Along with each element in the table is given a char- 
acteristic number, whose significance will be explained later, and also a 
symbol used to designate the element. It consists of the initial letter of 
the name, followed, where necessary to distinguish between several ele-' 
ments having the same initial, by a second appropriate letter. In many 
cases the symbol follows the Latin name of the element rather than the 
English name, allowing the symbols to be practically the same in all the 
principal languages. Thus, the symbol for iron is Fe from the Latin 
ferrum; that for gold is Au, from the Latin aurum, etc. 
Classification of the Elements. The elements may be classified in 
various ways, all of which might be useful for certain purposes. Some 
of them are solids, some liquids, and some gases, at ordinary temperature 
and pressure. This is, however, not a very fundamental distinction, as 
each element may be made to exist in any of these states by suitable 
alterations in temperature and pressure. A most useful classification is 
that furnished by the Periodic System, whereby the elements are arranged 
in families or groups which show similar properties. This arrangement 
will be used very extensively later, after we have explained the basis 
upon which it is made. One of the most significant of all divisions is 
that into metals and non-metals. The general characteristics of metals 
12 and non-metals are familiar to all. Metals are characterized by an appear- 
ance called metallic luster, by malleability and ductility, and by relatively 
high conductivity for heat and electricity. These physical differences are 
Table I. 
Symbol Atomic 
Weight 
Symbol Atomic 
Mo 
Weight 
96.0 
Aluminum 
Al 
27.1 
Molybdenum 
Antimony 
Sb 
120.2 
Neodymium 
Nd 
144.3 
Argon 
A 
39.88 
Neon 
Ne 
20.2 
Arsenic 
As 
74.96 
Nickel 
Ni 
58.68 
Barium 
Ba 
137.37 
Niton (radium emanationL.Nt 
222.4 
Beryllium 
Be 
9.1 
Nitrogen 
l.N 
14.01 
Bismuth 
....Bi 
208.0 
Osmium 
Os 
190.9 
Boron 
B 
11.0 
Oxygen 
O 
16.00 
Bromine 
Br 
79.92 
Palladium 
Pd 
106.7 
Cadmium 
Cd 
112.40 
Phosphorus 
P 
31.04 
Caesium 
..Cs 
132.81 
Platinum 
Pt 
195.2 
Ca 
40.07 
Potassium 
K 
-39.10 
c 
12.005 
Praseodymium 
Pr 
140.9 
Cerium 
Ce 
140.25 
Radium. 
Ra 
226.0 
Chlorine 
.Cl 
35.46 
Rhodium 
Rh 
102.9 
Chromium 
Cr 
52.0 
Rubidium 
Rb 
85.45 
Co 
58.97 
Ruthenium 
Ru 
101.7 
Columbium 
.Cb 
93.5 
Samarium 
150.4 
Copper 
Cu 
63.57 
Scandium 
Sc 
44.1 
Dysprosium 
Dy 
162.5 
Selenium 
Se 
79.2 
Erbium. 
Er 
167.7 
Silicon 
Si 
28.3 
Europium 
Eu 
152.0 
Silver 
Ag 
107.88 
Fluorine 
F 
19.0 
Sodium 
Na 
23.00 
Gadolinium. . 
. .Gd 
157.3 
Strontium 
Sr 
87.63 
Gallium 
Ga 
69.9 
Sulfur 
S 
32.06 
Germanium 
Ge 
72.5 
Tantalum 
Ta 
181.5 
Gold 
....Au 
197.2 
Tellurium 
Te 
127.5 
Helium 
He 
4.00 
Terbium 
Tb 
159.2 
Holmium 
Ho 
163.5 
Thallium 
T1 
204.0 
Hydrogen 
H 
1.008 
Thorium 
Th 
232.4 
Indium 
Jn 
114.8 
Thulium 
Tm 
168.5 
Iodine 
I 
126.92 
Tin 
Sn 
118.7 
Iridium 
Jr 
193.1 
Titanium 
Ti 
48 1 
Iron 
Fe 
55.84 
Tungsten 
W 
184.0 
Krypton 
Kr 
82.92 
Uranium.. 
.U 
238.2 
Lanthanum 
La 
139.0 
Vanadium 
V 
51.0 
Lead 
Pb 
207.20 
Xenon 
Xe 
130.2 
Lithium 
Li 
6.94 
Ytterbium (Neoytterbium).Yb 
173.5 
Lutecium 
Lu 
175.0 
Yttrium 
Yt 
88.7 
Magnesium 
Mg 
24.32 
Zinc 
Zn 
65.37 
Manganese 
Mn 
54.93 
Zirconium 
Zr 
90.6 
Mercury 
Hg 
200.6 
* Certain elements decompose into others by a different type of change called 
radioactive, to be treated in Chapter XVI. 
13 accompanied by fundamental differences in chemical behavior to be 
discussed later. 
From what has just been said it must not be supposed that the 
metallic state is peculiar to elements. There are many compounds which 
are metallic also. The further classification of compounds will, however, 
be deferred to a later chapter, where we will see that the metallic or 
non-metallic character of the component elements furnishes the basis 
for this classification. 
It must not be thought that all elements fall readily into these two 
classes. In fact the two classes shade into each other, as illustrated by 
carbon, which in the formi of graphite has some metallic luster and is a 
fair conductor of .electricity, or by arsenic, which has metallic luster but 
is very brittle. Most of our schemes of classification are arbitrary, for 
our own convenience, and are not due to sharp natural boundaries. 
Summary. The classification of materials that has been presented 
in this chapter may be summarized in the following table. It is not 
important to memorize the table, but it is important to understand the 
nature of the distinctions it indicates. 
Table II. 
KINDS OF MATTER. 
HETEROGENEOUS: 
Consisting of more than one phase. 
HOMOGENEOUS: 
Consisting of one phase. 
SOLUTIONS: 
Variable Composition. 
PURE SUBTANCES. 
COMPOUNDS. 
ELEMENTS. 
Metals. 
Non-metals. 
Exercises. 
1. What is the distinction between the terms "object" and "sub- 
stance" ? 
2. Name a number of properties which serve to identify and 
distinguish substances. 
3. How could you distinguish between nickel and silver? Gold and 
brass? Alcohol and water? Glass and rock crystal (quartz)? Starch 
and talc powder? Salt and baking soda? Solid and plated silverware? 
Glass and diamond? Ether and chloroform? 
4. Classify the following materials according to the above table, 
giving reasons; where you are not sure of the classification devise an 
experiment which would enable you to decide:-milk, bronze, illuminating- 
gas, glass, wheat flour, a five-cent piece, well-water, gasoline, cast iron, 
plaster of Paris, white lead, sugar, gold, sulfur. 
5. Define, or at least use correctly in a sentence each of the following 
terms: phase, element, symbol, metal, substance, solution, property, com- 
pound, density, heat of fusion. 
14 CHAPTER II. 
While discussing the general properties of substances it will be 
profitable to consider certain properties of gases, on account of the 
existence of important uniformities in their behavior. When solids and 
liquids are heated, in nearly all instances, they expand, but the rate of 
expansion with the temperature is a specific property of each substance. 
The same holds true for the compressibility of solids and liquids. How- 
ever, when we investigate the behavior of gases with respect to alterations 
in temperature and pressure, we find that the effect on the volume is 
almost the same no matter what gas is being examined. One does not 
need to know the name of the gas in order to predict how its volume 
will change under the influence of changing pressure or temperature. 
In making such predictions we may have in mind either certain formulas 
which express the results of experiments upon gases, or else a theory 
as to the nature of gases which enables us to predict the behavior actually 
found by experiment. This twofold approach to scientific truth has been 
quite general, sometimes experiment leading to the discovery of new 
laws to be afterward explained by a theory, sometimes theory suggesting 
laws whose truth must be tested experimentally. In the present instance 
there are advantages in approaching the problem from the theoretical 
standpoint, since the theory is so simple and leads so obviously to the 
correct mathematics, making it unnecessary to commit to memory formulas 
almost sure to be forgotten or misapplied. 
Kinetic Theory. The fundamental difference between gases on the 
one hand, and liquids and solids on the other, is the tendency of gases 
to expand indefinitely, so as to fill completely any vessel in which they 
are contained. If a quantity of gas is introduced into an evacuated vessel, 
it distributes itself almost instantly throughout the whole space at its 
disposal, showing a tendency to still further expansion by exerting a 
pressure on all parts of the walls of the vessel. 
If the vessel already contains one gas, and another gas is admitted, 
it will be found, after allowing sufficient time to elapse, that both gases 
are equally distributed throughout the vessel. For example, if ammonia 
is exposed to the air in one part of a room in which the air is perfectly 
still, a person in a different part of the room will soon be able to smell 
the ammonia. This process, whereby one substance works its way 
through another, is called diffusion. 
Now, let us ask the question, what sort of a structure must we think 
of a gas as possessing in order to account for these phenomena? The 
most satisfactory picture is that furnished by the kinetic theory, which 
assumes that a gas is composed of particles called molecules, whose size 
is very small compared with the distance between them, and that, each 
molecule is in rapid motion, colliding with other molecules and with the 
walls of the containing vessel.* It is these impacts which produce the 
observed pressure, a series of rapid impacts obviously having the effect 
THE GAS LAWS AND THE KINETIC THEORY. 
* An idea of the actual values calculated may be gotten from the following 
figures for oxygen. At one atmosphere pressure and 0°C, one cubic centimeter 
of oxygen contains 28 X 1018 molecules, moving with an average velocity of 0.66 
kilometers per second (do about % mile per second). Each molecule travels on an 
average about 1000 times its own diameter before colliding with another molecule. 
15 of a steady pressure. The further assumption is made that the average 
velocity of the molecules, and hence the pressure produced upon the 
walls of the vessel, depends upon the temperature, increasing as the 
latter increases. The molecules are in such rapid motion, and are 
most of the time so far apart, relatively, that their mutual attractions 
have very little effect. 
This picture of the condition of gases has proven to be of such 
immense service in explaining and predicting their properties that it 
is hardly questioned by scientists at the present time. 
Relation Between Quantity and Pressure When Volume and Tem- 
perature Are Constant. Let us see first, how, on the basis of the theory 
just given, we should expect changes in the amount of a gas to be 
related numerically to changes in pressure. Imagine a gas enclosed in 
a vessel provided with a cock, as represented in Fig. 5, a. The vessel 
is supposed to be immersed in a bath to preserve constant temperature. 
Fig 5 
Let us have a very few molecules in the vessel, and visualize them by 
the dots in the figure. These molecules will all be in rapid rectilinear 
motion, some moving faster, some more slowly, between impacts, but 
with a certain average velocity depending on the temperature, producing 
a definite pressure on the vessel walls as they rebound from them. 
Suppose, now, we introduce an equal number of the same kind of mole- 
cules through the valve, producing the condition illustrated in Fig. 5, b. 
It is obvious that since the volume of the vessel remains the same, and 
the average speed of the molecules is kept the same by the action of the 
bath in maintaining constant temperature, the only effect of the doubling 
of the amount of gas is merely to double the number of impacts in a given 
time per unit of area of the vessel walls, and hence to double the 
pressure. Evidently we would expect any change in the number of 
molecules in the vessel to produce a proportional change in the pressure, 
and since the number of molecules depends on the amount of the gas, 
we can make the general deduction that when temperature and volume 
are kept constant the pressure of a gas is proportional to the amount 
present. 
16 As an example of how this may be applied let us consider a tank, 
with an attached pressure gauge, containing any gas, and let gas escape 
until the pressure falls from 100 lbs. per square inch to 30 lbs. It is 
evident, since the pressure has fallen to 0.3 of its initial value, that 
there remains in the tank only 0.3 of the gas originally present. 
Relation Between Pressure and Volume When Temperature and 
Amount Are Constant. Let us imagine, next, that the gas is contained 
in a cylinder like that of a steam engine, with a movable piston, as in 
Fig. 6, so that the volume of the enclosed gas may be altered without 
changing the amount. If, now, the piston is pushed down from the posi- 
tion represented by a, in the figure, to that represented by b, where the 
volume is half as great, the number of impacts on a given area of the 
walls in a given time will be doubled. While the piston is being pushed 
down the molecules striking it will rebound a little faster than before, 
causing a slight increase in temperature,* so that we must wait a moment 
Fig. 6. 
till this temperature has been lowered to that of the bath surrounding the 
cylinder, when the average speed of the molecules will be the same as 
before the piston was pushed down. Since the number of impacts on 
a given area of the walls has been doubled, the pressure is likewise twice 
what it was before the piston was moved. Evidently if the volume were 
made V3 as great as at first the pressure would become three times as 
great, etc., so that we may conclude in general that when the temperature 
and amount of gas are constant the pressure is inversely proportional to 
the volume. 
As an example let us calculate the final volume when 10 liters of gas 
at a pressume of 76 cm. of mercury is subjected to a pressure of 19 cm. 
Since the pressure is decreased in the ratio 19/76, the volume will be 
* This effect may be noticed in the heating of the pump when a bicycle or 
automobile tire is pumped up. 
17 increased in the same ratio, becoming 76/19 of 10 liters, the original volume, 
or 40 liters. 
Relation Between Pressure and Concentration When Temperature Is 
Constant. The ratio between the amount of material and the volume in 
which it is contained, or the amount in unit volume, is called its con- 
centration. Evidently it is the concentration which determines the number 
of impacts, and hence the pressure so long as the speed of the molecules 
is unaltered. If, for example, 2 gm. of a certain gas in 5 liters exerted a 
pressure of 2 atmospheres, then 1 gm. of the same gas at the same tem- 
perature in a volume of 10 1. would be only */4 as ocncentrated, and 
would exert only x/4 of the pressure, which would be 1/2 atmosphere. In 
general we may say that the pressure of a gas is proportioned to its 
concentration when the temperature is constant. 
Relation Between the Pressure, Mass and Speed of the Molecules. 
Although, as has been said, the pressure exerted by a gas depends upon 
the velocity of its molecules, it is not directly proportional to the speed, 
for if the speed is doubled, not only does each molecule hit the vessel 
walls twice as often, but also twice as hard, for the momentum of each 
molecule is doubled by doubling the speed. The pressure is thus propor- 
tional to the square of the average velocity of the molecules, and hence is 
proportional to their mean kinetic energy. Moreover, the momentum of 
each impact is proportional to the mass of the molecules. A heavy 
molecule would exert more pressure than would a light one moving at 
the same speed. In a mixture of gases the different kinds of molecules 
must all have the same kinetic energy, hence light molecules like 
those of hydrogen must move much faster than heavier ones like those 
of oxygen. 
Fig. 7. 
Fig. 8. 
Effect of Temperature, Amount Constant. Let us next see how a gas 
would be affected by a change in the temperature. It will be simplest, 
first, to consider the effect of this change upon the pressure and volume 
separately, keeping one constant while the other is allowed to vary. We 
will imagine a constant weight on the piston of the cylinder containing 
the gas, so that the volume may vary while the pressure remains the 
same. We may now vary the temperature by altering that of the bath 
in which the cylinder is immersed. If the temperature of the gas is 
increased we would expect, in terms of our theory, that the molecules 
would gain in kinetic energy, moving faster and hitting the vessel walls 
and the piston both harder and more frequently, forcing the latter up- 
wards until the reduction in the number of impacts in a given time 
compensates for the greater force of each impact. It is found by 
experiment that this increase in volume is uniform, as expressed by the 
plot in Fig. 7. 
18 Absolute Zero of Temperature. Similarly, if the piston is not permitted 
to move, so that the volume remains constant, an increase in tempera- 
ture would be expected to increase the kinetic energy of the molecules 
resulting in harder and more frequent impacts, and hence an increase in 
pressure. It is found that the increase in pressure with temperature is 
uniform, as expressed by the plot in Fig. 8, similar to the previous case. 
By careful measurements the increase is found to amount to 1/273 of the 
value of the pressure at 0°C. We are led naturally to inquire the effect 
of a continued decrease in temperature. Diminishing the pressure 1/273 
of its value at 0°C for every degree the temperature is lowered, would 
give no pressure at all at -273°C, (more exactly, -273.09°C). 
What does this mean in terms of the kinetic theory? Since we have 
attributed gas pressure to the impacts of gas molecules moving with an 
energy depending on the temperature, when the pressure becomes zero we 
must conclude that the molecules are no longer in motion, and that we 
have reached the Absolute Zero of temperature. A lower temperature 
is inconceivable in terms of our theory. This remarkable conclusion is 
confirmed by the behavior of other properties, many of w'hich approach 
either zero or infinity, as the temperature approaches -273°C. 
It is both logical, therefore, and for many scientific purposes con- 
venient, to reckon temperature from the absolute zero, which is 273° below 
the centigrade zero. We denote absolute temperature by A. On this 
scale the melting point of ice, 0°C, is 4-2730A; 17°C is 273-{-17 or 
290°A, and in general, letting T and t stand for temperature on the 
absolute and centigrade scales respectively, 
T = t + 273. 
The most convenient way of expressing the effect of temperature on the 
pressure is to say that when the amount and volume of a gas are kept 
constant, the pressure is proportional to the absolute temperature, since 
both depend upon the kinetic energy of the molecules. 
As an example let us find the final pressure produced when a tank 
of gas under a pressure of 10 atmospheres is cooled from 40°C to 18°C. 
40°C is 313°A, and 18°C is 291 °A. Since the absolute temperature has 
been lowered to 291/313 of its initial value, the kinetic energy of the 
molecules will be lowered to 291/313 of its initial value, and hence the 
pressure will be lowered to the same fraction of its initial value, which 
is 2^1/si3 10 atmospheres or 9.3 atmospheres. 
When the amount of gas and the pressure are kept constant, 
any change in the absolute temperature, with its proportionate change 
in the kinetic energy of the molecules, must be compensated by a pro- 
portionate change in volume. 
Suppose, for example, that we have 250 cc. of gas at 27°C, and wish 
to determine at what temperature the volume will become 200 cc. The 
new volume will be 200/250 or 0-8 of the old volume, hence the absloute 
temperature will be reduced to 0.8 of its initial value, or from 300°A, 
to 240°A, which is -33°C. 
Effect of Changing Both Temperature and Pressure, Amount of 
Gas Constant. When any two of the factors pressure, temperature and 
volume are changed, the effect on the third can be calculated by separating 
the process into two steps similar to the above. For example, suppose 
we have 50.0 cc. of gas at 74 cm. pressure and 20° C, and wish to know 
the volume it would occupy under standard conditions, which are 1 
atmosphere (76 cm. of mercury) and 0°C. If we first change the 
pressure to the final value, since the pressure is increased, the volume 
19 will be diminished to 74/T6 of 50.0 cc. Next, keeping the pressure constant, 
let us change the temperature from 20°C, which is 293°A, to 0°C or 
273°A. This will cause a further diminution in volume in the ratio 
273/293> so that the result of both changes is 273/203X74/76X50.0 cc. or 
45.3" cc. 
The several gas laws which we have just deduced from the kinetic 
theory, and which are found to be in accord with experiment, may be 
expressed by algebraic equations, and may be combined into one funda- 
mental general equation. When one has to apply the gas laws frequently 
it is wise to use equations, as the mathematical expression of a law makes 
it is easier to use. As has been said "mathematics is a substitute for 
thought"; it enables one to represent briefly the result of thought, so 
that the process may be repeated in the future without repeating the 
reasoning. When, however, the formula is used without the previous 
thought process, as when it is copied out of a textbook, there is the 
danger that it will be misapplied or forgotten, and even when applied 
correctly the process is nearly devoid of educational value. Instead of 
depending on the memory to retain several algebraic equations we may 
use the kinetic theory, a very simple conception, and by the aid of easy 
logic solve any problem in the gas laws, as illustrated above. The only 
thing that needs to be remembered is that 0°C is 273°A. 
Tn drawing the various conclusions given above concerning the 
relations between pressure, volume, amount and temperature, it must 
be remembered that we have assumed that the gases are sufficiently 
expanded that the volume occupied by the molecules themselves is 
negligible compared with the volume of the vessel, also that they are 
so far apart and moving with such high velocity that their mutual 
attractions can be neglected. These assumptions correspond more nearly 
to the truth the higher the temperature and the lower the concentration. 
On the other hand, as the temperature is lowered and as the concen- 
tration is increased, increasing deviations are to be expected from the 
behavior previously deduced, becoming more and more marked until the 
attractive forces become sufficient to cause the molecules to condense to the 
liquid state. The conditions under which this condensation takes place 
obviously depend upon the kind of molecules involved. With the gas 
helium, at a pressure of one atmosphere, the temperature must be lowered 
to 4.5°A before liquefaction takes place. 
Partial Pressure of Different Gases in Mixtures. The kinetic 
theory also enables us to predict correctly certain properties of gaseous 
mixtures. In the first place, we would expect the gas laws to hold for 
mixtures of gases as well as for pure gases, as is indeed the case. In 
the second place, we would expect the molecules of one species to maintain 
the same average kinetic energy at the same temperature regardless of 
the presence of any other species of molecules, and therefore the part of 
the pressure which is due to the impacts on the vessel walls of one 
species, called the partial pressure of that gas, would be the same no 
matter what other gases are present. This may be expressed by saying 
that the total pressure in a mixture of gases equals the sum of the par- 
tial pressures. By way of illustration, let us suppose a closed vessel 
containing water and carbon dioxide gas maintained at a constant 
temperature. Some of the carbon dioxide will dissolve in the water. 
Suppose now that some other gas, say nitrogen, is injected into the same 
vessel. How would this affect the amount of carbon dioxide dissolved in 
the water? Without the aid of the kinetic theory one might suppose 
20 that more carbon dioxide would be forced into the water, but from the 
molecular-kinetic standpoint we see that the number and momentum of 
the carbon dioxide molecules striking the water surface, upon which 
alone the solubility of the gas depends, is practically unaltered by the 
presence of the nitrogen molecules. 
Again, suppose we consider two vessels of equal size, one evacuated 
and the other containing oxygen, both kept at the same temperature. If 
water is introduced into the first one a little of it will vaporize, since the 
molecules of the liquid are in rapid motion with a velocity depending 
on the temperature, and certain molecules at the surface which have 
unusually high velocities may be able to escape the attraction of the 
liquid and go into the vapor phase. As the number of molecules in the 
vapor phase increases, there is a constantly increasing chance that some 
of them, moving more slowly than usual, will be caught into the liquid 
again, instead of rebounding at the surface. The concentration of mole- 
Fig.- 9. 
cules in the vapor state thus tends to become adjusted so that there is 
an equilibrium or balance between the two phases such that the number 
of molecules of liquid vaporizing in a given time just equals the number 
of molecules of water condensing in the same time. The pressure of 
vapor necessary for this equilibrium depends on the nature of the liquid 
and on the temperature, increasing as indicated by the curve in Fig. 9. 
It is obvious that this pressure will not depend on the extent of the liquid 
surface, as an increase in surface has the same effect on the number of 
molecules vaporizing as it has on the number condensing, so that the 
net effect is zero. 
Suppose, now, we admit water to the second vessel, containing 
21 oxygen. Evidently the tendency of molecules of liquid to escape will 
be the same as before, but the vaporized molecules will be unable to 
distribute themselves so rapidly throughout the vessel since they must 
diffuse through oxygen molecules which collide with them and retard their 
progress. Hence the liquid will evaporate more slowly than in the pre- 
vious case, where oxygen was absent, but eventually the molecules of 
water vapor will be distributed throughout the vessel just as if the oxygen 
were absent, and the partial pressure of the water vapor will be inde- 
pendent of the other gas (or gases) present. 
Avogardo's Hypothesis. One more consequence of the kinetic 
theory should be given at this time, as it will play a very important part 
in the reasoning in Chapter III. 
When two different gases are at the same temperature the average 
kinetic energy of the two different kinds of molecules is the same. We 
conclude that this is true from the fact that when two different gases 
at the same temperature are mixed the total resulting pressure is the sum 
of the partial pressures each gas would have if the other were absent, 
hence the kinetic energy of neither is altered by the mixing, which can 
be the case only when they have the same kinetic energy before mixing. 
Let us then take, in two vessels of equal volume, such amounts of two 
different gases at the same temperature that their pressures will be the 
same. Now we have seen that the pressure of a gas depends upon the 
number of molecules, their kinetic energy and the volume they occupy, 
and upon no other factors. Since, therefore, we have chosen equal 
pressures, volumes and temperatures (and hence kinetic energies) the 
only other factor, the number of molecules in the two quantities of gas 
must likewise be the same. In other words, equal volumes of all gases, at 
the same temperature and pressure, contain the same number of molecules. 
This was first announced by an Italian physicist, Avogardo, in 1811, and 
is usually called Avogardro's Rule or Hypothesis. We will use it pres- 
ently to measure the relative numbers of molecules involved in chemical 
reactions. 
Exercises. 
1. Does the height of the mercury column of a barometer depend 
on the cross-section? Explain. 
2. If a tank contains gas at pressure of 100 lbs. per sq. in, and 
gas is allowed to escape till the pressure becomes 40 lbs. per sq. in., 
the temperature being kept the same, what proportion of the original 
amount of gas remains in the tank? 
3. A quantity of gas having a volume of 150 cc. and a pressure of 
800 mm. is expanded, the temperature being kept the same, to a volume of 
750 cc. What is the final pressure? Ans. 160 mm. 
4. What volume of oxygen, measured at 20° and 1 atmosphere 
would be required to fill a 10 liter tank at the same temperature to a 
pressure of 25 atmospheres? Ans. 250 1. 
5. A flask is heated to a temperature of 100° while open to the 
air, the barometric pressure being 755 mm. It is then closed and allowed 
to cool at 27°; what is the final pressure? Ans. 607 mm. 
6. A quantity of gas has a volume of 60cc. at 27°C. At what 
temperature will its volume become 50 cc? Ans. -23°C. 
7. A quantity of hydrogen measured 24.8 cc. at 18°C and 765 mm.. 
22 what volume would it occupy under standard conditions? (1 atm. and 
0°C.) Ans. 23.4 cc. 
8. 1 liter of carbon dioxide weighs 1.96 grams under standard 
conditions. What weight of the gas is contained at 18°C, in a tank 
having a capacity of 20 liters, the pressure being 5 atmospheres? 
Ans. 184 gm. 
9. If a 2 liter vessel containing oxygen at a pressure of 1 atmos. 
is connected with a 5 liter vessel containing nitrogen at a preassure of 
4 atmos., what will be the partial pressure of each gas after diffusion 
has taken place? What will be the total pressure? 
Ans. Oxygen, 0.286 atm.; nitrogen, 2.857 atm.; total, 3.143 atm. 
10. If a 5 liter vessel contains 12 grams of a gas at 2 atmospheres 
pressure, what will be the concentration if more of the same gas is 
pumped in till the pressure becomes 3 atmospheres, the temperature 
remaining the same? Ans. 3.6 gm. per 1. 
11. Formulas may be derived for the gas laws by carrying out the 
reasoning as given, for example, on p. 8, but using the general symbols 
instead of numbers. Derive the relation between pressure, volume and 
temperature in this way using p, v, and T for the initial values, respect- 
ively, and p', v', and T' for the final values. Derive a similar relation 
between the pressure temperature and amount of gas (w) when the 
volume is constant. 
23 CHAPTER III. 
WEIGHT RELATIONS IN CHEMICAL REACTIONS. ATOMIC THEORY. 
An examination of the amounts of material involved is fundamental 
to an understanding of the chemical reactions. The practical importance 
of this has been referred to already in Chapter I, but we shall soon see 
that further consideration of these relations gives rise to theories and 
laws which lie at the very foundation of the structure of modern chemical 
science. 
Law of the Conservation of Mass. Everyday experience teaches us 
that many changes in matter take place without any change in weight 
of the materials involved (weight being the usual measure of mass). Thus 
100 pounds of ice, on melting, give 100 pounds of water; 1 pound of 
salt dissolved in 10 pounds of water give 11 pounds of solution; the 
weight of slaked lime equals the weight of quicklime from which it is 
made plus the weight of the water added in slaking. It is true that 
there is a loss in weight when quicklime is made from limestone, but that 
is because a gas, carbon dioxide, is given off, and if this were caught, 
compressed into a cylinder as is often done, and weighed, its weight 
would be found exactly equal to the loss in weight when the limestone 
changed into quicklime. Similarly, when something burns it may seem 
to disappear, but that is only because gases are formed which may 
escape notice. If a candle is allowed to bum in a closed vessel no change 
in weight can be detected. 
As a result of a vast scientific experience we may assert with great 
confidence that in every chemical reaction the sum of the masses of the 
products is the same as that of the substances taken. This is known 
as the Law of the Conservation of Mass, or sometimes the Indestructi- 
bility of Matter. It has been tested with considerable accuracy in 
millions of cases, and with extreme accuracy in a number of them, 
so that there is hardly another scientific law in which we have such 
confidence. - 
Law of Simple Multiple Proportions.* When pure substances are 
analyzed there is frequently found more than one compound of the 
same elements, the difference being due to the different proportions in 
which the elements are united. For example, there are three oxides of 
lead: litharge, used in glazing pottery, red lead, used in certain paints, 
and lead dioxide, which gives the brown color to the heads of certain 
kinds of matches, and which forms the positive plate of a lead storage 
battery. The proportions of lead to oxygen in these three compounds 
is found by analysis to be as follows: 
Litharge 1 part lead to 0.0772 part oxygen. 
Red Lead, 1 part lead to 0.1029 (=4/3X0.0772) part oxygen. 
Lead Dioxide, 1 part lead to 0.1542 (=2X0.0772) part oxygen. 
It is very striking that the proportions of oxygen in the red lead and 
lead dioxide are simple multiples of that in litharge, 4/3 and 2, respectively. 
Again, in addition to water, we know of another compound of 
* It is customary to consider a "Law of Definite Proportions," but since the 
identification of the pure substances to which alone it can be applied involves 
constancy of composition, as explained in Chap. I, it should be evident that this 
"law" is simply the definition of a pure substance. 
24 hydrogen and oxygen, hydrogen peroxide, a solution of which in water 
is found in almost every medicine closet for use as an antiseptic, cleans- 
ing, or bleaching agent. This substance is unstable, and readily gives 
up half of the oxygen it contains, leaving water. The relative amounts of 
hydrogen and oxygen in water and in hydrogen peroxide are as follows: 
Water, 1 part hydrogen to 15.88 parts oxygen. 
Hydrogen peroxide, 1 part hydrogen to 31.76 parts oxygen. 
In proportion to the hydrogen there is just twice as much oxygen in the 
second compound as in the first. An almost endless number of illus- 
trations of this nature might be given, and we would find that whenever 
the same elements form more than one compound, if the weights are in 
the ratio a to b in one they will be in some simple multiple of this in 
another, such as a to 2b, a to 36, 2a to 36, 3a to 46. The same 
simplicity is apparent when more than two elements combine to form 
different compounds. This experience is summarized in the Law of 
Simple Multiple Proportions, which may be stated as follows: Whenever 
the same elements form more than one compound, their proportions in 
the different compounds are simple multiples of each other. 
The Atomic Theory. The existence of the law of simple multiple 
proportions suggests the question, what fundamental quality of matter 
is responsible for this behavior? Why is the proportion of an element in 
a second compound always twice, three times, etc., what it is in the first, 
never, for example, 1.376 times what it is in the first? This state of 
affairs could only well be true if each chemical element were composed 
of minute particles, which we call atoms, all alike for the same ele- 
ment, compounds being formed by simple combinations of different kinds 
of atoms. For example, an atom of element A, having weight a, might 
combine with one atom of element B, having a weight 6, to form a 
particle of a compound AB. It is obvious that no matter how many 
of these particles we take, the ratio of the weights of the elements in the 
compound would be a to 6. The same elements might also form a 
different compound by a different combination of atoms, such as by 
one atom of A combining with two atoms of B, to form a particle which 
may be designated by AB2. Evidently the ratio of weights of the 
elements A and B in a particle of this compound, and hence in any 
number of particles, would be a to 26. There is here twice as much 
of element B with respect to A in the second compound, AB,, as in 
the first, AB, in accordance with the experience summarized in the law 
of simple multiple proportions. If still other compounds of A with B 
are formed we would expect other simple combinations of atoms, like 
A,B, AB3, A,B3, A3B2, etc. Accordingly, if the ratio of the weights of 
A and B in one compound is a/b, we would expect to find in another 
compound some such ratio as 2a/b, a/3b, 2a/3b, ?>a/2b. 
There are two ways in which we may imagine matter to be con- 
stituted : either it is continuous, as it appears to the eye, so that we 
might go on subdividing it indefinitely and never reach any portion 
however small that could not be still further subdivided without losing 
its identity, or else matter may have a granular structure, so that when 
we have subdivided it until these ultimate grains are reached no further 
subdivision is possible, without fundamentally altering the nature of 
the material. That the latter assumption corresponds to the facts is 
strongly indicated by the success of the kinetic theory in explaining the 
physical properties of gases, and now we find that a study of chemical 
changes points to the same conclusion. If we add to this the evidence 
25 furnished by radioactive phenomena and colloids, to be discussed in later 
chapters, we may well feel that the atomic structure of matter is more 
than a mere theory, and rather an established fact. 
Atoms and Molecules. Formulas. We may now consider a little 
more closely the molecules implied by the kinetic theory and regard them 
as little groups of atoms, whose kind, number and arrangement in the 
molecule determine the nture of the substance. Thus water, which we 
can easily prove to be composed of the elements of hydrogen and oxygen 
in the approximate weight ratio 1 to 8, consists of molecules, each com- 
posed of atoms of hydrogen and oxygen. We will now extend the 
meaning of the symbols H and O to signify not only the elements of 
hydrogen and oxygen respectively, as in Chap. I, but also one atom 
of each element. We cannot be sure without the aid of the subject treated 
in Chapter IV just how many atoms of each kind there are in a molecule 
of water, and we may make several assumptions. There may be one atom 
of each element in a molecule of water, indicated by the combination of 
the symbols of the elements H and O into the formula of the molecule 
of the compound HO. If this represents the molecule of water, then we 
piust represent that of hydrogen peroxide, which contains twice as much 
oxygen in proportion to the hydrogen, by the formula HO2, and the 
atom of oxygen must weigh 8 times as much as the atom of hydrogen. 
Again, the formula of water may be H.,O, indicating two atoms of 
hydrogen and one of oxygen in the molecule, in which case the molecule 
of hydrogen peroxide would have to be HO, or H2O2, etc., the atom 
of oxygen being 16 times as heavy as that of hydrogen, in order to account 
for the observed composition of these two substances. Still other simple 
assumptions with corresponding relative atomic weights might be made. 
As set forth in the following chapter we have satisfactory evidence that 
the second of the above assumptions, H,O, is the correct one, and that 
the relative weights of the hydrogen and oxygen atoms are as 1 to 16. 
We have reason also to believe that the molecule of litharge contains 1 
atom, each of lead and oxygen, as represented by the formula PbO, the 
weights being in the ratio of 207 to 16, corresponding to the ratio 1 to 
0.0772 given on p. 2. Red lead, where we have 3 times this amount of 
lead and 4 times this amount of oxygen, would accordingly be represented 
by the formula Pb3O4, and lead dioxide, where the proportion of oxygen 
to lead is twice that in litharge, would correspond to the formula PbO.,. 
Numerical Values of Atomic Weights. The actual weights of the 
atoms are exceedingly small, compressed in our ordinary units, such as 
grams,* and it is far more convenient to select a much smaller weight as 
unit, such as the weight of one of the atoms themselves. It is purely arbi- 
trary what atom we choose and what number we assign to it, so long as we 
select corresponding values for the other atoms. Thus, knowing that the 
molecule of water contains 8 times as much oxygen as hydrogen, and 
granting that H2O is the correct formula for it, we might call O = 100, 
when H must be 6.25; or again, if we let O-l, then H = 0.0625; etc. 
The choice appeared a very simple one to most early chemists, who, finding 
that the atoms of hydrogen are lighter than those of any other known 
element, decided to call its weight 1, making O = 16, and making the 
weights of the atoms of other substances to be expressed by numbers 
greater than 1. More recently it has been found that the ratio of weights 
of the atoms of hydrogen and oxygen is not exactly 1 to 16, but 1 to 
* It would take approximately 3 X 1023 atoms of hydrogen to weigh 1 gram. 
26 15.88, and that if we take O = 16 as our unit, making H = 1.0076, the 
atomic weights of most of the other elements come out much nearer 
whole numbers than if we let H = 1 and O = 15.88. This is illustrated 
by the following sets of values based on the two assumptions. 
Hydrogen 1.0076 1.000 
Oxygen 16.00 15.88 
Calcium 40.07 39.77 
Carbon 12.01 ' 11.91 
Lithium 6.94 6.89 
Magnesium 24.32 24.14 
Nitrogen 14.01 13.90 
Phosphorus 31.04 30.81 
Potassium 39.10 38.80 
Sodium 23.00 22.83 
Sulfur 32.06 31.76 
A still greater reason for the choice of O = 16.00 lies in the fact 
that but few of the elements form compounds with hydrogen, while 
nearly all form compounds with oxygen, and hence the atomic weight 
ratio between an element and oxygen can be determined directly, while 
the atomic weight ratio between the element and hydrogen could only be 
indirectly determined, and would therefore be subject to more frequent 
change. 
Interpretation of Formulas. M are now prepared to interpret the 
notation used by chemists to indicate the composition of pure substances, 
though we have not yet shown that actual compounds have the composi- 
tion assigned to them. We find the atomic weights given with the names 
and symbols of the elements on p. 13, indicating the weight of each 
atom in terms of a unit which is 1/16 the weight of an oxygen atom. 
The smallest part of any compound, the molecule, is thus made up of 
atoms of its component elements, with corresponding weights. This 
molecule is indicated by a formula showing the number and kind of 
atoms of which it is composed. Thus the formula Fe2O3 indicates a 
molecule of a certain oxide of iron containing 2 atoms of the element 
iron, which constitute 2 X 56 parts by weight, together with 3 atoms of 
the element oxygen, which constitute 3 X 16 parts by weight. Another 
formula, Fe3O4, represents another oxide of iron, containing 3 atoms 
of iron to 4 atoms of oxygen, the weights being in the ratio 3 X 56 to 
4 X 16. Na2CO3 denotes 1 molecule of sodium carbonate or "washing 
soda," containing 2 atoms or 2 X 23 or 46 parts of sodium, 1 atom or 
12 parts of carbon, and 3 atoms or 3 X 16 or 48 parts of oxygen, the 
whole molecule representing 106 parts. 
Gram-Atoms and Gram-Molecules, or Mois. Although we reason 
in terms of atoms and molecules, we work with vastly larger amounts of 
material, and it is often convenient to give a slightly different meaning 
to our symbols and formulas so as to express parts by weight in terms 
of units actually used. Accordingly the symbol Fe is often used to 
denote not a single atom of iron, but 56 grams of iron, called a gram-atom 
of iron. Likewise the formula Fe2O3 denotes not only a single molecule 
of an oxide of iron, but a gram-molecule, or mol of the oxide, weighing 
150 grams. Since one molecule of the oxide contains 112 parts of iron 
to 48 parts of oxygen, then no matter how many molecules of the oxide 
we take, the proportion of iron to oxygen is still 112 to 48, hence the 
same weight relations that apply to atoms and molecules apply to gram- 
27 atoms and gram molecules. The adjective moled is used in referring to 
the properties of the mol. Thus we would refer to the molal weight 
of Fe2O3 as being 150 grams. 
This numerical significance of chemical formulas aids us in solving 
problems involving the relative weights of compounds and their com- 
ponent elements. Suppose that we wished to calculate the amount of 
litharge that might be made from 1 kilogram (1000 grams) of lead. 
Since one atom of lead, representing 207 parts by weight, combines with 
one atom of oxygen, representing 16 parts by weight to give a molecule 
of litharge, PbO, representing 223 parts by weight we can say that 
207 parts of lead will give 223 parts of litharge, 
and accordingly 
207 grams of lead will give 223 grams of litharge, 
and 1 gram of lead will give 1.0772 grams of lithargs, 
and 1000 grams of lead will give 1000 X 1.0772 grams of litharge, or 
1077.2 grams. 
Chemical Equations. When a chemical reaction takes place there 
is a rearrangement of atoms into new groups, or molecules, forming new' 
substances. This is indicated by a chemical equation, in which the 
rearrangement is showm by the regrouping of the atomic symbols to 
represent the new molecular species formed. For example, the forma- 
tion of iron and carbon monoxide, CO, from the ore Fe2O3 by the action 
of carbon is represented by the equation, 
Fe2O3 -f- 3C - 2Fe 3CO. 
The following things may be noted about such an equation: 
First, an equation tells nothing about the conditions necessary for 
carrying out the reaction, and the writing of an equation does not imply 
that the reaction it represents w'ill actually be realized. 
Second, it must represent the fact that no matter disappears, by 
showing the same number of atoms of each element on both sides of the 
equation, though they are arranged in different molecules. 
Third, it represents the relative number of molecules of each sub- 
stance taking part in the reaction. 
Fourth, it represents the relative weights of each substance, with the 
aid of the table of atomic weights. 
The above equation, then, may be read as follows: 1 molecule of 
ferric oxide reacts with 3 molecules (or atoms) of carbon to yield 2 
molecules (or atoms) of iron and 3 molecules of carbon dioxide; also, 
2 X 56 3 X 16 parts (by weight) of iron oxide react with 3 X 12 
parts of carbon, to give 2X56 parts of iron and 3(12 -|- 16) parts of 
carbon monoxide. These parts by w'eight may be in any kind of weight 
unit, pounds, tons, ounces, grams, etc. In accordance with our double 
usage of the symbols and formulas to represent both single atoms and 
molecule, and also gram-atoms and gram-molecules, or moles, as ex- 
plained on p. 8, we often read an equation in terms of grams; hence, 160 
gm. of iron oxide when reacting with 36 gm. of carbon give 112 gm. of 
iron and 84 gm. of carbon monoxide, which quantities represent 1 mol, 
3 mols, 2 mols, and 3 mols respectively. 
This may be summarized briefly by w'riting under each formula in 
the equation its complete meaning in both of the above senses. 
Calculation of Weight Relations. The meaning thus attached to 
chemical formulas and equations makes it possible to calculate the 
28 Fe2O3 + 3C = 2Fe + 3CO 
(1) Number of molecules 13 2 3 
Weight, in terms of 1/16 the weight of 
an oxygen atom 160 36 112 84 
(2) Number of mols 13 2 3 
Weight 160 gm. 36 gm. 112 gm. 84 gm. 
weights of all the other substances involved in a chemical reaction from 
the weight given for any one of them. Suppose, for example, that it 
is required to find the weight of iron that should be obtained from 100 
pounds of the oxide of iron considered above. From the weights implied 
by the formulas, as explained above, we may say, 
160 parts (pounds) of Fe2O3 will give 112 parts (pounds) of iron, 
hence 1 part (pound) of Fe2O3 will give 112/160 parts (pounds) of iron, 
and 100 parts (pounds) of Fe2O3 will give 100xll2/160 parts (pounds) of 
iron or 70 parts (pounds) of iron. 
Again, suppose we wish to find the amount of coke, containing 80% 
carbon, the rest being ash, that is required to reduce to iron the above 
amount of ore. From the weights implied in the equation, we can say 
that 
160 pounds of oxide will require 36 pounds of carbon, hence 1 
pound of oxide will require 36/160 pounds of carbon, and 100 pounds 
of oxide will require 100 X 36/160, or 22.5 pounds of carbon. Since, 
further, this amount of carbon is 80% of the coke required, the amount 
of the latter is 100/80 of 22.5 pounds or 28.2 pounds. 
This connection between atomic weights, formulas and relative 
weights of substances may be used not only to calculate the relative 
weights from the formulas and atomic weights, as above, but also to 
calculate the atomic weights when the formulas and suitable weight 
relations are known, or again, to determine formulas when the other 
two factors are known. 
Calculation of Atomic Weights. As an example of the calculation 
of atomic weights let us suppose that we know the formula of a certain 
oxide of copper to be Cu2O, and find on analysis that 0.5120 gm. of it 
contain 0.4548 gm. of copper, and wish to calculate the atomic weight of 
copper. The formula tells us that 2 atoms of copper are combined with 
1 atom of oxygen, hence there would be 2 gram-atoms of copper to 1 
gram-atom of oxygen. Since there are, according to the analysis, 0.4548 
gm. of copper to 0.0572 gm. of oxygen, there would be 16xo.4548/o gm. 
or 127.2 gm. of copper to 1 gram-atom (16. gm.) of oxygen. Since 172.2 
gm. of copper is 2 gram-atoms, 1 gram-atom is 63.6 grams. Hence the 
atomic weight of copper is 63.6. 
Calculation of Formula. The third case consists in calculating the 
formula when the other two factors, atomic weight and composition, are 
known. Suppose that we find on analysis that a certain chloride of 
arsenic contains 58.68% of chlorine, and know the atomic weight of 
arsenic, As = 75.0, and chlorine, Cl = 35.5, and wish to determine the 
formula of the compound. Now the formula indicates primarily the 
relative number of atoms, or of gram-atoms in the molecule, or mol, 
respectively, so that we may proceed to determine the number of gram- 
atoms of one element combined with 1 gram-atom of the other. Since 
58.68 gm. of chlorine would combine with 41.32 gm. of arsenic, 35.5 gm. 
of chlorine, 1 gram-atom would combine with 35,5x41-32/58.68 grams, or 
29 25.0 grams of arsenic, which is V3 gram-atom of arsenic. We cannot 
have 1/3 atom of arsenic, so that we conclude, from the fact that 1 gram- 
atom of chlorine would combine with 1/3 gram-atom of arsenic, that 3 
atoms of chlorine would combine with 1 atom of arsenic, hence the 
formula as AsC13 or some multiple of it, such as As2C18, As3C19, etc. 
To distinguish between these multiples another kind of data is necessary. 
This will be considered in Chapter IV. 
Exercises. 
1. What do you understand by the law of simple multiple propor- 
tions ? 
2. How does the atomic theory explain this law? 
3. Write out in words all that the following formlas mean to you: 
H2O; H2O2; Ca(OH)2; FeSO4. 
4. Define symbol; formula; chemical reaction; chemical equation; 
mol; molecule; atom. 
5. How many pounds of iron are there in 100 pounds of magnetic 
iron ore, Fe3O4? Ans. 72.4 lbs. 
6. How many grams of sodium peroxide (Na2O2) can be obtained, 
from 115 gm. of sodium? Ans. 195 gm. 
7. When treated with hydrochloric acid, barium peroxide yields 
hydrogen peroxide according to the following reaction: 
BaO2 4- 2HC1 = H2O2 + BaCl2. 
How much BaO, is required to make 10 lbs. of a 5% solution of H.,Oo? 
Ans. 21//lbs. 
8. On heating 58 gm. of silver oxide, Ag,O, it was decomposed 
into oxygen gas and 54 gm. of metallic silver; what is the atomic weight 
of silver, if that of oxygen is 16? Ans. 108. 
9. How many grams are there in 1 mol of each of the following 
substances: Ca (OH)2; HC1; NaCl? Ans. 74 gm.; 36.5 gm.; 58.5 gm. 
10. 10 gm. of limestone, CaCO3, were put into a closed vessel and 
heated to such a temperature that 2% of it decomposed according to 
the equation: CaCO3 = CaO 4~ CO2f How many grams of CO2 were 
produced? Ans. 0.088 gm. 
11. A certain oxide of nitrogen was found by analysis to contain 
74.1% of oxygen. Work out its simplest formula. 
30 CHAPTER IV. 
VOLUME OF GASES. MOLECULAR WEIGHTS. 
Choice of Formulas and Corresponding Atomic Weights. In the 
last chapter we showed how the Atomic Theory, deduced from the Law 
of Conservation of Mass and the Law of Simple Multiple Proportions 
by weight, could be expressed in a simple and comprehensive notation by 
the aid of which any one of the three factors, relative weights, formula, 
and atomic weights could be calculated from the other two. Now the 
relative weights may be found by experiment, but in order to get either 
of the other factors one of them must be known. Although we practiced 
with a number of formulas to become familiar with the meaning of the 
notation used, we did not show that the molecules of the substances 
referred to actually contained the numbers of atoms of the constituent 
elements signified by the formulas used, e. g., FLO, Fe2O3, AsC13. 
To illustrate the problem we now face let us consider the formula 
of water. We know thus far only that the ratio of the weight of hydro- 
gen to that of oxygen in water is approximately 1 to 8. This we learn 
by the analysis of water. If, now, the formula of water is HO, the 
relative weights of the atoms of hydrogen and oxygen are 1 to 8. But 
if the molecule of water is to be represented by H2O, the relative weights 
of the single atoms would have to be 1 to 16. Again, if the formula 
is HO2, then the relative atomic weights must be 1 to 4. Evidently 
we must bring to bear some criterion other than those thus far considered, 
in order to select the correct formula and atomic weight ratio. 
If we try to represent the formation of water from its elements on 
the basis of the various assumptions as to its formula, we should have, 
as probable guesses, equations such as the following: 
H+ O = HO 
2H + O = H,O 
H 4- 20 = HO2, etc. 
By suitable choice of atomic weights any of these, or others, as just 
explained, would conform to the experimental weight ratio. The essential 
point of difference lies in the number of molecules of the three substances. 
Now we have, from the kinetic theory, a means of measuring the relative 
numbers of molecules for gaseous substances. It was shown in Chapter 
II that, by Avogadro's Rule, at equal temperatures and pressures, equal 
volumes of all gases contain the same number of molecules. Accordingly, 
if we determine the relative volumes of the gases taking part in a chemical 
reaction, the temperature and the pressure being the same before and 
after the experiment, we may conclude that the relative numbers of 
molecules are in the same proportion as the volumes. If the above reac- 
tion is carried out at a temperature high enough to keep the water in 
the vapor state it will be found that 2 volumes of hydrogen and 1 
volume of oxygen give two volumes of steam. Accordingly, we conclude 
that two molecules of hydrogen and one molecule of oxygen would give 
two molecules of steam. Our equation must now be written so as to 
express these numbers of molecules. Since the oxygen in one molecule is 
enough to give two molecules of steam, each containing oxygen, there 
must be more than one atom of oxygen in its molecule, hence its formula 
must be not O but O2, or O4, etc. However, as we consider other 
31 reactions, we never find that one molecule of oxygen gives more than 
two molecules of any other substances, so we assume the simplest formula, 
O,. This limits the above equation to the following: 
2Hn + O2 = 2HnO, 
where n is some whole number. 
We still have to determine the number of atoms in the molecule of 
hydrogen. This is done, as in the case of oxygen, by finding the greatest 
number of other molecules, each containing hydrogen, that can be gotten 
from 1 molecule of hydrogen. This number has never been found 
greater than 2, but often equal to 2, hence we write H,. The union of 
hydrogen and chlorine to form hydrogen chloride gas may be used for 
illustration. Experiment shows that 1 volume of hydrogen and 1 volume 
of chlorine give 2 volumes of hydrogen chloride, hence 1 molecule of 
hydrogen and 1 molecule of chlorine give 2 molecules of hydrogen 
chloride, and the equation would be, 
H2 + Cl2 = 2HC1. 
Using this formula for hydrogen in our previous example, we write the 
equation for the formation of water 
2H2 + O2 = 2H2O. 
Finally, knowing the relative weights by experiment, and having 
deduced the formulas by the aid of Avogadro's rule, we can assign values 
to the atomic weights. As explained before, one of these must be chosen 
arbitrarily, and we have seen how the number 16 has been assigned to 
the oxygen atom. Using this value, the weight of the oxygen molecule, 
O2 becomes 32. By analysis of water we find tht 32 parts of oxygen 
combine with 4.030 of hydrogen. The above equation shows that this 
represents 4 gram-atoms of hydrogen, hence 1 gram-atom of hydrogen 
is 1.008 gram, and this number represents the atomic weight of hydrogen. 
It is not long since the time when the simple way just explained 
for deciding upon the correct formula and atomic weight ratio w'as not 
understood. Many chemists as late as 1858 used HO for the formula 
of water, while others used H2O, and much confusion existed. Many 
criteria were used, other than that furnished by gas volumes, to determine 
the true formulas, but with only partial success. The table of atomic 
weights is thus the product of the thought and experimental work of a 
great many men. 
We have seen that the molecules of hydrogen, oxygen and chlorine 
each contain two atoms. The same is true for nitrogen, fluorine, bromine 
and iodine. 
Another variety of oxygen exists, called ozone, formed from oxygen 
by the silent electric discharge, with a diminution in volume in the ratio 
of 3 to 2, indicating that the equation must be written, 
3O2 = 2O3. 
The molecule of ozone, therefore, contains three oxygen atoms. The 
molecule of phosphorus vapor at low temperatures is P4, at higher tem- 
perature P2, that of sulfur may be S8, S6, S2 or S, according to the 
temperature. When metals are vaporized their molecules seem always 
to consist of single atoms. The same is true of argon, and certain 
other gases recently discovered existing in small amounts in the atmos- 
phere. 
* 
Changes in Volume or Pressure in Gas Reactions. The connection 
we have traced between the volumes of gases in chemical reactions and 
the relative number of molecules may be used conversely to deduce the 
change of volume or pressure to be expected when reactions take place be- 
32 tween substances whose formulas are known. Consider, for example, the 
reaction represented by the equation 
2SO2 + O2 = 2SO3. 
If this reaction is carried out at a temperature at which all of the sub- 
stances represented are gaseous, then we see that three molecules have 
formed two. If the temperature and volume have been kept the same 
during the reaction, this will result in a decrease of the pressure to 2/3 of 
its former value. If the temperature and pressure are the same as before, 
then the volume will be 2/3 of its former value. 
It is important, whenever an equation will be used to deduce relative 
volumes, to use the correct formulas of the substances in the gaseous 
form. For example, the same weight, 28 gm., is indicated both by N2 
and 2N, but the second formula is incorrect, and would lead us to expect 
twice the volume which would actually be involved. 
Volume of One Mol of Any Gas. The double meaning attached to 
chemical symbols and formulas was explained in Chapter II, one referring 
to single atoms and molecules, and furnishing a basis for reasoning, and 
the other referring to the gram-atom and gram-molecule or mol, and 
furnishing a basis for experimental work. Having determined that equal 
numbers of gaseous molecules are contained in equal volumes, at equal 
temperatures and pressures, and that the mol of all substances consists 
of the same number of molecules, it becomes important to ask what is 
the actual volume occupied by the mol of gas. This rests upon the 
measurement of gas densities, as illustrated by the following table: 
Wt. of 1 liter 
Vol. of 1 gm. 
Gm. per mol 
Vol. of 1 mol 
Gas 
O°C, 1 atm. 
O°C, 1 atm. 
O°C, 1 atm. 
h2 
0.08988 
11.126 1. 
2.016 
22.43 1. 
o2 
1.4291 
0.6998 1. 
32.00 
22.39 1. 
N2 
1.2507 
0.7995 1. 
28.02 
22.40 1. 
nh3 
0.7621 
1.3122 1. 
17.03 
22.35 1. 
These values for the volume of the mol of gas under standard 
conditions, as well as others that might be added, are all very close to 
22.4 liters, which may therefore be selected as the molal volume of a 
gas under standard conditions. (This is the volume of a cube 28.2 cm. 
or IP/s inches on the side). We may take this volume as the basis 
of our working definition of molecular weight, saying that the molecular 
weight of a substance is the number of grams of it which occupy 22.4 
liters when it is in the gaseous form and under standard conditions. Very 
frequently, of course, substances cannot exist in the vapor state under 
these conditions, so that the relation between the weight and the volume 
of vapor must be found^experimentally at some higher temperature, and 
perhaps lower pressure after which it is possible to calculate the weight 
that 22.4 liters of the vapor would have at 0°C and 1 atmosphere, if no 
condensation took place. 
We may illustrate by examples the various types of problems it is 
possible to solve with the aid of this relationship. 
1. Calculation of the Weight of a Given Volume or the Volume 
of a Given Weight of a Gas. Example: What is the volume of 10 gm. 
of O2 at 127°C and 0.5 atmospheres? The answer to this is obtained 
by the following obvious steps. 
32 gm. of O, occupy 22.4 1. at 273°A and 1 atm. 
32 gm. of O2 occupy 44.8 1. at 273°A and 0.5 atm. 
32 gm. of O2 occupy 400/273 X 44.8 1. at 400°A (127°C) and 0.5 atm. 
33 1 gm. of O2 occupies 1/32 X400/273 X 44.8 1., at 400°A and 0.5 atm. 
10 gm. of O, occupy 10/32 X 400/273 X 44.8 1. or 20.5 1. at 400°A and 
0.5 atm. 
Example: What is the weight of 50 1. of CO2 measured at 25°C and 
3 atm.? 
At 273° and 1 atm. 22.4 1. of CO, weigh 44 gm. 
At 273°A and 3 atm. 1/3 X 22.4 1. of CO, weigh 44 gm. 
At 298°A (25°C) and 3tm. 298/273 X 7S X 22.4 1. or 8.16 1. of CO, 
weigh 44 gm. 
At 298°A and 3 atm. 1 1. of CO, weighs ^/s.ie gm. 
At 298°A and 3 atm. 50 1. of CO2 weighs 50X 44/8>10 gm- or 270 gm. 
2. Determination of the Relative Weights of a Gas and Air. It is 
often important to know whether a gas is heavier or lighter than air. 
The proportion of nitrogen to oxygen in air is very nearly 4 to 1 by 
volume, so that in 22.4 liters of air V- of the molecules would be nitrogen, 
and 1/5 would be oxygen. Taking 4/- of 22.4 liters of nitrogen under 
standard conditions, weighing 4/- of 80 gm., and mixing it with Vs of 
22.4 liters of oxygen, weighing 1/5 of 32.0 gm., we would get 22.4 1. of 
are weighing 28.8 gm., under standard conditions. Suppose we wish to 
know the relative density of the following gases with respect to aid: 
CO2, NH3, HC1, Cl2. We can find their molecular weights with the aid 
of the atomic weight table, and conclude that 22.4 1. under standard 
conditions would have the following weights: CO2, 44 gm.; NH3, 17 gm.; 
HC1, 36.5 gm.; Cl,, 70 gm.; air, 28.8 gm. This shows at once the relative 
weight of each with respect to the air. 
3. Determination of Molecular Weight. Example: What is the 
molecular weight of phosphorus vapor and what is its formula, given 
the atomic weight, P = 31.0, and the experimental determination that a 
flask having a volume of 583 cc. was filled with the vapor at 310°C and a 
pressure of 756 mm., and that it was found on cooling to contain 1.49 
gm. of phosphorus? The molecular weight is found from the number of 
grams occupying 22.4 1. at 273°A and 760 mm., which must therefore 
be calculated. 
Since 1.49 gm. of phosphorus vapor would occupy 583 cc. at 756 mm. and 
583° A, 
1.49 gm. of phosphorus vapor would occupy 755/760 X 583 cc. at 
760 mm. and 583°A, 
1.49 gm. of phosphorus vapor would occupy 273/583 X 756/tgo X 583 
cc. at 760 mm. and 273°A, 
1.49 gm. of phosphorus vapor would occupy 0.2715 1. at 760 mm. 
and 273°A. 
22-4/0.27i5 X 1-49 gm. or 123 gm. of phosphorus vapor would occupy 
22.4 1. at 760 mm. and 273°A. 
Since 123 gm. would occupy 22.4 1. under standard conditions, if it 
did not condense, 123 is approximately the molecular weight. The atomic 
weight being 31. there are evidently 4 atoms in the molecule, so that 
the formula is P4. 
As another example we may determine the formula of a compound, 
found by analysis to contain 40% of carbon, 6.67% of hydrogen, and 
53.33% of oxygen; and of which 0.50 gm. gave 328 cc. of vapor at 
200°C and 750 mm. We will first find the relative number of atoms of 
each element in the molecule of the compound. From the per cents 
given we may say that 
34 40 gm. of carbon combine with 6.67 gm. of hydrogen and 53.33 gm. of 
oxygen, therefore 12 gm. of carbon combine with 12/40 of 6.67 gm. or 2 
gm. of hydrogen and 12/40 of 53.33 gm. or 16 gm. of oxygen. Since 12 
gm. of carbon, 2 gm. of hydrogen and 16 gm. of oxygen are 1 gram-atom of 
carbon, 2 gram-atoms of hydrogen, and 1 gram-atom of oxygen respect- 
ively, these represent the relative numbers of atoms of the three elements 
in the molecule. The formula might therefore be CH2O. However, the 
same proportions by weight would be found if the formula were C2H4O2 
or C3H0O3, or any other multiple of CH2O. In order to distinguish 
between them we ntust therefore determine the number of grams per mol, 
which is numerically equal to the molecular weight. 
Since 0.5 gm. occupy 328 cc. at 750 mm. and 473°A (200°C), 
0.5 gm. would occupy 75O/76O X 328 cc. at 760 mm. and 473° A, and 
0.5 gm. would occupy 273/473 X75O/7eo X 328 cc. at 760 mm. and 273°A, 
or 187 cc. 
Now the weight of gas contained in 22.4 liters, or 22400 cc. is 1 mol, and 
since 187 cc. weigh 0.5 gm., 1 cc. would weigh 0.0026/ gm. and 22400 
cc. would weigh 22400 X 0.00267 gm. or 60 gm. Of the various multiples 
of CH2O the one having a molal weight of 60 gm. is C2H4O2, hence 
this is the correct formula. 
4. Direct Relation Between Volumes of Gases and IVeights of Other 
Substances in Reactions. Example: When limestone is heated to a 
sufficiently high temperature it decomposes as represented by the following 
equation: 
CaCO3 = CaO + CO2. 
Suppose we wish to calculate how many liters of CO2, measured at 
20°C and 10 atmospheres pressure, should be obtained from 2 kilograms 
(2000 gm.) of CaCO3. It is not necessary to calculate the weight of CO2 
since the equation can be interpreted so as to give directly the relation 
between the weight of CaCO3 and the volume of CO2. From 1 mol of 
CaCO3 weighing 100 gm .there can be obtained 1 mol of CO2 occupying 
22.4 1. at 273°A and 1 atmosphere. Accordingly, 2000 gm. CaCO3 would 
give 448 1. CO2 at 273°A and 1 atm., or 293/273 X 448 1. at 293°A and 
1 atm., or 1/10 X 293/273 X 448 1. or 41.7 1. of CO2 at 293°A and 10 atm. 
Complete Interpretation of Chemical Equations. It is possible now 
to summarize what has been presented in this chapter and the preceding 
one concerning the interpretation of chemical equations, both as to 
weight and volume relations, bearing in mind that a formula always 
signifies a certain weight of some substance, and, when the substance 
exists in the gaseous state, a certain volume under standard conditions. 
We may recall that formulas are interpreted in two ways, one referring 
to atoms and molecules, the other to gram-atoms and gram-molecules, 
or mols. This double interpretation may be illustrated by using an 
equation such as the following, writing under each formula • its signifi- 
cance in both senses. 
CO2(gas) + C(solid) =2CO(gas) 
The last row of figures is added for the sake of interest, not because 
it will be used in problems. The other data represent all that is 
necessary in order to solve any problem involving weights, or volumes 
of gases, provided that one knows how to apply the gas laws to get the 
relation between the volumes of gases at standard conditions and at 
other conditions. It must be noted that this generalization connecting 
35 weights and volumes is possible only where gases are involved. With 
liquids and solids the densities of the particular substances must be 
determined by experiment. 
(1) Weights in terms of oxygen 
atom ==16 44 12 56 
Number of molecules 112 
Relative volumes . 1 0-|- 2 
(2) Weight in grams 44 12 56 
Number of mols 112 
12 
Volume at 0°C, 1 atm. 22.4 1. density cc 44 g j 
12 
Number of rrtolecules 6 X 1023 6 X 1023 12 X 1023 
Molecular Weights of Dissolved Substances. It is frequently impos- 
sible to determine the molecular weight of a substance from the volume 
of its vapor, owing to decomposition on heating. Sugar, for example, 
decomposes instead of vaporizing when heated, turning first to caramel 
and then charring. Some other property than the density in the vapor 
state is therefore necessary in order to determine its molecular weight 
and to decide which multiple of its empirical formula, C12H22O11 is the 
correct one. Hydrogen peroxide, likewise, is very unstable, except when 
in a dilute solution, decomposing often with explosive violence when in 
the pure state. Analysis shows that it contains twice as much oxygen 
in proportion to the hydrogen as does water, hence, having decided that 
water is represented by the formula H,O we may conclude that hydrogen 
peroxide must be represented by one of the formulas HO, H2O2, H3O3, 
etc. These are alike in the proportion of hydrogen to oxygen, 1 to 16, 
but differ in the total weight of the molecule, i. e., the molecular weight. 
To distinguish between these possible formulas, therefore, we must find 
out whether the mol weighs 17 gm. or 34 gm. or 51 gm., etc. For such 
substances the molecular weight must be determined under conditions 
such that the substance does not decompose, which is the case often in 
dilute solutions. 
Among the properties of a solution which depend upon the molecular 
weight of the dissolved substance (called the solute) are the vapor 
pressure, the boiling-point and the freezing-point. To understand the 
nature of the effect of a solute we may apply again the kinetic theory. 
If we have a pure liquid in contact with its vapor in a closed vessel, we 
imagine that there is a constant interchange of molecules between the 
two phases, molecules going from the liquid into the vapor phase at the 
same rate that others go from the vapor phase back into the liquid. If 
now, keeping the temperature constant, we add to the liquid in the closed 
vessel some molecules of another kind which are not volatile, and so 
cannot escape into the vapor phase, the number of molecules capable of 
escaping from the solution is now less than before, resulting in a smaller 
pressure of vapor above the liquid. If the new rrtolecules introduced 
are attracted in about the same way as the molecules of solvent originally 
present, then we might expect that if the volatile molecules of solvent 
constitute, say, only 0.9 of the total number of molecules, their vapor 
pressure would be only 0.9 of what it is when the solvent is pure, 
a conclusion borne out by experiment. This may be used to determine 
molecular weights, for if we find that a certain number of grams of solute 
36 in a certain number of mols of solvent lower the vapor pressure of the 
latter to, say, 0.9 of its value when the solvent is pure, then 0.9 of the 
total number of molecules are of the solvent and 0.1 are of the solute, 
from which we can calculate the number of grams per mol of solute. 
Example: Pure carbon disulfide, CS2, has a vapor pressure of 297 mm. 
at 20°C. It was found by experiment that 1.27 gm. of iodine dissolved 
in 38 gm. of CS2 lowered the vapor pressure of the latter to 270 mm., 
and we wish to know the formula of dissolved iodine, given the atomic 
weight, 127. Since 270 mm. is 0.909 of 297 mm., the number of mols 
of CS2 is now 0.909 of the total number, and the number of mols of 
iodine is therefore 0.091 of the total number, or 0.1 of the number of 
mols of CS2. One mol of CS2 weighs 76 gm., and therefore 38 gm. is 
0.5 mol. Since there are 0.1 as many mols of iodine as of CS2, or 0.05 
mols, then 127 gm. of iodine is 0.5 mol, and 254 gm. is 1 mol. The 
gram-atom of iodine being 127 gm., there are two gram-atoms per mol, 
and hence 2 atoms per molecule, so that the formula is I2. 
Instead of measuring the number of dissolved mols by the lowering 
of the vapor pressure of the solvent, the temperature being kept the 
same, we may, on the other hand, keep the pressure the same, say 1 
atmosphere, and determine how much the temperature must be increased 
in order to make the vapor pressure of the solvent great enough to con- 
tinu boiling, i. e., we may measure the rise in the boiling point produced 
by the added solute. Although the relation between this rise and the 
number of mols of solute is not so obvious as that between the lowering of 
vapor pressure and the number of mols of solute, it is probably 
evident that it is only the number of added mols and not their kind 
that determines the rise in the boiling point, so that having determined 
this effect for one solute of known molecular weight, we can use it 
to determine the molecular weight of other solutes. Just as it is 
harder to vaporize the solvent from a solution than it is from the 
pure solvent, so it is harder to remove it by any other means, such 
as freezing; and just as the boiling point of a solution of a non-volatile 
substance is higher than that of the solvent, so the freezing point of 
the solution will be lower, and will depend, likewise, only on the 
number of mols of solute in a given amount of solution. In addition, 
therefore, to defining a mol experimentally as the number of grams which 
in the vapor state and under standard conditions, would occupy 22.4 1., 
we may also define a mol as the number of grams which, dissolved in a 
specified amount of some solvent, will cause a certain rise in the boiling 
point or lowering in the freezing point. For water as solvent, the rise 
in boiling point is 0.52° and the lowering in the freezing point is 1.86° 
when 1 mol of any solute is contained in 1000 gm. of water. For other 
concentrations . the effect is nearly proportional. These numbers have 
been calculated from the heats of vaporization and fusion, but the ex- 
planation of the method lies beyond the scope of this book. Confirmation 
is found in the fact that a substance like alcohol, which can exist both 
as vapor and in aqueous solution, gives approximately, the same molal 
weight by both methods. 
Example: It was found that a solution of 7.29 gm. of sugar in 
100 gm. of water had a freezing point 0.395° lower than that of 
water; what is the molecular weight of the sugar? In 1000 gm of 
a solution of the same concentration there would be 72.9 gm. of 
sugar. Now, 1 mol of sugar in this amount of water would produce a 
lowering of 7.86°, and since 72.9 gm. produces a lowering of only 0.395°, 
37 which is 0-395/i.86 or 0.213 of the lowering produced by one mol, there must 
be only 0.213 of a mol present. If, then, 72.9 gr. is 0.213 of a mol, 1 
mol is ^2,9/o.2i3 Sm- or 342 gm. Comparing this with the empirical formula 
of sugar given on p. 36. C12H22O11, as determined by analysis, we see 
that the simplest formula is the correct one. 
Exercises. 
1. It is found by experiment that 3 vols. of hydrogen and 1 vol. 
of nitrogen are produced by the decomposition of 2 vols. of ammonia. 
If the molecules of hydrogen and nitrogen have been determined to be 
H, and N2, respectively, what must be the formula of ammonia? 
2. What are the relative volumes of the gases represented by the 
following equations: 
2CO + O. = 2CO,; 
2H,S + 3O2 = 2H2O + 2SO2. 
3. Write equations expressing the combustion, to CO2 and H2O, 
of two samples of gasoline consisting supposedly of C6H14 and C8H1S 
respectively. (a) What volume of oxygen is required per volume of 
vapor, in each case? (b) What approximate volume of air? 
Ans. (a) 9.5; 13.5; (b) 47.5; 67.5. 
4. What is the weight of 22.4 liters of each of the following gases 
under standard conditions: N2, CO, H2S, NH3, Hcl NO? 
Ans. 28 gm.; 28 gm.; 34 gm.; 17 gm.; 36.5 gm.; 30 gm. 
5. Considering air to be a mixture of Vs oxygen and 4/- nitrogen, 
by volume, what is the weight of 22.4 liters of it? Ans. 28.8 gms. 
6. Which of the gases in Question 4 are heavier and which lighter 
than air? 
7. What is the volume of 10 grams of hydrogen (a) under stand- 
ard conditions; (b) at 546°C and 10 atmos.? 
Ans (a) 112 1.; (b) 33.6 1. 
8. How many grams of oxygen are in a tank having a volume of 
5 liters, the pressure being 10 atmos., and the temperature 27°C? 
Ans. 65 gm. 
9. How many liters of CO2, measured at 77°C. and 2 atmos. could 
be produced from 2 kilograms of CaCO3? Ans. 287 1. 
10. A bulb having a capacity of 450 cc. weighs 20.155 grams when 
empty. It is then filled with a certain gas at a temperature of 177°C 
and a pressure of one atmosphere. The weight when filled with gas is 
found to be 21.170 grams. What is the molecular weight of the gas? 
Ans. 83. 
11. A certain gas is known to have a formula of the type CH, 
C2H2, etc. Given that 295 cc. of it weigh 0.317 grams, at 22°C and 1 
atmosphere; which formula is correct? 
12. What would be the relative efficiency of hydrogen, H2, and 
methane, CH4, in raising a balloon. Ans. 2.1 to 1. 
13. Hydrogen peroxide contains 1 part by weight of hydrogen to 
16 parts of oxygen, and a solution of 0.369 gm. of it in 25 gm of water 
had a freezing point of --0.805°. What is its formula? Ans. H,O2. 
14. The vapor pressure of pure chloroform, CHC1S is 370 mm. at 
40°C. What will it become is 6.40 gm. of naphthalene, C10H8, is dissolved 
in 239 gm. of CHC13? Ans. 361 mm. 
38 CHAPTER V. 
TYPES OF CHEMICAL COMPOUNDS. 
Relation Between Stability of Compounds and Metallic or Non- 
Metallic Character of Their Constituent Elements. On examining the 
known chemical compounds containing but two elements, the binary 
compounds, it is found that they represent the large majority of the 
possible pairs of elements, so that nearly all of the elements are capable 
of chemical union with each other. There is, however, a great difference 
in the nature and stability of the compounds formed, depending on the 
nature of the constituent elements. For example, oxygen and fluorine 
form no compound, oxygen and chlorine form two unstable compounds, 
easily decomposed by heating, oxygen and zinc form a very stable 
compound. On examining a large number of binary compounds in this 
way we find usually that those which are most stable are composed 
of a metallic and a non-metallic element. In Chapter I the physical 
properties differentiating metals and non-metals were discussed. We will 
now see that these physical differences are accompanied by fundamental 
differences in chemical behavior. As was pointed out, there is no sharp 
dividing line between the metals and non-metals, but rather they grade 
into each other. Accordingly it is quite logical to speak of one element 
being more or less metallic than another, and we might make an attempt 
to arrange them in the order of their metallic character. It is impossible 
to make such an arrangement very strict, for the order would vary some- 
what according to the criterion chosen. This difficulty does not, however, 
prevent the notion of the role of metallic character in chemical combina- 
tion from being an exceedingly useful one. As an illustration we may give 
the following list which contains a few of the commoner elements, ranging 
from those which are extremely metallic, at the top, to the opposite ex- 
treme of those least mjetallic at the bottom.* 
The "Alkali Metals," Potassium, Sodium, etc. 
The "Alakaline Earth Metals," Barium, Calcium, etc. 
Aluminum, Magnesium, Zinc. 
Tron, Nickel, Tin, Lead. 
Hydrogen. 
The "noble metals," Silver, Gold. 
Sulfur. 
Iodine. 
Chlorine. 
Oxygen, Fluorine. 
The most stable compounds are those formed between elements far 
apart on this list. The alkali metals unite with elements like oxygen and 
chlorine very eagerly, with the evolution of a great deal of heat, and 
these compounds require a great deal of energy for their decompo- 
sition.. The resulting compounds, furthermore, are very unlike their 
component elements. Sodium chloride, common salt, bears no resem- 
blance, physical or chemical, to its component elements. On the other 
hand, elements near together, like chlorine and oxygen, or chlorine 
and iodine, form compounds very similar in properties to their component 
* A fuller discussion of the relative metallic character of the individual elements 
will be found in Chapters XVI and XVII. 
39 elements, and which are readily broken down. The metals form numerous 
compounds with each other, but -very little energy is involved in the 
process, usually, unless the metals are very different, like mercury and 
sodium, and in many cases the compounds are much like the constituent 
elements in character. Hydrogen is rather more like a metal, chemically, 
than it is like a non-metal, as it forms compounds only with the most 
strongly metallic metals, and these are easily decomposed, whereas it 
combines with the non-metals to form compounds of increasing stability 
as we descend the list. 
Since it is usually possible to regard one of the elements in a binary 
compound as more metallic and the other as less metallic, or more non- 
metallic, it is useful to be able to represent this distinction in less cum- 
bersome terms, which we do by calling one constituent positive and the 
other negative. The choice of which to consider positive, instead of being 
entirely arbitrary, is made to conform to the behavior of the compounds 
in so far as they may be decomposed by electrolysis. When an electric 
current is passed through a solution of a binary compound, or through 
the liquid formed by its fusion, it may often be shown that the more 
metallic constituent travels with the positive current and may often be 
discharged as the free element at the kathode, or negative pole, while the 
non-metallic constituent may travel with the negative current, and may 
often be discharged at the anode, or positive pole. Accordingly, it is 
customary to refer to the metallic constituent as the positive element in 
the compound, and to the other as the negative element. Thus, when 
a water solution of hydrogen chloride gas is electrolyzed, using inert 
electrodes which are not acted upon chemically, such as graphite or 
platinum, hydrogen gas is liberated at the negative pole, whence it is con- 
sidered the positive or metallic element, while chlorine gas is liberated at 
the positive pole, so it is considered the negative element. Similarly, 
aluminum oxide may be dissolved in certain fused salts and decomposed by 
an electric current, giving aluminum, the positive element at the negative 
pole, and oxygen the negative or non-metallic element at the positive pole. 
In writing formulas of compounds the usual practice is to put first 
the symbol of the metallic or positive element, e. g., CuCl2, H2O, NaCl, 
PbS, S2C12. Sometimes the same element may be the more negative 
element in one compound and the more positive in another as is the case, 
apparently, with sulfur in the two compounds, respectively, just given. 
Unfortunately this convention is transgressed in a few cases like that 
of ammonia, usually written NHa, where the convention would require 
H3N. 
It is interesting to note that where more than one compound of the 
same element exists, the physical characteristics of the compounds seem 
to approach those of the element which is in excess. For example, there 
are a number of sulfides of sodium and potassium. Those containing least 
sulfur, as K2S, are white in color, but the color becomes increasingly 
yellow as more sulfur is present, and the higher sulfides act chemically 
very much like pure sulfur. Again, where more than one chloride of 
a metal exists, as illustrated by SnCl2 and SnCl4, FeCl2 and FeCl3, 
PbCl2 and PbCl4, the one containing more chlorine is more volatile, which 
we may connect with the fact that chlorine is a gas. 
Oxides and Sulfides. Among the most important of all chemical 
compounds are the oxides, largely because we live in an atmosphere 
containing oxygen, so that there is great opportunity for oxides to form. 
All the elements form oxides, except fluorine and the rare gases of the 
40 atmosphere, argon, neon, etc. Most of these oxides may be formed directly 
by burning the elements in the air, and this process, especially when light 
and heat are involved, is called combustion. Most fuels consist of carbon 
or hydrogen or their compounds. Coke and anthracite coal are fairly 
pure carbon, and burn to give carbon monoxide, CO, when the supply 
of air is insufficient, and to form carbon dioxide, CO2, when the air 
supply is sufficient. Other fuels, like gas, kerosene, gasoline, wood, 
contain either hydrogen or compounds of hydrogen with carbon, and, 
in the case of wood, oxygen. In all of these cases we have the oxidation 
of the hydrogen, on burning, to form water, along with the formation of 
CO2. Certain metals, called the base metals, oxidize quite readily. Sodium 
burns easily in air. Magnesium, when powdered, is an exceedingly com- 
bustible substance, giving off a very bright light, which makes it a useful 
constituent of flash-powders. We are accustomed to think of metals 
like zinc and aluminum as incombustible, but that is because the solid 
oxides formed, unlike the gaseous oxides of carbon, hydrogen, sulfur, 
etc., coat the metal, preventing access of the oxygen to the latter. If 
these metals are finely divided, as in turnings, foil or powder, they 
burn quite readily. A still less positive metal, like iron, will not burn 
by itself in the air, but when in the form of wire will burn vigorously 
in pure oxygen and will rust in moist air forming an iron oxide, a process 
similar to ordinary combustion. A nobler metal like copper cannot be 
ignited, even in pure oxygen, but if it is kept sufficiently hot by holding 
it in an ordinary flame, it will gradually change to oxide. The noblest 
metals, like silver, gold and platinum', will not combine with oxygen 
directly. Their oxides must be prepared by roundabout means, and are 
easily decomposed. They are the metals which occur free in nature. 
Sulfur is an element which is capable of forming several types of 
compounds, as we shall see later, but we may note now that in one of 
these types it closely resembles oxygen. Accordingly we have sulfides 
analogous to oxides in nearly all cases. The following formulas may be 
given to show this resemblance: 
H2O, Cu2O, CuO, ZnO, A1,O3, CO,, SnO, SnO,, Ag,, Ag,O, Na,O2, 
H2S, Cu2S, CuS, ZnS, A12S2, CS2, SnS, SnS2, Ag2S, Na2S2. 
Nearly all the oxides and sulfides of the metals are insoluble in 
water and have very high melting points. 
Halides. Next to the oxides, probably the halides of the metals are 
most important. By this name we designate the binary compounds of 
the elements known as the halogens, i. e., fluorine, chlorine, bromine 
and iodine. (The name comes from the Greek for sea-salt, sodium 
chloride being typical of the group). The halides of a given metal are 
usually all very much alike as to solubility, appearance, etc. Those 
of fluorine show the least resemblance to the others. Nearly all of them 
are easily dissolved by water, and most of them can be melted without 
decomposition at moderate temperatures. 
The Relation of the Relative Positive and Negative Characters of 
Elements to the Reactions of Their Compounds, (a) Replacement. The 
fruitfulness of the conception of the relative positive and negative char- 
acters of the elements is illustrated by its application in helping to 
determine the course of reactions. We may consider, first, reactions in 
which one element replaces another. What has been said regarding the 
relation between the stability of compounds and the positive and nega- 
tive character of their component elements would lead us to expect that 
41 a more postive element would replace a less positive one in a compound. 
This is usually the case, as illustrated by the following reactions: 
3Na + A1C13 = 3NaCl + Al 
Fe 4- CuCL = FeCl2 + Cu 
Cu+ HgCl2'= CuCl2 4- Hg. 
Similarly, a more negative element may replace a less negative one., 
as illustrated by the reaction 
Cal2 4- Cl, = CaCl2 + I2. 
Carbon has a much greater power to combine with oxygen than would 
be inferred from its non-metallic character, and does not lend itself to 
inclusion in the arrangement previously given. We may, however, com- 
pare the relative stability of various metallic oxides by considering the 
relative ease with which they may be reduced to metal by carbon. 
Aluminum oxide, A12O3, for example, is so stable that carbon is unable 
to remove the oxygen from it. The oxide of zinc, a less positive metal, 
can be so reduced, the zinc of commerce being obtained in this way. 
ZnO 4- C = Zn 4- CO. 
The oxides of iron, a still less positive metal, can be reduced more easily, 
as is done in the blast furnace. 
Fe2O3 4- 3C = 3CO 4- 2Fe. 
The dioxide of tin is still more easily reduced, as tin is nobler than 
iron. 
SnO2 4~ 2C = Sn 4~ 2CO. 
(b) M etathesis. When two compounds are brought together there 
occurs frequently a type of reaction known as a metathetical reaction, in 
which certain elements or groups of elements change places. The follow- 
ing equations represent reactions of this type: 
AgCl 4- KI = Agl 4- KC1. 
CaO + 2HC1 = CaCk 4- H2O. 
A12S3 4- 3H2O = A12O3 4- 3H2S. 
The important thing to note is that the more metallic element of one 
compound changes place with the more metallic element of the other. (In 
each case the positive element is written first in the formula, according 
to the convention stated previously). We seldom find such a reaction as 
AgCl 4- KI =AgK 4- C1I, 
(which would represent the formation of compounds containing like 
elements) when unlike elements have the opportunity to combine, as 
represented above. Of course, it may not be possible to tell whether the 
most stable pair of compounds would be silver chloride and potassium 
iodide or silver iodide and potassium chloride. This would have to be 
determined by experiment, or by the aid of criteria to be discussed in 
Chapters XII and XIII. 
(c) Addition. Another type of reaction between binary compounds 
occurring frequently where the two compounds have one element in 
common, consists simply of a combination to form a less simple type of 
compound. The most interesting of these combinations are those taking 
place between oxides. We may, in order to understand these combina- 
tions, regard the oxides as possessing a residual positive or negative 
nature dependent on the character and the amounts of the elements in 
the oxide. Thus, the more positive elements would yield more positive 
oxides, and vice-persa, and also, where two oxides of the same element 
exist, the one containing more of the highly negative element, oxygen, 
would be more negative. To illustrate we give a list of oxides arranged 
42 according to these ideas, with the more positive oxides at the top, as we did 
before when considering the elements. 
K2O, Na2O 
CaO, MgO 
ALOS, ZnO 
H.,O 
CO., 
SO, 
so3 
In accordance with this arrangement, and its governing idea, we 
find that the oxides far apart show great tendency towards combination. 
Plaster of Paris, which may be considered as a compound of CaO and 
SO3 is a very stable substance. Limestone, CaO'CO, (usually written 
CaCOg) is less stable and is broken down at a red heat into the gas 
CO2 and CaO, quicklime. When water, H2O, is added to quicklime, 
slaked lime is formed, CaO. H,O (also written Ca(OH)2), and this is 
broken down (into quicklime and steam) more easily than is limestone. 
The combination between H,O and the oxide of a still more positive 
metal than calcium gives, as we should expect, a more stable compound 
than Ca(OH)2. Illustrations might be multiplied ad libitum. Enough 
has been said to show what oxides might be expected to combine most 
readily, and to show how compounds of this sort might be classified 
according to the oxides of which we may consider them to be composed. 
As might be expected from the similarity of sulfur and oxygen 
compounds, many double sulfides exist, similar to the compounds between 
the oxides just discussed. These are numerous in the mineral world, as 
exemplified by 3Cu2S-Sb2S3, PbS*Sb,S3, 3Ag2S'As2S3. 
Many combinations of halides exist corresponding to those of oxides, 
though they are not so numerous or important. The chlorides of the 
base metals and those of the noble metals fulfill our expectations in fur- 
nishing compounds of this nature, some of which, like KC1 * AuC13, 
2KC1 'PtCl^ are very stable. With iodides we have, similarly, 2KI 'HgL, 
etc. The fluorides are especially prone to form double compounds, as 
witnessed by HF'KF, 2HF • SiF4, BaF • SiF4. 
Classification of Compounds. We have seen that binary compounds 
may be classified very simply according to the positive and negative 
elements of which they are composed. In this sense we might investigate 
and compare the various silver compounds and their reactions, on the one 
hand, or the various chlorides, on the other hand. When we come to 
compounds containing more than two elements we have seen, further, 
how we may classify them according to the oxides, chlorides, sulfides, 
fluorides, etc., of which we may consider them composed. For many 
purposes this is a most useful classification. There is, however, another 
way in which we may regard their constitution, a way especially useful 
for reactions occurring in water solution. 
We have seen how a compound like silver sulfate, Ag2SO+, may be 
regarded as composed of the oxides Ag2O and SO3, as is indicated by 
writing the formula Ag2O • SO3. Many of its reactions could profitably 
be represented in this way. However, if we mix a solution of this com- 
pound with one of potassium chloride, KC1, instead of an interchange of 
K and Ag2O, we have an interchange of K and Ag, represented by the 
equation: 
2KC1 + Ag2SO4 = K2SO4 + 2AgCl. 
It would evidently be more satisfactory, in this reaction, to regard silver 
43 sulfate as composed of Ag and SO4. This group, SO4, is called the 
sulfate radical. In the reaction between calcium sulfate, CaSO4, and 
sodium carbonate, NaoC03, we might still use the oxide notation, writing: 
CaO • SO3 4- Na,O • CO2 = CaO • CO2 + Na2O • SO3. 
Or, we might consider the reaction to be an interchange of metals rather 
than an interchange of oxides, writing: 
CaSO4 + Na2CO3 = CaCO3 4- Na2SO4. 
When, however, we use calcium chloride, CaCl2, instead of CaSO4, an 
interchange of metals represents the only type of reaction possible, so 
that we must write 
CaCl2 4" Na2CO3 = CaCO3 4~ 2NaCl. 
From this and other examples that might be given, we see that the 
second method of considering the constitution of these compounds is 
much more inclusive, besides allowing us to consider both types of com- 
pounds as made up of the same fundamental constituents. The same 
applies to other double compounds, such as chlorides, as well as to 
oxides. If we mix solutions of Na,PtCle and KQSO4, the reaction is: 
Na2PtCl6 4- K2SO4 = K2PtCl6 4- Na2SO4, 
implying that the reaction represents an interchange between sodium and 
potassium and the radicals PtCl6 and SO4. If the constitution of the 
new compounds were that represented by 2NaCl-PtCl4 and K,O' SO3, 
we should be led to expect the wrong type of interchange entirely. Our 
conclusion must be that for some reactions one point of view is more 
useful, and for other reactions the other is better. Thus, when dealing 
with the decomposition of calcium carbonate by heat it is useful to 
think of it as a compound of CaO and CO2, whereas its reactions in 
the presence of water correspond better with the idea that it is composed 
of calcium and carbonate radical. Corresponding to this latter concep- 
tion we would classify compounds according to the constituent elements 
or radicals which take part in metathetical or interchange reactions.. We 
will proceed to do this so as to point out three very important types, 
i. e., acids, bases and salts. 
We have seen that it is usually better not to attempt to distinguish 
between compounds of KC1 and K2O, K2S, etc., but to consider them all 
as compounds of potassium, from which standpoint we would investigate 
the common properties of potassium compounds. Similarly we might 
investigate in groups, calcium compounds, lead compounds, and those 
of all the postive elements or metals. Then, in grouping die negative 
constituents, we would have, in addition to single elements, like chlorine, 
etc., negative groups, or radicals, containing more than one element, like 
sulfate radical, SO4, carbonate radical, CO3, nitrate radical, NO3, etc., 
and we would seek to determine what properties were common to all 
sulfates, carbonates, and nitrates, respectively. (It is seldom that the 
positive constituent consists of more than a single metallic element). The 
first of these classes for us to investigate is that composed of compounds 
in which the positive element is hydrogen. 
Acids. The most obvious property of many compounds in which 
hydrogen is the positive constituent is a sour taste. This taste we are 
accustomed, in everyday life, to ascribe to substances we call acids, and 
since only certain hydrogen compounds give such a taste, we may 
consider, it the property of this kind of combined hydrogen. There are 
a number of more or less familiar acids which we may recall and char- 
acterize briefly. Hydrochloric acid, HC1, is a gas when in the pure 
state, but it is very soluble in water, forming the solution sometimes 
44 known as "muriatic acid.'' Sulfuric acid, H2SO4, is a dense syrupy 
liquid, in the absence of water, and has the common name of "oil of 
vitriol." It is manufactured and used in larger quantities than any other 
substance for purely chemical purposes. Nitric acid, HNO3, is a fuming 
corrosive liquid extensively used to dissolve the nobler metals and in 
the manufacture of explosives. Acetic acid HC2H3O2 (only one of the 
hydrogen atoms has acid properties), is the acid of vinegar. Phosphoric 
acid, H3PO4, is used in the phosphate drink at the soda fountain. Boric 
acid H3BO3, or "boracic acid," is a solid giving a solution of very faint 
acid properties. It finds use as a mild antiseptic. Oxalic acid, H2C2O4, a 
solid in the pure state, finds frequent use as a chemical reagent. The 
various fruit juices contain characteristic acids, such as citric acid, 
found in lemons, tartaric acid in grapes, malic acid in apples, etc. 
Whenever we have the sour taste of an acid we may notice certain 
other properties which are, therefore, likewise characteristic of acids. 
The hydrogen of an acid may be replaced by a metal, indirectly, using 
a suitable compound of the metal, as we shall see later, when discussing 
bases, and directly, using the free metal if the acid is not too weak and if 
the metal is a base one such as zinc or magnesium. Ordinary acids, then, 
if not too weak, will dissolve base metals like zinc. The following 
equation represents a typical reaction, zinc and sulfuric acid forming zinc 
sulfate and hydrogen gas: 
Zn -f- H2SO4 = ZnSO4 -j- H2. 
Again, unless an acid is extremely weak in its acid properties, when it is 
brought in contact with a carbonate, such as sodium carbonate, Na2CO3, 
"washing soda," it liberates the weak carbonic acid, H2CO3, which breaks 
down into water and carbon dioxide gas, producing an effervescence 
unless a great deal of water is present. To illustrate this, we may give 
the equation for the action of sulfuric acid on sodium carbonate: 
H2SO4 + Na2CO3 = Na2SO4 + H2O + CO2. 
One of the most useful ways of detecting acids, especially when they 
are very dilute, is by means of certain natural and artificial coloring 
matters called "indicators," because they indicate whether or not a 
solution is acid. Among these may be mentioned litmus, a plant extract 
giving a blue color in pure water which is changed to red by the minutest 
traces of acid. Another indicator, an artificial dye called methyl orange, 
gives a yellow solution in water, changing to orange and then red as acid 
is added. It takes more acid in a given amount of water to change 
the color of this indicator than it does to turn litmus red. Another, 
methyl violet, requires a still more concentrated acid to change it suc- 
cessively from violet, through blue, green, and yellow as more acid is 
added. 
Bases. There is another class of compounds which are capable of 
neutralizing or destroying the above properties shown by acids in water 
solution. These are all characterized by the radical OH, called hydroxyl, 
being compounds of this group with metals. They are called bases. 
Most of them are insoluble in water, or nearly so, the chief exceptions 
being sodium hydroxide, NaOH, sometimes called "caustic soda," 
potassium hydroxide, KOH, sometimes called "caustic potash," its solu- 
tion being known often as "lye," and barium hydroxide, Ba(OH)2. Cal- 
cium hydroxide, Ca(OH)2, "slaked lime," is sparingly soluble, its 
solution being known as "lime water." To these should be added ammo- 
nium hydroxide, NH4OH, formed when ammonia gas, NH3, is dissolved 
in water. The group NH4 acts like a metal in combination. The soluble 
45 bases, especially the first two, are often called "alkalis," and their solutions 
"alkaline." The property fundamental to bases is their ability to neutralize 
acids. This reaction is an exchange such that the hydrogen of the acid 
arid the hydroxyl of the base unite to form water. The equations repre- 
senting neutralization are among the easiest to write, for we need one 
OH group for every hydrogen atom. The following equations show 
examples of neutralization: 
NaOH + HC1 = HOH (or H„O) + NaCl. 
2KOH 4- H„SO4 = 2HoO + K2SO4. 
Ca(OH)2 + 2HNOo = 2H2O + Ca(NO3)„. 
Al (OH), 4- 3HC1 = 3H.O 4- A1C13. 
2A1(OH)3 4- 3H2SO4 = 6H2O 4- A1,(SO4),. 
Besides the ability to neutralize acids, and their common constitu- 
tion as hydroxides of metals, which characterize all bases, there are 
certain additional important qualities of the solutions of soluble bases, or 
alkalis we may regard as the properties of the hydroxyl radical dissolved 
in water. One of these is a peculiar taste, for which we have no word as 
"sour" for the corresponding taste of acids. It is the taste of lime 
water. Again, alkaline solutions have an emulsifying action on fats, and 
a tendency to produce soap with grease, which gives a soapy feeling when 
such a solution is rubbed between the fingers. Most conclusive, however, 
are the colors given to indicators. When litmus has been turned red by 
acid the addition of alkali will restore the blue color. Similar changes 
are produced with the other indicators previously mentioned. A very 
useful indicator is phenolphthalein, which is colorless in pure water or 
in acid, but is turned a deep red by very dilute alkali. Still others, such 
as tri-nitrobenzene, require considerable amounts of alkali in order to 
produce a change in color. By a suitable choice of indicators it is 
possible to cover the entire range from- strongly acid solutions, through 
neutral, to strongly alkaline solutions. 
Salts. The result of the neutralization of acids and bases is the 
formation of water, together with a substance called a salt. The name 
is derived from the typical substance common salt, sodium chloride; 
and salts in general, like it, are usually solid crystalline substances at 
ordinary temperatures. The number of salts known is nearly equal to 
the product of the number of existing positive constituents of bases and 
the number of negative constituents of acids. The natural classi- 
fication of all salts is according to these constituents or radicals. 
Thus, we might make a study of all the salts of zinc, on the one 
hand, or of all sulfates, on the other. In a sense the acids are simply 
the salts of hydrogen, and the bases those of hydroxyl; the extremely 
important role played by acids and bases, however, justifies their separate 
classification. We see that the idea of a salt is a more general one than 
the idea that a substance may be made up of positive and negative oxides 
or halides, etc., for it includes all substances which contain decidedly 
positive and negative constituents. The question as to how positive and 
negative these constituents must be in order that the compound may be 
considered a salt is no more capable of an arbitrary answer than is the 
question as to how sour a solution of a substance must be in order that 
the substance may be called acid. We have all gradations between un- 
doubted salts like common salt, though substances of feeble salt-like 
nature, such as tetrachloride of tin, to sulfur or iodine chlorides, 
which are not usually considered as salts at all. Sharp boundaries 
46 of this sort are not very frequent in nature, and we must always 
remember that many of our classifications are purely for the sake of 
convenience and not to be overworked. We may take definitions of 
such terms as metal, salt, acid, etc., but when we come to fit actual sub- 
stances to our definitions we are confronted by such problems as how 
good a conductor, or how malleable, etc., an element must be to be 
called a metal, or again, how much difference there must be in the posi- 
tive and negative character of the elements in a compound before we will 
be willing to class it as a salt. History shows must waste of paper and 
good temper by futile controversies of this sort. 
Anhydrides of Acids and Bases. From what has been said above 
it is evident that a base may be looked upon either as a compound of a 
metallic element with the hydroxyl radical, or as a compound of a metallic 
oxide with water. Several further examples may be given in the following 
equations: 
Na.O 4- H.O = 2NaOH 
CaO 4 H.O = Ca(OH). 
Cu(OH)„ = CuO + H4 
2A1(OH)3 =A12O3 4 3H2O. 
On the basis of this relation between the hydroxides and the oxides 
the latter are often called basic anhydrides, the latter word signifying the 
removal of water. Similarly, acids which contain oxygen as well as 
rogen may be dehydrated by the removal of water, or may be formed 
from the non-metallic oxide and water, as the case may be. The follow- 
will serve as examples: 
H.O 42CO2 = HoCO, 
H.O 4 N.6„ = 2HNO3. 
H2O 4 so2 = h,so3 
H.O 4 SO3 = H.SO4 
H.O 4 P2O-, = 2HPO3 
2H.O 4 P265 = H4P.O7 
3H2O 4 P2O5 = 2H3PO4 
The oxides from which these acids may be formed are likewise 
often called acid-anhydrides. The important thing to note is the similarity 
of the chemical behavior of the acid and basic anhydrides to that of their 
respective acids and bases. When acids and bases neutralize each other 
they produce a salt and water. Since the difference between the acid 
or base and its anhydride is simply water, the difference between their 
reactions with each other and ordinary neutralization is simply in the 
amount of water produced. We may write a series of reactions to 
illustrate this relationship: 
Ca(OH)2 4 H.SO3 = CaSO3 4 2H2O 
CaO 4 H.SO3 = CaSO3'4 H»O 
Ca(OH)2 4 SO2 = CaSO3 4 H2O 
CaO 4 SO. = CaSO3 
We see here the fusion of the two ideas of the classification of a 
salt like CaSO3 either as composed of a positive metal and a negative acid 
radical or else as a compound of a positive and relatively negative oxide. 
We see also that a metallic oxide will neuralize an acid, or an acid 
anhydride will neutralize a base, hence we must expect the chemistry 
of the anhydrides and of their acids and bases to be very similar. 
47 Exercises. 
1. Make a list of the characteristics of acids and bases. 
2. What is meant by each of the terms, alkali; alkaline reaction; 
radical; metathesis; neutralization; anhydride? 
3. Determine, by subtracting H2O in proper proportion, what are 
the anhydrides of the following substances: HNO3; HnCO3; Fe(OH)3; 
Fe(OH)2; H3PO4; HPO3; H3PO3; CuOH; H2SO3. 
4. From your knowledge of the metallic or non-metallic nature of 
the elements involved, state which constituents of the following pairs of 
compounds might be expected to exchange if they react with each other: 
(1) AgoS04 and Na2CO3; (2) CuCl2 and K2S; (3) ZnS and 2HC1; 
(4) Ca(OH)2 and Na2CO3; (5) Ca(OH)2 and H2SO4. 
5. Which substance in each of the following pairs would decompose 
more easily, on heating, in the manner represented by the respective 
equations: 
(1) CuCO3 - CuO + CO2 or 
CaCO3 = CaO + CO2; 
(2) CuCO3 = CuO + CO2 or 
CuSO4 = CuO -|- SO3; 
(3) Ca(OH)2 = CaO + H2O or 
2A1(OH)3 = A12O3 + 3H2O? 
48 CHAPTER VI. 
VALENCE. WRITING EQUATIONS. 
For the purpose of writing formulas and equations it is very im- 
portant to know the number of atoms of substances that are likely to 
combine to form compounds, without being under the necessity of per- 
forming an analysis of each substance or of remembering its composition 
as an isolated fact. It is possible to do this by the aid of a number, 
or numbers, characteristic of each element, which represent rather 
uniformly the relative number of atoms of the different elements that 
may be found in a molecule. The number expressing this combining 
value for an element is called its valence. We will first see how the 
valence of an element may be determined and then see how it may be 
used in predicting formulas and writing equations. 
Assignment of Valence Numbers. If we classify the binary com- 
pounds of hydrogen according to the numbers of atoms of hydrogen per 
atom of the other element we will find them dividing into groups as 
follows: 
HC1 
HBr 
HI 
HoO 
H2S 
H3N 
h3p 
H3As 
H4C 
H4Si 
We notice that the elements of these compounds may be divided according 
to their power to combine with hydrogen. One group of elements, Cl, 
Br, I, are able to combine atom for atom with hydrogen; another, rep- 
resented by O and S, have atoms which can combine with two atoms of 
hydrogen; still another group of atoms can combine each with three 
hydrogen atoms, another with four. Evidently the elements in the second, 
third and fourth columns have respectively 2, 3 and 4 times the combining 
power, measured with hydrogen, that the elements in the first group 
have. Moreover, we do not find elements having a smaller combining 
power than hydrogen. We may thus take the hydrogen atom as our unit 
for determining the combining value or valence. We will find, then, 
that we are able to measure the combining value, or valence, of any 
kind of atom by determining from the formulas of some of its com- 
pounds its power either to combine with hydrogen or to replace it in 
some compound such as those given in the table. Thus an atom of oxygen, 
O, can combine with two atoms of hydrogen, as in H2O, while an atom 
of calcium, Ca, can replace two atoms of hydrogen, as is seen by com- 
paring CaO with H2O. It is, however, important to distinguish between 
power to combine with hydrogen and power to replace hydrogen, for 
atoms that can replace hydrogen can replace each other, but as a rule 
cannot combine with each other, while atoms that can combine with 
hydrogen can also replace each other but cannot as a rule combine 
with each other. Calcium and oxygen are both related to two atoms of 
hydrogen, but in a different way, the former replacing hydrogen, the 
latter combining with two atoms of hydrogen. We must, therefore, dis- 
tinguish between two kinds of valence, as well as between different 
numerical values. This we will do by calling one positive and the other 
negative. The decision as to which shall be postive is based on the 
electrical behavior, to be explained more fully in Chapter X, and follows 
the metallic or non-metallic character of the elements discussed in 
49 Chapter V, where we called the more metallic element in the compound 
the positive element and vice-versa. In the above compounds the hydrogen 
appears to be the positive constituent, so that we will call its valence 
4-1. Any element, furthermore, which replaces hydrogen will be con- 
sidered as showing a positive valence, while any element combining with 
hydrogen or its equivalent will be considered as showing negative valence. 
Accordingly, in the above list of compounds, the valence of Cl, Br and 
I is considered -1, that of O and S -2, that of N, P, and As, -3. In 
CaO the valence of Ca is 4-2, since that of O is taken as -2, also since it 
appears, in comparing H2O and CaO, that one atom of calcium replaces 
two of hydrogen. In NaCl the valence of Na is 4-1 and that of Cl -1, 
as can be seen by comparing NaCl with HO, from which it may be 
obtained. In ALO3 the valence of Al is 3, because the valence of 
oxygen is -2 and three atoms represent -3, hence the two atoms of 
aluminum represent 4-6, each atom being -|-3. As further illustrations we 
may give a number of compounds, writing under each element the valence 
per atom, and indicating the multiple of this corresponding io the number 
of atoms in the molecule. The sum of the total positive and negative 
valence is, of course, zero. 
MgCl2 
+ 2,2(-l) 
SnCl4 
+4,4(-l) 
H3BO3 
3(4-1), 4-3, 3(-2) 
NaOH 
+1, -2, 4-1 
Na.S 
2(4-1)', -2 
* PbO, 
+4, 2(-2) 
It is possible to consider the valence of the radicals discussed in 
Chapter V, without paying, any attention to the valence of the atoms 
within the radical. In H2SO4 it is obvious that the valence of the SO4 
radical is -2, and in K2PtClt) that the valence of the- PtClQ radical is -2. 
We may, similarly, speak of the valence of the hydroxyl radical as -1, 
of that of the Fe(CN)6 radical in K4Fe(CN)6 as -4. Where the radical 
in question remains intact in a reaction, it is unnecessary to consider the 
valence of its component elements. 
Elements Exhibiting More Than One Valence. Many of the 
elements show more than one valence in their various compounds. Tin, 
for example, gives two series of compounds, the stannous compounds, in 
which the valence of tin is 2, and the stannic compounds in which it is 
4, as exemplified by the following: 
Stannous 
SnO, 
+2, 2(-l) 
SnO 
4-2, -2 
SnS 
4-2, -2 
Sn(OH), 
4-2, (-2,4-1) 
Stannic 
SnCl4 
+4,4(-l) 
SnOo 
+4, 2(-2) 
SnS2 
+4, 2(-2) 
Sn(OH)4 
4-4, 4(-2, 4-1) 
In this case we have no doubt that the tin is the element that changes 
valence rather than the other elements present in the above compound 
because the other elements react in the same way in both cases. For 
example, upon the addition of silver nitrate solution to a solution of either 
of the chlorides of tin the same silver chloride precipitate is produced. 
We must, therefore, have the same kind of chlorine in both compounds. 
Again both sulfides give hydrogen sulfide when acted upon by hydro- 
chloric acid, as represented by the equations: 
SnS + 2HC1 = SnCl2 + H2S 
SnS2 + 4HC1 = SnCl4 + 2H2S. 
As contrasted with the tin compounds*we have two oxides of barium, 
BaO and BaO2. These appear analogous to the oxides of tin given just 
mentioned, and we might be inclined to say that there must be two series 
of barium compounds in which barium would have a valence of 2 and 4 
50 respectively. However, we fail to find any other barium compounds in 
which the valence appears to be other than 2. Also, on investigating the 
chemical reactions of BaO2 we find that while it reacts to give barium 
compounds in which the valence of barium is evidently 2, it gives oxygen 
compounds quite different from the ordinary oxygen compounds. The 
following equations will bring out this contrast in the behaviors of SnO2 
and BaO.,: 
SnO9 4- 4HC1 = SnCl4 + 2H.,O 
BaO2 + 2HC1 = BaCl2 + H2O2. 
There are no other barium compounds and no other hydrogen compounds 
in which we would be tempted to assign to those elements any but the 
ordinary valences, 2 and 1 respectively. We find by experiment, however, 
that the oxygen in these compounds, together with a few others, cannot 
be classed with the oxygen in most compounds. It behaves quite differ- 
ently, and it is therefore logical to ascribe to it a different valence. We 
therefore conclude that the 2 atoms of oxygen in these compounds (to 
all of which the term peroxide is given) have together a valence of -2. 
It is impossible to decide whether both of these two atoms have the same 
valence, -1, or different valences, like 0 and -2, whose mean is -1. 
For practical purposes, however, it is unnecessary to distinguish these 
atoms from each other, so that we may consider them both to have a 
valence of - 1. In all other oxygen compounds it is satisfactory to 
ascribe to the oxygen a valence of -2. A similar difficulty is encountered 
with respect to the valence of sulfur in sodium thiosulfate, Na2S2O3, 
the "hypo" of the photographer. Since Na and O have valences respect- 
ively of 1 and -2, as usual, that of the two sulfur atoms together must 
be -|-4. This does not necessarily mean, however, that both of these 
atoms have the same valence, 2, for a total of 4 would be obtained if 
one valence were 3 and the other 1, or if one were 5 and the other -1, 
etc. The probable valence, again, can be determined with any degree of 
certainty only by finding what kinds of compounds the two atoms of 
sulfur produce when thiosulfates are decomposed. 
Some of the elements, especially some of the non-metals, exhibit 
even more than two valences. In the following compounds, KC1, KC1O, 
KC1O2, KC1O3 and KC1O4, the potassium and the oxygen are believed 
to have the usual valences, 1 and -2, respectively, which indicates that 
the valences of the chlorine are respectively -1, 1, 3, 5 and 7. Likewise 
with phosphorus we have such compounds as H3P (usually written PH3, 
though the hydrogen is doubtless the positive element) H3PO2, H3PO3 and 
H3PO4. In these compounds the valences of phosphorus seem to be 
respectively -3, 1, 3 and 5. 
Use of Valence in Classifying Compounds. Where an element shows 
a varying valence, as in the instances just given, the determination of the 
valence gives a means of classifying together those compounds in which 
the element has the same valence, so that we may be led to anticipate 
similarities in. chemical behavior within the same class of compounds. In 
the compound NaPO3, which has no clear connection with any of the 
other compounds above, a simple calculation shows that the phosphorus 
has a valence of 5, wherefore it is to be classed with H3PO4. The same 
applies to an apparently different compound, Mg2P2O7 (Cf. MgCl2). 
On the other hand, in the compound Na4P2O5, the valence of the 
phosphorus places it with H3PO3. A proper understanding of the reac- 
tions of these different classes of conipounds is aided by considering 
valence. Thus PC13, where the valence is 3, yields H3PO3 with water, 
51 while PC15 yields H3PO4. Again, a regard for valence makes it evident 
that neither P2O3 or P2O5, in reacting with water, could be expected to 
give a compound such as H2PO3. 
Writing Formulas With the Aid of Valence. It has just been shown 
that the valence of the elements can be determined by examining a few of 
the compounds of each element. Having done this, it is possible to pre- 
dict the formulas of other compounds. In the table below are given 
valences of positive and negative constituents of compounds as determined 
on the previous pages, together with the formulas of the compounds that 
mip'ht he exnected tn result from their combination 
Constituent 
Cl 
Br 
OH 
O 
S 
SO4 
PO4 
Valence 
-1 
-1 
-1 
-2 
-2 
-2 
-3 
H 1 
HC1 
HBr 
HOH 
HoO 
HoS 
h„so4 
H3PO4 
Na 1 
NaCl 
NaBr 
NaOH 
Na,O 
Na,S 
Na2SO4 
Na3PO4 
K 1 
KC1 
KBr 
KOH 
KoO 
K„S 
k2so4 
K3PO4 
Ca 2 
CaCL 
CaBr2 
Ca(OH)2 
CaO 
CaS 
CaSO4 
Ca3(PO4)2 
Ba 2 
BaCl2 
BaBro 
Ba(OH)., 
BaO 
BaS 
BaSO4 
Ba3(PO4)2 
Al 3 
A1CL 
AlBr.' 
A1(OH)3" 
A12O3 
als3 
A12(SO4)3 
A1PO4 
$n 2 
SnCl2 
SnBr2 
Sn(OH)2 
SnO 
SnS 
SnSO4 
Sn3(PO4)2 
Sn 4 
SnCl4 
SnBr4 
Sn(OH)4 
SnOo 
SnSo 
Sn(S04)o 
Sn3(PO4)4 
In this table we have given a large number of compounds not pre- 
viously considered, solely on the basis of valences deduced from the few 
that have been considered. Illustrations might, of course, be multiplied 
indefinitely. 
Writing Equations. In connection with chemical equations it is 
important to realize that the mere writing of the equation confers no ob- 
ligation upon the substance involved to behave as indicated by the equation. 
The obligation is rather upon the one who writes the equation to endeavor 
to make it correspond to an actual or possible reality. A complete 
justification of an equation would require the experimental determination 
both of the substances produced, their formulas and their relative amounts. 
To avoid this labor and to aid in predicting possible reactions we may 
avail ourselves of the results of other experimenters which enable us to 
determine valence. We may expect that where a reaction takes place at all 
it will tend to produce compounds between elements or radicals having 
valence of different signs, and also in the proportions demanded by the 
numerical magnitude of the valences. It then becomes necessary to take 
the proper number of molecules of each substance so that there will be 
the same total number of atoms of each element on both sides, since 
it is this equality that justifies our use of the term equation. These steps 
will be best illustrated by examples. 
Let us consider, first, the reaction between aluminum and Fe3O4 
to form aluminum oxide and iron. To be absolutely certain of the 
formula of the aluminum oxide produced we would have to determine 
it by analysis. We do not, however, find aluminum compounds in which 
the valence of aluminum is other than 3, as we see it in A1C13, which 
would lead us to anticipate that the oxide formed would be A12O3, which 
is correct. Accordingly we have Al and Fe3O4 giving A12O3 and Fe. 
We must next adjust the number of molecules so that there will be the 
same number of atoms of each element on each side of the equation. 
It is obvious that we must take 3Fe3O4 to get an integral number, 4, of 
A12O3 molecules. Adjusting the number of atoms of Al and Fe we have 
as the equation 
3Fe3O4 + 8A1 = 9Fe + 4A12O3. 
52 It is always desirable to check the equation written by comparing the 
number of atoms of each element on each side of the equation. They 
should, of course, be equal. 
The products of the reaction between solutions of BaCl2 and 
A12(SO4)3 would evidently be those formed by an interchange of the 
positive constituents so as to form aluminum chloride and barium sulfate. 
From the valences of the four constituent elements and radicals we 
would predict the formulas of the latter to be A1C13 and BaSO4. Our 
equation is therefore, in part, 
BaCL + A12(SO4)3 = BaSO4 + A1C13. 
Adjusting the number of molecules of each substance so as to make the 
equation balance we have finally 
3BaCl2 + A12(SO4)3 = 3BaSO4 + 2A1C13. 
When solutions of Fe(NO3)3 and Ba(OH)2 are mixed a precipitate 
is produced which we might expect to be formed by the inter change of 
the constituents of the same sign, i. e., Fe and Ba. The valence of the 
Fe in this compound is seen to be 3 if we compare it with HNO3. The 
valence of each of the other constituents is already familiar with the pre- 
vious pages, so that the formulas of the compoundc produced might be 
predicted as Fe(OH)3 and Ba(NO3)2. Adjusting the number of 
molecules of these four substances so as to represent no gain or loss of 
atoms, we write as the equation: 
2Fe(NO3)3 + 3Ba(OH)2 = 2Fe(OH)3 4- 3Ba(NO3)2. 
Equations for the neutralization of acids and bases are very easily 
written by regarding the union of one H of the acid with one OH of 
the base. The numbers of molecules of acid and base must therefore be 
such that they will yield equal numbers of these constituents of water. 
Examples of neutralization equations illustrating this may be found on pp. 
Oxidation and Reduction. We have certain reactions in which there 
is a change of valence on the part of two or more elements. In both 
of the following reactions, 
2SnO -j- O2 =2SnO2 and 
SnCl2 + Cl2 = SnCl4 
the tin increases in valence, going from a compound of one class to that 
of another, a different reaction from those illustrated, on page -. In 
the first we have the addition of oxygen to the tin, wherefore we say that 
the tin has undergone oxidation. In the second, though no oxygen is in- 
volved, the same change of valence on the part of the tin takes place, 
so that the same name is applies to the process. The -term oxidation is 
applied whenever valence takes on a more positive (or less negative) value. 
The opposite process, the removal of oxygen or, in general, the decrease 
in valence, is called by the more obviously general name reduction. In 
the two following equations, starting with metallic tin, we evidently have 
oxidation of the tin likewise: 
Sn 4- Cl2 = SnCl2 
2Sn -|- O2 = 2SnO. 
In order to make this accord with our definition the valence of metallic 
tin must be zero. It is important to remember this throughout in dealing 
with the free elements. When we speak of the valence of an element we 
must refer to its valence in a particular compound or class of compounds. 
Accordingly the oxygen and chlorine in the above equations are both 
reduced. 
In balancing equations involving oxidation and reduction the principle 
invoked is to make the total increase in valence of all the atoms oxidized 
53 equal to the total decrease in valence of all the atoms reduced. Further 
discussion of this type of equation will be resumed later in Chapter XV. 
Exercises. 
1. What is the probable valence of each radical and element in each 
of the following compounds: CuSO4; Cu(NO3)2; Pb(NO3)2; PbCO3; 
CrCl3; MgO; NH4C1; CuS; FeCO3; KC1O3; K2CrO4; Fe(OH)2? 
2. Write the formulas of a number of the compounds that might be 
formed from the constituent parts of the above compounds (omitting the 
atoms within the radicals). 
3. Explain the terms: valence, radical, oxidation, reduction. 
4. Classify according to the valence of arsenic the following arsenic 
compounds: AsC13, As2O3, As2Or„ As2S5, NaAsS2 (consider the valence 
of S to be -2), H3As, H3AsO4, Ag3As. 
5. Complete and balance the following incomplete equations: 
Fe,O3 + C = CO + 
Mg(OH)2 + HC1 = 
MgO + HNO3 = 
Cr(OH)3 + H„SO4 = 
CrCl3 4- KOH = 
CrCl3 + AgNOo = 
Al + HgCl2 = Hg + A1C13 
NH4C1 + Ca(OH),= 
6. In each of the following reactions determine what elements are 
oxidized and what reduced: 
2HgCl, 4- SnCl2.= 2HgCl 4- SnCl4 
2AgNO3 4- Zn = 2Ag 4- Zn (NO3) 2 
Zn 4- NaOH 4- H2O = NaHZnO2 4- H2. 
54 CHAPTER VIL 
CHEMICAL NOMENCLATURE. 
The common names of most familiar substances, given before the 
development of chemistry as a science, usually give no clue to their 
composition. However, as the number of known substances has increased 
it has become necessary to use names which are based upon some system. 
The multiplication of terms such as "oil of vitriol" (which is not an oil), 
"copperas" (which contains no copper), "Glauber's salt," "cream of 
tartar," "muriatic acid," etc., would place an almost impossible task upon 
the memory if it were extended to the hundred thousand or more known 
substances. To avoid such a situation as is found in mineralogy, where the 
terms give little clue to the composition of the substance, chemists have 
evolved a nomenclature which is based upon the elements composing the 
compound. This nomenclature has the advantage of being almost inter- 
national, so far as the European languages are concerned. 
Names of the Elements. The only generalization possible in the 
nomenclature of the elements is in regard to* the names of the metals. 
The effort has been to give the Latin neuter suffix, -um, to the names of 
metals, as illustrated by potassium, magnesium, platinum. Where the ele- 
ment was commonly known before the introduction of this systematic 
nomenclature we may have still surviving the common name without this 
suffix, as with iron, silver, copper, gold. In such cases the Latin name, 
from which the symbol is derived, may, however, often be used to designate 
compounds. For example, instead of copper sulfate we may say cupric 
sulfate, from the Latin cuprum. The reason for this use of the Latin 
name will become apparent later. 
The group NH4, which gives compounds analogous to those of the 
metal potassium, is called ammonium. 
The Most Positive Element Named First. It is customary to give 
the name of the most metallic or positive element first in naming the 
compound (corresponding to the order used in writing the formula), as 
is seen in the names of aluminum oxide, sodium chloride, copper sulfide, 
magneisum sulfate, potassium nitrate, hydrogen fluoride, carbon dioxide, 
etc. 
Binary Compounds Designated by Suffix -Ide. When the compound 
consists of but two elements it is customary to suffix -ide to the name 
of the second or less mtaallic element in naming the compound. Thus 
all compounds of a single more positive element with oxygen are called 
oxides, thoseof chlorine are called chlorides, those of phosphorus, phos- 
phides, those of carbon, carbides, etc. 
In a few cases this suffix is used where the compound contains more 
than two elements because of the existence of certain groups acting much 
like a single element. For example, the cyanide group, CN, acts very 
much like the halogens, hence KCN is called potassium cyanide, by 
analogy with KC1, potassium chloride. 
Designation of Compounds Where the Positive Element May Show 
Different Valences. As was pointed out in the chapter on Valence, there 
are many elements showing more than one valence in their compounds, 
so that the simple naming of the constituents as above leaves an ambiguity 
55 as to which of the possible compounds is meant. There are, for example, 
two chlorides, oxides, etc., of mercury, iron, tin, lead, copper, etc., so 
that it is not sufficient, in these cases, to speak of mercury chloride, iron 
chloride, tin oxide, copper sulfide, etc. In most cases the positive elements 
exhibit but two valences in their compounds, and it is possible to distin- 
guish their two series of compounds by the suffixes -ous and -ic to the 
name of the positive element, signifying the lower and higher valences 
respectively. Accordingly we designate HgCl as mercurous chloride and 
HgCl2 as mercuric chloride; As2O3 as arsenous oxide, and As2O5 as arsenic 
oxide. In many cases the common English name of the element does not 
lend itself to this usage, as with the elements iron, copper, tin, lead. In 
such instances the Latin name of the element is used, so that we have 
FeCl2, ferrous chloride, and FeCl3, ferric chloride; Cu2O, cuprous oxide, 
and CuO, cupric oxide; SnS, stannous sulfide, and SnS2, stannic sulfide, 
etc. 
Another way of distinguishing between different compounds of the 
same element is by the use of the Latin and Greek numeral prefixes to 
the names of the elements. These prefixes, up to six, are as follows: 
Latin 
Greek 
1 
Uni- 
Mono- 
2 
Bi- 
Di- 
3 
Ter- 
Tri- 
4 
Quadri- 
Tetra- 
5 
Quinque- 
Penta- 
6 
Sexa- 
Hexa- 
In addition to these we may give the Latin sesqui- meaning one and a 
half. By the use of these prefixes we can often designate the compound 
with less ambiguity than by means of the suffixes -ous and -ic, especially 
when more than two compounds exist. As examples of the use of these 
prefixes we may mention CO, carbon monoxide, and CO2, carbon dioxide; 
CS2, called both carbon bisulfide and carbon disulfide; PC13, phosphorus 
tri-chloride, and PC15, phosphorus pentachloride; Cr2O3, chromium ses- 
quioxide, and CrO3, chromium trioxide; Sb2S3, (di)antimony trisulfide 
(antimonous sulfide), and Sb2S.-„ (di-)antimony pentasulfide (antimonic 
sulfide). 
The prefix per- is used in binary compounds chiefly to denote oxides 
in which the oxygen is related to the oxygen in hydrogen peroxide, H2O2, 
with an average valence of -1. The most important besides hydrogen 
peroxide are barium peroxide, BaO2, and sodium peroxide, Na2O2. It is 
doubtless wise to use the term only with oxides of this type, though this 
is by no means a uniform practice. 
A compound like Fe3O4, which is doubtless composed of FeO, 
ferrous oxide, and Fe2O3, ferric oxide, may be designated as ferrous- 
ferric oxide, though it is usually referred to as "magnetic iron oxide." 
Similarly, Pb3O4 is usually known by its commercial name "red-lead." 
Acids consisting of hydrogen and one other element are designated by 
the prefix hydro- and the suffix -ic, like HO, hydrochloric acid; HN3, 
hydronitric acid; H2S, hydrosulfuric acid. 
Oxy-Acids and Salts. Certain elements capable of assuming more 
than one valence form two or more acids containing oxygen. In such 
cases the significance -ous and -ic in distinguishing higher and lower 
valence is similar to that used with binary compounds. We have H2SO3, 
sulfurous acid, and H2SO4, sulfuric acid; HCrO2, chromous acid, H2CrO4, 
chromic acid; H3AsO3, arsenous acid, and H3AsO4, arsenic acid; HNO2, 
nitrous acid and HNO3, nitric acid. When more than two oxygen acids 
of the same elements exist the prefixes hypo- and per- are used, as 
illustrated in the following table, which shows at the same time the salt 
56 designations corresponding to the different types of acid, using the acids 
of chlorine for illustration: 
Acid 
HC1 Hydrochloric 
HC1O Hypochlorous 
HClOo Chlorous 
HC1O3 Chloric 
HC1O4 Perchloric 
Corresponding salt 
KC1 Potassium chloride 
KC1O Potassium hypochlorite 
KC1O2 Potassium chlorite 
KC1O3 Potassium chlorate 
KC1O4 Potassium perchlorate 
Sulfo-Acids and Salts. There are many cases where sulfur replaces 
oxygen in salts, though the corresponding acids are not capable of existing 
free. These are designated in the same way as the oxy-salts, using the 
prefix sulf(o)-, or sometimes thio-. We have KSbO,, potassium anti- 
monate, and KSbS2, potassium sulfantomonite; KSbO3, potassium anti- 
monate and KSbS3, potassium sulfantimonate. 
In the case of Na2S2O3, which seems to be sulfuric acid with one 
oxygen atom replaced by a sulfur atom, the name thiosulfate is given. 
Complex Halogen Acids and Salts. Similar to the oxy- and sulfo- 
acids we have others containing the halogens, like H2PtCl6, chlor-platinic 
acid, giving salts called chlor-platinates; H2SiF6, fluo-silicic acid (cf 
H2SiO3, silicic acid) ; HAu Cl, chlor-aurous acid, and HAuC14, chlor-auric 
acid. 
Acid and Basic Salts. Many acids have more than one replaceable 
hydrogen atom in the molecule, and many bases have more than one re- 
placeable hydroxyl group in the molecule. Sulfuric acid, for example, 
is called a diabasic acid on account of its ability to neutralize two 
hydroxyl groups. Similarly, H3PO4, phosphoric acid, is a tribasic acid. 
On the other hand, Mg(OH)2, magnesium hydroxide, is a diacid base 
for similar reasons. In such cases it is usually possible to have partial 
neutralization of either acid or base so as to give either an acid salt or 
a basic salt. The following equations represent the formation of such 
salts * 
NaOH 4- HoSO4 = NaHSO4 + H.O 
KOH + H3PO4 = KH2PO4 4- H2O 
2K0H 4- H3PO4 = K2HPO4 4- H2O 
Mg(OH)2 4- HC1 = Mg(OH)Cl 4- H2O. 
The salts produced in these reactions would be called respectively: 
sodium acid (or hydrogen) sulfate, potassium di-hydrogen phosphate, 
di-potassium' hydrogen phosphate, and magnesium hydroxychloride or 
basic magnesium chloride. 
Sometimes a basic or hydroxy-salt may lose water, becoming an 
oxy-salt. With bismuth hydroxide, for example, Bi(OH)3, the progressive 
neutralization with HC1 would give first Bi(OH)2Cl, then Bi(OH)Cl2, 
finally BiCl3. The second of these loses water, becoming BiOCl, a well 
known substance called bismuth oxychloride. In naming these two kinds 
of basic salts chemists are not always careful to distinguish between an 
oxy-salt and a hydroxy-salt, but there is some justification in this for the 
uncertainty in many cases as to whether the water present in the precipi- 
tate is actually combined or not. 
Sometimes this oxy-group plays the part of a radical, existing through- 
out various metathetical reactions. The radical is then usually designated 
by the suffix -yl, as in BiOCl, bismuthyl chloride, CrO2Cl, chromyl 
chloride, UO2C12, uranyl chloride. 
Acid salts may likewise lose water, but the possible types of com- 
57 pounds can best be understood in connection with the various types of 
partially dehydrated acids. 
Partially Dehydrated Acids and Their Salts. If an acid salt like 
KHSO4, potassium acid sulfate, is heated it loses water according to the 
equation: 
2KHSO4 = H2O + K2S2O7. 
The same result is gotten by adding SO3 to K2SO4. With reference to its 
mode of formation this salt is called potassium pyro-sulfate (from the 
Greek pyr, meaning fire). It may also be called potassium bi-sulfate on 
account of the extra SO3 present. This is more evident if the formulas 
of both the sulfate and the bisulfate are expressed as compounds of acid 
and basic anhydrides, as was done in Chapter V for many salts, giving 
K2O • SO3, and K2 • 2SO3, respectively. There are other examples of this 
type. When acid is added to a chromate, as K,CrO4, instead of getting 
an acid salt KHCrO4, there results a bichromate, as represented by the 
equation: 
K2CrO4 + HC1 = (KC1 + KHCrO4) = KC1 + %H2O -j- ^K2Cr2O7. 
There is no great distinction between a bi-salt and an acid salt; in solu- 
tion it is generally impossible to distinguish them, so that the former 
name is applied somewhat indiscriminately. Thus NaHCO3, sodium acid 
carbonate, is often called sodium bicarbonate, though the real bicarbonate, 
which would beNa2C2O-„ does not exist. Similarly KHSOt may be called 
potassium bisulfate. Sometimes it is possib'e to have still more of the 
acid anhydride present, as with K2Cr3O10 and K2Cr4O]3, called potassium 
trichromate and tetrachromate, respectively, or in general, polychromate 
(from the Greek poly, meaning many). This again becomes most obvious 
when the formulas are written in terms of the acid and basic anhydrides, 
giving K2O • 3CrO3 and K2O • 4CrO3. 
With a tribasic acid like phosphoric acid, H3PO4, there exist several 
partial anhydrides, like HPO3 and H4P2O7. The relations between these 
acids is best brought out by the equations 
3H2O + PoO, = 2H3PO4 
Hod + P2O5 = 2HPO3 = 2H3PO4 - 2H„O 
2H2O + 2P2O5 = H3P2O7 = 2H3PO4 - H2O. 
In the first, H3PO4, we have the maximum amount of water, and it is 
called ortho-phosphoric acid (from the Greek meaning straight, direct). 
The second, meta-phosphoric acid (from the Greek meaning after), may 
be derived by the simple abstraction of water from the ortho-acid. The 
third is called pyro-phosphoric acid. The application of these terms to 
several other acids, as in the following table, will make themselves clearer: 
Acids 
Ortho 
Meta 
Pyro 
Boric 
H3BO3 
HBO„ 
H2B4O7 
(also called 
tetra-) 
Silicic 
H4SiO4 
H2SiO3 
(Many poly- 
silicic acids. 
H2O nSiO2) 
Phosphorus 
H3PO3 
(hpo.,) 
H4P.,Ov 
Stannic 
H4SnO4 
H2SnO3 
It is evident that there is no distinction between a pyro-acid and its 
salts and a poly-acid, and where there are several it is not customary to 
use the prefix pyro- with its ambiguous significance, for while there is 
but one pyro-phosphoric acid there are a number of poly-silicic acids which 
require such distinguishing terms as di-meta-silicic acids for H2Si2O5; 
tri-meta-silicic acid for H2Si2O7; di-ortho-silicic acid for H4Si2O6, etc. 
The salts are, of course, designated as for the simpler acids; thus 
Na2H2P2O7 is di-sodium-di-hydrogen pyro-phosphate; Na2B4O7 (borax) 
is sodium tetra-(or pyro) borate; K2SnO3 is potassium metastannate. 
58 Exercises. 
1. How should each of the following substances be named: CaH„, 
H2S, PbS, Ni(OH)2, Ni(OH)3, Hg2SO4) HgSO4, SbCl3, SbCl„ PbO, 
Pb3O4, CuCO3, H2CO3, Hg(NO3)2, Hg(OH)NO3, NH4C1, NH4OH, 
NH4HSO4, KHNH4PO4, SbCl3, SbCl5, SbOCl, NaCl, NaClO, A1C1S, 
A1(OH)3, NaH2AlO3, NaAlO2, SOo, H2SO3 (Cf. HoS04, sulfuric acid), 
NaHSO3, Na2S2O5, SO3, S2O3? 
2. If H3AsO4 is called arsenic acid, what should each of the 
following be called: HAsO3, H3AsO3, H4As2O7, PbHAsO4, KAsS3? 
3. If K2MnO4 is called potassium maganate, what should KMnO4 
be called ? 
4. If H4SnO4 is called ortho-stannic acid, what would you call 
each of the following: H2SnO2, NaHSnO3, NaHSnO2, Na2SnS3? 
5. If KCNO is the only oxygen salt of its type, what should it be 
called? (CN is the "cyanide" group). What should each of the follow- 
ing be called: KCNS, HCNO, KCN, HCN? 
6. The iron in the salts K4Fe(CN)6 and K3Fe(CN)6 is bivalent 
and tri valent, respectively. Which should be called potassium ferro(us) 
cyanide, and which ferri(c) cyanide? 
59 CHAPTER VIII. 
QUANTITATIVE RELATIONS INVOVLING SOLUTIONS OF KNOWN 
CONCENTRATION. 
In quantitative analysis the most useful reactions are usually those 
occurring in solution. In such cases the quantities may be determined, not 
by weighing the pure substance, but by measuring the volume of the 
solution containing a definite amount of the substance. Thus if a large 
quantity of a solution of sulfuric acid is made containing, say, 60 grams 
per liter, the amount of the acid taking part in a certain reaction is very 
simply found by measuring the volume of this solution required in the 
reaction. For example, if 25.0 cc. of this solution were used in a certain 
reaction, then the actual amount of sulfuric acid used would be 0.025 
of 60 gm. or 1.25 gm. Before proceeding on this basis, however, it will 
be profitable to pause and consider whether the amount of the dissolved 
substance had best be expressed in grams per liter. 
Concentration in Terms of Mois per Liter. As a type of the reac- 
tions in solution used in quantitative analysis, we will consider reactions 
in which acids and bases neutralize each other. It has been explained 
already how such equations may easily be written if we are guided 
by the reaction that is fundamental to all neutralization acids and bases, 
i. e., the formation of water from the H of the acid and the OH of the 
base. Such equations may then be interpreted in terms either of grams 
or mols, and correct calculations of weight relations may be made in 
either terms. It is obvious, however, that simpler numerical relations are 
obtained using mols rather than grams. In a large number of cases 
1 mol of acid will neutralize 1 mol of base; in other cases, like the 
reaction between NaOH and H2SO4, 1 mol of acid neutralizes 2 mols of 
base. We find it advantageous, therefore to express concentration in 
terms of mols per liter, and to indicate this briefly by a suitable 
prefix to the formula of the substance. Thus 2M-H2SO4 indicates a 
solution containing 2 mols of sulfuric acid per liter; 0.5M-H2SO4, or % 
M-H2SO4 or M/2 H2SO4 all denote a solution containing 0.5 mol of 
sulfuric acid per liter. Similar designations are applied to solutions of 
bases and salts, and even to dissolved substances in general. 
Titration. Suppose that it were desired to determine the amount of 
KOH in a sample of lye solution. We would take a definite amount of 
the solution, say 4.52 gm, and titrate with a suitable acid of known 
concentration, say 0.5 M-H2SO4, measuring, by means of a buret, the 
volume of acid required to neutralize exactly the lye solution, a suitable 
indicator being used for this purpose. This process is called titration. 
Suppose that 45.80 cc. of the acid is used; then, since 1 liter of 0.5 
M-H2SO4 contains 0.5 mol of H2SO4, 45.80 cc. will contain 0.0458 X 0.5 
mol, or 0.0229 mol of HoS04. Now we see from the equation, 
2KOH 4- H2SO4 = 2H2 + K2SO4, 
that 1 mol of H2SO4 neutralizes 2 mols of KOH, hence, 0.0229 mols of 
H2SO4 will neutralize 0.0458 mols of KOH. Since 1 mol of KOH is 
56.1 gm., or 0.458 mols is 0.0458X56.1 gm., or 2.57 gm. Since, further, 4.52 
gm. of the solution contain 2.57 gm. of KOH, 57.0% of the lye solution 
is KOH. 
60 Standardizing Solutions. In making solutions of known concen- 
tration for titration it is necessary to start with some substance that can 
be weighed, such as benzoic acid, HC7H5O2 (only one of the hydrogen 
atoms has acid properties), or sodium carbonate, Na2CO3. If it were 
desired to make 0.05 M-H2SO4, which contains 0.05 mol or 4.9 gm. of 
H2SO4 per liter, we would make up a quantity of the solution a little 
more concentrated than 0.05 M by dissolving about 3 cc. of concentrated 
acid (density 1.8) in a liter of water. We would next determine accurately 
the concentration of acid in this solution by titrating a known weight of 
Na2CO3. Suppose that it takes 43.0 cc. of our solution to react with 
0.248 gm. of Na2CO3, according to the equation 
Na2CO3 + H2SO4 = H2O + CO2 + 2NaCl. 
Since 1 mol of sodium carbonate reacts with 1 mol of the acid, and since 
0.248 gm. of Na2CO3 is 0.00234 mol, the amount of H2SO4 is also 
0.00234 mol, and as this is dissolved in 43.0 cc. the amount in one liter 
is ^^°^/43 X 0.00234 mols, or 0.0545 mols. This solution may be used 
in titrations just as it is, or the concentration may be adjusted to exactly 
0.05 molal, by diluting with the right amount of water. Since 0.05 mols 
per liter is the same as 0.0545 mols in 1090 cc., if water is added to a 
liter of the 0.0545 molal solution until the volume becomes 1090 cc., then 
the concentration of the resulting solution is 0.05 molal, as desired. From 
this an alkaline solution of desired concentration may then be made up 
by a similar process. 
The Equivalent-. Normal Concentration. It is very convenient, in 
titration, to have various solutions of the same neutralizing power. 
While it is true that solutions of such acids as HC1 and HNO3 of the same 
molal concentration may be used interchangeably in titrating bases, solu- 
tions of molal HC1 and molal H2SO4, for example, could not be used 
indiscriminately, since 1 mol of H2SO4 has twice the power that 1 mol 
of HO has to neutralize alkali. Similarly, while 10 cc. of 0.1 M HC1 
would neutralize 10 cc. of 0.1 M NaOH, it would neutralize only 5 cc. 
of 0.1 M Ba(OH)2. Since, in titrating acids and bases, we are inter- 
ested in the hydrogen of the acid and the hydroxyl of the base, the simplest 
basis for making and using such solutions is the amount of acid hydrogen 
and basic hydroxyl in a liter of the solution. Solutions of NaOH and 
Ba (OH)2 would have the same neutralizing power if equal volumes 
contained, not 1 mol of each, but 1 mol and % mol respectively. The 
amounts of acids and bases to be dissolved in equal volumes of solution are 
the amounts which, in the particular reactions involved, will yield 1 gram 
of acid hydrogen and 17 grams of basic hydroxyl, respectively. This 
amount is called inequivalent, and concentration expressed in equivalents 
is called normal concentration, and denoted by the letter N, in the same 
way that molal concentration is indicated by M. Thus 2N-H2SO4 denotes 
2 equivalents of this acid per liter; 0.5N-KOH denotes 0.5 equivalent of 
KOH per liter. Obviously, whenever the acid and base contain one acid 
hydrogen atom and one hydroxyl group in the molecule, the mol and 
equivalent are the same, and so molal and normal concentrations of such 
acids and bases are the same. 
Summary. By way of summary it may be stated that any kind of 
problem involved in titrating acids and bases may be solved by under- 
standing the three following factors, utilizing them in the connection 
demanded by the problem: 
First, the number of equivalents and the number of grams per mol 
61 of substance, as indicated by the formula; e. g., Ca(OH)2 denotes 1 mol, 
2 equivalents, 74 gr. 
Second, the meaning of the symbols expressing concentration; e. g., 
0.01 M-Ca(OH)2 denotes a solution of 0.01 mol (hence, by the first step, 
0.02 equivalent or 0.74 gr.) of Ca(OH)2 per liter. 
Third, the relation between the amount of acid and base used, which 
is given by the equation, if amounts are expressed in mols or grams, or 
without the necessity of considering the equation if equivalents are used, 
since 1 equivalent of any acid neutralizes 1 equivalent of any base, by 
the definition of the term equivalent. 
Examples. What is (1) the molal concentration and (2) the normal 
concentration when 7.4 gr. of Ca(OH)2 is in two liters of solution? One 
mol of Ca(OH)2 is two equivalents, and weighs 74 gr., hence 7.4 gr. 
is 0.1 mol and 0.2 equivalent. Since this is dissolved in two liters, one 
liter would contain 0.05 mol or 0.1 equivalent, so that the solution would 
be 0.05 molal or 0.1 normal, expressed briefly as 0.05 M-Ca(OH), and 
0.1 N-Ca(OH)2, respectively. 
How many cc. of 0.2 N-HC1 will be required to neutralize 40.0 
cc. of 0.5 N-NaOH? By 0.5 N-NaOH we mean 0.5 equivalent of NaOH 
in 1 liter of solution, hence 40.0 cc. of it, which is 0.040 1, will contain 
0.040 X 0.5 equivalent, or 0.20 equivalent of NaOH. Since one equiva- 
lent of any acid neutralizes one equivalent of any base, the amount of 
acid neutralized by 0.02 equivalent of NaOH is 0.02 equivalent. Since 
0.2 N-HC1 means 0.2 equivalent of it per liter of solution, to get 0.02 
equivalent we would have to take 0.1 liter or 100 cc. 
It was found that the acetic acid in 10 cc. of vinegar was neutralized 
by 14.0 cc. of 0.5 M-KOH solution; how many grams of acetic acid 
were in 100 cc. of the vinegar, the formula of acetic acid being HC2H3O2, 
and one only of the hydrogen atoms in the molecule having acid properties ? 
By 0.5 M-KOH we mean a solution containing 0.5 mol of KOH per 
liter. Since 14 cc. is 0.014 liter, this volume of solution will contain 0.014 
of 0.5 mol or 0.007 mol of KOH. By writing the equation, 
KOH 4- HC2H3O2 = H2O + KC2H3O2, 
we see that 1 mol of the acid neutralizes 1 mol of the base, so that 0.007 
mol of the base will be neutralized by 0.007 mol of the acid, which is 
the amount present in 10 cc. of vinegar. Since 1 mol of acetic acid 
weighs 60 gr., 0.007 mol weighs 0.42 gr. Since this amlount is contained 
in lOcc. of vinegar, 100 cc. would contain 4.2 gr. 
If 0.654 gr. of zinc is completely dissolved in 100. cc. of 0.5 
M-H2SO4, what will be the concentration of the acid after the reaction in 
terms of (1) mols per liter, (2) equivalents per liter? The atomic weight 
of zinc is 65.4, so that 0.654 gr. is 0.01 gram-atom of zinc. From the 
equation of the reaction, 
Zn + H2SO4 = ZnSO4 + H2, 
we see that 1 gram-atom of zinc will use up 1 mol of acid, hence, 0.01 
gram-atom of zinc would use up 0.01 mol of the acid. Now 1 liter of 
0.5 M-H,SO4 contains 0.5 mol, and 100 cc. would contain 0.05 mol. The 
zinc used" up 0.01 mol of this, leaving 0.04 mol of H2SO4 in 100 cc., which 
corresponds to 0.4 mol per liter. Since 1 mol of H2SO4 = 2 equivalents, 
0.4 mol = 0.8 equivalent and the solution may be called 0.8 normal. 
Titration Involving Other Types of Reaction. It will be seen from 
the last example above that solutions of known concentration may be 
used to measure amounts of material in other reactions besides those 
62 where acids and bases neutralize each other. The speed with which 
volumes of solutions can be measured in titrations, the process called 
volumetric analysis, compared with that required for getting the products 
of reactions into weighable form, gravimetric analysis, makes the former 
desirable wherever speed is desired. When neutralization of acids and 
bases does not take place some other means of indicating the "end-point" 
of the reaction is necessary. Where a highly colored substance disappears 
in the reaction this may be used to indicate when the exact amount re- 
quired has been added, any excess making itself evident by reason of the 
color. One such substance frequently used is potassium permanganate, 
KMnO4, a substance with a deep purplish red color, and which is a very 
strong oxidizing agent. As an example of the type of reaction now being 
considered, let us determine how much iron is in an ore, 0.590 gr. of 
which, when dissolved in acid so as to give a solution containing FeSO4, 
required 45.0 cc. of 0.2 M-KMnO4 to change the FeSO4 to Fe2(SO4)3 
according to the equation: 
10 FeSO4 + 2 KMnO4 + 8 H.,SO4 = 5 Fe2(SO4)3+2 MnSO4+ 
K2SO4 4-8 H2O. 
The actual amount of KMnO4 used is evidently 0.045 of 0.5 mol, or 0.0225 
mol. According to the equation 2 m'ols of KMnO4 react with 10 mols of 
FeSO4, so that 0.0225 mol of KMnO4 would react with 5 X 0.0225 mol, 
or 0.1125 mol of FeSO4. Now 0.1125 mol of FeSO4 contains 0.1125 
gram-atom of iron, and since 1 gram-atom of it is 55.8 gm., 0.1125 gram- 
atom would weigh 0.1125-X 55.9 gm. or 0.626 gm. Since, finally, this 
amount of iron was found from 1.016 gm. of the ore, the iron is 61.7% of 
the ore. 
Questions. 
1. Write equations expressing the neutralization of sodium hydroxide 
with hydrochloric acid; sodium hydroxide with sulfuric acid; calcium 
hydroxide with nitric acid; barium hydroxide with sulfuric acid. 
2. Write under each substance involved in the above equation the 
number of mols, grams, and equivalents represented. 
3. How many grams are there in 1 equivalent of each of the follow- 
ing substances: KOH; H2SO4; H,S; A1(OH)3? From your answers 
state how many grams of aluminum hydroxide and sulfuric acid would 
combine to form aluminum sulfate. 
4. How many mols of sulfuric acid will be neutralized by 4 equiva- 
lents of sodium hydroxide? Ans. 2 mols. 
5. If 0.98 gm. of sulfuric acid are in 1 liter of solution, what is 
(a) the molal concentration? (b) the normal concentration? 
Ans. (a) 0.01 M; (b) 0.02 N. 
6. When 3.7 gm. of calcium hydroxide are in 5 liters of solution, 
what is (a) the molal concentration? (b) the normal concentration? 
Ans. (a) 0.01 M; (b) 0.02 N. 
7. (a) How many equivalents of barium hydroxide are in 200 cc. of 
0.1 normal solution? (b) How many mols? (c) How many equivalents 
of hydrochloric acid could this solution neutralize? (d) How many 
grams ? 
8. If 10 cc. of a certain solution of sulfuric acid neutralized 20 
cc. of a normal solution of potassium hydroxide, what was the concen- 
tration of the former in (a) equivalents per liter; (b) mols per liter; (c) 
grams per liter? 
Ans. (a) 2N; (b) IM; (c) 98 gm. per 1. 
63 9. If 25 cc. of 0.2 normal acid neutralized some base, how many 
equivalents of base were there? How many grams of base were there 
if the base was (a) sodium hydroxide; (b) calcium hydroxide? 
Ans. 0.0005; (a) 0.2 gm.; (b) 0.185 gm. 
10. If 10 cc. of N/5 ammonium hydroxide neutralizes 20cc. of 
sulfuric acid, what was the concentration of the latter. Ans. 0.1N. 
11. How many cc. of 0.2 normal acid are necessary to neutralize 
25cc. of 0.5 normal alkali? Ans. 62.5 cc. 
12. How many grams of hydrogen peroxide, H2O2, were present in 
a solution, 10 cc. of which required 40.0 cc. of 0.05 KMnO4 to react 
according to the following equation: 
5 H2O2 + 2 KMnO4+3 H,SO4 = 5 O2 + 8 H2O + K2SO4 + 2 MnSO4? 
Ans. 0.17 gr. 
64 CHAPTER IX. 
THERMOCHEMISTRY. 
There are many chemical reactions in which our interest lies ordi- 
narily not in the substances produced but in the heat evolved or absorbed. 
The burning of fuels is a chemical reaction, but our purpose in burning 
them is not to produce the carbon dioxide and water that usually result. 
These substances we take pains to dispose of by means of chimneys and 
ventilation. Our purpose is rather to obtain the heat or light that 
accompany the reaction. In buying fuels, therefore, the important consid- 
eration is not the relative cost of the various available fuels but the 
relative cost of the heat as obtained from each. The chemist should 
therefore know how the heat of such a reaction may be measured and 
expressed. 
Again, in many reactions of importance a knowledge of whether heat 
is absorbed or evolved helps to determine the most desirable condiitons 
for carrying it out, as will be explained in Chapter XIV. Also in technical 
processes, if heat is absorbed during a reaction this may have to be 
considered in estimating the cost, since the heat thus absorbed may have 
to be supplied by means of fuel or electrical energy. 
The Calorie. The unit in which heat is most conveniently measured 
is the calorie (abbreviated cal.). This is the amount of heat required 
to raise the temperature of 1 gram of water 1°C.* Since many heats 
of reaction involve many thousand calories it is also customary to use a 
larger unit, the kilogram calorie (abbreviated Kal.). This is the heat 
required to raise the temperature of 1 kilogram of water 1°C. Obviously 
1 Kal. = 1000 cal. 
Experimental Determination of Heats of Reaction. The amount of 
heat involved in a chemical reaction is determined in an apparatus known 
as a calorimeter. It consists of a vessel holding perhaps a liter of water, 
provided with a stirrer and a more or less delicate thermometer. This 
vessel must be well surrounded by some effective heat insulator so that 
the temperature of the water within shall be affected only to the least 
possible extent by the room temperature. This insulation may be fur- 
nished by felt, a water jacket, or, often best of all, by a vacuum jacket 
as in a Dewar flask. (This is the vessel that has been adapted to house- 
hold and picnic use and sold under such names as "Thermos," "Caloris," 
etc.). The reacting substances may be introduced into the calorimeter in 
various ways. When the heat of solution of a solid is to be measured 
the substance may be sealed in a thin glass bulb immersed in water in 
the calorimeter. After uniform temperature has been reached the bulb 
* In engineering practice in England and America it is customary to use the 
"British Thermal Unit" (B. T. U.) which is the amount of heat required to raise 
the temperature of 1 pound of water 1° Fahrenheit. 1 B. T. U. = cals. In very 
accurate work it is necessary to specify the temperature of the water, for the amount 
of heat required to produce the same elevation in temprature is not the same 
at all temperatures. We thus have the "zero calorie," the "15° calorie" and the 
"mean calorie," which is the average value between 0° and 100°C. 
It is possible also to express the quantity of heat in terms of other energy 
units, since the relation between these different kinds of energy is known. A unit 
frequently used is the joule, which is 107 ergs, or 1 volt-colounb, or 1/4 cal. 
65 is broken and the solid allowed to dissolve. If the reaction is one like the 
precipitation of an insoluble substance from dilute solutions, or the 
neutralization of an acid by a base one liquid may be enclosed in a thin 
glass vessel or in one with a valve, and immersed in the other liquid. 
The liquids may then be mixed at the proper time. If the reaction is one 
taking place at a high temperature, it may be carried on within a "bomb" 
immersed in the water of the calorimeter. A small weighed portion of 
the material under investigation is placed within this bomb which is then 
closed and filled, through a tube in the top, with oxygen under high 
pressure. It is then placed in the calorimeter and when uniform tempera- 
ture is attained the combustion is brought about by means of an electrical 
connection through the lid of the bomb. 
The amount of heat liberated by the reaction is measured by the rise 
in temperature of the calorimeter contents and the heat capacity of 
the latter. If water alone were involved the number of calories liberated 
would be found simply by multiplying the rise in temperature by the 
weight of the water. Various other substances are present, however, 
such as the stirrer, the containing vessel, the thermometer, etc., so that 
their power of absorbing heat compared with that of water must be 
measured or calculated and added to the weight of the water as the 
"water equivalent" of the calorimeter. This may be measured by allowing 
a reaction evolving a known amount of heat to take place in the calori- 
meter, or else calculated, by knowing the heat capacities or specific heats, 
together with the weights of the various substances in the calorimeter. 
The Numerical Representation of Heats of Reaction. In expressing 
the results of thermochemical measurements it is usual to include in the 
equation the number of calories evolved or absorbed. The amounts of 
substances represented by the equation are understood to be expressed 
in mols. Thus the equation, 
C 4- O2 = CO2 4- 96500 cal., 
indicates that when 1 mol (12 gm.) of carbon burns with 1 mol of oxygen 
to form 1 mol (44 gm.) of carbon dioxide, 96500 cal. of heat are evolved. 
This amount of heat is also called the "heat of formation" of carbon 
dioxide from its elements. 
Similarly, the equation, 
N2 + O2 = 2NO - 43200 cal. 
means that when 1 mol (28 gm.) of nitrogen reacts with 1 mol (16 gm.) 
of oxygen to form 2 mols (30 gm.) of nitric oxide, 43200 cal. of heat 
are absorbed. This reaction might equally well be written 
N2 4- O2 4- 43200 cal. = 2NO. 
The ordinary rules of algebra apply to the sign of the heat if the number 
is thus transposed. Either w^ay of writing the equation signifies that nitro- 
gen, oxygen and heat are all used up in the formation of nitric oxide. 
The term exothermic is often applied to reactions which evolve heat 
and endothermic to those which absorb heat. It may be mentioned that 
the former predominate in reactions occurring at ordinary temperatures 
and the latter in those occurring at very high temperatures, in the electric 
furnace. 
Effect of the State of the Reacting Substances. The heat of a reac- 
tion depends upon the temperature at which it is carried out and upon 
the state of the substances involved. For example, in the union of 
hydrogen and oxygen to form wrater, it makes a difference whether the 
water produced is at a temperature at which it is liquid or vapor. Since 
66 the condensation of steam to liquid takes place with the evolution of heat 
it is evident that where the final product is liquid the heat of formation 
will be greater than when it is a gas. To make this clear we will need 
to specify in the equation the state of the substance, which we would do 
in this case as follows: 
2H2 + O2 =2H2O( gas) + 67800 cal., or =2H2O (liq.) +6840 cal. 
Or, again, we might find it desirable to express the states of all substances 
in the equation where ambiguity might otherwise arise, as in the equation, 
Hg(liq.) + Cl2(gas) =HgCl2(solid) + 53300 cal. 
It makes a difference also whether or not a substance is in solution, 
and what its concentration in the solution is. We might, therefore, have 
an equation like the following: 
NaOH(0.1 molal) + HC1 (0.1 molal )= H„O + NaCl (0.05 molal) + 
13700 cal. 
Where the concentration is small and does not need to be specified exactly 
it may suffice to use the abbreviation aq. after the formula of the dissolved 
substance, which signifies the presence of a large amount of water. (From 
the Latin aqua, water.). 
In a gaseous reaction, in which the number of molecules of gaseous 
substances changes, the heat will be somewhat different if the reaction 
is carried out in a closed vessel, so that the volume remains constant, 
from what it will be if the reaction is carried out so as to keep the 
pressure constant. Since a gas does work in expanding against an ex- 
ternal pressure, heat must be absorbed during the expansion in order to 
maintain the original molecular velocity and hence the original pressure. 
This behavior finds illustration in the fact that a ball thrown against a 
yielding surface rebounds with less speed than when thrown against a 
rigid surface. The heat of a chemical reaction at constant pressure will 
therefore be greater or less than that of the same reaction at constant 
volume according as the number of gaseous molecules decreases or 
increases during the reaction. The difference can be calculated from the 
work of expansion, a change of volume of a liter against a pressure of 
one atmosphere being equal to 24.2 cal. 
In the reaction 
2CO + O2 = 2CO2 + 135000 cal. (constant volume, 20°C.) 
all of the substances being gases, there is a decrease in the number of mols 
from 3 to 2. If, therefore, this reaction were allowed to take place so 
that the pressure remained constant, which would require a decrease in 
the volume, work would be done on the reacting mixture by the atmos- 
phere, which would appear as additional heat of reaction. The amount of 
this additional heat can be calculated as follows. The volume of one mol 
of gas at 20°C (293°A) is 293/273 X 22.4 1, or 24 1. Since the work done 
when the volume changes by 1 liter against a pressure of 1 atmosphere it 
may be taken as 24.2 cal., the heat evolved in this reaction when it takes 
place at constant pressure is greater by 24 X 24.2 cal. or 580 cal. than 
it is at constant volume. 
Indirect Determination of Heats of Reaction. In many cases the 
heat of a reaction cannot be determined directly on account of the diffi- 
culty of carrying out the reaction in a calorimeter. In such cases it is 
usually possible to determine it indirectly. Moreover, even though the 
reactions may be easily carried out in the calorimeter it is unnecessary 
to measure the heats of all, since some may be calculated from others. 
To do this we use a well-founded law stating that all heats of reaction 
depend only on the state of the initial or final substances and not upon 
67 the steps into which the reaction may be divided. For example, ZnSO4 
might be formed from its elements in either of the following ways: 
First Way. 
Zn 4-%O2 = ZnO 4-85 Kal. 
S 4- O2 = SO2 4- 70 Kal. 
SO2 4- %O2 = SO3 4- 22 Kal. 
ZnO 4- SO3 = ZnSO4 4" 53 Kal. 
Total: Zn 4- S 4" 2O2 = ZnSO4 4~ 230 Kal. 
Second Way. 
Zn 4~ S = ZnS 4~ 40 Kal. 
ZnS 4- 2O2 = ZnSO4 4- 190 Kal. 
Total: Zn 4-S 4- 2O2 = ZnSO4 4- 230 Kal. 
The total result in either case is the formation of ZnS.O4, from Zn, S, 
and 2O2 with the evolution of heat. Our law tells us that the total heat 
liberated is the same in either case, as shown by the above figures. 
As a consequence of this, thermochemical equations may be added 
or subtracted like any algebraic equation. Suppose that the following 
heats of reaction have been measured: 
C + O2 = CO2 + 96.5 Kal, and (1) 
2CO + O2 = 2CO2 + 135.0 Kal, (2) 
and we wish to learn the heat of the reaction 
2C + O2 = 2CO. 
All that is necessary is to eliminate CO2 from equations (1) and (2) by 
multiplying (1) by two and subtracting (2) from it, obtaining 
2C - 2CO + O2 = 58 Kal, 
or, transposing, 
2C + O2 = 2CO + 58 Kal. 
If, on the other hand, we eliminate O2 by subtracting (1) from (2), 
we obtain the equation for a different reaction, 
2CO= C + CO2 + 38.5 Kal, 
which, since it usually reacts in the reverse direction, we may prefer 
to write 
C 4- CO2 = 2CO - 38.5 Kal. 
Again, it is not possible to synthesize H2O2 aq. from its elements in 
the calorimeter, but the heat that would be evolved if the reaction were 
possible may be calculated from the following: 
H2O2 aq. = HoO + k2O2 + aq. 4- 23.0 Kal. 
H2 + %O2 = H2O + 68.4 Kal. 
Subtracting the first from the second'to eliminate H2O we obtain 
H2 -f- z^O2 - H2O, aq. =-%O2 - aq.- 45.4 Kal. 
or, transposing to make the signs positive, 
H2 4- O2 4- aq. = H2O2 aq.- 45.4 Kal. 
From this we learn that H2O2 is an enthodermic compound, formed from 
its elements by the absorption of heat. 
Fuel Value of Foods. The heat evolved by the combustion of dry 
food materials in a calorimeter bomb serves for an approximate valuation 
of their ability to furnish the heat and muscular energy needed by the 
body. Accordingly the amount of food materials needed per day can be 
estimated roughly by the aid of a knowledge of the fuel values of the foods. 
68 An average adult needs from 2000 to 3500 cal. per day, according to 
the amount of muscular activity. This makes it possible, for example, 
to calculate roughly the quantity of food needed and the relative costs of 
different foods, which aids in reducing costs of food in the home and in 
institutions, and furnishes a valuable means of determining the least 
weight of food necessary for a camping or exploring expedition. It may 
be of interest to have the few figures given in the following table: 
Approximate number 
Foods- cal. per pound. 
Starchy foods, cereals, flour, rice, beans, etc 1650 
Sugar 1860 
Fats and oils 3650 
Cheese 2000 
Chocolate 2850 
Milk, evaporated, unsweetened 800 
Milk, condensed, sweetened 1500 
Dried fruit 1300 
Fish 1000 
Meat 1600 
Exercises. 
1. What is the distinction between heat and temperature? 
2. Calculate the heat of the reaction 
CaO A CO2='CaCO3 
from the following: 
COo A Ca(OH), = CaCO3 A H2O A 30.5 Kal, 
CaO A H2O = Ca(OH)2 A 11.5 Kal. Ans. 42.0 Kal. 
3. Calculate the heat of the reaction 
Zn -p Cl2 A aq. = ZnCl2aq. 
from the following determinations: 
Zn A 2HC1 aq.= ZnCl2 aq. A Ho A 36 Kal, 
H, ACl2=2HCl(gas)A44 Kal, 
HCl(gas) A aq. = HC1 aq. A 17 Kal. 
4. (a) Calculate the heat of the reaction 
N2 A 3H2 = 2NH3 
from the following: 
4NH3 A 3Oo = 2N2 A 6H2O(liq.) A 364 Kal, 
H2 A %O2 = H2O(liq.) A 68.4 Kal. Ans. 23.2 Kal. 
(b) The above figures are for constant pressure; would the heat of 
formation of NH3 be greater or less at constant volume? 
5. Given that C A O2 = CO2 A 96.5 Kal, and that anthracite coal, 
considered as 95% carbon costs $10 per ton (=1000 kg.), determine the 
cost of 1000 Kal. from the burning of this coal. Ans. 0.13 cents. 
6. Considering "water gas" to consist of CO and H2 only as indicated 
by the equation of its formation, 
C A H2O = CO a h2, 
and given the heats of combustion of CO and H2, i. e, 
CO A %O2 = CO2 A 68.0 Kal, 
H2 A %O2 = H2O A 68.5 Kal, 
calculate the cost of 1000 Kal. from this gas if it costs 90 cents per 1000 
cu. ft. (=28300 1.) when measured at 20°C and 1 atm. pressure. 
Ans. 1.12 cents. 
7. If kerosene were pure C10H22, whose heat of combusion is 1626 
69 Kal., and costs 15 cents per gallon (3.781.) and has a density of 0.79 gm. 
per cc., calculate the cost of 1000 Kal. obtained by burning kerosene. 
Ans. 0.44 cents. 
8. From the table on p. 69 make out a "grub-list" for two men on a 
ten-day walking tour, endeavoring to get the minimum weight consistent 
with variety and palatibility. 
70 CHAPTER X. 
BEHAVIOR OF SUBSTANCES DISSOLVED IN WATER. 
IONIC THEORY. 
Such a large proportion of the reactions with which we concern 
ourselves take place in water solution that the behavior of substances 
dissolved in water is a subject of great importance. We will consider 
a number of the properties and reactions of substances in aqueous solution, 
and see what conclusions may be drawn as to the nature of these solutions. 
Abnormally Great Lowering of the Freezing Point. It has been 
pointed out in Chapter III that the molecular weight of substances may 
be determined by two essentially different methods. One depends upon 
the property of a mol of any substance in the gaseous form to occupy 
22.4 liters at 0°C and 1 atmosphere. The other depends upon the ability 
of 1 mol of any substance, when dissolved in a given amount of a certain 
solvent, to exert the same effect upon the escaping tendency of the solvent, 
as measured by the freezing point, boiling point or vapor pressure. For 
example, we may define a mol, experimentally, as the number of grams 
of substance which, in the vapor state at 0°C and 1 atmosphere, will 
occupy 22.4 1., or else as the number of grams which, dissolved in 1000 
gr. of water, will lower the freezing point of the water 1.86°. If we 
determine the effect of various substances on the freezing point of water 
we find that a large number of substances show concordant behavior. 
Other substances, including acids, bases and salts, show an abnormally 
great lowering of the freezing point. This is illustrated in the following 
table: 
Freezing points of 0.1 
Substance molal solution. 
Sugar 0.188 
Glycerine 0.186 
Alcohol 0.184 
NaCl 0.348 
HC1 0.356 
KNO3 0.330 
BaCE 0.478 
Na2SO4 0.459 
It will be seen that the first three substances in the table behave as 
would be expected, but that the others give an abnormally great lowering 
of the freezing point. There are two ways of accounting for this abnorm- 
ality. Either the law of the uniform molal lowering does not hold 
uniformly, or else acids, bases and salts dissociate into two or more 
independent parts when dissolved in water. If this law still holds, then 
1 mol of NaCl, HC1, or KNO3, when dissolved in a large quantity of 
water, seems to yield nearly two mols of dissolved substances, while 1 
mol of BaCl2 or of Na2SO4 seems to yield nearly 3 mols of dissolved 
substances. Now we have, with some substances, examples of an 
abnormally great molal volume when in the vapor state. When 1 mol 
of NH4C1 is vaporized at 1 atmosphere and 35O°C, instead of occupying 
a volume 623/273 X 22.4 liters, or 51.2 liters, it occupies twice that volume. 
In this case it is possible to show that the vapor contains HC1 and NH3, 
uncombined, so that dissociation of the ammonium chloride molecule into 
two other molecules has actually taken place. Before abandoning, there- 
71 fore, the similar conclusion in connection with the effect of dissolved 
substances on the freezing point of the solvent, let us see whether there 
are any other phenomena which indicate that acids, bases and salts are 
dissociated into simpler parts when dissolved in water. 
If such dissociation occurs what can be the nature of these resultant 
mols? Does a solution of NaCl contain sodium and chlorine? The 
properties of elementary sodium and chlorine are evidently not those of 
a solution of sodium chloride, sodium being a metal that reacts violently 
with water, and chlorine being a very corrosive poisonous substance whose 
aqueous solution has a green color. When the dissociation theory was 
first advanced, some chemists made such objections with much effect. 
Let us, however, consider some oher properties of these solutions to see if 
we can deduce a more plausible idea of the nature of the products of the 
dissociation indicated by freezing point lowering. 
Independent Migration of Ions in Electrolysis. It has been found 
that all dissolved substances that give the abnormal effect on the freezing 
point of water also give solutions which conduct the electric current, 
depositing materials at the electrodes, or else dissolving the electrodes. 
On account of this behavior they are called electrolytes. Those conducting 
very well, the strong acids and alkalis, and most salts, are called strong 
electrolytes. The only satisfactory way to account for this conducting 
power is to assume that the solution contains charged substances which 
travel towards the positive and negative poles respectively; and that on 
their arrival at the electrodes their charges are neutralized, so that 
uncharged substances are there deposited. It is indeed possible to detect 
this migration experimentally. If, for example, we have a U-tube contain- 
ing, say, dilute copper sulfate in the bottom with dilute potassium nitrate 
carefully superimposed, as shown in Fig. 10a, and allow an electric current 
to pass through the solution for a time, we will find that the blue color 
characteristic of copper salts travels towards the negative pole, while the 
sulfate radical travels towards the positive pole, as might be shown by 
testing layers of solution with a solution of barium chloride. The state 
of affairs after the passage of current is depicted in the figure at b. 
Whenever electrolysis takes place it is possible to show such independent 
migration of the two parts of the salt, acid or base by appropriate experi- 
ments similar to the above. 
The Dissolved Ions Carry Definite Electric Charges. The salt, acid 
or base is thus found to consist of two parts, one of which travels to 
the negative pole, or kathode, the other to the positive pole or anode. 
On account of their migration these parts are called ions, from the 
Greek word meaning to go. The ion which travels toward the kathode is 
called the kation, the one traveling towards the anode, the anion. On 
account of this migration in electrolysis we are led to conclude that the 
products of dissociation are charged electrically, the kation being positive 
and the anion negative. The difficulty of imagining such substances in 
solution as sodium and chlorine is overcome by this discovery of positively 
charged sodium, called sodium ion, a totally different substance, and 
negatively charged chlorine ion, a substance quite different from free 
chlorine. The ions must not be confused, therefore, with the uncharged 
metals or non-metals, or radicals, which are quite different substances, as 
seen by their fundamfentally different properties. Thus copper ion, a 
greenish-blue substance existing only in solution, giving precipitates with 
dilute sodium hydroxide or carbonate, etc., is as different from copper, 
a reddish, rather heavy metal, as both are from the third substance, copper 
bromide, a brown salt. 
72 Fig. 10 
We can prove further that copper ion and copper differ by a positive 
charge by actually changing one into the other by adding or removing 
the charge. If the blue solution of a copper salt is electrolyzed, copper 
deposits on the cathode, where the positive charge on the copper ion is 
neutralized. Again, if copper is made the anode ,or positive pole, in a 
solution of a salt such as sodium chloride or sulfate, the postive charge 
sent to the anode changes the copper to copper ion which goes into 
solution giving it a blue color. 
The amount of the charge on the various ions is not always the 
same, but a simple and important relationship exists between the charge 
and the amount of substance. To deposit one gram-atom of any univa- 
lent substance, like hydrogen, sodium, silver, chlorine, etc., requires 96540 
coulombs of electricity.* The importance of this quantity is therefore very 
great, and it is called a -faraday of electricity. For exmple, since 108 gm. 
of silver equal a gram-atom, to deposit 1 gm. would require of 96540 
coulombs, and to deposit 5 gm. would require 5/10s of 96540 coulombs or 
4470 coulombs. If this were done by a current of two amperes (2 
coulombs per second) the time required would be 2235 seconds, or 37 
minutes,, 35 seconds. However, to deposit 63.6 grams of copper, from 
ordinary copper salts, 64.5 grams of zinc, etc., requires twice as much 
electricity, or two faradays. Similarly, certain other ions require three 
* A coulomb is the amount of electricity carried by a current of one ampere 
in 1 second; e. g., 10 amperes flowing for 1 minute would amount to 600 coulombs, 
and if 1000 coulombs are delivered in 50 minutes, the current strength is 20 amperes. 
73 faradays per gram-ion. We will therefore indicate in the formula of the 
ion the amount of electricity carried by it, and designate the foregoing 
ions by the following formulas: H+, Cl-, Ag+, Cu++, Zn++, where the 
number of charges written indicates the number of faradays of electricity 
per gram-atom, and the sign indicates whether the charge is positive or 
negative. Since we start with an uncharged salt, the total charges on the 
ions produced when it dissolves in water must add up to zero. 
CuSO4 = Cu++ 4- SO4" 
HC1 = H+ 4- Cl- 
KC1 = K+ 4- Cl- 
KOH = K+ 4- OH- 
KoSO4 = 2K+4- SO4-- 
Cu (NO3) 2 = Cu++ 4- 2NO3- 
In general the metal constituent of a salt or base and the hydrogen 
of acids yield the kations, or positive ions, while the remaining portion 
of the compound, the hydroxyl (OH) of bases or the "acid radical'' of 
acids and salts yield the anions or negative ions. The signs of the electric 
charges correspond to and justify the differentiation into positive and 
negative constituents made in Chapter VI in assigning algebraic signs 
to valences. The valence must be identical both as to magnitude and 
sign with the charge on the ion. Hydrogen in practically all of its com- 
pounds has a valence of -Ch and when these compounds dissociate on 
dissolving in water they yield hydrogen ion with one positive charge, H+. 
Chlorine in the metallic chlorides has a valence of -1, both from the 
point of view of its relation to hydrochloric acid, HC1, and also from 
the formula of chloride ion, Cl-. In A1C13 aluminum is considered 
trivalent, either from the fact that it has replaced 3 hydrogen atoms, or 
from the formula of the ion AF++ which results on dissolving the chloride 
in water. 
The formulas of the two ions formed from any electrolyte on disso- 
ciation and the formula of the undissociated substance are connected so 
that any two of them will serve to predict the third. Thus, knowing the 
formula of sulfuric acid, H2SO4, and that of hydrogen ion, H+, we can 
determine the formula of the sulfate ion by writing the equation for 
the dissociation, remembering that the total amount of positive and 
negative electricity in the solution must be the same, since the pure sub- 
stances which form the ions on dissolving are uncharged. Taking away 
two positive hydrogen ions from H2SO4 would leave therefore SO4--. 
Knowing, now, the formula of sulfate ion, and given that of another 
sulfate, like ferric sulfate, Fe(SO4)3, we get 
Fe2(SO4)3 = 2Fe+++ 4- 3SO4--. 
Again, if we know the following ionic formulas, K+, Ca++, Fe+++, 
SO4--, PO4---, by remembering that the solid compounds formed from 
them would be uncharged, we can write the formulas of the compounds 
of these ions, i. e., K„SO4, K3PO4, CaSO4, Ca3(PO4)2, Fe0(S04)3 and 
FePO4. 
The Atom of Electricity. We see that the amount of electricity com- 
bined with any gram-atom in an ion is always some simple multiple of 
the faraday. Furthermore, we have different ions of the same element, 
in some cases, in which the number of faradays per gram-atom in one 
is a simple multiple of what is in the other. This is exemplified by 
Fe++ and Fe+++, Hg+ and Hg++, Sn++ and Sn++++, Cu+ and Cu++. The two 
ions of each metal are radically different substances. We thus have a 
law corresponding to the law of simple multiple proportions by weight, as 
74 explained in Chapter III, and just as that law led logically to the atomic 
theory, so the similar law of simple multiple proportions of electricity 
per gram-atom of substance leads logically to an atomic theory for elec- 
tricity, a theory which has been abundantly confirmed by independent 
evidence. 
We way thus regard the number of plus or minus signs written in 
the formulas of ions as indicating the number of atoms of electricity, 
positive or negative, in a single ion. We will find later in Chapter XVI, 
on the Electron, that there seems to be but one kind of electricity, 
negative electricity, involved in these changes, and that a positive electric 
charge on an atom is the result of the removal of negative electricity. At 
present it seems that we must regard an ion like chloride ion as a com- 
pound of a chlorine atom with an atom of (negative) electricity, called an 
electron, and a positive ion like that of silver as the result of removing an 
atom of electricity, an electron, from an atom of silver. An atom of 
metallic silver is therefore regarded as a combination of electrons with an 
equivalent amount of positive electricity. In fact it is these electrons, 
free to move through the metal, that are supposed to explain the electric 
conductivity of metallic silver. As confirmation we may mention that 
metals which are good conductors, in which the electrons are held loosely, 
are the ones in general which are most easily changed to positive ions in 
solution. If we adopt the symbol (-) to denote the electron we may 
represent by equations such as the following what is supposed to happen 
when ions change into elements or into other ions, etc.: 
Cu = Cu++2 (-) 
Cl2-p2(-)=2Cl- 
Fe+++ + (-)= Fe++ 
2OH- = 2(-)4-H2O + %O2. 
Having presented evidence that electrolytes are dissociated into charged 
ions in dilute aqueous solution, and having found how their formulas 
should be written in order to express their electrical behavior, let us 
examine other properties of these solutions to see whether there is general 
confirmation of the conclusions thus far reached. 
The Properties of Dilute Aqueous Solutions of Strong Electrolytes 
are the Sum of Independent Sets of Properties, Hence Independent Sub- 
stances Are Present. It is found that all of the properties of dilute 
solutions of strong electrolytes are the sum of independent sets of prop- 
erties belonging to the different kinds of ion present. For example, the 
colors of salts of a colored with a colorless constituent may be different 
when the salts are in the solid state, but the color becomes the same 
when the substance is dissolved in a large amount of water. Solid CuCl2, 
CuBr2 and Cu(NO3)2 are colored green, brown and blue, respectively, but 
in dilute aqueous solution all assume the same greenish blue color. Since 
we recognize substances by their properties we conclude that we have 
the same substance. Again, the electrical conductivity, refractive index, 
density, etc., of these solutions depend only on the ions present, and not 
upon the substances giving the ions. If, say, 0.1 mol of NaCl and 0.1 
mol of KNO3 are dissolved together in a liter of water, the properties 
of the solution are the same as if the solution were made from. 0.1 mol 
of KC1 and 0.1 mol of NaNO3 dissolved in the same amount of water. 
Most striking of all, perhaps, are the chemical reactions. It is not 
possible to predict the properties of gaseous HC1 from the properties of 
hydrogen and those of chlorine, but it is possible to predict the properties 
of a dilute aqueous solution of HC1 from the properties of all other acids, 
75 on the one hand, and those of all soluble chlorides on the other hand. 
The hydrogen ion in this solution reacts in the same way and has the 
same properties as the hydrogen ion obtained from any other acid, as 
explained in Chapter V. It affects indicators in the same way, it reacts 
with carbonates, dissolves zinc, tastes sour, just as it always does unless 
the other ion present exerts some complicating effect on these tests. 
The chloride ion gives precipitates with the same reagents, such as 
AgNO3 or HgNO3 solutions, as are obtained from all other chloride 
electrolytes. Knowing the properties of chloride ion, as obtained from 
som'e chloride, and knowing the properties of some kation, say barium 
ion, it is possible to predict the properties of dilute solutions of barium 
chloride, without the aid of direct experiment. This makes it possible 
to simplify enormously the labor of learning the chemical reactions of 
electrolytes upon each other. It is necessary only to know the reactions 
given by the important ions in order to predict the reactions of the 
enormous number of electrolytes which might be obtained oy combinations 
of these ions. Instead of remembering the action of a large number of 
sulfates individually upon all soluble barium salts it suffices to remember 
that barium ion, Ba++, reacts with sulfate ion, SO4-", to give a very 
soluble white precipitate of barium sulfate, BaSO4. Such generaliza- 
tions cannot be made with any degree of assurance with substances that 
are un-ionized. Alcohol and phenol (carbolic acid) both contain the 
hydroxyl radical (OH), their formulas being C2H5OH and C6H5OH 
respectively, but these are not appreciably ionized, and when dissolved 
in water react very differently, having no set of properties in common. 
Again, it is not correct to conclude that silver nitrate is a "test for 
chlorine" as is sometimes stated. A substance like chloroform, CHC13, 
will give no such test unless it is heated with silver nitrate for a long 
time. A solution of potassium chlorate gives no precipitate with silver 
ion, Ag+, because, although it contains chlorine, it yields no chloride ion, 
Cl-, but, instead, chlorate ion, C1O3~. 
The value of the ionic theory is very great in reducing equations to 
their simplest terms and focusing the attention on the essential reaction 
taking place. 
Heats of Reaction in Dilute Solutions of Strong Electrolytes Depend 
on the Reacting Ions Only. Heats of reaction between ions in dilute 
solution are independent of the nature of any other ions present. This 
is not true for substances not in solution. Solid NaOH and solid KOH 
do not give the same heat of neutralization with concentrated sulfuric 
or nitric acid, but when both acid and base are dilute, the amount of 
heat liberated is always the same, 13700 calories per equivalent. Since 
the heat liberated is the same the reaction liberating it is likely to be the 
same, and the ionic theory indicates that this should be the case. Instead 
of writing 
NaOH + HC1 = H2O and BaCl2, and 
KOH 4- HNO3 = H2O + KNO3, 
for the reaction in dilute solution, which would lead us to expect different 
heats of reaction, we should, according to previous evidence, write, 
Na+ 4- OH- 4- H+ 4- Cl" = H2O 4" Na+ 4- Cl- 4- 13,700 cal., and 
K+ 4- OH- 4- H+ 4- NO3- = H2O 4- K+ 4- NO3- 4- 13,700 cal. 
The ionic form of the equations indicates that nothing happens to anything 
except H+ and OH-, which unite to form H2O; the other ions, being 
unaffected, may be cancelled from the equation, leaving simply 
76 as the fundamental and only reaction in both cases. 
The same would apply to precipitations. Any soluble barium salt 
would give the same heat effect per mol with any soluble sulfate, since, 
no matter what other ions are present in the solution, the essential 
reaction is 
Ba++ -f- SO4" = BaS04 -|- 5600 cal. 
If, however, Ba(OH)2 and H2SO4 solutions were used the heat of reac- 
tion would, of course, be no longer the same as above, for then we would 
have occurring simultaneously two independent reactions, each with its 
own heat effect, and the ionic equation 
Ba++ + 2OH- -ff 2H+ +SO4~ = BaSO4 + 2H2O + 2 X 13700 cal. 
+ 5600 cal. 
shows that nothing should be cancelled out. 
The Degree of Ionization of Strong Electrolytes.. We have seen the 
evidence that the substances we call strong electrolytes are highly ionized 
in aqueous solution, and we may ask the question how completely does 
this dissociation take place? The following table shows the fraction of 
the solute ionized (the "degree of ionization") according to the evidence 
of conductivity and freezing point. In calculating from electrical con- 
ductivity the idea used is to compare the actual conductivity with that 
calculated on the assumption of complete ionization. We have not space 
to go into the explanation of this method. In interpreting the freezing 
point measurements the idea is to compare the actual lowering with that 
which would be expected if no ionization took place, and to calculate the 
increase in the number of mols due to dissociation from the abnormality 
in the freezing point lowering. It will be seen that the degree of ioniza- 
tion calculated by these two methods is fairly concordant. The differences 
are probably due, in part to the fact that the law for the freezing point 
lowering is not an exact law and, in part, to ambiguity as to just how far 
apart the ions must be before they are to be considered dissociated. Some 
molecules pass for dissociated according to one method which do not 
according to the other. For all practical qualitative purposes in predicting 
the results of chemical reactions, as will be done in later chapters, it 
suffices to Consider these substances as completely ionized. 
H+ + OH- = H2O + 13,700 cal. 
Fraction ionized in 
0.05 molal solution. 
By freezing point 
By electrical 
Substance. 
lowering. 
conductivity. 
HC1 
0.93 
0.94 
HNO3 
91 
.94 
NaCl 
89 
.88 
KNO3 
84 
.88 
Na2SO4 
74 
.70 
CaCl2 
1 .82 
.76 
Weak Acids and Bases Are Less Ionized Than Strong Acids and 
Bases in Solutions of the same Concentration. There are some acids and 
bases which show the properties of hydrogen and hydroxide ions, 
respectively, to a much less extent than do strong acids and bases in 
solutions of the same concentration. Solutions of HC1, HNO3, and 
H2SO4, at the same concentration, say 0.01 normal, would have about 
the same sour taste, but 0.01 normal acetic acid, HC2H3O2, would not taste 
nearly so sour. The first three acids would show almost identical colors 
77 with an indicator like methyl violet, sensitive to that degree of acidity, 
but it would require approximately normal acetic acid to show the same 
color. The same weakness on the part of acetic acid is shown in power 
to conduct current. While, therefore, the properties of acetic acid 
indicate that it yields hydrogen ion in water, they also indicate that only 
a little of the hydrogen of the acid is ionized. Again, while the effect of 
acetic acid on the freezing point of water is greater than that of an 
un-ionized substance like sugar, it is not double the latter as is approxi- 
mately the case with an acid like hydrochloric acid. Its heat of neutrali- 
zation is no longer 13700 cal., as with a strong acid, but has a different 
value and includes the heat of dissociation of the acid during its neutraliza- 
tiqn. 
There are a number of such weak acids, including carbonic acid, 
H2CO3, boric acid, H3BO3, hydrogen sulfide (hydrosulfuric acid), H2S, 
silicic acid, H2SiO;( (also many poly-sicilic acids), nitrous acid, HNO,, 
arsenous acid, H3AsO3, arsenic acid, H3AsO4 phosphorous acid, H3PO8, 
hypophosphorous acid, H3PO2, hypochlorous acid, HC1O, etc. They vary 
in strength considerably, as indicated by the figures in the table on p. 79 
giving the "degree of ionization" or fraction ionized in 0.1 normal solution. 
The most important weak base ordinarily encountered is ammonium 
hydroxide, NH4OH. It is this quality which makes it useful in cleansing, 
for while it does not yield enough free hydroxide ion to be injurious in 
washing, yet, if the little it does yield is used up, more can dissociate, 
the undissociated portion acting as a sort of reserve for hydroxide ion. 
Weak Salts. Most salts are highly ionized when in solution, even 
though they be salts of weak acids or bases like sodium acetate or 
ammonium chloride. There are, however, a few exceptions, including lead 
acetate, Pb(C2H3O2)2, the halides of cadmium, CdCl2, CdBr2 and 
Cdl2, mercuric chloride, HgCl,, mercuric cyanide, Hg(CN)2, and ferric 
sulfocyanate, Fe(CNS)3. These salts are poor electrolytes, and their 
reactions in solution show the presence of relatively few ions. For 
example, all ordinary soluble chlorides give a precipitate of lead chloride, 
PbCl2, on the addition of lead ion to their solutions. With mercuric 
chloride, however, no precipitate is obtained, indicating that there is less 
chloride ion in a concentrated solution of mercuric chloride than in a 
saturated solution of the rather insoluble lead chloride. If, however, 
silver ion is taken, a precipitate is formed, for silver chloride, AgCl, 
is much less soluble than lead chloride, and HgCl2 yields enough chloride 
ion to precipitate it. 
Ionisation of Weak Poly basic Acids in Steps. It is found that the 
two or more hydrogen atoms in the molecule of a polybasic acid like 
H,CO3, or H2S or H3PO4, tend to dissociate with unequal readiness into 
ions. While the first hydrogen atom of H2CO3 is but slightly ionized, 
it is possible to neutralize it with OH" forming water and HCO3", before 
the neutralization of the second hydrogen atom, in the HCO3", begins. 
This explains the ease with which it is possible to form acid salts in such 
cases. The first atom of hydrogen in H3PO4 ionizes readily, corresponding 
to a moderately strong acid, the second ionizes from H2PO4~ with much 
less ease, so that the ion H2PO4~ is a rather weak acid, while the third 
is hard to neutralize, HPO4"~ being a very weak acid. Conversely, when 
hydrogen ion, H+, is added to an ion like PO4 it is taken up very com- 
pletely at first until it is converted into HPO4~~, then less completely 
until H2PO4~ is formed, and then only a little to form H3PO4. It takes, 
78 likewise, a smaller concentration of H+ to change CO3 to HCO3' than it 
does to change HCO3~ to H,CO3. 
Fraction ionized in 0.1 
Weak Electrolytes. Molal solution. 
Acetic acid, HC2H3O2 - 0.013 
Carbonic acid, H2CO3 0.0017 
Dihydrogen phosphate ion, H2PO4' 0.0014 
Hydrogen sulfide, H2S 0.001 
Hydrocyanic acid, HCN 0.0001 
Acid carbonate ion, HCO3~ 0.00002 
Monohydrogen phosphate ion, HPO4" 0.000002 
Ammonium, hydroxide, NH4OH 0.017 
Cadmium iodide, Cdl2 0.10 
Mercuric chloride, HgCl2 0.00006 
Ionization in Other Solvents Than Water. Water is the chief solvent 
in which we are interested, but there are others which can dissolve elec- 
trolytes with ionization. The most important of these are liquid ammonia 
and the alcohols. Ionization in all other solvents that have been investi- 
gated is less than in water, and seems to be roughly proportioned to the 
dielectric constant of the liquid. The following values for the dielectric 
constants for several liquids are of interest in this connection: 
Water 80 
Methyl (wood) alcohol 31 
Ethyl alcohol 26 
Acetone 22 
Ammonia 16 
Chloroform 5 
Ether 4 
Ionization of Fused Salts. Fused salts appear to be highly ionized, 
conducting the electric current very rapidly. Those most highly ionized 
are probably those in which there is the greatest difference in metallic 
character between the two constituent parts. On account of this high 
conductivity many substances are obtained industrially by the electrolysis 
of fused salts rather than of aqueous solutions. This is the case especially 
with metals like sodium, calcium and aluminum, where, instead of the 
desired metal, the hydrogen from the water would be deposited at the 
cathode. Sodium is obtained by the electrolysis of fused NaOH, 
steam and oxygen being produced at the anode. Calcium is obtained, along 
with chlorine gas, by the electrolysis of fused calcium chloride. Aluminum, 
used at the present time in such large amounts, is obtained by the elec- 
trolysis of A12O3 dissolved in fused cryolite, Na3AlFG. 
Ionization of Water. Water itself is very slightly ionized. The 
purest water that can be prepared has a very slight conductivity due to 
H+ and OH-. Knowing the conductivity of these ions in more concen- 
trated solution, as we do from measurements with acids and bases, it 
is possible to calculate the concentration of these ions in water to be 
approximately 10'7 molal at ordinary temperatures. This corresponds to 
0.1 milligram of H+ and 1.7 milligram of OH' per ton of water. Other 
methods of measurement confirm these figures. 
Exercises. 
1. How would you define acid, base, salt, in terms of the ionic 
theory? 
2. Summarize in writing the evidence in favor of the theory of 
electrolytic dissociation. 
3. State in writing all the evidence you have learned, showing that 
79 a solution of acetic acid is much less ionized than one of hydrochloric 
acid of the same concentration. 
4. Divide the following substances into two lists, one for those 
slightly ionized in water, the other for those highly ionized: ammonium 
hydroxide, ammonium chloride, hydrochloric acid, sodium acetate, car- 
bonic acid, sodium chloride, mercuric chloride. 
5. Write the formula's of all the salts that can be formed from the 
following ions: Ca++, Mg++, Fe+++, Al+++, SO4_", Cl-. 
6. What are the properties of H2 and H+? How might each one 
be converted into the other? 
7. What are the properties of the following substances: Cu; Cu'+; 
Ag; Ag+; Cl2; Cl"? 
8. Describe some of the properties of the dilute aqueous solutions 
of the following: HC1; KNO3; H2SO4; Na2SO4; KC1; CuCl2; CuSO4. 
9. How does the evidence obtained from the heat of neutralization 
support the ionic theory? 
10. Write the principal reaction occurring when (a) NaCl dissolves 
in water; (b) a solution of sodium sulfate is evaporated; (c) when 
concentrated HCl-solution is warmed. 
11. In writing the following equations represent all substances by 
formulas indicating that they are mainly ionized or un-ionized, as the 
case may be: a solution of acetic acid is neutralized by one of sodium 
hydroxide; silver nitrate and potassium chloride solutions are mixed, 
giving a precipitate of silver chloride; dilute hydrochloric acid acts on a 
solution of sodium carbonate giving CO2 gas, water and sodium chloride 
solution; zinc is put into a solution of copper sulfate, giving a precipitate 
of metallic copper; concentrated sulfuric acid acts on solid sodium 
chloride, giving hydrogen chloride gas and solid sodium sulfate; mercuric 
chloride solution is treated with hydrogen sulfide gas, giving a precipitate 
of mercuric sulfide, etc. 
12. What is the concentration of OH~ in (a) equivalents per liter, 
and (b) grams per liter, when 1.48 gm. of Ca(OH)2 is dissolved in 500 
cc. ? Ans. (a) 0.08 N; (b) 1.36 gm. per 1. 
13. What is the valence of each element in the following substances: 
Cu++, Ag+, OH", Al (OH) 3, Cl2, C1-, HgCl2, H2S, NH4+, H2O? 
14. Define or explain the following terms: ion, electrolyte, degree of 
ionization, faraday, electron. 
15. What is the degree of ionization of a weak monacid base if the 
concentration of the OH" in a 0.5 normal solution is 0.002 normal? 
Ans. 0.004. 
16. (a) What is the concentration of H+ in a 2 normal solution of 
acetic acid if it is 0.006 ionized? (b) What is the concentration of Ac-? 
(c) Of the un-ionized HAc? Ans. (a) 0.012N; (b) 0.012N; (c) 1.988N. 
17. How many (a) faradays, (b) coulombs are required to deposit 
1 gram of each of the following ions at the appropriate electrode; and 
(c) how long will it take using a current of 2 amperes: (1) H+, (2) Ni++, 
(3) CD, (4) Al-? 
Ans. (a) (1) 1.00, (2) 0.340, (3) 0.0282, (4) 0.112 faradays; 
(b) (1) 96540, (2) 3280, (3) 2720, (4) 10700 coulombs; 
(c) (1) 26 h. 49 min., (2) 54 min. 40 sec., (3) 45 min. 20 sec., 
(4) 2 h. 58 min. 20 sec. 
80 CHAPTER XI. 
THE SPEED OF CHEMICAL REACTIONS. 
The previous chapters have been taken up with various aspects of 
the nature and composition of substances and with the representation of 
substances and their reactions by means of formulas and equations. We 
are now ready to inquire into the means of controlling chemical reactions, 
so as to realize desirable possibilities or prevent undesirable ones. There 
are two factors involved in this control, the direction and the speed. 
For example at high temperatures the following reaction takes place 
readily: 
2NH3 = N2 4- 3H2. 
The theory to be developed in Chapter XV shows that it should be 
possible to reverse this reaction at low temperatures and make ammonia 
from its elements. We find by experiment, however, that the reaction 
proceeds so slowly at low temperatures as to be utterly useless as a means 
of making ammonia, and that we must seek some means for its acceler- 
ation in order to make it of any practical use. The means of altering 
the direction of a reaction will be discussed in the following chapters, 
this one being devoted to the question of speed in reactions, where the 
direction is already assured. Sometimes we may wish to increase the 
speed of a reaction, as in the cooking of food, and sometimes we may 
wish to retard an undesirable reaction, such as the rusting of iron or the 
decay of wood. How may this be done? 
Application of the Kinetic Theory. We have seen, in connection 
with the behavior of gases and solutions, Chapters II and IV, how fruitful 
is the kinetic theory of matter. We may apply it with equal success to 
the question of the means for controlling the speed of chemical reactions. 
In order for reactions between different substances to take place their 
molecules must come together, or collide. We should therefore expect 
any means of increasing the number (and perhaps the force) of these 
collisions to be effective in increasing the speed. What means are available 
for changing the number and speed of these collisions? 
Effect of Temperature. Since the kinetic theory connects the tem- 
perature of any body with the velocity of the molecules composing it, 
one way indicated for increasing the number and also the force of the 
molecular collisions is to increase the temperature. We find, as a matter 
of fact, that the effect of increasing the temperature is always to increase 
the reaction velocity. For reactions occurring within a single phase, 
e. g., between dissolved or gaseous substances, it is found that the effect 
of increasing the temperature 10° is usually to increase the speed between 
two and three times. Where the reaction occurs at the boundary line 
between two phases, as in dissolving a solid, the effect of temperature is 
usually somewhat less. 
Chemists constantly make use of the accelerating effect of higher 
temperature in chemical reactions, and we see it constantly in everyday 
life. The reactions responsible for the growth of plants, for the decay 
of dead animal and plant substances, for the souring of milk, are all 
greatly influenced by the temperature, so that the importance of warmth 
for growth and of cold storage for preservation is well known to all. 
The housewife sets her bread to rise in a warm place to hasten the 
81 fermentation which produces the carbon dioxide. Those who have lived 
or camped at high altitudes know of the difficulty of cooking certain 
foods at the lower temperature at which water boils under such condi- 
tions. The difficulty is sometimes solved by means of "pressure cookers," 
which hold the steam under pressure allowing a higher temperature to be 
reached. Such an apparatus may also be used at ordinary altitudes, 
resulting in a great saving of fuel when cooking such things as beans 
or the tougher cuts of meat, for it requires scarcely any more fuel to 
maintain a temperature of 120° than it does to maintain one of 100°, 
while a reaction that requires 5 hours at the latter temperature might be 
completed in less than 1 hour at the former. In such processes as 
extracting glue, gelatin and fats "superheated steam" is used with similar 
effect. The speed of cooking such a thing as a potato is also increased 
by cutting it into small pieces, so that all parts of it are quickly brought 
to the temperature of the boiling water in which it is immersed. 
The process of hardening steel is an interesting case of retarding a 
reaction by lowering the temperature. Above 725° a steel containing 
0.9% carbon consists of a hard, tough variety of iron holding the carbon 
in solid solution, but below that temperature it tends to change into a 
heterogeneous mixture of soft iron (like wrought iron) and a hard, 
brittle iron carbide, Fe3C, known as cementite. However, this transition 
requires an appreciable time, and if the steel, heated above 725°, is 
suddenly cooled by quenching in water, so that low temperatures are 
quickly reached, the tough solid solution may be gotten at ordinary 
temperatures, where the velocity of the change is practically negligible. 
A solution which required 10 seconds for transition to the variety stable 
at ordinary temperatures at 720° would require 300,000 billion years at 
20° if the reaction velocity were halved for every 10° fall in temperature. 
Hardened steel is thus an unstable substance at ordinary temperatures but 
we can keep it almost indefinitely by reason of the extreme slowness of 
its change into the stable variety. If we cool the steel slowly, i. e., 
anneal it, or if we allow the hardened steel to get too warm for a while, 
as by too rapid grinding or cutting with a tool, it changes into the stable 
soft mixture, losing its "temper." 
Many chemical substances are similarly unstable al ordinary tem- 
peratures, existing only because the speed at which they decompose is 
small. In this class are included ozone, O3, hydrogen peroxide, H2O2, 
nitric oxide, NO, and all explosives. Nitric oxide is stable in the true 
sense only at the temperature of a very hot electric arc and it is gotten 
at ordinary temperatures by a process of rapid cooling similar in principle 
to the hardening of steel. 
Spontaneous combustion is the result of the preliminary slow oxida- 
tion which many combustible substances undergo in contact with air. 
Ordinarily the heat of these reactions is liberated so slowly that it has 
time to be conducted away, so that no perceptible rise in temperature of 
the oxidizing material takes place. When, however, the material is a 
poor conductor of heat, like oil-soaked cotton waste, and is in large 
quantities, so that the heat produced is partly retained, then the rise in 
temperature causes an increase in the rate of oxidation, with a further 
gradual rise in temperature, until the temperature of ordinary rapid com- 
bustion is attained, and the material bursts into flames. 
Effect of Concentration. Another means of increasing the number 
of favorable collisions between molecules is to increase the concentration 
of the reacting substances. The more concentrated a reacting substance 
82 is the more chances there are of its molecules colliding with those of other 
substances, hence an increase in the concentration of each kind of reacting 
molecule should increase the speed of reaction. If, for example, two 
molecules have to collide in order to bring about a certain reaction, then 
doubling the concentration of either should double the speed of reaction, 
and doubling the concentration of both should quadruple the speed. Of 
course, as the reacting substances are used up, the speed diminishes. 
Numerical Relation Between Concentration and Speed. If cn c2, c;>, 
etc., denote the concentrations of the reacting molecules then the velocity 
of the reaction, v, as measured by the rate of disappearance of one of 
the substances or the formation of another is proportional to, or equal 
to a constant times, the product of the concentration of the reacting 
molecules, i. e., v = kc1c2c3 .... The so-called "order" of the reaction 
is the number of reacting molecules which determine its speed. If there 
is only one kind of molecule reacting, two of which have to collide to 
react, then c1 = c2, and v = kcx2. This would be called a reaction of the 
second order, or a bimolecular reaction. In a reaction like 2Fe+++ -|- 
Sn++= 2Fe++-J-Sn++++, if we let ct denote the concentration of ferric 
ion, Fe+++, and c2, denote that of the stannous ion, Sn++, we have as the 
expression for the velocity v = kc^c,. This would be a reaction of the 
third order, or a trimolecular reaction. When the reaction appears to take 
place between several molecules we usually find it taking place in stages, 
since the chances of a large number of molecules colliding simultaneously 
are vastly less than for a few molecules. Thus, the equation for the 
burning of acetylene, 
2C2H2 + 5O2 = 4CO2 + 2H2O, 
appears to depend on the collision of seven molecules, which would make 
it a reaction of the seventh order, but measurement shows it to be of 
the second order. Evidently there oq,cur simultaneous reactions, a slower 
one, probably. 
C2H2 + O2 = 2CO + H2, 
which determines the speed of the total reaction, and probably two more 
rapid reactions, 
2CO + O2 = 2CO2, and 2H2 + O2 = 2H2O. 
Measurements of the. rate at which the speed falls off, enable us to 
determine the order of the reaction and hence, often, to determine what 
steps occur in complicated reactions. 
Effect of Stirring and of Contact Surface in Reactions Occurring 
at the Boundaries Between Phases. If a reaction occurs at the boundary 
between phases, as in the solution of a solid in a liquid, it is evident 
that stirring would have great effect in bringing fresh reacting molecules 
to the surface of contact, also in removing the products of reaction from 
the scene of action. The natural tendency of the molecules to diffuse can 
thus be greatly aided. 
Contact between the materials in the two phases can be likewise 
aided by having the contact surface large. For this reason powdered 
sugar will dissolve more rapidly than granulated or lump sugar. 
We may summarize and illustrate the foregoing by stating the con- 
ditions for the rapid solution of a metal in acid as follows: high tem- 
perature, concentrated acid (unless the salt produced would be thereby 
rendered insoluble, preventing its diffusion away from the surface of the 
metal), the use of a strong acid like hydrochloric acid giving a high 
concentration of hydrogen ion, instead of a weak acid like acetic acid, 
giving a small concentration of hydrogen ion (assuming that the metal 
83 is a base metal which will be dissolved by hydrogen ion), stirring the 
solution, and using the metal in rather finely divided condition rather 
than in large pieces. 
An undesirable reaction may, of course, be hindered by the converse 
of the above means. 
Effect of Catalysts. Many reactions can be accelerated or retarded 
by the presence of substances which are not themselves permanently used 
up by the reaction, and which need be present only in small amount to 
affect the speed for a large amount of reacting mixture. Such substances 
are called catalysts. One class of catalysts acts by furnishing a surface 
at which the reacting molecules are condensed and concentrated, causing 
them to come into more intimate contact. Platinum and palladium act in 
this way for a number of gas reactions. Hydrogen and oxygen scarcely 
react at all at ordinary temperatures, but platinum and palladium have the 
power to absorb these gases, especially the hydrogen, and allow them to 
react with each other. Under suitable conditions the heat produced by 
this union may be sufficient to raise the temperature of the metal to a 
point where it can ignite the remaining gas mixture. Some gas and 
cigar lighters are constructed on this principle, and when the platinum 
or palladium is held in the stream of gas mixed with air, or in the 
escaping vapors of wood alcohol from a wick, the gas or vapor is ignited 
by the heat generated at first in the catalyst. Platinum finds extensive 
use in the "contact process" for the manufacture of sulfuric acid. As 
will be explained in Chapter XIV, the reaction, 2SO2 + O2 = 2SO3 
(sulfuric acid is obtained by dissolving SO3 in water) occurs in the 
desired direction only at lower temperatures, where the reaction proceeds 
too slowly to be of use. When, however, the gases are passed over 
platinum at about 500°C they react readily, so that the process has 
supplanted the older "lead-chamber process" to a considerable extent. 
When platinum is used to catalyze reactions its effectiveness may be 
increased by increasing the surface as much as possible, since the reactions 
occur at the surface of the metal. This is particularly important on 
account of the high cost of platinum. The usual form in which it is used 
is "platinized asbestos," made by wetting asbestos fiber with a solution 
of a platinum salt and then heating it so as to decompose the salt and 
yield finely divided metallic platinum. 
The union of nitrogen and hydrogen directly to give ammonia, 
referred to at the beginning of this chapter, is brought about by the aid 
of a similar catalyst, uranium, to whose aid we owe this important means 
for the fixation of atmospheric nitrogen and its application as fertilizer. 
Hydrogen peroxide is a very unstable substance, tending to decompose 
giving oxygen and water. This decomposition is aided by the presence 
of finely divided solids, such as manganese dioxide, MnO2. This oxide 
likewise catalyzes other reactions where oxygen is evolved, notably 
2KC1O3 = 2KC1 + 3O2. 
In making oxygen by heating potassium chlorate it is customary to add 
powdered manganese dioxide so that the oxygen will be evolved more 
smoothly and at a lower temperature. 
To prevent the decomposition of hydrogen peroxide, besides keeping 
it dilute and cool, as the kinetic theory would require, a negative or 
retarding catalyst is frequently added in the form of acetanilid 
C6H5NHOC2H3. 
Many catalysts evidently act as "carriers," by the formation of an 
intermediate product which then decomposes readily. Such a catalyst 
84 is NO in the "lead-chamber process" for the manufacture of sulfuric 
acid. The reaction, 
SO2 4" H2O 4* //2O2 = H2SO4, 
does not occur with sufficient speed, but since the reaction 
NO + %O2 = NO2 
occurs readily in both directions, the oxides of nitrogen make the oxygen 
available in the following way, 
SO2 + H2O + NO2 = H2SO4 + NO. 
The NO2 is regenerated as shown above, so that a little of it suffices for 
the manufacture of an almost indefinite quantity of sulfuric acid. 
An interesting reaction is the "inversion" of cane sugar, which con- 
sists in its splitting up, with the addition of a molecule of water, into 
two molecules of simpler sugars, fructose and glucose, containing the 
same atoms, but in different arrangement. The reaction is 
C12H22OU 4- H2O = C6H12O6(glucose)+ C6H12O8(fructose). 
The resulting mixture of sugars is often called invert sugar. This reaction 
is catalyzed by the hydrogen ion of acids, the speed of inversion being 
proportional to the concentration of hydrogen ion, which doubtless forms 
an intermediate product. Since the resulting invert sugar does not easily 
crystallize, and is, moreover, less sweet than the equivalent amount of 
cane sugar, it is important in the refining of sugar, to avoid all acidity 
of the sugar solutions. When stewed fruit is sweetened the effect of the 
sugar will be greater if the fruit is first allowed to cool, retarding the 
inversion of the sugar by the acid of the fruit. The altered taste of 
lemonade that has stood for some hours is due to this inversion of its 
sugar. It is of interest to note that the sugar in honey is invert sugar. 
The inversion doubtless takes place in the body of the bee when it is 
gathered, and prevents subsequent crystallization in the comb. 
By a process analogous to the inversion of sugar, starch may be made 
to take up water becoming transformed into a mixture of various sugars 
to which is given the commercial name "glucose." Large quantities of 
syrup and sugar are made from corn-starch by digestion with acid which 
acts as the catalyst. 
Water is a catalyst for many reactions. Absolutely dry CO will not 
burn. Dry AgNO3 and HC1 dissolved in ether form no precipitate of 
AgCl. When ordinary NH4C1 is heated it sublimes, dissociating into 
NH3 and HC1 gases. If, however, the salt is quite dry it vaporizes 
without dissociation. 
Enzymes. An important class of catalysts, known as enzymes, are 
produced by living organisms and catalyze many reactions. The ptyalin 
of saliva, whose function is to convert starch into sugar, the diastase 
of malt, which has a similar action, and the zymase of the yeast plant, 
which converts certain sugars into alcohol and carbon dioxide, are 
examples of important enzymes. Similar bodies are thought to play 
fundamental roles in the processes of growth and nutrition of plants and 
animals. 
Exercises. 
1. Define catalyst. 
2. Divide the catalysts mentioned in the text into surface catalysts 
and dissolved catalysts. 
3. What effect may a double boiler have on the speed of cooking 
food? 
85 4. What means could you employ for rapidly washing photographic 
prints free from the "hypo" of the fixing bath? 
5. How might you alter the time required for developing a photo- 
graphic negative? 
6. How might you endeavor to increase the speed of the reaction: 
BrO3~ + 61- + 6H+ = 3H2O + Br- + 3I2 ? 
What effect would acetic acid have upon the speed as compared with 
hydrochloric acid of the same concentration? 
7. The organism existing in "mother of vinegar" converts alcohol 
into acetic acid by the reaction 
C2H5OH 4- O2 = C2H4O2 +h2o. 
By what means could the manufacture of vinegar be hastened? 
86 CHAPTER XII. 
CHEMICAL EQUILIBRIUM. THE EFFECT OF CONCENTRATION. 
Reversibility of Chemical Reactions. The problem of the control 
of chemical reactions involves not only the speed, as discussed in Chapter 
XI, but also the.fact that most reactions are reversible. We have, there- 
fore, two problems as distinct from each other as are the speed and desti- 
nation of a train. To reach a desired destination it is not sufficient to 
get on a fast train, the train must also move in the right direction. With 
chemical reactions, similarly, the conditions accelerating the reaction 
often alter the direction at the same time. For example, the velocity of 
formation of SO3 from SO2 and O2 increases with the temperature, but 
the velocity of decomposition of SO3 back into SO2 and O2 is likewise 
increased thereby, and more rapidly than the velocity of its formation, so 
that SO3 becomes increasingly unstable at higher temperatures. We will 
therefore turn our attention to the problem of the direction of chemical 
reactions, assuming that all the time necessary for them to take place 
is allowed. We will finally, in Chapter XIV, consider both effects together. 
Most chemical reactions are reversible, that is, the products of a 
reaction may usually be made to react with each other to produce the 
original substances by suitable choice of conditions. Many examples of 
this may be given. A liquid may be made to evaporate by raising the 
temperature or by reducing the pressure, and its vapor, conversely, may 
be recondensed to liquid by lowering the temperature or increasing the 
pressure. Solids may be liquified by increasing the temperature and 
liquids may be solidified by lowering the temperature. Gases may either 
be dissolved in liquids or driven out of solution by suitable changes in 
pressure or temperature. The naturally occurring mineral, gypsum, 
CaSO4 • 2H2O, may be deprived of part of its water by moderate heating, 
giving plaster of Paris, which is approximately CaSO2 • J^H2O. When 
this is mixed with water the original dihydrate is re-formed, which causes 
the plaster to "set," due to the interlocking of the growing crystals of 
CaSO4 • 2H2O. Nitric oxide, NO, is decomposed by gentle heating into 
nitrogen and oxygen, and yet, at the high temperature of the electric 
arc a considerable amount of nitric oxide may be formed from the 
nitrogen and oxygen of air. Oxygen, O2, is changed into ozone, O3, by 
the silent electric discharge, and then gradually reverts to oxygen again. 
When a lead storage battery is discharged the lead dioxide of the positive 
plate, the lead of the negative plate, and the sulfuric acid in the solution 
react as follows: 
PbO2 4- Pb + 4H+ 4- 2SO4" = 2PbSO4 + 2H2O, 
and when the battery is recharged the reverse reaction takes place, cor- 
responding to the equation as read from right to left. When water 
charged with carbon dioxide passes over limestone rock the following 
reaction occurs: 
CaCO3 + H2O 4- CO2 = Ca++ 4- 2HCO3-. 
The calcium ion dissolved in the water produces what is called "hard 
water." and interferes with the cleansing power of soap by precipitating 
insoluble calcium soap. When this water is boiled, however, the CO2 
is driven off and the calcium carbonate is re-precipitated, corresponding 
to the equation as read from right to left. When steam is passed through 
87 a tube containing iron filings or nails, heated to dull redness, the following 
reaction occurs: 
3Fe + 4H,O = Fe3O4 + 4H2. 
When hydrogen is passed over the iron oxide, Fe3O4, under similar con- 
ditions, some of it is changed to steam with the reduction of the oxide 
to iron, exactly the reverse of the above. Illustrations of reversibility 
might be multiplied ad libitum. 
The experimental proof of reversibility consists simply in mixing each 
set of substances, as indicated by the two sides of the equation, and seeking 
conditions under which the other set will be formed. The disappearance 
of any of the substances taken or the appearance of any of the products 
is usually sufficient proof that a reaction has taken place. What particular 
substance to test for is determined by the ease and decisiveness of the 
available tests. Thus the oxidation of mercury, . represented by the 
equation 2Hg + O2 = 2HgO, could be proven if conditions could be 
found under which mercury in contact with oxygen would increase in 
weight. A change in appearance corresponding to the red color of the 
oxide, or a decrease in the amount of the gaseous oxygen would suffice 
equally well. The reverse reaction, the dissociation of mercuric oxide into 
mercury and oxygen, could be proven by taking the red mercuric oxide 
and discovering conditions of pressure and temperature under which there 
could be detected either the appearance of mercury or oxygen, or else a 
loss in weight. 
It is Possible to Have All of the Substances Involved in Reversible 
Reactions Present Together in Chemical Equilibrium with Each Other. 
Instead of passing steam over heated iron in a tube, as in the experiment 
described above, we might simply heat them together in a closed vessel. 
We would expect, as before, the formation of some hydrogen and iron 
oxide, corresponding to the equation as read from left to right. Likewise, 
if we heat hydrogen and iron oxide together at the same temperature and 
in a similar vessel, we would expect the formation of iron and steam, 
corresponding to the equation as read from right to left. If we allow 
sufficient time for reaction in both cases how will the final state of affairs 
in one vessel compare with that in the other? It is obviously absurd to 
expect either reaction to go to completion, just as it would be absurd to 
expect a ball placed on a certain table to roll all the way from a to b 
when we have previously found that, when placed at b it rolls towards a. 
Obviously there must be some intermediate position on the table at which 
it tends to come to rest no matter at which end it is started. Again, if 
it is found that one body becomes cooler when brought into a certain room, 
and another becomes warmer when brought into the same room, it is 
evident that each body will come eventually to the same temperature, 
that of the room, which must lie somewhere between the initial tempera- 
tures of the two bodies. By the same kind of reasoning, it should make no 
difference whether we start with 3Fe-|- 4H2O, or with Fe3O4 -j- 4H2, 
if we heat them to a temperature at which they are able to react, in identical 
vessels, we will eventually get a mixture of all four substances in chemical 
equilibrium. The term equilibrium signifies an adjustment of the amounts 
of the reacting substances, such that there is no further tendency to react, 
just as in mechanical equilibrium we have an adjustment of forces such 
that no further action takes place. This adjustment of the amounts of 
substances in chemical equilibrium may be thought of as the result of an 
adjustment between chemical forces. Just as a balance loaded with certain 
weights will finally come to rest, provided that the balance is free to 
88 swing, at a position dependent only upon the weights on the pans and not 
upon the direction in which an original displacement took place, so in 
a reversible chemical reaction the final equilibrium between the amounts 
of the reacting substances, provided that they are free to react (suitable 
temperature, catalyst, etc.), will depend only on the relative amounts of 
the substances present, and not upon which particula rset of substances, 
as represented by the two sides of the equation, is chosen. (There is, 
of course, no oscillation in reaching chemical equilibrium comparable to 
that of a balance in coming to rest.) 
From a kinetic standpoint chemical equilibrium must be regarded as 
the result of two opposite reactions proceeding with equal speed. When, 
for example, we start with iron and steam, at a sufficiently high tempera- 
ture, they react with a velocity depending on the concentration of steam 
molecules and on the surface of the iron. As these are used up the reac- 
tion gradually becomes slower. At the same time, as more hydrogen and 
iron oxide are produced, they begin to react with each other to regenerate 
the oxide and steam with an increasing velocity until the rate of one 
reaction just equals that of the other. This conception of equilibrium 
was used in connection with simpler phenomena in Chapter II, (p. 15) 
and has an important bearing on what is here to follow. 
The Effect on Chemical Equilibrium of Changing the Concentration 
of the Reacting Substances. If a little weight is added to or taken from 
one pan of a balance in equilibrium the pointer swings a little to one side 
or the other until a new position of equilibrium is reached. Similarly, 
when a chemical equilibrium exists it is possible to disturb it by any 
means which changes the relative speeds of the two opposing reactions. 
These rates may be affected by changing the temperature, or the total 
pressure on the reacting system, or the concentration of the substances 
involved. The discussion of the effect of changing the temperature or total 
pressure will be taken up in Chapter XIV; and we will turn our attention 
first to changes in concentration. 
If we recall the kinetic conception of chemical equilibrium we will see 
that if we increase the concentration of one of the reacting substances 
there are then more molecules of that substance present in a given space, 
so that their ability to find and react with other molecules is increased. 
This will cause a temporary gain in the rate of the reaction using up this 
kind of molecule, until more of the products of its reaction, represented on 
the other side of the equation, are produced, so as to equalize again the 
rates of the two reactions. The net result is that the reaction proceeds 
to some extent in the direction which will use up the substance whose 
concentration is thus increased. For illustration let us consider a closed 
vessel containing SO3, SO2 and O2 in chemical equilibrium to attain 
which requires a sufficiently high temperature and perhaps a catalyst. 
When equilibrium is attained we must imagine that the two reactions, rep- 
resented by the equation 2SO2 + O2 = 2SO3, read in both directions, are 
taking place with equal velocities, with a net result of no change in the 
relative amounts of the three gases present. Suppose, now, that more 
oxygen is introduced into the vessel. The increased number of oxygen 
molecules will render it easier for them to collide with sulfur dioxide 
molecules, resulting in a more rapid formation of sulfur trroxide molecules. 
As the number of the latter increase their own speed of decomposition 
will increase until the two reactions again take place at equal rates. The 
net effect of this increase in the amount of oxygen is to cause more of it 
to be used up, with a decrease in the amount of sulfur dioxide and an 
89 increase in the amount of sulfur trioxide. The effect of adding or 
removing any of the three substances present can be predicted on the 
above basis. 
From a technical point of view, in making sulfur trioxide, the im- 
portant thing is to use up the sulfur dioxide as completely as possible, 
since the latter costs money whereas the supply of oxygen is free. To 
accomplish this the kinetic theory indicates that it is desirable to have 
the oxygen in excess, since otherwise considerable sulfur dioxide would 
remain uncombined when equilibrium has been reached. 
In general, after equilibrium has been reached, the effect of any 
change in the concentration of the reacting substances is to cause that 
reaction to take place which neutralizes the effect of the change; that is, 
if the concentration of any substance is increased, that reaction tends to 
take place which uses up that substance, and where the concentration of 
any substance is decreased, that reaction tends to take place which will 
replace that substance. This conclusion is a special case of a more general 
one, called the Theorem of Le Chotelier, which states that when equili- 
brium has been reached a change in any of the factors affecting equilibrium 
tends to make that reaction take place which will neutralize the effect of 
the change. The discussion of other factors will be taken up in Chapter 
XIV. 
The conclusion, in so far as concentration alone is concerned, is 
called the "Law of Mass Action," or the "Mass Law," which is slightly 
misleading, since it is not the masses of the substances present, but their 
concentrations, or masses per unit volume, which determine equilibrium. 
This point was discussed in Chapter II, where it was shown that the 
depth of the water in equilibrium with its vapor, or the extent of the 
surface between the two phases has no effect upon the equilibrium. On 
the other hand, a change in the concentration of the vapor, as by changing 
its volume, or in the concentration of the liquid, as by adding some other 
substance to it, would cause some reaction to take place. Likewise, in 
the equilibrium discussed above, between iron, its oxide, steam and 
hydrogen, the relative amounts of the solids has no more effect than 
would the addition of more salt to a saturated salt solution. The reaction 
takes place where the three phases, the two solids and the gas phase, are 
in contact, and its rate does not depend on the excess of either solid. 
These same conclusions apply to reactions occurring in solution, 
including those between ionized substances. For example, that the 
reaction 
HC2H8O2 = H+ + C2H3O2 - 
will take place as read from left to right may be shown by taking pure 
acetic acid (which is un-ionized, as shown by its being a non-conductor) 
and dissolving it in water, when the solution will have the properties of 
hydrogen ion. That the reaction will also take place as read from right 
to left may be shown by mixing a solution of a strong acid, like hydro- 
chloric acid, giving a large concentration of hydrogen ion, with one of 
sodium acetate, which gives a large concentration of acetate ion, when we 
find that hydrogen ion disappears partly from the solution, also that acetic 
acid appears, as is evident by its odor. It makes no difference therefore, 
whether we start with a mol of acetic acid in a liter of water, or with a 
mol each of hydrogen and acetate ions in a liter of water, the final 
state of affairs is the same, which is that we will have present about 0.01 
mol of each of the two ions and 0.99 mol of the undissociated molecule. 
If, subsequently, we add another mol of acetate ion (in solution with, say, 
90 sodium ion, which affects none of the substances present) some of the 
acetate ion will be able to use up more of the hydrogen ion, becoming 
acetic acid, resulting in a considerable decrease in the acidity (hydrogen 
ion concentration) of the solution. In like manner, solutions of the 
weak base ammonium hydroxide, NH4OH, are much less alkaline in the 
presence of any substance which dissolves to yield an excess of ammonium 
ion, NH4+. 
When experiment shows that one set of substances react almost com- 
pletely to give certain products, then it must be evident that these products 
themselves, when mixed under the same conditions, will give but little 
of the original substances at equilibrium. When, for example, we have 
found that carbonic acid is extremely weak, dissociating but slightly accord- 
ing to the equation: 
H2CO3 = H+ + HCO3-; 
then we should be able to conclude at once that when the ions are mixed, 
as in mixing hydrochloric acid and sodium acid carbonate solutions, they 
will unite almost completely to form carbonic acid. 
Quantitative Expression of the Mass Law. In the above discussion 
of the effect of concentration on chemical equilibrium we have been con- 
tent with predicting the direction of the effect of altering concentration, 
the qualitative effect, without inquiring how much a given equilibrium 
would be disturbed by a certain change in concentration of one or more 
of the reacting substances. This would be called the quantitative effect. 
While the qualitative prediction suffices in most cases, it is nevertheless 
often desirable to make a quantitative prediction. This is possible by com- 
bining the quantitative expression for the speed of a chemical reaction, 
given in Chapter XI, with the idea introduced in this chapter, that at 
equilibrium the two opposite reactions are proceeding with equal velocity. 
For the sulfur trioxide reaction discussed above we would have for 
the velocity of formation v4 = kxc2So2 ' c01>, where kt is a constant 
Similarly, for the decomposition of the trioxide back into the dioxide and 
oxygen we would have for the velocity v2 = k2c2So3, at equilibrium 
the two velocities are equal, so that k1c2So2 ' co2 = k2c2So3- Since kx 
and k2 are constants their quotient is a constant, K, so that we may write 
C SO2Co2 
= K. 
c2so3 
By means of this expression we are able to make qualitative as well 
as quantitative predictions. Since K is constant at a given temperature, 
the values of the three concentrations must always adjust themselves 
accordingly. For example, suppose that it were found by measurement, at 
a certain temperature, that when equilibrium was reached the concen- 
trations were as follows: SO2, 3 mols per liter, O2, 1 mol per liter, SO3, 
2 mols per liter. By substituting these values in the above expression we 
obtain for the value of the constant K, 2.25. From this we can calculate 
another set of equilibrium concentrations. Thus, if the concentration of 
O2 were 5 mols per liter, and that of SO3 2 mols per liter, we would have 
c2so2 X 5 
= 2.25 
4 
from which we would find the concentration of SO2 in equilibrium with 
sulfur trioxide and oxygen at the new concentrations to be 1.34 mols 
per liter. 
91 It is also the custom to denote concentrations by the formula of the 
substance enclosed in brackets or parenthesis. Using this notation we 
would write the above expression 
(SO2)2(O2) 
=K, 
(SO3)2 
which has exactly the same significance as before when c was used to 
denote concentration. 
It may be worth while to give examples of the equilibrium equation 
for several other reactions. 
Reaction. Equilibrium equation. 
(NO)2 
N2 +O2 = 2NO = K 
(N2)(O2) 
(NH3)2 
N2 + 3H. = 2NH3 = K 
(N2)(H2)3 
(CO)2 
CO, + C( solid )=2CO = K 
(CO2) 
(Since carbon is solid its concentration is constant and 
need not be expressed in the equilibrium equation. 
(H2O)' (H2O) 
Fe3O4(solid)+ 4H2 = 3Fe (solid)+4H,O = K or = K 
(H2)4 (H2) 
(H+) (C2H3O2~) 
H+ + C2H3O2- = HC2H3O, = K 
(HC2H3O2) 
(H+)(HCO3") 
H+ + HCO3- = H2CO3 = K 
(H2CO3) 
(H+)(CO3") 
H+ + CO3 -- = HCO3- = K 
(HCO3-) 
Since the partial pressure of a gas is proportional to the number of 
molecules in a given space, and hence to the concentration, it is possible, 
where gases are concerned, to write partial pressure instead of concen- 
trations and have constant a similar function of partial pressures. This 
constant is not necessarily the same as that where concentrations are used, 
so that we will express it as Kp. We would write, accordingly, 
P2no P2nhs 
= Kp, = Kp, etc. 
Pn2 ' Po2 Pn2 ' P3h2 
Dissociation Constants of Electrolytes. The equilibrium constants for 
the dissociation of an electrolyte, as with acetic and carbonic acids above, 
are often called dissociation constants or ionization constants, and express 
the strength of an acid or base in the most general way. It is evident that 
92 the stronger the electrolyte the larger are the concentrations written in 
the numerator of the fraction, hence the larger the dissociation constant. 
Instead of stating the degree of dissociation for an electrolyte at a certain 
concentration, as was done in Chapter X, p. 16, it is more satisfactory to 
give the dissociation constant, from, which the degree of dissociation at 
any concentration may be calculated. The following values for certain 
familiar substances may be given: 
Dissociation constant 
Substance. 
Reaction. at room temperature. 
Acetic acid 
...,HC2H3O2 = H+ 4- C2H3O2- 
1.8 X 10-5 
Dihydrogen phosphate ion... 
...h,po4-=H+ + hpo4- - 
2 X IO'7 
Carbonic acid 
.. .H2CO3 = H+ HCO3- 
3 X IO"7 
Hydrogen sulfide 
...H2S = H+ + HS- 
io-7 
Hydrocyanic acid 
_HCN = H+'+ CN- 
io-9 
Bicarbonate ion 
...,hco3-=h+ + co3-- 
4 X 10-11 
Hydrosulfide ion 
...HS- = H+ + S" 
10-15 
Ammonium hydroxide 
...NH4OH = NH4+ + OH- 
2.3 X 10-5 
The concentration of hydrogen ion in a solution of acetic acid, for 
example, may be calculated from the dissociation constant of the acid as 
follows. If the acid alone is present in the solution the hydrogen ion and 
the acetate ion are at the same concentration, hence (H+)=(C2H3O2_)• 
Suppose that the total concentration of acetic acid is 0.1 molal. The 
amount ionizing is so slight that the un-ionized acid may be considered 
as 0.1 molal without appreciable error, i. e., (HC2H3O2)=0.1. Substi- 
tuting these values in the equilibrium equation, we obtain, 
(H+)2 
=1.8X 10-5, and hence (H+)=0.0013. 
0.1 
Again, suppose that 0.05 molal acetate ion is present in the same 
solution, then (C2H3O2-)= 0.05, (HC2H3O2)= 0.1, and (H+)= 0.000036. 
It will be observed that the acidity of this solution is vastly less than 
that of the former. 
Strong electrolytes do not obey the Mass Law, for reasons which 
there is not space here to discuss, so that calculations such as the above 
cannot be made with them. However, little is lost on this account, as they 
may be considered as completely dissociated, with sufficient accuracy for 
most purposes. 
Chemical Equations Do Not Indicate Amounts Present at Equilibrium. 
It must be remembered, as was pointed out in Chapter III, that a chemical 
equation does not indicate that the reaction it represents will actually take 
place, and says nothing about the necessary conditions. When we write 
the equation 
2SO2 + O2 = 2SO3 
we should not read it "sulfur dioxide reacts with oxygen, etc.," but 
rather, if sulfur dioxide and oxygen can be made to react to form sulfur 
trioxide, they will do so in certain proportions by weight (expressed in 
mols or grams) and volume (if the temperature is such that they are 
in the gaseous state). As a matter of fact, in order to convert 1 mol 
of sulfur dioxide into sulfur trioxide in any reasonable time several things 
are necessary which are not indicated by the equation. The proper 
temperature is required; a catalyst must be used; and an excess of both 
sulfur dioxide and oxygen must be present. In reactions which do not 
naturally run to completion, but which reach equilibrium with consider- 
93 able amounts of all the substances still present, an excess of one or more 
of the reacting substances must be taken in order to get the desired 
amount of product, and this excess is in no way indicated by the equation, 
nor does the presence of this excess change the proportions in which the 
substances react. We may illustrate this important point by imagining a 
Western celebration in which a large number of cowboys are to ride 
wild horses. Now in order to have fifty mounted cowboys it would be 
necessary to have more than fifty cowboys and fifty horses present, for, 
we will assume cowboys are constantly being unhorsed. It might be 
necessary to have, say, seventy cowboys in order to have all the horses 
mounted most of the time. This does not affect the fundamental reversible 
reaction which is 
Man 4~ Horse = Mounted horseman. 
If the men are in excess of the horses we will not have two men trying 
to ride one horse, nor, if the horses are in excess, will one man try to 
ride two horses. After equilibrium has been reached, in which there are, 
on the average, a certain number of each of the three varieties represented 
by the equation, if more men arrive, the reaction corresponding to the 
equation as read from left to right will take place in a few more cases, 
shifting the equilibrium so as to use up riderless horses, while if some 
of the unmounted men leave for the nearest saloon, the net result will 
be that the reaction as read from right to left will take place until a 
new condition of equilibrium is reached. Again, starting with sixty 
horses, in order to get fifty mounted horsemen, it might be necessary 
to have several hundred men, if bank-clerks were substituted for cow- 
boys. In analogous fashion, it would require more than 1 mol of H2S 
in order to precipitate 1 mol of Cd" as CdS by the reaction 
Cd++ + H2S = CdS + 2H+. 
Though the amounts reacting do so in the proportion indicated by the 
equation, an excess of both Cd++ and H2S have to be present in order to 
make these amounts react. If, instead of Cd++ we take Zn++, we find 
that a still larger excess of H2S must be present in order to produce 1 mol 
of sulfide. In either case, after equilibrium has been reached, it may 
be shifted so as to dissolve sulfide either by the addition of more H+ or by 
the removal of H,S, as by boiling the mixture; or it might, on the otljer 
hand, be shifted so as to precipitate more of the Cd++ if H2S were added 
at a greater partial pressure (or concentration) or if H+ were removed, 
as by the addition of OH- or C2H3O2~. 
Exercises. 
1. Describe experiments which would prove whether or not the 
following reactions can be made to take place in both directions: 
2BaO(solid)4- Oo - 2BaCC(solid) ; 
H„S + S (liquid )== H,S; 
Mg++ + 2NH4OH = Mg(OH)2 (precipitate) 4- 2NHf; 
2CrO4-(yellow)+ 2H+ = H2O + Cr 2O7"(red) ; 
NH4+ 4- ci- 4- H2O = NH4OH 4- H+ + ci-. 
2. Hard water may be softened by means of an insoluble substance 
called by the commercial name, "permutite," Na(H6AlSiO7), which ex- 
changes the calcium (or magnesium) ion in the water for sodium ion 
as follows: 
Ca++ 4- 2Na (H„AlSiO7) (solid )= 2Na+ 4- Ca (H6AlSiO7) 2 (solid). 
Can you suggest a way of restoring the sodium permutite after it has 
all been changed to calcium or magnesium permutite? 
94 3. Describe in words the effect of adding solid sodium acetate to 
a solution of acetic acid, and write an equation for the reaction. 
4. The precipitation of manganese ion as manganese sulfide, accord- 
ing to the equation, 
Mn++ + H2S(gas)= MnS + 2H+, 
is never complete. How would you make the amount precipitated as 
great as possible? How would you redissolve all of the precipitate? 
5. Write equations showing what happens when 
(a) 2 mols of hydrogen are mixed with 4 mols of oxygen and an 
electric spark passed; 
(b) dilute solutions containing respectively 1 mol of hydrochloric 
acid and 2 mols of sodium acetate are mixed; 
6. (a) Which solution tastes more sour, vinegar or plain soda 
water? 
(b) Which, then, of the two following reactions takes place more 
completely, as read from left to right: 
HC2H3O2 = H+ + C2H3O2-, or H2CO3 = H+ + HCO3-? 
(c) Which of the following reactions would yield the greater 
concentration of OH-, C2H302_+Ho0 = HC9H3O2-|-OH-, or 
HCO3- + H2O = H2CO3 + OH-? 
7. Using the equilibrium constants given on page 12 calculate the 
(H+) for acids and the (OH-) for bases when the substance is present 
in 0.05 molal solution. 
95 CHAPTER XIII. 
CHEMICAL EQUILIBRIUM. PROPERTIES WHICH MAY AFFECT 
CONCENTRATION. 
We have seen, in Chapter XII, that it is possible to control reactions 
which come to equilibrium by changing the concentrations of the reacting 
substances. A reaction may be made more complete by using the reacting 
substances at greater concentration or by removing one or more of the 
products of the reaction. We may now ask the question, in what ways 
may substances be added to or taken away from reacting mixtures, and 
what properties of substances determine whether a reaction proceeds 
more completely in one direction or another? The ability to answer 
these questions enables one to predict and control reactions in a large 
number of cases without previous experiment. It is important that the 
amount of experimental material to be memorized should be the minimum 
necessary to deal with the vast number of reactions likely to be encount- 
ered. How this material may be arranged and applied will be shown in 
the following pages. 
Volatility. One of the properties of substances that may be utilized 
in bringing about reactions is volatility. If a certain reacting mixture 
is in an enclosed space, equilibrium may be reached long before all of 
the desired products are obtained, but if one. of the substances produced 
is volatile, at the temperature of the reaction, it may be allowed to escape, 
or be pumped off, and its removal will allow more of it to be formed, 
according to the principle set forth in the last chapter. If this removal 
is continued it may be possible to make the reaction go to completion. For 
example, if solutions of sodium chloride and sulfuric acid are mixed 
there is only a small tendency to form hydrochloric acid, since the latter 
is a little stronger than the former. However, since hydrogen chloride, 
in the absence of water, is a gas, it is possible to produce it by the reaction 
between solid sodium chloride and concentrated sulfuric acid, as follows: 
NaCl + H2SO4 = NaHSO4 + HC1. 
If this were done in a closed vessel equilibrium would be reached before 
much of the salt and sulfuric acid had reacted in this way, but if the 
vessel is open, so that the HC1 gas can escape, then it can all be removed, 
and the reaction may go to completion. Further heating, if enough salt 
is present, will give neutral sodium sulfate, NaoSO4, the second step being 
NaCl 4- NaHSO4 = Na2SO4'+ HC1. 
The more volatile acid is thus driven out of its salt by the less volatile 
one. This principle is used in the manufacture of hydrochloric acid, the 
gas being caught and dissolved in water. A number of other acids are 
obtained commercially from their salts by heating them with sulfuric 
acid in the same way, among them nitric acid, HNO3, from the naturally 
occurring NaNO3; hydrofluoric acid, H2F2, from fluorspar, CaF2; acetic 
acid from calcium acetate,* etc. 
Sodium sulfate, on the other hand, may be changed to sodium 
phosphate by using P2O5, which is less volatile than SO3, the reaction 
being, 
* When wood is distilled the acetic acid is separated from the wood alcohol, 
acetone, etc., occurring in the distillate by using slaked lime, Ca(OH)„, in order 
to convert it into the non-volatile calcium acetate. 
96 3Na2SO4 + P2O3 = 2Na3PO4 + 3SO3. 
This is a case where it may be more satisfactory to consider the salts as 
combinations of the metallic and non-metallic oxides, as explained in 
Chapter V, writing the equation 
3Na2O • SO3 + P2O5 =(Na2O)3 • P2O5 + 3SO3. 
Again, since SiO2 is still less volatile than P2O5, a phosphate may be 
changed to a silicate by heating it to a sufficiently high temperature with 
SiO2, as follows: 
(Na2O)3 • P2O5 + 3SiO2 = 3Na2O • SiO2 + P2O5. 
In solution, where the difference in volatility of the above substances 
would not be evident, the reverse changes would only take place, since 
silicic, phosphoric and sulfuric acids are successively stronger, as will 
be explained presently. 
It is well to remember that ammonium salts are easily volatilized 
and that many chlorides are rather volatile, especially HgCl2, FeCl3, 
A1C13, SbCl3, SbCl5, AsC13, SnCl4. 
Solubility. The solubility of substances has an important bearing 
on the course of reactions. When a rather insoluble salt is put into water 
a little of it dissolves, for almost no salts are so insoluble that the amount 
going into solution cannot be measured. The small quantity which does 
dissolve will be ionized, since nearly all salts are ionized in so far as 
they will go into solution. The equation expressing the reaction when 
such a salt is put into water should express this ionization, as exemplified 
by the following: 
BaSO4(solid)= Ba++ + SO4" 
AgCl(solid)=Ag+ + Cl- 
CaCO3 (solid)== Ca++ CO3-- 
PbCrO4(solid)= Pb++ + CrO4" 
When the solution is saturated there will be equilibrium between the solid 
salt, represented on the left side of each equation, and its ions in solution, 
represented on the right. The observed fact that all the above salts are 
only very slightly soluble means that the solutions will contain but very 
little of their ions. It also enables us to predict that if we mix the ions 
of any of the salts at any considerable concentration they will react almost 
completely to form a precipitate (from a Latin word meaning to throw 
down) of the corresponding salt. A knowledge of the solubilities of salts, 
therefore, enables us to tell what ions will precipitate each other from 
solution. No matter what other ions are present, no matter from what 
compounds they are obtained, in any appreciable concentration, Ba++ and 
SO4"~ will always form a precipitate of BaSO4. Moreover, if it is desired 
to remove Ba++ from solution most completely, our knowledge of equi- 
librium indicates that we should add an excess of SO4~~. This effect is 
constantly sought in quantitative analysis. 
Conversely, if we wish to dissolve a precipitate we must endeavor 
to remove from the solution one of its ions. This removal, in the 
case of ions, is not so simple as is the removal of a gas, but requires 
a knowledge of other combinations the ion is capable of forming, as will 
be discussed later. 
Which of the two salts of a common ion will be precipitated depends 
on their relative solubility and the relative concentrations of the ions in 
question. If a mixture of chloride and iodide ions is added to a solution 
of silver ion, silver iodide, yellow, is formed instead of silver chloride, 
white. This shows that Agl is less soluble than AgCl, or, in other words, 
97 I" removes Ag+ from solution more completely than does Cl" at the same 
concentration. This would also enable us to predict that if we start with 
AgCl we might bring about the following transformation quite readily: 
AgCl + I- = AgI + Cl-, 
whereas the reverse reaction could only be brought about by keeping the 
concentration of T very small and that of Cl" very large as would be the 
case in continual washing of the Agl precipitate on a filter with a solution 
of Cl". If, subsequently, we find that Ag2CO3, silver carbonate, is readily 
transposed into AgCl by allowing Cl- to act upon it, then we know that 
Ag2CO3 is more soluble than either of the others, or, in other words, 
carbonate ion, CO3~_, removes Ag+ from solution less completely than 
either Cl- or I-, and without trying it, we could predict with assurance that 
the following reaction would take place very readily as read from left 
to right, but not in the reverse direction: 
Ag2CO3 + 2l" = 2AgI + CO3" 
Again, on finding by experiment that Agl is readily changed into Ag2S 
by SH_, we could conclude that both of the other precipitates could be 
transposed into Ag2S still more easily. It will be seen that with the aid 
of the ideas of equilibrium we are able to predict a large number of 
reactions on the basis of a few well chosen experiments, and thus utilize 
our experimental knowledge to the best advantage. 
The same considerations apply to insoluble acids and bases. Mag- 
nesium hydroxide, Mg(OH)2, when put into water dissolves slightly 
forming its ions as follows: 
Mg(OH)2 = Mg++ + 2OH-. 
Therefore if solutions of, say, magnesium chloride and sodium hydroxide 
are mixed, a precipitate of Mg(OH)2 will be formed. 
A knowledge of the solubilities of salts, acids and bases in water 
is evidently extremely important in enabling one to predict and control 
reactions. This subject will be taken up more fully in Chapter XVII, 
but some useful generalizations may be given at this time. 
As far as the relation between solubility and other properties is con- 
cerned, we may note that, other things being equal, the higher the melting 
point of the compound the less soluble it will be, not only in water, but 
in any solvent. Also, where the melting points are approximately the 
same, the compounds most soluble in water will be those in which there is 
the greatest difference in the positive and negative character of the con- 
stituent parts (see Chapter V), so that the salts of a metal like silver 
will be less soluble than those of a metal like potassium. These rough 
conclusions are illustrated by the values in the following table: . 
Solubility at 20°, in 
mols per 1000 grams of 
Substance. 
Melting point. 
solution. 
CaCl2: 6H0O 
30° 
1.9 
H3BO3 
185° 
0.077 
AgNO3 
218° 
4.3 
NaNO3 
333° 
5.5 
AgCl 
455° 
0.00001 
KC1 
722° 
3.3 
BaCOs 
795° 
0.0001 
Ag2S 
830° 
0.00001 
PbSO4 
1100° 
0.00015 
CaF2 
1400° 
0.0002 
CaSiO3 
1510° 
BaSO4 
1580° 
0.00001 
98 A number of statements may be made regarding the compounds of 
particular ions such as the following; where it must be understood that 
only the commoner ones are considered: 
All nitrates are soluble. 
All acetates are soluble (AgC2H3O2 only moderately). 
All chlorides are soluble, except AgCl, HgCl, PbCl2 (the last is spar- 
ingly soluble in cold water, moderately soluble in hot. 
All sulfates are soluble, except BaSO4, PbSO4. (CaSO4 and AgSO4 
are sparingly soluble). 
All carbonates and phosphates are insoluble, except those of sodium, 
potassium and ammonium. (Many acid phosphates are soluble, e. g. 
Mg(H2PO4)2, Ca(H2PO4)2 and Ba(H2PO4)2.) 
All hydroxides are insoluble, except NaOH, KOH, NH4OH and 
Ba(OH)2; Ca(OH)2 is sparingly soluble. 
All sulfides are insoluble, except those of sodium, potassium and 
ammonium. Those of magnesium, calcium, barium and aluminum are 
not precipitated from solution because they are decomposed by water. 
All salts of sodium, ammonium and potassium are soluble, except 
K2PtCl6, (NH4)2PtCl6, K3Co(NO2)6, Na4Sb2O7. 
All silver salts are insoluble, except AgNO3, AgC2H3O2, Ag2SO4 
(the last two are only moderately soluble). 
It is upon the basis of solubilities that different ions present in the 
same solution may be separated. Suppose, for example, that a solution 
contained the nitrates of silver, barium, zinc and potassium. The addition 
of chloride ion, using, say a solution of ammonium chloride, would pre- 
cipitate the silver ion as AgCl, which could be filtered out. The addition, 
then, of sulfate ion, as by using ammonium sulfate solution, would pre- 
cipitate the barium ion as BaSO4. After this is filtered out, the zinc ion 
could be precipitated as sulfide, using ammonium sulfide. The excess of 
ammonium salts in the solution could finally be removed by evaporating 
the solution to dryness and heating the residue, which would volatilize 
the ammonium salts leaving only the potassium salts. 
Ionization of Water. It was shown in Chapter X that water is an 
extremely weak electrolyte, dissociating to a minute extent as follows: 
H2O = H+ + OH-. 
In pure water we have (H+)=(OH~)= IO-7 mols per liter. This is 
very slight indeed, but we shall see that it is very important. If an excess 
of either of the ions H+ or OH' is added to water, as in dissolving an 
acid or base in it, the result is a diminution in the concentration of the 
other. If, for example, 0.001 M HC1 is present, the concentration of the 
hydrogen ion is increased 10,000 times, which will cause that of the 
hydroxide ion to decrease 10,000 times, becoming 10'11. 
The concentration of H+ and OH- in a solution may be estimated 
approximately by the aid of indicators (see Chap. V, p. 45). The fol- 
lowing table gives the concentrations of H+ and OH' between normal 
H+ and normal OH', together with the corresponding colors shown by 
several of the common indicators: 
Acid 
Concentr. of H+ 1 0.1 0.01 IO'3 IO'4 IO'5 IO-6 
Concentr. of OH- IO"14 10~13 IO-12 10'11 10"10 IO'9 IO'8 
Methyl violet yellow green blue violet 
Methyl orange red orange yellow 
99 "Neutral 
Point" Alkaline 
Concentr. of H+ IO"7 IO"8 IO'9 IO"10 IO'11 IO-12 IO'13 IO"14 
Concentr. of OH"' IO"7 IO-6 IO"5 IO"4 IO"3 0.01 0.1 1 
Litmus red purple blue 
Phenolphthalein colorless violet colorless 
Tri-nitrobenzene colorless orange 
We may conclude, from the slightness of this ionization, that the 
reverse reaction will take place almost completely, so that whenever an 
acid and a base are mixed they will react almost completely to form water, 
leaving the other ions in solution, unless they happen to form an insoluble 
substance. We must therefore regard neutralization of acids and bases 
as taking place because of the fact that water is almost entirely undisso- 
ciated. We will return later to this topic in connection with the neutral- 
ization of weak acids and bases. 
Ionization of Weak Acids, Bases and Salts. The fact that many acids, 
bases and salts are but slightly ionized, as set forth in Chapter X, furnishes 
other reasons to expect certain reactions to take place, for the principles 
of equilibrium tell us that in such cases the corresponding ions will unite, 
when brought together, to an extent depending on the ionization of the 
substance. For example, because we know that acetic acid is weak, that 
is, but slightly ionized, we know that when hydrochloric acid, which gives 
a high concentration of H+, is mixed with sodium acetate, which gives 
a high concentration of C2H3O2~, these ions will combine till one or both 
is almost used up. Although sodium acetate is not a base, but a salt, we 
see that it has the effect of neutralizing high acidity. Similarly, because 
NH4OH is a weak base it can be prepared from any ammonium salt by 
the action of any strong base, since NH4+ and OH- unite rather completely. 
This property is used as a test for NH4+ and also in the manufacture of 
NH3. Considerable ammonia is given off when coal is distilled, as in 
making gas and coke. This is extracted from the gas by washing in 
sulfuric acid, forming (NH4)2SO4. When this salt is acted upon by 
slaked lime, Ca(OH)2, which gives a sufficiently high concentration of 
OH", there is set free NH4OH, which breaks up at the temperature applied 
into NH3 gas and H2O. It is evident, in general, that weak acids should 
be liberated from their salts in solution by the action of stronger acids, 
and that weak bases should be liberated in like manner by strong bases. 
An important application of this principle is in making "super- 
phosphate" for fertilizer. Tri-calcium phosphate (normal calcium phos- 
phate), Ca3(PO4)2 occurs in bones and mineral deposits, but is so 
insoluble that even though finely ground it does not furnish plants at all 
freely with the phosphate necessary for their growth. Accordingly the 
more soluble di-calcium phosphate is produced by the action of concen- 
trated sulfuric acid, as represented by the equation 
Ca3(PO4)2 + H2SO4 = 2CaHPO4 4- CaSO4. 
Where the ions of several weak acids are competing for an insufficient 
amount of hydrogen ion, it is evident that the weakest acid will be 
formed first, the next weakest second, etc. In a mixture of C2H3O,~ and 
CO3"", the values on page 77, Chap. X, or page 93, Chap. XII, make 
it evident that if H+ is added to the solution gradually, the first reaction 
to take place will be 
H+ + CO3" = HCO3- 
followed by H+ + HCO3" = H2CO3, 
followed by H+ + C2H3Oo" = HC2H.,O.,. 
100 Complex Ions. As a rule the positive ions in solution are very simple, 
consisting nearly always only of a metallic atom, with its ionic charge 
or charges. There are, however, some cases where these simple ions 
can unite with other ions or neutral molecules and still remain in solution 
as a more complex ion. They may best be considered in groups, accord- 
ing to the substances which are prone to form them. 
a. Ammonia complexes. A solution of ammonia in water yields 
all of the substances denoted in the equation 
NH3 + H2O = NH4OH = NH4+ +OH-. 
Ordinarily, when this solution is added to one containing an ion of a 
heavy metal there is enough OH" present to precipitate the metallic 
hydroxide. If, however, more of the ammonia solution is added, the 
concentration of NH3 will increase faster than that of OH", and there 
results, in the case of certain metallic ions, notably Cu++, Ni++, Cd++, Zn++, 
Ag+, Cu+, a solution which contains a complex positive ion containing both 
metal and ammonia. With the divalent cations the complex ions seem 
all to contain 4NH3, while those of the mono-valent, ions contain 2NH3, 
so that the formulas are as follows: Cu(NH3)4++, Ni(NH3)4++, 
CdNH3)4++, Ag(NH3)2+, Cu(NH3)2++. Since the NH3 is a neutral 
molecule its presence does not alter the original ionic charge. In the 
cases of copper and nickel these ammonia complexes are a deep blue in 
color, furnishing rather delicate tests for the presence of these metals. 
On account of this behavior the effect of an excess of ammonia is always 
to use up most of the free ions of these metals, with corresponding effects 
upon other equilibria, as will be explained later. 
b. Cyanide complexes. Cyanide ion has the power to form com- 
plexes with a large number of metallic ions. In this case the metals are 
contained in the cation, as illustrated by the equation for the formation 
of the silver cvanide complex ion: 
Ag+ + 2CN" = Ag(CN)2". 
(When Ag+ is present in any considerable amount AgCN is first precipi- 
tated, dissolving in excess of CN".) 
Among these complexes may be mentioned the following: Cu(CN)2", 
Zn(CN)4"", Au(CN)2", Au(CN)4", Pt(CN)4"", Pt(CN)6", Fe(CN)6"" 
Fe(CN)6-, Ni(CN)4-, Co(CN)e~", Co(CN)6-. 
c. Complex halides. Many metallic ions, especially those of the 
noble metals, have a tendency to form complex halides. This is particu- 
larly strong in the case of gold, platinum and the other "platinum metals," 
as shown in the following: AuC12", AuC14", PtCl6"~. 
d. Complex oxalates. Certain metallic ions form complex oxalates 
with oxalate ion, C2O4"~, as exemplified by Fe(C2O4)3 . 
Competition Between the Foregoing Factors. We have seen in a 
few instances how reactions are determined by competition between the 
various ions of weak acids, or the relative solubilities of various salts of 
the same ion. We are now prepared to discuss competition between the 
different factors, volatility, solubility, ionization of water, weak acids, bases 
and salts, complex ions, and to consider general examples of controlling 
reactions. 
a. Hydrolysis. When the salt of a weak acid is dissolved in water 
we have present in the solution an ion that has a great tendency to com- 
bine with hydrogen ion. Now water, though it is such a weak electrolyte, 
does yield some hydrogen ion. Consequently, there is a slight formation 
101 of the weak acid with a resulting decrease in the concentration of the 
hydrogen ion and a corresponding increase in the concentration of the 
hydroxide ion. To say this concretely, let us consider a solution of 
sodium acetate, which gives a large concentration of Na+ and C2H3O2~. 
Now, since the water yields a trace of H+ and OH", and since HC2H3O2 
is a weak acid, there will be some union of H+ and C2H3O2" to form the 
acid, liberating an excess of OH", so that the solution will be slightly 
alkaline. This may be expressed by an equation as follows: 
C2H3O2- + H2O = HC2H3O2 + OH- 
It will be observed that this reaction is the reverse of the neutralization 
of acetic acid with a strong base, and since the latter reaction is nearly 
complete, of course the former can take place but slightly under similar 
conditions. In writing the equation we show the net effect of what has 
happened, as expressed previously in words. We must guard against 
the error of thinking that the formulas on the right hand side of the 
equation tell what is in the solution after equilibrium is reached. The 
solution still contains mostly C2H3O2" and H2O, which fact is expressed 
by saying that the above substances react only slightly before equilibrium 
is reached, not by writing their formulas on both sides of the equation. 
Moreover, we should not write water as ionized in a single equation like 
this, for the bulk of it is un-ionized. As a matter of fact, there are 
two simultaneous reactions occuring, as follows: 
C2H3O2" + H+= HC2H3O2, and H,O = H+ + OH". 
The sum of these two give the former. 
Obviously there is a competition between C2H3O2" and OH" for the 
H+, as might be represented in this way: 
HC2H3O2 = C2H3O2- + H+ 
+ 
OH" 
H2O. 
That the OH" should get or keep most of the H+ is to be expected from 
the fact that water is vastly less ionized than is acetic acid. It is this 
disparity between the ionization of the two substances that enables 
C2H3O2~ to set free but a slight amount of OH~ from water, on the one 
hand, and which causes neutralization of HC2H3O2 by OH" to be nearly 
complete, on the other hand. 
If, instead of an acetate, we dissolve in the water the salt of some 
weaker acid, like sodium phenolate, NaOC0Hr„ (the "phenol sodique" of 
the drug store) we have in phenolate ion, OCGH5", one which will unite 
with the hydrogen ion of water more completely than will acetate ion, 
forming HOCGH5, the acid known as phenol, or carbolic acid, and 
setting free from the water more OH" than in the case of sodium acetate. 
The competition for H+ is here between OH" and OC6H5", and though the 
latter gets but little of it, since carbolic acid is much more ionized than 
water, it nevertheless gets more than does acetate ion. We may express 
this competition in the same manner as before, 
HOC6H5 = OC6H5- + H+ 
+ • 
OH" 
h2o. 
The net result we may write, as before, 
102 OC6H5- + H2O = HOC6H5 + OH-. 
The discussion shows why "phenol sodique" can be used as a mild anti- 
septic. Its solution always contains free carbolic acid, which is an 
antiseptic, and though its concentration is small, if it is used up more 
can be formed from the above equilibrium to replace it. 
A process such as this is called "hydrolysis," the idea underlying 
the name being that a salt is split up by water into free acid and base. 
We see that salts of weak acids always hydrolize in aqueous solution on 
account of the possibility of forming the free weak acid. Of course, the 
ions giving weak bases also take part in hydrolysis. All ammonium salts 
are hydrolyzed, on account of the formation of the weak base, ammonium 
hydroxide, as illustrated by the equation: 
NH4+ + H2O = NH4OH + H+ 
where the solution becomes slightly acid. If the salt is one of both , a 
weak acid and a weak base, like ammonium acetate, NH4C2H3O2, both of 
its ions take part in hydrolysis, as shown by the equation: 
NH4+ + C2H3O2- + h2o = NH4OH +HC2H3O2. 
In a case like this the acidity or alkalinity of the resulting solution depends 
upon whether the acid or the base is weaker. In this case they are of 
almost identical strength, as may be seen from the values on page 16, 
Chap. X, so that the solution is almost neutral, though it does contain 
more NH4OH than a solution of NH4C1, and more free HC2H3O2 than 
one of NaC2H3O2, since both of the ions present aid in splitting up the 
water. 
Hydrolysis plays an important part in many reactions, a few examples 
of which will be given. It is desirable that solutions to be used for 
household cleaning should have a slight alkaline reaction, since this aids 
in emulsifying grease and removing dirt, and also tends to soften hard 
water, as explained later. It is not desirable, however, to use a strong 
base like sodium hydroxide, because this gives too great a concentration 
of OH-, which has injurious effects on the hands and on some fabrics, 
unless used in very dilute solution, where the slight amount of OH- 
would soon be used up. What is wanted is a solution containing but 
little free OH-, but which is able to yield more should this be used up. 
Such is the case with ammonia, so that it finds large use in the home. 
We see also that salts of weak acids fulfill these conditions, and several 
of them find extensive application as cleansing agents. Sodium carbonate, 
Na2CO3, is the salt of a very weak acid, and its solutions give a very 
marked alkaline reaction, as follows: 
CO3" + H2O = HCO3- + OH-. 
On this account it is extensively known and used as "washing soda." 
There are two sodium carbonates, this one and NaHCO3, sodium acid 
carbonate. The solution of the latter also undergoes hydrolysis, 
HCO3- 4- H2O = H2CO3 + OH-. 
However, since H2CO3 is considerably more dissociated than HCO3-, 
according to the figures given on page 77, Chap. X, the former will be 
formed from HCO3- and H+ less completely than the latter will be from 
CO3-- and H+. Hence the solution of Na2CO3 is much more alkaline 
than one of NaHCO3, and is used as "washing soda," while the latter is 
used only where a much less alkaline reaction is desired, as for internal 
use. The NaHCO3, on the other hand, is "baking soda," because it can 
yield much more CO2 for the same weight both of itself and acid. 
Borax, sodium borate, is the salt of the very weak boric acid, and 
103 has the same effect in cleansing as sodium carbonate. Sodium silicate, 
the salt of weak silicic acid, behaves in the same way, and is a constituent 
of laundry soaps. 
Aluminum hydroxide, A1(OH)3, is not only insoluble, but is a weak 
base as well, so that aluminum ion, Al+++, in water forms some of the 
hydroxide, liberating the hydrogen ion and causing the solution to become 
distinctly acid. On this account aluminum salts can be used to liberate 
carbon dioxide from carbonates, as is done in the alum baking powders. 
A similar reaction is used by oil companies in a scheme for extinguishing 
gasoline or petroleum fires on oil tanks. The sodium carbonate solution 
and one of aluminum sulfate, containing glue to produce a froth, which 
retains the CO2 gas, are projected from adjacent pipes over the burning 
tank, when the froth produced spreads over the fire and quenches it. 
On account of hydrolysis of aluminum ion with ions of very weak acids 
aluminum salts of these cannot exist in the presence of water. When 
CO3", is added to Al+++, A1(OH)3 is precipitated instead of a car- 
bonate, and aluminum sulfide, prepared directly from the elements, decom- 
poses completely when put into water, as follows: 
A12S3 + 6H2O = 2A1(OH)3 4-6H2S. 
There are a number of other compounds that undergo complete 
hydrolysis, such as the chlorides of the non-metals. When PC13, phos- 
phorus tri-chloride, is put into water the following reaction takes place: 
PC13 + 3H2O = P(OH)3 4- 3H+ 4- 3C1". 
The hydroxide of phosphorus is so completely an acid rather than a base 
that the above reaction is complete. 
Bismuth chloride undergoes hydrolysis readily. In a case like this 
the hydrolysis doubtless takes place in steps, and we have successively, 
Bi+++, Bi(OH)++, Bi(OH)2+, or its anhydride, BiO+, and finally Bi(OH)3, 
or its anhydride. When bismuth chloride solution is used the chloride ion 
is capable of forming an insoluble precipitate with the bismuthyl ion, 
BiO+, so that the hydrolysis takes place as follows: 
BP++ 4- Cl- 4- H2O = BiOCl 4- 2H+. 
The addition of sufficient H+ of course reverses the reaction and dissolves 
the precipitate. Similar reactions are obtained with bismuth nitrate and 
with antimonous and antimonic chlorides. 
In titrating a weak acid with a strong base, or vice versa, it is 
necessary to determine, not when the solution is neutral, but when it 
contains equivalent amounts of acid and base, a minute proportion of 
which will be free, giving an alkaline or acid reaction, as the case may 
be. If sodium, acetate, for example, is dissolved in water, the solution is 
slightly alkaline, although equivalent amounts of acid and base are present. 
To reproduce this condition, when acetic acid is titrated with sodium 
hydroxide, one should use an indicator like phenolphthalein, which, as 
shown on page 8, changes color in a solution which is faintly alkaline, so 
that a slight excess of either acid or base would affect the indicator. 
Similarly, in titrating ammonia with hydrochloric acid, we wish to end 
with the solution not really neutral but slightly acid, since a solution of 
ammonium chloride reacts slightly acid. Hence an indicator like methyl 
orange, according to the table on page 99, would be suitable for this 
purpose. In general, the proper indicator to use in a titration may be 
found by taking the normal salt which will result from the titration, 
dissolving it in water, and determining what indicator will change color 
when a drop of solution of the free acid or base is added in excess. 
104 b. The Solution of Hydroxides. We have seen that when a more 
or less insoluble metallic hydroxide is in contact with water it gives to the 
water a certain amount of its ions, equilibrium being reached the sooner 
the less soluble the hydroxide. As an example we may consider magne- 
siub hydroxide, Mg(OH)2, which gives a small concentration of ion 
as follows: 
Mg (OH) 2 = Mg++ + 2OH-. 
We have seen also that hydrogen ion has a very great tendency to unite 
with hydroxide ion. If, therefore, H+ is added to the precipitate of 
magnesium hydroxide suspended in water, there will be a competition 
between the H+ and the Mg++ for the OH", which may be represented, as 
before, by the double equation: 
Mg(OH)2 = Mg++ + 2OH- 
+ 
2H+ 
2H2O. 
Now, since water is so extremely undissociated we might expect that 
the H+ added would be successful in stealing the OH- away from the 
magnesium ion, liberating the latter, which would thus go into solution so 
that the effect of the acid would be to dissolve the precipitate. The total 
reaction is therefore 
Mg (OH) 2 + 2H+ = Mg++ + 2H2O. 
As a matter of fact, water is formed so completely from its ions that it 
takes but a moderate concentration of acid to dissolve practically all 
hydroxides. 
Our principles of equilibrium would lead us to expect that any 
substance that would unite with OH- would have the effect of dissolving 
insoluble bases. Besides H+, we have found that NH4+ is a substance 
which has a tendency to unite with OH- to form the weak base NH4OH. 
We know, however, that these two ions do not unite nearly so completely 
as do H+ and OH-, so that we should not expect NH4+ to be nearly as 
effective as H+ in dissolving bases. It is true that the solubility of any 
base will be increased by the presence of NH4+, but if the base is very 
insoluble to begin with, the increase may not be sufficient to bring the 
amount in solution up to a value such that we would call it soluble, 
suppose, for example, that the solubilities of two bases in water are 0.001 
and 0.000,001 molal respectively, and that the addition of a certain amount 
of NH4+ is enough to increase each 1000 times. The effect would be to 
make the solubilities of the two bases now 1 molal and 0.001 molal, 
respectively. We would then call the first soluble in NH4+, but the 
second, insoluble. Among the insoluble bases which may be thus dis- 
solved in NH4+ and Mg(OH)2, Ca(OH)2, Zn(OH)2, Fe(OH)2, 
Mn(OH)2, Ni(OH)2, Co(OH)2. 
c. The Solution of Oxides. Oxides of metals are much like the 
corresponding hydroxides, and may either be converted into hydroxides, 
in some cases, or produced from them by heating. The solubility of 
hydroxides in H+ leads us to expect that oxides will likewise dissolve in 
acids. We may think here of the formation of water from the H+ of 
the acid and the oxygen of the oxide. Most oxides, indeed, are easily 
dissolved in acids, though more difficultly than are the corresponding 
hydroxides. This difficulty may be largely a matter of the speed of the 
reaction. In a few instances, notably Fe2O3, Cr2O3, and A12O3, if they 
105 have been heated to a very high temperature they become practically 
insoluble in acids. 
d. Amphoteric hydroxides. There are several metallic hydroxides 
that are capable of acting either as weak bases or weak acids. Any 
substances that behave in this way are called amphoteric. Such a sub- 
stance may give to water not only the metallic kation and the OH- of a 
base, but also H+ and a corresponding anion. Taking A1(OH)3 as an 
example, we may express this double power of dissociating as follows: 
A1++++3OH-=A1(OH)3 (or H3A1O3 = H++ H2A1O3- or H2O 4-AlO2"). 
As is the case with most weak polybasic acids, the various hydrogen ions 
dissociate in steps, and since H3A1O3 is an exceedingly weak acid, even the 
first H+ is ionized with difficulty, and the second and third are practically 
un-ionized unless a large excess of OH" is added. From the double 
equilibrium that we have represented it will be seen that the precipitate 
can be dissolved not only by the addition of H+, which would act as 
follows: 
A1(OH)3 + 3H+ = Al+++ + 3H2O, 
but also by the addition of OH", which would act in the following way, 
H3A1O3 + OH- = H2O 4- H2A1O3". 
It is possible, therefore, to have aluminum present either in an alkaline 
or in an acid solution, while in a solution approximately neutral it cannot 
exist, but will be precipitated as A1(OH)3. Other hydroxides behaving 
in this way are those of zinc, chromium, lead and tin (both stannous and 
stannic). This amphoteric character may be made use of in separating 
from each other metallic ions existing in the same solution. Thus Fe++* 
is commonly separated from Al+++and Cr+++ by adding concentrated OH" 
in excess, which leaves the first precipitated as Fe(OH)3 but dissolves 
the others as aluminate and chromite ions respectively. 
e. The Solution of Sulfides in Hydrogen Ion. The similarity 
between oxides and sulfides has been previously pointed out, page 6, 
Chap. V, also the fact that hydrogen sulfide is a very weak electrolyte, 
like water. It is but a slight transition, therefore, from the consideration 
of oxides to that of sulfides. The chief difference to be noted is that, 
whereas reactions like the following are very numerous, 
Zn++ + 2OH" = Zn(OH)2, 
the corresponding reaction, when sulfur is substituted for oxygen ,is the 
formation of a sulfide, rather than a hydrosulfide, as shown by the reaction 
Zn+++ 2SH"= ZnS 4-H2S. 
This is analogous to what takes place with OH" in the case of Ag+, and, 
when the solution is boiled, with Cu++, as shown by 
2Ag+ 4- 2OH" = Ag2O 4- H2O, and Cu++4-2OH" = CuO 4-H,O. 
The solution of a sulfide in an acid depends upon the removal of the 
sulfide ion, given to the water by the precipitate, by the hydrogen ion of 
the acid to form hydrogen sulfide, H2S, and its precipitation depends upon 
the presence of sufficient sulfide ion. .This competition between the 
metallic ion and the hydrogen ion for the sulfide ion may be represented 
by the scheme used before, 
ZnS = Zn++ 4- S"~ 
4" 
2H+ 
H2S. 
106 Some sulfides are so insoluble that the- metallic ion finds enough 
sulfide ion even when the solution contains a moderately concentrated 
acid. These sulfides, conversely, will not be dissolved by hydrogen ion 
in moderate concentrations. The sulfides which behave in this way 
include Ag2S, HgS, PbS, Bi2S3, CuS, As2S3, As,S5. Next, we have several 
that will dissolve if the H+ is concentrated but not if dilute, including CdS, 
Sb2S3, Sb2S5, SnS, SnS2. Then follow ZnS, MnS, FeS, NiS, CoS, which 
are so soluble that dilute H+ is sufficient to prevent their precipitation. 
The cases of NiS and CoS are peculiar in that they are not precipitated in 
acid solution, but nevertheless will not dissolve at all rapidly in dilute 
acid if first precipitated from an alkaline solution. This seems to be 
due to the existence of two modifications of the solids, so that when 
the more soluble modification is once precipitated it changes over into 
the insoluble modification, in the class with CuS, etc. In order to pre- 
cipitate this group of sulfides it is necessary to have a neutral or slightly 
alkaline solution so that sufficient sulfide ion can be present. The usual 
reagent is a solution of ammonium sulfide, made by passing H,S into 
NH4OH solution. If acetic acid is used, it is possible to reduce the con- 
centration of H+ sufficiently, by adding acetate ion, to allow the precipita- 
tion of ZnS. 
In dissolving a sulfide in H+, as represented by the equation, 
Sb2S3 + 6H+ = 2Sb+++ + 3H2S, 
we have been considering the effect of the concentration of H+, and seen 
that the less soluble the sulfide the more concentrated must be the H+ in 
order to dissolve it. Another factor may be introduced by virtue of the 
volatility of H2S. If the solution is boiled the H2S is more readily 
removed, so that more can form, resulting in an easier and more rapid 
solution of the sulfide. To dissolve the less soluble ones the use of con- 
concentrated H+, together with boiling the solution, is insufficient, and we 
shall see in Chapter XV that it is necessary to destroy the H2S by an 
oxidizing agent. 
The order of solubility of the above sulfides in water is roughly as 
follows, beginning with the most soluble: MnS, FeS, ZnS, NiS, CoS, 
SnS2, SnS, Sb2S3, CdS, PbS, Bi2S3, CuS, As2S5, Ag2S, HgS. 
f. Amphoteric sulfides. We have seen above how several hydroxides 
can dissociate in two ways, so as to be able to dissolve either in H+ or in 
OH-. There are, likewise, three sulfides which can dissolve in both OH- 
and SH-, namely, As2S3, As2S5, Sb2S3, Sb2S5, SnS2. These same elements, 
it will be recalled ,give oxides and hydroxides which dissolve in alkalies. 
The analogy between the oxides and sulfides may be brought out in 
the following parallel reactions: 
H3AsO4 4- 3OH-= AsO4-+ 3H2O; 
As2O5+ 6OH-= 2AsO4--f- 3H2O; As2S5+ 6SH-= 2AsS4-+ 3H2S; 
As2O5+ 2OH-= 2AsO3~ + H2O; As2S5+2SH-=2AsS3- +H2S. 
Since there is undoubtedly more S in an alkaline sulfide solution 
than there is O-- in the presence of OH-, we sometimes prefer to express 
the reaction for the solution of these sulfides by equations such as the 
following, which are a little simpler than the above, though they do not 
bring out the analogy with the corresponding oxygen compounds quite as 
well: 
As2S54-3S" = 2AsS4- or As2S5 4-S" = 2AsS3- 
As2S3 4- 3S~ = 2AsS3- or As2S3 4- S-- = 2AsS2- 
SnS2 4- S--=SnS3". 
107 The compounds in which the metals have the lower valence are 
always less acidic, both in the case of the oxides and the sulfides, and 
Sb2S3 dissolves with some difficulty in ammonium sulfide ,and SnS .is 
practically insoluble in the same. In dissolving both of these sulfides 
yellow ammonium sulfide is used, which contains polysulfide ions, like 
S2~. This is capable of dissolving the lower sulfide of all three metals, 
changing it to the sulfo-ion of the higher valence, thus, 
Sb2S3 + 2S2- + S" = 2SbS4-, or Sb2S3 +S3" = 2SbS3- 
SnS + S2--= SnS3". 
When acid is added to the solutions of the above sulfo-ions the weak 
acid, H,S, is produced, reprecipitating the sulfides, just as the addition of 
acid to stannate ion reprecipitates the stannic hydroxide. This reprecipi- 
tation consists simply in the reversal of the reactions whereby these 
sulfo-ions are produced. 
The different behaviors of sulfides towards the H+ and SH" or S" 
is made the basis of important separations in both qualitative and quan- 
titative analysis. If a solution contained the following ions, Pb++, Bi+++, 
Cu++, As+++, Sb+++, Sn++, Zn++, Mn++, Mg++, they could be divided into groups 
by controlling the concentrations of the substances in the equilibrium. 
H2S = H+ + SH- = 2H+ 4- S-. 
This separation can best be outlined by the aid of a diagram. 
Have dilute H+ (HC1) present, and pass in H2S. 
Precipitate insoluble 
PbS, Bi2S3, CuS, As2S3, Sb2S3, SnS. 
Treat with S2" 
(yellow ammonium sulfide) 
In solution, 
Zn++, Mn++, Mg++ 
Make alkaline, add S-_(SH_) 
(ammonium sulfide) 
Left insoluble, 
PbS, Bi2S3, CuS. 
In Solution, 
AsS3", SbS3", SnS3"" 
Reprecipitate the sulfides 
and heat with 12NHC1 
Precipitate 
ZnS, MnS. 
In solution 
Mg++. 
Left 
insoluble, 
As2SB 
In solution, 
Sn++++, Sb+++, 
dilute to 2 N H+, 
pass in H2S cold, 
ppt. Sb2S3. 
g. Salts of other weak acids. Salts of all weak acids are made 
more soluble by the addition of H+. This increase in solubility, in the 
case of carbonates, is enough to bring all insoluble carbonates into 
solution in dilute H+. A very interesting case of the solution of a 
carbonate was alluded to on page 87, Chap. XII. It is the reaction 
whereby hard water is produced from limestone rock. If very dilute 
H+, as is present in very dilute HO, is allowed to act upon CaCO, the 
CO3~- which the latter gives to the solution in very slight amount is con- 
verted into HCO3_ only, and not into H2CO3, as would be the case if 
more concentrated acid were used. The equation for this reaction is 
as follows: 
CaCO3 + H+ + Cl- = Ca++ + HCO3" + Cl". 
If H2CO3 is used instead of HC1, it also can furnish sufficient H+ to 
bring about essentially the same reaction, 
CaCO3 + H2CO3 = Ca++ + 2HCO3-. 
Since most surface water contains some dissolved carbon dioxide, the 
above reaction takes place whenever such water flows over limestone. 
108 A quite analogous one occurs with MgCO3. The reaction can be 
reversed by removing the H2CO3, either by boiling the solution, as 
occurs in a tea-kettle, whence the name "temporary hardness" or else 
by neutralizing it with some alkali, or some substance giving OH" by 
hydrolysis, like sodium carbonate or borax. When this is done on a 
large scale, the water is analyzed and just the right amount of Ca(OH)2 
is added. Although more Ca++ is added in this way, there is produced 
enough CO3~~ to precipitate all of it, the total reaction being represented 
by the equation 
2Ca++ + 2HCO3- + 2OH- =.2CaCO3 + 2H2O. 
When the hardness in water is produced by dissolving a salt like 
CaSO4, then it is necessary to add CO3"" (as Na2CO3) to precipitate 
the Ca++. 
When the solubility of the hydroxide of a metal is about as small 
as the solubility of its carbonate, the latter will hydrolize partly to give 
a basic carbonate. This is the case with the carbonates of copper, lead 
and mercury, as illustrated by Pb(OH)2 • 2PbCO3, the "white lead" of 
the painter. 
There are no phosphates so insoluble in water but that they are 
dissolved by moderately concentrated H+. Although PO4 is the ion 
of the very weak acid, HPO4__, nevertheless dilute H+ is not sufficient 
to dissolve all of the normal phosphates, both because some are so 
insoluble to begin with that the increase is not sufficient to make them 
soluble in the usual sense, and also because some mono-hydrogen phos- 
phates, like CaHPO4 are rather insoluble, and even though Ca3(PO4)2 
be converted into the former, solution does not result. 
As a rule the acid salts are more soluble than the normal salts, as 
exemplified by the series of calcium phosphates. Ca(H2PO4)2 is soluble, 
CaHPO4 is insoluble, and Ca3(PO4)2 is very insoluble in water. The 
chief exception to this rule is furnished by the sodium carbonates, where 
the acid salt is the least soluble. 
The industrial preparation of sodium carbonate by the Solvay process 
furnishes an interesting application of the principles here discussed. The 
normal salt, Na2CO3, is very soluble, but the acid salt, NaHCO3, is 
not very soluble in the cold. The latter can therefore be precipitated by 
bringing together the ions Na+ and HCO3" in sufficient concentration. To 
do this NH3 gas and CO2 gas are led into a cold concentrated solution of 
NaCl. The acid and base produced by the solution of the gases in water 
react to give NH4+ and HCO3_, provided they are used in the right pro- 
portion, and when the concentration of HCO3_ reaches a sufficient value 
it begins to precipitate the Na+ as NaHCO3.* 
h. Solution by forming weak salts. In the previous cases we have 
treated compounds which could be brought into solution by adding some 
substance, especially H+, which would use up the anion of the precipitate. 
It is equally possible to dissolve a compound by adding some reagent 
that will use up the free kation. Thus lead sulfate, PbSO4, may be dis- 
solved by the addition of an acetate, because lead acetate is a soluble 
weak salt, and the acetate ion, if in sufficient concentration ,can compete 
* This solid is separated from the solution and heated to produce Na2CO3 
(cf. p. 113). The CO2 given off during the heating is returned to the process. A 
fresh supply is obtained by heating CaCO3, the CaO resulting being slaked to give 
Ca(OH)2, which is used to recover the NH3 from the NH4+. 
109 successfully with the sulfate ion for the lead ion, as shown by the equation 
PbSO4( solid)+ 2C2H3O2- = Pb(C2H3O2)2 + SO4'- 
Similar effects are shown, at least to some extent, wherever weak 
salts, such as are mentioned on page 78, can be formed. We find, for 
example, that Cdl2 is a moderately weak salt, consequently the following 
reaction can be made to proceed to some extent by using a sufficient con- 
centration of 1": 
CdS + 2I-=CdI24-S". 
i. Solution by forming complex ions. Wherever the kation is capable 
of forming a complex ion, a precipitate can be rendered more soluble than 
in water. Thus, a salt like ZnCO3 can be dissolved, not only by using up 
the anion, as in adding H+, but also by using up the kation, as when 
Zn(NH3)4++ is formed by the addition of ammonia solution. The com- 
petition between the CO3-- and the NH3 may be represented as before: 
ZnCO3 = Zn++ -|- CO3-- 
+ 
4NH3 
Zn(NH3)4- 
If the concentration of the ammonia is sufficient the complex ion is 
formed at the expense of the insoluble carbonate, resulting in the net 
reaction 
ZnCO3 + 4NH3 = Zn(NH3)4++ + CO3". 
In the case of ZnS, which is much less soluble, the formation of the 
complex ion cannot take place to an appreciable extent, so that ZnS is 
not dissolved by ammonia. The copper ammonia complex ion seems to be 
formed more completely than the corresponding zinc ion, so that traces 
of Cu(NH3)4++ can be formed from CuS if concentrated ammonia is 
used, in spite of the smaller solubility of CuS. 
j. Solution by forming ions of amphoteric hydroxides. The effect 
of a considerable concentration of OH' may be to dissolve the relatively 
insoluble salts of certain metals, in a fashion similar to that whereby 
complex ions are formed. In fact, these ions may be considered as complex 
ions. Most of the lead salts, for example, are dissolved by OH-, as ex- 
emplified by the reaction 
PbSO4 + 2OH- = H2O + HPbO2- + SO4". 
Even ZnS is not so insoluble but that it can be dissolved by concentrated 
OH-. 
k. Systematizing the reactions of a given ion. It is not necessary 
to perform and remember all of the reactions of a certain ion with a set 
of reagents in order to tell what they will be. It suffices to make an 
intelligent selection of a few of the possible reactions and arrange the 
resulting information systematically. As an example of what is meant 
let us consider a number of the reactions of silver ion, Ag+, with the 
following: CO3", OH-, CN-, Cl-, Br-, I-, S--, NH3, S2O3--. Let us 
assume that we have performed experiments which show that the reac- 
tions take place which are represented by the following equations as 
read from left to right: 
AgCl + Br- = AgBr + Cl" 
AgCl + 2NH3 = Ag (NH3) / + Cl- 
' Ag(NH3)2+ + I- =AgI4-2NH3 
AgoO + H2O + 2C1- = 2AgCl + 2OH- 
Agl + 2CN- =Ag(CN)2- + I- 
110 Ag,O 4- H2O 4- CO3-- = Ag2CO3 4- 2OH- (easily reversed) 
AgBr 4- 2S2O3- = Ag(S2O3)2 4~ Br 
Ag (S2O3) 2- 4-1' = Agl 4- 2S2O3- 
AgCl 4- CN- = AgCN 4- Cl- 
AgBr 4- 2NH3 = Ag(NH3)2+4-Br- (easily reversed) 
2Ag(CN)2"4- S~ =AgoS4-HCN- 
Ag2CO3 4- 2C1" = 2AgCl 4- CO3-- 
AgCN'4- 2NH3 = Ag(NH3)2+ 4- CN- 
From these observations it is possible to arrange the following list 
of substances in the order showing the completeness with which they 
unite with Ag+ to form a precipitate or a complex ion as the case may be, 
it being understood that the reagents forming complexes are taken in 
moderate excess: 
Substance. 
OH- 
co3-- 
ci- 
CN- 
nh3 
Br- 
S2O3 
I- 
CN- 
s- 
To form 
Ag2O 
Ag2CO3 
AgCl 
AgCN 
Ag(NH3)2+ 
AgBr 
Ag(S2O3)2™ 
Agl 
Ag(CN)2- 
Ag2S 
In order to make this list it is necessary to make eight or more 
observations, and from it it is possible to predict the direction of forty-five 
reactions. The list was made in such a way that each substance in the 
first line will steal Ag+ away from its combinations with any substance 
to the left, because less Ag+ can exist in the same solution with the former 
at a given concentration than with the latter at the same concentration. 
Thus AgBr gives to the solution more Ag+ than can exist together with 
I-, hence the following reaction will readily take place: 
AgBr -f- 1" = Agl Br-. 
For the same reason, Agl will dissolve in CN-, and also any precipitate 
in the above list can be transposed into Agl by I- except Ag2S. Where 
two substances have about the same power to combine with Ag+, as is the 
case with NH3 and Br-, it is possible to reverse the reaction by altering 
the concentrations. Thus the reaction 
AgBr 4- 2NH3 = Ag(NH3)2+ 4- Br-, 
can be made to take place as read from left to right or from right to left, 
if the concentration of NH3 or that of Br-, respectively, is taken in 
sufficient excess. 
By adding to the information given by the above list a few facts 
about the other reactions of the substances, we are able to predict a very 
large number of reactions. For example, HCN is a very weak acid, 
hence AgCN can be dissolved by H+. For the same reason, H+ will 
liberate Ag+ from a solution of Ag(CN)2-, and if Cl- is present along 
with the H+, then AgCl will be precipitated as follows: 
Ag(CN)2- + 2H+ 4- Cl-= AgCl 4- 2HCN. 
In similar fashion, AgCl will be dissolved by NH3, according to the 
reaction represented earlier, but will be reprecipitated on the addition of 
H+, which takes up the NH3 to form NH4+, allowing the Ag+ to recombine 
with the Cl- left from the previous reaction, as shown by the equation: 
111 Ag(NH3)2+ + 2H+ 4- Cl- = AgCl + 2NH3. 
Again, since Ag2SO4 is moderately soluble (see page 7) we see that 
it can be readily transposed to AgCl by dissolving it in water and adding 
Cl", as AgCl is insoluble. In order to reverse this transposition we see 
that some other principle besides solubility must be invoked. If water is 
absent, then we can take advantage of the greater volatility of HC1 as 
compared with HQSO4, so that the reaction, 
' 2AgCl + H2SO4 = Ag,SO4 + 2HC1, 
can be brought about by heating AgCl with concentrated H2SO4. 
Miscellaneous Transformations. The general application of the above 
principles and information can be illustrated by further miscellaneous 
examples. 
To change CaSO4 to CaCO3 we can take advantage of the greater 
solubility of the former (see page 99) and treat it with a solution of a 
soluble carbonate, like Na2CO3, which gives CO3--, when the insoluble 
CaCO3 would be precipitated. The manner of expressing this in an 
equation would depend upon whether the amount of water present is 
such that the CaSO4 is mainly present as solid or as dissolved ions. These 
cases would be represented respectively as follows: 
CaSO4 + CO3" = CaCO3 + SO4" 
Ca++ 4- SO4- + CO3" = CaCO3 + SO4" . 
To bring about the reverse transformation we would have to make 
up for the greater solubility of CaSO4, which shifts the equilibrium 
towards CaCO3, by adding H+, which would use up CO3-- and reverse 
the reaction. 
To change ZnSO4 to ZnCl2 it is necessary to replace SO4__ by Cl-. 
This might be done by adding BaCl2 to a solution of the ZnSO4, giving the 
reaction, 
Zn++ 4- SO4" 4- Ba++ 4- 2CB = BaSO4 4- Zn++ 4- 2C1". 
If exactly the right amount of BaCl2 were used, all of the Ba++ and SO4~~ 
present would precipitate each other, leaving in solution Zn++ and Cl-, and 
if the precipitate were filtered out and the solution evaporated these ions 
would combine to give solid ZnCl2, as desired. However, it is difficult 
to add just the equivalent amount of BaCl,, and it is accordingly better 
to use a different method, with reagents an excess of which can be easily 
removed. Since Zn(OH)2 is insoluble, it can be precipitated by adding 
OH- (e. g., NaOH solution) to the solution of ZnSO4. If the precipitate 
is now filtered out, the SO4" is disposed of. The precipitate of Zn(OH)2 
can now be dissolved by HC1 solution, since the H+ of the acid unites 
so strongly with the OH" of the base, giving a solution containing Zn++, 
with H+ and Cl- in excess. If, now, this solution is evaporated, the excess 
of H+ and Cl~ go off along with the water, as HC1 gas, leaving finally 
solid ZnCI2. By precipitating the carbonate, ZnCO3, instead of the 
hydroxide, the same transformation can be made ,since ZnCO3 is 
equally soluble in the stronger acids. In fact, a soluble salt of any metal 
which has an insoluble hydroxide or carbonate can be transposed to 
another soluble salt of the same metal by precipitation of one of these 
and dissolving it in the acid corresponding to the desired salt. 
Aluminum chloride, A1C13, can be changed to aluminum oxide, A12O3, 
by first precipitating A1(OH)3, using a soluble hydroxide, e. g., NaOH, 
NH4OH, and then heating, which aids the decomposition of the hydroxide 
into the oxide and the volatile substance steam. The two processes are 
represented by the equations 
112 Al++++ 3OH_ = A1(OH)3, and 2A1(OH)3 = A12O3 + 3H2O. 
To change sodium acid carbonate, NaHCO3, into sodium carbonate, 
Na2CO3, two methods are available. The salt can be dissolved in water 
and treated with 1 equivalent of NaOH, when the HCO3~ will be 
neutralized by the added OH", as shown by the equation, 
Na+ + HCO3- + Na+ + OH- = 2Na+ + CO3" " + H2O. 
The resulting solution yields solid Na2CO3 on evaporation. Again, the 
acid carbonate is capable of another kind of decomposition, represented 
by the equation 
2NaHCO3 = Na2CO3 + H2O + CO2. 
Since two of the substances produced by this reaction are volatile, the 
reaction can be made to proceed so as to produce them by heating the 
mixture in an open vessel so that they can escape. 
To change Na2CO3 to NaOH, it is necessary to remove the CO3"- 
and supply OH". Since both sodium compounds are soluble, the insolu- 
bility of some other substances must be invoked. We need a hydroxide 
of some metal which is more soluble than the corresponding carbonate. 
The best one for this purpose is Ca(OH)2, which is itself scarcely soluble, 
especially in a solution of NaOH. The insolubility of CaCO3 causes 
the following reaction to take place, while any evcess of Ca(OH),, 
together with the CaCO3 produced, can be filtered off: 
Ca(OH)2 + 2Na+ + CO3"= CaCO3 + 2Na+ + 2OH~. 
Iron rust can be dissolved by moderately dilute hydrochloric acid, as 
shown by the equation: 
Fe2O3 + 6H+ = 2Fe+++ + 3H2O. 
In dissolving a rust or ink stain off a delicate fabric it is not desirable, 
however, to use such concentrated acid. More dilute H+ can be used if 
the reaction is aided by the removal of one of the products. This is 
accomplished by using oxalic acid, .which both furnishes dilute H+, since 
it is a rather weak acid, and also oxalic ion, C2O4"_, which unites with 
the free ferric ion to form a complex ion (see page 101), the reaction 
being 
Fe2O3 + 6H2C2O41 = Fe(C2O4)3- + 3H2O + 6H+. 
Exercises. 
1. Under what conditions would you expect the following reactions 
to take place: 
CaCO3 = CaO + CO2. 
Fe2(SO4)3 + 6NaCl = 3Na2SO4 + 2FeCl3. 
2. If a solution contains the following ions, how can they be 
separated: Hg+, Pb++, Cu++, Ca++ ? 
3. How can the following ions be separated when in the same 
solution: SO4~_, Cl-, C2H3O2", CO3-~? 
4. What is formed when ammonia is added, first in small amount, 
then in excess, to each of the following ions: cupric, zinc, aluminum and 
silver? 
5. Name all the substances present in a dilute solution of sodium 
acetate, and tell whether each is present in relatively large or small 
amounts. 
6. What class of salts gives an alkaline reaction when dissolved in 
water ,and what class gives an acid reaction? Give examples of each. 
7. State which of the following ions would take part in hydrolysis, 
and write equations for two cases: Cl-, SO4", NH4+, Al+++, Ba++, CO3" ". 
113 8. Which of the two salts, NaH2PO4, and Na2HPO4, would be more 
hydrolyzed when dissolved in water? 
9. Which would be more hydrolyzed, a solution of sodium carbonate, 
or one of sodium acid carbonate of the same concentration? 
10. What is the effect on a water glass (sodium silicate) solution 
of (a) exposure to the CO2 in the air; (b) the addition of a solution of 
ammonium chloride? 
11. The following substances are relatively insoluble in water; which 
would dissolve in a solution of HC1? (Give equations and explain your 
answers.) Mg(OH)2; BaSO4; CaCO3. 
12. How would you prepare a pure silver compound and a pure 
copper compound from a mixture of solid silver chloride and cupric 
hydroxide? Write equations. 
13. In the following list of ions - H+, Ba++, Ag+, NH4+, SO4--, 
Ac-, Cl-, CO3--, CN-, OH- -state each pair that would tend to combine 
largely if brought together in dilute solution. In each case give reason 
for your answer and write equation. 
14. What is an amphoteric substance? 
15. Explain the solubility of lead sulfate in (a) OH-; (b) Ac-. 
16. Can the following substances be present in moderate concen- 
tration in the same solution? If not, what is formed? H+ and NO3-; 
H+ and OH-; H+ and SO4--; H+ and CO3--; H+ and C2O4--; Ca++ and 
CO3--; Ca++ and H2CO3; OH- and H2CO3; H2CO3 and CO3--; H+ and 
H,CO3; HAc and CO3--; Al++ and OH-; Cu++ and NH4OH; Cu++ and 
NO3-; Ag+ and NH3; Hg++ and Cl-. 
17. What is the effect of (1) NH4C1 solution, (2) HNO3 solution, 
(3) H2CO3 solution, upon a precipitate of Ca(OH)2? Explain in each 
case. 
18. If, when 2 mols of NaCl and 1 mol of H2SO4 are brought 
together, the reaction ceases when the HC1 gas produced exerts a pressure 
of 10 atmospheres at 77° in a volume of 2 liters, (a) how many mols 
each of NaCl and H2SO4 have reacted; (b) what fractioon of each has 
reacted; (c) how might the quantities reacting be increased? 
19. How do you account for the fact that antimony sulfide is soluble 
in concentrated hydrochloric acid, while arsenic sulfide is not? 
20. How could you separate arsenic from silver in a mixture of 
arsenous sulfide and silver sulfide? 
21. Which of the precipitates, AgCl, Agl, Ag„O, Ag2S, are soluble 
in (a) HNO3; (b) NH3-aq; (c) CN-? 
22. Cupric hydroxide is readily soluble in ammonia while cupric 
sulfide is only slightly soluble in the same reagtent. Discuss the reactions, 
and explain the difference in behavior. 
23. Starting with metallic calcium, show by means of equations how 
you would transform it into calcium oxide, the oxide into chloride, chloride 
into carbonate, and carbonate again into oxide. State, with each reaction, 
the conditions for carrying it out. 
24. Stating with cupric nitrate, how would you prepare in succession: 
cupric sulfate, cupric oxide, cupric chloride, cupric carbonate, cupric 
hydroxide ? 
25. Tell how you would determine, in each of the following cases, 
whether the impurity named is present or absent. Name reagents used 
and describe the essential observation in your experiments: 
Sulfuric acid in vinegar. 
KC1 in common salt. 
114 (NH4)2SO4 in concentrated H2SO4. 
Dissolved zinc in a solution of NaOH. 
BaSO4 in ZnO. 
26. Tell how you would make each of the following transformations, 
getting the final product as a pure solid. State the reagents used and 
explain why you would expect your method to work: 
NaCl to NaoS04. 
CuSO4 to CuCl2 
Ca++ + 2HCO3- to CaCO3. 
Na2CO3 to NaOH. 
Ag(NH3)2+ + NO3- to AgCl. 
NaOH to Na2CO3. 
Zn(NO3)2 to ZnO. 
BaCl2 to Ba(C2H3Oo),. 
NH4C1 to NaCl. 
115 CHAPTER XIV. 
CHEMICAL EQUILIBRIUM. THE EFFECT OF PRESSURE AND 
TEMPERATURE. 
When chemical equilibrium has been reached it is possible, as we have 
seen in the two previous chapters, to bring about a further reaction by 
altering the concentration of one or more of the reacting substances, which 
causes that reaction to proceed which will tend to restore the original 
concentration, or, as we may say, to neutralize the effect of the change. 
It is likewise possible to affect equilibrium by changing the total pressure 
to which the system is subject, and by changing its temperature. We 
will now inquire into the direction in which a reaction must proceed in 
order to restore equilibrium when subjected to changes in pressure or 
temperature. 
Effect of Changing Total Pressure upon a System in Equilibrium. 
If we have an equilibrium between a liquid and its vapor, as discussed 
in Chapter II, we have seen that an increase in the pressure tends 
to cause vapor to condense, since the number of gaseous molecules is 
thereby diminished, relieving the pressure. We have seen, also, that this 
is in harmony with the kinetic theory. Now, the same effect is always 
produced, no matter how complex the equilibrium. An increase in 
pressure always tends to produce the system having the smallest volume, 
a decrease in volume tending to relieve the increase in applied pressure. 
Whenever the total pressure is incrased it increases the concentration of 
all of the substances, but this increase in concentration is greatest for 
those substances having naturally the largest volume, hence their tendency 
to react is increased to a greater extent than is that of the substances 
having a smaller volume, so that the former react more rapidly than the 
latter, resulting in a shift of equilibrium whereby the set of substances 
having the largest volume tends to be used up. 
Suppose, for example, that we have a closed vessel containing SO3, 
SO2, and O2 in equilibrium, which has been reached by the substances 
reacting in one direction or the other as represented by the equation 
2SO2 + O2 = 2SO3. 
We see from the equation that when SO3 is formed there is a change 
in the number of molecules in the proportion 3 to 2. If pressure is 
applied to the three gases in equilibrium it will compress them, resulting 
in a decrease in volume. A further decrease is possible, however, by 
the union of some of the SO2 with some of the O2, and the consequent 
diminution in the number of molecules, relieving somewhat the pressure 
applied. Starting, in two experiments, with identical amounts of SO2 
and O2 at the same temperature, with a catalyst present, or else waiting 
for equilibrium to be established, more SO3 will be produced when 
equilibrium is finally attained if the pressure is high than if it is low. 
The same conclusion may be drawn with respect to the reaction 
N2 + 3H2 = 2NH3, 
which shows that ammonia is formed with a diminution in the number 
of molecules. A high pressure would favor that reaction which would 
yield the smallest volume, thus relieving, as far as possible, the high 
pressure applied. 
On the other hand, with the equilibria 
116 N2 + O2 = 2NO, and H2-ff Cl2 = 2HC1, 
there being no change in the number of gaseous molecules during a 
reaction, pressure will be without effect on the amounts formed. 
In predicting the effect of pressure we must have regard to the 
physical states of the substances involved under the conditions of the ex- 
periment. It is possible, for instance, to make H2S directly from its 
elements, using either liquid sulfur or sulfur vapor. The respective 
reactions may be represented by the equations 
Ho + S(liquid)= HoS 
2H2 + S2(vapor)= 2H2S. 
In the first reaction there is practically no change in volume, if the 
reaction is done at constant pressure, as the volume of the liquid sulfur 
is negligible, so that the relative amounts of H2 and H2S present at 
equilibrium will be practically independent of the pressure applied, or of 
the volume of the vessel in which the reaction takes place. In the second 
process, where the conditions are such that the sulfur is in the vapor 
state, the formation of H2S at constant pressure results in a decrease 
in volume, hence the higher the pressure the greater the amount of H2S 
formed. 
Not only is the effect on the direction indicated by the above 
discussion, but its magnitude also depends upon the magnitude of the 
volume change during the reaction. Thus a tenfold increase in total 
pressure will produce a much greater increase in the amount of NH3 
produced from N2 and H2 than in the amount of SO3 produced from 
SO^ and O2, since the decreases in volume in the two reactions are 
respectively from 4 to 2 and from 3 to 2. For the same reason very 
great changes in pressure have to be applied in order to have much 
effect on equilibria between liquids and solids only, where only small 
changes in volume occur. 
Effect on Equilibrium of Changing the Temperature. In order to 
appreciate the effect of changing temperature upon a system in equilib- 
rium, we must realize that the immediate effect of any difference in 
temperature is to determine the flow of heat. If we have a tank of 
water at 20°C, and put into it a piece of iron at a temperature of 30°, 
heat will flow from the iron to the water in the tank, tending to equalize 
the temperature. If, on the other hand, the piece o.f iron is at a temper- 
ature of 10°, when it is put into the water heat will flow from the water 
into the iron. Suppose, now, that we have a closed vessel containing 
SO2, O2 and SO3, in equilibrium, in an oven at a temperature of 500°, 
and we remove it to another oven in which the temperature is 550°. 
Just as with the piece of iron in the former case, so here heat will flow 
into the vessel and its reacting mixture. The immediate effect of this 
heat will be to increase the temperature, the heat content, of the enclosed 
gases. There is, however, another way in which heat may be absorbed 
by the mixture, that is by the reaction preceding in the direction which 
absorbs heat.. The direction of the reaction for heat absorption is shown 
by the complete thermochemical equation 
2SO2 ff- O2 = 2SO3 + 45000 cal. 
Since heat is evolved by the formation of SO3, it will be absorbed by the 
reverse reaction, in which SO3 is dissociated into SO2 and O2, hence the 
effect of bringing the reacting mixture to the higher temperature is to 
shift the equilibrium so as to form less SO3. 
If, on the contrary, the vessel with its equilibrium mixture is put 
into an over at 450°, the heat will flow out of it, which heat will be 
117 furnished not only by the cooling of the contained gases, but also by 
their reacting to some extent to give off heat, accordingly more SO3 is 
present at equilibrium at the lower temperature. 
In the formation of NO from its elements heat is absorbed, as shown 
by the equation 
N2 + O2 = 2NO - 44000 cal. 
Consequently, if we have these gases in equilibrium, a rise in temperature 
will favor the absorption of heat in every possible way, which will not 
only cause the temperature of the gases to increase, but will also tend 
to form more NO. We see here that the higher the temperature the more 
NO will be obtained at equilibrium, while with SO3 the reverse is true, 
the higher the temperature the less SO3 is obtained at equilibrium. 
Considerations of this sort apply to every sort of equilibrium and 
enable us either to predict the effect of changing the temperature from 
the sign (and magnitude) of the heat of reaction, or, conversely, to predict 
the latter from the former. As an illustration of this reverse process, 
let us consider a saturated salt solution where we find that the higher 
the temperature the more soluble is the salt, and ask ourselves whether 
this salt evolves or absorbs heat when more dissolves under equilibrium 
conditions. We have found by experiment that raising the temperature 
causes more salt to dissolve. From our theory we conclude that raising 
the temperature also favors any process that can take place with absorption 
of heat. Putting these two statements together we conclude that this 
salt absorbs heat on dissolving, since its solubility increases with the 
temperature. By similar reasoning, any salt that became less soluble 
as the temperature increased must have positive heat of solution, that is, 
it must evolve heat on dissolving. 
Simultaneous Consideration of All Factors Governing Reactions. 
In Chapter XI we considered the velocity of chemical reactions, as it is 
affected by concentration, temperature and the presence of catalysts. We 
have now considered separately the effect of concentration, pressure and 
temperature upon chemical equilibrium and are now prepared to consider 
the simultaneous effect of these factors in controlling chemical reactions. 
We must again emphasize that the question of velocity is quite distinct 
from that of equilibrium. If we should put some calcium fluoride, 
CaF2, and some calcium sulfide, CaS, into separate beakers of water, 
we might conclude from a hasty examination that both are very insoluble. 
As a matter of fact, the small amount of the solid going into solution is 
due to very different reasons in the two cases. Calcium fluoride is truly 
insoluble, and no amount of time or stirring would cause more than a 
very small amout of it to dissolve. On the other hand, calcium sulfide 
appears to be insoluble because, in order to dissolve, it must hydrolyze, 
giving Ca(OH)2 and Ca(SH)2, a process which takes place very slowly. 
Again, calcium chromate, CaCrO4, is more soluble in cold water than 
in hot, if sufficient time is allowed for the solution to become saturated, 
but it will dissolve much faster in hot water than in cold, so that after 
a given time more might be found in a hot than in a cold solution, although 
eventually the latter would contain more. It is thus very important, in 
attempting to realize difficult reactions, to remember that speed and equlib- 
rium are two different considerations, and that conditions which favor 
the desired equilibrium may not favor its rapid attainment, and vice 
versa. A discussion of the complete conditions for realizing certain 
important reactions will make this clearer. 
a. The "Contact Process" for Making Sulfuric Acid. We have 
118 seen above that to make SO3 from SO2 and O2 a low temperature is 
desirable in order to give the most complete yield of SO3. By recalling 
the discussion in Chapter XI, however, we will see that the lower the 
temperature the more slowly will the desired equilibrium be reached. 
We must therefore distinguish between the amount of SO3 which would 
be produced in infinite time and the amount that would be produced in 
a given time. This difference may be brought out by the curves in the 
accompanying figure, Fig. 11, where the heavy line indicates the percent 
Fig. 11. 
of the total possible amount of SO3, which would be obtained at equilib- 
rium, provided that sufficient time were allowed. This time necessary 
for reaching equilibrium grows enormously as the temperature is lowered, 
with the result that it is practically impossible to reach equilibrium 
except at high temperatures, where the amount of SO2 converted into 
SO3 is so Small that it would not be profitable to use the reaction. The 
dotted lines indicate the actual approach to equilibrium in a given time, 
showing how, by allowing more time, more SO3 is obtained. At the 
higher temperatures, such as tT and t2 in the figure, where the reaction 
occurs with measurable speed, no amount of waiting will give a satisfac- 
tory yield of SO3. In order to work at temperatures at which the 
equilibrium is favorable a new factor must be invoked to increase 'the 
velocity. This is attained in practice by the use of catalysts, Fe2O3, and 
finely divided platinum. This allows the desired velocity to be attained 
at a much lower temperature, t3 in the figure, where the equilibrium is 
favorable to a good yield of SO3. The contact of the gases with this 
solid catalyst gives the name "contact process." 
The speed of the reaction, and also the desired equilibrium might be 
influenced in the right direction by having the reaction take place under 
high pressure. This involves difficulties in the way of more complicated 
machinery that it is preferred to avoid in practice. Since the SO2 must 
be paid for, while the O2 is gotten free from the air, it is more important 
to use up the SO2 completely than the O2. This is done by taking the 
119 air in considerable excess of the equivalent amount required. We may 
summarize in tabular form our conclusions concerning the best conditions 
for realizing this reaction. 
To favor large amount 
To favor high speed 
either of formation 
Total pressure 
of SO3 at equilibrium. 
High 
or decomposition. 
High 
Temperature 
Low 
High 
Other factors 
Excess of O2 
Catalyst 
To illustrate the conditions to be used in practice, there may be 
added a table showing the yield of SO3 that has been calculated, using a 
gas mixture at a total pressure of 1 atmosphere, with the following partial 
pressures: SO2, 0.07 atm.; O2, 0.10 atm.; N2, 0.83 atm. 
Temperature 434° 550° 645° 
Maximum yield, %SO3 99 85 60 
Sulfuric acid is made by dissolving the sulfur trioxide in water. 
b. The Synthesis of Ammonia. If we compare the reaction for the 
synthesis of ammonia, 
N2 + 3H2 = 2NH2 + 24000 cal., 
with the above reaction for making SO3, we see that the same conditions 
are indicated for realizing this reaction as for the former. This reaction, 
however, is exceedingly slow, so that even with the best catalyst that 
has been discovered after extensive search, uranium, advantage must be 
taken of high pressure in aiding both the speed and the equilibrium at 
the temperatures at which the reaction will take place with measurable 
velocity. 
The effect of both temperature and pressure on the yield of NH3 
at equilibrium is indicated in the following figures: 
• 
Yield of NH 
3, in per cent. 
Temperature. 
At 1 atmosphere. 
At 100 atmospheres 
800° 
0.011 
1.1 
700° 
0.021 
2.1 
600° 
0.048 
4.5 
500° 
0.13 
10.8 
A further lowering of temperature below 500° to improve the yield 
would give too slow a reaction, so that the process is worked at about 
this temperature, and under pressures up to 200 atmospheres. As the 
N2 and H2 pass over the catalyst they unite partially to form NH3. The 
gases are then passed through a cooling coil in which the NH3 is condensed 
to liquid and removed. The N2 and H2 again pass over the heated catalyst, 
again forming the equilibrium amount of NH3 and so on. The removal 
of NH3 is compensated by pumping in fresh H2 and N2. 
This process is 'now on a working basis and helping to solve the 
problem of making the inert nitrogen of the atmosphere available for 
our needs for fertilizer and explosives. It is of interest to note that 
it was not the discovery of the so-called "practical man," but was the 
result of extended and difficult scientific research along the lines here 
indicated. The conditions employed are such as would not occur to one 
not well versed in the theory, and call for pressures such as were never 
before used in a similar process. 
c. The Synthesis of Nitric Oxide. We have already referred briefly 
to the reaction, 
N2 4- O2 = 2NO - 44000 cal. 
120 We see that the formation of NO would be aided by the following condi- 
tions, which we will tabulate as before. 
To favor high speed 
To favor large amount 
either of formation 
of NO at equilibrium. 
or decomposition. 
Total pressure 
No effect 
High 
Temperature 
High 
High 
cvs a consequence oi these conclusions, we see that INO should De 
formed more rapidly and also more completely at high temperatures. 
That this is true is seen from the following figures on the basis of ex- 
periments : 
% NO at equilibrium, starting with air 0.37 0.97 5.0 10 
Temperature 1540 1920 2930 3930 
Obviously NO is only stable, in a true sense, at extremely high tem- 
peratures, and both a favorable equilibrium and high speed in attaining 
it are aided by high temperatures such as those given by an electric arc. 
The problem is then to get it down to ordinary temperature without 
letting the equilibrium shift back to N2 and O2 as it cools. This is done 
by cooling it sufficiently rapidly as the gas mixture leaves the arc. There 
is also an electrical effect, due to the arc tending to give a greater yield 
than is indicated by the above conditions of ordinary equilibrium. 
At ordinary temperatures NO unites with more O2 readily to form 
N2O4, which will also dissolve in water forming a mixture of nitric and 
nitrous acids. The N2O4 may also be passed over slaked lime giving a 
mixture of calcium nitrate and nitrate, which is used as fertilizer to 
supply nitrogen to the soil. This process was first used successfully on 
a commercial scale in Norway, and is now responsible for the fixation of 
enormous quantities of nitrogen. 
d. The "Cracking" of Petroleum. The principles above set forth 
have been of great aid in "cracking" the heavier constituents of petroleum 
so as to increase the yield of the more valuable lighter portions, such 
as benzene, CGHG, toluene, CGH5CH3, and those like C5H12, C6H14, C7H16, 
etc., which are contained in gasolene. It has been found possible to con- 
trol the nature of the products to a remarkable degree by designing 
apparatus that would permit of the independent control of temperature 
and pressure. The same substance is capable of a number of different 
transformations which would be influenced in different ways by changing 
the pressure and temperature. The probable effect of these factors 
may be predicted by the aid of the complete equations showing both 
the volume changes and the heats of reaction. The subject is too com- 
plicated for further discussion here, and we will merely give for illus- 
tration several equations showing possible reactions for C12H2G, a promi- 
nent constituent of kerosene. 
2C12H26 = CH4 + C23H48 - 500 cal. 
C12H26 = C2H2 + 2C5H12 «- 47,000 cal. 
C12Ho6 = 6C2HO + 7H2 - 381,000 cal. 
3C2H2 = C6H6 + 137,700 cal. 
It is evident that the equilibria in these reactions would be influenced 
in very different ways by pressure and temperature. 
Exercises. 
1. If the volume of a solution is greater than that of the salt and 
water from which it is made, how will the solubility of the salt change 
when pressure is applied to the saturated solution? 
121 2. In which direction would each of the following equilibria be 
displaced by decreasing the total pressure: (explain). 
N2O4 = 2NO2 
2HC1 = H, + Cl2 
C (solid) + S2(vapor)= CS2 (vapor). 
3. Ammonium chloride dissolves with considerable absorption of 
heat; how is its solubility affected by the temperature? 
4. The solubility of sodium chloride increases very slightly with 
temperature; what can be concluded concerning its heat of solution? 
5. When AgCl is prepared by mixing its ions heat is evolved; is 
the precipitate more soluble in hot or in cold water? (Explain). 
6. Explain how temperature would affect the production of H,S 
according to the reaction: 
H2 + S = H2S + heat. 
7. The formation of benzene from acetylene according to the reac- 
tion 3C2H2 = C6Hg 131,000 cal. takes place slowly at a high temper- 
ature (benzene being a vapor) and is reversible. 
(a) What weight of benzene could be made by the complete con- 
version of 49.3 liters of acetylene, measured at 27°C and 1 atmosphere? 
(b) State and explain the experimental conditions you would adopt 
to secure the maximum yield of benzene from acetylene in a given time. 
8. Is the peculiar taste of stale bread due to its becoming dry? Why 
do crackers taste fresh as long as they are kept dry? How can the 
fresh taste be restored in either bread or crackers? Would bread keep 
fresh longer in a refrigerator than at room temperature? 
122 CHAPTER XV. 
OXIDATION AND REDUCTION. 
In the previous chapters our attention has been largely confined to 
reactions in which there are no changes in valence. In Chapter V, 
however, brief mention was made of reactions in which such changes 
do occur, and it was there pointed out that oxidation consists essentially 
of an increase of valence and reduction of a decrease of valence. We 
are now prepared to give fuller attention to such reactions, considering, 
first, the writing of equations representing them, and, second, the oxidizing 
and reducing powers of the substances involved. 
Whenever an element is oxidized some element must be reduced. The 
substance containing the element which is oxidized is called the reducing 
agent, since it is responsible for the reduction of the other element, while, 
conversely, the substance which causes an increase in valence is the 
oxidising agent, and contains an element which is reduced. The following 
table illustrates these terms: 
Cu + 2H2SO4 =CuSO4 4- SO2 + 2H2O 
Valences 0 6 2,6 4 
Element oxidized, Cu. Oxidizing agent, H2SO4. 
Element reduced, S. Reducing agent, Cu. 
H2S + I2 = S + 2H+ + 21- 
Valences 1,-2 0 0 -1 
Element oxidized, S. Oxidizing agent, I2. 
Element reduced, I. Reducing agent, H2S. 
2Fe+++ + H2SO3 + H2O = 2Fe++ + SO4" + 4H+ 
Valences 3 4 2 6 
Element oxidized, S. Oxidizing agent, Fe+++. 
Element reduced, Fe. Reducing agent, H2SO3. 
MnO4- + 5Fe++ + 8H+ =Mn++ + 5Fe+++ + 4H2O 
Valences 7 2 2 3 
Element oxidized, Fe. Oxidizing agent, MnO4_. 
Element reduced, Mn. Reducing agent, Fe++. 
Writing Equations for Reactions Involving Oxidation and Reduction. 
(a). The Substance Produced. Before an equation can be written for 
any reaction it is necessary to know what substances are produced. This 
may be determined by experiment or predicted, in most cases, from a 
knowledge of the reacting substances. For example, when hot concen- 
trated sulfuric acid reacts with copper it is possible to detect the sulfur 
dioxide evolved by its odor, and the copper sulfate produced by the blue 
color it gives on the addition of water. When sulfurous acid acts upon 
ferric ion, the ferrous ion produced shows its presence by its faint green 
color, or by other characteristic tests, while the formation of sulfate ion 
can be proved by adding barium ion. One sufficiently familiar with the 
chemistry of iron and of sulfur would not need these tests, for he knows 
that when ferric ion decreases in valence in acid solution it would usually 
123 become ferrous ion, and that the corresponding increase of valence on 
the part of sulfurous acid would give sulfate ion. It is therefore important 
to know the valences the elements are capable of assuming ,and the com- 
pounds characteristic of each valence. 
For example, many of the metals, including the alkali and alkaline 
earth metals, silver, zinc, cadmium and aluminum exhibit only one valence 
in addition to that of the free metal which is zero. Therefore a substance 
like aluminum ion, Al+++, cannot be oxidized further, nor can it be reduced 
to metal in water solution. Hence we wouuld not expect it to take part 
in any oxidation or reduction process in aqueous solution. Chlorine shows 
the following valences, - 1, 0, 4-1, 4~ 3, 4" 5, +7, which correspond 
to the compounds tabulated on page 57. Consequently, we would never 
find the chlorine in chloride ion, Cl", acting as an oxidizing agent, because 
it is already as completely reduced as it is capable of being. Likewise, 
if chlorine, Cl2, acts as an oxidizing agent it must be reduced, and the 
only possible product would be chloride ion or one of its compounds. 
Again, knowing that chlorine reacts with hydroxide ion as follows: 
Cl2 4- 2OH- = Cl" 4- CIO- 4- H2O, 
we know that when CIO- acts as an oxidizing agent in alkaline solution 
the chlorine could not be reduced to the next lower valence, zero, corre- 
sponding to Cl2, but would be reduced all the way to chloride ion. 
The following table represents the substances characteristic of 
chromium in the valences of three and six respectively, and in acid, neutral 
or alkaline solution: 
Acid 
Neutral 
Alkaline 
Valence 3 
Cr++ 
Cr(OH)3 
CrO2" 
Valence 6 
Cr2O7-_ 
CrO4" 
CrO4" 
In accordance with the knowledge thus summarized, we would expect 
Cr+++ to be produced when a reducing agent acts upon Cr2O7~". Similar 
information should be remembered for the other common elements giving 
compounds in which they show more than one valence. 
(b) The Assignment of Valence Numbers to the Elements Oxidized 
and Reduced. Having decided upon the substances produced by the 
reaction, the first step is to write their formulas, together with those of 
the substances reacting, upon the appropriate side of the equation, and 
to assign a valence number to the elements oxidized and reduced, as 
explained in Chapter V. Let us illustrate the process by balancing the 
equations given in the preceding table. Omitting the valence numbers 
for the elements obviously not oxidized or reduced during the reaction 
we would have the following incomplete equations: 
Cu + HoS04 = SO2 + CuSO4 
0 6 4 2 
H2S + I2 = i- + S 
-20 -10 
Fe+++ + H2SO3 = Fe++ 4- SO4~ 
3 4 2 6 
MnO4~ + Fe++ = Mn++ + Fe+++ 
7 2 2 3 
(c) The Determination of the Change in Valence per Atom of Each 
Element Oxidized and Reduced is the next step. This may be conveniently 
indicated by marking it upon a line connecting the elements in question, 
as follows: 
124 Cu 4- H2SO4 = SO2 4- CuSO4 
0 '6 4 " 2 
I I -2 I I 
I L -* I 
I 4-2 | 
i i 
H2S 4- I2 = I" 4- S 
- 2 0 -1 0 
11-11 I 
I L J I 
I 4-2 | 
i i 
Fe+++ 4~ H2SO3 = Fe++ 4~ SO4-- 
3 4 2 6 
I I 14-2 1 
| - 1 L + J 
I 1 
MnO4~ + Fe++ = Mn++ 4- Fe+++ 
7 2 2 3 
I I 14-11 
I -5 l- 4- J 
i . i 
(d) The next step is the adjustment of the molecules of these sub- 
stances so that the total increase in valence of the atoms oxidized equals 
the total decrease in valence of the atoms reduced. In the first equation 
it is seen that the production of one molecule of SO2 will correspond to 
the oxidation of one atom of copper, hence no further change is necessary 
under this heading. 
In the second, obviously, one atom of sulfur can oxidize two atoms 
of iodine, so that this equation should be altered to read 
H2S 4- I2 = 21- + S 
- 2 6 2(-1) 0 
I I 2(-l) | | 
I 1 1 I 
I +2 | 
i i 
In the third two atoms of iron can be reduced when one atom of 
sulfur is oxidized, so that we should write: 
2Fe+++ + H2SO3 = 2Fe++ +SO4" 
2(3) 4 2(2) 6 
I I I +2 I 
I 2(-l) L + J 
i i 
In the fourth every atom of manganese, in being reduced, can oxidize 
five atoms of iron, hence we write 
MnO4~ + 5Fe++ = Mn++ 4~ 5Fe+++ 
7 5(2) 2 5(3) 
I I I +5 | 
| -5 l- + J 
i 1 
125 e. The final balancing of the equation usually requires the addition 
to one side or the other of appropriate substances. Thus, in the first 
example it is evident that another molecule of sulfuric acid is needed on 
the left to furnish the sulfate radical for the copper sulfate, so that we 
will have 
Cu + 2H2SO4 = SO2 -f-CuSO4. 
Now we have four atoms of hydrogen and two extra atoms of oxygen 
on the left which must be balanced by adding 2H,O on the right, so that 
the completed equation will be 
Cu + 2H2SO4 = SO2+ CuSO4 + 2H2O. 
It is desirable to test the accuracy with which the above steps have 
been made by seeing if the total number of atoms of each element on 
one side of the equation equals that on the other. 
In the second equation we see that we have, so far, two atoms of 
hydrogen on the left and two negative charges on the right with no 
counterpart on the other side. By adding on the right two hydrogen ions 
the equation will obviously be balanced, reading finally as follows: 
H2S + I2 = 2I-+S + 2H+. 
In the final testing of an equation like this the total charges on the 
two sides, cancelling those of opposite sign, must be the same. 
In the third equation we have one atom of oxygen in excess on the 
right, and two atoms of hydrogen and four positive charges in excess 
on the left. In such cases probably it is simplest to balance the charges 
first, by adding to one side or the other the proper amount of H+ or OH". 
In the above instance the charges could obviously be balanced either by 
adding 4H+ on the right or 4OH~ on the left. Either would enable us 
to balance the equation, but from a chemical standpoint it would be 
incorrect to add 4OH", for the solution is not alkaline, Fe+++ and H2SO3 
being unable to exist in alkaline solution. Accordingly we aid 4H+ on the 
right, whereby the equation becomes 
2Fe+++ + H2SO3 = 2Fe++ + SO4" " + 4H+. 
All that is necessary now, finally, to balance the equation is to add 
H2O on the left, getting 
2Fe+++ -f- H2SO3 + H2O = 2Fe++ + SO4" + 4H+. 
In the fourth case, evidently an acid solution likewise, it is necessary, 
in order to balance the charges, to add 8H+ on the left, getting 
MnO4" + 5Fe++ + 8H+ = Mn++ + 5Fe+++. 
This requires only the addition of 4H2O on the right to become complete. 
MnO4" 4- 5Fe++ 4- 8H+ = Mn++ 4- 5Fe+++ 4- 4H2O. 
As an example of balancing an equation for a reaction occurring in 
alkaline solution we may take the following, where the successive steps 
just explained are indicated successively. 
(a) CrO2- + CIO- = Cl" + CrO4" 
(b) 3 1-16 
(c) | +3 | -2 | | 
I L J I 
i 1 
(d) 2CrO2 + 3C1O' = 3C1- + 2CrO4"" 
2(3) 3(1) 3(-l) 2(6) 
I I 3(-2) | | 
I 2(4-3) l J | 
i i 
126 (e) 2CrO2" + 3C1O" + 2OH" = 3CF + 2CrO4" 
2CrO2- + 3C1O- + 2OH- = 2OH~ + 2CrO4~ + H2O. 
The Relative Oxidizing and Reducing Powers of Various Substances. 
a. The Metals and their Ions. It was pointed out in Chapter V that 
the more positive elements can replace the less positive ones in their com- 
pounds. This is illustrated by the following reactions, all of which take 
place quite readily : ; 
2Na 4- Zn++= Zn 4~ 2Na+ 
• Zn -j- Fe++- Fe 4- Zn++ 
Fe 4- Pb++= Pb 4- Fe++ 
Pb 4- Cu++= Cu 4- Pb++ 
Cu 4- 2Hg+= 2Hg 4- Cu++ 
Hg + Ag+= Ag 4- Hg+ 
In each of these cases the baser metal is oxidized by the ion of the 
nobler one, and it is exident that zinc will reduce not only ferrous ion, 
but lead, copper and silver ions as well. Likewise, silver ion can act as 
an oxidizing agent, not only to copper, but to all of the other metals 
represented. In all cases the metal acts as the reducing agent, their 
reducing power decreasing as we descend the list, while -the ions act as 
oxidizing agents, their oxidizing power increasing as we descend the list. 
This behavior mav be summarized in the following1- table: 
a 
.2 
' Weak 
Oxidizing agents 
Strong 
<T> 
, Na+ 
Zn++ 
Fe+ 
+ Pb++ Cu++ 
Hg+ 
Ag+ 
CL 
c 
O 
Na 
1 Strong 
Zn 
Fe 
Pb Cu 
Reducing agents 
Hg 
Ag 
Weak 
o' 
This table has been arranged so that any ion will be reduced if 
brought in contact with a metal to the left in the table. In order, however, 
for this arrangement to hold, the ions must all be in approximately equal 
concentration. If this is not the case the positions of some may be 
reversed. Thus Hg+ and Ag+ have about the same oxidizing power, 
hence the reaction 
Hg+ + Ag = Ag+ + Hg 
is easily reversible by changing the relative concentrations. The principles 
of equilibrium set forth in Chapter XII apply here also, so that a high 
concentration of Hg+ and a low concentration of Ag+ favor the reaction 
as read from left to right, whereas the reverse reaction is favored by 
reversing the relative concentration. With most pairs of metals the 
corresponding reversal requires greater differences in concentration. 
Many interesting applications of this oxidation-reduction series are 
possible. For example, it js evident that if a piece of zinc is put into 
a solution of ferrous ion there is a tendency for the latter to be reduced. 
Accordingly, if zinc and iron are in contact in an oxidizing solution the 
former will be dissolved before the latter. This is the reason for the 
protective action of the zinc coating on "galvanized" iron. 
When a piece of zinc is dropped into a solution of one of the ions 
to the right of it in the table, a reaction takes place such as the following: 
Zn Cu++ = Cu + Zn++. 
Now this may be regarded as a transfer of positive electricity from the 
copper ion to the zinc, or, better, as the transfer of negative electrons 
from the zinc to the copper ion. If this transfer of electricity can be 
made to take place through a wire we will have an electric current and 
the reaction can be made to do work. To do this the zinc must not be 
127 put directly into the solution of copper ion, but electrodes of the two 
metals must dip into solutions of their own ions, the solutions being 
separated by a porous partition or by the aid of their difference in 
density, as in the case in the well known "gravity cell." At the zinc 
electrode the following process tends to take place: 
Zn = Zn++ 4~ 2(-) 
while at the copper electrode the reaction is 
Cu++ + 2(-)=Cu. 
Each of these processes can continue only if the stream of negative 
electrons, or atoms of electricity, can flow from the zinc to the copper, 
while the positive and negative ions in the solution migrate towards their 
respective poles. For every mol of zinc that dissolves one mol of copper 
will be deposited and two faradays of electricity (see Chap. X) will 
travel through the wire. The electromotive force between the electrodes 
depends upon the difference between the tendencies of the metals to 
become ions, and also upon the relative concentration of the ions in the 
solution. If the concentration of zinc ion is diminished, er that of copper 
ion is increased, the reaction 
Zn Cu++ = Cu + Zn++ 
will have a greater tendency to take place and will give a larger electro- 
motive force when occurring in an electric battery. If silver ion and 
silver are substituted for copper ion and copper we would have a stronger 
oxidizing agent, a greater tendency to take up the electrons given off by 
the zinc, and a battery of higher electromotive force. 
The electromotive force of such a battery is thus a measure of the 
tendency of the reaction to take place, or the chemical affinity of the 
reaction, a term much used in the development of chemistry, but also 
much abused, for it has often been used very loosely. It is only in 
comparatively recent years that it has been given an exact definition. On 
this account it is desirable to substitute the term "free energy" for affinity. 
The electromotive force of the copper-zinc cell when normal solutions of 
copper and zinc sulfates are used is about 1.09 volts. The free energy 
of the reaction at these concentrations is the work it can do, which would 
be 0.09x2x96540 volt-coulombs or joules per mol of reacting substances. 
The reaction between zinc and silver ion may be utilized in cleaning 
the tarnish from silverware. The tarnish consists of silver sulfide, Ag2S, 
and though this gives but a small concentration of silver ion, the free 
energy of the reaction is still sufficient for it to take place quite readily. 
If, therefore, a piece of zinc (aluminum will also serve) is placed in 
electrical contact with a piece of tarnished silver, both dipping into 
water containing a little salt to render it conducting, and hot to increase 
the speed of reaction, the following reaction takes place: 
AgS 4- Zn = Ag 4- Zn++ 4- S". 
The silver is thus restored instead of being rubbed off as in ordinary 
polishing. 
If a solution containing the ions of two metals is electrolyzed it is 
possible to reduce at the kathode, first the ion of the baser metal, and 
afterwards that of the nobler metal, by proper regulation of the electro- 
motive force. It is possible thus to make an electrolytic separation of 
different metals which is to be preferred, in many cases, to the ordinary 
methods of separation based upon the precipitation of insoluble com- 
pounds. 
b. The Solubility of Metals in Water, Acids and Alkalis. When a 
128 metal dissolves in water, in alkali, or in dilute acids the substance reduced 
is usually hydrogen ion. Accordingly it is desirable to introduce into the 
table previously given H+ and H2. In attempting to do this we find an 
apparent discrepancy in the reactions of the metals with hydrogen ion, 
for iron will dissolve in dilute hydrogen ion while zinc, if very pure, 
and especially if wet with mercury, will not. We find, however, that 
this is due to the difficulty which hydrogen has in depositing on a surface 
of pure zinc or mercury, for if a piece of platinum is dipped into the 
acid and brought in contact with the zinc we find that the latter dissolves 
quite readily, the hydrogen depositing on the platinum, while the electrons 
left behind when the zinc becomes ion travel through the metals to the 
platinum where they neutralize the hydrogen ions, changing them into 
hydrogen gas. If we test the solubility of the various metals in dilute 
acids in this way we find that all those to the left of copper in the table 
will dissolve, although lead and zinc will do so only with great difficulty, 
when pure and in the absence of some more favorable surface for the 
hydrogen to deposit upon. The same difficulty is encountered in reducing 
hydrogen ion at certain cathodes, a higher electromotive force being 
required to deposit hydrogen at a certain rate on some surfaces than on 
others. A successively higher electromotive force is required with 
kathodes of the following substances: Pt (rough, Pt (smooth), Ag, Cu, 
Sn, Pb, Zn, Hg. Hydrogen thus takes the place given in the more 
extended table on page 131, provided that a platinum surface is in electrical 
contact with the metal dissolving. We see, then, that copper, silver, 
mercury and the other noble metals will not dissolve in acids like dilute 
hydrochloric and sulfuric under any circumstances, while lead and the 
baser metals will dissolve in dilute hydrogen ion, provided that platinum 
is in contact with the metal, but that lead and zinc, or other base metals 
if amalgamated with mercury, may dissolve very slowly, or practically 
not at all. When impure zinc is used, the particles of foreign substances, 
like iron and carbon, furnish nuclei upon which the hydrogen can deposit 
with more ease, hence such zinc will dissolve at a reasonable rate. 
As we approach the left side of the table, where the metals become 
baser, and better reducing agents, the hydrogen ion does not need to be 
so concentrated in order to act upon the metal, and even the extremely 
dilute hydrogen ion of water is able to oxidize and dissolve metals like 
calcium and sodium. 
Aluminum dissolves easily in dilute hydrogen ion, the reaction being 
Al + 3H+ = AW+ + 3/2H2. 
If hydroxide ion is used instead of hydrogen ion, although the decrease 
in the concentration of the latter would tend, by itself, to stop the solution 
of the aluminum, at the same time the concentration of aluminum ion is 
decreased very greatly on account of the formation of the aluminate ion, 
a fact which aids the reaction, and compensates for the reduction in the 
hydrogen ion concentration. Hence aluminum will dissolve in alkali 
as well as in acid. The same is true of zinc. 
For the metals to the right of hydrogen in the table a stronger 
oxidizing agent than hydrogen ion is necessary. This is found, for some, 
in nitric acid, whose oxidizing action is due to the reduction of nitrogen 
rather than hydrogen. Copper, silver and mercury can accordingly be 
dissolved in nitric acid, a typical reaction being the following: 
3Cu + 2NO3- + 8H+ = Cu++ + 2NO + 4H2O. 
Sulfuric acid, if hot and concentrated, can oxidize copper, silver and 
mercury, its sulfur being reduced to SO2. 
129 When we come to the noblest metals, like gold and platinum, even the 
oxidizing power of nitric acid becomes insufficient. Instead, however, of 
seeking a stronger oxidizing agent, another principle is invoked. The 
following reaction will not take place with nitric acide alone: 
2Pt 4- 4NO3- + 16H+ = 2Pt++++ + 4NO + 8H2O. 
Platinic ion, however, is prone to form a complex acid with hydrochloric 
acid, H2PtClfi, chlor-platinic acid (see page 101), and hence, if a mixture 
of hydrochloric and nitric acids, known as "aqua regia," is used instead 
of the nitric acid alone, the platinic ion, one of the products of the 
oxidation, is used up sufficiently to enable the reaction to proceed. This 
action of aqua regia is not due, as is often erroneously stated, to the 
chlorine liberated by the aqua regia (4H+4-NO3" 4~ ^Cl" = Cl2 4~ NOCI 
4~2H2O) being a stronger oxidizing agent than the nitric acid, for if this 
were the case chloride ion would not be oxidized by nitric acid. Chlorine 
is, however, a more rapid oxidizing agent, and the effectiveness of the 
aqua regia is doubtless due to this fact as well as to the formation of the 
complex chlor-platinic acid. The action of aqua regia on gold is quite 
analogous to that on platinum, the complex chlor-auric acid, HAuC14 
being formed. 
Cyanide ion, CN", forms such a firm complex with both aurous and 
auric ions that when it is present the oxygen of the air is sufficient to 
oxidize gold, the reaction being 
2Au + %O2 4- H2O 4- 4CN- = 2Au(CN)2- 4- 2OH-. 
This reaction is used in the "cyanide process" for the extraction of 
gold from its ores. 
c. General Table of Oxidizing and Reducing Agents. We may consider 
the oxidizing and reducing powers involved in changes other than those 
between the metals and their ions just considered. We find, for example, 
that chlorine will liberate bromine from bromide ion, and that bromine 
will liberate iodine from iodide ion. Accordingly we may conclude- that 
chlorine will liberate iodine from iodine ion. This information may be 
tabulated in the same way as was used previously for the metals, as 
follows: 
W eak 
Oxidizing Agents 
Strong 
I2 
Br2 
CL 
I- 
Br 
ci- 
Strong 
Reducing Agents 
Weak 
By the aid of appropriate experiments we may find the place of other 
common substances in such a table, and also combine it with the table 
previously given. The following table gives an illustration of such 
arrangement. 
130 Table of Oxidizing and Reducing Powers of Substances. 
Oxidation 
Zn 
Fe 
Sn 
Pb 
H2 on Pt 
Sn++ 
H2S only 
Zn++ 
Fe++ 
Sn++ 
Pb++ 
H+ 
Sn++++ 
S 
Strong 
Weak 
Reducing Agents 
Oxidizing Agents 
SO2 
only 
SO4- 
I- 
Fe++ 
Ag 
Br- 
I2 
SO. 
I2 
Fe+++ 
Ag+ 
Br. 
IO3" 
H2SO4, hot 
concentrated 
HNO3 
NO 
only 
Cr+++ 
Cr2O7 -f- concen- 
trated H+ 
MnO4" + H+ 
PbO2 + concentrated 
hvhnoa 
Strong 
Weak 
Mn++ 
Pb 
Reduction 
It should be borne in mind that such an arrangement is not altogether 
strict, and that it is affected by the concentrations of the reacting sub- 
stances. According to the table, for example, one might infer that the 
following reaction would occur: 
Ag+-F I-= Ag-f-%I2. 
This would undoubtedly take place if Ag+ and I" could act upon each 
other at ordinary concentrations without precipitating insoluble Agl. 
The same applies to the action of H2S on Ag+. 
The hydrogen ion concentration affects the equilibrium in many cases. 
For example, the reaction 
SO. 4- I2 4- H2O = SO4" + 21- 4- 4H+ 
proceeds readily in dilute solution. It will be seen, however, that by 
diminishing the amount of water present and increasing the relative 
amount of acid the equilibrium is shifted in favor of those substances 
on the left, and, as a matter of fact, concentrated sulfuric acid readily 
liberates iodine from solid iodides. Similarly, the following reaction 
accords with the positions in the table of the substances involved: 
Cr2O7" 4- 6C1- 4- 14H+ = 2CF++ 4- 3C12 4- 7H2O, 
and in concentrated acid the reaction proceeds in this way. The large 
amount of H+ used up, however, indicated that the equilibrium would be 
enormously affected by changing its concentration. In fact the reaction 
proceeds in the other direction in alkaline solution, the chromium being 
oxidized and the chlorine reduced. The transition from one to the 
other may be made by adding 16OFF to both sides of the equation. Of 
131 these, 14OH" will react with the 14H+, 2OH" with Cr2O7 giving 
2CrO4" + H2O, while on the right we will have 2Cr+++ 4- 8OH = 2CrO," 
+ 4H2O, and 3C12 + 6OH~ = 6C1O~ + 3H2O, with 2OH~ in excess. 
Cancelling out H2O and Cl", and transposing the two sides, we have the 
equation previously considered on page -. 
2CrO2" + 3C1O" + 2OH- = 3Cb + 2CrO4". 
Although this is apparently very different from the equation given above, 
it involves essentially the same oxidation and reduction steps. On account 
of the great effect on many of these reactions of varying the concentration 
of the ions of water, it is well to include in a separate table the reagents 
which act in alkaline solution. The following brief table may serve for 
illustration: 
Weak 
Oxidizing Agents 
Strong 
HSnOo" 
CrO4- 
C1O~ 
HSnO3- 
CrO2' 
C1-. 
Strong 
Reducing Agents 
Weak 
Oxidizing Power and Speed of Oxidation. There are some sub- 
stances,, like hydrogen peroxide, H2O2> which would have great oxidizing 
power if they could react rapidly enough. It is important to distinguish 
between an oxidizing agent that is "strong" because it is rapid, and one 
which is strong but slow. Thus iodine will often react more rapidly than 
chlorine, and hence produce more oxidation in a given time. Its true 
oxidizing power, however, is much less, as shown by the fact that it 
will not oxidize ferrous ion completely, whereas chlorine will. The oxygen 
acids, like nitric and sulfuric acids, are not rapidly oxidizing agents. 
This seems to be connected with the fact that they are highly ionized. 
Nitrous acid, which is not a stronger oxidizing agent in the true sense 
of the term, is nevertheless a much more rapid oxidizing agent, and is, 
at the same time, much less dissociated. As we compare the oxygen 
acids of chlorine we find that while their reactivity as oxidizing agents 
decreases in going from HC1O to HC1O3, their strength as acids increases 
greatly. In a reaction such as the following, 
C1O3- + s + H2O = Cl- + SO4" + 2H+, 
although hydrogen ion appears, whence we might expect it to be favored 
by decreasing the concentration of hydrogen ion, as a matter of fact it 
will not take place unless the solution is strongly acid. We may infer 
from this that C1O8" is very unreactive, and a large amount of H+ must 
be present in order to give a certain amount of the undissociated strong 
acid, HC1O3, which is reactive. With HC1O, however, reduction to Cl" 
will take place rapidly without adding H+, because HC1O is such a weak 
acid that most of it is undissociated even in neutral solution. The 
difference between nitrous and nitric acids may be explained similarly; the 
latter, being a strong acid, must be rather concentrated in order to give 
enough undissociated molecules of HNO3 to react rapidly. In general, 
when substances react slowly their true place in the table may not be 
apparent from their ordinary reactions, and the place assigned may not 
correspond to a reversible reaction. 
Substances that can act both as Oxidizing and Reducing Agents. 
Certain substances appear at two places in the table, corresponding to 
their power to act either as oxidizing or reducing agents. The way in 
which they will react depends, of course, upon the nature of the substances 
with which they are mixed. The ferrous ion can be reduced to iron by 
the action of a very powerful reducing agent like zinc, or it can be 
132 oxidized to ferric ion by the action of a strong oxidizing agent like 
chlorine. Iodine may be reduced to iodide or oxidized to iodate according, 
respectively, as it reacts with a reducing agent stronger than iodine or 
an oxidizing agent stronger than iodate ion. Hydrogen peroxide, not 
shown in the table because of the irreversibility and slowness of its reac- 
tions, may act as an oxidizing agent with some substances themselves 
capable of oxidation, the oxygen in the hydrogen peroxide being reduced, 
as illustrated by the reaction 
2H+ + H2O2 4- 2Fe++ = 2Fe+++ + 2H2O. . 
On the other hand, in the presence of a sufficiently strong oxidizing agent, 
the oxygen in the peroxide may be oxidized to free oxygen, as illustrated 
by the reaction 
2MnO4~ + 5H2O2 + 6H+ = 2Mn++ + 5O2 + 8H2O. 
The Solution of Insoluble Sulfides. In Chapter XIII it was shown 
how certain metallic sulfides could be separated by controlling the hydrogen 
ion concentration, and it was pointed out that there is a group of sulfides 
so insoluble in water that even moderately concentrated hydrochloric or 
sulfuric acids will not sufficiently remove the sulfide ion, S"", to make the 
metallic ion concentration large. To dissolve such sulfides it is necessary 
to remove the sulfide ion more completely. This is possible by the use 
of a strong oxidizing agent, usually nitric acid, which changes the S"" com- 
pletely into S or even SO4-". Moderately concentrated hot nitric acid, 
accordingly, will dissolve CuS, Bi2S3, PbS and Ag2S. The solubility of 
mercuric sulfide is so extremely low, however, that nitric acid is prac- 
tically without effect. To dissolve it aqua regia is necessary, the effect- 
iveness of which depends upon the greater speed of reaction of chlorine 
as compared with hydrochloric acid, and also upon the presence of the 
chloride ion, which unites with the mercuric ion to form undissociated 
mercuric chloride. The aqua regia thus attacks both constituents of the 
mercuric sulfide instead of only one. 
Storage Batteries. Other reactions besides the simple change between 
metals and their ions can be utilized in generating an electric current. A 
very important one is that occurring in the lead storage battery. The 
reaction utilized is the following, 
Pb + PbO2 + 4H+ + 2SO4; - = 2PbSO4 + 2H2O. 
It is carried out so that the oxidation of the lead to lead sulfate 
Pb + SO4" = PbSO4 + 2(-) 
tages place at one electrode and the reduction of lead doixide to lead 
sulfate, 
PbO2 + 4H+ + SO4- - + 2<-)= PbSO4 + 2H2O, 
takes place at the other electrode, the electrons traveling through the 
metallic part of the circuit from one electrode to the other. To charge 
the cell the reaction is reversed by sending a current through the cell 
in the opposite direction. From the equations it may be seen that the 
more concentrated is the acid the more is the equilibrium shifted in favor 
of PbSO4, hence the larger the free energy of the reaction and electro- 
motive force of the cell. There are other considerations, however, which 
we have not room to consider here, which argue against too concentrated 
an acid. 
The nickel-iron storage battery utilizes the following reaction, taking 
place in alkaline solution: 
2Ni(OH)3 + Fe = 2Ni(OH)2 4- Fe(OH)2. 
133 The electrode reactions are 
2Ni(OH)3 + 2(-)= 2Ni(OH)„+ 2OH", and 
Fe + 2OH- = Fe(OH)2 + 2(-). 
Exercises. 
1. In the following reactions state which elements are oxidized and 
which are reduced: 
2Ag+ + Zn = 2Ag + Zn++ 
Ag+ + Cl- = AgCl 
Zn + OH- + H,0 = HZnO." + H, 
2Hg+ + H2S = HgS + Hg + 2H+. 
2HgCl2 + Sn++ = 2HgCl + Sn^ + 2CF. 
2. Write equations for the following reactions: zinc plus chlorine 
to give zinc ion plus chloride ion; stannous ion plus bromine to give 
stannic ion plus bromide ion; ferric ion plus stannous ion to give ferrous 
ion plus stannic ion; hydrogen sulfide (H2S) plus iodine to give sulfur 
plus hydrogen ion plus iodide ion; cupric ion plus iodide ion to give 
cuprous iodide (ppt.) plus iodine. 
3. Write equations for the reactions between: (a) cupric oxide 
and hydrogen; (b) cupric oxide and hydrogen ion; (c) cuprous oxide 
and oxygen; (d) magnesium and hydrogen ion; (e) cupric ion and 
zinc to give copper; (f) copper and cupric ion; (g) silver ion and zinc. 
State in each case what elements are oxidized and which are reduced. 
4. Complete the following incomplete equations: 
(a) H+ + NO3- + Ag = Ag+ + NO 
(b) H+ -j- NO3~ -j- Sn = HoSn03 + N2O3 
(c) H+ + NO3' + CuS = Cu++ + SO4-' + NO 
(d) H+ + NO3- + Zn = Zn++ + NH4+ 
(e) Bi(OH)3 4- HSnO2-=Bi+ 
(f) HAuC14 + Fe++ = Au + 
(g) P4 + OH- = PH3 + H2PO2- 
(h) Aso03 + H+ 4-NO3- = H3AsO4 + N2O3 
(i) CIO- + AsO2" = As63- + 
5. What is meant by "noble" metals and "base" metals? 
6. From what metals given in the table, oage -, construct the 
battery that would give the greatest electromotive force. 
7. Explain how contact with zinc protects iron from corrosion. 
8. How can you tell when a metal will dissolve in any of the 
following reagents: water, hydrochloric acid, nitric acid, sulfuric acid, 
or sodium hydroxide? 
9. How is the oxidizing power of H+ affected by its concentration? 
10. Explain the cleaning of silver by zinc or by aluminum. 
11. How would you prepare each of the following: 
(a) Br- from Br2 
. (b) HO gas from Cl2 
(c) I2 from Nai 
(d) PbCrO4 from Pb and CrCl3 
(e) Fe from FeCl3 
(f) Fe(OH)2 from Fe 
(g) Fe2O3 from FeSO4 
(h) Mn(OH)2 from MnO4~ 
(i) HgCl from HgCL 
(j) Hg from HgCl2 
12. How many grams of hydrogen peroxide are necessary to reduce 
134 10 grams of MnO (OH)2 to Mn++ in acid solution? Ans. 3.24 gm. 
13. A solution of KMnO4 is standardized by titrating a ferrous 
solution made by dissolving 1.12 grams of pure iron. The volume of 
KMnO4 solution required is 42.0 cc. The same permanganate solution 
is then used to determine the strength of a solution of hydrogen peroxide, 
50 grams of the latter requiring 34.0 cc. of the former. What per cent of 
hydrogen peroxide is in the solution? (In this reaction Mn++ and O2 are 
produced). Ans. 0.55%. 
14. Manganese in the form of manganous ion can be titrated with 
a solution of potassium permanganate, MnO2 being formed. Write the 
equation. 
15. How many grams of K2Cr2O7, in solution, would it take to 
oxidize the same amount of Fe++ as would be oxidized by 1.00 gram of 
KMnO4, in acid solution ? 
16. Starting with metallic tin, how would you prepare in succession: 
SnCl2; SnCl4; SnO2; SnO; SnS; SnS2; SnCl4? 
17. How could you separate the metallic elements in a precipitate 
consisting of PbS, ZnS and SnS? 
18. Why cannot HI be prepared by the action of concentrated 
H2SO4 on Nai? 
135 CHAPTER XVI. 
PERIODIC SYSTEM. 
It became evident to chemists a long time ago that there are certain 
elements which may be grouped together by reason of their similarity. 
Among such groups may be mentioned lithium, sodium and potassium, 
commonly called the alkali metals, calcium, strontium and barium, the 
alkaline earth metals, fluorine, chlorine, bromine and iodine, called 
the halogens. In 1866 Newlands read a paper before the London Chemical 
Society in which he showed that if the then known elements were arranged 
in the order of increasing atomic weight, as follows: 
H Li Be B C N O 
1 7 9 11 12 14 16 
; F Na Mg Al Si P S 
i 19 23 24 277 28 31 32 
1 Cl K Ca 
i 35.5 39 40 
there is a similarity between every eighth element in the series, so that 
if it is divided into groups of eight and these placed under each other 
in successive rows, as follows, 
H 
F 
Cl 
Li 
Na 
K 
Be 
Me 
Ca 
B 
Al 
etc. 
C 
Si 
N 
P 
O 
S 
then the elements standing in the same vertical column are those which 
correspond to each other in chemical nature. It is evident from the 
table that this is the case. Strange as it may seem this discovery was 
received with some ridicule, and its value not appreciated till later. In 
1869 this same idea was more fully elaborated independently by Mendeleeff 
and Lothar Meyer. Each of these chemists tabulated the elements substan- 
tially as is here shown in Table I. 
The fundamental idea expressed by such an arrangement is that of 
the periodic recurrence of properties as the atomic weights increase. This 
periodicity is shown graphically in Fig. 12, where atomic volumes (the 
136 number of cubic centimeters occupied by a gram-atom of the element) 
are plotted, not against atomic weights, however, but against the atomic 
numbers. The atomic numbers are whole numbers houghly half the 
atomic weights, which have been determined by X-ray spectra, and 
which are more fundamental than atomic weights. It will be noticed 
that at the successive peaks are found similar elements, those the alkali 
metals. The next element to the left of each peak is a membr of the 
group of rare inert gases, the next element to the right of each peak is 
an alkaline earth metal, and so on. The curve is thus made up of a 
series of roughly similar sections each of which corresponds to a row 
of elements in the table. If the absolute melting points are plotted against 
the atomic numbers we obtain the curve shown in Fig. 2, in which we 
see a similar periodicity. Here the peaks are occupied by the elements 
of the group of inert gases. This curve, likewise, could be cut up into 
a number of similar sections. 
137 0 
1 F 
2 2' 
3 3' 
4 4' 
5' 5 
6' 6 
7' 7 
8 
9 
10 
H 
1 
He 
Li 
Be 
B 
C 
N 
O 
F 
2 
3 
4 
5 
6 
7 
8 
9 
Ne 
Na 
Mg 
Al 
Si 
P 
S 
Cl 
10 
11 
12 
13 
14 
15 
16 
17 
A 
K 
Ca 
Sc 
Ti 
V 
Cr 
Mn 
Fe 
Co 
Ni 
18 
19 
20 
21 
22 
23 
24 
25 
26 
27 
28 
Cu 
Zn 
Ga 
Ge 
, As 
Se 
Br 
29 
30 
31 
32 
33 
34 
35 
Kr 
Rb 
Sr 
Y 
Zr 
Cb . 
Mo 
- 
Ru 
Rh 
Pd 
36 
37 
38 
39 
40 
41 
42 
43 
44 
45 
46 
Ag 
Cd 
In 
Sn 
Sb 
Te 
I 
47 
48 
49 
50 
51 
52 
53 
Xe 
Cs 
Ba 
La 
Ce 
(Rare earths, 59-72, unclassified). 
54 
55 
56 
57 
58 
Ta 
W 
- 
Os 
Ir 
Pt 
• 
73 
74 
75 
76 
77 
78 
Au 
Hg 
T1 
Pb 
Bi 
Po 
- 
79 
80 
81 
82 
83 
84 
85 
Nt 
- 
Ra 
Ac 
Th 
UX„ 
U 
86 
87 
88 
89 
90 
91 
92 
TABLE I. 
Periodic System.-Mendeleeff. Not only do the measurable physical properites, but also the chemical 
characteristics show a similar periodicity. If we used figures expressing 
the acidic or basic characteristics of the hydroxides of the elements we 
would find a similar periodic curve. 
The periodic classification of the elements is exceedingly useful in 
correlating a vast number of facts regarding the elements and their 
compounds so that they may be more easily remembered and also pre- 
dicted. However, before discussing specific properties it will be well 
to consider certain defects in the Mendeleeff table, and other ways of 
representing the actual periodic behavior. In the first place, the Mendeleeff 
table throws into the same group such dissimilar elements as sodium and 
potassium, on the one hand, and copper, silver and gold, on the other 
hand. The only justification for this is the fact that the latter elements 
show the same valence in some of their compounds as do the former. In 
the second group zinc cadmium and mercury show a little more resem- 
blance to the main members of the group, though it is by no means 
striking. In each group there is a branching, after the second period, 
into a main group, whose elements resemble closely the first two elements 
of the group, and a sub-group, whose resemblance to the first two 
elements is less marked. This is shown in the table by setting the sub- 
groups to one side. As we go towards the middle of the table we find 
the sub-groups remembling the main group more and more closely, until 
we find, in the fourth group, that the two branches bear about equal 
resemblance to the first two elements. It is rather useful, for example, 
to have vanadium placed, in the same group as phosphorous, less so to 
have chromium classed with sulfur, and of very little use to have man- 
ganese classed with chlorine. The placing of the sub-groups more and 
more to one side as the ends of the table are approached is an attempt 
to avoid this defect in the table by representing this growing dissimi- 
larity, and is very significant in the discussion of properties to be taken 
up later. 
Again, Mendeleeff placed iron cobalt, and nickel and the platinum 
metals in the same group, regarding them as triplets really belonging in 
the same space on account of their nearly equal atomic weights. When 
the atomic numbers are used instead of atomic weights it is found that 
these are not triplets at all, and really belong in separate groups, as we 
have here indicated by numbering them 8, 9 and 10, instead of including 
all these elements in group 8 as Mendeleeff did. 
When atomic weights are used as the basis for the table there appears 
an anomaly due to the fact that the atomic weight of potassium is less 
than that of argon, though the latter should evidently come first. The 
same is true of tellurium and iodine, cobalt and nickel. The use of the 
atomic numbers, however, places all of these in their proper order. 
Another arrangement of the Periodic System is indicated by the 
curves of properties, such as given in Figs. 12 and 13, and regards the 
elements of the sub-group as a new series intruding after the second 
period. This arrangement, first given by the Danish chemist Thomsen, is 
shown in Table II, and though it is not quite so easily remembered, it 
has advantages over the first form. It does not imply relationships between 
the sub-groups and the main groups which are slight or non-existent, 
and it places together such similar elements as nickel and copper, platinum 
and gold, etc. The slight relationships over-emphasized by the Mendeleeff 
table between the main and sub-groups are here indicated by dotted lines. 
If three dimensions are used instead of two it is possible to give a 
139 TABLE II. 
Periodic System-Thomsen. 
GROUPS. 
much more adequate representation of the periodic relationships. The 
Mendeleeff table would then be represented by a spiral on which the 
homologous elements fall below one another. The Thomsen table would 
then show the interpolated groups as branching off from the main spiral 
as shown in Table III, which is drawn according to an idea put forth 
by Soddy. This shows very clearly the nearly equal relationship of the 
two branches of Group 4 to carbon and silicon; the greater removal of 
sub-groups 3' and 5' from their main groups, 3 and 5 respectively, the 
still greater divergence of sub-groups 2' and 6' from the corresponding 
main groups, and the almost entire lack of relationship of groups 
1' and 7' with 1 and 7 respectively. On the other hand it brings out the 
fact shown by all curves of properties, i. e., that the properties of the 
elements in the neighborhood of iron and platinum are farthest removed 
from those of the inert gases in Group 0. The peak represented by Group 
0 is intended to indicate the abrupt change between Group 7 and Group 1. 
After becoming familiar with the basis of the Periodic System the 
student should familiarize himself with the position of the commoner 
140 Periodic System-Soddy. 
TABLE III. 
Loop containirTg 
M rare earth 
elements. elements in some one of the three forms of the system that have been 
given. The designation of the groups and sub-groups has been made 
the same in all of the forms, so that the following discussion will be 
equally intelligible whichever form of the table the student decides to 
adopt. It is far more important to become thoroughly familiar with one 
of the forms than it is to attempt to learn two or three of them with 
possible confusion. 
We will now proceed to a discussion of the relationship between the 
position of an element in the table and its behavior. 
\Density. We have seen that as we go from one period to the next, 
i. e., from top to bottom of each group, there is nearly always a small 
increase in the atomic volume, and that the elements with largest atomic 
volume are in Group 1, while those of smallest atomic volume are among 
the iron and platinum metals. This corresponds to the fact that in general 
the elements increase in density as we go from the top to the bottom 
of each group, and that the heaviest elements are among the platinum 
metals, as illustrated by the following values: 
Group 1. 
Li 0.53 
Na 0.99 
K 0.86 
Rb 1.52 
Cs 1.90 
Group 2. 
Be 1.73 
Mg 1.74 
Ca 1.53 
Sr 2.55 
Ba 3.78 
Densities of Solid Elements. 
Group 2' 
Zn 7.1 
Cd 8.6 
Hg 13.5 
Group 1' 
Cu 8.9 
Ag 10.5 
Au 19.2 
Group 8 
Fe 7.9 
Ru 12.0 
Os 22.5 
Group 10. 
Ni 8.7 
Pd 11.5 
Pt 21.4 
Melting Point. As shown by the curve in Fig. 13, the elements with 
the lowest melting point are in the Group 0. The adjacent groups, 1 
and 7, contain elements of low melting points, increasing as we approach 
Group 4, where we find carbon at the top with the highest of all melting 
points. Another maxim is found with tungsten, whose neighbors also 
have high melting points. Among the metals of the main groups the 
melting points decrease as we descend the group. The following values 
may serve as illustrations: 
Melting Points, Absolute. 
Group 1. 
Li 459 
Na 371 
K . 335 
Rb 311 
Cs 299 
Group 2. 
Be 1573 
Mg 924 
Ca 1083 
Ba 1123 
Group 2' 
Zn 692 
Cd 594 
Hg 234 
Group 1' 
Cu 1356 
Ag 1234 
Au 1336 
Group 8. 
Fe 1800 
Ru 2570 
Os 2970 
Group 10. 
Ni 1725 
Pd 1820 
Pe 2028 
Tensile Strength. The metals with the greatest tensile strength are 
found among the sub-groups, starting with 5' and extending to 10. These 
groups include iron and the metals alloyed with it to make the modern 
alloy steels of such remarkable properties, V, Cr, Mo, W, U, Mn, and Ni. 
Valence. Very striking relationships are found for the valence 
of the elements in the main groups. The inert gases are commonly said 
to have zero valence, corresponding to their group number. The alkali 
metals in group one have a valence of -|-1, the alkaline earth metals 
a valence of -j-2, and the valence of the elements in each succeeding 
group increases with the group number as illustrated by the formulas of 
the oxides and chlorides. 
142 LioO 
LiCl 
Na„O 
NaCl 
BeO 
BeCl2 
MgO" 
MgCl2 
b203 
bci3 
ALO3 
A1C1, 
co2 
CC14 
SiO2 
SiCl4 
N2O5 
p2o5 
PC15 
so3 
C12O7 
With the fourth group and those following lower valences also appear, 
as in CO, P2O3, SO2, etc., and a negative valence becomes evident as 
shown in their compounds with hydrogen. 
ch4 
NH3 
'PH3 
AsH3 
H„O 
H„S 
H„Se 
HF 
HC1 
HBr 
The limiting values of the valences in the main groups seem, therefore, 
to be as follows: 
Group 0 1 2 3 4 5 6'7 
Valence 0 4-1 4-2 4-3 +5 4-6 4-7 
-4 -3 -2 -1 
Other valences lying between these extreme values may be found, as 
shown by the following: 
Valence of P 
P2O5 4-5 
PoO2 4-3 
H3PO2 4-1 
H3P -3 
Valence of Cl 
KC1O4 4-7 
KC1O3 +5 
KC1O2 +3 
KC1O +1 
KC1 -1 
In the sub-groups valences corresponding to the group number 
usually appear, but they are not always the most characteristic, and other 
valences appear which might not be expected on the basis of the Periodic 
System. The following table illustrates this fact: 
Valences 
Cr 2, 3, 6. 
Fe 2, 3. 
Mn 2, 3, 4, 6, 7. 
Pt 2, 4. 
Os 2, 4, 8. 
The horizontal relationships are rather more important than the ver- 
tical relationships with these elements, as illustrated by the fact that the 
valence 4-2 is characteristic of the successive elements, Cr, Mn, Fe, Co, 
Ni, Cu, Zn. Again Cu and Ni are more alike than Ni and Pd, or Cu 
and Ag, and Au is rather more like Pt than Ag. 
It is rather striking that the rare earth elements, which seem to be 
out of harmony with the rest of the system, all have 4" 3 as their most 
characteristic valence. 
Noble and Base Metals. In the main groups there are no noble 
metals, all of the metals in these groups being easily oxidized. The noble 
metals are found rather in the sub-groups and at the bottom. Thus in 
group 2' we have first zinc, a base metal, cadmium, a less base one, and 
finally mercury, a rather noble metal. In group 1' we have copper, 
intermediate between the two classes, then silver, about as noble as 
mercury, and finally gold, one of the noblest of all. At the bottom of the 
adjacent groups 10 and 9 we find respectively platinum and iridium, 
metals which are exceedingly difficult to attack. 
Metals and Non-Metals. It has been pointed out in Chapter V that 
the distinction between metals and non-metals is of fundamental import- 
143 ance in connection with the chemical behavior of the elements. We have 
seen, further, that this distinction may be expressed by saying that the 
more metallic an element is the more prone it is to take on a positive 
valence, while the more non-metallic it is the more prone it is to take on a 
negative valence. Or, we may say more briefly, the more metallic elements 
are more electropositive, while the less metallic elements are more electro- 
negative. 
It must be noted that this distinction is different from that between 
noble and base metals. A highly metallic element is one which is a good 
conductor of heat and electricity, has a high reflective power for light, 
is ductile and malleable, and tends to form compounds in which it is the 
positive element. According to these criteria, sodium, which is so base 
that it rapidly oxidizes in the air and decomposes water vigorously, 
is a more metallic element than the much nobler metal, tin. 
The most electropositive elements are in Group 1, and the most 
electronegative of these is cesium, at the bottom. In all of the main 
groups the most metallic or electropositive element is at the bottom. The 
metallic character decreases as we go from group to group, until, in 
Group 7 we find elements which are all highly electronegative. In Group 
6 we have the non-metallic elements, oxygen and sulfur, at the top but 
the weakly metallic tellurium at the bottom. In Group 5 we find nitrogen 
at the top entirely non-metallic in character, phosphorus non-metallic but 
with a weakly metallic allotropic form, arsenic with metallic luster but 
brittle and volatile, antimony a fairly good metal, and bismuth at the 
bottom quite metallic in nature, forming no compounds in which it shows 
a negative valence as do the others, (e. g. AsH3). In Group 4 we find 
at the top the non-metal carbon with an allotropic modification, graphite, 
in which it shows faint metallic properties. As we approach the bottom 
the metallic character quickly predominates, as in tin and lead. 
This change in metallic character from group to group and from 
top to bottom of each group serves to correlate a vast array of facts 
which we have space only to illustrate. 
Strength of Acids and Bases. One of the most important chemical 
characteristics of the metals is that their hydroxides are bases, as discussed 
in Chapter V. Since we have elements of all degrees of electropositive 
character we have bases of all degrees of strength. The strongly positive 
alkali metals give hydroxides which are powerful bases, and the strongest 
of these is, as we should expect, cesium hydroxide. The hydroxides of 
the metals of Group 2 are weaker, that of beryllium, at the top being a 
rather weak base, though the strength increases as we descend the group, 
until, in barium hydroxide, we have a base approaching those of Group 1 
in strength. 
When a base is sufficiently weak it may often show the complimentary 
behavior of an acid, and in the presence of very concentrated alkali 
beryllium hydroxide will dissolve, so that we may regard it as an ex- 
tremely weak acid. In Group 3 this acid character becomes more 
pronounced, and B(OH)3 (or H3BO3) is a strong enough acid to turn 
blue litmus paper red. The basic character, however, is practically absent. 
There exists, it is true, BC13, which may be regarded as the salt of 
B(OH3 with HC1, but it is completely hydrolyzed by water, and hence 
cannot be formed in aqueous solution. A1(OH3 is, as we should expect, 
a much weaker acid, and is, at the same time a weak base. 
The corresponding compounds with carbon and silicon, H2CO3 and 
144 H2SiO3 (or SiO(OH)2 or Si(OH)4, etc.) are weak acids, H2TiO3 is 
amphoteric, i. e., both acidic and basic), while Th(OH)4 is a base only. 
The hydroxides of tin and lead are amphoteric, with the basic character 
most marked in the latter. As a consequence salts of Sn++ are more 
hydrolyzed then those of Pb++. 
Comparing similar compounds in Group 5 we find that HNO3 is a 
strong acid, HPO3 (or H3PO4) is a moderately strong acid, H3AsO4 is a 
weak acid, H3SbO4 is a very weak acid and a very weak base. With 
the corresponding trivalent compounds we go from the moderately weak 
acid HNO2 to the weak base Bi(OH)3 exactly as we should expect. 
It may aid somewhat in understanding this connection between 
metallic character and the behavior of the hydroxide to compare such 
compounds as KOH. potassium hydroxide, and C1OH (or HC1O) 
hypochlorous acid. In these compounds the potassium and chlorine are 
both postive, but the latter can be only weakly positive since its normal 
tendency is to become highly negative, as in HC1 and KC1. If we attempt 
to represent the relative positive and negative characters of the elements 
involved by the size of the charges we write under them we will get some 
such results as the following: 
H ; C1 C1 O : H 
+ : + ! + 
K i Cl K: O H 
Now when dissociation takes place it tends to break up a compound 
into parts having charges of as nearly equal intensity as possible, accord- 
ingly there is a _weak tendency for C1OH to ionize into CIO- and H+, 
which makes it a weak acid, while KOH ionizes strongly into K+ and OH-, 
making it a strong base. In a compound like AIO(OH), aluminum 
hydroxide, where the aluminum and the hydrogen have about the same 
"positivity" OH- or H+ may split off with nearly equal readiness, so that 
the substance is'amphoteric. 
Amphoteric Sulfides. The relationships involving the solubility of 
sulfides pointed out in Chapter XII may also be correlated in accordance 
with the Periodic System. Thus, of the three sulfides, As2S3, Sb2S3 and 
Bi2S3, the first, being the sulfide of a rather non-metallic element, is suffi- 
ciently acidic to dissolve easily in sulfide ion, S--, forming AsS2-; the 
second, Sb2S3, dissolves with difficulty while Bi2S3 is insoluble in 
this reagtent. 
Stability of Hydrides. In the hydrogen compounds of the elements 
in groups from 5, 6 and 7, the hydrogen is doubtless always the positive 
element, and therefore the more electronegative the other element the 
more stable is the compound. Nitrogn, for example, is non-metallic and 
assumes the negative valence rather easily, hence ammonia, NH3, is a 
moderately stable compound which will withstand moderate heating 
145 without decomposing. As we descend the group we find that PH3, AsH3 
and SbH3 become increasingly easy to decompose by heating, while 
the last element of the group, bismuth, is so metallic that it will not 
assume the negative valence at all, hence gives no compound with hydrogen. 
Oxygen, the first member of Group 6, is far more electronegative than 
nitrogen, and therefore gives a much more stable hydride, water, which 
will break down into its elements only at extremely high temperatures, 
being only about 2% dissociated at 2000°C. Hydrogen sulfide is con- 
siderably less stable, and hydrogen telluride, finally, decomposes spon- 
taneously, and is prepared with difficulty. 
A similar decrease in stability is found in Group 7 as we descend 
from the very stable hydrogen fluoride and hydrogen chloride to the 
rather easily decomposable hydrogen iodide. 
Comparison of Oxides and Halides of Other Positive Elements. The 
relationships with compounds of the metals with the members of Groups 
5, 6 and 7 are parallel to those already noted with the hydrides. For 
example, cupric iodide does not exist, while cupric chloride and bromide 
do. Copper is not sufficiently electropositive and iodine sufficiently elec- 
tronegative for this compound to be stable, consequently the attempt to 
prepare it results instead in the formation of cuprous iodide and iodine. 
Cu++ -|- 21" = Cui 
Molybdenum trioxide exists, MoO3, but chlorine, which is less electro- 
negative than oxygen, cannot call forth the highest valence of molybdenum, 
which forms only a pentachloride, MoC15. With fluorine, however, which 
is the most electronegative of all the elements, the molybdenum can again 
be hexavalent, and we have the hexafluoride, MoF6. 
If we compare the compounds of sulfur with bromine, chlorine, 
oxygen and fluorine, which is the order corresponding to increasing electro- 
iegative character, we find that highest bromide has the formula S2Br2; 
a chloride exists, SC14, in which the sulfur has a higher valence; in 
sulfur trioxide, SO3, sulfur has its maximum positive valence, but the 
compound breaks down rather completely above 700°C into SO2 and O2; 
and finally the fluoride, SF6, is an exceedingly stable inert substance. 
Illustrations of this sort might be multiplied almost indefinitely. 
Solubility of Salts. The solubility of salts changes in regular order 
according to the Perodic System in nearly all instances. This may be 
illustrated for a single group, the second, by the following table: 
Hydroxide 
Sulfate 
Chloride 
• 
Magnesium 
Insoluble 
Soluble 
Deliquescent 
Calcium 
Sparingly 
Soluble 
Sparingly 
Soluble 
Deliquescent 
Strontium 
Moderately 
Soluble 
Insoluble 
Soluble 
Barium 
Soluble 
Very 
Less Soluble 
Insoluble 
It may be mentioned that a similar order is shown by the more soluble 
salts of the alkali metals, where, for example, the least soluble sulfate 
is that of cesium, at the bottom of the group. 
Elements that Form Complex Ions. The elements- most prone to form 
complex ions are in the sub-groups and especially in the neighborhood of 
146 Group 1' in the second and third forms of the table. Complex ions 
with ammonia are given by Zn, Cd, Cu, Ag, Ni and Co. Complex 
cyanides are given by the same elements, also by the platinum metals and 
iron. Complex chlorides are given chiefly by the platinum metals and 
gold, as illustrated by PtCl6~" and AuC14~. 
Resemblance Between the First Member of a Group and the Second 
Member of the Succeeding Group. There appears a striking similarity 
between the first element in Groups 1 and 3 and the second element in 
the succeeding group. Thus lithium, while an alkali metal, resembles 
magnesium in having a hydroxide and carbonate which are not very 
soluble, a rather insoluble phosphate and a deliquescent chloride 
Beryllium is so much like aluminum that the two elements are difficult 
to separate. Boron is like silicon in physical properties and gives similar 
compounds in many respects. The borates of all but the alkali metals, 
like the corresponding silicates, are insoluble in water, and when fused 
tend to give glasses on cooling. 
Radioactive Elements. As might be expected, it is the heaviest and 
most complex atoms which tend to break down into simpler ones, and 
all of the elements of higher atomic weight than bismuth are radioactive. 
These are, accordingly, found at the bottom of the table. 
Prediction o£ Unknown Elements. The value of the Periodic System 
was demonstrated in very striking fashion by the prediction of unknown 
elements by Mendeleeff. At the time he constructed his table it was 
obvious that several gaps should be left in order to have the succeeding 
elements fall into their proper groups. Such gaps were to be found at 
the places now filled by scandium, gallium and germanium. From the 
properties of the adjacent elements Mendeleeff predicted the properties 
of the unknown elements. The remarkabl accuracy of these predictions 
is illustrated by the following comparison of the unknown element called 
by Mendeleeff "ekasilicon," with germanium, discovered later by Winkler. 
Ekasilicon. 
At. wt. 72, density 5.5 
Oxide EsO2, density 4.7 
Chloride EsC14, liquid, boiling 
slightly below 100°, density 1.9. 
Ethide, Es(C2H,)4, liquid, boiling 
at 160°, density 0.96. 
Fluoride, EsF4, not gaseous. 
Germanium. 
At. Wt. 72.5, density 5.46. 
Oxide GeO2, density 4.7. 
Chloride, GeCl4, liquid, boiling at 
86°, density 1.887. 
Ethide, Ge(C2H5)4, liquid, boiling 
at 160°, density slightly less than 
1. 
Fluoride, GeF4, while, solid. 
The foregoing comparison should serve as an illustration of the 
value of the Periodic System to the student of chemistry, for if an 
undiscovered element can be foretold so brilliantly by its aid, then any 
knowledge a student may have some elements may be similarly projected 
to deduce the proprties of adjacent elements. The use of the Periodic 
System, therefore enormously increases the effectiveness of a given amount 
of mental effort, and it is very important to form the habit of using it 
constantly. 
147 CHAPTER XVII. 
DISPERSED SYSTEMS. 
In Chapter II a distinction was made between heterogeneous and 
homogeneous material, between mechanical mixtures and solutions. If 
we have a bottle containing water in the lower portion and moist air in 
the upper we would not hesitate to call the system heterogeneous. Even 
if the water were blown into the air from an atomizer, we would doubtless 
still considr the system heterogeneous. Suppose, however, we could break 
up the droplets into successively smaller and smaller ones, ending finally 
with molecules of water vapor in air. The resulting system we would 
have to consider homogeneous, a solution. When, we may well ask, 
did the system cease to be heterogeneous and become homogeneous? We 
might imagine a similar gradation starting with sand and water, which 
would be heterogeneous, through successively finer and finer suspensions 
of silica powder in water until we would have a solution of silica molecules 
in water. There is, evidently, no sharp boundary between systems cer- 
tainly heterogeneous, on the one hand, and solutions on the other, and 
between these two extremes there are systems which should have some 
of the charactristics of both. These we will call dispersed systems, one 
phase being highly dispersed in another. The possible systems may be 
clasified according to the nature of the dispersed phase and the enclosing 
phase as follows: 
Dispersed Phase 
Liquid 
Gas 
Solid 
Gas 
Liquid 
Solid 
Liquid 
Solid 
Enclosing Phase 
Gas 
Liquid 
Gas 
Solid 
Liquid 
Liquid 
Solid 
Solid 
Type 
Fog 
Foam 
Smoke 
Solid foam 
(e. g. pumice) 
Emulsion 
Suspension 
The most important of these from a chemical standpoint are emulsions 
and suspensions, and our attention will be largely confined to them. 
Brownian Movement. As the dispersed phase in suspension and 
emulsions becomes more and more finely divided it is possible to get 
systems in which the two phases do not separate on standing under the 
influence of gravity. A relatively large body like a grain of sand, dropped 
into water, will fall through the water at a rate of speed dependent on 
its size, the difference in density between it and water, and the viscosity 
of the water. A very small sand particle will fall more slowly through 
the water, because as the size diminishes the surface, upon which the 
resistance depends, does not decrease as fast as the weight. Fine 
suspensions, therefore, settle out more slowly than coarse ones. When 
a sufficiently minute particle is observed with a microscope it may be 
seen that it no longer falls in a straight line, but follows an irregular 
zigzag path. The reason for this is as follows. The molecules of water, 
according to the kinetic theory, are in rapid motion with a mean kinetic 
energy proportonal to the absolute temperature. Now in disordered 
148 motion of this sort some molecules, at a given instant ,will be moving 
more rapidly and some more slowly than the average. Now when the 
particle of sand is so large compared with the water molecules that it 
is being bombarded by a very large number of the latter, the effect of 
their difference in kinetic energy will be neutralized; but when the 
particle gets sufficiently minute an especially hard impact on one side 
of it may not be neutralized by the impacts of other molecules on the 
other side, hence the particle will be displaced slightly. Since these 
abnormallc hard impacts are constantly occurring irregularly from 
different directions the result is the irregular zigzag movement of the 
particle above referred to. As smaller and smaller particles are taken 
the chances for inequality in the momentum or the impacts upon them 
increase, and also their resulting displacement increases, for a molecule 
of water striking a small body will have more effect on it than upon 
a larger one. In fact it can be seen by the aid of the microscope that the 
particles of a fine suspension are in much more violent motion than 
those of a coarse suspension. This movement is known from its discov- 
erer as the Brownian movement, and gives us very direct evidence of 
the truth of the kinetic theory. 
When sufficiently fine particles are taken the Brownian movement 
overcomes the effect of gravity so that the particles remain in permanent 
suspension. The system is then known as a colloidal suspension, or 
colloidal solution. There will be, in such a suspension, if time is allowed 
an increasing concentration of particles from the upper to the lower 
portions of the vessel, analogous to the increasing concentration of the 
atmosphere from high altitudes to lower. In face a French physical 
chemist Perrin, showed that the same rate of increase of the particles in 
.'suspension obeys the same laws that govern the change in barometric 
pressure with altitude, a most striking proof of the truth of the kinetic 
theory and of the real existence of molecules. 
The dividing line between suspensions which settle out on standing 
and colloidal suspensions, which remain permanently suspended, is purely 
one of convenience, for by applying centrifugal force it is possible to 
separate suspensions that would not do so under the influence of gravity 
alone. In general colloidal suspensions cannot be separated by filtration 
through ordinary filters. 
It is possible to distinguish in practice between true solutions and 
colloidal solutions or suspensions by reason of the great difference they 
show in the rate of diffusion. The Brownian movement of colloidal 
particles is so much less than the movement of the much smaller molecules 
of substances in true solution that the latter diffuse from regions of 
greater to those of smaller concentration with much greater rapidity. The 
first distinction between the two types was made upon this basis by 
Graham, in 1862. He obtained the following figures for the time required 
for the equal diffusion of various substances: 
HC1 1 
NaCl 2.3 
Sugar 7 
MgSO4 7 
Albumen 49 
Caramel 98 
To the class of substances which diffuse slowly he gave the name 
colloids, from the Greek name for glue, a member of this class. 
The molecular weight can be calculated from the rate of diffusion, 
and also from the lowering in vapor pressure or in freezing point, as 
explained in Chapter IV. These methods all agree in ascribing very 
149 high molecular weights to substances in colloidal solution as is to be 
expected. The following figures illustrate the values found: 
Substance 
Molecular Weight. 
Egg albumen 
20,000 
Gelatin 
40,000 
Pepsin 
12,000 
Optical Properties of Colloidal Suspensions. When a colloidal sus- 
pension is viewed by transmitted light it usually appears homogeneous, 
like a true solution, its color depending on the size of the particles and 
on the nature of the material. By reflected light we usually notice a 
somewhat different color and an opaque muddy appearance similar to 
that of ordinary suspensions. The microscope will detect particles as 
small as O.OOO25 millimeter. Smaller particles than this cannot be 
viewed directly as they are smaller than the wave length of light, so that 
the latter cannot be regularly reflected from them. It is possible, however, 
to detect the presence of particles as small as 0.000006 millimeter, by 
looking at them in a strong transverse beam of light against a dark 
background. (The diameters of the molecules of most ordinary substances 
are less than 0.000001 millimeter.) Although no real image of the 
particle may be seen there is a scattering of the light by it so that a 
bright speck of light can be seen in the microscope. The same effect is 
responsible for the visibility of a beam of light in a dark room through 
illumination of dust particles invisible, perhaps, even in a microscope. 
A microscope arranged for use with powerful transverse illumination 
against a dark background is called an ultramicroscope. By knowing 
the concentration of the material in a colloidal suspension, and counting 
the number of particles in a tiny beam of known dimensions it is possible 
to determine the average size of the particles. 
Adsorption. Solid surfaces have the power of condensing gases upon 
them. The amount of gas so adsorbed depends upon the nature and 
area of the surface, the temperature, and the nature of the gas. The 
lower temperature the more slowly do the gaseous molecules move, and 
the more easily are they held by the attraction exerted by the solid surface, 
hence the gases which are most easily condensed to the liquid state are 
also more highly adsorbed by solid surfaces. Of all gases helium, which 
boils at 4.5°A, is least adsorbed, followed by hydrogen, whose boiling 
point is 20.5°A, then by gases like nitrogen, oxygen, etc., then by 
ammonia, water vapor, etc. It is a fact familiar to chemical analysts 
that considerable amounts of water may be condensed upon the surfaces 
of pieces of apparatus like crucibles and glass beakers, so that it is 
necessary to have them in a dry atmosphere in order to get exact and 
reproducible weights. If the solid is porous, like charcoal, or finely 
powdered, its surface is vastly increased, allowing the same amount to 
adsorb very much more of the gas. The use of charcoal to adsorb odorous 
gases in this way is very familiar. A similar adsorption of dissolved 
substances can take place from solutions upon solid surfaces immersed 
in them. "Bone black," formed by the strong heating of ground bone, 
turning its animal matter into charcoal, is used in the refining of sugar 
to remove the yellow coloring matter present. The colored bodies are 
mainly substances with large molecules which are readily adsorbed on 
the large surface of the bone-black, from which the solution of sugar 
can be filtered clear and colorless. Cotton seed oil is decolorized by 
fuller's earth in the same way. On account of the large surface even 
150 a small amount of solid in colloidal suspension (a cube having a diameter 
of 1 cm. and a surface of 6 cm.2, is divided into cubes with a diameter 
of 0.0001 cm. would have a surface of 6000 cm.2) the adsorption of 
dissolved substances upon the colloidal particles exerts an important effect 
upon their properties. 
Electrical Migration of Colloidal Particles. An especially important 
effect upon the properties of fine suspensions is exerted by adsorbed 
ions. It is seldom that electrolytes are absent from aqueous suspensions, 
and even where such is the case, the ions of water may be preferentially 
adsorbed. Certain kinds of suspensions are prone to adsorb positive 
ions in excess, whereby the particles acquire a positive charge, while 
others, in a similar way, become negative. As a result they are able to 
migrate with an electric current flowing through the solution, the direction 
depending upon the sign of the adsorbed ions. The speed with which 
they move is comparable to that of the slower ions, about 7 mm. per hour 
under a drop of potential through the solution of 1 volt per cm. It is 
thus possible to classify suspensions as positive or negative, just as the 
ions themselves were classified in Chapter X. Among the positive 
suspensions are those of the metallic hydroxides and sulfides, while the 
negative suspensions include arsenous sulfide, silicic acid, graphite and 
the-noble metals. 
Relation Between Charge and Stability. Coagulation. Even though 
a suspended particle may be so small that its Brownian movement prevents 
its settling out under the influence of gravity, a suspension of such par- 
ticles may not be stable, because of the tendency of the particles to 
adhere to one another on impact, forming successively larger aggregates 
until settling out can occur. This process may be seen when silver 
chloride is formed by mixing its ions. At first a very fine suspension 
is formed, which, if not too concentrated, appears quite stable, and may 
run through an ordinary filter. On standing, however, and more rapidly 
if the .solution is heated and stirred, these extremely fine particles adhere 
to each other, forming eventually the familiar curdy precipitate of silver 
chloride. Now the effect of electrolytes upon the stability of such a 
suspension is very great. The presence of a charge upon the particles 
tends to make them repel each other and to prevent coagulation, so that 
a small amount of the proper electrolyte increases the stability. Suppose 
that we are dealing with a negative colloid like As2S3. If some electrolyte 
is added whose positive ion tends to be highly adsorbed it may partly dis- 
place the ion giving the negative charge, leaving the particle neutral and 
free to coagulate with others like it. In general hydrogen and hydroxide 
ions are highly adsorbed, so that the addition of a small amount of acid 
to a suspension of arsenous sulfide suffices to coagulate it while the 
addition of a small amount of alkali to a positive suspension like that 
of ferric hydroxide coagulates it completely. In general, then negative 
suspensions are coagulated by acids and positive suspensions by alkalis. 
There appear to be no striking differences in the extent to which most 
other ions are adsorbed, but the effect of the valence of the adsorbed 
ion upon its ability to coagulate a suspension is very great. Let us sup- 
pose that Na+, Ca++ and Al+++ are adsorbed by suspended As2S3 to the 
same degree. It is evidently not the amount of substance adsorbed but 
the amount of positive charge that it carries with it that determines its 
ability to neutralize a negative suspension. To have the same amount 
of electricity adsorbed we will need but half as much of the divalent 
151 Ca++ and but one-third as much of the trivalent Al+++ as we will of the 
univalent Na+. Hence it will take very much less Ca++ and still less 
Al+++ to have the same effect as a given amount of Na+. Moreover 
it is the amount adsorbed, and not the amount in solution that determines 
the coagulating effect, and since the concentration of a substance in 
solution increases more rapidly than the amount adsorbed, the concen- 
tration of Na+ in the solution must be considerably more than three times 
that of the Al+++ in order that the amount of the former adsorbed be 
three times the latter. 
Use is made of the great coagulating power of aluminum ion upon 
the negative suspension of mud in water in purifying it for water 
supplies. Aluminum sulfate is extensively used in connection with 
filtration plants to increase the efficiency of the filtration. The aluminum 
hydroxide formed by the hydrolysis of the aluminum sulfate is itself a 
positive colloid, and aids considerably in the coagulation, since a positive 
colloid acts like a positive ion in neutralizing the charge on a negative 
suspension. 
The ions in sea water have a similar coagulating effect upon muddy 
river water discharged into it, a factor which aids in the formation of 
bars at river mouths. 
Most dyes are very fine suspensions which are themselves adsorbed 
or absorbed in animal or vegetable fibres. When the fiber is unable to 
hold the dye permanently the color may frequently be made "fast" by 
using a coagulant, called a "mordant." The mordant is the salt of a 
polyvalent ion like Al+++ or Sn++++, which can diffuse into the fiber and 
hydrolyze giving a positive hydroxide, which can hold the dye firmly. 
A "lake" is a pigment made by such coagulum of a hydroxide like 
Al (OH) 3 with a suitable dye. The famous "purple of Cassius" is a 
similar coagulum of gold and stannic hydroxide, formed simultaneously 
in the reaction of stannous ion with chlor-aurate ion, as follows: 
2AuC14- + Sn++ 4- 12H2O = 2Au + 3Sn(OH)4 + 12H+ + 8C1-. 
In washing precipitates it is often found that the precipitate tends to 
become suspended and run through the filter. In such cases it is possible 
to prevent this re-suspension by washing with water containing a little 
acid, or ammonium salt, or ammonia, etc., as the case requires. 
The effect of the valence of negative ions in coagulating positive 
suspensions follows the principles laid down for the corresponding effect 
upon negative colloids. Ferric hydroxide suspension, for example, is 
more easily coagulated by sodium sulfate than by sodium chloride, and 
still more easily by sodium phosphate. 
Emulsions.. We have thus far been considering dispersed systems 
in which a solid is suspended in a liquid. We will now consider those 
systems in which the suspended phase is a liquid, called emulsions. Much 
of what has been said in regard to solid suspensions applies here. When 
the suspended droplets are small enough, as in cream, they are subject 
to Brownian movement, though in this case not sufficient to prevent their 
rising to the top by reason of gravity. The optical properties are similar 
to those suspensions. It is possible, however, to have stable emulsions with 
the suspended material less highly dispersed and in larger amounts than 
is the case ordinarily with solid suspensions, as illustrated by cream and 
mayonnaise dressing. In order to do this, however, an emulsifying agent 
must be used, for two pure liquids will not form a stable emulsion. The 
reason for this is as follows. 
152 Whenever it is possible for two suspended droplets to come together 
they will coalesce into one, for the surface tension operates to make the 
surface as small as possible, and one spherical drop will have less surface 
than two of the same total volume. In order to prevent their coalescing, 
it is desirable to have the liquid in which they are suspended viscous, but 
especially is it necessary that the film of liquid which separates them 
when they are close' together should be difficult of rupture, like the soap 
bubble film which prevents two adjacent air bubbles from coalescing. 
Now a stable film can be produced by some substance which greatly 
lowers the surface tension of a liquid. An exact proof of this could be 
given, but the following simpler explanation may suffice. The diminu- 
tion in the surface tension is an evidence that the substance added 
diminishes the attractive forces of the molecules within the liquid, conse- 
quently there will be a tendency for molecules of the solvent to restore 
the condition of the pure solvent, where the molecules are probably closer 
together, by squeezing out, as it were, the molecules of the solute, which 
will threfore tend to be concentrated at the surface, still further lowering 
the surface tension. In fact we find that the surface tension of soap 
solutions is much less than that of water, and also that the tension at a 
surface that has stood for a short time is less than it is at a fresh surface, 
as illustrated by the following figures: 
• Surface Tension 
Fresh Surface Old Surface 
Water 79 79 
1.25% Soap Solution 62 26 
Suppose, now, that we consider a film of such a solution, representing 
in Fig. 14a the greater concentration of the soap at the surface by the 
Fig. 14. 
shading. If a strain is put upon the film so that it starts to rupture at 
a certain place, represented at b, the effect is to bring fresh solution from 
the interior to the surface at this point, solution in which the soap is at 
first less concentrated, and whose surface tension therefore greater than 
that of the old surface. Consequently, whenever rupture threatens the 
film becomes automatically stronger, and hence it is stable. Sufficient 
soap dissolved in water will therefore give both a stable foam, if shaken 
with air, or a stable emulsion if shaken with oil. Other substances like 
saponin, gelatin, albumen, the casein of milk, etc., may have the same 
effect in greater or less degree. The substances in the egg act in this 
way in making mayonnaise. Milk, when concentrated in the form of 
evaporated milk, can serve also for making mayonnaise. 
It is possible to have either phase the enclosed phase, the oil being 
the enclosing phase in mayonnaise that has "separated," and in the water- 
petroleum emulsions that often give trouble in oil fields. In this case 
the crude oil is the solution, and hence can give stable films separating 
water drops. An emulsion of water in benzene can be made by dissolving 
in the latter a magnesium soap which lowers its surface tension. There 
is a tendency, also, for the phase present in smallest amount to be the 
153 enclosed phase, and one of the secrets of success in making mayonnaise 
is accordingly to wet the sides of the vessel with egg, and to add the oil 
slowly at first. 
Jellies. When the enclosed phase of an emulsion is present in 
sufficient amount the droplets begin to press on each other interfering 
with their spherical tendency and substituting a structure which would 
appear, in cross section, somewhat like a honeycomb. The droplets can 
no longer slide over each other rapidly, and any stress tending to deform 
the mass results in deforming the droplets, increasing their surface. On 
the removal of the stress their surface tension operates to restore their 
symmetry. Consequently such an emulsion is elastic, and represents one 
type of jelly. 
Emulsion Colloids. Allied in many respects to emulsions are solutions 
of many substances produced by animal and vegetable life, including glue, 
gelatin, albumens, tannin, gums and starch. These properties may be 
deduced, to a considerable extent, from the supposition that they are 
highly dispersed emulsions, the two phases, unlike the familiar oil-water 
emulsions, being soluble in each other to a considerable extent, so that 
the surface tension between them is very slight. On this account there 
is little adsorption of ions at their bounding surfaces, and their stability 
is little affected by ionic charges. This distribution of water between the 
two phases may, however, be considerably changed by the addition of elec- 
trolytes, even to the extent of "salting out" the colloid, as is done in 
making soap when common salt is used to make the soap separate from 
the solution. The effectiveness of various electrolytes in salting out 
colloids follows roughly the order in which they affect the surface tension 
of water. The effectiveness of various anions decreases in the following 
order: citrate, sulfate, acetate, chloride, nitrate, iodide, sulfocyanate, so 
that sodium sulfate, for example, would be far more effective than sodium 
nitrate. 
On account of the ease with which most colloids of this type remain 
in solution in water in large concentrations, they are often called hydro- 
philic colloids, those which "like water," as distinguished from the rela- 
tively unstable hydrophobic colloids, those which "hate water," which we 
considered earlier. 
Just as a fine emulsion is viscous, if at all concentrated, so these 
emulsion colloids, or hydrophilic colloids, are often very viscous, as illus- 
trated by glue, or gelatine solutions. Furthermore, when an emulsion 
contains enough of the enclosed phase so that the droplets deform each 
other, we begin to get the properties of a jelly. Similarly, by lowering 
the temperature, or on evaporation of water, the two phases apparently 
present in solutions of emulsion colloids usually become of such relative 
amounts as to produce gelatinization. The swelling or shrinking of these 
jellies when placed in solutions of different electrolytes plays an important 
part in biological processes, since much of the material of the living 
organism is colloidal in nature. 
As one phase of an emulsion becomes more viscous it approaches a 
solid in properties, hence there is no sharp distinction between a solid 
suspension and an emulsion. The corresponding jelly, moreover, may be 
less like the jelly of a true emulsion, like stiff mayonnaise, and more like 
wet sponge, where we have a liquid dispersed within an elastic fibrous 
solid. Silicic acid jelly seems to be of this nature. 
Protective Colloids. If an oil emulsion were mixed with a metallic 
154 suspension the oil drops- would tend to enclose the metal particles, and the 
resulting system would behave essentially like a pure emulsion. Similarly, 
when an emulsion colloid, or hydrophilic colloid, is mixed with a hydro- 
phobic colloid the mixture behaves like a colloid of the former class. 
Since the hydrophilic colloids are relatively stable, a hydrophobic colloid 
may be made stable in this way. For example, if gelatin is present in 
a solution in which silver chloride is formed from its ions, each particfe 
of the silver chloride is coated with gelatin so as to prevent the coagulation 
that would otherwise ensue. The retention of this highly dispersed state 
in the solidified gelatin makes possible the photographic dry plate. 
Again, when gelatin is added to a suspension of arsenous sulfide it 
requires much more concentrated acid to coagulate it. 
Preparation of Colloidal Solutions. In addition to the colloids 
obtained from biological sources, like gelatin, albumens, starches, etc., 
it is possible to prepare highly dispersed systems of others by various 
methods which may be classified as dispersion or condensation methods, 
according, respectively, as we start with undispersed material or with 
molecularly dispersed material, molecules and ions. 
Dispersion Methods. Colloidal solutions of the nobler metals may 
be prepared by passing an arc between electrodes of that metal under 
water. The metal is vaporized in the arc and the vapor is condensed to 
solid particles in a highly dispersed state. 
Previously coagulated material may often be redispersed by washing 
out the coagulant, as occurs in washing certain precipitates. Basic ferric 
acetate, precipitated by boiling a solution containing ferric acetate, may 
be redispersed by washing with water containing a trace of acid. 
Some substances can be dispersed by the aid of a "peptonizer," 
usually a protective colloid, taking its name from the action of pepsin in 
digestion in dispersing aluminous material in the process of digestion. 
Thus graphite can be peptonized, or suspended, by grinding itx with 
tannin. The resulting solution is known commercially as "aquadag," 
and can be used as a lubricant. The cleansing action of soap, which is in 
the colloidal state when dissolved in water (not when in alcohol), is of 
this nature. 
Condensation Methods. By starting with ordinary solutions it is 
possible to get suspensions by suitable means. Thus a solution of a ferric 
chloride always contains free H+, Cl- and Fe(OH)3 (or basic chloride), 
the last of which is partly agglomerated, so that it will not pass through 
a very fine membrane like parchment. Accordingly, if such a solution 
is separated by parchment from pure water the H+ and Cl~ can diffuse 
out, the Fe(OH)3 -accumulating as the hydrolysis proceeds, as represented 
by the equation 
Fe^ + 3H2O = Fe(OH)3 4- 3H+ + 3C1" 
By continually renewing the water into which the hydrochloric acid diffuses 
the hydrolysis may be made complete, leaving a colloidal solution of ferric 
hydroxide. Such a process of separation by differential diffusion is called 
dialysis. Other hydroxides can be gotten in colloidal solution by the 
same method. That a colloidal solution is obtained by dialysis and not 
by the addition of alkali is due to the coagulating effect of the larger 
amount of electrolyte, chloride and hydroxide ions especially, which would 
be present under the latter circumstances. 
Whenever a relatively insoluble substance can be produced in the 
155 absence of any considerable amount of electrolytes, it is possible to 
prepare it as a colloid*. Thus the action of H2S on a solution of 
As2O3 will give a suspension of As2S3 since the substances involved 
the still practically un-ionized. A suspension of HgS can be gotten by the 
action of H2S on Hg(CN)2 for the same reason. Colloidal gold can 
be prepared by the reduction of a very dilute solution of HAuC14 by some 
non-ionized reducing agent such as formaldehyde, CH2O, or a solution of 
phosphorus in ether. The addition of some protective colloid, like gelatin, 
allows the solution of the metal to be prepared in much more stable form, 
and more concentrated. The ''argyrol," used in curing cold in the head 
is a concentrated protected solution of colloidal silver. 
Exercises. 
1. Discuss the conditions necessary for the existence of stable foams 
and emulsions. 
2. Summarize the characteristics of hydrophilic and hydrophobic 
colloids, respectively. 
3. It having been found that a certain colloid migrates with the 
positive current, decide whether it would be coagulated more easily by 
HC1 or NaOH. 
4. When a substance is dispersed in a liquid the behavoar of the 
system will depend partly upon the following factors: (a) the size of 
the dispersed particles, (b) the magniture of the surface tension between 
the phases, (c) whether the dispersed phase is solid or liquid, (d) whether 
the dispersed particles adsorb positive or negative ions more readily. 
Decide what effect, if any, each of the above factors would have upon each 
of the following phenomena: (1) the Brownian movement, (2) the 
electric migration, (3) the rate of diffusion, (4) coagulation, (5) separa- 
tion on standing, (6) adsorbing power, (7) protective action on other 
dispersed substances. 
5. What conditions should be observed in the formation of colloidal 
solutions? 
6. Outline experiments to enable you to determine the most effective 
means for coagulating a colloidal suspension of an unfamiliar substance. 
7. Make a classification of dispersed systems that will, in your 
opinion, best express their various behaviors. 
156 INDEX 
Acetanilide 84 
Acetates 97, 102 
Acetic acid 45, 77, 79, 90, 93, 96 
Acetone 79 
Acids 44, 47, 78 
nomenclature 56 
Acid salts 57, 58, 78 
Acetylene 83, 121 
Addition compounds 42 
Adsorption 1 148 
Air 11 
specific gravity of gases referred to 34 
Albumen 147, 152 
Alcohol t 76, 79 
Alkali metals ...39, 142 
Alkaline earth metals 39 
Alkalis 46, 142 
Aluminum 39, 79 
chloride 112 
hydroxide 102, 112, 142 
oxide 105, 112 
solution by acids and alkalis 129 
sulfate 150 
use in cleaning silver 128 
Ammonia 45, 100, 143 
complex ions with 101, 110, 145 
synthesis .-. 81, 84, 116, 120 
Ammonium 
hydroxide 45, 78, 91, 93 
ion, solution of bases by 105 
sulfide 107 
Amphoteric hydroxides 106, 129 
Amphoteric sulfides 107, 143 
Analysis 11 
groups in , 108 
volumetric .62 
Anhydrides 47 
Antimonic acid 143 
Antimony 142 
sulfides 107, 143 
Aqua regia 130, 133 
Argyrol 154 
Arsenic 142 
oxides 107 
sulfides.... 107, 143, 149 
Arsenic acid : 59, 143 
Arsine 144 
Atom of electricity 74 
Atomic numbers 137 
Atomic theory 25 
Atomic volumes... 136 
Atomic weights, 
calculation of 29 
choice of 26 
table of 13 
Atoms 26 
Avogadro's Rule 22, 31 
Baking soda 7, 103 
Barium 39 
chloride 144 
hydroxide 45, 144 
oxide 50 
peroxide 50 
sulfate.. 97, 99, 112, 144 
Base metals 40, 141 
Bases 45, 142 
Basic salts .57, 109 
Batteries ; 8, 128, 133 
Benzene 121, 151 
Beryllium 142, 145 
Bicarbonate ion 93, 103, 108, 109 
Bichromates 58, 124, 131 
Binary compounds 55 
Bismuth 142, 144 
chloride , 57, 104 
hydroxide 57, 104, 143 
oxychloride 57, 104 
sulfide 107, 133, 143 
Bisulfates 58 
Boneblack 148 
Borax 58, 103, 109 
Boric acid 45, 58, 78, 103, 142 
Boron 145 
Bromide ion..... 110, 131 
Bromine 41, 131, 144 
Brownian movement......... 146 
Cadmium, 
complex ions 101, 145 
halides 78, 79, 110 
sulfide 107, 110 
Calcium 39, 78 
carbonate 43, 44, 87, 108, 112, 113 
chloride 44, 144 
hydroxide 45, 113, 144 
nitrate 121 
phosphates ..100, 109 
sulfate 44, 109, 112, 144 
Calorie 65 
Caramel 147 
Carbolic acid 76 
Carbon ...42, 142, 152 
dioxide 43, 45 
monoxide .41 
Carbonic acid 45, 78, 91, 93, 108, 142 
Casein 151 
Cast iron- 11 
Catalysis 84, 119 
Charcoal 148 
Chemical affinity i . 128 
Chemical reactions, 
control 8, 81, 87, 116 
speed of ..82, 118 
types 41 
weight relations in .29 
volume relations in ..32 
Chlorate ion 132 
Chlorine.... 539, 41, 124, 130, 132 
Chloroform 76, 79 
Chromate ion ..124, 132 
Chromic ion 124, 131 
Chromic oxide 105 
Chromium 140 
hydroxide 106 
Classification 
of compounds ..43 
of elements 13, 39, 136 
of reactions 41, 42 
Cleaning silver 128 
Coagulation of colloids.... ...148 
157 Cobalt, 
complex ions 101, 145 
sulfide 107 
Coke 41 
Colloids 147 
Combustion 41 
Complex acids - 57 
Complex ions 101, 110, 130, 144 
Compounds 11 
classification of 51 
Concentration, 
effect on equilibrium 87 
effect on reaction velocity 82 
of gases 18 
molal 60 
normal 61 
"Contact process" 118 
Copper 72, 75 
complex ions of 101, 110, 145 
oxidation 127, 144 
sulfide 107, 110, 133 
Cotton-seed oil 148 
Coulomb 73 
Cryolite 79 
Cupric ion 72, 75 
Cyanide ion 101, 110, 130 
Cyanides 55, 101, 110, 130, 145 
"Cyanide process" 130 
Densities of the elements 140 
Diffusion 15 
Dispersed systems 146 
Dissociation, 
constants 92 
electrolytic 71 
thermal 43 
Dry cell 8 
Dyes 150 
Electricity, quantity per gram-ion....73 
Electrolysis 12, 40 
Electrolytic dissociation 71 
Electromotive force 128 
Electron 74 
Electropositive and electronegative 
character of elements, 39, 49, 74, 141 
Elements 11, 12 
classification 13, 136 
melting points 137, 140 
nomenclature 55 
Emulsions 46, 146, 150 
Endothermic reactions 66, 117 
Enzymes 85 
Equations, 
interpretation 28, 35 
writing 49, 52,' 123 
Equilibrium, effect of, 
concentration •_ 87 
pressure 116 
temperature 117 
Equivalent 61 
electrochemical 73 
Ether 79 
Ethyl alcohol 79 
Exothermic reactions 66 
Faraday of electricity... 73, 128 
Ferric 
ion 131, 132 
oxalate complex 101, 113 
oxide 105, 113 
Ferricyanides.... 59, 101, 145 
Ferrocyanides 59, 101, 145 
Ferrous 
ion.. 131, 132 
sulfide 107 
Fertilizers . 100, 120, 121 
Films, stability of liquid 151 
Fire extinguishers 104 
Flash powders 41 
Fluorine 41. 144 
Foams 146, 151 
Fog 146 
Foods 68 
Formulas 
determination of 26 
meaning 26 
writing 40, 52 
Free energy .-128 
Freezing point and molecular weights 
37, 71, 77, 147 
Fuels 41, 65 
Fuel value of foods 68 
Fuller's earth..... 148 
Fused salts 79 
Galvanized iron 127 
Gases 15 
adsorption 148 
change in pressure or volume 
during reactions 32 
volume per md 33 
Gasolene 41, 121 
Gelatin 148, 151, 152 
Germanium 145 
Glass 145 
Glue 147, 152 
Gold . 39 
colloidal 150 
solution of 130 
Graham 147 
Gram-atoms 27 
Gram-molecules -27, 33 
Graphite... 142, 149, 153 
Gravity battery 152 
Halogens 41 
Hard water 87 
Heat of reaction ...8, 65, 76, 117 
Helium 148 
Heterogeneous systems 9, 146 
Homogeneous systems .....9, 146 
Hydrides 143 
Hydrochloric acid .-..40, 77, 83, 96 
Hydrocyanic acid .. 93, 111 
Hydrofluoric acid 96 
Hydrogen 12 
deposition on metals 129 
formula 32 
stability of compounds of 40 
Hydrogen ion 74, 79, 90, 93, 99 
Hydrogen peroxide 25, 36, 82, 84 132 
Hydrogen sulfide 78, 93 
reducing power 131 
stability 144 
Hydrolysis 101, 143 
Hydrophilic colloids 152 
Hydrophobic colloids 152 
Hydrosulfide ion 93 
Hydroxides 45 
solution by hydrogen ion 105 
Hydroxyl 45, 79, 99 
Hypochlorous acid....57, 78, 124, 132, 143 
158 Iodate ion . 131 
Iodide ion 130, 144 
Iodine 41, 130, 144 
Ionic theory 71 
Ionization 71, 100 
constants 92 
degree of 77 
Ions, 
charges on 72 
complex 101, 110, 130, 144 
migration of. 72 
Indicators 45, 46, 99, 100, 104 
Iron 11, 39 
oxides... .....105, 113 
reducing power 127 
rust ..113 
sulfide 107 
Jellies 152 
Kerosene 41 
Kinetic theory 15, 81, 89, 146 
Lakes 150 
Law of conservation of mass .24 
Law of mass action.. ..90 
Law of simple multiple proportions....24 
Lead 39, 127, 130, 142 
acetate 78, 109 
carbonate 109 
hydroxide 106, 143 
oxides 11, 24, 131, 133 
sulfate.. 99, 109, 110, 133 
sulfide 107, 133 
LeChatelier's Theorem 90, 116 
Limestone 35, 43, 87, 108 
Litharge 11 
Lithium 145 
Litmus ; 45, 99 
Magnesium 41, 145 
chloride 144 
hydroxide 98, 105, 144 
sulfate 144 
Manganese 140 
dioxide 84 
ion 131 
sulfide 107 
Marble 7 
Mass law 90 
Mayonnaise 151 
Melting points of the elements .....140 
Mendeleeff 136 
Mercuric 
chloride 78, 79, 97 
cyanide .78 
sulfide 107, 133 
sulfocyanate .78 
Mercurous 
chloride 99 
ion 127 
Mercury 1 127 
Metals 12 
noble and base... 39, 141 
reducing power 127 
solution in acids 83, 128 
Metals and non-metals 12, 14, 39, 98 
Metathesis 42 
Methyl orange 45, 99, 104 
Methyl violet.... 45, 99 
Meyer 136 
Migration 
of colloidal particles , 148 
of ions , 72 
Mol 27, 33 
Molal concentration i 60 
Molecular weights 34, 36 
Molecules ...15, 26, 148 
Molybdenum 140, 144 
chloride .....144 
fluoride 144 
oxide 144 
Mordants 150 
Neutralization 46, 60 
Newlands 136 
Nickel 39, 140 
complex ions 101, 145 
hydroxides 133 
sulfide 107 
Nitrates 99 
Nitric acid 45 
oxidation by 130, 131, 132, 133 
preparation 96, 121 
strength 77, 143 
synthesis 121 
volatility 96 
Nitric oxide. 82, 117, 118, 120, 131 
Nitrogen 11 
Nitrous acid 56, 78, 132, 143 
Noble metals 39, 141 
Nomenclature 55 
Non-metals 13, 39, 141 
Normal concentration ....61 
Oxalates 101, 113 
Oxalic acid 45 
Oxidation 41, 53, 123 
Oxides 40, 42 
solution by acids 105 
Oxidizing power .....127, 130 
Oxygen 11 
electronegative character.. 39, 144 
formula 31 
standard for atomic weights.... ...26 
Oxy-acids 56 
Ozone r 82 
Partial pressure in gas mixtures 20 
Pepsin 148 
Peptonization 153 
Periodic system 136 
Permanganate ion 131, 133 
Permutite - 94 
Peroxides 51 
Perrin 147 
Petroleum . 121 
Phenol 76, 102 
Phenolphthalein 46, 99, 104 
Phosphates- 99, 109 
Phosphine 144 
Phosphoric acid 45, 93, 143 
Phosphorus acid 58, 104 
Phosphorus 142 
trichloride 104 
valence 141 
Platinum, 
complex ions 101, 130, 145 
depostion of hydrogen on..... 129 
solution by aqua regia 130 
Plumbate ion 110 
Polybasic acids 78 
Polysulfide ions ..107 
159 Positive and negative character of 
elements ..40, 46, 98 
Potassium 39 
chlorate 57, 76, 84 
chloride 143 
hydroxide 45, 143 
oxide 43 
Pressure, 
change during reactions 32 
effect of temperature on 18 
effect on equilibrium 116 
of gases 15 
partial 20 
Properties 8, 75 
Protective colloids 152 
Pure substances 9 
Purple of Cassius 150 
Pyro-acids 58 
Radicals .44, 52, 57 
Radioactivity 145 
Red lead , 11 
Reducing power 127, 130 
Reduction 53, 123 
of metals by carbon 42 
Replacement ; 41 
Salts ...46 
fused . 79 • 
nomenclature . 56 
solubilities 97 
solubility and temperature .118 
weak 78 
Silicates 145 
Silicic acid 58, 78, 143, 149 
Silicon . 145 
Silver 39 
ammonia complex 101, 110, 145 
bromide 110 
carbonate 98, 111 
chloride 97, 99, 110, 149 
cleaning 128 
cyanide 110 
cyanide complex .....101, 110, 145 
iodide. 97, 110 
oxidation 127 
sulfate 99, 112 
sulfide 107, 114, 128, 133 
thiosulfate 110 
Simple multiple proportions, law of....24 
Slaked lime 7 
Soap. 46, 104, 151, 153 
Soddy 139 
Sodium 40, 79, 127, 142 
acid carbonate 91, 103, 109, 113 
borate 103, 108 
carbonate 45, 103, 109, 113 
hydroxide 45, 79, 113 
peroxide 51 
phenolate 102 
silicate 103 
thiosulfate ("hyposulfite") 51 
Solid solutions 10 
Solubility 97, 144 
Solutions 9 
colloidal 147 
freezing point lowering in 
37, 71, 77, 147 
ionization in 71 
Solvay process ..109 
Speed of reactions 81, 118, 132 
Spontaneous combustion 82' 
Stability of compounds 39, 143, 144 
Standard solutions . 61 
Stannate ion 132 
Stannic 
acid... 58, 106, 150 
chloride 46 
compounds 50 
ion 131 
sulfide 107 
Stannite ion 132 
Stannous compounds, 
compounds v - 50 
hydroxide 106 
ion 131, 150 
sulfide 107 
Steel 10, 82 
Stibine 144 
Stirring, effect on reaction 
velocity 83 
Storage batteries 133 
Strontium, 
chloride 144 
hydroxide „ 144 
sulfate 144 
Substances 7, 8, 72, 73 
Sugar 85, 69, 71, 147 
Sulfides ..40 
solubility in acids 106, 133 
solubility in sulfide ion 107, 143 
solubility in water 99 
Sulfo-acids 57, 107 
Sulfur dioxide 
43, 47, 89, 91, 116, 117, 118, 131 
Sulfur trioxide 
........43, 47, 89, 91, 116, 117, 118, 144 
Sulfuric acid 45, 47 
oxidizing power 131 
manufacture 85, 118 
volatility 96 
Superphosphate 100 
Surface, 
effect on adsorption 148 
effect on reaction velocity. 83 
Suspensions 146 
Symbols 12 
Synthesis 11 
Tellurium 142 
Temperature, 
absolute scale 19 
effect on chemical equilibrium 117 
effect on pressure and volume of 
gases 16 
effect on speed of reactions 81 
Tensile strength 140 
Thiosulfates 51, 110 
Thomsen 138 
Thorium hydroxide 143 
Tin (see also stannic stannous).... 
1 39, 131, 142 
hydroxides 143 
Titanic acid 143 
Titration 60 
Toluene 121 
Trinitrobenzene 100 
Tungsten 140 
Ultramicroscope 148 
Uranium 84, 140 
Valence 49, 55, 74, 123, 140, 149 
160 Vanadium 140 
Vapor pressure 21, 36 
Velocity of gas molecules 15 
Velocity of reactions 81, 118, 132 
Volatility 1 96 
Volume of gases 17, 31 
change during reactions 32 
of a given weight.... 33 
of one mol 33 
Volumes, atomic 136 
Washing soda 45, 103 
W ater, 
composition 11 
dielectric constant 79 
formula 24, 31 
hardness 87, 94, 108 
ionization 79, 99 
purification 150 
stability 144 
Weak acids and bases 77, 108 
Weak salts ..78, 79, 109 
Weight of a given volume of gas 33 
Weight relations in chemical re- 
actions 29 
White lead ....: 109 
Wood alcohol 79 
Yeast 85 
Zinc 39 
carbonate 110, 112 
chloride 112 
complex ions 101, 110, 145 
hydroxide 106, 112 
reducing power 127 
solution by acids and alkalis 129 
sulfate 112 
sulfide ... 106, 110 
use in cleaning silver 128