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Title: An Elementary Study of Chemistry

Author: William McPherson
William Edwards Henderson

Release date: March 18, 2007 [eBook #20848]

Language: English

Other information and formats: www.gutenberg.org/ebooks/20848

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*** START OF THE PROJECT GUTENBERG EBOOK AN ELEMENTARY
STUDY OF CHEMISTRY ***

ANTOINE LAURENT LAVOISIER

Famous for his care in quantitative
experiments, for demonstrating the true
nature of combustion, for introducing system
into the naming and grouping of chemical
substances. Executed (1794) during the French
Revolution because of his connection with the
government

This picture is taken from a French engraving
of 1799. The panel represents Lavoisier as he
is being arrested in his laboratory by the
Revolutionary Committee

AN ELEMENTARY

STUDY OF CHEMISTRY
BY

WILLIAM McPHERSON, PH.D.
PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY

AND

WILLIAM EDWARDS HENDERSON, PH.D.
ASSOCIATE PROFESSOR OF CHEMISTRY, OHIO STATE UNIVERSITY

REVISED EDITION

GINN & COMPANY
BOSTON * NEW YORK * CHICAGO * LONDON
COPYRIGHT, 1905, 1906, BY
WILLIAM MCPHERSON AND WILLIAM E. HENDERSON

ALL RIGHTS RESERVED

The Athenæum Press
GINN & COMPANY * PROPRIETORS * BOSTON * U.S.A.

Transcriber's note: Minor typos have been corrected.

PREFACE
In offering this book to teachers of elementary chemistry the authors lay no claim
to any great originality. It has been their aim to prepare a text-book constructed
along lines which have become recognized as best suited to an elementary
treatment of the subject. At the same time they have made a consistent effort to
make the text clear in outline, simple in style and language, conservatively modern
in point of view, and thoroughly teachable.
The question as to what shall be included in an elementary text on chemistry is
perhaps the most perplexing one which an author must answer. While an
enthusiastic chemist with a broad understanding of the science is very apt to go
beyond the capacity of the elementary student, the authors of this text, after an
experience of many years, cannot help believing that the tendency has been rather
in the other direction. In many texts no mention at all is made of fundamental laws
of chemical action because their complete presentation is quite beyond the
comprehension of the student, whereas in many cases it is possible to present the
essential features of these laws in a way that will be of real assistance in the
understanding of the science. For example, it is a difficult matter to deduce the law
of mass action in any very simple way; yet the elementary student can readily
comprehend that reactions are reversible, and that the point of equilibrium depends
upon, rather simple conditions. The authors believe that it is worth while to present
such principles in even an elementary and partial manner because they are of great
assistance to the general student, and because they make a foundation upon which
the student who continues his studies to more advanced courses can securely build.
The authors have no apologies to make for the extent to which they have made use
of the theory of electrolytic dissociation. It is inevitable that in any rapidly
developing science there will be differences of opinion in regard to the value of
certain theories. There can be no question, however, that the outline of the theory
of dissociation here presented is in accord with the views of the very great
majority of the chemists of the present time. Moreover, its introduction to the
extent to which the authors have presented it simplifies rather than increases the
difficulties with which the development of the principles of the science is attended.
The oxygen standard for atomic weights has been adopted throughout the text. The
International Committee, to which is assigned the duty of yearly reporting a
revised list of the atomic weights of the elements, has adopted this standard for
their report, and there is no longer any authority for the older hydrogen standard.

The authors do not believe that the adoption of the oxygen standard introduces any
real difficulties in making perfectly clear the methods by which atomic weights are
calculated.
The problems appended to the various chapters have been chosen with a view not
only of fixing the principles developed in the text in the mind of the student, but
also of enabling him to answer such questions as arise in his laboratory work. They
are, therefore, more or less practical in character. It is not necessary that all of
them should be solved, though with few exceptions the lists are not long. The
answers to the questions are not directly given in the text as a rule, but can be
inferred from the statements made. They therefore require independent thought on
the part of the student.
With very few exceptions only such experiments are included in the text as cannot
be easily carried out by the student. It is expected that these will be performed by
the teacher at the lecture table. Directions for laboratory work by the student are
published in a separate volume.
While the authors believe that the most important function of the elementary text
is to develop the principles of the science, they recognize the importance of some
discussion of the practical application of these principles to our everyday life.
Considerable space is therefore devoted to this phase of chemistry. The teacher
should supplement this discussion whenever possible by having the class visit
different factories where chemical processes are employed.
Although this text is now for the first time offered to teachers of elementary
chemistry, it has nevertheless been used by a number of teachers during the past
three years. The present edition has been largely rewritten in the light of the
criticisms offered, and we desire to express our thanks to the many teachers who
have helped us in this respect, especially to Dr. William Lloyd Evans of this
laboratory, a teacher of wide experience, for his continued interest and helpfulness.
We also very cordially solicit correspondence with teachers who may find
difficulties or inaccuracies in the text.
The authors wish to make acknowledgments for the photographs and engravings
of eminent chemists from which the cuts included in the text were taken; to
Messrs. Elliott and Fry, London, England, for that of Ramsay; to The Macmillan
Company for those of Davy and Dalton, taken from the Century Science Series; to
the L. E. Knott Apparatus Company, Boston, for that of Bunsen.
T
HE AUTHORS

OHIO STATE UNIVERSITY
COLUMBUS, OHIO

CONTENTS
CHAPTER PAGE
I. INTRODUCTION 1
II. OXYGEN 13
III. HYDROGEN 28
IV. WATER AND HYDROGEN DIOXIDE 40
V. THE ATOMIC THEORY 59
VI. CHEMICAL EQUATIONS AND CALCULATIONS 68
VII. NITROGEN AND THE RARE ELEMENTS IN THE
78
ATMOSPHERE
VIII. THE ATMOSPHERE 83
IX. SOLUTIONS 94
X. ACIDS, BASES, AND SALTS; NEUTRALIZATION 106
XI. VALENCE 116
XII. COMPOUNDS OF NITROGEN 122
XIII. REVERSIBLE REACTIONS AND CHEMICAL
137
EQUILIBRIUM
XIV. SULPHUR AND ITS COMPOUNDS 143
XV. PERIODIC LAW 165
XVI. THE CHLORINE FAMILY 174
XVII. CARBON AND SOME OF ITS SIMPLER
196
COMPOUNDS
XVIII. FLAMES,—ILLUMINANTS 213
XIX. MOLECULAR WEIGHTS, ATOMIC WEIGHTS,
223
FORMULAS
XX. THE PHOSPHORUS FAMILY 238
XXI. SILICON, TITANIUM, BORON 257
XXII. THE METALS 267
XXIII. THE ALKALI METALS 274
XXIV. THE ALKALINE-EARTH FAMILY 300
XXV. THE MAGNESIUM FAMILY 316
XXVI. THE ALUMINIUM FAMILY 327
XXVII. THE IRON FAMILY 338
XXVIII. COPPER, MERCURY, AND SILVER 356

XXIX. TIN AND LEAD 370
XXX. MANGANESE AND CHROMIUM 379
XXXI. GOLD AND THE PLATINUM FAMILY 390
XXXII. SOME SIMPLE ORGANIC COMPOUNDS 397
INDEX 421
Facing back
APPENDIX A
cover
Inside back
APPENDIX B
cover

LIST OF FULL-PAGE ILLUSTRATIONS
PAGE
ANTOINE LAURENT LAVOISIER Frontispiece
JOSEPH PRIESTLEY 14
JOHN DALTON 60
WILLIAM RAMSAY 82
DMITRI IVANOVITCH MENDELÉEFF 166
HENRI MOISSAN 176
SIR HUMPHRY DAVY 276
ROBERT WILHELM BUNSEN 298

AN ELEMENTARY STUDY OF CHEMISTRY

CHAPTER I
INTRODUCTION

The natural sciences. Before we advance very far in the study of nature, it
becomes evident that the one large study must be divided into a number of more
limited ones for the convenience of the investigator as well as of the student. These
more limited studies are called the natural sciences.
Since the study of nature is divided in this way for mere convenience, and not
because there is any division in nature itself, it often happens that the different
sciences are very intimately related, and a thorough knowledge of any one of them
involves a considerable acquaintance with several others. Thus the botanist must
know something about animals as well as about plants; the student of human
physiology must know something about physics as well as about the parts of the
body.
Intimate relation of chemistry and physics. Physics and chemistry are two
sciences related in this close way, and it is not easy to make a precise distinction
between them. In a general way it may be said that they are both concerned with
inanimate matter rather than with living, and more particularly with the changes
which such matter may be made to undergo. These changes must be considered
more closely before a definition of the two sciences can be given.
Physical changes. One class of changes is not accompanied by an alteration in the
composition of matter. When a lump of coal is broken the pieces do not differ from
the original lump save in size. A rod of iron may be broken into pieces; it may be
magnetized; it may be heated until it glows; it may be melted. In none of these
changes has the composition of the iron been affected. The pieces of iron, the
magnetized iron, the glowing iron, the melted iron, are just as truly iron as was the
original rod. Sugar may be dissolved in water, but neither the sugar nor the water
is changed in composition. The resulting liquid has the sweet taste of sugar;
moreover the water may be evaporated by heating and the sugar recovered
unchanged. Such changes are called physical changes.
DEFINITION: Physical changes are those which do not involve a change in the
composition of the matter.

Chemical changes. Matter may undergo other changes in which its composition is
altered. When a lump of coal is burned ashes and invisible gases are formed which
are entirely different in composition and properties from the original coal. A rod of
iron when exposed to moist air is gradually changed into rust, which is entirely
different from the original iron. When sugar is heated a black substance is formed
which is neither sweet nor soluble in water. Such changes are evidently quite
different from the physical changes just described, for in them new substances are
formed in place of the ones undergoing change. Changes of this kind are called
chemical changes.
DEFINITION: Chemical changes are those which involve a change in the
composition of the matter.
How to distinguish between physical and chemical changes. It is not always
easy to tell to which class a given change belongs, and many cases will require
careful thought on the part of the student. The test question in all cases is, Has the
composition of the substance been changed? Usually this can be answered by a
study of the properties of the substance before and after the change, since a change
in composition is attended by a change in properties. In some cases, however, only
a trained observer can decide the question.
Changes in physical state. One class of physical changes should be noted with
especial care, since it is likely to prove misleading. It is a familiar fact that ice is
changed into water, and water into steam, by heating. Here we have three different
substances,—the solid ice, the liquid water, and the gaseous steam,—the properties
of which differ widely. The chemist can readily show, however, that these three
bodies have exactly the same composition, being composed of the same substances
in the same proportion. Hence the change from one of these substances into
another is a physical change. Many other substances may, under suitable
conditions, be changed from solids into liquids, or from liquids into gases, without
change in composition. Thus butter and wax will melt when heated; alcohol and
gasoline will evaporate when exposed to the air. The three states—solid, liquid,
and gas—are called the three physical states of matter.
Physical and chemical properties. Many properties of a substance can be noted
without causing the substance to undergo chemical change, and are therefore
called its physical properties. Among these are its physical state, color, odor, taste,
size, shape, weight. Other properties are only discovered when the substance
undergoes chemical change. These are called its chemical properties. Thus we find
that coal burns in air, gunpowder explodes when ignited, milk sours when exposed
to air.

Definition of physics and chemistry. It is now possible to make a general
distinction between physics and chemistry.
DEFINITION: Physics is the science which deals with those changes in matter
which do not involve a change in composition.
DEFINITION: Chemistry is the science which deals with those changes in matter
which do involve a change in composition.
Two factors in all changes. In all the changes which matter can undergo, whether
physical or chemical, two factors must be taken into account, namely, energy and
matter.
Energy. It is a familiar fact that certain bodies have the power to do work. Thus
water falling from a height upon a water wheel turns the wheel and in this way
does the work of the mills. Magnetized iron attracts iron to itself and the motion of
the iron as it moves towards the magnet can be made to do work. When coal is
burned it causes the engine to move and transports the loaded cars from place to
place. When a body has this power to do work it is said to possess energy.
Law of conservation of energy. Careful experiments have shown that when one
body parts with its energy the energy is not destroyed but is transferred to another
body or system of bodies. Just as energy cannot be destroyed, neither can it be
created. If one body gains a certain amount of energy, some other body has lost an
equivalent amount. These facts are summed up in the law of conservation of
energy which may be stated thus: While energy can be changed from one form into
another, it cannot be created or destroyed.
Transformations of energy. Although energy can neither be created nor
destroyed, it is evident that it may assume many different forms. Thus the falling
water may turn the electric generator and produce a current of electricity. The
energy lost by the falling water is thus transformed into the energy of the electric
current. This in turn may be changed into the energy of motion, as when the
current is used for propelling the cars, or into the energy of heat and light, as when
it is used for heating and lighting the cars. Again, the energy of coal may be
converted into energy of heat and subsequently of motion, as when it is used as a
fuel in steam engines.
Since the energy possessed by coal only becomes available when the coal is made
to undergo a chemical change, it is sometimes called chemical energy. It is this
form of energy in which we are especially interested in the study of chemistry.

Matter. Matter may be defined as that which occupies space and possesses weight.
Like energy, matter may be changed oftentimes from one form into another; and
since in these transformations all the other physical properties of a substance save
weight are likely to change, the inquiry arises, Does the weight also change? Much
careful experimenting has shown that it does not. The weight of the products
formed in any change in matter always equals the weight of the substances
undergoing change.
Law of conservation of matter. The important truth just stated is frequently
referred to as the law of conservation of matter, and this law may be briefly stated
thus: Matter can neither be created nor destroyed, though it can be changed from
one form into another.
Classification of matter. At first sight there appears to be no limit to the varieties
of matter of which the world is made. For convenience in study we may classify
all these varieties under three heads, namely, mechanical mixtures, chemical
compounds, and elements.
Mechanical mixtures. If equal bulks of
common salt and iron filings are thoroughly
mixed together, a product is obtained
which, judging by its appearance, is a new
substance. If it is examined more closely,
however, it will be seen to be merely a
mixture of the salt and iron, each of which
substances retains its own peculiar
properties. The mixture tastes just like salt;
the iron particles can be seen and their
gritty character detected. A magnet rubbed
in the mixture draws out the iron just as if
the salt were not there. On the other hand,
the salt can be separated from the iron quite
easily. Thus, if several grams of the mixture
are placed in a test tube, and the tube half
filled with water and thoroughly shaken,
the salt dissolves in the water. The iron
particles can then be filtered from the liquid
Fig. 1
by pouring the entire mixture upon a piece
of filter paper folded so as to fit into the
interior of a funnel (Fig. 1). The paper retains the solid but allows the clear liquid,
known as the filtrate, to drain through. The iron particles left upon the filter paper

will be found to be identical with the original iron. The salt can be recovered from
the filtrate by evaporation of the water. To accomplish this the filtrate is poured
into a small evaporating dish and gently heated (Fig. 2) until the water has
disappeared, or evaporated. The solid left in the dish is identical in every way with
the original salt. Both the iron and the salt have thus been recovered in their
original condition. It is evident that no new substance has been formed by rubbing
the salt and iron together. The product is called a mechanical mixture. Such
mixtures are very common in nature, almost all minerals, sands, and soils being
examples of this class of substances. It is at once apparent that there is no law
regulating the composition of a mechanical mixture, and no two mixtures are
likely to have exactly the same composition. The ingredients of a mechanical
mixture can usually be separated by mechanical means, such as sifting, sorting,
magnetic attraction, or by dissolving one constituent and leaving the other
unchanged.
DEFINITION: A mechanical mixture is one in
which the constituents retain their original
properties, no chemical action having taken
place when they were brought together.
Chemical compounds. If iron filings and
powdered sulphur are thoroughly ground
together in a mortar, a yellowish-green substance
results. It might easily be taken to be a new
body; but as in the case of the iron and salt, the
ingredients can readily be separated. A magnet
draws out the iron. Water does not dissolve the
sulphur, but other liquids do, as, for example, the
liquid called carbon disulphide. When the
mixture is treated with carbon disulphide the
iron is left unchanged, and the sulphur can be
obtained again, after filtering off the iron, by
evaporating the liquid. The substance is,
therefore, a mechanical mixture.
Fig. 2
If now a new portion of the mixture is placed in
a dry test tube and carefully heated in the flame
of a Bunsen burner, as shown in Fig. 3, a striking change takes place. The mixture
begins to glow at some point, the glow rapidly extending throughout the whole
mass. If the test tube is now broken and the product examined, it will be found to
be a hard, black, brittle substance, in no way recalling the iron or the sulphur. The

magnet no longer attracts it; carbon disulphide will not dissolve sulphur from it. It
is a new substance with new properties, resulting from the chemical union of iron
and sulphur, and is called iron sulphide. Such substances are called chemical
compounds, and differ from mechanical mixtures in that the substances producing
them lose their own characteristic properties. We shall see later that the two also
differ in that the composition of a chemical compound never varies.

Fig. 3
DEFINITION: A chemical compound is a substance the constituents of which have
lost their own characteristic properties, and which cannot be separated save by a
chemical change.
Elements. It has been seen that iron sulphide is composed of two entirely different
substances,—iron and sulphur. The question arises, Do these substances in turn
contain other substances, that is, are they also chemical compounds? Chemists
have tried in a great many ways to decompose them, but all their efforts have
failed. Substances which have resisted all efforts to decompose them into other
substances are called elements. It is not always easy to prove that a given
substance is really an element. Some way as yet untried may be successful in
decomposing it into other simpler forms of matter, and the supposed element will
then prove to be a compound. Water, lime, and many other familiar compounds
were at one time thought to be elements.

DEFINITION: An element is a substance which cannot be separated into simpler
substances by any known means.
Kinds of matter. While matter has been grouped in three classes for the purpose
of study, it will be apparent that there are really but two distinct kinds of matter,
namely, compounds and elements. A mechanical mixture is not a third distinct kind
of matter, but is made up of varying quantities of either compounds or elements or
both.
Alchemy. In olden times it was thought that some way could be found to change
one element into another, and a great many efforts were made to accomplish this
transformation. Most of these efforts were directed toward changing the commoner
metals into gold, and many fanciful ways for doing this were described. The
chemists of that time were called alchemists, and the art which they practiced was
called alchemy. The alchemists gradually became convinced that the only way
common metals could be changed into gold was by the wonderful power of a
magic substance which they called the philosopher's stone, which would
accomplish this transformation by its mere touch and would in addition give
perpetual youth to its fortunate possessor. No one has ever found such a stone, and
no one has succeeded in changing one metal into another.
Number of elements. The number of substances now considered to be elements is
not large—about eighty in all. Many of these are rare, and very few of them make
any large fraction of the materials in the earth's crust. Clarke gives the following
estimate of the composition of the earth's crust:

Oxygen 47.0% Calcium 3.5%
Silicon 27.9 Magnesium 2.5
Aluminium 8.1 Sodium 2.7
Iron 4.7 Potassium 2.4
Other elements 1.2%

A complete list of the elements is given in the Appendix. In this list the more
common of the elements are marked with an asterisk. It is not necessary to study
more than a third of the total number of elements to gain a very good knowledge
of chemistry.
Physical state of the elements. About ten of the elements are gases at ordinary
temperatures. Two—mercury and bromine—are liquids. The others are all solids,
though their melting points vary through wide limits, from cæsium which melts at
26° to elements which do not melt save in the intense heat of the electric furnace.

Occurrence of the elements. Comparatively few of the elements occur as
uncombined substances in nature, most of them being found in the form of
chemical compounds. When an element does occur by itself, as is the case with
gold, we say that it occurs in the free state or native; when it is combined with
other substances in the form of compounds, we say that it occurs in the combined
state, or in combination. In the latter case there is usually little about the
compound to suggest that the element is present in it; for we have seen that
elements lose their own peculiar properties when they enter into combination with
other elements. It would never be suspected, for example, that the reddish, earthy-
looking iron ore contains iron.
Names of elements. The names given to the elements have been selected in a great
many different ways. (1) Some names are very old and their original meaning is
obscure. Such names are iron, gold, and copper. (2) Many names indicate some
striking physical property of the element. The name bromine, for example, is
derived from a Greek word meaning a stench, referring to the extremely
unpleasant odor of the substance. The name iodine comes from a word meaning
violet, alluding to the beautiful color of iodine vapor. (3) Some names indicate
prominent chemical properties of the elements. Thus, nitrogen means the producer
of niter, nitrogen being a constituent of niter or saltpeter. Hydrogen means water
former, signifying its presence in water. Argon means lazy or inert, the element
being so named because of its inactivity. (4) Other elements are named from
countries or localities, as germanium and scandium.
Symbols. In indicating the elements found in compounds it is inconvenient to use
such long names, and hence chemists have adopted a system of abbreviations.
These abbreviations are known as symbols, each element having a distinctive
symbol. (1) Sometimes the initial letter of the name will suffice to indicate the
element. Thus I stands for iodine, C for carbon. (2) Usually it is necessary to add
some other characteristic letter to the symbol, since several names may begin with
the same letter. Thus C stands for carbon, Cl for chlorine, Cd for cadmium, Ce for
cerium, Cb for columbium. (3) Sometimes the symbol is an abbreviation of the old
Latin name. In this way Fe (ferrum) indicates iron, Cu (cuprum), copper, Au
(aurum), gold. The symbols are included in the list of elements given in the
Appendix. They will become familiar through constant use.
Chemical affinity the cause of chemical combination. The agency which causes
substances to combine and which holds them together when combined is called
chemical affinity. The experiments described in this chapter, however, show that
heat is often necessary to bring about chemical action. The distinction between the
cause producing chemical action and the circumstances favoring it must be clearly

made. Chemical affinity is always the cause of chemical union. Many agencies
may make it possible for chemical affinity to act by overcoming circumstances
which stand in its way. Among these agencies are heat, light, and electricity. As a
rule, solution also promotes action between two substances. Sometimes these
agencies may overcome chemical attraction and so occasion the decomposition of
a compound.

EXERCISES

1. To what class of changes do the following belong? (a) The melting of ice; (b)
the souring of milk; (c) the burning of a candle; (d) the explosion of gunpowder;
(e) the corrosion of metals. What test question must be applied in each of the
above cases?
2. Give two additional examples (a) of chemical changes; (b) of physical changes.
3. Is a chemical change always accompanied by a physical change? Is a physical
change always accompanied by a chemical change?
4. Give two or more characteristics of a chemical change.
5. (a) When a given weight of water freezes, does it absorb or evolve heat? (b)
When the resulting ice melts, is the total heat change the same or different from
that of freezing?
6. Give three examples of each of the following: (a) mechanical mixtures; (b)
chemical compounds; (c) elements.
7. Give the derivation of the names of the following elements: thorium, gallium,
selenium, uranium. (Consult dictionary.)
8. Give examples of chemical changes which are produced through the agency of
heat; of light; of electricity.

CHAPTER II
OXYGEN

History. The discovery of oxygen is generally attributed to the English chemist
Priestley, who in 1774 obtained the element by heating a compound of mercury
and oxygen, known as red oxide of mercury. It is probable, however, that the
Swedish chemist Scheele had previously obtained it, although an account of his
experiments was not published until 1777. The name oxygen signifies acid former.
It was given to the element by the French chemist Lavoisier, since he believed that
all acids owe their characteristic properties to the presence of oxygen. This view
we now know to be incorrect.
Occurrence. Oxygen is by far the most abundant of all the elements. It occurs
both in the free and in the combined state. In the free state it occurs in the air, 100
volumes of dry air containing about 21 volumes of oxygen. In the combined state
it forms eight ninths of water and nearly one half of the rocks composing the
earth's crust. It is also an important constituent of the compounds which compose
plant and animal tissues; for example, about 66% by weight of the human body is
oxygen.
Preparation. Although oxygen occurs in the free state in the atmosphere, its
separation from the nitrogen and other gases with which it is mixed is such a
difficult matter that in the laboratory it has been found more convenient to prepare
it from its compounds. The most important of the laboratory methods are the
following:
1. Preparation from water. Water is a compound, consisting of 11.18% hydrogen
and 88.82% oxygen. It is easily separated into these constituents by passing an
electric current through it under suitable conditions. The process will be described
in the chapter on water. While this method of preparation is a simple one, it is not
economical.
2. Preparation from mercuric oxide. This method is of interest, since it is the one
which led to the discovery of oxygen. The oxide, which consists of 7.4% oxygen
and 92.6% mercury, is placed in a small, glass test tube and heated. The compound
is in this way decomposed into mercury which collects on the sides of the glass
tube, forming a silvery mirror, and oxygen which, being a gas, escapes from the
tube. The presence of the oxygen is shown by lighting the end of a splint,

extinguishing the flame and bringing the glowing coal into the mouth of the tube.
The oxygen causes the glowing coal to burst into a flame.
In a similar way oxygen may be obtained from its compounds with some of the
other elements. Thus manganese dioxide, a black compound of manganese and
oxygen, when heated to about 700°, loses one third of its oxygen, while barium
dioxide, when heated, loses one half of its oxygen.
3. Preparation from potassium chlorate (usual laboratory method). Potassium
chlorate is a white solid which consists of 31.9% potassium, 28.9% chlorine, and
39.2% oxygen. When heated it undergoes a series of changes in which all the
oxygen is finally set free, leaving a compound of potassium and chlorine called
potassium chloride. The change may be represented as follows:

/ potassium \ (potassium /potassium \ (potassium
{ chlorine } chlorate) = { } chloride) + oxygen
\ oxygen / \ chlorine /

JOSEPH PRIESTLEY (English) (1733-1804)

School-teacher, theologian, philosopher,
scientist; friend of Benjamin Franklin;
discoverer of oxygen; defender of the
phlogiston theory; the first to use mercury in a
pneumatic trough, by which means he first
isolated in gaseous form hydrochloric acid,
sulphur dioxide, and ammonia
The evolution of the oxygen begins at about 400°. It has been found, however, that
if the potassium chlorate is mixed with about one fourth its weight of manganese
dioxide, the oxygen is given off at a much lower temperature. Just how the
manganese dioxide brings about this result is not definitely known. The amount of
oxygen obtained from a given weight of potassium chlorate is exactly the same
whether the manganese dioxide is present or not. So far as can be detected the
manganese dioxide undergoes no change.

Fig. 4
Directions for preparing oxygen. The manner of preparing oxygen from
potassium chlorate is illustrated in the accompanying diagram (Fig. 4). A mixture
consisting of one part of manganese dioxide and four parts of potassium chlorate is
placed in the flask A and gently heated. The oxygen is evolved and escapes
through the tube B. It is collected by bringing over the end of the tube the mouth of
a bottle completely filled with water and inverted in a vessel of water, as shown in

the figure. The gas rises in the bottle and displaces the water. In the preparation of
large quantities of oxygen, a copper retort (Fig. 5) is often substituted for the glass
flask.
In the preparation of oxygen from potassium chlorate
and manganese dioxide, the materials used must be
pure, otherwise a violent explosion may occur. The
purity of the materials is tested by heating a small
amount of the mixture in a test tube.
The collection of gases. The method used for
collecting oxygen illustrates the general method used
Fig. 5 for collecting such gases as are insoluble in water or
nearly so. The vessel C (Fig. 4), containing the water in
which the bottles are inverted, is called a pneumatic trough.
Commercial methods of preparation. Oxygen can now be purchased stored
under great pressure in strong steel cylinders (Fig. 6). It is prepared either by
heating a mixture of potassium chlorate and manganese dioxide, or by separating it
from the nitrogen and other gases with which it is mixed in the atmosphere. The
methods employed for effecting this separation will be described in subsequent
chapters.
Physical properties. Oxygen is a colorless, odorless, tasteless gas, slightly heavier
than air. One liter of it, measured at a temperature of 0° and under a pressure of
one atmosphere, weighs 1.4285 g., while under similar conditions one liter of air
weighs 1.2923 g. It is but slightly soluble in water. Oxygen, like other gases, may
be liquefied by applying very great pressure to the highly cooled gas. When the
pressure is removed the liquid oxygen passes again into the gaseous state, since its
boiling point under ordinary atmospheric pressure is -182.5°.
Chemical properties. At ordinary temperatures oxygen is not very active
chemically. Most substances are either not at all affected by it, or the action is so
slow as to escape notice. At higher temperatures, however, it is very active, and
unites directly with most of the elements. This activity may be shown by heating
various substances until just ignited and then bringing them into vessels of the gas,
when they will burn with great brilliancy. Thus a glowing splint introduced into a
jar of oxygen bursts into flame. Sulphur burns in the air with a very weak flame
and feeble light; in oxygen, however, the flame is increased in size and brightness.
Substances which readily burn in air, such as phosphorus, burn in oxygen with
dazzling brilliancy. Even substances which burn in air with great difficulty, such as
iron, readily burn in oxygen.

The burning of a substance in oxygen is due to the rapid
combination of the substance or of the elements composing it with
the oxygen. Thus, when sulphur burns both the oxygen and sulphur
disappear as such and there is formed a compound of the two,
which is an invisible gas, having the characteristic odor of burning
sulphur. Similarly, phosphorus on burning forms a white solid
compound of phosphorus and oxygen, while iron forms a reddish-
black compound of iron and oxygen.
Oxidation. The term oxidation is applied to the chemical change
which takes place when a substance, or one of its constituent parts,
combines with oxygen. This process may take place rapidly, as in
the burning of phosphorus, or slowly, as in the oxidation (or
rusting) of iron when exposed to the air. It is always accompanied
by the liberation of heat. The amount of heat liberated by the
oxidation of a definite weight of any given substance is always the
same, being entirely independent of the rapidity of the process. If
the oxidation takes place slowly, the heat is generated so slowly
that it is difficult to detect it. If the oxidation takes place rapidly,
however, the heat is generated in such a short interval of time that Fig. 6
the substance may become white hot or burst into a flame.
Combustion; kindling temperature. When oxidation takes place so rapidly that
the heat generated is sufficient to cause the substance to glow or burst into a flame
the process is called combustion. In order that any substance may undergo
combustion, it is necessary that it should be heated to a certain temperature, known
as the kindling temperature. This temperature varies widely for different bodies,
but is always definite for the same body. Thus the kindling temperature of
phosphorus is far lower than that of iron, but is definite for each. When any portion
of a substance is heated until it begins to burn the combustion will continue
without the further application of heat, provided the heat generated by the process
is sufficient to bring other parts of the substance to the kindling temperature. On
the other hand, if the heat generated is not sufficient to maintain the kindling
temperature, combustion ceases.
Oxides. The compounds formed by the oxidation of any element are called oxides.
Thus in the combustion of sulphur, phosphorus, and iron, the compounds formed
are called respectively oxide of sulphur, oxide of phosphorus, and oxide of iron. In
general, then, an oxide is a compound of oxygen with another element. A great
many substances of this class are known; in fact, the oxides of all the common
elements have been prepared, with the exception of those of fluorine and bromine.

Some of these are familiar compounds. Water, for example, is an oxide of
hydrogen, and lime an oxide of the metal calcium.
Products of combustion. The particular oxides formed by the combustion of any
substance are called products of combustion of that substance. Thus oxide of
sulphur is the product of the combustion of sulphur; oxide of iron is the product of
the combustion of iron. It is evident that the products of the combustion of any
substance must weigh more than the original substance, the increase in weight
corresponding to the amount of oxygen taken up in the act of combustion. For
example, when iron burns the oxide of iron formed weighs more than the original
iron.
In some cases the products of combustion are invisible gases, so that the substance
undergoing combustion is apparently destroyed. Thus, when a candle burns it is
consumed, and so far as the eye can judge nothing is formed during combustion.
That invisible gases are formed, however, and that the weight of these is greater
than the weight of the candle may be shown by the following experiment.

Fig. 7
A lamp chimney is filled with sticks of the compound known as sodium hydroxide
(caustic soda), and suspended from the beam of the balance, as shown in Fig. 7. A
piece of candle is placed on the balance pan so that the wick comes just below the
chimney, and the balance is brought to a level by adding weights to the other pan.
The candle is then lighted. The products formed pass up through the chimney and

are absorbed by the sodium hydroxide. Although the candle burns away, the pan
upon which it rests slowly sinks, showing that the combustion is attended by an
increase in weight.
Combustion in air and in oxygen. Combustion in air and in oxygen differs only
in rapidity, the products formed being exactly the same. That the process should
take place less rapidly in the former is readily understood, for the air is only about
one fifth oxygen, the remaining four fifths being inert gases. Not only is less
oxygen available, but much of the heat is absorbed in raising the temperature of
the inert gases surrounding the substance undergoing combustion, and the
temperature reached in the combustion is therefore less.
Phlogiston theory of combustion. The French chemist Lavoisier (1743-1794),
who gave to oxygen its name was the first to show that combustion is due to union
with oxygen. Previous to his time combustion was supposed to be due to the
presence of a substance or principle called phlogiston. One substance was thought
to be more combustible than another because it contained more phlogiston. Coal,
for example, was thought to be very rich in phlogiston. The ashes left after
combustion would not burn because all the phlogiston had escaped. If the
phlogiston could be restored in any way, the substance would then become
combustible again. Although this view seems absurd to us in the light of our
present knowledge, it formerly had general acceptance. The discovery of oxygen
led Lavoisier to investigate the subject, and through his experiments he arrived at
the true explanation of combustion. The discovery of oxygen together with the part
it plays in combustion is generally regarded as the most important discovery in the
history of chemistry. It marked the dawn of a new period in the growth of the
science.
Combustion in the broad sense. According to the definition given above, the
presence of oxygen is necessary for combustion. The term is sometimes used,
however, in a broader sense to designate any chemical change attended by the
evolution of heat and light. Thus iron and sulphur, or hydrogen and chlorine under
certain conditions, will combine so rapidly that light is evolved, and the action is
called a combustion. Whenever combustion takes place in the air, however, the
process is one of oxidation.
Spontaneous combustion. The temperature reached in a given chemical action,
such as oxidation, depends upon the rate at which the reaction takes place. This
rate is usually increased by raising the temperature of the substances taking part in
the action.

When a slow oxidation takes place under such conditions that the heat generated is
not lost by being conducted away, the temperature of the substance undergoing
oxidation is raised, and this in turn hastens the rate of oxidation. The rise in
temperature may continue in this way until the kindling temperature of the
substance is reached, when combustion begins. Combustion occurring in this way
is called spontaneous combustion.
Certain oils, such as the linseed oil used in paints, slowly undergo oxidation at
ordinary temperatures, and not infrequently the origin of fires has been traced to
the spontaneous combustion of oily rags. The spontaneous combustion of hay has
been known to set barns on fire. Heaps of coal have been found to be on fire when
spontaneous combustion offered the only possible explanation.
Importance of oxygen. 1. Oxygen is essential to life. Among living organisms
only certain minute forms of plant life can exist without it. In the process of
respiration the air is taken into the lungs where a certain amount of oxygen is
absorbed by the blood. It is then carried to all parts of the body, oxidizing the
worn-out tissues and changing them into substances which may readily be
eliminated from the body. The heat generated by this oxidation is the source of the
heat of the body. The small amount of oxygen which water dissolves from the air
supports all the varied forms of aquatic animals.
2. Oxygen is also essential to decay. The process of decay is really a kind of
oxidation, but it will only take place in the presence of certain minute forms of life
known as bacteria. Just how these assist in the oxidation is not known. By this
process the dead products of animal and vegetable life which collect on the surface
of the earth are slowly oxidized and so converted into harmless substances. In this
way oxygen acts as a great purifying agent.
3. Oxygen is also used in the treatment of certain diseases in which the patient is
unable to inhale sufficient air to supply the necessary amount of oxygen.

OZONE

Preparation. When electric sparks are passed through oxygen or air a small
percentage of the oxygen is converted into a substance called ozone, which differs
greatly from oxygen in its properties. The same change can also be brought about
by certain chemical processes. Thus, if some pieces of phosphorus are placed in a
bottle and partially covered with water, the presence of ozone may soon be
detected in the air contained in the bottle. The conversion of oxygen into ozone is
attended by a change in volume, 3 volumes of oxygen forming 2 volumes of
ozone. If the resulting ozone is heated to about 300°, the reverse change takes

place, the 2 volumes of ozone being changed back into 3 volumes of oxygen. It is
possible that traces of ozone exist in the atmosphere, although its presence there
has not been definitely proved, the tests formerly used for its detection having
been shown to be unreliable.
Properties. As commonly prepared, ozone is mixed with a large excess of oxygen.
It is possible, however, to separate the ozone and thus obtain it in pure form. The
gas so obtained has the characteristic odor noticed about electrical machines when
in operation. By subjecting it to great pressure and a low temperature, the gas
condenses to a bluish liquid, boiling at -119°. When unmixed with other gases
ozone is very explosive, changing back into oxygen with the liberation of heat. Its
chemical properties are similar to those of oxygen except that it is far more active.
Air or oxygen containing a small amount of ozone is now used in place of oxygen
in certain manufacturing processes.
The difference between oxygen and ozone. Experiments show that in changing
oxygen into ozone no other kind of matter is either added to the oxygen or
withdrawn from it. The question arises then, How can we account for the
difference in their properties? It must be remembered that in all changes we have
to take into account energy as well as matter. By changing the amount of energy in
a substance we change its properties. That oxygen and ozone contain different
amounts of energy may be shown in a number of ways; for example, by the fact
that the conversion of ozone into oxygen is attended by the liberation of heat. The
passage of the electric sparks through oxygen has in some way changed the energy
content of the element and thus it has acquired new properties. Oxygen and ozone
must, therefore, be regarded as identical so far as the kind of matter of which they
are composed is concerned. Their different properties are due to their different
energy contents.
Allotropic states or forms of matter. Other elements besides oxygen may exist in
more than one form. These different forms of the same element are called
allotropic states or forms of the element. These forms differ not only in physical
properties but also in their energy contents. Elements often exist in a variety of
forms which look quite different. These differences may be due to accidental
causes, such as the size or shape of the particles or the way in which the element
was prepared. Only such forms, however, as have different energy contents are
properly called allotropic forms.

MEASUREMENT OF GAS VOLUMES

Standard conditions. It is a well-known fact that the volume occupied by a
definite weight of any gas can be altered by changing the temperature of the gas or
the pressure to which it is subjected. In measuring the volume of gases it is
therefore necessary, for the sake of accuracy, to adopt some standard conditions of
temperature and pressure. The conditions agreed upon are (1) a temperature of 0°,
and (2) a pressure equal to the average pressure exerted by the atmosphere at the
sea level, that is, 1033.3 g. per square centimeter. These conditions of temperature
and pressure are known as the standard conditions, and when the volume of a gas
is given it is understood that the measurement was made under these conditions,
unless it is expressly stated otherwise. For example, the weight of a liter of oxygen
has been given as 1.4285 g. This means that one liter of oxygen, measured at a
temperature of 0° and under a pressure of 1033.3 g. per square centimeter, weighs
1.4285 g.
The conditions which prevail in the laboratory are never the standard conditions. It
becomes necessary, therefore, to find a way to calculate the volume which a gas
will occupy under standard conditions from the volume which it occupies under
any other conditions. This may be done in accordance with the following laws.
Law of Charles. This law expresses the effect which a change in the temperature
of a gas has upon its volume. It may be stated as follows: For every degree the
temperature of a gas rises above zero the volume of the gas is increased by 1/273
of the volume which it occupies at zero; likewise for every degree the temperature
of the gas falls below zero the volume of the gas is decreased by 1/273 of the
volume which it occupies at zero, provided in both cases that the pressure to which
the gas is subjected remains constant.
If V represents the volume of gas at 0°, then the volume at 1° will be V + 1/273 V;
at 2° it will be V + 2/273 V; or, in general, the volume v, at the temperature t, will
be expressed by the formula

(1) v = V + t/273 V,

or (2) v = V(1 + (t/273)).

Since 1/273 = 0.00366, the formula may be written

(3) v = V(1 + 0.00366t).

Since the value of V (volume under standard conditions) is the one usually sought,
it is convenient to transpose the equation to the following form:

(4) V = v/(1 + 0.00366t).

The following problem will serve as an illustration of the application of this
equation.
The volume of a gas at 20° is 750 cc.; find the volume it will occupy at 0°, the
pressure remaining constant.
In this case, v = 750 cc. and t = 20. By substituting these values, equation (4)
becomes

V = 750/(1 + 0.00366 × 20) = 698.9 cc.

Law of Boyle. This law expresses the relation between the volume occupied by a
gas and the pressure to which it is subjected. It may be stated as follows: The
volume of a gas is inversely proportional to the pressure under which it is
measured, provided the temperature of the gas remains constant.
If V represents the volume when subjected to a pressure P and v represents its
volume when the pressure is changed to p, then, in accordance with the above law,
V : v :: p : P, or VP = vp. In other words, for a given weight of a gas the product of
the numbers representing its volume and the pressure to which it is subjected is a
constant.
Since the pressure of the atmosphere at any point is indicated by the barometric
reading, it is convenient in the solution of the problems to substitute the latter for
the pressure measured in grams per square centimeter. The average reading of the
barometer at the sea level is 760 mm., which corresponds to a pressure of 1033.3
g. per square centimeter. The following problem will serve as an illustration of the
application of Boyle's law.
A gas occupies a volume of 500 cc. in a laboratory where the barometric reading is
740 mm. What volume would it occupy if the atmospheric pressure changed so
that the reading became 750 mm.?
Substituting the values in the equation VP = vp, we have 500 × 740 = v × 750, or v
= 493.3 cc.
Variations in the volume of a gas due to changes both in temperature and
pressure. Inasmuch as corrections must be made as a rule for both temperature
and pressure, it is convenient to combine the equations given above for the
corrections for each, so that the two corrections may be made in one operation.
The following equation is thus obtained:

(5) Vs = vp/(760(1 + 0.00366t)),

in which Vs represents the volume of a gas under standard conditions and v, p, and
t the volume, pressure, and temperature respectively at which the gas was actually
measured.
The following problem will serve to illustrate the application of this equation.
A gas having a temperature of 20° occupies a volume of 500 cc. when subjected to
a pressure indicated by a barometric reading of 740 mm. What volume would this
gas occupy under standard conditions?
In this problem v = 500, p = 740, and t = 20. Substituting these values in the above
equation, we get

Vs = (500 × 740)/(760 (1 + 0.00366 × 20)) = 453.6 cc.

Variations in the volume of a gas due to the pressure of
aqueous vapor. In many cases gases are collected over water,
as explained under the preparation of oxygen. In such cases
there is present in the gas a certain amount of water vapor.
This vapor exerts a definite pressure, which acts in opposition
to the atmospheric pressure and which therefore must be
subtracted from the latter in determining the effective pressure
upon the gas. Thus, suppose we wish to determine the
pressure to which the gas in tube A (Fig. 8) is subjected. The
tube is raised or lowered until the level of the water inside and
outside the tube is the same. The atmosphere presses down
upon the surface of the water (as indicated by the arrows),
thus forcing the water upward within the tube with a pressure
equal to the atmospheric pressure. The full force of this
upward pressure, however, is not spent in compressing the gas
within the tube, for since it is collected over water it contains
a certain amount of water vapor. This water vapor exerts a
pressure (as indicated by the arrow within the tube) in
opposition to the upward pressure. It is plain, therefore, that
Fig. 8
the effective pressure upon the gas is equal to the atmospheric
pressure less the pressure exerted by the aqueous vapor. The
pressure exerted by the aqueous vapor increases with the temperature. The figures
representing the extent of this pressure (often called the tension of aqueous vapor)
are given in the Appendix. They express the pressure or tension in millimeters of

mercury, just as the atmospheric pressure is expressed in millimeters of mercury.
Representing the pressure of the aqueous vapor by a, formula (5) becomes

(6) Vs = v(p - a)/(760(1 + 0.00366t)).

The following problem will serve to illustrate the method of applying the
correction for the pressure of the aqueous vapor.
The volume of a gas measured over water in a laboratory where the temperature is
20° and the barometric reading is 740 mm. is 500 cc. What volume would this
occupy under standard conditions?
The pressure exerted by the aqueous vapor at 20° (see table in Appendix) is equal
to the pressure exerted by a column of mercury 17.4 mm. in height. Substituting
the values of v, t, p, and a in formula (6), we have

(6) Vs = 500(740 - 17.4)/(760(1 + 0.00366 × 20)) = 442.9 cc.

Adjustment of tubes before reading gas volumes. In measuring the volumes of
gases collected in graduated tubes or other receivers, over a liquid as illustrated in
Fig. 8, the reading should be taken after raising or lowering the tube containing the
gas until the level of the liquid inside and outside the tube is the same; for it is only
under these conditions that the upward pressure within the tube is the same as the
atmospheric pressure.

EXERCISES

1. What is the meaning of the following words? phlogiston, ozone, phosphorus.
(Consult dictionary.)
2. Can combustion take place without the emission of light?
3. Is the evolution of light always produced by combustion?
4. (a) What weight of oxygen can be obtained from 100 g. of water? (b) What
volume would this occupy under standard conditions?
5. (a) What weight of oxygen can be obtained from 500g. of mercuric oxide? (b)
What volume would this occupy under standard conditions?
6. What weight of each of the following compounds is necessary to prepare 50 l. of
oxygen? (a) water; (b) mercuric oxide; (c) potassium chlorate.

7. Reduce the following volumes to 0°, the pressure remaining constant: (a) 150
cc. at 10°; (b) 840 cc. at 273°.
8. A certain volume of gas is measured when the temperature is 20°. At what
temperature will its volume be doubled?
9. Reduce the following volumes to standard conditions of pressure, the
temperature remaining constant: (a) 200 cc. at 740 mm.; (b) 500 l. at 380 mm.
10. What is the weight of 1 l. of oxygen when the pressure is 750 mm. and the
temperature 0°?
11. Reduce the following volumes to standard conditions of temperature and
pressure: (a) 340 cc. at 12° and 753 mm; (b) 500 cc. at 15° and 740 mm.
12. What weight of potassium chlorate is necessary to prepare 250 l. of oxygen at
20° and 750 mm.?
13. Assuming the cost of potassium chlorate and mercuric oxide to be respectively
$0.50 and $1.50 per kilogram, calculate the cost of materials necessary for the
preparation of 50 l. of oxygen from each of the above compounds.
14. 100 g. of potassium chlorate and 25 g. of manganese dioxide were heated in
the preparation of oxygen. What products were left in the flask, and how much of
each was present?

CHAPTER III
HYDROGEN

Historical. The element hydrogen was first clearly recognized as a distinct
substance by the English investigator Cavendish, who in 1766 obtained it in a pure
state, and showed it to be different from the other inflammable airs or gases which
had long been known. Lavoisier gave it the name hydrogen, signifying water
former, since it had been found to be a constituent of water.
Occurrence. In the free state hydrogen is found in the atmosphere, but only in
traces. In the combined state it is widely distributed, being a constituent of water
as well as of all living organisms, and the products derived from them, such as
starch and sugar. About 10% of the human body is hydrogen. Combined with
carbon, it forms the substances which constitute petroleum and natural gas.
It is an interesting fact that while hydrogen in the free state occurs only in traces
on the earth, it occurs in enormous quantities in the gaseous matter surrounding the
sun and certain other stars.
Preparation from water. Hydrogen can be prepared from water by several
methods, the most important of which are the following.
1. By the electric current. As has been indicated in the preparation of oxygen,
water is easily separated into its constituents, hydrogen and oxygen, by passing an
electric current through it under certain conditions.
2. By the action of certain metals. When brought into contact with certain metals
under appropriate conditions, water gives up a portion or the whole of its
hydrogen, its place being taken by the metal. In the case of a few of the metals this
change occurs at ordinary temperatures. Thus, if a bit of sodium is thrown on
water, an action is seen to take place at once, sufficient heat being generated to
melt the sodium, which runs about on the surface of the water. The change which
takes place consists in the displacement of one half of the hydrogen of the water
by the sodium, and may be represented as follows:

_ _ _ _
| hydrogen | | sodium |
sodium + | hydrogen |(water) = | hydrogen |(sodium hydroxide) +

hydrogen
|_oxygen _| |_oxygen _|

The sodium hydroxide formed is a white solid which remains dissolved in the
undecomposed water, and may be obtained by evaporating the solution to dryness.
The hydrogen is evolved as a gas and may be collected by suitable apparatus.
Other metals, such as magnesium and iron, decompose water rapidly, but only at
higher temperatures. When steam is passed over hot iron, for example, the iron
combines with the oxygen of the steam, thus displacing the hydrogen. Experiments
show that the change may be represented as follows:

_ _
| hydrogen | _ _ _ _
iron + | hydrogen |(water) = | iron |(iron oxide) + | hydrogen |
|_oxygen _| |_oxygen _| |_hydrogen_|

The iron oxide formed is a reddish-black compound, identical with that obtained
by the combustion of iron in oxygen.
Directions for preparing hydrogen by the action of steam on iron. The
apparatus used in the preparation of hydrogen from iron and steam is shown in
Fig. 9. A porcelain or iron tube B, about 50 cm. in length and 2 cm. or 3 cm. in
diameter, is partially filled with fine iron wire or tacks and connected as shown in
the figure. The tube B is heated, slowly at first, until the iron is red-hot. Steam is
then conducted through the tube by boiling the water in the flask A. The hot iron
combines with the oxygen in the steam, setting free the hydrogen, which is
collected over water. The gas which first passes over is mixed with the air
previously contained in the flask and tube, and is allowed to escape, since a
mixture of hydrogen with oxygen or air explodes violently when brought in contact
with a flame. It is evident that the flask A must be disconnected from the tube
before the heat is withdrawn.
That the gas obtained is different from air and oxygen may be shown by holding a
bottle of it mouth downward and bringing a lighted splint into it. The hydrogen is
ignited and burns with an almost colorless flame.

Fig. 9
Preparation from acids (usual laboratory method). While hydrogen can be
prepared from water, either by the action of the electric current or by the action of
certain metals, these methods are not economical and are therefore but little used.
In the laboratory hydrogen is generally prepared from compounds known as acids,
all of which contain hydrogen. When acids are brought in contact with certain
metals, the metals dissolve and set free the hydrogen of the acid. Although this
reaction is a quite general one, it has been found most convenient in preparing
hydrogen by this method to use either zinc or iron as the metal and either
hydrochloric or sulphuric acid as the acid. Hydrochloric acid is a compound
consisting of 2.77% hydrogen and 97.23% chlorine, while sulphuric acid consists
of 2.05% hydrogen, 32.70% sulphur, and 65.25% oxygen.
The changes which take place in the preparation of hydrogen from zinc and
sulphuric acid (diluted with water) may be represented as follows:

_ _ _ _
| hydrogen |(sulphuric | zinc |(zinc
zinc + | sulphur | acid) = | sulphur | sulphate) + hydrogen
|_oxygen _| |_oxygen _|

In other words, the zinc has taken the place of the hydrogen in sulphuric acid. The
resulting compound contains zinc, sulphur, and oxygen, and is known as zinc
sulphate. This remains dissolved in the water present in the acid. It may be
obtained in the form of a white solid by evaporating the liquid left after the metal
has passed into solution.
When zinc and hydrochloric acid are used the following changes take place:

_ _ _ _

| hydrogen |(hydrochloric | zinc |(zinc
zinc + |_chlorine_| acid) = |_chlorine_| chloride) + hydrogen

When iron is used the changes which take place are exactly similar to those just
given for zinc.

Fig. 10.
Directions for preparing hydrogen from acids. The preparation of hydrogen
from acids is carried out in the laboratory as follows: The metal is placed in a flask
or wide-mouthed bottle A (Fig. 10) and the acid is added slowly through the funnel
tube B. The metal dissolves in the acid, while the hydrogen which is liberated
escapes through the exit tube C and is collected over water. It is evident that the
hydrogen which passes over first is mixed with the air from the bottle A. Hence
care must be taken not to bring a flame near the exit tube, since, as has been stated
previously, such a mixture explodes with great violence when brought in contact
with a flame.
Precautions. Both sulphuric acid and zinc, if impure, are likely to contain small
amounts of arsenic. Such materials should not be used in preparing hydrogen,
since the arsenic present combines with a portion of the hydrogen to form a very
poisonous gas known as arsine. On the other hand, chemically pure sulphuric acid,
i.e. sulphuric acid that is entirely free from impurities, will not act upon chemically
pure zinc. The reaction may be started, however, by the addition of a few drops of
a solution of copper sulphate or platinum tetrachloride.
Physical properties. Hydrogen is similar to oxygen in that it is a colorless,
tasteless, odorless gas. It is characterized by its extreme lightness, being the
lightest of all known substances. One liter of the gas weighs only 0.08984 g. On
comparing this weight with that of an equal volume of oxygen, viz., 1.4285 g., the
latter is found to be 15.88 times as heavy as hydrogen. Similarly, air is found to be

14.38 times as heavy as hydrogen. Soap bubbles blown with hydrogen rapidly rise
in the air. On account of its lightness it is possible to pour it upward from one
bottle into another. Thus, if the bottle A (Fig. 11) is filled with hydrogen, placed
mouth downward by the side of bottle B, filled with air, and is then gradually
inverted under B as indicated in the figure, the hydrogen will flow upward into
bottle B, displacing the air. Its presence in bottle B may then be shown by bringing
a lighted splint to the mouth of the bottle, when the hydrogen will be ignited by the
flame. It is evident, from this experiment, that in order to retain the gas in an open
bottle the bottle must be placed mouth downward.
Hydrogen is far more difficult to liquefy than any other
gas, with the exception of helium, a rare element
recently found to exist in the atmosphere. The English
scientist Dewar, however, in 1898 succeeded not only
in obtaining hydrogen in liquid state but also as a solid.
Liquid hydrogen is colorless and has a density of only
0.07. Its boiling point under atmospheric pressure is
-252°. Under diminished pressure the temperature has
Fig. 11 been reduced to -262°. The solubility of hydrogen in
water is very slight, being still less than that of oxygen.
Pure hydrogen produces no injurious results when inhaled. Of course one could
not live in an atmosphere of the gas, since oxygen is essential to respiration.
Chemical properties. At ordinary temperatures hydrogen is not an active element.
A mixture of hydrogen and chlorine, however, will combine with explosive
violence at ordinary temperature if exposed to the sunlight. The union can be
brought about also by heating. The product formed in either case is hydrochloric
acid. Under suitable conditions hydrogen combines with nitrogen to form
ammonia, and with sulphur to form the foul-smelling gas, hydrogen sulphide. The
affinity of hydrogen for oxygen is so great that a mixture of hydrogen and oxygen
or hydrogen and air explodes with great violence when heated to the kindling
temperature (about 612°). Nevertheless under proper conditions hydrogen may be
made to burn quietly in either oxygen or air. The resulting hydrogen flame is
almost colorless and is very hot. The combustion of the hydrogen is, of course, due
to its union with oxygen. The product of the combustion is therefore a compound
of hydrogen and oxygen. That this compound is water may be shown easily by
experiment.

Fig. 12
Directions for burning hydrogen in air. The combustion of hydrogen in air may
be carried out safely as follows: The hydrogen is generated in the bottle A (Fig.
12), is dried by conducting it through the tube X, filled with some substance
(generally calcium chloride) which has a great attraction for moisture, and escapes
through the tube T, the end of which is drawn out to a jet. The hydrogen first
liberated mixes with the air contained in the generator. If a flame is brought near
the jet before this mixture has all escaped, a violent and very dangerous explosion
results, since the entire apparatus is filled with the explosive mixture. On the other
hand, if the flame is not applied until all the air has been expelled, the hydrogen is
ignited and burns quietly, since only the small amount of it which escapes from the
jet can come in contact with the oxygen of the air at any one time. By holding a
cold, dry bell jar or bottle over the flame, in the manner shown in the figure, the
steam formed by the combustion of the hydrogen is condensed, the water
collecting in drops on the sides of the jar.
Precautions. In order to avoid danger it is absolutely necessary to prove that the
hydrogen is free from air before igniting it. This can be done by testing small
amounts of the escaping gas. A convenient and safe method of doing this is to fill a
test tube with the gas by inverting it over the jet. The hydrogen, on account of its
lightness, collects in the tube, displacing the air. After holding it over the jet for a
few moments in order that it may be filled with the gas, the tube is gently brought,
mouth downward, to the flame of a burner placed not nearer than an arm's length
from the jet. If the hydrogen is mixed with air a slight explosion occurs, but if pure
it burns quietly in the tube. The operation is repeated until the gas burns quietly,
when the tube is quickly brought back over the jet for an instant, whereby the
escaping hydrogen is ignited by the flame in the tube.
A mixture of hydrogen and oxygen is explosive. That a mixture of hydrogen and
air is explosive may be shown safely as follows: A cork through which passes a
short glass tube about 1 cm. in diameter is fitted air-tight into the tubule of a bell

jar of 2 l. or 3 l. capacity. (A thick glass bottle with
bottom removed may be used.) The tube is closed with a
small rubber stopper and the bell jar filled with
hydrogen, the gas being collected over water. When
entirely filled with the gas the jar is removed from the
water and supported by blocks of wood in order to leave
the bottom of the jar open, as shown in Fig. 13. The
stopper is now removed from the tube in the cork, and
the hydrogen, which on account of its lightness escapes
from the tube, is at once lighted. As the hydrogen
escapes, the air flows in at the bottom of the jar and
mixes with the remaining portion of the hydrogen, so . Fig. 13
that a mixture of the two soon forms, and a loud
explosion results. The explosion is not dangerous, since the bottom of the jar is
open, thus leaving room for the expansion of the hot gas.
Since air is only one fifth oxygen, the remainder being inert gases, it may readily
be inferred that a mixture of hydrogen with pure oxygen would be far more
explosive than a mixture of hydrogen with air. Such mixtures should not be made
except in small quantities and by experienced workers.
Hydrogen does not support combustion. While hydrogen is readily combustible,
it is not a supporter of combustion. In other words, substances will not burn in it.
This may be shown by bringing a lighted candle supported by a stiff wire into a
bottle or cylinder of the pure gas, as shown in Fig. 14. The hydrogen is ignited by
the flame of the candle and burns at the mouth of the bottle, where it comes in
contact with the oxygen in the air. When the candle is thrust up into the gas, its
flame is extinguished on account of the absence of oxygen. If slowly withdrawn,
the candle is relighted as it passes through the layer of burning hydrogen.

Fig. 14

Fig. 15
Reduction. On account of its great affinity for oxygen, hydrogen has the power of
abstracting it from many of its compounds. Thus, if a stream of hydrogen, dried by
passing through the tube B (Fig. 15), filled with calcium chloride, is conducted
through the tube C containing some copper oxide, heated to a moderate
temperature, the hydrogen abstracts the oxygen from the copper oxide. The change
may be represented as follows:

hydrogen + {copper} {hydrogen}

{oxygen}(copper oxide) = {oxygen
}(water) + copper
The water formed collects in the cold portions of the tube C near its end. In this
experiment the copper oxide is said to undergo reduction. Reduction may therefore
be defined as the process of withdrawing oxygen from a compound.
Relation of reduction to oxidation. At the same time that the copper oxide is
reduced it is clear that the hydrogen is oxidized, for it combines with the oxygen
given up by the copper oxide. The two processes are therefore very closely related,
and it usually happens that when one substance is oxidized some other substance is
reduced. That substance which gives up its oxygen is called an oxidizing agent,
while the substance which unites with the oxygen is called a reducing agent.
The oxyhydrogen blowpipe. This is a form of apparatus used for burning
hydrogen in pure oxygen. As has been previously stated, the flame produced by
the combustion of hydrogen in the air is very hot. It is evident that if pure oxygen
is substituted for air, the temperature reached will be much higher, since there are
no inert gases to absorb the heat. The oxyhydrogen blowpipe, used to effect this
combination, consists of a small tube placed within a larger one, as shown in Fig.
16.

Fig. 16
The hydrogen, stored under pressure, generally in steel cylinders, is first passed
through the outer tube and ignited at the open end of the tube. The oxygen from a
similar cylinder is then conducted through the inner tube, and mixes with the
hydrogen at the end of the tube. In order to produce the maximum heat, the
hydrogen and oxygen must be admitted to the blowpipe in the exact proportion in
which they combine, viz., 2 volumes of hydrogen to 1 of oxygen, or by weight, 1
part of hydrogen to 7.94 parts of oxygen. The intensity of the heat may be shown
by bringing into the flame pieces of metal such as iron wire or zinc. These burn
with great brilliancy. Even platinum, having a melting point of 1779°, may be
melted by the heat of the flame.

While the oxyhydrogen flame is intensely hot, it is almost non-luminous. If
directed against some infusible substance like ordinary lime (calcium oxide), the
heat is so intense that the lime becomes incandescent and glows with a brilliant
light. This is sometimes used as a source of light, under the name of Drummond or
lime light.

Fig. 17
The blast lamp. A similar form of apparatus is commonly used in the laboratory
as a source of heat under the name blast lamp (Fig. 17). This differs from the
oxyhydrogen blowpipe only in the size of the tubes. In place of the hydrogen and
oxygen the more accessible coal gas and air are respectively used. The former is
composed largely of a mixture of free hydrogen and gaseous compounds of carbon
and hydrogen. While the temperature of the flame is not so high as that of the
oxyhydrogen blowpipe, it nevertheless suffices for most chemical operations
carried out in the laboratory.
Uses of hydrogen. On account of its cost, hydrogen is but little used for
commercial purposes. It is sometimes used as a material for the inflation of
balloons, but usually the much cheaper coal gas is substituted for it. Even hot air is
often used when the duration of ascension is very short. It has been used also as a
source of heat and light in the oxyhydrogen blowpipe. Where the electric current is
available, however, this form of apparatus has been displaced almost entirely by

the electric light and electric furnace, which are much more economical and more
powerful sources of light and heat.

EXERCISES

1. Will a definite weight of iron decompose an unlimited weight of steam?
2. Why is oxygen passed through the inner tube of the oxyhydrogen blowpipe
rather than the outer?
3. In Fig. 14, will the flame remain at the mouth of the tube?
4. From Fig. 15, suggest a way for determining experimentally the quantity of
water formed in the reaction.
5. Distinguish clearly between the following terms: oxidation, reduction,
combustion, and kindling temperature.
6. Is oxidation always accompanied by reduction?
7. What is the source of heat in the lime light? What is the exact use of lime in this
instrument?
8. In Fig. 12, why is it necessary to dry the hydrogen by means of the calcium
chloride in the tube X?
9. At what pressure would the weight of 1 l. of hydrogen be equal to that of
oxygen under standard conditions?
10. (a) What weight of hydrogen can be obtained from 150 g. of sulphuric acid?
(b) What volume would this occupy under standard conditions? (c) The density of
sulphuric acid is 1.84. What volume would the 150 g. of the acid occupy?
11. How many liters of hydrogen can be obtained from 50 cc. of sulphuric acid
having a density of 1.84?
12. Suppose you wish to fill five liter bottles with hydrogen, the gas to be collected
over water in your laboratory, how many cubic centimeters of sulphuric acid
would be required?

CHAPTER IV
COMPOUNDS OF HYDROGEN AND OXYGEN; WATER AND
HYDROGEN DIOXIDE

WATER

Historical. Water was long regarded as an element. In 1781 Cavendish showed
that it is formed by the union of hydrogen and oxygen. Being a believer in the
phlogiston theory, however, he failed to interpret his results correctly. A few years
later Lavoisier repeated Cavendish's experiments and showed that water must be
regarded as a compound of hydrogen and oxygen.
General methods employed for the determination of the composition of a
compound. The composition of a compound may be determined by either of two
general processes these are known as analysis and synthesis.
1. Analysis is the process of decomposing a compound into its constituents and
determining what these constituents are. The analysis is qualitative when it results
in merely determining what elements compose the compound; it is quantitative
when the exact percentage of each constituent is determined. Qualitative analysis
must therefore precede quantitative analysis, for it must be known what elements,
are in a compound before a method can be devised for determining exactly how
much of each is present.
2. Synthesis is the process of forming a compound from its constituent parts. It is
therefore the reverse of analysis. Like analysis, it may be either qualitative or
quantitative.
Application of these methods to the determination of the composition of
water. The determination of the composition of water is a matter of great interest
not only because of the importance of the compound but also because the methods
employed illustrate the general methods of analysis and synthesis.
Methods based on analysis. The methods based on analysis may be either
qualitative or quantitative in character.
1. Qualitative analysis. As was stated in the study of oxygen, water may be
separated into its component parts by means of the electric current. The form of

apparatus ordinarily used for effecting this
analysis is shown in Fig. 18. A platinum
wire, to the end of which is attached a small
piece of platinum foil (about 15 mm. by 25
mm.), is fused through each of the tubes B
and D, as shown in the figure. The stopcocks
at the ends of these tubes are opened and
water, to which has been added about one
tenth of its volume of sulphuric acid, is
poured into the tube A until the side tubes B
and D are completely filled. The stopcocks
are then closed. The platinum wires
extending into the tubes B and D are now
connected with the wires leading from two or
three dichromate cells joined in series. The
pieces of platinum foil within the tubes thus
become the electrodes, and the current flows
from one to the other through the acidulated
water. As soon as the current passes, bubbles
Fig. 18
of gas rise from each of the electrodes and
collect in the upper part of the tubes. The gas
rising from the negative electrode is found to be hydrogen, while that from the
positive electrode is oxygen. It will be seen that the volume of the hydrogen is
approximately double that of the oxygen. Oxygen is more soluble in water than
hydrogen, and a very little of it is also lost by being converted into ozone and other
substances. It has been found that when the necessary corrections are made for the
error due to these facts, the volume of the hydrogen is exactly double that of the
oxygen.
Fig. 19 illustrates a simpler form of apparatus, which may be used in place of that
shown in Fig. 18. A glass or porcelain dish is partially filled with water to which
has been added the proper amount of acid. Two tubes filled with the same liquid
are inverted over the electrodes. The gases resulting from the decomposition of the
water collect in the tubes.
2. Quantitative analysis. The analysis just described is purely qualitative and
simply shows that water contains hydrogen and oxygen. It does not prove the
absence of other elements; indeed it does not prove that the hydrogen and oxygen
are present in the proportion in which they are liberated by the electric current. The
method may be made quantitative, however, by weighing the water decomposed
and also the hydrogen and oxygen obtained in its decomposition. If the combined

weights of the hydrogen and oxygen exactly
equal the weight of the water decomposed, then
it would be proved that the water consists of
hydrogen and oxygen in the proportion in which
they are liberated by the electric current. This
experiment is difficult to carry out, however, so
that the more accurate methods based on
synthesis are used.
Methods based on synthesis. Two steps are
necessary to ascertain the exact composition of
water by synthesis: (1) to show by qualitative
synthesis that water is formed by the union of
oxygen with hydrogen; (2) to determine by
Fig. 19
quantitative synthesis in what proportion the two
elements unite to form water. The fact that water
is formed by the combination of oxygen with hydrogen was proved in the
preceding chapter. The quantitative synthesis may be made as follows:
The combination of the two gases is brought about in a
tube called a eudiometer. This is a graduated tube about
60 cm. long and 2 cm. wide, closed at one end (Fig. 20).
Near the closed end two platinum wires are fused through
the glass, the ends of the wires within the tube being
separated by a space of 2 mm or 3 mm. The tube is
entirely filled with mercury and inverted in a vessel of the
same liquid. Pure hydrogen is passed into the tube until it
is about one fourth filled. The volume of the gas is then
read off on the scale and reduced to standard conditions.
Approximately an equal volume of pure oxygen is then
introduced and the volume again read off and reduced to
standard conditions. This gives the total volume of the
two gases. From this the volume of the oxygen introduced
may be determined by subtracting from it the volume of
the hydrogen. The combination of the two gases is now
brought about by connecting the two platinum wires with
an induction coil and passing a spark from one wire to the
Fig. 20
other. Immediately a slight explosion occurs. The mercury
in the tube is at first depressed because of the expansion
of the gases due to the heat generated, but at once rebounds, taking the place of the
gases which have combined to form water. The volume of the water in the liquid

state is so small that it may be disregarded in the calculations. In order that the
temperature of the residual gas and the mercury may become uniform, the
apparatus is allowed to stand for a few minutes. The volume of the gas is then read
off and reduced to standard conditions, so that it may be compared with the
volumes of the hydrogen and oxygen originally taken. The residual gas is then
tested in order to ascertain whether it is hydrogen or oxygen, experiments having
proved that it is never a mixture of the two. From the information thus obtained the
composition of the water may be calculated. Thus, suppose the readings were as
follows:

Volume of hydrogen taken 20.3 cc.
Volume of hydrogen and oxygen 38.7
Volume of oxygen 18.4
Volume of gas left after combination has taken place (oxygen) 8.3

The 20.3 cc. of hydrogen have combined with 18.4 cc. minus 8.3 cc. (or 10.1 cc.)
of oxygen; or approximately 2 volumes of hydrogen have combined with 1 of
oxygen. Since oxygen is 15.88 times as heavy as hydrogen, the proportion by
weight in which the two gases combine is 1 part of hydrogen to 7.94 of oxygen.
Precaution. If the two gases are introduced into the eudiometer in the exact
proportions in which they combine, after the combination has taken place the
liquid will rise and completely fill the tube. Under these conditions, however, the
tube is very likely to be broken by the sudden upward rush of the liquid. Hence in
performing the experiment care is taken to introduce an excess of one of the gases.
A more convenient form of eudiometer. A form of eudiometer (Fig. 21) different
from that shown on page 43 is sometimes used to avoid the calculations necessary
in reducing the volumes of the gases to the same conditions of temperature and
pressure in order to make comparisons. With this apparatus it is possible to take
the readings of the volumes under the same conditions of temperature and
pressure, and thus compare them directly. The apparatus (Fig. 21) is filled with
mercury and the gases introduced into the tube A. The experiment is carried out as
in the preceding one, except that before taking the reading of the gas volumes,
mercury is either added to the tube B or withdrawn from it by means of the
stopcock C, until it stands at exactly the same height in both tubes. The gas
inclosed in tube A is then under atmospheric pressure; and since but a few minutes
are required for performing the experiment, the conditions of temperature and
pressure may be regarded as constant. Hence the volumes of the hydrogen and
oxygen and of the residual gas may be read off from the tube and directly
compared.

Fig. 21
Method used by Berzelius and Dumas. The method used by these investigators
enables us to determine directly the proportion by weight in which the hydrogen
and oxygen combine. Fig. 22 illustrates the apparatus used in making this
determination. B is a glass tube containing copper oxide. C and D are glass tubes
filled with calcium chloride, a substance which has great affinity for water. The
tubes B and C, including their contents, are carefully weighed, and the apparatus
connected as shown in the figure. A slow current of pure hydrogen is then passed
through A, and that part of the tube B which contains copper oxide is carefully
heated. The hydrogen combines with the oxygen present in the copper oxide to
form water, which is absorbed by the calcium chloride in tube C. The calcium
chloride in tube D prevents any moisture entering tube C from the air. The
operation is continued until an appreciable amount of water has been formed. The
tubes B and C are then weighed once more. The loss of weight in the tube B will
exactly equal the weight of oxygen taken up from the copper oxide in the
formation of the water. The gain in weight in the tube C will exactly equal the
weight of the water formed. The difference in these weights will of course equal
the weight of the hydrogen present in the water formed.

Fig. 22
Dumas' results. The above method for the determination of the composition of
water was first used by Berzelius in 1820. The work was repeated in 1843 by
Dumas, the average of whose results is as follows:

Weight of water formed 236.36 g.
Oxygen given up by the copper oxide 210.04
———
Weight of hydrogen present in water 26.32

According to this experiment the ratio of hydrogen to oxygen in water is therefore
26.32 to 210.04, or as l to 7.98
Morley's results. The American chemist Morley has recently determined the
composition of water, extreme precautions being taken to use pure materials and to
eliminate all sources of error. The hydrogen and oxygen which combined, as well
as the water formed, were all accurately weighed. According to Morley's results, 1
part of hydrogen by weight combines with 7.94 parts of oxygen to form water.
Comparison of results obtained. From the above discussions it is easy to see that
it is by experiment alone that the composition of a compound can be determined.
Different methods may lead to slightly different results. The more accurate the
method chosen and the greater the skill with which the experiment is carried out,
the more accurate will be the results. It is generally conceded by chemists that the
results obtained by Morley in reference to the composition of water are the most
accurate ones. In accordance with these results, then, water must be regarded as a
compound containing hydrogen and oxygen in the proportion of 1 part by weight
of hydrogen to 7.94 parts by weight of oxygen.
Relation between the volume of aqueous vapor and the volumes of the
hydrogen and oxygen which combine to form it. When the quantitative
synthesis of water is carried out in the eudiometer as described above, the water
vapor formed by the union of the hydrogen and oxygen at once condenses. The
volume of the resulting liquid is so small that it may be disregarded in making the
calculations. If, however, the experiment is carried out at a temperature of 100° or

above, the water-vapor formed is not condensed and it thus becomes possible to
compare the volume of the vapor with the volumes of hydrogen and oxygen which
combined to form it. This can be accomplished by surrounding the arm A of the
eudiometer (Fig. 23) with the tube B through which is passed the vapor obtained
by boiling some liquid which has a boiling point above 100°. In this way it has
been proved that 2 volumes of hydrogen and 1 volume of oxygen combine to form
exactly 2 volumes of water vapor, the volumes all being measured under the same
conditions of temperature and pressure. It will be noted that the relation between
these volumes may be expressed by whole numbers. The significance of this very
important fact will be discussed in a subsequent chapter.
Occurrence of water. Water not only covers
about three fourths of the surface of the earth, and
is present in the atmosphere in the form of
moisture, but it is also a common constituent of
the soil and rocks and of almost every form of
animal and vegetable organism. The human body
is nearly 70% water. This is derived not only from
the water which we drink but also from the food
which we eat, most of which contains a large
percentage of water. Thus potatoes contain about
78% of water, milk 85%, beef over 50%, apples
84%, tomatoes 94%.
Impurities in water. Chemically pure water
contains only hydrogen and oxygen. Such a water
never occurs in nature, however, for being a good
solvent, it takes up certain substances from the
rocks and soil with which it comes in contact.
When such waters are evaporated these substances
are deposited in the form of a residue. Even rain Fig. 23
water, which is the purest form occurring in
nature, contains dust particles and gases dissolved from the atmosphere. The
foreign matter in water is of two kinds, namely, mineral, such as common salt and
limestone, and organic, that is the products of animal and vegetable life.
Mineral matter in water. The amount and nature of the mineral matter present in
different waters vary greatly, depending on the character of the rocks and soil with
which the waters come in contact. The more common of the substances present are
common salt and compounds of calcium, magnesium, and iron. One liter of the
average river water contains about 175 mg. of mineral matter. Water from deep

wells naturally contains more mineral matter than river water, generally two or
three times as much, while sea water contains as much as 35,000 mg. to the liter.
Effect of impurities on health. The mineral matter in water does not, save in very
exceptional cases, render the water injurious to the human system. In fact the
presence of a certain amount of such matter is advantageous, supplying the
mineral constituents necessary for the formation of the solid tissues of the body.
The presence of organic matter, on the other hand, must always be regarded with
suspicion. This organic matter may consist not only of the products of animal and
vegetable life but also of certain microscopic forms of living organisms which are
likely to accompany such products. Contagious diseases are known to be due to
the presence in the body of minute living organisms or germs. Each disease is
caused by its own particular kind of germ. Through sewage these germs may find
their way from persons afflicted with disease into the water supply, and it is
principally through the drinking water that certain of these diseases, especially
typhoid fever, are spread. It becomes of great importance, therefore, to be able to
detect such matter when present in drinking water as well as to devise methods
whereby it can be removed or at least rendered harmless.
Analysis of water. The mineral analysis of a water is, as the name suggests,
simply the determination of the mineral matter present. Sanitary analysis, on the
other hand, is the determination of the organic matter present. The physical
properties of a water give no conclusive evidence as to its purity, since a water
may be unfit for drinking purposes and yet be perfectly clear and odorless. Neither
can any reliance be placed on the simple methods often given for testing the purity
of water. Only the trained chemist can carry out such methods of analysis as can be
relied upon.

Fig. 24

Purification of water. Three general methods are used for the purification of
water, namely, distillation, filtration, and boiling.
1. Distillation. The most effective way of purifying natural waters is by the process
of distillation. This consists in boiling the water and condensing the steam. Fig. 24
illustrates the process of distillation, as commonly conducted in the laboratory.
Ordinary water is poured into the flask A and boiled. The steam is conducted
through the condenser B, which consists essentially of a narrow glass tube sealed
within a larger one, the space between the two being filled with cold water, which
is admitted at C and escapes at D. The inner tube is thus kept cool and the steam in
passing through it is condensed. The water formed by the condensation of the
steam collects in the receiver E and is known as distilled water. Such water is
practically pure, since the impurities are nonvolatile and remain in the flask A.
Commercial distillation. In preparing distilled water on a large scale, the steam is
generated in a boiler or other metal container and condensed by passing it through
a pipe made of metal, generally tin. This pipe is wound into a spiral and is
surrounded by a current of cold water. Distilled water is used by the chemist in
almost all of his work. It is also used in the manufacture of artificial ice and for
drinking water.
Fractional distillation. In preparing distilled water, it is evident that if the natural
water contains some substance which is volatile its vapor will pass over and be
condensed with the steam, so that the distillate will not be pure water. Even such
mixtures, however, may generally be separated by repeated distillation. Thus, if a
mixture of water (boiling point 100°) and alcohol (boiling point 78°) is distilled,
the alcohol, having the lower boiling point, tends to distill first, followed by the
water. The separation of the two is not perfect, however, but may be made nearly
so by repeated distillations. The process of separating a mixture of volatile
substances by distillation is known as fractional distillation.
2. Filtration. The process of distillation practically removes all nonvolatile foreign
matter, mineral as well as organic. In purifying water for drinking purposes,
however, it is only necessary to eliminate the latter or to render it harmless. This is
ordinarily done either by filtration or boiling. In filtration the water is passed
through some medium which will retain the organic matter. Ordinary charcoal is a
porous substance and will condense within its pores the organic matter in water if
brought in contact with it. It is therefore well adapted to the construction of filters.
Such filters to be effective must be kept clean, since it is evident that the charcoal
is useless after its pores are filled. A more effective type of filter is the
Chamberlain-Pasteur filter. In this the water is forced through a porous cylindrical
cup, the pores being so minute as to strain out the organic matter.

City filtration beds. For purifying the water supply of cities, large filtration beds
are prepared from sand and gravel, and the water is allowed to filter through these.
Some of the impurities are strained out by the filter, while others are decomposed
by the action of certain kinds of bacteria present in the sand. Fig. 25 shows a cross
section of a portion of the filter used in purifying the water supply of Philadelphia.
The water filters through the sand and gravel and passes into the porous pipe A,
from which it is pumped into the city mains. The filters are covered to prevent the
water from freezing in cold weather.

Fig. 25
3. Boiling. A simpler and equally efficient method for purifying water for drinking
purposes consists in boiling the water. It is the germs in water that render it
dangerous to health. These germs are living forms of matter. If the water is boiled,
the germs are killed and the water rendered safe. While these germs are destroyed
by heat, cold has little effect upon them. Thus Dewar, in working with liquid
hydrogen, exposed some of these minute forms of life to the temperature of boiling
hydrogen (-252°) without killing them.
Self-purification of water. It has long been known that water contaminated with
organic matter tends to purify itself when exposed to the air. This is due to the fact
that the water takes up a small amount of oxygen from the air, which gradually
oxidizes the organic matter present in the water. While water is undoubtedly
purified in this way, the method cannot be relied upon to purify a contaminated
water so as to render it safe for drinking purposes.
Physical properties. Pure water is an odorless and tasteless liquid, colorless in
thin layers, but having a bluish tinge when observed through a considerable

thickness. It solidifies at 0° and boils at 100° under the normal pressure of one
atmosphere. If the pressure is increased, the boiling point is raised. When water is
cooled it steadily contracts until the temperature of 4° is reached: it then expands.
Water is remarkable for its ability to dissolve other substances, and is the best
solvent known. Solutions of solids in water are more frequently employed in
chemical work than are the solid substances, for chemical action between
substances goes on more readily when they are in solution than it does when they
are in the solid state.
Chemical properties. Water is a very stable substance, or, in other words, it does
not undergo decomposition readily. To decompose it into its elements by heat
alone requires a very high temperature; at 2500°, for example, only about 5% of
the entire amount is decomposed. Though very stable towards heat, water can be
decomposed in other ways, as by the action of the electrical current or by certain
metals.
Heat of formation and heat of decomposition are equal. The fact that a very
high temperature is necessary to decompose water into hydrogen and oxygen is in
accord with the fact that a great deal of heat is evolved by the union of hydrogen
and oxygen; for it has been proved that the heat necessary to decompose a
compound into its elements (heat of decomposition) is equal to the heat evolved in
the formation of a compound from its elements (heat of formation).
Water of crystallization. When a solid is dissolved in water and the resulting
solution is allowed to evaporate, the solid separates out, often in the form of
crystals. It has been found that the crystals of many compounds, although perfectly
dry, give up a definite amount of water when heated, the substance at the same
time losing its crystalline form. Such water is called water of crystallization. This
varies in amount with different compounds, but is perfectly definite for the same
compound. Thus, if a perfectly dry crystal of copper sulphate is strongly heated in
a tube, water is evolved and condenses on the sides of the tube, the crystal
crumbling to a light powder. The weight of the water evolved is always equal to
exactly 36.07% of the weight of copper sulphate crystals heated. The water must
therefore be in chemical combination with the substance composing the crystal;
for if simply mixed with it or adhering to it, not only would the substance appear
moist but the amount present would undoubtedly vary. The combination, however,
must be a very weak one, since the water is often expelled by even a gentle heat.
Indeed, in some cases the water is given up on simple exposure to air. Such
compounds are said to be efflorescent. Thus a crystal of sodium sulphate
(Glauber's salt) on exposure to air crumbles to a fine powder, owing to the escape
of its water of crystallization. Other substances have just the opposite property:

they absorb moisture when exposed to the air. For example, if a bit of dry calcium
chloride is placed in moist air, in the course of a few hours it will have absorbed
sufficient moisture to dissolve it. Such substances are said to be deliquescent. A
deliquescent body serves as a good drying or desiccating agent. We have already
employed calcium chloride as an agent for absorbing the moisture from hydrogen.
Many substances, as for example quartz, form crystals which contain no water of
crystallization.
Mechanically inclosed water. Water of crystallization must be carefully
distinguished from water which is mechanically inclosed in a crystal and which
can be removed by powdering the crystal and drying. Thus, when crystals of
common salt are heated, the water inclosed in the crystal is changed into steam and
bursts the crystal with a crackling sound. Such crystals are said to decrepitate.
That this water is not combined is proved by the fact that the amount present
varies and that it has all the properties of water.
Uses of water. The importance of water in its relation to life and commerce is too
well known to require comment. Its importance to the chemist has also been
pointed out. It remains to call attention to the fact that it is used as a standard in
many physical measurements. Thus 0° and 100° on the centigrade scale are
respectively the freezing and the boiling points of water under normal pressure.
The weight of 1 cc. of water at its point of greatest density is the unit of weight in
the metric system, namely, the gram. It is also taken as the unit for the
determination of the density of liquids and solids as well as for the measurement of
amounts of heat.

HYDROGEN DIOXIDE

Composition. As has been shown, 1 part by weight of hydrogen combines with
7.94 parts by weight of oxygen to form water. It is possible, however, to obtain a
second compound of hydrogen and oxygen differing from water in composition in
that 1 part by weight of hydrogen is combined with 2 × 7.94, or 15.88 parts, of
oxygen. This compound is called hydrogen dioxide or hydrogen peroxide, the
prefixes di- and per- signifying that it contains more oxygen than hydrogen oxide,
which is the chemical name for water.
Preparation. Hydrogen dioxide cannot be prepared cheaply by the direct union of
hydrogen and oxygen, and indirect methods must therefore be used. It is
commonly prepared by the action of a solution of sulphuric acid on barium
dioxide. The change which takes place may be indicated as follows:

sulphuric acid + barium dioxide = barium sulphate + hydrogen dioxide
——————— ——————— ——————— ———————
hydrogen barium barium hydrogen
sulphur oxygen sulphur oxygen
oxygen oxygen

In other words, the barium and hydrogen in the two compounds exchange places.
By this method a dilute solution of the dioxide in water is obtained. It is possible
to separate the dioxide from the water by fractional distillation. This is attended
with great difficulties, however, since the pure dioxide is explosive. The
distillation is carried on under diminished pressure so as to lower the boiling
points as much as possible; otherwise the high temperature would decompose the
dioxide.
Properties. Pure hydrogen dioxide is a colorless sirupy liquid having a density of
1.49. Its most characteristic property is the ease with which it decomposes into
water and oxygen. One part by weight of hydrogen is capable of holding firmly
only 7.94 parts of oxygen. The additional 7.94 parts of oxygen present in hydrogen
dioxide are therefore easily evolved, the compound breaking down into water and
oxygen. This decomposition is attended by the generation of considerable heat. In
dilute solution hydrogen dioxide is fairly stable, although such a solution should
be kept in a dark, cool place, since both heat and light aid in the decomposition of
the dioxide.
Uses. Solutions of hydrogen dioxide are used largely as oxidizing agents. The
solution sold by druggists contains 3% of the dioxide and is used in medicine as an
antiseptic. Its use as an antiseptic depends upon its oxidizing properties.

EXERCISES

1. Why does the chemist use distilled water in making solutions, rather than
filtered water?
2. How could you determine the total amount of solid matter dissolved in a sample
of water?
3. How could you determine whether a given sample of water is distilled water?
4. How could the presence of air dissolved in water be detected?
5. How could the amount of water in a food such as bread or potato be
determined?

6. Would ice frozen from impure water necessarily be free from disease germs?
7. Suppose that the maximum density of water were at 0° in place of 4°; what
effect would this have on the formation of ice on bodies of water?
8. Is it possible for a substance to contain both mechanically inclosed water and
water of crystallization?
9. If steam is heated to 2000° and again cooled, has any chemical change taken
place in the steam?
10. Why is cold water passed into C instead of D (Fig. 24)?
11. Mention at least two advantages that a metal condenser has over a glass
condenser.
12. Draw a diagram of the apparatus used in your laboratory for supplying distilled
water.
13. 20 cc. of hydrogen and 7 cc. of oxygen are placed in a eudiometer and the
mixture exploded. (a) How many cubic centimeters of aqueous vapor are formed?
(b) What gas and how much of it remains in excess?
14. (a) What weight of water can be formed by the combustion of 100 L of
hydrogen, measured under standard conditions? (b)What volume of oxygen would
be required in (a)? (c)What weight of potassium chlorate is necessary to prepare
this amount of oxygen?
15. What weight of oxygen is present in 1 kg. of the ordinary hydrogen dioxide
solution? In the decomposition of this weight of the dioxide into water and
oxygen, what volume of oxygen (measured under standard conditions) is evolved?

CHAPTER V
THE ATOMIC THEORY

Three fundamental laws of matter. Before we can gain any very definite idea in
regard to the structure of matter, and the way in which different kinds of
substances act chemically upon each other, it is necessary to have clearly in view
three fundamental laws of matter. These laws have been established by
experiment, and any conception which may be formed concerning matter must
therefore be in harmony with them. The laws are as follows:
Law of conservation of matter. This law has already been touched upon in the
introductory chapter, and needs no further discussion. It will be recalled that it may
be stated thus: Matter can neither be created nor destroyed, though it can be
changed from one form into another.
Law of definite composition. In the earlier days of chemistry there was much
discussion as to whether the composition of a given compound is always precisely
the same or whether it is subject to some variation. Two Frenchmen, Berthollet and
Proust, were the leaders in this discussion, and a great deal of most useful
experimenting was done to decide the question. Their experiments, as well as all
succeeding ones, have shown that the composition of a pure chemical compound is
always exactly the same. Water obtained by melting pure ice, condensing steam,
burning hydrogen in oxygen, has always 11.18% hydrogen and 88.82% oxygen in
it. Red oxide of mercury, from whatever source it is obtained, contains 92.6%
mercury and 7.4% oxygen. This truth is known as the law of definite composition,
and may be stated thus: The composition of a chemical compound never varies.
Law of multiple proportion. It has already been noted, however, that hydrogen
and oxygen combine in two different ratios to form water and hydrogen dioxide
respectively. It will be observed that this fact does not contradict the law of
definite composition, for entirely different substances are formed. These
compounds differ from each other in composition, but the composition of each one
is always constant. This ability of two elements to unite in more than one ratio is
very frequently observed. Carbon and oxygen combine in two different ratios;
nitrogen and oxygen combine to form as many as five distinct compounds, each
with its own precise composition.

In the first decade of the last century John Dalton, an English school-teacher and
philosopher, endeavored to find some rule which holds between the ratios in which
two given substances combine. His studies brought to light a very simple relation,
which the following examples will make clear. In water the hydrogen and oxygen
are combined in the ratio of 1 part by weight of hydrogen to 7.94 parts by weight
of oxygen. In hydrogen dioxide the 1 part by weight of hydrogen is combined with
15.88 parts by weight of oxygen. The ratio between the amounts of oxygen which
combine with the same amount of hydrogen to form water and hydrogen dioxide
respectively is therefore 7.94: 15.88, or 1: 2.

JOHN DALTON (English) (1766-1844)

Developed the atomic theory; made many
studies on the properties and the composition
of gases. His book entitled "A New System of
Chemical Philosophy" had a large influence
on the development of chemistry
Similarly, the element iron combines with oxygen to form two oxides, one
of which is black and the other red. By analysis it has been shown that the
former contains 1 part by weight of iron combined with 0.286 parts by
weight of oxygen, while the latter contains 1 part by weight of iron
combined with 0.429 parts by weight of oxygen. Here again we find that the
amounts of oxygen which combine with the same fixed amount of iron to
form the two compounds are in the ratio of small whole numbers, viz., 2:3.

Many other examples of this simple relation might be given, since it has
been found to hold true in all cases where more than one compound is,
formed from the same elements. Dalton's law of multiple proportion states
these facts as follows: When any two elements, A and B, combine to form
more than one compound, the amounts of B which unite with any fixed
amount of A bear the ratio of small whole numbers to each other.
Hypothesis necessary to explain the laws of matter. These three
generalizations are called laws, because they express in concise language
truths which are found by careful experiment to hold good in all cases. They
do not offer any explanation of the facts, but merely state them. The human
mind, however, does not rest content with the mere bare facts, but seeks
ever to learn the explanation of the facts. A suggestion which is offered to
explain such a set of facts is called an hypothesis. The suggestion which
Dalton offered to explain the three laws of matter, called the atomic
hypothesis, was prompted by his view of the constitution of matter, and it
involves three distinct assumptions in regard to the nature of matter and
chemical action. Dalton could not prove these assumptions to be true, but
he saw that if they were true the laws of matter become very easy to
understand.
Dalton's atomic hypothesis. The three assumptions which Dalton made in
regard to the nature of matter, and which together constitute the atomic
hypothesis, are these:
1. All elements are made up of minute, independent particles which Dalton
designated as atoms.
2. All atoms of the same element have equal masses; those of different
elements have different masses; in any change to which an atom is
subjected its mass does not change.
3. When two or more elements unite to form a compound, the action
consists in the union of a definite small number of atoms of each element to
form a small particle of the compound. The smallest particles of a given
compound are therefore exactly alike in the number and kinds of atoms
which they contain, and larger masses of the substances are simply
aggregations of these least particles.

Molecules and atoms. Dalton applied the name atom not only to the
minute particles of the elements but also to the least particles of
compounds. Later Avogadro, an Italian scientist, pointed out the fact that
the two are different, since the smallest particle of an element is a unit,
while that of a compound must have at least two units in it. He suggested
the name molecule for the least particle of a compound which can exist,
retaining the name atom for the smallest particle of an element. In
accordance with this distinction, we may define the atom and the molecule
as follows: An atom is the smallest particle of an element which can exist. A
molecule is the smallest particle of a compound which can exist. It will be
shown in a subsequent chapter that sometimes two or more atoms of the
same element unite with each other to form molecules of the element.
While the term atom, therefore, is applicable only to elements, the term
molecule is applicable both to elements and compounds.
The atomic hypothesis and the laws of matter. Supposing the atomic
hypothesis to be true, let us now see if it is in harmony with the laws of
matter.
1. The atomic hypothesis and the law of conservation of matter. It is evident
that if the atoms never change their masses in any change which they
undergo, the total quantity of matter can never change and the law of
conservation of matter must follow.
2. The atomic hypothesis and the law of definite composition. According to
the third supposition, when iron combines with sulphur the union is
between definite numbers of the two kinds of atoms. In the simplest case
one atom of the one element combines with one atom of the other. If the
sulphur and the iron atoms never change their respective masses when they
unite to form a molecule of iron sulphide, all iron sulphide molecules will
have equal amounts of iron in them and also of sulphur. Consequently any
mass made up of iron sulphide molecules will have the same fraction of
iron by weight as do the individual iron sulphide molecules. Iron sulphide,
from whatever source, will have the same composition, which is in
accordance with the law of definite composition.
3. The atomic hypothesis and the law of multiple proportion. But this
simplest case may not always be the only one. Under other conditions one

atom of iron might combine with two of sulphur to form a molecule of a
second compound. In such a case the one atom of iron would be in
combination with twice the mass of sulphur that is in the first compound,
since the sulphur atoms all have equal masses. What is true for one
molecule will be true for any number of them; consequently when such
quantities of these two compounds are selected as are found to contain the
same amount of iron, the one will contain twice as much sulphur as the
other.
The combination between the atoms may of course take place in other
simple ratios. For example, two atoms of one element might combine with
three or with five of the other. In all such cases it is clear that the law of
multiple proportion must hold true. For on selecting such numbers of the
two kinds of molecules as have the same number of the one kind of atoms,
the numbers of the other kind of atoms will stand in some simple ratio to
each other, and their weights will therefore stand in the same simple ratio.
Testing the hypothesis. Efforts have been made to find compounds which
do not conform to these laws, but all such attempts have resulted in failure.
If such compounds should be found, the laws would be no longer true, and
the hypothesis of Dalton would cease to possess value. When an hypothesis
has been tested in every way in which experiment can test it, and is still
found to be in harmony with the facts in the case, it is termed a theory. We
now speak of the atomic theory rather than of the atomic hypothesis.
Value of a theory. The value of a theory is twofold. It aids in the clear
understanding of the laws of nature because it gives an intelligent idea as to
why these laws should be in operation.
A theory also leads to discoveries. It usually happens that in testing a theory
much valuable work is done, and many new facts are discovered. Almost
any theory in explaining given laws will involve a number of consequences
apart from the laws it seeks to explain. Experiment will soon show whether
these facts are as the theory predicts they will be. Thus Dalton's atomic
theory predicted many properties of gases which experiment has since
verified.
Atomic weights. It would be of great advantage in the study of chemistry if
we could determine the weights of the different kinds of atoms. It is evident

that this cannot be done directly. They are so small that they cannot be seen
even with a most powerful microscope. It is calculated that it would take
200,000,000 hydrogen atoms placed side by side to make a row one
centimeter long. No balance can weigh such minute objects. It is possible,
however, to determine their relative weights,—that is, how much heavier
one is than another. These relative weights of the atoms are spoken of as the
atomic weights of the elements.
If elements were able to combine in only one way,—one atom of one with
one atom of another,—the problem of determining the atomic weights
would be very simple. We should merely have to take some one convenient
element as a standard, and find by experiment how much of each other
element would combine with a fixed weight of it. The ratios thus found
would be the same ratios as those between the atoms of the elements, and
thus we should have their relative atomic weights. The law of multiple
proportion calls attention to the fact that the atoms combine in other ratios
than 1: 1, and there is no direct way of telling which one, if any, of the
several compounds in a given case is the one consisting of a single atom of
each element.
If some way were to be found of telling how much heavier the entire
molecule of a compound is than the atom chosen as a standard,—that is, of
determining the molecular weights of compounds,—the problem could be
solved, though its solution would not be an entirely simple matter. There are
ways of determining the molecular weights of compounds, and there are
other experiments which throw light directly upon the relative weights of
the atoms. These methods cannot be described until the facts upon which
they rest have been studied. It will be sufficient for the present to assume
that these methods are trustworthy.
Standard for atomic weights. Since the atomic weights are merely relative
to some one element chosen as a standard, it is evident that any one of the
elements may serve as this standard and that any convenient value may be
assigned to its atom. At one time oxygen was taken as this standard, with
the value 100, and the atomic weights of the other elements were expressed
in terms of this standard. It would seem more rational to take the element of
smallest atomic weight as the standard and give it unit value; accordingly
hydrogen was taken as the standard with an atomic weight of 1. Very

recently, however, this unit has been replaced by oxygen, with an atomic
weight of 16.
Why oxygen is chosen as the standard for atomic weights. In the
determination of the atomic weight of an element it is necessary to find the
weight of the element which combines with a definite weight of another
element, preferably the element chosen as the standard. Since oxygen
combines with the elements far more readily than does hydrogen to form
definite compounds, it is far better adapted for the standard element, and
has accordingly replaced hydrogen as the standard. Any definite value
might be given to the weight of the oxygen atom. In assigning a value to it,
however, it is convenient to choose a whole number, and as small a number
as possible without making the atomic weight of any other element less
than unity. For these reasons the number 16 has been chosen as the atomic
weight of oxygen. This makes the atomic weight of hydrogen equal to
1.008, so that there is but little difference between taking oxygen as 16 and
hydrogen as 1 for the unit.
The atomic weights of the elements are given in the Appendix.

EXERCISES

1. Two compounds were found to have the following compositions: (a)
oxygen = 69.53%, nitrogen = 30.47%; (b) oxygen = 53.27%, nitrogen =
46.73%. Show that the law of multiple proportion holds in this case.
2. Two compounds were found to have the following compositions: (a)
oxygen = 43.64%, phosphorus = 56.36%; (b) oxygen = 56.35%, phosphorus
= 43.65%. Show that the law of multiple proportion holds in this case.
3. Why did Dalton assume that all the atoms of a given element have the
same weight?

CHAPTER VI
CHEMICAL EQUATIONS AND CALCULATIONS

Formulas. Since the molecule of any chemical compound consists of a
definite number of atoms, and this number never changes without
destroying the identity of the compound, it is very convenient to represent
the composition of a compound by indicating the composition of its
molecules. This can be done very easily by using the symbols of the atoms
to indicate the number and the kind of the atoms which constitute the
molecule. HgO will in this way represent mercuric oxide, a molecule of
which has been found to contain 1 atom each of mercury and oxygen. H2O
will represent water, the molecules of which consist of 1 atom of oxygen
and 2 of hydrogen, the subscript figure indicating the number of the atoms
of the element whose symbol precedes it. H2SO4 will stand for sulphuric
acid, the molecules of which contain 2 atoms of hydrogen, 1 of sulphur, and
4 of oxygen. The combination of symbols which represents the molecule of
a substance is called its formula.
Equations. When a given substance undergoes a chemical change it is
possible to represent this change by the use of such symbols and formulas.
In a former chapter it was shown that mercuric oxide decomposes when
heated to form mercury and oxygen. This may be expressed very briefly in
the form of the equation

(1) HgO = Hg + O.

When water is electrolyzed two new substances, hydrogen and oxygen, are
formed from it. This statement in the form of an equation is

(2) H2O = 2H + O.

The coefficient before the symbol for hydrogen indicates that a single
molecule of water yields two atoms of hydrogen on decomposition.

In like manner the combination of sulphur with iron is expressed by the
equation

(3) Fe + S = FeS.

The decomposition of potassium chlorate by heat takes place as represented
by the equation

(4) KClO3 = KCl + 3O.

Reading of equations. Since equations are simply a kind of shorthand way
of indicating chemical changes which occur under certain conditions, in
reading an equation the full statement for which it stands should be given.
Equation (1) should be read, "Mercuric oxide when heated gives mercury
and oxygen"; equation (2) is equivalent to the statement, "When
electrolyzed, water produces hydrogen and oxygen"; equation (3), "When
heated together iron and sulphur unite to form iron sulphide"; equation (4),
"Potassium chlorate when heated yields potassium chloride and oxygen."
Knowledge required for writing equations. In order to write such
equations correctly, a considerable amount of exact knowledge is required.
Thus, in equation (1) the fact that red oxide of mercury has the composition
represented by the formula HgO, that it is decomposed by heat, that in this
decomposition mercury and oxygen are formed and no other products,—all
these facts must be ascertained by exact experiment before the equation can
be written. An equation expressing these facts will then have much value.
Having obtained an equation describing the conduct of mercuric oxide on
being heated, it will not do to assume that other oxides will behave in like
manner. Iron oxide (FeO) resembles mercuric oxide in many respects, but it
undergoes no change at all when heated. Manganese dioxide, the black
substance used in the preparation of oxygen, has the formula MnO2. When
this substance is heated oxygen is set free, but the metal manganese is not
liberated; instead, a different oxide of manganese containing less oxygen is
produced. The equation representing the reaction is

3MnO2 = Mn3O4 + 2O.

Classes of reactions. When a chemical change takes place in a substance
the substance is said to undergo a reaction. Although a great many different
reactions will be met in the study of chemistry, they may all be grouped
under the following heads.
1. Addition. This is the simplest kind of chemical action. It consists in the
union of two or more substances to produce a new substance. The
combination of iron with sulphur is an example:

Fe + S = FeS.

2. Decomposition. This is the reverse of addition, the substance undergoing
reaction being parted into its constituents. The decomposition of mercuric
oxide is an example: HgO = Hg + O.
3. Substitution. It is sometimes possible for an element in the free state to
act upon a compound in such a way that it takes the place of one of the
elements of the compound, liberating it in turn. In the study of the element
hydrogen it was pointed out that hydrogen is most conveniently prepared by
the action of sulphuric or hydrochloric acid upon zinc. When sulphuric acid
is used a substance called zinc sulphate, having the composition represented
by the formula ZnSO4, is formed together with hydrogen. The equation is

Zn + H2SO4 = ZnSO4 + 2H.

When hydrochloric acid is used zinc chloride and hydrogen are the products
of reaction:

Zn + 2HCl = ZnCl2 + 2H.

When iron is used in place of zinc the equation is

Fe + H2SO4 = FeSO4 + 2H.

These reactions are quite similar, as is apparent from an examination of the
equations. In each case 1 atom of the metal replaces 2 atoms of hydrogen in
the acid, and the hydrogen escapes as a gas. When an element in the free
state, such as the zinc in the equations just given, takes the place of some

one element in a compound, setting it free from chemical combination, the
act is called substitution.
Other reactions illustrating substitution are the action of sodium on water,

Na + H2O = NaOH + H;

and the action of heated iron upon water,

3Fe + 4H2O = Fe3O4 + 8H.

4. Double decomposition. When barium dioxide (BaO2) is treated with
sulphuric acid two compounds are formed, namely, hydrogen dioxide
(H2O2) and barium sulphate (BaSO4). The equation is

BaO2 + H2SO4 = BaSO4 + H2O2.

In this reaction it will be seen that the two elements barium and hydrogen
simply exchange places. Such a reaction is called a double decomposition.
We shall meet with many examples of this kind of chemical reactions.
Chemical equations are quantitative. The use of symbols and formulas in
expressing chemical changes has another great advantage. Thus, according
to the equation

H2O = 2H + O,

1 molecule of water is decomposed into 2 atoms of hydrogen and 1 atom of
oxygen. But, as we have seen, the relative weights of the atoms are known,
that of hydrogen being 1.008, while that of oxygen is 16. The molecule of
water, being composed of 2 atoms of hydrogen and 1 atom of oxygen, must
therefore weigh relatively 2.016 + 16, or 18.016. The amount of hydrogen
in this molecule must be 2.016/18.016, or 11.18% of the whole, while the
amount of oxygen must be 16/18.018, or 88.82% of the whole. Now, since
any definite quantity of water is simply the sum of a great many molecules
of water, it is plain that the fractions representing the relative amounts of
hydrogen and oxygen present in a molecule must likewise express the
relative amounts of hydrogen and oxygen present in any quantity of water.

Thus, for example, in 20 g. of water there are 2.016/18.016 × 20, or 2.238 g.
of hydrogen, and 16/18.016 × 20, or 17.762 g. of oxygen. These results in
reference to the composition of water of course agree exactly with the facts
obtained by the experiments described in the chapter on water, for it is
because of those experiments that the values 1.008 and 16 are given to
hydrogen and oxygen respectively.
It is often easier to make calculations of this kind in the form of a
proportion rather than by fractions. Since the molecule of water and the two
atoms of hydrogen which it contains have the ratio by weight of 18.016:
2.016, any mass of water has the same ratio between its total weight and the
weight of the hydrogen in it. Hence, to find the number of grams (x) of
hydrogen in 20 g. of water, we have the proportion

18.016 : 2.016 :: 20 g. : x (grams of hydrogen).

Solving for x, we get 2.238 for the number of grams of hydrogen. Similarly,
to find the amount (x) of oxygen present in the 20 g. of water, we have the
proportion

18.016 : 16 :: 20 : x

from which we find that x = 17.762 g.
Again, suppose we wish to find what weight of oxygen can be obtained
from 15 g. of mercuric oxide. The equation representing the decomposition
of mercuric oxide is

HgO = Hg + O.

The relative weights of the mercury and oxygen atoms are respectively 200
and 16. The relative weight of the mercuric oxide molecule must therefore
be the sum of these, or 216. The molecule of mercuric oxide and the atom
of oxygen which it contains have the ratio 216: 16. This same ratio must
therefore hold between the weight of any given quantity of mercuric oxide
and that of the oxygen which it contains. Hence, to find the weight of
oxygen in 15 g. of mercuric oxide, we have the proportion

216 : 16 :: 15 : x (grams of oxygen).

On the other hand, suppose we wish to prepare, say, 20 g. of oxygen. The
problem is to find out what weight of mercuric oxide will yield 20 g. of
oxygen. The following proportion evidently holds

216 : 16 :: x (grams of mercuric oxide) : 20;

from which we get x = 270.
In the preparation of hydrogen by the action of sulphuric acid upon zinc,
according to the equation,

Zn + H2SO4 = ZnSO4 + 2 H,

suppose that 50 g. of zinc are available; let it be required to calculate the
weight of hydrogen which can be obtained. It will be seen that 1 atom of
zinc will liberate 2 atoms of hydrogen. The ratio by weight of a zinc to an
hydrogen atom is 65.4: 1.008; of 1 zinc atom to 2 hydrogen atoms, 65.4:
2.016. Zinc and hydrogen will be related in this reaction in this same ratio,
however many atoms of zinc are concerned. Consequently in the proportion

65.4 : 2.016 :: 50 : x,

x will be the weight of hydrogen set free by 50 g. of zinc. The weight of
zinc sulphate produced at the same time can be found from the proportion

65.4 : 161.46 :: 50 : x;

where 161.46 is the molecular weight of the zinc sulphate, and x the weight
of zinc sulphate formed. In like manner, the weight of sulphuric acid used
up can be calculated from the proportion

65.4 : 98.076 :: 50 : x.

These simple calculations are possible because the symbols and formulas in
the equations represent the relative weights of the substances concerned in a
chemical reaction. When once the relative weights of the atoms have been
determined, and it has been agreed to allow the symbols to stand for these
relative weights, an equation or formula making use of the symbols

becomes a statement of a definite numerical fact, and calculations can be
based on it.
Chemical equations not algebraic. Although chemical equations are
quantitative, it must be clearly understood that they are not algebraic. A
glance at the equations

7 + 4 = 11, 8 + 5 = 9 + 4

will show at once that they are true. The equations

HgO = Hg + O, FeO = Fe + O

are equally true in an algebraic sense, but experiment shows that only the
first is true chemically, for iron oxide (FeO) cannot be directly decomposed
into iron and oxygen. Only such equations as have been found by careful
experiment to express a real chemical transformation, true both for the
kinds of substances as well as for the weights, have any value.
Chemical formulas and equations, therefore, are a concise way of
representing qualitatively and quantitatively facts which have been found by
experiment to be true in reference to the composition of substances and the
changes which they undergo.
Formulas representing water of crystallization. An examination of
substances containing water of crystallization has shown that in every case
the water is present in such proportion by weight as can readily be
represented by a formula. For example, copper sulphate (CuSO4) and water
combine in the ratio of 1 molecule of the sulphate to 5 of water; calcium
sulphate (CaSO4) and water combine in the ratio 1: 2 to form gypsum.
These facts are expressed by writing the formulas for the two substances
with a period between them. Thus the formula for crystallized copper
sulphate is CuSO4·5H2O; that of gypsum is CaSO4·2H2O.

Heat of reaction. Attention has frequently been directed to the fact that
chemical changes are usually accompanied by heat changes. In general it
has been found that in every chemical action heat is either absorbed or
given off. By adopting a suitable unit for the measurement of heat, the heat

change during a chemical reaction can be expressed in the equation for the
reaction.
Heat cannot be measured by the use of a thermometer alone, since the
thermometer measures the intensity of heat, not its quantity. The easiest
way to measure a quantity of heat is to note how warm it will make a
definite amount of a given substance chosen as a standard. Water has been
chosen as the standard, and the unit of heat is called a calorie. A calorie is
defined as the amount of heat required to raise the temperature of one gram
of water one degree.
By means of this unit it is easy to indicate the heat changes in a given
chemical reaction. The equation

2H + O = H2O + 68,300 cal.

means that when 2.016 g. of hydrogen combine with 16 g. of oxygen,
18.016 g. of water are formed and 68,300 cal. are set free.

C + 2S = CS2 - 19,000 cal.

means that an expenditure of 19,000 cal. is required to cause 12 g. of carbon
to unite with 64.12 g. of sulphur to form 76.12 g. of carbon disulphide. In
these equations it will be noted that the symbols stand for as many grams of
the substance as there are units in the weights of the atoms represented by
the symbols. This is always understood to be the case in equations where
the heat of reaction is given.
Conditions of a chemical action are not indicated by equations.
Equations do not tell the conditions under which a reaction will take place.
The equation

HgO = Hg + O

does not tell us that it is necessary to keep the mercuric oxide at a high
temperature in order that the decomposition may go on. The equation

Zn + 2HCl = ZnCl2 + 2H

in no way indicates the fact that the hydrochloric acid must be dissolved in
water before it will act upon the zinc. From the equation

H + Cl = HCl

it would not be suspected that the two gases hydrogen and chlorine will
unite instantly in the sunlight, but will stand mixed in the dark a long time
without change. It will therefore be necessary to pay much attention to the
details of the conditions under which a given reaction occurs, as well as to
the expression of the reaction in the form of an equation.

EXERCISES

1. Calculate the percentage composition of the following substances: (a)
mercuric oxide; (b) potassium chlorate; (c) hydrochloric acid; (d) sulphuric
acid. Compare the results obtained with the compositions as given in
Chapters II and III.
2. Determine the percentage of copper, sulphur, oxygen, and water in
copper sulphate crystals. What weight of water can be obtained from 150 g.
of this substance?
3. What weight of zinc can be dissolved in 10 g. of sulphuric acid? How
much zinc sulphate will be formed?
4. How many liters of hydrogen measured under standard conditions can be
obtained from the action of 8 g. of iron on 10 g. of sulphuric acid? How
much iron sulphate (FeSO4) will be formed?

5. 10 g. of zinc were used in the preparation of hydrogen; what weight of
iron will be required to prepare an equal volume?
6. How many grams of barium dioxide will be required to prepare 1 kg. of
common hydrogen dioxide solution? What weight of barium sulphate will
be formed at the same time?
7. What weight of the compound Mn3O4 will be formed by strongly heating
25 g. of manganese dioxide? What volume of oxygen will be given off at
the same time, measured under standard conditions?

8. (a) What is the weight of 100 l. of hydrogen measured in a laboratory in
which the temperature is 20° and pressure 750 mm.? (b) What weight of
sulphuric acid is necessary to prepare this amount of hydrogen? (c) The
density of sulphuric acid is 1.84. Express the acid required in (b) in cubic
centimeters.
9. What weight of potassium chlorate is necessary to furnish sufficient
oxygen to fill four 200 cc. bottles in your laboratory (the gas to be collected
over water)?

CHAPTER VII
NITROGEN AND THE RARE ELEMENTS: ARGON,
HELIUM, NEON, KRYPTON, XENON

Historical. Nitrogen was discovered by the English chemist Rutherford in
1772. A little later Scheele showed it to be a constituent of air, and
Lavoisier gave it the name azote, signifying that it would not support life.
The name nitrogen was afterwards given it because of its presence in
saltpeter or niter. The term azote and symbol Az are still retained by the
French chemists.
Occurrence. Air is composed principally of oxygen and nitrogen in the free
state, about 78 parts by volume out of every 100 parts being nitrogen.
Nitrogen also occurs in nature in the form of potassium nitrate (KNO3)—
commonly called saltpeter or niter—as well as in sodium nitrate (NaNO3).
Nitrogen is also an essential constituent of all living organisms; for
example, the human body contains about 2.4% of nitrogen.
Preparation from air. Nitrogen can be prepared from air by the action of
some substance which will combine with the oxygen, leaving the nitrogen
free. Such a substance must be chosen, however, as will combine with the
oxygen to form a product which is not a gas, and which can be readily
separated from the nitrogen. The substances most commonly used for this
purpose are phosphorus and copper.
1. By the action of phosphorus. The method used for the preparation of
nitrogen by the action of phosphorus is as follows:
The phosphorus is placed in a little porcelain dish, supported on a cork and
floated on water (Fig. 26). It is then ignited by contact with a hot wire, and
immediately a bell jar or bottle is brought over it so as to confine a portion
of the air. The phosphorus combines with the oxygen to form an oxide of
phosphorus, known as phosphorus pentoxide. This is a white solid which
floats about in the bell jar, but in a short time it is all absorbed by the water,

leaving the nitrogen. The withdrawal of the oxygen is indicated by the
rising of the water in the bell jar.
2. By the action of
copper. The oxygen
present in the air may
also be removed by
passing air slowly
through a heated tube
containing copper. The
copper combines with
the oxygen to form
copper oxide, which is
a solid. The nitrogen
passes on and may be
collected over water.
Fig. 26 Nitrogen obtained
from air is not pure.
Inasmuch as air, in addition to oxygen and nitrogen, contains small amounts
of other gases, and since the phosphorus as well as the copper removes only
the oxygen, it is evident that the nitrogen obtained by these methods is
never quite pure. About 1% of the product is composed of other gases, from
which it is very difficult to separate the nitrogen. The impure nitrogen so
obtained may, however, be used for a study of most of the properties of
nitrogen, since these are not materially affected by the presence of the other
gases.
Preparation from compounds of nitrogen. Pure nitrogen may be obtained
from certain compounds of the element. Thus, if heat is applied to the
compound ammonium nitrite (NH4NO2), the change represented in the
following equation takes place:

NH4NO2 = 2H2O + 2N.

Physical properties. Nitrogen is similar to oxygen and hydrogen in that it
is a colorless, odorless, and tasteless gas. One liter of nitrogen weighs
1.2501 g. It is almost insoluble in water. It can be obtained in the form of a

colorless liquid having a boiling point of -195° at ordinary pressure. At
-214° it solidifies.
Chemical properties. Nitrogen is characterized by its inertness. It is neither
combustible nor a supporter of combustion. At ordinary temperatures it will
not combine directly with any of the elements except under rare conditions.
At higher temperatures it combines with magnesium, lithium, titanium, and
a number of other elements. The compounds formed are called nitrides, just
as compounds of an element with oxygen are called oxides. When it is
mixed with oxygen and subjected to the action of electric sparks, the two
gases slowly combine forming oxides of nitrogen. A mixture of nitrogen
and hydrogen when treated similarly forms ammonia, a gaseous compound
of nitrogen and hydrogen. Since we are constantly inhaling nitrogen, it is
evident that it is not poisonous. Nevertheless life would be impossible in an
atmosphere of pure nitrogen on account of the exclusion of the necessary
oxygen.
Argon, helium, neon, krypton, xenon. These are all rare elements
occurring in the air in very small quantities. Argon, discovered in 1894, was
the first one obtained. Lord Rayleigh, an English scientist, while engaged in
determining the exact weights of various gases, observed that the nitrogen
obtained from the air is slightly heavier than pure nitrogen obtained from its
compounds. After repeating his experiments many times, always with the
same results, Rayleigh finally concluded that the nitrogen which he had
obtained from the air was not pure, but was mixed with a small amount of
some unknown gas, the density of which is greater than that of nitrogen.
Acting on this assumption, Rayleigh, together with the English chemist
Ramsay, attempted to separate the nitrogen from the unknown gas.
Knowing that nitrogen would combine with magnesium, they passed the
nitrogen obtained from the air and freed from all known substances through
tubes containing magnesium heated to the necessary temperature. After
repeating this operation, they finally succeeded in obtaining from the
atmospheric nitrogen a small volume of gas which would not combine with
magnesium and hence could not be nitrogen. This proved to be a new
element, to which they gave the name argon. As predicted, this new
element was found to be heavier than nitrogen, its density as compared with
hydrogen as a standard being approximately 20, that of nitrogen being only
14. About 1% of the atmospheric nitrogen proved to be argon. The new

element is characterized by having no affinity for other elements. Even
under the most favorable conditions it has not been made to combine with
any other element. On this account it was given the name argon, signifying
lazy or idle. Like nitrogen, it is colorless, odorless, and tasteless. It has been
liquefied and solidified. Its boiling point is -187°.
Helium was first found in the gases expelled from certain minerals by
heating. Through the agency of the spectroscope it had been known to exist
in the sun long before its presence on the earth had been demonstrated,—a
fact suggested by the name helium, signifying the sun. Its existence in
traces in the atmosphere has also been proven. It was first liquefied by
Onnes in July, 1908. Its boiling point, namely -269°, is the lowest
temperature yet reached.
The remaining elements of this group—neon, krypton, and xenon—have
been obtained from liquid air. When liquid air is allowed to boil, the
constituents which are the most difficult to liquefy, and which therefore
have the lowest boiling points, vaporize first, followed by the others in the
order of their boiling points. It is possible in this way to make at least a
partial separation of the air into its constituents, and Ramsay thus succeeded
in obtaining from liquid air not only the known constituents, including
argon and helium, but also the new elements, neon, krypton, and xenon.
These elements, as well as helium, all proved to be similar to argon in that
they are without chemical activity, apparently forming no compounds
whatever. The percentages present in the air are very small. The names,
neon, krypton, xenon, signify respectively, new, hidden, stranger.

EXERCISES

1. How could you distinguish between oxygen, hydrogen, and nitrogen?
2. Calculate the relative weights of nitrogen and oxygen; of nitrogen and
hydrogen.
3. In the preparation of nitrogen from the air, how would hydrogen do as a
substance for the removal of the oxygen?
4. What weight of nitrogen can be obtained from 10 l. of air measured under
the conditions of temperature and pressure which prevail in your

laboratory?
5. How many grams of ammonium nitrite are necessary in the preparation
of 20 l. of nitrogen measured over water under the conditions of
temperature and pressure which prevail in your laboratory?
6. If 10 l. of air, measured under standard conditions, is passed over 100 g.
of hot copper, how much will the copper gain in weight?

WILLIAM RAMSAY (Scotch) (1855-)

Has made many studies in the physical
properties of substances; discovered helium;
together with Lord Rayleigh and others he
discovered argon, krypton, xenon, and neon;
has contributed largely to the knowledge of

radio-active substances, showing that radium
gradually gives rise to helium; professor at
University College, London

CHAPTER VIII
THE ATMOSPHERE

Atmosphere and air. The term atmosphere is applied to the gaseous
envelope surrounding the earth. The term air is generally applied to a
limited portion of this envelope, although the two words are often used
interchangeably. Many references have already been made to the
composition and properties of the atmosphere. These statements must now
be collected and discussed somewhat more in detail.
Air formerly regarded as an element. Like water, air was at first regarded
as elementary in character. Near the close of the eighteenth century Scheele,
Priestley, and Lavoisier showed by their experiments that it is a mixture of
at least two gases,—those which we now call oxygen and nitrogen. By
burning substances in an inclosed volume of air and noting the contraction
in volume due to the removal of the oxygen, they were able to determine
with some accuracy the relative volumes of oxygen and nitrogen present in
the air.
The constituents of the atmosphere. The constituents of the atmosphere
may be divided into two general groups: those which are essential to life
and those which are not essential.
1. Constituents essential to life. In addition to oxygen and nitrogen at least
two other substances, namely, carbon dioxide and water vapor, must be
present in the atmosphere in order that life may exist. The former of these is
a gaseous compound of carbon and oxygen having the formula CO2. Its
properties will be discussed in detail in the chapter on the compounds of
carbon. Its presence in the air may be shown by causing the air to bubble
through a solution of calcium hydroxide (Ca(OH)2), commonly called lime
water. The carbon dioxide combines with the calcium hydroxide in
accordance with the following equation:

Ca(OH)2 + CO2 = CaCO3 + H2O.

The resulting calcium carbonate (CaCO3) is insoluble in water and
separates in the form of a white powder, which causes the solution to appear
milky.
The presence of water vapor is readily shown by its condensation on cold
objects as well as by the fact that a bit of calcium chloride when exposed to
the air becomes moist, and may even dissolve in the water absorbed from
the air.
2. Constituents not essential to life. In addition to the essential constituents,
the air contains small percentages of various other gases, the presence of
which so far as is known is not essential to life. This list includes the rare
elements, argon, helium, neon, krypton, and xenon; also hydrogen,
ammonia, hydrogen dioxide, and probably ozone. Certain minute forms of
life (germs) are also present, the decay of organic matter being due to their
presence.
Function of each of the essential constituents. (1) The oxygen directly
supports life through respiration. (2) The nitrogen, on account of its
inactivity, serves to dilute the oxygen, and while contrary to the older
views, it is possible that life might continue to exist in the absence of the
atmospheric nitrogen, yet the conditions of life would be entirely changed.
Moreover, nitrogen is an essential constituent of all animal and plant life. It
was formerly supposed that neither animals nor plants could assimilate the
free nitrogen, but it has been shown recently that the plants of at least one
natural order, the Leguminosæ, to which belong the beans, peas, and clover,
have the power of directly assimilating the free nitrogen from the
atmosphere. This is accomplished through the agency of groups of bacteria,
which form colonies in little tubercles on the roots of the plants. These
bacteria probably assist in the absorption of nitrogen by changing the free
nitrogen into compounds which can be assimilated by the plant. Fig. 27
shows the tubercles on the roots of a variety of bean. (3) The presence of
water vapor in the air is necessary to prevent excessive evaporation from
both plants and animals. (4) Carbon dioxide is an essential plant food.

Fig. 27
The quantitative analysis of air. A number of different methods have been
devised for the determination of the percentages of the constituents present
in the atmosphere. Among these are the following.
1. Determination of oxygen. (1) The oxygen is withdrawn from a measured
volume of air inclosed in a tube, by means of phosphorus.
To make the determination, a graduated tube is filled with water and
inverted in a vessel of water. Air is introduced into the tube until it is
partially filled with the gas. The volume of the inclosed air is carefully
noted and reduced to standard conditions. A small piece of phosphorus is
attached to a wire and brought within the tube as shown in Fig. 28. After a
few hours the oxygen in the inclosed air will have combined with the
phosphorus, the water rising to take its place. The phosphorus is removed
and the volume is again noted and reduced to standard conditions. The
contraction in the volume of the air is equal to the volume of oxygen
absorbed.
(2) The oxygen may also be estimated by passing a measured volume of air
through a tube containing copper heated to a high temperature. The oxygen
in the air combines with the copper to form copper oxide (CuO). Hence the
increase in the weight of the copper equals the weight of the oxygen in the
volume of air taken.

(3) A more accurate method is the following. A
eudiometer tube is filled with mercury and
inverted in a vessel of the same liquid. A
convenient amount of air is then introduced into
the tube and its volume accurately noted. There
is then introduced more than sufficient
hydrogen to combine with the oxygen present
in the inclosed air, and the volume is again
accurately noted. The mixture is then exploded
by an electric spark, and the volume is once
more taken. By subtracting this volume from
the total volume of the air and hydrogen there is
obtained the contraction in volume due to the
union of the oxygen and hydrogen. The volume
occupied by the water formed by the union of
the two gases is so small that it may be
disregarded in the calculation. Since oxygen
and hydrogen combine in the ratio 1: 2 by
volume, it is evident that the contraction in
Fig. 28
volume due to the combination is equal to the
volume occupied by the oxygen in the air
contained in the tube, plus twice this volume of hydrogen. In other words,
one third of the total contraction is equal to the volume occupied by the
oxygen in the inclosed air. The following example will make this clear:

Volume of air in tube 50.0 cc.
Volume after introducing hydrogen 80.0
Volume after combination of oxygen and hydrogen 48.5
Contraction in volume due to combination (80 cc.-48.5 cc.) 31.5
Volume of oxygen in 50 cc. of air (1/3 of 31.5) 10.5

All these methods agree in showing that 100 volumes of dry air contain
approximately 21 volumes of oxygen.
2. Determination of nitrogen. If the gas left after the removal of oxygen
from a portion of air is passed over heated magnesium, the nitrogen is
withdrawn, argon and the other rare elements being left. It may thus be

shown that of the 79 volumes of gas left after the removal of the oxygen
from 100 volumes of air, approximately 78 are nitrogen and 0.93 argon. The
other elements are present in such small quantities that they may be
neglected.
3. Determination of carbon dioxide. The percentage of carbon dioxide in
any given volume of air may be determined by passing the air over calcium
hydroxide or some other compound which will combine with the carbon
dioxide. The increase in the weight of the hydroxide equals the weight of
the carbon dioxide absorbed. The amount present in the open normal air is
from 3 to 4 parts by volume in 10,000 volumes of air, or about 0.04%.
4. Determination of water vapor. The water vapor present in a given volume
of air may be determined by passing the air over calcium chloride (or some
other compound which has a strong affinity for water), and noting the
increase in the weight of the chloride. The amount present varies not only
with the locality, but there is a wide variation from day to day in the same
locality because of the winds and changes in temperature.
Processes affecting the composition of the air. The most important of
these processes are the following.
1. Respiration. In the process of respiration some of the oxygen in the
inhaled air is absorbed by the blood and carried to all parts of the body,
where it combines with the carbon of the worn-out tissues. The products of
oxidation are carried back to the lungs and exhaled in the form of carbon
dioxide. The amount exhaled by an adult averages about 20 l. per hour.
Hence in a poorly ventilated room occupied by a number of people the
amount of carbon dioxide rapidly increases. While this gas is not poisonous
unless present in large amounts, nevertheless air containing more than 15
parts in 10,000 is not fit for respiration.
2. Combustion. All of the ordinary forms of fuel contain large percentages
of carbon. On burning, this carbon combines with oxygen in the air,
forming carbon dioxide. Combustion and respiration, therefore, tend to
diminish the amount of oxygen in the air and to increase the amount of
carbon dioxide.

3. Action of plants. Plants have the power, when in the sunlight, of
absorbing carbon dioxide from the air, retaining the carbon and returning at
least a portion of the oxygen to the air. It will be observed that these
changes are just the opposite of those brought about by the processes of
respiration and combustion.
Poisonous effect of exhaled air. The differences in the percentages of
oxygen, carbon dioxide, and moisture present in inhaled air and exhaled air
are shown in the following analyses.

INHALED AIR EXHALED AIR
Oxygen 21.00% 16.00%
Carbon dioxide 0.04 4.38
Moisture variable saturated

The foul odor of respired air is due to the presence of a certain amount of
organic matter. It is possible that this organic matter rather than the carbon
dioxide is responsible for the injurious effects which follow the respiration
of impure air. The extent of such organic impurities present may be judged,
however, by the amount of carbon dioxide present, since the two are
exhaled together.
The cycle of carbon in nature. Under the influence of sunlight, the carbon
dioxide absorbed from the air by plants reacts with water and small
amounts of other substances absorbed from the soil to form complex
compounds of carbon which constitute the essential part of the plant tissue.
This reaction is attended by the evolution of oxygen, which is restored to
the air. The compounds resulting from these changes are much richer in
their energy content than are the substances from which they are formed;
hence a certain amount of energy must have been absorbed in their
formation. The source of this energy is the sun's rays.
If the plant is burned, the changes which took place in the formation of the
compounds present are largely reversed. The carbon and hydrogen present
combine with oxygen taken from the air to form carbon dioxide and water,
while the energy absorbed from the sun's rays is liberated in the form of
energy of heat. If, on the other hand, the plant is used as food, the
compounds present are used in building up the tissues of the body. When

this tissue breaks down, the changes which it undergoes are very similar to
those which take place when the plant is burned. The carbon and hydrogen
combine with the inhaled oxygen to form carbon dioxide and water, which
are exhaled. The energy possessed by the complex substances is liberated
partly in the form of energy of heat, which maintains the heat of the body,
and partly in the various forms of muscular energy. The carbon originally
absorbed from the air by the plant in the form of carbon dioxide is thus
restored to the air and is ready to repeat the cycle of changes.
The composition of the air is constant. Notwithstanding the changes
constantly taking place which tend to alter the composition of the air, the
results of a great many analyses of air collected in the open fields show that
the percentages of oxygen and nitrogen as well as of carbon dioxide are
very nearly constant. Indeed, so constant are the percentages of oxygen and
nitrogen that the question has arisen, whether these two elements are not
combined in the air, forming a definite chemical compound. That the two
are not combined but are simply mixed together can be shown in a number
of ways, among which are the following.
1. When air dissolves in water it has been found that the ratio of oxygen to
nitrogen in the dissolved air is no longer 21: 78, but more nearly 35: 65. If it
were a chemical compound, the ratio of oxygen to nitrogen would not be
changed by solution in water.
2. A chemical compound in the form of a liquid has a definite boiling point.
Water, for example, boils at 100°. Moreover the steam which is thus formed
has the same composition as the water. The boiling point of liquid air, on
the other hand, gradually rises as the liquid boils, the nitrogen escaping first
followed by the oxygen. If the two were combined, they would pass off
together in the ratio in which they are found in the air.
Why the air has a constant composition. If air is a mixture and changes
are constantly taking place which tend to modify its composition, how,
then, do we account for the constancy of composition which the analyses
reveal? This is explained by several facts. (1) The changes which are caused
by the processes of combustion and respiration, on the one hand, and the
action of plants, on the other, tend to equalize each other. (2) The winds
keep the air in constant motion and so prevent local changes. (3) The

volume of the air is so vast and the changes which occur are so small
compared with the total amount of air that they cannot be readily detected.
(4) Finally it must be noted that only air collected in the open fields shows
this constancy in composition. The air in a poorly ventilated room occupied
by a number of people rapidly changes in composition.
The properties of the air. Inasmuch as air is composed principally of a
mixture of oxygen and nitrogen, which elements have already been
discussed, its properties may be inferred largely from those of the two
gases. One liter weighs 1.2923 g. It is thus 14.38 times as heavy as
hydrogen. At the sea level it exerts an average pressure sufficient to sustain
a column of mercury 760 mm. in height. This is taken as the standard
pressure in determining the volumes of gases as well as the boiling points of
liquids. Water may be made to boil at any temperature between 0° and
considerably above 100° by simply varying the pressure. It is only when the
pressure upon it is equal to the normal pressure of the atmosphere at the sea
level, as indicated by a barometric reading of 760 mm., that it boils at 100°.
Preparation of liquid air. Attention has been called to the fact that both
oxygen and nitrogen can be obtained in the liquid state by strongly cooling
the gases and applying great pressure to them. Since air is largely a mixture
of these two gases, it can be liquefied by the same methods.
The methods for liquefying air have been simplified greatly in that the low
temperature required is obtained by allowing a portion of the compressed
air to expand. The expansion of a gas is always attended by the absorption
of heat. In liquefying air the apparatus is so constructed that the heat
absorbed is withdrawn from air already under great pressure. This process is
continued until the temperature is lowered to the point of liquefaction.
The Dewar bulb. It is not possible to preserve air in the liquid state in a
closed vessel, on account of the enormous pressure exerted by it in its
tendency to pass into the gaseous state. It may however be preserved for
some hours or even days before it will completely evaporate, by simply
placing it in an open vessel surrounded by a nonconducting material. The
most efficient vessel for this purpose is the Dewar bulb shown in Fig. 29.
The air is withdrawn from the space between the two walls, thus making it
nonconducting.

Properties and uses of liquid air.
When first prepared, liquid air is cloudy
because of the presence of particles of
solid carbon dioxide. These may be
filtered off, leaving a liquid of slightly
bluish color. It begins to boil at about
-190°, the nitrogen passing off first,
gradually followed by the oxygen, the
last portions being nearly pure oxygen.
To a certain extent oxygen is now
prepared in this way for commercial
purposes.
The extremely low temperature of liquid
air may be inferred from the fact that
mercury when cooled by it is frozen to a
mass so hard that it may be used for
driving nails.

Fig. 29 Liquid air is used in the preparation of
oxygen and as a cooling agent in the
study of the properties of matter at low
temperatures. It has thus been found that elements at extremely low
temperatures largely lose their chemical activity.

EXERCISES

1. When oxygen and nitrogen are mixed in the proportion in which they
exist in the atmosphere, heat is neither evolved nor absorbed by the process.
What important point does this suggest?
2. What essential constituent of the air is found in larger amount in
manufacturing districts than in the open country?
3. Can you suggest any reason why the growth of clover in a field improves
the soil?
4. Why are the inner walls of a Dewar bulb sometimes coated with a film of
silver?

5. To what is the blue color of liquid air due? Does this color increase in
intensity on standing?
6. When ice is placed in a vessel containing liquid air, the latter boils
violently. Explain.
7. Taking the volumes of the oxygen and nitrogen in 100 volumes of air as
21 and 78 respectively, calculate the percentages of these elements present
by weight.
8. Would combustion be more intense in liquid air than in the gaseous
substance?
9. A tube containing calcium chloride was found to weigh 30.1293 g. A
volume of air which weighed 15.2134 g. was passed through, after which
the weight of the tube was found to be 30.3405 g. What was the percentage
amount of moisture present in the air?
10. 10 l. of air measured at 20° and 740 mm. passed through lime water
caused the precipitation of 0.0102 g. of CaCO3. Find the number of
volumes of carbon dioxide in 10,000 volumes of the air.

CHAPTER IX
SOLUTIONS

Definitions. When a substance disappears in a liquid in such a way as to
thoroughly mix with it and to be lost to sight as an individual body, the
resulting liquid is called a solution. The liquid in which the substance
dissolves is called the solvent, while the dissolved substance is called the
solute.
Classes of solutions. Matter in any one of its physical states may dissolve
in a liquid, so that we may have solutions of gases, of liquids, and of solids.
Solutions of liquids in liquids are not often mentioned in the following
pages, but the other two classes will become very familiar in the course of
our study, and deserve special attention.

SOLUTION OF GASES IN LIQUIDS

It has already been
stated that oxygen,
hydrogen, and nitrogen
are slightly soluble in
water. Accurate study
has led to the
conclusion that all
gases are soluble to
some extent not only in
water but in many
other liquids. The
amount of a gas which
will dissolve in a liquid
depends upon a number
of conditions, and these Fig. 30
can best be understood

by supposing a vessel B (Fig. 30), to be filled with the gas and inverted over
the liquid. Under these circumstances the gas cannot escape or become
mixed with another gas.
Circumstances affecting the solubility of gases. A number of
circumstances affect the solubility of a gas in a liquid.
1. Nature of the gas. Other conditions being equal, each gas has its own
peculiar solubility, just as it has its own special taste or odor. The solubility
of gases varies between wide limits, as will be seen from the following
table, but as a rule a given volume of a liquid will not dissolve more than
two or three times its own volume of a gas.

Solubility of Gases in Water

1 l. of water at 760 mm. pressure and at 0° will dissolve:

Ammonia 1148.00 l.
Hydrochloric acid 503.00
Sulphur dioxide 79.79
Carbon dioxide 1.80
Oxygen 41.14 cc.
Hydrogen 21.15
Nitrogen 20.03

In the case of very soluble gases, such as the first three in the table, it is
probable that chemical combination between the liquid and the gas takes
place.
2. Nature of the liquid. The character of the liquid has much influence upon
the solubility of a gas. Water, alcohol, and ether have each its own peculiar
solvent power. From the solubility of a gas in water, no prediction can be
made as to its solubility in other liquids.
3. Influence of pressure. It has been found that the weight of gas which
dissolves in a given case is proportional to the pressure exerted upon the
gas. If the pressure is doubled, the weight of gas going into solution is

doubled; if the pressure is diminished to one half of its original value, half
of the dissolved gas will escape. Under high pressure, large quantities of gas
can be dissolved in a liquid, and when the pressure is removed the gas
escapes, causing the liquid to foam or effervesce.
4. Influence of temperature. In general, the lower the temperature of the
liquid, the larger the quantity of gas which it can dissolve. 1000 volumes of
water at 0° will dissolve 41.14 volumes of oxygen; at 50°, 18.37 volumes;
at 100° none at all. While most gases can be expelled from a liquid by
boiling the solution, some cannot. For example, it is not possible to expel
hydrochloric acid gas completely from its solution by boiling.

SOLUTION OF SOLIDS IN LIQUIDS

This is the most familiar class of solutions, since in the laboratory
substances are much more frequently used in the form of solutions than in
the solid state.
Circumstances affecting the solubility of a solid. The solubility of a solid
in a liquid depends upon several factors.
1. Nature of the solid. Other conditions being the same, solids vary greatly
in their solubility in liquids. This is illustrated in the following table:

Table of Solubility of Solids at 18°

100 cc. of water will dissolve:
Calcium chloride 71.0 g.
Sodium chloride 35.9
Potassium nitrate 29.1
Copper sulphate 21.4
Calcium sulphate 0.207

No solids are absolutely insoluble, but the amount dissolved may be so
small as to be of no significance for most purposes. Thus barium sulphate,
one of the most insoluble of common substances, dissolves in water to the
extent of 1 part in 400,000.

2. Nature of the solvent. Liquids vary much in their power to dissolve
solids. Some are said to be good solvents, since they dissolve a great variety
of substances and considerable quantities of them. Others have small
solvent power, dissolving few substances, and those to a slight extent only.
Broadly speaking, water is the most general solvent, and alcohol is perhaps
second in solvent power.
3. Temperature. The weight of a solid which a given liquid can dissolve
varies with the temperature. Usually it increases rapidly as the temperature
rises, so that the boiling liquid dissolves several times the weight which the
cold liquid will dissolve. In some instances, as in the case of common salt
dissolved in water, the temperature has little influence upon the solubility,
and a few solids are more soluble in cold water than in hot. The following
examples will serve as illustrations:

Table of Solubility at 0° and at 100°

100 cc. of water will dissolve:

At 0° At 100°
Calcium chloride 49.6 g. 155.0 g.
Sodium chloride 35.7 39.8
Potassium nitrate 13.3 247.0
Copper sulphate 15.5 73.5
Calcium sulphate 0.205 0.217
Calcium hydroxide 0.173 0.079

Saturated solutions. A liquid will not dissolve an unlimited quantity of a
solid. On adding the solid to the liquid in small portions at a time, it will be
found that a point is reached at which the liquid will not dissolve more of
the solid at that temperature. The solid and the solution remain in contact
with each other unchanged. This condition may be described by saying that
they are in equilibrium with each other. A solution is said to be saturated
when it remains unchanged in concentration in contact with some of the
solid. The weight of the solid which will completely saturate a definite

volume of a liquid at a given temperature is called the solubility of the
substance at that temperature.
Supersaturated solutions. When a solution, saturated at a given
temperature, is allowed to cool it sometimes happens that no solid
crystallizes out. This is very likely to occur when the vessel used is
perfectly smooth and the solution is not disturbed in any way. Such a
solution is said to be supersaturated. That this condition is unstable can be
shown by adding a crystal of the solid to the solution. All of the solid in
excess of the quantity required to saturate the solution at this temperature
will at once crystallize out, leaving the solution saturated. Supersaturation
may also be overcome in many cases by vigorously shaking or stirring the
solution.
General physical properties of solutions. A few general statements may
be made in reference to the physical properties of solutions.
1. Distribution of the solid in the liquid. A solid, when dissolved, tends to
distribute itself uniformly through the liquid, so that every part of the
solution has the same concentration. The process goes on very slowly
unless hastened by stirring or shaking the solution. Thus, if a few crystals of
a highly colored substance such as copper sulphate are placed in the bottom
of a tall vessel full of water, it will take weeks for the solution to become
uniformly colored.
2. Boiling points of solutions. The boiling point of a liquid is raised by the
presence of a substance dissolved in it. In general the extent to which the
boiling point of a solvent is raised by a given substance is proportional to
the concentration of the solution, that is, to the weight of the substance
dissolved in a definite weight of the solvent.
3. Freezing points of solutions. A solution freezes at a lower temperature
than the pure solvent. The lowering of the freezing point obeys the same
law which holds for the raising of the boiling point: the extent of lowering
is proportional to the weight of dissolved substance, that is, to the
concentration of the solution.
Electrolysis of solutions. Pure water does not appreciably conduct the
electric current. If, however, certain substances such as common salt are

dissolved in the water, the resulting solutions are found to be conductors of
electricity. Such solutions are called electrolytes. When the current passes
through an electrolyte some chemical change always takes place. This
change is called electrolysis.

Fig. 31
The general method used in the electrolysis of a solution is illustrated in
Fig. 31. The vessel D contains the electrolyte. Two plates or rods, A and B,
made of suitable material, are connected with the wires from a battery (or
dynamo) and dipped into the electrolyte, as shown in the figure. These
plates or rods are called electrodes. The electrode connected with the zinc
plate of the battery is the negative electrode or cathode, while that
connected with the carbon plate is the positive electrode or anode.
Theory of electrolytic dissociation. The facts which have just been
described in connection with solutions, together with many others, have led
chemists to adopt a theory of solutions called the theory of electrolytic
dissociation. The main assumptions in this theory are the following.
1. Formation of ions. Many compounds when dissolved in water undergo
an important change. A portion of their molecules fall apart, or dissociate,
into two or more parts, called ions. Thus sodium nitrate (NaNO3)
dissociates into the ions Na and NO3; sodium chloride, into the ions Na and
Cl. These ions are free to move about in the solution independently of each
other like independent molecules, and for this reason were given the name
ion, which signifies a wanderer.

2. The electrical charge of ions. Each ion carries a heavy electrical charge,
and in this respect differs from an atom or molecule. It is evident that the
sodium in the form of an ion must differ in some important way from
ordinary sodium, for sodium ions, formed from sodium nitrate, give no
visible evidence of their presence in water, whereas metallic sodium at once
decomposes the water. The electrical charge, therefore, greatly modifies the
usual chemical properties of the element.
3. The positive charges equal the negative charges. The ions formed by the
dissociation of any molecule are of two kinds. One kind is charged with
positive electricity and the other with negative electricity; moreover the
sum of all the positive charges is always equal to the sum of all the negative
charges. The solution as a whole is therefore electrically neutral. If we
represent dissociation by the usual chemical equations, with the electrical
charges indicated by + and - signs following the symbols, the dissociation
of sodium chloride molecules is represented thus:

NaCl --> Na+, Cl-.

The positive charge on each sodium ion exactly equals the negative charge
on each chlorine ion. Sodium sulphate dissociates, as shown in the equation

Na2SO4 --> 2Na+, SO4-.

Here the positive charge on the two sodium ions equals the double negative
charge on the SO4 ion.

4. Not all compounds dissociate. Only those compounds dissociate whose
solutions form electrolytes. Thus salt dissociates when dissolved in water,
the resulting solution being an electrolyte. Sugar, on the other hand, does
not dissociate and its solution is not a conductor of the electric current.
5. Extent of dissociation differs in different liquids. While compounds most
readily undergo dissociation in water, yet dissociation often occurs to a
limited extent when solution takes place in liquids other than water. In the
discussion of solutions it will be understood that the solvent is water unless
otherwise noted.

The theory of electrolytic dissociation and the properties of solutions. In
order to be of value, this theory must give a reasonable explanation of the
properties of solutions. Let us now see if the theory is in harmony with
certain of these properties.
The theory of electrolytic dissociation and the boiling and freezing
points of solutions. We have seen that the boiling point of a solution of a
substance is raised in proportion to the concentration of the dissolved
substance. This is but another way of saying that the change in the boiling
point of the solution is proportional to the number of molecules of the
dissolved substance present in the solution.
It has been found, however, that in the case of electrolytes the boiling point
is raised more than it should be to conform to this law. If the solute
dissociates into ions, the reason for this becomes clear. Each ion has the
same effect on the boiling point as a molecule, and since their number is
greater than the number of molecules from which they were formed, the
effect on the boiling point is abnormally great.
In a similar way, the theory furnishes an explanation of the abnormal
lowering of the freezing point of electrolytes.
The theory of electrolytic dissociation and electrolysis. The changes
taking place during electrolysis harmonize very completely with the theory
of dissociation. This will become clear from a study of the following
examples.

Fig. 32
1. Electrolysis of sodium chloride. Fig. 32 represents a vessel in which the
electrolyte is a solution of sodium chloride (NaCl). According to the

dissociation theory the molecules of sodium chloride dissociate into the
ions Na+ and Cl-. The Na+ ions are attracted to the cathode owing to its
large negative charge. On coming into contact with the cathode, the Na+
ions give up their positive charge and are then ordinary sodium atoms. They
immediately decompose the water according to the equation

Na + H2O = NaOH + H,

and hydrogen is evolved about the cathode.
The chlorine ions on being discharged at the anode in similar manner may
either be given off as chlorine gas, or may attack the water, as represented
in the equation

2Cl + H2O = 2HCl + O.

2. Electrolysis of water. The reason for the addition of sulphuric acid to
water in the preparation of oxygen and hydrogen by electrolysis will now
be clear. Water itself is not an electrolyte to an appreciable extent; that is, it
does not form enough ions to carry a current. Sulphuric acid dissolved in
water is an electrolyte, and dissociates into the ions 2 H+ and SO4—. In the
process of electrolysis of the solution, the hydrogen ions travel to the
cathode, and on being discharged escape as hydrogen gas. The SO4 ions,
when discharged at the anode, act upon water, setting free oxygen and once
more forming sulphuric acid:

SO4 + H2O = H2SO4 + O.

The sulphuric acid can again dissociate and the process repeat itself as long
as any water is left. Hence the hydrogen and oxygen set free in the
electrolysis of water really come directly from the acid but indirectly from
the water.
3. Electrolysis of sodium sulphate. In a similar way, sodium sulphate
(Na2SO4), when in solution, gives the ions 2 Na+ and SO4—. On being
discharged, the sodium atoms decompose water about the cathode, as in the

case of sodium chloride, while the SO4 ions when discharged at the anode
decompose the water, as represented in the equation

SO4 + H2O = H2SO4 + O

That new substances are formed
at the cathode and anode may be
shown in the following way. A U-
tube, such as is represented in Fig.
33, is partially filled with a
solution of sodium sulphate, and
the liquid in one arm is colored
with red litmus, that in the other
with blue litmus. An electrode
placed in the red solution is made
to serve as cathode, while one in
the blue solution is made the
anode. On allowing the current to
pass, the blue solution turns red,
while the red solution turns blue.
These are exactly the changes
which would take place if sodium
hydroxide and sulphuric acid were
to be set free at the electrodes, as
Fig. 33
required by the theory.
The properties of electrolytes depend upon the ions present. When a
substance capable of dissociating into ions is dissolved in water, the
properties of the solution will depend upon two factors: (1) the ions formed
from the substance; (2) the undissociated molecules. Since the ions are
usually more active chemically than the molecules, most of the chemical
properties of an electrolyte are due to the ions rather than to the molecules.
The solutions of any two substances which give the same ion will have
certain properties in common. Thus all solutions containing the copper ion
(Cu++) are blue, unless the color is modified by the presence of ions or
molecules having some other color.

EXERCISES

1. Distinguish clearly between the following terms: electrolysis, electrolyte,
electrolytic dissociation, ions, solute, solvent, solution, saturated solution,
and supersaturated solution.
2. Why does the water from some natural springs effervesce?
3. (a) Why does not the water of the ocean freeze? (b) Why will ice and salt
produce a lower temperature than ice alone?
4. Why does shaking or stirring make a solid dissolve more rapidly in a
liquid?
5. By experiment it was found that a certain volume of water was saturated
at 100° with 114 g. of potassium nitrate. On cooling to 0° a portion of the
substance crystallized. (a) How many grams of the substance remained in
solution? (b) What was the strength of the solution at 18°? (c) How much
water had been used in the experiment?
6. (a) 10 g. of common salt were dissolved in water and the solution
evaporated to dryness; what weight of solid was left? (b) 10 g. of zinc were
dissolved in hydrochloric acid and the solution evaporated to dryness; what
weight of solid was left?
7. Account for the fact that sugar sometimes deposits from molasses, even
when no evaporation has taken place.
8. (a) From the standpoint of the theory of electrolytic dissociation, write
the simple equation for a dilute solution of copper sulphate (CuSO4); this
solution is blue. (b) In the same manner, write one for sodium sulphate; this
solution is colorless. (c) How would you account for the color of the copper
sulphate solution?
9. (a) As in the preceding exercise, write a simple equation for a dilute
solution of copper chloride (CuCl2); this solution is blue. (b) In the same
manner, write one for sodium chloride; this solution is colorless. To what is
the blue color due?

10. What component is present in concentrated sulphuric acid that is almost
wanting in very dilute sulphuric acid?
11. Why will vegetables cook faster when boiled in strong salt water than
when boiled in pure water?
12. How do you explain the foaming of soda water?

CHAPTER X
ACIDS, BASES, AND SALTS; NEUTRALIZATION

Acids, bases, and salts. The three classes of compounds known
respectively as acids, bases, and salts include the great majority of the
compounds with which we shall have to deal. It is important, therefore, for
us to consider each of these classes in a systematic way. The individual
members belonging to each class will be discussed in detail in the
appropriate places, but a few representatives of each class will be described
in this chapter with special reference to the common properties in
accordance with which they are classified.
The familiar acids. Hydrochloric acidis a gas composed of hydrogen and
chlorine, and has the formula HCl. The substance is very soluble in water,
and it is this solution which is usually called hydrochloric acid. Nitric acid
is a liquid composed of hydrogen, nitrogen, and oxygen, having the formula
HNO3. As sold commercially it is mixed with about 32% of water.
Sulphuric acid, whose composition is represented by the formula H2SO4, is
an oily liquid nearly twice as heavy as water, and is commonly called oil of
vitriol.
Characteristics of acids. (1) All acids contain hydrogen. (2) When
dissolved in water the molecules of the acid dissociate into two kinds of
ions. One of these is always hydrogen and is the cation (+), while the other
consists of the remainder of the molecule and is the anion (-). (3) The
solution tastes sour. (4) It has the power to change the color of certain
substances called indicators. Thus blue litmus is changed to red, and yellow
methyl orange is changed to red. Since all acids produce hydrogen cations,
while the anions of each are different, the properties which all acids have in
common when in solution, such as taste and action on indicators, must be
attributed to the hydrogen ions.
DEFINITION: An acid is a substance which produces hydrogen ions when
dissolved in water or other dissociating liquids.

Undissociated acids. When acids are perfectly free from water, or are
dissolved in liquids like benzene which do not have the power of
dissociating them into ions, they should have no real acid properties. This is
found to be the case. Under these circumstances they do not affect the color
of indicators or have any of the properties characteristic of acids.
The familiar bases. The bases most used in the laboratory are sodium
hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide
(Ca(OH)2). These are white solids, soluble in water, the latter sparingly so.
Some bases are very difficultly soluble in water. The very soluble ones with
most pronounced basic properties are sometimes called the alkalis.
Characteristics of bases. (1) All bases contain hydrogen and oxygen. (2)
When dissolved in water the molecules of the base dissociate into two kinds
of ions. One of these is always composed of oxygen and hydrogen and is
the anion. It has the formula OH and is called the hydroxyl ion. The
remainder of the molecule, which usually consists of a single atom, is the
cation. (3) The solution of a base has a soapy feel and a brackish taste. (4) It
reverses the color change produced in indicators by acids, turning red
litmus blue, and red methyl orange yellow. Since all bases produce
hydroxyl anions, while the cations of each are different, the properties
which all bases have in common when in solution must be due to the
hydroxyl ions.
DEFINITION: A base is a substance which produces hydroxyl ions when
dissolved in water or other dissociating liquids.
Undissociated bases. Bases, in the absence of water or when dissolved in
liquids which do not dissociate them, should have none of the properties
characteristic of this class of substances. This has been found to be the case.
For example, they have no effect upon indicators under these
circumstances.
Neutralization. When an acid and a base are brought together in solution in
proper proportion, the characteristic properties of each disappear. The
solution tastes neither sour nor brackish; it has no effect upon indicators.
There can therefore be neither hydrogen nor hydroxyl ions present in the
solution. A study of reactions of this kind has shown that the hydrogen ions
of the acid combine with the hydroxyl ions of the base to form molecules of

water, water being a substance which is not appreciably dissociated into
ions. This action of an acid on a base is called neutralization. The following
equations express the neutralization of the three acids by three bases, water
being formed in each case.

Na+, OH- + H+, Cl- = Na+, Cl- + H2O.

K+, OH- + H+, NO3- = K+, NO3- + H2O.

Ca++, (OH)2— + H2++, SO4- = Ca++, SO4— + 2H2O.

DEFINITION: Neutralization consists in the union of the hydrogen ion of
an acid with the hydroxyl ion of a base to form water.
Salts. It will be noticed that in neutralization the anion of the acid and the
cation of the base are not changed. If, however, the water is expelled by
evaporation, these two ions slowly unite, and when the water becomes
saturated with the substance so produced, it separates in the form of a solid
called a salt.
DEFINITION: A salt is a substance formed by the union of the anion of an
acid with the cation of a base.
Characteristics of salts. (1) From the definition of a salt it will be seen that
there is no element or group of elements which characterize salts. (2) Salts
as a class have no peculiar taste. (3) In the absence of all other substances
they are without action on indicators. (4) When dissolved in water they
form two kinds of ions.
Heat of neutralization. If neutralization is due to the union of hydrogen
ions with hydroxyl ions, and nothing more, it follows that when a given
weight of water is formed in neutralization, the heat set free should always
be the same, no matter from what acid and base the two kinds of ions have
been supplied. Careful experiments have shown that this is the case,
provided no other reactions take place at the same time. When 18g. of water
are formed in neutralization, 13,700 cal. of heat are set free. This is
represented in the equations

Na+, OH- + H+, Cl- = Na+, Cl- + H2O + 13,700 cal.

K+, OH- + H+, NO3- = K+, NO3- + H2O + 13,700 cal.

Ca++, (OH)2- + H2++, SO4- = Ca++, SO4- + 2H2O + 2 ×
13,700 cal.

Neutralization a quantitative act. Since neutralization is a definite
chemical act, each acid will require a perfectly definite weight of each base
for its neutralization. For example, a given weight of sulphuric acid will
always require a definite weight of sodium hydroxide, in accordance with
the equation

H2, SO4 + 2Na, OH = Na2, SO4 + 2H2O.

Determination of the ratio in neutralization. The quantities of acid and
base required in neutralization may be determined in the following way.
Dilute solutions of the two substances are prepared, the sulphuric acid being
placed in one of the burettes (Fig. 34) and the sodium hydroxide in the
other. The levels of the two liquids are then brought to the zero marks of the
burettes by means of the stopcocks. A measured volume of the acid is
drawn off into a beaker, a few drops of litmus solution added, and the
sodium hydroxide is run in drop by drop until the red litmus just turns blue.
The volume of the sodium hydroxide consumed is then noted. If the
concentrations of the two solutions are known, it is easy to calculate what
weight of sodium hydroxide is required to neutralize a given weight of
sulphuric acid. By evaporating the neutralized solution to dryness, the
weight of the sodium sulphate formed can be determined directly.
Experiment shows that the weights are always in accordance with the
equation in the preceding paragraph.
Extent of dissociation. The question will naturally arise, When an acid,
base, or salt dissolves in water, do all the molecules dissociate into ions, or
only a part of them? The experiments by which this question can be
answered cannot be described here. It has been found, however, that only a
fraction of the molecules dissociate. The percentage which will dissociate in
a given case depends upon several conditions, the chief of which are: (1)

The concentration of the solution. In
concentrated solutions only a very small
percentage of dissociation occurs. As the
solution is diluted the percentage increases, and
in very dilute solutions it may be very large,
though it is never complete in any ordinary
solution. (2) The nature of the dissolved
compound. At equal concentrations substances
differ much among themselves in the
percentage of dissociation. The great majority
of salts are about equally dissociated. Acids
and bases, on the contrary, show great
differences. Some are freely dissociated, while
others are dissociated to but a slight extent.
Strength of acids and bases. Since acid and
basic properties are due to hydrogen and
hydroxyl ions respectively, the acid or base
which will produce the greatest percentage of
Fig. 34 these ions at a given concentration must be
regarded as the strongest representative of its
class. The acids and bases described in the foregoing paragraphs are all
quite strong. In 10% solutions they are dissociated to about 50%, and this is
also approximately the extent to which most salts are dissociated at this
same concentration.
Partial neutralization. 1. Basic salts. The chemical action between an acid
and a base is not always as complete as has been represented in the
foregoing paragraphs. For example, if the base magnesium hydroxide
(Mg(OH)2) and hydrochloric acid (HCl) are brought together in the ratio of
an equal number of molecules of each, there will be only half enough
hydrogen ions for the hydroxyl ions present.
Mg, (OH)2 + H, Cl = Mg, OH, Cl + H2O.

Magnesium, hydroxyl, and chlorine ions are left at the close of the reaction,
and under the proper conditions unite to form molecules of the compound
Mg(OH)Cl. This compound, when dissolved, can form hydroxyl ions and

therefore possesses basic properties; it can also form the ions of a salt (Mg
and Cl), and has properties characteristic of salts. Substances of this kind
are called basic salts.
DEFINITION: A basic salt is a substance which can give the ions both of a
base and of a salt when dissolved in water.
2. Acid salts. In a similar way, when sulphuric acid and sodium hydroxide
are brought together in the ratio of equal numbers of the molecules of each,
it is possible to have a reaction expressed by the equation

Na, OH + H2, SO4 = Na, H, SO4 + H2O.

The ions remaining after all the hydroxyl ions have been used up are those
of an acid (H) and those of a salt (Na and SO4). These unite to form the
substance NaHSO4, and as the solution becomes saturated with this
substance through evaporation, it separates in the form of crystals. In
solution this substance can give hydrogen ions, and therefore possesses acid
properties; it can also give the ions characteristic of a salt. It is therefore
called an acid salt.
DEFINITION: An acid salt is one which can give the ions of an acid and of
a salt when in solution.
3. Normal salts. Salts which are the products of complete neutralization,
such as Na2SO4, and which in solution can give neither hydrogen nor
hydroxyl ions, but only the ions of a salt, are called normal salts to
distinguish them from acid and basic salts.
Methods of expressing reactions between compounds in solution.
Chemical equations representing reactions between substances in solution
may represent the details of the reaction, or they may simply indicate the
final products formed. In the latter case the formation of ions is not
indicated. Thus, if we wish to call attention to the details of the reaction
between sodium hydroxide and hydrochloric acid in solution, the equation
is written as follows:

Na+, OH- + H+, Cl- = Na+, Cl- + H2O.

On the other hand, if we wish simply to represent the final products formed,
the following is used.

NaOH + HCl = NaCl + H2O.

Both of these methods will therefore be used:
Radicals. It has been emphasized that the hydroxyl group (OH) always
forms the anion of a base, while the group NO3 forms the anion of nitric
acid and sodium nitrate; the group SO4, the anion of sulphuric acid and
calcium sulphate. A group of elements which in this way constitutes a part
of a molecule, acting as a unit in a chemical change, or forming ions in
solution, is called a radical. Some of these radicals have been given special
names, the names signifying the elements present in the radical. Thus we
have the hydroxyl radical (OH) and the nitrate radical (NO3).

DEFINITION: A radical is a group of elements forming part of a molecule,
and acting as a unit in chemical reactions.
Names of acids, bases, and salts. Since acids, bases, and salts are so
intimately related to each other, it is very advantageous to give names to the
three classes in accordance with some fixed system. The system universally
adopted is as follows:
Naming of bases. All bases are called hydroxides. They are distinguished
from each other by prefixing the name of the element which is in
combination with the hydroxyl group. Examples: sodium hydroxide
(NaOH); calcium hydroxide (Ca(OH)2); copper hydroxide (Cu(OH)2).

Naming of acids. The method of naming acids depends upon whether the
acid consists of two elements or three.
1. Binary acids. Acids containing only one element in addition to hydrogen
are called binary acids. They are given names consisting of the prefix
hydro-, the name of the second element present, and the termination -ic.
Examples: hydrochloric acid (HCl); hydrosulphuric acid (H2S).

2. Ternary acids. In addition to the two elements present in binary acids, the
great majority of acids also contain oxygen. They therefore consist of three

elements and are called ternary acids. It usually happens that the same three
elements can unite in different proportions to make several different acids.
The most familiar one of these is given a name ending in the suffix -ic,
while the one with less oxygen is given a similar name, but ending in the
suffix -ous. Examples: nitric acid (HNO3); nitrous acid (HNO2). In cases
where more than two acids are known, use is made of prefixes in addition to
the two suffixes -icand -ous. Thus the prefix per- signifies an acid still
richer in oxygen; the prefix hypo- signifies one with less oxygen.
Naming of salts. A salt derived from a binary acid is given a name
consisting of the names of the two elements composing it, with the
termination -ide. Example: sodium chloride (NaCl). All other binary
compounds are named in the same way.
A salt of a ternary acid is named in accordance with the acid from which it
is derived. A ternary acid with the termination -ic gives a salt with the name
ending in -ate, while an acid with termination -ous gives a salt with the
name ending in -ite. The following table will make the application of these
principles clear:

ACIDS SYMBOL SALTS SYMBOL
Hydrochloric HCl Sodium chloride NaCl
Hypochlorous HClO Sodium hypochlorite NaClO
Chlorous HClO2 Sodium chlorite NaClO2
Chloric HClO3 Sodium chlorate NaClO3
Perchloric HClO4 Sodium perchlorate NaClO4

EXERCISES

1. 25 cc. of a solution containing 40 g. of sodium hydroxide per liter was
found to neutralize 25 cc. of a solution of hydrochloric acid. What was the
strength of the acid solution?
2. After neutralizing a solution of sodium hydroxide with nitric acid, there
remained after evaporation 100 g. of sodium nitrate. How much of each
substance had been used?

3. A solution contains 18 g. of hydrochloric acid per 100 cc. It required 25
cc. of this solution to neutralize 30 cc. of a solution of sodium hydroxide.
What was the strength of the sodium hydroxide solution in parts per
hundred?
4. When perfectly dry sulphuric acid is treated with perfectly dry sodium
hydroxide, no chemical change takes place. Explain.
5. When cold, concentrated sulphuric acid is added to zinc, no change takes
place. Recall the action of dilute sulphuric acid on the same metal. How do
you account for the difference?
6. A solution of hydrochloric acid in benzene does not conduct the electric
current. When this solution is treated with zinc, will hydrogen be evolved?
Explain.
7. (a) Write equation for preparation of hydrogen from zinc and dilute
sulphuric acid. (b) Rewrite the same equation from the standpoint of the
theory of electrolytic dissociation, (c) Subtract the common SO4 ion from
both members of the equation, (d) From the resulting equation, explain in
what the preparation of hydrogen consists when examined from the
standpoint of this theory.
8. In the same manner as in the preceding exercise, explain in what the
action of sodium on water to give hydrogen consists.

CHAPTER XI
VALENCE

Definition of valence. A study of the formulas of various binary
compounds shows that the elements differ between themselves in the
number of atoms of other elements which they are able to hold in
combination. This is illustrated in the formulas

HCl, H2O, H3N, H4C.
(hydrochloric acid) (water)(ammonia) (marsh gas)

It will be noticed that while one atom of chlorine combines with one atom
of hydrogen, an atom of oxygen combines with two, an atom of nitrogen
with three, one of carbon with four. The number which expresses this
combining ratio between atoms is a definite property of each element and is
called its valence.
DEFINITION: The valence of an element is that property which determines
the number of the atoms of another element which its atom can hold in
combination.
Valence a numerical property. Valence is therefore merely a numerical
relation and does not convey any information in regard to the intensity of
the affinity between atoms. Judging by the heat liberated in their union,
oxygen has a far stronger affinity for hydrogen than does nitrogen, but an
atom of oxygen can combine with two atoms only of hydrogen, while an
atom of nitrogen can combine with three.
Measure of valence. In expressing the valence of an element we must
select some standard for comparison, just as in the measurement of any
other numerical quantity. It has been found that an atom of hydrogen is
never able to hold in combination more than one atom of any other element.
Hydrogen is therefore taken as the standard, and other elements are
compared with it in determining their valence. A number of other elements

are like hydrogen in being able to combine with at most one atom of other
elements, and such elements are called univalent. Among these are chlorine,
iodine, and sodium. Elements such as oxygen, calcium, and zinc, which can
combine with two atoms of hydrogen or other univalent elements, are said
to be divalent. Similarly, we have trivalent, tetravalent, pentavalent
elements. None have a valence of more than 8.
Indirect measure of valence. Many elements, especially among the metals,
do not readily form compounds with hydrogen, and their valence is not easy
to determine by direct comparison with the standard element. These
elements, however, combine with other univalent elements, such as
chlorine, and their valence can be determined from the compounds so
formed.
Variable valence. Many elements are able to exert different valences under
differing circumstances. Thus we have the compounds Cu2O and CuO, CO
and CO2, FeCl2 and FeCl3. It is not always possible to assign a fixed
valence to an element. Nevertheless each element tends to exert some
normal valence, and the compounds in which it has a valence different from
this are apt to be unstable and easily changed into compounds in which the
valence of the element is normal. The valences of the various elements will
become familiar as the elements are studied in detail.
Valence and combining ratios. When elements combine to form
compounds, the ratio in which they combine will be determined by their
valences. In those compounds which consist of two elements directly
combined, the union is between such numbers of the two atoms as have
equal valences. Elements of the same valence will therefore combine atom
for atom. Designating the valence of the atoms by Roman numerals placed
above their symbols, we have the formulas

II II II III I II IV IV
HCl, ZnO, BN, CSi.

A divalent element, on the other hand, will combine with two atoms of a
univalent element. Thus we have

II II II II

ZnCl2 and H2O

(the numerals above each symbol representing the sum of the valences of
the atoms of the element present). A trivalent atom will combine with three
atoms of a univalent element, as in the compound

III III
H3N.

If a trivalent element combines with a divalent element, the union will be
between two atoms of the trivalent element and three of the divalent
element, since these numbers are the smallest which have equal valences.
Thus the oxide of the trivalent metal aluminium has the formula Al2O3.
Finally one atom of a tetravalent element such as carbon will combine with
four atoms of a univalent element, as in the compound CH4, or with two
atoms of a divalent element, as in the compound CO2.

We have no knowledge as to why elements differ in their combining power,
and there is no way to determine their valences save by experiment.
Valence and the structure of compounds. Compounds will be met from
time to time which are apparent exceptions to the general statements just
made in regard to valence. Thus, from the formula for hydrogen dioxide
(H2O2), it might be supposed that the oxygen is univalent; yet it is certainly
divalent in water (H2O). That it may also be divalent in H2O2 may be made
clear as follows: The unit valence of each element may be represented
graphically by a line attached to its symbol. Univalent hydrogen and
divalent oxygen will then have the symbols H- and -O-. When atoms
combine, each unit valence of one atom combines with a unit valence of
another atom. Thus the composition of water may be expressed by the
formula H-O-H, which is meant to show that each of the unit valences of
oxygen is satisfied with the unit valence of a single hydrogen atom.
The chemical conduct of hydrogen dioxide leads to the conclusion that the
two oxygen atoms of its molecule are in direct combination with each other,
and in addition each is in combination with a hydrogen atom. This may be
expressed by the formula H-O-O-H. The oxygen in the compound is

therefore divalent, just as it is in water. It will thus be seen that the structure
of a compound must be known before the valences of the atoms making up
the compound can be definitely decided upon.
Such formulas as H-O-H and H-O-O-H are known as structural formulas,
because they are intended to show what is known in regard to the
arrangement of the atoms in the molecules.
Valence and the replacing power of atoms. Just as elements having the
same valence combine with each other atom for atom, so if they replace
each other in a chemical reaction they will do so in the same ratio. This is
seen in the following equations, in which a univalent hydrogen atom is
replaced by a univalent sodium atom:

NaOH + HCl = NaCl + H2O.

2NaOH + H2SO4 = Na2SO4 + 2H2O.

Na + H2O = NaOH + H.

Similarly, one atom of divalent calcium will replace two atoms of univalent
hydrogen or one of divalent zinc:

Ca(OH)2 + 2 HCl = CaCl2 + 2H2O.

CaCl2 + ZnSO4 = CaSO4 + ZnCl2.

In like manner, one atom of a trivalent element will replace three of a
univalent element, or two atoms will replace three atoms of a divalent
element.
Valence and its applications to formulas of salts. While the true nature of
valence is not understood and many questions connected with the subject
remain unanswered, yet many of the main facts are of much help to the
student. Thus the formula of a salt, differs from that of the acid from which
it is derived in that the hydrogen of the acid has been replaced by a metal.
If, then, it is known that a given metal forms a normal salt with a certain
acid, the formula of the salt can at once be determined if the valence of the

metal is known. Since sodium is univalent, the sodium salts of the acids
HCl and H2SO4 will be respectively NaCl and Na2SO4. One atom of
divalent zinc will replace 2 hydrogen atoms, so that the corresponding zinc
salts will be ZnCl2 and ZnSO4.

The formula for aluminium sulphate is somewhat more difficult to
determine. Aluminium is trivalent, and the simplest ratio in which the
aluminium atom can replace the hydrogen in sulphuric acid is 2 atoms of
aluminium (6 valences) to 3 molecules of sulphuric acid (6 hydrogen
atoms). The formula of the sulphate will then be Al2(SO4)3.

Valence and its application to equation writing. It will be readily seen
that a knowledge of valence is also of very great assistance in writing the
equations for reactions of double decomposition. Thus, in the general
reaction between an acid and a base, the essential action is between the
univalent hydrogen ion and the univalent hydroxyl ion. The base and the
acid must always be taken in such proportions as to secure an equal number
of each of these ions. Thus, in the reaction between ferric hydroxide
(Fe(OH)3) and sulphuric acid (H2SO4), it will be necessary to take 2
molecules of the former and 3 of the latter in order to have an equal number
of the two ions, namely, 6. The equation will then be

2Fe(OH)3 + 3H2SO4 = Fe2(SO4)3 + 6H2O.

Under certain conditions the salts Al2(SO4)3 and CaCl2 undergo double
decomposition, the two metals, aluminium and calcium, exchanging places.
The simplest ratio of exchange in this case is 2 atoms of aluminium (6
valences) and 3 atoms of calcium (6 valences). The reaction will therefore
take place between 1 molecule of Al2(SO4)3 and 3 of CaCl2, and the
equation is as follows:

Al2(SO4)3 + 3 CaCl2 = 3CaSO4 + 2AlCl3.

EXERCISES

1. Sodium, calcium, and aluminium have valences of 1, 2, and 3
respectively; write the formulas of their chlorides, sulphates, and

phosphates (phosphoric acid = H3PO4), on the supposition that they form
salts having the normal composition.
2. Iron forms one series of salts in which it has a valence of 2, and another
series in which it has a valence of 3; write the formulas for the two
chlorides of iron, also for the two sulphates, on the supposition that these
have the normal composition.
3. Write the equation representing the neutralization of each of the
following bases by each of the acids whose formulas are given:

NaOH HCl
Ba(OH)2 H2SO4
Al(OH)3 H3PO4

4. Silver acts as a univalent element and calcium as a divalent element in
the formation of their respective nitrates and chlorides. (a) Write the
formula for silver nitrate; for calcium chloride. (b) When solutions of these
two salts are mixed, the two metals, silver and calcium, exchange places;
write the equation for the reaction.
5. Antimony acts as a trivalent element in the formation of a chloride. (a)
What is the formula for antimony chloride? (b) When hydrosulphuric acid
(H2S) is passed into a solution of this chloride the hydrogen and antimony
exchange places; write the equation for the reaction.
6. Lead has a valence of 2 and iron of 3 in the compounds known
respectively as lead nitrate and ferric sulphate. (a) Write the formulas for
these two compounds. (b) When their solutions are mixed the two metals
exchange places; write the equation for the reaction.

CHAPTER XII
COMPOUNDS OF NITROGEN

Occurrence. As has been stated in a former chapter, nitrogen constitutes a
large fraction of the atmosphere. The compounds of nitrogen, however,
cannot readily be obtained from this source, since at any ordinary
temperature nitrogen is able to combine directly with very few of the
elements.
In certain forms of combination nitrogen occurs in the soil from which it is
taken up by plants and built into complex substances composed chiefly of
carbon, hydrogen, oxygen, and nitrogen. Animals feeding on these plants
assimilate the nitrogenous matter, so that this element is an essential
constituent of both plants and animals.
Decomposition of organic matter by bacteria. When living matter dies
and undergoes decay complicated chemical reactions take place, one result
of which is that the nitrogen of the organic matter is set free either as the
element nitrogen, or in the form of simple compounds, such as ammonia
(NH3) or oxides of nitrogen. Experiment has shown that all such processes
of decay are due to the action of different kinds of bacteria, each particular
kind effecting a different change.
Decomposition of organic matter by heat. When organic matter is
strongly heated decomposition into simpler substances takes place in much
the same way as in the case of bacterial decomposition. Coal is a complex
substance of vegetable origin, consisting largely of carbon, but also
containing hydrogen, oxygen, and nitrogen. When this is heated in a closed
vessel so that air is excluded, about one seventh of the nitrogen is converted
into ammonia, and this is the chief source from which ammonia and its
compounds are obtained.

COMPOUNDS OF NITROGEN WITH HYDROGEN

Ammonia (NH3). Several compounds consisting exclusively of nitrogen
and hydrogen are known, but only one, ammonia, need be considered here.
Preparation of ammonia. Ammonia is prepared in the laboratory by a
different method from the one which is used commercially.
1. Laboratory method. In the laboratory ammonia is prepared from
ammonium chloride, a compound having the formula NH4Cl, and obtained
in the manufacture of coal gas. As will be shown later in the chapter, the
group NH4 in this compound acts as a univalent radical and is known as
ammonium. When ammonium chloride is warmed with sodium hydroxide,
the ammonium and sodium change places, the reaction being expressed in
the following equation.

NH4Cl + NaOH = NaCl + NH4OH.

The ammonium hydroxide (NH4OH) so formed is unstable and breaks
down into water and ammonia.

NH4OH = NH3 + H2O.

Calcium hydroxide (Ca(OH)2) is frequently used in place of the more
expensive sodium hydroxide, the equations being

2NH4Cl + Ca(OH)2 = CaCl2 + 2NH4OH,

2NH4OH = 2H2O + 2NH3.

In the preparation, the ammonium chloride and calcium hydroxide are
mixed together and placed in a flask arranged as shown in Fig. 35. The
mixture is gently warmed, when ammonia is evolved as a gas and is
collected by displacement of air.

Fig. 35
2. Commercial method. Nearly all the ammonia of commerce comes from
the gasworks. Ordinary illuminating gas is made by distilling coal, as will
be explained later, and among the products of this distillation a solution of
ammonia in water is obtained. This solution, known as gas liquor, contains
not only ammonia but other soluble substances. Most of these combine
chemically with lime, while ammonia does not; if then lime is added to the
gas liquor and the liquor is heated, the ammonia is driven out from the
mixture. It may be dissolved again in pure, cold water, forming aqua
ammonia, or the ammonia water of commerce.
Preparation from hydrogen and nitrogen. When electric sparks are
passed for some time through a mixture of hydrogen and nitrogen, a small
percentage of the two elements in the mixture is changed into ammonia.
The action soon ceases, however, for the reason that ammonia is
decomposed by the electric discharge. The reaction expressed in the
equation
N + 3H = NH3

can therefore go in either direction depending upon the relative quantities of
the substances present. This recalls the similar change from oxygen into
ozone, which soon ceases because the ozone is in turn decomposed into
oxygen.
Physical properties. Under ordinary conditions ammonia is a gas whose
density is 0.59. It is therefore little more than half as heavy as air. It is easily
condensed into a colorless liquid, and can now be purchased in liquid form
in steel cylinders. The gas is colorless and has a strong, suffocating odor. It
is extremely soluble in water, 1 l. of water at 0° and 760 mm. pressure
dissolving 1148 l. of the gas. In dissolving this large volume of gas the
water expands considerably, so that the density of the solution is less than
that of water, the strongest solutions having a density of 0.88.
Chemical properties. Ammonia will not support combustion, nor will it
burn under ordinary conditions. In an atmosphere of oxygen it burns with a
feeble, yellowish flame. When quite dry it is not a very active substance,
but when moist it combines with a great many substances, particularly with
acids.
Uses. It has been stated that ammonia can be condensed to a liquid by the
application of pressure. If the pressure is removed from the liquid so
obtained, it rapidly passes again into the gaseous state and in so doing
absorbs a large amount of heat. Advantage is taken of this fact in the
preparation of artificial ice. Large quantities of ammonia are also used in
the preparation of ammonium compounds.
The manufacture of artificial ice. Fig. 36 illustrates the method of
preparing artificial ice. The ammonia gas is liquefied in the pipes X by
means of the pump Y. The heat generated is absorbed by water flowing over
the pipes. The pipes lead into a large brine tank, a cross section of which is
shown in the figure. Into the brine (concentrated solution of common salt)
contained in this tank are dipped the vessels A, B, C, filled with pure water.
The pressure is removed from the liquid ammonia as it passes into the pipes
immersed in the brine, and the heat absorbed by the rapid evaporation of the
liquid lowers the temperature of the brine below zero. The water in A, B, C
is thereby frozen into cakes of ice. The gaseous ammonia resulting from the

evaporation of the liquid ammonia is again condensed, so that the process is
continuous.

Fig. 36
Ammonium hydroxide (NH4OH). The solution of ammonia in water is
found to have strong basic properties and therefore contains hydroxyl ions.
It turns red litmus blue; it has a soapy feel; it neutralizes acids, forming salts
with them. It seems probable, therefore, that when ammonia dissolves in
water it combines chemically with it according to the equation

NH3 + H2O = NH4OH,

and that it is the substance NH4OH, called ammonium hydroxide, which
has the basic properties, dissociating into the ions NH4 and OH.
Ammonium hydroxide has never been obtained in a pure state. At every
attempt to isolate it the substance breaks up into water and ammonia,—

NH4OH = NH3 + H2O.

The ammonium radical. The radical NH4 plays the part of a metal in many chemical reactions
and is called ammonium. The ending -ium is given to the name to indicate the metallic properties
of the substance, since the names of the metals in general have that ending. The salts formed by
the action of the base ammonium hydroxide on acids are called ammonium salts. Thus, with
hydrochloric acid, ammonium chloride is formed in accordance with the equation

NH4OH + HCl = NH4Cl + H2O.

Similarly, with nitric acid, ammonium nitrate (NH4NO3) is formed, and with sulphuric acid,
ammonium sulphate ((NH4)2S04).

It will be noticed that in the neutralization of ammonium hydroxide by acids the group NH4
replaces one hydrogen atom of the acid, just as sodium does. The group therefore acts as a
univalent metal.
Combination of nitrogen with hydrogen by volume. Under suitable conditions ammonia can be
decomposed into nitrogen and hydrogen by passing electric sparks through the gas. Accurate
measurement has shown that when ammonia is decomposed, two volumes of the gas yield one
volume of nitrogen and three volumes of hydrogen. Consequently, if the two elements were to
combine directly, one volume of nitrogen would combine with three volumes of hydrogen to form
two volumes of ammonia. Here, as in the formation of steam from hydrogen and oxygen, small
whole numbers serve to indicate the relation between the volumes of combining gases and that of
the gaseous product.

COMPOUNDS OF NITROGEN WITH OXYGEN AND HYDROGEN

In addition to ammonium hydroxide, nitrogen forms several compounds with hydrogen and
oxygen, of which nitric acid (HNO3) and nitrous acid (HNO2) are the most familiar.

Nitric acid (HNO3). Nitric acid is not found to any extent in nature, but some of its salts,
especially sodium nitrate (NaNO3) and potassium nitrate (KNO3) are found in large quantities.
From these salts nitric acid can be obtained.

Fig. 37
Preparation of nitric acid. When sodium nitrate is treated with concentrated cold sulphuric acid,
no chemical action seems to take place. If, however, the mixture is heated in a retort, nitric acid is
given off as a vapor and may be easily condensed to a liquid by passing the vapor into a tube
surrounded by cold water, as shown in Fig. 37. An examination of the liquid left in the retort
shows that it contains sodium acid sulphate (NaHSO4), so that the reaction may be represented by
the equation

NaNO3 + H2SO4 = NaHSO4 + HNO3.

If a smaller quantity of sulphuric acid is taken and the mixture is heated to a high temperature,
normal sodium sulphate is formed:
2NaNO3 + H2SO4 = Na2SO4 + 2HNO3.

In this case, however, the higher temperature required decomposes a part of the nitric acid.
The commercial preparation of nitric acid. Fig. 38 illustrates a form of apparatus used in the
preparation of nitric acid on a large scale. Sodium nitrate and sulphuric acid are heated in the iron
retort A. The resulting acid vapors pass in the direction indicated by the arrows, and are condensed
in the glass tubes B, which are covered with cloth kept cool by streams of water. These tubes are
inclined so that the liquid resulting from the condensation of the vapors runs back into C and is
drawn off into large vessels (D).

Fig. 38
Physical properties of nitric acid. Pure nitric acid is a colorless liquid, which boils at about 86°
and has a density of 1.56. The concentrated acid of commerce contains about 68% of the acid, the
remainder being water. Such a mixture has a density of 1.4. The concentrated acid fumes
somewhat in moist air, and has a sharp choking odor.
Chemical properties. The most important chemical properties of nitric acid are the following.
1. Acid properties. As the name indicates, this substance is an acid, and has all the properties of
that class of substances. It changes blue litmus red and has a sour taste in dilute solutions. It forms
hydrogen ions in solution and neutralizes bases forming salts. It also acts upon the oxides of most
metals, forming a salt and water. It is one of the strongest acids.
2. Decomposition on heating. When boiled, or exposed for some time to sunlight, it suffers a
partial decomposition according to the equation

2HNO3 = H2O + 2NO2 + O.

The substance NO2, called nitrogen peroxide, is a brownish gas, which is readily soluble in water
and in nitric acid. It therefore dissolves in the undecomposed acid, and imparts a yellowish or
reddish color to it. Concentrated nitric acid highly charged with this substance is called fuming
nitric acid.
3. Oxidizing action. According to its formula, nitric acid contains a large percentage of oxygen,
and the reaction just mentioned shows that the compound is not a very stable one, easily
undergoing decomposition. These properties should make it a good oxidizing agent, and we find
that this is the case. Under ordinary circumstances, when acting as an oxidizing agent, it is
decomposed according to the equation

2HNO3 = H2O + 2NO + 3O.

The oxygen is taken up by the substance oxidized, and not set free, as is indicated in the equation.
Thus, if carbon is oxidized by nitric acid, the oxygen combines with carbon, forming carbon
dioxide (CO2):

C + 2O = CO2.

4. Action on metals. We have seen that when an acid acts upon a metal hydrogen is set free.
Accordingly, when nitric acid acts upon a metal, such as copper, we should expect the reaction to
take place which is expressed in the equation

Cu + 2HNO3 = Cu(NO3)2 + 2H.

This reaction does take place, but the hydrogen set free is immediately oxidized to water by
another portion of the nitric acid according to the equation

HNO3 + 3H = 2H2O + NO.

As these two equations are written, two atoms of hydrogen are given off in the first equation, while
three are used up in the second. In order that the hydrogen may be equal in the two equations, we
must multiply the first by 3 and the second by 2. We shall then have

3Cu + 6HNO3 = 3Cu(NO3)2 + 6H,

2HNO3 + 6H = 4H2O + 2NO.

The two equations may now be combined into one by adding the quantities on each side of the
equality sign, canceling the hydrogen which is given off in the one reaction and used up in the
other. We shall then have the equation

3Cu + 8HNO3 = 3Cu(NO3)2 + 2NO + 4H2O.

A number of other reactions may take place when nitric acid acts upon metals, resulting in the
formation of other oxides of nitrogen, free nitrogen, or even ammonia. The reaction just given is,
however, the usual one.
Importance of steps in a reaction. This complete equation has the advantage of making it
possible to calculate very easily the proportions in which the various substances enter into the
reaction or are formed in it. It is unsatisfactory in that it does not give full information about the
way in which the reaction takes place. For example, it does not suggest that hydrogen is at first
formed, and subsequently transformed into water. It is always much more important to remember
the steps in a chemical reaction than to remember the equation expressing the complete action; for
if these steps in the reaction are understood, the complete equation is easily obtained in the manner
just described.
Salts of nitric acid,—nitrates. The salts of nitric acid are called nitrates. Many of these salts will
be described in the study of the metals. They are all soluble in water, and when heated to a high
temperature undergo decomposition. In a few cases a nitrate on being heated evolves oxygen,
forming a nitrite:

NaNO3 = NaNO2 + O.

In other cases the decomposition goes further, and the metal is left as oxide:

Cu(NO3)2 = CuO + 2NO2 + O.

Nitrous acid (HNO2). It is an easy matter to obtain sodium nitrite (NaNO2), as the reaction given
on the previous page indicates. Instead of merely heating the nitrate, it is better to heat it together
with a mild reducing agent, such as lead, when the reaction takes place which is expressed by the
equation

NaNO3 + Pb = PbO + NaNO2.

When sodium nitrite is treated with an acid, such as sulphuric acid, it is decomposed and nitrous
acid is set free:

NaNO2 + H2SO4 = NaHSO4 + HNO2.

The acid is very unstable, however, and decomposes readily into water and nitrogen trioxide
(N2O3):

2HNO2 = H2O + N2O3.

Dilute solutions of the acid, however, can be obtained.

COMPOUNDS OF NITROGEN WITH OXYGEN

Nitrogen combines with oxygen to form five different oxides. The formulas and names of these are
as follows:

N2O nitrous oxide.
NO nitric oxide.
NO2 nitrogen peroxide.
N2O3 nitrogen trioxide, or nitrous anhydride.
N2O5 nitrogen pentoxide, or nitric anhydride.

These will now be briefly discussed.
Nitrous oxide (laughing gas) (N2O). Ammonium nitrate, like all nitrates, undergoes
decomposition when heated; and owing to the fact that it contains no metal, but does contain both
oxygen and hydrogen, the reaction is a peculiar one. It is represented by the equation

NH4NO3 = 2H2O + N2O.

The oxide of nitrogen so formed is called nitrous oxide or laughing gas. It is a colorless gas having
a slight odor. It is somewhat soluble in water, and in solution has a slightly sweetish taste. It is
easily converted into a liquid and can be purchased in this form. When inhaled it produces a kind
of hysteria (hence the name "laughing gas"), and even unconsciousness and insensibility to pain if
taken in large amounts. It has long been used as an anæsthetic for minor surgical operations, such
as those of dentistry, but owing to its unpleasant after effects it is not so much in use now as
formerly.

Chemically, nitrous oxide is remarkable for the fact that it is a very energetic oxidizing agent.
Substances such as carbon, sulphur, iron, and phosphorus burn in it almost as brilliantly as in
oxygen, forming oxides and setting free nitrogen. Evidently the oxygen in nitrous oxide cannot be
held in very firm combination by the nitrogen.
Nitric oxide (NO). We have seen that when nitric acid acts upon
metals, such as copper, the reaction represented by the following
equation takes place:

3Cu + 8HNO3 = 3Cu(NO3)3 + 2NO + 4H2O.

Nitric oxide is most conveniently prepared in this way. The metal is
placed in the flask A (Fig. 39) and the acid added slowly through the
funnel tube B. The gas escapes through C and is collected over water.
Pure nitric oxide is a colorless gas, slightly heavier than air, and is
practically insoluble in water. It is a difficult gas to liquefy. Unlike
nitrous oxide, nitric oxide does not part with its oxygen easily, and
burning substances introduced into this gas are usually extinguished.
A few substances like phosphorus, which have a very strong affinity
for oxygen and which are burning energetically in the air, will
continue to burn in an atmosphere of nitric oxide. In this case the
nitric oxide loses all of its oxygen and the nitrogen is set free as gas.
Action of nitric oxide with oxygen. When nitric oxide comes into Fig. 39
contact with oxygen or with the air, it at once combines with the
oxygen even at ordinary temperatures, forming a reddish-yellow gas
of the formula NO2, which is called nitrogen peroxide. This action is not energetic enough to
produce a flame, though considerable heat is set free.
Nitrogen peroxide (NO2). This gas, as we have just seen, is formed by allowing nitric oxide to
come into contact with oxygen. It can also be made by heating certain nitrates, such as lead nitrate:

Pb(NO3)2 = PbO + 2NO2 + O.

It is a reddish-yellow gas of unpleasant odor, which is quite poisonous when inhaled. It is heavier
than air and is easily condensed to a liquid. It dissolves in water, but this solution is not a mere
physical solution; the nitrogen peroxide is decomposed, forming a mixture of nitric and nitrous
acids:

2NO2 + H2O = HNO2 + HNO3.

Nitrogen peroxide will not combine with more oxygen; it will, however, give up a part of its
oxygen to burning substances, acting as an oxidizing agent:

NO2 = NO + O.

Acid anhydrides. The oxides N2O3 (nitrogen trioxide) and N2O5 (nitrogen pentoxide) are rarely
prepared and need not be separately described. They bear a very interesting relation to the acids of

nitrogen. When dissolved in water they combine with the water, forming acids:

N2O3 + H2O = 2HNO2,

N2O5 + H2O = 2HNO3.

On the other hand, nitrous acid very easily decomposes, yielding water and nitrogen trioxide, and
by suitable means nitric acid likewise may be decomposed into water and nitrogen pentoxide:

2HNO2 = H2O + N2O3,

2HNO3 = H2O + N2O5.

In view of the close relation between these oxides and the corresponding acids, they are called
anhydrides of the acids, N2O3 being nitrous anhydride and N2O5 nitric anhydride.

DEFINITION: Any oxide which will combine with water to form an acid, or which together with
water is formed by the decomposition of an acid, is called an anhydride of that acid.

EXERCISES

1. Perfectly dry ammonia does not affect litmus paper. Explain.
2. Can ammonia be dried by passing the gas through concentrated sulphuric acid? Explain.
3. Ammonium hydroxide is a weak base, i.e. it is not highly dissociated. When it is neutralized by
strong acids the heat of reaction is less than when strong bases are so neutralized. Suggest some
possible cause for this.
4. Why is brine used in the manufacture of artificial ice?
5. Discuss the energy changes which take place in the manufacture of artificial ice.
6. What weight of ammonium chloride is necessary to furnish enough ammonia to saturate 1 l. of
water at 0° and 760 mm.?
7. What weight of sodium nitrate is necessary to prepare 100 cc. of commercial nitric acid? What
weight of potassium nitrate is necessary to furnish the same weight of acid?
8. 100 l. of nitrogen peroxide were dissolved in water and neutralized with sodium hydroxide.
What substances were formed and how much of each?(1 l. nitrogen peroxide weighs 2.05 grams.)
9. How many liters of nitrous oxide, measured under standard conditions, can be prepared from 10
g. of ammonium nitrate?
10. What weight of copper is necessary to prepare 50 l. of nitric oxide under standard conditions?
11. (a) Calculate the percentage composition of the oxides of nitrogen. (b) What important law
does this series of substances illustrate?
12. Write the equations representing the reactions between ammonium hydroxide, and sulphuric
acid and nitric acid respectively, in accordance with the theory of electrolytic dissociation.

13. In the same way, write the equations representing the reactions between nitric acid and each of
the following bases: NaOH, KOH, NH4OH, Ca(OH)2.

CHAPTER XIII
REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM

Reversible reactions. The reactions so far considered have been represented as continuing, when
once started, until one or the other substance taking part in the reaction has been used up. In some
reactions this is not the case. For example, we have seen that when steam is passed over hot iron
the reaction is represented by the equation

3Fe + 4H2O = Fe3O4 + 8H.

On the other hand, when hydrogen is passed over hot iron oxide the reverse reaction takes place:

Fe3O4 +8H = 3Fe + 4H2O.

The reaction can therefore go in either direction, depending upon the conditions of the experiment.
Such a reaction is called a reversible reaction. It is represented by an equation with double arrows
in place of the equality sign, thus:

3Fe + 4H2O <--> Fe3O4 + 8H.

In a similar way, the equation

N + 3H <--> NH3

expresses the fact that under some conditions nitrogen may unite with hydrogen to form ammonia,
while under other conditions ammonia decomposes into nitrogen and hydrogen.
The conversion of oxygen into ozone is also reversible and may be represented thus:

oxygen <--> ozone.

Chemical equilibrium. Reversible reactions do not usually go on to completion in one direction
unless the conditions under which the reaction takes place are very carefully chosen. Thus, if iron
and steam are confined in a heated tube, the steam acts upon the iron, producing iron oxide and
hydrogen. But these substances in turn act upon each other to form iron and steam once more.
When these two opposite reactions go on at such rates that the weight of the iron changed into iron
oxide is just balanced by the weight of the iron oxide changed into iron, there will be no further
change in the relative weights of the four substances present in the tube. The reaction is then said
to have reached an equilibrium.
Factors which determine the point of equilibrium. There are two factors which have a great
deal of influence in determining the point at which a given reaction will reach equilibrium.
1. Influence of the chemical nature of the substances. If two reversible reactions of the same
general kind are selected, it has been found that the point of equilibrium is different in the two
cases. For example, in the reactions represented by the equations

3Fe + 4H2O <--> Fe3O4 + 8H,

Zn + H2O <--> ZnO + 2H,

the equilibrium will be reached when very different quantities of the iron and zinc have been
changed into oxides. The individual chemical properties of the iron and zinc have therefore
marked influence upon the point at which equilibrium will be reached.
2. Influence of relative mass. If the tube in which the reaction

3Fe + 4H2O <--> Fe3O4 + 8H

has come to an equilibrium is opened and more steam is admitted, an additional quantity of the
iron will be changed into iron oxide. If more hydrogen is admitted, some of the oxide will be
reduced to metal. The point of equilibrium is therefore dependent upon the relative masses of the
substances taking part in the reaction. When one of the substances is a solid, however, its mass has
little influence, since it is only the extent of its surface which can affect the reaction.
Conditions under which reversible reactions are complete. If, when the equilibrium between
iron and steam has been reached, the tube is opened and a current of steam is passed in, the
hydrogen is swept away as fast as it is formed. The opposing reaction of hydrogen upon iron oxide
must therefore cease, and the action of steam on the iron will go on until all of the iron has been
transformed into iron oxide.
On the other hand, if a current of hydrogen is admitted into the tube, the steam will be swept away
by the hydrogen, and all of the iron oxide will be reduced to iron. A reversible reaction can
therefore be completed in either direction when one of the products of the reaction is removed as
fast as it is formed.
Equilibrium in solution. When reactions take place in solution in water the same general
principles hold good. The matter is not so simple, however, as in the case just described, owing to
the fact that many of the reactions in solution are due to the presence of ions. The substances most
commonly employed in solution are acids, bases, or salts, and all of these undergo dissociation.
Any equilibrium which may be reached in solutions of these substances must take place between
the various ions formed, on the one hand, and the undissociated molecules, on the other. Thus,
when nitric acid is dissolved in water, equilibrium is reached in accordance with the equation

H+ + NO3- <--> HNO3.

Conditions under which reversible reactions in solution are complete. The equilibrium
between substances in solution may be disturbed and the reaction caused to go on in one direction
to completion in either of three ways.
1. A gas may be formed which escapes from the solution. When sodium nitrate and sulphuric acid
are brought together in solution all four ions, Na+, NO3-, H+, SO4-, are formed. These ions are free
to rearrange themselves in various combinations. For example, the H+ and the NO3- ions will reach
the equilibrium

H+ + NO3- <--> HNO3.

If the experiment is performed with very little water present, as is the case in the preparation of
nitric acid, the equilibrium will be reached when most of the H+ and the NO3- ions have combined
to form undissociated HNO3.

Finally, if the mixture is now heated above the boiling point of nitric acid, the acid distills away as
fast as it is formed. More and more H+ and NO3- ions will then combine, and the process will
continue until one or the other of them has all been removed from the solution. The substance
remaining is sodium acid sulphate (NaHSO4), and the reaction can therefore be expressed by the
equation

NaNO3 + H2SO4 = NaHSO4 + HNO3.

2. An insoluble solid may be formed. When hydrochloric acid (HCl) and silver nitrate (AgNO3) are
brought together in solution the following ions will be present: H+, Cl-, Ag+, NO3-. The ions Ag+
and Cl- will then set up the equilibrium

Ag+ + Cl- <--> AgCl.

But silver chloride (AgCl) is almost completely insoluble in water, and as soon as a very little of it
has formed the solution becomes supersaturated, and the excess of the salt precipitates. More silver
and chlorine ions then unite, and this continues until practically all of the silver or the chlorine ions
have been removed from the solution. We then say that the following reaction is complete:

AgNO3 + HCl = AgCl + HNO3.

3. Two different ions may form undissociated molecules. In the neutralization of sodium hydroxide
by hydrochloric acid the ions H+ and OH- come to the equilibrium

H+ + OH- <--> H2O.

But since water is almost entirely undissociated, equilibrium can only be reached when there are
very few hydroxyl or hydrogen ions present. Consequently the two ions keep uniting until one or
the other of them is practically removed from the solution. When this occurs the neutralization
expressed in the following equation is complete:

NaOH + HCl = H2O + NaCl.

Preparation of acids. The principle of reversible reactions finds practical application in the
preparation of most of the common acids. An acid is usually prepared by treating the most
common of its salts with some other acid of high boiling point. The mixture is then heated until the
lower boiling acid desired distills out. Owing to its high boiling point (338°), sulphuric acid is
usually employed for this purpose, most other acids boiling below that temperature.

EXERCISES

1. What would take place when solutions of silver nitrate and sodium chloride are brought
together? What other chlorides would act in the same way?

2. Is the reaction expressed by the equation NH3 + H2O = NH4OH reversible? If so, state the
conditions under which it will go in each direction.
3. Is the reaction expressed by the equation 2H + O = H2O reversible? If so, state the conditions
under which it will go in each direction.
4. Suggest a method for the preparation of hydrochloric acid.

CHAPTER XIV
SULPHUR AND ITS COMPOUNDS

Occurrence. The element sulphur has been known from the earliest times, since it is widely
distributed in nature and occurs in large quantities in the uncombined form, especially in the
neighborhood of volcanoes. Sicily has long been famous for its sulphur mines, and smaller
deposits are found in Italy, Iceland, Mexico, and especially in Louisiana, where it is mined
extensively. In combination, sulphur occurs abundantly in the form of sulphides and sulphates. In
smaller amounts it is found in a great variety of minerals, and it is a constituent of many animal
and vegetable substances.
Extraction of sulphur. Sulphur is prepared from the native substance, the separation of crude
sulphur from the rock and earthy materials with which it is mixed being a very simple process. The
ore from the mines is merely heated until the sulphur melts and drains away from the earthy
impurities. The crude sulphur obtained in this way is distilled in a retort-shaped vessel made of
iron, the exit tube of which opens into a cooling chamber of brickwork. When the sulphur vapor
first enters the cooling chamber it condenses as a fine crystalline powder called flowers of sulphur.
As the condensing chamber becomes warm, the sulphur collects as a liquid in it, and is drawn off
into cylindrical molds, the product being called roll sulphur or brimstone.
Physical properties. Roll sulphur is a pale yellow, crystalline solid, without marked taste and with
but a faint odor. It is insoluble in water, but is freely soluble in a few liquids, notably in carbon
disulphide. Roll sulphur melts at 114.8°. Just above the melting point it forms a rather thin, straw-
colored liquid. As the temperature is raised, this liquid turns darker in color and becomes thicker,
until at about 235° it is almost black and is so thick that the vessel containing it can be inverted
without danger of the liquid running out. At higher temperatures it becomes thin once more, and
boils at 448°, forming a yellowish vapor. On cooling the same changes take place in reverse order.
Varieties of sulphur. Sulphur is known in two general forms, crystalline and amorphous. Each of
these forms exists in definite modifications.
Crystalline sulphur. Sulphur occurs in two crystalline forms, namely, rhombic sulphur and
monoclinic sulphur.
1. Rhombic sulphur. When sulphur crystallizes from its solution in carbon disulphide it separates in
crystals which have the same color and melting point as roll sulphur, and are rhombic in shape.
Roll sulphur is made up of minute rhombic crystals.
2. Monoclinic sulphur. When melted sulphur is allowed to cool until a part of the liquid has
solidified, and the remaining liquid is then poured off, it is found that the solid sulphur remaining
in the vessel has assumed the form of fine needle-shaped crystals. These differ much in appearance
from the rhombic crystals obtained by crystallizing sulphur from its solution in carbon disulphide.
The needle-shaped form is called monoclinic sulphur. The two varieties differ also in density and
in melting point, the monoclinic sulphur melting at 120°.
Monoclinic and rhombic sulphur remain unchanged in contact with each other at 96°. Above this
temperature the rhombic changes into monoclinic; at lower temperatures the monoclinic changes

into rhombic. The temperature 96° is therefore called the transition point of sulphur. Heat is set
free when monoclinic sulphur changes into rhombic.
Amorphous sulphur. Two varieties of amorphous sulphur can be readily obtained. These are
white sulphur and plastic sulphur.
1. White sulphur. Flowers of sulphur, the preparation of which has been described, consists of a
mixture of rhombic crystals and amorphous particles. When treated with carbon disulphide, the
crystals dissolve, leaving the amorphous particles as a white residue.
2. Plastic sulphur. When boiling sulphur is poured into cold water it assumes a gummy, doughlike
form, which is quite elastic. This can be seen in a very striking manner by distilling sulphur from a
small, short-necked retort, such as is represented in Fig. 40, and allowing the liquid to run directly
into water. In a few days it becomes quite brittle and passes over into ordinary rhombic sulphur.
Chemical properties of sulphur. When sulphur is
heated to its kindling temperature in oxygen or in the
air it burns with a pale blue flame, forming sulphur
dioxide (SO2). Small quantities of sulphur trioxide
(SO3) may also be formed in the combustion of
sulphur. Most metals when heated with sulphur
combine directly with it, forming metallic sulphides.
In some cases the action is so energetic that the mass
becomes incandescent, as has been seen in the case of
iron uniting with sulphur. This property recalls the
action of oxygen upon metals, and in general the
metals which combine readily with oxygen are apt to
combine quite readily with sulphur.
Uses of sulphur. Large quantities of sulphur are used
as a germicide in vineyards, also in the manufacture of
gunpowder, matches, vulcanized rubber, and sulphuric
acid.
Fig. 40

COMPOUNDS OF SULPHUR WITH HYDROGEN

Hydrosulphuric acid (H2S). This substance is a gas having the composition expressed by the
formula H2S and is commonly called hydrogen sulphide. It is found in the vapors issuing from
volcanoes, and in solution in the so-called sulphur waters of many springs. It is formed when
organic matter containing sulphur undergoes decay, just as ammonia is formed under similar
circumstances from nitrogenous matter.
Preparation. Hydrosulphuric acid is prepared in the laboratory by treating a sulphide with an acid.
Iron sulphide (FeS) is usually employed:

FeS + 2HCl = FeCl2 + H2S.

A convenient apparatus is shown in Fig. 41. A few lumps of iron sulphide are placed in the bottle
A, and dilute acid is added in small quantities at a time through the funnel tube B, the gas escaping
through the tube C.
Explanation of the reaction. Iron sulphide is a salt of hydrosulphuric
acid, and this reaction is therefore similar to the one which takes place
when sulphuric acid acts upon a nitrate. In both cases a salt and an
acid are brought together, and there is a tendency for the reaction to go
on until a state of equilibrium is reached. This equilibrium is
constantly disturbed by the escape of the gaseous acid set free, so that
the reaction goes on until all of the original salt has been decomposed.
The two reactions differ in that the first one is complete at ordinary
temperatures, while in the case of sulphuric acid acting upon sodium
nitrate, the reacting substances must be heated so as to secure a
temperature at which nitric acid is a gas.
Physical properties. Hydrosulphuric acid is a colorless gas, having a
weak, disagreeable taste and an exceedingly offensive odor. It is rather
sparingly soluble in water at ordinary temperatures, about three
volumes dissolving in one of water. In boiling water it is not soluble at
all. In pure form it acts as a violent poison, and even when diluted
largely with air produces headache, dizziness, and nausea. It is a little
heavier than air, having a density of 1.18.
Fig. 41
Chemical properties. The most important chemical properties of
hydrosulphuric acid are the following:
1. Acid properties. Hydrosulphuric acid is a weak acid. In solution in water it turns blue litmus red
and neutralizes bases, forming salts called sulphides.
2. Action on oxygen. The elements composing hydrosulphuric acid have each a strong affinity for
oxygen, and are not held together very firmly. Consequently the gas burns readily in oxygen or the
air, according to the equation

H2S + 3O = H2O + SO2.

When there is not enough oxygen for both the sulphur and the hydrogen, the latter element
combines with the oxygen and the sulphur is set free:

H2S + O = H2O + S.

3. Reducing action. Owing to the ease with which hydrosulphuric acid decomposes and the strong
affinity of both sulphur and hydrogen for oxygen, the substance is a strong reducing agent, taking
oxygen away from many substances which contain it.
4. Action on metals. Hydrosulphuric acid acts towards metals in a way very similar to water. Thus,
when it is passed over heated iron in a tube, the reaction is represented by the equation

3Fe + 4H2S = Fe3S4 + 8H.

Water in the form of steam, under similar circumstances, acts according to the equation

3Fe + 4H2O = Fe3O4 + 8H.

Salts of hydrosulphuric acid,—sulphides. The salts of hydrosulphuric acid, called sulphides,
form an important class of salts. Many of them are found abundantly in nature, and some of them
are important ores. They will be frequently mentioned in connection with the metals.
Most of the sulphides are insoluble in water, and some of them are insoluble in acids.
Consequently, when hydrosulphuric acid is passed into a solution of a salt, it often happens that a
sulphide is precipitated. With copper chloride the equation is

CuCl2 + H2S = CuS + 2HCl.

Because of the fact that some metals are precipitated in this way as sulphides while others are not,
hydrosulphuric acid is extensively used in the separation of the metals in the laboratory.
Explanation of the reaction. When hydrosulphuric acid and copper chloride are brought together
in solution, both copper and sulphur ions are present, and these will come to an equilibrium, as
represented in the equation

Cu+ + S- <--> CuS.
Since copper sulphide is almost insoluble in water, as soon as a very small quantity has formed the
solution becomes supersaturated, and the excess keeps precipitating until nearly all the copper or
sulphur ions have been removed from the solution. With some other ions, such as iron, the
sulphide formed does not saturate the solution, and no precipitate results.

OXIDES OF SULPHUR

Sulphur forms two well-known compounds with oxygen: sulphur dioxide (SO2), sometimes called
sulphurous anhydride; and sulphur trioxide (SO3), frequently called sulphuric anhydride.

Sulphur dioxide (SO2). Sulphur dioxide occurs in nature in the gases issuing from volcanoes, and
in solution in the water of many springs. It is likely to be found wherever sulphur compounds are
undergoing oxidation.
Preparation. Three general ways may be mentioned for the preparation of sulphur dioxide:
1. By the combustion of sulphur. Sulphur dioxide is readily formed by the combustion of sulphur in
oxygen or the air:

S + 2O = SO2.

It is also formed when substances containing sulphur are burned:

ZnS + 3O = ZnO + SO2.

2. By the reduction of sulphuric acid. When concentrated sulphuric acid is heated with certain
metals, such as copper, part of the acid is changed into copper sulphate, and part is reduced to

sulphurous acid. The latter then decomposes into sulphur dioxide and water, the complete equation
being

Cu + 2H2SO4 = CuSO4 + SO2 + 2H2O.

3. By the action of an acid on a sulphite. Sulphites are salts of sulphurous acid (H2SO3). When a
sulphite is treated with an acid, sulphurous acid is set free, and being very unstable, decomposes
into water and sulphur dioxide. These reactions are expressed in the equations

Na2SO3 + 2HCl = 2NaCl + H2SO3,

H2SO3 = H2O + SO2.

Explanation of the reaction. In this case we have two reversible reactions depending on each
other. In the first reaction,

(1) Na2SO3 + 2HCl <--> 2NaCl + H2SO3,

we should expect an equilibrium to result, for none of the four substances in the equation are
insoluble or volatile when water is present to hold them in solution. But the quantity of the H2SO3
is constantly diminishing, owing to the fact that it decomposes, as represented in the equation

(2) H2SO3 <--> H2O + SO2,

and the sulphur dioxide, being a gas, escapes. No equilibrium can therefore result, since the
quantity of the sulphurous acid is constantly being diminished because of the escape of sulphur
dioxide.
Physical properties. Sulphur dioxide is a colorless gas, which at ordinary temperatures is 2.2
times as heavy as air. It has a peculiar, irritating odor. The gas is very soluble in water, one volume
of water dissolving eighty of the gas under standard conditions. It is easily condensed to a
colorless liquid, and can be purchased in this condition stored in strong bottles, such as the one
represented in Fig. 42.
Chemical properties. Sulphur dioxide has a marked tendency to combine with other substances,
and is therefore an active substance chemically. It combines with oxygen gas, but not very easily.
It can, however, take oxygen away from some other substances, and is therefore a good reducing
agent. Its most marked chemical property is its ability to combine with water to form sulphurous
acid (H2SO3).

Sulphurous acid (H2SO3). When sulphur dioxide dissolves in water it combines chemically with
it to form sulphurous acid, an unstable substance having the formula H3SO3. It is impossible to
prepare this acid in pure form, as it breaks down very easily into water and sulphur dioxide. The
reaction is therefore reversible, and is expressed by the equation

H2O + SO2 <--> H2SO3.

Solutions of the acid in water have a number of interesting properties.

1. Acid properties. The solution has all the properties typical of an
acid. When neutralized by bases, sulphurous acid yields a series of
salts called sulphites.
2. Reducing properties. Solutions of sulphurous acid act as good
reducing agents. This is due to the fact that sulphurous acid has the
power of taking up oxygen from the air, or from substances rich in
oxygen, and is changed by this reaction into sulphuric acid:

H2SO3 + O = H2SO4,

H2SO3 + H2O2 = H2S04 + H2O.

3. Bleaching properties. Sulphurous acid has strong bleaching
properties, acting upon many colored substances in such a way as to
destroy their color. It is on this account used to bleach paper, straw
goods, and even such foods as canned corn.
4. Antiseptic properties. Sulphurous acid has marked antiseptic
properties, and on this account has the power of arresting Fig. 42
fermentation. It is therefore used as a preservative.
Salts of sulphurous acid,—sulphites. The sulphites, like sulphurous acid, have the power of
taking up oxygen very readily, and are good reducing agents. On account of this tendency,
commercial sulphites are often contaminated with sulphates. A great deal of sodium sulphite is
used in the bleaching industry, and as a reagent for softening paper pulp.
Sulphur trioxide (SO3). When sulphur dioxide and oxygen are heated together at a rather high
temperature, a small amount of sulphur trioxide (SO3) is formed, but the reaction is slow and
incomplete. If, however, the heating takes place in the presence of very fine platinum dust, the
reaction is rapid and nearly complete.

Fig. 43
Experimental preparation of sulphur trioxide. The experiment can be performed by the use of
the apparatus shown in Fig. 43, the fine platinum being secured by moistening asbestos fiber with
a solution of platinum chloride and igniting it in a flame. The fiber, covered with fine platinum, is
placed in a tube of hard glass, which is then heated with a burner to about 350°, while sulphur
dioxide and air are passed into the tube. Union takes place at once, and the strongly fuming
sulphur trioxide escapes from the jet at the end of the tube, and may be condensed by surrounding
the receiving tube with a freezing mixture.
Properties of sulphur trioxide. Sulphur trioxide is a colorless liquid, which solidifies at about 15°
and boils at 46°. A trace of moisture causes it to solidify into a mass of silky white crystals,

somewhat resembling asbestos fiber in appearance. In contact with the air it fumes strongly, and
when thrown upon water it dissolves with a hissing sound and the liberation of a great deal of heat.
The product of this reaction is sulphuric acid, so that sulphur trioxide is the anhydride of that acid:

SO3 + H2O = H2SO4.

Catalysis. It has been found that many chemical reactions, such as the union of sulphur dioxide
with oxygen, are much influenced by the presence of substances which do not themselves seem to
take a part in the reaction, and are left apparently unchanged after it has ceased. These reactions go
on very slowly under ordinary circumstances, but are greatly hastened by the presence of the
foreign substance. Substances which hasten very slow reactions in this way are said to act as
catalytic agents or catalyzers, and the action is called catalysis. Just how the action is brought
about is not well understood.
DEFINITION: A catalyzer is a substance which changes the velocity of a reaction, but does not
change its products.
Examples of Catalysis. We have already had several instances of such action. Oxygen and
hydrogen combine with each other at ordinary temperatures in the presence of platinum powder,
while if no catalytic agent is present they do not combine in appreciable quantities until a rather
high temperature is reached. Potassium chlorate, when heated with manganese dioxide, gives up
its oxygen at a much lower temperature than when heated alone. Hydrogen dioxide decomposes
very rapidly when powdered manganese dioxide is sifted into its concentrated solution.
On the other hand, the catalytic agent sometimes retards chemical action. For example, a solution
of hydrogen dioxide decomposes more slowly when it contains a little phosphoric acid than when
perfectly pure. For this reason commercial hydrogen dioxide always contains phosphoric acid.
Many reactions are brought about by the catalytic action of traces of water. For example,
phosphorus will not burn in oxygen in the absence of all moisture. Hydrochloric acid will not unite
with ammonia if the reagents are perfectly dry. It is probable that many of the chemical
transformations in physiological processes, such as digestion, are assisted by certain substances
acting as catalytic agents. The principle of catalysis is therefore very important.
Sulphuric acid (oil of vitriol) (H2SO4). Sulphuric acid is one of the most important of all
manufactured chemicals. Not only is it one of the most common reagents in the laboratory, but
enormous quantities of it are used in many of the industries, especially in the refining of
petroleum, the manufacture of nitroglycerin, sodium carbonate, and fertilizers.
Manufacture of sulphuric acid. 1. Contact process. The reactions taking place in this process are
represented by the following equations:

SO2 + O = SO3,

SO3 + H2O = H2SO4.

To bring about the first of these reactions rapidly, a catalyzer is employed, and the process is
carried out in the following way: Large iron tubes are packed with some porous material, such as
calcium and magnesium sulphates, which contains a suitable catalytic substance scattered through
it. The catalyzers most used are platinum powder, vanadium oxide, and iron oxide. Purified

sulphur dioxide and air are passed through the tubes, which are kept at a temperature of about
350°. Sulphur trioxide is formed, and as it issues from the tube it is absorbed in water or dilute
sulphuric acid. The process is continued until all the water in the absorbing vessel has been
changed into sulphuric acid, so that a very concentrated acid is made in this way. An excess of the
trioxide may dissolve in the strong sulphuric acid, forming what is known as fuming sulphuric
acid.
2. Chamber process. The method of manufacture exclusively employed until recent years, and still
in very extensive use, is much more complicated. The reactions are quite involved, but the
conversion of water, sulphur dioxide, and oxygen into sulphuric acid is accomplished by the
catalytic action of oxides of nitrogen. The reactions are brought about in large lead-lined
chambers, into which oxides of nitrogen, sulphur dioxide, steam, and air are introduced in suitable
proportions.
Reactions of the chamber process. In a very general way, the various reactions which take place
in the lead chambers may be expressed in two equations. In the first reaction sulphur dioxide,
nitrogen peroxide, steam, and oxygen unite, as shown in the equation
(1) 2SO2 + 2NO2 + H2O + O = 2SO2 (OH) (NO2).

The product formed in this reaction is called nitrosulphuric acid or "chamber crystals." It actually
separates on the walls of the chambers when the process is not working properly. Under normal
conditions, it is decomposed as fast as it is formed by the action of excess of steam, as shown in
the equation
(2) 2SO2 (OH) (NO2) + H2O + O = 2H2SO4 + 2NO2.

The nitrogen dioxide formed in this reaction can now enter into combination with a new quantity
of sulphur dioxide, steam, and oxygen, and the series of reactions go on indefinitely. Many other
reactions occur, but these two illustrate the principle of the process.
The relation between sulphuric acid and nitrosulphuric acid can be seen by comparing their
structural formulas:

O= -OH O= -OH
S S
O= -OH O= -NO2

The latter may be regarded as derived from the former by the substitution of the nitro group (NO2)
for the hydroxyl group (OH).

Fig. 44.
The sulphuric acid plant. Fig. 44 illustrates the simpler parts of a plant used in the manufacture
of sulphuric acid by the chamber process. Sulphur or some sulphide, as FeS2, is burned in furnace
A. The resulting sulphur dioxide, together with air and some nitrogen peroxide, are conducted into
the large chambers, the capacity of each chamber being about 75,000 cu. ft. Steam is also admitted
into these chambers at different points. These compounds react to form sulphuric acid, according
to the equations given above. The nitrogen left after the withdrawal of the oxygen from the
admitted air escapes through the Gay-Lussac tower X. In order to prevent the escape of the oxides
of nitrogen regenerated in the reaction, the tower is filled with lumps of coke, over which trickles
concentrated sulphuric acid admitted from Y. The nitrogen peroxide dissolves in the acid and the
resulting solution collects in H. This is pumped into E, where it is mixed with dilute acid and
allowed to trickle down through the chamber D (Glover tower), which is filled with some acid-
resisting rock. Here the nitrogen peroxide is expelled from the solution by the action of the hot
gases entering from A, and together with them enters the first chamber again. The acid from which
the nitrogen peroxide is expelled collects in F. Theoretically, a small amount of nitrogen peroxide
would suffice to prepare an unlimited amount of sulphuric acid; practically, some of it escapes, and
this is replaced by small amounts admitted at B.
The sulphuric acid so formed, together with the excess of condensed steam, collect upon the floor
of the chambers in the form of a liquid containing from 62% to 70% of sulphuric acid. The product
is called chamber acid and is quite impure; but for many purposes, such as the manufacture of
fertilizers, it needs no further treatment. It can be concentrated by boiling it in vessels made of iron
or platinum, which resist the action of the acid, nearly all the water boiling off. Pure concentrated
acid can be made best by the contact process, while the chamber process is cheaper for the dilute
impure acid.
Physical properties. Sulphuric acid is a colorless, oily liquid, nearly twice as heavy as water. The
ordinary concentrated acid contains about 2% of water, has a density of 1.84, and boils at 338°. It
is sometimes called oil of vitriol, since it was formerly made by distilling a substance called green
vitriol.
Chemical properties. Sulphuric acid possesses chemical properties which make it one of the most
important of chemical substances.
1. Action as an acid. In dilute solution sulphuric acid acts as any other acid, forming salts with
oxides and hydroxides.

2. Action as an oxidizing agent. Sulphuric acid contains a large percentage of oxygen and is, like
nitric acid, a very good oxidizing agent. When the concentrated acid is heated with sulphur,
carbon, and many other substances, oxidation takes place, the sulphuric acid decomposing
according to the equation

H2SO4 = H2SO3 + O.

3. Action on metals. In dilute solution sulphuric acid acts upon many metals, such as zinc, forming
a sulphate and liberating hydrogen. When the concentrated acid is employed the hydrogen set free
is oxidized by a new portion of the acid, with the liberation of sulphur dioxide. With copper the
reactions are expressed by the equations

(1) Cu + H2SO4 = CuSO4 + 2H,

(2) H2SO4 + 2H = H2SO3 + H2O,

(3) H2SO3 = H2O + SO2.

By combining these equations the following one is obtained:

Cu + 2H2SO4 = CuSO4 + SO2 + 2H2O.

4. Action on salts. We have repeatedly seen that an acid of high boiling point heated with the salt
of some acid of lower boiling point will drive out the low boiling acid. The boiling point of
sulphuric acid (338°) is higher than that of almost any common acid; hence it is used largely in the
preparation of other acids.
5. Action on water. Concentrated sulphuric acid has a very great affinity for water, and is therefore
an effective dehydrating agent. Gases which have no chemical action upon sulphuric acid can be
freed from water vapor by bubbling them through the strong acid. When the acid is diluted with
water much heat is set free, and care must be taken to keep the liquid thoroughly stirred during the
mixing, and to pour the acid into the water,—never the reverse.
Not only can sulphuric acid absorb water, but it will often withdraw the elements hydrogen and
oxygen from a compound containing them, decomposing the compound, and combining with the
water so formed. For this reason most organic substances, such as sugar, wood, cotton, and woolen
fiber, and even flesh, all of which contain much oxygen and hydrogen in addition to carbon, are
charred or burned by the action of the concentrated acid.
Salts of sulphuric acid,—sulphates. The sulphates form a very important class of salts, and many
of them have commercial uses. Copperas (iron sulphate), blue vitriol (copper sulphate), and Epsom
salt (magnesium sulphate) serve as examples. Many sulphates are important minerals, prominent
among these being gypsum (calcium sulphate) and barytes (barium sulphate).
Thiosulphuric acid (H2S2O3); Thiosulphates. Many other acids of sulphur containing oxygen are
known, but none of them are of great importance. Most of them cannot be prepared in a pure state,
and are known only through their salts. The most important of these is thiosulphuric acid.
When sodium sulphite is boiled with sulphur the two substances combine, forming a salt which
has the composition represented in the formula Na2S2O3:

Na2SO3 + S = Na2S2O3.

The substance is called sodium thiosulphate, and is a salt of the easily decomposed acid H2S2O3,
called thiosulphuric acid. This reaction is quite similar to the action of oxygen upon sulphites:
Na2SO3 + O = Na2SO4.

More commonly the salt is called sodium hyposulphite, or merely "hypo." It is a white solid and is
extensively used in photography, in the bleaching industry, and as a disinfectant.
Monobasic and dibasic acids. Such acids as hydrochloric and nitric acids, which have only one
replaceable hydrogen atom in the molecule, or in other words yield one hydrogen ion in solution,
are called monobasic acids. Acids yielding two hydrogen ions in solution are called dibasic acids.
Similarly, we may have tribasic and tetrabasic acids. The three acids of sulphur are dibasic acids. It
is therefore possible for each of them to form both normal and acid salts. The acid salts can be
made in two ways: the acid may be treated with only half enough base to neutralize it,—

NaOH + H2SO4 = NaHSO4 + H2O;

or a normal salt may be treated with the free acid,—

Na2SO4 + H2SO4 = 2NaHSO4.

Acid sulphites and sulphides may be made in the same ways.
Carbon disulphide (CS2). When sulphur vapor is passed over highly heated carbon the two
elements combine, forming carbon disulphide (CS2), just as oxygen and carbon unite to form
carbon dioxide (CO2). The substance is a heavy, colorless liquid, possessing, when pure, a pleasant
ethereal odor. On standing for some time, especially when exposed to sunlight, it undergoes a
slight decomposition and acquires a most disagreeable, rancid odor. It has the property of
dissolving many substances, such as gums, resins, and waxes, which are insoluble in most liquids,
and it is extensively used as a solvent for such substances. It is also used as an insecticide. It boils
at a low temperature (46°), and its vapor is very inflammable, burning in the air to form carbon
dioxide and sulphur dioxide, according to the equation

CS2 + 6O = CO2 + 2SO2.

Commercial preparation of carbon disulphide. In the preparation of carbon disulphide an
electrical furnace is employed, such as is represented in Fig. 45. The furnace is packed with carbon
C, and this is fed in through the hoppers B, as fast as that which is present in the hearth of the
furnace is used up. Sulphur is introduced at A, and at the lower ends of the tubes it is melted by the
heat of the furnace and flows into the hearth as a liquid. An electrical current is passed through the
carbon and melted sulphur from the electrodes E, heating the charge. The vapors of carbon
disulphide pass up through the furnace and escape at D, from which they pass to a suitable
condensing apparatus.
Comparison of sulphur and oxygen. A comparison of the formulas and the chemical properties
of corresponding compounds of oxygen and sulphur brings to light many striking similarities. The
conduct of hydrosulphuric acid and water toward many substances has been seen to be very

similar; the oxides and sulphides of the metals
have analogous formulas and undergo many
parallel reactions. Carbon dioxide and
disulphide are prepared in similar ways and
undergo many analogous reactions. It is clear,
therefore, that these two elements are far more
closely related to each other than to any of the
other elements so far studied.
Selenium and tellurium. These two very
uncommon elements are still more closely
related to sulphur than is oxygen. They occur
in comparatively small quantities and are
usually found associated with sulphur and
sulphides, either as the free elements or more
commonly in combination with metals. They
form compounds with hydrogen of the
formulas H2Se and H2Te; these bodies are
gases with properties very similar to those of
H2S. They also form oxides and oxygen acids Fig. 45
which resemble the corresponding sulphur
compounds. The elements even have allotropic forms corresponding very closely to those of
sulphur. Tellurium is sometimes found in combination with gold and copper, and occasions some
difficulties in the refining of these metals. The elements have very few practical applications.
Crystallography. In order to understand the difference between the two kinds of sulphur crystals,
it is necessary to know something about crystals in general and the forms which they may assume.
An examination of a large number of crystals has shown that although they may differ much in
geometric form, they can all be considered as modifications of a few simple plans. The best way to
understand the relation of one crystal to another is to look upon every crystal as having its faces
and angles arranged in definite fashion about certain imaginary lines drawn through the crystal.
These lines are called axes, and bear much the same relation to a crystal as do the axis and
parallels of latitude and longitude to the earth and a geographical study of it. All crystals can be
referred to one of six simple plans or systems, which have their axes as shown in the following
drawings.
The names and characteristics of these systems are as follows:
1. Isometric or regular system (Fig. 46). Three equal axes, all at right angles.

Fig. 46

2. Tetragonal system (Fig. 47). Two equal axes and one of different length, all at right angles to
each other.

Fig. 47
3. Orthorhombic system (Fig. 48). Three unequal axes, all at right angles to each other.

Fig. 48
4. Monoclinic system (Fig. 49). Two axes at right angles, and a third at right angles to one of these,
but inclined to the other.

Fig. 49
5. Triclinic system (Fig. 50). Three axes, all inclined to each other.

Fig. 50
6. Hexagonal system (Fig. 51). Three equal axes in the same plane intersecting at angles of 60°,
and a fourth at right angles to all of these.

Fig. 51
Every crystal can be imagined to have its faces and angles arranged in a definite way around one
of these systems of axes. A cube, for instance, is referred to Plan 1, an axis ending in the center of
each face; while in a regular octohedron an axis ends in each solid angle. These forms are shown
in Fig. 46. It will be seen that both of these figures belong to the same system, though they are
very different in appearance. In the same way, many geometric forms may be derived from each of
the systems, and the light lines about the axes in the drawings show two of the simplest forms of
each of the systems.
In general a given substance always crystallizes in the same system, and two corresponding faces
of each crystal of it always make the same angle with each other. A few substances, of which
sulphur is an example, crystallize in two different systems, and the crystals differ in such physical
properties as melting point and density. Such substances are said to be dimorphous.

EXERCISES

1. (a) Would the same amount of heat be generated by the combustion of 1 g. of each of the
allotropic modifications of sulphur? (b) Would the same amount of sulphur dioxide be formed in
each case?
2. Is the equation for the preparation of hydrosulphuric acid a reversible one? As ordinarily carried
out, does the reaction complete itself?

3. Suppose that hydrosulphuric acid were a liquid, would it be necessary to modify the method of
preparation?
4. Can sulphuric acid be used to dry hydrosulphuric acid? Give reason for answer.
5. Does dry hydrosulphuric acid react with litmus paper? State reason for answer.
6. How many grams of iron sulphide are necessary to prepare 100 l. of hydrosulphuric acid when
the laboratory conditions are 17° and 740 mm. pressure?
7. Suppose that the hydrogen in 1 l. of hydrosulphuric acid were liberated; what volume would it
occupy, the gases being measured under the same conditions?
8. Write the equations representing the reaction between hydrosulphuric acid and sodium
hydroxide and ammonium hydroxide respectively.
9. Show that the preparation of sulphur dioxide from a sulphite is similar in principle to the
preparation of hydrogen sulphide.
10. (a) Does dry sulphur dioxide react with litmus paper? (b) How can it be shown that a solution
of sulphur dioxide in water acts like an acid?
11. (a) Calculate the percentage composition of sulphurous anhydride and sulphuric anhydride. (b)
Show how these two substances are in harmony with the law of multiple proportion.
12. How many pounds of sulphur would be necessary in the preparation of 100 lb. of 98%
sulphuric acid?
13. What weight of sulphur dioxide is necessary in the preparation of 1 kg. of sodium sulphite?
14. What weight of copper sulphate crystals can be obtained by dissolving 1 kg. of copper in
sulphuric acid and crystallizing the product from water?
15. Write the names and formulas of the oxides and oxygen acids of selenium and tellurium.
16. In the commercial preparation of carbon disulphide, what is the function of the electric
current?
17. If the Gay-Lussac tower were omitted from the sulphuric acid factory, what effect would this
have on the cost of production of sulphuric acid?

CHAPTER XV
PERIODIC LAW

A number of the elements have now been studied somewhat closely. The first three of these,
oxygen, hydrogen, and nitrogen, while having some physical properties in common with each
other, have almost no point of similarity as regards their chemical conduct. On the other hand,
oxygen and sulphur, while quite different physically, have much in common in their chemical
properties.
About eighty elements are now known. If all of these should have properties as diverse as do
oxygen, hydrogen, and nitrogen, the study of chemistry would plainly be a very difficult and
complicated one. If, however, the elements can be classified in groups, the members of which have
very similar properties, the study will be very much simplified.
Earlier classification of the elements. Even at an early period efforts were made to discover some
natural principle in accordance with which the elements could be classified. Two of these
classifications may be mentioned here.
1. Classification into metals and non-metals. The classification into metals and non-metals most
naturally suggested itself. This grouping was based largely on physical properties, the metals being
heavy, lustrous, malleable, ductile, and good conductors of heat and electricity. Elements
possessing these properties are usually base-forming in character, and the ability to form bases
came to be regarded as a characteristic property of the metals. The non-metals possessed physical
properties which were the reverse of those of the metals, and were acid-forming in character.
Not much was gained by this classification, and it was very imperfect. Some metals, such as
potassium, are very light; some non-metals, such as iodine, have a high luster; some elements can
form either an acid or a base.
2. Classification into triad families. In 1825 Döbereiner observed that an interesting relation exists
between the atomic weights of chemically similar elements. To illustrate, lithium, sodium, and
potassium resemble each other very closely, and the atomic weight of sodium is almost exactly an
arithmetical mean between those of the other two: (7.03 + 39.15)/2 = 23.09. In many chemical and
physical properties sodium is midway between the other two.
A number of triad families were found, but among eighty elements, whose atomic weights range
all the way from 1 to 240, such agreements might be mere chance. Moreover many elements did
not appear to belong to such families.
Periodic division. In 1869 the Russian chemist Mendeléeff devised an arrangement of the
elements based on their atomic weights, which has proved to be of great service in the comparative
study of the elements. A few months later the German, Lothar Meyer, independently suggested the
same ideas. This arrangement brought to light a great generalization, now known as the periodic
law. An exact statement of the law will be given after the method of arranging the elements has
been described.

DMITRI IVANOVITCH MENDELÉEFF
(Russian) (1834-1907)

Author of the periodic law; made many
investigations on the physical constants of
elements and compounds; wrote an important
book entitled "Principles of Chemistry";
university professor and government official
Arrangement of the periodic table. The arrangement suggested by Mendeléeff, modified
somewhat by more recent investigations, is as follows: Beginning with lithium, which has an
atomic weight of 7, the elements are arranged in a horizontal row in the order of their atomic
weights, thus:
Li (7.03), Be (9.1), B (11), C (12), N (14.04), O (16), F (19).
These seven elements all differ markedly from each other. The eighth element, sodium, is very
similar to lithium. It is placed just under lithium, and a new row follows:
Na(23.05), Mg (24.36), Al (27.1), Si (28.4), P (31), S (32.06), Cl(35.45).
When the fifteenth element, potassium, is reached, it is placed under sodium, to which it is very
similar, and serves to begin a third row:
K (39.15), Ca (40.1), Sc (44.1,) Ti (48.1), V (51.2), Cr (52.1), Mn(55).

Not only is there a strong similarity between lithium, sodium, and potassium, which have been
placed in a vertical row because of this resemblance, but the elements in the other vertical rows
exhibit much of the same kind of similarity among themselves, and evidently form little natural
groups.
The three elements following manganese, namely, iron, nickel, and cobalt, have atomic weights
near together, and are very similar chemically. They do not strongly resemble any of the elements
so far considered, and are accordingly placed in a group by themselves, following manganese. A
new row is begun with copper, which somewhat resembles the elements of the first vertical
column. Following the fifth and seventh rows are groups of three closely related elements, so that
the completed arrangement has the appearance represented in the table on page 168.

THE PERIODIC ARRANGEMENT OF THE ELEMENTS
Place of the atmospheric elements. When argon was discovered it was seen at once that there
was no place in the table for an element of atomic weight approximately 40. When the other
inactive elements were found, however, it became apparent that they form a group just preceding
Group 1. They are accordingly arranged in this way in Group 0 (see table on opposite page). A
study of this table brings to light certain very striking facts.
Properties of elements vary with atomic weights. There is evidently a close relation between the
properties of an element and its atomic weight. Lithium, at the beginning of the first group, is a
very strong base-forming element, with pronounced metallic properties. Beryllium, following
lithium, is less strongly base-forming, while boron has some base-forming and some acid-forming
properties. In carbon all base-forming properties have disappeared, and the acid-forming properties

are more marked than in boron. These become still more emphasized as we pass through nitrogen
and oxygen, until on reaching fluorine we have one of the strongest acid-forming elements. The
properties of these seven elements therefore vary regularly with their atomic weights, or, in
mathematical language, are regular functions of them.
Periodic law. The properties of the first seven elements vary continuously—that is steadily—away
from base-forming and toward acid-forming properties. If lithium had the smallest atomic weight
of any of the elements, and fluorine the greatest, so that in passing from one to the other we had
included all the elements, we could say that the properties of elements are continuous functions of
their atomic weights. But fluorine is an element of small atomic weight, and the one following it,
sodium, breaks the regular order, for in it reappear all the characteristic properties of lithium.
Magnesium, following sodium, bears much the same relation to beryllium that sodium does to
lithium, and the properties of the elements in the second row vary much as they do in the first row
until potassium is reached, when another repetition begins. The properties of the elements do not
vary continuously, therefore, with atomic weights, but at regular intervals there is a repetition, or
period. This generalization is known as the periodic law, and may be stated thus: The properties of
elements are periodic functions of their atomic weights.
The two families in a group. While all the elements in a given vertical column bear a general
resemblance to each other, it has been noticed that those belonging to periods having even
numbers are very strikingly similar to each other. They are placed at the left side of the group
columns. In like manner, the elements belonging to the odd periods are very similar and are
arranged at the right side of the group columns. Thus calcium, strontium, and barium are very
much alike; so, too, are magnesium, zinc, and cadmium. The resemblance between calcium and
magnesium, or strontium and zinc, is much less marked. This method of arrangement therefore
divides each group into two families, each containing four or five members, between which there
is a great similarity.
Family resemblances. Let us now inquire more closely in what respects the elements of a family
resemble each other.
1. Valence. In general the valence of the elements in a family is the same, and the formulas of their
compounds are therefore similar. If we know that the formula of sodium chloride is NaCl, it is
pretty certain that the formula of potassium chloride will be KCl—not KCl2 or KCl3. The general
formulas R2O, RO, etc., placed below the columns show the formulas of the oxides of the
elements in the column provided they form oxides. In like manner the formulas RH, RH2, etc.,
show the composition of the compounds formed with hydrogen or chlorine.
2. Chemical properties. The chemical properties of the members of a family are quite similar. If
one member is a metal, the others usually are; if one is a non-metal, so, too, are the others. The
families in the first two columns consist of metals, while the elements found in the last two
columns form acids. There is in addition a certain regularity in properties of the elements in each
family. If the element at the head of the family is a strong acid-forming element, this property is
likely to diminish gradually, as we pass to the members of the family with higher atomic weights.
Thus phosphorus is strongly acid-forming, arsenic less so, antimony still less so, while bismuth
has almost no acid-forming properties. We shall meet with many illustrations of this fact.
3. Physical properties. In the same way, the physical properties of the members of a family are in
general somewhat similar, and show a regular gradation as we pass from element to element in the
family. Thus the densities of the members of the magnesium family are

Mg = 1.75, Zn = 7.00, Cd = 8.67, Hg = 13.6.

Their melting points are

Mg = 750°, Zn = 420°, Cd = 320°, Hg = -39.5°.

Value of the periodic law. The periodic law has proved of much value in the development of the
science of chemistry.
1. It simplifies study. It is at once evident that such regularities very much simplify the study of
chemistry. A thorough study of one element of a family makes the study of the other members a
much easier task, since so many of the properties and chemical reactions of the elements are
similar. Thus, having studied the element sulphur in some detail, it is not necessary to study
selenium and tellurium so closely, for most of their properties can be predicted from the relation
which they sustain to sulphur.
2. It predicts new elements. When the periodic law was first formulated there were a number of
vacant places in the table which evidently belonged to elements at that time unknown. From their
position in the table, Mendeléeff predicted with great precision the properties of the elements
which he felt sure would one day be discovered to fill these places. Three of them, scandium,
germanium, and gallium, were found within fifteen years, and their properties agreed in a
remarkable way with the predictions of Mendeléeff. There are still some vacant places in the table,
especially among the heavier elements.
3. It corrects errors. The physical constants of many of the elements did not at first agree with
those demanded by the periodic law, and a further study of many such cases showed that errors
had been made. The law has therefore done much service in indicating probable error.
Imperfections of the law. There still remain a good many features which must be regarded as
imperfections in the law. Most conspicuous is the fact that the element hydrogen has no place in
the table. In some of the groups elements appear in one of the families, while all of their properties
show that they belong in the other. Thus sodium belongs with lithium and not with copper; fluorine
belongs with chlorine and not with manganese. There are two instances where the elements must
be transposed in order to make them fit into their proper group. According to their atomic weights,
tellurium should follow iodine, and argon should follow potassium. Their properties show in each
case that this order must be reversed. The table separates some elements altogether which, in many
respects have closely agreeing properties. Iron, chromium, and manganese are all in different
groups, although they are similar in many respects.
The system is therefore to be regarded as but a partial and imperfect expression of some very
important and fundamental relation between the substances which we know as elements, the exact
nature of this relation being as yet not completely clear to us.

EXERCISES

1. Suppose that an element were discovered that filled the blank in Group O, Period 5; what
properties would it probably have?
2. Suppose that an element were discovered that filled the blank in Group VI, Period 9, family B;
what properties would it have?

3. Sulphur and oxygen both belong in Group VI, although in different families; in what respects
are the two similar?

CHAPTER XVI
THE CHLORINE FAMILY

ATOMIC WEIGHT MELTING POINT BOILING POINT COLOR AND STATE
Fluorine (F) 19.00 -223° -187° Pale yellowish gas.
Chlorine (Cl) 35.45 -102° -33.6° Greenish-yellow gas.
Bromine (Br) 79.96 -7° 59° Red liquid.
Iodine (I) 126.97 107° 175° Purplish-black solid.

The family. The four elements named in the above table form a strongly marked family of
elements and illustrate very clearly the way in which the members of a family in a periodic group
resemble each other, as well as the character of the differences which we may expect to find
between the individual members.
1. Occurrence. These elements do not occur in nature in the free state. The compounds of the last
three elements of the family are found extensively in sea water, and on this account the name
halogens, signifying "producers of sea salt," is sometimes applied to the family.
2. Properties. As will be seen by reference to the table, the melting points and boiling points of the
elements of the family increase with their atomic weights. A somewhat similar gradation is noted
in their color and state. One atom of each of the elements combines with one atom of hydrogen to
form acids, which are gases very soluble in water. The affinity of the elements for hydrogen is in
the inverse order of their atomic weights, fluorine having the strongest affinity and iodine the
weakest. Only chlorine and iodine form oxides, and those of the former element are very unstable.
The elements of the group are univalent in their compounds with hydrogen and the metals.

FLUORINE

Occurrence. The element fluorine occurs in nature most abundantly as the mineral fluorspar
(CaF2), as cryolite (Na3AlF6), and in the complex mineral apatite (3 Ca3(PO4)2·CaF2).

Preparation. All attempts to isolate the element resulted in failure until recent years. Methods
similar to those which succeed in the preparation of the other elements of the family cannot be
used; for as soon as the fluorine is liberated it combines with the materials of which the apparatus
is made or with the hydrogen of the water which is always present. The preparation of fluorine was
finally accomplished by the French chemist Moissan by the electrolysis of hydrofluoric acid.
Perfectly dry hydrofluoric acid (HF) was condensed to a liquid and placed in a U-shaped tube
made of platinum (or copper), which was furnished with electrodes and delivery tubes, as shown in
Fig. 52. This liquid is not an electrolyte, but becomes such when potassium fluoride is dissolved in
it. When this solution was electrolyzed hydrogen was set free at the cathode and fluorine at the
anode.
Properties. Fluorine is a gas of slightly yellowish color, and can be condensed to a liquid boiling
at -187° under atmospheric pressure. It solidifies at -223°. It is extremely active chemically, being
the most active of all the elements at ordinary temperatures.

It combines with all the common elements save oxygen, very
often with incandescence and the liberation of much heat. It has a
strong affinity for hydrogen and is able to withdraw it from its
compounds with other elements. Because of its great activity it is
extremely poisonous. Fluorine does not form any oxides, neither
does it form any oxygen acids, in which respects it differs from
the other members of the family.
Hydrofluoric acid (HF). Hydrofluoric acid is readily obtained
from fluorspar by the action of concentrated sulphuric acid. The
equation is

CaF2 + H2SO4 = CaSO4 + 2HF.

In its physical properties it resembles the binary acids of the
other elements of this family, being, however, more easily
condensed to a liquid. The anhydrous acid boils at 19° and can
therefore be prepared at ordinary pressures. It is soluble in all
proportions in water, and a concentrated solution—about 50%—
Fig. 52 is prepared for the market. Its fumes are exceedingly irritating to
the respiratory organs, and several chemists have lost their lives
by accidentally breathing them.

HENRI MOISSAN (French) (1853-1907)

Famous for his work with the electric furnace
at high temperatures; prepared artificial
diamonds, together with many new binary

compounds such as carbides, silicides, borides,
and nitrides; isolated fluorine and studied its
properties and its compounds very thoroughly
Chemical properties. Hydrofluoric acid, like other strong acids, readily acts on bases and metallic
oxides and forms the corresponding fluorides. It also dissolves certain metals such as silver and
copper. It acts very vigorously upon organic matter, a single drop of the concentrated acid making
a sore on the skin which is very painful and slow in healing. Its most characteristic property is its
action upon silicon dioxide (SiO2), with which it forms water and the gas silicon tetrafluoride
(SiF4), as shown in the equation

SiO2 + 4HF = SiF4 + 2H2O.

Glass consists of certain compounds of silicon, which are likewise acted on by the acid so that it
cannot be kept in glass bottles. It is preserved in flasks made of wax or gutta-percha.
Etching. Advantage is taken of this reaction in etching designs upon glass. The glass vessel is
painted over with a protective paint upon which the acid will not act, the parts which it is desired
to make opaque being left unprotected. A mixture of fluorspar and sulphuric acid is then painted
over the vessel and after a few minutes the vessel is washed clean. Wherever the hydrofluoric acid
comes in contact with the glass it acts upon it, destroying its luster and making it opaque, so that
the exposed design will be etched upon the clear glass. Frosted glass globes are often made in this
way.
The etching may also be effected by covering the glass with a thin layer of paraffin, cutting the
design through the wax and then exposing the glass to the fumes of the acid.
Salts of hydrofluoric acid,—fluorides. A number of the fluorides are known, but only one of
them, calcium fluoride (CaF2), is of importance. This is the well-known mineral fluorspar.

CHLORINE

Historical. While studying the action of hydrochloric acid upon the mineral pyrolusite, in 1774,
Scheele obtained a yellowish, gaseous substance to which he gave a name in keeping with the
phlogiston theory then current. Later it was supposed to be a compound containing oxygen. In
1810, however, the English chemist Sir Humphry Davy proved it to be an element and named it
chlorine.
Occurrence. Chlorine does not occur free in nature, but its compounds are widely distributed. For
the most part it occurs in combination with the metals in the form of chlorides, those of sodium,
potassium, and magnesium being most abundant. Nearly all salt water contains these substances,
particularly sodium chloride, and very large salt beds consisting of chlorides are found in many
parts of the world.
Preparation. Two general methods of preparing chlorine may be mentioned, namely, the
laboratory method and the electrolytic method.
1. Laboratory method. In the laboratory chlorine is made by warming the mineral pyrolusite
(manganese dioxide, MnO2) with concentrated hydrochloric acid. The first reaction, which seems
to be similar to the action of acids upon oxides in general, is expressed in the equation

MnO2 + 4HCl = MnCl4 + 2H2O.

The manganese compound so formed is very unstable, however, and breaks clown according to the
equation

MnCl4 = MnCl2 + 2Cl.

Instead of using hydrochloric acid in the preparation of chlorine it will serve just as well to use a
mixture of sodium chloride and sulphuric acid, since these two react to form hydrochloric acid.
The following equations will then express the changes:

(1) 2NaCl + H2SO4 = Na2SO4 + 2HCl.

(2) MnO2 + 4 HCl = MnCl2 + 2Cl + 2H2O.

(3) MnCl2 + H2SO4 = MnSO4 + 2HCl.

Combining these equations, the following equation expressing the complete reaction is obtained:

2NaCl + MnO2 + 2H2SO4 = MnSO4 + Na2SO4 + 2H2O + 2Cl.

Since the hydrochloric acid liberated in the third equation is free to act upon manganese dioxide, it
will be seen that all of the chlorine originally present in the sodium chloride is set free.
The manganese dioxide and the hydrochloric acid are brought together in a flask, as represented in
Fig. 53, and a gentle heat is applied. The rate of evolution of the gas is regulated by the amount of
heat applied, and the gas is collected by displacement of air. As the equations show, only half of
the chlorine present in the hydrochloric acid is liberated.

Fig. 53

2. Electrolytic method. Under the discussion of electrolysis (p. 102) it was shown that when a
solution of sodium chloride is electrolyzed chlorine is evolved at the anode, while the sodium set
free at the cathode reacts with the water to form hydrogen, which is evolved, and sodium
hydroxide, which remains in solution. A great deal of the chlorine required in the chemical
industries is now made in this way in connection with the manufacture of sodium hydroxide.
Physical properties. Chlorine is a greenish-yellow gas, which has a peculiar suffocating odor and
produces a very violent effect upon the throat and lungs. Even when inhaled in small quantities it
often produces all the symptoms of a hard cold, and in larger quantities may have serious and even
fatal action. It is quite heavy (density = 2.45) and can therefore be collected by displacement of air.
One volume of water under ordinary conditions dissolves about three volumes of chlorine. The gas
is readily liquefied, a pressure of six atmospheres serving to liquefy it at 0°. It forms a yellowish
liquid which solidifies at -102°.
Chemical properties. At ordinary temperatures chlorine is far more active chemically than any of
the elements we have so far considered, with the exception of fluorine; indeed, it is one of the most
active of all elements.
1. Action on metals. A great many metals combine directly with chlorine, especially when hot. A
strip of copper foil heated in a burner flame and then dropped into chlorine burns with
incandescence. Sodium burns brilliantly when heated strongly in slightly moist chlorine. Gold and
silver are quickly tarnished by the gas.
2. Action on non-metals. Chlorine has likewise a strong affinity for many of the non-metals. Thus
phosphorus burns in a current of the gas, while antimony and arsenic in the form of a fine powder
at once burst into flame when dropped into jars of the gas. The products formed in all cases where
chlorine combines with another element are called chlorides.
3. Action on hydrogen. Chlorine has a strong affinity for hydrogen, uniting with it to form
hydrochloric acid. A jet of hydrogen burning in the air continues to burn when introduced into a jar
of chlorine, giving a somewhat luminous flame. A mixture of the two gases explodes violently
when a spark is passed through it or when it is exposed to bright sunlight. In the latter case it is the
light and not the heat which starts the action.
4. Action on substances containing hydrogen. Not only will chlorine combine directly with free
hydrogen but it will often abstract the element from its compounds. Thus, when chlorine is passed
into a solution containing hydrosulphuric acid, sulphur is precipitated and Hydrochloric acid
formed. The reaction is shown by the following equation:

H2S + 2Cl = 2HCl + S.

With ammonia the action is similar:

NH3 + 3Cl = 3HCl + N.

The same tendency is very strikingly seen in the action of chlorine upon turpentine. The latter
substance is largely made up of compounds having the composition represented by the formula
C10H16. When a strip of paper moistened with warm turpentine is placed in a jar of chlorine dense
fumes of hydrochloric acid appear and a black deposit of carbon is formed. Even water, which is a

very stable compound, can be decomposed by chlorine, the oxygen being liberated. This may be
shown in the following way:
If a long tube of rather large diameter is filled with a strong solution of
chlorine in water and inverted in a vessel of the same solution, as shown in
Fig. 54, and the apparatus is placed in bright sunlight, very soon bubbles of a
gas will be observed to rise through the solution and collect in the tube. An
examination of this gas will show that it is oxygen. It is liberated from water in
accordance with the following equation:
H2O + 2Cl = 2HCl + O.

5. Action on color substances,—bleaching action. If strips of brightly colored
cloth or some highly colored flowers are placed in quite dry chlorine, no
marked change in color is noticed as a rule. If, however, the cloth and flowers
are first moistened, the color rapidly disappears, that is, the objects are
bleached. Evidently the moisture as well as the chlorine is concerned in the
action, and a study of the case shows that the chlorine has combined with the
hydrogen of the water. The oxygen set free oxidizes the color substance,
converting it into a colorless compound. It is evident from this explanation that
chlorine will only bleach those substances which are changed into colorless
compounds by oxidation.
6. Action as a disinfectant. Chlorine has also marked germicidal properties,
and the free element, as well as compounds from which it is easily liberated,
Fig. 54
are used as disinfectants.
Nascent state. It will be noticed that oxygen when set free from water by chlorine is able to do
what ordinary oxygen cannot do, for both the cloth and the flowers are unchanged in the air which
contains oxygen. It is generally true that the activity of an element is greatest at the instant of
liberation from its compounds. To express this fact elements at the instant of liberation are said to
be in the nascent state. It is nascent oxygen which does the bleaching.
Hydrochloric acid (muriatic acid) (HCl). The preparation of hydrochloric acid may be discussed
under two general heads:
1. Laboratory preparation. The product formed by the burning of hydrogen in chlorine is the gas
hydrochloric acid. This substance is much more easily obtained, however, by treating common salt
(sodium chloride) with sulphuric acid. The following equation shows the reaction:

2NaCl + H2SO4 = Na2SO4 + 2HCl.

The dry salt is placed in a flask furnished with a funnel tube and an exit tube, the sulphuric acid is
added, and the flask gently warmed. The hydrochloric acid gas is rapidly given off and can be
collected by displacement of air. The same apparatus can be used as was employed in the
preparation of chlorine (Fig. 53).
When a solution of salt is treated with sulphuric acid there is no very marked action. The
hydrochloric acid formed is very soluble in water, and so does not escape from the solution; hence
a state of equilibrium is soon reached between the four substances represented in the equation.

When concentrated sulphuric acid, in which hydrochloric acid is not soluble, is poured upon dry
salt the reaction is complete.
2. Commercial preparation. Commercially, hydrochloric acid is prepared in connection with the
manufacture of sodium sulphate, the reaction being the same as that just given. The reaction is
carried out in a furnace, and the hydrochloric acid as it escapes in the form of gas is passed into
water in which it dissolves, the solution forming the hydrochloric acid of commerce. When the
materials are pure a colorless solution is obtained. The most concentrated solution has a density of
1.2 and contains 40% HCl. The commercial acid, often called muriatic acid, is usually colored
yellow by impurities.
Composition of hydrochloric acid. When a solution of hydrochloric acid is electrolyzed in an
apparatus similar to the one in which water was electrolyzed (Fig. 18), chlorine collects at the
anode and hydrogen at the cathode. At first the chlorine dissolves in the water, but soon the water
in the one tube becomes saturated with it, and if the stopcocks are left open until this is the case,
and are then closed, it will be seen that the two gases are set free in equal volumes.
When measured volumes of the two gases are caused to unite it is found that one volume of
hydrogen combines with one of chlorine. Other experiments show that the volume of hydrochloric
acid formed is just equal to the sum of the volumes of hydrogen and chlorine. Therefore one
volume of hydrogen combines with one volume of chlorine to form two volumes of hydrochloric
acid gas. Since chlorine is 35.18 times as heavy as hydrogen, it follows that one part of hydrogen
by weight combines with 35.18 parts of chlorine to form 36.18 parts of hydrochloric acid.
Physical properties. Hydrochloric acid is a colorless gas which has an irritating effect when
inhaled, and possesses a sour, biting taste, but no marked odor. It is heavier than air (density =
1.26) and is very soluble in water. Under standard conditions 1 volume of water dissolves about
500 volumes of the gas. On warming such a solution the gas escapes, until at the boiling point the
solution contains about 20% by weight of HCl. Further boiling will not drive out any more acid,
but the solution will distill with unchanged concentration. A more dilute solution than this will lose
water on boiling until it has reached the same concentration, 20%, and will then distill unchanged.
Under high pressure the gas can be liquefied, 28 atmospheres being required at 0°. Under these
conditions it forms a colorless liquid which is not very active chemically. It boils at -80° and
solidifies at -113°. The solution of the gas in water is used almost entirely in the place of the gas
itself, since it is not only far more convenient but also more active.
Chemical properties. The most important chemical properties of hydrochloric acid are the
following:
1. Action as an acid. In aqueous solution hydrochloric acid has very strong acid properties; indeed,
it is one of the strongest acids. It acts upon oxides and hydroxides, converting them into salts:

NaOH + HCl = NaCl + H2O,
CuO + 2HCl = CuCl2 + H2O.

It acts upon many metals, forming chlorides and liberating hydrogen:

Zn + 2HCl = ZnCl2 + 2H,
Al + 3HCl = AlCl3 + 3H.

Unlike nitric and sulphuric acids it has no oxidizing action, so that when it acts on metals
hydrogen is always given off.
2. Relation to combustion. Hydrochloric acid gas is not readily decomposed, and is therefore
neither combustible nor a supporter of combustion.
3. Action on oxidizing agents. Although hydrochloric acid is incombustible, it can be oxidized
under some circumstances, in which case the hydrogen combines with oxygen, while the chlorine
is set free. Thus, when a solution of hydrochloric acid acts upon manganese dioxide part of the
chlorine is set free:

MnO2 + 4HCl = MnCl2 + 2H2O + 2Cl.

Aqua regia. It has been seen that when nitric acid acts as an oxidizing agent it usually
decomposes, as represented in the equation

2HNO3 = H2O + 2NO + 3O.

The oxygen so set free may act on hydrochloric acid:

6HCl + 3O = 3H2O + 6Cl.

The complete equation therefore is

2HNO3 + 6HCl = 4H2O + 2NO + 6Cl.

When concentrated nitric and hydrochloric acids are mixed this reaction goes on slowly, chlorine
and some other substances not represented in the equation being formed. The mixture is known as
aqua regia and is commonly prepared by adding one volume of nitric acid to three volumes of
hydrochloric acid. It acts more powerfully upon metals and other substances than either of the
acids separately, and owes its strength not to acid properties but to the action of the nascent
chlorine which it liberates. Consequently, when it acts upon metals such as gold it converts them
into chlorides, and the reaction can be represented by such equations as

Au + 3Cl = AuCl3.

Salts of hydrochloric acid,—chlorides. The chlorides of all the metals are known and many of
them are very important compounds. Some of them are found in nature, and all can be prepared by
the general method of preparing salts. Silver chloride, lead chloride, and mercurous chloride are
insoluble in water and acids, and can be prepared by adding hydrochloric acid to solutions of
compounds of the respective elements. While the chlorides have formulas similar to the fluorides,
their properties are often quite different. This is seen in the solubility of the salts. Those metals
whose chlorides are insoluble form soluble fluorides, while many of the metals which form soluble
chlorides form insoluble fluorides.
Compounds of chlorine with oxygen and hydrogen. Chlorine combines with oxygen and
hydrogen to form four different acids. They are all quite unstable, and most of them cannot be
prepared in pure form; their salts can easily be made, however, and some of them will be met with
in the study of the metals. The formulas and names of these acids are as follows:

HClO hypochlorous acid.
HClO2 chlorous acid.
HClO3 chloric acid.
HClO4 perchloric acid.

Oxides of chlorine. Two oxides are known, having the formulas Cl2O and ClO2. They decompose
very easily and are good oxidizing agents.

BROMINE

Historical. Bromine was discovered in 1826 by the French chemist Ballard, who isolated it from
sea salt. He named it bromine (stench) because of its unbearable fumes.
Occurrence. Bromine occurs almost entirely in the form of bromides, especially as sodium
bromide and magnesium bromide, which are found in many salt springs and salt deposits. The
Stassfurt deposits in Germany and the salt waters of Ohio and Michigan are especially rich in
bromides.
Preparation of bromine. The laboratory method of preparing bromine is essentially different
from the commercial method.

Fig. 55
1. Laboratory method. As in the case of chlorine, bromine can be prepared by the action of
hydrobromic acid (HBr) on manganese dioxide. Since hydrobromic acid is not an article of
commerce, a mixture of sulphuric acid and a bromide is commonly substituted for it. The materials
are placed in a retort arranged as shown in Fig. 55. The end of the retort just touches the surface of
the water in the test tube. On heating, the bromine distills over and is collected in the cold receiver.
The equation is

2NaBr + 2H2SO4 + MnO2 = Na2SO4 + MnSO4 + 2H2O + 2Br.

2. Commercial method. Bromine is prepared commercially from the waters of salt wells which are
especially rich in bromides. On passing a current of electricity through such waters the bromine is
first liberated. Any chlorine liberated, however, will assist in the reaction, since free chlorine
decomposes bromides, as shown in the equation

NaBr + Cl = NaCl + Br.

When the water containing the bromine is heated, the liberated bromine distills over into the
receiver.
Physical properties. Bromine is a dark red liquid about three times as heavy as water. Its vapor
has a very offensive odor and is most irritating to the eyes and throat. The liquid boils at 59° and
solidifies at -7°; but even at ordinary temperatures it evaporates rapidly, forming a reddish-brown
gas very similar to nitrogen peroxide in appearance. Bromine is somewhat soluble in water, 100
volumes of water under ordinary conditions dissolving 1 volume of the liquid. It is readily soluble
in carbon disulphide, forming a yellow solution.
Chemical properties and uses. In chemical action bromine is very similar to chlorine. It
combines directly with many of the same elements with which chlorine unites, but with less
energy. It combines with hydrogen and takes away the latter element from some of its compounds,
but not so readily as does chlorine. Its bleaching properties are also less marked.
Bromine finds many uses in the manufacture of organic drugs and dyestuffs and in the preparation
of bromides.
Hydrobromic acid (HBr). When sulphuric acid acts upon a bromide hydrobromic acid is set free:

2NaBr + H2SO4 = Na2SO4 + 2HBr.

At the same time some bromine is set free, as may be seen from the red fumes which appear, and
from the odor. The explanation of this is found in the fact that hydrobromic acid is much less
stable than hydrochloric acid, and is therefore more easily oxidized. Concentrated sulphuric acid is
a good oxidizing agent, and oxidizes a part of the hydrobromic acid, liberating bromine:

H2SO4 + 2HBr = 2H2O + SO2 + 2Br.

Preparation of pure hydrobromic acid. A convenient way to make pure hydrobromic acid is by
the action of bromine upon moist red phosphorus. This can be done with the apparatus shown in
Fig. 56. Bromine is put into the dropping funnel A, and red phosphorus, together with enough
water to cover it, is placed in the flask B. By means of the stopcock the bromine is allowed to flow
drop by drop into the flask, the reaction taking place without the application of heat. The equations
are

(1) P + 3Br = PBr3,

(2) PBr3 + 3H2O = P(OH)3 + 3HBr.

Fig. 56
The U-tube C contains glass beads which have been moistened with water and rubbed in red
phosphorus. Any bromine escaping action in the flask acts upon the phosphorus in the U-tube. The
hydrobromic acid is collected in the same way as hydrochloric acid.
Properties. Hydrobromic acid very strikingly resembles hydrochloric acid in physical and
chemical properties. It is a colorless, strongly fuming gas, heavier than hydrochloric acid and, like
it, is very soluble in water. Under standard conditions 1 volume of water dissolves 610 volumes of
the gas. Chemically, the chief point in which it differs from hydrochloric acid is in the fact that it is
much more easily oxidized, so that bromine is more readily set free from it than chlorine is from
hydrochloric acid.
Salts of hydrobromic acid,—bromides. The bromides are very similar to the chlorides in their
properties. Chlorine acts upon both bromides and free hydrobromic acid, liberating bromine from
them:

KBr + Cl = KCl + Br,

HBr + Cl = HCl + Br.

Silver bromide is extensively used in photography, and the bromides of sodium and potassium are
used as drugs.
Oxygen compounds. No oxides of bromine are surely known, and bromine does not form so
many oxygen acids as chlorine does. Salts of hypobromous acid (HBrO) and bromic acid (HBrO3)
are known.

IODINE

Historical. Iodine was discovered in 1812 by Courtois in the ashes of certain sea plants. Its
presence was revealed by its beautiful violet vapor, and this suggested the name iodine (from the
Greek for violet appearance).
Occurrence. In the combined state iodine occurs in very small quantities in sea water, from which
it is absorbed by certain sea plants, so that it is found in their ashes. It occurs along with bromine

in salt springs and beds, and is also found in Chili saltpeter.
Preparation. Iodine may be prepared in a number of ways, the principal methods being the
following:
1. Laboratory method. Iodine can readily be prepared in the laboratory from an iodide by the
method used in preparing bromine, except that sodium iodide is substituted for sodium bromide. It
can also be made by passing chlorine into a solution of an iodide.
2. Commercial method. Commercially iodine was
formerly prepared from seaweed (kelp), but is
now obtained almost entirely from the deposits of
Chili saltpeter. The crude saltpeter is dissolved in
water and the solution evaporated until the
saltpeter crystallizes. The remaining liquors,
known as the "mother liquors," contain sodium
iodate (NaIO3), in which form the iodine is
present in the saltpeter. The chemical reaction by
which the iodine is liberated from this compound
is a complicated one, depending on the fact that
sulphurous acid acts upon iodic acid, setting
iodine free. This reaction is shown as follows:

2HIO3 + 5H2SO3 = 5H2SO4 + H2O + 2I.

Purification of iodine. Iodine can be purified
very conveniently in the following way. The crude
iodine is placed in an evaporating dish E (Fig. 57), Fig. 57
and the dish is set upon the sand bath S. The
iodine is covered with the inverted funnel F, and
the sand bath is gently heated with a Bunsen burner. As the dish becomes warm the iodine rapidly
evaporates and condenses again on the cold surface of the funnel in shining crystals.
This process, in which a solid is converted into a vapor and is again condensed into a solid without
passing through the liquid state, is called sublimation.
Physical properties. Iodine is a purplish-black, shining, heavy solid which crystallizes in brilliant
plates. Even at ordinary temperatures it gives off a beautiful violet vapor, which increases in
amount as heat is applied. It melts at 107° and boils at 175°. It is slightly soluble in water, but
readily dissolves in alcohol, forming a brown solution (tincture of iodine), and in carbon
disulphide, forming a violet solution. The element has a strong, unpleasant odor, though by no
means as irritating as that of chlorine and bromine.
Chemical properties. Chemically iodine is quite similar to chlorine and bromine, but is still less
active than bromine. It combines directly with many elements at ordinary temperatures. At
elevated temperatures it combines with hydrogen, but the reaction is reversible and the compound
formed is quite easily decomposed. Both chlorine and bromine displace it from its salts:

KI + Br = KBr + I,

KI + Cl = KCl + I.
When even minute traces of iodine are added to thin starch paste a very intense blue color
develops, and this reaction forms a delicate test for iodine. Iodine is extensively used in medicine,
especially in the form of a tincture. It is also largely used in the preparation of dyes and organic
drugs, iodoform, a substance used as an antiseptic, has the formula CHI3.

Hydriodic acid (HI). This acid cannot be prepared in pure condition by the action of sulphuric
acid upon an iodide, since the hydriodic acid set free is oxidized by the sulphuric acid just as in the
case of hydrobromic acid, but to a much greater extent. It can be prepared in exactly the same way
as hydrobromic acid, iodine being substituted for bromine. It can also be prepared by passing
hydrosulphuric acid into water in which iodine is suspended. The equation is

H2S + 2I = 2HI + S.

The hydriodic acid formed in this way dissolves in the water.
Properties and uses. Hydriodic acid resembles the corresponding acids of
chlorine and bromine in physical properties, being a strongly fuming, colorless
gas, readily soluble in water. Under standard conditions 1 volume of water
dissolves about 460 volumes of the gas. It is, however, more unstable than
either hydrochloric or hydrobromic acids, and on exposure to the air it
gradually decomposes in accordance with the equation

2HI + O = H2O + 2I.

Owing to the slight affinity between iodine and hydrogen the acid easily gives
up its hydrogen and is therefore a strong reducing agent. This is seen in its
action on sulphuric acid.
The salts of hydriodic acid, the iodides, are, in general, similar to the chlorides
and bromides. Potassium iodide (KI) is the most familiar of the iodides and is
largely used in medicine.
Oxygen compounds. Iodine has a much greater affinity for oxygen than has
either chlorine or bromine. When heated with nitric acid it forms a stable oxide
(I2O5). Salts of iodic acid (HIO3) and periodic acid (HIO4) are easily prepared,
and the free acids are much more stable than the corresponding acids of the
other members of this family.

GAY-LUSSAC'S LAW OF VOLUMES

In the discussion of the composition of hydrochloric acid it was stated that one
volume of hydrogen combines with one volume of chlorine to form two
volumes of hydrochloric acid. With bromine and iodine similar combining
ratios hold good. These facts recall the simple volume relations already noted
in the study of the composition of steam and ammonia. These relations may be
represented graphically in the following way:

Graph
In the early part of the past century Gay-Lussac, a distinguished French
chemist, studied the volume relations of many combining gases, and concluded
that similar relations always hold. His observations are summed up in the
following law: When two gases combine chemically there is always a simple
ratio between their volumes, and between the volume of either one of them and
that of the product, provided it is a gas. By a simple ratio is meant of course
the ratio of small whole numbers, as 1 : 2, 2 : 3.

EXERCISES

1. How do we account for the fact that liquid hydrofluoric acid is not an
electrolyte?
2. Why does sulphuric acid liberate hydrofluoric acid from its salts?
3. In the preparation of chlorine, what advantages are there in treating
manganese dioxide with a mixture of sodium chloride and sulphuric acid
rather than with hydrochloric acid?
4. Why must chlorine water be kept in the dark?
5. What is the derivation of the word nascent?
6. What substances studied are used as bleaching agents? To what is the
bleaching action due in each case?
7. What substances studied are used as disinfecting agents?
8. What is meant by the statement that hydrochloric acid is one of the strongest
acids?
9. What is the meaning of the phrase aqua regia?
10. Cl2O is the anhydride of what acid?

11. A solution of hydriodic acid on standing turns brown. How is this
accounted for?
12. How can bromine vapor and nitrogen peroxide be distinguished from each
other?
13. Write the equations for the reaction taking place when hydriodic acid is
prepared from iodine, phosphorus, and water.
14. From their behavior toward sulphuric acid, to what class of agents do
hydrobromic and hydriodic acids belong?
15. Give the derivation of the names of the elements of the chlorine family.
16. Write the names and formulas for the binary acids of the group in the order
of the stability of the acids.
17. What is formed when a metal dissolves in each of the following? nitric
acid; dilute sulphuric acid; concentrated sulphuric acid; hydrochloric acid;
aqua regia.
18. How could you distinguish between a chloride, a bromide, and an iodide?
19. What weight of sodium chloride is necessary to prepare sufficient
hydrochloric acid to saturate 1 l. of water under standard conditions?
20. On decomposition 100 l. of hydrochloric acid would yield how many liters
of hydrogen and chlorine respectively, the gases being measured under the
same conditions? Are your results in accord with the experimental facts?

CHAPTER XVII
CARBON AND SOME OF ITS SIMPLER COMPOUNDS

The family. Carbon stands at the head of a family of elements in the fourth
group in the periodic table. The resemblances between the elements of this
family, while quite marked, are not so striking as in the case of the elements of
the chlorine family. With the exception of carbon, these elements are
comparatively rare, and need not be taken up in detail in this chapter. Titanium
will be referred to again in connection with silicon which it very closely
resembles.
Occurrence. Carbon is found in nature in the uncombined state in several
forms. The diamond is practically pure carbon, while graphite and coal are
largely carbon, but contain small amounts of other substances. Its natural
compounds are exceedingly numerous and occur as gases, liquids, and solids.
Carbon dioxide is its most familiar gaseous compound. Natural gas and
petroleum are largely compounds of carbon with hydrogen. The carbonates,
especially calcium carbonate, constitute great strata of rocks, and are found in
almost every locality. All living organisms, both plant and animal, contain a
large percentage of this element, and the number of its compounds which go to
make up all the vast variety of animate nature is almost limitless. Over one
hundred thousand definite compounds containing carbon have been prepared.
In the free state carbon occurs in three allotropic forms, two of which are
crystalline and one amorphous.
Crystalline carbon. Crystalline carbon occurs in two forms,—diamond and
graphite.
1. Diamond. Diamonds are found in considerable quantities in several
localities, especially in South Africa, the East Indies, and Brazil. The crystals
belong to the regular system, but the natural stones do not show this very
clearly. When found they are usually covered with a rough coating which is
removed in the process of cutting. Diamond cutting is carried on most
extensively in Holland.

The density of the diamond is 3.5, and, though brittle, it is one of the hardest of
substances. Black diamonds, as well as broken and imperfect stones which are
valueless as gems, are used for grinding hard substances. Few chemical
reagents have any action on the diamond, but when heated in oxygen or the air
it blackens and burns, forming carbon dioxide.
Lavoisier first showed that carbon dioxide is formed by the combustion of the
diamond; and Sir Humphry Davy in 1814 showed that this is the only product
of combustion, and that the diamond is pure carbon.
The diamond as a gem. The pure diamond is perfectly transparent and
colorless, but many are tinted a variety of colors by traces of foreign
substances. Usually the colorless ones are the most highly prized, although in
some instances the color adds to the value; thus the famous Hope diamond is a
beautiful blue. Light passing through a diamond is very much refracted, and to
this fact the stone owes its brilliancy and sparkle.
Artificial preparation of diamonds. Many attempts have been made to
produce diamonds artificially, but for a long time these always ended in
failure, graphite and not diamonds being the product obtained. The French
chemist Moissan, in his extended study of chemistry at high temperatures,
finally succeeded (1893) in making some small ones. He accomplished this by
dissolving carbon in boiling iron and plunging the crucible containing the
mixture into water, as shown in Fig. 58. Under these conditions the carbon
crystallized in the iron in the form of the diamond. The diamonds were then
obtained by dissolving away the iron in hydrochloric acid.

Fig. 58
2. Graphite. This form of carbon is found in large quantities, especially in
Ceylon, Siberia, and in some localities of the United States and Canada. It is a
shining black substance, very soft and greasy to the touch. Its density is about
2.15. It varies somewhat in properties according to the locality in which it is
found, and is more easily attacked by reagents than is the diamond. It is also
manufactured by heating carbon with a small amount of iron (3%) in an
electric furnace. It is used in the manufacture of lead pencils and crucibles, as a
lubricant, and as a protective covering for iron in the form of a polish or a
paint.
Amorphous carbon. Although there are many varieties of amorphous carbon
known, they are not true allotropic modifications. They differ merely in their
degree of purity, their fineness of division, and in their mode of preparation.
These substances are of the greatest importance, owing to their many uses in
the arts and industries. As they occur in nature, or are made artificially, they
are nearly all impure carbon, the impurity depending on the particular
substance in question.
1. Pure carbon. Pure amorphous carbon is best prepared by charring sugar.
This is a substance consisting of carbon, hydrogen, and oxygen, the latter two

elements being present in the ratio of one oxygen atom to two of hydrogen.
When sugar is strongly heated the oxygen and hydrogen are driven off in the
form of water and pure carbon is left behind. Prepared in this way it is a soft,
lustrous, very bulky, black powder.
2. Coal and coke. Coals of various kinds were probably formed from vast
accumulations of vegetable matter in former ages, which became covered over
with earthy material and were thus protected from rapid decay. Under various
natural agencies the organic matter was slowly changed into coal. In anthracite
these changes have gone the farthest, and this variety of coal is nearly pure
carbon. Soft or bituminous coals contain considerable organic matter besides
carbon and mineral substances. When heated strongly out of contact with air
the organic matter is decomposed and the resulting volatile matter is driven off
in the form of gases and vapors, and only the mineral matter and carbon
remain behind. The gaseous product is chiefly illuminating gas and the solid
residue is coke. Some of the coke is found as a dense cake on the sides and
roof of the retort. This is called retort carbon and is quite pure.
3. Charcoal. This is prepared from wood in the same way that coke is made
from coal. When the process is carried on in retorts the products expelled by
the heat are saved. Among these are many valuable substances such as wood
alcohol and acetic acid. Where timber is abundant the process is carried out in
a wasteful way, by merely covering piles of wood with sod and setting the
wood on fire. Some wood burns and the heat from this decomposes the wood
not burned, forming charcoal from it. The charcoal, of course, contains the
mineral part of the wood from which it is formed.
4. Bone black. This is sometimes called animal charcoal, and is made by
charring bones and animal refuse. The organic part of the materials is thus
decomposed and carbon is left in a very finely divided state, scattered through
the mineral part which consists largely of calcium phosphate. For some uses
this mineral part is removed by treatment with hydrochloric acid and
prolonged washing.
5. Lampblack. Lampblack and soot are products of imperfect combustion of oil
and coal, and are deposited from a smoky flame on a cold surface. The carbon
in this form is very finely divided and usually contains various oily materials.
Properties. While the various forms of carbon differ in many properties,
especially in color and hardness, yet they are all odorless, tasteless solids,

insoluble in water and characterized by their stability towards heat. Only in the
intense heat of the electric arc does carbon volatilize, passing directly from the
solid state into a vapor. Owing to this fact the inside surface of an incandescent
light bulb after being used for some time becomes coated with a dark film of
carbon. It is not acted on at ordinary temperatures by most reagents, but at a
higher temperature it combines directly with many of the elements, forming
compounds called carbides. When heated in the presence of sufficient oxygen
it burns, forming carbon dioxide.
Uses of carbon. The chief use of amorphous carbon is for fuel to furnish heat
and power for all the uses of civilization. An enormous quantity of carbon in
the form of the purer coals, coke, and charcoal is used as a reducing agent in
the manufacture of the various metals, especially in the metallurgy of iron.
Most of the metals are found in nature as oxides, or in forms which can readily
be converted into oxides. When these oxides are heated with carbon the
oxygen is abstracted, leaving the metal. Retort carbon and coke are used to
make electric light carbons and battery plates, while lampblack is used for
indelible inks, printer's ink, and black varnishes. Bone black and charcoal have
the property of absorbing large volumes of certain gases, as well as smaller
amounts of organic matter; hence they are used in filters to remove noxious
gases and objectionable colors and odors from water. Bone black is used
extensively in the sugar refineries to remove coloring matter from the impure
sugars.
Chemistry of carbon compounds. Carbon is remarkable for the very large
number of compounds which it forms with the other elements, especially with
oxygen and hydrogen. Compounds containing carbon are more numerous than
all others put together, and the chemistry of these substances presents
peculiarities not met with in the study of other substances. For these reasons
the systematic study of carbon compounds, or of organic chemistryas it is
usually called, must be deferred until the student has gained some knowledge
of the chemistry of other elements. An acquaintance with a few of the most
familiar carbon compounds is, however, essential for the understanding of the
general principles of chemistry.
Compounds of carbon with hydrogen,—the hydrocarbons. Carbon unites
with hydrogen to form a very large number of compounds called
hydrocarbons. Petroleum and natural gas are essentially mixtures of a great
variety of these hydrocarbons. Many others are found in living plants, and still

others are produced by the decay of organic matter in the absence of air. Only
two of them, methane and acetylene, will be discussed here.
Methane (marsh gas) (CH4). This is one of the most important of these
hydrocarbons, and constitutes about nine tenths of natural gas. As its name
suggests, it is formed in marshes by the decay of vegetable matter under water,
and bubbles of the gas are often seen to rise when the dead leaves on the
bottom of pools are stirred. It also collects in mines, and, when mixed with air,
is called fire damp by the miners because of its great inflammability, damp
being an old name for a gas. It is formed when organic matter, such as coal or
wood, is heated in closed vessels, and is therefore a principal constituent of
coal gas.
Preparation. Methane is prepared in the laboratory by heating sodium or
calcium acetate with soda-lime. Equal weights of fused sodium acetate and
soda-lime are thoroughly dried, then mixed and placed in a good-sized, hard-
glass test tube fitted with a one-holed stopper and delivery tube. The mixture is
gradually heated, and when the air has been displaced from the tube the gas is
collected in bottles by displacement of water. Soda-lime is a mixture of sodium
and calcium hydroxides. Regarding it as sodium hydroxide alone, the equation
is

NaC2H3O2 + NaOH = Na2CO3 + CH4.

Properties. Methane is a colorless, odorless gas whose density is 0.55. It is
difficult to liquefy, boiling at -155° under standard pressure, and is almost
insoluble in water. It burns with a pale blue flame, liberating much heat, and
when mixed with oxygen is very explosive.
Davy's safety lamp. In 1815 Sir Humphry Davy invented a lamp for the use
of miners, to prevent the dreadful mine explosions then common, due to
methane mixed with air. The invention consisted in surrounding the upper part
of the common miner's lamp with a mantle of wire gauze and the lower part
with glass (Fig. 59). It has been seen that two gases will not combine until
raised to their kindling temperature, and if while combining they are cooled
below this point, the combination ceases. A flame will not pass through a wire
gauze because the metal, being a good conductor of heat, takes away so much
heat from the flame that the gases are cooled below the kindling temperature.
When a lamp so protected is brought into an explosive mixture the gases inside

the wire mantle burn in a series of little explosions, giving warning to the
miner that the air is unsafe.
Acetylene (C2H2). This is a colorless gas usually
having a disagreeable odor due to impurities. It is
now made in large quantities from calcium carbide
(CaC2). This substance is formed when coal and lime
are heated together in an electric furnace. When
treated with water the carbide is decomposed,
yielding acetylene:

CaC2 + 2H2O = C2H2 + Ca(OH)2.

Under ordinary conditions the gas burns with a very
smoky flame; in burners constructed so as to secure a
large amount of oxygen it burns with a very brilliant
white light, and hence is used as an illuminant.
Laboratory preparation. The gas can be prepared
readily in a generator such as is shown in Fig. 60. The
inner tube contains fragments of calcium carbide,
while the outer one is filled with water. As long as the
stopcock is closed the water cannot rise in the inner Fig. 59
tube. When the stopcock is open the water rises, and,
coming into contact with the carbide in the inner tube, generates acetylene.
This escapes through the stopcock, and after the air has been expelled may be
lighted as it issues from the burner.
Carbon forms two oxides, namely, carbon dioxide (CO2) and carbon monoxide
(CO).
Carbon dioxide (CO2). Carbon dioxide is present in the air to the extent of
about 3 parts in 10,000, and this apparently small amount is of fundamental
importance in nature. In some localities it escapes from the earth in great
quantities, and many spring waters carry large amounts of it in solution. When
these highly charged spring waters reach the surface of the earth, and the
pressure on them is removed, the carbon dioxide escapes with effervescence. It
is a product of the oxidation of all organic matter, and is therefore formed in
fires as well as in the process of decay. It is thrown off from the lungs of all

animals in respiration, and is a product of many
fermentation processes such as vinegar making and
brewing. Combined with metallic oxides it forms vast
deposits of carbonates in nature.
Preparation. In the laboratory carbon dioxide is always
prepared by the action of an acid upon a carbonate,
usually calcium carbonate, the apparatus shown in Fig. 39
serving the purpose very well. This reaction might be
expected to produce carbonic acid, thus:

CaCO3 + 2HCl = CaCl2 + H2CO3.

Carbonic acid is very unstable, however, and decomposes
into its anhydride, CO2, and water, thus:

H2CO3 = H2O + CO2.

The complete reaction is represented by the equation
Fig. 60
CaCO3 + 2HCl = CaCl2 + CO2 + H2O.

Physical properties. Carbon dioxide is a colorless, practically odorless gas
whose density is 1.5. Its weight may be inferred from the fact that it can be
siphoned, or poured like water, from one vessel downward into another. At 15°
and under ordinary pressure it dissolves in its own volume of water and
imparts a somewhat biting, pungent taste to it. It is easily condensed, and is
now prepared commercially in this form by pumping the gas into steel
cylinders (see Fig. 6) which are kept cold during the process. When the liquid
is permitted to escape into the air part of it instantly evaporates, and in so
doing absorbs so much heat that another portion is solidified, the solid form
strikingly resembling snow in appearance. This snow is very cold and mercury
can easily be frozen with it.
Solid carbon dioxide. Cylinders of liquid carbon dioxide are inexpensive, and
should be available in every school. To demonstrate the properties of solid
carbon dioxide, the cylinder should be placed across the table and supported in
such a way that the stopcock end is several inches lower than the other end. A
loose bag is made by holding the corners of a handkerchief around the neck of

the stopcock, and the cock is then turned on so that the gas rushes out in large
quantities. Very quickly a considerable quantity of the snow collects in the
handkerchief. To freeze mercury, press a piece of filter paper into a small
evaporating dish and pour the mercury upon it. Coil a flat spiral upon the end
of a wire, and dip the spiral into the mercury. Place a quantity of solid carbon
dioxide upon the mercury and pour 10 cc.-15 cc. of ether over it. In a minute
or two the mercury will solidify and may be removed from the dish by the wire
serving as a handle. The filter paper is to prevent the mercury from sticking to
the dish; the ether dissolves the solid carbon dioxide and promotes its rapid
conversion into gas.
Chemical properties. Carbon dioxide is incombustible, since it is, like water,
a product of combustion. It does not support combustion, as does nitrogen
peroxide, because the oxygen in it is held in very firm chemical union with the
carbon. Very strong reducing agents, such as highly heated carbon, can take
away half of its oxygen:

CO2 + C = 2CO.

Uses. The relation of carbon dioxide to plant life has been discussed in a
previous chapter. Water highly charged with carbon dioxide is used for making
soda water and similar beverages. Since it is a non-supporter of combustion
and can be generated readily, carbon dioxide is also used as a fire extinguisher.
Some of the portable fire extinguishers are simply devices for generating large
amounts of the gas. It is not necessary that all the oxygen should be kept away
from the fire in order to smother it. A burning candle is extinguished in air
which contains only 2.5% of carbon dioxide.
Carbonic acid (H2CO3). Like most of the oxides of the non-metallic elements,
carbon dioxide is an acid anhydride. It combines with water to form an acid of
the formula H2CO3, called carbonic acid:

H2O + CO2 = H2CO3.

The acid is, however, very unstable and cannot be isolated. Only a very small
amount of it is actually formed when carbon dioxide is passed into water, as is
evident from the small solubility of the gas. If, however, a base is present in
the water, salts of carbonic acid are formed, and these are quite stable:

2NaOH + H2O + CO2 = Na2CO3 + 2H2O.

Action of carbon dioxide on bases. This conduct is explained by the
principles of reversible reactions. The equation

H2O +CO2 <--> H2CO3

is a reversible equation, and the extent to which the reaction progresses
depends upon the relative concentrations of each of the three factors in it.
Equilibrium is ordinarily reached when very little H2CO3 is formed. If a base
is present in the water to combine with the H2CO3 as fast as it is formed, all of
the CO2 is converted into H2CO3, and thence into a carbonate.

Salts of carbonic acid,—carbonates. The carbonates form a very important
class of salts. They are found in large quantities in nature, and are often used in
chemical processes. Only the carbonates of sodium, potassium, and
ammonium are soluble, and these can be made by the action of carbon dioxide
on solutions of the bases, as has just been explained.
The insoluble carbonates are formed as precipitates when soluble salts are
treated with a solution of a soluble carbonate. Thus the insoluble calcium
carbonate can be made by bringing together solutions of calcium chloride and
sodium carbonate:

CaCl2 + Na2CO3 = CaCO3 + 2NaCl.

Most of the carbonates are decomposed by heat, yielding an oxide of the metal
and carbon dioxide. Thus lime (calcium oxide) is made by strongly heating
calcium carbonate:

CaCO3 = CaO + CO2.

Acid carbonates. Like all acids containing two acid hydrogen atoms, carbonic
acid can form both normal and acid salts. The acid carbonates are made by
treating a normal carbonate with an excess of carbonic acid. With few
exceptions they are very unstable, heat decomposing them even when in
solution.

Action of carbon dioxide on calcium hydroxide. If carbon dioxide is passed
into clear lime water, calcium carbonate is at first precipitated:

H2O + CO2 = H2CO3,

Ca(OH)2 + H2CO3 = CaCO3 + 2H2O.

Advantage is taken of this reaction in testing for the presence of carbon
dioxide, as already explained in the chapter on the atmosphere. If the current
of carbon dioxide is continued, the precipitate soon dissolves, because the
excess of carbonic acid forms calcium acid carbonate which is soluble:

CaCO3 + H2CO3 = Ca(HCO3)2.

If now the solution is heated, the acid carbonate is decomposed and calcium
carbonate once more precipitated:

Ca(HCO3)2 = CaCO3 + H2CO3.

Carbon monoxide (CO). Carbon monoxide can be made in a number of ways,
the most important of which are the three following:
1. By the partial oxidation of carbon. If a slow current of air is conducted over
highly heated carbon, the monoxide is formed, thus:

C + O = CO

It is therefore often formed in stoves when the air draught is insufficient. Water
gas, which contains large amounts of carbon monoxide, is made by partially
oxidizing carbon with steam:

C + H2O = CO + 2H.

2. By the partial reduction of carbon dioxide. When carbon dioxide is
conducted over highly heated carbon it is reduced to carbon monoxide by the
excess of carbon:

CO2 + C = 2CO.

When coal is burning in a stove or grate carbon dioxide is at first formed in the
free supply of air, but as the hot gas rises through the glowing coal it is
reduced to carbon monoxide. When the carbon monoxide reaches the free air
above the coal it takes up oxygen to form carbon dioxide, burning with the
blue flame so familiar above a bed of coals, especially in the case of hard
coals.
3. By the decomposition of oxalic acid. In the laboratory carbon monoxide is
usually prepared by the action of concentrated sulphuric acid upon oxalic acid.
The latter substance has the formula C2H2O4. The sulphuric acid, owing to its
affinity for water, decomposes the oxalic acid, as represented in the equation

C2H2O4 + (H2SO4) = (H2SO4) + H2O + CO2 + CO.

Properties. Carbon monoxide is a light, colorless, almost odorless gas, very
difficult to liquefy. Chemically it is very active, combining directly with a
great many substances. It has a great affinity for oxygen and is therefore
combustible and a good reducing agent. Thus, if carbon monoxide is passed
over hot copper oxide, the copper is reduced to the metallic state:

CuO + CO = Cu + CO2.

When inhaled it combines with the red coloring matter of the blood and in this
way prevents the absorption of oxygen, so that even a small quantity of the gas
may prove fatal.

Fig. 61
The reducing power of carbon monoxide. Fig. 61 illustrates a method of
showing the reducing power of carbon monoxide. The gas is generated by

gently heating 7 or 8 g. of oxalic acid with 25 cc. of concentrated sulphuric
acid in a 200 cc. flask A. The bottle B contains a solution of sodium hydroxide,
which removes the carbon dioxide formed along with the monoxide. C
contains a solution of calcium hydroxide to show that the carbon dioxide is
completely removed. E is a hard-glass tube containing 1 or 2 g. of copper
oxide, which is heated by a burner. The black copper oxide is reduced to
reddish metallic copper by the carbon monoxide, which is thereby changed to
carbon dioxide. The presence of the carbon dioxide is shown by the precipitate
in the calcium hydroxide solution in D. Any unchanged carbon monoxide is
collected over water in F.
Carbon disulphide (CS2). Just as carbon combines with oxygen to form
carbon dioxide, so it combines with sulphur to form carbon disulphide (CS2).
This compound has been described in the chapter on sulphur.
Hydrocyanic acid (prussic acid)(HCN). Under the proper conditions carbon
unites with nitrogen and hydrogen to form the acid HCN, called hydrocyanic
acid. It is a weak, volatile acid, and is therefore easily prepared by treating its
salts with sulphuric acid:

KCN + H2SO4 = KHSO4 + HCN.

It is most familiar as a gas, though it condenses to a colorless liquid boiling at
26°. It has a peculiar odor, suggesting bitter almonds, and is extremely
poisonous either when inhaled or when taken into the stomach. A single drop
may cause death. It dissolves readily in water, its solution being commonly
called prussic acid.
The salts of hydrocyanic acid are called cyanides, the cyanides of sodium and
potassium being the best known. These are white solids and are extremely
poisonous.
Solutions of potassium cyanide are alkaline. A solution of potassium cyanide
turns red litmus blue, and must therefore contain hydroxyl ions. The presence
of these ions is accounted for in the following way.

Although water is so little dissociated into its ions H+ and OH- that for most
purposes we may neglect the dissociation, it is nevertheless measurably
dissociated. Hydrocyanic acid is one of the weakest of acids, and dissociates to
an extremely slight extent. When a cyanide such as potassium cyanide

dissolves it freely dissociates, and the CN- ions must come to an equilibrium
with the H+ ions derived from the water:

H+ + CN- <--> HCN.

The result of this equilibrium is that quite a number of H+ ions from the water
are converted into undissociated HCN molecules. But for every H+ ion so
removed an OH- ion remains free, and this will give the solution alkaline
properties.

EXERCISES

1. How can you prove that the composition of the different allotropic forms of
carbon is the same?
2. Are lampblack and bone black allotropic forms of carbon? Will equal
amounts of heat be liberated in the combustion of 1 g. of each?
3. How could you judge of the relative purity of different forms of carbon?
4. Apart from its color, why should carbon be useful in the preparation of inks
and paints?
5. Could asbestos fibers be used to replace the wire in a safety lamp?
6. Why do most acids decompose carbonates?
7. What effect would doubling the pressure have upon the solubility of carbon
dioxide in water?
8. What compound would be formed by passing carbon dioxide into a solution
of ammonium hydroxide? Write the equation.
9. Write equations for the preparation of K2CO3; of BaCO3; of MgCO3.

10. In what respects are carbonic and sulphurous acids similar?
11. Give three reasons why the reaction which takes place when a solution of
calcium acid carbonate is heated, completes itself.
12. How could you distinguish between carbonates and sulphites?

13. How could you distinguish between oxygen, hydrogen, nitrogen, nitrous
oxide, and carbon dioxide?
14. Could a solution of sodium hydroxide be substituted for the solution of
calcium hydroxide in testing for carbon dioxide?
15. What weight of sodium hydroxide is necessary to neutralize the carbonic
acid formed by the action of hydrochloric acid on 100 g. of calcium carbonate?
16. What weight of calcium carbonate would be necessary to prepare sufficient
carbon dioxide to saturate 10 l. of water at 15° and under ordinary pressure?
17. On the supposition that calcium carbide costs 12 cents a kilogram, what
would be the cost of an amount sufficient to generate 100 l. of acetylene
measured at 20° and 740 mm.?
18. How would the volume of a definite amount of carbon monoxide compare
with the volume of carbon dioxide formed by its combustion, the
measurements being made under the same conditions?

CHAPTER XVIII
FLAMES,—ILLUMINANTS

Conditions necessary for flames. It has been seen that when two substances
unite chemically, with the production of light and heat, the act of union is
called combustion. When one of the substances undergoing combustion
remains solid at the temperature occasioned by the combustion, light may be
given off, but there is no flame. Thus iron wire burning in oxygen throws off a
shower of sparks and is brilliantly incandescent, but no flame is seen. When,
however, both of the substances are gases or vapors at the temperature reached
in the combustion, the act of union is accompanied by a flame.
Flames from burning liquids or solids. Many substances which are liquids or
solids at ordinary temperatures burn with a flame because the heat of
combustion vaporizes them slowly, and the flame is due to the union of this
vapor with the gas supporting the combustion.
Supporter of combustion. That gas which surrounds the flame and constitutes
the atmosphere in which the combustion occurs is said to support the
combustion. The other gas which issues into this atmosphere is said to be the
combustible gas. Thus, in the ordinary combustion of coal gas in the air the
coal gas is said to be combustible, while the air is regarded as the supporter of
combustion. These terms are entirely relative, however, for a jet of air issuing
into an atmosphere of coal gas will burn when ignited, the coal gas supporting
the combustion. Ordinarily, when we say that a gas is combustible we mean
that it is combustible in an atmosphere of air.
Either gas may be the supporter of combustion. That the terms combustible
and supporter of combustion are merely relative may be shown in the
following way: A lamp chimney A is fitted with a cork and glass tubes, as
shown in Fig. 62. The tube C should have a diameter of from 12 to 15 mm. A
thin sheet of asbestos in which is cut a circular opening about 2 cm. in
diameter is placed over the top of the chimney. The opening in the asbestos is
closed with the palm of the hand, and gas is admitted to the chimney through
the tube B. The air in the chimney is soon expelled through the tube C, and the
gas itself is then lighted at the lower end of this tube. The hand is now

removed from the opening in the asbestos, when
the flame at the end of the tube at once rises and
appears at the end within the chimney, as shown
in the figure. The excess of coal gas now escapes
from the opening in the asbestos and may be
lighted. The flame at the top of the asbestos
board is due to the combustion of coal gas in air,
while the flame within the chimney is due to the
combustion of air in coal gas, the air being
drawn up through the tube by the escaping gas.
Appearance of flames. The flame caused by the
union of hydrogen and oxygen is almost
colorless and invisible. Chlorine and hydrogen
combine with a pale violet flame, carbon
monoxide burns in oxygen with a blue flame,
while ammonia burns with a deep yellow flame.
The color and appearance of flames are therefore
often quite characteristic of the particular
combustion which occasions them. Fig. 62

Structure of flames. When the gas undergoing
combustion issues from a round opening into an atmosphere of the gas
supporting combustion, as is the case with the burning Bunsen burner (Fig.
63), the flame is generally conical in outline. It consists of several distinct
cones, one within the other, the boundary between them being marked by
differences of color or luminosity. In the simplest flame, of which hydrogen
burning in oxygen is a good example, these cones are two in number,—an
inner one, formed by unburned gas, and an outer one, usually more or less
luminous, consisting of the combining gases. This outer one is in turn
surrounded by a third envelope of the products of combustion; this envelope is
sometimes invisible, as in the present case, but is sometimes faintly luminous.
The lower part of the inner cone of the flame is quite cool and consists of
unburned gas. Toward the top of the inner cone the gas has become heated to a
high temperature by the burning envelope surrounding it. On reaching the
supporter of combustion on the outside it is far above its kindling temperature,
and combustion follows with the evolution of much heat. The region of
combustion just outside the inner cone is therefore the hottest part of the flame.

Oxidizing and reducing flames. Since the tip of the outside
cone consists of very hot products of combustion mixed with
oxygen from the air, a substance capable of oxidation placed
in this part of the flame becomes very hot and is easily
oxidized. The oxygen with which it combines comes, of
course, from the atmosphere, and not from the products of
combustion. This outer tip of the flame is called the oxidizing
flame.
At the tip of the inner cone the conditions are quite different.
This region consists of a highly heated combustible gas, which
has not yet reached a supply of oxygen.
If a substance rich in oxygen, such as a metallic oxide, is
placed in this region of the flame, the heated gases combine
with its oxygen and the substance is reduced. This part of the
flame is called the reducing flame. These flames are used in
testing certain substances, especially minerals. For this
purpose they are produced by blowing into a small luminous
Fig. 63 Bunsen flame from one side through a blowpipe. This is a
tube of the shape shown in Fig. 64. The flame is directed in
any desired way and has the oxidizing and reducing regions
very clearly marked (Fig. 65). It is non-luminous from the same causes which
render the open Bunsen burner flame non-luminous, the gases from the lungs
serving to furnish oxygen and to dilute the combustible gas.

Fig. 65

Luminosity of flames. The luminosity of flames is
due to a number of distinct causes, and may therefore
be increased or diminished in several ways.
1. Presence of solid matter. The most obvious of these
causes is the presence in the flame of incandescent
solid matter. Thus chalk dust sifted into a non-
luminous flame renders it luminous. When
hydrocarbons form a part of the combustible gas, as
they do in nearly all illuminating gases and oils, some
carbon is usually set free in the process of combustion.
This is made very hot by the flame and becomes
incandescent, giving out light. In a well-regulated
flame it is afterward burned up, but when the supply of
oxygen is insufficient it escapes from the flame as
lampblack or soot. That it is temporarily present in a
well-burning luminous flame may be demonstrated by
holding a cold object, such as a small evaporating
dish, in the flame for a few seconds. This cold object
cools the carbon below its kindling temperature, and it Fig. 64
is deposited on the object as soot.
2. Pressure. A second factor in the luminosity of flames is the pressure under
which the gases are burning. Under increased pressure there is more matter in
a given volume of a gas, and the chemical action is more energetic than when
the gases are rarefied. Consequently there is more heat and light. A candle
burning on a high mountain gives less light than when it burns at the sea level.
If the gas is diluted with a non-combustible gas, the effect is the same as if it is
rarefied, for under these conditions there is less combustible gas in a given
volume.
3. Temperature. The luminosity also depends upon the temperature attained in
the combustion. In general the hotter the flame the greater the luminosity;
hence cooling the gases before combustion diminishes the luminosity of the
flame they will make, because it diminishes the temperature attained in the
combustion. Thus the luminosity of the Bunsen flame is largely diminished by
the air drawn up with the gas. This is due in part to the fact that the burning gas
is diluted and cooled by the air drawn in. The oxygen thus introduced into the

flame also causes the combustion of the hot particles of carbon which would
otherwise tend to make the flame luminous.
Illuminating and fuel gases. A number of mixtures of combustible gases,
consisting largely of carbon compounds and hydrogen, find extensive use for
the production of light and heat. The three chief varieties are coal gas, water
gas, and natural gas. The use of acetylene gas has already been referred to.
Coal gas. Coal gas is made by heating bituminous coal in large retorts out of
contact with the air. Soft or bituminous coal contains, in addition to large
amounts of carbon, considerable quantities of compounds of hydrogen,
oxygen, nitrogen, and sulphur. When distilled the nitrogen is liberated partly in
the form of ammonia and cyanides and partly as free nitrogen gas; the sulphur
is converted into hydrogen sulphide, carbon disulphide, and oxides of sulphur;
the oxygen into water and oxides of carbon. The remaining hydrogen is set
free partly as hydrogen and partly in combination with carbon in the form of
hydrocarbons. The most important of these is methane, with smaller quantities
of many others, some of which are liquids or solids at ordinary temperatures.
The great bulk of the carbon remains behind as coke and retort carbon.
The manufacture of coal gas. In the manufacture of coal gas it is necessary to
separate from the volatile constituents formed by the heating of the coal all
those substances which are either solid or liquid at ordinary temperature, since
these would clog the gas pipes. Certain gaseous constituents, such as hydrogen
sulphide and ammonia, must also be removed. The method used to accomplish
this is shown in Fig. 66. The coal is heated in air-tight retorts illustrated by A.
The volatile products escape through the pipe X and bubble into the tarry
liquid in the large pipe B, known as the hydraulic main, which runs at right
angles to the retorts. Here is deposited the greater portion of the solid and
liquid products, forming a tarry mass known as coal tar. Much of the ammonia
also remains dissolved in this liquid. The partially purified gas then passes into
the pipes C, which serve to cool it and further remove the solid and liquid
matter. The gas then passes into D, which is filled with coke over which a jet
of water is sprayed. The water still further cools the gas and at the same time
partially removes such gaseous products as hydrogen sulphide and ammonia,
which are soluble in water. In E the gas passes over some material such as
lime, which removes the last portions of the sulphur compounds as well as
much of the carbon dioxide present. From E the gas passes into the large gas
holder F, from which it is distributed through pipes to the places where it is
burned.

Fig. 66
One ton of good gas coal yields approximately 10,000 cu. ft. of gas, 1400 lb.
of coke, 120 lb. of tar, and 20 gal. of ammoniacal liquor.
Not only is the ammonia obtained in the manufacture of the gas of great
importance, but the coal tar also serves as the source of many very useful
substances, as will be explained in Chapter XXXII.
Water gas. Water gas is essentially a mixture of carbon monoxide and
hydrogen. It is made by passing steam over very hot anthracite coal, when the
reaction shown in the following equation takes place:

C + H2O = CO + 2H.

When required merely to produce heat the gas is at once ready for use. When
made for illuminating purposes it must be enriched, that is, illuminants must be
added, since both carbon monoxide and hydrogen burn with non-luminous
flames. This is accomplished by passing it into heaters containing highly
heated petroleum oils. The gas takes up hydrocarbon gases formed in the
decomposition of the petroleum oils, which make it burn with a luminous
flame.
Water gas is very effective as a fuel, since both carbon monoxide and hydrogen
burn with very hot flames. It has little odor and is very poisonous. Its use is
therefore attended with some risk, since leaks in pipes are very likely to escape
notice.
Natural gas. This substance, so abundant in many localities, varies much in
composition, but is composed principally of methane. When used for lighting

purposes it is usually burned in a burner resembling an open Bunsen, the
illumination being furnished by an incandescent mantle. This is the case in the
familiar Welsbach burner. Contrary to statements frequently made, natural gas
contains no free hydrogen.

TABLE SHOWING COMPOSITION OF GASES

PENNSYLVANIA COAL WATER ENRICHED
NATURAL GAS GAS GAS WATER GAS
Hydrogen 41.3 52.88 30.00
Methane 90.64 43.6 2.16 24.00
Illuminants 3.9 12.05
Carbon
6.4 36.80 29.00
monoxide
Carbon dioxide 0.30 2.0 3.47 0.30
Nitrogen 9.06 1.2 4.69 2.50
Oxygen 0.3 1.50
Hydrocarbon
1.5 1.50
vapors

These are analyses of actual samples, and may be taken as about the average
for the various kinds of gases. Any one of these may vary considerably. The
nitrogen and oxygen in most cases is due to a slight admixture of air which is
difficult to exclude entirely in the manufacture and handling of gases.
Fuels. A variety of substances are used as fuels, the most important of them
being wood, coal, and the various gases mentioned above. Wood consists
mainly of compounds of carbon, hydrogen, and oxygen. The composition of
coal and the fuel gases has been given. Since these fuels are composed
principally of carbon and hydrogen or their compounds, the chief products of
combustion are carbon dioxide and water. The practice of heating rooms with
portable gas or oil stoves with no provision for removing the products of
combustion is to be condemned, since the carbon dioxide is generated in
sufficient quantities to render the air unfit for breathing. Rooms so heated also
become very damp from the large amount of water vapor formed in the
combustion, and which in cold weather condenses on the window glass,
causing the glass to "sweat." Both coal and wood contain a certain amount of
mineral substances which constitute the ashes.

The electric furnace. In recent years electric furnaces have come into wide
use in operations requiring a very high temperature. Temperatures as high as
3500° can be easily reached, whereas the hottest oxyhydrogen flame is not
much above 2000°. These furnaces are constructed on one of two general
principles.

Fig. 67
1. Arc furnaces. In the one type the source of heat is an electric arc formed
between carbon electrodes separated a little from each other, as shown in Fig.
67. The substance to be heated is placed in a vessel, usually a graphite
crucible, just below the arc. The electrodes and crucible are surrounded by
materials which fuse with great difficulty, such as magnesium oxide, the walls
of the furnace being so shaped as to reflect the heat downwards upon the
contents of the crucible.

Fig. 68
2. Resistance furnaces. In the other type of furnace the heat is generated by the
resistance offered to the current in its passage through the furnace. In its
simplest form it may be represented by Fig. 68. The furnace is merely a

rectangular box built up of loose bricks. The electrodes E, each consisting of a
bundle of carbon rods, are introduced through the sides of the furnace. The
materials to be heated, C, are filled into the furnace up to the electrodes, and a
layer of broken coke is arranged so as to extend from one electrode to the
other. More of the charge is then placed on top of the coke. In passing through
the broken coke the electrical current encounters great resistance. This
generates great heat, and the charge surrounding the coke is brought to a very
high temperature. The advantage of this type of furnace is that the temperature
can be regulated to any desired intensity.

EXERCISES

1. Why does charcoal usually burn with no flame? How do you account for the
flame sometimes observed when it burns?
2. How do you account for the fact that a candle burns with a flame?
3. What two properties must the mantle used in the Welsbach lamp possess?
4. (a) In what respects does the use of the Welsbach mantle resemble that of
lime in the calcium light? (b) If the mantle were made of carbon, would it
serve the same purpose?
5. Would anthracite coal be suitable for the manufacture of coal gas?
6. How could you prove the formation of carbon dioxide and water in the
combustion of illuminating gases?
7. Suggest a probable way in which natural gas has been formed.
8. Coal frequently contains a sulphide of iron. (a) What two sulphur
compounds are likely to be formed when gas is made from such coal? (b)
Suggest some suitable method for the removal of these compounds.
9. Why does the use of the bellows on the blacksmith's forge cause a more
intense heat?
10. What volume of oxygen is necessary to burn 100 l. of marsh gas and what
volume of carbon dioxide would be formed, all of the gases being measured
under standard conditions?

11. Suppose a cubic meter of Pennsylvania natural gas, measured under
standard conditions, were to be burned. How much water by weight would
result?

CHAPTER XIX
MOLECULAR WEIGHTS, ATOMIC WEIGHTS, FORMULAS

Introduction. In the chapter on The Atomic Theory, it was shown that if it
were true that two elements uniting to form a compound always combined in
the ratio of one atom of one element to one atom of the other element, it would
be a very easy matter to decide upon figures which would represent the
relative weights of the different atoms. It would only be necessary to select
some one element as a standard and determine the weight of every element
which combines with a definite weight (say 1 g.) of the standard element. The
figures so obtained would evidently represent the relative weights of the
atoms.
But the law of multiple proportion at once reminds us that two elements may
unite in several proportions; and there is no simple way to determine the
number of atoms present in the molecule of any compound. Consequently the
problem of deciding upon the relative atomic weights is not an easy one. To
the solution of this problem we must now turn.
Dalton's method of determining atomic weights. When Dalton first
advanced the atomic theory he attempted to solve this problem by very simple
methods. He thought that when only one compound of two elements is known
it is reasonable to suppose that it contains one atom of each element. He
therefore gave the formula HO to water, and HN to ammonia. When more than
two compounds were known he assumed that the most familiar or the most
stable one had the simple formula. He then determined the atomic weight as
explained above. The results he obtained were contradictory and very far from
satisfactory, and it was soon seen that some other method, resting on much
more scientific grounds, must be found to decide what compounds, if any, have
a single atom of each element present.
Determination of atomic weights. Three distinct steps are involved in the
determination of the atomic weight of an element: (1) determination of the
equivalent, (2) determination of molecular weights of its compounds, and (3)
deduction of the exact atomic weight from the equivalent and molecular
weights.

1. Determination of the equivalent. By the equivalent of an element is meant
the weight of the element which will combine with a fixed weight of some
other element chosen as a standard. It has already been explained that oxygen
has been selected as the standard element for atomic weights, with a weight of
16. This same standard will serve very well as a standard for equivalents. The
equivalent of an element is the weight of the element which will combine with
16 g. of oxygen. Thus 16 g. of oxygen combines with 16.03 g. of sulphur, 65.4
g. of zinc, 215.86 g. of silver, 70.9 g. of chlorine. These figures, therefore,
represent the equivalent weights of these elements.
Relation of atomic weights to equivalents. According to the atomic theory
combination always takes place between whole numbers of atoms. Thus one
atom unites with one other, or with two or three; or two atoms may unite with
three, or three with five, and so on.
When oxygen combines with zinc the combination must be between definite
numbers of the two kinds of atoms. Experiment shows that these two elements
combine in the ratio of 16 g. of oxygen to 65.4 g. of zinc. If one atom of
oxygen combines with one atom of zinc, then this ratio must be the ratio
between the weights of the two atoms. If one atom of oxygen combines with
two atoms of zinc, then the ratio between the weights of the two atoms will be
16: 32.7. If two atoms of oxygen combine with one atom of zinc, the ratio by
weight between the two atoms will be 8: 65.4. It is evident, therefore, that the
real atomic weight of an element must be some multiple or submultiple of the
equivalent; in other words, the equivalent multiplied by 1/2, 1, 2, or 3 will give
the atomic weight.
Combining weights. A very interesting relation holds good between the
equivalents of the various elements. We have just seen that the figures 16.03,
65.4, 215.86, and 70.9 are the equivalents respectively of sulphur, zinc, silver,
and chlorine. These same figures represent the ratios by weight in which these
elements combine among themselves. Thus 215.86 g. of silver combine with
70.9 g. of chlorine and with 2 × 16.03 g. of sulphur. 65.4 g. of zinc combine
with 70.9 g. of chlorine and 2 × 16.03 g. of sulphur.
By taking the equivalent or some multiple of it a value can be obtained for
each element which will represent its combining value, and for this reason is
called its combining weight. It is important to notice that the fact that a
combining weight can be obtained for each element is not a part of a theory,
but is the direct result of experiment.

Elements with more than one equivalent. It will be remembered that oxygen
combines with hydrogen in two ratios. In one case 16 g. of oxygen combine
with 2.016 g. of hydrogen to form water; in the other 16 g. of oxygen combine
with 1.008 g. of hydrogen to form hydrogen dioxide. The equivalents of
hydrogen are therefore 2.016 and 1.008. Barium combines with oxygen in two
proportions: in barium oxide the proportion is 16 g. of oxygen to 137.4 g. of
barium; in barium dioxide the proportion is 16 g. of oxygen to 68.7 g. of
barium.
In each case one equivalent is a simple multiple of the other, so the fact that
there may be two equivalents does not add to the uncertainty. All we knew
before was that the true atomic weight is some multiple of the equivalent.
2. The determination of molecular weights. To decide the question as to
which multiple of the equivalent correctly represents the atomic weight of an
element, it has been found necessary to devise a method of determining the
molecular weights of compounds containing the element in question. Since the
molecular weight of a compound is merely the sum of the weights of all the
atoms present in it, it would seem to be impossible to determine the molecular
weight of a compound without first knowing the atomic weights of the
constituent atoms, and how many atoms of each element are present in the
molecule. But certain facts have been discovered which suggest a way in
which this can be done.
Avogadro's hypothesis. We have seen that the laws of Boyle, Charles, and
Gay-Lussac apply to all gases irrespective of their chemical character. This
would lead to the inference that the structure of gases must be quite simple,
and that it is much the same in all gases.
In 1811 Avogadro, an Italian physicist, suggested that if we assume all gases
under the same conditions of temperature and pressure to have the same
number of molecules in a given volume, we shall have a probable explanation
of the simplicity of the gas laws. It is difficult to prove the truth of this
hypothesis by a simple experiment, but there are so many facts known which
are in complete harmony with this suggestion that there is little doubt that it
expresses the truth. Avogadro's hypothesis may be stated thus: Equal volumes
of all gases under the same conditions of temperature and pressure contain the
same number of molecules.

Avogadro's hypothesis and molecular weights. Assuming that Avogadro's
hypothesis is correct, we have a very simple means for deciding upon the
relative weights of molecules; for if equal volumes of two gases contain the
same number of molecules, the weights of the two volumes must be in the
same ratio as the weights of the individual molecules which they contain. If we
adopt some one gas as a standard, we can express the weights of all other gases
as compared with this one, and the same figures will express the relative
weights of the molecules of which the gases are composed.
Oxygen as the standard. It is important that the same standard should be
adopted for the determination of molecular weights as has been decided upon
for atomic weights and equivalents, so that the three values may be in harmony
with each other. Accordingly it is best to adopt oxygen as the standard element
with which to compare the molecular weights of other gases, being careful to
keep the oxygen atom equal to 16.
The oxygen molecule contains two atoms. One point must not be
overlooked, however. We desire to have our unit, the oxygen atom, equal to 16.
The method of comparing the weights of gases just suggested compares the
molecules of the gases with the molecule of oxygen. Is the molecule and the
atom of oxygen the same thing? This question is answered by the following
considerations.
We have seen that when steam is formed by the union of oxygen and
hydrogen, two volumes of hydrogen combine with one volume of oxygen to
form two volumes of steam. Let us suppose that the one volume of oxygen
contains 100 molecules; then the two volumes of steam must, according to
Avogadro's hypothesis, contain 200 molecules. But each of these 200
molecules must contain at least one atom of oxygen, or 200 in all, and these
200 atoms came from 100 molecules of oxygen. It follows that each molecule
of oxygen must contain at least two atoms of oxygen.
Evidently this reasoning merely shows that there are at least two atoms in the
oxygen molecule. There may be more than that, but as there is no evidence to
this effect, we assume that the molecule contains two atoms only.
It is evident that if we wish to retain the value 16 for the atom of oxygen we
must take twice this value, or 32, for the value of the oxygen molecule, when
using it as a standard for molecular weights.

Determination of the molecular weights of gases from their weights
compared with oxygen. Assuming the molecular weight of oxygen to be 32,
Avogadro's hypothesis gives us a ready means for determining the molecular
weight of any other gas, for all that is required is to know its weight compared
with that of an equal volume of oxygen. For example, 1 l. of chlorine is found
by experiment to weigh 2.216 times as much as 1 l. of oxygen. The molecular
weight of chlorine must therefore be 2.216 ×32, or 70.91.
If, instead of comparing the relative weights of 1 l. of the two gases, we select
such a volume of oxygen as will weigh 32 g., or the weight in grams
corresponding to the molecular weight of the gas, the calculation is much
simplified. It has been found that 32 g. of oxygen, under standard conditions,
measure 22.4 l. This same volume of hydrogen weighs 2.019 g.; of chlorine
70.9 g.; of hydrochloric acid 36.458 g. The weights of these equal volumes
must be proportional to their molecular weights, and since the weight of the
oxygen is the same as the value of its molecular weight, so too will the weights
of the 22.4 l. of the other gases be equal to the value of their molecular
weights.
As a summary we can then make the following statement: The molecular
weight of any gas may be determined by calculating the weight of 22.4 l. of the
gas, measured under standard conditions.
Determination of molecular weights from density of gases. In an actual
experiment it is easier to determine the density of a gas than the weight of a
definite volume of it. The density of a gas is usually defined as its weight
compared with that of an equal volume of air. Having determined the density
of a gas, its weight compared with oxygen may be determined by multiplying
its density by the ratio between the weights of air and oxygen. This ratio is
0.9046. To compare it with our standard for atomic weights we must further
multiply it by 32, since the standard is 1/32 the weight of oxygen molecules.
The steps then are these:
1. Determine the density of the gas (its weight compared with air).
2. Multiply by 0.9046 to make the comparison with oxygen molecules.
3. Multiply by 32 to make the comparison with the unit for atomic weights.
We have, then, the formula:

molecular weight = density × 0.9046 × 32;

or, still more briefly,

M. = D. × 28.9.

The value found by this method for the determination of molecular weights
will of course agree with those found by calculating the weight of 22.4 l. of the
gas, since both methods depend on the same principles.
Determination of densities of
gases. The relative weights of
equal volumes of two gases can
be easily determined. The
following is one of the methods
used. A small flask, such as is
shown in Fig. 69, is filled with
one of the gases, and after the
temperature and pressure have
been noted the flask is sealed up
and weighed. The tip of the
sealed end is then broken off, the
flask filled with the second gas,
and its weight determined. If the
weight of the empty flask is
subtracted from these two
weighings, the relative weights Fig. 69
of the gases is readily found.
3. Deduction of atomic weights from molecular weights and equivalents.
We have now seen how the equivalent of an element and the molecular weight
of compounds containing the element can be obtained. Let us see how it is
possible to decide which multiple of the equivalent really is the true atomic
weight. As an example, let us suppose that the equivalent of nitrogen has been
found to be 7.02 and that it is desired to obtain its atomic weight. The next step
is to obtain the molecular weights of a large number of compounds containing
nitrogen. The following will serve:

DENSITY BY APPROXIMATE PERCENTAGE PART OF
EXPERIMENT MOLECULAR OF MOLECULAR

WEIGHT (D. × NITROGEN WEIGHT
28.9) BY DUE TO
EXPERIMENT NITROGEN
Nitrogen gas 0.9671 27.95 100.00 27.95
Nitrous
1.527 44.13 63.70 27.11
oxide
Nitric oxide 1.0384 30.00 46.74 14.02
Nitrogen
1.580 45.66 30.49 13.90
peroxide
Ammonia 0.591 17.05 82.28 14.03
Nitric acid 2.180 63.06 22.27 14.03
Hydrocyanic
0.930 26.87 51.90 13.94
acid

Method of calculation. The densities of the various gases in the first column
of this table are determined by experiment, and are fairly accurate but not
entirely so. By multiplying these densities by 28.9 the molecular weights of the
compounds as given in the second column are obtained. By chemical analysis
it is possible to determine the percentage composition of these substances, and
the percentages of nitrogen in them as determined by analysis are given in the
third column. If each of these molecular weights is multiplied in turn by the
percentage of nitrogen in the compound, the product will be the weight of the
nitrogen in the molecular weight of the compound. This will be the sum of the
weights of the nitrogen atoms in the molecule. These values are given in the
fourth column in the table.
If a large number of compounds containing nitrogen are studied in this way, it
is probable that there will be included in the list at least one substance whose
molecule contains a single nitrogen atom. In this case the number in the fourth
column will be the approximate atomic weight of nitrogen. On comparing the
values for nitrogen in the table it will be seen that a number which is
approximately 14 is the smallest, and that the others are multiples of this.
These compounds of higher value, therefore, contain more than one nitrogen
atom in the molecule.
Accurate determination of atomic weights. Molecular weights cannot be
determined very accurately, and consequently the part in them due to nitrogen
is a little uncertain, as will be seen in the table. All we can tell by this method
is that the true weight is very near 14. The equivalent can however be

determined very accurately, and we have seen that it is some multiple or
submultiple of the true atomic weight. Since molecular-weight determinations
have shown that in the case of nitrogen the atomic weight is near 14, and we
have found the equivalent to be 7.02, it is evident that the true atomic weight is
twice the equivalent, or 7.02 × 2 = 14.04.
Summary. These, then, are the steps necessary to establish the atomic weight
of an element.
1. Determine the equivalent accurately by analysis.
2. Determine the molecular weight of a large number of compounds of the
element, and by analysis the part of the molecular weight due to the element.
The smallest number so obtained will be approximately the atomic weight.
3. Multiply the equivalent by the small whole number (usually 1, 2, or 3),
which will make a number very close to the approximate atomic weight. The
figure so obtained will be the true atomic weight.
Molecular weights of the elements. It will be noticed that the molecular
weight of nitrogen obtained by multiplying its density by 28.9 is 28.08. Yet the
atomic weight of nitrogen as deduced from a study of its gaseous compounds
is 14.04. The simplest explanation that can be given for this is that the gaseous
nitrogen is made up of molecules, each of which contains two atoms. In this
respect it resembles oxygen; for we have seen that an entirely different line of
reasoning leads us to believe that the molecule of oxygen contains two atoms.
When we wish to indicate molecules of these gases the symbols N2 and O2
should be used. When we desire to merely show the weights taking part in a
reaction this is not necessary.
The vapor densities of many of the elements show that, like oxygen and
nitrogen, their molecules consist of two atoms. In other cases, particularly
among the metals, the molecule and the atom are identical. Still other elements
have four atoms in their molecules.
While oxygen contains two atoms in its molecules, a study of ozone has led to
the conclusion that it has three. The formation of ozone from oxygen can
therefore be represented by the equation

3O2 = 2O3.

Other methods of determining molecular weights. It will be noticed that
Avogadro's law gives us a method by which we can determine the relative
weights of the molecules of two gases because it enables us to tell when we are
dealing with an equal number of the two kinds of molecules. If by any other
means we can get this information, we can make use of the knowledge so
gained to determine the molecular weights of the two substances.
Raoult's laws. Two laws have been discovered which give us just such
information. They are known as Raoult's laws, and can be stated as follows:
1. When weights of substances which are proportional to their molecular
weights are dissolved in the same weight of solvent, the rise of the boiling
point is the same in each case.
2. When weights of substances which are proportional to their molecular
weights are dissolved in the same weight of solvent, the lowering of the
freezing point is the same in each case.
By taking advantage of these laws it is possible to determine when two
solutions contain the same number of molecules of two dissolved substances,
and consequently the relative molecular weights of the two substances.
Law of Dulong and Petit. In 1819 Dulong and Petit discovered a very
interesting relation between the atomic weight of an element and its specific
heat, which holds true for elements in the solid state. If equal weights of two
solids, say, lead and silver, are heated through the same range of temperature,
as from 10° to 20°, it is found that very different amounts of heat are required.
The amount of heat required to change the temperature of a solid or a liquid by
a definite amount compared with the amount required to change the
temperature of an equal weight of water by the same amount is called its
specific heat. Dulong and Petit discovered the following law: The specific heat
of an element in the solid form multiplied by its atomic weight is approximately
equal to the constant 6.25. That is,

at. wt. × sp. ht. = 6.25.

Consequently,

6.25
at. wt. = ————

sp. ht.

This law is not very accurate, but it is often possible by means of it to decide
upon what multiple of the equivalent is the real atomic weight. Thus the
specific heat of iron is found by experiment to be 0.112, and its equivalent is
27.95. 6.25 ÷ 0.112 = 55.8. We see, therefore, that the atomic weight is twice
the equivalent, or 55.9.
How formulas are determined. It will be well in connection with molecular
weights to consider how the formula of a compound is decided upon, for the
two subjects are very closely associated. Some examples will make clear the
method followed.
The molecular weight of a substance containing hydrogen and chlorine was
36.4. By analysis 36.4 parts of the substance was found to contain 1 part of
hydrogen and 35.4 parts of chlorine. As these are the simple atomic weights of
the two elements, the formula of the compound must be HCl.
A substance consisting of oxygen and hydrogen was found to have a molecular
weight of 34. Analysis showed that in 34 parts of the substance there were 2
parts of hydrogen and 32 parts of oxygen. Dividing these figures by the atomic
weights of the two elements, we get 2 ÷ 1 = 2 for H; 32 ÷ 16 = 2 for O. The
formula is therefore H2O2.

A substance containing 2.04% H, 32.6% S, and 65.3% O was found to have a
molecular weight of 98. In these 98 parts of the substance there are 98 × 2.04%
= 2 parts of H, 98 × 32.6% = 32 parts of S, and 98 × 65.3% = 64 parts of O. If
the molecule weighs 98, the hydrogen atoms present must together weigh 2,
the sulphur atoms 32, and the oxygen atoms 64. Dividing these figures by the
respective atomic weights of the three elements, we have, for H, 2 ÷ 1 = 2
atoms; for S, 32 ÷ 32 = 1 atom; for O, 64 ÷ 16 = 4 atoms. Hence the formula is
H2SO4.

We have, then, this general procedure: Find the percentage composition of the
substance and also its molecular weight. Multiply the molecular weight
successively by the percentage of each element present, to find the amount of
the element in the molecular weight of the compound. The figures so obtained
will be the respective parts of the molecular weight due to the several atoms.
Divide by the atomic weights of the respective elements, and the quotient will
be the number of atoms present.

Avogadro's hypothesis and chemical calculations. This law simplifies many
chemical calculations.
1. Application to volume relations in gaseous reactions. Since equal volumes
of gases contain an equal number of molecules, it follows that when an equal
number of gaseous molecules of two or more gases take part in a reaction, the
reaction will involve equal volumes of the gases. In the equation

C2H2O4 = H2O + CO2 + CO,

since 1 molecule of each of the gases CO2 and CO is set free from each
molecule of oxalic acid, the two substances must always be set free in equal
volumes.
Acetylene burns in accordance with the equation

2C2H2 + 5O2 = 4CO2 + 2H2O.

Hence 2 volumes of acetylene will react with 5 volumes of oxygen to form 4
volumes of carbon dioxide and 2 volumes of steam. That the volume relations
may be correct a gaseous element must be given its molecular formula. Thus
oxygen must be written O2 and not 2O.

2. Application to weights of gases. It will be recalled that the molecular weight
of a gas is determined by ascertaining the weight of 22.4 l. of the gas. This
weight in grams is called the gram-molecular weight of a gas. If the molecular
weight of any gas is known, the weight of a liter of the gas under standard
conditions may be determined by dividing its gram-molecular weight by 22.4.
Thus the gram-molecular weight of a hydrochloric acid gas is 36.458. A liter of
the gas will therefore weigh 36.458 ÷ 22.4 = 1.627 g.

EXERCISES

1. From the following data calculate the atomic weight of sulphur. The
equivalent, as obtained by an analysis of sulphur dioxide, is 16.03. The
densities and compositions of a number of compounds containing sulphur are
as follows:

NAME DENSITY COMPOSITION BY PERCENTAGE

Hydrosulphuric acid 1.1791 S = 94.11 H = 5.89
Sulphur dioxide 2.222 S = 50.05 O = 49.95
Sulphur trioxide 2.74 S = 40.05 O = 59.95
Sulphur chloride 4.70 S = 47.48 Cl = 52.52
Sulphuryl chloride 4.64 S = 23.75 Cl = 52.53 O = 23.70
Carbon disulphide 2.68 S = 84.24 C = 15.76

2. Calculate the formulas for compounds of the following compositions:

MOLECULAR WEIGHT
(1) S = 39.07% O = 58.49% H = 2.44% 81.0
(2) Ca = 29.40 S = 23.56 O = 47.04 136.2
(3) K = 38.67 N = 13.88 O = 47.45 101.2

3. The molecular weight of ammonia is 17.06; of sulphur dioxide is 64.06; of
chlorine is 70.9. From the molecular weight calculate the weight of 1 l. of each
of these gases. Compare your results with the table on the back cover of the
book.
4. From the molecular weight of the same gases calculate the density of each,
referred to air as a standard.
5. A mixture of 50 cc. of carbon monoxide and 50 cc. of oxygen was exploded
in a eudiometer, (a) What gases remained in the tube after the explosion? (b)
What was the volume of each?
6. In what proportion must acetylene and oxygen be mixed to produce the
greatest explosion?
7. Solve Problem 18, Chapter XVII, without using molecular weights.
Compare your results.
8. Solve Problem 10, Chapter XVIII, without using molecular weights.
Compare your results.
9. The specific heat of aluminium is 0.214; of lead is 0.031. From these
specific heats calculate the atomic weights of each of the elements.

CHAPTER XX
THE PHOSPHORUS FAMILY

SYMBOL ATOMIC WEIGHT DENSITY MELTING POINT
Phosphorus P 31.0 1.8 43.3°
Arsenic As 75.0 5.73 —
Antimony Sb 120.2 6.7 432°
Bismuth Bi 208.5 9.8 270°

The family. The elements constituting this family belong in the same group
with nitrogen and therefore resemble it in a general way. They exhibit a regular
gradation of physical properties, as is shown in the above table. The same
general gradation is also found in their chemical properties, phosphorus being
an acid-forming element, while bismuth is essentially a metal. The other two
elements are intermediate in properties.
Compounds. In general the elements of the family form compounds having
similar composition, as is shown in the following table:

PH3 PCl3 PCl5 P2O3 P2O5
AsH3 AsCl3 AsCl5 As2O3 As2O5
SbH3 SbCl3 SbCl5 Sb2O3 Sb2O5
BiCl3 BiCl5 Bi2O3 Bi2O5

In the case of phosphorus, arsenic, and antimony the oxides are acid
anhydrides. Salts of at least four acids of each of these three elements are
known, the free acid in some instances being unstable. The relation of these
acids to the corresponding anhydrides may be illustrated as follows,
phosphorus being taken as an example:

P2O3 + 3H2O = 2H3PO3 (phosphorous acid).

P2O5 + 3H2O = 2H3PO4 (phosphoric acid).

P2O5 + 2H2O = H4P2O7 (pyrophosphoric acid).

P2O5 + H2O = 2HPO3 (metaphosphoric acid).

PHOSPHORUS

History. The element phosphorus was discovered by the alchemist Brand, of
Hamburg, in 1669, while searching for the philosopher's stone. Owing to its
peculiar properties and the secrecy which was maintained about its
preparation, it remained a very rare and costly substance until the demand for
it in the manufacture of matches brought about its production on a large scale.
Occurrence. Owing to its great chemical activity phosphorus never occurs
free in nature. In the form of phosphates it is very abundant and widely
distributed. Phosphorite and sombrerite are mineral forms of calcium
phosphate, while apatite consists of calcium phosphate together with calcium
fluoride or chloride. These minerals form very large deposits and are
extensively mined for use as fertilizers. Calcium phosphate is a constituent of
all fertile soil, having been supplied to the soil by the disintegration of rocks
containing it. It is the chief mineral constituent of bones of animals, and bone
ash is therefore nearly pure calcium phosphate.
Preparation. Phosphorus is now manufactured from bone ash or a pure
mineral phosphate by heating the phosphate with sand and carbon in an
electric furnace. The materials are fed in at M (Fig. 70) by the feed screw F.
The phosphorus vapor escapes at P and is condensed under water, while the
calcium silicate is tapped off as a liquid at S. The phosphorus obtained in this
way is quite impure, and is purified by distillation.

Fig. 70
Explanation of the reaction. To understand the reaction which occurs, it must
be remembered that a volatile acid anhydride is expelled from its salts when
heated with an anhydride which is not volatile. Thus, when sodium carbonate
and silicon dioxide are heated together the following reaction takes place:
Na2CO3 + SiO2 = Na2SiO3 + CO2.

Silicon dioxide is a less volatile anhydride than phosphoric anhydride (P2O5),
and when strongly heated with a phosphate the phosphoric anhydride is driven
out, thus:
Ca3(PO4)2 + 3SiO2 = 3CaSiO3 + P2O5.

If carbon is added before the heat is applied, the P2O5 is reduced to phosphorus
at the same time, according to the equation
P2O5 + 5C = 2P + 5CO.

Physical properties. The purified phosphorus is a pale yellowish, translucent,
waxy solid which melts at 43.3° and boils at 269°. It can therefore be cast into

any convenient form under warm water, and is usually sold in the market in the
form of sticks. It is quite soft and can be easily cut with a knife, but this must
always be done while the element is covered with water, since it is extremely
inflammable, and the friction of the knife blade is almost sure to set it on fire if
cut in the air. It is not soluble in water, but is freely soluble in some other
liquids, notably in carbon disulphide. Its density is 1.8.
Chemical properties. Exposed to the air phosphorus slowly combines with
oxygen, and in so doing emits a pale light, or phosphorescence, which can be
seen only in a dark place. The heat of the room may easily raise the
temperature to the kindling point of phosphorus, when it burns with a
sputtering flame, giving off dense fumes of oxide of phosphorus. It burns with
dazzling brilliancy in oxygen, and combines directly with many other
elements, especially with sulphur and the halogens. On account of its great
affinity for oxygen it is always preserved under water.
Phosphorus is very poisonous, from 0.2 to 0.3 gram being a fatal dose. Ground
up with flour and water or similar substances, it is often used as a poison for
rats and other vermin.
Precaution. The heat of the body is sufficient to raise phosphorus above its
kindling temperature, and for this reason it should always be handled with
forceps and never with the bare fingers. Burns occasioned by it are very
painful and slow in healing.
Red phosphorus. On standing, yellow phosphorus gradually undergoes a
remarkable change, being converted into a dark red powder which has a
density of 2.1. It no longer takes fire easily, neither does it dissolve in carbon
disulphide. It is not poisonous and, in fact, seems to be an entirely different
substance. The velocity of this change increases with rise in temperature, and
the red phosphorus is therefore prepared by heating the yellow just below the
boiling point (250°-300°). When distilled and quickly condensed the red form
changes back to the yellow. This is in accordance with the general rule that
when a substance capable of existing in several allotropic forms is condensed
from a gas or crystallized from the liquid state, the more unstable variety forms
first, and this then passes into the more stable forms.
Matches. The chief use of phosphorus is in the manufacture of matches.
Common matches are made by first dipping the match sticks into some
inflammable substance, such as melted paraffin, and afterward into a paste

consisting of (1) phosphorus, (2) some oxidizing substance, such as manganese
dioxide or potassium chlorate, and (3) a binding material, usually some kind of
glue. On friction the phosphorus is ignited, the combustion being sustained by
the oxidizing agent and communicated to the wood by the burning paraffin. In
sulphur matches the paraffin is replaced by sulphur.
In safety matches red phosphorus, an oxidizing agent, and some gritty material
such as emery is placed on the side of the box, while the match tip is provided
as before with an oxidizing agent and an easily oxidized substance, usually
antimony sulphide. The match cannot be ignited easily by friction, save on the
prepared surface.
Compounds of phosphorus with hydrogen. Phosphorus forms several
compounds with hydrogen, the best known of which is phosphine (PH3)
analogous to ammonia (NH3).

Preparation of phosphine. Phosphine is usually made by heating phosphorus
with a strong solution of potassium hydroxide, the reaction being a
complicated one.

Fig. 71

The experiment can be conveniently made in the apparatus shown in Fig. 71. A
strong solution of potassium hydroxide together with several small bits of
phosphorus are placed in the flask A, and a current of coal gas is passed into
the flask through the tube B until all the air has been displaced. The gas is then
turned off and the flask is heated. Phosphine is formed in small quantities and
escapes through the delivery tube, the exit of which is just covered by the
water in the vessel C. Each bubble of the gas as it escapes into the air takes
fire, and the product of combustion (P2O5) forms beautiful small rings, which
float unbroken for a considerable time in quiet air. The pure phosphine does
not take fire spontaneously. When prepared as directed above, impurities are
present which impart this property.
Properties. Phosphine is a gas of unpleasant odor and is exceedingly
poisonous. Like ammonia it forms salts with the halogen acids. Thus we have
phosphonium chloride (PH4Cl) analogous to ammonium chloride (NH4Cl).
The phosphonium salts are of but little importance.
Oxides of phosphorus. Phosphorus forms two well-known oxides,—the
trioxide (P2O3) and the pentoxide (P2O5), sometimes called phosphoric
anhydride. When phosphorus burns in an insufficient supply of air the product
is partially the trioxide; in oxygen or an excess of air the pentoxide is formed.
The pentoxide is much the better known of the two. It is a snow-white,
voluminous powder whose most marked property is its great attraction for
water. It has no chemical action upon most gases, so that they can be very
thoroughly dried by allowing them to pass through properly arranged vessels
containing phosphorus pentoxide.
Acids of phosphorus. The important acids of phosphorus are the following:

H3PO3 phosphorous acid.
H3PO4 phosphoric acid.
H4P2O7 pyrophosphoric acid.
HPO3 metaphosphoric acid.

These may be regarded as combinations of the oxides of phosphorus with
water according to the equations given in the discussion of the characteristics
of the family.

1. Phosphorous acid (H3PO3). Neither the acid nor its salts are at all frequently
met with in chemical operations. It can be easily obtained, however, in the
form of transparent crystals when phosphorus trichloride is treated with water
and the resulting solution is evaporated:

PCl3 + 3H2O = H3PO3 + 3HCl.

Its most interesting property is its tendency to take up oxygen and pass over
into phosphoric acid.
2. Orthophosphoric acid (phosphoric acid) (H3PO4). This acid can be obtained
by dissolving phosphorus pentoxide in boiling water, as represented in the
equation

P2O5 + 3H2O = 2H3PO4.

It is usually made by treating calcium phosphate with concentrated sulphuric
acid. The calcium sulphate produced in the reaction is nearly insoluble, and
can be filtered off, leaving the phosphoric acid in solution. Very pure acid is
made by oxidizing phosphorus with nitric acid. It forms large colorless crystals
which are exceedingly soluble in water. Being a tribasic acid, it forms acid as
well as normal salts. Thus the following compounds of sodium are known:

NaH2PO4 monosodium hydrogen phosphate.
Na2HPO4 disodium hydrogen phosphate.
Na3PO4 normal sodium phosphate.

These salts are sometimes called respectively primary, secondary, and tertiary
phosphates. They may be prepared by bringing together phosphoric acid and
appropriate quantities of sodium hydroxide. Phosphoric acid also forms mixed
salts, that is, salts containing two different metals. The most familiar
compound of this kind is microcosmic salt, which has the formula
Na(NH4)HPO4.

Orthophosphates. The orthophosphates form an important class of salts. The
normal salts are nearly all insoluble and many of them occur in nature. The
secondary phosphates are as a rule insoluble, while most of the primary salts
are soluble.

3. Pyrophosphoric acid (H4P2O7). On heating orthophosphoric acid to about
225° pyrophosphoric acid is formed in accordance with the following
equation:

2H3PO4 = H4P2O7 + H2O.

It is a white crystalline solid. Its salts can be prepared by heating a secondary
phosphate:

2Na2HPO4 = Na4P2O7 + H2O.

4. Metaphosphoric acid (glacial phosphoric acid) (HPO3). This acid is formed
when orthophosphoric acid is heated above 400°:

H3PO4 = HPO3 + H2O.

It is also formed when phosphorus pentoxide is treated with cold water:

P2O5 + H2O = 2HPO3.

It is a white crystalline solid, and is so stable towards heat that it can be fused
and even volatilized without decomposition. On cooling from the fused state it
forms a glassy solid, and on this account is often called glacial phosphoric
acid. It possesses the property of dissolving small quantities of metallic oxides,
with the formation of compounds which, in the case of certain metals, have
characteristic colors. It is therefore used in the detection of these metals.
While the secondary phosphates, on heating, give salts of pyrophosphoric acid,
the primary phosphates yield salts of metaphosphoric acid. The equations
representing these reactions are as follows:

2Na2HPO4 = Na4P3O7 + H2O,

NaH2PO4 = NaPO3 + H2O.

Fertilizers. When crops are produced year after year on the same field certain
constituents of the soil essential to plant growth are removed, and the soil
becomes impoverished and unproductive. To make the land once more fertile
these constituents must be replaced. The calcium phosphate of the mineral

deposits or of bone ash serves well as a material for restoring phosphorus to
soils exhausted of that essential element; but a more soluble substance, which
the plants can more readily assimilate, is desirable. It is better, therefore, to
convert the insoluble calcium phosphate into the soluble primary phosphate
before it is applied as fertilizer. It will be seen by reference to the formulas for
the orthophosphates (see page 244) that in a primary phosphate only one
hydrogen atom of phosphoric acid is replaced by a metal. Since the calcium
atom always replaces two hydrogen atoms, it might be thought that there could
be no primary calcium phosphate; but if the calcium atom replaces one
hydrogen atom from each of two molecules of phosphoric acid, the salt
Ca(H2PO4)2 will result, and this is a primary phosphate. It can be made by
treatment of the normal phosphate with the necessary amount of sulphuric
acid, calcium sulphate being formed at the same time, thus:

Ca3(PO4)2 + 2H2SO4 = Ca(H2PO4)2 + 2CaSO4.

The resulting mixture is a powder, which is sold as a fertilizer under the name
of "superphosphate of lime."

ARSENIC

Occurrence. Arsenic occurs in considerable quantities in nature as the native
element, as the sulphides realgar (As2S2) and orpiment (As2S3), as oxide
(As2O3), and as a constituent of many metallic sulphides, such as arsenopyrite
(FeAsS).
Preparation. The element is prepared by purifying the native arsenic, or by
heating the arsenopyrite in iron tubes, out of contact with air, when the
reaction expressed by the following equation occurs:

FeAsS = FeS + As.

The arsenic, being volatile, condenses in chambers connected with the heated
tubes. It is also made from the oxide by reduction with carbon:

2As2O3 + 3C = 4As + 3CO2.

Properties. Arsenic is a steel-gray, metallic-looking substance of density 5.73.
Though resembling metals in appearance, it is quite brittle, being easily

powdered in a mortar. When strongly heated it sublimes, that is, it passes into a
vapor without melting, and condenses again to a crystalline solid when the
vapor is cooled. Like phosphorus it can be obtained in several allotropic forms.
It alloys readily with some of the metals, and finds its chief use as an alloy
with lead, which is used for making shot, the alloy being harder than pure lead.
When heated on charcoal with the blowpipe it is converted into an oxide which
volatilizes, leaving the charcoal unstained by any oxide coating. It burns
readily in chlorine gas, forming arsenic trichloride,—

As + 3Cl = AsCl3.

Unlike most of its compounds, the element itself is not poisonous.
Arsine (AsH3). When any compound containing arsenic is brought into the
presence of nascent hydrogen, arsine (AsH3), corresponding to phosphine and
ammonia, is formed. The reaction when oxide of arsenic is so treated is

As2O3 + 12H = 2AsH3 + 3H2O.

Arsine is a gas with a peculiar garlic-like odor, and is intensely poisonous. A
single bubble of pure gas has been known to prove fatal. It is an unstable
compound, decomposing into its elements when heated to a moderate
temperature. It is combustible, burning with a pale bluish-white flame to form
arsenic trioxide and water when air is in excess:

2AsH3 + 6O = As2O3 + 3H2O.

When the supply of air is deficient water and metallic arsenic are formed:

2AsH3 + 3O = 3H2O + 2As.

These reactions make the detection of even minute quantities of arsenic a very
easy problem.

Fig. 72
Marsh's test for arsenic. The method devised by Marsh for detecting arsenic
is most frequently used, the apparatus being shown in Fig. 72. Hydrogen is
generated in the flask A by the action of dilute sulphuric acid on zinc, is dried
by passing over calcium chloride in the tube B, and after passing through the
hard-glass tube C is ignited at the jet D. If a substance containing arsenic is
now introduced into the generator A, the arsenic is converted into arsine by the
action of the nascent hydrogen, and passes to the jet along with the hydrogen.
If the tube C is strongly heated at some point near the middle, the arsine is
decomposed while passing this point and the arsenic is deposited just beyond
the heated point in the form of a shining, brownish-black mirror. If the tube is
not heated, the arsine burns along with the hydrogen at the jet. Under these
conditions a small porcelain dish crowded down into the flame is blackened by
a spot of metallic arsenic, for the arsine is decomposed by the heat of the
flame, and the arsenic, cooled below its kindling temperature by the cold
porcelain, deposits upon it as a black spot. Antimony conducts itself in the
same way as arsenic, but the antimony deposit is more sooty in appearance.
The two can also be distinguished by the fact that sodium hypochlorite
(NaClO) dissolves the arsenic deposit, but not that formed by antimony.
Oxides of arsenic. Arsenic forms two oxides, As2O3 and As2O5,
corresponding to those of phosphorus. Of these arsenious oxide, or arsenic
trioxide (As2O3), is much better known, and is the substance usually called
white arsenic, or merely arsenic. It is found as a mineral, but is usually
obtained as a by-product in burning pyrite in the sulphuric-acid industry. The

pyrite has a small amount of arsenopyrite in it, and when this is burned
arsenious oxide is formed as a vapor together with sulphur dioxide:

2FeAsS + 10O = Fe2O3 + As2O3 + 2SO2.

The arsenious oxide is condensed in appropriate chambers. It is a rather heavy
substance, obtained either as a crystalline powder or as large, vitreous lumps,
resembling lumps of porcelain in appearance. It is very poisonous, from 0.2 to
0.3 g. being a fatal dose. It is frequently given as a poison, since it is nearly
tasteless and does not act very rapidly. This slow action is due to the fact that it
is not very soluble, and hence is absorbed slowly by the system. Arsenious
oxide is also used as a chemical reagent in glass making and in the dye
industry.
Acids of arsenic. Like the corresponding oxides of phosphorus, the oxides of
arsenic are acid anhydrides. In solution they combine with bases to form salts,
corresponding to the salts of the acids of phosphorus. Thus we have salts of the
following acids:

H3AsO3 arsenious acid.
H3AsO4 orthoarsenic acid.
H4As2O3 pyroarsenic acid.
HAsO3 metarsenic acid.

Several other acids of arsenic are also known. Not all of these can be obtained
as free acids, since they tend to lose water and form the oxides. Thus, instead
of obtaining arsenious acid (H3AsO3), the oxide As2O3 is obtained:

2H3AsO3 = As2O3 + 3H2O.

Salts of all the acids are known, however, and some of them have commercial
value. Most of them are insoluble, and some of the copper salts, which are
green, are used as pigments. Paris green, which has a complicated formula, is a
well-known insecticide.
Antidote for arsenical poisoning. The most efficient antidote for arsenic
poisoning is ferric hydroxide. It is prepared as needed, according to the
equation

Fe2(SO4)3 + 3Mg(OH)2 = 2Fe(OH)3 + 3MgSO4.

Sulphides of arsenic. When hydrogen sulphide is passed into an acidified
solution containing an arsenic compound the arsenic is precipitated as a bright
yellow sulphide, thus:

2H3AsO3 + 3H2S = As2S3 + 6H2O,

2H3AsO4 + 5H2S = As2S5 + 8H2O.

In this respect arsenic resembles the metallic elements, many of which produce
sulphides under similar conditions. The sulphides of arsenic, both those
produced artificially and those found in nature, are used as yellow pigments.

ANTIMONY

Occurrence. Antimony occurs in nature chiefly as the sulphide (Sb2S3), called
stibnite, though it is also found as oxide and as a constituent of many complex
minerals.
Preparation. Antimony is prepared from the sulphide in a very simple manner.
The sulphide is melted with scrap iron in a furnace, when the iron combines
with the sulphur to form a slag, or liquid layer of melted iron sulphide, while
the heavier liquid, antimony, settles to the bottom and is drawn off from time
to time. The reaction involved is represented by the equation

Sb2S3 + 3Fe = 2Sb + 3FeS.

Physical properties. Antimony is a bluish-white, metallic-looking substance
whose density is 6.7. It is highly crystalline, hard, and very brittle. It has a
rather low melting point (432°) and expands very noticeably on solidifying.
Chemical properties. In chemical properties antimony resembles arsenic in
many particulars. It forms the oxides Sb2O3 and Sb2O5, and in addition Sb2O4.
It combines with the halogen elements with great energy, burning brilliantly in
chlorine to form antimony trichloride (SbCl3). When heated on charcoal with
the blowpipe it is oxidized and forms a coating of antimony oxide on the
charcoal which has a characteristic bluish-white color.

Stibine (SbH3). The gas stibine (SbH3) is formed under conditions which are
very similar to those which produce arsine, and it closely resembles the latter
compound, though it is still less stable. It is very poisonous.
Acids of antimony. The oxides Sb_{2}O_{3} and Sb_{2}O_{5} are weak
acid anhydrides and are capable of forming two series of acids corresponding
in formulas to the acids of phosphorus and arsenic. They are much weaker,
however, and are of little practical importance.
Sulphides of antimony. Antimony resembles arsenic in that hydrogen
sulphide precipitates it as a sulphide when conducted into an acidified solution
containing an antimony compound:

2SbCl3 + 3H2S = Sb2S3 + 6HCl,

2SbCl5 + 5H2S = Sb2S5 + 10HCl.

The two sulphides of antimony are called the trisulphide and the pentasulphide
respectively. When prepared in this way they are orange-colored substances,
though the mineral stibnite is black.
Metallic properties of antimony. The physical properties of the element are
those of a metal, and the fact that its sulphide is precipitated by hydrogen
sulphide shows that it acts like a metal in a chemical way. Many other
reactions show that antimony has more of the properties of a metal than of a
non-metal. The compound Sb(OH)3, corresponding to arsenious acid, while
able to act as a weak acid is also able to act as a weak base with strong acids.
For example, when treated with concentrated hydrochloric acid antimony
chloride is formed:

Sb(OH)3 + 3HCl = SbCl3 + 3H2O.

A number of elements act in this same way, their hydroxides under some
conditions being weak acids and under others weak bases.

ALLOYS

Some metals when melted together thoroughly intermix, and on cooling form a
homogeneous, metallic-appearing substance called an alloy. Not all metals will

mix in this way, and in some cases definite chemical compounds are formed
and separate out as the mixture solidifies, thus destroying the uniform quality
of the alloy. In general the melting point of the alloy is below the average of
the melting points of its constituents, and it is often lower than any one of
them.
Antimony forms alloys with many of the metals, and its chief commercial use
is for such purposes. It imparts to its alloys high density, rather low melting
point, and the property of expanding on solidification. Such an alloy is
especially useful in type founding, where fine lines are to be reproduced on a
cast. Type metal consists of antimony, lead, and tin. Babbitt metal, used for
journal bearings in machinery, contains the same metals in a different
proportion together with a small percentage of copper.

BISMUTH

Occurrence. Bismuth is usually found in the uncombined form in nature. It
also occurs as oxide and sulphide. Most of the bismuth of commerce comes
from Saxony, and from Mexico and Colorado, but it is not an abundant
element.
Preparation. It is prepared by merely heating the ore containing the native
bismuth and allowing the melted metal to run out into suitable vessels. Other
ores are converted into oxides and reduced by heating with carbon.
Physical properties. Bismuth is a heavy, crystalline, brittle metal nearly the
color of silver, but with a slightly rosy tint which distinguishes it from other
metals. It melts at a low temperature (270°) and has a density of 9.8. It is not
acted upon by the air at ordinary temperatures.
Chemical properties. When heated with the blowpipe on charcoal, bismuth
gives a coating of the oxide Bi2O3. This has a yellowish-brown color which
easily distinguishes it from the oxides formed by other metals. It combines
very readily with the halogen elements, powdered bismuth burning readily in
chlorine. It is not very easily acted upon by hydrochloric acid, but nitric and
sulphuric acids act upon it in the same way that they do upon copper.
Uses. Bismuth finds its chief use as a constituent of alloys, particularly in
those of low melting point. Some of these melt in hot water. For example,
Wood's metal, consisting of bismuth, lead, tin, and cadmium, melts at 60.5°.

Compounds of bismuth. Unlike the other elements of this group, bismuth has
almost no acid properties. Its chief oxide, Bi2O3, is basic in its properties. It
dissolves in strong acids and forms salts of bismuth:

Bi2O3 + 6HCl = 2BiCl3 + 3H2O,

Bi2O3 + 6HNO3 = 2Bi(NO3)3 + 3H2O.

The nitrate and chloride of bismuth can be obtained as well-formed colorless
crystals. When treated with water the salts are decomposed in the manner
explained in the following paragraph.

HYDROLYSIS

Many salts such as those of antimony and bismuth form solutions which are
somewhat acid in reaction, and must therefore contain hydrogen ions. This is
accounted for by the same principle suggested to explain the fact that solutions
of potassium cyanide are alkaline in reaction (p. 210). Water forms an
appreciable number of hydrogen and hydroxyl ions, and very weak bases such
as bismuth hydroxide are dissociated to but a very slight extent. When Bi+++
ions from bismuth chloride, which dissociates very readily, are brought in
contact with the OH- ions from water, the two come to the equilibrium
expressed in the equation

Bi+++ + 3OH- <--> Bi(OH)3.

For every hydroxyl ion removed from the solution in this way a hydrogen ion
is left free, and the solution becomes acid in reaction.
Reactions of this kind and that described under potassium cyanide are called
hydrolysis.
DEFINITION: Hydrolysis is the action of water upon a salt to form an acid
and a base, one of which is very slightly dissociated.
Conditions favoring hydrolysis. While hydrolysis is primarily due to the
slight extent to which either the acid or the base formed is dissociated, several
other factors have an influence upon the extent to which it will take place.

1. Influence of mass. Since hydrolysis is a reversible reaction, the relative
masses of the reacting substances influence the point at which equilibrium will
be reached. In the equilibrium

BiCl3 + 3H2O <--> Bi(OH)3 + 3HCl

the addition of more water will result in the formation of more bismuth
hydroxide and hydrochloric acid. The addition of more hydrochloric acid will
convert some of the bismuth hydroxide into bismuth chloride.
2. Formation of insoluble substances. When one of the products of hydrolysis
is nearly insoluble in water the solution will become saturated with it as soon
as a very little has been formed. All in excess of this will precipitate, and the
reaction will go on until the acid set free increases sufficiently to bring about
an equilibrium. Thus a considerable amount of bismuth and antimony
hydroxides are precipitated when water is added to the chlorides of these
elements. The greater the dilution the more hydroxide precipitates. The
addition of hydrochloric acid in considerable quantity will, however,
redissolve the precipitate.
Partial hydrolysis. In many cases the hydrolysis of a salt is only partial,
resulting in the formation of basic salts instead of the free base. Most of these
basic salts are insoluble in water, which accounts for their ready formation.
Thus bismuth chloride may hydrolyze by successive steps, as shown in the
equations

BiCl3 + H2O = Bi(OH)Cl2 + HCl,

BiCl3 + 2H2O = Bi(OH)2Cl + 2HCl,

BiCl3 + 3H2O = Bi(OH)3 + 3HCl.

The basic salt so formed may also lose water, as shown in the equation

Bi(OH)2Cl = BiOCl + H2O.

The salt represented in the last equation is sometimes called bismuth
oxychloride, or bismuthyl chloride. The corresponding nitrate, BiONO3, is
largely used in medicine under the name of subnitrate of bismuth. In these two

compounds the group of atoms, BiO, acts as a univalent metallic radical and is
called bismuthyl. Similar basic salts are formed by the hydrolysis of antimony
salts.

EXERCISES

1. Name all the elements so far studied which possess allotropic forms.
2. What compounds would you expect phosphorus to form with bromine and
iodine? Write the equations showing the action of water on these compounds.
3. In the preparation of phosphine, why is coal gas passed into the flask? What
other gases would serve the same purpose?
4. Give the formula for the salt which phosphine forms with hydriodic acid.
Give the name of the compound.
5. Could phosphoric acid be substituted for sulphuric acid in the preparation of
the common acids?
6. Write the equations for the preparation of the three sodium salts of
orthophosphoric acid.
7. Why does a solution of disodium hydrogen phosphate react alkaline?
8. On the supposition that bone ash is pure calcium phosphate, what weight of
it would be required in the preparation of 1 kg. of phosphorus?
9. If arsenopyrite is heated in a current of air, what products are formed?
10. (a) Write equations for the complete combustion of hydrosulphuric acid,
methane, and arsine. (b) In what respects are the reactions similar?
11. Write the equations for all the reactions involved in Marsh's test for
arsenic.
12. Write the names and formulas for the acids of antimony.
13. Write the equations showing the hydrolysis of antimony trichloride; of
bismuth nitrate.
14. In what respects does nitrogen resemble the members of the phosphorus
family?

CHAPTER XXI
SILICON, TITANIUM, BORON

SYMBOL ATOMIC WEIGHT DENSITYCHLORIDES OXIDES
Silicon Si 28.4 2.35 SiCl4 SiO
Titanium Ti 48.1 3.5 TiCl4 TiO
Boron B 11.0 2.45 BCl3 B2O3

General. Each of the three elements, silicon, titanium, and boron, belongs to a
separate periodic family, but they occur near together in the periodic grouping and are
very similar in both physical and chemical properties. Since the other elements in their
families are either so rare that they cannot be studied in detail, or are best understood
in connection with other elements, it is convenient to consider these three together at
this point.
The three elements are very difficult to obtain in the free state, owing to their strong
attraction for other elements. They can be prepared by the action of aluminium or
magnesium on their oxides and in impure state by reduction with carbon in an electric
furnace. They are very hard and melt only at the highest temperatures. At ordinary
temperatures they are not attacked by oxygen, but when strongly heated they burn
with great brilliancy. Silicon and boron are not attacked by acids under ordinary
conditions; titanium is easily dissolved by them.

SILICON

Occurrence. Next to oxygen silicon is the most abundant element. It does not occur
free in nature, but its compounds are very abundant and of the greatest importance. It
occurs almost entirely in combination with oxygen as silicon dioxide (SiO2), often
called silica, or with oxygen and various metals in the form of salts of silicic acids, or
silicates. These compounds form a large fraction of the earth's crust. Most plants
absorb small amounts of silica from the soil, and it is also found in minute quantities
in animal organisms.
Preparation. The element is most easily prepared by reducing pure powdered quartz
with magnesium powder:

SiO2 + 2Mg = 2MgO + Si.

Properties. As would be expected from its place in the periodic table, silicon
resembles carbon in many respects. It can be obtained in several allotropic forms,
corresponding to those of carbon. The crystallized form is very hard, and is inactive
toward reagents. The amorphous variety has, in general, properties more similar to
charcoal.

Compounds of silicon with hydrogen and the halogens. Silicon hydride (SiH4)
corresponds in formula to methane (CH4), but its properties are more like those of
phosphine (PH3). It is a very inflammable gas of disagreeable odor, and, as ordinarily
prepared, takes fire spontaneously on account of the presence of impurities.
Silicon combines with the elements of the chlorine family to form such compounds as
SiCl4 and SiF4. Of these silicon fluoride is the most familiar and interesting. As stated
in the discussion of fluorine, it is formed when hydrofluoric acid acts upon silicon
dioxide or a silicate. With silica the reaction is thus expressed:

SiO2 + 4HF = SiF4 + 2H2O.

It is a very volatile, invisible, poisonous gas. In contact with water it is partially
decomposed, as shown in the equation

SiF4 + 4H2O = 4HF + Si(OH)4.

The hydrofluoric acid so formed combines with an additional amount of silicon
fluoride, forming the complex fluosilicic acid (H2SiF6), thus:

2HF + SiF4 = H2SiF6.

Silicides. As the name indicates, silicides are binary compounds consisting of silicon
and some other element. They are very stable at high temperatures, and are usually
made by heating the appropriate substances in an electric furnace. The most important
one is carborundum, which is a silicide of carbon of the formula CSi. It is made by
heating coke and sand, which is a form of silicon dioxide, in an electric furnace, the
process being extensively carried on at Niagara Falls. The following equation
represents the reaction

SiO2 + 3C = CSi + 2CO.

The substance so prepared consists of beautiful purplish-black crystals, which are very
hard. Carborundum is used as an abrasive, that is, as a material for grinding and
polishing very hard substances. Ferrosilicon is a silicide of iron alloyed with an excess
of iron, which finds extensive use in the manufacture of certain kinds of steel.
Manufacture of carborundum. The mixture of materials is heated in a large
resistance furnace for about thirty-six hours. After the reaction is completed there is
left a core of graphite G. Surrounding this core is a layer of crystallized carborundum
C, about 16 in. thick. Outside this is a shell of amorphous carborundum A. The
remaining materials M are unchanged and are used for a new charge.

Fig. 73
Silicon dioxide (silica) (SiO2). This substance is found in a great variety of forms in
nature, both in the amorphous and in the crystalline condition. In the form of quartz it
is found in beautifully formed six-sided prisms, sometimes of great size. When pure it
is perfectly transparent and colorless. Some colored varieties are given special names,
as amethyst (violet), rose quartz (pale pink), smoky or milky quartz (colored and
opaque). Other varieties of silicon dioxide, some of which also contain water, are
chalcedony, onyx, jasper, opal, agate, and flint. Sand and sandstone are largely silicon
dioxide.
Properties. As obtained by chemical processes silicon dioxide is an amorphous white
powder. In the crystallized state it is very hard and has a density of 2.6. It is insoluble
in water and in most chemical reagents, and requires the hottest oxyhydrogen flame
for fusion. Acids, excepting hydrofluoric acid, have little action on it, and it requires
the most energetic reducing agents to deprive it of oxygen. It is the anhydride of an
acid, and consequently it dissolves in fused alkalis to form silicates. Being
nonvolatile, it will drive out most other anhydrides when heated to a high temperature
with their salts, especially when the silicates so formed are fusible. The following
equations illustrate this property:

Na2CO3 + SiO2 = Na2SiO3 + CO2,

Na2SO4 + SiO2 = Na2SiO3 + SO3.

Silicic acids. Silicon forms two simple acids, orthosilicic acid (H4SiO4) and
metasilicic acid (H2SiO3). Orthosilicic acid is formed as a jelly-like mass when
orthosilicates are treated with strong acids such as hydrochloric. On attempting to dry
this acid it loses water, passing into metasilicic or common silicic acid:

H4SiO4 = H2SiO3 + H2O.

Metasilicic acid when heated breaks up into silica and water, thus:

H2SiO3 = H2O + SiO2.

Salts of silicic acids,—silicates. A number of salts of the orthosilicic and metasilicic
acids occur in nature. Thus mica (KAlSiO4) is a salt of orthosilicic acid.

Polysilicic acids. Silicon has the power to form a great many complex acids which
may be regarded as derived from the union of several molecules of the orthosilicic
acid, with the loss of water. Thus we have

3H4SiO4 = H4Si3O8 + 4H2O.

These acids cannot be prepared in the pure state, but their salts form many of the
crystalline rocks in nature. Feldspar, for example, has the formula KAlSi3O8, and is a
mixed salt of the acid H4Si3O8, whose formation is represented in the equation above.
Kaolin has the formula Al2Si2O7·2H2O. Many other examples will be met in the study
of the metals.
Glass. When sodium and calcium silicates, together with silicon dioxide, are heated to
a very high temperature, the mixture slowly fuses to a transparent liquid, which on
cooling passes into the solid called glass. Instead of starting with sodium and calcium
silicates it is more convenient and economical to heat sodium carbonate (or sulphate)
and lime with an excess of clean sand, the silicates being formed during the heating:

Na2CO3 + SiO2 = Na2SiO3 + CO2,

CaO + SiO2 = CaSiO3.

Fig. 74
The mixture is heated below the fusing point for some time, so that the escaping
carbon dioxide may not spatter the hot liquid; the heat is then increased and the
mixture kept in a state of fusion until all gases formed in the reaction have escaped.
Molding and blowing of glass. The way in which the melted mixture is handled in the
glass factory depends upon the character of the article to be made. Many articles, such
as bottles, are made by blowing the plastic glass into hollow molds of the desired
shape. The mold is first opened, as shown in Fig. 74. A lump of plastic glass A on the
hollow rod B is lowered into the mold, which is then closed by the handles C. By
blowing into the tube the glass is blown into the shape of the mold. The mold is then
opened and the bottle lifted out. The neck of the bottle must be cut off at the proper
place and the sharp edges rounded off in a flame.
Other objects, such as lamp chimneys, are made by getting a lump of plastic glass on
the end of a hollow iron rod and blowing it into the desired shape without the help of a
mold, great skill being required in the manipulation of the glass. Window glass is
made by blowing large hollow cylinders about 6 ft. long and 1-1/2 ft. in diameter.
These are cut longitudinally, and are then placed in an oven and heated until they
soften, when they are flattened out into plates (Fig. 75). Plate glass is cast into flat
slabs, which are then ground and polished to perfectly plane surfaces.
Varieties of glass. The ingredients mentioned above make a soft, easily fusible glass.
If potassium carbonate is substituted for the sodium carbonate, the glass is much
harder and less easily fused; increasing the amount of sand has somewhat the same
effect. Potassium glass is largely used in making chemical glassware, since it resists
the action of reagents better than the softer sodium glass. If lead oxide is substituted
for the whole or a part of the lime, the glass is very soft, but has a high index of
refraction and is valuable for making optical instruments and artificial jewels.

Fig. 75
Coloring of glass. Various substances fused along with the glass mixture give
characteristic colors. The amber color of common bottles is due to iron compounds in
the glass; in other cases iron colors the glass green. Cobalt compounds color it deep
blue; those of manganese give it an amethyst tint and uranium compounds impart a
peculiar yellowish green color. Since iron is nearly always present in the ingredients,

glass is usually slightly yellow. This color can be removed by adding the proper
amount of manganese dioxide, for the amethyst color of manganese and the yellow of
iron together produce white light.
Nature of glass. Glass is not a definite chemical compound and its composition varies
between wide limits. Fused glass is really a solution of various silicates, such as those
of calcium and lead, in fused sodium or potassium silicate. A certain amount of silicon
dioxide is also present. This solution is then allowed to solidify under such conditions
of cooling that the dissolved substances do not separate from the solvent. The
compounds which are used to color the glass are sometimes converted into silicates,
which then dissolve in the glass, giving it a uniform color. In other cases, as in the
milky glasses which resemble porcelain in appearance, the color or opaqueness is due
to the finely divided color material evenly distributed throughout the glass, but not
dissolved in it. Milky glass is made by mixing calcium fluoride, tin oxide, or some
other insoluble substance in the melted glass. Copper or gold in metallic form
scattered through glass gives it shades of red.

TITANIUM

Titanium is a very widely distributed element in nature, being found in almost all
soils, in many rocks, and even in plant and animal tissues. It is not very abundant in
any one locality, and it possesses little commercial value save in connection with the
iron industry. Its most common ore is rutile (TiO2), which resembles silica in many
respects.
In both physical and chemical properties titanium resembles silicon, though it is
somewhat more metallic in character. This resemblance is most marked in the acids of
titanium. It not only forms metatitanic and orthotitanic acids but a great variety of
polytitanic acids as well.

BORON

Occurrence. Boron is never found free in nature. It occurs as boric acid (H3BO3), and
in salts of polyboric acids, which usually have very complicated formulas.
Preparation and properties. Boron can be prepared from its oxide by reduction with
magnesium, exactly as in the case of silicon. It resembles silicon very strikingly in its
properties. It occurs in several allotropic forms, is very hard when crystallized, and is
rather inactive toward reagents. It forms a hydride, BH3, and combines directly with
the elements of the chlorine family. Boron fluoride (BF3) is very similar to silicon
fluoride in its mode of formation and chemical properties.

Boric oxide (B2O3). Boron forms one well-known oxide, B2O3, called boric
anhydride. It is formed as a glassy mass by heating boric acid to a high temperature. It
absorbs water very readily, uniting with it to form boric acid again:

B2O3 + 3H2O = 2H3BO3.

In this respect it differs from silicon dioxide, which will not combine directly with
water.
Boric acid (H3BO3). This is found in nature in considerable quantities and forms one
of the chief sources of boron compounds. It is found dissolved in the water of hot
springs in some localities, particularly in Italy. Being volatile with steam, the vapor
which escapes from these springs has some boric acid in it. It is easily obtained from
these sources by condensation and evaporation, the necessary heat being supplied by
other hot springs.
Boric acid crystallizes in pearly flakes, which are greasy to the touch. In the laboratory
it is easily prepared by treating a strong, hot solution of borax with sulphuric acid.
Boric acid being sparingly soluble in water crystallizes out on cooling:

Na2B4O7 + 5H2O + H2SO4 = Na2SO4 + 4H3BO3.

The substance is a mild antiseptic, and on this account is often used in medicine and
as a preservative for canned foods and milk.
Metaboric and polyboric acids. When boric acid is gently heated it is converted into
metaboric acid (HBO2):

H3BO3 = HBO2 + H2O.

On heating metaboric acid to a somewhat higher temperature tetraboric acid (H2B4O7)
is formed:

4HBO2 = H2B4O7 + H2O.

Many other complex acids of boron are known.
Borax. Borax is the sodium salt of tetraboric acid, having the formula Na2B4O7·10
H2O. It is found in some arid countries, as southern California and Tibet, but is now
made commercially from the mineral colemanite, which is the calcium salt of a
complex boric acid. When this is treated with a solution of sodium carbonate, calcium
carbonate is precipitated and borax crystallizes from the solution.

When heated borax at first swells up greatly, owing to the expulsion of the water of
crystallization, and then melts to a clear glass. This glass has the property of easily
dissolving many metallic oxides, and on this account borax is used as a flux in
soldering, for the purpose of removing from the metallic surfaces to be soldered the
film of oxide with which they are likely to be covered. These oxides often give a
characteristic color to the clear borax glass, and borax beads are therefore often used
in testing for the presence of metals, instead of the metaphosphoric acid bead already
described.
The reason that metallic oxides dissolve in borax is that borax contains an excess of
acid anhydride, as can be more easily seen if its formula is written 2NaBO2 + B2O3.
The metallic oxide combines with this excess of acid anhydride, forming a mixed salt
of metaboric acid.
Borax is extensively used as a constituent of enamels and glazes for both metal ware
and pottery. It is also used as a flux in soldering and brazing, and in domestic ways it
serves as a mild alkali, as a preservative for meats, and in a great variety of less
important applications.

EXERCISES

1. Account for the fact that a solution of borax in water is alkaline.
2. What weight of water of crystallization does 1 kg. of borax contain?
3. When a concentrated solution of borax acts on silver nitrate a borate of silver is
formed. If the solution of borax is dilute, however, an hydroxide of silver forms.
Account for this difference in behavior.

CHAPTER XXII
THE METALS

The metals. The elements which remain to be considered are known collectively as
the metals. They are also called the base-forming elements, since their hydroxides are
bases. A metal may therefore be defined as an element whose hydroxide is a base.
When a base dissolves in water the hydroxyl groups form the anions, while the
metallic element forms the cations. From this standpoint a metal can be defined as an
element capable of forming simple cations in solution.
The distinction between a metal and a non-metal is not a very sharp one, since the
hydroxides of a number of elements act as bases under some conditions and as acids
under others. We have seen that antimony is an element of this kind.
Occurrence of metals in nature. A few of the metals are found in nature in the free
state. Among these are gold, platinum, and frequently copper. They are usually found
combined with other elements in the form of oxides or salts of various acids. Silicates,
carbonates, sulphides, and sulphates are the most abundant salts. All inorganic
substances occurring in nature, whether they contain a metal or not, are called
minerals. Those minerals from which a useful substance can be extracted are called
ores of the substance. These two terms are most frequently used in connection with
the metals.
Extraction of metals,—metallurgy. The process of extracting a metal from its ores is
called the metallurgy of the metal. The metallurgy of each metal presents peculiarities
of its own, but there are several methods of general application which are very
frequently employed.
1. Reduction of an oxide with carbon. Many of the metals occur in nature in the form
of oxides. When these oxides are heated to a high temperature with carbon the oxygen
combines with it and the metal is set free. Iron, for example, occurs largely in the form
of the oxide Fe2O3. When this is heated with carbon the reaction expressed in the
following equation takes place:

Fe2O3 + 3 C = 2 Fe + 3 CO.

Many ores other than oxides may be changed into oxides which can then be reduced
by carbon. The conversion of such ores into oxides is generally accomplished by
heating, and this process is called roasting. Many carbonates and hydroxides
decompose directly into the oxide on heating. Sulphides, on the other hand, must be

heated in a current of air, the oxygen of the air entering into the reaction. The
following equations will serve to illustrate these changes in the case of the ores of
iron:

FeCO3 = FeO + CO2,

2Fe(OH)3 = Fe2O3 + 3H2O,

2FeS2 + 11O = Fe2O3 + 4SO2.

2. Reduction of an oxide with aluminium. Not all oxides, however, can be reduced by
carbon. In such cases aluminium may be used. Thus chromium may be obtained in
accordance with the following equation:

Cr2O3 + 2 Al = 2 Cr + Al2O3.

This method is a comparatively new one, having been brought into use by the German
chemist Goldschmidt; hence it is sometimes called the Goldschmidt method.
3. Electrolysis. In recent years increasing use is being made of the electric current in
the preparation of metals. In some cases the separation of the metal from its
compounds is accomplished by passing the current through a solution of a suitable salt
of the metal, the metal usually being deposited upon the cathode. In other cases the
current is passed through a fused salt of the metal, the chloride being best adapted to
this purpose.
Electro-chemical industries. Most of the electro-chemical industries of the country
are carried on where water power is abundant, since this furnishes the cheapest means
for the generation of electrical energy. Niagara Falls is the most important locality in
this country for such industries, and many different electro-chemical products are
manufactured there. Some industries depend upon electrolytic processes, while in
others the electrical energy is used merely as a source of heat in electric furnaces.
Preparation of compounds of the metals. Since the compounds of the metals are so
numerous and varied in character, there are many ways of preparing them. In many
cases the properties of the substance to be prepared, or the material available for its
preparation, suggest a rather unusual way. There are, however, a number of general
principles which are constantly applied in the preparation of the compounds of the
metals, and a clear understanding of them will save much time and effort in
remembering the details in any given case. The most important of these general
methods for the preparation of compounds are the following:
1. By direct union of two elements. This is usually accomplished by heating the two
elements together. Thus the sulphides, chlorides, and oxides of a metal can generally

be obtained in this way. The following equations serve as examples of this method:

Fe + S = FeS,

Mg + O = MgO,

Cu + 2Cl = CuCl2.

2. By the decomposition of a compound. This decomposition may be brought about
either by heat alone or by the combined action of heat and a reducing agent. Thus
when the nitrate of a metal is heated the oxide of the metal is usually obtained. Copper
nitrate, for example, decomposes as follows:

Cu(NO3)2 = CuO + 2NO2 + O.

Similarly the carbonates of the metals yield oxides, thus:

CaCO3 = CaO + CO2.

Most of the hydroxides form an oxide and water when heated:

2Al(OH)3 = Al2O3 + 3H2O.

When heated with carbon, sulphates are reduced to sulphides, thus:

BaSO4 + 2C = BaS + 2CO2.

3. Methods based on equilibrium in solution. In the preparation of compounds the first
requisite is that the reactions chosen shall be of such a kind as will go on to
completion. In the chapter on chemical equilibrium it was shown that reactions in
solution may become complete in either of three ways: (1) a gas may be formed which
escapes from solution; (2) an insoluble solid may be formed which precipitates; (3)
two different ions may combine to form undissociated molecules. By the judicious
selection of materials these principles may be applied to the preparation of a great
variety of compounds, and illustrations of such methods will very frequently be found
in the subsequent pages.
4. By fusion methods. It sometimes happens that substances which are insoluble in
water and in acids, and which cannot therefore be brought into double decomposition
in the usual way, are soluble in other liquids, and when dissolved in them can be
decomposed and converted into other desired compounds. Thus barium sulphate is not
soluble in water, and sulphuric acid, being less volatile than most other acids, cannot
easily be driven out from this salt When brought into contact with melted sodium

carbonate, however, it dissolves in it, and since barium carbonate is insoluble in
melted sodium carbonate, double decomposition takes place:

Na2CO3 + BaSO4 = BaCO3 + Na2SO4.

On dissolving the cooled mixture in water the sodium sulphate formed in the reaction,
together with any excess of sodium carbonate which may be present, dissolves. The
barium carbonate can then be filtered off and converted into any desired salt by the
processes already described.
5. By the action of metals on salts of other metals. When a strip of zinc is placed in a
solution of a copper salt the copper is precipitated and an equivalent quantity of zinc
passes into solution:

Zn + CuSO4 = Cu + ZnSO4.

In like manner copper will precipitate silver from its salts:

Cu + Ag2SO4 = 2Ag + CuSO4.

It is possible to tabulate the metals in such a way that any one of them in the table will
precipitate any one following it from its salts. The following is a list of some of the
commoner metals arranged in this way:

Zinc
Iron
Tin
Lead
Copper
Bismuth
Mercury
Silver
Gold

According to this table copper will precipitate bismuth, mercury, silver, or gold from
their salts, and will in turn be precipitated by zinc, iron, tin, or lead. Advantage is
taken of this principle in the purification of some of the metals, and occasionally in the
preparation of metals and their compounds.
Important insoluble compounds. Since precipitates play so important a part in the
reactions which substances undergo, as well as in the preparation of many chemical
compounds, it is important to know what substances are insoluble. Knowing this, we
can in many cases predict reactions under certain conditions, and are assisted in