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INORGANIC CHEMISTRY

WORKS BY
G. S. NEWTH, F.I.C., F.CS-,

Late Senior Demonstrator in the Royal
College of Science, London.

Chemical Lecture Experiments. With
230 Illustrations. Crown 8vo.

Elementary Practical Chemistry. With
108 Illustrations and 254 Experiments.
Crown 8vo.

A Manual of Chemical Analysis, Quali-
tative and Quantitative. With 102
Illustrations. Crown Svo,

Smaller Chemical Analysis. Crown Svo.
A Text-book of Inorganic Chemistry.

With 155 Illustrations. Crown Svo.

LONGMANS, GREEN AND CO.,

LONDON, NEW YORK, BOMBAY,
CALCUTTA, AND MADRAS.

A TEXT-BOOK

INORGANIC CHEMISTRY

G. S. NEWTH, F.I.C., F.C.S.

NEW IMPRESSION

LONGMANS, GREEN AND CO.

39 PATERNOSTER ROW, LONDON
FOURTH AVENUE & 3oTH STREET, NEW YORK
BOMBAY, CALCUTTA, AND MADRAS

1920.
AH rights reserved

PREFACE

IN drawing up a systematic course of elementary chemical
instruction based upon the periodic classification of the ele-
ments, whether it be as a course of lectures, or as a text-book,
a number of serious difficulties are at once encountered.
These possibly are sufficient to account .for the fact, that
although twenty-five years have elapsed since 'Mendelejeff
published this natural system of classification, the method has
not been generally adopted as the basis of English elementary
text-books.

I have endeavoured to obviate many of these difficulties,
while still making the periodic system the foundation upon
which this little book is based, by dividing the book into
three parts. Part I. contains a brief sketch of the funda-
mental principles and theories upon which the science of
modem chemistry is built. Into this portion of the book I
have introduced, necessarily in briefest outlines, some of the
more recent developments of the science in a physico-chemical
direction, of which it is desirable that the student should gain
some knowledge, even early in his career.

Part II. consists of the £tudy of the four typical elements,
hydrogen, oxygen, nitrogen, and carbon, and of their more
important compounds. By dissociating these four elements
from their position in the periodic system, and treating them
separately, the student is early brought into contact with many
of the simpler and more familiar portions of the science. Such

v! Preface

subjects as water, the atmosphere, and combustion, to which it
is desirable that he should be introduced at an early stage in
his studies, are thus brought much more forward than would
otherwise be the case.

In Part III. the elements are treated systematically, accord-
ing to the periodic classification. In this manner, while
avoiding a sharp separation of the elements into the two arbi-
trary classes of metals and non-metals, it has been possible to
so far conform to the prevailing methods of instruction, that
all those elements which are usually regarded as non-metals
(with the two exceptions of boron and silicon) are treated in
the earlier portion of the book:

The science of chemistry has of recent years developed and
become extended to such a degree, that the difficulty of giving
a fairly balanced treatment of the subject, within the limits of
a small text-book, is an ever-increasing one, and it necessarily
resolves itself into a question of the judicious selection of
matter. In making such a selection, I have endeavoured, as
far as possible, to keep in view the requirements of students
at the present time, without, however, following any examina-
tion syllabus.

Acting upon this principle, I have omitted all detailed
description of the rare elements and their compounds, con-
fining myself merely to a short mention of them in a few
general remarks at the commencement of the various chapters.

Although from a purely scientific standpoint many of these
rare substances are of the greatest interest and importance,
it must be admitted that they stand quite outside the range
of all the customary courses of chemical instruction; and so
far as the wants of the ordinary student are concerned, the
space which would be occupied by an account of these
elements is more advantageously devoted to such matters

Preface vii

as are discussed in the Introductory Outlines. Moreover, it
is a matter of common observation that text-books, even
upon the shelves of reference libraries, and which bear un-
mistakable evidence of much use, are frequently uncut in those
portions which treat of these elements.

Details of metallurgical processes, also, are out of place
in a text-book of chemistry, and must be sought in metal-
lurgical text-books. Only such condensed outlines, therefore,
have been given as are sufficient to explain the chemica)
changes that a're involved in these operations.

The great importance to the student, of himself perform-
ing experiments illustrating the preparation and properties of
many of the substances treated of in his text-book, cannot
well be over-estimated. If he be in attendance upon a course
of chemical lectures, opportunity should be given to him for
repeating the simpler experiments he may see performed
upon the lecture table : if he be not attending lectures, the
necessity for this practical work on his part is greater still.
Instead of burdening this text-book with specific directions
for carrying out such elementary experiments, frequent refer-
ences have been made to my "Chemical Lecture Experi-
ments," where minute directions are given for carrying out
a large number of experiments, many of which may be easily
performed, and with the very simplest of apparatus.

Several of the woodcuts have been borrowed from existing
modern works, such as Thorpe's " Dictionary of Applied
Chemistry," Mendelejeff s " Principles of Chemistry," Ost-
wald's "Solutions," and others. Care has been taken, how-
ever, to exclude all antiquated cuts, and a large number of
the illustrations are from original drawings and photographs.

G. S. N.

SOUTH KENSINGTON.

HINTS TO STUDENTS

FOR the help of students who may use this book at the
commencement of their chemical studies, and especially for
those who may not be working under the immediate guidance
of a teacher, the following hints are given : —

Begin by carefully reading the first four chapters (pages
1-24). Then pass on to Part II. (page 171), and begin
the study of the four typical elements, hydrogen, oxygen,
nitrogen, and carbon, and their compounds, in the order in
which they are treated. Accompany your reading by per-
forming as many of the experiments referred to as possible,
in order that you may become practically familiar with the
substances you are studying.

During the time occupied in the study of these four
elements and their compounds, again read Chapters I. to IV.,
and slowly and carefully continue reading Part L, so that
by the time Part III. is reached, you may have fairly mastered
at least the first thirteen chapters of the Introductory Out-
lines.

The order in which the elements are treated in Part III.
is based upon the periodic classification, therefore read the
short introductory remarks at the commencement of the
various chapters, in the light of the table on page 118.

Throughout the book temperatures are given in degrees
of the Centigrade thermometer. i° Centigrade equals J.8*

x Hints to Students

Fahrenheit, and as the zero of the latter scale is 32° below
that of the Centigrade, temperatures given in degrees of one
scale are readily translated into degrees of the other, by
the simple formula —

The abbreviation mm. stands for millimetre; the y^Vff part
of a metre (i 1116^6 = 39.370113 inches; or roughly, 25
mm. = i inch). The abbrevation c.c. signifies cubic centi-
metre; the j-oVu part of a "cubic decimetre, or litre (i
litre = 1.76077 pints).

i gramme (the weight of i c.c. of distilled water, taken at
its point of maximum density) = 15.43235 English grains.

TABLE OF CONTENTS

PART 1
INTRODUCTORY OUTLINES

CHAP. |.AGjL

I. Chemical Change — The Constitution , of Matter — Molecules —

Atoms . , * . ~i . • .... i
II. Elements and Compounds — Mixtures — Chemical "Affinity —

Modes of Chemical Action 6

III. Chemical Nomenclature 15

IV. Chemical Symbols . , . . . . .21
V. The Atomic Theory — Laws of Chemical Action 25

VI. Atomic Weights— Modes of Determining Atomic Weights . 34

VII. Quantitative Chemical Notation -53

VIII. Valency of the Elements . . . . . . -59

IX. General Properties of Gases — Relation to Heat and Pressure —

Liquefaction — Diffusion — The Kinetic Theory . . 69
X. Dissociation — Reversible or Balanced Actions ... 88
XL Electrolysis — Electrolytic Dissociation — The Ionic Theory , 96
XII. Classification of the Elements — The Periodic System . . 112
Kill. General Properties of Liquids — Evaporation, Boiling, Vapour
Pressure of Solutions — The Passage of Liquids into Solids
— Freezing Point of Solutions — Raoult's Method . .126
XIV. Solution — Gases in Liquids — Liquids in Liquids — Solids in

Liquids — Osmotic Pressure — Crystalline Forms . ... 142

XV. Thermo-chemistry . 163

PART II

THE STUDY OF FOUR TYPICAL ELEMENTS
Hydrogen — Oxygen — Nitrogen — Carbon,

AND THEIR MORE IMPORTANT COMPOUNDS.

I. Hydrogen — Hydrogenium . . » . • * 171

II. Oxygen — Allotropy — Ozone . . „ • « . . 181

III Compounds of Hydrogen with Oxygen . , . • . . . 203

xii Contents

CHAP. PAGE

IV. Nitrogen ...'', . . . , « , . 229

V. Oxides and Oxy-acids of Nitrogen ...... 234

VI. The Atmosphere and the Argon Group of Elements . . 252
VII. Compounds of Nitrogen and Hydrogen — Hydroxylamine —

Ammon-sulphonates ; Halogen Compounds of Nitrogen . 272

VIII. Carbon 285

IX. Carbon Monoxide — Carbon Dioxide — Carbonates . . . 295
X. Compounds of Carbon with Hydrogen — Methane — Ethylene —

Acetylene 312

XI. Combustion — Heat of Combustion — Ignition Point — Flame —
Structure of Flame — Ca'use of Luminosity of Flames —
The Bunsen Flame . . . Vv •':" • V . .321

PART III

THE SYSTEMATIC STUDY OF THE ELEMENTS, BASED
UPON THE PERIODIC CLASSIFICATION

L ELEMENTS OF GROUP VII. (FAMILY B.)

Fluorine : Hydrofluoric Acid. Chlorine : Hydrochloric
Acid — Oxides and Oxyacids of Chlorine. Bromine:
Hydrobromic Acid — Oxyacids of Bromine. Iodine:
Hydriodic Acid — Oxyacids of Iodine — Periodates . . 34$

II. ELEMENTS OF GROUP VI. (FAMILY B.)

Sulphur: Compounds of Sulphur with Hydrogen — Com-
pounds with Chlorine — Oxides and Oxyacids of Sulphur
— Oxychlorides — Carbon Bisulphide. Selenium — Tellu-
rium -597

III. ELEMENTS OF GROUP V. (FAMILY B.)

Phosphorus: Compounds with Hydrogen — Compounds with
the Halogens — Oxides and Oxyacids. Arsenic ; Arsenu-
retted Hydrogen — Halogen Compounds — Oxides and
Oxyacids— Sulphides. Antimony: Antimony Hydride —
Halogen Compounds — Oxides and Acids — Sulphides.
Bismuth : Bismuth and Halogens — Oxides — Sulphides . 450

IV. ELEMENTS OF GROUP I. (FAMILY A.)

Potassium — Sodium — Lithium — Rubidium — Ammonium

Salts ..... r .... 505
V. ELEMENTS OF GROUP I. (FAMILY B.)

Copper — Silver — Gold ... . c , . . « 549

Contents xiii

CHAP. PAGE

VI. ELEMENTS OF GROUP II. (FAMILY A.)

Beryllium — Magnesium — Calcium — Strontium — Barium 570

VII. ELEMENTS OF GROUP II. (FAMILY B.)

Zinc — Cadmium — Mercury . , • 5 . coo

VIII. ELEMENTS OF GROUP III.

FAM I LY A. : Scandium — Yttrium — Lanthanum — Ytter-
bium.
FAMILY B. : Boron — Aluminium — Gallium — Indium —

Thallium ••*•••;'. 606

IX. ELEMENTS OF GROUP IV.

FAMILY A. : Titanium — Zirconium — Cerium — Thorium-
FAMILY B. : Silicon — Germanium — Tin — Lead . . 627

X. ELEMENTS. OF GROUP V. (FAMILY A.)

Vanadium — Niobium — Tantalum 655

XI. ELEMENTS OF GROUP VI. (FAMILY A.)

Chromittm — Molybdenum — Tungsten — Uranium . • 657

XII. ELEMENTS OF GROUP VII. (FAMILY A.)

Manganese .......... 666

XIII. TRANSITIONAL ELEMENTS OF THE FIRST LONG PERIOD.

Iron— Cobalt— Nickel . .671

"XIV. TRANSITIONAL ELEMENTS OF THE SECOND AND FOURTH

LONG PERIOD.

Ruthenium — Rhodium — Palladium — Osmium — Iridium —
Platinum — Argon — Helium 690

APPENDIX : RADIUM, AND RADIOACTIVE ELEMENTS . 697

INDEX . . . . . . . . . . . » , 705

INORGANIC CHEMISTRY

PAET I
INTRODUCTORY OUTLINES

CHAPTER I
CONSTITUTION OF MATTER

THE science of chemistry may be described as the study of a
certain class of changes which matter is capable of undergoing.
Matter is susceptible of a variety of changes, some of which are
regarded as physical and others as chemical. Thus, when a steel
knitting-needle is rubbed upon a magnet, the needle undergoes a
change, by virtue of which it becomes endowed with the power
of attracting to itself iron filings or nails ; and when an ordinary
lucifer match is rubbed upon a match-box the match undergoes a
change, resulting in the production of flame. In the first case the
change is said to be a physical one, while the ignition and com-
bustion of the match is a chemical change.

When a fragment of ice is gently warmed, it is changed from a
hard, brittle solid to a mobile, transparent liquid ; and when white
of egg is gently heated, it changes from a transparent, colourless
liquid to an opaque white solid. These changes, which appear at
first sight to be of a similar order, are in reality essentially different
in their nature : the transformation of solid ice into liquid water
is a physical change, the coagulation of albumen is a chemical
change.

Again, when certain substances (such as the materials which
constitute the so-called luminous paint} are exposed to a bright
light, they undergo a change whereby they become invested with

2 Introductory Outlines

the power to emit a feeble light when seen in the dark. A stick of
phosphorus also emits a very similar light when seen in the dark.
The glowing of these materials under these circumstances might
readily be regarded as the result of the same kind of change in
both cases ; but in reality the luminosity of the phosphorus is due
to a chemical change taking place upon the surface of that sub-
stance, while the emission of light from the luminous paint is a
purely physical phenomenon.

The two sciences, chemistry and physics, are so closely related
and interdependent upon each other, that no sharp distinction or
line of separation between them is possible. Every chemical
change that takes place is attended by some physical change, and
it often happens that this accompanying physical change forms
the only indication of the chemical change that has taken place.
In certain important points, however, a chemical change is very
different from one that is purely physical : in the latter case no
material alteration in the essential nature of the substance takes
place. This will be seen in the examples quoted. The steel
needle remains unaltered in its essence, although by magnetisation
it has acquired a new property — a property which it again loses,
and which can be again and again imparted to it. The match, on
the other hand, when ignited has undergone a material and per-
manent change: the combustible substance is now no longer
combustible, neither will it ever return to its original state. The
solid water, in being transformed to liquid water, has not under-
gone any vital change ; in essence it is the same substance merely
endowed with a new property of liquidity, a property which it loses
again when cooled, and which can be again and again imparted to
it. On the other hand, the coagulated albumen has undergone a
complete and lasting change, and never returns to its original
condition.

In the same way, the luminous paint gradually ceases to emit
light, and returns to its original state ; it may be exposed to the
influence of light, when it once more acquires the property of
phosphorescence, and this change may be brought about indefi-
nitely, without altering the intrinsic nature of the substance. The
glowing phosphorus, on the other hand, is gradually changed into
a white substance, which escapes from it as a smoke or fume ; in
the act of glowing the phosphorus is undergoing a process of slow
burning, and if allowed to remain will continue glowing and burn-
ing until the whole of it has disappeared in the form of smoke.

Molecules 3

The Constitution of Matter. Molecules.— Matter is regarded
by the chemist and physicist as being composed of aggregations
of minute particles ; every substance, whether it be solid, liquid,
or gaseous, presents the appearance to his mind of a vast number
of extremely minute particles. To these particles the name mole-
cules (" little masses ") has been given. The particles or molecules
of any particular substance are all alike : thus in sulphur the
molecules are all of one kind, while in water they are all of another
kind ; the chemical properties associated with sulphur are the pro-
perties of the individual sulphur molecules, while those belonging
to water are the properties of the molecules of that substance. All
matter, therefore, is to be conceived as having what may be called a
grained structure. The actual sizes of molecules is a matter which
has not yet been determined with exactness ; they are orders of
magnitude which are as difficult for the mind to grasp on account
of their minuteness, as many astronomical measurements are by
reason of their vastness. It is certain that their size is less than
half a single wave-length of light,* and that therefore they are
beyond the visual limits of the microscope. Some general idea
of their order of magnitude may be gathered from Lord Kelvin's
calculation, that if a single drop of water were magnified to the size
of the earth, each molecule/being proportionately enlarged, the
grained appearance which the mass would present would probably
be finer than that of a heap of cricket-balls, but coarser than a
heap of small shot.

It will be evident, therefore-, that in the strictest sense matter is
not homogeneous^ since it consists of aggregations of molecules,
between which there exist certain interspaces.

The forces which similar molecules exert upon each other are
regarded as physical^ in contradistinction to chemical. These
forces are either attractive in their nature, or repellent. The
attractive forces tend to draw the molecules closer together, and
thus to cause the substance to assume the solid state ; while
repellent forces, on the other hand, tend to separate the molecules
and to make the substance pass into the gaseous condition.
Changes which matter undergoes by the action of these forces are
physical changes ; they do not affect the chemical nature and
properties of the substance, which properties, as already stated,
reside in the molecules themselves."

* The wave-length of the blue ray (G) = 0.0004311 millimetre, or
0.0000169 inch.

4 Introductory Outlines

In each of the three states of matter, viz. solid, liquid, or gaseous,
the molecules are conceived as being in a state of motion ; they
are regarded as executing some vibratory movement within the
spaces that divide them. In the solid state this movement is
usually the most restricted, for the reason that in this case the
intermolecular spaces are as a rule the smallest. In the gaseous
condition, however, the attractive force between the molecules
has been almost entirely overcome by the operation of the
repellent forces. The molecules are therefore widely sepa-
rated, and consequently permit of a much greater freedom of
movement.

Such changes in matter, 'which are merely the result of
alterations in the motions of the molecules, are likewise purely
physical changes.

Molecules may be defined as the smallest weight of matter
in which the original properties of the matter are retained.

Atoms. — It is the belief of chemists that most molecules are
possessed of a structure. That is to say, they are not simple,
single, indivisible masses, but themselves consist of aggregations
of still smaller particles, which are held together by the opera-
tions of some other force. These particles of which molecules
are composed are termed atoms, and the force which holds them
together is called chemical affinity, or chemical attraction. To
the mind of the chemist, such molecules are little systems, con-
sisting of a number of atoms which are attracted to each other
by this particular force ; in the ordinary movements of the mole-
cule, the system moves about as a whole. In this respect it bears
some analogy, on an infinitely minute scale, to a solar system.
The atoms of a molecule are regarded as in a state of motion as
respects one another, possibly revolving about one another, while
the entire system, or molecule, at the same time performs its in-
dependent movements, just as in a solar system the various
members perform various movements towards each other, while
at the same time the whole system travels upon its prescribed
orbit. In the case of the heavenly bodies, the force which regulates
the movements of the individual members of the system amongst
themselves is the same force that controls the motion of the united
system, namely, gravitation. What is the precise relation, or
difference, if any, between the forces which control the move-
ments of molecules, and those which operate between the atoms
of the molecule, is not known j but as the effects produced are

Molecules and Atoms 5

different, the latter force is distinguished by the name of chemical
affinity.

Any change which matter undergoes, in which the integrity of
the molecules is not destroyed, is regarded as a physical change ;
while any change which arises from an alteration in the structure
of the molecule is a chemical change. For example, the molecules
of water consist of three separate atoms, one of oxygen and two
of hydrogen ; any change which water can be made to undergo,
in which these three atoms still remain associated together as the
molecule, is a physical change. The water may be converted into
ice, or it may be changed into steam ; but these alterations still
leave the molecules intact — the three atoms still remain united as
an unbroken system, and so long as this is the case chemical
change has not taken place.

Suppose now the molecules of water are heated to a much
higher temperature than that which is necessary to convert the
water into steam, by passing electric sparks through the steam.
It will then be found that a very different kind of change has come
over the substance. The steam, after being so heated, no longer
condenses to water again when cooled ; it has been changed into
a gas which can be bubbled through water and collected in an
inverted vessel filled with water standing in a pneumatic trough,
and if a flame be applied to this gas a sharp explosion takes place.
The change in this case is a chemical change, for the integrity
of the molecules of water has been destroyed. The two atoms
of hydrogen have become detached from the oxygen atom, and
the original triune structure of the system is destroyed.

Atoms are therefore defined as the smallest particles of matter
which can take part in a chemical change* *

* The study of the phenomena of radioactivity has led to the belief that
atoms are not indivisible particles of matter, but that they are themselves
systems, which under certain circumstances are capable of undergoing change
by ejecting from themselves relatively minute portions of the system called
electrons. (See p. 104; also Appendix.) The precise nature of these electrons
still belongs to the realm of speculation, and the changes resulting from their
movements do not belong to the category of "chemical change" as the term
is here employed.

CHAPTER II
ELEMENTS AND COMPOUNDS

THERE are certain molecules in which all the atoms present are
of the same kind, and there are other molecules which are com-
posed of atoms which differ from one another. Thus, in the
substance sulphur, all the atoms composing the molecules are
alike ; while in water, as already mentioned, there are two distinct
kinds of atoms in the molecule. Matter, therefore, is divided into
two classes, according as to whether its molecules are composed of
similar or of dissimilar atoms. Molecules consisting of atoms of
the same kind are termed elementary molecules, and substances
whose molecules are so constituted are known as elements ; mole-
cules, on the other hand, which contain dissimilar elements are
called compound molecules, and substances whose molecules are
thus composed are distinguished as compounds.

Sulphur, therefore, is an element, and water is a compound. It
will be evident that in the case of elementary molecules, whatever
processes they may be subjected to, only one kind of matter can
be obtained from them ; while in the case of compounds, the
molecules consisting of dissimilar atoms, a;- many different kinds
of matter can be obtained as there are different atoms present.
By appropriate means the atoms of hydrogen and oxygen in water
molecules can be separated, and two totally different kinds of
matter, namely, hydrogen and oxygen, can be obtained from this
compound.

At the present time there are about seventy substances known to
chemists which are believed to be elements. In the history of the
science it has frequently happened that substances which were
considered to be elements have proved, when subjected to new
methods of investigation, to be in reality compound bodies : thus,
prior to the year 1783, water was thought to be an elementary
substance ; it was indeed regarded as the very type of an element,
until Cavendish and Lavoisier proved that it was composed of
two entirely different kinds of matter. In the year 1807 Sir

Elements and Compounds 7

Humphry Davy showed that the substances known as potash
and soda, which were believed to be elements, were in reality
compound bodies, and he succeeded in separating the constituent
atoms in the molecules of these substances, and in obtaining from
them two essentially different kinds of matter. It is therefore
quite possible, perhaps even probable, that some at least of the
forms of matter which are now held to be elements may yet prove
to be compound bodies. On the other hand, the list is from time
to time extended by the discovery of new elements. Thus during
the last few years at least five new members have been added to
the number.

The number of compounds is practically infinite.

The elements are very unequally distributed in nature, and are
of very different degrees of importance to mankind. Some are
absolutely essential to life as it is constituted, while others might
be blotted out of creation without, so far as is known, their absence
being appreciated. The following thirty elements include all the
most important (for the complete list see page 22) : —

Aluminium. Gold. Oxygen.

Antimony. Hydrogen. Phosphorus.

Arsenic. Iodine. Platinum.

Bismuth. Iron. Potassium.

Bromine. Lead. Silicon.

Calcium. Magnesium. Silver.

Carbon. Manganese. Sodium.

Chlorine. Mercury. Sulphur.

Copper. Nickel. Tin.

Fluorine. Nitrogen. Zinc.

On account of certain properties common to a large number of
the elements, and more or less absent in others, properties which
are for the most part physical in character, the elements are
divided into two classes, known as metals and non-metals. The
metals generally are opaque, and their smoothed surfaces reflect
light to a high degree, thus giving them the appearance known as
metallic lustre. They also conduct heat and electricity. Gold,
silver, copper, iron, are metals ; sulphur, bromine, oxygen, phos-
phorus, are non-metals. These two classes, however, gradually
merge into one another, and certain elements are sometimes
placed in one division and sometimes in the other, depending
upon whether the distinction is based more upon their physical

8 Introductory Outlines

or their chemical properties : thus, the element arsenic possesses
many of the physical properties of a metal, but in its chemical
relations it is more allied to the non-metals ; such elements as
these are often distinguished by the name metalloids. The follow-
ing list embraces all those elements which by common consent
are regarded as non-metals and metalloids, including the recently
discovered elements of the argon group, which are here printed in
italics : —

Arsenic.
Boron.
Bromine.
Carbon.
Chlorine.

Fluorine.
Hydrogen.
Iodine.
Nitrogen.
Oxygen.

Phosphorus.
Selenium.
Silicon.
Sulphur.
Tellurium.

Helium.
Neon.
Argon.
Krypton.
Xenon.

The number of atoms which compose the various elementary
molecules is not the same in 'all cases : thus in the elements
sodium, potassium, cadmium, mercury, and zinc, the molecules,
when the elements are in a state of vapour, consist of only one
atom. The same is true also of the newly discovered elements in
the last column. The molecules of all these substances are single
particles of matter. The terms molecule and atom, therefore, as
applied to these elements, are synonymous. Such molecules as
these are called mono-atomic molecules. In many cases elemen-
tary molecules consist of two atoms ; such is the case with the
elements hydrogen, bromine, chlorine, oxygen, nitrogen, and
others. Elementary molecules of this twin or dual nature are
known as di-atomic molecules. Only one instance is known in
which an elementary molecule consists of a trio of atoms, namely,
the molecule of ozone, which is an aggregation of three oxygen
atoms. This molecule is said to be tri-atomic. In two cases,
namely, arsenic and phosphorus, the molecules are composed of
a quartette of atoms, and these elements, therefore, are said to
form tetr-atomic molecules. In a large number of instances the
atomic constitution of the molecule of the elements is not known.
These terms, mono-atomic^ di-atomic, &c., are applied exclu-
sively to molecules of elements, and are not used in reference
to compounds, where the molecules are composed of dissimilar
atoms.

Mechanical Mixtures.— When .molecules of different kinds of
matter are brought together, one of two results may follow : either
they will merely mingle together without losing their identity, that

Mechanical Mixtures 9

is to say, the atoms composing the individual molecules will still
remain associated together as before, or the atoms in the molecules
of one kind will attach themselves to certain atoms present in
molecules of another kind to form still different molecules ; in other
words, there will be a redistribution of the atoms, whereby diffe-
rent systems or molecules are produced.

In the first case the result is said to be a simple or mechani-
cal mixture, in the second it is the formation of a chemical
compound.

In a simple mixture the ingredients can be again separated by
purely mechanical methods ; and as the properties of a substance
are the properties of the molecules of that substance, it follows that
if the integrity of the molecules is not broken, the properties of a
mechanical mixture will be those of the ingredients. For example,
oxygen is a colourless gas without taste or smell ; hydrogen also is
a colourless gas without taste or smell : when these two gases are
mixed together, the mixture is gaseous, is colourless, and tasteless,
and, being only a mixture, the molecules of one gas can be readily
sifted away from the other.

Again, charcoal is a black solid, insoluble in water ; sulphur is a
yellow solid, also insoluble in water ; nitre is a white solid, readily
dissolved by water : when these three substances are finely
powdered and mixed together, the result is a mechanical mixture,
which is solid, and which is dark grey or nearly black in colour.
If this mixture be placed in water, the nitre is dissolved away and
the charcoal and sulphur are left.

When, however, the integrity of the molecules is disturbed, and a
rearrangement of the atoms takes place, resulting in the formation
of new molecules, then it is said that chemical action has taken
place.

Chemical action, therefore, always results in the formation of
new molecules — new molecules which are endowed with their
own special properties, differing often in the most remarkable and
quite inexplicable manner from those of the original molecules.
One or two examples may be quoted in order to illustrate this
extraordinary modifying effect of chemical action. The two
colourless gases, oxygen and hydrogen, when simply mixed to-
gether, give rise, as already mentioned, to a colourless, gaseous
mixture, in which the dual molecules of hydrogen and the simi-
larly constituted oxygen molecules move about freely amongst

IO Introductory Outlines

each other. By suitable means chemical action may be made
to take place between these two elements, whereby a complete
rearrangement of the atoms takes place, resulting in the formation
of molecules of water— molecules in which, as has been already
mentioned, one atom of oxygen is associated with two atoms of
hydrogen. The product of the chemical action is therefore water,
while both the forms of matter of which it is composed are
gaseous.

The air we breathe, and which is necessary to life, consists of
a simple mixture of two colourless gases, viz., oxygen and nitrogen.
When chemical action takes place between these substances, a
brown-coloured gas is produced in which no animal or vegetable
life could exist for many minutes, on account of its suffocating
nature.

Common salt, which is a white solid substance, and not only
harmless, but even a necessary article of food, contains two atoms
in its molecules — one an atom of chlorine, and the other an atom
of sodium. Chlorine is a yellow gas, intensely suffocating and
poisonous; and sodium is a soft, silver-like metal, which takes
fire in contact with water.

Why it is that a molecule, consisting of an atom of chlorine
and an atom of sodium held together by chemical affinity,
should be endowed with properties so totally different from
those of the contained elements, is altogether unknown ; and
similarly, it is quite impossible to predicate from the properties
of any compound what are the particular elements of which it is
composed. Thus, sugar is a white crystalline solid, soluble in
water, and possessing a sweet taste ; but no one would have
ventured to predict that the molecules of this substance were com-
posed of atoms of carbon, a black, tasteless, insoluble solid ;
hydrogen, a colourless, tasteless gas ; and oxygen^ another colour-
less, tasteless gas.

Chemical Affinity.— When molecules, consisting of two atoms,
say A B, come in contact with molecules consisting of other two
atoms, C D, and a chemical change takes place resulting in the
formation of new molecules, A C and B D, the question naturally
arises, Why does the atom A leave the atom B and attach itself to
C ? In other words, what determines the rearrangement of the
atoms into new molecules ?

At present no exact answer can be given to this question.
Chemists express the fact by saying that the chemical affinity

Chemical Affinity It

existing between A and C is greater than that exerted by B upon
A. This remarkable selective power possessed by the atoms of
different elements lies at the root of all chemical phenomena, and
it differs between the various elements to an extraordinary degree.
For example, the atom of chlorine possesses a very powerful
chemical affinity for the atom of hydrogen : when hydrogen mole-
cules, which consist of two atoms, are mixed with chlorine mole-
cules, which are also aggregations of two atoms, at first a simple
mechanical mixture is obtained, the two different kinds of mole-
cules move amongst each other without undergoing change. On
very small provocation, however, the affinity of the hydrogen atoms
for the chlorine atoms can be caused to exert itself; by merely
momentarily exposing the mixture to sunlight a complete redistri-
bution of the atoms suddenly takes place with explosive violence
and new molecules are formed, each containing one atom of
hydrogen and one atom of chlorine.

Again, an atom of nitrogen is capable of associating itself in
chemical union with three atoms of the element chlorine, forming
a compound whose molecules therefore contain four atoms. The
chemical affinity between the atoms of chlorine and nitrogen is
so feeble, the system is, so to speak, in a state of such unstable
equilibrium, that the very slightest causes are sufficient to instantly
separate the atoms in the most violently explosive manner, and
so break up the compound molecules into separate molecules of
chlorine and nitrogen. In this case the affinity between one
chlorine atom and another chlorine atom is greater than that
between chlorine and nitrogen, consequently the redistribution
that results is of the opposite order to that of the former
example.

As a rule, those elements which the more closely resemble each
other in their chemical habits have the least affinity for each other,
while the greatest affinity usually exists between those which are
most dissimilar,

Chemical Action.— The actual process of redistribution of the
atoms that takes place when molecules of different kinds of matter
are brought together is called chemical action. In many cases
chemical action takes place when the substances are merely
brought together, while in others it is necessary to expose the
bodies to the influence of some external energy : thus chemical
action is brought about in a great number of instances by the
application of heat to the substances. In some cases the influence

12 Introductory Outlines

of light has the effect of causing chemical action to take place ;
for example, when the gases chlorine and hydrogen are mingled
together, no chemical action takes place in the dark, but on
exposing the mixture to light the hydrogen and chlorine combine,
and form the compound hydrochloric acid. It is upon the effect
of light in causing chemical action to take place that the art
of photography depends.

Chemical action may sometimes be induced by the influence of
pressure ; thus, when the two gases, hydrochloric acid and phos-
phoretted hydrogen, are subjected to increased pressure they
combine together to form a -crystalline solid compound known as
phosphonium chloride. In the same way, by very great mechanical
pressure, a mixture of powdered lead and sulphur can be caused
to combine together, when they form the compound, lead sulphide.
There are also a number of chemical actions that are only able
to proceed in the presence of small quantities (often extremely
small) of a third substance, which itself remains unchanged at the
conclusion of the action. These cases are generally included
under the name of catalytic actions : in some of them the modus
operandi of the third substance can be traced (see Oxygen, Modes
of Formation ; also Chlorine, Deacon's Process), while in others
it is not understood. Thus it is found that a number of chemical
actions are quite unable to take place if the materials are abso-
lutely dry ; for example, the element chlorine has a powerful
affinity for the metal sodium, and when these substances are
brought together under ordinary conditions, chemical action in-
stantly takes place, and the compound known as sodium chloride
(common salt) is produced. If, however, every trace of moisture
be perfectly removed from both the sodium and the chlorine, no
action between these elements takes place when they are brought
together, and so long as they remain in this state of perfect dryness
no chemical change takes place. The admission into the mixture
of the minutest trace of the vapour of water, however, at once
induces chemical action between the chlorine and the sodium, but
the exact part that the trace of moisture plays in producing this
effect is not known with certainty. (See also foot-note, page 89.)

A few interesting cases are known in which chemical action is
brought about by the vibration caused by a loud sound or note ;
for example, the molecules of the gas acetylene consist of two
atoms of carbon associated with two of hydrogen. When a quantity
of this gas is exposed to the report produced by the detonation of

Chemical Action 13

mercury fulminate, the mere shock of the explosion causes a re-
distribution of the atoms whereby solid carbon is deposited and
hydrogen set free. We may suppose that the particular vibration
produced by the detonation of the fulminate exercises a disturbing
effect upon the motions of the atoms constituting the molecules of
acetylene, and thereby causes them to swing beyond the sphere of
their mutual attractions, and thus the system undergoes disruption
and rearrangement.

All known instances of chemical action can be referred to one
of three modes, in which the rearrangement of the atoms can take
place.

(i.) By the direct union of two molecules to form a more
complex molecule. Thus, if CO and C1C1 represent two mole-
cules between which chemical action takes place according to
this mode, they unite to form a molecule containing the four
atoms COC1C1.

(2.) By an exchange of atoms taking place between different
molecules. In its simplest form this is illustrated in the action
of one element upon another to form a compound. Thus, if H H
and C1C1 stand for two elementary molecules between which
chemical action takes place, the result is the formation of the two
molecules HC1 HC1. Such a process as this, in which a com-
pound substance is produced directly from the elements which
compose it, is termed synthesis.

The same mode of chemical action may also be exemplified by
the exact opposite to this process, namely, the resolution of a
compound into its constituent elements. Thus, if OHH OHH
represent two molecules of the same compound, when chemical
action takes place it will result in the formation of the three
elementary molecules OO, HH, and HH. Such a process as
this, in which a compound is resolved into its elements, is known
as analysts.*

(3.) By a rearrangement of the atoms contained in a molecule.
There are a number of instances of chemical change, in which the
molecules of the substance do not undergo any alteration in their
composition — that is to say, no atoms leave the molecule, nor are
any added to it. The molecule still consists of the same atoms
after the change as it did before, but the chemical action has

* It will be seen that in each of the examples here given, the process of
rearrangement involves first the decomposition of one or both of the reacting
molecules, and then the combination of the atoms to form different molecules.

14 Introductory Outlines

caused them to assume new relative positions, or different relative
motions with respect to each other. For example, the substances
known to chemists as ammonium cyanate and urea are two totally
different and distinct kinds of matter. These molecules, however,
each contain the same atoms and in the same number ; they each
consist of aggregations of one atom of carbon, one atom of oxygen,
two atoms of nitrogen, and four atoms of hydrogen. When am-
monium cyanate is gently warmed, the eight atoms composing
the molecules undergo this process of rearrangement, and the
substance is changed into urea.

When chemical action takes place between two substances, say A and B,
in ordinary language we say that A acts upon B. Such a statement, however,
must not be understood to imply that A takes the initiative, so to speak, and
that B is in any way less responsible for the action. It is equally true to say
that B acts upon A. For instance, we commonly say nitric acid acts upon
copper, hydrochloric acid acts upon zinc, nitric acid has no action upon gold,
and so on; but it is equally true to say copper acts upon nitric acid, zinc
acts upon hydrochloric acid, gold has no action upon nitric acid. A more
strictly scientific expression would be A and B react, or do not react, as the
case may be. Thus, nitric a'cid and copper react, gold and nitric acid do not
react.

CHAPTER III
CHEMICAL NOMENCLATURE

THE names which have been given to the various elementary forms
of matter are not based upon any scientific system. The names of
some have their origin in mythology. Others have received names
which are indicative of some characteristic property, while those of
several bear reference to some special circumstance connected with
their discovery. It has been the custom in modern times to dis-
tinguish metals from non-metals by applying to the former names
ending in the letters urn, and consequently such metals as are of
more recent discovery all have names with this termination. The
common metals, however, which have been known since earlier
times, such as gold, silver, tin, and copper, keep their old names.
The two elements selenium and tellurium were at the time of their
discovery thought to be metals, and they consequently received
names with the terminal um ; these substances strongly resemble
metals in many of their physical properties, but in their chemical
relations they are so closely similar to the non-metal sulphur, that
they are by general consent classed among the non-metals ; they are
examples of those elements which are distinguished as metalloids.
On this account selenium is by some chemists termed selenion.

In naming chemical compounds, the chemist endeavours that
the names employed shall not only serve to identify the sub-
stances, but shall as far as possible indicate their composition.
The simplest chemical compounds are those composed of only
two different elements ; such are spoken of as binary compounds*
and their names are made up of the names of the two elements
composing them, thus —

* This expression is now sometimes used in a somewhat modified sense.
Thus in the language of the ionic theory (p. 107) the term Unary compound is
used to denote a substance which dissociates into two ions, quite irrespective
of the number of elements it may contain. It is to be regretted that under
these circumstances a new word was not coined to denote the newer idea.

1 6 Introductory Outlines

The compound formed by the chemical union of —

Hydrogen with sulphur is called hydrogen sulphide.
Sodium ,, chlorine ,, sodium chloride.
Copper ,, oxygen ,, copper oxide.
Calcium ,, fluorine ,, calcium fluoride.
Potassium ,, iodine ,, potassium iodide.

It continually happens, however, that the same two elements
combine together in more than one proportion, giving rise to as
many different compounds, in which case it becomes necessary to
so modify the names that each of the compounds may be dis-
tinguished. This is accomplished by the use of certain terminal
letters or of certain prefixes ; for example, the element phos-
phorus combines with chlorine in two proportions, forming two
different compounds — in one the molecules contain one atom of
phosphorus united to three atoms of chlorine, in the other the
molecules consist of one atom of phosphorus associated with five
of chlorine. These two compounds may be distinguished in the
following ways : —

i atom of phosphorus with 3 atoms of chlorine forms phosphorow^ chloride,
i ,, ,, ,, 5 ,, ,, M phosphom: chloride.

or —

latom of phosphorus with3 atomsof chlorine forms phosphorus ^-/chloride,
i ,, ,, ,, 5 ,, ,, ,, phosphorus /*>»ta chloride.

The latter method of distinction is the more general, thus —

i atom of sulphur with 2 atoms of oxygen forms sulphur dioxide,
i ,, ,, „ 3 ,, ,, ,, sulphur trioxide.
i atom of carbon with i atom of oxygen forms carbon monoxide,
i ,, ,, ,, 2 atoms ,, ,, carbon dioxide.

Occasionally the prefixes sub vca&proto are employed to denote
these differences of composition, but their use is more limited, and
is becoming out of vogue. When more than two compounds are
formed by the union of the same two elements, the additional
prefixes hypo, under, and/^r, over, are sometimes used.

In a considerable number of instances the systematic names of
familiar compounds give way to the vulgar or common names by
which they are known, thus —

/•Ammonia . . . Hydrogen nitride ^
Common I Hydrochloric acid . Hydrogen chloride I Systematic

Sulphuretted hydrogen . Hydrogen sulphide I names.
Water , Hydrogen monoxide'

Chemical Nomenclature \j

Binary compounds consisting of elements united with oxygen
are called the ffxides of those elements. Certain of these oxides
are capable of reacting with water, giving rise to substances known
as acids; such oxides are distinguished as acid-forming oxides, or
acidic oxides. They are also sometimes termed anhydrides. All
the non-metallic elements, except hydrogen and the members of
the argon group, form oxides of this order, and the acids derived
from them are known as the oxy-acids or hydroxy-acids.

Certain other oxides also unite with water, but give rise to com-
pounds known as hydroxides. When such oxides, which are all
derived from the metallic elements, are brought into contact with
acids, chemical action takes place, and a compound termed a salt'^
formed, together with water. Such" oxides are distinguished as salt-
forming or basic oxides. There are also oxides which are neither
acidic nor basic. The names of oxy-acids are derived from the
name of the particular oxide from which they are formed, thus —

Carbon dioxide gives carbonic acid.
Silicon dioxide „ silicic acid.

When the same element forms two acid-forming oxides, the
terminals ic and ous are applied to the acids to denote respectively
the one with the greater and the less proportion of oxygen, thus —

Sulphur /rzoxide gives sulphunV acid.
Sulphur ^fzbxide gives sulphurous acid.
Nitrogen ^<?;//oxide gives nitric acid.
Nitrogen /r/oxide gives nitrous acid.

When more than two such acids are known, the additional
prefixes hypo or per are made use of. Thus /^rsulphuric acid
denotes an acid containing the highest quantity of oxygen, while
4y/0nitrous acid stands for an acid containing less oxygen than is
present in nitrous acid.

There is a class of binary compounds formed by the combination
of a large number of the elements with sulphur ; these are known as
sulphides. Certain of these sulphides are also capable of forming
acids which are analogous in their constitution to oxy-acids, but in
which the oxygen atoms are substituted by atoms of sulphur.
These acids are known as thio acids (sometimes sulpho acids),
and the same system of nomenclature is adopted to distinguish
these : thus we have thio-arsenitfWJ acid, thio-arsenzV acid, denoting

iS Introductory Outlines

respectively the acid with the smaller and the larger proportion of
sulphur.

It was at one time believed that all acids contained oxygen, that
indeed this element was essential to an acid. The name oxygen
indicates this belief, the word signifying "the acid-producer."
This view is now seen to have been incorrect, for many acids are
known in which oxygen is not one of the constituents. Thus the
elements fluorine, chlorine, bromine, and iodine, which constitute
the so-called Halogen group of elements, each combines with
hydrogen, giving rise respectively to hydrofluoric, hydrochloric,
hydrobromic, and hydriodic acids.

All known acids contain hydrogen as one of their constituents.

As already stated, when chemical action takes place between an
acid and a base* a salt is formed. Oxy-acids in this way give rise
to oxy-salts, thio-acids to thio-salts, and halogen acids to haloid
salts.

The latter salts being binary compounds, their names are given
according to the system already explained, such, for example, as
calcium fluoride, sodi'im chloride, potassium bromide, silver iodide.

In the case of the oxy-salts and thio-salts, the names are made
up from the names of the acid and of the metal contained in the
base, with the addition of certain distinctive suffixes : thus if the
acid be one whose name carries the terminal ous, its salts will be

* The word base is unfortunately employed by different chemists in different
senses, so that it is scarcely possible to give a precise definition of it. Originally,
no doubt, the term was employed simply to denote the idea of foundation, and
was applied to the metal or the oxide of the metal entering into the composition
of a salt ; which being the more tangible constituent was thus regarded as the
more important one, or the basis of the salt. At the present day the word
base is used in INORGANIC chemistry chiefly to denote that class of compounds
described on page 17 as hydroxides, while the qxides from which these
hydroxides are derived are spoken of as basic oxides. Besides this class, it
includes ammonia and a few other compounds which like ammonia are not
derived from metallic oxides. The ORGANIC chemist, on the other hand, regards
ammonia as the true type of a base ; and all organic compounds which can be
regarded as " derivatives " of ammonia are called bases. Not only so, but the
term is even extended so as to include similar " derivatives " of the phosphorus,
arsenic and antimony analogues of ammonia, thus giving rise to the expressions
nitrogen bases, phosphorus bases, &c.

Again, in the language of the modern theory of ionic dissociation, a base is
defined as a compound in which the only negative ions are the hydroxide ions
(page 107). This definition includes the class of hydroxides above mentioned,
but does not include ammonia gas.

Chemical Nomenclatu

distinguished by the suffix tie, while the names of the salts derived
from acids whose names end in ic are terminated* by the letters
ate.

Nitrous acid and potassium oxide give potassium nitx*/*.
Sulphurous acid „ „ „ sulphite.

NitnVacid „ „ „ nitrate.

SulphunVacid „ „ „ sulphate.

The formation of a salt by the. action of an acid upon a base is
due to the redistribution of the atoms composing the molecules of
the two compounds, in such a manner that some or all of the
hydrogen atoms in the acid molecules exchange places with certain
metallic atoms from the molecules of the base. Acids which con-
tain only one atom of hydrogen so capable of becoming exchanged
for a metal are termed mono-basic acids ; those with two, three, or
four such hydrogen atoms are distinguished respectively as di-basic,
tri-basic, and tetra-basic acids.

If the whole of the displaceable hydrogen in an acid becomes
replaced by the base, the salt formed is known as a normal salt.
On the other hand, when only a portion of the hydrogen atoms
is displaced by the base, the salt is distinguished as an adt(__
salt. Thus sulphuric acid contains two atoms of hydrogen in its
molecule (associated with one of sulphur and four of oxygen) ; if
both the hydrogen atoms are exchanged for potassium, the salt
obtained is normal potassium sulphate, and when only one is so
replaced the salt is known as acid potassium sulphate. By the
term acid salt, therefore, must be understood not a substance
having the familiar properties of an acid, such as a sour taste and
the power to redden litmus, but a salt in which one or more of the
hydrogen atoms of the original acid are still left in the. molecule.*
It is quite true that some of the salts of this class do possess acid
qualities and will redden litmus, but this is due to what may be
regarded as merely the accidental circumstance of the acidic
portion of the molecule being derived from a strong acid. Many
substances belonging to the class of acid salts are perfectly neutral
in their behaviour towards litmus, while, on the other hand, some
are strongly alkaline'. For example, acid potassium sulphate is acid

* Some chemists prefer to regard the acids themselves as the hydrogen salts ;
accordingly they apply to nitric acid, sulphuric acid, nitrous acid, sulphurous
acid, &c., the names hydrogen nitrate, hydrogen sulphate, hydrogen nitrite,
hydrogen sulphite, &c., respectively.

2O introductory Outlines

to test paper, acid calcium carbonate is neutral, while acid sodium
carbonate is alkaline.

A third class of salts is formed by the association of one or
more molecules of normal salt, with one or more additional mole-
cules of the base : these are known as basic salts. Thus, carbonic
acid and the base lead hydroxide form such a salt known as basic
lead carbonate.

CHAPTER IV
CHEMICAL SYMBOLS

CHEMISTS are agreed in adopting certain symbols to denote the
atoms of the various elementary forms of matter. The table on
page 22 contains the names of the elements at present recognised,
and in the second column are given the symbols which are em-
ployed to represent their atoms. The names of the rare elements
are printed in italics.

In a number of instances the atomic symbol is the initial letter
of the ordinary name of the element : thus Boron, B ; Carbon, C ;
Fluorine, F ; Hydrogen, H ; Oxygen, O ; Sulphur, S.

When more than one element has the same initial, either the
first two letters of the name, or the first and another that is pro-
minently heard in pronouncing the word are employed, as Bromine,
Br ; Cobalt, Co ; Chlorine, Cl ; Platinum, Pt. In some cases
letters taken from the Latin names for the elements are used, such
as Antimony (Stifo'um}, Sb ; Gold (Aurum}, Au ; Silver (Argentum\
Ag ; Lead (Plumbum\ Pb ; and Iron (Ferrum\ Fe.

These symbols are not intended to be employed as mere short-
hand signs, to be substituted as abbreviations for the full names
of the elements, but in every case they denote one atom of the
element. The symbol H stands for one atom of hydrogen, the
symbol O stands for one atom of oxygen ; Cl means one atom of
chlorine, and Ag represents one atom of silver. No other use of
these symbols is legitimate.

It has been already mentioned (page 8) that the molecules of
the different elements are composed of different numbers of atoms ;
for example, the molecule of hydrogen consists of two atoms, and
ordinary oxygen also forms diatomic molecules. These facts are
expressed in chemical notation by the use of small numerals placed
immediately after the symbol of the atom, thus H2 denotes a mole-
cule of hydrogen, O2 a molecule of oxygen. The molecule of ozone
consists of an aggregation of three atoms of oxygen, and is

Introductory Outlines

Atomic Weights.
2 3

6 $

jil

.§ 8

lla

Name.

rt£

Name.

C/3

oj;^

£ o •

ill

M< Ii

M< H

Aluminium

27.1

Molybdemtm .

Antimony(^^^'«'w)

120

120.2

Neodymium

144

J44-3

Ar2"on

39.88

20. 2

Arsenic .

74.96

Nickel . . .

58.68

Barium .

137

137.4

Nitrogen

14.01

Beryllium (G ucinum

Osmium . . i.

191

190.9

Bismuth

208

208.0

Oxygen . . .

16.00

Boron .

Palladium

106

106.7

Bromine

79.92

Phosphorus . ,

31.04

Cadmium

112

II2.4

Platinum . . .

195

Caesium .
Calcium .

Cs
Ca

133
40

132.8
40.07

Potassium (Kal- \
ium) . . . }

39.10

Carbon .

•1?

12.00

Praseodymium

140-5

140.6

Cerium .

140

140.25

Radium ....

226-4

226.4

Chlorine

35-5

35-45

Rhodium . . .

103

102.9

Chromium

Rubidium .

85.45

Cobalt .

58.97

Ruthenium . .

1017

IOI-7

Columbium )
(Niobium) . )

93-5

93- S

Samarium .
Scandium .

Sm
Sc

150
44

150.4
44-1

Copper (Cuprum)
Erbium ....

63-5
166

63.6
167.4

Selenium . . .
Silicon ....

Se
Si

79
28

79-2
28.3

Fluorine . . .

Silver (Argentum]

108

107.88

Gallium . .

Sodium (Natrium\

23.00

Germanium

72.5

Strontium . . .

87-6

87.6

Gold (Aurum)

197

197.2

Sulphur ....

32.07

Helium ....

3-99

Tantalum .

181

181.5

Hydrogen . . .

1.008

Tellurium .

127

127.5

Indium . .

115

114.8

Thallium

204

204.0

Iodine ....

127

126.92

Thorium . . .

232

232.4

Iridium ....

193

Tin (Stannum) .

118

119

Iron (Ferrum)

55-85

Titanium . .

48.1

Krypton ....

82.92

Timgsten

184

TMnthanum . ,

139

X39

Uranium .

238-5

238.5

Lead (Plumbum)

207

207. 10

Vanadium .

Lithium ....

6.94

Xenon ....

130

130.2

Magnesium '.' .

24.32

Ytterbium . . .

172

Manganese . . .

54-93

Yttrium . . .

Mercury (Hydr- )

argyrum) . . f

200

200. 6

Zirconium . . .

907

90.6

represented by the symbol O3, while the tetr-atomic character of
the phosphorus molecule is expressed in the symbol P4. The
composition of compound molecules is expressed by placing the

Chemical Symbols 23

symbols of the atoms which compose such molecules in juxta-
position : thus a molecule consisting of one atom of sodium (symbol
Na) and one atom of chlorine (symbol Cl) is represented by the
united symbols of these two elements, NaCl ; a compound con-
sisting of one atom of carbon and one atom of oxygen by the
symbols of these two atoms, CO. Such arrangements of symbols
representing molecules are termed molecular formula, or, simply,
formula:.

When the molecule contains more than one atom of any parti-
cular element, this fact is indicated by the use of numerals placed
immediately after the symbol to be multiplied : thus, a molecule of
water consists of two atoms of hydrogen and one atom of oxygen,
\hzformula for water is therefore H2O. One molecule of ammonia,
consisting of an atom of nitrogen with three atoms of hydrogen,
is represented by the formula NH3 ; and a molecule of sulphuric
acid, which is an aggregation of two atoms of hydrogen, one
atom of sulphur, and four atoms of oxygen, has the formula
H2SO4.

It is sometimes necessary to represent the presence in a mole-
cule of certain groups of atoms, groups which seem to hold together
and often to function as a single atom. This is accomplished by
the use of brackets : thus (NH4)2SO4 is the formula for a molecule
containing one atom of sulphur, four atoms of oxygen, eight atoms
of hydrogen, and two atoms of nitrogen ; the nitrogen and hydrogen
atoms being present as two groups, in each of which one nitrogen
atom is associated with four hydrogen atoms. Such groups of
atoms are termed compound radicals.

When it is required to indicate more than one molecule of the
same substance, numerals are placed immediately in front of the
formula : thus 2H2O signifies two molecules of water, and 3NH3
expresses three molecules of ammonia.

By means of these symbols and formulae, chemists are enabled
to represent, in a concise manner, the various chemical changes
which it is the province of chemistry to examine. Such changes
are usually termed chemical reactions, and they are represented
in the form of equations in which the symbols and formulae of
the reacting substances as they are before the change are placed
on the left, and those of the substances which result from the
change upon the right, thus —

H2 + Cl2=2HCl

24 . " • Introductory Outlines

The sign + has a different significance as used on the left side
of the equation to that which it bears upon the right. On the
left hand il implies that chemical action takes place between the
substances, while on the opposite side it has the simple algebraic
meaning. Thus, the second of the above equations is understood
to mean, that when the compounds, mercuric chloride and potassium
iodide, are brought together in such a way that chemical action
results, a redistribution of the atoms will take place, resulting in
the formation of mercury iodide and also potassium chloride.

As further illustrations of the use of chemical symbols, the
following three examples may be given as exemplifying the three
modes of chemical action mentioned on page 13 : —

(1) NH3+ HC1 = NH4C1.

Ammonia combines with hydrochloric acid, and gives ammonium
chloride.

(2) H2S04 + Na2CO3 = Na2SO4 + CO2 + H2O.

Sulphuric acid reacts with normal -sodium carbonate, and yields
normal sodium sulphate, carbon dioxide, and water.

(3)(CN)0(NH4) = (NH2)2CO.
Ammonium cyanate is converted into urea.

In all cases where the nature of the chemical change is under-
stood, it is capable of expression by such equations, and as matter
is indestructible, every atom present in the interacting molecules
upon the left of the expression reappears on the right-hand side
in some fresh association of atoms.*

* See also Chemical Notation, chapter vii.

CHAPTER V

THE ATOMIC THEORY

THE atomic view as to the constitution of matter, briefly sketched
out in Chapter I., forms a part of what is to-day known as the
atomic theory.

When chemical changes were carefully studied from a quantita-
tive standpoint, four laws were discovered in obedience to which
chemical action takes place. These laws are distinguished as
the laws of chemical combination. Three of these generalisations
refer to quantitative relations as respects weight; while one ex-
presses quantitative relations with regard to volume, and only
relates to matter in the gaseous state.

I. Law of Constant Proportion.— The same compound always
contains the same elements combined together in ;the same proportion
by weight; or expressed in other words, The weights of the con-
stituent elements of every compound bear an unalterable ratio to
each other, and to the weight of the compound formed.

II. Law of Multiple Proportions. — When the same two
elements combine together to form more than one compound, the
different weights of one of the elements which unite with a constant
weight of the other bear a simple ratio to one another; or this law
may be stated thus : When one element unites with another in
two or more different proportions by weight, these proportions are
simple multiples of a common factor.

III. Law of Reciprocal Proportions, or Law of Equivalent
Proportions. — The weights of different elements which combine
separately with one and the same weight of another element, are
either the same as, or are simple multiples of, the weights of these
different elements which combine with each other; or in other
words, The relative proportions by weight in which the elements,
A, B, C, D, &C., combine with a constant weight of another
element, X, are the same for their combinations with any other
element^ F.

26 Introductory Outlines

IV. Law of Gaseous Volumes, or The Law of Gay-Lussae.

— When chemical action takes place between gases, either elements
or compounds, the volume of the gaseous product bears a simple
relation to the volumes of the reacting gases.

These four laws are the foundations upon which the whole
superstructure of modern chemistry rests.

(i.) The Law of Constant Proportions.— When two sub'
stances are mingled together, and remain as a mere mechanical
mixture, they may obviously be present in any proportion, and it
was at one time thought that when two substances entered into
chemical combination with each other, they could do so also in
any proportion, and that the composition of the resulting com-
pound would vary from this cause. This belief was finally
disproved, and the law of constant proportions definitely estab-
lished by Proust in the year 1806. The same compound, therefore,
however made, and from whatever source obtained, is always
found to contain the same elements united together in the same
proportion by weight. Thus, common salt, or, to adopt its
systematic name, sodium chloride, which is a compound of the
two elements sodium and chlorine, may be made by bringing the
metal sodium into contact with chlorine gas, when the two
elements unite and form this compound. It can also be made
by the action of hydrochloric acid upon the metal sodium, or by
adding hydrochloric acid to sodium carbonate, and by a variety
of other chemical reactions. When the sodium chloride obtained
by any or all of these processes is analysed, it is invariably found
to contain the elements chlorine and sodium in the proportion by
weight of 1 10x6479, or, expressed centesimally —

Sodium . . 39.32
Chlorine 60.68

IOO.OO

and when this is compared with the sodium chloride as found in
nature, obtained either from the salt-mines of Cheshire, or the
celebrated mines in Galicia, or by evaporating sea-water, it is
found that the composition of the compound in all cases is exactly
the same. In the same way the compound water, consisting of
the two elements hydrogen and oxygen, whether it be prepared
synthetically by causing the two elements to unite directly, or
from any natural source, as rain, or spring, or river, is

The Atomic Theory 27

found to contain its constituent elements hydrogen and oxygen in
the ratio by weight. of i : 8, or,

Hydrogen . . 11.12 —
Oxygen . . 88.88

" ." 100.00

If in the formation of sodium chloride by the direct combination
of its constituent elements, an excess of either one or other be
present beyond the proportions 39.32 per cent, of sodium and 60.68
per cent, of chlorine, that excess will simply remain unacted upon.
If eight parts by weight of hydrogen and eight parts by weight
of oxygen be brought together under conditions that will cause
chemical action, the eight parts of oxygen will unite with one part
of hydrogen, and the other seven parts of hydrogen merely remain
unchanged. This fact, that elements are only capable of uniting
with each other in certain definite proportions, marks one of the
most characteristic differences between chemical affinity and those
other forces, such as gravitation, that are usually distinguished as
physical forces ; for although there are many instances known in
which the extent to which a chemical action may proceed (that is,
the particular proportion of the reacting bodies which will undergo
the permutation that results in the formation of different mole-
cules) is influenced by the mass of the acting substances, it never
governs the proportion in which the elements combine in these
compounds.

It follows from the law of constant composition that the sum of
the weights of the products of a chemical action will be equal to
that of the interacting bodies ; and upon the validity of this law
depend all processes of quantitative analyses.

(2.) The Law of Multiple Proportions was first recognised
by Dalton, who investigated certain cases where the same two
elements combine together in different proportions, giving rise to
as many totally distinct compounds. These proportions, however,
were always found to be constant for each compound so produced,
so that this law formed no contradiction to the law of constant
composition. The simple numerical relation existing between the
numbers representing the composition of such compounds will be
evident from the following examples. The two* compounds of

* In Dalton's day these two substances were the only known compounds of
cajbon with hydrogen.

28 v Introductory Outlines

carbon with hydrogen, known as marsh gas and ethylene, are
found to contain these elements in the proportions —

Marsh gas . . i part by weight of hydrogen with 3 parts of carbon.
Ethylene . . i ,, ., ,, 6 ., ,,

The two compounds of carbon with oxygen contain these ele-
ments in the proportion —

Carbon monoxide . i part of carbon with 1.334 parts of oxygen by weight.
Carbon dioxide . i ,, ,, 2.667 »» >» »»

The elements nitrogen and oxygen form as many as five different
compounds, in which the two elements are present in the propor-
tions— •

Nitrous oxide . . part of nitrogen with 0.571 parts of oxygen by weight.
Nitric oxide. . . ,, ,, i-I43 , . . . »

Nitrogen trioxide . ,, ,, 1.714 «> » »

Nitrogen peroxide ,, ,, 2.286 ,, ,, ,,

Nitrogen pentoxide ,, ,, 2-857 ,, ,, ,,

The relative propr nons of carbon combining with a constant
weight of hydrogen in the two first compounds are as i : 2.

Those of oxygen uniting with a constant weight of carbon in the
second example are also as i : 2, while in the nitrogen series the
relative proportions of oxygen in combination with a constant
weight of nitrogen are as i 12:3:4:5.

(3.) Law of Reciprocal Proportions.— Known also as the law
of proportionality, or the law of equivalent proportions. When
the weights of various elements, which were capable of uniting
separately with a given mass of another element, were compared
together, it was seen that these weights bore a simple relation to
the proportions in which these elements combined amongst them-
selves. For example, the elements chlorine and hydrogen each
separately combine with the same weight of phosphorus, the pro-
portions being —

Phosphorus : chlorine = i : 3.43
Phosphorus : hydrogen = i : 0.097

The elements chlorine and hydrogen can combine together, and
they do so in the proportion —

Chlorine : hydrogen = 35.5 : i
but 35 : i = 3.43 : 0.097

Therefore the proportions by weight in which chlorine and

'fhe Atomic Theory £§

hydrogen separately combine with phosphorus is a measure of the

proportion in which they will unite together.

Again, the two elements carbon and sulphur each separately

combine with the same weight of oxygen, the proportion being-
Oxygen : carbon = i : 0.375
Oxygen : sulphur = i : i

But the elements carbon and sulphur themselves unite together,
and in the proportion —

Carbon : sulphur = 0.1875 ' l
but 0.1875 : l — °-375 : 2

Therefore the proportion by weight in which carbon and sulphur
separately unite with the same mass of oxygen is a simple multiple
of that in which these two elements combine together. These
remarkable numerical relations will be rendered still more evident
by comparing the proportions in which the members of a series of
elements combine with a constant weight of various other elements :
thus—

Hydrogen. Sodium. Potassium. Silver. Mercury. Chlorine.

0.02817 0.6479 1.02 3.04 2.816 unite separately with i part.

It will be seen that the proportion in which these numbers stand
to each other is as —

i : 23 : 39 : 107 : 100 : 35.5

Let us now compare these proportions with those in which the
same elements unite with a constant weight of the element
bromine —

Hydrogen. Sodium. Potassium. Silver. Mercury. Bromine.

0.0125 0.2875 0.4875 1.34 1.25 unite with i part,

or as —

i : 23 : 39 : 107 : 100 : 80

Each of these five elements in like manner combines with
oxygen, and the weights which are found to unite with a constant
mass of oxygen are —

Hydrogen. Sodium. Potassium. Silver. Mercury. Oxygen.

0.125 2.875 4-875 J3-38 I2-S unite with i part,

again as —

i : 23 : 39 107 : 100 : 8

30 Introductory Outlines

The same relation will appear in the case of the combination of
these five elements with a constant weight of sulphur —

Hydrogen. Sodium. Potassium. Silver. Mercury. Sulphur.

0.0625 1.4375 2.4375 6.69 6.25 unite with i part.

or as —

i : 23 : 39 : 107 : 100 : 16

It is thus evident that the proportions in which the members of
such a series combine with a constant weight of one element is the
same as that in which they unite with a constant mass of another
element. One part by weight of hydrogen combines with 35.5
parts of chlorine, 80 parts of bromine, 8 parts of oxygen, and 16
parts of sulphur — that is to say, these proportions of these four
elements satisfy the chemical affinity of i part of hydrogen ; they
are therefore said to be equivalent. Twenty-three parts of sodium
is likewise equivalent to 35.5 parts of chlorine, 80 parts of bromine,
8 parts of oxygen, and 16 parts of sulphur, and by the same
reasoning it is also equivalent to i part of hydrogen, 39 parts of
potassium, 107 parts of silver, and 100 parts of mercury. These
numbers, therefore, are known as the equivalent weights or the
equivalents of the elements, or their combining proportions, and the
combining weight of an element may therefore be defined as the
smallest weight of that element which will combine with i part by
weight of hvdrogen.

This law of proportionality, or reciprocal proportion, was dis-
covered by Richter, but it was left for Dalton to trace the connec-
tion between these three generalisations. Dalton adopted and
adapted an ancient theory concerning the ultimate constitution of
matter which was expounded by certain of the early Greek philo-
sophers. The exponents of this theory held that matter is built up
of vast numbers of minute indivisible particles, in opposition to the
antagonistic theory believed by others, namely, that matter was
absolutely homogeneous and capable of infinite subdivision.

Dalton embraced the ancient doctrine of atoms, and extended it
into the scientific theory which is to-day known as Dalton's atomic
theory, and is accepted as a fundamental creed by modern chemists.

According to this theory, matter consists of aggregations of minute
particles, or atoms, which are chemically indivisible. Dalton
conceived that chemical combination takes place between atoms—
that is to say, when chemical action takes place between two
elements, it is due to the union of their atoms ; the atoms, coming
into juxtaposition with each other under the influence of chemical

The Atomic Theory $t

affinity, are held together by the operation of this force. He further
assumed that the atoms of the various elements possessed different
relative weights, and that the relations existing between these
weights was the same as that between the weights in which experi-
ment had shown the elements to be capable of combining together.
In other words, he said that the numbers representing the combin-
ing proportion of the elements expressed also the relative weights
of the atoms.

Let us now see how this theory satisfies and explains the first
three laws of chemical combination.

(i.) The Law of Constant Composition.— It has already been
shown (p. 26) that the compound sodium chloride, wheresoever and
howsoever obtained, contains the elements chlorine and sodium
in the proportion — •

Chlqrine : sodium = I : 0.6479.

These numbers have been shown on p. 29 to represent the com-
bining proportions —

Chlorine : sodium = 35.5 : 23.

Now the atomic theory states, that sodium chloride is formed by
the union of atoms of chlorine with atoms of sodium, and that the
relative weights of these atoms is expressed by the combining
weights of the elements, namely, 35.5 and 23. If therefore, sodium
is to combine with chlorine, since atoms are indivisible masses, it
follows that the compound produced by the union of one atom of
each of these two elements must always have the same composi-
tion.

(2.) The Law of Multiple Proportions.— The ratio in which
oxygen combines with hydrogen to form the compound water is
seen on p. 27 to be as 8 : i. This number 8, therefore, we will
for the present argument regard as the relative weight of the atom
of oxygen.*

Oxygen combines with carbon as already mentioned, forming
two different compounds ; in the first, the elements are present in
the proportion —

Carbon : oxygen = i : 1.334 = 6:8,

* For reasons which will be explained later, chemists now regard the number
16 as representing (in ipund numbers) the relative weight of the atom oi
oxygen.

$2 Tntroductory Outlines

that is to say, in the proportion of one atom of carbon to one atom
of oxygen. According to the theory, if the atom of carbon unites
with more oxygen than one atom, it must at least be with two
atoms. It may be with three or with four, but as the compound
must be formed by the accretion of these indivisible atoms, the
increment of oxygen must take place by multiples of 8. When the
second compound is examined it is found to contain its constituent
elements in the proportion —

Carbon : oxygen = i : 2.667 = 6 : 16,

that is to say, in the proportion of one atom of carbon to two
atoms of oxygen. This information respecting the composition of
these two compounds is conveyed both in their names and their
formulae. The first is termed carbon monoxide, and its formula is
expressed by the symbol CO ; while the second is distinguished as
carbon dzoxide, and has the formula CO2.

The difference in the composition of the five compounds that
nitrogen forms by union with oxygen will be. made evident by the
aid of this theory. The proportion of nitrogen to oxygen in these
compounds is —

(i.) Nitrogen : oxygen = i : 0.571 = 14 : 8

(2.) Nitrogen : oxygen = i : 1.143 = !4 '• 16

(3.) Nitrogen : oxygen = i : 1.714 = 14 : 24

(4.) Nitrogen : oxygen = i : 2.268 = 14 : 32

(5.) Nitrogen : oxygen = i : 2.857 = 14 : 40

And it will be seen that the increase in the proportion of oxygen in
the compounds takes place by the regular addition of a weight of
that element equal to 8, which at the present stage of the argument
we are regarding as representing the relative weight of the atom
of oxygen.

(3.) The Law of Reciprocal Proportions.— If the illustrations
given on p. 28 of the operation of this law be examined in the light
of the atomic theory, their explanation will be evident : thus, the
relative proportions in which hydrogen and chlorine separately
combine with phosphorus is 0.097 : 3.43, and the ratio between these
numbers is as i : 35.5, which is the proportion in which these two
elements are known to unite together to form hydrochloric acid.
These numbers, however, represent the relative weights of the
atoms of these elements, therefore hydrochloric acid may be sup-
posed to be formed by the union of one atom of hydrogen with
one atom of chlorine.

The * A tomic Theory 3 3

Again, the relative weights of carbon and sulphur which sepa-
rately combine with a constant weight of oxygen are — carbon, 0.375 ;
sulphur, i ; and the ratio between these numbers is as 6 : 16.

Carbon and sulphur, however, unite together in the relative
proportion —

Carbon : sulphur = 0.1875 : l = ° : 32<

Therefore the compound they produce may be supposed to consist
of one atom of carbon, having the relative weight 6, and two atoms
of sulphur, each with the relative weight 16.

CHAPTER VI
ATOMIC WEIGHTS

IN the third column of the table on page 22, the numbers are
given which are at the present time generally accepted by chemists
as representing the approximate atomic weights of the elements.
These numbers depart, in many instances, from those arrived at
by Dalton's methods : thus, the relative weights of carbon, oxygen,
nitrogen, and sulphur, which were found to be equivalent to one
part of hydrogen, are— carbon = 6,* oxygen = 8, nitrogen = 4.66,
sulphur = 1 6 ; while the figures given as the approximate atomic
weights of these elements in the table are — carbon = 12, oxygen
— 16, nitrogen = 14, sulphur = 32. We must now discuss some
of the chief reasons for these departures. In the two compounds
of carbon and hydrogen known to Dalton, namely, marsh gas and
ethylene, the proportions of carbon to hydrogen are —

In ethylene . . . Carbon : hydrogen = 6 : i.
In marsh gas . . Carbon : hydrogen = 6:2.

Dalton therefore concluded that ethylene was a compound con-
taining i atom of carbon united with I atom of hydrogen, and to
which, therefore, he gave the formula CH ; and that marsh gas
consisted of i atom of carbon combined with 2 atoms of hydrogen,
and which he accordingly represented by the formula CH2.

There was, however, nothing to prove that the weight of carbon
was constant in the two compounds, for it will be obvious that the
same ratio between the weight of carbon and hydrogen will still
be maintained by assuming that the hydrogen is constant, and
that the carbon varies, thus —

In marsh gas . . Hydrogen : carbon : : i : 3.
In ethylene . . Hydrogen : carbon : :i 13x2.

* These are the numbers which Dalton ought to have obtained had his
methods of determination been more exact. The figures he actually found for
the combining weights of these four elements were respectively. £, 7, 5, 13.

Atomic Weights 35

That is to say, the ratios are not disturbed by the assumption
that in marsh gas we have I atom of hydrogen combined with i
atom of carbon, having the relative combining weight of 3, and in
ethylene I atom of hydrogen united with 2 atoms of carbon.

It will be evident, however, that if we could gain any exact
information as to the actual number of atoms which are present
in these various molecules, this difficulty would no longer exist.

For example, suppose it were possible to ascertain that in the
molecule of marsh gas there were 4 atoms of hydrogen, then as
the relative weights of hydrogen and carbon in this compound are
as £ : 3, the weight of the carbon atom would obviously have to
be raised from 3 to 12 ; and if it could be determined that in the
ethylene molecule there were also 4 atoms of hydrogen, then
seeing that the ratio of hydrogen to carbon in this substance is
as i : 6, we should conclude that it contained 2 atoms of carbon,
of the relative weight not less than 1 2, and the composition of the
two compounds would be expressed by the formulae, marsh gas
CH4, ethylene C2H4.

Again, the relative weights of hydrogen and oxygen in water
are as i : 8. If the molecule of water contains only i atom of
hydrogen, then we conclude that 8 represents the relative weight
of the oxygen atom, and the formula for water will be HO. But
suppose it to be discovered that there are two atoms of hydrogen in
a molecule of this compound, then it becomes necessary, in order
to retain the ratio between the weight of these constituents (a
ratio ascertained by. analysis), to double the number assigned to
the oxygen atom and to regard its weight as 16, as compared with
i atom of hydrogen, and the formula for water in this case would
be H2O.

The compound ammonia contains the elements hydrogen and
nitrogen in the ratio —

Hydrogen : nitrogen : : i : 4.66.

If the molecule of ammonia contains only i atom of hydrogen,
then 4.66 represents the relative weight of the nitrogen atom, and
the formula will be NH ; but if it should be found that there are
3 atoms of hydrogen in this molecule, then again the relative
weight assigned to the nitrogen must be trebled in order to pre-
serve the ratio, and it will have to be raised from 4.66 to 14 (in
round numbers), and the formula for ammonia will be NH3.

From these considerations it will be evident, that it is of the

36 Introductory Outlines

highest importance to gain accurate knowledge as to the actual
number of atoms which are contained in the molecules of matter —
in other words, to learn the true atomic composition and structure
of molecules ; and it may be said that this problem has occupied
the minds of chemists from the time that Dalton published his
atomic weights, in the year 1808, down to the present time. There
is no single method of general application, by means of which
chemists are able to determine the atomic weight of an element ;
but they are guided by a number of independent considerations,
some of which are chemical in their character, while others are of
a physical nature ; and that particular number which is in accord
with the most of these considerations, or with what are judged to
be the most important of them, is accepted as the true atomic
weight.

The chief methods employed for determining atomic weights
may be arranged under the following four heads :—

1. Purely chemical methods.

2. Methods based upon volumetric relations.

3. Methods based upon the specific heats of the elements.

4. Method based upon the isomorphism of compounds.

I. As an illustration of the chemical processes from which
atomic weights may be deduced, the following examples may be
given, namely, the case of the two elements oxygen and carbon.

Oxygen combines, as already stated, with hydrogen in the
proportion —

Hydrogen : oxygen = i : 8.

When water is acted upon by the element sodium, the compound
is decomposed and hydrogen is evolved ; and it is found that if
1 8 grammes of water are so acted on, I gramme of hydrogen is
evolved,- and 40 grammes of a compound are formed, which
contains sodium, together with all the oxygen originally in the
\S grammes of water, and some hydrogen. This compound, under
suitable conditions, can be acted upon by metallic zinc, and when
these 40 grammes are so acted on, I gramme of hydrogen is again
evolved, and 72.5 grammes are obtained of a compound containing
no hydrogen, but sodium and zinc combined with all the oxygen
originally contained in the 1 8 grammes of water.

It will be evident, therefore, that the hydrogen contained in
water can be expelled in two equal moieties ; there must, therefore,
be two atoms of hydrogen in this compound. By no known

Atomic Weights 37

process can the oxygen be withdrawn from water in two stages :
thus, if 1 8 grammes of water are acted upon by chlorine, under the
conditions in which chemical action can take place, 73 grammes of
a compound containing only chlorine and hydrogen are formed, and
the whole of the oxygen is thrown out of combination and evolved
as gas. It is therefore concluded that water contains in its mole-
cule 2 atoms of hydrogen and i atom of oxygen, and as they are
combined in the relative proportion of I : 8, the atomic weight of
oxygen cannot be less than 16.

No compounds have been found in which a smaller weight of
oxygen, relative to one atom of hydrogen, than is represented by
the number 16 (approximately), is known to take part in a chemical
change.

The compound marsh gas contains hydrogen and carbon in
the proportion by weight of 1:3. By acting on this compound
with chlorine, it is possible to remove the hydrogen from it in
four separate portions.

By the first action of chlorine upon 16 grammes of marsh gas,
i gramme of hydrogen is removed in combination with 35.5
grammes of chlorine, and a compound containing carbon, hydrogen,
and chlorine, in the ratio 12 : 3 : 35.5, is formed.

By the successive action of chlorine, three other moieties of
hydrogen can be thus withdrawn, each being in combination with
its equivalent (35.5 parts) of chlorine. The second and third com-
pounds that are formed contain carbon, hydrogen, and chlorine in
the ratios 12:2: (35.5 x 2) and 12:1: (35.5 x 3).

The compound produced by the fourth action of chlorine, which
withdraws the fourth portion of hydrogen, contains only carbon
and chlorine, in the ratio 12 :(35-5 x 4). From the fact that the
hydrogen contained in marsh gas can thus be removed in four
separate portions, the molecule must contain four hydrogen atoms,
and therefore the atomic weight of carbon must be at least 12. No
compounds of carbon are known in which a smaller weight of
carbon, relative to one atom of hydrogen, than is represented by
the number 12, takes part in a chemical change.

The definition of atomic weight, furnished by considerations
of a chemical nature, may be thus stated : the atomic weight of an
element, is the number which represents how many times heavier
the smallest mass of that element capable of taking part in a
chemical change is, than the smallest weight of hydrogen which
can so function.

38 Introductory Outlines

The choice of hydrogen as the unit of atomic weights is a purely arbitrary
selection ; but since atomic weight values can only be determined relatively, it
becomes necessary to select some one element and to assign to its atom some
particular number to serve as a standard. As hydrogen is the lightest of all
elements, Dalton originally adopted it, and arbitrarily fixed unity as the
number which should stand for its atomic weight. The disadvantages of this
particular unit are twofold : in the first place the number of elements that form
hydrogen compounds that are suitable for atomic weight determinations is very
small, whereas nearly all the elements form convenient oxygen compounds, or
compounds with elements whose atomic weights with reference to oxygen are
accurately known, and in actual practice such compounds are almost always
made use of for such determinations. In the second place, the exact ratio of
the weights of an atom of hydrogen and oxygen is not known with certainty, so
that in calculating atomic weights that are determined with reference to oxygen,
possible errors may arise. The ratio Hydrogen : Oxygen is not exactly i : 16.
Various values have been obtained by different experimenters, and at the present
time i : 15.88 is accepted as more nearly the truth.

On account of the extreme difficulty of exactly determining this ratio,
chemists are now generally agreed in adopting as the unit in all exact determi-
nations of atomic weights a number which is ^th the weight of the atom of
oxygen : that is to say, the atomic weight of oxygen is in reality the standard,
and is fixed as 16, and the unit, instead of being the weight of one atom of
hydrogen, is T^th of this number.

The effect of this change is only of importance in cases of chemical investiga-
tion where a high degree of exactitude is required ; for purposes of ordinary
analyses and chemical calculations the difference that it makes is practically nil.
Fixing the atomic weight of oxygen at 16 merely raises the atomic weight of
hydrogen from i to 1.008. As the use of small decimal fractions introduces
unnecessary complications which tend to obscure simple processes of reasoning,
the -approximate atomic weights given in the third column of page 22 will be
employed for the most part in the following Introductory chapters.

The student will frequently meet with slight discrepancies between the
numbers given as the atomic weights of various elements by different writers.
Such discrepancies are often due to the fact that in some cases H = i is used
as the standard, and in others O = 16. For example, the atomic weight of
gold will be 195.7 m the first case, and 197.2 in the second; while with the
lighter metal aluminium the numbers will be 26.9 as against 27.1.

The discrepancy may also arise from the fact that the determination of
atomic weights by different experimenters often vary very considerably. With
a view to arrive at some uniformity, a conference of representative chemists
was held to consider the subject, and the atomic weights finally decided upon
by them were published under the title of International Atomic Weights. A
revised list of these weights is published annually in the Berichte, and in the
fourth column of the table on p. 22 will be found the latest values (1912).

2. Determination of Atomic Weights from Considerations
based upon Volumetric Relations. The Law of Gaseous
Volumes. — In the year 1805 the fact was discovered by Gay-
Lussac and, Humboldt, that when I litre of oxygen combines with

Atomic Weights 39

2 litres of hydrogen the vapour of water (or steam) which was
produced occupied 2 litres, the volumes in all cases being measured
under the same conditions of temperature and pressure.* This
fact led to the discovery of the simple relation existing between
the volumes of other reacting gases and the volume of the products :
thus it was found that —

i vol. of hydrogen unites with i vol. of chlorine, and gives
2 vols. of hydrochloric acid.

1 vol. of hydrogen unites with i vol. of bromine vapour, and

gives 2 vols. of hydrobromic acid.

2 vols. of hydrogen unite with i vol. of oxygen, and give

2 vols. of steam.

2 vols. of carbon monoxide unite with i vol. of oxygen, and
give 2 vols. of carbon dioxide.

1 vol. of carbon monoxide unites with i vol. of chlorine, and

gives i vol. of phosgene gas.

In the same way with compounds that cannot be obtained by
the direct union of their constituent elements, it is found that on
being subjected to processes of decomposition similar simple
volumetric relations exist : thus by suitable methods of decom-
position— : *

2 vols. of ammonia gas yield i vol. of nitrogen and 3 vols. of

hydrogen.
2 vols. of nitrous oxide yield 2 vols. of nitrogen and i vol. of

oxygen.
2 vols. of nitric oxide yield i vol. of nitrogen and i vol. of

oxygen.
I vol. of marsh gas yields 2 vols. of hydrogen and some solid

carbon, which cannot be volatilised, and therefore its

vapour volume is unknown.
i vol of ethylene yields 2 vols. of hydrogen and solid carbon

as in the preceding.

The observations of these and similar facts gave rise to the law
of Gay-Lussac, and it will be seen that there is evidently a close
connection between the simple -volumetric relations and those
existing between the multiple proportions by weight, in which one

* For the relations of gaseous volumes to temperature and pressure the
student is referred tc chapter ix. , on the general properties of gases.

4.O Introductory Outlines

element unites with another. For example, in the two oxides of
nitrogen the ratios of the two elements by weight are —

Nitrous oxide . . .. Nitrogen : oxygen = 28 : 16.
Nitric oxide . .v . Nitrogen : oxygen =14 : 16,

while the volumetric relation in which the two constituents are
present is —

Nitrous oxide . . . Nitrogen : oxygen = 2 : I.

Nitric oxide . . . Nitrogen : oxygen =i : i.

In other words, there is twice as much nitrogen by weight in the
one compound as in the other, and there is twice as much nitrogen
by volume in the one as compared to the other. Moreover, if 14
and 1 6 respectively represent the relative weights of atoms of nitro-
gen and oxygen, then the numbers representing the relative
'volumes in which these elements unite will also express the number
of atoms of each in the molecule.

The connection existing between the proportions in which
elements unite by weight, and by volume, was first explained by
the Italian physicist and chemist Avogadro, who in the year
1811 advanced the theory now recognised as a fundamental prin-
ciple, and known as Avogadro's hypothesis. This theory may be
thus stated : Equal volumes of all gases or vapours, under the
same conditions of temperature and pressure, contain an equal
number of molecules. If this be true, if there are the same
number of molecules in equal volumes of all gases, it must follow
that the ratio between the weights of equal volumes of any two
gases will be the same as that between the single molecules of the
particular gases. If a litre of oxygen be found to weigh sixteen
times as much as a litre of hydrogen (under like conditions of tem-
perature and pressure), inasmuch as there are the same number
of molecules in each, the oxygen molecule must be sixteen times
heavier than that of hydrogen ; and therefore by the comparatively
simple method of weighing equal volumes of different gases, it
becomes possible to arrive at the relative weights of their molecules.

The relative weights of equal volumes of gases and vapours, in
terms of a given unit, are known as their densities or specific
gravities. Sometimes densities are referred to air as the unit, but
more often hydrogen, as being the lightest gas, is taken as the
standard. Taking hydrogen as the unit, the density or specific
gravity of a gas is the weight of a given volume of it, as compared

Atomic Weights 41

with the weight of the same volume of hydrogen — or in other
words, the ratio between the weight of a molecule of that gas and
a molecule of hydrogen. The ratio that exists between the weight
of a gaseous molecule and half the weight of a molecule of hydrogen,
chemists term the molecular weight of that gas ; hence it will be
obvious that the number which represents the molecular weight of
a gas is double that of its density or specific gravity.

If i litre of hydrogen and I litre of chlorine be caused to combine,
2 litres of gaseous hydrochloric acid are formed. As equal volumes
of all gases (under like conditions) contain the same number of
molecules, in the 2 litres of hydrochloric acid there must be twice
as many molecules of that compound as there were of hydrogen
molecules in the I litre, or of chlorine molecules in the other.
• But each molecule of hydrochloric acid is composed of chlorine
and hydrogen (from other considerations one atom of each element),
therefore there must have been at least twice as many atoms
of hydrogen in the litre of that gas as there were molecules ;
and by the same reasoning, twice as many chlorine atoms in the
litre of chlorine as there were molecules : in other words, both
hydrogen and chlorine molecules consist of two atoms. The
molecular weight of hydrogen therefore is 2 ; that is, its molecule
is twice as heavy as its atom. The atom of hydrogen is the unit
to which molecular weights are referred, while the weight of the
molecule of hydrogen is taken as the standard of densities or
specific gravities.

In order, therefore, to find the molecular weight of any gas or
vapour, it is necessary to learn its density — that is, to ascertain
how many times a given volume of it is heavier than the same
volume of hydrogen,* and to double the number so obtained.t

The following table gives the densities or specific gravities of all
the elements whose vapour densities have been determined. The
list includes all those elements which are gases at the ordinary
temperature, and those that can be vaporised under conditions

* Certain exceptions to this rule are discussed under the subject of Dissocia-
tion, chap. x. p. 88.

f The specific gravity of hydrogen, as compared with air taken as unity,
is 0.0695, or air is 14.3875 times heavier than hydrogen. If, therefore, it be
desired to find the molecular weight of a given gas, whose density as compared
with air is known, it is only necessary to multiply its density (air=i) by the
number 14.3875, which gives its density as compared with hydrogen, and then
to double the number so obtained.

42 Introductory Outlines

which render such determinations experimentally possible. (Hy-
drogen being taken as unity, the other numbers are the approxi-
mate values, which for purposes of discussion are more suitable
than figures that run to two or three decimal places.)

Hydrogen ... . i

Helium . * . 2

Neon . . » .10

Nitrogen . . . 14

Oxygen .... . 16

Fluorine ... • 19

Argon , . . 20

Sulphur . . -32

Chlorine . . . 35.5

Krypton . . . 41

Xenon . . . . 64.0

Selenium ... 79
Bromine . . .80
Iodine . . . 127
Sodium . . . 11.5
Potassium . . . 19.5
Zinc .... 32.5
Cadmium . . .56
Mercury . . . 100
Phosphorus . . 62
Arsenic . . .150

Let us now consider how the knowledge of the relative weights
of gaseous molecules is utilised in assigning a particular number
as the atomic weight of an element.

The molecular weight of chlorine is 71. It has been shown that
the molecule certainly contains more than I atom, and probably 2,
in which case 35.5 would represent the relative weight of the
atom.

The compound hydrochloric acid has the molecular weight 36.5.
It has been already proved that this compound contains i atom of
hydrogen, therefore 36.5 - i = 35.5.

The compound carbon tetrachloride gives a molecular weight
154. Analysis shows that this compound contains 12 parts of
carbon in 154 parts, therefore 154—12=142 = 35.5x4.

In these three molecules the weights of chlorine relative to the
weight of i atom of hydrogen are 142, 35.5, and 71, the greatest
common divisor of which is 35.5. This number, therefore, is
selected as the atomic weight of chlorine.

Again, it has been shown that by the action of metals upon
water, the hydrogen contained in the water could be expelled in two
separate portions, thus proving that there must be 2 atoms of
hydrogen in the molecule of that compound.

The molecular weight of water is found to be 18 ; deducting from
this the weight of the two hydrogen atoms we get 18-2 = 16.

The molecular weight of carbon monoxide is 28 ; 28 parts of
this compound contain 12 parts of carbon, therefore 28- 12 = 16.

Atomic Weights 43

The molecular weight of carbon dioxide is 44 ; 44 parts of this
compound also contain 12 parts of carbon, therefore 44- 12 = 32.

When I litre of oxygen combines with two litres of hydrogen,
2 litres of water vapour are formed ; there are therefore twice the
number of water molecules produced as there are oxygen mole-
cules (since by Avogadro's hypothesis 2 litres contain twice as many
molecules as I litre). But each water molecule contains certainly
i atom of oxygen, therefore the original oxygen molecules must
have consisted of not less than 2 atoms. When the density of
oxygen is determined it is found to be 16, its molecular weight
therefore is 32.

In these four various molecules the weights of oxygen relative to
the weight of i atom of hydrogen are 16, 16, 32, 32, the greatest
common divisor of which is 16. This number, therefore, is selected
as the atomic weight of oxygen.

Again, it has already been shown that in the compound ammonia,
the hydrogen can be removed in three separate moieties, proving
that there must be three atoms of that element in the molecule.
The molecular weight of ammonia is found to be 17, therefore
17-3 = 14, which is the weight of the nitrogen.

The molecular weight of nitrous oxide is 44 ; 44 parts of this
compound are found to contain 16 parts of oxygen and 28 parts of
nitrogen.

The molecular weight of nitric oxide is 30 ; 30 parts of this
compound contain 16 parts of oxygen and 14 parts of nitrogen.

The molecular weight of nitrogen is found to be 28.

In these four different molecules the weights of nitrogen relative
to the weight of i atom of hydrogen are 14, 28, 14, 28, the
greatest common divisor of which is 14. The atomic weight of
nitrogen, therefore, is regarded as 14.

These three examples, namely, chlorine, oxygen, and nitrogen
are instances of elements which are gaseous at ordinary tempera-
tures ; but the same methods are applicable in the case of the non-
volatile elements, such as carbon, provided they furnish a number
of compounds that are readily volatile.

On comparing the numbers in the foregoing table (p. 42),
representing the densities of various elements, with the atomic
weights of those elements as given on p. 22, it will be seen
that in several cases the numbers given are approximately
the same. This agreement is merely because the molecules
of these elements consist of two atoms. The molecules of

44 Introductory Outlines

helium, neon, argon, krypton, xenon, sodium, potassium, zinc,
cadmium, and mercury consist of only one atom ; their atomic
weights, therefore, will be the same as their molecular weights, that
is, twice their densities. The elements arsenic and phosphorus, on
the other hand, contain in their molecules four atoms — that is to
say, the number which represents the smallest weight of phosphorus
and of arsenic, capable of taking part in a chemical change, is only
half the density, and therefore a fourth of the molecular weight.

The definition of atomic weight that is furnished by the con-
sideration of volumetric relations may be thus stated. The atomic
weight is the smallest weight of an element that is ever found in a
•volume of any gas or vapour equal to the volume occupied by one
molecule of hydrogen at the same temperature and pressure.

The volume occupied by one molecule of hydrogen is regarded
as the standard molecular volume, while that occupied by an atom
of hydrogen— or, in other words, the atomic volume of hydrogen — is
called the unit volume. The standard molecular volume, therefore,
is said to be two unit volumes; and as, from Avogadro's law, all
gaseous molecules have the same volume, it follows that the mole-
cules of all gases and vapours occupy two unit volumes. Atomic
weight may therefore be defined as the smallest weight of an
element ever found in two unit volumes of any gas or vapour.

The molecular volume of a gas is its molecular weight divided
by its relative density, a ratio which in all cases will obviously
equal 2, that is, two unit volumes.

The atomic volume of an element in the state of vapour is its
atomic weight divided by its relative density. In the case of such
elements as chlorine, nitrogen, oxygen, &c., whose molecules are
diatomic, the quotient will be I — that is to say, the'atomic volume
of these elements is equal to I unit volume. In the case of mer-

, atomic weight = 200

cury vapour, however, we have = : — " =2.

density =100

The atomic volume of mercury vapour, therefore, is equal to 2
unit volumes, and is identical with its molecular volume.

On the other hand, with the element phosphorus the atomic

. atomic weight=3i
volume is - — ^ensit =62 '$> or one'half tne umt volume,

and therefore one-fourth the molecular volume ; consequently, four
atoms exist in this molecule.

The method of determining atomic weights based upon volu-
metric relations, when taken by itself, is not an absolutely coi'tain

Atomic Weights 45

criterion, for although the atomic weight of an element cannot be
greater than the smallest mass that enters into the composition of
the molecules of any of its known compounds, it might be less than
this, as there is always the possibility of a new compound being
discovered, in which the relative weight of an element is such as to
make it necessary to halve the previously accepted atomic weight.
3. Determination of Atomic Weight from the Specific
Heat of Elements in the Solid State.— When equal weights of
different substances are heated through the same range of tempera-
ture, it is found that they absorb very different quantities of heat,
and on again cooling to the original temperature, they consequently
give out different amounts of heat. Thus, if I kilogramme of water,
and i kilogramme of mercury be each heated to a temperature of
100°, and then each be poured into a separate kilogramme of water
at o°, in the first case the resultant mixture will have a temperature
of 50°, while in the second it will only reach the temperature of 3.2°;
that is to say, while the water in cooling through 50° has raised the
temperature of an equal weight of water from o° to 50°, the amount
of heat in I kilogramme of mercury at 100° has only raised the
temperature of an equal weight of water from o° to 3.2°, and in so
doing has itself become loweredin temperature 100 - 3.2 = 96.8°. The
amount of heat contained, therefore, in equal weights of water and of
mercury at the same temperature, as shown by these figures, is as —

52. 3£_ i .
50-96.8" '™'

therefore it requires 30 times as much heat to raise a given weight
of water through a given number of degrees as to raise an equal
weight of mercury through the same interval of temperature, or
the thermal capacity of mercury is ^th that of water.

The specific heat of a substance is the ratio of its thermal
capacity to that of an equal weight of water ; or, the ratio between
the amount of heat necessary to raise a unit weight of the sub-
stance from o° to i °, and that required to raise the same weight
of water from o° to i° ; thus, the specific heat of mercury is ^j, or
0.033. Water is chosen as the standard of comparison because it
possesses the highest thermal capacity of all known substances ;
the numbers, therefore, which express the specific heats of other
substances are all less than unity.

Dulong and Petit were the first to draw attention (1819) to a
remarkable relation which exists between the specific heats and
the atomic weights of various solid elementst whose specific heats

46 Introductory Outlines

they themselves had determined. They found that the specific
heats of the solid elements were inversely as their atomic weights ;
that is to say, the capacity for heat of masses of the elements pro-
portional to their atomic weight was equal. This law, known as
the law of Dulong and Petit, may be thus stated : The thermal
capacities of atoms of all elements in the solid state are equal.

The thermal capacity of an atom is termed its atomic heat;
hence the law may be more briefly stated, all elements in the
solid state have the same atomic heat. This important constant
is the product of the atomic weight into the specific heat. From
the following table it will be seen that the number expressing
the atomic heat is not perfectly constant : the departures from the
mean 6.4 are, as a rule, only slight, and may be attributed to
the fact that the determinations are not always made upon the
elements under conditions that are strictly comparable. At the
end of the table, however, there are certain elements which appear
to present marked exceptions to the law.

Element.

Specific
Heat.

Atomic Atomic
Weight. Heat.

Lithium . •.-*$&

0.94

7 =

6.6

Sodium .

0.29

23 =

6-7

Potassium

0.166

39 =

6.5

Manganese

0.122

55 =

6.7

Iron.

O.II2

56 -

6-3

Silver

0.057

1 08 -

6.1

Gold

0.032

196 =

6.2

Mercury (solid)

0.032

200 =

6.4

Lead

0.031

206.4 =

6.5

/'Beryllium

0.41

9.1 =

3-7

J Boron (cryst.) .

0.25

ii =

2.75

| Carbon (diamond) .

0.147

12 =

1.76

V Silicon (cryst.).

0.177

28 =

4-95

It will be seen that, relatively speaking, the four elements
which show a considerable departure from the law of Dulong are
elements with low atomic weights. Low atomic weight, however,
is not always accompanied by such deviation, as is shown in the
case of lithium and sodium.

When the different allotropes of carbon are experimented upon,
it is found that the departure is not the same for each modification
of the element, thus —

Atomic Weights 47

Element. Specific Atomic Atomic

Heat. Weight. Heat.

Diamond . . . 0.147 x 12 = 1.76
Graphite . . . 0.200 x 12 = 2.40
Charcoal . . . 0.241 x 12 = 2.90

It has been observed that, as a general rule, the specific heat of
an element is slightly higher at higher temperatures ; but in the
case of the four elements showing abnormal atomic heats, this
increase rises rapidly with increased temperature, until a certain
point is reached, when it remains practically constant, and repre-
sents an atomic heat which closely approximates to the normal
value ; thus in the case of diamond, the specific heat at increasing
temperatures is —

Specific Atomic Atomic
Heat. Weight. Heat.

Diamond at 10.7° . . . 0.1128 x 12 = 1.35
„ 45° • • • 0.1470 x 12 = 1.76

„ 206° . . . 0.2733 x 12 = 3.28

„ 607° . . . 0.4408 X 12 = 5.30

„ 806° . . . 0.4489 X 12 = 5.4

„ 985° • • 0.4589 X 12 - 5.5

The same result is seen in the case of graphite, and it is also to
be remarked, that while at low temperatures there exists a wide
difference between the specific heats of these two modifications of
carbon, this difference vanishes at a temperature of about 600°.

Specific Atomic Atomic
Heat. Weight. Heat.

Graphite at 1 0.8° . . . 0.1604 x 12 = 1.93
„ ' 61.3° . . . 0.1990 x 12 = 2.39
„ 642° . . . 0.4454 x 12 = 5.35

„ 978° . . . 0.4670 X 12 = 5.50

Both the elements boron and silicon are found to follow the
same rule, and at moderate temperatures their atomic heats nearly
approximate the normal constant.

The case of the somewhat rare element beryllium is of special
interest from another point of view, which will be referred to when
treating of the natural classification of the elements : from the
following numbers* it will be seen that its atomic heat very
rapidly rises with moderate increase of temperature.
* Humpidge.

48 Introductory Outlines

Specific Atomic Atomic
Heat. Weight. Heat.

Beryllium at 100° . . . 0.4702 x 9.1 = 4.28

„ 200° . . . 0.5420 x 9.1 = 4.93

„ 400° . . . 0.6172 x 9.1 = 5.61

„ 500° . . . 0.6206 x 9.1 = 5.65

The relation between atomic weight and specific heat, established
by Dulong and Petit, is of service in the determination of atomic
weights, not as a method of ascertaining the exact value with any
degree of refinement, but rather as a means of deciding between
two numbers which are multiples of a common factor.

If specific heat x atomic weight = atomic heat, it will be obvious
that, if we experimentally determine the specific heat, and divide
that value into the constant atomic heat, 6.4, we obtain the
approximate atomic weight.

The two following examples will serve to illustrate the applica-
tion of the method.

The element indium combines with chlorine in the proportion —

Indium : chlorine = 37.8 : 35.5.

If InCl is the formula, then 37.8 is the atomic weight of indium :
but from the chemical similarity between indium and zinc (whose
chloride has the formula ZnCl2), it was believed that the formula
for indium chloride was InCl2, in which case, in order to preserve
the ratio between the two elements, the atomic weight would have
to be 37.8 x 2 = 75.6.

When the specific heat of indium was determined,* it was found
to be 0.057.

Therefore the atomic weight must be raised by one-half, from
75.6 to 113.4, and the formula for the chloride will be InCl3.

The element thallium combines with chlorine in the proportion—

Thallium : chlorine = 203.6 : 35.5.

In some of its compounds thallium exhibits a strong resemblance
to potassium, the chloride of which has the formula KC1. If the
formula for the thallium chloride is T1C1, the atomic weight of the
metal must be 203.6.

In many respects thallium exhibits a striking analogy with lead,
* Bunsen, 1870.

Atomic Weights 49

the chloride of which has the formula, PbCl2. If thallium chloride
has a corresponding formula, T1C12) then the atomic weight of
thallium must be raised to 407.2.

When the specific heat of thallium was ascertained,* it was found
to be 0.0335.

6.4

This result shows that the number 203.6 and not 407.2 is the
atomic weight of thallium, and that the chloride has the formula
T1C1.

Molecular Heat Of Compounds.— The capacity for heat of an
atom undergoes no alteration when the atom enters into combina-
tion with different atoms — in other words, the atomic heat of an
element is the same in its compounds. The molecular heat of a
compound (that is, the product of the molecular weight into the
specific heat) will therefore be the sum of the atomic heats of its
constituent elements. Hence it is possible to calculate what will
be the atomic heat of an element which does not exist as a solid
under ordinary conditions ; and therefore the atomic weight of
such an element, as deduced from other considerations, is capable
of verification, by determinations of the molecular heat of various
of its compounds : thus —

The specific heat of silver chloride, AgCl, is 0.089 : —

Specific Molecular Molecular

Heat. Weight. Heat.

0.089 * 143.5 = J2-77.

The atomic heat of silver = 6.1, therefore, as deduced from this
compound, the atomic heat of chlorine is 12.77 ~ 6.1 = 6.6.
Again, the specific heat of stannous chloride, SnCl2, is 0.1016: —

Specific Molecular Molecular

Heat. Weight. Heat.

0.1016 x 189 = 19.2.

The atomic heat of tin is 6.6, therefore the atomic heat of two
atoms of chlorine, as deduced from this compound, is 19.2 — 6.6 =
12.6, giving 6.3 as the atomic heat of chlorine.

The differences that appear in the value, as deduced from
various compounds, are lessened, because the errors of the
method are more equally distributed, if we divide the molecular
heat by the number of atoms in the molecule. Thus, in the

* Regnault.

50 Introductory Outlines

two examples quoted, silver chloride consists of two atoms, while
the molecule of stannous chloride contains three ; if, therefore, the
molecular heats of these two compounds are divided respectively
by 2 and by 3 we get—

as the value representing the atomic heat of chlorine.

The element calcium combines with chlorine in the proportion-

Calcium : chlorine = 20 : 35.5.

If the atomic weight of calcium is 20, the formula will be CaCl,
whereas if 40 is the atomic weight of the metal, the compound
must be represented by the formula CaCl2.

The molecular weight of CaCl would be 55.5, that of CaCl2 n i.o.

When the specific heat of .the compound was determined, it
was found to be 0.1642. In order, therefore, to decide between
the two values for the atomic weight of calcium, we calculate the
molecular heat from both of the molecular weights, and divide the
result by the number of atoms in the molecule in each case.

On the supposition that Ca = 2O, and that CaCl represents the
chloride : —

Cad . . . 0-1642^ 55.5=455.

Or, if Ca = 40, and CaCl2 is the formula for the chloride, then—

~ ~, o.i642X 1 1 i.o

CaCl2. . . - 2 - _

The number 6.07, which nearly agrees with the constant 6.4,
decides the value 40 as the atomic weight of calcium. The
element calcium is one of those metals which it is very difficult to
isolate and obtain in a state of purity, but when in recent years
the specific heat of this metal was experimentally determined,*
it was found to be 0.1704 : —

0.1704x40=6.8.

Thus affording direct confirmation of the value 40 for the atomic
weight of calcium, which had been deduced from the molecular
heat of its compounds.

* Bunsen,

Atomic Weights 51

Deductions based upon molecular heats of compounds are only
trustworthy in the case of the most simply constituted compounds.

4. Determination of Atomic Weight from Considerations
based on Isomorphism. — It was early observed that certain rela-
tions existed between the crystalline forms of compounds and their
chemical composition. Mitscherlich found that certain substances
having an analogous chemical composition, as, for example, sodium
phosphate and sodium arsenate, crystallised in the same geometric
form. In the year 1821 he stated his law of isomorphism as follows :
" The same number of atoms, combined in the same way, give rise
to the same crystalline form, which is independent of the chemical
nature of the atoms, being influenced only by their number and
mode of arrangement." Subsequent investigations, however, have
shown that this statement is too general.

In its broad sense as signifying the same crystalline form,
isomorphism is found to exist —

i. Between compounds containing the same number* of atoms
similarly combined, and which bear close chemical analogies to
each other.

, ( Zinc sulphate .... ZnSO4,7H2O.
15 {Magnesium sulphate. . . MgSO4,7H2O.

J Hydrogen disodium phosphate . HNa2PO4,12H2O.
13 I Hydrogen disodium arsenate . HNa2AsO4,12H2O.

/ Rubidium alum . . . . Rb2SO4,Al2(SO4)3,24H2O.
Isomor hous J Potassium chrome alum . . K2SO4,Cr2(SO4)3l24H2O.

13 ) Potassium aluminium selenium 1 v ^^(SeO^UHzO.

2. Between compounds containing a different number of atoms,
but which also bear close chemical analogies to one another.

Potassium chloride .

Potassium sulphate . . .

3. Between compounds containing either the same or a different
number of atoms, and which exhibit little or no chemical analogies.

( Sodium nitrate .... NaNO3.
Isomorphous < . ~ nr.

I Calcium carbonate . . . CaCO3.

Isomorphous { Sodium sulphate (anhydrous) . Na2SO4.
I Barium permanganate . . BaMn2O8.

52 Introductory Outlines

Isomorphism of this order, where little or no chemical relations
exist between the compounds, is sometimes distinguished as
isogonism. It must not be supposed, that because two chemically
analogous compounds contain the same number of atoms, they will
necessarily crystallise in the same form : there are indeed a large
number of similarly constituted analogous compounds that do not
exhibit isomorphism.

No simple definition of isomorphism is possible, but the following
test is generally accepted as a criterion, namely, the power to form
either mixed crystals or layer crystals. Thus, when two substances
are mixed in a state of liquidity, and allowed to crystallise, if the
crystals are perfectly homogeneous, they are known as mixed
crystals, and the substances are regarded as isomorphous.

Or when a crystal of one compound is placed in a solution of
another compound, and the crystal continues to grow regularly
in the liquid, the compounds are isomorphous. Thus, if a crystal
of potassfam alum (white) be placed in a solution of manganese
alum, the crystal continues to grow without change of form, and
a layer of amethyst-coloured manganese alum is deposited upon it.

In making use of the law of isomorphism in the determination of
atomic weights, it is assumed that the weights of different atoms
that can mutually replace each other without altering the crystal-
line form are proportional to their atomic weights."*

Thus, if we suppose that, in the case of the sulphates of zinc
and magnesium, the atomic weight of zinc is known, viz., 65, and
that of magnesium is doubtful ; from the fact of the isomorphism
of the sulphates it may be premised that the elements are present in
proportions relative to their atomic weights. Analysis shows that
the proportion is 24 of magnesium to 65 of zinc, therefore 24 is pre-
sumably the atomic weight of magnesium.

In this way Berzelius corrected many of the atomic weights
which in his day had been assigned to the elements.

* The group (NH4) may be regarded as an atom, having the relative weight 18.

CHAPTER VII
QUANTITATIVE CHEMICAL NOTATION

THE use of chemical symbols and formulae, as a convenient means
of representing concisely the qualitative nature of chemical changes,
has been explained in chapter iv. We are now in a position to
read into these symbols a quantitative significance, which at that
stage it would have been premature to explain.

The symbol of an element stands for an atom ; but, as we have
now learnt, the atoms of the various elements have different relative
weights, hence these symbols represent relative weights of matter.
The symbol Na signifies 23 relative parts by weight of sodium, O
stands for 16 relative parts by weight of oxygen, H for i part of
hydrogen ; in other words, the weight of sodium represented by
the symbol Na is 23 times as heavy as that which is conveyed
by a symbol H. A chemical equation, therefore, is a strictly
quantitative expression, in which certain definite weights of matter
are present in the form of the reacting substances, and which
reappear without loss or gain in the compounds resulting from the
change. In this sense a chemical equation is a mathematical
expression. Thus, the equation —

Na + Cl = NaCl,

not only means that an atom of sodium combines with an atom of
chlorine and forms i molecule of sodium chloride, but it also means

23 + 35-5 = 58.5
Na Cl NaCl.

In other words, that sodium and chlorine unite in the relative pro-
portion of 23 parts of the former and 35.5 parts of chlorine, and
produce 58.5 parts of sodium chloride.

In the same way, into the equation which expresses the action of

54 Introductory Outlines

sulphuric acid upon sodium carbonate, we read the quantitative
meaning of the symbols—

H2SO4 + Na2CO3 = Na2S04 + CO2 4- H2O.
2 46 46

32 12 32 12 2

64 48 64 32 16

98* + 106 = 142 4- 44 + 18

That is to say, 98 parts by weight of sulphuric acid act upon
106 parts of sodium carbonate, producing 142 parts of sodium
sulphate, 44 parts of carbon dioxide, and 18 parts of water. It will
be evident that it becomes a matter of the simplest arithmetic to
calculate the weight of any product that can be obtained from a
given weight of the reacting substances ; or vice versa, to find
the weight of any reacting substance which would be required to
produce a given weight of the product of the action.

Not only is information respecting the quantitative relations
by weight embodied in a chemical equation, but when gaseous
substances are reacting, the equation also represents the volu-
metric relation between the gases. In order that the volumetric
relations may be more manifest, the equations expressing the re-
actions are written in such a manner as to represent the molecules
of the substances.

H + Cl = HC1

is an atomic equation, but as the molecule is the smallest particle
which can exist alone, a more exact statement of the chemical
change is made, by representing the action as taking place between
molecules, thus —

H2 + C12 = 2HC1.

From such an equation we see that i molecule of hydrogen, or
2 unit volumes, unites with i molecule or 2 unit volumes of chlorine,
and forms 2 molecules or 4 unit volumes of hydrochloric acid :
or again —

O2 + 2H2 = 2H2O.

One molecule, or 2 unit volumes of oxygen, unite with 2 mole-
cules, or 4 unit volumes of hydrogen, and produce 2 molecules of

* The number obtained by adding together the weights of the atoms in a
formula is known as a " formula weight," thus 98 is the formula weight of
sulphuric acid.

Quantitative Notation 55

water, which when vaporised, and measured under the same con-
ditions of temperature and pressure, occupy 4 unit volumes. In
other words, the number of molecules, in all cases * where gases
and vapours are concerned, represent exactly the volumetric
relations. In the cases quoted, it will be observed, the same ratio
also subsists between the number of atoms of the reacting gases
and the molecules of the compound, but this is not always the
case, for example —

Atomic equation, Hg + 2C1 = HgCl2.

In this equation 3 atoms unite to produce i molecule, but the
ratio between the volumes is not represented by the statement,

1 volume of mercury vapour and 2 volumes of chlorine produce

2 volumes of vapour of mercury chloride.

Molecular equation, Hg + C12 = HgCl2.

By this we see that i molecule t (2 unit volumes) of mercury
vapour and i molecule (2 unit volumes) of chlorine give i mole-
cule (2 unit volumes) of vapour of mercury chloride.

Again,

P + 3C1 = PC13

is an atomic equation, showing that i atom of phosphorus unites
with 3 atoms of chlorine ; but it is not true that the ratio between
the volumes is represented by the statement, i volume of phos-
phorus vapour combines with 3 volumes of chlorine and gives 2
volumes of the vapour of phosphorus trichloride, as will be seen
by comparison with the molecular formulae —
P4 + 6C12 = 4PC13.

This equation tells us that i molecule J (2 unit volumes) of phos-
phorus vapour combines with 6 molecules (12 unit volumes) of
chlorine, producing 4 molecules (8 unit volumes) of phosphorus
trichloride vapour.

Knowing the relative densities of gases compared with hydro-
gen, it is obviously possible, by ascertaining the actual weight in
grammes of some definite volume of hydrogen, to calculate the
actual weight of any given volume of any other gas.

Two units are in common use, namely —

* See Dissociation, where apparent exceptions are explained.

t The atomic volume of mercury vapour being equal to 2 unit volumes (p. 44).

t The atomic volume of phosphorus is .5 of a unit volume (p. 44).

56 Introductory Outlines

(i.) The weight of i litre of hydrogen, measured at a temperature
of o° C., and under a pressure of 760 mm. of mercury.*

(2.) The volume occupied by i gramme of hydrogen, measured
under the same conditions.

I. One litre of hydrogen, measured at the standard temperature
and pressure, weighs .0896 grammes.t This number is known as
the crith;\ and by means of it the weight of i litre, and therefore
any given volume, of any gas can be deduced : thus, the relative
densities of oxygen, nitrogen, and chlorine are 16, 14, and 35.5
respectively, therefore i litre of these gases (measured always at
the standard temperature and pressure) weighs 16 criths, 14 criths,
and 35.5 criths respectively, or —

i litre of oxygen weighs 16 x. 0896= 1.4336 grammes,
i „ nitrogen „ 14 x .0896= 1.2544 „
I „ chlorine „ 35,5 X. 0896 = 3. 1808 „

So also with reference to compound gases, where in each case
the density is represented by the half of the molecular weight.
Thus, the relative densities of hydrochloric acid, ammonia, and
carbon dioxide are —

i+35-i

HCl1^'^ 18.25,

NH3 ^5=8.5,
C02^±32 = 22,

and the weights of i litre of these gases are therefore —

I litre of hydrochloric acid= 18.25 x -0896= 1.6352 gramme.
i „ ammonia = 8.5 x. 0896=0. 7610 „

I „ carbon dioxide =22.0 x. 0896 =1.97 12 „

II. The volume occupied by i gramme of hydrogen at the
standard temperature and pressure is 11.127 litres. As the rela-
tive density of oxygen is 16, it obviously follows that 16 grammes

* This temperature and pressure is chosen as the standard at which volumes
of gases are compared. See General Properties of Gases, chapter ix.

f From time to time slightly different values have been given for this
constant. The most recent determinations give the number .089873.

J From the Greek, signifying a barley-corn, and used symbolically to denot*
a little weight.

Quantitative Notation 57

of this gas will also occupy 11.127 litres; in other words, this
number 11.127 represents the volume in litres of any gas, which
will be occupied by the number of grammes corresponding to its
relative density, thus —

1 4 grammes of nitrogen . . occupy 11.127 litres.
35.5 „ chlorine . :. „ 11.127 „
18.25 55 hydrochloric acid „ 11.127 »
22.0 „ carbon dioxide . „ 11.127 „

The number of grammes of a substance, equal to the number
which represents its molecular weight, is spoken of as ^gramme-
molecule. The molecular weight of hydrogen = 2, therefore the
gramme-molecule of hydrogen (that is, 2 grammes of hydrogen)
will occupy 11.127x2 = 22.25 litres. The molecular weight of
oxygen = 32, therefore 32 grammes of oxygen will occupy 22.25
litres ; in other words, 22.25 litres is the volume which will be
occupied by the gramme-molecule of any gas.

By means of this important constant, 22.25, the volume of any,
or all, of the gaseous products of a chemical change (when
measured at the standard temperature and pressure) can be de-
duced directly from the equation representing the change, thus —

Zn + H2SO4=ZnSO4+H2

expresses the reaction taking place when zinc is dissolved in
sulphuric acid. Just as in the former illustrations it carries the
information that 65 grammes of zinc + 98 grammes of sulphuric
acid produce 161 grammes of zinc sulphate and 2 grammes of
hydrogen. But 2 grammes of hydrogen occupy 22.25 litres, there-
fore by the solution of 65 grammes of zinc, the volume of hydrogen
obtained will be 22.25 litres.

So also in the following equation, which represents the formation
of carbon dioxide from chalk (calcium carbonate) by the action
upon it of hydrochloric acid—

CaCO3 + 2HC1 = CaCl2 •+ H2O + CO2.

40+12 + 48 2(1 + 35.5) 40 + 71 2 + 16 12 + 32

100 + 73 = in + 18 + 44

100 grammes of chalk, when acted upon by 73 grammes of hydro-
chloric acid, yield in grammes of calcium chloride and 18
grammes of water, and 44 grammes of carbon dioxide.

Carbon dioxide is gaseous, therefore 44 grammes (the gramme-

58 Introductory Outlines

molecule) will occupy, at the standard temperature and pressure,
22.25 litres ; hence, by the decomposition of 100 grammes of
chalk, 22.25 litres of carbon dioxide are produced.

This chapter may be concluded with one illustration of the
methods employed in the exact determination of atomic weights
which depends essentially upon the quantitative character of
chemical reactions. By the three following processes the atomic
weights of chlorine, potassium, and silver may be deduced.

1. By heating a known weight of potassium chlorate, the formula
weight of potassium chloride is found —

KC1O3 = KC1 + 3O.

50 grammes of potassium chlorate when heated left a residue
of potassium chloride weighing 30.395 grammes. 50 - 30.395 =
19.605 e= grammes of oxygen evolved.

As potassium chlorate contains in its formula weight 3 atoms
of oxygen (16 x 3 = 48), we get the expression —

19.605 : 30.395 = 48 : 744o=formula weight of potassium chloride.

2. By dissolving a known weight of potassium chloride, and
adding to it excess of silver nitrate, silver chloride is precipitated,
which can be washed and dried and weighed, and from which the
formula weight of silver chloride is obtained —

KC1 + AgN03 = AgCl + KN03.

10 grammes of potassium chloride were found to yield 19.225
grammes of silver chloride ; therefore,

10 : 19.225 = 74.40 : 143.03 = formula weight of silver chloride.

3. By the direct combination of silver and chlorine, by heating
the metal in a stream of the gas, the ratio of chlorine to silver in
silver chloride is found :

10 grammes of silver so treated yielded 13.285 grammes of silver
chloride ; therefore,

13.285 : 10 = 143.03 : 107.66 = atomic weight of silver.
Since the formula weight of silver chloride, AgCl = 143.03,

therefore, 143.03- 107.66 = 35«37 = atomic weight of chlorine.
And since the formula weight of potassium chloride, KC1 = 74.40,

therefore, 74.40 — 35.37 = 39.03 = atomic weight of potassium.

CHAPTER VIII

VALENCY OF THE ELEMENTS

WHEN chlorine unites with hydrogen, the combination takes place
between one atom of chlorine (relative weight = 35.5) and one
atom of hydrogen (relative weight = i) ; but when oxygen com-
bines with hydrogen, one atom of oxygen unites with two atoms
of hydrogen. The compound ammonia consists of one atom of
nitrogen, combined with three atoms of hydrogen ; while one atom
of carbon, on the other hand, can unite with four atoms of
hydrogen.

One atom of chlorine never combines with more than one atom
of hydrogen ; its affinity for that element is satisfied, or saturated^
by union with one atom.

The affinity of one atom of oxygen for hydrogen, however, is
not satisfied by one atom of that element, but requires two atoms
for its saturation ; while nitrogen requires three, and carbon four
hydrogen atoms, in order to satisfy their respective affinities for
this element. . •

This varying power of combining with hydrogen is seen in a
number of other instances : thus, the elements fluorine, bromine,
and iodine, resemble chlorine in being only able to unite with one
atom of hydrogen. Sulphur, like oxygen, has its affinity for
hydrogen saturated by two atoms of that element. Phosphorus
and arsenic require three atoms of hydrogen in order to saturate
their combining capacity, while silicon resembles carbon in com-
bining with four hydrogen atoms. This combining capacity of
an element is termed its valency. Elements like chlorine,
fluorine, bromine, and iodine, whose atoms are only capable
of uniting with one atom of hydrogen, are called monovalent
(or sometimes monad) elements ; while those whose atoms com-
bine with two, three, or four hydrogen atoms, are distinguished
as di-valent (or dyad), tri-valent (or triad), and tetra-valent (or
tetrad) elements. All elements, however, are not capable of

60 Introductory Outlines

entering into combination with hydrogen ; in which case, their
valency is measured by the number of atoms of some other
monovalent element which is capable of satisfying their com-
bining capacity. Thus : —

atom of sodium combines with i atom of chlorine, forming NaCl.
,, calcium ,, , 2 atoms , ,, CaCl2.

boron ,,
tin ,,

phosphorus "c
tungsten , ,

BC13.
SnCl4.
PC15.
WC16.

In the combinations of elements with hydrogen alone, no in-
stances are known in which a higher valency is exhibited than
that of four ; but with chlorine, as here seen, cases are known in
which elements exhibit pentavalent and hexavalent characters.

Measured by their combining capacity for hydrogen and chlorine,
elements do not, however, always exhibit the same valency : thus,
the affinity of phosphorus for hydrogen is satisfied by three hydrogen
atoms, whereas one atom of this element can unite with five atoms
of chlorine.

As measured by hydrogen, the valency of sulphur is two, the
compound that it forms with hydrogen being expressed by the
formula SH2, while, as estimated by its capacity for chlorine, it
becomes tetravalent, as seen in the compound SC14. As a general
rule, however, the highest number of monovalent atoms with which
one atom of an element is capable of combining is accepted as
representing the valency of that element. Thu%, one atom of
phosphorus not only combines with five atoms of chlorine, but
also with five atoms of fluorine ; phosphorus is therefore a penta-
valent element.

As measured by hydrogen alone, or by chlorine alone, nitrogen
is a trivalent element, for the largest number of these atoms with
which one atom of nitrogen can unite is three, as seen in the
compounds having the composition NH3 and NC13 ; neverthe-
less, one atom of nitrogen is capable of combining with four
atoms of hydrogen and one of chlorine, forming the compound
NH4C1, ammonium chloride, in which the nitrogen atom is penta-
valent.

This rule, however, is not always followed ; for example, one
atom of iodine will unite with three atoms of chlorine, forming the

* Phosphorus also combines with hydrogen.

Valency of the Elements 61

compound IC13, but iodine is not generally regarded as a trivalent
element.*

In symbolic notation, this power possessed by an atom, of uniting
to itself monovalent atoms, is often represented by lines, each line
signifying the power of combination with one monovalent atom.
Thus, in the symbol H— Cl, the line is intended to give a concrete
expression to the fact that both hydrogen and chlorine are mono-
valent elements, and that the affinity of each element for the
other is satisfied when one atom of the one unites with one atom of
the ether. The symbol H— O— H, in like manner, signifies that
the oxygen atom is divalent, that its affinity for hydrogen is satisfied
only when it has united with two monad atoms. In the same way
we may express the facts that nitrogen and carbon, in their com-
binations with hydrogen, are respectively trivalent and tetravalent,

by the symbols H— N— H, and H— C— H. These lines are merely

H H

a convenient symbolic expression for the operation of the force of
chemical affinity ; their length and direction bear no meaning.!
The power to combine with one monovalent atom is sometimes
spoken of simply as one affinity : thus it is said that in the com-
pound having the composition PH3, or H — P — H, three of the

affinities of the phosphorus atom are saturated, and that two
affinities still remain unsatisfied, phosphorus, as already stated,
being a pentavalent element.

* See Iodine, Compounds.

f The student cannot be too often warned against attaching any materialistic
significance to these lines. The use of this convention is always attended with
the danger that the beginner is liable to fall into the error of regarding these
lines as representing in some manner fixed points of attachment, or links,
between the atoms. It must be remembered, therefore, that these lines not only
have no materialistic signification, but they must not even be regarded as convey-
ing any statical meaning. The atoms are undergoing rapid movements with
respect to each other, which movements are in some way governed by the
chemically attractive force exerted by the individual atoms upon one another ;
and the molecule will be more correctly considered, if we regard its atoms as
being held together in a manner resembling that by which the numbers of a
cosmical system are bound together. The lines simply denote that the atoms
are held to each other by the attractive force which we call chemical affinity.

62 Introductory Outlines

Compounds of this order, in which one of the elements has still
unsatisfied affinities, are called unsaturated compounds.

In its power to satisfy the affinities of an element, a divalent
atom is equal to two monovalent atoms : thus, when the affinities
of the tetravalent carbon atom are saturated with oxygen, the mole-
cule contains two atoms of oxygen, which may be symbolically
expressed thus, O = C = O, in which the four affinities of the
carbon (represented by the four lines) are satisfied by the two
divalent atoms of oxygen. Carbon, however, combines with a
smaller proportion of oxygen, forming the compound carbon mon-
oxide, CO. The carbon atom in this case is divalent, as" expressed
by the formula C = O, and this substance. is also an unsaturated
compound.

The number of divalent atoms with which an element can unite
cannot, however, be taken as a safe criterion or measure of the
valency of that element in cases where that number is greater
than i ; for example, in such a compound as calcium oxide, CaO,
we regard the two affinities of the divalent atom of oxygen as being
satisfied by two affinities possessed by the calcium, and express this
belief in the formula Ca = O, and regard the calcium as divalent.
In the same way, in carbon monoxide, CO, the carbon being united
with one atom of the divalent element oxygen is itself divalent in
this compound ; but in the case of carbon dioxide, where the carbon
atom is united with two atoms of divalent oxygen, we are not
justified in asserting that the atoms are united, as represented by
the formula O = C = O, in which the four affinities of carbon
are represented as saturated with oxygen. There exists no posi-
tive proof that the carbon is not divalent in this compound, and
that the molecule does not consist of three divalent atoms united,

as shown in the formula /\. From the fact, however, that car-
O O

bon forms a compound with four atoms of hydrogen, and another
with four atoms of chlorine, we know that this element is tetra-
valent, and therefore we believe that in carbon dioxide it is also
tetravalent.

Again, as measured by its compound with hydrogen, sulphur is
divalent ; while with chlorine it forms SC14. But sulphur unites
with oxygen, forming the two compounds sulphur dioxide, SO2, and
sulphur trioxide, SO3. If it be assumed that in these molecules the

Valency of the Elements 63

whole of the oxygen affinities are satisfied with sulphur, then the
symbolic representation of these oxides will be O = S = O, and
O = S = O, the sulphur being in one case tetravalent and in the

other hexavalent. There is, however, no positive proof that the
affinities of one oxygen atom are not partially satisfied by union
with another oxygen atom, and that the valency of the sulphur is
higher than either two or four, as seen in the alternative formulae,

S °\ S

S02 /\ S03

= O; or

0—0 Q/ 0-0-0

Chemists believe, however, that in these two oxides the
sulphur functions in the one case as a tetravalent, and in
the other as a hexavalent element ; and this belief is strengthened
by the recent discovery (Moissan) of a fluoride having the com-
position SF6, in which the hexavalent character of sulphur is
unquestionable.

It will be evident from these considerations, that in many cases
the valency of an element is a variable quantity, depending partly
upon the particular atoms with which it unites. It is also found
that it is dependent in many instances upon temperature and
upon pressure. Thus, between a certain limited range of
temperature, one atcm of phosphorus combines with five atoms
of chlorine in the compound PC15, but above that limit two atoms
of chlorine leave the molecule, and the phosphorus becomes tri-
valent. Again, if hydrogen phosphide, PH3, be mixed with hydro-
chloric acid, HC1, and the mixed gases be subjected to increased
pressure, the gases combine and form a solid crystalline com-
pound known as phosphonium chloride, PH4C1, in which the
phosphorus atom, being united with five monovalent atoms, is
pentavalent. When the pressure is released an atom of hydrogen
and an atom of chlorine leave the molecule, and the phosphorus
returns to its trivalent condition.

A compound, in whose molecules there is an atom which for the
time being is not functioning in its highest recognised valency,
often exhibits a readiness to unite with additional atoms to form

64 Introductory Outlines

new compounds : thus ammonia combines eagerly with hydro-
chloric acid, forming ammonium chloride —

NH3 + HC1 = NH4C1.

Carbon monoxide unites directly with chlorine to form carbonyl
chloride —

= COC12.

Carbon monoxide also combines with an additional atom of
oxygen, and gives carbon dioxide, thus —

2CO + O2=2CO2.

In this last action it will be- seen that the molecule of carbon
monoxide, in being converted into the dioxide, takes up one atom
of oxygen ; but as the molecule of oxygen is the smallest isolated
particle, it follows that the two atoms contained in such a molecule
must first separate, and each one then furnishes the requisite
additional oxygen for one molecule of carbon monoxide. In the
union of carbon monoxide with chlorine, and of ammonia with
hydrochloric acid, are we to suppose that the same action takes
place ? That is to say, do the two atoms in the molecule of
chlorine separate from each other and unite with carbon, thereby
satisfying its tetrad valency, in the manner here expressed ? —

Ck
Cl— — C1 + CO = >C = O.

Cl/

And in the case of ammonia and hydrochloric acid, do the
hydrogen and chlorine atoms part., and each unite with the
nitrogen atom, thereby raising it from the trivalent to the penta-
valent condition ? thus —

Cl H

H— ;— C1 + H— N— H = H— N— H.
H H

Valency of the Elements 65

Or are we to suppose that the two molecules, without losing their
integrity, become held together as independent molecules, by
virtue of the unsatisfied affinities of the carbon, or the nitrogen,
as the case may be, in which case the compounds might be repre-
sented thus —

Cl H — Cl

I C = O H — N — H.

Cl |

This question would be settled by determining the vapour-
density of the compound. If, for instance, we were to find the
vapour-density of ammonium chloride to be 26.75, then the com-
pound having the composition NH4C1 would have the normal
molecular volume, that is, its molecule would occupy two unit
volumes,* and the conclusion would be that the vapour consisted
of single molecules of the composition represented by the formula
NH4C1. But ammonium chloride at ordinary temperatures is a
solid, and when heated to the temperature necessary to convert it
into vapour its molecules break up into separated molecules of the
two original gases — ammonia, NH3, and hydrochloric acid, HCl.t
So that we are unable to gain any information in this direction
as to the mode in which the atoms are disposed in the compound.
When the two gases are brought together under ordinary con-
ditions, they combine with the evolution of considerable heat,
owing to loss of energy ; this is taken as evidence that true
chemical action, in the sense of atomic rearrangement, has re-
sulted, hence it is believed that in this compound the nitrogen
is united with the five monovalent atoms, and consequently is
pentavalent.

In the case of carbonyl chloride, COC12, the vapour-density can
be ascertained, this compound existing in the gaseous condition
at the ordinary temperature. Its vapour-density, determined by
experiment, is found to be 50.6. This number, divided into the
molecular weight of the compound having the composition COC12,
gives practically the number 2 as the molecular volume of the
compound. Hence we conclude that these four atoms constitute
a single molecule.

There are a number of combinations, however, in which mole-

* See p. 43. f See Dissociation, p. 89.

66 Introductory Outlines

cules of different compounds unite, that do not so readily admit of
explanation, because in neither of the molecules is there any
atom functioning in a lower state of valency than that which
it is known to be capable of. For example, the monovalent
elements fluorine and hydrogen form the compound hydrofluoric
acid, HF ; fluorine also combines with the monovalent element
potassium, forming potassium fluoride, KF. Both of these com-
pounds come under the head of saturated compounds, in the sense
that neither of them contains an atom which is known to be
capable of exercising a higher valency than it exhibits in these
compounds. Nevertheless these two molecules unite together and
form a definite chemical compound, known as hydrogen-potassium
fluoride.

Again, the divalent element zinc combines with two atoms of
the monad element chlorine, forming zinc chloride, ZnCl.2 ; the
two monovalent elements sodium and chlorine also combine,
giving the compound sodium chloride, NaCl. Both of these
substances must be regarded as saturated compounds, and yet
they unite with each other, forming a distinct chemical compound,
known as sodium zinc chloride. Such compounds as these are
known as double salts, and examples might be multiplied almost
indefinitely. A similar union of molecules, where the recognised
valency of the atoms is all satisfied, is seen in a large number
of compounds containing water of crystallisation ; * for example,
the divalent element copper, in combination with two atoms of
chlorine, forms cupric chloride, CuCl2. The divalent element
oxygen, in combination with two hydrogen atoms, forms water,
H2O. When cupric chloride crystallises from aqueous solution,
each molecule of the chloride unites to itself two molecules of
water, which is therefore termed water of crystallisation.

In chemical notation, it is usual to represent compounds of this
order by placing the formulae of the different molecules that have
entered into union in juxtaposition, with a comma between ;
accordingly,. the examples here quoted would be indicated thus —

Hydrogen potassium fluoride . . HF,KF.
Sodium zinc chloride . .... ZnCl2,NaCl.
Crystallised cupric chloride . . CuCl2,2H2O.

Combinations of this order are by no means confined to the
* See page 216.

Valency of the Elements 67

union of two kinds of molecules, as the following examples will
serve to show : —

Platinum sodium chloride . . PtCl4,2NaCl,6H2O.
Mercuric potassium chloride . 2HgCl2,KCl,2H2O.

At the present time our knowledge of the nature of the union
between these various molecules is too imperfect to admit of any
precise explanation ; such compounds are frequently distinguished
as molecular combinations.

It is quite possible that the unit which has been adopted for estimating
valency, namely, i monovalent atom, is after all only an extremely rough and
crude measure, which is incapable of appreciating smaller differences of com-
bining capacity that may, and most probably do, exist. Its use may be com-
pared to the adoption of a single unit, say i gramme, for the estimation of
mass or weight ; when, if a given quantity of matter has a weight equal to i
gramme, but less than 2 grammes, its weight would be i; if greater than 2
grammes, but less than 3, then its weight would be 2 — a method of estimating
which tacitly assumes that no intermediate weights of matter between the
various multiples of the selected unit are possible. There is no evidence to
show that the combining capacity of an element is exactly expressed by simple
multiples of a monovalent atom.

For example, i hydrogen atom unites with i chlorine atom, that is to say,
with a mass of chlorine weighing 35.5 times its own weight ; and we say that
the mutual affinities of these atoms are satisfied. But for anything we know
to the contrary, an atom of hydrogen may have an affinity for chlorine which
would enable it to unite with a mass of chlorine weighing 40 or 45 or 50 times
its own weight, but not a mass weighing 71 (35.5 x 2) times its own. But since
a mass of chlorine 35. 5 times the weight of a hydrogen atom is the smallest
quantity that is ever known to take part in a chemical change, is the chemically
indivisible mass we call an atom, it follows that as the hydrogen atom has not
sufficient combining capacity to unite with 2 atoms, it is compelled to be
satisfied with i. It might still, however, retain a residual combining capacity.
Or the residual combining capacity may be lodged in the chlorine atom,
which may be conceived as being able to unite with a greater weight of
hydrogen than is represented by i atom, but not so much as that of 2
atoms.

Each of the elements fluorine, chlorine, bromine, and iodine unites with
i atom of hydrogen, and we represent their compounds in a similar manner,
thus —

H - F ; H - Cl ; H - Br , H - I ;

but we make an enormous assumption if we suppose that in each of these
compounds the mutual affinities of the atoms is equally satisfied.

The trend of modern thought, however, lies in the direction of an electrical
interpretation of valency. The fact that atoms are always associated with
fixed and definite charges of electricity, that valency, indeed, could be measured
in terms of electric units (the outcome of Faraday's Law, chap, xi.) seemed

68 Introductory Outlines

at one time only to emphasis* the difficulty of explaining such cases as those
above mentioned ; but the more recent developments in this region of physics
have led to modified views as to the nature of the bond which unites atoms
together. Stated in briefest outline, this chemical " bond" or unit of affinity,
which formerly has been regarded in the light of a single line of force — a
fraction of a bond being considered as altogether inadmissible — is now regarded
as a bundle of lines of force (a Faraday bundle). Under appropriate condi-
tions, such as the proximity of suitable molecules or ions, it is conceived that
some strands of the bundle may become loosened from one of the attached
atoms and thus become available for attraction by similar wandering strands
from other molecules. Obviously, therefore, this view admits of practically an
unbroken gradation in degrees of chemical affinity. Instead, therefore, of
residual affinity, we have varying fractions of the total bundle of lines of
force which in its entirety constitutes the chemical " bond."

A modification of this view, recently advanced by Sir W. Ramsay,* substitutes
electrons f for this " bundle of lines of force. ' ' Atoms are regarded as carrying
with them a " reserve of electrons" electrons which may be inactive, or latent.
Thus taking chlorine as an example, he says: "It appears likely that each
atom of chlorine carries with it no fewer than seven electrons, . . . latent as it
were, not revealing themselves in such a compound as common salt. . . .
These valencies are manifested in such compounds as perchloric acid."

* Presidential address, Chem. Soc., 1909. f See page 104.

CHAPTER IX
GENERAL PROPERTIES OF GASES

UNDER the head of the general properties of gases it will be con-
venient to consider the following subjects : *—

1. The relation of gases to heat.

2. The relation of gases to pressure.

3. The liquefaction of gases.

4. Diffusion of gases.

5. The kinetic theory of gases.

The Relation of Gases to Heat.— The fact that substances
expand when heated, and again contract upon being cooled, was
observed in very early times. The fact also that all substances do
not undergo the same alterations in volume when subjected to the
same changes of temperature has been long known ; but it was not
until the beginning of the nineteenth century that it was proved by
Charles and Gay-Lussac that all gases expanded and contracted
equally when exposed to the same alterations of temperature.
This law is generally known as the Law of Charles, and may be
thus stated : When a gas is heated, the pressure being constant, it
increases in vohime to the same extent whatever the gas may be.

The increase in bulk suffered by I volume of a gas in being
heated from o° to i° is termed the coefficient of expansion, and if
the law of Charles is true all gases will have the same coefficient.

Modern research has shown that the law of Charles is not abso-
lutely true, and the extent to which gases deviate from the strict
expression will be seen from the coefficients of expansion given in
the following table : —

* The study of these subjects belongs more especially to the science of
physics or chemico-physics. For fuller information on these points than can
be included within the scope of this book students are referred to special

treatises on physics.

£6 Introductory Outlines

Air . 003665*1

Hydrogen 003667!

Carbon monoxide . .003667 j

Nitrogen . . ' . ,, . . . .003668]

Nitrous oxide . .003676

Carbon dioxide ..... .003688

Cyanogen. . . . . .003829

Sulphur dioxide . . . . .003845

It will be noticed that the first four gases have almost the same
coefficient of expansion : these gases are all very difficult of lique-
faction, and it will be seen that the coefficient rapidly rises in the
case of the other gases, which are easily liquefied.

For purposes of ordinary calculation it is usual to adopt the
coefficient of expansion of air as applicable to all gases. It will
be obvious that since the volume of a gas is affected by alterations
of temperature, it becomes necessary, when measuring the volume
of a gas, to have regard to the particular temperature at which the
measurement is made, and in order to compare volumetric measures
they must be all referred to some standard temperature. This
standard temperature is by general consent o° C.

Taking the fraction .003665, therefore, for the coefficient—

I volume of a gas at o° becomes i + .003665 volumes at i*

I „ „ o° „ i + .003665 x 2 „ 2°

or I „ „ o° „ i + .003665 / „ f

Therefore the volume at /° equals the volume at o° multiplied by
i + .003665 /. Let v be the volume at /°, and v0 the volume at o°,
then—

v = v0(i + .003665 /),

and conversely the volume at o° equals the volume at /e divided by
I + .003665 /—

v
~ i + .003665 /

The vulgar fraction equivalent to .003665 is ^. 273 volumes
at o° become 273 + / at /°.

What is known as the absolute temperature of a substance is the
number of degrees above — 273° C. Taking this point as the zero,
the absolute temperature of melting' ice, for example, will be 273°.
Charles' law, therefore, may be thus stated : The volume of an)

Relation of Gases to Pressure J\

gas. under constant pressure, is proportional to the absolute tern-
perature.

The Relation of Gases to Pressure.— The effect of increase
of pressure upon a gas is to diminish its volume. The law which
connects the volume occupied by a gas, with the pressure to which it
is subjected, was discovered by Robert Boyle (1661), and is known
as Boyle's Law. It may be thus stated : The volume occupied by
a given weight of any gas is inversely as the pressure. The
general truth of this law may readily be illustrated by subjecting a
gas to varying pressures, and it will be seen that when the pressure
is doubled the volume of gas is reduced to one-half, and so on.

Just as in the case of the law of Charles, modern investigations
have shown that the law of Boyle is not a mathematical truth. It
is found not to be absolutely true of any gas, for, with the exception
of hydrogen, all gases are more compressible than is demanded by
the law. Hydrogen deviates from the law in an opposite sense, in
that it requires a higher pressure than the law would indicate, in
order to reduce a volume of it to a given point. These deviations
from Boyle's law are explained by the operation of two causes ;
first, the attraction exerted by gaseous particles upon each other ;
second, the fact that increased pressure diminishes the space
between the molecules, and not the actual space occupied by the
molecules of a gas. When the former cause predominates, the
gas deviates from the law by being more compressible ; in the case
of hydrogen the second cause operates more powerfully. (See
Kinetic Theory of Gases.) For ordinary purposes of calculation
the law of Boyle may be regarded as true.

As the volume of a given weight of gas is so intimately related
to the pressure, and as the atmospheric pressure is variable, it
becomes necessary, in all quantitative manipulation with gases, to
know the actual pressure under which the gas is at the time of
measurement, and to refer the volume to a standard pressure.
The pressure that has been adopted as the standard is that of a
column of mercury 760 mm. in height. (See Atmosphere.)

If ?/ equals the volume of gas measured at / pressure, and v0
the volume at the standard pressure, then

In practice it is most usual tc make both correction for tempe

Introductory Outlines

rature and pressure together ; then i>0 being the volume at the
standard temperature and pressure, we get

v
i +'.003665 t

_£*
760'

The Liquefaction of Gases.— Under certain conditions of tem-
perature and pressure, the law of Charles and the law of Boyle both
completely break down. According to
the law of Charles, 100 c.c. of a gas at
o° C. should occupy 96.4 c.c. if the tem-
perature were lowered to — 10°. If 100
c.c. of the gas sulphur dioxide at o° C.
be confined in a glass tube standing in
mercury, and the gas be cooled to — 10°
by surrounding the tube with a freezing
mixture^ it will be found that the volume
of gas, instead of occupying 96.4 c.c.,
has been reduced to a few cubic centi-
metres only, and that the surface of the
mercury in the tube is wet owing to the
presence of a minute layer of a colourless
liquid upon it. In this case the law of
Charles has broken down, and the sul-
phur dioxide has passed from the gaseous
^frV to the liquid state.

~^' rO Similarly, according to the law of
Boyle, 100 c.c. of a gas measured at the
standard pressure should occupy 25 c.c.
when exposed to a pressure of four additional atmospheres. If
loo c.c. of the gas sulphur dioxide be enclosed in one limb of a long
U-tube, as shown in Fig. i, the other limb being filled with air,
and the two gases be simultaneously exposed to increased pressure
by raising the mercury reservoir, it will be seen that at first the
gases in both tubes are compressed equally. As the pressure
approaches three atmospheres, however, the mercury will be seen

* The student should familiarise himself with the method of calculating the
changes of volume suffered by gases, by changes of temperature and pressure,
by working out a number of examples such as the following : —

1. If 30 litres of gas are cooled from 25° to o°, what is the diminution in
volume, the pressure being constant? Ans. 2.51 litres.

2. If a litre of air at o° weighs 1.293 grammes when the barometer is at

FIG. i.

Liquefaction of Gases 73

to rise much more rapidly in the tube containing the sulphur
dioxide, and when the mercury reservoir has been raised to such a
height that the gases are subjected to four atmospheres, the sulphur
dioxide will have completely broken down, and will be entirely con-
verted into a few drops of liquid, which appear upon the surface of
the mercury. The air meantime, in the other limb, will be found to
occupy 25 c.c., as that gas at that pressure obeys Boyle's law almost
absolutely. We see, therefore, that at a certain temperature and at
a certain pressure the gas sulphur dioxide begins rapidly to depart
from the laws of Charles and Boyle, and ultimately passes into the
liquid condition.

All gases, when exposed to certain conditions of temperature
and pressure, conditions which are special for each different
gas, will pass from the gaseous to the liquid state ; and
as the point at which liquefaction takes place is approached,
the departures from Boyle's law become more and more pro-
nounced.

The first substance, recognised as being under ordinary condi-
tions a true gas, that was transformed into the liquid condition
was chlorine, which was liquefied in the year 1806 by Northmore.
The true nature of this liquid was
not understood until Faraday inves-
tigated the subject.

In his earlier experiments Fara-
day's method consisted in sealing
into a bent glass tube (Fig. 2) sub-
stances which, when heated, would
yield the gas ; the substances being
contained in one limb of the tube,
and the empty limb being immersed FIG. 2.

in ice. The pressure exerted by the gas thus generated in a con-
fined space was sufficient to cause a portion of it to condense to

760 mm., what will be the weight of a litre of air at 27°, the barometer
standing at the same height? Ans. 1.177 grammes.

3. What will be the weight of a litre of air at 42° when the barometer stands
at 735 mm, ? Ans. 1.084 grammes.

4. Air at a temperature of 15° is enclosed in a vessel and heated to 93°.
Compare the pressure of the enclosed air with that of the atmosphere. Ans.
As 61 : 48.

5. What will be the volume, at the standard temperature and pressure, of
500 c.c, of hydrogen, measured at 20°, and under a pressure of 800 mm.?
Ans. 490 c.c.

Jfy. Introductory Outlines

the liquid state, and the liquid collected in the cooled limb. Irt
this way Faraday liquefied such gases as chlorine, sulphur dioxide,
ammonia, cyanogen. In his later experiments Faraday compressed
the gas by means of a small compression pump, and at the same
time applied a low degree of cold, and by so doing he succeeded
in liquefying carbon dioxide, hydrochloric acid, nitrous oxide, and
other gases. There were a number of gases, however, which Fara-
day found it impossible to liquefy, such as hydrogen, oxygen, nitro-
gen, marsh gas, nitric oxide, carbon monoxide, &c. It became the
custom to call these permanent gases, and this term was applied to
them until the year 1877.

In that year it was proved by Pictet, and independently by Cail-
letet, that under sufficiently strong pressure, and a sufficiently low
degree of cold, the so-called permanent gases could in the same
way be reduced to the liquid condition. Pictet's method was in
principle the same as that employed by Faraday, the difference
being that with the machinery at his disposal he was able to
employ enormously increased pressure and a greater degree of
cold. For the liquefaction of oxygen, a quantity of potassium
chlorate was heated in a strong wrought-iron retort, to which was
connected a long horizontal copper tube of great strength and small
bore. At the extreme end of this tube there was a pressure gauge
capable of indicating pressures up to 800 atmospheres, and a stop-
cock. The tube was cooled by being contained in a wider tube,
through which a constant stream of liquid carbon dioxide, at a tem-
perature of — 120° to — 140°, was caused to flow.

The machinery employed to maintain this flow of liquefied car-
bon dioxide was somewhat elaborate, consisting of condensing and
exhaust pumps for liquefying and rapidly evaporating sulphur
dioxide, and similar condensing and exhaust pumps for liquefying
and rapidly evaporating carbon dioxide : the sulphur dioxide being
merely the refrigerating agent used to assist the liquefaction of
the carbon dioxide. This machinery was driven by two eight-
horse-power engines. As the potassium chlorate was heated
and oxygen evolved, the internal pressure in the retort and
copper tube rapidly rose, and its amount was indicated by the
gauge.

When the stop-cock upon the end of the tube was opened, liquid
oxygen was forcibly driven out in the form of a jet.

In the method employed by Cailletet, the pressure to which the
gas is subjected is obtained by purely mechanical meaas. The

Liquefaction of Gases

gas to be liquefied is introduced into a glass tube (Fig. 3), the
narrow end of which consists of a strong capillary tube. The tube
carries a metal collar, which enables it to be secured in position
in the strong steel bottle (Fig. 4), by means of a nut, E' (Fig. 5),
which screws into the mouth. The bottle, which is partially filled
with mercury, is connected by means of a flexible copper tube of
fine bore with a small hydraulic pump, by means of which water
is forced into the steel bottle. The water so driven in forces the

FIG. 3.

FIG. 4.

FIG. 5.

mercury up into the glass tube T, and thereby compresses the
contained gas. In this way a pressure of several hundred atmos-
pheres may be applied to the gas. In his earlier experiments
Cailletet depended almost entirely for the refrigeration he required
upon the fact, that when a gas is allowed suddenly to expand it
undergoes a great reduction in temperature. This method of
cooling may be termed internal refrigeration. In the case of
cxygen, the gas was first subjected to a pressure of 300 to 400

»5 introductory Outlines

atmospheres, and was then allowed suddenly to expand by a rapid
release of the pressure. The result of the sudden expansion was
to momentarily lower the temperature of the gas to such a point
that the tube was filled with a fog, or mist, consisting of liquid
particles of oxygen.

This principle, namely, the self-cooling of a gas by its own
sudden expansion, has recently been applied for the liquefaction
of oxygen in large quantities. When oxygen under considerable
pressure, say 120 atmospheres, i« allowed to escape from a fine
orifice at the end of a long pipe, the issuing gas suddenly expands,
and thereby its temperature is greatly lowered. If this self-cooled

gas is made to
sweep over the pipe
from which it is
escaping, it will
cool the pipe, and
therefore lower the
temperature of the
remaining gas be-
fore it issues. In
this way the cooling
effect becomes cu-
mulative, the initial
temperature of the
gas before it es-
capes being con-
tinually brought
lower and lower,
until at last the
point is reached at
which the oxygen
is liquefied.*

If the oxygen be
first cooled to about
— 80° by means of
solid carbon di-
oxide, then in a few

minutes, by the further cooling due to its own expansion, the tem-
perature will fall below the boiling-point of oxygen, and the
liquefied gas be obtained.

The apparatus for the purpose is shown in Fig. 6.t Oxygen

* Linde, The Engineer ; Oct. 4, 1895.

f Designed bv Dewar ; constructed by Messrs. Lennox, Reynolds & Fyfe.

FIG. 6.

Liquefaction of Gases 77

under a pressure of 120 to 140 atmospheres is passed through a
series of spirals of fine copper pipe contained in the chamber C,
which is encased in a non-conducting jacket of cork-dust. The
gas enters by the pipe O (seen in the enlarged section), and passes
through the spiral S S, which is immersed in a mixture of alcohol
and solid carbon dioxide (the liquid carbon dioxide from the reservoir
being admitted into the alcohol through the valve W, which is regu-
lated by the screw B). The oxygen thus cooled passes through the
double spiral pipe D D, which ultimately extends through the
bottom of the chamber, and terminates in a stirrup, U, the short
end of which is closed. In the bend of this stirrup there is a fine
hole, which can be closed or opened at will by the pointed end
of the rod V, connected to the screw A. On opening this valve,
the oxygen, already cooled to about — 80°, escapes from the hole
under a pressure of 120 to 140 atmospheres. It instantly expands,
and is thereby cooled still lower. This cold gas is prevented from
escaping at once into the atmosphere by the glass tube G, but is
compelled to rush upwards (as shown by the arrows), and, sweep-
ing past the double spiral D D, cools this pipe, and therefore the
succeeding portions of issuing oxygen. In a few minutes the tem-
perature of this pipe is thereby brought so low, that the further,
cooling of the gas by its expansion causes the liquefaction of a
portion of it, and a fine spray of liquid is seen to spurt out from
the hole. This spray quickly increases in quantity, and rapidly
collects as a clear liquid in the glass tube G. This tube is double-
walled, the space between the walls being perfectly vacuous. In
such a vessel the liquid oxygen may be kept for a considerable
time, evaporating only very slowly in spite of its extremely low
boiling-point, as it has been found that such a vacuous envelope
forms the most perfect non-conductor.

The instruments designed by Linde in Germany, and by Hamp-
son in England, and known as air-liquefierS) are constructed on
precisely similar principles. In this case, however, the preliminary
cooling by means of solid carbon dioxide is dispensed with ; for
instead of a limited and comparatively small supply of com-
pressed gas in a steel cylinder, an unlimited supply of air is
delivered into the machine, under a pressure of 120 to 160
atmospheres, by means of powerful compression pumps driven
by an engine.

By an extension of the same principles hydrogen was first suc-
cessfully liquefied in 1898. In this case, however, the gas requires

78 Introductory Outlines

to be previously cooled to about — 200° before expansion is allowed to
take place. By utilising the low temperatures which can be obtained
by means of boiling liquefied gases, it has now become possible
to liquefy all the known gases by cold alone, that is, without the
application of pressure ; in other words, their temperatures can
be brought down below their boiling-points, under which circum-
stances they must obviously assume the liquid state. For example,
liquefied ethylene boils at — 103.5° ; if, therefore, a stream of nitrous
oxide is passed through a tube immersed in a bath of liquid ethy-
lene, the nitrous oxide will be cooled below its boiling-point ( - 89.80)>
and will consequently be reduced at once to the liquid state.
Again, liquid oxygen boils at —182.5°.
This boiling liquid therefore is sufficiently
cold to cool marsh gas below its boiling-
point, namely, — 164.7°, and therefore to
cause its liquefaction.

,v Moreover, by the rapid evaporation of liquid
oxygen the temperature may readily be
lowered to the point at which air will liquefy.
Thus, if a quantity of liquid oxygen in the
glass tube O (Fig. 7), which is provided with
a vacuous envelope, V, be made to boil
rapidly by putting the pipe P in connection
with an exhaust - pump, the temperature
q'uickly falls to —200°, when air itself be-
comes liquefied without the application of
pressure ; and drops of liquid air quickly
collect upon the walls of the inner empty tube,
N, which is freely open to the atmosphere.
,In this way considerable quantities of lique-
fied air can be collected in a few minutes.
By means of boiling liquid hydrogen the low temperature of
— 253° has been reached, at which temperature all other known
gases, except helium, are frozen to the solid state. The lowest
temperature yet obtained by the rapid evaporation of solid
hydrogen is -260° (Dewar).

The Critical Point— As far back as the year 1869, it was
shown by Andrews that when liquid carbon dioxide was heated
to a particular temperature, it passed from the liquid to the gaseous
state, and that no additional pressure was able to condense it again
so long as the temperature remained at or above that point. This

..0

.— N

FIG. 7.

Critical Temperature of Gases 79

particular temperature is called the critical point, or the critical
temperature of the gas. In the case of carbon dioxide this critical
temperature is 31.35°, and in order that this gas may be liquefied by
pressure, it is an essential condition that the temperature be below
that point ; above 32° no pressure is capable of bringing about
liquefaction. All gases have a critical temperature, which is special
for each gas, and until the temperature of the gas be lowered to
that point, liquefaction is impossible. The critical temperatures
of the different gases vary through a very wide range : thus,
the critical temperature of hydrogen is as low as — 238°, while that
of sulphur dioxide is + 155.4°. In the third column of the table
of physical constants on page 80 the critical temperatures of a
number of the more common gases are given.*

The gases in this list, from ethylene downwards, all have their
critical temperatures so high that there is no difficulty in cooling
them below these points. These are the gases which were first
reduced to the liquid state. The first five upon the list have
very low critical temperatures ; these are the very gases which for
so long resisted all attempts to liquefy them, and which were on
that account called permanent gases. We now know that the
failure to obtain them in the liquid state was owing to the fact
that the relation between the critical temperature and the point
of liquefaction was not fully realised. Just as carbon dioxide
cannot be liquefied unless its temperature be brought down to
31.35°, so oxygen resists liquefaction under the highest possible
pressures, until its temperature be lowered to — 118.8°, the critical
temperature of oxygen.

The critical temperature of a gas is sometimes spoken of as the
absolute boiling-point.

Critical Pressure.— The particular pressure that is required
to liquefy a gas at its critical temperature is called the critical
pressure. Thus the pressure necessary to liquefy oxygen, when
the temperature has been lowered to - 118.8°, is 58 atmospheres ;
while that required to condense chlorine at its critical point, viz.,
+ 141°, is 84 atmospheres. At temperatures below the critical
temperatures a gas liquefies under less pressure than the critical

* For the constants for the gases of the Argon family see page 271. It may
be well to remind the student that such constants as are here tabulated
are obtained from measurements involving very great experimental difficulties,
and that consequently they are always liable to revision. The values here
given are from the most recent determinations.

Introductory Outlines

pressure, until when the temperature is lowered to the boiling-
point of the gas it passes into the liquid state without the applica-
tion of any external pressure. The following table contains the
most recently determined physical constams of a number of
common gases : —

TABLE OF PHYSICAL CONSTANTS.

Boiling-
Point.

Melting-
Point.

Critical
Temp.

Critical
Pressure.

Density
at Boiling-
Point.

Hydrogen . .

-253°

-257°

-238°

15.3 Ats.

0.06

Nitrogen .

~I95-S

-213°

-149°

27-5

0.791

Carbon monoxide

-190°

-207°

-I36°

33-5

Oxygen

-182.5°

-223°

-1x8.8°

58.0

I.I3I

Methane (marsh gas

-164.7°

-184°

- 82°

55-8

0.416

Ethylene .

-103-5°

-169°

+ 9°

58.0

0.571

Nitrous oxide .

- 89-8°

— 102. 7°

+ 37°

. .

Acetylene .

- 827°

+ 35°

Carbon dioxide

- 80°

+ 3J-35°

72.3

Ammonia .

- 38.5°

- 75-5°

+ 131°

i*3

Chlorine . . ,

- 33-4°

4-141°

I-5°7

Sulphur dioxida

- 10°

-ri55-4°

78.9

From the figures given in thi? table it will be seen that the
critical pressure (which is the pressure required to liquefy a gas at
the highest temperature at which pressure can possibly cause lique*
faction) is in most cases comparatively small. In only one instance,
namely, ammonia, is it over 100 atmospheres, and falling in the
case of hydrogen as low as 15.3 atmospheres. The enormous
pressures, therefore, amounting often to many hundred atmospheres,
which some of the earlier experimenters employed in attempting to
effect the liquefaction of the so-called permanent gases, are thus
seen to have been efforts in an entirely wrong direction. It was not
greater pressure that was required, but the means of cooling the
gases to a sufficiently low temperature.

In ordinary language such a gas as chlorine is spoken of as an
easily liquefied gas, while oxygen would be described as a difficultly
liquefied gas. Strictly speaking, however, and considering them
from a comparable standpoint, it would perhaps be more correct
to regard them in exactly the opposite light. Thus, taken at their
respective critical temperatures, oxygen is liquefied by a pressure
of 58 atmospheres ; while at the critical temperature of chlorine this

Diffusion of Gases 81

gas requires a pressure of 84 atmospheres to reduce it to the liquid
state. At o° it is true chlorine may be liquefied by a pressure of
only 6 atmospheres, but it must be remembered that o° is 141
degrees below the critical temperature of this gas. Long before
oxygen has been cooled 141 degrees below its critical temperature,
which would be down to -254°, it not only passes into the liquid
state without the application of any external pressure at all, but is
frozen to the solid state.

Diffusion of Gases.— If a jar filled with hydrogen be placed
mouth to mouth with a jar of air, the hydrogen being uppermost,
it will be found that after the lapse of a few minutes some of the
hydrogen will have passed into the bottom jar containing air, and
some of the air will have made its way up into the hydrogen jar.
The light gas hydrogen does not, as might have been supposed,
remain floating upon the air, which is 14.44 times as heavy, but
gradually escapes into the lower jar ; and the heavier gas finds its
way, in opposition to gravitation, into the upper jar. This process
goes on until there is a uniform mixture of air and hydrogen in both
jars, and the gases never separate again according to their
densities.

This transmigration of gases will take place even through tubes
of considerable length : thus, if two soda-water bottles be filled one
with hydrogen and the other with oxygen, and the two bottles be
connected by a piece of glass tube a metre in length, the system
being held in a vertical position with the light hydrogen upper-
most, it will be found after an hour or two that the two gases
have become mixed. Some of the hydrogen will have descended
through the long tube into the lower bottle, and in like manner
a portion of the oxygen, although nearly sixteen times as heavy
as hydrogen, will have travelled up into the top bottle. That the
gases have so mixed may be readily shown by applying a lighted
taper to the mouth of each bottle, the detonation which then takes
place proving that the bottles contain a mixture of oxygen and
hydrogen. This passage of one gas into another is called the
diffusion of gases. It was observed by Graham that when the
two gases were separated from each other by a thin porous
septum, such, for instance, as a piece of unglazed porcelain (so-
called " biscuit "), or plaster of Paris, the pressure of the gas on
the two sides of the porous partition did not remain the same
during the process of diffusion : that is to say, one gas made its
way through the partition faster than the other, and it was noticed

Introductory Outlines

that the lighter the gas the more rapidly was it able to transpire
or diffuse through the porous medium. This fact, viz., that a light
gas diffuses more rapidly than a heavier one, may be observed
in a variety of ways.* The apparatus seen in Fig. 8 is a modified
form of Graham's diffusiometer. It consists of a long glass tube
with an enlargement or bulb near to one end. Into the short neck
of this bulb there is fastened a thin diaphragm of stucco, or other
porous material. If the apparatus be filled with hydrogen by dis-
placement, the short neck being closed by a cork, and the long
limb be immersed in water, it will be seen, upon the withdrawal

FIG. 8.

FIG. 9.

of the cork, that the water rapidly rises in the long tube. The
hydrogen diffusing out through the diaphragm so much more
rapidly than air can make its way in, a diminution in pressure
within the apparatus results, and this causes the water to ascend
in the tube. The same phenomenon may be seen even more
strikingly by means of the apparatus, Fig. 9, which consists of a
tall glass U-tube, upon the end of one limb of which there is
fastened, by means of a cork, a porous cylindrical pot, such as

* See Experiments Nos. 350-359, Newth's "Chemical Lecture Experiments,"
newed.

Diffusion of Gases §3

is used in ah ordinary Bunsen battery. The U-tube is half
filled with coloured water. Under ordinary circumstances air is
continually diffusing through the porous pot, but as it passes at
an equal rate in both directions, there is no disturbance of the
pressure, and consequently the coloured water remains level in
the two limbs. If now a beaker containing hydrogen be brought
over the apparatus, as seen in the figure, the hydrogen will stream
through the porous pot so much more rapidly than the air in the
pot can make its way out, that there will be an increase in the
total amount of gas inside the apparatus, which will be instantly
rendered evident by the change of level of the liquid in the U-tube,
the water being forcibly driven down the tube which carries the
porous pot. Upon removing the beaker the reverse operation
will at once take place ; the hydrogen inside the apparatus now
rapidly diffuses out, and much more quickly than air can pass in,
consequently a reduction of pressure within the apparatus results,
which is indicated by a disturbance of the level of the water
in the tube, in the opposite direction to that which occurred at
first.

The Law Of Gaseous Diffusion.— Graham established the law
according to which the diffusion of gases is regulated, and it may
be thus stated : The relative velocities of diffusion of any two
gases are inversely as the square roots of their densities.

The density of hydiogen being i, that of air is 14.44, the velocity
of the diffusion of hydrogen, therefore, as compared with that of
air, will be in the ratio of ^14.44 to Ji. VJ4-44 = 3-8, *Ji = i.
Therefore hydrogen diffuses 3.8 times faster than air ; or 3.8 yolumes
of hydrogen will pass out through a porous septum, while only i
volume of air can enter.

\id= the density of a gas, air being unity, and v = the volume
of the gas which diffuses in the same time as I volume of air, then

The following table gives in the last column the results obtained
by Graham, which will be seen to accord very closely with the cal-
culated numbers demanded by the law of diffusion :—

Introductory Outlines

Volume of Gas

Name of Gas.

Density of Gas
compared with

which Diffused in
the same Time as

Air = rf.

one Volume of

Air.

Hydrogen
Marsh gas

0.06926

0-559

3-7794
1-3375

3-83
1-344

Carbon monoxide

0.9678

1.0165

1.0149

Nitrogen . . •
Oxygen .
Sulphuretted hydrogen
Carbon dioxide

0.9713
1.1056
1.1912
1.5290

1.0147
0.9510
0.9162
0.8087

1.0143
0.9487

o-95
0.812

Sulphur dioxide .

2.247

0.6671

0.68

The property of diffusion is sometimes made use of in order to
separate gases, having different densities, from gaseous mixtures.
This process of separation by diffusion is known as atmolysis.
The principle may readily be illustrated by causing a mixture of

oxygen and hydrogen, in
proportion to form an ex-
plosive mixture, to slowly
traverse tubes made of
porous material, such as
ordinary tobacco pipes.
Two such pipes may be
arranged as shown in Fig.
10, and the gaseous mix-
ture passed through in the
direction indicated by the
arrow. On collecting the

* . , * •

issuing gas over water in a
pneumatic trough, it will
be found to have so far
lost the hydrogen, by dif-
fusion through the tube,
that a glowing splint of
wood when introduced into it will be reignited.

From the rate of diffusion of ozone, in a mixture of ozone and
oxygen, Soret was able to calculate the density of this allotropic
form of oxygen, and so confirm the result he had previously ob-
tained by other methods (see Ozone).

Attempts have been made to utilise this principle in order to
obtain oxygen from the air. The relative densities of oxygen and

FIG. 10.

The Kinetic Theory 85

nitrogen are as 16 to 14 ; the rate of diffusion, therefore, of nitrogen
is slightly greater than that of oxygen.

Effusion is the term applied by Graham to the passage of gases
through a fine opening in a very thin wall, and he found that it
followed the same law as diffusion. Bunsen utilised this principle
for determining the density, and therefore the molecular weights,
of certain gases. The method, in essence, is as follows :— A
straight glass eudiometer is so constructed, that a gas contained
in it can be put into communication with the outer air through a
minute pin-hole in a thin platinum plate. The gas is confined in
the tube, which is placed in a cylindrical mercury trough, by
means of a stop-cock at the top. When the tube is depressed
in the mercury, and the cock opened, the gas escapes through
the minute perforation in the platinum plate, and its rate of effu-
sion is determined by the time occupied by a glass float placed
in the tube in rising a graduated distance within the eudiometer.

The flow of gases through capillary tubes is called transpiration
of gases. In this case the friction between the gas and the tubes
becomes a factor in the movement, so that this phenomenon is
not governed by the same law as gaseous diffusion.

The Kinetic Theory of Gases.— The term kinetic signifies
motion, and as applied to this theory it expresses the modern
views of physicists concerning matter in the gaseous state, and
serves to harmonise and explain the physical laws relating to
the properties of gases. Matter in the state of gas or vapour
is regarded as an aggregation of molecules in which the attractive
forces which tend to hold them together are reduced to a minimum,
and in which the spaces that separate them are at a maximum.
These molecules are in a state of rapid motion, each one moving
in a straight line until it strikes some other molecule, or rebounds
from the walls of the containing vessel, when it continues its move-
ment in another direction until it is once more diverted by another
encounter. As they constantly encounter and rebound from each
other, it will be evident that at any given instant some will be
moving with a greater speed than others ; the majority, however,
will have an average velocity. In these encounters no loss of
energy results so long as the temperature remains constant, but
any change of temperature results in a change in the velocity of
movement of the molecules, the speed being increased with
increased heat. The actual volume of the molecules is very small
as compared with the space occupied by the mass ; the space

86 Introductory Outlines

between the molecules, therefore, in which they pass to and fro,
is relatively very great. As the molecules are constantly colliding
and rebounding, the distances between them, as well as their speed,
will be sometimes greater and sometimes less ; but there will be
an average distance, which is known as the mean free path of the
molecule.

The pressure exerted by a gas, or its elastic force, is the combined
effect of the bombardment of its molecules against the containing
vessel ; in other words, the pressure of a gas is proportional to the
sum of the products obtained by multiplying the mass of each
molecule by half the square of its velocity. It will be obvious
that if the space within which a given mass of gas is confined be
reduced, the number of impacts of the molecules against the walls
of the containing vessel, in a given time, will be increased, and
therefore the pressure it exerts, or its elastic force, will also be
increased. If the space be reduced to one-half the original, the
number of these impacts will be doubled, or in other words, the
number of impacts in a given time is inversely as the volume.
This statement is simply the law of Boyle stated in the language
of the kinetic theory.

When a given mass of gas contained in a confined space is
heated, the pressure it exerts, or its elastic force, is increased. But
as the number of molecules present has not been increased by
raising the temperature of the gas (provided no chemical decom-
position of the gas is brought about by the change of temperature),
the increased pressure can only have resulted from the greater
frequency, and greater energy, of the impacts of the molecules
against the walls of the vessel, owing to their greater velocity.

Two equal volumes of different gases under the same conditions
of temperature and pressure, exert the same elastic force upon the
containing vessels, that is to say, the kinetic energy in each volume
is the same. According to Avogadro's hypothesis, equal volumes
of all gases under the same conditions of temperature and pressure
contain an equal number of molecules, however much the weight
of these molecules may vary ; therefore the average kinetic energy
of each individual molecule will be the same. It follows from this
that the mean velocities of different molecules must vary, and the
calculated numbers representing the actual velocities of movement
of the molecules of different gases show that these rates are pro-
portional to the inverse square roots of their respective densities.
Put according- to the law of gaseous diffusion (Graham's, law), the

The Kinetic Theory 87

relative rapidity of diffusion of gases is inversely proportional to
the square roots of their densities, hence by purely mathematical
processes, based upon the kinetic theory of gases, the law of
gaseous diffusion is proved to be true. Similarly, the kinetic theory
is applicable to the consideration of the phenomena of evaporation
and condensation (see page 126), and to the processes of solution
(page 148).

The deviations from the laws of Boyle and Charles, already
referred to,* are also explained by the dynamical theory of gases,
from considerations of the following order : —

1. That the molecules themselves are not mathematical points,
but occupy a space ; in other words, the space occupied by the
actual particles of matter is not infinitely small as compared with
the entire volume of the gas, i.e. the bulk of the particle plus the
intermolecular spaces.

While the pressure upon a gas is only slight, and therefore the
total volume occupied by a given mass of the gas is great, the bulk
of the actual particles themselves becomes a vanishing quantity in
comparison with the total volume (i.e. the space occupied by
particles, plus the intermolecular spaces), and the gas under these
circumstances tends to approach more nearly to the conditions of
an ideal gas. But when the pressure is increased, and the total
volume thereby greatly reduced, then the bulk of the particles
themselves begins to bear an appreciable proportion to the total
volume cccupied by the gas.

2. That the impact of the molecules against, each other and
against Ihe containing envelope occupies time ; or, in other words,
the time occupied by the impacts is not infinitely small compared
with the time elapsing between the impacts.

3. That the molecules themselves are not entirely without attrac-
tion for each other ; that is to say, although the attractive force
between the molecules which holds them together in the liquid
and solid states of matter is at a minimum in the case of gases, it
is not entirely absent.

* See page yi.

CHAPTER X

DISSOCIATION— REVERSIBLE OR BALANCED
ACTIONS

DISSOCIATION is the term employed to denote a special class of
chemical decompositions. When potassium chlorate is heated it
breaks up into potassium chloride and oxygen, thus —

2KC1O3 = 2KC1 + 302,

and when calcium carbonate (chalk) is heated it breaks up into
calcium oxide (lime) and carbon dioxide —

CaCO3 = CaO + CO2.

In the first case the oxygen is incapable of reuniting with the
potassium chloride, but in the second, the carbon dioxide can
recombine with the lime and reproduce calcium carbonate : there-
fore both the following expressions are possible —

"CaCO3 = CaO + CO2,
and

CaO + CO2 = CaCO3.

Reactions of this order are known as reversible or balanced actions,
and the breaking up of calcium carbonate by the action of heat is
termed dissociation, while that of the potassium chloride under
similar circumstances is simple decomposition.

When ammonia is passed through a tube heated to a dull red
heat, the gas is decomposed into nitrogen and hydrogen —

2NH3 = N2 + 3H2,

and the two gases pass out of the heated tube as separated gases,
and do not recombine again.*

But when steam is strongly heated it is dissociated into oxygen

* Nitrogen and hydrogen can be caused to unite under suitable conditions
(see Ammonia).

Dissociation 89

and hydrogen, and as these separated gases pass away from the
heated region they reunite, forming molecules of water vapour.
Such a reversible reaction may be thus expressed

' 2H20 ^ 2H2 + 02.

Again, when the gases ammonia and hydrochloric acid are brought
together at the ordinary temperature, they unite to form solid
ammonium chloride, and when ammonium chloride is heated it
dissociates into its two generators,* hence we have the expression—

NH3 + HC1 ^ NH4C1.

The corresponding compound containing phosphorus in the place
of nitrogen dissociates at a temperature as low as - 20°, hence
when hydrogen phosphide and hydrochloric acid are mixed at
ordinary temperatures no combination takes place, the separate
molecules are in the same relation to one another as those of
ammonia and hydrochloric acid at a high temperature. When,
however, the mixture of gases is cooled below - 20°, union takes
place and crystals of phosphonium chloride are formed, which at
once begin to dissociate into the original gases as the temperature
again rises. The change, as before, may be represented as a
reversible one —

PH3 + HC1 ;± PH4C1.

In such cases of dissociation as that of calcium carbonate, where
one of the products is gaseous and the other solid, no difficulty
exists in separating the simpler compounds that result from the
decomposition ; but where the products are entirely gaseous, special
methods have to be adopted to withdraw the one from the other,
while they still exist as separate molecules, and before they reunite
again. One such method, which is well adapted for the quali-
tative illustration of dissociation, is based on the law of gaseous
diffusion. If when ammonium chloride is heated it is dissociated
into ammonia, NH3, and hydrochloric acid, HC1, these two gases,
having the relative densities of 8.5 and 18.25, will diffuse through
a porous medium at very different rates. According to the law of
diffusion, these rates will be inversely as the square roots of the
densities of the gases ; if, therefore, the conditions are so arranged

* Baker has shown (May 1894) that when absolutely dry, these gases do not
combine ; and also, that when aqueous vapour is entirely absent, ammonium
chloride does not undergo this dissociatipn.

go Introductory Outlines

that the heating of the ammonium chloride takes place in the
neighbourhood of a porous diaphragm, more of the light ammonia
gas will diffuse through in a given time than of the heavier hydro-
chloric acid, so that a partial separation of these gases will be
effected. Fig. 1 1 shows a convenient arrangement for carrying out
the experiment. A fragment of ammonium chloride is heated in a
short glass tube through which passes the stem of an ordinary clay
tobacco pipe. As the dissociation takes place, both of the gaseous
products begin to diffuse into the interior of the porous clay pipe,
but owing to their greater rate of diffusion, a larger number of am-
monia molecules will pass in, than of hydrochloric acid, in the same
time ; consequently, when the gases pass away from the heated
region and once more recombine, there will be a surplus of am-
monia molecules within the porous pipe, and for the same reason
an excess of hydro.chloric acid molecules outside. If the gaseous
contents of the porous tube be driven out by means of a stream of

FIG. n.

air from an ordinary bellows, the presence of the free ammonia may
be recognised by allowing the air to impinge upon a piece of paper,
coloured yellow with turmeric, which is instantly turned brown by
ammonia. The excess of hydrochloric acid within the glass tube
may also be proved by placing a piece of blue litmus paper in the
tube before heating the compound, and it will be reddened by the
free hydrochloric acid.

In all cases of dissociation we may imagine two opposing forces
in operation, one being the external force supplying the energy
which tends to bring about the disruption of the molecules, and
the other being the force of the chemical affinity existing between
the disunited portions of the molecule, which tends to bring about
their reunion. When these forces are equally balanced, the same
number of molecules are dissociated as are recornbined in a given

Dissociation 91

unit of time, and the system is said to be in a state of equilibrium.
If by any means the balance between the two opposing forces is
disturbed, by augmenting or lessening either one or the other of
them, the equilibrium of the system will also be disturbed and a
new condition of equilibrium will be set up, in which again an equal
number of molecules undergo dissociation and combination in a
given time, but in which the ratio of the number of united and dis-
united molecules is different from that which obtained under the
former condition of equilibrium. The relation between these two
forces may be most readily disturbed, by either a change of tempe-
rature or pressure. Thus, in the case of nitrogen peroxide, N2O4,
when this gas is at a temperature of 26.7°, 20 per cent, of it is
dissociated into molecules having the composition NO2 ; and so
long as this temperature is maintained this ratio of the weight of
the dissociated molecules to the total weight of the system (known
as the fraction of dissociation) still subsists.

When the temperature of the gas is raised to 60.2°, the state of
equilibrium existing at the lower temperature is disturbed, and the
system gradually assumes a new condition of equilibrium, where
once more the actual number of molecules undergoing dissociation
and recombination in a given unit of time is the same, but where
the percentage of dissociated molecules in the gaseous mixture is
now 52.04.

It might at first be supposed when such a gas is heated, and a
temperature is reached at which the molecules are dissociated, that
they would all dissociate, and that the process once begun would
rapidly proceed until the decomposition was complete ; instead of
which, we find a definite fraction of dissociation corresponding to a
particular temperature. This may be explained on the basis of the
kinetic molecular theory. Let us imagine the gas nitrogen per-
oxide to be at a temperature below that at which dissociation
begins, when all the molecules will have the composition N2O4.
The molecules of the gas are in a state of rapid movement, and the
rapidity of their movement is increased by rise of temperature.
But the molecules in a given volume of the gas do not all move
at the same velocity, and therefore they have not all the same
temperature. On account of the infinite complications in their
movements, caused by their impacts against one another, some will
be moving at a speed considerably greater than that of the average,
and will have a temperature proportionally higher, wfrle others
again will have a velocity and a temperature below the average.

92 Introductory Outlines

The observed temperature of the gas, therefore, is not that of the
molecules having the highest or the lowest velocity and tempera-
ture, but is the average or mean temperature between, possibly, a
very wide range.

On the application of heat to the gas, the observed or mean
temperature rises, but the velocity of some of the molecules, and
consequently their temperature, may have been thereby raised to
the point at which dissociation takes place, and they consequently
separate into the simpler molecules. Let us suppose that the
observed temperature of the nitrogen peroxide is 26.7°, and that it
is maintained at this point. Although this temperature may be
below the dissociation temperature of the molecules, it must be
remembered that it only represents the mean temperature, and that
while some of the molecules have a lower, some also have a higher
temperature. As already mentioned, at the temperature of 26.7°,
20 per cent, of the molecules afe dissociated ; that is to say, at
any given instant one-fifth of the total number of molecules reach
a velocity which causes them to break down into the simpler NO2
molecules, which themselves then take up independent movements.
If, in the process of their movements, two of these disunited mole-
cules come into contact with each other at a moment when their
velocities are lower than that at which they dissociated, they at
once reunite, so that at the same instant some are uniting and
others are dissociating, and, the two processes going on equally,
the percentage of disunited molecules at any moment is the same,
although the actual molecules which are dissociated at one point
of time may not be the identical ones that are in this state at
another time. Let us now suppose the gas to be heated until the
registered (i.e. the mean) temperature reaches 60.2°, and that it be
maintained at this point. At this higher temperature a much
larger proportion of the molecules will acquire a velocity at which
they are unable to hold together, namely, 52.04 per cent.; but the
remainder, amounting to nearly one-half, are still at a temperature
below that at which dissociation takes place. Under these altered
conditions a greater number of disunions and reunions takes place
during a given interval of time, but the numbers are equal, and
therefore the equilibrium exists. If once more the gas be further
heated, until the indicated temperature is 140°, then it is found
that the whole of the N2O4 molecules have dissociated into NO2
molecules ; that is to say, when the mean temperature has reached
140°, then even those molecules that are moving with the slowest

Balanced Actions 93

s]peed have reached the temperature of dissociation. It will be
evident that the rate at which the fraction of dissociation in-
creases, as the temperature of a gas is gradually raised, will be
greatest when the mean temperature approaches the real dissocia-
tion temperature of the gas, for the temperature of the greater
number of the molecules will be coincident with, or very closely
approximating to, that point.

The vapour density of nitrogen peroxide, if it could be ascertained
when all the gaseous molecules had the composition N2O4, would
be 46 ; while that of the gas, when entirely dissociated into NO2
molecules, is 23. At temperatures between these extremes, the gas,
consisting of mixtures of both molecules, will have a density lying
between these figures, thus at 27.6° and 60.2° the density is 38.3 and
30.1 (see Nitrogen Peroxide, and also Phosphorus Pentachloride).

The effect of increased pressure upon a gas being to diminish
the mean free path of the molecules, and thereby increase the
number of molecules in a given space, the number of impacts
between the molecules in a given time will be increased. If,
therefore, while the nitrogen peroxide is maintained at a constant
temperature, say 62.2°, the pressure be increased, the dissociated
molecules, having shorter distances to travel, and making more
frequent impacts in a given time, will unite more quickly than
others are being disunited, and a fresh condition of equilibrium
will be established for any particular pressure.

The case of phosphonium chloride already mentioned may
be referred to as an illustration. This compound is completely
dissociated into molecules of hydrogen phosphide, PH3, and
hydrochloric acid, below a temperature of o°. If, while at this
temperature, it be subjected to pressure, the dissociated molecules
are caused to unite, and at a pressure of thirteen atmospheres the
union is complete, the whole of the disunited molecules having
combined to form molecules of phosphonium chloride, PH4C1.

If in the process of dissociation one of the products be with-
drawn from the sphere of action, then the process may be carried
on to completion. For example, in the case of calcium carbonate
already quoted, if this substance is heated in such a manner that
as fast as it dissociates, the gaseous product, namely the carbon
dioxide, is allowed to escape and so pass away from the sphere of
action, the change expressed by the equation

CaCO3 = CaO + CO2
will proceed until the whole of the carbonate has been converted

£4 Introductory Outlines

into oxide. But if, on the other hand, the action is made to take
place in a closed vessel, so that the carbon dioxide remains in
contact with the lime, then the reverse action comes into operation,
namely —

CaO + CO2 = CaCO3,

and a condition is arrived at in which the one action proceeds at
the same rate as the other. The pressure exerted by the carbon
dioxide under these circumstances is spoken of as the dissocia-
tion pressure of the calcium carbonate for that particular tem-
perature.

If, now, when this condition of equilibrium is established the
temperature be raised, the balance will be disturbed, and the
materials will readjust themselves to a fresh condition of equilibrium
at the higher temperature in which the dissociation pressure will
also be greater. For any given temperature, therefore, the dis-
sociation pressure is the only possible pressure at which a state of
equilibrium can be established between carbon dioxide, calcium
carbonate, and calcium oxide ; for if while the temperature is con-
stant the pressure upon the gas were to be increased by external
means and maintained at a higher point, union between the carbon
dioxide and lime would proceed until the whole of the lime was
converted into the carbonate. On the other hand, if the pressure
were to be reduced and maintained at a lower point, then dis-
sociation would go on until the action was complete and once
moje one of the three interacting substances would cease to
exist.

Increasing and diminishing the pressure upon a gas is obviously
synonymous with increasing and diminishing the number of mole-
cules in a given volume. This in modern phraseology is called
the molecular concentration of the gas, which embodies the same
idea as the expression active mass. From the above illustration,
therefore, it will be clear that there is some connection between
the molecular concentration (or active mass) of the carbon dioxide
and the rate of the chemical actions in question. This connection
is thus formulated (Guldberg and Waage) : the rate of chemical
action is proportional to the active mass (molecular concentration)
of each of the reacting substances. Advantage is sometimes
taken of these facts in determining the vapour-density of a sub-
stance which when heated dissociates into two gaseous con-
stituents. For example, phosphorus penlachloride when heated

Balanced Actions £5

dissociates into phosphorus trichloride and chlorine (see page 466),
according to the equation —

PC15 £ PC13 + C12.

But if the active mass of either the chlorine or the trichloride be
increased by adding more molecules of either one of these sub-
stances from some other source, the extent to which dissociation
takes place will be proportionally diminished. Hence, by heating
the pentachloride in an atmosphere of chlorine and thereby greatly
increasing the molecular concentration of this gas, dissociation may
be so far prevented that the density of the vapour is found to have
practically the normal value for the compound PC16.

CHAPTER XI
ELECTROLYSIS AND ELECTROLYTIC DISSOCIATION

IF a strip of pure zinc and a strip of platinum be together dipped
into a vessel containing dilute sulphuric acid, neither metal is
affected by the acid, so long as the metals do not touch each other.
If the ends of the strips outside the liquid be joined by means of a
metal wire, the zinc gradually dissolves in the acid, and bubbles
of hydrogen are disengaged from the liquid in contact with the
surface of the platinum plate (which itself is otherwise unaffected
by the acid), and at the same time an electric current passes
through the wire. So long as the chemical action of the sulphuric
acid upon the zinc proceeds, so long will the electric current con-
tinue to pass ; in other words, chemical energy will be transformed
into electrical energy. If the wire be severed, the electric current
can no longer pass, and the chemical action at once stops.

Such an arrangement constitutes a galvanic or voltaic element
or cell, and a series of such cells forms a galvanic battery. The
zinc plate, or the end of a wire that may be connected to it, is
termed the negative pole of the battery, while the end of a wire
attached to the platinum plate is the positive pole. Other arrange-
ments can be employed for generating a galvanic current, but in
all cases the electrical energy is derived ultimately from chemical
action.

If the two poles of a battery are connected together by placing
them both in contact with various different substances, it is seen
that in some cases the electric current passes, and in others not.
For instance, if the poles are joined by placing them both in contact
with a bar of sulphur, no current passes, whereas when connected
by a rod of graphite the current freely passes. Substances which
behave in this respect like the sulphur are said to be non-con-
ductors of electricity, while those that allow the current to pass
are distinguished as conductors. Substances capable of conducting

electricity are of two kinds, namely, those which are merely heated,

Electrolysis and Electrolytic Dissociation 97

and those which undergo a chemical change in consequence. All
the metals, and a few of the non-metals, belong to the first of these
classes ; while the second includes a large number of compound
substances, which are either in the liquid state or in solution in
some solvent. Thus, if the poles of a battery are immersed in pure
water, practically no current passes, because this liquid is a non-con-
ductor ; but if a quantity of hydrochloric acid (HC1) be dissolved in
the water, the solution at once becomes a conductor, and it is seen
that gas is disengaged from the liquid upon the surface of each
wire. If the solution of hydrochloric acid is moderately strong,
it will be found, upon examination, that the gas evolved at the
negative pole is hydrogen, while that from the positive pole is
chlorine : the hydrochloric acid, therefore, is separated into its
elements by the passage of an electric current through its
aqueous solution. Such a process is termed electrolysis; and
the conducting liquid is known as an electrolyte.

The poles or terminals that are introduced into the electrolyte
are called electrodes^ the negative electrode being termed the
cathode, and the positive electrode the anode.

Liquids which do not conduct electricity, or conduct only with
extreme difficulty, such as water, benzene, aqueous solutions of
alcohol or of sugar, are called non-electrolytes; while those which
are good conductors, such as aqueous solutions of hydrochloric
acid or of sodium chloride, are called electrolytes. Other liquids
range themselves between these two extremes with respect to
their conductivity, but those which may be said to fall about
midway are sometimes spoken of as half-electrolytes. These
terms, strictly speaking, apply to the actual liquids or solutions ;
thus in the above examples it is the aqueous solution of sugar
which is the non-electrolyte, and the aqueous solution of sodium
chloride which is the electrolyte. For brevity, however, it is usual
to apply the terms to the substance in solution, and to understand
that an aqueous solution is meant unless another solvent is specially
mentioned. Thus, when we say that sugar is a non-electrolyte,
and sodium chloride an electrolyte, it is the aqueous solutions of
these substances that are referred to.

In the class of electrolytes are included the strong acids, such
as nitric, hydrochloric, and sulphuric acids j the strong bases, such
as the hydroxides of the alkali metals, and almost all the class of
substances known as salts, irrespective of whether the acids and
oases they are composed of are electrolytes or half-electrolytes.

98 Introductory Outlines

The half-electrolytes are the weak acids, such as acetic, tartaric,
and oxalic acids, and the weak bases, as ammonium hydroxide and
the hydroxides of divalent metals other than the alkaline earth
metals. Non-electrolytes are substances of a neutral character
such .as sugar, this class including the large majority of organic
compounds which do not happen to fall under the category of
acids, bases, and salts.

In a great number of instances the electrolytic separation is
accompanied by certain secondary reactions, caused by the
action of the primary products of the electrolysis upon either the
electrolyte or the solvent ; for example, when a solution of sodium
chloride (NaCl) is electrolysed, the primary products are sodium and
chlorine, the latter appearing at the anode and the sodium making
its appearance at the cathode. The sodium, however, in contact
with the water in the neighbourhood of the cathode at once reacts
with the liquid, with the liberation of its equivalent of hydrogen,
according to the equation —

2Na + 2H2O=2NaHO + H2.

Similarly, in the case of hydrochloric acid, if the solution is
sufficiently dilute the final products obtained by subjecting it to
electrolysis are not hydrogen and chlorine, but hydrogen and
oxygen. The primary products are the same as before, but under
the altered condition the chlorine which is discharged at the
anode acts upon the water, combining with the hydrogen, and
liberating an equivalent quantity of oxygen : the two actions
being expressed by the equations —

Again, when a dilute solution of sulphuric acid in water is
electrolysed, the acid separates into the two primary products
H2 and SO4. The hydrogen as before appears at the cathode,
while the group or radical SO4 passes to the anode, where it under-
goes decomposition in contact with the water, reforming sulphuric
acid, while oxygen escapes. Thus —

2H2SO4 = 2H2 + 2SO4
2SO4+2H2O = O2 + 2H2SO4.

It will be observed that the final products are oxygen and
hydrogen in the proportion of two volumes of hydrogen to one

Electrolysis and Electrolytic Dissociation 99

volume of oxygen ; that is, the proportion in which they exist
in water. This process is, in fact, the same as that frequently
spoken of as the " electrolysis of water."

If instead of a solution of sulphuric acid, a solution of sodium
sulphate, Na2SO4, is treated in the same way, this compound
separates into the two primary products 2Na and SO4 ; the
sodium appearing at the cathode and the SO4 at the anode.
The sodium in contact with the water reacts as explained above,
liberating an equivalent quantity of hydrogen ; while the SO4
group, as before, gives rise to the reformation of sulphuric
acid and the liberation of oxygen. The final products, there-
fore, are again hydrogen and oxygen in the same proportions
as before.

In the same way, when an aqueous solution of copper
sulphate (CuSO4) is submitted to electrolysis, the primary
products are copper, Cu, and the group SO4. The copper is
liberated at the cathode, and since it exerts no action upon
the water, it is deposited as a metallic film upon the electrode.*
The group SO4 again passes to the anode, where it undergoes
decomposition in the presence of the water, as in the former
cases. t

Faraday's Law.— When the same quantity of electricity is
passed through different electrolytes, the ratio between the
quantities of the liberated products of the electrolysis is the
same as that between their chemical equivalents.

Thus, if the two electrolytes, hydrochloric acid and dilute sul-
phuric acid, be introduced into the same electric circuit, hydrogen
and chlorine are evolved in the one case and hydrogen -and oxygen
in the other. If the gases be all collected. in separate measuring
vessels, it will be seen (i) that the hydrogen and chlorine evolved

* This is the essence of the process of electro-plating. The metal to be de-
posited, whether it be gold, silver, or nickel, &c. , in the form of a suitable salt
(usually a double cyanide) in aqueous solution, forms the electrolyte. The object
to be plated is made the cathode, that is, it is suspended in the liquid and is
connected to the negative electrode of a suitable battery. The anode consists
of a strip of the metal to be deposited. Thus in silver plating, a strip of silver
is employed, and in this way the acidic radical that is liberated at the anode
dissolves the metal, and thereby prevents the weakening of the solution,
which would otherwise result from the gradual deposition of silver upon the
cathode.

f For fuller explanation of these changes see page 207.

ioo Introductory Outlines

from the hydrochloric acid are equal in volume; (2) that the
volume of hydrogen collected from the other electrolyte is the same,
while that of the oxygen is equal to only one-half this amount.
Knowing the relative weights of equal volumes of these three gases
to be hydrogen, oxygen, chlorine, as I, 16, 35.5, we see that they
must have been liberated in the proportions by weight of—-
Hydrogen = i Oxygen = 8 Chlorine = 35.5.

Similarly, if the same quantity of electricity be passed through
aqueous solutions of hydrochloric acid (HC1), silver nitrate (AgNO3),
copper sulphate (CuSO4), and gold chloride (AuCl3), by the time
that i gramme of hydrogen has been liberated from the hydro-
chloric acid, there will be deposited upon the cathodes of the other
electrolytic cells 108 grammes of silver, 31.7 grammes of copper,
and 65.6 grammes of gold. These numbers, which are the electro-
chemical equivalents, are identical with the chemical equivalents of
those elements, the chemical equivalent of an element being its
atomic weight divided by its valency.

H. o. Ci. Ag. Cu. Au.

Atomic weights . . i 16 35.5 108 63.5 197

Valency . . . . I 2 i i 2 3

Regarding the quantity of electricity required to liberate I
gramme of hydrogen as the unit, we may say that 16 grammes ot
oxygen require 2 units of electricity for its liberation, 108 grammes
of silver i unit, 63.5 grammes of copper 2 units, and 197 grammes
of gold 3 units ; or, in other words, the number of units of
electricity required to liberate a gramme-atom is identical with
the number representing the valency of that atom in the particular
electrolyte employed.

Some metals, such as copper, mercury, tin, &c., are capable of
functioning with different degrees of valency. Thus copper is
divalent in copper sulphate and in cupric chloride, but mono-
valent in cuprous chloride. If, therefore, i unit of electricity be
passed through aqueous solutions of each of these copper chlorides,

in the case of cupric chloride — ^ = 31.7 grammes of copper will

be deposited, while in the cuprous chloride -3:5 = 63.5 grammes

are formed.
The Ionic Theory.— The modern theory now generally held,

The Ionic Theory ' ibir

to explain the phenomena of electrolysis, is known as the theory
of electrolytic dissociation or the ionic theory. The passage of
electricity through conductors of the two classes txbove mentioned,
that is, through conductors such as metals, and those which are
electrolytes, may be compared with the two ways by which heat
is transmitted, namely, by conduction and convection. When
a bar of metal is heated at one end, the heat travels along
the bar, the metal remaining stationary ; but when water is
contained in a tube which is heated at its lower end, the heated
particles of water travel along the tube, conveying the heat
to the other extremity. In a similar manner, when electricity
passes through a metallic conductor, the electricity travels through,
or along, the metal, which itself does not move;* but when
it is passed through an electrolyte, it is conveyed or transported
through the liquid by the moving particles, to which the name
ions (signifying wanderers) was first given by Faraday. One set
of ions charged with negative electricity travels towards the anode,
while another set conveying positive electricity moves towards the
cathode. Inasmuch as the negative ions appear at the anode they
are called anions, while the positively charged ions are distinguished
as cations. In the earlier stages of the development of the present
theory it was supposed that the electrolyte was only separated
into its ions as the electric current was passed into it, that the
electricity was the prime cause of the dissociation of the electro-
lyte, hence the expression electrolytic decomposition, still commonly
used. It was believed (Grotthus) that the first effect of the current
was to cause the molecules in the solution to take up positions
towards each other and the electrodes which may be crudely
represented by the top line in the following diagram, where the
molecules of hydrochloric acid, for example, are arranged with
their electro-negative constituents all directed to the anode, and
their electro-positive elements towards the cathode, precisely as a
number of separate cells in a battery would be arranged. Then
that a disruption of the molecules took place in which those
nearest to the electrodes parted with their positive and negative
ions to their respective electrodes (where they would be disengaged
as free hydrogen and chlorine in the case of hydrochloric acid),
while an exchange of partners between the other molecules all
along the line took place, as represented in the second line, result-
ing in the formation of fresh molecules of the original compound.

* In the language of the modern theory of the atomic nature of electricity,
it is the electrons which travel, while the metal ions remain (probably) stationary.

Introductory Outlines

These would then immediately assume the position of those in the
upper row. This theory, while affording an explanation of many
of the phenomena connected with electrolysis (such as the fact
that the ions are disengaged only at the surface of the electrodes,
and not in the intervening space ; that the appearance of the
liberated ions takes place simultaneously at the two electrodes,
however far removed from each other, &c.), was not capable of
satisfying all the facts of the case. It was pointed out (Clausius)
that if the electric current were the actual cause of the separation
of the molecules into their constituent ions, this ought to be made
manifest by the fact that the current would have to expend energy
in doing the work of effecting such decomposition. But exact
experiment shows that this is not the case. It is found that when
an electric current passes through an electrolyte, no electric energy
is absorbed in causing the dissociation of the molecules of the

HCl HCl HCl HCl HCl

(£P)

9 Q© 0® 0® ©0 0® 0

H CIH Cl H Cl H ClH Cl H Cl

FIG. 12.

dissolved substance ; but that the current is conducted by electro-
lytes with the same freedom as it is by metallic conductors. In
other words, it has been shown that Ohm's law is equally appli-
cable to electrolytes as it is to metals, namely, that the current is
proportional to the electro-motive force for all values of that force.
The theory of electrolytic dissociation, first proposed byArrhenius,
and now generally accepted by chemists and physicists, is that all
solutions which are capable of conducting electricity contain mole-
cules which are already in a state of dissociation. That is to say,
the electrolyte consists of molecules which are already dissociated
into their constituent ions to a greater or less extent. The simple
act of solution in water results in the dissociation of a portion of
the molecules into their positive and negative ions. For example,
a solution of sodium chloride is an electrolyte ; when, therefore,
this substance is dissolved in water a certain proportion of the
molecules immediately undergoes ionic dissociation, so that the
solution contains some molecules of sodium chloride, some sodium
ions, and some chlorine ions ; a state of balance or equilibrium
between the ions and the undissociated molecules being main-

The Ionic Theory 103

tamed, depending upon various conditions. In such solutions
it is the ions alone which take any part in the conduction of
the electric current, the undissociated molecules being entirely
inoperative. Obviously, therefore, when a substance dissolves
in water without undergoing ionic dissociation, the solution
will be a non-electrolyte j while if dissociation only takes place
to a limited extent the solution will come under the head of
the half-electrolytes. Strong acids, bases, and salts, which are
good electrolytes, are therefore the substances which undergo
dissociation to the greatest extent. For any given solution the
extent to which dissociation takes place increases as the solution
is diluted until a point is reached at which all the molecules are
dissociated into their ions.

At first it might appear contrary to established ideas that in
such a case as sodium chloride, for instance, the sodium and
chlorine in the free or separated state should be capable of exist-
ence side by side in the same liquid — a liquid, moreover, upon which
one of these elements, namely, the sodium, is under ordinary circum-
stances capable of exerting a chemical action. Similarly, that with
such a compound as sodium sulphate there should not only be the
same element, sodium, existing in contact with water, but also a
group of elements, or radical, SO4, which is not known in a state of
separate existence. These ions, however, whether elementary like
sodium or compound like the group SO-4, are all united with and
carry with them enormous electrical charges, positive or negative,
as the case may be ; and it is only so long as they retain their
electrical charges that they can retain an independent existence
and exhibit their own special properties. When the electrodes
from an electric battery are introduced into a solution of sodium
chloride, the sodium ions with their positive charges are attracted
to the cathode ; they there discharge their loads of electricity, and
thereupon become ordinary molecules of sodium, possessing the
properties usually associated with that metal. Hence, since
ordinary sodium cannot exist in contact with water, the metal
immediately upon its liberation at the cathode reacts upon the water
with which it is in contact in the manner usual to sodium. Similarly,
the chlorine ions with the negative electric charges are endowed with
their own characteristic properties, which are retained so long as
the atom is united to the electricity. So soon as it loses its charge,
which it does when it conveys it to the anode, the chloride ion
then becomes a chlorine atom, two of which immediately unite,
forming a molecule of the element possessing the ordinary proper-

104 Introductory Outlines

ties of chlorine gas. If, therefore, we use the term radical to
embrace single atoms as well as groups of atoms, we may describe
an ion as a radical united to an electric charge — a positive ion
being one which carries positive electricity, and a negative ion
being a radical which is united to a negative charge.

Indeed, instead of regarding this subject as one presenting a
new difficulty to the mind, we may even trace an analogy between
it and another set of ideas with which we are already quite fami-
liar. We know that when two elements enter into chemical union
with each other they lose their own characteristic properties, and
that the resulting compound is endowed with new and different
properties ; when an atom of sodium combines with an atom of
chlorine the sodium no longer exhibits the properties of metallic
sodium. Similarly, when an atom of sodium is combined with a
negative electric charge, the product of the union, namely, the ion^
possesses properties differing from those of metallic sodium. The
exact "how" and "why" are equally mysterious in both cases,
and in neither case are we able to explain the precise nature
of the union for which in both instances we employ the word
"combine." Since the immediate effect of passing an electric
current through an electrolyte is to cause the ions to travel to
their respective electrodes, and there becoming electrically dis-
charged to cease to exist as ions^ it will be evident that the
condition of equilibrium previously existing between the ions and
the undissociated molecules is at once disturbed. This disturbance,
however, immediately adjusts itself by the dissociation of more of
the molecules ; as fast as ions are removed fresh molecules dis-
sociate into ions. Hence, although the electric current is not the
prime cause in the production of the ions, it is in a sense an
indirect cause, since by bringing about the removal of the ions
previously present it induces conditions which allow more of the
molecules to dissociate into ions.

Atomic Electric Charges — Valency. — If we take as our unit
the amount of electricity which is carried by one atom of hydrogen,
then of all monovalent ions we may say that they convey one unit
of electricity, for all such ions are united to equal amounts of
electricity, whether they be simple or complex radicals. Divalent
and trivalent ions respectively are united to two and three units of
electricity. Valency may, in fact, be defined as the number of
unit electric charges which are united to an atom (or radical).
These electric charges are called electrons^ or atoms of electricity,
in accordance with the present-day views as to the nature of

The Ionic Theory 105

electricity. Electricity is now regarded as having an atomic
structure : it is believed to consist of indivisible and inde-
structible particles, positive electrons and negative electrons,
comparable in a measure with the atoms of monovalent chemical
elements. To denote these electrons, or atomic charges of elec-
tricity, the symbols + and - are employed j they represent one
"atom of electricity" (positive and negative respectively), just as
the symbol H stands for one atom of hydrogen.'55'

A positive electron combined with a positive chemical atom
or radical gives rise to a positive ion, or cation', while negative
elements or radicals united to negative electrons constitute nega-
tive ions or anions.

Ionic Notation. — In chemical notation it is usual to represent
ions by employing either the ordinary 0 and © signs, or more
commonly a dot (') and dash ('), in conjunction with the chemical

Symbol for the atom or radical. Thus Na or Na* signifies a sodium

ion, and Cl or Cl' represents the chloride ion.

The symbol Na' therefore conveys the information that the
sodium ion is a monovalent cation ; while Cl' indicates that the
chloride ion is a monovalent anion. SO4", in the same way, stands
for the sulphate ion, with its two negative charges, and Fe'" for
the trivalent ferric ion with its triple charge of positive electricity.
Sodium chloride in solution would be represented by the formula
Na'Cl', ferric chloride by Fe"'Cr3, potassium sulphate by K'2SO4",
and so on.

In the system of nomenclature of the ions now generally adopted,!
the names of the cations are formed by the addition of the termina-
tion ion to the stem of the chemical name of the element or radical ;
thus, hydrion, H', sodion, Na', ammonion, NH'4, calcion, Ca",
zincion, Zn", &c.

When it becomes necessary to indicate the number of unit
charges (i.e. the valency) of the radical, Greek numerals are pre-
fixed to the name. For example, diferrion, Fe" (the ions in ferrous
salts), triferrion, Fe'" (the ions in ferric salts) ; monocuprion, Cu*,
and dicuprion, Cu", for the cations in cuprous and cupric com-
pounds respectively.

* Negative electrons are known in the free state. The "cathode" rays
emitted from a Geissler vacuum tube consist of these negative electrons, and
they also form a part of the "radiation " emitted by the element radium (see
Appendix). So far positive electrons have not been isolated.

f First introduced by J. Walker.

106 Inorganic Chemistry

In the case of anions the names are formed by the use of one of
the three terminations — idion, anion, and oston, depending upon
whether the salt radical ends in zVfe, ate, or ite. For instance,
anions derived from chlor/V&f, brom/Vfe.?, hydroxzV/<?.y, sulphz'dk?, will
be chlor/</zVw Cl', brom/dfo?# Br', hydroxzV/z'072 OH', sulphz'dfow S"
respectively ; those from chlora/£r, sulph<?/<?.r, orthophosphc?/^,
&c., chloral?* C1O3', sulpha/uV?* SO/, orthophosprnzmV?;* PO4'",
&c. ; while those derived from such salts as mtrites and sulphz'/^
are termed \\itrosion NO2', su\plnosion SO3". These names are
employed precisely as ordinary chemical names are used, that is
to say, they apply to the material taken collectively, and not to
the particles themselves of which the material is composed.*

It is often convenient to regard the amount of electricity which
is carried by one gramme of hydrogen as the unit, instead of that
conveyed by one atom. The value of this unit is 96,550 coulombs.
Hence these dots and dashes signify that one, two, or three times
96,550 coulombs of electricity are carried by the gramme-molecule
(see p. 57) of the ion according to the number of these signs
attached to it. Thus 96 grammes of SO/ will carry 96,550x2
coulombs of negative electricity; 18 grammes of NH4* carries
96,550x1 coulombs of positive electricity, and 95 grammes of
PO/' conveys 96,550x3 coulombs, or 3 units of electricity. In
other words, each dot and dash attached to the formula signifies
one charge of 96,550 coulombs united to the gramme-molecule of
the ion.

* Just as the names sodium, hydrogen, chlorine, &c. , are used to denote
matter which is made up of atoms or molecules of sodium hydrogen or chlorine
respectively, so the terms sodion, hydrion, chloridion, are the names applied
to the matter which is composed of sodium ions, hydrogen ions, and chloride
ions respectively. We speak of a sodium atom, and of hydrogen molecules, so
also of a sodi-um ion and hydrogen ions. But to use such expressions as a
sodion, or hydrions, is as meaningless as to speak of a sodium or hydrogens.

The translation of the chemical equation

2HC1 = H2 + C12

is that hydrochloric acid is decomposed into hydrogen and chlorine — or that
two molecules of hydrogen chloride yield one molecule of hydrogen and one
molecule of chlorine— similarly the ionic equation

HC1 = H- + C1'

signifies that on solution in water hydrochloric acid is ionised into hydrion
and chloridion — or that a molecule of hydrogen chloride yields on ionisation
a hydrogen ion and a chloride ion.

The Ionic Theory to/

It will be evident that ionisation or electrolytic dissociation
is a phenomenon of a- different order from that which takes
place when a compound dissociates under the influence of heat,
as discussed in the previous chapter. Unde.r these circumstances
it was explained that the salt ammonium chloride, for example,
dissociates when heated into the two compounds NH3 and HC1 ;
whereas when it is dissolved in water it undergoes electrolytic
dissociation into the two ions NH4' and Cl' ; in the first case the
products are electrically neutral chemical compounds, while in the
latter they are electrically charged ions, or compounds of radicals
with electrons.

From the point of view of the ionic theory, acid, bases, and
salts all behave in a perfectly similar manner ; to the " ionist," as
such, there is no difference between these three kinds of sub-
stances ; it is therefore sometimes convenient to class them all
together as salts. Those which from a chemical point of view are
acids, from the ionic standpoint are salts of hydrogen^ that is, salts
in which all the positive ions are hydrogen'; while those which
are usually termed bases are spoken of as salts of hydroxyl, or salts
in which the only negative ions are hydroxide ions.*

Molecular Conductivity. — What is understood as the molecular
conductivity of a solution is its specific conductivity expressed in
the usual electrical units, divided into the number of gramme-
molecules of the dissolved substance contained in the solution ; or
what is the same, multiplied by the number of litres of the solution
which contains one gramme-molecule of the substance.

Now since it is the ions present in an electrolyte which alone
take any part in the conveyance of electricity, the undissociated
molecules present being inoperative, it will be obvious that the
molecular conductivity of an electrolyte will depend partly upon
the number of ions present— in other words, upon the extent to
which the electrolyte is dissociated— and partly upon the rate at
which the ions travel or migrate in the liquid.

It has been found (Hittorf) that different ions under the same
conditions travel at different rates. From determinations of the
changes in concentration which take place in the electrolyte

* The student will not fall into the error of supposing that it would be
either desirable or possible to abolish the classification of acids, bases, and
salts. From a purely chemical standpoint acids and bases are two perfectly
distinct classes of compounds, and these two terms will always be employed
to denote them.

io8 Introductory Outlines

immediately round the electrodes, it has been shown that in a
solution of given concentration and under the same electrical
conditions, all the ions of one kind travel with a constant velocity,
but that the rate differs for different kinds of ions. For example,
it is found that the ion H* migrates with a velocity about twice as
great as that at which the negative ion HO' travels, and about five
times the rate at which the cation K* migrates.

When, therefore, a solution is diluted, and its molecular con-
ductivity thereby increased, this increased conductivity will be due
partly to the greater rate of migration of the ions which follows
upon dilution, and partly to the increased number of ions present ;
for, as already stated, as the solution is diluted more and more, so
ionisation takes place to a greater extent.

It is found by experiment that as the solution is diluted, the
molecular conductivity at first rises somewhat rapidly, that is to
say, a moderate increase of dilution causes a considerable rise in
conductivity ; but after a certain dilution is reached, the rate of
increase of molecular conductivity is greatly diminished; and
after continuing slowly to increase on further dilution, a point is
at length reached beyond which no increase of conductivity follows
upon additional dilution. The conductivity at this latter point is
called the molecular conductivity at infinite dilution, and at this
point the whole of the electrolyte has become dissociated into its
ions. The point of dilution at which the rate of increase of mole-
cular conductivity makes the marked change may be regarded as
the point at which dilution ceases to influence the rate of migra-
tion of the ions.

Since the molecular conductivity is in this way dependent upon
two factors, namely, the speed of migration of the ions and the
degree of ionic dissociation, it will be obvious that it cannot by
itself afford a true measure of dissociation. The dissociation
coefficient^ or the fraction of the molecules of an electrolyte which
are dissociated into their ions at a given concentration, is the ratio
between the molecular conductivity at that concentration to the
molecular conductivity at infinite dilution. Hence, if m^ and
*#c are the molecular conductivities at definite dilution and at
concentration c respectively, then the coefficient of dissociation d
will be —

Some general idea of the degrees of dilution which are being
dealt with in these considerations may be gained from a single

The Ionic Theory 109

example. Thus in a solution of common salt, the strength of the
solution at which the rate of the migration of the ions is practically
unaffected by further dilution is such that one litre contains
about ^jth of a gramme-molecule of the salt, or 5.85 grammes ;
while a solution which has been diluted until its molecules are
wholly dissociated contains only about To^ortth of a gramme-
molecule per litre, or is a thousand times more dilute.

Some Applications of the Ionic Theory.— The ionic theory is in
harmony with and derives support from the laws which regulate
the influence of substances in solution upon osmotic pressure
(page 158), upon the lowering of the vapour-pressure (page 135),
and upon the lowering of the freezing-point of the solvent
(page 140). Dilute solutions of electrolytes are found to ex-
hibit deviations from these laws much in the same way that
gases which undergo dissociation depart from the usual gas laws.
Thus it is observed that in the case of dilute solutions of electro-
lytes, the osmotic pressure, the lowering of the vapour-pressure,
and the lowering of the freezing-point of the solvent, instead of
being proportional to the number of molecules of the dissolved
substance, are proportional to the number of dissociated ions.

Again, this theory affords an explanation of the fact that the
heat of neutralisation of one equivalent of strong acids and bases
(in dilute solution) is practically a. constant, namely, about 13,700
heat units or calories (see page 165). Now, in the neutralisation
of, say, nitric acid by potassium hydroxide, according to the ionic
theory these two reacting substances are in a state of dissociation
in the dilute solution ; moreover, the salt potassium nitrate, result-
ing from the interaction, will also be dissociated. The only product
of the chemical action which is not dissociated is the water, as this
compound is practically a non-electrolyte ; * hence the process of
neutralisation of this acid with this base resolves itself into the
union of H* ions with HO' ions to form molecules of H2O, as may
be seen by the equation

in which the formulae for the dissociated molecules are written with
their ions separated by a comma. It will be obvious, therefore,

* Probably there is no such thing as an absolutely perfect non-electrolyte.
In reality water itself undergoes ionic dissociation to a very slight extent. It
has been estimated that in ten million litres of water there will be about one
gramme-molecule in the ionic state.

HO Introductory Outlines

that the final result, namely, the union of H* with HO', will be the
same if we substitute other strong acids or bases, thus —

H-,HyS04" + K',HO' = H-,K',SO4"+H2O

Therefore the heat of neutralisation of dilute solutions of these
acids and bases is in reality the heat of formation of H2O mole-
cules by the union of H* ions with HO' ions.

Similarly, the ordinary "reactions " employed in chemical analysis,
when considered from the standpoint of the ionic theory, become
invested with a new meaning, and are often rendered more intel-
ligible : one or two examples may be given. When the metal tin
is precipitated from a solution of stannous chloride by means of
metallic zinc, the following ionic equation expresses the change : —

Sn",Cl',Cr + Zn = Sn + Zn",Cl',Cl'.

In other words, the two unit charges of positive electricity have
been discharged by the tin ion, which then ceases to be an ion,
but appears as ordinary metallic tin, and are transferred to the
metal zinc, which then ceases to be ordinary metallic zinc, but
passes into the solution as a zinc ion.

Again, the tests for iron in the ferric state are really tests for
triferrion Fe"*, and tests for this metal in the ferrous condition
are tests for diferrion Fe". But if a compound containing this
metal should dissociate in such a manner as to afford neither
Fef" nor Fe" ions, it will be evident that the usual reagents em-
ployed to detect these ions will yield no result. The salt potassium
ferro-cyanide, K4Fe(CN)6, is a case in point. On solution this
compound dissociates into the ions K' and Fe(CN)6iv, and the iron
in this solution, therefore, does not respond to the usual tests for
either triferrion or diferrion.

Again, the action of ammonium chloride in preventing the pre-
cipitation of magnesium as hydroxide by ammonia, is explained by
the fact that ammonium hydroxide being a comparatively weak
base undergoes dissociation to only a slight extent into ammonion
NH4f and hydroxidion OH' — to an extent far smaller than is the
case with sodium and potassium hydroxides. Upon the addition
of an ammonium salt of a strong acid, such as hydrochloric acid,
we are throwing into the solution a large number of ammonium
ions, which has the effect of causing the re-union of the hydroxide
ions until practically the whole of the ammonium hydroxide present

The Ionic Theory \ \ \

is in the undissociated state, and as there is now no hydroxidion
present no magnesium hydroxide can be formed.

When no ammonium chloride is added, partial precipitation of
magnesium hydroxide results —

Mg",Cr,Cl'+NH-4,OH'+NH-4,OH'
= Mg(HO)2 + NH-4,Cl'+NH'4,Cl'.

But this process results in the introduction into the solution of
NH'4 ions, and equilibrium is established when these are present
in sufficient quantity to prevent further production of hydroxidion
by the dissociation of any more of the ammonium hydroxide.

Similarly the behaviour of many salts in yielding, when dissolved
in water, solutions which are either acid or alkaline, admits of an
ionic explanation. Sodium nitrite may serve as an example.

Unlike sodium nitrate, which yields a neutral solution, this salt
when dissolved in water gives a solution which is alkaline, that is,
a solution containing hydroxidion. When dissolved, the salt is
first largely ionised into sodion and nitrosion,

NaNO2 = Na',NO2'.

Besides these ions, however, there are also present minute quantities
of hydrion H' and hydroxidion OH' due to the very slight ionisation
of the water itself, hence we have the ions

Since nitrous acid is a weak acid, i.e. one which is only slightly
ionised in solution, the NO2' and the H* ions tend to unite to form
molecules of undissociated nitrous acid, HNO2, thereby causing
more water molecules to become ionised, with the consequent
increase in the number of hydroxide ions present* This process
goes on until equilibrium is established, which may be thus
represented —

Sodium hydroxide being a strong base, the hydroxidion and sodion
do not unite to form molecules, but remain in the ionic state.

Processes of this order are spoken of as hydrolysis — the sodium
nitrite in this case is said to be hydrolysed. All salts of weak acids
with strong bases behave in a similar manner.

CHAPTER XII
CLASSIFICATION OF THE ELEMENTS

IT has already been mentioned (page 7), that the elements may
be classified under the two subdivisions, metals and non-metals.
Further classifications have from time to time been in use, based
upon other properties, such, for example, as the valency of the
elements.

Classified according to their valency, the elements fall into six
subdivisions, consisting of mono-, di-, tri-, tetra-, penta-, and hexa-
valent elements. This system of classification has now largely
fallen into disuse, owing partly to the difficulties arising out of the
variability of valency so often exhibited, but more especially to the
more recent development of another system, known as the natural
classification of the elements, or the periodic system, which practi-
cally absorbs and includes the older method.

Certain remarkable numerical relations have long been observed
to exist among the atomic weights of elements that closely re-
semble one another in their chemical habits. In such groups or
families it is frequently seen that the atomic weight of one mem-
ber is approximately thejarithmetic mean of the atomic weights of
those immediately before and after-it, when they are arranged in
order of their atomic weights. This will be seen from the following
examples : —

Li. Na. K.

7 23 39

K. Rb. Cs.

39 85 133 39±I33 =

P. As. Sb.

31 75 120

S. Se. Te.

32 77 125

78.5

The Periodic Classification 113

If the elements in these various families are so arranged, as
to bring out the differences between their atomic weights, the
striking fact will be observed that the_ increase in the atomic
weights in each group takes place by practically the same incre-
ment In the following table the elements belonging to the same
group are placed in vertical columns, the differences between the
various atomic weights being placed between them : —

F = i9

N = 14

0 = i6

Na = 23

Mg = 24

Difference . 16.5

Diff. . 17

Diff. . 16

Diff. . 1 6

Diff. . 16

C1 = 35-S

P = 3'

8 = 32

K = 39

Ca = 40

Difference . 44.5

Diff. . 44

Diff. . 47

Diff. . 46.2

Diff. . 47-3

Br = 80

As = 75

Se = 79

Rb = 85.2

Sr = 87.3

Difference . 47

Diff. . 45 Diff. . 46

Diff. . 47.8

Diff. . 49.7

I = 127

Sb = 120

Te =125

Cs = 133

Ba = 137

It will be seen that in each group the difference between the first >
and second number is about 16, while between all the others the ^
increase in weight takes place by a number which approximates )
to 16 x 3.

This numerical relation between the atomic weights of elements
of the same family, and between the various groups, is obviously
not a chance one, and chemists were led by it to believe that the
properties of the elements were in some way related to their atomic
weights. Newlands (1864) was the first to point out, that if the
elements are tabulated in the order of increasing atomic weights,
the properties belonging to each of the first seven elements reap-
peared in the second seven, and he applied to this relation the
name of the law of octaves. A more elaborated and systematic
representation of Newlands' law of octaves was afterwards deve-
loped by Mendelejeff (1869), and which is now generally known as
MendelejefPs periodic law. At the present time, owing to the
recent discovery of the argon family of elements, it is not until
eight elements have been traversed that the properties of the first
reappear ; the term " octaves " is therefore no longer strictly
applicable.*

* Unless, indeed, we stretch the musical simile somewhat and look upon
these five inert gases as " accidentals."

1 14 Introductory Outlines

If the sixteen elements with lowest atomic weights, after
hydrogen, be arranged in order of increasing atomic weights in
two horizontal rows of eight, soine of these relations will be
recognised —

He=4 Li =7 Be =9 B =11 C = i2 N = i4 O = i6 F=i9.
Ne = 2o Na = 23 Mg = 24 Al = 27 Si = 28 P=3i 8=3201 = 35.5.

In traversing the upper row from helium to fluorine, we meet with
certain characteristic properties belonging to each member, and
also a certain gradation in those properties that are common.
Coming to the second row, many of the characteristic properties
of the members of the first row again appear, and the same regular
modulation is met with in passing along the series : thus helium
exhibits a likeness to neon, lithium resembles sodium, carbon
corresponds to silicon, fluorine to chlorine, and so on. These
resemblances are seen both in the physical as well as the chemical
properties of the elements, thus .lithium and sodium are both soft
white metals, and are strongly electro-positive. Fluorine and
chlorine are both pungent corrosive gases, and are intensely electro-
negative ; while helium and neon are neither electro-positive nor
electro-negative, have no chemical properties whatever, and
therefore no valency. Taking their power of combining with
chlorine and with hydrogen as indicative of their valency, we see
that the change in this respect, as the two series are traversed, is
the same in each, thus —

LiCl BeCl2 BC13 CC14 CH4 NH3 OH2 FH.
NaCl MgCl2 (Aids), SiCl4 SiH4 PH3 SH2 C1H.

The gradation in properties exhibited by the elements in a series
is also seen in their power of combining with oxygen, which will
be more clearly brought out if the formulae of the compounds be
so written as to indicate the relative proportions of oxygen with
which two atoms of each element unite, thus—

Na20 (Mg802) A1203 (Si204) P2O6 (S2O6) C12O7
MgO SiO2 SO3

Regarding, then, the eight elements of the first row as ^.period, we
find that the various properties exhibited by the several members
are met with again in those of the second period.

The Periodic Classification 11$

Not only do the properties of the elements themselves reappear,
but also those possessed by the various compounds they form : thus
lithium chloride (LiCl) and sodium chloride (NaCl) strongly re-
semble one another. The oxides of beryllium and magnesium
(BeO and MgO) have similar properties. The compounds of fluo-
rine and chlorine with hydrogen (HF and HC1) closely resemble
each other, and so on.

This periodic reappearance of similar properties, exhibited by the
elements and their compounds as the atomic weights of the former
gradually increase, is thus stated by Mendelejeff in his law of
periodicity. The properties of the elements, as well as the proper-
ties of their compounds, form a periodic function of the atomic
weights of the elements.

When the tabulation of the elements according to this system is
continued (after the completion of the second period with chlorine),
it will be seen that, beginning with argon, eighteen elements have
to be arranged before we meet with the reappearance of those pro-
perties that belong to the first ; that is to say, there are two
"octaves," one containing eight members like the former ones, and
one containing seven, and three elements over, which in the follow-
ing table are placed within brackets : —

A.* K. Ca. Sc. Ti. V. Cr. Mn. (Fe. Co. Ni.)

40 39 40 44 48 51 52 55 (56 59 59)

Cu. Zn. Ga. Ge. As. Se. Br.

63.5 65 70 72 75 79 80

This constitutes what is known as a long period, in contradis-
tinction to the two first, which are distinguished as short periods.
In certain respects, however, the last seven elements in this long
period exhibit resemblances to the seven in the first portion (count-
ing after the first element, argon) ; that is to say, the properties
displayed by the members of the first period, which is known as
the typical period, reappear twice over in the long period. The
three elements within the brackets are termed by Mendelejeff
transitional elements. Continuing the arrangement from bromine,
another long period occurs, again containing three transitional
elements : —

* It will be noticed that the element argon, A, is placed before potassium, K,
although, according to the atomic weights here given, it would appear as
though they should be in the reverse order. This will be discussed later.

Introductory Outlines

Kr. Rb. Sr. Y. Zr. Cb. Mo. - (Ru. Rh- Pd.)

83 85 87-6 89 90.7 93.5 96 ? (101.7 103 106)

Ag. Cd. In. Sn. Sb. Te. I.

108 112 114 118 120 125? 127

It will be seen that a gap is left where the eighth member of
the first part of this period should be, an element which would
correspond, in this period, with manganese in the period above.
This element is at present unknown. The remaining elements
belong to three other long periods, in which, however, the number
of gaps is very considerable, thus —

X. Cs. Ba. La. Ce. — — — ( — — — )

13° J33 J37 139 HO

195)

— -

Yb.
172

— Ta. W. — (Os. Ir.

181 184 (191 193

Au.
I97

Hg.
200

Tl.
204

Pb. Bi.
2O7 2O8

Th. —

Ur. — ( —

232

238.5

Those elements that fall in the first eight places of the long
periods are termed the even series, while the last seven are dis-
tinguished as the odd series; arranging them, therefore, in such a
manner as to bring the odd and even series into columns, we get
the table on page 118.

In this manner the elements are arranged in nine groups.
The first of these groups contains the so-called " inert gases " —
the five new elements of recent discovery, which take their place
rather outside this classification scheme, regarding it from a purely
chemical standpoint. And as the system of numbering the groups
of elements in this periodic arrangement has become familiarised
by long use, this group containing the " inert gases " has been
numbered Group O, and the systematic numbering of the other
groups begins as usual. The last group contains the transitional
elements that come between the even and odd series of the long
periods.

In each of the remaining seven groups, the elements belonging

The Periodic Classification 117

to the even series of their respective long periods, are placed to the
left, while those belonging to the odd series are arranged on the
right-hand side of each vertical column. In this way the groups are
divided into the subdivisions A and B, in which the resemblance
between the members is most pronounced. Thus in Group II.,
although there are certain properties common to all the members,
there is a much closer similarity existing between the elements
calcium, strontium, and barium than between zinc and calcium, or
cadmium and barium.* The elements in the two short periods
have been placed in that subdivision or family with the members
of which they exhibit the closest resemblance. Thus, in Group I.
lithium and sodium are more allied to potassium, rubidium, and
caesium, than to copper, silver, and gold; while in Group VII.
fluorine and chlorine are placed in the same family with bromine
and iodine, with which they exhibit a close similarity.

In the eighth group, containing the transitional elements, the
families consist of the horizontal and not the vertical rows ; that is
to say, the closest resemblance is between the three transitional
elements in each series, elements whose atomic weights, instead of
exhibiting a regular increase, as in the other families, have almost
the same value, such as Fe = 56 ; Co = 59 ; Ni = 59.

A glance at the table shows that in the last three long periods
there is a large number of gaps. It is possible that these gaps
may represent elements which yet await discovery. This supposi-
tion gains considerable support from the fact, that at the time
MendelejefF first formulated the periodic law, there were three such
gaps in the first long period, which have since been filled up by the
subsequent discovery of three new elements ; these will be referred
to later.

The periodic recurrence of some of the chemical properties
is indicated in the lowest horizontal column, where the general
formulas of the oxygen compounds and the hydrides are given ; R
standing for one atom of any element in the group. As explained
on page 114, these formulas are so written as to show the relative
amount of oxygen to two atoms of element, in order to establish
the true relation between the different groups. For example, the

* This, however, is by no means uniformly the case ; thus the element copper
(Group I.) in many of its chemical attributes is much more closely allied to
mercury (Group II.) than to silver ; and silver, again, more strongly resembles
thallium (Group III.) than either copper or gold, with which it is associate^
in this system of classification.

118

Introductory Outlines

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P? P=H

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The Periodic Classification 119

oxides of the elements of Group I. contain two atoms of the element
to one of oxygen, as Li2O ; but those of the second group only con-
tain one atom of the element, as CaO : hence the general formula
is doubled, R2O2. It will be seen, therefore, that the proportion of
oxygen relative to two atoms of the element regularly increases
from the first group to the eighth. The oxides of the members of
Group I. are strongly basic in character, and in general this basic
nature gradually diminishes as we traverse the series, giving place
to acidic characteristics, which are strongly marked in the seventh
group.

The periodic reappearance of the physical properties of the
elements is seen in such points as their electrical characters, their
malleability, ductility, melting-points, &c., all of which are in
harmony with the periodic law ; but in none is it more strikingly
seen than in their atomic volumes in the solid state. The atomic
volumes of the elements are the relative volumes occupied by
quantities proportional to their atomic weights, or by gramme-
atoms ; and they are obtained by dividing the atomic weights of
the elements by their specific gravities. In the case of gases, as
has been already explained on page 40, the specific gravity is
the density referred to hydrogen as the unit : the atomic volume,
therefore, of such a gas as oxygen is —

16 = atomic weight _
1 6 = density

The specific gravities of solids (and also liquids) are referred to
water as the unit, and as I cubic centimetre of water weighs
I gramme, the specific gravity of a solid or liquid expresses the
weight in grammes of I cubic centimetre of the substance. Dividing
the atomic weight, expressed in grammes, by the weight in grammes
of I cubic centimetre (i.e. the specific gravity), the atomic volume
will be represented in cubic centimetres. It must be remembered
that the atomic volumes do not express the relative volumes that
are actually occupied by the atoms, they represent in reality the
relative volume of the atoms plus the unknown volumes of the
spaces that separate them.

The following table gives the specific gravities and the calculated
atomic volumes of the first and the middle elements of the two
short and two long periods, not counting the group of " inert "
elements :—

120

Introductory Outlines

Specific
Gravity.

Atomic
Weights.

Atomic
Volumes.

ist Period {^m. ,

0-S9
3.0

11.9
4

( Sodium . . . --':.- , -'

0.97

23.6

( Silicon . , ,

2.4

28.3

n.8

( Potassium
3rd Periodi Iron (Cobalt-Nickel)
., r,^ . j ( Rubidium
4th F< 3d I (Ruthenium-Rhodium) Palladium

0.865
7-79
1-52
11.4

106.2

45
7.2
56.0
9

Caesium . . . .

1.88

133

From the figures in the last column it will be seen, that beginning
with lithium, 11.9, the atomic volume falls as the middle element of
the period, namely, carbon, is reached ; after which it again rises
and reaches a maximum with the first member of the second period,
namely, sodium. In this period the same gradual fall in atomic
volume is again noticed until the middle element (silicon) is
reached, when the value of this function of the elements once more
rises, and a second maximum is attained with the first member
(potassium) of the third period. The two next are long periods, and
the atomic volumes steadily decrease until the middle three (transi-
tional) elements, after which they gradually increase again to a
maximum in rubidium, the starting-point of the fourth period. In
the fourth period the same thing once more occurs, the minimum
atomic volumes being those of the middle or transition elements,
after which a maximum is again reached in caesium.

This periodicity of the atomic volumes may be graphically
represented by a curve, where the ordinates represent atomic
volumes and the abscissas atomic weights. This curve, which was
first constructed by Lothar Meyer, is known as Lothar Meyer's
curve (page 121), and a comparison of it with MendelejerFs table
is most instructive.

The divisions indicated by the Roman numerals correspond to
the different periods: Groups I. and II. being the two short periods,
III. and IV. the two complete long periods, while V., VI., and
VII. correspond to the fragmentary portions of the last three
periods.

The transitional elements of periods III., IV., and VI. are all to
be found at the minima of the large hollows ; separating the even
series (situated on the descending portion of the curve) from the
pdd series which lie on the ascending slope. The elements belong'

The Periodic Classification

121

..Si. —

122 Introductory Outlines

ing to the different groups in MendelejefPs table are seen to occupy
the same relative positions upon the different portions of this curve.

Thus .in Group I. the elements Li, Na, K, Rb, Cs, are all found
upon the maxima of the curve, and Cu, Ag, and Au at those points
at the minima where the electro-negative properties reappear. The
halogen elements (chlorine, bromine, iodine) are seen in similar
positions upon the ascending, and the alkaline earths (beryllium,
magnesium, calcium, strontium, barium) on the descending
portions.

When the periodic law was first formulated by Mendelejeff
(1869), there were a number of instances in which the system did
not harmonise with the then accepted atomic weights of the
elements. The discoverer boldly asserted that the atomic weights,
and not the system, were at fault, and in almost every such case
the careful reinvestigation of the atomic weights by numerous
chemists has proved the correctness of the assertion. One or two
instances may be quoted. The element indium had assigned to it
the atomic weight 76. Its combining proportion is 38, and being
regarded as a divalent element, its oxide was believed to have
the formula InO. Having an atomic weight = 76, indium would
occupy a place between As = 75 and Se = 79 ; but in the system
(see table on page 118) there is no room for an element with such
an atomic weight ; and, moreover,' if indium be a divalent element
having this atomic weight, it should come between Zn = 65 and
Sr = 87 in Group II., where again there is no room. Mendelejeff
made the assumption that the oxide of indium had the formula
In2O3, believing the element to be an analogue of aluminium
(Group III.). If this be the true composition of the oxide, the
atomic weight of the element would be 38 x 3 = 1 14, and indium
would then take its place in Group III., between the elements
cadmium = 112 and Sn = 118, in the odd series of the second long
period. Bunsen afterwards determined the specific heat of indium
by means of his ice calorimeter, and found it to be 0.057 : —

Mean atomic heat jM_ = = atomic wei ht (see 8).

Specific heat . - 0.057

Hence 114 and not 76 is the accepted (approximate) atomic weight
of indium.

Again, the element beryllium (formerly known as glucinum) has
a combining proportion of 4.6. Its chloride was believed to have
tjie composition BeClg, and its oxide to be a sescjuipxjde haying

The Periodic Classification 123

the formula Be2O3. The atomic weight assigned to the element,
therefore, was 13.8.

With this atomic weight beryllium would take its place between
carbon = 12 and nitrogen = 14 ; but according to the periodic
classification there is no room for such an element, and, moreover,
in such a position it would be among elements with which it has
no properties in common. On the supposition that the oxide of
beryllium has the formula BeO, that is, that the element is divalent,
its atomic weight would have to be lowered from 13.8 to 9. i in
order to maintain the same ratio between the weights of metal and
oxygen in the compound. On this assumption, beryllium would
fall into the second place in the first series, between lithium = 7
and boron = 1 i, and in the same group as magnesium and zinc.

When the specific heat of beryllium was determined, it gave the
value 0.45, and this number divided into the atomic heat constant,
6.4, gave 14 as the atomic weight. In spite of this evidence in
favour of the higher value as the atomic weight of beryllium,
MendelejefT still regarded the lower number as correct, and it
was suggested that possibly beryllium, like carbon and boron
(elements also of very low atomic weight), had an abnormally
low specific heat at ordinary temperatures. This was found to be
the case (see page 48), and at 500° the specific heat of beryllium
was found to be 0.6206. This divided into 6.4 gives the value 10
as the atomic weight, which indicates that 9.1 and not 13.8 is in
reality the atomic weight of beryllium.

Not only has the periodic law been of service in bringing about
the correction of a number of doubtful atomic weights, but by
means of it its originator was enabled to predict with considerable
certainty the existence of hitherto undiscovered elements, and
even to predicate many of the properties of these elements. As
already mentioned, at the time when the periodic law was first
formulated, there were three gaps in the system in the first long
period, namely, No. 4 in the even series (now occupied by scandium),
and Nos. 3 and 4 in the odd series (now filled by gallium and
germanium). To the unknown elements which were destined to
occupy these positions, Mendelejeff gave the names eka-boron,
eka-aluminium, and eka-silicon (the prefix eka being the Sanscrit
numeral one\ and from the known properties of the neighbouring
elements of the series (horizontal rows in the table, page 118), and
also of those situated nearest in the same family (vertical columns),
he predicted spine of the prominent properties that would pro-

124 Introductory Outlines

bably be possessed by these elements. Thus in the case of eka-
aluminium, from the known properties of aluminium and indium,
the neighbouring elements in the same family, and from zinc, the
contiguous element in the same series (the 4th place in the series
being unoccupied), Mendelejeff deduced the following properties
for the unknown element that he called eka-aluminium : —

PREDICTED PROPERTIES OF EKA-ALUMINIUM (1871).

(i.) Should have an atomic weight about 69.

(2.) Will have a low melting-point.

(3.) Its specific gravity should be about 5.9.

(4. ) Will not be acted upon by the air.

(5.) Will decompose water at a red heat.

(6.) Will give an oxide £72O3, a chloride £72C16, and sulphate £72(SO4)3.

(7.) Will form a potassium alum, which will probably be more soluble and
less easily crystallisable than the corresponding aluminium alum.

(8. ) The oxide should be more easily reducible to the metal than alumina.
The metal will probably be more volatile than aiuminium, and therefore its
discovery by means of the spectroscope may be expected.

In the year 1875 Lecoq de Boisbaudran discovered a new
element in a certain specimen of zinc blende (zinc sulphide), the
individuality of which he first recognised by the spectroscope,
the spectrum being characterised by a brilliant violet line. This
element he named gallium. The properties of this metal, as they
were subsequently observed, showed that it was, in fact, the pre-
dicted eka-aluminium of Mendelejeff, as will at once be seen by a
comparison of the following facts.

PROPERTIES OF GALLIUM (discovered 1875).

(i.) Atomic weight -=69. 9.
(2.) Melting-point, 30.15°.
(3.) Specific gravity, 5.93.
(4.) Only slightly oxidised at a red heat.
(5.) Decomposes water at high temperatures.

(6.) Gallium oxide, Ga2O3. Gallium chloride, Ga2Cl6. Gallium sulphate,
Ga.2(S04)3.

(7. ) Forms a well-defined alum:
(8.) Is easily obtained by the electrolysis of alkaline solutions.

In a similar manner the properties of eka-boron and eka-silicon
were predicted, and the subsequent discovery of scandium (Nilson,
1879), and germanium (Winkler, 1886), whose properties were
found to closely accord with these hypothetical elements, formed
an additional demonstration of the truth of the periodic law.

The Periodic Classification 125

There are at present two elements, however, which appear not
to conform strictly to this periodic classification. These are the
elements argon and tellurium. The atomic weight of argon
according to most recent determination is 39.88, while that of
potassium is 39.15. Now the periodic system requires that the
atomic weight of argon shall be below and not above that of
potassium. Again, the latest determinations of the atomic weight
of tellurium give 127.5, as against 126.92 for iodine ; while in
order to conform to the periodic system the atomic weight of
tellurium should be below that of iodine. Whether these two
cases will prove to be true exceptions, or whether future investiga-
tions will show that the atomic weights here given are not the
true ones, time alone will show. It must be borne in mind, how-
ever, that both, argon and tellurium are elements which it is
extremely difficult to obtain in a state of absolute purity, and there
is considerable probability that in the latter case the element in a
pure state has never yet been obtained.

The position which should be given to hydrogen in the periodic
system has been the subject of much discussion. It will be
noticed that in the table it is placed with a query in Group I. and
again in Group VII. ; its univalent character suiting either position
equally well. The chief argument in favour of placing it in
Group I. is its electro-positive character, in which it strongly re-
sembles the elements lithium, sodium, potassium, &c., metals
which may be substituted for hydrogen atom for atom ; the
"salts of hydrogen" (i.e. acids), and the metallic salts resembling
each other when regarded from the ionic standpoint.

The arguments in favour of assigning it a position at the head
of Group VII. are more numerous, and may be briefly summarised
as follows :* —

1. Its gaseous character and low boiling-point.

2. Absence of any metallic properties.

3. The diatomic nature of its molecules H2 (while many of the alkali metals
are monatomic).

4. Readiness with which H is substituted by Cl, Br, or I, in organic com-
pounds.

5. If placed in Group I. a series of six blank spaces is left, for as many un-
known elements, whose atomic weights must all fall between H = i and He =3.99,

6. The numerical difference between H = i and F=ig is 18 units, which is
much closer to the average of about 16 units than that between H = i and
Li=7, which is only 6 units.

* Masson, Chem. News, vol. Ixxiii. p. 283.

CHAPTER XIII
GENERAL PROPERTIES OF LIQUIDS

UNDER this head the following subjects will be considered :—

1. The passage of liquids into vapours or gases.

2. The passage of liquids into solids.

3. Solution.

1. The Passage of Liquids into Gases. Evaporation and
Boiling. — Just as in the gaseous condition, so in the liquid state,
the molecules are in a state of motion : in the liquid state, however,
the mean kinetic energy of the molecules is unable to overcome the
force of their mutual attraction. Some of the molecules have a
smaller kinetic energy (that is, a lower temperature), and others
a greater kinetic energy, than the average ; and when in the course
of their movements the latter strike the surface of the liquid and
break through it, they continue their movements in the space
above, as gaseous molecules. If the space into which they wander
be unlimited, that is, if the liquid be freely exposed to the air, these
molecules escape away altogether, and consequently the liquid
diminishes in quantity. This process is known as evaporation^
and as the molecules which so leave the liquid are those having
the highest temperature, it follows that the temperature of the
liquid, which is the average temperature of the molecules, will fall.
The more completely the molecules that so escape from the surface
of a liquid are prevented from falling back, that is, the more rapidly
they are swept away from the immediate neighbourhood of the
liquid, the more quickly will this escape of molecules take place,
and therefore the greater will be the fall of temperature that results
from evaporation. Thus, if a quantity of liquid, say water, be
exposed in a dish so that a current of air is blown across the sur-
face, the rate of evaporation is increased, and the temperature con-
sequently falls lower than if the water be merely placed in a still

atmosphere ; similarly, if the water be placed in a vacuum the rate

126

A B

Evaporation

of evaporation is increased, because the molecules that escape from
the surface of the liquid are not impeded in their motions by
collisions with the molecules of air.

This fall of temperature resulting from evaporation may be
readily seen by enveloping the bulb of a thermometer in a piece of
thin muslin, and moistening it with water. If such a thermometer
be placed by the side of a naked thermometer, it will be seen that
the mercury will fall lower in the one that is moistened, and the
difference will be still more
marked if the instruments
are placed in a draught,
whereby the evaporation ot
the water from the muslin
is accelerated.

If the space above the
liquid be limited, molecules
still continue to escape
from the surface; but a
state of equilibrium is soon
established, when as many
are thrown back again by
rebounding from one an-
other and from the walls
of the containing vessel as
leave the surface in a given
time. Under these con-
ditions the enclosed space ^g
is said to be saturated ivitJi
the vapour of the liquid.
The number of molecules "
which escape from the sur- Fie. 13.

face depends upon the tem-
perature, and is independent of the pressure, for if the volume
of a saturated vapour be forcibly diminished, it merely results
in the condensation of a portion of the vapour ; and if ex-
panded, a corresponding vaporisation of an additional quantity
of the liquid, the pressure remaining always constant. The num-
ber of molecules that re-enter the liquid is' determined by the
number and the velocity of those that exist as gaseous molecules
in a unit volume. But the pressure exerted by a gas is caused by
the number and velocity of the molecules in a given volume, hence

128 Introductory Outlines

the condition of equilibrium is set up, when the vapour above the
liquid exerts a definite pressure, which pressure will be constant
for any given temperature. The pressure exerted by a vapour
under these conditions is termed the vapour-tension of the liquid.
The fact that the vapour given off from a liquid exerts pressure
may readily be experimentally illustrated by means of the apparatus
seen in Fig. 13. Three glass tubes, A, B, and c, about one metre
long, are completely filled with mercury and inverted in a trough
of the same liquid. The mercury will sink to the same level in
each tube, the length of the mercury column representing the
atmospheric pressure at the time. Into two of these barometer
tubes, B and C, a few drops of water are introduced, when it will be
found that the mercury is depressed, as indicated in B, below the
level at which it previously stood. This depression of the mercury
column represents the tension of the vapour of the water for the
particular temperature at which the experiment is made. If tube
C be surrounded by a wider glass tube, through which steam from
a small boiler is passed, it will be noticed that as the temperature
of the water in the tube rises, the mercury is more and more de-
pressed, thus showing that the tension of the vapour increases with
rise of temperature. As soon as the steam circulates freely and is
escaping at the bottom of the wide tube, in other words, as soon
as the temperature of the enclosed water in tube C reaches 100°,
i.e. the temperature of the steam surrounding it, the mercury in
the tube will be depressed to the level of that in the trough. The
tension of the vapour within the tube, under these circumstances,
is therefore equal to the atmospheric pressure.

If, instead of introducing water into the barometer tube, ether
were employed, and a stream of vapour from boiling ether were
passed through the outer tube, it would be seen that when the ether
within the tube reached the temperature of the vapour from the
boiling ether, namely, 35°, the mercury would again be depressed
to the level of that in the trough ; that is, the tension of the ether
vapour would then be equal to the pressure of the atmosphere. We
see, therefore, that when water is heated to its boiling-point, viz.,
100°, the tension of its vapour is equal to the atmospheric pressure ;
and when ether is heated to its boiling-point, viz., 35°, the pressure
exerted by its vapour is equal to the pressure of the atmosphere.
The boiling-point of a liquid may therefore be defined as the
temperature at which the vapour-pressure is equal to the pressure
of the atmosphere. As soon as this point is passed, the kinetic

Boiling- Points of Liquids 129

energy of the molecules has been so much augmented by the
supply of external heat, that it is able to overcome the force of
their mutual attractions, and, consequently, the molecules freely
pass away from the surface of the liquid.

As will be seen from the illustrations given, namely, water and
ether, the temperatures at which the vapours of different liquids
exert a pressure equal to that of the atmosphere are widely different.
This fact will be still more evident from the following table, giving
the temperatures at which the vapour pressure of various liquids is
equal to the standard atmospheric pressure : —

Liquid hydrogen .... — 253°

Liquid oxygen . . . . . -182.5°
Liquid nitrous oxide . . . . — 89.8°

Liquid sulphur dioxide . . . - 10°

Ethyl chloride . . . . . + n°
Carbon disulphide . . . 47°

Water 100°

Aniline . . • . . « . 182°

Mercury 358°

Since the boiling-point of a liquid is that temperature at which
its vapour-tension is equal to the atmospheric pressure, it will be
evident that, if the latter increases or decreases, the temperature
necessary to produce an equal vapour-pressure must also rise or
fall; in other words, the boiling-point of a liquid is dependent upon
the pressure. If a quantity of water, no warmer than the hand, be
placed beneath the receiver of an air-pump, which is then quickly
exhausted, the water will be seen to enter into violent ebullition.
It does this when the pressure within the receiver is reduced to
the point at which it is equal to the tension of aqueous vapour at
the temperature taken.

For this reason water boils at a lower temperature in high
altitudes than at the sea-level ; and as the vapour-tension of water
at various temperatures has been experimentally determined, we
can, by ascertaining the boiling-point of water at any particular
altitude, calculate the atmospheric pressure, and consequently the
height above the sea-level.

Many liquids when heated, especially in glass vessels that have
been carefully cleansed, may be raised several degrees above the
boiling-point without ebullition taking place. The liquid under
these circumstances assumes a pulsating movement, which con-

130 Introductory Outlines

tinues for a short time, when a burst of vapour is suddenly evolved
with violence, and the temperature at once drops to the boiling-
point. The liquid then becomes quiescent, and again as the
temperature rises the pulsating movement begins, ending once
more in an explosive evolution of vapour. This succtissive boiling,
or bumping, is sometimes sufficiently violent to cause the fracture ot
the vessel. In order to experimentally ascertain the boiling-point
of a liquid, the thermometer, for this reason, is not immersed in
the liquid, but is suspended in the vapour, the temperature of
which remains constant throughout these irregularities in the
boiling.

Latent Heat of Vaporisation.— When a liquid is heated, its
temperature rises, as indicated by the thermometer, until a certain
point is reached (the boiling-point of the liquid), when the con-
tinued application of heat causes no further rise of temperature.
Thermometers placed in the liquid, and in the vapour, indicate the
same temperature and remain' constant, and all further applica-
tion of heat is unappreciated by these instruments, and disappears
in changing the liquid into vapour. The heat which in this way is
absorbed during the vaporisation of a liquid is spoken of as the
latent heat of vaporisation ; and the same amount of heat which
thus disappears during the conversion of a liquid into a vapour
is again rendered sensible when the vapour passes back into the
liquid state.

The heat which is thus said to become latent is in reality con-
verted into kinetic energy ; it is expended in imparting to the
molecules the kinetic energy necessary to overcome the attractive
forces operating between them while in the
liquid state ; in other words, it is doing the
work of overcoming cohesion (internal work),
and also the external pressure on the vapour
(external work).

In order that a liquid may pass into a vapour
it is necessary that heat be absorbed.' We
have seen (page 126) that a liquid undergoing
spontaneous evaporation becomes colder (that
FIG. 14. is> heat is absorbed by the molecules that are

converted into the gaseous state), and also
that the more rapidly the liquid can be made to pass into the
vaporous condition, without supplying external heat, the lower
will its temperature fall. Upon this fact depend a number

Latent Heat of Vaporisation

131

of methods for the artificial production of low degrees of cold.
For example, ether boils at 35°, but if a small quantity of ether
be placed in a glass flask standing upon a wooden block, upon
which a few drops of water have been poured, and a current of air
from a bellows be briskly blown through the ether (Fig. 14),
the temperature of the ether will fall so rapidly that in a few
moments the flask will be frozen to the block. By the rapid eva-
poration of liquids with lower boiling-points, the extreme degrees of
cold necessary for the liquefaction of such gases as oxygen, carbon
monoxide, air, &c., are obtained. Thus, liquid methyl chloride

FIG. 15.

boils at - 23° j by causing it to rapidly vaporise, its temperature
can be reduced to - 70°. Liquid ethylene in the same way falls
to a temperature of - 120°, and liquid oxygen by rapid evaporation
gives a temperature as low as -210°.

The temperature of water, in like manner, maybe so lowered by
its own rapid evaporation, as to cause it to freeze. We have already
seen that by reducing the pressure, the boiling-point of a liquid is
lowered ; if, therefore, a quantity of water be placed in a vacuum,

132

Introductory Outlines

and methods be adopted to remove the water vapour as rapidly
as it is formed, the water will enter into rapid ebullition. The
evaporation will therefore proceed so rapidly, and consequently
absorb heat so quickly, that the temperature of the boiling liquid
will quickly fall to o° when it passes into the solid state. The
instrument known as Carre's freezing machine depends upon this
principle. The water to be frozen is placed in the glass bottle C
(Fig. 15), which is in connection with a metal reservoir R, half
filled with strong sulphuric acid. This in its turn is connected by
b with an air-pump P, worked by the lever
M, to which is also attached a connecting
rod /, so that a stirrer within the reservoir is
kept constantly in motion. As soon as the
apparatus is exhausted to a pressure of two
or three millimetres, the water begins rapidly
to boil, and, as the sulphuric acid absorbs the
water vapour as rapidly as it is given off, the
temperature quickly falls and the water
freezes.

Fig. 16 illustrates another method by which
the same result may be obtained. A tall
glass vessel is exhausted by means of an
ordinary air-pump, and water is allowed
slowly to enter from a stoppered funnel, upon
the end of which is secured a short string.
At the same time strong sulphuric acid is ad-
mitted by the second funnel, and caused to
flow down a glass red, round which is wound
a spiral of asbestos thread. The acid at once
absorbs the aqueous vapour from the evapo-
rating water, the temperature of which, there-
fore, falls below the freezing-point, and it
solidifies as it flows over the string into the
form of an icicle.

Just as diminution in pressure lowers the boiling-point of a
liquid, so increased pressure raises the boiling-point. If water be
heated in a closed iron vessel, as in a high-pressure steam boiler,
the pressure caused by its own vapour raises the boiling-point
many degrees above 100°. There is a definite temperature, how-
ever, for every liquid, beyond which the liquid state is impossible,
whatever may be the pressure ; that is to say, the liquid when

FIG. 16.

Vapour-Pressures of Solutions 133

heated beyond this fixed point passes into the gaseous state, how-
ever great the pressure may be. This temperature is the critical
temperature (see page 79). If a liquid be heated in a sealed and
strong glass tube, as the critical temperature is approached the
surface of the liquid gradually becomes ill-defined, and finally the
tube is completely occupied by transparent vapour. On again
cooling, as soon as the critical point is passed the contents of
the tube again separate into two distinct layers consisting of liquid
and gas.

VapOUP-Pressures Of Solutions.— The boiling-point of a liquid
is modified by the presence in the liquid of dissolved substances.
If the substance in the solution be less volatile than the liquid, the
boiling-point is raised. Thus, while the boiling-point of pure
water (under the normal atmospheric pressure) is 100°, the tem-
perature at which saturated aqueous solutions of salts boil, is
considerably higher, thus : —

Containing Grammes of

Water Saturated with Salt in 100 Grammes Boiling-point,

of Water.

Sodium chloride . . . 41.2 108.4°

Potassium nitrate . . . 335.1 115-9°

Potassium carbonate . . 205.0 i33-o°

Calcium chloride . . . 325.0 179-5°

The temperature of the steam of these boiling solutions, as
ascertained by suspending a thermometer in the vapour, appears
to be the same as that from pure water, as the thermometer in
all cases indicates 100°. In reality, however, the temperature is
higher, although not so high as that of the boiling liquid. The
reason that the thermometer indicates ico° in all cases is because
the water vapour continually condenses upon the bulb of the
instrument, covering it with a film of pure water, which boiling
off from the bulb indicates only the boiling-point of the pure
liquid. By special arrangements this condensation may be pre-
vented, when it has been shown (Magnus) that the temperature
of the vapour, from such boiling solutions, rises as the solutions
become more concentrated — that is, as the temperature of the
boiling liquids rise. It has been already explained that the boil-
ing-point of a liquid is that temperature at which the vapour
tension is equal to the atmospheric pressure ; since, then, the
presence of dissolved substances raises the boiling-point, it follows

1 34 Introductory Outlines

that it must lower the vapour-pressure, for (in the case of aqueous
solutions) when the temperature has reached 100° the vapour-
pressure is still below that of the atmosphere, for the liquid does
not enter into ebullition at that temperature. Lowering the
vapour-pressure, therefore, is synonymous with raising the boiling-
point. The extent to which the vapour-pressure of a liquid is
lowered (or its boiling-point raised) by dissolving in 100 grammes
of it i gramme-molecule of a given substance is called the mole-
cular lowering of the vapour-pressure, or the molecular elevation
of the boiling-point of that liquid. Now it has been found with
substances which do not undergo ionic dissociation in the solvent
employed, and also which do not themselves exert any appreciable
vapour-pressure at the boiling-point of the solvent, that this mole-
cular lowering of the vapour-pressure is practically a constant.
Thus, for water the molecular rise of boiling-point is 5.2° ; while for
benzene it is 27.0°.

For example, given two substances, say glycerol and sugar,
which, when dissolved in water, yield solutions which are non-
electrolytes (i.e. these compounds do not dissociate), and are also
themselves practically non-volatile at the boiling-point of water ;
then, if I gramme-molecule of each be separately dissolved in
100 grammes of water, the two solutions obtained will be found to
boil at about 105.2° instead of 100°. Or again, two substances
fulfilling the same conditions when dissolved in benzene would
send up the boiling-point of this liquid from 80.5° to 107.5°.

If, on the other hand, the substance is an electrolyte- — that is,
one which undergoes ionic dissociation in the solvent, then the
effect produced by the same weight of substance is greater, since
the ions behave as though they were molecules, and the result is
the same as though a larger number of molecules were present in
the solution. Obviously the increase in the effect produced will
depend upon the extent to which dissociation takes place.

The following general laws relating to the effect of dissolved
substances upon vapour-pressure have been established :—

1. The relation between the quantity of a substance in solution
and the diminution of the vapour-pressure below that of the pure
solvent is the same at all temperatures.

2. The diminution of the vapour-pressure of a liquid, by a dis-
solved substance, is proportional to the amount of the substance in
solution (provided the substance itself exerts no appreciable vapour-
pressure at the temperature of the experiment).

Vapour-Pressures of Solutions 135

3. The molecular lowering of vapour-pressure by chemically
similar substances is constant; that is to say, solutions containing
one molecular weight in grammes (one gramme-molecule} of such
substances in equal volumes of the solvent, give rise to the same
dimin ution of vapour-pressure.

PIG. 17.

4. The relative lowering of vapour-pressure is proportional to
the ratio oj the number of molecules of the dissolved substance, to
the total number of molecules in the solution, i.e. the sum of the
number of molecules of the dissolved substance and of the solvent*

* Except in the case of electrolytes. See page 109.

136 Introductory Outlines

Upon these considerations it becomes possible, by means of the
lowering of the vapour-pressure, to determine the molecular weight
of a substance that is capable of being dissolved in a volatile
liquid.

The apparatus in which such a determination is made is shown
in dissected form in Fig. 17. A weighed quantity of the solvent to
be employed is contained in the tube A which is inserted in the
vessel B, which in its turn is placed upon the asbestos support D,
and heated from below by means of small flames. As the liquid
in A boils, its vapour is condensed by the condenser indicated at
C1} and thereby returned to the vessel. The outer vessel B also
contains a small quantity of the same liquid which boils simulta-
neously, so that the inner tube is thus surrounded by a jacket filled
with the hot vapour of the same liquid as is boiling inside. The
vapour from the boiling liquid in this jacket vessel is condensed by
the condenser at C2 and constantly returned. By means of a
thermometer the exact temperature at which the liquid boils is
thus ascertained, after which a weighed quantity of the substance
whose molecular weight is to be determined is introduced and the
boiling-point again ascertained.

The result is calculated by the formula —

When C = Constant— namely, the molecular elevation of the boil-

ing-point of the solvent used ;
g = The percentage strength of the solution ; and
R = The observed rise of boiling-point.

The Passage of Liquids into Solids.— Most liquids, when
cooled to some specific temperature, pass into the solid state ; the
temperature at which this change takes place is termed the solidi-
fying point. Generally speaking, the temperature at which a liquid
solidifies is the same as that at which the solid again melts ; but as
the solidification of a liquid is subject to disturbances from causes
that do not affect the melting-point, this is not always the case.
Thus, water may be cooled many degrees below o° if it be pre-
viously freed from dissolved air, and be kept perfectly still. This
super-cooling of water may readily be illustrated by means of the
apparatus represented in Fig. 18. This consists of a thermometer
whose bulb is enclosed in a larger bulb containing water, which
before the bulb is sealed at a, is briskly boiled to expel all the air.

Solidifying Points of Liquids 13?

When the instrument is immersed in a freezing mixture the tem-
perature of the water may be lowered to —15° without congela-
tion taking place, but on the slightest agitation it at once solidifies
and the temperature rises to o°. It is on account of this property
of water to suspend its solidification, that in determining
the lower fixed point of a thermometer, the temperature
of melting ice, and not that of freezing water, is made
use of.

Many other liquids exhibit suspended solidification to
a very high degree ; thus glycerine may be cooled to — 30°
or - 40° without solidifying, but if a crystal of solid gly-
cerine be placed in the liquid the entire mass freezes, and
does not again melt until a temperature of 15.5° -is
reached. ||20

Change of Volume on Solidification.— Most liquids,
in the act of solidifying, contract ; that is to say, the solid
occupies a smaller volume than the liquid. Consequently
the solid is specifically denser, and sinks in the liquid.
Thus 100 volumes of liquid phosphorus at 44° (the melting-
point) when solidified occupy only 96.7 volumes. Water
expands upon solidification, hence ice is relatively lighter
than water, and floats upon the liquid. The reverse
change of volume accompanies the change of state in the
opposite direction.

Effect of Pressure upon the Solidifying- Point of
Liquids. — In the case of liquids that contract upon soli-
dification, increased pressure raises the point of solidifi-
cation, and consequently raises the melting-point of the
solid. The effect, however, is extremely small : thus the
solidifying-point (and melting-point) of spermaceti under
the standard atmospheric pressure is 47-7°, while under p g
a pressure of 156 atmospheres it is raised to 50.9°.

With liquids that expand on solidification, increased pressure has
the opposite effect, and lowers the solidifying point. Thus, water
under great pressure may be cooled below o° and still remain liquid ;
and in the same way ice may be liquefied by increased pressure
without altering its temperature. In the case of water it has been
found that an increased pressure of n atmospheres lowers the soli-
difying point by 0.0074^° ; hence under a pressure of 135 atmos-
pheres, the freezing-point of water (and the melting-point of ice) is
lowered i°. This lowering of the melting-point of ice under pres-

Introductory Outlines

sure may be illustrated by the experiment represented in Fig. 19.
Over a block of ice is slung a fine steel wire, to which are hung a
number of weights. The pressure thus exerted upon the ice, by
lowering the melting-point, causes the ice to liquefy immediately
beneath the wire, which therefore gradually cuts its way through
the block. But as the wire passes through the mass, each layer of
water behind it again resolidifies, being no longer subject to the
increased pressure ; hence, although the wire cuts its way com-
pletely through the ice, the block still remains intact.
Latent Heat Of Fusion.— When a liquid, at a temperature
above its solidifying point, is cooled,
a thermometer placed in the liquid
indicates its loss of heat until solidi-
fication begins. At this point the
temperature remains constant until
solidification is complete, when the
thermometer again begins to fall.
And again, when a solid, at a tem-
perature below its melting-point, is
heated, its temperature rises until
the melting begins, but no further
rise of temperature takes place by the
application of heat until liquefaction
is complete. The sensible heat that
so disappears during fusion is spoken
of as the latent heat of fusion. Just
as in the passage of liquids into
gases, this so-called latent heat re-
presents heat that has ceased to be
heat, but which is converted into
kinetic energy that is taken up by the
molecules : when the liquid passes
back into the solid state, this energy is again transformed into
sensible heat.

The fact that heat is thus changed into energy, and so rendered
insensible to the thermometer, may be seen by adding boiling water
to powdered ice. A thermometer placed in ice indicates the tem-
perature o°, and although boiling water is poured upon it, so long
as any ice remains unmelted no rise of temperature of the mixture
results, the heat contained in the boiling water being expended in
doing the work of liquefying the ice, and converting it into water at o°.

FIG. 19.

Solidifying Points of Liquids 139

When such an experiment is made more exactly, it is found that
i kilogramme of water at 80.25°, when mixed with I kilogramme of
ice at o°, gives 2 kilogrammes of water at o°. That is to say, the
amount of heat contained in a kilogramme of water at 80.25° ls
exactly capable of transforming an equal weight of ice at o° into
water at o°.

As the heat required to raise the temperature of i kilogramme of
water from o° to i° is the unit of heat, or major calorie, we say that
the latent heat effusion of ice is 80.25 thermal units or calories.

During the solidification of a liquid, the 'latent heat of fusion is
again given out. The solidification, therefore, only takes place
gradually, for the heat evolved by the congelation of one portion
is taken up by the neighbouring particles, whose solidification is
thereby retarded until this heat is dissipated. In the case of super-
cooled liquids and super-saturated saline solutions, the solidifica-
tion takes place more suddenly, and the evolution of the latent heat
is therefore manifest by a rise of temperature.

Effect of Substances in Solution upon the Solidifying Point
Of a Liquid. — It has long been known that a lower degree of cold
is necessary to freeze salt water than fresh ; and also that the water
obtained by remelting ice from frozen sea-water is so little salt. as
to be drinkable. Quantitative experiments show that water con-
taining i per cent, of common salt requires to be cooled to -0.6°
before the water begins to freeze ; and, moreover, that when such
a dilute solution begins to freeze, the solid which separates out is
not the salt, but is pure ice. This also holds in the case of all
other solvents that are capable of being solidified, the pure solidi-
fied solvent alone separating when the solution is frozen. For
instance, benzene freezes at 6° ; but if a small quantity of any sub-
stance which it is capable of dissolving be added (either a solid 01
liquid substance), it will be found necessary to cool the liquid below
6° before the benzene begins to freeze. The effect of dissolved
substances in lowering the solidifying point of the solvent was first
discovered by Blagden (1788), who formulated the law that the
depression of the freezing-point of aqueous solutions of the same
substance was proportional to the strength of the solution. By
referring the lowering of the solidifying point to quantities of the
dissolved substances that were in molecular proportions, instead
of to equal weights, it has been found that in the case of certain
chemically allied substances the following general law holds good :
Solutions containing in equal volumes of the solvent quantities oj

140 Introductory Outlines

dissolved substances proportional to their molecular weights have
the same point of solidification.

Thus, centi-normal solutions of sodium chloride and potassium
chloride (i.e. solutions containing 0.585 gramme NaCl and 0.746
gramme KC1 respectively in one litre of water) will begin to freeze at
the same fraction of a degree below o°. In other words, the depres-
sion of the freezing-point of the solvent is a function of the number
of molecules of the dissolved substance, irrespective of the nature
of the molecules. The extent to which the freezing-point of a
liquid would be depressed* by dissolving in 100 grammes of it one
gramme-molecule of any substance is called the molecular depres-
sion of the freezing-point of that liquid, and it is found that in the
case of all substances which are non-electrolytes, i.e. which do not
undergo ionisation, this molecular depression for a given liquid is
practically a constant. Thus in the case of water, when the sub-
stance dissolved is a non-electrolyte, the molecular depression is
about 18.5°.

In the case of substances which dissociate into their ions in the
solution, the molecular depression will be greater, depending upon
the degree of ionisation. Thus in the case of strong acids, bases,
and salts, that is, " electrolytes " which undergo dissociation to the
highest degree, it is found that the molecular depression is practi-
cally double that given by non-electrolytes. The ions in the liquid
acting as independent molecules, it will be obvious that if dissocia-
tion is complete there will be twice as many ions as there were
molecules of the compound, and therefore the effect produced in
respect of lowering the freezing-point should be twice as great.
The relations thus established between the molecular weight of a
compound and its influence in lowering the freezing-point of a
solvent form the basis of a method for the determination of mole-
cular weights (Raoult's method).

The process is carried out in a tube quite similar to tube A, Fig.
17 (the side tube in this case being merely closed with a cork).
A weighed quantity of the solvent is introduced into this tube,
which is then carefully cooled in a freezing-mixture, the liquid
being gently stirred by means of a wire passing through a hole in
the top cork. The temperature at which freezing begins to take
place is noted. The tube is then withdrawn from the freezing-

* In actually determining depressions of freezing-point, solutions so strong
as this cannot be used. The determination is made with dilute solutions, and
the molecular depression obtained by calculation.

Solidifying Points of Liquids 141

mixture and the solidified portion allowed to melt, when a weighed
quantity of the substance whose molecular weight is to be deter-
mined is introduced, and the operation repeated. The molecular
depression is calculated from the formula —

where C= constant — the molecular depression of the freezing-point;
g= grammes of substances in 100 grammes of the solvent ;
and /=the observed depression of the freezing-point

CHAPTER XIV
SOLUTION

A SOLUTION may be defined as a homogeneous mixture of either a
gas, a liquid, or a solid with a liquid, this liquid being termed the
solvent*

Substances that are capable of forming such homogeneous mix-
tures with a solvent are said to be soluble in that liquid. The
solution of matter in its three states will be treated separately.

1. Solution Of Gases in Liquids.— When a gas is dissolved
by a liquid, the liquid is said to absorb the gas, and although it is
held that most liquids are capable of absorbing most gases to a
greater or less degree, most of the investigations in this direction
have been made with the two liquids, water and alcohol, by Bunsen.

The quantity of a gas which a liquid is capable of absorbing
depends upon four factors — (i) the specific nature of the liquid ;
(2) the nature of the gas ; (3) the temperature of the liquid j (4)
the pressure.

(i.) The influence of the solvent may be seen by a comparison
of the quantities of the same gas which equal volumes of water and
of alcohol are capable of dissolving, thus —

ioo volumes of water at o° dissolve 179.6 volumes of carbon dioxide,
while loo ,, alcohol ,, 432.9 ,, ,,

(2.) The various quantities of different gases which the same
liquid will absorb are found to extend over a very wide range,
thus —

ioo vojumes of water at o° dissolve 4. 114 volumes of oxygen,

while ioo ,, ,, ,, 114800.0 ,, ammonia.

* Mixtures of gases are sometimes regarded as solutions, one gas being said
to be dissolved in the other. Gases also are sometimes spoken of as dissolving
liquids and solids, when liquid and solid substances directly vaporise into
them.

143

Henry's Law

143

(3.) The volume of any gas which a liquid can absorb diminishes
with a rise of temperature.* This will be seen from the following
table, where the volumes of different gases are given which 100
volumes of water will absorb at various temperatures.

Temperature.

Carbon Dioxide.

Nitrous Oxide.

Oxygen.

Nitrogen.

179.6

130.5

4.11

2.03

144.9

109-3

3-62

1.79

' 10

118.4

91.9

3-25

1.60

90.1

67.0

2.83

1.40

It was at one time believed that the solvent power of water for
hydrogen was the same at all temperatures between o° and 25°.
Recent experiments have shown, however, that there is no excep-
tion to the general law in this case ; thus it has been found that
100 volumes of water —

At o° dissolve 2. 1 5 volumes of hydrogen.
At 5° „ 2.06 „ „

At 10° „ 1.98 „ „

At 20° „ 1.84 „ „

When a solution of a gas in water is heated, the gas being less
soluble at the higher temperature is expelled, and in most cases
the whole of the gas is driven off at the boiling temperature.
This, however, is not invariably the case ; for example, the solution
of hydrochloric acid in water, when boiled, will distil, without further
evolution of gas, when a solution of definite strength is reached
(see Hydrochloric Acid).

(4.) The influence of pressure upon the volume of a given gas
which a liquid can absorb was discovered by Henry (1803), and is
known as Henry's law, namely, The volume of the gas absorbed by
a liquid is directly proportional to the pressure of the gas. If the
pressure be doubled, the same volume of liquid will dissolve twice
the volume of the gas, the volume in each case being measured
at o° and 760 mm. But since, according to Boyle's law, the
volume of a gas is inversely as the pressure, this law may be thus
stated : A given volume of a liquid will absorb the same volume of
a gas at all pressures.

* Helium, between certain limits of temperature, is an exception.

144 Introductory Outlines

Thus, if 100 volumes of water at o° dissolve 2.03 volumes of
nitrogen, under the standard atmospheric pressure (the volume of
the gas being measured at o° and 760 mm.), under twice this
pressure, i.e. two atmospheres, the same volume will absorb twice
the volume of nitrogen, viz., 4.06 volumes measured at o° and 760
mm. But 4.06 volumes of gas measured at o° and 760 mm. occupy
2.03 volumes under a pressure of two atmospheres, therefore the
liquid dissolves the same volume of compressed gas as of gas
under ordinary pressure.

Henry's law is sometimes stated in a slightly altered form. If
the quantity of gas present in a unit volume of both the liquid and
the space above it be called the concentration of the gas, then the
law may be expressed by saying that under all pressures, the ratio
of the concentrations of the gas in the liquid, and in the space above
it) remains constant. This ratio is termed the coefficient of solu-
bility, or the " solubility " of the gas in the particular liquid.

The term coefficient of absorption, first introduced by Bunsen, is
the volume of the gas measured at o° and 760 rnm., which is
absorbed by I cubic centimetre of a liquid at the same tem-
perature and pressure ; and it is therefore simply the volume
representing the " solubility " of the gas, reduced to o°.

The solubility of gases in liquids is measured by agitating a
known volume of liquid with a measured volume of the gas, under
determinate conditions of temperature and pressure. The apparatus
employed by Bunsen, and known as Bunsen's absorptiometer, is
shown in Fig. 20. It consists of a graduated tube e, into which
known volumes of the gas and liquid are introduced. The lower
end of this tube is furnished with an iron screw, by means of
which it can be securely screwed down upon an indiarubber pad,
in order to completely close the tube (seen in the side figure).
The tube containing the gas and liquid under examination is
lowered into a tall cylinder g g, in the bottom of which is a
quantity of mercury. The cylinder is then filled with water, and
the cap p screwed down. The thermometer k registers the tem-
perature. The apparatus is then briskly shaken, in order that the
liquid in the eudiometer may exert its full solvent action upon the
gas, and on slightly unscrewing the tube from the caoutchouc pad,
mercury enters to take the place of the dissolved gas. The tube
is again closed and the shaking repeated, and these operations are
continued until no further absorption results. Finally, the volume
pf gas is measured, the temperature noted, and the pressure

FIG. 20.

146

Introductory Outlines

ascertained by reading the position of the mercury within the tube,
and deducting the height of the column from b to the surface of
the mercury a, from the barometric pressure at the time of making
the experiment. The temperature of the water in the cylinder
may be varied, and the coefficient of absorption at different tem-
peratures can thus be determined.

Fig. 21 represents a more modern absorptiometer, being a modi-
fied form of Heidenhain and Meyer's apparatus. In this instru-
ment the measuring tube and the absorption vessel are separate,
and it admits of the use of much larger volumes of liquid. By
means of the three-way cock a, the gas to be experimented upon
is introduced into A by first raising and
then lowering B ; and the volume is
measured when the levels of the mer-
cury in A and B are coincident. By
means of the three-way cock <£, the
vessel C, of known capacity, and which
is connected with A by means of a
flexible metal capillary tube, is filled
with the desired liquid. The vessels A
and C are then put into communica-
tion, and by raising B and opening the
tap c a definite volume of the liquid is
run out into a measuring vessel, which
represents the volume of gas that enters.
The gas and liquid are then thoroughly
agitated, after which the gas is passed
back into A by lowering B^ and, when
A and C are in communication, open-
ing the tap c beneath mercury. By
measuring the diminution in volume

suffered by the gas, the volume absorbed by the known volume
of liquid is obtained. The measuring tube and absorption vessel
are kept constant at any desired temperature by surrounding
them by water, or with vapours at known temperatures.

Solubility Of Mixed Gases.— When two gases are mixed
together, the . pressure exerted by each is the same as would be
exerted if the other were absent and the entire space were
occupied by the same mass of the one. Thus, if a mixture of
two gases are in the proportion of two volumes of one and one
volume of the other, the pressure exerted by the one present in

FIG. 21.

The Law of Partial Pressures 147

larger proportion will be twice as great as that of the other ; this
pressure is termed the partial pressure of the gas under the
circumstances, and obviously the total pressure of the mixture
will be the sum of the partial pressures of the constituents. As the
solubility of a gas in a liquid is proportional to the pressure, the
solubility of the gases in a gaseous mixture'will be influenced by
the proportions in which they are present in the mixture. This
is known as Dalton's law of partial pressures, which may be thus
stated : The solubility of a gas in a gaseous mixture is proportional
to its partial pressure. For example, the atmosphere consists of
a mixture of oxygen and nitrogen, in the proportion of four volumes
of nitrogen to one volume of oxygen (in round numbers). The
partial pressure exerted by the oxygen is therefore only one-fifth of
the total atmospheric pressure, and consequently the amount of
oxygen which a given volume of a liquid is capable of dissolving
from the atmosphere is only about one-fifth of that which it will
absorb from pure oxygen — in other words, will be one-fifth the
absorption coefficient of oxygen for that liquid.

The application of the law of partial pressures will be seen in
the solvent action of water upon the atmosphere. Taking the
coefficients of absorption of oxygen and nitrogen for water as
given by Bunsen —

Oxygen = .04114 ; Nitrogen = .02035,

and the proportion of oxygen to nitrogen in the air as one to four,
by volume, we get —

= . 00823, and 'O2°^ X 4 = .01628,

for the number of cubic centimetres of oxygen and nitrogen which
will be dissolved from the atmosphere by i cubic centimetre of
water at o°.

One hundred volumes of water, therefore, will dissolve 2.451
volumes of air, of which .823 volume is oxygen and 1.628 volumes
is nitrogen ; and if this dissolved air be again expelled from the
water by boiling, the air so obtained will contain oxygen and
nitrogen in the proportions —

Oxygen . . . V,, 33-6
Nitrogen ... • 66.4

100.0

148 Introductory Outlines

If a mixture of oxygen and nitrogen in this proportion be once
more dissolved in water, since the percentage of oxygen has risen
from 20 to 33.6, and the partial pressure proportionately increased,
the mixture of the two gases that will be dissolved will be still
richer in oxygen ; and after solution in water for the third time the
boiled-out air will be found to contain as much as 75 per cent,
of oxygen. It will be obvious that the partial pressure which de-
termines the extent to which the separate gases in a mixture are
dissolved is not represented by the proportion in which the gases
are present before solution, but that in which they exist in the
gaseous mixture after the solvent has become saturated.

Henry's law does not hold good in the case of such very soluble
gases as ammonia, hydrochloric acid, &c. These gases appear to
enter into a true chemical union with the water, and in most of
these cases the act of solution is attended with considerable evolu-
tion of heat. In some of these instances the deviation from the
law diminishes with rise of temperature ; thus at temperatures
above 40° the absorption of sulphur dioxide obeys the law, while
in the case of ammonia conformity to the law is observed at 100°.

The gases dissolved by a liquid are not only expelled by boiling,
but are withdrawn by placing the solution in a vacuum. This, in-
deed, follows from Henry's law, for if the solubility is proportional
to the pressure, and the pressure is nil, the amount of gas dissolved
must also be nil.

The molecules of gas dissolved by a liquid are regarded as being
held by some attractive forces exerted between them and the mole-
cules of the liquid ; in the course of their movements, gas molecules
are constantly leaving and entering the liquid, and equilibrium is
established when the same number enter and escape from the
surface of the liquid in the same time. When the pressure is in-
creased, more gas molecules strike the surface in a unit of time, and
consequently a greater volume is absorbed. When a solution of a
soluble gas is placed in an atmosphere of another gas, the dissolved
gas continues to leave the liquid until equilibrium is established
between the pressure exerted by the gas so leaving and the amount
remaining in solution. For this reason a solution of ammonia,
when left exposed to the air, rapidly becomes weaker, owing to
the escape of the dissolved gas into the atmosphere. This process
is accelerated if a stream of a less soluble gas be caused to bubble
through the solution.

Solubility of Liquids in Liquids.— The solubility of liquids in

Solution 149

liquids may be divided into two orders. First, cases in which the
degree of solubility of one in the other is unlimited ; and second,
cases where the extent of the solubility is limited, or where the
liquids are said to be partially miscible. Two liquids whose
solubility in each other is unlimited are said to be miscible in all
proportions ; thus alcohol and water are capable of forming a
homogeneous mixture when added together in any proportion.

In the second class, where the solubility of two liquids for each
other is limited, it is found that each liquid is capable of dissolving
some of the other. Thus, if equal volumes of ether and water are
shaken together, the liquids will afterwards separate out into two
distinct layers, one floating upon the other, " The heavier layer at
the bottom is an aqueous solution of ether, containing about 10 per
cent, of ether ; while the upper liquid is an ethereal solution of water
containing about 3 per cent, of water. The presence of ether
dissolved in the water may be proved by separating the two layers
and gently heating the aqueous liquid in a small flask, when the
dissolved ether will be expelled and can be inflamed. The pre-
sence of the water in the ether is also readily proved, either by
introducing into the liquid a small quantity of dehydrated copper
sulphate, which will rehydrate itself at the expense of the water in
the ether, and be changed from white to blue ; or by placing in the
ethereal liquid a fragment of sodium, which decomposes the dis-
solved water with the liberation of hydrogen.

Another illustration of two partially miscible liquids is seen in
the case of a strong aqueous solution of potassium carbonate
and strong ammonia, which is of special interest as being the
only example at present known of two aqueous solutions of in-
organic substances which exhibit this phenomenon.* Thus, when
strong aqueous ammonia (sp. gr. 0.880) is added to a concen-
trated solution of potassium carbonate, the two liquids separate
from each other in two distinct layers, the upper layer consisting
of ammonia which has taken up a certain amount of potassium
carbonate, while the lower liquid consists of a solution of potassium
carbonate which has dissolved a definite quantity of ammonia.

In most cases the solubility of liquids in liquids is increased by
rise of temperature, although in some it is decreased. As an
example of the former, the case of these two aqueous liquids may
be quoted. If the temperature be raised then the solubility oi
each of these solutions in the other steadily increases, and the

* Newth, Trans. Chem. Soc. , 1900, p. 775.

150 Introductory Outlines

composition of the two layers will therefore gradually approximate
until a point is reached at which they become identical. This
point is arrived at when the temperature reaches about 43°, and at
this temperature, therefore, the two liquids are miscible in all
proportions. If this liquid be now cooled below this temperature,
separation into the two phases, as it is termed, at once begins, and
the liquid gradually becomes milky or turbid owing to the pre-
cipitation from it of the heavier solution in minute drops.

An instance of decreased solubility by rise of temperature is
seen in the case of a mixture of triethylamine and water. If equal
volumes of these liquids be mixed together, at a temperature below
20°, complete solution takes place, and a single homogeneous
liquid results. On warming the solution it becomes turbid, owing
to the separation of the liquid into two portions, which ultimately
settle out as two distinct layers. As the temperature of the solu-
tion approaches 20°, the liquid becomes very sensitive to a slight
rise of temperature, the heat of the hand being sufficient to cause
turbidity in the solution.

It will be evident, therefore, from these considerations that the
distinction between liquids which are miscible in all proportions
and those which are only partially miscible is after all only an
arbitrary one, the difference being simply a function of the tem-
perature. It is, nevertheless, a convenient distinction to make, so
long as we understand that it refers to liquids at the ordinary
temperature.

Solution Of Solids in Liquids.— When a solid is immersed in
a liquid, the forces which oppose the solution of the solid are the
attractive forces exerted by the molecules of the solid upon each
other and those of the liquid upon themselves. The forces that
tend to effect solution are the attractive forces exerted by the
molecules of the liquid upon the molecules of the solid, and the
kinetic energy of the molecules.

By the action of the liquid, the attractive force between the mole-
cules of the solid is diminished, and those molecules nearest the
surface, by their own energy and the attraction exerted by the
liquid, pass into and through the liquid. In the course of their
movements, these sometimes return to the solid, and a condition
of equilibrium is finally established when as many molecules leave
the surface of the solid as return to it in a given time. Under these
circumstances the solution is said to be satyrated with respect tp
the particular solid.

Solution 151

Saturated Solutions.— The amount of solid held in solution by
the liquid when the latter is saturated depends upon the tempera-
ture, for if the temperature be raised, the kinetic energy of the
molecules is increased, and consequently an increased number will
become detached from the solid. As a general rule, therefore, the
solubility of a solid in a liquid is increased by rise of temperature.
A saturated solution at a given temperature may be obtained in
two ways, namely, by maintaining the liquid at that temperature
and stirring into it an excess of the solid, until no more of it is dis-
solved ; or by dissolving a larger quantity of the solid at a higher
temperature, and allowing the solution to stand in contact with an
excess of undissolved solid, until the temperature falls to the specified
point. During the cooling the amount of solid that the liquid had
taken up, over and above that which was necessary to saturation
at the lower temperature, is deposited.

Supersaturated Solutions.— The condition of saturation can
only be determined when an excess of the undissolved solid is
present in the liquid ; for when a solution, which is not in contact
with any of the undissolved solid, is brought to the point of satura-
tion, either, by cooling or by evaporation of the liquid, it frequently
happens that no separation of solid takes place. Solutions can in
this way be obtained, in which a larger amount of the solid remains
dissolved at' a given temperature than corresponds to the amount
required to form a saturated solution at that temperature : such
solutions are said to be supersaturated. If into such a supersatu-
rated solution a fragment of the solid be introduced, molecules of
the dissolved solid at once deposit themselves upon it, and this
separation of the dissolved substance continues until the solution
reaches a state of concentration corresponding to its normal satura-
tion at the particular temperature. The introduction into a super-
saturated solution of a particle of the solid, in respect to which the
solution is supersaturated, is the only sure method of bringing
about the separation of the excess of the dissolved substance ; such
a solution, therefore, may be preserved for an indefinite time, if it be
kept in an hermetically sealed vessel. Minute particles of the solid
towards which a solution is supersaturated, that might be present
in the dust of the air, falling into such a solution, will determine
the deposition of the dissolved solid.

The phenomenon of supersaturation is strictly analogous to that
of supercooling, or the suspended solidification of fused solids, and
is exhibited most readily by salts containing water of crvstallisa-

152

Introductory Outlines

tion, such as sodium acetate, NaC2H3O2,3H2O ; sodium thiosul-
phate, Na2S2O355H2O ; and sodium sulphate, Na2SO4,10H2O.
Thus, if a small quantity of water be poured into a flask nearly
filled with crystallised sodium thiosulphate (the so-called " Hypo "
of the photographer), and the mixture be warmed by immersion in

250

O* 20' 20" &0" 40' 50° 63°. 7O° QO° 90' 2OO*

Temp erature .
FIG. 22.

hot water, the whole of the salt will dissolve ; and if the solution be
then allowed to cool undisturbed, it will assume the ordinary tem-
perature, and still remain fluid. If into the supersaturated solution
a crystal of the salt be dropped, the excess of salt present in solution

Solution 153

beyond the normal quantity for saturation at that temperature will
crystallise out, and so great is this excess that the contents of the
flask will appear practically solid.

The different solubility of various solids in the same liquid, and
the increased solubility by rise of temperature, is graphically shown
in Fig. 22, where the solubility curves of five salts in water are
represented. The abscissas indicate temperatures, and the ordi-
nates the number of parts of salt dissolved by loo parts of water.

Thus at o° 1 80 grammes of water will dissolve 35.7 parts of
sodium chloride, and as the temperature is raised the quantity of
salt which the water will dissolve very slowly increases, until at 100°
the amount is nearly 40 parts. : sodium chloride is therefore nearly
equally soluble in water at all temperatures.

In the case of potassium nitrate, 100 grammes of water at o° will
only dissolve 13.3 grammes of the solid, but as the temperature
rises the amount capable of being dissolved by this quantity of
water very rapidly increases, until at 75° 150 grammes are dissolved.
Lead nitrate is more soluble than potassium nitrate between o° and
50°, but above this point it is not so soluble as the other, hence the
two curves intersect at that temperature. The solubility of sodium
sulphate in water appears at first sight lo be anomalous. The
solubility at first rapidly increases with rise of temperature from
o°, and reaches a maximum at a point between 33° and 34°, when
it gradually diminishes with further rise of temperature. This
behaviour is in reality due to the fact that we are not dealing with
one and the same substance throughout the experiment Sodium
sulphate exists as a solid in at least three forms, namely, the
decahydrate, Na2SO4,10H2O (ordinary Glauber's salt) ; the hepta-
hydrate, Na2SO4,7H2O ; and the anhydrous salt, Na2SO4. The
first portion of the curve (Fig. 23) represents the solubility of
Glauber's salt ; thus, at 20° such an amount of this decahydrated
salt is dissolved, that the solution contains 2o*grammes of Na2SO4
in 100 grammes of water. The solubility of this salt rapidly rises
until 34° is reached, at which temperature the salt melts, and is
then miscible with water in all proportions. The melted salt con-
tains 78.8 parts of Na2SO4 in 100 parts of water, which is indicated
as the highest point upon its curve : —

Na2S04. IOH.,0.

23 + 23 + 32 + 64 (2 + 16) x 10

•" v • ' 1 80 : 100 = 142 : 78.5.

142 |3Q

154

Introductory Outlines

The decahydrated salt is unable to exist as such at temperatures
higher than 34°, and when the melted salt is heated above this
point it is converted into the anhydrous salt and water satu-
rated with the salt ; therefore above 34° it is not possible to have
a solution of sodium sulphate in contact with' solid Glauber's
salt. It can, however, be in contact with the anhydrous salt, and
the second portion of the curve expresses the solubility of this com-
pound in water, which slowly diminishes as the temperature rises.

§
f 80

!„

7*1

% »

•s

I-
>»

JS J0

•*^.

"•• —

*•• —

— —

• •

•« •

Te mp e r a/ tur e .

FIG. 23.

Paradoxical as it may at first appear, it is possible, by gradually
cooling solutions of a salt in water, to cause them to become either
more concentrated or more dilute according to circumstances. It
has been already explained (page 139) that when a dilute solution
of a salt in water is cooled below o°, ice only separates out.
Obviously, therefore, the solution that remains is more con-
centrated than at first, and its freezing-point will consequently be
lowered. If the cooling be continued, more and more ice separates
out, and the remaining liquid becomes gradually more and more
concentrated until at length a point is reached when the solution
is saturated for that particular temperature. If cooled below this
point ice still separates, but as the solution would then be super

Osmotic Pressure 155

saturated, salt also separates out ; and the composition of the
mixture of ice and salt which thus separates is the same as that of
the remaining solution ; in other words, the solution freezes as
though it were a pure chemical compound of the water and the
salt in solution.* Such a solution is known as a constant-freezing
solution, or sometimes a cryohydric solution, and is comparable
with such constant-boiling mixtures as are obtained by distilling
either nitric or hydrochloric acids (see pages 239 and 367).

If now, instead of starting with a dilute solution, a concentrated
solution is gradually cooled, at some particular temperature
(depending upon the degree of concentration at first) the solu-
tion will become saturated for that temperature. Further cooling
below this point will then cause the solution to deposit some of the
salt ; and the liquid, although still a saturated one as respects this
lower temperature, will be more dilute. As the cooling continues,
the separation of the salt continues, and the solution therefore
becomes more and more dilute (still remaining saturated for the
lower temperatures) until the point is reached when the solution is
of such a strength that any further separation of salt (i.e. dilution)
would yield a liquid which is below its own freezing-point. That
is to say, the water itself now begins to freeze and separate along
with the salt, and at this point the solution has reached the same
constant-freezing condition as in the former case.

Osmotic Pressure.— When a dilute solution of a substance in
water is placed in a vessel closed with an animal membrane, such
as bladder (M, Fig. 24), and the whole is immersed in water to
such a depth that the level of the water outside is coincident with
that of the solution within, it is found that the liquid in the inner
vessel increases in volume, as seen by the fact that it gradually
rises in the narrow stem of the apparatus. Water, therefore, from
the outer vessel must have passed in through the membrane, and
inasmuch as some of the dissolved substance is found in the water
of the outer vessel, some of the solution must at the same time
have made its escape through the membrane. After the liquid has
risen to a certain height in the narrow tube, it again begins to
fall, as the contained solution continues to penetrate the mem-
brane. This process is known as endosmose, and the instrument
described is called an endosmometer.

* At one time, indeed, such solutions* were believed to contain definite
chemical compounds of the salt with water, which were called cryohydrati$
(Guthrie),

156

Introductory Outlines

Many attempts were made to establish general relations between
the height to which the liquid rose in the narrow tube and -the
quantities of substance in the solution, but it was found impossible
to obtain accurate or comparable measurements, for not only were
the results disturbed by the effect of the
constantly changing pressure upon the rate
at which the dissolved substance escaped
through the membrane, but different ani-
mal membranes yielded different results.

Semipermeable Membranes.— It was
first discovered by Traube (1867), and
afterwards extended by Pfeffer (1877), that
artificial membranes, or pellicles, could be
obtained, which, while allowing of the pass-
age of water through them just as in the
case of animal membranes, unlike these
materials, they offered a perfect barrier to
the passage of many substances in solu-
tion in the water. Such pellicles are
known as semipermeable membranes.
The material that has been found most
suitable is precipitated copper ferro-
cyanide. If a solution of copper sulphate
(CuSO4) be brought cautiously in contact
with a solution of potassium ferrocyanide
(K4Fe(CN)6), at the point where the two
liquids meet, a film or pellicle of precipitated copper ferrocyanide
(Cu2Fe(CN)6) is produced. In order to make use of this extremely
fragile membrane, Pfeffer devised the plan of precipitating it within
the walls of a vessel made of unglazed porcelain. A small clay
cylindrical cell, after thorough cleansing, was filled with a dilute
solution of potassium ferrocyanide, and immersed in dilute copper
sulphate. As these solutions entered the pores of the clay, and there
met, a membrane, consisting of copper ferrocyanide, was formed
within the walls, which under these circumstances was sufficiently
strong to withstand a pressure of five or six atmospheres.

If such a cell, furnished with a semipermeable membrane, be
employed as an endosmometer, and a dilute solution, say of sugar,
be placed within the apparatus, which is then immersed in water,
it is found that the liquid rises in the narrow tube to a certain
height above the level of the water in the outside vessel, and

FIG. 24.

Osmotic Pressure

157

remains stationary. Water passes through the membrane, but no
dissolved substance passes out. At first more water penetrates
inwards than passes out, hence the increased volume of liquid in
the cell ; but when a certain
pressure is reached, repre-
sented by the height to which
the liquid rises in the narrow
tube, equilibrium is estab-
lished, and water then passes
in each direction at equal
rates. The pressure at which
this equilibrium is established
is called the osmotic pressure
of the solution.

Fig. 25 shows the apparatus
employed by Pfeffer. z is the
porous cell, in the walls
of which the semipermeable
membrane is precipitated.
Into this are cemented the
glass tubes v and /, the latter
being attached, in the manner
indicated, to a mercury mano-
meter, m. When the cell con-
taining a solution is immersed
in water, the increased volume
of the contained liquid that
results causes a compression
of the air enclosed in the
upper part of the apparatus,
which consequently drives up
the mercury in the little mano-
meter, which thus affords a
means of measuring the os-
motic pressure of the solution
under examination. a\

The following laws in rela-
tion to osmotic pressure have
i , i. * ,

been established : —

J. Temperature and concentration being the same, different
substances when in solution exert different pressures.

I $8 Introductory Outlines

2. For one and the same substance, at constant temperature, the
pressure exerted is proportional to the concentration.

3. The pressure for a solution of a given concentration is pro-
portional to the absolute temperature,* the volume being kept
constant.

4. Equimolecular quantities of different substances (i.e. quanti-
ties in the ratio of their gramme-molecule weights), when dissolved
in the same volume of solvent, exert equal pressures at the same
temperature.t

The analogy between these laws and those relating to gaseous
pressure is very close. Thus the second statement corresponds
with Boyle's law, when we consider the term concentration to
denote the quantity of gas, that is, the number of molecules, in a
given space ; for if the number of molecules in a unit space be
doubled, the gaseous pressure is doubled, and if the number of
molecules of dissolved substance' in a given volume of water be
doubled, the osmotic pressure is doubled.

The third statement corresponds with the law of Charles : the
volume of a gas is proportional to the absolute temperature ; or, if
the volume be maintained constant, the pressure exerted by a gas
is proportional to the absolute temperature.

Osmotic pressure, therefore, just as gaseous pressure, increases
with rise of temperature and diminishes with fall of temperature.

Again, in the last of these laws, we see the extension of Avogadro's
hypothesis into the region of solution. Avogadro's hypothesis
states that equal volumes of all gases contain (under similar con-
ditions) an equal number of molecules — that is to say, an equal
number of molecules at equal temperatures exert the same pressure.
But an equal number of molecules of different gases represents an
amount of the gases in the ratio of their molecular weights, hence
Avogadro's hypothesis may be stated : equimolecular quantities of
gases at the same temperature exert equal pressures ; and this
statement, as we have seen, is only true of molecules which do
not dissociate when they pass into the gaseous state.

This close analogy between the gaseous laws and those regulat-
ing the behaviour of substances in dilute solution is explained on

* By absolute temperature is meant the number of degrees above - 273°»C.

f This is only true of those substances whose molecules neither dissociate
into simpler forms (i.e. non-electrolytes), nor associate into more complex groups
when in solution.

Osmotic Pressure

the assumption that the molecules of the dissolved body in a dilute
solution are so far apart that their mutual attractive forces are
reduced to a minimum, just as they are in the case of gaseous
molecules, and that only such properties are exhibited by them
as depend upon their number in a unit space. Further, it has
been shown in the case of a dilute solution of sugar that the
osmotic pressure (experimentally determined) is the same as the
gaseous pressure that would be exerted by the weight of sugar
present in the solution, if it were converted into gas and made to
occupy the same volume as that occupied by the solution at the
same temperature ; hence the general statement
that the pressure exerted by a substance in dilute
solution (its osmotic pressure) is the same as
would be exerted by the same amount of the sub-
stance if it existed as gas and occupied the same
volume at the same temperature.
Diffusion of Dissolved Substances.— If a

quantity of a soluble solid substance be placed
at the bottom of a vessel, which is then filled
with water, the solid dissolves, and a layer of a
strong solution is formed at the bottom. In
time, however, the dissolved substance' gradually
diffuses throughout the liquid. This process of
diffusion may be illustrated by means of the
experiment represented in Fig. 26. At the
bottom of the tall cylinder is placed a layer of
a strong solution of ferric chloride, and upon
this is carefully poured a quantity of water until
the cylinder is nearly full. Upon the top of the
water is then floated a solution of potassium
thiocyanate in alcohol, and the whole is allowed to remain undis-
turbed. The ferric chloride will gradually diffuse up into the water,
and the dissolved thiocyanate will diffuse down, and at the point
where these salts meet they will interact chemically upon each,
giving rise to a blood-red coloured solution, which will appear as
a ring about midway down the cylinder.

This phenomenon of the diffusion of dissolved substances is
strictly comparable with the diffusion of gases, although in the
former case the operation proceeds with extreme slowness. The
force which impels the molecules of dissolved substances to diffuse
is the osmotic pressure of the substance in solution.

FIG. 26.

160 Introductory Outlines

The extension of the gaseous laws into the domain of solutions
necessitates the hypothesis that in the case of some solutions the
molecules of the dissolved substance unite to form more compli-
cated molecular associations ; while in other cases (including those
substances which are electrolytes, such as the solutions of strong
acids, bases, and salts) the molecules of the substances undergo
dissociation into their ions. For, just as in the case of gases where
departures from the strict gaseous laws are seen to take place, on
account of the dissociation in some instances, and the association
in others, of the various molecules, so it is believed that the de-
viations from the strict continuity of the ideal gaseous laws into the
realm of solution are due to the operation of similar causes.

CRYSTALLINE FORMS.

When a saturated solution of a solid in a liquid is either cooled
or allowed to evaporate, the dissolved solid begins to deposit itself
out of the liquid, and it does so in most cases in definite geometric
shapes, termed crystals. (Solids which exhibit no crystalline
structure are said to be amorphous.}

The same arrangement of molecules into geometric forms often
takes place also when substances in a state of fusion (as distin-
guished from solution} pass into the solid condition, as, for example,
when melted sulphur, or mercury, or water are cooled to their re-
spective solidifying points ; and it also frequently takes place when
vapours are condensed to the solid state. Speaking generally, the
more slowly the process of solidification takes place, the larger and
more geometrically perfect will be the crystals that are formed.

All the varieties of crystalline forms, both naturally occurring
and artificially produced, are susceptible of classification into thirty-
two classes,* based upon their symmetrical development with respect
to certain imaginary planes, lines, and points, called respectively
planes of symmetry^ lines of symmetry, and centres of symmetry.

Planes of symmetry are planes cut through the crystal in such a
direction that the two divided portions are the mirrored reflections
the one of the other, the mirror being the plane itself. Crystals
may have from o to 9 planes of symmetry ; a cube, for example, has
nine such planes.

Axes of symmetry are imaginary straight lines passing through
the crystal in such a manner that when the crystal is rotated upon
one of them there will be a complete recurrence of similar faces

* The study of this classification belongs to the science of crystallography,
and falls outside the scope of a general chemical text-book ; it is therefore here
treated oply in broadest outline.

Crystalline Forms 161

and angles at least once before an entire revolution has been made.
For instance, if a tube is rotated upon an axes passing through the
centre of one face at right angles to the face, it will obviously
present the same appearance four times during a complete revolu-
tion. In thus being rotated through 360° crystals may exhibit this
periodic reappearance of the same aspect, either two, three, four,
or six times, and the axes are spoken of as binary^ trigonal, tetra-
gonal^ and hexagonal respectively.* Crystals may possess from
o to 13 axes of symmetry ; the cube, for example, has thirteen such
axes, "viz. six binary, four trigonal, and three tetragonal.

Centres of symmetry. A crystal has a centre of symmetry when
opposite to every face there is a precisely similar face parallel to it
on the other side of the crystal.

Based upon these three orders of symmetry there are mathe-
matically possible thirty-two classes into which crystals can be
ranged ; and with two or three exceptions only, crystals are known
belonging to each class.

Crystallographic Systems. These thirty-two classes are suscep-
tible of a further classification into the following six systems, based
upon the relations of their crystallographic axes : t—

I. Cubic (Regular or Isometric) System. Crystals of this system ace re-
ferred to three axes at right angles to each other, and all equal in
length. The system includes five classes. Of these, that one which
exhibits the highest degree of symmetry (spoken of as the holohedral
or normal class) has nine planes of symmetry, thirteen axes of sym-
metry, and a centre of symmetry.

1 1. Hexagonal System. Forms of this system are referred to four axes :
three are equal in length and intersect at angles of 120°, while the
fourth, known as the principal axis, is different in length and is
vertical to the others. Twelve classes are included in this system. %
The normal or holohedral class has seven planes of symmetry, seven
axes of symmetry, and a centre of symmetry.

III. Tetragonal System. Crystals belonging to this system are referable to

three axes at right angles to each other, two being equal in length.
The system includes seven classes, the holohedral class having five
planes, five axes, and a centre of symmetry.

IV. Orthorhombic (Rhombic or Prismatic) System. Crystals are referred to

three axes at right angles to each other, and all unequal in length.
Three classes are included in this system, the holohedral class having
three planes, three axes, and a centre of symmetry.

* Sometimes the terms diad, triad, &c. , are employed.

f Crystallographic axes do not necessarily correspond with axes of symmetry,
although they are made to do so whenever possible.

J Seven of these classes consist of rhombohedral forms of this system, in
which the principal axis is a trigonal axis of symmetry instead of being one
of hexagonal symmetry. They are thus regarded as hemihedral (half the
number of faces) modifications of the hexagonal form. Some crystallo-

fraphers classify them as a separate system under the name of the Rhombo-
edral system,

162 Introductory Outlines

V. Monosymmetric (or Monoclinic] System. In this system the forms are
referable to three axes of unequal lengths, two of which intersect at
an acute angle, while the third is at right angles to the other two.
Two classes belong to this system, the holohedral class having one
plane, one axis, and a centre of symmetry.

VI. Asymmetric (or Triclinic] System. Crystals are referred to three axes
of unequal lengths, intersecting one another at oblique angles. Two
classes are included in the system, the holohedral class having no
planes or axes of symmetry, but a centre of symmetry only, while the
second class has no element of symmetry at all.

One of the simplest forms in each of these six systems is the double
pyramid (see Fig. 147 A). Thus there is the tetragonal pyramid,
the hexagonal pyramid, and so on. In the case of the isometric
or cubic system this double pyramid is called the Octahedron*

Another frequently recurring form common to all the systems
except the cubic, is that of the prism, giving rise to tetragonal
prisms, hexagonal prisms,. &c.f Fig. 148 represents a group of
natural crystals in the form of hexagonal prisms terminating in
hexagonal prisms. Crystals, whether naturally occurring or artifi-
cially obtained, very seldom exhibit the perfect geometric shape
of the ideal form, but usually exhibit more or less distortion.
Fig. 1 10 (the left crystal) illustrates distortion in an orthorhombic
pyramid. Fig. 147 A represents an octahedron which has de-
veloped into the perfect ideal form, but it is only by the greatest
care in regulating their growth that such perfect crystals are
obtained. In Fig. 147 B is seen the development of what is known
as twin crystals. It very often happens that what would be an
edge or a solid angle in the ideal crystal is replaced by a plane or
planes. Such variations are called truncations. Illustrations of
truncated crystals are seen in Fig. no. In the right-hand crystal
both apexes are truncated by planes or bases ; while in the other
crystal one apex is truncated by a pyramid.

Two or more substances which crystallise in the same form are
said to be isomorphous ; and on the other hand a substance which
is capable of crystallising in two forms which do not belong to the
same system is termed a dimorphous substance. Thus sulphur is
dimorphous because it is capable of crystallising in orthorhombic
pyramids and in monosymmetric prisms. Occasionally a dimor-
phous substance is isomorphous with another dimorphous substance
in both its forms ; to this double isomorphism the term isodimor-
phism is applied.

* The double pyramids of some of the other systems are also octahedra, in
the sense that they possess eight faces, but in modern nomenclature the term
Octahedron is reserved exclusively for the isometric pyramid.

f It will be obvious that a description of a crystal merely as being prismatic
is incomplete without reference to the system to which it belongs.

CHAPTER XV
THERMO-CHEMISTR Y

WE have seen that by means of symbols and formulae chemists
express, in the form of equations, a certain amount of information
respecting chemical changes: thus by the equation C + O2=CO2
there are conveyed the facts, that carbon unites with oxygen to
form carbon dioxide, that 12 grammes of carbon combine with 32
grammes of oxygen, yielding 44 grammes of carbon dioxide, and
that the volume of the gaseous carbon dioxide obtained is the
same as that of the oxygen taking part in its formation. All such
equations bear upon the face of them the truth, that matter can
neither be destroyed nor created. The total quantity of matter
taking part in the action is unaltered by the process, although it
appears in altered form in the products of the reaction.

In all chemical changes, besides matter, energy also takes a
part ; not only do the materials concerned undergo rearrangement
or readjustment, but at the same time there is a rearrangement or
readjustment of energy. This energy change is not expressed by
the ordinary symbolic equation. Thus in the equation —

SO3 + H2O = HoSO4

the fact is embodied that 80 grammes of sulphur trioxide combine
with 1 8 grammes of water and form 98 grammes of sulphuric acid;
but the equation takes no cognisance of the fact, that when these
weights of these two substances unite to form 98 grammes of sul-
phuric acid an amount of energy, in the form of heat, is disengaged
that would raise the temperature of 213 grammes of water from o°
to the boiling-point.

Similarly, in the equation 2NC13=N2 + 3C12 there is no recogni-
tion of the fact that during this change an enormous amount of
energy leaves the system in the form of external work (over-
coming the atmospheric pressure) ; in other words, that the con-
version of nitrogen trichloride into its constituent elements is

attended with the most violent explosion.

163

164 Introductory Outlines

Energy, like matter, can neither be created nor destroyed, but as
a result of chemical action it reappears as energy in another form.
Thus it may appear as heat, as electrical energy, as kinetic
energy, or as chemical energy ; and just as the total amount of
matter taking part in a chemical change reappears in altered form
in the products of the change, so the disappearance of energy in
any of its forms gives rise to the reappearance of a proportionate
amount of energy in another form. This is the law of the conserva-
tion of energy , which may be thus stated : * " The total energy of any
material system is a quantity which can neither be increased nor
diminished by any action between the parts of the system, although
it may be transformed into any of the forms of which energy is
susceptible?

Chemical energy, or that form of energy that is set free during
chemical processes, cannot be measured by any direct method.
This energy, however, is generally transformed, during chemi-
cal change, into heat, and may therefore be measured by, and
expressed in, heat units. Thermo-chemistry may therefore be
defined as the science of the thermal changes which accompany
chemical changes.

All matter is regarded as containing a certain amount of energy
in spme form, and the purpose of thermo-chemistry is, by measur-
ing the thermal disturbance that is conditioned by a chemical
change, to ascertain the difference between the amount of energ^
contained in a system before and after such a change.

If all the energy of a system in its original state (i.e. before the
chemical change takes place) that undergoes transformation into
other forms of energy passes into heat ; if none of it leaves the
system as energy in some other form, and thereby escapes mea-
surement ; then the difference between the amount of energy
contained 'in the system in its original and its final state may be
ascertained. It by no means follows, however, that this represents
the chemical energy alone ; it has already been explained that
chemical changes are always attended by physical changes, such
as change of volume, of physical state, and so on, and we have
also learned that such physical changes are likewise accompanied by
thermal changes; the problem, therefore, is often a complicated one,
and it is not always possible to differentiate between the chemical
and the physical causes that may be operating simultaneously, and
to decide what share of the final result is due to the chemical phase
* Clerk Maxwell, " Matter and Motion."

Thermo- Chem is try 1 6 5

of the change, and what to the physical change that simultaneously
takes place.

As an illustration of the complex nature of chemical reactions
when considered from a thermal standpoint, and of the disturbing
effect of the accompanying physical changes, we may take the case
of the action of aqueous hydrochloric acid upon crystallised sodium
sulphate, Na2SO4,10H2O—

Na2SO4,10H2O + 2HC1 = 2NaCl + H2SO4 + 10H2O.

The chemical action here consists of (i) the decomposition of
sodium sulphate, (2) the decomposition of hydrochloric acid, (3)
the formation of sodium chloride, (4) the formation of sulphuric
acid. Heat is absorbed by the first two portions of the action, and
heat is evolved by the other two. The physical changes include
the passage of ten molecules of water of crystallisation {i.e. solid
water) into liquid water, and the solution of sodium chloride in
water. These changes are attended with absorption of heat, and
the net result of the entire change is the disappearance of a con-
siderable amount of heat, that is to say, the thermal value of the
reaction is a negative quantity.

The methods adopted in order to express thermo-chemical
reactions are quite simple. The ordinary chemical symbols and
formulae are used, and represent, in all cases, quantities in grammes
corresponding to the formula-weights of the substances. Thus Cl
represents 35.5 grammes of chlorine ; H2O stands for 18 grammes
of water, and so on. The chemical equation is followed by a
number representing the quantity of heat, expressed in heat units,
which is either produced or which disappears as a result of the
change. The unit .of heat is the calorie, or the quantity of heat
that is capable of raising the temperature of I gramme of water from
o° to i°. Sometimes the unit employed is the quantity of heat
required to raise I gramme of water from o° to 100°, and this unit
(which is 100 times greater than the calorie) is indicated usually
by the letter K. When heat is produced by a chemical change,
the sign + is placed in front of the number of units, and when
heat disappears, the fact is indicated by the sign - .

Thus the equation —

H2+C12 = 2HC1 +44,000 cal.
or H2+C12=2HC1 + 440 K,

means that when 2 grammes of hydrogen combine with 71

1 66 Introductory Outlines

grammes of chlorine to form gaseous hydrochloric acid, heat is
disengaged to the amount of 44,000 calories, or 440 of the larger
units, K. Or, in other words, that when these quantities of these
substances combine, an amount of energy is lost to the system,
represented by 44,000 calories. Therefore the energy possessed
by 2 grammes of hydrogen and 71 grammes of chlorine is greater
than that possessed by 73 grammes of hydrochloric acid gas by an
amount which is represented by 44,000 gramme-units of heat.
Hence the equation may be written —

2H Cl = H2 + C12 - 44,000 cal.

which signifies that when 73 grammes of gaseous hydrochloric acid
are decomposed into chlorine and hydrogen, it is necessary to
supply an amount of energy equal to 44,000 calories.

In order to indicate the state of aggregation of the different sub-
stances, the method introduced by Ostwald consists in the use of
different type, thick type being employed to denote solids, ordinary
type indicating liquids, and italics signifying gases, thus —

C + Oz = CO2 + 97,000 cal.

means that the total energies of 12 grammes of solid carbon and
32 grammes of gaseous oxygen is greater than the energy pos-
sessed by 44 grammes of gaseous carbon dioxide by an amount
equivalent to 97,000 calories.
Or, again, the equation —

cal.

signifies that 80 grammes of solid sulphur trioxide unites with 18
grammes of liquid water and forms 98 grammes of liquid sulphuric
acid, with the liberation of 21,300 gramme-units of heat.

Similarly, the heat evolved by the passage of water into ice, and
the heat that disappears when water passes into steam, may be
expressed by the equations —

H2O = H20+ 1440 cal.

when water takes a direct part in the chemical change, as, for
example, in the action of sulphur trioxide and water already quoted,
the formula represents a gramme-molecule just as in all other

Thermo- Chemistry 1 67

Cases ; but where the presence of a large quantity of water affects the
thermal result of the chemical change, by exerting, for example, a
solvent action, the symbol Aq is employed to signify that the pre-
sence of the water is considered in the thermal expression.
Thus the expression —

//^r+Aq = HBrAq+i9,9oo cal.

signifies that when 81 grammes of gaseous hydrobromic acid are
dissolved in a large excess of water, 19,900 calories are evolved.
Again, the equation —

means that when 160 grammes of gaseous bromine combine with
2 grammes of hydrogen, and the product is dissolved in an excess
of water (i.e. such a quantity of water that no thermal change is
produced by the addition of any further quantity), 64,000 calories
are disengaged. Of this 64,000 calories, 19,900x2 = 39,800 are
due to the solution of the twice 81 grammes of hydrobromic acid,
and the difference, viz., 24,000 calories, represent the heat produced
by the combination of 2 grammes of hydrogen with 160 grammes
of bromine.

If water is formed as one of the products of the chemical reaction
taking place in the case of substances in aqueous solution, such as
when a solution of hydrochloric acid is added to a solution ot
sodium hydroxide, HCl + NaHO = NaCl + H2O, as the water so
produced simply mixes with the water in which the materials are
dissolved, without producing any thermal effects by so doing, it is
usually neglected in energy equations ; although, as already stated
(page 109), when explained from the standpoint of the ionic theory,
the heat of neutralisation is here due to the formation of mole-
cules of water by the union of H'ions with HO' ions. Thus the
above action may be expressed —

HClAq+NaHOAq=NaClAq+ 13,736 cal.

The heat that is produced, or that disappears, in a chemical
change which results in the formation of a particular compound
is termed the heat of formation of that compound. Thus in the
equation —

o cal.

the heat of formation of 73 grammes of hydrochloric acid is 44,000

1 68* Introductory Outlines

thermal units. This number, however, is in reality the algebraic sum .
of three quantities. It does not express merely the heat developed
by the simple union of chlorine and hydrogen. The chemical
change expressed by the equation consists in reality of three
operations —

(i.) H2=H + H. (2.)C12 = C1 + C1. (3.)C1 + C1 + H + H = 2HCL

Each of these operations represents a distinct thermal effect ; in
Nos. (i) and (2) heat is absorbed, in No. (3) heat is evolved, and
calling these values h^ h^ and /&3, we have as the net result
/'s ~ (^1+^2)= 44,000 cal. ^

The number of heat-units, therefore, which expresses the heat of
formation of hydrochloric acid is the heat produced by the union of
two atoms of hydrogen with two atoms of chlorine, minus the heat
absorbed in the decomposition of one hydrogen and one chlorine
molecule.

Compounds such as hydrochloric acid, in the formation of which
heat is developed, are termed exothermic compounds, the reaction
by which they are produced being an exothermic change ; com-
pounds, on the other hand, whose heats of formation are expressed
by a negative sign, that is, in whose formation heat disappears, are
distinguished as endothermic compounds, and the reactions by
which they are formed are endothermic reactions.

Thus C + S2 = CS2- 19,600 cal.,

signifies that in the formation of carbon disuiphide heat is absorbed,
and the compound is therefore an endothermic compound.

Thermo-chemical determinations are made by means of instru-
ments termed calorimeters. These are of great variety, although
the principle involved is the same. The chemical reaction is caused
to take place under such circumstances, that the whole of the heat
that is liberated shall be communicated to a known volume of
water, at a known temperature.*

Direct determinations of the thermal value of chemical changes
have hitherto been made in only a limited number of comparatively
simple cases ; it is possible, however, from a few known data, to cal-
culate the thermal values of a number of changes which cannot be
directly measured. This depends upon the fundamental principle

* For descriptions of the various calorimeters, see treatises on PhysicSc

Thermo- Chemistry 169

of thermochemistry, which is itself the corollary of the law of the
conservation of energy, and which was first experimentally proved
by Hess (1840). This principle, which is sometimes termed the law
of constant heat consummation, or the law of equivalence of heat
and chemical change^ may be thus stated : The amount of heat
that is liberated or absorbed, during a chemical process, is de-
pendent solely upon the initial and final states of the system, and is
independent of the intermediate stages. The following examples
will serve to explain the application of the principle : —

1. Let us suppose it is desired to find the heat of formation of
carbon monoxide, the data at our disposal being (i) the heat pro-
duced when carbon unites with oxygen to form carbon dioxide ; and
(2) the heat formed by the combustion of carbon monoxide to
carbon dioxide. The thermal equations are —

(1) C+02 = £02 + 97,000 cal. "^

(2) 2C0+02=2C02 + 136,000 cal.

Halving the second equation, in order to get the heat produced
in the formation of 44 grammes of carbon dioxide (i.e. the same
weight as in the first), we may represent the equation as —

CO + O = C02 + 68,000 cal.*

The difference between the two values 97,000 and 68,000 will be
the heat of formation of carbon monoxide, therefore we get the
equation —

C + O = CO + 29,000 cal.

2. The compound, methane (marsh gas), CH4, cannot be formed
by the direct union of its elements, but its heat of formation can
be calculated by the application of this principle. The data in
this case are the ascertained heats of formation of carbon dioxide

* It must be remembered that this equation does not express the whole
truth : as it here stands it would imply that 68,000 calories represent the heat
formed by the simple chemical union of 28 grammes of carbon monoxide with
16 grammes of oxygen. In reality this number is half the sum of the two
values, namely, the heat of combination of 56 grammes of carbon monoxide
with 32 grammes of oxygen, minus the heat absorbed by the decomposition of
32 grammes of oxygen molecules into their constituent atoms. The oxygen
atom does not exist alone, and whenever free oxygen takes part in a chemical
change the molecules of the element are first separated into their atoms.

Introductory Outlines

and of water, and the heat produced by the combustion of marsh
gas, the thermal equations being —

(1) C + O2= CO2 + 97,000 cal

(2) 2tfa+08 = 2#-a0+ 136,800 cal.

(3) Ctf4 + 2O2=C02 + 2H2O + 212,000 cal.

The difference between the thermal value of the last process and
the sum of the first and second represents the heat of formation of
marsh gas —

97,000+136,800-212,000=21,800,

hence we get the expression —

C + 2//2 = C7/4 + 2 1 ,800 caL

PART II

THE STUDY OP POUR TYPICAL ELEMENTS
HYDROGEN-OXYGEN-NITROGEN-CARBON

AND THEIR MORE IMPORTANT COMPOUNDS

CHAPTER I

HYDROGEN

Symbol, H. Atomic weight— 1.008. Molecular weight =2.016.
Density = i. 008.

History. — The existence of hydrogen as an individual sub-
stance was first established by Cavendish (1766), who applied to it
the name inflammable air. He obtained the gas by acting upon
certain metals, as iron, tin, and zinc, with either sulphuric or hydro-
chloric acid.

Occurrence. — In the free state hydrogen occurs only in small
.quantities upon the earth. It is evolved with other volcanic gases,
and is present in the gases which escape from petroleum wells.
It is evolved also during the fermentation and decomposition of
certain organic compounds, and is therefore present in the breath
and the intestinal gases of animals. From these sources it finds
its way into the atmosphere, where it is present to the extent of
about .04 volumes in 1000 volumes of air. Hydrogen has also
been found in many specimens of meteoric iron, and also in
certain rocks, where it is present as occluded gas.

Hydrogen in the uncombined state exists in enormous masses
upon the sun, and is present in certain stars and nebulae. The
so-called prominences which are seen projecting from the sun's
disk to a distance of many thousands of miles, and which were

Inorganic Chemistry

first observed during solar eclipses, consist of vast masses of in-
candescent hydrogen.

In combination with other elements hydrogen is extremely
abundant ; its commonest compound is water, which consists of
one part by weight of this element combined with eight parts of
oxygen. In combination with chlorine, as hydrochloric acid, with
carbon as marsh gas, and with sulphur as sulphuretted hydrogen,
this element also occurs in large quantities. All known acids

FIG. 27.

contain hydrogen as one of their constituents, and it is present in
almost all organic compounds.

Modes Of Formation.— ( i.) Hydrogen may be obtained from
water by the action of various metals upon that compound under
certain conditions. The metals sodium and potassium will decom-
pose water at the ordinary temperatures ; when, therefore, a frag-
ment of either of these metals is thrown upon water, the latter is
decomposed and hydrogen set free : —

Hydrogen

173

The metals, being lighter than water, float upon its surface, and,
owing to the heat of the reaction, melt and roll about upon the liquid
as molten globules. With potassium, the temperature developed
is sufficiently high to cause the hydrogen to inflame, and it burns
with a flame coloured violet by the vapour of the metal. The
hydroxide of the metal, which is the second product of the action,
dissolves in the excess of water, rendering the liquid alkaline.
The alkalinity of the solution may be made evident by the addition of
a reddened solution of litmus, which will be turned blue by the alkali.

In order to collect the hydrogen evolved by the action of sodium
upon water, the metal is placed in a short piece of lead tube closed
at one end, which causes it to sink in the liquid, and an inverted
glass cylinder filled with water is placed over it, as shown in Fig. 27.
The evolved hydrogen then rises as a stream of bubbles into the
cylinder and displaces the water.*

FIG. 28.

(2.) Water may be readily decomposed at the boiling-point,
by means of zinc, if the metal be previously coated with a thin
film of copper by immersion in a dilute solution of copper sul-
phate. When this copper-coated zinc (known as zinc -copper
couple) is heated in a small flask filled with water, and provided
with a delivery tube, the oxygen of the water combines with the
zinc forming zinc oxide, and hydrogen is evolved, which may be
collected over water at the pneumatic trough : * —

* For detailed description of these experiments, see Newth's "Chemical
Lecture Experiments," p. 2.

174 Inorganic Chemistry

(3.) At a still higher temperature, water in the state of steam can
be readily decomposed by the metal magnesium, magnesium oxide
being formed and hydrogen liberated : —

For this purpose the magnesium is strongly heated in a glass
bulb (Fig. 28), while steam from a small boiler is passed over it.
As the temperature of the metal approaches a red heat it bursts
into flame, and the issuing hydrogen may be ignited as it escapes
from the end of the tube.

(4.) If iron be heated to bright redness and steam be passed
over it, the water is decomposed, the oxygen uniting with the iron

FIG. 29.

to form an oxide known as triferric tetroxide, or magnetic oxide of
iron, thus —

This method is employed on a large scale for the preparation of
hydrogen for commercial purposes. Iron borings or turnings are
packed into an iron tube, which is strongly heated in a furnace,
and steam from a boiler is passed through the tube.

(5.) For laboratory purposes hydrogen is most conveniently pre-
pared by the action of dilute sulphuric acid upon zinc : —

Zn + H2SO4=ZnS(^4+ H2.
For this purpose granulated zinc (i.e. zinc which has been melted

Hydrogen 175

and poured into water) is placed in a two-necked Woulf's bottle
(Fig. 29), and a quantity of sulphuric acid, previously diluted with
six times its volume of water, is introduced by means of the funnel.
A brisk action sets in, and hydrogen is rapidly disengaged. After
the lapse of a few minutes the air within the apparatus will be
swept out by the hydrogen, when the gas may be collected over
water in the pneumatic trough.

The hydrogen so obtained is never absolutely pure ; it is liable to contain
traces of arsenic hydride, hydrogen sulphide, hydrogen phosphide, oxides of
nitrogen, and nitrogen. The nitrogen is derived from the air, which finds
its way through joints in the apparatus, and also from air dissolved in the
acid. There is no known process for removing this impurity. The other
gases are due to impurities in the zinc and the sulphuric acid, and can be
removed, if required, by passing the hydrogen through a series of tubes con-
taining absorbents (see page 210).

Absolutely pure sulphuric acid, even when diluted with water,
has no action upon perfectly pure zinc.

Scrap iron may be substituted for zinc, but the hydrogen so

obtained is much less pure, and is accompanied by compounds of

carbon (derived from the carbon in the iron), which impart to the

gas an unpleasant smell ; the reaction in this case is the following : —

Fe + H2SO4 = FeSO4+H2.

Hydrochloric acid can be employed in place of sulphuric acid
with either zinc or iron, the reaction then being : —
Zn + 2HCl = ZnCl2+H2.

These actions of acids upon metals when expressed in the form of
ionic equations will each be seen to consist simply of the transference
of two positive atomic charges from two hydrogen ions to an atom
of the metal,whereby the latter is converted into a divalent ion, thus —

(6.) Hydrogen in a high degree of purity is conveniently prepared
in small quantity by the electrolysis of water acidulated with sul-
phuric acid (see page 207).

(7.) Hydrogen is disengaged when certain metals, such as zinc,
iron, and aluminium, are boiled with an aqueous solution of potas-
sium or sodium hydroxide. Thus, in the case of zinc, when this
metal in the form of filings is boiled with a solution of potassium
hydroxide, hydrogen is evolved, and a compound of zinc, potassium,
and oxygen remains in solution, namely, potassium zinc oxide (or
potassium zincate), thus : —

(8.) Hydrogen is also obtained by heating alkaline oxalates, or

176

Inorganic Chemistry

formates, with either potassium or sodium hydroxide, with the
simultaneous formation of an alkaline carbonate ; thus with sodium
oxalate : —

Na2C2O4 + 2NaHO = H2 + 2Na2CO3.

Properties. — Hydrogen is a colourless gas, and has neither
taste nor smell. It is the lightest known substance, being 14.3875
times lighter than air. Its specific gravity is 0.0695 (air=i).
One litre of the gas at o° C., and under a pressure of 760 mm. of
mercury (i.e. the standard temperature and pressure) weighs
0.089873 gramme ; or I gramme of hydrogen at the standard tem-
perature and pressure occupies 11.127 litres.

On account of its extreme lightness, hydrogen may be poured

FIG. 30.

upwards from one vessel to another. If a large beaker be sus-
pended mouth downwards from the arm of a balance and counter-
poised, and the contents of a jar of hydrogen be poured upwards
into the beaker, the equilibrium of the system will be disturbed,
and the arm carrying the beaker will rise.

The lightness of hydrogen can also be shown by causing a
stream of the gas to issue from a tube placed in such a position
that its shadow is cast upon a white screen by means of a powerful
electric light. When the gas is streaming from the tube, its up-
ward rush will be visible upon the screen as a distinct shadow,
caused by the difference between the refractive power of air and
hydrogen (Fig. 30).

Hydrogen

Hydrogen is inflammable and burns with a non-luminous flame,
the temperature of which is very high. The product of the com-
bustion of hydrogen is water, and if a jet of the gas be burned
beneath the apparatus seen in Fig. 31, considerable quantities of
water may be collected in the bulb. In the act of combustion the
hydrogen combines with the oxygen of the air, forming the oxide
of hydrogen, namely, water:*—

If hydrogen be mixed with the requisite quantity of air, or oxygen,
and a light applied to the mixture, the
combination of the two gases takes
place instantly with a violent explo-
sion ; t hence the necessity of care-
fully expelling all the air from the
apparatus in which hydrogen is being
generated before applying a flame to
the issuing gas.

Hydrogen will not support the com-
bustion of ordinary combustibles ; thus,
if a burning taper be thrust into a jar of
the gas, the hydrogen itself will be
ignited at the mouth of the jar, which
must be held in an inverted position,
but the taper will be extinguished ; on
withdrawing the taper it may be re-
ignited by the burning hydrogen.

Although hydrogen is not poisonous,
it is incapable of supporting animal life
owing simply to the exclusion of oxygen.
When mixed with air and inhaled, it raises the .pitch of the voice

FIG. 31.

* From this fact the name Hydrogen (signifying the water producer) is derived.

f Baker has recently shown (Jour. Chem. Soc. , April 1902) that if the two
gases are perfectly pure and dry, they may be strongly heated without uniting.
In these experiments a coil of silver wire suspended in the gases was heated by
means of an electric current until the silver melted, that is, above 1000°; but
no chemical union of the oxygen and hydrogen took place, although the ordi-
nary temperature of explosion is 615° (V. Meyer). Baker has also shown that if
a mixture of these two gases, which has not been specially dried, be exposed to
sunlight, combination slowly takes place; whereas with the perfectly dry
gases no measurable combination occurs.

178 Inorganic Chemistry

almost to a falsetto. The same effect may be seen by sounding a
pitch-pipe, or organ-pipe, by means of a stream of hydrogen
instead of ordinary air, when it will be noticed that the note given
out is greatly raised in pitch.

Hydrogen is very slightly soluble in water. It was formerly
believed that this gas formed an exception to the rule that the
solubility of gases in water diminishes with rise of temperature,
and it was supposed that the solubility of hydrogen was constant
between the temperatures o° and 25°. More recent experiments
have shown that this is not the case. The solubility of this gas, as
determined by W. Timofejeff (1890), is seen in the table on p. 143.

Hydrogen was first liquefied on May 10, 1898, by Dewar. Prior
to this time it had never been obtained as a coherent or static
liquid — that is, a liquid with a meniscus — although momentary
indications of its liquefaction had been obtained by Olszewski as
far back as 1895. The critical temperature of hydrogen (namely,

— 238°) being below the lowest point obtainable by the rapid
ebullition of liquid oxygen or air, no external refrigerating agent
is available which is capable of cooling the gas below its critical
point, and therefore of causing its liquefaction. By an extension
of the principle of self-cooling explained on p. 76, however, namely,
by causing a jet of the gas previously cooled to — 205° to continu-
ously escape from a fine orifice under a pressure of 1 80 atmos-
pheres, Dewar has succeeded in collecting considerable quantities
of liquid hydrogen in specially constructed vacuum -jacketed
vessels.

Liquid hydrogen is clear and colourless as water, thus disposing
of the theory once advocated that if obtained in the liquid state
hydrogen would be found to exhibit metallic properties. The
boiling-point of the liquid is —253° (Dewar), at which temperature
air is immediately solidified. Thus, if a tube sealed at one end,
but freely open to the air at the other, be immersed in liquid
hydrogen, the cooled end of the tube quickly becomes filled with
solidified air. Similarly, oxygen is frozen to a pale-blue solid.

The specific gravity of liquid hydrogen is about 0.07 ; that is to
say, it is only about ^jth the density of water, or about 14 c.c. of the
liquid weighs only i gramme. By its own rapid evaporation liquid
hydrogen has been frozen to a white solid mass, which melts at

— 257° ; and by the rapid evaporation of this solid a temperature
of — 260° has been obtained, which is the lowest degree of cold
ever reached. By means of liquid hydrogen as a refrigerating

Hydrogen 1 79

agent, all the known gases, have been condensed to the liquid
state.

Occluded Hydrogen.— Certain metals, such as iron, platinum,
and notably palladium, possess the property when heated of
absorbing a large quantity of hydrogen, and of retaining it when
cold. Graham found that at a red heat palladium absorbed, or
occluded, about 900 times its own volume of hydrogen, while
even at ordinary temperatures it was able to absorb as much
as 376 times its volume.* Graham believed that the hydrogen
so occluded assumed the solid form, and was alloyed with the
palladium, and to denote the metallic nature of the gas he gave
to it the name hydrogenium. From later experiments of Troost
and Hautefeuille, it seems probable that a definite compound of
hydrogen and palladium exists, of the composition of Pd2H.

After its absorption of hydrogen the metal presents the same
appearance as before, although some of its physical properties
have become slightly modified; thus it is more magnetic than
ordinary palladium, and its electric conductivity is considerably
reduced.

In view of our present knowledge of the entire absence of any
metallic characters in liquid or solid hydrogen (gained, however,
entirely since Graham's time), the view that this is an alloy is no
longer tenable, as this term is only strictly applicable to the union
of metals.

The absorption of hydrogen by palladium is readily seen by
making a strip of palladium foil the negative electrode in an
electrolytic cell containing acidulated water, the positive pole
being of platinum. Oxygen will be evolved from the latter
electrode, while for some time no gas will be disengaged from
the surface of the palladium, the hydrogen being completely
absorbed by the metal. During the absorption of hydrogen the
palladium undergoes an increase in volume : Graham observed the
increase in length of a palladium wire to be equal to 1.6 per cent.

* According to Neumann and Strientz (Zeitschrift fur Analytische Chemief
vol. 32), one volume of various metals in a fine state of division is capable of
absorbing the following amounts of hydrogen : —

Palladium, black . 502.35 vols.
Platinum, sponge . 49.3 ,,
Gold .... 46,3
Iron . ... 19.17 „

Nickel . . . 17. 57 vols.

Copper . . . 4.5 „

Aluminium . . 2.72 ,,

Lead. . • . 0.1.5 »

I So Inorganic Chemistry

This change in volume suffered by the metal may be strikingly
demonstrated by employing two strips of palladium foil, protected
/>n one side by a varnish, as the electrodes in the electrolytic cell.
On passing the current the negative electrode immediately begins
to bend over towards the varnished side : when the current is
reversed it again uncurls ; and the other, being now the negative
pole, at once begins to perform the same curling movements.

Hydrogen, which is thus occluded in the metal palladium, is
capable of bringing about a number of chemical changes which
ordinary hydrogen is unable to effect : thus, when a strip of
hydrogenised palladium is immersed in a solution of a ferric salt
a portion of the iron is reduced to the ferrous state.*

* See "Chemical Lecture Experiments," Nos. 27, 28, 29.

CHAPTER II
OXYGEN

Symbol, O. Atomic weight ='16.00. Molecular weight = 32.

History. — Oxygen was discovered by Priestley (1774). He ob-
tained it by heating the red oxide of mercury (known in those days
as mercurius calcinatus, per se) by concentrating the sun's rays
upon it by means of a powerful lens. Priestley applied to the gas
the name dephlogistigated air. Oxygen was independently dis-
covered by Scheele. Scheele's discovery of oxygen was published
in 1775, but recent research among his original papers has .brought
to light the fact that the discovery was actually made in 1773, prior
therefore to Priestley's discovery. Scheele called the gas empyreal
air, on account of its property of supporting combustion. Lavoisief
subsequently applied to this gas the name " oxygene " (from o£v$-,
sour ; and yewao>, I produce), to denote the fact that in many
instances the products obtained by the combustion of substances
in the gas were endowed with acid properties. Oxygen, indeed,
came to be regarded as an essential constituent of acids, and was
looked upon as the " acidifying principle." The subsequent deve-
lopment of the science has shown that this idea is erroneous, and
that oxygen is not a necessary constituent of all acids.

Occurrence. — In the free state oxygen occurs in the atmos-
phere, mechanically mixed with about four times its volume of
nitrogen. In combination with other elements it is found in
enormous quantities. Thus it constitutes eight-ninths by weight
of water, and nearly one-half by weight of the rocks of which the
earth's crust is mainly composed.

The following table (Bunsen) gives the average composition of
the earth's solid crust, so far as it has been penetrated by man.
It must be remembered, however, that the greatest depth to which
man has examined, when compared with the diameter of the earth,
is after all only, as it were, a mere scratch,

181

1 82 Inorganic Chemistry

Average Composition of the Earths 'Crust.

Oxygen . . . . . . 44.0 to 48.7

Silicon . ..... . . 22.8 „ 36.2

Aluminium . . • . ' , . 9.9 „ 6.1

Iron . .: r . 9.9 „ 2.4

Calcium . • , , , , 6.6 „ 0.9

Magnesium , . '.".'• « '- 2.7 „ o.i

Sodium. . «... - 2.4 „ 2.5

Potassium . . . ''•".•-' . 1.7 „ 3.1

IOO.OO IOO.OO

Modes Of Formation. — (i.) Oxygen may readily be obtained
by a slight modification of Priestley's original method, namely, by
heating mercuric oxide in a glass tube, by means of a Bunsen
flame. The red oxide of mercury first darkens in colour, and is
decomposed by the action of the heat into mercury and oxygen,
thus—

The evolved oxygen may be collected over water in the pneumatic
trough, while the mercury condenses in the form of metallic
globules upon the cooler parts of the tube. This method of
obtaining oxygen is never employed when any quantity of the
gas is required — it is chiefly of historic interest.

(2.) For experimental purposes oxygen is best prepared from
potassium chlorate. When this salt is heated it melts, and at
about 400° decomposes with brisk effervescence due to the evolution
of oxygen, while potassium chloride remains : * —

KC1O3 = KC1 + 3O.

If the potassium chlorate be previously mixed with about one-
fourth of its weight of manganese dioxide, it gives up the whole of
its oxygen at a temperature considerably below the melting-point
of the salt, and at a greatly accelerated rate. When, therefore, the
oxygen is not required to be perfectly pure, a mixture of these two

* The mechanism of this reaction is more complex than is represented by
this equation. It has been shown (P. F. Frankland) that during the decom-
position potassium perchlorate, KC1O4, is continuously being formed, and
again resolved into KC1O3 and O. The extent to which this takes place depend-
ing upon the temperature,

Oxygen 183

substances is usually employed. The mixture may be conveniently
heated in a " Florence " flask, supported in the position shown in
the figure, and gently heated with a Bunsen flame. The gas is
washed by being passed through water, and then collected either
at the pneumatic trough or in a gas-holder.

The manganese dioxide is found at the end of the reaction to be
unchanged : the part it plays in the decomposition belongs to a
class of phenomena to which the name catalysis is applied ; the
manganese dioxide, in this instance, being the catalytic agent. It
was at one time supposed that by its mere presence, itself under-
going no change, the manganese dioxide enabled the potassium
chlorate to give up its oxygen more readily and at a lower tempera-
ture ; but the accumulated evidence which has been collected by
the study of an increasing number of similar cases of catalytic
action leads to the conclusion that the manganese dioxide is here

FIG. 32.

playing a more distinctly chemical part in the reaction. So far as
is known, in all phenomena of this order, the catalytic agent is a
substance which possesses a certain degree of chemical affinity for
one of the constituents of the body to be decomposed, and the
influence of this attraction is a necessary factor in determining the
splitting up of the compound. Owing, however, to certain condi-
tions which are present, such, for example, as the particular
temperature at which the reaction is conducted, the catalytic agent
is unable to actually combine with the constituent for which it has
this affinity, or if it combines, the combination it forms is unable to
exist and is instantly resolved again : hence the catalytic agent
comes out of the action in the same state as it was at the com-
mencement.

In the case before us, it is now believed* that the action of the
manganese dioxide in facilitating the evolution of oxygen from
* JSodeau, Trans, Chem, Soc.t 1902, vol. ii. p. 1066,

184 Inorganic Chemistry

potassium chlorate is due to the formation of a higher oxide of
manganese by the oxidising action of the chlorate, which oxide
being unstable under the existing conditions, subsequently breaks
up into oxygen and the original oxide.*

The temperature at which this reaction takes place is below that
which is necessary for the formation of potassium perchlorate,
hence under these conditions this salt is not produced.

(3.) When manganese dioxide itself is heated to bright redness, it
parts with one-third of its oxygen and is converted into trimanganic
tetroxide.

3MnO2= Mn3O4+ O2.

(4.) Other peroxides, when heated, similarly yield a portion of the
oxygen they contain. One of these, namely, barium peroxide, is
now largely employed for the preparation of oxygen upon a manu-
facturing scale. This method, known as Brings process, from the
name of the inventor, is based upon the fact that when barium
oxide (BaO) is heated in contact with air, it unites with an additional
atom of oxygen, forming barium peroxide, thus —

And that when this substance is still further heated, it again parts
with the additional oxygen and is reconverted into the monoxide —

The process, therefore, is only an indirect method of obtaining
oxygen from the air, the same quantity of barium monoxide being
employed over and over again. In practice it was found that
instead of effecting the two reactions by altering the temperature,
which involved loss of time and considerable expense, the same
result could be obtained by altering the pressure and keeping the
temperature constant. If the monoxide be heated to the lower
temperature, at which the first reaction takes place, and air be
passed over it at the ordinary atmospheric pressure, atmospheric
oxygen is taken up and barium peroxide is formed. If the pressure

* Secondary reactions simultaneously take place, resulting in the formation
of small quantities of potassium permanganate, and the evolution of traces of
ghlprine,

Oxygen

I85

be then slightly reduced by suitable exhaust pumps, the peroxide
immediately gives up one atom of oxygen without any further
application of heat, and is retransformed into the monoxide. In.

Fg=JL — a

this way, by alternately sending air through the heated retorts
containing the oxide and then exhausting the retorts, a continuous
process is obtained without change of temperature.
The modus operandi of the process will be seen from Fig. 33,

1 86 Inorganic Chemistry

which represents the general arrangement of the apparatus. A
number of retorts, R, consisting of long narrow iron pipes, are
arranged vertically in rows in the furnace, where they are heated
by means of "producer-gas" (i.e. carbon monoxide with atmospheric
nitrogen, obtained by the regulated combustion of coke).

By means of the pump P, air is drawn in at the " air intake," and
forced through purifiers in order to withdraw atmospheric carbon
dioxide ; the complete removal of this impurity being essential to
the successful carrying out of the operation. The purifiers are
so arranged that any of them can be thrown out of use at
will.

By means of automatic gear the purified air is sent through pipe
J to the distributing valve X, from which it passes by the pipe Y into
the retorts, being made to pass down through one row and up through
the other. The oxygen is then absorbed, and the accumulating
nitrogen escapes by the relief valve W. When the absorption of
oxygen by the barium monoxide in the retorts has continued for
ten or fifteen minutes, the automatic reversing gear comes into
operation. The relief valve W is thereby closed, communication
with the purifiers is cut off, and the action of the pumps at once
causes a reduction of pressure within the retorts. When the pres-
sure falls to about 660 mm. (26 inches, or about 13 Ibs. on the
square inch), the peroxide gives up oxygen, and is reduced to the
monoxide. The oxygen is drawn away by the pipe J and is passed
on to a gas-holder. The first portions of gas that are drawn out
of the retorts will obviously be mixed with the atmospheric nitrogen
which was there present ; in order that this shall be got rid of, the
automatic gear is so arranged that communication with the pipe
leading to the gas-holder is not opened until a few seconds after
the reversing gear is in operation, and the first portions of gas that
are pumped out are made to escape into the air by a snifting valve
S, which is automatically opened and closed.

(5.) Oxygen may be obtained by heating manganese dioxide with
sulphuric acid, the dioxide parting with the half of its oxygen, and
a sulphate of the lower oxide being formed — *

(6.) Similarly, potassium dichromate (a salt containing chromium

* In order to avoid unnecessarily complicating chemical equations, it is some-
times convenient to represent them atomically. Moreover, by so doing the
mechanism of the reaction is often rendered more clear,

Oxygen 187

trioxide, CrO3), when heated with sulphuric acid, yields oxygen ; the
chromium at the same time being reduced to a lower state of oxi-
dation, viz., Cr2O3, in which condition it unites with sulphuric acid,
forming chromium sulphate —

K2Cr2Or + 4H2SO4= K2SO4 + Cr2(SO4)3 + 4H2O + 3O.

During the reaction the red colour of the dichromate changes to
the deep olive-green colour possessed by chromium sulphate.

(7.) Many other highly oxidised salts yield oxygen when acted
upon by sulphuric acid ; thus, with potassium permanganate the
following action takes place : —

(8.) If hydrogen peroxide be added to dilute sulphuric acid, and
the mixture dropped upon a solution of potassium permanganate
contained in a suitable generating flask, a rapid evolution of oxygen
takes place at the ordinary temperature, thus —

K2Mn2O8 + 3H2SO4+5H2O2=K2SO4

(9.) When strong sulphuric acid is dropped upon fragments of
brick or pumice-stone, contained in an earthenware or platinum
retort and maintained at a bright red heat, the acid is decomposed
into water, sulphur dioxide, and oxygen —

H2SO4 = H2O + SO2+O.

The products of the decomposition are passtd through water, which
absorbs the sulphur dioxide, and also arrests any undecomposed
sulphuric acid, and the oxygen is collected over water. When this
process is used on a large scale, the sulphur dioxide is absorbed by
being passed through a tower filled with coke, and down which a
stream of water is allowed to trickle, and the solution so obtained
can be utilised in the manufacture of sulphuric acid.

(10.) Oxygen can be obtained from bleaching-powder by methods
which afford interesting instances of catalytic action.* The
composition of bleaching-powder is expressed by the formula
Ca(OCl)Cl. If this substance be mixed with water, and a small
quantity of precipitated cobalt oxide added, and the mixture gently
warmed, oxygen is rapidly evolved. The cobalt oxide, CoO, is the
catalytic agent ; it is able to combine with more oxygen to form

* Experiments 35, 36, 37, 152, " Chemical Lecture Experiments," new ed,

1 88 Inorganic Chemistry

Co2O3, but this compound is reduced as fast as it is formed, and
the oxygen is evolved as gas —

(1) Ca(OCl)Cl -1- 2CoO = Co2O3+ CaCl2.

(2) Co2O3=2CoO + O.

A solution of calcium hypochlorite, which may be obtained from
bleaching-powder (see Bleaching-powder), behaves in the same
way ; and, as in the above reaction, nickel oxide may be substi-
tuted for cobalt—

£a(OCl)2=CaCl2+O2.

(11.) A similar instance of catalysis, by which oxygen may be
obtained, is seen when a stream of chlorine gas is passed through
boiling milk of lime, to which a small quantity of the oxide of
cobalt or nickel has been added—

CaH2O2+ C12 = CaCl2 + H2O + O.

A reaction of the same order takes place when the milk of lime is
replaced by either potassium or sodium hydroxide —

= 2NaCH-H2O + O.

(12.) When a mixture of steam and chlorine gas is heated to bright
redness, the steam is decomposed, the hydrogen combining with
the chlorine to form hydrogen chloride (hydrochloric acid), and
the oxygen is set free —

H2O + C12 = 2HC1 + O.

In order to prepare oxygen by this reaction, chlorine gas is caused
to bubble through water which is briskly boiling in a glass flask,
F (Fig. 34). The mixture of chlorine and steam is then passed
through a porcelain tube filled with fragments of porcelain, and
maintained at a bright red heat in a furnace. The issuing gases
are passed through a Woulfs bottle, containing a solution of
sodium hydroxide, in order to absorb the hydrochloric acid, and
the oxygen is collected at the pneumatic trough.

(13.) Oxygen is formed on a large scale in nature by the decom-
position of atmospheric carbon dioxide by the green leaves of
plants, under the influence of light. The carbon dioxide is decprn-

Oxygen

i89

posed into carbon, which is assimilated by the plant, and into
oxygen, which is thrown into the atmosphere. It has been esti-
mated that i square metre of green leaf is able, under the influ-
ence of sunlight, to decompose more than I litre of carbon dioxide
per hour.

(14.) Of the many other methods by which it has been proposed,
from time to time, to manufacture oxygen on a large scale, may be
mentioned one, known as the Tessie dti Motay process, from the
name of the inventor. This method consists in the alternate for-
mation and decomposition of sodium manganate. The process
consists of two operations, which are carried out at different tem-
peratures. When a current of air is passed over a moderately

FIG. 34.

heated mixture of manganese dioxide and sodium hydroxide,
sodium manganate is formed —

2MnO2 + 4NaHO + O2=2H2O + 2Na2MnO4.

And if this sodium manganate be heated to bright redness, and a
current of steam at the same time passed over it, the manganate is
reduced to dimanganic trioxide, sodium hydroxide is reformed, and
oxygen evolved, thus —

On again passing air over the residue, after allowing the tempera

190 Inorganic Chemistry

ture of the mass to fall to that at which the first reaction was
conducted, sodium manganate is once more reformed —

Properties. — Oxygen is a colourless gas, having no taste or
smell. It is slightly heavier than air, its specific gravity being
1.1056 (air= i). Onejitre of the gas, at the standard temperature
and pressure, weighs 1.429 grammes. Oxygen is slightly soluble
in water. I c.c. of water at o° C. dissolves 0.0489 c.c. of oxygen
measured at o° C. and 760 mm. pressure. The solubility of oxygen
in water diminishes as the temperature rises in accordance with
the interpolation formula (Winkler) : —

£=0.0489- .001 341 3/+. 0000283^ ~ .00000029 5 34/3.

Fish are dependent upon the dissolved oxygen in water for their
supply of this gas for respiration. Oxygen is also soluble in molten
silver, which is capable of absorbing about twenty times its own
volume of this gas (see Silver).

Oxygen is endowed with very powerful chemical affinities. Even
at the ordinary temperature it is able to combine with such elements
as phosphorus, sodium, potassium, and iron. Most of the chemical
phenomena exhibited by the atmosphere are due to the presence in
it of free oxygen, the atmosphere being practically oxygen diluted
with four times its volume of nitrogen. Thus, when a piece of bright
metallic sodium is exposed to the air, the surface becomes instantly
tarnished and coated over with a film of oxide : when iron rusts, it
in the same way is being acted upon by the oxygen of the air forming
an oxide (or hydrated oxide) of iron ; in these cases the metals are
said to become oxidised. If the metal be obtained in a sufficiently
finely divided condition before being exposed to the air, or to pure
oxygen, this process of oxidation may proceed so rapidly that the
heat developed by the combination will cause the metal to burn.
When the process of oxidation is accompanied by light and heat,
the phenomenon is known as combustion^ the oxygen being spoken
of as the supporter of combustion : bodies which burn in the air,
therefore, are simply undergoing rapid combination with oxygen.
It will obviously follow, that bodies which are capable of burning
in the air will burn with greatly increased rapidity and brilliancy
when their combustion is carried on in pure or undiluted oxygen.
A glowing chip of wood, or a taper with a spark still upon the

Oxygen

191

wick, when plunged into pure oxygen, will be instantly rekindled.
Such substances as sulphur, charcoal, phosphorus, which readily
burn in air, when burnt in pure oxygen carry on their combustion
with greatly increased brilliancy. Many substances which are not
usually regarded as combustible bodies will burn in oxygen if their
temperature be raised sufficiently high to initiate the combustion ;
thus a steel watch-spring, or a bundle of steel wires, if strongly
heated at one end, will burn in oxygen, throwing out brilliant
scintillations. This experiment is most readily shown by project-
ing a spirit-lamp flame upon the ends of a bundle of steel wire, by
means of a stream of oxygen, as shown in Fig. 35. As soon as the
ends of the wire are sufficiently heated, and begin to burn, the
lamp may be withdrawn and the wire held in the issuing stream

FIG. 35.

of oxygen, in which it will continue its combustion with great
brilliancy.*

It is a remarkable fact, and one which has not yet received any
entirely satisfactory explanation, that these instances of combustion
in oxygen will not take place if both the gas and the material be
absolutely dry. It has been shown that phosphorus, sealed up in a
tube with oxygen which has been absolutely freed from aqueous
vapour, may even be distilled in the gas without any combination
taking place. A mixture of oxygen and carbon monoxide, which
under ordinary circumstances explodes when heated, is found when
perfectly dry to remain unaffected by the passage of an electric
spark (p. 298). Similarly, perfectly dry chlorine is without action
upon metals such as copper or sodium, while under common con-
ditions it combines with them with the greatest readiness. In
all these cases where the absolutely dry materials are incapable

* See also Experiments 48-52, " Chemical Lecture Experiments,"

102 Inorganic Chemistry

of acting upon each other, the introduction of the minutest trace
of moisture is sufficient to allow the action to proceed, but the
exact way in which this operates in causing the effect is at present
not known with certainty.*

Oxygen is the only gas which is capable of supporting respira-
tion : an animal placed in any gas or gaseous mixture containing
no free oxygen rapidly dies. Undiluted oxygen may be breathed
with impunity for a short time, but its continued inhalation soon
produces febrile symptoms. The inhalation of oxygen is occa-
sionally had recourse to in cases of asphyxiation, or under

* It would seem evident if it is the presence of water, and water only as the
third substance, which is the necessary condition to bring about chemical
action in such cases as these, that however completely a mixture of oxygen
and hydrogen were dried, it would explode when heated above the temperature
at which union begins ; because the product of the combination of a minute
portion of the mixture would furnish sufficient water to determine the ex-
plosion of the remainder. The researches of V. Meyer, Dixon, and Baker,
extending over the last decade, seemed to entirely confirm this view, for they
could detect no diminution in the velocity of the union of these gases even
when most carefully dried.

Quite recently, however (April 1902), Baker has shown that if these gases
be perfectly pure (see p. 208), as well as perfectly dry, the mixture may be
heated to temperatures much higher than that «tt which combination usually
takes place without exploding. Moreover, if the pure gases are heated before
the drying operation has been carried to its highest degree, it is found that
although union begins to take place and water is actually formed in quantity
greatly in excess of that which would be necessary to bring about the action
had the gases not been pure, nevertheless no explosion of the mixture takes
place. The gases being absolutely pure to start with, the water produced bj»
their union will also be pure ; and it would appear from these experiments that
perfectly pure water alone is not capable of bringing about these chemical
combinations.

These new and most interesting results lend support to the hypothesis which
has been put forward in order to explain the influence of water in determining
such chemical actions as these, namely, that chemical action cannot take place
without the presence of an electrolyte ; that the removal of water is in reality
the removal of any possibility of an electrolyte being present. If the water
present is absolutely pure, since pure water is a non-electrolyte (p. 109), it
should therefore not be able to operate in causing the chemical action to take
place. On the other hand, any impurity present in the water would at once
cause it to become an electrolyte, in which case it would be able to bring
about the chemical action.

Experiments have been made with a view to determine whether or not the
gases themselves undergo any dissociation in the moist condition or during the
process of drying. But the results so far only show that if any dissociation
takes place, the extent to which it occurs is beyond the limits of measurement
by any volumetric methods.

Oxygen £93

circumstances of great bodily prostration, where the necessary
oxygenation of the blood cannot take place on account of the
enfeebled action of the lungs.

Compressed oxygen acts upon the animal economy as a poison :
an animal placed in oxygen gas under a pressure of only a few
atmospheres quickly dies.

During the respiration of man, air is drawn into the lungs, and
is there deprived of 4 to 5 per cent, of its oxygen, and gains 3 to 4
per cent, of carbon dioxide. The oxygen that is withdrawn from
the inhaled air by means of the lungs is absorbed by the blood.
The power to absorb this oxygen is believed to reside in a crystal-
line substance contained in the corpuscles of the blood, called
hcemoglobin, with which it enters into feeble chemical union, form-
ing the substance known as oxyhamoglobin. This substance is
red, and imparts to arterial blood its well-known colour. During
its circulation in the system the oxyhaemoglobin parts with its
oxygen, and is reconverted into the purple-coloured haemoglobin.
Under normal conditions the whole of the oxyhaemoglobin is not
so reduced, for venous blood is found still to contain it to some
extent. The amount of carbon dioxide exhaled is diminished
during sleep, and to a still greater extent during hibernation.

Oxygen can be liquefied at very low temperatures by the appli-
cation of moderate pressure (see Liquefaction of Gases). It was
first liquefied in 1877 by Cailletet, and independently by Pictet.
Its critical temperature is — 118.8°, at which point a pressure of
58 atmospheres is required to bring about its liquefaction.

Liquid oxygen is a pale steel-blue, mobile liquid, which boils at
-182.5°. Its specific gravity at -182.5° *s I-I3I5- The liquid
expands when warmed much more rapidly than gases do for the
same increment of temperature, and its density diminishes in pro-
portion, thus —

At -182.5° density =1.124.
„ -139° „ = -877-
„ -134° „ = -806.
,? -129° „ = .755°

Liquid oxygen is strongly magnetic. If a quantity of the liquid
be placed in a dish between the poles of a powerful electro-magnet,
the liquid will be drawn up to the magnet the instant the latter is
excited.

Inorganic Chemistry

ISOMERISM— POLYMERISM— ALLOTROPY.

Isomerism. — It is frequently found that two different compounds have the
same composition ; that is, their molecules are composed of the same number
of the same atoms, and yet the substances have different properties. Such
compounds are said to be isomeric, the one is an isomer of the other, and the
phenomenon is called isomerism. Cases of isomerism are so numerous among
the compounds of carbon (i.e. in the realm of organic chemistry, see Carbon,
p. 295), that it has been found convenient to classify them. The term
isomerism, therefore, is frequently restricted to cases in which the compounds
have the same percentage composition, the same molecular weight, and belong
to the same chemical type or class of substances. Thus, the two compounds
dimethyl benzene and ethyl benzene are both expressed by the formula
C8H10. The molecules in each case contain 8 atoms of carbon and 10 atoms
of hydrogen, they therefore have the same molecular weight and the same
percentage composition ; and as they both belong to the same type or family,
they are said to be isomeric with each other. The difference in the properties
of these compounds is due to a difference in the arrangement of the atoms
within the molecules, and this difference is expressed in their formulae in the
following manner : —

Dimethyl benzene, C6H4(CH3)2. Ethyl benzne, C6H5(C2H5).

Different compounds having the same molecular weight and the same per-
centage composition, but which do not belong to the same family of compounds,
are distinguished as metamers. Thus, the two compounds acetone and allyl
alcohol are each expressed by the formula C3H6O. They have the same
molecular weight and the same percentage composition, but belong to two
widely different types of compounds; they are therefore called metameric
compounds. The difference between them is again due to a difference in
molecular structure, and they are distinguished by formulae which convey this
difference, thus : —

Acetone, COfCHg)^ Allyl alcohol, C3H5(HO).

Polymerism. — This term is employed to denote those cases in which dif-
ferent compounds belonging to the same family have the same percentage
composition, but differ in molecular weight ; that is to say, their molecules are
composed of the same elements, which are present in the sa.me proportion ; but
they do npt contain the same actual numbers of the various atoms, and therefore
have different weights. Thus, the compounds ethylene (C2H4), propylene
(C3H6), butylene (C4H8), belong to the same family, and have each the same
percentage composition, but they differ in molecular weight. These sub-
stances are said to be polymers of one another.

Allotropy may be regarded as a special case of polymerism. In its widest
sense the term is sometimes used to denote polymerism in general, but it is
usually restricted to those instances of polymerism which are exhibited by
elementary bodies only. Many of the elements are capable, under special

Ozone 195

conditions, of assuming such totally different habits and properties, that they
appear to be entirely different substances. Thus, the element sulphur, as
usually seen, is a primrose-yellow, opaque, solid substance, extremely brittle,
and readily dissolved by carbon disulphide. Under certain circumstances it
may be made to appear a totally different thing ; it is then a translucent amber-
coloured substance, soft and elastic like indiarubber, and insoluble in carbon
disulphide ; it is still sulphur, and sulphur only. Phosphorus, again, as usually
known, is a nearly colourless, translucent, wax-like solid, which melts at a
temperature only slightly above that of the hand, and which takes fire a few
degrees higher; it is also extremely poisonous. Under special influences
phosphorus can be made to assume the following properties : — A dark reddish-
brown powder, resembling chocolate, which may be heated to 250° without
taking fire, and which is non-poisonous. The substance is still phosphorus,
and phosphorus only. This property possessed by certain of the elements of
appearing in more than one form, of assuming, as it were, an alias, is called
allot ropy ; the more uncommon form being spoken of as the allotropic modifica-
tion, or the allotrope of the other.

From a study of the best known instances of this phenomenon, it is believed
that allotropy, in all cases, is due to a difference in the number of atoms of the
element that are contained in the molecule. In the case of ozone, which is
the allotrope of oxygen, this is known to be the case. Ordinary oxygen
molecules consist of two atoms, while the molecule of ozone is an aggregation
of three oxygen atoms.

OZONE.

Molecular symbol, O3. Molecular weight = 48. Density =24.

History. — When an electrical machine is in operation a peculiar
and characteristic smell is noticed in its vicinity. The same smell
is sometimes observed in and about buildings, or other objects, when
struck by lightning. In 1785 it was observed by Van Marum that
when electric sparks were passed in oxygen, the oxygen acquired
this peculiar smell. Schonbein (1840) showed that the oxygen
obtained by the electrolysis of water also contained this substance
having a smell, and he gave to it the name ozone, signifying a smell.
Schonbein made a careful study of the substance, and found that
it might be obtained by various other methods. The more recent
work of Andrews, Soret, and Brodie has brought our knowledge of
the constitution of ozone to its present state.

Occurrence.— Ozone is present in the atmosphere irr extremely
small quantities (see Atmospheric Ozone).

Modes Of Formation.— (i.) Mixed with an excess of oxygen,
ozone is best obtained by exposing pure dry oxygen to the influence
of the silent electric discharge. This may be effected by means of
the instrument shown in Fig. 36, known as " Siemens' ozone tube.''

196

Inorganic Chemistry

It consists of two concentric glass tubes, A and B. Tube A is coated
upon its inner surface with tinfoil, which is brought into metallic
contact with the binding screw D, as shown in the figure. Tube B
is coated upon the outer surface, also with tinfoil, which is in
metallic connection with binding screw C. These two surfaces of
tinfoil are connected by means of their respective binding screws

FIG. 36.

to the terminals of a Ruhmkorf coil, and the slow stream of oxygen
which is admitted at E, and which passes along the annular space
between the two tubes, is there exposed to the action of the silent
electric discharge. A small portion of the oxygen so passing
becomes converted into the allotropic modification, and the mixture

FIG. 37.

of oxygen and ozone issues from the narrow tube at the opposite
end of the .apparatus.

For general purposes of illustration, a very simple arrangement
may be substituted for the above. It consists, as shown in Fig. 37,
of a straight length of narrow glass tube having a piece of platinum
wire down the inside, which passes out through the walls of the
tube near to one end, and is there sealed to the glass. A second

Ozone 197

platinum wire is coiled round the outside of the tube, and these two
wires are connected to the induction coil. On passing a slow
stream of oxygen through the tube, the issuing gas will be found to
be highly charged with ozone.

(2.) Ozone is also formed when an electric current is passed
through water acidulated with sulphuric acid. Thus, in the ordinary
electrolysis of water the oxygen evolved from the positive electrode
is found to contain ozone in sufficient quantity to be readily detected,
both by its odour and by other tests.

(3.) During many processes of slow oxidation at ordinary tempera-
tures, ozone is formed in varying quantities. Thus, when phos-
phorus is exposed to the air an appreciable amount of ozone is
formed. One or two short sticks of freshly scraped phosphorus
are for this purpose put into a stoppered bottle containing air, and
allowed to remain for a short time, when the air will be found to
contain ozone.

(4.) Ozone is also formed during the combustion of ether upon
the surface of red-hot platinum. When a spiral of platinum wire is
warmed in a gas-flame, and while hot is suspended over a small
quantity of ether contained in a beaker, the mixture of ether vapour
and air undergoes combustion upon the surface of the platinum,
which continues in an incandescent state so long as any ether
remains. During this process of combustion a considerable quantity
of ozone is formed. (See also Peroxide of Hydrogen.)

(5.) Ozone is formed during the liberation of oxygen in a number
of the reactions by which that gas is obtained ; thus, from manga-
nese dioxide and sulphuric acid the oxygen that is evolved contains
sufficient ozone to answer to the ordinary test. In the same way,
by the action of sulphuric acid upon barium peroxide or potassium
permanganate, this allotrope is present with the ordinary oxygen
that is evolved.

Properties. — As prepared by any of the methods described,
ozone is always mixed with a large excess of unaltered oxygen,
probably never less than about 80 per cent, of the latter gas being
present. Even in this state of dilution it has a strong and rather
unpleasant smell, which rapidly induces headache. When inhaled
it irritates the mucous membranes, and is rather suggestive of
dilute chlorine.

Ozone is a most powerful oxidising substance ; it attacks and
rapidly destroys organic matter : on this account ozonised oxygen
cannot be passed through the ordinary caoutchouc tubes, as these

198 Inorganic Chemistry

are immediately destroyed by it. It bleaches vegetable colours,
and most metals are at once acted upon by it. Even metals like
mercury, which are entirely unaltered by ordinary oxygen, are
attacked by ozone. Its action upon mercury is so marked in its
result, that the presence of exceedingly small traces of ozone can be
detected by it ; the mercury is seen to lose its condition of perfect
liquidity, and adheres to the surface of the glass vessel containing
it, leaving " tails " upon the glass. Ozone converts lead sulphide
(PbS) into lead sulphate (PbSO4), and liberates iodine from potas-
sium iodide —

This property is generally made use of for detecting the presence
of ozone, advantage being taken of the fact that iodine, when set
free from combination in the presence of starch, gives rise to a
deep blue-coloured compound, the reaction being one of extreme
delicacy. In order to apply this test for ozone, strips of paper are
dipped in an emulsion of starch to which a small quantity of potas-
sium iodide has been added. These papers may be dried and
preserved, and are usually spoken of as ozone test papers. When
one of these papers is moistened with water, and placed in air
containing ozone, the iodine is liberated from the potassium iodide,
and being in the presence of starch, the paper instantly becomes
blue by the formation of the coloured compound of starch. It will
be obvious that this method of testing for ozone can only be relied
upon when there is no other substance present which is able to
decompose potassium iodide ; for example, when testing for ozone
in the atmosphere, the presence of oxides of nitrogen or peroxide
of hydrogen (both of which are capable of liberating iodine, and
are liable to be present in the air) would materially vitiate the
result (see also Atmospheric Ozone). The above decomposition
of potassium iodide by ozone may be made use of as a test for
ozone in another way, which, although less delicate, is also less
likely to be vitiated by the presence of other substances. Blue
litmus papers are dipped into water which has been rendered very
feebly acid, and to which a small quantity of potassium iodide has
been added. The papers may be dried and preserved. On
moistening one of these papers with water and exposing it to
ozone the iodide is decomposed as in the former case, and the
potassium hydroxide which is formed, being a powerfully alkaline
substance, converts the colour of the litmus from red to blue.

Ozone 199

When heated to a temperature of about 250°, ozone is retrans-
formed into ordinary oxygen ; if, therefore, the ozonised gas
obtained by means of the Siemens' ozone tube be passed through
a glass tube heated by means of a Bunsen flame, the whole of the
ozone will be decomposed, and the issuing gas will therefore be
found to be without action upon the ozone test papers.

Ozone is also decomposed by certain metallic oxides, such as
those of manganese, copper, and silver. The action appears to be
one of alternate reduction and oxidation, the metallic oxide remain-
ing unaltered at the conclusion, thus —

The oxidising power of ozone is due to the instability of the mole-
cule and the readiness with which it loses an atom of oxygen,
leaving a molecule of ordinary oxygen, thus —

The oxygen molecule is comparatively inert, but the liberated atom
in its nascent state is endowed with great chemical activity. No
change of volume accompanies these processes of oxidation by
ozone, as the volume of the oxygen molecule (O2) is the same as
that of the ozone molecule (O3), the third atom of oxygen being that
which enters into new combination with the oxidised substance.

Ozone is soluble to a slight extent in water, imparting to the
solution its own peculiar smell. 1000 c.c. of water dissolve about
4.5 c.c. of ozone.

Under the influence of extreme cold, ozone condenses to liquid
having an intense blue colour. So deep is the colour, that a layer
of it 2 mm. in thickness is opaque. This liquid is obtained by
passing ozonised oxygen through a tube which is cooled by being
immersed in boiling liquid oxygen, which has a temperature of
— 182.5°. At this temperature the ozone liquefies, but most of the
oxygen with which it was mixed passes on. In a higher state of
purity it has been more recently obtained by first liquefying ozonised
oxygen, and then separating the more volatile oxygen by fractional
distillation. Liquid ozone boils at -119°. It is described by
Olszewski and Dewar as an extremely explosive substance.

Constitution of Ozone. — The fundamental difference between
ordinary oxygen and its allotrope ozone lies in the fact that the
molecule of the latter contains three atoms, while that of ordinary

200

Inorganic Chemistry

oxygen consists of only two. Ozone, therefore, is a polymer of
oxygen ; its molecule is more condensed, three atoms occupying
two unit volumes. This conclusion as to the constitution of ozone
has been arrived at from the consideration of a number of experi-
mental facts.

(i.) When oxygen is subjected to the action of the electric dis-
charge, it is found to undergo a diminution in volume.* This was
shown by Andrews and Tait by means of the tube seen in Fig. 38.
The tube was rilled with dry oxygen, which was prevented from
escaping by means of the sulphuric acid contained in the bent por-
tion of the narrow tube, which served as a manometer. When the
silent discharge was passed through the oxygen, a
contraction in the volume took place, indicated by
a disturbance of the level of the acid in the syphon.
When the tube was afterwards heated to about
300° C. and allowed to cool, the gas was found to
. have returned to its original volume, and to be
devoid of ozone. This could be repeated inde-
finitely, the gas contracting when ozonised and re-
expanding when the ozone was converted by heat
into ordinary oxygen. As only a very small propor-
tion of the oxygen was converted into ozone, this
experiment alone afforded no clue as to the rela-
tion between the change of volume and the extent
to which this conversion took place.

(2.) A small sealed glass bulb, containing a solu-
tion of potassium iodide, was placed in the tube
before the experiment. The oxygen was ozonised,
and the usual contraction noticed. The bulb was
then broken, and on coming in contact with the ozone present
the potassium iodide was decomposed, iodine being liberated.
No further contraction, however, followed ; and, further, when the
tube was subsequently heated to 300° and cooled, the gas suffered
no increase in volume. By carefully estimating the amount of
iodine that was liberated by the ozone, the actual amount of oxygen
which had caused this liberation could be determined according
to the equation —

2KI-hH2O + 0 = I2 + 2KHO,

and it was found that the volume of oxygen so used up was exactly
* " Chemical Lecture Experiments," new ed., Nos. 63, 64,

FIG. 38.

Ozone

201

equal to the contraction which first resulted on the ozonisation of
the oxygen.

These facts proved that when potassium iodide was oxidised by
ozone a certain volume of ordinary oxygen was liberated, which
was equal to the volume of ozone ;
and a certain volume was used up,
which was equal to the original
contraction.

These facts were explained by
the supposition that ozone was repre-
sented by the molecular formula O3 ;
and its action upon potassium iodide

may be expressed as follows — HIHV :fi-^4\ C

2KI + H2O + O3= O2 + I2 + 2KHO.

(3.) To prove the correctness of
this supposition, however, it was
necessary to learn the exact relation
between these two volumes. This
Soret did, by making use of the pro-
perty possessed by turpentine (and
other essential oils) of absorbing
ozone without decomposing it ; and
he found that the diminution in
volume which took place by absorb-
ing ozone from ozonised oxygen was
exactly twice as great as the increase
in volume that resulted when the
same volume of ozonised oxygen
was heated.

This fact may be shown by means
of the apparatus, Fig. 39.* The
oxygen to be ozonised is contained
in the annular space between
the elongated hollow stopper, which
reaches nearly to the bottom, and
the outer tube. The turpentine is
contained in a little sealed thin glass tube d, almost capillary in bore,
which is held in position between four little projecting glass points a
and b upon the stopper and outer tube. The temperature is main-

* ^ewth, Trans. Chem. Soc., 1896, p. 1298.

FiG. 39-

2O2 Inorganic Chemistry

tained constant throughout the experiment by placing the apparatus
in melting ice. One wire from the induction coil is dipped into the
ice water, while the other passes into the dilute acid contained in the
stopper. When the electric discharge is passed a portion of the
oxygen is ozonised, resulting in a contraction in the volume which
is indicated by a rise of the liquid in the gauge ; when sufficient
contraction has taken place the discharge is interrupted, and the
contents of the capillary tube brought into contact with the gas.
This is done by a slight twist of the stopper, which thereby crushes
the little tube and throws out the' turpentine. Immediately a
further contraction takes place, due to the absorption of the ozone
by the reagent, and if the gauge be graduated it will be seen that
this second contraction is twice as great as the first.

(4.) If the molecule of ozone be correctly represented by O3, its
density will be 24, as against 16 for oxygen ; and its rate of diffu-
sion will be proportionately slower in accordance with the law of
gaseous diffusion (see Diffusion of Gases, p. 84). Soret found that
this was actually the case, and from his experiments the number 24
for the density of ozone receives conclusive confirmation (see also
p. 228)'.

CHAPTER III

COMPOUNDS OF HYDROGEN WITH OXYGEN
THERE are two oxides of hydrogen known, viz. : —

Hydrogen monoxide, or water . . , . H2Q
Hydrogen dioxide , . : , .v ' • S . H2O2

WATER.

Formula, H2O. Molecular weight = 18.02.

Until the time of Cavendish, water was considered to be an
elementary substance. Priestley had noticed that when hydrogen
and oxygen were mixed and inflamed, moisture was produced,
and he had also observed that
the water so obtained was some-
times acid. Cavendish showed
that the water was actually the
product of the chemical union of
hydrogen with oxygen, and he
also discovered that the acidity
which this water sometimes pos-
sessed was due to the presence
of small quantities of nitric acid ;
and he traced the formation of
this acid to the accidental pre-
sence of nitrogen (from the at-
mosphere) with which the gases

were sometimes contaminated.

Cavendish filled a graduated
bell-jar with a mixture of hydro- Fio. 40.

gen and oxygen, in the propor-
tion of two volumes of the former to one of oxygen, and he attached
to the bell-jar a stout glass vessel resembling the pear-shaped
apparatus shown in Fig. 40, which was perfectly dry and rendered

204 Inorganic Chemistry

vacuous. On opening the stop-cocks, gas entered the exhausted
tube, which was furnished at the top with two platinum wires. The
cocks were again closed and an electric spark passed through the
mixed gases, thereby causing their explosion, when the interior
surface of the previously dry glass vessel was found to be dimmed
with a film of moisture. On again opening the stop-cocks more
gas was drawn into the upper vessel, the same volume passing in
as originally entered the evacuated apparatus. This showed that
the two gases in their combination with each other had entirely
disappeared. By repeatedly filling the vessel with the mixed gases
and causing them to unite in this way, Cavendish succeeded in
collecting sufficient of the water to identify the liquid, and prove
that it was in reality pure water.

The more exact volumetric proportion in which oxygen and
hydrogen combine to form water has been determined by modern
eudiometric methods which have been developed from Cavendish's
experiment. Accurately measured volumes of the two gases are
introduced into a long graduated glass tube standing in the
mercurial trough and provided with two platinum wires, by means
of which an electric spark can be passed. The gases are caused to
unite by means of the spark, and the contraction in volume is
carefully observed. Fig. 41 shows the apparatus for this purpose.
The long glass tube A having a millimetre scale graduated upon it
and having two platinum wires sealed into the glass near the upper
and closed end, is completely filled with mercury and inverted in
the trough of the same liquid : this tube is known as a eudiometer.
A quantity of pure oxygen is then introduced into the tube, and
the volume occupied by the gas carefully read off upon the gradua-
tions. Seeing that the volume occupied by a given mass of gas is
dependent both upon the temperature and the pressure, each of
these factors has to be taken into account in the process of this
experiment. The temperature is ascertained by the attached
thermometer T. The pressure under which the gas is, will be the
atmospheric pressure at the time (ascertained by the barometer B
placed near the apparatus) minus the pressure of a column of
mercury, equal to the height of the mercury within the eudiometer
above the level of that in the trough. This height is obtained in
millimetres by carefully reading upon the graduated scale the level
of the mercury in the trough and the top of the column in the
tube, and the number of millimetres so obtained is deducted from
{he barometric reading. These observations are made by means of

Water

205

a telescope placed at such a convenient distance that the heat of
the body may not introduce disturbances.

The data obtained give the volume of gas at a particular tem-
perature, and under a pressure less than that of the atmosphere
By the process of calculation explained under the general pro-
perties of gases (p. 69), this is reduced to the standard temperature
and pressure, viz., o° and 760 mm.

A quantity of hydrogen is then introduced into the eudiometer,
considerably in excess of that required for complete combination

FIG. 41.

with the oxygen, and the volume again ascertained with the above
precautions and corrections.

The difference between the first and second reading will give the
volume of hydrogen which has been added.

The eudiometer is then firmly held down against a pad of caout-
chouc upon the bottom of the trough, and the gases fired by an
electric spark from a Ruhmkorff coil. A bright flash of light
passes down the tube, and on releasing it from the indiarubber bed,
mercury enters to fill the space previously occupied by the gases
which have combined.

2o6 Inorganic Chemistry

On allowing the instrument to once more acquire the tempera-
ture of the surrounding atmosphere, the residual volume is read off
and corrected for temperature and pressure.

The following data have now been obtained : —

(i.) The volume of oxygen, corrected for temperature and
pressure.

(2.) The volume of mixed oxygen and hydrogen, corrected for
temperature and pressure.

(3.) The volume of residual hydrogen, corrected for tempera-
ture and pressure.

A concrete example will explain how the result is deduced from
these observations : —

Corrected volume of oxygen used 45-35

Corrected volume after the addition of hydrogen . . 256.05
Corrected volume of residual hydrogen .... 120. 10

256.05— 45. 35 = 210. 70= total volume of hydrogen employed.
210.70-120.10= 90. 60= volume of hydrogen which has combined with

45-35 volumes of oxygen.

•'• 45- 35:1: '•90.60:1.997.
.'. One volume of oxygen has combined with 1.997 volume of hydrogen

to form water.*

The volume composition of water may be shown by analytical
processes, as well as the synthetical
method described above. This decom-
position of water is most conveniently
effected by means of an electric cur-
rent. If the two terminals from a gal-
vanic battery are connected to two
pieces of platinum wire or foil, and
these are dipped into acidulated water,
bubbles of gas make their appearance
upon each of the wires. If these two
strips of platinum be so arranged in a
FIG. 42. bottle that all the gas evolved escapes

by a delivery-tube (Fig. 42), it will
be found that the gas explodes violently on the application to it

* In accurate experiments the volume occupied by the minute quantity of
water formed has to be taken into account, and a number of other corrections
have to be made that are not mentioned in this outline description of the

Water

207

of a lighted taper, showing it to be a mixture of oxygen and hy-
drogen. By modifying the apparatus in
such a way that the gas from each platinum
plate shall be collected in separate tubes,
so arranged that the volumes of the gases
can be measured, it is found that twice as
much hydrogen is evolved, in a given time,
as oxygen. A convenient form of volta-
meter is seen in Fig. 43, where the two
measuring tubes are suspended over the
platinum plates contained in a glass basin.
The electrode, which is connected with the
negative terminal of the battery, is the one
from which the larger volume of gas, viz.,
the hydrogen, is evolved, while the oxygen
is liberated at the positive plate.

When the volumes of the gases are care-
fully measured, it is found that they are not
exactly in the proportion of two of hydro-
gen to one of oxygen, but that the oxygen
is in deficit of this proportion. This is due,
in the first place, to the greater solubility
of oxygen in water than of hydrogen ; and,
secondly, to the formation of a certain
quantity of ozone during the electrolysis,
whereby there is a shrinking of volume in
the proportion of three to two. FIG. 43.

This process of electrolysis has already been partially explained on pp. 88,
89. In the dilute sulphuric acid employed the ions present are mainly H and

SO4 (see persulphuric acid, p. 424). The H ions convey the current to the
cathode, where they discharge their electricity and unite to form molecules of

ordinary hydrogen gas. The SO± ions travel to the anode, but instead of
regarding them as becoming discharged and then interacting with water
molecules to form H2SO4 with liberation of oxygen gas, the view now taken
is that the minute proportion of dissociated water molecules plays a part in
what actually goes on. In the neighbourhood of the anode there are a few

O ions due to this dissociation of water molecules, and it is believed that

these, and not the 804 ions, give up their charges and escape as oxygen
molecules ; while as fast as they do so more H2O molecules dissociate in

order to furnish sufficient H ions to establish equilibrium with the SO4 ions.

208 Inorganic Chemistry

Instead, therefore, of the chemical equation

2S04 + 2H20=2H2S04-|-02
we may substitute the ionic equation

Similarly when " secondary reactions " take place at the cathode, as, for in-
stance, when sodium chloride or sodium sulphate are electrolysed, it is believed

+ +

that as water is also slightly dissociated into H and OH ions, it is these H
ions which actually discharge and escape as hydrogen gas, and as fast as they

do so more water molecules dissociate so as to supply OH ions to establish equi-

librium with the K ions. The two conceptions of the mechanism of the action
are expressed by the equations —

2K + 2H20=2KHO+H2.

2K+2;H ,OH: = 2;K ,OH; + H2.

It has been recently shown* that the purest "electrolytic gas,"
as this mixture of hydrogen and oxygen is called, is obtained by
the electrolysis of pure barium hydroxide. Under these circum-
stances the oxygen contains no ozone or hydrogen peroxide.

The Volumetric Composition of Steam.— When a mixture of
oxygen and hydrogen is exploded in a eudiometer, we have seen
that a certain contraction of volume follows, due to the formation
of water by the uniting gases. The oxygen and hydrogen that
have entered into combination have disappeared as gases, the
volume of the resultant water being practically negligible. It is
important to know what relation exists between the volume of
the uniting gases and the volume of the product of their combina-
tion when in a state of vapour — that is to say, what volume of
steam is produced by the union of one volume of oxygen with
two volumes of hydrogen ; in other words, whether there is any
molecular contraction in the formation of steam.

To ascertain this, the mixed gases, in the exact proportions to
form water, must be made to combine under such circumstances
that the product shall remain in a state of gas or vapour, so that
its volume and that of the mixed gases may be measured under
comparable conditions. For this purpose a mixture of oxygen and
hydrogen, obtained by the electrolysis of acidulated water, is in-
troduced into the closed limb of the U-shaped eudiometer shown
in Fig. 44. t This tube is graduated into three equal divisions,
indicated by the broad black bands, and is furnished with two
platinum wires at the closed end. It is also surrounded by an
outer tube, so that a stream of vapour from some liquid, boiling

* Baker, Jour. Chem. Soc., April 1902.

+ See Experiments Nos. 74 and 75, " Chem. Lecture Experiments," newed.

Water

209

above the boiling-point of water, can be made to circulate. The
most convenient liquid for the purpose is amyl alcohol, which
boils at 130°. In this way the eudiometer and the contained gases
will be maintained at a constant temperature high enough to keep
the water formed by their combination in the state of vapour.

The amyl alcohol is briskly boiled in the flask, and its vapour is
led into the tube surrounding the eudiometer. The temperature of
the mixed gases is thereby raised to 130°, and they occupy the

FIG. 44.

three divisions of the tube when the mercury in the open limb is
at the same level, that is, when the gases are under atmospheric
pressure. The amyl alcohol vapour leaves the apparatus by the
glass tube at the bottom, and is conveyed away and condensed.
An electric spark is then passed through the gases by means of the
induction coil. (In order to prevent the mercury from being
forcibly ejected from the open limb of the U-tube at the moment of
explosion, an additional quantity of mercury is poured in, and the
open end is closed by the thumb when the spark is passed.) On
bringing the enclosed gas again to the atmospheric pressure, by

2IO

Inorganic Chemistry

adjusting the level of the mercury until it is once more at the same
height in each limb, it will be found that the mercury in the eudio-
meter is now standing at the second band ; that is to say, the three
volumes of gas originally present have now become two volumes of
steam. This condensation is expressed in the molecular equation —

The Gravimetric Composition of Water.— Having learned
the composition of water by volume, and knowing also that the
relative weights of equal volumes of oxygen and hydrogen are as
15.88 : i, the composition by weight can readily be calculated, thus —

1 volume of oxygen = 15.88

2 volumes of hydrogen = 2.00

17.88
17.88 parts by weight of water are composed of 2.00 parts by

FIG. 45.

weight of hydrogen and 15.88 parts of oxygen, or, expressed centesi-
mally, we have —

Oxygen .
Hydrogen

88.8

11.2

100.0

The composition of water by weight has been experimentally
determined with great care by a number of chemists.

The apparatus shown in Fig. 45 represents the method employed
by Dumas (1843). When copper oxide is heated in a stream of

Water 2 1 1

hydrogen, the copper oxide is deprived of its oxygen, which unites
with the hydrogen to form water —

Dumas' method is based upon this reaction. A weighed quantity
of perfectly dry copper oxide was heated in the bulb A, in a current
of hydrogen generated from zinc and sulphuric acid in the bottle H,
and rendered absolutely pure and dry by its passage through a
series of tubes containing absorbents. The water formed by the
union of the hydrogen with the oxygen of the copper oxide was
collected in the second bulb, B, previously weighed ; and the un-
condensed aqueous vapour which was carried forward in the stream
of hydrogen was arrested in the weighed tubes which follow. The
increase in weight of the bulb B and the weighed tubes gave the
total weight of water produced ; while the loss of weight suffered
by the copper oxide gave the weight of oxygen contained in that
water. The difference between these two weights is the weight
of the hydrogen that entered into combination with the oxygen.

As a mean of many experiments it was found that in the forma-
tion of 236.36 grammes of water the oxygen given up by the
copper oxide was 210.06 grammes.

236.36-210.06 = 26.30,
therefore 236.36 grammes of water were made up of

Hydrogen = 26.30
Oxygen =210.06

236.36
The ratio of hydrogen to oxygen is therefore as 2 : 1 5.88.

Hydrogen prepared from zinc and sulphuric acid is liable to contain traces of
(i.) Hydrogen sulphide. This is absorbed in the first tube containing

broken glass moistened with a solution of lead nitrate.

(2.) Arsenic hydride (absorbed in the second tube, filled with glass

(3. ) Hydrogen phosphide ( moistened with silver sulphate.

/"absorbed in the third tube, containing in one limb
(4.) Sulphur dioxide) pumice moistened with a solution of potassium
(5,) Carbon dioxide 1 hydroxide, and in the other fragments of solid

I potassium hydroxide.

Tubes 4, 5, 6, and 7, containing solid potassium hydroxide and phosphorus
pentoxide (the two latter being placed in a freezing-mixture), are for the pur-
pose of withdrawing every trace of aqueous vapour. Tube 8 was weighed before

212

Inorganic Chemistry

and after the experiment in order to test the absolute dryness of the hydrogen
that entered the bulb. In order to get rid of dissolved air, the dilute sulphuric
acid used was previously boiled. Tubes 9, 10, n were weighed both before and
after the experiment ; while tube 12, which was not weighed, was placed at the
end to prevent any absorption of atmospheric moisture by the weighed tubes.
Since the time of Dumas this subject has been reinvestigated by other
experimenters, who have introduced various modifications into the process ;
thus, with a view to finding the weight of hydrogen directly and of eliminating
many of the possible sources of error arising from the presence of impurities in

ftb. 46.

the hydrogen, the hydrogen has been absorbed by palladium. The metal so
charged with hydrogen can be weighed before and after the experiment, and
the actual weight of hydrogen used directly ascertained.

Most recently the matter has been investigated by Scott and Rayleigh, and
the results obtained show only the slightest departure from the numbers
obtained by Dumas.

Properties Of Water.— Pure water is a tasteless and odourless
liquid. When seen in moderate quantities it appears to be colour-
less, but when viewed through a stratum of considerable thickness
it presents a beautiful greenish-blue colour. This colour may be

Water

seen by filling a horizontal tube about 15 feet long with the purest
water, and passing a strong beam of light through it. It may also
be perceived by directing a ray of light through a tall cylinder of
water in the manner shown in the figure, andcausing it to be reflected
up through the water from the surface of a layer of mercury at the
bottom ; the immerging ray, being then reflected upon a screen,
shows the characteristic colour of the water. By intercepting the
ray by a hand mirror at A, the white light can be thrown upon the
screen as a contrast to the greenish-blue tint.

Aitkin has recently shown that the presence of extremely finely
divided suspended matters in water will give to the liquid the appear-
ance of a blue colour. Thus, in tanks where water is being softened
by the addition of milk of lime, after the bulk of the precipitated chalk
has settled, and only the finest particles still remain suspended in
the liquid, it is often noticed that the water appears to have a rich
blue colour. The wonderful blue colour of the waters of many of
the Swiss lakes is probably due in part to this optical phenomenon
as well as to the intrinsic colour of the water, "\yhen a mass of
pure snow, such as falls in high mountainous regions, is broken
open in such a way that the light is reflected from side to side
of the small crevice, the true greenish-blue colour of the water is
very manifest.

Water is compressible to only a very slight extent ; thus, under
an additional pressure of one atmosphere, 1000 volumes of water
become 999.95 volumes.

Small as this compressibility is, it exerts an important influence upon the
distribution of land and water upon the earth. It has been calculated that
owing to this compression, where the ocean has a depth of six miles, its surface
is lower by 620 feet than it would be if water were absolutely non-compressible;
and, calculated from the average depth of the sea, its average level is depressed
116 feet. The effect of this depression of the sea-level is that 2,000,000 square
miles of land are now uncovered which would otherwise be submerged beneath
the ocean.

Water is an extremely bad conductor of heat. A quantity of water
contained in a tube held obliquely may be boiled by the application
of heat to the upper layers without appreciably affecting the
temperature of the water at the bottom ; a fragment of weighted
ice sunk to the bottom will remain for a long time unmelted, while
the water a few inches above it is vigorously boiling. This low
conductivity for heat is shared in common by all liquids that are
not metallic. Indeed, Guthrie has shown, that water conducts heat

2 14 Tnorganic Chemistry

better than any other substance which is liquid at the ordinary
temperature, with the exception of mercury.

Steam. — Under a pressure of 760 mm., water boils at 100°
(see p. 128), and is converted into a colourless and invisible gas
or vapour. The visible effect that is observed when steam is
allowed to issue into the atmosphere is due to the condensation of
the steam in the form of minute drops of water. What is popularly
called steam is in reality, therefore, not steam, but an aggrega-
tion of small particles of liquid water. The invisibility of steam
is readily demonstrated by boiling a small quantity of water in a
capacious flask ; as the steam issues from the neck it condenses
in contact with the cool air* and presents the familiar appearance,
but within the flask it will be perfectly transparent and invisible.

ICG. — At a temperature of o° water solidifies to a transparent
crystalline mass. In the act of solidification the water expands
by nearly j^th of its volume, 10 volumes of water become 10.908
volumes of ice : solid water, therefore, is specifically lighter than
liquid water, and floats upon its surface. Water in this respect is
anomalous, for in the case of most other substances the solid form
is denser than the liquid. The disruptive force exerted by water at
the moment of freezing is the cause of the bursting of pipes and
other vessels containing water during winter; and it is also an
important factor in the economy of nature in the disintegration of
rocks and of soil. Under certain conditions water may be cooled
many degrees below o° without solidification taking place. Thus,
if a small quantity of water contained in a vacuous tube be care*
fully cooled without being subjected to vibration, its temperature
may be lowered to —15° without it solidifying ; a slight shock,
however, at once causes it to pass into the solid state, when its
temperature instantly rises to o° (see p. 137). Although the exact
temperature at which water freezes is liable to uncertainty from
this cause, the point at which ice melts is, under ordinary cir-
cumstances, constant, viz., o°. Under increased pressure ice
will melt at temperatures below o° ; thus, Mousson found that,
under a pressure of 13,000 atmospheres, ice melted at — 18°. The
melting-point of ice is lowered by about 0.0074° by each additional
atmosphere of pressure (see p. 137).

Between the temperatures of +4° and 100°, water follows the
ordinary laws that govern the expansion and contraction of liquids
due to change of temperature ; if water be cooled from 100°, it
gradually contracts until the temperature reaches 4°, Between

Water

this point and o° it forms a remarkable exception to the general
law, for, when cooled below 4°, it slowly expands instead of con-
tracting, and continues expanding until o° is reached, when it
solidifies. At 4°, therefore, water expands whether it be heated
or cooled ; consequently, at this point it is denser than at any
other temperature. This temperature is known as its point of
maximum density. (The most accurate observations fix the exact
point at 3.945°-)

. The following table shows the change of volume suffered by
water on being heated from o° to 8° :—

i. oooooo volumes at o° becomes

0-9999I5 » +2° „

0.999870 „ 4° „

0.999900 „ 6° „

i.oooooo „ 8°

One cubic centimetre of water, measured at its point of maxi-
mum density and at 760 mm., is the unit of weight of the metrical
system, and is called a gramme*

It is also at this temperature that water is taken as the unit for
comparison of the densities of other liquids and of solids ; thus,
when it is stated that the density or specific gravity of diamond
is 3.5, it is meant that diamond is 3.5 times as heavy as an equal
bulk of water measured at its point of maximum density.

The fact that water has a point of maximum density remote from
its freezing-point is one of far-reaching consequences in the opera-
tions of nature.

When a mass of water, such as a lake, is exposed to the influence
of a cold wind, the superficial layer of water is.cooled, and thereby
becoming specifically denser, it sinks to the bottom and exposes a
fresh surface. This in its turn has its temperature lowered, and in
like manner falls to the bottom. A circulation of the water in this
way is set up until the entire mass reaches a temperature of 4°.
At this point the further cooling of the surface-layer causes expan-
sion instead of contraction, and the colder water becoming speci-
fically lighter now floats upon the top, where it remains until it
congeals. If water continued to contract as its temperature was
reduced below 4°, the circulatory motion would continue until the
whole body of the water was cooled to o°, when solidification of the
entire mass would take place. The reason that certain very deep
* At the time this standard was first adopted, methods of measurement were
less refined than at present. In reality the gramme is not exactly the weight
of i c.c. of water at 'its point of maximum density.

2i6 Inorganic Chemistry

waters seldom or never freeze is because the duration of the cold
is not long enough to bring the temperature of the entire mass
of the water down to 4°, and until that is effected no ice can form
upon the surface.

The Solvent Power of Water — Water is possessed of more
general solvent powers than any other liquid ; that is to say, a larger
number of substances are dissolved by water than by any other
liquid. The solvent action of water upon gases, liquids, and solids,
in so far as it is shared by other liquids, has been dealt with under
the General Properties of Liquids (Part" I., chap. xiii.).

Water of Crystallisation.— When solid substances are dis-
solved in water, and the water afterwards evaporated, the dissolved
substance is frequently deposited in definite crystalline shapes.
Many salts owe their crystalline nature to the fact that a certain
number of molecules of water have solidified along with molecules
of the salt, each molecule of the salt being associated with a defi-
nite number of molecules of solid water. The water molecules
must be regarded as having entered into a feeble chemical union
with the salt molecule, but a union which is of a somewhat diffe-
rent order from that which holds together the atoms of oxygen and
hydrogen in the water molecules, or the atoms composing the salt
in the salt molecule (see p. 66). Thus copper sulphate crystallises
associated with five molecules of water, CuSO4,5H2O ; magnesium
sulphate with seven, MgSO4,7H2O. Water so associated with
crystals is known as water of crystallisation, and the compound
is called a hydrate.

Many salts are capable of crystallising with more than one defi-
nite number of molecules of water, depending upon the temperature
at which the crystallisation takes place : thus sodium carbonate,
crystallised at the ordinary temperature, has the composition
Na2CO3,10H2O ; while at temperatures between 30° and 50° the salt
that is deposited contains seven molecules of water, Na2CO3,7H2O.
Sodium chloride, crystallised from solution at —7°, has the compo-
sition, NaCl,2H2O ; while the crystals that are deposited at —23°
contain ten molecules of water, NaCl,10H2O.

In such cases as these, the particular crystalline form of the salt
differs with the different degrees of hydration.

Many crystalline salts, when exposed to the air, lose either some
or all of their water of crystallisation, and in so doing lose .their
particular geometric form. Thus the salt. Na2CO3,10H2O (ordinary
washing soda), when freely exposed, gradually loses its crystalline

Water 217

form and falls down to a soft white powder, which consists of small
crystals of another form, having the composition Na2CO3,H2O.
This process is known as efflorescence, the crystals being said to
effloresce. Other crystals undergo exactly the reverse change ; they
combine with moisture from the air, and pass into other crystalline
forms containing more water of crystallisation, or in some cases
they absorb sufficient moisture to cause them to liquefy. Such
substances are said to deliquesce. This property of certain salts is
made use of for withdrawing traces of water from either liquids or
gases. Thus, such a liquid as ether may be freed from dissolved
water by adding to it copper sulphate containing one molecule of
water of crystallisation, CuSO4,H2O ; this compound takes up water
and passes into CuSO4,5H2O} and thereby has the effect of drying
the ether. Gases in the same way are frequently dried by being
passed through tubes containing calcium chloride from which the
water of crystallisation has been removed. This substance absorbs
water with avidity, passing into the hydrated salt CaCl2,6H2O.

The characteristic colours of certain salts are in many cases
dependent upon the amount of water of crystallisation they contain.
Thus cobalt chloride, CoCl2,6H2O, is a pink salt. If it be gently
heated to 120° it loses its water and becomes CoCl2, which has a
rich blue colour. Solutions of this salt have been employed for
the so-called sympathetic- inks. The faint colour of the pink salt
renders words written upon paper with its dilute solution prac-
tically invisible ; but on warming the paper, and thereby expelling
the water from the salt, the written characters appear in a blue
colour, which again disappears as the salt is allowed to rehydrate
itself by exposure to the air.

One of the most striking examples of this change of colour
resulting from varying proportions of water of crystallisation is
seen in the salt magnesium platino-cyanide, which crystallises under
ordinary circumstances as a bright scarlet salt with seven molecules
of water, MgPt(CN)4,7H2O. When this salt is heated to about 50°
it loses two molecules of water, and is converted into a canary-
yellow salt, MgPt(CN)4,5H2O. If the temperature be raised to
100° the yellow salt becomes white by the loss of three more mole-
cules, the composition of the white salt being MgPt(CN)4,2H2O.
When a solution of the salt is carefully evaporated to dryness in
a dish and gently warmed, these colour changes will be rendered
evident ; and upon exposing the dried and white residue to the air,
or by gently breathing into the dish, the salt rehydrates itselft and

2i8 Inorganic Chemistry

is converted into the crimson compound having seven molecules
of water.

Many salts can have their combined water withdrawn by power-
ful dehydrating agents ; thus, if a crystal of copper sulphate ("blue
vitriol," CuSO4,5H2O) be immersed in strong sulphuric acid, the
acid abstracts four out of the five molecules from the salt, leaving
the nearly white salt CuSO4,HaO ; or when alcohol is added to a
solution of cobalt chloride, or to crystals of the salt, CoCl2,6H2O,
the alcohol abstracts water, and the solution becomes blue.

When salts containing water of crystallisation are heated, it
frequently happens that a portion of the water is more easily parted
with than the remainder. Thus copper sulphate, CuSO4,5H2O,
when heated to 100°, parts with four molecules of water, leaving the
salt CuSO4,H2O ; and in order to drive off this one remaining mole-
cule, the temperature must be raised above 200°. Zinc sulphate
(or white vitriol), ZnSO4,7H2O, in like manner loses six molecules of
water at 100°, but retains the seventh until a temperature of 240° is
reached. In order, therefore, to distinguish between the water that
is more firmly held and that which is readily parted with, the term
water of constitution is frequently applied to the former, and the
fact is sometimes expressed in notation in the following manner:—

CuSO4H2O,4H2O ; ZnSO4H2O,6H2O.

Natural Waters. — On account of the great solvent powers of
water, this compound is never found upon the earth in a state of
absolute purity ; even rain, as it falls in regions far removed from
the dirty atmosphere of towns, not only dissolves the gases of the
atmosphere, but also small quantities of those suspended matters
which are always present in the air. As soon as the rain reaches
the earth, the water at once exerts its solvent action upon the
mineral matter constituting the portion of the earth's crust over
which it flows, and through which it percolates, and the liquid is
rapidly rendered less and less pure as it travels on its course to
lake or ocean.

Natural waters may be broadly divided into two classes, based
upon the amount of dissolved impurities they contain. If the sub-
stances in solution are present in excessive quantities, or to such an
extent as to be perceptible to the taste, the water is said to be a
mineral water; while, on the other hand, waters that are not so
rich in dissolved impurities are known a&jfctti waters.

Natural Waters 219

Mineral Waters. — The most exaggerated examples of mineral
waters are to be found in sea.-water and in the waters of certain
lakes, which, having no outlet, are fulfilling the purpose of enormous
evaporating basins, in which the waters that flow into them are
undergoing evaporation and therefore concentration ; such, for
example, as the salt lakes of Egypt, the Elton lake in Russia, and
the Dead Sea. In waters of this description the total quantity of
dissolved solid matter is very considerable, and, as in the case of
the Dead Sea, is often deposited in crystalline masses round the
shores of the lake. The following table gives the total amount of
dissolved saline matter contained in zooo grammes of certain of
these waters : —

Irish Sea . . 33.86 I Dead Sea . . 228.57

Mediterranean Sea . 40.0 Elton Lake . . 271.43

As a typical example of a sea water, the composition of the
water of the English Channel may be quoted ; 1000 grammes of
this water contain —

Sodium chloride .

Magnesium chloride
Magnesium sulphate
Calcium sulphate .
Potassium chloride
Calcium carbonate

. 3.666

2. 296
1.406
. 0.766
. 0.033

Water . ...

35-2S5
964-74S

IOOO.OOO

Passing from these highly concentrated mineral waters, we find
a large number of spring waters which are classed as mineral, not
because the total quantity of foreign matter in solution is excessive,
but rather because they contain an abnormally large proportion of
a few special substances. Thus, large quantities of magnesium
sulphate and chloride are found in such springs as those at
Epsom and Friedrichshall. Others are found to contain consider-
able quantities of sodium sulphate and sodium carbonate ; while
those known] as chalybeate waters contain ferrous carbonate in
solution. Spring waters that are charged with unusual quantities of
soluble gases are likewise placed in the category of mineral waters,
such as the waters of Apollinaris and Seltzer, containing large
quantities of carbon dioxide ; and the sulphur springs at Harrogate
and Aachen, which hold in solution sulphuretted hydrogen as well
as alkaline sulphides.

22Q Inorganic Chemistry

Fresh Waters. — The purest form of natural water is rain-water.
The average weight of solid matter dissolved in rain-water, col-
lected in the country and in perfectly clean vessels upon which it
exerts no solvent action, is found to be 0.0295 Part m Ioo° parts
of water. Collected in or near towns, rain-water always contains
a larger amount of dissolved impurities, such as nitrates, sulphates,
ammoniaca! salts, and often considerable quantities of sulphuric
acid : it is the acid nature of the rain that causes so much damage
to stone buildings.

The nature and extent of the contamination that rain-water
suffers after it has fallen must obviously depend very largely upon
geographical and geological circumstances, and therefore there are
no special features that are distinctly characteristic of waters from
rivers, lakes, or springs.

Thus, the total solid impurity in 1000 parts of water from the
river Dee at Aberdeen is 0.057, while that contained in the
Thames is 0.30 parts.

The water of Loch Katrine only contains 0.032 part of solid
matter dissolved in 1000 parts, while that of Elton lake contains
as much as 271.43.

The same wide differences are also seen in spring waters from
different geological strata. Spring waters from granite and gneiss
rocks contain on an average 0.059 Part °f dissolved solid matter
in looo parts, while those from magnesian limestone average as
much as 0.665 Par^ As a broad general rule, river waters are
found to contain less solid matter in solution than spring waters,
and these in their turn less than deep well waters. Thus, com-
paring waters from different sources, and selecting only such
samples as are known to be free from pollution from either sewage
matter or other abnormal impurities, it will be seen that, with
regard to the dissolved solid matter they contain, they fall in the
following ordejr : —

Total Solid Impurity Dissolved in 1000 Parts of
Unpolluted Waters.

Rain-water (average of 39 samples) . . .0295

Rivers and lakes (average of 195 samples) . .0967

Spring waters (average of 198 samples) . . .2820

Deep well waters (average of 157 samples) . .4378

Natural Waters 221

Hardness of Water.— Certain of the salts that are very fre-
quently present as impurities in natural waters give to these
waters the property that is known as hardness. The chief com-
pounds that produce this effect are the salts of calcium and
magnesium. The term hardness is applied to such waters on
account of the difficulty of obtaining a lather, with soap, in the
ordinary process of washing. Pure soap may be regarded as a
mixture of the sodium salts of certain fatty acids (oleic, stearic,
palmitic, £c.), which are soluble in pure water. In the presence
of salts of calcium and magnesium the soap is decomposed, and
an insoluble curdy precipitate is formed by the union of the fatty
acid of the soap with the calcium and magnesium of the salts.
Until the whole of the hardening salts have in this way been
thrown out of solution, no lather can be obtained, and the soap is
useless as a cleansing agent ; but as soon as this point is reached,
the addition of any further quantity of soap at once raises a lather
on the water, and the soap is capable of acting as a detergent.
This process of precipitating the salts of calcium and magnesium
is known as softening, and in this instance the water is softened at
the expense of the soap.

Hard waters often become less hard after being boiled for a
short time, and this hardness which is so removed is termed the
temporary hardness. The degree of hardness which the water still
possesses after prolonged boiling is distinguished by the term
permanent hardness. The diminution of the total hardness of a
water by boiling is due to the fact that the soluble acid carbonates
of calcium and magnesium are decomposed during this process
into water, carbon dioxide (which escapes as gas), and the prac-
tically insoluble normal carbonates of these metals ; thus, in the
case of the calcium salt —

CaH2(CO3)2= H2O + CO2 + CaCO3.

When such a water is boiled, the calcium carbonate is thrown
down as a white precipitate, which gradually collects upon the
bottom of the containing vessel. The " furring " of kettles, and the
formation of calcareous deposits in boilers, is largely due to this
cause.

In the case of waters that are highly charged with calcium car-
bonate, held in solution by dissolved carbonic acid, this deposition
of calcium carbonate may even take place at the ordinary tempe-
rature, owing to the diffusion of the dissolved carbon dioxide into
the air. It is in this way that those remarkable, and often beauti-

222 Inorganic Chemistry

fully fantastic formations, known as stalactites, have been produced
in certain subterranean caves. Water charged with the soluble
calcium carbonate, in slowly dropping from the roof of such a cave,
loses a portion of its dissolved carbon dioxide, and, in consequence,
deposits a certain amount of the calcium carbonate which was in
solution. Each drop, as it slowly forms, adds its little share of
calcium carbonate to the deposit, which thereby gradually grows,
much as an icicle grows, as a dependent mass called a stalactite.
Whether the water that drops from the stalactite has deposited
the whole of its calcium carbonate, will depend largely upon the
time occupied by each drop in gathering and dropping ; if, as often
happens, the whole has not been precipitated, the remainder is
deposited upon the floor of the cave, and a growing column of
calcium carbonate, called a stalagmite, gradually rises from the
ground until it ultimately meets the stalactite.

Clark's Process for Softening- Water.— Waters whose hard-
ness is due to the presence of the carbonates of calcium and
magnesium can be deprived of their hardness by the addition to
them of lime. The amount of hardness is first estimated, and such
an amount of milk of lime is then added as is demanded by the
following equation : —

CaH2(CO3)2 + CaO = H2O + 2CaCO3.

In this way the soluble calcium salt is converted into the insoluble
normal carbonate, which settles to the bottom of the tank.

The salts, which are mainly instrumental in causing the per-
manent hardness, are the sulphates of calcium and magnesium.
The degree of hardness and its particular order, that is, whether
temporary or permanent, will obviously be determined entirely by
the particular geological formation from which the water is derived.

The Permutit System of water softening is a recently in-
vented process, based upon the fact that if water containing
calcium carbonate in solution is made to filter through a stratum
of certain silicates containing sodium silicate, the following inter-
action takes place : —

CaCO3+Na2SiO4=CaSiO4 + Na2CO3.

The insoluble calcium silicate remains in the filter while the
soluble sodium carbonate passes into the water. The silicate
actually employed is an artificially produced compound to which
the coined name permutit has been given. When in course of
time this substance becomes exhausted, it may be regenerated

Natural Waters 223

by passing a solution of brine through the filter ; this reacts with
the calcium silicate, reforming the sodium compound, while soluble
calcium chloride passes away. It is claimed that this process
of regeneration may be repeated indefinitely.

Potable Waters. — Undoubtedly the most important use to
which water is put is its employment as an article of food to man,
and since it has been proved beyond dispute that many virulent
diseases, such as cholera, typhoid fever, and others, are propagated
through the medium of drinking-water, it becomes a matter of the
greatest sanitary importance that the waters supplied for this pur-
pose should be as pure as possible. Excepting in very rare in-
stances, where poisonous mineral matters accidentally gain access
to drinking-water (as, for example, in the case of certain waters
which are capable of attacking, and to a slight extent dissolving,
the lead of the pipes through which they may be passed), the solid
matters that are usually found in waters are not injurious to health.
The living germs or bacilli, through whose agency zymotic diseases
are caused, cannot be detected in a sample of water by any direct
chemical analysis. A specimen of pure distilled water might
be artificially contaminated with such organisms so as to con-
stitute it a most virulent poison, and still chemical analysis
would fail to detect the danger, and the water would be pronounced
pure. Chemical analysis can, however, reveal the presence of
excrementitious matter, and also of the characteristic products re-
sulting from its decomposition : it can with certainty detect in the
water the evidence of recent contamination with sewage matters,
and it can also, with considerable precision, trace the evidences
of its having been so contaminated at an earlier stage of its history.
It cannot, however, distinguish between pollution with healthy and
with infected excreta, and therefore it is necessary to regard with
the greatest suspicion any water to which sewage has at any time
gained access. Waters that are made us6 of for drinking purposes
may be classified in the following order : —

/ i. Spring water.
Safe . . < 2. Deep well water.

( 3. Mountain rivers and lakes.

Suspicious | 4' Stored rain-water.

( 5. Surface water from cultivated land.

224 Inorganic Chemistry

HYDROGEN PEROXIDE.

Formula, H2O2.

Occurrence. — This compound is occasionally found in small
quantities in the atmosphere, and also in dew and rain.

Modes Of Formation.— ( i.) Hydrogen peroxide is produced in
small quantities during the burning of hydrogen in the air. If a
jet of burning hydrogen be caused to impinge upon the surface of
water, the temperature of which is not allowed to rise above 20°,
the water will be found, after a short time, to contain hydrogen
peroxide.*

(2.) This compound is also produced by the decomposition of
barium peroxide by carbonic acid. For this purpose a stream of
carbon dioxide is passed through ice-cold water, into which from
time to time small quantities of barium peroxide are stirred.
Baiiiim carbonate is precipitated, and a dilute aqueous solution
of hydrogen peroxide is obtained.

Ba02+H2C03 = BaC03+H202.

(3.) Barium peroxide may be decomposed by either hydrochloric,
sulphuric, silicofluoric, or phosphoric acid. Whichever acid be
employed, the barium peroxide, previously mixed with a small
quantity of water, is added gradually to the acid ; which, in the
case of either hydrochloric or sulphuric acid, should be diluted
with from five to ten times its volume of water. The temperature
of the mixture is not allowed to rise above 20°. Thus, in the case
of hydrochloric acid —

BaO2 + 2HC1 = BaCl2 + H2O2,

the soluble barium chloride is removed by the addition of sulphuric
acid, whereby barium sulphate is precipitated and hydrochloric
acid formed —

BaCl2 + H2SO4= BaSO4 + 2HC1.

The hydrochloric acid may be removed by adding a solution of
silver sulphate, which precipitates silver chloride, leaving sulphuric
acid in solution —

And, lastly, the free sulphuric acid is withdrawn by the addition of
barium carbonate —

When sulphuric acid is employed for the decomposition of barium
peroxide, the crystallised, or hydrated peroxide (BaO2, 8H2O), is
* See " Chemical Lecture Experiments," new ed. , p. 74.

Hydrogen Peroxide 225

most advantageous for the purpose. This salt, made^into a paste
with water, is gradually added to the diluted and cooled acid, until
the acid is nearly but not quite neutralised. The slight excess of
acid is removed by the addition of the exact quantity of barium
hydroxide (baryta- water) necessary to neutralise it, and the insoluble
barium sulphate is removed by filtration. On a large scale silico-
fluoric acid or phosphoric acid is usually employed, preferably the
latter, as it is found that small quantities of free phosphoric acid
in hydrogen peroxide greatly retard its decomposition.

(4.) Hydrogen peroxide is also readily obtained by decomposing
potassium peroxide by means of tartaric acid. The potassium
peroxide is added to a cooled strong aqueous solution of tartaric
acid, when potassium tartrate separates out, and an aqueous solu-
tion of hydrogen peroxide is obtained.

(5.) When small quantities of hydrogen peroxide are required
for the purpose of illustrating its properties, it is most conveniently
obtained by adding sodium peroxide to dilute and well-cooled
hydrochloric acid, whereby sodium chloride , and hydrogen per-
are formed, both of which remain in solution —

(6.) Hydrogen peroxide is formed in considerable quantity when
ozone is passed through ether floating upon water. Probably a
peroxidised compound of ether is first produced, which is then
decomposed by the water. This production of hydrogen peroxide
may readily be demonstrated by placing a small quantity of water
and ether in a beaker, and suspending into the vapour a spiral of
platinum wire which has been gently heated. The combustion of
the ether vapour upon the wire, whereby the latter is maintained
at a red heat, is attended with the formation of ozone, and this
acting upon the ether, as already described, results in the pro-
duction of hydrogen peroxide, which may be detected in solution
in the water.

(7.) In small quantities, hydrogen peroxide is produced when
moist ether is exposed to the action of oxygen, under the prolonged
influence of sunlight.

Properties.— The dilute aqueous solution of hydrogen peroxide,
obtained by the foregoing methods, is concentrated by evaporation
over sulphuric acid in vacuo. In the pure condition it is a colour-
less and odourless, syrupy liquid, having an extremely bitter and

226 Inorganic Chemistry

metallic taste. The specific gravity of the liquid is 1.4532. The
substance is extremely unstable, giving up some of its oxygen even
at temperatures as low as — 20°, and decomposing with explosive
violence when heated to 100°. Hydrogen peroxide bleaches
organic colours, but less rapidly than chlorine. When placed
upon the skin it destroys the colour, and gives rise to an irritating
blister. When diluted with water, and especially if rendered acid,
the compound is far more stable, and in this condition may be
preserved at the ordinary temperature for a considerable length of
time. When such an aqueous solution-is strongly cooled, it deposits
ice, and in this way, by the removal of the frozen water, the solu-
tion may be concentrated. Hydrogen peroxide- itself solidifies
between — 20° and — 23°. When heated the solution is decom-
posed into water and oxygen —

H202=H2O + O.

Owing to the readiness with which hydrogen peroxide gives up
the half of its oxygen and is converted into water, its properties
are generally those of a powerful oxidising agent. It liberates
iodine from potassium iodide ; it converts sulphurous acid into
sulphuric acid, and oxidises lead sulphide into lead sulphate. Its
action upon lead sulphide is made use of in restoring something
of the original brilliancy to oil paintings that have become dis-
coloured. The " white-lead " used in oil paints is gradually con-
verted into lead sulphide when such paintings are exposed to air,
especially the air of towns, which is liable to contain small
quantities of sulphuretted hydrogen. Lead sulphide being black,
the picture slowly assumes a uniformly dark colour. When such
a discoloured picture is washed over with dilute hydrogen peroxide,
the black sulphide is oxidised into the white lead sulphate —

This compound is employed for bleaching articles that would
suffer injury by the use of other bleaching agents, such as ivory,
feathers, and even the teeth.

Hydrogen peroxide is also capable of oxidising hydrogen.
Thus when a dilute acidulated solution of the peroxide is electro-
lysed, oxygen is evolved at the anode, but no gas escapes from
the cathode ; the nascent hydrogen being oxidised to water —

Ffydrogen Peroxide 227

Hydrogen peroxide, in many of its reactions, appears to act as a
deoxidising agent ; thus, manganese dioxide in contact with this
substance is reduced to manganous oxide —

MnO2 + H2O2= MnO + O2 + H2O.

Similarly, silver oxide is reduced to metallic silver with the
evolution of oxygen —

Ag2O + H2O2 = Ag2 + O2 + H2O.

In like manner, when ozone is acted upon by hydrogen per-
oxide, a reaction takes place exactly analogous to that with silver
oxide, which will be the more obvious if the formula for ozone be
written O2O instead of O3, thus —

O2O + H2O2= O2+ O2 + H2O.

Although, in a sense, these reactions may be regarded as reduc-
ing^ or deoxidising, actions, in essence they are not different from
those which have been given as illustrative of the oxidising power
of hydrogen peroxide. It will be seen that they all depend upon
the readiness with which the compound parts with an atom of
oxygen, but that in these latter cases the oxygen that is so given
up is engaged in oxidising another atom of oxygen, contained in the
other compound. Thus, in the case of silver oxide, its atom of
oxygen is oxidised by the liberated oxygen from the hydrogen
peroxide, and converted into the complete molecule of oxygen.
By these reactions Brodie first demonstrated the dual, or di-
atomic, character of the molecule of oxygen.

When hydrogen peroxide is added to a dilute acidulated solution
of potassium dichromate, a deep azure-blue solution is obtained
(see Chromium), which affords a delicate test for this com-
pound. To apply the test, the dilute hydrogen peroxide is shaken
up with ether, a few drops of acidulated potassium dichromate
are then added, and the mixture again shaken. The blue com-
pound being more soluble in ether than in water, the ethereal
liquid will separate as a blue layer. In this way, the presence of
0.00025 grammes of hydrogen peroxide in 20 c.c. of water can be
detected.

Hydrogen peroxide is decomposed by contact with many sub-
stances which themselves do not combine with the oxygen ; thus
charcoal, finely divided palladium, platinum, mercury, and notably
silver, when brought into hydrogen peroxide, determine its decom-

228 Inorganic Chemistry

position into water and oxygen, the rapidity of the action being
increased if the liquid be made alkaline. The action is doubtless
catalytic, although in all cases the exact modus operandi is not
clearly understood. In the case of silver it is believed that silver
oxide (perhaps peroxide) is first formed, and then decomposed,

Ag2 + H2O2=H2O + Ag2O
Ag2O + H2O2 = H2O + O2 + Ag2.

When hydrogen peroxide is added to solutions of the hydroxides
of barium, strontium, or calcium, the peroxide of the metal is
precipitated —

Ba(HO)2 + H20.=2H20 + BaO2.

The compound is deposited in crystals having the composition

With the hydroxides of the alkali metals, the peroxide (which is
soluble in water) may be precipitated by the addition of alcohol ;
when in the case of sodium peroxide, crystals are obtained of
Na2O2,8H2O.

Hydrogen peroxide is a useful antiseptic ; it possesses the ad-
vantages of being free from smell, without poisonous or injurious
action upon the system, and of leaving as a residue, after having
furnished its available oxygen, only water.

The constitution of hydrogen peroxide is usually expressed by the formula
H-O-O-H, but the accumulating evidence that oxygen is capable of functioning
as a quadrivalent element has led to the view that its constitution is better

H\
represented by the formula ^O : O.

Those metallic peroxides which yield hydrogen peroxide on treatment with
dilute acids may also be regarded as similarly constituted, thus
Na\

)>O : O and Ba : O : O
Na/

Ba : O : O + 2HC1= BaCl2 + H2O : O

while those peroxides which yield oxygen (or its equivalent of chlorine) under
similar treatment, contain only divalent oxygen atoms, e.g.

O : Pb : 0 + 2HCl=PbCl2 + H2O + O.
Ozone may be regarded as peroxide of oxygen, and expressed by the formula

O - O
O : O : O instead of the more familiar formula \ / in which each atom is

represented as being divalent.

CHAPTER IV
NITROGEN

Symbol, N. Atomic weight = 14.01. Molecular weight = 28.02.

History. — Nitrogen was discovered by Rutherford in 1772. He
showed that when an animal is placed in a confined volume of air
for some time, and the air afterwards treated with caustic potash,
to absorb from it the carbon dioxide ("fixed air"), there still
remained a gas which was incapable of supporting either respira-
tion or combustion. He called the gas mephitic air. Scheele was
the first to recognise that this gas was a constituent of the air.
Lavoisier applied the name azote to the gas, to denote its inability
to support life. The name nitrogen, signifying the nitre-producer,
was suggested by Chaptal, from the fact that the gas was a con-
stituent of nitre.

Occurrence.— In the free state nitrogen is present in the atmos-
phere, of which it forms about four-fifths. Certain nebulce have
been shown by spectroscopic observation to contain nitrogen in
the uncombined condition. In combination, nitrogen is found in
ammonia, in nitre (potassium nitrate), and in a great number of
animal and vegetable compounds.

Modes of Formation. — (i.) Nitrogen is very readily obtained
from the atmosphere by the abstraction of the oxygen with which
it is there mixed.* This is conveniently done by burning a piece
of phosphorus in air, confined over water. The phosphorus in
burning combines with the oxygen, forming dense white fumes of
phosphorus pentoxide, which gradually dissolve in the water, and
nitrogen remains in the vessel. The nitrogen obtained in this way
is never quite pure, for the phosphorus becomes extinguished
before the oxygen is entirely removed. It is also admixed with
the other gases present in the atmosphere (argon, carbon dioxide,
&c. ; see Atmosphere), amounting in all to about I per cent, of
the total.

(2.) Nitrogen in a purer state can be prepared from the atmos-

* Experiments 254, 255, " Chemical Lecture Experiments," new ed.
229

230 Inorganic Chemistry

phere by passing a stream of pure air over metallic copper cok
tained in a combustion tube, and heated to redness in a furnace.
The air is contained in a gas-holder, and is passed through two
U -tubes, the first containing potassium hydroxide (caustic potash),
in order to absorb the carbon dioxide ; and the second filled with
fragments of pumice moistened with sulphuric acid, in order to
arrest the aqueous vapour. The purified air, on passing over the
heated copper, is deprived of the whole of its oxygen, cupric oxide,
CuO, being formed, while the nitrogen passes on and may be
collected. This gas contains small quantities of argon (p. 256).

(3.) Oxygen is rapidly absorbed by a solution of cuprous chloride
in hydrochloric acid ; a ready method, therefore, of obtaining
nitrogen from the air is to place a quantity of this solution in a
stoppered bottle, and shake it up with the contained air. The
colourless cuprous chloride solution quickly absorbs the oxygen,
becoming dark in colour, and being converted into cupric chloride,,
the nitrogen of the air remaining in the bottle —

Cu2Cl2+2HCl + 0 = H20 + 2CuCl2.

(4.) Nitrogen is obtained by heating a strong solution of ammo-
nium nitrite in a flask, the salt splitting up into water and nitrogen —

NH4NO2 = 2H2O + N2.

In practice it is found more convenient to employ a mixture of
ammonium chloride and sodium nitrite —

NH4Cl + NaNO2=NaCl + 2H2O + N2.

(5.) By heating a mixture of ammonium nitrate and ammonium
chloride, a mixture of nitrogen and chlorine is evolved ; the latter
gas may be absorbed, by passing the mixture through either milk
of lime or a solution of sodium hydroxide —

2NH4NO3+NH4C1 = 5N + C1+6H2O.

(6.) Nitrogen is also evolved when ammonium chromate, or a
mixture of potassium dichromate and ammonium chloride, is
heated—

(NH4)2Cr207 = Cr203+4H2Q_+N2,
or—

K2Cr2O7 + 2N H4C1 = Cr2O3 + 2KC1 + 4H3

Nitrogen

231

(7.) When ammonia is acted upon by chlorine it is decomposed,
the chlorine combining with the hydrogen to form hydrochloric
acid, and the nitrogen being liberated —

If the chlorine be passed into a strong solution of ammonia, the
hydrochloric acid which is produced combines with the excess of
ammonia, forming ammonium chloride ; thus —

8N H3 + 3C12 = 6N H4C1 + Ng.

The chlorine, after being washed by passing through water, is
bubbled through strong aqueous ammonia contained in a Woulf s
bottle. As each bubble of chlorine enters into the ammonia, the

FIG. 47.

combination is attended by a feeble yellowish flash of light, and a
rapid stream of nitrogen is evolved. The nitrogen, which carries
with it dense white fumes of ammonium chloride, should be scrubbed
by being passed through a second bottle filled with fragments of
broken glass moistened with water, and it can then be collected
over water in the ordinary way, as shown in Fig, 47.* In prepar-
ing nitrogen by this reaction it is very necessary that the ammonia
should be in considerable excess, otherwise there is liable to be
formed the dangerously explosive compound of nitrogen and chlo-
rine (see Nitrogen Trichloride).

* Experiment 261.

Inorganic Chemistry

Properties. — Nitrogen is a colourless gas without taste or
smell. It is slightly Mghter than air, its specific gravity being
°-973 (air — 0- One litre of the gas at o° C. and 760 mm. weighs
14 criths, or 1.250 grammes.

Nitrogen is only very slightly soluble in water, its coefficient of
absorption at o° C. being 0.020346.

Nitrogen will not burn, neither will it support the combustion of
ordinary combustibles. It is not
poisonous, but is incapable of sup-
porting respiration.

Nitrogen is an extremely inert
substance, combining directly, and
with difficulty, with only a very few
elements. Under the influence of
the high temperature of the electric
spark it can be made to unite directly
with oxygen (see p. 235). Certain
metals also combine directly with it,
forming nitrides. Thus, when lithium
or magnesium are heated in nitrogen,
they form respectively NLi3 and
N2Mg3. This reaction may be con-
veniently shown by means of the
apparatus seen in Fig. 48. A small
quantity of powdered magnesium is

B i il \ placed in a hard glass tube, which

// ^g^r;.,.i^fc>- x^— _, js connected to a long narrow tube
dipping into water, and a stream of
nitrogen is passed through. When
the air is all displaced the passage of

the nitrogen is stopped and the magnesium strongly heated. At a
red heat the nitrogen will be rapidly absorbed, and the water will
be seen to rise in the long tube.

This property of nitrogen of uniting directly with magnesium
was utilised in effecting the separation of the nitrogen of the air
from the small quantities of argon and other " inert gases " con-
tained in the atmosphere.

Although it is true that nitrogen in the elemental condition is an
inert substance, the element itself is in reality possessed of strong
chemical affinities. Indeed, the very inertness of its molecules,
may be regarded as an indication of the strong affinity between

FIG. 48.

Nitrogen . £33

the two atoms which constitute the molecule. Nitrogen enters
into the composition of an enormous number of compounds, and
its atoms must be regarded as possessing great chemical activity.
The formation of such compounds as the nitrides above mentioned
may be quoted as an illustration. Although elementary nitrogen
combines directly with comparatively few metals, and with most
of these only at somewhat high temperatures, these compounds
are readily produced if, instead of elementary nitrogen, nitro-
gen in combination with hydrogen (ammonia) be employed (see
p. 278).

The critical temperature of nitrogen is - 149°, and when cooled
to this point a pressure of 27.5 atmospheres causes its liquefaction.
Under ordinary atmospheric pressure the liquid boils at — 195.5° >
the gas, therefore, can be liquefied by the cold obtained by the
rapid evaporation of liquid oxygen (see p. 78).

CHAPTER V

OXIDES AND OXY-ACIDS OF NITROGEN
NITROGEN combines with oxygen, forming five oxides: —

(i.) Nitrous oxide (hyponitrous anhydride) . N2O.

(2.) Nitric oxide NO.

(3.) Nitrogen trioxide (nitrous anhydride) . N2O3.

(4.) Nitrogen peroxide * .... NO2 and N2O4.

(5.) Nitrogen pentoxide (nitric anhydride) . N2O6.

Three oxy-acids of nitrogen are known, corresponding to the
three oxides, Nos. i, 3, 5 : —

Hyponitrous acid . . . H2N2O2.
Nitrous acid . „ . . . HNO2.
Nitric acid . „ .. « . HNO3.

The most important of all these compounds, and the one from
which all the others are directly or indirectly obtained, is nitric
acid.

NITRIC ACID.

Formula, HNO3. Molecular weight =63.02.

History. — Nitric acid, or aquafortis, was a well-known and
valued liquid to the alchemists. Down to the time of Lavoisier

* This name is usually applied to this substance both at low temperatures
when its composition is expressed by the formula N2O4, and also at higher
temperatures when the molecules have dissociated into the simpler molecules
NO2. In the strictest sense, however, they may be regarded as two oxides,
and it has been suggested to name the one (N2O4) nitrogen tetroxide, and the
other nitrogen peroxide.

234

Nitric Acid 235

(1776) its true nature was not known ; he showed that oxygen was
one of its constituents, but as to its other components he was un-
certain. Its exact composition was determined by Cavendish.

Modes of Formation.— ( i.) When an electric spark is passed
through a detonating mixture of oxygen and hydrogen with which
a certain quantity of air or nitrogen is mixed, the water that is
produced by the union of the oxygen and hydrogen is found to
contain nitric acid. This fact was first observed by Cavendish in
the course of his investigations on the composition of water, when,
owing to the accidental admixture of air with the mixed gases,
oxygen and hydrogen, he found that the water resulting from the
union was sometimes acid.

The direct union of nitrogen and oxygen may be brought about
by allowing a series of electric sparks to pass between platinum
wires in a confined volume of air, contained in a glass globe, as
shown in Fig. 49. In a short time the air in the globe will become
distinctly reddish in colour, owing to the formation of nitrogen
peroxide. The rapidity of the formation of the red fumes will
be greatly increased by compressing the air within the globe by
means of a small compression pump, as indicated in the figure.

If a small quantity of water be introduced, and the contents of
the globe shaken up, the red gas will be seen to dissolve in the
water, which will then acquire an acid reaction, owing to the forma-
tion of nitric acid.

Similarly, when a jet of hydrogen is allowed to burn in air to
which additional oxygen has been added, considerable quantities
of nitrogen peroxide are formed. The hydrogen may be burnt
from a jet, surrounded by a glass tube, as shown in Fig. 50, into
which oxygen can be passed by means of the small bent tube at
the bottom. On holding a clean dry cylinder over the flame,
sufficient of the products of combustion will collect in a few seconds
to show the presence of nitrogen peroxide.

This direct union of atmospheric oxygen and nitrogen has
recently been made the basis of a manufacturing process. A
stream of air is caused to pass through the electric arc at a rate
sufficiently rapid to sweep away the products of the action and so
prevent their dissociation. The nitrogen peroxide which is formed
is condensed to the liquid state, and thereby separated from the
other gases, by passing the mixture through a refrigerator.

(2.) Nitric acid is formed when nitrogenous animal matter under-
goes slow oxidation in the air, in the presence of water and an

236

Inorganic Chemistry

alkali, the nitric acid combining with the alkali to form a nitrate.
In this way nitrates are found in the soil, and from the soil often
find their way into shallow well-waters of towns. In hot and rain-
less countries these nitrates are sometimes found as crystalline
deposits on the surface of the soil, as in Chili and India (see
Potassium Nitrate).

(3.) Nitric acid is prepared by acting upon potassium nitrate
(nitre-saltpetre} with sulphuric acid. The nitre is placed in a glass
retort, together with an equal weight of sulphuric acid, and the
mixture gently heated. The nitric acid readily distils over, and

FIG. 49-

FIG. 50.

may be collected in a cooled receiver. The residue in the retort
consists of hydrogen potassium sulphate —

KNO3 + H2SO4=KHSO4-f HNO3.

The acid so obtained is not entirely free from water, and contains
nitrqgen peroxide in solution, which imparts to it a yellowish-red
colour. To purify it, it is again distilled with an equal volume of
sulphuric acid ; and the redistilled acid is deprived of the last traces
of dissolved peroxide of nitrogen, by causing a stream of dry air to

Nitric Acid 237

bubble through it while slightly warm. Nitric acid so prepared
may contain as much as 99.8 per cent, of anhydrous acid, HNO3.

(4.) Nitric acid is an article of commercial manufacture. In this
process potassium nitrate is replaced by the sodium salt, as being
the cheaper material. The proportion of acid to sodium nitrate
employed was formerly arranged in accordance with the equation —

2NaNO3+ H2SO4= Na2SO4 + 2HNO3.

It will be seen that the whole of the hydrogen of the sulphuric
acid is thus replaced by the alkali metal derived from two molecules
of the nitrate, and that two molecules of nitric acid result.

This reaction takes place in two stages ; in the first we have —

(i) NaNO3 + H2S04 = NaHSO4 + HNO3.

And then, as the temperature is raised, the hydrogen sodium
sulphate reacts upon a second molecule of the nitrate, thus —

(2) NaNO3 + NaHSO4 = Na2SO4+HN03.

The temperature necessary to effect this second stage, however,
causes the decomposition of a certain quantity of the nitric acid —

And for this and other reasons, most modern manufacturers
work only to equation No. i.

The retorts usually employed for the manufacture of this acid
are large cast-iron stills, which are sometimes lined, either entirely
or in part, with fireclay, and which are built into a furnace in such
a manner as to allow of their being heated as uniformly as possible.
The charge of sodium nitrate (12 to 14 cwts.) and sulphuric acid is
introduced, 'and the vapours carried off through an earthenware
pipe (c, Fig. 51), connected to a series of earthenware pots, b, in the
manner shown in the figure. The last of these jars is connected
with a tower, filled with coke, down which water is caused to per-
colate, and any peroxide of nitrogen which escapes is thereby
absorbed. The most modern form of still is not cylindrical, as
shown in Fig. 51, but takes the shape of an enormous crucible
with a dome-shaped lid ; and is furnished with an exit pipe at the
bottom, from which the liquid sodium bisulphate is run off.

Properties.— Nitric acid is a colourless liquid having a specific
gravity of 1.53. It fumes strongly in the air, and has a peculiar
and choking smell. It is extremely hygroscopic, absorbing moisture

238

Inorganic Chemistry

from the air with great readiness. Nitric acid is an intensely
corrosive liquid : the strongest acid, when brought in contact with
the skin, causes painful wounds, while in more dilute conditions

it stains the skin and other organic materials a bright yellow
colour. A quantity of strong nitric acid thrown upon sawdust
causes it to burst into flame. When nitric acid is distilled it first

Nitric Add 239

begins to boil at 86°, at the same time it is partially decomposed
into water, nitrogen peroxide, and oxygen ; the distillate, therefore,
gradually becomes weaker, and the boiling-point gradually rises.
This continues until a certain point is reached, when both the
temperature of the boiling liquid and the strength of the distillate
remain constant. If, on the other hand, a weak acid be distilled,
the distillate gradually increases in strength, until, when the same
point is reached, the boiling liquid has again the same temperature.

This constant boiling-point is 120.5°, and the distillate which
comes over at that temperature contains 68 per cent, of HNO3.
Whatever the strength of the acid, therefore, on being boiled it
loses either nitric acid or water until the strength reaches 68 per
cent., and this liquid boils at 120° C. The specific gravity of this
acid at 15° is 1.414. It was formerly supposed that the acid of this
strength constituted a definite hydrate, but Roscoe has shown that
the strength of the acid is purely a function of the pressure, for by
varying the pressure under which the distillation is conducted,
acids of various compositions can be caused to distil at a constant
temperature. Mixed liquids of this nature are known as constant-
boiling mixtures, and are strictly analogous to constant-freezing
mixtures (page 155).

When nitric acid is mixed with water there is a rise in tempera-
ture and a contraction in volume, the maximum effect being pro-
duced when the mixture is made in the proportion of three molecules
of water with one molecule of acid.

Nitric acid is a powerful oxidising agent, on account of the readi-
ness with which it parts with oxygen. Elements such as sulphur
and phosphorus are oxidised into sulphuric and phosphoric acids ;
arsenious oxide into arsenic acid ; and many protosalts are con-
verted into persalts. It attacks a large number of metals, forming
in many cases the nitrate. Its action upon metals is often of
a complicated nature, and depends not only upon the particular
metal, but also upon the strength of the acid, the temperature,
and the presence of the saline products of the reaction ; thus,
when nitric acid acts upon copper, the following reaction takes
place — . .

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

It is found, however, that as the amount of copper nitrate accu-
mulates, the nitric oxide which is evolved is mixed more and more
largely with nitrous oxide, N2O, apd even with nitrogen.

240 Inorganic Chemistry

Again, when dilute nitric acid acts upon zinc, nitrous oxide is
produced, according to the following equation —

4Zn + 10HNO3=4Zn(NO3)2 + 5H2O + N2O.

When, however, strong nitric acid i.s employed, ammonia is formed,
which combines with the excess of acid —

4Zn + 9HNO3 = 4Zn(NO3)2 + 3H2O + NH3.

In some cases, as with copper and silver, the presence of nitrous
acid (either as an impurity in the nitric acid, or as a first product
of its attack upon the metal) is believed to be a necessary condition
of the action.

Owing to the strong oxidising properties of nitric acid, hydro-
gen is rarely isolated by the action of metals upon this acid, the
hydrogen which is displaced from the acid being converted into
water. With magnesium, however, free hydrogen is evolved.

The chief reactions of nitric acid may be broadly divided into
three classes : —

(i.) With metallic oxides its behaviour is in common with other
acids. It exchanges its hydrogen for an equivalent quantity of the
metal, forming a nitrate, with the elimination of water, e.g. —

Ag2O-t-2HNO3 = 2AgNO3 + H2O.

(2.) Reactions in which it acts as an oxidising agent ; as an
example, its action upon iodine, which is converted into iodic acid,
may be cited —

i I + 3HN03=HIO3+H2O + NO + 2NO2.

(3.) Actions in which hydrogen in an organic compound is
replaced by the elements NO2, with the elimination of H2O, no
gas being evolved. The conversion of cotton-wool, or cellulose,
C]2H20O10, into gun-cotton, or nitro-cellulose, C12H14O10(NO2)6, is
an illustration of this class of reactions —

C12H20010-I-6HN03 = 6H20 + C12H14010(N02)6.

Nitric acid is without action upon the so-called noble metals,
gold and platinum.

Commercial nitric acid, which is of a reddish colour, is liable
to contain many impurities : chlorine and iodic acid, derived from
the Chili saltpetre ; iron, sulphuric acid, and sodium sulphate,

Nitrogen Pentoxide 241

carried mechanically over from the retorts ; and nitrogen peroxide,
from the decomposition of the acid. From these it is purified by
redistillation.

Nitric acid is a monobasic acid ; the salts of which, known as
the nitrates, are for the most part readily soluble in water, and
crystallise in well-defined forms. They are all decomposed at a
high temperature, evolving oxygen and nitrogen peroxide, or oxy-
gen and nitrogen, leaving an oxide of the metal.

The presence of a nitrate in solution is easily recognised by the
following characteristic test. A solution of ferrous sulphate is first
added to the solution containing the nitrate, and concentrated sul-
phuric acid is then cautiously poured down the side of the test-
tube, held in a sloping position, so as to fall to the bottom without
mixing with the solution. The sulphuric acid acting upon the
nitrate liberates nitric acid ; this is reduced by the ferrous sulphate
to nitric oxide, which, dissolving in the ferrous sulphate, forms a
brown-coloured solution at the point where the two layers of liquid
meet (see Nitric Oxide).

When nitric acid is added to hydrochloric acid, a mixture is
obtained which is known by the name of aqua regia. This name
was applied to it by the alchemists on account of its power of dis-
solving gold. Aqua regia is used in the laboratory for dissolving
gold, platinum, and certain ores. Its solvent power depends upon
the free chlorine which is evolved from the mixture —

NITROGEN PENTOXIDE (Nitric Anhydride).
Formula, N2O5. Molecular weight = 108. 02.

Modes Of Formation.— ( i.) By withdrawing from nitric acid
the elements of water, by means of phosphorus pentoxide —

2HN03+P205=2HP03 + N206.

For this purpose the strongest nitric acid is cautiously added to
phosphorus pentoxide in a cooled retort, in the proportion de-
manded by the equation ; the mixture being made as far as possible
without rise of temperature. The pasty mass is then gently heated,
when the nitrogen pentoxide distils over, and, if collected in a well-
cooled receiver, at once crystallises.

(2.) The method adopted by Devi lie, who discovered this com-

242 Inorganic Chemistry

pound (1849), was by passing dry chlorine over dry silver nitrate
contained in a U-tube, which was kept at the desired temperature
by being immersed in a water-bath. The following equation ex-
presses the final result of the action —

Properties. — Nitrogen pentoxide is a white solid substance,
crystallising in brilliant prismatic crystals, which melt at 30°
with partial decomposition. Between 45° and 50° it undergoes
rapid decomposition, evolving brown fumes. It is a very unstable
compound ; when suddenly heated it decomposes with explosive
violence, and even at ordinary temperatures decomposition slowly
takes place. It absorbs moisture rapidly, and when thrown into
water it dissolves with the evolution of great heat —

N2O6+H2O = 2HNO3.

When nitrogen pentoxide is gradually mixed with nitric acid, a
compound is formed having the composition 2N2O6,H2O ; which
separates, on cooling, as a definite crystalline hydrate

NITROGEN PEROXIDE.

Formula, NO2 and N2O4. Molecular weight = 46.01 and 92.02.
Density =23. oo and 46.01.

Modes Of Formation.— ( i.) This compound may be prepared
by mixing one volume of oxygen with two volumes of nitric oxide,
and passing the red gas so obtained through a tube surrounded
by a freezing-mixture —

2NO.

(2.) The nitrates of certain metals, when heated, are decomposed
into nitrogen peroxide, oxygen, and an oxide of the metal ; thus,
if dry lead nitrate be heated in a retort and the gaseous products
of decomposition are conducted into a U-tube placed in a freezing-
mixture, the nitrogen peroxide collects in the tube —

Pb(N03)2 =

(3.) When arsenious oxide is gently warmed with nitric acid, a
mixture of nitric oxide, NO, and peroxide, NO2, is evolved, and if
this gaseous mixture be passed through a cooled tube, it condenses

Nitrogen Peroxide 243

to a blue liquid. On passing a stream of oxygen through this
liquid it loses its blue colour, and is converted into a yellowish
liquid which consists of nitrogen peroxide.

Properties.— At low temperatures nitrogen peroxide is a colour-
less crystalline compound. It melts at - 9°, but requires a tem-
perature as low as — 30° to solidify it. At a temperature slightly
above its melting-point the liquid begins to acquire a pale yellowish
tint, which rapidly deepens until at the ordinary temperature it is
a full orange colour. The liquid boils at 22°, and gives a vapour
having a reddish-brown colour. The colour of the vapour also
becomes deeper as its temperature is raised, until at 40° it is a
dark chocolate brown, and almost opaque. On allowing the
vapour to cool the reverse changes take place. This change of
colour, as the temperature rises, is accompanied by a steady
change in the density of the gas, as will be seen from the table : —

Tenure. Denshy.

26.7° 38.3 20.00

60.2° 30.1 50.04

100.1° 24.3 79.23

I35.00 23.1 98.96'

140.0° 23.00 100.00

The density required by the formula N2O4 is 46.04, while that
demanded by the formula NO2 is 23.02 ; hence as the temperature
rises a process of dissociation goes on in which N2O4 molecules
are broken down into molecules of the simpler composition.
At 140° this process is complete, and the gas is entirely re-
solved into NO2. It is believed that at low temperatures nitrogen
peroxide has the composition represented by the formula N2O4,
out that dissociation begins to take place even during the state of
liquidity, as indicated by the gradual change of colour ; and there-
fore at temperatures between the boiling-point of the liquid, viz.,
22°, and 140°, the gas consists of mixtures of molecules of NO2 and
N2O4. The calculated percentage of NO2 molecules, which the
gas contains at the temperatures at which the above densities are
taken, are given in the third column.

Nitrogen peroxide is decomposed by water. At low tempera*
tures, and with small quantities of water, nitric and nitrous acids
are the products of the action, thus —

N204+H20 = HN03+HN02.

244 Inorganic Chemistry

At the ordinary temperature, and in the presence of an excess of
water, the following reaction takes place —

3NO2+H2O = 2HNO3 + NO.

Gaseous nitrogen peroxide is incapable of supporting the com-
bustion of a taper. Phosphorus, when strongly burning and
plunged into the gas, continues its combustion with brilliancy,
the temperature of the burning phosphorus being sufficiently high
to effect the decomposition of the gas. Nitrogen peroxide is a
suffocating and highly poisonous gas, and even when largely
diluted with air rapidly produces headache and sickness.

Nitrogen peroxide unites directly with certain metals, giving rise to a re-
markable series of compounds, to which the name nitro-metals, or metallic
nitroxyls, may be given (Sabatier and Senderens).* Thus, when the vapour
of nitrogen peroxide is passed over metallic copper (obtained by the reduction
of copper oxide in a stream of hydrogen), the gas is rapidly absorbed by the
metal with considerable rise of temperature, and a solid brown compound is
formed. This substance is the copper-uitroxyl, and its composition is ex-
pressed by the formula Cu2NO2.

Copper-nitroxyl is a fairly stable compound, and is unacted upon by dry air.
It is decomposed by water and by nitric acid, hence in its preparation care
must be taken to free the nitrogen peroxide from these substances.

At a temperature of about 90° copper-nitroxyl is decomposed into copper
and nitrogen peroxide. If, therefore, a quantity of the compound be sealed
up in a bent glass tube, and the empty limb of the tube be immersed in a
freezing-mixture while the compound is gently warmed, the nitrogen peroxide
which is evolved will be condensed in the cold portion of the tube.

Similar compounds are formed with the metals cobalt, nickel, and iron.

Nitrous Aeid, HNO2.— This substance is not known in the pure
state. Even in dilute aqueous solution it rapidly decomposes into
nitric acid, nitric oxide, and water —

3HNO2=HNO3 + 2NO + H2O.

The solution of this acid sometimes acts as a reducing agent,
taking up oxygen from such highly oxidised compounds as per-
manganates or chromates and passing into nitric acid —

HNO2 + O = HNO3.

Under other conditions it exerts an oxidising action, as when it
bleaches indigo, or liberates iodine from potassium iodide, being

* Bulletin de la Socidtt Chimigue, September 1893.

Nitrous Acid 245

itself reduced to nitric oxide and water, with the elimination of
oxygen —

2HNO2 = 2NO + H2O + O.

The salts of nitrous acids, viz.;the nitrites, are stable compounds.
The alkali nitrites may be prepared by carefully heating the nitrates
above their fusion points —

KN03=KN02 + 0.
At a higher temperature the nitrite is also decomposed.

On the manufacturing scale the nitrate is reduced by fusion with
metallic lead —

NaNO3 + Pb = PbO + NaNO2.

Nitrites are decomposed by dilute acids evolving brown vapours,
and in this way are at once distinguished from nitrates.

Nitrogen Trioxide, N2O3.— This compound is obtained by
passing a mixture of nitric oxide and nitric peroxide through a
tube cooled to about - 20°, when it condenses in the form of a
bluish green liquid.

Nitrogen trioxide so obtained is stable only at low temperatures.
When the temperature is allowed to rise the liquid itself undergoes
dissociation into its two generators, and as the more volatile oxide,
NO, escapes, the green colour of the liquid gradually disappears,
leaving a yellow liquid which finally passes off as gaseous NO2.
The reaction —

is therefore a reversible reaction in which the change as read
from right to left is practically complete at ordinary temperatures.
There has long been considerable doubt as to whether N2O3
had any existence as a gas. All the ordinary chemical reactions
in which this compound was at one time thought to be evolved as
gas, were found in reality to yield mixtures of NO and NO2 ; thus
in the case of the action of nitric acid upon arsenious oxide —

It has now, however, been shown by Baker (Jour. Chem. Soc.,
Nov. 1907) that if the liquid N2O3 is rendered absolutely dry by
prolonged exposure to phosphoric oxide, it then vaporises com-
pletely without dissociating, yielding a vapour whose density was
in no case below 38 (that demanded by the formula N2O3), and
which in many experiments was considerably above this figure. This
latter fact would seem to indicate that under these conditions the
gas may to some extent be polymerised, probably into N4OG, cor-
responding to the analogous oxides of phosphorus and arsenic.

246 Inorganic Chemistry

NITRIC OXIDE.

Formula, NO. Molecular weight =30. 04. Density =15. 02.

History.— Nitric oxide was first obtained by Van Helmont.
Priestley, however, was the first to investigate this gas, which he
termed nitrous airt and which was employed by him in his analysis
of air.

Modes of Formation.— ( i.) This gas is obtained by the action
of nitric acid of specific gravity 1.2 upon copper or mercury. In
practice copper is always employed.* The action may be repre-
sented thus —

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

The gas obtained by this method is always liable to contain
nitrous oxide and even free nitrogen ; the amount of these im-
purities rapidly increasing if the temperature be allowed to rise,
and still more so as the amount of copper nitrate in solution
increases.

(2.) Pure nitric oxide is readily obtained by the action of nitric
acid upon ferrous sulphate. The reaction is best applied by gene-
rating the nitric acid from potassium nitrate and sulphuric acid in
the presence of ferrous sulphate. A mixture of the two salts, in
the proportion of about one part of nitre to four of ferrous sulphate,
is introduced into a flask, with a small quantity of water. Strong
sulphuric acid is dropped upon the mixture by means of a drop-
ping funnel, and the mixture gently warmed, when a steady stream
of pure nitric oxide is evolved —

2KNO3+5H2SO4 + 6FeSO4=2HKSO4+3Fe2(S04)3+4H20 + 2NO.

A precisely similar result may be obtained by the reduction of
potassium nitrate by means of ferrous chloride in the presence of
hydrochloric acid, thus—

KNO3+3FeCl2 + 4HCl = 3FeCl3 + KCl + 2H2O + NO.

Properties. — Nitric oxide is a colourless gas, having a specific
gravity of 1.039. When brought into the air, it combines with the
atmospheric oxygen, forming red-brown vapours, consisting of

* Experiment 314, " Chemical Lecture Experiments," new ed.

Nitric Oxide

247

nitrogen peroxide, the combination being attended with a rise
of temperature. The formation of these red fumes in contact
with oxygen is characteristic of this gas, thereby distinguishing
it from all other gases. This property of nitric oxide renders
it impossible to ascertain whether this gas has any smell, or
is possessed of any toxicological action. Nitric oxide is only
very sparingly soluble in water. It is the most stable of all the
oxides of nitrogen, being able to stand a dull red heat without
decomposition. It is not a supporter of combustion. A lighted
taper, or a burning piece of sulphur, when introduced into the gas,
are extinguished. If the temperature of the burning substance is
sufficiently high to decompose the gas, com-
bustion then continues at the expense of the
liberated oxygen : thus, if a piece of phos-
phorus, which is freely burning in the air, be
plunged into this gas, it continues its com-
bustion with great brilliancy ; if, however,
the phosphorus be only feebly burning when
thrust into the gas, it is at once extinguished.
A mixture of carbon disulphide vapour and
nitric oxide, obtained by allowing a few drops
of the liquid to fall into a cylinder of the gas,
burns, when inflamed, with an intensely vivid
bluish flame, which is especially rich in the
violet or actinic rays, and has on this ac-
count been sometimes employed by photo-
graphers to illuminate dark interiors. Nitric

oxide is soluble in a solution of ferrous sulphate, forming a dark-
brown solution, containing an unstable compound of ferrous sul-
phate and nitric oxide, 2FeSO4,NO. This compound is readily
decomposed by heat, nitric oxide being evolved. • By means of this
reaction, nitric oxide may be separated from other gases. Nitric
oxide is a difficultly liquefiable gas, its critical temperature being
-93.5°: at this temperature a pressure of 71.2 atmospheres is
required to liquefy it.

The composition of nitric oxide may be proved by heating a
spiral of iron wire by means of an electric current in a measured
volume of the gas (as shown in Fig. 52).* As the metal becomes
red hot the gas is gradually decomposed and the oxygen combines

FIG. 52.

* No. 321, "Chemical Lecture Experiments,"

248 Inorganic Chemistry

with the iron to form ferric oxide. The residual nitrogen will
be found to occupy one-half the original volume.

Two vols. of nitric oxide, weighing 30.04
Contain i vol. of nitrogen, weighing 14.04

16.00= weight of i vol. of oxygen.

Therefore we learn that two volumes of nitric oxide consist of
one volume of nitrogen and one volume of oxygen united without
condensation.

NITROUS OXIDE (Hyponitrous anhydride, Laughing gas),
Formula, N2O. Molecular weight = 44. 08. Density =22. 04.

History. — This gas was discovered by Priestley, and called by
him dephlogisticated nitrous air.

Modes Of Formation. — (i.) Nitrous oxide is formed by the
reduction of nitric acid by certain metals, as zinc or copper, under
special conditions (see Nitric Acid). These reactions, however,
are never made use of for the preparation of the gas for experi-
mental purposes.

(2.) The most convenient method for obtaining this compound
is by the decomposition of ammonium nitrate. A quantity of the
dry salt is gently heated in a flask fitted with a cork and delivery-
tube. The salt rapidly melts and splits up into nitrous oxide and
water —

The heat should be carefully regulated, or the decomposition is
liable to become violent, in which case nitric oxide is also evolved.
Nitrous oxide being rather soluble in cold water, the gas should
be collected either over mercury or over hot water.

When the gas is to be used- for anaesthetic purposes, it should be purified
by being passed first through a solution of ferrous sulphate to absorb any nitric
oxide, and afterwards through caustic soda, to remove any chlorine which may
have been derived from the presence of ammonium chloride in the nitrate.

Properties. — Nitrous oxide is a colourless gas, having a faint
and not unpleasant smell, and a peculiar sweetish taste. Its
specific gravity is 1.52. The gas is somewhat soluble in water, its
coefficient of absorption at o° being 1.3052. The solubility rapidly

Nitrous Oxide 249

decreases as the temperature rises, as will be seen by the follow-
ing table (Carius) : —

i c.c. Water at c.c. N2O at o9 C.

760 mm. Dissolves and 760 mm.

At o° rv r- , , . . 1.3052
„ 10° .- ^ *-' , . 0.9196

„ 20° ...... 0.6700

,,25° . . . - . . 0.5962

The loss of gas during its collection over water in the pneumatic
trough, arising from its solubility in that liquid, is therefore greatly
lessened by using warm water. Nitrous oxide is much more
readily decomposed than nitric oxide ; a red-hot splint of wood is
instantly rekindled, and bursts into flame when plunged into the
gas. Phosphorus burns in it with a brilliancy scarcely perceptibly
less dazzling than in pure oxygen. If a piece of sulphur which is
only feebly burning be thrust into a jar of this gas, the sulphur is
extinguished, the temperature of the flame not being sufficiently
high to decompose the gas. When, however, the sulphur is
allowed to get into active combustion before being placed in the
gas, the combustion continues with greatly increased brilliancy.
In all cases of combustion in nitrous oxide, the combustion is
simply the union of the burning body with oxygen, the nitrogen
being eliminated. From its behaviour towards combustibles,
nitrous oxide might readily be mistaken for oxygen ; it can, how-
ever, be easily distinguished from that gas by the fact that when
added to nitric oxide it does not produce red vapours, whereas
when oxygen is mixed with nitric oxide these coloured fumes are
instantly formed.

When equal volumes of nitrous oxide and hydrogen are mixed
in a eudiometer, and an electric spark passed through the mix-
ture, the gases combine with explosion, water being produced and
nitrogen set free ; the volume of nitrogen so resulting being equal
to that of the nitrous oxide employed. This compound, therefore,
contains its own volume of nitrogen, and half its own volume of
oxygen. Nitrous oxide, when inhaled, exerts a remarkable action
upon the animal organism. This fact was first observed by Davy.
If breathed for a short time, the gas induces a condition of hysterical
excitement, often accompanied by boisterous laughter, hence the
name laughing gas. If the inhalation be continued, this is followed
by a condition of complete insensibility, and ultimately by death.

250 Inorganic Chemistry

On account of the ease with which the state of insensibility can be
brought about, this gas is extensively employed as an anaesthetic,
especially in dentistry.

Nitrous oxide is a gas which is moderately easily liquefied ; at
o° C. a pressure of thirty atmospheres is required to effect its
liquefaction.

Liquid nitrous oxide is colourless and mobile ; it boils at - 89.8°,
and when dropped upon the skin produces painful blisters.
When thrown upon water, a quantity of the water is at once con-
verted into ice ; mercury poured into a tube containing a small
quantity of the liquid is instantly frozen. An ignited fragment of
charcoal thrown upon the liquid floats upon the surface, at the
same time burning with brilliancy. If the liquid be mixed with
carbon disulphide, and placed in vacuo, the temperature falls
to - 140°. By strongly cooling the liquid, contained in a sealed
tube, Faraday succeeded in solidifying it ; this may also be
effected by the rapid evaporation of the liquid. The solid melts
at - 102.7°, and if placed upon the hand causes a painful blister j
in this respect it differs from solid carbon dioxide, which gasifies
without previous liquefaction.

Hyponitrous Acid, H2N2O? or N2(HO)2.

When a solution of potassium nitrate or nitrite is acted upon by sodium
amalgam the salt is reduced by the nascent hydrogen (evolved by the action
of the amalgam upon the water) with the formation of potassium hyponitrite-—

=2H2O-}-K2N2O2 or

The solution, which is alkaline, is neutralised with acetic acid, and silver
nitrate added, which causes the precipitation of yellow insoluble silver
hyponitrite, Ag2N2O2.

The acid itself is obtained from the silver salt by the action of dilute hydro-
chloric acid. By employing an ethereal solution of hydrogen chloride and
evaporating the solution, the hyponitrous acid is obtained in the form of white
deliquescent crystals. The pure acid is very unstable, exploding spontaneously
even below o° C. The aqueous solution when gently warmed is quickly
broken up into water and nitrous oxide.

Formerly nitrous oxide was regarded as containing nitrogen in the mono-
valent condition, and 'its constitution was represented by the formula
N-O-N. It is more probable, however, that the molecule contains the
two nitrogen atoms doubly linked, and that its formation by the decomposition
of hyponitrous acid is expressed by the equation —

N : N

Nitrosyl Chloride 251

NitrosylChloride,* NOC1. — This compound may be obtained by the direct
combination of nitric oxide with chlorine—

It is also formed by the action of phosphorus pentachloride upon potassium
nitrite, thus —

PC15+KN02 = NOCI+POC13+KC1.

Nitrosyl chloride is formed together with chlorine when a mixture of nitric
•ind hydrochloric acids is gently heated —

HNO3+3HC1=NOC1 + C12+2H2O.

Nitrosyl chloride is also readily prepared by the action of nitrosyl hydrogen
sulphate upon dry sodium chloride, thus —

(NO) HSO4 + NaCt= NOC1 + NaHSO4.

Properties.— Nitrosyl chloride is an orange-yellow gas, which easily con-
denses when passed through a tube immersed in a freezing-mixture, to an
orange-yellow liquid, which boils at about - 8°. It is decomposed by water
into nitrous acid and hydrochloric acid —

NOC1 + H20=HNO2+HC1.

In a similar manner it is decomposed by metallic oxides and hydroxides,
thus —

NOC1 + 2KHO= KN0.2+ KC1 + H20.

Nitrosyl chloride has no action upon gold and platinum, but it attacks
mercury with the formation of merctirous chloride and the liberation of nitric
oxide —

2NOCl + 2Hg=Hg2Cl2+2NO.

Nitrosyl perchlorate, NOC1O4, 2H2O, is obtained as a white crystalline
hydroscopic salt by the action of the mixed oxides of nitrogen (evolved by the
interaction of nitric acid and sodium nitrite) upon concentrated perchloric
acid.

Nitrogen Oxyfluorides. — Two of these compounds have been obtained.
Nitroxyl fluoride, NO2F, is obtained by the direct union of fluorine and nitric
oxide at the temperature of liquid oxygen (Moissan). It is a colourless gas
to a liquid which boils at -63.5. Water decomposes it into nitric and
hydrofluoric acid —

N02F+H2O=HN03+HF.

Nitrosyl fluoride, NOF, is produced by the action of nitrosyl chloride upon
silver fluoride (Ruff and Stauber)—

NOC1 + AgF = NOF + AgCl.

* Tilden has; shown that the compound represented by the formula NO2C1
does not exist.

CHAPTER VI
THE ATMOSPHERE

THE atmosphere is the name applied to the gaseous mixture
which envelops the earth, and which is commonly called the air.
The older chemists used the word air much as in modern times
the word gas is employed ; thus^ they spoke of inflammable air,
dephlogisticated air, alkaline air, and so on.

The air consists of a mixture of gases, the two chief ingredients
being nitrogen and oxygen. Lavoisier was the first to clearly
prove that oxygen was a constituent of the air, although Robert
Boyle and others before him had shown that air was absorbed by
metals in the process of forming a calx> and that the metal gained
weight as the calx formed. When the fact that the air was com-
posed of oxygen and nitrogen became established, various de-
vices were adopted to determine the proportion of oxygen in it.
Priestley's method was by means of nitric oxide. It depended
upon the fact that when nitric oxide is mixed with air it combines
with the oxygen, forming brown fumes which dissolve in the waten
A contraction in volume therefore takes place, from which the
volume of oxygen may be calculated. This method yielded results
which seemed to show that there was considerable variation in the
proportion of oxygen present in different samples of air, and the
idea arose that the wholesomeness or goodness of the air was
dependent upon the quantity of oxygen which it contained. Hence
arose the term eudiometry^ signifying to measure the goodness.
Cavendish, on the other hand, as the result of a large number of
experiments made by him, came to the conclusion that there was
no difference in the'samples of air that he experimented upon.

Since the time of Cavendish, eudiometric analysis has been
brought to a state of great perfection and accuracy by Bunsen,
Regnault, Frankland, and others. The conclusion to be drawn
from the extended researches of these chemists is, that although
the atmosphere certainly shows a remarkable uniformity of com-
position, there do exist perceptible, though very slight, variations

352

The Atmosphere

253

!n the amount of oxygen present at different places and at different
times. Samples of air collected from all parts of the globe, from

mid ocean, from high mountain peak, American prairie, and
crowded cities, show a variation in the proportion of oxygen rang-
ing from 20.99 to 20.86. Angus Smith has shown that in foggy

254 Inorganic Chemistry

weather the oxygen in the air in towns sometimes falls as low as
20.82. Samples of air taken from crowded theatres have been
found to contain as little as 20.28, while in many mines the amount
averages as low as 20.26.

The mean proportions of oxygen and nitrogen in the atmosphere
may be given as —

Oxygen 20.96 parts by volume.

Nitrogen* . . . . 79-Q4 „ „

100.00

The composition of the atmosphere by weight was determined
by Dumas and Boussingault (1841). In their method, air which was
freed from carbon dioxide and moisture was slowly drawn through
a glass tube containing a known weight of metallic copper, heated
to redness. The oxygen combined with the copper, forming copper
oxide, which was afterwards weighed, and the nitrogen passed into
a vacuous flask, and was also weighed. The apparatus as em-
ployed by Dumas is seen in Fig. 53. B is a glass flask having a
capacity of 10 to 15 litres, which was exhausted and then weighed.
It was then attached, as shown, to the tube T, containing a known
weight of metallic copper, and which was also exhausted. The
bulbs L contained a solution of potassium hydroxide, and the tubes
/j solid potash, for the removal of atmospheric carbon dioxide.
The bulbs O contained strong sulphuric acid, and the tubes / were
filled with pumice moistened with the same acid, by means of
which the moisture was withdrawn from the air. When the copper
was heated and the cocks partially opened, air, free from carbon
dioxide and moisture, was slowly drawn over the heated metal,
which was thereby converted into the oxide. At the conclusion
of the experiment the globe and the tube T were reweighed. The
nitrogen remaining in tube T was then pumped out and the tube
once more weighed. The difference between the two last weigh-
ings of the tube, added to the gain in weight suffered by the globe,
gave the nitrogen ; while the difference between the original and
final weights of the tube gave the increase of weight suffered by
the copper, that is, the amount of oxygen. The result of numerous
experiments gave the mean composition —

Oxygen . . » . . 23 parts by weight.

Nitrogen* ' •. . / •••*'{ . . 77 » „

100
* The small percentage of argon present is here included with the nitrogen.

The Atmosphere

255

A more modern method for estimating the amounts of oxygen
and nitrogen in the air, based upon the same principle, namely, the
absorption of the oxygen by heated metallic copper, is illustrated
in Fig. 54 (known as Jolly's apparatus). The sample of air to be
examined is allowed to enter the glass globe A (whose capacity is
about 100 c.c., and which has been previously exhausted) by means
of the three-way cock b. (The air is first dried, by being drawn
through tubes rilled with pumice moistened with sulphuric acid, on

FIG. 54.

its way into the apparatus.) The bulb is tnen surrounded by the
metal jacket B, which is filled with broken ice, and when the tem-
perature has fallen to o° the bulb is put into communication with
the tube d by means of the three-way cock. The tube g is then
raised or lowered, so as to bring the mercury in d to a fixed point
in the tube at m, and the tension of the enclosed air is ascertained
by the graduated scale behind tube g. The ice-jacket is then
removed, and the spiral of copper wire within the bulb is heated
to redness by the passage through it of an electric current. The

256 Inorganic Chemistry

copper combines under these conditions with the oxygen^ form-
ing copper oxide, thereby reducing the volume of the contained
gas. The globe is again cooled, and the tube g lowered to such
a position that when communication is once more made between
the globe and tube d, the mercury shall stand at the same point in.
From the observed tension of the gas before and after the
experiment, the volume relations of the two constituents can be
calculated. Thus, suppose the tension of the enclosed air to be
720.25 mm., and that of the residual nitrogen 569.28 mm., then for
I volume of air the reduction would be —

569.28

7~2^ = -7904 vols.

Therefore in 100 volumes the composition would be —

Nitrogen* = 79.04
Oxygen =20.96

100.00

Besides oxygen and nitrogen, the air contains variable quantities
of the ' following gases: aqueous vapour, carbon dioxide, argon,
hydrogen, ammonia, ozone, nitric acid. With the exception of
aqueous vapour, these substances are present only in relatively
small proportions, and with some of them the amount is liable
to considerable variation. Especially is this the case with the
aqueous vapour, as the amount of this constituent present at any
time is largely influenced by the temperature. The average com-
position of normal air may be taken as follows : —

Vols. per 1000.

Nitrogen . - , . . 769.6500

. Oxygen . , „ -» ,. 206.5940

Aqueous vapour . V ."" -V~ 14.0000

Argon t . -',-; .:..•;.. . 9-37QO

Carbon dioxide . . . . 0.3360

Hydrogen . „ * ,, . 0.0400

Ammonia . •"• / > » » 0.0080

Ozone . -v ,-. * » . 0.0015

Nitric acid . , . .>:• ... 0.0005

IOOO.OOOO

* The small percentage of argon present is here included with the nitrogen,
f The other four gases of the argon group taken together come to about
0.012 parts per 1000 (see page 270).

The Atmosphere 257

Aqueous Vapour. — For any given temperature there is a
maximum amount of aqueous vapour which a given volume of air
is capable of taking up : under these conditions the air is said to
be saturated with moisture at the particular temperature. Thus I
cubic metre of air is saturated with moisture at the various tempe-
ratures stated, when it has taken up the following weights of
water :—

4.871 grammes.

10° . 9.362

At 20° . 17.157 grammes.
„ 30° . 30.095 „

When air, saturated with moisture at say 20°, is cooled to 10°,
the excess of water beyond 9.362 (the maximum for 10°) is deposited
either as mist or rain. The temperature at which air thus begins
to deposit moisture is called the dew-point. The deposition of
moisture from the air caused by the lowering of the temperature
is a matter of everyday observation. A glass vessel containing
iced water becomes bedewed with moisture upon the outside as
the air in its immediate vicinity is cooled. When a season of
severe frost is suddenly followed by a warm wind, highly charged
with aqueous vapour, it is not unusual to see condensed moisture
collecting upon and streaming down the cold surface of walls.
For the same reason, after the sun has set, and the heat from the
ground has radiated, leaving the ground colder than the atmos-
phere, the temperature of the air is lowered, and it begins to
deposit its aqueous vapour in the form of dew.

The amount of aqueous vapour in the air, or the humidity of the
air, is estimated by meteorologists by means of an instrument
called the wet and dry bulb thermometer.

Carbon Dioxide. — The proportion of this gas present in the
air is also liable to considerable variation, although not through
such a wide range as the aqueous vapour. The processes of
respiration, combustion, and putrefaction are attended by the
evolution of carbon dioxide, hence the amount of this gas present
in closed inhabited places is greater than that in the open air ; in
badly ventilated and crowded rooms the proportion sometimes
rises to three parts in 1000 vols. Frankland has found that at high
elevations the amount of carbon dioxide in the air is often, although
not invariably, considerably above the normal.

At Chamounix (3000 feet) the amount of carbon dioxide was 0.63 per 1000 vols.
,, Grands Mulcts (11,000 feet) ,, „ ,, i.n ,, ,r

., Mont Blanc (15,732 feet) ., ,, „ 0.61 „ „

258 Inorganic Chemistry

This fact is probably due to the absence, in these hign regions,
of the vegetation which is one of the chief natural causes operating
to remove atmospheric carbonic dioxide (see Oxygen, page 188).

The amount of carbon dioxide is slightly higher during the
night, and often rises considerably during foggy weather. Thorpe
has shown that near the surface of the sea the amount of carbon
dioxide in the air is slightly less, being on an average 0.300 volume
per looo.

Ammonia in the atmosphere is derived from the decomposition
of nitrogenous organic matter. Although present in relatively
very small quantities, it varies in amount very considerably. From
the experiments of Angus Smith, 1000 grammes of air from various
sources were found to contain the following amounts of ammonia : —

London .. _ » . 0.05 gramme.
Glasgow . . . 0.06 „
Manchester. . "... , o.io „

The proportion of ammonia appears to be higher during the
night than in the daytime, and immediately after heavy rain the
amount is perceptibly diminished.

Rain-water always contains ammonia, although the amount
varies greatly with changing atmospheric and climatic conditions.
Lawes and Gilbert, Angus Smith, and others, have made a large
number of estimations of the amount of ammonia in rain-water at
various places and seasons, and under many different conditions.

Nitric Acid is produced in the atmosphere by the direct
union of oxygen and nitrogen whenever a lightning flash passes
through the air (see Nitric Acid). Rain which falls during or
immediately after a thunderstorm is found to contain nitrates and
nitrites.

These two nitrogenous compounds, ammonia and nitric acid,
although present only in such small proportion in the atmosphere,
fulfil a most important function in the economy of nature. From
the experiments of Lawes and Gilbert, and others, it has been shown
that most plants are unable to draw upon the free nitrogen of the
atmosphere for the supply of that element which they require for
the development of their structure and fruit.* Although they are
surrounded by, and bathed in nitrogen, they cannot assimilate it.
Plants that are growing in unmanured soil, therefore, derive their

* Leguminous plants, such as clovers, vetches, beans, peas, which develop
root-nodules or tubercles, are exceptions.

The Atmosphere 259

nitrogen from the ammonia and nitric acid which are present in
the air, and which are washed into the ground by the rain. It has
been found that a plant grown under such experimental conditions,
as to exclude the possibility of its obtaining supplies of these nitro-
genous compounds, will yield upon analysis exactly the same
amount of nitrogen as was originally contained in the seed from
which it grew.

Ozone. — The causes which operate in the formation of this sub-
stance in the air are at present imperfectly known ; it is supposed
that >ts occurrence is related to the development of electricity in
the atmosphere. On account of the powerful oxidising character
of ozone, its presence can never be detected in the air where much
organic matter of an oxidisable nature is present, as is the case in
the air of such places as malarial swamps, dwelling-houses, and
large towns.

The amount of ozone in pure country air has been found to vary
with the time of year, reaching a maximum in the spring-time, and
gradually falling towards winter. Thorpe has found that in sea
air the amount of ozone is practically constant during all seasons.

The usual method which is available for the detection and
estimation of ozone in the air is extremely crude. It consists in
exposing ozone test papers (see Ozone) to the air for a certain time,
and comparing the colour that is produced with a standard scale
of tints ; moreover, other substances than ozone, which may be
present in the atmosphere, will also liberate iodine from potassium
iodide, and these are therefore measured as ozone. Besides the
higher oxides of nitrogen, which, as we have seen, are formed in the
atmosphere, and which liberate iodine from potassium iodide, it has
been shown that peroxide of hydrogen is also present. The state of
our knowledge at present, therefore, respecting the exact amount of
atmospheric ozone and its variation is far from satisfactory ; it is,
indeed, quite possible that many of the effects which have been attri-
buted to ozone are in reality due to peroxide of hydrogen. Thus it
has been shown by Schonbein that this compound is formed during
the evaporation of water, and this statement probably derives con-
firmation from the fact that its presence may be detected in rain-
water. The salubrity of the air of the sea-shore, where large areas
of wet sand and stones offer the most perfect conditions for the
rapid evaporation of water, and consequently, for the formation of
peroxide of hydrogen, may therefore be attributable as much to the
presence of this substance as to the proverbial ozone.

2<5o Inorganic Chemistry

Hydrogen. — The presence of this gas in sensible quantities as
a constituent of normal air appears to have escaped notice until
quite recently. When a quantity of liquefied air is subjected to
fractional distillation, the first and most volatile portion which
collects is found to be very rich in hydrogen,*

The elaborate researches of Gautier t show also that this hydro-
gen is not only present in the air of towns, but that it is a con-
stituent of country air, air from high mountain regions, and of sea
air. The chief sources of this hydrogen are indicated on page
171.

The various gases of which the air is composed are not com-
bined, but are merely mingled together. The remarkable con-
stancy of its composition, as regards the oxygen and nitrogen, led
chemists at one time to suppose that these gases were in chemical
union with each other in the atmosphere ; but a number of facts
which have since been learnt respecting these gases prove with-
out doubt that this is not the case, and that the air is simply a
mechanical mixture. This evidence may be briefly summed up
as follows : —

(i.) When oxygen and nitrogen are mixed together in the
proportion in which they occur in air, the resulting mixture
behaves in all respects like ordinary air, and the mixing of the
gases is not attended by any volumetric or thermal disturbance,
such as would be expected to accompany the chemical union of
two elements.

(2.) The degree to which air is capable of refracting light is
found to be the mean of the refractive power of oxygen and
nitrogen. Were these gases chemically combined, the compound
should behave in this respect as other compound gases, where it
is found that the refractive index is always either greater or less
than the mean of that of the constituents.

(3.) According to a fundamental law of chemical science, the
composition of a chemical compound is constant. Such a thing as
variability in the composition of a compound is unknown. The
proportion of oxygen and nitrogen, as we have seen, does vary in
the air, although through only small limits, hence they cannot be
united to form a compound.

(4.) The proportion by weight in which oxygen and nitrogen are

* Dewar, Nature, December 20, 1900.

f Gautier, Annales de Chimie et de Physique, January 1901.

The Atmosphere 261

present in air bears no simple relation to the atomic weights of
these elements.

(5.) When air is dissolved in water, the oxygen and nitrogen
dissolve as from a simple mixture of these gases, in accordance to
the law of partial pressures (see page 147).

(6.) The oxygen and nitrogen can be partially separated, by
taking advantage of the different rates of diffusion of these two
gases (see Diffusion of Gases, page 83).

The various gases of the atmosphere are maintained in a state
of uniform admixture, in spite of their widely different densities,
by the operation of two causes : first, air currents, which effect the
rapid removal of large masses of air from place to place ; and,
second, their own molecular movements, which bring about the
phenomena of gaseous diffusion.

Suspended Impurities in the Atmosphere.— Besides the
gaseous constituents of the air, there is always present a certain
quantity of suspended matter, both liquid and solid. The exist-
ence of this suspended matter in the air can be rendered evident
from the fact that these minute particles are capable of reflecting
light ; if, therefore, a strong beam of light be passed through a
darkened room, the track of the beam is distinctly visible, on
account of its being reflected from innumerable particles floating
about in the air, many of them appearing quite large. Pasteur has
shown that this suspended matter can be removed by filtration
through cotton wool.* Tyndall also has shown that in undis-
turbed air the suspended matter settles in the course of a few
hours, leaving the air almost entirely free from this impurity.
For this purpose the floor of a large oblong glass box was
smeared over with glycerine. The box, after being hermetically
closed, was then allowed to stand for twenty-four hours, during
which time the suspended matter subsided and adhered to the
glycerine. When a beam of light is allowed to pass through air
that has been thus freed from suspended matter, there being
nothing present to reflect the light, the beam cannot be seen ;
its track will be evident in the air of the room as it enters and
leaves the box, but within the box it will be invisible (as repre-
sented in Fig. 55). To air in which a beam of light is in this way
invisible, Tyndall has applied the term " optically pure."

The suspended matters are partly mineral and partly organic.
Of the mineral matters, sodium chloride and certain sulphates
* See Experiments 334 to 341, " Chemical Lecture Experiments," new ed.

262 Inorganic Chemistry

are present in greatest quantity. These are thrown into the air
in the sea-spray, and as the small globules of water evaporate
they leave minute residual particles of saline matter, which,
being driven by the wind, remain floating in the atmosphere. It
is only very rarely, even at far inland places in Europe, that spec-
troscopic examination fails to detect the presence of sodium com-
pounds in the air. In the air of islands, such as England, it is
never absent. Sulphates are also produced by the oxidation and
combustion of sulphuretted compounds ; the amount of these, there-
fore, is greatly increased in the neighbourhood of towns.

The organic suspended matter of the air has of late years been
made the subject of extended research. Pasteur has shown that
amongst these organic substances are the germs and organisms
which produce fermentation, putrefaction, and disease. Putresciblo

FIG. 55.

substances, such as milk, urine, flesh, &c., if themselves carefully
freed from all such germs, may be preserved unchanged, for ap-
parently any length of time, in air that has been deprived of all
suspended matter. It is highly probable that the salubrity or
otherwise of different places is associated with the nature and
amount of the organic matter in the air, and it is certain that
these organisms play a most important part in relation to the life
and health of man. The feelings of lassitude and headache, which
result from the prolonged breathing of the air of rooms containing
many people, are brought about more by the poisonous effects of
the organic emanations evolved during respiration than by any
diminution in the supply of oxygen, or increase in the proportion
of carbon dioxide in the air. The well-known and unpleasant
smell that is perceived on first entering a crowded room is also
due to the same cause, and it has been shown that the moisture
which condenses from such an atmosphere upon a cold object, if

The Atmosphere 263

preserved for a short time, rapidly becomes putrescent, owing to
the decomposition of this organic matter.

The presence of suspended matter in the air appears to exert a
remarkable influence upon the formation and character of fogs.
Aitkin has shown that those conditions which result in the forma-
tion of a fog in ordinary air are incapable of producing that effect
in air that has been freed from suspended matter. It would appear
that the suspended particles act as innumerable points, or nuclei,
which facilitate the deposition of moisture, much in the same way
as the crystallisation of a salt, from its solution, is known to start
from any minute particles of foreign matter that may be floating in
the liquid.

The height to which the atmosphere extends has been variously
estimated. From observation of the flight of meteorites, it ap-
pears that even at a height of seventy to seventy-five miles the
air still has a sensible degree of density. The air being elastic,
and subject to the law of gravitation, its density, which is greatest
at the earth's surface, rapidly diminishes as the altitude increases ;
thus, at about three and a half miles the density is only one-half,
and at seven miles one-third, of that which obtains at the sea-level.
From a consideration of the physical properties of gases, there is
every reason to believe that in an extremely attenuated condition
the atmosphere extends far into space, and it has been calculated
that the pressure exerted by our atmosphere upon the surface of
the moon is equal to about I mm. of mercury.

The density of the atmosphere varies at different points of the
earth's surface, and at the same point at different times. The
pressure exerted by the atmosphere is measured by the height
of a column of mercury which it is capable of supporting, the
instrument employed for the purpose being called the barometer.
At the sea-level in the latitude of London, the average weight of
the atmosphere is equal to that of a column of mercury 760 mm.
at o°, and this is taken as the standard pressure of the atmos-
phere.

THE ARGON GROUP OF ATMOSPHERIC GASES

History. — More than a hundred years ago Cavendish observed
that when a mixture of nitrogen {phlogisticated air} and oxygen
(dephlogisticated air) was confined in a glass tube over mercury
along with a solution of caustic potash, and the gases exposed to

264 Inorganic Chemistry

the continued action of electric sparks, there was a small residue
of gas (amounting to about T^th of the volume of the nitrogen)
which was not absorbed, and he raised the question as to whether
the " phlogisticated air" of our atmosphere is entirely of one kind.*

This observation and speculation of Cavendish's remained buried
until 1894, when Lord Rayleigh and Professor Ramsay announced
to the world the discovery of a new gaseous constituent of the
atmosphere.

In making exact determinations of the densities of gases, Lord
Rayleigh found that nitrogen obtained from atmospheric sources
always gave a slightly higher number than that obtained for
nitrogen which was prepared from chemical compounds. On
careful investigation, in conjunction with Professor Ramsay, it was
found that this higher density of "atmospheric nitrogen" was due
to the presence in the air of a hitherto unknown gas, which they
succeeded in isolating, and to which they gave the name Argon

(1894).

In the following year, in searching for probable sources of
argon, Ramsay was led to examine the gas which was known to
be occluded in certain minerals, notably in the rare minerals
cl^veite and broggerite. This gas, which had hitherto been regarded
as nitrogen, was found to give a spectrum the most characteristic
line of which was a remarkably brilliant one in the yellow. The
position of this yellow line proved to be coincident with the
line D3 of the solar spectrum, which is the characteristic line of a
hitherto unknown solar element first observed by M. Janssen of
Paris in 1868, the spectrum of which was studied by Frankland
and Lockyer, who applied the name " helium " (" the sun ") to the
element. Subsequently Ramsay has shown that helium is present
in the atmosphere, although in much smaller quantities than argon.

In the year 1898 the discoverer of argon announced to the Royal
Society the discovery of two other gases which were associated
with argon, to which he gave the names neon ("the new one")
and krypton ("the hidden one"), and subsequently he discovered
still another and denser gas, which has been called xenon (" the
stranger"). The five gases, therefore, belonging to this group, in
the order of their densities, are : helium, 2 ; neon, 10; argon, 19.95 ;
krypton, 41.50; xenon, 65.35.

Not only are these five new gases elementary substances, but
they possess many properties in common. They all are extremely

* "Experiments on Air," Phil. Trans., 75, 372, 1785.

Argon 265

inert elements, having apparently no chemical activities whatever.
No compounds are known in which any one of them exists as a
chemical constituent, and they have resisted ail attempts to cause
them to enter into chemical combination with any other element.
Hence there is at present no chemistry of these strange substances.
This being the case, the only light which can be thrown upon the
complexity of the molecules of these elements is by the determina-
tion of the ratio of their specific heats at constant pressure and at
constant volume, deduced from determinations of the wave-length
of sound. This ratio is found to be 1.66, which is the same as that
obtaining in the case of mercury vapour, the only other monatomic
gas in which this ratio has been determined ; whereas with diatomic
gases, such as hydrogen, nitrogen, and oxygen, the ratio is 1.4.

ARGON.

Symbol, A. Density, 19.94. Atomic weight, 39.88.

Occurrence. — Argon is present in the atmosphere,* where it
exists to the extent of 0.937 per cent., or rather more than i per
cent, of the "atmospheric nitrogen" is argon. It is also found in
the occluded gases of certain specimens of meteoric iron, and in
minute quantities in almost ail natural waters, derived doubtless
by solution from the atmosphere. Argon has not been met with
in chemical combination with other elements, and no compounds
containing this element are known.

Modes of Preparation.— ( i.) Argon may be obtained from the
atmosphere by sparking a mixture of air and oxygen. The nitrogen
combines with the oxygen, and the oxidised product is absorbed
by potash. When no further contraction of volume is obtained,
the excess of oxygen is removed by alkaline pyrogallate, and the
residual gas is the argon. Unless a high-tension alternating
electric discharge is employed the process is extremely slow.

(2.) Argon may also be separated from the other atmospheric
gases by first withdrawing the oxygen by means of red-hot copper,
and after removing the carbon dioxide and aqueous vapour, passing
the remaining gas over strongly heated magnesium turnings. The
magnesium combines with the nitrogen (p. 232) and leaves the
argon. In order to effect the complete absorption of every trace
of nitrogen, the gas is passed backwards and forwards over the
heated magnesium for many hours.

Recently it has been found that the metal calcium is a more

266 Inorganic Chemistry

efficient agent for the absorption of nitrogen. If, therefore, the
" atmospheric nitrogen " be passed over a heated mixture of mag-
nesium filings and pure dry lime the magnesium and lime interact,
forming magnesia and calcium, which latter absorbs the nitrogen
very rapidly and at a lowe- temperature than that required by
metallic magnesium.

The purification of the argon thus obtained, and its complete
separation from the other gases with which it is associated, was a
problem which was only solved as a result of the later achievements
by Dewar of obtaining liquid hydrogen in quantity. By means of
the intense cold obtainable by liquid hydrogen, comparatively
large quantities of argon were liquefied, and the liquid so obtained
was then submitted to a process of fractional distillation. The
liquid gases having the lowest boiling-points, namely helium and
neon, are the first to evaporate, and by careful adjustment of the
temperature of the refrigerating bath the denser gases, krypton
and xenon, may be maintained even in the solidified state, while the
whole of trie argon in a state of practical purity can be distilled away.

Properties.— Argon is remarkable for its extraordinary inert-
ness, a property which is indicated by its name, argon signifying
" inactive." As already mentioned, it has hitherto resisted all
attempts to cause it to unite chemically with any other element.
The density of the gas is 19.94, and therefore its molecular weight
is 39.88 ; and since argon is a monatomic element, its atomic
weight and its molecular weight are the same.

Argon is about two and a half times as soluble in water as nitro-
gen, 100 volumes of water at 15° dissolving 4.1 volumes of argon.
Owing to this superior solubility, the gases which are expelled
from rain-water by boiling are slightly richer in argon than the
original air before solution. The critical temperature of argon is
— 117.4°, at which temperature the gas is liquefied by a pressure of
about fifty- three atmospheres (or 40.20 metres of mercury). Liquid
argon has a specific gravity of 1.212, and boils at - 186.9°. The
boiling-point of argon, therefore, lies between those of the two chief
constituents of the atmosphere, namely, oxygen and nitrogen, while
its critical temperature is slightly above that of oxygen, as may be
seen by the following comparison : —

Boiling-point. Critical Temperature.

Oxygen . . . . -182.5° -118.8°

Argon , -, . • . .: -186.9° -U7-4"

Nitrogen • > ' - I9S-S° -U9*

Helium 267

A very slight reduction of temperature below its boiling-point is
sufficient to freeze argon to a white solid, the melting-point of
which is — 187.9° 5 tnat is> ^ess tnan two degrees below the boiling-
point. The spectrum of argon is very complex. The most cha-
racteristic lines are two in the red (less refrangible than the red
lines of either hydrogen or lithium), a bright yellow line (more
refrangible than the sodium line), a group of bright green lines,
and another group of strong lines in the violet. The general cha-
racter of the spectrum depends upon the nature of the electric
discharge employed. With an intermittent discharge the lines in
the red and pale green are the most prominent, while with a
Leyden-jar discharge the red and light green lines almost entirely
disappear, giving place to lines in the dark green, blue, and violet.

HELIUM.

Symbol, He. Density, 1.995. Atomic weight, 3.99.

Occurrence. — The existence of this element in the universe
may be said to have been first discovered by Janssen, who during
a solar eclipse in 1 868 observed a certain line in the yellow of the
spectrum of the sun's chromosphere which was not coincident with
that of any known terrestrial element. This unknown element
was afterwards named helium by Frankland and Lockyer.

Terrestrial helium was discovered by Ramsay in 1895, in the gas
which is contained in certain rare minerals, and which is evolved
from them either when they are heated or when they are treated
with dilute sulphuric acid. Chief among these minerals are
cleveite, broggerite, and uraninite, all of them minerals containing
the metal uranium.*

Helium is present in minute quantities in the atmosphere,
namely, to the extent of about I or 2 volumes in 1,000,000 volumes
of air, as estimated by its discoverer.

Helium also occurs in certain natural waters, notably in the
water from the Bath springs, which has been found to contain
argon mixed with about 8 per cent, of its volume of helium.

Method Of Preparation.— Helium is isolated from the atmos-
phere by a method consisting firstly of what may be described as
fractional liquefaction, followed by fractional evaporation or distilla-
tion. When air is liquefied by the so-called self-cooling or recupe-

* The amount of helium contained in i gramme of cleveite is about 3.2 c. a
(Ramsay), only about one half of which is given off by heat alone.

268 Inorganic Chemistry

rative method in any of the modern air-liquefiers based upon
Linde's original apparatus (see page 77), those portions of the air
which escape liquefaction and pass out of the apparatus will
obviously contain most of the constituents having the lowest boil-
ing-points. Therefore, by collecting the gas which escapes from
the air-liquefier under these circumstances, and compressing it into
a vessel cooled by liquid air, a liquid is obtained which contains
most of the more volatile constituents (namely, the helium and
neon), with, of course, argon and some nitrogen. Thus, by this
process of fractional liquefaction liquid air is divided into two
fractions, one containing practically all the denser and least vola-
tile constituents, namely, the krypton and xenon, the other con-
taining the helium and neoi .

The separation of the gases contained in the more volatile
fraction is accomplished by fractional distillation or evaporation.
At the low temperature obtainable by means of liquid hydrogen
both argon and neon exert no vapour-pressure, being reduced to
the state of non-volatile solids, and the helium in a state of purity
can be pumped away from the mixture.

Properties.— Next to hydrogen, helium is the lightest known
gas, its density being 2. Like all the other gases of this group
its molecules are monatomic, its atomic and molecular weight
therefore is 4. Helium is much less soluble in water than
argon. The solubility of this gas in water forms an exception to
the usual behaviour of gases, for it has been found that while its
solubility diminishes with rise of temperature up to about 15°,
between 25° and 50° the solubility slightly increases, as is shown
by the following table.*

100 volumes of water at 760 mm. dissolve —

o° 10° 23* 30° 40* 50*

0.01500 0.01442 0.01386 0.01382 0.01387 0.01404 vols.

When induction sparks are passed through rarefied helium, the gas
emits a brilliant yellow light with a tinge of apricot colour. When
viewed through the spectroscope the most prominent and charac-
teristic line is the intense yellow line D3.x which is accompanied by
one bright red line, two in the green, and two in the blue. On
reducing the pressure in the tube, the yellow light due to line D3
gradually changes to a green, owing to the light from one of the
green lines becoming greatly intensified.

* Estreicher, " Z. Phys. Ch.," 31, 176.

Neon 269

Helium is the last of the known gases to be liquefied, having
resisted all attempts upon it until July 1908, when its liquefac-
tion was successfully accomplished by Prof. Onnes. 200 litres of
the purest helium after circulating for some hours through the
" liquefier " cooled by liquid hydrogen boiling under reduced pres-
sure, yielded as much as 60 c.c. of liquid helium. Liquid helium
boils about - 268.7° or 4.3° absolute, its critical temperature being
about 5° absolute.

Like all the other gases of this group helium is chemically inactive.

NEON.

Symbol, Ne. Density=io.i. Atomic weight = 20. 2.

History. — From analogy with other natural families of elements
and the numerical relations of the atomic weights of the different
members, the discoverer of argon and helium was led to believe
that another element should exist having an atomic weight between
those of these two elements, and about sixteen units higher than
that of helium. The long and careful search for this unknown
element was at last rewarded by the discovery of neon, whose
atomic weight was found to be 20, or exactly sixteen units above
that of helium.

Although only present in minute quantities in the atmosphere,
the discoverer estimates the amount as about ten times that of
helium ; that is to say, r or 2 parts of neon in 100,000 parts of air.

Neon is more readily liquefied than helium, but no exact deter-
mination of its boiling-point has yet been made.

The colour emitted by this gas when induction sparks are passed
through it is a brilliant orange-pink. Its spectrum is characterised
by a bright yellow line, D6, and a great cluster of lines in the more
orange part of the red end. It also exhibits other fainter lines
throughout the spectrum.

KRYPTON AND XENON.

Krypton, symbol Kr. Density = 4t.46. Atomic v eight = 82. 92.
Xenon ,, X. ,, =65.1. ,, ,, =130.2.

These two denser gases are obtained by the fractional distillation
of the heavier portion of liquid air obtained in the air-liquefier
(see Helium). Large quantities of this liquid, amounting to 30
litres, were carefully evaporated, and the residual portion, after
being entirely freed from nitrogen and oxygen, was again liquefied
by means of liquid air. The constituents of this liquid were then
separated by fractionation. As soon as most of the argon was

270 Inorganic Chemistry

removed, the residue consisting of krypton and xenon was easily
solidified. Under these circumstances it was found that the krypton
could be withdrawn by pumping, for at the temperature of liquid
air solid krypton is appreciably volatile, while the solidified xenon
is practically non-volatile. In the estimation of the discoverer, air
contains only about one part of krypton in one million parts ;
while of xenon the proportion is about one part in twenty millions.
The boiling and melting points of krypton are — 1 5 1.7° and — 169° re-
spectively, while those of xenon were found to be — 109.1° and — 140°.

The light emitted by krypton, under the influence of the induc-
tion spark, is a yellowish-green colour, while that given by xenon
under the same circumstances is more of a sky-blue.

The most prominent lines in the spectrum of krypton are two
very near together in the red, one bright yellow line and one strong
green line ; besides which there are a few in the blue and violet.
The brilliant green line (wave length, 5570.5) has attracted special
notice, as it is considered highly probable that this line may prove
to be coincident with the chief line in the aurora spectrum.

The spectrum of xenon, like that of argon, is markedly different
as the electric discharge is modified. With the intermittent dis-
charge the prominent lines are four in the red end, and a number
of strong lines in the blue and greenish-blue. With the "jar"
discharge the red and blue lines become very reduced or alto-
gether disappear, and their place is taken by a number of lines in
the bright green.

The relative proportion in which these gases are believed to be
present in the atmosphere has been estimated provisionally by the
discoverer as follows : —

Helium, I part per 250,000 of air.

Neon, i to 2 parts „ 100,000 „
Argon, 0.937 part „ 100 „

Krypton, i „ „ 1,000,000 „

Xenon, i „ „ 20,000,000 „

Or, expressed in parts per 1000, to compare more readily with
the figures in the table on page 256 : —

Argon. . . • <> 9-37

Neon ..... o.oi

Helium . . . . 0.004

Krypton . . ; . o.ooi

Xenon. , . > • 0.00005

The Inert Gases

271

The following table gives the latest physical constants for the
members of this strange group of new elements.

Den-
sity.

Atomic
Weight.

Boiling-
Point.

Melting-
Point.

Critical
Temp.

Critical
Pressure.

Helium

1-995

3-99

-268.7°

-268°

Neon .

10. 1

20. 2

...

Argon .

19.94

39.88

-186.9°

-189.6°

-117.4°

52.9 Ats.

Krypton

41.46

82.92

-ISI-70

-169°

- 62.5°

54-2 ,,

Xenon .

65-1

130.2

— 109. 1°

-140°

+ r4.75°

57-2 „

The members of this group, while exhibiting many close re-
semblances, such as the monatomic nature of their molecules,
their remarkable inertness, &c., show also that gradation of pro-
perties which is met with in other natural groups of elements.
This appears by the results tabulated above, as well as by such
other properties as the refractive indices, atomic volumes, &c.

CHAPTER VII
COMPOUNDS OF NITROGEN AND HYDROGEN

FOUR compounds of nitrogen with hydrogen have been pre-
pared, namely : —

Ammonia . . .' . • NH3.

Hydrazine . . . . . N2H4 or (NH2)2.

Hydrazoic acid ... . . N3H or HN3.

Ammonium hydrazoate , ., . N4H4 or NH4N3.

AMMONIA.

Formula, NH3. Molecular weight= 17.04. Density=8.S2. ..

History. — Ammonium salts, and also the aqueous solution of
ammonia, were known to the alchemists. It was termed by
Glauber, spiritus volatilis salts armoniaci, being obtained by the
action of an alkali upon sal-armoniacum. Subsequently the name
spirits of hartshorn was applied to the ammoniacal liquid obtained
by the destructive distillation of such refuse as hoofs and horns of
animals. The actual discovery of gaseous ammonia was made by
Priestley (1774), when he collected the gas, evolved by the action
of lime upon sal-ammoniac, by means of his mercurial pneumatic
trough. Priestley named the gas alkaline air.

Occurrence. — In combination as ammonium carbonate it is pre-
sent in small quantities in the air, derived by the decay of nitro-
genous animal and vegetable matter. As nitrate and nitrite it is
found in rain-water. It is evolved, along with boric acid, from the
fumaroles of Tuscany (see Boric Acid), and is found as ammonium
chloride antl sulphate in the vicinity of active volcanoes.

Modes Of Formation. — (i.) Ammonia can be synthetically pro-
duced by submitting a mixture of nitrogen and hydrogen to the
influence of the silent electric discharge (Donkin). The amount
of ammonia so obtained, however, is extremely small, and can best
be shown by passing the gases, as they issue from the "ozone
tube," through a cylinder containing a small quantity of Nessler's

272

Ammonia 273

solution.* In a short time the solution will begin to show a
yellowish-brown colour, indicating the presence of traces of
ammonia.

(2.) Ammonia may be prepared by gently heating any of the
ammonium salts, with either of the caustic alkalies, potash or
soda, or with slaked lime. The salt most commonly employed is
the chloride. When this is mixed with an excess of slaked lime,
and the mixture gently heated in a flask, ammonia is evolved, and
calcium chloride and water are formed—

The gas may be dried by being passed through a cylinder con-
taining lumps of quicklime,t and may then be collected either by
upward displacement or in the mercurial trough. On account of
its extreme solubility it cannot be collected over water.

(3.) Ammonia is formed by the action of nascent hydrogen upon
salts of nitrous and nitric acid, thus —

NaNO3 + 4H2=NaHO + 2H2O + NH3.

This method is often made use of in the quantitative estimation
of nitrates in drinking water.

(4.) When nitrogenous organic matter is subjected to destruc-
tive distillation, that is, strongly heated out of contact with air,
ammonia is formed ; hence when coal, which usually contains about
2 per cent, of nitrogen, is distilled in the process of the manu-
facture of ordinary illuminating gas, one of the products of the
decomposition is ammonia. The "ammoniacal liquor" of the
gas works is the source of all ammonia salts at the present day.
The liquor is boiled with milk of lime, and the ammonia thus
expelled is absorbed by sulphuric acid. The ammonium sulphate
so obtained is purified by recrystallisation.

Properties. — Ammonia is a colourless gas, having a powerfully
pungent smell, and a strong caustic taste. It is lighter than air,
its density being 0.589 (air= i). Ammonia possesses the property
of alkalinity in a very high degree ; it turns red litmus blue, and
yellow turmeric brown. The gas is unable to support combustion,

* A solution of mercuric iodide in potassium iodide, rendered alkaline with
potassium hydroxide.

f The usual desiccating agents, namely, sulphuric acid, or phosphorus
pentoxide, are inadmissible in the case of ammonia, as this gas at once unites
with such compounds.

274

Inorganic Chemistry

and is irrespirable. Under ordinary conditions ammonia *s not
combustible, but if the air be heated or if the amount of oxygen be
increased the gas will then burn with a flame of a characteristic
yellow-ochre colour. This behaviour of ammonia as regards com-
bustibility is most conveniently
illustrated by means of the ap-
paratus shown in Fig. 56. A
stream of the gas obtained by
gently heating a quantity of the
strong aqueous solution in a
small flask is delivered through
a tube which is surrounded by
a wider glass tube. Through
the cork which carries this tube
a second tube passes, through
which a supply of oxygen can
be passed. On applying a
lighted taper to the jet of am-
monia as it issues from the tube
it will be noticed that the gas
burns in the heated air round
the flame of the taper, but is
unable to continue burning when
the taper is withdrawn. If now
a gentle stream of oxygen be
admitted into the annular space
between the two tubes the ammonia readily ignites, and continues
to burn with its characteristic flame. On cutting off the supply
of oxygen the flame of the burning ammonia languishes and
dies out.

Ammonia is extremely soluble in water ; i c.c. of water at o° C.,
and at the standard pressure, dissolves 1148 c.c. of ammonia,
measured at o° C. and 760 mm. The solubility rapidly decreases
as the temperature rises, as will be seen by the following table : —

FIG. c6.

i c.c. of Water at
760 mm. Dissolves

At o° .
„ 8°.
,. 16° .
„ 30° >
» 50° -

Grammes, NH3.
. 0.875 .

0.582 .
0.403 .
0.229 •

c.c. at o° C. and
760 mm.

. 1148

• 923

. 764

. 529

. 306

Ammonia

275

When a solution of ammonia is heated the gas is rapidly evolved,
and at the boiling temperature the whole of it is given up.

The great solubility of this gas in water may be shown by filling
a large bolt-head flask with ammonia by displacement, the flask
being closed by means of a cork through which a long tube passes,
as shown in Fig. 57. On removing the cork from the end of the
tube water slowly rises until it reaches the top, and as soon as the
first drops enter the globe the absorption
proceeds with great rapidity, the water being
forced up the tube in the form of a fountain,
which continues until the flask is filled.

Commercial liquor ammotiice is prepared
by passing ammonia gas into water ; the
strongest solution has a specific gravity of
0.882 at 15°, and contains 35 per cent, of
ammonia. During the process of solution
heat is liberated, and when the gas is again
expelled the same amount of heat is reab-
sorbed. If a rapid stream of air be driven
through a quantity of strong ammonia solu-
tion, contained in a glass flask, the ammonia
gas is quickly expelled ; and if the flask
be placed upon a wooden block, as seen in
Fig. 58, upon which a few drops of water
have been poured, it will be found that after
a few moments the flask will have become

firmly frozen to the block. By the rapid evaporation of ammonia
in this way it is possible to lower the temperature to —40° C.

Ammonia is an easily liquefiable gas ; thus at 15.5° it requires a
pressure of 6.9 atmospheres, and at o° only 4.2 atmospheres, in order
to liquefy it. The gas was first liquefied by Faraday (1823) by heat-
ing in one limb of a closed and bent glass tube (see Fig. 2) a quantity
of a compound of ammonia with silver chloride, the other limb of
the tube being immersed in a freezing-mixture. The experiment
may be made in a tube constructed as seen in Fig. 59- The wide
limb is nearly filled with dry precipitated silver chloride which has
been saturated with ammonia gas. This compound melts at
about 38°, and at a somewhat higher temperature it gives up
its ammonia. If the narrow limb of the tube be immersed in a
freezing-mixture while the compound is being heated, the com-
bined influence of the cold and the pressure exerted by the

FIG. 57.

276 Inorganic Chemistry

evolved ammonia will cause the gas to liquefy and collect in the
cold portion of the tube. On removing the .tube from the freezing-
mixture and allowing the other end to cool, the liquid ammonia
will boil off and be reabsorbed by the silver chloride, reforming
the original compound.

Liquid ammonia is easily obtained in larger quantity by passing
the gas through a glass tube immersed in a bath of solid carbonic
acid and ether. Liquid ammonia is a colourless, mobile, and
highly refracting liquid, boiling at —33.7°, and having a specific
gravity at o° of 0.6234. When cooled below —75° it solidifies to a
mass of white crystals.

Liquid ammonia dissolves the metals sodium and potassium, the
solution in each case being of an intense blue colour. On the
evaporation of the liquid the metal is deposited unchanged.

FIG. 58. FIG. 59.

During the evaporation of liquid ammonia, boiling as it does at
so low a temperature as — 33.7°, a rapid absorption of heat takes
place, and as this substance is so easily obtained it was one of the
earliest liquids employed for the artificial production of ice. Various
ice-making machines have been invented by M. Carre, in which the
reduction of temperature required is obtained by the evaporation
of liquid ammonia.

Ammonia is decomposed into its elements at a temperature
below a red heat. In this decomposition two volumes of ammonia
give one volume of nitrogen and three volumes of hydrogen. The
gaseous products, therefore, obtained by passing ammonia through
a red-hot tube are inflammable. In the same way, when electric
sparks are passed through ammonia, the gas is resolved into its
constituents. By performing this experiment upon a measured
volume of ammonia confined in a eudiometer over mercury, it will
be found that after the passage of the sparks for a short time and

Ammonia

277

the readjustment of the levels of mercury, the original volume of
the gas has been doubled.

The fact that the hydrogen and nitrogen are present in ammonia
in the proportion of three volumes of hydrogen to one of nitrogen
can be shown by taking advantage of the fact that ammonia is
decomposed by chlorine, the latter combining with the hydrogen
to form hydrochloric acid and the nitrogen being set free. This is
effected by means of the apparatus shown in Fig. 60. The long
glass tube, divided into three equal divi-
sions, is rilled with chlorine and closed by
a cork carrying a small dropping funnel. A
few cubic centimetres of strong aqueous
ammonia are poured into the funnel and
allowed to enter the tube drop by drop.
As the first two or three drops fall into
the chlorine it will be seen that the com-
bination is attended with a feeble flash of
light, and fumes of ammonium chloride are
formed. When the reaction is complete
the whole of the chlorine will have com-
bined with hydrogen derived from the
ammonia to form hydrochloric acid, and
this in its turn will combine with the excess
of ammonia added, forming ammonium
chloride. This substance dissolves in the
water. A small quantity of dilute sulphuric
acid is next introduced by means of the
dropping funnel in order to absorb the
remaining excess of ammonia. The at-
mospheric pressure is then once more re-
stored by attaching to the funnel a bent
tube, dipping into a beaker of water, as
shown in the figure, and when the water is allowed to enter
it will be found to flow into the tube until it reaches the second
graduation. The gas which is left and which occupies one of the
divisions of the tube is found on examination to be nitrogen.
This one measure of nitrogen, therefore, has been eliminated from
that amount of ammonia which has been decomposed by the
chlorine with which the tube was originally filled. Now chlorine
combines with its own volume of hydrogen, therefore the volume
of hydrogen which was in combination with the one measure of

FIG. 60.

278 Inorganic Chemistry

nitrogen is equal to the volume of chlorine contained in the tube,
that is to say, it was three measures. We have, therefore, one
measure of nitrogen and three measures of hydrogen, or, in other
words, ammonia is a combination of nitrogen and hydrogen in the
proportion of one volume of nitrogen to three volumes of hydrogen.

In contact with many metals at a moderately high temperature
ammonia is decomposed into its elements, and a compound of
the metal with nitrogen is formed. In this way, at temperatures
ranging between about 400° and 800° a number of metallic nitrides
have been obtained.* These compounds are produced by passing
a rapid stream of ammonia gas through heated porcelain tubes
containing the metal in the form of either wire, foil, or fine powder.
When heated in an atmosphere of hydrogen, these nitrides are
decomposed into nitrogen and the respective metal, hence they
can only be produced in the presence of a large excess of am-
monia gas.

Ammonia combines directly with acids forming salts, known
as ammonium salts, in which the nitrogen functions as a pentad
element ; thus with hydrochloric and sulphuric acids it forms respec-
tively ammonium chloride and ammonium sulphate —

NH3 + HC1 = (NH4)C1.
2NH3+H2SO4=(NH4)2SO4.

(The ammonium salts will be described with the compounds of
the alkali metals.)

Hydrazine (diamidogen), NH2- NH2 or N2H4.— This compound was first
prepared by Curtius (1887). It is obtained from a salt of an organic acid

N\
known as diazo-acetic acid, || CH'COOH. When the ethereal salt of this

acid is acted upon by potassium hydroxide, the potassium salt of another
acid is formed, namely triazo-acetic acid. This we may regard as merely a
polymer of the first acid, and represent its formula (N2 : CH-COOH)3.
When this compound is digested with dilute sulphuric acid it is converted
into hydrazine sulphate and oxalic acid. Thus, employing the simple formula
for the add-

Hydrazine may also be prepared from purely inorganic sources. When
hydrogen potassium sulphite is acted upon by potassium nitrite, a compound
known as potassium dinitroso-sulphonate is produced, O : N'N-OK-KSO3.

* Beilby and Henderson, Jour. Chem. Soc. , November 1901.

Hydrazine 279

The mechanism of the reaction will be made clearer it' the formula for the
nitrite be written O : N'OK. Thus —

2O:N<)K + 2HKSO3=O:N-N-OK-KSO3+K2SO4.

By the action of nascent hydrogen (from sodium amalgam) at the temperature
of ice, this compound is coaverted into the potassium salt of hydrazine
sulphonate —

0:N-N-OK-KS03+6H = H2N-NH-KS03+KHO + H20.

And this compound on distillation with potassium hydroxide yields hydrazine—

H2N-NH-KSO3+HKO = H2N-NH2+K2SO4.

The base itself may also be obtained by heating together in a sealed tube, to
a temperature of 170°, hydrazine hydrate, N2H4, H2O, and barium monoxide —

BaO + N2H4H20 = Ba(HO)2 + N2H4.

More recently it has been obtained (Raschig. Ber. 1910) by distilling hydrazine
hydrate with sodium hydroxide.

Anhydrous hydrazine at the ordinary temperature is a colourless strongly
fuming liquid. It combines with water with great readiness, and boils
at 113°.

Hydrazine Hydrate, N2H4H2O.— The compound formed by the combina-
tion of hydrazine with water is obtained by distilling hydrazine sulphate,
N2H4H2SO4, with an aqueous solution of potassium hydroxide (caustic potash)
in a vessel of silver. It is a colourless, fuming, powerfully corrosive liquid,
which boils at 118.5°. It attacks glass, cork, and indiarubber, and can only
be prepared in vessels of silver or platinum which are screwed together at their
junctions. With the halogen acids it forms two series of salts, in which either
one or two molecules of the halogen acid are present : thus with hydrochloric
acid we have —

Hydrazine monohydrochloride . . . N2H4HC1.
Hydrazine dihydrochloride .... N2H42HC1.

Hydrazine and its salts act as powerful reducing agents, and give the charac-
teristic red precipitate of cuprous oxide when added to Fdhling's solution.
This reaction serves to immediately distinguish these compounds from
ammonium salts.

/N
Hydrazoic Acid or Azoimide, HN3 or HN II .—Discovered by Curtius

^1890), The sodium salt is prepared by boiling benzoylazo-imide with sodium
hydroxide, when sodium benzoate and sodium hydrazoate are formed, thus —

/N /N

C6H6CO-N || +2NaHO=C6H6CO'ONa+Na-N || +H8O.
\N \N

280 Inorganic Chemistry

It is also produced when sodamide (obtained by heating sodium in dry
ammonia gas) is heated to 200° in a stream of nitrous oxide * —

2NH2Na+N2O=NaN3+NaHO + NH3.

The sodium hydrazoate so obtained is then gently warmed with dilute sulphuric
acid, when sodium sulphate and hydrazoic acid are formed, thus —

2NaN3 + H2SO4=Na2SO4+2HN3.

Properties. — This compound is a colourless volatile liquid, boiling at 37°.
The vapour possesses a most unpleasant and powerfully penetrating odour.
If inhaled, even when largely diluted with air, it exerts an irritating action
upon the mucous membrane. As its name denotes, it is an acid substance,
and in many of its properties it strongly resembles the halogen acids. The
compound is extremely soluble in water, and forms a strongly acid liquid
which smells of the vapour. This solution when boiled, finally assumes a
definite strength, and yields on distillation an aqueous acid of constant com-
position, in this respect resembling aqueous hydrochloric acid, q.v.

In its constitution this acid may be compared with hydrocyanic acid, and
with the halogen acids —

H(N3) ; H(CN) ; H(C1) ; H(Br),

in which the radical cyanogen (CN), or the halogen elements, Cl and Br, are
replaced by the group consisting of three nitrogen atoms.

When a solution of hydrazoic acid is added to a solution of silver nitrate, a
white precipitate of silver hydrazoate is formed, strongly resembling silver
cyanide and silver chloride in appearance. This silver salt, however, is not
acted upon by light in the way the chloride is, and it differs also in being
extremely explosive. A minute quantity of the compound, when touched with
a hot wire, detonates violently.

This instability and tendency to explode is characteristic of the acid and
most of its salts. The sodium salt, however, may be heated to about 250°
before it decomposes.

When gaseous hydrazoic acid is mixed with gaseous ammonia, dense white
fumes are formed, consisting of ammonium hydrazoate. These two hydrides
of nitrogen, apparently so similar, but in reality so widely different, unite to
form the ammonium salt, just as gaseous hydrochloric acid and ammonia
combine to form ammonium chloride, thus —

NH3+H(N3) = NH4(N3).
NH8 + HC1=NH4C1.

The alkaline hydride of nitrogen, ammonia, combines with the acid hydride of
nitrogen, hydrazoic acid, and forms the salt ammonium hydrazoate NH4N3
or N4H4. The salts of this acid are sometimes called nitrides, thus sodium
nitride, NaNs.

* See Experiment 298, " Chemical Lecture Experiments," new ed.

Hydroxylamine 281

HYDROXYLAMINE.

Formula, NH2(OH).

Discovered by Lessen in 1865.

Modes of Formation.— ( i.) By the action of nascent hydrogen
upon nitric oxide, nitric acid, or certain nitrates —

2NO + 3H2 = 2NH2(OH).
HNO3 + 3H2 = 2H2O + NH2(OH).

The nascent hydrogen is evolved from tin and hydrochloric acid,
and a stream of nitric oxide passed through the mixture. The
hydrochloride of hydroxylamine is thus obtained. The dissolved
tin is then precipitated as sulphide by means of hydrogen sulphide,
the filtered liquid evaporated to dryness, and the hydroxylamine
hydrochloride dissolved out of the residue with absolute alcohol.
On evaporating this solution the salt is obtained in the form of
white crystals.

(2.) By the interaction of alkali nitrites and metasulphites, and
the subsequent prolonged boiling of the hydroxylamine disulphonate
so obtained with water —

2KN 02 + 3K2SO3,SO2 + H2O = 2K2SO3 + 2N(OH)(SO3K)2.

2N(OH)(SO3K)2 + 4H2O = (NH3OH)2SO4 + 2K2SO4 + H2S04.

The potassium sulphate is separated from the hydroxylamine
sulphate by crystallisation.

Hydroxylamine itself, in aqueous solution, may be obtained from
the sulphate by the addition of baryta water. In the anhydrous
state it is produced by the distillation of hydroxylamine phosphate
under reduced pressure, the distillate being solidified by immersion
in ice. If this is dissolved in absolute alcohol and the solution
cooled to about — 18° the pure compound separates out in white
scales or needles.

Properties.— Hydroxylamine melts at 33°, and under reduced
pressure may be boiled and distilled ; but although tolerably stable
at the ordinary temperature it decomposes with explosion when
heated above about 100° at atmospheric pressure. It is readily
soluble in water yielding a strongly alkaline solution, which pre-
cipitates silver and mercury from solutions of their salts, and
which reduces cupric salts with precipitation of red cuprous oxide.

Hydroxylamine is a base, and may be regarded as ammonia, in
which one of the hydrogen atoms has been replaced by the monad
group hydroxyl (OH). Like ammonia it unites with acids forming
salts, without the elimination of water.

= NH3OHC1 (or NH2OH,HC1).
(NH3OH)2SO4 (or 2NH2OH,H2SO4).

The salts of hydroxylamine all decompose on the application of

282 Inorganic Chemistry

heat, with a more or less sudden and violent evolution of gas ; thus
the nitrate breaks up with almost explosive violence into nitric
oxide and water —

NH2OH'HNO3

AMMON-SULPHONATES.

These compounds may be regarded as derived from ammonia, by the
gradual replacement of the hydrogen by the group SO3H or SO..OH.

Ammon-sulphonic acid . . /- NH2(SO3H).

Ammon-disulphonic acid . . NH(SO3H)2.

Ammon-trisulphonic acid . . N(SO3H)3.

Potassium ammon-trisulphonate is precipitated as a crystalline salt when
excess of a solution of potassium sulphite is added to a solution of potassium
nitrite —

Prolonged boiling with water converts it first into the ammon-disulphonate —

N(SO3K)3+H2O=NH(SO3K)2 + HK.SO4,
arid finally into ammon-sulphonate—

NH(SOsK)a+H2O=NH2(SO3K) + HKSO4.

Ammon-sulphonic acid is a stable crystalline body ; the other two acids are
only known in their salts.

When an ice-cold solution of sodium nitrite is added to hydrogen sodium
sulphite, a compound is obtained which may be regarded as derived from
ammon-trisulphonic acid by the replacement of one of the groups, SO3H, by
hydroxyl, OH—

NaNO2+2NaHSO3=N(OH)(SO3Na)2

On the addition of a saturated solution of potassium chloride in the cold, the
sodium salt is converted into the potassium salt, which slowly crystallises from
the solution, with two molecules of water, N(OH)(SO3K)2,2H2O.

This potassium hydroxylamine disulphonate is an unstable compound, and
on boiling with water the two SO3K groups are replaced by hydrogen, forming
first potassium hydroxylamine monosulphonate, NH(OH)SO3K, and finally
hydroxylamine, NH2OH.

COMPOUNDS OF NITROGEN WITH THE HALOGEN ELEMENTS.

Nitrogen Trichloride, NC13. — This compound was discovered by Dulong
(1811). Its true composition was proved by Gattermann (1888).

Mode of Formation.— Nitrogen trichloride is obtained by the action of
chlorine upon ammonium chloride —

NH4Cl + 3Cl2=NCl3-r4HCL

Nitrogen Iodide 283

When a solution of ammonium chloride is electrolysed, the chlorine, which is
evolved at the positive electrode, acts upon the ammonium chloride, forming the
trichloride of nitrogen.*

Properties. — Nitrogen trichloride is a thin oily liquid, of a pale-yellow
colour, and having a specific gravity of 1.65. It is very volatile, and has an
unpleasant pungent smell, the vapour being extremely irritating to the eyes. It
is the most dangerously explosive compound known, and when suddenly
heated, or brought into contact with grease, turpentine, or phosphorus, it at once
explodes. It also explodes on exposure to sunlight. At a temperature of 71°
it may be distilled, but the operation is one of the utmost danger. Nitrogen
trichloride is decomposed by ammonia, forming ammonium chloride and free
nitrogen ; hence in the preparation of nitrogen by the action of chlorine upon
ammonia, the presence of an excess of ammonia prevents the formation of this
dangerous compound.

Nitrogen Tribromide, NBr3. — When potassium bromide is added to nitro-
gen trichloride beneath water, a red, oily, hignly explosive substance is
obtained, believed to be the tribromide of nitrogen.

Nitrogen Iodide, N2H3I3. — When strong aqueous ammonia is added to
powdered iodine, a brown-coloured powder is formed which has violently
explosive properties. Also when alcoholic solutions of iodine and of ammonia
are mixed, a brown and highly explosive compound is produced.

Curtois, who first prepared the substance, believed it to have the composi-
tion NI3, and this view was held by Gay-Lussac and others. Gladstone and
others considered that the substance contained one atom of hydrogen, and that
the formula NHI3 expressed the composition. The investigations of Szuhay
(1893) also led him to believe that the compound obtained by the addition
of an excess of aqueous ammonia to a concentrated solution of iodine in
potassium iodide has the composition NHI2.

The subject has recently been reinvestigated by Chattaway (Proc. Chem. Soc. ,
1899), who for the first time appears to have obtained the compound in a state
of purity by the addition of ammonia to a solution of potassium hypoiodite.
Under these circumstances the substance separates out in the form of well-
defined crystals having a composition expressed by the formula N2H3I3, which
may be regarded either as NI3,NH3 or NHI2,NH2I. The equations represent-
ing the formation of the compound may be thus expressed —

KIO + NH4HO=NH4IO+KHO,

3NH4IO=N2H3I3+NH3+3H2O.

The reaction which takes place when the compound is obtained by the
action of iodine upon strong ammonia appears also to involve the first forma-
tion of the unstable ammonium hypoiodite, thus —

I2+2NH4HO=NH4IO + NH4I + H20,

which then breaks up as shown above.

Properties.— Nitrogen iodide is a copper-coloured glittering crystalline com-
pound, appearing red by transmitted light. In the amorphous state, as obtained
by the action of iodine upon strong ammonia, it presents the appearance of

* See "Chemical Lecture Experiments," new ed., No. 301.

284 Inorganic Chemistry

a dark chocolate-brown powder. When moist it may be handled without much
risk of explosion, although it has been known to explode even under water.
When dry the substance is extremely explosive ; the shock caused by the tread
of a fly upon it is more than sufficient to explode it ; even falling dust particles
will sometimes cause it to explode.

WThen nitrogen iodide is placed in dilute aqueous ammonia, and exposed to
bright light, it is decomposed, and bubbles of nitrogen are seen escaping from
the compound —

N2H3I3=N2+3HI,

the hydriodic acid being neutralised by the, ammonia present. At the same
time a small quantity of the compound is converted into ammonium hypoiodite,
which being unstable slowly passes into the iodate, thus—

N2H.iI3 + 3H20-»-NH3=3NH4IOl
3NH JO=NH4IO3 + 2NH4I.

CHAPTER VIII
CARBON

Symbol, C. Atomic weight = 12.00.

Occurrence.— This element is capable of assuming three allo-
tropic forms, and it occurs free in nature in each of these modifica-
tions, viz., diamond, graphite, and charcoal.

In combination with oxygen, carbon occurs in carbon dioxide, a
gas which is present in the air, being a constant product of com-
bustion and respiration. In combination with hydrogen it occurs
as marsh gas. Carbon is a constituent of all the natural car-
bonates, such as limestone, dolomite, &c, which form an important
fraction of the earth's crust, and it is also an essential constituent
of all organic substances.

DIAMOND.

Occurrence. — This substance has been known and prized from
the remotest antiquity. It is found in various parts of India,
mostly in river gravels and superficial deposits, in Brazil, South
Africa, Australia, and various parts of the United States. The
diamond has also recently been obtained from extra-terrestrial
sources. In a meteorite which fell in Russia on September 22,
1886, carbon was found, partly as amorphous and partly as ada-
mantine carbon.

The diamond form of carbon is found of various colours ; some-
times it is dark grey, or even black, stones of these colours being
known as carbonado and bort. The former of these is extremely
hard, and is of great value for use in rock-boring and drilling
instruments. Bort is used in the crushed condition by lapidaries
for grinding and polishing.

Occasionally the diamond is found coloured blue, or red, or
green by traces of foreign materials. Some of these coloured

a8S

286 Inorganic Chemistry

stones are of great value as gems: the well-known "Hope'1
diamond, a stone weighing 44^ carats, has a fine sapphire
colour:

The origin of the diamond is unknown, although many theories
have been put forward to explain its formation. Newton's famous
suggestion, that diamond was " an unctuous substance coagulated,"
was based upon its remarkably high refractive index. The cellular
structure which is sometimes to be seen in the ash that is left when
the diamond is burnt seems to indicate that it is of vegetable
origin.

Modes Of Formation.— Innumerable attempts have been made
to effect the crystallisation of carbon in the adamantine form ; but
while it is readily possible to convert this variety of carbon into its
allotropes graphite and charcoal, the transformation of these back
again to the diamond is a problem that is beset with the greatest
difficulties. Moissan has recently shown * that the carbon, which
is capable of being dissolved in molten iron, and which is usually
deposited in the graphitic form on cooling, can, under certain
conditions, be caused to take up the adamantine form.

By raising the temperature of the iron to about 3000° by means
of an electric furnace, and then suddenly cooling the molten mass
by plunging the crucible into water or molten lead, until the cooled
and solidified surface is at a dull red heat, an enormous pressure is
brought to bear upon the interior and still liquid portion. Under
these circumstances, a part of the carbon which is deposited by the
slowly cooling mass was found by Moissan to be in the adamantine
form. On dissolving the iron in hydrochloric acid, amongst the
carbonaceous residue were found fragments having a specific
gravity between 3.0 and 3.5, and sufficiently hard to scratch ruby.
Some of the fragments were the black or carbonado variety, while
others were transparent. On combustion in oxygen, Moissan
proved that these were really carbon in the diamond form.

Properties. — The diamond in its purest condition is a colourless
crystalline substance. Its crystalline forms belong to the cubic
system, and appear to some extent to be characteristic of the
locality in which the element occurs. It is extremely hard and
moderately brittle. When struck with a hammer the diamond not
only splits along its cleavage-planes, but also in other directions,
with a conch oidal fracture. It does 'not conduct electricity The

* Comptes Rendus de V Acadtmie des Sciences, vol. cxvi. p. 218.

Carbon

287

specific gravity of diamond varies slightly in different specimens,
the mean being about 3.5. Its refractive index is higher than that
of any other substance, and it is this property which gives its
peculiar beauty and brilliancy to the cut stone.

The value of diamond as a gem depends largely upon its colour-
lessness, except in the case of those rare instances where the
colour is quite definite and also pleasing, such as distinct red, blue,
or green.

When diamond is strongly heated it becomes black, and in-
creases in bulk, being converted into a substance having the
properties of coke. Lavoisier (1772) was the first to show that
the diamond was a combustible body, and that it yielded carbon
dioxide. Davy (1814) showed that carbon dioxide was the only
product of its combustion, and proved that
diamond was pure carbon. t^~]r^=d

The combustion of diamond in oxygen may '^ lp
readily be accomplished by means of the ap-
paratus shown in Fig. 61. A fragment of
diamond is supported upon a small gutter of
platinum foil, which bridges across two stout
copper wires, A. These wires pass through a
cork in a perforated glass plate, and are
lowered into a cylinder of oxygen. By the
passage of an electric current the little plati-
num boat can be strongly heated, when the FIG. 61.
diamond will become ignited, and continue to
burn brilliantly in the oxygen, with the formation of carbon dioxide.
The ash, which is always left after a diamond has been burnt,
varies from 0.2 to 0.05 per cent, of the stone. It is found usually
to contain ferric oxide and silica.

GRAPHITE.

Occurrence. — This second allotrope of carbon is much more
plentiful in nature than the first. It is found in large quantities
in Siberia, Ceylon, and various parts of India. In England the
chief source of graphite has been the mines at Borrowdale in
Cumberland ; this supply is now practically exhausted. Enor-
mous quantities of very pure graphite are now obtained from the
Eureka Black-Lead Mines in California. Graphite also occurs
in many specimens of meteoric iron.

288 Inorganic Chemistry

Mode of Formation. — Molten iron, especially when it contains
silicon, is capable of dissolving a considerable amount of carbon,
which, on cooling, is deposited in the form of black shining crystals
of graphite. Occasionally considerable quantities of graphite are
found deposited in this way in iron-smelting furnaces, to which the
name " kish >J has been applied.

Graphite is now manufactured by heating a mixture of 97 parts
of amorphous carbon (charcoal or coke) and 3 parts of iron in an
electric furnace. It was formerly believed that at the high tem-
perature of the electric arc amorphous carbon was converted
directly into the graphitic modification ; but it has recently been
shown (Acheson) that pure charcoal does not by itself undergo
this transformation ; that the change, in reality, takes place
through the intermediate formation of a metallic carbide. The
product obtained is practically free from iron, as the metal is
volatilised at the high temperature.

Properties. — Graphite is a soft, shiny, greyish-black substance,
which is smooth and soapy to the touch. It is usually found in
compact laminated masses, but sometimes crystallised in six-sided
plates. Its specific gravity varies in different specimens, aver-
aging about 2.5. Graphite is a good conductor of both heat and
electricity.

When strongly heated in oxygen, graphite takes fire and burns,
forming carbon dioxide, and leaving an ash consisting of silica,
alumina, and oxide of iron. Graphite has been found by Regnault
to contain, usually, traces of hydrogen. Graphite is employed for
the manufacture of ordinary lead pencils ; for, on account of its
softness, it leaves a black mark upon paper when drawn across it.
For the purposes of the pencil manufacture the natural graphite is
ground to powder and carefully washed free from gritty matter.
It is then mixed with the finest washed clay, and the pasty mass is
forced by hydraulic pressure through perforated plates. The name
" graphite," from the Greek to write^ is given to this substance on
account of its use for this purpose. It was formerly supposed
that this material contained lead, hence the names black-lead and
plumbago.

When powdered graphite is subjected to prolonged treatment with boiling
nitric acid and potassium chlorate it undergoes partial oxidation, and is con-
verted into a greyish crystalline substance which was termed by its discoverer
(JBrodie) graphitic acid. It contains carbon, hydrogen, and oxygen, and is

Carbon 289

believed to have a composition represented by the formula H4CnO5. When
heated, this compound undergoes a very curious transformation. If a frag-
ment about the size of a pea is heated in the bottom of a test-tube, feeble signs
of visible combustion are seen, and a light, porous black mass is produced
which fills and overflows the tube. Tnis porous mass appears to be pure
graphite. At the same time a little moisture condenses upon the tube.

Graphite is largely employed, on account of its refractoriness,
for the manufacture of the so-called plumbago crucibles, which
consist of fireclay mixed with finely-ground graphite.

Other uses to which graphite is put are for glazing or polishing
gunpowder, especially the larger grained varieties ; as a lubricant
for machinery, where oil is inadmissible on account of high tem-
perature ; for electrotyping processes, and also as a coating for
ironwork, to prevent rusting.

AMORPHOUS CARBON.

This non-crystalline form of carbon may be obtained by the
decomposition of a great variety of carbon compounds, by the
process known as destructive distillation. The carbon so obtained
differs very much as regards its purity, according to the particular
organic compound used for its preparation. The commonest forms
of amorphous carbon to be met with are lampblack or soot, gas
carbon, coke, charcoal, animal charcoal or bone-black. None
of these substances is pure carbon ; animal charcoal, for example,
usually containing only about 10 per cent, of carbon.

Lampblack. — This substance is manufactured by burning sub-
stances rich in carbon, and which burn with a smoky flame (as
turpentine, petroleum, or tar), with a limited supply of air. The
smoke is passed into chambers in which are suspended coarse
blankets, upon which the soot collects. The lampblack always
contains hydrogen in the form of hydrocarbons. If the soot be
heated to redness in a stream of chlorine, this hydrogen can be
removed, and pure amorphous carbon will be left.

Lampblack is used for printers' ink and for black paint.

Gas Carbon. —This form of carbon is obtained by the destruc-
tive distillation of coal in the manufacture of illuminating gas.
It remains in the retorts as an extremely hard deposit, lining the
roof and sides. It is a very pure carbon, coming second to puri-
fied lampblack. Its specific gravity is about 2.35. Gas carbon is
a good conductor of electricity, and is extensively used for the
manufacture of carbon rods for the arc light.

290 Inorganic Chemistry

Coke. — This substance differs from gas carbon, although it also
is obtained in the process of coal distilling. It contains all the
inorganic matter which constitutes the ash of the coal, and also
small quantities of hydrogen, nitrogen, and oxygen. The average
amount of carbon in coke is about 91 per cent.

Charcoal. — The purest form of charcoal is obtained by the
carbonisation of pure white sugar and the subsequent ignition
of the charcoal in a stream of chlorine gas. Charcoal so ob-
tained has a specific gravity of 1.57. Charcoal in a much less
pure condition is manufactured from wood. The methods
by which the carbonisation of wood is carried out are, broadly,
of two kinds : first, those in which access of air is permitted
to the burning material ; and, second, those in which air is ex-
cluded.

The first of these, and the most ancient, is generally carried on
in more primitive parts, where wood is plentiful. The wood is
piled into mounds or stacks, which are built with some care.
They are set on fire in the interior by means of a lighted bundle
of brushwood, which is introduced through a vertical opening or
chimney, left for this purpose in the centre of the mound during its
construction. The outside of the heap is covered with brushwood,
and finally with turf, in order to regulate the access of air to the
interior, and therefore to control the rate of combustion of the wood.
The object of the charcoal-burner is to carbonise the wood as
slowly as possible. In this process there is great liability to loss,
by the too rapid combustion of the wood ; and, in addition, it pos--
sesses the disadvantage that the secondary products, such as the
pyroligneous acid, tar, &c., are entirely lost.

Various modifications have been introduced into the method of
coaling -wood, as the process is termed, with a view to collect these
products.

In the second general process of carbonising wood, the material
is heated in ovens or retorts from the outside, no air being admitted
to the wood. The operation is very similar to that employed in
the destructive distillation of coal, in the manufacture of coal gas.
In these methods all the volatile and condensable products are
collected ; among these are water, pyroligneous acid, wood spirit,
acetone, and fatty oils. The non-condensable products consist
mainly of such gases as hydrogen, carbon monoxide, carbon di-
oxide, marsh gas, and acetylene.

Animal Charcoal. — Bone-black is obtained by the carbonisa-

Charcoal 291

tion of bones in iron retorts. This variety of charcoal is the least
pure of all the ordinary forms of amorphous carbon.

Bone contains only about 30 per cent, of organic matter, the
other 70 per cent, consisting chiefly of calcium phosphate, asso-
ciated with small quantities of magnesium phosphate and calcium
carbonate. It will be obvious, therefore^ that as the carbon is
derived from the organic matter, the amount of it in carbonised
bones must be small. The average composition of animal char-
coal is found to be —

Carbon , ,-.»> .,"•>.'•.• '*nr .«.':/? iv"lV '•'• Iao
Calcium phosphate ...;.' . . . 88.0
Other saline substances '_„, . .\ ,,^\. ,. .2.0

i oo.o

Although containing relatively so small an amount of carbon,
animal charcoal possesses many of the valuable properties of
charcoal in a highly marked degree, owing to the fact that it con-
tains its carbon disseminated throughout an extremely porous mass
of calcium phosphate.

Properties Of Charcoal.— Charcoal varies very considerably in
its properties, depending upon the particular wood from which it
is obtained, and the method by which it is prepared. Thus, char-
coal obtained at 300° is a soft, brownish-black, very friable material,
having an igniting point as low as 380°. On the other hand,
charcoal prepared at very high temperatures is black and com-
paratively dense, and requires to be strongly heated in order to
ignite it.

Under ordinary circumstances, charcoal burns in air without the
formation of a flame, or the production of smoke. At high tem-
peratures, however, the combustion of charcoal is seen to be
attended by a flame. This is probably accounted for by the fact,
that as the temperature at which the combustion of carbon takes
place is raised above 700°, the amount of carbon monoxide which
is formed increases, and the carbon dioxide decreases.*'

When charcoal is thrown -upon water it floats, on account of the
air which is enclosed within its pores. The specific gravity of
charcoal when thus filled with air varies from 0.106 (charcoal made

* Ernst, Chemisches Repertoriumt vol. xvii. p. 2.

292 Inorganic Chemistry

from the ash) to 0.203 (charcoal from the birch). If the air be
withdrawn from charcoal it sinks in water, the average specific
gravity of charcoal itself being 1.5.

Ordinary charcoal is a bad conductor of electricity, but its con-
ductivity is greatly increased by strongly heating the charcoal in
closed vessels.

Charcoal has the power of absorbing gases and vapours to a
remarkable extent: this power, which is exhibited to a dif-
ferent degree by the various kinds of charcoal, is due to the
porosity of the material, whereby it' exposes a very large sur-
face ; and it belongs to a class of phenomena known as surface
action.

If a fragment of charcoal, recently strongly heated to expel the
air from its pores, be passed up into a cylinder of ammonia gas,
standing in a trough of mercury, the ammonia will be gradually
absorbed by the charcoal, and the mercury will ascend in the
cylinder. Saussure found that recently heated beech-wood char-
coal was capable of absorbing ninety times its own volume of
ammonia gas ; while Hunter, by employing charcoal made from
cocoa-nut shell, found that 171.7 volumes of ammonia were absorbed
by i volume of charcoal The results of both of these experi-
ments show that those gases are absorbed in the largest quantities
which are the most readily liquefiable. The gas so held by the
charcoal is in a highly condensed condition upon the surface of
the porous mass. Probably in the case of easily liquefied gases,
such as ammonia, sulphur dioxide, and others, the gases are par-
tially liquefied upon the surface of the charcoal. In this condensed
state the gas is more chemically active than under ordinary condi-
tions, and charcoal is therefore able to induce many striking com-
binations to take place. Thus, if charcoal be allowed to absorb
chlorine, and dry hydrogen be then passed over it, the chlorine is
capable of combining with the hydrogen even in the dark, with the
formation of hydrochloric acid. This chemical activity of gases,
when absorbed by charcoal, is strikingly exemplified in the case
of sulphuretted hydrogen. If a quantity of powdered charcoal,
which has been saturated with sulphuretted hydrogen, be brought
into oxygen, the rapid combination of the two gases is attended
with the development of so much heat that the charcoal bursts
into active combustion. In the same way a mixture of air, with
10 or 15 per cent, of sulphuretted hydrogen, may be passed rapidly
through a tube, about a metre in length, filled with charcoal,

Coal 293

without a trace of sulphuretted hydrogen escaping at the end.*
Owing to this property charcoal is largely employed to absorb
noxious gases, the atmospheric oxygen which is condensed in the
pores of the charcoal oxidising these offensive and injurious com-
pounds ; thus sewer ventilators are often trapped with a layer of
charcoal, which effectually arrests all bad-smelling gases.

Charcoal also has the power of absorbing colouring matters from
solution : thus, if water which has been tinted with an organic
colouring matter be shaken up with powdered charcoal and filtered,
the solution will be found to be entirely decolourised. The variety
of charcoal which possesses this property in the highest degree is
animal charcoal, or bone-black, and this substance is largely em-
ployed in many manufacturing processes, such as sugar-refining,
in order to remove all colouring matter from the liquid.

Charcoal under ordinary conditions is unacted upon by the air,
but when the temperature is raised it enters into active combus-
tion, forming carbon dioxide. In an extremely divided condition,
however, carbon is capable of combining spontaneously with the
oxygen of the air, and with so much energy as to take fire.

Coal. — The carbonaceous minerals that are included under the
name coal are an impure form of carbon, containing compounds
of carbon with hydrogen and oxygen. Coal is the final result of a
series of decomposition changes which have been undergone by
vegetable matter of the remote past, the process having extended
over long geological periods. During this prolonged process a por-
tion of the carbon and hydrogen is eliminated as marsh gas, and
large quantities of this gas are found associated with, and occluded
in, coal.

Broadly speaking, the numerous varieties of coal maybe divided
into soft or bituminous, and hard or anthracitic.

The former are employed for the manufacture of coal gas and
for ordinary domestic purposes ; they burn with a smoky flame, and
evolve large quantities of gases and volatile vapours on combus-
tion or distillation. Anthracite coal is much harder, ignites with
more difficulty, and burns with the formation of very little flame or
smoke. It contains a higher percentage of carbon, and gives out
great heat on combustion, and is employed largely as a steam-
coal.

* " Chemical Lecture Experiments," 394-396, new ed. „

294

Inorganic Chemistry

The following table shows the average composition of coals from
various sources, and the general difference between coals of the
two main classes : —

Locality.

bfl

3
ff

. /^Northumberland

|§] Wales . .;•"•

81.41
83.78

5.83

4-79

7.90
4-15

2.05
0.98

0.74

1-43

2.07
4.91

^35

66.70
72.60

^Staffordshire

78.57

5-29

12.88

1.84

0-39

i.o3

11.29

57-21

5 . fS Wales

92 56

0 00

2 ?Q

1 58

s «J 1

% ° (^Pennsylvania

QO.4?

2.4°,

2.4C

4.6

CHAPTER IX
CARBON COMPOUNDS

THE compounds of the element carbon are so numerous that it
has been found convenient to constitute the study of these sub-
stances a separate branch of chemistry. In the early history of
the science it was believed that there were a large number of
substances which could only be obtained as the product of living
organisms. They were known to be elaborated by the action of
life, or, as it was termed, the vital force •, and it was believed that
owing to some inherent specific quality belonging to this vital force
the substances produced by its action were distinct from such
substances as could be prepared by any laboratory processes. To
denote this distinction, the term organic was applied to those things
which were known to be the products of living organisms, and
other compounds were distinguished as inorganic substances. This
distinction received its deathblow in 1828, when Wohler produced,
by purely laboratory processes, one of the most typical of all organic
compounds, namely, urea. The names " organic " and " inorganic "
chemistry are still retained, but their old significance is entirely
gone, as no distinction is to-day recognised between products elabo-
rated by the action of life and those which can be synthetically
produced.

Speaking broadly, organic chemistry may be described as the
chemistry of the carbon compounds. Nevertheless, although it is
quite true that all " organic compounds " contain carbon, it has not
been found expedient to include in the category of organic com-
pounds all compounds containing carbon. Not because there is
any intrinsic difference in these compounds, but merely from con-
siderations of convenience. The following may be mentioned as
examples of such compounds as are not regarded as belonging to
the " organic " division : compounds of carbon with the metals,
namely, the so-called carbides, of which cast iron and calcium

carbide are familiar cases ; the compounds of carbon with sulphur

295

296 Inorganic Chemistry

and the extensive series of thio-carbonates ; carbon monoxide
and the compounds formed by its direct union with non-metals
{e.g. carbonyl chloride, &c.) and with metals (e.g. nickel carbonyl,
&c.) ; and lastly, carbon dioxide and all the multitude of metallic
carbonates. Obviously, therefore, the broad distinction above
mentioned must not be regarded as a definition. Indeed, it
may be said that no exact definition of an " organic " compound has
ever been framed, and we have to accept the general statement
that " organic" chemistry is the chemistry of the carbon com-
pounds with certain generally acknowledged exceptions.

Amongst the compounds of carbon which will be briefly treated
of in these chapters, there will be included three which all chemists
agree to regard as organic substances : these are methane (marsh
gas), CH4 ; ethylene, C2H4 ; and acetylene, CaH2. These three com-
pounds play an important part in our ordinary illuminating flames
and in coal gas.

COMPOUNDS OF CARBON WITH OXYGEN.

There are two well-known oxides of carbon, both 'of which are
colourless gases, viz. : —

Carbon monoxide . ._ ; »'••!• t * CO.
Carbon dioxide . . . -..-,: :.;;>>- . CO2.

CARBON MONOXIDE.

Formula, CO. Molecular weight =28. Density =14.

Modes of Formation.— ( i.) Carbon] monoxide is formed when
carbon dioxide is passed over charcoal heated to bright redness —

CO2 + C = 2CO.

The same result is obtained when a slow stream of air or oxygen
is passed over red-hot charcoal contained in a tube. The first
action of the air on coming in contact with the carbon is to form
carbon dioxide, which, passing over the remainder of the heated
material, is deprived of a portion of its oxygen according to the
above equation. This operation goes on in an ordinary fire-grate :
the air, on first gaining access to the burning coal or coke, causes
the complete oxidation of a portion of the carbon to carbon dioxide ;
and as this gas passes through the mass of red-hot carbon it is
reduced to the lower oxide, which either escapes with the other

Carbon Monoxide

products of combustion or becomes ignited and burns with a
lambent bluish flame such as may frequently be noticed upon the
top of a " clear " fire.

(2.) When steam is passed over strongly heated carbon a mixture
of carbon monoxide and hydrogen is produced. This mixture,
known as water gas, is employed in many manufacturing processes
as a gaseous fuel —

H2O + C = CO + H2.

(3.) Carbon monoxide is also formed by the action of carbon
dioxide upon red-hot ^ron —

4CO2 + 3Fe = Fe3O4+4CO.

(4.) Or by strongly heating either carbon or iron with a car-
bonate, such as calcium carbonate, which is capable of yielding
carbon dioxide, thus —

CaCO3 + C = CaO + 2CO.
4CaCO3 + 3Fe = Fe3O4 + 4CaO + 4CO.

(5.) Carbon monoxide is most conveniently prepared, by the
decomposition of certain organic compounds by means of sulphuric
acid. Thus, when formic acid, or a formate, is acted upon by sul-
phuric acid, the sulphuric acid withdraws the elements of water
from the molecule of formic acid, and leaves carbon monoxide —

H'COOH-H2O = CO.

(6.) By a similar decomposition, oxalic acid yields a mixture of
carbon monoxide and carbon dioxide in equal volumes —

C2H204-H20 = C02 + CO.

The carbon dioxide is readily removed from the mixture, by passing
the gases through a solution of sodium hydroxide (caustic soda),
in which carbon dioxide is absorbed with the formation of sodium
carbonate,

(7.) The method usually employed when carbon monoxide is
required for experimental purposes consists in heating a mixture
of one part by weight of crystallised potassium ferrocyanide (yellow
prussiate of potash) with ten parts of strong sulphuric acid in a
capacious flask, when the following reaction takes place —

K4FeC6N6 + 6H2SO4 + 6H2O = 2K2SO4 + FeSO4

298 Inorganic Chemistry

The six molecules of water required by the reaction are derived
partly from the acid employed and partly from the salt, which
contains three molecules of water of crystallisation.*

Properties. — Carbon monoxide is a colourless, tasteless gas,
having a faint smell. It is only slightly soluble in water, its co-
efficient of absorption at o° being 0.03287. It burns in the air with
a characteristic pale-blue flame, forming carbon dioxide —

When mixed with half its own volume of oxygen, and inflamed,
the mixture explodes with some violence. t If the two gases be
confined in a eudiometer standing over mercury, and be rendered
absolutely free from aqueous vapour by powerful desiccating agents,
no explosion will take place upon the passage of an electric spark
through the mixture. And in the same way, if carbon monoxide,
which has been deprived of all aqueous vapour, be burned from a
jet in the air, and the jet be lowered into a cylinder containing air
which has been similarly dried, the flame will be extinguished
(see page 191).

Carbon monoxide is an extremely poisonous gas : very small
quantities present in the air rapidly give rise to headache and
giddiness, and if inhaled for a length of time, or if taken into the
lungs in a less dilute condition, insensibility and death quickly
follow. The deaths that have resulted from the use of unventi-
lated fires — either of charcoal or coke, or in some cases of coal gas
— in dwelling-rooms, have been due to the escape of this poisonous
gas into the air. The extremely deadly nature of the after-damp
resulting from a colliery explosion is due to the presence of carbon
monoxide in the carbon dioxide which is formed as a product of
the combustion.

The poisonous action of this gas is due to its absorption by the blood, with
the formation of a bright red compound, to which the name carboxy-hauno-
globin is applied. Blood so charged appears to be unable to fulfil its function
of absorbing and distributing oxygen throughout the system. This carboxy-
hasmoglobin gives a characteristic absorption spectrum, which furnishes a
ready method of detection in cases of poisoning from the inhalation of carbon
monoxide.

* " Chemical Lecture Experiments," new ed., 435-439.

f The rate at which the combustion is propagated throughout a mixture of
carbon monoxide and oxygen is much slower than through hydrogen and
oxygen. Bunsen has estimated it at less than i metre per second.

Carbon Monoxide 299

Carbon monoxide is one of the most difficultly liquefiable gases,
its critical temperature being - 1 36°.

At high temperatures this gas is a powerful reducing agent,
uniting with another atom of oxygen to form carbon dioxide.
This fact is made use of in many metallurgical processes for re-
ducing the oxides of the metals to the metallic state.

Carbon monoxide is absorbed at ordinary temperatures by a
solution of cuprous chloride, forming the compound COCu2Cl2.

At a temperature of boiling water, carbon monoxide is slowly
absorbed by solid potassium hydroxide, with the formation of
potassium formate —

H-COOK.

Carbon monoxide unites directly with chlorine, under the in-
fluence of sunlight, forming the compound known as phosgene gasy
or carbonyl chloride —

CO + C12 = COC12.

If the two gases are mixed in equal volumes, and kept in the
dark, no action takes place, but on exposure to sunlight they com-
bine, and the yellowish colour due to the chlorine will disappear.
On opening the vessel in moist air, clouds of hydrochloric acid
are formed, owing to the decomposition of carbonyl chloride by
the moisture, according to the equation —

Carbonyl chloride may be readily condensed to a liquid, its
boiling-point being + 8°r

Carbon monoxide unites directly with certain metals, giving rise
to compounds which possess some very remarkable properties,
and to which the name metallic carbonyls has been applied by
their discoverer.*

When carbon monoxide is allowed to stream slowly over metallic
nickel (obtained by the reduction of nickel oxide in a stream of
hydrogen), the gas is absorbed by the finely-divided metal, forming
a compound having the composition Ni(CO)4. If the issuing gas
be passed through a cooled tube, the nickel carbonyl condenses
to a colourless, mobile, highly refracting liquid, having a specific
gravity at o° of 1.356, and boiling at 43° under a pressure of
751 mm.t

* Mond, 1890.
t See "Chemical Lecture Experiments," new ed., 446-448.

306 Inorganic Chemistry

Nickel carbonyl vapour burns with a luminous flame, which
produces a black deposit of metallic nickel when a cold porcelain
dish is depressed upon the flame. The gas is decomposed into
nickel and carbon monoxide if passed through a hot glass tube,
the nickel being deposited as a bright metallic mirror upon the
glass—

Ni(CO)4 =

A similar compound of carbon monoxide/and iron has also been obtained,
having the composition Fe(CO)5. Iron carbonyl is a pale-yellow, viscous
liquid, boiling at 102. 8° under a pressure of 749 mm. Its specific gravity at
18° is 1.4664. When heated to 180° the vapour is decomposed, iron being
deposited and carbon monoxide being evolved. This compound has been found
in iron cylinders in which the so-called -water gas (a mixture of H and CO) has
been stored under pressure for a length of time ; it is also said to be present in
minute quantities in coal gas.

CARBON DIOXIDE.

Formula, CO2. Molecular weight =44. Density =22.

History. — Van Helmont in the seventeenth century was the first
to distinguish between this gas and ordinary air : he observed that it
was formed during the processes of combustion and fermentation,
and he applied to it the name gas sylvestre. Black showed that
this gas was a constituent of what in his day were known as the
mild alkalis (alkaline carbonates), and on account of its being so
combined, or fixed, in these substances, he named the gas fixed
air. Lavoisier first proved its true chemical composition to be
that of an oxide of carbon.

Oeeurrenee. — Carbon dioxide is a constant constituent of the
atmosphere, being present to the extent of about 3 volumes in
10,000 volumes of air. It is also found in solution in all spring-
water, which is sometimes so highly charged with this gas under
pressure that the water is effervescent, or " sparkling," from the
escape of the gas. Carbon dioxide is evolved in large quantities
from vents and fissures in the earth in volcanic districts. The
well-known Poison Valley in Java, which is an old volcanic crater,
and the Grotto del Cane near Naples, owe their peculiar pro-
perties to the discharge into them of large quantities of carbon
dioxide from such subterranean sources.

Carbon Dioxide

3oi

Modes Of Formation.— (i.) Carbon dioxide* is produced when
carbon is burnt with a free supply of air or oxygen _

If an insufficient supply of oxygen be employed, carbon mon-
oxide is produced at the same time.

(2.) When limestone or chalk is strongly heated, as in the
process of burning lime, carbon dioxide is evolved in large
quantities —

(3.) In the ordinary processes of fermentation, and during the
decay of many organic substances, carbon dioxide is also formed.

FIG. 62.

Thus, when sugar undergoes alcoholic fermentation by means of
yeast, the sugar is converted into alcohol and carbon dioxide —

(4.) Carbon dioxide is formed during the process of respiration ;
also by the combustion of all ordinary fuels, and of any compound
containing carbon, such as candles, oils, gas> &c.

(5.) For experimental purposes, carbon dioxide is most readily
obtained by the decomposition of a carbonate by means of a
stronger acid. The effervescence that results from the action of
tartaric acid upon sodium bicarbonate, in an ordinary Seidlitz

* Experiments on carbon dioxide, Nos. 400-434, " Chemical Lecture Ex-
periments," new ed,

302

Tnorganic Chemistry

powder, is due to the disengagement of this gas. The most con-
venient carbonate for the preparation of this gas is calcium
carbonate, in one of its many naturally occurring forms, such as
marble, limestone, or chalk. f Fragments of marble are for this
purpose placed in a two-necked bottle (Fig. 62), with a quantity of
water, and strong hydrochloric acid is added by means of the
funnel-tube. A rapid effervescence takes place owing to the

FIG. 63.

FIG. 64.

elimination of the gas, and a solution of calcium chloride remains
in the bottle—

CaCO3 + 2HCl = CaCl2 + H2O + CO2.

If sulphuric acid be substituted for hydrochloric acid, the frag-
ments of marble rapidly become coated with a crust of insoluble
calcium sulphate, which soon prevents the further action of the
acid, and therefore puts an end to the reaction : by employing
finely powdered chalk, however, instead of lumps of calcium car-
bonate, this difficulty is obviated. This gas is largely manu-
factured from these materials.

Properties. — Carbon dioxide is a colourless gas, having a feeble

Carbon Dioxide 303

acid taste and a faint and pleasantly pungent smell. It is incap-
able of supporting either combustion or respiration; a burning
taper is instantly extinguished, and an animal speedily dies when
introduced into this gas. Although carbon dioxide is not such a
poisonous compound as the monoxide, it nevertheless does exert
a direct poisonous effect upon the system, and death caused by
this gas is not merely due to the absence of oxygen. The pro-
longed inhalation of air containing only a very slightly increased
amount of carbon dioxide has a distinctly lowering effect upon
the vitality.

Carbon dioxide is a heavy gas, being about one and a half
times heavier than air. On this account it may readily be col-
lected by displacement. By virtue of its great density it may be
poured from one vessel to another, much in the same way as an
ordinary liquid : thus, if a large bell-jar be filled with the gas by
displacement, a beaker-full may be drawn up, as water " from a
well (Fig. 63). If the gas so drawn up be poured into a similar
beaker, suspended from the arm of a balance, and counterpoised,
the weight of the gas will be evident by the disturbance of the
equilibrium of the system.

If a soap bubble be allowed to fall into a large jar filled with
carbon dioxide, it will be seen to float upon the surface of the
dense gas (Fig. 64). The power of carbon dioxide to extinguish
flame is so great, that a taper will not burn in air in which this gas
is present to the extent of 2.5 per cent., and in which the oxygen
is reduced to 18.5 per cent. For this reason a comparatively small
quantity of carbon dioxide, brought into the air surrounding a burn-
ing body, is capable of extinguishing the flame. This property has
been put to valuable service in the construction of numerous con-
trivances for extinguishing fire, such as the " extincteur." This is
a metal vessel containing carbon dioxide under pressure, the gas
having been generated within the closed apparatus by the action
of dilute sulphuric acid upon sodium carbonate. A stream of the
gas, projected judiciously upon a moderate conflagration in a
dwelling, readily extinguishes the fire. This property may be
illustrated by inflaming a quantity of turpentine in a dish, and
pouring upon the flames a quantity of carbon dioxide contained
in a large bell-jar (Fig. 65), when it will instantly extinguish the
conflagration.

Although carbon dioxide is incapable of supporting combus-
tion in the ordinary sense, certain metals are capable of burn-

304 Inorganic Chemistry

ing in this gas. Thus, a fragment of potassium when heated
in this gas burns brightly, forming potassium carbonate with
the deposition of carbon —

When carbon dioxide is passed into a solution of calcium
hydroxide (lime water) a turbidity at once results, owing to the
precipitation of insoluble calcium carbonate or chalk —

CaH2O2 + CO2= CaCO3 + H2O.
This reaction furnishes the readiest means for the detection of

FIG. 65.

carbon dioxide. Thus, if the gas obtained by any of the modes of
formation described be passed into clear lime water, the formation
of this white precipitate of chalk is proof that the gas is carbon
dioxide. By this test it may readily be shown that carbon dioxide
is a product of respiration, by merely causing the exhaled breath
to bubble through a quantity of lime water, which will quickly be
rendered turbid.

Carbon dioxide is moderately soluble in water. At the ordinary
temperature water dissolves about its own volume of this gas.

The coefficient of absorption at o° is 1.7967, the solubility de-
creasing with rise of temperature in accordance with the inter-
polation formula —

<r= 1.7967 -0.07761 /+ 0.00 1 6424/2,

Carbon Dioxide 305

Carbon dioxide shows a slight departure from Henry's law
(see page 143), when the pressures are greater than that of the
atmosphere. Thus, when the pressure is doubled, the amount dis-
solved is slightly more than doubled. The solubility of carbon
dioxide in water, and its increased solubility under pressure, is
illustrated in the ordinary aerated waters. Water under a pres-
sure of several atmospheres is saturated with the gas, and upon
the release of this pressure by the withdrawal of the cork the
excess of gas, over and above that which the water can dissolve at
the ordinary pressure, escapes with the familiar effervescence. In
a similar manner the natural aerated waters have thus become
charged with carbon dioxide, under subterranean pressure, and
when such waters come to the surface the dissolved gas begins
to make its escape.

The solution of carbon dioxide in water is feebly acid, turning
blue litmus to a port-wine red colour, characteristically different
from the scarlet red given by stronger acids. This acid may be
regarded as the true carbonic acid —

CO2 + H2O = H2CO3.

A recently-made sample of aerated water is seen to effervesce
more briskly and give off the dissolved gas more rapidly than
specimens that have been long preserved. In process of time the
dissolved carbon dioxide gradually combines with the water, with
the formation of carbonic acid, an unstable compound which slowly
decomposes into carbon dioxide and water, especially at a slight
elevation of temperature. Many of the naturally occurring aerated
waters, such as Apollinaris, when opened exhibit scarcely any
effervescence, but give off carbon dioxide gradually. Such waters
have in all probability been exposed to pressure for a great length
of time, and their dissolved carbon dioxide has almost entirely
combined to form carbonic acid. When such a water is gently
warmed a rapid stream of gas is evolved.

When carbon dioxide is strongly heated, as by the passage of
electric sparks, it is partially dissociated into carbon monoxide
and oxygen. This decomposition is never complete ; for when
the amount of these two gases in the mixture reaches a certain
proportion, they reunite to form carbon dioxide, and a point of
equilibrium is reached when as many molecules are united as are
dissociated in the same time.

Liquid Carbon Dioxide. — Carbon dioxide is easily liquefied.

306 Inorganic Chemistry

At - 5° it requires a pressure of 30.8 atmospheres ; at + 5°, 40.4
atmospheres; while at +15° a pressure of 52.1 atmospheres is
required.

Faraday first liquefied this gas, by introducing into a strong bent
glass tube a quantity of sulphuric acid and a few lumps of ammo-
nium carbonate, which were prevented from touching the acid by
means of a plug of platinum foil. The tube was then hermetically
sealed, and the acid allowed gently to come in contact with the
carbonate, which was at once decomposed with the formation of
ammonium sulphate and carbon dioxide. By the internal pres-
sure exerted by the evolved gas, aided by the application of cold
to one end of the bent tube, the gas condensed to a colourless
liquid.

Large quantities of this liquefied gas were obtained by Thilorier
by a precisely similar method, the experiment being perfonned in
strong wrought-iron vessels.

Liquid carbon dioxide is to-day manufactured on a large scale,
by pumping the gas into steel cylinders by means of powerful
compression pumps. The enormous volumes of carbon dioxide
evolved in the process of brewing, and which until quite recently
were allowed to escape into the atmosphere, are now utilised for
this purpose. The gas, as it is evolved from the fermenting vats,
is washed and purified, and pumped into steel bottles for the
market. In this form the gas is largely employed by manufac-
turers of aerated waters, and also as the refrigerating agent in
" cold storage."

Liquid carbon dioxide is a colourless and extremely mobile
liquid, which floats upon water without mixing. It boils at - 80°
under atmospheric pressure.

When heated, liquid carbon dioxide expands at a more rapid
rate than a gas, its coefficient of expansion being greater than that
of any known substance. Its rapid change of volume is seen by
the following figures : —

95 volumes at — 10° become
100 „ „ o° „
106 „ „ + 10° „
114 „ „ +20°

The critical temperature of carbon dioxide is 31.35*." If the liquid
be heated to this point, it passes into the gaseous state without any
change of volume. The line of demarcation between the liquid

Carbon Dioxide

307

and gas in the tube gradually fades away, and the tube appears
filled with gas. Above this temperature no additional pressure is
able to liquefy the gas. On once more cooling the tube, when the
critical point is passed the liquid again appears, and the dividing
line between it and the gas is once more sharply defined.

Solid Carbon Dioxide.— When liquid carbon dioxide is allowed
to escape into the air the absorption of heat due to its rapid eva-
poration causes a portion of the liquid to solidify. This solid is
most conveniently collected by allowing the jet of liquid to stream
into a round metal box (Fig. 66), in which it is caused to revolve by
being made to impinge upon the curved tongue of metal. The box
is furnished with hollow wooden handles, through which the gas
makes its escape. Considerable quantities of the frozen carbon
dioxide can in this way be collected in a few minutes.

On a larger scale the brass box is
substituted by a canvas bag, which is
simply tied over the nozzle of the cylinder
containing the liquefied gas, and a rapid
stream of the liquid allowed to escape
into it.

Solid carbon dioxide is a soft, white,
snow-like substance. When exposed to
the air it quickly passes into gas, without
going through the intermediate state of
liquidity.

Solid carbon dioxide is readily soluble
in ether, and this solution constitutes one
of the most convenient sources of cold.
A large number of gases can readily be
liquefied by being passed through tubes
immersed in this freezing-mixture. When this ethereal solution
is rapidly evaporated its temperature can be lowered to - no*.

" Carbonic acid snow," as this substance is sometimes termed, is
now an article of commerce, the compound being sent into the
market in this form to avoid the cost of the carriage of the neces-
sarily heavy steel bottles containing the liquid.

Composition of Carbon Dioxide.— When carbon burns in
oxygen, the oxygen undergoes no change in volume in being trans-
formed into carbon dioxide. Thevolume of carbon dioxide produced
is the same as that of the oxygen which is required for its produc-
tion. This may be shown by means of the apparatus (Fig. 67),

FIG. 66.

308

Inorganic Chemistry

The bulb of the U-tube is filled with oxygen, and the stopper,
which carries a small bone-ash crucible upon which a fragment
of charcoal is placed, is lowered into position. The charcoal is
ignited by means of a thin loop of platinum wire, as shown in the
figure, which can be heated by an electric current. As the carbon
burns the heat causes a temporary expansion of the included gas ;
but after the combustion is complete and the apparatus has
cooled, the level of mercury will be found to be undisturbed.
Carbon dioxide, therefore, contains its own volume of oxygen.
From this experiment the composition of carbon dioxide by weight
can be deduced. One litre of carbon dioxide weighs 22 criths ;
deducting from this the weight of I litre of oxygen, viz., 16
criths, we get 6 as a remainder. Six parts by
weight of carbon, therefore, combine with 16
parts by weight of oxygen to form 22 parts
of carbon dioxide : expressing this proportion
atomically, the proportion of carbon to oxygen
is 12 to 32.

The gravimetric composition of carbon
dioxide may be directly determined by the
combustion of a known weight of pure carbon
in a stream of oxygen gas, and absorbing and
weighing the carbon dioxide that is formed.
This was done with great care and accuracy
by Dumas and Stas in the experiments by
which they determined the atomic weight of
carbon. Fig. 68 represents the apparatus em.
ployed for this purpose. A weighed quantity
of diamond, contained in a small platinum boat, was introduced into
a porcelain tube, which could be strongly heated in a furnace. The
oxygen for its combustion was contained in a glass bottle, from
which it could be expelled by allowing water to enter through the
funnel. As it was necessary that the oxygen should be absolutely
free from any carbon dioxide, the water used in the little gas-
holder contained potassium hydroxide in solution. The oxygen
was then passed through the tubes A, B, C, in order to deprive it
of carbon dioxide and moisture, and lastly through a small desic-
cating tube, d, which was weighed before and after the experiment.
The pure dry oxygen then entered the strongly heated tube, and
the carbon there burnt away to carbon dioxide, leaving a minute
quantity of ash, which was carefully weighed at the conclusion of

FIG. 67.

Carfion Dioxide

309

the experiment. A small layer of copper oxide was placed in the
tube, in the position indicated
in the figure, in order to oxidise
any traces of carbon monoxide
which were liable to be formed
into the dioxide. The product
of the combustion was carried
forward by the stream of oxygen
through a series of tubes ; d'
is a small weighed desiccating
tube, the weight of which, if
the diamond used contained no
hydrogen, should remain un-
changed. It then passes through
the bulbs F and G, where the
carbon dioxide is entirely ab-
sorbed. To arrest aqueous
vapour, which would be carried
away from the solution in these
bulbs by the escaping oxygen,
the gas is passed through H,
containing fragments of solid
potassium hydroxide ; this tube
is weighed along with the potash
bulbs. K is a guard tube con-
taining fragments of solid potas-
sium hydroxide, in order to
prevent atmospheric carbon
dioxide and moisture from gain-
ing access to the weighed por-
tions of the apparatus.

The weight of the diamond
minus the weight of the ash
which was left gave the actual
weight of the carbon burnt ;
the increase in weight of the
tubes gave the weight of the
carbon dioxide which was pro-
duced, and this weight, minus
the weight of carbon used, gave
the weight of oxygen that wa§

3 TO Inorganic Chemistry

consumed. As a mean of a number of experiments, Dumas and
Stas found that 80 parts of oxygen by weight combined with 29.99
parts of carbon.

From a knowledge of the density of carbon dioxide and the
volume of oxygen it contains, we know that the molecule of this
gas contains two atoms ; therefore, by the simple equation —

80 : 32 : : 29.99 : 11.99,

11.99 parts of carbon combine with 32 parts of oxygen, and the
number 1 1.99 is therefore the atomic weight of carbon as determined
by these chemists.

The Carbonates.— Although carbonic acid, H2CO3, is a very
unstable compound, the salts it forms are stable. Being a dibasic
acid, it is capable of forming salts in which either one or both of
the hydrogen atoms have been replaced by an equivalent of a
metal : thus in the case of sodium we have —

(1) Disodium carbonate (normal sodium carbonate)

(2) Hydrogen sodium carbonate (bicarbonate of soda) . HNaCO3.

Similarly, with the divalent metal calcium, it is possible to form—

(1) Normal calcium carbonate . . " -; . . ' " > ; . CaCO3,
and —

(2) Hydrogen calcium carbonate (bicarbonate of lime) . CaH2(CO3)2.

The formation of carbonates by the action of carbon dioxide upon
the hydroxides may be illustrated by the following equations : —

2KHO + C02=K2C03+ H20.
CaH2O2 + CO2=.CaCO3 + H2O.

The first of these changes is the one that takes place when
carbon dioxide is absorbed by the potassium hydroxide employed
by Dumas and Stas in the course of their experiments already
described. The second equation represents the reaction which
results when carbon dioxide is passed into lime water. In this
latter case, if the gas be passed through the turbid solution for
some time, the turbidity will gradually disappear, and the solution
once more become clear. The normal calcium carbonate (CaCO3)
which is first formed, and which is insoluble, is converted into
the soluble bicarbonate, CaH2(CO3)2. If this solution be boiled,
this unstable salt is decomposed with the evolution of carbon

Carbonates 311

dioxide and water and the reprecipitation of the normal calcium
carbonate —

CaH2(CO3)2 = CaCO3 + H2O + CO2.

The presence of this compound in natural waters is associated
with the property known as the hardness of water (see Natural
Waters, p. 221).

When one volume of dry carbon dioxide is mixed with two volumes of dry
ammonia, the two gases unite, forming a compound known as ammonium
carbamate —

C02+2NH3=C02,2NH3 or £}g*o JCO,

which is the ammonium salt of the unknown carbamic acid, tjQ2 CO.

The relation between this compound and carbamide or urea will be obvious
by an inspection of the formula j^pj2 i CO. •

This substance was the first " organic " compound which was ever obtained
from purely inorganic sources (page 295). It can be obtained by the action of
carbonyl chloride upon ammonia —

COC12+4NH3=CO(NH2)2+2NH4C1.

Carbon Suboxide, C3O2. — By the action of phosphoric acid upon ethyl-
malonate under reduced pressure, and at a temperature about 300°, a mixture
of ethylene and carbon suboxide is obtained *—

CH2(COOC2H5)2=2C2H4+ C302 + 2H2O.

The products of the action are condensed in a vessel cooled by liquid air.
The ethylene is then allowed to vaporise at the ordinary temperature, leaving
the liquid suboxide behind. This is purified by volatilisation and recon •
densing the vapour in a tube cooled to about - 65°.

The compound is a colourless mobile liquid having a powerful acrid smell.
It boils at 7°, and the vapour burns with a blue but smoky flame, yielding
carbon dioxide.

The substance may be regarded as the anhydride of malonic acid. It
readily combines with water, forming malonic acid —

C3O2 + 2H2O=C3H4O4, or

O : C : C : C : O + 2H2O=HOOC • CH2 • COOH.

Diels and Wolf. Ber. , 1906.

CHAPTER X
COMPOUNDS OF CARBON WITH HYDROGEN

THESE two elements unite together in various proportions, form-
ing an enormous number of compounds known generally under
the name of the hydrocarbons. The reason for the existence of so
great a number of compounds of these two elements is to be found
in the fact that the atoms of carbon possess, in a very high degree,
the property of uniting amongst themselves. This property of
carbon gives rise to the formation of a number of groups or series
of compounds the members of which are related to each other
and to the simplest member of the series. Thus the compound
methane, CH4, is the simplest member, or the " foundation-stone,"
of a series of hydrocarbons of which the following are the first
four : —

Methane . . . CH4 I Propane . . . C3H8
Ethane. . . . C2H6 | Butane . . . C4H10

It will at once be seen that each compound differs in composi-
tion from its predecessor by an increment of CH2, and that each
may be expressed by the general formula, CnH2n+2.

Such a series of compounds is known as a homologous series^ and
any one member is called a homologue of any other.

In the following chapter the three hydrocarbons, methane,
ethylene, and acetylene, will be briefly studied. Each of these is
a " foundation-stone," or starting-point, of a series similar to the
one already mentioned ; thus —

Methane, CH4, first member of the CnHgn+2 series of hydrocarbons.
Ethylene, C2H4, ,, „ CnH211 „

Acetylene, C2H2 ,, ,, CnH«n-2 .- ••

METHANE (Marsh Gas— Fire-Damp}.
Formula, CH4. Molecular weight = 16.4. Density=8.2.
Occurrence. —Methane is found in the free state in large quan«
tities in nature. It is one of the products of the decompositions

Methane

3*3

which has resulted in the formation of the coal-measures. It is
therefore found in enormous quantities in coal mines, where it
not only occurs in vast pent-up volumes, under great pressure,
which escape with a rushing sound when the coal is being hewn ;
but it is also occluded within the pores of the coal. Methane is
also evolved from petroleum springs.

The name marsh gas has been given to this compound, on
account of its occurrence in marshy places by the decomposition of
vegetable matter. The bubbles of gas which rise to the surface
when the mud at the bottom of a pond is gently disturbed consist
largely of marsh gas.

Modes of Formation. — (i.) When a mixture of sodium acetate
and sodium hydroxide is strongly heated
in a copper retort, sodium carbonate is
produced and marsh gas is evolved —

The gas obtained by this reaction always
contains more or less hydrogen.

(2.) Pure methane may be obtained by
the decomposition of zinc methyl, by means
of water —

Zn(CH3)2 + 2H2O = ZnH2O2 + 2CH4.

(3.) The most convenient method for
preparing methane is by the action of zinc-
copper couple upon methyl iodide.* For

FIG. 69.

this purpose the zinc-copper couple is placed in a small flask, and
a mixture of equal volumes of methyl iodide and methyl alcohol is
introduced by means of the stoppered funnel (Fig. 69). The gas
is caused to pass through a tube rilled with the zinc-copper couple,
whereby it is deprived of any vapour of the volatile methyl iodide,
and is collected over water in the pneumatic trough.

The reaction which takes place is essentially a reduction of the
iodide by means of the nascent hydrogen produced by the action
of the zinc-copper couple upon the alcohol or the water present,
and may therefore be represented by the equation —

CH3I+2H

HI.

* " Chemical Lecture Experiments," new ed. , No.

314 Inorganic Chemistry

The hydriodic acid must not be regarded as escaping as such,
but in the presence of the zinc forming a compound with it. If water
only is present, the compound Znl'HO is formed ; while if methyl
alcohol is employed the zinc compound will have the composition
ZnI'CH3O : the complete equation (omitting the copper which does
not enter into the chemical change) being —

Marsh gas is formed during the process of the distillation of coal, and is
therefore a large constituent of coal gas, j,he amount varying from 35 to 40
per cent.

Properties. — Methane is a colourless gas, having no taste or
smell. It burns with a pale, feebly luminous flame. When mixed
with air or oxygen and ignited the mixture explodes with violence.
The products of its combustion are water and carbon dioxide,
the methane requiring twice its own volume of oxygen for its
complete combustion, and yielding its own volume of carbon
dioxide —

Methane is only about one-half as heavy as air, its specific
gravity being 0.55 (air= i). The fire-damp of coal mines is nearly
pure methane, its average composition being —

Methane . ,H -.-••.'''""» '•'••' - v* '^ 96.0
Carbon dioxide . . . •* . \ 0.5
Nitrogen .- .. •:'^;< . . -3-5

100.0

ETHYLENE (Olefiant Gas}.
Formula, C2H4. Molecular weight =28. 4. Density =14. 2.

Modes of Formation.—(i.) This compound is obtained when
ethyl iodide is acted upon by an alcoholic solution of potassium
hydroxide —

C2H5I + KHO = KI + H2O + C2H4.

(2.) It is also formed when ethylene dibromide is brought in con-
tact with zinc-copper couple, the ethylene dibromide being diluted
with its own volume of alcohol—

C2H4Br2 + Zn = ZnBr2 + C2H4.

(3.) Ethylene may be prepared by acting upon alcohol with
certain powerful dehydrating agents, such as phosphoric pent-

Ethylene 315

oxide or sulphuric acid, the latter being most commonly em-
ployed. The mixture of alcohol and sulphuric acid is heated in
a flask to about 165°, when a brisk effervescence takes place. From
the point of view of the final products, the reaction may be re-
garded as the abstraction of the elements of water from the
alcohol,* thus —

C2H60-H20 = C2H4.

The action, however, is always accompanied by secondary re-
actions, which result in the rapid blackening of the mixture owing
to the separation of carbon. The sulphuric acid then acts upon
this carbon with the evolution of carbon dioxide and sulphur
dioxide. Hence the ethylene that is obtained by this process is
always contaminated with considerable quantities of these gases,
from which it must be purified by being passed through a solution
of sodium hydroxide.

(4.) Pure ethylene is most readily prepared by the action of
syrupy phosphoric acid (the ordinary tribasic acid) upon alcohol.t
About 50 or 60 c.c. of the acid are placed in a small Wurtz flask
of about 1 80 c.c. capacity. The flask is fitted with a cork carrying
a thermometer and a dropping-tube (Fig. 70), the end of the latter
being drawn out to a fine point and reaching to the bottom of the
flask. The acid is boiled for a few minutes until the concentration
reaches such a point that the temperature rises to 200°, when the
alcohol is allowed to enter drop by drop; the rate at which the alcohol
is admitted being visible in the dropping-bulb. By keeping the tem-
perature between 200° and 220°, a steady and continuous stream
of gas is evolved; which after being deprived of the small quantities
of ether and undecomposed alcohol with which it is accompanied,
by being passed through a small Woulf's bottle standing in ice,
is practically pure ethylene. The action of the phosphoric acid is
the same as that of sulphuric acid, the first action being the
formation of phosphovinic acid, which is subsequently decomposed
in a similar manner to the sulphovinic acid.

* In reality the action is more complex, and takes place in two stages, the
first being the formation of ethyl hydrogen sulphate or sulphovinic acid,
C2H5'HSO4— a compound which is analogous to hydrogen potassium sulphate,
KHSO4 ; and the second being the decomposition of this compound when
heated either alone or in the presence of sulphuric acid —
C2 H5 -OH + H2S04 = QHg -HSO4 + H2O,

C2H3-HS04=C2H4+H2S04.
f Newth, Jour. Chern. Soc.t 1901.

Inorganic Chemistry

Properties. — Ethylene is a colourless gas, having a somewhat
pleasant ethereal smell ; it burns with a highly luminous flame,
forming carbon dioxide and water, one volume of the gas requiring
three volumes of oxygen for its complete combustion, and produc-
ing twice its own volume of carbon dioxide —

C2H4 + 3O2 = 2CO2 + 2H2O.

If mixed with oxygen in this proportion and inflamed, the mix*
.1

ture explodes with great violence.

FIG. 70.

When mixed with twice its volume of chlorine and ignited, the
mixture burns rapidly with a lurid flame, with the formation of
hydrochloric acid and deposition of carbon —

Acetylene

Ethylene is rapidly absorbed by fuming sulphuric acid (more
slowly by the ordinary strong acid), forming ethyl hydrogen
sulphate —

C2H4+ H2SO4=C2H6'HSO4,

and from this compound, by distillation with water, alcohol may
be produced —

C2H6'HSO4 + H2O = C2H6'OH + H2SO4.

Ethylene is reduced to the liquid state, at a temperature of o°,
by a pressure of 41 atmospheres ; the critical temperature of
the gas is + 10.1°, at which point a pressure of 51 atmospheres is
required to liquefy it. Liquefied ethylene boils at - 103°, and by
increasing its rate of evaporation temperatures as low as - 140°
can readily be obtained. Ethylene (together with higher members
of the same series) constitutes the chief illuminating constituent of
ordinary coal gas, of which it forms from 4 to 10 per cent.

ACETYLENE.

Formula, C2H2. Molecular weight =26. 2. Density =13.1.

Modes Of Formation.— ( i.) Acetylene is capable of being syn-
thetically formed by the direct union of its elements. For this

FIG. 71.

purpose a stream of hydrogen is passed through a three-way globe,
in which an electric arc is burning between two carbon rods,
arranged as seen in Fig. 71 (a quantity of sand being placed in
the globe to prevent fracture from falling fragments of red-hot
carbon). Under these circumstances a small quantity of the
carbon and hydrogen unites to form acetylene, which is swept out
of the globe by the current of hydrogen.*

* The formation of acetylene appears to be a secondary result, due io the
high temperature decomposition of methane which is first produced (Boner
Jour. Chem. Soc.t 1897).

318

Inorganic Chemistry

(2.) Acetylene may be obtained by the action of alcoholic potash
upon ethylene dibromide. Alcoholic potash is heated in a flask,
and ethylene dibromide dropped upon it from a stoppered funnel,
when the following reaction takes place —

(3.) Acetylene is formed when marsh gas or coal gas is burned
with an insufficient supply of air for complete combustion ; thus,
when a Bunsen lamp becomes accidentally ignited at the base of
the chimney, the peculiar and unpleasant smell that is perceived is
partly, though not entirely, due to the formation of acetylene.

The formation of acetylene by the imperfect combustion of coal
gas is readily shown by causing a jet of air to burn in an atmos-

FIG. 72.

phere of coal gas, and aspirating the products of combustion
through a cylinder containing an ammoniacal solution of cuprous
chloride, as shown in Fig. 72. The acetylene is absorbed by the
ammoniacal cuprous chloride, forming a deep-red coloured com-
pound known as cuprous acetylide —

Cu2Cl2,2NH3 + H20 + C2H2 = 2NH4C1 + C2H2Cu2O *

When this compound is acted upon by hydrochloric acid, it is
decomposed with the evolution of acetylene, thus —

C2H2Cu2O + 2H Cl = Cu2Cl2 + H2O + C2H2.

* Keiser has shown that, when perfectly dry, the compound loses a mole-
cule of water, and has the composition C2Cu2, and not C2H2Cu2O (or
HaO) ; in fact, that the compound is a carbide of copper.

Acetylene

Formerly this method was commonly practised when any quantity
of acetylene was required.

(4.) For all practical purposes acetylene is now always prepared
by the action of water upon calcium carbide. The carbide may be
placed in a dry flask furnished with a dropping funnel and de-
livery tube, and on gradually admitting water drop by drop a
rapid evolution of nearly pure acetylene at once takes place —

CaC2 + 2H20 = Ca(HO)2 + C2H2.

Properties. — Acetylene is a colourless gas having an extremely
offensive smell, which rapidly induces headache ; when inhaled in
an undiluted state it is poisonous. The gas burns with a highly
luminous and smoky flame. When burnt from specially constructed
jets it gives a pure white light of [great intensity, and on this
account is a most important illuminant. Acetylene is present in
small quantities in ordinary coal gas, and its presence may be
detected by the formation of the red precipitate of cuprous acety-
lide when coal gas is allowed to bubble through an ammoniacal
solution of cuprous chloride. This reagent furnishes not only a
ready and delicate test for the presence of acetylene, but also
provides a means of removing this gas from admixture with other
gases. Thus, in the synthetic formation described above, the
gases issuing from the globe are passed into a flask containing
this solution, which immediately absorbs the acetylene. When
acetylene is subjected by prolonged heating to a temperature short
of a red heat, it undergoes polymerisation, and is converted into
liquid hydrocarbons, of which benzene, C6H6, is one.

Nascent hydrogen converts acetylene into ethylene —

From acetylene, therefore (a compound which can be syn-
thetically prepared from its elements, carbon and hydrogen), a
great number of " organic " compounds can be built up, for, as has
been already explained (page 317), from ethylene it is easy to
obtain alcohol, which opens the door to the preparation of a vast
number of other organic compounds.

Coal Gas.— When coal is distilled, the volatile products obtained
are : (i) coal tar ; (2) an aqueous liquid containing ammonia and
other products, and known as ammoniacal liquor; (3) coal gas.

Inorganic Chemistry

Coal gas, after being subjected to ordinary purification, is a
mixture of gases which maybe divided into three classes, namely :
illuminantS) diluents, and impurities. The most important of
these substances are —

f Ethylene, C2H4 ; propylene, C3H6 ; \

I butylene, C4H8 .... (CnH2n) (About 6.5

nts'\ Acetylene, C2H2 ; allylene, C3H4 . (CnH2n_2) (percent.

I Benzene, C6H6 (CnH2n_6) J

Diluents. — Hydrogen, marsh gas, carbon monoxide . About 90 per cent.
Impurities.— Nitrogen, carbon dioxide, 'sulphuretted

hydrogen . . . ; ~ . ,; . . About 3.5 per cent.

The composition of the gas is largely determined by the nature
of the coal employed, as may be seen from the following analyses
of gas from bituminous and from cannel coal :—

From Bituminous Coal.

London (Frankland).

From Cannel Coal.

Manchester

Hydrogen .
Marsh gas .
Carbon monoxide
Illuminants
Nitrogen .
Carbon dioxide .
Sulphuretted hydrogei

. 50.05
32.87
12.89
3.87

0.30

51.24
35.28
7.40

3.56

2.24

0.28

35.94
41.99

10.07
10.81

1.19

1 00. OP

100.00

Roscoe).

45.58

34.90

6.64

6.46

2.46

3.67

0.29

ICO.OO

CHAPTER 'XI
COMBUSTION

WHEN chemical action is accompanied by light and heat, the
phenomenon is called combustion. All exhibitions of light and
heat are not necessarily instances of combustion ; thus, when an
electric current is passed through a spiral of platinum wire, or
through a carbon thread in a vacuous bulb (as in the familiar
"glow" lamps), these substances become hot, and emit a bright
light. Neither the platinum nor the carbon, however, is under-
going any chemical change, and therefore the phenomenon is not
one of combustion. The materials are simply being heated to a
state of incandescence by external causes, and as soon as these
cease to operate, the glowing substances return to their original
condition unchanged.

Combustion may be defined as the chemical union of two sub-
stances^ taking place with sufficient energy to develop light and
heat. When the amount of light and heat are feeble, the combus-
tion is described as slow or incipient; while, on the other hand,
when they are considerable, the combustion is said to be rapid or
active. The true nature of combustion was not understood until
after the discovery of oxygen in 1775. From about the year 1650
until after that important discovery, the phlogistic theory was
universally adopted. According to this view, a combustible body
was one which contained, as one of its constituents, a substance or
principle to which the name phlogiston was applied. Easily com-
bustible substances were considered to be rich in phlogiston, while
those that were less inflammable were held to contain but little of
this ingredient. The act of combustion was regarded as the
escape of this principle from the burning substance. Thus, when
a metal was burnt in the air, it was considered to be giving off its
phlogiston, and the material that was left after the combustion
(which we now know to be the oxide of the metal) was regarded as
the other constituent of the metal, and was called the calx. The

32* ^

322 Inorganic Chemistry

metal, therefore, was supposed to be a compound of a calx with
phlogiston. By heating a calx along with some substance rich in
phlogiston, the former again combined with this principle and the
metal was once more produced. Thus, when the calx of lead was
heated with charcoal (a substance pre-eminently rich in phlo-
giston), the charcoal supplied the calx with the necessary amount of
phlogiston to produce the compound of calx of lead and phlogiston,
which was metallic lead. This theory of combustion, after sustain-
ing many severe shocks (from such experiments as those of Boyle
and others, who showed that the calx of a metal was heavier than
the metal used in its formation), received its death-blow on the
discovery of the compound nature of water, and that this substance
was produced by the combustion of hydrogen in oxygen.

In all processes of combustion it is customary to regard one of the
substances taking part in the chemical change as the combustible,
and the other as .the supporter of combustion. Usually that sub-
stance which surrounds or envelops the other is called the sup-
porter of combustion. Thus, when a jet of burning hydrogen is
introduced into a jar of chlorine, or when a fragment of charcoal
burns in oxygen, the chlorine and the oxygen are spoken of as the
supporters of combustion, while the hydrogen and carbon are termed
the combustibles.

In all the more familiar processes of combustion the atmosphere
itself is the enveloping medium, and the air is therefore, par excel-
lence, the supporter of combustion ; and in ordinary language the
terms combustible and incombustible are applied to denote sub-
stances which burn, or do not burn, in the air. By a similar
process of limitation, it has become customary to speak of other
gases as supporters or non-supporters of combustion, if they behave
towards ordinary combustibles as air does. Thus we say of hydro-
gen, or marsh gas, or coal gas, that they are combustible, but do
not support combustion ; and of oxygen, or chlorine, or nitrous
oxide, that they do not burn, but will support combustion ; and
lastly, of such gases as ammonia, or carbon dioxide, or sulphur
dioxide, that they neither burn nor support combustion.

This distinction, however, is a purely conventional one, and has
little or no scientific significance ; for, by a slight modification of the
conditions, either hydrogen, marsh gas, or coal gas may become
supporters of combustion, and oxygen, chlorine, or nitrous oxide
the combustible substances. Thus, when a jet of hydrogen burns
in oxygen, we say that the hydrogen is the combustible, and the

Combustion

323

oxygen the supporter of combustion (Fig. 73, A) ; but if a jet of
oxygen be thrust up into a jar of hydrogen (Fig. 73, B), it ignites
as it passes the burning hydrogen, and continues to burn in the
hydrogen.

By means of the apparatus shown in Fig. 74, this may be still
more strikingly shown.* A stream of hydrogen is passed into the
lamp chimney by the tube H, and the issuing gas inflamed as it
escapes at the top. Oxygen is admitted through the tube O, and
the jet of gas ignited by pushing the long tube up into the burning

FIG. 73-

FIG. 74.

hydrogen at the top, and then drawing it down to the position
shown in the figure, where the jet of oxygen continues to burn in
the atmosphere of hydrogen.

By means of the same apparatus, oxygen, or chlorine, or nitrous
oxide may be caused to burn in either hydrogen, marsh gas, or
coal gas. Ammonia, which, as already mentioned, is usually
described as being neither combustible nor a supporter of com-
bustion, when surrounded by an atmosphere of oxygen is readily
inflammable, and will as readily support the combustion of oxygen.

The atmosphere itself becomes the combustible body when the

* " Chemical Lecture Experiments," new ed., No. 367.

324

Inorganic Chemistry

usual conditions of combustion are reversed. Thus, if a stream of
coal gas be passed through a similar lamp glass, through the cork
of which a short straight glass tube passes (Fig. 75), air will be
drawn up through this tube, and may be inflamed by passing up a
lighted taper. The jet of air will then continue to burn as a non-
luminous flame. The air is here the combustible, and the coal gas
the supporter of combustion. If the excess of coal gas be inflamed
as it escapes from the top, the opposite conditions will be fulfilled,
the air being the supporter of combustion, and the coal gas the
combustible.

This interchangeableness of the terms combustible and sup-
porter of combustion applies
also to substances that are
liquid or even solid at the
ordinary temperature. If a
small quantity of some inflam-
mable liquid, as ether, carbon
disulphide, turpentine, &c., be
boiled in a flask, and the issu-
ing vapour inflamed, a jet of
oxygen gas when lowered into
the flask will ignite as it passes
the flame, and continue to burn
in the vapour of the liquid.
In the same way, sulphur,
which is a combustible solid,
and whose vapour is inflam-
mable in the air, is capable in
the state of vapour of support-
ing the combustion of oxygen.

FIG. 75.

Since combustion is the result
of energetic chemical union,
and since also it is a mere condition of experiment which of the
two acting substances shall function as the environment of the
other, it will be seen that the terms " combustible " and " supporter
of combustion," as applied to a chemical substance, do not express
any definite or characteristic property of that body.

It was demonstrated by Boyle, that when a metal is burnt in
the air, the calx (or oxide) that is obtained weighs more than the
metal employed, instead of less, as the phlogistic theory seemed to
demand. This fact, which the upholders of phlogiston found it so

Combustion

325

difficult to reconcile, is seen to be a necessary consequence of
combustion, considered from the modern point of view. In all
instances of combustion the weight of the products of the action
is equal to the total weight of each of the two substances taking
part in the chemical combination. When, for example, the metal
magnesium burns in the air, the weight of the product of the com-
bustion is equal to the weight of the metal, plus the weight of a
certain amount of oxygen with which it united in the act of burn-
ing. This gain in weight during combustion may be demonstrated
in a number of ways. Thus, if a small heap of finely divided iron,
obtained by the reduction of
the oxide, be counterpoised
upon the pan of a balance, and
then ignited, the iron will be
seen to burn, and as it burns
the balance will show that the
smouldering mass is increasing
in weight. In this case the
sole product of the combustion
is a solid substance, namely,
iron oxide, which remains upon
the pan of the balance j but
the same result follows when
the product of the action is
gaseous. Thus, for instance,
when a fragment of sulphur is
burnt, although it disappears
from sight, it, like the iron,
combines with oxygen to form
an oxide. This oxide, however,
being a gas, escapes into the
atmosphere. If the sulphur be burnt in such a manner that the
sulphur dioxide is collected and weighed, it also will be found to
be heavier than the original sulphur. In the process of burning,
i gramme of sulphur unites with about I gramme of oxygen, and
the product therefore weighs 2 grammes. By causing an ordinary
candle to burn in the apparatus shown in Fig. 76, where the in-
visible products of its combustion are arrested, the increase in
weight may easily be seen. The candle being essentially a com-
pound of carbon and hydrogen, the products of its burning will be
carbon dioxide and water, both of which will be absorbed by the

FIG. 76.

326 Inorganic Chemistry

sodium hydroxide in the upper part of the tube. Consequently,
as the candle burns away, the arrangement gradually gains in
weight ; the increase being the weight of the atmospheric oxygen
which has combined with the carbon and the hydrogen to form
the compounds carbon dioxide and water.

Heat Of Combustion. — During the process of combustion, a
certain amount of heat is evolved, and a certain temperature is
attained — two results which are quite distinct. The temperature is
measured by thermometers or pyrometers, while the amount oj
heat is measured in terms of the calorie, or heat unit.*

The amount of heat produced by 'the combustion of any sub-
stance is the same, whether it burns rapidly or slowly, provided
always that the same final products are formed in each case.
Thus, when I gramme of phosphorus burns in the air to form
phosphorus pentoxide, it evolves 5747 calories ; and when the
same weight of phosphorus is burnt in oxygen, although the com-
bustion is. much more rapid and energetic, and the temperature
consequently rises higher, the amount of heat evolved is precisely
the same.

Again, when iron is heated in oxygen it burns with great bril-
liancy, and with evolution of much heat ; if, however, the same
weight of iron be allowed slowly to combine with oxygen, even
without any manifestation of combustion, it is found that the
amount of heat produced in forming the same oxide is absolutely
the same.

So far, therefore, as the quantity of heat produced is concerned,
there is no difference between active combustion and slow com-
bustion, or (confining ourselves to the case of combinations with
oxygen) between active combustion and the ordinary process of
spontaneous oxidation at ordinary temperatures. In the latter
case the heat is given out slowly — so slowly that it is conveyed
away by conduction and radiation as fast as it is produced, and
consequently the temperature of the material undergoes no per-
ceptible change. In the case of active combustion, the action is
crowded into a few minutes or seconds, and, as all the heat de-
veloped is evolved in this short space of time, the temperature
of the substances rapidly rises to the point at which light is
emitted.

That heat is developed during the process of spontaneous oxida-

* The major calorie sometimes used is equal to 1000 calories. See Thermo-
chemistry, Part I. chap. xv.

Heat of Combustion 327

tion is readily shown. Thus, if a small heap of fragments of
phosphorus be exposed to the air, it will be evident from the
formation of fumes of oxide that it is undergoing oxidation. As
the action proceeds, and as the heat produced by the oxidation is
developed more rapidly than it is radiated away (especially from
the interior portions of the heap), it will be seen that the phos-
phorus quickly begins to melt, and finally the temperature will
rise to the point at which active combustion begins, when the mass
will burst into flame.

It has been shown that many destructive fires have arisen from
masses of combustible material, such as heaps of oily cotton waste,
undergoing this process of spontaneous oxidation, until the heat
developed within the mass has risen sufficiently high to inflame
the material. To the operation of the same causes is to be
referred the spontaneous firing of haystacks which have been
built with damp hay, and also the spontaneous inflammation of
coal in the holds of ships.

As the temperature produced by combustion is augmented by
increasing the rapidity with which the chemical action takes place,
it will be at once obvious why substances which burn in the air,
burn with increased brilliancy and with higher temperature in pure
oxygen. In the air every molecule of oxygen is surrounded by
four molecules of nitrogen, therefore for every one molecule of
oxygen that comes in contact with the burning substance, four
molecules of this inert element strike it ; and by so doing they not
only prevent the contact of so much oxygen in a given interval of
time, but they themselves have their temperature raised at the
expense of the heat of the burning material. The number of
oxygen molecules coming in contact with a substance burning in
the air, in a given time, may be increased by artificially setting the
air in rapid motion : hence the increased rapidity of combustion
(and consequent rise of temperature) that is effected by the use of
bellows, or by increasing the draught by means of chimneys and
dampers.

The augmentation of temperature obtained by the substitution
of pure oxygen for air is well illustrated in the case of burning
hydrogen. The temperature of the flame of hydrogen burning
in oxygen, known as the oxy-hydrogen flame, is extremely high,
and when allowed to impinge upon a fragment of lime, it quickly
raises the temperature of that substance to an intense white heat,
when it emits a powerful light — the so-called oxy-hydrogen limelight.

328

Inorganic Chemistry

The following results obtained by Bunsen show the temperatures
reached by the combustion of hydrogen, and of carbon monoxide,
in air and in oxygen —

The flame of hydrogen burning in air . » . 2024°

„ „ „ oxygen . . 2844°

„ carbon monoxide burning in air . 1997°

» » » oxygen 3003°

It will be seen that whereas the flame of hydrogen in air is hotter than that
of carbon monoxide in air, when these gases burn in oxygen the temperature

' FIG. 77.

of the carbon monoxide flame is higher than that of hydrogen. This is due
to the partial dissociation of the water which results from the combustion of
the latter. It has been shown that when a mixture of hydrogen and oxygen, in
the proportion to form water, is ignited, the temperature produced by the
union of a portion of the mixture rises above the point at which water dis-
sociates ; and consequently for a certain small interval of time a condition of
equilibrium obtains, during which as many molecules of water are dissociated
as are formed : during this state the temperature falls, when rapid combus-
tion once more proceeds. It will be seen, therefore, that the limits to the
temperature which can be reached by combustion are influenced by the
points at which the products of combustion undergo dissociation.

Igmtion Point 329

Ignition Point. — The temperature to which a substance must
be raised in order that combustion may take place is called its
ignition point. Every combustible substance has its own ignition
temperature. If this point be below the ordinary temperature
the substance will obviously take fire when brought into the air,
without the application of heat ; such substances are said to be
spontaneously inflammable, and must necessarily be preserved out
of contact with air.

Passing from cases of spontaneous inflammability, we find a
very wide range existing between the igniting points of different
substances. Thus, a jet of gaseous phosphoretted hydrogen may
be ignited by causing it to impinge upon a test-tube containing
boiling water ; carbon disulphide vapour is inflamed by a glass
rod heated to 120°, while the diamond requires to be raised nearly
to a white heat before combustion begins.

The difference between the temperatures of ignition of hydrogen
and marsh gas may be
well seen by means of the

old steel mill of the miner ,/ .js^ If7 ^|i

(Fig. 77). By causing the " ~";\|f"s';
steel disk to revolve at a
high speed, while a frag-
ment of flint is lightly
pressed against its edge, a
shower of sparks is thrown
out; and on directing a jet FIG. 78.

of hydrogen upon these

sparks the gas is instantly ignited, while they may be projected
into a stream of marsh gas without causing its inflammation.
The same fact is also made strikingly apparent by depressing
a piece of fine wire gauze upon flames of marsh gas (or coal
gas) and hydrogen. In the former case the flame will not pass
through the gauze, although it may be shown that marsh gas
is making its way through by applying a lighted taper imme-
diately above the wire. If the gauze be held over the issuing jet
of gas the latter may be ignited by a taper upon the upper side of
the gauze, but the combustion will not be communicated to the
inflammable gas beneath (Fig. 78). The gauze conducts the heat
away from the flame so rapidly that the temperature of the metal
does not rise to the ignition point of the marsh gas on the other
side, and therefore the combustion cannot be propagated through

33O Inorganic Chemistry

the gauze. In the case of hydrogen, however, it will be found that
the instant the gas upon the upper side of the gauze is inflamed
the flame passes through and ignites the hydrogen beneath.*

It is upon this principle that the safety of the "Davy lamp " depends.
This consists of an ordinary oil lamp, the flame of which is surrounded by
a cylinder of wire gauze (usually made double at the top), through which air
to supply the flame freely passes in and the products of combustion pass out.
When such a lamp is taken into an atmosphere in which marsh gas is pre-
sent, this gas, entering through the gauze, becomes ignited within the chimney,
producing a very characteristic effect upon the lamp flame. According to the
amount of marsh gas present the flame is seen to become more and more
extended, at the same time becoming less luminous, until the whole interior
of the gauze cylinder is filled with the burning gas, emitting a faint bluish
light, known among the miners as the corpse-light. The burning marsh gas
is unable to communicate its combustion to the inflammable mixture outside,
for the same reason that the flame, in the experiment already referred to, was
unable to pass through the wire gauze. If from any cause the flame should
heat any spot of the gauze chimney to a temperature above the ignition point
of marsh gas, the outside combustible mixture will become ignited. It has
been shown that by exposing the lamp to a strong air draught the flame may
be so driven against the gauze as to unduly heat the metal. It has also been
proved that the same result frequently follows from the explosive wave that
is produced in a mine when, from some accidental cause, the operation"of
blasting (or shot-firing] results, not in the splitting of the rock, but in merely
blowing out the " tamping." The violent concussion to the air which follows
such a blown-out shot has been known to blow the flames of the Davy lamps,
even in remote parts of the workings, bodily through the gauze; and if such
lamps are burning at the time in an inflammable mixture, it would thereby be
fired.

By the behaviour of the flame of a Davy lamp when placed into an atmos-
phere containing marsh gas, it is possible to estimate, with a rough degree of
accuracy, the percentage amount of that gas which is present. For this pur-
pose the flame is turned down as low as possible, and the height to which the
burning marsh gas extends (the so-called fire-damp cap] is measured against a
scale graduated in tenths of inches. Fig. 79 (two-thirds the actual size) shows
the " caps " obtained by the presence of 4, 5, and 6 per cent, of marsh gas.f

When the ignition point of a substance is lower than the tem-
perature produced by its combustion, such a substance, when

* Recent experiments of Dixon and Coward (Trans. Chem. Soc., 1909) >
upon the ignition temperature of explosive gaseous mixtures, give the follow-
ing results:- !„ Oxygen. In Air.

Hydrogen . . . 580° to 590° 580° to 590°

Carbon monoxide . 637° „ 658° 644° ,, 658°

Methane . . . 556° „ 700° 700° ,, 750°

f In a recent development of this method of testing, a small hydrogen
flame is substituted for the oil-lamp flame, whereby it is possible to detect the
presence of 0.25 per cent, of marsh gas (Clowes).

Ignition Point 331

ignited, will continue to burn without further application of ex-
ternal heat, the inflammation being propagated from particle to
particle by the heat developed by their own combustion. All the
ordinary processes of combustion are actions of this order, and
belong to the class of chemical reactions known as exothermic^
that is to say, reactions which are accompanied by an evolution
of heat (page 168).

If, on the other hand, the ignition point be higher than the heat
produced by chemical union, combustion cannot proceed without
the continuous application of external heat. The igniting point of

FIG. 79.

nitrogen in oxygen, for example, is higher than the temperature
produced by the union of these elements ; therefore, although the
nitrogen may be ignited by the heat of the electric spark, it is
unable to communicate its combustion to contiguous particles, and
the inflammation does not spread. If the ignition point of nitro-
gen in oxygen had been lower instead of higher than the heat of
the chemical union of these elements, the first flash of lightning that
discharged into the air would have initiated a conflagration, which
would have extended through the whole atmosphere, and resulted
in the removal of the oxygen and its replacement by oxides of
nitrogen.

332 Inorganic Chemistry

The production of acetylene by the combination of carbon with hydrogen
under the influence of high temperature, and the formation of cyanogen and
carbon disulphide, by the union of the same element with nitrogen and with
sulphur respectively, are illustrations of the same class of action : phenomena
of this order being known as endothermic reactions, that is, reactions that are
attended with an absorption of heat (page 168).

Flame. — When both the substances taking part in combustion
are gases or vapours, the sphere of the chemical action assumes
the character of flame ; while, on the other hand, if one of the
materials is a solid which is not volatile at the temperature of its
combustion, no flame accompanies its burning. Such solids as
sulphur, phosphorus, camphor, wax, &c., during combustion in air,
undergo vaporisation, and consequently burn with the formation of
flame ; while such substances as iron, copper, carbon,* &c., which
do not pass into vapour at the temperature produced by their com-
bustion in oxygen, burn in this gas without giving rise to a flame.

Flames differ very widely in their general appearance, and in
the majority of cases are distinctly characteristic : thus, hydrogen
burns in air with a flame that is almost absolutely colourless, and
is scarcely visible in bright daylight ; sulphur burning in air pro-
duces a pale blue flame ; ammonia in oxygen a flame having a
yellow-ochre colour ; carbon monoxide a rich blue flame ; while
cyanogen burns with a flame having the delicate colour of the
peach blossom. Other flames are characterised by their luminosity.
Thus, phosphorus burning in oxygen emits a dazzling yellow light,
that is almost blinding to the eyes ; magnesium burns in the air
with an intense bluish-white light ; the flame produced by the
combustion of the vapour of nickel carbonyl in air emits a bright
white light ; and the flames that are produced by most hydro-
carbons during their combustion give a characteristic yellowish-
white light.

The General Structure of Flame.— The simplest form of
flame is one that is obtained by the combustion of a substance
which itself undergoes no decomposition, and in which the product
of combustion is arrived at in a single stage. Such flames, for
example, as that of hydrogen burning in chlorine or in air, or of
carbon monoxide burning in air. In the case of hydrogen burning
in air, the materials taking part in the process being elementary

* Under certain conditions the combustion of carbon in oxygen is accom-
panied by flame ; but it has been shown that at the temperature at which this
occurs carbon monoxide is being formed.

Flame

333

bodies, no complications arising from decomposition are possible ;
and although carbon monoxide is a compound, it unites with
oxygen without itself undergoing any decomposition, and passes
directly into carbon dioxide. Such flames as these, when burning
from the end of a tube, consist of a single hollow conical sheath
of actively burning gas. Fig. 80 represents a fla.me of burning
hydrogen : the darker region d is the hollow space within the flame,
consisting of unburnt hydrogen ; while the flame proper, the actual
burning portion, is the sheath £, which appears practically uniform
throughout. That the flame-cone is hollow may be proved by a
variety of experiments. Thus, if a sheet of white paper be quickly
depressed into a flame, a charred impression of the section of the
cone will be obtained, as shown in Fig. 81, from which it will be

FIG. 80.

FIG. 81.

seen that no combustion is taking place within the cone. In the
same way, an ordinary lucifer match may be suspended within the
flame, where it will remain without ignition so long as the burning
walls of the flame do not touch it. The shape of a flame is due to
the fact, that as the gas issues, the layer nearest to the walls of the
tube burn round the orifice of the tube as a ring, consequently the
next layer has to reach 'up above this ring before it can meet with
air for its combustion, and each successive layer has to pass up
higher and higher in order to find its supply of air, and in this way
the burning area is built up into the form of a cone. To show that
the hollow space consists of unburnt gas, it is only necessary to
insert a tube into the interior of the flame in such a way as to

334

Inorganic Chemistry

draw off a portion of the gas, when it will be found that the gas so
withdrawn will burn.

Passing from this simplest type to substances that undergo
decomposition during combustion, or which yield the final product
of oxidation by successive stages, it is found that the flames they
give rise to are less simple in structure.

As illustrations of various degrees of complexity, the following
examples may be mentioned : —

(i.) Ammonia burning in oxygen. ,This flame (Fig. 82) is very
characteristic, and on inspection it is at once obvious that it has a
less simple structure than the hydrogen flame. In this case the
inner hollow portion d is surrounded by a double flame-cone, the
inner cone a having a yellow-ochre colour, and
the outer portion b possessing a much paler colour,
and tending to green. During the combustion
of ammonia, the compound undergoes decomposi-
tion into nitrogen and hydrogen. This decom-
position, which begins in the hollow region d,
takes place mainly in the inner cone a, and the
hydrogen which escapes combustion in this
region passes to the outside, and there burns,
forming the outer cone. Probably there is also
a partial combustion of the nitrogen.

(2.) Carbon disulphide burning in air. This
flame, like the ammonia flame, consists of a
double flame-cone, consisting of an inner lilac-
' coloured cone, surrounded by an outer region
having a deeper blue colour. During combus-
tion carbon disulphide, like ammonia, is decomposed, but in this
case not only are both of the constituents readily combustible, but
the carbon passes into its final state of oxidation in two stages,
forming first carbon monoxide and afterwards carbon dioxide.

(3.) Hydrocarbons burning in air. The flames produced by the
combustion of these compounds include those which are commonly
employed for illuminating purposes, such as candle, gas, and oil
flames, and in all essential points of construction they are practi-
cally identical. This may be seen to be the case by a comparison
of the flames of a candle and of coal gas (Figs. 83 and 84). In
these flames, as in the former cases, there is the dark hollow space
4 consisting of heated unburnt gas (in the candle flame this gas
is generated by the vaporisation of the materials of the candle,

FIG. 82.

Flame

335

which in the melted condition are drawn up the wick by capillary
action). Above this there is a region, a, which, in comparison
with the rest of the flame, appears almost opaque, and which
emits a bright yellow light. This luminous area constitutes rela-
tively the largest part of the flame, and in flames that are used for
light-giving purposes it is intentionally made as large as possible
by means of various devices. At the base of the flame there is
a small region, r, which appears bright blue in colour, and is non-

FIG. 83.

FIG, 84.

luminous ; and surrounding the entire flame there will be seen a
faintly luminous mantle, b.

The flame proper, therefore, consists of three distinct parts,
namely : (i) the blue region r,at the base ; (2) the faintly luminous
mantle b ; and (3) the yellow, brightly luminous region a. These
three parts constitute the flame-cone, the actual area of combustion,
which envelops the dark region d ; this, as already stated, consists
of unburnt gas, and therefore is not, strictly speaking, a part of the
flame.

If the supply of gas to a flame, burning as represented in Fig. 84,
be diminished, or if air be slowly admitted to the interior, the flame

336 Inorganic Chemistry

will shrink down, and the luminous area become less and less,
until it finally disappears altogether. The flame-cone will then be
found to consist of two parts, resembling in structure the double
cone of the ammonia flame, Fig. 82. The blue region c, Fig. 84,
which is only fragmentary in the flame as there represented, will
have become continuous, and now constitutes the inner cone ;
while the mantle b forms the outer cone, the flame presenting the
appearance seen in Fig. 85. The region d^ as before, consists of
unburnt gas.

It has been shown, in the case of coal gas flames burning in this
manner, that in the inner cone c, the changes going on result
mainly in the formation of carbon monoxide and water, together
with small quantities of carbon dioxide and
hydrogen ; and that in the outer cone, or
mantle, the carbon monoxide and hydrogen
are burning to carbon dioxide and water.
In the inner cone, therefore, the carbon is
burnt to its first stage of oxidation, and a
portion of the hydrogen is oxidised to water ;
in the outer cone, the second stage of oxida-
pIG 85. tion of the carbon takes place by the com-

bustion of the carbon monoxide to carbon
dioxide, and the hydrogen which escapes combustion in the inner
cone is also burnt.

It has been known since the time of Dalton, that when certain
hydrocarbons are burnt with an insufficient amount of oxygen for
the complete oxidation of both the hydrogen and carbon, carbon
monoxide, water, and hydrogen are produced. This result is pro-
bably due to a secondary reaction ; the first stage being the com-
bustion of hydrogen to form water, which at the high temperature
is then decomposed, either by the carbon or the hydrocarbons,
according to the following equations —

CH4 + O2=2H2O + C = CO + H2+ H2O.

The various parts of an ordinary gas or candle flame, therefore,
are due to the different chemical reactions that are taking place in
these areas ; these changes are not of such a nature that they can
in all cases be perfectly traced, neither is one set of reactions
exclusively confined to each area, but rather is it the case that

Flame 337

certain chemical actions predominate in each particular part of the
flame.

In the blue region ^, Figs. 83 and 84, the main reactions going
forward are those already indicated, by which carbon monoxide,
water, and hydrogen are produced. In the faintly luminous
mantle £, carbon monoxide and hydrogen are burning, together
with small quantities of hydrocarbons which may have escaped
combustion and decomposition in the luminous region. The non-
luminous character of this mantle is due to the cooling effect of the
air which is drawn into the flame, and which even extinguishes
combustion upon the outer limits of the flame before every trace of
combustible material is burnt ; for it has been shown that small
quantities of carbon monoxide, marsh gas, and even hydrogen
escape unburnt from a gas flame.

The chemical decompositions which go on in the luminous area
cannot be said to have been thoroughly established. It has been
shown that very early in its passage up the flame a certain amount
of the marsh gas and ethylene present is converted into acetylene,
the change taking place as the result of heat alone. The gases
ascending the dark region d are surrounded on all sides by a wall
of burning material, and are thereby raised in temperature to the
point at which the marsh gas and ethylene suffer decomposition
into acetylene and hydrogen —

2CH4=C2H2+3H2.

The following table (Lewes) shows the gradual development of
acetylene in such a flame : —

Total Unsaturated

Hydrocarbons. Acetylene.

Per Cent. Per Cent.

Gas in burner ., . . « 4.38 0.035

\ inch above rim of burner . . 4.00 0.340

\\ inch above rim . . . 1.53 0.560

Tip of dark region ',;• *, . 1.98 1.410

Centre of luminous area . . . . 0.45 0.045

Tip of luminous area . . . o.oo o.oo

Therefore, by the time the gases have reached the tip of the dark
region, the effect of heat upon them has been to raise the amount
of acetylene to over 70 per cent, of the total unsaturated hydro-
carbons present. As the acetylene and other hydrocarbons pass
on through the flame along with steam, carbon dioxide, and

338 Inorganic Chemistry

carbon monoxide, other and more complex changes go on whereby
denser hydrocarbons are formed, and carbon itself is precipitated.

The formation of acetylene in that region of the flame where the
coal gas is in excess is well exemplified in the case of air burning
in an atmosphere of coal gas (see Fig. 75). In this flame the air
is in the inside and the coal gas upon the outside ; it is, in effect, an
ordinary coal gas flame turned inside out. The formation of acety-
lene, instead of taking place 'within the flame (in which case it has
to pass through the heated area and is thereby decomposed),
takes place upon the outer surface or periphery of the flame, and
therefore largely escapes combustion and decomposition and passes
away into the coal gas atmosphere. (See Acetylene, where this
method is described for the preparation of this compound.)

The Cause of Luminosity in Flames.— The light-giving property of a flame
is not due to the operation of any one simple cause. It was at one time sup-
posed that the luminosity of a flame depended solely upon the presence in it
of suspended solid matter resulting from the chemical decompositions going
on during combustion. It has been shown, however, that this general state-
ment does not satisfy all cases, as there are a number of highly luminous
flames in which, from the known properties of the products of combustion,
there cannot possibly be any solid matter present. Thus, for example,
phosphorus burning in air gives a flame of a high degree of luminosity ; but
the phosphorus pentoxide which is the product of combustion, although solid
at ordinary temperatures, is volatile at a temperature far below that of the
flame. The same may be said of the luminous flame of arsenic burning in
oxygen, where the still more volatile arsenious oxide is the product.

When carbon disulphide burns in oxygen or in nitric oxide, a well-known and
intensely luminous flame is obtained, in which only gaseous products of com-
bustion can be present ; and, lastly, the flame of hydrogen burning in oxygen
can be made under certain circumstances to emit a bright light : thus, when a
mixture of these gases is ignited in a closed eudiometer, their combustion is
attended with a brilliant flash of light, the only product being water.

There are three causes which may operate, either separately or together, in
imparting luminosity to a flame or in increasing its light-giving power : these
are— (i. ) The temperature of the flame ; (2. ) the density of the flame gases ; and
(3. ) the introduction into the flame of solid matter. These three causes will
be treated separately and illustrations given, which, so far as our knowledge
extends, can be directly traced to the independent operation of each.

(i.) The effect of temperature.

(a.) Upon flames in which solid matter is known to be absent.

When phosphorus is introduced into chlorine, it spontaneously inflames and
burns with a flame of such extremely feeble luminosity that it may be regarded
as non-luminous ; if, however, the chlorine be previously strongly heated by
being passed through a red-hot tube, and the phosphorus be boittng when it
comes in contact with the gas, the combustion thus started upon a higher

The Luminosity of Flames • 339

platform of temperature is accompanied by a flame of very considerable
luminosity.

The flame of carbon disulphide burning in air emits but a feeble light ; but
when this substance burns in pure oxygen, its temperature of combustion is
greatly raised and the luminosity of the flame is enormously increased.

Phosphoretted hydrogen burning in air gives a flame of considerable lumi-
nosity ; but when this flame is fed with pure oxygen, and its temperature
thereby raised, it becomes intensely luminous.

03.) Upon flames in which solid matter is known to be present.

The flames produced by the combustion of zinc or magnesium in the air,
and in which the solid oxides are present, have their luminosity greatly in-
creased when pure oxygen is substituted for air and the temperature of com-
bustion thereby augmented.

The same result is seen in the case of flames in which the solid matter is
artificially introduced, as in the familiar Welsbach burner, where a solid gauze
mantle, composed of an alkaline earth, is placed in the flame-cone of a non-
luminous gas flame, thereby rendering it luminous. If the temperature of
this flame be augmented by feeding it with oxygen, the light emitted by the
incandescent solid is greatly increased.

(7.) Upon flames in which solid matter is believed to be present, such as
candle, gas, and other hydrocarbon flames.

When a candle or gas flame is introduced into oxygen, although it shrinks
in size, its luminosity is increased. It has also been shown that when a coal
gas flame is chilled by causing it to spread against a cold surface, its luminosity
is diminished or destroyed altogether ; and, conversely, if the gas and the air
supplying the flame be strongly heated before combustion, the luminosity is
greatly increased. In this case, however, the direct effect of change of tem-
perature is complicated by the decompositions going on in the flame ; for, as
already mentioned, the conversion of the non-illuminating marsh gas into the
highly illuminating gas acetylene is a function of the temperature.

The increase of light obtained from a gas flame by previously heating
the gas and air is the principle underlying all the so-called recuperative
burners.

It is evident, therefore, that most flames gain luminosity by having their
temperature raised. There are, however, cases in which increase of tempera-
ture alone appears to exert no influence upon the luminosity. The flame of
hydrogen, for example, which is practically non-luminous when burning in
air, does not become more luminous when burnt in oxygen, although its
temperature is greatly increased.

(2.) The influence of the density of the flame gases.

It has been shown by Frankland * that the luminosity of flame is intimately
associated with the pressure to which it is subjected, or with the density of the
flame gases. Thus, it is found that a gas or candle flame, when burnt either
at high altitudes cr in artificially rarefied atmospheres, has its luminosity
greatly reduced ; and,/<?r contra, when caused to burn under increased pres-
sure, the luminosity is increased. In the case of hydrocarbons, complication
arises from the fact that the temperature of the flame is changed by alterations

* Phil. Trans., vol. cli. p. 629; Proc. Royal Society, vol. xvi. p. 419.

340

Inorganic Chemistry

of pressure. Under diminished pressure the temperature falls, and although
there is less loss of heat by radiation in rarefied air than in air at the ordinary
pressure, it is possible that the general lowering of the temperature of the
flame may modify the chemical decompositions in the direction already re-
ferred to.

Flames other than those of hydrocarbons, however, and in which no solid
matter can exist, are found to become luminous when the density of the flame
gas is increased by pressure. Thus, the flame of carbon monoxide in oxygen
at ordinary pressures emits a moderate light ; but when exposed to a pressure
of two atmospheres the luminosity is greatly increased. Even the non-luminous
flame of hydrogen burning in oxygen becomes luminous under a pressure of
two atmospheres, and when examined by the spectroscope is found to give a

FiG. 86.

FIG. 87.

continuous spectrum. It has been found, as a general rule, that dense gases
and vapours, when heated, become incandescent or luminous at much lower
temperatures than those of low specific gravity ; thus, if different gases be
raised to incandescence by the passage through them of electric sparks, under
similar conditions, ft is seen that the light emitted by the glowing vapour
varies with the density of the gas. The luminosity of glowing oxygen (density,
16) is greatly superior to that of hydrogen (density, i), while the light emitted
when the sparks are passed through chlorine (density, 35.5) is considerably^
advance of either. And it is found that in one and the same gas the luminosity
of the spark increases as the density is increased by artificial compression.
Other things being equal, it may be said that the denser the vapours which are
present the more luminous is the flame.
(3.) The introduction of solid matter into flames.

The Bunsen Flame 341

Non-luminous flames may be rendered luminous by the intentional introduc-
tion into them of solid matter, which, by being raised to a sufficiently high
temperature, will become strongly incandescent. Thus, the ordinary lime-
light owes its luminosity to the incandescence of the fragment of lime, which
is raised to a bright white heat by the high temperature of the non-luminous
oxy-hydrogen flame. The lime is not vaporised at the temperature of the
flame, the light being entirely due to the glowing solid matter.

The " Welsbach " burner, already referred to, is another example of the same
order, the luminosity in this case being due to the introduction into an
ordinary non-luminous Bunsen flame of a fine gauze mantle made of thoria
or other metallic oxide (Fig. 86). When such a mantle is raised to incandes-
cence by the heat of the gas flame, it emits a bright white light, strongly
resembling that of an ordinary Argand gas flame. A flame may also be
rendered luminous by the intentional precipitation within it of carbon, which,
by its ignition and its combustion, produces a high degree of luminosity. Thus,
if a small quantity of alcohol be boiled in a flask, and a jet from which chlorine
is issuing be then lowered through the burning vapour into the flask, as shown
in Fig. 87, the chlorine will burn in the alcohol vapour with a luminous flame ;
and the precipitated carbon (which is thrown out of combination by the action
of the chlorine upon the alcohol), ascending into the previously non-luminous
alcohol flame, will render it brightly luminous.

From these considerations it will be evident that the luminosity of a flame
may be due, first, to the presence of vapours sufficiently dense to become
incandescent at the temperature of the flame ; or, second, to the presence of
solids rendered incandescent, either by the heat of the flame gases alone, or
in conjunction with their own combustion ; or, third, from the simultaneous
operation of all these causes. Ordinary gas and candle flames come under
the last of these heads. The decompositions that go forward in these flames
not only give rise to dense vapours which become incandescent, but also to
the precipitation of solid carbon, which by its ignition and combustion adds
to the luminosity of the flame.

The Bunsen Flame. — The construction of the Bunsen lamp is too well
known to need description. The gas, issuing from a small jet situated at the
base of a metal tube, and mixing with air which is drawn in through openings in
the tube, burns at the top of the chimney with the familiar non-luminous flame.
The existence of this flame in its ordinary condition depends upon two main
causes ; first, upon the fact that in the immediate neighbourhood of a jet of
gas issuing from a small orifice, there is a reduction of pressure ; and, second,
upon the relation between the velocity at which the gases pass up the tube
and the rate of propagation of combustion in the mixture of air and coal gas.
Upon the first of these causes depends the entrance of air into the " air-holes "
of the lamp, and upon the second depends the continuance of the flame in its
position upon the top of the tube.

As the coal gas issues from the small jet at the base of the chimney, instead
of the gas escaping through the side-holes, air is drawn into the tube by virtue
of the reduced pressure produced immediately round the jet. That this area of
reduced pressure actually exists in the neighbourhood of the jet of a Bunsen
may be proved by attaching a delicate manometer to the air-hole of such a
lamp, as shown in Fig. 88. As the gas is turned on, the liquid in the horizontal

342 Inorganic Chemistry

tube will be sucked towards the lamp, showing that the issuing gas causes a
partial vacuum in its immediate neighbourhood.*

In order that the flame shall remain at the top of the tube, there must be a
certain relation between the velocity of the issuing gases and the rate of pro-
pagation of combustion in the mixture ; for if the latter be greater than the
former, the flame will travel down the tube and ignite the gas at the jet below.
By gradually reducing the supply of gas to the flame, and so altering the pro-
portion of gas and air ascending the tube, the mixture becomes more and
more explosive, until a point is reached when the velocity of inflammation is
greater than the rate of efflux of the gases, and the flame travels down the tube,
and the familiar effect of the flame " striking down" is obtained.

The same result may be brought about, an-d the effect more closely observed,
by extending the chimney of the lamp by means of a wide glass tube. As the
supply of gas is reduced, or the quantity of air introduced is increased, the
flame will be seen to shrink in size and finally descend the tube. By adjust-
ment it may be caused either to ex-
plode rapidly down the tube or to
travel quite slowly, or even to remain
stationary at some point in the tube,
which is slightly constricted, and where,
therefore, the flow of the issuing gas is
slightly accelerated, f

The non-luminosity of a Bunsen
flame is due to the combined opera-
tion of three causes, namely, oxidation,
dilution, and cooling. It was formerly
supposed that the destruction of the
luminosity of a gas flame by the ad-
mixture of air with the gas before
burning was entirely owing to the
influence of the oxygen in bringing
FiG. 88. about a more rapid and complete

state of oxidation, that the hydro-
carbons were at once completely burnt up by the additional supply of oxygen
so provided. It has been shown, however, that not only is this effect brought
about by air, but also by the use of such inert gases as nitrogen, carbon dioxide,
and even steam. The following table (Lewes) shows the relative volumes of
various gases that are required to destroy the luminosity of a gas flame :—

I volume of coal gas requires 0.5 volumes of oxygen.

»» » i, 1.26 ,, carbon dioxide.

«• .. >. 2.27 ,, air.

*» »i » 2.30 ,, nitrogen.

i» 1 1 ii 5.11 ,, carbon monoxide.

That the atmospheric oxygen effects the result by a direct oxidising action,
and is not acting merely as nitrogen does, is proved by the fact that mixtures
of oxygen and nitrogen, containing a higher proportion of oxygen than is

* See " Chemical Lecture Experiments," new ed., 498-502. f Ibid., 506.

The Bunsen Flame

343

present in air, destroy the luminosity more rapidly than is effected by air.
Thus, when mixtures containing nitrogen and oxygen in the proportion of 3 to
i, 2 to i, i to i by volume are employed, the volumes of the mixtures required
to destroy the luminosity of one volume of coal gas are respectively 2.02, 1.49,
and i. oo.

It has been shown that when coal gas is diluted with nitrogen a higher
temperature is necessary to effect its decomposition ; hence the action of the
atmospheric nitrogen in causing the loss of luminosity of a gas flame is in part
due to the higher temperature that is required for the formation of acetylene,
which, as already mentioned, is the first step in the decomposition and con-
densation of the hydrocarbons in the gas.

As already mentioned, the luminosity of a flame is very much influenced by
alterations of temperature ; and just as the non-luminosity of the outer mantle
of an ordinary flame is partly due to the cooling action of the air which is
dragged into the flame from the outside, so the want of luminosity of the
Bunsen flame is in part due to the cooling influence of the large volume of air
that is drawn up into the interior of the flame. That the gases which are
drawn into a flame reduce the luminosity by virtue of their cooling action is
borne out by the fact that the higher the specific heat of the diluent (and
therefore the greater its power to abstract heat from the flame) the less of
it is required to effect the destruction of the luminosity ; thus, as already men-
tioned, less carbon dioxide than nitrogen is necessary to render a flame non-
luminous : the specific heat of nitrogen is 0.2370, while that of carbon dioxide
is 0.3307.

The specific heat of oxygen is also slightly greater than that of nitrogen,
being 0.2405; but the cooling effect of dilution with this gas is enormously
overpowered by the increased temperature due to its oxidising action upon
the combustible materials of the flame.

Experiments made upon the actual temperatures of various regions of a
Bunsen flame, rendered non-luminous by admixture with different gases, the
results of which are seen in the following table (Lewes), show the cooling effect
of these diluents upon the flame : —

Temperature of Flame from Bunsen Burner, burning 6 cubic feet of Coal
Gas per Hour.

Flame rendered Non-

luminous by

Region in Flame.

Flame.

Air.

Nitrogen.

Carbon
Dioxide.

Degrees.

\ inch above burner , v

I3S

i£ inch above burner . .
Tip of inner cone

421
913

175
1090

III

444

70
393

Centre of outer cone . " V

1328

J533

999

770

Tip of outer cone

728

"75

«5i

95i

Side of outer cone, level with tip of )
inner cone . ... . }

1236

1333

1236

970

344 Inorganic Chemistry

In the case of air, it will be seen that the first effect is to cool the flame ;
but in the upper region, where the oxidising action of the oxygen is felt, the
temperature rapidly rises to a maximum at a point about half-way between
the tip of the inner and outer cones. In the flames rendered non-luminous
by the two inert gases, the highest temperature is only reached at the outer
limit, where the full amount of oxygen for combustion is obtained from the
outer atmosphere.

On account of the wide range of temperature exhibited by the various
regions of a Bunsen flame, it constitutes a most valuable analytical instru-
ment, for, by the judicious use of the different parts of the flame, it is often
possible to detect the presence of several flame-colouring substances in a
mixture. Thus, if a mixture of sodium and potassium salts be introduced
upon platinum wire into the cooler region of the flame near its base, the more
volatile potassium compound will impart its characteristic violet tint to the
flame before the sodium salt is volatilised sufficiently to mask the colour, by
the strong yellow it itself gives to the flame. In this way many mixtures may
readily be differentiated.

If a piece of copper wire be held horizontally across a Bunsen flame, so as
to cut the inner cone, it will be seen that the wire in contact with the edges
of the flame becomes coated with copper oxide, while the portion in the centre
remains bright. On moving the wire so as to bring the oxidised portion into
the inner region, the oxide will be reduced, the metal once more becoming
bright. The outer area of a flame, where oxygen is in excess, is called the
oxidising flame ; while the inner region, in which heated and unburnt hydro-
gen or hydrocarbons exist, is spoken of as the reducing flame. These regions
exist in all ordinary flames. The oxidising action of the outer flame of a
candle, for example, is illustrated in the behaviour of the wick. So long as
the wick remains in the inner region of the flame it is not burnt ; and in the
early aays of candles, as the tallow gradually consumed, the wick remained
standing straight up, and by degrees extended into the luminous area of the
flame, where, owing to the deposition of soot upon it, it frequently developed
a cauliflower-like accretion, which greatly impaired the luminosity of the
flame, and which necessitated the use of snuffers. In the modern candle,
owing to a method of plaiting the wick, it is caused to bend over (as shown
in Fig. 83), and so thrusts its point into the oxidising region, where it is
continually burnt away.

PART III

THE SYSTEMATIC STUDY OP THE ELEMENTS,
BASED UPON THE PERIODIC CLASSIFICA-
TION.

CHAPTER I

THE ELEMENTS OP GROUP VII. (FAMILY B.)

Fluorine, F . f . ,19.00 I Bromine, Br . . 79.92
Chlorine, Cl . .-'.'" . 35.45 | Iodine, I . . . 126.92

THE first to be discovered, and the most important element of
the group, is chlorine, which is a constituent of sea s'alt (sodium
chloride). The term halogen^ signifying sea salt producer, has
been applied to this family of elements, on account of the close
resemblance between their sodium salts and sea salt. This family
exhibits, in a marked manner, many of the features which are
found to exist in most chemical families of elements.

In their general behaviour they strongly resemble one another,
and readily displace each other in combinations without producing
any very marked change upon the character of the compounds.
They each unite with hydrogen, giving rise respectively to hydro-
fluoric acid, HF ; hydrochloric acid, HC1 ; hydrobromic acid,
HBr ; hydriodic acid, HI.*

These hydrogen compounds are all colourless gases, which fume
strongly in the air; they are extremely soluble in water, and are
strongly acid in character. In combination with potassium and
with sodium, the halogens form a series of compounds, which are
similarly constituted, and which closely resemble each other in their

* Some chemists name these compounds hydrogen fluoride, hydrogen
chloride, hydrogen bromide, hydrogen iodide respectively, and employ the
names hydrofluoric acid, hydrochloric acid, &c. , to denote the aqueous
solutions only.

345

346 Inorganic Chemistry

habits. Their similarity of composition is expressed in the fol-
lowing formulae : —

Compounds with potassium, KF, KC1, KBr, KI.

Compounds with sodium, NaF, NaCl, NaBr, Nal.

The physical properties of the elements exhibit a regular grada-
tion with increasing atomic weight ; thus, fluorine and chlorine are
gases, bromine is liquid, while iodine is solid at ordinary tempera-
tures. In their chemical activity they also show the same gradual
change ; thus, in the case of their combination with hydrogen,
when fluorine and hydrogen are brought together, combination
instantly takes place with explosion, even in the dark. Chlorine
and hydrogen do not combine in the dark, but in diffused daylight
they unite slowly, and in direct sunlight their combination takes
place suddenly with explosion.

Bromine vapour and hydrogen do not combine even in direct
sunlight, but a mixture of the two gases ignites in contact with a
flame, yielding hydrobromic acid, while iodine vapour and hydro-
gen require to be strongly heated in contact with spongy platinum
to effect their combination. This difference in the activity of the
halogens towards hydrogen is seen by a comparison of the heats
of formation of their hydrogen compounds, thus —

H + F =HF + 38,500 cal.

H + C1 = HC1+ 22,000 „

H + Br=HBr+ 8,440 „ *

H + I =HI - 6,040 „ t

Although a strong resemblance exists between all the members
of the halogen family, the element fluorine, which is the typical
member (see page 115), stands marked off from the others in many
of its attributes. Thus fluorine exhibits a great tendency to form
double salts which have no counterpart among the compounds of
the other elements of the family, and at temperatures below 32°
the molecule of hydrofluoric acid consists of two atoms of hydrogen
and two of fluorine, having the composition HgFg.

FLUORINE.

Symbol, F. Atomic weight =19.

History. — This element, the first of the halogen series, was the
most recent to be isolated, it having baffled all attempts to
* This value refers to bromine in the liquid state,
f Iodine as solid.

Fluorine 347

obtain it until the year 1886, when Moissan succeeded in solving
the problem.

Occurrence. — Fluorine occurs in considerable quantities in com-
bination with calcium in the mineral fluor spar (CaF2), which is
found in cubical crystals. On account of the occurrence of this
mineral in large quantities in Derbyshire it is frequently termed
Derbyshire spar. It is a constituent also of cryolite^ Na3AlF6, fltwr-
apatite^ 3P2O8Ca3,CaF2, and many others. In small quantities
fluorine is found in bones, in the enamel of teeth, and also in
certain mineral waters.

Mode of Formation.— When an electric current is passed into
an aqueous solution of hydrochloric acid, the acid is decomposed
into its elements, chlorine being liberated at the positive electrode,
while hydrogen is evolved at the negative. When aqueous hydro-
fluoric acid is treated in the same way, the water only is decom-
posed, oxygen and hydrogen being liberated. Davy found that the
more nearly the acid approached the anhydrous condition, the less
easily did it conduct electricity ; and that in the perfectly pure
state, that is, entirely free from water, hydrofluoric acid was a non-
conductor. Moissan's recent success in the isolation of fluorine
depends upon the discovery that a solution of the acid potassium
fluoride, HF,KF, in anhydrous hydrofluoric acid is an electrolyte)
and that by the passage of an electric current through this solution
fluorine is disengaged at the anode, or positive electrode, and
hydrogen is evolved at the cathode.

The primary products of the electrolysis are potassium (at the
cathode) and fluorine at the anode. The potassium then reacts
with the hydrofluoric acid, re-forming potassium fluoride and
liberating an equivalent of hydrogen —

Or, expressed in the form of ionic equations —

The reaction is performed in a U-tube made of an alloy of
platinum and iridium, a material which is less acted upon by the

348

Inorganic Chemistry

liberated fluorine than platinum alone. The apparatus has two
side-tubes (Fig. 89), which can be either closed with a screw cap, c,
or connected to platinum delivery tubes by means of the union D.
The two limbs of the tube are closed by means of stoppers made
of fluor spar, shown in section at S, and which can be securely
screwed into the tube. These serve to insulate the electrodes,
which are constructed of the same platinum-iridium alloy. The
anhydrous hydrofluoric acid is introduced into the apparatus, and
about 25 per cent, of its weight of the acid potassium fluoride is
added, which readily dissolves in the liquid. The tube is immersed
in a bath of methyl chloride (M, Fig. 90), which boils at - 23° ; the

supply being continuously re-
plenished from the reservoir B,
while the vapour is dravrn away
by the pipe C. On passing a
current from 20 to 25 Grove's
cells through the apparatus,
fluorine is evolved at the posi-
tive electrode, and hydrogen is
liberated at the negative.*

Properties.— Fluorine is, of
all known elements, the most
chemically active. It is on
account of its intense chemical
affinities that it so long resisted
all attempts to isolate it, as
when liberated from combina-
tion it instantly combined with
FlG- 89- the materials of the vessels in

which the reactions were made.

It is impossible to collect this gas by any of the usual methods, for
it decomposes water and instantly combines with mercury. It also
attacks glass, so that it can only be collected by displacement of air
in vessels of platinum. Fluorine is a pale yellowish-coloured gas,
appearing almost colourless when viewed in small quantities. The
smell of the gas is very characteristic — it is irritating to the mucous
membranes, and is not unlike the odour of the mixture of chlorine
and chlorine peroxide, evolved from potassium chlorate and hydro-

* More recently Moissan employs a copper tube of 300 c.c. capacity, fitted
with large platinum electrodes. By keeping the temperature about -50°,
and using a current of 15 amperes, he obtains the gas in large quantities.

Fluorine

349

chloric acid. Whether the smell actually perceived is the true
smell of fluorine is doubtful, for when fluorine comes into contact
with the moisture in the nostrils water is decomposed, with the for-
mation of ozonised oxygen and hydrofluoric acid.

Fluorine not only decomposes potassium iodide, with liberation
of iodine, but also displaces chlorine from sodium chloride.

It combines directly with a large number of elements with
intense energy ; in contact with hydrogen it instantly explodes.
Iodine, sulphur, and phosphorus first melt, and then take fire in
fluorine. Crystals of silicon, when brought into the gas, spontane-

'^iiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiiM

FIG. 90.

ously inflame, and burn with brilliancy. All of the metals are acted
upon by fluorine ; some, when finely divided, undergoing spontane-
ous inflammation when thrown into the gas. Even gold and plati-
num are attacked by fluorine, especially if gently warmed ; its action
upon the latter metal being seen by the corrosion of the apparatus,
and especially the positive electrode employed in its preparation.

Organic compounds are attacked by fluorine with violence, and
often inflamed.

When fluorine is cooled to a temperature about - 187° (*•*. a few

350 Inorganic Chemistry

degrees below the temperature of boiling oxygen, obtained by
boiling the oxygen under slightly reduced pressure) it condenses
to the liquid state.* Liquid' fluorine is a mobile yellow liquid,
resembling liquid chlorine. Its specific gravity is 1.14. It is
without action upon silicon, phosphorus, sulphur, or glass ; it can
therefore be produced and contained in glass vessels. Even at
this low temperature, however, fluorine attacks hydrogen and
hydrocarbons. When cooled by liquid hydrogen it forms a pale
yellow solid, melting at -223°. On cooling the solid to —252° it
loses its yellow colour and appears perfectly white.

HYDROFLUORIC ACID (Hydrogen Fluoride].
Formula, HF. Molecular weight =20.01. Density=io.

Modes of Formation.— ( I.) Hydrofluoric acid is produced when
powdered calcium fluoride (fluor spar) is acted upon by strong
sulphuric acid —

CaF2 + H2SO4 = CaSO4+2HF.

This method is employed for the commercial preparation of
aqueous solutions of hydrofluoric acid. The mixture of fluor spar
and sulphuric acid is gently warmed in a leaden retort, and the
gaseous acid passed into water contained in leaden bottles. This
aqueous acid is sent into the market in gutta-percha bottles.

(2.) The anhydrous acid is prepared by heating hydrogen potas-
sium fluoride (acid potassium fluoride) in a platinum retort. . The
double fluoride of potassium and hydrogen splits up into normal
potassium fluoride and hydrofluoric acid —

HF,KF = KF + HF.

For this purpose the perfectly dry double fluoride is placed in
a platinum retort, which is screwed to a platinum condensing
arrangement, as seen in Fig. 91. The wooden trough through
which the long tube passes is filled with a freezing-mixture, and
the platinum bottle is also surrounded by a similar mixture.

Properties. — Anhydrous hydrofluoric acid is a colourless,
limpid, strongly fuming liquid, which boils at 19.5°. It has a
powerful affinity for water, and can only be preserved in perfectly
stoppered platinum vessels, which are kept in a cool place. The
acid at once attacks gutta-percha. Gore found that the anhydrous
acid was without action upon glass.

Pure hydrofluoric acid is an extremely dangerous substance to
manipulate ; its vapour, even when diluted with air, has a most
* Moissan, May 1897.

Hydrofluoric Acid 351

irritating and injurious effect upon the respiratory organs, and if
inhaled in the pure state causes death.

A single drop of the liquid upon the skin causes the most painful
ulcerated sores, accompanied by distressing aching pains through-
out the whole body. The metals potassium and sodium dissolve in
pure hydrofluoric acid, with the formation of fluorides and evolution
of hydrogen.

At temperatures above 88° the vapour-density of hydrofluoric
acid corresponds to the formula HF. As the temperature is
lowered the molecules aggregate together, and the density of the
vapour steadily rises, until at a few degrees above the boiling-

FIG. 91.

point it approaches what would be required for molecules of H3F3.
At about 32° the density is 20 ; but whether this signifies the exis-
vtence of molecules having the composition H2F2, or whether it
merely represents a certain mixture of more complex molecules,
HnFn, with molecules of HF, has not been definitely determined.

Gaseous hydrofluoric acid rapidly attacks glass, and it is largely
employed for etching purposes, both for obtaining designs upon
glass and for the purpose of etching graduations upon glass mea-
suring instruments. The object to be etched is first coated with
wax, and the design or other marks cut upon the wax by means of
a pointed steel tool. In this way the surface of the glass is laid

352 Inorganic Chemistry

bare in parts, and on exposing the object to the action of the acid.
either as gas or aqueous solution, the glass is rapidly eaten into,
where the surface has been exposed. Its action upon glass is
due to the readiness with which it attacks silicates, the fluorine
combining with the silicon to form silicon tetrafluoride —

Crystallised silicon, when gently heated, takes fire in gaseous
hydrofluoric acid, giving silicon fluoride and hydrogen.

Hydrofluoric acid is extremely soluble in water, forming a
strongly acid corrosive liquid, which readily dissolves many of
the metals with evolution of hydrogen —

' = FeF2+H2.
Silver and copper are also dissolved by this acid.

CHLORINE.

Symbol, Cl. Atomic weight =35. 45. Molecular weight =70. 90.

History. — Chlorine was discovered by Scheele (1774), but was
regarded by him as a compound substance. He applied to it the
name of dephlogisticaled marine acid air, having obtained it by
the action of hydrochloric acid upon ores of manganese. The
belief that chlorine was a compound of oxygen and hydrochloric
acid was generally held until Davy's time, and gave rise to the
name of oxy 'muriatic acid.

The elementary nature of chlorine was proved by Davy (1810),
who gave to it the name chlorine, in allusion to the greenish-yellow
colour of the gas.

Occurrence. — In the uncombined condition chlorine does not
occur in nature. In combination with metals, as chlorides, chlorine
is very abundant, the commonest chloride being sodium chloride
(common salt).

Many of the salts found in the Stassfurt deposits consist largely
of chlorides (see Alkali Metals). Chlorides of the alkali metals
are also found in animal secretions and in certain plants. Chlorine
occurs in combination with hydrogen, as hydrochloric acid, in
volcanic gases, and also in the gastric juice.

Modes Of Formation.— ( i.) When hydrochloric acid is poured
upon manganese dioxide, and the mixture kept cool, a dark-brown

Chlorine

353

solution is obtained, which rapidly decomposes at a slight rise of
temperature with the evolution of chlorine.

It has not yet been clearly established whether this brown solu-
tion consists of the compound MnCl4 or MnCl3, formed according
to one of the equations —

or —

When this dark-brown solution is gently warmed, the higher
chloride breaks up into manganous chloride (MnCl2) and chlorine ;
the complete reaction being expressed by the equation —

MnO2 + 4H Cl = 2H2O + MnCl2 + C12.

The experiment is conveniently carried out in the apparatus seen
in Fig. 92. The mixture of manganese dioxide and hydrochloric

FIG. 92.

acid is gently heated in a large flask, and the gas, after being
passed through water in the Woulf's bottle, may be collected by
downward displacement, as shown in the figure.*

(2.) Instead of employing hydrochloric acid, the materials from
which this compound is prepared, namely, sodium chloride and
sulphuric acid, may be used. Thus, if a mixture of sodium

* See Experiment 154, " Chemical Lecture Experiments," new ed.

354 ' Inorganic Chemistry

chloride, manganese dioxide, and sulphuric acid be gently warmed,
chlorine is readily evolved —

SNaCl + MnO2 + 2H2SO4 = Na2SO4 + MnSO4 + 2H2O + C12.

It will be seen that by this reaction the whole of the chlorine con •
tained ,in the reacting compounds is evolved as gas, while in th<j
former case a part of it remains in combination with the man-'
ganese.

(3.) Many other highly oxygenised compounds, when acted upon
oy hydrochloric acid, evolve chlorine ; thus, when crystals of potas-
sium dichromate are drenched with hydrochloric acid and the
mixture heated, a rapid stream of chlorine takes place, thus —

(4.) When crystals of potassium chlorate are similarly treated, a
mixture of chlorine and chlorine peroxide is evolved, even without
the application of heat—

4KC1O3 + 12H Cl = 4KC1 + 6H2O + 3C1O2 + 9C1.

(5.) Red lead (Pb3O4), when treated with hydrochloric acid,
reacts in a manner similar to manganese dioxide and many other
peroxides. In the case of lead, however, there is no intermediate
chloride formed —

Pb3O4 + 8HC1 = 3PbCl2 + 4H2O + C12.

(6.) Manufacturing Processes — Deacon's Process. — This

method for the preparation of chlorine is by the oxidation of the
hydrogen in hydrochloric acid by atmospheric oxygen. It will be
seen that in the foregoing methods the oxidation of this hydrogen
is carried on at the expense of the oxygen contained in either the
metallic peroxide or the highly oxygenated salt used ; in the
Deacon process atmospheric oxygen is made use of. When a
mixture of gaseous hydrochloric acid and oxygen is heated, a
slight decomposition takes place ; but if these gases be heated in
the presence of a third substance which acts as a catalytic agent,
the decomposition of the hydrochloric acid is much more readily
effected. The catalytic agent employed in the Deacon process is
cuprous chloride (Cu2C\^). This substance is capable of taking up

Chlorine

355

ftn additional quantity of chlorine, and of being converted into
cupric chloride (CuCl2), thus —

Cu2Cl2 + Cl2 = 2CuCl2.

If, therefore, a mixture of hydrochloric acid and oxygen be
passed over fragments of pumice impregnated with cuprous
chloride contained in a tube which is heated to dull redness, the
hydrochloric acid will be decomposed. We may suppose that the
affinity of oxygen for the hydrogen in hydrochloric acid is un-
able to overcome the affinity existing between the hydrogen and
chlorine, but that the additional pull exerted upon the molecules
of hydrochloric acid by the cuprous chloride is sufficient to dis-
turb the equilibrium and rupture the molecule—

0<-H:C1->CuCl

*~H:ci^CU2Cl2'

The result of the action being H2O 4- 2CuCl2.

FIG. 93.

At the temperature at which the reaction is carried on, however,
the compound CuCl2 cannot exist ; two molecules of it are con-
verted into one of Cu2Cl2, and a molecule of chlorine is evolved.
The final result, therefore, of the reaction may be thus expressed —

O + 2HCH- Cu2Cl2=H2O + C12 + Cu2Cl2.

In reality the action is rather more complex, there being an intermediate
compound formed by the combination of cuprous chloride with oxygen. This

356

Inorganic Chemistry

oxychloride of copper then acts upon the hydrochloric acid, as seen in the
following equations : —

(i)

(2)

(3) 2CuCl2=Cu2Cl2+Cl2.

This reaction may be made on a small scale by means of the
apparatus shown in Fig. 93. Hydrochloric acid is generated from
salt and sulphuric acid in the flask, and a stream of the gas passed
through the Woulf's bottle, into which also enters a stream of
oxygen. The mixed gases are then passed through the bulb-tube,
containing fragments of pumice which have previously been soaked

FIG. 94.

in a solution of cupric chloride and dried. On heating the bulb by
means of a Bunsen flame, chlorine will issue from the end of the
tube. When chlorine is manufactured on an industrial scale by
the Deacon process, the mixture of hydrochloric acid and air (in
the proportion of four volumes of the latter to one volume of hydro-
chloric acid) is drawn by means of a Root's blower first through
iron pipes, which are heated to a temperature of about 500°, and
then the hot gases pass on through the decomposer. This consists
of a cylinder of cast iron containing masses of broken brick or
burnt clay impregnated with cupric chloride, and so arranged that
the gases are drawn through the mass.

Chlorine 357

The gas leaving the decomposer consists of a mixture of chlorine,
imdecomposed hydrochloric acid, and atmospheric nitrogen and
oxygen. By passing them through water, the hydrochloric acid is
removed, and the chlorine is usually converted at once into bleach-
ing-powder.

The process by which chlorine is usually made on a manufactur-
ing scale is by the action of hydrochloric acid upon manganese
dioxide. The best ore for the purpose is pyrolusite. The process
is conducted in stills made of thick slabs of stone, usually " York-
shire flag," which are fitted and luted together, and securely bound
by cast-iron clamps. Fig. 94 shows such a chlorine still, repre-
sented as cut across the centre.

The charge of manganese is placed upon the false bottom a, and
the acid is run in through the funnel tube £, which, dipping into a
small pot, does not allow the gas to escape. As the action begins
to slacken, steam is cautiously blown in from time to time. The
chlorine escapes by the pipe g> and passes from thence into the
main h.

The reaction that goes on in the still is the same as that given
above in the first mode of formation, except that as pyrolusite is
not pure MnO2, small quantities of other compounds are formed.
The following analysis, by Black, of still-liquor from stone stills,
shows the general composition of this substance : —

MnCl2 . , * . . ". . 10.5700

A12C16 . V . . . *; 0.6200

Fe2Cl6 . _.'.;. . : v ,.. . 0.4551

HC1 (undecomposed) . V . 6.6220

H2O . . . .' . • . 81.7329

100.0000

(7.) The Weldon Process, although indirectly a method for
making chlorine, is in reality a process for recovering the man-
ganese contained in the still-liquors as manganous chloride, and of
reconverting it into available manganese dioxide. The manganese
so recovered, however, is again utilised for the preparation of
chlorine by the decomposition of a further quantity of hydrochloric
acid. The essence of the process is the following : — The still-
liquor is mixed with ground chalk, or limestone dust, in large tanks
or wells, and the mixture thoroughly stirred by agitators. One of
these wells, A, is shown in the diagrammatic figure. By this opera-

3S8

tnvrganic Chemistry

tion the free acid is neutralised, and the iron precipitated as
hydrated oxide. The neutral liquor, consisting of manganous
chloride and calcium chloride, is then pumped into large tanks,
where it is allowed to settle ; one of these " settlers," B, is shown
in the figure. By means of a pipe upon a swivel-joint,^ the clear
liquid from the settler can be drawn off without disturbing the
sediment, and run into the oxidiser C. The oxidiser is merely a
flat-bottomed iron cylinder, open at the top. Milk of lime from

0 D

the tank E, where lime and water are stirred together, is pumped
into the oxidiser as required.

The milk of lime is added in quantity more than sufficient to
precipitate the manganese as manganous hydroxide, MnH2O2.
Into this mixture, which consists of manganous hydroxide and
calcium hydroxide (milk of lime) in suspension, and to a smaller
extent in solution in the calcium chloride which is present, a
stream of compressed air is forced by means of the pipe A, which

Chlorine 359

passes to the bottom of the oxidiser, where it ends in perforated
branches. During this process the manganese becomes oxidised
and is converted mainly into calcium manganite, a compound of
manganese dioxide with calcium oxide, CaO,MnO2, or CaMnO3.
By a further addition of the neutral liquor from tank B, and by
raising the temperature within the oxidiser by injecting steam, a
portion of the calcium manganite is converted into a compound
having the composition CaO,2MnO2.

When the operation is complete, the contents of the oxidiser are
run out into a series of tanks called mud settlers, of which one
is shown at D in the figure. The product here settles as a thin
black mud, known as the Weldon mudj and this is ultimately
drawn from the settlers, and run direct into chlorine stills, where
it is at once treated with hydrochloric acid for the preparation
of chlorine. The Weldon stills are similar to the ordinary chlorine
stills, but are much larger, and usually octagonal in shape.

(8.) Electrolytic Methods.— Of late years, since the applica-
tion of electricity on a commercial scale has become possible,
manufacturing processes for obtaining chlorine by the electrolysis
of a solution of common salt have begun to compete with the
older methods. By the electrolysis of brine, the sodium chloride
is separated into its two elements ; the chlorine is evolved at the
anode, and the sodium which is liberated at the cathode there acts
upon the water present, generating sodium hydroxide (see Caustic
Soda ; also Sodium Carbonate).

Properties. — Chlorine is a greenish-yellow coloured gas, with a
strong suffocating smell. It is quite irrespirable, and if inhaled in
the pure state causes death. Even when largely diluted with air
it is extremely disagreeable and injurious, as it acts rapidly upon
the mucous membranes of the nose and throat, causing*irritation
and inflammation, which usually result in severe catarrh. A few
bubbles of chlorine allowed to escape and diffuse into the air of a
room give to the air a distinct and rather pleasant smell. Chlorine

is an extremely heavy gas, being *fe» = 2.45 times heavier than

14.44
air. One litre of the gas, measured under the standard conditions

of temperature and pressure, weighs 3.168 grammes. The density
of chlorine, taken at all temperatures, does not exactly agree with
that which is required for the molecular formula C12. At tempe-
ratures above 1200° the density is markedly less than theory
demands, showing that partial dissociation of the chlorine mole-

360 inorganic Chemistry

cules into single atoms has taken place. (Compare Bromine and
Iodine.)

On account of its heaviness, chlorine. is readily collected by dis-
placement ; it cannot be collected over mercury, as it attacks that
metal, and in water it is considerably soluble. It may, however,
be collected over a strong brine, as it is much less soluble in this
solution than in water.

Chlorine is not inflammable, but it supports the combustion of
many burning bodies. It is possessed of such extremely powerful
chemical affinities that it acts upon a large number of substances
at ordinary temperatures, and in many cases the combination is
sufficiently energetic to result in the inflammation of the bodies.
Phosphorus, when introduced into chlorine, first melts and then
spontaneously inflames, burning with a somewhat feeble light to
form phosphorus trichloride (PC13) and phosphorus pentachloride
(PC16). The elements arsenic and antimony, when finely powdered
and dusted into a vessel of chlorine at once take fire and burn,
forming their respective chlorides. Many metals, when finely
divided, or in the form of thin leaf, such as ordinary Dutch metal,
instantly take fire when brought into chlorine. If a quantity of
sodium be heated in a deflagrating spoon until it begins to burn in
the air, and be then plunged into chlorine, the sodium continues to
burn in the gas with dazzling brilliancy, forming sodium chloride.

Although under ordinary circumstances chlorine unites with
metals with great readiness, it has been shown that this action will
not take place if the chlorine be absolutely dry. Thus, if chlorine
which has been completely freed from aqueous vapour be passed
into a tube containing bright metallic sodium, and the tube sealed,
the sodium not only remains bright and unaffected by the gas, but
may even be melted in the atmosphere of chlorine without any
action taking place. Similarly, dry chlorine, when allowed to enter
a flask filled with Dutch metal, has no action upon it ; but upon the
introduction of the smallest trace of moisture the metal at once
takes fire.* These facts are of the same order as those mentioned
in connection with oxygen (see page 191).

Chlorine is not capable of direct combination with carbon ; ordi-
nary combustibles, therefore, which consist of hydrocarbons, burn
in chlorine by virtue of the combination of their hydrogen with the
gas, and they burn with a lurid smoky flame, owing to the elimina-
tion of their carbon in the form of soot. A burning taper 01'

* See Experiments 174, 175, " Chemical Lecture Experiments," new. ed.

Chlorine 361

ordinary gas flame when introduced into chlorine burns in this
manner, emitting a dense smoke and forming fumes of hydro-
chloric acid.

Chlorine has a most powerful affinity for hydrogen ; a jet of
hydrogen burns freely in chlorine, with the formation of hydro-
chloric acid. A mixture of hydrogen and chlorine unites with
explosion on the application of a flame. This combination takes
place also under the influence of light (see Hydrochloric Acid).
The affinity shown by chlorine for hydrogen is seen in its action
upon many of the compounds of hydrogen and carbon. If one
volume of ethylene (C2H4) be mixed with two volumes of chlorine,
and the mixture ignited, the carbon is instantly thrown out of com-
bination as a black smoke, while the hydrogen unites with the
chlorine, forming a cloud of hydrochloric acid. Similarly, if a
liquid hydrocarbon, such as tuipentine (C10H16), be poured upon a
piece of filter paper, and the paper be thrust into a jar of chlorine,
instant inflammation takes place, with deposition of a large quantity
of carbon.

Chlorine possesses strong bleaching properties, which depend
upon its power of combining with hydrogen, for it is an essential
condition that water shall be present. The chlorine unites with
the hydrogen of the water, and the liberated oxygen oxidises the
colouring matter. If chlorine be bubbled into liquids coloured
with any vegetable colouring matter, or if a dyed rag be dipped into
chlorine water, the colour will be rapidly discharged. Ordinary
writing-ink (which usually consists of a compound of iron with
tannic and gallic acids) is readily bleached by chlorine ; while
printer's ink, which consists mainly of carbon, in the form of lamp-
black, is not acted upon by this gas. If, therefore, a piece of printed
paper be brushed over with writing-ink so as to completely obli-
terate the print, and the blackened paper be immersed in chlorine
water, the writing-ink will be rapidly bleached away, leaving the
print unchanged.

The bleaching power of chlorine constitutes its most valuable
property from an industrial point of view ; the chlorine for this
purpose is combined with lime to form the substance known as
bleaching-powder (see Calcium Compounds).

Chlorine is soluble in water to a considerable extent. One
volume of water at 10° absorbs 3.0361 volumes of chlorine
measured at o° and under 760 mm. pressure. This solution,
known as chlorine water, has the same colour as the gas, and

362 Inorganic Chemistry

smells strongly of chlorine. If exposed to the air, the chlorine
rapidly diffuses out of the solution. Chlorine water cannot be
preserved for any length of time, as it slowly undergoes de-
composition, the chlorine combining with the hydrogen of the
water, forming hydrochloric acid, which remains in solution, and
the oxygen being liberated, thus —

H2O + C12=2HC1 + O.

This action, which proceeds slowly under ordinary conditions,
is greatly accelerated by the influence of light, and if exposed
to direct sunlight the decomposition is very rapid.

If chlorine water be cooled to within one or two degrees of the
freezing-point of water, or if chlorine be passed into ice-cold water,
a solid crystalline compound of chlorine with water is deposited.
This substance is termed chlorine hydrate, and has a composition
expressed by the formula C12,10H2O. The compound is very un-
stable, and when exposed to the air it melts and rapidly gives otf
chlorine. If the crystals are quickly freed from adhering water,
and are then sealed up in a glass tube, they may be heated to
a temperature of 38° before being decomposed. Faraday made
use of this compound in order to obtain liquefied chlorine. A
quantity of the hydrate was sealed up in one limb of a bent
tube and was gently warmed, the compound dissociated into
water and chlorine, and the internal pressure caused the condensa-
tion of the chlorine to the liquid condition.

Liquid Chlorine. — Under the ordinary atmospheric pressure,
chlorine may be liquefied by lowering its temperature to — 34°.

At a temperature of o° the pressure required to effect its lique-
faction is equal to six atmospheres. When, therefore, the liquid
is obtained by heating the crystalline hydrate, as in Faraday's
method, one limb of the tube should be cooled by being placed
in ice.

The critical temperature of chlorine is 141°, and the pressure
required to effect its liquefaction at that point, or its critical
pressure, is 84 atmospheres (see Liquefaction of Gases).

Liquid chlorine has a bright golden-yellow colour, entirely free
from the greenish tint possessed by the gas. Its specific gravity
is 1.33, and it boils at —33.7°. When cooled to a temperature of
- 101.5°, it freezes to a yellow crystalline mass. Liquid chlorine
is now an article of commerce. It is contained in iron bottles
lined with lead, and is largely exported in this form, for use in the

Hydrochloric Acid

363

extraction of gold, to parts of the world where the carriage of the
plant and materials necessary for generating large quantities of
chlorine would be attended with great difficulties.

HYDEOCHLORIC ACID (Hydrogen Chloride).
Formula, HC1. Molecular weight =36. 46. Density =18.23.

History.— In solution in water this compound was known to
the early alchemists, and the mixture of this solution with nitric
acid constituted the valued liquid known as aqua regia. The

FIG. 96,

preparation of hydrochloric acid from common salt is associated
with the name of Glauber (1650), who obtained it by the action
of sulphuric acid upon sodium chloride (common salt). Gaseous
hydrochloric acid was first collected and examined by Priestley,
who collected it over mercury in the mercurial pneumatic trough
invented by him. He named the gas marine acid air.

Occurrence.— Gaseous hydrochloric acid is evolved in consider-
able quantities from volcanoes during active eruption.

ModGS of Population. — (i.) Hydrogen chloride may be syn-
thetically produced directly from its elements ; thus, this compound
is formed when a jet of hydrogen is caused to burn in an atmos-

364 Inorganic Chemistry

phere of chlorine. If a mixture of chlorine and hydrogen IDC
ignited, the union takes place instantaneously with explosion,
and hydrogen chloride is produced. The union of hydrogen with
chlorine will also take place under the influence of light ; thus, if a
mixture of these two gases be exposed to even diffused daylight for
a few hours the greenish colour imparted to the mixture by the
chlorine will gradually disappear, and on examination it is found
that the tube contains hydrogen chloride. This combination, which
is only gradual when the mixture is exposed to diffused
daylight, becomes explosively sudden if the mixed gases
are exposed to direct sunlight, or any artificial light which
is rich in rays of high refrangibility — the so-called actinic
rays. If, therefore, a glass vessel be filled with a mixture
of these gases in equal volumes, and the mixture be placed
in bright sunshine, a violent explosion will result, and
hydrogen chloride will be produced. This phenomenon
is best illustrated by filling small thin glass bulbs with a
mixture of the two gases obtained by the electrolysis of
aqueous hydrochloric acid. The bulbs when filled can be
hermetically sealed before the blowpipe without causing
the combination of the gases,* and if kept in the dark may
be preserved indefinitely.

On exposing one of these bulbs to the light of burning
magnesium the combination of the two gases instantly
takes place, with a sharp explosion which shatters the bulb
to powder. The bulb should therefore be screened, as
shown in Fig. 96.

The rays of light which are capable of causing this
combination are those which compose the blue and violet
end of the spectrum ; if these particular rays are absorbed
from the light by means of ruby glass, the mixture of
FIG. 97. gases may be exposed to the red light so obtained with-
out any action taking place.t

The combination of chlorine with hydrogen is not attended by
any alteration in volume ; one volume of chlorine combines with one
volume of hydrogen, and the resultant hydrogen chloride occupies
two volumes. This may be readily proved by filling a stout glass
tube, provided with a stop-cock at each end, with a mixture of the two
gases in exactly equal volumes, and causing them to combine either

* See Experiment 178, " Chemical Lecture Experiments," new ed.
f Ibid., p. 93.

Hydrochloric Acid

365

by the influence of light or by the passage of an electric spark by
means of the platinum wires sealed into the tube (Fig. 97). On
opening one of the stop-cocks under mercury it will be seen that
no mercury is drawn in, neither does any gas pass out from the
tube, thus showing that the union has taken place without any
alteration in the volume. If one of the cocks be now opened
beneath water, the hydrogen chloride which has resulted from
the union of the hydrogen and chlorine, being extremely soluble
in water, the liquid will rush up into the tube and completely fill
it, showing that no free hydrogen or chlorine remains in the tube.

FIG. 98.

(2.) For all ordinary purposes, hydrogen chloride is always
obtained by the action of sulphuric acid upon sodium chloride.
For laboratory uses the apparatus seen in Fig. 98 may be con-
veniently employed. Sulphuric acid, previously diluted with rather
less than its own volume of water, is placed in the flask, and a
quantity of common salt is added. On the application of a gentle
heat a steady stream of gas is evolved, which may be dried by
being passed through the tubulated bottle, containing pumice
moistened with strong sulphuric acid. The gas is then collected
either over mercury or by displacement. The reaction which
takes place is expressed by the equation —

NaCl + H2SO4=NaHSO4 + HCl.

If strong sulphuric acid be employed along with an excess of
salt, both of the atoms of hydrogen can be displaced from -the

366 Inorganic Chemistry

acid ; and instead of the hydrogen sodium sulphate there is
formed the normal sodium sulphate —

2NaCl + H2SO4= Na2SO4 + 2HC1.

A much higher temperature is necessary in order to complete the
reaction indicated by this' equation.

Properties. — Hydrogen chloride is a colourless gas with a
choking, pungent odour. In contact with the moist air it forms
dense fumes, consisting of minute globules of a solution of the
gas in the atmospheric aqueous vapour. The gas does not burn,
neither does it support ordinary combustion.

It is heavier than air, its specific gravity being —

14.44

Hence the gas is readily collected by displacement. One litre of
the gas weighs 18.185 criths.

Hydrogen chloride is extremely soluble in water ; I volume of
water at o° and under a pressure of 760 mm. is capable of dis-
solving 503 volumes of gaseous hydrochloric acid, measured at
o° and 760 mm. As the temperature rises the solubility diminishes,
as seen by the following table : —

Temperature. Coefficient of Absorption.

o° ....... 5°3

30° . . 411

50° . 364

The solubility of hydrogen chloride may be illustrated by com-
pletely filling a large globular flask with the gas, by displacement,
the flask being provided with a long tube passing through the cork,
as seen in Fig. 99. On opening this tube beneath water, the gas
begins to dissolve, and the liquid rises slowly in the tube until it
reaches the top. As soon as the first few drops enter'the globe,
they rapidly absorb the gas, thereby causing a partial vacuum in
the vessel, so that the water is driven up the tube with consider-
able force, forming a fountain, which continues until the globe is
nearly filled with liquid. If the water in the dish is rendered blue
by the addition of litmus solution, the acid nature of the solution
of the gas will be evident by the reddening of the liquid as it
enters the globe,

Hydrochloric Acid

367

When a weak aqueous solution of hydrochloric acid is boiled
it loses water and becomes stronger ; while, on the other hand,
if a strong solution be heated, it loses gas and becomes weaker,
until in both cases an acid containing 20.24 per cent, of HC1 is
produced which boils at 1 10°. This strength of acid corresponds
to a composition expressed by the formula HC1 + 8H2O, and it
was at one time supposed to represent a definite compound.
Roscoe and Dittmar have shown, however, that, as with nitric
acid, the composition of the liquid which
boils at a constant temperature is simply a
function of the pressure. (Compare Nitric
Acid, page 239.)

The strongest aqueous solution of hydro-
chloric acid at 1 5° C. has a specific gravity
of i. 212, and contains 42.9 per cent, of HC1.

Hydrogen chloride gas is readily liquefied
by pressure. At 10° a pressure of 40 atmos-
pheres will effect its liquefaction, while at
— 16° the same result is obtained by a
pressure of 20 atmospheres. The critical
temperature of the gas is 52.3°.

The liquefied gas boils at —83.7°. At
low temperatures it freezes to a white solid
which melts at — 1 10°. The liquid is with-
out action on most of the metals, which are
readily dissolved by the aqueous acid.

The composition of hydrogen chloride may be experimentally
proved by a number of methods. It may be shown synthetically
by the volumetric experiment referred to above (page 364).

The volumetric proportion of hydrogen contained in the gas may
be shown by means of sodium amalgam. The sodium in the
amalgam reacts with the gas, combining with the chlorine, and
liberating the hydrogen —

FIG. 99.

For this purpose gaseous hydrochloric acid is introduced into one
limb of the U-shaped eudiometer (Fig. 100), and its volume indicated
by means of a ring upon the tube, the mercury being level in both
limbs. A second ring marks exactly half the volume. A quantity
of liquid sodium amalgam is then poured into the open limb until
it is completely filled, and on being closed by the thumb the tube

368

Inorganic Chemistry

can be inverted so as to decant the gas into this limb. After being
bubbled once or twice through the amalgam, the gas Is again
returned to its former place ; and by drawing mercury from the
branch tube, the levels in each limb can be again adjusted, when it
will be found that the gas remaining in the tube occupies the space
exactly down to the upper ring, that is to say, two volumes of
hydrochloric acid contain one volume of hydrogen. That the gas

FIG. 100.

FIG. 101.

FIG 102.

is hydrogen can be shown by again filling up the open limb with
mercury, and driving the gas out of the stop-cock, where it can be
inflamed as it escapes.

The fact that hydrogen chloride contains the same volume of
chlorine as of hydrogen may also be demonstrated by collecting
the mixed gases, evolved by the electrolysis of the aqueous acid, in
a long tube provided with a stoppered funnel, as shown in Fig. 101.

Hydrochloric Acid 369

The gases may be collected over a saturated solution of salt in
water, and the tube filled to the lower ring. On allowing a solution
of potassium iodide to enter by means of the funnel, the chlorine is
absorbed with the liberation of iodine, which partially dissolves
and partly separates as a solid. When the absorption of the
chlorine is complete, the water will have risen to the second band
placed half-way up the tube, showing that one-half of the gaseous
mixture consists of chlorine. The former experiment proved that
hydrochloric acid contained half its volume of hydrogen, therefore
the two elements, in uniting to form this compound, do so in equal
volumes and without any contraction in volume.

When aqueous hydrochloric acid is subjected to electrolysis, the
hydrochloric acid is decomposed, hydrogen being evolved at the
negative electrode and chlorine at the positive. At first the
liberated chlorine is dissolved in the solution, but after the liquid
has become saturated with the gas, the whole of the chlorine is
liberated. By conducting this decomposition in the apparatus
seen in Fig. 102, and continuing the passage of the electric current
until the liquid in one limb is saturated with chlorine before closing
the stop-cocks, it will be seen, when the gases are collected in the
tubes, that they are evolved in equal volumes.

The Manufacture of Hydrochloric Acid. — The aqueous
solution of hydrochloric acid is an object of commercial manu-
facture, which is carried out on an enormous scale. It is obtained
by the decomposition of common salt by means of sulphuric acid,
according to the reaction —

2NaCl + H2SO4= Na2SO4 + 2HC1.

Formerly hydrochloric acid was a waste product obtained in the manufacture
of sodium carbonate by the method known as the Leblanc process, the first
stage in this process being the conversion of sodium chloride into sodium
sulphate by the action upon it of sulphuric acid. The hydrochloric acid
evolved as gas in this reaction was allowed to escape into the atmosphere.
The nuisance caused by this acid gas being thrown into the air, ultimately
resulted in the "Alkali Act," which compelled manufacturers to absorb this
waste acid. Since that time the Leblanc process for the manufacture of sodium
carbonate has had a formidable rival in another method, known as the
ammonia-soda process (see Sodium Compounds), which would probably have
completely driven the older method out of the field, but for the commercial
value of the hydrochloric acid which is obtained as a secondary product in
the Leblanc process. The hydrochloric acid, therefore, which formerly was
thrown away as a waste product, is now the salvation of the process, and the
utmost care is taken to prevent any of it from escaping, not now by com-
pulsion of the Alkali Act so much as from purely economic reasons.

2 A

370

Inorganic Chemistry

The charge of salt and sulphuric acid is heated in an enormous
hemispherical cast-iron pan, built into a brickwork chamber, so
that it can be heated by a fire beneath, and so that the evolved
gas can be conveyed away by brick or earthenware flues. The gas
evolved by the reaction is led into towers which are filled with coke
or bricks, and down which water is made to percolate, the water
being caused to flow equally over the mass by means of special
distributing contrivances. As the gaseous hydrochloric acid passes
up the towers and meets the descending stream of water it is en-
tirely dissolved, and the aqueous acid becomes nearly saturated as
it reaches the bottom of the tower.

In works where the condensers or towers are not of great

FIG. 103.

height, it is usual either to cool the gas before admitting it into
the towers, or to pass it through a series of jars resembling gigantic
Woulf s bottles (Fig. 103).

The water in these bottles is made to flow steadily from one to
the other by the side pipes c, c (in the direction from left to right),
while the gas passes through the system in the opposite direction.
In this way a constantly changing surface of water is exposed to
the gas, and a very strong solution is obtained.

Commercial hydrochloric acid is generally yellow in colour,
owing to the presence of iron as an impurity ; and it is always
liable to contain sulphuric acid, free chlorine, arsenic, and some-

Chlorine Monoxide 37*

times sulphur dioxide. This aqueous solution of hydrochloric
acid is also known under the names of "spirits of salt" and
muriatic acid.

OXIDES AND OXYACIDS OF CHLORINE.

The elements oxygen and chlorine have never been made to
unite together directly : three compounds, however, of these ele-
ments can be obtained by indirect methods ; these are —

Chlorine monoxide (hypochlorous anhydride) . C12O.
Chlorine peroxide . . . ... . . C1O2.

Chlorine heptoxide ...... C12O7.

Three oxyacids are known, viz. : —

Hypochlorous acid Jr- . '. . . . HC1O.
Chloric acid . . v. : ; . . :J . HC1O3.
Perchloric acid . .' .. ; . . . HC1O4.

CHLORINE MONOXIDE (Hypochlorous Anhydride].
Formula, C12O. Molecular weight =86. 90. Density =43. 45.

Mode Of Formation. — This compound is obtained by passing
dry chlorine over dry precipitated mercuric oxide contained in a
glass tube, the temperature of which is not allowed to rise. The
chlorine combines with the mercuric oxide, forming mercuric oxy-
chloride, and chlorine monoxide is liberated —

2HgO + 2C12 = HgO,HgCl2+ C12O.

Properties. — At ordinary temperatures chlorine monoxide is a
pale yellow gas, without the greenish tint possessed by chlorine.
Its smell strongly suggests chlorine, but is readily distinguishable
from it. It is a very unstable compound, decomposing with more
or less violence with moderate rise of temperature. When strongly
cooled it is condensed to an orange-yellow coloured liquid, which
boils at about —20°. This liquid is extremely unstable, exploding
with great violence on the gentlest application of heat, and some-
times on merely being poured from one vessel to another. When
exposed to direct sunlight it also explodes with violence.

Gaseous chlorine monoxide is considerably soluble in water, one

372 Inorganic Chemistry

volume dissolving about 100 volumes of the gas, forming hypo-
chlorous acid —

HO = 2HC1O.

CHLORINE PEROXIDE.

Formula, C1O2. Molecular weight =67. 45. Den 5^7=33.72.

Modes of Formation.— ( i.) By the action of sulphuric acid
upon potassium chlorate —

3KC1O3 + 2H2SO4=KC1O4 + 2HKSO4 + H2O + 2C1O2.

Finely powdered potassium chlorate is added little by little to
concentrated sulphuric acid in a small retort. The salt dissolves
with the formation of a reddish liquid, and if the temperature is
not allowed to rise, no gas is evolved. On very cautiously warm-
ing the retort by means of warm water, taking care not to heat
the glass above the level of the liquid in the retort, the chlorine
peroxide is evolved.

(2.) A mixture of chlorine peroxide and carbon dioxide, in equal
volumes, is obtained by heating a mixture of powdered potassium
chlorate and oxalic acid to a temperature of 70° in a water-bath —

2KC1O3 + 2H2C2O4= K2C2O4 + 2H2O + 2CO2 + 2C1O2.

(3.) Chlorine peroxide, mixed with chlorine, is evolved by the
action of hydrochloric acid upon potassium chlorate —

This mixture of gases was formerly supposed to be a definite
compound of oxygen and chlorine, and received the name of
euchlorine.

Properties. — Chlorine peroxide is a heavy gas, with a deep
yellow colour. It has an intensely unpleasant smell, and if in-
haled, even when largely diluted with air, produces headache.
The gas attacks mercury, and is soluble in water, so that it can
only be collected by displacement. Chlorine peroxide is an ex-
tremely unstable compound, it is gradually resolved into its ele-
ments by the influence of light ; the passage of an electric spark,
or the introduction into it of a hot wire, causes it to decompose
with violent explosion. It is a powerful oxidising compound ; a

Hypochlorous Acid 373

piece of phosphorus introduced into the gas takes fire spontane-
ously. If a jet of sulphuretted hydrogen be lowered into a jar of
chlorine peroxide, the sulphuretted hydrogen ignites spontaneously
and continues burning in the gas.

Its oxidising action upon organic matter may be shown by
liberating the gas in the presence of such a substance as sugar,
by adding a drop of sulphuric acid to a mixture of powdered sugar
and potassium chlorate. The chlorine peroxide, liberated by the
action of the acid upon the chlorate, ignites the mixture, and the
entire mass then bursts into flame.

When chlorine peroxide is strongly cooled it condenses to a
dark red liquid, which is even more explosive than the gas.

Chlorine Heptoxide, Cl.,O7.— This compound is obtained by
the cautiously regulated action of phosphoric oxide upon perchloric
acid,* whereby the elements of water are withdrawn from two
molecules of the acid —

2HC1O4-H2O = C12O7.

The operation is attended with some danger, although the heptoxide
when isolated is described as less unstable than either of the othei
oxides.

HYPOCHLOROUS ACID.

Formula, HC1O.

Modes Of Formation.— ( i.) As already mentioned, this acid is
formed when chlorine monoxide is dissolved in water.

(2.) It may readily be obtained in dilute solution by passing an
excess of chlorine through water in which precipitated mercuric
oxide is suspended —

HgO + H2O + 2Cl2=HgCl2 + 2HClO.

On distilling the liquid, th,e dilute acid passes over as a colourless
distillate.

(3.) In dilute solution, hypochlorous acid may be obtained by the
decomposition of a hypochlorite by a very dilute mineral acid, and
subsequent distillation of the mixture ; thus, if to a solution of
calcium hypochlorite (obtained by treating bleaching-powder with
water and filtering the solution) very dilute nitric acid be added
and the solution distilled, a dilute colourless acid is obtained —

Ca(ClO)2 + 2HNO3 = Ca(NO3)2 + 2HClO.

(4.) This compound is also formed when a stream of chlorine is
* Michael and Conn., Am. Chem, Journ., 1900.

374 Inorganic Chemistry

passed through water containing precipitated calcium carbonate in
suspension —

Properties.— Pure hypochlorous acid, free from water, has
never been obtained. The acid produced by the solution in water
of chlorine monoxide has a pale straw-yellow colour, and a very
characteristic chlorous smell. Dilute solutions of this acid are
moderately stable, while more concentrated solutions readily
undergo spontaneous decomposition.

Hypochlorous acid is a powerful oxidising and bleaching agent,
as it readily gives up its oxygen, and is resolved into hydrochloric
acid —

As an oxidising agent it is twice as effective as an equivalent
quantity of chlorine in chlorine water, for two atoms of chlorine are
here necessary for the liberation of one atom of oxygen —

C12+H20 = 2HC1 + 0.

Hypochlorous acid is decomposed by hydrochloric acid with
the evolution of chlorine —

It is also decomposed by silver oxide, oxygen being liberated —

The salts of hypochlorous acid may be obtained by the action of
the acid upon the hydroxides of the metals ; thus —

HC1O + KHO = KC1O + H2O.

The most important salt of this acid is bleaching-powder (see
Calcium Salts).

CHLORIC ACID.

Formula, HC1O3.

Mode Of Formation.— This compound is best obtained by
decomposing barium chlorate with an exact equivalent of sulphuric
acid, previously diluted with water —

Ba(ClO3)2 + H2SO4=BaSO4 + 2HClO3.

The clear liquid is decanted from the precipitated barium sul-
phate, and is then concentrated by evaporation in vacuo.

Perchloric Acid 375

The strongest acid that can be obtained still contains 80 per
cent, of water. Attempts to concentrate it further result in its
decomposition into free chlorine and oxygen, with the formation
of perchloric acid and water.

Properties. — The strong aqueous acid has powerful oxidising
properties ; many organic substances, as wood or paper, are so
rapidly oxidised by it that when the acid is dropped upon them
they are frequently inflamed.

The acid even in dilute solution has strong bleaching powers.

The salts of chloric acid are far more stable than the acid, and
some of them are of considerable technical importance. The
chlorates are all soluble in water, and all yield oxygen on being
heated. Chloric acid is a monobasic acid ; the chlorates, there-
fore, have the general formula M'C1O3 and M"(C1O3)2, where M'
and M" stand for monovalent and divalent metals respectively.

Of all the chlorates, potassium chlorate, KC1O3, is by far the
most important (see Potassium Compounds).

PERCHLORIC ACID.

Formula, HC1O4.

Mode Of Formation.— Perchloric acid is best prepared by the
Action of strong sulphuric acid upon potassium perchlorate—

2KC1O4+H2SO4 = K2SO4 + 2HC1O4.

Pure and dry potassium perchlorate is mixed with four times its
weight of concentrated sulphuric acid, and the mixture gently dis-
tilled in a small retort. The distillate at first consists of perchloric
acid ; but as the operation proceeds a portion of the perchloric acid
is decomposed into lower oxides of chlorine and water, and the
latter, combining with the first portions of the distillate, forms a
white crystalline compound, having the composition HC1O4H2O.
This body, when gently heated, gives off perchloric acid ; it may,
therefore, be employed for the preparation of the acid in a state of
purity.

Properties.— Perchloric acid is a colourless, volatile, and strongly
fuming liquid, having a specific gravity of 1.782 at 15°. It .is an
extremely powerful oxidising substance ; a drop of the liquid
allowed to fall upon paper, wood, or charcoal is instantly decom-
posed, sometimes with a violent explosion. In contact with the

376 Inorganic Chemistry

skin it produces most painful wounds ; when allowed to drop into
water it produces a hissing sound, owing to the energy of the
combination.

The salts of this acid are the perchlorates, of which the most
important is potassium perchlorate ; they are all soluble in
water.

Constitution of the Oxides and Oxyacids of Chlorine.— On the assump-
tion that chlorine is a monovalent element, the constitution of these compounds
may be thus represented : —

Chlorine monoxide, Cl - O - Cl. I Hypochlorous acid, Cl ^ O - H.

Chlorine peroxide, Cl - O - O - . I Chloric acid, Cl-O-O-O-H.

Perchloric acid, Cl-O-O-O-O-H.

It is possible, however, that in some of these compounds the chlorine
functions as a trivalent element, and that these compounds have a constitution
similar to the oxides and oxyacids of nitrogen, thus : —

Chlorine monoxide, Cl - O - Cl. Nitrogen monoxide, N — O - N.

/° /°

Chlorine peroxide, - Cl<( | . Nitrogen peroxide, — NC | .

\O \O

Hypochlorous acid, Cl— O — H. Hyponitrous acid, N — O — H.

/O /O

Chloric acid, H-O-C1<( J . Nitric acid, H - O - N< | .

\0 \O

Perchloric acid, H-O-C1

There are several facts which point to the belief that not only chlorine, but
also bromine and iodine, are capable of fulfilling the functions of a trivalent
element. The existence, for example, of such a compound as trichloride of
iodine, IC13, is difficult to explain on any other assumption than that iodine is
here a trivalent element.

Indeed, from a consideration of the salts of periodic acids, some chemists are
in favour of assigning to iodine even a still higher valency, and of regarding it
as a heptad element in these compounds (see Periodates, page 394). The
constitution of such molecules as those of hydrofluoric acid at low temperatures,
namely, H2F2, or H3F3, or HnFn, and of the acid fluoride of potassium,
HF.KF, is readily understood if we regard the fluorine as functioning in
these compounds as a trivalent element, thus —

H-F=F-H; H-F — F-H: H-F=F-K.

Bromine 377

BROMINE.

Symbol, Br. Atomic weight = 79. 92. Molecular weight = 159. 84.
Vapour density = 79. 92.

History. — This element was discovered by Balard (1826), in the
mother-liquor obtained after the crystallisation of salt from con-
centrated sea-water. He applied the name bromine (signifying a
stench") to the element, in allusion to its unpleasant smell.

Occurrence. — Bromine is never found in the uncombined state
in nature. In combination chiefly with the metals potassium,
sodium, and magnesium, it occurs in small quantities in all sea-
water, and more abundantly in many mineral waters and salt
springs. The saline deposits of Stassfurt contain notable quantities
of bromides, and the main supply of bromine for the market is
manufactured from this source.

Modes of Formation. — (i.) Bromine may be obtained from a
bromide by displacement with chlorine. If to a solution of mag-
nesium bromide, chlorine water is added, the chlorine combines
with the magnesium and the bromine is liberated —

On distilling the liquid the bromine is driven off, and can be
collected in a well-cooled receiver. The addition of any excess of
chlorine results in the formation of bromide of chlorine, and is
therefore to be avoided.

(2.) Bromine is readily obtained from potassium bromide by the
action of manganese dioxide and sulphuric acid, a reaction exactly
analogous to that by which chlorine is obtained from sodium
chloride —

The mixture is gently distilled from a retort into a receiver kept
cold by means of ice.

(3.) Manufacturing Methods.— Practically all the bromine
that is required at the present day is manufactured from crude
carnallite obtained at Stassfurt (see Alkali Metals). This salt
contains bromine combined with magnesium, the magnesium
bromide forming about I per cent, of the magnesium chloride in
tlie crude substance. The final mother-liquors from the manufac-

378

Inorganic Chemistry

ture of potassium chloride, and which were formerly run to waste,
are found to contain about .25 per cent, of bromine as magnesium
bromide, and these liquors are now utilised for the manufacture of
bromine.

The bromine is liberated from its combination with magnesium,
by means of chlorine. In some processes the mother-liquor is
mixed with manganese dioxide and sulphuric acid in a stone vessel

FIG. 104.

resembling an ordinary chlorine still. The magnesium chloride in
the liquor is acted upon by the manganese dioxide and sulphuric
acid with the evolution of chlorine, and this decomposes, the
bromide present displacing the bromine —

The bromine that is driven out is condensed by means of a worrn
condenser.

Bromine

379

Instead of the chlorine being generated within the mother-liquor,
it is now more usually produced in a separate chlorine still, and
passed into the liquor. Fig. 104 shows in diagrammatic form the
method employed. The hot mother-liquor is admitted by the pipe
A into the tower T, which is filled with earthenware balls, between
which the liquid percolates. It leaves the tower by the pipe B,
and flows into the tank W, which is provided with shelves in such
a way that the liquid must circulate through it in the direction
indicated by the arrows. The exit-pipe from this tank empties
into a waste, placed at such a height that the tank is always nearly
full. The liquid in the tank is kept at, or near, the boiling-point, by
means of a current of steam blown in through S. Chlorine from a

FIG. 105.

still is admitted by the pipe L, and passing into the tower by the
pipe B, travels in an opposite direction to the current of liquid.
As the chlorine passes up the tower it meets the descending mother-
liquor, and decomposes the magnesium bromide contained in it
with the liberation of bromine. The bromine vapour leaves the
tower by the pipe C, and is conveyed to a worm (Fig. 105), where it
is condensed Any bromine which dissolves in the water in the
tower is again expelled from solution by the steam as the liquid
traverses the tank W, and is swept up into the tower by the current
of chloride. The condensed bromine as it leaves the worm is
collected in a tubulated bottle, and any vapour which escapes con-

380 Inorganic Chemistry

densation is arrested by the vessel F (Fig. 105). This tube is filled
with iron borings, kept moist by the constant dropping of water
upon them, and any bromine or bromide of chlorine is there con-
verted into iron compounds, which are dissolved by the water,
and flow away into the receiver. The bromine is purified by
redistillation.

Just as in the case of chlorine, these older methods of manu-
facture seem destined to give place to electrolytic processes. The
method now being introduced for the extraction of the bromine
from the Stassfurt liquors depends upon the fact that when a
solution containing a chloride and a bromide is submitted to
electrolysis, the bromine is liberated first, before any chlorine
escapes. Hence, by subjecting the liquors to electrolysis, under
suitable conditions, the whole of the bromine is readily separated.

Properties. — Bromine is a heavy but mobile liquid of a deep
reddish-brown colour. Except in extremely thin layers it is opaque.
It is the only non-metallic element which is liquid at the ordinary
temperature. Bromine boils at 59°, but being a very volatile liquid
it gives off vapour rapidly at the ordinary temperature. A drop of
bromine allowed to fall into a flask immediately evaporates and
fills the vessel with a dark red-brown vapour. The specific gravity
of the liquid at o° is 3.188. At -7° bromine solidifies to a crystal-
line mass. Bromine has a powerful and disagreeable smell. When
the vapour, largely diluted with air, is inhaled, it suggests chlorine
by its smell and by its action upon the mucous membrane of the
throat and nose ; it has in addition, however, a most irritating
action upon the eyes. It is very poisonous, and the liquid exerts a
corrosive action upon the skin ; it produces a yellow colour when
brought in contact with starch.

The vapour-density of bromine, taken at moderately high tem-
peratures, gradually becomes less than is demanded by the formula
Br2, showing that dissociation takes place. In the case of bromine
this is more marked than with chlorine.

Bromine is soluble in water, imparting its own colour to the
solution which is known as bromine water. 100 grammes of water
at o° dissolve 3.60 grammes of bromine. The solubility steadily
diminishes as the temperature rises : at 20° it is 3.208, and at 30°
it is 3.126.

When bromine water is cooled to o° it deposits a crystalline
hydrate similar in composition to the hydrate of chlorine, B^10H2O.

Bromine resembles chlorine in its chemical attributes ; it coin-

Hydrobromic Acid 381

bines directly with metals and many other elements, although with
less energy than is exhibited by chlorine. A fragment of arsenic,
for example, when dropped upon bromine, ignites and burns upon
the surface of the liquid.

Like chlorine, it has bleaching properties, due to its power of
combining with hydrogen.

HYDBOBROMIC ACID (Hydrogen Bromide}.
Formula, HBr. Molecular weight =80. 93. Density =40. 46.

Modes of Formation.— ( i.) Hydrobromic acid can be obtained
by the direct union of its elements. Bromine vapour and hydrogen,
when mixed, do not combine under the influence of light ; neither
does such a mixture explode when a light is applied to it. The
mixture, however, may be caused to burn, when hydrobromic acid
is formed ; or, if the mixed gases be passed through a red-hot tube,
the same result follows. A simple method of preparing hydro-
bromic acid synthetically consists in passing a mixture of hydrogen
and bromine vapour over a spiral of platinum wire, maintained at a
red heat by means of an electric current*

(2.) The best method for the preparation of gaseous hydrobromic
acid consists in dropping bromine upon red phosphorus which has
been moistened with a small quantity of water, when tribasic
phosphoric acid is formed, and hydrobromic acid is liberated —

P + 4H2O + 5Br=H3PO4 + 5HBr.

We may suppose that in this reaction the bromides of phosphorus
are formed and simultaneously decomposed, the action of water
upon these compounds being thus expressed —

= H3PO4+5HBr.

(3.) Hydrobromic acid may be obtained by the action of phos-
phoric acid upon potassium bromide —

3KBr+ H3PO4=K3PO4 + 3HBr.

(4.) If sulphuric acid be employed (as in the formation of hydro-
chloric acid from a chloride), free bromine is simultaneously pro-

* See "Chemical Lecture Experiments," new ed., No. 225.

Inorganic Chemistry

duced, owing to the reduction of a portion of the sulphuric acid
by the hydrobromic acid which is first evolved, thus —

H2SO4 + 2HBr=SO2 + 2H2O-fBr2.

(5.) A dilute aqueous solution of hydrobromic acid may also be
obtained by passing a stream of sulphuretted hydrogen through
bromine water —

= S + 2HBr.

(6.) Hydrobromic acid is readily obtained by the action of
bromine upon certain hydrocarbons, such as turpentine or melted
paraffin. The action is one of substitution, one atom of bromine
replacing one atom of hydrogen in the compound, and the hydrogen
so displaced combining with a second bromine atom to form hydro-
bromic acid. Thus, if the hydrocarbon be represented by the
general formula, CnH2n + 2, the action of bromine will be repre-
sented thus —

CnH2n+2 + Br2= CnH2n+1Br + HBr.

Properties. — Hydrobromic acid is a colourless, pungent-smell-
ing gas, which fumes strongly in the air. It is extremely soluble
in water, forming an acid liquid strongly resembling aqueous
hydrochloric acid.

When boiled, this solution loses either acid or water, until it
reaches a degree of concentration at which it contains 48 per cent.
of hydrobromic acid. The acid of this strength then continues to
boil unchanged at 1 26°. As with hydrochloric acid, the strength
of the liquid which boils at a constant temperature depends upon
the pressure.

Hydrobromic acid is decomposed by chlorine, with the liberation
of bromine —

2HBr+Cl2=2HCl + Br2.

In its chemical behaviour, hydrobromic acid closely resembles
hydrochloric acid, and this resemblance is extended to the bromides.
All bromides are soluble in water, except mercurous bromide,
silver bromide, and. lead bromide, the latter being slightly
soluble-

Hypobromous Acid 383

OXY ACIDS OF BROMINE.

No oxides of bromine corresponding with the oxides of chlorine
have as yet been obtained ; two oxyacids, however, are known, viz. : —

Hypobromous acid . . . HBrO.
Bromic acid HBrO3.

HYPOBROMOUS ACID.

Formula, HBrO.

Mode Of Formation. — An aqueous solution of hypobromous
acid may be obtained by shaking together a mixture of bromine
water and precipitated mercuric oxide, the reaction being ana-
logous to that by which hypochlorous acid is prepared —
HgO + H2O + 2Br2 = HgBr2 + 2HBrO.

Properties. — Hypobromous acid is an unstable compound ; it
breaks up on distillation into oxygen and bromine. By heating to
40° in vacuo, however, it can be distilled without decomposition.
The aqueous liquid so obtained has a pale-yellow colour. It
readily gives up its oxygen, and is a strong bleaching agent ;
when heated to about 60° it decomposes.

Bromous Acid, HBrO2. — When bromine is shaken up with a saturated
solution of silver nitrate, the resulting liquid is believed to contain bromous
acid, probably produced by the formation first- of hypobromous acid, and its
subsequent oxidation to bromous acid * —

r2+H2O=HBrO + AgBr+HNO3.
.2+2AgNO3+Br2=HBrO2+2HNO3+2AgBr.

The acid has not been isolated, nor have any of its salts been obtained.

BROMIC ACID.

Formula, HBrO3.

Modes Of Formation. — (i.) This acid is only known in aqueous
solution ; in this form it may be obtained by the action of bromine
upon silver bromate in the presence of water —

5 AgBrO3 + 3Br2 -f- 3H2O = 5AgBr + 6HBrO3.
The insoluble silver bromide separates out, and the aqueous
acid can be decanted from the precipitate.

(2.) A solution of this acid, mixed with hydrochloric acid, is also
formed when chlorine is passed through bromine water —

* Richards, Jour. Soc. Chem. Ind. , 1906.

384 Inorganic Chemistry

(3.) The decomposition of barium bromate by the requisite
weight of sulphuric acid affords the best method for the preparation
of a pure aqueous solution of bromic acid —

Ba(BrO3)2 + H2SO4=BaSO4 + 2HBrO3.

Properties. — Bromic acid is an unstable, strongly acid sub-
stance, closely resembling chloric acid. The aqueous solution
may be concentrated in vacuo until it contains about 50 per cent,
of bromic acid, representing a composition of I molecule of the
acid to 7 of water. Beyond this degree of concentration, or if
heated to 100°, the acid decomposes into bromine, oxygen, and
water.

The bromates are formed by reactions similar to those by which
the chlorates are produced ; thus, by adding bromine to a solution
of potassium hydroxide, a mixture of potassium bromide and
bromate is obtained —

KBrO3 + 3H2O.

(\nd the two salts can be separated by crystallisation, owing to the
greater solubility of the bromide.

The bromates decompose on being heated, some with the
liberation of oxygen and formation of bromide —

KBrO3=KBr + 3O,

but without the intermediate production of a perbromate. Others
give off their bromine as well as a part of the oxygen they contain,
leaving the metal in combination with oxygen —

Mg(Br03)2 = MgO + Br2 + 50.

IODINE.

Symbol, I. Atomic weight = 126. 92. Molecular weight =
Vapour density = 126. 92.

History. — Iodine was discovered by Courtois (1812), who ob-
served that a beautiful violet vapour was evolved during his
endeavours to prepare nitre from liquors obtained by lixiviating
the ashes of burnt seaweed. The substance was subsequently
investigated by Gay-Lussac.

Iodine 385

Occurrence.— Like all the other members of this group of ele-
ments, iodine is not found in nature in the' uncombined condition.*
In combination it occurs associated principally with potassium,
sodium, magnesium, and calcium, as iodides and iodates.

Iodine is a widely distributed element, although not occurring
in more than small quantities in any particular source. Thus it is
found in small quantities in sea-water and in both marine plants
and animals. The amount of iodine in seaweed varies with diffe-
rent plants ; generally speaking, those from greater depths contain
more than weeds which grow in comparatively shallow waters.

Dry Weed,

Drift weed { Laminaria digitata (stem) . 0.4535
( Laminaria stenophylla . 0.4777

£' , ( Fucus serratus . . . • .-.-* - 0.0856

Cut weed < .

( Ascophyllum nodosum . . 0.0572

Iodine is also found in small quantities in many mineral waters
and medicinal springs.

In small quantities iodine is present in the natural sodium nitrate
of Chili and Peru, known as Chili saltpetre, and at the present day
this constitutes the most abundant source of this element.

Mode Of Formation.— Iodine may be readily obtained by a
precisely similar reaction to that by which both bromine and
chlorine are produced ; thus, if potassium iodide be mixed with
manganese dioxide and sulphuric acid, and the mixture gently
heated in a retort, iodine distils over and condenses in the form of
greyish-black crystals — •

SKI + MnO2 + 2H2SO4=K2SO4+ MnSO4 + 2H2O + 12.

Manufacturing1 Processes. — On an industrial scale iodine is
obtained from two sources, namely, from seaweed and from caliche
(Chili saltpetre).

(i.) From seaweed. The weeds chiefly employed are the Lami-
naria digitata and Laminaria stenophylla. The weed is burnt in
shallow pits, care being taken to avoid too high a temperature ; the
maximum yield of iodine being obtained if the ash is not allowed
to fuse. This ash is technically known as kelp^ and if the weed is
properly burnt, it should yield a kelp containing from 25 to 30 Ibs.
of iodine per ton. The kelpers, however, usually lose about half

* It is on record (Wanklyn, Chem. News, 54) that a minute quantity of free
iodine was found in the water from Woodhall Spa.

2 B

386 Inorganic Chemistry

the iodine on account of burning the weed at too high a tempera-
ture, thereby fusing the ash into a hard slag, instead of obtaining
a porous residue.

An improved process of carbonising the weed was introduced by
Stanford (1863), in which it was heated in large retorts, whereby
the volatile products of the distillation, consisting largely of tar
and ammoniacal liquor, could be collected. The kelp obtained
by this method is in a very porous condition, and contains the
whole of the iodine originally present in the weed.

A still more recent process for extracting the iodine from sea-
weed, and at the same time obtaining other useful materials, has
since been discovered by Stanford. The weed is boiled with
sodium carbonate and filtered : the residue consists of a substance
called algulose. Hydrochloric acid is added to the filtered liquid,
which precipitates a compound known as alginic acid, and this
is again separated by filtration. The liquor is neutralised with
sodium hydroxide, evaporated to dryness and carbonised. The
residue, which is known as "kelp substitute," contains all the
iodine, as well as the potash salts, and should yield about 30 Ibs.
of iodine per ton.

[The alginic acid obtained in this process is purified and converted into the
sodium salt, which constitutes the commercial " algin," a material of a gelatin-
ous or albuminous nature which has lately been put to a number of useful .
applications.]

The kelp obtained by either of these methods is lixiviated with
water in large iron vats, whereby all the soluble salts are extracted.
This aqueous liquid is concentrated in large open boiling pans,
and the less soluble salts, viz., the alkaline sulphates, carbonates,
and chlorides, are allowed to crystallise. The mother-liquor is
then mixed with sulphuric acid and allowed to stand. The sul-
phuric acid decomposes any sulphides and sulphites which may
be present, with the separation of sulphur ; it also converts the
bromides and iodides into the corresponding sulphates, with the
liberation of hydrobromic and hydriodic acids which remain in
solution, while the alkaline sulphates are deposited from the liquid,
and are technically known as plate sulphate. The liquor is then
transferred to the iodine still^ which is an iron pot furnished with
a leaden cover into which two exit-pipes are fixed (Fig. 106).
These are connected to a series (usually ten in each row) of large
earthenware jars or aludels. A gentle heat is applied, and

Iodine

387

manganese dioxide is introduced from time to time through
the opening. The iodine is evolved according to the following
equation —

2HI + MnO2+H2SO4=MnSO4

and condenses in the jars. These vessels are also furnished with a
tubulus upon their under side, so that the water which is evolved
during the distillation can drain out, and run off down the trough in
which the jars are resting.

(2.) From Chili saltpetre. The crude sodium nitrate of Chili

FIG. 106.

and Peru, known as caliche, contains small quantities of iodine,
chiefly as sodium iodate. Although the amount of iodine in
caliche is only very small, averaging about 0.2 per cent., in view
of the enormous quantity of nitrate that is turned out, the aggre-
gate amount of iodine is very great. The iodine is now extracted,
and the supply of this element that is now manufactured from this
source is more than the total consumption of iodine in the whole
world. The process is based upon the fact, that when a solution
of hydrogen sodium sulphite (sodium bisulphite) is added to a
solution of an iodate, iodine is precipitated, thus —

The final mother-liquor from the sodium nitrate, or caliche, in
which all the iodate has concentrated, contains as much as 22
per cent, of this salt. This liquor is mixed with the requisite
proportion of the hydrogen sodium sulphite solution, in large lead-

388 Inorganic Chemistry

lined vats, and the precipitated iodine allowed to settle. It is
then washed and pressed into blocks, and is found to contain
from 80 to 85 per cent, of iodine. This impure product is then
distilled at a gentle heat from iron retorts, the vapour being con-
densed in a series of earthenware receivers much as in the older
method.

Properties. — Iodine is a bluish-black shining solid, somewhat
resembling graphite in lustre and general outward appearance. It
crystallises in large brilliant plates, which have a specific gravity
of 4.95. When heated to 107° iodine melts and gives off vapour
having a beautiful violet colour. Its boiling-point is about 175°.

Iodine vaporises slowly at ordinary temperatures, and sublimes
from one part to another of a bottle in which a small quantity of
it is contained. The smell of iodine vapour is somewhat irritating
and unpleasant, recalling the smell of moderately diluted chlorine.
When iodine vapour is heated it passes from a violet colour to a
deep indigo blue.* This change in the colour is accompanied
by a diminution of the vapour-density. Up to a temperature of
700° the density of iodine corresponds to the formula I2 ; as the
temperature is raised the density gradually diminishes, until at
1468° it is reduced to less than two-thirds. At this point, 73.1 per
cent, of the iodine molecules have become dissociated into single
atoms.

Iodine is slightly soluble in water, I gramme of iodine being
dissolved by 5.524 litres of water at 10°. This dilute solution,
however, has a perceptible brown colour. Iodine is freely soluble
in aqueous potassium iodide solution, in alcohol, ether, and aqueous
hydriodic acid ; in all these solvents it dissolves to a dark reddish-
brown solution. In chloroform, carbon disulphide, and many
liquid hydrocarbons, iodine is also soluble, but in these solvents
it dissolves to a deep violet solution resembling the colour of the
vapour.

When iodine is brought into contact with starch it forms an
intense blue colour. This reaction is so extremely delicate that
it is capable of revealing the minutest trace of iodine. The exact
nature of this blue compound is not known. The colour disappears
when the liquid is heated to about 80°, but returns on cooling ;
continued boiling destroys it permanently.

In its chemical relations iodine resembles chlorine and bromine,
but with a lesser degree of energy. Both these elements are
* " Chemical Lecture Experiments," new ed., No. 231,

Hydriodic Acid 389

capable of displacing iodine from its combinations with electro-
positive elements, thus —

KI + Br=KBr+I.
KI + C1=KC1 + I.

Iodine combines with many elements, both metals and non-
metals, forming iodides. Phosphorus, when brought in contact
with iodine, at once melts and inflames ; antimony powder dropped
into iodine vapour also spontaneously inflames. When mercury
and iodine are gently heated, energetic combination takes place,
and mercuric iodide is formed.

HYDRIODIC ACID (Hydrogen Iodide).
Formula HI. Molecular weight = 127. 93. Density =63. 96.
Modes of Formation.— ( i.) Hydriodic acid can be obtained

FIG. 107.

synthetically by passing a mixture of hydrogen and iodine vapour
over strongly heated, finely divided platinum.

(2.) It is also obtained by the action of phosphoric acid upon
sodium or potassium iodide (see Hydrobromic Acid).

As in the case of the corresponding bromine compound, sul-

3QO Inorganic Chemistry

phuric acid cannot be employed, as by its action upon the iodide,
iodine and sulphur dioxide are liberated, thus —

2KI +SH2SO4=2HKSO4 + 2H2O + SO2 + I2.

(3.) Hydriodic acid is produced by the action of sulphuretted
hydrogen upon iodine (p. 410). At the ordinary temperature,
and in the absence of water, these two substances do not react,
hydriodic acid being an endothermic compound (p. 168) ; but if
the iodine be suspended in water and sulphuretted hydrogen
passed through, the heat of solution of the hydriodic acid supplies
the necessary energy to enable the action to proceed. When,
however, the solution reaches a sp. gr. of 1.56 the action ceases,
because, as Naumann has shown, the heat produced by the solution
of the product is insufficient to carry on the process beyond this
degree of concentration.

(4.) Hydriodic acid is most readily prepared by the action of
phosphorus upon iodine in the presence of water —

P + 51 + 4H2O = H3PO4 + 5H I.

The red phosphorus and iodine for this reaction may be placed
in a dry flask, and water gradually dropped upon the mixture, when
hydriodic acid is rapidly evolved. The gas is allowed to pass
through a U-tube containing red phosphorus, in order to arrest any
iodine vapour which may accompany it. Hydriodic acid may be
collected over mercury or by displacement, as shown in Fig. 107.

Properties. — Hydriodic acid is a colourless, pungent-smelling
gas, which fumes strongly on coming into the air. The gas is
readily decomposed by heat into hydrogen and iodine. Thus, if a
heated wire be thrust into the gas, or if a spiral of platinum wire
be heated in the gas by means of an electric current, the violet
vapour of iodine at once makes its appearance.

When mixed with chlorine, hydriodic acid is at once decomposed,
with the liberation of iodine, thus —

2HI+C12 = 2HC1 + I2.

Hydriodic acid is one of the most readily liquefied gases ; at o°,
and under a pressure of four atmospheres, it condenses to a colour-
less liquid.

The gas is extremely soluble in water. An aqueous solution of
it is readily produced by allowing the gas, obtained by the method

Iodine Pentoxide

391

of preparation above described, to pass into water. In order to

prevent the water from being drawn back into the generating flask,

it is convenient to pass the gas through a retort arranged in the

position seen in Fig. 108. Should there be

any back rush of water, owing to the inter-

mission of the evolution of gas in the ap-

paratus, the liquid in the beaker will be

drawn up into the retort and there lodge,

leaving the end of the neck open to the

air.

A saturated aqueous solution of hydri-
odic acid at o° has a specific gravity
of 2. At the ordinary pressure the
strongest acid that can be obtained by
distillation has a specific gravity of 1.67,
and contains 57.7 per cent, of hydriodic
acid. This solution boils at 127°. As
in the case of the corresponding bromine
and chlorine compounds, the particular

strength of acid which has a constant boiling-point is a function
of the pressure.

Aqueous hydriodic acid, when freshly prepared, is colourless, but
it rapidly turns brown, owing to the oxidation of the compound
and the solution of the liberated iodine in the acid —

FIG. 108.

OXIDE AND OXYACIDS OF IODINE.

One * compound of iodine with oxygen is well known, and three
oxyacids, viz. : —

Iodine pentoxide .... I2O5.
lodic acid . . . . . HIO3.
Periodic acid . . . . HIO4.
Hypoiodous acid "• •.-.» .-..'W . HIO.

IODINE PENTOXIDE (lodic Anhydride}.

Formula, I2O5.

Mode of Formation. — When iodic acid is heated to 170°, it
loses water and is converted into the pentoxide —
2HIO3 = H2O + I2O6.

* A second oxide, IO2 or I2O4, has recently been described by Muir, four.
Chem, Soc., April 1909.

392 Inorganic Chemistry

Properties. — Iodine pentoxide is a white crystalline solid body.
It is soluble in water, combining with a molecule of the water to
form iodic acid. Iodine pentoxide is more stable than any of the
oxides of the other halogens, but, when heated to a temperature of
30x3°, it decomposes into its elements.

IODIC ACID.

Formula, HIO3.

Modes of Formation.— ( i.) Iodic acid can be prepared by
adding to a solution of barium iodate the requisite amount of
sulphuric acid demanded by the equation —

Ba(IO3)2+H2SO4=BaSO4+2HIO3.

The aqueous solution of iodic acid is decanted from the preci-
pitated barium sulphate, and may be concentrated at 100° without
being decomposed.

(2.) When chlorine is passed through water in which powdered
iodine is suspended, a mixture of iodic acid and hydrochloric acid
is produced —

3H2O + I + 5C1 = 5HC1 + HIO3.

The hydrochloric acid may be removed by the addition of preci-
pitated silver oxide to the solution, and separating the precipitated
silver chloride by filtration.

(3.) Iodic acid is most conveniently prepared by heating iodine
with nitric acid, whereby the iodine is oxidised and a mixture of
oxides of nitrogen is evolved as dense red vapours —

Properties. — Iodic acid is a white crystalline solid, soluble in
water. The aqueous solution shows an acid reaction with litmus,
but the colour is ultimately discharged by the bleaching action of
the compound. Iodic acid does not form any blue colour with
stnrch ; being, however, an oxidising substance, it readily gives up
oxygen to such reducing agents as sulphur dioxide, sulphuretted
hydrogen, or hydriodic acid, with the liberation of iodine, thus —

Periodic Acid 393

If, therefore, a small quantity of sulphurous acid be added to a
dilute solution of iodic acid, previously mixed with starch, the blue
iodide of starch will be formed. This reaction affords an excellent
illustration of the time required for certain chemical changes to go
forward. It is readily possible to obtain an interval of 30 to 60
seconds between the addition of the sulphurous acid and appear-
ance of any visible result, when at the expiration of that time the
whole mass of the liquid suddenly turns blue.*

lodates. — When iodine is dissolved in potassium hydroxide, a
mixture of potassium iodide and iodate is produced by an analogous
reaction to that which takes place with either bromine or chlorine —

With the exception of the iodates of the alkali metals, the iodates
are for the most part insoluble in water. On being heated they
behave in a similar 'manner to the bromates, some being decom-
posed into an iodide and oxygen, while others leave a metallic
oxide and evolve iodine as well as oxygen. The alkaline iodates
are capable of uniting with iodic acid, forming salts which are
termed acid and di-acid iodates, thus —

Normal potassium iodate . . . KIO3.
Acid potassium iodate . . . KIO3,HIO3.
Di-acid potassium iodate . . . KIO3,2HIO3.

PERIODIC ACID.

Formula, HIO4,2H2O or H5IO6.

Modes of Formation. — (i.) The compound represented by the
formula HIO4has never been obtained; when aqueous solutions
of periodic acid are evaporated, the compound which crystallises
out has the composition HIO4,2H2O, or H5IO6.

It may be obtained by boiling silver periodate with water, when
an insoluble basic silver salt is produced —

2AgI04+4H20 = Ag2H3I06+HI04,2H20.

The silver periodate is prepared by passing chlorine into an aqueous solu-

* See Experiment 246. " Chemical Lecture Experiments," new ed.

394 Inorganic Chemistry

tion of sodium iodate and sodium hydroxide, when the sparingly soluble
disodium periodate separates out —

NaIO3 + SNaHO + Cl2=2NaCl + Na^IO,,.

This sodium salt is then dissolved in nitric acid and silver nitrate added,
whereby AgIO4 is formed, which crystallises out on concentration —

= 2NaNO3 + 4H2O + 2NaIO4.
\ 2NaIO4+2AgN03=2NaN03+2AgI04.

(2.) Periodic acid is also formed by the addition of iodine to
an aqueous solution of perchloric acid —

2HC104+2H20 + I2=C12 + 2HI04,2H20.

Properties. — The acid having the composition HIO4,2H2O
is a colourless, crystalline, deliquescent substance. It melts at
133°, and at 150° is decomposed into iodine pentoxide, water, and
oxygen —

2H5I06=I205 + 5H20 + 02.

The acid cannot be converted into HIO4 by heat, for oxygen is
evolved as soon as water begins to be given off.

The Periodates constitute a numerous class of salts, many of them being of
a highly complex composition. On the assumption that iodine is monovalent
in these compounds, their classification is somewhat difficult, and they must be
represented as associations of molecules of salts of the unknown monobasic
periodic acid, HIO4, with metallic oxide and water in various proportions —
thus, the silver periodate in the foregoing equation, Ag2H3IOg, would lie
expressed by the formula, 2AgIO4,Ag2O,2H2O.

The classification of these compounds is much simplified if we regard iodine
as here functioning as a heptavalent element. On this assumption the perio-
dates may be considered as the salts of various hypothetical acids, which are
all derived from the compound I(HO)7 (itself hypothetical) by the withdrawal
of varying quantities of water. Thus, by the successive removal of one mole-
cule of water, the following three acids would be formed—

I(HO)7 -H2O=IO(HO)5 . -.'••"'. H5I06. (i.)
IO(HO)5 -H20=I02(HO)3 . '. . H3I05. (2.)
I02(HO)3-H20=I03(HO) . . . HI04. (3.)

From these three acids the following salts may be regarded as being
derived —

(i.) Na2H3I06; Ag2H3IO6 ; Ag5IO6 ;
(2.) Ag3I05; Pb3(I05)2.
(3.) KI04 ; AgI04.

Hypoiodous Acid and Hypoiodites 395

By the abstraction of one molecule of water from two molecules of these
acids, still more complex acids would be derived, thus —

IO(HO)4

IO(HO)4
I02(H02

I02(HO)2

. And from these two acids the following periodates may be regarded as being
derived —

(4.) Zn4I2On ;
(5.) Ag4I209;

HYPOIODOUS ACID AND HYPOIODITES.

When an aqueous solution of iodine is added to either ammonia, potassium,
or sodium hydroxides, lime-water or baryta- water, a colourless solution is
obtained which possesses bleaching properties. The liquid is a dilute solution
of the hypoiodite and iodide of the alkali used. Somewhat stronger solutions
may be produced by adding small quantities of powdered iodine to the
mixture —

2KHO + I2+Aq=KIO+KI + H2O+Aq.

A dilute solution of the acid itself is obtained by shaking mercuric oxide
with iodine and water (see Hypochlorous Acid, p. 373).

The solution of the alkaline hypoiodite obtained by the above reaction pos-
sesses well-marked bleaching properties. When freshly prepared it is without
action upon starch, but is immediately decomposed by even so feeble an acid
as carbonic acid, when the blue starch compound is at once formed.

A compound of iodine with lime, analogous to bleaching powder, has been
obtained by shaking powdered iodine with milk of lime. The compound in
the presence of water appears to behave in the same way as bleaching powder,
yielding a solution of calcium hypoiodite and calcium iodide —

2Ca(OI)I = Ca(OI)2+CaI2.

On filtering the mixture a colourless liquid is obtained, which gives no reaction
with starch, but which yields iodine when treated with an acid.

Neither the acid nor any of its salts has been isolated, being known only in
dilute solution. The compounds are all extremely unstable, decomposing at
the ordinary temperature in a few hours, and in a few minutes when the
solutions are boiled ; the salts passing into iodides and iodates—

3KIO=

while the acid decomposes first into hydiiodic and iodic acids, which then
react upon each other with elimination of free iodine.

396 Inorganic Chemistry

COMPOUNDS OF THE HALOGENS WITH EACH OTHER.

Chlorine unites both with bromine and with iodine, and the two latter
elements combine with each other.

(i.) Chlorine and. Bromine. —Bromine monochloride. This substance is
obtained as a reddish-yellow liquid, when chlorine gas is passed into bromine.
The compound is believed to have the composition BrCl.

(2.) Chlorine and Iodine. — Iodine monochloride, IC1. When dry chlorine
is passed over iodine, the latter rapidly melts, forming a dark reddish-brown
liquid, strongly resembling bromine in appearance. The liquid solidifies to a
mass of red prismatic crystals, which melt at 25°. It is decomposed by water
into iodic and hydrochloric acids, and iodine is liberated —

5ICH-3H2O=HIO3+5HC1+2I2.

Iodine trichloride, IC13. This compound is formed by passing an excess of
chlorine over iodine, or by passing chlorine through iodine monochloride. It
is also formed when hydriodic acid is acted upon by an excess of chlorine —

HI+2C12=HC1 + IC13.

Iodine trichloride is a yellow solid substance, crystallising in long brilliant
needle-shaped crystals, which sublime at the ordinary temperature. When
gently warmed it melts, at the same time dissociating into chlorine and the
monochloride ; on cooling, reunion takes place with the reformation of IC13.

(3.) Bromine and Iodine. — Two compounds of these elements are believed
to exist, viz., a crystalline solid and a deep-coloured liquid. Their composition
is probably expressed by the formulas, IBr and IBr3.

CHAPTER II
THE ELEMENTS OF GROUP VI. (FAMILY 5.)

Oxygen, O . . 16.00 Selenium, Se . . 79.2

Sulphur, S . . 32.07 I Tellurium, Te . . 127.5

THE relation in which oxygen, the typical element, stands to the
remaining members of the family is very similar to that between
fluorine and the other halogens.

All the elements of this family unite with hydrogen, forming
compounds of the type RH2 —

OH2, SH2, SeH2, TeH2 ;

but the hydride of oxygen stands apart from the others in many of
its attributes. Thus at ordinary temperatures it is a colourless and
odourless liquid, while the remaining compounds are all foetid-
smelling and poisonous gases.

Sulphur, selenium, and tellurium each combines with oxygen,
forming respectively SO3, SeO3, and TeO3, while none of these
elements in a divalent capacity forms a similar compound ; that is
to say, no such combinations are known as OS3, or OSe3, although
amongst themselves they unite, forming SeS2 and TeS2.

Sulphur, selenium, and tellurium also unite with oxygen, forming
dioxides, SO2, SeO2, and TeO2, in which these elements are pos-
sibly tetravalent, in which case the constitution of the compounds
will be represented thus, O = S = O ; O = Se = O.

We may, however, consider them as functioning in a divalent

/° /°

capacity, and regard the oxides as constituted thus, S<^ | ; Se<( | ,

\O NO

in which case we may look upon ozone as being the corresponding

/°

oxygen compound, OO2, O<^ | .

397

398 Inorganic Chemistry

All the elements of this family combine with chlorine, producing
compounds .having the following composition : —

Oxygen. Sulphur. Selenium. Tellurium.

O2C1

S2C12 Se2Cl2

OC12 SC12 TeCl2

SCJ4 SeCl4 TeCl4

Oxygen again differs from the other members by alone forming
a compound of the type, R2C1. This element also shows no ten-
dency to function with a higher atomicity than that of a divalent ;
while the others unite with four atoms of the halogen, thereby
exhibiting their tetravalent nature.

The members of this family pass by a regular gradation from
the strongly electro-negative, gaseous, non-metal oxygen to the
feebly negative and slightly basic element tellurium, which possesses
many of the properties of a true metal. Selenium and tellurium
are both elements which lie very close to that ill-defined boundary
between the metals and non-metals, and are on this account some-
times termed metalloids. In tellurous oxide, TeO2, we have a
compound which is both an acid-forming and a salt-forming oxide,
its acidic and basic properties being nearly equally balanced. Thus,
it unites with water, forming tellurous acid, H2TeO3, corresponding
to sulphurous acid, H2"SO3 ; while tellurium replaces hydrogen in
sulphuric acid, forming tellurium sulphate, Te(SO4)2.

Of the four elements of this family, oxygen is by far the most
abundant, both in combination and in the free state ; sulphur is
more plentiful than the other two, and tellurium occurs in the
smallest quantity.

The element oxygen has already been treated in Part 1 1.

SULPHUR.

Symbol, S. Atomic weight =32. 07. Molecular weight =64. 14.

Occurrence. — In the free state this element occurs chiefly in
volcanic districts. In Italy and Sicily large quantities of native
sulphur are found, which have long been the most important
European sources of this substance. Large deposits are to be met
with in Transylvania and in Iceland, and it also occurs in beds,

Sulphur 399

often of great thickness, in parts of China, India, California, and
the Yellowstone district of the Rocky Mountains. These natural
deposits are sometimes found stratified with beds of clay or rock, '
but they often occur as what are known as " living " beds, in which
the sulphur is continuously being formed as the result of chemical
decompositions which are at present at work. Such a " living "
sulphur bed is known as a solfatara, and, as in the case of the
Iceland deposits, they are usually found associated with geysers,
fumaroles, and other signs of volcanic action.

In combination with hydrogen, sulphur occurs as sulphuretted
hydrogen. Enormous quantities of sulphur are found combined
with various metals, constituting the important class of substances
known as sulphides ; as, for example, galena^ or lead sulphide,
PbS ; zinc blende ', or zinc sulphide, ZnS ; pyrites, or iron sulphide,
FeSa ; copper pyrites ; or copper iron sulphide, Cu2Fe2S4 ; stibnite^
or antimony sulphide, Sb2Ss ; cinnabar ; or mercury sulphide, HgS.

In combination with metals and oxygen, sulphur occurs in
sulphates, such as gypsum^ CaSO4,2H2O ; heavy spar^ BaSO4 ;
kieserite, MgSO4,H2O.

Modes of Formation. — (i.) Sulphur is formed when sulphu-
retted hydrogen is brought in contact with sulphur dioxide ; the
two gases mutually decompose one another with the formation of
water and the precipitation of sulphur —

2H2S4 SO2=2H2O + 3S.

(2.) It is also produced when sulphuretted hydrogen is burnt
with an insufficient supply of air —

This reaction probably takes place in two stages, a portion
of the sulphuretted hydrogen burning to sulphur dioxide, and this
then reacting upon a further quantity of sulphuretted hydrogen,
thus—

(a) H2S + 3O = H2O + SO2.

(b) 2H2S + SO2=2H2O + 3S.

It is supposed that some of the free sulphur found in volcanic
regions has been produced by this action of these two gases upon
one another.

Extraction of Sulphur from Native Sulphur.— Natural
sulphur is always more or less mixed with earthy or mineral

400 Inorganic Chemistry

matters, from which it is necessary to free it. This is usually
effected by melting the sulphur and allowing it to flow away from
the accompanying impurities. The crude sulphur rock is stacked
in brick kilns having a sloping floor, and the mass ignited by
introducing through openings in the heap burning faggots of
brushwood. The heat produced by the combustion of a part of the
sulphur causes the remainder to melt and collect upon the sloping
floor of the kiln, from which it can be drawn off into rough mouldr,.
The loss of sulphur by this method is very considerable, usually
not more than two-thirds of the total amount contained in the rock
being obtained.

(3.) Sulphur may be obtained by heating certain metallic sul-
phides ; thus when iron pyrites is heated it yields one-third of its
sulphur —

If the pyrites be roasted in kilns, the whole of the sulphur is
obtained, partly as free sulphur and partly as sulphur dioxide—

This method was at one time rather extensively employed for
the preparation of sulphur on a manufacturing scale, but has now
practically gone out of use, the pyrites being usually roasted with
excess of air, whereby the whole of the sulphur is converted into
sulphur dioxide for use in the manufacture of sulphuric acid.

By a similar process, sulphur is obtained as a bye-product during
the roasting of copper pyrites in the first stage of the operation of
copper-smelting.

(4.) Large quantities of sulphur are now extracted from the vat-
waste or alkali-waste ', obtained in the manufacture of sodium
carbonate by the Leblanc process. This material consists largely
of an insoluble oxy-sulphide of lime, a compound containing calcium
sulphide (CaS) and calcium oxide (CaO) in varying proportions.
Montfs process ', which, however, has now been entirely superseded
by Chances process (p. 411), is the following : A current of air is
blown through the compound, whereby the calcium sulphide it
contains is ultimately converted into a mixture of calcium hydro-
sulphide (CaH2S2), thiosulphate (CaS2O3), and polysulphide (CaS6),
according to the following equations —

Sulphur 401

This reaction goes forward in several stages, in the course of
which a quantity of sulphur is set free ; this is then acted upon by
the calcium hydroxide, with the formation of calcium polysulphide
and calcium thiosulphate, thus —

(2.) 3CaH2O2 + 125 = 2CaS5 + CaS2O3 + 3H2O.

The materials are alternately oxidised and lixiviated several
times, and the liquor is then treated with excess of hydrochloric
acid at a temperature of about 60°, which decomposes the various
sulphur compounds according to the following equations —

(a.) CaH2S2 + 2HCl

CaCl+SO

The best results are obtained when the sulphur compounds are
present in such proportions that the SO2 evolved by reaction c is
sufficient to decompose the whole of the SH2 produced by the
other two reactions, so that neither gas escapes —

(5.) Sulphur is also obtained from the spent oxide of iron which
has been used in the " purifiers " employed upon gas-works. Coal
gas contains considerable quantities of sulphuretted hydrogen,
which are removed from the gas by passing it through hydrated
ferric oxide (Fe2H6O6), which absorbs the whole of the sulphuretted
hydrogen, thus —

When the compound has lost its power to absorb sulphuretted
hydrogen, the material is thrown out of the purifiers and exposed
to air and moisture, when the iron becomes reconverted into the
hydrated oxide, and the sulphur is set free —

2FeS + 3O + 3H2O = Fe2H6OG + 2S.

This revivified material is then employed for the purification of
a further quantity of gas. It is found that after a certain number
of revivifying operations the substance begins to lose its power of
absorbing any additional sulphuretted hydrogen, and as it then

402

Inorganic Chemistry

contains nearly half its weight of sulphur, it becomes a valuable
source of this element. The sulphur is obtained from it by distil-
lation, or the material may be roasted in special kilns, whereby the
sulphur is converted into sulphur dioxide, and employed for the
manufacture of sulphuric acid.

Purification. — The crude sulphur obtained by the foregoing
methods is purified by distillation, the process being carried out in
the arrangement shown in Fig. 109. The sulphur is first melted
in an iron pot d, and the liquid substance drawn off at will by

FIG. 109.

means of the pipe F into the retort B. The sulphur is there
boiled by means of the fire, and the vapour allowed to issue into the
large brickwork chamber G. As the vapour enters the chamber, it
condenses upon the walls and floor in the fonn of a light, powdery
deposit, consisting of minute crystals, and constituting the flowers
of sulphur of commerce. As the process continues, and the brick-
work becomes hot, this soft powder melts and collects upon the
floor as an amber- coloured liquid? which is run out from time to

Sulphur 403

time from the opening at H, and cast either into large blocks or
into cylindrical rods, by means of wooden moulds. In the latter
form it is known as roll sulphur.

When the sulphur vapour first enters the chamber and mixes
with the air, the mixture frequently ignites with a feeble explosion ;
the chamber, therefore, is furnished with a valve, S, at the top,
whereby the pressure developed at the moment of combustion may
be relieved.

Properties. — Sulphur, as ordinarily seen, is a pale-yellow brittle
crystalline solid. It is insoluble in water, but readily dissolves in
carbon disulphide, and to a greater or less degree in turpentine,
benzene, chloroform, sulphur chloride, and many other solvents.
It is a non-conductor of electricity, and an extremely bad con-
ductor of heat. A piece of sulphur on being very gently warmed,

FIG. no.

even by being grasped in the hand, may be heard to crack by the
unequal heating, and will ultimately fall to pieces. At a tem-
perature of 114.5° sulphur melts to a clear amber-coloured and
moderately mobile liquid ; on raising the temperature of this
liquid its colour rapidly darkens, and at the same time it loses its
mobility, until at a temperature of about 230° the mass appears
almost black, and is so viscous that it can no longer be poured
from the vessel. As the temperature is still further raised, the
substance, while retaining its dark colour, again becomes liquid,
although it does not regain its original limpidity. At 448° the
liquid boils, and is converted into a pale yellowish-brown coloured
vapour. On allowing the boiling sulphur to cool, it passes through
the same changes in reverse order until it solidifies.

When the vapour of sulphur is heated to 1000°, it is converted

404

Inorganic Chemistry

into a true gas, and has a density of 32, one litre of the gas weigh-
ing 32 criths.

Sulphur is known to exist in four allotropic modifications, two of
which are crystalline and two amorphous.

(a) "Rhombic" Sulphur. — Of the two crystalline varieties this is
the more stable. Sulphur, therefore, that occurs native is found
crystallised in this form, namely, orthorhombic pyramids. It may
be obtained by allowing a solution of sulphur in carbon disulphide
to slowly evaporate. Fig. no represents two large crystals of
sulphur obtained in this way.

Orthorhombic crystals of sulphur can also be obtained under
certain conditions when melted sulphur is allowed to crystallise.

FIG. in.

Sulphur in the liquid condition exhibits the phenomenon of sus-
pended solidification to a very high degree, and if the liquid be
carefully cooled out of contact with air, the temperature may fall
to 90° before solidification takes place. If into the liquid in this
state a crystal of the rhombic variety be dropped, the sulphur
begins to solidify in crystals of that form. If the superfused sulphur
be contained in a hermetically closed flask, the liquid frequently
deposits orthorhombic crystals, and by allowing the mass to
partially solidify, and quickly inverting the flask, the crystals may
be seen upon the bottom of the vessel.

The specific gravity of this form of sulphur is 2.05.

(fi) "Prismatic" Sulphur. — When melted sulphur is allowed

Sulphur 405

to cool under ordinary conditions, such as in a crucible or
beaker, it crystallises in the form of prismatic needles, belong-
ing to the monoclinic or monosymmetric system. By allowing
the mass to partially solidify, and pouring off the still liquid por-
tion, these crystals will be seen lining the inside of the beaker
as long translucent prisms. Fig. in shows such a mass of
crystals. Prismatic crystals of sulphur are also obtained when
this element is crystallised from a hot solution in oil of tur-
pentine.

The specific gravity of this form of sulphur is less than that of
the orthorhombic variety, being 1.98.

At ordinary temperatures this modification is unstable, and in
the course of a day or two the crystals lose their translucent
appearance, owing to their becoming broken down into a number
of smaller crystals of the rhombic variety, and present the opaque
yellow appearance of ordinary roll sulphur. This change from
the prismatic to the " rhombic " variety, which takes place more
quickly when the crystals are scratched or subjected to vibration,
is attended with evolution of heat. When monoclinic sulphur is
thrown into carbon disulphide, its transformation into the stable
modification takes place rapidly, and in this way, by means of a
thermopile, the heat evolved by the change may be rendered
evident. As carbon disulphide, however, at once exerts its solvent
action upon the "rhombic" sulphur the moment it is formed, the
reduction of temperature resulting from this cause would com-
pletely overbalance and mask the more feeble heat effect produced
by the passage of the sulphur from the unstable to the stable form.
In order, therefore, to render evident the heat resulting from the
change of crystalline form, the carbon disulphide must be pre-
viously allowed to dissolve as much sulphur as it can take up. If
a small quantity of carbon disulphide, so saturated with sulphur,
be placed in a corked flask, and stood upon the face of a thermo-
electric pile* in connection with a galvanometer, and a quantity
of prismatic crystals of sulphur be quickly thrown into the liquid,
a sensible deflection of the galvanometer needle will be seen in the
direction caused by heat.

Although under ordinary conditions monoclinic sulphur is un-
stable and passes into the " rhombic " form, at temperatures between

* The thermo-electric pile is a delicate physical instrument employed for
registering slight changes of temperature ; for descriptions of the apparatus
the student must consult text-books on physics.

406 Inorganic Chemistry

100° and 114° it appears to be the more stable variety, for at this
temperature " rhombic " sulphur passes into the monoclinic variety.
(y) Plastic Sulphur. — When sulphur which has been heated
until it reaches the viscous condition is suddenly plunged into
water, or when boiling sulphur is poured in a thin stream into
water, the substance solidifies to a tough elastic material some-
what resembling indiarubber. The sulphur in this form is known
as plastic sulphur. This variety is best obtained by distilling a
quantity of ordinary sulphur from a glass retort (Fig. 112), and
allowing the distilled liquid to flow in a fine stream into cold
water placed for its reception. As the liquid sulphur falls into
the water, it congeals to the plastic condition as a continuous

thread, which winds itself in
a regular manner into beauti-
ful coils of a delicate trans-
lucent amber colour. The
specific gravity of plastic sul-
phur is 1.95, and it is not
soluble in carbon disulphide.
At ordinary temperatures this
allotrope of sulphur is gra-
dually transformed into the
stable "rhombic" variety ; in
the course of a few days it
FIG. 112. loses its transparency and

elasticity, and becomes con-
verted into the ordinary lemon-yellow brittle condition of common
sulphur. This change takes place more quickly if the plastic
material be stretched and worked between the fingers, and still
more readily by heating it for a few moments to 100°, and allowing
it again to cool.

(8) White Amorphous Sulphur. — When sulphur is heated, and
the vapour condensed upon a cool surface, as in the formation of
ordinary flowers of sulphur, although the greater portion of the
sulphur is sublimed in the orthorhombic form, the sublimate con-
tains a small amount of sulphur in the form of an amorphous
powder, which is almost milk-white in colour.

This modification is best obtained by treating flowers of sulphur,
which usually contains as much as 5 or 6 per cent, of amorphous
sulphur, with carbon disulphide, whereby the orthorhombic variety
is dissolved, and the white amorphous substance, which is insoluble

Sulphur 407

in that liquid, is left behind. By filtering the liquid and washing
the residue with carbon disulphide until the whole of the soluble
sulphur is removed, the amorphous powder may be obtained in a
state of purity.

This amorphous substance is also produced in small quantity, by
the action of light upon a solution of sulphur in carbon disulphide.
Thus, if a perfectly clear solution of sulphur in this liquid be placed
for even a few minutes in the path of a beam of electric light, the
solution will be seen to become rapidly turbid, owing to the forma-
tion of this insoluble modification.

This milk-white amorphous modification is stable at the ordinary
temperature, and therefore does not pass spontaneously into the
rhombic variety. When heated to a temperature of 100°, it quickly
becomes yellow in colour, and is then readily soluble in carbon
disulphide, having been transformed at that temperature into the
ordinary stable form.

Milk Of Sulphur. — This substance is a medicinal preparation,
obtained by precipitating sulphur from a polysulphide of lime by
means of hydrochloric acid. Flowers of sulphur and milk of lime
are boiled together for some time, and after settling the clear
reddish liquid containing the calcium polysulphides is decanted off,
and hydrochloric acid added to it ; calcium chloride is formed,
and sulphur in a fine state of subdivision is precipitated —

The product so obtained is pale yellow in colour, and consists of
ordinary sulphur often contaminated with considerable quantities
of calcium sulphate, derived from sulphuric acid present in the
hydrochloric acid employed in the precipitation.

When sulphur in any of its modifications is heated in the air, it
takes fire and burns with a pale blue flame, giving rise to sulphur
dioxide ; when burnt in oxygen a small quantity of sulphur tri-
oxide is at the same time produced.

Finely divided sulphur, when exposed to air and moisture, under-
goes slow oxidation even at ordinary temperatures, with the forma-
tion of sulphuric acid. Thus, if flowers of sulphur be moistened
with water and freely exposed to the air, in a short time the water
will be distinctly acid. On this account sulphur that is used for
pyrotechnic purposes is thoroughly washed and dried, and pre-
served in warm dry places.

408 Inorganic Chemistry

Sulphur combines directly with many metals under the influence
of heat, forming sulphides, the union in many cases being accom-
panied by vivid combustion. Thus, a strip of copper, when intro-
duced into sulphur vapour, burns brilliantly with the formation of
copper sulphide, and a red-hot bar of iron, when pressed against
a roll of sulphur, burns in the vapour which is generated, and the
molten sulphide falls in scintillating masses through the air —

= FeS.

Heated with sodium or potassium, the alkaline sulphides are
formed with deflagration —

K2 + S = K2S.

COMPOUNDS OF SULPHUR WITH HYDROGEN.

Two compounds of these elements are known, namely —
Hydrogen sulphide or sulphuretted hydrogen . H2S.
Hydrogen persulphide ... . ..,,.* f, H2Ss.

HYDROGEN SULPHIDE.

Formula, H2S. Molecular weight = 34.08. Density = 17.04.

Occurrence. — This gas is evolved in volcanic regions, and is
met with in solution in sulphur mineral waters.

Modes Of Formation.— ( i.) Sulphuretted hydrogen may be
formed by the direct union of its elements, by passing a mixture of
hydrogen and the vapour of sulphur through a strongly heated
tube. In small quantity it is produced when hydrogen is passed
into boiling sulphur, or over certain heated metallic sulphides.

(2.) Sulphuretted hydrogen is most readily obtained by the
action of either hydrochloric or sulphuric acid upon ferrous sul-
phide, thus —

FeS + 2HCl = FeCl2+H2S.
FeS + H2SO4=FeSO4 + H2S.

The ferrous sulphide in broken fragments is placed in a two-
necked bottle, similar to the apparatus (Fig. 29) employed for
the preparation of hydrogen, and the dilute acid poured upon it.
The gas is rapidly evolved without the application of heat. The
gas obtained by this method always contains free hydrogen,
owing to the presence of uncombined iron in the ferrous sulphide.

Hydrogen Sulphide 409

(3.) Pure hydrogen sulphide may be obtained by heating anti-
monytrisulphide(grey antimony ore)with strong hydrochloric acid—

(4.) Also by the action of water upon aluminium sulphide
(p. 623).

(5.) Sulphuretted hydrogen is produced during the putrefaction
of organic substances containing sulphur, the offensive smell of a
decomposing egg being due to the presence of this gas. It is also
produced during the destructive distillation of coal, by the direct
union of hydrogen with the sulphur contained in the pyrites, hence
coal gas always contains sulphuretted hydrogen amongst its im-
purities.

Properties. — Sulphuretted hydrogen is a colourless gas, having a
somewhat sickly sweetish taste and an extremely offensive odour.
It acts as a powerful poison when inhaled in the pure state, and
even when very largely diluted with air it gives rise to dizziness
and headache. Its poisonous effects are more marked upon some
animals than others : thus, a bird was found to die in an atmosphere
containing only T-5Joo °f sulphuretted hydrogen, while it required an
amount equal to ^j to poison a hare ; and again, cold-blooded
animals are in no way affected by inhaling these proportions of
the gas. Sulphuretted hydrogen is moderately soluble in water ;
at ordinary temperatures water dissolves about three times its own
volume of the gas. In collecting it over water, therefore, consider-
able loss results unless the water be warm. The coefficient of
absorption by water at o° is 4.3706.

The aqueous solution gives an acid reaction with litmus, and
possesses the taste and smell of the gas. It quickly decomposes
on exposure to air, the hydrogen of the sulphuretted hydrogen
combines with oxygen, and the liquid becomes turbid by the preci-
pitation of sulphur. Hydrogen sulphide is an inflammable gas,
burning with a bluish flame, and producing sulphur dioxide and
water —

= 2SO+2HO.

If mixed with oxygen in the proportion demanded by this equa-
tion, viz., two volumes of sulphuretted .hydrogen and three volumes
of oxygen, and ignited, the mixture explodes with violence. When
the gas is burned with an insufficient supply of air or oxygen for
its complete combustion, the sulphur is deposited.

Sulphuretted hydrogen is decomposed by the halogens, with the

4io Inorganic Chemistry

deposition of sulphur, and the formation of the hydrogen compound
of the halogen element thus —

H2S + C12 = 2HC1 + S.

Fluorine, chlorine, and bromine are capable of bringing about
this decomposition at ordinary temperatures ; in the case of iodine,
the reaction is attended with absorption of heat, which may be
supplied by passing the mixture of iodine vapour and sulphuretted
hydrogen through a hot tube, or by causing the action to take
place in 'the presence of water. In the latter case the heat of solu-
tion of the hydriodic acid determines the reaction.

When passed into sulphuric acid, reduction of the acid takes
place with the precipitation of sulphur —

H2SO4+H2S = SO2 + 2H2O + S.

Sulphuretted hydrogen, therefore, cannot be dried by means of
sulphuric acid.

The gas acts upon many metals with the formation of sulphides ;
thus, when potassium is heated in a stream of hydrogen sulphide,
it readily burns and produces potassium hydrosulphide —

Such metals as tin, lead, silver, &c., are rapidly tarnished in
contact with this gas. On this account articles of silver, when
exposed to the air of towns, quickly become covered with a film
of sulphide, which first appears yellowish-brown, and gradually
becomes black. The discoloration of a silver spoon, when intro-
duced into an egg which is partially decomposed, is due to the
same cause.

Sulphuretted hydrogen also acts upon metallic salts, combining
with the metal to form a sulphide. The " white-lead " employed
in ordinary paint is gradually blackened on prolonged exposure
to the air by the formation of lead sulphide.

Hydrogen sulphide is rapidly absorbed by lime, with the forma-
tion of calcium hydrosulphide —

CaH2O2 + 2H2S = CaH2S2 + 2H2O.

It is also absorbed by calcium sulphide, yielding the same
compound. This reaction is employed in the method known as

Hydrogen Sulphide 411

Chances process, for utilising the sulphur of the vat-waste of the
alkali manufacture. This consists in passing lime-kiln gases
through a series of vessels containing the waste mixed with water.
In the first vessels the carbon dioxide is absorbed, and sulphuretted
hydrogen evolved. This, passing into the later vessels, is absorbed
by the vat-waste, forming calcium hydrosulphide, which in its
turn is decomposed by carbon dioxide, with the evolution of twice
the volume of sulphuretted hydrogen for a given volume of carbon
dioxide, as in the first reaction —

(i)

(2) CaS + H2S=CaH2S

(3)

The sulphuretted hydrogen, mixed with atmospheric nitrogen
and a small quantity of carbon dioxide, is' sufficiently rich to burn,
yielding sulphur dioxide, which can then be employed for the
manufacture of sulphuric acid.

Sulphuretted hydrogen is also decomposed by ferric hydroxide,
with the formation of ferrous sulphide and water, and the deposi-
tion of sulphur, as described on page 401. This action takes
place with the evolution of considerable heat, the temperature
rising high enough to ignite a mixture of sulphuretted hydrogen
and oxygen.*

Sulphuretted hydrogen is a valuable laboratory reagent, on
account of the general behaviour of certain classes of sulphides.
Thus, the sulphides of certain metals, being insoluble in dilute
acids, are precipitated from acid solutions ; for example —

CuSO4+H2S =

CdCl2 + H2S = CdS + 2HCl.

Others are soluble in acids, but insoluble in alkaline liquids, and
are therefore precipitated by sulphuretted hydrogen in the presence
of ammonia, or by the addition of ammonium sulphide, thus —

ZnSO4 + (NH4)2S = ZnS + (NH4)2SO4.

A third group of metals yield sulphides that are soluble in water,

and therefore are not separated either in acid or alkaline solutions.

Many of the metallic sulphides are also possessed of charac-

* "Chemical Lecture Experiments," new ed., Nos. 577, 578.

412 Inorganic Chemistry

teristic colours, which readily serve for their identification. Thus
arsenious sulphide is pale yellow, and cadmium sulphide golden
yellow. Antimonious sulphide has a bright red colour, while zinc
sulphide is white.

This behaviour of metals towards sulphuretted hydrogen is
the basis upon which certain methods of qualitative analysis are
founded.

HYDROGEN PERSULPHIDE.

Formula, H2S2.

Modes Of Formation.— ( i.) This substance, which stands in the
same relation to hydrogen sulphide as hydrogen peroxide does to
water, may be obtained by slowly pouring a solution of calcium
or sodium disulphide into diluted hydrochloric acid cooled by a
freezing mixture, the liquids being rapidly stirred during the pro-
cess of mixing, and the acid being kept in considerable excess —

CaS2 + 2HCl = CaCl2+H2S2.

The hydrogen persulphide separates out as a heavy, pale-yellow,
oily compound, which settles to the bottom of the liquid.

Properties. — Hydrogen persulphide or hydrogen disulphide is
an oily liquid having a specific gravity of 1.376. It has a pungent
smell, accompanied by the odour of sulphuretted hydrogen, due
probably to the partial decomposition of the compound, and its
vapour is irritating to the eyes. It is an unstable substance,
decomposing at ordinary temperatures into sulphur and sulphu-
retted hydrogen ; when heated this decomposition takes place
rapidly. It is immediately decomposed by alkalis, but is more
stable in the presence of dilute hydrochloric acid. It is in-
soluble in water, but dissolves readily in carbon disulphide,
the solution in this liquid being more stable than the liquid
substance itself.

Hydrogen persulphide burns with a blue flame, yielding sulphur
dioxide and water. It possesses feeble bleaching properties, and,
like its oxygen analogue, it is decomposed by certain metallic
oxides, with the evolution of sulphuretted hydrogen.

Hydrogen Persulphide 413

Hydrogen persulphide readily dissolves sulphur, and owing to
the fact that sulphur is always liable to be precipitated along with
the persulphide in its preparation, and also to the instability of the
compound, its exact composition has been the subject of some
doubt. By subjecting the crude oil as first precipitated to careful
fractional distillation under greatly reduced pressure, Bloch and
Horm (Ber., 1908) have not only succeeded in obtaining the disul-
phide in a sufficiently pure state to establish its formula, but have
also isolated a second sulphide having the composition H2S3.
This compound, hydrogen trisulphide, is somewhat denser, specific
gravity 1.496, and less volatile than the disulphide, but otherwise
closely resembles it in properties.

COMPOUNDS OF SULPHUR WITH CHLORINE.

Two of these compounds exist at ordinary temperatures, while a
third is only known at temperatures below - 22°.

1. Disulphur dichloride or sulphothionyl chloride S2C12.

2. Sulphur dichloride ...... SC12.

3. Sulphur tetrachloride ..... SC14.

Disulphur Bichloride, S2C12. — This substance is obtained by
passing dry chlorine over the surface of heated sulphur, contained
in a retort ; the compound, which distils away as fast as it is
formed, condenses in the receiver as a yellow liquid —

S2C12.

Properties. — The redistilled liquid is an amber-coloured fuming
substance with a disagreeable penetrating odour, the vapour of
which irritates the eyes. Its specific gravity is 1.709, and it boils
at 138.1°. In contact with water it gradually decomposes into
hydrochloric .acid and sulphur dioxide, with the precipitation of
sulphur. The action takes place in two stages, thiosulphuric acid
being formed as an intermediate product, thus —

(a) 2S2C12 + 3H2O = 4HC1 + S2 + H2S2O3.
0) H2S203=H2S03 + S.

414 Inorganic Chemistry

Disulphur dichloride dissolves sulphur with great readiness, and
the solution so obtained is largely employed in the process of
vulcanising indiarubber.

This compound is the most stable of the three chlorides of
sulphur. From the fact that it contains chlorine and sulphur in
the proportion of one atom of each element, it is sometimes called
sulphur monochloride ; but as its vapour-density (67.5) shows that
it contains two atoms of each element in the molecule, the use of
the word monochloride is calculated to mislead. The name sul-
phothionyl chloride indicates its analogy to thionyl chloride, SOC12,
from which it may be regarded as being derived, by the replace-
ment of the oxygen by an atom of sulphur, O = S<^ Q ; S = S<^ ^

Sulphur Dichloride, SC12.— This compound is obtained by passing a stream
of dry chlorine into disulphur dichloride at a temperature not above o°. When
the maximum amount of chlorine is absorbed, the liquid assumes a dark
reddish-brown colour. Excess of chlorine is removed by passing a stream of
carbon dioxide through the liquid.

Sulphur dichloride rapidly dissociates with rise of temperature into free
chlorine and disulphur dichloride; at +20° this decomposition amounts to
6.5 per cent., at 50°, 24.59 Per cent., and at 100°, 80.85 Per cent- On boiling
the compound, therefore, chlorine is evolved, and the disulphur dichloride
remains behind.

In contact with water it is decomposed in the same manner as the more
stable compound.

Sulphur Tetrachloride, SC14. — This compound only exists at temperatures
below -22°, and is produced by saturating sulphur dichloride with chlorine at
that temperature. It dissociates very rapidly as the temperature rises ; thus,
at 7° above the temperature at which it is formed, viz., at -15°, this decom-
position amounts to 58.05 per cent. At -2°, 88.07 Per cent- of the compound
dissociates, while at +6.2° the percentage rises to 97.57.

The compound is decomposed by water with violence into sulphur dioxide
and hydrochloric acid —

SC14+2H20=SO2+4HC1.

Compounds of Sulphur with Bromine and Iodine have been obtained,
corresponding to S2C12. S2Br2 as a red-coloured liquid, boiling with partial
decomposition at 200° ; and S2I2 as a dark-grey crystalline solid, which melts
at a temperature about 60°.

OXIDES AND OXY ACIDS OF SULPHUR.
Four oxides of sulphur are known, namely —

(i.) Sulphur sesquioxide (hyposulphurous anhydride) S2O3.

(2.) Sulphur dioxide (sulphurous anhydride) . . SO2.

(3.) Sulphur trioxide (sulphuric anhydride) . . SO3.

(4.) Persulphuric anhydride ... > ' . . S2O7.

Sulphur Dioxide 415

These four oxides give rise respectively to the acids, hypo-
sulphurous, sulphurous, sulphuric, and persulphuric, besides
which several other sulphur acids are known —

HO'SO

Hyposulphurous acid . . H2S2O4 urvsn

HO,

Sulphurous acid . . . H2SO3 ^>SO

HO'
HOX

Sulphuric acid . . . H2SO4 >SO2

HO/
HO.

Permonosulphuric acid . . H2SO5 /SO2

HO-0/
/OH HO,

Perdisulphuric acid . . H2S2O8 SO2< >SO2

Pyrosulphuric acid (Nord- ) „ Q n Qn X°H nH°N
hausen sulphuric acid) . \ ^Wr >U2—

HOX
Thiosulphuric acid * . ' '. H2S2O3 /SO2

HS '

Besides these acids, there is a series known under the general
name of the polythionic acids. They may be regarded as being
derived from dithionic acid, which is the first of the series, by the
absorption into the molecule of various quantities of sulphur.
Four of these acids are believed to exist, viz. : —

Dithionic acid (sometimes called 1 j^ g Q HO'SO2 )

hyposulphuric acid) | 226 HO'SO2 J

Trithionic acid ... . H2S3O6 Hol82 j S<

Tetrathionic acid . • -. . H2S4O6 Holo2 i S2-

Pentathionic acid ; . , H2S6O6 HO'SO2 1 S3-

SULPHUR DIOXIDE.

Formula, SO2. Molecular weight =64. 07.' Density=32.o35.

Occurrence. — This compound is met with in the gaseous
emanations from volcanoes, and in solution in certain volcanic
springs. It is also present in the air of towns, being derived
from the combustion of the sulphur compounds present in coal.

* This acid is sometimes incorrectly called hyposulphurous acid, its sodium
salt being known as sodium hyposulphite : the so-called ' ' hypo " of the photo-
graphers.

416 Inorganic Chemistry

Modes of Formation.— ( i.) Sulphur dioxide is formed when
sulphur burns in air or oxygen —

S + O2=SO2.

At the same time small quantities of sulphur trioxide are formed,
which render the gas obtained by this combustion more or less
foggy.

(2.) Sulphur dioxide may also be obtained by heating sulphur
with a metallic peroxide, such as manganese dioxide, thus —

S2+ MnO2 = SO2+ MnS.

(3.) It is obtained when such metallic sulphides as copper pyrites
or iron pyrites are roasted in a current of air, the metal being con-
verted into oxide, thus —

2FeS2 + 11O = Fe2O3 + 4SO2.

(4.) The most convenient laboratory process for preparing sul-
phur dioxide consists in heating sulphuric acid with copper, the
final products of the reaction being copper sulphate, water, and
sulphur dioxide *—

The metals mercury or silver may be substituted for copper, but
in practice the latter metal is usually employed.

(5.) Sulphur dioxide is also formed when sulphuric acid is heated
with sulphur, the oxidation of the sulphur and the reduction of the
sulphuric acid going on simultaneously —

S + 2H2SO4 = 2H2O + 3SO2.

(6.) The reduction of sulphuric acid may be brought about by
means of carbon ; thus, if sulphuric acid be heated with carbon, the
latter is oxidised to carbon dioxide, and the acid is reduced to
sulphur dioxide —

This method is employed on a large scale for the preparation of
alkaline sulphides. The carbon dioxide which accompanies the
sulphur dioxide, not being soluble to any extent in water containing
sulphurous acid, is not in any way detrimental.

* Secondary reactions go on simultaneously, resulting in the formation of
black cuprous sulphide —

5Cu + 4H^04=3CuS04+4H90+ CuaS.

Sulphur Dioxide 417

(7.) Sulphur dioxide is formed by the decomposition of a sulphite
by dilute sulphuric acid, thus —

+ H2SO4=Na2SO4 +

Properties. — Sulphur dioxide is a colourless gas, having the
well-known suffocating smell usually associated with burning
sulphur. The gas will not burn in the air, nor will it support the
combustion of ordinary combustibles : a taper introduced into the
gas is instantly extinguished. Sulphur dioxide is more than twice
as heavy as air, its specific gravity being 2.211 (air=i). On this
account it is readily collected by displacement ; it cannot be
collected over water on account of its solubility in that liquid, but
may be collected over mercury. The solubility of sulphur dioxide
in water at various temperatures is seen by the following figures —

i vol. of water at o° dissolves 79.789 vols. SO2.

» » 20° » 39-374 „

„ „ 40° „ 18.766 „

The solution is strongly acid, and is regarded as sulphurous
acid, the gas having entered into chemical union with the water —

On cooling a saturated solution of sulphur dioxide to o°, a solid
crystalline hydrate is deposited, having the composition H2SO3,
8H20.*

When the solution is boiled the whole of the sulphur dioxide is
expelled.

Sulphur dioxide is an easily liquefied gas. At o° a pressure of
1.53 atmospheres is sufficient to condense it, while at ordinary
pressures it may be liquefied by a cold of - 10°. Its critical
temperature is 1554°.

To obtain liquid sulphur dioxide, the gas, as evolved from the
action of sulphuric acid upon copper, is dried by being passed
through a bottle containing sulphuric acid, and is then passed
through a gas-condensing tube (Fig. 113) immersed in a freezing-
mixture. The gas at once condenses in the bulb of the apparatus
as a colourless, transparent, mobile liquid, which boils at —8°.
When the liquid is cooled to — 76° it solidifies to a transparent, ice-
iike mass.

* Several hydrates of sulphurous acid have been obtained, H2SO3,6H2O;
H2SO310H2O; H2SO3.14HJ.C.

2 D

4i 8 Inorganic Chemistry

Liquid sulphur dioxide is largely employed as a refrigerating
agent, low temperatures being obtained by its rapid evaporation
under reduced pressure. The liquid dissolves phosphorus, iodine,
sulphur, and many resins. When thrown upon water a portion of
the liquid dissolves, and owing to the reduction of temperature
caused by the rapid evaporation of the remainder, a quantity of
the water is frozen. The ice so produced contains a large pro-
portion of the solid hydrate, H2SO3,8H2O.

Although sulphur dioxide is incapable of supporting the com-

FIG. 113.

bustion of ordinary combustibles, many metals will take fire and
burn when heated in the gas. Thus, when finely divided iron is
heated in a stream of sulphur dioxide it burns, forming sulphide
and oxide of the metal.

It also unites with many metallic peroxides, and often with so
much energy as to give rise to light and heat. Thus, when passed
over peroxide of lead, the mass glows spontaneously in the gas,
and lead sulphate is produced —

Sulphur Dioxide 419

Or if sodium peroxide is dusted into a cylinder of the gas, the
peroxide burns with a brilliant light, yielding sodium sulphate —
Na2O2 + SO2 = Na2SO4.

Sulphur dioxide is decomposed by the influence of strong light.
If a concentrated beam of electric light be passed through a vessel
filled with gaseous sulphur dioxide, the gas at first will appear
perfectly transparent and clear ; but in the course of a few minutes
the track of the beam will become more and more visible as it
traverses the gas, owing to the formation of thin clouds of sulphur
trioxide and sulphur, until the atmosphere of the vessel appears
to be filled with fog (Fig. 114) —

After the lapse of a short time, if the vessel be removed from

FIG. 114.

the strong light, the atmosphere will once more become clear,
owing to the reformation of sulphur dioxide.

Sulphur dioxide possesses powerful bleaching properties when
in the presence of water. Its bleaching action is due to its
absorption of oxygen from water, and consequent liberation of
hydrogen, thus —

SO2 + 2H2O = H2SO4+H2.

The hydrogen so set free reduces the colouring-matter, with the
formation of colourless compounds : the action in this case being
the reverse to that which takes place with chlorine. In some
instances, the bleaching is due to the formation of a colourless
compound, by the direct combination of sulphur dioxide with the
colouring-matter, as the original colour may often be restored by
treatment with dilute sulphuric acid, or by weak alkaline solutions.

420 Inorganic Chemistry

Thus, by passing sulphur dioxide into an infusion of rose leaves,
the red colour of the liquid is quickly discharged, but on the addi-
tion of a small quantity of sulphuric acid the colour is restored.

Sulphur dioxide is employed in bleaching materials that would
be injured by exposure to chlorine, such as straw, silk, wool,
sponge, &c., and the familiar yellow colour which gradually comes
over a sponge or a piece of bleached flannel when it is washed with
soap is an illustration of the power of alkalies to restore the original
colour to materials that have been bleached by this substance.

In the presence of water sulphur dioxide- converts chlorine into
hydrochloric acid, and on this account is employed as an " anti-
chlor •"—

In the same way it acts upon iodine, with the formation of
hydriodic acid —

SO2 + 2H2O + I2 = 2H I + H2SO4.

In the case of iodine^ however, this reaction only takes place
when a certain degree of dilution is maintained, for in a more
concentrated condition sulphuric acid is reduced by hydriodic acid
into sulphur dioxide, according to the reverse equation, thus —

It has been shown* that'aqueous sulphurous acid can only be
completely oxidised by iodine, as indicated in the former equation,
when the proportion of sulphur dioxide does not exceed 0.05 per
cent. ; when the amount exceeds this proportion the second reaction
comes into operation.

Sulphur dioxide brought into contact with iodic acid, or an
iodate, is oxidised into sulphuric acid and liberates iodine, thus —
2HIO3 + 4H2O + 5SO2=5H2SO4+I2.

This reaction is made use of as a method for the detection of
the presence of sulphur dioxide. Paper which has been moistened
with a solution of potassium iodate and starch, on exposure to
sulphur dioxide is at once turned blue, owing to the liberated
iodine combining with the starch.

The composition of sulphur dioxide may be determined by the
combustion of sulphur in a measured volume of oxygen, in the
apparatus employed for showing the volume composition of carbon
dioxide (Fig. 67). After the fragment of sulphur has burnt, and
the apparatus has been allowed to cool, it will be seen that there
is no alteration in the volume of the contained gas — the sulphur

* Bunsen.

Sulphur Trioxide 421

dioxide produced occupying the same volume as the oxygen used
in its formation. Sulphur dioxide, in other words, contains its
own volume of oxygen. One molecule, therefore, of sulphur
dioxide contains one molecule of oxygen, weighing 32. But th6
molecular weight of sulphur dioxide is 64.06 ; therefore 64.06 - 32 =
32.06 = the weight of sulphur contained in the molecule of the gas.
Sulphur dioxide, therefore, contains in the molecule one atom of
sulphur combined with two atoms of oxygen, hence its composition
is expressed by the formula SO2.

Sulphurous Acid and Sulphites.— Sulphurous acid is only
known in solution and in its hydrates. The solution smells
of sulphur dioxide, and gradually undergoes decomposition by
absorption of oxygen. The acid is dibasic, having two atoms of
hydrogen replaceable by metals ; it is therefore capable of form-
ing two series of salts, according to whether one or both of the
hydrogen atoms are replaced. Thus, by its action upon potassium
hydroxide, when the acid is in excess, the so-called acid potassium
sulphite, or hydrogen potassium sulphite, is obtained —
KHO + H2SO3=H2O + HKSO3.

Whereas, if the metallic hydroxide be in excess, the normal
potassium sulphite is formed —

2KHO + H2SO3 = 2H20 + K2SO3.

The alkaline sulphites are readily soluble in water, all other
normal sulphites being either difficult of solution or insoluble.

SULPHUR TRIOXIDE.

Formula, SO3. Molecular weight= 80.07. Vapour density =40. 03,
Modes Of Formation.— ( i.) This compound is produced when
a mixture of sulphur dioxide and oxygen is passed over heated
spongy platinum or platinised asbestos —
2SO2+02=2S03.

On leading the product through a well-cooled receiver, the sulphur
trioxide condenses in white silky needles. This method has been
successfully employed on a commercial scale. The mixture of
sulphur dioxide and oxygen is obtained by allowing ordinary strong
sulphuric acid to drop into earthenware retorts heated to bright
redness, whereby it is almost entirely broken up into these two
gases and water, thus —

H2SO4=SO2 + O + H2O.

The gases are then deprived of the water, by passage first through
a condenser, and then through a leaden tower containing coke

422 Inorganic Chemistry

moistened with sulphuric acid, and are finally passed over heated
platinised asbestos contained in glazed earthenware pipes.

(2.) Sulphur trioxide is most conveniently obtained by gently
heating pyrosulphuric acid in a glass retort. The trioxide distils
over and may be collected in a well-cooled receiver —

H2S,07=H2SO4 + SO3.

(3.) It may also be obtained by heating sodium pyrosulphate to
bright redness —

Na2S2O7 = Na2SO4 + SO3.

The sodium pyrosulphate is produced when hydrogen sodium sul-
phate (so-called bisulphatc of soda) is heated to about 300°, thus —

2HNaSO4 = H2O + Na2S2Or.

And on account of this origin it is sometimes termed anhydrous
sodium bisulphate.

(4.) Sulphur trioxide can also be produced by the action of
phosphorus pentoxide upon sulphuric acid. This most powerful
dehydrating substance withdraws from the sulphuric acid the ele-
ments of water when gently heated, thus —

The trioxide is distilled from the mixture, and the metaphosphoric
acid remains in the retort.

Properties.— Sulphur trioxide is a white, silky-looking, crystal-
line substance, which melts at 14.8° and boils at 46°. It is very
volatile, and gives off dense white fumes in contact with air, owing
to the combination of its vapour with atmospheric moisture to form
sulphuric acid. It combines with water with great energy to form
sulphuric acid ; a fragment of the compound dropped into water
dissolves with a hissing sound resembling the quenching of red-hot
iron —

S03 + H20 = H2S04.

When brought in contact with the skin, or other organic matter
containing hydrogen and oxygen, it abstracts these elements and
produces a burnt or charred effect upon the substance. Sulphur
trioxide unites directly with barium oxide, BaO, and if the baryta
be dry the mass becomes incandescent, owing to the heat of the
union, and barium sulphate is formed —

BaO + SO3 = BaSO4.

When the vapour of sulphur trioyide is passed through a red-hot
tube, it is broken down into sulphur dioxide and oxygen.

When the trioxide is heated, it melts to a colourless liquid, which

Sulphur Trioxide 423

exhibits a remarkably high rate of expansion by heat ; between
25° and 45° its mean coefficient of expansion is 0.0027, nearly three-
fourths of the expansion coefficient of a gas.

Sulphur Sesquioxide, S2OS. — A solution of this compound in fuming
sulphuric acid was obtained early in the century by heating flowers of sulphur
with Nordhausen sulphuric acid, whereby a blue solution was obtained. The
substance may be prepared by the gradual addition of dry flowers of sulphur
to melted sulphur trioxide, at a temperature just above its melting-point, when
a malachite-green crystalline solid separates out.

The compound is unstable at ordinary temperatures, being resolved into
sulphur dioxide and sulphur, the decomposition taking place rapidly upon
gently warming —

2S203=S + 3S02.

If the sesquioxide be sealed up in a bent glass tube and gently warmed, the
sulphur dioxide may be obtained liquid in one limb of the tube.

Hyposulplmrous Acid,* HaS^.— This compound was discovered by
Schutzenberger, who gave to it the formula H2SO.>. The later investigations
of Bernsthen prove that its composition is expressed by the foimula H2S2O4.
It is obtained by the reduction of sulphurous acid by means of nascent
hydrogen. Thus, when zinc is acted upon by an aqueous solution of sul-
phurous acid, no hydrogen is evolved, as the nascent gas combines with
oxygen of the acid to form water—

2H2S03+2H=

The solution so obtained has a yellowish colour, and possesses powerful
reducing and bleaching properties.

Sodium hyposulphite (Na2S2O4) may be obtained by the action of zinc upon
a cooled concentrated solution of hydrogen sodium sulphite (HNaSOs), air
being carefully excluded, a double sulphite of sodium and zinc being at the
same time produced, thus — •

4HNaSO3+Zn=Na2S204+Na2S03,ZnSO3+2H20.

The greater part of the double sodium-zinc sulphite is deposited as crystals,
the rest is removed by adding to the mother-liquor about four times its volume
of alcohol, in a closed flask. The double salt, being less soluble in alcohol
than the hyposulphite, is first deposited, and the clear liquid after being poured
off and corked up, is cooled, when it solidifies to a mass of crystals of nearly
pure sodium hyposulphite. The crystals in the wet condition are rapidly
oxidised on exposure to air ; but if quickly pressed between blotting-paper and
dried in vacuo, the dry salt is not acted upon by atmospheric oxygen.

The acid is obtained from the sodium salt by the action upon it of oxalic acid.
' Sodium hyposulphite is also formed when a solution of hydrogen sodium
sulphite is subjected to electrolysis, the nascent hydrogen developed at the
negative electrode reducing the sulphite to hyposulphite by the abstraction of
one atom of oxygen.

Properties. — Hyposulphurous acid is an extremely unstable, yellow-coloured
liquid which rapidly decomposes into sulphur dioxide, water, and sulphur —

2H2S204=3S02+2H20 + S,

a certain quantity of thiosuiphuric acid being formed as an intermediate
product at the same time—

2H2S204 = H2S203 + 2S02 + H2O.

* This compound (sometimes called hydrosulphurous acid) must not be con-
founded with thiosuiphuric acid, which is often incorrectly called hypo-
sulphurous acid (page 435).

424 Inorganic Chemistry

The acid reduces salts of silver or mercury, with precipitation of the metal,
thus —

Hga2+2H20 + H2S204=Hg+2HCl+2H2S03.

The sodium salt possesses the same bleaching and reducing powers as the
acid, and when wet or in solution it rapidly absorbs oxygen from the air and
is converted into a compound known as sodium metabisulphite—

Na.jS2O4 + O = Na2S.p5.

The relation in which these two compounds stand to each other will perhaps
be more evident if their formulae are written thus —
NaOSO ) , n_ NaO'SO ) n
NaO-SOf +U~ NaO-SO }U-

Persulphuric Anhydride, S2O7, is formed when a mixture of sulphur
trioxide and oxygen is subjected to the prolonged action of the silent electric
discharge. The compound is very unstable, breaking up even at compara-
tively low temperatures into its generators.

Perdisulphuric Acid (Persulphuric Acid], H2S2O8. — When sulphuric

+
acid is diluted with water it not only dissociates into hydrion H and sul-

phanion SO4, but also into hydrion and hydrosulphanion HSO4, the extent to
which dissociation in these two directions takes place depending upon the
degree of dilution. Within certain limits the less the acid is diluted the

larger is the proportion of HSO4 ions present

When, therefore, moderately strong sulphuric acid is electrolysed, the HSO4
ions on discharging at the anode unite to form molecules of perdisulphuric
acid (HSO4)2 or FL^CK,. The pure acid has not been isolated, its aqueous
solutions rapidly decomposing into sulphuric acid and oxygen. The acid is
dibasic, and its salts (usually called persulphates, but more correctly per-
disulphates) have the general formula M'2S.2O8.

The potassium salt is obtained by the electrolysis of a strong solution of

potassium hydrogen sulphate, which will contain the ions K, HSO4. The
acid which is first formed by the union of the discharged negative ions at the
anode, there interacts with the potassium hydrogen sulphate forming -the
sparingly soluble persulphate which separates out. The salt may be freed
from the acid sulphate by recrystallisation.

Barium perdisulphate, BaS^Og, 4H2O, is soluble in water ; 100 parts of
water at o° dissolving 52-2 parts of the salt. On this account barium chloride
gives no precipitate with a solution of a persulphate, which distinguishes these
salts from the sulphates. On warming, however, the persulphate is decom-
posed into sulphate and oxygen —

BaS2O8 + HaO = BaSO4 + H2SO4 + O .

In the solid state the persulphates are stable salts, but their aqueous solu-
tions gradually evolve oxygen and pass into sulphates. The reactions of both
the acid and its salts are therefore those of strong oxidising agents.

Permonosulphuric Acid (Cards Acid], H2SO5, is obtained when potassium

Sulphuric Acid 425

perdisulphate is treated with strong sulphuric acid and the mixture poured
upon crushed ice, both operations being conducted in a freezing mixture —
K2S.,08+ H2S04= K2S04+ H2S208.
H2S208+ H20= H2S04+ H2S05.

The relation in which these two persulphuric acids stand to each other will
be seen by a consideration of their constitutional formulae —

/OH /OH HO\

SO/ ; . S02< )S02.

\0— OH \0 - CK

Permonosulphuric acid is regarded as a monobasic acid, the "acidic" OH
group being the one directly associated with the sulphur. No pure salts have
yet been obtained.

SULPHURIC ACID.

Formula, H2SO4.

Modes of Formation. — ( i .) This acid is formed when sulphur
trioxide is dissolved in water —

SO3+H2O = H2S04.

(2.) It is also formed by the direct union of sulphur dioxide with
hydrogen peroxide —

S02+H202=H2S04.

(3.) An aqueous solution of sulphur dioxide gradually absorbs
oxygen, and is converted into sulphuric acid —

(4.) Manufacture of Sulphuric Acid.— Sulphur dioxide is un-
able to absorb an additional atom of oxygen, and so pass into
sulphur trioxide, without the aid of some third substance which
can act as a catalytic agent or a carrier of oxygen. The material
which is employed for this purpose in the process by which sul-
phuric acid is manufactured is one of the oxides of nitrogen, which
is capable of giving up oxygen to the sulphur dioxide, and of again
taking up oxygen from the air. Thus, nitrogen peroxide (NO2),
by the loss of one atom of oxygen, is reduced to nitric oxide, NO ;
which in its turn combines with atmospheric oxygen and is re-
converted into nitrogen peroxide. Therefore, when sulphur dioxide
and oxygen are mixed with nitrogen peroxide in the presence of
steam, a series of reactions takes place, the final result* of which
is that the oxygen is caused to combine with the sulphur dioxide
and water, with the formation of sulphuric acid —

The nitrogen peroxide at the end of the reaction is unchanged,
and is able to react in tjie same series of changes over and over
again, thus transforming, theoretically, an unlimited, and, in

426 Inorganic Chemistry

practice, a relatively large quantity of sulphur dioxide into sul-
phuric acid.

The series of changes that gives rise to the ultimate product is
the following : — The sulphur dioxide, nitrogen peroxide, and water
give rise, in the first place, to the formation of nitro-sulphonic acid
and a molecule of nitric oxide —

(i.) 2SO2 + 3NO2+H2O = 2H(NO)SO4

Nitro-sulphonic acid (sometimes called nitro-sulphuric acid^ and
nitrosyl sulphate) may be regarded as sulphuric acid in which one
of the hydrogen atoms is replaced by the group (NO), thus,

NO in which case the nitrogen is linked to the sulphur

by the intervention of oxygen ; or it may be considered as derived
from sulphuric acid by the replacement of one of the groups (HO)

/OH
by the group NO2, SO2<^ N^ when the nitrogen is directly

attached to the sulphur. The substance is a white crystalline
compound which in the presence of water is instantly decomposed
into sulphuric acid and a mixture of nitric oxide and nitrogen
peroxide, thus —

(2.) 2SO2(HO)(NO2) + H2O = 2H2s64-fNO + NO2.

The nitric oxide in this and the former reaction, on coming in
contact with the atmospheric oxygen, is at once reconverted into
nitrogen peroxide —

(3.) NO + 0 = N02.

In the process of the manufacture the crystalline compound
SO2(HO)(NO2) (known as chamber crystals) is not actually isolated,
unless from accidental causes the supply of water is in deficit, the
production of these crystals being regarded as an indication that
the process is not being well carried out.

The formation of sulphuric acid by these reactions, with the
intermediate production of the chamber crystals, may be carried
out on a. small scale by means of the apparatus shown in Fig. 115.
A large flask, F, is fitted with a cork, through which pass five
tubes : three of these are connected to separate two-necked bottles
containing sulphuric acid, through which can be delivered respec-
tively, nitric oxide, sulphur dioxide, and oxygen. The fourth tube
is attached to a flask in which water may be boiled, and through
which oxygen can be passed, and the fifth tube (not shown in

Sulphuric Acid

427

the figure) serves as an exit. A quantity of oxygen is first passed
into the large flask through the drying-bottle D, and sufficient
nitric oxide is then allowed to enter, to form deep red vapours ; at
the same time sulphur dioxide is passed in through the bottle S.
In order to introduce a small quantity of moisture, oxygen is
allowed to enter through the flask of boiling water, and in a few
moments large white crystals begin to form all over the interior
of the flask, and rapidly spread until the whole surface is
covered.

In order to show the second reaction in the cycle, the gaseous

FIG. 115.

contents of the flask may be swept out by means of a rapid stream
of oxygen passed in through the drying-bottle D ; and when the
atmosphere within the apparatus is colourless, a quantity of steam
is driven in from the small flask. The chamber crystals will
be seen to dissolve with effervescence, and the flask once more
becomes filled with brown fumes. The nitric oxide evolved by
the decomposition of the nitrosyl sulphate, coming in contact with
the oxygen within the flask, at once regenerates nitrogen peroxide,
in accordance with equation No. 3.

The solution formed in the flask will be found to yield a pre-
cipitate of barium sulphate, on the addition to it of a soluble
barium salt.

On a manufacturing scale, the combination of the reacting gases
and vapours which gives rise to the sulphuric acid takes place in

428 Inorganic Chemistry

large leaden chambers, usually about 100 feet long, 25 feet wide,
and 20 feet high, having therefore a capacity of 50,000 cubic feet,
several of such chambers being placed in series. Into these
chambers there is delivered sulphur dioxide, air, oxides of nitrogen,
and steam.

The plant employed for the manufacture of sulphuric acid con-
sists broadly of four parts, i. Apparatus for generating sulphur
dioxide. 2. Apparatus for producing oxides of nitrogen. 3. Appa-
ratus for absorbing oxides of nitrogen from the gases leaving the
chambers. 4. The chambers in which the reactions are made.

(i.) Pyrites Burners. — The sulphur dioxide is obtained either by
burning native sulphur, or roasting the " spent oxide " of the gas
works (see Sulphur), or by roasting pyrites, the latter being the
most general method. The pyrites burner (Fig. 116, B) is essen-
tially a small furnace or kiln in which the ore is heated, and in
which the admission of air can be duly regulated, as not only is it
necessary to admit sufficient air to completely burn the whole of
the sulphur, and so prevent any volatilisation of it in an unburnt
condition, but also to supply the requisite volume of oxygen for
the requirements of the reactions which are to go on within the
chamber. Too large a volume of air must be avoided, in order not
to unduly dilute the chamber gases.

(2.) If no loss of nitrogen peroxide took place during the cycle
of changes, the same quantity of this gas would convert an infinite
amount of sulphur dioxide and water into sulphuric acid, but in
practice, owing to leakage, defective absorption, and the reduction
of a certain percentage of this compound into nitrous oxide, it is
necessary to constantly replenish the supply. This is usually done
by generating a small quantity of nitric acid (by the action of
sulphuric acid upon nitre) in earthenware pots, which are usually
placed in an enlarged part of the flue of the pyrites burner, known
as the "nitre oven," and which is provided with a door for the
introduction of the pots (Fig. 116, N). The heated gases playing
upon these pots promotes the evolution of the nitric acid, which in
contact with sulphur dioxide is at once decomposed according to
the equation —

It is found that to make up for the loss of nitrogen peroxide,
about three to four parts of nitre are required for every 100 parts
of sulphur burnt as pyrites.

Sulphuric Acid 429

(3.) The apparatus for the absorption of the nitrogen peroxide
from the gases that are drawn from the chamber at the end of the
series is known as the " Gay-Lussac Tower" (Fig. 116, T). This
consists of a square leaden tower filled with fragments of coke,
and down which there is caused to slowly percolate a stream of
cold strong sulphuric acid, the acid being evenly spread over the
mass of coke by a special distributing arrangement. The nitrogen
peroxide is absorbed by the acid, with the formation of nitro-
sulphonic acid, SO2(HO)(NO2). In order to make use of the
absorbed nitroxygen compound, the acid which flows from the
Gay-Lussac tower is pumped to the top of another very similar
tower, situated between the " burners " and the first of the cham-
bers, and known as the " Glover Tower," G. The hot gases from
the burners, consisting of sulphur dioxide, nitrogen, and oxygen,
together with the small quantities of nitrogen peroxide from the
nitre pots, are made to pass up this tower on their way to the first
chamber, and meeting with the descending stream of nitro-sul-
phonic acid as it runs over the bricks or flints with which the tower
is filled, denitrification of the latter takes place, thus —

or 3H2SO4.

The nitric oxide thus evolved, in presence of the atmospheric
oxygen, is converted into nitrogen peroxide, and swept along with
the other gases into the chambers.

In practice it is usual to deliver down the Glover tower, besides
the nitro-sulphonic acid, a quantity of "chamber acid"** from a
separate tank. The effect of the heated gases upon this dilute
acid is to remove a portion of the water from it, thereby effecting
its partial concentration, and furnishing the water demanded by
the above equation. It will be seen, therefore, that there is a
scrubber tower at each end of the series of chambers, the
"Gay-Lussac" at the exit, where nitrogen peroxide is absorbed;
and the "Glover" at the commencement, where the dissolved
nitrogen compound is again liberated and returned to the
chambers.

(4.) The chambers are made of sheet lead, connected togethei
by fusing the edges by means cf an oxyhydrogen flame, without
the intervention of solder, as the presence of another metal gives
rise to the rapid corrosion of the lead on account of galvanic

43° Inorganic Chemistry

action being set up ; this method of joining the lead is known as
autogenous soldering. The enormous leaden chamber is supported
in a framework of wood, to which the lead is secured by bands of

FIG. 116.

B. — Double row of pyrites burners, placed
back to back : one being shown
open.

N.— Hearth where the nitre pots are
placed : one shown as open.

G. — Glover Tower, with two tanks at top:
one for the nitro-sulphuric acid de-
rived from the Gay-Lussac tower,
the other for the "chamber acid."

These acids are forced up from the

leaden vessels E, called " eggs."
C— Leaden chamber, of which there are

three shown in the figure.
p'~ pip<; conveying the gases from the

third chamber to the Gay-Lussac

tower.
T.— Gay-Lussac Tower. The tanks at the

top of this and the Glover tower

are enclosed in wooden sheds.

the same metal, and the whole is usually supported on iron or
brick pillars.

The general arrangement of a modern sulphuric acid works is

Sulphuric Acid 431

seen in Fig. 116. The gases from the double row of pyrites
burners B are led through the Glover tower G, where they
effect the denitrification of the nitro-sulphonic acid, as already
explained. From this tower they are delivered into the series of
chambers, where they meet with the necessary supply of steam.
The acid collects upon the floor of the chambers, and samples are
constantly drawn off by means of an arrangement known as a
drip-pipe, which, acting in a manner similar to a rain gauge, indi-
cates the progress of the processes going on within. The gases,
after being drawn through the entire series of chambers by means
of the draught caused by the tall chimney, are finally passed up
the Gay-Lussac tower T, where all the nitrogen peroxide is ab-
sorbed, and returned to the chambers through the intervention of
the Glover tower G, as above described.

The acid which collects in the chambers is usually not permitted
to reach a higher specific gravity than about 1.6, when it contains
about 68 per cent, of sulphuric acid ; for if the strength be allowed
to exceed this, the acid not only begins to dissolve the nitrogen
peroxide in the chamber, but exerts a corrosive action upon the
lead of which the chamber is constructed. It is therefore with-
drawn, and the first stage in the further concentration is effected
either by the action of the Glover tower, or by evaporation in
shallow leaden pans.

In order to bring up the strength of the acid to that of " oil of
vitriol," that is, to about 98 per cent., the acid from the Glover
tower or the leaden pans is heated in either glass or platinum
stills.

Sulphuric acid, unless specially purified, is liable to contain a
number of impurities, such as lead sulphate, derived from the
action of the acid upon the chamber ; arsenic, from the pyrites
employed ; oxides of nitrogen, and sulphur dioxide. From most
of the impurities, except the arsenic, the acid may be purified
by the addition of ammonium sulphate, and subsequent redis-
tillation—

4 + 2S02(HO)(NO2) =

Arsenic is removed by precipitation of the sulphide, by means
of sulphuretted nydrogen, from tne acid m a moderately dilute
state.

432 Inorganic Chemistry

(5.) The " Contact Process? As already stated, when a mixture
of sulphur dioxide and oxygen is brought into contact with finely
divided platinum, the metal acts the part of a carrier, or catalytic
agent, and causes the union of the gases. By absorbing the sulphur
trioxide so produced in water, sulphuric acid is obtained.

The chief obstacles to the successful utilisation of this re-
action on a manufacturing scale are the impurities present in
the sulphur dioxide derived from the pyrites burners, which are
found to rapidly destroy the effectiveness of the platinum. By the
system of purification now adopted these difficulties have been
removed, and the operation is being conducted on a successful
manufacturing scale.

In outline the process is the following. The mixture of sulphur
dioxide and air drawn from the pyrites burners is first passed
through a chamber called the "dust chamber," into which jets of
steam can be injected. This serves the twofold purpose of remov-
ing dust carried mechanically from the burner, and of diluting and
partially removing the sulphuric acid which is also a product of the
burner. The gases after being sufficiently cooled are then made
to pass up through a series of towers (resembling Glover towers)
where they meet a descending spray of water. They are next
dried by passing up another tower (which may be compared to the
Gay-Lussac tower), where they meet a descending stream of strong
sulphuric acid. The gases are then admitted to the contact chamber,
which consists of a vessel containing a number of small perforated
shelves upon which is spread a layer of platinised asbestos.* The
shelves are arranged one above the other in tall narrow columns
which are separated from each other in the chamber by a slight
space, the object being to prevent the mass from locally overheating
during the operation.

At the commencement the vessel is gently heated by gas jets,
but when the operation has once started external heat is with-
drawn, and care is then required to prevent the temperature rising
above about 350° (which is found to be the most favourable tempera-
ture) owing to the heat of union of sulphur dioxide and oxygen.

Properties. — Sulphuric acid is a perfectly colourless, heavy,
oily liquid. The acid obtained by distillation always contains
about 2 per cent, of water ; stronger than this it cannot be prepared

* Ferric oxide may be substituted for platinum, but the percentage yield is
smaller.

Sulphuric Acid 433

by evaporation or distillation. When, however, acid of this strength
is cooled to o°, colourless crystals of pure sulphuric acid, containing
loo per cent. H2SO4, are deposited. The crystals melt at 10.5°,
and remain liquid at temperatures much below this point. The
specific gravity of the pure acid is 1.854 at o°. When boiled it
gives off sulphur trioxide until the amount of water in it rises
to 1.5 per cent, when it distils unchanged at a temperature of
.338°.

Sulphuric acid has a powerful affinity for water, and absorbs
moisture from the atmosphere with great readiness. On this
account it constitutes one of the most valuable desiccating agents,
and is constantly made use of for depriving gases, upon which it
exerts no chemical action, of aqueous vapour. Owing to its strong
affinity for water it decomposes many organic substances contain-
ing hydrogen and oxygen, withdrawing from the compounds these
elements in the proportion to yield water ; its action upon formic
acid, oxalic acid (see Carbon Monoxide), and alcohol (see Ethy-
lene) are examples of this action.

When the acid is poured upon such substances as wood or sugar
the elements composing water are withdrawn and the carbon is
liberated, with the result that the compounds are blackened or
charred.

When sulphuric acid is mixed with water considerable heat is
disengaged, the temperature often rising to the boiling-point of
water, and at the same time a diminution in volume takes place.
The maximum contraction is obtained upon mixing the materials
in the proportion of one molecule of acid to two molecules of water.
The diminution in volume in this case amounts to 8 per cent., and
the composition of the acid produced corresponds to the formula
H2S04,2H20.

Sulphuric acid combines with water in various proportions, form-
ing a number of hydrates of a more or less definite character. The
best known are those represented by the formulae H2SO4,H2O and
H2SO4,2H2O. These compounds may be regarded as respectively
tetrabasic and hexabasic sulphuric acid, and their relation to
the ordinary dibasic acid may be expressed by the following
formulae —

H2SO4 . . . orSO2(HO)2.
H4SO5 or H2SO4,H2O „ SO(HO)4.
H6S06 „ H2S04,2H20 „ S(HO)6.

2 E

4.34 Inorganic Chemistry

Salts of each of these acids are known —

Hydrogen potassium sulphate . . HKSOX \

Normal potassium sulphate . . K2SO4 > Derived from H2SO4.

Barium sulphate BaSO4 )

Tetrabasic lead sulphate . . . Pb2SO5 ,,. „ H4SO5.

Hexabasic mercuric sulphate ) v\n^.c\ u cr*

/^ . t\ f Hg3bO6 .1 »» H6SO6.

(Turpeth mineral) j

Most sulphates are soluble in water ; those of lead, calcium, arid
strontium are only very sparingly soluble, whilst barium sulphate is
insoluble both in water and acids. The presence of sulphuric acid
or a sulphate may therefore be readily detected by the addition of
a soluble barium salt, which causes the immediate precipitation of
white barium sulphate, insoluble in hydrochloric acid.

PTBOSULPHURIC ACID (Nordhausen Add; Fuming Sulphuric Acid].
Formula, H2S2O7 or

Modes of Formation.— ( I.) This acid may be obtained by dis-
solving sulphur trioxide in ordinary sulphuric acid —

H2SO4 + SO3=H2S2O7.

On cooling the solution to o° the pyrosulphuric acid separates out
in the form of large colourless crystals.

(2.) Pyrosulphuric acid is manufactured by the distillation of
ferrous sulphate in clay retorts, mounted in series in a large
" galley " furnace. The first action of heat upon crystallised ferrous
sulphate (green vitriol) is to expel six molecules of water of crystal-
lisation, leaving the salt of the composition FeSO4,H2O. When
this substance is further heated it is decomposed finally into ferric
oxide, with the formation of sulphur trioxide, water, and sulphur
dioxide, thus —

2FeSO4,H2O = Fe2

The decomposition takes place in two stages, the sulphur dioxide
and water being evolved in the first part of the process with the
formation of ferric sulphate, which is afterwards broken up in the
manner shown in the following equation —

(i.) 6FeSO4,H2O = Fe2(SO4)3 + 2Fe2O3+3SO2 + 6H2O.
(2.) Fe2(S04)3=Fe203 + 3S03.

Thiosulphuric Acid 435

The sulphur trioxide is condensed in receivers containing either
a small quantity of water or a charge of sulphuric acid.

(3.) Pyrosulphuric acid may also be obtained by decomposing
sodium pyrosulphate (Na2S2O7), either by heating it to a high
temperature (see Sulphur Trioxide, page 422), or by acting upon it
with sulphuric acid, thus —

The sulphur trioxide obtained is dissolved in sulphuric acid, as in
the former methods ; and the hydrogen sodium sulphate, when
gently heated to about 300°, is reconverted into pyrosulphate by
the loss of a molecule of water (page 422).

Properties. — Pyrosulphuric acid is a colourless, strongly fuming
liquid, having a specific gravity of 1.88. When cooled, it solidifies
to a crystalline mass, which melts at 35°. The compound may be
regarded as consisting of one molecule of sulphuric acid plus a
molecule of sulphur trioxide, H2SO4,SO3 ; or, as being derived
from two molecules of sulphuric acid, by the withdrawal of one
molecule of water, thus —

Pyrosulphuric acid forms a stable series of salts, of which the
sodium compound already mentioned is a typical example. These
salts are sometimes spoken of as the disulphates^ and are analogous
to the dichromates (q>v.\

Two other definite compounds of sulphur trioxide and sulphuric acid are
known to exist, both of which are fuming acids. The composition of these
substances is expressed by the formulae—

or H2S4O13 ; and SH^C^SOg, or H«S4O16.

THIOSULPHURIC ACID.

Formula, H2S2O3.

This acid has never been obtained in the free state, as it decom-
poses almost as soon as liberated from its salts into sulphur dioxide
and water, with precipitation of sulphur —

436 Inorganic Chemistry

The thiosulphates, however, are stable and important salts, the
sodium salt being largely used in photography under the name of
hyposulphite of soda, or " hypo."

Modes cf Formation of Thiosulphates.— (i.) These salts
may be obtained by digesting flowers of sulphur with solutions of
the sulphites, thus —

Na2SO3 + S = Na2S2O3.

(2.) Sodium thiosulphate is also formed when sulphur dioxide is
passed into a solution of sodium sulphide. The reaction may be
regarded as taking place in three steps, in which sodium sulphite
and sulphuretted hydrogen are the first products. The latter com-
pound is then acted upon by sulphur dioxide, with the precipitation
of sulphur, thus —

SO + H2O + Na2S = Na2SO3 + H2S.

And the sulphur reacts with the already formed sulphite, as indi-
cated in the equation given above.

(3.) When sulphur is boiled with sodium hydroxide, or with milk
of lime, mixtures of sulphides and thiosulphates are obtained in
both cases —

3Ca(HO)2 + 125 = CaS2O3 + 2CaS6 + 3H2O.

The sodium sulphide can be converted into thiosulphate by the
reactions given above. Calcium pentasulphide, on exposure to air,
absorbs oxygen and forms a further quantity of thiosulphate with
precipitation of sulphur —

CaS5+3O = 3S + CaS2O3.

The thiosulphates are decomposed by most acids, with the libera-
tion of sulphur dioxide, and precipitation of sulphur. They show a
great tendency to form double salts, many of which are soluble
in water ; thus sodium thiosulphate, in contact with either silver
chloride, bromide, or iodide, forms the soluble double sodium-silver
thiosulphate, NaAgS2O3 —

Na2S203 + AgCl = NaCl + NaAgS2O8.

Trithionic Acid 437

The employment of sodium thiosulphate in photography, for
" fixing " negatives or silver prints, depends upon this property.

Thiosulphuric acid may be regarded as being derived from sul-
phuric acid by the replacement of one of the (HO) or hydroxyl
groups, by an equivalent of (HS) or hydrosulphyl —

Dithionic Acid, H2S2O6 or Ho'sQ2 }*~~This comPound is prepared by
passing a stream of sulphur dioxide through water in which manganese
dioxide is suspended, whereby manganese dithionate is formed ; while at the
same time a portion of the salt is acted upon by manganese dioxide, and con-
verted into manganous sulphate, thus —

2SO2 + MnO2 = MnS2O6.
MnS2O6+ MnO2=2MnSO4.

On the addition of barium hydroxide to the solution, barium dithionate,
barium sulphate, and manganous hydrate are formed —

MnS206 + Ba(HO)2=BaS206+Mn(HO)2.

Barium dithionate, being soluble, is separated by nitration, and upon
evaporation separates out in crystals of the composition BaS2O6,2H2O.

Upon the addition of dilute sulphuric acid in amount demanded by the
equation — '

BaS206+ H2SO4= BaSO4+ H2S2O6,

the acid itself is obtained. The solution may be concentrated in vacuo until
it reaches a specific gravity of 1.347. Further concentration results in its de-
composition into sulphuric acid and sulphur dioxide —

Dithionic acid forms well-defined crystalline salts, which on heating decom-
pose into sulphates with evolution of sulphur dioxide.

Dithionic acid was formerly called hyposulphuric acid, and its salts are still
sometimes referred to as hyposulphates.

Trithionic Acid, H2S3O6 or Hol82 }S'~The Potassium salt of this acid
may be obtained by passing sulphur dioxide through a strong solution of
potassium thiosulphate —

3S02+ 2K2S203=S + 2K2S3O(!.

It is also formed when a solution of potassium silver thiosulphate is boiled—
KO

KO-SOs

«wT—

438 Inorganic Chemistry

The sodium salt may be obtained by the addition of iodine to a mixture of
sodium sulphite and thiosulphate —

NaO i

NaS |

The acid itself is obtained by the addition of fluosilicic acid to a solution of
the potassium salt, when insoluble potassium fluosilicate is precipitated.

Both the acid itself and its salts are readily decomposed into sulphur dioxide,
sulphur, and either sulphuric acid or a sulphate, thus —

K2S3O6= K2SO4 + SO2 + S.

When acted upon by sodium amalgam, sodium trithionate is converted
back again into its generators, sodium sulphite and thiosulphate, thus—

NaO ) „,

Tetrathionic Acid, H2S4O6 or 2 s2.— The sodium salt is obtained

riO'ovA^ )

by the action of iodine upon sodium thiosulphate —

The barium salt, from which the acid itself is most readily obtained, is pre-
pared by the gradual addition of iodine to barium thiosulphate in water —

2BaS2O3 + 12= BaI2 + BaS4O6.

The barium tetratbionate is separated by the addition of alcohol, which dis-
solves the iodide and excess of iodine, leaving the tetrathionate. By the
addition of dilute sulphuric acid to an aqueous solution of this salt, in
amount demanded by the equation —

BaS4O6 + H2SO4= H2S4O6 + BaSO4,

a dilute aqueous solution of the acid may be obtained. The dilute acid may
be boiled without decomposition ; but when concentrated, it readily passes
into sulphuric acid, sulphur dioxide, and sulphur.

Sodium amalgam decomposes the sodium salt into two molecules of thio-
sulphate, reversing the reaction by which it is produced.

Pentathionic Acid, H2S5O6 or ^"f^2 }s3.— This acid is prepared by

rivJ °o(J2 '

passing sulphuretted hydrogen into a strong aqueous solution of sulphur
dioxide —

or—

5H2S03 + 5H2S = H2S506 + 5S + 9H2O.

The solution contains, however, more or less of the other thionic acids, but
as the passage of sulphuretted hydrogen is continued, these are gradually

Oxychlorides of Sulphur 439

decomposed, and ultimately the pentathionic acid also, so that the final
products of the action of excess of this gas will be sulphur and water —

H2S5O6 + 5H2S = 6H2O + 10S.

The solution obtained by the first action may be concentrated by cautious
evaporation in vacuo, until a specific gravity of 1.46 is obtained, when on partial
saturation with potassium hydroxide and fi'tration, a solution is obtained which
on spontaneous evaporation deposits crystals of potassium pentathionate, having
the composition K2S5O6,3H2O. On heating, the salt splits up into potassium
sulphate, sulphur dioxide, and sulphur.

OXYCHLORIDES OF SULPHUR.

Four of these compounds are known, all of which may be re-
garded as being derived from the oxyacids by the replacement of
hydroxyl (HO) by its equivalent of chlorine.

i. Thionyl chloride, or Cl ) so corresponding oxyacid ™ 1 SO, sulPhu;ous
Sulphurous chloride Cl J 3 HO j acid.

2. Sulphuryl chloride, or Cl )gO

. . \

Sulphuric chloride Cl J 2

IHO.U

,.~ sulphuric

3. Sulphuric chlorhydrate, or Cl ) gQ

/nor

'°2« acid.

Chlorosulphonic acid HO f "

4. Disulphuryl chloride, or C1'SO2 | n

HO'S02)(

-. pyrosulphu-

Pyrosulphuric chloride C1'SO2 J "

"HO'S02J

' ric acid.

Thionyl Chloride, SOCLj, molecular weight =118.96, is obtained by the
action of phosphorus pentachloride upon sodium sulphite —

SO( NaO)2 + 2PC15 = SOC12 + 2POC13 + 2NaCl.

It is also obtained when dry sulphur dioxide is passed over phosphorus
pentachloride —

\ SO2 + PC15 = SOC12 + POClg.

Properties. — Thionyl chloride is a colourless and highly refractive liquid
which fumes in moist air, and has a pungent unpleasant smell. It boils at 78°,
and is at once decomposed by water into its corresponding oxyacid, with for-
mation of hydrochloric acid—

SOC12 + 2H2O = H2SO3 + HC1.

Sulphuryl Chloride, SO2C12 ; molecular weight = 134. 96. This compound
(sometimes known as chlorosulphuric acid) can be obtained by the direct
union of chlorine and sulphur dioxide, under the prolonged influence of bright
sunlight—

SO2+C12=SO2C12.

It is also formed by the action of heat upon sulphuric chlorhydrate,

44O • Inorganic Chemistry

This substance, on being simply heated to 180° in sealed tubes for a lew hours,
breaks up into sulphuryl chloride and sulphuric acid —

Properties. — Sulphuryl chloride is a colourless liquid, which fumes in moist
air, and has a specific gravity of 1.66. It boils at 70°, and is decomposed by
water, with formation of sulphuric acid and hydrochloric acid —

£} }S02+2H20=2HC1+HO }SO2-

Sulphuric CMorhydrate, SO2C1(HO).— This compound is the first pro-
duct of the replacement of the (HO) groups in sulphuric acid by chlorine,
and is formed by the direct combination of sulphur trioxide and hydrochloric
acid—

SO3+HC1=HC1SO3 or SO2C1(HO).

It may be obtained by distilling sulphuric acid with phosphorus oxychloride —

2 Ho}s°2 + pOCl3=2 ci°}s02+HCl+HP03.
Or by passing dry gaseous hydrochloric acid into melted pyrosulphuric acid—

Properties. — Sulphuric chlorhydrate is a colourless fuming liquid, having a
specific gravity of 1.76, and boiling at i49°-i5i°, with partial dissociation into
its generators, sulphur trioxide and hydrochloric acid. In contact with water
it is decomposed with considerable violence, with formation of sulphuric and
hydrochloric acids —

Disulphuryl Chloride (pyrosulphuric chloride}, QSOo1 } ° or ^OgClg. This
substance is obtained by the action of sulphur trioxide, or sulphuric chlor-
hydrate, upon phosphorus pentachloride —

2S03 + PC15 = POC13 + S205C12.
2S02C1( HO) + PC15 = POC13 + 2HC1 + S^Cl,.

It is also produced by the action of sulphur trioxide upon sulphur
dichloride —

5SO3 + S2C12 = S2O5C12 + 5SO2.

Or by the action of sulph ur trioxide upon sulphuric chloride —

Carbon Disulphide 441

Properties. — Pyrosulphuric chloride is a heavy, oily, fuming liquid, resem-
bling pyrosulphuric acid in appearance. It has a specific gravity of 1.819, and
boils at 146°. When mixed with water it slowly decomposes into sulphuric
and hydrochloric acids, showing a marked difference in this respect from
sulphuric chlorhydrate —

S2O5C12 + 3H2O =2H2SO4 + 2HC1.

COMPOUNDS OF SULPHUR WITH FLUORINE.

Perfluoride of Sulphur, SFg. — This compound has been recently obtained
(Moissan) by passing fluorine over sulphur.

Properties. — Sulphur perfluoride is a colourless inodorous gas, very soluble
in water, and incombustible in air. It is a comparatively inactive compound.

Thionyl Fluoride, SOF2, is obtained by the action of fluorine upon thionyl
chloride ; also by the action of arsenic trifluoride upon thionyl chloride —

2AsF3 + 3SOC12 = 3SOF2 + AsCl3.

Properties. — Thionyl fluoride is a colourless gas which fumes strongly in
moist air. It is immediately decomposed by water—

CARBON DISULPHIDE.

Formula, CS.J. Molecular weight=76.i2. Vapour density =38. 06.

History. — This compound was accidentally produced by Lam-
padius (1796) when heating a mixture of charcoal and pyrites.

Mode Of Formation. — Carbon disulphide is prepared by passing
the vapour of sulphur over red-hot charcoal, when the two elements
unite and form the volatile product, which is condensed in vessels
surrounded with cold water —

C + S2=CS2.

The product is always contaminated with free sulphur, which
volatilises, and is also accompanied by considerable quantities of
sulphuretted hydrogen, formed by the action of sulphur upon the
hydrogen contained in the charcoal.

When carbon disulphide is prepared on a manufacturing scale,
the charcoal is heated in a vertical cast-iron or earthenware retort,
C (Fig. 117), having an elliptical section, and provided with three
openings. The retort is built into a suitable furnace, whereby it
can be uniformly heated to redness. A quantity of sulphur, con-
tained m the pot S, kept liquid by the heat of the furnace, is

442

Inorganic Chemistry

allowed to enter at intervals through the pipe B. As the vapour
comes in contact with the red-hot charcoal, combination ensues,
and the carbon disulphide escapes through the pipe D, which is
inclined to the retort so as to allow condensed sulphur to run back.
Sulphur which escapes condensation in this pipe, collects, for the
most part, in the vessel E, which is closed by water seals, as seen
in the figure. The volatile compounds are then passed through a
Liebig's condenser about 30 ft. long, and the crude disulphide so
condensed is collected in a receiver. Any vapour of carbon disulphide

which is carried on by the sul-
phuretted hydrogen is absorbed
by passing the gas through a
scrubber containing oil ; and
finally the sulphuretted hydrogen
is absorbed in a lime purifier,
similar to those employed for the
purification of coal gas. The
ashes are withdrawn from the
retort through the wide tube B ;
and the fresh charcoal is intro-
duced through the opening A.
In order to prevent the escape
of the unpleasant and injurious
vapours from A during the addi-
tion of fresh charcoal, the opening
A' is put into communication with the chimney of the furnace.
The sulphur which flows back into the retort from D is conveyed,
by means of the pipe^J nearly to the bottom of the mass of heated
charcoal, so that its vapour shall once more be made to pass over
the carbon.

At the present day, since the application of electrical heating to
manufacturing processes, the mixture, instead of being heated from
the outside by fuel, is heated inside in vessels of modified form by
means of the electric arc.

The crude product is purified by distillation and subsequent
agitation with mercury.

Properties. — Carbon disulphide is a colourless, mobile, and
highly refracting liquid. When perfectly pure it possesses a
sweetish, and not unpleasant, ethereal smell, but as usually met
with the odour is decidedly foetid.

Its specific gravity at o° is 1.292, and it boils at 46°. The vapour

FIG. 117.

Carbon Disulphide 443

of carbon disulphide has a very low igniting-point (see page 329).
It burns with a blue flame, which, when fed with oxygen, emits a
dazzling blue light. When carbon disulphide vapour is mixed
with three times its volume of oxygen, and a light applied, the
mixture explodes with violence ; the products of the combustion
being carbon dioxide and sulphur dioxide —

The vapour of carbon disulphide, when constantly inhaled in
small quantities, has an injurious effect upon the health, and if
breathed in large quantities is a powerful poison.

When heated to a bright red heat, carbon disulphide vapour is
decomposed into its constituent elements : on this account, in the
manufacture of this compound, care is taken that the temperature
does not rise too high.

The vapour of carbon disulphide is decomposed by potassium,
which, when heated, burns in the vapour, forming potassium
sulphide, and liberating carbon —

CS2 + 2K2 = C + 2K2S.

When passed over heated slaked lime, carbon disulphide vapout
is converted into carbon dioxide and sulphuretted hydrogen —

CS2 + 2CaH2O2=2CaO + CO2 + 2H2S.

This reaction is made use of for converting the carbon disul-
phide, which is always present in coal gas, into the two easily
removed substances, carbon dioxide and sulphuretted hydrogen.

When a mixture of carbon disulphide vapour and sulphuretted
hydrogen is passed over heated copper, marsh gas is formed —

Carbon disulphide is soluble to a minute extent in water ; i
volume of water dissolves .001 volume of this liquid, and the
solution possesses the taste and the smell of the disulphide. It
mixes in all proportions with alcohol, ether, the hydrocarbons of
the benzene family, and most essential oils. It also dissolves
sulphur, phosphorus, iodine, bromine, caoutchouc, and most fats ;
and is largely used in the arts, both as a solvent for caoutchouc
and in extracting essential oils, spices, and perfumes.

Thioearbonie Acid.— Carbon disulphide is the sulphur ana-
logue of carbon dioxide, CS2 ; CO2. Like the oxygen compound,

Inorganic Chemistry

It forms a feeble acid, which has received the name thiocarbonic
acid, H2CS3 ; carbonic acid, H2CO3.

The thiocarbonates are produced by reactions analogous to
those by which carbonates are formed. Thus, when carbon disul-
phide is brought into contact with potassium sulphide, potassium
thiocarbonate is obtained —

Thiocarbonates are likewise formed by the action of carbon
disulphide upon metallic hydroxides —

The acid itself is obtained as a yellow oil, having an unpleasant
odour by the decomposition of a thiocarbonate by dilute hydro-
chloric acid.

A large number of compounds are known in which divalent
sulphur replaces oxygen, and which therefore stand in the same
relation to the oxygen compounds as thiocarbonic acid stands to
carbonic acid ; for example —

Thiocarbamic acid, CS2,NH3, or J^g2 | CS ;
Carbamic acid, CO2,NH3, or ^2 j- CO.

Other Compounds of Carbon and Sulphur.— When carbon disulphide is
exposed to the influence of light, there is gradually formed upon the glass
vessel containing it a brown deposit, which is believed to be carbon mono-
sulphide, CS ; the sulphur analogue of carbon monoxide. When electric
sparks from carbon poles are passed thrc agh the vapour of carbon disulphide,
or when the electric arc is produced in the vapour, an offensive smelling liquid
is obtained, which exerts a most irritating and tear-producing effect upon the
eyes. This liquid has been shown to have the composition C3S2.*

SELENIUM.

Symbol, Se. Atomic weight =79.1. Molecular weight =158. 2.

History. — This element was discovered by Berzelius (1817), who gave it
the name selenium (signifying the moon) on account of its close analogy with
the previously discovered element tellurium (signifying the earth).

Occurrence.— Selenium is occasionally met with associated with native
sulphur, probably as a selenide of sulphur. In a few minerals of considerable

* Von Lengyel, 1894.

Selenium 445

rarity, selenium is met with in the form of selenides of such metals as mercury,
lead, silver. It occurs in very small quantities in a large number of metallic
sulphides.

Modes Of Formation.— ( i.) When pyrites containing selenium is employed
in the manufacture of sulphuric acid, the selenium is oxidised by the atmos-
pheric oxygen into selenium dioxide, which is carried forward with the sulphur
dioxide. Selenium dioxide, being a solid substance, is partly deposited in the
flues, and in the Glover tower, and partly carried forward into the chambers,
where it forms a red-coloured deposit. To obtain the selenium, either the flue
dust or the chamber deposit is first boiled with dilute sulphuric acid, and either
nitric acid or potassium chlorate added, in order to oxidise it completely into
selenic acid, H2SeO4. The solution is then boiled with strong hydrochloric
acid, whereby'it is reduced to selenious acid, H2SeO3, when a stream of sulphur
dioxide is passed through it which precipitates the selenium as a red powder —

H2SeO3 + 2SO2 + H20 = Se + 2H2SO4.

(2.) A second method for the preparation of selenium from the chamber
deposit consists in digesting the substance with potassium cyanide, whereby
it is converted into soluble potassium selenocyanide, SeK(CN). On the
addition of hydrochloric acid to this solution, the element is precipitated as a
red amorphous powder, and hydrocyanic acid and potassium chloride go into
solution —

SeK(CN) + HCl=Se + KCl+H(CN).

Properties. — Selenium is known in various allotropic modifications.

1. Soluble in carbon disulphide. — a. Brick-red amorphous powder, obtained
by precipitation with acids, or reduction of selenious acid, in the cold, by
sulphur dioxide.

|8. Black crystalline powder, obtained by reduction of hot selenious acid by
sulphur dioxide.

•y. Dark red translucent monoclinic crystals, specific gravity 4.5, deposited
from solution in carbon disulphide.

5. Black, shining, brittle amorphous mass, having a conchoidal fracture,
and a specific gravity of 4.3, obtained by rapidly cooling melted selenium.

2. Insoluble in carbon disulphide. — Black, metallic-looking crystalline mass,
having a granular fracture. Obtained by quickly cooling melted selenium to
210° and keeping it for some time at that temperature, when the mass solidifies
with rise of temperature to 217°. This insoluble variety, sometimes called
metallic selenium, is also formed as a deposit of minute black crystals, when
concentrated solutions of sodium or potassium selenide are exposed to the air.

This modification has a specific gravity of 4.5, and melts at 217°.

Selenium boils at 680°, forming a dark-red vapour which condenses in the
form of flowers of selenium, having a scarlet-red colour.

At high temperatures the vapour of selenium, like that of sulphur, becomes
a true gas ; thus at 1420°, the vapour-density is found to be 81.5, approaching
very closely to the normal density demanded by the molecule Se2.

"Metallic" selenium conducts electricity, and the element exhibits the
remarkable property of having its conductivity increased by light ; the con-
ductivity of selenium when exposed to diffused daylight being about twice as

446 Inorganic Chemistry

great as when in the dark. This alteration in the electrical resistance with
varying intensities of light, is a property of selenium that was made use of in
the construction of an instrument known as the photophone, but it has not as
yet been put to any practical use. When selenium is heated in the air, it
burns with a blue flame, with the formation of selenium dioxide, and at the same
time emits a powerful and characteristic smell resembling rotten horse-radish.

When selenium is heated in a tube filled with an indifferent gas, it sublimes
in the form of a red deposit ; but when heated in hydrogen, the sublimate is in
the form of black shining crystals. The formation of these crystals is due to
the fact that selenium combines with the hydrogen, and the hydrogen selenide
is again decomposed by the heat.

Hydrogen Selenide (selenuretted hydrogen], H2Se ; molecular weight=8i. 12.
Hydrogen selenide is formed when selenium is heated in hydrogen.

This compound is also obtained by the action of dilute hydrochloric or sul-
phuric acid upon either potassium selenide or ferrous selenide —

FeSe + H2SO4 = FeSO4 + H2Se.

Properties. — Hydrogen selenide is a colourless gas, strongly resembling
sulphuretted hydrogen, both in its smell and in its chemical behaviour. It is
readily soluble in water, and when passed through metallic solutions precipi-
tates insoluble selenides of most of the heavy metals. Hydrogen selenide
burns with a blue flame, with the production of water and selenium dioxide.
Its smell, although resembling that of its sulphur analogue, is more unpleasant,
and its effects upon the system are more persistent and injurious. A single small
bubble inhaled through the nostril produces temporary paralysis of the olfac-
tory nerves, accompanied by inflammation of the mucous membrane.

No compound of selenium corresponding to hydrogen disulphide is known.

COMPOUNDS WITH HALOGENS.

Diselenium Bichloride, S^Og, is obtained by passing chlorine over
selenium, or by passing gaseous hydrochloric acid through a solution of
selenium in nitric acid.

Properties. — Selenium chloride is a brown oily liquid, in which selenium
itself is readily soluble, and from which the element is deposited in the form
which is insoluble in carbon disulphide. It is slowly decomposed by water,
thus —

2Se2Cl2 + 3H2O = H2SeO3 + 3Se + 4HC1.

Corresponding bromine and iodine compounds are known, Se2Br2, and
Se2I2.

Selenium Tetrachloride, SeCl4, is prepared either by the action of chlorine
upon selenium chloride —

Se2Cl2+3Cl2=2SeCl4,

or by heating a mixture of selenium dioxide and phosphorus pentachloride—
3Se02+ 3PCl5=3SeCl4 + P2O5 + POC13.

Properties. —Selenium tetrachloride is a white, crystalline, volatile com-
pound; which may be sublimed without decomposition and without fusion.

Selenic Acid 447

When the vapour is heated above 200° it begins to dissociate into selenium
and chlorine. It dissolves in water, with decomposition into hydrochloric and
selenious acids — '

l4 + 3H2O=

Corresponding bromine and iodine compounds are known, SeBr4 and Sel

OXIDES AND OXYACIDS OF SELENIUM.

Only one oxide of selenium is known, namely, selenium dioxide, SeO2,
although a second oxide of unknown composition is believed to exist, and to
constitute the peculiar smelling substance which is always formed when
selenium is burnt in the air.

Selenium Dioxide is prepared by burning selenium in a stream of oxygen
in a glass tube ; the element burns in the gas with a blue flame, and the oxide
condenses upon the distant portions of the tube, as a white crystalline deposit.

Properties. — Selenium dioxide crystallises in long white prisms, which when
heated readily sublime without passing through the state of liquidity. It dis-
solves in water and gives rise to selenious acid.

The following oxyacids of selenium are known —

Selenious acid, H2SeO3, corresponding to sulphurous acid, H2SO3.
Selenic acid, H2SeO4, corresponding to sulphuric acid, H2SO4.
Selenosulphuric HO ) crx corresponding ( thiosulphuric HO ) e/->
acid HSeP°2. to } acid> HS;b°2'

Selenious Acid, H2SeO3, is obtained as a white crystalline compound, when
the dioxide is dissolved in hot water, and the solution allowed to cool. The
acid is dibasic, and forms both acid and normal selenites, corresponding to
the sulphites : it also forms a series of so-called superacid salts, containing a
molecule of the acid salt combined with a molecule of acid, thus —

HKSe03,H2SeO3.

Selenic Acid, H2SeO4.— This acid is best prepared by the addition of
bromine to silver selenite suspended in water, when insoluble silver bromide is
formed and selenlc acid is left in solution —

Ag2SeO3 + H2O + Br2= 2AgBr + H2SeO4.

The solution may be evaporated by heating until it contains 94 per cent, of
selenic acid, and still further evaporated in vacuo until it reaches 97.4 per cent.,
when its specific gravity is 2.627. When heated to 280° it decomposes into
selenium dioxide, water, and selenium.

Properties. — Selenic acid in its most concentrated condition is a colourless,
strongly-acid liquid, which mixes with water with the development of con-
siderable heat. It dissolves iron and zinc with evolution of hydrogen ; and when
heated dissolves copper with formation of selenious acid.

The selenates closely resemble the sulphates. Barium selenate, like the

Inorganic Chemistry

sulphate, is quite insoluble in water, but differs from that compound in being
converted by boiling hydrochloric acid into barium selenite, which is soluble.

Selenium al-^» r^i^ a compound with oxygen and chlorine, selenium oxy-
chloride, or ^leriyl chloride, SeOCLj, corresponding with thionyl chloride,
SOC12.

TELLURIUM.

Symbol, Te. Atomic weight* =127.5,

Occurrence. — In the free state small quantities of this element have been
found as crystals, consisting of almost pure tellurium. In combination it is
met with in a few rare minerals, such as iellurite (TeO2), and, more commonly,
tetradymite (Bi2Te3). Some specimens of pyrites contain small quantities of
this element, hence it is found in the deposit from the vitriol chambers, from
which selenium is obtained.

Mode of Formation. — Tellurium is obtained from bismuth telluride, Bi2Te3,
by fusion with an intimate mixture of sodium carbonate and carbon. The
mass on treatment with water yields a solution containing a mixture of sodium
telluride and sodium sulphide, which on exposure to the air deposits tellurium
as a grey powder. The element is purified by distillation in a stream of
hydrogen.

Properties. — Tellurium is a bluish-white, silver-like solid, possessing metallic
lustre. It conducts heat and electricity, although badly, and is very brittle.
Its specific gravity is 6.26, and it melts at 452°. When melted tellurium is
slowly cooled, it forms rhombohedral crystals. When heated in the air it burns
with a blue flame, and forms tellurium dioxide, TeO2. When heated in a
sealed tube with hydrogen, tellurium sublimes in the form of brilliant prismatic
crystals.

Hydrogen Telluride (telluretted hydrogen), . H2Te. — When tellurium
is heated in hydrogen the elements combine, forming hydrogen telluride,
which exhibits the same phenomenon as is shown by hydrogen selenide of
being decomposed by heat, and depositing the element as a crystalline
sublimate.

Hydrogen telluride is obtained by the action of hydrochloric acid upon zinc
telluride—

ZnTe + 2HC1 = ZnCl2 + H2Te.

Properties. — Hydrogen telluride is a most offensive smelling and highly
poisonous gas. It behaves like sulphuretted hydrogen in precipitating metals
from solutions. It is soluble in water, and the solution gradually absorbs
oxygen and deposits tellurium.

* Various numbers have been obtained by different observers for the
atomic weight of tellurium. Some of these numbers are higher than the atomic
weight of iodine, which would make it impossible to give to tellurium a posi-
tion between antimony (atomie weight=i2o) and iodine (atomic weight=
126.92) as demanded by the periodic law. Brauner, who has spent many
years investigating this point, considers that hitherto pure tellurium has nevef
been obtained. The most recent determinations give the number 197.5.

Telluric Acid 449

COMPOUNDS WITH THE HALOGENS.

Two chlorides of tellurium are known, namely, tellurium dichloride, TeCl2,
and tellurium tetrachloride, TeCl4. It will be noticed that the composition of
the dichloride is not analogous with the lower chloride of either selenium
(Se2Cl2) or sulphur (^CLj).

Two bromides, TeBr2 and TeBr4, and corresponding iodides are known.

OXIDES AND OXYACIDS OF TELLURIUM.

Two oxides of tellurium are known with certainty, namely, tellurium dioxide,
TeO2, and tellurium trioxide, TeO3, which give rise respectively to the two
acids, tellurous acid, H2TeO3, and telluric acid, H2TeO4.

Tellurous Acid is obtained by pouring a solution of tellurium in nitric acid
into an excess of water. The acid is precipitated as a white amorphous
powder. When strongly heated it is converted into the dioxide and water.

Tellurous acid, like sulphurous acid, is dibasic, and gives rise to both acid
and normal salts : thus, with potassium it forms hydrogen potassium tellurite,
HKTeO3, and dipotassium tellurite, K2TeO3. It also forms superacid salts
such as —

Quadracid potassium tellurite . . . HKTeO3,H2TeO3.
Potassium tetratellurite .... K2TeO3,3TeO2.

Telluric Acid is prepared by fusing either tellurium or tellurium dioxide
with a mixture of potassium nitrate and carbonate —

Te + K2C03+2KN03=2K2Te04+N2+CO.

The fused mass, after solution in water, is mixed with a solution of barium
chloride, which precipitates barium tellurate ; this is then decomposed by the
addition of the exact amount of sulphuric acid, and after nitration the clear
solution deposits crystals of telluric acid, H2TeO4,2H,O. When these crystals
are heated to 160° the water is expelled, and the anhydrous acid in the form cf
a white powder is left. On strongly heating, telluric acid decomposes into
water and tellurium trioxide, which at a higher temperature splits up into the
dioxide and oxygen.

Like tellurous acid, telluric acid forms not only normal and acid salts, but a
number of more complex superacid salts —

Normal potassium tellurate " " '. . . K2TeO4,5H2O.

Hydrogen potassium tellurate . i - .. HKTeO4.

Quadracid potassium tellurate . . . K2TeO4,H2TeO4,3H2O.

Potassium tetratellurate .... K2TeO4,3H.,TeO4,H2O.

CHAPTER III
THE ELEMENTS OF GROUP V. (FAMILY B.)

Nitrogen, N . . 14.01
Phosphorus, P . .31.0
Arsenic, As . . 75.0

Antimony, Sb . . 120.2
Bismuth, Bi . . 208

IN this family of elements we have a gradual transition from the
non-metals to the metals. Nitrogen and phosphorus may be con-
sidered as typical non-metallic elements, both as regards their
physical and chemical properties. The third member, arsenic,
begins to exhibit metalline properties ; its specific gravity is more
than three times as high as that of phosphorus, and it possesses
considerable metallic lustre ; arsenic is called a metalloid on this
account. Antimony is still more metallic in its character, possess-
ing most of the physical attributes of a true metal, while in bismuth
all non-metallic properties cease altogether to exist.

All these elements form more than one compound with oxygen,
of which the following may be compared —

N203 ; (P203)2 ; (As2O3)2 ; Sb2O3 ; Bi2O3.
N204 ; P204 ; Sb204 ; Bi2O4.

N205 ; P206 ; As2O5 ; Sb2O5 ; Bi2O5.

The oxides (which in the case of nitrogen and phosphorus are
strongly acidic in their nature, combining with water to form acids)
gradually become less and less acidic and more basic as the series
is traversed.

Thus, nitrogen pentoxide, N2O5, unites violently with water to
form nitric acid, which with bases yields nitrates. Antimony pent-
oxide is insoluble in water, and no antimonic acid has been isolated,
although its salts, the antimonates, are known. The oxides of
antimony, on the other hand, begin to exhibit basic properties and
unite with acids, forming salts in which the antimony functions as

the base.

450

Phosphorus 451

In the case of the last element the acidic nature of the oxides is
entirely lost ; no bismuth compounds being known corresponding
to antimonates or arsenates, while these oxides unite with acids in
the capacity of bases, giving rise to bismuth salts.

Four of the elements of this group unite with hydrogen, forming
similarly constituted compounds, NH3, PH3, AsH3, SbH3.

The stability of these compounds gradually decreases as we pass
from nitrogen to antimony. Antimony hydride has never been ob-
tained free from other gases, while no similar bismuth compound is
known. Ammonia is alkaline and strongly basic, and unites readily
with acids to form ammonium salts. Phosphorus hydride has no
alkaline character, and is only feebly basic. It combines, however,
with the halogen acids to form phosphonium chloride, bromide,
and iodide, PH4C1, PH4Br, PH4I, analogous to ammonium chloride,
bromide, and iodide. The hydrides of arsenic and antimony ex-
hibit no basic character. All the elements of this group unite with
chlorine, giving rise to the compounds —

NC13, PC13, AsCl3, SbCl3, BiCl3,

which also exhibit a gradation in their properties ; thus, nitrogen
trichloride is an extremely unstable liquid, exploding with extra-
ordinary violence upon very slight causes, while the analogous
bismuth compound is a perfectly stable solid.

The boiling-points of these compounds show a gradual increase
with the increasing atomic weight of the element ; thus, nitrogen
chloride boils at 71°, ph'osphorus trichloride at 78°, arsenic tri-
chloride at 130.2°, and antimony trichloride at 200°.

The elements arsenic, antimony, and bismuth are isomorphous,
and their corresponding compounds are also isomorphous.

The first member of this family, namely, nitrogen, has been
already treated in Part II. as one of the four typical elements
studied in that section of the book. It occupies a position in
relation to the other members of the family very similar to that of
oxygen towards sulphur, selenium, and tellurium.

PHOSPHOEUS.

Symbol, P. Atomic weight = 31.0. Vapour density =62.0.
Molecular weight =124.0.

History. — Phosphorus was first discovered by the alchemist
Brand of Hamburg (1669), who obtained it by distilling a mixture

45 2 Inorganic Chemistry

of sand with urine which had been evaporated to a thick syrup.
The process, however, was kept secret. Robert Boyle (1680) dis-
covered the process of obtaining this element, but the method was
not published till after his death. Until the year 1771, when
Scheele published a method by which phosphorus could be ob-
tained from bone ash, this element was looked upon as a rare
chemical curiosity. The name phosphorus was not first coined for
this element : it had been in previous use to denote various sub-
stances known at that time, which had the property of glowing in
the dark. To distinguish the element it was called Brand* s phos-
phorus, or English phosphorus.

Occurrence. — Phosphorus is not found in nature in the free
state.* In combination with oxygen and metals, as phosphates,
it is very widely distributed, especially as calcium phosphate.
The following are some of the commonest natural phosphates —

Sombrerite, or estramadurite . Ca3(PO4)2.
Apatite . . * .... . . 3Ca3(PO4)2,CaCl2.

Wavellite . ^,^ . , 2A12(PO4)2,A12(HO)6,9H2O.

Calcium phosphate is present in all fertile soils, being derived
from the disintegration of rocks : the presence of phosphates in
soil has been shown to be essential to the growth of plants. From
the vegetable it passes into the animal kingdom, where it is chiefly
present in the urine, brain, and bones. Bones contain about 60
per cent, of calcium phosphate, to which they entirely owe their
rigidity.

Mode Of Formation.— Manufacture. The chief source of
phosphorus is bone ash, a material obtained by burning bones,
and which consists of nearly pure calcium phosphate, Ca3(PO4).,.
Other varieties of calcium phosphate, such as sombrerite and
apatite, are also employed, as well as phosphates of other metals,
such as the Redonda phosphates, which consist of phosphates of
iron and alumina. The bone ash, in fine powder, is first decom-
posed by means of sulphuric acid, specific gravity 1.5 to 1.6. This
operation is performed in large circular wooden vessels, resem-
bling a brewer's " mash tun," provided with an agitator, and into
which high pressure steam can be driven. Finely-ground bone ash
and sulphuric acid, in charges of a few cwts. at a time, are alter-
nately stirred into the decomposer, until from four to five tons of

* Farrington (Am. Jour. Science, vol. xv., 1903) records having discovered
small quantities of free phosphorus enclosed, or occluded in a meteorite.

Phosphorus

453

phosphate have been introduced, with sufficient acid to convert the
whole of the lime into calcium sulphate, according to the equa-
tion—

Ca3(PO4)2 + 3H2SO4 = 3CaSO4 + 2H3PO4.

The contents of the decomposer are next run out into filtering
tanks, and the phosphoric acid is then concentrated to a syrup, in
large lead-lined pans through which steam-coils pass, the liquor
being constantly agitated by a mechanical stirrer.

The concentrated liquor is next mixed, either with sawdust, or
with coarsely-ground charcoal, or coke, and the mixture com-
pletely dried by being heated in a cast-iron pot, or in a muffle, to a

Fi(5. 118.

dull red heat. During this process the tribasic phosphoric acid
(or orthophosphoric acid), H3PO4, is converted by loss of water
into metaphosphoric acid, HPO3 —

H3PO4=H2O + HPO3.

The charred mixture is then distilled in bottle-shaped retorts of
Stourbridge clay, about 3 feet long, and having an internal diameter
of 8 inches. A number of these retorts, usually twenty-four, are
arranged in two tiers, in a galley furnace, as seen in section in Fig.
1 1 8. The empty retorts are first gradually raised to a bright red
heat, and a charge of the mixture is then quickly introduced. Bent

454

Inorganic Chemistry

pieces of 2-inch malleable iron pipe are then luted into the
mouths of the retorts connecting them with the pipes, D D'.
These pipes dip into troughs of water, E E', which run along the
entire length of the furnace, and in which the phosphorus con-
denses. The temperature of the furnace is then raised to a white
heat, when decomposition of the metaphosphoric acid commences,
and phosphorus begins to distil over. The process is continued
for about sixteen hours. The change that goes on is mainly
represented by the following equation—

The crude product, which is usually dark red or black in appear-

FIG. 119.

ance, is first melted under hot water and thoroughly stirred, in
order to allow the greater part of the rougher suspended matters to
rise to the surface. The mass is then allowed to resolidify. The
exact processes by which phosphorus is further purified on a manu-
facturing scale are guarded as trade secrets ; one method that has
been in use consists in treating the phosphorus while melted under
water with a mixture of potassium dichromate and sulphuric acid,
whereby some of the impurities are oxidised and others are caused

Phosphorus 4f> $

to rise to the surface as a scum, leaving the phosphorus as a dear
liquid beneath.

Since the advent of the electric furnace, phosphorus is now being
manufactured direct from calcium phosphate by a process which
threatens to entirely supersede the method of distillation already
described. The calcium phosphate is mixed with carbon and
simply heated in the electric furnace. At the high temperature of
the electric arc the calcium phosphate is decomposed, the calcium
uniting with the carbon to form calcium carbide, while the phos-
phorus in the state of vapour escapes along with carbon monoxide
by the pipe P (Fig. 119), and is condensed in suitable condensers —

The molten calcium carbide is tapped off from the furnace from
QC

time to time as fresh charges of phosphate and carbon are
introduced.

Phosphorus usually comes into commerce either in the form of
wedges or as sticks. The operation of casting the phosphorus into
sticks is performed beneath water. A quantity of phosphorus
beneath a shallow layer of water is placed in the vessel C (Fig. 120),
which is contained in a tank of water through which a steam-coil
passes. Connected to the phosphorus reservoir is a glass tube, G,
which passes into a second shallow tank of cold water. On open-
ing the cock D, the liquid phosphorus flows into the cold glass tube
where it congeals, and it may then be drawn through as a continuous
rod of phosphorus if care be taken not to draw it out faster than it
solidifies. It is the custom to adopt a uniform length and thick-
ness of stick, namely, 7^ inches long and \ inch diameter. Nine
such sticks weigh I Ib.

Properties. — When freshly prepared and kept in the dark,
phosphorus is a translucent, almost colourless, wax-like solid.
Even in the dark it soon loses its transparency and becomes

45 6 Inorganic Chemistry

coated with an opaque white film ; while if exposed to the light
the film that forms becomes first yellow, then brown, and in time
the phosphorus assumes a red and even a black colour throughout
its entire mass. Its specific gravity at 16° is 1.82. At o° phos-
phorus becomes moderately brittle, and a stick of it may be readily
snapped, when its crystalline character will be seen. At 1 5° it
becomes soft, and may be cut with a knife like wax. Phosphorus
melts under water at 43.3°, and the liquid exhibits the property of
suspended solidification. If the melted material, which has been
cooled below its solidifying point, be touched with a fragment of
phosphorus upon the end of a capillary glass tube, the mass
instantly congeals with rise of temperature.*

Phosphorus contained in a closed vessel without water melts
at as low a temperature as 30°,! and when heated in air to 34° it
takes fire. At a temperature of 269° phosphorus boils and forms
a colourless vapour.

Phosphorus is volatile at ordinary temperatures : if a small
quantity of phosphorus be sealed in a vacuous tube, and the tube
be placed in the dark, the phosphorus will slowly vaporise ; and if
one end of the tube be kept slightly cooler than the rest, the phos-
phorus will sublime upon that part in the form of brilliant, colour-
less_, and highly refracting rhombic crystals, which retain their
beauty as long as . they are kept in the dark. The density of
the vapour of phosphorus is 62.0, giving a molecular weight of
124.0, which is four times the atomic weight, showing that the
molecule of phosphorus contains four atoms. Even at temperatures
as high as 1040° these tetratomic molecules are stable, but it has
been shown that at high temperatures dissociation begins to take
place.

On account of its ready inflammability, phosphorus is always
preserved under water, which exerts practically no solvent action
upon it. It is extremely soluble in carbon disulphide, i part of this
liquid dissolving 9.26 parts of phosphorus. On evaporation, the
element is deposited in the form of colourless crystals. Phosphorus
is also soluble, but to a less extent, in chloroform, benzene, turpen-
tine, alcohol, olive oil, and many other solvents. A solution of
phosphorus in carbon disulphide, when allowed to evaporate upon
a piece of blotting-paper, leaves the element in so finely divided a
condition, that its rapid oxidation almost immediately raises the

* See "Chemical Lecture Experiments," new ed., Nos. 528, 529.
f Read man.

Phosphorus

temperature to the ignition point of the phosphorus, when it
takes fire.

On exposure to moist air in the dark, phosphorus appears faintly
luminous, emitting a pale greenish-white light, and at the same
time evolving white fumes which possess an unpleasant, garlic-like
smell, and are poisonous. These fumes consist mainly of phos-
phorus oxide, P4O6, and the glowing of the phosphorus is the
result of its oxidation ; phosphorus does not glow when placed
in an inert gas which is perfectly free from admixed oxygen,
although the presence of very small traces of free oxygen in such a
gas is sufficient to cause the phosphorescence. At a few de-
grees below o°, phosphorus ceases to glow in the air. Although
the glowing is due to oxidation, phosphorus does not appear
luminous in pure oxygen at temperatures below about 15°. If,
therefore, a stick of phosphorus which is glowing in the air,
be immersed in a jar of oxygen, its phosphorescence is at once
stopped. If, however, the oxygen be slightly rarefied, the phos-
phorus again becomes luminous. Similarly, the phosphorescence
that is exhibited in air is stopped if the air be compressed.* The
glow of phosphorus is believed to be associated with the formation
of ozone, for the presence in the air of traces of such gases and
vapours as ethylene, turpentine, or ether, which are known to
possess the power of destroying ozone, at once stops the glowing
of a stick of phosphorus.

Phosphorus is incapable of uniting with oxygen if the gas be
perfectly pure and free from aqueous vapour. It has been shown
that in oxygen which has been dried by prolonged exposure to the
desiccating action of phosphorus pentoxide, phosphorus may not
only be melted, but even distilled, without any combination with
the oxygen taking place.

If water, beneath which is a small quantity of melted phosphorus,
be boiled, the phosphorus vaporises with the steam, and renders
the steam luminous : use is made of this property, as a means
of detecting free phosphorus, in toxicological analysis.

Phosphorus is a powerfully poisonous substance ; in large doses
it causes death in a few hours, in smaller quantities it produces
stomachic pains and sickness, usually ending in convulsion.
Persons constantly exposed to the vapours arising from the hand-
ling of phosphorus, either in its manufacture or in the manufacture
of matches, are very liable to suffer from caries of the bones of the
* " Chemical Lecture Experiments," new ed., Nos. 530 to 534.

Inorganic Chemistry

jaw and nose ; it is believed that this injurious effect is caused by
the white fumes which are the product of oxidation, and not by
the actual vapour of phosphorus.

Red Phosphorus. — When phosphorus is heated to a tempera-
ture between 240° and 250°, out of contact with air, it passes into
an allotropic modification.* The same transformation takes place
when phosphorus is heated to 200° with an extremely small pro-
portion of iodine.

Red phosphorus is manufactured by heating ordinary phosphorus

FIG. lax.

in a cast-iron pot, provided with a cover, through which passes a
short open pipe, D (Fig. 121). The pot is carefully and uniformly
heated to between 240° and 250°, as indicated by the thermometers
C C', which are encased in metal tubes, to prevent the phosphorus
from attacking the glass. A small quantity of the phosphorus
becomes oxidised by the air within the vessel, but after this atmos-
pheric oxygen is used up, no further oxidation takes place. If the

* There is diversity of opinion as to whether ordinary red phosphorus is an
individual allotrope, or a mixture of two modifications.

Red Phosphorus 45 0

temperature be allowed to rise above 260°, the red phosphorus is
reconverted into the ordinary modification, and with the evolution
of so much heat, that unless the open tube be provided, as a safety-
valve, the iron vessel is liable to burst. The material that is
obtained at the end of the operation is in the form of hard, solid
lumps, which still contain a certain amount of the unchanged phos-
phorus mixed with them. It is first ground to powder beneath water,
and then boiled with a solution of sodium hydroxide (caustic soda),
to remove the ordinary phosphorus, and finally washed and dried.

Properties. — Red phosphorus, as usually sent into commerce,
is a chocolate-red powder, having a specific gravity of 2.25. It
is not luminous in the dark, and has no taste or smell. It is
not poisonous, and when taken into the system is excreted f un-
changed. It is not soluble in carbon disulphide, or in any of the
solvents which dissolve ordinary phosphorus. Red phosphorus is
unaffected by exposure to dry air or oxygen, but in the presence
of moisture it is very slowly oxidised. If red phosphorus which
has been perfectly freed from ordinary phosphorus, and carefully
washed and dried, be exposed to air and moisture, it is found after
the lapse of some time to have become acid, owing to slight oxida-
tion into phosphoric acid. When heated in contact with air, red
phosphorus does not ignite below a temperature of 240°. Red
phosphorus may be obtained in the form of rhombohedral crystals
by heating the substance underpressure to a temperature of 580°.

The chief use of phosphorus is in the manufacture of matches.
When ordinary phosphorus is employed, the bundles of wooden
splints are first tipped with melted paraffin wax, and afterwards
dipped into a paste, made of an emulsion of phosphorus, chlorate
of potash, and glue. Matches so made ignite when rubbed upon
any rough surface ; the paraffin (which is sometimes replaced by
sulphur) serving to transmit the combustion from the phosphorus
to the wood. Since the discovery of red phosphorus, and its non-
injurious properties, the old phosphorus match has been largely
superseded by the so-called safety matches. In these matches the
splints are tipped with a mixture of potassium chlorate, potassium
dichromate, red lead, and antimony sulphide, and they are ignited
by being rubbed upon a prepared surface consisting of a mixture
of antimony sulphide and red phosphorus. Although these matches
will not ignite by ordinary friction upon any but the specially
prepared surface, they may be inflamed by being swiftly drawn
along a sheet of ground glass or strip of linoleum.

466 Inorganic Chemistry

COMPOUNDS OF PHOSPHORUS WITH HYDROGEN.

Three compounds of phosphorus and hydrogen are known,
namely —

PH3 (gaseous) ; P2H4 (liquid) ; and (P4H2)3 (solid).

GASEOUS HYDROGEN PHOSPHIDE (Phosphoretted Hydrogen :
Phosphine).

Formula, PH3. Molecular weight =34. 03. Density =17.01 5.

Modes of Formation.— ( i.) This substance is formed when red
phosphorus is gently heated in a stream of hydrogen.

(2.) It may be prepared by boiling phosphorus with a solution of
potassium or sodium hydroxide —

= 3NaH2PO2

In this reaction a small quantity of the liquid hydride (P2H4) is
simultaneously formed, which imparts to the gas the property of
spontaneous inflammability. It also contains a certain quantity
of free hydrogen, produced by the action of the caustic alkali upon
the sodium hypophosphite, thus —

To obtain the gas by this method, a quantity of a strong solution
of caustic soda, and a few fragments of phosphorus, are placed in
a flask, fitted as shown in Fig. 122. A stream of coal gas is passed
through the apparatus, in order to displace the air, and the solution
is gently heated. Hydrogen phosphide is readily disengaged, and
as each bubble escapes into the air, it bursts into flame, and forms
a vortex ring of white smoke of phosphoric acid.

If alcoholic potash be substituted for the aqueous solution in
this reaction, the liquid hydrogen phosphide is dissolved in the
alcohol, and the gas which is evolved is therefore not spontaneously
inflammable.*

(3.) Hydrogen phosphide is also produced by the action of
water upon calcium phosphide —

= 6Ca(HO)

* See "Chemical Lecture Experiments," new ed., No. 545.

Hydrogen Phosphide 461

A secondary reaction, by which liquid hydrogen phosphide is
formed, goes on simultaneously —

P2Ca2 + 4H2O = 2Ca(HO)2 + P2H4.

The gas, therefore, that is evolved is spontaneously inflammable.

(4.) Pure gaseous hydrogen phosphide may be prepared by the
action of potassium hydroxide upon phosphonium iodide —

Properties. — Gaseous hydrogen phosphide, or phosphine, is a
colourless gas, having an offensive smell resembling rotten fish.
It is not spontaneously inflammable, but its ignition temperature is

FIG. 122.

below 100° C. (see p. 329). The gas burns with a brightly luminous
flame, producing water and metaphosphoric acid —

PH3 + 2O2=HPO3+H2O.

When burnt in oxygen the flame is extremely dazzling.

The gas is not acted upon by oxygen at ordinary temperatures
and pressures, but if a mixture of these gases be suddenly rarefied,
combination at once takes place with explosion. Hydrogen phos-
phide is decomposed by chlorine or bromine, a jet of the gas
spontaneously igniting when introduced into chlorine or the vapour
of bromine —

The gas is also decomposed by iodine, but in this case the
action is less energetic, and a portion of the hydriodic acid which

462 Inorganic Chemistry

is produced unites with the phosphine and forms phosphonium
iodide, thus —

(i.) PH3 + 3I2=PI3 + 3HI.

(2.) PH3+HI = PH4I.

Phosphine is a highly poisonous gas, and the inhalation of even
small quantities of it is attended with injurious effects. It is
slightly soluble in water, and imparts its own smell and an un-
pleasant taste to the liquid. The solution decomposes after a
short time, especially in the light, and deposits red phosphorus.

Hydrogen phosphide has no action upon either litmus or tur-
meric paper, but it resembles its nitrogen analogue, ammonia,
in combining with hydrochloric, hydrobromic, and hydriodic acids,
forming respectively phosphonium chloride, bromide, and iodide.

Phosphonium Chloride, PH4C1.— When a mixture of phos-
phine and gaseous hydrochloric acid is passed through a tube
immersed in a freezing-mixture, the gases unite and form a white
crystalline incrustation upon the tube. If the tube be afterwards
sealed up, the compound may be sublimed from one part of the
tube to another, when it crystallises in large, brilliant, transparent
cubes. If the tube be opened, the compound rapidly dissociates
into its two generators. This compound may also be obtained by
subjecting a mixture of the two gases to pressure. Under a pres-
sure of about eighteen atmospheres at the ordinary temperature,
crystals of phosphonium chloride are deposited ; and as the pres-
sure is released the crystals gradually dissociate again.

Phosphonium Bromide, PH4Br.— Hydrogen phosphide com-
bines with hydrobromic acid at ordinary temperatures and pres-
sures, but as the compound begins to dissociate at the ordinary
temperature, the combination is only completely brought about by
cooling the gases. Phosphonium bromide may be readily pre-
pared by passing the two gases into a flask immersed in a mode-
rate freezing-mixture. The salt may be obtained in the form of
large transparent cubical crystals by sublimation in a sealed
vessel.

Phosphonium Iodide, PH4L— This compound may be obtained
by a method similar to that given for the preparation of the bro-
mide. It is also produced when hydrogen phosphide is passed
over iodine, as already mentioned. It is most readily prepared by
the action of water upon a mixture of phosphorus and iodine. For
this purpose ten parts of phosphorus are dissolved in carbon disul-

Liquid Hydrogen Phosphide 463

phide in a tubulated retort, to which seventeen parts of iodine are
gradually added, the retort being kept cold. The carbon disul-
phide is then distilled off from a water-bath, a stream of carbon
dioxide being passed through the apparatus towards the end of
the distillation to assist in expelling .the last traces of the disul-
phide.

Six parts of water are then gradually introduced from a dropping
funnel, when a brisk action takes place, and the phosphonium
iodide produced is volatilised, and may be condensed in a long
wide glass tube connected to the retort. Hydriodic acid is at the
same time formed —

The phosphonium iodide condenses in the form of brilliant
quadratic prisms.

Liquid Hydrogen Phosphide, P2H4. — This compound is
obtained in small quantities when phosphorus is boiled with a

FIG. 123.

solution of caustic soda. It is obtained in large quantities by the
decomposition of calcium phosphide with water by the reaction
already mentioned. In order to collect the compound a quantity
of calcium phosphide is introduced into a flask provided with a
dropping funnel and exit tube. After displacing the air from the
apparatus by an inert gas water is gradually introduced from the
funnel, and the products of the reaction, after passing through a
small empty tube where water is arrested, are passed through a
U-tube immersed in a freezing-mixture, where the liquid hydrogen
phosphide condenses.

Properties. — Liquid hydrogen phosphide is a colourless, highly
refracting, and spontaneously inflammable liquid. On exposure
to light, or when brought into contact with granulated calcium
chloride, it is quickly decomposed into the gaseous and the solid
hydrides of phosphorus —

464 Inorganic Chemistry

The formation of a spontaneously inflammable gas by the action
of water upon calcium phosphide has found a practical application
in the marine appliance known as " Holmes' signal." This con-
sists of a tin canister filled with lumps of calcium phosphide. A
metal tube, closed at the bottom with a piece of block tin, enters the
canister from "below, and a short cone of the same soft metal is
soldered upon the top. When the signal is to be used it is securely
fixed into a wooden float (Fig. 123). The cone is cut off and a
hole punctured through the bottom of the tube below, and the
apparatus thrown into the sea. The hydrogen phosphide spon-
taneously ignites and burns with a large brilliant flame from the
top of the tin, illuminating a considerable area.

Solid Hydrogen Phosphide, (P4H2)3 or P]2H6, is a yellow
powder, obtained, as already mentioned, by the spontaneous
decomposition of the liquid compound. Recent determinations*
of its molecular weight prove that its molecular composition is
expressed by the formula P12H6.

COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS.
Phosphorus combines with all the halogen elements, forming the
following compounds —

PF3 PC13 PBr3 PI3.

PF5 PC15 . PBr5 P2I4.

Phosphorus Trifluoride, PF3, is obtained by the action of
arsenic trifluoride upon phosphorus trichloride —

AsF3 + PC13 = PF3 + AsCl3.

It is more conveniently prepared by gently heating a mixture of
zinc fluoride and phosphorus tribromide —

Properties. — Phosphorus trifluoride is a colourless, pungent-
smelling gas. It has no action upon glass in the cold, but when
heated it forms silicon fluoride and phosphorus. It is moderately
soluble in water. Phosphorus trifluoride unites directly with
bromine, forming the compound PF3Br2.

Phosphorus Pentafluoride, PF6.— This compound is formed
when phosphorus burns in fluorine. It is best prepared by the
action of arsenic trifluoride upon phosphorus pentachloride —
5AsF3 + 3PCl5 = 3PF6 + 5AsCl3.
* Schenck and Buck, Berichte, 1904.

Phosphorus Pentachloride 465

Properties. — Phosphorus pentafluoride is a heavy, colourless
gas, which fumes strongly in moist air, being decomposed by water
into hydrofluoric and phosphoric acids —

PF5 + 4H2O = 5HF + H3P04.

Owing to this decomposition it has a pungent and irritating
effect upon the mucous membrane.

It is not acted upon by oxygen, but unites directly with dry
gaseous ammonia, forming a white solid compound having the
composition 2PF5,5NH3.

Phosphorus pentafluoride is an extremely stable compound, being
capable of withstanding a very high temperature without dissocia-
tion. On this account it is of special interest, as affording an
example of a compound in which phosphorus is united to five
monovalent atoms to form a stable substance. The corresponding
chlorine and bromine compounds readily dissociate, when heated,
into compounds containing trivalent phosphorus and the free
halogen.

Phosphorus Trichloride, PC13.— This compound is prepared
by passing dry chlorine over red phosphorus, gently heated in a
tubulated retort. The two elements readily combine, and the
volatile trichloride, mixed with more or less of the pentachloride,
distils off, and is collected in a well-cooled receiver. The product
is freed from the higher chloride by redistillation over ordinary
phosphorus.

Properties. — Phosphorus trichloride is a colourless, mobile
liquid, which boils at 75.95°. It has a pungent smell, and fumes
strongly in moist air. Water at once decomposes it into hydro-
chloric and phosphorous acids —

Phosphorus Pentaehloride, PC15.— This compound is formed
when phosphorus burns in excess of chlorine. It is prepared by
the action of chlorine upon the trichloride. Dry chlorine is passed
on to the surface of a quantity of the trichloride, contained in a
flask which is kept cool. The absorption of the chlorine is attended
with considerable rise of temperature, and the contents of the flask
rapidly become converted into a dry, pale-yellow solid.

Phosphorus pentachloride is conveniently obtained by passing

2 G

466 Inorganic Chemistry

chlorine through a solution of phosphorus in carbon di sulphide,
the solution being kept cold.

Properties. — Phosphorus pentachloride is a yellowish-white,
crystalline solid, having a pungent and irritating odour. It fumes
strongly in contact with moist air, being decomposed by moisture
into hydrochloric acid and phosphorus oxychloride—

PC15+H2O = 2HC1 + POC13.

With excess of water, both phosphorus oxychloride and phos-
phorus pentachloride dissolve with evolution of heat, forming
hydrochloric and phosphoric acids — •

POC13 + 3H2O = H3P04 + 3HC1.
PC15 + 4H2O = H3P04 + 5HC1.

Phosphorus pentachloride readily sublimes, without melting, at
a temperature below that of boiling water. It can only be melted
by being heated under pressure to a temperature of 148°.

As the vapour of phosphorus pentachloride is heated, the com-
pound dissociates into phosphorus trichloride and free chlorine.
At 300° this dissociation is complete, and the vapour consists of
equal molecules of the trichloride and chlorine. The gradual
breaking down of the molecules of pentachloride is seen from the
following table, which gives the densities of the gas at different
temperatures —

Temperatures 182° 200° 250° 300°

Density . . 72.5 69.2 57.0 52.06

At 300° it consists of molecules of PC13 (molecular weight =
137.35) and molecules of chlorine (molecular weight = 70.90) in
equal numbers, which theoretically gives the molecular weight —

137

Phosphorus pentachloride is an important chemical reagent, in-
asmuch as by its action upon oxyacids, both inorganic and organic,
the (HO) group in the acid can be replaced by chlorine. Thus
with sulphuric acid, chlorosulphuric acid is formed —

HO] S02 + PC16= HQ! j SOa+POCl3

Phosphorus Pentachloricte 467

With acetic acid it yields acetyl chloride —

}

It also effects the replacement of (HO) by chlorine, in alcohols.
Thus, with ethyl alcohol (spirits of wine) it forms ethyl chloride —

CH8+POCl

Phosphorus Tribromide, PBr3, is best prepared by dropping bromine upon
an excess of red phosphorus. It forms a colourless pungent-smelling liquid,
which boils at 172.9°.

Phosphorus Pentabromide, PBr5, is prepared by adding bromine to the
tribromide. It is a yellow solid, which melts to a reddish liquid. It is very
unstable, being dissociated below 100° into its generators, the tribromide and
bromine.

Diphosphorus Tetriodide (phosphorus di-iodide], P2I4.— This substance is
prepared by the gradual addition of 8.2 parts of iodine to i part of phosphorus
dissolved in carbon dismphide. On gently distilling off the carbon disulphide,
the iodide is left as a yellow crystalline solid. The compound melts at 110°.

Phosphorus Tri-iodide, PI3, is obtained by employing a larger proportion of
iodine in the above reaction. It is a solid substance, crystallising in red six-
sided prisms, which are decomposed by water into hydriodic and phosphorus
acids.

OXY AND THIO COMPOUNDS OF PHOSPHORUS
AND THE HALOGENS.

The following compounds are known, containing phosphorus
combined with the halogens, and either oxygen or sulphur —

POF3; POC13; P2OSC14 ; POBrCl2 ; POBr3.
PSF3; PSC13; — PSBr3.

These compounds may be regarded as derived from the haloid
compounds, by the replacement of two atoms of the halogen by an
equivalent of oxygen or of divalent sulphur j or they may be viewed
as derivatives of phosphoric acid, by the substitution of halogen
elements in the place of (HO) groups. The tribasic phosphoric
acid, PO(HO)3, may be regarded as giving rise to the compounds
POF3, POC13, &c. ; while the compound P2O3C14 may be viewed
as a derivative of pyrophosphoric acid, P2O3(HO)4.

468 Inorganic Chemistry

Phosphoryl Fluoride (phosphorus oxyjluoride), POF3, may be obtained by
the action of phosphoryl chloride (POC13) upon zinc fluoride —

3ZnF2 + 2POC13 = 2POF3 + 3ZnCl2.

It may also be prepared by gently heating a mixture of finely powdered
cryolite and phosphorus pentoxide —

2(AlF8,3NaF) + 2P2O5 = 4POF3 -f A^Og + 3Na2O.

Phosphoryl fluoride is a colourless gas, which fumes in the air, and is de-
composed by water. The gas in a dry condition does not attack glass.

Thiophosphoryl Fluoride, PSF3, is most readily prepared by gently heat-
ing in a leaden tube a mixture of dry lead fluoride and phosphorus penta-
sulphide—

+ P2S5=

The gas may be collected over mercury.

Thiophosphoryl fluoride is a colourless gas, which spontaneously inflames
when a jet of it is allowed to escape into the air. It burns with a pale greenish
non-luminous flame. In pure oxygen the gas burns with a yellow and more
luminous flame. The gas is decomposed by heat into phosphorus fluoride,
phosphorus, and sulphur. When heated in a glass vessel, sulphur and phos-
phorus are deposited, and silicon tetrafluoride is formed —

4PSF3 + 3Si = 3SiF4 + 4P + 4S.

Phosphoryl Chloride (phosphorus oxychloride), POC13.— This
compound is formed by the first action of water upon phosphorus
pentachloride (g.v.). It is also obtained when phosphorus penta-
chloride and pentoxide are heated together in a sealed tube —

3PC15 + P2O5 = 5POC13.

It is most readily prepared by heating phosphorus pentachloride
with either oxalic acid or boric acid, thus —

O3. •

Properties. — Phosphoryl chloride is a colourless fuming liquid,
which boils at 107.23.° When cooled to about - 10° it solidifies to
a white crystalline mass, which melts at 0.8°. It is decomposed by
water with formation of tribasic phosphoric acid and hydrochloric
acid —

Oxides and Oxy acids of Phosphorus 469

Pyrophosphoryl Chloride, P2O3C14, is obtained by passing nitrogen peroxide
through phosphorus trichloride, and subsequently distilling the liquid. The
reaction is complicated, and cannot be expressed by a single equation ; nitro-
gen is evolved, and phosphorus pentoxide, nitrosyl chloride, and phosphoryl
chloride are simultaneously formed. Pyrophosphoryl chloride is a colourless
fuming liquid, boiling between 210° and 215°. It is decomposed by water,
and forms hydrochloric acid and orthophosphoric acid (not pyrophosphoric
acid}—

P203C14 + 5H40=2H3P04+4HC1.

It is converted by phosphorus pentachloride into phosphoryl chloride —
P203C14+PC15=3POC13.

Thiophosphoryl Chloride, PSC13, is prepared by heating a mixture of
phosphorus pentasulphide and pentachloride —

3PC15+P2S5=5PSC13.

It is a colourless liquid, boiling at 125°. It fumes in moist air, being de-
composed by water into sulphuretted hydrogen, phosphoric and hydrochloric
acids —

PSC13-MH2O=H2S+H3PO4+3HC1.

OXIDES AND OXY AC IDS OF PHOSPHORUS.

Four compounds of phosphorus and oxygen are known, all of
which are formed when phosphorus is burned in a limited supply
of air —

Phosphorus monoxide . .' * :. . 7>4 O?

Phosphorous oxide (phosphorus trioxide) . P4O6.

Phosphorus tetroxide . . . • P2O4.

Phosphoric oxide (phosphorus pentoxide) . P2O5.

The two compounds, phosphorus trioxide and pentoxide, are the
best known of these oxides, and they give rise respectively to
phosphorous and phosphoric acids. The following oxyacids of
phosphorus are known —

Corresponding
Oxide.

Hypophosphorous acid H3PO2

Phosphorous acid . . H3PO3 or P(HO)3 . P4O6

Orthophosphoric acid . H3PO4 „ PO(HO)3 }

Pyrophosphoric acid . H4P2O7>, P2O3(HO)4 P2O5

Metaphosphoric acid . HPO? ,, PO^HO)

470

Inorganic Chemistry

When phosphorus is dissolved in a solution of aqueous alcoholic potash, and
dilute hydrochloric acid is added, a yellow or reddish precipitate is obtained
which was believed to have the composition P4O. Recent investigations, how-
ever, seem to prove that the substance so obtained is identical with red phos-
phorus. (Chem. Soc, Journal, Nov. 1899, and Nov. 1901.)

Phosphorous Oxide (phosphorous anhydride}, P4O6 ; molecular
weight = 220. — This oxide is obtained, mixed with a large excess
of the pentoxide, when phosphorus is burned in a tube through
which a regulated stream of air is passed. In order to obtain the
compound in a state of purity, the following method is employed.
A quantity of phosphorus is introduced into a glass tube bent in the

FIG. 124.

manner indicated in Fig. 124, and fitted into one end of a long,
wide, Liebig's condenser. Into the end of the condenser nearest to
the U-tube there is introduced a loose plug of glass wool, which
serves to arrest the pentoxide. The phosphorus is ignited at the
open end of the glass tube, and a stream of air drawn through
the apparatus by means of an aspirator. A stream of water, at
60°, is circulated through the condenser, when the easily fusible
phosphorous oxide is carried over, and condenses in the U-tube,
which is immersed in a freezing-mixture.

Properties. — Phosphorous oxide, as it collects in the cooled
tube, is a snow-white crystalline solid, which melts at 22.5° to a
colourless liquid. The liquid solidifies at 2 1° to a white, waxy-looking
mass, consisting of monoclinic prismatic crystals. The liquid boils
at I73-1- It possesses an unpleasant garlic smell, and is highly
poisonous. Phosphorous oxide is only very slowly acted upon by

Phosphoric Oxide

471

cold water, which gradually dissolves it, forming phosphorous
acid —

P4O6 + 6H2O=4H3PO3.

In contact with hot water a violent action takes place, in which
spontaneously inflammable phosphoretted hydrogen is evolved, and
a red deposit, consisting of red phosphorus, is formed.

When heated in a sealed tube to a temperature of 440°, phos-
phorous oxide is decomposed into phosphorus
tetroxide, and red phosphorus —

2P406=3P204 + 2P.

When exposed to air or oxygen, phosphorous
oxide is gradually oxidised into phosphorus
pentoxide, but when placed in warm oxygen it
bursts into flame. When brought into chlorine
it also spontaneously inflames.

Phosphorus Tetroxide, P2O4. — This substance is
obtained when phosphorous oxide is heated in a sealed
tube to a temperature of 440°. It forms brilliant trans-
parent ciystals, which appear as a sublimate in the tube.
This oxide is highly deliquescent, and dissolves in water
with evolution of heat.

FIG. 125.

Phosphoric Oxide, P2O5 (or P4O10).— This oxide is the main
product of the combustion of phosphorus in air or oxygen. It may
readily be obtained by burning a quantity of phosphorus in a small
capsule, and covering the whole with a large bell-jar (Fig. 125).
The white clouds of phosphoric oxide (or phosphorus pentoxide)
collect as a soft snow-like substance.

Properties. — Phosphorus pentoxide is a white, amorphous, and
very voluminous powder. It is without smell, although as usually
prepared it frequently possesses a slight garlic odour, owing to the
presence of phosphorous oxide. At a temperature short of a red
heat this oxide vaporises, and recent determinations of its vapour-
density point to the conclusion that the compound under these
conditions has a composition expressed by the formula P4O10.

Phosphoric oxide is extremely hygroscopic, absorbing moisture
from the air with great rapidity. It must therefore be preserved
either in well-fitting stopper bottles or in hermetically sealed
vessels. Its affinity for water constitutes it the most useful de-
siccating agent known to chemists : prolonged exposure to phos-

472 Inorganic Chemistry

phoric oxide removes the last traces of aqueous vapour from
gases.

When thrown into water, phosphoric oxide is dissolved with a
hissing sound resembling the quenching of hot iron, and forms
metaphosphoric acid —

P205 + H20 = 2HP03,

which gradually passes into orthophosphoric acid —
HPO3 + H2O = H3PO4.

Phosphoric oxide reacts with a number of substances, both
inorganic and organic, removing oxygen and hydrogen from them
in the proportion in which these elements form water ; thus, it
converts nitric acid into nitrogen pentoxide —

2HN03-H2O = N2O5.

In the same way it withdraws the elements of water from alcohol,
with the evolution of ethylene —

C2H6O-H2O = C2H4.

Hypophosphorous Acid, H3PO2.— This acid is prepared by
the action of sulphuric acid upon the barium salt —

Ba(H2PO2)2 + H2SO4 = BaSO4 + 2H3PO2.

The solution, after the removal of the barium sulphate by filtra-
tion, is gently heated until its temperature rises to 130°, when it will
be sufficiently concentrated to deposit crystals when cooled to o°.

The barium hypophosphite is obtained by boiling phosphorus
with a solution of barium hydroxide —

3Ba(HO)2 + 8P + 6H2O = 2PH3 + 3Ba(H2PO2)2.

Properties. — Hypophosphorous acid is a white crystalline com-
pound which melts at 17.4°. When strongly heated it is converted
into orthophosphoric acid, with the evolution of gaseous hydrogen
phosphide —

2H3PO2=H3P04+PHo.

Hypophosphorous acid acts as a powerful reducing agent, on
account of the readiness with which it absorbs oxygen and is con-
verted into orthophosphoric acid —

HO) H) HO)

=HO[PO; or HO [-P + O2 = HO Y

HOj HOJ HOJ

Phosphorous Acid 473

Thus if hypophosphorous acid or the sodium salt in solution be
added to a solution of copper sulphate, and the mixture gently
warmed, the copper is reduced even a stage further than to the
metallic state, and a dark red-brown precipitate of copper hydride^
Cu2H2, is obtained, thus —

> = 3H3PO4+H2SO4
+ 3HNaSO4+2Cu2H2.

This constitutes a characteristic reaction for hypophosphites.
Hypophosphorous acid is a feeble monobasic acid ; its salts with
monovalent metals being represented by the formula MH2PO2.

It is customary to express the basicity of oxyacids by the number of (HO)
groups that are contained in the molecule, and as this acid is monobasic its
constitution would be expressed by the formula POH2(HO). Many of the
oxyacids of phosphorus, however, show a tendency to exhibit a lower degree
of basicity than is represented by the number of (HO) groups they contain ;
thus, orthophosphoric acid, PO(HO)3, which is tribasic, and forms the salt
trisodium phosphate, PO(NaO)3, holds the third atom of the metal so loosely
that even such a feeble acid as carbonic acid is capable of expelling it —

PO(NaO)3+CO2+H2O=PO(HO)(NaO)2+HNaCO3.
or —

Na3PO4 + CO2+ H2O=HNa2PO4+ HNaCO3.

The weaker acid, phosphorous acid, is also tribasic, P(HO)3, and forms
trisodium phosphite, P(NaO)3, or Na3PO3. But this salt is even decomposed
by water, into the disodium phosphite, P(HO)(NaO)2, or HNa^PO^

Hypophosphorous acid being a still weaker acid, its acidic power is destroyed
as soon as one atom of hydrogen is replaced by a base, and its constitution may,

in harmony with these facts, be expressed by the formula PH(HO)o, or HO

Phosphorous Acid, H3PO3, or P(HO)3.— As already mentioned,
this acid is formed when phosphorous oxide is dissolved in cold
water.

It is most readily prepared by the action of water upon phos-
phorus trichloride —

PC13 + 3H2O = 3HC1 + P(HO)3.

The production and decomposition of the phosphorus trichloride
may be carried on simultaneously, by passing a stream of chlorine
through phosphorus which is melted beneath water. The solution
is evaporated until its temperature rises to 180°, when the liquid

474 Inorganic Chemistry

will have become so concentrated that on cooling it solidifies to a
crystalline mass.

Properties. — Phosphorous acid is a white crystalline substance
which melts at 70.1°. When heated, it decomposes into ortho-
phosphoric acid, with evolution of hydrogen phosphide —

4H3P03 = 3H3PO4 + PH3.

Like hypophosphorous acid, this compound absorbs oxygen, and
therefore is a powerful reducing agent ; silver salts are reduced
to metallic silver, and mercuric salts are reduced to mercurous
salts. By the absorption of oxygen it is converted into ortho-
phosphoric acid —

H3PO3 + O = H3PO4.

Although a tribasic acid, its tribasic salts are unstable ; the
sodium compound, Na3PO3, which is the most stable inorganic
salt, is decomposed by water into the dibasic salt —

Na3PO3 + H2O = HNa2PO3 + NaHO.
NaCM HO ^

or NaO ^P + H2O = NaO lP + NaHO.

NaO J NaO J

Orthophosphorie Aeid, H3PO4, or PO(HO)3.— This acid is
formed when phosphorus pentoxide is dissolved in boiling water,
or when the solution of the oxide in cold water is boiled —

P2O5+3H2O = 2H3PO4.

Orthophosphoric acid is readily obtained by the oxidation of red
phosphorus with nitric acid. Copious red fumes are evolved, and
phosphoric acid remains in solution.

Phosphoric acid is prepared on a large scale by the action of
sulphuric acid upon bone ash, as in the process for the manu-
facture of phosphorus —

Ca3(PO4)2 + 3H2SO4=3CaSO4+2H3PO4.

The calcium sulphate is removed by filtration, and the solution
evaporated to a syrup. Prepared in this way the acid usually
contains arsenic. This is removed by first reducing it to arsenious
oxide by means of sulphur dioxide, and after boiling off the excess
of sulphur dioxide, precipitating the arsenic as sulphide by means
pf sulphuretted, hydrogen.

Pyrophosphoric . A cid 47 5

Properties. — The solution obtained by these methods is either
concentrated in vacuo or heated to a temperature of 140°, and
allowed to cool, when the acid is obtained in the form of trans-
parent six-sided prismatic crystals belonging to the rhombic
system. The substance is deliquescent, and melts at 38.6°.

Phosphoric acid is trfbasic, and forms three series of salts,
according as one, two, or three of its hydrogen atoms are replaced
by an equivalent of metal. Thus, with the metal sodium the three
salts are known —

Dihydrogen sodium phosphate .... H2NaPO4.
Hydrogen disodium phosphate . .. . . HNa2PO4.
Trisodium phosphate (normal sodium phosphate) Na3PO4.

The hydrogen may be replaced by its equivalent of more than
one base. Thus, the well-known compound, microcosmic saltt is
hydrogen sodium ammonium phosphate, HNa(NH4)PO4,4H2O.
The salt, which is precipitated when magnesium sulphate (in the
presence of ammonium chloride and ammonia) is added to a solution
of a phosphate, consists of the compound ammonium magnesium
phosphate (NH4)MgPO4,6H2O.

The heavy metals usually only form normal phosphates. Thus,
on the addition of silver nitrate to a solution of either of the three
sodium salts, the same silver salt is precipitated, namely, tri-
argentic phosphate.

Na3P04 + 3AgNO3 = Ag3PO4 + 3NaNO3.
*HNa2PO4+3AgNO3 = Ag3P04 + 2NaNO3 + HNO3.
= Ag3P04+NaNO3 +2HNO3.

Pyrophosphorie Acid, H4P2Or, or P2O3(HO)4.— This acid is
derived from orthophosphoric acid by the withdrawal of one
molecule of water from two molecules of the acid. This change is
effected by heating the ortho acid to 213° —

2H3PO4-H2O = H4P2O7.

* Hydrogen disodium phosphate, although belonging to that class of com-
pounds commonly called acid salts, on account of the fact that it still retains
a portion of the replaceable hydrogen of the acid, is strongly alkaline in its
action upon litmus ; silver nitrate is a neutral compound, hence in this reaction,
by mixing an alkaline and a neutral liquid, an acid liquid is obtained, on
account of the molecule of nitric acid that is set free.

476 Inorganic Chemistry

The formation of this acid from two molecules of orthophos-
phoric acid will be made clear by the following formulas —

HO HO HO HO HO HO HO HO

O=P - ;o - H H; - o - P = O = H2o + O=P - o- p=o.

Pyrophosphates are formed when monohydrogen orthophos-
phates are heated. Thus, by heating hydrogen disodium ortho-
phosphate, sodium pyrophosphate is formed —

2H Na2PO4 - H2O = Na4P2O7.

When ammonium magnesium phosphate (see above) is heated
in the same way it loses water and ammonia, and is transformed
into magnesium pyrophosphate, thus —

2(NH4)MgPO4=Mg2P2O7 + H2O + 2NH3.

Properties. — Pyrophosphoric acid is an opaque white crystal-
line solid, readily soluble in water. Its aqueous solution passes
slowly into orthophosphoric acid, the change taking place rapidly
on boiling ; a solution of this acid, therefore, cannot be concen-
trated by boiling.

The pyrophosphates are stable salts, and their solutions may be
boiled without change ; by boiling with acids, however, they are
converted into orthophosphates.

Metaphosphorie Aeid, HPO3 or PO2(HO).— This acid is formed
when phosphorus pentoxide is allowed to deliquesce. It may be
obtained by .the abstraction of one molecule of water from one
molecule of orthophosphoric acid, which is brought about by heat-
ing the tribasic acid to redness —

H3P04-H2O = HPO3.

It is also obtained by strongly heating pyrophosphoric acid —
H4P2Or-H2O = 2HPO3.

The scdium salt is obtained by strongly igniting either dihydrogen
sodium phosphate, H2NaPO4, or hydrogen sodium ammonium
phosphate (microcosmic salf)^ HNa(NH4)PO4 ; or dihydrogen
sodium pyrophosphate, H2Na2P2Or.

Properties. — Metaphosphorie acid is a transparent vitreous
solid (frequently termed glacial phosphoric acid}. It is readily

Metaphosphoric Acid

477

fusible, and is usually cast into sticks. At a high temperature it
may be volatilised. Metaphosphoric acid is easily soluble in
water, and its solution is slowly transformed into orthophosphoric
acid ; this change takes place rapidly on boiling, and the acid
passes directly into the tribasic acid without the intermediate
formation of pyrophosphoric acid —

HPO3 + H2O = H3PO4.

Metaphosphoric acid is monobasic, but it possesses the remark-
able property of forming a number of salts which may be regarded
as derived from several hypothetical polymeric varieties of the
acid.

Monometaphosphoric acid, HPO3, forms monometaphosphates, NaPO3.

Dimetaphosphoric acid, (HPO3)2
Trimetaphosphoric acid, (HPO3)3
Tetrametaphosphoric acid, (HPO3)4,
Hexarnetaphosphoric acid, (HPO3)6,

dimetaphosphates, K2P2O6.
trimetaphosphates, Na3P3O9.
tetrametaphosphates, Pb2P4O12.
hexametaphosphates, Na6P6Ols.

The three compounds, ortho-, pyro-, and metaphosphoric acids,
are readily distinguished from each other by means of silver nitrate
and their action upon albumen : —

Reagent.

Orthophosphoric
Acid.

Pyrophosphoric
Acid.

Metaphosphoric
Acid.

Silver nitrate .
Albumen . .

Canary yellow
precipitate of
Ag3P04

No action

White crystalline
precipitate of
Ag4P207

No action

White gelatinous
precipitate of
AgP03

Coagulates

Orthophosphoric acid is also distinguished by giving a yellow
precipitate of ammonium phospho-molybdate upon the addition of
excess of a solution of ammonium molybdate in nitric acid (see
Molybdenum).

Compounds of Phosphorus containing Nitrogen. — By the action of am-
monia upon phosphorus pentachloride, and upon phosphoryl chloride (POC13),
a number of nitrogen derivatives are obtained. Thus, when gaseous ammonia
is passed over phosphorus pentachloride, and the solid mass so obtained is
heated in a stream of an inert gas until the ammonium chloride is driven
off, a white insoluble powder remains having the composition represented by
the formula PN(NH), to which the name phospham has been given—

PC1B + 7NH8=5NH4C1 + PN(NH).

478 Inorganic Chemistry

Phosphoryl Triamide, PO(NH2)3, is obtained by the action of gaseous
ammonia upon phosphoryl chloride —

POC13+6NH3=PO(NH2)3+3NH4C1.

When heated out of contact with air, phosphoryl triamide yields ammonia
and phosphoryl nitride, thus —

PO(NH2)3=2NH3+ PON.

Pyropliospliamic Acids. — Three of these compounds are known, which
may be regarded as pyrophosphoric acid, in which i, 2, and 3 of the (HO)
groups have been replaced respectively by the group (NH2), thus —

Pyrophosphoric acid .... P2O3(HO)4.

Pyrophosphamic acid . . . ' . P2O3(HO)3(NH2).

Pyrophosphodiamic acid . ......... P2O3(HO)2(NH2)2.

Pyrophosphotriamic acid . . . P2O3(HO)(NH2)3.

Compounds of Phosphorus with Sulphur.— A number of compounds of
phosphorus and sulphur have been obtained by heating together varying pro-
portions of sulphur and red phosphorus. The following compounds are
known : —

Analogous Oxides.

Phosphorus monosulphide . . . P4S . . —
Phosphorus sesquisulphide . . P4S3 . . —
Phosphorus trisulphide . . ; ; . PgSs . . P4Og
Phosphorus tetrasulphide (?) . . P2S4 . . P2O4
Phosphorus pentasulphide . . P.jS5 . . P2O5

Phosphorus Pentasulphide, P2S5. — This- compound is-%the best-known
member of the series. It is prepared by gently heating red phosphorus and
fragments of sulphur, in the proportion required by the formula, in a flask.
The elements combine with energy, and on cooling a solid mass is obtained.
This solid material is then distilled in a current of carbon dioxide, when the
pentasulphide is obtained in the form of yellow crystals. The compound may
also be obtained by dissolving ordinary phosphorus and sulphur in the proper
proportions jn carbon disulphide and heating the solution in sealed tubes
to 210°. On allowing the solution to cool, yellow crystals of the pentasul-
phide are deposited. Phosphorus pentasulphide is decomposed by water with
the formation of orthophosphoric acid and the evolution of sulphuretted
hydrogen —

ARSENIC.

Symbol, As. Atomic weight = 75. Molecular weight = 300.
Vapour density = 150.

. — Arsenic is found in the free state in nature,
usually in the form of small nodules, more rarely as distinct crystals.
In combination with sulphur it constitutes the minerals realgar, or

Arsenic 479

ruby sulphur •, As2S2, and orpiment, As2S3. In combination with
metals, as arsenides, it occurs widely distributed, the commonest
ores being arsenical iron, FeAs2 and Fe4As3 ; kupfernickel, NiAs
and NiAs2 ; and tin white cobalt, CoAs2. With metals and sulphur
it is met with in such minerals as arsenical pyrites, mispickel,
or white mundic, FeS2, FeAs2 ; cobalt glance, CoS2,CoAs2 ; nickel
glance, NiS2,NiAs2. Arsenic is present in small quantities in most
samples of iron pyrites, hence it finds its way into sulphuric acid
manufactured from pyrites. It also occurs in coal smoke, being
derived from the pyrites contained in coal, and hence is present in
the atmosphere : during the prevalence of yellow fogs the amount
of arsenic present is very appreciable.

Modes Of Formation.— On the small scale, arsenic is obtained
by heating a mixture of arsenious oxide, As4O6, with powdered
charcoal —

On a larger scale it is usually obtained either from native arsenic
or from arsenical pyrites ; the latter substance, when heated, gives
up arsenic, and ferrous sulphide is left behind —

FeS2,FeAs2 = 2As + 2FeS.

The mineral is heated in long narrow horizontal earthenware
retorts, into whose mouths are fitted earthenware receivers. The
arsenic volatilises and condenses in these receivers as a compact
crystalline solid. It is purified by redistillation.

Properties.— Arsenic which has been resublimed is a brilliant
steel-grey metallic-looking substance, forming hexagonal rhombo-
hedral crystals, having a specific gravity of 5.62 to 5.96. It is very
brittle, and is a good conductor of heat and electricity. Arsenic
begins to volatilise at 100°, and rapidly vaporises at a dark-
red heat, passing from the solid to the vaporous states without
liquefying. The vapour has a yellow colour and an unpleasant
garlic smell. When heated under pressure arsenic melts at 5oo°j
and on cooling solidifies to a compact crystalline mass. When
arsenic is vaporised in a glass tube, in a current of hydrogen, it
condenses along the tube in three distinct conditions : that which
is deposited nearest to the heated portion of the tube is .in the form
of rhombohedral crystals ; that which sublimes a little farthei
along, and condenses at a point where the temperature is about
210^-2.20°, consists of a black shining amorphous deposit ; while at a

480 Inorganic Chemistry

still more distant and cooler portion of the tube a grey crystalline
sublimate is formed. These are regarded as allotropic modifica-
tions of arsenic. The amorphous variety is also formed, when
arsenic hydride is decomposed by being passed through a heated
tube (q-v.). Amorphous arsenic is unacted upon by air at ordi-
nary temperatures, and only slightly oxidised at 80°. The grey
crystalline variety is readily oxidised on exposure to air at ordinary
temperatures.

Amorphous arsenic, when heated out of contact v/ith air to 360°,
is converted into the rhombohedral variety.

Arsenic, like phosphorus, forms tetratomic molecules, its mole-
cular weight as deduced from its vapour-density being 75 x 4 = 300.

When heated in oxygen arsenic burns with a bright bluish-white
flame, forming arsenious oxide, As4OG. It is oxidised by sulphuric
acid, nitric acid, and other oxidising agents. It combines readily
with chlorine, and when thrown into this gas in the condition
of powder it spontaneously inflames, forming arsenic trichloride.
Thrown into bromine a fragment of arsenic spontaneously in-
flames, and burns as it floats about upon the surface of the liquid.

Arsenic, in many of its characteristics, resembles the true metals ;
it is one of those elements lying on the borderland between true
metals and non-metals, to which the name metalloid is applied. It
is capable of forming alloys with metals, and an alloy of this
element with lead is employed for the manufacture of shot. It is
found that by the addition of a small proportion of arsenic to lead
the melted metal is more fluid, and therefore more readily assumes
the spheroidal form when projected from the shot tower, and on
solidification the alloy is considerably harder than pure lead.

ARSENIC HYDRIDE (Arsenuretted Hydrogen. Arsine}.
Formula, AsH3. Molecular weight— 78.03. Density =39. 015.

Modes Of Formation. — (i.) Arsenic hydride is formed when
soluble arsenic compounds are exposed to the action of nascent
hydrogen : thus, when a solution of arsenious oxide is introduced
into a mixture from which hydrogen is being generated, such as
zinc or iron and dilute hydrochloric or sulphuric acid, arsenic
hydride is obtained, mixed with free hydrogen —

As4O6 + 24H = 4AsH3 + 6H2O.

Arsenic Hydride 481

(2.) By the same action of nascent hydrogen, arsenic hydride
is formed when a solution of either arsenious oxide, As4O6, or
arsenic oxide, As2O5, is subjected to electrolysis.

(3.) Arsenic hydride is also formed when arsenical compounds
are in contact with organic matter which is undergoing decom-
position. During the growth of certain moulds and fungi a small
quantity of hydrogen is evolved, which by its action upon the
arsenic compound, gives rise to the formation of arsenic hydride.
By this action arsenic hydride is sometimes formed in dwelling-
houses where arsenical wall-papers are employed, and where, from
dampness or other causes, mould develops.*

(4.) Pure arsenic hydride is prepared by the action of dilute
hydrochloric or sulphuric acid upon an alloy of arsenic and zinc —

As2Zn3 + 3H2SO4 = 2AsH3 + 3ZnSO4,

or by the action of either water or dilute acid upon an alloy of
arsenic and sodium, prepared by heating sodium in the impure
arsenic hydride obtained by method No. i.

Properties. — Arsenic hydride is a colourless, offensive-smell-
ing, and highly poisonous gas. Under pressure it condenses to a
colourless liquid, which boils at -54.8° and solidifies at -113.5°.
The gas burns with a lilac-coloured flame, forming water and white
fumes of arsenious oxide —

4AsH3 + 6O2 = As4O6 + 6H2O.

When the supply of air to the flame is limited, as when a cold
surface is depressed upon it, water is formed and arsenic is deposited
as a shining black amorphous film —

4AsH3 + 3O2 = As4-f6H2O.

Arsenic hydride is readily decomposed by heat into its elements :
thus, when the gas is passed through a glass tube, which is
heated at one point by a Bunsen flame, arsenic in the amor-
phous condition is deposited upon the tube immediately beyond
the heated spot. Even when greatly diluted with hydrogen this
reaction takes place, and it therefore affords a delicate test for the
presence of exceedingly small quantities of arsenic. This method

* Extensive experiments on this subject by C. R. Sanger (Proc. American
Academy} have led him to believe that volatile organic arsenical compounds
are produced under these circumstances. No compound was isolated how-
ever.

2 H

482 Inorganic Chemistry

for the detection of arsenical compounds is known as Marsh's test,
and may be carried out by means of the apparatus seen in Fig.
126. Hydrogen is generated in the two-necked bottle from zinc
and dilute sulphuric acid (which are themselves free from arsenic),
and the arsenic in the form of an oxygen or a haloid compound is
introduced.* On igniting the issuing gas, and depressing a white
porcelain capsule into the flame, black stains of amorphous arsenic
are produced ; and if the tube be heated as shown in the figure,
the arsenic is deposited as a black film. The corresponding anti-
mony compound, SbH3 (q.v.\ gives rise to a similar deposit of
metallic antimony, when treated in the same way ; but the arsenic
deposit is readily distinguished by being easily soluble in a solu-

tion of calcium hypochlo-
rite. Many metals, such
as sodium or potassium,
when heated in arsenic
hydride, form alloys with
the arsenic, and hydrogen
is set at liberty ; while
metallic oxides when simi-

larly treated form metallic
FIG. 126. J .

arsenides and water.

Arsenic hydride is slightly soluble in water, but the solution on
exposure to air deposits arsenic.

When passed into a solution of silver nitrate, metallic silver is
precipitated, and a solution of arsenious oxide (the hypothetical
arsenious acid, H3AsO3) is obtained, thus —

When the gas is passed into copper sulphate solution, cuprous
arsenide is precipitated —

Arsenic hydride is decomposed by the halogens with energy,
forming the haloid compound of-arsenic and the halogen acid —

'•' When minute traces of arsenic have to be detected, as in food analyses,
the material is introduced into the cathode compartment of a small specially
constructed electrolytic cell (Thorpe) in which pure dilute sulphuric acid is
electrolysed.

Arsenic Chloride 483

Solid Arsenic Hydride. — When arsenide of potassium or sodium is acted
upon by water, a soft brown solid substance separates, which contains equal
atomic proportions of arsenic and hydrogen. Its molecular weight is unknown ;
its composition is therefore expressed by the formula, (AsH)n.

COMPOUNDS OF ARSENIC WITH THE HALOGENS.
The following compounds are known —

AsF3; AsCl3; AsBr3 ; Asl^

Two other compounds with iodine have been described contain-
ing the elements in the proportion represented by the formulae,
AsI2 and As2I5, the molecular weights of which are unknown.

Arsenic Fluoride, AsF3, molecular weight = 132, is formed when
sodium fluoride is heated with arsenic chloride —

3NaF + AsCl3 = 3NaCt + AsF3.

It is best obtained by distilling a mixture of arsenious oxide,
powdered fluor spar, and sulphuric acid in a leaden retort. The
hydrofluoric acid generated by the action of the acid upon the
calcium fluoride reacts upon the arsenious oxide, thus —

As406+12HF = 4AsF3 + 6H2O.

Properties. — Arsenic fluoride is a colourless fuming liquid,
boiling at 60.4°. It is rapidly decomposed by water into arsenious
oxide and hydrofluoric acid. On this account it forms painful
wounds when brought into contact with the skin.

Arsenic Chloride, AsCl3, molecular weight- 181.35, is ob-
tained when arsenic burns in chlorine, or when chlorine is passed
over fragments of arsenic in a tube.

It is also produced when either arsenic or arsenious sulphide is
distilled with mercuric chloride —

+ 2AsCl3.

It is readily prepared by the action of hydrochloric acid upon
arsenious oxide ; for which purpose sodium chloride, arsenious
oxide, and sulphuric acid are gently heated together in a retort
connected with a well-cooled receiver —

484 • Inorganic Chemistry

Properties. — Arsenic chloride is a colourless, fuming, and some-
what oily liquid which boils at 130.2°, and is extremely poisonous.
In the presence of excess of water, or when added to warm water,
it is decomposed into arsenious oxide and hydrochloric acid.
With a small quantity of water a solid crystalline arsenic chlor
hydroxide is formed, As(HO)2Cl —

AsCl3 + 2H2O = 2HCl + As(HO)2Cl.

Arsenious Bromide, AsBr3.— This compound is formed by the direct union
of arsenic with bromine, and is prepared by adding powdered arsenic to a
solution of bromine in carbon disulphide. On evaporation the compound is
deposited in the form of colourless deliquescent crystals, which melt at 20° to
25° to a straw-coloured liquid.

Arsenious Iodide, AsI3, is obtained by heating a mixture of arsenic and
iodine. It is most conveniently prepared by digesting a saturated ethereal
solution of iodine with powdered arsenic in a flask with a reflux condenser.
On filtering and cooling, the iodide'deposits in the form of lustrous red hexa-
gonal crystals.

OXIDES AND OXYACIDS OF ARSENIC.

Two oxides of arsenic are known, both of which act as anhy-
drides—

Arsenious oxide . . .... ,.' •. • As4O6.

Arsenic oxide (arsenic pentoxide) . ^. . As2O5.

No acid corresponding to arsenious oxide is known in the free
state, although the arsenites constitute a class of stable salts.

Three arsenic acids, derived from arsenic pentoxide, are known,
analogous in constitution to the three phosphoric acids, namely —

Ortho-arsenic acid . . H3AsO4 or AsO(HO)3.
Pyro-arsenic acid . . . H4As2O7 or As2O3(HO)4.
Metarsenic acid ... HAsO3 or AsO2(HO).

ARSENIOUS OXIDE.
Formula, As4O6. Molecular weight =396.

Mode of Formation. — Arsenious oxide is formed when arsenic
burns in air or in oxygen, or when arsenic minerals are roasted in
a current of air. On a small scale it may be produced by burning
arsenic in a hard glass tube in a stream of oxygen, and allowing

Arsenious Oxide

485

the white fumes of arsenious oxide to pass into a glass cylinder (as
shown in Fig. 127), where the greater part condenses, while the rest
is led into a draught flue.

Arsenious oxide is obtained as a secondary product, in the
metallurgical process of roasting arsenical ores of nickel, cobalt,
tin, silver, and others, for the extraction of these metals. It
is also obtained as a principal product by roasting arsenical
pyrites. The ore is heated either upon the hearth of a rever-
beratory furnace, where h is raked over from time to time, or
it is introduced by means of a hopper into one end of a long clay-
lined iron cylinder, placed at an incline of about i in 18, and caused
slowly to revolve about its longitudinal axis (Fig. 128). The lower
end of this cylinder enters a furnace, the upper end is connected to
a series of brickwork flues. The ore is delivered into the upper
end of the revolving cylinder, and as it gradually gravitates down

FIG. 127.

the incline, it is completely roasted by the furnace flames which
pass over it, and finally falls out into a chamber beneath. The
fumes of arsenious oxide pass through a series of chambers or
flues, so arranged as to present an extensive condensing surface to
the gases, and the crude product, known as arsenical soot, is from
time to time collected. This is known as Oxland and Hocking's
revolving calciner.

Properties. — Arsenious oxide, known familiarly as white arsenic,
or simply arsenic, is known in three modifications —

(i.) Amorphous. .

(2.) Octahedral crystals of the cubic or regular system.

(3.) Prismatic crystals of the monosymmetric system.

486

Inorganic Chemistry

Amorphous Arsenious Oxide is a colourless, transparent, vitreous

substance, which is obtained
when the vapour of the oxide
is condensed at a temperature
only slightly below its vaporis-
ing point. On exposure it gra-
dually becomes opaque, being
transformed into the regular octa-
hedral variety. This change
takes place from the outside, and
lumps of opaque " white arsenic,"
when broken, often show a
nucleus of the vitreous modifica-
tion. Amorphous arsenious oxide
may be preserved unchanged
in sealed tubes. The change
from the vitreous to the crystal-
line form is attended with evolu-
tion of heat, and a diminution of
specific gravity from 3.738 to
3-689-

Amorphous arsenious oxide,
when heated to about 200°, melts,
and at a higher temperature
vaporises. It is soluble in 108
parts of cold water.

Octahedral Arsenious Oxide. —
The vitreous form passes spon-
taneously into this variety. It
is obtained directly, by quickly
cooling the vapour of arsenious
oxide, or by crystallisation from
the aqueous solution of either
form of the oxide. Arsenious
oxide is also deposited in this
form from solution in hydro-
chloric acid.

Octahedral arsenious oxide is
less soluble in water than the
amorphous variety, i part requir-
ing 355 Parts of water for its

Arsenic Pentoxide 487

solution. When heated, the crystals vaporise without fusion, but
when heated under pressure they melt, and are converted into the
vitreous form.

Prismatic Arsenious Oxide is obtained by crystallisation from a
hot saturated solution of arsenious oxide'in potassium hydroxide.

Aqueous solutions of arsenious oxide possess a feeble acid re-
action, probably due to the formation of unstable arsenious acid,
H3AsO3. The acid has not been isolated, and on concentration
the solution deposits crystals of arsenious oxide.

Arsenious oxide is a powerful poison : from 2 to 4 grains usually
prove fatal. It is possible, however, by the habitual use of it, to
so accustom the system to this poison, that doses sufficiently large
to cause certain death to one unused to it may be taken with
apparent impunity. The use of arsenic is said to beautify the
complexion, and to improve the wind. The men who are em-
ployed upon arsenic works are constantly liable to swallow doses
of arsenious oxide which would cause death to one unaccustomed
to the occupation.

Arsenites. — Three classes of arsenites are known, which may
be regarded as being derived from the three hypothetical acids—

/Silver ortho-arsenite, Ag3AsO3.
Ortho-arsenious acid, H3AsO3> or As(HO)3 s

\ green),

{Calcium pyro-arsenite,Ca2As2O5.
Barium pyro-arsenite, Ba2As2O5.
Ammonium 1 /XTTT .
pyro-arsenite, }(NH4)4As2O5.

/ Potassium inetarsenite, KAsO2.

Metarsenious acid, HAsO2, or AsO(HO) J Acid'potassium j R AsO HAsO

| metarsenite, /
' Lead metarsenite, Pb(AsO2)2.

The pigment known as Schweinfurt green is a double metar-
senite and acetate of copper —

3Cu(As02)2,Cu(C2H302)2.

All arsenites, except those of the alkali metals, are insoluble in
water. When heated, most arsenites are converted into arsena/^j
and arsenic ; and when heated with charcoal the whole of the
arsenic is reduced.

Arsenic Pentoxide, As2O5.— This oxide is not formed when
arsenic burns in oxygen. It is obtained by the oxidation of ar-

488 Inorganic Chemistry

senious oxide by nitric acid, and subsequently heating the arsenic
acid so produced, to a dark-red heat —

2H3AsO4 = 3H2O + As2O6.

Properties. — Arsenic pentoxide is a white deliquescent solid,
completely soluble in water, with the formation of arsenic acid.

When strongly heated it breaks up into arsenious oxide and
oxygen—

2As2O6=As4O6+2O2.

ARSENIC ACIDS AND ARSENATES.

When arsenic pentoxide is dissolved in water, crystals are ob-
tained having the composition 2AsO(HO)3,H2O. At 100° these
melt and lose water, leaving ortho-arsenic acid, H3AsO4. By the
withdrawal of water from this acid, both pyro- and metarsenic acid
are obtained.

Heated between 140° and 180°, two molecules of the "ortho"
acid lose one of water —

2H3AsO4 = H4As2O7 + H2O.

And by heating the pyro-arsenic acid so obtained to 200°, another
molecule of water is expelled, with the formation of metarsenic acid
(compare corresponding acids of phosphorus) —

H4As2O7 = 2HAsO3+ H2O.

Pyro- and metarsenic acids are both crystalline solids, which
dissolve in water with the evolution of heat and formation of ortho-
arsenic acid ; aqueous solutions of these two acids, therefore,
cannot exist. In this respect they differ from the corresponding
phosphorus acids, both of which can be obtained in aqueous
solution.

Each of the three arsenic acids forms salts, of which the following
are examples : —

Trisodium ortho-arsenate .''•*•• • Na3AsO4.

Hydrogen disodium ortho-arsenate . HNa2AsO4.

Dihydrogen sodium ortho-arsenate . H2NaAsO4.

Ammonium magnesium ortho-arsenate. (NH4)MgAsO^.

Sodium pyro-arsenate . . . . Na4As2O7.

Sodium metarsenate . . . . NaAsO3.

Arsenic Trisulphide 489

The salts of pyro- and metarsenic acids, like the acids them-
selves, only exist in the solid state ; when dissolved in water they
pass into the ortho-compounds.

The arsenates are isomorphous with the corresponding phos«
phates.

COMPOUNDS OF ARSENIC WITH SULPHUR.
Three sulphides of arsenic are known, namely —

Arsenic disulphide (found native as Realgar) . As2S2.
Arsenic trisulphide (found native as Orpiment} . As2S3.
Arsenic pentasulphide As2S6.

Arsenic Disulphide, As2S2, is formed when sulphur and arsenic,
or arsenic trisulphide and arsenic, are heated together ; or by
heating arsenious oxide and sulphur —

As4O6+7S = 2As2S2+3SO2.

It is prepared on a large scale by distilling a mixture of iron
pyrites and arsenical pyrites —

FeS2,FeAs2+ 2FeS2= As2S2 + 4FeS.

Properties. — Arsenic disulphide is a red, vitreous, brittle solid,
having a specific gravity of 3.5. It is readily fusible, and sublimes
unchanged. Heated in air or oxygen, it burns with a blue flame,
forming arsenious oxide and sulphur dioxide—

2As2S2 + 7O2 = 4SO2 + As4O6.

Arsenic disulphide is employed in pyrotechny. So-called Bengal
fire consists of a mixture of realgar, sulphur, and nitre.

Arsenic Trisulphide, As2S3, is obtained by heating sulphur
and arsenic in the proportion required by the formula, and sublim-
ing the compound.

It may readily be produced by passing sulphuretted hydrogen
through a solution of arsenious oxide in hydrochloric acid —

As4O6 + 6H2S = 2As2S3 + 6H2O.

Properties. — The compound, as obtained by precipitation with
sulphuretted hydrogen, is a pure canary-yellow solid, which easily

490 Inorganic Chemistry

melts, and on again cooling forms a brittle crystalline mass. It
volatilises and sublimes unchanged, but when heated in air or
oxygen it burns with formation of arsenious oxide and sulphur
dioxide.

Arsenic trisulphide may be regarded as a thio-anhydride, as it gives rise to
a series of salts known as thio-arsenites , or sulpharsenites. Thus, when arsenic
trisulphide is brought into a solution of a caustic alkali, such as potassium
hydroxide, the sulphide readily dissolves with'the formation of an arsenite and
thio-arsenite, thus —

As2S3+4KHO=HK2AsO3+ HK2AsS3+ H2O.

Upon the addition of an acid, the salts are decomposed and arsenic tri-
sulphide reprecipitated —

Thio-arsenites. —These salts may be looked upon as being derived from
three hypothetical thio-arsenious acids, corresponding to the oxyacids —

Ortho-thio-arsenious acid, H3AsS3. Potassium ortho-thio-arsenite, K3AsS3.

T Ammonium pyro - thio - arsenite,
Pyro-thio-arsenious acid, H^AsgSj. -j (NH^As^g.

^ Lead pyro-thio-arsemte, Pb2As2S5.
Meta-thio-arsenious acid, HAsS2, Potassium meta-thio-arsenite, KAsS2.

Thio-arsenites of the alkali metals, the metals of the alkaline earths, and of
magnesium, are soluble in water, but decompose on boiling. Their solutions
are decomposed by acids, with evolution of sulphuretted hydrogen and pre-
cipitation of arsenic trisulphide, thus —

2K3AsS3+6HCl=

Arsenic Pentasulphide, As2S5. — This compound is prepared
by adding an acid to a solution of a thio-arsenate, thus —

Arsenic pentasulphide is a yellow, easily fusible solid. It is
readily soluble in caustic alkalies, forming an arsenate and a thio-
arsenate —

Arsenic pentasulphide, like the trisulphide, gives rise to a series of salts
known as thio-arsenates. These may be regarded as being derived from thr
three hypothetical thio-arsenic acids —

fTripotassium ortho-thio-arsenate, K3AsS4,
Ortho-thio-arsenic acid, H3AsS4. -| Hydrogen disodium ortho - thio - arsenate,

I. HNa2AsS4.

Pyro-thio-arsenic acid, H4As2S7. Magnesium pyro-thio-arsenate, Mg2As2S7.
Meta-thio-arsenic acid, HAsS3. Ammonium meta-thip-arsenate (NH4)AsS3.

Antimony 49 r

ANTIMONY.

Symbol, Sb. Atomic weight =120.2

Occurrence. — Antimony in the uncombined state is found in
small quantities in various parts of the world, and notably in
Borneo. In combination with oxygen, as Sb2O3, it constitutes
the mineral antimony bloom, or white antimony; and as Sb2O4 it
occurs m antimony ochre. In combination with sulphur, as Sb2S3,
it occurs as the mineral stibnite, or grey antimony ore, which is the
most important source of the metal ; and with both oxygen and
sulphur, as Sb2O3,2Sb2S3, it constitutes the mineral antimony blende,
or red antimony.

It also occurs in combination with sulphur and with metals, in
the form of thio-antimonites.

Modes of Formation. — (i.) Antimony is obtained from the
native sulphide by one of the two following methods. The
broken-up ore is heated in plumbago crucibles along with scrap
iron. As the mass melts, the sulphur combines with the iron,
forming a slag of iron sulphide, and the liberated antimony settles
out beneath —

Sb2S3 + 3Fe = 2Sb + 3FeS.

(2.) The crude sulphide is first liquated, or melted in such a
manner as to separate the sulphide from the rocky matter associated
with it. The liquated sulphide is then mixed with about half its
weight of charcoal, in order to prevent the mass from caking, and
carefully roasted. During this process the antimony sulphide is
partially converted into antimony trioxide (Sb2O3)2, which passes
into flues, and is there condensed, leaving a mixture containing
antimony tetroxide (Sb2O4), and unchanged sulphide. Most of
the arsenic present is also oxidised, and volatilises with the anti-
mony trioxide, while sulphur dioxide escapes. The residue, con-
sisting of the tetroxide and sulphide (known as antimony ash) is
mixed with an additional quantity of charcoal and with sodium
carbonate, and heated to redness in a crucible, when the changes
represented by the following equations take place—

(i.)
By the action of the carbon upon the sodium carbonate, sodium

492 Inorganic Chemistry

is liberated, which combines with the sulphur of the trisulphide,
forming sodium sulphide and metallic antimony —

(2.)
(3.)

The sodium sulphide in its turn unites with a further quantity of
antimony sulphide, forming a double sulphide of sodium and anti-
mony, which, mixed with the sodium carbonate and charcoal,
constitutes the slag. The metal obtained by either process is
subsequently refined.

Properties. — Antimony is a bright, highly crystalline, and very
brittle metal, possessing a bluish- white colour, and having a specific
gravity of 6.7 to 6.8. It is unacted upon by air or oxygen at the
ordinary temperature, but when heated it burns brilliantly, forming
antimony trioxide. The metal melts at 630° ; and when allowed to
solidify, its crystalline character is seen by the fern-like appearance
of its surface. If a quantity of the molten metal be allowed slowly
to cool, and when partially solidified the remaining liquid portion
be poured off, the interior of the mass is found to be lined with
well-formed rhombohedral crystals, isomorphous with arsenic. In
the act of solidification antimony expands, a property which it
imparts to its alloys, thus giving to them the valuable quality oi
taking very fine and sharp castings. The most important of these
alloys are type metal (lead 75, antimony 20, tin 5) ; stereotype metal
(lead 112, antimony 18, tin 3) ; Britannia metal (tin 140, copper 3,
antimony 9). Regarded as a metal, antimony is a bad conductor
of heat and electricity.

Dilute sulphuric and hydrochloric acids are without action upon
antimony. The concentrated acids convert it into sulphate and
chloride respectively —

2Sb + 6H2SO4 = 3SO2 + 6H2O + Sb2(SO4)3.

Antimony is oxidised by nitric acid, dilute acid converting it into
antimony trioxide or a compound of the oxide with nitrogen pent-
cxide, Sb2O3,3N2O5, while strong acid oxidises it chiefly into anti-
mony tetroxide and pentoxide.

Powdered antimony, when thrown into chlorine, takes fire
spontaneously and forms antimony trichloride.

Amorphous Antimony.— Antimony is obtained in an amor-
phous form by the electrolysis of a solution of tartar emetic in
antimony trichloride.

Antimony Hydride 493

Properties.— Amorphous antimony presents the appearance of
a smooth polished rod of graphite, and has a specific gravity of
5.78. It always contains a certain quantity of antimony trichloride
(from 4 to 12 per cent.) ; but whether this is in chemical union or
merely mechanically retained by the metal is not known. Amor-
phous antimony is very unstable, and readily passes into the
crystalline modification ; a slight blow, even a scratch with a
needle, causes it instantly to transform itself into the stable form
with explosive violence, the temperature at the same moment
rising to 250°, and clouds of the vapour of antimony trichloride
being evolved. ,

ANTIMONY HYDRIDE (Antimoniuretted Hydrogen).
Symbol, SbH3.

Modes Of Formation.— ( i.) This compound is formed when a
solution of an antimony compound is introduced into a mixture
generating hydrogen, such as zinc and sulphuric acid.

(2.) Hydrogen containing as much as 1 1 per cent, of antimony
hydride can be obtained by the regulated action of an alloy of
antimony and magnesium upon dilute hydrochloric acid ; and on
cooling this gaseous mixture in liquid air the antimony hydride
freezes out while the hydrogen passes on.

Properties.— Antimony hydride melts at -88° and boils at
-17°. At ordinary temperatures it is a colourless, offensive-
smelling, and poisonous gas, closely resembling the correspond-
ing arsenic compound in its general behaviour. It burns with
a violet-tinted flame, forming water and antimony trioxide —

When the supply of air is limited water is formed and antimony
is deposited ; when, therefore, a cold object is depressed upon the
flame black stains of metallic antimony are obtained. The gas is
easily decomposed by heat, and if passed through a glass tube
heated at one point a black deposit of antimony is formed upon
the glass. The antimony so deposited is insoluble in a solution of
bleaching powder (see Arsenic Hydride). Antimony hydride is
decomposed by the halogen elements, with the formation of the
halogen hydride, and the halogeri compound of antimony —

+ 3Cl2=3HCH-SbCl3

494 Inorganic Chemistry

Sulphuretted hydrogen, under the influence of sunshine, converts
antimony hydride into antimony trisulphide —

2SbH3 + 3H2S = Sb2S3 + 6H2.

When passed into silver nitrate solution the antimony is preci-
pitated in combination with silver, in this way differing from the
arsenic analogue —

COMPOUNDS OF ANTIMONY WITH THE HALOGENS.

The compounds represented by the following formulae are
known —

SbF3 ; SbCl3 ; SbBr3 ; SbI3
SbF6; SbCl6.

Antimony Trifluoride, SbF3> is prepared by dissolving the trioxide in
aqueous hydrofluoric acid. From the concentrated solution it is deposited in
the form of white deliquescent crystals. It dissolves in water, and is gradually
converted into an oxyfluoride.

Antimony Pentafluoride, SbF5, is obtained when hydrated antimony pent-
oxide is dissolved in aqueous hydrofluoric acid. When the solution is evapo-
rated the compound remains as an amorphous gum-like residue.

Both of these fluorides exhibit a great tendency to unite with alkaline
fluorides, forming double salts, such as SbF3,2KF ; SbF3,2NH4F, in the case
of the trifluoride ; and SbF5, KF ; SbF5,2KF, with the pentafluoride.

Antimony Trichloride, SbCl3, is formed when chlorine is
passed over metallic antimony, or antimony trisulphide —

2Sb + 3Cl2=2SbCl3.
2Sb2S3 + 9C12 = 4SbCl3 + 3S2C12.

It may also be obtained by the action of boiling hydrochloric
acid, containing a small quantity of nitric acid, upon either metallic
antimony, antimony trioxide, or trisulphide —

Sb2S3 + 6H Cl = 2SbCl3 + 3H2S.

Properties. — Antimony trichloride is a colourless, deliquescent,
crystalline substance, melting at 73.2° to an oily liquid, which
again solidifies to a soft translucent mass. It is soluble in alcohol
and in carbon disulphide, and from the latter may be crystallised.
It may be dissolved in a small quantity of water unchanged. Thus,

Antimony Pentachloride 495

if allowed to deliquesce it liquefies in the water it absorbs, forming
a colourless solution, which, upon evaporation over sulphuric acid,
again deposits crystals of the trichloride. The addition of larger
quantities of water results in the formation of oxychlorides * —

(i.) SbCl3+ H20 = 2HCH SbOCl.
(2.)

Continued boiling with water removes the whole of the chlorine,
forming the trioxide —

Sb4O5Cl2 + H2O = 2HC1 + Sb4O6.

Antimony chloride unites with alkaline chlorides, forming double salts (see
Antimony Fluoride), such as SbCl3,2NH4Cl ; SbCl3,8KCl. With potassium
bromide it forms the compound SbCl3,3KBr, which, strangely enough, appears
to be identical with the double compound of antimony tribromide with
potassium chloride, SbBr3,3KCl.

Antimony Pentaehloride, SbCl5, is obtained by passing excess
of dry chlorine over metallic antimony, or antimony trichloride, in
a retort, when antimony pentachloride distils over in the excess of
chlorine —

Sba3 + Cl2=SbCl6.

Properties. — Antimony pentachloride is a nearly colourless,
strongly-fuming liquid. It solidifies, when cooled, to a mass of
colourless crystals, which remelt at —6°. Under the ordinary
atmospheric pressure the pentachloride dissociates, when heated,
into the trichloride and chlorine, but under reduced pressure it
may be boiled and distilled. Thus, under a pressure of 22 mm.
it boils at 79°.

By the regulated action of ice-cold water, oxychlorides are
formed —

(i.)

(2.) SbOCl3+H2O

Antimony pentachloride, and also the oxychlorides, are con-
verted by hot water into pyro-antimonic acid (analogous to pyro-
arsenic and pyro-phosphoric acids) —

2SbCl6 + 7H2O = H4Sb2O7 + 10H Cl.
2SbO2Cl + 3H2O = H4Sb2O7 + 2HC1.

* The mixed product obtained by the action of water upon antimony tri-
chloride is known as powder of Algaroth.

496 Inorganic Chemistry

Sulphuretted hydrogen (the sulphur analogue of water) acts
upon antimony pentachloride, forming antimony sulphotrichloride,
corresponding to the oxytrichloride —

SbCl6+H2S = SbSCl3 + 2HCl.

Antimony tribromide, SbBr3, and antimony tri-iodide, SbI3, are obtained by
adding powdered antimony to solutions of the halogens in carbon disulphide,
from which liquid the compounds are crystallised : the bromide as colourless
deliquescent crystals, and the iodide as hexagonal ruby-red crystals. Both of
these compounds are similarly acted upon by water, forming the oxybromides
SbOBr ; Sb4O5Br2, and the oxyiodides SbOI ; Sb4O5I2.

CTXIDES AND OXYACIDS OF ANTIMONY.
Three oxides of antimony are known —

Antimony trioxide (antimonious oxide) . (Sb2O3)2 or Sb4O6.
Antimony tetroxide . ;'*•'"• • • Sb2O4.
Antimony pentoxide .... Sb2O5.

No acids are known corresponding to the trioxide, although a sodium salt
of the hypothetical metantimonious acid, HSbO2, has been described, having
the composition NaSbO2,3H2O.

Three acids are known derived from antimony pentoxide which
are analogous to the three arsenic and phosphoric acids —

Orthoantimonic acid .••;•• • : • • H3SbO4.
Pyroantimonic acid . v -'.: ... . H4Sb2O7.
Metantimonic acid . •.- i > •• HSbO3.

Antimony Trioxide, Sb4O6, may be prepared by the addition of
hot water to a solution of either antimony trichloride or antimony
sulphate, and washing the precipitated oxide with a solution of
sodium carbonate to remove the free acid —

4SbCl3 + 6H2O = Sb4O6+12HCl.

Properties. — Antimonious oxide is a white powder, which,
when volatilised, condenses in two distinct forms, namely, pris-
matic crystals of the trimetric system and regular octahedra. The
former are deposited nearest to the heated material, the latter in
more remote and cooler regions. (See Arsenious Oxide, with
which antimonious oxide is iso dimorphous.} Antimonious oxide
is only very slightly soluble in water, and the solution is without

Antimony Pentoxide 497

action upon litmus. It is insoluble in nitric or sulphuric acid, but
is dissolved by hydrochloric acid with formation of the trichloride.
It is readily soluble in tartaric acid, and in a boiling solution
of hydrogen potassium tartrate (cream of tartar), giving rise to
potassium antimony tartrate, or tartar emetic —

4H K(C4H4O6) + Sb4O6 = 4(SbO)K(C4H4O6) + 2H2O.
Antimonious oxide burns in the air, forming the tetroxide —

Antimony Tetroxide, Sb2O4, is formed when the trioxide burns
in air. It may be prepared by strongly heating antimony pent-
oxide —

Properties. — Antimony tetroxide is a white non-volatile powder
which is insoluble in water. It is decomposed by boiling hydrogen
potassium tartrate, forming tartar emetic and metantimonic acid,
thus—

H K(C4H4O6) + Sb2O4 = (SbO)K(C4H4O6) + HSbO3.

Antimony Pentoxide, Sb2O6, is obtained by oxidising metallic
antimony with nitric acid, and heating the antimonic acid so
obtained to a temperature not exceeding 275°.

Properties. — Antimony pentoxide is a straw-coloured powder,
insoluble in water. When heated to 300° it gives up oxygen and
is converted into the tetroxide. Its feeble acidic character is seen
by its formation of an alkaline metantimonate when fused with an
alkaline carbonate —

Sb2O6+Na2CO3 = CO2 + 2NaSbO3.

Antimonic Acids and Antimonates.— None of the three antimonic acids
can be obtained by the action of water upon the oxide. Pyro-antimonic acid
is formed when antimony pentachloride is treated with hot water, and the
precipitate dried at 100° —

2SbCl5 + 7H2O = H4Sb2O7 + 10HC1.

Pyro-antimonic acid readily passes, by loss of water, into metantimonic
acid —

H4Sb2Or - H20=2HSb03.

Metantimonic acid is also formed by oxidising metallic antimony by means
of nitric acid —

2Sb+4HN03=2HSb03 + NO2 + 3NO + H2O,

2 I

498 Inorganic Chemistry

or by the decomposition of an aqueous solution of a metantimonate by means

of nitric acid —

KSbO3+HNO3=B

On allowing the precipitated metantimonic acid to remain for a long time
in contact with water it is converted into ortho-antimonic acid, H3SbO4—

HSb03 + H20 = H3Sb04.

No salts of ortho-antimonic acid, H3SbO4, are known ; the antimonates,
therefore, belong to the two acids, pyro-antimonic acid and metantimonic acid —

Pyro-antimonates.* Metantimonates.*

Normal potassium pyro-anti- Potassium metantimonate, KSbO3.

monate, K4Sb2O7.
Hydrogen potassium pyro- Barium metantimonate, Ba(SbO3)a.

antimonate, H2K2Sb2O7.

COMPOUNDS OF ANTIMONY WITH SULPHUR.
Two sulphides of antimony are known, namely —

Antimony trisulphide .... Sb2S3,
Antimony pentasulphide . . . . Sb2S5.

Antimony Trisulphide, Sb2S3.— This compound occurs native
as the mineral stibnite^ or grey antimony ore. It is prepared by
heating a mixture of powdered antimony and sulphur (in propor-
tion required by the formula) beneath a layer of fused sodium
chloride in a crucible. It is also formed when sulphuretted hydro-
gen is passed through a solution of antimony trichloride, or a
solution of tartar emetic —

Properties. — Antimony trisulphide as it occurs native, and as
obtained by the direct union of antimony and sulphur, is a grey-

* As only two types of antimonates are known, and as the salts of the type
MSbO3 are the best known, the name antimonates was formerly applied to
them, and the term metantimonates was given to the salts belonging to the
other class. It is better, however, to adopt the same system of nomenclature
for the antimony compounds as that which is in use for the similarly constituted
arsenic and phosphorus compounds —

Phosphates. Arsenates. Antimonates.

Ortho . . M3PO4 M3AsO4

Pyro . ! M4P20- M4As2O7 M^O?

Meta . . MPO» MAsO8 MSbOa

Antimony Pentasulphide 499

black crystalline substance ; as prepared by precipitation with
sulphuretted hydrogen, and subsequently drying at 200°, it is a
brick-red amorphous powder, which, when melted and slowly
cooled, solidifies in the crystalline form. Antimony sulphide sub-
limes unchanged when heated in an inert gas, but when heated in
air sulphur dioxide is evolved, and antimonious oxide and tetroxide
are formed. Heated with hydrochloric acid, it evolves sulphuretted
hydrogen, and forms antimony trichloride —

Antimony Pentasulphide, Sb2S5, is obtained when antimony
pentachloride is mixed with water, and sulphuretted hydrogen
passed through the liquid —

or—

2SbO2Cl + 5 H2S = Sb2S5 + 4H2O + 2H Cl.

Properties. — Antimony pentasulphide is a dark, orange-red
powder, which, on being heated, is decomposed into the trisulphide
and free sulphur.

Both of these antimony sulphides may be regarded as thio-anhydrides, for
although no thio-acids derived from them are known, salts have been produced
which may be viewed as derivatives of hypothetical thio-acids. When the
trisulphide is either fused with caustic potash, or boiled in an aqueous solution,
potassium thio-antimonite is formed —

2Sb2S3+4KHO=3KSbS2+ KSbO2+2H2O.

Similarly, when antimony pentasulphide is dissolved in potassium hydroxide,
a mixture of antimonate and thio-antimonate is obtained —

The following are illustrations of the thio-salts of antimony —
Sulphide. HTdds!iCal Salts.

( (Ortho) H3SbS3. Potassium thio-antimonite, K3SbS3.
< (Meta) HSbS2. Silver thio-antimonite, AgSbS^

l(Pyro) H4Sb2S5. Lead thio-antimonite, Pb2Sb2S5.

Potassium thio-antimonate, K3SbS4.

Barium thio-antimonate, Ba3(SbS4)2.
Only ortno-thio-antimonafe* are known.

5OO Inorganic Chemistry

BISMUTH.

Symbol, Bi. Atomic weight =208

Occurrence. — Bismuth occurs most commonly in the uncom-
bined condition. It is met with in combination with oxygen, as
Bi2O3, in bismuth ochre; and in combination with sulphur, as
Bi2S3, in bismuth glance.

Mode of Formation. — Bismuth is principally obtained from
the native metal, and from ores with which metallic bismuth is
associated. The broken-up ore is liquated by being heated in
inclined iron pipes, when the bismuth readily melts and drains
away.

Pure bismuth can be prepared from the crude metal thus ob-
tained, by first dissolving it in nitric acid, forming bismuth nitrate
Bi(NO3)3, and then precipitating the basic nitrate by the addition
of water —

= (BiO)NO3,H2

The basic nitrate is next dried and heated in a crucible with
charcoal ; the salt is first converted into the trioxide by the action
of heat, and the oxide is then reduced by the carbon —

2(BiO)NO3,H2O = Bi2O3 + N2O4 + O + 2H2O.

Properties. — Bismuth is a lustrous white metal with a faint
reddish tinge. It melts at 268.3°. If the molten metal be allowed
to cool until partially solidified, and the remaining liquid be then
poured off, obtuse rhombohedral crystals (belonging to the hexa-
gonal system), closely approaching to the cube, are obtained.

The specific gravity of bismuth is 9.823 ; it is extremely brittle,
and a poor conductor of electricity. Bismuth is unacted upon by
dry air at ordinary temperatures ; moist air tarnishes its surface.
Heated in air or oxygen it burns, forming the trioxide. It is only
slightly attacked by hydrochloric acid, but is converted by hot
sulphuric acid into a basic sulphate.

Bismuth readily forms alloys with other metals, and imparts to
them the useful properties of ready fusibility and hardness. The
alloys known by the general name of fusible metal contain bismuth ;
thus, Wood's fusible metal, which melts at 65°, consists of 4 parts
of bismuth, 2 of lead, I of tin, and i of cadmium.

Bismuth Bichloride 501

COMPOUNDS OF BISMUTH WITH THE HALOGENS.
Compounds represented by the following formulae are known —
BiF3 BiCl3 BiBr3 BiI3.

(BiCl2)2 (BiBr2)2?

Bismuth Trichloride, BiCl3, may be prepared by passing dry
chlorine over powdered bismuth gently heated in a retort. A
yellow liquid is first formed, after which the stream of chlorine is
stopped and the liquid distilled, when the trichloride sublimes in
the form of crystals. It may also be obtained by distilling a mix-
ture of powdered bismuth and mercuric chloride —

Properties. — Bismuth trichloride is a white, extremely deli-
quescent crystalline compound. Heated in an atmosphere of
chlorine, it melts to a yellow liquid. It is decomposed by water
with the precipitation of bismuth oxychloride- —

Bismuth Diehloride (BiCl2)2, is obtained by the prolonged
heating of mercurous chloride and finely powdered bismuth to 230°
in a sealed tube. The mixture melts, and mercury collects at
the bottom, and on cooling the dichloride solidifies as a black,
extremely deliquescent solid upon the surface of the mercury.
When heated above 300° the dichloride is resolved into the tri-
chloride and metallic bismuth. The molecular weight of the
compound is unknown.

Bismuth Tribromide, BiBrg, is prepared by gradually adding bromine to
powdered bismuth and slightly warming the mixture for some time. The
bromide sublimes in the form of golden-yellow, deliquescent crystals, which
are decomposed by water, forming oxybromide, BiOBr.

Bismuth Tri-iodide, BiI3, is prepared by subliming a mixture of iodine and
bismuth. The sublimate is afterwards finely powdered and again sublimed,
and the product finally distilled in a stream of carbon dioxide, when it forms
dark grey crystals with a bright metallic lustre. Boiling water decomposes
the compound, with formation of bismuth oxyiodide, BiOI.

COMPOUNDS OF BISMUTH WITH OXYGEN.
Four oxides of bismuth are known, namely —

Bismuth dioxide (Hypobismuthous oxide) . . BiaO^

„ trioxide (Bismuthous oxide) . . Bi2O3.

„ tetroxide (Hypobismuthic oxide) . . Bi2O4.

„ pentoxide (Bismuthic oxide) . . Bi2O6e

5O2 Inorganic Chemistry

None c\f these compounds is an acid-forming oxide, although,
with the exception of the first, they all form hydrated oxides.
These hydrated oxides have no acidic properties, and no salts
have been obtained in which the acidic or negative portion of the
molecule consists of bismuth and oxygen. All the four oxides,
when acted upon by acids, yield the same series of salts in which
the bismuth fulfils the functions of a trivalent element, replacing
three atoms of hydrogen. In the case of the dioxide, metallic
bismuth is deposited, thus —

While with the higher oxides oxygen is evolved —

2Bi(NO3

Bismuth trioxide is the most stable and the most important of
the oxides ; when heated in air, the remaining three compounds
are converted into the trioxide : the dioxide by oxidation, and the
tetroxide and pentoxide by loss of oxygen. The trioxide alone is
unchanged on being heated in air or oxygen.

Bismuth Dioxide, Bi2O2. — This oxide is prepared by adding a mixed solu-
tion of bismuth trichloride and stannous chloride to an excess of a 10 per cent.
solution of caustic potash, air being excluded : potassium stannate is formed,
and bismuth dioxide is precipitated —

+ 10KHO=Bi2O2+8KCl+K2SnO3+5H20.

Properties. — The precipitated compound, after being washed in dilute
caustic potash and dried in vacuo, is obtained as a black crystalline powder.
When heated in air it smoulders, uniting with oxygenate form the trioxide.
When moist it oxidises spontaneously —

Bi2O24-O=Bi2O3.

Bismuth Trioxide, Bi2O3, is formed when the metal is burnt in
air or oxygen. It may also be obtained by heating the hydrated
oxides, the carbonate, or basic nitrate, thus —

Bi2O3,H20 = Bi2
Bi2O3,CO2 = Bi2O3 + CO2.
2(BiO)NO3,H2O = Bi2O3 + 2NO2 + O + 2H2O.

Properties. — Bismuth trioxide is a cream-coloured powder,
insoluble in and unacted upon bv water, and is the only oxide of

Bismuth Tetr oxide 5°3

bismuth which is unchanged when heated in the air or in oxygen.
It dissolves in acids, forming salts of bismuth —

3H2O-f2Bi(NO3)3.
4 = 3H20 + Bi2(S04)3.

With small quantities of hydrochloric acid it first forms bismuth
oxychloride, BiOCl, which dissolves in additional acid, yielding
the trichloride —

Bi2O3 + 2HCl = H2O + 2BiOCl.
Bibci + 2HCl = H2O + BiCl3.

None of these compounds is soluble in water without the presence
of excess of the acid. Water alone converts them into insoluble
basic salts and free acid, which in the state of extreme dilution
is unable to exert any solvent action. Thus, in the case of the
nitrate when water is added, this compound is decomposed into the
basic nitrate and free nitric acid —

Bi(N03)3 + 2H20 = (BiO)N03,H20 + 2HN03.

Bismuth trioxide forms three hydrates, represented by the
formulae —

Bi2O3,H2O. Bi2O3,2H2O. Bi2O3,3H2O.

These hydrates have no acid properties, and are incapable of
combining with bases to form salts, but themselves play the part
of a base, uniting wit]? acids to form bismuth salts.

The trihydrate is obtained by pouring an acid solution of bis-
muth nitrate into an excess of strong aqueous ammonia —

2Bi(N03)3 + 6NH3H2O = 6NH4(NO3) + Bi2O3,3H2O.

Heated to 100° it is converted by loss of water into the mono-
hydrate —

Bi2O3,3H2O = Bi2O3,H2O + 2H2O.

Bismuth Tetroxide, Bi2O4, is formed by the action of potassium
hypochlorite upon the trioxide, the product being dried at 180° —

Bi2O3 + KC1O = Bi2O4 + KC1.

Properties. — Bismuth tetroxide is a brownish-yellow powder,
which readily parts with an atom of oxygen and passes into the
trioxide.

504 Inorganic Chemistry

Bismuth Pentoxide, Bi2O6, is prepared by passing chlorine into
a nearly boiling solution of caustic potash in which is suspended a
quantity of bismuth trioxide —

Properties. — Bismuth pentoxide is a red powder, which is
readily deoxidised into the tetroxide and trioxide by heat. It com-
bines with water, forming the hydrate Bi2O5,H2O, but with excess
of water it is gradually deoxidised into hydrates of the tetroxide or
trioxide.

Bismuth pentoxide is reduced, with evolution of oxygen, by both
nitric and sulphuric acids —

Bi2O6 + 3H2SO4 = Bi2(SO4)3 + 3H2O + O2.

With hydrochloric acid it behaves in the usual manner of
peroxides, causing the evolution of chlorine —

Bi2O6 + 10HC1 = 2BiCl3 + 5H2O + C12.

Bismuth Trisulphide, Bi2S3.— This compound is the only com-
pound of bismuth with sulphur that is known with certainty. It
occurs native as the mineral bismuth glance.

It is precipitated when sulphuretted hydrogen is passed into a
solution of a bismuth salt —

2Bi(NO3)3 + 3H2S = Bi2S3 + 6HNO3.

It is also obtained by heating together the requisite proportions
of bismuth and sulphur.

Properties. — As obtained by precipitation, bismuth sulphide is
a dark brown, almost black powder; the native sulphide forms
steel-grey lustrous crystals.

It is decomposed, when strongly heated, into its constituent
elements. Bismuth sulphide differs from the corresponding anti-
mony and arsenic compound in not being dissolved by alkaline
hydrates or sulphides.

CHAPTER IV
THE ELEMENTS OF GROUP L (FAMILY A.)

THIS family comprises the following five elements, known as the
alkali metals —

Atomic Weights. Melting-points.

Lithium (Li) . . *. ', * . 7.00 . . . . 180°
Sodium (Na) - " • , .. . , 23.00 .... 95-6°
Potassium (K) . . . 39.10 . . . . 62.5°
Rubidium (Rb) . . . 85.45 . . . . 38.5°
Caesium (Cs) .... 132.9 .... 26.5°

The most important and the most abundant of these elements
are potassium and sodium, which also were the first to be dis-
covered, having been isolated by Davy in the year 1807. The
element lithium, although widely distributed in nature, is for the
most part found only in minute quantities; the element was first
isolated by Bunsen in the year 1855. The two remaining elements
are still rarer substances, usually met with in very minute quantities
accompanying sodium and potassium. Both of these elements
were discovered by Bunsen by means of the spectroscope — caesium
in 1860 and rubidium in the following year.

All these elements are soft, silvery-white metals, which may be
readily cut with a knife, and which rapidly tarnish in the air.
They all decompose water at the ordinary temperature. The
members of this family exhibit that gradation in properties which
is met with in all similar families. Thus, their melting-points
gradually decrease as their atomic weights rise, as will be seen
from the figures given above. Their chemical activity also steadily
increases as we pass from lithium to caesium. Thus, in the case

506 Inorganic Chemistry

of their behaviour in contact with water : potassium, when thrown
upon cold water, decomposes that liquid with sufficient energy to
cause the ignition of the hydrogen which is evolved ; sodium
under the same conditions melts and floats about upon the sur-
face, but the action is not sufficiently energetic to effect the
inflammation of the gas, unless the water be previously heated ;
while with lithium, even with boiling water, the temperature
produced by the reaction does not rise to the ignition-point of
hydrogen. The same is also seen in the spontaneous oxidation
of these elements when they are exposed to the air. Thus,
lithium when cut with a knife, although it is soon covered with
a film of oxide, nevertheless retains its bright metallic surface for
some seconds ; sodium tarnishes so much more quickly, that the
film of oxide appears almost to follow the knife. When potassium
is cut the bright surface can scarcely be seen, so rapid is the
oxidation, and if left exposed a fragment of the metal soon begins
to melt by the heat of its own oxidation, and frequently spon-
taneously ignites. With rubidium and caesium the oxidation is
even more rapid, and a fragment of these metals freely exposed to
the air very rapidly takes fire spontaneously.

The electro-positive character of these elements gradually in-
creases from lithium to caesium, which is the most electro-positive
of all the known elements.

The term alkali^ applied to metals of this family, was originally
used (before any distinction was made between potash and soda) to
de-note the salt obtained by treating the ashes of plants with water.
Later on, in order to distinguish between this substance and what
became known as the volatile alkali (i.e. ammonium carbonate),
it was termed the fixed alkali. The first distinction between
potash and soda was based upon the erroneous belief that the
former was entirely of vegetable origin, while the latter was only
to be found in the mineral kingdom ; hence the names vegetable
alkali and mineral alkali were used to denote these two sub-
stances, both of which were regarded as elementary bodies until
1807, when Davy showed that they contained the two metals
potassium and sodium.

The Alkali Metals

507

The resemblance between the different members of this family
and between their compounds is very close ; so much so, that in the
case of sodium, potassium, rubidium, and caesium, there are scarcely
any ordinary chemical reactions by which they can be distinguished.
They are all readily identified, however, by means of the spectro-
scope. When a minute quantity of a lithium salt is introduced
upon a loop of platinum wire into the non-luminous Bunsen flame,
the latter is tinged a brilliant crimson-red colour ; a potassium salt

FIG. 129.

similarly treated colours the flame a delicate lilac, while a sodium
compound gives a brilliant daffodil-yellow colour. The colour
imparted to a flame by rubidium and caesium salts is indistinguish-
able by the eye from that given by potassium compounds ; and,
moreover, when any of these are mixed with a sodium salt the
intense yellow emitted by the latter completely masks the colours
given by the others. By means of the spectroscope, not only are
the apparently similar colours given by potassium, rubidium, and

508 Inorganic Chemistry

caesium readily distinguished, but the presence of any or all of
them is easily detected, even when admixed with sodium salts.
Spectrum analysis is based upon the fact that light of different
colours has different degrees of refrangibility, and therefore when
passed through a prism the different coloured rays are bent out
of their straight course at different angles. Ordinary white light
is composed of rays of all degrees of refrangibility, i.e. rays of all
colours; hence, when a^beam of such light is passed through a
prism, the various coloured rays are separated and become spread
out in the order of their refrangibility, from the least refrangible
red at the one extreme to the deep violet at the other. This
familiar "rainbow" coloured band of light is termed the con-
tinuous spectrum. «

A simple form of spectroscope is seen in Fig. 129. The light is
caused to pass through a narrow slit at the end of the fixed tube B,
known as the collimator tube. If the prism P be removed and the
telescope A be moved round so as to be in a continuous line with B,
a magnified image of the slit is seen by the observer. When the
prism is replaced, and A is moved into such a position that the bent
rays fall upon its lens, the continuous spectrum is seen, which is an
infinite number of strips of light (corresponding to the image of the
slit) of all colours arranged side by side. If the light to be
examined, instead of being ordinary white light, were composed of
rays all of one degree of infrangibility (i.e. monochromatic light),
there would be produced only a single image of the slit, which
would fall in that position corresponding to the particular degree
of refrangibility of the light. Such a monochromatic light is pro-
duced when a sodium salt is heated in a Bunsen flame ; if, there-
fore, a salt of this metal be introduced upon a loop of platinum
wire into the non-luminous flame G, and the light, after passing
through the prism, be observed through A, instead of a continuous
spectrum, there will be seen a single image of the slit, falling in
the brightest yellow part of the spectrum. When the sodium salt
is replaced by a lithium salt, it is seen that two images of the slit
are obtained, one in the red and the other in the yellow regions of
the spectrum. The light emitted from this element consists of rays

The Alkali Metals 509

of two degrees of refrangibility. We say, therefore, that the
spectrum of sodium is one yellow line,* and that of lithium con-
sists of one red and one yellow line. In order to distinguish the
positions of, for example, the yellow lithium line and that given
by sodium, an image of a graduated scale, illuminated by the
candle flame F, is also thrown into the telescope A.

If salts of sodium and lithium mixed together be introduced into
the flame G, then three images of the slit are seen, namely, the
yellow line given by the sodium, the yellow line of the lithium,
situated slightly nearer the red, and the lithium red line.

Potassium, like lithium, gives a light of two degrees of refrangi-
bility, forming consequently two images of the slit, one in the
deep red and the other in the deep violet ; if, therefore, lithium,

FIG. 130.

sodium, and potassium salts are mixed, and examined by the
spectroscope, five lines are seen (Fig. 130), namely, two red (one
belonging to lithium and one to potassium), two yellow (one
belonging to lithium and one to sodium), and the violet line of
potassium.

When analysed in this manner, the lights emitted by rubidium
and caesium compounds are seen to be totally different from each
other, and from potassium. The spectrum of rubidium consists of
two prominent lines in the violet (nearer the blue region than that
belonging to potassium), two brilliant red lines (very near the
potassium red line), and a number of less brilliant lines in the

* In reality, when examined by a higher dispersive power, the sodium line
is seen to be a group of lines.

510 Inorganic Chemistry

orange, yellow, and green. That of caesium consists of two bril-
liant blue lines, two bright red lines (near the lithium red line), and
a number of less prominent lines in the yellow and green. It will
be seen, therefore, that the three elements potassium, rubidium,
and caesium may be at once sharply distinguished by this optical
method of analysis, although they so closely resemble one another
in their chemical behaviour, as to render it highly probable that
the separate existence of the two latter would never have been dis-
covered by chemical methods alone.

Indeed, before the discovery of caesium by Bunsen, a rare
mineral known as Pollux (now known to contain caesium) was
mistaken for a potassium mineral.*

The element lithium, the member of the family that belongs to
the Typical series, exhibits certain characteristic differences from
the other members. This is seen particularly in the case of the
carbonate and phosphate of this element. Lithium carbonate is so
little soluble in water, that it is precipitated by the addition of
carbonate of either sodium or potassium to a solution of a lithium
compound. The phosphates of all the other members are readily
soluble in water, while lithium phosphate is almost insoluble, and
is precipitated from solutions of a lithium salt by the phosphates of
either sodium or potassium. In these two compounds, the car-
bonate and phosphate, lithium behaves more like one of the metals
of the alkaline earths.

All the metals of this family are monovalent, and replace each
other, atom for atom, in chemical compounds.

POTASSIUM.

Symbol, K. Atomic weight = 39.00.

Occurrence.— In combination this element is widely distributed
in nature. It forms an essential constituent of many of the com-
mon silicates and rocks which form the earth's crust. From
these rocks, by processes of disintegration, the potassium com-

* The student should consult special works on spectrum analysis.

Potassium 5 1 1

pounds find their way into the soil, from whence they are absorbed
by plants, which can only flourish in a soil that contains com-
pounds of potassium. Most of the potassium found in plants is
present in combination with organic acids.

From the vegetable kingdom, potash passes directly into the
bodies of animals. The material known as suint^ which is the

FIG. 131.

oily perspiration of the sheep, that accumulates in, and is extracted
from the wool, consists of the potassium salt of an organic acid
(sudoric acid). In the form of chloride and sulphate, potassium
is present in sea-water and many mineral springs. As nitrate it is
found as a crystallised efflorescence upon the soil, notably In Peru
and Chili, where it is associated with sodium nitrate. The largest

512 Inorganic Chemistry

supplies of potassium compounds are met with in the great saline
deposits of Stassfurt, where the element is found as chloride (KC1)
in sylvine^ as a double chloride of potassium and magnesium
(KCl,MgCl2,6H2O) in carnallite^ and as a mixed sulphate in kainite
(K2S04,MgS04,MgCl256H20).

Modes Of Formation.— ( i.) The method by which Davy first
effected the isolation of potassium was by the electrolysis of
potassium hydroxide : the method may be illustrated by the ex-
periment represented in Fig. 131. A small quantity of potassium
hydroxide is gently heated in a platinum capsule, which is con-
nected to the positive terminal of a powerful battery. A stout
platinum wire, flattened out at one end, is made the cathode.
When this is introduced into the fused potash, a brisk evolution
of gas takes place, and minute beads of metallic potassium make
their appearance in the liquid and upon the negative electrode,
some of which ignite upon the surface. The decomposition takes
place according to the equation —

(2.) Potassium may also be obtained by allowing melted potassium
hydroxide to pass over iron turnings heated to whiteness, when the
magnetic oxide of iron is formed —

This is known as Gay-Lussac and Thenard's method.

(3.) The method devised by Brunner, and modified by Wohler,
Deville, and others, consisted in heating to whiteness an intimate
mixture of potassium carbonate and carbon. This mixture was
obtained by first igniting in a covered iron pot crude tartar (hydro-
gen potassium tartrate, or cream of tartar), which was thereby
decomposed as indicated by the equation —

2HKC4H4O6= K2CO3 + 3C + 5H2O + 4CO.

The charred mass was then introduced into an iron retort, and
strongly heated in a furnace, when the potassium carbonate was
reduced by the carbon, as follows- —

In this process there was frequently formed variable quantities
of a highly explosive compound, owing to the union of potassium
with carbon monoxide, believed to have the composition K6(CO)C.

Potassium

513

(4.) Castner's process for the manufacture of potassium (1886)
consisted in strongly heating potassium hydroxide with a carbide
of iron, having approximately the composition CFe2.

The potassium hydroxide, with the powdered carbide of iron, was
introduced into large egg-shaped retorts, one of which is repre-
sented in Fig. 132. These retorts \vere placed upon hydraulic
lifts, so that they could be lowered away from their covers, to the
ground-level, in order to be discharged at the end of the distilla-
tion. The retorts were heated by gaseous fuel, and the metal, as

FIG. 132.

it distilled, was passed into long narrow cast-iron condensers, from
which it dropped into iron pots, and was protected from oxidation
by mineral oil. The reaction which takes place may be represented
by the equation —

(5.) At the present time potassium is obtained almost exclusively
by a modernised form of Davy's original method, namely, by
the electrolysis of fused potassium hydroxide. The process

Inorganic Chemistry

is conducted precisely as described for the manufacture of
sodium.

Properties. — Potassium is a lustrous white metal, which at
ordinary temperatures is sufficiently soft to be moulded between
the fingers ; at o° it is brittle, and shows a crystalline fracture.
The metal is readily crystallised by melting a quantity of it in a
vacuous tube, and when it has partially solidified, pouring the still
liquid portion to the other end of the tube. Potassium melts at
62.5°, and when boiled gives an emerald-green vapour. The
metal is rapidly acted on by ordinary air, its freshly cut surface
becoming instantly covered with a film of oxide, which, by absorp-
tion of atmospheric moisture and carbon dioxide, passes first into
the hydroxide and finally into the carbonate. Potassium is there-
fore usually preserved beneath naphtha, or some other liquid
devoid of oxygen.

When potassium is volatilised in a vacuous tube, the thin film of
metal which condenses upon the cool portion of the tube is seen to
possess a rich violet-blue colour, when viewed by transmitted light.
The density of potassium vapour is about 2o(Dewar and Scott), show-
ing that in the vaporous condition the molecules are monatomic.

Potassium dissolves in liquefied ammonia, forming a deep indigo
solution (page 276). When potassium is thrown upon water, that
liquid is decomposed with sufficient energy to cause the ignition
of the liberated hydrogen (page 172). When heated in carbon
dioxide, potassium takes fire, forming potassium carbonate and
carbon (page 304). Heated in carbon monoxide, it forms the ex-
plosive compound already mentioned. Potassium takes fire
spontaneously in contact with the halogens, forming the haloid
compounds of the metal. When heated in hydrogen to 360°,
potassium hydride, KH, is obtained as a white crystalline compound,
which is decomposed by moisture with evolution of hydrogen, and
takes fire spontaneously in oxygen.

Oxides Of Potassium. — When potassium is heated in ordinary
air, it takes .fire and burns, giving rise to a mixture of the oxides
of the metal. Perfectly dry air or oxygen is without action upon
potassium. The most stable oxide is the peroxide, K2O4.

Potassium Peroxide, K2O4, is formed when potassium is burnt
in oxygen. It may also be obtained by heating the metal in
nitrous oxide. It is a yellow powder which, when strongly heated,
evolves oxygen and is converted into a lower oxide. When thrown
into water a violent action takes place, oxygen and potassium

Potassium Fluoride 515

hydroxide being produced ; but by the regulated action in the
cold hydrogen peroxide is also formed —

Potassium Oxide, K2O, is obtained as a greyish-white mass by
the partial oxidation of potassium and subsequent removal of the
excess of metal by distillation in vacuo.*

By the regulated combustion of potassium in nitrous oxide Holt and Sims
have obtained compounds having the composition K2O2 and K^O3.

Potassium Hydroxide (caustic potash}, KHO, is prepared by
adding lime to a dilute boiling solution of potassium carbonate, in
iron vessels, when calcium carbonate is precipitated and potassium
hydroxide remains in solution —

K2CO3 + Ca(HO)2 = CaCO3 + 2KHO,

the reaction being complete when the addition of an acid to a
small test sample of the clear liquor produces no effervescence.
This reaction is a reversible one, and if the concentration is beyond
a certain limit, the potassium hydroxide reacts upon the calcium
carbonate, reforming potassium carbonate. The liquid is therefore
constantly maintained at a certain state of dilution during the
reaction, at the completion of which the mixture is allowed to
settle, and the clear solution is then partially concentrated in
iron vessels, and finally in silver, until on cooling the substance
solidifies. It is then usually cast into sticks. Potassium hydroxide
is now also manufactured by the electrolytic method (see Sodium
hydroxide).

Caustic potash is a white brittle solid ; it is extremely deliques-
cent, and dissolves in water with evolution of heat, forming a
highly caustic liquid. The solid, as well as the solution, readily
absorbs carbon dioxide, and is employed in the laboratory for this
purpose when it is desired to deprive a gas of the last traces of any
admixed carbon dioxide. A hot saturated solution of potassium
hydroxide, when cooled, deposits crystals of a hydrate having the
composition KHO,2H2O.

Potassium Fluoride, KF.— This salt is prepared by neutralising
aqueous hydrofluoric acid with potassium carbonate, and evaporat-
ing the solution in a platinum vessel, when the salt is obtained in
the form of deliquescent cubical crystals. Potassium fluoride dis-
solves in aqueous hydrofluoric acid with evolution of heat, forming
the acid fluoride of potassium, HF,KF, which is obtained as an
anhydrous salt when the solution is evaporated to dryness and

* Rengade, Compt. Rend., 1906.

516 Inorganic Chemistry

heated to 110°. This salt is not deliquescent. When heated to a
dull red heat it decomposes into the normal salt and hydrofluoric
acid (see p. 350).

Potassium Chloride, KC1. — This salt is found in sea-water, and
was at one time obtained as a secondary product in the manufacture
of bromine from sea salt, and of iodine from seaweed, as well as in
various other industrial processes. At the present day it is almost
exclusively obtained from the enormous deposits of carnallite at
Stassfurt. The method by which potassium chloride is obtained
from this double salt, KCl,MgCl2,6H2O, is based upon the fact,
that when dissolved in water the salt dissociates into its two
constituents ; and when the solution is concentrated, the more
insoluble potassium chloride first separates out, leaving the mag-
nesium chloride in solution.

In practice, the crushed crude carnallite is treated with boiling
mother-liquors from previous operations, in large tanks into which
steam can be driven. These mother-liquors are practically a
strong solution of magnesium chloride, and it is found that while
potassium chloride is readily soluble in this liquid, the sodium
chloride and magnesium sulphate which are present in the crude
carnallite are only slightly dissolved by it, and are therefore left
behind in the residue.

The muddy liquid is allowed to settle for about an hour, when it
is drawn off into large iron crystallising tanks. The salt which is
then deposited contains from 80 to 90 per cent, of potassium chloride,
the remainder being mainly sodium and magnesium chlorides.

The mother-liquor from these crystallising tanks is either used
again for treating a fresh charge of mineral, or is further evaporated,
when crystals of carnallite separate out ; for it is found that when
the amount of magnesium chloride present is greater than three
times the proportion of potassium chloride in the solution, the liquid
on crystallising deposits the double chloride of the two metals. The
impure potassium chloride from the crystallising tanks is purified by
washing with cold water, in which the salt is only slightly soluble,
and by subsequent recrystallisation. Potassium chloride crystal-
lises, like the chlorides of sodium, rubidium, and caesium, in cubes.

Potassium Chlorate, KC1O3. — When chlorine is passed into a
solution of potassium hydroxide, a mixture of potassium chlorate
and chloride is obtained, thus —

Potassium Chlorate

517

The two salts in solution may be separated by crystallisation,
the chlorate being much less soluble in cold water than the
chloride.

On the manufacturing scale, potassium chlorate is obtained by
passing chlorine "into milk of lime, when a mixture of calcium
chlorate and chloride is formed —

6Ca(H O)2 + 6C12 = Ca(ClO3)2 + 5CaCl2 + 6H2O.

The operation is conducted in cast-iron cylinders connected
in series, one of which is shown in section in Fig. 133, furnished

FIG. 133.

with mechanical stirring gear, a, b, b. The shaft and the pipes
conveying the chlorine into and from the vessel are connected to
it by means of the water-sealed joints, c, *?, e. The manhole/ is a
short wide leaden pipe, dipping a few inches into the liquid, which
allows of the periodic withdrawal of samples for examination.
Several reactions are involved in the final formation of the calcium
chlorate ; in the first case calcium hypochlorite is produced,
thus—

+ CaCl2+2H2O.

The calcium hypochlorite then passes into a mixture of chlorate
and chloride in accordance with the equation —

3Ca(OCl)2 = Ca(C103)2 + 2CaCl2.
The second change is brought about by the operation of two

518 Inorganic Chemistry

causes, namely, rise of temperature and the presence of excess of
chlorine. Heat alone is incapable of converting more than a
small proportion of the hypochlorite into chlorate, for the former
compound is at the same time decomposed into calcium chloride
and free oxygen. The excess of chlorine is believed to act, through
the intervention of hypochlorous acid, HOC1, merely as a carrier
of oxygen, reducing two molecules of calcium hypochlorite to
chloride, and oxidising the third to chlorate, thus —

2CaCl2 + Ca(ClO3)2 + 2C12 + 2 H2O.

The absorption of chlorine by the milk of lime is attended with
evolution of heat ; care is taken to prevent the temperature from
rising above about 70°, otherwise loss results by the decomposition
of hypochlorite with evolution of oxygen, thus —

Ca(OCl)2 = CaCl2 + O2.

When the formation of calcium chlorate is complete, the liquid
is allowed to settle, and is then run into concentrating pans, where
the requisite amount of potassium chloride in solution demanded
by the following equation is added —

Ca(ClO3)2 + 2KC1 = CaCl2 + 2KC1O3.

The liquid is then concentrated in iron pans and allowed to
crystallise, when the moderately soluble potassium chlorate sepa-
rates out, leaving the very soluble calcium chloride in solution.
The chlorate is afterwards purified by recrystallisation.

Potassium chlorate, although only moderately soluble in water, is much
more soluble in a strong solution of calcium chloride, hence there is always a
loss (usually about 10 per cent. ) of chlorate in this process. Pe'chiney's pro-
cess for obviating this consists in concentrating the liquid obtained by the
chlorination of the lime to a definite specific gravity, and then cooling it to
12°, when about 78 per cent, of the calcium chloride crystallises out. The
mother-liquor, containing all the calcium chlorate and only the comparatively
small proportion of calcium chloride, is then treated with potassium chloride
as usual.

Like so many of the older manufacturing processes, this for
the preparation of potassium chlorate is now being displaced by
modern electrolytic methods. A solution of potassium chloride is

Potassium Per chlorate 519

electrolysed in an undivided cell ; the anode consisting of a thin
sheet of platinum, and the cathode being a vertical grid of copper
wire. The solution is caused to flow continuously through the
electrolytic cell, the rate of flow being so regulated that the tem-
perature of the liquid is maintained at about 50° C., and that the
proportion of chlorate produced does not rise above 3 per cent, in
the liquid. The dilute liquor is passed into suitable refrigerators,
where the sparingly soluble chlorate crystallises out. The chemical
action may be regarded as taking place in stages ; the chlorine
liberated at the anode there unites with oxygen and water, yielding
hypochlorous acid —

At the same time potassium hydroxide is produced at the cathode,
with elimination of hydrogen. The caustic potash coming in contact
with hypochlorous acid, or with chlorine, gives rise to potassium
hypochlorite, which reacting with hypochlorous acid produces
potassium chlorate —

KC1O + 2HC1O = KC1O3+2HCI.

Potassium chlorate crystallises in white tables, belonging to the
monosymmetric system, which when of large size often exhibit fine
iridescent colours. 100 parts of water at o° dissolve 3.3 parts of
the salt ; while at 100°, 59 parts are dissolved. •

Potassium chlorate is used largely in the manufacture of matches,
on account of the ease with which it gives up its oxygen : thus, if a
small quantity of the finely powdered salt be carefully mixed with
an equally small amount of red phosphorus, the friction caused by
lightly rubbing it with a spatula is sufficient to cause the mixture
to detonate violently. Similarly, when powdered potassium chlo-
rate and sulphur are rubbed together in a mortar, the mixture
explodes with violence. Potassium chlorate is also largely em-
ployed in pyrotechny, especially in the production of coloured
effects, where a fiercely burning mixture is required.

Potassium chlorate melts between 360° and 370°, and at a tern-
perature about 380° begins to evolve oxygen.

Potassium Perehlorate, KC1O4. — When the chlorate is heated.
it first melts and begins to give off oxygen ; but it soon begins to
partially solidify, owing to the formation of potassium perchlorate,
and the evolution, of oxygen stops unless a stronger heai be

520 Inorganic Chemistry

applied. The reaction at this stage may be expressed by the

equation —

8KC1O3 = 5KC1O4 + 3KC1 + 2O2.

The evolution of oxygen, however, is not an essential condition of
the formation of the perchlorate. By careful regulation of the tem-
perature the following decomposition can be made to take place —
4KC1O3 = KC1 + 3KC1O4.

The perchlorate is separated by first treating the residue with
cold water, which dissolves the greater part of the chloride, and
afterwards with warm hydrochloric acid, which decomposes any
remaining chlorate. The salt is then purified by crystallisation.

Potassium perchlorate is very slightly soluble in cold water, 100
parts of water at o° dissolving only 0.7 part of the salt ; while at
100°, 20 parts are dissolved.

Potassium Bromide, KBr, and Iodide, Kl.— These two salts
are obtained by similar methods. When bromine or iodine is
added to a solution of potassium hydroxide, the reaction which
takes place is exactly analogous to that in the case of chlorine
(see Potassium Chlorate, above) —

If the solution so obtained be evaporated to dryness, and the
dry residue ignited, the bromate (or iodate) is decomposed, just as
potassium chlorate is decomposed by heat, giving off its oxygen,
and being converted into bromide (or iodide)—
KBr03 = KBr + 3O.

The residue, on being dissolved in water and recrystallised,
yields pure potassium bromide (or iodide).

These salts are manufactured by decomposing ferrous bromide,
Fe3Br8 (or iodide, Fe3I8), with potassium carbonate, thus —
Fe3Br8 + 4K2CO3 = Fe3O4 + 8KBr + 4CO2.

The ferrous bromide is obtained by adding bromine to moistened
iron borings (see Manufacture of Bromine).

Potassium iodide and bromide both crystallise in cubes, and are
both readily soluble in water. These salts are chiefly used for
medicinal and photographic purposes.

Potassium Sulphate, K2SO4.— This salt is present in the Stass-
furt deposits principally as kainite, K2SO4,MgSO4,MgC]2,GH2O,
and as polyhalite, K2SO4,MgSO4,2CaSO4,2H2O. When kainite
is treated with small quantities of water, or mother-liquors from
other processes, the extremely soluble magnesium chloride is

Potassium Carbonate 521

removed, leaving the potassium magnesium sulphate ; and on
adding to this the requisite amount of potassium chloride, the
following change takes place —

K2SO4,MgSO4+3KCl = 2K2SO4 + KCl,MgCl2.

From this solution the potassium sulphate crystallises out.

Potassium sulphate is also obtained by the action of sulphuric
acid upon the chloride, by a process corresponding exactly to the
first stage in the Leblanc soda process (q.v.} —

2KC1 + H2SO4=K2SO4 + 2HC1.

Potassium sulphate forms colourless rhombic crystals, contain-
ing no water of crystallisation, therein differing from sodium
sulphate, which crystallises with ten molecules of water.

Potassium sulphate is largely used for agricultural purposes.

Potassium Carbonate, K2CO3. — This salt was formerly obtained
exclusively from the ashes of wood and other land plants, and was
known under the name of pot-ashes. The process is still carried
on in parts of Canada and the United States. The wood is burned
in pits, and the ashes are collected and lixiviated with water
(with the addition of a small quantity of lime) in wooden tubs
with perforated false bottoms. The liquid which is drawn off is
evaporated to dryness, and usually calcined to burn away the
organic matter. This material, known as American pot-ashes,
contains varying quantities of caustic potash, on account of the
previously added lime. The so-called American pearl-ash is a
purer product, obtained by concentrating the liquor from the
lixiviating tubs until the less soluble impurities crystallise out,
and finally evaporating the mother-liquor, containing the potassium
carbonate, to dryness, and calcining the residue.

Potassium carbonate is also obtained from beet-root molasses,
an uncrystallisable residue obtained in the manufacture of beet
sugar, carried on chiefly in France. The syrup is fermented with
yeast, whereby the sugar it contains is converted into alcohoL and
then distilled. The residual liquid, known as vinasse, is evaporated
to dryness ; and from the black residue, termed " vinasse cinder,"
the potassium carbonate is extracted.

Potassium carbonate is obtained also from suint^ which, as
already stated, contains considerable quantities of potassium in
the form of potassium sudorate. The sheep's wool is lixiviated

5 22 Inorganic Chemistry

with water, and the solution evaporated to dry ness. The residue
is heated in iron retorts, whereby the organic potassium salts are
converted into carbonate, while, at the same time, ammonia and
an illuminating gas are evolved. The carbonaceous residue is
extracted with water, and the potassium carbonate separated by
crystallisation.

Since the development of the Stassfurt potash supplies, these
sources of potassium carbonate are rapidly sinking into the back-
ground, and the bulk of this compound is now manufactured from
potassium sulphate by a process similar to the Leblanc soda
process (g.v.).

Potassium carbonate is not manufactured by a method analogous
to the ammonia-soda process (Solvay), on account of the too great
solubility of potassium bicarbonate (hydrogen potassium carbonate).

Pure potassium carbonate may be obtained by igniting cream
of tartar (see page 512), and extracting with water ; or by heating
hydrogen potassium carbonate, which gives up water and carbon
dioxide, thus —

2HKCO3=K2CO3

Potassium carbonate forms long prismatic crystals belonging to
the monosymmetric system, and containing three molecules of
water, K2CO3,3H2O. The anhydrous salt is highly deliquescent,
and very soluble in water.

Hydrogen Potassium Carbonate (bicarbonate of potash\
HKCO3, is produced by passing carbon dioxide into an aqueous
solution of the normal carbonate, thus —

This salt is much less soluble in water than the normal salt, and is
readily purified by crystallisation.

Potassium Nitrate (nitre, saltpetre}, KNO3.— This salt has
been known since very early times. It occurs as an efflorescence
upon the earth, as a result of the oxidation of organic nitrogenous
matter in the presence of the potash in the soil, and is found in
the neighbourhood of villages, more especially in hot climates,
where urine and other readily decomposable organic matters rich
in nitrogen find their way into the surface soil. It has been shown
that the process of nitrification which results in the formation of
nitre under these circumstances is due to the action of specific
organisms, or microbes, and never takes place in their absence.

Potassium Nitrate 523

At various times this natural process has been artificially carried
on, by mixing manure and other decomposing refuse with porous
soil, lime, and wood ashes, and exposing the mixture in heaps
which were moistened from time to time with drainage from
manure. The saltpetre earth, collected from the natural sources
or from the artificial nitre plantations, on lixiviation with water, and
subsequent evaporation, yielded crystals of potassium nitrate.

At the present time potassium nitrate is almost exclusively ob-
tained from sodium nitrate (Chili saltpetre), by treatment with
potassium chloride derived from the Stassfurt supplies. The requi-
site quantities of the two solutions are run into a tank, and heated
by means of steam, when the following double decomposition takes
place —

NaNO3+KCl = NaCl + KNO3.

The greater part of the sodium chloride is at once precipitated,
and is removed by canvas filters. The clear liquid is then allowed
to crystallise in tanks furnished with stirring gear, in order to
cause the formation of small crystals, and the nitre-meal so ob-
tained is purified by recrystallisation.

Potassium nitrate crystallises usually in rhombic prisms, but it
can also be obtained in the form of small rhombohedral crystals,
isomorphous with sodium nitrate.

The solubility of potassium nitrate rapidly increases with rise of
temperature (see Solubility Curve, p. 152). 100 parts of water at
o° dissolve 13.3 parts ; at 50°, 86 parts ; and at 100°, 247 parts.

Nitre melts at 339°, and at a higher temperature loses oxygen
and is converted into potassium nitrite ; on this account it readily
oxidises many of the elements when heated in contact with them.
Thus, a fragment of charcoal or sulphur thrown upon melted nitre
takes fire and burns with great energy ; in the one case with forma-
tion of potassium carbonate and carbon dioxide, and in the other
of potassium sulphate and sulphur dioxide —

= 2K2CO3 + 3CO2+2N2.
2KNO3+2S= K2SO4+ SO2 + N2.

Nitre is chiefly used in the manufacture of gunpowder and in
pyrotechny.

Gunpowder is a mixture of nitre, charcoal, and sulphur. The proportions
in which these ingredients are present varies, within small limits, according

524 Inorganic Chemistry

to the special kind of powder, as will be seen from the following table (Abel
and Nobel), giving analyses 'of various powders manufactured at Waltham
Abbey.

' •

Fine-Grain.

Rifle
Fine-Grain.

Rifle
Large-Grain.

Pebble
Powder.

Potassium nitrate

73-55

75-04

74-95

74.67

,, sulphate

0.36

0.14

0.15

0.09

Sulphur

IO.O2

9-93

10.27

10.07

Charcoal .

14-59

14.09

I3-52

14.22

Water

;

I.48

0.80

I. II

0-95

These proportions are very close to those which would be demanded by the
equation —

which was at one time supposed to represent the change which takes place
when gunpowder is exploded. In reality the decomposition is much more
complex, and it has been shown that the solid products consist of mixtures of
the following substances in varying proportions, depending upon the particular
powder, and the conditions of firing —

Potassium nitrate.
,, oxide.

Ammonium sesquicarbonate.
Carbon.
Sulphur.

Potassium carbonate.
, , sulphate.
,, sulphide.
,, thiosulphate.
, , thiocyanate.

While the gases that are evolved consist of—

Carbon dioxide. Marsh gas.

Nitrogen. Oxygen.

Carbon monoxide. Hydrogen.

Sulphuretted hydrogen.

From the combustion of one gramme of powder the total weight of solids
ranges from 0.55 to 0.58 gramme, and the total weight of the gaseous products
from 0.45 to 0.42 gramme.

Potassium Cyanide, KCN.— This salt is manufactured from
potassium ferrocyanide, which is first obtained by heating in an
iron pot a mixture of scrap iron, crude potashes, and waste animal
refuse, such as hoofs, horns, hide, £c. The complex changes which
take place do not at once result in the formation of the ferrocyanide,
as this salt is unstable at high temperatures, but in the production
of various compounds (the very stable salt potassium cyanide
amongst them) which, when the mass is subsequently treated with
water, interact, with the formation of potassium ferrocyanide. The

Compounds of Potassium with Sulphur 525

aqueous extract is allowed to crystallise, and the ferrocyanide is
obtained as large lemon-yellow prisms, with three molecules of
water. When this compound is dried and heated alone it decom-
poses into potassium cyanide, free nitrogen, and a carbide of iron —

K4Fe(CN)6 = 4KCN + N2 + FeC2.

By heating the ferrocyanide with potassium carbonate a larger
yield of the cyanide is obtained, mixed with potassium cyanate —

K4Fe(CN)6+K2CO3 = 5KCN + KCNO + Fe + CO2.

For many commercial uses for which potassium cyanide is
required the presence of this cyanate is not detrimental.

If potassium ferrocyanide be heated with metallic sodium the
whole of the cyanogen it contains is converted into alkali cyanide
(Erlenmeyer) —

K4Fe(CN)6 + 2Na = 4KCN + 2NaCN + Fe.

The mixed potassium and sodium cyanide thus obtained is well
suited for the technical processes for which cyanide is required.

Potassium cyanide is a white solid which is extremely soluble in
water, from which it crystallises in white anhydrous plates. When
heated the salt readily fuses, but is stable at very high tempera-
tures, being capable of being volatilised without decomposition. In
the fused state it is a powerful reducing agent, taking up oxygen
to yield potassium cyanate, KCNO.

COMPOUNDS OF POTASSIUM WITH SULPHUR.
Four sulphides of potassium have been obtained, namely —

Potassium monosulphide . . V ... . K2S

Potassium trisulphide .-'•.- . . . . K2SS

Potassium tetrasulphide . . . . . . K2S4

Potassium pentasulphide . . . . . . K2S5

Just as potassium decomposes water with evolution of hydrogen
and formation of potassium hydroxide, so also, when heated in
sulphuretted hydrogen (the sulphur analogue of water) it forms
potassium hydrosulphide (the analogue of potassium hydroxide)
and liberates hydrogen, thus —

526 Inorganic Chemistry

When potassium hydroxide and hydrosulphide are mixed in equi-
molecular proportions, potassium monosulphide and water are
formed—

The liquid, on evaporation in vacuo, deposits reddish prismatic
deliquescent crystals having the composition K2S,5H2O.

When potassium carbonate and sulphur are heated together a
mixture of the higher sulphides of potassium with potassium thio-
sulphate is obtained, thus —

3K2CO3+ 8S = 2K2S3 + K2S2O3
3K2CO3H-12S = 2K2S6+K2S2O3

The reddish - brown solid product was named by -the early
chemists hepar sulphuris, or " liver of sulphur."

SODIUM.

Symbol, Na = 23.00.

Occurrence. — The most abundant natural compound of sodium
is the chloride, which is present in sea-water and in many salt
lakes and springs. Enormous deposits of sodium chloride or
rock-salt are found in Cheshire, Lancashire, and other parts of
the world. As nitrate, this element occurs in large quantities in
Chili and Peru, and in combination with silicic acid it is a con-
stituent of many rocks.

Modes Of Formation. — Sodium was first isolated by Davy, by
the electrolysis of sodium hydroxide. On a manufacturing scale
it has been obtained by the various processes described under
potassium, the history of the commercial preparation of these two
elements being practically identical. Sodium, however, does not
form any explosive compound with carbon monoxide, so that the
manufacture in this case has been free from this difficulty.

At the present time sodium (and also potassium) is almost ex-
clusively obtained by electrolytic methods.

(i.) Castner's Process. — This method consists in the electrolysis
of fused sodium -hydroxide, and is, in fact, simply the original
process by which Davy first obtained the metal adapted to modern
resources of electrical power. The apparatus employed *s shown
in section in Fig. 134. The caustic soda is contained in an iron

Sodium

527

pot P, set in suitable brick-work, and is kept in a melted state by
a ring of gas flames below. Through the bottom of this vessel
passes the cathode, which is maintained steady in its position by
the caustic soda in the lower and narrow part of vessel P being in
the solidified state. The anodes A are suspended from above
round the cathode, and are prevented from touching it by means
of a wirework cylinder which hangs from the vessel V. This vessel
is an iron cylinder having a lid at the top, and is the receiver in
which the sodium collects.

The products of the electrolysis are oxygen, hydrogen, and
sodium. The oxygen liberated at the anodes escapes by the
opening O in the lid. The
sodium floats up to the sur-
face of the molten caustic
in the receiver V, and is
withdrawn from time to
time by means of a per-
forated ladle, which allows
the caustic to drain through,
but holds the liquid metal,
owing to the extremely high
surface tension of the latter.
The hydrogen which is also
liberated at the cathode
escapes through the loosely
fitting lid of the receiver.*

(2.) Borcherf Process. —
It will be evident from an

economic (and therefore the manufacturer's) point of view that the
hydrogen liberated in the above process represents wasted electrical
energy. Many attempts, therefore, have been made to substitute
fused sodium chloride for the hydroxide. The practical difficulties to
be overcome in this case are more serious, owing partly to the higher
temperature required, and also to the corrosive action exerted by
the fused chloride upon the materials of which the vessels are
constructed. On the other hand, it will be evident that both the
products of the electrolysis in this case will have commercial
value. Borchers' apparatus is shown in section in Fig. 135. It

FIG. 134.

* This process is extensively employed at Oldbury^ near Birmingham, and
at the works of the Niagara Electrical Company.

528

Inorganic Chemistry

consists essentially of a U-tube made in two parts, the wide limb
being of fireclay and the narrow part of iron, the two parts being
clamped together. To prevent leakage at the joint, a square
tube W is interposed between the two parts, which is kept cool by
a flow of water through it. This causes the solidification of the
sodium chloride in the form of a layer all round the junction.
Chlorine is liberated at the anode and escapes by the pipe P.
The narrow limb C is itself the cathode, and the sodium there

FIG. 135.

produced overflows down the side pipe into a suitable receiver.
Fresh sodium chloride is added as required through the tube D.

Properties. — Sodium closely resembles potassium in its general
properties. It is a soft, white metal which can be readily moulded
by the fingers, and is easily pressed into wire. At - 20° it is hard.
The colour of sodium vapour is violet, while the colour exhibited
by a thin film of the metal, obtained by sublimation in vacuo, is
greenish-blue. The vapour-density of sodium is about 12 (Dewar
and Scott), showing that this metal in the vaporous state is
monatomic.

Like potassium, sodium dissolves in liquid ammonia, yielding a

Sodium Peroxide 529

blue solution. When heated in the air, sodium burns, forming
the peroxide, Na2O2. Perfectly dry air or oxygen is without
action upon the metal.

When heated in hydrogen, sodium forms the hydride, NaH>
analogous to the potassium compound, but not spontaneously
inflammable in air. When this is heated to about 300° in vacuo
the whole of the hydrogen is evolved.

Alloy of Sodium and Potassium.— When these two metals are
melted together beneath petroleum an alloy is obtained which is
liquid at ordinary temperatures. When prepared and preserved*
out of contact with air the alloy resembles mercury in appearance.
This alloy is employed in the construction of thermometers for
registering high temperatures, where mercury would be inad-
missible.

Oxides Of Sodium. — Two oxides are known, viz., sodium
monoxide, Na2O, and sodium dioxide, or peroxide, Na2O2. Of
these the peroxide is the more important.

Sodium Oxide (or sodium monoxide), Na2O, is produced as a
white amorphous compound when sodium is partially oxidised in
a limited supply of oxygen, and the excess of the metal subse-
quently removed by distillation in vacuo.*

Sodium Peroxide, Na2O2, is obtained by allowing sodium to
burn briskly in oxygen. It is a yellowish-white solid, which de-
composes in contact with water, with considerable rise of tem-
perature and evolution of oxygen —

Na2O2+H2O = 2NaHO + O.

The oxygen which is evolved contains appreciable quantities of
ozone. When sodium peroxide is slowly added to water or to dilute
hydrochloric acid in the cold, hydrogen peroxide is formed —

Owing to the readiness with which it gives up oxygen, sodium
peroxide is a powerful oxidising agent, and as such finds many
uses in the laboratory. Thus it readily converts chromic com-
pounds into chromates.

Sodium peroxide forms a crystalline hydrate of the composition,
Na2O2,8H2O (page 228). When heated in either nitrous or nitric

* Rengade, Compt. Rend.) ic,o6.

2 L

530

Inorganic Chemistry

oxides it yields sodium nitrite ; in the former case with the elimina-
tion of nitrogen —

Na2O2-f2N2O = 2NaNO2 + N2.
Na2O2 + 2NO = 2NaNO2.

Sodium Hydroxide (caustic soda), NaHO.— This compound is

produced when sodium is brought into contact with water, and also

when either sodium monoxide or peroxide is dissolved in water.

On the large scale caustic soda is prepared by the action of lime

.upon a boiling solution of sodium carbonate (see Caustic Potash).

The so-called tank liquors (obtained in the manufacture of sodium
carbonate by the Leblanc process, q.v.) are heated to the boiling-
point, and an excess of lime is stirred into the mixture. The
sodium sulphide present in the tank liquor is oxidised into sulphate

FIG. 136.

by the combined action of air injected into the mixture, and ot
sodium nitrate, which is added for this purpose. The liquor, after
being causticised, is decanted or filtered from the precipitated
calcium carbonate, and is concentrated in large cast-iron hemi-
spherical pans. The decomposition suffered by the sodium nitrate
depends upon the temperature and concentration of the liquid ; at
300° to 360° the change may be expressed by the equation —

NaNO3-l-2H2O = NaHO4-NH3 + 4O.

The liberated oxygen oxidises the sulphides to sulphates.

Caustic soda is now being manufactured by the electrolysis of
brine. The apparatus devised by Castner for this purpose is seen
in Fig. 136. It consists of a rectangular vessel divided into three
compartments. Upon the floor of the vessel there is a layer of

Sodium Chloride 53 T

mercury about £th of an inch deep. The partitions, which are
non-porous, dip into narrow gutters across the bottom, but do not
actually touch the» bottom, so that when the tank is gently oscil-
lated the mercury can flow from one compartment to the other,
while the liquid above is prevented from so doing. The two
outside compartments are rilled with brine, while the centre one
contains water ; and in this is placed the cathode, consisting of a
number of metal plates. Since the partitions are non-porous the
current will pass from the carbon anodes through the salt solution
to the mercury, which in the two extreme compartments then
becomes the cathode. It then passes from the mercury in the
middle space, which now becomes the anode of this cell, through
the aqueous liquid to the metal cathode which is there suspended.
In the outside compartments the sodium chloride is electrolysed ;
the chlorine discharged at the carbon anodes escapes by the pipes
P P, while the sodium dissolves in the mercury cathodes. During
the process a slow rocking movement is given to the tank by means
of the excentric represented at E, whereby the mercury is caused
to flow to and fro along the bottom. In the middle compartment
the sodium contained in the amalgam is transported to the cathode,
where it dissolves in the water, forming sodium hydroxide.

Sodium hydroxide is a white, strongly caustic, and highly de-
liquescent solid. It is soluble in water, with considerable rise of
temperature, and a concentrated aqueous solution when cooled
to -8°, deposits a crystalline hydrate, having the composition
2NaHO,7H2O.

Sodium Chloride, NaCl. — Of the compounds of sodium with
the halogens the chloride is the most important. In warm
climates, as upon the shores of the Mediterranean, sodium chloride
is obtained by the evaporation of sea-water in large shallow basins
or pools, constructed upon the sea-shore and exposed to the sun's
heat. As the brine concentrates in these salterns, the crystals
of salt are raked off the liquid and allowed to drain in heaps at
the side of the pools. The mother-liquors, known as bittern,
were formerly utilised for the extraction of the bromine which they
contain.

Salt is obtained from salt-beds, where it is found in enormous
deposits, either by direct mining operations, when the salt is
sufficiently pure, or by first dissolving the material in water,
whereby insoluble admixed impurities are removed, and afterwards
evaporating the brine so obtained. The latter method is carried

532

Inorganic Chemistry

out by sinking borings through the upper strata of rock, and
sending water down to the salt-beds beneath. The brine is then
pumped up and the salt obtained by evaporation. The first stage
of the concentrating process, especially where the brine is not very
strong, is in some parts carried on by exposing the liquid to the

FIG. 137.

wind. This is effected by causing the solution to trickle over
erections of brushwood known as graduators (Fig. 137), which are
built so that the prevailing winds blow across them. The brine is
pumped up into the wooden troughs running along the top, from
which it escapes by a number of openings, a, a, a, and flows over
the pile of brushwood down into the reservoir upon which the

Sodium Chloride

533

erection is constructed. In this way the solution is made to
expose a large surface to the air, and it quickly reaches a concen-
tration when it contains 20
per cent, of salt in the
solution. The liquor is then
evaporated in shallow iron
pans by means of artificial
heat, and as the salt crys-
tallises it is lifted out by
means of perforated iron
skimmers. Salt obtained
in this manner always con-
tains small quantities of
other salts, such as sodium
sulphate, calcium sulphate,
calcium and magnesium
chlorides. The presence
of chlorides of magnesium
or calcium causes the salt
to become moist, especially
in damp weather.

Pure sodium chloride
may be prepared by add-
ing hydrochloric acid to a
strong aqueous solution of
salt ; the sodium chloride is
thereby precipitated, while
the other salts remain in
solution.

Sodium chloride forms
colourless, cubical crystals,
which are anhydrous. If
deposited at — 10° it crystal-
lises in monosymmetric
prisms, with two molecules
of water of crystallisation,
which at the ordinary temperature lose their water and break up
into minute cubes. '

Sodium chloride is a necessary article of food for man and other
animals ; it is estimated that about 20 Ibs. of salt per head of
population is annually used, directly or indirectly, for this purpose.

534 Inorganic Chemistry

The hydrochloric acid present in the gastric and other acid fluids
of the stomach is derived from the decomposition of sodium chloride
which is taken into the organism.

Enormous quantities of sodium chloride are employed in the
alkali industry, and all the chlorine that is manufactured is derived
primarily from this compound.

Sodium Bromide, NaBr, and Sodium Iodide, Nal, are pre-
pared by methods similar to those for obtaining the potassium
compounds. They are both isomorphous with sodium chloride,
and when deposited at low temperatures they form monosymmetric
crystals containing two molecules of water.

Sodium Carbonate, Na2CO3.— The preparation of this com-
pound is carried on by three methods, and constitutes that important
industry, the alkali manufacture. Two of these processes are known
by the names of their respective discoverers, namely, the Leblanc
process and the Solvay process, the latter being also known as
the ammonia-soda process. The third is a modern electrolytic
method.

I. The Leblanc method of manufacture consists essentially of
three processes, namely —

(i.) The conversion of sodium chloride into sodium sulphate
by the action of sulphuric acid, known as the salt-cake
process. Two chemical reactions are involved in the
process —

NaCl + H2SO4 =NaHS04 + HCl.
NaCl + NaHSO4=Na2SO4 +HC1.

(2.) The decomposition of sodium sulphate, salt-cake^ by means
of calcium carbonate (limestone) and coal, at a high
temperature, whereby a crude mixture of sodium car-
bonate and calcium sulphide is obtained, known as
black-ash. This black-ash process takes place in accord-
ance with the following equation —

Na2SO4+CaCO3 + 2C = Na2CO3 + Ca$ + 2CO2.

The change may be conveniently regarded as taking place in two
stages, which proceed simultaneously according to the equations —

Na2SO4+2C = Na2S-f 2CO2. .
Na2S + CaCO3=CaS + Na2CO3.

fc.) The process of extracting and purifying the sodium car*
bonate contained in the black-ash.

Sodium Carbonate

535

(I.) The Salt-cake Process.— The first stage of this process is
usually carried on in a large cast-
iron pan (Dt Fig. 138), built into
a furnace in such a manner that
it shall be heated as uniformly as
possible. The charge of common
salt is placed in the covered pan,
and the requisite quantity of sul-
phuric acid is then run in. Hydro-
chloric acid is given off in tor-
rents, according to the first of the
above equations, and the gas is
led away by the pipe E in the
arched roof to the condensing-
towers, where it is absorbed by
water (see Hydrochloric Acid,
page 369). The mixture is heated
until it begins to stiffen into a
solid mass, when the damper h
is raised and the mass is raked
out of the pan on to the hearth
of the roaster or reverberatory
furnace, b. Here it is exposed
to the hot gases from the coke fire
«, which sweep over it and ulti-
mately raise its temperature nearly
to a red heat, whereby the second
of the above reactions is com-
pleted. The acid gas, together
with the fire gases, leave the
roaster by the chimney ^, and are
also led to condensing - towers,
where the hydrochloric acid is
absorbed. The mass is from time
to time raked or worked by means
of side - openings or "working
doors " in the roaster, and as soon
as the operation is completed the
salt-cake is withdrawn. The salt-
cake so obtained usually contains
from 95 to 96 per cent, of normal

536

Inorganic Chemistry

sodium sulphate, Na.2SO4 ; the remaining 4 or 5 per cent, consist-
ing of hydrogen sodium sulphate, NaHSO4, undecoinposed sodium

chloride, and such impurities as werfe
originally present in the salt.

(2.) The Black-ash Process.— The
salt-cake is mixed with limestone (or
chalk) and coal dust (slack\ and
heated in a reverberatory furnace
known as the black -ash or balling
furnace. As the mixture softens with
the heat it requires to be thoroughly
mixed together, which, in the older
forms of furnace (still used in many
places), is accomplished by manual
labour. Fig. 139 shows such a furnace
in section. The materials are intro-
duced by the hopper k on to the
hearth /, where they are exposed to
the hot gases from the fire a ; and
as the decomposition proceeds they
are raked along to the more strongly-
heated front portion of the hearth h.
During this process carbon dioxide
is freely evolved, the escaping bubbles
of gas giving the semi-fluid mass
the appearance of boiling. As the
temperature rises and the process
approaches completion, the mass
thickens, when it is worked up into
large balls by means of rakes or
paddles. At this stage carbon mon-
oxide begins to be evolved, the bubbles
of which, bursting from the doughy
material, become ignited and burn
upon its surface as small jets of flame
coloured yellow by the soda. As soon
as these appear the ball is quickly withdrawn from the furnace.
The formation of carbon monoxide at the high temperature reached
at this point in the process is due to the action of carbon upon the
limestone according to the equation —

CaCO3 + C = CaO + 2CO,

Sodium Carbonate 537

excess of these materials being intentionally present in the mixture.
The effect of the escaping carbon monoxide at this point in the
process, in rendering the black-ash light and porous (an important
consideration in view of the next operation), is similar to that of
baking-powder when used for cooking purposes. The heated
gases from the furnace are made to pass over large evaporating
pans, P, where liquors from a subsequent process are concen-
trated.

In the more modern forms of black-ash furnace, the mixing and
working up of the materials is accomplished mechanically by
means of a revolving hearth. Fig. 140 shows the general arrange-
ment of a revolving black-ash furnace. The mixture is placed in
the cylinder ^, which is made to slowly revolve upon its horizontal
axis. The heated gases from the fire a pass through this revolv-
ing hearth ; they are then conveyed through a dust-chamber, m, and
finally over concentrating-pans. Limestone and two-thirds of the
coal are first thrown into the furnace and heated until the blue
flame of burning carbon monoxide makes its appearance, when
the salt-cake along with the rest of the coal is added, and the
process continued until the yellow flames appear upon the surface
of the mass. The contents of the cylinder are then thrown out
into iron trucks beneath.

Black-ash is a mixture of variable composition, containing—

Sodium carbonate, Na2CO3 . from 40 to 45 per cent.
Calcium sulphide, CaS . . „ 30 „ 33 „
Calcium carbonate, CaCO3 . „ 6 „ 10 „
Coke . . . ., ,. „ 4 „ 7 „

Calcium oxide, CaO A- '•-.. ..'. „ 2 „ 6 „

And smaller quantities of sodium chloride, sodium sulphate, sodium
sulphite, sodium sulphide, sodium thiosulphate, oxides of iron
and alumina.

(3.) Lixiviation of Black-ash. — The lixiviation of black-ash is
carried on in a series of tanks, so arranged that the liquid can be
made to pass from one to the other. The action of water upon
the black-ash is more than a simple process of dissolving the
sodium carbonate from the mixture, for in the presence of water
chemical action takes place between some of the ingredients.
Thus the lime reacts upon sodium carbonate, forming sodium
hydroxide, hence the tank liquor always contains caustic soda in
varying quantities. Under certain conditions of temperature and

538 Inorganic Chemistry

dilution, the calcium sulphide also reacts upon the sodium car
bonate, fortning sodium sulphide and calcium carbonate, thus —

CaS + Na2C03 = CaCO3 + Na2S.

Also by the oxidising influence of atmospheric oxygen, calcium
sulphide, CaS, is converted into calcium sulphate, CaSO4, which
in its turn is acted upon by the sodium carbonate, involving loss
of this product —

CaSO4+Na2CO3 = CaCO3 + Na2SO4.

The process of lixiviation is carried on as rapidly as possible,
and at temperatures ranging from about 30° (for the dilute liquors)
to about 60° (for those more concentrated) ; for the formation of
sodium sulphide diminishes as the concentration of the liquid
increases. The tank liquor, after settling, is then either at once
concentrated by evaporation, when the soda crystallises out, leav-
ing the caustic soda in the mother-liquor, or it is submitted to the
action of carbon dioxide, whereby, both the caustic soda and the
sodium sulphide are converted into sodium carbonate, thus —

= Na2CO3+H2O.
Na2S + CO2 + H2O = Na2CO3 + H2S.

The concentration of the tank liquor is accomplished in the
shallow pans above mentioned, by means of the waste heat from
the black-ash furnace ; and the product obtained by evaporating the
liquid is usually calcined at a red heat in an ordinary reverberatory
furnace. This substance is known as soda-ash, and when dissolved
in water, and the solution allowed to crystallise, the so-called soda
crystals are obtained, having the composition Na2CO3,10H2O.

II. The Ammonia- Soda Process. — This process is based upon the
fact, that hydrogen ammonium carbonate (bicarbonate of ammonia)
is decomposed by a strong solution of sodium chloride, according
to the equation —

H(NH4)CO3+NaCl = HNaCO3 + NH4Cl.

In practice the brine is first saturated with ammonia gas, and
the cooled ammoniacal liquid is then charged with carbon dioxide,
under moderate pressure, in carbonating towers.

The hydrogen sodium carbonate (bicarbonate of soda], being
much less soluble, separates out, leaving the more soluble am-

Sodium Carbonate 539

monium chloride in solution, from which the ammonia is recovered
by subsequent treatment with lime.

The hydrogen sodium carbonate is converted into normal sodium
carbonate by calcination, and the carbon dioxide evolved is again
utilised in carbonating a further quantity of ammoniacal brine —

2HNaCO3= Na2CO3 + CO2+ H2O.

These two processes, namely, the Leblanc and the ammonia-soda
process, have been keen competitors for a number of years ; and
a glance at the figures giving the annual output from the two
sources shows how rapidly and steadily the younger process has
gained upon its older rival. Indeed, there can be little doubt that
but for the value of the hydrochloric acid which is simultaneously
produced in the Leblanc process, this method would before now
have ceased to exist as a manufacture. Now, however, both of
these processes are threatened by the advent of a new and formid-
able rival in the electrolytic method.

III. The Electrolytic Process (Hargreaves-Bird). — In this
method a solution of sodium chloride (brine, pumped direct from
the salt-beds) is submitted to electrolysis in a cell of special con-
struction. This consists of an oblong box divided longitudinally
into three compartments, the centre one being comparatively
large, while the two extreme compartments are quite narrow. The
partitions which divide the box in this manner are made of a
"composition" consisting largely of asbestos ; and are of such a
nature that when the middle compartment is filled with brine,
none of the liquid percolates or oozes through into the side cham-
bers. These asbestos diaphragms are backed on their outer sides
by a network of copper wire which is made the cathode in the
system. The anode consists of pieces of gas-carbon which are
suspended in the brine in the centre chamber. Although the
asbestos diaphragms are water-tight, in the sense that they do not
allow the brine to pass from the middle to the outer compartments,
they are nevertheless sufficiently porous to keep the copper wire
cathodes moist, and to allow therefore of the passage of the
current. Chlorine is evolved at the anode, and is conveyed away
directly to lime chambers and converted into bleach ing-powder.
The sodium ions pass freely through the asbestos partitions to the
cathodes, there generating sodium hydroxide ; while a stream of
steam and carbon dioxide which is passed through the narrow

54O Inorganic Chemistry

compartments immediately converts the hydroxide into carbonate,
which is thus washed away from the cathodes as fast as it is
formed. The solution so obtained is sufficiently concentrated to
deposit crystals of sodium carbonate on cooling.

The " soda" obtairied by this process, which is now being carried
out on an extensive scale at Middlewich, Cheshire, is extremely
pure, containing from 97 to 98 per cent, of sodium carbonate, and
only about i per cent, of sodium chloride.

Sodium carbonate crystallises in large, transparent, monosym-
metric crystals, commonly known as "soda" or "washing-soda,"
having the composition Na2CO3,10H2O. On exposure to the air
the crystals give up water, and become effloresced upon the surface,
and finally fall to powder, having the composition Na2CO3,H2O.
When crystallised from hot solutions, it forms rhombic crystals,
containing 7H2O. The solubility of sodium carbonate in water
increases with rise of temperature, reaching a maximum at 32.5°,
when ico parts of water dissolve 59 parts of the salt. Above this
temperature the solubility falls, and at 100° the amount dissolved
is 45.4 parts.

Hydrogen Sodium Carbonate (bicarbonate ofsoda\ HNaCO3,
may be obtained by the action of carbon dioxide upon the normal
carbonate, either in solution or as crystals —

Na2CO3,10H2

The greater part of the bicarbonate of soda of commerce is
obtained in the ammonia-soda process above described.

This salt is less soluble in water than the normal carbonate.
Thus, 100 parts of water at different temperatures dissolve the
following quantities of these compounds —

10°. 20°. 30°. 40°.

Na2CO3. . . 12.6 21.4 38.1 50 parts.

HNaCO3 . . 8.8 9.8 10.8 11.7 „

When a solution of hydrogen sodium carbonate is heated, the salt
gives off a portion of its carbon dioxide, and on cooling the solution
deposits crystals having the composition Na2CO3,2HNaCO3,2H2O,
known as sodium sesquicarbonate. On continued boiling, the salt
is completely converted into the normal carbonate. Sodium
sesquicarbonate occurs as a natural deposit in Egypt, Africa,
South America, and elsewhere, known as trona^ from which the
name natrium is derived.

Sodium Nitrate 541

Sodium Sulphate (Glauber's salf), Na2SO4,10H2O, occurs native
in the anhydrous condition as the mineral thenardite, and as a
double sulphate of sodium and calcium, Na2SO4,CaSO4, in the
mineral Glattberite.

It is manufactured in immense quantities in the first (salt-cake)
process in the alkali manufacture, by the Leblanc method.

It is also obtained in large supplies from the Stassfurt deposits,
by double decomposition between magnesium sulphate (from
kieserite) and sodium chloride.

The solution of the mixed salts, when cooled a few degrees
below o°, deposits sodium sulphate, and the soluble magnesium
chloride remains in solution —

2NaCl + MgSO4= Na2SO4 + MgCl2.

Sodium sulphate is also manufactured , by the action of sulphur
dioxide and oxygen upon sodium chloride. This is known as
Hargreavtfs process. The reaction is expressed by the equation —

This process is, in essence, the production of sodium sulphate
from sodium chloride and the constituents of sulphuric acid, with-
out the intermediate manufacture of the acid. The gases from
pyrites burners, similar to those used by the "vitriol" manufacturer,
together with steam, are passed through a series of cast-iron
cylinders containing sodium chloride, and maintained at a tem-
perature of 500° to 550°. Many days are required for the com-
plete conversion of the chloride into sulphate by this process.

Sodium sulphate crystallises in colourless prisms belonging to
the monosymmetric system, containing ten molecules of water ;
when exposed to the air the crystals effloresce, and when heated
to 33° they melt in their own water of crystallisation (see page

153)-

When sodium sulphate is heated with sulphuric acid, in the pro-

portions required by the following equation, hydrogen sodium
sulphate is formed—

Na2SO4 + H2SO4=2HNaSO4.

Sodium Nitrate, NaNO3, occurs associated with other salts in
Bolivia and Peru, as cubical nitre, or Chili saltpetre. The crude

542 Inorganic Chemistry

salt is purified by solution in water, and crystallisation. It forms
rhombohedral crystals, isomorphous with calcspar.

Sodium nitrate is very soluble in water. 100 parts of water dis-
solve at o°, 68.8 parts ; at 40°, 102 parts ; and at 100°, 180 parts of
the salt. When exposed to the air, the salt absorbs moisture, and
on this account cannot be employed as a substitute for potassium
nitrate in the manufacture of gunpowder, or in pyrotechny. Its
chief uses are for the manufacture of nitric acid ; for the manufacture
of potassium nitrate by double decomposition with potassium
chloride ; and as an ingredient in artificial manures.

Sodium Phosphates. — The most important of these compounds
is the hydrogen disodium orthophosphate, or common phosphate
of soda, HNa2PO4. This salt is prepared on a large scale, by
adding sodium carbonate to phosphoric acid until the solution is
alkaline, and then filtering and evaporating the solution, when
large transparent prisms, belonging to the monosymmetric system,
are deposited, having the composition HNa2PO4,12H2O. Exposed
to the air the crystals effloresce, and when heated become an-
hydrous. The salt melts at 35°.

100 parts of water at 10° dissolve 4.1 parts ; at 50°, 43.3 parts ;
and at 100°, 108.2 parts of the anhydrous salt.

Normal Sodium Orthophosphate, Na3PO4, is obtained from
hydrogen disodium phosphate, by evaporating a solution of the
latter salt with sodium hydroxide, until the liquid crystallises —

HNa2PO4+ NaHO = Na3PO4 + H2O.

This salt contains twelve molecules of water, and forms thin
six-sided prisms. Its aqueous solution is strongly alkaline, and
absorbs atmospheric carbon dioxide, with the formation of hydrogen
sodium carbonate and hydrogen disodium phosphate, thus—

Na3PO4 + CO2 + H2Q=HNa2PO4 + HNaCO3.

Dihydrogen Sodium Orthophosphate, H2NaPO4, is obtained
when phosphoric acid is added to ordinary phosphate of soda, until
the liquid gives no precipitate with barium chloride. On evapo-
rating the solution, the salt crystallises —

HNa2P04 + H3P04 = 2H2NaP04.
The aqueous solution of this salt is acid.

Lithium 543

Hydrogen Sodium Ammonium Phosphate (microcosmic salt},
3Na(NH4)PO4,4H2O, is 6btained by adding- a strong solution of
common sodium phosphate to ammonium chloride

HNa2P04 + NH4Cl = NaCl + HNa(NH4)P04.

The orthophosphates are readily converted into pyro- and meta-
phosphates (see page 476).

LITHIUM.

Symbol, Li. Atomic weight=7.oo.

Occurrence.— Lithium is only found in combination with other
elements. It is a constituent of a few somewhat rare minerals, as
petalite, 30SiO2,4Al2O3,Na2O,2Li2O ; spodumene, 15SiO2,4Al2O3,
3Li2O ; lepidolite, or lithium mica, 9SiO2,3Al2O3,K2O,4LiF.

By means of the spectroscope, lithium compounds have been
detected in sea- water, and in most spring and river waters. In a
few cases spring waters are met with which contain considerable
quantities of lithium salts. Thus, W. A. Miller found as much as
0.372 gramme of lithium chloride in i litre of the water of a spring
near Redruth in Cornwall.

Mode Of Formation. — Lithium is obtained by the electrolytic
decomposition of the fused chloride. For this purpose the dry
salt is heated in a porcelain crucible, when it melts at a low red
heat to a mobile liquid. A rod of gas carbon is made the positive
electrode ; and a stout iron wire, one end of which is flattened out,
is used for the negative pole, upon which the lithium is collected.
On passing an electric current through the molten chloride, the
metal forms as a bright globule upon the negative electrode. The
wire is withdrawn and quickly dipped beneath petroleum, and the
solidified globule of lithium is then cut off with a knife. The
reduced metal, in its passage from the crucible to the petroleum,
is protected from oxidation by the film of fused chloride which
coats it.

Properties. — Lithium is a soft, silver-white metal, which soon
tarnishes on exposure to the air. It is easily cut with a knife,
being softer than lead, but harder than sodium. It may be pressed
into wire, and two pieces of the metal may be made to adhere,
or welded together, at the ordinary temperature. Lithium is the
lightest known solid, its specific gravity being 0.59. Its extreme

544 Inorganic Chemistry

lightness is illustrated by the fact that the metal floats upon
petroleum, a liquid which itself floats upon water. Lithium melts
at 1 80°, and at a higher temperature it takes fire and burns with
a bright white light. Lithium decomposes water at the ordinary
temperature, liberating hydrogen and forming lithium hydroxide,
Li HO ; but when a fragment of the metal is thrown upon cold
water it does not melt, and even with boiling water the action is
not attended by inflammation of the hydrogen.

When strongly heated in nitrogen the two elements unite, with
feeble combustion, forming lithium nitride, NLi3.

Lithium Oxide, Li2O, is formed when the metal burns in the
air. It is also obtained by heating the nitrate. It dissolves in
water, forming lithium hydroxide, LiHO.

Lithium Hydroxide is produced by the prolonged boiling of
lithium carbonate with milk of lime, the carbonate of this metal,
unlike potassium and sodium carbonates, being only very slightly
soluble in water.

Lithium Carbonate, Li2CO3, is obtained as a white precipitate
when a solution of either potassium, sodium, or ammonium car-
bonate is added to a solution of either chloride or nitrate of
lithium. The compound is only slightly soluble in cold water, 100
parts of water at 13° dissolving 0.77 part of the carbonate.

Lithium Phosphate, Li3PO4, is precipitated as a crystalline
powder, by the addition of hydrogen disodium phosphate to a
solution of a lithium salt. In the presence of sodium hydroxide the
precipitation is complete, and the formation of this compound is
employed as a quantitative method for estimating lithium. The
crystals contain 2H2O, which they lose when heated. Lithium
phosphate is soluble in nitric, hydrochloric, and phosphoric acidsi
and from the latter solution, on evaporation, the dihydrogen
phosphate is deposited (H2LiPO4) as deliquescent and very soluble
crystals. The chloride, nitrate, and sulphate of lithium are obtained
by dissolving the carbonate in the respective acids. The salts
are readily soluble in water.

Rubidium and Caesium.*— These two rare elements, which were first dis-
covered by Bunsen in the waters of Diirkheim, in the years 1 860-61, are met
with, associated with sodium and potassium, in certain minerals, such as
lepidolites (lithium mica), porphyrites, and in carnallite. They are also found

* For detailed descriptions of these elements and their compounds, the
student is referred to larger works.

Ammonium Salts 545

in many mineral waters, in the mother-liquors from the evaporation of sea-
water, and in the ashes of plants. Although widely distributed, the quantities
present are extremely minute, one of the richest lepidolites in which these
metals occur containing only 0.24 per cent, of rubidium oxide.

The rare mineral pollux, a silicate of aluminium and caesium, containing
also iron calcium and sodium, is the only known mineral in which either of
these two elements occurs as an essential constituent. The analysis of Pisani
(1864) gives 34.07 per cent, of caesium oxide in this substance.

Rubidium is obtained by heating the carbonate with carbon (the charred
tartrate), as in the older method for the preparation of sodium and potassium.

Caesium cannot be isolated by this reaction, but is obtained by the electro-
lysis of the fused cyanide, Cs(CN) (mixed with barium cyanide in order to
render it more readily fusible). Rubidium melts at 38.5°, caesium at 26.5°.

Rubidium gives a green vapour, and when sublimed in a vacuous tube yields
a thin film of metal, which appears deep blue by transmitted light : when
slowly sublimed in this way the metal forms small needle-shaped crystals.
The compounds of these metals closely resemble those of potassium, from
which they can only be distinguished by the different spectra they give.

AMMONIUM SALTS.

The monovalent group or radical (NH4) is capable of replacing
one atom of hydrogen in acids, thereby giving rise to a series of
salts which are closely analogous to, and are isomorphous with,
those of potassium. The radical (NH4), to which the name
ammonium is given, has never been isolated. United to an atomic
electric charge it constitutes the anion NH4', ammonion, which
closely resembles sodion and potassion. When an amalgam of
sodium and mercury is thrown into a solution of ammonium
chloride, the mercury swells up into a honeycombed or sponge-
like mass, which floats upon the surface of the liquid. This so-
called ammonium amalgam was at one time thought to be a true
amalgam of mercury with the metallic radical ammonium. It is
now generally believed to consist of mercury which is simply
inflated by the evolution of hydrogen and ammonia gas. When
this sponge-like substance is subjected to changes of pressure, it
is found to contract and expand in conformity to Boyle's law : its
formation may be. represented by the equation—

In the course of a few minutes the inflated mass shrinks down,
and ordinary mercury remains at the bottom of the solution,
hydrogen and ammonia having been rapidly evolved.

The ammonium salts are obtained for the most part from the
ammoniacal liquor of the gasworks. The material is treated with

2 M

546 Inorganic Chemistry

lime, and distilled ; and the ammonia so driven off is absorbed in
sulphuric or hydrochloric acid, giving rise to ammonium sulphate
or chloride.

Ammonium Chloride (sal ammoniac), NH4C1.— The product
obtained by absorbing ammonia from gas liquor in hydrochloric
acid is purified by sublimation. The crude material is heated
in large iron pots, covered with iron dome-shaped vessels, into

FIG. 141.

which the substance sublimes. Ammonium chloride crystallises in
arborescent or fern-like crystals (Fig. 141), consisting of groups of
small octahedra belonging to the regular system.

loo parts of water at 10° dissolve 32.8 parts, and at 100°, 77 parts
of the salt. On boiling the aqueous solution, dissociation to a
small extent takes place, and a portion of the ammonia escapes
with the steam ; the. solution at the same time becoming slightly
acid.

Ammonium Sulphate (NH4)2SO4.— The product obtained by

Ammonium Carbonate 547

the absorption of ammonia obtained from gas liquors by sul-
phuric acid is purified by recrystallisation, when it forms colourless
rhombic crystals, isomorphous with potassium sulphate. 100 parts
of water at the ordinary temperature dissolve 50 parts of the salt.
The chief use of ammonium sulphate is for agricultural purposes,
as a manure ; and for this use the crude salt, as first obtained,
which is usually more or less coloured with tarry matters, is em-
ployed. Ammonium sulphate is also used for the preparation of
ammonia alum and other ammonium compounds, as well as in
the ammonia-soda process.

Ammonium Carbonates. — Commercial ammonium carbonate
(sal volatile) is obtained by heating a mixture of ammonium
sulphate and ground chalk to redness in horizontal iron retorts or
cylinders, and conducting the vapours into leaden receivers or
chambers, where the carbonate condenses as a solid crust. It is
afterwards purified by resublimation, when it is obtained as a
white fibrous mass. This substance is a mixture of hydrogen
ammonium carbonate, H(NH4)CO3, and ammonium carbamate,
(NH4)CO2(NH2), and smells strongly ammoniacal. When treated
with alcohol the ammonium carbamate is dissolved, leaving the
carbonate behind.

Normal Ammonium Carbonate, (NH4)2CO3, is obtained from
the commercial compound, by passing ammonia gas into a strong
aqueous solution, or by digesting the compound in strong aqueous
ammonia. The carbamate present is converted into normal car-
bonate by the action of the water, thus^

(NH4)C02(NH2) + H20 = (NH4)C03(NH4)=(NH4)2C03;

and the ammonia converts the bicarbonate into the normal salt,

thus—

H(NH4)CO3 + NH3 = (NH4)2CO3.

Normal ammonium carbonate on exposure to the air gives oft
ammonia, and passes back into hydrogen ammonium carbonate.
When heated to 60° the salt breaks up into carbon dioxide,
ammonia, and water.

Hydrogen Ammonium Carbonate, H(NH4)CO3, may also be
obtained by passing carbon dioxide into a solution of the normal

salt—

(NH4)2CO3 + CO2+H2O = 2H(NH4)CO3.

It forms large lustrous crystals belonging to the orthorhombic

548 Inorganic Chemistry

system, which, when dry, do not smell of ammonia. 100 parts of
water at 15° dissolve 12.5 parts of this salt. At ordinary tempera-
tures this solution on exposure to the air slowly gives off carbon
dioxide, and becomes alkaline ; and when heated above 36° the
liquid begins to effervesce, owing to the rapid evolution of carbon
dioxide. This salt forms with the normal carbonate a double salt
analogous to sodium sesquicarbonate, and having the composition
CNH4)2CO3r2H(NH4)CO3,H2O.

Ammonium Thioeyanate, NH4S(CN), is prepared by adding
aqueous ammonia to an alcoholic solution of carbon disulphide,
and allowing the mixture to stand, when ammonium thiocarbonate
is formed, thus —

On heating this solution, the ammonium thiocarbonate is de-
composed with evolution of sulphuretted hydrogen —

(NH4)2CS3=2H2S + NH4S(CN).

Ammonium thiocyanate (known also as ammonium sulpJw-
cyanate) forms colourless crystals, which are extremely soluble
both in water and alcohol. The solution in water is attended with
considerable absorption of heat : thus, if 20 grammes of the salt
be dissolved in 25 cubic centimetres of water at 18°, the temperature
of the liquid falls to -13°.

CHAPTER V
THE ELEMENTS OF GROUP I. (FAMILY B.)

Copper, Cu . , . „--,/> , . 63.6
Silver, Ag . . . ... . . 107.88

Gold, Au . . . ... . . . 197.2

THE elements of this family present many striking contrasts to
those of the other family belonging to the first group. These
three metals are not acted upon by oxygen, or by water, at
ordinary temperatures ; they are all found native in the un-
combined state, and on this account are amongst the earliest
metals known to man. The alkali metals, on the other hand, are
instantly oxidised on exposure to air, they decompose water at
the ordinary temperature, are never found native, and are amongst
the most recently discovered metals. With the exception of
sodium and potassium, which are used in a few manufacturing
processes, the alkali metals, as such, are of little practical service
to mankind, whilst the metals of this family are amongst the most
useful of all the metals, and are the three universally adopted for
coinage. Many of the compounds of the elements of this family
are similarly constituted to those of the alkali metals : thus, with
oxygen and with sulphur we have Cu2O, Ag2O, Au2O, and Cu2S,
Ag2S, Au2S, corresponding to Li2O and K2S.

With the halogens they all form compounds of the type
RX. Although the three elements, copper, silver, and gold, fall
into the same family upon the basis of the periodic classification
of the elements, they are in many respects widely dissimilar.
Thus, silver is consistently monovalent, while copper is divalent,
forming compounds of the type CuX2, and gold is trivalent, giving
compounds AuX3. The chlorides, AgCl and Cu2Cl2, on the other
hand, are both insoluble in water, are both soluble in ammonia,
and both absorb ammonia.

In many of their physical attributes, these metals show a regular

549

55O Inorganic Chemistry

gradation iu their properties. Thus, as regards malleability and
ductility, silver is intermediate between copper and gold, the
latter possessing these properties in the highest degree. With
respect to their tenacity, silver is again intermediate, copper being
the most, and gold the least tenacious of the three.

COPPER.

Symbol, Cu. Atomic weight =63. 6.

Occurrence.— Copper is found in the elementary condition in
various parts of the world, notably in the neighbourhood of Lake
Superior, where native copper occurs in enormous masses. In
combination, copper is a very abundant element, and is widely
distributed, the most important of these natural compounds being
the following —
Ruby ore . . . Cu2O.
Copper glance . . Cu2S.

Copper pyrites

(Cu2S,Fe2S3or
' I CuFeS2.

Purple copper ( SC^S.FegSg or

ore . . .{ Cu3FeS3.
Malachite . . CuCO3,Cu(HO)2.

Azurite . . 2CuCO3,Cu(HO)2.

Modes of Formation.— The methods by which copper is
obtained from its ores vary with the nature of the ore. From
ores containing no sulphur, such as the carbonates and oxide>
the metal may be obtained by a method known as the reducing
process, which consists in smelting down the ore in a blast-furnace
with coal or coke, when the metal is reduced according to the
equation —

In the case of mixed ores, containing sulphides, the process
(known as the English method) consists of six distinct stages —

(i.) The ores, which contain on an average 30 per cent, of iron
and 13 of copper (the remainder being chiefly sulphur and silica),
are first calcined ; usually in a reverberatory furnace, whereby a
portion of the sulphur is burnt to sulphur dioxide, and the metals
are partially oxidised.

(2.) The second step consists in fusing the calcined ore ; when
the copper oxides, formed during calcination, react upon a portion
of the ferrous sulphide, with the formation of cuprous sulphide
and ferrous oxide, thus —

Cu2O + FeS = Cu2S + FeO.
2CuO + 2FeS = Cu2S + 2FeO + S

Copper $51

The oxide of iron combines with the silica already present (or
which is added in the form of metal-slag obtained from the fourth
step) to form a fusible silicate of iron, or slag, which contains
little or no copper. This is run off, and a fused regulus remains,
consisting of cuprous and ferrous sulphides, known as coarse-metal,
and containing from 30 to 35 per cent, of copper. This molten
regulus, which has a composition very similar to copper pyrites,
is usually allowed to flow into water, whereby it is obtained in a
granulated condition favourable for the next operation.

(3.) The third step consists in calcining the granulated coarse-
metal ; the result, as in the first calcination, being the removal of
a part of the sulphur as sulphur dioxide, and the partial oxidation
of the metals.

(4.) The calcined mass is next fused along with refinery -slag,
which results in the production of a regulus consisting of nearly
pure cuprous sulphide, the greater part of the iron having p?ssed
into the slag (known as metal-slag). This regulus, called fine-
metal, or white-metal, contains from 60 to 75 per cent, of copper.

(5.) The fifth operation consists in roasting the "white-metal''
in a reverberatory furnace. A portion of the cuprous sulphide is
here oxidised into cuprous oxide, which, as the temperature rises,
reacts upon another portion of cuprous sulphide, thus —

At the same time any remaining ferrous sulphide is converted into
oxide, thus —

3Cu2O + FeS = 6Cu + FeO + SO2,

The metallic copper so obtained presents a blistered appearance,
and on this account is known as blister-copper.

(6.) This impure copper is lastly subjected to a refining process.
For this purpose it is melted down upon the hearth of a reverbera-
tory furnace, in an oxidising atmosphere. The impurities present
in the metal, such as iron, lead, and arsenic, are the first to oxidise ;
and the oxides either volatilise or combine with the siliceous matter
of which the furnace bed is composed, forming a slag, which is
removed. The oxidation is continued until the copper itself begins
to oxidise, when the oxide so formed reacts upon any remaining
cuprous sulphide with the reduction of copper and the evolution of
sulphur dioxide, according to the above equation. The metal at
this sta'ge is termed dry 'copper ; and in order to reduce the copper

Inorganic Chemistry

oxide which it still contains, the molten mass is stirred with poles
of wood, and a quantity of anthracite is thrown upon the surface to
complete the reducing process.

Wet Process. — Copper is extracted from the burnt pyrites,
obtained in enormous quantities in the manufacture of sulphuric
acid, which contains about 3 per cent, of copper. Although too
poor in copper to be submitted to the smelting process, it is
found that when calcined with 12 to 15 per cent, of common salt,
the copper is all converted into cupric chloride. On lixiviating the
calcined mass with water, the cupric chloride goes into solution, and
metallic copper can be precipitated from it by means of scrap-iron
or by electrolysis.

Properties. — Copper is a lustrous metal, having a characteristic
reddish-brown colour. The peculiar copper-red colour of the metal
is best seen by causing the light to be several times reflected from
the surface before reaching the eye.

Native copper is occasionally found crystallised in regular octa-
hedra, and small crystals of the same form may be artificially
obtained by the slow deposition of the metal from solutions of its
salts by processes of reduction.

Copper is an extremely tough metal, and admits of being drawn
into fine wire, and hammered out into thin leaf. Its ductility and
malleability are greatly diminished by admixture with even minute
quantities of impurities. When heated nearly to its melting-point,
copper becomes sufficiently brittle to be powdered. The specific
gravity of pure copper, electrolytically deposited, is 8.945, which
by hammering is increased to 8.95.

Copper is only slowly acted upon by exposure to dry air
at ordinary temperatures ; but in the presence of atmospheric
moisture and carbon dioxide it becomes coated with a greenish
basic carbonate. When heated in air or oxygen, it is converted
into black cupric oxide, which flakes off the surface in the form of
scales. When volatilised in the electric arc, copper gives a vapour
having a rich emerald-green colour.

Copper is readily attacked by nitric acid, either dilute or con-
centrated, with the formation of copper nitrate and oxides of
nitrogen (page 246).

Dilute hydrochloric and sulphuric acids are without action upon
copper when air is excluded, but slowly attack it in the presence
of air, or in contact with platinum. Cold concentrated sulphuric
acid does not act upon copper ; but when heated, copper sulphate

Cuprous Oxide $53

and sulphur dioxide are formed, with the simultaneous production
of varying quantities of cuprous and cupric sulphides, which
remain as a black residue (page 416).

Finely divided copper is slowly dissolved by boiling concen-
trated hydrochloric acid, with evolution of hydrogen and formation
of cuprous chloride —

CuCl+H or

In the presence of air, copper is acted upon by a solution of
ammonia, the oxide dissolving in the ammonia forming a deep
blue solution.

Copper is an extremely good electric conductor, being only
second to silver in this respect ; it is therefore extensively em-
ployed for cables, or leads, for purposes of telegraphy and electric
lighting.

Copper possesses the property, in a high degree, of being de-
posited in a coherent form by the electrolysis of solutions of its
salts. On this account it is extensively used in processes ot
electrotyping.

Alloys of Copper. — The most extensive use of copper is in
the formation of certain alloys, many of which are of great technical
value. The following are among the most important : —

English brass . . . Copper 2 parts Zinc part

Dutch brass (Tombac} . ,, 5 „ „

Muntz metal . . . „ 3 „ „

Gun metal . ^ .: . . „ 9 „ Tin

Aluminium bronze . . „ 9 „ Aluminium

Oxides of Copper. — Two oxides of copper are well known,
namely, cuprous oxide (copper sub-oxide\ Cu2O, and cupric oxide
(copper monoxide)^ CuO.

Cuprous Oxide, Cu2O, occurs native as red copper ore. It is
formed when finely divided copper is gently heated in a current
of air, or when a mixture of cuprous chloride and sodium carbonate
is gently heated in a covered crucible.

Cu2Cl2 + Na2CO3 = 2NaCl + CO2 + Cu2O.

Cuprous oxide is also obtained when an alkaline solution of a
copper salt is reduced by grape sugar.

554 Inorganic Chemistry

Cuprous oxide is insoluble in water ; it is converted into cuprous
chloride by strong hydrochloric acid. Nitric acid converts it into
cupric nitrate with the evolution of oxides of nitrogen. When acted
upon by dilute sulphuric acid, it is partly reduced to metallic copper
and partly oxidised into copper sulphate, thus —

When heated with the strong acid it is entirely oxidised, thus —

Cuprous oxide fuses at a red heat, and when melted with glass,
imparts to the latter a rich ruby-red colour.

Cuprie Oxide, CuO, occurs as the rather rare mineral, tenorite.
It is formed when copper is strongly heated in the air or in oxygen,
or by gently igniting either the nitrate, carbonate, or hydroxide.
It is a black powder, which rapidly absorbs moisture from the
air. When heated, it first cakes together and finally fuses,
giving up a part of its oxygen, and leaving a residue consisting
ofCuO,2Cu2O.

When heated in a stream of carbon monoxide, marsh gas, or
hydrogen, it is reduced to the metallic state. Similarly, when
mixed with organic compounds containing carbon and hydrogen,
it oxidises these elements to carbon dioxide and water, itself being
reduced : on this property depends its use in the ultimate analysis
of organic compounds.

Cuprie Hydroxide, Cu(HO)2, is the pale blue precipitate pro-
duced when sodium or potassium hydroxide is added in excess to a
solution of a copper salt. The compound, when washed, may be
dried at 100° without parting with water ; but if the liquid in which
it is precipitated be boiled, the compound blackens, and is con-
verted into a hydrate having the composition Cu(HO)2,2CuO.
Cupric hydrate dissolves in ammonia, forming a deep blue liquid,
which possesses the property of dissolving cellulose (cotton wool,
filter paper, &c.).

Salts of Copper. — Copper forms two elementary ions, monocu
prion Cu* and dicuprion Cu", giving rise to two series of salts,
namely, cuprous and cupric salts. The former, which are colour-
less, readily pass by oxidation into cupric salts, and serve therefore
as powerful reducing agents, and are mostly insoluble in water.
The cupric salts in the hydrated condition are either blue or
green in colour ; the anhydrous cupric salts 'are colourless or

Cupric Chloride 555

yellow. The normal salts are mostly soluble in water. Copper
salts impart to a non-luminous flame a blue or green colour, and
on this account are employed in pyrotechny. The soluble salts
are poisonous.

Cuprous Chloride, Cu2Cl2, may be obtained by dissolving
cuprous oxide in hydrochloric acid. It is more readily prepared
by boiling a solution of cupric chloride in hydrochloric acid,
with copper turnings or foil. The nascent hydrogen, liberated by
the action of the hydrochloric acid upon the copper, reduces the
cupric chloride to cuprous chloride. The liquid is then poured
into water, which causes the precipitation of the cuprous chloride
as a white crystalline powder.

A mixture of zinc dust and copper oxide added to strong hydro-
chloric acid also yields cuprous chloride, the nascent hydrogen in
this case being derived from the action of the acid upon the zinc,
and this causes the reduction of cupric chloride formed by the
action of the acid upon the cupric oxide.

Cuprous chloride melts when heated, and volatilises without
decomposition. It is insoluble in water, but dissolves in hydro-
chloric acid, ammonia, and alkaline chlorides. These solutions, on
exposure to the air, absorb oxygen, turning first brown, and fin-
ally depositing a greenish-blue precipitate of copper oxychloride,
CuCl2,3CuO,4H2O. This compound occurs native as the mineral
atacamite. Solutions of cuprous chloride also absorb carbon
monoxide, forming a crystalline compound, believed to have the
composition, COCu2Cl2,2H2O. They also absorb acetylene (see
page 318).

Cuprous bromide, Cu2Br2 ; iodide, Cu2I2 ; and fluoride, Cu2F2,
are also known.

Cuprie Chloride, CuCl2. — This compound is formed when
copper is dissolved in nitro-hydrochloric acid, or when cupric
oxide, carbonate, or hydroxide are dissolved in hydrochloric acid.
It is also produced when copper is burnt in chlorine.

Cupric chloride is readily soluble in water, forming a deep green
solution, which, on being largely diluted, turns blue. The salt
crystallises in green rhombic prisms, with 2H2O. When heated
it loses its water, and at a dull red heat is converted into cuprous
chloride, with evolution of chlorine (see page 355).

Cupric chloride forms three compounds with ammonia. The
anhydrous salt absorbs ammonia gas, forming a blue compound,
CuCl2,6NH3, When ammonia is passed into aqueous cupric

556 Inorganic Chemistry

chloride, the solution deposits deep blue crystals (tetragonal
pyramids) of the compound, CuCl2,4NH3,H2O. Both these sub-
stances, when moderately heated, yield the green compound
CuCl2,2NH3, which at a higher temperature is decomposed,
thus—

6(CuCl2J2NH3) = 3Cu2Cl2+6NH4Cl + 4NH3+N2.

Cupric bromide, CuBr2, and fluoride, CuF2, are known, but the
iodide is unknown.

Cuprie Nitrate, Cu(NO3)2,3H2O, may be obtained by the
action of nitric acid upon cupric oxide, hydroxide, carbonate, or
the metal itself. It is deposited from the solution in deep blue
deliquescent crystals, soluble in alcohol. When heated to about
65° the crystals lose nitric acid and water, and are converted into
the basic nitrate, Cu(NO3)2,3Cu(HO)2. The normal salt, there-
fore, cannot be obtained anhydrous. Cupric nitrate is a caustic,
powerfully oxidising substance. If the moist salt be rubbed in a
mortar with a quantity of tinfoil, the tin is quickly converted into
oxide, with considerable rise of temperature. When a solution
containing copper nitrate and ammonium nitrate is evaporated, the
mixture suddenly deflagrates when a certain degree of concentra-
tion is reached.

Cuprie Sulphate (blue vitriol), CuSO4,5H2O, is the most
important of all the copper salts. It is formed when either the
metal or the oxide is dissolved in sulphuric acid. On a com-
mercial scale it is obtained from waste copper by first converting
the metal into sulphide by heating it in a furnace, and throwing
sulphur upon the red-hot metal. Air is then admitted, and the
sulphide is thereby oxidised into sulphate, which is dissolved in
water and crystallised.

It is also manufactured from the sulphur ores of copper, by
roasting them under such conditions that the iron is for the most
part converted into oxide, while the copper is oxidised to sulphate.
On lixiviating the roasted mass the copper sulphate, with a certain
amount of ferrous sulphate, is dissolved out. The ores may
also be roasted so as to convert both the metals into oxides ;
the mass is then treated with "chamber acid," which dissolves
copper oxide, leaving the iron oxide for the most part unacted
upon.

Cupric and ferrous sulphates cannot be entirely separated by
Crystallisation, as a solution of these salts deposits a double

Copper Sulphides 557

sulphate of the two metals. If, however, the amount of iron pre-
sent is comparatively small, the first crop of crystals obtained is
moderately pure copper sulphate. The copper is removed from
the mother-liquors by precipitation upon plates of iron, and the
copper so obtained is converted into sulphide, as above described.

Copper sulphate forms large blue asymmetric (tridinic) crystals,
with 5H2O. At 100° it is converted into a bluish- white salt,
CuSO4,H2O, and at 220° to 240° it becomes anhydrous. The
anhydrous salt is white, and extremely hygroscopic, and is used
both for the detection and removal of small quantities of water
in organic liquids.

100 parts of water at 10° dissolve 36.6 parts, and at 100°, 203.3
parts of the crystallised salt.

Several basic sulphates of copper are known : thus, when the
normal salt is submitted to prolonged heating, it is converted into
an amorphous yellow powder, consisting of CuSO4,CuO, which
when thrown into cold water forms an insoluble green compound,
CuSO4,3Cu(HO)2, and on treatment with boiling water yields
CuSO4,2Cu(HO)2. Copper sulphate forms several compounds
with ammonia. Thus, the anhydrous salt readily absorbs ammonia
gas, forming the compound, CuSO4,5NH3. When excess of
ammonia is added to a solution of copper sulphate, the deep blue
solution deposits blue crystals of CuSO4,H2O,4NH3. At 150°
this compound is converted into CuSO4,2NH3, and at 200° it loses
one more molecule of ammonia, leaving CuSO4,NH3.

CupriC Carbonates. — The normal carbonate has not been
obtained. The two most important basic carbonates are (i)
CuCO3,Cu(HO)2, occurring native as malachite, and obtained when
sodium carbonate is added to a solution of copper sulphate (the
green deposit which appears upon copper when exposed to atmos-
pheric moisture and carbon dioxide (verdigris) is the same com-
pound) ; and (2) 2CuCO3,Cu(HO)2, occurring as the mineral azurite.

Sulphides Of Copper. — Two sulphides are known, correspond-
ing to the two oxides.

Cuprous Sulphide, Cu2S, occurs in nature as copper glance,
in the form of -grey metallic-looking rhombic crystals. It is pro-
duced when copper burns in sulphur vapour, or when an excess
of copper filings is heated with sulphur.

Cupric Sulphide, CuS, is met with in nature as the mineral indigo-
copper. It is obtained when either copper or cuprous sulphide is
heated with sulphur to a temperature not beyond 1 14° \ so obtained,

558 Inorganic Chemistry

the compound is blue. As a black precipitate, it is formed when
sulphuretted hydrogen is passed into solutions of cupric salts.

SILVER.

Symbol, Ag. Atomic weight =107. 88.

Occurrence.— Silver is found uncombined, occasionally in
masses weighing several cwts. Such native silver usually contains
copper, gold, and other metals.

Amongst the more important natural compounds of silver are
the following : —

Argentite, or silver glance .. . Ag2S.

Pyrargyrite, or ruby silver ore . 3Ag2S,Sb2S3, or Ag3SbS3.

Proustite, or light red silver ore . 3Ag2S,As2S3 „ Ag3AsS3.

Stephanite . . .. • . . 5Ag2S,Sb2S3 „ Ag5SbS4.

Polybasite . . . >• . . 9(Ag2S,Cu2S),Sb2S3,As2S3.

Stromeyerite Ag2S,Cu2S.

Horn silver . . . . AgCl.

Silver is present also in most ores of lead, notably with galena
(lead sulphide) ; argentiferous lead ores constituting one of the
main supplies of silver.

Modes of Formation. — This element may be obtained from
its salts by the electrolysis of their aqueous solutions. The metal
is so readily reduced from its compounds, that many organic
substances, such as grape sugar, aldehyde, certain tartrates, &c.,
are capable of effecting its deposition. When a strip of zinc is
introduced into silver nitrate solution, the silver is at once de-
posited upon the zinc as a crystalline mass.

Pure silver for analytical purposes may be prepared by pre-
cipitating silver chloride, by the addition of hydrochloric acid to
a solution of the nitrate, and reducing the chloride by boiling with
sodium hydroxide and sugar, or by means of metallic zinc. In
this way the metal is obtained as a fine grey, powder. The
chloride may also be reduced by fusion with sodium carbonate,
when the silver is obtained as a button at the bottom of the
crucible. The methods by which silver is obtained from its ores
are very varied ; they may, however, be classed under three heads,
namely —

Silver

559

1. Processes involving the use of mercury. (Amalgamation
processes.)

2. Processes by means of lead.

3. Wet processes.

(i.) Amalgamation Processes.— These depend upon the fact
that certain compounds of silver are reduced by mercury. The
reduced silver then dissolves in the mercury, forming an amalgam,
from which the silver is obtained, and the mercury recovered by
distillation. The process, as still carried on in Mexico and South
America, is the following. The ore is first crushed and then
ground to a fine powder with water, and the mud so obtained is
mixed with 3 to 5 per cent, of common salt, and spread upon the
floor of a circular paved space, the mixing being effected by the
treading of mules. After the lapse of a day, mercury is added,
together with a quantity of roasted pyrites (known as magistral,
and consisting of a crude mixture of cupric and ferric sulphates
and oxides), and the materials thoroughly incorporated. Fresh
mercury is added from time to time, during the several days
required for the completion of the chemical decompositions that
take place. The exact nature of these changes is not thoroughly
understood, but it is probable that they involve first the formation
of copper chlorides, by double decomposition between the copper
sulphate and sodium chloride, and the subsequent action of these
upon the silver sulphide present in the ore, thus —

2CuCl2 + Ag2S = 2AgCl + Cu2Cl2 + S.
Cu2Cl2 + Ag2S = 2AgCl + Cu2S.

The silver chloride dissolves in the sodium chloride present, and
is reduced by the mercury, with the production of mercurous
chloride (calomel}, which is ultimately lost in the washing —

2AgCl + 2Hg= Hg2Cl2 + 2Ag.

The amalgam is first washed, and freed from adhering particles
of mineral, and is then filtered through canvas bags, whereby the
excess of mercury is removed. The solid residue, containing the
silver, is then submitted to distillation.

In other amalgamation processes the ore is first roasted with
salt, in order to convert the silver into chloride. The roasted
ore is reduced to fine powder with water, and introduced into

560 Inorganic Chemistry

revolving casks along with scrap iron, when the chloride is reduced
according to the equation —

2AgCl + Fe = 2Ag + FeCl2,

and the reduced silver is then extracted by the addition of mercury,
with which it amalgamates.

In the modern amalgamation process the finely crushed ore, with
water, is placed in iron pans provided with revolving machinery,
which serves the purpose of further grinding, and also of mixing.
When the ore is reduced to an almost impalpable powder, mercury
is added, and fhe machinery is kept in operation for a few hours,
when the amalgamation is complete ; sometimes common salt and
copper sulphate are added, either together or singly. Their pre-
sence does not appear to be necessary to the process, except in so
far as they aid in keeping the surface of the mercury clean, or
" quick" ; for in the extremely finely divided condition to which the
ore is reduced in this "pan" amalgamation process, the silver
sulphide is readily acted upon by mercury, with the formation of
mercuric sulphide —

and the silver so reduced dissolves in the excess of mercury, from
which it is finally separated by distillation.

(2.) Processes by Means Of Lead.— When silver ores are
smelted with lead, or with materials which yield metallic lead ; in
other words, when silver ores are smelted with lead ores, an alloy of
silver and lead is obtained, from which the silver can be separated.
When the argentiferous lead is rich in silver, the alloy is submitted
to cupeUation^ which consists in heating the metal in a reverbera-
tory furnace, the hearth of which consists of a movable, oval-shapeds
shallow dish, made of bone ash, known as a cupel, or test. The
alloy is fed into this cupel from a melting-pot, and a blast of air is
projected upon the surface of the molten metal. The lead is thus
converted into litharge, and the melted oxide, by the force of the
blast, is made to overflow into iron pots. As the oxidation of the
lead reaches completion, the thin film of litharge begins to exhibit
iridescent interference colours, which presently disappear, leaving
the brilliant surface of the melted silver. The sudden appearance
of the bright metallic surface is known as \\\e flashing of silver.

In the case of argentiferous lead too poor in silver tofce directly

Silver 561

cupelled, the alloy is submitted to one of two processes of con-
centration, namely, the Pattinson process , or the Parkes's process.

The Pattinson process for desilverising lead depends upon the
fact that alloys of silver and lead have a lower melting-point than
pure lead, and therefore when argentiferous lead is melted and
allowed to cool, the crystals which first form consist of lead which
is nearly or quite pure, and the greater part of the silver is in
the still liquid portion. The operation is carried out in a row of
iron pots. A quantity of the metal is melted in one pot, and as
it cools the crystals which begin to form are removed by means
of a perforated iron ladle and transferred to the next pot on
one side. This operation is continued until a definite proportion
(either two-thirds or seven-eighths, depending upon the propor-
tion of silver) has been removed. The residue is then transferred
to the neighbouring pot on the opposite side, and a second charge
melted up in the first pot. As the neighbouring pots fill up they
are similarly treated, and in this way an alloy, gradually becoming
richer and richer in silver, is passed along in one direction, and
purer and purer lead is sent in the opposite way. The rich alloy
is then cupelled.

The Parkers process depends upon the fact that when zinc is
added to a melted alloy of lead and silver, the zinc deprives the
lead of the silver, and itself forms an alloy with it. The alloy of
zinc and silver rises to the surface and is the first portion to solidify,
and can be removed. The operation is carried out in iron pots.
The argentiferous lead is melted and a quantity of zinc is
thoroughly stirred into the molten mass, the amount of zinc
depending upon the richness of the lead. As the mixture cools,
the first portions to solidify are skimmed off with a ladle and
transferred to another pot. These skimmings, consisting of zinc,
silver, and lead, are first liquated ; that is, carefully heated to such
a temperature that the adhering lead melts and flows away from
the less fusible zinc silver alloy. The solid alloy is then distilled,
and the residue, consisting of silver and lead, is submitted to
cupellation.

(3.) Wet Processes (Ztervogel process). — When argentiferous
pyrites, or an artificially formed regulus containing sulphides of
silver, copper, and iron is roasted, the sulphides are first converted
into sulphates ; and, as the roasting continues, first the iron, then
the copper, and lastly the silver sulphate, is converted into oxide.
By careful regulation the process is continued until the whole

2 N

562

1'norganic Chemistry

of the iron and a part of the copper sulphates are decomposed
On lixiviating the roasted mass with water, the silver sulphate,
together with the remaining copper sulphate, dissolves. From
this solution the silver is precipitated by scrap copper.

The copper is recovered from the solution by precipitation with
iron.

The Percy-Patera Process. — In this method the ore is roasted
with salt and the silver chloride so formed is then extracted by
means of sodium thiosulphate—

Na2S2O3 + AgCl = NaCl + NaAgS2O3.

To the solution so obtained sodium or calcium sulphide is added,
which precipitates silver sulphide —

2NaAgS2O3+ Na2S = Ag2S + 2Na2S2O3.

The silver sulphide is then reduced by being roasted in a rever-

beratory furnace.

Properties. — Silver is a lustrous
white metal which appears yellow
when the light is reflected many
times from its surface before reach-
ing the eye. It is unacted upon by
atmospheric oxygen, but quickly
becomes tarnished by traces of
sulphuretted hydrogen in the air.
Silver has the highest conductivity
for heat and electricity of all the
metals. It is extremely malleable and ductile, being second only
to gold. Thin films of silver appear blue by transmitted light.
Silver melts at about 1000°, and when heated by the oxy hydrogen
flame may be readily made to boil and distil. The pure metal
employed by Stas for the determination of the atomic weight was
obtained by distillation in this way. When volatilised in the
electric arc, the vapour of silver has a brilliant green colour.
Molten silver absorbs as much as twenty-two times its volume of
oxygen, which it gives up again (with the exception of 0.7 volume)
on solidification. As the mass cools, the oxygen evolved often
bursts through the outer crust of solidified metal with consider-
able violence, ejecting portions of the still liquid silver as irregular
excrescences, as seen in Fig. 142. This phenomenon is known

Silver Oxides 563

as the " spitting " of silver. Small quantities of admixed metals
prevent the absorption of oxygen.

Silver is readily soluble in nitric acid, forming argentic nitrate, with
liberation of oxides of nitrogen. Hot concentrated sulphuric acid
converts it into argentic sulphate, with formation of sulphur dioxide
(the reactions in both cases being similar to those with copper).

Silver Alloys. — Silver, alloyed with copper, is largely employed
for coinage and for ornamental purposes. English standard silvef
contains 925 parts of silver per 1000. It is said, therefore, to
have a fineness of 925. In France three standards are used.
That for coinage contains 900 parts per 1000. For medals and
plate the silver has a fineness of 950, while for jewellery it con-
tains only 800 parts per 1000.

Silver-plating. — For purposes of electro-plating, a solution of silver cyanide
in potassium cyanide is used. When a feeble electric current is passed
through this solution (the article to be silvered being the negative electrode,
and a plate of silver the positive), silver in a coherent form is precipitated
upon the negative electrode, thereby coating the object ; and cyanogen is dis-
engaged at the positive pole, where it dissolves the electrode, reforming silver
cyanide.

Silver is reduced from solutions and deposited as a coherent film by a
variety of organic compounds ; and various methods based upon this property
are in use for obtaining mirrors and silvered glass specula for optical pur-
poses. One such method is the following. Two solutions are prepared,
thus—

(i.) Ten grammes of silver nitrate are dissolved in a small quantity of
water, and ammonia added until the precipitate dissolves. The liquid is then
filtered and diluted up to one litre.

(2.) Two grammes of silver nitrate are dissolved in a litre of boiling water,
and 1.66 gramme of Rochelle salt (sodium potassium tartrate, NaKC4H4O6)
are added and the liquid filtered. Equal volumes of these two solutions are
poured into a shallow dish, and the glass to be silvered (after being perfectly
cleaned) is laid in the solution. In about twenty minutes the silver will have
formed a brilliant mirror upon the glass.*

Oxides of Silver.— Three oxides are believed to exist, namely—-

Silver monoxide , . . . Ag2O.
Silver peroxide . • . . Ag2O2
Silver suboxide -...'.• • • Ag4O ?

* By the reduction of silver solutions in the presence of certain organic
compounds, Carey Lea has obtained the metal in the form of a dark bronze
powder, which, when dry, resembles burnished gold. He has also obtained
it exhibiting bluish-green and ruby-red colours. The material differs in
many of its properties from ordinary silver, and is regarded by its discoverer
as an allotropic form of silver (American Journal of Science, 1891).

564 Inorganic Chemistry

Silver Monoxide (argentic oxide\ Ag2O, is obtained by adding
sodium or potassium hydroxide to a solution of silver nitrate. A
brown precipitate consisting of hydrated oxide is obtained which,
when heated, is converted into the anhydrous compound. It is
also formed when silver chloride is boiled with a strong solution of
potassium hydroxide —

Silver oxide is a black amorphous powder which, when heated
to 260°, begins to give off oxygen, and become reduced to metallic
silver. It is a powerful oxidising substance, and when rubbed
with sulphur, red phosphorus, sulphides of antimony or arsenic, or
other readily oxidised substances, it causes them to ignite.

Silver oxide, although only very slightly soluble in water (i part in
about 3000), imparts to the solution a distinct metallic taste and an
alkaline reaction.

It is reduced by hydrogen at 100°, with formation of water and
metallic silver ; and when brought into contact with peroxide of
hydrogen, oxygen is evolved and metallic silver formed (see p. 227).

Silver oxide is soluble in strong ammonia, and, on standing, the
solution deposits black shining crystals of the so -called fulminating
silver. When dry this compound is extremely explosive, and it
often explodes when wet. Fulminating silver is believed to be the
nitride, with the composition NAg3.

Silver Peroxide, Ag2O2. — When a solution of silver nitrate is submitted to
electrolysis, a black powder, consisting of small octahedral crystals, is deposited
upon the positive electrode. The same compound is obtained when a plate of
silver is made the positive electrode in the electrolysis of acidulated water,
and also when silver is acted upon by ozone.

It readily parts with oxygen, and is a still more powerful oxidising agent
than the monoxide. It dissolves in aqueous ammonia with the evolution of
nitrogen —

3Ag202+2NH3=3Ag20 + 3H20 + N2.

Silver SubOXide, Ag4O(?). — The black powder, obtained when silver citrate
is reduced in a current of hydrogen at 100°, and potassium hydroxide is
added to the aqueous solution of the residue, is believed to have the composition
Ag40.

Silver Chloride, AgCl, is obtained as a white, bulky, curdy
precipitate when a soluble chloride is added to silver nitrate. It
melts at 451° to a yellowish liquid, which, on cooling, congeals to a

Silver Fluoride 565

tough horny mass (hence the name horn silver^ as applied to the
native silver chloride). The precipitated chloride is soluble to a
slight extent in strong hydrochloric acid, but readily soluble in
alkaline chlorides, in ammonia, and in sodium thiosulphate. ' Potas-
sium cyanide converts silver chloride into silver cyanide, which
dissolves in the excess of alkaline cyanide, forming the double
cyanide KCN,AgCN. When exposed to the light, silver chloride
darkens in colour, assuming first a violet tint, and finally becoming
dark brown or black (see Photo-salts, p. 566).

Silver chloride absorbs large volumes of ammonia, forming the
compound 2AgCl,3NH3 (see p. 275).

Silver Bromide, AgBr, is prepared similarly to the chloride,
the precipitated compound having a pale yellow colour. It is less
soluble in ammonia than silver chloride ; in dilute ammonia it is
nearly insoluble. Silver bromide is decomposed by chlorine, and
at a temperature of 100° by hydrochloric acid. At ordinary tem-
peratures this reaction is reversed, hydrobromic acid converting
silver chloride into the bromide.

Dry silver bromide does not absorb gaseous ammonia. Silver
bromide is extremely sensitive to the action of light, and is the
chief silver compound used in dry-plate photography.

Silver Iodide, Agl, may be obtained by precipitation from silver
nitrate, with a soluble iodide ; or by dissolving silver in strong
hydriodic acid. As obtained by precipitation it is an amorphous
yellow substance, less soluble in ammonia than either the bromide
or chloride. It dissolves in hot hydriodic acid, which on cooling
deposits colourless crystals of Agl, HI ; the addition of water fo
the solution precipitates the normal iodide, Agl. Silver iodide
absorbs gaseous ammonia, forming a white compound, 2AgI,NH3,
which, on free exposure to the air, evolves ammonia, and is recon-
verted into the yellow iodide.

Silver iodide is the most stable of the three halogen compounds.
When either the chloride or bromide is treated with hydriodic acid
or potassium iodide, iodine replaces the other halogens, forming
silver iodide.

Silver Fluoride, AgF. — This compound is markedly different
in many respects from the other halogen silver salts. It is obtained
by dissolving silver oxide or carbonate in hydrofluoric acid, and
is deposited from the solution in colourless, tetragonal pyramids,
AgF,H2O, or in prisms, AgF,2H26. The salt is extremely deli-
quescent, and very soluble in water. When dried in vacuo? the?

566 Inorganic Chemistry

salt AgF,H2O undergoes partial decomposition, leaving a brownish
residue. When heated, it is partially decomposed, according to
the equation —

2AgF,H2O = 2Ag + 2HF + H2O + O.

The dry salt absorbs gaseous ammonia in large quantities, more
than 800 times its own volume being taken up by the powdered
substance.

Silver Nitrate, AgNO3, is obtained by dissolving silver in
nitric acid. It forms large colourless rhombic tables, which melt
at 218°, and resolidify to a white, fibrous, crystalline mass, known
as lunar caustic. Below a red heat it gives off oxygen, and forms
silver nitrite ; and at higher temperatures it is decomposed into
metallic silver and oxides of nitrogen. 100 parts of water at o°
dissolve 121.9 parts, and at 100°, mo parts of the crystallised
salt ; the solution is neutral. In contact with organic matter,
silver nitrate is blackened on exposure to light. Thus, when the
skin is touched with a solution of this salt, a few seconds' exposure
to light causes a brown or black stain. Owing to this property,
silver nitrate is employed for marking-inks. Silver nitrate absorbs
gaseous ammonia, forming the compound AgNO3,3NH3, the ab-
sorption being accompanied with considerable rise of temperature.
The compound AgNO3,2NH3 is deposited as rhombic prisms when
aqueous silver nitrate is saturated with ammonia.

Silver Sulphate, Ag2SO4, is formed when silver, sjlver carbo-
nate, or silver oxide is dissolved in sulphuric acid. It crystallises
in rhombic prisms, isomorphous with sodium sulphate. With
aluminium sulphate it forms an alum, in which the monovalent
element silver takes the place of potassium in common alum,
Ag2S04,Al2(S04)3,24H20.

Photo-salts. — This name has been applied by Carey Lea to the coloured
compounds formed by the action of light upon silver chloride, bromide, and
iodide. The exact composition of the compounds that are formed when these
silver salts are exposed to light is not definitely known. The change which they
undergo has been attributed (i) to the partial reduction to metallic silver;
(2) to the formation of sub-salts, such as Ag2Cl, Ag2Br, with elimination ot
chlorine or bromine ; (3) to the formation of oxychloride or oxybromide ;
(4) to the production of double compounds of variable composition, of the
sub-salt with the normal salt,

Gold . * 567

GOLD.

Symbol, Au. Atomic weight = 197. 2.

Occurrence. — This element occurs in nature almost exclusively
in the uncombined condition, chiefly in quartz veins and in alluvial
deposits formed by the disintegration of auriferous rocks. It is
present in small quantities in many specimens of iron pyrites,
copper pyrites, and many lead ores, from which it is often
profitably extracted.

Gold is also met with in the form of an amalgam, and in com-
bination with the element tellurium in the minerals petzite^
(AgAu)2Te, and sylvanitc, (AgAu)Te2.

Extraction. — Gold is extracted from auriferous quartz by caus-
ing the finely-crushed substance to flow, by means of a stream
of water, over amalgamated copper plates. The gold particles
adhere to the mercury, with which they amalgamate, and the
amalgam so obtained is carefully removed and distilled.

From alluvial deposits, the native gold is separated by me-
chanical washing.

Gold is extracted from auriferous pyrites by means of chlorine.
The ore is first carefully roasted, and, after being wetted, is exposed
to the action of chlorine gas. The gold is thereby converted into
the soluble auric chloride, AuCl3, which is extracted by lixiviation,
and precipitated by the addition of ferrous sulphate —

2AuCl3 + 6FeSO4 = 2Au + Fe2Cl6 + 2Fe2(SO4)3.

Native gold usually contains silver, from which it may be sepa-
rated by passing chlorine over the molten metal, in crucibles glazed
with borax. The fused chloride of silver rises to the surface, and
is prevented from volatilising by a covering of melted borax.
When the operation is complete, the crucible is allowed to cool,
when the gold solidifies, and the still liquid silver chloride is
poured off.

The Cyanide Process.— Increasing quantities of gold are at
the present time extracted by solution in potassium cyanide. The
method is specially advantageous in cases where the gold is present
in the ore in a very finely divided condition, and it also possesses
the advantage over the " chlorination process," that the preliminary
operation of roasting is obviated. The crushed ore is treated with
a dilute solution of potassium cyanide (containing from 0.25 to

568 Inorganic Chemistry

i per cent, of potassium cyanide), with free exposure to the atmos-
phere, since it has been shown that atmospheric oxygen takes a
necessary part in the action. The gold is dissolved m the form of
a double cyanide, according to the equation —

From this solution the gold is precipitated either by means of
metallic zinc (usually in the form of fine turnings) or by electro-
lytic deposition. The precipitation by means of zinc takes place
according to the equation —

2K AuCy2 + Zn = K2ZnCy4 + 2Au.

The deposit, after being freed as far as possible from zinc, is
melted down with a suitable flux, and yields an alloy containing
70 to 80 per cent, of gold.

When the gold is precipitated electrolytically, the anodes em-
ployed are of lead foil. These are finally melted down and cupelled,
yielding gold of a high degree of purity.

Properties. — Gold is a soft yellow metal, which, when seen by
light many times reflected from its surface, appears red. It is not
acted upon by air or oxygen at any temperature, and does not
decompose steam. No single acid is capable of attacking it
(except selenic acid) ; but it is dissolved by aqua-regia, with for-
mation ot auric chloride. Gold is the most malleable and ductile
of all the metals, and when beaten into very thin leaf, it appears
green by transmitted light.

Gold is most easily reduced from its combinations. Most
metals, when placed in a solution of a gold salt, precipitate the
gold, and the most feeble reducing agents bring about the same
result. On this account a solution of auric chloride is used for
toning photographs. All the compounds of gold, when ignited in
the air, are reduced to metallic gold. Gold is readily deposited
upon other metals by the process of electro-gilding, the most
suitable solution being that of the double cyanide of gold and
potassium, Au(CN)3,KCN.

Gold Alloys. — Alloys of gold with copper and with silver
are used for coinage and for ornamental purposes, pure gold
being too soft for these purposes. Silver gives the alloy a paler
colour than that of pure gold, while copper imparts to it a reddish
tinge. The alloy used for English gold coin consists of gold, 1 1
parts ; copper, i part. The proportion of gold in alloys is usually

Gold Sulphides 569

expressed in parts per 24 (instead of in percentages), these parts
being termed carats. Thus, pure gold is said to be 24-carat gold ;
i8-carat gold contains, therefore, 18 parts of gold and 6 parts of
copper or silver. Most countries have their own legal standards.
In England the legal standard for gold coinage is 22-carats.

Compounds of Gold. — Gold forms two series of compounds, namely, aurous t
in which the metal is monovalent, and auric, in which it is trivalent.

The composition of aurous compounds corresponds to that of the silver
compounds. They are very readily decomposed. Thus, aurous chloride
cannot exist in the presence of water, being decomposed into auric chloride
and metallic gold. For this reason, when aurous oxide, Au2O, is acted upon
by aqueous hydrochloride acid it forms auric, and not aurous chloride, thus —

3Au20 + 6HCl=2AuCl3+3H20 + 4Au.

With iodine geld forms only aurous iodide, Aul ; therefore, when auric oxide
is acted upon by hydriodic acid, aurous iodide and free iodine are formed,
thus—

Au203+6HI=2AuI + 2I2+3H20.

Auric Chloride, AuCls, is obtained by dissolving gold in aqua-regia, and
evaporating the solution to dryness. When the residue is dissolved in water the
concentrated solution deposits reddish crystals of the composition AuCl32H2O.
These lose their water when carefully heated, leaving a brown mass of deliques-
cent crystals. Auric chloride forms double chlorides with the alkaline chlorides,
and with hydrochloric acid, which may be obtained as crystalline compounds.
Thus, the compound AuCl3HCl,3H2O is deposited from a strong solution oi
gold in aqua-regia. This substance is sometimes termed chloro-auric acid,
and the double compounds with metallic chlorides, such as AuCl3,NaCl,2H2O
and ( AuCl3, KC1)2H2O, are known as chloro-aurates.

Auric Oxide, Au2O3, is obtained as a brown powder when the hydrated
oxide, Au.2O3,3H2O (or Au(HO)3), is gently warmed. At 100° it begins to de-
compose, and at higher temperatures is completely converted into oxygen and
metallic gold.

Auric oxide is feebly basic, forming a few unstable salts, in which gold
replaces the hydrogen in acids. It is also a feeble acid-forming oxide, and
forms salts called aurates, such as potassium aurate, KAuO2,3H2O, which may
be regarded as being derived from an acid of the composition HAuO2.

Auric oxide forms a compound with ammonia, known B& fulminating go f^
the exact composition of which is not known. It explodes with violence when
dry if struck or gently warmed.

Gold Sulphides. — Two sulphides of gold have been obtained, aurous
sulphide, Au2S, and auro-auric sulphide, Au2S,Au2S3 (or AuS). The latter is
formed when sulphuretted hydrogen is passed into a cold solution of auric
chloride —

8AuCl3+9H3S + 4H20=2(Au2S, Au2S3) + 24HC1+ HgSO*.

CHAPTER VI
ELEMENTS OF GROUP II. (FAMILY A.)

Beryllium, Be
Magnesium, Mg .
Calcium, Ca

Atomic
Weights.
. 9.1
. 24.32'
. 40.1

Strontium, Sr
Barium, Ba .

Atomic

Weights.

87.6

137.4

WITH the exception of the rare element beryllium, these metals
were first obtained (although not in the pure state) by Davy, who,
soon after his discovery of the metals potassium and sodium,
showed that the so-called earths were not elementary bodies as
had been supposed, but were compounds of different metals with
oxygen.

The element beryllium is of later discovery, for although as
early as 1798 it had been shown by Vanquelin that the particular
" earth " in the mineral beryl was different from any other known
earth, it was not until 1827 that the metal it contained was iso-
lated by Wohler. In a state approaching to purity, beryllium was
first prepared by Humpidge, 1885.

None of the elements of this family occurs in nature in the un-
combined condition ; and, with the exception of magnesium and
calcium, the metals themselves, in their isolated condition, are at
present little more than chemical curiosities. In the case of
beryllium this is due to the comparative rarity of its compounds ;
but with strontium, and barium, whose compounds are extremely
abundant, it is owing partly to the difficulty of isolating the metals
in a pure state, and also to the fact that hitherto they have re-
ceived no useful application. Beryllium and magnesium are
white metals, which retain their lustre in the air. Calcium, stron-
tium, and barium on exposure to air quickly become converted
into oxide.

The metals calcium and strontium, as obtained by earlier experi-
menters, presented a pale yellow colour (it is doubtful whether

579

Metals of the Alkaline Earths 571

the metal barium was actually obtained by these chemists). But
the calcium which has recently been obtained in considerable
masses is found to be a silver-white metal.

All these metals form an oxide of the type RO. Beryllium oxide
is insoluble in water ; magnesium oxide is very slightly soluble
(i part in 55,000 or 100,000 parts of water), but the solution
shows a feeble alkaline reaction. The calcium, strontium, and
barium oxides show increasing solubility, and stronger alkalinity
and causticity. On this account these elements are known as the
metals of the alkaline earths. These three elements also form
peroxides of the type RO2.

All the monoxides are basic, and combine with acids to form
salts of the types RC12, RSO4, R(NO3)2.

The element beryllium (the typical element) stands apart from
the others of this family in many of its chemical relations. Thus
the oxide BeO, unlike the corresponding compounds of the other
elements, does not combine with water to form the hydroxide.
The hydroxide Be(HO)2 is soluble in sodium and potassium
hydroxide. In this respect beryllium exhibits its resemblance to
zinc. The chloride also differs from the other chlorides in being
volatile.

In its permanence in air, its colour, its high melting-point,
the solubility of its sulphate, and the readiness with which its
hydroxide is converted by heat into the oxide, beryllium ex-
hibits a close similarity to magnesium. In the solubility of
its hydroxide in potassium hydroxide, and in its inability to
decompose water, beryllium also shows a marked resemblance
to zinc.

The three elements, calcium, strontium, and barium, ex-
hibit a closer resemblance to each other in most of their
physical and chemical relations, than to either magnesium or
beryllium.

They are readily distinguished by their different spectra.
Barium salts, when heated in a non-luminous flame, impart to it
a green colour. Calcium and strontium, under the same cir-
cumstances, each give a red colour ; but the red imparted by
strontium compounds is more brilliant, and less orange, than that
of calcium salts. When the flames are examined by the spectro-
scope, the most characteristic lines given by barium are two in the
bright green (Baa and Ba/3). These are accompanied by a number
qf less Brilliant lines. The spectrum of strontium consists of four

572 Inorganic Chemistry

specially prominent lines, one in the bright blue (SrS), one in the
orange (Sra), and two in the red (Sr/3 and Sry), with others less
pronounced ; while that of calcium contains one brilliant green
line (Ca/2) and one equally brilliant orange line (Caa), besides a
large number of less prominent lines.

BERYLLIUM (Glucinum).
Symbol Be. Atomic weight =9.1.

Occurrence. — This element occurs principally in the mineral beryl, a double
silicate of the composition 3BeO,Al2O3,6SiO2. The transparent varieties are
used as gems, the transparent green beryl being the precious emerald.

Phenacite is beryllium silicate Be2SiO4, while chrysoberyl has the compo-
sition BeO,Al2O3.

Formation. — The element is obtained by heating sodium in the vapour of
beryllium chloride, all air having been previously replaced by hydrogen. The
product is afterwards melted beneath fused sodium chloride, when it is
obtained as a coherent solid metal. It may also be obtained by the electro-
lysis of the fused mixed chlorides of beryllium and potassium.

Properties. — Beryllium is a white metal resembling magnesium. It has a
specific gravity of 2.1, and is moderately malleable. It does not readily
tarnish in the air at ordinary temperatures, but when strongly heated, be-
comes coated with a protecting film of oxide. The powdered metal, when
heated, takes fire, and burns with a bright light. It has no action upon
water, even at the boiling temperature.

Beryllium is easily dissolved by dilute hydrochloric acid, with evolution of
hydrogen. Cold dilute sulphuric acid is without action, but when heated
slowly dissolves it. Nitric acid slowly attacks it when concentrated and hot.
It readily dissolves in potassium hydroxide, with evolution of hydrogen.

Beryllium Compounds.— The best known are the oxide (berylla], BeO, a
white infusible powder, insoluble in water, soluble in acids ; the chloride,
BeCl2, obtained by heating the oxide with charcoal in a stream of chlorine, a
white crystalline solid, readily fused and volatilised.

Beryllium compounds do not impart any colour to a Bunsen flame. They
are characterised by possessing a sweet taste, hence the name of glucinum
originally given to this element.

MAGNESIUM.

Symbol, Mg, Atomic weight =24. 32.

Occurrence. — Magnesium is not found in the uncombined state.
In combination it is widely distributed, and is extremely abundant.
In the mineral dolomite, associated with calcium as carbonate, it
occurs in mountainous masses.

tylagnesite, MgCO3 ; kieserite, MgSO4,H2O ; carnallite, MgCl2,

Magnesium 573

KC1,6H2O, are amongst the commoner naturally occurring magne-
sium compounds. It is also a constituent of asbestos, meerschaum,
serpentine, talc, and a large number of other silicates. As sulphate
and chloride it is met with in sea-water and many saline springs.

Modes of Formation. — Magnesium was obtained by Bunsen
by the electrolysis of fused magnesium chloride ; and later by
Matthiessen by electrolysing the fused double chloride of mag-
nesium and potassium (carnallite).

On a manufacturing scale it was later produced by the reduction
of magnesium chloride by means of sodium. A mixture of
anhydrous magnesium chloride (or fused mixed chlorides of mag-
nesium and sodium, or potassium), powdered cryolite, and sodium
is thrown into a red-hot crucible, which is quickly closed. A
violent reaction takes place, at the conclusion of which the melted
mixture is stirred with an iron rod to cause the globules of mag-
nesium to run together.

The crude metal is afterwards purified by distillation.

At the present time magnesium is manufactured by a process
which is practically that formerly employed by Matthiessen on a
small scale, but modified in detail to suit modern electrical re-
sources. An iron crucible or melting pot is used, which is made
the cathode, and the double magnesium potassium chloride
(carnallite) is maintained at a temperature about 700° — i.e. a dull
red heat — by means of gaseous fuel. The anode consists of
a stout carbon rod which dips into the molten material, and
is surrounded by a porcelain cylinder which conveys away the
chlorine.

Properties. — Magnesium is a silvery-white metal, which does
not tarnish in dry air, but becomes coated with a film of oxide
when exposed to air and moisture. At a red heat it melts, and at
higher temperatures may be distilled. When heated in the air it
takes fire, and burns with a dazzling white light, which is extremely
rich in the chemically active rays. The flash of light, obtained by
projecting a small quantity of magnesium filings into a spirit flame,
is used for photographic purposes. Magnesium is only moderately
malleable, and is only ductile at high temperatures ; it is readily
pressed into the form of wire at a temperature slightly below its
melting-point. Magnesium only slightly decomposes water even at
the boiling-point ; but when strongly heated in a current of steam, the
metal takes fire (p. 173). Magnesium is rapidly dissolved by dilute
acids, with brisk evolution of hydrogen, but solutions of caustic

574 Inorganic Chemistry

alkalies are unacted upon by it (compare Zinc). When heated with
aqueous solutions of ammonium salts, hydrogen is evolved, and a
double salt of magnesium and ammonium is found in the solution.

Magnesium combines directly with nitrogen, when strongly
heated in that gas, forming magnesium nitride, N2Mg3 (p. 232).

On account of the brilliant light emitted by burning magnesium,
it is employed for signalling purposes, and also in pyrotechny.

Magnesium Oxide (magnesia), MgO, is found native as the
mineral periclase* It is formed when magnesium burns in the air,
or when magnesium carbonate is submitted to prolonged gentle
calcination, when it is obtained as a white bulky powder, known in
commerce as calcined magnesia or magnesia usta.

Magnesia is extensively manufactured from the magnesium
chloride occurring in the Stassfurt deposits, by first converting the
chloride into carbonate and subjecting this to calcination. Mag-
nesia has been obtained in the crystalline form, identical with that
of periclase, by heating the amorphous compound in a stream of
gaseous hydrochloric acid. It may be fused in the oxyhydrogen
flame, and on cooling it solidifies to a vitreous mass which is suffi-
ciently hard to cut glass. On account of its extreme refractoriness,
magnesia is used for a variety of metallurgical purposes, such as
the manufacture of crucibles, cupels, &c.

Magnesium Hydroxide, Mg(HO)2, is found in nature as the
mineral brucite. It is prepared by precipitating a magnesium salt
by sodium or potassium hydroxide. At a dull red heat it loses
water, and is converted into the oxide, and the magnesia so
obtained has the property of rehydrating itself in contact with
water, with evolution of heat.

Magnesium hydroxide slowly absorbs carbon dioxide, forming
the carbonate ; owing to this fact, and to the property it possesses
of rehydration, magnesia that has been prepared by calcination at
a low temperature can be employed as a cement. Thus, if calcined
magnesite be made into a paste with water, the mixture is found to
harden in about twelve hours, and ultimately to acquire a hardness
equal to that of Portland cement.

Magnesium Chloride, MgCl2. — This salt is formed when mag-
nesia, or magnesium carbonate, or the metal itself, is dissolved in
hydrochloric acid. From this solution monosymmetric crystals of
the composition MgCl2,6H2O are deposited. When this salt is
heated it loses water, and at the same time is partially decomposed
into hydrochloric acid and magnesia ; in order, therefore, to pre-

Magnesium Sulphate

pare the pure anhydrous compound, the double magnesium ammo-
nium chloride is first formed, by adding ammonium chloride to a
solution of magnesium chloride. On evaporation, the double salt
separates out, MgCl2,NH4Cl,6H2O. This salt allows itself to be
dehydrated by heating, without any decomposition of the magne-
sium chloride. When the dried salt is more strongly heated,
ammonium chloride volatilises and leaves the anhydrous magnesium
chloride as a fused mass, which congeals to a white crystalline
solid. Magnesium chloride is deliquescent, and dissolves in water
with evolution of heat. With alkaline chlorides it forms double
salts, as the ammonium salt above mentioned. The potassium
salt, MgCl2,KCl,6H2O, occurs in large quantities as the mineral
carnallite; and the calcium salt, 2MgCl2,CaCl2,12H2O, as tachy-
drite^ in the Stassfurt deposits. When a strong solution of
magnesium chloride is made into a thick paste with calcined
magnesia, the mass quickly sets and hardens, like plaster of Paris,
and is found to contain an oxychloride having the composition
MgCl2,5MgO, associated with varying quantities of water. The
white deposit which forms in bottles containing the solution known
as magnesia mixture consists of MgCl2,5MgO,13H2O.

When magnesium oxychloride is heated to redness in a current
of air, the magnesium is converted into oxide, and a mixture of
chlorine and hydrochloric acid is evolved. The reaction may be
represented as taking place as follows —

The W7eldon-Pechiney process for manufacturing chlorine is
based upon this reaction.

Magnesium Sulphate, MgSO4,7H2O (Epsom salts), is met with
in many mineral springs, and in large quantities as the mineral
kieserite, MgSO4,H.,O.

Magnesium sulphate may be obtained by decomposing dolomite,
(CaMg)CO3, with sulphuric acid, the nearly insoluble calcium
sulphate being readily removed from the soluble magnesium salt
Magnesium sulphate is now very largely manufactured from
kieserite, which in contact with water is converted from the slightly
soluble monohydrated salt into MgSO4,7H2O, which is readily
soluble, and is purified by recrystallisation. As usually obtained,
crystallised magnesium sulphate, MgSO4,7H2O, forms colourless
rhombic prisms ; but when deposited from a cold supersaturated

576 Inorganic Chemistry

solution, it sometimes forms prisms belonging to the monosymmetric
(monoclinic) system, having the same degree of hydration. Above
50°, monosymmetric prisms of the composition MgSO4,6H2O are
deposited.

When the ordinary salt, MgSO4,7H2O, is placed over sulphuric
acid, it loses two molecules of water : when heated to 1 50° it loses
six molecules, and at 200° it becomes anhydrous. At the ordinary
temperature, 100 parts of water dissolve 126 parts of crystallised
magnesium sulphate ; the solution has a bitter taste, and acts as a
purgative. With alkaline sulphates, magnesium sulphate forms a
series of double salts, having the general formula MgSO4,R2SO4,
6H2O. They are ismorphous with each other, crystallising in
monosymmetric prisms. The potassium salt occurs in the Stassfurt
deposits as sckonite.

When anhydrous magnesium sulphate is dissolved in hot sul-
phuric acid, two acid sulphates are obtained. One, having the com-
position MgSO4H2SO4, is deposited from the hot solution ; while
from the cold liquid the salt that crystallises has the composition
MgSO4,3H2SO4. They are at once decomposed by water.

Magnesium Carbonate, MgCO3, occurs as the mineral magne-
site, which is sometimes found as rhombohedral crystals, isomor-
phous with crystals of caldte (CaCO3). Magnesium exhibits a
great tendency to form basic and hydrated carbonates ; the normal
carbonate, MgCO3, is therefore not obtained by precipitating a
magnesium salt with an alkaline carbonate ; the white precipitate
formed under these circumstances is a basic carbonate, whose
composition varies with the conditions of precipitation. If, how-
ever, this precipitate be suspended in water, and the liquid saturated
with carbon dioxide, the compound dissolves (more readily under
increased pressure), and when the solution is heated to 300° under
pressure, in such a manner that the evolved carbon dioxide can
escape, the normal anhydrous carbonate is deposited in rhombo-
hedral crystals isomorphous with calcite. If the solution be
evaporated to dryness, the normal carbonate is deposited in
rhombic crystals isomorphous with arragonile (CaCO3). Magne-
sium and calcium carbonates are therefore isodimorphous. «

Basic Carbonates.— The mineral hydromagnesite is a basic
carbonate of the composition 3MgCO3,Mg(HO)2,3H9O. A number
of basic carbonates are formed by the precipitation of a magnesium
salt with sodium carbonate. Thus, under ordinary conditions a
white bulky precipitate is obtained, known in pharmacy as magnesia

Calcium 577

alba levis. Its composition, although liable to vary through the
presence of other basic carbonates, is in the main the same as that
of hydromagnesite.

If the precipitation be made with boiling solutions, and the pre-
cipitate so obtained be dried at 100°, a denser carbonate is ob-
tained, termed magnesia alba ponderosa, 4MgCO3,Mg(HO)2,4H2O.

When an excess of sodium carbonate is employed, and the
mixture is subjected to prolonged boiling, a carbonate is obtained
having the composition 2MgCO3,Mg(HO)2,2H2O.

CALCIUM.

Symbol, Ca. Atomic weight =40.07.

Occurrence. — Calcium is only met with in nature in combina-
tion. It occurs in enormous quantities as the carbonate in a great
variety of different minerals, such as marble, limestone, calcspar,
and also as coral; and with carbonate of magnesium as dolomite,
or inagnesian limestone. In the form of sulphate, calcium
occurs as gypsum and selenite, CaSO4,2H2O, and as anhydrite,
CaSO4. The fluoride CaF2 occurs &<=> flttorspar, and the various
silicious rocks contain compound silicates of calcium and other
metals. The carbonate and sulphate are present in most spring
and river waters. Calcium compounds are also present in all
vegetable and animal organisms. Thus, bones consist largely of
calcium phosphate.

Modes Of Formation. — Although calcium compounds are so
extremely abundant, the metal itself, until quite recently, was
scarcely more than a chemical curiosity. The element was first
isolated in an impure state by Davy (1808).

More recently Moissan obtained the metal in the form of crystals
by heating together sodium and calcium iodide —

The calcium dissolves in the excess of sodium, and on cooling it cry-
stallises out. The sodium is removed by solution in absolute alcohol.
At the present time calcium is obtained commercially by the
electrolysis of the fused chloride, the success of the process de-
pending upon the device adopted for removing the metal, as it is
reduced, from the action of the fused electrolyte. The cathode
employed is a rod of iron which is brought just to the surface of
the melted chloride. As soon as a small quantity of the metal
calcium collects beneath the end of the cathode, the latter is very

c^8 Inorganic Chemistry

slowly raised by a suitable mechanical arrangement, so that the
calcium may solidify upon the end of the iron rod without any
interruption of the electrolysis. As this process of continuously
raising the cathode proceeds, a rugged rod or bar of calcium
weighing several pounds may be gradually built up.

Properties. — Calcium is a silver-white metal having a specific
gravity 1.85, and melting about 760° C. It is moderately soft and
malleable. The metal is readily oxidised by moist air, and
decomposes water at the ordinary temperature. When heated
in air it takes fire and burns. When heated in hydrogen it
forms calcium hydride, CaH2.

Oxides of Calcium. — Two oxides are known, namely, calcium
monoxide, CaO, and calcium dioxide, CaO2.

Calcium Oxide (lime, quicklime), CaO, is obtained by heating
calcium carbonate to redness —

CaCO3 = CO2+CaO.

On a large scale lime is manufactured by burning limestone or
chalk in kilns with coal. If much clay be present with the lime-
stone, care is required to prevent the mass from fusing when it is
said to be dead burnt. Lime is a white amorphous substance,
which is infusible by the oxyhydrogen flame, but which, when so
heated, emits a bright light, known as the oxyhydrogen limelight.
It absorbs moisture and carbon dioxide from the air. On account
of its power of absorbing moisture, lime is frequently employed as
a dehydrating agent. Thus, gases which cannot be dried by means
of sulphuric acid (e.g. ammonia) may be deprived of moisture by
being passed over calcium oxide. It is also used for withdrawing
water from alcohol in the preparation of absolute alcohol. When
a small quantity of water is poured upon lime the mass rapidly
becomes hot, and volumes of steam are given off, the lime at the
same time swelling up and crumbling to a soft, dry powder. This
process is known as the slaking of lime, and the product is termed
slaked lime, in contradistinction to quick lime. The lime enters
into chemical union with water, forming calcium hydroxide, thus —

CaO + H20 = Ca(HO)2.

Calcium Hydroxide, Ca(HO)2, is a white amorphous powder,
sparingly soluble in water, and, unlike the majority of solids, it is
less soluble in hot than in cold water. 100 parts of water at the

Calcium Chloride 579

ordinary temperature dissolve 0.14 part of calcium hydroxide,
while at 100° the same volume of water dissolves about half that
amount. This solution, known as lime-water, has an alkaline
reaction, and absorbs carbon dioxide, . with the precipitation of
calcium carbonate.

Milk Of Lime is the name given to a mixture of lime with less
water than will dissolve it, whereby an emulsion of lime is obtained.
When a thick paste of lime and water is exposed to the atmos-
phere, in a few days it sets, and continues gradually to harden.
On this account lime is used for mortars and cements. Mortar
consists of a mixture of lime and sand with water. The sand
serves the double purpose of preventing shrinkage on drying, and
also of rendering the mass more permeable to atmospheric carbon
dioxide. The setting of mortar is due to the combined action of
evaporation arid absorption of carbon dioxide.

Calcium Dioxide, CaO2, is obtained by adding lime-water to
hydrogen peroxide, or to sodium peroxide acidulated with dilute
nitric acid ; sparingly soluble crystals of CaO2,8H2O separate out,
which at 130° lose their water. When more strongly heated the
monoxide is formed with evolution of oxygen.

Calcium Chloride, CaCl2, occurs in sea and river waters,
and is present in tachydrite of the Stassfurt deposits. It
is obtained in large quantities as a bye-product in many manu-
facturing processes, such as that of potassium chlorate, ammonia
from ammonium chloride, &c. It may be obtained by the action
of hydrochloric acid upon calcium carbonate, and is deposited on
concentration, in large colourless, deliquescent, hexagonal prisms,
CaCl2,6H2O, which melt at 29° in their water of crystallisation.
When heated below 200° the crystals part with four molecules of
water, and above 200° become anhydrous. As thus obtained the
anhydrous salt is a porous mass, which is extremely hygroscopic,
and on this account is used as a desiccating agent, both for gases
and liquids. At a red heat it fuses, and on cooling solidifies
to a crystalline, deliquescent mass. Calcium chloride combines
with ammonia, forming the compound CaCl2,8NH3. Calcium
chloride, therefore, cannot be employed for drying gaseous
ammonia. ,

Crystallised calcium chloride is extremely soluble in water;
100 parts of water at 16° dissolve 400 parts of the salt, the solu-
tion being attended with considerable absorption of heat. When
mixed with powdered ice or snow liquefaction of both the solids

580 Inorganic Chemistry

rapidly takes place, and the consequent absorption of heat lowers
the temperature of the mixture to - 40°.

Bleaehing-Powder (chloride of lime\ Ca(OCl)Cl.— This im-
portant compound is manufactured on a large scale by the action
of chlorine upon slaked lime. The hydrated lime is spread upon
the floor of the bleach ing-powder chambers to a depth of three or
four inches, and raked into ridges or furrows with a special wooden
rake. Chlorine is then led into the chambers, which are provided
with glass windows to enable the operator to examine the colour
of the atmosphere within. At first the absorption of the chlorine is
rapid, but as the reaction proceeds it becomes slower, and the lime
is from time to time raked over to expose a fresh surface. The
lime is left in contact with the gas for twelve to twenty-four hours.
The excess of chlorine is absorbed by projecting into the chamber
a shower of fine lime dust by means of a mechanical fan-distributor.
This, in settling, rapidly absorbs all the chlorine, and the chambers
can then be opened without any unpleasant smell of chlorine being
perceptible.

The reaction which takes place is expressed by the equation —

Ca(H O)2 + C12 = Ca(OCl)Cl + H2O.

It was formerly believed that bleaching-powder was a mechani-
cal mixture of calcium chloride, CaCl2, and calcium hypochlorite,
Ca(OCl)2, but it has been conclusively shown that the substance
does not contain any free calcium chloride. It may, however, be
regarded as a compound consisting of equivalent proportions of
these two salts, and its composition may be expressed by the for-
mula Ca(OCl)2,CaCl2, which corresponds to 2Ca(OCl)Cl.

The relation in which bleaching-powder stands to calcium chlo-
ride on the one hand and calcium hypochlorite on the other will
be seen by the following formulae —

Calcium Chloride. Calcium Hypochlorite,

Cl— Ca— Cl CIO— Ca— OC1 Cl— Ca— OC1.

In practice the absorption of chlorine by the lime is never as
complete as is represented by the above equation, and the com-
mercial value of the product depends upon the amount of available
chlorine it contains, i.e. chlorine which is evolved on treating the
compound with hydrochloric or sulphuric acid. This ranges from
30 to 38 per cent.

Plaster of Paris 581

When treated with water, bleaching-powder is converted into
calcium chloride and hypochlorite, thus —

2Ca(OCl)Cl = CaCl2 + Ca(OCl)2.

Exposure to atmospheric moisture and carbon dioxide decom-
poses it : —
(i)

2HC1O + 2HC1 = 2H2O + 2C12.
When acted upon by acids chlorine is evolved, thus —
Ca(OCl)Cl + 2HCl =CaCl2+ H2O + C12
Ca(OCl)CJ + H2SO4 - CaSO4 + H2O + C12.

When a solution of bleaching-powder is treated with very dilute
acids, hypochlorous acid is first liberated, which in contact with
hydrochloric acid yields chlorine —

(1) Ca(OCl)2.Aq+2HCl.Aq = CaCl2+2HClO.Aq.

(2) HC1O + HC1 = H2O + C12.

In the process of bleaching, the material is first steeped in a
dilute solution of bleaching-powder and then in dilute acid. The
hypochlorous acid first formed is decomposed in the presence of
excess of hydrochloric acid, generating chlorine within the fibres
of the wet cloth.

Calcium Sulphate, CaSO4) occurs as the mineral anhydrite,
and in the hydrated condition as gypsum, CaSO4,2H2O, of which
satinspar (or fibrous gypsum), alabaster, and selenite are different
varieties. It is obtained in the hydrated condition by precipita-
tion from a solution of calcium chloride, on the addition of sul-
phuric acid or a soluble sulphate. When dried at 110° to 120° it
loses a portion of its water, leaving the hydrate, (CaSO4)2,H2O ; at
200° it becomes anhydrous. Calcium sulphate, in the hydrated
condition, is slightly soluble in water, the solubility reaching a
maximum at 35°, when i part of the compound requires 432 parts
of water for its solution ; above this temperature the solubility
again diminishes. Its solubility is increased by the presence of
alkaline chlorides and free hydrochloric acid.

When boiled in strong sulphuric acid calcium sulphate partially
dissolves, and on cooling an acid sulphate crystallises out, having
the composition CaSO4,H2SO4.

Plaster of Paris is calcium sulphate which has been partially-
deprived of its water of hydration by heat, and converted into the

5 82 Inorganic Chemistry

hydrate, (CaSO4)2,H2O. It is manufactured by burning gypsum in
a kiln or oven in such a way that the carbonaceous fuel does
not come in contact with the sulphate, which would result in its
reduction to sulphide ; the temperature is not allowed to exceed
about 130°. If heated more strongly (above 200°) the sulphate
becomes anhydrous, and is said to be dead burnt; in this con-
dition its property of setting when mixed with water is greatly
impaired. When plaster of Paris is made into a paste with water
it rapidly sets to a hard mass ; this setting is due to its rehydra-
tion, whereby gypsum is reformed, thus —

Calcium Carbonate, CaCO3.— This compound is extensively
met with in nature, as limestone, chalk) marble, and innumerable
varieties of calcspar. It is formed when lime is exposed to atmos-
pheric carbon dioxide. It is obtained when an alkaline carbonate
is added to a soluble calcium salt.

Calcium carbonate is dimorphous ; it occurs as arragonite in
crystals belonging to the orthorhombic system, and as calcspar in
crystals belonging to the hexagonal system. Both these crystal-
line varieties can be artificially obtained ; when deposited from
solutions at the ordinary temperature the crystals are identical
with calcite ; but when crystallised from hot solutions, they form
rhombic crystals corresponding to arragonite.

Calcium carbonate is nearly insoluble in water ; 1000 grammes
of water dissolve .0018 gramme of the compound. It is more
soluble in water charged with carbon dioxide, forming the acid
carbonate of lime, CaCO3,H2CO3, or H2Ca(CO3)2.

looo grammes of water saturated with carbon dioxide will dis-
solve, at o°, 0.7 gramme of calcium carbonate. By increasing the
pressure (thereby increasing the amount of dissolved gas) as much
as 3 grammes of calcium carbonate may be dissolved. When this
solution is boiled the acid carbonate is decomposed (page 221).

Calcium Phosphate (tricalcium orthophosphate\ Ca3(PO4)2, is
the most important of the phosphates of calcium. It. is found as
the mineral osteolite, Ca3(PO4)2,2H2O, and also as sombrerite,
estramadurite, and coprolites. Apatite consists of phosphate and
fluoride, 3Ca3(PO4)2,CaF2 5 and the mineral constituents of bones
consist chiefly of calcium phosphate.

It is obtained in a pure state by the addition of ordinary sodium

Calcium Sulphide 583

phosphate to a solution of calcium chloride in the presence of
ammonia. The precipitate is decomposed on boiling into an
insoluble basic salt and a soluble acid salt. Although nearly
insoluble in pure water, calcium phosphate dissolves in water con-
taining salts in solution, such as sodium chloride or nitrate, or
even dissolved carbon dioxide. On this fact depends the readi-
ness with which this substance is absorbed by the roots of plants.
Calcium phosphate is readily soluble in both nitric and hydro-
chloric acids. It is decomposed by sulphuric acid, with the forma-
tion of monocalcium orthophosphate and calcium sulphate, thus —

Ca3(PO4)2+2HS2O4=2CaSO4+H4Ca(PO4)2.

This mixture of calcium sulphate and monocalcium phosphate
is known as superphosphate of lime > and is largely used as an arti-
ficial manure.

With a larger quantity of sulphuric acid the phosphate is con-
verted into tribasic phosphoric acid. (See Phosphorus, page 453.)

Calcium Carbide, CaC2. — This compound is produced when
lime or chalk is heated with carbon in the electric furnace. It is
also obtained as a second product in the manufacture of phosphorus
when calcium phosphate is heated with carbon (see Phosphorus).
Calcium carbide is manufactured on an extensive scale for use in
the preparation of acetylene (page 318).

Calcium Sulphide, CaS, is formed when sulphuretted hydrogen
is passed over heated lime —

Ca(HO)2 + H2S

Or by heating calcium sulphate with carbon —

CaSO4 + 4C = CaS + 4CO.

Calcium sulphide is decomposed on boiling with water, forming
calcium hydroxide and hydrosulphide, thus —

2CaS + 2H2O = Ca(HO)2 + Ca(HS)2.

Calcium sulphide (in common with barium and strontium sul-
phides), as ttsually obtained, possesses the property of emitting a
feeble light (or phosphorescence) in the dark, after being previously
exposed to a bright light. The light emitted gradually diminishes
in intensity but on re-exposing the compound to the light its

584 Inorganic Chemistry

luminosity is again restored. This property has been long known>
and calcium sulphide was formerly termed Canton's phosphorus.
The material formerly known as Bononian (or Bologniati) phos-
phorus is the corresponding barium compound.

These various sulphides are now manufactured for the preparation of
so-called luminous paint. The phosphorescence of these compounds appears
to be due to the presence of small quantities of foreign substances ; thus, not
only is the particular colour of the light emitted changed by the intentional
introduction of minute traces of bismuth, cadmium, manganese, zinc, and
many other metals, but it has been shown, in the case of calcium sulphide,
that the perfectly pure substance does not exhibit phosphorescence.

STRONTIUM.

Formula, Sr. Atomic weight =87. 6.

Occurrence. — The chief natural compounds of this element are
strontianite, SrCO3, and celestine, SrSO4.

Modes of Formation.— The metal was. first obtained in small
quantity by Davy, by the electrolysis of the hydroxide, or chloride,
moistened with water.

It is more advantageously obtained by electrolysing the fused
chloride. Pure strontium has recently been prepared * by
strongly heating strontium hydride, SrH2, in vacuo.

Properties. — Strontium is a silver-white metal which melts
about 800°. It is readily oxidised by air, and decomposes water
at ordinary temperatures. When heated in hydrogen it forms
strontium hydride, SrH2. At -60° it combines with dry ammonia,
forming red-brown crystals of a compound known as " strontium-
ammonium," believed to have the composition Sr,6NH3.

Oxides of Strontium.— Two oxides, corresponding to those of
calcium, are known, namely, strontium monoxide, SrO, and dioxide.
Sr02.

Strontium Monoxide (strontia\ SrO, is obtained by heating the
nitrate or carbonate. It is prepared on a large scale by decompos-
ing strontium carbonate by superheated steam ; carbon dioxide is
evolved, and strontium hydroxide remains, which on ignition forms
the monoxide. Strontia strongly resembles lime. When treated
with water it slakes with evolution of heat, forming strontium
hydroxide, Sr(HO)2. The hydroxide is more soluble in water than
the lime compound, and the solution on cooling deposits tetragonal
crystals, Sr(HO)2,8H2O. The solution is strongly alkaline.

* Guntz and Roederer, Compt. Rend., 1906.

Strontium Nitrate 585

Strontium hydroxide combines with sugar, forming a saccharate
of strontia, which is readily decomposed by carbon dioxide. On
this account it is prepared on a large scale for use in the manu-
facture of beet-sugar. One process by which it is obtained on a
commercial scale consists in first forming strontium sulphide, by
reducing the natural sulphate with carbon, and treating the solution
of the sulphide with sodium hydroxide, thus —

SrS + NaHO + H2O = Sr(HO)2+NaHS.

Strontium Dioxide, SrO2. — When hydrogen peroxide is added
to a solution of strontium hydroxide, a hydrate of the peroxide
separates out in the form of pearly crystals, SrO2,8H2O. On gently
heating this compound, it is converted into the anhydrous peroxide.
On heating to redness it evolves oxygen, and is converted into the
monoxide.

Strontium Chloride, SrCl2, is obtained from strontianite by the
action of hydrochloric acid. The salt deposits from the solution in
deliquescent hexagonal prisms, SrCl2,6H2O, isomorphous with the
corresponding calcium compound.

Strontium Sulphate, SrSO4. — The native compound celestine
occurs in amorphous fibrous masses, and also in rhombic crystals.
The name of the mineral is derived from the fact that it usually
has a light blue colour. It is produced by precipitation from a
strontium salt by sulphuric acid. It is only slightly soluble in cold
water, and still less in hot. When boiled with solutions of alkaline
carbonates, strontium sulphate is completely converted into stron-
tium carbonate —

SrSO4+ Na2CO3= SrCO3 + Na2SO4.

In this respect strontium sulphate differs from barium sulphate,
which under these conditions remains unchanged. On treatment
with strong sulphuric acid, strontium sulphate forms SrSO4,H2SO4,
which, like the corresponding calcium compound, is converted by
water into sulphuric acid and the normal sulphate.

Strontium Nitrate, Sr(NO3)2, is obtained by dissolving the
natural carbonate in dilute nitric acid. On concentration, the
anhydrous salt separates out in octahedrons. From dilute solu-
tion, on cooling, it forms monosymmetric prisms, Sr(NO3)2.
4H2O, which effloresce on exposure to the air. When heated with
carbon, or other readily combustible substances, the mixture in-

586 Inorganic Chemistry

flames and burns with the red colour characteristic of strontium
compounds ; strontium nitrate is therefore largely used in pyro-
techny for the production of red fire. This property is most
readily illustrated by mixing dry powdered strontium nitrate with
ammonium picrate, and igniting the mixture, which burns with a
brilliant red light.

BARIUM.

Symbol, Ba. Atomic weight= 137.4.

Occurrence. — The most abundant natural compounds of barium
are heavy spar, BaSO4, and witherite, BaCO3. It occurs also,
associated with calcium, in the mineral barytocalcite, BaCO3,CaCO3.

Modes of Formation. — The element barium is more difficult to isolate than
either strontium or calcium, and it is doubtful whether pure barium has ever
been obtained. Davy electrolysed various barium salts, made into a thick
paste with water, using mercury as the negative electrode : in this way an
amalgam of barium was formed, from which, on distilling away the mercury,
a dark porous mass was obtained. Amalgams of barium and mercury have
been prepared in other ways, but it has been shown that the product obtained
after distilling the mercury from these is not pure barium, but is a solid alloy
or compound of barium with mercury.

By the electrolysis of the fused chloride, Matthiessen obtained small globules
of metal, which on exposure to the air rapidly oxidised. More recent experi-
menters fail to obtain the metal by this process (Limb., Compt. Rend.t 112).

Oxides of Barium. — Two oxides are known, namely, barium
monoxide, BaO, and dioxide, BaO2.

Barium Monoxide (baryta\ BaO, is usually prepared by heat-
ing the nitrate. The mass fuses and evolves oxygen and oxides of
nitrogen, leaving a greyish white friable residue of the oxide. It
may also be obtained by heating the carbonate ; but as the tem-
perature necessary to expel the carbon dioxide is very high, it is
usual to mix the carbonate with lampblack, tar, or other sub-
stances which on heating will yield carbon, when the conversion
takes place more readily, carbon monoxide being evolved, thus —

BaCO3 + C = BaO + SCO.

Small quantities may readily be obtained by heating barium
iodate in a porcelain crucible, when the iodate is decomposed as
follows —

Ba(IO3)2

Barium Dioxide 587

Barium oxide is a strongly caustic and alkaline compound ; in
contact with water it slakes with evolution of so much heat that
the mass may become visibly red hot if too much water be not
added.

When heated to a dull red heat in oxygen, or air, it takes up an
additional atom of oxygen and forms the dioxide (see p. 184).

Barium Hydroxide, Ba(HO)2, is obtained when the monoxide
is slaked with water. It is manufactured by first heating the
powdered native sulphate with coal, when a crude barium sulphide
is formed. This is then heated in a stream of moist carbon
dioxide, whereby it is converted into the carbonate, and super-
heated steam is then passed over the heated carbonate —

= BaCO3+H2S.
= Ba(HO)2 + CO2.

Barium hydroxide is soluble in water : the solution, known as
baryta-water, absorbs carbon dioxide, with the precipitation of
barium carbonate.

The aqueous solution deposits crystals of hydrated barium
hydroxide, Ba(HO)2,8H2O, in the form of colourless tetragonal
prisms, which on exposure to the air lose seven molecules of water.

Barium hydroxide, when heated in a current of air, yields barium
dioxide.

Barium hydroxide was formerly employed in sugar-refining, but
owing to its poisonous nature it has been superseded by strontium
hydroxide (^.^.).

Barium Dioxide (barium peroxide), BaO2. — This oxide is
obtained by heating the monoxide to a low red heat in a stream of
oxygen, or of air which has been deprived of atmospheric carbon
dioxide.

The pure compound may be obtained by adding an excess of
baryta-water to hydrogen peroxide, when hydrated barium per-
oxide separates out in crystalline scales —

Ba(HO)2+H2O2 + 6H2O = BaO2,8H2O.

On drying in vacuo at 130° this compound loses water and is
converted into the anhydrous peroxide.

The commercial peroxide may be purified by treatment with
dilute hydrochloric acid, whereby barium chloride and hydrogen
peroxide are formed. After the removal of insoluble impurities by

588 Inorganic Chemistry

filtration, baryta-water is cautiously added, which causes the pre-
cipitation of ferric oxide and silica. The liquid is then filtered,
and to the clear liquid, consisting of a solution of barium chloride
and hydrogen peroxide, an excess of strong baryta- water is added,
when the hydrated barium peroxide is precipitated, as already
explained.

Barium peroxide is a grey powder, which on being heated to a
bright red heat gives up oxygen and forms the monoxide (p. 184).

Dilute acids decompose barium peroxide, with formation of
hydrogen peroxide and a barium salt. Concentrated sulphuric
acid forms barium sulphate and ozonised oxygen. When gently
warmed in a stream of sulphur dioxide, the mass becomes incan-
descent and forms barium sulphate —

BaO2 + SO2 = BaSO4.

Barium Chloride, BaCl2, may be obtained by dissolving the
natural carbonate in hydrochloric acid. It may be obtained from
the natural sulphate, either by first converting it into the sulphide,
and decomposing that with hydrochloric acid, or by roasting the
mineral with powdered coal, limestone, and calcium chloride, when
the following reactions take place —

BaS + CaCl2 = BaCl2 + CaS.

The barium chloride is dissolved in water, and an insoluble oxy-
sulphide of calcium remains.

Barium chloride forms colourless rhombic tables, BaCl2,2H2O,
which at 15.6° are soluble to the extent of 43.5 parts in 100 parts
of water. The salt is nearly insoluble in hydrochloric acid, and
may therefore be precipitated from an aqueous solution by the
addition of this acid.

Barium chloride, in common with all the soluble salts of this
element, is highly poisonous.

Barium Sulphate, BaSO4, is the most abundant naturally
occurring barium compound. It is frequently met with as large
rhombic crystals. The specific gravity of the mineral is 4.3 to
4.7 ; and on account of its high specific gravity it received the
name of barytes, or heavy spar.

It is formed as a heavy white precipitate when sulphuric acid,
or a soluble sulphate, is added to a solution of a barium salt. It is
insoluble in water and only very slightly soluble in dilute acids,

Barium Sulphide 589

It is soluble in hot concentrated sulphuric acid, especially when
freshly precipitated ; and the solution deposits, on cooling, an acid
sulphate, BaSO4,H2SO4. On exposure to moisture the solution
deposits crystals of BaSO4,H2SO4,2H2O. Both of these com-
pounds, in contact with water, yield insoluble normal barium
sulphate and sulphuric acid.

Precipitated barium sulphate is largely used as a pigment,
known as permanent white.

Barium Nitrate, Ba(NO3)2, is obtained by dissolving the native
carbonate, or the sulphide, in dilute nitric acid. It is also formed
by double decomposition, when hot saturated solutions of sodium
nitrate and barium chloride are mixed. The salt crystallises in
large colourless octahedrons. 160 parts of water at the ordinary
temperature dissolve 9 parts, and at 100°, 32.2 parts of barium
nitrate. When strongly heated it is converted into barium oxide,
with the evolution of nitrogen peroxide, oxygen, and nitrogen.

Barium nitrate is used in pyrotechny, in the preparation of
mixtures for green fire.

Barium Sulphide, BaS, is obtained by methods analogous to
those for preparing calcium sulphide (page 583), which it closely
resembles in its properties.

CHAPTER VII
ELEMENTS OF GROUP II. (FAMILY B.)

Zinc, Zn . . . . ._ . . . 65.4
Cadmium, Cd . . . • .' . . 112.4
Mercury, Hg ...... 20x3

THE three elements composing this family do not exhibit such
a close resemblance to each other as exists between barium,
strontium, and calcium ; for although zinc and cadmium are very
closely related, mercury in many respects differs widely from these,
and from all the other elements in the same group.

Cadmium and zinc are almost invariably found associated
together in nature, they are both fairly permanent in the air,
and both readily take fire and burn when strongly heated,
forming the oxides. Both are acted upon by dilute hydrochloric
and sulphuric acids, with evolution of hydrogen, and most of their
salts are isomorphous.

Mercury is peculiar in being liquid at ordinary temperatures.
Zinc and cadmium melt at 430° and 320° respectively, while
mercury melts at — 38.8°. It is quite unacted upon by oxygen at
ordinary temperatures, and combines with extreme slowness when
heated. Its oxide, also, is readily decomposed by heat into its
elements.

Dilute hydrochloric and sulphuric acids are entirely without
action upon it, and it forms no hydroxide.

Mercury also differs from zinc and cadmium in forming two
elementary ions, giving rise to mercurous and mercuric salts.
Both zinc and cadmium have only one ion and form only one
series of salts.

The hydroxide of zinc, Zn(HO)2, differs from the corresponding
cadmium compound, in being soluble in alkaline hydroxides.

These three elements resemble each other, and differ from
those of family A of this group, in that they can be volatilised,
mercury at a temperature about 357°, cadmium and zinc at
temperatures approaching 1000°.

These three elements are also alike, in that their vapours con-
sist of mono-atomic molecules.

390

Zinc 591

ZINC.

Symbol, Zn. Atomic weight =65.37.

Occurrence. — Zinc is stated to have been found in Australia in
the uncombined condition ; with this exception, it is always met
with in combination, chiefly as carbonate in calamine or zinc-spar,
ZnCO3, and as sulphide in zinc-blende, or black-jack, ZnS. Other
ores are red zinc ore, ZnO ; 2&&franklinite, (ZnFe)O,Fe2O3.

Gahnite, or zinc-spinnelle, has the composition ZnO,Al2O3.

Modes Of Formation. — The ores chiefly employed for the pre-
paration of zinc are the carbonate and sulphide, although in New
Jersey the red oxide and franklinite are used. The process con-
sists of two operations, namely, first, the conversion of the ore into
oxide of zinc, by calcination ; and, second, the reduction of the oxide
by means of coal at a high temperature. The calcination of the
natural carbonate is readily accomplished, this compound merely
giving up its carbon dioxide at a high temperature —

ZnCO3 = ZnO + CO2.

In the case of zinc-blende, the operation consists in the oxida-
tion of both the sulphur and the zinc by atmospheric oxygen, thus —

Considerable care has to be exercised in order to prevent the
formation of zinc sulphate, which, in the subsequent operation,
would be reconverted into sulphide, and so lost. The finely
crushed calcined ore is mixed with coke or coal and heated to
bright redness in earthenware retorts, when the oxide is reduced,
with the formation of carbon monoxide, and the metal distils and
is collected in iron receivers. Zinc ores frequently contain small
quantities of cadmium, and as this metal is more readily volatilised
than zinc, it passes over in the first portions of the distilled
product.

The two processes now almost exclusively in use for the reduc-
tion of zinc, known as the Silesian and the Belgian process,*
differ only in metallurgical details, &c.

* The old method, known as the English process, or distillation per
descensum, is entirely obsolete. For details of this and all other metallurgical
processes, the student is referred to treatises on metallurgy, such as Percy.

592 Inorganic Chemistry

Commercial zinc usually contains carbon, iron, and lead, and
occasionally arsenic and cadmium. It may be obtained in a higher
degree of purity by careful distillation, but pure zinc is best ob-
tained by first preparing the pure carbonate by precipitation, and
then calcining and finally reducing with charcoal obtained from
sugar.

Properties.— Zinc is a bluish-white, highly crystalline, and
brittle metal. At a temperature of 300° it can be readily powdered
in a mortar, while between 100° and 150° it admits of being drawn
into wire or rolled into thin sheet. The presence of a small
quantity of lead greatly enhances this property, but is detrimental
when the zinc is required for making brass. Zinc which has been
either rolled or drawn no longer becomes brittle when cold, but
retains its malleability.

Zinc melts at a temperature about 430,° and when heated in air
much beyond this point the metal takes fire and burns with a bluish-
white flame, the brilliancy of which becomes dazzling if a stream of
oxygen be projected upon the burning mass. The product of its
combustion is zinc oxide, ZnO, which forms a soft, white, flocculent
substance resembling wool, and formerly known as philosophers
wool.

The boiling-point of zinc is about 930°.

Zinc is permanent in dry air at ordinary temperatures, but when
exposed to moist air it tarnishes superficially ; it is also unattacked
by water at the boiling temperature. It is soluble in a hot solution
of sodium or potassium hydroxide, with evolution of hydrogen

(P- 175)-

Pure zinc is scarcely acted upon by pure sulphuric or hydrochloric
acid, either dilute or strong. The presence of small quantities of
impurities, however, determines the solution of the metal with the
rapid evolution of hydrogen, hence ordinary commercial zinc is
readily attacked by these acids, and also decomposes water at the
boiling-point, with the evolution of hydrogen.*

* The difference between the behaviour of acids towards pure and com-
mercial zinc was formerly explained on the ground that the impurities present
formed with the zinc a voltaic couple, whereby local electric currents were set;
up, while in the case of pure zinc no such action took place. The recent
observations of Pullinger (Chem. Soc., 57) and Weeren (Berichte, 24) show that
this is not a complete explanation. Weeren concludes that the insolubility of
pure zinc in dilute acids is due to the formation of a film of condensed hydrogen
upon the surface of the metal, which stops all further action. The addition of
oxidising agents, such as hydrogen oeroxide, or dilute sulphuric acid which has

Zinc Oxide 593

Zinc is extensively used in the process of galvanising iron, which
consists in coating iron with a film of zinc, not by electrical deposi-
tion, as would be implied by the name, but by dipping the iron
into a bath of molten zinc. The layer of zinc preserves the iron
from rusting. Galvanised iron is better able to withstand the
action of air and moisture than tinned iron, hence it is extensively
used for wire netting, corrugated roofing, water tanks, and other
purposes where the metal is exposed to the oxidising influence of
air and water.

Alloys of Zine. — Zinc forms a number of useful alloys, the most
important of which are the various forms of brass (see Copper).
With certain metals, such as tin, copper, and antimony, zinc will
mix in all proportions ; while with others, such as lead and bismuth,
it is only possible to obtain solid alloys of definite composition.
When, therefore, lead and zinc are melted together, although in
the molten condition the mixture is homogeneous, on cooling the
metals separate into two layers, the lighter zinc rising to the surface.
The separation of the metals, however, is not perfect, for the zinc
will have dissolved a certain quantity of the lead (1.2 per cent.),
and the lower layer of lead is found to have dissolved a small
proportion of zinc (1.6 per cent.), just as water and ether, when
shaken together, separate into two layers, the uppermost being an
ethereal solution of water, and the lower an aqueous solution of
ether.

This property is made use of in the extraction of silver from lead
(see p. 561).

The so-called German silver t or nickel silver^ is a nearly white
alloy of copper, nickel, and zinc.

Bronze coinage consists of 95 parts of copper, 4 of tin, and I of
zinc, the small proportion of zinc giving to the alloy an increased
hardness and durability.

Zinc Oxide, ZnO, the only oxide of zinc, occurs native as red
zinc ore, the colour being due to the presence of manganese. It is

been electrolysed, and therefore contains presulphuric acid, tends to destroy
this film by oxidising the hydrogen, and therefore promotes the solution of the
zinc. He also finds, that by mechanically removing this layer Of hydrogen,
either by constantly brushing the metallic surface or placing the materials
under reduced pressure, the solution of the zinc by the acid is promoted. It is
also found that the character of the surface of the metal, whether smooth or
rough, affects the result ; zinc that is unacted upon when its surface is perfectly
smooth is more readily attacked by the dilute acid when its surface is rough.

594 Inorganic Chemistry

formed as a soft white substance when zinc is burnt in the air. It
is manufactured under the name of zinc white by the combustion
of zinc, the fumes being led into condensing-chambers, where the
oxide collects.

Zinc oxide is a pure white substance, which when heated becomes
yellow, but again becomes white on cooling. When strongly heated
in oxygen, it may be obtained in the form of hexagonal crystals ;
such crystals are occasionally found in the cooler parts of zinc
furnaces. The oxide does not fuse in the oxyhydrogen flame, but,
like lime, under these circumstances it becomes intensely incan-
descent ; for some time after being so heated it appears phos-
phorescent in the dark. It is insoluble in water, and does not
combine directly with water to form the hydroxide. It dissolves
in acids, giving rise to the different zinc salts. Zinc oxide is largely
used in the place of " white lead " as a pigment ; although it does
not equal white lead in covering power, or body^ it possesses the
advantage of not being blackened by exposure to atmospheric
sulphuretted hydrogen.

Zine Hydroxide, Zn(HO)2, is formed as a white flocculent pre-
cipitate, when either sodium or potassium hydroxide, or a solution
of ammonia, is added to a solution of zinc sulphate. The compound
is soluble in an excess of either alkali, and is deposited from a
strong solution in regular octahedra of the hydrated hydroxide,
Zn(HO)2,H2O. Both of these compounds on heating readily lose
water, and are converted into the oxide.

Zine Chloride, ZnCl2, is formed by the direct combination of zinc
with chlorine, or by the action of hydrochloric acid upon the metal.
It is also obtained in the anhydrous state by distilling a mixture of
mercuric chloride and zinc, or a mixture of anhydrous zinc sulphate
and calcium chloride.

It is usually prepared on a large scale by dissolving zinc in
hydrochloric acid, and after precipitating any manganese and iron,
the liquid is boiled down in enamelled iron vessels, until on cooling
it solidifies ; it is usually cast into sticks.

Zinc chloride is a soft, white, easily fusible solid, which volatilises
and distils without decomposition. It is extremely deliquescent,
and readily soluble in water and in alcohol, its solution being
powerfully caustic. From a strong aqueous solution deliquescent
crystals are deposited, having the composition ZnCl2,H2O.

When the aqueous solution is evaporated, partial decomposition
takes place, hydrochloric acid being evolved and basic compounds

Zinc Sulphide 595

being precipitated, consisting of combinations of the chloride and
oxide. Hence, during the concentration of the liquid in the pre-
paration of zinc chloride, hydrochloric acid is added to redissolve
this compound.

A paste made by moistening zinc oxide with zinc chloride rapidly
sets to a hard mass ; this mixture, under the name of oxychloride
of zinc, is employed in dentistry as a filling or stopping for teeth.

Zinc chloride unites with alkaline chlorides, forming a series of
crystalline double salts having the general formula ZnCl2,2RCl.

Zine Sulphate, ZnSO4, is formed when zinc is dissolved in
sulphuric acid. It is obtained on a large scale by roasting the
natural sulphide, whereby it is partially converted into the sulphate,
which is then extracted with water.

The salt crystallises from its aqueous solution at ordinary tem-
peratures in colourless rhombic prisms, ZnSO4,7H2O, isomorphous
with MgSO4,7H2O. It is extremely soluble in water : 100 parts of
water at the ordinary temperature dissolve i£o parts, and at 100°,
653.6 parts of the crystalline salts. When exposed to the air, the
crystals slowly effloresce, and if placed in vacuo over sulphuric
acid, or if heated to 100°, they lose six molecules of water, leaving
the monohydrated salt ZnSO4,H2O. At a temperature about 300°
this is converted into the anhydrous compound, and at a white
heat it gives off sulphur dioxide and oxygen, leaving the oxide.

The hydrated salt, ZnSO4,6H2O, is obtained in the form of mono-
symmetric crystals, when the salt is deposited at temperatures
above 40°. This compound is isomorphous with MgSO4,6H2O.

Zinc sulphate combines with alkaline sulphates, forming a series
of double salts, having the general formula ZnSO4,R2SO4,6H2O,
which are also isomorphous with the corresponding magnesium
compounds (page 576).

Zinc sulphate, in common with all the soluble salts of zinc, has
an astringent taste, and is poisonous.

Zinc Sulphide, ZnS. — The natural compound, zinc-blende, is
usually dark-brown or black, and exhibits crystalline forms belong-
ing to the regular system. The mineral wurtzite is a less common
variety of zinc sulphide, crystallising in hexagonal prisms. Zinc
sulphide is obtained as a white amorphous precipitate when an
alkaline sulphide is added to a solution of a zinc salt, or when
sulphuretted hydrogen is passed through an alkaline solution of a
zinc salt.

Precipitated zinc sulphide is insoluble in acetic acid, but readily

596 Inorganic Chemistry

dissolves in dilute mineral acids, with evolution of sulphuretted
hydrogen ; hence the compound is not formed when sulphuretted
hydrogen is passed through a solution of a zinc salt containing a
free mineral acid.

Zinc Carbonate, ZnCO3, is obtained as a white powder when
hydrogen sodium carbonate is added to a solution of zinc sulphate.

If normal sodium carbonate be employed, the precipitated zinc
compound consists of a basic carbonate, whose composition varies
with the conditions of temperature and concentration of the liquids.

A basic carbonate, having the composition ZnCO3,2Zn(HO)2,H2O,
is employed as a pharmaceutical preparation under the name zinci
carbonas.

CADMIUM.

Symbol, Cd. Atomic weight= 112.4.

Occurrence. — Cadmium is never found in the uncombined state.
The only natural compound of which cadmium is the chief con-
stituent is the extremely rare mineral greenockite, which is the
sulphide, CdS. Cadmium occurs in small quantities in many zinc
ores, such as the sulphide and carbonate ; and in the process of
extracting zinc from these ores, the cadmium is obtained in the
first portions of the product of the distillation, partly as metal
and partly as oxide.

Mode of Formation.— The crude product of distillation is dis-
solved in dilute sulphuric or hydrochloric acid, and the cadmium
precipitated as sulphide by means of sulphuretted hydrogen. The
cadmium sulphide is then dissolved in strong hydrochloric acid,
and precipitated as carbonate by means of ammonium carbonate.
The washed and dried carbonate is first converted into oxide by
calcination, and finally mixed with charcoal and distilled.

Properties. — Cadmium is a bluish- white metal resembling zinc
in appearance, but much more malleable and ductile. It tarnishes
superficially on exposure to the air, and, when strongly heated,
burns with the formation of a brown smoke of cadmium oxide,
CdO. The metal is attacked by dilute hydrochloric and sulphuric
acids, with the evolution of hydrogen. It readily dissolves in nitric
acid, yielding the nitrate, with the formation of oxides of nitrogen.
Cadmium is less electro-positive than zinc, and is precipitated in
the metallic condition from its solutions by that metal.

Cadmium melts at 320°, and boils about 745°. When volatilised

Mercury 597

in an atmosphere of hydrogen, it forms crystals belonging to the
regular system.

Cadmium Oxide, CdO, is formed as a brown fume or smoke
when cadmium burns in the air. It may be obtained by heating
the carbonate or nitrate. That obtained by the ignition of the
latter salt is in the form of minute crystals, having a bluish-black
appearance. Cadmium oxide is insoluble in water, but dissolves
in acids yielding cadmium salts. It is infusible in the oxyhydrogen
flame, but is readily reduced when heated on charcoal before the
blowpipe ; and the reduced metal, as it volatilises and burns, forms
a characteristic brown incrustation of oxide upon the charcoal.

Cadmium Chloride, CdCl2, is obtained by the action of hydro-
chloric acid upon the metal or the oxide. The salt is deposited
from the solution in white silky crystals, having the composition
CdCl2,2H2O. On exposure to the air the crystals effloresce, and
when heated become anhydrous.

Cadmium Sulphide, CdS, is obtained as a bright yellow preci-
pitate when sulphuretted hydrogen is passed through a solution of
a cadmium salt. The precipitate is soluble in concentrated hydro-
chloric and nitric acids, and in warm dilute sulphuric acid. Cad-
mium sulphide is insoluble in ammonium sulphide ; this property
readily distinguishes it from arsenious sulphide, which in colour
it closely resembles.

Cadmium sulphide is used as a pigment, both in oil and water-
colours.

MERCURY.

Symbol, Hg. Atomic weight = 200.

Occurrence. — In the uncombined state mercury is met with in
small globules, disseminated through its ores, especially the sul-
phide. It is also occasionally found as an amalgam with silver
and gold. The principal ore is cinnabar ; HgS, and the chief
mines of this ore are those of Almaden (Spain), Idria (Carniola),
California, and the Bavarian Palatinate.

Modes Of Formation.— Mercury may be obtained from the
natural sulphides by either roasting the ore, whereby the sulphur
is oxidised to sulphur dioxide and the metal liberated, or by dis-
tillation in closed retorts with lime, when calcium sulphide and
sulphate are formed, and the mercury set free. The first method
is almost exclusively employed.

598

Inorganic Chemistry

At Idria the crude ore, consisting of cinnabar mixed with shale
and earthy matters, is roasted in a furnace, upon perforated arches,
n, ri \ p-tp'i Fig. 143. The action of the fire and heated air is to
oxidise the sulphur and volatilise the mercury, and the gases and
vapours together pass through a series of flues or chambers, C, C,
where the mercury condenses.

By the use of a reverberatory furnace (the Albert! furnace), the
process can be made continuous The ore is fed into the furnace

FIG.

through a hopper, and the calcined residue is raked out through
an opening at the opposite end of the hearth. The gases are
passed first through iron pipes kept cool by water, and then
through a series of chambers where the remaining metal is
condensed. . ;

The method adopted at Almaden is essentially the same as the

Idrian process, except that the

FIG. 144. ware vessels, called aludels^ which

are connected together as shown

in Fig. 144. Usually six rows of forty-seven such aludels are
connected with six openings in a chamber immediately above the
furnace.

The impure mercury is freed from mechanically mixed impurities
by straining or filtering through chamois leather, but from metals
in solution, such as zinc, tin, lead, and others, it is purified by
distillation. For laboratory purposes, pure mercury is best ob-
tained by distillation in vacuo, by means of the apparatus shown in

Mercury

599

Fig. 145 (Clarke). In this arrangement the mercury is distilled
in a Sprengel vacuum. The mercury (previously cleaned by being
thoroughly agitated with mercuric nitrate) is placed in the reser-
voir R, which is then placed upon the upper shelf S, and by means
of the clamp, mercury is allowed to pass into the long wide tube T,
and up into the bulb. The air in the tube and bulb escapes down
the narrow inner tube, which reaches nearly to the top of the bulbj

FIG. 145.

as seen in the enlarged detail, /. The mercury is allowed to rise
in the bulb and fall down the long inner tube, after the manner of
the Sprengel pump. The reservoir is then placed upon the lower
adjustable stand, and its height so arranged that the mercury in
the bulb falls to the position shown in the figure. This space is a
Torricellian vacuum. The mercury is then heated by a ring-
burner B, and the whole is protected from draught by the hood h.

6oo Inorganic Chemistry

As the mercury distils, it passes down the inner tube, and by its
fall continues to preserve the Sprengel vacuum within the bulb.

Properties. — At ordinary temperatures mercury is a bright, silver-
white liquid metal (hence its old name quicksilver^ i.e. live silver).
When cooled to - 38.8° it solidifies to a highly crystalline solid,
which is ductile and malleable, and softer than lead. When the
liquid is cooled it contracts uniformly until the solidifying point
is reached, when considerable contraction takes place. Solid
mercury, therefore, is denser than the liquid metal, and sinks in
it. The specific gravity of liquid mercury at o° is 13.596, while
that of the solid at its melting-point is 14.193. Mercury in
extremely thin films appears a violet colour by transmitted light.

Under a pressure of 760 mm. mercury boils at 357.25°, giving
a colourless vapour. The density of mercury vapour referred to
hydrogen is 100.15 > hence this element, like its associates in the
family to which it belongs, consists of mono-atomic molecules when
in a state of vapour. Mercury gives off vapour even at ordinary
temperatures, and a gold leaf suspended over mercury in a stop-
pered bottle gradually becomes white upon the surface, owing to
its amalgamation with the mercurial vapour.

The vapour of mercury is poisonous, giving rise to salivation.

Mercury does not tarnish on exposure to the air, and is unacted
upon by a large number of gases ; hence this liquid is invaluable
to the chemist, affording a means of collecting and measuring
gases which are soluble in water.

When submitted to prolonged heating in the air it is slowly
converted into the red oxide, which at a higher temperature is
again decomposed into its elements.

Mercury is obtained in the form of a dull-grey powder when it
is shaken up with oil or triturated with sugar, chalk, or lard. This
operation is known as deadening^ and is made use of in the pre-
paration of mercurial ointment. The grey powder consists simply
of very finely divided mercury in the form of minute globules.

Mercury is not attacked by hydrochloric acid. Strong sulphuric
acid is without action upon it in the cold, but when heated the
metal dissolves, with evolution of sulphur dioxide. Strong nitric
acid rapidly attacks it, with formation of mercuric nitrate and
oxides of nitrogen. Cold dilute nitric acid slowly dissolves it,
forming mercurous nitrate.

Alloys Of Mercury. — When mercury is one of the constituents
of an alloy the mixture is called an amalgam. Most metals will

Salts of Mercury 60 1

form an amalgam with mercury. In some cases, aS with the
alkali metals, the union is attended with great rise of temperature.
In other cases, as with tin, an absorption of heat takes place.

Sodium and potassium amalgams are obtained by dissolving
various amounts of the metals in mercury. In contact with water
they are decomposed, hydrogen being evolved and the alkaline
hydroxide formed. On this account sodium amalgam is frequently
used in the laboratory as a reducing agent. When heated to
440° these amalgams leave behind crystalline compounds, K2Hg
and Na3Hg, which spontaneously inflame in contact with the air.

Zinc amalgams are only very slowly acted upon by dilute sul-
phuric acid ; therefore, by the superficial amalgamation of the
zinc plates used for galvanic batteries, the same result is ob-
tained as though the zinc were perfectly pure (see page 592)> and
no solution of zinc takes place until the electric circuit is closed.

Tin amalgams are employed for the construction of ordinary
mirrors.

Amalgams of gold, and also copper and zinc, are used in
dentistry as a filling or stopping for teeth.

Oxides of Mercury. — Two oxides are known, namely, mercu-
rous oxide, Hg2O, and mercuric oxide, HgO.

Mereurous Oxide, Hg2O, is obtained as an unstable dark-brown
or black powder when sodium hydroxide is added to mercurous
chloride. When exposed to the light, or when gently heated, it is
converted into mercuric oxide and mercury.

Mercuric Oxide, HgO, is produced in small quantity by the pro-
longed heating of mercury in contact with air, or by igniting the
nitrate. It is prepared on a large scale by heating an intimate
mixture of mercuric nitrate and mercury. Obtained by these
methods, it is a brick-red crystalline powder ; but when sodium
hydroxide is added to a solution of a mercuric salt, the oxide is pre-
cipitated as an orange-yellow amorphous powder. When heated,
mercuric oxide first darkens in colour, and gradually becomes
almost black, but returns to its original bright red colour on cool-
ing. At a red heat it is completely decomposed into its elements.

Salts Of Mercury. — Two series of salts, corresponding to the
two oxides, are known — (a) mercurous salts, in which two atoms of
the hydrogen of the acids are replaced by the divalent radical
or double atom (Hg2) ; and (ft) mercuric salts, in which the same
amount of hydrogen is replaced by the single divalent atom (Hg).
All the mercury salts are poisonous.

602 Inorganic Chemistry

(a) MERCUROUS SALTS.

Mereurous Chloride, Hg2Cl2 (calomel\ is met with in small
quantities as the mineral horn mercury. It may be obtained by
the addition of sodium chloride or hydrochloric acid to a solution
of mercurous nitrate. On a large scale it is usually prepared by
heating a mixture of mercuric chloride and mercury, when the
mercurous chloride sublimes as a white or translucent fibrous
cake.

When a mixture of mercuric sulphate, common salt, and mercury
is heated, mercurous chloride is also obtained, thus —

HgSO4+2NaCH-Hg = Na2SO4 + Hg2Cl2.

Calomel is perfectly tasteless, and is insoluble in water. When
heated it vaporises without fusing. The density of the vapour
that is formed by heating mercurous chloride is 117.87, which is
half that demanded by the formula Hg2Cl2. It has been shown,
however, that the compound dissociates when vaporised into
mercuric chloride and mercury.* Boiling hydrochloric acid de-
composes mercurous chloride into mercury, which separates out,
and mercuric chloride, which dissolves.

Mereurous Nitrate, Hg2(NO3)2, is deposited in the form of
colourless monosymmetric crystals containing 2H2O, from a solu-
tion of mercury in cold dilute nitric acid. The salt is soluble in
water acidulated with nitric acid, but an excess of water causes
the precipitation of a basic nitrate having the composition —

Hg2(N03)2,Hg20,H20 (or 2Hg2(NO3)(HO)),

which, on boiling, is converted into mercuric nitrate and mercury.
If either this or the normal salt be boiled in the presence of an
excess of mercury, a basic nitrate of the composition —

3Hg2(N03)2,2Hg20,2H20 (or Hg2(NO3)2,4Hg,(NO3)(HO)),

is obtained.

Mereurous Sulphate, Hg2SO4, is obtained as a white crystalline
precipitate when dilute sulphuric acid is added to a solution of
mercurous nitrate. It is very slightly soluble in water.

* Harris and Meyer, Berichte, June 1894.

Mercuric Iodide 603

(/S) MERCURIC SALTS.

Mercuric Chloride, HgC\2 (corrosive sublimate] , is formed when
chlorine is passed over heated mercury. It is prepared on a large
scale by heating a mixture of mercuric sulphate and common salt,
a small quantity of manganese dioxide being added to prevent,
as far as possible, the formation of mercurous chloride. The
mercuric chloride sublimes as a white translucent mass. It dis-
solves in water to the extent of 6.57 parts in 100 parts of water at
10°, and 54 parts in the same volume of water at 100°, forming an
acid solution from which the salt is deposited in long white silky
needles. It readily melts, and volatilises unchanged. It dissolves
without decomposition in nitric acid and in sulphuric acid, and
volatilises unchanged from its solution in the latter acid on
boiling.

Mercuric chloride is a violent poison : the best antidote is albu-
men, with which it forms an insoluble compound. It has also
strong antiseptic properties, and on this account is largely used by
taxidermists.

With hydrochloric acid, mercuric chloride forms two crystalline
double chlorides, HgCl2,HCl and 2HgCl2,HCl; and with the
alkaline chlorides it forms a number of similar double salts, of
which the ammonium compound, HgCl2,2NH4Cl,H2O, was known
to the early chemists under the name sal alembroth.

Mercuric Iodide, HgI2. — When mercury and iodine are rubbed
together in a mortar, and moistened with a small quantity of
alcohol, the red mercuric iodide is formed. It is also obtained by
precipitation from a solution of mercuric chloride, upon the addition
of potassium iodide. The precipitate first appears yellow, but in a
few seconds becomes scarlet.

Mercuric iodide is insoluble in water, but readily dissolves in
either mercuric chloride or potassium iodide, and also in alcohol
and in nitric acid. From its solutions it is deposited in scarlet
tetragonal pyramids.

Mercuric iodide is dimorphous ; when heated to about 150° the
scarlet crystals are changed into bright yellow orthorhombic
prisms. At ordinary temperatures this yellow form is unstable,
and on being lightly touched it is at once retransformed into
the red modification. At very low temperatures, however, the
yellow variety is the more stable : thus, when the red crystals

604 Inorganic Chemistry

are exposed to the temperature of evaporating liquid oxygen, they
pass into the yellow variety.

Mercuric Nitrate, Hg(NO3)2, is prepared by boiling nitric acid
with mercury, until sodium chloride produces no precipitate with a
sample of the liquid. If this solution be evaporated over sulphuric
acid, deliquescent crystals are obtained of 2Hg(NO3)2,H2O, while
the mother-liquor has the composition Hg(NO3)2,2H2O.

Mercuric nitrate exhibits a great tendency to form basic salts :
thus, when this mother-liquor is boiled, the compound Hg(NO3)2,
HgO,2H2O is precipitated. When this compound, or the normal
nitrate, is treated with an excess of cold water, there is formed the
still more basic salt Hg(NO3)2,2HgO,H2O.

Mercuric Sulphide, HgS (cinnabar).— When mercury and
sulphur are triturated together in a mortar, or when excess of
sulphuretted hydrogen is passed into a solution of a mercuric salt,
mercuric sulphide is obtained as a black amorphous powder. If
this be sublimed, it is obtained as a red crystalline substance.

Mercuric sulphide in the red condition is also obtained by
digesting the black amorphous product for some hours in alkaline
sulphides. A soluble double sulphide is first formed, which when
heated is decomposed, with the deposition of red mercuric sulphide.
This compound is manufactured on a large scale for use as the
pigment vermilion.

Mercuric sulphide is insoluble in either nitric, hydrochloric, or
sulphuric acid. In the presence of an alkali it is soluble in sodium
or potassium sulphide, and deposits crystals from these solutions
having the composition HgS,Na2S,8H2O, and HgS,K2S,5H2O
respectively.

Ammoniaeal Mercury Compounds.— These may be regarded
as ammonium salts, in which two atoms of hydrogen in ammonium
(NH4) have been replaced by either (Hg2) in the mercurous, or by
(Hg) in the mercuric compounds ; the two atoms so replaced being
either drawn from one and the same ammonium group, or from two.

(a) MERCUROUS COMPOUNDS.

Mereurous Ammonium Chloride, (NH2Hg2)Cl, is the black
powder produced by the action of aqueous ammonia upon calomel,
thus—

Mercuric Compounds 605

Mereurous Ammonium Nitrate, (NH2Hg2)NO3, is formed, to-
gether with other compounds, when aqueous ammonia is added to
mercurous nitrate.

Mereurous Diammonium Chloride, ^H3ci \ Hg2 or (NHs)2
Hg2Cl2, is obtained when calomel absorbs dry gaseous ammonia.
On exposure to the air it gives up its ammonia, and is reconverted
into mercurous chloride.

(j8) MERCURIC COMPOUNDS.

Mercuric Ammonium Chloride, (NH2Hg)Cl (infusible white
pretipitate\ is formed when ammonia is added to a solution of
mercuric chloride —

HgCl2-f2NH3=(NH2Hg)Cl + NH4Cl.

Dimereurie Ammonium Chloride, (NHg2)Cl, is obtained by
the action of water on the preceding compound.

Mercuric Diammonium Chloride, £J ^3£ j j Hg, or (N H3)2HgCl2

(fusible white precipitate), is obtained by adding mercuric chloride
to a boiling aqueous solution of ammonium chloride and ammonia,
until the precipitate which first forms no longer dissolves. On
cooling, the solution deposits small crystals belonging to the
regular system.

Oxy-dimereurie Ammonium Iodide, (NH2Hg)l,HgO, is pro-
duced by the action of aqueous ammonia upcn mercuric iodide,
thus—

4NH3 + 2HgI2 + H2O = (NH2Hg)I,HgO + 3NH4I.

It is readily produced as a brown precipitate by adding ammonia
to a solution of mercuric iodide in potassium iodide containing an
excess of potassium hydroxide.

The alkaline solution of potassium mercuric iodide is known as
NessleSs solution, and constitutes a delicate reagent for detecting
the presence of ammonia. Minute traces of free ammonia in solu-
tion produce a yellow or brown coloration with this test.

CHAPTER VIII
THE ELEMENTS OF GROUP III

Family A. Family B.

Scandium, Sc . . 44.1 Boron, B

Yttrium, Y . , 89
Lanthanum, La . . 139
Ytterbium, Yb . . 172

Aluminium, Al . . 27.1

Gallium, Ga -. . 70

Indium, In . . 115

Thallium, Tl . . 204

WITH the exception of boron, aluminium, and thallium, the mem-
bers of this group are amongst the rarest of the elements.* Some
of these occur only in minute traces in certain ores of other metals :
such is the case with the elements gallium and indium, which are
met with in certain specimens of zinc-blende, the ore being con-
sidered rich in gallium if it contains as much as 0.002 per cent, of
this element. Both gallium and indium were discovered by means
of the spectroscope ; the latter by Reich and Richter (1863),. and
named indium on account of two characteristic lines in the indigo-
blue part of the spectrum ; gallium by Lecocq de Boisbaudran
(1875), and named after his own country. The spectrum of this
metal is characterised by two violet lines. One of the most
remarkable properties of gallium is its extremely low fusing-point,
the metal melting at 30.15°. (For a comparison of the properties
of gallium with Mendelejeff's eka-aluuiinium, see page 124.)

Others of these elements are met with in certain rare minerals,
thus, lanthanum occurs in the mineral orthite (from Greenland) ;
and both yttrium and lanthanum (associated also with the rare
elements cerium and erbium) are found in gadolinite or ytterbite
(from Ytterby).

Boron (the typical element of the group) is the only non-metal :
all the others exhibit well-marked metallic properties. They all
yield sesquioxides of the type R2O3 ; in the case of boron this
oxide, B2O3, is acidic.

* For detailed descriptions of the rare elements, the student is referred to
larger treatises, or to chemical dictionaries.

606

Boron 607

Thallium in many respects is peculiar, It forms two series of
compounds ; in one class it functions as a monovalent, and in the
other as a trivalent element. In some of its properties it exhibits a
close analogy to the alkali metals ; thus, it forms a soluble strongly
alkaline hydroxide, T1HO, corresponding to KHO. And many of
its salts, such as the sulphate, T12SO4 ; perchlorate, T1C1O4, arid
the phosphates, are isomorphous with the corresponding potassium
compounds.

Thallium also shows many properties in common with lead,
which in the periodic system is the next element in the series
(the fourth long series). Thus, the chloride, like lead chloride,
is thrown down as a white curdy precipitate on the addition of
hydrochloric acid to a soluble salt of the metal, and like lead
chloride, thallous chloride is soluble in hot water. Thallous
iodide also closely resembles lead iodide, being formed as a yellow
crystalline precipitate when potassium iodide is added to a soluble
thallous salt.

Metallic thallium also bears the closest resemblance to metallic
lead.

In the thallic compounds this element is more closely related to
the other members of this family : thus, thallic oxide, T12O3 ; thallic
chloride, T1C13 ; and thallic sulphide, T12S3, are analogous to the
corresponding boron compounds, B2O3, BC13, B2S3.

BORON,

Symbol, B. Atomic weight =11,

Occur renee.— The element boron has never been found in the
free state. In combination it occurs principally as boric acid in
volcanic steam, and as metallic borates, of which the commonest
are tincal, a crude sodium borate, or borax, Na2B4O7,10H2O ;
boracite and colemanite, or borate spar, Ca2B6On ; and boronatro-
caltite, or ulexite, Ca2B6On,Na2B4O7,16H2O.

Modes of Formation. — (i.) Boron may be prepared by heating
boron trioxide with either sodium or potassium in a covered
crucible —

2B2O3 + 6Na = 3Na2O2 + 4B.

The fused mass is boiled with dilute hydrochloric acid, and the

608 Inorganic Chemistry

boron, which is in the form of a dark-brown powder, is separated
by filtration.

(2.) The element may also be obtained by heating potassium
borofluoride with potassium —

BF3,KF + 3K = 4KF + B.

(3.) Boron is also formed when potassium is heated in the
vapour of boron trichloride —

BC13 + 3K = 3KC1 + B.

Properties. — Boron, as obtained by these methods, is a dark
greenish-brown powder. When strongly heated in air it burns,
uniting both with oxygen and nitrogen, forming a mixture of boron
trioxide, B2O3, and boron nitride, BN. It is unacted upon by air
at ordinary temperatures.

Boron has no action upon boiling water, but cold nitric acid
converts it into boric acid —

B + 3HNO3=H3BO3
When heated with sulphuric acid it is similarly oxidised —

When fused with alkaline carbonates, nitrates, sulphates, and
hydroxides it forms borates of the alkali metals, thus —

2B + 3Na3CO3 = 2Na3BO3

2B + 6KHO =2K3BO3 + 3H2.

Boron dissolves in molten aluminium, which on cooling deposits
crystals of a compound of aluminium and boron.*

Boron Trioxide, B2O3, is formed when boron burns in the air or
in oxygen. The readiest method for its preparation consists in
heating boric acid to redness, when it fuses and gives up water —

2B(HO)3 = 3H2O + B2O3.

Properties. — The fused mass solidifies to a transparent, colour

* This compound was at one time mistaken for an allotropic modification of
boron.

Orthoboric Acid 609

less, vitreous solid, which gradually absorbs atmospheric moisiure
and becomes opaque. It is not volatile below a white heat, and
on this account, although only a feeble acid, it is capable at high
temperatures of displacing strong acids which are volatile from
their combinations ; thus, when boron trioxide is fused with potas-
sium sulphate, potassium borate is formed and sulphur trioxide
expelled —

B2O3 + 3K2SO4=2B(KO)3 + 3SO3.

Boron trioxide at a high temperature is capable of dissolving
many metallic oxides, some of which impart to the fused mass a
characteristic colour.

Boron forms three oxyacids, namely —

Orthoboric acid, B(HO)3, or H3BO3.

Metaboric acid, B2O2(HO)2, or H2B2O4, or B2O3,H2O.

Pyroboric acid, B4O6(HO)2, or H2B4O7, or Sfc/^HjO.

Orthoborie Acid, or Boric Acid, B(HO)3, occurs naturally,
both in the waters and in the jets of steam which issue from the
ground in many volcanic districts, notably in Tuscany.

The actual amount of boric acid in these natural jets of steam
or soffioni is very small ; but as the steam becomes condensed in
the pools of water or lagoons which often surround the jets, the
amount of boric acid with which the water becomes charged is suffi-
cient to constitute this a profitable source of supply. To obtain the
acid, large brick-work basins are built round the steam jets in such
a manner that the liquid can be caused to flow from one to another.
Water is placed in the highest basin, and after the steam from the
fumaroles beneath it has blown through for twenty-four hours the
liquid is passed on to the second basin, and a fresh supply of water
is run into the first. In this way the water passes on through a
series of four or five such basins, receiving the steam of the soffioni
for twenty-four hours in each. The muddy liquor, after passing
through a settling reservoir, is concentrated by evaporation, the
heat from the natural steam being utilised. The concentrated
liquor, having a specific gravity about 1.07, is allowed to cool
in lead-lined tanks ; and the crystals, after being drained, are
dried upon the floor of a chamber, also heated by the natural
steam. The crude boric acid thus obtained is purified by recrys-
tallisation.

6iO Inorganic Chemistry

Boric acid may be prepared by the action of sulphuric acid or
hydrochloric acid upon a strong solution of borax —

Properties. — Boric acid crystallises in lustrous white laminae,
which are soft and soapy to the touch. 100 parts of water at 18°
dissolve 3.9 parts of the acid. The aqueous solution turns blue
litmus to a port wine red, similar to the colour produced by car-
bonic acid. In contact with turmeric paper it gives a brown
stain resembling that caused by alkalies, but readily distinguished
by not being destroyed by acids and by being turned black in
contact with a solution of sodium hydroxide. Boric acid is more
soluble in alcohol than in water, and when this solution is boiled
a portion of the boric acid volatilises with the alcohol and imparts
a green colour to the flame of the burning vapour.

The orthoborates are mostly unstable salts.

Metaborie Acid, H2B2O4, is obtained when boric acid is heated
to 1 00° —

2H3BO3=2H20 + H2B2O4.

The metaborates are more stable salts than the orthoborates
The acid is dibasic, and forms normal and acid salts as well as
super-acid salts, thus —

Normal potassium metaborate . . K2B2O4.
Acid potassium metaborate . . HKB2O4.
Super-acid potassium metaborate . HKB2O4,H2B2O4.

Pyroborie Aeid, H2B4O7, is obtained by heating either meta-
boric acid or orthoboric acid to 140° for some time —

2H2B2O4= H2O + H2B4O7.
4H3BO3=5H2O + H2B4O7.

Borax. — The most important salt of pyroboric acid is the sodium
salt, ordinary borax, Na2B4Or,10H2O. This compound occurs
naturally as the mineral tincal. It is manufactured from boric
acid by double decomposition with sodium carbonate —

Anhydrous sodium carbonate is added to a boiling solution of

Boron Trifluoride 611

boric acid, and the liquid is then allowed to crystallise, when it
forms large transparent prisms belonging to the mono-symmetric
system of the composition Na2B4O7,10H2O.

The chief source of borax, however, is furnished by the natural
deposits of borate of lime in Bolivia. The powdered mineral is
boiled with water, and soda ash is added to the mixture, when
calcium carbonate is precipitated, and a mixture of borax an/d
sodium metaborate is formed —

Ca2B6On + 2Na2CO3 = 2CaCO3 + Na2B4O7 + Na2B2O4.

On crystallisation the borax deposits, and the more soluble
metaborate remains in the mother-liquor. On concentrating
these mother-liquors and blowing carbon dioxide through the
solution, the metaborate is converted into borax, which is pre-
cipitated as a fine meal, leaving sodium carbonate in solution —

2Na2B2O4 + CO2= Na2CO3 + Na2B4O7.

When heated, borax loses its water of crystallisation and swells
up, forming a white porous mass, which finally melts to a clear glass.

100 parts of water at 10° dissolve 4.6 parts of crystallised
borax, and at 100°, 201.4 parts; the solution possesses a feeble
alkaline reaction.

When deposited slowly from warm solutions (i.e. above about
50° C), borax crystallises in octahedrons having the composition
Na2B4O7,5H2O ; but when crystallised without any special precau-
tions it forms prismatic crystals containing 10 molecules of water.
This is the ordinary form in which borax is met with.

Boron Trifluoride, BF3, is formed when boron is brought into
fluorine ; the boron takes fire spontaneously in the gas, forming
the trifluoride.

It is also produced when a mixture of dry powdered fluorspar
and boron trioxide is heated to redness in an iron vessel, calcium
borate being at the same time produced —

2B2O3 + 3CaF2 = Ca3B2O6 + 2BF3.

It is more conveniently prepared by heating together fluorspar,
boron trioxide, and sulphuric acid. The reaction may be regarded
as taking place in two stages, thus —

(i.)
(2.)

6l2 Inorganic Chemistry

Properties. —Boron trifluoride is a colourless, pungent-smelling
gas, which fumes strongly in moist air on account of its powerful
affinity for water. So great is this affinity, that a strip of paper
introduced into the gas is charred, by the abstraction of the
elements of water.

Boron fluoride neither burns nor supports the combustion of
ordinary combustibles. When potassium is heated in the gas it
burns brilliantly, forming the borofluoride.

At o° one volume of water dissolves about 1000 volumes of the
gas, the absorption being attended with rise of temperature.

When the gas is passed into water until the solution is distinctly acid, a
mixture of metaboric acid and hydrofluoboric acid is obtained; the former
separates out, while the latter remains in solution—

8BF3+4H2O=H2B.2O4+6HBF4.

When the gas is passed into water until the latter is saturated , a syrup-like
liquid is obtained which chars organic matter and is strongly corrosive. This
liquid is sometimes called fluoboric acid, and contains boron trifluoride and
water in the proportions represented by the formula 2BF3,4H2O ; or it may
be regarded as consisting of metaboric acid and hydrofluoric acid, as ex-
pressed by the formula H2B2O4,6HF.* In presence of an excess of water,
this substance is decomposed into metaboric acid and hydrofluoboric acid.
i When mixed with its own volume of dry ammonia gas, boron fluoride forms
a white crystalline compound, having the composition represented by the
formula BF3,NH3. This substance may be sublimed without change. Two
other compounds with ammonia are known, namely BF3,2NH3, and
BF3,3NH3. These are both colourless liquids, which on being heated give off
ammonia, leaving the solid BF3,NH3.

The salts of hydrofluoboric acid, HBF4> are known as borofluorides (sometimes
ftuoboratcs) , and are formed by the actic i of the acid upon metallic hydroxides —

HBF4+KHO=H2

In many instances their aqueous solutions redden litmus; this is the case
with ammonium borofluoride, NH4BF4, and calcium borofluoride, Ca(BF4)2.

Boron Trichloride, BC13, is produced when boron is heated in
a stream of dry chlorine.

It is most readily prepared by passing dry chlorine over an
intimate mixture of boron trioxide and charcoal, heated to redness
in a porcelain tube. The volatile product is condensed in a tube
immersed in a freezing mixture —

* It is considered very doubtful whether 'this substance can be regarded
as a definite compound.

Boron Sulphide 613

Properties. — Boron trichloride is a mobile, colourless liquid,
boiling at 18.23. ^ fumes in moist air, being decomposed in
contact with water, with formation of boric and hydrochloric
acids —

2O = B(HO)3+3HC1.

Boron trichloride unites directly with dry gaseous ammonia,
with evolution of considerable heat, forming a white crystalline
compound, having the composition 2BC13,3NH3.

Boron Hydride, BH3. — This compound has never been obtained in a state
of purity. When magnesium boride (an impure substance obtained by fusing
boron trioxide and magnesium in a covered crucible) is acted upon by
hydrochloric acid, a gas is evolved which has a characteristic and unpleasant
smell, and which produces headache and sickness when inhaled. The ga«
is largely hydrogen, containing, however, a certain quantity of boron hydride,
which imparts to the flame a green colour, and produces boron trioxide.
When passed through a heated tube, boron is deposited as a brown film.
When burnt with a limited supply of air, or when a cold porcelain dish is
depressed into the flame of the burning gas, a brown stain of boron is
deposited.

Boron Nitride, BN, is formed when boron is strongly heated in nitrogen
or in ammonia. It is best obtained by heating, in a covered platinum
crucible, a mixture of one part of dehydrated borax, and two parts of
ammonium chloride —

Boron nitride is a white amorphous powder. It is insoluble in water, but
is slowly acted upon by boiling Caustic alkalies, with evolution of ammonia —

= K3B03+NH3.
Heated in a current of steam it forms boron trioxide and ammonia—

Boron Sulphide, B2S3, is prepared by heating a mixture of boron trioxide
and carbon (made by mixing boron trioxide and soot with oil, and heating
the pellets out of contact with air) to bright redness in a stream cf vapoui
of carbon disulphide —

=6CO + 26283.

Boron sulphide is a yellowish solid, consisting of small crystals. I* ' ^s
a strong unpleasant smell, and its vapour attacks the eyes. It is immediately
decomposed by water, being converted into boric acid and sulphuretted
bydrogep —

6 14 Inorganic Chemistry

ALUMINIUM.

Symbol, Al. Atomic weight =27.1.

Occurrence. — Aluminium is one of the most abundant of all
the elements, although it has never been found in the uncombined
state. In combination with oxygen as A12O3, it constitutes such
minerals as corundum, ruby, sapphire. As the hydrated oxide,
A12O3,2H2O, it occurs associated with iron oxide in the mineral
bauxite, which constitutes the chief source from which the metal
itself is obtained. As a double fluoride of aluminium and sodium,
AlF3,3NaF, it occurs in the mineral cryolite, and as a hydrated
phosphate in the various forms of turquoise. Aluminium is met
with in enormous quantities in the form of silicate, constituting
the various clays; and as compound silicates in the felspars, and
other common minerals constituting a large proportion of the
solid crust of the earth.

Mode Of Formation.— Prior to the advent of the electric
furnace as a manufacturing agent, aluminium was obtained from
the mineral bauxite by the following method : — The process was
conducted in four stages — (i.) and (2.) The preparation of pure
aluminium oxide, free from iron. (3.) The preparation of a double
chloride of aluminium and sodium. (4.) The reduction of the
double chloride by means of sodium.

(i.) The powdered bauxite (usually, containing about 50 per
cent, of alumina) was mixed with sodium carbonate and heated for
five or six hours in a reverberatory furnace, when carbon dioxide
is evolved and sodium aluminate is formed — •

A12O3 + Na2CO3 = 2NaAlO2 + CO2.

(2.) The sodium aluminate was extracted with water, leaving the
iron in the form of insoluble oxide. Through the filtered liquid
a stream of carbon dioxide was then passed, which decomposes
sodium aluminate, regenerating sodium carbonate, and precipi-
tating hydrated aluminium oxide —

(3.) The purified alumina, after being washed and dried, was mixed
with sodium chloride and powdered wood charcoal, and sufficient
water added to enable the mixture to be worked up into balls,

Aluminium

These were dried and packed into a vertical fireclay cylinder, and
strongly heated in a stream of chlorine —

A12O3 + 3C + 3C12 = SCO + 2 A1C13.

The aluminium chloride combines with the sodium chloride
present in the mixture, forming the double chloride, AlCl3,NaCl,
which volatilises from the retort, and is condensed in a receiver.

FIG. 146.

(4.) Finally the double chloride of aluminium and sodium was
strongly heated with metallic sodium and powdered cryolite (to
serve as a flux) —

AlCl3,NaCl + 3Na = 4NaCl + Al.

Electrolytic Method. — At the present time aluminium is ex-
clusively obtained by means of the electric furnace. The process
is an electrolytic one, the electrolyte being a solution of alumina
in a bath of molten cryolite. One of the most modern forms of

616 Inorganic Chemistry

apparatus for the purpose (Borcher's) is shown in section in
Fig. 146.

It consists of an iron cylinder or crucible C, with a fireclay
bottom F, and thickly lined throughout with alumina, L. The
cathode consists of a steel plate S, let into the bottom of the
crucible, into which is screwed the copper tube T. To prevent
the steel plate from becoming too much heated, and in consequence
combining with the aluminium, an arrangement is made to cir-
culate water through the pipe T.

The anode consists of a thick carbon rod, or bundle of rods,
which can be raised or lowered at will. A few fragments of
aluminium, together with a small quantity of cryolite, is first placed
in the crucible, and melted by bringing the anode down upon it.
The fused button of aluminium then becomes the cathode. The
crucible is then gradually filled up with its charge of cryolite and
bauxite until the entire mass is in a molten state. The aluminium
oxide alone is decomposed in the process, the oxygen escaping
through an opening in the lid, while the metal collects at the bottom
and is drawn off at the tap-hole. Fresh bauxite is added in small
quantities at a time as the action continues. It is found that the
lining of the crucible, although of alumina, is not dissolved to a
very great extent, owing to the cooling of the surface by outside
exposure to the air.

Properties. — Aluminium is a tin-white metal, possessing great
tensile strength. It is very ductile and malleable, but requires
frequent annealing during the process of drawing or hammering.
Its specific gravity is 2.58 ; by hammering and rolling it may be
raised to 2.68. Its power of conducting heat and electricity is
about one-third that of silver. Aluminium is an extremely sono-
rous metal, and when struck it emits a clear and sustained note.
It is not tarnished by air under ordinary circumstances, but when
strongly heated it becomes oxidised ; and in the condition of thin
foil it readily burns in oxygen, forming alumina, A12O3. The metal
melts at a temperature about 655°. Aluminium is scarcely acted
upon by nitric acid of any strength, but readily dissolves in hydro-
chloric acid, and in solutions of sodium or potassium hydroxide
with elimination of hydrogen. When heated with strong sulphuric
acid, aluminium sulphate is formed, and sulphur dioxide is
evolved.

Organic acids are almost without action upon aluminium, but
in the presence of sodium chloride they are capable of dissolv-

Aluminium Oxide 617

ing it to a slight extent. Pure aluminium is scarcely acted upon
by water or steam, but the presence of impurities such as usually
occur in the commercial metal renders it much more readily
oxidised.

Aluminium is a highly electro-positive element, and is capable
of reducing a number of other metals from their combinations with
oxygen or sulphur. Thus, when finely divided aluminium is heated
with the oxides of such metals as manganese, chromium, tungsten,
uranium, along with lime to form a slag, an energetic action takes
place, in which the aluminium combines with the oxygen, and the
metals are thrown out of combination, and are obtained as a
coherent mass. Similarly, iron pyrites is reduced to the condition
of metallic iron, with the formation of aluminium sulphide.

The extreme readiness with which aluminium is able to effect such
reduction, and the exceedingly high temperature which is reached
by the action, have led to some useful applications, such as the
welding of iron, &c. The property may readily be demonstrated
by heating upon a spatula or an iron plate a small quantity of a
mixture of powdered aluminium and copper oxide. When such a
mixture is brought into a Bunsen flame the action takes place
immediately, and is accompanied by an instantaneous and vivid
flash of light.

Alloys Of Aluminium.— The most important of these is an
alloy with copper, containing 10 per cent, of aluminium, and
known as aluminium bronze. This alloy has a yellow colour,
resembling that of gold ; it is scarcely tarnished by exposure to
air, and is susceptible of a high polish. Its specific gravity is 7.69,
and it possesses a tenacity equal to that of steel, and more than
twice that of the best gun-metal. The alloy is malleable, and
yields good castings, and on account of its many valuable pro-
perties it is employed for a variety of purposes.

Aluminium Oxide (alumina), A12O3, occurs native in a colour-
less crystalline condition as corundum, and coloured by traces of
various metallic oxides in such precious stones as ruby, sapphire,
and amethyst. In a less pure condition, it occurs in large quantities
as emery. These naturally occurring crystalline forms of alumina
are extremely hard, ranking second only to diamond. Alumina is
obtained in an amorphous condition by igniting either the pre-
cipitated hydroxide or ammonia alum, thus —

2A1(HO)3 = 3H20 + A1203.

618 Inorganic Chemistry.

It is also obtained by the action of carbon dioxide upon sodium
aluminate (p. 614). In the crystalline form it is obtained by
strongly heating a mixture of aluminium fluoride and boron tri-
oxide —

The boron trifluoride volatilises, leaving alumina in the form of
rhombohedral crystals. Artificial rubies have been obtained by
heating barium fluoride with alumina, and adding a trace of
potassium dichromate.

Amorphous alumina is a soft white powder, insoluble in water,
but dissolved by acids with the formation of aluminium salts ;
after being strongly heated, however, alumina is attacked only
with slowness by hydrochloric or sulphuric acid.

Aluminium Hydroxides. — When ammonia is added to a
solution of an aluminium salt, a white gelatinous precipitate is
obtained, consisting of the trihydrate, A12O33H2O, or A1(HO)3.
If this be heated to 300° it loses water, and is converted into the
mono-hydrate, A12O3,H2O, or AIO(HO). If the precipitation be
made in boiling solution, or if the trihydrate be heated to 100°,
a compound is obtained having a composition expressed by the
formula A12O3,2H2O, or A12O(HO)4.

These compounds are soluble in acids, and all yield the same
aluminium salts.

Aluminium hydroxide unites with many soluble organic colour-
ing-matters, and precipitates them from solution as Ic.kes. Upon
this property depends the use of aluminium salts as mordants in
dyeing and calico-printing : the colouring-matter being held in the
fibres of the material by the aluminium hydroxide, which is pre-
viously precipitated upon the fabric.

Aluminates. — Alumina is capable of acting as a feeble acidic
oxide : thus, the hydroxides are dissolved by sodium or potassium
hydroxide, yielding salts known as aluminates. Certain alu-
minates occur native, such as spinelle (magnesium aluminate),
Al2O3,MgO, or Mg(AlO2)2, and chrysoberyl (beryllium aluminate),
Al2O3,BeO, or Be(AlO2)2. Sodium aluminate is manufactured by
fusing bauxite with sodium carbonate (p. 614), or by boiling
powdered cryolite with milk of lime—

It is also produced when a mixture of cryolite and lime is heated

Aluminium Sulphate 619

to redness. Under these conditions the 07-/$<?aluminate is formed,
which when acted upon by water undergoes hydrolysis into the
w^aluminate and free alkali —

AlF3,3NaF + 3CaO = 3CaF2 + Na3AlO3.
Na3AlO3+ H2O = NaAlO2 + 2NaHO.

Sodium aluminate is largely employed as a mordant in dyeing
and calico-printing owing to the ease with which it is decomposed,
yielding alumina. Thus it is hydrolysed to a considerable extent
by water, the action being greatly accelerated by the addition of
a small quantity of aluminium hydroxide —

It is decomposed by carbon dioxide (p. 614) and also by any
halogen compound of aluminium : powdered cryolite being often
employed for this purpose —

3NaAlO2 + AlF3,3NaF + 6H2O = 6NaF + 4Al(HO)3.

Aluminium Sulphate, A12(SO4)3,18H2O, is found native as the
minerals hair salt and aluminite, the latter being a basic salt

FIG. 147.

having the composition A12O3SO3,9H2O. Large quantities of
commercial aluminium sulphate are made by directly dissolving
bauxite in sulphuric acid. The product, however, contains iron,

62O Inorganic Chemistry

which is detrimental to the technical uses to which the sulphate is
applied, and from which therefore it must be carefully purified.
Pure aluminium sulphate is prepared by dissolving the hydroxide
in sulphuric acid. It forms a white difficultly crystallisable solid.

The Alums. — Aluminium sulphate unites with certain other
sulphates, forming double salts, which belong to a class of com-
pounds known as the alums. The most important of these
compounds is the double sulphate of aluminium and potassium,
A12(SO4)3,K2SO1,24H20, known as potassium alum, or simply
alum.

The alums have the general formula R2(SO4)3,M2SO4,24H2O,
in which R may be either aluminium, iron, chromium, manganese
(indium or gallium), and M a monovalent element or group, such
as sodium, potassium, or ammonium.

These compounds are all isomorphous, crystallising in the
regular system (usually in cubes or octahedra) with twenty-four
molecules of water. Fig. 147 represents a crystal of potassium
alum (A) and potassium chromium alum (B). In naming the
alums* it is usual, when the salt contains aluminium, only to
introduce the name of the monovalent element or group : thus,
ammonium alum, or potassium alum, signifies the double sulphate
of ammonium, or potassium, and aluminium. If, on the other hand,
the compound contains no aluminium, the names of both metals
are used, thus, potassium chromium alum, ammonium iron alum.

A second class of double sulphates is known, which resemble the alums,
although they are not isomorphous with them. These are termed pseudo-
alums. They may be regarded as alums, in which the two atoms of the
monovalent element are replaced by one atom of a divalent element, thus —

Manganese aluminium, pseudo-alum . . Al.2(SO4)3MnSO4,24H2O.

Iron aluminium, pseudo-alum . . . Al2(SO4)3FeSO4,24H2O.

Copper iron, pseudo-alum -' . . . . Fe2(SO4)8CuSO4,24H2O.

Zinc iron, pseudo-alum .- . . . . Feo(SO4)8ZnSO4,24H2O.

Magnesium manganese, pseudo-alum . . . Mn2(SO4)3MgSO4,24H2O.

The alums are all soluble in water, and their solutions have an

* Selenic acid (the selenium analogue of sulphuric acid) forms a similarly
constituted series of double selenates, crystallising in the same form, and
with the same number of molecules of water. The system of nomenclature
adopted for these compounds is the same : thus, ammonium selenio-alum
signifies the double selenate of ammonium and aluminium, while potassium
chromium selenio-alum represents the double selenate of ootassium and
chromium.

Alum 62 1

acid reaction and possess an astringent taste. When heated
they gradually part with water, and at higher temperatures are
broken up into oxides and alkaline sulphates ; in the case of
ammonium alums, leaving only the metallic oxide.

Potassium Alum, Al2(SO4)3,K2SO4,24H.p, is prepared by
the addition of the requisite quantity of potassium sulphate to
aluminium sulphate. A considerable quantity of alum is also
obtained from a naturally occurring basic potassium alum, known
as alum stone , or alunite, which has the composition A12(SO4)3,
K2SO4,2A12O3,8H2O. At Tolfa this is first calcined, and after-
wards lixiviated with water, which dissolves the potassium alum,
leaving alumina undissolved. The alum so obtained is known as
Roman alum; and although it has a reddish colour, due to the pre-
sence of iron, this iron is present only as the insoluble oxide, which
is readily removed, and the salt is in reality extremely pure.

Alunite is also converted into alum, by treating the calcined
mineral with sulphuric acid, and adding the requisite quantity of
potassium sulphate. A large quantity of alum is manufactured
from alum shale, which is a bituminous mineral, consisting chiefly
of aluminium silicate, with finely divided iron pyrites dissemi-
nated throughout the mass. The shale is usually first roasted,
and is then exposed to the action of air and moisture, whereby
the' oxidation of the pyrites is completed. The result of this
oxidation is the formation of sulphuric acid, which, acting upon
the aluminium silicate, forms aluminium sulphate, while the iron
is converted into ferrous and ferric sulphates, and ferric oxide. The
oxidised mass is then lixiviated with water, and, after concentra-
tion, the requisite quantity of potassium chloride or sulphate is
added to the hot liquor. (The use of potassium chloride is pre-
ferable, as by double decomposition the ferrous and ferric sulphates
are converted into the very soluble chlorides, and an equivalent
amount of potassium sulphate is formed.) The liquor is stirred
mechanically during its cooling, whereby the alum is deposited in
small crystals known as alum mea/y which permit of its more
ready purification by recrystallisation.

Alum crystallises in fine colourless regular octahedra, which, on
exposure to the air, become coated with a white efflorescence, due
not to loss of water, but to absorption of atmospheric ammonia, and
the formation of a basic salt.

The solubility of alum in water increases rapidly with rise of
temperature. Thus, TOO parts of water at o° dissolve 3.9 parts of

622 Inorganic Chemistry

alum; at 50°, 44.1 parts; and at 100°, 357.5 parts. Alum is in-
sbluble in alcohol.

When heated to 42°, alum loses 1 1 molecules of water ; and
when heated to 61° in a closed vessel over sulphuric acid, it parts
with 1 8 molecules.

On the application of heat, alum first melts in its own water of
crystallisation, which is gradually expelled, until at a dull red heat
the salt is converted into a white porous mass, known as burnt
alum. At a still higher temperature it is broken up into potassium
sulphate, alumina, and sulphur trioxide. Burnt alum is only very
slowly dissolved by water. The chief use of alum is as a mordant
in dyeing, alum being a salt which is much more easily obtained
in a state of purity than aluminium sulphate. By the addition of
sodium hydroxide or carbonate to a solution of alum, until the
precipitate first thrown down is just redissolved, a basic alum is
produced known as neutral alum —

2Al2(SO4)3,K2SO4+6NaHO = Al2(SO4)3,Al2(HO)6,K2S04 +
3Na2SO4 + K2SO4.

This solution gives up its alumina to the fabric with great ease,
and on this account is used by dyers and calico-printers as a
-mordant. • /•'-

When this solution is heated to 40°, ordinary alum is reformed,
and a precipitate is obtained consisting of another basic salt, hav-
ing the same composition as alunite^ thus —

2Al2(S04)3,Al2(HO)6,K2S04=Al2(S04)3,K2S04)-f
A12(S04)3,2A1203,K2S04 + 6H20.

Aluminium Fluoride, A1F3.— This compound may be prepared by parsing
gaseous hydrochloric acid over a mixture of fluorspar and alumina heated to
whiteness in a graphite tube, when aluminium fluoride volatilises, leaving
calcium chloride —

3CaF2+Al2O3+GHCl=3H2O + 3CaCl2 + 2AlF3,

In the form of a crystalline hydrate it may be obtained by dissolving alumina
in aqueous hydrofluoric acid —

A12O3 + 6HF+H20=2A1F3,7H20.

Aluminium fluoride fcrms colourless rhombohedral crystals, which are in-
soluble in water. It combines with alkali fluorides, forming insoluble double
fluorides, of which the sodium compound is the most important, AlF3,3NaF.
This compound occurs native as the mineral cryolite.

Thallium 623

Aluminium Chloride, A1C13.— This compound is produced
when powdered aluminium is strongly heated in chlorine, or with
certain metallic chlorides, such as zinc chloride. It is best
obtained by passing chlorine over a strongly heated mixture of
alumina and charcoal.

An aqueous solution of aluminium chloride may be obtained by
dissolving alumina in hydrochloric acid. On evaporation the
solution deposits crystals of a hydrate, A1C13,6H2O.

Aluminium chloride forms white hexagonal crystals, which fume
strongly in moist air. When gently heated it vaporises, and sub-
limes without fusion. When heated under pressure of its own
vapour, the compound melts. It dissolves in water with the
evolution of heat, and the solution on evaporation deposits the
hydrated chloride, which, on being heated, breaks up into hydro-
chloric acid, water, and alumina —

A1C13,6H2O = 3HC1 + 3H20 + A1(HO)3.

Aluminium chloride unites with other metallic chlorides, forming
double salts, of which the sodium compound AlCl3,NaCl (p. 615)
is the most important. It also combines with ammonia, forming
the compounds A1C13,6NH3 and A1C13,NH3.

Aluminium Sulphide, A12S3.— When finely divided aluminium
is heated with iron pyrites, an energetic reaction takes place ;
metallic iron being reduced, and aluminium sulphide being formed.
The same compound is produced when sulphur is thrown upon
strongly heated aluminium. As obtained by these methods,
aluminium sulphide is a greyish-black solid, which, when thrown
into water, is converted into the oxide with evolution of sul-
phuretted hydrogen —

The compound is decomposed in the same manner by atmos-
pheric moisture when exposed to the air.

THALLIUM.

Formula, Tl. Atomic weight = 204.

History.— Thallium was discovered by Crookes (1861) in the
seleniferous deposit from a sulphuric acid manufactory. In the

624 Inorganic Chemistry

spectroscopic examination of certain residues obtained in the ex-
traction of selenium from this deposit, the presence of an unknown
element was manifested, by the appearance of one bright green
line. From its characteristic spectrum, the name thallium (signi-
fying a green twig) was given to the element

Occurrence. — Thallium is found in small quantities in many
varieties of iron pyrites, and when these are employed in the
manufacture of sulphuric acid, oxide of thallium collects in the
flue dust of the pyrites burners. Thallium also occurs associated
with copper, selenium, and silver, in the rare mineral crookesite.

Mode of Formation. — The metal is obtained by reducing the
sulphate, by immersing strips of zinc into the solution. The
thallium is deposited upon the zinc, as a spongy or crystalline mass,
which is then pressed together and fused beneath potassium
cyanide in a crucible.

Properties. — Thallium is a soft heavy metal, resembling Jead.
It is readily cut with a knife, and leaves a streak when drawn
across paper. When preserved out of contact with air it is a tin-
white lustrous metal ; but on exposure to the air it tarnishes
upon its surface, with the formation of black thallous oxide. Its
specific gravity is n.8, and it melts at 290°.

When exposed to air and moisture, or when placed in water
which is free to absorb atmospheric oxygen, the metal is slowly
converted into thallous hydroxide, which is soluble in water, and
imparts to the liquid a strong alkaline reaction. The solution
absorbs carbon dioxide, with the formation of thallous carbonate.
When heated in the air thallium melts, and rapidly oxidises to
thallium trioxide, T12O3 ; heated in oxygen it burns, forming the
same oxide. It readily burns when heated in chlorine, producing
thallous chloride, T1C1. The metal is soluble in dilute acids.

Oxides of Thallium. — TWO oxides are known, namely, thallous
oxide, T12O, and thallic oxide, T12O3.

ThallOUS Oxide, T12O, forms as a dark grey film upon the
surface of the metal, on exposure to the air. It may also be
obtained by heating the hydroxide to 100°. It dissolves in water,
forming the hydroxide.

ThallOUS Hydroxide is obtained by the addition of barium
hydroxide to a solution of thallous sulphate, the precipitated barium
sulphate being removed by filtration —

Tl2SO4 + Ba(HO)2=BaSO4

Thallic Chloride 625

The solution, on concentration, deposits yellow needle-shaped
crystals of T1HO,H2O. Thallous hydroxide is soluble in water,
yielding an alkaline solution which gives a brown stain upon
turmeric paper. The stain soon disappears, owing to the de-
struction of the colouring-matter, and is thereby distinguished
from the similar stains produced by sodium and potassium
hydroxides.

Thallic Oxide, T12O3, is obtained when thallium burns in the
air, or when thallium oxyhydroxide, TIO(HO), is heated to 100°.
It forms a dark reddish powder, insoluble in water. In warm
dilute sulphuric acid it dissolves, forming thallic sulphate —

but with hot concentrated acid oxygen is evolved, and thallous
sulphate formed —

At a red heat thallic oxide is converted into thallous oxide with
loss of oxygen.

Thallium Oxyhydroxide, TIO(HO), is formed by the action of
potassium hydroxide upon thallium trichloride —

Thallous Chloride, T1C1, is obtained as a white curdy precipi-
tate when hydrochloric acid is added to a solution of a thallous
salt. It is considerably more soluble in hot than, m cold water :
100 parts of water at 16° dissolve 0.265 Par^ > and at 100°, 1.427
part of thallous chloride.

Thallie Chloride, T1C13, is formed by passing chlorine through
water in which thallous chloride is suspended. The solution so
obtained, on evaporation in vacuo, deposits colourless transparent
crystals of T1C13,2H2O.

When either thallium or thallous chloride is gently heated in a
stream of chlorine, a compound is obtained, having the composition
T1C13,T1C1, or T12C14. If this be further heated, it loses chlorine,
and is converted into a yellow crystalline compound of the com-
position T1C13,3T1C1, or T14C18, thus—

2T12C]4=C12

2 R

626 Inorganic Chemistry

ThallOUS OxysaltS.— The sulphate T12SO4, and nitrate T1NO3,
are best obtained by dissolving the metal in the respective acids.
Both salts are soluble in water.

ThallOUS Carbonate, T12CO3, is prepared by saturating a solu-
tion of thallous hydroxide with carbon dioxide. The salt forms
long white prismatic (monosymmetric) crystals, which are mode-
rately soluble in water, giving an alkaline solution.

ThallOUS Phosphate, T13PO4, is obtained by precipitation from
a thallous solution, by the corresponding potassium phosphate.
The monohydrogen phosphate, HT12PO4, on being heated to 200°,
is converted into pyrophosphate —

2HTl2P04=H20 + Tl4P2Or, -

and the dihydrogen salt, on being ignited, yields the metaphos-
phate —

H2T1PO4 = H2O + T1PO3.

ThalliC OxysaltS.— The chief of these are thallic sulphate,
T12(SO4)3 ; and thallic nitrate, T1(NO3)3. They are obtained by the
action of sulphuric acid and nitric acid respectively upon thallic
oxide T12O3. Thallic sulphate forms colourless crystals of the
composition T12(SO4)3,7H2O. It is decomposed by excess of water,
with precipitation of the hydrated oxide ; and when heated yields
thallous sulphate, sulphur trioxide, and oxygen —

T12(SO4)3=T12SO4 + 2S03 + O2.

Thallicnitrate is depositedin colourless crystalsof T1(NO3)3,8H2O.
which are decomposed in the presence of much water.

CHAPTER IX

THE ELEMENTS OF GROUP IV

Family A.
Titanium, Ti .
Zirconium, Zr .
Cerium, Ce
Thorium, Th

48.1

90.6

140.25

232-5

Family B,

Carbon, C . . . 12.00
Silicon, Si . . . 28.3
Germanium, Ge . . 72.5
Tin, Sn . . .119
Lead, Pb . 207.1

Family A consists of four rare elements.* Titanium, as the
oxide TiO2, occurs in the three rare minerals — rutile, brookite, and
anastase. The metal is extremely difficult to isolate in a pure
state, owing to the fact that it unites directly with nitrogen, forming
a nitride.

Zirconium is met with as the silicate ZrSiO4 (or ZrO2,SiO2) in
the mineral zircon. Like silicon, it has been obtained in two
forms, crystalline and amorphous. The latter variety, when gently
heated, burns in the air, while the crystalline variety requires the
high temperature of the oxyhydrogen flame for its ignition.

Cerium occurs associated with lanthanum in the rare minerals
cerite and orthitey and with yttrium and ytterbium in gadolinite and
ivohlerite.

Thorium is found in the extremely rare minerals, thorite and
orangeite, met with in Norway.

Family B.— In this family the rare element germanium forms
a link between carbon and silicon on the one hand, and tin and
lead on the other.

Carbon (the typical element) is essentially non-metallic, and
forms an acidic oxide. Silicon approaches more nearly to the
metals in its physical properties, but its oxide is still acidic, and
but few compounds are known in which silicon funcdons as a basic
element. Germanium is both metallic and non-metallic ; its oxide

* For descriptions of these rare elements the student is referred to larger
treatises.

627

628 Inorganic Chemistry

unites with acids ; and it also combines with alkaline hydroxides,
forming germanates corresponding to silicates. Tin is a still more
basic element, forming well-marked salts with acids ; but it is also
acidic, and with alkalies forms stannates.

Carbon and silicon exhibit a close relationship. They both
form allotropes, which correspond in many respects. They both
unite with hydrogen, forming the analogous compounds CH4 and
SiH4 ; and with hydrogen and chlorine they form the similarly con-
stituted compounds, chloroform, CHC13 ; and silicon chloroform,
SiHCl3.

Tin and lead approach more nearly to each other, especially in
their physical properties, than to the other members of the family.
They both form compounds, in which the metals function both
as divalent and tetravalent elements ; although in the case of
lead (as often happens with the heaviest metals of a family), the
element exhibits much greater readiness to act in the lower state
of atomicity. Until quite recently (1893) no compound was known
in which an atom of lead is united with four monovalent atoms,
although lead ethide, Pb(C2H5)4, had been obtained. Now, how-
ever, the compound PbCl4 has been produced, corresponding to
SnCl4, which it resembles in many respects ; and still more
recently (1894) the tetrafluoride has been obtained.

Carbon, as usual with the typical elements, stands apart from
the other members of the family in many of its attributes. Thus,
its oxides are both gaseous ; it also forms a vast number of com-
pounds with hydrogen, oxygen, and nitrogen, the study of which
constitutes the science of organic chemistry. This element has
already been treated in Part II. (page 285).

SILICON.

Symbol, Si. Atomic weight =28.3.

Occurrence. — Silicon is not known to occur in the uncombined
state, although in combination it is the most abundant and widely
distributed of all the elements, with the exception of oxygen. In
combination with oxygen, as silicon dioxide or silica^ SiO2, it
occurs as flint, sand^ quart z^ rock crystal, and chalcedony / while
in combination with oxygen and such metals as calcium, magnesium,
and aluminium, it occurs in clay and soil, and constitutes a large
number of the rocks which make up the earth's crust. Silicon, in

Silicon 629

combination with oxygen, is also met with in the vegetable kingdom,
being absorbed by plants from the soil.

Modes Of Formation. — (i.) Silicon maybe obtained by strongly
heating a mixture of potassium silico-fluoride and potassium —

The mass, after cooling, is treated with water, which dissolves
the potassium fluoride, leaving the liberated silicon.

(2 ) This element may also be prepared by heating sodium in a
stream of the vapour of silicon tetrachloride —

iCl4 + 2Na2=Si

(3.) In an impure state, mixed with magnesium silicide, it may
also be obtained by heating a mixture of dry white sand with
about four times its weight of dry magnesium powder in a hard
glass tube.

As obtained by cither of these methods the silicon is in the form
of an amorphous, dark-brown powder.

(4.) Silicon is obtained in a crystalline condition by passing a
slow stream of the vapour of silicon tetrachloride over aluminium,
previously melted in a current of hydrogen ; the volatile aluminium
chloride passes on in the stream of gas, and the liberated silicon
dissolves in the excess of aluminium —

3SiCl4 + 4 Al = 3Si + 2A12C16.

As the mass cools, silicon is deposited in the form of long, lustrous,
needle-shaped crystals.

(5.) The most convenient method for the preparation of crystal-
lised silicon consists in heating in a crucible a mixture of 3 parts
of potassium silico-fluoride, I part of sodium, and 4 parts of granu-
lated zinc. The regulus so obtained contains crystallised silicon.
It is gently heated, and the excess of zinc drained away, the re-
mainder being removed by treatment with acids.

Properties.— Amorphous Silicon, as obtained by the reactions
Nos. I and 2, is a dark-brown amorphous powder, having a specific
gravity of 2.15. When heated in the air it burns with the forma-
tion of silicon dioxide, which, being non-volatile, coats the particles
of the element and protects it from complete oxidation. It burns
when heated in a stream of chlorine, with formation of silicon
tetrachloride. It is insoluble in water, and in all acids except

630 Inorganic Chemistry

hydrofluoric acid, in which it dissolves, with the formation of
silico-fluoric (or hydrofluosilicic) acid and evolution of hydrogen —

On boiling with potassium hydroxide it forms potassium silicate
and hydrogen —

Crystallised Silicon. — As obtained by reactions Nos. 4 and 5,
silicon is a brilliant, steely-grey substance, crystallised in needles
derived from the orthorhombic pyramid. The specific gravity of
the crystals is 2.34 to 2.49. Crystallised silicon does not burn in
oxygen, even when strongly heated ; it burns when heated in
chlorine, and takes fire spontaneously when brought into fluorine.
It is not soluble in any acid except a mixture of nitric and hydro-
fluoric acids. Crystallised silicon is very hard, being capable of
scratching glass. When silicon is exposed to a high temperature,
out of contact with air, it becomes denser and harder, and has
been obtained in the form of small, steel-grey nodules, showing a
crystalline structure, and having a specific gravity as high as 3.0.*

Silicon Hydride, SiH4. — This compound is evolved at the
negative electrode (along with hydrogen) when dilute sulphuric
acid is electrolysed, the electrodes consisting of aluminium con-
taining silicon.

In ah impure condition, also mixed with hydrogen, this gas may

* Although silicon in combination is such an abundant element, constituting,
as it does, about one-fourth of the total weight of the solid crust of the earth,
in the free state it must still be regarded as somewhat of a rarity, and con-
sequently a good deal of uncertainty exists as to its properties. From differ-
ences that have been observed in the substance, as obtained by different
methods, and from the close analogy that exists between silicon and carbon,
it was at one time believed that three allotropes of this element existed, corre-
sponding to those of carbon. Amorphous silicon was considered to represent
charcoal. A crystalline substance obtained by Wohler, by heating potassium
silico-fluoride with aluminium, has been regarded as corresponding to graphite,
and called graphitic silicon ; while the octahedral crystals of silicon prepared
by reactions 4 and 5 given above (Deville) were thought to be the analogue of
diamond ; and this substance has, therefore, been called diamond or adaman-
toid silicon. There is considerable doubt as to whether the silicon obtained
by all these various methods was sufficiently pure to warrant this classification,
and this doubt is not diminished by the recently discovered fact that silicon
unites with carbon, forming a hard crystalline substance which has received
the name carborundum.

Liquid Silicon Hydride 631

be obtained by the action of hydrochloric acid upon magnesium
silicide —

SiMg2+4HCl = 2MgCl2+SiH4.

(Magnesium silicide for this reaction may be prepared by fusing
together in a covered crucible a mixture of dry magnesium chloride
40 parts, dry sodium chloride 10 parts, sodium silico-fluoride 35
parts, and metallic sodium 20 parts.)

Pure silicon hydride is prepared by acting upon triethyl silico-
formate with metallic sodium. The mode of action of the -sodium
is not known ; the ethyl silico-formate breaks up into silicon hydride
and ethyl silicate —

4SiH(OC2H5)3=SiH4 + 3Si(OC2H5)4.

Properties. — Silicon hydride is a colourless gas. As obtained
by the first two methods it inflames spontaneously. The pu -e
gas does not possess this property. Its ignition-point, however,
is very low, and if the gas be slightly warmed, or if a jet of it be
caused to impinge upon an object a few degrees above the ordinary
temperature, the gas at once takes fire and burns with a brightly
luminous flame : it is also rendered spontaneously inflammable
by reduction of pressure or by admixture with hydrogen. When
brought into chlorine the gas takes fire, with formation of silicon
chloride and hydrochloric acid.

When treated with an aqueous solution of sodium or potassium
hydroxide, silicon hydride is decomposed, giving the alkaline
silicate and evolving four times its own volume of hydrogen —

2O = SiO(NaO)2

Liquid Silicon Hydride, Si2HG.— This compound, which has
quite recently been discovered,* is obtained by passing the pro-
ducts from the action of dilute hydrochloric acid upon magnesium
silicide through a vessel cooled by liquid air or oxygen, and
separating the condensed products by fractionation.

Properties. — Liquid silicon hydride is a colourless mobile liquid
boiling at + 52°. It may be frozen by means of liquid air to a
white crystalline solid, melting at - 138°. The liquid is spontaneously
inflammable in air at the ordinary temperature, burning with a bright
white flame and depositing amorphous silicon and silicon dioxide.

* Moissan and Smiles, Comptes Rendus, March 1902.

632 Inorganic Chemistry

If a small quantity of the liquid be vaporised into an atmosphere
of hydrogen, the hydrogen acquires the property of spontaneous
inflammability in contact with the air. Liquid silicon hydride is
immediately attacked by an aqueous solution of potash, yielding
potassium silicate and hydrogen —

Silicon Fluoride, SiF4. — This compound is formed when silicon
is brought into fluorine, the silicon taking fire spontaneously in
the gas.

It is prepared by the action of sulphuric acid upon a mixture o*
powdered fluorspar and sand —

2CaF2+2H2SO4 + SiO2=2CaSO4 + 2H2O + SiF4.

Properties. — Silicon fluoride is a colourless, fuming gas. It is
not inflammable, and does not support combustion. It is decom-
posed by water into hydrofluosilicic acid and silicic acid, hence the
gas cannot be collected over water —

3SiF4 + 3H2O = 2H2SiF0+ H2SiO3.

The silicic acid is precipitated as a gelatinous mass. Each
bubble of gas as it comes in contact with the water is at once
decomposed, and a little sack-like envelope of silicic acid is
formed round it. On filtering the liquid, a solution of hydrofluo-
silicic acid is obtained. When silicon fluoride is passed over
strongly heated silicon, a white powder is obtained having the
composition Si2F6.

Silicon Chloride, SiCl4, is formed when silicon is heated in a
stream of chlorine. Under these circumstances the silicon burns
m the gas.

It is obtained by heating an intimate mixture of silica and
carbon in a stream of chlorine, and passing the products through
a cooled tube —

Properties. — Silicon chloride is a colourless liquid which fumes
strongly in moist air and boils at 58.3°. It is decomposed by
watev into silicic and hydrochloric acids —

= Si(HO)4 + 4HCl,

Silicon Dioxide 633

and the silicic acid so formed passes either entirely or in part
into the dibasic acid, thus —

Si(HO)4=S5O(HO)2+H2O.

Disilicon Hexachloride (also known as silicon trichloride), Si2Cl6, is formed
when the vapour of silicon tetrachloride is passed over strongly heated
silicon —

l4 + Si=2Si2Cl6.

It may be prepared by gently heating the corresponding iodine compound
with mercuric chloride—

Si2I6+3HgCl2=Si2Cl6 f 3HgI2.

Properties. — Disilicon hexachloride is a mobile, colourless, fuming liquid,
which boils at 147° and crystallises at - i°. When the liquid is boiled and the
hot vapour allowed to escape into the air, it spontaneously ignites.

Silicon forms two compounds with bromine and with iodine corresponding
to the chlorides, namely —

SiBr4; Si2Br6 ; SiI4; Si2I6.

Silicon Dioxide, SiO2, occurs in nature in a more or less pure
form * in a large number of minerals, some of which have already
been alluded to as natural compounds of silicon. Silicon dioxide
in an amorphous form is met with in the different varieties of
opal) and in enormous quantities in the deposit known as kiesel-
guhr. This substance consists of the remains of extinct dia-
tomaceas, and is met with in various parts of Germany. In a
crystalline condition silica occurs as quartz or rock crystal, and
also in a rarer form as tridymite.

Modes Of Formation. — (i.) Silicon dioxide is formed when
amorphous silicon is burnt in air or oxygen.

(2.) It may be prepared by heating silicic acid, which readily parts
with water and leaves pure silicon dioxide as a light white amor-
phous powder —

Si(HO)4=SiO2+2H2O; or
SiO(HO)2 = Si02 + H2O.

(3.) in minute crystals, silicon dioxide is obtained by strongly
heating a solution of an alkaline silicate in a sealed glass tube,
whereby a portion of the silica of the glass is dissolved. When
this solution is cooled, silicon dioxide is deposited, if the crys-
tallisation takes place above a temperature of 180°, crystals of
quartz are obtained ; if below this point, it deposits crystals of

634 'Inorganic Chemistry

tridymite, while at ordinary temperatures the silica is deposited
in the amorphous condition. Much larger quartz. crystals have
been obtained by the prolonged heating to 250° of a 10 per cent,
aqueous solution of silicic acid (obtained by dialysis) in stout
sealed glass flasks.

Properties. — In the crystalline condition as quartz, silicon
dioxide forms prismatic crystals belonging to the hexagonal sys-
tem, terminating in hexa-
gonal pyramids. Fig. 148
represents a mass of quartz
or rock crystal.

The purest forms of rock
crystal are perfectly colour-
less, having a specific gravity
of 2.69, and are sufficiently
hard to cut glass. When cut
and polished, it exhibits a
brilliancy not far inferior to
that of the diamond, and is
occasionally substituted for
this gem.

Quartz is often found
coloured by the presence of
small quantities of impurities,
as in the varieties known as
amethyst quartz and smoky
qttartZ) and in great quanti-
ties as milky quartz.

The variety of silicon di-
oxide known as tridymite is
found as minute crystals in cavities in certain specimens of trachytic
rocks. The crystalline form of tridymite, although belonging to
the hexagonal system, is distinct from that of quartz, and the crystals
are frequently met with grown together in the manner known as
twin-crystals.

Amorphous silicon dioxide, as it occurs in nature, is a translu-
cent substance having a conchoidal or vitreous fracture ; its specific
gravity is 2.3. As artificially prepared, it is a soft white powder
whose specific gravity is 2.2. At the temperature of the oxy-
hydrogen flame, silicon dioxide melts to a transparent glass-like
substance which is capable of being drawn out into fine threads

FIG. 148.

Silicic Acids 635

resembling spun glass. These fibres possess many valuable pro-
perties, and are employed by physicists in delicate instruments of
precision.

Silicon dioxide is insoluble in water and in all acids with the
exception of hydrofluoric acid. It dissolves in alkalies, and the
amorphous powder can be dissolved in a boiling solution of sodium
carbonate. Many natural hot springs contain silica held in solu-
tion as an alkaline silicate, and on exposure to atmospheric carbon
dioxide the silicate is decomposed with the deposition of silica and
the reformation of an alkaline carbonate. The enormous quantities
of siliceous sinter deposited by geysers at Rotomahama, New Zea-
land, were formed in this way. When fused with sodium carbo-
nate, silicon dioxide is converted into soluble sodium silicate —

SiO2+2Na2CO3=2C02+Si(NaO)4.

Silicic Acids. — Silicon dioxide is capable of forming weak
polybasic acids, but from the readiness with which they give up
water it is probable that none have ever been obtained in a state
of purity. The compound represented by the formula Si(HO)4 is
known as orthosilicic acid, and is tetrabasic. By the loss of one
molecule of water it forms metasilicic acid, SiO(HO)2. When
hydrochloric acid is added to a solution of an alkaline silicate, a
gelatinous precipitate is obtained, which consists of the dibasic
acid SiO(HO)2, or H2SiO3—

SiO(NaO)2+2HCl = SiO(HO)2 + 2NaCl.

If, on the other hand, the solution of alkaline silicate be added
cautiously to an excess of hydrochloric acid, the silicic acid remains
in solution, and is probably present as orthosilicic acid, Si(HO)4, or
H4SiO4—

SiO(NaO)2 + 2HCl + H2O = Si(HO)4 + 2NaCl.

The sodium chloride in the solution may be removed by a pro-
cess of separation known as dialysis. This process, discovered by
Graham, is based upon a property belonging to certain classes of
substances, of passing when in solution through certain mem-
branes. The mixture is placed in an apparatus resembling a
small tambourine (Fig. 149) (made by stretching either parch-
ment or parchment paper over a wooden hoop), which is then
floated upon water. The sodium chloride passes through the

636 Inorganic Chemistry

membrane, while the silioic acid remains behind in the dialyser
as a dilute aqueous solution. Substances in solution which are
capable of readily diffusing through such a membrane were termed
by Graham crystalloids; while others, such as the silicic acid,
which either do not pass through or only do so with difficulty, are
known as colloids.

This aqueous solution of silicic acid may be concentrated by
boiling, and further by evaporation in vacuo over sulphuric acid,
until it contains about 21 per cent, of tetrabasic silicic acid, or 14
per cent, of silicon dioxide. In this condition it is a tasteless
liquid, having a feeble acid reaction. It cannot be preserved, as

FIG. 149.

on standing it solidifies to a transparent gelatinous mass, which
has approximately the composition H2SiO3.

Silicates. — The silicates constitute a large class of important minerals,
many of which are of extremely complex composition. Some of the simplest
of these silicates are derived from the dibasic and tetrabasic acids already
described, while others may be regarded as the salts of a number of hypo-
thetical polybasic silicic acids, derived from metasilicic acid by the gradual
elimination of water. Thus, by the withdrawal of one molecule of water from
two molecules of metasilicic acid, an acid known as disilicic acid is obtained,
having the composition Si2O3(HO)2, or 2SiO.2,H2O, or H2Si2O5— -

2SiO(HO)2= H20 -f Si203(HO)2.

By the abstraction of on-3 molecule of water from two molecules of ortho
silicic acid another disilicic acid is similarly derived —

2Si(HO)4=H20-rSi20(HO)6, or 2SiO2,3H2O, or H6Si2O7.

Tin 637

By the partial withdrawal of water from three molecules of silicic acid a
number of hypothetical trisilicic acids may be derived, such as —

3SiO2,2H2O or H4Si3O3 ; 3SiO2,5H2O or H10Si3On ;
3SiO2,7H2O or H14Si3O13.

Silicates derived from an acid containing one atom of silicon are termed
monosilicates ; those from acids with two or three atoms of silicon respec-
tively, disilicates and trisilirates.

Thus, the mineral peridote is a monosilicate, Mg2SiC>4.

Serpentine is a disilicate, Mg3Si2O7, and

Felspar, or orthoclase, is a trisilicate, Al2K2(Si3O8)2.

TIN.

Symbol, Sn. Atomic weight =119.

Occurrence. — Tin does not occur in nature in the uncombined
state.* It is met with chiefly as the oxide SnO2 in the mineral
tin-stone or cassiteritejt which is found in immense deposits,
although in comparatively few localities. It is usually associated
with arsenical ores, copper pyrites, wolfram (a tungstate of iron
and manganese), and other minerals. Occasionally it is met with
in nodules of nearly pure oxide, known as stream-tin.

Mode Of Formation.— Tin is obtained exclusively from tin-
stonej' and the process with ordinary ore consists of three opera-
tions, namely — (i) calcining, (2) washing, (3) reducing or smelting.
If the ore be nearly pure tin-stone it may be at once smelted.

The finely crushed ore, after being washed from earthy matters,
is calcined in a reverberatory furnace. The sulphur and arsenic
pass away as sulphur dioxide and arsenious oxide, and are led into
condensing flues, where the arsenic deposits and is collected. The
iron and copper are oxidised to oxide and sulphate. This calcina-
tion is sometimes conducted in the revolving calciner, shown on
page 486. The calcined ore is next washed, whereby copper
sulphate is dissolved, and the iron oxide and other light matters
are removed. The purified ore is then mixed with powdered
anthracite and smelted in a reverberatory furnace —

* Metallic tin has been found in Bolivia, but its origin, whether natural or
artificial, is doubtful.

f Cassiterides, the ancient name for the British Isles, is derived from the
fact that tin-stone was found in large quantities in Devonshire and Cornwall.

638 Inorganic Chemistry

The metal so obtained is purified by first heating it upon the
hearth of a similar furnace until the more readily fusible tin melts
and flows away from the associated alloys ; and afterwards by
stirring into the molten tin so separated billets of green wood,
which results in the separation of a scum or dross carrying with it
the impurities.

Properties. — Tin is a bright white metal, which retains its
lustre unimpaired in the air. It is sufficiently soft to be cut with a
knife, but is harder than lead, although less hard than zinc. At
ordinary temperatures it is readily beaten out into leaf (known as
tinfoil), and may be drawn into wire ; but at temperatures a little
below its melting-point (228°) it becomes brittle and may be
powdered. Tin may be obtained in the form of crystals by melt-
ing a quantity of the metal in a crucible, and when partially
solidified pouring out the remaining liquid portion. Its crystalline
character is also seen by pouring over the surface of a block of
cast tin or a sheet of ordinary tinned iron a quantity of warm
dilute aqua regia, when the surface of the metal will immediately
exhibit a beautiful crystalline appearance.

When a bar of tin is bent it emits a faint crackling sound, and
if quickly bent backwards and forwards two or three times the
metal becomes perceptibly hot at the point of flexure. These
phenomena are due to the friction of the crystalline particles.
Ordinary tin has a specific gravity about 7.2 ; but if the metal be
exposed to the prolonged influence of very low temperatures, it
loses its crystalline character and appears of a grey colour. In
this condition its specific gravity is 5.8 ; and it is believed to be
an allotropic modification of the element. When strongly heated
tin takes fire and burns, forming stannic oxide, SnO2. It is
oxidised by both sulphuric and nitric acids ; thus, when heated
with strong sulphuric acid, stannous sulphate and sulphur dioxide
are produced —

Sn + 2H2SO4 = SnSO4+ SO2+2H2O.

The strongest nitric acid (specific gravity, 1.5) is without action
upon tin. Ordinary concentrated nitric acid (specific gravity, 1.24)
attacks it with violence, forming metastannic acid (page 640), while
in cold dilute acid it slowly dissolves with the production of stannous
nitrate — ^-— £

4Sn + 9HNO8=4Sn(NO8)a + 3H2O + NHj.

Stannous Oxide 639

The ammonia unites with another portion of nitric acid, forming
ammonium nitrate. Strong hydrochloric acid converts it into
stannous chloride, with evolution of hydrogen.

Tin is extensively employed in the process of tinning, which
consists in coating other metals with a thiri film of tin by dipping
into a bath of the molten metal. Ordinary tin-plate (or in common
parlance, " tin," the material of which articles generally called
" tins " are made) is thin sheet-iron which has been thus super-
ficially coated with tin.

Alloys of Tin. — Tin enters into the composition of a large
number of useful alloys. With lead, tin will mix in all proportions,
and many alloys are in use consisting of these two metals. They
are all white, and melt at temperatures lower than that of either
constituent.

Pewter contains 3 parts of tin to i part of lead. Common
solder consists of I part tin and I part lead, while coarse and fine
solder contain half and twice this proportion of tin respectively.
With copper, the most important alloys are the various brasses
and bronzes. Britannia metal contains tin 84 parts, antimony 10
parts, copper 4 parts, and bismuth 2 parts. Tin is a constituent
also of the so-called fusible alloys (see Bismuth, page 500).

Oxides Of Tin. — Two oxides are definitely known, namely,
stannous oxide, SnO, and stannic oxide, SnO2. The monoxide is
a base, yielding the stannous salts j the dioxide is both a basic and
an acidic oxide.

Stannous Oxide, SnO, is obtained by heating stannous oxalate
out of contact with air, thus —

SnC2O4= SnO + CO2 + CO.

When sodium carbonate and stannous chloride are mixed, carbon
dioxide is evolved, and the white hydrated oxide is precipitated,
thus—

When this hydrated oxide is boiled with insufficient caustic
alkali to dissolve it, the undissolved portion is dehydrated and
converte.d into the black monoxide.

When heated in the air, stannous oxide becomes incandescent,
burning to the dioxide. It is soluble in acids, forming stannous
salts, The solution of stannous oxide in sodium hydroxide is

640 Inorganic Chemistry

used by the calico-printer, and is known commercially as sodium
stannite.

Stannic Oxide, SnO2 (tin dioxide), is the chief ore of tin. It is
formed where the metal is burnt in the air, but is most readily pre-
pared by igniting metastannic acid.

It is a white amorphous powder, which changes to yellow and
brown on heating, but returns to its original condition on cooling.
When strongly heated in a stream of gaseous hydrochloric acid, it
may be obtained in small crystals, identical with the natural com-
pound. Stannic oxide is unacted upon by acids or alkalies, but
in contact with fused potassium hydroxide it is converted into
potassium starinate.

Stannic Acid, H2SnO3, or SnO2,H2O, is obtained in a hydrated
condition, as a white gelatinous precipitate, when calcium car-
bonate is added to stannic chloride in insufficient quantity for
complete precipitation. When the precipitate is "dried in vacuo,
it has the composition H2SnO3. The equation representing its
formation may be expressed thus —

2CaCO3 + SnCl4+ H2O = 2CaCl2 + 2CO2 + H2SnO3.

Stannic acid forms a number of salts, of which sodium and
potassium stannates are the most important — the former being
extensively employed as a mordant in dyeing, under the name of
preparing salt. The salts have the composition Na2SnO3,3H2O,
and K2SnO3,3H2O respectively, and are both soluble in water.

Metastannie Acid, H10Sn6O16, is obtained as a white amorphous
powder when tin is acted upon by strong nitric acid ; the reaction
may be represented thus —

5Sn + 20HNO3=H10Sn6O16 + 5H2O + 20NO2.

The composition of the compound depends upon the particular
temperature at which it is dried. This acid is sometimes regarded
as a polymer of stannic acid, which may be expressed by the
formula 5(H2SnO3) ; metastannic acid, however, appears to be
dibasic, forming salts in which two only of the hydrogen atoms
are replaced ; its 'composition may therefore be conveniently ex-
pressed thus —

HaSnO3,4SnO2,4H2O, or H,Sn6Ou,4H,Or

Stannous Chloride 641

Potassium and sodium metastannates are the best known salts,
their formulas being —

K2SnO3,4SnO2,4H2O, and Na2SnO3,4SnO2,4H2O.

Stannous Chloride, SnCl2, is obtained by dissolving tin in
hydrochloric acid, and evaporating the solution, when monosym-
metric prisms separate out, having the composition SnCl2,2H2O.
When dried in vacuo they become anhydrous. The anhydrous
chloride is directly obtained when tin filings and mercuric chloride
are heated together —

HgCl2 + Sn = SnCl2 + Hg.

The reduced mercury volatilises and leaves the chloride, which
at a higher temperature may be distilled.

Stannous chloride dissolves in a small quantity of water, but
with an excess of water, or on exposure to the air, an oxychloride
(or basic chloride) is precipitated, with simultaneous elimination of
hydrochloric acid, thus —

2SnCl2 + 2H2O = SnCl2,SnO,H2O + 2HCl.

The composition of this oxychloride may also be expressed by
either of the following formulas —

Sn2OCl2,H2O, or 2Sn(OH)Cl, or 2(SnO,HCl).

Stannous chloride is a powerful reducing agent, as it readily
combines with either oxygen or chlorine ; thus, when added to a
solution of mercuric chloride, the latter is first reduced to mer-
curous chloride, which, on being gently warmed, is reduced to
metallic mercury —

2HgCl2 + SnCla = Hg2Cl2 + SnCl4.
Hg2Cl2 + SnCl2=2Hg+ SnCl4.

By the absorption of oxygen, the above oxychloride and stannic
chloride are formed, thus—

3SnCl2 + O + H2O = SnCl2,SnO,H2O + SnCl4.

Stannous chloride boils at a temperature about 606°. The
density of the vapour only agrees with the formula SnCl2 at tem-
peratures above 900°, at lower temperatures its vapour-density
approaches more nearly to that required by the formula Sn2Cl4.

642 Inorganic Chemistry

Stannic Chloride, SnCl4, is obtained by passing a stream of
dry chlorine over melted tin in a glass retort, or by heating a
mixture of powdered tin with an excess of mercuric chloride, when
the anhydrous chloride distils over as a colourless, mobile, fuming
liquid, which boils at 113.9°. It unites with water with evolution
of heat, forming hydrated compounds of the composition SnCl4,
3H2O ; SnCl4,5H2O, and SnCl4,8H2O. The compound containing
5HCO is employed as a mordant, and is commercially known as
oxy muriate of tin.

Stannic chloride combines with alkaline chlorides, forming
double chlorides (sometimes called chloro-stannates], such as
SnCl4,2NH4Cl, and SnCl4,2KCl.

Stannous Sulphide, SnS.— When tinfoil is introduced into
sulphur vapour the metal takes fire, and yields a leaden-coloured
mass of stannous sulphide.

In the hydrated condition stannous sulphide is precipitated as
a brown powder when sulphuretted hydrogen is passed through
stannous chloride ; on drying, this becomes black and anhydrous.

Stannous sulphide dissolves in hot concentrated hydrochloric
acid. It is also soluble, in alkaline poly sulphides, forming thio-
stannates, thus —

(i)

(2) SnS2 + K2S

On the addition of hydrochloric acid to the solution; stannic
sulphide is precipitated —

K2SnS3 + 2HCl = 2KCl-f H2S + SnS2.

Stannic Sulphide, SnS2. — This compound cannot be formed
by heating tin and sulphur alone, as the heat of the reaction is
greater than that at which stannic sulphide is resolved into
stannous sulphide and sulphur. It is obtained by heating tin
amalgam, sulphur and ammonium chloride, in a retort. The action
that takes place is a complicated one, various products being
volatilised, and stannic sulphide remaining in the retort as a mass
of golden-yellow scales. Amongst the products expelled during
the process are ammonium chloride, sulphur, mercuric chloride,
mercuric sulphide, and sulphuretted hydrogen. The ammonium
chloride present probably acts by the formation of ammonium
stannous chloride, as an intermediate product, which is then de*

Lead 643

•

composed with the production of stannic sulphide and ammonium
stannic chloride thus —

2SnCl2,2N H4C1 + 2S = SnS2 + N H4C1 + SnCl4,2N H4G1.

Stannic sulphide is a golden yellow crystalline substance which,
when heated, partially sublimes as such, but is for the most part
decomposed into the monosulphide and free sulphur. It is largely
used as a pigment known as mosaic gold.

LEAD.

Symbol, Pb. Atomic weight =207.1.

Occurrence. — Lead has been found in small quantities in the
uncombined state, probably reduced from its ores by volcanic
action.

In combination with sulphur it occurs in enormous quantities in
the mineral galena^ PbS, which is the ore from which the metal is
chiefly obtained. Large quantities are also met with as carbonate
in the mineral cerussite, PbCO3. Other natural compounds are
anglesite, PbSO4 ; lanarkite, PbSO4,PbO ; matlockite, PbCl2,PbO;
pyromorphite, 3Pb3P2O8,PbCl2.

Modes Of Formation.— Lead is very readily reduced from its
compounds, and on this account was one of the earliest known
metals. It was termed by the Romans plumbum nigrum.

Two general processes are made use of for the reduction of lead
from its ores : —

In the first method (known as the reduction process) the lead
sulphide is reduced by double decomposition with lead oxide and
sulphate, which are formed by roasting the ore.

In the second (called the precipitation process) the sulphide is
reduced by metallic iron.

(i.) The galena is introduced into a reverberatory furnace, where
it is partially roasted, whereby a portion of the sulphide is oxidised
to sulphate and oxide —

=PbSO4.
2PbS + 3O2 = 2PbO + 2SO2.

The temperature is then raised, when the oxide and sulphate

644

Inorganic Chemistry

react upon a further portion of the sulphide, with the formation of
metallic lead and the evolution of sulphur dioxide —

This method of lead smelting is followed when the ore is fairly
free from other metallic sulphides. The reverberatory furnace
usually employed (known as the Flintshire furnace) has a con-
siderable depression, or well, in the hearth, where the metallic

FIG. 150.

lead collects during the process, and from which it is drawn off
into a metal pot.

The same process is carried out in the North of England, and
in Scotland, where a very pure lead ore is employed, upon open
shallow hearths (known as the ore hearth^ or Scotch hearth}, built
under a brickwork hood "or chimney in such a manner that the
fumes of lead which escape are caused to pass into condensing
chambers. Fig. 150 shows such a hearth in section. The fire of
peat and coal is urged by "a small blast admitted from behind, and
the ore is added in small quantities at a time. The reduced metal,
sinking to the bottom, runs under the fire-bar and overflows
from the shallow hearth down a channel upon an inclined stone
surface S (called the work-stone\ into an iron pot P, which is gently

Lead 645

heated by a small fire to enable the operator to ladle the metal out
into moulds.

(2.) This method of lead smelting depends upon the fact that at
a high temperature metallic iron, in contact with lead sulphide, is
converted into ferrous sulphide, with separation of lead—

The ores (either in the raw state, or after previous calcination)
are smelted in a blast-furnace with coke and either metallic iron
or such materials as will yield iron under the furnace conditions.
The sulphide of iron, along with other metallic sulphides, rises to
the top of the molten lead as a matt or regulus, while above this a
fusible slag collects, consisting chiefly of silicate of iron.

The lead first obtained by any of these processes usually con-
tains antimony, tin, copper, and other metals. These impurities
are removed by heating the metal in a shallow, flat-bottomed
reverberatory furnace. Most of the admixed metals oxidise before
the lead, and collect in the dross which forms upon the surface.
This process is known as the softening of lead. The silver, how-
ever, which is always present, is not removed by this operation,
but is extracted by one of the methods for desilverising lead
described under silver, page 560.

Properties. — Lead is a soft, bluish- white metal, which when
freshly cut exhibits a bright metallic lustre. On exposure to the
air its bright surface becomes quickly covered with a film of oxide.
Lead is sufficiently soft to be scratched with the finger nail, and
it leaves a black streak when drawn across paper. It cannot be
hammered into foil or drawn into wire, but may readily be obtained
in these forms by rolling and pressing. When a quantity of
melted lead is allowed partially to resolidify, and the still liquid
portion poured off, the metal is obtained in the form of octahedral
crystals belonging to the regular system. Its crystalline nature is
also readily seen by submitting a solution of a lead salt to electro-
lysis, when the metal is deposited upon the negative electrode in
beautiful arborescent crystals with a brilliant metallic lustre (Fig.
151). It is deposited in a similar form, known as the lead tree ', by
suspending a strip of zinc in such a solution. The specific gravity
of lead is 11.3 ; it melts at 330° to 335°, and becomes covered with
a black film of the suboxide, Pb2O : when more strongly heated
it is oxidised to the monoxide, PbO.

Lead is rapidly dissolved by nitric acid, but hydrochloric and

646 Inorganic Chemistry

sulphuric acids are almost without action upon it in the cold. Hot
concentrated hydrochloric acid, however, slowly converts it into
lead chloride.

Lead is unacted upon by pure water in the absence of air ; but
in contact with air lead hydroxide is formed, which is slightly
soluble in water. By the action of atmospheric carbon dioxide upon
this solution, a basic carbonate is precipitated, having the com-
position 2PbCO3,Pb(HO)2. The solvent action of water upon lead
is greatly influenced by the presence of various dissolved sub-
stances in the water ; thus, water containing small quantities of

FIG. 151.

ammoniacal salts, notably the nitrate, dissolves lead much more
rapidly, and the same is the case with water charged with carbon
dioxide under pressure. In the latter case the action is probably
due to the formation of a soluble acid carbonate.

Water, on the other hand, containing small quantities of phos-
phates and carbonates, especially the acid calcium carbonate, are
almost entirely without action upon lead. Certain drinking waters
(such as the Loch Katrine water), which on account of their purity
exert a solvent action upon the lead pipes through which they are

Plumbic Oxide 647

conveyed, are rendered incapable of acting upon the lead by being
first filtered through chalk or animal charcoal, which enables them
to take up sufficient calcium carbonate or phosphate to prevent
this action.

On account of the exhaustive methods of desilverisation to which
the lead is subjected, commercial lead possesses a degree of purity
not found in any other metal as commonly met with ; the total
amount of foreign metals present in ordinary commercial lead
ranges from o. i to 0.006 per cent.

Lead is put to a large number of uses in the arts, on account of
the ease with which it can be worked, and its power of resisting
the action of water and many acids. In the manufacture of lead
pipes advantage is taken of the extreme softness of the metal and
the readiness with which it can be pressed into shape ; the lead,
in a pasty or semi-molten condition, being merely squeezed, or
squirted, through a steel die by hydraulic pressure.

Lead bullets are also made by squeezing the metal-into moulds ;
for as lead contracts on solidification, bullets made by casting
always contain a small cavity, which (unless it happens to form
exactly at the point of centre of gravity) renders the flight of the
bullet untrue.

Oxides of Lead. — Five oxides of lead are known, kaving the
composition Pb2O, PbO, Pb2O3, Pb3O4, PbO2.

Lead Suboxide (plumbous oxide}, Pb2O, is the black compound
which is formed when lead is heated to its melting-point. It is
obtained by heating plumbic oxalate to about 300° in a glass tube
or retort —

2PbC204 = CO + 3CO2+ Pb2O.

. jf* * - :

When heated in the air it burns, forming plumbic oxide ; in the
absence of air it is decomposed into the same oxide and metallic
lead, the reactions being —

Pb2O + O = 2PbO.
Pb2O = Pb + PbO.

In contact with acids it decomposes in the same manner, lead
being deposited, and the plumbic oxide dissolving in the acid to
form a plumbic salt.

Plumbic Oxide (lead monoxide, litharge, massicot}, PbO, is
formed when lead is strongly heated in the air, and is obtained in
large quantities in the cupellation of argentiferous lead. It may

648 Inorganic Chemistry

be obtained by heating lead nitrate or carbonate, and it is produced
when any of the other oxides are heated.

Plumbic oxide is a yellowish powder, known commercially as
massicot, which, when melted and resolidified, is obtained as a
crystalline mass, known as litharge. Plumbic oxide is very slightly
soluble in water, I part dissolving in 7000 parts of water : this
solution is alkaline, and on exposure to the air absorbs carbon
dioxide, forming an insoluble basic carbonate. Plumbic oxide is
dissolved by acids, with formation of the salts of lead ; it also
dissolves in warm potassium or sodium hydroxide.

This oxide forms two hydrated compounds, having the com-
position 2PbO,H2O and 3PbO,H2O. The former is obtained as a
white precipitate when ammonia is added to a solution of lead
acetate ; the second, by the action of ammonia on basic lead
acetate at 25°.

Lead Sesquioxide, Pb2O3, is obtained as an orange-coloured
precipitate by adding sodium hypochlorite to a solution of plumbic
oxide in potassium hydroxide. Heat decomposes it into oxygen
and plumbic oxide. Acids convert it into the monoxide and
dioxide, the former dissolving and yielding a salt of lead. This
oxide may be regarded as a compound of two oxides, PbO,PbO2.

Triplumbie Tetroxide (red lead, minium\ Pb3O4, is obtained
when lead carbonate, or monoxide, is subjected to prolonged
heating in contact with air, at a temperature not above 450°. At
higher temperatures it again gives up oxygen. It is a scarlet
crystalline powder, varying somewhat in colour, according to its
mode of preparation. Dilute acids convert it into PbO2 and
2PbO, the latter oxide dissolving to yield lead salts. With strong
hydrochloric acid and sulphuric acid the molecule of lead
dioxide is acted upon, with evolution of chlorine and oxyger
respectively—

Pb3O4 + 8H Cl = 4H2O + 3PbCl2 + C12.

When red lead is added in small quantities at a time to hot
glacial acetic acid, it dissolves entirely in the acid, and the liquid
contains lead tetracetate— a salt which is of interest as being one of
the few salts known containing tetravalent lead (see page 628). On
cooling, the tetracetate separates out as pale greenish- white needles.
The salt is immediately decomposed by water ; if therefore the

Plumbic Chloride 649

acid solution of it be poured into water, a brown precipitate of
lead peroxide is thrown down, and acetic acid regenerated —

Red lead* is employed as a pigment, and also in the manu-
facture of flint glass.

Plumbic Peroxide (lead dioxide}, PbO2, may be obtained by
the action of dilute nitric acid upon red lead —

Pb3O4 (or PbO2,2PbO) + 4HNO3 = PbO2 + 2Pb(NO3)2 + 2H2O.

Or it may be prepared by the action of oxidising agents upon
the monoxide. Thus, when chlorine is passed through an alkaline
solution, in which the monoxide is suspended, or when bleaching-
powder is added to a solution of lead acetate, the dioxide is
produced.

The dark-brown deposit which forms upon the positive electrode
when a solution of a lead salt is electrolysed, also consists of the
dioxide.

Plumbic peroxide is a brown or puce-coloured powder. It is a
powerful oxidising substance, and when gently rubbed with flowers
of sulphur in a warm mortar the mass suddenly inflames. When
a stream of sulphur dioxide is passed over the peroxide in a tube,
the two compounds unite to form lead sulphate, the mass becom-
ing incandescent. Nitric acid is without action upon it, but
hydrochloric and sulphuric acids act upon it in the same manner
as upon red lead. When strongly heated the peroxide gives up
oxygen, and is converted into the monoxide.

When plumbic peroxide is boiled with strong aqueous potassium
hydroxide it dissolves, and the solution deposits crystals of potas-
sium plumbate, K2PbO3,3H2O. This compound corresponds with
potassium stannate, K2SnO3,3H2O, and its existence shows that
lead possesses, although to a very feeble extent, the acidic properties
exhibited by the other members of the same family of elements.

Plumbic Chloride (lead dichloride), PbCl2, is obtained as a
white curdy precipitate when hydrochloric acid, or a soluble
chloride, is added to a solution of a lead salt. It is also produced

* Commercial red lead varies considerably in composition, and although it
has been shown that a definite compound exists of the composition Pb3O4
(which may also be expressed by the formula 2PbO,PbO2), it is still uncer-
tain whether there are not other compounds consisting of these two oxides
united in different proportions.

650 Inorganic Chemistry

by the action of boiling hydrochloric acid upon lead in the pre-
sence of air- It is best prepared by dissolving lead oxide or
carbonate in hot hydrochloric acid, when the lead chloride sepa-
rates out on cooling in long white, lustrous, needle-shaped crystals
belonging to the rhombic system. Lead chloride is soluble in
boiling water to the extent of about 4 parts in 100 parts of water.
On cooling the solution the greater part of the salt separates out,
and at o° the liquid contains 0.8 part in solution. The presence
of hydrochloric acid and soluble chlorides diminishes the solu-
bility of lead chloride.

When heated in contact with air it is converted into an oxy-
chloride, of the composition Pb2OCl2, or PbCl2,PbO, corresponding
with the natural compound matlockite. This compound, in the
hydra ted condition, Pb2OCl2,H2O, is prepared on a large scale by
the addition of lime-water to a solution of lead chloride, and is
employed as a white pigment, known as Pattinsorfs white lead.

Cassel yellow is an oxychloride of lead of the composition
PbCl2,7PbO, obtained by heating lead oxide and ammonium
chloride.

Lead Tetraehloride (lead perchloride), PbCl4. — When plumbic
peroxide is dissolved in cold concentrated hydrochloric acid, a
yellow liquid is obtained, which, on warming, yields chlorine,
with precipitation of lead dichloride. This liquid contains the
tetrachloride of lead in solution.

When lead dichloride is suspended in hydrochloric acid, and
chlorine is passed through the mixture, a solution of lead tetra-
chloride is obtained ; and on the addition of ammonium chloride,
ammonium plumbic chloride, PbCl4,2NH4Cl (corresponding to
ammonium stannic chloride), separates out. When this compound
is acted upon with strong sulphuric acid, in the cold, lead tetra-
chloride separates out as a yellow oily liquid.

Lead tetrachloride is a yellow, highly-refracting, fuming liquid,
which decomposes in contact with moisture into lead dichloride
and chlorine. It may be preserved beneath concentrated sul-
phuric acid. With small quantities of water it forms a hydrated
compound, but excess of water decomposes it into hydrochloric
acid and lead peroxide —

PbCl4 + 2H2O = PbO2 + 4H Cl.

When heated with strong sulphuric acid to about 105°, it suddenly
decomposes with explosion.

Lead Carbonate 651

Lead Nitrate, Pb(NO3)2, is obtained by dissolving litharge in
nitric acid. The salt is deposited from the solution in the form of
regular octahedral crystals. It is soluble in water to the extent of
50 parts in 100 parts of water at the ordinary temperature. When
heated it evolves nitrogen peroxide and oxygen, leaving plumbic
oxide (page 242).

On boiling an aqueous solution of lead nitrate with lead oxide,
the latter dissolves, and the solution on cooling deposits crystals
of a basic nitrate, Pb(NO3)HO or Pb(NO3)2,PbO,H2O. By the
addition of ammonia to a solution of lead nitrate, other basic
nitrates are obtained, which may be regarded as consisting of
compounds of Pb(NO3)HO with PbO, or of Pb(NO3)2 with PbO
and H2O in varying proportions.

Lead Carbonate, PbCO3, is obtained as a white crystalline
powder by the addition
of ammonium sesquicar-
bonate to a solution of
lead nitrate. It occurs in
the form of transparent
rhombic crystals in the
mineral cerussite, isomor-
phous with arragonite.
Lead carbonate is almost
insoluble in water, but is
appreciably dissolved in
water charged with carbon pIGi Ie2.

dioxide. When sodium or

potassium carbonate is added to a solution of lead nitrate, basic
carbonates of lead are precipitated, varying in composition with the
conditions of temperature. The most important of the basic car-
bonates is white lead, 2PbCO3,Pb(HO)2. This compound is manu-
factured on a large scale by several processes for use as a pigment.
The oldest process, and that which yields the best product, is known
as the Dutch method. It depends upon the action of acetic acid
upon metallic lead, in the presence of moist air and carbon
dioxide. The lead, cast into rough gratings in order to expose
a large surface, is placed in earthenware pots, as shown in Fig.
152. A small quantity of dilute acetic acid (in the old Dutch
process, vinegat*) is placed in the pots, and the gratings of lead,
which rest upon the shoulder of the pot, are piled one upon the
other. These pots are then placed upon a thick bed of spent tan-

652 Inorganic Chemistry

bark (in the original method, dung), upon the floor of a shed,
and covered with planks. Upon these another layer of tan-
bark is spread, and a second row of pots similarly charged. In
this manner the layers of pots are built up to the roof of
the shed, and the whole allowed to remain for about three
months. Such a stack will contain many tons of lead, and
about 65 gallons of dilute acetic acid to the ton of metal.
The acid is gradually vaporised by the heat developed by the
fermenting tan-bark, which results first in the formation of a
basic lead acetate —

2H(C2H302) + 2Pb + 02 = Pb(C2H302)2,Pb(HO)2.

This basic acetate is then acted upon by the carbon dioxide
evolved during the fermentation, with the production of a mixture
of normal lead acetate and basic lead carbonate, thus —

3{Pb(C2H302)2,Pb(HO)2}+2C02=3Pb(C2H302)2+2PbC03,Pb(HO)2 + 2H20.

And the lead acetate, in the presence of air and moisture, reacts
upon a further portion of the metal, regenerating the basic acetate,
which is once more decomposed by carbon dioxide —

Pb(C2H30.,)2 + Pb + O + H20 = {Pb(C2H302)2,Pb(HO)2}.

In this cycle of reactions, therefore, the acetic acid acts as a
carrier, a comparatively small quantity being able to convert an
indefinite amount of lead into white lead.

White lead is also prepared by passing carbon dioxide into a
solution of the basic acetate, obtained by boiling plumbic oxide
(litharge) with lead acetate. The product, however, is not so
opaque as that obtained by the former method, and is therefore
not so valuable as a pigment. (This method is known as the
Clichy^ or Thenard's process.)

Milner's process consists in grinding together litharge, sodium
chloride, and water, whereby a mixture of an oxychloride of lead
and sodium hydroxide is formed —

4PbO + 2NaCl + 5H2O = PbCl2,3PbO,4H2O + 2NaHO,

Lead Sulphate 653

and then passing carbon dioxide into the mixture, which converts
it into white lead and sodium chloride, thus —

3[PbCl2,3PbO,4H2OJ + 6NaHO + 8CO2=6NaCl
+ 4[2PbCO3,Pb(HO)2] + llH2O.

White lead is a heavy, amorphous powder, whose value as a
pigment, or body colour, depends upon its opacity and density.
Although this compound labours under the disadvantages of being
extremely poisonous, and of becoming blackened by sulphuretted
hydrogen, no substitute for it has yet been found which possesses
the same "body " or covering power.

Lead Sulphate, PbSO4.— The mineral anglesite, PbSO4, occurs
in the form of rhombic crystals, isomorphous with. strontium and
barium sulphates. Lead sulphate is obtained as a white powder?
by precipitating a lead salt with sulphuric acid or a soluble
sulphate. It is soluble in water only to an extremely slight extent,
and still less in dilute sulphuric acid, but strong sulphuric acid
dissolves it readily. It also dissolves in potassium hydroxide, and
in many ammoniacal salts, notably the acetate, and in sodium
thiosulphate.

An acid sulphate, of the composition PbSO4,H2SO4,H2O, is
obtained by boiling the normal sulphate with sulphuric acid ; and
a basic sulphate, PbSO4,PbO, is formed by the actior\^>f ammonia
upon the normal salt.

Lead Disulphate, Pb(SO4)2.— This substance is obtained by the
electrolysis of sulphuric acid of sp. gr. 1.7 to 1.8 at a temperature
not above 30°, employing an anode of lead. The cell is divided
by a porous partition, and the compound collects as a muddy
deposit in the anode compartment. Lead disulphate is a cry-
stalline substance having a faint greenish colour. It is immediately
decomposed by water into lead peroxide and sulphuric acid —

Sulphuric acid of a sp. gr. less than 1.65 decomposes it in a
similar manner, but in concentrated sulphuric acid it is slightly
soluble, 100 c.c. acid at 30° dissolving 0.345 gram of the com-

654 Inorganic Chemistry

pound. Concentrated hydrochloric acid and glacial acetic acid
convert it respectively into lead tetrachloride, PbCl4, and lead
tetracetate, Pb(C2H3O2)4. Each of these compounds, like the
disulphate, is decomposed by water into lead peroxide and
the respective acid. Double salts, such as K2Pb(SO4)3 and
(NH4)2Pb(SO4)3, have been prepared, which are more stable
than the disulphate itself.

Lead Sulphide, PbS.— The natural sulphide, galena, is found
in the form of cubical crystals, possessing very much the colour
and the metallic lustre of freshly cut lead. It is artificially formed
when lead is heated in sulphur vapour, or when sulphuretted
hydrogen is passed through a solution of a lead salt.

When heated in vacuo, or in a stream of an inert gas, lead
sulphide melts, and sublimes in the form of small cubes. When
heated with free access of air it is converted into lead sulphate.

Boiling dilute nitric acid converts lead sulphide into the nitrate,
with separation of sulphur; but strong nitric acid oxidises it into
lead sulphate. It is decomposed by hot concentrated hydrochloric
acid, with evolution of sulphuretted hydrogen.

When sulphuretted hydrogen is passed into a solution of lead
chloride, the precipitate which forms is first yellow, then reddish-
brown, and finally black ; the yellow and red precipitates are com-
pounds of lead chloride and lead sulphide, termed sulphochlorides,
having the composition, PbS,PbCl2, and 3PbS,PbCl2.

The compounds of lead are powerful poisons, and when con-
tinuously taken into the system in small quantities, they act as
cumulative poisons. Painters and others who constantly handle
white lead are liable to suffer from chronic lead poisoning.

CHAPTER X
ELEMENTS OF GROUP V. (FAMILY A.)

Vanadium, ¥=51.2; Columbium, Cb=93-5; Tantalum, Ta=i8r.

THE three rare metals comprising this family are closely related to each
other, and also to the elements of family B of the same group, namely, the
nitrogen and phosphorus series.

Vanadium occurs in a few rare minerals, as vanadite, 3Pb3(VO4)2,PbCl2
(the vanadium analogue of pyromorphite) ; fucherite, BiVO4 ; mottramite,
(PbCu)3(VO4)3,2(PbCu)(HO)2. Small quantities also occur in certain iron
ores, the vanadium ultimately finding its way into the Bessemer slag, in
which it has been found concentrated to the extent of i. 5 per cent.

Metallic vanadium was first isolated by Roscoe (1867), although its existence
was previously discovered by Del Rio (1801). The metal is extremely difficult
to obtain, as at a red heat it combines with oxygen with great readiness,
yielding the pentoxide V2O5, and also with nitrogen, forming the nitride VN.
The element is prepared by heating the dichloride in a stream of perfectly
pure hydrogen —

VCl2+H2=2HCl + V.

Vanadium is unacted upon by air at ordinary temperatures, but when
heated burns brilliantly to the pentoxide.

Columbium and tantalum are found associated together in the rare mineral
tantalite or columbite. The first to be discovered was tantalum (Hatchett,
1801), and was originally named columbium; and the name niobium (from
Niobe, the daughter of Tantalus) was given to the allied element by Rose
(1846). More recently, however, the name of this second element has been
changed to columbium, although it may still be met with under its original
name of niobium. Columbium is obtained by heating the trichloride, CbCl3,
in a stream of hydrogen.

Vanadium forms five oxides, corresponding to the oxides of nitrogen, while
three oxides of columbium and two of tantalum are known : —

V2O ; V2O2(orVO)
CbO

V203 ; V204(or VO2) ;

'5-

— ; TaO2 ; Ta2O5.
The pentoxides are obtained when the metals are burned in air or oxygen.
They give rise respectively to vanadates, columbates, and tantalates, cor-
responding to nitrates and metaphosphates, thus —

Sodium nitrate, NaNO3. Sodium metacolumbate, NaCbO3.

Sodium metaphosphate, NaPO3. Sodium metatantalate, NaTaO3.
Sodium metavanadate, NaVO3.

655

656 Inorganic Chemistry

The closer relation of these elements to phosphorus than to nitrogen is seen
in the formation of salts derived from ortho- and pyro-acids, corresponding
to orthophosphates and pyrophosphates. The naturally occurring vanadium
compounds • above mentioned are vanadates derived from the hypothetical
orthovanadic acid, H3VO4. Both metavanadic acid, HVO3, and pyrovanadic
acid, H4V2O7, have been obtained. Unlike the phosphorus compounds, the
metavanadates are the most stable of the three classes of salts, and the
orthovanadates the least stable. The most important of these salts is the
ammonium metavanadate, NH4VO3, which is prepared by dissolving the
pentoxide in ammonia. This salt is insoluble in ammonium chloride, and
use is made of this property in the preparation of vanadium compounds
from the mineral mottramite. When ammonium metavanadate is ignited,
vanadium pentoxide is obtained —

2NH4V03= V2O5 + 2NH3 + H2O.

Vanadium acts also as a feeble base. Thus, when the tetroxide, or hypo-
vanadic oxide, is dissolved in sulphuric acid, hypovanadic sulphate, V2O2(SO4)2t
is formed. The solution of this salt possesses a rich blue colour.

Vanadium forms three chlorides, having the composition —

VCL2 (or V2C14) ; VC13 (or V2C16) ; VC14.

Columbium gives a trichloride, CbCl3, and pentachloride, CbCl5, while only
the pentachloride of tantalum is known, TaCl5.

Vanadium forms a number of compounds with oxygen and chlorine. Thus,
when vanadium tetrachloride is acted upon by water, it yields hypovanadic
chloride, V2O4Cl2, which dissolves in the water, giving a blue solution.

Vanadium oxychloride, or vanadyl trichloride, VOC13, corresponds to phos-
phorus oxychloride, POC13. From vanadyl trichloride, by treatment with
zinc, vanadyl dichloride is obtained, VOC12, and by the action of hydrogen at
a high temperature upon this, both vanadyl monochloride, VOC1, and divanadyl-
monochloride, V3O2C1, are formed.

CHAPTER XI
ELEMENTS OF GROUP VI. (FAMILY A.}

Chromium, Cr 52.1 Tungsten, W 184

Molybdenum, Mo . . .96 Uranium, U 239-S

CHROMIUM.

Symbol, Cr. Atomic weight = 52. i.

Occurrence. — Chromium does not occur in nature in the un-
combined state. In combination with oxygen and associated
with iron it is met with in considerable quantities in the mineral
chrome iron ore, or chromite, Cr2O3,FeO. This ore is the chief
source of chromium compounds. Other natural compounds are
crocoisite, PbCrO4, and chrome-ochre, Cr2O3. Traces of chromium
are present in various minerals, such as the emerald and green
serpentine, and impart to them their green colour.

Modes of Formation. — Until quite recently metallic chromium
was a mere chemical curiosity. It may be obtained by the re-
duction of the oxide, Cr2O3, by means of carbon at the high tempera-
ture of the electric furnace. The metal so produced, however,
always contains carbon.

It is now produced on a manufacturing scale, by reducing the
oxide by means of metallic aluminium. The powdered oxide
mixed with the requisite quantity of powdered aluminium is placed
in a refractory crucible, and the mixture ignited by means of a fuse.

The ignition temperature of this mixture being very high, the
most suitable fuse for the purpose consists of a mixture of barium
peroxide and powdered aluminium. A small quantity of this
mixture is placed in a depression made in the surface of the charge
in the crucible, and a piece of magnesium ribbon inserted into it.
When the magnesium is ignited it immediately fires the fuse, which
in its turn communicates its combustion to the charge. The con-
tents of the crucible undergo rapid vivid combustion, and the tem-
perature of the entire mass rises sufficiently high to melt the reduced
chromium.

657

658 Inorganic Chemistry

Properties. — Chromium is a hard, steel-grey metal, which is
not oxidised in dry air. Its melting-point is about 2000°, being
somewhat higher than that of platinum. The metal has no mag-
netic properties. It dissolves in dilute hydrochloric and sulphuric
acids, with evolution of hydrogen. When placed in nitric acid
chromium assumes the so-called passive condition, and while in
this state it is unacted upon by the acids which dissolve it under
normal conditions.

Metallic chromium when added to steel imparts to the latter a
high degree of hardness and tenacity, and it is now largely em-
ployed in the production of these "chrome steels," which contain
from 0.4 up to as much as 2 or 3 per cent, of chromium.

Oxides of Chromium. — Two oxides of chromium are definitely
known, namely —

Chromium sesquioxide (chromic oxide) . . Cr2O3.
Chromium trioxide (chromic anhydride] . . CrO3.

The first is a basic, and the second an acidic oxide. Besides
these two compounds, a hydrated oxide, derived from the unknown
chromous oxide, also exists, having the composition CrO,H2O, or
Cr(HO)2. It is obtained as a yellowish precipitate by adding potas-
sium hydroxide to a solution of chromium dichloride (chromous
chloride), with the exclusion of air. It rapidly absorbs oxygen,
turning dark brown. When heated out of contact with air it is
converted into the sesquioxide, with evolution of hydrogen —

2CrO,H2O = Cr2O3 + H2O + H2.

Other compounds of chromium and oxygen are described, whose composi-
tion, however, is not definitely established ; thus, the product obtained as a
brown powder, either by the partial reduction of the trioxide or the oxidation
of the sesquioxide, is regarded by some chemists as chromium dioxide, CrO2,
and by others as chromium chromate, Cr2O3,CrO;}. It is readily obtained by
passing nitric oxide into a solution of potassium dichromate.

Chromium Sesquioxide, Cr2O3, is obtained as a grey-green
powder, when either the hydroxide, or the trioxide, or ammonium
dichromate is ignited (see page 230).

When the vapour of chromyl dichloride, CrO2Cl2, is passed
through a red-hot tube, chromic oxide is deposited in the form of
daik-green hexagonal crystals. Chromic oxide which has been
strongly ignited is nearly insoluble in a-cids. It is used under the
name of clirome green as a pigment, and for giving a green colour
to glass.

Chromic Hydroxides.— Chromic oxide yields a number of

Chromium Trioxide 659

hydrated compounds. When ammonia is added to a solution of
chromic chloride, or other chromic salt, free from alkali, a light
blue compound is precipitated, which, when dried over sulphuric
acid, has the composition Cr(HO)3,2H2O (or Cr2O3,7H2O). When
this is dried in vacuo it loses water, and becomes 2[Cr(HO)3],H2O
(or Cr2O3,4H2O) ; and on being heated at 200°, it again parts with
water, and has the composition CrO(HO) (or Cr2O3,H2O).

When potassium dichromate and boric acid are heated to dull
redness, and the mass treated with water, a rich green residue
is obtained, having the composition Cr2O(HO)4(or Cr2O3,2H2O).
This compound, known as Guignefs green ^ is employed as a
pigment.

The first two of these compounds, which may be looked upon
as consisting of the hydroxide Cr(HO)3 in a hydrated condition,
are readily soluble in acids, yielding the chromic salts.

Chromium Trioxide (chromic anhydride] CrO3.— When strong
sulphuric acid is added to a cold saturated solution of potassium
dichromate, the trioxide separates out in long, red, needle-shaped
crystals —

K2Cr2O7 + H2SO4 = K2SO4+ H2O + 2CrO3.

The liquid is decanted from the crystals, which are drained
upon porous tiles, and the adhering sulphuric acid and potassium
sulphate washed away by strong nitric acid. The crystals are
finally heated upon a sand-bath, whereby the nitric acid is
evaporated.

Chromium trioxide dissolves in water to the extent of 62 parts
in loo parts of water at 26°. It melts at a temperature about 192°.
At 250° it begins to give off oxygen, and is ultimately converted
into the sesquioxide —

Chromium trioxide is a powerful oxidising agent, and in contact
with most organic substances it is reduced. In the preparation
of the compound, therefore, the liquid cannot be filtered through
paper in the usual way. Warm alcohol dropped upon the trioxide
at once takes fire, while in a more diluted condition it is oxidised
to acetic acid ; and the reduction of the chromium trioxide is made
evident by the change of colour of the liquid, from red or yellow
to. olive green.

Gaseous ammonia reduces the trioxide to the sesquioxide, with
formation of water and nitrogen —

660 Inorganic Chemistry

the reaction being accompanied with the evolution of so much heat
that the chromic oxide produced becomes incandescent.

When hydrogen peroxide is added to a dilute solution of
chromium trioxide, or to a dilute solution of potassium dichromate,
acidified with sulphuric acid, a deep indigo-blue solution is ob-
tained.. This blue compound is believed to contain perchromic
acid) but its composition has not been definitely established. It
may be regarded as a compound of chromium trioxide, CrO3, or of
perchromic acid, HCrO4, with hydrogen peroxide, H2O2, in unde-
termined proportions.

In aqueous solution the blue colour quickly disappears, oxygen
being eliminated. The compound is soluble in ether ; and there-
fore, when the aqueous solution is shaken up with that liquid, a
deep blue ethereal solution rises to the top. In this solution the
compound is more stable, but when evaporated it evolves oxygen,
leaving chromium trioxide. It is decomposed by alkalies, forming
alkaline chromates with evolution of oxygen. The formation of
this compound constitutes a delicate test for either chromium
trioxide or hydrogen peroxide (see Hydrogen Peroxide, page 227).

Chromous Compounds. — These correspond to chromous hydrate, Cr(HO)2,
in which the chromium functions as a divalent element. Comparatively few
of these salts are known.

Chromous Chloride, CrCl2, is formed when the metal dissolves in hydro-
chloric acid. It is prepared in the anhydrous state by gently heating chromic
chloride in a current of pure hydrogen. It is a white crystalline compound,
soluble in water to a blue solution, which rapidly absorbs oxygen.

Chromous Sulphate, CrSO4,7H2O, is obtained by dissolving chromous
acetate in dilute sulphuric acid. It is deposited from the solution in blue
crystals, isomorphous with ferrous sulphate, FeSO4,7H2O.

Chromic Compounds. — These are derived from chromic oxide,
the oxide acting as a base.

Chromic Chloride, CrCl3, or Cr2Cl6, is prepared by strongly
heating a mixture of chromic oxide, Cr2O3, and carbon in a stream
of dry chlorine. The chromic chloride sublimes in the form of
scales, having a reddish-pink colour. The molecular weight of
chromic chloride is 158.45, showing that in the vaporous state its
molecules have the formula CrCl3.

It is nearly insoluble in water, but readily dissolves in water
containing minute traces of chromous chloride, forming a green
solution. The same solution is obtained by dissolving hydrated
chromic hydroxide, Cr(HO)3,2H2O, in hydrochloric acid, and if
this solution be slowly evaporated, very soluble green crystals
separate out, having the composition CrCl3,6H2O. If strongly
heated in the air, this compound gives off water and hydrochloric

Chrome Alum 66 1

acid, leaving chromic oxide, Cr2O3 ; but when heated to 250°, in
either gaseous hydrochloric acid or chlorine, it is converted' into
the pink anhydrous chromic chloride, which redissolves in water
to the green solution. If heated strongly and sublimed, the com-
pound obtained is nearly insoluble in water.

Chromic Sulphate, Cr2(SO4)3, is obtained by dissolving chro-
mium hydroxide in concentrated sulphuric acid, when a green
solution is formed, which on standing changes to blue, and slowly
deposits violet-blue crystals. The salt may be purified by dis-
solving in cold water and precipitating with alcohol. If insufficient
alcohol be added to cause immediate precipitation, the salt slowly
deposits from the dilute spirit in blue octahedrons, belonging to
the regular system.

A cold aqueous solution, which has a violet colour, becomes
green when boiled.

Chromic sulphate forms double salts with the sulphates of the
alkalies, which belong to the alums.

Potassium Chromium Alum (chrome alum\ K2SO4,Cr2(SO4)3,
24H2O. — This double sulphate is formed when solutions of potas-
sium and chromium sulphates are mixed together in molecular
proportions. It is most conveniently prepared by the addition of
the requisite amount of sulphuric acid to an aqueous solution of
potassium dichromate, and reducing the chromic oxide by passing
sulphur dioxide through the liquid —

(1) K2Cr2O7 + H2SO4

(2) 2Cr03 + 3S02 =Cr2(S04)3.

The resulting solution, containing the two sulphates in mole-
cular proportions, deposits crystals of the double sulphate, in the
form of dark plum-coloured octahedrons (Fig. 147, B, page 619),
which appear red by transmitted light.

Chrome alum dissolves in water, yielding a plum-coloured solu-
tion, which on boiling turns green, but on long standing returns to
its original colour.

Sodium chromium alum is more soluble, and ammonium chro-
mium alum is less soluble, than the potassium salt.

Cliromites. — Chromic oxide acts also as a weak acid, and combines with
other oxides, forming compounds resembling the aluminates. When potas-
sium hydroxide is added to a solution of a chromic salt, the green hydrated
oxide which is precipitated contains alkali which cannot be removed by hot

662 Inorganic Chemistry

water ; this is present in the form of potassium chromite. The best known
chromites are zinc chromite, Cr2O3,ZnO ; manganous chromite, Cr2O3,MnO,
and ferrous chromite, Cr2O3,FeO; the latter occurs naturally as chrome
iron ore.

Chromates. — When chromium trioxide is dissolved in water,
the solution is believed to contain chromic acid, H2CrO4, or
dichromic acid, H2Cr2O7 ; when the solution is evaporated, how-
ever, the trioxide alone is left. Red crystals have been obtained,
by cooling a hot saturated solution of the trioxide, which are
believed to be chromic acid.

Potassium Chromate, K2CrO4, is prepared by adding potas-
sium hydroxide to a solution of the dichromate —

On evaporation, the yellow chromate of potash separates out in
rhombic crystals, isomorphous with potassium sulphate. It is
soluble in water at the ordinary temperature to the extent of 60
parts in 100 parts of water, forming a yellow solution having an
alkaline reaction.

Potassium Bichromate, K2Cr2O7, is manufactured from chrome
iron ore by roasting the finely crushed ore with potassium car-
bonate and lime in a reverberatory furnace ; the mass being
frequently raked over to expose fresh portions to the oxidising
action of the flames. In this way a mixture of calcium and potas-
sium chromates is produced —

2Cr2O3, FeO + 3KrCO3 + CaO + 7O = CaCrO4 + 3K2CrO4 + Fe2O3 + 3CO2.

The yellow mass, when cold, is broken up and lixiviated with a
hot solution of potassium sulphate, which, by double decomposition
with the calcium chromate, forms potassium chromate and precipi-
tates calcium sulphate. The solution after settling is treated with
the requisite quantity of sulphuric acid to convert the chromate
into the dichromate, thus —

The dichromate being much less soluble than the normal chro-
mate, a large proportion of it at once deposits as the solution cools ;
and the mother-liquor containing potassium sulphate is used again
to lixiviate a fresh quantity of the roasted mixture.

Potassium dichromate forms large red prisms or tables, belong-

Chromyl Chloride 663

ing to the asymmetric (triclinic) system. It is soluble in water at
the ordinary temperature to the extent of 10 parts in 100 parts of
water, yielding an acid solution, which is extremely poisonous.
When a film of gelatine is impregnated with potassium dichromate
and exposed to light, a reduction of the chromium to chromic
oxide takes place, and at the same time the gelatine is rendered
insoluble. This property is utiiised in photographic processes.*

Potassium dichromate is also known under the misnomer bichromate of
potash, which would suggest that the salt was in reality hydrogen potassium
chromate, corresponding to bisulphate of potash, HKSO4. Such a chromium
compound does not exist. The dichromates correspond to the disulphates (or
pyrosulphates), see page 435.

Potassium Trichromate, K2Cr3O10 (or K2CrO4,2CrO3), and Potassium
Tetrachromate, K2Cr4O13 (or K2CrO4l3CrO3), are also known.

Lead Chromate, PbCrO4, is found as tne mineral crocoisite.
It is produced by precipitation from a lead salt, with either
potassium chromate or dichromate. It forms a bright yellow
powder, known as chrome-yellow ', and is employed as a pigment.
It melts without decomposition, and resolidifies on cooling to a
brown crystalline solid. At higher temperatures it gives off
oxygen, and is converted into chromic oxide and a basic lead
chromate. When heated with organic compounds, the latter are
completely oxidised ; lead chromate is therefore employed in
organic analyses.

When lead chromate is digested with sodium hydroxide, or with
normal potassium chromate, a basic lead chromate is obtained
as a rich red powder —

2PbCrO4 + 2NaHO = Na2CrO4 + H2O + Pb2CrO5.

This compound is known as chrome-red.

Chromyl Chloride, CrO2Cl2.— This compound is prepared by
distilling a mixture of potassium dichromate and sodium chloride
with strong sulphuric acid. Chromyl chloride is a deep red, mobile,
strongly fuming liquid. It is decomposed by water into hydro-
chloric acid and chromium trioxide, and acts as a powerful
oxidising substance. When dropped upon phosphorus it explodes.
When heated in sealed tubes it is converted into trichromyl
chloride with loss of chlorine, (£rO2)3Cl2.

Chromyl chloride may be regarded as being derived from

* Abney, " Treatise on Photography."

664 Inorganic Chemistry

chromic acid, CrO2(HO)2, by the complete substitution of
(HO) by Cl. The intermediate compound, chloro-chromic acid,
CrO2(HO)Cl, is -unknown, although its salts have been pre-
pared ; thus, by the gentle action of hydrochloric acid upon
potassium dichromate, potassium chloro-chromate is obtained as
a red crystalline salt —

Molybdenum, Mo=96; Tungsten, W=i84; Uranium, 11=238.5.

These three somewhat rare elements are closely related to chromium.

Molybdenum occurs in the mineral molybdenite, MoS2 (which strongly re-
sembles graphite in appearance), and more rarely as molybdemim ochre, MoO3,
and wulfenite, PbMoO4.

Tungsten is found chiefly in wolfram, 2FeWO4,3MnWO4 (occurring in
the Cornish tin mines) -, more rarely as scheelinite, PbWO4, and wolfram
ochre, WO3.

Uranium occurs as an oxide, UO2,2UO3, in pitchblende (a considerable
quantity of which, associated with other uranium compounds, has recently
been discovered at St. Stephens, Cornwall).

Molybdenum is obtained by the action of hydrogen upon the heated oxide
or chloride ; uranium, by the action of sodium upon the chloride ; while
tungsten has been obtained by both methods. In their specific gravities,
tungsten and uranium exhibit a marked difference from chromium and
molybdenum; thus, Cr, sp. gr. =6; Mo, sp. gr. =8.6; while W, sp. gr. =
19.1 ; U, sp. gr. =18.7.

Molybdenum and uranium form a large number of oxides, some of which
are regarded as definite oxides, while others are looked upon as combinations
of two oxides. Only two oxides of tungsten are known. The composition of
these compounds is as follows —

MoO — —

Mo.2O3

Mo02 WO2 U02 TT _

MnO wn no* U2u5-uu2,uo3.

Mo03 W03 UOs U3o8=Uo2,2U03.

The trioxide of each metal is an acid oxide ; uranium trioxide, however, is
both acidic and basic. They are insoluble in water, but by the action of
alkalies they yield molybdates, tungstates, and uranates. Molybdates and
tungstates, derived from the acids H2MoO4,2H2O and H2WO4,2H2O (corre-
sponding to chromic acid), are known. And all three oxides yield salts
corresponding to potassium dichromate, thus —

Sodium Dimolybdate. Sodium Ditungstate. Sodium Diuranate.

NaaMo2O7

Molybdic and tungstic acids also form numerous polymolybdates and poly-

Molybdenum, Tungsten, Uranium 66$

tungstates, by the absorption of varying quantities of the trioxide into the
molecule of the normal salt (see Chromates, page 663). And in the case of
tungsten, the compound metatungstic acid, HgWjOuiTH^O, is known.

Uranium dioxide and trioxide are both basic oxides, the former yielding the
unstable uranous salts, such as uranous sulphate, U(SO4)2; and the latter
producing the uranyl salts, of which the sulphate, (UO2)SO4, and nitrate,
(UO2)(NO3)2, are well known.

Uranium peroxide, UO4, is an acid oxide which yields per-uranates.

Both molybdic and tungstic acids form complex compounds with phos-
phoric acid, known as phospho-molybdic and phospho-tungstic acids : thus,
when a nitric acid solution of ammonium molybdate (NH4)2MoO4> is added
in excess to a solution of orthophosphoric acid or an orthophosphate, a
canary- yellow crystalline precipitate of ammonium phospho-molybdate,
2(NH4)3PO4,22MoO3,12H2O, is obtained (see page 477). It is soluble in
alkalies and in excess of phosphoric acid, but insoluble in dilute mineral acids.
When this compound is dissolved in aqua-regia the solution deposits yellow
crystals of phospho-molybdic acid, 2H3PO4,22MoO3.

Compounds with chlorine having the following composition are known—

MoCl2 WC12
Mod, or Mo2Cl6

MoCl4 WC14 UC14.

MoCl8 WC15 UClg,

- WC16

CHAPTER XII

GROUP VII. (FAMILY A.)

MANGANESE.

Symbol, Mn. Atomic weight = 54. 93.

Occurrence. — This element is never found in nature in the free
state. It is widely distributed in combination with oxygen, as
pyrolusite^ MnO2 ; braunite, Mn2O3 ; and hausmanniie, Mn3O4.
Also as a hydrated oxide in manganite, Mn2O3,H2O. It is met
with also as carbonate in manganese spar, MnCO3 ; and as sul-
phide in manganese blende, MnS.

Modes Of Formation. — Manganese may be obtained by the
reduction of the oxide by means of carbon at a very high tempera-
ture, as obtained in the electric furnace. The product, however,
contains carbon. In a purer state it may be prepared by the re-
duction of fused anhydrous manganous chloride by means of metallic
magnesium. At the present time, however, it is obtained by reduc-
tion of the oxide by means of aluminium, in a manner precisely
similar to that employed for the manufacture of chromium.

Properties. — Manganese is a hard,4brittle metal, the colour of
which is grey with a slight tinge of red. Its melting-point is below
that of chromium but higher than that of iron, being about 1900°.
It slowly oxidises on exposure to moist air, and is readily dissolved
by dilute sulphuric, hydrochloric, or even acetic acid, with evolution
of hydrogen. The chief use of manganese is in the iron industry
(see Iron).

Oxides Of Manganese. — The four most important of these are —
Manganous oxide (manganese monoxide) . . MnO.
Red manganese oxide . . . . . Mn3O4.
Manganic oxide (manganese sesquiooci.de) . . Mn2O3.
Manganese dioxide „ MnO2.

The monoxide and sesquioxide are basic, giving rise to man-
ganous and manganic salts respectively. The oxide, Mn3O4, is
also basic, but yields with acids both manganous and manganic

666

Manganese Dioxide 667

salts. Manganese dioxide or peroxide, MnO2, gives manganous
salts with elimination of available oxygen. It also combines with
certain more basic oxides, forming unstable compounds known as
manganites.

Manganese trioxide, MnO3, and hept-oxide, Mn2O7, have also
been obtained. They are both acid oxides, giving rise respec-
tively to the manganates and permanganates.

Manganous Oxide, MnO, is obtained by heating any of the
higher oxides in a stream of hydrogen, or by igniting a mixture of
manganous chloride, sodium carbonate, and ammonium chloride.
It is a light green powder, which, if prepared at a low tempera-
ture, oxidises in the air. When perfectly air-free solutions of
potassium hydroxide and a manganous salt are mixed, with exclu-
sion of air, hydrated manganous oxide, or manganous hydroxide,
Mn(HO)2, is obtained as a white precipitate, which rapidly oxidises
on exposure to air.

Red Manganese Oxide (mangano-manganic oxide\ Mn3O4, is
the most stable of the oxides of manganese, being formed when
both the higher or lower oxides are strongly heated. Thus, in the
preparation of oxygen by heating the dioxide, this compound
remains (page 184). With cold sulphuric acid it yields a mix-
ture of manganous and manganic sulphates, but when heated with
dilute acid, manganous sulphate and dioxide are formed —

Mn3O4 + 2H2SO4 = 2MnSO4+Mn02 + 2H2O.

Manganic Oxide (manganese sesquioxide\ Mn2O3, occurs native
as braunite,2cn& in the hydrated condition as manganite, Mn2O3,H2O.
The hydrated oxide is formed by the spontaneous oxidation of man-
ganous hydroxide, and when gently heated it yields the oxide.
Both the oxide and the hydrate, on treatment with warm nitric
acid, yield manganous nitrate and manganese dioxide.

Manganese Dioxide, MnO2, is the most important of the man-
ganese ores. It may be obtained by the cautious ignition of
manganous nitrate —

Manganese dioxide is a hard black solid which conducts electri-
city and is strongly electro-negative to metals. On this account
it is employed in certain forms of voltaic battery. When heated
it loses oxygen, and forms first the sesquioxide and finally Mn3O4.

668 Inorganic Chemistry

Manganese dioxide dissolves in cold concentrated hydrochloric
acid, forming a dark-brown solution which is believed to contain
the compound MnCl3. On warming it evolves chlorine, and leaves
manganous chloride, MnCl2.

Manganites.^-Manganese dioxide combines with certain me-
tallic oxides, forming unstable compound oxides. Thus, with lime
it forms CaO,MnO2 ; CaO,2MnO2, and CaO,5MnO2. These com-
pounds are produced in the Weldon recovery process (page 359).

MANGANOUS SALTS.

Manganous Chloride, MnCl2, is the only chloride of this metal
that has been isolated. It is obtained by dissolving any of the
oxides or the carbonate in hydrochloric acid, and on evaporation
is deposited as pink crystals of MnCl2,4H2O. The anhydrous
salt is prepared by heating the crystals in a stream of hydro-
chloric acid. Manganese chloride forms double salts with chlo-
rides of the alkalies, the ammonium salt MnCl2,2NH4Cl,H2O being
the best known.

Manganous Sulphate, MnSO-i, is prepared by strongly heating
a pasty mixture of the dioxide and strong sulphuric acid. The
iron present is thereby converted into ferric oxide, and on treat-
ing the calcined mass with water manganous sulphate dissolves.
The solution on evaporation deposits, at ordinary temperatures,
large pink crystals of MnSO4,5H2O (isomorphous with copper
sulphate). Below 6° rhombic crystals are formed (also pink) of the
composition MnSO4,7H2O (isomorphous with ferrous sulphate).

When these salts are heated to 200°, or when their solutions are
boiled, the anhydrous salt is formed. With sulphates of the
alkalies, manganous sulphate forms double salts, as potassium
manganous sulphate, K2SO4,MnSO4,6H2O ; and with aluminium
sulphate it yields z. pseudo-alum (see page 620), MnSO4,Al2(SO4)3,

MANGANIC SALTS.

Manganic Chloride is obtained as a dark-brown solution
when the dioxide is dissolved in cold hydrochloric acid. It has
never been isolated, and is believed to have the composition
MnCl3.

Manganic Sulphate, Mn2(SO4)3, is obtained as a green deli-
quescent powder by the action of sulphuric acid upon the pre-

Permanganates 669

cipitated peroxide. On exposure to the air the deliquesced mass
becomes muddy, by the precipitation of hydrated manganic oxide,
thus—

Mn2(S04)3 + 4H20 = 3H2SO4 + Mn2O3,H2O.

On the addition of potassium sulphate to a solution of manganic
sulphate in dilute sulphuric acidj potassium manganese alum is
obtained, K2SO4,Mn2(SO4)3,24H2O, which deposits in violet regular
octahedra. In the presence of much water the salt is decomposed,
and deposits the hydrated manganic oxide.

MANGANATES.

These salts are derived from the hypothetical manganic acid,
H2MnO4. The oxide corresponding to this acid is known, viz.,
MnO3. It is an unstable compound, obtained as a reddish amor-
phous mass, by adding a solution of potassium permanganate in
sulphuric acid to dry sodium carbonate.

The manganates of the alkalies are obtained by fusing manganese
dioxide with potassium or sodium hydroxide. If air be excluded
the following reaction takes place —

In the presence of air or oxygen, or by the addition of potassium
nitrate or chlorate, more of the manganese is converted into man-
ganate. The fused mass has a dark-green colour, and dissolves in
a small quantity of cold water to a deep green solution, which is
only stable in the presence of free alkali.

When a solution of potassium manganate is largely diluted or
gently warmed, it changes from green to pink, owing to the con-
version of the manganate into permanganate, thus —

3K2MnO4 + 2H2O = 2KMnO4+MnO2 + 4KHO.

The same change takes place when carbon dioxide is passed
through the solution.

PERMANGANATES.

These salts are derived from permanganic acid, HMnO4. When
potassium permanganate is cautiously added to cold strong sul-
phuric acid, green oily drops of the unstable manganese heptoxide

670 Inorganic Chemistry

(or permanganic anhydride) are obtained, Mn2O7. This compound
dissolves in a small quantity of water to a purple solution, which
contains the unstable acid, Mn2O7,H2O, or H2Mn.>O8 = 2HMnO4
The solution evolves oxygen and deposits manganese dioxide.

Potassium Permanganate, KMnO4, is the most important salt
of this class. It is prepared by fusing the dioxide with potassium
hydroxide and potassium chlorate, dissolving the manganate so
obtained in water, and passing carbon dioxide through the solu-
tion. The filtered solution, on evaporation, deposits dark purple
rhombic prisms, which appear deep red by transmitted light.
Potassium permanganate is isomorphous with potassium per-
chlorate, KC1O4 ; it dissolves in water, forming a rich purple
solution. When boiled with strong caustic alkalies it loses oxygen
and forms the green potassium manganate —

2KMnO4 + 2KHO = 2K2MnO4+ H2O + O.

It readily gives up oxygen to oxidisable and organic compounds,
and on this account is used both as a laboratory oxidising agent
and as a disinfectant. The crude sodium salt is largely employed,
under the name of Candy's Disinfecting Fluid, for this purpose.
When solid potassium permanganate is heated to 240° it evolves
oxygen, and forms potassium manganate and manganese dioxide —

2KMnO4= K2MnO4+ MnO2 + Ojj.

CHAPTER XIII

THE TRANSITIONAL ELEMENTS OF THE FIRST
LONG PERIOD

Iron, Fe=55-85. Cobalt, €0 = 58.97. Nickel, Ni=s8.7.

THESE three elements belonging to Group VIII. (see classifica-
tion, page 118) stand in a different relation to each other than the
members of the other seven groups.

Iron, cobalt, and nickel belong to the same period, being the
transitional elements falling between the first and second series of
the first long period. They are related, on the one hand, through
iron, to the preceding metals manganese and chromium (see such
compounds as ferrates, manganates, chromates) ; while, on the other
hand, through nickel, they approach the metal copper, which is the
next following in the period.

Iron, cobalt, and nickel are closely related elements ; in nature
they are usually associated together. They are all attracted by
the magnet, and are nearly white, hard, and difficultly fusible
metals. In their chemical habits, however, they exhibit a gradual
transition in their properties. Thus, iron forms two basic oxides,
yielding two series of stable salts, viz., ferrous and ferric. Cobalt
also has two basic oxides, but the basicity of the sesquioxide is
very feeble, and cobaltzV salts (except double salts) are unstable,
and are only known in solution. Nickel only forms one basic
oxide, and yields only one series of salts corresponding to the
ferrous salts, the sesquioxide of nickel behaving with acids as a
peroxide,

IRON.

Symbol, Fe. Atomic weight =55. 85.

Occurrence. — Iron is one of the most abundant and widely
distributed elements. It occurs in the uncombined state in small

particles disseminated through certain basalts, and also in meteoric

671

6/2 Inorganic Chemistry

iron, where it is usually associated with nickel, cobalt, and copper.
Masses of iron have also been found which have been formed by
the reduction of iron ores, owing to the firing of coal pits : such
iron is known as natural steel.

The chief ores of iron are red hcematite and specular iron ore,
Fe^Og ; brown h^matite^ 2Fe2O3,3H2O ; magnetic iron ore (load-
stone), Fe3O4 ; spathic iron ore> FeCO3 ; clay iron stone consists
of spathose iron mixed with clay ; and blackband is clay iron
stone containing from 20 to 25 per cent, of coal.

Iron is also found in combination with sulphur, as iron pyrites^
FeS2, and with iron and copper in copper pyrites, Cu2S,Fe2S3, but
these compounds are not employed in the metallurgy of iron.

Modes of Formation. — Iron is readily reduced from its com-
pounds. Thus, if ferric oxide or oxalate be gently heated in a
stream of hydrogen, the metal is obtained as a black powder,
which spontaneously oxidises with incandescence when brought
into the air. On the industrial scale the reduction is effected by
means of coke and limestone. The ore is first calcined, whereby
water and carbon dioxide are expelled, and any sulphides present
are oxidised, with the expulsion of sulphur dioxide. By this pro-
cess also the ore is rendered more porous. The calcined ore is
then smelted in a blast-furnace, with limestone and coke. Fig. 1 53
shows in section a modern blast-furnace. The charge is admitted
at the top by means of the cup and cone arrangement, which closes
the furnace, and a powerful hot-blast is forced through tuveres
placed round the base of the furnace. The furnace gases are led off
by the side pipe at the top, and are utilised for heating the blast.

The chemical reactions which take place in a blast-furnace are
many and complex, and differ in different parts of the furnace.
In the main, the following are the changes which occur. The
atmospheric oxygen of the hot-blast, on coming in contact with
the carbon, forms carbon monoxide (at the high temperature
carbon dioxide is probably not first formed). As the charges of
ore gradually work their way down the furnace, they soon arrive
at a point where the ferric oxide begins to be reduced by the
heated carbon monoxide, first to ferrous oxide, and then to a
spongy or porous mass of metallic iron. The region where this
takes place is termed the zone of reduction-*-

Fe2O3 + SCO = 3CO2 + 2Fe.
In the early stages of its descent through the furnace, the lime-

Iron

673

stone is converted into carbon dioxide and lime. The reduced
spongy metal, as it passes down through the hotter regions of the
furnace, begins to take up carbon. It is probable that carbon

15-0 •

FIG. 153.

monoxide first combines with the reduced iron, forming iron
carbonyl (see page 300), which at a higher temperature is decom-
posed, with the precipitation of finely divided carbon within the
pores of the mass. More and more carbon is taken up by the iron

2 u

674 Inorganic Chemistry

as it descends, until it passes from a pasty condition to a state of
complete fusion, when it collects upon the bottom, or hearth, of the
furnace. In passing through the hottest regions the lime combines
with the siliceous materials originally present in the ore to form a
fusible slag, beneath which the molten iron collects. Other re-
actions which go on in various regions of the furnace are the reduc-
tion of sulphur compounds, and of phosphates and silicates, with
the absorption into the iron of a certain amount of sulphur, phos-
phorus, and silicon. The precise nature of the changes suffered
by the. gases in the various regions of the furnace is still obscure.
The cyanogen formed by the direct union of atmospheric nitrogen
with carbon, and also the hydrocarbons present, doubtless undergo
a chemical change in contact with the heated iron, and probably
aid in its carburisation. The molten iron is drawn off at intervals
from a tap-hole into moulds, and is known as cast iron or pig iron.
The slag as it accumulates overflows in a regular stream through
an opening known as the slag hole. When such a furnace is in full
blast, fresh charges of materials are introduced at regular intervals,
and the process continues uninterruptedly for years. The metal
obtained from the blast-furnace is far from pure iron, but contains
varying quantities of carbon, silicon, phosphorus, sulphur, and
manganese.

The carbon may be present either in combination with iron as
a carbide, or distributed throughout the metal as fine particles of
graphite, or in both of these forms. White cast iron contains its
carbon in the combined form, while grey cast iron owes its grey
colour to the presence of minute crystals of graphite disseminated
throughout the metal. When grey cast iron is dissolved in hydro-
chloric acid, the graphite remains behind as a black powder ; but
on similarly treating iron containing combined carbon, the carbon
unites with the hydrogen, forming various hydrocarbons, which
impart to the escaping gas a characteristic and unpleasant smell.
Average cast iron contains from 90 to 95 per cent, of iron, and 3 to
5 per cent, of carbon. Spiegel is a variety of white cast iron con-
taining 3.5 to 6 per cent, of carbon, and from 5 to 20 per cent, of
manganese. With more than 20 per cent, of manganese, the
metal is termed ferro-manganesc.

Purification. — The properties of iron are greatly modified by
the presence of various impurities, especially carbon, and for
different purposes for which iron is used, metal of different degrees
of purity is required. The purest form of ordinary commercial

Iron 675

iron is known as wrought iron, while steel is intermediate between
this and cast iron.

The process by which cast iron is converted into wrought iron
is termed puddling; and the method is called either dry puddling
or pig-boiling, depending upon whether the cast iron is subjected
to a preliminary refining or not. The chemical reactions in both
cases are the same, and consist in the oxidation of the impurities ;
the carbon being expelled as carbon dioxide, while the oxides of
silicon, phosphorus, and manganese pass into the slag. The
method of pig-boiling is almost exclusively adopted.

The cast iron is melted in a reverberatory furnace, the working
bottom of which, as well as the lining (or fettling}, consists of a
layer of ferric oxide. The decarburisation of the iron is mainly
effected by means of the oxide of iron derived from the fettling ;
and for some time the molten mass appears to boil, owing to the
escape of carbon monoxide. As the impurities are oxidised and
removed, the mass becomes pasty (owing to the fact that the
melting-point of pure iron is much higher than that of cast iron),
and is then worked up into lumps, or blooms, which are ultimately
removed and placed under a steam hammer, whereby admixed slag
is squeezed out, and the metal is welded into a solid mass.

Wrought iron contains from 0.06 to 0.15 per cent, of carbon.

Steel may be produced either from wrought iron, by adding
carbon, or from cast iron by removing that impurity. Formerly
steel was exclusively obtained by the first method, by what is
known as the cementation process. This simply consists in heating
the bars of iron, buried in broken charcoal, for several days to a red
heat. The precise nature of the chemical change which results in
the carburisation of the iron is not definitely established. In all
probability the carbon is conveyed into the body of the metal
(which is not even heated to the softening point) by the intervention
of iron carbonyl ; the carbon monoxide being formed by the union
of the carbon with the air retained within the layer of charcoal.
At the conclusion of the operation the iron presents a blistered
appearance, and on this account is termed blister-steel.

At the present time steel is mostly produced by the Bessemer
process, which consists in oxidising the impurities present in cast
iron, by blowing through the molten metal a blast of air. This
operation is performed in a large pear-shaped vessel, known as a
converter, which is mounted on trunnions, and through the bottom
of which a powerful air blast can be admitted. The converter is

Inorganic Chemistry

tilted into a horizontal position, and a quantity of molten cast iron
is run in. The air blast is then started and the converter immedi-
ately swung back into a vertical position. In the course of a very
short time the whole of the impurities are burnt away, and the
stage at which the operation is complete is sharply marked, by the
sudden disappearance of the flame from the open mouth of the
converter. The converter is once more swung into a horizontal
position, and the blast is stopped. The exact quantity of molten
spiegel is then added to supply the carbon required to convert the
entire charge into steel, and the blast is turned on for a few
moments in order to thoroughly mix the materials, after which the
contents are poured out into the casting ladie.

The comparative purity of the three forms of iron will be seen
from the three fo1 lowing typical examples :—

Cast Iron. Steel. Wrought Iron.

Carbon . . 3.81 0.65 o.io

Silicon . .1.68 0.07 0.05

Phosphorus . 0.70 0.03 0.15

Sulphur . . 0.60 0.02 0.05

Manganese . 0.41 0.40 0.07

7.20 — 1.17 0.42

Iron . . 92.80 98.83 99.58

100.00 100.00 100.00

Properties. — Pure iron is a white lustrous metal, capable of
taking a high polish. Its specific gravity is 7.84 to 8.139. It is
more difficultly fusible and more malleable than wrought iron, but
at a red heat it becomes soft and can be welded. The physical
properties usually associated with iron are in reality those of
iron containing varying amounts of impurities : thus, pure iron
when rendered magnetic quickly loses this property, whereas
steel retains its magnetism at ordinary temperatures, losing it,
however, when heated. Pure iron, when heated and suddenly
cooled, does not take a temper^ while steel when so treated be-
comes extremely hard and brittle.

Iron is unacted upon by dry air at ordinary temperatures,
but in moist air it quickly becomes coated with rust. The pre-
sence of the atmospheric carbon dioxide appears to be indispensable
to the rusting of iron, the first stage in the process being the for-

Oxides of Iron 677

mation of ferrous carbonate.* Dilute nitric acid dissolves it,
forming ferrous nitrate and ammonium nitrate ; with stronger
nitric acid, ferric nitrate and oxides of nitrogen are formed.

Concentrated nitric acid (specific gravity, 1.45) is without solvent
action upon iron. A strip of iron which has been immersed in
such strong acid is unacted upon when afterwards dipped into
the more dilute acid, and is also incapable of precipitating metallic
copper from a solution of copper sulphate. Iron in this condition
is said to \>e passive. Other oxidising agents, as chromic acid, or
hydrogen peroxide, are capable of bringing about the same result.
It is believed that this condition is due to the formation of a film
of the oxide Fe3O4 upon the surface.

Finely divided iron takes fire spontaneously in chlorine ; and
when gently warmed in sulphur dioxide it combines with that gas
with incandescence. It absorbs carbon monoxide with formation
of iron carbonyl, Fe(CO)5. When heated in ammonia it forms a
nitride, Fe4N2 (see page 278).

Oxides of Iron. — Three oxides of iron are known, namely : —

Ferrous oxide (iron monoxide) . . FeO.
Ferric oxide (iron sesquioxide) . . Fe2O3.
Ferroso-ferric oxide (magnetic oxide) . Fe3O4, or Fe2O3,FeO.

The two first are basic oxides, giving rise respectively to ferrous
and ferric salts ; the third yields both ferrous and ferric salts.

Ferric oxide combines with certain more basic oxides, form-
ing compounds analogous to Fe2O3,FeO ; such as Fe2O3,CaO,
Fe2O3,ZnO. These are known asferrt'tes.

Ferrous Oxide (protoxide of iron), FeO, is formed as an inter-
mediate product during the reduction of ferric oxide by hydrogen
or carbon monoxide ; but it is difficult to obtain it free from either
the higher oxide or the metal. It is also formed when ferrous
oxalate is heated out of contact with air. It is a black powder,
which oxidises in the air, and which dissolves in acids yielding
ferrous salts.

Ferrous Hydroxide, Fe(HO)2, or FeO,H2O, is obtained as a
white precipitate when potassium hydroxide is added to a solution
of a ferrous salt with entire exclusion of air. In the presence of
air it is green. It readily absorbs oxygen and passes into ferric
oxide.

* Moody, Jour. Chem. Soc., 1906.

6; 8 Inorganic Chemistry

Ferric Oxide (sesquioxide of iron), Fe2O3, occurs in brilliant
black crystals belonging to the hexagonal system, in specular
iron ore. It is obtained as a red amorphous powder by heating
hydrated ferric oxide, ferrous sulphate, or ferrous carbonate. In
a crystalline condition it may be produced by carefully heating a
mixture of ferrous sulphate and common salt, or by heating the
amorphous oxide in gaseous hydrochloric acid. The natural com-
pound, and also the artificial substance after strong ignition, is
only slowly dissolved by acids. Ferric oxide is extremely hygro-
scopic. When strongly heated it is partially converted into Fe3O4.
The amorphous substance, obtained by distilling ferrous sulphate
for the manufacture of Nordhausen sulphuric acid, is employed as
a red pigment and a polishing powder under the name of rouge.

Ferric Hydroxide, or Hydrated Ferric Oxide, Fe2(HO)6, or
Fe2O3,3H2O. — When an excess of ammonia is added to a solution
of ferric chloride, and the voluminous brown precipitate is dried
at a moderate temperature, it has the composition Fe2O3,3H2O.
On exposure to various temperatures, or by precipitation under
various conditions, hydrated oxides of the composition Fe2O3,
2H2O ; Fe2O3,H2O, and others, have been obtained ; and several
of these compounds occur in nature. Ordinary rust of iron consists
of a mixture of hydrated ferric and ferrous oxides with ferrous
carbonate, in varying proportions depending upon conditions.
Exposure to air gradually oxidises the ferrous oxide and carbonate
into ferric oxide.

The monohydrate Fe2O3,H2O has been obtained as a soluble modification,
by heating an acetic acid solution of precipitated ferric hydroxide to 100° in
sealed vessels. On the addition of sulphuric acid, a brown precipitate is
obtained, having the composition Fe2O3,H2O, which is insoluble in acids,
but soluble in water. The solution gives no reaction with potassium ferro-
cyanide. Another soluble hydroxide is produced by dissolving the ordinary
precipitated hydroxide in ferric chloride, and subjecting the solution to
dialysis. This solution is employed in medicine under the name of dialysed
iron.

Ferroso-ferrie Oxide, Fe3O4, occurs native as magnetite and
magnetic oxide of iron', the magnetic variety being known also
as loadstone. When iron is heated in the air, the black film
which forms (the so-called iron-scale or hammer-scale) consists of
the oxide Fe3O4, with more or less ferric oxide, Fe2O3, upon the
outer surface. It is also produced when steam or carbon dioxide
is passed over heated iron, with evolution of hydrogen and carbon
monoxide respectively, these reactions being the reverse of those by
which oxides of iron are reduced by hydrogen or carbon monoxide.

Ferrous Sulphate 679

This oxide is also formed as a black precipitate when ammonia
is added to a solution containing mixed ferrous and ferric salts, and
the mixture gently warmed.

Ferrates. — These compounds correspond to the manganates,
but neither the acid H2FeO4 nor the oxide FeO3 are known.
Potassium ferrate, K2FeO4, is formed when chlorine is passed
through a solution of potassium hydroxide in which ferric hydroxide
is suspended.

FERROUS SALTS.

Ferrous Chloride, FeCl2.— The anhydrous compound is ob-
tained by heating iron wire in gaseous hydrochloric acid, when the
salt sublimes in the form of white deliquescent crystals. In aqueous
solution it is obtained when iron is dissolved in hydrochloric acid,
and is deposited in pale blue-green crystals ofFeCl2,4H2O.

When heated in the air it is converted into ferric oxide and
chloride, the latter volatilising —

6FeCl2 + 3O = Fe2O3 + 4FeCl3.

When volatilised in an atmosphere of hydrochloric acid its
vapour-density at high temperatures corresponds to the formula
FeCl2 ; at lower temperatures it lies between the values required
for FeCl2 and Fe2Cl4.

When strongly heated in a current of steam it is decomposed as
follows —

3FeCl2 + 4H2O = Fe3O4 + H2 + 6H Cl.

Ferrous Sulphate (green vitriol), FeSO4,7H2O, is obtained
when iron is dissolved in sulphuric acid. It is prepared on a
large scale by exposing heaps of iron pyrites, FeS2, to the action
of air and moisture. The liquor which drains away contains
ferrous sulphate and sulphuric acid, and the latter is converted
into ferrous sulphate by the introduction of scrap iron.

Ferrous sulphate forms pale green monosymmetric crystals,
which effloresce on exposure to the air. They are soluble in water
to the extent of 70 parts in 100 parts of water at 15°, and 370 parts
in 100 parts at 90°. At 100° the crystals lose 6H2O, being con-
verted into FeSO4,H2CX

If a crystal of zinc sulphate be thrown into a supersaturated
solution of ferrous sulphate, the iron salt is deposited in rhombic

680 Inorganic Chemistry

prisms (isomorphous with zinc sulphate). On the other hand, if a
crystal of copper sulphate be added, asymmetric (triclinic) crystals
of FeSO4,5H2O (isomorphous with copper sulphate) are formed.

Ferrous sulphate forms double salts with the sulphates of the
alkalies. Thus, when mixed with ammonium sulphate in the re-
quisite proportions, ammonium ferrous sulphate, FeSO4,(NH4)2SO4,
6H2O, is obtained. This salt is less readily oxidised on exposure
to air than ferrous sulphate itself.

Ferrous salts give, with potassium ferrocyanide (K4Fe(CN),.,
or 4KCN,Fe(CN)2), a white precipitate of potassium ferrous ferro-
cyanide (FeK2Fe(CN)6, or 2KCN,2Fe(CN)2). The precipitate is
quickly oxidised, and becomes blue. With potassium ferricyanide
(K3Fe(CN)6, or 3KCN,Fe(CN)3), ferrous salts yield a blue pre-
cipitate of ferrous ferricyanide (Turnbull^s blue) (Fe3{Fe(CN)6}2, or
3Fe(CN)2,2Fe(CN)3); thus—

4 + 2K3Fe(CN)6=Fe3{Fe(CN)6}2

FERRIC SALTS.

Ferric Chloride, FeCl3, is prepared in the anhydrous state by
passing dry chlorine over heated iron wire. In solution it may
be obtained by dissolving iron in aqua regia; or ferric oxide in
hydrochloric acid. The anhydrous salt forms nearly black crystals,
appearing deep red by transmitted light. It readily volatilises, and
at temperatures above 700° the density of its vapour corresponds to
the formula FeCl3, while at lower temperatures its density agrees
more nearly with the formula Fe2Cl6.

Ferric chloride is extremely deliquescent, and readily dissolves in
water. When the solution is slowly evaporated, yellow crystals are
deposited, having the composition Fe2Cl6,12H2O (or FeCl3,6H2O).
When a dilute solution of ferric chloride is boiled, it decomposes,
forming either an insoluble oxychloride or a soluble hydroxide and
free hydrochloric acid (depending upon the strength of the solution).

Ferric Sulphate, Fe2(SO4)3, is prepared by the addition of sul-
phuric or nitric acids to a solution of ferrous sulphate —

The brown solution, on evaporation, leaves the anhydrous salt
as a white mass. When the requisite quantity of potassium sul-

Sulphides of Iron 68 1

phate is dissolved in a strong solution of ferric sulphate at o°,
the double potassium iron sulphate (iron alum), K2SO4,Fe2(SO4)3,
24H2O, separates out in the form of violet octahedrons.

Ferric salts give, with potassium ferrocyanide (K4Fe(CN)6, or
4KCN,Fe(CN)2), a dark blue precipitate of ferric ferrocyanide
(Prussian blue\ 4Fe(CN)3,3Fe(CN)2) or Fe4{Fe(CN)6}3—
4FeCl3 + 3K4Fe(CN)6=Fe4{Fe(CN)6}3 + 12KCl.

With potassium ferricyanide ferric salts give no precipitate.

SULPHIDES OF IRON.

Ferrous Sulphide, FeS.— When a white-hot bar of wrought
iron is dipped into melted sulphur, the elements unite ; and the
readily fusible monosulphide of iron falls to the bottom. It may
be prepared by throwing into a red-hot crucible a mixture of iron
filings and sulphur. So obtained, it is a dark, yellowish-grey,
metallic-looking mass. When heated out of contact with air, it
does nojt part with sulphur, but in the presence of air is converted
into ferric oxide and sulphur dioxide. Ferrous sulphide is pre-
cipitated from either ferrous or ferric solutions, by alkaline sul-
phides, as a black amorphous powder, which in the moist state is
quickly oxidised by the air. Dilute sulphuric acid, or hydrochloric
acid, decomposes ferrous sulphide, with evolution of sulphuretted
hydrogen.

Iron Sesquisulphide, Fe2S3, is formed when equal weights of
iron and sulphur are heated to a moderate temperature. It can-
not be obtained by precipitation from a ferric salt, as the product
so formed consists of ferrous sulphide and sulphur —

Fe2Cl6 + 3(N H4)2S = 6N H4C1 + 2FeS + S.

It is a yellow, metallic-looking solid, which is decomposed by
dilute hydrochloric acid, yielding sulphuretted hydrogen.

Ferric Bisulphide, FeS2, occurs in nature in large quantities as
iron Pyrites, sometimes in the massive condition, and at others in
the form of brass-yellow cubical crystals. In many cases the
native compound bears the impression, or assumes the shape,
of various organised forms, such as wood, ammonites, &c., the
mineral having been formed by the reducing action of the organic
matter upon ferrous sulphate in solution. Ferric disulphide is
also found in the form of brass-like, rhombic crystals in radiated
pyrites.

682 Inorganic Chemistry

The compound may be prepared by heating to a low red heat
a mixture of ferrous sulphide and sulphur.

Ferric disulphide is unacted upon by dilute acids : hot con-
centrated hydrochloric acid decomposes it, with liberation of sul-
phur and sulphuretted hydrogen. When heated in hydrogen,
sulphur is evolved (which partly combines with the hydrogen),
and ferrous sulphide remains. When heated in the air, ferric
oxide and sulphur dioxide are formed.

Ferroso-ferrie Sulphide (magnetic pyrites), Fe3S4, occurs in
the form of hexagonal crystals. Like the corresponding oxide,
this compound is attracted by the magnet, and is itself sometimes
magnetic.

COBALT.

Symbol, Co. Atomic weight =58. 97.

Occurrence.— Cobalt is not found uncombined in nature. Its
chief natural compounds, which are only sparsely distributed, are
spelts-cobalt, or smaltine, CoAs2 ; cobalt glance, CoAsS, in both of
which the cobalt is partially replaced by nickel and iron ; and
cobalt-bloom^ Co3(AsO4)2,8H2O.

Modes Of Formation. — Cobalt is obtained by reducing the
oxide, or the chloride, in a stream of hydrogen, or by strongly
heating cobalt oxalate in a closed crucible. It is also readily
obtained by reduction of its oxide with powdered aluminium.

Properties.— Cobalt is an almost white, hard metal, which,
when polished, resembles nickel, but is slightly bluer. It is
malleable, and when heated is very ductile. Like both iron and
nickel, it is attracted by the magnet ; but unlike these, it retains
this property, even at a red heat. In the massive form, cobalt is
unacted upon by the air ; but the finely-powdered metal, obtained
by the reduction of the oxide in hydrogen^ rapidly oxidises on
exposure to the air, sometimes with incandescence. When heated
in the air, it forms the oxide Co3O4. Cobalt decomposes steam at
a red heat, yielding cobaltous oxide, CoO.

Oxides Of Cobalt.— Three oxides of cobalt are recognised,
namely, cobaltous oxide, CoO ; cobaltic oxide, Co2O3 ; and cobalto-
cobaltic oxide, Co3O4.

Four other oxides are known, which are regarded as compounds of the two
first, having the composition 2CoO,Co2O3; 3CoO,Co2O3 ; 4CoO,Co2O3;
6CoO,Co2O3.

Cobaltous Chloride 683

The monoxide, CoO, is basic, and yields the cobaltous salts.
The sesquioxide, Co2O3, is feebly basic, forming only unstable
salts. Stable double salts, however, corresponding to this oxide
are known.

Cobaltous Oxide (cobalt monoxide), CoO, is formed when the
sesquioxide is heated to redness in a stream of carbon dioxide, or
gently heated in hydrogen. It is also obtained when the carbo-
nate or hydroxide is heated in the absence of air. It forms a drab-
coloured powder, which is unacted upon by the air, but when heated,
forms Co3O4. When heated in either hydrogen or carbon mon-
oxide, it is reduced to metallic cobalt.

Cobaltous Hydroxide, Co(HO)2.— When potassium hydroxide
is added to a solution of a cobaltous salt, a blue basic hydrate is
precipitated, which, on boiling, is converted into the pink hydroxide
CXHO)^ It turns brown on exposure to the air, by the absorp-
tion of oxygen. Both the oxide and hydroxide are really soluble
in acids, giving cobaltous salts.

Cobaltie Oxide (cobalt sesquioxide), Co2O3, is obtained by care-
fully heating cobaltous nitrate until red fumes cease to be evolved.
It is a dark grey powder, which, when strongly heated, is con-
verted into the intermediate black oxide, Co3O4. Cobaltie oxide
dissolves in cold acids, forming brown solutions, which contain
unstable cobaltic salts. When warmed, these are converted into
cobaltous salts, with evolution of oxygen in the case of oxy-salts,
and of the halogen from haloid salts. This sesquioxide, therefore,
behaves as a peroxide.

Cobaltie Hydroxide, Co2(HO)6, or Co2O3,3H2O, is obtained as a
nearly black precipitate, by the addition of an alkaline hypochlorite
to a cobaltous salt. With acids it behaves as the oxide.

Cobalto-Cobaltie Oxide, Co3O4, is formed as a black powder,
when the sesquioxide is strongly heated in air.

COBALTOUS SALTS.

Cobaltous Chloride, CoCl2.— When the carbonate5 or any of
the oxides, are dissolved in hydrochloric acid, the concentrated
solution deposits dark red prisms (monosymmetric), having the
composition CoCl2,6H2O. When exposed over sulphuric acid, they
lose 4H2O, and are converted into a rose-red salt, CoCl2,2H2O,
which reabsorbs moisture from the air to form the hexahydrate.

684 Inorganic Chemistry

When the dihydrate is heated to about 100°, it Is converted into
violet-blue crystals of CoCl2,H2O ; and at 120° it becomes an-
hydrous, and is blue. The blue salts, on exposure to the air,
rapidly rehydrate themselves, and become pink.

Cobaltous chloride dissolves in alcohol, giving a deep blue solu-
tion, which, on the addition of water, also becomes pink. This
property of forming pink hydrated salts, which become blue or
green when nearly or quite anhydrous, is common to most cobal-
tous salts. Thus, the iodide CoI2,6H2O forms rose-coloured
crystals. When gently heated, it changes to a moss-green salt,
CoI2,3H2O, which, when dehydrated, becomes nearly black.

Cobaltous Sulphate, CoSO4,7H2O, is obtained by dissolving
the carbonate or oxides in sulphuric acid, and is deposited from
the solution in dark red crystals, isomorphous with ferrous sul-
phate. Cobalt sulphate, like the sulphates of iron and nickel,
forms double salts with alkaline sulphates, of which cobalt potas-
sium sulphate, CoSO4,K2SO4,6H2O, is the best known.

Cobaltie Salts. — Single salts corresponding to cobalt sesqui-
oxide are unstable, and exist only in solution. More stable double
salts are known. Thus, when potassium nitrite is added to an
acetic acid solution of cobalt chloride, a yellow crystalline precipi-
tate is obtained, consisting of the double nitrite of cobalt and
potassium —

= Co2(NO2)6,6KNO2 +
2NO + 4KC1 + 2H2O.

The formation of this compound is made use of for separating cobalt from
nickel, the latter element yielding no corresponding double nitrite. In the
presence, however, of salts of barium, strontium, or calcium, nickel forms,
with potassium nitrite, triple salts, such as Ni(NO2)2,Ba(NOa)..,,2KNOo, which
are precipitated as yellow crystalline powders. Hence, in the presence of
metals of the alkaline earths, nickel and cobalt cannot be separated by this
method.

SULPHIDES OF COBALT.

Cobaltous Sulphide, CoS, is obtained by heating cobaltous
oxide with sulphur, or by fusing a mixture of cobalt sulphate,
barium sulphide, and common salt. It forms bronze-coloured
crystals, which are soluble in strong hydrochloric acid. Cobalt
sulphide is precipitated as a black amorphous powder when
ammonium sulphide is added to a cobalt solution. The precipi-
tate slowly dissolves in dilute mineral acids, but is insoluble in

Cob alt am ines 685

acetic acid. When heated in a stream of sulphuretted hydrogen,
it yields the sesquisulphide Co2S3 ; and if mixed with sulphur, and
heated in a current of hydrogen, it forms the disulphide CoS2.

Cobaltamines (ammoniacal cobalt compounds *). Cobalt forms
a large number of complex ammoniacal salts. A few of these
contain the metal in the divalent condition, and are known as
ammonio-cobaltous salts; but by far the larger number contain
the hexavalent double atom Co2, and are termed ammonio-cobaltic
compounds. These compounds are classified as follows t : —

Animonio-Cobaltous Salts are formed by the absorption of gaseous am-
monia by anhydrous cobaltous salts, or by dissolving the salts in strong
aqueous ammonia, with exclusion of air. In this way the following salts have
been obtained —

^ ™ *XTTT f which, at 120°, is converted
Ammonio-cobaltous chloride, GoCl2,6NH3 •! jntoCoCL2NH

Ammonio-cobaltous sulphate, CoSO4,6NHo.
Ammonio-cobaltous nitrate, Co(NO3)2,6NH3,2H2O.

Ammonio-Cobaltic Salts.— These may be arranged under the following
classes and subdivisions : —

I. Hexammonio Salts. — General formula, Co2(NH3)6'R6, where R equals
a monacid radical, or its equivalent of di or tri acid radicals.

( Hexammonio-cobaltic chloride (dichro-cobaltic chloride)
Examples -j Co2-(NH3)6'Cl6,2H2O.

^ Hexammonio-cobaltic sulphate, Co2'(NH3)6-(SO4)3,6H2O.

II. Octammonio Salts —

(a.) Praseo% Salts. — General formula, Co2'(NH3)8'R6.

( Praseo-cobaltic chloride, Co2(NH3)s'Cl6,2H2O.
Examples \ Praseo-cobaltic chloro-nitrate, Co2(NH3)8-Cl4-(NO3)2,
I 2H20.

(p.) Fusco Salts.— General formula, Co2(NH3)s(HO)2-R4.

f Fusco-cobaltic chloride, Co2(NH3)8(HO)2-Cl4,2H2O.
Examples \ Fusco - cobaltic sulphate, Co2(NH3)s(HO)2'(SO4)2,
t 2H20.

* For details respecting the preparation and properties of these salts the
student is referred to larger works.

f On the constitution of metallammonium compounds generally, see Werner,
7.eitschrift fur Anorganische Chemie, 1893, vol. iii.

t These names denote the characteristic colours of the salts, thus— -prasimus,
leek-green ; fuscus, swarthy ; crocus, yellow, &c.

686 Inorganic Chemistry

(7.) Croceo Salts.— General formula, Co2(NH3)8(NO2)4'R2.

Examples / Croceo-cobaltic chloride, j.

I Croceo-cobaltic sulphate, Co2(NH3)8(NO2)4'SO4.

III. Decammonio Salts —

(a.) Roseo Salts.— General formula, Co2(NH3)10(H2O)2R6.

f Roseo-cobaltic chloride, Co2(NH3)10(H2O)2Cl6.
Examples \ Roseo-cobaltic sulphate, Co2(NH3)10(H2O)2-(SO4)3,
I 3H20.

8.) Purpureo Salts.— General formula, Co2(NH3)10X2R4

(where X and R are either the same or different acid radicals).

IChloro-purpureo-cobaltic chloride,Co2(NH3)10Cl2'Cl4.
Chloro-purpureo-cobaltic sulphate, Co2(NH3)10Cl2'
(SO4)2.
Bromo - purpureo - cobaltic nitrate, Co2(NH3)10Br2'
(N03)4.

(7.) Xantho Salts.— General formula, Co2(NH3)10(NO2)2'R4.

( Xantho-cobaltic chloride, Co2(NH3)10(NO2)2-Cl4.
Examples -j Xantho-cobaltic bromo-nitrate, Co2(NH3)10(NOr>)2'
I BV(N03)2.

IV. Oxy-decammonio Salts.— General formula, Co2(NH3)10R4-X'O(HO)

(where X is either (HO) or an acid radical either the same as,
or different from, R).

/'Oxy-decammonio cobaltic chloride, Co2(NH3)10Cl4*

Examples} (HO)'O-(HO).

| Anhydro - oxy - decammonio cobalt chloride,
( Co2(NH3)10Cl4-CVO'(HO).

V. Dodecammonio Salts (luteo-cobaltic salts). —General formula,Co2(HN3)12R6.

/Luteo-cobaltic chloride, Co2(NH3)12Cl6.

i. Luteo-cobaltic sulphate, Co2(NH3)12(SO4)3,5H2O.

When cobalt compounds are fused with borax, a clear blue
vitreous mass is obtained, which contains a borate of cobalt. A
similar blue colour is imparted to ordinary potash glass when a
small quantity of a cobalt salt is added to the molten material,
owing to the formation of a silicate of cobalt. Under the name
of smalty this substance has been manufactured for use as a pig-
ment, by fusing the roasted cobalt ore with quartz sand and pearl-
ash. The fused mass of deep blue glass is then finely ground
beneath water.

Nickel Alloys 687

NICKEL.

Symbol, Ni. Atomic weight =58. 7.

Occurrence. — Nickel occurs chiefly in combination with arsenic
as kupfer nickel? Ni2As2 ; white nickel, NiAs2 ; nickel glance,
Ni2(AsS)2, also as nickel blende, NiS. Nickel ore almost invari-
ably contains cobalt, and frequently antimony and bismuth.

Modes of Formation. — Nickel is obtained by reducing the
oxide with carbon at a high temperature. It may be obtained as a
black powder by reducing nickelous oxide in a stream of hydrogen,
or by heating nickelous oxalate out of contact with air. It is also
obtained as a lustrous coherent deposit by the electrolysis of an
ammoniacal solution of the double sulphate of nickel and ammonia.

Nickel in a high state of purity is now being made on a com-
mercial scale by what is known as the " Mond's " process. This
consists in passing carbon monoxide over gently heated niclcel
oxide, whereby the nickel is first reduced and is then taken up by
the carbon monoxide to form nickel carbonyl, Ni(CO)4 (see p. 299).
This volatile compound is then passed through tubes which are
more strongly heated, which causes the compound to decompose
into carbon monoxide (which can be again utilised) and metallic
nickel. In this way the metal is deposited in the form of a coherent
solid, entirely free from cobalt, with which nickel is always associated
in its ores.

Properties.— Nickel is a lustrous white metal, with a faint
yellow tinge when compared with silver. It is ductile and malle-
able, and at the same time very hard and tenacious. It is sus-
ceptible of a very high polish. Nickel is attracted by the magnet,
but loses this property when moderately heated. When obtained
by reduction with charcoal, the metal contains a certain amount of
carbon (like cast iron), which renders it less malleable, and when
produced by reduction of the oxalate at a low temperature the
powder is pyrophoric.

In the massive form, nickel is unacted upon by moderately dry
air, but in moist air it tarnishes, and becomes covered with a film

* Kupfer nickel signifies the false copper, and was applied by the Germans
in the Middle Ages to this ore, which resembled a copper ore, because they
tried in vain to extract copper from it. It is probable that this ore had been
smelted along with copper ores, under the belief that it contained copper, by
the early andents. Thus, a coin, 235 B.C., has been found to contain 20 per
cent, of nickel.

688 Inorganic Chemistry

of nickelous oxide. It decomposes steam only slowly at a red
heat, and is slowly attacked by dilute hydrochloric or sulphuric
acid (contrast iron).

Nickel is largely used for electro-plating iron and steel articles.
It is also employed on an extensive scale in the production of nickel
steel for modern armour plate, in which the proportion of nickel
reaches 20 or even 30 per cent.

Nickel Alloys. — With copper, and with copper and zinc, nickel
furnishes several important alloys. The small coinage in use in
Belgium, Germany, and the United States consists of i part of
nickel and 3 parts of copper ; while the so-called German silver,
or nickel-silver, contains in addition about 1.5 part of zinc.

Oxides of Nickel, — Three oxides of nickel have been obtained,
namely,, nickelous oxide, NiO ; nickelic oxide, Ni2O3 ; and nickelo-
nickelic oxide, Ni3O4. The first alone is basic.

.Niekelous Oxide (nickel monoxide), NiO, is obtained as a
greenish powder by heating nickel carbonate or hydroxide out of
contact with air. It is dissolved by acids yielding nickel salts.
When heated in hydrogen or carbon monoxide it is readily re-
duced to the metallic state.

Niekelous Hydroxide, Ni(HO)2, is obtained in a pale green
precipitate when potassium hydroxide is added to a solution of a
nickel salt ; the precipitate has the composition 4Ni(HO)2,H2O.
When strongly heated it is converted into nickelous oxide and
water. It is readily soluble in acids, forming the nickel salts, and
it also dissolves in ammonia and in solutions of ammonium salts.

Nickel Sesquioxide, Ni2O3, is obtained as a black powder
when the nitrate is decomposed by heat at the lowest temperature.
With hydrochloric acid and sulphuric acid it behaves like a per-
oxide, yielding nickel salts, with the elimination of chlorine and
oxygen respectively —

Ni.2O3 + 6HCl =2NiCl2 +3H2O + C12. •

It is soluble in ammonia, with evolution of nitrogen —

Hydrated Sesquioxide of Nickel, Ni2(HO6), or Ni2O3,3H2O.

When chlorine is passed through water or sodium hydroxide, in which
nickelous hydroxide, Ni(HO)2, is suspended, a black powder is ob-
tained having the composition Ni2O3,3H2O. The same compound
is obtained when a nickel salt is added to a solution of bleaching-
powder. In contact with acids and ammonia it behaveslike the oxide.

Nickelous Sulphide 689

Niekelo-niekelie Oxide, Ni3O4, is obtained as a grey metallic-
looking mass, when nickel chloride is heated to about 400° in a
stream of oxygen.

Nickel Salts. — Nickel forms only one series of salts, corre-
sponding to the monoxide. In the anhydrous state these are
usually yellowish, while in the hydrated condition they are green.

Nickel Chloride, NiCl2, is obtained as a yellow amorphous
mass, by dissolving the oxide or carbonate in hydrochloric acid,
and evaporating the solution to dryness. When heated in a
current of chlorine it sublimes in the form of lustrous golden
yellow scales, which dissolve in water, forming a green solution.
Fr6m the aqueous solution, green crystals of the composition
NiCl2,6H2O are deposited.

Anhydrous nickel chloride absorbs gaseous ammonia, forming the
compound NiCl2,6NH3, which when deposited from an aqueous
solution forms blue octahedrons.

Nickel Sulphate, NiSO4,7H2O, is produced when the metal,
the carbonate, or the oxide is dissolved in dilute sulphuric acid,
and the concentrated solution is allowed to crystallise at the ordi-
nary temperature. It forms green crystals, isomorphous with
magnesium sulphate. When heated to 100° the crystals lose
6H2O, and above 300° the salt becomes anhydrous. The anhy-
drous salt absorbs gaseous ammonia, being converted into a pale
violet powder having the composition NiSO4,6NH3. When nickel
sulphate is dissolved in strong aqueous ammonia, the solution
deposits dark blue tetragonal crystals of NiSO,4,4NH3,2H2O.

With sulphates of the alkalies, nickel sulphate forms double
salts, of which the ammonium salt is the most important, NiSO4,
(NH4)2SO4,6H2O. It is obtained by mixing concentrated solu-
tions of the two sulphates in the requisite proportions. This salt
is employed in the process of nickel-plating.

Niekelous Sulphide (nickel moncsulphide}^ NiS, occurs as the
mineral capillary pyrites. It is obtained as a bronze-like mass,
insoluble in hydrochloric acid, by heating sulphur and nickel
together. In the hydrated condition nickel sulphide is precipitated
as an amorphous black powder, on the addition of ammonium
sulphide to a nickel salt. The precipitate is scarcely soluble in
hydrochloric acid, but partially dissolves in excess of ammonium
sulphide, forming a brown solution. Three other sulphides have
been obtained, having the composition Ni2S, NiS2, and Ni3S4.

2 X

CHAPTER XIV

THE TRANSITIONAL ELEMENTS OF THE SECOND
AND FOURTH LONG PERIOD

Ruthenium, Ru= 101.7. Rhodium, Rh — 103. Palladium, 106.

Osmium, 03 = 191. Iridium, Ir=i93. Platinum, 194.8.

THESE elements, although constituting two transitional groups, are very closely
related to each other. In nature they all occur associated together in what is
commonly known "as, platinum ore, and they are on this account usually spoken
of as the platinum metals.

Platinum ore, or native platinum, contains all these elements in the metallic
state. It is found in small grains, sometimes in nuggets, in alluvial deposits and
river sand, principally in Brazil, Borneo, California, Australia, and the Urals.
Native platinum contains from 60 to 86 per cent, of platinum, the remainder
consisting of the other five metals of the group, together with varying quan-
tities of gold, copper, and iron. Amongst the grains of platinum ore there
are also found grains which consist essentially of an alloy of platinum and
iridium (containing from 30 to 75 per cent, of iridium) known as platin-
iridium : and also particles of an alloy of osmium and iridium (called osmiri-
diurn], which contain from 30 to 40" per cent, of osmium, as well as small
quantities of rhodium and ruthenium.

They are all white lustrous metals, having high melting-points. They are
unacted upon by air or oxygen at ordinary temperatures ; and, with the excep-
tion of osmium (which burns when strongly heated, forming the tetroxide),
they are scarcely oxidised by oxygen at any temperature.

With the exception of palladium, which readily dissolves in hot nitric acid,
these metals are unacted upon by ordinary acids. Aqua regia converts
osmium into the tetroxide ; it dissolves platinum with formation of the tetra-
chloride, * and slowly acts upon ruthenium, but is without action upon
rhodium and iridium.

The specific gravities of the metals of the first group, although very close to
one another, are widely different from those of the second group ; and it will
be seen that the specific gravities fall, with increasing atomic weights, thus —

Ru, sp. gr. =12.26. Rh, sp. gr. = 12.1. Pd, sp. gr. = 11.4.

Os, ,, =22.47. Ir> »> =22.38. Pt, ,, =21.5.

The element osmium is the heaviest known substance.

The most easily fusible of these metals is palladium, which melts about the
temperature of wrought iron. The melting-point of platinum is somewhat
higher, but it may be boiled by the oxyhydrogen flame. Rhodium and

690

Platinum 691

indium come next in order of fusibility, the latter metal being just fusible by
the oxyhydrogen flame, while ruthenium has a still higher melting-point.
Osmium has not been melted. When heated to the melting-point of iridium,
osmium volatilises ; and if air be present, it burns.
The following oxides of these metals are known —

- Pd20

RuO OsO RhO PdO PtO

RuA Os2O3 Rh2O3 Ir2O3

RuO2 OsOo RhO2 IrO2 PdO2 PtO2

RuO4 OsO4

Ruthenium, osmium, rhodium, and iridium form salts corresponding to the
sesquioxide, such as ruthenious chloride, Ru2Cl6; rhodium sulphate, Rh2(SO4)3;
iridious chloride, Ir2Cl6.

With the exception of rhodium, they all form chlorides, corresponding to
the dioxides, thus — ruthenic chloride, RuCl4; iridic chloride, IrCl4; platinic
chloride, PtCl4, while palladium and platinum yield pa.lla.dous and platinowj
compounds, corresponding to their monoxides.

The tetr oxides of ruthenium and osmium are remarkable in melting at
an extremely low temperature (about 40°), and boiling about 100°. They
yield intensely irritating vapours, which, in the case of osmium tetroxide,
exerts a most injurious effect upon the eyes, and is extremely poisonous.
(Osmium tetroxide is commonly known as osmic acid. ) Osmium and ruthenium
also exhibit a non-metallic character in forming compounds derived from the
unknown ruthenic and osmic trioxides, such as potassium ruthenate, K2RuO4,
and potassium osrnate, KoOsO4 (the corresponding ruthenic and osmic acids
are unknown). Ruthenium also forms potassium per-ruthenate, KRuO4
(analogous to permanganate), although the corresponding acid and peroxide,
Ru2O7, are unknown. The most important of these elements is platinum.

PLATINUM.

Symbol, Pt. Atomic weight =194. 8.

In order to separate platinum from the other metals with which
the native platinum (see page 690) is mixed, the ore is digested in
dilute aqua regiay under slightly increased pressure. The solution
so obtained contains the higher chlorides of platinum, palladium,
rhodjum, and iridium (for although in the pure state the last two
named metals are scarcely attacked by aqua regia, when alloyed with
much platinum they dissolve). The solution is evaporated to dry-
ness, and heated to 125°, whereby the palladium and rhodium are
obtained in the form of their lower chlorides, PdCl2 and Rh2Cl6
(the latter of which, in the anhydrous condition, is insoluble in
water). The residue is extracted with water, and to the clear solu-

692 Inorganic Cfremistry

tion, acidified with hydrochloric acid, ammonium chloride is added.
The double chloride of platinum and ammonium (PtCl4,2NH4Gl),
separates out as yellow crystals, while the corresponding iridium
salt, being more soluble, remains for the most part in solution,
and may be obtained- by concentrating the mother-liquor. The
ammonium platinic chloride, on being ignited, loses ammonium
chloride and chlorine, leaving the metal in the form of a black
spongy mass known as spongy platinum, which is then melted by
means of the oxyhydrogen flame in a lime crucible. The platinum
so obtained usually contains small quantities of iridium and traces
of associated metals.

Pure platinum is obtained by alloying commercial platinum
with pure lead, and treating the alloy first with nitric acid, which
dissolves any copper and iron, a part of the palladium and rhodium,
and most of the lead ; and then with dilute aqua regta^ which dis-
solves the whole of the platinum and the remaining lead, with
traces of rhodium. From this solution the lead is precipitated as
sulphate, and the platinum is then precipitated as the double
chloride, by ammonium chloride. To remove traces of rhodium
which are present, the dried double chloride is ignited with
hydrogen potassium sulphate, whereby the rhodium is converted
into a soluble double sulphate of rhodium and potassium, while
the platinum is reduced to the condition of the spongy metal.

Properties. — Platinum is a lustrous, greyish-white, malleable,
and ductile metal. At a red heat it may be welded with great
ease. It is melted by the oxyhydrogen flame, and vessels of
platinum are readily made by fusing the metal together in this
way. Heated platinum absorbs large quantities of hydrogen
(see page 179) ; and when the metal is melted in the oxyhydrogen
flame, it exhibits the phenomenon of "spitting," when it again
solidifies (see Silver, page 562). Platinum does not combine with
oxygen at any temperature, neither does the heated metal absorb
this gas ; but it has the property, when cold, of condensing oxygen
upon its surface. A piece of clean platinum foil or wire, when
introduced into a mixture of oxygen, and a readily inflammable
gas or vapour (such as hydrogen, ether, alcohol, £c.), causes their
combination ; and occasionally the metal becomes red hot, and
ignites the mixture. This action is more rapid in the case of
platinum sponge, when a larger surface is brought into play, and
a fragment of this material introduced into a detonating mixture
of oxygen and hydrogen at once determines its explosion.

Platinum Dichloride 693

Platinum is not acted upon by either nitric or hydrochloric acid.
It is oxidised when fused with caustic alkalies, or with potassium
nitrate, and is also attacked by fused alkaline cyanides. In the
form of sponge, it is dissolved by boiling potassium cyanide, with
the evolution of hydrogen and formation of a double cyanide.

Platinum readily combines with phosphorus, silicon, and carbon.
The carbide of platinum is formed when the metal is continuously
heated by a smoky flame, or one in which combustion is incom-
plete, hence care is necessary in the use of platinum vessels.

Platinum Black is the name given to the finely divided metal
obtained by precipitating platinum from its solutions by reducing
agents or by metals. It is a soft, black powder, which is capable
of absorbing, or condensing upon its surface, large quantities of
oxygen. It therefore acts as a powerful oxidising agent.

Platinum Alloys. — Platinum readily alloys with many metals ;
hence compounds of easily reducible metals should not be heated
in vessels of platinum. The most important alloys are those with
iridium. The addition of 2 per cent of iridium is found greatly to
increase the hardness and raise the melting-point of platinum.
A.n alloy containing 10 per cent, of iridium resists the corrosive
action of chemical reagents to a greater extent than pure platinum
(see Fluorine, page 348).

Oxides Of Platinum.— Platinous oxide, PtO, and platinic oxide,
PtO2, are obtained in the form of dark grey or black powders by
gently heating the corresponding hydroxides. When strongly
heated they are converted into the metal.

Platinous Hydroxide, Pt(HO)2, is obtained by the action of
potassium hydroxide upon platinum dichloride. It is a black
powder, which dissolves in the halogen acids, yielding platmous
compounds.

Platinie Hydroxide, Pt(HO)4, is prepared by adding boiling
potassium hydroxide to a solution of platinum tetrachloride, and
treating the precipitate with acetic acid to remove the potash.
When dried it forms a yellowish powder, which is soluble in acids
to form platinic salts. Platinic hydroxide behaves both as a weak
base and a feeble acid. With stronger bases it forms compounds
known as platinates, which are yellow crystalline salts. The
sodium salt has the composition Na2O,3PtO2,6H2O.

Platinum Diehloride (platinous chloride), PtCl2, is produced
when platinum tetrachloride is heated to about 250°. It forms a
greenish powder, insoluble in water. It dissolves in hydrochloric

694 Inorganic Chemistry

acid, giving a reddish-brown solution which is believed to contain
the double compound PtCl2,2HCl, or H2PtCl4, to which the name
chloro-platinous acid has been given. The compound has never
been isolated, but a number of double salts of platinous chloride
with other chlorides are known, which may be regarded as
derivatives of this acid, and which are therefore termed chloro-
platinites ; thus, potassium platinous chloride, 2KCl,PtCl2, or
potassium chloro-platinite, K2PtCl4, is obtained as fine red crystals,
by adding potassium chloride to a solution of platinous chloride
in hydrochloric acid. This salt is used in the platinotype photo-
graphic process.

Platinum Tetraehloride (platinic chloride}, PtCl4, is obtained
by dissolving the metal in aqua regia, and removing the excess
of the acids by evaporating to dryness and gently heating the
residue. From its aqueous solution, the salt deposits in large
red crystals having the composition PtCl4,5H2O, which are not
deliquescent. When the salt is crystallised from a hydrochloric
acid solution, or when the aqua regia solution is evaporated to
expel the nitric acid, with frequent addition of hydrochloric acid,
the double compound of platinic chloride and hydrochloric acid is
formed, PtCl4,2HCl, which is deposited as reddish-brown deli-
quescent crystals, with 6H2O. To this substance (which is
commonly called platinic chloride), the name chloro-platinic acid
has been given, and the double salts of platinic chloride and
various chlorides are regarded as salts of this acid. The most
important of these chloro-platinates are those of the alkali metals,
their different solubilities being made the basis for the separation
of these metals.

Potassium Chloro-platinate (or potassium platinic chloride),
2KCl,PtCl4 or K2PtCl6, is obtained as a yellow crystalline pre-
cipitate by adding potassium chloride to platinic chloride. It is
soluble in 100 parts of water at the ordinary temperature to the
extent of i.i part, and at 100°, 5.18 parts. It is insoluble in alcohol.

The rubidium and caesium compounds are very similar, but are
still less soluble in water, 100 parts of water at 20° dissolving 0.141
of the rubidium and 0.07 of the caesium salt.

Ammonium Chloro-platinate, 2NH4Cl,PtCl4, closely resembles
the potassium salt, being slightly less soluble, but more so than
the rubidium compound.

Sodium Chloro-platinate, 2NaCl,PtCl4,6H2O, is a reddish^
yellow salt, readily soluble in both water and alcohol.

Platinamines 695

Platino-eyanides. — Just as platinous chloride combines with
metallic chlorides to form chloro-platinites, so platinous cyanide,
Pt(CN)2, unites with other cyanides, forming similarly constituted
double compounds, known as platino-cyanides.*

Potassium platino-cyanide, K2Pt(CN)4, or 2KCN,Pt(CN)2, is
formed when spongy platinum is dissolved in boiling potassium
cyanide. The platino-cyanides may be regarded as the salts of
platino-cyanic acid, H2Pt(CN)4. Both the acid and the salts are
characterised by the wonderful play of colours they exhibit when
viewed in different lights, and by forming different coloured
crystals with varying quantities of water of crystallisation (see
page 217).

Sulphides Of Platinum.— Platinous sulphide, PtS, and platinic
sulphide, PtS2, are obtained as amorphous black powders by the
action of sulphuretted hydrogen upon the respective chlorides.

OxysaltS Of Platinum.— Few well-defined single salts of
platinum with oxyacids are known. This element, however,
exhibits a great tendency to form complex double salts. One such
series of compounds is seen in the plaUno-nitrites, which may be
regarded as the salts of platino-nitrous acid, H2Pt(NO2)4.

These salts are remarkable, in that the platinum they contain
cannot be detected by the ordinary tests for that metal ; just as
the iron present in ferro-cyanides is not detected by the ordinary
reagents used in testing for that metal.

Ammoniacal Platinum Bases, or Platinamines.

Like cobalt, platinum forms a large number of basic compounds with
ammonia, many of which are of extremely complex composition. The first
of these to be discovered was a bright green salt, obtained by the action
of ammonia upon platinous chloride, having the composition PtCl2,2NH3, or
Pt(NH3).2Cl2, and known as the green salt of Magnus. Many of the platina-
mines exhibit isomerism ; thus, a compound known as the chloride of Reisefs
second base is a yellow crystalline salt having the same composition as Magnus's
green salt. Twelve distinct series of ammoniacal platinum compounds are
known, four of which are derived from platmous and the remainder from
platimV salts; the former are termed platoso ammonium compounds, while
the latter are distinguished as \\\^.platino compounds, f

* The name Cyano-platinites might with advantage be applied to these
compounds.

f For detailed descriptions of these compounds, the student is referred to
larger works on chemistry ; and on the constitution of these, and metallam-
monium compounds generally, the article by Werner, in the '/.eitschrift far
Anorganische Chemie, 1893, vol. iii. p. 267, may be consulted.

APPENDIX

RADIUM, AND RADIOACTIVE ELEMENTS

As far back as the year 1896, Becquerel discovered that the element uranium
and its salts possess the remarkable property of emitting rays somewhat
similar in character to the now familiar Rontgen or "X" rays; resembling
these rays in their penetrating power, their photographic action, and their
action upon electrified gases. These peculiar rays were distinguished from
the Rontgen rays by being called the "uranium," or the " Becquerel" rays.
Somewhat later it was found that the element thorium and its compounds
were likewise possessed of the property of emitting rays, which, while differing
from both the "X" and the "uranium" rays in some respects, closely
resembled them in others. To denote this property, the term radioactivity
has been coined, and substances possessing the property are said to be
radioactive bodies.

In 1898 it was announced that M. and Mme. Curie had discovered a new
radioactive substance contained in pitchblende, a mineral consisting essentially
of uranium oxide. From researches already made, it had been shown that
the radioactivity of uranium compounds is roughly proportional to the amount
of the metal present, but it was found that in the case of certain specimens of
pitchblende this was not the case, but that the radioactivity was greatly in
excess of that calculated from the percentage of uranium in the mineral. This
fact suggested the presence of some new substance of superior radioactivity to
that possessed by uranium. It was found in the ordinary process of separation
of the metals by precipitation from an acid solution by sulphuretted hydrogen,
that this new active substance was thrown down along with the sulphides, and
finally was separated from the copper and arsenic, &c. , and remained associ-
ated with the bismuth. No isolation of the new substance was effected, but
from its greatly superior radioactivity the discoverers concluded that there was
sufficient evidence of the presence of a new element to warrant them in giving
it a name. They therefore called it polonium, from the country from which
the pitchblende was obtained.* (Compt. rend. 127, p. 175.)

Following up their investigations, the same workers very shortly afterwards
discovered in the same mineral another radioactive body of still far greater
activity. This new substance, they found, is not precipitated by either
sulphuretted hydrogen, ammonium sulphide, or ammonia, but is associated

* Although the name polonium is still met with in the literature of the
subject, no further evidence has been produced in proof of the existence of a
new element corresponding to the name. The name is used rather to denote
the radioactivity which appears to be associated with the element bismuth.

697

698 Appendix

with and accompanies barium in the various chemical reactions the latter
element undergoes. Thus, when barium sulphate or carbonate is precipitated
from a solution of the chloride, the precipitated barium compound is accom-
panied by the radioactive material ; or when the chloride itself is precipitated
either by strong hydrochloric acid or by alcohol, the "aotive" substance is
thrown down along with it.

By the careful fractional precipitation of the chloride with alcohol it was
found possible to gradually concentrate the radioactive substance in the
barium chloride, and in this way a product was obtained possessing a radio-
activity 900 times greater than that of uranium. In view of the intensity
of its "activity," the discoverers gave the name radium to the new element
which they believed to be present, although in almost infinitely minute
quantities. (Compt. rend. 127, p. 1215.)

The spectrum exhibited by this "active" barium chloride also confirmed
the presence of a new element, for besides the lines belonging to barium it
contained a well-defined line which had never previously been observed in the
spectra of any of the known elements.

Determinations of the atomic weight of the metal (barium) in the speci-
mens of barium chloride which contained the radioactive element to an
extent sufficient to show an "activity" 900 times greater than that of
uranium, gave values practically the same as those of ordinary barium, namely
137.4. That is to say, the actual amount of radium which gave rise to so
high an "activity" in the barium chloride was too small to influence the
atomic weight determination. When, however, the concentration of the.
radium chloride in the barium chloride was considerably increased by a con-
tinuation of the fractionating process, the atomic weight of the metal was
found gradually to rise. Thus, when the intensity of the radioactivity reached
3000 times that of uranium, the atomic weight of the " barium " rose to 140 ;
while with a concentration representing a radioactivity 7500 times that of
uranium, the atomic weight of the metal present was found to be 145.8. From
these determinations it was evident that radium would be found to be an
element of very high atomic weight, and in the course of time when it became
possible to obtain small quantities of radium compounds— such as the chloride
and bromide — in a state of comparative purity, this was found to be the case.
The latest determinations which have been made by Mme. Curie and others,
have assigned the number 226.4 as tne atomic weight of this new element.

The metal itself was not isolated* until late in the year 1910, when Mme.
Curie obtained it by the electrolysis of a solution of radium chloride, employ-
ing mercury as the cathode. A liquid amalgam of radium was in this way
obtained, from which, by a process of very careful distillation, the mercury
was removed.

Radium is a bright white metal, melting at 700°, and more volatile than
barium. On exposure to air it turns black, and is believed to form a nitride.
It acts readily upon water, forming a soluble radium hydroxide.

The element radium closely resembles barium in its chemical relations.
Thus the sulphate is insoluble in water and in acids ; the carbonate is insoluble
in water, and the chloride is precipitated by both strong hydrochloric acid
and alcohol.

* In the literature of the subject the name radium is constantly employed
when in reality a radium salt is intended.

Appendix: 699

As seen in the Bunsen flame, the strongest and most permanent line pro-
duced by radium bromide is the blue line 4826.

The chief interest attaching to this new element is associated with the
strange property it possesses in such a high degree of emitting "radiations."
Radium bromide, for example, is self-luminous in the dark ; the rays it emits
are capable of acting upon a photographic plate, much as the Rontgen rays
affect it. They cause phosphorescence upon a screen of barium platino-
cyanide, and produce radiographic effects similar to those given by the "X"
rays. They are capable of penetrating metals, and will discharge an electro-
scope not only through considerable intervals of space, but also through
screens of various materials.

Most mysterious of all, they appear to possess the power of exciting a
temporary radioactivity in other substances otherwise inactive. Thus, if a
solution of a radium salt and some distilled water are placed in separate
dishes in a perfectly closed space, radioactivity is communicated to the water.
The water, however, gradually loses this power even in a closed space, while
it rapidly loses it if exposed. It has been found also that the intensity of this
"induced" radioactivity is the same for all substances, -under the same con-
ditions, irrespective of their chemical nature.

Concerning the nature and the cause of the radiations emitted by radium
and the other two well-defined radioactive elements uranium and thorium,* a
large amount of experimental work has been done, and much speculation put
forward. As the outcome of the former it has been established that at least
four distinct, and to some extent separable emissions, may go to make up what
is included in the term "radiations." These are distinguished as a, /S, and y
rays, and " radioactive emanation."

i. The a Rays. — These rays are very easily absorbed by thin layers of
matter. Thus, a thickness of aluminium 0.0005 cm. reduces their intensity
to one-half. To them is mainly attributable the property of causing the
ionisation of a gas, whereby its electrical conductivity is increased. They
are deviated by a very strong magnetic field, the deviation being in the
opposite direction to that exhibited by "cathode" rays. These a rays are
not waves like ordinary light rays, but consist of actual matter, which is being
projected at an enormous velocity, and is highly charged with positive electricity.
They are described as a "flight of material particles," having a mass of the
same order as the atoms of hydrogen, f and travelling with a velocity about
one-tenth that of light. J These particles carry with them a relatively enor-
mous amount of energy, each particle apparently having sufficient energy
associated with it to excite phosphorescence visible to the eye. Thus, Crookes
has shown that when a fragment of solid radium nitrate is brought near to a
screen of " Sidot's hexagonal blende " (zinc sulphide), and the phosphorescent
surface of the screen is examined with a pocket lens, it is seen to be dotted all

* Polonium, and the still more recent actinium, are at present too undefined
to be included as elements.

f i.e. , the ratio of the charge of the carrier to its mass is — = 6 x io3.

% That is, about 2.$x io9 cms. per sec. (Rutherford and Soddy, Phil. Mag.,
Feb. 1903).

7oo Appendix

over with brilliant specks of green light. In proportion as the radium salt
is brought closer to the screen, these flashes or scintillations become more
brilliant and more numerous, following each other with such rapidity that the
surface presents the appearance of a " turbulent luminous sea."

" It seems probable that we are here actually -witnessing the bom-
bardment of the screen by the electrons hurled off by the radium"
(Crookes).

2. The /3 Rays. — These rays are readily deviated by the magnetic field;
and also differ from the a rays in their greater penetrating powers. Thus,
while the intensity of the latter is reduced to one-half by passing through
0.0005 cm- °f aluminium, the /3 rays are able to traverse a thickness ot
0.05 cm. of this metal before their intensity is halved. A sheet of mica o.oi cm.
thick will completely absorb all the a rays, while it transmits the (3 and
also the y rays without appreciable diminution. The /3 rays, like the a
rays, also consist of projected particles with a high velocity, but in this case
they carry a negative electric charge, and their mass is believed to be greatly
less than that of the particles constituting the a rays, namely, about the T^^of
that of the hydrogen atom (Rutherford and Soddy, Phil. Mag., May 1903).
(3 rays are similar in all respects to the " cathode " rays emitted from a vacuum
tube, except that the velocity of the particles is greater and consequently they
are more penetrative. Their velocity is estimated to be between 2 x io10 and
3 x io10.

3. The <y Rays.— These are non-deviable by the magnetic field, and closely
resemble the Rontgen or "X'1 rays. They are far more penetrating than
either the a or /3 rays, being capable of penetrating a thickness of 8.0 cms.
of aluminium before their intensity is reduced to one-half. These rays are
believed to be a -wave motion, and not to consist of projected particles of
matter.

4. " Radioactive Emanation." — The elements thorium and radium * possess
the property of emitting something which has the power of imparting
radioactivity to any substance in their immediate neighbourhood. The
radioactivity thus imparted or excited is only of a temporary character,
its intensity diminishing and dying away when the substance is removed
from the influence of the original radioactive body. Experiments seem
to prove that these effects are not produced by any of the rays already
described, but are due to some other distinct emission, and the term
"radioactive emanation," or shortly "emanation," has been adopted to
denote this.

The radioactivity which is thus imparted to substances in the proximity
of these radioactive elements (usually spoken of as excited radioactivity) is
believed to be caused by the deposition upon their surface of radioactive
matter, which is transmitted by positively charged carriers ; while the radio-
activity of the " emanation " itself is believed to be due to the emission from
it of a rays only. When a small quantity of thorium oxide f is placed in a
tube (the oxide being enveloped in material capable of intercepting the

* Uranium appears not to share this property.

f Most of the earlier work by Rutherford and Soddy (Phil. Mag. , 1902) in
this connection was done with thorium compounds.

Appendix 701

ordinary radiations) and a stream of air is passed over it, the air is found
to carry with it the "emanation" which the thorium oxide gives out; and
the issuing stream of air, even after being conveyed through many feet of
tube, is capable of discharging an electroscope. In the case of radium
compounds the amount of this "emanation" was found to be comparatively
small when the radium compound is employed in the solid state, but when
the radium salt is dissolved in water, the "emanation" appears to ba given
off in a sudden rush, as it were, and the solution continues to emit this
"emanation" in amount many hundred times as great as was produced by
the solid salt. A similar enormous increase also takes place when the radium
compound is heated. These observations have led to the belief that the
"emanation " is actually occluded by the solid compound, f

In many other respects this "emanation" behaves like an inert gas. Thus
if the stream of air carrying the "emanation" is passed through a tube
plugged with cotton wool, nothing is arrested or filtered out by the wool and
the radioactivity of the air as it issues is not diminished. Neither is it affected
by the air being bubbled through strong sulphuric acid, or passed through a
red-hot platinum tube. When air conveying " emanation " is slowly passed
through a U-tube cooled by liquid air, the "emanation" is completely con-
densed, and the air which passes out is entirely free from all trace of this
substance. If a glass tube is employed, and the air current is sufficiently
slow, the progress of the condensation can be traced by the fluorescent
appearance of the glass, showing that the condensation has all taken place
upon the first portions of the tube traversed by the stream of air. If now the
tube is closed at both ends and the temperature allowed to rise above a
certain point, the condensed "emanation" appears to vaporise again, and
the fluorescence extends throughout the entire length of the tube. The
volatilisation point of radium emanation appears to be about - 150°, while
that of the thorium emanation is given as about —120° (Phil. Mag,, 1903,

P- 575).

In the case of thorium, the "emanation" loses its radioactivity, or decays,
much more rapidly than the radium emanation. Thus, while the activity
of thorium emanation falls to half its intensity in the space of one minute,
the intensity of the radium emanation only sinks to half its value in the space
of four days, while still retaining sufficient activity to be detected after the
lapse of one month. This rate of decay of the radioactivity of the
"emanation" is the same even at the low temperature of liquid air, and it
is considered probable that the marked difference in the rates of decay of the
"emanation" from thorium and radium may account for the difference
observed in their vaporisation temperatures.

It was at one time supposed that the radioactivity of these radioactive
elements was not a property intrinsic to the elements themselves, but was due
to the presence in small and varying quantity of some unknown substance.
Crookes found (Proc. Royal Soc., 1900) that by processes of a purely chemical
nature he was able to separate from uranium nitrate small quantities of
material which seemed to possess all the radioactivity, leaving the bulk
of the uranium compound inactive. He applied the name Uranium X to this

* Rutherford and Soddy, Phil. Mag., 1903, p. 449.

702 Appendix

"unknown uranium." Similarly in the case of thorium ; when the hydroxide
was precipitated by ammonia, and the nitrate (which chemically should contain
no thorium) was evaporated to dryness and ignited to expel ammonia salts,
minute residues were obtained which were many hundred times more active
than an equal weight of thorium oxide (Rutherford and Soddy, Phil. Mag.,
September 1902). The precipitated hydroxide, although not entirely robbed
of radioactivity, was found to have its activity greatly reduced. This sup-
posed "active " constituent was therefore called Th X. Later investigations,
however, revealed the remarkable fact that the thorium compound which had
thus been partially deprived of its radioactivity gradually regained it when
left to itself ; white the separated Th X gradually lost it. Moreover, it was
found that the two processes went on exactly at the same rate, that the rate
of decay of the activity of Th X was the same as the rate of recovery of activity
of original thorium compound. From this it would appear that two opposing
processes are simultaneously going forward in a radioactive substance, namely,
the continual production of fresh radioactive material and the constant decay
of the radiating power of the active material. In other words, what may
be called the normal radioactivity is a condition of equilibrium, where the
rate of increase of activity due to the production of fresh active material
balances the rate of the decay of the activity in the radioactive material
already formed.

The views now generally held are that the phenomena of radioactivity are
due to atomic changes, but changes of a character altogether different from
any that have previously been dealt with in chemistry. It is believed that the
atoms of these radioactive elements (which, it will be noted, are possessed of
the highest atomic weights of all the elements) are undergoing a -process of
disintegration or degradation : that in the course of their movements, owing
to some combination of conditions about which at present we know nothing,
the kinetic energy of some of the atoms reaches a point beyond which the
stability of the atom is no longer possible. Under these circumstances the
atom breaks up, throwing off some matter from itself, and assumes a more
stable configuration. The particles or fragments of the original atoms them-
selves undergo further changes, giving off other particles, thus giving rise to
the various phenomena of radioactivity.

In the case of radium the fact seems now to be thoroughly established that
one of the final products of the radioactive change is the element helium. It
has also been clearly proved that it is the a-particles themselves which are
identical with atoms of helium. Thus it is found that when a quantity of
radium emanation is enclosed in a tube of extremely thin glass which is sur-
rounded by a vacuum jacket, the a-particles expelled from the emanation
penetrate the thin glass into the vacuous space ; and if, after the lapse of some
time, the contents of this space are swept out by mercury into a minute spec-
trum tube, the gas thus collected will give the characteristic spectrum of helium.

It is also now generally believed that radium itself is a product of the
radioactivity of uranium ; not necessarily a first product, but probably through
one or more intermediate stages.

If these are the true interpretation of the phenomena of radioactivity we are
undoubtedly face to face with an actual instance of the ' ' transmutation of ihe

Appendix 703

elements," which has hitherto been regarded only as an idle dream of the
alchemist. It may, indeed, be that in these radioactive processes we have as
it were a peep into the unknown region of the " evolution of the elements." *

The energy which is liberated during this process of atomic disintegration is
enormous, taking into account the minute quantities of matter concerned.
M. and Mme. Curie have shown that a sample of a radium salt gave out
energy sufficient to melt half its own weight of ice per hour. This energy,
which is stored up in the atoms of these elements, the ' ' internal energy of the
chemical atom," as it has been termed, and which is set free during radioactive
change, is of an entirely different order of magnitude from that which is dis-
engaged during any processes of ordinary chemical change. It has been
calculated, indeed, that the energy of radioactive change is many thousand
times, or even a million times, as great as that of any known chemical
change, when equal weights of matter are concerned. Ramsay has recently
shown f that this energy is capable of bringing about certain ordinary chemical
changes. Thus radium "emanation" is able both to decompose water into
its elements, and also to cause the recombination of oxygen and hydrogen.
It is always found, however, that the mixed gases resulting from the action of
the "emanation" upon water, contain a slight excess of hydrogen, the exact
reason of which is at present unknown.

The idea of an atom as a system, and, moreover, one capable of under-
going changes into simpler systems, is a view which, to the chemist, may at
first seem strangely heterodox, and one altogether opposed to fundamental
doctrines of chemistry. In reality, however, this new view as to the con-
stitution of an atom does not touch the question of the indivisibility of the
atom in the purely chemical sense. From this point of view the chemical
atom still retains its position as the lowest stage in the complexity of matter,
and may still be defined as the smallest particle of matter which can take
part in a chemical change. The chemical atoms of these radioactive elements
are not divisible into what may be called " chemical fragments." If the atom
is a system, then in all chemical reactions and changes the system in its
entirety takes part. When it is borne in mind that the weight of matter
which the atom, regarded as a changing system, throws off in the form of
"radiations," "emanation," or "electrons" is so infinitely minute, that it has
been estimated that it would require many hundreds, if not thousands of
years before enough of it could be collected to be detected by the most deli-
cate balance, it will be evident that we are dealing with phenomena of a totally
different order from those in which the relative weights of matter entering into
chemical combination are concerned.

* A comprehensive theory of the evolution and devolution of the elements,
by Jessop, is to be found in the Phil. Mag. , Jan. 1908.
t Jour. Chem. Soc.t 1907.

INDEX

ABSOLUTE boiling-point, 79

Acid, m

,, temperature, 70

Absorptiometer (Bunsen), 144

Absorption of gases by charcoal, 292

, ,

Acetylene, 317

Acetylide of copper, 318

Acid, ammon-sulphonic, 282

, antimonic, 497

, arsenic, 488

, arsenious, 487

, boric, 609

, bromic, 383

, carbamic, 311, 444

, carbonic, 305

chloric, 374

, chloro-auric, 569

( (

, chlorochromic, 664

• i '

, chloroplatinic, 694

ii <

, chloroplatinous, 694

.i <

, chlorosulphuric, 439

ii <

, chlorosulphonic, 439

ii <

, chromic, 662

dithionic, 437

ii <

, hydrazoic, 279

M 1

, hydriodic, 389

•. ]

, hydrobromic, 381

» ]

, hydrochloric, 363

M ]

, hydrofluoboric, 612

M ]

, hydrofluoric, 350

'- 1

, hydrofluosilicic, 632

" 1

, hydrosulphurous, 423

» 1

, hypobromous, 383

» ]

, hypochlorous, 373

» ]

hypoiodous, 395

•- ]

, hyponitrous, 250

I. I

, hypophosphorous, 472

>. ]

, hyposulphuric, 437

» I

, hyposulphurous, 423

» I

, iodic, 392

" I

manganic, 660
metaboric, 610
metantimonic, 497
metaphosphoric, 476
metarsenic, 488
metasilicic, 635
metastannic, 640
metatungstic, 665
metavanadic, 656
molybdic, 664
muriatic, 371
nitric, 234
nitrosulphuric, 426
nitrous, 244

Nordhausen sulphuric, 434
ortho-antimonic, 496
ortho-arsenic, 488
ortho-arsenious, 487
orthoboric, 609
orthophosphoric, 474
orthosilicic, 635
osmic, 691
oxymuriatic, 352
pentathionic, 438
perchloric, 375
perchromic, 660
periodic, 393
permanganic, 669
persulphuric, 424
phosphomolybdic, 665
phosphoric, 474
phosphoric (glacial), 476
phosphorous, 473
pyrc-antimonic, 496
pyro-arsenic, 488
pyro-arsenious, 487
pyroboric, 610
pyrophosphamic, 478
pyrophosphodiamic, 478
2 Y

Index

Acid, pyrophosphoric, 475

,, pyrophosphotriamic, 478

,, pyrosulphuric, 434

,, pyrovanadic, 656

,, selenic, 447

,, selenious, 447

,, silicic, 635

,, stannic, 640

,, sulphovinic, 315

,, sulphuric, 425
sulphurous, 421

,, telluric, 449

,, tellurous, 449

,, tetrathionic, 438

,, thiocarbamic, 444

,, thiocarbonic, 443

,, thiosulphuric, 435

,, trithionic, 437

,, tungstic, 664
Acid-forming oxides, 17
Acids, dibasic, 19

,, mono-, tetra-, and tribasic, 19
Active mass, 94
Affinities, 61

Affinity, chemical, 10, 61
After-damp, 298
Air-liquefiers, 77
Alabaster, 581
Algin, 386
Alkali manufacture, 534

,, metals, 505
Alkali-waste, 400
Alkaline earths, 571
Allotropy, 194
Aludels, 386, 598
Alum, 620

,, burnt, 622
meal, 621

,, shale, 621

,, stone, 621
Alumina, 617
Aluminates, 618
Aluminite, 619
Aluminium, 614

,, alloys, 617

bronze, 553, 617

,, chloride, 623

,, fluoride, 622

,, hydroxides, 618

,, sodium chloride, 623

Aluminium sulphate, 619
,, sulphide, 617, 623
Alums, 620
Alunite, 621

Amalgamation process (silver), 559
Amalgams, 6co
American pot-ashes, 521
Amethyst, 617
Ammonia, 272

,, solubility of, in water, 275
Ammonia-soda process, 538
Ammoniacal cobalt compounds, 685
,, liquor, 319
,, mercury compounds, 604
platinum compounds, 695
Ammonium, 545
,, alum, 620
amalgam, 545
bprofluoride, 612
,, carbamate, 311, 547

carbonate, 547
,, chloride, 546

,, dissociation of, 89
,, chloroplatinate, 694

chromate, 230
,, cyanate, 14, 24
,, ferrous sulphate, 680
hydrazoate, 280
hypoiodite, 283
,, iron alum, 620
,, magnesium arsenate, 488
,, magnesium phosphate, 475

manganous chloride, 668
,, meta-thio-arsenate, 490
metavanadate, 656
molybdate, 665
nitrate, 248
,, nitrite, 230

phosphomolybdate, 477, 665
,, plumbic chloride, 650

pyro-arsenite, 487
,, pyro-thio-arsenite, 490
,, salts, 545
,, sesquicarbonate, 548
,, sodium phosphate, 475
,, stannic chloride, 642
,, sulphate, 546
,, thiocyanate, 548
Ammon-sulphonates, 282
Amorphous silicon, 629

Index

707

Analysis, 13
Anastase, 627
Anglesite, 643
Anhydrides, 17
Anhydrite, 577
Animal charcoal, 290
Anions, 99, 105
Anodes, 97 •
Anthracite, 293
Antimonates, 497
Antimonious oxide, 496
Antimony, 491

,, amorphous, 492

, , blende, 491

,, bloom, 491

,, compounds with halogens,

494

chlorides, 495
hydride ,493
ochre, 491

oxides and oxyacids, 496
oxychlorides, 495
sulphides, 498
sulpho-trichloride, 496
tetroxide, 497
trioxide, 496
Apatite, 347, 582
Apollinaris water, 220, 305
Aquafortis, 234
Aqua regia, 241

Aqueous vapour (atmospheric), 257
Argentic compounds (see Silver),

558

Argentiferous lead, 560
Argentite, 558
Argon, 265

Argon group of gases, the, 263
Arragonite, 582
Arsenates, 488
Arsenic, 478

,, allotropic modifications of,

480

,, chlorhydroxide, 484
,, chloride, 483
,, compounds with halogens,

483

,, fluoride, 483
,, hydride, 480
„ oxides and oxyacids, 484
, , pentoxide, 487
„ sulphides, 489

Arsenical iron, 479

,, pyrites, 479
Arsenious bromide, 484

,, iodide, 484

,, oxide, 484
Arsenites , 487

Arsenuretted hydrogen, 480
Arsine, 480
Asbestos, 573
Asymmetric system, 162
Atacamite, 555
Atmolysis, 84
Atmosphere, 252

,, composition of, 256

, , height of, 263

,, suspended impurities in, 261
Atmospheric ammonia, 258

,, aqueous vapour, 257

,, carbon dioxide, 257

,, gases mechanically mixed, 260

,, ,, argon group of, 263

' ,, hydrogen, 260
nitric acid, 258

,, ozone, 259
Atomic electric charge, 104
"heat, 46

,, theory, 25

,, volumes, 44, 119

,, weight, definitions of, 37, 44

,, weight, determination of, by
chemical methods, 36, 58

,, weight, determination of, by
means of isomorphism, 51

,, weight, determination of, by
means of specific heat, 45

,, weight, determination of .from
volumetric relations, 38

,, weights, list of, 22

,, ,, international, 22, 38

Atoms, 4
Aurates, 569
Auric chloride, 569

,, oxide, 569
Auro-auric sulphide, 569
Aurous iodide, 569
Autogenous soldering, 430
Avogadro's hypothesis, 40
Axes of symmetry, 160
Azoimide, 279
Azote, 229
Azurite, 550

Index

BALANCED actions, 88
Balling furnace, 536
Barium, 586

,, amalgam, 586
,, bromate, 384
,, carbonate, 586
chlorate, 374
chloride, 588
dioxide, 184, 587
dithionate, 437
hydroxide, 587
hypophosphite, 472
iodate, 586
monoxide, 586
nitrate, 589
oxides, 586
peroxide, 587
sulphate, 588
sulphide, 589
tetrathionate, 438
thiosulphate, 438
Baryta, 586

,, water, 587
Barytocalcite, 586
Base, 18
Basic oxides, 17
,, salts, 20

Basicity of acids, the, 18, 473
Battery, galvanic, 96
Bauxite, 614
Beryl, 572
Berylla, 572
Beryllium, 572

,, aluminate, 618
,, compounds, 572
,, specific heat of, 48
,, Bessemer process (steel), 675
Binary compounds, 15
Bismuth, 500
,, alloys, 500
,, carbonate, 502
,, compounds with halogens,

,, dichloride, 501

,, dioxide, 502

,, glance, 500

,, nitrate, 503

,, nitrate, basic, 503

,, ochre, 500
oxides, 501

,, oxychloride, 503

Bismuth pentoxide, 504
tetroxide, 503

,, tribromide, 501

,, trichloride, 501

,, tri-iodide, 501

,, trioxide, 502

,, trisulphide, 504
Bismuthic oxide, 501
Bismuthous oxide, 501
Bisulphate of soda, 422
Bittern, 531
Bituminous coal, 293
Black ash, composition of, 537

,, furnace, 535

,, revolving furnace, 536
Black-band, 672
Black-jack, 591
Blacklead, 288
Blast-furnace, 673
Bleaching-powder, 187, 373, 580
Blister copper, 551 .

,, steel, 675
Blue vitriol, 556
Boiling-point, absolute, 79

,, definition of, 128

,, molecular elevation of, 134
Boiling-points, 129

,, effect of pressure upon, 129

,, effect of dissolved substances
upon, 134

,, of saturated saline solutions,

133

Bolognian phosphorus, 584
Bone ash, 452

,, black, 290
Bones, composition of, 291
Boracite, 607
Borate spar, 607
Borates, 610
Borax, 610
Borofluorides, 612
Boron, 607

hydride, 613

,, nitride, 613

,, sulphide, 613

,, trichloride, 612

,, trifluoride, 611

,, trioxide, 6c8
Boronatrocalcite, 6oy
Bort, 285
Boyle, law of, 71

Index

709

Brass, 553
Braunite, 666

Erin's process (oxygen), 184
Britannia metal, 492, 639
Broggerite, 264
Bromates, 384
Bromides, 382
Bromine, 377

,, electrolytic manufacturing

process, 380
hydrate, 380

,, monochloride, 396

,, oxyacids, 383

,, water, 380
Bromous acid, 383
Bronze, 639
Brookite, 627
Brown haematite, 672
Brown iron ore, 672
Brucite, 574
Bunsen flame, the, 341

,, non-luminosity of, 342

,, temperature of, 343
Burnt alum, 622

CADMIUM, 596

,, chloride, 597
,, oxide, 597
,, sulphide, 597
Caesium, 505, 544

,, spectrum of, 510
Cailletet's apparatus, 75
Calamine, 591
Calcined magnesia, 574
Calcite, 576
Calcium, 577

bicarbonate, 22T, 311, 582

borate, 611

borofluoride, 612

carbide, 319, 455

carbonate, 582

chlorate, 517

chloride, 579

chloro-hypochlorite, 580

dioxide, 579

fluoride, 347

hydride, 578

hydroxide, 578

hypochlorite, 517, 580

manganite, 359, 668

oxides, 578

Calcium phosphate, 452, 582
,, phosphide, 460

,, sulphate, 581

,, sulphide, 400, 410, 583
Calc-spar, 577
Caliche, 387
Calomel, 602
Calorie, 165, 326
Calx, 252, 321
Candle-flame, 335
Canton's phosphorus, 584
Capillary pyrites, 689
Carat, definition of, 569
Carbides, 295

,, of iron, 513
Carbon, 285, 627

,, compounds, 295

,, dioxide, 300

,, ,, atmospheric, 2^7

,, ,, composition of, 307

,, ,, solid, 307

,, disulphide, 441

,, hydrogen, compounds of, 312

,, monoxide, 296

,, oxides of, 296

,, specific heat of, 47

,, suboxide, 311
Carbonado, 285
Carbonates, 310
Carbonyl chloride, 299
Carbonyls, metallic, 299
Carborundum, 630
Carboxy-haemoglobin, 298
Carnallite, 512, 579
Caro's acid, 425
Carry's freezing-machine, 132
Cassiterite, 637
Cast iron , 674
Catalysis, 183

Catalytic action, 12, 183, 354
Cathodes, 97
Cations, 99, 105
Caustic potash, 515

,, soda, 530
Celestine, 584
Cellulose, 240

Cementation process (steel), 67?
Centres of symmetry, 160
Cerite, 627
Cerium, 627
Cerussite, 643

7io

Index

Chalcedony, 628

Chalk, 582

Chalybeate waters, 220

Chamber acid, 429

Chamber crystals, 426

Chance's process 411

Change of volume on solidification,

137

Charcoal, 290

,, absorption of gases by, 292

,, animal, 290

specific heat of, 47
Charles' law, 69
Chemical action, n

,, affinity, 10

,, combination, laws of, 25

,, equations, 23

,, formulae, 23

,, > modes of, 13

,, nomenclature, 15

,, notation, quantitative, 53

,, reactions, 23

,, symbols, 22
Chili saltpetre, 541
Chlorates, 375
Chloride ions, 104
Chloride of lime, 580
Chlorine, 352

,, heptoxide, 373

,, hydrate, 362

,, liquefaction of, 73, 362

,, liquid, 362

, , manufacturing processes, 354

,, monoxide, 371

, , oxides and oxyacids, 371

,, peroxide, 372

» water, 361
Chloro-aurates, 569
Chloro-chromates, 664
Chloro-stannates, 642
Chromates, 662
Chrome alum, 66 1

,, green, 658

,, iron ore, 657

,, ochre, 657

,, red, 663

,, yellow, 663 •
Chromic anhydride, 659

,, chloride, 660

,, hydroxides, 658

,, sulphate, 66 1
Chromite, 657

Chromites, 661
Chromium, 657

,, anhydride, 659

,, chromate, 658
dioxide, 658

,, oxides of, 658

,, sesquioxide, 658

,, trioxide, 659
Chromous chloride, 660

,, hydrated oxide, 658

,, sulphate, 660
Chromyl chloride, 663
Chrysoberyl, 572, 618
Cinnabar, 597
Clark's process for softening water,

222

Classification of elements, 112
Clay, 614

,, ironstone, 672
Cleveite, 264
Coal, 293

., gas, 319

Coarse metal (copper), 551
Cobalt, 682

,, bloom, 682

,, glance, 479, 682

,, oxides of, 682
Cobaltamines, 685
Cobaltic hydroxide, 683

,, oxide, 683
Cobalto-cobaltic oxide, 683
Cobaltous chloride, 683

,, hydroxide, 683

,, oxide, 683

,, sulphate, 684

,, sulphide, 684
Coefficient of absorption, 144

,, solubility, 144
Coefficients of expansion of gases, 69
Coke, 290
Colemanite, 607
Colloids, 636
Columbite, 655
Columbium, 655
Combining proportions, 30
Combustibles, 322
Combustion, 321

,, gain in weight by, 325

,, heat of, 326

, , supporters of, 322
Common salt, 531
Compound radicals, 23

Index

711

Compounds, 7
Conductivity, molecular, 107
Condy's fluid, 670
Constant-boiling mixtures, 155
Constant-freezing solution, 155
Constant composition, law of, 25, 31
Constitution of matter, 3
Contact process, sulphuric acid, 432
Copper, 550

,, acetylide, 318

„ alloys, 553
arsenite, 487

,, bromide, 555

,, carbide, 318
carbonates, 557
chlorides, 555

,, ferrocyanide, 156

,, fluoride, 555
glance, 550
hydride, 473

,, hydroxide, 554
-,„ nitrate, 556

,, nitroxyl, 244
oxides, 553

,, oxychloride, 555

,, pyrites, 550

,, sulphate, 556

,, sulphides, 557
Coprolites, 582
Coral, 577
Corpse light, 330
Corrosive sublimate, 603
Corundum, 614
Cream of tartar, 497, 512
Crith, 56
Critical pressure, 79

„• temperature, 78, 133
Croceo-cobaltic salts, 686
Crocoisite, 657
Crookesite, 624
Cryohydric solutions, 155
Cryolite, 347, 614
Crystalline forms, 160
Crystallisation, suspended, 137, 151

,, water of, 216
Crystalloids, 636
Cubic system, 161
Cubical nitre, 541
Cupel, 560
Cupellation process (silver), 560

Cupric carbonates, 557
,, chloride, 555
,, hydroxide, 554
,, nitrate, 556
,, oxide, 554
, , sulphate, 556
sulphide, 557
Cuprous acetylide, 318
,, chloride, 555 A
oxide, 553
sulphide, 557
Cyanide process (gold), 567

DALTON, atomic theory, 30
Davy lamp, 330
Deacon's process, 354
Dead Sea, solid matter in, 219
Deep well waters, 220
Deliquescence, 217
Dephlogistigated air, 181

,, marine acid air, 352
Dew-point, 257
Diamidogen, 278
Diatomic molecules, 8
Dialysed iron, 678
Dialysis, 635
Diamond, 285

,, combustion of, 287

,, specific heat of, 47
Diffusiometer, 82
Diffusion of gases, 81

law of, 83

,, of dissolved substances,

159

Dimorphism, 162
Dissociation, 88

,, coefficient, 108

,, electrolytic, 96

,, pressure, 94
Disulphates, 435
Disulphur dichloride, 413
Disulphuryl chloride, 440
Dithionates, 437
Divalent elements, 59
Dolomite, 572 ,
Dry copper, 551
Dulong and Petit, law of, 46
Dutch brass, 553

,, metal, 360
Dyad elements, 59

712

Index

EARTH'S crust, composition of, 182

Ebullition, 129

Efflorescence, 217

Effusion of gases, 85

Eka-aluminium, 123

Eka-boron, 123

Eka-silicon, 123

Electric furnace, 454, 615

Electro-chemical equivalents, 100

Electro-gilding, 568

Electrolysis, 96

Electrolytes, 97

Electrolytic dissociation, 96, 101

Electrons, 104

Electroplating, 99, 563

Elements and compounds, 6

,, classification of, 112

,, list of, 22

,, non-metallic, 8
Elton Lake, water or, 219
Emerald, 572
Emery, 617
Empyreal air, 181
Endosmometer, 155
Endosmose, 155
Endothermic compounds, 168
English brass, 553

,, Channel, composition of, 219
Epsom salts, 575
Equations, chemical, 23
Equivalents, chemical, 30

,» electro-chemical, 100
Estramadurite, 452, 582
Ethyl hydrogen sulphate, 315

,, silicate, 631
Ethylene, 314

,, dibromide, 314
Euchlorine, 372
Eudiometry, 252
Evaporation, 126

,, cold produced by, 78, 130, 276
Exothermic compounds, 168, 331
Expansion by heat of liquid carbon

dioxide, 308
Expansion by heat of liquid oxygen,

193

Extincteur, 303

FARADAY'S law, 100
Felspar, 614, 637

Ferrates, 679
Ferric chloride, 680

,, ferrocyanide, 681

,, hydroxide, 678

,, ,, soluble, 678

,, oxide, 678

,, sulphate, 680

,, sulphide, 681
Ferrites, 677
Ferro-manganese, 674
Ferroso-ferric oxide, 678

,, sulphide, 682
Ferrous bromide, 520

,, chloride, 679

,, chromite, 662

,, ferricyanide, 680

,, ferrocyanide, 680

,, hydroxide, 677

,, oxide, 677

,, sulphate, 679

,, sulphide, 681
Fettling, 675
Fine metal (copper), 551
Fire-damp, 314
Fire-damp caps, 331
Fixed air, 300
Fixed alkali, 506
Flame, 332

candle, 335

,, the Bunsen, 341

,, structure of, 332
Flames cause of luminosity of

338

Flint, 628

Flintshire furnace, 644
Fluorapatite, 347
Fluorides, 350
Fluorine, 346
Fluor-plumbates, 348
Fluor-spar, 347, 577
Forces, chemical and physical, 3
Formula weight, 54
Formulas, 23

Fraction of dissociation, the, 91
Franklinite, 591
Fulminating gold, 569

,, silver, 564
Fusco-cobaltic salts, 685
Fusible metal, 500
Fusion, latent heat of, 138

Index

713

GADOLINITE, 606

Gahnite, 591

Galena, 558, 643

Gallium, 124, 606

Galvanised iron, 593

Gas carbon, 289

Gases, absorption by charcoal, 292

,, coefficients of expansion of, 69

, , critical pressure, 79

critical temperature of, 79

, , diffusion of, 81
effusion of, 85

, , kinetic theory of, 85

, , liquefaction of, 72

,, occlusion of, 179

, , relation to heat, 69

, , relation to pressure, 71

,, solubility of, in liquids, 142

,, transpiration of, 85
Gastric juice, 534
Gay-Lussac, law of, 26, 38
General properties of gases, 69

,, liquids, 126
German silver, 593
Germanium, 124, 627
Gilding, 568
Glauberite, 541
Glauber's salt, 541
Glucinum, 572
Gold, 567

,, alloys, 568

compounds of, 569

,, fineness of,. 568

,, fulminating, 569
Graduators, 532
Graham's law, 83
Gramme-molecule, 57
Graphite, 287

,, specific heat of, 47
Greenockite, 596
" Green salt of Magnus," 695
Green vitriol, 434, 679
Grey antimony ore, 498

,, cast iron, 674
Guignet's green, 659
Gun-cotton, 240
Gun-metal, 553
Gunpowder, 523

, , products of combustion of, 524
Gypsum, 581

,, fibrous, 581

HEMATITE, 672
Haemoglobin, 193
Hair salt, 619
Half-electrolytes, 97
Halogens, 18, 345
Haloid salts, 18
Hardness (water), 221
Hargreaves' process, 539
Hausmannite, 666
Heat, atomic, 46
,, molecular, 49
,, of combustion, 326
,, of formation, 167

specific, 45

,, specific, table of, 46
,, units, 165
Heavy spar, 586
Helium, 267
Henry's law, 143
Hepar sulphuris, 526
Hexagonal system, 161
Holmes's signal, 464
Horn mercury, 602
Horn silver, 565
Hydrazine, 278

,, hydrochloride, 279
,, hydrate, 279
,, sulphate, 278
Hydrocarbons, 312
Hydrofluosilicic acid, 630
Hydrogen, 171

,, atmospheric, 260

,, chloride, 363

,, compounds with oxygen,

203

,, dioxide, 223
,, disodium phosphate, 475
,, displaceable, 19
,, liquid, 178
,, monoxide, 203
,, nitrate, 19
,, occlusion of, 171, 179
,, peroxide, 224
,, persulphide, 412
,, phosphide, gaseous, 460
,, ,, liquid, 463

,, ,, solid, 464

,, position of, in the periodic

classification, 125
,, potassium fluoride, 347
M •> sulphate, 434

2 Y 2

Hydrogen sodium ammonium phos-
phate, 475
sulphate, 19

,, sulphide, 408

,, telluride, 448

,, trisulphide, 413
Hydrogenium, 179
Hydrolysis, in
Hydromagnesite, 576
Hydroxides, 17
Hydroxyl, 281
Hydroxylamine, 281

,, disulphonate, 281

,, hydrochloride, 281

,, mono-sulphonate, 282
Hypobismuthic oxide, 501
Hypobismuthous oxide, 501
Hypochlorites, 373
Hypochlorous anhydride, 371
Hypoiodous acid, 395
Hyponitrous anhydride, 248
Hypophosphites, 472
Hypovanadic chloride, 656

,, oxide, 656

,, sulphate, 656

ICE, 214

,, effect of pressure upon, 214

, , the melting-point of, 138
Icicle, 132
Ignition-point, 329
Indigo-copper, 557
Indium, 122, 606
Inflammable air, 171
International atomic weights, 22, 38
Intestinal gases, hydrogen in, 171
lodates, 393
lodic anhydride, 391
Iodine, 384

,, bromides, 396

,, chlorides, 396

,, pentoxide, 391
Ionic notation, 105

,, theory, the, 101
lonisation, 107
Ions, 99

migration of, 108
Indium, 690

,, chlorides, 691

,, oxides, 691

Irish Sea, solrd impurity in, ,219
Iron, 671

Index

Iron alum, 68 1

,, carbide, 513, 674

,, carbonyl, 300

,, magnetic oxide of, 678

,, monoxide, 677

, , oxides of, 677
passive, 677

,, pyrites, 400, 672

,, sesquioxide, 678

,, sesquisulphide, 681

,, sulphides of, 681
Isodimorphism, 162
Isogonism, 52
Isomerism, 194
Isometric system, 161
Isomorphism, 51, 162

, , law of, 51

JOLLY'S apparatus, 255

KAINITE, 512
Kelp, 385 .
Kelp substitute, 386
Kiesel-guhr, 633
Kieserite, 572, 575
Kinetic theory, 85
Kish, 288
Krypton, 269
Kupfernickel, 479, 687

LAGOONS (boric acid), 609

Lakes, 618

Laminaria digitata, 385
,, stenophylla, 385

Lamp-black, 289

Lanarkite, 643

Lanthanum, 606

Latent heat effusion, 138

,, ,, vaporisation, 130

Laughing-gas, 248

Law of Boyle, 71

Charles, 69
,, constant heat consummation,

169

,, constant proportion, 25, 26
,, Duiong and Petit, 45
,, gaseous diffusion, 83
,, Gay-Lussac, 26, 38
,, multiple proportionst 25,

27
-, octaves 113

Index

Law of partial pressures, 147
,, periodic, 112
,, reciprocal proportions, 25, 28
Layer crystals, 52
Lead, 643

,, acetate, 652

,, action of water upon, 646

carbonate, 651
, , chromate , 663
,, composition of commercial, 647
,, desilverisation of, 561
,, dichloride, 649
dioxide, 649
disulphate, 653
ethide, 628
nitrate, 651
oxides of, 647
,, oxychloride, 650
,, sesquioxide, 648
, , softening of, 645
,, squirted, 647
,, suboxide, 647
sulphate, 653
sulphide, 654
,, sulphochlorides, 654
,, tetracetate, 654
,, tetrachloride, 650
,, tree, 645
,, white, 651
Leblanc process, 534
Leguminous plants, 258
Lepidolite, 543
Light red silver ore, 558
Lime, 578

chloride of, 580
,, dead burnt, 578

milk of, 579
,, quick, 578

slaked, 578

,, superphosphate of, 583
Limestone, 577
Lines of symmetry, 160
Liquefaction of air, 77-

,, of gases, 72

Liquids, general properties of, 126
Liquor ammonia, 275
Litharge, 647
Lithium, 543

,, carbonate, 544
,, hydroxide, 544
„ mica, 543

Lithium nitride, 232

oxide, 544

phosphate, 544
,, spectrum of, 509
Liver of sulphur, 526
Load-stone, 672, 678
Lothar Meyer's curve, 120
Lucifer matches, 459
Luminous paint, 584
Lunar caustic, 566
Luteo-cobaltic salts, 686

MAGISTRAL, 559
Magnesia, 574
Magnesia alba levis, 576
,, ponderosa, 577

usta, 574

Magnesia mixture, 575
Magnesian limestone, 220, 577
Magnesite, 576
Magnesium, 572

,, aluminate, 618
,, ammonium chloride, 575
,, ammonium phosphate, 475
,, boride, 613
bromate, 384
calcium chloride, 575
carbonates, 576
chloride, 574

combustion of, in steam, 174
hydroxide, 574
nitride, 232, 574
oxide, 574
oxychloride, 575
phosphate, 475
platinocyanide, 217
, , potassium chloride, 57
,, pyrophosphate, 476
,, silicide, 629
,, sulphate, 575
Magnetic iron ore, 672
,, oxide of iron, 678
,, pyrites, 682
Magnetite, 678
Malachite, 550
Manganates, 669
Manganese, 666
,, blende, 666
,, dioxide, 667
,, monoxide, 666
,, oxides of, 666

7i6

Manganese sesquioxide, 667

,, spar, 666
Manganic chloride, 668

,, oxide, 667

,, sulphate, 668
Manganite, 666
Manganites, 668
Mangano-manganic oxide, 667
Manganous chloride, 668

,, chromite, 662

,, hydroxide, 667

,, sulphate, 668
Marble, 582
Marine acid air, 363
Marsh gas, 312

,, synthesis of, 443
Marsh's test, 482
Massicot, 647
Matches, 459
Matlockite, 643
Mechanical mixtures, 8
Mediterranean Sea, 219
Meerschaum, 573
MendelejefFs periodic law, 113
Mephitic air, 229
Mercuric ammonium chloride, 605

,, chloride, 603

,, iodide, 603

,j, oxide, 601

., potassium chloride, 67
Mercurius calcinatus per se, 181
Mercurous chloride, 602

,, nitrate, 602

,, oxide, 601

,, sulphate, 602
Mercury, 597

„ alloys of (amalgams), 600

,, deadening of, 600

,; distillation of, 598

, , oxides of, 601
Metal slag (copper), 551
Metallic carbonyls, 299

,, nitroxyls, 244
Metalloids, 8
Metals and non-metals, 7
Metameric compounds, 194
Metantimonates, 497
Metaphosphates, 476, 498
Metarsenates, 488, 498
Metarsenites, 487

Index

Metastannates, 641

Metavanadates, 655

Meteoric iron, 171

Methane, 312

Meyer, Lothar, curve of atomic

volumes, 120
Microcosmic salt, 543
Migration of ions, 108
Milk of lime, 579

,, sulphur, 407
Milky quartz, 634
Mineral alkali, 506
Minium, 648
Mispickel, 479
Mixed crystals, 52
Modes of chemical action, 13
Molecular combinations, 67

,, conductivity, 107

,, concentration, 94

,, depression of the freezing,
point, 140

, , elevation of the boiling-point,
134

,, equations, 55
formulae, 23
heats, 49

,, lowering of vapour pressure,

,, volume, 44

,, weight, 41 -^

,, weight, determination of, by
the depression of freezing,
point, 140
Molecules, 3

,, compound, 6

,, definition of, 4

,, elementary, 6

,, mean free path of, 86

,, size of, 3
Molybdates, 664
Molybdenite, 664
Molybdenum, 664

,, chlorides, 665

,, ochre, 664

,, oxides, 664
Monad elements, 59
Mono-atomic molecules, 8
Monoclinic system, 162
Monosymmetric system, 162
Monovalent elements, 59

Index

717

Mordants, 618
Mortar, 579

the setting of, 579
Mosaic gold, 643
Mottramite, 655
Mundic, 479
Muntz metal, 553
Multiple proportions, law of, 25, 31

NATURAL waters, 218.
Natural steel, 624
Neon, 269

Nessler's solution, 272, 605
Neutral alum, 622
Nickel, 687

alloys of, 688
blende, 687

,, car bony 1, 299
chloride, 689
glance, 479, 687
monosulphide, 689

,, monoxide, 688
oxides of, 688

,, sesquioxide, 688
. ,, silver, 593

,, sulphate, 689
Nickelo-nickelic oxide, 689
Nickelous oxide, 688

,, sulphide, 689
Niobates, 655
Niobium, 655

,, oxides of, 655
Nitrates, 241

,, detection of, 241
Nitre, 522

,, plantations, 523
Nitric acid, manufacture from atmos-
pheric nitrogen, 235

„ anhydride, 241

,, oxide, 246
Nitrides, 278 280
Nitrification, 522
Nitrites, 245
Nitro-cellulose, 240
Nitrogen, 229

,, iodide, 283

oxides and oxyacids of, 234

,, oxyfluorides, 251

,, pentoxide, 241

,, peroxide, 242

,, tribromide, 283

,, trichloride, 282
Nitro-metals, 244

Nitro-sulphuric acid, 426
Nitrosyl chloride, 250

,, fluoride, 251

,, hydrogen sulphate, 251

,, perchlorate, 251

,, sulphate, 426
Nitroxyl fluoride, 251
Nitrous anhydride, 234
Noble metals, 240
Nomenclature, 15

,, of ions, 105
Non-electrolytes, 97
Non-metals, 7
" Nordhausen " acid, 434
Notation, chemical, 21, 53

OCCLUDED hydrogen, 179

Occlusion of gases, 171

Olefiant gas, 314

Opal, 633

Ore hearth, 644

Orangeite, 627

Organic chemistry, definition of, 296

Orpiment, 479

Orthite, 606

Orthoclase, 637

Orthorhombic system, 161

Osmiridium, 690

Osmium, 690

,, oxides of, 691

,, tetroxide, 691
Osmotic pressure, 155
Osteolite, 582
Oxides, 17
Oxygen, 181

,, allotropic, 195

,, Brin's process, 184

,, Tessie" du Motay process, 189
Oxyhcemoglobin, 193
Oxyhydrogen flame, 327
Oxymuriatic acid, 352
Ozone, 195

,, atmospheric, 256

,, constitution of, 199

,, tube, Siemens', 195

,, ,, Andrews', 200

PALLADIUM, 690

,, absorption of hydrogen by,

179

, , chlorides, 691
,, hydride, 179
,, oxides, 691

718

Index

Parkes's process, 561
Partial pressures, law of, 147
Partially miscible liquids, 149
Passive iron, 677
Pattinson's process, 561

,, white lead, 650
Pearl-ash, 521
Perchlorates, 375
Percy-Patera process, 562
Perdisulphuric acid, 424
Periclase, 574
Peridote, 637
Periodates, 393
Periodic classification, 112
Permanent white, 589

,, hardness, 221
Permanganates, 669
Permanganic anhydride, 670
Permonosulphuric acid, 425
Permutit, 222
Persulphates, 425
Persulphuric anhydride, 424
Petalite, 543
Petzite, 567
Pewter, 639
Phenacite, 572
Phlogiston, 321
Phosgene gas, 299
Phospham, 477
Phosphates, 474
Phosphine, 460
Phosphites, 473
Phosphonium bromide, 462

,, chloride, 462

,, iodide, 462
Phosphoretted hydrogen, gaseous,

460

Phosphorous oxide, 470
Phosphorus, 451

,, allctropic, 458

,, compounds with sulphur, 478

,, manufacture of, 452

,, ,, by electric furnace,

454

,, oxides and oxyacids, 469

„ oxychloride, 468

,, oxyfluoride, 468

,, pentabromide, 467

,, pentachloride, 465

„ pentafluoride, 464

„ pentasulphide, 478

Phosphorus pentoxide, 471

„ red, 458

,, tetr iodide, 467

,, tribromide, 467

,, trichloride, 465

,, trifluoride, 464

,, triodide, 467
Phosphoryl chloride, 468

,, fluoride, 468
nitride, 478

,, triamide, 478
Photo-salts, 566
Physical constants of gases, 80
Pig-boiling, 675
Pig iron, 674
Pitchblende, 664
Planes of symmetry, 160
Plaster of Paris, 581
Plastic sulphur, 406
Plate sulphate, 386
Platinamines, 695
Platinates, 693
Platinic hydroxide, 693

,, chloride, 694
Platiniridium, 690
Platino-chlorides, 693

,, cyanides, 695

,, nitrites, 695
Platinotype process, 694
Platinous chloride, 693

,, hydroxide, 693
Platinum, 691

,, alloys, 693

,, black, 693
oxides of, 693

,, oxysalts, 695

,, sodium chloride, 67
spongy, 692

,, sulphides of, 695

,, tetrachloride, 694 '..
Platoso-ammonium compounds, 695
Plumbago, 288
Plumbic chloride, 649

,, oxalate, 647

,, oxide, 647

,, peroxide, 649
Plumbous oxide, 647
Plumbum nigrum, 643
Pollux, 510
Polybasite, 558
Polyhalite, 520

Index

719

Polymerism, 194
Pot-ashes, 521
Potash, caustic, 515
Potassium, 510

,, alum, 621

,, aluminate, 618

,, antimonate, 498

,, borofluoride, 608, 612

,, bromate, 384, 520

,, bromide, 520

,, carbonate, 521

,,• chlorate, 516

,, ,, electrolytic manu-

facture of, 518

,, chloride, 516

,, chlorochromate, 664

,, chloroplatinate, 694

,, chloroplatinite, 694

,, chr ornate, 662

,, chromium alum, 620, 661

,, dichromate, 662

,, ferrate, 679

,, ferricyanide, 681

,, ferrocyanide, 297, 680

,, fluoride, 515

fluor-plurnbate, 348

,, hydride, 514
hydroxide, 515

,, hypoiodite, 283

,, hyponitrite, 250

,, iodate, 393

,, iodide, 520

„ manganate, 669

,, metaborate, 610

,, metantimonate, 498

,, metarsenite, 487

,, metastannate, 641

,, meta-thio-arsenite, 490

,, nitrate, 522

,, nitrite, 245

,, osmate, 691

,, oxides of, 514

,, ortho-thio-antimonate, 499

,, ortho-thio-antimonite, 499

,, ortho-thio-arsenate, 490

,, ortho-thio-arsenite, 490

,, pentasulphide, 525

., pentathionate, 439

,, perchlorate, 519

,, periodate, 394

Potassium permanganate, 670
,, peroxide, 514
,, platinic chloride, 694
,, platino-cyanide, 695
,, platinous chloride, 694
,, plumbate, 649

pyro-antimonate, 498
,, ruthenate, 691
,, silico-fluoride, 629
,, silver thiosulphate, 437
,, stannate, 640
,, sulphate, 520
sulphite, 421
,, sulphides of, 525
,, tetrachromate, 663
,, trichromate, 663
,, zinc oxide, 175

Powder of Algaroth, 495

Praseo-cobaltic salts, 685

Preparing salt, 640

Producer gas, 186

Proustite, 558

Prussian blue, 681

Pseudo-alums, 620

Pucherite, 655

Puddling, 675

Purple copper ore, 550

Purpureo-cobaltic salts, 686

Pyrargyrite, 558

Pyrites burners, 428

Pyrolusite, 666

Pyromorphite, 643

Pyrophosphates, 476

Pyrosulphuric chloride, 440

QUANTITATIVE notation, 53
Quartz, 633
Quicklime, 578

RADIATED pyrites, 681

Radicals, compound, 23

Radium, 697

Rain water, solid impurity in, 220

Raoult's method, 140

Realgar, 489

Red antimony, 491

,, copper ore, 553

,, haematite, 672

,, lead, 648

,, manganese oxide, 667

Index

Red phosphorus, 458

,, zinc ore, 591
Refinery slag (copper), 551
Regular system, 161
Reiset's second base, chloride of, 695
Relation of gases to heat, 69

,, ,, ,, pressure, 71

Reversible reactions, 88
Rhodium, 690
Rhpmbic system, 161
Rochelle salt, 563
Rock crystal, 633
Rock salt, 526
Rodonda phosphates, 452
Roll sulphur, 403
Roman alum, 621
Roseo-cobaltic salts, 686
Rouge, 678
Rubidium, 544
Rubies, artificial, 618
Ruby, 617
Ruby ore, 550
„ silver ore, 558
,, sulphur, 479
Rust of iron, 676, 678
Ruthenium, 691

,, chlorides of, 691

,, oxides, 691
Rutile, 627

SAL alembroth, 603

,, ammonia, 546
Salt-cake process, 534
Salt-forming oxides, 17
Salterns, 531
Saltpetre, 522
Salts, acid, 19

,, basic, 20

,, haloid, 18

,, normal, 19

,, of hydrogen, 107

,, of hydroxyl, 107

,, oxy-, 18

,, thio-, 18
Sand, 628
Sapphire, 617
Satinspar, 581
Saturated solutions, 151

,, vapours, 127
Scandium, 606

Scheele's green, 487
Scheelinite, 664
Schlippe's salt, 499
Schonite, 576
Schweinfurt green, 487
Scotch hearth, 644
Seaweed, iodine in, 385
Selenite, 581
Selenium, 444

,, alums, 620

,, dichloride, 446

dioxide, 447

Selenuretted hydrogen, 446
Seltzer water, 220
Semipermeable membranes, 156
Serpentine, 573, 637
Siemens' ozone tube, 195
Silica, 633
Silicates, 636

Siliciuretted hydrogen, 630
Silicon, 628

,, chloride, 632

,, chloroform, 628

,, dioxide, 633
fluoride, 632

,, hexachloride, 633

,, hexafluoride, 632

,, hydride, 630

,, liquid, 631
Silver, 558

,, allotropic, 563
alloys, 563

,, alum, 566

,, bromide, 565

,, chloride, 564

,, flashing of, 560

,, fluoride, 565

,, fulminating, 564
glance, 558
iodide, 565

,, nitrate, 566
oxides, 563

,, oxybromide, 566

,, oxychloride, 566

,, periodate, 394

,, phosphates, 475, 477
plating, $63

,, spitting of, 563

,, standards, 563

,, suboxide, 564

Index

721

Silver sulphate, 522, 566

,, sulphide, 558
Slaked lime, 578
Smalt, 686
Smaltine, 682
Smoky quartz, 634
Soda, 540
Soda-ash, 538

caustic, 530
,, crystals, 540
Sodium, 526

,, acetate, 313

,, alloy with potassium, 529

,, aluminate, 614

,, aluminium chloride, 615

,, amalgam, 601

„ antimonate, 497

,, antimonite, 496

,, arsenate, 488

,, benzoate, 279

,, bicarbonate, 540

,, bromide, 534

,, carbonate, 534

,, ,, electrolytic manu-

facture "of, 539

,, chloride, 531

,, chloro-platinate, 694

,, electrolytic manufacture of,
526

J( electrolytic manufacture of
(Borchers' process), 527

,, hydrazoate, 279

„ hydride, 529

,, hydroxide, 530

,, hypophosphite, 473

„ hyposulphite, 436

„ iodide, 534

,, metabisulphite, 424

,, metaniobate, 655

,, meta phosphate, 476

,, metastannate, 641

,, metatantalate, 655

,, metavanadate, 655
nitrate, 541

,, nitride, 280
oxalate, 175
oxides, 529

,, permanganate, 670

,, phosphates, 542

Sodium pyro-arsenate, 488
pyrophosphate, 476
sesquicar bonate, 540
silicate, 635
silver thiosulphate, 562
stannite, 640
sulphate, 541

solubility curve, 153
sulphide, 534
thio-antimonate, 499
thiosulphate, 436
tungstate, 664
uranate, 664
zinc chloride, 66
Soffioni, 609
Solar prominences, 171
Solder, 639
Solfatara, 399

Solidification, suspended, 137, 404, 456
Solidifying points of liquids, 137

points of liquids, effect of dis-
solved substances upon, 139
,, points of liquids, effect of

pressure on, 137
Solubilities, diagram of, 152
Solubility of gases in liquids, 142
,, of liquids in liquids, 148
, , of mixed gases, 146
,, of solids in liquids, 150
Solution, 142
Solutions, saturated, 151

,, supersaturated, 151
Sombrerite, 452, 582
Spathic iron ore, 672
Specific gravity of gases, 40

,, ,, liquids and solids, 119

,, heat, 45
,, heats, tables, 46
Spectra of alkali metals, 505
Spectroscope, 507
Specular iron ore, 672
Speiss-cobalt, 682
Spiegel, 674
Spinelle, 618
Spirits of hartshorn, 272
Spitting of silver, the, 563
Spodumene, 543
Spring water, 219, 220
Stalactites, 222
Stalagmites, 222

722

Index

Standard temperature and pressure,

69, 71

Stannates, 640
Stannic chloride, 642
,, sulphide, 642
Stannous chloride, 641

hydrated oxide, 639
,, nitrate, 638
,, oxide, 639

oxychloride, 641
,, sulphate, 638
,, sulphide, 642
Stassfurt deposits, 512, 520, 574
Steam, 214

,, volume, composition of, 208
Steel, 675
Steel mill, 329
Stephanite, 558
Stereotype metal, 492
Stibnite, 498

Still-liquor, composition of, 357
Stream-tin, 637
Stromeyerite, 558
Strontia, 584
Strontianite, 584
Strontium, 584

,, ammonium, 584
„ chloride, 585
,, dioxide, 585
,, hydride, 584
,, hydroxide, 584

nitrate, 585
,, oxides, 584
,, sulphate, 585
Substitution, 382
Suint, 511
Sulphates, 434
Sulphides, 410
Sulphion, 105
Sulphites, 421
Sulpho-acids, 17
Sulpho-thionyl chloride, 414
Sulphovinic acid, 315
Sulphur, 398

allotropic modifications, 404

chlorides of, 413

dioxide, 415

flowers of, 402

milk of, 407

oxides and oxyacids of, 414

oxychlorides of, 439

perfluoride, 441

Sulphur, plastic, 406
,, prismatic, 404
, recovery of, from alkali- waste,

400
, , recovery of (Chance's process),

411

rhombic, 404
,, sesquioxide, 423
,, tetrachloride, 414
,, trioxide, 421
Sulphuretted hydrogen, 408
Sulphuric acid, contact process, 432
,, ,, manufacture of, 428
Sulphuric anhydride, 414
,, chlorhydrate, 440
Sulphurous anhydride, 414
Sulphuryl chloride, 439
Supercooling of water, 136
Superphosphate of lime, 583
Supersaturated solutions, 151
Suspended solidification, 137, 404, 456
Sylvanite, 567
Sylvine, 512
Symbols, 21 .
Sympathetic inks, 217
Synthesis, 13

TACHYDRITE, 575
Talc, 573
Tank liquor, 537
Tantalite, 655
Tantalum, 655

,, oxides of, 655
Tartar emetic, 497
Tellurates, 449
Telluretted hydrogen, 448
Tellurites, 449
Tellurium, 448
Temporary hardness, 221
Tenorite, 554

Tessi6 du Motay process, 189
Tetradymite, 448
Tetragonal system, 161
Tetratomic molecules, 8
Tetravalent elements, 59
Thallic chloride, 625

,, nitrate1, 626

,, oxide, 607, 625

, , sulphate, 626

,, sulphide, 607
Thallium, 623

,5 oxides of, 624

Index

723

Thallium oxy hydroxide, 625

,, perchlorate, 607

,, sulphate, 607
Thallous carbonate, 626

,, chloride, 625

,, hydroxide, 624

,, iodide, 607

,, oxide, 624

,, phosphate, 626
The"nardite, 541
Thermochemistry, 163
Thio-acids, 17
Thio-antimonates, 499
Thio-antimonites, 499
Thio-arsenates, 490
Thio-arsenites, 490
Thiocarbonates, 443
Thionyl chloride, 439
1 hiophosphoryl chlorides 469

,, fluoride, 468
Thorite, 627
Thorium, 627
Tin, 637

,, alloys of, 639

,, dioxide, 640

,, oxides of, 639

,, oxy muriate, 642
Tin-pi at&, 639
Tin-stone, 637
Tin- white cobalt, 479
Tincal, 607
Tinning, 639
Titanium, 627
Tombac, 553

Transitional elements, 115, 671
Transpiration of gases, 85
Triad elements, 59
Triclinic system, 162
Tridymite, 633
Triethylamine, 150
Triethyl silico-formate, 631
Trivalent elements, 59
Trona, 540

Truncated crystals, 162
Tungstates, 664
Tungsten, 664

,, chlorides, 665

,, oxides, 664
Turnbull's blue, 6bo
Turpeth mineral, 434

Turquoise, 614
Type metal, 492
Typical elements, 115
Twin crystals, 634

ULEXITE, 607

Unit of heat, 165, 326

,, volume, 44
Unsaturated compounds, 62
Uranates, 664
Uraninite, 267
Uranium, 664

,, chlorides, 665

,, oxides, 664
Uranous salts, 665

,, sulphate, 665
Uranyl salts, 665
Urea, 13, 24, 295

VALENCY, 59
Vanadates, 655
Vanadite, 655
Vanadium, 655

,, chlorides of, 656

,, oxides of, 655

,, oxychlorides of , 656
Vaporisation, latent heat of, 130
Vapour densities of elements, 42

,, pressures of solutions, 133
Vapour tension, 128
Verdigris, 557
Vermilion, 604
Vinasse, 521

,, cinder, 521
Vital force, 295
Vitriol chambers, 430
Volatile alkali, 506

WATER, 203

Clark's process for softening,

222

colour of, 212
,, compressibility of, 213

electrolysis of, 207
,, freezing of, 131
,, gas, 297
. ,, gravimetric composition oft

210
,, hardness of, 221

maximum density of; 214

724

Index

Water of constitution, 218

Wood's fusible metal, 500

,, of crystallisation, 216

Wrought iron, 675

,, rain, 220

Wulfenite, 664

,, solubility of gases in, 147

Wurtzite, 595

i) » » salts in, 150

,, solvent power of, 216

XANTHO-COBALTIC salts, 686

supercooling of, 136

Xenon, 269

,, volumetric composition of,

206

YTTERBITE, 606

Waters, chalybeate, 220

Ytterbium, 606

dangerous, 223

Yttrium, 606

deep well, 220

,, fresh, 220

ZIERVOGEL process, 561

,, hard, 221

Zinc, 591

mineral, 219

alloys of, 593

,', natural, 218

,, aluminate, 591

potable, 222

amalgam, 60 1

,, river, 220

,, blende, 591

,, safe, 223

,, carbonate, 596

,, sea, 219

chloride, 594

spring, 219

chromite, 662

,, suspicious, 223

granulated, 174

Wavellite, 452 s

,, hydroxide, 594

Weldon's process, 357

methyl, 313

Welsbach burner, 339

,, nitrate, 240

White arsenic, 485

oxide, 593

cast iron, 674

,, spar, 591

,, lead, 651

spinnelle, 591

,, metal (copper), 551

,, sulphate, 595

nickel, 687

,, sulphide, 595

,, vitriol, 218

,, white, 594

Witherite, 586

Z,inci carbonas, 5g6

Wohlerite, 627

Zinc-copper couple, 173, 313

Wolfram, 664

Zircon, 627

„ ochre, 664

Zirconium, 627

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