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CMji.^x^<^ 3 0-0 Co . 3

Harvard College
Library

FROM THE BEQ.UEST OF

Daniel Treadwell

RUMFORD PROFESSOR AND LECTURER

ON THE APPLICATION OF SCIENCE

TO THE USEFUL ARTS

1834-1845

SCIENCE CENTER LIBRARY

A B C D E r

Solar Spectruin

Fig. 33. Various spectra.
See page J 07,

FIRST PRINCIPLES

CHEMISTRY

RAYMOND B. |ROWNLEE

FAR ROCKAWAY HIGH SCHOOL

WILLIAM J. HANCOCK

ERASMUS HALL HIGH SCHOOL

ROBERT W. FULLER

STUYVESANT HIGH SCHOOL

MICHAEL D. SOHON

MORRIS HIGH SCHOOL

JESSE E. WHITSIT

DE WITT CLINTON HIGH SCHOOL

ALL OF NEW YORK CITY

Boston

ALLYN AND BACON

1907

C-<'V.'i.VN. "i old , i

Tjx.jL<xJjuirdLL -^<yv^

'h
^

PREFACE

In the early spring of 1904, the authors were appointed a
joint committee of two organizations of New York science
teachers, to prepare a syllabus for a first course in Chem-
istry. The fruit of their labor during the succeeding year was
adopted in substance by the Department of Education as the
syllabus for New York State. In planning this outline, the
guiding principles were to make the experimental evidence
precede the chemical theory, to follow as far as feasible the
historical order in the development of the theory, and to
emphasize the practical aspects of the science.

As there was no existing text-book written from the point
of view contemplated by the syllabus, the preparation of the
present book was begun in September, 1905. During the writ-
ing of the book, the authors have been constantly in close
touch with one another, and the book is in every part the joint
product of their efforts.

In selecting their material, the authors have been governed
wholly by what they considered of intrinsic value to the
elementary student, without reference to its traditional place
in a text-book. This has led to the omission of some subjects
commonly found in books for beginners. To the subjects
selected they have striven to give a discussion simple enough
to be readily comprehended by the beginner, and complete
enough to furnish him with a clear idea of the underlying
principles of Chemistry and a definite knowledge of its more
important facts.

To this end, the experimental determination of chemical
facts is emphasized from the first. When sufficient facts have
been given to make explanation necessary, the generalizations

IV PREFACE

of the science have been introduced. In some of the theoreti-
cal chapters, particularly those on solution and ionization,
it may be advisable to omit certain portions at first and to
take them up afterwards as need arises. The authors have
attempted to bring out the fundamental principles first by a
simple statement, which is later developed and driven home
by illustrations, exercises, and problems, all designed to stimu-
late the pupil to think for himself, and constantly to connect
his new facts with the facts and principles already learned.

In order to give the pupil some idea of the great commercial
importance of Chemistry, a number of typical manufacturing
processes have been described and illustrated. Where a sub-
stance is manufactured in sever.al ways, the authors have en-
deavored to avoid confusion by giving a description of one
process only, selecting the one which they believe is, or will
become, most extensively used in this country. For example,
the time-honored chamber process for the manufacture of
sulphuric acid has been omitted, because it is giving way to
the contact process. The commercial production of copper,
aluminum, iron, and carborundum have been described some-
what in detail, as they are notable examples of modern chemi-
cal processes. ,

The thanks of the authors are very gratefully given to all
those who have so generously assisted them in securing de-
scriptions and illustrations of chemical processes as they are
actually carried on. Mr. Allen B. Doggett has rendered great
assistance in photographic work. We are especially indebted
to Mr. C. D. McArthur, for the drawing and description of the
lampblack process; to our students, for the execution of many of
the line drawings; to the Carborundum Company, the American
Smelting and Eefining Company, the National Lead Company,
and to the Scientific Americany for assistance which they have
rendered.

New Yoek, August, 1907.

II.

III.

IV.

VI.

VII.

VIII.

IX.

XI.

XII.

xin.

XIV.

XV.

XVI.

XVII.

XVIII.

XIX.

XX.

XXI.

CONTENTS.

Introduction . . .

Gases and their Measurement .

Oxygen .....

Hydrogen

Composition of Water and Combining Weights

Water and Solution .

Atoms and Molecules

Chlorine

Hydrochloric Acid .

Molecular Composition

Atomic and Molecular Weights ^

Symbols and Formulas

Chemical Equations .

Sodium and Potassium

Solution

Sodium and Potassium Compounds

Sulphur and Hydrogen Sulphide

Oxides ana Acids of Sulphur ""7 ~
Nitrogen and the Atmosphere .
Nitrogen Compounds
Elements of the Nitrogen Group

PAGE
1

103

109

128

141

164

167

179

200

CONTENTS.

OHAPTKB

XXII.

PAex

The Halogens 208

XXIII.

Carbon

. 227

XXIV.

Oxides of Carbon ....

. 244

XXV.

Silicon and Boron .

. 255

XXVI.

Calcium and its Compounds .

. 266

XXVII.

Magnesium, Zinc, and Mercury

. 278

XXVIII.

Copper and its Compounds .

. 291

XXIX.

Silver, Gold, and Platinum .

. 304

XXX.

Aluminum and its Compounds

. 317

XXXI.

Iron, Cobalt, and Nickel

. 328

XXXII.

Tin and Lead ....

. 350

XXXIII.

Manganese and Chromium

. 360

XXXIV.

The Periodic Law .

, .365

XXXV.

Compounds of Carbon .

. 374

Appendix
I.

Physical Constants of the Elements . .396

II.

Table of Solubilities

398

III.

Greneral Eules for Solubility .

399

IV.

Volatility of Compounds

399

Weight of One Liter of Common Gases

399

VI.

Pressure of Water Vapor

400

VII.

The Kinetic Theory of Gasea.

401

Ikdbx

405

PRINCIPLES OF CHEMISTRY.

CHAPTER I.
nrTEODUOTIOF.

1. Phjnncal Change. — When we notice the things about
us, we see that they undergo changes : a piece of wood
bends under a weight, or warps when wet ; a rod lengthens
when heated ; a piece of iron placed near a magnet at-
tracts another piece of iron. If we remove the weight
from the stick, it straightens ; the iron removed from the
magnet loses its power of attraction. In such changes,
although the object may be considerably altered, we still
recognize the pieces of the stick as wood, as we do the
fragments of a broken tumbler as glass; that is, the
material has not lost or changed those peculiar properties
or the characteristics by which we identify it. Such
changes are called physical changes; they result usually
in a change of such properties as size, shape, or color.

2. Chemical Change. — Another kind of change is seen in
the burning of wood. Iron in rusting falls to a red powder ;
mortar and cement change from a plastic to a stone-like
condition ; fruits and vegetables decay ; meat spoils ;
milk sours ; fruit juices ferment. In all these changes
the materials have apparently lost their identity, and sub-
stances with new characteristics are formed. We do not

2 INTRODUCTION.

find any resemblance between iron and rust or between
wood and what remains after it is burned. The iron and
the wood have apparently disappeared or changed and we
have a new substance in their places. Such alterations
are called chemical changes. Chemistry is the name given
to the science which has grown out of the study of chem-
ical changes and the effort to control or modify them.

The most striking chemical change that goes on about
us is that which occurs when a substance burns. If the
burning substance is coal or wood, it seems to disap-
pear except for a small quantity of ashes. In studying
this change, several questions suggest themselves. What
has happened to produce the heat that is given off?
What has become of the great mass of substance that has
apparently disappeared? Why does the substance burn
more brightly when the air is blown on it ? Why does it
stop burning when the air is shut off, as in smothering the
flame ? If we find answers to these questions, we shall
have explained the phenomena of burning.

Early investigators did not discover the true nature of
this process. They adopted a wrong explanation, and, as
a result, chemistry, as a science, did not advance. A true
explanation was stated only a little more than a hundred
years ago. Since that time chemistry has progressed at
an astonishing rate. The true explanation of burning was
found by studying the change that many metals undergo
when heated in air.

3. Heating Metals in Air. — A few metals — for example,
magnesium — burn when heated in the air. Most metals
undergo a similar, but much slower, change, without the
production of light. If a piece of bright copper is heated,
it assumes a black color ; on bending or scraping it gently

LAVOISIER'S EXPERIMENT. 8

a black powder separates from it. If the metal is heated
again, another layer of the black substance forms. By re-
peating the process a sufficient number of times, the piece
of copper can be entirely changed into the black powder.
Since the black powder in no way resembles the copper
from which it was prepared, a chemical change has taken
place. Iron heated in a similar manner gives a somewhat
similar result. Lead after melting gives a yellowish
powder ; zinc, if in the form of- a powder, will take fire
and give a white powder. Metals, in general, when thus
heated give powdered substances which bear no resem-
blance to the original metal.

Several things may be thought of as explanations of the
change which these metals undergo. It may be that by
the effect of the heat, without the aid of any substance,
the metal is transformed into a new kind of matter ; pos-
sibly the metal in being heated has lost some of its sub-
stance, which has passed off as gas ; possibly the metal has
absorbed something from the air.

As an aid in testing these possible explanations, it will
be advisable to weigh the metal before and after it is
heated. When this is done, it will be found that the pow-
der alwat/8 weighs more than the metal from which it was
formed. This seems to indicate that during the heating
the metal adds to itself more substance and that this sub-
stance is taken from the air. To further test this conclu-
sion, a piece of metal can be sealed in a glass tube from
which the air has been exhausted; heated under these con-
ditions^ the metal is not changed.

4. Lavoisier's Experiment. — Another conclusive experi-

.ment is one that was performed by Lavoisier, the French

chemist, to whom is given the credit of discovering the

4 INTRODUCTION.

nature of this kind of chemical change. He put some tin
in a good-sized glass flask and sealed it so that the air
could neither enter nor leave it. He then heated the flask
carefully for several days. At the end of this time he
noticed that a certain amount" of white powder had been
formed. He next ascertained that the flask with its con-
tents had not changed in weight. He then opened the
neck of the flask and noticed that air rushed in. On
again weighing the flask he found that it had increased in
weight, and that this increase was equal to the increase
which the tin had undergone on being converted into the
white powder. He explained these facts as follows: the
tin on being heated combined with some of the air in
the flask, producing the white powder. The flask as a
whole did not increase in weight because no air entered
the flask to take the place of that which had combined
with the tin. When the flask was opened, the air entered,
causing the increase in weight.

From these facts it appears probable that the powdered
substances are more complex than the metals from which
they are formed; that is, they contain the metal plus
something which has been taken from the air. Lavoisier
undertook to find out the nature of the substance which
was taken from the air.

5. Heating Mercury in Air. — Mercury, heated in the
air, undergoes a much less rapid change than the metals of
which we have been speaking. By keeping it at a tem-
perature near its boiling-point for several days, a small
quantity of a red powder is gradually formed (Fig. 1). If
a quantity of this powder is heated in a glass tube to a
temperature somewhat higher than that at which it was ^
formed, a colorless gas is given off. If a glowing splinter

LAVOISIER'S EXPERIMENT. 6

is inserted into the tube, it bursts into flame and burns
brilliantly. The gas cannot be ordinary air, for a splinter
does not behave so in air. A considerable quantity of the

Fig. 1.
Lavoisier*s apparatus for heating mercury. ^

gas can be collected and shown to be very different from
ordinary air, in the fact that substances burn in it with
extraordinary vigor.

On examining the tube it will be found that a part or
all of the red powder has disappeared and that drops of
mercury have collected on the sides of the tube. It
appears from this experiment that the red powder has
decomposed into mercury and a gas which readily supports
combustion. Lavoisier named this gas oxygen.

Since the red powder was made by heating mercury in
the air, and was not formed unless air was present, the
oxygen must have come from the air. Hence air must con-
tain oxygen. That air is not all oxygen is shown by the
fact that only about one-fifth, and not all, of the air was

6 INTRODUCTION.

absorbed in Lavoisier's experiments; and also by the fact
that substances do not burn as readily in air as in oxygen.

6. Burning. — The burning of wood or other substances
is a process that closely resembles the change of a metal
into a compound of the metal and oxygen. In the case of
ordinary combustible materials, the products are chiefly
gases which pass off unseen. By the use of suitable appa-
ratus the products formed in the burning of a candle can
be collected, and it is found that their weight is greater
than the weight of the candle burned. As in the case of
the metals, this increase in weight is due to the oxygen
taken up from the air. If air is excluded, the burning
substance is extinguished because it can no longer combine
with oxygen. On account of the tendency of the sub-
stances to unite, energy is evolved in the form of heat.

7. Compomids and Elements. — We have shown that the
red substance contains oxygen and mercury. The sub-
stance formed on heating copper in the air contains oxygen
and copper. As these substances formed are composed of
more than one kind of material, they are called compounds.
No one has changed copper or mercury into anything else,
without adding something. So far as we know, gold, iron,
oxygen, and about eighty other things are not composed of
anything else. They are simple, or, as we say, elements.

Summary.

Chemical changes involve change in the identity of the material.
The composition of the substance is usually altered, and energy
changes are also involved. •

A compound is a substance that can be separated into two or
more substances. An element is a substance which has not been
separated into other substances.

EXERCISES. 7

When a substance bums in air, it combines with oxygen, fonn-
ing a new compound.

Lavoisier obtained oxygen from air by heating mercury in it
and then decomposing the material obtained.

Bxercises.

1. Air and water were formerly called elements ; why are
they not now ?

2. How could you prove that air contains oxygen ?

3. What kind of change is involved in the withering of a
leaf? making cloth from wool? baking bread? burning
coal ? extinguishing the coal ?

*. Why is a candle extinguished by blowing ?

CHAPTER II.

GASES AITD THTJTK MEASTTBEHENT.

8. Solids, Liquids, and Oases. — Matter is generally de-
fined as anything that occupies space. The different
kinds of matter are called svistances. Substances differ
in the way they fill space, and it is this difference that
determines their physical state. The three physical states
of matter are the solid, the liquid, and the gaseous.

A solid has a definite shape or form^ and therefore a
definite volume. A liquid has no definite form, but has,
however, a definite volume. It can fill a vessel only to
the extent of its volume and takes the shape of the con-
taining vessel so far as it fills it. Q-a%eB have neither a
definite form nor a definite volume. They tend to dis-
tribute themselves in all directions and fill completely
any vessel into which they are brought. Their only
boundaries are the containing walls.

9. Effect of Temperatnre and Fressnre Changes on. Volumes
of Oases. — The measurement of the volume of gases
usually involves a correction of the gas volume. This is
necessary because the volume of a given quantity of gas
is considerably affected by even slight changes in tem-
perature and pressure. If the room gets warmer, the
volume will be larger ; if it gets colder, the volume will
be less. Changes in atmospheric pressure will also cause
the volume to vary. An increased pressure will mean a

CHABLM8' LAW.

OENXIGLRADB

ABSCCDTB

^" ofWatet ^^*

80- --BoonS/Demperatiue - -j

Freezing Poiat
"""of Water

--273''

diminished volume, and a decreased pressure an increased
volume.

The measurement of gases in experiments like the
analysis of air will be of little value for accurate work
unless account is taken of
the temperature and press-
ure changes. For this
reason it becomes necessary
to know to what extent
these aflfect the volumes.
This is not a diflBcult mat-
ter, for it is found that all
gases contract or expand to
almost exactly the same de-
gree when they are sub-
jected to the same changes.
This regularity is some-
what surprising. A similar
thing is not at all true for
solids or liquids. Any ob-
served regularity of this
sort is called a law.

16. Charles' Law. — It is
found that if a certain quan-
tity of any gas is made to
have a temperature of 0° C,
and then is warmed one
degree, the gas expands ^^
of its volume. Warmed to
10**, it expands ^i^^ of its
volume. Heated to 273°, its volume will be doubled.
On cooling the gas, we find that it contracts ^^ of its

-278

Boiling ft)lnt.pf^
Hydrogen

- -Absolute Zero — •-

Fig. 2.

Centigrade and absolute

temperature scales.

80.6"
0**

10 QA8E8 AND THEIR MEASUREMENT.

volume at 0° C. for each degree. The [^ aeral statement
of these facts is known as Charles' Law. The best
statement of it involves the use of the term absolute
temperature.

By the absolute temperature is meant the temperature
reckoned from the absolute zero, which is placed 273
degrees below the freezing-point of water (Fig. 2). The
conception of the absolute zero has come from Charles'
Law. A gas that followed this law perfectly would have
zero volume at — 273°. As a matter of fact, all known
gases become liquids before this temperature is reached.
Charles' Law may be stated thus :

The pressure remaining the same^ the volume of a gas
varies directly as the absolute temperature.

A Centigrade temperature is changed to an absolute
temperature by adding it algebraically to 273. Thus 24"*
C. becomes 297° absolute (273 + 24) ; - 12° C. becomes
261° absolute (273 - 12).

11. Use of Charles' Law in correcting Oas Volumes. — By

using this law we can calculate what will be the volume of
a gas at a temperature differing from that under which it
is measured. For example, a quantity of air measures
25.6 c.c. at a temperature of 21°. Find its volume at 0°.
21° C. = 294° abs.
0° C. = 273°.
If the temperature of the gas were actually changed
from 294° to 273°, it would be cooled, and would therefore
contract. The volume at 273° will be less than the orig-
inal volume. According to the law, it will be ^^ of its
former volume. Hence :

a: = 25.6 c.c. xfll
= 23.7 c.c.

BOYLE'S LAW. 11

The temperature of 0** C. ( = 273** abs.) is chosen as the
standard temperature for the measurement of gas vol-
umes. The operation of finding the volume at the stand-
Correcting ^^^ temperature is called correcting the volume
fortem- for .temperature. Sometimes it is necessary to
perature. ^^^ ^j^^ volume at a temperature other than
the standard temperature. The operation is a similar
one.

Example : A quantity of gas has a volume of 75 c.c. at
a temperature of 24°. What will be its volume at 100° ?

Since the temperature is increased, the volume will also
be increased. The fraction by which the original volume
is to be multiplied is therefore greater than one.

a; = 75 c.c. x^f
= 94.2 c.c.

12. Boyle's Law. — Experiment shows that if the pressure
on any gas is doubled and the temperature kept constant,
the resulting volume will be one-half the original volume.
Under a pressure three times as great, the volume is one-
third. In general, the greater the pressure, the less the
volume in a proportional degree. This generalization,
known as Boyle's Law, is usually stated thus :

The temperature remaining the same^ the volume of a gas
varies inversely as the pressure exerted upon it

13. Use of Boyle's Law in the Correction of Oas Volumes.
— Boyle's Law, like Charles' Law, enables us to calculate
the volume of a given quantity of gas under new condi-
tions. For example, a quantity of gas has a volume of
120 c.c, the barometer standing at 740 mm. What will
be the volume when the atmospheric pressure has increased
until the barometer stands at 760 mm.?

GA8E8 AND THEIB MEASUREMENT.

It-

r760

The numbers 740 mm. and 760 mm. are measures of
the two pressures. The new volume will be found by
multiplying the original volume by the ratio
of these two numbers. It is apparent that
the gas will be subjected to a greater press-
ure under the new condition. According
to the law its volume will be less. The
fraction will therefore have the less number
as the numerator. Hence :

2: = 120x|t^
= 116.8 c.c.

The standard pressure for measuring
gases is the pressure that the atmosphere
exerts when the barometer stands at
760 mm. This is the average height of
the barometer at sea-level. The opera-
tion of finding the volume of a gas at
this pressure is called correcting the gas
for pressure. The volume of a gas at
any pressure whatever is found in a similar
manner.

Example: A quantity of air measures
82.2 c.c. at 520 mm. pressure. What will
be the volume at 800 mm., the temperature
remaining constant ?

It is evident that the resulting volume
will be less than the original, since the press-
ure under the new condition is increased.
Hence the ratio by which the original volume is multi-
plied must be less than one.

a; = 82.2xj
= 63.4 c.c.

Fig. 3.
Barometer.

BOYLE" 8 LAW.

Simnltaneoiu Correetion for Temperatote and Pressure. —

These two corrections can be carried out in one arith-
metical operation, for the temperature effect and the
pressure effect are entirely independent of each other.
For example, a quantity of gas measures 206 c.c. at a
temperature of 22° and a pressure of 760 mm. What
will be the volume of the gas under standard conditions
of temperature and pressure ?

Temperatare Pressure
correction, correction.

a: = 206 X
= 188.1 c.c.

m X m

14. Correetion for Difference in Level. — Gases are usually
enclosed in bottles or. tubes that stand over liquids. The
liquid, as a rule, is either water
or mercury. In order that the
pressure of the gas enclosed under
these conditions shall be equal to
the atmospheric pressure, the
levels of the liquid outside and
inside the tube must be the same
(Fig. 4). This condition is
usually realized by adjusting the
apparatus. Sometimes this is im-
possible, and then it is necessary to correct for the differ-
ence in level. This is done by adding to or subtracting
from the height of the barometer a suitable number. When
the inside level is the higher, the pressure on the enclosed
gas is less than atmospheric (Fig. 6); when the inside
level is the lower, the pressure is greater than atmospheric
(Fig. 6). For mercury, the actual difference in milli-
meters is added or subtracted ; • for water, one-thirteenth

Fig. 4. Fig. 5. Fig. 6.

GASES AND THEIR MEASUREMENT.

value is used, since water is about one-thirteenth
as heavy as mercury.

Example: A volume of gas is enclosed
in a tube over mercury (Fig. 7). The
volume of gas measures 68.3 c.c, and
the level of the mercury inside the tube
is 114 mm. above the level in the dish.
The thermometer reads 20° C. and the
barometer 766 mm. Find the volume
of the gas at standard conditions.

The corrected pressure is found by
subtracting 114 mm. from 766 mm.,

ytUmfn

Fig. 7.

766 - 114 = 652 mm.

The gas volume will be corrected to
standard conditions as follows:

= 54.6 C.C.

15. Correction for Pressure of Water Vapor. — A gas

becomes saturated with water vapor if it is in contact
with water. In such a case the pressure of the water .
vapor makes an appreciable part of the pressure that we
are measuring, and its value must be subtracted from
the observed barometric pressure in order to determine
the pressure of the dry gas; The pressure of water
vapor depends only on temperature, and not on
any other conditions of the experiment. Hence it is
always the same for the same temperature. Tables of
these values for different temperatures have been,
prepared as the result of careful experiments. (See
page 18.)

COBBECTION OF GAS VOLUMES.

Corrections for difference in level and for the pressure
of water vapor (sometimes called aqueous tension) are
both pressure corrections. They are made by adding to
or subtracting from the observed barometric pressure
suitable numbers. They are parts, then, of the pressure
correction.

Example: 24.6 c.c. of nitrogen is contained in a tube
over water. The level of the water inside the tube is
%1 mm. above the outside level. The barometer stands
at 762 mm., and the thermometer at 23°. What is the
corrected pressure ? On consulting a table we find that
the pressure of aqueous vapor at 23° is approximately 21
mm. The corrected pressure is therefore

Difference
in level.

Aqneous
tension.

762 - \i - *21 = 739 mm.

16. The following example will illustrate in full the
operation of correcting gas volumes.

Volume of air 79.3 c.c.

Pressure uncorrected 764 mm.

Temperature ^1°

Difference in level (water) . . . -f 41 mm.

Aqueous tension at 21° .... 18 mm.

Corrected pressure

Corrected volume of air ...

The corrected pressure is

Difference Aqueous
in level. tension.

764 - 41 - 18 = 743 mm.
21" C. = 294° abs.

16 (iASES AND TBEIB MEASUREMENT.

If the temperature of the gas were changed from 294^
absolute to 273° absolute, its volume would become less.
Hence the ratio for the temperature correction is |-J^.
Changing the pressure from 743 to 760 mm. would also
tend to diminish the volume ; the pressure correction ratio
is, therefore, f|^.

'^ = 79.3xn|xm
= 71.9c.c.

Problems.

1. A quantity of hydrogen measures 53 c.c. at a temperature
of 20^ What would it measure at 28° ?

2. 80.2 c.c. of air stand in a tube over water, the barometer
stands at 768 mm.; the next day it reads 755 mm. What
volume would the air then have?

3. 151 c.c. of nitrogen stand in a tube over water, with the
inside level 139 mm. above the outside level. What volume
would the gas have if the two levels were the same? The
barometer stands at 754 mm.

4. How much would 52.2 c.c. of air measure if the barometric
pressure changed from 750 mm. to 762 mm. ? If the tempera-
ture also changed from 18° to 25° ?

5. A quantity of air and water vapor, standing over
water in a gas-measuring tube, levels adjusted, has a vol-
ume of 31.8 c.c. The temperature is 26°; the barometer
stands at 737.6 mm. Correct the volume of air to standard
conditions.

6. A quantity of air and water vapor in a tube over
water, levels adjusted, measures 43 c.c. The thermometer
stands at 24°, the barometer at 770 mm. Correct to standard
conditions.

PROBLEMS.

In the following cases correct the gas volume to standard
conditions :

VOLUMK.

CONDITTONS.

Tbmpbe-

ATITRR.

Baromktkr.

152 C.C.

Over mercury ;
levels the same.

27°

755 mm.

1.26 C.C.

Over water ;
levels the same.

20°

748 mm.

210 c.c.

Over water ;
inside level 80 mm.
above outside level.

22°

764 mm.

10.

15.2 c.c.

Over mercury ;
inside level 30 mm.
above outside level.

21°

760 mm.

11.

129 C.C.

Over- water ;
levels ihe same.

17°

770 mm.

12. A. volume of gas (dry) measures 58.5 q.c. at a tempera-
ture of 183** and a barometric pressure of 759 mm. Find the
volume of the gas under standard conditions.

13. In determining the percentage composition of air, a pupil
has the following data:

Volume

Temperature

Barometer

Original.

99.8 c.c.
24°
763 mm.

Final.

77.0 c.c.
19°
750 mm.

The air stands over water, and the levels are adjusted in
reading both volumes. Determine the per cent of oxygen in
the air.

GASES AND THEIR MEASUREMENT.

PuKssuRE OF Water Vapor or Aqueous Tension
(ill millimeters of mercury)

TXMPKRATURK.

PRB88URK.

Tbmpbrature.

Prxsbure.

10.0° c.

9.2 mm.

20.0° C.

17.4 mm.

10.5

9.5

20.5

17.9

11.

9.8

21.

18.5

11.5

10.1

21.5

19.1

12.

10.5

22.

19.7

12.5

10.8

22.5

20.3

13.

11.2

23.

20.9

13.5

11.5

23.5

21.5

14.

11.9

24.

22.1

14.5

12.3

24.5

22.8

15.

12.7

25.

23.5

15.5

13.1

25.5

24.2

16.

13.5

26.

25.0

16.5

14.0

26.5

25.7

17.

14.4

27.

26:5

17.5

14.9

27.5

27.3

18.

15.4

28.

28.1

18.5

15.9

28.5

28.9

19.

16.4

29.

29.8

19.5

16.9

29.5

30.7

30.

31.6

CHAPTER III.

OXTGEir.

17. Preparation. — In 1774, Priestley obtained oxygen
from a red powder prepared by heating mercury in the air.
If this powder be heated at a temperature somewhat
higher than that at which it was prepared, it is decom-
posed into a gas (oxygen) and metallic mercury.

When pure oxygen is desired in quantity, it is usually
prepared by heating potassium chlorate (Fig. 8), a com-

FV^

'^^

Fig. 8. Preparation of Oxygen.

i4. tube containing potassium chlorate and manganese dioxide ; B, safety
tube ; C, pneumatic trough ; D, D, bottles for collecting the gas.

pound of potassium, chlorine, and oxygen; this, when
heated, melts and gives oxygen gas and a residue of potas-
sium chloride. In the laboratory it is customary to mix the
potassium chlorate with manganese dioxide, as it is found
that the decomposition is more regular and takes place at
a lower temperature. A material which aids the decom-

20 OXYGEN.

position of another, without itself changing, is called a
catalytic agent. To free the oxygen from dust and other
impurities, it is allowed to bubble through water.

Oxygen may be prepared from water by passing an
electric current through it.

18. Physical Properties. — Pure oxygen is a gas without
color, taste, or odor. It is slightly more dense than air.
It dissolves somewhat in water ; under ordinary conditions,
100 volumes of water dissolve 3 volumes of oxygen. If
ordinary faucet water be allowed to stand in a glass, or if
the water be warmed, bubbles will be observed collected
against the sides of the glass before the water actually
boils. Such bubbles are largely oxygen, which was dis-
solved in the waters If cooled sufficiently, oxygen con-
denses to a pale blue liquid, and, on still further cooling,
solidifies.

19. Chemical Properties. — The most noticeable chemi-
cal property of oxygen is its tendency to combine with other
elements. At ordinary temperatures it does not readily
react with many substances, but at higher temperatures
its action is rapid, and is usually accompanied by heat and
light. Nearly all the elements combine readily with oxy-
gen to form compounds known as oxides.

Combustion is a chemical action by which heat and
light are evolved. Lavoisier, in 1786, was the first to
Gombua- explain ordinary burning as the combining of
*^**^- a substance with oxygen. Heat is usually

given off during such combinations. When the action
takes place rapidly, the increase in temperature may be
appreciable, and light may result. Thus, when a piece of
coal burns, the carbon of the coal combines with the
oxygen df the air to form carbon dioxide, a gas which

COMBUSTION.

passes o£f unseen; at the same time a considerable
quantity of heat is evolved, and the neighboring par-
ticles of fuel become
red-hot.

As the air is only
about one-fifth oxygen,
substances do not burn
as readily in it as in
pure oxygen. A glow-
ing splinter plunged
into oxygen bursts into
flame. Charcoal, which
only glows in the air,
burns rapidly in OXy- ^'^•'^' Phosphorus combining with oxygen.

gen. Sulphur burns in air with a pale blue flame, in oxygen
brilliantly. Iron can be made to burn in oxygen with
dazzling scintillations. A convenient method for burning
solid substances in oxygen is to use a quick-sealing fruit
jar- and deflagrating spoon (Fig. 10).

The combining action is not always
\ ^ ' ^" — "^ accompanied by light or even by notice-
j r able heat. Thus, when iron siow ozida-

' ^ rusts, it slowly combines with *io»-

oxygen ; when wood decays, the materials
produced are nearly the same as are
formed when it burns. The total
amount of heat is the same in both
cases, but in the decay the change takes
so long a time that there is no appreci-
able change of temperature. A match
gently rubbed in the dark, appears lumi-
nous without flaming. Such changes are termed slow
oxidation^ as distinguished from burning.

Fig. 10.

22 OXYGEN.

20. We know that some substances burn more easily
than others; heat must be applied to raise them to the
temperature at which they take fire and begin to burn.
Kindiinif '^^^® hindling temperature varies with different
tempera- substances ; the kindling temperature of phos-
*"'®' phorus is but little above the ordinary labora-
tory temperature, but the temperature produced by the
burning is high. If the burning material is a good con-
ductor, as iron, the heat is conducted away so rapidly
that the temperature falls below the kindling tempera-
ture and the fire goes out. Similarly, gas lighted above

H an iron gauze (a. Fig. 11) does not catch

I fire below the gauze, because the heat of

I the flame is conducted away by the iron.

I JK When the material is in small pieces, or

^9\ ^\, J(tr%.^^ is powdered, there is more surface exposed
to the oxygen, so that the burning can
proceed more rapidly ; thus, finely divided
iron will burn, since there is a large sur-
face exposed, and there is no large mass
to withdraw the heat. In the case of
iron, the oxide produced is a solid which
^^^' ^^' remains, and may cover the iron and

prevent its coming in contact with the oxygen, thus stop-
ping further action. Iron is artificially coated with a thin,
regular film of oxide to protect the sheet against rusting.
Iron so protected is known as Russia iron.

21. Many oils, such as are used in paints, absorb

oxygen. Linseed oil absorbs oxygen and forms
ouB com- a tough, resinous substance, the skin seen on the
buBtion. surface of paint. On painted surfaces this skin
holds the coloring matter and protects the material be-

RELATION TO LIFE. 23

neath. The heat generated in its formation is dissipated in
the air. If rags or waste, greasy with such oils, are left
4ying about, oxidation takes place, and since the mate-
rials are usually poor conductors and their form prevents
sufficient circulation of the air to keep them cool, the heat
does not escape, but accumulates until the temperature
rises high enough for the stuff to take fire. Such cases of
burning, started by the accumulation of the heat of a slow
oxidation, are often called spontaneous combustion. It is
especially liable to occur in poorly ventilated places and
with greasy cloth and waste such as is used about ma-
chinery. Coal-dust often takes fire in this way in coal
bunkers.

Since all common cases of burning depend on the pres-
ence of oxygen, the gas is said to support combustion.

22. Occnrrence.— Oxygen is the most abundant element:
about one-half of the solid crust of the earth, eight-ninths
of the water, and one-fifth of the air, is oxygen. Lime-
stone, marble, clay, quartz, and sand are nearly half oxygen,
and it comprises a large proportion of animal and vegeta-
ble matter.

23. Oxygen in Relation to Life. — All animals need oxy-
gen for the carrying on of their life processes. The air
supplies this needed oxygen to land animals, while fishes
obtain it from the dissolved oxygen which water absorbs
from the air. The oxygen is taken in during the process
of breathing, absorbed by the blood, and carried to all
parts of the body. The various tissues are slowly oxi-
dized, heat being liberated by the action. It is this heat
which keeps the bodies of the higher animals continually
warmer than the surrounding air. One of the chief

24 OXYGEN.

products of this oxidation is carbon dioxide, which is
carried by the blood to the lungs and there exhaled.

Plants feed on carbon dioxide, which they absorb from
the air through their leaves. The carbon of this com-
pound is retained in the tissues of the plants, but the
oxygen for the most part is returned to the air. Thus
plants and animals mutually assist in keeping the quan-
tity of oxygen in the air a constant quantity. Plants also
inhale a small quantity of oxygen directly from the air,
and exhale a little carbon dioxide.

24. Ozone. — If electric sparks are passed through oxy-
gen, or better, if it be subjected to a " silent discharge,"
it is changed to another form of oxygen, which is more
active, and which has the irritating smell noticeable where
electrical machinery is working. This form of oxygen is
known as ozone.

Ozone is also produced when turpentine slowly evapo-
rates and by the slow oxidation of phosphorus in moist air.

Silver, which is not affected by oxygen, is rapidly tar-
nished (oxidized) by ozone. Many colors are bleached
by it (notice the cork of a turpentine bottle). Could it
be cheaply prepared in quantity, it would be a useful dis-
infectant.

When ozone is heated to 250° C. it is changed to oxygen,
two volumes of ozone yielding three volumes of oxygen,
so that the ozone is one and a half times as dense as
oxygen.

The atmosphere sometimes contains a small amount of
ozone near the seashore and in the open country. The
bleaching of flax and linen by exposing the goods on the
grass in the early morning is attributed to the ozone dis-
solved in the dew.

SUMMARY. 25

Summary.

A catalytic agent is a material which aids chemical action with-
out itself being permanently changed.

An oodde is a compound of oxygen and another element.

Ccmbustion is a chemical action in which heat and light are
evolved.

Oxidation is the combination of a substance with oxygen.

Slow oopidation is the combination of a substance with oxygen
without the accompaniment of light or noticeable heat.

The kindling temperature of a substance is the lowest tempera-
ture at which it takes fire and bums.

Bxercises.

1. How would the production of oxygen be affected if
potassium chlorate was heated without a catalytic agent ?

2. What would happen if a lighted candle was lowered into
a jar of oxygen ? Why ?

3. Explain why paper, wood, and coal are used in making
a coal fire.

4. Why are met,al cans provided for the oily waste in wood-
turning shops?

5. Why should not the cloths used in wiping oily lamps be
thrown into a closet?

6. Why may a spark in a flour mill produce an explosion ?

7. How would you prove* that water from a stream or a
pond contains dissolved oxygen ?

8. Account for the peculiar odor in the air after a thunder-
storm.

9. Explain why polishing stoves prevents rusting.

10. Explain why a candle goes out if a wire gauze is slowly
lowered till it touches the wick.

11. What two gases in the air are required by plants?
Which in the larger amount?

CHAPTER IV.

HTDBOaEN.

25. Preparation. — If tjie two wires from a battery be
placed in pure water, it will be found that practically no
current passes. Wateras a very poor conductor of elec-
tricity. If a small quantity of sulphuric acid is added
I Electroly- ^^ ^^^ water, the solution is a good conductor.
During the passage of the current, bubbles form
at the ends of the wires: at the positive electrode
(anode) small bubbles of oxygen appear ; at the negative
electro^Q (cathode) there is a rapid evolution of hydrogen

j| 8i8 of
I water,

Fig. 12. Electrolysis of water. '
a, platinum electrodes showing contact through mercury.

(Fig. 12). If the volumes of the gases be compared, it
will be found that there has been twice as much hydrogen

PREPARATION.

set free as oxygen. The sulphuric acid is found unchanged
in amount at the end of the experiment, while some of the
water has disappeared. Other catalytic agents may be
used instead of sulphuric acid. The electrolysis has
practically converted the water into oxygeij and hydrogen,
the volume of which is very great compared with the
volume of the water decomposed.

If a piece of potassium is placed on water, it skims
rapidly back and forth over the surface, decomposing the
water so rapidly that if a large piece of the metal j^^^qj^ qi i
is used, the action is dangerously violent. A mefiOB o^
great deal of heat is generated ; usually the hydro- ^«•^•'•
gen set free ignites if air is present (Fig. 13). If sodium

be used, although the action is very rapid, the heat gener-
ated is not usually sufficient to ignite the hydrogen unless

28 HYDROGEN.

the water is warm. The metal sets free only one-half of
the hydrogen of the water and combines with the remaining
half and all of the oxygen to form the hydroxide of the
metal. This dissolves in the excess of the water.
Potassium + water — >• potassium hydroxide + hydrogen
Sodium + water — >• sodium hydroxide + hydrogen
If calcium be used, the action is quiet, and the calcium
hydroxide is not all dissolved. Magnesium will only act
rapidly if the water is hot. If steam is passed through a
heated iron pipe, filled with nails, an abundant supply of
hydrogen can be obtained, all the oxygen of the steam
combining with the iron.

Iron + water (steam) — >• iron oxide + hydrogen

In the seventeenth century, Paracelsus observed that
when iron dissolved in acids a gas was evolved.

All acids contain hydrogen and generally give it up in
exchange for a metal, acting in this way somewhat like
^ , water. A water solution of hydrochloric or of

ment in sulphuric acid is commonly used, and the metal
acids by usually employed is zinc (Fig. 14). In the
reaction, the hydrogen is set free and the metal
combines with the acid residue to form a new material.
Thus :

hydrogen ,

+ hydrogen

This is the most convenient method, as the action is
quiet and takes place at the ordinary temperature. The

Sulphuric acid +

zinc — >.

zinc sulphate +

hydrogen
sulphur

[zinc
sulphur

oxygen

.oxygen

Hydrochloric acid

zinc

— >. zinc chloride

hydrogen
chlorine

fzinc
1 chlorine

PHYSICAL PB0PEBTIE8.

rapidity of the action depends on the temperature, the
concentration of the solution, the surface of the metal
exposed, and the purity of the
materials. If commercial zinc
be used, some of the impuri-
ties are carried along with the
hydrogen, giving it a peculiar,
disagreeable smell. If iron is
used instead of zinc, the unpleas-
ant odor is more noticeable.

Fig. 14.
Hydrogen generator.

26. Physical Properties. — Hy-
drogen is without color, taste,
or smell. It is the lightest, sub-
stance known. Its specific gravity is very low, and it is
scarcely soluble in water. The rate of escape, of gases
through minute apertures (effusion of gases) varies in-
versely as the square roots of the densities, and hydrogen,
being the lightest gas, escapes more rapidly than any other.
A small rubber balloon filled with hydrogen collapses
more rapidly than a similar balloon filled with illumi-
nating gas.

Certain metals, as platinum or palladium, have the power
of absorbing a large volume of hydrogen. The hydrogen
can be expelled from the metal by warming it. Such an
absorption of a gas by a solid is called oedusion. While
the gas is being absorbed, considerable heat is set free, and
if oxygen is present, the hydrogen may ignite. The occlud-
ing action of such a metal is utilized in self-lighting
burners and mantles.

Hydrogen has been liquefied and solidified. The liquid
is one-fourteenth as dense as water, and is the lightest
liquid known.

HYDROGEN.

27. Chemical Properties. — The most important chemical
property of hydrogen is its combustibility. Cavendish,
in 1783, showed that hydrogen burning in air formed
water (steam). The flame of hydrogen is blue, almost

Fig. 15.

Metallic cone and condenser to cool water from burning

hydrogen.

invisible in daylight, but very hot. Hydrogen, in burning,
gives as much heat as five times its weight of coal. The
water formed in the combustion usually passes off as steam,
but may be condensed on a cool surface (Fig. 15).

If hydrogen and oxygen are mixed, and the mixture is

raised to the kindling temperature, or a spark is applied,

combustion takes place throughout the entire

mixture almost instantaneously. Such a rapid

combustion is called an explosion.

USES.

A jet of oxygen can be made to burn in hydrogen, thus
showing that the gases take equal parts in the action. A
lighted candle is extinguished in hydrogen.

When heated, hydrogen will combine with the oxygen
of many oxides (Fig. 16), forming water (steam) and
the metal, thus:

Copper oxide + hydrogen — ► copper + water

This process of taking oxygen away from a substance
is called reduction, and substances that take oxygen away
are called reducing agents. As the hydrogen is q^^^^^
oxidized in the process, we see that oxidation and and
reduction go on together and are opposite pro- '•*^<'**^-
cesses. Hydrogen is one ol the most energetic reducing
agents, and the great amount of energy involved on the
formation of water explains the stability of water. Energy
equivalent to 3800 calories of heat must be used* to de-
compose one gram of water into hydrogen and oxygen.

28. Uses. — The low den-
sity of hydrogen permits its
use in balloons. The high
temperature of the flame is
used in the oxy-hydrogen
blowpipe. This consists of

two tubes as shown ^ . _^

Oxy-nyaro^

in Figure 17. The gen blow-
hydrogen passes ^^^•'
through the outer and

Fig. 16.

Reduction of hot copper oxide by

hydrogen.

a, hydrogen generator ; b, drying tube ; c,

copper oxide ; d, anhydrous copper sul-

phate, test for water. lighted at

the tip, then the oxygen is turned on
through the inner tube. As the gases
are supplied under pressure, a blast is
formed which gives an intense heat. This flame is used

82 HYDROGEN.

to melt platinum and other refractory materials. When
a stick of quicklime is placed in the tip of the flame,
it does not melt, but becomes white-hot, giving an intense
white light. This is known as the Drummond, lime, or
calcium light.

The process of joining sheets of lead, edge to edge,
known as lead-burning, consists in laying the sheets in
the position desired and melting the edges together with
a hydrogen flame.

Ordinary water-gas, used in illuminating, contains about
forty per cent of hydrogen. It is prepared by blowing
steam through a mass of incandescent anthracite coal or
coke.

Summary.

Hydrogen is commonly prepared by :

(1) the electrolysis of water ;

(2) the reaction between water and a metal ;

(3) replacement in an acid by a metal. This is the most

convenient method.

A liter of hydrogen, under standard conditions, weighs 0.09
gram. A liter of water at 20*^ dissolves 18.4 c.c. of hydrogen.
Liquid hydrogen boils at — 9>59>.5P and solidifies at — %5^,

Hydrogen burns in oxygen or air, forming water. It is a power-
ful reducing agent. The chief uses of hydrogen are for balloons
and fuel.

Exercises.

1. What would be the result of collecting together the
gases formed by the electrolysis of water and applying a light
to the mixture ?

2. Would you use water or sand to extinguish burning
potassium ? Why ?

EXERCISES. 38

3. Would you use zinc or iron for making hydrogen to fill
a large balloon ? Why ?

4. What becomes of the product, other than hydrogen,
formed when zinc and sulphuric acid react ?

5. Is water an oxide ?

6. How would a soap-bubble behave if filled with hydrogen
instead of air ?

7. Why must all the air be expelled from a hydrogen
generator before the gas is lighted at the end of the delivery
tube?

8. Should vessels containing hydrogen be kept mouth up-
ward or mouth downward ?

9. Why would pure hydrogen not make a good illuminating
gas? P?6>

10. Could hydrogen be substituted for illuminating gas in a
gas stove ? j^ "^ O

11. What is formed when iron oxide is heated in a current
of hydrogen?

12. Would a Welsbach burner supplied with hydrogen give
light ?

13. Why is it particularly important that all the joints of a
hydrogen apparatus be tight ?

14. Would a bottle of hydrogen, closed with an ordinary
cork, remain full after standing overnight ?

15. Would a bottle of hydrogen remain full if left inverted
overnight with its mouth under water ?

CHAPTER V.
OOMPOSmOU OF WATEE AND OOMBIUING WEIGHTS.

Analysis.

29. Determination of Compositioji. — In the electrolysis of
water we showed that water could be separated into two
parts hydrogen and one part oxygen by volume.
Such a separation is called an analysis. The
combining of these substances is called a synthesis of water.
If a known volume of hydrogen and oxygen are intro-
duced into a tube inverted over mercury, and exploded by
Synthesis- ^^ electric spark between platinum wires fused
Yoinmetric. through the glass, it is found' that the volumes
of the gases used up are two volumes of hydrogen to one
of oxygen, and that any excess of
either gas is left unchanged. It is
only when two volumes of hydrogen
nre mixed with one of oxygen that
the two gases totally disappear. A
little mist is seen on the tube, which
ia the moisture formed, and the mer-
cury rises and fills the tube. This,
again, shows that water consists of
J) two parts hydrogen and one part
oxygen. Another form of appa-
ratus (Fig. 18), differing in the
shape of the tube from that just
described, facilitates the adjust-
ment and reading of the mercury
mercury, if needed, can be poured
a4

levels.

Fig. 18.
Additional

DETERMINATION OF COMPOSITION.

through the open arm B and an excess drawn off through
the lower stopcock D. The gases used in A can be drawn
in easily through the three-way stopcock C shown in de-
tail at ^.

If dry hydrogen is passed over a weighed quantity
of copper oxide which is heated, steam and copper
result (Fig. 19). The water can be collected and weighed
in a tube containing a drying agent. The syntheBis:
weight lost by the copper oxide is the weight gravimetric,
of the oxygen. The difference between the weight of the
oxygen and the weight of the water formed is the weight
of the hydrogen. Cu ^ -f f^^ - ^ -/ /■< .^

Fig. 19.

Composition of water by weight.

a, hydrogen generator ; b, b, drying bottles containing concentrated sulphuric acid ;

c. ignition tube containing copper oxide ; d, d, apparatus for collecting water formed.

The ratio between the weights of the oxygen and
hydrogen is found to be 7.94 : 1. This relation is un-
varying. Experience has shown that every i^^of
compound has a definite composition by weight, definite
This is known as Dalton's first law, or the law P'^opo'^<»«-
of definite proportions.

30. Combining Weights. — Experience has shown that the

36 WATER AND COMBINING WEIGHTS.

knowledge of the composition by weight of chemical com-
pounds is very useful. For example, we can determine
what weights of substances are needed for a given chemical
action, and what weight of the products will be formed.
Moreover, a study of these weights reveals some surprising
regularities. Let us consider a few simple cases : hydro-
gen combines readily with chlorine and bromine. The
ratios of the combining weights in the two cases are :

(a) weight of hydrogen : weiofht of chlorine : : 1 : 35.5.
(6) weight of hydrogen : weight of bromine : : 1 : 80.

Sodium and potassium ilso form compounds with chlo-
rine and bromine ; the ratios in these cases are :

((?) weight of sodium : weight of chlorine : : 1 : 1.54.
(c?) weight of sodium : weight of bromine : : 1 : 3.47.
(^) weight of potassium : weight of chlorine : : 1 : 0.91.
(f^ weight of potassium : weight of bromine : : 1 : 2.05.

In combining with hydrogen (see a and 6) :
weight of bromine : weight of chlorine : : 80 : 35.5, or 2.^5 : 1.

In combining with sodium (see c and rf) :
weight of bromine; weight of chlorine : : 3.47 : 1.54, or ^.^5 : i.

In combining with potassium (see e and /) :
weightof bromine : weight of chlorine : : 2.05 : 0.91, ox 2,25 : 1.

An examination shows the ratio of bromine to chlorine to
be the same in each of the three kinds of compounds, that
is, 2.25:1. This suggests the value of reducing all the
ratios to a common standard, so that the regularities will
be apparent at a glance. Hydrogen enters into combina-
tion in the least part by weight of any of the elements.

COMBINING WEIGHTS. 37

Let U9^ therefore^ U9e 1 for the combining weight of this
element. It follows, then, from (a) and (b) that the num-
ber for chlorine will be 35.6; for bromine, 80. In the
case of the sodium compounds it will be necessary to
multiply the ratio (c) by such a number that the number
for chlorine becomes 85.5 ; ratio (<f) by such a number
that the number for bromine becomes 80. (The multi-
plier for ratio ((?) is found by dividing 35.5 by 1.54,
which gives 23; the multiplier for ratio (rf) by dividing
80 by 3.47, which gives 23.)

Multiplying ratios (c) and (d) we have

(^) weight of sodium.: weight of chlorine : : 23 : 35.5.
(A) weight of sodium : weight of bromine : : 23 : 80.

In a similar manner from ratios (e) and (/) we can de-
rive the ratios:

(i) weight of potassium : weight of chlorine: : 39: 35.5.

(y) weight of potassium : weight of bromine : : 39 : 80.

Thus we see that the combining number is found to be
the same for each of these elements, regardless of the
other element in the compound. Experiment shows that
twenty-three grams of sodium or thirty-nine grams of
potassium are required to liberate one gram of hydrogen .
from water or any other hydrogen compound. From this
it appe^rrs that each element enters into chemical action
in a definite number of parts by weight, and if we estab-
lish these numbers on a relative scale, the number for an
element is the same in all its compounds. This number
is called the equivalent or reacting weight of the element.

31. Beacting Weight. — It is found that such a number
can be assigned to every element. The number is found
by determining the number of parts by weight of the given

88 WATEB AND COMBINING WEIGHTS.

element which unite with, or replace, one part of hydro-
gen, or its equivalent. Thus we see that all reacting
weights are relative numbers, and they refer or relate to
the combining weight of hydrogen which is taken as unity.
It frequently happens that more than one reacting
weight can be assigned to a given element. For instance,
oxygen combines in two different proportions with hydro-
gen, forming two different compounds. In water the
ratio is 8 to 1, in the other compound the ratio is 16 to 1.
In such cases, one number is always a multiple of the
other.

32. Hethod of determining Eeacting Weights. —The re-
acting weight is determined by an analysis of the hydro-
gen compound, if one exists. In some other cases the
value is determined by finding the weight of the element
that replaces 1 gram of hydrogen. In still other cases,
the number expressing the weight, of the element that
combines with 35.5 grams of chlorine or 8 grams of oxy-
gen is taken as the reacting weight.

Summary.

The composition of water can be shown by analysis and by
synthesis. Two volumes of hydrogen unite with 1 volume of oxy-
gen to form 2 volumes of steam.

Water consists of 1 part by weight of hydrogen combined with
8 parts of oxygen ; and it illustrates the law of definite proportions,
since it always has this composition.

The number of parts by weight of an element which react with
one part by weight of hydrogen, or its equivalent, is called the
reiicting weight of that element. When an element has more than
one reacting weight, the numbers expressing these weights are
always multiples of the smallest number.

EXERCISES 89

Bxercises.

1. When sugar is heated sufficiently to char it, water is
driven off. What two elements besides carbon must sugar
contain ?

2. Why does a thin film of water collect on the inside of a
lamp chimney when the lamp is first lighted ? Why does the
moisture soon disappear ?

3. Dry hydrogen was passed over heated copper oxide and
the water formed absorbed by fused calcium chloride. The fol-
lowing results were obtained :

Wt. of copper oxide tube before the experiment ... 70 g.
Wt. of copper oxide tube after the experiment .... 66 g.
Wt. of calcium chloride tube after the experiment . 106.6 g.
Wt. of calcium chloride tube before the experiment . 102 g.
From the above data calculate the weight composition of
water.

4. 15 c.c. of oxygen were collected in a eudiometer over mer-
cury. Dry hydrogen was passed into the eudiometer until the
volume of the mixed gas was 22.4 c.c. A spark was then
passed through the mixture. What gas was left in the eudi-
ometer ? How would you prove your answer ? What would
be the volume of the remaining gas ?

5. Mention three ways by which water can be decomposed.

6. What does the analysis of water show its composition to
be by volume ?

7. Mention two methods for the synthesis of water.

CHAPTER VI.
WATEE Aim SOLUTION.

33. Physical Properties. — Pure water is an odorless
liquid. Small quantities appear to be colorless, although
large masses show a distinct blue color. Water is usually
taken as the standard in comparisons of physical properties
of liquids and solids. The zero of the Centigrade ther-
mometer registers the position of the top of the mercury
column when the thermometer is placed in melting ice, and
since a pure substance on being warmed always melts at
the same temperature as that at which it would solidify
if cooled, the zero of the Centigrade thermometer is the
freezing-point for water.

When pure water is heated to 100° C. it boils; if we con-
tinue to apply heat, the temperature does not rise higher,
provided the steam i^ allowed to escape. The heat used
in converting water into steam is known as the heat of
vaporization^ or the latent heat of steam; it is given off
when the steam condenses. Seventy-nine heat units are
required to change a unit mass of ice to water, and 536 heat
units are needed to change the same mass of water to
steam. The heat unit is yJi^ of the quantity of heat re-
quired to raise a unit mass of water from 0° C. to 100° C.

Both the freezing and boiling tiemperatures change when
the pressure changes ; increased pressure raises the boiling-
point and lowers the freezing-point, in both cases tending
to keep it in the liquid form. Any dissolved solid acts in
the same way.

DISTILLATION.

34. Distillation. — All natural water contains dissolved
substances. ' It is therefore necessary to boil water and
then condense the steam to make it fit for chemical use;
this process is called distillation (Fig. 20). Solids and

Fig. 20.

Distilling apparatus.

a, boiling flask ; b, condenser.

liquids with boiling-points higher than that of water would
be left behind, as the temperature of the steam remains the
same during the distillation. Materials having lower boil-
ing-points than water would be distilled before or with the
water ; such impurities, as ammonia, are found in the first
portion of the condensed steam, and this is rejected.

35. Steam. -^ Steam is water in the gaseous state ; at ordi-
nary pressure it condenses to liquid at 100° C. ; if the press-
ure were removed, it would remain in the gaseous condi-
tion at lower temperatures. The volume of steam is about
1200 times that of the water from which it was formed.

42 WATER AND SOLUTION.

36. loo. — If the temperature of the water is lowered to
0° C, and energy removed, it solidifies to ice, usually
crystallizing in hexagonal clusters of needles. There is
considerable expansion during the solidification, and the
density of the ice is only 0.91 that of water.

Water requires more heat to raise its temperature than
do most substances; thel-ef ore its temperature changes more
slowly than most objects, and large masses of water have a
determining influence on the climate of the neighboring
land.

87. Solution. — The most important property of water is
its ability to dissolve substances. A substance is said to
be in solution in a liquid when it is distributed uniformly
through the liquid in a state of such fine division that its
particles cannot be seen, and do not settle out on standing.
When the particles are visible, the substance is said to be
in suspension, and will usually settle quickly. A liquid
used to dissolve a substance is called a solvent; the dis-
solved substance is termed the solute. A solution will not
boil at the same temperature as the solvent, nor will theii*
freezing-points be the same.

Water is a solvent for a large number of substances and
this use is most important. It dissolves both gases and
solids and mixes with many liquids. Liquids which do
not separate but form a uniform mixture when brought
together, as alcohol and water, or glycerine and water, are
said to be misdble.

Saturation. — A solution is not a definite compound. A
small portion of salt may be dissolved in a large quantity
of water; such a solution is said to be dilute. In a dilute
solution, the substance is as uniformly distributed in all

SOLUBILITY FACTORS. 48

parts of the liquid as it is in one containing a much larger
proportion of the dissolved substance. A definite amount
of water will dissolve any amount of a given solid up to a
fixed quantity. If a liter of water at 20° C. is taken, it is
possible to dissolve in it any weight of salt up to 360 grams.
When the water has dissolved all the salt it can under
given conditions, it is said to be saturated with salt at the
temperature mentioned. A similar statement can be made
concerning the solubility of any solid in any liquid. When
any solvent has dissolved all of a given solute it can, under
definite conditions, it is said to be saturated with respect
to that substance under the conditions named, A solution
saturated with one substance may dissolve other sub-
stances. Thus, water saturated with respect to salt can
dissolve saltpeter.

'h 38. Solubility Factors. — The solubility of most substances
is decidedly affected by the temperature. Solids are usu-
ally, but not always^ more soluble in liquids at high than
at low temperatures. Sugar and alum are more soluble in
hot water than in cold. Salt dissolves nearly as well in
cold as in hot water. Calcium hydroxide, used in the
preparation of lime-water, is more soluble in cold water
than in warm.

Unlike solids, the solubility of gases in liquids decreases
as the temperature rises. Ammonia and carbon dioxide
are less soluble in hot water than they are in cold. Dif-
ferent substances differ very much in their solubility in a
given solvent, and different solvents differ in their power
to dissolve the same substance.

39. Freezing Mixtures. — There are important energy
changes during solution. When a solid is dissolved.

WATER ANB SOLITTIOW.

Fig. 2i.
Crystal Usatlon of a supersaturated sclution.

SUPERS ATURATION. 45

energy is absorbed and there is generally a fall in tem-
perature. This is made use of in freezing mixtures.
When ice and salt are mixed, some of the ice melts and
the salt dissolves in the water. Both processes result in
the absorption of heat, and the temperature of the mix-
ture falls considerably below the freezing point of pure
water. When equal parts of ammonium nitrate and
water are mixed, at 0° C, the temperature falls to — 15° C.

40. Snpersatnration. — If a solution is saturated at a high
temperature and then allowed to cool slowly without any
disturbance, it will often cool to a much lower tempera-
ture without depositing any of the substance dissolved
(Fig. 21, a). But if a particle of the dissolved sub-
stance is dropped into the solution, a sudden crystalliza-
tion takes place, accompanied by an evolution of heat
(Fig. 21, J, c, d). Such a solution is said to have been
supersaturated at the lower temperature. Any disturbance
is liable to produce the crystallization.

41. Crystals. — The fact that the solubility varies with
the temperature is made use of in separating solids from
solution. If a solution which is saturated at a high tem-
perature be allowed to cool slowly, the dissolved substance
will often separate into definite forms called crystals.
Crystals are usually transparent and symmetrical (Fig.
22). By the evaporation of the solvent, crystals may be
obtained from the dilute solution of a solid.

42. Water of Crystallization. — Many substances in crys-
tallizing from aqueous solutions unite with a definite
quantity of water which is necessary to the shape of the
crystal. This water is called water of crystallization.
Copper sulphate or blue vitriol contains water of crystal-

WATER AND SOLUTION.

Fig. 22.

Crystals of familiar substances.

I. Quartz (ideal). 2. Quartz (actual). 3. Galena or lead sulphide. 4. Garnet- 5. Alum.

HYDROGEN PEROXIDE. 47

lization, and if it is heated in a test-tube, moisture will be
seen on the cooler portions of the tube and the blue crystal
will change to a white powder. The heating has driven
off the water of crystallization.

43. Effiorescenoe and Deliquesoenoe. — If a crystal of
washing soda is exposed to the air in a dry place, it will
lose its water of crystallization and become covered with
a fine powder. Such a material is said to be efflorescent.

Many materials, as lime, calcium chloride, and caustic
potash, usually absorb moisture from the air and are there-
fore said to be hygroscopic. If they absorb sufficient
moisture to dissolve them or to become weU they are said
to be deliquescent. Such materials are useful in drying
others. Whether a substance will give up its moisture
to the air or will absorb moisture, depends largely on the
amount of moisture already in the air and also on the
temperature.

HYDROGEN PEROXIDE. ^J-^

Hydrogen and oxygen form a compound other than
water in which the weights of hydrogen and oxygen are
as 1 to 16. As it contains more oxygen for a given
amount of hydrogen than water, it is called hydrogen per-
oxide, or hydrogen dioxide.

44. Preparation. — Hydrogen peroxide is prepared by
treating barium dioxide with dilute sulphuric acid. The
reaction taking place may be represented :

barium dioxide + sulphuric acid — > *

hydrogen peroxide + barium sulphate

45. Properties. — Hydrogen peroxide is a clear, syrupy
liquid, heavier than water and miscible with it. Its most

48 WATER AND SOLUTION.

important chemical property is the ease with which it
gives up part of its oxygen. The other product of the
decomposition is water.

hydrogen peroxide — ►- water + oxygen

46. XTses. — As hydrogen peroxide gives up its oxygen
so readil}', it is a valuable oxidizing agent, and most of its
uses depend on this fact. It is seldom prepared pure,
and is always used in solution. The commercial solution
contains about 3 per cent of the dioxide.

It is used as a disinfectant, in washing wounds and sores,
as it oxidizes dead and decomposing matter. It has but
little action on living tissue, and as only water remains
after its decomposition, it does not produce irritation and
poisoning as do many other disinfectants.

It destroys the coloring matter of hair and woollen
goods and in this way bleaches them. In many of its
actions it resembles ozone.

47. Law of Multiple Proportions. — In water the weights
of the hydrogen and oxygen are in the ratio of one to
eight. In hydrogen peroxide the ratio is 1 to 16.
Thus the hydrogen in the peroxide is combined with twice
as much oxygen as the hydrogen of the water. A similar
relation is found in many cases. WTienever two substances^
A and JB, unite to form more than one compound^ if we con-
sider a fixed weight of A^ the weights of.B which combine
with it^ are integral multiples of one another. This is
known as the law of multiple proportions or Dalton's
second law.

Summary.

Water is the standard for specific gravity and for the specific
heat of liquids and solids. Its freezing-point and its boiling-point
are respectively 0° and 100° on the Centigrade thermometer.

EXEBCI8E8. 49

Water can be purified by filtratioii, distillation, and freezing.

It IS the most common solvent. The amount of a solute in a
given quantity of a solvent causes a solution to be either un-
saturated, saturated, or supersaturated. Important temperature
changes take place during solution.

Most substances are either hygroscopic, deliquescent, or efflores-
cent. Many compounds contain water of crystallization.

Hydrogen Peroxide can be prepared by the addition of barium
peroxide to cold dilute sulphuric acid or to dilute hydrochloric add.
Hydrogen dioxide is a strong oxidizing agent and is used as a
germicide and for bleaching.

Hydrogen peroxide consists of 1 part by weight of hydrogen
combined with 16 parts by weight of oxygen.

The composition of water and of hydrogen dioxide illustrate the
law of muUiple proportions,

Bzercisea.

1. Why does water put out fire ?

2. Is ammonia more soluble in cold or in hot water?

3. Water is saturated with soda at a high temperature and
the solution allowed to cool. Would the solution then be
saturated?

4. How could you determine whether a certain solution is
saturated, unsaturated, or supersaturated ?

5. How could a supersaturated solution of "hypo'' be pre-
pared?

6. Why do crystals of washing soda become covered with a
coating of white powder when exposed to air ?

7. How does a hygroscopic substance differ from a deliques-
cent substance ?

a Why is fused calcium chloride used as a drying agent?
9. How does a solution differ from a chemical compound?

50 WATER AND SOLUTION.

10. When sea-water is evaporated, why does one of the sub-
stances in solution commence to separate before the others?

11. How would you show that any natural water is a dilute
solution?

12. Mention three ways by which water can be purified.

13. If sold at the same price per pound, would it be more
economical to buy washing soda before or after it has been
exposed to the air for some time?

14. Show how the composition of water and hydrogen per-
oxide illustrate the law of multiple proportions.

15. Why does not a solution of hydrogen peroxide keep well
when exposed to the air ?

CHAPTER VII.
ATOMS AITD MOLEOULES.

48. Law of Conservation of Matter. — We have studied
several substances and some of the laws governing the
quantities of matter that take part in chemical actions,
without attempting any description of the structure or
make-up of the materials used.

Matter is generally defined as anything which takes up
room. The different kinds of matter are called substances.
So far as we know, matter is indestructible^ nor has any one
succeeded in making something from nothing. We may
change its properties^ but we always have the same amount
of matter after the change as before.

49. Atomic Hypothesis. — We found that the combining
or reacting weights are different for various elements but
are constant or unchanging for each element. There is
apparently something significant in the fact that in the
compounds of oxygen the amount of oxygen combined with
a given weight of hydrogen is eight, or twice eight, times
the weight of the hydrogen.

Since water is composed of eight parts oxygen and one
part of hydrogen, the smallest masses of water must have
this composition. For the same reason, the smallest
masses of hydrogen peroxide must contain sixteen parts
of oxygen to one of hydrogen. There must be some
reason why this number eight is characteristic of oxygen,
and why there is no compound of these elements in which
the ratio is twelve to one or twenty to one.

62 ATOMS AND MOLECULES.

John Dalton in 1806 made certain assumptions, known
as the atomic hypothesis^ by which we can readily explain
these facts. These assumptions were:

1st, matter is made up of small particles ;

2d, these particles possess the power of attracting or
holding on to other particles ;

8d, these particles do not subdivide in taking part in
chemical changes.

These particles, which do not divide in chemical
changes, are called atoms. Different kinds of atoms
may differ in mass, form and combining power,
but all atoms of the same material must be
alike. All the atoms of hydrogen are alike, all the atoms
of oxygen are alike. We have found that when oxygen
and hydrogen combine a substance is formed which pos-
sesses properties differing from either. The smallest con-
ceivable quantity of oxygen will possess the same
properties as a mass of oxygen which we can observe;
the same will be true of the smallest mass of hydrogen,
and of the smallest mass of the product, water. If we
assume that each atom of oxygen is accompanied by an
atom of hydrogen that always holds on to it; the mass
made up of such a pair of minute particles does not
have the properties of hydrogen or of oxygen. It is
a new kind of substance — an oxide of hydrogen. The
smallest conceivable quantity of this substance contains
both hydrogen and oxygen.

The smallest quantity of a substance, haying the prop-
erties of the mass is called a molecule. An atom is the
smallest subdivision of an element. Molecules
are usually aggregations of atoms. The mole-
cule is the physical unit of the mass, as the atoms compris-
ing it do not usually separate during physical changes.

LAW OF DEFINITE PROPORTIONS. 53

50. Explanation of fhe Law of Definite Proportions. —

Suppose one atom of an element, e.g. hydrogen, combines
with one atom of another element, e.g. oxygen, to form
a molecule; then any considerable quantity of hydrogen
oxide would be made up of a great number of such
molecules. The weight of the hydrogen oxide would be
the sum of the weights of the atoms composing it.

If we assume that the mass of the oxygen atom is eight
times that of the hydrogen atom, it follows that in the
mass of the oxide the ratio of the oxygen to the hydrogen
must be eight to one, since equal numbers of the atoms of
each element were used in the combination. Suppose,
however, an effort is made to cause 8.3 grams of oxygen
to unite with 1 gram of hydrogen. These weights will
not contain equal numbers of atoms; the mass of the
oxygen will contain the larger number. Consequently,
when combination takes place, a number of oxygen atoms
will remain unused. The mass of oxygen that has com-
bined will weigh exactly eight times as much as the
hydrogen. The 0.3 gram excess of oxygen will remain
uncombined.

Whatever the weight of the atoms may be, chemical
action must take place between definite masses of sub-
stances, and the composition of a compound must be
definite. The law of definite proportions, then, is ex-
plained by assuming that chemical combinations always
take place between atoms.

51. Explanation of the Law of Multiple Proportions. — If

the hydrogen oxide molecule is composed of one atom of
hydrogen and one atom of oxygen, we can imagine com-
binations of one atom of hydrogen with two, three^or more
oxygen atoms. Whatever the combination may be, it is

54 ATOMS AND MOLECULES.

eyident from the atomic hypothesis that the weight of
oxygen combined with a certain quantity of hydrogen must
be an integral multiple of the amount which combines with
the hydrogen to form hydrogen oxide.

The reacting weights are ratios between the weights of
different kinds of atoms, or multiples of these weights.
If we knew that in water one atom of oxygen was combined
with one atom of hydrogen, as we assumed^ the weight of
the oxygen atom would be eight times that of the hydrogen
atom. If, however, there are two atoms of hydrogen to each
oxygen atom, the one atom of oxygen must weigh sixteen
times as much as one atom of hydrogen. If there are two
oxygen atoms to each hydrogen atom, each oxygen atom
would be four times as heavy as the one hydrogen atom.

If we know how many of each kind of atom there are in
a molecule, we can find the relative weights of the atoms.
Such determinations have been made by comparison of
physical properties.

52. Value of Atomic Hypothesis. — The atomic hypothesis
gives a convenient way of explaining the facts upon
which the laws of definite and multiple proportions are
based. We must pot forget, however, that the laws are
statements of facts ^ based on experimental evidence^ while
the atomic hypothesis is used in the attempt to picture a
structure or process which would agree with the facts.
We do not know that this is the way that matter is made
up. Perhaps in time a better explanation, based on dif-
ferent suppositions, may be offered, but we do know
that it has proved useful in explaining a wide variety of
facts and has done more than any other theory for the ad-
vancement of chemistry. Practically all scientific explana-
tions of chemical phenomena are based on this hypothesis.

SUMMARY. 55

Sumznajy.

Matter is anything that takes up room. It is indestructible.
Its properties may be changed, but there is always the same
amount of matter after a change as before.

The study of the weight relations of chemical changes shows
that each element has its definite combining or reacting weight.
The amount of any element found in chemical compounds is either
this reacting weight or some multiple of it.

These facts are explained by the atomic hypothesu. This
assumes matter to be made up of small particles which attract
or hold on to other particles, but which do not subdivide in
chemical changes.

Atoms are the particles indivisible in chemical changes. All the
atoms of an element are alike and possess the characteristic
properties of that element, but differ from the atoms of all other
elements. A molecxde is the smallest quantity of a substance
having the properties of the mass.

The atomic hypothesis gives a convenient explanation of the
facts upon which the laws of definite and multiple proportions are
based. It has been the most valuable theory in the establishment
of chemistry as a science. Sometime a better explanation may
replace this hypothesis.

BzerciseB.

1. Why must an extended study of the composition of
substances be made before the atomic hypothesis can be ac-
cepted ?

2. Mercury is put into a glass flask which is then sealed,
weighed, heated, and weighed again. Why is there no change
in the weight although the mercury turns to a red powder?

3. Why is it that the attempt to make 35.5 grams of
chlorine combine with 24 grams sodium, always leaves 1 gram
of sodium uncombined ?

56 ATOMS AND MOLECULES.

4. Dalton knew that one oxide of carbon contained 2| parts
of oxygen to 1 part of carbon and that another oxide was com-
posed of 1^ parts of oxygen to 1 part of carbon. What law do
these two facts illustrate? Explain them according to the
atomic hypothesis.

5. Why is the molecule of more importance in physics than
in chemistry ?

6. Why was not the present atomic hypothesis evolved
before the time of Lavoisier ?

7. Dalton showed that for one part by weight of hydrogen,
defiant gas contained twice as many parts by weight of car-
bon as marsh gas. Explain these facts according to the atomic
hypothesis.

8. Explain this statement: ''Without the atomic concep-
tion, chemistry would be a chaos of unrelated facts ; with the
theory, it has become an orderly science " (T. W. Clarke).

CHAPTER VTTI.

OHLOsnrE.

53. Chlorine may be said to be a typical non-metallic
element. It displays in a marked degree those properties
which are regarded as characteristic of the non-metals.
The most abundant compound of chlorine found in nature
is sodium chloride, common salt. Sodium chloride is a
very stable compound ; heat does not decompose it except
at an extremely high temperature. Chlorine can be ob-
tained from it in several ways.

54. Preparation. — An electric current can be passed

through a solution of common salt, using apparatus

similar to that used in the electrolysis of water. The

electrodes in this case, however, should be of carbon, since

platinum might combine with the chlorine which is

evolved. The apparatus is filled with a con- EiectroiyBis

centrated solution of salt. When the current Jf**?™

uon 01 com-
passes, chlorine is evolved as a gas at the monsait.

anode and hydrogen at the cathode. Sodium is probably
first liberated at the cathode; but since this element
reacts rapidly with water, it is impossible for it to accu-
mulate. Hydrogen is set free as a result of the action of
sodium with water.

Sodium chloride — >- sodium + chlorine
Sodium 4- water — >- hydrogen + sodium hydroxide
As the final products we have the two gases, hydrogen
and chlorine, and sodium hydroxide which is dissolved in
the water.

CHLORINE,

Hydrochloric acid is a compound of hydrogen and
chlorine. The chlorine might b^ separated by electrolysis,
but it is more usual to take advantage of the fact that
hydrogen has a great tendency to combine with oxygen ;
Oxidation ^^ ^^^^ ^^ ^^ oxidize hydrochloric acid, the
of hydro- hydrogen will combine with the oxygen to
c oncaa . ^^^^ water, and free chlorine will be obtained.
Oxygen from the air might be used. Hydrochloric acid
(gas) and air are passed through a heated tube containing
a catalytic agent. The action is slow and can be well
carried out only on a large scale.

In the laboratory, manganese dioxide is the oxidizing
agent usually employed. Concentrated hydrochloric acid
solution is mixed with manganese dioxide ; when the mix-
ture is warmed, chlorine is evolved (Fig. 23). The

Fig. 23. Preparation of chlorine.
<7, generating flask ; b, bottles for collection of gas.

hydrogen of the acid combines with the oxygen of the
dioxide, forming water. The manganese combines with

PHYSICAL PROPERTIES. 69

half the chlorine of the acid, forming manganese chlo-
ride, which dissolves in^the water; the remaining por-
tion of the chlorine is evolved as a gas.

Hydrochloric acid + manganese dioxide — ►
water + manganese chloride + chlorine

The chlorine is not usually collected over water, since
dry chlorine is desirable for many experiments. It is
commonly collected by displacement of air, or over salt
water.

A mixture of salt, sulphuric acid, and manganese dioxide
is often used. The salt and sulphuric acid react and form
hydrochloric acid, which is then oxidized by the manga-
nese dioxide.

55. Physical Properties. — Chlorine is a greenish yellow
gas, nearly 2^ times as dense as air ; it dissolves slightly
in water; its density and color, however, render its col-
lection by downward displacement a simple matter.

Chlorine has an intensely disagreeable odor, and attacks
the membrane of the nasal passages and lungs, producing
somewhat the effects of a bad cold. It is very poison-
ous, a full breath of the pure gas would probably cause
death. Inhaling ammonia or alcohol will counteract
some of the effects. It should be prepared and handled
with caution to prevent its escape.

56. Chemical Properties. — Chlorine is a very active
element. It combines directly with many other elements,
especially metals, forming chlorides. When powdered
antimony is sprinkled into a jar of chlorine, Action with
brilliant sparks are seen and a white cloud of metals,
antimony chloride is produced. Zinc, copper, and iron,

60 CHLORINE.

especially when heated, also unite readily with chlorine,
with the formation of chlorides.

Antimony + chlorine — >- antimony chloride
Zinc + chlorine — >- zinc chloride

Iron + chlorine — >- iron chloride

These are true cases of combustion, since heat and light
appear. So we may say chlorine supports combustion, and
thus resembles oxygen.

When molten sodium comes in contact with chlorine, it
blazfes with a dazzling light, sodium chloride (common salt)
being formed. To one who for the first time observes the
change, it seems almost incredible that a harmless, house-
hold necessity like common salt could result from the
union of a gas possessing the disagreeable poisonous prop-
erties of chlorine, with a metal which has sufl&cient energy
to decompose water.

If a jet of hydrogen is ignited in the air and lowered
into a jar of chlorine,. the flame will assume a pale white
Action with appearance ; the color of the chlorine will dis-
hydrogen. appear, and in the jar we will find a colorless
gas, hydrogen chloride, which fumes strongly in moist air.
Much heat is given off in the union of chlorine with
hydrogen, another analogy between chlorine and oxygen.
A mixture of chlorine and hydrogen will not combine in
the dark ; in diffused daylight they combine slowly, and
explode when exposed to direct sunlight or other bright
light.

The greiat tendency of chlorine to combine with hydro-
gen is shown by the fact that it will abstract hydrogen
from many compounds. Turpentine is a compound of
carbon and hydrogen. If a piece of paper is moistened
with warm turpentine and thrown into a jar of chlorine, a

CHEMICAL PROPERTIES. 61

violent action occurs, often with the production of a flame,
and a heavy deposit of soot (carbon) forms on the side of
the bottle. If the breath is blown into the bottle, the
moisture will cause the hydrogen chloride there to fume.
An action similar to that with the turpentine is seen in
the burning of a wax taper in chlorine. Paraffin wax,
like turpentine, contains carbon and hydrogen, and only
the latter combines with the chlorine.

Although water is a very stable substance, under cer-
tain circumstances chlorine will react with it, combining
with the hydrogen to form hydrochloric Action wltli
acid and setting the oxygen free. If a tube ▼a*®'^-
is filled with a solution of chlorine in water and is
allowed to stand in the sunlight, oxygen is slowly
formed and collects at the top of the tube (Fig. 24) :

Water + chlorine — >- oxygen + hydrochloric acid

The acid formed is dissolved by the water. Chemical
actions brought about by the action of light are nut
uncommon ; an important example is the formation <jf
starch in the green leaves of plants under the influ-
ence of sunlight. The photographic process also de-
pends on the effect of light on chemical action.

Chlorine is able to decompose water in the absence
of light, provided there is present an oxidizable sulj-
stance. For this reason chlorine water is a fairly
good oxidizing agent; the chlorine combines with
the hydrogen of the water, and the oxygen set free
combines with the other material present.

57. XTses. — The chief commercial use of chlorine \ ' --J
is as a bleaching agent, especially for cotton goods. Fig724,
Cotton fibre is not naturally white. If unbleached or

CHLORINE.

certain colored goods are placed in a jar of chlorine, no
action takes place if the cloth is dry ; but if moist, the

Fig. 25. Bleaching with chlorine,
fl, dry colored cloth ; b, wet cloth ; c, c, calcium chloride to keep moisture from dry doth.

A , color is quickly destroyed (Fig. 25). Many

dyes and the coloring-matter of many fibres
are easily oxidizable materials ; so that when the chlorine
acts with the water, forming hydrochloric acid, the oxygen
set free changes the coloring-matter to colorless compounds.
Chlorine will bleach some colored compounds by decom-
posing them, combining with the hydrogen of the dye.

In bleaching it is not usual to use chlorine gas, but
bleaching-powder, a compound obtained by absorbing
chlorine in slaked lime. The cotton cloth is soaked in a
solution of this, and then in dilute acid to liberate the
chlorine, and finally thoroughly washed to remove the
chemicals (Fig. 26).

USES.

In the bleaching action the destruction of the color was
attributed to the oxygen ; but oxygen does not ordinarily

Fig. 26. Diagrammatic representation of bleaching.

a, cloth ; b. b, bleaching powder solutions ; c, c, acid solutions ; d, " anti-chlor " (sodium
sulphite solution) ; e, water : /, drying^ and ironing rolls.

bleach even weak dyes. It is found that elements in
Hascent general are more active, that is, have a greater
rtate. tendency to combine with another substance, if

they come in contact with it at the moment of liberation.
An element acting under these conditions is said to act in
the nascent (just born) state.

Nascent oxygen will readily oxidize and kill microscopic
organisms, such as disease germs. Hence chlorine is a
Diain. good disinfectant. Bleaching-powder (chloride

fectant. of lime) affords a convenient source of chlorine
for this purpose ; on standing exposed to air, chlorine is
slowly given off. The gas can be more quickly liberated
by the addition of an acid.

Summary.

Chlorine occurs in nature combined with metals, the most
important compound being salt.

Chlorine is prepared: (1) by electrolysis of brine; (2) by oxida-
tion of hydrochloric acid; (3) by the action of salt with a mixture
of manganese dioxide and sulphuric acid. The first and last
methods are the most common.

Atomic weight, 35.5. Density, 3.19 grams per liter. One vol-
ume of water at ordinary temperatures dissolves about three
volumes of chlorine.

64 CHLORINE.

Chlorine is a greenish yellow, poisonous gas characterized by a
pungent odor and its chemical activity. It reacts with metals to
form chlorides, and with hydrogen and many hydrogen compounds
to form hydrogen chloride. Its reaction with water, yielding
nascent oxygen, is utilized in bleaching cotton goods.

The principal uses of chlorine are for bleaching and disinfecting.

Bzercises.

1. Melted sodium chloride on being electrolyzed gives
sodium and chlorine. Why does not the solution yield the
same products?

2. In the mixture of salt, sulphuric acid, and manganese
dioxide, used in the preparation of chlorine, what is the use of
each?

3. If a solution of chlorine is allowed to stand in the sun-
light, bubbles collect and the color of the solution fades.
Why?

4. Cotton cloth soaked for a long time in chlorine bleaching
solution falls to pieces. Why ?

5. Chlorine injures wool. What substance, already studied,
is used to bleach wool ?

6. Describe a case of combustion in which oxygen is not
involved.

CHAPTER IX,

HTDBOOHLOBIO AOID.

58. Preparation. — One of the most important com-
pounds of chlorine is hydrochloric acid, or hydrogen chlo-
ride. As its name implies, it may be made by the direct
union of hydrogen and chlorine, but
the combination is so violent that
only small quantities can be made at
a time. It may be more conveniently
prepared by taking a chloride, e.g.
sodium chloride, and adding concen-
trated sulphuric acid (Fig. 27). The
action begins immediately and the
gaseous hydrogen chloride is evolved
so easily that little heating is neces-
sary. Too violent action may be
avoided by the successive additions
of small quantities of the sulphu-
ric acid to the chloride, using a dropping funnel. The
hydrogen chloride gas may be collected by the downward
displacement of the air, or, better, over mercury, since this
metal is not attacked by the gas. More frequently, how-
ever, the gas is dissolved in water and the solution used.
The action may be represented thus:

sodium chloride -f- sulphuric acid — >-

sodium sulphate -f- hydrochloric acid

The chlorine of the salt combines with the hydrogen from

66 HTDROCHLOBIC ACID.

the sulphuric lacid and the sodium with the other part of
the sulphuric acid — that is, the part which is not hydro-
gen.

59. General Method for Preparing Acids. — The preparation
of hydrochloric acid illustrates a general method for prepar-
ing volatile acids. Sulphuric acid is used because it boils
(vaporizes) at a comparatively high temperature (338° C),
while hydrochloric acid vaporizes at a much lower temper-
ature. When the sulphuric acid comes in contact with a
chloride, a reaction occurs and some hydrochloric acid is
formed. The excess of sulphuric acid and the newly formed
hydrochloric acid are then both present in the mixture.
The lower boiling hydrochloric acid, however, is soon va-
porized, since its boiling-point is many degrees below the
temperature at which the operation is conducted. The
higher boiling sulphuric acid remains behind and gradually
completes its reaction with the sodium chloride. Finally
all the hydrochloric acid is driven off and any excess of
sulphuric acid remains mixed with the sodium sulphate.
The sulphuric acid furnishes the hydrogen for the hydro- .
chloric acid, while chlorine is obtained from the sodium
chloride. Sulphuric acid is generally lised to prepare acids
having a boiling-point lower than 338°.

60. Physical Properties. — Hydrogen chloride is a color-
less gas with a sharp, penetrating odor. It is slightly
heavier than air.

Its solubility in water is most striking, between four and
five hundred volumes of the gas dissolving in one volume
of water at the room temperature. This solution, com-
monly known as hydrochloric acid or muriatic
SolubiUty. ..-^ ^ . . \ oorrf v. • 1... r xi.

acid, contains about 38% by weight of the

hydrogen chloride. The high solubility of the gas causes

CHEMICAL PROPERTIES. 67

it to unite with the moisture of the air, condensation
occurs, and the minute particles of the resulting liquid
appear as a white mist or fumes which can be -^.
seen when a concentrated solution of hydro-
chloric acid is exposed to the air. The fuming is still more
marked when the moist breath is blown across the mouth
of a tube from which hydrogen chloride gas is issuing.

Hydrogen chloride can be liquefied and also solidified
at low temperatures with increased pressure.

61. Chemical Properties. — Neither liquid hydrogen chlo-
ride nor the gas, when perfectly dry, shows the chemical
properties characteristic of the acids. These properties
belong to the water solution. Hydrochloric acid, then, is
the aqueous solution of hydrogen chloride. The water
solution has a sour taste, changes blue litmus typical
to red, and reacts with many metals, e,ff. zinc, »«id.
iron, or magnesium, setting free hydrogen and forming a
compound of the metals which is usually soluble. The
equations are :

Zinc + hydrochloric acid — >- zinc chloride -H hydrogen
Iron + hydrochloric acid — >- iron chloride -H hydrogen

In these actions the metal replaces the hydrogen in the
acid, forming a chloride. Such actions are characteristic
of a number of compounds called acids. All acids contain
hydrogen which may he replaced hy metals. Substances,
like sugar and glycerine, whose hj^drogen cannot be
replaced by metals, are not classed as acids. In general,

metal + acid — >- salt of the metal + hydrogen

When an acid reacts with a metal, hydrogen is liberated
and is generally evolved as a gas unless there is an oxidiz-
ing agent in the solution, in which case the hydrogen may

68 HTDROCULOBIC ACID.

be oxidized to water. The compound formed by the

replacement of the hydrogen of an acid by a

metal is called a salt. The salt is usually

found dissolved in the water which was used to dilute

the acid.

The sour taste of acids is an interesting but not an im-
portant distinguishing property. Many fruits owe their
taste to the presence of acids. Vinegar is hardly more
than a dilute solution of acetic acid. The change in
color of litmus and of other organic coloring-matters is a
convenient way of recognizing acids, but is not reliable in
all cases.

Hydrochloric acid, like chlorine, reacts with
many metals, forming chlorides.

Metal + hydrochloric acid — >- metallic chloride H- hydrogen
Metal + chlorine — >■ metallic chloride

All the commofa chlorides are readily soluble in water
except three: silver chloride, mercurous chloride, and lead
chloride. 77ie metals having insoluble chlorides do not react
with the acid.

The insolubility of silver chloride is used as a means of
identifying soluble chlorides. If a solution of a silver
nitrate is added to a solution of a chloride, a white, curdy
solid separates ; this precipitate darkens in the light.

Chloride of a metal + silver nitrate — >-

nitrate of a metal + silver chloride

Addition of silver nitrate causes a white precipitate in
many other solutions, but the silver chloride is insoluble in
dilute nitric acid^ but dissolves in ammonium hydroxide. As
hydrochloric acid is a solution of hydrogen chloride^ the
same test together with the litmus test serves to identify it.

COMPOSITION BY VOLUME.

62. Uses. — Very small quantities of hydrochloric acid
are found in the gastric juice and are necessary in the
gastric digestion. It is often given as a medicine in
certain cases of indigestion. Large quantities of hydro-
chloric acid are employed in the preparation of chlorine
to be used in the manufacture of bleaching-powder. It is
also used in the making of chlorides, in cleaning metals,
and in the manufacture of gelatine.

63. Composition by Volnme. — When sodium is placed in
hydrogen chloride, a violent reaction occurs, during which
the sodium replaces the hydrogen. The reaction can be

Fig. 28. Volume composition of hydrogen chloride.

a, cork to prevent heating tube while handling ; b, sodiiim amalgam ; c, rubber band ;

d, rubber stopper.

made less energetic by using sodium amalgam instead of
sodium. Sodium chloride, mercury, and hydrogen result
from the reaction. The volume of the hydrogen remain-
ing after the reaction is found to be one-half that of the
hydrogen chloride taken.

Sodium Amalgam Method. — The experiment can be per-
formed in the following manner: hydrogen chloride is

BTDROCHLOBIC ACID.

generated by causing sulphuric acid to drop slowly into
concentrated hydrochloric acid (Fig. 28, A), It is then
dried by being made to pass through concentrated sul-
phuric acid. A glass tube, about 70 cm. long and 1.5 cm.
in diameter, is filled with the dry hydrogen chloride by
the displacement of mercury.

Sodium amalgam is dropped into the tube of hydrogen
chloride and the mouth of the tube instantly closed with
a stopper (Fig. 28, J?). The tube is then inverted sev-
eral times in succession, its mouth placed under some
water in a tall cylinder, and the stopper removed (Fig.
28, C7). Water rushes into the tube.

The remaining gas (hy-

Fig. 29.
Electrolysis of hydrochloric acid.

«► drogen) is brought under
d ^ r atmospheric pressure by
W raising or lowering the tube
in the cylinder until the
liquid on the inside and out-
side of the tube are at the
same level. A small rub-
ber band is then placed on
the tube at the surface of
the liquid.

The volume occupied by
the hydrogen chloride and
that occupied by the hy-
drogen are determined by
pouring water into the
tube from a graduate to the
levels marked by the rubber
bands.

Electrolytic Method. — The composition by volume can
also be shown by the use of the electrolysis apparatus

8UMMABT. 71

shown in Figure 29. Hydrochloric acid, having a specific
gravity of 1.1, is placed in the tubes a. The three-way
stopcocks b are turned so that there is a passage from c
to d and a saturated solution of sodium chloride is drawn
from the dishes i into the collecting tubes e until they
are filled. The stopcocks are then turned so that there is
a passage from / to d. The current is turned on, and as
soon as the hydrochloric acid above the anode is saturated
with chlorine, the stopcocks are turned so that the hydro-
gen and chlorine will pass into the collecting tubes e.
When the upper surfaces of the sodium chloride solution
are just above the support ^, it is inclined, if need be, so
as to mark the relative height of the solution in the col-
lecting tubes. The lower support h is then made parallel
with ff. The solution between g and h is displaced in the
same time, showing that equal volumes of hydrogen and
chlorine are obtained by the electrolysis of hydrochloric
acid.

Summary.

Hydrogen chloride may be prepared : (1) by direct union
of its elements; (2) by the action of sulphuric acid with a chloride.
The latter is the common method.

It is a gas with a pungent odor. One liter under standard con-
ditions weighs 1.64 grams. One liter of water at 20® dissolves 450
liters of hydrogen chloride.

The dry gas is inactive; its water solution is a typical acid.
The replacement of the hydrogen by a metal gives a chloride. All
but three of the common chlorides are soluble in water.

Two liters of hydrogen chloride, when decomposed, yield one
hter of hydrogen and one liter of chlorine.

The chief uses of hydrochloric acid are for the preparation of
chlorine and chlorides, and for cleansing metals.

72 HTDBOCHLOBIC ACID.

Bzerclses.

1. Why is not the direct union of hydrogen, and chlorine a
practical method of making hydrogen chloride ?

2. Acetic acid boils at 118**. Could hydrogen chloride be
produced by the action of such acid on sodium chloride ?

3. Should hydrogen chloride be collected by upward or
downward displacement?

4. Why is tin moistened with a solution containing hydro-
chloric acid before being soldered?

5. When chlorine is brought in contact with ammonia,
which is a compound of hydrogen and nitrogen, a reaction
occurs. Name one compound formed.

6. How would you determine whether a gas is hydrogen
chloride?

7. What products are formed when metallic magnesium is
treated with hydrochloric acid ?

8. What is formed when an amalgam of potassium and
mercury is exposed to hydrogen chloride?

CHAPTER X.

MOLEOULAB OOHFOSITIOIT.

^64. Volume Selations of Oases. — It has been shoWn that
1 volume of oxygen with 2 volumes of hydrogen give

2 volumes of steam, and that
1 volume of chlorine with 1 volume of hydrogen give
2 volumes of hydrogen chloride.
The study of the actions of other gases gives similar
results; thus:

1 volume of nitrogen with 3 volumes of hydrogen give
2 volumes of ammonia.
In these cases the ratio of the volumes of the gases which
combine may be expressed in whole numbers ; this is also
true of the ratio of the volume of each of the combining
gases to the volume of the product.

Law of Oay-Lussac. — These relations were first stated by
Gay-Lussac in his law of volumes : The relative combining
volumes of gases and the volume of the product^ if gaseous^
may be expressed by small whole numbers. Two other
generalizations relative to gases are :

Boyle's law : the volume of any gas varies inversely as

the pressure ; and
Charles' Law : the volume of any gas varies directly as
the absolute temperature.

65. Eeacting Weights and Volume Weights of Oases. — We

found that a volume of chlorine weighs 35.5 times as much
as an equal volume of hydrogen if the comparison is made

74 MOLECULAR COMPOSITION.

under similar conditions of temperature and pressure.
Similarly we found that oxygen weighs 16 times as much
as hydrogen. The weights of equal volumes of oxygen
and chlorine are, then, as 16 to 35.6. It will be noticed
that these numbers are the reacting weights of the
elements. A similar regularity is found in the case of
other gaseous elements. Hence we make the general
statement that the weights of equal volumes of gaseous
elements are to each other as their reacting weights.

66. Avogadro's Hypothesis. — These uniformities in the
behavior of gases were explained by Avogadro in 1811
by the following hypothesis : Equal volumes of gases
under like conditions of temperature and pressure contain
the same number of molecules. That is to say, a liter of
hydrogen contains just as many molecules as a liter of
oxygen, a liter of chlorine, a liter Of hydrogen chloride, or
a liter of any other gas measured under the same condi-
tions of temperature and pressure.

67. Nnmber of Atoms in the Molecnles of Gaseous Ele-
ments. — In a former chapter (cf . § 61) we showed that if
we knew the number of atdtos of each element in a mole-
cule, we could determine the relative weight of the atoms.
We cannot count the number of atoms in a molecule, but
by means of Avogadro's hypothesis we can arrive at a
definite belief in the matter.

Experiment shows that one volume of chlorine and one
volume of hydrogen combine to form two volumes of
hydrogen chloride. Suppose the given volume of hydro-
gen contains 1000 molecules, then by Avogadro's hypoth-
esis, 1000 molecules must also be contained in the volume
of chlorine; and, since the hydrogen chloride occupies
twice the space of the hydrogen, the volume of hydrogen

MOLECULES. 75

chloride resulting from the combination must contain
2000 molecules. Or, briefly stated :

1 volume of hydrogen + 1 volume of chlorine

— >■ 2 volumes of hydrogen chloride
1000 molecules of hydrogen + 1000 molecules of

chlorine — ►- 2000 molecules of hydrogen chloride

In each of these 2000 molecules of hydrogen chloride
there must be some hydrogen, at least one atom (cf. § 49).
At least 2000 atoms of hydrogen have, there- Hydrogwi
fore, been developed from the 1000 molecules ™<>i*««l«-
of hydrogen. Consequently, each hydrogen molecule
contains at least two atoms of hydrogen. Similar reason-
ing shows that the chlorine molecule also contains at least
two atoms.

It is to be noted that any even number might be used
instead of two, but since there is no chemical action known
in which either the hydrogen or the chlorine molecule
seems to divide into more than two parts, it is not proba-
ble that there are more than two atoms in either of these
molecules.

Let us consider the composition of water. Experiment
shows that two volumes of hydrogen with one volume of
oxygen gives two volumes of steam. Following the
method of reasoning used with hydrogen chloride, two
molecules of hydrogen plus one molecule of oxygen
oxygen give two molecules of steam. There moiocui*.
must be at least one atom of oxygen in each of the mole-
cules of steam, so the original molecule of oxygen must
contain at least two atoms. It has been shown that the
hydrogen molecule contains at least two atoms. The
steam molecule, then, must contain at least one oxygen
and two hydrogen atoms.

MOLECULAR COMPOSITION.

While the molecules of all the common gaseous elements
contain two atoms, this is not true of all elements in the
gaseous state. For example, mercury and zinc have each
one atom to the molecule; phosphorus has four; and
sulphur eight, six, or two according to the temperature.

68. Oraphic Method of showing Number of Atoms to Mole-
cule. — The following is an alternative method for arriving
at a belief as to the composition of the molecules of the
common gaseous elements.

As. all our chemical expressions and calculations are
based on the hypothesis of Avogadro, its application
merits a fuller study. The number of atoms in the mole-
cules of gases can be demonstrated graphically.

Let us consider the space occupied by four particles of
hydrogen, representing the particles by four

spheres. A similar space would contain, ac-
cording to Avogadro's hypothesis, four parti-
cles of chlorine. When we say that molecules
occupy a given space, we do not mean that they fill that
space, but that they occupy it in the sense in
which persons are said to occupy a room. I •!

Suppose our volume of hydrogen to unite
with the volume of chlorine ; if one particle of
hydrogen combines with one particle of chlorine, it is evi-
dent that we should have /our pairs; that m^ four particles

1 volume hydrogen 1 volume chlorine

1 volume hydrogen
chloride

of hydrogen chloride. These four particles of hydrogen
chloride would occupy the same volume as four particles

SUMMARY.

of hydrogen, or of chlorine, since equal numbers of par-
ticles of gases occupy equal spaces. We should then
eocpect one volume of hydrogen chloride to be formed.
When one volume of hydrogen actually unites with one
volume of chlorine, two volumes and not one of hydrogen
chloride result. The volume of the acid is twice that of
the hydrogen. Eacli of these two volumes must, accord-
ing to Avogadro's hypothesis, contain four particles of the
acid, or eight in all, so that in the eight particles of the
acid there must be eight particles of hydrogen and eight
particles of chlorine.

1 Yoloine of hydrogen 1 yolume of chlorine

2 TolnmeB hydrogen
chloride

Each of the original four particles must have been com-
posed of two smaller particles. The larger particles are
molecules; the smaller, atoms. Therefore, the hydrogen
molecules and the chlorine molecules consist each of two
atoms.

Summary.

The uniformities in the behavior of all gases, independent of
their chemical composition, leads to the belief that equal volumes of
gases, under the same conditions of temperature and pressure, con-
tain the same number of molecules. (Avogadro.)

It follows from this assumption, and from the volumetric
composition of certain substances, that the elements hydrogen,
oxygen, chlorine, and nitrogen have each two atoms to the
molecule. Zinc and mercury have one atom to the molecule,
phosphorus and arsenic have four.

78 MOLECULAR COMPOSITION.

Exercises.

1. What uniformities (laws) are known about the physical
and chemical behavior of gases?

2. What are the evidences in support of Avogadro^s hypoth-
esis?

3. The gas carbon monoxide combines readily with oxygen.
What relation exists between the volumes of the two gases that
will combine with each other?

4. One liter of marsh gas in burning combines with two
liters of oxygen. How many molecules of oxygen are needed
to react with one molecule of marsh gas ?

5. What volume of air is needed for the complete combustion
of 100 C.C. of marsh gas?

6. About what are the relative numbers of molecules of
oxygen and nitrogen in air ?

7. Two molecules of nitric oxide (gas) unite with one mole-
cule of oxygen when the two are brought together. How many
cubic centimeters of oxygen would be needed for complete
reaction with 64 c.c. of nitric oxide ? How much air would be
needed for the same purpose?

8. Two tanks of equal capacity contain oxygen. The gas in
the first is under atmospheric pressure ; that in the second is
under 3.2 atmospheres' pressure. How does the weight of
oxygen in the second tank compare with the weight of that in
the first?

9. One volume of hydrogen unites with one volume of
bromine gas, forming two volumes of hydrogen bromide.
How many atoms are there in the molecule of bromine?

CHAPTER XI.
ATOmO Aim MOLEOTTLAB WEIGHTS.

69. Atomic Weights. — Since' it has been shown that
there are twice as many hydrogen atoms as oxygen atoms
in the molecule of water, the weight of the oxygen atom
relative to the weight of the hydrogen atom can now be
determined. The oxygen in water weighs eight times as
much as the hydrogen. The weight of the hydrogen atom
is taken as the unit in comparing the weights of the atoms
of different elements ; therefore, the two hydrogen atoms
contained in a molecule of water must have a weight of 2.
The one oxygen atom combined with the two hydrogen
atoms must weigh 8x2, or 16. The atomic weight of
hydrogen is 1 ; of oxygen, 16. The atomic weight of an
element i8 a number which expresses how many times its
atom is as heavy as the hydrogen atom.

70. Density and Specific Gravity. — By the density of a
substance is meant the number of units of mass which
occupy a unit volume. In scientific work the gram is the
unit of mass, and the cubic centimeter the unit of volume.
The density of a substance, then, is the number of grams
of that substance occupying one cubic centimeter. One
gram of water at 4** C. occupies one cubic centimeter.

The specific gravity of a substance is the weight of that
substance divided by the weight of an equal volume of
some substance taken as a standard. Water is taken as
the standard of specific gravity for liquids and solids.

80 ATOMIC AND MOLECULAR WEIGHTS.

71. Specific Oravity of Oases. — There are two standards
for the specific gravity of gases, air and hydrogen. For
experimental purposes air is commonly used ; for purposes
of calculation hydrogen is more convenient. The specific
gravity of a gas with respect to hydrogen is found by
determining the weight of a liter of the gas and compar-
ing this weight with the weight of a liter of hydrogen
under the same conditions of temperature and pressure.
Such comparisons of the weights of equal volumes of
gases can be made at any convenient temperature and
pressure. The usual custom, however, is to compare
weights of equal volumes at 0° C. and 760 mm. ; that is,
at standard conditions. In the determination of specific
gravity, great care must be taken to have the gases pure
and dry. Vapor density is another name sometimes used
for the specific gravity of gases with respect to hydrogen.
The term density is often used for specific gravity.

72. Determination of Molecular Weights. — It follows

from Avogadro's hypothesis that the weights of equal

volumes of two gases will have the same ratio as the

weights of their molecules. This can be shown by the

following reasoning. One liter of hydrogen, measured

under standard conditions, weighs 0.09 gram; one liter

of oxygen, measured under similar conditions, weighs

1 43
1.43 grams. The liter of oxygen is * » or almost 16

times as heavy as a liter of hydrogen. If there are n
molecules in a liter of oxygen, there must be, according
to Avogadro's hypothesis, n molecules in a liter of hydro-
gen. Since n molecules of oxygen weigh 16 times as
much as n molecules of hydrogen, one molecule of oxygen
must weigh 16 times as much as one molecule of hydro-

MOLECULAR WEIGHT.

gen. Tins affords a means of getting the relative weights
of the molecules, or what are called molecu- MolaenUr
lar weights. The molecular weight of a substance "^^^^
is a number which expresses how many times its molecule is
as heavy as the hydrogen atom.

The hydrogen molecule we showed to contain two
atoms, hence its molecular weight is 2. Since a liter
of oxygen contains the same number of molecules as a
liter of hydrogen, the molecule of oxygen must weigh 16
times as much as the molecule of hydrogen, or 32. Thus
we see that the molecular weight of a gas must be twice its
specific gravity referred to hydrogen as a standard. As
chlorine is 35.5 as dense as hydrogen, its molecule weighs
71 times as much as an atom of hydrogen. We have
proved that the chlorine molecule contains two atoms,
so the atom of chlorine
weighs 35.5 times as
much as an atom of hy-
drogen.

73. Altematiye Method
for Determination of
Molecular Weights.^ —
Since one liter of hy-
drogen weighs, under
standard conditions,
0.09 gram, two grams
of hydrogen will occupy
2 -+- 0.09, or 22.22
liters (Fig. 30). The hydrogen molecule we showed to
contain two atoms, hence its molecular weight is 2 (cf.

^ The instructor is advised to have class omit either paragraph 72 or
paragraph 73. The use of both methods will confuse beginners.

Fig. 30. Gram molecular volume.

82 ATOMIC AND MOLECULAR WEIGHTS.

§ 67). There are the same number of molecules in 22.22
liters of any gas (Avogadro's hypothesis). Since the
weight of this volume of hydrogen is expressed by a
number that is the molecular weight of hydrogen, it fol-
lows that the number eocpreBsing the weight in grarns of 22.22
litevB of any gas will be the molecular weight of that ga%,

22.22 liters of hydrogen weigh 2 grams.

22.22 liters of oxygen weigh 32 grams.

22.22 liters of chlorine weigh 71 grams.

22.22 liters of hydrogen chloride weigh 36.5 grams.
Therefore, 32, 71, and 36.5 are respectively the molecular
weights of oxygen, chlorine, and hydrogen chloride.

If 82 c.c. of carbon dioxide weigh 0.1623 gram, we can
find the weight of 22.22 liters by the proportion :
0.082: 22.22 :: 0.1623 ix
X = 44: grams.
Therefore, the weight of the carbon dioxide molecule is 44.
It is 44 times as heavy as the hydrogen atom.

74. Determination of the ITumber of Atoms in the Molecule
of a Compound. — We have shown how the number of atoms
in a molecule of a gaseous element is determined in the
cases of hydrogen, oxygen, and chlorine (cf. § 67).
When the gas is a compound, we can find how many of
each kind of atoms are present by ascertaining
(a) the composition by weight;
(S) the molecular weight.
In the case of carbon dioxide:

(a) its composition by weight is 27.3 % carbon, 72.7 %

oxygen;
(S) its molecular weight is 44.
The weight of the oxygen in the molecule is 72.7 % of 44,
or 32.

&UMMABT. 88

We have shown that the atomic weight of oxygen is 16,
so there must be two atoms of oxygen in each molecule of
carbon dioxide.

The carbon in each molecule will be 27.3% of 44, or 12.
Carbon cannot be vaporized, consequently we cannot de-
termine its atomic weight by the method used for oxygen.
However, many compounds of carbon are gases, and in ho
ease does the carbon furnish less than twelve parts of the
molecular weight. That is, the smallest portion of carbon
that enters into chemical combination (the atorn) weighs
twelve times as much as the hydrogen atom.

The carbon dioxide molecule is composed, therefore, as
its name indicates, of one atom of carbon (weighing 12)
and two atoms of oxygen.

Summary.

It follows from Avogadro's hypothesis that the specific gravities
of gases are in the same ratio as their molecular weights.

The molecular weight of a gas can be calculated by multi-
plying the molecular weight of hydrogen, two, by the specific
gravity of the gas, hydrogen standard. The specific gravity of a
gas is determined experimentally.

The molecular weight of a gas can also be determined by
making use of the fact that the molecular weights in grams of all
gases have the same volume, 22.22 liters. Knowing the weight of
any given volume of the gas, the required molecular weight can
be found from a proportion in which the weight and volume
and the number 22.22 are the three known quantities.

TTie molecular weight of hydrogen is 2, of oxygen 32, of
chlorine 71, of nitrogen 28.

Problems.

1. A liter of bromine gas, at standard conditions, would
weigh 7.2 grams. What is its specific gravity (hydrogen

84 ATOMIC AND MOLECULAR WEIGHTS.

standard)? Using the answer to question 9, chapter X,
determine the atomic weight of bromine.

2. A liter of marsh gas, at st;andard conditions, weighs 0.72
gram. What is the molecular weight of marsh gas ?

3. Methane gas is composed of carbon 75 %, hydrogen 25 %,
and its molecular weight is 16. What part of the molecular
weight of the compound is carbon? What part is hydrogen?
The atomic weight of carbon is 12 ; how many atoms of each
element are there in a molecule of the compound?

4. 0.58 gram of acetylene gas has a volume of 495.7 c.c,
standard conditions. What is the specific gravity of acety-
lene ? What is its molecular weight?

5. Determine the molecular weight of the following sub-
stances :

GA8.

Wbtoht data.

Nitrous oxide

290 C.C. weighs 0.674 g.

93 C.C. weighs 0.631 g.

230 C.C. weighs 0.766 g.

Hydriodic acid

Ether (gaseous conditions)

6. Air is 14.44 times as heavy as hydrogen. Compute the
specific gravity, air standard, of the gases mentioned in the
above problems.

7. What is the numerical ratio between the molecular weight
of a gas and its specific gravity, hydrogen standard ? What
is the ratio between the molecular weight and the specific
gravity, air standard ?

8. Determine the molecular weights of the following gases :

Gas.

8p. G. (air).

Ammonia . . .
Carbon monoxide .
Cyanogen . . .

0.697
0.968
1.806

CHAPTER XII.
SYMBOLS Aim FOBMULAS.

75. It is a great conyenience to have a short, accurate
method of representing the chemical changes taking place
in a chemical action. We have been expressing those
changes in the form of equations, in which we have on
one side the names of the substances which enter into the
action, and on the other the names of the products
formed. According to the atomic theory, chemical action
takes place between molecules, by the rearrangement and
redistribution of the atoms. If we express the action in
terms of molecules and atoms, it becomes much more sig-
nificant. Symbols are used to represent chemical action
as simply as possible.

76. Significance of the Symbol. — The symbol of an
element is usually the initial letter of the name capital-
ized. Thus H means one atom of hydrogen, and as the
atom has mass, it represents also a definite mass of hydro-
gen, one part by weight. The symbol thus means not only
the substance but a definite quantity of the substance.
O means one atom of oxygen, also sixteen parts by
weight of oxygen.

When several substances have the same initial, another
letter conspicuous in the name is added, but not capital-
ized, as C (carbon); Ca (calcium); Cd (cadmium); CI
(chlorine). In some cases the symbol is derived from
the Latin name, as Fe Qferrum^ iron); Cu (jsuprum^ cop-

86 SYMBOLS AND FORMULAS.

per); Na (natrium^ sodium); and K (kaliwm^ potas-
sium).

77. Significance of the Formula. — The formula of a
molecule is formed by grouping together the symbols
of the atoms composing it. The molecule of hydrochloric
acid was found to consist of one atom of hydrogen and
one atom of chlorine. Its formula, therefore, is HCl.
This means

(1) one molecule of hydrochloric acid ;

(2) one molecule of hydrochloric acid containing one

atom of hydrogen and one atom of chlorine ;

(3) one molecule of hydrochloric acid composed of 1 part

by weight of hydrogen and 35.5 parts by weight of
chlorine ;

(4) 36.5 parts of hydrochloric acid by weight;

(5) 1 part by volume of hydrochloric acid (Avogadro's

hypothesis).

When a molecule contains more than one atom of the
same kind, the symbol is not usually repeated, but the
number of the atoms is written as a subscript to the sym-
bol. The formula of water is usually written HgO and
not HOH. As stated above, HgO means

(1) one molecule of water ;

(2) one molecule of water composed of two atoms of

hydrogen and one atom of oxygen ;

(3) one molecule of water composed of 2 parts by

weight of hydrogen and 16 parts by weight of
oxygen ;

(4) 18 parts by weight of water ;

(5) 1 part by volume of steam.

CALCULATION OF PERCENTAOE COMPOSITION. 87

In many cases the molecular weight of a substance can
be determined only indirectly ; in other cases it cannot be
determined at all. In these cases the formula expresses
the simplest combination that will agree with the percent-
age composition. Thus, NaCl represents one molecule of
sodium chloride as composed of one atom of sodium and
one atom of chlorine; that is, 2»3 parts by weight sodium
and 35.5 parts chlorine, or a total weight of 68.5 sodium
chloride. It also represents one volume of the gaseous
salt, but not of the solid.

The number of molecules which take part in a reaction
is represented by means of coeflBcients.

2 HCl = 2 molecules hydrochloric acid

3 HgO = 3 molecules water

78. Calculation of the Percentage Composition from the
Formula. — If the formula of a compound is known, and
also the atomic weights of the elements composing it,
the percentage composition can be calculated. Thus the
formula of an iron oxide is FcjOg ; the atomic weight of
iron is 56, of oxygen 16.

The formula represents :

2 X 56, or 112 parts by weight of iron, and

3 X 16, or 48 parts by weight of oxygen, making

160 parts by weight of iron oxide.

■^J^, or 70 per cent by weight is iron, and
^5^, or 30 per cent by weight is oxygen.

79. Calcidation of the Formula from the Percentage Com-
position. — If we know the percentage composition and the
atomic weights, we can calculate the formula. Thus,
alcohol is composed of: carbon, 52.17%; hydrogen,

88 SYMBOLS AND FORMULAS.

18.04%; oxygen, 84.78%. Dividing each of these by
the weight of one atom of the element, we get:

Carbon ^|1I = 4.84;

Hydrogen 1^ = 18.04;

Oxygen ?i|§ = 2.17. ^

The numbers 4.34, 13.04, and 2.17 bear the same relation
to each other as do the respective number of carbon,
hydrogen, and oxygen atoms in a molecule of the com-
pound. Since, however, atoms are indivisible, the num-
bers expressing them must be whole numbers. To reduce
the numbers in this case to integers, divide each by 2.17,
the smallest number, viz. :

Carbon tvi^^'

Hydrogen __ = 6;

r, 2.17 -,

Oxygen ^^=1.

The simplest formula, based on these relative numbers of
atoms, is CgH^O. This requires a whole molecular weight
of

(2 X 12) + (6 X 1) + (1 X 16) = 46.

The specific gravity of alcohol vapor referred to hydrogen
is found to be 23, which requires a molecular weight of
46 (cf . § 72) ; the molecular weight calculated from the
simplest formula agrees with this. Therefore, CgH^O is
not only the simplest but the correct formula for a mole-
cule of alcohol.

PROBLEMS. 89

For acetylene :

peroentage atomic atomio limplMt

by weight. weight ratio. atomio ratio.

Carbon 92.30 -f- 12 « 7.69 1

Hydrogen 7.69 n- 1 = 7.69 1

The simplest formula agreeing with the analysis would
be CH ; this requires a molecular weight of (12 + 1), 18.
The specific gravity is found to be 18, so that the molecu-
lar weight must be 26 (cf . § 72), and the formula of the
molecule must be CjHj.

For sulphuric acid :

by weight.

Hydrogen 2.041
Sulphur 82.658
Oxygen 65.306

atomic atomio
weight. ratio.

+ 1 = 2.04
H- 32 = 1.02
+ 16 = 4.08

•ImplMt
stomle ntto.

1
4

The simplest formula agreeing with this composition is

Problems.

1. State the meaning of every symbol and figure in each
of the following formulas :

HCl

H,SO«

SCO,

Ca(NO,),

CuS04.6H,0

2. How many atoms of hydrogen in each of the following?

HBr

H,SO,

NH,

NH4C,H,0,

(NH4),Fe,(S04)4.24H,0

90 SYMBOLS AND FORMULAS.

3. Making use of the table of atomic weights in the Appen-
dix, calculate the molecular weights of the following com-
pounds:

CuO

H2SO4
KClOa
ZnCla
NaOH

4. Determine the specific gravity (hydrogen standard) of
each of the following gases :

HCl
CO,
NHs

5. Calculate the weight of a liter of each of the following

CO,
NHs
SO2
CO

6. Acetylene gas has the formula CgHj. What is the
weight of a liter of it ?

7. What per cent of potassium chlorate, KClOa, is oxygen ?

8. If a sample of washing soda has a composition repre-
sented by the formula NagCOg . 10 H2O, what per cent of it is
water ?

9. A hundred grams of a compound contain 30.43 grams
of nitrogen and 69.57 grams 'of oxygen. What per cent of the
compound is nitrogen and what per cent is oxygen ? What
is the ratio between the number of nitrogen atoms and the
number of oxygen atoms ? What is the simplest formula that
could be used to express the composition of the compound ?

• PBOBLEMS. 91

10. Calculate the. percentage composition of the compound
whose molecule is represented by the formula Ca(N08)2.

11. A substance on analysis was found to contain carbon
40%, hydrogen 6.67%, and oxygen 53.33%. What is the
simplest formula that could be used to represent such a sub-
stance?

12. Calculate the empirical (simplest) formula of a com-
pound containing calcium 29.41%, oxygen 47.06%, and sul-
phur 23.53%.

13. The specific gravity of a certain gas, hydrogen standard,
is 14. What is the molecular weight of the gas ? It is com-
posed of carbon 42.8% and oxygen 57.1%. What is its
formula ?

14. Alcohol, a liquid at ordinary temperatures, is readily
converted into a gas; 0.247 gram of the gas has a volume of
184.9 c.c. at a temperature of 150® C. What is the specific
gravity of the gas, hydrogen standard? What is the molecu-
lar weight of alcohol? Alcohol is composed of carbon 52.2%,
hydrogen 13.0 %, oxygen 34.8 %. Determine its formula.

15. 0.55 gram of a certain gas has a volume of 277.7 c.c. at
standard conditions. The gas is composed of nitrogen 63.6 %,
oxygen 36.3 %. What is the formula of the substance ?^ (At.
wt. N = 14.)

16. 0.35 gram of a liquid that is easily vaporized has, in the
gaseous form, a volume of 99.7 c.c. (corrected). The substance
is composed of carbon 92.3%, hydrogen 7.7%. Determine
the formula of the substance.

CHAPTER XIII.
OHEMIOAL EQUATIOirS.

80. CSiemioal Seaotions Sepresented by Equations. — Since
the symbol of an element and the formula of a compound
represent more than the name, we may use them instead
of the names in the equations we have employed, and then
the equation will represent definite numbers and masses
as taking part in the reaction. The reacting substances
are usually written first, on the left ; the products on the
right ; the arrow (or the equality sign) is not to be read
as " equal to," but as give or form ; the addition sign, as
with. As the change may occur under different conditions,
no attempt is made to represent how the action occurred.

These equations are not algebraic; they cannot be cal-
culated or predicted with certainty. They represent
changes which actually take place. When we know by
experiment: • •

(1) that substances will react ;

(2) the composition of each substance ;

(3) all the products formed ;

(4) the composition of each product ;

we can represent the reaction by an equation, and calcu-
late the relative quantities involved.

The fundamental principle upon which chemical calcula-
tions depend is the indestructibility of matter, so the
equation must represent the same amount of each element
after the change as before. There must be the same

FOBMATION OF 0ZIDE8. 98

number of atoms of each element represented on each
side of the equation. On heating potassium chlorate,
only two products result, potassium chloride and oxygen.
Potassium chlorate has the composition shown by the
formula KCIO3; potassium chloride, KCl; and we have
shown (page 75) that there are two atoms in the mole-
cule of oxygen, Oj. Using these formulas, we have:
KCIO3 — >- KCl 4- Oj ; but it will be seen that there are
three atoms of oxygen on the left and only two on the
right-hand side of the equation. In order that the quan-
tities and compositions shall be correctly represented, the
equation must be written :

2KCIO3— )-2KCl-|-3 0,

In this equation the number of atoms of each element is

the same on the two sides of the equation. So, also, we

must write :

2H2 4-Oa^^2H20

Materials which are present, but which undergo no
change, such as water in which the substances are dis-
solved, catalytic agents, etc., are not expressed in the
equation.

81. Equations for Seaetions Already Studied. — The heat-
ing of copper, magnesium, and zinc in the air result in
the formation of the oxide of the metal heated in each
case. The equations are :

2Cu +O2 — ^2CuO

2Mg 4-O2 — ^2MgO

2Zn 4-02-^-2ZnO

The slow oxidation of phosphorus, made use of in the
analysis of air, and the burning of phosphorus in oxygen
are both represented by the equation :

4P4.5O2— ^2P205

94 CHEMICAL EQUATIONS.

Thus we find that the product, phosphorus pentoxide,
is the same whether the phosphorus combines with the
oxygen slowly and quietly, or rapidly and violently. All
the above equations illustrate the process of direct combi-
nation or synthesis.

The production of oxygen from mercuric oxide and
from potassium chlorate illustrates the opposite process,
decomposition or analysis. The equations are :

2HgO — ^2Hg 4-O2
2KC108— ^2KCl + 302

82. The equation for the preparation of hydrogen by
electrolysis of water is :

2H3O— ^2Ha-h02

When sodium reacts on water, we have :

2Na 4- 2H2O — ^ 2NaOH + H^

For the formation of hydrogen by the reaction between
metals and acids, we have the following equations :

Zn + H2SO4 -^ ZnSO^ 4- Hg
Mg 4- 2 HCl — >- MgCla 4- H^

The formulas of sulphuric acid and zinc sulphate show
very clearly how the latter is produced by the replacement
of the hydrogen of the acid by the zinc. The formation
of sodium hydroxide and magnesium chloride in the equa-
tions given above furnipb oth«r examples of replacement.

83. The formation of water by the burning of hydro-
gen in oxygen or in air is represented by the equation :

2H2 + O2— ^2H20

It will be noticed that this is the exact reverse of the
equation given above for the decomposition of water;

PREPARATION OF CHLORINE. 96

a large proportion of our equations are reversible ; the
direction in which the reaction proceeds depends upon
the conditions. This may be shown by the use of the
double arrow ; so we may write the equation :

The reduction of copper oxide by hydrogen is expressed
by the equation :

CuO + Ha:5±:Cu + H20

This equation is reversed with the liberation of hydrogen,
when steam is passed over heated copper.

84. The equation for the preparation of chlorine by
electrolysis of brine is :

2 NaCl + 2 HjO — )- 2 NaOH + Ha + Cla

The equation represents the final result of the reaction,
and does not show the intermediate steps; viz. the separa-
tion of the sodium and chlorine, and the reaction of the
latter with the watejp.

The liberation of chlorine by the oxidation of hydro-
chloric acid in the presence of a catalytic agent is repre-
sented by the equation :

4HCl4-Oa^2H20 + 2Cl2

When manganese dioxide is used as the oxidizing agent,
the equations are :

MnOj + 4 HCl — ^ MnClj 4- 2 Hfi + Cl^
MnOa + 2 NaCl -t- 2 li^S6\ — ^ MnSO^ 4- Na^SO^

4-2H20-t-Clj

A comparison of these three equations shows the forma-
tion of water and chlorine in each case. We shall see

96 CHEMICAL EQUATION 8.

from the equation for the formation of hydrochloric acid
given below, that in all three methods we may regard the
liberation of the chlorine as the result of the oxidation of
the hydrogen of hydrochloric acid.

85. The following equations represent the reaction of
chlorine with various substances, resulting in the forma-
tion of chlorides :

+ C1,

:i^

2HC1

2Sb

+ 3C1,

, — >-

2SbCl,

+ C1,

— >-

ZnClj

HjO + CI,

— >-

2CHC1-I- O

The last equation represents the liberation of nascent
oxygen, in the bleaching by chlorine, so we write the
symbol O, indicating the oxygen atom, and not Oj, indicat-
ing the oxygen molecule. It will be seen that this
equation is the reverse of that given above for the prepa-
ration of chlorine.

86. Two methods for producing hydrochloric acid are
indicated in the equations just given. We may represent
its formation from salt and sulphuric acid as follows :

2 NaCl + HjSO^ — )- Na^SO^ + 2 HCl

This equation illustrates double replacement, or, as it is
more often called, double decomposition ; each compound
apparently breaks up into two parts, each of which unites
with a different part of the other compound. Such reac-
tions can only be prevented from becoming reversible by
the removal of one of the products from the field of
action; in this case hydrochloric acid is driven off as a
gas. Double replacements are common in solutions.

CALCULATION OF RELATIVE WEIGHT 8. 97

87. Calonlation of SelatiYe Weights from the Equation. —
In the equation

2KCIO3— ^2KCl4-30a,

the molecule of potassium chlorate weighs

894-35.5 + 3x16(122.5);

the molecule of potassium chloride weighs

39 + 36.5(74.5);

the oxygen molecule weighs

2x16(32).

That is, 2 x 122.5, or 245 parts by weight of potassium
chlorate, on being decomposed, give 2 x 74.5 (149) parts
of potassium chloride ; and 3 x 32 (96) parts of oxygen.

2KC10,— »-2KCl

30,

2 X 122.5 2 X 74.5

8x82

245 149

Suppose we wish to prepare 20 grams of oxygen; how
much potassium chlorate must be used ? From the num-
bers we find that for every 96 parts of oxygen produced,
245 parts of potassium chlorate were used; so we may
obtain the number sought for from the proportion:

96: 245:: 20 grams: 2;
96aj = 4900.

2; = 51.0 grams, the potassium chlorate needed.

What quantity of sulphuric acid (HjSO^) is needed to
exactly decompose 100 grams of sodium chloride, when
sodium sulphate and hydrochloric acid are formed ?

98 CHEMICAL EQUATIONS.

2NaCl + H^04 — ^ Na^04 + 2HC1

2(234-36.5) (2+32+4x16) (2x23+32+4x16) 2(1+35.Q

117 98 142 73

The weight of the sodium chloride decomposed and the
weight of the sulphuric acid needed will be as 117 : 98.
Forming the proportion:

117: 98:: 100 grams: 2?

2; = 83.7 grams, sulphuric acid required.

Similarly the weights of the sodium sulphate and of
the hydrochloric acid might be found :
for the sodium sulphate,

117: 142:: 100 grams: a:
for the hydrochloric acid,

117: 78:: 100 grams: a;

In any reaction the weights of the molecules involved
have the same ratio as the weights of the reacting sub-
stances. We have used this relation in the calculations
above.

88. Calculation of Yolnme Belations and Volumes from the
Equation. — What has been said in the preceding para*
graphs applies to all substances, solid, liquid, or gaseous.
But in the case of gases, we found the formula has a
meaning that did not apply to liquids or solids (cf. § 77).

In the equation :

2H2+Oj— ^2H20

it appears that two molecules of hydrogen react with
one molecule of oxygen to form two molecules of steam.
Since equal numbers of molecules occupy equal volumes,
the volume of the hydrogen must be twice that of the
oxygen and equal to that of steam.

CALCULATION OF VOLUME RELATIONS. 99

In the equation :

2H2O + 2CI2— ^4HCl + 0a

we see that two molecules of chlorme are used to liberate
one molecule of oxygen, hence two volumes of chlorine
will furnish one volume of oxygen. Therefore, in the
case of ga%e%^ the coefficients represent the relative volumes
of the substances.
The equation

Ha + Cla^2HCl

may be read : 1 part by volume of hydrogen with 1 part
by volume of chlorine will give 2 parts by volume of hy-
drochloric acid ; and also 2 parts by weight of hydrogen
and 71 parts by weight of chlorine give 78 parts by
weight of hydrochloric acid.

Moreover, since the molecular weight in grams (gram
molecular weight) of any gas has a volume of 22.22 liters
at standard conditions, the volumes of gases reacting can
be directly calculated from the equation.

In the decomposition of potassium chlorate:

2KC108— ^2KCl + 3 0a

only one of the substances is gaseous; 245 grams of
potassium chlorate give 3 gram molecules of oxygen (3 O2),
which occupy 3 x 22.22 liters. If we wish to produce 100
liters of oxygen, we can find the weight of potassium chlorate
needed, by the following proportion :

66.66 liters : 100 liters : : 245 grams : x

x= 367.5 grams, potassium chlorate needed.

100 CHEMICAL EQUATIONS.

If we wish to produce 50 liters of hydrogen by the
reaction

Zn -t- H2SO4 — ^ ZnSO^ + H^

65 grams of zinc give one gram molecule of hydrogen
(22.22 liters), therefore :

22.22: 50:: 65 grams: a?

X = 146.2 grams, zinc needed.

Summary.

Chemical reactions are represented by equations in which the
reacting substances are written on the left and the products on the
right, separated by an equality sign or an arrow.

Equations represent actual chemical changes and must indi-
cate the same amount of each element after the change as before.
The composition of each reacting substance and of all the products
must be known before the equations can be written. Solvents
which are unafiPected and catalytic agents are not expressed in
the equation. The conditions of the chemical action are not
indicated by the equation.

Equations may represent processes of :

(a) direct combination;

(b) decomposition;

(d) double replacement;

(e) oxidation and reduction.

The relative weights of the different substances in a chemical
equation can be calculated. Such calculations deal with but two
of the substances at a time. The weights of the substances are in
the same ratio as the weights of the molecules involved. With
the weight in grams of one of the substances known, and the
weights of the molecules ascertained, a proportion is formed with
these three quantities. The weight in grams of the second sub-
stance is foimd by solving the proportion.

EXEBC1SE8. 101

Coefficients of molecules of gaseous substances^ as expressed in an
equation, represent relative volumes. The gram mdectdar weight
of any gas occupies 22.22 liters at standard conditions. By mak-
ing use of these two facts, the volumes of gases as well as their
weight can be calculated from chemical equations.

Bzexcises.

1. Write an equation for (a) a synthesis, (6) an analysis.

2. Write the equation for what takes place when potas-
sium, a metal similar to sodium, reacts with water. What
process does the equation represent ?

3. Write a reversible equation involving hydrogen and
chlorine. Tell how you can control the direction in which the
action proceeds.

4. Write the equation for the reaction that might be
reasonably expected to occur if sodium chlorate, NaClOs, is
heated.

5. Write the equations for the reaction of potassium
chloride, KCl, with sulphuric acid. Name the process and
the products. Explain why the reaction is not ordinarily
reversible.

6. Write the equation for a laboratory preparation of chlo-
rine.

7. Write the equation expressing the reaction of magne-
sium with oxygen; with sulphuric acid; with chlorine; and
with hydrochloric acid.

Name the products and the process illustrated in each case.

8. Calculate the weight of oxygen obtained from heating
20 grams of mercury oxide.

9. How many grams of copper were heated to form 2.64
grams of copper oxide ?

102 CHEMICAL EQUATIONS.

10. 2.4 grams of zinc were treated with an excess of dilute
sulphuric acid. Calculate weight of each product formed.

11. Find how many grams (a) of potassium chloride, KCl,
and (6) of sulphuric acid are needed to produce 2.8 grams of
hydrogen chloride.

12. Calculate the weight of manganese dioxide and that of
hydrogen chloride used to produce 4.8 grams of chlorine.
How many grams of manganese chloride were formed ?

13. How many liters of oxygen at standard conditions can
be obtained by heating 8.4 grams of potassium chlorate ?

14. State the relative volumes of each gaseous substance
indicated by the equations representing :

(a) hydrogen combining with chlorine ;

(b) chlorine combining with water ;

(d) electrolysis of sodium chloride.

15. How many kilograms of iron would be required to fur-
nish hydrogen enough to fill a balloon of 6360 cubic meters
capacity ?

Fe + H2SO4 — >- FeSO* + Hj

CHAPTER XIV.

SODIUM An) POTASSIUM.

89. Preparation of Sodium. — The great chemical activ-
ity of sodium and the stability of its compounds made
the preparation of metallic sodium a difficult chemical
problem. Like many others, it has been solved by the
application of electricity. Sir Humphry Davy, in 1807,
obtained both sodium and potassium by the electrolysis of
the melted hydroxides, and the metals are now prepared
commercially by this method (Fig. 81).

When the current is passed
through the melted hydroxide,
the sodium and hydrogen appear
at the cathode and the oxygen
at the anode. The hydroxide is
contained in an iron cylinder
and melted by gas flames be-
neath. The cathode is a carbon
cylinder coming up through the
bottom. The anode is an iron
or nickel cylinder coming down
from the top and surrounding
the cathode. Between the elec-
trodes is suspended a cylinder
of wire gauze (tf, c) through
which the fused hydroxide (a)
can pass, but which prevents the
passage of bubbles of the gas or globules of

103

104 SODIUM AND POTASSIUM.

dium. The sodium, being lighter than the hydroxide,
rises to the surface (6) above the cathode, and is ladled
oflF from time to time. The hydrogen escapes through
holes in the cover and the oxygen is led off through a
pipe from the side. The reasons for these precautions
will be readily understood when we keep in mind the
energy with which oxygen combines with both sodium
and hydrogen. The heat generated by the passage of the
current is sufficient to keep the hydroxide molten after
the action starts.

90. Phyiical Properties of BocUnm. — Sodium is a silver-
white metal, possessing a brilliant lustre when in a pure
state. It is soft enough at ordinary temperatures to be
readily cut with a knife and to be moulded by the fingers.
It may be formed into wire by pressing it through a hole
in a metal plate. It is a good conductor of heat and elec-
tricity ; silver, copper, and gold are the only metals that
surpass it in this respect. "Metallic" lustre and con-
ductivity are characteristic physical properties of metals.
It is a very light metal, slightly less dense than water.

91. Chemical Properties of Sodium. — Sodium is in gen-
eral a very active element chemically. It burns readily in
oxygen and in chlorine, and is an energetic reducing agent.
The color of its flame is bright yellow, and this color is
imparted to a non-luminous flame when any sodium com-
pound is heated in it.

The most striking chemical property of sodium is its
action in water. Exposed to moist air, it tarnishes almost
Action with instantly, on account of the formation of a
wat«r. layer of sodium hydroxide. In perfectly dry

air it remains unchanged at ordinary temperatures.

SODIUM HYDROXIDE AS A TYPICAL BASE. 105

When thrown on water, it skims over the surface with
a hissing sound. The water is rapidly decomposed, one-
half the hydrogen being set free and the sodium com-
bining with the oxygen and the other half of the hydrogen
to form sodium hydroxide :

2Na + 2H20 — )- 2NaOH + H,

The hydroxide dissolves in water and may be obtained by
evaporation. A large amount of energy is liberated in
the decomposition of water by sodium, which may be
readily shown by dropping a piece of sodium on a moist
piece of filter paper. In this way the heat is all liberated
at one place and is sufficient to ignite the hydrogen.
On account of its ready action with water, the sodium
is always kept under kerosene or some other oil containing
no oxygen.

92. Sodium Hydroxide as a Typioal Base. — A solution
of sodium hydroxide turns red litmus blue, an action
exactly the reverse of that of an acid. If we mix solu-
tions containing weights of hydrochloric acid and sodium
hydroxide proportional to their molecular weights, there
is a rise of temperature, and the resulting solution affects
neither red nor blue litmus. This solution contains
sodium chloride, a fact which is clearly indicated by its
taste. As the acid and the hydroxide have both lost their
characteristic properties, the resulting solution is said to
be neutral and the process is known as neutralization.
The change may be expressed by the equation :

HCl + NaOH — ^ HOH +NaCl

A compound whose solution turns red litmus blue and
neutralizes acids is called a ha%e. Sodium hydroxide,
then, is a typical base.

106 SODIUM AND POTASSIUM.

93. Speotmm Analyns. — The colors imparted to flames
by different elements furnish a simple and yaluable
method of analysis. The different colors found in light
are bent to different degrees in passing through a prism,
and are so separated from each other. The band of light
thus produced is known as a spectrum^ and the instrument
used to produce and view the spectrum is called a spectro-
scope. The light to be studied is admitted through a
narrow slit (^A) in the end of a tube, in such a way that it
will fall in parallel rays on a prism ((7) with its edges
parallel to the slit. When the beam emerges from the
prism, the different colors are separated and the spectrum
is viewed through lenses ( J^) placed at the end of another
tube (Fig. 82).

Fig. 32. Diagrammatic representation of spectroscope.

White light, produced by an incandescent solid, gives a
spectrum consisting of a continuous band of color, shad-
ing from red through orange, yellow, green, blue, and
indigo to violet. When the light is due to incandescent
vapors, the spectrum consists of a series of bright lines,
the color and position of which differ for each element.
The yellow color spoken of above is due to the presence
of sodium vapor in the flame. This may be most conven-
iently produced by placing in a flame a platinum wire
moistened in a solution of a sodium compound. Such a

PROPERTIES OF POTASSIUM. 107

flame when viewed with a spectroscope shows two yellow
lines very close together. The spectrum of potassium con-
sists of a double line in the violet end and a line in the red
end of the spectrum. Lithium is a rare metal closely re-
lated to sodium and potassium. It was thought to exist
in very few minerals until the spectroscope showed that
small quantities of it were widely distributed through-
out nature. Its spectrum consists of a bright red line
and a very faint yellow line. The presence of one-
millionth of a milligram of lithium can be shown by
means of the spectroscope. (For spectra of some familiar
elements, see Fig. 33, frontispiece.)

94. Potassium and its Preparation. — Potassium resem-
bles sodium in so many respects that it may be very
briefly discussed.

Its preparation is similar to that of sodium, substituting
potassium hydroxide for sodium hydroxide.

95. Properties of Potassium. — Potassium is a very silvery
white metal with a slight bluish tinge. It is softer than
sodium, lighter, and melts at a lower temperature. Al-
though sodium and potassium are solids at ordinary tem-
peratures, an alloy of the two can be prepared which is a
liquid.

The chemical properties of potassium closely resemble
those of sodium, but it is more active. It decomposes
water, forming potassium hydroxide and liberating hydro-
gen. The energy produced is sufficient to ignite the
hydrogen, as the potassium skims over the surface (Fig.
13, page 27). Potassium imparts a reddish violet color to
the flame. As the presence of a slight trace of sodium
obscures the potassium flame, several thicknesses of cobalt

108

SODIUM AND POTASSIUM.

blue glass should be interposed between the flame and the
eye to absorb the yellow sodium light.

Summary.

Atomio Wt.

8pioino Or.

MlLTIHG PT.

BoiUMO Ft.

Sodium
Potassium

23.0
39.1

.97

.87

97.6°
62.6°

877^
767°

Both metals are soft and light. They react with water to form
the hydroxides, and with acids to form salts.

Their compounds are characterized by the colors they impart to
a non-luminous flame. Most of the compounds are colorless
(white when powdered), and soluble in water, the potassium conci-
pounds being more soluble.

The hydroxides are typical bases — caustic, alkaline, and neutral-
izing acids. They are used in the preparation of soaps and bleach-
ing solutions, in oil-refining, and in glass-making.

Qzercises.

1. Compare the properties of sodium with the corresponding
properties of metals with which you are more familiar, for
example, iron, copper, silver. Do you find any resemblances ?

2. Why is sodium classed as a metal ?

3. What are the most essential properties of metals from
the chemical point of view ?

4. Calculate the quantity of sodium that could be obtained
from 1 kilogram of pure sodium hydroxide.

5. Calculate the quantity of sodium hydroxide that would
be formed by the action of 5 grams of sodium on water.
What weight of hydrogen would be evolved? What volume
would the hydrogen have under standard conditions ?

6. What weight of hydrogen chloride would be required for
the complete neutralization of 2.63 grams of sodium hydroxide?

CHAPTER XV.

SOLUnOH.

Solutions have already been defined (chapter VI) and
some of their properties discussed. We have found that
many of the reactions studied take place only in the
presence of water, even though the water itself does not
react. The purpose of the present chapter is to furnish
an explanation of these various phenomena.

96. Eleetrol]^ and STon-electrol]^. — Solutions differ
from each other greatly in their power to conduct elec-
tricity. We have already seen that water solutions of
sodium chloride and sulphuric acid (electrolysis of water)

Fig. 34. Determination of electrolytes.

tf , b, c, beakers containing solutions to be tested ; d, cork carrying electrodes with mer-
cury contacts.

readily permit the passage of the current. We may test
other solutions by arranging the following circuit, '^fcclud-
ing a source of current, the substance to be tested^'' and
some instrument to detect the passage of the current
(Fig. 34).

109

110 SOLUTION.

The mcandescent lamp serves two purposes. Its resist-
ance cuts down the current to a strength suitable for
passing through a solution. Secondly, if the lamp lights,
the current must be passing through the liquid in the
circuit, and the solution is a conductor. The brilliancy
with which the lamp glows roughly indicates how well
the solution in the circuit conducts.

When distilled water is put into the circuit, the lamp
does not even glow ; hence, in those water solutions which
are conductors the current must be carried by the aid of
the dissolved substance. A solution of hydrochloric acid
conducts the current readily ; a water solution of sugar
fails to conduct, but sodium hydroxide, a typical base,
proves as good a conductor as hydrochloric acid. Solu-
tions of sulphuric acid, potassium liydroxide, sodium
chloride, sodium sulphate, and copper sulphate all allow
the lamp to glow brilliantly. Glycerine, alcohol, and
many similar compounds are found to be non-conductors.

Our circuit includes twf> different classes of conductors.
The first consists of solid conductors, chiefly metallic,
GiaBBes of s^^ch as the copper wire and the carbon filament
conductors. Qf the lamp. These conductors undergo no
permanent change in carrying the current. The other
class includes those liquids which we have found to be
conductors, the components of which are free to move
toward the electrodes, which are the points where the
current enters and leaves the liquid. In these liquid
conductors the transmission of electricity is attended by
the decomposition of the conductors, components of which
may be liberated at the electrodes. A liquid which is de-
composed during the passage of an electric current is
called an electrolyte. The name given to the process is
electrolysis.

EFFECT OF DISSOLVED SOLIDS. Ill

Arranging the electrolytes and non-electrolytes in two
colnmns, we have

ElEOTBOLTTU Noir-SLSOTBOLTTIS

hydrochloric acid distilled water

sulphuric acid

sodium hydroxide sugar

potassium hydroxide

sodium chloride glycerine

sodium sulphate

copper sulphate alcohol

An examination of the column of electrolytes shows that
it contains only acids, bases, and salts. All electrolytes
belong to these classes of compounds. The non-elec-
trolytes in the second column are not generally placed
in any of these classes of compounds. Why is it that
water solutions of acids, bases, and salts conduct elec-
tricity, while other substances fail? This question has
led to a more careful study of the properties of water
solutions.

97. Effect of Dissolved Solids on the Freezing-point. — Under
ordinary conditions of pressure pure water freezes at
0° C. and boils at 100° C. The addition of any soluble
solid lowers the freezing-point and raises the boiling-
point. We shall discuss only the freezing-point. All
water solutions freeze at a lower temperature than pure
water. Careful measurements show that the amount of
the lowering depends on the mass of the substance added.
A solution containing 10 grams of sugar to one liter of
water will freeze at a point twice as far below 0° as one
containing 5 grams of sugar to 1 liter of water. Ac-
cording to the atomic theory 10 grams of sugar contain
twice as many molecules as 5 grams. Therefore, the

112

SOLUTION.

lowering of the freezing-point of the pure solvent is pro-
portional to the number of sugar molecules dissolved.

SThe depression of the freezing-point of
water (Fig. 35) ^ produced by equal num-
bers of molecules of different substances
has been carefully determined. To get
equal numbers of molecules^ a weight in
grams equal to the molecular weight of
each substance was dissolved in one liter of
water. Thus 46 grams of alcohol, C^HgOH,
and 342 grams of sugar, C^^U^O^i, were
dissolved in a liter of water. In both cases
_ ^ the resulting solutions freeze at — 1.86°.

fit was further found that a solution of
any non-electrolyte containing the molec-
ular weight in grams of the solute to the
liter gave this same lowering of the freez-
ing-point. If weights of non-electrolytes
proportional to their molecular weights are
dissolved in equal volumes of the same
solvent, the solutions will freeze at the same
temperature. This is strictly true for di-
lute solutions only. This indicates that
thie depression of the freezing-point is de-
pendent on the number of molecules present
and not on the weight or kind of these molecules. But if
common salt is used, we find that, in dilute solution, the

Fig. 35.

1 Figure 85 represents an apparatus for determining the depression of the
freezing-point. A Beckmann thermometer (a) is used and the liquid to
be frozen is held in the inner tube (6). Through the side tube (e) is
introduced the substance whose effect on the freezing-point of the pure
solvent is to be determined. A stirring rod (d) moves up and down in
the solution which is cooled by a freezing mixture contained in the outer
vessel (6).

CHEMICAL ACTIVITY OF ELECTROLYTES. 113

freezing-point is lowered nearly twice as ranch as in the
case of an equimolecular solution of a non-electrolyte. . This
can only be explained by assuming the presence of twice as
many particles in a solution containing a molecular weight
of salt as in the same volume of a solution containing a
molecular weight of sugar. In other words, we have
reason to think that all or nearly all of the salt molecules
are separated, each into two particles, when they are dis-
solved in water. Other electrolytes behave as salt does,
while non-electrolytes act like sugar.

A careful study of the boiling-point shows twice as
great an elevation in the case of electrolytes as in non-
electrolytes, in proportion to the number of dissolved
molecules. Similar evidence is afforded by the
phenomena of osmotic pressure. The differences ^©n of
in the freezing-point, boiling-point, and osmotic «l«ctro-
pressure produced, by equal numbers of mole-
cules of electrolytes and non-electrolytes may all be ex-
plained • by assuming that the dissolved molecules in
electrolytes are broken up or dissociated into parts. In
non-electrolytes the molecules of the dissolved substance
aire not dissociated.

98. Chemical Activity of Electrol]^. — We have already
noted the fact that acids, bases, and salts are electrolytes ;
these are the substances that we have made use of in the
greater part of our chemical actions so far. Nearly all
chemical actions require the presence of water. When
two solutions of electrolytes are mixed, action takes place
at once if one of the products of the reaction is either in-
soluble or gaseous. So we may say that in general elec-
trolytes are very active chemically. Non-electrolytes
show very little activity.

114 SOLUTION.

99. Ions. — We find in the electrolysis of a dilute solu-
tion of hydrochloric acid that the hydrogen is liberated
at the cathode and chlorine at the anode. To these mov-
ing particles in the electrolyte Faraday gave the name
ion«, meaning wanderers. We have just seen that the
hydrochloric acid molecule is broken up, in the act of dis-
solving, into two particles, and since atoms are chemically
indivisible, these must be the hydrogen and chlorine
atoms. Why do they move in opposite directions ? Bodies
with electrical charges of opposite kind attract each other.
Since the hydrogen particles in a water solution are at-
tracted toward the cathode, or negative pole, when an
electric current is passed, these hydrogen particles must
be positively charged. Similarly the chlorine particles
must have negative charges, since they are drawn toward
the positive pole or anode.

When a solution of copper sulphate is electrolyzed,
copper appears at the cathode and experimental evidence
shows the presence of sulphur and oxygen at the anode.
The copper sulphate must therefore dissociate into Cu, a
positive (4-) ion, and SO^, a negative ( — ) ion. An ion^
then, is either an atom or group of atoms carrying an electric
charge,

100. Explanation of Electrolysis — Sodium chloride disso-
ciates in water into sodium and chlorine ions. From the
intensity with which sodium ordinarily reacts with water,
we might expect the instant formation of sodium hydrox-
ide. But this cannot be present, for the solution is per-
fectly neutral in reaction. The sodium ion does not react
with water. But when a current is passed through a solu-
tion, the positive charges of the sodium ions are neutralized
by the negative electrification of the cathode ; the sodium

DIFFERENCES BETWEEN ION AND ATOM. 115

ions then become sodium atoms and react with the water
surrounding the cathode, forming sodium hydroxide. It
should be understood, however, that the electric current
is not the cause of the dissociation into charged particles,
but simply determines the direction in which they move.
The dissociation took place while the substance was dis-
solving.

The chlorine ions give up their negative charges the
instant they touch the anode, neutralizing in part the
positive charge on that electrode. The particles of chlorine
without their charges are ordinary atoms, which unite in
pairs to form molecules of chlorine gas which bubbles oflF
at the anode.

-^' 101. Bifferenoes between Ion and Atom. — The existence
of an electric charge on an atom entirely changes its
properties, as may be seen from the electrolysis of sodium
chloride just described. Chlorine atoms decompose water
by combining with the hydrogen and setting free the
oxygen. Chlorine ions, however, move freely about in
water without decomposing it. Atoms of chlorine tend
to combine in pairs to form molecules of chlorine gas;
ions of chlorine repel each other because they possess like
charges.

The differences between ion and atom may now be
stated : first, an ion has an electric charge, an atom has
not ; second, a single ion often includes several atoms of
different elements. The electric charges on ions are very
heavy, hence their energy supply is very different from
that of atoms. The properties of bodies depend upon
their energy supply, and this accounts for the differ-
ence in activity of the ion and the atom of the same
element.

116 SOLUTION.

102. lonizatioiL of Acids and Bases. — When an electrolyte
is dissolved in water, some of its molecules enter the solu-
tion undissociated, while the other molecules separate into
particles which become electrically charged in the process.
On passing the current through a solution of hydrochloric
acid, we find hydrogen collecting at the cathode and chlo-
rine at the anode. In the solution there must be positively
charged hydrogen ions and negatively charged chlorine
ions. The electrolysis of nitric acid reveals the presence
of hydrogen ions and NOg ions; that of sulphuric acid
shows two hydrogen ions and one SO^ ion for each mole-
cule dissociated. The electrolysis of other acids gives, in
every case, hydrogen at the cathode and so indicates the
presence of hydrogen ions. To these are ascribed the
characteristic properties of acids. This hydrogen ion,
common to all acids, is responsible for the sour taste, and
the red effect on litmus. The presence of hydrogen in a
compound does not make it an acid ; the compound must
give hydrogen ions in water solution. Liquid hydrogen
chloride, free from water, fails to act as an acid, because
none of its molecules are separated into hydrogen ions.
An acid is a hydrogen compound whose water solution con-
tains hydrogen ions.

When bases are dissolved in water, the metallic atom
in the molecule becomes the positive ion and the hydroxyl
group the negative ion. The hydroxyl ions are the only
ions common to all bases, and to them the characteristic
properties of bases are attributed. Hence our definition :
A base is a hydroxide whose water solution contains hydroxyl
ions.

103. Effect of Dilution on Ionization. — It will be readily
seen that the more ions there are present in a given solu-

ACTIVITY OF ACIDS AND BASES. 117

tion, the better it will conduct the current, for the ions
act as carriers. So we can measure the degree of ioniza-
tion by the conducting power of the solution, provided
there is the same weight of solute between the electrodes
in every case. Conductivity under these conditions is
called molecular conductivity. We find, as we continue
to dilute an electrolyte, the molecular conductivity in-
creases up to a certain point, beyond which it does not
change. This might be expected, for at a certain dilution
all the molecules are dissociated, and a further dilution
causes no further dissociation. The effect of dilution
on the ionization of chloracetic acid is shown in the
following table :

DiLVTioir. Ionization.

47.25 grams per liter 6.4 %

11.81 grams per liter 10.67 %

2. 96 grams per liter 20. 00 %

104. Actiyity of Acids and Bases. — The chemical activity
of an acid depends upon the extent to which its dissolved
molecules are dissociated into ions. Hydrochloric acid is
almost wholly dissociated into its ions in dilute solution.
This is why it acts vigorously on metals and neutralizes
bases. Such an acid is spoken of as a strong acid. Nitric
acid is another example of a strong acid. Sulphuric acid,
which is dissociated bujt two-thirds as much as the acids just
mentioned, is not so strong an acid. Acetic acid is the
type of a weak acid, because it is so slightly dissociated
(less than 2% in dilute solution) that it does not act
vigorously on most metals, and its reactions with other
substances are slow.

The term strength of acid must not be confused with

118 SOLUTION.

the term concentrated acid. Concentrated sulphuric acid
usually contains about 98% HgSO^, while concentrated
hydrochloric contains but 37% HCl. Sulphuric acid,
then, is usually the more concentrated, but it is the
weaker acid of the two. Hydrochloric acid is the more
reactive, since it is dissociated to a greater extent in water
solution. The reason for the displacement of another
acid by sulphuric acid is not that sulphuric acid is a
stronger acid, but that it has a higher boiling-point.

As in the case of acids, the strength of a base depends
upon the degree of ionization it undergoes in solution.
Sodium hydroxide and potassium hydroxide are almost
completely dissociated in dilute solutions. They are types
of strong bases.

Its. Explanation of STentralization. — When we mix dilute
solutions of hydrochloric acid and sodium hydroxide, we
have present in the mixture the positive ions, hydrogen
and sodium, and the negative ions, chlorine and hydroxyl.
Any change that takes place while the acid and base are in
solution must consist in the union of oppositely charged
ions to form undissociated or electrically neutral mole-
cules. We have already seen that water is a non-electro-
lyte and can only be ionized to a very small extent. Now
we have present in our solution hydrogen ions, each with
one positive charge, and hydroxyl ions, each with an equal
charge of negative electricity. Each positive hydrogen
ion will attract a negative hydroxyl ion, and vice vena.
The two ions combine, the equal opposite charges neutral-
ize each other, and a molecule of undissociated water
results. Indicating the sign of the charge by a -i- or —
above and to the right of the symbol, we may express this
charge by the equation:

PB0DUCT8 OF NEUTRALIZATION. 119

H++OH-— ^HOH

As the water formed is practically undissociated (only two
molecules in a billion), we may consider it as completely
removed from the action as if it had formed an insoluble
compoimd or precipitate. In the neutralization, for each
H+ion withdrawn in this way a Cl~ion is left, and
for each OH~ion a Na'^ion remains. These will not unite
permanently so long as water is present, for the sodium
chloride is dissociated as fast as it is formed. When
the neutralization is complete, there are neither H"*" nor
OH~ ions left to give an acid or a basic reaction, and the
solution contains equal numbers of Na"*" and CI" ions. The
equation showing these facts is:

H+ + CI- + Na+ + OH- — ^ HOH + Na++ Cl"

If the solution is concentrated, the degree of dissociation
is lessened, and Na"^and CI" ions unite to form undis-
sociated sodium chloride. When evaporation is com-
plete, we find that they have completely united, forming
crystals of salt.

106. Products of STeutralization. — We get similar results
from the use of other acids and bases. Potassium hydrox-
ide and hydrochloric acid give water, K"*" ions and Cl"
ions. Nitric acid and sodium hydroxide yield undisso-
ciated water, Na^ ions and NOf ions.

K+-i-OH- + H++Cl- — >-HOH-i-K+ -i-Cl"
Na+ + OH- -i- H+H- NOg" —^ HOH + Na+ + NO^

In every case of neutralization the products are:

(1) undissociated water ;

(2) a solution containing positive ions from the base

and negative ions from the acid;

(3) energy in the form of heat.

120 SOLUTION.

During the evaporation of the solvent these ions unite to
form a compound known as a %alt.

107. Heat of STeutralization. — Neutralization, as a chemical
process, is essentially the formation of undissociated water
molecules. It is always accompanied by the liberation of
heat. When dilute solutions, containing equal amounts of
a strong acid, are neutralized by different strong bases, the
same quantity of heat is produced in each case. This heat
is known as the heat of nevtralization.

Table showing heat of neutralization expressed in calo-
ries:

HCl HNOg

NaOH 13,700 cal. 13,700 cal.

KOH 13,700 cal. 13,700 cal.

This uniformity in the heat of neutralization indicates an
action that is common to all these cases. The only com-
mon product is water. Therefore^ the heat of neutraliza-
tion is the heat of formation of water from hydrogen and
hydroxyl ions,

108. Other Actions between Ions. — Most chemical actions
used in the laboratory take place in solution or in the pres-
ence of water. These are usually the reactions between
the ions of acids, bases, and salts which are largely disso-
ciated in solution. Such reactions are widely used to
produce characteristic insoluble compounds in analytical
chemistry.

The test for hydrochloric acid or a chloride is a search
for the presence of chlorine ions. The solution of silver
nitrate used in the test contains silver ions, viz. ;

AgNOgZ^ZtAg^ + NOi"

OTHER ACTIONti BETWEEN I0N8. 121

These positive silver ions will encounter negative chlorine
ions if the solution tested contains a chloride. Silver
chloride is formed. This compound, being practically in-
soluble and hence undissociated, separates as a precipi-
tate. Therefore the equation is :

Ag+ + Cl-"— ^AgCl

A solution of potassium chlorate also contains chlorine,
not as a simple ion, but as part of the ion ClOg". So when
we mix this solution with the silver nitrate, we have the
ions Ag+, NO^, K+, ClOg-. Here no precipitation takes
place, since the compounds, silver chlorate, AgClOg, and
potassium nitrate, KNOg, that would be likely to form, are
both soluble. Silver nitrate solution, then, is the test
for the chlorine zow, and not for the chlorine atom.

The test for a sulphate depends upon the combination
of barium ions, Ba"^"*" (from the barium chloride added),
with the SO^ ion of the sulphate. The barium sul-
phate formed, is insoluble and sepamtes as undissociated
molecules. The equation is :

Ba++ -h SO^-— >. BaSO^

When the combination of ions in a chemical action
produces an insoluble compound, the action tends to
complete itself, or "run to an end. " The precipitation of
the undissociated molecules gradually removes the com-
bining ions from solution and allows the further dissocia-
tion of the compounds producing them. The reaction is
at an end when the supply of either of the combining ions
is exhausted.

A reaction also completes itself when a combination of
ions produces a substance, volatile under the existing
conditions. The combining ions leave the solution in the
undissociated molecules of the volatile substance.

122

SOLUTION.

109. Chargei carried by loiu. — We have seen that one
molecule of hydrochloric acid neutralizes one molecule of
sodium hydroxide, producing one molecule of water. But
we find that to neutralize one molecule of barium hydrox-
ide, Ba(0H)3, two molecules of hydrochloric acid are re-
quired, according to the equation:

Ba(0H)3 -h 2 HCl -^ 2 HOH -h BaCl^
When neutralization is complete, the barium chloride is
largely dissociated into barium and chlorine ions. In any
solution the number of positive charges must equal the num-
ber of negative charges. Therefore, each barium ion must
contain two positive charges to neutralize the negative
charges on the two chlorine ions, with which it will unite
when the solution is evaporated to get the salt. Sulphuric
acid dissociates into H"*", H% and SO^ ions; the
SO4 ion must carry two negative charges to balance
the positive charges of the two hydrogen ions.
H2SO4 11^ H+ -h H+ + SO4—
Some other ions carry three and even four charges.

No. OP

CHABGB8.

P08ITITB Ions.

Nboativk Ions.

H, Na, K, Li, Ag, NH^,
Cu (cuprous), Hg (mercurous)

Ca, Sr, Ba, Mg, Zn, Hg, Co, Ni,
Cu (cupric), Fe (ferrous),
Sn (stannous)

Al, Bi, Sb, Fe (ferric)

Sn (stannic)

OH, F, CI, Br, I, NOj,
ClOg, and other nega-
tive ions of mono-
basic acids

S, SO4, COj, and other
negative ions of di-
basic acids

PO4 and other negative

ions of tribasie acids
.J

Si04 and negative ions
of tetrabasic acids

VALENCE. 123

The preceding table shows the common ions with the
number of charges carried by each and the signs of the
charges.

110. Valence. — Atoms differ with respect to the num-
ber of atoms of other elements with which they com-
bine. The difference in this combining power is indicated
by the term valence. We have already seen that the hy-
drochloric acid molecule consists of one atom of hydrogen
and one atom of chlorine. As this is the simplest combin-
ing relation possible, we take the valence of each of these
elements as unity and call them univalent elements, since
no other elements are found to have a less combining
power. When one atom of any element combines with
one atom of hydrogen or of chlorine, the valence of the
element is one. The formulas NaCl, KCl, HBr, HI, show
that sodium, potassium, bromine, and iodine are univalent
elements. One atom of oxygen unites with two atoms of
hydrogen, so the valence of oxygen is two. An atom of
calcium replaces two atoms of hydrogen in hydrochloric
acid and unites with two atoms of chlorine;

Ca -h 2 HCl -^ CaClg -h H^

Hence the valence of calcium is two^ or it is bivalent. We
see in this case that the number of atoms of hydrogen
replaced by one atom of an element measures the valence
of that element. The valence of an element is the number
of hydrogen atoms which its atom will unite with or will
replace.

In ammonia, NHg, one atom of nitrogen is combined
with three hyarogen atoms ; in this compound, nitrogen .
has a valence of three, or it is trivalent. Marsh gas, CH^,
contains four atoms of hydrogen combined with one atom

124 SOLUTION.

of carbon; therefore carbon has a valence oifour; it is
quadrivalefU.

In electrolytes, the valence of each ion is numerically
equal to the number of charges carried by it. This, of
course, does not explain the valence of atoms in non-
electrolytes, but does indicate a connection between the
valence and the electrical capacity of an atom.

111. VariationB in Valence. — An element may have
more than one valence, according to the element with
which it combines and the conditions under which com-
bination takes place. The valence of hydrogen is always
regarded as one, and that of oxygen generallj'^ as two.
Sulphur has a valence of two in hydrogen sulphide, HgS ;
in sulphur dioxide, SOj, its valence is four^ being twice
that of oxygen ; in sulphur trioxide, SOg, its valence is
six.

The changes in valence, which take place during certain
reactions, give an extended meaning to the , terms oodda-
tion B,nd reduction^ because such changes in valence are
often brought about by the action of oxidizing and reduc-^
ing agents :

SO2-I-O — ^-SOg (oxidation) ^
2 BaO, — >-2 BaO + O^j (reduction)
Hence the terms oxidation and reduction are often used
Oxidatioii *^ express increase or decrease in valence. In this
and redue- sense, oxidation means an increase in the valence^
^^* and reduction a decrease in the valence^ of an di^r

ment. The terms, then, may be applied to reactions ip
which oxygen take's Ho part* Thus if ferric chloride,
FeClg, is treated with nascent hydrogen, ferrous chlQridj^,
.FeCl^^ and hydrochloric acid are produced: - :

FeGlg + H— ^FeClj-hHCl : . .

8UMMABT. 125

The iron atom is reduced from a valence of three in the
ferric compound to a valence of two in the ferrous com-
pound. By adding an oxidizing agent, such as nitric
acid, to the mixture of ferrous chloride and hydrochloric
acid, the ferrous chloride is oxidized to ferric chloride, by
the addition of an atom of chlorine, thus increasing the
valence of the iron :

2 FeClg + 2 HCl -h O -^ 2 FeQl, + H,0

112. Application of Valence. — Valence is useful in
writing formulas and equations. Thus, knowing that
both hydrogen and sodium are univalent, we know that
it will take two sodium atoms to replace the two hydrogen
atoms in sulphuric acid, HgSO^, and that sodium sulphate
will have the formula NajSO^. Similarly, as the valence
of zinc is two, zinc sulphate is represented by the formula
ZnSO^.

Summary.

Solutions are classified as electrolytes and non-electrolytes ^ accord-
ing to their conducting power. The electrolytes include solutions
of adds, bases, and salts.

The effect of dissolved substances on the boiling-point and freez-
ing-point indicate that the molecules of the solute are dissociated
in the case of electrolytes. Electrolytes are more active chemically
than non-electrolytes.

The portions into which a molecule dissociates are ions and
carry equal and opposite electric charges. When an electric current
is passed through an electrol3rte, each ion passes to the electrode
of opposite sign and is there discharged and liberated. Acids fur-
nish hydrogen ions in solution ; bases^ hydroxyl ions, and salts fur-
nish other ions. The percentage of ionization increases with the
dSution.

The strong or a^ive acids and bases are those which are highly

126 SOLUTION.

dissodated. The essential action in netUralization is the union of
the hydrogen and hydroxyl ions to form undissoeiated water ; dur-
ing the evaporation, the union of the other ions to form a salt is
completed. The strong acids and bases unite with the same heat
of neutralization.

Chemical tests are often tests for ions. Hydrogen and metallic
ions carry positive charges ; the non-metallic ions negative charges.

Elements differ in their valence. The hydrogen atom is the
standard of valente. An element may have more than one
valence under different conditions. The terms oxidation and
reduction are sometimes applied to changes of valence.

BzerclBes.

1. How do you determine whether a given solution is an
electrolyte ?

2. When a beaker of acetic acid is put in circuit with an
incandescent lamp, the lamp glows feebly, while the solution
of sodium acetate allows it to glow brightly. What does this
show?

3. Compare the freezing-point of sea-water with that of
ordinary rain-water.

4. Why will substances often react with each other in
solution while they will not in a dry state?

5. What ions are present in solutions of the following:
KCl, ZnS04, KClOs, NaOH.

6. Distinguish carefully between an ion of potassium and
an atom of potassium.

7. Why is nitric acid a more active acid than sulphuric acid?

8. Give the changes that take place and the products
formed in the following cases:

(a) when nitric acid is added to water;

(b) when a piece of caustic potash is dissolved in water ;

EXERCISES. 127

9. A piece of red litmvis turns blue in a solution of sodium
carbonate (KagCOa). The presence of what ions is shown by
this test ?

10. Write equations, indicating ions, for the neutralization
of potassium hydroxide with sulphuric acid; of nitric acid
with ammonium hydroxide.

11. Will sodium chlorate (NaClOj) give the test for chlo-
rine ions with silver nitrate ?

12. Give two conditions under which a reaction goes on to
completion, and explain each.

CHAPTER XVI.
SODIUM AND POTASSIUM OOMPOUNDS.

113. General Properties. — Just as sodium and potassium
resemble each other very strongly, so the corresponding
compounds which the two elements form have many points
of similarity. For this reason these compounds can be
advantageously studied together.

Most sodium and potassium compounds are white crys-
talline substances; practically all of them are soluble in
water. With the exception of the hydroxides, all of those
that we shall study are salts, possessing in a marked degree
those properties which are characteristic of this class of
bodies. As a rule they are very stable compounds. They
are among our most common and useful substances.

Sodium compounds are generally less soluble in water
than are the corresponding potassium compounds; they
are, therefore, not so satisfactory for certain uses. On the
other hand, sodium salts are usually cheaper than those of
potassium. Moreover, since the atomic weight of sodium
is 28, while that of potassium is 39, a gram of sodium salt
contains a greater number of molecules than a gram of
the corresponding potassium compound. Consequently a
gram of sodium salt will "go farther" than a gram of the
same potassium salt. For these reasons sodium com-
pounds are generally used in manufacturing operations in
preference to those of potassium.

114. Preparation of the Hydroxides (ITaOH, KOH). —

Two methods are in use for the manufacture of these

128

PREPARATION OF THE HTBB0XIDE8.

129

hydroxides. In describing the operations, the sodium
compound will be taken as a type for both.

We saw (page 67) that a water solution of sodium chlo-
ride, on being decomposed by an electric current, gives
chlorine, hydrogen, and sodium hydroxide as products :

2 NaCl + 2 H20-^2 NaOH + CI, + H,

The importance of sodium hydroxide is such that this
operation is now carried out on a manufacturing scale.
The appcwatus shown in the diagram (Fig. 86) Bj^^^yQi.
is used. The anodes are placed in solutions of sis of
salt in compartments AA^ and the cathodes in ™^*'
B. A layer of mercury (MHP) permits the passage of a
current between A and -B, but prevents the mixing of
the solutions.

Fig. 36. Electrolytic preparation of sodium hydroxide.

On passing the current, chlorine is liberated at the anode,
and passes off through a pipe, not shown in the figure.
The positive sodium ions, moving toward the cathode,
discharge against the mercury, and the metallic sodium,
dissolving in the mercury, forms an amalgam which floats
on the surface of the mercury.

180 SODIUM AND POTASSIUM COMPOUNDS.

By rocking the tank by means of the cam (7, the
sodium amalgam is passed into the middle compartment
-B, where it reacts with the water, forming sodium hydrox-
ide. The hydrogen which is set free is withdrawn by a
pipe not shown in the figure.

In this way a strong solution of pure sodium hydroxide
is obtained. The solution is drawn off, the water evapo-
rated, and the sodium hydroxide obtained as a white solid.
This is melted and cast into sticks or run into iron
drums. The chlorine which is obtained as a by-product
is used to make bleaching powder.

Calcium hydroxide, Ca(0H)2 (ordinary slaked lime), is
Lye made from calcium oxide, CaO (unslaked lime),

prooeu. by adding water :

CaO -h H2O -^ Ca (OH),

Calcium hydroxide is slightly soluble in water. If a solu-
tion of it is boiled with sodium carbonate, NagCOg, the
following reaction occurs :

Na^COg + Ca(0H)2 — ^ 2NaOH -f- CaCOg

The calcium carbonate, CaCOg, formed is insoluble in
water. It is this fact that makes the action possible.
The solution of sodium hydroxide is separated from the
precipitated calcium carbonate and evaporated to dry-
ness.

115. Properties of the Hydroxides. — The hydroxides are
very strong bases. As we have seen, they neutralize
acids, forming salts and water. These bases have such
a corrosive action on animal and vegetable matter that
they are called eaustie alkalies. Glass is attacked by
them; and although the action is somewhat slow, a solu-

SOURCES OF SODIUM CHLORIDE. 131

tion of either hydroxide on standing in a glass bottle
becomes quite impure.

Sodium hydroxide, eaustic 9odc^ is manufactured in enor-
mous quantities for use in soap-making. It is also used in
making bleaching solutions and in numerous
other operations. Potassium hydroxide, caustic
potash^ is not so extensively used in manufacturing opera-
tions. It is used to make other compounds of potassium,
and in the preparation of some soaps.

116. Sources of Sodium Chloride. — Sodium chloride,
NaCl, common salt^ is the most abundant sodium compound
found in nature. Rock salt, or halite, is found in many
countries, but the largest deposits are those in
Louisiana, Austria, Germany, and Spain. Often
in these beds the salt is of such purity that it has only to
be mined and crushed to be ready for use.

Much of the salt in this country is obtained from salt
wells in New York, Michigan, Ohio, and several other
states. A boring is made, and water sent down
to the salt bed. There brine is formed and
pumped to the surface. The earthy impurities in the brine
settle, and then it is evaporated. A fairly pure salt is
thus obtained.

The total amount of salt found in deposits, however, is
insignificant compared with the quantity contained in the
seas and oceans. The percentage of salt in sea-water
is small, yet it has been computed from the average
per cent that the total quantity in the sea is
36,000,000,000^000,000 tons. The percentage "®*-^*^'^-
of salt is not uniform for all seas and oceans. These
variai^ions depend upon the ratio existing between the
ambunt of water delivered by rivers to a sea and the

122

SOLUTION.

109. Charges carried by Ions. — We have seen that one
molecule of hydrochloric acid neutralizes one molecule of
sodium hydroxide, producing one molecule of water. But
we find that to neutralize one molecule of barium hydrox-
ide, Ba(0H)2, two molecules of hydrochloric acid are re-
quired, according to the equation:

Ba(0H)2 -h 2 HCl -^ 2 HOH + BaCl^
When neutralization is complete, the barium chloride is
largely dissociated into barium and chlorine ions. In any
solution the number of positive charges must equal the num-
ber of negative charges. Therefore, each barium ion must
contain two positive charges to neutralize the negative
charges on the two chlorine ions, with which it will unite
when the solution is evaporated to get the salt. Sulphuric
acid dissociates into H"^, H% and SO4 ions; the
SO4 ion must carry two negative charges to balance
the positive charges of the two hydrogen ions.
H2SO4 11^ H+ + H+ + SO4—
Some other ions carry three and even four charges.

No. OP

CHAEGB8.

POSITIVB ION8.

Negative Ions.

3
4

H, Na, K, Li, Ag, NH„
Cu (cuprous), Hg (mercurous)

Ca, Sr, Ba, Mg, Zn, Hg, Co, Ni,
Cu (cupric), Fe (ferrous),
Sn (stannous)

Al, Bi, Sb, Fe (ferric)
Sn (stannic)

OH, F, CI, Br, I, NOg,
ClOg, and other nega-
tive ions of mono-
basic acids

S, SO4, CO3, and other
negative ions of di-
basic acids

PO4 and other negative
ions of tribasie acids

SiO* and negative ions
of tetrabasic acids

VALENCE. 123

The preceding table shows the common ions with the
number of charges carried by each and the signs of the
charges.

110. Valence. — Atoms differ with respect to the num-
ber of atoms of other elements with which they com-
bine. The difference in this combining power is indicated
by the term valence. We have already seen that the hy-
drochloric acid molecule consists of one atom of hydrogen
and one atom of chlorine. As this is the simplest combin-
ing relation possible, we take the valence of each of these
elements as unity and call them univalent elements, since
no other elements are found to have a less combining
power. When one atom of any element combines with
one atom of hydrogen or of chlorine, the valence of the
element is one. The formulas NaCl, KCl, HBr, HI, show
ihat sodium, potassium, bromine, and iodine are univalent
elements. One atom of oxygen unites with two atoms of
hydrogen, so the valence of oxygen is two. An atom of
calcium replaces two atoms of hydrogen in hydrochloric
acid and unites with two atoms of chlorine:

Ca -h 2 HCl -^ CaClj + Hg

Hence the valence 6f calcium is two^ or it is bivalent. We
see in this case that the number of atoms of hydrogen
replaced by one atom of an element measures the valence
of that element. The valence of an element is the number
of hydrogen atoms which its atom will unite with or will
replace.

In ammonia, NHg, one atom of nitrogen is combined
with three hyarogen atoms ; in this compound, nitrogen .
has a valence of three^ or it is trivalent. Marsh gas, CH^,
contains four atoms of hydrogen combined with one aton

124 SOLUTION.

of carbon ; therefore carbon has a valence of four ; it is
quadrivalent.

111. VariationB in Valence. — An element may have
more than one valence, according to the element with
which it combines and the conditions under which com-
bination takes place. The valence of hydrogen is always
regarded as one, and that of oxygen generally as two.
Sulphur has a valence of two in hydrogen sulphide, HgS ;
in sulphur dioxide, SOj, its valence is four^ being twice
that of oxygen; in sulphur trioxide, SOg^ its valence is
9ix.

The changes in valence, which take place during certain
reactions, give an extended meaning to the , terms oxida-
tion a,nd reduction^ because such changes in valence are
often brought about by the action of oxidizing and. reduc^
ing agents :

SO2+O — ^-SOg (oxidation) ^
2 BaO, — >- 2 BaO + O^j (reduction)
Hence the terms oxidation md reduction are often used
Oxidatioii *^ express increase or decrease in yalence. In this
aAd redue- sense, oxidation means an increase in the valence^
^^' and reduction a S^orease in the valence^ of an deT

ment. The terms, then, may be applied to reactions in
which oxygen take's tio part. Thus if ferric chloride,
FeClg, is treated with nascent hydrogen, ferrous ehlqridj^,
FeGl^, and hydrochloric ^cid are produced: :: '

FeOlj + H— ^FeCljjH-HCl .. , .

SUMMART. 125

2 FeClj + 2 HCl -h O -^ 2 FeQl, + H,0

112. Application of Valenoe. — Valence is useful in
writing formulas and equations. Thus, knowing that
both hydrogen and sodium are univalent, we know that
it will take two sodium atoms to replace the two hydrogen
atoms in sulphuric acid, HjSO^, and that sodium sulphate
will have the formula NajSO^. Similarly, as the valence
of zinc is two, zinc sulphate is represented by the formula
ZnSO^.

Summary.

Solutions are classified as electrolytes and non-electrolytes , accord-
ing to their conducting power. The electrolytes include solutions
0^ adds, bases, and salts.

The effect of dissolved svhstances on the boiling-point and freez-
ing-point indicate that the molecules of the solute are dissociated
in the case of electrolytes. Electrolytes are more active chemically
than non-electrolytes.

The portions into which a molecule dissociates are ions and
carry equal and opposite electric charges. When an electric current
is passed through an electrolyte, each ion passes to the electrode
of opposite sign and is there discharged and liberated. Acids fur-
nish hydrogen ions in solution ; bcLses, hydroxyl ions^ and saUs fur-
nish other ions. The percentage of ionization increases with the
dMution.

The strong or a>ctive adds and ba^es are those which are highly

126 SOLUTION.

dissodated. The essential action in neutralization is the union of
the hydrogen and hydroxyl ions to form undissociated water ; dur-
ing the evaporation, the union of the other ions to form a salt is
completed. The strong acids and bases unite with the same heat
of neutralization.

Chemical tests are often tests for ions. Hydrogen and metallic
ions carry positive charges ; the non-metallic ions negative charges.

BzerclBes.

1. How do you determine whether a given solution is an
electrolyte ?

2. When a beaker of acetic acid is put in circuit with an
incandescent lamp, the lamp glows feebly, while the solution
of sodium acetate allows it to glow brightly. What does this
show?

3. Compare the freezing-point of sea-water with that of
ordinary rain-water.

4. Why will substances often react with each other in
solution while they will not in a dry state?

5. What ions are present in solutions of the following:
KCl, ZnS04, KClOa, NaOH.

6. Distinguish carefully between an ion of potassium and
an atom of potassium.

7. Why is nitric acid a more active acid than sulphuric acid?

8. Give the changes that take place and the products
formed in the following cases:

(a) when nitric acid is added to water ;

(b) when a piece of caustic potash is dissolved in water ;

EXERCISES. 127

9. A piece of red litmus turns blue in a solution of sodium
carbonate (Na2C08). The presence of what ions is shown by
this test ?

10. Write equations, indicating ions, for the neutralization
of potassium hydroxide with sulphuric acid; of nitric acid
with ammonium hydroxide.

11. Will sodium chlorate (NaClOs) give the test for chlo-
rine ions with silver nitrate ?

12. Give two conditions under which a reaction goes on to
completion, and explain each.

138 SODIUM AND POTASSIUM COMPOUNDS.

western coast of South America. These beds lie near
the boundary lines of Peru, Chile, and Bolivia and have
been the cause of many disputes between these
countries. The boundary lines have now been
so adjusted that Chile owns the greater portion of these
extremely valuable deposits. The crude nitrate is obtained
by crushing the loose, rocky material of the beds and boil-
ing it in hot water. The liquor containing the nitrate is
run oflf and allowed to crystallize. The product, crude
Chile saltpeter, contains 94 % to 98 % of sodium nitrate.
A purer quality is obtained by recrystallization.

124. Mannfactnre of Potassium Nitrate. — Most of the
potassium nitrate now used is prepared from sodium
nitrate. The potassium compound is made by mixing
hot, concentrated solutions of sodium nitrate and potas-
sium chloride. The equation for the reaction is:

KCl + NaNOg-^ KNO3 + NaCl

On evaporation the boiling mixture first deposits common
salt, since this substance is less soluble in boiling water
than is potassium nitrate. On cooling the solution, how-
ever, the potassium nitrate crystallizes out, because it is
far less soluble in cold than in hot water. A purer quality
of potassium nitrate can be obtained by recrystallization.

125. Properties and Uses of the Nitrates. — Sodium and
potassium nitrates are white, soluble salts. Sodium nitrate
differs from the potassium compound in crystalline form
and in being hygroscopic. Both nitrates give off oxygen
when heated, leaving compounds containing less oxygen,
known as nitrites.

The principal use of potassium nitrate (ordinary salt-
peter) is in the manufacture of black gunpowder. Sodium

8UMMABT. 139

nitrate is too hygroscopic for this purpose. Potassium
nitrate is used to preserve meat, and corned beef owes its
red color to this treatment. The cheaper sodium nitrate
has replaced potassium nitrate in the manufacture of nitric
acid and its derivatives. Chile saltpeter is also used as a
fertilizer, and in the manufacture of sodium nitrite, a most
important substance in the manufacture of aniline dyes.

Summary.

The hydroxides of sodium and potassium are prepared by the
electrolysis of solutions of the chlorides, or by the reaction of the
carbonate with slaked lime.

The chlorides of these two metals occur in nature.

They are typical salts. They have a saline taste, are neutral,
and result from the action of the corresponding acid and base.

Their uses are dietary and preservative and for the preparation
of other sodium and potassium compounds.

The carbonates occur in plant ashes : sodium carbonate in marine
plants; potassium carbonate in land plants.

They are commercially prepared by the decomposition of the
bicarbonates by heat.

The carbonates are used in the preparation of soaps, washing-
powders, glass, and other compounds.

The bicarbonates are made commercially by the action of carbon
dioxide with ammoniacal brines.

They are only slightly soluble and are less corrosive than the
carbonates.

The bicarbonates are used in the preparation of the carbonates.
Sodium bicarbonate is a constituent of all baking powders and is
used in the preparation of some aerated liquors.

Sodium nitrate is obtained from Chile. Potassium nitrate is
formed by the reaction of potassium chloride with sodium nitrate.

Sodium nitrate is used to prepare potassium nitrate, also as a

140 SOLIUM AND POTASSIUM COMPOUNDS.

fertilizer and as a source of nitric acid. Potassium nitrate is used
in gunpowder.

Bxercises.

1. How could you tell potassium nitrate from sodium nitrate ?

2. How are the nitrates produced in nature ?

3. How could you show that there are potassium compounds
in plants ?

4. Why was sodium carbonate obtained from the ashes of
sea plants and not from the water directly?

5. Write the equations for tfie preparation of potassium car-
bonate by the Solvay process.

6. In the Solvay process, why is not the normal carbonate
produced directly, instead of forming the bicarbonate and then
decomposing it ?

7. What advantage has baking soda over potassium bicar-
bonate?

a Equal weights being taken, which will neutralize the larger
quantity of acid, sodium hydroxide or potassium hydroxide ?
sodium carbonate or bicarbonate ?

9. Why is " soda " used in cleaning ?

CHAPTER XVII.

SlTLPHim AND HTDBOGEV SULPHIDE.
SULPHUR.

126. Oocnrrence. — Sulphur is found in nature both free
and in combination with other elements. In the uncom-
bined or native state it is found in volcanic regions.
A specimen of native sulphur shows the element permeat-
ing a mass of calcium sulphate or calcium carbonate.
Sulphur is also formed by chemical decomposition brought
about by bacteria. Important deposits of this character
occur in Louisiana.

Sulphur is found combined with many difiPerent metals,
as sulphides ; those of iron, copper, lead, and zinc are the
most abundant. The sulphates of a few metals are found
in considerable quantities.

127. Preparation. — The presence of uncombined sul-
phur in the neighborhood of volcanoes is probably ac-
counted for by the interaction of two gases, hydrogen
sulphide and sulphur dioxide, both of which are likely to
be produced by volcanic action :

SO2+ H3S-^2S + H20

This reaction can easily be shown in the laboratory by
bringing bottles of the two moist gases mouth to mouth.
An action immediately occurs, and as a result sulphur is
deposited on the sides of the bottles. The same thing
can be shown in a m&re marked manner by using the
apparatus represented in Figure 38.

141

142

SULPHUR AND HYDROGEN SULPHIDE.

Commercial Forms,
commerce. It can
sulphur and flowers
pared in the same
traction of sulphur
is associated. The

, — Sulphur is an important article of
be bought in two different forms, roll
of sulphur. The two forms are pre-
operation. This begins with the ex- <
from the rocky material with which it
ore is heaped in piles and heated with-

Fig. 38. Sulphur deposition.

fl, hydrogen sulphide generator; b, moistened glass wool; c, jar with deposited sulphur;
d, sulphur dioxide generator.

out access of air, until the sulphur melts and thus sepa-
rates from the rock (Fig. 39). The sulphur is then
further purified by heating it in iron vessels (retorts) with-
out access of air, until it is vaporized (Fig. 40). The
vapor is led into brick-lined rooms. A part of the gas is
suddenly cooled and is deposited as a fine powder on the
sides of the room. This form is known as flowers of sul-

ALLOTROPIC FORMS.

148

phur. Most of the sulphur vapor condenses in the liquid
form, and flows down to the
outlet of the condensing
chamber. It is then run
into wooden moulds, where
it solidifies into cylindrical-
shaped rods about an inch
and a half in diameter. This
form is called roll sulphur or
brimstone. Flowers of sul-
phur differ slightly in chem-
ical properties from roll
sulphur.

Fig. 39. Extraction of sulphur.
a, covering of spent ore ; b, draught holes;
c, slope down which sulphur runs ; d, col-
lecting pool

Allotropic Forms. — No element displays a greater

variety of allotropic forms
than sulphur. Three of
these are well known and
easily obtained by the fol-
lowing methods:

(1) Crystallization
from carbon disulphide.

Carbon disulphide dis-
solves powdered roll sul-
phur very readily. The
solvent evaporates
quickly at ordinary tem-
peratures; so that if
such a solution stands,
the sulphur soon begins
to be deposited, just as
3~ salt is deposited when a
Fig. 40. Purification of sulphur. solution of salt is evapo-

144

SULPHUR AND HYDROGEN SULPHIDE.

Fig. 41. Rhombic sulphur crystals.

rated. The sulphur is deposited in crystals of a beau-
tifully regular octahedral shape. This form is ortho-
rhombic sulphur (Fig. 41). It is soluble in carbon
disulphide, and has a density of 2.01.

(2) Crystallization of melted sulphur.

Sulphur is melted and

then allowed to cool un-
til crystals appear at the
surface ; on pouring off
the still liquid sulphur,
the solid part is found in
crystals shaped like long,
narrow prisms with
sharp ends (Fig. 42).
This kind of sulphur is
known as prismatic sul-
phur. Its density is 1.96, and it differs in other prop-
erties from the rhombic variety. The prismatic form is
unstable. On standing a few days, its crystals lose their
transparency, become more brittle, and increase in density.
Examination by polarized light shows that the long, nar-
row prisms have broken up into minute rhombic crystals.

(3) Sudden cooling of boiling sulphur.

By application of considerable heat sulphur can be made
to boil. Before it reaches its boiling-point (450°) it goes
through some interesting and unusual changes* When
just above its melting-point, sulphur is a mobile liquid of
a light amber color. As the temperature rises it -darkens
rapidly and thickens so that it can hardly be poured from
the inverted test-tube ; on further heating the sulphur
again becomes less viscous, and finally boils, forming a
pale yellowish-brown vapor.

When boiling sulphur is poured into cold water, the

ALLOTROPIC FORMS.

145

cooled sulphur assumes a form (Fig. 43) altogether diflfer-
rent in appearance from those already described. It is
without crystalline form, of a rubber-like consistency, and
light amber in color. Because of its lack of crystalline
form it is called amorphous sulphur. This differs from
the crystalline form in being insoluble in carbon disul-
phide. It is sometimes spoken of as plastic sulphur.

Fig. 42. Prismatic sulphur.

Like prismatic sulphur, it is unstable and is changed
in the course of a few days into the stable rhombic form.
In this change the amorphous sulphur loses its plastic
character and becomes soluble in carbon disulphide.

If sulphur is dissolved in some alkali, as sodium hy-
droxide, and hydrochloric acid added to the solution, a
white, finely divided precipitate is obtained. This precipi-

146

SULPHUR AND HYDROGEN SULPHIDE.

tate is a form of amorphous sulphur. When shaken
with water it gives a fluid known as milk of sulphur.

129. Stability of the Allotropic Forms. — At temperatures
up to 96° C. all forms of sulphur tend to assume the rhom-
bic form. From 96° to 114°
the stable form is the pris-
matic variety. If crystals
of the rhombic variety are
put into a test-tube and
kept standing in boiling
water for several days, the
form will be changed to
minute crystals of the pris-
matic variety. Thus it is
seen that the form which
the sulphur assumes depends
on the temperature. It is
generally true that each of

the allotropic forms of any element is stable under certain
definite conditions.

Roll sulphur consists entirely of the rhombic variety.
Flowers of sulphur are made up of the rhombic form to-
gether with a smaller proportion of the amorphous form,
as is shown by the fact that flowers of sulphur are only
partly soluble in carbon disulphide.

130. Chemical Properties. — All forms of sulphur burn
readily when heated in the presence of oxygen; sulphur
dioxide is formed as a product of the action. Sulphur is
capable of taking oxygen away from compounds:

S + 02-^802

It does not, however, make a very satisfactory reducing

Fig. 43. Amorphous sulphur.

SULPHIDES.

147

agent, because it is apt to combine with the products of
the reduction.

Sulphur is a moderately active element. It combines
readily with many metals, often with the evolution of
light and heat. In a sense, then, it may be said to support
combustion. This can be shown by heating sulphur in a
test-tube until it boils and the sulphur vapor nearly fills
the tube ; on placing a strip of very thin copper in the
tube, it takes fire and burns, copper sulphide being formed:

Cu + S— ^CuS
Powdered zinc and sulphur, if mixed in certain propor-
tions, will unite almost explosively, with the formation of
zinc sulphide (Fig. 44).

Zn-hS — >-ZnS
Sulphur does not unite very readily with non-metals.
Such compounds are, in general, not very stable.

131. Sulphides. — The metallic sulphides form a very im-
portant class of compounds. Many metals are extracted
from sulphide ores. Most
of the sulphides, excepting
those of the extremely me-
tallic elements, are insolu-
ble in water. They usually
have characteristic colors.
These facts are useful to
the analytical chemist, for
he can recognize the pres-
ence of certain metals by
the color of their sulphides.

The blackening or tar-
nishing of metals is sometimes due to the action of sul-
phur. This is particularly true of silver. Sulphur or

Fig. 44.

148 SULPHUR AND HYDROGEN SULPHIDE.

sulphur compounds get into the air from illuminating gas
or burning coal and, acting on silver, produces the black
sulphide. Brass and copper are also readily tarnished by
sulphur and some of its compounds.

132. Besemblances of Sulphur to Other Elements. — In the
ease with which it combines with other elements, particu-
larly metals, sulphur resembles both chlorine and oxygen.
Generally it displays the closer resemblance to oxygen.
This is shown in comparing the solubility in water of
oxides and sulphides. Those metals which form insoluble
oxides also, as a rule, form insoluble sulphides.

133. Uses. — Sulphur is used in making certain kinds of
matches ; to furnish sulphur dioxide for bleaching or dis-
infecting purposes ; and in gunpowder. This explosive
is composed of potassium nitrate, sulphur, and carbon.
The chemical action that occurs in the act of explosion is,
in the main, an oxidizing and reducing action. The num-
ber of products formed is very large. The reaction is
partly represented by the equation:

2 KNOg + S + 2 C-^ K2SO4 + N2 -h 2 CO
or 2 KNOg + S -h 3 C -^ KgS + 3 COj + Ng

Sulphur or a compound of sulphur, the monochloride, is
used in the hardening or vulcanizing of india-rubber.

HYDROGEN SULPHIDE.

134. Formation. — Hydrogen sulphide, H2S, is formed
slowly and in small quantity when hydrogen is passed
over heated sulphur (Fig. 46). The presence of the
gas is shown by the blackening of a strip of paper (a a)
moistened with lead acetate. The difl&culty of forming
hydrogen sulphide by the direct union of the elements is

PBEPABATION OF HYDROGEN SULPHIDE. 149

Fig. 45.

in strong contrast to the great readiness with which the

hydrogen compounds, so far studied, are

formed. The lack of a strong tendency

of hydrogen and sulphur to combine

indicates the unstable character of their

product.

When albumen or other organic mat-
ter, containing sulphur, decays, hydrogen sulphide is one
of the substances formed. If a little sulphur is added
to fermenting sugar solution, hydrogen sulphide is
produced.

136. Preparation. — Hydrogen sulphide is readily formed
by the action of dilute hydrochloric acid on ferrous sul-
phide (Fig. 46):

FeS -h 2 HCl -^ FeClg + H2S

The hydrogen sulphide passes off
readily as a gas ; for this reason
the double decomposition pro-
ceeds rapidly. The sulphides of
some other metals could be used
in place of the ferrous sulphide,
and instead of hydrochloric acid
dilute sulphuric acid might be
. used. In the figure, a is the test-
tube generator containing iron
sulphide and dilute hydrochloric
acid, and h is the tube for col-
lecting the gas by downward dis-
placement.

V_y 136. Physical Properties. — Hy-

Fig. 46. drogen sulphide is a colorless gas.

150 SULPHUR AND HYDBOQEN SULPHIDE.

slightly heavier than air, having a characteristic odor
resembling rotten eggs. It dissolves to some extent in
water, and therefore is not usually collected over this
liquid. It is generally collected by the displacement of
air. A water solution of the gas is sometimes used in
the laboratory, but in such a solution the gas is decom-
posed in a few days by action with the oxygen of the air.
Besides being unpleasant, it is injurious to inhale the gas,
as headache and sickness are apt to result.

137. Chemical Properties. — The gas burns readily, if
there is an ample supply of oxygen, forming water and
sulphur dioxide ; with a limited quantity of oxygen, sul-
phur is formed :

2 H2S + 3 02-^2 H2O -h 2 SO2
2H2S-I- O2— ^2H20 + 2S

On being heated moderately, hydrogen sulphide breaks up
into hydrogen and sulphur.

The fact that water solutions of hydrogen sulphide
are poor conductors of electricity, shows it to be a weak
acid. The solution reddens blue litmus very slowly, and
the color is never as deep as that produced by solutions of
strong acids.

Hydrogen sulphide acts on most metals, forming sul-
phides and liberating hydrogen :

Pormation Cu + H2S — >- CuS + H2

of snlphides. 2 Ag + H2S — >- AggS + H2

If hydrogen sulphide is passed into a solution of copper
sulphate, copper sulphide is formed :

CUSO4 + HaS-^CuS + H2SO4

A similar reaction occurs with the dissolved salts of many
metals, as lead, silver, and tin. In such cases the deter-

8ULPHUR SPRING a.

151

mining factor that makes the action possible is the
insolubility of the metallic sulphide either in water or in
the dilute acid that is formed as a result of the action.
We have seen that metallic sulphides can generally be
formed in three different ways, for example :

2Ag+S— ^AggS
2Ag + H2S-^Ag2S + H2
2 AgNOs -f H2S -^ AgaS + 2 HNO3

The possibility of precipitating metallic sulphides by this
third method is of great value to the analytical chemist in
determining the composition of unknown substances.
The identification of the metal in a particular case is made
by means of the color of the sulphide, its solubility in
various reagents, or other characteristic reactions.

TABLE OF SULPHIDES.

Name.

Formula.

Color.

Dissolves in

Mercuric sulphide

HgS

Black

Aqua regia

Copper sulphide

CuS

Black

Concentrated nitric acid

Cadmium sulphide

CdS

Yellow

Concentrated nitric acid

Arsenic sulphide

A82S8

Yellow

Yellow ammonium sulphide

Antimony sulphide

SbaSs

Orange

Yellow ammonium sulphide

Zinc sulphide

ZnS

White

Dilute acids

138. Sulphur Springy. — Hydrogen sulphide is produced
in nature in several ways. We have mentioned its
production during the decay of certain organic mat-
ter. Calcium sulphide is formed when decaying matter
reduces calcium sulphate. Water reacts with calcium
sulphide, forming calcium hydroxide and hydrogen sul-
phide; the calcium hydroxide is converted into the car-
bonate by the action of water containing carbon dioxide-

152

8ULPHUB AND HYDROGEN SULPHIDE.

CaS04-40 — >-CaS
CaS + 2 H2O — ^ H2S 4- Ca(0H)2
Ca(0H)2 + H2CO3— ^CaCOs + 2H2O

The presence of hydrogen sulphide in the waters of some
sulphur springs is accounted for by these reactions.

Summary.
Svlphur is found native.

Commercial forms are roll sulphur and flowers of sulphur.
Rhombic sulphur, prismatic sulphur, and plastic or amorphous
sulphur are allotropic forms.

Rhombic Sulphur.

Pbismatio Sulphur.

Specific gravity

2.07

1.96

Solubility

Soluble in carbon
disulphide

Insoluble in carbon
disulphide

Stability

Stable below 96°

Stable between 96°
and 114°

Sulphur melts at 114° and boils at 450® ; it is a non-conductor
of electricity.

Sulphur is a moderately active element, bums in air, and unites
directly with many elements.

Sulphur is used in the vulcanizing of rubber, in the manufacture
of matches, and in the preparation of sulphur dioxide, sulphites,
and carbon disulphide.

Hydrogen sulphide is prepared by the action of hydrochloric
acid or of sulphuric acid with ferrous sulphide.

It occurs in volcanic gases and in water of sulphur springs, and
is formed in nature by decay of organic matter containing sulphur.

Hydrogen sulphide is a colorless gas, poisonous, and possessing
a characteristic disgusting odor.

EXERCISES. 153

One volume of water at 20° dissolves 3.1 volumes of hydrogen
sulphide.

Its water solution is a weak acid which forms sulphides with
most metallic ions.

It bums in a limited supply of air to fonn water and sulphur,
and in an excess of air forms water and sulphur dioxide.

Hydrogen sulphide is used by chemists to precipitate certain
metals, as sulphides, from solutions.

Bxerclaes.

1. What conditions determine whether sulphur is to assume
the prismatic or rhombic form ?

2. Why is sulphur used in the manufacture of matches ?

3. Write the formulas of five sulphides and the formulas of
the corresponding oxides.

4. Starting with iron, sulphur, and hydrochloric acid, give
directions for the preparation of hydrogen sulphide.

5. Calculate the weight of a liter of hydrogen sulphide
measured under standard conditions.

6. How many grams of ferrous sulphide, would be required
for the preparation of five liters of hydrogen sulphide ?

7. Under what conditions is sulphur deposited when hydro-
gen sulphide burns ?

8. How many liters of oxygen would be required for the
complete combustion of four liters of hydrogen sulphide ?

9. How many grams of hydrogen sulphide would be required
to combine with the copper in 0.1 gram of copper sulphate ?

CHAPTER XVIII.

OXIDES Aim AOIDS OF SULFHTJE.

SULPHUR DIOXIDE.

i 139. Preparation. — When sulphur burns in oxygen or
in air, sulphur dioxide is formed:

The odor of burning sulphur is due to the dioxide formed,
mrect Many ores are sulphides of metals, and large

combinatioiL. quantities of sulphur dioxide are prepared by
roasting such ores in air.

Fig. 47. Preparation of sulphur dioxide.
a, generator; b, safety bottle-

When copper, silver, or mercury is treated with hot,
concentrated sulphuric acid, sulphur dioxide is formed
(Fig. 47). The chemical action is complicated, but it is

154

PREPARATION OF SULPHUR DIOXIDE. 156

probable that the acid first acts with the metal, and hy-
drogen is displaced, according to the equation:
Cu + H3SO4 — >- CUSO4+ 2 H

The hydrogen is then oxidized to water by the deduction
excess of hot, concentrated sulphuric acid. This of luiphu-
oxidation is accomplished at the expense of the ^^ •^*-
reduction of the sulphuric acid to sulphurous acid, accord-
ing to the equation:

H2SO4 + 2 H — ^ HaSOg + H3O
The sulphurous acid decomposes, as soon as it is formed,
into water and sulphur dioxide, according to the equation:

HaSOg-^HaO + SOa
The changes indicated in the three equations given above
go on together. The total change may be represented by
the one equation:

Cu + 2 H2SO4 — ^ CuSO^ + 2 H2O + SO3

When a mixture of sodium sulphite and dilute sulphuric
acid is gently heated in a flask, sulphur dioxide is evolved,
according to the equation: Deoompoii-

Na^SO, + H,SO, -^ Na^SO, + H^O + SO, ^^^ /

Possibly sulphurous acid is first formed, but if so, this
unstable compound immediately decomposes into water
and sulphur dioxide. The equations are:

NajSOg + H2SO4 -^ Na2S04 + H2SO3
HjSOg— ^HgO + SOa

The evolution of the gas may be made very regular by
using sodium acid sulphite and allowing sulphuric acid
(1 to 1) to fall on it, drop by drop (P'ig. 48).

140. Physical Properties. — Pure sulphur dioxide is a
colorless gas, with a suffocating odor, and is more than

156

OXIDES AND ACIDS OF SULPHUR.

twice as heavy as air. It is very soluble in water; one
volume of water dissolves many times its own volume of

the gas at ordinary tempera-
tures. The gas may be com-
pletely expelled from the solu-
tion by boiling.

Sulphur dioxide is one of the
easiest gases to liquefy. At
ordinary temperatures a pres-
sure of but two atmospheres is
required ; at the temperature of
a freezing mixture of ice and
salt, the ordinary atmospheric
pressure is sufficient to change
the gas to a liquid. Liquid
sulphur dioxide is sold in glass
or metal siphons (Fig. 49).

Fig. 48.

141. Chemical Properties. — The solution of sulphur diox-
ide has an acid reaction and neutralizes bases, forming
sulphites. Thus with sodium hydroxide the reaction re-
sults in the formation of sodium sulphite and water. This
indicates the presence in the solution of hydrogen and
sulphite ions, or, in other words, the solution contains
sulphurous acid. This acid, however, has never been
separated from the solution. The acid is formed accord-
ing to the equation:

H2O+SO2— ►-HaSOg
Since sulphur dioxide unites with water to form an acid,
Acid it is called an aeid anhydride. An anhydride is

anhydride, named from the acid it forms with water. Hence
sulphur dioxide is called sulphurous anhydride.

If a solution of sulphurous acid is allowed to stand

USES.

157

in contact with air, it gradually takes up oxygen, forming
sulphuric acid, according to the equation:

2H2SO8+O2— ^2H2SO^

Sulphites are likewise changed to sulphates.

Since sulphurous acid is so readily oxidized, it acts as
a reducing agent in many cases. Potassium Bedudng
permanganate is reduced to potassium sulphate *c*ioa-
and manganese sulphate :

2 KMnO^ + 5 HgSOg— ►- K3SO4 + 2 MnSO^
+ 2H2SO4 + 3H2O

Chromic acid is reduced to chromium sulphate. This
reducing action, in the presence
of water, probably explains the
bleaching action of sulphur diox-
ide in some cases. The coloring
matter is reduced, thus decolor-
izing the substance. In other
cases the sulphur dioxide unites
directly with the coloring mat-
ter, producing a colorless com-
pound. The color of materials
bleached by sulphur dioxide will
often return onq^posure to the
air. Bleaching by chlorine is
more permanent. Fig. 49.

142. Uses. — Great quantities of sulphur dioxide or of
sulphites are used in the bleaching of many organic coloring
matters. Straw, silk, wool, and other materials, which
chlorine bleaching makes hard and brittle, are decolorized
by sulphur dioxide (Fig. 50). Cherries are sometimes
bleached with sulphurous acid and then colored with the

158

OXIDES AND ACIDS OF SULPHUR.

bright shades that are seen in the canned goods of

commerce.

Sulphur dioxide is also used as a disinfeetant. For this

purpose sulphur is burned, or liquid sulphur dioxide is

allowed to evaporate in the
room to be disinfected. In
either case the room should
be tightly closed and the air
in the room be kept moist,
so that the disease germs
may be killed by the sul-
phurous acid formed. This
power of sulphurous acid to
kill lower organisms is the
reason for its use in the cur-
ing of wine and beer. The
^^^' ^* growth of mould is thus

AcamatloninaIr(.)andinsulphurdioxide(«. prevented. An increasing

but questionable use of the acid is its employment as
a food preservative. Sulphur dioxide not only attacks
lower organisms, but is decidedly injurious to higher
forms of life. Vegetable growth is blighted in the neigh-
borhood of smelters and chemical works where the gas is
generated. In cities the gases from the burning of coal
containing sulphur have a like effect.

The most important use of sulphur dioxide, however, is
in the manufacture of sulphuric acid. Immense quantities
are used for this purpose.

143. Snlphor Trioxide. — Sulphur trioxide is prepared by
the oxidation of sulphur dioxide. It is formed in small
quantities when sulphur burns, and its presence causes the
cloudiness often observed in sulphur dioxide. This oxida-

SULPHUR TBIOXIDE.

159

tion may be made more rapid by the presence of catalytic
agents. Finely divided platinum and ferric oxide are the
most important ones. Platinized asbestos is prepared by
soaking asbestos in platinum chloride and heating until
finely divided platinum is left in the asbestos. A mixture
of sulphur dioxide and air is passed over platinized asbes-
tos (Fig. 51). The temperature must be carefully regu-
lated, as the trioxide decomposes at a temperature only a
little higher than that necessary for its formation. The
reaction is then reversible and may be represented ;

2802 + 02:^2808

Aspirator

Fig. 51. Laboratory preparation of sulphur trioxide.

Sulphur trioxide (sulphuric anhydride) dissolves in
water, forming sulphuric acid :

808+H2 0-^H2 804
The laboratory preparation of sulphur trioxide and sul-

160 OXIDES AND ACIDS OF SULPHUR.

phuric acid by the contact method is represented in Fig 51,
Sulphur (a) burns in the air drawn into the apparatus by
an aspirator, and the sulphur dioxide formed passes with
the excess of air into the U-tube (J) which is filled with
some loose material to rid the gases of sulphur dust. In
the bottle (c) the gases lose their moisture as they make
their way through the glass beads drenched with concen-
trated sulphuric acid. The platinized asbestos in the
tube (d) is gently heated, and brings about the union of
the sulphur dioxide with the oxygen of the air. The sul-
phur trioxide resulting from the action is absorbed by con-
centrated sulphuric acid trickling down through the
apparatus («). Water may be used for this last absorption,
but this makes the process longer.

SULPHURIC ACID.

144. Preparation. — Commercially sulphuric acid is pre-
pared by the method just stated (cf. 143). Sulphur,
or ores containing sulphur, are burned in air. The sul-
phur dioxide and the other gaseous products of the com-
bustion are cooled, freed from dust, arsenic, and most of
Contact the moisture, then mixed with air and passed
procMi. through a tube containing a catalytic agent,
usually finely divided platinum (Fig. 52). The oxygen
of the air combines with the sulphur dioxide and forms
the trioxide, which is then combined with water to form
sulphuric acid :

S-fOa— ►-SOg
2 802 + O2— ►-2S08
SOg-f H3O— ^HaSO^

In practice sulphur trioxide is passed into concentrated
sulphuric acid, as it does not dissolve readily in water.

SULPHURIC ACID.

161

The solution is then diluted :

f H3S04 4-S03-^H3S04 . SOg
HaSO^. SOj + HjO— ^2HaS0^

146. Physical Properties. — Sulphuric acid is a heavy,
oily liquid. Ordinary commercial sulphuric acid, called
oil of vitriol, is nearly twice as heavy as water. It boils
at a higher temperature (338°) than most of the common
acids, and many of its uses depend on this fact.

146. Chemical Properties. — Sulphuric acid mixes with
water in all proportions; during the mixing considerable
heat is evolved. If such a mixture is made, the acid
should be slowly poured and stirred into the water. If the
water is poured into the heavier acid, they cannot mix so
readily, and a sudden generation of steam may cause spat-
tering or break the vessel.

.=j ^

r^:

Fig. 52. Contact process (diagrammatic).

A, blower; B, pyrites burner; C. dry scrubber filled with coke ; D, wet scrubber filled
with coke wet with sulphuric acid; E, arsenic purifier: /=, heater; G, contact cham-
ber ; H, absorber (concentrated sulphuric acid).

162 OXIDES AND ACIDS OF SULPHUR.

Concentrated sulphuric acid absorbs moisture from the
air, and this tendency of the acid to take up water explains
Dehydrating many of its actions. Wood, paper, sugar, and
action. similar substances, containing hydrogen and

oxygen, are charred by sulphuric acid. The acid removes
the hydrogen and oxygen to form water, leaving a residue
consisting largely of carbon. On the flesh it acts simi-
larly, and a painful wound results.

With metals the acid acts in two ways. If the action
takes place at a low temperature, hydrogen is evolved,
Action with provided sufficient water is present to dissolve
metals. the metallic sulphate formed:

Zn + H2SO4 -^ ZnSO^ + H3
Fe + H2SO4 — >■ FeSO^ + H^

Mercury, silver, and copper are not acted on by the cold
acid, but if concentrated acid is used and the temperature
raised sufficiently, they react, reducing part of the sul-
phuric acid, forming sulphur dioxide, water, and metallic
sulphates:

Cu + H2SO4 — ^ CuSO^ + 2 H
HjSO^ + 2 H ^ 2 H2O -t- SO2

Thus at ordinary temperatures sulphuric acid acts like
hydrochloric acid, exchanging its hydrogen for metals,
but when hot and concentrated, it acts also as an oxidizing
agent.

With bases and metallic oxides it reacts, forming water
and sulphates:

Action with 2KOH-|-H2S04^K2S04H-2H20
iwwes. Ca(0H)2 + H2SO4 — ^ CaSO^-h 2 H2O

ZnO + HgSO^ — ►- ZnSO^ + HgO
FcgOg + 3 H2SO4 ^ Fe2(S04)3 + 3 H2O

U8E8 OF SULPHURIC ACID. 168

147. Test for a Sulphate. — The sulphates are all soluble
except four, the sulphates of barium, strontium, calcium,
and lead. In detecting the SO4 ion, a solution of barium
chloride is usually employed. Representing by M "^"•^ any
ion carrying two positive charges :

M + + SO4— +Ba++Cl2--— ^M + ^Clg-'+BaSO^
The barium sulphate is easily identified, because it is
white and insoluble in water, dilute acids, and alkalies.
The method is to test, with dilute hydrochloric acid, the
solubility of the precipitate formed with the barium
chloride.

148. Uses. — The absorption of water by sulphuric acid
renders it a good dehydrating agent, and in the laboratory
gases are dried by being made to bubble through it
(Fig. 19, J). In the manufacture of sulphuric acid the
air and sulphur dioxide employed are dried by contact
with sulphuric acid. In the purification of petroleum
products, kerosene, etc., it is used to remove, by charring,
materials which would give offensive odors in burning.
In the preparation of nitroglycerine it removes water
from the nitric acid and glycerine.

As sulphuric acid has a higher boiling-point than most
acids, it is used in their preparation. Examples of this
action have already been studied (§§ 58, 135).

On account of the conductivity of its solutions, sul-
phuric acid is used in storage batteries and in plating. It
is used also as a catalytic agent in the production of
glucose from starch and water.

As a solvent it is used in cleaning metals, previous to
tinning or galvanizing, and certain phosphate rocks are
decomposed and made soluble and suitable for fertilizers
by it. Enormous quantities of it are used in these opera-

164 OXIDES AND ACIDS OF SULPHUR.

tions, and in hundreds of others. There are few materials
in common itse by civilized man with which sulphuric acid has
not been directly or indirectly connected.

Summary.

Sulphur dioxide can be prepared in several ways:
Direct combination of oxygen with free sulphur or with sulphur
in sulphides; reduction of sulphuric acid; decomposition of sulphites
with adds.

The characteristic odor, the weight, and the solubility in water
are three striking physical properties of sulphur dioxide. Chem-
ically it is an acid anhydride, forming sulphurous acid, which is a
powerful reducing agent.

Sulphur dioxide is used in bleaching, as a disinfectant, as a
food preservative, and, most important of all, in the manufacture
of sulphuric acid.

Sulphur trioodde is prepared by the oxidation of sulphur dioxide
by means of a catalytic agent.

Sulphur trioxide is the anhydride of sulphuric add. It combines
energetically with water.

Sulphuric acid is now manufactured by the " contact process,"
consisting of the following steps:

(1) oxidation of sulphur to the dioxide;

(2) catalytic oxidation of the sulphur dioxide to the trioxide;

(3) dissolving the trioxide in concentrated sulphuric add;

(4) dilution of the last solution.

Sulphuric add is a heavy, oily liquid of high boiling-point.

With metals sulphuric acid acts in two ways. At low tempera-
tures and when dilute, hydrogen is evolved and the sulphate of the
metal formed. When hot and concentrated, it acts on certain
metals as an oxidizing agent, forming sulphur dioxide, water, and
metallic sulphates. Sulphuric acid acts on bases and metallic
oxides as a typical add, forming water and a sulphate.

All sulphates are soluble in watery except those of lead, barium.

EXERCISES. 165

strontium, and calcium. To test for a sulphate, add a solution of ^
barium chloride; a white, granular precipitate, insoluble in dilute
hydrochloric acid, indicates the presence of sulphate ions.

Sulphuric acid is used as a dehydrating agent, in the prepara-
tion of other adds, and in a wide range of industrial applications.

Exercises.

1. Which of the laboratory methods would you use for pre-
paring pure sulphur dioxide ? Why ?

2. If a bottle partly filled with concentrated sulphuric acid
is left open to the air, the liquid contents increase. Explain.

3. What advantages has sulphur dioxide over chlorine as a
bleaching agent? What disadvantage?

4. Explain why concentrated sulphuric acid must be poured
slowly into water when the two liquids are mixed.

5. Account for the darkened rings formed on wood where
bottles of concentrated sulphuric acid have been standing.

6. Why can either hydrochloric or sulphuric acid be used in
the preparation of hydrosulphuric acid? Explain which of
these two acids must be taken for the preparation of nitric
acid.

7. Explain why boiling concentrated sulphuric acid produces
such frightful burns.

8. Show how concentrated sulphuric acid acts as an oxidizing
agent on metallic silver.

9. What effect would you expect if a strip of lead were placed
in dilute sulphuric acid ? Explain.

10. Why is a dish containing sulphuric acid put inside the
case of a delicate balance ? Why are clocks for keeping exact
time similarly treated?

11. Why is the civilization of a country said to be indicated
by the amount of sulphuric #cid it uses ?

166 OXIDES AND ACIDS OF SULPHUH.

12. Compare the chemical actions in chlorine and sulphur
dioxide bleaching.

13. What is an acid anhydride ? Name two anhydrides con-
taining sulphur, and give their formulas.

14. How many pounds of sulphuric acid could be manufac-
tured from 120 pounds of pure sulphur ?

15. How many liters (standard conditions) of sulphur diox-
ide would result from the reaction of 12 grams of copper with
concentrated sulphuric acid?

16. What weight of sodium sulphite must be decomposed to
furnish 3.6 liters sulphur dioxide (standard conditions) ?

17. Calculate how many grams (a) of silver sulphate and
(&) of copper sulphate you could make from a dime which is
10 % copper. A dime weighs 2.48 grams.

CHAPTER XIX.
iriTBOOEN AND THE ATHOSFHEBE.

149. Occurrence. — Nitrogen has already been mentioned
as constituting a large portion of the atmosphere. It
is also found in a few mineral compounds, many of which,
however, are the result of the activity of animal and vege-
table organisms. Nitrogenous organic compounds exist
in great variety;* and one class, the proteids, of which the
white of egg is an example, are directly concerned with
the life processes. In fact, nitrogen is perhaps the most
characteristic element in living organisms, since the pro-
teids make up the living matter of the muscles and the
protoplasm of the cells. Life without nitrogen would be
impossible. In the decay of organic matter much of its
nitrogen is returned to the atmosphere in the form of the
familiar gas ammonia.

150. Preparation. — Nitrogen may be prepared from air
by causing oxygen to combine with phosphorus in the
presence of water. Phosphorus is employed because its
great tendency to combine with oxygen insures the com-
pleteness of the reaction, even at ordinary temperatures,
and because its oxides have a great tendency to combine
with water and so are rapidly removed from the vessel.
In Fig. 53, a is a small crucible, floating on water, and
containing phosphorus.

Other reducing agents may be used, provided the oxide
formed is easily separated from the nitrogen. If air is

167

<i;

168 NITROGEN AND THE ATMOSPHERE.

passed through a strongly heated tube containing reduced
copper or fine-meshed copper gauze, nearly pure nitrogen
results (Figs. 54 and 55). The reason
for the use of copper is that its oxide is
a non- volatile solid. Nitrogen pre-
pared from air always contains argon
and other impurities.

The oxidation of ammonia is a con-
~ IT ^ venient method for preparing pure
nitrogen. Ammonia gas is passed over
strongly heated copper oxide. The hydrogen is oxidized
to water, and the nitrogen remains. Heat alone will
liberate nitrogen from its compounds. * Ammonium ni-
trite, gently heated, decomposes into water and nitrogen.
Owing to the unstable nature of ammonium nitrite, a mix-
ture of ammonium chloride and sodium nitrite is used.
Ammonium nitrite is probably formed and then decom-
poses, as just stated.

NH4CI + KNO2 — ^ KCl -I- NH4NO2
NH4NO2— ^N2 + 2H20

151. Physical Properties. — Nitrogen is slightly lighter
than air, as we should expect from the fact that oxygen,
the other chief constituent, is heavier. It is without color,
odor, or taste. Nitrogen is less soluble in water than
oxygen, so that the bubbles of gas given off, when ordinary
water is warmed, contain a smaller proportion of nitrogen
than air. Cooled to a very low temperature under atmos-
pheric pressure, nitrogen becomes a colorless liquid; on
further cooling the liquid freezes to a white solid.

152. Chemical Properties. — The large amount of nitrogen
in the air is due to its inertiiBss : it does not combine

CHEMICAL PROPERTIES.

169

readily with many substances, and its compounds are easily
decomposed. It unites directly with few elements and
with these only at high temperatures ; sometimes the
electric spark is necessary to cause combination. The
ease and violence with which its compounds decompose is
well illustrated by nitroglycerine and guncotton.

Nitrogen may be caused to combine slowly with oxygen
by passing electric sparks through the mixture and remov-
ing the oxides by dissolving them in water as Aotion with
fast as they are formed. If they were not so oxygen,
removed, they would be decomposed by the heat of the

Air

, ^ m?i

Copper Gauze

P@^-"J^

Fig. 54. Preparation of nitrogen.

succeeding sparks. Nitrogen will not burn in oxygen
without a continual supply of external energy, as the tem-
perature of the combustion is lower than the kindling-point
of nitrogen. Such a combination of the oxygen and nitro-
gen of the air takes place in a flaming electric arc between
metal poles.

Ammonia, (NHg), can be formed by the passage of
sparks through a mixture Qf hydrogen and nitrogen. In

170 NITROGEN AND THE ATMOSPHERE.

this case, as in the similar production of the oxide, the
ammonia must be removed as formed, since the reaction is
reversible and a point of equilibrium is reached, at which
it proceeds as rapidly in one direction as in the other :

N2 + 3H2::;:^2NH8

A few nitrideB are known, of which the principal ones
are those of lithium, calcium, magnesium, and boron.
These require a red heat for their formation. Magnesium
nitride, the most common one, is a yellow powder formed
by igniting magnesium and nitrogen.

While nitrogen does not react readily, many reactions
are affected by its presence. Thus burning cannot be so
vigorous in the air as in oxygen, since the large propor-
tion of nitrogen dilutes the oxygen, preventing a rapid
contact with the combustible material. Some heat is also
employed in raising the temperature of the nitrogen; the
temperature of combustion is lower than would be the
case were nitrogen absent.

153. Composition of the Air. — The average proportions of
the chief constituents of the air are as follows :

GOMPOBinOH.

By Yolume.

By velfht

Nitrogen .

. 78.06

76.6

Oxygen

. 21.00

23.2

Argon

. 0.94

1.3

Carbon dioxide

. 0.04

0.06

Traces of other substances are often present, but under
the term air we usually include only the nitrogen, oxygen,
and argon. The relative amounts of these are practically
constant, except in certain localities, as in cities, and in
poorly ventilated places.

PROOFS THAT AIR 18 A MIXTURE. 171

154. Proofs that Air is a Mixture. — Air must be regarded
as a mixture rather than a compound for several reasons :

(1) The composition is not- absolutely uniform. This
is not in agreement with the law of definite proportions.
While the differences in composition are slight, they are
greater than those found in different samples of a pure
chemical compound.

(2) If the various components of the air are mixed in
the proportions in which they are found in the atmosphere,
there is no evidence of reaction. We have found that
when a chemical change takes place, there is usually a
change in the temperature caused by the absorption or
liberation of heat. Other energy changes, such as the pro-
duction of light and sound (explosions), often accompany
reactions. None of these energy changes occur in this
case, hence there is no probability of a reaction.

(3) If air is allowed to pass through an unglazed por-
celain tube, it is found that the lighter nitrogen diffuses
through the porcelain walls more rapidly than the oxygen;
were they combined in molecules of a compound, they
would go through with equal velocity.

(4) If air is cooled under pressure, it is found that the
argon and oxygen liquefy bef ore*the nitrogen, and if the
liquid air is allowed to evaporate, the nitrogen vaporizes
more rapidly than the oxygen. If air were a compound,
it would have a definite boiling-point, at which it would
vaporize unchanged.

(5) When air is brought in contact with water, nitrogen
and oxygen dissolve in the proportion of 63:34; while
iii atmospheric air the proportion is about 4: 1 by volume.
• The molecular motion of the gases and the winds suf-
fice to keep the composition of the atmosphere practically
constant. Local conditions may slightly affect the com-

172

NITROGEN AND THE ATMOSPHERE.

position, especially in ill-ventilated places, but the total
quantity of the air is so great — 15 pounds resting on each
square inch of the earth — that even a large city produces
scarcely any noticeable effect on the composition.

The constituents of air may be successively removed,
so as to leave the nitrogen, by the apparatus represented
in Fig. 65. The bottle (a) serves as an aspirator to draw
air through the apparatus and also to collect the residual

£. ^^

' miyyyyyyymnf '

I-^IM^

d d e
Fig. 55. Separation of the components of air.

nitrogen. The oxygen is removed by combining it with
copper (gauze) in the hard glass tube (c), which is heated
by the combustion furnace (6). Before reaching the com-
bustion tube, however, the air has to pass through the
bottles (ee) containing a concentrated solution of potas-
sium hydroxide to take out the carbon dioxide, and
through the bottles ((iS) containing concentrated sul-
phuric acid to remove the water vapor or moisture. The
nitrogen collected in (a) is purer than that obtained by
method shown in Fig. 64.

155. Water Vapor of the Air. — Some water vapor, de-
rived from evaporation, is always present in the air ; the
amount usually increases with the temperature; thus warm
breezes blowing over bodies of water are moist. When
cooled, the vapor may condense as fog or rain. The air in
desert regions though warm is dry, because the air before

CARBON DIOXIDE AND NITROGEN CYCLES. 178

being warmed has passed over a cool, mountainous region
and has deposited its moisture.

The amount of water in the air, relative to the amount
necessary to saturate the air under g^yen conditions, is
known as the relative humidity. This is high when the
air is nearly saturated, and low when the air is very dry.
If the air is warm and damp, it is sultry and uncomfort-
able; when the air is half saturated, it is comfortable;
when the air is cool and damp, it is chilly and disagreeable.

156. Carbon Dioxide and ITitrogen Cycles. — Carbon diox-
ide is always present in the air, though in a very small
proportion. In normal outdoor air about 4 parts in 10,000,
or four-hundredths of 1 per cent, are present. The
proportion may rise as high as 1 per cent in a crowded
room from the exhalations of the people present. Carbon
dioxide is continually given off to the air in the exhala-
tions of animals and in combustion, but as it' is taken up
from the air by plants, the amount in the air remains
practically constant. The very small percentage of car-
bon 'dioxide in the air furnishes all the carbon needed for
the growth of plants. ^

The oxygen in the air is removed by animals and re-
placed by plant life ; thus the plant and animal life pre-
serve the balance, maintaining the atmosphere at a
constant composition.

The nitrogen removed from the air, to form soluble
compounds in the soil, is taken up by plants and con-
verted into proteids. These proteids are the source of
the protoplasm of animals. These unstable proteids break
up both during the life of the plants and animals and
after their death, and the nitrogen finally makes its way
back to the air.

174 NITROGEN AND THE ATMOSPHERE.

157. Other Constituents of the Air. — Other materials are
found in small amounts, argon and helium (inert gases
about 1 %), traces of ammonia, sulphur compounds, and
fine dust particles, which depend on local conditions and
which often produce climatic effects. These dust par-
ticles include a great variety of materials — steel, stone,
soil, and coal dust. The organic particles include pollen
grains and spores of plants, germs and bacteria of dis-
ease, which are always present, shreds of various fabrics,
as cotton and woollen cloth, and dried bits of refuse of
all sorts.

THE INERT GASES.

158. Discovery of Argon. — The discovery and investi-
gation of the inert gases in the air have afforded one of
the most brilliant and interesting chapters in the history
of chemistry. In 1892 Rayleigh, an English scientist,
noticed that nitrogen from the air was a trifle heavier
than that obtained from nitrogen compounds. This
meant that the supposedly pure nitrogen from the air
contained some gas, heavier than nitrogen, which had
remained undetected dgspite the careful study of the
atmosphere for more than a century.

A small amount of the hitherto unknown gas was ob-
tained by Ramsay, an English chemist, who passed nitro-
gen from the air over heated magnesium which combined
with the nitrogen, forming magnesium nitride, a yellowish
solid. This method yielded but a trace of the new gas,
and a better way was soon devised by Lord Rayleigh.
Even this, however, was slow and required many pre-
cautions to secure a very small sample of the new ma-
terial.

The new substance was found to constitute about 1 % of

ISOLATION OF THE OTHER INERT QA8E8. 175

the air. It was one-fourth heavier than oxygen and over
one-third heavier than nitrogen. All attempts to make
the gas enter into chemical combination failed, and hence
it was given the name argon^ signifying inactive.

159. Isolation of the Other Inert Oases. — Certain irregu-
larities in the properties of argon led Rayleigh and Ram-
say to suspect that this new gas was not itself pure. By
means of liquid air the argon obtained from the atmos-
phere was liquefied, and, at the low temperatures ob-
tained, repeated processes of fractional evaporation and
liquefaction were carried on. The argon was found to
contain minute amounts of other inert gases. Two of
these could be separated only by using the extremely low
temperature possible with liquid hydrogen. Three of the
new inert gases were given names which bring to mind
the long, baffling search for them. NeoTL means new ;
xenon ^ stranger ; and krypton^ hidden. Besides these
three gases a trace of helium was found. This element
was formerly supposed to exist only in the sun.

160. Properties. — Neon, xenon, and krypton closely re-
semble argon, but each was found to have its peculiar
spectrum and all except neon a very low but definite boil-
ing-point. Thus they were elements and formed a very
closely related group with argon. Certain considerations
have led us to believe that all these elements contain but
one atom to the molecule. Their inertness with respect
to chemical combination explains why no compounds con-
taining them are known, and why they were overlooked
until recently.

161. Helimn. — In 1869 Lockyer noticed some lines in
the sun's spectrum which did not correspond with those

176 NITROGEN AND THE ATMOSPHERE.

of any other element known on earth. This element was
concluded to exist in considerable quantities in the sun.
In 1896 Ramsay, in searching for sources of argon, ex-
amined the gases given off by certain rare minerals, as
clevite, when heated. In some cases a gas was obtained
which gave a spectrum identical with that of the supposed
element in the sun, and hence was given the name helium.
The new element has since been obtained from the waters
of certain mineral springs and exists in minute quantities
in the atmosphere.

Helium is a very light gas, being only twice as heavy
as hydrogen. Its properties resemble those of argon, and
it is therefore classed with the other inert gases.

Recent researches have proved that helium results from
the decomposition of radium, which was thought to be an
element.

Summary.

Nitrogen constitutes the larger part of the air. It is a constitu-
ent of protoplasm and of proteids, hence is essential to vital pro-

It is prepared by

(1) the oxidation of ammonium compounds ;

(2) removing the oxygen from the air by phosphorus; this is

sufficiently pure for ordinary use.

One liter of nitrogen weighs 1.26 grams. Its atomic weight is 14.
The nitrogen molecule contains 2 atoms (N2).

Nitrogen is generally inert ; under electric stress it reacts tardily
with oxygen and with hydrogen. A few bacteria are capable of
assimilating it.

Air is essentially nitrogen, oxygen, and argon, with var3ring
amounts of water vapor, carbon dioxide, and compounds of nitro-
gen and sulphur. Nearly all the constituents of the air are pro-
duced by and used in vital processes.

SUMMARY. 177

CARBON OXYGEN CYCLE.
oxygen -<

Animal Plant

Stbucturb ^^^ Strdotubb

• carbon dioxide-

NITROGEN CYCLE.

Ais

oxygen nitrogen

Water

nitrates

ammonia

Animal Stsucture < Plant Structure

The amount of the air is so enormous that local conditions have
little or no appreciable effect on its composition. The important
factor determining the composition is the balance maintained be-
tween plant and animal life.

The variation in composition indicates that air is merely a
mixture. Chemical reactions are always accompanied by thermal
changes.

Ezerclaes.

1. Why are so few mineral compounds of nitrogen found in
nature ?

2. How was it shown that the material in the air, formerly
known as nitrogen, was not a pure substance?

178 NITBOGEN AND THE ATMOSPHERE.

3. Under what influences does nitrogen react?

4. Wliat chemical reactions take place in the air during a
thunderstorm ?

5. From what sources are the principal constituents of the
air continually derived ? By what means are they removed ?
Why is the composition of the air so nearly constant all over
the earth?

6. What is the weight of air over a city lot 26 x 100 ft. ?
How much of it is oxygen ?

CHAPTER XX.

HITBOOEH OOMFOinrDS.

AMMONIA.

162. Hatural Formation. — The most important constitu-
ent of all living organisms is protoplasm, a complex sub-
stance containing nitrogen, carbon, hydrogen, oxygen, and
other elements. When a plant or animal dies and decom-
position sets in, the protoplasm breaks up very quickly,
yielding simpler compounds. The nitrogen unites with
the hydrogen to form the gas ammonia, the molecule of
which contains one atom of nitrogen and three of hydro-
gen. Its formula, therefore, is NHg. Its characteristic
odor can often be noted in the vicinity of heaps of decom-
posing animal or vegetable refuse.

163. Preparation. — Ammonia is obtained commercially
as one of the products of the distillation of coal, in the
manufacture of coal gas (Fig. 66). Soft coal is heated in
iron retorts at an intense heat. Moisture, volatile matter,
and gases are driven off, coke remaining in the retort. The
gases are cooled in pipes, and coal-tar condenses, commercial
Then the gases are passed into a " scrubber " method,
where they come in contact with water, and here the
ammonia dissolves. The gases pass on to be further puri-
fied from sulphur compounds by passing through calcium
and iron oxides in the purifier, previous to distribution
as illuminating gas. The water containing ammonia com-

179

180

NITROGEN COMPOUNDS.

i i - i!! h ' nTi:-

j' l rr II

^j—y

[||.i.iJiK. [oik

PBEPABATION.

181

pounds is * boiled with milk of lime, and the expelled
ammonia is mixed with sulphuric acid ; tarry materials
are separated, and the solution of ammonium sulphate
is evaporated and crystallized.

The dried, crystallized ammonium sulphate is mixed
with slaked lime in an iron retort and heated. Ammonia
gas and water are g^ven off, and calcium sulphate remains:

Ca(0H)2 + (NH^)^ SO4 — ^ CaSO^ + H^O + 2 NH3

The ammonia is dissolved in water, forming ammonia
water, or spirits of hartshorn, or it may be dried by pass-
ing through quicklime (CaO) and compressed in tanks.

In the laboratory, ammonia is usually pre- Laboratory
pared (Fig. 57). by heating ammonium chloride method,
(sal ammoniac) with calcium hy-
droxide (slaked lime) :

Ca(0H)2 + 2 NH4CI — ^ CaClj
-I-2NH8+H2O
In this preparation any ammonium
salt can be substituted for ammonium
chloride, and any non-volatile base
for the calcium hydroxide. A typi-
cal reaction probably proceeds as
follows :

(NH^)^ SO4 -h 2 NaOH — ^ NaaSO^

+ 2NH4OH

NH4OH— ^NHg + HaO

That is, ammonium hydroxide is
first formed and breaks up at once
into ammonia and water. Since ammonia is a gas, a vola-
tile product can be formed as a result of the reaction
between ammonium salts and bases. This is analogous
to the fact that many acids are formed by the action of

Fig. 57.

182

NITROGEN COMPOUNDS.

sulphuric acid upon their salts, because they have lower
boiling-points than sulphuric acid.

Ammonia can also be obtained by warming a strong
ammonium hydroxide solution :

NH^OH— ^NHa + HgO

164. Physical Properties. — Ammonia is a colorless gas
with peculiar odor, lighter than air, and exceedingly sol-
uble in water. At 0° C. one volume of water will hold
in solution over 1000 volumes of the gas ;
at ordinary temperatures about 700 vol-
umes. This solution is known as am-
monia water, or ammonium hydroxide.
On heating or on standing exposed to air,
it gives off ammonia.

The great solubility of this gas is strik-
ingly shown by the "ammonia fountain"
(Fig. 68). A flask is filled with dry am-
monia, and inverted over water. As soon
as the clip (not shown in the figure) is
removed from the rubber tubing, the
water rushes in to dissolve the gas.

The gas is easily liquefied ; at ordinary
temperatures a pressure of 4.6 atmo-
spheres is needed.
Fig. 58.

165. Chemical Properties. — Pure, dry

ammonia is not an active substance; it is not readily

combustible in air, but can be burned in oxygen. When

ammonia is passed over heated copper oxide, water and

^nitrogen are obtained:

2 NHg + 3 CuO — ^ 3 Cu + 3 HgO -h N^

The most important chemical property of ammonia is

AMMONIA. 188

the basic character of its water solution. This solution,
which is often incorrectly called ammonia, turns red litmus
blue, neutralizes acids, and conducts electricity; it behaves
like the solution of a base. When this solution is neutra-
lized with hydrochloric acid, a salt is formed whose com-
position is represented by the formula NH^Cl, and similar
salts are formed with other acids:

NH4OH -h HCl — ^ NH4CI + HgO
2 NH4OH + H2SO4 — ^ (NH4)2S04 -h 2 HjO

We may assume, therefore, the existence of NH^ ions
and 0H~ ions in the solution of ammonia. The group
NH^ is known as the ammonium radical. The solution,
then, contains ammonium hydroxide, NH^OH. AmmonivnL
There are a large number of ammonium salts, radical,
but ammonium has never been obtained in a free
state. •

Ammonium salts react similarly to the compounds of
sodium and potassium, and they may be considered as
substances in which the group of atoms NH^ (ammonium
radical) takes the same part as an atom of hydrogen or
potassium. Thus as potassium chloride dissociates into
K"*" and Cl~ ions, ammonium chloride, NH4CI, dissociates
into NHJ and Cl~ ions.

If an electric current is passed through a solution of
ammonium chloride, we might expect to obtain ammo-
nium and chlorine, since these are the ions formed.
The chlorine, however, liberated at the anode reacts with
the ammonium salt present in the solution, forming hydro-
chloric acid and nitrogen. At the cathode the NHJ ion,
on discharging, decomposes into ammonia and hydrogen,
the ammonia dissolving in the water. The equations for
the electrolysis are:

184

NITROGEN COMPOUNDS.

6 NH4CI — ^ 6 NH4 -h 3 CI2

6 NH4 -h 6 HjO — ^ 6 NH4OH + 3 H2
3 CI2 + 2 NH^Cl — ^ 8 HCl -h N2

Thus by electrolysis we get one part of nitrogen and
three parts of hydrogen.

166. XTaes — The most important uses of ammonia are as
a refrigerating agent and for* the preparation of ammonia
water. When a gas is liquefied, heat is liberated, and when
the liquid retnirns to the gaseous state, heat is absorbed.
In the manufacture of artificial ice (Fig. 59), ammonia is

Fig. 59. Refrigerating plant.

liquefied by being compressed by powerful pumps; then
the liquid ammonia is cooled by passing cold water over
the pipes containing it. The liquid ammonia is distrib-
uted through pipes, where it evaporates rapidly. The gas
is drawn back by the pump, condensed to a liquid, and
used again. The pipes in which the evaporation takes

NITBOUS OXIDE. 186

place are immersed in a strong salt solution, which, by
furnishing heat for evaporation, is cooled to a point
below the freezing-point of water. Cans of water are
placed in the cooled brine, and the water is frozen in
from 24 to 36 hours. Cold-storage rooms may be kept
cool by distributing the cold brine to the apartments to
be cooled, where it is passed through coils near the ceil-
ing.

The value of ammonia water as a cleansing agent is due
to its ability to dissolve grease. Its basic properties also
give it a use in the laboratory, whenever a volatile alkali
is desirable. Large quantities of ammonia are used in the
manufacture of sodium carbonate by the Solvay process.

OXIDES OF NITROGEN.

Nitrogen combines with oxygen in five proportions, cor-
responding to the formulas: NgO, nitrous oxide; NO, nitric
oxide; NgOg, nitrous anhydride; NOg, nitrogen peroxide;
NgOg, nitric anhydride.

167. Hitroug Oxide. — Nitrous oxide (NgO), laughing gas,
is prepared by heating ammonium nitrate (Fig. 60) :

NH^NOg — ^ NgO -h 2 HgO

The nitrate melts and soon begins to decompose with
effervescence. The heat must be carefully regulated or
an explosion may occur. The nitrous oxide is a colorless
gas with a slightly sweet taste; inhaled, it produces uncon-
sciousness; for this purpose it is used in minor surgical
operations. It was the first of modern anesthetics and was
discovered by Sir Humphry Davy.

Nitrous oxide supports combustion almost as well as
oxygen, but, unlike oxygen, it does not react with nitric

186

NITROGEN COMPOUNDS.

oxide or support the combiistion of sulphur which is not
burning vigorously.

168. Vitric Oxide. — Nitric oxide is a colorless gas; it is
generally formed in the action of dilute nitric acid with
metals. In the laboratory copper and nitric acid are used:

3 Cu -h 8 HNOg— ^ 3 Cu(N05)2 -h 2 NO -h 4 HgO

It does not support combustion, being more stable than
nitrous oxide, but readily combines with oxygen, forming
nitrogen peroxide, with a slight rise of temperature:

2NO-h02— ^2N02
This action makes it useful as a catalytic agent in one
process for the manufacture of sulphuric acid.

Fig. 60. Preparation of nitrous oxide.

a, flask containing melted ammonium nitrate ; b, catch bottle for water formed ;
c. collecting bottle.

169. Nitrogen Peroxide. — Nitrogen peroxide, NOg, is a
heavy red-brown gas of disagreeable odor. It is formed
immediately whenever nitric oxide is brought in contact

NITROGEN PEROXIDE.

187

with oxygen or with air (Fig. 61). It dissolves in water,
the solution has an acid reaction and contains nitrous and
nitric acids:

2 NOj -h HgO — ^ HNO2 -h HNO3

Thus the fumes from nitric acid, containing oxides of
nitrogen, form nitric acid with water, and cause the corro-
sion usually observed on metal objects near which nitric
acid is kept.

Nitrogen peroxide is read-
ily liquefied and solidified,
the liquid being yellow and
the solid colorless. When
the liquid vaporizes, the va-
por given off at the boiling-
point is light brown and
grows darker as the temper-
ature rises. Vapor density
determinations indicate that
vapor given off has a compo-
sition represented by the
formula NjO^, part of the
molecules of which immediately dissociate into NOg mole-
cules, so that the light-colored gas is a mixture of the
two oxides. As the temperature rises, more molecules
dissociate, and the dark gas at high temperatures is
chiefly NO^. Thepe changes are represented by the equar
tion:

170. Other Oxides. — Nitrogen trioxide (NjOg) and ni-
trogen pentoxide (NgOg) are unstable substances of no
particular importance. They unite with water, forming
acids:

Fig. 61. Nitric oxide.
a, closed ; b, open to air.

188

NITROGEN COMPOUNDS.

H2O + N2O8— ^2HN02

H2O + NA— ^2HN03

Hence the trioxide is termed nitrous anhydride and the
pentoxide is known as nitric anhydride.

NITRIC ACID.

Nitric acid was known to alchemists, who called it aqiia
fortis (strong water), because of the great chemical activ-
ity it displays. They prepared it by heating a mixture
of potassium nitrate, copper sulphate, and potassium alu-
minum sulphate. The last two of these substances con-
tain water of crystallization, and from this came the
hydrogen which the acid contains.

171. Preparation. — Both commercially and in the labor-
atory, nitric acid is prepared by
heating a mixture of concentrated
sulphuric acid and sodium nitrate
(Fig. 62). The latter substance
is found in considerable quantities
in certain parts of Chile and in the
western United States. Other
nitrates might be used; for in-
stance, potassium nitrate, which is also found in nature,
though in much smaller quantity than sodium nitrate.

The reaction may proceed in two stages. The first re-
action is :

NaNOg + H2SO4 — ^ NaHSO^ + HNO3

If there is an excess of acid, the reaction does not proceed
beyond this point. If, on the other hand, there is an ex-
cess of sodium nitrate, H;he sodium hydrogen sulphate
that is formed in the first action reacts at a higher

Fig. 62.

NITRIC ACID. 189

temperature with more sodium nitrate, according to the
equation :

NaNOg + NaHSO^ — ^ Na^SO^ + HNO3

Writing one equation to show the final results of the two
stages of the reaction, we have :

2 NaNOg + H2SO4 — ^ Na^SO^ -h 2 HNO3

Since the second action requires a higher temperature
than the first, and since nitric acid undergoes considerable
decomposition at the higher temperature, it is customary
to use enough sulphuric acid to give only the first reac-
tion. Sulphuric acid is used in this operation for the
reason that its boiling-point is higher than that of nitric
acid. Very few acids could be substituted for sulphuric
acid because most of them have too low boiling-points. In
the laboratory preparation of nitric acid, the distilled acid
is usually collected in a test-tube or other receiver, kept
cool by water in a battery jar (a, Fig. 62).

Nitric acid is an important article of commerce ; so that
the reaction that has been described is carried out on a
large scale. Iron retorts are used, and the acid is con-
densed and collected in a series of earthenware vessels.

1 72. Physical Properties. — Nitric acid is a colorless liquid
at ordinary temperatures, boiling, if free from water, at
86° C. The diluted acid has a higher boiling point, varying
with the dilution. A mixture that contains 68% of this
acid boils constantly at 120°.

As it is ordinarily prepared, nitric acid contains con-
siderable water and is colored yellow by the presence of
dissolved oxides of nitrogen, which result from the decom-
position of the acid by the heat used in its preparation. It
is usual to distil the acid in an apparatus in which the

190 NITROGEN COMPOUNDS.

pressure is less than that of the atmosphere. In this way
the distillation can be carried on at a lower tempera-
ture and the undesirable decomposition is avoided.

173. Chemical Properties. — The chemical behavior of
nitric acid is v6ry interesting. Generally its action is not
a simple one. This is because it possesses two distinct
chemical characteristics, both of which it displays in a
marked degree.

First, it is a very strong acid. This is because it is
highly dissociated into ions when dissolved in water, even
in concentrated solution. The hydrogen ions, being present
in large numbers, produce all the actions that are charac-
teristic of acids, such as the formation of salts with bases
and the transference of the electric charge of the hydrogen
ion to form metallic ions when the acid is brought in con-
tact with a metal.

Second, nitric acid is a powerful oxidizing agent. This
can be shown in a number of ways : charcoal can be made
to burn in nitric acid ; horsehair will take fire if put into
the gaseous substance ; both the coloring matter and fabric
of cotton or woollen goods are quickly destroyed by it.

When nitric acid does oxidizing work, it is itself re-
duced. There are various reduction products of the acid.
Bednction The product formed depends on a number of
products, conditions, particularly on the temperature and
the degree of dilution of the acid. In any case there are
several reduction products, though usually one is found
in excess of the others. If the acid is moderately dilute
and acts at ordinary temperatures, the reduction product
is commonly nitric oxide. From concentrated nitric acid
a large quantity of nitrogen peroxide is always obtained.
From very dilute acid the reduction product may be

ACTION WITH METALS. 191

nitroas oxide, hydrogen, or even ammonia. Thus we see
that the more dilute the acid, the farther the reduction is
carried. This does not mean that the more dilute acid
is the stronger oxidizing agent; on the contrary, it is
because the concentrated acid is such a powerful oxidiz-
ing agent that the lower reduction products cannot escape
from the acid without being themselves oxidized to a
certain extent.

As we should expect from its being so strong an oxidiz-
ing agent, nitric acid is a rather unstable substance, tend-
ing to give up part of its oxygen to form more stable
substances. It will do this under the influence of light, or
more readily if some oxidizable substance is present.

174. Action with Metals. — Nitric acid acts with many of
the metals, but owing to its dual chemical character, it
does not act on them in the same way that other acids do.
Hydrogen is seldom evolved by the action of nitric acid
on metals. The gases that are given off are the reduction
products of nitric acid.

The action of moderately dilute nitric acid on copper
can be taken as a type of its action on the heavy metals,
as silver, mercury, and lead. Experiment shows that the
products of this action are copper nitrate, nitric oxide, and
water :

8 Cu -h SHNOg — ^ 3 Cu(N08)2 + 2 NO + 4 HgO

This equation represents the results of an action that takes
place in two stages. It is probable that, as a result of the
first stage of the action, there is a tendency to liberate
hydrogen, according to the equation :

Cu + 2 HNOa— ^ Cu(N08)2 + 2 H

192 NITROGEN COMPOUNDS.

But the nascent hydrogen is at once oxidized by nitric
acid:

3H + HNO3— ^2H20 + NO

As a final result of these actions we get the products
shown in the first equation.

The balancing of such an equation as this, involving
oxidation and reduction, is a somewhat difficult matter.
It will probably be found convenient to remember the
numbers 3 and 8 in this reaction.

If concentrated nitric acid acts on copper, nitrogen
peroxide is formed in considerable quantity, as well as
some nitric oxide :

Cu + 4 HNOg— ^ Cu(N08)2 + 2 NOg + 2 H^O

If very dilute nitric acid acts on zinc, or metals like it,
the nitrogen of the acid is reduced to ammonia, which
then combines with more of the acid, forming ammonium
nitrate :

4 Zn + 10 HNOg — ^ 4 Zn(N08)2 + NH4NO3 + 3 HgO

Nitric acid does not act on platinum.

175. XTses. — Nitric acid dissolves silver, but does not act
on gold ; hence it is sometimes used to separate these two
metals. The chief uses of nitric acid depend upon its
ability to form unstable salts with organic bases (com-
pounds containing hydrogen and carbon). Two of these
products are nitroglycerine and guncotton. Celluloid is
a mixture of guncotton and camphor.

Aqua regia is a mixture of nitric and hydrochloric acids.
It dissolves gold and platinum. The fact that the mixture
of the acids does what neither acting alone can do, is ex-
plained by the liberation of nascent chlorine by the inter-

NITRATES. 193

action of the two acids. The hydrogen of the hydrochloric
acid is oxidized by the nitric acid :

3 HCl + HNO3 — ^ 3 CI + 2 HjO + NO

Nitric acid is also extensively used in the manufacture
of many dyes and drugs.

NITRATES.

176. Sodinm and Potassium Nitrates. — The salts formed
by the replacement of the hydrogen of nitric acid are
called nitrates. The nitrates of sodium and potassium are
the only ones found in nature in any considerable quantity.
Potassium nitrate, ordinary saltpeter, is manufactured in
a manner analogous to that by which it is produced in
nature. Excreted animal matter which contains nitrogen
decomposes under the influence of certain minute organ-
isms and bases, and has its nitrogen transformed into
nitrates. As a result of this action potassium nitrate
gathers around stables. The potassium carbonate that is
involved in the action is taken from the soil.

Sodium nitrate is found in large quantities in Chile,
from which fact it gets the name Chile saltpeter. Nitric
acid is made from it. Because of its great abundance
sodium nitrate is cheaper than potassium nitrate. The
following reaction will take place in hot concentrated
solution:

NaNOg + KCl — ^ KNOg + NaCl

Advantage is taken of the fact to prepare the more
expensive potassium nitrate (cf. p. 138).

One of its chief uses is for the manufacture of gun-
powder (page 138). Sodium nitrate cannot be used for
this purpose, because it is slightly deliquescent. Potas-
sium nitrate is used as a preservative in the making of
corned beef.

194 NITROGEN COMPOUNDS.

177. Preparation of Nitrates. — Nitrates, like the salts of
the other common acids, can be made in several simple
ways in the laboratory:

(a) By the action of nitric acid on metals :

3Ag -h 4HN08— ^ SAgNOg + NO + 2H2O

As has been pointed out, hydrogen is seldom a product in
the action of nitric acid on metals.

(6) By the action of nitric acid on oxides or hydroxides
of metals:

ZnO -h 2HN08 — ^ Zn(N08)2 + H2O

Zn(0H)2 -h 2HN08— ^ Zn(N08)2 + H2O

(<?) By the action of nitric acid on salts that give vola-
tile products with this aci5:

ZnCOg + 2HN08 — ^ ZnCNOg)^ + CO2 + H2O

Nitrates cannot be prepared by precipitation, because
nitrates of all metals are soluble in water. This fact also
prevents the use of a precipitation method as a test for a
nitrate.

178. Test for Nitrates. — The test for the NOg ion de-
pends upon the oxidizing power of the NOg group. The
substance to be tested is mixed with a solution of ferrous
sulphate. Concentrated sulphuric acid is then added, so
as to form a layer below the mixed solution. Nitric acid
oxidizes ferrous sulphate, and, at the same time, nitric
oxide, NO, is formed as a reduction product. This com-
bines with some of the unchanged ferrous sulphate, pro-
ducing a characteristic unstable compound (whose formula
is probably 2 FeSO^ .NO), which appears as a dark brown
coloration or ring at the line of contact of the heavier sul-
phuric acid and the mixed solution above (Fig. 63).

All nitrates are decomposed by heat. The sodium and

NITBIFICATION. 196

potassium salts, when thus treated, give up oxygen and are
converted into nitrites. Other nitrates yield oxygen and
nitrogen peroxide, and the oxide of a metal.

179. iritriflcation. — Nitrogen compounds are invariably
found in certain tissues of both plants and animals. The
nitrogen which helps to form these compounds comes from
the soil, since neither

plants uor animals, with
one exception about to
be noted, can take nitro-
gen from the air. Tha
problem how to oiain-
tain the supply n^tptrr
gen compi>uiif***^^ tl

th(?

Fig, 63.

^t^l,>^>^^f ^^^li somewhat
i^ifficult to solve. The
renewal takes place slo ,

tile because of the ' ^ ^^' "^ become infer-

difiSculty is usual' ^^^^ ^* nitrogen compounds. The
which ca«e th^ '^^ overcome by manuring the fields, in
nitrogen tCv ^ decomposing animal matter gives up its

j^ L,, 'O the soil as ammonia.
^y* ^ds been recently discovered that certain bacteria,
.xiich are found in the tubercles on the roots of legumi-
nous plants, such as peas and clover, have the power of
taking nitrogen from the atmosphere and converting it
into nitrates of bases found in the soil. This discovery has
been of great importance in rendering productive soils
which had remained infertile from the lack of nitrogen
compounds. The process of converting nitrogen from air
into nitric acid or nitrates is called nitrification. It is
difficult to accomplish because of the inactive character of
nitrogen. Nitrification is brought about, to a small extent.

196

NITROGEN COMPOUNDS.

Fig. 64.

by passing electric sparks
through air. The oxygen and
nitrogen unite, forming nitro-
gen peroxide, which in turn
forms nitric acid on dissolving
in water. During a thunder-
storm a certain amount of ni-
tric acid is formed in this way.
Attempts are now made to
use these reactions for the
manufacture of the acid.

180. Explosives — Certain
nitrogen compounds that con-
tain^3)2 -n oxygen, and hy-

^^ are so uft^'^^^^abl® ^^^^
th^^ompose under the im-

pulsexslight shock, form-
ing gas^ products. ^^^
this reason ^e compounds
are powerful e^sives. Ni-
troglycerine and giiP>\^^ *^®
examples of this class o^^"
pounds. Nitroglycerine is
made by treating glycerine, an
organic base, with the mixture
of nitric and sulphuric acids.
The sulphuric acid serves to
absorb the water that is formed
by the reaction of the other
two substances:

C3H5(OH)3 + 3HN03— ^
C3H5(N03)3 + 3 HgO

auMMAnr. 197

The nitroglycerine molecule is evidently a very unstable
one that can rearrange itself into new and more stable
molecules under the impulse of a slight shock. Dyna-
mite is a mixture of some inert materials and nitro-
glycerine.

Guncotton (nitrocellulose) is made by treating cotton
fibre (cellulose) with a mixture of nitric and sulphuric
acids. Cellulose, like glycerine, is an organic base, and
nitrocellulose is an unstable salt. One variety of smoke-
less powder is a special form of guncotton.

Some forms of explosives are shown on the opposite
page : a is granular gunpowder ; 6, cordite ; c, giant pow-
der ; d, brown prismatic powder ; e, high explosive.

Summary.

Ammonia is formed in nature as a decomposition product from
protoplasm.

It is obtained commercially as a by-product from the distilla-
tion of coal. It may be produced by the action of a base on an
ammonium salt.

Ammonia is a gas with a pungent odor; its specific gravity,
relative to hydrogen, is 8.5, It is very soluble in water, 1 liter
of water at 15° dissolves 720 liters of ammonia.

The solution is basic and contains NH4'*' and OH" ions. It
, reacts with acids with the formation of ammonium salts.

Ammonia is used as a refrigerating agent, and in the prepara-
tion of sodium bicarbonate and of ammonia water.

Nitrogen forms five oxides.

Nitrous oxide, N2O, is made by heating ammonium nitrate. It
is a good supporter of combustion, and is used as an anesthetic
(" laughing gas ")•

Nitric oxide, NO, is formed by the action of diluted nitric acid
on metals. It unites with oxygen at ordinary temperatines, forming

198 NITROGEN COMPOUNDS.

nitrogen "peroxide^ NOj. This is a brown, poisonous gas, soluble
in water.

NUroua anhydride, N2O8, and nitric anhydride, N2O5, are unim-
portant.

Nitric acid is prepared by the reaction between sulphuric add
and a nitrate.

When pure, it is a colorless liquid, with a specific gravity of 1.53.

It is a powerful oxidizing agent, and when it reacts with metals,
the hydrogen is oxidized to water and nitrogen oxides are liberated.
A mixture of nitric and hydrochloric add is aqua regia; this fur-
nishes nascent chlorine.

Nitric acid is used in the preparation of nitrates and explosives.

The nitrates of sodium and potassium are the most important.
Nitrates may be prepared by the action of nitric add on :

(1) metals;

(2) oxides or hydroxides ;

(3) salts yielding volatile products.

Potassium nitrate is used in gunpowder and as a meat preserva-
tive; sodium nitrate as a fertilizer and for the production of
potassium nitrate.

Nitrogen compounds are formed in the soils by bacteria, which
cause the nitrogen of the air to enter into combination.

Many nitrogen compounds are used in explosives. Nitroglycerine
and guncotton are made by the action of nitric acid on glycerine
and cotton respectively. Dynamite is a mixture of inert materiab
with nitroglycerine.

Exercises.

1. How many liters of ammonia can be obtained by the
action of lime on 50 grams of ammonium chloride ?

2. Why is ammonia called " spirits of hartshorn " ? why the
volatile alkali ?

3. What method would you use to get a few cubic centi-
meters of ammonia gas for use in the laboratory ?

EXERCISES. 199

4. Explain what is meant by the ammonium theory.

5. Show how the nitrogen oxides illustrate the law of mul-
tiple proportions.

6. By what tests would you distinguish between oxygen and
nitrous oxide ?

7. What volume of air would convert 100 c.c. of nitric oxide,
NO, into nitrogen peroxide, NO2 ?

8. Compare nitric acid with sulphuric acid and with hydro-
chloric acid in regard to its action with metals.

9. Explain the natural formation of (a) ammonia, (b) nitric
acid, (c) nitrates.

10. Why is nitric acid a better solvent than hydrochloric
acid for silver, mercury, and lead ?

11. What is aqua fortis f aqua regia f sal ammoniac f

12. Upon what properties of nitrie acid do most of its uses
depend ? Illustrate.

13. How would you test an unknown substance for the
nitrate ion ? for the ammonium ion ?

14. Explain the significance of the statement, "No life
without nitrogen."

15. State the substances and conditions necessary to yield
each of the following products from nitrogen compounds:
oxygen, hydrogen, nitric oxide, nitrogen peroxide.

16. Why has it become necessary to devote much attention
to the artificial production of fertilizers containing nitrogen ?

CHAPTER XXI.

ELEMEITTS OF THE NITBOaEK aSOlTF.

PHOSPHORUS.

181. Phosphorus, like nitrogen, is found in compounds
associated with all living matter, in the protoplasm and
bony structure of animals. A few, but widely distributed,
minerals contain it, and most of these seem to have been
derived from organic sources. Phosphorus was discovered
in 1669 by Brand, an alchemist of Hamburg, while dis-
tilling urine in the course of his attempts to find the
philosopher's stone. Scheele, the Swedish chemist, pre-
pared it from bones in 1771.

182. Preparation. — Phosphorus is prepared from bone-
ash or other phosphates. The phosphate is ground and
mixed in a vat with warm sulphuric acid, forming phos-
phoric acid and calcium sulphate:

Ca8(P04)2 -h 3 H2SO4 — ^ 2 HgPO^ + 3 CaSO^

The insoluble calcium sulphate is removed from the solu-
tion by filtering through coke. The phosphoric acid solu-
tion is concentrated, and absorbed in sawdust or coke, and
dried:

HgPO^— ^HPOg + HaO

The dry carbonaceous mass is then distilled in clay retorts
(Fig. 65) at a high temperature, phosphorus, hydrogen,
and carbon monoxide being produced :

2 HPOg-h 6 C— >.6 CO + 2 P -»- Hj
200

PREPARATION OF PHOSPHORUS.

201

The phosphorus vapor is condensed in water and solidifies
on cooling.

The electric furnace is used in a recent method for mak-
ing phosphorus. The phosphate is mixed with fine sand
and coke or coal and fed into an electric furnace. The
non-volatile product, calcium silicate, settles to the bottom
of the furnace, while
the phosphorus va-
por and carbon mo-
noxide are passed
into water and col-
lected :
Ca8(P04)2 + 8Si02

+ 5C— ^SCaSiOg

-h5CO + 2P

To purify the phos-
phorus, it is again
distilled or filtered
through bone-ash
and then run into
cylindrical moulds.

Fig. 65. Phosphorus distillation.
Qt retorts ; b, condensing vessels.

183. Physical Prop-
erties. — Phosphorus,
like sulphur, occurs in several allotropic forms. Prepared
by the methods given above, it is ordinary or yellow phos-
phorus, a straw-colored; translucent solid, about as hard
as beeswax, with a peculiar odor, due in part to ozone
formed by its action on the oxygen of the air. It melts
easily in water, but does not dissolve. In carbon disul-
phide, however, it is readily soluble.

Heated in a closed vessel to about 250° C, it becomes red
phosphorus^ which is red brown in color, opaque, and odor-

202 ELEMENTS OF THE NITROGEN GROUP.

less. Red phosphorus is not easily ignited and does not
dissolve in carbon disulphide. The red form is more dense
and more stable than the yellow form. Red phosphorus
can be converted into the yellow by distillation.

184. Chemical Properties. — Yellow phosphorus takes fire
at 35°, burning with a brilliant flame and forming dense
fumes of the oxide:

P4+5O2— ^2P206

In moist air it glows and slowly oxidizes. The ease with
which it burns makes it a very dangerous material. Yel-
low phosphorus is always kept under water.

Burns produced by phosphorus are very painful and
diflScult to heal. It is very poisonous; the vapor inhaled
by workmen in the factories produces an incurable dis-
ease, characterized by ulceration of the jawbones.

Phosphorus combines readily with chlorine, iodine, and
sulphur. Red phosphorus is less active than yellow; it
burns -with more diflSculty, evolving less heat. All its
actions indicate that the red variety has less energy than
the yellow. Heat is evolved in the conversion of the yel-
low to the red, so that the red is the more stable form.

185. Uses. — A little phosphorus is used in medicine and
in vermin poisons, but the larger part is consumed in the
manufacture of matches. A common friction match con-
sists of a stick of soft wood, tipped with sulphur, paraffin,
or other easily combustible material, and a head composed
of an oxidizing material, — as potassium chlorate or lead
oxide, — phosphorus and glue, with coloring matter. The
stick is dipped by machinery into melted paraffin which
soaks into the wood, water-proofing it. Then it is dipped
into a paste of the other materials and dried. The glue

ARSENIC. 208

protects the phosphorus from the air, but on rubbing, suffi-
cient heat is generated to ignite the phosphorus in contact
with the oxidizing material. This combustion will raise
the temperature sufficiently to ignite the paraffin, and the
burning of this will bring the wood to its kindling tem-
perature.

As such matches are poisonous and very easily ignited,
they are prohibited in many places and replaced by safety
matches^ the head of which consists of glue, antimony
sulphide, and an oxidizing agent. The box against which
they are rubbed has a surface of red phosphorus and
powdered glass and glue. As the head of the match is
soft, it will rub off on a rough surface and not burn,
but it will usually ignite on a hard smooth surface, like
glass or a slate blackboard.

186. Compounds — Phosphorus, like nitrogen, forms sev-
eral oxides and acids. Phosphoric oxide, P20g, is formed
when phosphorus burns with a sufficient supply of air or
oxygen. It is a white solid, which combines energetically
with water, forming phosphoric acid.

Phosphorus oxide^ ^2^9,^ forms when phosphorus bums
with a limited supply of oxygen. This white solid com-
bines with water, forming phosphorous acid.

Phosphoric acid^ HgPO^, has several salts of common
occurrence: ordinary sodium phosphate, NagHPO^, used
in medicine; calcium phosphate, Ca3(P04)2, the principal
mineral constituent of the bones ; calcium acid phosphate,
CaH4(P04)3, used in baking powder and fertilizers.

ARSENIC.

187. Arsenic is generally found in nature combined
with sulphur, associated with iron and copper. The ore

204 ELEMENTS OF THE NITROGEN GROUP.

is roasted, forming arsenious oxide, which is then reduced
with carbon:

AsaOg-h 3 C— >-2 As + 3C0

188. Properties. — Arsenic is a brittle, steel-gray, crystal-
line solid, with a metallic lustre, and tarnishes rapidly in
the air. It volatilizes without melting at the ordinary
atmospheric pressure and has an odor like garlic. It burns
with a bluish flame, forming the oxide, AsgOg. In its physi-
cal properties, arsenic resembles the metals, but in its ac-
tions> it resembles the non-metals, especially phosphorus.

189. XJses. — Arsenic is added to lead in the manufacture
of shot. The melted metal in a strainer or coUander is
dropped from a height into water. The arsenic lowers
the melting-point of the lead and makes it more fluid,
so that the shot becomes spherical before cooling. The
arsenic also makes the shot harder than pure lead.

190. Componnds. — Arsenious oxide^. AsgOg, is a white
crystalline powder, slightly soluble in water, and, like all
compounds of arsenic, poisonous. It is used in the manu-
facture of certain colors, also in medicine, and as a poison.
Arsenic sulphides^ realgar, AsgSg, and orpiment, AsgSg, are
used as pigments. Paris green is a copper and arsenic
compound used as pigment and as insecticide.

ANTIMONY.

191. Antimony is found combined with sulphur. It is
prepared in a manner analogous to that for arsenic, or by
heating the sulphide with iron:

SbaSg -f 3Fe — ^ 2Sb -f 3FeS

BISMUTH.

205

192. Properties and Uses. — Antimony is a handsome,
silver-white, crystalline, brittle solid, with a pronounced
metallic lustre. It does not tarnish in air, but when
heated in the air, burns, forming the oxide, SbgOg. As it
does not change in air, it is used to cover other mate-
rials, as brass and lead alloys. Antimony black is finely
powdered metal used to coat plaster casts imitating metal.
Antimony alloys are usually hard. Britannia metal and
pewter contain copper, tin, and antimony. Babbitt metal
and other anti-friction alloys generally contain antimony.

Lead contracts on solidifying ; sm alloy of antimony and
lead expands on solidifying and is hard. This alloy is
used for type metal.

BISMUTH.

193. Properties and Uses. — Bismuth resembles antimony,
but is more metallic. It has a red tinge, is brittle, crys-
talline, heavy, and tarnishes slowly in moist air.

Bismuth alloys are remarkable for their low melting-
points. Bismuth melts at 270° C. The two most common
alloys. Wood's metal and Rose's metal, have the following
composition :

Such easily melted alloys are extensively used: as fuses
in electric connections, in fire alarms, in safety plugs, in
boilers, and in automatic sprinklers in buildings. When

206

ELEMENTS OF THE NITROGEN GROUP,

the fusible plug of a sprinkler (Fig. 66, a)
melts, the water rushes out from the
main and strikes a cap &, scattering the
water in all directions. A piece of
Wood's metal is sometimes placed in the
gas-pipe where it enters the building, so
that in case of fire the alloy will melt
and stop the flow of gas.

Fig. 66.

194. Comparison of the Nitrogen Group

Atomic
Wbight.

Mbltino-

POINT.

Htdbidks.

OXIDU.

Acids.

Nitrogen

-210°

NHj

NjO, NO,

NA
NOs, N,Oj

Strong acid

HNOa

HNO3

Phosphorus

440

PH3

PA. PA

H3PO3, H3PO,
Weak acid

Arsenic

185°

A8H3

AsjOg. ASjO,

H3As03,H3A804

Antimony

120

630°

SbH3

SbjOg, SbjO,

H3SbO,

Bismuth

208

270°

BiA

Bi(OH),^ Base

Sximmary.

Phosphorus exists in two allotropic forms. Both varieties, espe-
cially the yellow, have a great tendency to unite with oxygen. It
also reacts readily with other non-metals.

Phosphorus is extracted from bone-ash.

Phosphorus is chiefly used for making matches.

Arsenic^ though a non-metal, shows some of the characteristics
of metals. It alloys with other metals and is used to make shot
hard. Some of its compounds are valuable as paints.

Antimony shows the characteristics of both metals and non-
metals. It is a constituent of type metal and other alloys.

Bismvih is a metal used in many alloys. These alloys generally
have low melting-points.

EXERCISES. 207

The dements in this group resemble each other in properties to a
considerable d^ee. It is approximately true that in going through
the group, a given property changes steadily in one direction as the
atomic weights increase. Thus nitrogen is a colorless gas ; phos-
phorus is a waxlike solid; arsenic is a dark gray solid with some-
thing of the appearance of a metal ; antimony has a distinctly
metallic appearance; bismuth is a metal. Their respective specific
gravities are: 0.97 * 1.8, 5.7, 6.7, 9.7.

The elements of the family form many compounds similar in
character and formula. AU except bismuth form compounds of
the type XHs, where X stands for the symbol of any element in
the family. There are two oxides, XsOs and XjO^, which are the
anhydrides of the adds HXO2 and HXOa respectively. In the
cases of phosphorus, arsenic and antimony, the acid formulas are
HgXOs and H8XO4, showing the addition of three molecules of water
to the anhydride instead of one.

Bzercises.

1. What is the per cent of phosphorus in calcium phosphate
of the composition Ca8(P04)a?

2. What weight of oxygen would be used in combining with
0.5 gram of phosphorus ? What would be the volume of the
oxygen at standard conditions?

3. What weight of phosphorus would be necessary to re-
move the oxygen from 10 liters of air (measured at standard
conditions)?

4. Why is yellow phosphorus always covered with water
in the vessels in which it is kept?

5. Why do matches ignite on being rubbed?

* Air = 1 ; the other specific gravities refer to water.

CHAPTER XXII.

THE HALOaENS.

The elements fluorine, chlorine, bromine, and iodine are
called halogens (salt formers), because they unite directly
with a large number of metallic elements to form salts.

BROMINE.

195. Occurrence. — Bromine was discovered in 1826 by
Ballard, who separated it from the mother-liquor of sea-
salt.

Large deposits of crude salt, impure sodium chloride,
occur in the states of Michigan, West Virginia, Ohio, and
Pennsylvania. Magnesium bromide and sodium bromide
are two of the substances mixed with the sodium' chloride.
When the brines from the deposits mentioned are evapo-
rated, nearly all of the sodium chloride crystallizes out
before the magnesium salts begin to separate in an appre-
ciable quantity. The liquid remaining, after a portion of
the substances contained in the original solution has crys-
tallized, is known by the technical term of mother-liquor.
The compounds of magnesium remaining in the mother-
liquor of salt works impart to it a bitter taste and cause it
to be known as bittern. Extensive deposits of magnesium
salts containing bromides are found in the almost inex-
haustible salt beds at Stassfurt, Germany. The United
States and Germany furnish a large percentage of all the
bromine used. .

208

BROMINE.

209

196. Preparation. — Bromine can be prepared from the
bromides by a method analogous to one of the methods
described for the preparation of chlorine ; namely, by heat-
ing a mixture of a bromide,
manganese dioxide, and
sulphuric acid (Fig. 67).
Bromine, having a low boil-
ing-point, passes off in the
state of vapor, which can
easily be liquefied by keep-
ing the receiver cool.

If a small quantity is
made, the bromine vapor
can be condensed in a test-
tube partly filled with water
(see a and b in Fig. 67).

The reaction may be con-
sidered as taking place in
three steps. .

Sulphuric acid reacts
with potassium bromide to
produce potassium sulphate and hydrobromic acid:

2 KBr -h H2SO4 — >- K2SO4 -f 2 HBr

When warm sulphuric acid is added to manganese diox-
ide, manganese sulphate, water, and oxygen are formed:

MnOa + H2SO4 — ^ MnSO^ -h H2O + O

Nascent oxygen converts hydrobromic acid into water
and bromine:

2HBr + 0— ^H20-fBr2

The equation for the complete reaction is:

2KBr-fMn02-h2H2S04— ^K2S04+MnS04+2H20-hBr2

Fig. 67. Preparation of bromine.

210 THE HAL00EN8.

197. Physioal Properties — Bromine is a dark brownish-
red liquid, about three times as dense as water. It is the
only non-metallic element that, under ordinary conditions,
exists in the state of a liquid. Bromine has an odor some-
what resembling that of chlorine; its name is derived from
a Greek word meaning stench.

When a bottle of bromine is opened, the brownish-red
vapor of bromine can be seen issuing from its mouth. If
a few drops of bromine are poured into a large bottle filled
with air, the vapor is seen first at the bottom of the bottle,
and it diffuses slowly until it fills the bottle.

The vapor of bromine has a strong corrosive action on
the mucous membrane. When it comes into contact with
the eyes, the irritation is sufficient to cause a copious flow
of tears. Great care should be taken not to inhale bromine
vapor, and never to allow the bromine to come in contact
with the skin. If bromine is inhaled, the irritation can be
lessened by smelling of chloroform or alcohol. When
bromine comes in contact with the skin, the injured part
should be washed freely with water and then covered with
a paste made by mixing sodium bicarbonate with water,
or better with some oil, olive or cocoanut.

Bromine is somewhat soluble in water ; the solution is
called bromine water. Bromine is more soluble in aqueous
solutions of the bromides than it is in pure water. It is
very soluble in chloroform and in carbon disulphide.

198. Chemical Properties. — The chemical behavior of
bromine very closely resembles that of chlorine. Bromine
is, however, not so active* an element as chlorine. We can
illustrate this fact by comparing the action between chlo-
rine and hydrogen with that between bromine and hy-
drogen. We have already seen that when a mixture of

HYDBOBBOMIC ACID. 211

chlorine and hydrogen is placed in the sunlight, the ele-
ments combine with explosive violence to form hydrogen
chloride. Under similar conditions, bromine vapor and
hydrogen enter into only a partial combination, without
any display of energy. Aqueous solutions of bromine
bleach many dyes, but the action is not as rapid as in
the case of chlorine.

Bromine combines directly with a number of elements, as
phosphorus, antimony, copper, and iroi;i, forming bromides.

199. Uses. — Bromine is used in the manufacture of
many organic compounds, especially certain of the aniline
dyes, and in the preparation of bromides. Its water solu-
tion is used in the laboratory as an oxidizing agent.

HYDROBROMIC ACID.

200. Preparation. — Hydrobromic acid can be prepared
by the direct combination of bromine with hydrogen; the
method is of no practical importance.

The addition of sulphuric acid to a bromide would prob-
ably appear to be a convenient method for the preparation
of hydrobromic acid. A dilute solution of hydrobromic
acid is prepared, on a commercial scale, by the reaction be-
tween diluted sulphuric acid and a solution of potassium
bromide. The actual carrying-out of the process requires
considerable time and most careful attention.

When concentrated sulphuric acid is added to potassium
bromide, hydrogen bromide appears as a gas which fumes
as soon as it comes in contact with the air. Other gases
are formed at the same time ; the odor of sulphur, dioxide
can generally be detected, and sometimes that of hydrogen
sulphide. The products formed vary with the concentra-
tion of the sulphuric acid and the temperature at which

212 THE HALOGENS.

the reaction takes place. The more concentrated the sul-
phuric acid, and the higher the temperature, the less will
be the amount of the hydrogen bromide produced.

Let us consider the reactions involved when sulphur
dioxide is formed. Sulphuric acid reacts with potassium
bromide to form potassium sulphate and hydrobromic
acid:

2 KBr + H2SO4 — >- KgSO^ -f 2 HBr

The excess of concentrated sulphuric acid, however, oxi-
dizes the hydrobromic acid, the result of the oxidization
being water and bromine:

2HBr + H2SO4— ^2H20 -f- SO2 +Br2

This is similar to the action of hot, concentrated sul-
phuric acid with copper (§ 139), in which a portion of the
sulphuric acid is reduced to sulphur dioxide, and at the
same time water is formed. The fact should be con-
stantly kept in mind that when one substance is oxi-
dized, some other substance is reduced.

If we use one equation to represent the formation of
bromine and sulphur dioxide, by the method just consid-
ered, we obta^in:

2 KBr H- 2H2SO4— ^K2S04 -f 2H2O + SO2 4- Br2

201. Properties. — Hydrobromic acid is a colorless gas,
readily soluble in water, and its solution possesses the
characteristic properties of a strong acid. Solutions of
it are easily oxidized by the oxygen of the air, water and
bromine resulting from the oxidation.

Dilute solutions of hydrobromic acid are used to some
extent for medicine, and the bromides are an important
series of salts.

Silver bromide separates as a yellowish white precipitate

REPLACEMENT OF BROMINE. 213

when a solution of silver nitrate is added to a solution of a
bromide. It is insoluble in nitric acid, slightly soluble in
dilute ammonium hydroxide, and more readily soluble in
concentrated ammonium hydroxide. These reactions are
sometimes used as a test for a bromide.

202. Keplacement of Bromine. — When chlorine is added
to a solution of a bromide, free bromine appears and chlo-
rine molecules pass into chlorine ions. The solution of po-
tassium bromide contains potassium ions, bromine ions, and
molecules of potassium bromide; the undissociated and dis-
sociated potassium bromide being in equilibrium. As soon
as chlorine is added, the bromine ions give their negative
charge of electricity to the chlorine molecules, which
then dissociate into ions. The bromine ions, having lost
their charge of electricity, unite to form bromine mole-
cules. As soon as some of the bromine ions pass out of
solution, the equilibrium between the dissociated and the
undissociated potassium bromide is destroyed, and more
molecules of potassium bromide dissociate. If enough
chlorine is added, all the bromine ions will finally appear
as bromine molecules, and the solution will contain potas-
sium ions, chlorine ions, molecules of potassium chloride,
and molecules of bromine.

2K-^ + 2Br- 4- Clg— ^2K+ -f 201" -h Bv^
or 2 KBr + CI2 — ^ 2 KCl + Bv^

If there is too little bromine to be noticed in the pres-
ence of the yellow color produced by an excess of chlorine,
its presence can be shown by adding a little chloroform or
carbon disulphide, and shaking. Water and chloroform
are not miscible (§ 37), and bromine is much more soluble
in chloroform than it is in water. The bromine wiU be

214

TBM HAL0QBN8.

distributed between the water and the chloroform in pro-
portion to the ratio of its solubility in the two liquids.

The solution of bromine in chlo-
roform has a characteristic
color ; the bromine must be free,
for combined bromine does not
produce the characteristic color
with chloroform, as is shown
by shaking this liquid witli a
solution of potassium bromide
(Fig. 68, a). Since a small
quantity of chloroform can be
used to remove nearly all of the
free bromine from a compara-
tively large quantity of water
by shaking, the process is called
shaking out.

The liberation of bromine by
chlorine, followed by shaking out with chloroform or car-
bon disulphide, is used as a test for bromine ions. If we
add chlorine water to a solution of a bromide, and then
shake with chloroform, the latter dissolves the free bro-
mine, acquiring the characteristic reddish yellow colora-
tion (Fig. 68, 6).

Fig. 68.

IODINE.

203. Iodine was discovered by Courtois in 1812 while
trying to prepare potassium nitrate from liquors obtained
by washing the ashes of burnt seaweed. During his exper-
iments Courtois observed the violet color of the vapor of
iodine, but the properties of the element were first care-
fully studied by Gay-Lussac.

IODINE.

215

204. Preparation. — When seaweed (kelp) is burned at
a low temperature, the ash contains considerable quanti-
ties of the iodides of potassium and sodium. As both of
these salts are readily soluble in water, they can be sepa-
rated from the insoluble portion of the ash by leaching
or liodviation; that is, by allowing water to slowly pass
through the kelp.

Iodine is obtained from potassium iodide by a process
analogous to that described for the preparation of bromine
(§ 196); the iodide is warmed with manganese dioxide
and sulphuric acid. Iodine passes off in the form of a
vapor:

2 KI-hMnOg-f 2 HgSO^— ^K2S04+ MnSO^ + 2 HgO + I2
Large deposits of impure sodium nitrate are found in the
dry region west of the Andes. Compounds of iodine occur
in these deposits and most of the iodine used to-day is
obtained from the mother-liquor of the sodium nitrate
works.

20s. Physical Properties. — Iodine is a steel-gray solid
(Fig. 69), very slightly
soluble in water, but
readily soluble in alco-
hol, chloroform, carbon
disulphide, and in aque-
ous solutions of potas-
sium iodide. A solu-
tion of iodine in alcohol
is called tincture of
iodine. Solutions of io-
dine in chloroform and
in carbon disulphide
possess a characteristic violet color ; iodine vapor has the

Fig. 69. Iodine crystals.

216 THE HALOGENS.

same color. Iodine vaporizes slowly at ordinary tem-
peratures. When the solid is warmed, the change
takes place rapidly, and the vapor on being cooled passes
directly to the state of a solid. Such distillation of a
solid is called Bvhlimation^ and may be used to purify
solids that can be sublimed, as distillation is used to
purify liquids.

The fact that iodine is more soluble in a solution of an
iodide than in pure water is explained by the supposition
that iodine, ions combine with the iodine molecules to form
triiodine ions:

The difference between the color of a solution of iodine in
chloroform and the color of a solution of iodine in an aque-
ous solution of an iodide is due probably to the difference
in the number of atoms in the particles of iodine entering
the solution.

206. Chemical Properties. — Iodine unites directly with
many elements to form iodides. The reactions are not as
energetic as in the case of either chlorine or bromine.
When a piece of yellow phosphorus and a piece of iodine
are brought together, they combine to form an iodide with-
out the application of heat (Fig. 70). Iodine and iron
unite when heated.

When a dilute solution of iodine is mixed with a dilute
solution of starch paste, a characteristic blue color is pro-
duced. The reaction is made use of in testing for both
iodine and starch.

Iodine is used in the manufacture of certain compounds
which are used in medicine, in photography, and for
dyeing.

HTDRIODIC ACID.

217

HYDRIODIC ACID.

207. Preparation. — If concentrated sulphuric acid is
added to an iodide, the odor of hydrogen sulphide is very
noticeable. More hydrogen
sulphide is produced than
was formed when sulphuric
acid was added to a bromide.
This means that hydriodic
acid is more easily oxi-
dized (or is a better re-
ducer) than hydrobromic
acid.

The formation of iodine
by the action of sulphuric
acid with potassium iodide
can be represented by the
following equations :

8 KI + 4 H2SO4— ^iEgSO^
4- SHI

H2SO4 + SHI— ^ HgS + 4H2O + 81

The equation for the complete reaction is:

SKI + 5H2SO4 —^4X2804 -h 4H2O + H2S + 81

Hydriodic acid can be readily prepared by the reaction
of water with iodine and red phosphorus:

Fig. 70.

P + 31
Pig + 3H2O

Pis

HgPOg + 3HI

208. Iodides. — The iodides are important compounds,
finding extensive use in medicine.

Both chlotine and bromine liberate iodine from the
iodides. The presence of free iodine can be determined

218

THE HALOGENS.

by the starch test or by shaking out with chlorofoim or
carbon disulphide.

THE HALOGENS AS A GROUP.

209. A comparison of the properties of the three halo-
gens, chlorine, bromine, and iodine, leads to some interest-
ing results:

Name of
Elsvknt.

Atomic
Weight.

State.

Solubility in

ONE PABT OF
WaTBB AT

16° C.

Heat of Forma-
tion OF Hydro-
gen COMP.

Heat of For-
mation OF Po-
tassium COMP.

Chlorine

35.5

Gas

2.4 vol.

22,000 caL

104,300 cal.

Bromine

Liquid

0.032 pts.

8,400 cal.

95,100 caL

Iodine

127

Solid

0.00015 pts.

-7,000 cal.

80,100 cal.

Chlorine, bromine, and iodine form a natural group of
elements. The difference between the atomic weight of
bromine and that of chlorine is nearly the same as the
difference between the atomic weights of iodine and
bromine. An examination of the table will show that the
tabulated properties of the elements considered vary in
degree with the atomic weights. Chlorine, bromine, and
iodine very closely resemble each other in their chemical
behavior. The chemical activity of bromine is less than
that of chlorine and greater than that of iodine.

210. Heat of Formation. — When a chemical change
occurs without the addition of energy, the substances
resulting from the reaction usually contain less energy
than the original constituents. Chemical energy has been
transformed into some other kind of energy. In order to

HEAT OF FORMATION. 219

gain information in regard to chemical energy, it must be
changed to some form of energy which can be measured.
Heat is the form of energy most frequently considered in
connection with chemical changes.

When elements unite to form chemical compounds, the
heat evolved or absorbed is called the heat of formation of
the compound in question. In measuring the heat of
formation of any compound, weights of the substances
equal to their reacting weights expressed in grams are
■ considered, and the quantity of heat is commonly expressed
in calories.

Hydrogen unites with chlorine to form hydrogen chlo-
ride. The simplest equation representing the reaction is:

H + CI ^ HCl

This shows that 1 gram of hydrogen unites with 35.5
grams of chlorine to form 36.5 grams of hydrogen chloride.
During the combination, 22,000 calories of heat are evolved.
The thermal equation for the formation of hydrogen
chloride is:

H + CI — ^ HCl + 22,000 calories

The heat of formation of hydrogen chloride is 22,000
calories.

When hydrogen combines with iodine to form hydrogen
iodide, an absorption of heat occurs. The thermal equa-
tion reads:

H + I — ^ HI - 7000 calories

This shows that when 1 gram of hydrogen unites with 127
grams of iodine 7,000 calories of heat are absorbed. The
heat of formation of hydrogen iodide is —7000 calories.

211. Eelative Eeplacement and Heats of Formation. — Let

us use the double arrow in the following equation to indi-

220 THE HALOGENS.

cate the two possible directions in whieli the reaction
might proceed:

HCl + Br :5±: HBr + CI

Would bromine replace chlorine or would the reverse be
true? We may ask whether the displacement which
would actually take place would not be in accordance
with some general law governing chemical action. The
following generalization has been developed from the
study of the heat effects of many chemical changes.

• When a chemical reaction takes place without the addition '
of heat from an external source^ those substances which have
the greatest heat of formation will tend to form.

The heat of formation of hydrogen bromide is 8400 cal-
ories; that of hydrogen chloride is 22,000 calories. Thus
more heat is liberated when hydrogen unites with chlorine
than is liberated when hydrogen combines with bromine.
We should therefore expect chlorine to liberate bromine
from hydrogen bromide. Chlorine does liberate bromine
from hydrogen bromide. Chlorine and bromine liberate
iodine from hydrogen iodide, as the heats of formation of
hydrogen chloride, bromide, and iodide would lead us to
expect. A study of the heats of formation of chemical
compounds has been of value in the prediction of chemical
reactions.

It should be remembered that heat is not the only form
of energy into which chemical energy is converted, and in
cases of solution in which chemical compounds are disso-
ciated, the energy necessary to dissociate the compounds is
an important factor in the thermal equation.

FLUORINE.

212. Activity. — Fluorine belongs to the halogen group,
but does not so closely resemble the other members of the

FLUORINE.

221

group as they resemble each other. A consideration of
fluorine has therefore been made to follow a study of the
other members of the group.

Fluorine is an element of unusual chemical activity;
few substances are not attacked by it. It cannot be iso-
lated in the presence of water, as it unites with the hydro-
gen in the solution and liberates the oxygen. Further-
more, it cannot be prepared in glass vessels, as it reacts
with the glass. The heats of formation of the fluorides are
too great to permit of their being easily decomposed by
heat. From the statements just made it will be seen that
fluorine cannot be prepared by the methods generally em-
ployed in the preparation of the other halogens.

213. Preparation. — The problem of isolating pure fluo-
rine puzzled chemists until 1886, when Moissan discovered
that a solution of potassium fluoride in liquid hydrofluoric
acid conducted the electric current.
The apparatus used by Moissan to
carry on the electrolysis consisted
of a U-tube made of an alloy of
platinum and iridium, carrying
electrodes composed of the same s
material, which were insulated from
the U-tube by calcium fluoride stop-
pers (Fig. 71).

The solution of potassium fluo-
ride in hydrofluoric acid was placed
in the U-tube and kept at a tem-
perature near —23° C. during the
electrolysis. Moissan subsequently
found that a U-tube made of copper could be substituted
for the one composed of the expensive alloy mentioned.

Fig. 71.

222 THE HALOGENS.

During the electrolysis, fluorine is liberated at the anode
and passes off through the side arm of the tube surround-
ing it. Hydrogen is liberated at the cathode. Potassiiun
passes to the cathode, but, on giving up its electric charge,
instantly unites with fluorine, forming potassium fluoride,
which dissolves in the excess of hydrofluoric acid. The
result of this electrolysis is that only the hydrofluoric acid
is permanently decomposed.

214. Properties. — At ordinary temperatures fluorine is
a nearly colorless gas, much more poisonous than chlorine.
Liquid fluorine combines energetically with hydrogen, sul-
phur, phosphorus, arsenic, some other elements, and mauy
compounds, showing that violent chemical action can take
place at a very low temperature.

Under ordinary conditions, fluorine has a greater ten-
dency to form compounds than any other element. Copper
when placed in fluorine becomes coated with an insoluble
coating of copper fluoride. The fluorides of calcium,
strontium, and barium are insoluble. Silver fluoride is
soluble. Gold and platinum, which readily form com-
pounds with nascent chlorine, are very slowly attacked by
fluorine. No oxide of fluorine is known. It is interest-
ing to compare the properties of fluorine just mentioned
with those of chlorine, bromine, and iodine.

Two compounds of fluorine which occur in nature are of
importance: calcium fluoride or fluor-spar, CaFj, and cryo-
lite, 2NaF.AlFg. Cryolite is used in the preparation of
aluminum by the electrolytic process.

215. Hydrofluoric Acid. — Hydrofluoric acid is prepared
by the action of sulphuric acid with fluor-spar:

CaFa + H2SO4 — ^ CaSO^ + 2 HF

BTDnoPLUOniC ACID.

228

SThe reaction is commonly carried on in a lead or platinum
dish.

Pure hydrofluoric acid is a colorless liquid which fumes
strongly in air. It dissolves readily in water and aque-
ous solutions of it are sold in wax bottles. Great care
should be exercised in using this acid, as painful sores,
difficult to heal, are produced when it comes in contact
with the skin.

The chief use of hydrofluoric acid is in the etching of
glass. Glass is composed of silicates, and hydrofluoric
acid converts the silica, silicon dioxide, of the glass into
a gas, silicon fluoride, and water:

Si02 + 4 HF — ^ SiF^ + 2 HgO

Glass is prepared for etching by covering it with a coat-
ing of some substance
which is not attacked
by hydrofluoric acid,
such as paraffin or a
mixture of beeswax and
rosin, and then remov-
ing the coating from the
part to be etched. The
etching is accomplished
by subjecting the pre-
pared piece to hydroflu-
oric acid vapor, or by
applying to it a water solution of the acid (Fig. 72).
When the gas is used the surface of the etching is left
dull, while with a water solution it is left glossy. Hydro-
fluoric acid is used in the finishing of cheap cut glass, and
for the removal of sand from castings.

Fig. 72. Etching with hydrofluoric acid.

224 THE HALOGENS.

Summaiy.

Bromine, atomic weight, 80, resembles chlorine very closely. The
chief points of difference are that it is a dark-colored liquid and
that it is less active. In general, bromine reactions are like those
of chlorine except for the difference in intensity.

Bromides are found in nature associated with chlorides. Bromine
is prepared from bromides by a process exactly analogous to that
used in the preparation of chlorine from sodium chloride. A
mixture of a bromide and manganese dioxide is treated . with
concentrated sulphuric add. The manganese dioxide acts as an
oxidizing agent on the hydrobromic add that is formed by the
action of the sulphuric add with the bromide.

Hydrobromic add is formed by the action of moderately concen-
trated sulphuric add with potassium bromide. There is a marked
difference here from the action that occurs in the preparation of
hydrochloric add. The difference is due to the fact that the hydro-
bromic add is less stable {i,e. has a less heat of formation) than
hydrochloric acid. The secondary products formed in the case of
hydrobromic add are mainly .sulphur dioxide and free bromine,
resulting from the oxidizing action of the sulphuric add.

Bromine and its compounds are not of great commerdal impor-
tance. Bromides are used in medicine, and free bromine is used
in the preparation of certain dyes.

Since bromides have smaller heats of formation than the corre-
sponding chlorides, free chlorine will displace bromine from bro-
mides. The presence of free bromine in solution may be detected
by the color it imparts to carbon disulphide or chloroform.

Iodine, atomic weight, 127, is found as iodide in small quantities
in the ashes of sea-plants, and associated with sodium and potas-
sium compounds. It is a dark gray solid with the suggestion of a
metallic appearance. In its chemical properties it is like chlorine
and bromine, but it is less active ihan either. Thus we find that
with these three elements, as the atomic weight increases, the

EXERCISES.

225

activity diminishes, and the elements lose something of their non-
metallic character.

Iodine may be prepared by a process analogous to that used in
the preparation of chlorine or bromine.

Hydriodic add results from the action of moderately concentrated
sulphuric acid with potassium iodide, but only a small part of the
expected quantity is obtained. The heat of formation of hydriodic
add is so low that it is very readily oxidized by the sulphuric add.
Sulphur dioxide, free sulphur, hydrogen sulphide, water, and free
iodine may all be formed in this secondary action.

Iodides have a less heat of formation than bromides. Conse-
quently free bromine vnU displace iodine from iodides. Chlorine
will do the same thing, even more readily.

Fluorine, atomic weight, 19, is the most intensely active element
known. Hydrofluoric add is a stable compound used in etching

The four halogens, considered as a group, may be regarded as
the most nearly perfect representative of a chemical family. The
properties change in a very definite and regular way with the
change in atomic weight.

Bzercises.

1. Prepare a table of the halogens according to the following
form :
a
Element

h
Atomic
weight

/
Chemical
properties

Compounds
found in
nature
(Formulas)

Equations
for the
preparation
of element

Relative

replacing

power

Equations
for the
preparation
of the
hydrogen
compound

Physical
properties

Stability
of the
hydrogen
compound

226 THE HALOGENS.

2. Write equations for the preparation of potassium bromide
by three different methods.

3. A solution contains potassium chloride and potassium
bromide. How would you show the presence of the two halo-
gens?

4. What weight of iodine could be obtained from 150 grams
of potassium iodide? How much sulphuric acid would be
needed for the action, assuming that the acid sulphate of potas-
sium is produced?

5. Solutions of hydriodic acid become dark colored on
standing. To what substance is the dark color probably due ?
Explain the action that produces it.

6. A solution of sodium bromide is treated with an excess of
silver nittate, and 0.65 gram of silver bromide is precipitated.
What weight of sodium bromide was contained in the original
solution ?

7. Why should bottles of bromine water be kept in a dark
place ?

8. Write the equations representing the reactions of bromine
with zinc, water, antimony, and hydrogen. Name the products
in each case.

CHAPTER XXIII.
OABBON.

216. UniuiLal Character. — Carbon, though a non-metal,
dififers in several important respects from the other ele-
ments of its class. The chief difference is that it forms
almost numberless compounds with hydrogen. These sub-
stances are called hydrocarbons. In addition to these, there
are many compounds that are composed of carbon, hydro-
gen, and oxygen; others are known which contain in addi-
tion chlorine, bromine, iodine, nitrogen, or sulphur. There
are so many of them that the branch of chemistry which
deals with them has a special name, that of organic chem-
i%try. The compounds of this sort that were first known
were found only as products of animal or vegetable life;
that is, of organized bodies — hence the term organic. It
was formerly believed that such substances could not be
made artificially, but to-day many of them are manufac-
tured in large quantities. A very large number of organic
compounds that are not known in nature have been made
in the laboratory.

217. Importance of Carbon in Nature. — Every living
thing, plant or animal, contains carbon in its tissues in the
form of organic compounds. The higher plants always
contain a large amount of cellulose, which is composed of
carbon, hydrogen, and oxygen. The woody part, which
furnishes the strength necessary to keep the plant erect,
is cellulose. All plant fibres, such as cotton and hemp,

227

2:28 CARBON.

are composed of cellulose. Cotton fibre is the purest form
of cellulose that is known.

The muscular tissue of animals is composed of a class
of bodies known as proteids. They contain carbon, hydro-
gen, oxygen, and nitrogen, with a very small quantity of
other elements. The fat of animals is composed of car-
bon, hydrogen, and oxygen. The starting material for
the formation of all these complicated compounds is the
carbon dioxide of the air. The green leaves of plants con-
tain chlorophyl bodies, substances by which starch is made
from carbon dioxide and water under the influence of sun-
light as a source of the necessary energy. This synthesis
has never been accomplished in the laboratory. The
starch is converted in the plant into other compounds.
These serve in turn as the food of animals, and furnish
material for the building of animal tissues.

218. Occnrrence. — It is apparent from what has been
said that a very large number of compounds of carbon are
found in the realm of organic nature. The element is also
found in carbon dioxide of the atmosphere, and very
abundantly in the carbonates of certain metals, especially
calcium and magnesium. Marble and limestone are two
different varieties of calcium carbonate. A considerable
part of the earth's crust is made up of these materials.

In the uncombined form carbon is found as coal, graph-
ite, and diamond.

219. Coal. — There are two chief forms of coal and sev-
eral less important varieties. None of them is pure car-
bon. Anthracite or hard coal (Fig. 73, d) contains 90 %
to 98 % of uncombined carbon. Bituminous or soft coal
(Fig. 73, c) contains 60 % to 70 % carbon. The remain-

COAL.

229

Fig. 73. Fuels :
J, peat ; b, lignite ; c, soft coal ; d, anthracite.

230 CARBON.

ing 40 % to 30 % is chiefly hydrocarbons. These can be
driven off as gases by heating the coal without access of
air. This operation furnishes a means of making one
kind of illuminating gas.

Coal is fossil vegetable matter. During a part of the
earth's history, known in geology as the carboniferous
period, vegetation flourished to a remarkable extent.
Much of this vegetable matter was buried under beds of
mud and sand. In this condition it underwent very slow
partial decomposition. A large part of the hydrogen and
oxygen was driven off, and the remainder, consisting of
uncombined carbon, was left as coal. In anthracite coal
the decomposition has proceeded further than in bitumi-
nous coal.

Some forms of coal contain even less uncombined carbon
than does the bituminous variety. Cannel coal and lignite
belong to this class. Lignite exhibits much of the struc-
ture of the wood from which it was derived (Fig. 73, 6).
Peat is moss or other loose vegetable matter that, to a
slight extent, has undergone a change like that by which
coal was formed (Fig 73, a).

OTHER COMMERCIAL FORMS OF CARBON.

220. Lampblack. — Lampblack, or soot, is practically
pure carbon. It is best made by burning hydrocarbon
oils with a limited supply of air. The operation is
analogous to that which takes place when a kerosene lamp
smokes. The hydrogen of the oil butns, but much of the
carbon remains unburned because of the lack of oxygen.
The unburned carbon is deposited as a soft, amorphous
powder. Lampblack is used in making paint and print-
ers' ink.

LAMPBLACK.

281

Large quantities of lampblack are made from natural
gas by using machines of special design. One of these
lampblack machines (Salsburgh patent) is shown in Fig-
ure 74.

The cast-iron disk (2>), about four feet in diameter, is
mounted on a hollow shaft. The burner ((3^) is made of
one-inch gaspipe
in which are
placed lava tips
about two inches
apart. The burn-
er is mounted
close to the disk
in such a position
that the flames
from the tip will
be divided, half of
the flame deposit-
ing a thin coating
of soot, or lamp-
black, on the con-
cave surface, and the other half of the flame depositing a
thin coating on the convex surface of the disk.

As the disk revolves slowly in the direction of the
arrow, the scraper (/S) removes the lampblack from the
disk and the black powder falls through the hopper (^H)
to a spiral conveyor not shown in the cut. The conveyor
delivers the lampblack to the bolters, where it is sifted and
prepared for sacking.

The top of the disk is kept full of running water, which
cools the disk and then empties into the hollow shaft.
The disks are run in groups of fifteen in rooms kept
closed so as to avoid draughts.

Fig. 74. Manufacture of lampblack.

232 CARBON.

221. Wood Charcoal. — Wood charcoal is made by heat-
ing wood without access of air. To a certain extent
this operation resembles the natural process by which coal
was formed. The hydrogen, oxygen, nitrogen, and some
of the carbon which the wood contains pass off in the
form of various compounds. The operation is known as
destructive distillation (Fig. 120, page 876). The char-
coal which remains is uncrystallized in form, but it retains
the cell structure of the wood from which it was formed.
Charcoal is not pure carbon, as is shown by the fact that
it leaves an ash on burning.

222. Bonebla'bk. — Boneblack is a form of charcoal that
is obtained by heating bones without access of air. Bones
contain both mineral and animal matter. The animal
matter, consisting of carbon compounds, is converted into
charcoal by destructive distillation; the mineral matter,
which is calcium phosphate, is separated from the charcoal
by being dissolved in hydrochloric acid. Boneblack is
extensively used for filtering purposes. It has the power
of absorbing both coloring-matter and gases from solution.
Other forms of carbon have the same property in a less
degree. It is attributed to the porous character of char-
coal. The most important use of boneblack is in the
decolorization of sugar and oils.

223. Coke and Oas Carbon. — Coke is an amorphous form
of carbon obtained by the destructive distillation of bitu-
minous coal. It is obtained as a by-product in the manu-
facture of one kind of illuminating gas.

Gas carbon is found on the top and sides of iron retorts
used for the destructive distillation of bituminous coal in
the manufacture of illuminating gas. It differs from coke

ALLOTROPIC FORMS. 288

in being nearly pure carbon ; coke leaves coHsiderable ash
on burning. Gas carbon is a fairly good conductor of
electricity and is used for making the carbon rods for arc
lights.

ALLOTROPIC FORMS.

224. Among all the varieties of carbon, three distinct
allotropic forms are recognized. These are amorphous
carbon, of which lampblack is the purest form ordinarily
obtained, graphite, and diamond. The different varieties
of charcoal, coal, and coke are impure forms of amorphous
carbon.

The three allotropic forms can each be burned in oxygen,
if raised to a suflBciently high temperature, with the forma-
tion of nothing but carbon dioxide. This fact proves that
the three substances are different forms of the same ele-
ment. They differ widely in physical properties and in
the ease with which they bum.

225. Amorphous Carbon. — As the name implies, amor-
phous carbon is without crystalline structure. Its density
is low; it is soft, and it does not conduct an electric current
very well. These properties cannot be specified definitely
because they vary according to the temperature to which
this element has been subjected. Amorphous carbon bums
with comparative ease.

226. Graphite. — Graphite is a crystalline form of carbon.
There are a number of natural deposits of graphite, the
best coming from Ceylon and Siberia. Graphite is now
produced artificially by heating coke with a small percent-
age of iron in an electric furnace. The coke is converted
into graphite, and the iron is vaporized. The crystals of
graphite appear as minute scales or plates, which present

234 CARBON.

a very good sliding surface. To this property is due the
use of graphite as a lubricant. It differs from other forms
of carbon in being very soft, in conducting the electric
current readily, and in having a very high kindling tem-
perature. Like all forms of carbon, it has a very high
melting-point. The refractory nature of graphite permits
its use in crucibles for melting metals. Its resistance to
heat and its conducting power cause it to be extensively
used for the melting-pots and electrodes of electric fur-
naces. Graphite is also used for making lead pencils.
The varying degrees of hardness in the pencil are secured
by mixtures of graphite and clay.

227. Diamond. — Diamond, the third form, is also a crys-
tallized variety of carbon. It is found as octahedral crys-
tals, sometimes colorless and transparent, sometimes tinted
or black. The diamond differs from the other allotropic
forms in being extremely hard. It is the hardest sub-
stance found in nature. Like graphite, it burns only
when heated to a high temperature. The use of dia-
monds as gems is due to their great rarity, and to the
fact that light in passing through them is highly re-
fracted so that a sparkling play of color results. This
effect is heightened by cutting the surface of the stone
into numerous facets. The value of a diamond depends
largely on its color and its brilliancy. Black stones are
used only for cutting instruments for use on hard mate-
rials.

Microscopic diamonds have been made artificially. Car-
bon dissolves in melted iron. If such a solution is heated
to a very high temperature, in an electric furnace, and
then suddenly cooled in water or melted lead, the outer
part of the iron solidifies first, producing an enormous

CHEMICAL PROPERTIES. 235

pressure on the still liquid interior part. The carbon, in
separating from the coeling solution under this great pres-
sure, takes the form of minute diamond crystals.

228. Chemioal Properties — The physical properties of
carbon have been described in discussing the allotropic
forms of the element.

None of the forms of carbon reacts with either acids or
bases. They are all insoluble in ordinary solvents. Melted
iron dissolves carbon to a certain extent.

Carbon has a great tendency to unite with oxygen at
high temperatures. It burns in the gas and also unites
with the oxygen that is held in combination with other
elements. On this account carbon makes an excellent re-
ducing agent, and since it is cheap, it is the reducing agent
most extensively used. The reaction of carbon with cop-
per oxide illustrates this property. When the mixture is
heated, metallic copper is produced ; the other product is
carbon dioxide or carbon monoxide, according to the relsr
tive amounts of the substances used:

2CuO + C-^2Cu + COj
CuO + C -^ Cu + CO

Enormous quantities of coke are used in the reduction of
metals, particularly iron, from their ores.

If a mixture of copper oxide and charcoal is heated in
a test-tube (Fig. 75 a) and the gas led into lime water
(Fig. 75 6), a white precipitate of calcium carbonate indi-
cates the presence of carbon dioxide.

Carbon unites directly with only a few elements. If
heated in the presence of sulphur vapor, carbon disulphide
is produced:

C + 2S-^CS2

236

CARBON.

In spite of the fact that carbon forms so many different
compounds with hydrogen, it unites directly with this ele-
ment only at the tempera-
ture of an electric arc, and
then very slowly. By us-
ing the high temperature
of an electric furnace, car-
bon can be made to unite
with many metals and some
non-metals. The com-
pounds made in this man-
ner are called carbides.

229. Carbides and the
Electric Furnace. — The in-
tense heat of an electric
furnace makes carbon react
with lime (calcium oxide),
and calcium carbide is
Fig- 75. formed:

CaO -h 3C — ^ CaCa -h CO

Calcium carbide is used to prepare acetylene^ a gas which
makes, an excellent illuminant. The gas is generated by
the action of calcium carbide with water at ordinary tem-
peratures:

CaCj + 2H2O -^ C2H2 + Ca(0H)2

Acetylene is used in bicycle and automobile lamps and
in places where the gas supply of a city system is not
available.

The electric furnace (Fig. 76) is a device by which a
very high temperature is obtained by using a large arc
between carbon terminals. The arc consists of carbon

CARBIDES AND THE ELECTRIC FURNACE. 237

vapor, which conducts the current over the gap between
the carbons and the resistance is so great that much
heat is developed. A temperature of about 3000° C. is
obtained. The reactions that occur in the electric furnace
are due to the high temperatiire. The electric current
does nt>t directly play any part in the reactions.

Fig. 76. Electric furnace.

Carborundum, silicon carbide, is made by heating sand,
silicon dioxide, and carbon in an electric furnace (Fig. 87) :

SiOj + 3C-^SiC + 2C0

230. 11868. — The uses of carbon have been pointed out
in the discussion of the properties of the element. Coal is
the principal source of artificial heat. Anthracite is, in
general, the most desirable kind of coal, because it burns
with practically no flame and without the formation of
soot. It burns, however, more slowly and gives a more
uniform heat than does bituminous coal. Coke burns like
anthracite coal and is sometimes used in its place. In
warm countries charcoal is often the only fuel used.

238 CARBON.

Many metals are separated from their ores by the aid of
coke as the reducing agent. The making of coke from
bituminous coal for this purpose is an important industry.
Iron is obtained in this way, usually by the reduction of
hematite, ferric oxide. Enough coke is used to give car-
bon monoxide as the reduction product, and this gas is used
in other parts of the plant as fuel.

Charcoal and boneblack are both used to remove gases
and coloring-matter from solutions. Impure water is
sometimes passed through charcoal filters for purification.
This method, while it may suffice for the removal of im-
purities which can be seen or smelled, is not to be relied
upon for the removal of disease germs.

231. Flame of Candle. — We have seen numerous cases
of combustion or burijing. When the fuel is a solid, par-
ticles of the fuel or of the ash are often heated to incan-
descence and glow. When a stream of gas bums, a flame
is produced. This phenomenon can best be studied in
the flame of a common candle.

When a light is applied to the tip of a candle wick, a
flame appears and slowly spreads down the wick toward
the candle and upward about twice the height of the wick.
In a few moments the flame becomes constant in size and
position (Fig. 77). Starting about one-eighth of an inch
from the candle, the flame begins to be visible, rounded at
the base and tapering to a sharp point above. It entirely
encloses the wick except at the base and at the extreme
tip.

If we examine the flame closely, we see four parts. The
first is a greenish blue portion of the base^ shaped much like
the cup of an acorn. This follows the wick downward a
short distance, and its upward extension is hidden by the

FLAME OF CANDLE. 239

glare of light within. Second, above and within the first
green portion we find, immediately surrounding the wick, a
region dark, compared with the more brilliant part above,
but in reality transparent and colorless, as is shown by
placing an object behind it. This region tapers to a cone
above and is termed the non-luminous cone. Covering the
cone just mentioned is a bright cap known as the luminous
cone. This is the third and most conspicuous region of
the flame. Outside the luminous cone we can find the
fourth region, which is the faint and
scarcely visible blue mantle.

The candle is composed of a mixture
of paraffin and stearic acid, both of
which contain carbon and hydrogen.
The upper portion of the candle, heated
by the flame above, melts, so that the
top becomes a cup holding a small por-
tion of the melted fuel. This melted
portion is drawn up the wick by capil-
larity, and vaporizing, forms the non-
luminous cone immediately surrounding
the wick. If we insert a tube into this
portion of the flame, we can lead out a

quantity of vapor which condenses to a pig. 77.

solid similar to that composing the
candle. If we extend a thin piece of wire across the
flame, we find that it is not very warm near the wick ;
in fact the head of a match can be placed there without
igniting. There is no combustion in this non-luminous
cone since there is neither air nor sufficient heat.

As the vapor is carried upward by the draft it becomes
mixed with air and combustion ensues, and this region of
active burning is brilliant and hot, but not transparent.

240 CARBON.

If we place a piece of cold porcelain in this portion of the
flame, or lead off some of the materials through a tube,
we find a considerable portion of lampblack (carbon) and
moisture. The carbon while in the luminous cone is
heated to incandescence.

Since the combustible vapors are slowly mixed with air,
the materials may move considerable distances from the
vvick before they meet sufficient oxygen to burn; hence
when a large wick is used, the bulk of vapor makes an ex-
tensive region of flame. Indeed, the upper extremity is
often so far removed that its temperature falls below the
kindling temperature of the combustible materials. These
escape unburned or the flame smokes.

In the outer portion of the flame the conditions are re-
versed, in that the air is in excess. Here the combustion
is complete, but usually only a small amount of combusti-
ble material reaches this pale outer portion. This material,
if drawn out by a tube, is found to be principally air mixed
with products of combustion. The pale green cup below is
similarly constituted ; but from its position below the prin-
cipal region of combustion and its nearness to the cool wick
and candle, it is not so warm as the blue mantle above.

The wick is made of cotton and is proportioned to the
quantity of material to be burned. The cool stream of
liquid drawn up the wick protects it from burning, and,
at the same time, prevents the flame extending down to the
candle. As the candle material is vaporized, the wick chars
somewhat but does not burn, since there is no oxygen in
its immediate vicinity. In braiding the wick, one of the
threads is drawn tighter than the others, so that the wick,
when free from the candle stuff, is drawn over to one side
of the flame. This braiding of the wick brings the tip
into the outer zone of combustion, where it bums so far as

GAS FLAMES. 241

air is available, rotating as the candle shortens and prac-
tically maintaining a uniform length. This is why modern
candles do not need to be trimmed.

232. (Jas Flames. — In the ordinary gas flame the follow-
ing four regions are clearly marked: the region of fuel,
or non -luminous cone; the luminous cone, or region of
luminosity; outside and lower regions of combustion with
an excess of air. In a Bunsen burner we can adjust the
proportions of gas and air so that there shall be no excess
of either. Then the flame will appear to have two regions,
a lower cone where no combustion is apparent and the cap
above. The combustion is more rapid in that the fuel
does not get far from the burner before burning. There-
fore, the flame is smaller and for this reason hotter. It is
not luminous, owing to the rapid and complete burning
and to the dilution of the materials by the nitrogen.

In the ordinary flame, we find the position of maximum
temperature somewhat above and outside the light-giving
materials. A very large part of the energy passes off as
heat and only a very small per cent is converted into
light. If the solids giving light could be placed above or
in the region of maximum temperature, they would evi-
dently be hotter and give more light. Auer von Welsbach
accomplished this by placing a mantle composed of a net
of infusible and incombustible materials in the outer por-
tion of the flame of a Bunsen burner. The materials are
heated to incandescence and produce the brilliant light
given by the Welsbach burner.

Stimmary.
Carbon is a non-metallic element, characterized by the enormous
number of compounds it forms with the non-metallic elements,
espedallj hydrogen, oxygen, and nitrogen.

242 CARBON.

Carbon is an inert dement at ordinary temperatures; at higher
temperatures it combines readily with oxygen and with a few other
elements ; at the temperat we of the electric furnace it reacts with
lime, forming calcium carbide.

Amorphous carbon, graphite, and diamond are the three aUo-
tropic forms of carbon. These allotropic forms differ widely in
physical properties, but all yield the same product when burned in
sufficient oxygen.

Carbon is found combined in the tissues of every living thing,
in the carbon dioxide of the atmosphere and in metallic carbon-
ates. Uncombined it occurs as coal, graphite, and diamond.

Anthracite coal is nearly all carbon ; bituminous coal is about
two-thirds carbon and one-third hydrocarbons; cannel coal and
lignite are poor in uncombined carbon.

The uses of the forms of carbon are :
Coal, fuel and manufacture of illuminating gas;
Lampblack, paint and printers' ink;
Wood charcoal, fuel and filtering;

BonebUick, filtering and decolorizer of sugar and of oils;
Coke, fuel, ore reducer, and manufacture of water-gas;
Graphite, lead pencils, lubricant, crucibles, and electrodes for high

temperatures ;
Diamond, gem and abrasive.

The four portions of a candle flame are the greenish blue r^on
at base, the non-luminous cone, the luminous cone, and the blue
mantle. The ordinary gas flame has four similar portions, while
that of the Bunsen burner appears to have but two — the lower
cone and the surrounding conical cap above.

Exercises.

1. Wohler is said to have broken down the barrier between
organic and inorganic chemistry when, in 1828, he prepared
from chemicals, urea, a waste product of the body. Why ?

2. Why are the decomposition products of animal and vege-
table matter so similar ?

EXERCISES. 243

3. Tell how the carbon in your muscular tissue is indirectly
derived from the air.

4. Why does soft coal make such a smoky fire ?

5. Why is anthracite coal preferred for household use ?

6. Explain why fence posts are sometimes charred at the
end before being placed in the ground.

7. What properties of lampblack make it suitable for
printers' ink ?

a What kind of carbon is used in batteries ?

9. Write the equation for the chemical reaction when an
excess of carbon is heated with zinc oxide, ZnO.

10. What three conditions are necessary for ordinary burn-
ing?

11. Why does water put out a fire ?

12. If a taper is held over a gas burner, a flame may be
maintained several inches above the burner without " striking
back " to the burner. Explain.

13. Why does not the flame of a Bunsen burner extend
down the tube to the base ?

14. In working with compressed air it was found that com-
bustible materials burn with more flame than usual. Explain.

15. Why does the light increase with the pressure under the
condition given in question 14 ?

16. Why is the flame of a Bunsen burner hotter with the
holes open than when they are closed ? Is the total heat in-
creased ?

CHAPTER XXIV.
OXIDES OF OABBON.

233. Natural FormatioiL — There are three very impor-
tant sources of carbon dioxide : the decay of vegetable and
animal matter; the oxidation constantly going on in
animals, and to a much smaller extent in plants ; the com-
bustion of all ordinary fuels, such as wood, coal, and gas.
These fuels consist largely of carbon; this, on burning,
combines with oxygen, forming carbon dioxide :

C + O2— ^CO^

Hence the gas is always present in the air, usually to the
amount of four parts in ten thousand (§ 156). This per-
centage would be much higher were it not for the fact
that plants are constantly taking it in, building the carbon
into their tissues and returning the oxygen to the air.
Natural waters also very commonly contain it; in some
spring waters the gas is dissolved in such quantities that
they are effervescent, that is, they give off gas in bubbles
unless kept in tightly closed vessels. It is also given off
by volcanoes and from other subterranean sources, and
from fermenting liquids. It is found in mines, where it
is known as ch>he damp.

234. Preparation. — Carbon dioxide is most conveniently
prepared in a pure state by the action of an acid on a
carbonate (Fig. 78). Calcium carbonate is generally used
on account of its abundance, as marble, limestone, and chalk

244

PHYSICAL PROPERTIES.

245

consist almost entirely of this compound. Hydrochloric
acid is commonly used in the laboratory on account of the
solubility of calcium chloride which it forms. The equa-
tion representing the action is :

CaCOg + 2 HCl — ^ CaCLj + H^O + CO,

To remove any hydrochloric acid that may be carried over
the gas is allowed to bubble through water. The gas is
collected by downward displacement of air or over water.

235. Physical Properties. — Carbon dioxide is a colorless
gas with little taste or odor. It is about one and one-half
times as dense as air, so that
it can be poured from one
vessel to another like water.
In wells, caves, or mines,
carbon dioxide is often set
free from soil or water, and
if there is poor ventilation,
it will accumulate.

It is soluble in water,
which dissolves about its
own volume at ordinary tem-
peratures. Under increased
pressure, water dissolves
more of the gas; if the pressure is removed, the gas is
slowly given off. Soda water is water into which carbon
dioxide has been forced under pressure (60 to 150 lb.)
and effervesces when drawn. Liquids bottled during fer-
mentation, as champagne and kumiss, effervesce for a
similar reason. The liquefied gas is also used like am-
monia in refrigerating and ice-making processes, particu-
larly on ocean vessels where ammonia might be dangerous.

Fig. 78. Preparation of carbon

dioxide.

a, generator ; b, collecting bottle.

246

OXIDES OF CARBON.

236. Chemical Properties. — Carbon dioxide is chemically
inactive toward most substances, being the product of the
complete oxidation of carbon. It does not burn nor sup-
port combustion (Fig. 79). Burning potassium, sodium,
and magnesium, however, will decompose it, uniting with
the oxygen and setting the carbon free. The presence of
a small percentage of carbon dioxide in the air extinguishes
combustibles and causes death by suffocation. A candle
will go out in air containing 4 % of the gas, and when a
little larger amount is present, life is impossible.

Fig. 79. Candles successively extinguished by pouring carbon dioxide
into V-shaped trough.

The most important reaction of carbon dioxide is its
decomposition, under the influence of light and chlorophyl,
whereby oxygen is set free :

6 CO2 + 5 H2O — ^ 6 O2 + CgHioOg (starch)

HARD WATERS. 247

The water in this reaction comes up from the roots of the
plant and the carbon dioxide comes from the air.

The solution of carbon dioxide has a slightly acid reac-
tion and forms carbonates with bases. Carbonic acid, like
ammonium hydroxide, has never been isolated, but the
formation of carbonates indicates the presence of hydrogen
ions and COg ions. We may write the equation:

H2O + CO2— ^HjCOg

Carbon dioxide, then, passed into a solution of a base
produces the corresponding carbonate. As calcium car-
bonate is insoluble, calcium hydroxide (lime-water) is used
as a test for the presence of carbon dioxide. When carbon
dioxide is first passed into lime-water, it becomes milky,
and on standing a precipitate of calcium carbonate sepa-
rates *

HgCOg + Ca(0H)2 —^ CaCOg + 2 K^O

If the passage of carbon dioxide is continued, the precipi-
tate dissolves. This is because it has been converted into
calcium bicarbonate, which is soluble in water:

HjCOg-l- CaCOg — ^ CaHaCCOg)^

237. Hard Waters. — The solubility of calcium carbonate
in water containing an excess of carbon dioxide explains
the formation of one kind of hard water. ' In regions
where limestone is abundant, the surface water, becoming
charged with carbon dioxide from decay going on in the
soil, dissolves some of the limestone, converting it into
calcium bicarbonate. If an attempt is made to use soap
with such water, it is found difl&cult to obtain a lather, as
the dissolved limestone reacts with the soap to form an
insoluble compound. By boiling hard water of this kind,

248

OXIDES OF CARBON.

part of the carbon dioxide is driven off and the calcium
bicarbonate is converted into calcium carbonate:

CaH2(C03)a —>- CaCOg + COg + HOj

The carbon dioxide escapes and the calcium carbonate is
precipitated. If the rocks contain magnesium carbonate,
this may be converted into the bicarbonate and dissolved
in the same way as with the calcium compound, and the
water may be softened by the same means. Water con-
taining such dissolved bicarbonates is called water of tem-
porary hardness. Water of permanent hardness contains
sulphates of calcium and magnesium and cannot be softened
by boiling.

TTses. — Many uses of carbon dioxide have been
alluded to. As an easily soluble, non-poisonous gas, it is
extensively employed in the manufacture of
beverages. Seltzer, vichy, and other min-
eral waters are produced artificially by
charging solutions whose composition is
similar to that of the original spring. Fire
extinguishers (Fig. 80) often contain a
device for generating carbon dioxide
rapidly, as by mixing sulphuric acid (6) or
alum solution with sodium carbonate (a).
In chemical engines the pressure of the gas
itself is sometimes used to throw a stream
of water.

The raising of bread depends on the expansion of bub-
bles of carbon dioxide by heat, the carbon dioxide being
generated by yeast or by the reaction of sodium bicarbon-
ate (baking soda) and a material of acid reaction; for ex-
ample, acid potassium tartrate (cream of tartar).

Fig. 80.

CARBON MONOXIDE.

249

NaHCOg -I- HKC^H^Oe —>• NaKC^H^O^ + HjO + CO,

The compound, sodium potassium tartrate, NaKC^H^O^,
is known as Rochelle salt.

CARBON MONOXIDE.

239. Preparation. — As carbon monoxide differs from
carbon dioxide in containing one atom of oxygen instead
of two, it can be prepared by the reduction of carbon
dioxide. This is best accomplished by passing a current
of carbon dioxide over carbon heated to redness in a hard
glass or iron tube (Fig. 81). The carbon acts as a reduc-
ing agent:

COj + C— ^2C0

Fig. 81. Preparation of carbon monoxide.

a, carbon dioxide generator ; b, combustion furnace ; c, tube containing red-hot carbon ;
d, collecting bottle.

This action takes place in a coal fire where the carbon
dioxide, formed near the bottom, is reduced by the hot coal
above to carbon monoxide, which may be seen burning at
the top of the fire.

250

OXIDES OF CARBON.

In the laboratory carbon monoxide is commonly made
by the decomposition of oxalic acid (Fig. 82). This,
heated with concentrated sulphuric acid, decomposes:

C^HgO^ —>- COg + HjO + CO

The purpose of the sulphuric acid is to extract the water
formed in the decomposition of the oxalic acid. The car-
bon dioxide is removed by passing the gases through a
concentrated potassium hydroxide solution.

Fig. 82. Preparation of carbon monoxide.

a, generator containing oxalic and concentrated sulphuric acid ; b, bottles containing con-
centrated solution of potassium hydroxide ; c, collecting bottle.

240. Physical Properties. — Carbon monoxide is a color-
less gas with a slight odor. It is slightly lighter than air
and is nearly insoluble in water. It is extremely poison-
ous, a very small percentage in the air causing headache,
and a larger amount death.

CHEMICAL PROPEETIEa. 251

Chemical Properties. — Under ordinary conditions the
monoxide burns in air with a blue flame. If both the car-
bon monoxide and the air are dry, however, combustion
does not take place. Carbon monoxide may act as a
powerful reducing agent. The reduction of oxides of
iron in the blast furnace is largely due to it. The heated
carbon monoxide, formed by a blast of hot air forced
through coke, diffuses through the furnace, and reacts
with the iron oxide, producing carbon dioxide and me-
tallic iron, which melts and runs to the bottom of the
furnace.

Carbon monoxide prevents the oxygen of the air from
uniting with the haemoglobin of the red blood corpuscles,
entering itself into combination with the haemoglobin.
It is because of this action that a comparatively small
amount of the gas will cause death. The compound
formed in the corpuscles has a brilliant red color, and
is so stable that it can be detected in a body years after
death. Carbon monoxide is the chief poison in illuminat-
ing gas and the gas escaping from stoves.

241. TTses. — The chief use of carbon monoxide is in
water ga%. This is produced by the reduction of steam by
heated coal:

C + H2O — ^ CO -h Hg

The process (Fig. 83) consists essentially in alternately
forcing air and steam through incandescent anthracite coal
or coke contained in an iron cylinder {^producer). Air is
forced through until combustion is vigorous. Then the
air is shut off and the outlet of the chimney closed;
superheated steam is next forced in. This reacts as we
have seen above and the carbon monoxide and the hydro-
gen pass together into another cylinder (carburetter)^

252

OXIDES OF CARBON.

where the gas is mixed with the vapor of hydrocarbons.
This mixture is thoroughly blended in the superheater
and then goes through various purifiers {scrubhera) to the
storage tank or holder. The object of the addition of the
hydrocarbons is to give the mixture of hydrogen and
carbon monoxide substances that will produce light in
burning, as the original gases burn with non-luminous
flames.

Carbon monoxide is the principal heat-giving constitu-
ent of another gas known as producer ga%. This is made
by forcing air over a bed of incandescent coke in a furnace
of special form. Although producer gas contains about
63% of nitrogen, it is extensively used as a fuel on ac-
count of its cheapness, cleanliness, and the even tempera-
ture of its burning.

Eloldflf

riA\^

Fig. 83. Manufacture of water gas (diagrammatic).

SUMMARY. 263

Summary.

Carbon dioxide is formed in the vital processes of plants and
animals and in ordinary combustion.

It is prepared commercially by the reaction of acids with car-
bonates.

Carbon dioxide is colorless, slightly soluble in water, and suffo-
cating, but not poisonous.

One liter (standard conditions) weighs 1.98 grams.

It reacts with soluble bases, forming carbonates. Small quanti-
ties render Hme-water turbid, excess of the dioxide causes the pre-
cipitate to dissolve. Such water is temporarily hard; boiling expels
the excess of carbon dioxide, and the calcium carbonate again
separates.

Carbon dioxide is used in charging beverages, in fire extinguish-
ers, and in ice machines.

Carbon monoxide is formed by :

(1) the incomplete combustion of carbon;

(2) the reduction of the dioxide;

(3) the reaction of steam and red-hot coal.

Carbon monoxide is lighter than air, has a slight smell, and is
very poisonous. One Kter (standard conditions) weighs 1.26
grams.

It burns with a pale blue flame, forming the dioxide.

Carbon monoxide is need as a fuel in producer gas, and is one
of the constituents of water gas.

Exercises.

1. How can it be shown that there is carbon dioxide in the
air?

2. Why is hydrochloric acid preferred to sulphuric acid in
preparing large quantities of carbon dioxide from calcium car-
bonate ?

3t Calculate how many liters of carbon dioxide, at standard

254 OXIDES OF CARBON.

conditions, can be obtained by treating 45 grams of pure marble
with acid.

4. Why does soda water effervesce when taken from the
tank?

5. Explain how a fire extinguisher puts out a small fire.

6. Why is rain-water preferred for washing purposes in
limestone regions ?

7. It has been calculated that an average man exhales
464.5 liters (standard conditions) of carbon dioxide in a day.
Calculate how many grams of starch a plant could make from
this.

8. How could you prove that there is carbon in alcohol ?

9. How could you distinguish carbon dioxide from the
monoxide.

10. Explain why cases of asphyxiation occur after coal stoves
have been filled and left for the night.

11. Calculate the weight of steam that could be decom-
posed by a ton of incandescent coke containing 90 % carbon.

12. Write the equation for the reaction you would expect if
carbon monoxide was passed over hot copper oxide.

13. H6w would you determine whether a gas was hydrogen
or carbon monoxide ?

14. Why i^ it that such a large percentage of the cases of
asphyxiation from water gas result fatally ?

CHAPTER XXV. ^

SILICON AND BOSON.

242. Silicon. — Silicon, next to pxygen, is the most abun-
dant element in the earth's crust. Nearly all the common
rocks are silicon compounds. Silicon, as an element, is of
little practical importance. Like carbon, it exists in three
allotropic forms: a brown amorphous powder; a dull gray
crystalline form, like graphite; and in octahedral crystals.

243. Varieties of Silicon Dioxide. — Silicon dioxide, or
silica^ SiOg, is the most common compound of silicon. It is
found in many varieties, which differ in color and struc-
ture owing to minute quantities of impurities, and to condi-
tions under which it is formed. Quartz, the most common
form, crystallizes in hexagonal prisms, surmounted by a
pyramid (Figs. 22 and 84). Clear crystalline varieties are
known as rock crystal; purple varieties as amethyst.
Rose quartz, milky quartz, and smoky quartz, or cairn-
gorm stone, are other colored varieties of silicon dioxide,
which, indeed, is found in all shades and tints. Chal-
cedony is cryptocrystalline and waxlike. Onyx and agate
are varieties of chalcedony. Jasper and flint are other
forms of silica. Opal is a hydrated form.

The shells of diatoms and many other microscopic or-
ganisms are siliceous, and deposits of these comprise the
infusorial or diatomaceous earth. Sand is water-worn
silicon dioxide, and sandstone consists of particles of sand
cemented together.

255

256

SILICON AND BORON.

244. Properties and Uses of Silicon Dioxide. — Silica is
harder than glass; it is insoluble in ordinary reagents,
but will dissolve in melted alkalies. Melted in the oxy-
hydrogen flame, quartz can be drawn into delicate elastic
threads, which are used in scientific instruments.

Fig. 84. Mass of quartz crystals.

White sand, which is nearly pure silica, is used in mak-
ing glass and porcelain, dommon sand is discolored by
impurities, and if the particles are sufficiently irregular
and angular it can be used in sandpaper and mortar.
Sandstone is used for building; hard varieties are used
for grindstones and millstones. Ground glass is glass
roughened by blowing sand against it by means of a blast
of air. Many clear varieties of quartz are cut and polished
for jewelry, as amethyst, agate, carnelian, false topaz, and
imitation diamond. The clear rock crystal is cut for
lenses. Petrified wood has been formed by the gradual

SILICATE 8. 257

replacement of the woody fibre by silica, preserving the
woody structure. Cut and polished petrified wood is
used as an ornamental stone. The fine varieties of infu-
sorial earth are used as abrasives in polishing-powders,
and are also used in cements, in refractory fire-brick, and
as an absorbent in dynamite.

Silica is found in the ashes of most plants, especially
in the ashes of stalks and stems. It imparts firmness to
the stems and to the resistant exterior coating of straws
and bamboo. The quills of feathers, claws, and nails con-
tain considerable silica.

245. Silicates. — Silicon dioxide, like carbon dioxide, re-
acts with alkalies to form salts called silicates:

SiOa + 2 KOH — ^ KgSiOg + HgO

Sodium and potassium silicates are soluble in water; nearly
all the other silicates are insoluble, stable compounds,
which comprise the larger part of the earth. Such min-
erals as felspar, mica, hornblende, clay, and talc are
silicates, various mixtures of which comprise the common
rocks, as granite, gneiss, and slate.

Sodium and potassium silicates are made by melting the
metallic hydroxides, or carbonates, with silicon dioxide.
The thick water solution is used in filling soaps, in mak-
ing artificial stone and cement, in wall coloring, calico
printing, and fireproofing wood and textiles.

246. Glass. — Glass is a mixture of silicates. Common,
crown, or window-glass, consists of silicates of sodium
and calcium; Bohemian glass, of potassium and calcium
silicates; flint glass contains silicates of lead and potas-
sium.

Glass is made by melting together sand, an alkali, and

258

SILICON AND BORON.

Fig. 85. Glass furnace :
a, fire-boxes ; b, melting-pots.

calcium carbonate in pots of fire-clay (Fig. 85). The
alkali may be sodium or potassium carbonate or a mix-
ture of these. An oxidizing agent, as potassium nitrate or
manganese dioxide, may be added to remove the green

color due to iron com-
pounds. The mixture is
heated to a high tempera-
ture and thoroughly melted,
the gases given off aid the
mixing, and any infusible
impurities coming to the top
are skimmed off. When the
mass is cooled to a pasty
condition, it may be blown
or moulded.
Window-glass is made by the workman taking a mass of
the molten glass on the end of a long iron blowpipe, and
blowing it into a large bubble. This is drawn out into a
cylinder by swinging it and rolling it on a plate. The ends
of the cylinder are cut off, a cut is made lengthwise, and
the glass is spread out flat. Plate glass is made by pour-
ing the molten glass on a bronze table, rolling it with a
hot iron cylinder (Fig. 86), and finally polishing it. Cut
glass is flint glass moulded to the desired shape; the design
is cut by a wheel, and the glass polished with rouge or
putty powder.

Cheap glass dishes and similar objects are^%iade by
pressing the plastic glass in a die. Bottles are blown in a
mould. If the glass is cooled rapidly, it is hard, brittle,
and liable to break under a shock; to overcome this it is
annealed. The glass is passed slowly through a long fur-
nace from the hot to the cooler end, so that the tempera-
ture is very gradually lowered.

GLA88. 269

Crown fflass is a colorless window-glass used for convex
lenses. Bohemian glass is harder and less fusible, and is
used for chemical apparatus. Flint glass is brilliant,
heavy, and soft, and is used for concave lenses, lamp
chimneys, and globes, and for imitation gems, as paste
diamonds.

Glass is colored by dissolving various substances in the
melted mass. The green color of common glass is due to

By courtesy of TAe Scientific American.

Fig. 86. Rolling out plate-glass.

iron compounds in the sand and limestone; chromium
compounds give a rich green. Compounds of copper and
cobalt give blue color; manganese, pink to violet; man-
ganese with iron, yellow to brown; silver, yellow; gold,
ruby red; calcium fluoride, white and translucent.

All the silicon compounds mentioned contain oxygen,
and while silicon forms many other and more simple com-
pounds, few are of practical importance.

260

SILICON AND BORON.

247. Silicon Carbide, or Carbornndiim. — Silicon carbide,
carborundum, is a crystallized solid varying in color, and
often brilliant and iridescent. It is extremely hard, and

is used as a substi-
tute for emery for
grinding and polish-
ing in wheels, hones,
and carborundum
cloth (Fig. 87).

Carborundum is
made in an oblong
electric furnace, at
the ends of which
are metal plates to
which are attached
the heavy carbon electrodes projecting into the furnace.
The electric connection between the electrodes is through
a mass of granulated coke. Sand is mixed with coke, a
little salt, and sawdust. The salt is used to aid fusion, and
the sawdust to make the mass porous. This mixture is
piled around the central core of coke and held in place by-
side walls of loosely piled bricks (Fig. 88). The action
in the furnace is not electrolytic, but is due to the heat
generated through the resistance of the coke to the cur-
rent. The carbon reacts with the melted sand to form
carbon monoxide and carborundum:

Fig. 87. Carborundum products.

SiOg + 3 C

SiC + 2 CO

The action continues for about eight hours. When the
furnace has cooled, the sides are torn down and the car-
borundum removed (Fig. 89). The best crystals are
found around the central core. The crystals are crushed,

BORON AND BORIC ACID. 261

washed with sulphuric acid, dried, and graded according
to size.

Silicon hydride^ SiH^, is made by blowing coal-gas
through an electric furnace charged with silica. The
hydride is a strong reducing agent, and burns, forming
water and silicon dioxide.

Silieon fluoride^ SiF^, is a colorless gas formed in the
reaction of hydrofluoric acid, HF, with silica, SiOg, or
glass. It decomposes in water, forming hydrofluosilicic
acid, H2SiFg, and silicic acid, H^SiOg.

Fig. 88. Carborundum furnace, showing charge.

248. Boron and Boric Acid. — The element boron is of
little importance. It is a brown powder, soluble in melted
aluminum, and infusible at the temperature of the electric
arc. Its important compounds are boric acid and borax.

Boric acid, HgBOg, occurs in minute quantities in vapors
arising from the earth in the volcanic regions of Tuscany,

262

SILICON AND BORON.

in Italy. Although the amount of boric acid contained
in the steam is small, it is retained by constructing shal-
low reservoirs surrounding the vents from which the steam
escapes. The steam condensed in these reservoirs con-
tains a considerable amount of boric acid. The reservoirs
are arranged one below another on the hillside. After the
steam has been allowed to condense in each for twenty-
four hours, the solution is run into the one below and the
process repeated. After the boric acid solution has become
concentrated, it is evaporated in shallow pans, heated by

Fig. 89. Carborundum furnace after burning (sides removed).

steam from other jets. The boric acid separates and is
purified by recrystallization. Most of the boric acid used
in the United States is made by the reaction of borax with
an acid.

Boric acid is obtained in fine crystalline scales. It is
a weak acid, sparingly soluble in water. It is used as an
antiseptic and as a preservative.

249. Borax. — When boric acid is heated, it loses water,
forming successively metaboric acid, HgBgO^, and pyroboric

BORAX. 268

acid, HgB^Oy. Sodium pyroborate, NajB^O^, is the familiar
compound, borax. There are large deposits of borax and
calcium borate in California, which supply this country.
Borax is also obtained from the boric acid of Tuscany.

Borax is obtained from solutions in large crystals, con-
taining either 5 or 10 molecules of water of crystallization,
according to the temperature at which they are deposited.
Ordinary borax has the composition NagB^O^ • 10 HgO.
When heated, crystallized borax swells during the evapora-
tion of the water of crystallization and then melts to a clear
glassy mass. Fused borax dissolves metallic oxides, and
these often impart to the glassy mass a color Borax
characteristic of the metal. Thus, cobalt com- ^»d *••*«•
pounds give a blue color, and manganese compounds a
violet color, to a drop of fused borax heated in the oxidiz-
ing flame. A solution of borax has a feeble alkaline reac-
tion.

Borax is used in soldering and welding to dissolve the
metallic oxides formed. It is employed in large quantities
as an antiseptic and preservative.

Summary.

Silicon is a very abundant element of little practical importance.

Its most common compound is silicon dioxide, occurring as
quartz and sand and as a constituent of many rocks.

Silica is very hard and fuses only at high temperatures. It is
used in making glass, mortar, and polishing-powders.

Sodium and potassium silicates are soluble.

GUiss is a mixture of silicates. Three varieties are crown, flint,
and Bohemian glass. The materials are melted together and blown
or moulded into shape while plastic. Colors may be added to it
while in a melted state.

Silicon carbide is made by heating coke and sand in an electric
furnace. It is used as an abrasive.

264 SILICON AND BORON.

Silicon fluoride is produced by the action of hydrofluoric add on
glass.

B(nic acid is found in nature, but much of it is made from borax.

Borax occurs in large deposits in California.

It gives characteristic reactions with metallic oxides used as bead
tests.

Borax is used in soldering and welding, and as a preserTative.

Bzercises.

1. Why is not silicon found uncombined in nature ?

2. Give the most important uses of silicon dioxide.

3. Give the composition and uses of the chief varieties of
glass.

4. Why is sand the main final product of long-continued
disintegration of rock materials by water ?

5. What is water-glass ? How is it made, and what are its
uses?

6. Describe the manufacture of carborundum, and give its
uses.

7. Calculate the percentage of water of crystallization in
borax.

CHAPTER XXVI.

OALOIUH AND ITS OOHFOUimS.

CALCIUM.

250. — Although metallic calcium has been known for
many years, it is only recently that it has been made in
any quantity. Sir Humphry Davy was the first one to
see the metal, but failed to get enough to determine its
properties. The credit for the successful isolation of the
metal belongs to Dr. Robert Hare, a scientist of Philadel-
phia. His electrolytic method, with some modifications,
is the one now used to obtain the metal.

251. Preparation. — Metallic calcium is prepared by pass-
ing an electric current through fused calcium chloride con-
tained in a crucible of graphite, which acts as the anode
(Fig. 90). At first the cathode is an iron rod (5),
capable of being raised by a screw mechanism (^). The
calcium deposits on the end of the iron rod, solidifies,
and grows downward as an irregular cylinder. This rod
of calcium becomes the cathode ((7) as the iron rod is
gradually raised out of the molten chloride by the screw
mechanism. A coating of calcium chloride protects the
calcium from oxidation as it emerges from the molten
bath. By using suitable screw mechanism, pieces of cal-
cium several feet long can be obtained. The electric cur-
rent, by its passage, keeps the calcium chloride molten in
the graphite crucible, except at the bottom (2>), where it

265

266

CALCIUM AND ITS COMPOUNDS.

is kept cold, and solidified by water running through a
copper coil (JEE^, Most of the chloride adhering to

the sticks of calcium is
removed by hammer-
ing; the remainder is
dissolved off by alcohol.

252. Physical Proper^
ties. — Pure calcium is a
silver-white metal of
brilliant lustre, and is a
little lighter than mag-
nesium. It is harder
than lead or tin, almost
as hard as aluminum,
but softer than zinc.
At 300° to 400° C. it is
as soft as lead and can
Pj 90. easily be rolled or ham-

mered. Only four
metals surpass calcium as conductors of electricity.

253. Chemical Properties. — Calcium is a much less act-
ive metal than has been imagined. It keeps its lustre in
dry air and can be preserved without difficulty in a stop-
pered bottle. In moist air its surface becomes dulled.
The solid metal does not ignite when heated red-hot, but
when sent whizzing through the air against a brick wall
bursts into a brilliant white flame and burns violently, like
magnesium. Molten calcium burns vigorously in oxygen
and in chlorine.

Water is slowly decomposed by calcium, and hydrogen
is evolved at a rate very convenient for the collection

CALCIUM CARBONATE.

267

of the gas. As a reducing agent in the manufacture of
certain drugs and dyes, calcium is superior to sodium be-
cause it is more moderate in its action and less dangerous
to use.

254. Uses. — An authority on electrolytic processes states
that calcium could be prepared on a large scale for less
than ten cents a pound. At such a price it would no doubt
become available in important technical processes. Cal-
cium may form valuable alloys owing to its lightness,
strength, or electric conductivity. It has been recom-
mended for the dephosphorization of steel.

CALCIUM CARBONATE.

255. Occurrence. — Calcium carbonate, CaCOg, is one of
the most abundant compounds occurring in nature. In
the form of limestone it constitutes whole mountain ranges.
Marble, which exists in

enormous quantities in vari-
ous parts of the world, is a
purer form of calcium car-
bonate than limestone.
Marble was formed from
limestone by the action of
heat and pressure under
such conditions that carbon
dioxide, which generally is
given off when limestone is
heated, was prevented from making its escape. The
mineral matter in shells is chiefly calcium carbonate
derived from the water in which the animals lived. In
past ages, deposits of shells became cemented together into
rock materials. Coquina, or the loose shell rock of Flor-

Fig. 91. Coquina.

268

CALCIUM AND ITS COMPOUNDS.

ida, illustrates an early stage of this process (Fig. 91),
and limestone a more complete transformation.

Calcite is a pure crystalline form of calcium carbonate,
and one of its varieties, Iceland spar, gives a double
refraction of light (Fig. 92). In many other minerals,
such as chalk and dolomite, calcium carbonate is present.

ii .^L^^^T liJili

256. Properties. — When pure, calcium carbonate is a
white solid, often transparent. The color of limestone
and many varieties of marble is due to the presence of

impurities. Calcium carbo-
nate occurs naturally in
amorphous masses and in
crystals of different forms.
When precipitated from so-
lutions of calcium salts by
soluble carbonates, it comes
down at first as amorphous
scales and later as minute

Fig. 92. Iceland spar.

crystals. It is only very slightly soluble in pure water,
but, as we have already seen (§ 236, 237), it is more solu-
ble in water containing dissolved carbon dioxide. The
hardness of water in limestone regions causes it to form
a closely adhering deposit on all vessels in which it is
boiled, since boiling decomposes the calcium bicarbonate,
driving off part of its carbon dioxide and leaving the
insoluble carbonate. This deposit, known as boiler aeale,
clogs the tubes of steam boilers. Hence the water is
usually softened before being introduced into the boilers.
The formation of underground caverns in limestone
regions is due to the production, solution, and decomposi-
tion of the bicarbonate. Charged with carbon dioxide
from decaying organic materials, the soil water dissolves

CAVE FORMATION.

269

limestone and sinks through cracks in the rocks, widen-
ing them as it goes. Reaching a less soluble stratum
of rock, it flows along this, but dissolves the limestone
above. This action, continuing for centuries, finally ex-
cavates a cave, such as Mammoth Cave in Kentucky
and Luray Cavern in Virginia. As soon as these caves
are hollowed out, a new process sets in. The water,

Fig. 93. Stalactites and stalagmites in a limestone cave.

before it drops from the roof, loses some of its carbon
dioxide, and part of the bicarbonate is converted into the
insoluble carbonate. This is left behind on the roof of the
The drops that fall to the floor lose more carbon

cave.

dioxide and some water by evaporation, and likewise de-
posit calcium carbonate. The final result of the process
is the formation of hanging masses of calcium carbono-te.

270 CALCIUM AND ITS COMPOUNDS.

like icicles of stone, known as stalactites, and the forma-
tion of pointed mounds below called stalagmites (Fig. 93).
In this way the cave may become nearly filled again.

257. Usea. — Natural calcium carbonate has three very
important uses. Large quantities of limestone and mar-
ble are used as building stone. Enormous amounts of
limestone are yearly burned (heated to expel the carbon
dioxide) to form quicklime. It will be remembered that
the lime unites with sand when strongly heated to form
calcium silicate, a compound which is fusible at high tem-
peratures. The third use of limestone depends upon this
property. When iron ore contains silica, limestone is
mixed with the ore to act as a flux; that is, to make the
fusion more easy. When this mixture is heated with coke
in a blast-furnace, the silica unites with the lime to form
the slag, and the iron is obtained free.

CALCIUM OXIDE.

258. Manufacture of Lime. — Calcium oxide, or quick-
lime, is made by the decomposition of calcium carbonate
at a red heat :

CaCOg — ^ CaO -h COa

The furnace or kiln (Figs. 94, 96) in which the calcining
occurs is essentially a long shaft with a fire (a) near the
bottom, so arranged that only the flames and hot gases of
combustion come in contact with the charge of limestone,
marble, or shell in the shaft. The quicklime formed is
withdrawn from time to time at the bottom (by of the
shaft, thus causing a slow sinking of the charge. By with-
drawing the slide (<?) the lime tumbles into a chamber
below.

LIME.

271

fX ki

Many kilns in this country are made of brick or large
blocks of limestone. Several feet from the bottom of the
shaft a limestone arch walls off a fire-box. Numerous
small holes are left in the lime-
stone arch to allow the flames
and hot gases from the burning
fuel to enter the portion of the
shaft above. The lime made
by this long-flame process is not
made impure by the ash from
the fuel. If the limestone con-
tains impurities, as silica, iron,
or aluminum, in any considera-
ble amount, a poor quality of lime is obtained. The suc-
cessful operation of the kiln depends largely upon the efli-
cient removal of the waste gases from the shaft so that
the decomposition of the carbonate is not checked.

Fig. 94. Lime-kiln (section).

259. Properties and Uses.

Pure calcium oxide is a soft^
white, non-crystal-
line powder which
can only be fused
and vaporized at the
temperature of the
electric arc (3000°).
It slowly takes up
moisture, forming
the hydroxide :

Fig. 95. Lime-kiln at Yellow Springs, Ohio.

CaO + HgO —
Ca(OH)2

When water is put on lumps of quicklime, cracks soon
appear on the surface, the mass swells, and finally falls to

272

CALCIUM AND ITS COMPOUNDS.

Fig. 96. Lime-light burner.

a voluminous white powder. The heat of combination is
so great that the lime becomes hot, .and clouds of steam
arise. This energetic action or process is called slaking^

and the product of the reac-
tion, calcium hydroxide, is
called slaked lime. When
quicklime is left exposed to the
air, both water and carbon di-
oxide are taken up, with the
formation of calcium hydroxide
and calcium carbonate. This
process is known as air-slak-
ing.
On account of its infusibility and dazzling incandes-
cence in the oxy-hydrogen flame, the oxide is used in the
calcium or lime-light. A lime-light burner is shown in
Fig. 96 ; a is the burner tip for an oxy-hydrogen flame,
and b is the cylinder of quicklime. Many other uses
of calcium oxide are considered in connection with the
hydroxide.

CALCIUM HYDROXIDE.

260. Properties and Uses. — Calcium hydroxide, or slaked
lime, is a soft white solid when pure, and is sparingly
soluble in water, forming a solution called lime water.
Lime water, white with suspended but undissolved par-
ticles of the hydroxide, is known as milk of lime.

When heated, the hydroxide loses water and is recon-
verted into the oxide, showing the reaction to be a reversi-
ble one, according to the temperature:

Ca(OH)2::;;i^CaO + HgO

The water solution of calcium hydroxide is strongly basic,
which property has led to its wide use as a cheap alkalL

MOBTAB AND CEMENT. 273

In this respect it stands among the bases as sulphuric acid
does among the acids.

Lime, as a cheap base, is used in the manufacture of
alkalies and bleaching-powder, in glass-making, in the
removal of hair from hides, and in many other industries.
The chief use, however, is in the preparation of mortar
and cement.

261. Mortar and Cement. — Mortar is a slaked lime mixed
with water and three or four times its bulk of sand. The
hardening is due to the drying out of the water and the
slow absorption of carbon dioxide from the air:

Ca(0H)2 + CO2— ^CaCOg + H^O

The formation of the calcium carbonate occurs first in the
outer layer of the mortar, and years may be required to
affect the inner layers. The sand makes the mortar
porous, so carbon dioxide can gain access to form the car-
bonate, and also prevents shrinkage and cracking as the
mortar dries. Moreover, the minute crystals of calcium
carbonate formed become entangled with the sharp points
of the grains of sand, making a rigid, coherent mass.

Many cements are made by strongly heating mixtures
of limestone and clay and then grinding the product.
The burning forms complex compounds which have
the power of taking on water of crystallization, and set-
ting to a hard, rigid mass. The hardening of cement can
take place under the water, and the transformation occurs
throughout the whole mass within a short time. The
hardening of mortar, however, depends upon the access
and absorption of carbon dioxide, and does not begin until
most of the water has evaporated. Then the action pro-
ceeds very slowly inward.

274 CALCIUM AND ITS COMPOUNDS.

CALCIUM SULPHATE.

262. Varieties. — Calcium sulphate is, after the car-
bonate, the most abundant and widely distributed salt of
calcium. It occurs as the mineral anhydrite, CaSO^, and
as gypsum, CaSO^. 2 HgO. Satinspar, alabaster, and selen-
ite are varieties of gypsum. Selenite is often found in
large, transparent crystals that can be^ scratched with the
finger nail.

Properties and Uses. — Gypsum is but sparingly
soluble in water ; its solubility increases to 40° and then
decreases. When heated, gypsum loses three-quarters of its
water of crystallization, and the residue may be said to
have one molecule of water of crystallization to every two
molecules of calcium sulphate (CaS04)2. HgO. The chalky
Plaster powder resulting from the heating is tnown
of Paris. as plaster of Paris. On a large scale it is made
by heating in kilns a charge of gypsum broken into»small
lumps to insure evenness in "burning." Care is taken
not to overheat ; 126'^ C. is the most favorable tempera-
ture for the process.

When plaster of Paris is wet, water is again taken up,
forming needles of crystallized gypsum which, becoming
entangled, set or harden the mass. The hardening is also
Plaster accompanied by a slight increase in volume,
casts. This property explains the use of plaster of Paris

in making casts. The slight expansion secures a sharp
impression of the mould. The powder is mixed with about
a third its weight of water, the pasty mass put into the
mould, and in less than half an hour the plaster sets. The
ivory surface of casts is secured by dipping them in melted
paraffin or by painting them with a solution of paraffin in

CALCIUM PHOSPHATES. 275

petroleum ether. The solvent in the latter case evapo-
rates, leaving the waxy filling in the pores of the cast and
making it impervious to water. Many beautiful metallic
objects are now made by plating a thin film of metal on a
plaster of Paris base. Plaster of Paris is also used for
rigid bandages in surgery and as a surface coating for
walls. Stucco is plaster of Paris and rubble, mixed with
sizing or glue instead of water.

264. Caloinm Phosphates. — The phosphates of calcium
are of great importance to organic life. The bones of ani-
mals are essentially normal calcium phosphate^ CajCPO^)^.
This compound occurs as phosphorite which has been
derived from animal remains. Guano contains phosphor-
ite in addition to nitrogenous compounds. Phosphates are
an important plant food, but to be available must be in a
soluble form that can be taken up by the plants. The
soluble phosphates result from the decomposition of rocks
containing phosphates. Since the process is a slow one^
the- supply of phosphates in cultivated soils often becomes
exhausted. To supply this need, the manufacture of solu-
ble phosphates for fertilizers has grown to be an important
industry.

The superphosphate of lime^ CaH^(P0^)2 . 2 HgO, is the
most important artificial fertilizer. It is made by treat-
ing phosphorite or bone-ash with crude sulphuric acid.
The superphosphate formed is readily soluble in water,
and when spread upon the soil is available for plant use.

265. Bleaching-Powder. — Bleaching-powder, or chloride
of lime, is made by passing chlorine over freshly slaked
lime spread on the floors of a series of absorption cham-

- bers (Fig. 97). Chemists are still in doubt as to the

l-Tf'Ftiiriir

•'itnimrnJr'

276 CALCIUM AND ITS COMPOUNDS.

reactions involved in the process and the formula of the
product. The latter, however, is often represented as
CaOClg.

Bleaching-powder is an unstable white powder which
is slightly soluble in water. When bleaching-powder is

treated with acids, chlorine
is evolved. Hence the
powder is used as a source
of chlorine for bleaching
Fig. 97. purposes.

The cotton or linen to
be bleached is freed from grease and oil. The cloth is
next soaked in a solution of bleaching-powder, then dipped ,
in dilute acid, and finally thoroughly washed to remove
the chemicals (cf. § 57). The solutions used are very
weak, to prevent injury .to the fibre of the cloth.

When exposed to the air, bleaching-powder slowly de-
composes, with the absorption of carbon dioxide. As
a result chlorine is liberated, and for this reason th# pow-
der, often called "chloride of lime," is used as a disin-
fectant and germicide.

Summary.

Calcium, although very abundant in nature, is rarely seen as
metal.

It can be obtained by the electrolysis of fused calcium chloride.
The most important calcium compounds are :

the carbonate (limestone, chalk, marble) ;

the hydroxide (slaked lime) ;

the oxide (quicklime) ;

the phosphate (phosphorite) ;

the sulphate (gypsum).
Lims is made by heating calcium carbonate. Slaked lime is
made by adding water to quicklime.

EXERCISES. 277

Slaked lime is used in making mortar, which hardens by the
evaporation of water and the absorption of carbon dioxide.

Plaster of Paris is made by partly dehydrating gypsmn. It is
used in making plaster casts.

Calcium phosphates, derived from bone-ash and mineral beds,
are used in making fertilizers.

Bleaching-powder is made by passing chlorine over slaked lin^e.
It is used in bleaching and as a disinfectant.

Bxercises.

1. Why is the production of lime hastened by blowing air
or steam into a lime-kiln ?

2. How could quicklime be made from slaked lime ?

3. In the laboratory, loosely stoppered bottles that contain
quicklime are sometimes found with the sides broken out.
How would you account for this ?

4. Lime-water standing exposed to air becomes coated
with a film of insoluble substance. What is the substance ?
Explain its formation.

5. Compare the hardening of mortar with that of plaster
of Pai'is.

6. For what reason is normal calcium phosphate converted
into " superphosphate " in the manufacture of fertilizers ?

7. Calculate the weight of quicklime which a manufacturer
might expect to get from 1000 kilograms of pure limestone.
What weight of carbon dioxide would be given off during
the action ? What volume would the gas have, standard con-
ditions ?

8. What weight of nitric acid would be required to neu-
tralize 35 grams of calcium hydroxide ?

9. What weight of water enters into combination in
slaking 500 lb. of quicklime?

10. Compare the electrolytic processes for making sodium
and calcium.

QCj

CHAPTER XXVII.
MAaHESITJM, ZnrO, AND MESOTTBT.

MAGNESIUM.

266. Oocnrrenoe and Preparation. — Although magnesium
is of comparatively little importance in a commercial way,
its compounds are very abundant in nature. The most
important of these are dolomite, a double carbonate of
calcium and magnesium; magnesite, a carbonate; various
silicates, of which asbestos and hornblende are examples ;
carnallite, a double chloride of potassium and magnesium,
KCl . MgClg . 6 HgO. Magnesium sulphate is fpund in
certain mineral waters.

The metal is commercially obtained by the electrolysis
of carnallite. The salt is fused, together with some
common salt or cryolite, in an iron crucible which serves
as the cathode. A carbon rod serves as the anode.

267. Propertiefl and Uses. — Magnesium is a silvery white
metal of low specific gravity. It resembles both calcium
and zinc in its properties, and stands between them in me-
tallic character. It decomposes water slowly at 100°, but
does not affect it at ordinary temperatures. Moist air
acts on magnesium slowly. It burns with comparative
ease, with the evolution of a brilliant white light of great
actinic power. The oxide is formed in this reaction :

2Mg + 02— ^ 2MgO

Dilute acids react with magnesium very readily, hy-
drogen usually being evolved.

278

MAGNESIUM COMPOUNDS. 279

The common salts of magnesium are stable substances,
soluble in water, with the exception of the carbonate and
the phosphate.

Magnesium is used in flashlight preparations for photo-
graphic purposes because of the actinic power of the light
it gives in burning. It is also used in making fireworks.

268. Componnds of Magnesium. — The formulas of the
magnesium compounds show that the element has a va-
lence of two, or we may say it forms bivalent ions.

Many magnesium compounds, when heated in an
oxidizing flame, are converted into magnesium oxide.
This, when moistened with a solution of cobalt nitrate
and heated, yields a mass having a pale pink color.

Magnesium is one of the few elements that enter into
direct combination with nitrogen. When nitrogen is
passed over red-hot magnesium, magnesium nitride,
MggNg, is formed.

Mixtures of magnesium carbonate with magnesium
hydroxide are used in pharmaceutical preparations and
in face powders.

Several salts of magnesium are used medicinally. One
of these, the sulphate, is found in nature as Upsom salts
(MgSO^ . 7 HgO). Certain spring waters owe their laxa-
tive properties to the presence of magnesium salts.

ZINC.

269. Eztruotion from Ores. — Zinc is not found in the
uncombined state. Its common ores are zinc-blende,
ZnS; smithsonite, ZnCOg; zincite, ZnO.

To separate zinc from the oxide, the ore is finely pow-
dered and mixed with coal. The mixture is then heated

280

MAGNESIUM, ZINC, AND MERCURY.

in earthenware retorts (Fig. 98, a). The carbon reduces
the zinc oxide :

ZnO + C -^ Zn + CO

The temperature in the process is raised above the boiling-
point of the metal, 950°, which consequently passes off as
gas, and is condensed in earthenware or iron receivers (J).

Fig. 98.

When the ore is not an oxide, a preliminary operation
must precede the reduction. This consists in heating
the ore on grates in contact with air. The operation is
known as roasting, and converts the metal into an oxide.
In the case of zinc sulphide, care is taken not to convert
it into the sulphate.

ZnCOg — >. ZnO + COg
2 ZnS + 3 O2 — ^ 2 ZnO + 2 SOg
The oxide that results in these reactions is then reduced

PROPERTIES OF ZINC. 281

with the carbon in the manner that has been described.
The silicate is reduced directly. A rotary roasting-fur-
nace is shown in Fig. 99. The ore is dumped through
the hopper (J),and the flames from the fire-box (a) sweep
over the ore as the hearth slowly rotates; c, c are the
fume chambers.

Zinc, as extracted from its ores, usually contains carbon,
arsenic, cadmium, and other impurities. It is freed from
these by distillation.

270. Physical Properties. — Zinc is bluish white in
color. It comes into the market in the form of heavy
bars called ingots or spelter^
formed by pouring the
melted metal into moulds.
In this form the metal is
crystalline in structure and
rather brittle. Between Fig. 99.

100° and 160° it is malleable

and ductile, and can be rolled into sheets. After having
been obtained in this form, it remains malleable at ordi-
nary temperatures. At 300° it again becomes brittle and
can be powdered.

Granulated or mossy zinc is a form much used in the
laboratory. It is made by pouring the melted metal
into water. Zinc dust is obtained in the distillation of the
metal. As long as the receiver remains comparatively
cold, the distilled zinc collects in the form of a powder.
This operation is similar to the one by which sulphur is
obtained as flowers of sulphur. Zinc dust always contains
a certain amount of the oxide.

271. Chemical Properties. — Zinc is regarded as dis-
tinctly metallic, but it difl^ers considerably from such

282 MAGNESIUM, ZINC, AND MERCURY.

metals as calcium and sodium. It resembles cadmium and
magnesium more closely. Zinc does not act on water at
ordinary temperatures. Air attacks it slowly in the
presence of moisture, forming a basic carbonate, which
acts as a protective coating, so that only the outer layer
of the metal is affected. When zinc is heated in air or
oxygen, it burns with a bluish flame, forming zinc oxide.

2Zn+02 — ^2ZnO

Zinc reacts readily with dilute acids, forming zinc salts,
and, as a rule, liberates hydrogen :

Zn + 2 HCl -^ ZnCla + Hj
Zn + HjSO^ -^ ZnSO^ + H^

In acting on zinc (or other metals) nitric and sulphuric
acids, if concentrated, do not liberate hydrogen, since they
act as oxidizing agents and convert the hydrogen into
water. In these cases, the gases that are given off are
reduction products of the acids. In the case of sulphuric
acid, sulphur dioxide is obtained (§ 146) ; from nitric acid,
nitric oxide or nitrogen peroxide is produced (§ 173).

The action of dilute acids on zinc is hastened by the
presence of certain solid substances in contact with the
metal. Pure zinc will scarcely react with acids, but if
it is impure, solution takes place with great rapidity. A
similar effect is produced by the presence of a very small
amount of copper, or other metal, as a deposit on the sur-
face of the zinc. In these cases the particles of carbon
or copper act like cathode plates of a voltaic cell. Hence
the velocity of the evolution of hydrogen from acids is
increased by their presence.

Many compounds of zinc, when heated on charcoal or on
a plaster block before a blowpipe, yield zinc oxide, yellow

ZINC OXIDE AND HTDBOXIDE. 283

when hot, and white when cold. If the oxide is mois-
tened with a drop of a solution of cobalt nitrate, and again
heated, a bright green mass containing a compound of zinc
and cobalt oxide is obtained. Figure 100 shows the plaster
of Paris blocks with a hollow, made by the forceps, to
contain the compound to be tested by heating with cobalt
nitrate.

. 272. Uses. — Zinc is used in making several important
alloys. Brass is composed of copper and zinc ; Grerman
silver contains zinc, copper,
and nickel; bronze some-
times contains zinc in addi-
tion to copper and tin.

Q-alvanized iron is iron
covered with a thin layer of
zinc, which acts as a protec-
tive coating and prevents ^'
rusting. The iron is first cleaned and then dipped into
molten zinc. Of late an electrolytic method of galvaniz-
ing has come into use. The process is analogous to cop-
per or silver plating. Another important use of zinc
is for the anode plates of batteries.

273. Zinc Oxide and Hydroxide. — Zinc oxide, ZnO, is
much used as a base for white paints. It does not have as
great covering power as white lead, but it has the advan-
tage of not turning black from contact with hydrogen
sulphide. It can be made by burning zinc or by heating
zinc hydroxide or zinc carbonate:

Zn(0H)2 (heated) —^ ZnO + HgO
ZnCOg (heated) — ^ ZnO + COg

284 MAGNESIUM, ZINC, AND MERCURY.

On adding potassium hydroxide to a solution of a zinc
salt, zinc hydroxide is precipitated, since this substance is
insoluble in water. If an excess of potassium hydroxide
is added, the hydroxide is dissolved, forming potassium
zincate.

ZnClg + 2 KOH — ^ 2 KCl + Zn(0H)2

Zn(0H)2 + 2 KOH -^ KgZnOa + 2 HgO

274. Salts of Zinc. — Zinc chloride, ZnClg, is obtained
by the action of hydrochloric acid and zinc. It is an
extremely deliquescent substance, sometimes used as a
drying agent. It also has the power to dissolve metallic
oxides ; because of this property it makes a good flux for
soldering metals. Wood which has been soaked in a solu-
tion of zinc choride resists decay.

Zinc sulphate, ZnS04, is used in making battery solutions.

Zinc sulphide, ZnS, found in nature as zinc-blende, can
be precipitated from solutions of zinc salts by the addition
of hydrogen sulphide:

ZnOlg + HgS — ^ ZnS + 2HC1

But this reaction is reversible; that is, zinc sulphide will
dissolve in dilute hydrochloric acid with the formation of
zinc chloride and hydrogen sulphide:

ZnS + 2 HCl — ^ ZnCla + H^S

For this reason, the reaction shown in the first equation is
never complete; for when a certain amount of hydrochloric
acid has been formed, the second reaction begins to take
place. Hydrogen sulphide is only slightly ionized but its
ions are necessary for the first reaction. If, however,
an acid is present, its hydrogen ions force back the disso-
ciation of hydrogen sulphide to such an extent that the
number of sulphur ions present is very small. If we dis-

MERCURY. 285

pose of the hydrogen ions as fast as they are formed, by
adding such a substance as ammonium hydroxide, the pre-
cipitation of the zinc sulphide will be complete. If, on
the other hand, much acid is present in the solution, the
precipitation will be entirely prevented.
All soluble zinc salts are poisonouB.

MERCURY.

275. Oocnrrenoe and Separation. — Mercury is found only
in a few localities, the deposits of Spain and California be-
ing the most important. It occurs native in small drops
mixed with earthy materials, and in the sulphide, ciwndbar^
HgS. From the latter it is obtained by roasting, so as to
convert the sulphur into the dioxide and vaporize the mer-
cury, which is condensed and purified :

HgS + O2 -^ Hg + SO2

276. Physical Properties. — At ordinary temperatures,
mercury, commonly known as quicksilver^ is a silvery-
white liquid, with a brilliant metallic lustre. Its density
is greater than that of lead, so that iron easily floats on it.
Mercury solidifies to a substance resembling tin at about
— 40*^, and boils at a temperature below red heat, but
vaporizes slowly at ordinary temperatures. It is a good
conductor of electricity. The molecular weight of mer-
cury, as found from its vapor density, is the same as the
atomic weight, 200; hence there is one atom in the mole-
cule of mercury vapor.

Mercury has the power of dissolving many other metals,
forming alloys with them called amalgams. Mercury
dropped on a gold ring will whiten it by amal-
gamating with the gold. These amalgams are
not true compounds, as they may be formed in varying

286 MAGNESIUM^ ZINC, AND MERCURY.

proportions. When there is a large excess of mercury,
amalgams are liquid; otherwise they are solid.

The vapor of mercury is highly poisonous, as is the
metal itself when finely divided. Mercury can be ob-
tained as a fine gray powder by shaking it violently with
flour, grease, or any substance which will coat the minute
drops and prevent them from uniting to form a fluid mass.
This process, known as extinguishing^ is used in the prep-
aration of blue pills and mercurial ointments.

277. Chemioal Properties. — Mercury combines readily
with the halogens and sulphur. Oxygen does not com-
bine with it at ordinary temperatures, but at high temper-
atures it forms oxides, which at still higher temperatures
dissociate into mercury and oxygen. This is shown by
the reversible equation:

2Hg + 0,:^2HgO

Pure dilute acids do not attack mercury. Concentrated
nitric acid dissolves it readily, and dilute nitric acid also
attacks it in the presence of nitrogen peroxide.

278. Uses. — Mercury is used in important scientific
instruments, such as the thermometer, barometer, and
others. It is also used for the collection of gases soluble
in water. Its most important uses, however, are in the
various amalgams. Sodium amalgam finds considerable
use in the laboratory, and tin amalgam is used to coat the
back of mirrors. Amalgams of silver and other metals
are used in filling teeth. Gold and silver are extracted
by allowing the crushed ore to flow in a thin mud over
tables covered with mercury. The gold amalgamates
with the mercury, from which it can be separated by die-
tilling the mercury.

COMPOUNDS OF MEBCURT. 287

279. Componndfl of Meronry. — Mercury forms two series
of compounds, the mercurous and the mercuric. The
chlorides are the most important salts, and may be taken
as typical of the two series.

Mercurous chloride^ known as calomel^ has its composi-
tion represented by the formula HgCl. Since it is in-
soluble in water, it may be prepared by treating
a solution of a mercurous compound with a chlo-
ride. It is produced commercially by heating a mixture
of mercuric chloride and mercury, when it sublimes as a
white powder :

HgCla-hHg— 9^2HgCl

Exposed to the light, mercurous chloride slowly blackens
on account of the liberation of mercury by the reversal of
the above reaction. It is extensively used in medicine.

The common name of mercuric chloride is corrosive sub-
limate^ and its formula is HgClg. It is made by heating
a mixture of sodium chloride and mercuric corrosiye
sulphate ; the chloride sublimes, as its name BnWimate.
indicates. The sublimate is a white, translucent mass,
from which the salt can be obtained in silky needles by
dissolving in water and recrystallizing. It is slightly
soluble in water at ordinary temperatures, but at higher
temperatures it is more soluble. Corrosive sublimate
is a violent poison. It is also a powerful antiseptic
and germicide. For this purpose, very dilute solutions
are used (1 part to 1000 parts of water), as its antiseptic
properties increase with the ionization. Hence the dilute
solutions in which the salt is fully ionized are better
antiseptics than the more concentrated ones. With the
alkaline chlorides it foi*ms double salts more soluble than
mercuric chloride by itself, and much used in making anti-
septic solutions.

288 MAGNESIUM, ZINC, AND MERCURY.

MAGNESIUM.
Exercisea

1. How is magnesium prepared from camallite ?

2. For what is magnesium used ?

3. Magnesium oxide is slightly soluble in water. Would
the solution give an acid or an alkaline reaction ? Why ?

4. What reaction takes place when dilute sulphuric acid is
added to magnesium ?

^ 5. What is Epsom salts ?

6. Mention two ways by which carbon dioxide could be
obtained from magnesite.

7. How many grams of magnesia, MgO, could be prepared
by heating 20 lb. of magnesium carbonate ?

a When 0.362 gram of magnesium was added to an excess
of dilute acid, 365 c.c. of hydrogen were liberated. At the time
the measurement was made, the temperature was 21° C. and the
pressure was 770 mm. From the data given, calculate the hy-
drogen equivalent of magnesium.

9. Using your answer to 8, calculate the atomic weight of
magnesium.

10. What is a test for a magnesium compound ?

ZINC.

Summary.

Zinc is usually obtained by roasting the ore and then reducing
the oxide.

Spelter, sheet zinc, granulated or mossy zinc, and zinc dust are
commercial forms of zinc.

Brass and German silver are common alloys of zinc.

Zinc is bivalent.

Zinc hydroxide acts as a base in the presence of strong adds
and as an add in the presence of strong bases.

EXERCISES. 289

The carbonate, chloride, suphide, and sulphate are common
compounds.

ExerciaeB.

1. Starting with zinc carbonate, describe the preparation
of four commercial forms of zinc.

2. Name two alloys of zinc and tell what each contains.

3. Why does zinc corrode very slowly in air ?

4. Name a compound of zinc used as a paint base. Why
is it of value in making paint for use in chemical laboratories ?

5. Name a compound which would form zinc hydroxide on
the addition of the right amount of hydrochloric acid.

6. Write equations showing how hydrogen could be pro-
duced by the reaction of either an acid or a base with zinc.

7. Why is wood sometimes impregnated with a solution of
zinc chloride ?

a Why is tin-plate, preparatory to soldering, often wet
with a solution of zinc chloride ?

9. Why will not hydrogen sulphide completely precipitate
zinc, as zinc sulphide, from a solution of zinc sulphate ?

10. Explain the fact that water solutions of zinc sulphate
give an acid reaction.

11. Describe a test for zinc.

12. How many grams of zinc would be required to replace
the hydrogen contained in 15 grams of sulphuric acid ?

13. Mention two ways by which zinc sulphide can be formed.

MERCURY.
Summary.

Mercury is found as the sulphide from which it is extracted by
roasting.

It is a bright, silvery white fluid, solidifying at -39° C, boiling at

290 MAGNESIUM, ZINCy AND MERCURY.

357^ C, and with a specific gravity of 13.6. The atomic weight
of mercury is 200.

Mercury combines readily with sulphur and the halogens, but
with oxygen only at high temperatures. Concentrated nitric acid
is the only add that has much action upon it at ordinary tempera-
tures.

Mercury is used in scientific instruments, in the preparation of
amalgams for mirrors, and for the extraction of gold.

Mercurous chloride, calomel^ and mercuric chloride, corrosive
syblimate, are the two most important mercury compounds. Calo-
mel is used in medicine, and corrosive sublimate is a powerful

antiseptic.

Exercises.

1. Write the equation for the extraction of mercury from
cinnabar.

2. What properties fit mercury for use in thermometers ? in
barometers ?

3. What common metals would float on mercury ? Kame
those that would sink. How would lead act ?

4. Under what circumstances is mercury desirable in the
collection of gases ? What are its disadvantages ?

5. Why do surgeons dip their hands into a solution of
corrosive sublimate before performing an operation ?

6. What special name is given to the alloys of percury ?
Name two, and give a use of each.

7. Calculate the percentage composition of the two chlorides
of mercury. What law is illustrated by the composition of
these compounds ?

8. What would be a simple way of testing a solution for a
soluble salt of mercury ?

CHAPTER XXVIII.

2^0

OOPPEE AND ITS OOMPOUNDS.

280. Occurrence and Ores. — Copper is the only metal
which occurs free in large, widely distributed deposits.
For this reason, it was the first metal extensively used
by man. The copper age followed the stone age. The
island of Cyprus was noted in the time of the Romans for
its production of copper or Cyprian brass. We obtain the
symbol Cu from the Latin name, cuprum.

Fig. 101. Mass of native copper (3x2 ft.) in the University of Michigan

Museum.

The noted mines of native copper in Michigan, along
the southern shore of Lake Superior, were extensively
worked before Columbus discovered America. From
them masses of copper of enormous size, one of which

291

292 COPPER AND ITS COMPOUNDS.

weighed nearly five hundred tons, have been obtained.
These mines are still an important source of copper.

The ores of copper are numerous, and many of them
have a composition represented by complex formulas ; the
more important ores besides native copper are sulphides,
oxides, carbonates, and silicates. Much copper is obtained
from an ore named chalcopyrite, the composition of which
corresponds approximately to the formula CugS . FcgSg.
Malachite, a basic carbonate of copper (CuCOg . Cu(0H)2)
is of interest. Polished slabs of malachite often exhibit
variegated patterns of different shades of green which are
of great beauty, and the mineral is highly valued for orna-
mental purposes.

281. Metallurgy; Steps in Process. — The metallurgy of
copper is complex ; not only does the process vary with
the kind of ore used, but similar ores are seldom treated
in the same manner in different localities. In case the
ore contains much chalcopyrite, the process, as carried
out by one of the large copper companies, consists essen-
tially in:

1. Roasting a portion of the ore by which sulphur
dioxide is obtained and used in the manufacture of sul-
phuric acid.

2. The production of a complex sulphide called matte.

3. Converting the matte into blister copper.

4. Poling the blister copper and casting it into anode
plates.

5. Refining by electrolysis.

282. Roasting of Ore and Production of Matte. — When
chalcopyrite, CugS . FcgSg, is roasted, the following reac-
tions may be considered as taking place:

BOASTING OF ORE.

293

2 CuaS + 3 O2 = 2 Cu^O + 2 SO^
FejSj + 4 O2 = 2 FeO + 3 SO^

The roasted ore is then smelted in a blast-furnace for
the production of matte. A vertical section of one form
of blast-furnace is represented in Fig. 102. The cruci-
ble ((7) is the part of
the furnace in which the
molten matte and slag
collect. The body of the
furnace consists of two
concentric shells (c«),
made either of wrought
iron or of steel, between
which cold water ( W) is
caused to circulate to
prevent the inner shell
becoming heated to a
temperature sufficiently
high to injure it. Pipes
called tuyeres (^T) enter
the furnace a short dis-
tance above the hearth.
They are used to convey

a blast of air into the furnace. Above the body of the
furnace extend the hood (IT) and the stack (aS). A
door (2>), used in charging the furnace, is placed in the
hood.

The charge for the furnace consists of the roasted ore
mixed with a certain amount of unroasted (green) ore
and coke. The ores usually contain considerable silica.

When the furnace is in operation, the blast of air which
enters through the tuyeres causes a part of the coke to be

Fig. 102.

294

COPPER AND ITS COMPOUNDS.

converted into carbon monoxide ; this, together with the
hot carbon, reduces a part of the cuprous oxide to copper :

CuaO + CO— ^ 2 Cu + CO2

As copper has a greater tendency to combine with sul-
phur than has any other metal present, cuprous sulphide
is formed :

3 CugO + FcgSg— >-3 CujS + Fefi^

3 Cu + FcoS

2*^3 ■

.3CuS + 2Fe

A part of the iron unites with the silica which is present
in the ore and forms a fusible slag. Some of the iron
sulphide remains with the cuprous sulphide and forms
matte^ which is a more or less pure mixture of cuprous
and ferrous sulphides. In addition to the sulphides men-
tioned, matte contains all of the gold and silver present
in the ore, and generally arsenic and antimony.

The matte and slag are run into a forehearth, where
the lighter slag rises to the top and overflows through a
trough, while the matte collects at the bottom and is from
time to time drawn off through a pipe. Matte contains
from 45 % to 60 % of copper.

283. Conversion of Matte to
Blister Copper. — Matte is con-
verted into blister copper by
means of a process called Bes-
semerizing, in a furnace called
a converter (Fig. 103). This
consists of an iron shell (a)
lined with a thick layer of a
mixture of quartz and clay (J).

The converter is mounted so

Fig. 103. Converter (sectional), that it can be revolved around

CONVERSION OF MATTE TO BLISTER COPPER. 295

a horizontal axis. A blast of air enters through a pipe
(c) and is forced through openings near the bottom of the
converter. The matte is melted and poured into the con-
verter ; then air is blown through the molten mass (Fig.
104). Sulphur, iron, and other impurities are oxidized.
Those oxides which are volatile are driven off. The iron
oxide, however, unites with the quartz of the furnace
lining to form a slag. As sulphur and iron have a greater
tendency than copper to combine with oxygen, they are
the first to be oxidized. The operator endeavors to stop
the blast as soon as the impuri-
ties have been removed, and
thus prevent the oxidation of
the copper. This requires much
experience and great skill. The
copper obtained still contains all
of the gold and silver present in
the original ore, and small quan-
tities of other impurities. Mol-
ten copper dissolves consider-
able sulphur dioxide, which is
expelled when the mass cools,
and gives the copper the appear-
ance which causes it to be called
blister copper.

It would not be advisable to put the blister copper to
practical use, since the gold and silver are of too great
value to be allowed to remain with the copper. Besides,
large quantities of copper are used in the manufac-
ture of electric cables, and small quantities of impuri-
ties greatly reduce its conductivity. If the conductivity
of pure copper is considered as 100, copper containing
0.8% of arsenic has a conductivity of only 30, and

By courtesy of 'J he Scientific
A nterican*

Fig. 104.
Converter in operation.

296 COPPEB AND ITS COMPOUNDS.

copper containing 0.5 % of silicon has a conductivity
of 28.

284. Poling of Blister Copper and casting of Anode Plates.
— Blister copper is melted in a reverberatory furnace
(Fig. 105), and the molten mass is stirred by the
gases coming from a long pole or log of green wood,
which is forced into the metal. The hydrocarbons dis-

Fig. 105. Reverberatory furnace for poling copper.

tilled from the wood unite with the oxygen combined
with the copper. This process is called poling^ and has
for its object the reduction of the small amount of copper
oxide present to metallic copper. Any one who has seen
the interior of a poling furnace in operation will retain a
vivid impression of the seething mass of molten copper,
dazzling in its brilliancy of color. After being poled, the
copper is cast into anode plates to be refined by electrolysis.

REFINING BY ELECTROLYSIS.

297

\^n^

<^ if7-ia-n- £^. f

Fig. 106.
a, anodes ; c, cathodes.

285. Sefining by Electrolysis. — The anode plates from
the poling furnace are about | of an inch thick, 3 feet
wide, and 3 feet long.
Cathode plates of pure
copper about ^ of an inch
thick are made by cover-
ing a plate of impure cop-
per with a thin layer of
grease, and then making
it the cathode plate in a
bath of copper sulphate.
On the passage of the current, the anode plate dissolves,
and pure copper is deposited on the greased side of the
cathode plate. The layer of grease prevents the pure
copper from adhering firmly to the impure ; it can there-
fore be stripped off and used as a pure copper cathode.

The anode plates are arranged in series with the pure
cathode plates (Fig. 106). Some copper refiners do not use
pure copper cathode plates, but arrange the impure plates

in multiple (Fig. 107).
In such cases, the copper
from the front of one plate
is deposited on the back of
the next one. At the end
of the electrolysis the pure
copper is stripped off.
The plates, either arranged
in series or multiple, are
suspended in wooden tanks containing a warm solution of
copper sulphate acidulated with sulphuric acid (Fig. 108).
During the electrolysis the bath is kept slowly circulat-
ing, and at a definite concentration. When the current
passes, pure copper is deposited on the cathode. A part

w//w/Mm;mw/ww/m''''-.

Fig. 107.

298

COPPER AND ITS COMPOUNDS.

of the impurities enter the bath, while others, including
gold and silver, fall to the bottom of the tank and form a •
substance known as mud. The gold and silver are recov-
ered from the mud.

III. inn 1 1

j\M\f^^¥

By courtesy of TAe Scientific American.

Fig. 108. Tank-house for electrolytic copper refining.

286. Properties of Copper. — Copper has a characteristic
reddish color. Only two of the common metals, gold and
silver, surpass it in malleability and ductility. It stands
next to silver as a conductor of electricity.

On exposure to the atmosphere, copper is attacked by
carbon dioxide in the presence of moisture, and becomes
covered with a coating of a basic carbonate of a greenish
color. The coating, once formed, adheres to the copper
underneath and protects it. Copper is readily attacked
by nitric acid (preparation of nitric oxide, § 168), but
neither dilute hydrochloric acid nor dilute sulphuric acid

U8E8 OF COPPER. 299

attack it in the absence of air. It is readily acted upon
by hot, concentrated sulphuric acid (preparation of sul-
phur dioxide, § 139). Boiling concentrated hydrochloric
acid slowly converts copper into cuprous chloride, CuCl.

287. ITses of Copper. — Large quantities of copper are
used for a great variety of purposes. Among the more
important may be mentioned its use as wire and cables
for the transmission of electric currents ; its employment
in the manufacture of various articles for domestic and
scientific purposes, such as water-heaters, kettles, stills,
vacuum pans, etc. Much copper is also employed in the
manufacture of alloys. Brass is an alloy of Qopper and
zinc ; bronze, an alloy of copper, zinc, and tin ; and alumi-
num bronze, an alloy of copper and aluminum.

288; Compoimds of Copper. — Copper forms two kinds
of ions, cuprous, Cu+, and cupric, Cu"''"''. Its valence
may therefore be considered to be sometimes one and at
other times two. The monocuprion unites with negative
ions to form cuprous compounds, while the dicuprion
unites with the negative ions to form cupric compounds.

Oxide

Snlpbide

Chloride

Cuprous

CujO

CUjS

CuCl

Cupric

CuO

CuS

CuCl,

289. Oxides of Copper. — Cuprous oxide^ or red oxide of
copper, CugO, occurs in nature. When a strip of copper
is heated in air, a layer of cuprous oxide forms under
the layer of black cupric oxide. If a mixture of cupric
oxide and charcoal is heated, the cupric oxide is first
reduced to cuprous oxide, and then the cuprous oxide
is reduced to copper. Other reducing agents have a

800 COPPER AND ITS COMPOUNDS.

similar effect on cupric oxide. The formation of cuprous
oxide is utilized in testing for glucose by means of Feb-
ling's solution. Fehling's solution contains cupric sul-
phate, potassium hydroxide, and Rochelle salt. When
it is added to a solution containing glucose, or a similar
reducing agent, and the mixture is boiled, cuprous oxide
separates as a red precipitate. The Rochelle salt is added
to prevent the formation of cupric oxide, which, being
black, would hide the color of the cuprous oxide. Cuprous
oxide is used to give a beautiful red color to pottery.

Cupric oxide^ or black oxide of copper, CuO, can be pre-
pared by heating copper in air and also by heating cupric
hydroxide, nitrate, or carbonate. Many compounds con-
taining hydrogen are oxidized when heated with cupric
oxide, the hydrogen being converted into water. If car-
bon is present, it is converted into carbon dioxide. These
facts make cupric oxide a valuable substance to use in the
determination of the quantity of hydrogen and of carbon
present in compounds containing these elements.

290. Preparation of Copper Sulphate. — Crystallized cop-
per sulphate^, or blue vitriol, CuSO^ . 5 HgO, is prepared on
a large scale by placing coarsS copper shot in a perforated
lead basket, and then causing the basket and contents to
move up and down so that they will at one time be in the
air and at another time immersed in warm, dilute sul-
phuric acid. When the basket enters the acid, air is car-
ried into it with the shot. The action of the acid on the
copper in the presence of air results in the formation of
copper sulphate, which passes into solution:

2 Cu -h O2 -h 2 H2SO4 — ^ 2 CUSO4 -h 2 HgO

The solution, after being sufficiently concentrated, is

PROPERTIES OF COPPER SULPHATE.

301

allowed to stand in lead-lined vats in which are suspended
lead strips. Blue vitriol crystallizes on the lead and is
purified by recrystallization.

Blue vitriol is also obtained as a by-product in one
method used in separating gold from silver. The melted
alloy of these two metals is granulated by pouring it into
cold water. The granulated mass is boiled with concen-
trated sulphuric acid until the silver is dissolved as silver
sulphate. The gold remains undissolved and settles to
the bottom of the vat. The solution of silver sulphate is
removed to lead-lined vats and the silver separated by the
addition of copper:

Ag2S04 + Cu — >■ CUSO4 4- 2 Ag

291. Properties and ITses of Copper Snlphate. — Copper'
sulphate forms deep blue crystals (Fig. 109) which efflo-
resce in dry iair. Its water solution gives an acid reaction
with litmus (§ 122).

^■^^^ffPSjI^SBWyw^^^

Fig. 109. Mass of copper sulphate crystals.

For some time the great value of copper sulphate as a
fungicide has been recognized. A mixture of copper sul-
phate and slaked lime, known as the Bordeaux mixture,
is now extensively employed for this purpose. A thick
paste of calcium hydroxide and copper sulphate was first
used near the city of Medoc, France, to keep boys from
stealing grapes. When placed upon the trellises and vines

302 COPPER AND ITS COMPOUNDS.

it was conspicuous, and was believed to be poisonous. In
1882, Millardet, professor of botany in Bordeaux, visited
the vineyards near Medoc, and was informed by the grape
growers that portions of the vineyard which had been
treated with the paste were not attacked with mildew.
Much work has been done in the United States Department
of Agriculture in determining the value of the Bordeaux
mixture as a general fungicide. Its use in this country
has saved crops worth many thousands of dollars.

Dilute solutions of copper sulphate are used to moisten
seeds of cereals prior to sowing, to prevent the attack of
fungi called smuts.

Plants known as algse grow abundantly in the water of
londs and reservoirs. Some of them impart to the water
disagreeable odors; others produce effects equally unde-
sirable. Copper sulphate is added to the water of ponds
thus affected, in the proportion of one part of copper sul-
phate to from one to eight million parts of water, for the
purpose of destroying algae. The solution is too dilute to
kill fish. The copper sulphate appears to react with the
albumen of the algae to form an insoluble substance which
sinks to the bottom of the ponds.

Copper sulphate has many other important uses. It is
employed in batteries, in electroplating, as a mordant in
dyeing, and for making other compounds of copper.

Suminary.

Copper occurs as metal; this and the sulphides are its principal
sources.

It is separated from its ores by burning out the sulphur and re-
ducing the oxide by carbon. It is purified by electrolysis.

The atomic weight of copper is 6S.6 ; its specific gravity is 8.9,
and its melting point is 1065° C.

EXEBCI8E8. 303

Copper is durable under ordinary atmospheric conditions, and
IS used for protective coverings. Being ductile and a good con-
ductor, it is used for electric conductors. As a constituent of
many alloys, copper finds wide use.

Copper sulphate is the most important compound of copper. It
is used as a fungicide, for paints, for plating, and in some batteries.

Bzercisea.

1. What metals are usually found associated with copper ?

2. How is iron separated from copper ? how is silver ?

3. What copper compound is used in pottery and in glass ?
Why?

4. What is the result of the action of atmospheric agents
on copper?

5. Name three alloys of copper and their constituents.

6. Heavy electric cables often have iron wire above them to
which they are fastened. Why ?

7. How could you tell whether or not a given substance
is a compound of copper ?

8. What would be obtained if ammonia were passed over
heated copper oxide ?

9. State two uses for copper sulphate.

10. How could you tell whether a given substance was gilt
(brass) or gold ?

CHAPTER XXIX.

SILVEE, GOLD, AND PLATINUM.

SILVER.

292. Occurrence. — Silver is the most common of the
precious metals. Silver was known and valued from the
earliest times, as it frequently occurs free in rocks and is
easily separated.

Native silver is found in Arizona, Mexico,^ South Amer-
ica, and elsewhere, but much of the silver now used is
obtained from sulphide ores, usually associated with lead,
copper, arsenic, and gold. • Silver chloride (horn silver)
occurs in nature, and traces of silver compounds are found
in sea water. The principal supply of silver is from the
United States, Mexico, Australia, and Germany.

293. MetaUurg^. — Some of the ores of silver are so
complex that various processes are employed in the sepa-
ration of the metal, but since so large a proportion of the
commercial metal is obtained from lead ores, only the
method used for these (Parkes' process) will be described.

The ore, largely lead sulphide, is roasted to remove
sulphur, and then reduced as described under the metal-
lurgy of lead (§ 349). The crude metal is heated in a
reverberatory furnace and stirred. Such metals as
copper, antimony, and arsenic are oxidized, forming a
scum on the surface of the lead, and this is skimmed off.
The molten metal is now run into iron pots and a small

304

METALLUBGY OF SILVER.

305

percentage of zinc is stirred into it. As the mixture
cools, an alloy of zinc with silver and gold conies to the
top and is skimmed off, but little of the lead being re-
moved. If there is much silver in the crude lead, the
treatment with zinc may be repeated.

The skimmings containing zinc, lead, silver, and gold
are now heated in a retort and the zinc removed by
distillation.

The residue, containing lead, silver, and gold, is then
heated, cupelled^ in a shallow furnace exposed to the air.
The lead oxidizes, and the melted lead oxide flows oflf and
is. recovered. The melted silver and gold remaining is
then poured into moulds. A small
cupellation furnace is shown in Fig.
110 ; a is the muffle and h the cupel in
which the silver or gold is finally left
as a metallic button.

The gold is separated from the silver
by treating the alloy with hot concen-
trated sulphuric or nitric acid. The
silver dissolves as the sulphate or the
nitrate, but the gold isnot affected, and
after washing is melted and cast into bars.

The silver is recovered from the solu-
tion by hanging in it plates of copper :

Fig. 110.

2 AgNOg + Cu — ^ CuCNOg)^ + 2 Ag

The silver is deposited in a fine crystalline form known
as cement silver.

Another method of separating gold from silver is by
electrolysis. The alloy is made the anode in a dilute
nitric acid solution of silver nitrate, the anode sheet being
enclosed in a canvas bag. With a current of low voltage,

306 SILVER, GOLD, AND PLATINUM.

silver is dissolved from the anode as the nitrate and rede-
posited on the cathode as practically pure metal. The
gold is unaffected and remains at the anode, and as the
anode disintegrates, collects as a mud in the canvas bag.

294. Physical Properties. — Silver is a white metal,
fairly hard, capable of receiving and retaining a high pol-
ish. It is the best conductor of heat and of electricity.
Being ductile and malleable it is readily worked into .
various shapes.

295. Chemical Properties. — Silver does not change in
air, but darkens readily in the presence of sulphur com-
pounds, showing such stains as are seen on silver spoons
that have been used with eggs or mustard, on coins carried
in the pocket, or on silverware about the house.

Silver does not oxidize on being heated. Alkalies do
not affect it. Nitric and sulphuric acids react with it as
they do with copper :

2 Ag + 2 H^SO^ — ^ AgjSO^ + 2U^0 + SO,
8 Ag + 4HN08 — ^ 3 AgNOg + 2 H^O + NO

296. Uses. — Since pure silver is not hard enough to
stand the wear and tear of constant use, it is alloyed with
other metals, for instance copper. The silver coins of the
United States contain 90% of silver and are said to be
900 fine. British coins are 925 fine, and this is the grade
known as sterling silver. On account of its durability
and lustre it has long been used for jewellery and orna-
ments. An amalgam of silver is often used by dentists,
and silver is used on the back of glass for mirrors, and as
a plating on cheaper metals.

Some mirrors are made by depositing a layer of silver on

COMPOUNDS OF SILVER.

307

the polished glass. A solution of silver nitrate to which
has been added some ammonia and a reducing agent is
flowed over the glass and gently warmed. The silver is
reduced and deposited as a bright film on the glass. This
is washed, dried, and varnished to protect it.

Silver plating is usually done by electrolysis. To
secure a firm, uniform deposit, the electrolyte is a solution
of silver and potassium cyanides
(Fig. Ill, 6) made by adding
potassium cyanide solution to a
solution of silver nitrate until the
precipitated silver cyanide is dis-
solved. A bar or sheet of silver is
used as the anode (a) and the
object to be plated as the cathode
((?), a rather weak current being
employed. The positive silver
ions are discharged and deposited
on the cathode. The negative ions,
discharged on the anode, combine with the silver, forming
silver cyanide. This, on dissolving, is dissociated. The
amount of silver in the solution is unchanged, for silver
is dissolved from the anode and deposited on the cathode.

Fig. 111. Silver plating.

—Silver nitrate^ AgNOg, is the
It is prepared by dissolving

297. Componnds of Silver,
most common compound,
silver in nitric acid :

3 Ag + 4 HNOg^ 3 AgNOg + 2 H2O + NO
It is very soluble in water and crystallizes from it in flat,
rhomboidal, transparent forms. In contact with organic
matter and exposed to the light, it darkens. Moulded
into sticks, silver nitrate is used as a cauterizing agent for
warts, wounds, and sores, and is known as lunar caustic.

308 8ILVEB, GOLD, AND PLATINUM.

Silver nitrate is the most important compound of silver
because most of the other silver compounds are made
from it.

Silver chloride, AgCl, is made by adding a solution of a
chloride to a solution of silver nitrate :

AgNOg + KCl— ^ AgCl + KNOg
The silver chloride separates as a white, curdy, insolu-
ble solid. Silver chloride does not dissolve in acids, but
dissolves in ammonia and in sodium thiosulphate, the
hypo of the photographer.

Silver bromide, AgBr, and silver iodide, Agl, resemble
the chloride ; they have a yellowish tinge, and are more
easily changed in the light and are less soluble. Like
the chloride, they are extensively used in photography.

298. Photography.— The preparation of the photographic
negative involves these processes: the exposure, the devel-
opment, the fixing, and the washing. The photographic
plate consists of glass or transparent celluloid coated with
a film of gelatine containing very finely divided silver
bromide, which, as we have stated, is sensitive to light in
that it becomes somewhat less soluble and more easily
reduced. In the camera the plate is exposed to light,
and the change in the silver bromide is produced, strongly
where the light is bright, less intensely in the shadows.
The exposure is very short in the camera, and produces no
visible change in the plate.

The next operation is to develop the picture. As the
exposed silver compound is a little more easily affected
than the unexposed compound, it is possible to change the
one without materially affecting the other. For this, the
developer is used. The developer is a reducing agent of
such strength that it is capable of continuing the change

PHOTOQHAPHY. 809

begun by the light, but is not capable of initiating the
change in the unaffected parts of the plate. Ferrous sul-
phate, pyrogallol, hydroquinone, and many other reducing
agents are used as developers to reduce the silver com-
pounds to metal :

2 AgBr + HjO — ^ 2 Ag + 2HBr + O

The acid is neutralized by the alkali added in the devel-
oper, and the oxygen removed by the reducing agent.
Where the plate has been exposed to the light, there will
be a deposit of silver, which appears (Jark because it is
very finely divided. Where no light acted, the-^ver
compound is unchanged.

When it is seen that the picture is developed sufficiently,
it is placed in the fixing hath. This is a solution of sodium
thiosulphate, NagSgOg, commonly called hyposulphite of
soda, and is capable of dissolving many silver compounds^
such as the silver bromide, which remains unreduced in
those parts of the picture where the light has acted least.
This unchanged silver bromide is dissolved, and the glass
remains clear in these places. As all the material sensitive
to light has been removed, the plate is said to be fia:ed; it
is then thoroughly washed and dried.

On the fixed plate, those parts of the scene which are
brightest, that is, those parts which are white or blue, are
represented by a dark deposit of silver; the dark parts of
the scene are clear, so that shades are reversed; hence it is
called a negative (Fig. 112, 6).

The finished picture on paper is made from the negative.
The paper is sensitized^ as was the plate, by a film of silver
chloride or bromide. It is exposed to the light under the
negative. Now those parts of the paper under the clear
parts of the negative will be affected most by the light and

810

SILVER, GOLD, AND PLATINUM.

will be the darkest on reduction; the parts under the heavy
deposits will be little affected and appear light, as they do
in the object, so that the print, being the reverse of the
negative, is a positive (Fig. 112, a), in that its shades agree
with those of the object.

The positive may be developed in the same manner as
the negative, but in many cases, as printing out papers, the
developer is in the paper, or film, so that the reduction
occurs and the picture appears during the exposure. The

a, Positive.

Fig. 1 12.

b. Negative.

print is fixed and washed as the plate was, and to render it
more permanent and to improve the color, the print is toned
by immersing it in a solution of gold chloride, so that some
of the silver of which the picture is composed is replaced
by gold, giving it a warmer tone. Platinum and lead
compounds are also used in toning. A melanotype or tin-
type is a whitened negative on a polished black surface.
Other materials might be used besides silver compounds,

OCCURRENCE OF GOLD. 311

but these are the most sensitive to slight variations in light
and are the most easily controlled.

Ferrotypes, or blue prints^ are the simplest substitutes
for silver printing papers. The paper is coated with a
mixture of a ferric salt, a reducing agent, and potassium
ferricyanide, which has been applied and dried in the
dark. On exposure to light, some of the ferric salt is re-
duced to the ferrous compound, so that when the paper is
immersed in water, an insoluble blue is formed where it has
been exposed to the light. Where it has been protected
from the light, the materials are unchanged, and are re-
moved in the washing.

GOLD.

299. Ooourrence and Separation. — Gold has been known
from the earliest times. It commonly occurs native,
or alloyed with silver and other metals. It also occurs
combined with tellurium, an element closely related to
sulphur. Native gold is found in veins running through
quartz rock and also in the beds of streams whose sands
have been formed from the disintegration of such gold-
bearing quartz. It often occurs in nuggets varying in
size from that of a tiny pebble to a mass weighing over
a hundred pounds.

From river sands and gravel gold is separated by wash-
ing with water, when the lighter rock particles are washed
off, leaving the gold. Partially disintegrated rock and
coarse gravels are sometimes mined by washing them
down with powerful streams of water and causing the
resulting torrent to flow through long flumes or troughs,
with transverse cleats along the bottom. These retain
the heavy gold and permit the soil and gravel to be swept

812 SILVER, GOLD, AND PLATINUM.

on. This last process, hydraulic mining, is forbidden in
many states on account of its destructive effects.

300. Metallurgy. — When gold occurs in veins in mas-
sive rock, the rock is blasted with dynamite. The broken
rock is crushed to small pieces by powerful iron crushers
and is then pounded into fine powder by heavy iron
stamps working in iron troughs. Water is kept flowing
through these troughs, and the gold and rock leave them
as a thin mud. This is caused to flow over silver-plated
copper plates coated with mercury. The mercury amal-
gamates with the gold^ and when a sufficient quantity
accumulates, the amalgam is scraped off the plates and
freed from mercury by distillation.

The gold that escapes amalgamation is extracted by
means of potassium cyanide. It is allowed to stand in a
weak solution of cyanide exposed to air for days or
weeks and a double cyanide of gold and potassium is
formed. The gold is precipitated from this solution by
zinc or is extracted by electrolysis. This cyanide process
is also applied directly to ores poor in gold and to
tellurides.

The separation of gold from copper slimes has already
been mentioned (§ 285).

301. Properties. — Gold is soft and heavy and is the most
malleable and ductile of metals. The presence of a small
amount of other metals, however, often makes it brittle.
Gold leaf has been made ^ g ^^^ ^ ^ of an inch thick. Gold
leaf transmits green light, while finely divided gold, when
suspended in liquids, appears purple by reflected light,
and blue by transmitted light.

Gold is unaffected by air or water at any temperature.

U8E8 OF GOLD. 818

Ordinary acids do not act on it, but it is dissolved by
aqua regia^ with the fonnation of auric chloride, AuClg.

302. Uses. — Pure gold is used as gold leaf. The metal
is too soft to be used alone for other purposes and is
alloyed with silver or copper. The proportion of gold is
always indicated by the number of carats fineness : pure
gold is 24 carats fine, 18-carat gold contains 18 parts by
weight of gold and 6 parts of other metals. The gold
coin of the United States is 90 % gold and 10 % copper.
Articles are gold plated by an electroplating process with
a bath of double cyanide of gold and potassium.

PLATINUM.

Platinum occurs native, alloyed with osmium and iri-
dium, which closely resemble it, and with other metals.
The most important deposits of platinum are in the Ural
Mountains; it is also found in California, Australia, and a
few other places. The separation of platinum from the
metals alloyed with it is a complicated process.

803. Properties. — Platinum is a white, lustrous metal,
about twice as heavy as lead. It is very malleable and
ductile and is infusible except in the oxyhydrogen flame
or the electric arc. It absorbs or occludes large quantities
of hydrogen when hot, with an increase in temperature,
and releases it on cooling. It does not absorb oxygen
when hot, but condenses it on the surface when cold.
It may be obtained as a fine black powder, platinum blacky
by the action of a reducing agent on a solution of one of
its salts and as a %pongy platinum by igniting the double
chloride of platinum and ammonium. Both of these
forms of platinum act as powerful catalytic oxidizing

814 SILVER, GOLD, AND PLATINUM.

agents on account of the large surface they possess in pro-
portion to their mass.

Platinum is not attacked by air or water at any tem-
perature, and is not affected by acids, except aqua regia.
Caustic alkalies, phosphorus, silicon, and carbon attack it
when hot, so none of these substances should be heated
in platinum vessels. Platinum should never be heated in
a smoky flame, on account of its tendency to form a
carbide, nor should metals be heated in platinum vessels.

304. ITses. — The inf usibility of platinum and its chem-
ical indifference toward the great majority of elements
and compounds render it invaluable in chemical opera-
tions. It finds extensive use in the laboratory, in dishes,
wire, and foil. On account of its cost, which is more than
that of gold, it is used only to a limited extent in chemical
manufactures. The great expense, however, is partly
compensated for by its indestructibility. It is used for
stills for the concentration of sulphuric acid, and large
quantities of platinum black are employed as a catalytic
agent in the manufacture of sulphuric acid by the contact
process. It is a good conductor of electricity and ex-
pands with heat at the same rate as glass. On account of
these properties it is used to connect the filaments of in-
candescent lamps with the wires outside the exhausted
bulbs. The power of platinum to cause the ignition of
inflammable gases mixed with air is utilized in self-light-
ing burners and mantles. Its alloy with iridium is hard
and unalterable in air and is used for the manufacture
of government standards of length.

305. Componnds of Platinum. — The most important
platinum compound is chlorplatinic add, HgPtCl^, which

SUMMARY. 815

is formed by dissolving platinum in aqua regia. This
forms chlorplatinates with metallic compounds. It is used
in toning platinum photographs and as a test for potas-
sium compounds. The test depends upon the fact that
potassium chlorplatinate is but slightly soluble in water
or alcohol, while the corresponding sodium compound is
decidedly soluble.

Summary.

Silver occurs native, and with other metals in complex sulphides
and as the chloride.

It is extracted from lead bullion by means of zinc, and purified
by electrolysis or treatment with add.

Silver has a specific gravity of 10.5. It is the best conductor
of heat and electricity.

* Silver is unaltered by pure air, but is tarnished by sulphur
compounds.

Silver is alloyed with copper for most uses. A double cyanide
of potassium and silver is used for electroplating.

Silver nitrate is made by dissolving silver in nitric acid. It is
the basis of other silver compounds.

The diver halides are insoluble compounds, and are made
more easily reducible by the action of light. Photographic plates
are coated with silver bromide, which after exposure to the Ught
can be reduced by a developer. The unreduced silver bromide is
dissolved by sodium thiosulphate.

Prints are made by exposing to Ught, under the n^ative, paper
coated with silver bromide or chloride. Toning is the replace-
ment of the deposited silver of the print by gold or platinum, by
means of a chloride solution.

Exercises.

1. Why does not silver occur as the oxide ?

2. Why is not silver commonly used as an electric con-
ductor ?

816 SILVER, GOLD, AND PLATINUM.

3. Does coating glass with silver or with tin amalgam pro-
duce better mirrors ? Discuss.

4. How could silver be cleaned that was so tarnished as to
resist ordinary silver polish ?

5. What is horn silver ? " hypo " ? sterling silver ? lunar
caustic ? oxidized silver ?

a Grive the action of the developer, the fixing bath, and the
toning solution in photography.

7. How much silver nitrate can be made from a dime weigh-
ing 2.45 grarbs ?

CHAPTER XXX.

ALTJMUniM AM) 'its OOMPOUITOS.

306. Occnrrence. — Aluminum never occurs in a free
state, although it is one of the most abundant and widely
distributed elements. Emery, corundum, ruby, and sap-
phire are more or less pure forms of aluminum oxide.
Clay and the rocks by the decomposition of which it is
formed consist chiefly of aluminum silicate. Two min-
erals of great importance in the preparation of metallic
aluminum are bauxite, a hydrate of aluminum, and cryo-
lite, a fluoride of sodium and aluminum.

307. Preparation. — The only process used at present
for the extraction of aluminum is an electrolytic one.
The electrolyte consists of a solution of aluminum oxide
in melted cryolite. The mineral bauxite is used to fur-
nish the oxide. The cryolite is fused and kept liquid by
the heat generated during the passage of the current ; the
aluminum oxide is dissolved by the cryolite and decom-
posed by the current. The aluminum collects as a
molten mass in the bottom of the melting pot ; the oxy-
gen is liberated at the anodes, which are oxidized by it.

The apparatus consists of a rectangular iron box, lined
with a thick layer of carbon which constitutes the cathode
(Fig. 113, J). The inside dimensions are approximately
4J feet long, 2J feet wide, and 6 inches deep. Carbon
rods about 3 inches in diameter and 18 inches long, placed
in rows and supported by copper rods, serve as the anodes

317

318

ALUMINUM AND ITS COMPOUNDS.

(Fig. 113, d). These are so arranged that they can be
lowered into the bath. The aluminum is allowed to run

off at the base from time
to time. The process is
made continuous by the
addition of fresh supplies
of bauxite as needed. A
shunt connects the circuit
with an incandescent
lamp. The resistance of
the bath increases as the
oxide is decomposed.
When more aluminum oxide should be added to the bath,
the current through the shunt becomes strong enough to
make the light glow brightly, and thus notifies the attend-
ant that the bath needs aluminum oxide. A layer of fine
coal, which covers the bath, facilitates the addition of the
oxide and shields the eyes of the workmen.

Fig. 113.

Physical Properties. — Aluminum is a silver- white
metal, capable of a high polish. The dull surface usually
seen is the result of a thin coating of the oxide. It is
lighter than any other of the common metals, having
about the same density as glass. It is malleable and
ductile, but not very tenacious. It ranks next to silver
and copper in thermal and electrical conductivity.

Chemical Properties. — Pure aluminum is practically
unaltered in air. When aluminum powder or foil is
strongly heated, it burns with a very brilliant light, re-
sembling that of magnesium, and liberating a great deal
of heat.

Aluminum is scarcely affected by nitric acid at any
temperature. Dilute sulphuric acid acts very slowly

USES OF ALUMINUM. 319

with aluminum, with the liberation of hydrogen. With
the concentrated acid it behaves somewhat like copper,
liberating sulphur dioxide. It reacts with hydrochloric
acid, forming aluminum chloride. Aluminum is also dis-
solved by sodium and potassium hydroxides, with the
formation of the corresponding aluminates and the liber-
ation of hydrogen.

2 Al. + 6 KOH — >- 2 KgAlOg + 3 Hj

310. Uses. — Aluminum has a wide range of uses, al-
though the marked influence of a small amount of impur-
ities has made its application more restricted than was
thought probable at one time. Powdered aluminum is
extensively used as a paint to protect other metals from
corrosion. Aluminum foil is replacing tin foil to a consid-
erable extent. Many small useful and ornamental articles
are made of aluminum. It is difficult to solder, so the
parts of the larger articles are commonly welded together.
Aluminum cooking utensils, when made of the pure metal,
prove very satisfactory. When plated with platinum,
aluminum proves useful for many purposes in the labora-
tory, as for evaporating-dishes. Aluminum has a very im-
portant use in removing the last traces of oxygen from
molten iron. When a mixture of aluminum powder and
an oxide of manganese, chromium, or iron is ignited, a
rapid combustion -and a very high temperature result.
This fact is utilized in softening iron for welding and in
the production of metallic manganese and chromium :

CrgOa + 2 Al -^ AlgOg -h2 Cr

Aluminum is being used to a considerable extent in
place of copper as an electric conductor. An aluminum
wire, though larger than a copper wire of the same con-

820 ALUMINUM AND ITS COMPOUNDS.

ducting power, is lighter and does not produce so great a
strain on its supports.

Aluminum forms alloj/8 with many of the metals. The
most important is the alloy with copper, called aluminum
bronze. It is hard, elastic, unaltered in air, easily cut, and
has a color closely resembling gold. It has been success-
fully used in place of steel for small objects, such as watch-
springs and ball-bearings. Magnalium is the trade name
for an alloy of aluminum with magnesium and other
metals. The metals used with aluminum and their pro-
portions vary according to the use to be made of the alloy.
It always contains 90% aluminum and less than 2% mag-
nesium. The tensile strength of magnalium is much higher
than that of aluminum, and it can be turned in a lathe. It
is less corroded by air than aluminum, copper, zinc, or brass.
On account of its superior strength this alloy is replacing
aluminum for many purposes.

31 1. Aluminum Oxide. — The occurrence of this compound,
AlgOg, as corundum and emery has already been men-
tioned. It is easily formed as a white, amorphous powder
by igniting the hydroxide. It may be prepared in crys-
talline form by fusing it with lead oxide. The ruby and
sapphire can be artificially prepared by adding small
quantities of other salts; potassium dichromate gives the
ruby color, and a trace of a cobalt compound produces the
sapphire blue. The latter fact is used in testing for alu-
minum. The substance to be tested is first strongly heated
with a blowpipe on charcoal to reduce it to an oxide. It
is then moistened with a solution of cobalt nitrate and
heated; the appearance of a blue color indicates that the
substance contains aluminum.

Emery, on account of its great hardness, is extensively

ALUMS. 321

employed as an abrasive for grinding and polishing. An
artificial corundum, made by fusing bauxite in an electric
furnace, makes a better abrasive than the natural emery
and is manufactured and sold under the name of alundum.

312. Alums. — The first alums known were double sul-
phates of an alkali metal and aluminum; now similar
compounds containing iron or chromium instead of alu-
minum are called alums. The most common alums are
potassium alum and ammonium alum; the composition of
the former is represented by the formula, KAl (804)3.
12 HgO; ammonium alum is NH^ Al (804)3 . 12 H^O.

On being heated, the alum gives up its water of crys-
tallization and is converted into a glassy mass known as
burnt alum. The alums are much more soluble in hot
than in cold water, and are deposited from a cooling solu-
tion in well-marked crystals, usually octahedral or cubical
in form.

Alum is used for a styptic and in the preparation of
aluminum hydroxide, for use in mordanting cloth and in
clarifying water.

ALUMINUM HYDROXIDE.

313. Preparation. — This is prepared by the addition of
ammonium hydroxide to a solution of an aluminum com-
pound. For example:

AI2 (804)3 + 6NH4OH— ^2 Al (0H)8 + 3(NH4)2 8O4

Calcium hydroxide may also be used. If sodium or
potassium hydroxide is employed, they react in excess
with aluminum hydroxide, forming aluminates and water :

3K0H + A1(0H)8— ^^KgAlOg -f 3H2O

322 ALUMINUM AND ITS COMPOUNDS.

Aluminum hydroxide is a starchy, translucent substance,
insoluble in water. It reacts with acids, with the forma-
tion of the corresponding aluminum salts ; with strong
bases, it has, as we have just seen, the action of an acid.
On heating, aluminum hydroxide is converted into alumi-
num oxide :

2A1(0H)3 — >- Al^Og -h 3 H^O

314. Mordants and Lakes. — The gelatinous character of
the hydroxide renders it valuable in dyeing and water
purification. Dyestuffs do not readily enter the fibre of
cotton goods, so it is necessary to use some substance to
cause the dye to adhere and prevent it from washing off.
Such a substance is called a mordant. It is found that
when aluminum hydroxide is precipitated in a solution
containing coloring-matter, the color is carried down by
the hydroxide as it settles, leaving the solution clear.
The mixture of the color and aluminum hydroxide is
called a lake; lakes are used in dyeing and as pigments.
In dyeing cotton, aluminum hydroxide is precipitated on
the fibre, either by soaking the cloth first in alum and
then in ammonium hydroxide solution, or by impregnating
it with aluminum acetate, which yields the hydroxide on
heating. The cloth, mordanted with aluminum hydroxide,
is soaked in the dye, which forms an insoluble lake with the
mordant, and thus produces a fast color. Other gelatinous
hydroxides, as those of iron and copper, are used as mor-
dants. Many mordants react chemically with the dye,
producing new shades. So it is possible, by the use of the
proper mordants, to secure different colors from the same
dye. This is done in calico printing, where the pattern
is first stamped with a mordant. When the cloth is passed
through the dye, the mordanted portions take it up and

CLAY AND POTTERY. 828

retain it, while the color is removed from the unmordanted
parts by washing.

315. Coagnlnm in Water Purification. — The use of alumi-
num hydroxide in water purification is similar to that in
dyeing ; it carries down with it suspended particles of
foreign matter. The hydroxide is produced in this case
by adding proper proportions of aluminum sulphate and
lime :

8Ca(OH)2 + Al^CSO^g— >-2Al(OH)8 + SCaSO^
The aluminum is completely removed by precipitation ;
the precipitate carries with it the greater part of the sedi-
ment and disease germs. The calcium sulphate is partly
precipitated and partly dissolved, adding to the hardness of
the water. In water containing a considerable amount of
calcium bicarbonate in solution, this salt is sufficient to pre-
cipitate the h3'^droxide ; in such cases aluminum sulphate
or alum is added to the water and the precipitation takes
place without lime.

ALUMINUM SILICATES.

316. Clay and Pottery. — Ordinary clay is an impure
silicate of aluminum formed by the decomposition of
felspathic rock. Felspar is a silicate of aluminum and an
alkali metal. When such rock is exposed to the action of
the weather, the alkali silicate is removed by the water
and carbon dioxide, and the residue left is kaolin^ or clay
mixed with other rock materials, as sand and mica. Pure
kaolin is a white, pulverulent mass ; when wet, it is plas-
tic and can be moulded. When the moulded clay is heated
and dried, it shrinks. Iron compounds often give clay a
red color, seen in some bricks and in roofing and drain
tiles. Light brick is made from clay containing little or

324

ALUMINUM AND ITS COMPOUNDS.

no iron. Clay containing silica *is used for firebricks,
stove linings, and crucibles.

Bricks, earthenware, porcelain, and china are made hy
moulding the clay into the desired form and baking in a
furnace or kiln. In making common earthenware articles,
the baking temperature is not very high and the mass is
porous, as in flower pots. In making roofing tiles, jugs,
and drain pipes, salt is thrown into the fire ; it volatilizes
and forms on the surface of the articles a glaze impervious
to water, consisting of a fusible silicate of sodium and
aluminum.

Stone ware, granite ware, and crockery are made from
purer varieties of clay and are more carefully moulded.
In addition to the clay, they contain some fusible, harden-
ing material. A heavier and more durable glaze is used
than that on earthenware.

Porcelain and china are made from pure kaolin, felspar,
and quartz. The materials are ground fine, thoroughly

mixed and wet. The wet
mass is then modelled or
moulded in plaster of Paris
moulds and dried ; when
dry enough to handle, the
mould is removed and the
article smoothed. It is
then fired at a low tem-
perature which leaves it
firm and hard, but porous
and ready for the glaze.
This consists of felspar
and quartz ground fine and suspended in water. The
article is dipped into the mixture and dried. It is then
placed in a Beggar or fire-clay box (Fig. 114, J) and sup-

Fig. 114. Pottery Kiln.
a, Fire-box ; b, Seggars.

CEMENT AND CONCRETE. 325

ported on a tripod in such a way that it shall not touch
the box. This is placed in the oven, the temperature
of which is raised to a red heat in from twelve to
twenty-four hours. The temperature is then increased
for three days or more, and then allowed to fall very
slowly. The mass is then hard, dense, white, translucent,
thin, and not easily affected by chemicals, except alkalies.
The colors that are used in decorating china are com-
posed of substances that are stable at the high temper-
ature necessary for fusion of the glass.

317. Cement and Concrete. — Hydraulic cement results
from the heating of a mixture of limestone and clay until
they just begin to melt together. Many natural limestones
contain the aluminum silicate mixed with the calcium car-
bonate in nearly the required proportions. Such cement
rocks, when burned in a manner similar to that used in
making quicklime, yield natural cements. Artificial ce-
ments are made by grinding clay or shale with marl or
with limestone, then burning and powdering.

Cement is believed to consist of a mixture of calcium
silicate and calcium aluminate. When mixed with water,
it forms a mass that solidifies, or setSy in a short time. It
differs from lime in the fact that it sets under water. After
it has set, it slowly increases in firmness and tenacity, re-
quiring months and sometimes years to attain its greatest
strength. This process is known as hardening. The
chemistry of the hardening of cement is not very well
understood, but it is thought to be due to the formation
of crystals by the absorption of water of crystallization.
The compounds in the hardened cement are probably a
hydrated silicate and aluminate of calcium, of different
composition from those found in the dry cement.

826 ALUMINUM AND ITS COMPOUNDS.

The chief use of cement is in making concrete, a mix-
ture of cement, sand, and broken stone. Concrete is used
in the construction of buildings for the foundations, floors,
and walls by pouring the wet concrete into troughs of
board, around twisted rods of steel set in the centre of the
trough and parallel to its sides (reenforced concrete).
It is indispensable for bridge piers and other structures
below the water. It is also extensively used in side-
walks.

STimmary.

Aluminum does not occur native. Its oxides and silicates are
found widely distributed.

It is prepared by the electrolysis of oxide of aluminum dissolved
in cryolite.

Aluminum is a silver-white metal ; specific gravity, 2.6 ; melting-
point, 660®. It is a good conductor of electricity.

Aluminum dissolves in hydrochloric add and in potassium
hydroxide.

It is used for making aluminum bronze, cables for conducting
electricity, paint, flashlight powders, foil, ornamental articles, and
cooking utensils.

Corundum, ruby, and sapphire are nearly pure aluminum oxide,
Emery i^ corundum mixed with iron.

Common (potash) cdum has the formula KAl (804)3 . 12 HgO.

Aluminum hydroxide possesses the properties of a base and
those of an acid.

It is used as a mordant and in the purification of water.

Clay is an impure silicate of aluminiun and is used in the manu-
facture of bricks and earthenware.

Kaoliny nearly pure aluminiun silicate, is used in the manufac-
ture of porcelain and china.

Hydraulic cement is made by heating to indpient fusion a mix-
ture of limestone and clay. Cement hardens under water.

EXERCISE8. 827

Bxerclses.

1. Would you carry on the electrolysis of sodium chloride
in an aluminum dish? Would you concentrate a solution of
nitric acid in an aluminum vessel ? Would you concentrate a
solution of potassium hydroxide in an aluminum cup ? Ex-
plain.

2. For what purposes is aluminum bronze used ?

3. Write an equation showing how aluminum hydroxide can
play the part of a base; of an acid.

4. How could you show that common alum contains alumi-
num?

5. What would be the weight of a piece of aluminum con-
taining a cubic foot ? A cubic foot of water weighs 62.5 lb.

6. How much calcium hydroxide would be required to com-
bine with 20 kilos of aluminum sulphate ?

7. How much aluminum is contained in 200 tons of alumi-
num oxide?

8. What compound of aluminum is formed when a solution
of calcium carbonate reacts with a solution of aluminum sul-
phate ?

9. How many liters of hydrogen would be liberated by the
addition of 9 grams of aluminum to an excess of hydrochloric
acid?

CHAPTER XXXI.
lEOH, COBALT, AHD HIOKEL.

318. Occurrence of Iron. — A consideration of the enor-
mous quantity of iron used annually for an almost un-
limited number of purposes, will show how impossible our
present civilization would be without this metal.

Native iron occurs in igneous rocks in pieces varying
in size from small grains to that of a mass found in
Greenland which weighed more than a ton. As iron
rapidly corrodes when exposed to moist air, native iron
is not of common occurrence but compounds of iron are
common. The red and yellow colors of soils are generally
due to oxides and silicates of iron. Nearly all meteorites
contain iron alloyed with nickel. Ferric oxide is found
very widely distributed.

319. Formation of Iron Ores. — When water percolates
through a soil containing much vegetable matter, it takes
up substances capable of reducing ferric compounds
to ferrous compounds. When water containing carbon
dioxide comes in contact with the ferrous compounds,
acid ferrous carbonate, FeH2(C08)2» is formed, which is
soluble. In this way iron is dissolved out of the soil.
If water containing acid ferrous carbonate collects in a
warm place not in contact with air^ carbon dioxide is
driven off and ferrous carbonate (siderite), which is
insoluble in water, may be deposited. When water con-

FORMATION OF IRON ORES. 329

taming acid ferrous carbonate is exposed to the air, fer-
ric hydroxide is formed.

4 FeHa(C03)a+ 2 H^O + O^ — ^ 4 Fe(0H)3 + 8 CO^

This may be deposited, and on becoming dry may lose
sufficient oxygen and hydrogen in the form of water to
convert it into hydrated ferric oxide (limonite), 2Fe308.
3 H3O, or into ferric oxide (hematite), Fe203:

4 Fe(0H)8 — ^ 2 FejOg. 3 H^O + 3 H^O
2 Fe (0H)8 — ^ FejOg + 3 H^O V

Siderite, FeCOg, limonite, 2 FcjOg . 3 H2O, •hematite,
FcjOg, and magnetite, FcgO^, are the most important
ores of iron. Pyrite (FeSj), commonly called fools' gold,
is also widely distributed. Large quantities of it are
roasted for the production of sulphur dioxide and some
iron is smelted from the residue.

320. Manufacture of Cast Iron. — Iron ores are smelted in
a blast-furnace which is operated in a manner similar to
that employed in the production of copper matte (§ 282).
An iron blast-furnace (Fig. 115) consists of a steel shell
from 75 to 90 feet in height, lined with a thick layer
of fire-brick. Cold water is made to circulate through
hollow castings built into the fire-brick just above the
tuyeres. This portion of the shell surrounds the part of
the furnace where the most energetic chemical action
takes place during the smelting. The tuyeres are pipes
through which powerful blasts of hot air are forced.
If hematite containing sand is the ore used, the reduction
is accomplished in the following way. The ore mixed
with the right quantities of limestone (as a flux) and
coke (as fuel and as reducing agent) is dropped in at the
top of the furnace from time to time. The heated blast of

330 IRON, COBALT, AND NICKEL.

air forced through the tuyeres comes in contact with the
burning coke and forms carbon dioxide, which is immedi-
ately reduced to carbon monoxide by the excess of carbon.
The carbon monoxide reduces the iron oxide to iron:

C + O2 — ^ CO2
CO2+C— ^2C0
FejOg + 3C0 — >- 2Fe + 3 COj

Simultaneously with the reduction of the ferric oxide,
some of the silica is reduced to silicon, and sulphur and
phosphorus are reduced. The iron takes up these ele-
ments and from 4% to 7% of carbon. This impure iron
settles to the bottom of the furnace and is drawn off
through a hole as a white-hot stream of molten metal and
cast into ingots called pigs. The product is known, as
pig iron or cast iron. The flux of limestone, mentioned
above, combines with the silica, alumina, and other earthy
impurities in the ore, forming double silicates of calcium
and aluminum known as staff, a glassy substance which
remains fluid at the high temperature of the furnace.
Both slag and the molten cast iron collect in the crucible,
the slag floating on the heavier iron.

The desirability of cheap and rapid production of cast
iron has brought the blast-furnace to its present perfec-
tion. The opening at the top through which the charging
is done is closed by a cup and cone arrangement (Fig.
115). The best modern furnaces sometimes have an air-
lock at the top, closed above and below by a cup and
cone. In such a furnace the charge can be let into
the air-lock, and then, after the opening at the top is
closed, can be allowed to drop into the furnace. This
method prevents the escape of gas during charging. The
gases produced during the smelting, which contain about

MANUFACTURE OF CAST IRON.

331

20% carbon monoxide, are conveyed away from near the
top of the furnace through a large flue. This furnace
gas is burned to supply heat for the air-blast and under
the boilers to generate steam for the engines which com-
press the air for the tuyeres.

Fig. 115. Blast Furnace for Manufacture of Cast Iron.

The crucible in which the molten iron and slag collect
is about 16 feet in diameter. Holes are drilled through
its sides for the removal of slag and iron. The process is
called tapping the furnace. As soon as the iron or slag
has been removed, the tap-hole is closed by a clay plug
which is instantly hardened by the heat. The slag is
tapped off about every two hours, and the iron every
four to six hours. From 70 to 120 tons of metal are
drawn off at each cast. If the iron is to be used in mak-
ing steel in works near by, it is often carried directly
there in large ladles ; otherwise it is run into pigs.

332 IBON, COBALT, AND NICKEL.

321. Comporition and Properties of Cast Iron. — If the
iron, after being drawn from the blast-furnace, is suddenly
cooled, a white cast iron is obtained. The carbon in this
is in chemical combination with the iron as iron carbide,
FcgC. When the molten iron is cooled slowly, most of the
carbon separates in the form of graphite, and the product
is known as gray cast iron. As iron carbide is a very hard
compound, white cast iron is harder than the gray variety.

Cast iron always contains considerable carbon (4% to
7%) and is the most easily melted form of iron. It ex-
pands when it passes from a liquid to a solid. Therefore,
when molten cast iron is poured into a mould, and allowed
to solidify, the metal readily takes the shape of the mould.
Phosphorus and sulphur make cast iron brittle. Cast
iron is not malleable and can be neither welded nor tem-
pered. It is used for casting articles, such as stoves and
ornamental iron work, which are not to be subjected to
shock. Wrought iron and steel are made from cast iron.

822. Manufacture of Wrought Iron. — Wrought iron is
prepared by heating cast iron, placed on a layer of ferric
oxide in a reverberatory furnace (Fig. 118, p. 353). Flames
from the burning fuel play over the bed of the furnace
upon which rests the cast iron and oxide. When the
mass becomes molten, the carbon, silicon, sulphur, and
phosphorus contained in the cast iron are oxidized. The
carbon, changed into the monoxide, bubbles off, while the
other impurities pass into a slag. The iron becomes pasty
because pure iron has a higher melting-point than impure.
The molten mass in the furnace is stirred or puddled and
the pasty iron is gathered into large balls called blooms.
These are removed from the furnace and nearly freed
from slag by a process of squeezing and working under a

PROPERTIES AND USES OF WROUGHT IRON. 333

trip-hammer. The iron is then rolled so as to give the
finished product a fibrous structure. The small portion
of slag left in the iron aids in the process of welding.

323. Properties and Uses of Wrought Iron. — Wrought iron
is the purest form of commercial iron. Good varieties
contain not more than 0.3% of carbon. Piano wire is a
form of wrought iron containing less than 0.1% of carbon.

When wrought iron is heated, it becomes plastic before
melting. When in this condition, two pieces on being
hammered together adhere firmly, provided some substance,
such as borax, is placed on the iron to dissolve the thin
coating of iron oxide which forms on the heated sur-
face. This process is called welding. Plastic wrought
iron can be hammered into various shapes, rolled into bars,
and drawn into wire. Wrought iron is tough and can be
bent or stretched without breaking. It can very easily be
converted into a temporary magnet. Wrought iron is
used to make anchors, chains, wire, and other articles which
are intended to withstand sudden and severe strains. As
wrought iron can be readily forged and welded, it is the
iron used by the blacksmith. It cannot be tempered.

3214. Bessemer Iron and Steel. — Much cast iron is con-
verted into Bessemer steel, or more properly Bessemer iron.
About fifteen tons of molten cast iron are poured into a
Bessemer converter (Fig. 116), which ig an egg-shaped fur-
nace built of wrought iron plates and lined with a thick
layer of refractpry material. The bottom is perforated
with holes so that streams of air can be blown through
the molten metal. The blast lasts from eight to ten
minutes, during which the oxygen of the air unites
with the silicon, carbon, and other impurities in the cast

334

IRON, COBALT, AND NICKEL.

Fig. 1 16. Bessemer Converter.

iron, leaving nearly pure iron. The heat of combustion
raises the temperature of the metal to a high degree, pro-
ducing what is known as the
"boil." Just as the iron
commences to burn the blow
is stopped, and the desired
amount of a cast iron called
spiegeleisen, which is rich in
carbon and manganese, is
added. The blast of air is
forced through the mass for
a short time to thoroughly
mix the ingredients. The
spiegeleisen furnishes the de-
sired amount of carbon, and the manganese unites with
any dissolved oxygen present. The oxygen is thus pre-
vented from escaping and producing blow-holes when the
mass cools. The manganese also improves the quality of
the metal. At the end of the process, the converter is
turned over and the contents poured into a ladle and cast.
If the pig iron contains suflBcient sulphur and phospho-
rus to materially injure the quality of the Bessemer iron, the
converter is lined with basic material. The calcium oxide
formed when this is heated unites with the phosphorus and
sulphur, thus removing them from the iron. The use of a
basic lining was discovered by Thomas and Gilchrist, and
the process is named for them. The slag produced, known
as Thomas slag, contains basic phosphates and is of value
as a fertilizer.

325. Steel by Open Hearth Process. — Much steel is now
made by the open hearth process. An open hearth fur-
nace (Fig. 117) has a large bed (JE) lined with fire-brick

OPEN HEARTH STEEL.

335

and sand (fl^) on which the charge is placed. Gas is used
as fuel and is heated before entering the furnace by pass-
ing through a checkerwork of hot fire-brick (-A). The
heated gas is passed into the furnace through a pipe (a),
while through another flue (6) air enters which has been
similarly heated. The object is to obtain a higher tem-
perature than could be obtained by burning cool gas. The
burning gas passes over the charge on the furnace bed and
the hot, gaseous product escapes through checkerworks

Fig. 117. Open Hearth Furnace.

(Oi 2>) which are a duplicate of those used to heat the gas
and air. One set of checkerworks is thus raised to a high
temperature by combustion products, while the other is
being cooled as it heats the gas and air about to enter the
furnace. About every twenty minutes the direction of
the gas and air is reversed by means of a system of valves,
so that gas and air pass through the recently heated
checkerwork while the flame from the furnace
through the one just cooled.

836 IRON, COBALT, AND NICKEL.

The furnace charge consists of scrap steel, pig iron, and
iron ore. The scrap steel is placed on the bottom in order
to protect it from the oxidizing action of the flame.
The manganese and silicon are oxidized by the flame,
while the iron ore is active in furnishing oxygen to
consume the carbon. The process consumes from eight to
twelve hours and is watched and controlled most carefully
by the operator. Samples of metal are repeatedly taken
from the furnace and examined to determine when the
impurities have been removed and the carbon has been
reduced to the desired amount. In case the pig iron con-
tains much phosphorus, a basic furnace lining can be used
as in the Bessemer process. The steel produced is of
much better quality than the metal made by the Bessemer
process and is suitable for the manufacture of connecting
rods, shafts, armor-plate, heavy ordnance, etc., where
great strength and ability to stand vibration are required.
It can be forged and tempered. The percentage of carbon
varies from 0.6 in soft steel to 2.0 in hard.

326. Cmcible and Cementation Processes. — Crucible steel is
made by heating the best wrought iron mixed with a pure
cast iron in a graphite crucible. When pure iron is heated
to a temperature above 850°, it is converted into an allo-
tropic form of iron which is capable of absorbing carbon.
The carbon at first forms a solid solution with the iron, but
after the carbon has reached 0.8 % of the mass, it separates
as iron carbide, FcgC, forming a homogeneous mixture
with the iron. In the manufacture of crucible steel, the
heating is continued until sufficient carbon has been ab-
sorbed from the crucible to make the desired quality of
steel. A very high-grade steel, valuable for the manu-
facture of edged tools, or springs, is produced in this way.

TEMPERING OF STEEL. 837

Steel is made in the cementation process by heating
wrought-iron bars, packed with fine charcoal, in compart-
ments of fire-brick. The operation is continued at red heat
for a week or more and the carbon slowly penetrates the
iron. The cementation process, however, is expensive,
and is being replaced by the cheaper crucible process.

327. Special Steels. — Certain of the less familiar metals
when added in small quantities produce steel of great
hardness, toughness, or tensile strength. More than 8^
of manganese in a steel, or chromium in smaller amount,
gives great hardness to the steel. Tungsten and molybde-
num are used for making tools of self-hardening steel,
which is made without tempering and which retains its
cutting edge for a long time. Manganese, chrome, and
nickel steels are used for safes, armor plates, and parts of
machinery subject to great stress or vibration, as shafts
or automobile bearings. Deposits of iron ore which had
been regarded as too refractory to work have been made
valuable by the demand for these special steels.

328. Hardness and Tempering of Steel. — The hardness of
steel depends not only upon the amount of carbon which it
contains, but also upon the manner in which it is worked.
In general, the more iron carbide (FcgC) a steel contaihs,
the harder and more brittle it is. If a hard steel, a homo-
geneous mixture of iron carbide and iron, is cooled slowly
from a red heat, the iron carbide begins to decompose, at
a temperature below 670^^ into iron and graphite, giving a
softer steel. If a steel containing iron carbide is sud-
denly cooled from about 670°, the carbide has not time
to decompose, and a hard steel is obtained. If such a steel
is now heated to a temperature less than that required for

338

IRON, COBALT, AND NICKEL.

the formation of iron carbide (670°), the latter gradually
decomposes into iron and graphite. Hence, the reheating,
provided it is done below 670°, brings about the same de-
composition as the slow cooling mentioned above. Heat-
ing for a long time at moderate temperatures, or quick
heating to the higher temperatures below 670°, accom-
plishes the same result, viz. the decomposition of the
iron carbide. The process, however, can be stopped at
any time by a sudden cooling, and a steel of the desired
hardness obtained. The temperature can be estimated by
a play of colors ranging from yellow to brown, red, pur-
ple, violet, and blue, to gray, which appear when the metal
is heated. These colors are due to the formation of thin
layers of iron oxide, causing interference colors. The
colors follow each other when the metal is cooled, and as
soon as the right color is obtained the process is stopped
by dipping the article into either water or oil.

COLOB.

TXUPBRATURE.

Stiesl ttskd fob

Pale yellow

430°-450°

Razors

Fall yellow

470°

Penknives

Brown

490^-610°

Shears and tools for brasswork

Purple

520°

Table knives

Blue

530°-570°

Watch-springs and sword-blades

Blue-black

610°

Saws and other wood-working tools

The yellow tints give very hard but brittle steels, while,
as we proceed toward the blue, the steel is softer but
tougher. The films are usually removed by grinding, but
are seen on some saws and springs.

329. Classification of Iron and Steel — Commercial iron is
commonly classified as cast iron, wrought iron, and steel.
Cast iron contains the highest percentage of carbon.

PUBM IRON. 339

wrought iron the least, and steel usually stands between.
Chemical composition, however, is not a sure way of dis-
tinguishing these forms, as some steels contain less carbon
than certain wrought irons. Moreover, the condition the
carbon is in, whether free or combined with the iron, is as
important a factor as its percentage. Knowledge gained
from a careful study of the structure of a polished and
etched section under a microscope is used more and more
as a guide to the true nature of an iron or steel.
This and some recent principles developed by physical
chemistry have put the iron industry on a scientific
basis. At best, the classification of iron is difiicult, and
unless we classify a product according to the process by
which it is made, we shall be confused by the many varie-
ties of steel and iron that grade into one another.

330. Pure Iron. — Pure iron may be prepared by the
reduction of pure iron compounds, as the oxalate, in a
stream of hydrogen ; or electrolytic iron may be deposited
from solutions of certain iron salts.

Pure iron is a white, lustrous metal which is very tough
and which fuses only at a high temperature. It is malle-
able, ductile, and may be temporarily magnetized.

In dry air pure iron does not rust, but in moist air
rusting proceeds rapidly, particularly if carbon dioxide
is present. Iron decomposes water very slowly at ordi-
nary temperatures, but at higher temperatures the reac-
tion proceeds rapidly. With cold, dilute acids hydro-
gen is evolved and a salt is formed. With hot, or more
concentrated acids, certain reduction products may be pro-
duced. When cast iron is dissolved in acids, the unpleasant
smell is due to the formation of hydrocarbons and to sulphur
and phosphorus compounds from impurities in the iron.

340 IRON, COBALT, AND NICKEL.

IRON COMPOUNDS.

331. Iron forms two kinds of ions, — bivalent, Fe+"^,
and trivalent, Fe"^"*""^. Bivalent ions result from the dis-
sociation of ferrous compounds, and trivalent ions from
the dissociation of ferric compounds. Compounds in
which iron has a valence of two are known as ferrovs
compounds; those in which the valence of iron is three
are termed ferrie compounds.

332. Oxides of Iron. — FerrouB oxide, FeO, can be pre-
pared by the reduction of ferric oxide with hydrogen or
carbon monoxide. It is a black powder which cannot be
kept in air on account of the ease with which it passes
into ferric oxide.

Ferric oxide, FcgOg, forms the most important ore of
iron. It can be readily prepared by heating ferric hy-
droxide, ferrous carbonate, or ferrous sulphide. Ferric
oxide constitutes the coloring matter of such pigments as
Venetian red, Indian red, and light red. When ferrous
sulphate is calcined, a form of ferric oxide known as rov^e
is obtained. Rouge is used for polishing and as a pigment.
Limonite, or hydrated ferric oxide (2 FcgOg . 3 HgO), is
found in nature mixed with fine clay and sand. Such
mixtures constitute the pigment yellow ochre. When cal-
cined, various shades of yellow, orange, and brown are
obtained. These are sold as raw sienna, burnt sienna,
raw umber, and burnt umber.

The magnetic oxide of iron, FcgO^, occurs in nature as
lodestone. It is formed when ferric oxide is heated to a
high temperature and when iron is burned in oxygen
or air. It constitutes what is known as blacksmith's
scale, formed when red-hot iron is worked. When steam

HYDROXIDES OF IRON. 341

is passed over red-hot iron, hydrogen is liberated, and a
firmly adhering film of magnetic oxide is deposited on
the iron. This film prevents the rusting of the iron
under it. Musdia iron^ used as a covering for locomo-
tive boilers, etc., is iron which has been artificially coated
with magnetic oxide of iron.

333. Hydroxides of Iron. — Ferric hydroxide can be formed
by the addition of a base to a solution of a ferric salt :

FeClg + SNH^OH — ^ Fe(0H)8 + 3NH^C1

It is a reddish brown compound, which, on drying,
changes to iron-rust. Iron-rvst may be considered as a
double compound of ferric oxide and ferric hydroxide.
It is probably produced in a manner analogous to that
described for the natural formation of hematite. Iron
dissolves in water or moisture from the air containing
carbonic acid, forming acid ferrous carbonate :

Fe + 2 H^COg — ^ FeHaCCOg)^ + H^

This substance, on drying and further oxidation, is con-
verted into iron-rust. If we represent iron-rust by the
formula FegOg . 2 Fe(0H)8, the chemical equation would be

4 FeH2(G08)2+ O2 — ^ FegOg . 2 Fe(0H)8+ HgO -h 8 COg

A coating of rust does not. protect the iron under it,
because the coating does not adhere, but scales off.

When ferric hydroxide is heated to a red heat, it is con-
verted into ferric oxide and water:

2 Fe(0H)8 — ^ FegOg + 3 H^O

Ferrous hydroxide^ Fe(0H)2, appears white when first
precipitated, but soon changes to dull green and then to
brown,, by oxidation.

342 IRON, COBALT, AND NICKEL.

334. Chlorides of Iron. — FerrouB chloride, FeCl^, is formed
when iron is heated in hydrogen chloride. It is also formed
when iron is dissolved in hydrochloric acid in the absence
of air. It can be crystallized as a pale green compound,
FeClg . 4 HgO. It is unstable in air.

Ferric chloride, FeClg, can be prepared by passing chlo-
rine over hot iron ; by dissolving iron in aqua regia ; and
by dissolving ferric oxide in hydrochloric acid. When
hydrogen peroxide is added to a solution of ferrous
chloride containing hydrochloric acid, ferric chloride is
instantly formed. Ferric chloride is used in medicine,
and in general whenever a soluble ferric salt is required.

335. Sidphates of Iron. — Ferrom sulphate, FeSO^ . 7 HgO,

is a by-product in many industries. Scrap-iron is added
to sulphuric acid that has been used in the refining of
petroleum, or for cleaning iron. The iron is dissolved,
and on evaporating the solution, crystals of green vitriol,
sometimes called copperas, are obtained. Much ferrous
sulphate is used in the manufacture of blue pigments, as a
mordant, in the preparation of black inks, and for the
precipitation of gold from solutions of its chloride.

When solutions of ferrous sulphate and tannic acid are
mixed, ferrous tannate, a nearly colorless compound, is
formed. On exposure to the air, ferrous tannate is
changed to ferric tannate, which is insoluble, and has a
black color. Ferrous tannate is the chief ingredient of
iron inks. As the writing would not at first be visible if
a solution of pure ferrous tannate were used, some dye is
added to give the ink color. On exposure to the air, fer-
ric tannate is formed, which gives the ink its permanent
black color.

Ferric sulphate, Yq^(^0^^, is formed when ferrous sul-

FERR0CYANLDE8. 843

phate is oxidized in the presence of sulphuric acid. It is
used with ammonium sulphate in the preparation of ferric
ammonium alum, NH^Fe(S0^)2 . 12 HgO.

336. Ferrocyanides. — On igniting a mixture of nitroge-
nous organic matter, scrap-iron and potassium carbonate,
cooling the mass, and then treating it with hot water,
potas9iu/m ferrocyanide^ K^Fe(CN)g, passes into the solu-
tion. When the solution cools, crystals of ferrocyanide, or
yellow prussiate of potash, separate. These have a com-
position represented by the formula K^Fe(CN)g . 3 HgO.

A solution of potassium ferrocyanide does not give the
usual characteristic reactions of iron because the iron
exists as a part of a complex ion, Fe(CN)g"''"''"*"'*". When
solutions of ferric salts and potassium ferrocyanide are
brought together, a deep blue precipitate, ferric ferro-
cyanide, or Prussian blue, is formed. The reaction is
made use of in testing for ferric ions (Fe"*"*"*") :

4FeCl8 -h SK^FeCCN)^ — ^ Fe4[Fe(CN)e]8 + 12KC1

Prussian blue is an important pigment. Blueing, used
for laundry purposes, often contains Prussian blue. When
the clothes with which it is used are not thoroughly freed
from soap, the alkali of the soap decomposes the ferric
ferrocyanide, precipitating ferric hydroxide on the cloth
and producing spots of iron-rust :

Fe4(Fe(CN)e)8+12NaOH— ^3Na4Fe(CN)e-h4Fe(OH)8

337. Ferricyanides. — Potassitim ferricyanide^ or red
prussiate of potash, K3Fe(CN)g, can be prepared by
treating a solution of potassium ferrocyanide with
chlorine :

2 K^Fe(CN)e + Cl^ — ^ 2 KCl + 2 KgFeCCN)^

344 IBON, COBALT, AND NICKEL.

Solutions of potassium ferricyanide, when added to those
of ferrous compounds, give a blue precipitate, ferrous fer-
ricyanide, or TurnbuU's blue :

SFeCl^ + 2K3Fe(CN)g.— ^ 6KC1 + Fe8[Fe(CN)e]2
This compound constitutes the color of the blue-print.
When a solution containing ferric chloride and oxalic acid
is exposed to the sunlight, the ferric salt is reduced to a
ferrous salt :

2FeCl3 + H2C2O4 — ^ 2CO2 + 2HCH. 2FeCl2

When a sheet of paper is coated in a darkened room with
such a mixture, and then dried and exposed under a
negative to the sunlight, the greatest reduction will take
place where the light is brightest. On covering the ex-
posed paper with a solution of potassium ferricyanide,
Turnbull's blue will develop wherever ferrous iron exists,
and the depth of color will be proportional to the amount
of ferrous salt present. In other words, potassium ferri-
cyanide is in this case used as a developer. The picture
can be fixed by washing away the ferric chloride and the
excess of potassium ferricyanide.

Blue-print paper is generally coated with a mixture
of potassium ferricyanide and ammonium ferric citrate.
Such a paper, after exposure, is developed and fixed by
washing with water.

COBALT AND NICKEL.

Cobalt and nickel resemble iron in being attracted by
a magnet. Their chemical properties are also like those
of iron.

338. Cobalt Ores and Extraction. — Cobalt is found as a
minor constituent of ores of complex composition. They

PROPERTIES OF COBALT. 345

are usually sulphides or arsenides, in which iron is the
predominating metal, but which contain copper and nickel
as well. Cobalt spiess, Co Asg, is found in Saxony; cobalt-
glance, CoAsg' CoS2,in Norway and Sweden. The ores
are usually worked up to obtain cobalt compounds with-
out separating the element in the metallic state. They are
first roasted to remove sulphur and arsenic, and the re-
sulting oxides are then dissolved in acids.

339. Properties of Cobalt. — Cobalt is a hard metal, mal-
leable and ductile, and capable of receiving a high polish.
Its melting-point, like that of iron, is high. It dissolves
readily in nitric acid, but is acted on slowly by other acids.
Solutions of cobalt salts have a rose color. No practical
uses have yet been made of metallic cobalt.

340. Gompoimds of Cobalt. — The chloride, CoClg, and
the nitrate, Co(N08)2, are of some importance. The
chloride has a peculiar property of changing its color
when exposed to air of varying humidity. These changes
are explained by the fact that the substance forms a
number of different compounds with varying amounts
of water of crystallization which pass readily into one
another. The less hydrated forms are blue or lavender,
while the more hydrated are red. Heated or exposed to
dry air, the red salt loses water of crystallization, and is
changed to a blue, less hydrated form. Advantage is taken
of this fact to make " sympathetic ink," which is invisible
until heated, and for simple apparatus to detect the amount
of moisture in the air.

Cobalt nitrate^ Co(N03)2, is sometimes used in analytical
work in testing for certain metals. It unites with certain
metallic oxides, forming characteristically colored com-

346 IRON, COBALT, AND NICKEL.

pounds. Thus aluminum compounds, when converted into
the oxide by heating with the blowpipe, give a blue
coloration when further heated with cobalt nitrate solution.
Cobalt sulphide, CoS, black in color, is precipitated from
alkaline solutions of cobalt salts by hydrogen sulphide.
Like iron, cobalt forms two double cyanides with potas-
sium, K^CoCCN)^ and KgCoCCN)^,

341. Nickel Ores and Extraction. — Nickel is nearly always
a constituent of meteoric iron. The greater part of the
nickel that the world uses comes from the province of
Ontario in Canada, and from New Caledonia. The nickel
compounds in the ores form only a small part of the whole.
The Canadian ore is chiefly a sulphide of iron, containing
about 2f> each of nickel and copper. The treatment of the
ore in its first steps resembles that used in the separation
of copper from chalcopyrite. After the iron is removed
by Bessemerizing, a matte rich in copper and nickel is
obtained. This is roasted and the resulting oxides are
reduced with charcoal. An alloy of copper and nickel is
thus obtained. The two metals are separated by an
electrolytic process.

342. Properties of Nickel. — Nickel is a hard metal, mal-
leable, possessing a high melting-point, and resembling
silver in color. It is capable of receiving and retaining
a very high polish. Dry air does not attack it. Like
cobalt, it dissolves readily in nitric acid, but is only slowly
attacked by hydrochloric and sulphuric acids. Solutions
of nickel salts have a beautiful, characteristic green color.

343. Uses of Nickel. — Nickel is of considerable practical
importance because of its silver-white color and the fact
that it does not readily tarnish in air. It is chiefly used

COMPOUNDS OF NICKEL. 347

as a covering for other metals. It is deposited by an
electrolytic process similar to that used in silver or copper
plating. Nickel is a constituent of several important
alloys. Nickel steely which contains about 5% nickel, is
both hard and tough ; it is used in making armor-plates
for battle ships. Nickel coins contain about one part
nickel to three parts copper.

344. Compounds of Nickel — The sulphate, NiSO^, and a
double sulphate of nickel and ammonium are the salts
used as electrolytes in nickel-plating. A compound of
unusual character, known as nickel carbonyl, Ni(CO)^, is
formed when carbon monoxide is kept in contact with
nickel at a temperature of about 30^. In one method of
separating nickel from copper, advantage is taken of this
reaction.

Summary.

The principal ores of iron are the oxides, hematite and magne-
tite.

Ccut iron is made by reducing the ore in a blast-furnace with
coke as fuel; the earthy impurities pass into a slag.

Wrought iron is produced by burning out the carbon and im-
purities of cast iron in a reverberatory furnace.

Bessemer iron is made by blowing air through molten iron to
remove carbon and impurities, and then adding spiegeleisen to
furnish the desired percentage of carbon and manganese.

Sted is produced by melting together wrought iron and cast
iron, or by dissolving carbon in wrought iron.

Iron forms two series of compounds. Ferrous compounds usually
oxidize easily to ferric.

Ferric oxide is used as ore and as pigment.

Magnetic oxide is an ore. It is often produced artificially on
iron by the action of steam, to protect the metal.

348

IRON, COBALT, AND NICKEL.

Iron-rust is a hydrated oxide resulting from the action of carbon
dioxide and air on iron in the presence of moisture.

Ferric chloride is used in medicine.

Ferrous sulphate is used as a disinfectant, and as a reducing
agent and in the production of paint.

Potassium ferrocyanide is used in the production of Prussian
blue.

Cast Iron.

Steel.

Wrought Iron.

Carbon,
per cent

Structure

Melting-
point

Specific
gravity

Magnetic
properties

Hardness

Elasticity

Use

4 to 7

Crystalline
1200°

7.2 to 7.5

Difficultly
magnetized

Hard

Elastic and
brittle

Casting and
rigid structural
purposes

0.8 to 2.8

Granular or
fibrous

1400°
About 7.7

Retains
magnetism

Hard (if
tempered)

Elastic and
tenacious

Tools, springs

Less than 0.25

Fibrous
1600°

7.8 to 7.9

Temporarily
magnetized

Soft
Tenacious

Blacksmith's
iron, and wire

Bessemer iron resembles wrought iron in most of its properties,
and is used for bars, rails, and structures.

Open hearth steel is used particularly for structures subject to
vibrations.

SXERCI8E8. 349

Ezercises.

1. What becomes of the ashes which would ordinarily result
from the combustion of coke when the coke is burned in the
blast-furnace?

2. Why is it necessary to produce a slag in the working of
a blast-furnace ?

3. Why does the temperature of the converter rise when
cold air is blown through it ?

4. For what purposes is Bessemer iron and steel inferior to
crucible steel? Why?

5. Why are not cast-iron pillars used in the upper stories of
large buildings ?

6.. How is iron protected against corrosion?

7. Why was wrought iron probably the first form of iron
worked by man ?

a Iron is one of the most abundant elements. Why does
noi iron occur to any extent in the free state ?

9. What kind of iron would you use for the manufacture
of sheet iron ? dynamo cores ? nails ? magnets ? car wheels ?
steam radiators ?! ^

! ' , ■

CHAPTER XXXII.
TDT AND LEAD.

Tin was one of the earliest metals known. The Phoeni-
cians obtained it from the British Isles, which they called
Cassiterides, land of tin. As a constituent of bronze it
was used before iron.

345. Metallurgy of Tin. — Tin oxide, SnOj, is the only
available ore. The present commercial supply comes
from England, Germany, Australia, and the islands Bil-
liton and Banca, near Sumatra. Tin oxide is reduced by
heating it in a reverberatory furnace with coal:

SnOa + C— ^Sn + COa

The molten metal which collects at the bottom of the
furnace is drawn off and cast into ingots, known com-
mercially as block tin. It is purified by heating it on the
inclined hearth of a furnace. The less easily melted im-
purities remain, while the easily melted tin flows down
the hearth. It is further purified by being poled, in the
same way as blister copper.

346. Properties of Tin. — Tin is a white, lustrous metal,
capable of withstanding the ordinary atmospheric agents.
Being soft and malleable, it can be. cut and hammered.
Like zinc, it is crystalline in structure, and if a bar of
tin is bent, it makes a peculiar noise (tin cry), probably
caused by the friction of the crystals. Like zinc, its

360

USES OF TIN. 351

physical properties vary considerably with the tempera-
ture. It melts at a rather low temperature, and burns,
forming a white oxide.

With acids tin does not react like any ane of the other
metals ; with hot, concentrated hydrochloric acid, it forms
stannous chloride^ SnClj ; with sulphuric acid it reacts like
copper ; nitric acid oxidizes it to a white, insoluble solid
known as metastannic acid.

Tin can be separated from solutions of its compounds
as a gray, spongy mass, by immersing a strip of zinc in
the solution :

SnClj + Zn — >-Sn + ZnClj

347. Uses of Tin. — The resistance of tin to ordinary
corrosive agents is utilized in protecting other metals by
covering them with a layer of tin. Ordinary tinware is
sheet iron, which has been thoroughly cleaned and dipped
into melted tin. Copper vessels and brass pins are sim-
ilarly treated.

Tin foil is tin, hammered or rolled into thin sheets ;
cheaper grades contain some lead. Tin pipes are used
to convey soda water and beer from the tanks to the
faucet.

Many common alloys contain tin. Bronze contains copper,
tin, and often zinc. The one cent piece is bronze. Pew-
ter and solder contain tin and lead. Britannia metal and
white metal contain varying proportions of tin, antimony,
and copper. Anti-friction and fusible metals often con-
tain considerable tin.

348. Compounds of Tin. — Stannous chloride^ formed by
the reaction of tin and hydrochloric acid, is the only com-
mon compound. The hyd rated salt, SnClg. 2 H2O is tech-
nically known as tin crystals^ and is extensively used in

352 TIN AND LEAD.

mordanting. It produces more brilliant shades than the
aluminum compounds. Stannous chloride is a strong
reducing agent in acid or in alkaline solutions. Ferric
salts are reduced by it to ferrous compounds :

2 FeClg + SnClj— ^2 FeCla + SnCl^

Mercuric compounds are first reduced to mercurous
salts, and with an excess of the reagent, to metallic mer-
cury :

2 HgCla + SnCl^— ^ 2 HgCl + SnCl^

2 HgCl + SnCla— ^ SnCl^ + 2 Hg

The stannic chloride^ SnCl^, is a colorless, fuming liquid,
which is readily decomposed by water.

Thus the valence of tin may be two or four, as shown
by the existence of stannous and stannic compounds.
Stannous sulphide^ SnS, is a brown, insoluble compound.
Stannic sulphide^ SnSj, is a yellow, insoluble solid used
as a pigment.

LEAD.

Owing to the wide distribution of its compounds and
the ease of separation from its ores, lead has been used by
man from the earliest times.

349. Metallorgy. — The most common ore is the sul-
phide, galena, PbS, large deposits of which are found in
Missouri, Illinois, and Colorado. The method employed in
the extraction depends largely upon the purity of the ore.

Ores having a large percentage of lead are roasted in
a reverberatory furnace (Fig. 118) until part of the sul-
phide has been oxidized, forming lead oxide, sulphur di-
oxide, and some lead sulphate :

2 PbS + 3 O2— ^2 PbO 4- 2 SOa
and PbS + 2 O^ — ^ PbSO^

METALLURGY OF LEAD.

353

When the oxidation has proceeded far enough, the air is
shut off by closing the doors, and the mixture is heated to
a higher temperature. The remaining lead sulphide now
reacts with the lead oxide and sulphate, forming lead and
sulphur dioxide :

PbS + 2PbO — ^3Pb + SO3
PbS + PbSO^— ^ 2 Pb + 2 SO2
The lead is moulded into ingots known as pig lead. When
there is a considerable amount of precious metal in the
lead, it is known
SiB base bullion. The
working of this has
been described un-
der sUver (§ 293,
Parkes' process).

Ores poor in lead
are reduced in a
blast-furnace simi-
lar to that used for
copper; indeed,
they may be sepa-
rated from the ore
at the same opera-
tion, the heavy lead settling beneath the matte and slag.

Electrolytic reduction of galena is effected in a bath of
dilute sulphuric acid. The crushed galena is made the
cathode, the bottom of the pan the anode. The lead is
obtained as a spongy mass. The hydrogen sulphide pro-
duced is conducted away to a combustion chamber and
converted into sulphuric acid or sulphur.

350. Propertiefl of Lead. — Lead is a soft, bluish white
metal. The brilliant lustre, apparent when freshly cut,

Reverberatory Furnace.

354 . TIN AND LEAD.

soon disappears, owing to the formation of a thin film of
oxide. This coating, however, protects it from further
change. Lead is not very tenacious, but being soft it can
be rolled into sheets or forced through a die to form
pipe.

When heated in air, lead oxidizes. Cold hydrochloric
or sulphuric acid have very little effect on it. Nitric acid,
acetic acid (from vinegar) and many vegetable acids dis-
solve it, forming soluble salts. Water containing carbon
dioxide corrodes lead, hence the objection to lead water-
pipes which would be affected by such water.

All lead compounds are poisonous, and if taken into the
system cause serious illness. Even minute quantities in
the water will ultimately produce this result, for lead
compounds are excreted with difficulty, and therefore
accumulate in the body. Painter's colic is a form of
chronic lead-poisoning.

On immersing a strip of zinc in a solution of lead salt,
the lead separates in a characteristic crystalline deposit,
the lead-tree :

Pb(N03)2 + Zn_^Zn(N08)2 4- Pb

The formation of insoluble chrome yellow by addition of
potassium dichromate to a solution of a lead salt is another
characteristic property of lead salts :

Pb(N08)2 + KjCrO^— ^ PbCrO^ + 2 KNOj

351. Uses of Lead. — Lead is very extensively used for
pipes and as a sheathing for cables, as it is easily cut, bent,
or soldered. Lead pipe is now made by forcing the hot
lead through a die by a piston. The opening of the die
is partly obstructed by a solid cylindrical rod attached to
the upper surface of the piston. The rod moves upward

USES OF LEAD. 365

with the piston, and the pipe is formed by the lead being
squeezed out between the rod and the wall of the die.

Sheet lead was formerly used for roof covering much
more than at present. It is very widely used as a lining
for tanks, cisterns, and cells used in electrolytic opera-
tions. The Chinese have long used it for lining tea-chests.

Thin sheet lead, alloyed with tin, is often used instead of
pure tin-foil. Tt/pe metal contains lead with tin and an-
timony, which harden it and form an alloy that expands in
solidifying. Thus it fills the moulds and makes a clear-cut
type. Solder and fusible metals are largely lead and tin.
Such alloys are forced through a die in the same manner as
lead pipe, forming a wire used as fuse wire.

Large quantities of lead are used in the manufacture of
shot. As already stated, the shot contains a small amount
of arsenic. The molten metal is run into a perforated
vessel, and falls in streams a long distance into the water.
In falling, the streams separate into drops which solidify
before they reach the water. The sizes of shot are as-
sorted by allowing them to run down inclined planes or
screens of different meshes. The smallest shot fall through
the nearest (smallest) openings into the bins, the larger
shot going on to the larger holes. Irregular shaped pieces
will not roll well, and are finally pushed off at the end.
The shot are polished by tumbling them in a barrel or
drum with a little graphite.

COMPOUNDS OF LEAD.

352. Oxides. — Lead oxide^ PbO, is known as massicot
when of a yellowish tint ; when it solidifies from the molten
state it is buff-colored and crystalline, and is known as
litharge. The presence of bismuth sometimes gives the

356 TIN AND LEAD.

litharge a yellowish color. Litharge is made by heating
lead in the air. Considerable quantities are produced in
the cupellation of silver. It is largely used in the prepa-
ration of oils and varnishes, of glass and glazes, and of
other compounds of lead. A mixture of litharge and glyc-
erine is used as a cement, especially for stone and glass.

Red lead^ or minium, is a bright red powder, known as
American vermilion. It is prepared by heating lead or
lead oxide in the air, oxygen being absorbed in the opera-
tion. The tint and composition often vary with the ma-
nipulation. Its composition may be represented by the
formula: PbgO^ or (2 PhO . PbOg). It is used in making
flint glass and as a pigment, especially on ironwork. Be-
ing an oxidizing agent, it hastens the hardening of the oils
used in paint. On this account a mixture of red lead and
oil is used by plumbers and gas-fitters to make tight joints.

Lead dioodde^ PbOg, is a brown powder obtained by
treating red lead with nitric acid. It is used as an oxidiz-
ing agent on the negative plates of storage batteries.

353. White Lead. — Bcmc lead carbonate^ white lead, is
a heavy, white, opaque powder. It mixes well with lin-
seed oil and forms a valuable paint base. The body of
many paints is white lead, which furnishes opacity or hody^
the tint being furnished by admixed color. Owing to the
importance of white lead, many methods have been devised
for its production. ^

The Dvtch process of corrosion has been in use three hun-
dred years, and although details have been improved, re-
mains essentially the same. Ridged and perforated disks,
or "buckles," of lead (Fig. 119) are piled on a shelf in a
loosely covered earthenware pot, the lower part of which
contains a little dilute acetic acid. Such pots are placed

WHITE LEAD.

357

side by side and covered with tan-bark ; other layers of
pots are added to a considerable height. The decaying
mass generates heat and
carbon dioxide. The acetic
acid is volatilized, forming
basic lead acetate. The car-
bon dioxide resulting from
the fermentation changes
this to the basic carbonate.
Three or four months are
required for the complete

Fig. 119.

corrosion of the lead; the right-hand portion of Fig. 119
represents a jar broken open to show the lead buckles
after corrosion. The white lead is removed from the jars,
washed, ground, and sifted. Small pieces of unaltered
lead are removed by the sifting. The white lead is then
ground in oil and is ready for use. The corrosion process
requires much time but yields good paint.

In a recent electrolytic process a copper cathode and a
lead anode are placed in a solution of sodium nitrate.
The electrolyte furnishes sodium hydroxide at the cath-
ode and nitric acid at the anode. The nitric acid
reacts with the lead, forming lead nitrate. This lead
nitrate reacts with the sodium hydroxide, again forming
sodium nitrate and an insoluble lead hydroxide which
settles out, so that the process is continuous, only the
lead needing renewal:

CU--H2O-; Na^-NOg- -.-- Pb+

NaOH

-PbCNOe)^

NaNO„

Pb(0H)2

358 TIN AND LEAD.

The lead hydroxide is treated with sodium bicarbonate,
forming white lead of a different composition from that
made by the Dutch process.

2 Pb(OH)2+2 NaHCOg — ^

Pb(0H)2 . PbCOg -h Na^COg + 2 H^O

The process is rapid and is said to give a paint of good
covering power.

The Dutch process aims at a white lead with the com-
position 2 Pb(0H)2 . PbCOg, but the composition of its
product varies.

354. Chrome Yellow. — Lead chromate^ PbCrO^, is an in-
soluble, bright yellow powder, prepared by mixing solu-
tions of lead salts and chromates. It is used in dyeing and
painting.

Suxxinieay.

The chief ore of tin is the oxide, which is reduced by heating
with coal.

Tin is soft, malleable, and crystalline. Its specific gravity is
7.3, and it melts at 232° C.

It is unaltered by air at ordinary temperatures.

Tin is used as foil and as a coating for iron. It is a constitu-
ent of bronze, pewter, and white metal.

Stanrums chloride is formed by the action of hydrochloric add
on tin. It is a reducing agent.

Lead occurs chiefly as a sulphide. The ore is reduced in a
reverberatory furnace, or by electrolysis.

Lead is soft, malleable, and tenacious. Its specific gravity is
about 11.3 and its melting-point 327° C.

It oxidizes in air and dissolves in nitric and acetic adds. Water
containing carbon dioxide corrodes it, produdng poisonous com-
pounds.

EXERCISES. 359

Lead is used for pipe, as a lining and covering material, in
type-metal, solder, and shot.

Lead oxides are made by heating lead in air. They are used in
making varnishes and glass, and as pigments. Lead dioxide is
used in storage batteries.

White lead is basic lead carbonate and chrome yellow is lead
chromate.

Ezercises.

1. Which forms a better protective coating for iron, tin or
zinc?

2. State the relative advantages of lead and tin-plate as a
coating for roofs.

3. Why were lead and tin early obtained in the metallic
state?

4. Why is tin-foil superior to lead-foil for wrapping articles
of food ?

5. Why is arsenic put in shot ?

6. What is litharge ? red lead ? white lead ?

7. Why is red lead called a drier in paints and varnishes ?

8. What advantage has zinc-white over white lead as a
paint base ? white lead over zinc-white ?

9. What are the characteristic properties of glass contain-
ing lead ?

10. How much lead can be extracted from a ton of galena ?

CHAPTER XXXIII.
MANGANESE AND OHBOHmH.

355. Preparation and Properties of Manganese. — The most
important ore of manganese is pyrolusite, which is crude
manganese dioxide. The metal is obtained by igniting a
mixture of pyrolusite and aluminum powder. Heat is
applied at one point and the action spreads through the
whole mass :

3Mn02-h4Al— ^2Al208 + 3Mn

Manganese is a hard metal resembling steel in appear-
ance. It oxidizes in moist air and, when finely divided,
decomposes boiling water. It dissolves readily in sulphuria
and hydrochloric acids, with the liberation of hydrogen and
the formation of the corresponding salt of the manganous
ion, Mn^^:

Mn + H2SO4 — ^ MnSO^ + Hj

Its alloys, ferro-manganese and spiegeleisen, are used in
the production of Bessemer iron and steel.

356. Manganese Componnds. — Manganese forms several
oxides, of which the most important is the dioxide^ MnOj.
This is a hard, black solid which conducts electricity. It
is a powerful oxidizing agent, as we have already seen in
the preparation of chlorine from hydrochloric acid (§ 54).
Its conducting power and oxidizing action make it a valu-
able depolarizer in voltaic cells.

360

MANGANATES AND PERMANGANATES. 361

The manganous salts are stable compounds, whose
water solutions are pink. The salts of the manganic ion
]yju++++ ape unstable.

357. Manganates and Permanganates. — In addition to
the salts in which manganese occurs as a positive ion,
there are others in which it is a constituent of a negative
ion. The most important of these salts are the manga-
nates and the permanganates, both of which contain the
radical MnO^. In the manganates the ion is bivalent,
Mn04""; in the permanganates it carries only one charge,
MnO^". Corresponding to this difference in valence of
the ion are differences in the properties of the salts.
This can best be shown by comparing potassium manga-
nate and potassium permanganate.

Potassium manganate is made by fusing a manganese
compound with potassium hydroxide in the presence of
air or an oxidizing agent, dissolving the residue and
evaporating the solution in a vacuum. The crystals ob-
tained are dark green. They are decomposed by water,
with the liberation of manganese dioxide and the forma-
tion of potassium permanganate, KMnO^.

3 K2Mn04-|-2 H2O — ^ 2 KMnO^ + MnOa 4-4 KOH

Potassium permanganate is obtained as dark purple
crystals, which dissolve in water, yielding a violet solution.
It is a powerful oxidizing agent.

^\5^358. Chromiuin. — Chromium occurs chiefly as chromite,

O* iprjOg . FeOJ From this it is reduced by aluminum in a

^ manner analogous to that employed in the preparation of

manganese. It is a hard, steel-gray metal, unaltered by

the air. When very small quantities are added to steel,

the tenacity and hardness are increased.

362 MANGANESE AND CHROMIUM.

359. Oxides of Chromium. — The two important oxides
of chromium are chromic oxide^ CrgOg, and chromic anhy-
dride^ CrOg. The chromic salts are derived from chromic
oxide ; the most important is the double sulphate of po-
tassium and chromium, K2Cr(S04)2 . 12 HgO, known as
chrome alum. A solution of chromic anhydride yields
CrO^ — ions, but the acid is isolated with difficulty, ^as it
breaks up into chromic anhydride and water. The solu-
tion of the anhydride is a powerful oxidizing agent, and
its derivatives, the chromates and dichromates, resemble
it in this respect.

360. Chromates and Dichromates. — Potassium dichromate,
KjCrjO^, is the basis of most of the chromium salts. It
is prepared by heating chromite with potassium carbonate
and lime in a reverberatory furnace. It forms large red
crystals from solution or fusion ; these are somewhat sol-
uble in cold water, and their solubility increases rapidly
as the temperature rises. Potassium dichromate reacts
with sulphuric acid, with the formation of chromium sul-
phate and the liberation of oxygen. This oxidizing action
is frequently made use of in depolarizing voltaic cells.
Sodium dichromate is very similar to potassium dichro-
mate, but has the added advantage of being more soluble.
Most of the dichromates are red in solution.

Potassium chromatCy KjCrO^, is prepared by the addi-
tion of potassium hydrate to the dichromate :

KaCrgO^ + 2 KOH-^ 2 K^QtO^ + HgO
It forms yellow crystals, more soluble than those of the
dichromate. The relation between the chromate and the
dichromate may be seen if the formula of potassium di-
chromate is written KgCrO^. CrOg. Other complex chro-
mates are known containing more than one CrOg group.

RELATIONS OF CHROMIUM COMPOUNDS. 368

Lead chromate^ PbCrO^, made by treating a soluble
lead salt with a chromate or dichromate, is a yellow, in-
soluble compound, known as chrome yellow, and used as
a pigment.

361. Selationa of Chromium Componnds. — When chro-
mous compounds (^e,g. chromous chloride, CrClg) are
converted into chromic compounds (^e,g, chromic chloride,
CrClg), the valence of chromium changes from two to
three. If a chromic compound is fused with sodium bi-
carbonate and some oxidizing agent, such as potassium
chlorate, a chromate is obtained.

The valence of chromium in the chromates is six. The
changing of chromous into chromic compounds, and the
changing of the latter into chromates, are cases of oxida-
tion.

In the presence of acids, even weak ones, the soluble
chromates are converted into dichromates.

As the valence of chromium in each case can be shown to
be six, it appears unadvisable to consider this as a case of
oxidation. Here the hydrogen ion of the acid seems to be
the active agent.

In the presence of an excess of hydroxyl ions, that is,
in alkaline solutions, the soluble dichromates are changed
to chromates.

Cr207- - + 20H — ^2Cr04-- + U^O

This change should not be considered as a case of reduc-
tion.

When alcohol or some similar substance is added to a
mixture of sulphuric acid and potassium dichromate, the
change in color from orange to green shows that the

364 MANGANESE AND CHROMIUM.

chromium has been changed from a dichromate to a chro-
mic compound. The valence of chromium has been re-
duced from six to three. This is a case of reduction.

KaCr^O^ + 4 HaSO^ + 3 C^Hfi -^

KjSO^ + Cra(S04)3 + 3 CgH^O + 7 H2O

CHAPTER XXXIV.
THE PEEIODIO LAW.

362. Eaxly Attempts at ClassiflcatioiL. — The discovery
of new elements and the investigation of their properties
led the earlier chemists to recognize the existence of cer-
tain families or groups of elements. In 1829 Dobereiner
called attention to certain triads or groups of three ele-
ments in which the atomic weight of the second element
was the arithmetical mean of the first and third. He
also pointed out that the properties of the middle element
were intermediate between those of the other two. This
was the first attempt to show that a relation existed be-
tween the properties of elements and their atomic weights.
The elements chlorine, bromine, and iodine form a well-
marked triad :

35-5 + 127^81.2;

the atomic weight of bromine, 80, approximates the mean^^
81.2.

Other attempts to classify the elements were made from
time to time, but it was not until 1860-1870 that any sys-
. tem received recognition. In 1863-1864, Newlands, an Eng-
lishman, directed attention to the fact that the elements
showed surprising regularity when arranged in order of
their atomic weight. He said the properties of each
element seemed to be repeated in a measure by those of

365

366 THE PERIODIC LAW.

the eighth element following it. This relation is called
the law of octaves. Newlands' system of classification, al-
though it contained many of the principles we use to-day,
attracted little notice. Lacking a strong advocate to push
its claim, the new system was soon forgotten. In 1869
Mendelejeff, a Russian chemist, aroused great interest
in scientific circles by bringing forward a system of clas-
sification which for the first time brought all the ele-
ments into a comprehensive scheme of relationship based
upon their atomic weights. A few months later, Lothar
Meyer, a German, put forward a similar system which he
had worked out independent of Mendelejeff. Although
Meyer has done much to assist in classifying the elements,
it is now generally acknowledged that Mendelejeff is
entitled to the greater credit, and the system we use
to-day bears the name of the Russian chemist. It was
Mendelejeff who brought forward a system which he
elaborated and successfully defended against the many
attacks made upon it.

363. Periodic Law. — Beginning with lithium, let us
arrange the elements in the order of their atomic weights :

Lithium. Gluoikitm. Boson. Gasbok. Nitboosk. Oxtgsn. Fluobink.

7 9 11 12 14 16 19

Lithium is an element with strong metallic or basic prop-
erties ; glucinum. Be, is less metallic ; boron has some
metallic properties, but generally acts like a non-metal ;
carbon forms weak acids ; nitrogen shows stronger acid
properties ; oxygen is characteristically acid ; fluorine, at
the end, is the most pronounced acid element. Hence,
the seven elements show a gradation in properties from a
pronounced metal to an element of strongly acid charac-

LONG AND SHOUT PER10L8.

367

ter. A similar transition can be shown for other prop-
erties as we pass from lithium to fluorine. Thus the
properties seem to vary with the atomic weights, or, in
mathematical language, the properties are functions of
the atomic weights. Sodium, the eighth element after
lithium, closely resembles it, and may be placed directly
beneath as the beginning of another horizontal row :

Lithium

Glucinum

Boron

Carbon

Nitrogen

Oxygen

Fluorine

Sodium

Magnesium

Aluminum

Silicon

Phosphorus

Sulphur

Chlorine

35.5

Magnesium repeats the properties of glucinum, and the
characteristics of boron recur in aluminum. That is, the
eighth element repeats the properties of the one taken
as the first. Silicon, then, should be like carbon, and
phosphorus should resemble nitrogen. These we know
to be facts. Since the properties recur or are repeated at
regular intervals, the properties are said to be periodic ;
or, as Mendelejeff expressed it, " a periodic repetition of
properties is obtained if all the elements be arranged in
the order of the atomic weights."

364. Long and Short Periods. — The table on page
is arranged according to the principle of classification just
given. Omitting for the present the first vertical column
marked Series O, the seven elements, from lithium to
fluorine, form a horizontal series known as a bJioH period.
The set of elements from sodium to chlorine make the
second short period. Beginning in the next line with
potassium, it is found that the metallic properties do not
disappear so rapidly as in the first and second short
^riods. Manganese, the seventh element, has some well-

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FAMILIES OB GROUPS OF ELEMENTS. 369

marked metallic properties. Iron is not placed under po-
tassium, but is put in an eighth series together with cobalt
and nickel. There is a gradual increase in the metallic
properties as we pass thi'ough these three elements to the
more metallic copper. The elements from copper to
bromine show a gradual decline in the metallic properties
and an increase in the acid properties until the strongly
acid element bromine is reached. Hence we have a series
of seventeen elements, beginning with potassium and end-
ing with bromine. This is known as a long period. The
elements from rubidium to iodine constitute the second
long period. In this period the elements of the eighth
series, ruthenium, rhodium, and palladium, form a bridge
in the transition of properties from the seventh series to
the first.

365. Families or Oronps of Elements. — It is evident from
the periodic nature of the classification that all the ele-
ments in one of the vertical series have certain resem-
blances. The relationship, however, is much closer in
some cases than in others. Thus, in Series II, calcium,
strontium, and barium are more closely allied to each
other than they are to magnesium, zinc, cadmium, and
mercury. These last four elements form a closely re-
lated group. That is, the more closely related elements
are not successive, but alternate in a vertical series. The
result of this alternate arrangement is to divide each
vertical series into two families or groups. In Series VI,
chromium and molybdenum are in one family, while sul-
phur, selenium, and tellurium form the other. The halo-
gen elements, chlorine, bromine, and iodine, in Series
VII, afford one of the best examples of a closely related
group. Lithium, potassium, rubidium, and caesium, in

370

THE PElilODIC LAW.

Series I, are a group of soft, waxy metals of high lustre
and low boiling-points. They decompose water readily,
forming caustic bases. A study of their properties shows
that a gradual transition in properties accompanies the in-
crease in atomic weights.

A similar variation in properties in accordance with
the increase in atomic weight in a vertical series is well
brought out in the study of the elements of the halo-
gen group. At the head of each vertical series are placed
some general formulas for the oxides and hydroxides of the
elements in the series beneath. R is the general symbol
for an atom of the element under consideration. Thus in
Series I the general formula of the oxide is RgO, and we
have the oxides LigO, NagO, KgO, CugO, and so on.

These general formulas may be extended to include
other compounds, as the chloride, nitrate, and sulphate:

SERIES I.

SERIES n.

FOBMinjL.

Illustration.

Oxide RgO

KjO

R2O2 CaO
(=2R0)

Hydroxide ROH

KOH

R(0H)2 Mg(OH),

Chloride RCl

KCl

RCI2 HgCl,

Nitrate RNOg

KNO3

RCNOa), Zn(N08)a

Sulphate R^SO^

SERIES m.

RSO4 BaSOi

FOKHVI^.

Illvbtbatxon.

Oxide

RjOg

AlA

Hydroxide

R(0H)3

A1(0H)8

Chloride

RCI3

AICI3

Nitrate

RCNOg),

A1(N03)3

Sulphate

RaCSOJg

A1,(S003

POSITION OF THE INERT GASES. 371

It can be seen from the study of these general formulas
that there is a regular increase in valence as we proceed
from the first series to the seventh.

366. Position of the Inert Oases and of Hydrogen.— With
the discovery of argon and other inert gases, considerable
discussion arose as to their proper place in the periodic
system. Since no compounds of these elements are known,
they cannot be properly placed in any one of the vertical
series. Accordingly it has been deemed best to form a
separate vertical series for these elements at the beginning
of the classification, and mark it Series O.

It will be noticed that hydrogen is not placed in the
table given on page 368. Since it has the smallest atomic
weight, its natural position would be the beginning of the
classification. If, however, it were placed in Series O,
it would be classed with the inert elements from which it
differs decidedly in its properties. A similar difficulty
would arise if hydrogen was placed in Series I, which con-
tains the alkali metals. At present the position of hy-
drogen is so uncertain that it is left out of many periodic
tables.

367. Significance of Vacant Spaces in Table. — It will be
noticed that the series of elements is almost complete
until the atomic weight of 145 is reached, while among
the elements of a greater atomic weight many vacancies
exist. Considerable speculation has arisen as to the
meaning of these gaps. Perhaps in time other elements
will be discovered to fill in these blanks.

368. Valne of the Periodic System. — Mendelejeff^s system
has been of great value in predicting the discovery of new
elements. In fact, in the years immediately following the

372

THE PEmomC LAW.

announcement of the law, when its validity was so ques-
tioned, the fulfilment of Mendelejeff's predictions as to
the existence and properties of elements then unknown,
gave striking evidence of the correctness of the new system
of classification. The table following shows the predic-
tions and their verification in the case of an element which
Mendelejeff called eka'dluminvm^ and which is now known
as gallium:

Pbopkbtibs.

Pbxdiotbd.

DiSOOYSIUED.

Atomic weight

About 69

69.9

Melting-point

Low

30.1^

Specific gravity

About 5.9

5.93

Action of air

None

Slightly oxidized at red
heat

Action on water

Decomposes at red heat

Decomposes at high
temperatures

The predictions by MendelejefiE and their subsequent veri-
fication were equally striking in the cases of eka-boron
(scandium) and eka-silicon (germanium).

The second use of the classification is in the adjustment
and revision of atomic weights. In the early days of the
classification many of the elements were improperly
placed. It was suggested that this might be due to in-
correct values for the atomic weights. This led to more
accurate determinations of the atomic weights. In many
cases results were obtained which admitted of the ele-
ments being placed in the table according to their
proper relationships. Chemical research has been greatly
stimulated by these revisions of atomic weights.

The development of a systematic study of the elements has
been the greatest service of the periodic law. A knowledge

VALUE OF THE PERIODIC SYSTEM. 878

of relationships has simplified the determination of the
physical and chemical properties not only of the elements
but of their compounds. Although the system of Men-
delejeff is not perfect, and no exact numerical relations
have been found, the periodic classification is of great aid
to the student of descriptive chemistry.

CHAPTER XXXV.
OOHPOUNDS OF OASBOIT.

369. Most of the elements we have studied form com-
paratively few and simple compounds, which are gener-
ally ionized in solution. Carbon, on the other Ifand, forms
many compounds, often of complex structure and not
readily ionized. As many of these compounds are formed
in vital processes, the study of such materials was known
as organic chemistrt/. The term as now used means the
study of carbon compounds.

370. Sources. — Practically, carbon compounds are not
made from carbon, as carbon reacts indifferently with re-
agents. Certain easily obtainable compounds of carbon
which react more readily are used for the preparation of
others. A few compounds, as sugar, quinine, starch, and
turpentine, are obtained directly from plants. Many fats,
oils, and waxes, are obtained from plants and animals.
Many other compounds are obtained from the distillation
of petroleum and from the destructive distillation of
wood, coal, and bones. By such destructive distillation,
compounds are obtained which were formed during the
heating and which did not exist in the original material.

371. DeBtmctive Distillation of Wood. — When wood is
heated without access of air, volatile materials pass off
and charcoal remains. As these volatile materials contain
several valuable substances, the operation is carried out

374

BE8TBUCTIVE DISTILLATION OF SOFT COAL. 376

on a large scale. Wood is heated in cylindrical iron
retorts (Fig. 120, a). The volatile materials are cooled,
and while a portion remains gas-
eous and is used as fuel, a large
part is condensed (Fig. 120, 6)
to a dark, tarry liquid.

The tarry mass is distilled with
slaked lime at a low temperature,
and wood alcohol is obtained,
The object of adding the lime is
to neutralize acid present in the mass. The acids, prin-
cipally acetic acid, are recovered from the residue by the
usual method, distillation with sulphuric acid.

When bone is heated^ it undergoes changes analogous to
those of wood. Gas, volatile liquids, and boneblack are
obtained. The liquid bone-oil is a valuable source of car-
bon-nitrogen compounds.

Fig. 120.

372. Destructive Distillation of Soft Coal. — When soft coal
is distilled in the preparation of coal-gas (page 180, Fig. 66),
four products are obtained, — coal-gas, ammonia, coal-tar,
and coke. The coal is heated in horizontal pipes called
retorts^ and the volatile materials expelled. The non-
volatile residue is coke, principally carbon and the ash
constituents. The gaseous materials are cooled in pipes
where tarry materials condense and are collected. The
gases then pass to the scrubber^ where they come in contact
with water dripping over wooden lattice work. Here the
ammonia is dissolved in the water, and more tar separates.
The working of the ammoniacal liquid has been described
in the chapter on ammonia (§ 163). The gas is freed
from sulphur compounds by passing through boxes con-
taining slaked lime or iron oxide, and then passes to the

876 COMPOUNDS OF CABBON.

holders for distribution. Various compounds are obtained
from the tar by fractional distillation. Benzene, carbolic
acid, and naphthalene are some products from coal-tar;
they are invaluable in the preparation of many pharma-
ceutical materials and dyestuffs.

373. Petrolemn and its DistillatiozL. — Petroleum is a dark,
oily liquid occurring in the earth. It is believed to be
due to changes brought about in organic matter under
the influence of heat and pressure, as in the case of coal,
or, perhaps, by the reaction between water and metallic
carbides at great depths below the surface of the earth.
Petroleum is a mixture of many carbon-hydrogen com-
pounds. Practically no pure compounds are separated
from it, but the many commercial articles prepared from
petroleum are mixtures obtained in its distillation.

The crude petroleum is distilled in iron retorts (Fig.
121) connected with condensers and receiving-tanks.
As the petroleum is a mixture, the lighter substances with

lower boiling-points are first
volatilized, and the tempera-
ture gradually rises until
only a pitchy mass remains
Pj j2i^ in the retort. A hydrometer

indicates the specific gravity
of the liquid as it runs from the condenser, and, as the
density of the distillate varies, it is run to appropriate
tanks.

Gasolene^ naphtha^ benzine^ and kerosene are some of the

trade names of the fractions into which the separation

is made. They differ in density, volatility, and kindling

temperature, each being a miigbure of several compounds.

As in the case of kerosene, a further purification is

HYDBOCABBONS. 377

often needed. To rid the kerosene of materials which
would not burn well or might smell, the liquid is agitated
successively with concentrated sulphuric acid, sodium
carbonate solution, and water. The light petroleum prod-
ucts are used almost entirely for solvents and fuel. The
residue in the still, after the removal of the lighter oils,
is further distilled at a higher temperature in another
retort, and heavier materials obtained, such as the various
grades of lubricating oils, vaseline, and paraffin. The
mineral oils have largely displaced animal and vegetable
oils as lubricants. The paraffin is used in candles, as a
preservative coating, in waterproofing, and for many other
purposes. The final residue in the retort is coke, and is
used in the manufacture of electric-light carbons or for
fuel.

Natural gas is a mixture of combustible gases issuing
from the earth in many localities, often under consider-
able pressure. When petroleum is brought to the surface
and distilled, considerable gaseous material is separated.
The chief constituent of these natural gases is methane.
Where natural gas is obtained in quantity, it is, of course,
a valuable fuel.

HYDROCARBONS.

There are many compounds of carbon with hydrogen ;
these are called hydrocarbons.

374. Methane, or Harsh-gas. — Methane, CH^, is a col-
orless, odorless gas which, when pure, burns with a non-
luminous flame. It is often formed in the decomposition
of organic matter, as in swamps, hence its common name,
mar$h'ff(M. It is the principal constituent of natural gas.
In flof t-coal mines, the miners call it fire-damp^ as its mix-
ture with air is a serious source of danger.

378

COMPOUNDS OF CARBON.

It is formed in the production of water-gas when the
temperature of the furnace is low :

C + HaO— ^CO + Ha
3C + 2HaO— ^2C0-f-CH^
Other than fuel it is of no practical importance, as it re-
acts with very few materials.

375. ParafSn Series. — Methane is the simplest member
of a %erie8 of hydrocarbons, all of which resemble it in
their lack of chemical activity, whence the name paraffin
series.

PARAFFIN SERIES.

Formula..

MOLEOULAB

Weight.

BOILING-

Pomx.

Methane

CH,

-164°C

Ethane

C,H«

-89.5

—

Ordinarily

Propane

C8H«

-38

—

gaseous

Butane

C4H10

+ 1

—

Pentane

^6^12

—

Liquid

Hexane

^bHm

—

Hexadecane

^161184

226

288

18* ]

.Solid

Octodecane

CigHgg

254

317

28 J

It will be observed that the formula of each member dif-
fers from the preceding by CH^ ; such a series is called a
homologous series. The general formula for the series is
C^Hg^+g. With increasing molecular weight there will
be noticed a rising of the boiling-point and the tendency
to assume the solid form in the higher members. It is
mixtures of these compounds that occur in petroleum
products.

376. Benzene. — Benzene, or benzol, C^Hg, is a light, col-
orless, volatile liquid, having a peculiar odor. It is ob-

ACETYLENE. 379

tained from that portion of coal-tar that boils at 80®- 86°.
Benzene burns with a smoky flame. It is a good solvent
for resins and fats. Its principal use, however, is for the
production of more complex compounds. Unlike the par-
affin hydrocarbons, the coal hydrocarbons react with com-
parative ease, as with nitric and sulphuric acid, forming
important compounds used in the preparation of dyestuffs.
Benzene is the first member of a series, C^Hjn-^, analogous
to the paraffin series.

377. Acetylene. — Acetylene has been mentioned as re-
sulting from the reaction of calcium carbide and water :

CaCa -h 2 HjO — ^ CaCOH)^ + C^H^

It has a peculiar, disagreeable odor, noticeable when a
Bunsen burner is lighted at the base. It is here formed
by the incomplete combustion of the gas. Ordinarily it
burns with a smoky flame, but with a suitable burner it
furnishes a brilliant light which nearly approaches sunlight
in color. Acetylene is exploded by concussion or by an
electric spark, so that it is not often made in quantity, but
is made as used. Its chief use is as an illuminant.

SUBSTITUTION PRODUCTS.

378. BelatLon to Methane. — The numerous compounds of
carbon may appear puzzling, but a general view of them is
simplified if we regard them as substitution products of
methane.

If we consider any hydrogen atom of methane, CH^,
it is combined with a carbon atom and three other hydro-
gen atoms :

H-CH3 or ^>C<^

H H

880 COMPOUNDS OF CABBON.

The group (CHg) has evidently combining power equal to
one hydrogen atom, so that we may say its valence is one.
Now, we can conceive of one hydrogen atom in a molecule
of methane replaced by its equivalent (CHg) so that a
compound (CHg) (CHg) would result, having the com-
position of ethane, CjHg. In a similar manner, if one
atom of hydrogen in ethane is replaced by its equivalent,
CHg, we would have CgHg for the third member of the
series, and so on.

Chlorine is equivalent to hydrogen in valence (HCl).
If one atom of chlorine was substituted for one of the
hydrogen atoms of methane, we would have CHgCl; if two
were substituted, CHjClg ; similarly, CHClg and CCI4. In
like manner we could have :

CHgBr CHgBr^ CHBrg CBr^

CHgl CH2I2 CHTg CI4

All of these compounds are known, but it should be re-
membered that while the structure of such compounds is
explained on the supposition of substitution, they are
usually not made that way.

379. Monochlormethane. — Mpnochlormethane, methyl
chloride^ CHgCl, isa colorless gas having an ethereal odor.
Methane and chlorine react :

CH^ + Cla—^ CHgCl -h HCl

The reaction is too violent to be of practical use. The
chloride is commercially prepared from a by-product ob-
tained in sugar-refining. Methyl chloride is easily liquefied,
and the liquid is used as a local anesthetic, producing
insensibility by freezing. It has also been used in ice
machines.

CHLOROFORM AND IODOFORM. 381

380. Chloroform and Iodoform. — Trichlormethane, cUoro-
form^ CHClg, is a heavy, colorless, easily flowing liquid.
It has a peculiar odor and a sweet taste. It is scarcely
soluble in water. Chloroform is a most valuable anesthetic
and an important solvent. It is prepared by distilling
alcohol or acetone with a solution of bleaching-powder.

Tri-iodomethane, iodoform^ CHI3, is a light yellow pow-
der with a characteristic odor. It is useful as an antiseptic.
Iodoform may be prepared by the reaction of iodine and
alcohol rendered slightly alkaline.

ALCOHOLS.

Alcohols, as a general name, is applied to a class of
bodies resembling methyl and ethyl alcohols. They may
be briefly described as hydroxyl substitution products of
hydrocarbons.

381. Methyl or Wood Alcohol The hydrogen in hydro-
carbons can be substituted by other elements or groups.
This substitution often has to be accomplished by an
indirect process. For example, if monochlormethane is
heated with dilute potassium hydroxide, the following re-
action takes place :

CHgCl + KOH— ^CHgOH + KCl

The compound CH3OH is methyl hydroxide, wood alcohol.
From the reaction just given, it appears probable that this
compound is methane in which a hydroxyl group has been
substituted for a hydrogen atom. It is an organic hy-
droxide, and, as might be expected, has basic properties
to a certain limited extent. Wood alcohol is commercially
obtained by the destructive distillation of wood. It is a
colorless liquid of low boiling-point. It is used to a large

382 COMPOUNDS OF CARBON.

extent as a solvent in the manufacture of varnishes.
Wood alcohol is a very convenient fuel where small quan-
tities of heat are required, because it burns with a clean
flame of high heat value.

382. Ethyl or Grain Alcohol.— Ethyl hydroxide, CjHgOH,
ordinary alcohol^ can be made by heating moliochlorethane
with potassium hydroxide:

CaHgCl 4- KOH — ^ CjHgOH + KCl

This reaction indicates that ordinary alcohol is a hydroxyl
substitution product of ethane. It may also be regarded
as an oxidation product of ethane, since the two formulas
differ only in the presence of an oxygen atom in the for-
mula of alcohol.

Ordinary alcohol is made in large quantities from grain
or potatoes by the process of fermentation. Both of these
substances contain a large quantity of starch. This is
converted into glucose, a kind of sugar, when heated with
dilute sulphuric acid. The sulphuric acid acts merely as
a catalytic agent:

starch glacose

After the excess of sulphuric acid has been neutralized
with lime, yeast is added, B,nd fermentation occurs. Yeast
is a microscopic vegetable organism consisting of oval-
shaped cells arranged in chains. In solutions which con-
tain suitable food it multiplies rapidly and secretes a sub-
stance called zt/mase, which acts as a catalytic agent in
converting glucose into alcohol and carbon dioxide :

C^Hi^Og— ^2 C2H5OH + 2 CO3

The alcohol is separated from the resulting solution by
repeated distillations. One distillation is not enough to

ALCOHOLIC BEVERAGES. 383

accomplish the separation, because the boiling-point of
alcohol, 78°, is close to that of water. The first portions
distilled in each operation are relatively rich in alcohol.
By collecting the distillate in fractions, and redistilling
these, the separation is made nearly complete. The pro-
cess is known as fractional distillation. Ordinary com-
mercial alcohol contains from 90% to 95% of the pure
substance.

Ethyl alcohol resembles methyl alcohol in its properties.
It is a low-boiling liquid, an excellent solvent for organic
compounds, and it burns with a clean flame of high heat
value. As a constituent of alcoholic beverages it is manu-
factured in enormous quantities. These owe their in-
toxicating properties to the presence of alcohol. It is
oxidized when taken into the body, furnishing heat.

383. Alcoholic Beverages — Beer is the product obtained
by the fermentation of malt. Barley is placed in a warm
moist room until the kernels germinate. When the root-
let has grown to be two-thirds the length of the kernel,
the grain is heated to stop the growth. During the ger-
mination the starch in the barley is converted into a sugar
(maltose). The malt thus obtained is ground and boiled
with water. Yeast is added to ferment the malt sugar.
The fermented liquor is filtered and water added to pro-
duce a beer or ale of the desired concentration. Rice and
glucose are often used to replace barley. Hops and other
flavoring materials are also utilized. Beer contains from
3 to 5 per cent alcohol in addition to soluble materials
from the grain.

Wines are produced by the fermentation of fruit juices
which contain grape-sugar (glucose). After fermenting,
the liquor is allowed to settle and the clear liquid is drawn

884 COMPOUNDS OF CARBON.

off. The wine thus produced, in addition to 15 per cent
of alcohol, contains soluble materials derived from the
fruit or produced in the fermentation. Whiskey is made
by distilling a beer obtained from rye or corn, so that the
percentage of alcohol is increased to about 50 per cent.
Brandy, resulting from the distillation of wine, may con-
tain 70 per cent alcohol, and rum and gin, derived from
fermenting molasses, possess about 75 per cent. All of these
distilled liquors contain minute quantities of flavoring
materials and traces of acid.

384. Denatured Alcohol. — Denatured alcohol is ethyl
alcohol to which wood^ alcohol or other poisonous sub-
stances have been added in order to make its use impossi-
ble in beverages and medicines. In countries where a
tax is imposed on alcoholic liquors denatured alcohol
is often exempt so that the cost of the article in manu-
facturing operations shall not be prohibitive. Such an
exemption law has recently been passed in this country.
As a result of it, many kinds of chemical manufactures
will be stimulated, and alcohol will probably come into
more general use as a fuel.

In the United States methyl alcohol and benzene are
the denaturing agents authorized by the Commissioner of
Internal Revenue. The proportions by volume are as
follows :

100 parts ethyl alcohol (not less than 90% strength)
10 parts methyl (wood) alcohol
J part benzene

Such alcohol is classed as completely denatured, but there
are many formulas for denaturization to suit special pur-
poses.

ALDEHYDES. 385

385. Aldehydes. — Another series of compounds, known
as aldehydes, are closely related to alcohols. They can be
made by an oxidizing process which results in the taking
away of two hydrogen atoms from the molecule. Thus
formaldehyde is made by passing a mixture of methyl
alcohol vapor and air over a heated copper spiral or heated
asbestos which is sometimes platinized :

CHgOH + O — ^ CHjjO + H2O
Formaldehyde is a very valuable disinfectant and preserv-
ative. It is a gas at ordinary temperatures, but it comes
into the market as a water solution known q,^ formalin.

ORGANIC ACIDS.

386. These may be regarded as oxidation products of
the aldehydes. Thus, in oxidizing hydrocarbons, we
obtain successively alcohols, aldehydes, and acids :

CH4+ O — ^ CHgOH
CH80H+ O — ^ CHjO + H2O
CH2O + O — >- CH2O2 (formic acid)

387. Properties. — The organic acids have in a less
degree the characteristic properties of the familiar inor-
ganic acids. They neutralize bases, forming salts and
water, and they often act directly on metals, liberating
hydrogen. Their water solutions do not conduct the elec-
tric current very well, and hence they are weak acids.

388. Formic Acid. — Formic acid, CHgO^ or HCHOg, is
the lowest member of the series ; it is of no practical im-
portance. The sting of ants is said to be due to this
compound.

389. Acetic Acid and Vinegar. — Acetic acid, HC2Hg02,
is theoretically derived from the oxidation of ethane, with

386 COMPOUNDS OF CARBON.

the formation of ordinary alcohol and acetaldehyde as
intermediate products. Dilute solutions of alcohol, such
as are represented in weak wines, or hard cider, when
exposed to the air, undergo a fermentation which results
in the oxidation of alcohol to acetic acid :

C2H5OH + 03^- HC3H3O3 + HjO

Vinegar is the liquid that results from this action ; it is a
dilute solution of acetic acid containing impurities which
give it color and modify its flavor. Vinegar is some-
times manufactured by what is called the quiek vinegar
process. Dilute alcohol, to which a certain amount of beer
or malt extract has been added, is allowed to trickle over a
mass of wood shavings which have been previously treated
with vinegar in order to insure the presence of the fer-
menting organism. The porous mass of shavings makes
possible free contact with air, which furnishes the neces-
sary oxygen.

Pure acetic acid is a colorless liquid which freezes on
slight cooling. It has a corrosive action on the flesh.

ETHEREAL SALTS, OR ESTERS.

390. Formation and Uses. — Ethereal salts and water are
formed by the action of an acid with an alcohol. The
reaction is analogous to that which takes place during the
formation of a salt by neutralization. The alcohol may
therefore be considered as taking the part of a base :

NaOH + HCCaHgOa) —^ H^O + NaCCaHgOa)

sodium acetate

C2H5OH + HCCgHgO^) — ^ H^O + CgHgCCaHsO^)

ethyl acetate

NITROGLYCERINE. 387

The esters form an important group of compounds.
Some are employed in medicine, while others are used in
the preparation of perfumery. Many are used in making
artificial fruit flavors. The characteristic flavor of the
pineapple is due chiefly to ethyl butyrate ; oil of winter-
green is methyl salicylate.

391. Nitroglycerine. — Nitroglycerine, C8H5(N08)8, is an
ester of an alcohol (glycerine) and nitric acid. It is pre-
pared by the action of glycerine, C8Hg(OH)8, with a mix-
ture of concentrated nitric and sulphuric acids :

C3H6(OH)3 + 3 HNO, -^ 3 H3O + CgH^NOg),

The sulphuric acid aids the action. by uniting with the
water formed during the reaction. The concentration of
the free nitric acid is thus kept at maximum. Nitro-
glycerine is a highly explosive liquid at ordinary tem-
peratures. Dynamite is nitroglycerine which has been
absorbed by infusorial earth.

392. Oils, Fats, and Soaps. — Oils and fats are esters of
glycerine and various fatty acids. The chief constituent
of beef tallow is glyceryl stearate, an ester of glycerine
and stearic acid, commonly called stearin. It is formed
by the reaction between one molecule of glycerine and
three molecules of stearic acid.

When such a fat is boiled with a solution of sodium
hydroxide, a molecule of glycerine and three molecules of
sodium stearate, a hard soap, result from the reaction :

C8H6(Ci8H8502)8 + 3 NaOH ->C8H6(OH)s + SNaCigHscOa

Glyceryl stearate + sodium hydroxide — > glycerine + sodium stearate.

Common hard soap is a mixture of sodium salts of fatty
acids, chiefly stearic, palmitic, and oleic acids. Soft soap

888 COMPOUNDS OF CARBON.

is a mixture of the potassium salts of the fatty acids.
The term soap is applied in general to any metallic salt
of a fatty acid.

393. Soap-making. — No technical process depends more
upon the skill of the operator than the manufacture of
soap. In general, soaps may be classed as boiled, semi-
boiled, or cold process. Boiled soaps are produced by
boiling fats with sodium hydroxide and carbonate. They
are often called settled or grain soaps because, during the
process of manufacture, the glycerine is separated from
the soap. Semi-boiled soaps contain all of the glycerine
derived from the fats. Cold process soaps are formed by
the direct combination of the fat and alkali, without the
aid of external heat.

Most hard soaps used for household purposes are boiled
soaps. Soap is made in large iron kettles, fitted with two
sets of steam pipes ; one a closed coil to supply heat, and
the other an open coil to deliver steam through the charge
to keep it stirred. Melted fat and about one-fourth the
quantity of alkali required for complete saponification are
run into the kettle and the steam turned on. The con-
centration of the alkali is regulated by the kind of fat
used. When the mixture has become homogeneous, a more
concentrated alkali is added and the boiling continued
until a sample shows that the product has the desired con-
sistency. Salt is then added, and the soap being insoluble
in brine separates. The kettle is allowed to remain quiet
for several hours and the soap collects on top of the liquid.
This liquid is called spent lye. It contains glycerine, water,
salt, and impurities from the alkali and fat. The spent
lye is drained off, the salt and glycerine are separated from
it, and the layer of soap is boiled with sufficient alkali to

HABD WATSR AND SOAP. 889

complete the saponification. During this boiling, resin is
sometimes added. Pure resin soaps have strong detergent
properties, but are too soft and sticky for general use.

The soap when taken from the kettle is often mixed with
one or more of various fillers, such as sodium carbonate,
borax, and sodium silicate. Coloring materials and per-
fumes may also be added. Soaps that float are made light
by having air forced through them while they are in the
pasty condition. If a good quality of soap is dissolved in
alcohol and then dried, a transparent soap is obtained.
Sugar and glycerine are often used in the manufacture
of transparent soaps. Most soap powders consist essen-
tially of hard soap ground with sodium carbonate.

394. Hard Water and Soap. — When a soluble soap, for
example sodium stearate, is used with a water containing
calcium ions, an insoluble calcium soap, calcium stearate,
is formed :

2 NaCCigHggOa) + CaSO^ — ^ CaCCigHgsO^)^ + Na^SO,

The soap is said to be destroyed, as good suds cannot
be formed until the calcium ions are removed from
solution. This explains why hard water is not desirable
for washing purposes. The hardness of water is measured
by its soap-destroying power. This is commonly due to
the presence of calcium and magnesium ions in the water.

395. Ether. — Ordinary ether, C^H^qO, may be regarded
as ethyl oxide, (C2Hg)20. It is prepared by treating
alcohol with a dehydi'ating agent, such as sulphuric or
phosphoric acid.

2 C^HgOH — ^ CC^llsW + H^O
Ether is a volatile, inflammable liquid, boiling at 35^.
It is used as a solvent and as an anesthetic.

390 COMPOUNDS OF CARBON.

CARBOHYDRATES.

The carbohydrates are chemical compounds composed
of carbon united to hydrogen and oxygen ; the last two
elements being in the same proportion as in water.

396. Cellulose. — The cell walls of plants are composed
of cellulose, a compound having a percentage composition
corresponding to the formula CgHj^Og. Absorbent cot-
ton and the better grades of filter-paper are pure cellulose.
It is the chief constituent of straw and wood.

When boiled with acids, cellulose is slowly converted
into a sugar called glucose. Cellulose dissolves without
change in an ammoniacal solution of cupric hydroxide,
known as Schweitzer's reagent. The cellulose can be
precipitated from such a solution by the addition of hy-
drochloric acid. Cellulose is dissolved in zinc chloride
in making the filaments for incandescent electric light
bulbs.

If the formula for cellulose is considered to be
(CgHiQOg)^, from two to six nitro (NOg) groups can be
introduced into the molecule. Thus, when pure cotton
fibre is treated with a mixture of nitric and sulphuric
acids, products are obtained which may contain two,
three, four, five, or six nitro groups ; the number depend-
ing upon the concentration of the acids and the time
during which they are allowed to act.

The di-, tri-, tetra-, and penta-nitrocelluloses are known
as soluble guncotton. Hexanitrocellulose is insoluble gun-
cotton. Collodion is a solution of soluble guncotton in a
mixture of alcohol and ether. When such a solution is to be
used as liquid court plaster, about 5 % of Venice turpen-
tine and 3 % of castor oil are added to prevent the shrink-

PAPER MAKING. 391

age of the film and to make it more flexible. Celluloid is a
guncotton incorporated with camphor. Explosive gelatine
consists of guncotton dissolved in nitroglycerine. Bal-
listite contains equal parts of soluble nitrocellulose and
nitroglycerine, to which is added a small quantity of a
substance (diphenylamine) to increase its stability. Cor-
dite contains nitroglycerine, guncotton, and vaseline.
Explosive gelatine is used in blasting ; ballistite and
cordite are examples of smokeless powders.

397. Paper Making. — Much wood is used in the manu-
facture of pulp from which paper is made. The wood is
finely shredded, then boiled under pressure in a solution
of either sodium hydroxide or a mixture of calcium and
magnesium bisulphites. The fibre is bleached with
chlorine and pressed into sheets called pulp. As the
fibreflS^l^d are very short, the paper made from wood
is notrougn^

Better grades of paper are made from cotton and linen
rags. The rags are shredded, cleansed in a solution of
sodium hydroxide, and bleached with chlorine obtained
from bleaching-powder. The excess of chloride of lime
is removed by sodium sulphite. Loading material (ba-
rium or calcium carbonate) and size (starch or similar
material) are added to give stiffness to the fibre and to
overcome capillarity.

After the fibre, loading material, and size have been
thoroughly mixed in water, the paper is laid by shaking it
on fine screens. The watery liquid passes through the
screen, while the fibre is left matted in a uniform layer.
The layer of matted fibre is pressed between felt, dried,
and finally calendered by being passed between heated rolls.

Filters and blotting-paper contain neither loading ma-

892 COMPOUNDS OF CARBON.

terial nor size. Newspaper is made from pulp, and may
not have been bleached by chlorine. The best paper is
made from linen.

• 398. Starch. — Starch (CgHiQOg)^ is found as granules
in the cells of plants. These granules consist of a wall
of starch cellulose within which is soluble starch. Starch
is insoluble in cold water, but boiling water causes the
walls of the starch granules to burst, and the soluble
starch enters solution. Much starch is obtained from
potatoes and corn.

No matter what the source, the method of preparing
starch is the same. The material is crushed, then mac-
erated with water, and the milk-colored liquid filtered
through cloth fine enough to prevent all besides the water
and starch from passing. The starch is allowed to settle
to the bottom of the containing vessel, fronawhich it is
afterwards removed and dried. ^^^^w

Dilute acids convert starch into glucose, and much
starch is used for this purpose. Dextrine is prepared by
heating dry starch to about 250° C. It is a valuable
constituent of food, and is used in making paste similar
to that on the back of postage stamps.

Sprouting barley contains an enzyme (ferment) known as
diastase^ which is capable of converting starch into a sugar
named maltose. In the manufacture of malt, the grain is
allowed to germinate to produce the enzyme, after which
the process is stopped by heating the barley to 60®.^ At
a temperature of about 70® C, the diastase rapidly converts
the starch which the grain contains into maltose and dex-
trose.

399. Sugars. — A very large number of sugars are
known. Fructose, or fruit sugar ; glucose, or grape sugar;

8U0AR MANUFACTURE AND REFINING. 393

and saccharose^ or cane sugar, are among the more impor-
tant. Fructose and glucose have the empirical formula
CgHijOg ; the formula for saccharose is Ci^U^On.

Glucose is converted by a ferment (zymase) secreted
by the yeast plant into alcohol and carbon dioxide :

CgHigOg + zymase— ^2 CjHgOH + 2 COj,

This fermentation is made use of in the raising of bread
and in the preparation of alcohol. It also is the cause of
the formation of hard cider.

The manufacture of glucose from starch has already
been referred to. Large quantities of glucose are used in
making candies and table syrups. Common sugar, saccha-
rose, is obtained from the sap of the sugar-cane, sugar-
beet, and sorghum. It is also the principal constituent of
maple sugar.

400. Sugar Mannf aotore and Eefining. — Sugar is obtained
from sugar-cane or beets by macerating the fibre. Slaked
lime is added to the juice to prevent fermentation and
precipitate the albuminous substances coming from the
plant cells. The solution is then filtered through cloth
and evaporated in a vacuum pan at a temperature of about
66^. If the evaporation were carried on under ordinary
pressure, the temperature would become sufficiently high
to convert the saccharose into a mixture of glucose and
fructose.

As soon as a sample taken from the vacuum pan shows
that sugar will crystallize when the syrup cools, the solu-
tion is removed from the pan and allowed to cool. The
crystals are dried in centrifugal machines. The product
obtained is usually raw sugar which must be refined be-
fore being placed on the market. Nearly all sugar refin-
eries are in the northern states.

394 COMPOUNDS OF CARBON.

The raw sugar is dissolved in large vats and the syrup
pumped to the top of high buildings, where it is mixed
with lime and a little boneblack to precipitate any albu-
minous material that may have been left in the sugar.
The syrup is then filtered through long sacks, called bag-
filters, to remove the coarse impurities that are suspended
in the solution. After this, the liquid is filtered through
boneblack to remove the coloring-matter. The purified
syrup is boiled in vacuum pans as in the case of raw sugar.

Crystals of pure sugar have a pale, yellowish tint. As
most people are ignorant of this fact and demand that a
white sugar be sold them, the sugar refiners add some
blue pigment to the sugar, for example ultramarine. The
blue counteracts the yellow and causes the sugar to appear
white.

Saccharose is. converted into glucose and fructose by
boiling:

^12^22^11 + HgO — >- C^H^jOg 4- C^HijOg

Dilute acids hasten this action, which is known as inver-
sion. For this reason vinegar is often added to sugar
during the making of candy that is to be pulled.

Zymase does not convert saccharose into alcohol and
carbon dioxide. However, the yeast plant secretes an-
other ferment called invertase which changes saccharose
to a mixture of glucose and fructose. These, as has al-
ready been mentioned, can be fermented by zymase.

Common sugar melts at 160° C. to a colorless liquid which
solidifies on cooling to a transparent amber-colored mass,
called barley sugar. When sugar is heated to 215*' C, some
water is expelled and a brown mass, caramel, is obtained.

APPENDIX

APPENDIX.

I. PHTSJOAL OOirSTAHTS OF TEE IHFOBTAHT ELEMENTS.

EUEMBNT.

Atomic Weights.

Valbnoe.

SPBOino GnAiniT.

MSLTINO

Ponrr.

Boiling

POIMT.

a
&

Approx-
imate.

Exact
= 16.

Water =1.

Air=l.

OO.

oQ.

Aluminum

27.1

2.6

657

1500-
1700

Antimony

120

120.2

niy

6.6

630

1500-
1700

Argon

39.9

1.38

-188

-186

Arsenic

75.0

niv

5.7

<360

volatile

Barium

137

137.4

3.8

850

950

Bismuth

208

208.0

niv

9.7

1435

Boron

11.0

III

2.6

infiuifl 3500

Bromine

79.96

3.1

-1 59

Cadmium

112

112.4

8.6

322

about

778

Calcium

40.1

1.5

800

Carbon

12.00

1.7-2.1

sublimes

3500

Chlorine

35.6

35.45

2.49

-102

-33.6

Chromium

52.1

nmvi

6.9

1515

Cobalt

59.0

8.7

1530

. . .

Copper

63.6

III

8.9

1065

2100

Fluorine

19.0

1.31

-223

-187

Gold

197

197.2

19.3

1065

...

Helium

4.0

0.13

-271

-267

Hydrogen

1.008

0.07

-256.5

-252.6

Iodine

127

126.97

4.9

114

184

Iron

55.9

nin v:

7.8

1950

. . .

Lead

207

206.9

II IV

11.3

327

1400-
1600

396

PHYSICAL CONSTANTS.

897

EUUfXMT.

Atomic Wwghtb.

Yalsnob.

Spxoifio Gravity.

MSLTIMG
POIMT.

BoiLDve
Point.

Approx-
imate.

Exact
= 16.

Water = 1.

Alr»l.

oc.

OO.

Lithium

7.03

0.59

186

<1400

Magnesium

94.8

24.36

1.7

632

1100

Manganese

56.0

niv

7.4

1245

. . •

Mercury

800

200.0

18.6

-38.8

357

Nickel

68.7

8.7

1484

...

Nitrogen

14.01

inv

0.96

-210

-195

Oxygen

16.00

1.10

<-230

-182

Phofiphorus

31.0

inv

yeflow
1.8

yel
44.2

low
290

Platinum

196

194.8

21.5

1710-
1780

. . .

Potassium

39.15

0.87

62.6

767

Silicon

28.4

2.0

...

3600.

Silver

108

107.93

10.6

961

2060

Sodium

23.06

0.97

97.6

877

Strontium

87.6

2.5

900
rhombic

Sulphur

32.06

nivvi

2.0

114.6

444.6

Tin

119

119.0

niv

7.3

232

1460-

1600

918

Zinc

65.4 n

7.1

419

398

APPENDIX.

M!»

^ 1

H
SO

ouiz

oooDi-'aoaD |»^QQaD»-«»-«aD»-«fV<

aD^»-Hao^t-i |^aQ«-<i^^,5''^

ranipog

aQaDQQOQaQQQaQQCaQaDaQaQQQCG

inni8finno<{

ooaoaDoQaDaQaDaoaQaDaDaDaoao

I93PIK

aQaD«-i |ao |HHaQaD»-'»-JaD-^t-«

(++3h) ounoj9K

cQ(i,HHaQaQPk>-'»-'aQ'-«>-'aQ* |

(+Sh) CTKunowpi

Pt,>HI-taDI-lOH |H-.aQHH»HPH^ |

089aeda«p(

OQODHH |QQ|»-iaoao«-«PHao«-« |

inni89nS«i^

oQoQi-'coaQaQMQQaQMt-iaQGOPH

PB91

aQPU»-.aDPH*HH»-iaQ»-'«-<4S>5'-'

(-H-+9J) 0U-19J

|aD| |cQ|»-"|aDHHMaD| |

(++0J) 8n0JJ9J

aQaD«-< loQ |.i-(coaQ«-<»-'aDOHpk

j9ddoo

QQaQ»-'aQCQ«-<^ IcQHHHHOQ^aO

iJlBqoo

oQOD'^aQaD'^'^aDaQi-^'-'aQ'-''-'

urapiBO

aQao»-<QQaD»-"pLiaQaQPH>-'^fr<ft<

ranirapBO

oQOQi-'aoaQl HHcoaDH-iHHQQ^aQ

q^iira8ia

|pUtHH |pL,l-tMl-ICOI-IH-<aD^|

mnuvg

GOcQHHaQOQt-'aoPkaDaDH-ij^aOH-i

snomasjy

|C«| loQlcOd^lOHl Ml

jfaooii^ay

1 Ph 1 1 P^ 1 f^ c^ 1 - 1 M 1

mn^nomniy

GCaDaQODGCSOODOQaQ laDODODOQ

nroaimn^

CCOD laQOQ If-iOQOQi-iHHaQ 1 1

Acetate .
Bromide .
Carbonate
Chlorate
Chloride
Chromate
Hydroxide
Iodide . .
Nitrate .
Oxide .
Phosphate
- Sulphate
Sulphide
Sulphite

SOLUBLE AND VOLATILE COMPOUNDS.

899

in. aENEBAL BULES FOB SOLITBILITT.

Certain generalizations can be made concerning compounds shown
in the table on the opposite page. The exceptions to these general-
izations are few and unimportant.

1. All sodium, potassium, and ammonium compounds are soluble in
water.

2. All nitrates, chlorates, and acetates are soluble in water.

3. All chlorides are soluble, except those of silver, mercury (mej>
curous),and lead (lead slightly soluble).

4. All sulphates are soluble, except those of barium, lead, and
calcium (calcium slightly soluble V The silver and the mercurous
sulphates are only moderately soluole.

5. All carbonates are insoluble, except those of sodium, potassium,
and ammonium.

6. All oxides and hydroxides are insoluble, except those of am-
monium, sodium, potassium, and barium; calcium hydroxide is
slightly soluble.

IV. VOLATILITY OF 00MP0TOD8 THAT MAY BE8ULT
FBOM DOITBLE DEOOMPOSITIOITS.

1. Compounds volatile at ordinary temperatures :

HCl HBr HF H^S

2. Compounds decomposing at ordinary temperatures yielding
volatile products :

HoCOo (HgO + COo)
HgSOs (H2O + SO2)
NH^H (H2O + NH3)

3. Compounds volatile at varying temperatures below 338*' (boiling-
point of sulphuric acid) :

BoiLmChPOTNT. BoiLme-POINT.

H2O, 100° HNO3, 86°

HCl (aqueous solution), 110° HNO., (aqueous solution), 120°
HBr (aqueous solution), 126° HC2H3O2, 118°

V. WEIGHT OF ONE LITEB OF OOMMON OASES UUDEB
STANDABD OONDITIOUS.

J^Cai

Acetylene,

1.162 grams Hydrogen sulphide,

1.523 grams

Ammonia,

0.762 '

< Marsh gas.

0.716 "

Carbon dioxide,

1.965 *

' Nitrogen,

1.254 «

^arbon monoxide,

1.250 *

' Nitric oxide.

1.340 «

IChlorine,

3.166 *

* Nitrous oxide.

1.968 «

hydrogen chloride,

1.628 '

' Oxygen,

* Sulphur dioxide,

1.429 «

:Hydrogen,

0.0898 <

2.861 «

400

APPENDIX.

VI. FBESSTJBE OF WATEB YAFOB, OB AQUEOUS TEHSION.
(In millimeters of mercniy.)

TXMPXKATnSS.

PKX88UBB.

Pbmsubb.

o.o«c.

4.6 mm.

21.5<>C.

19.1 mm.

6.5

19.7

9.2

22.5

20.3

10.5

9.5

20.9

9.8

23.5

21.5

11.5

10.1

22.1

10.5

24.5

22.8

12.5

10.8

23.5

11.2

25.5

24.2

13.5

11.5

25.0

11.9

26.5

25.7

14.5

12.3

26.5

12.7

27.5

27.3

15.5

13.1

28.1

13.5

28.5

28.9

16.5

14.0

29.8

14.4

29.5

30.7

17.5

14.9

31.6

15.4

54.9

18.5

15.9

92.1

16.4

149.2

19.5

16.9

233.8

17.4

355.4

20.5

17.9

526.0

18.5

100

760.0

KINETIC THEORY OF GASES. 401

Vn. THE KIHETIO THEOEY OF GASES.

The physical structure made evident in the uniform
behavior of gases, under changes of temperature and pres-
sure, is explained by a simple mechanical conception, the
kinetic theory of gases.

According to this hypothesis a gas consists of many
minute particles, exerting practically no attraction for one
another, but moving freely with great velocity. If we
consider a cube to contain a number of such particles,
moving in all directions, it is evident that they must col-
lide among themselves and against the walls of the vessel,
in either case rebounding and continuing their motion in
Ihe new direction until the next collision.

This continual bombardment exerts a force (pressure)
against the walls, which must depend upon the number of
blows, the mass and the speed of the particles. If, now,
the particles (molecules) in our cube were confined to the
lower half, the number of blows against the surface would
be doubled, and the mass and speed being the same, the
pressure would be doubled. That is, the pressure would
vary inversely as the volume.

Let us express this relation mathematically. If m is the
mass of the molecule, n the number of molecules, v their
speed, I the edge of the cube, a molecule would strike the

wall - times a second. At each collision its momentum
would be changed 2 mv^ so that the total force exerted on
the sur

the surface of the cube by n molecules would be w-2 wv-- or

- As the cube has six faces, the total surface of
It

the cube is 6 P. The force per unit area, or pressure (P),

402 APPENDIX.

would be ^^ ^ 6 P or ?^?^. As P is the volume ( T),

or Pr= ^ = a constant. (2)

If the number, mass, and speed of the molecules do not

change, the expression — — is a constant. Hence the prod-

uct of the pressure and volume of a gas at a given tem-
perature is a constant. This is Boyle's law.

Heat is defined as molecular motion. The speed of the
molecules determines the temperature. If the speed of
the particles in the cube increases, it is evident that the
number of blows will increase; that is, the volume of
the gas remaining constant, the pressure increases with the

temperature. In equation (2) PF= ^^ , if t; varies, V

or P (or both) must vary. That is, if the pressure re-
main constant, the volume of the gas will vary as the
temperature (Charles' law). If the volume of the gas re-
mains constant, the pressure will vary as the temperature.
Two masses have the same temperature when they do
not impart energy to each other. Two molecules, then,
at the same temperature, must have equal kinetic energy,
otherwise they would be hastened or retarded on collision.

^mv =-m'v'^. (3)

From (1) it follows that two gases would have equal
pressures when

KINETIC THEORY OF 0A8E8. 403

If the temperatures are equal, it follows from (3) that
mv^ = m'v^\ (5)

Eliminating the equal factors from both members of
(4) we have

n Til ,^.

If the volumes are equal, F=F^, hence n^v! , That is,
equal volumes of gases under similar conditions of tem-
perature and pressure, contain equal numbers of mole-
cules. This is Avogadro's hypothesis.

•These equations give the explanation of other physical
phenomena. Thus, from (4) we get

v':v::yj^: y^, (7)

mn being the total mass of the gas 5 ^ is the density ;

hence the velocity varies inversely as the square root of
the density. (Law of diffusion.)

We can also find the absolute speed. From (1) : •

^=V-

3Pr

(8)
mn

One liter of oxygen (F) weighs 1.428 grams (mri)
under standard conditions. Standard pressure of 76 cm.
of mercury is equivalent to about 1000000 dynes per
square centimeter. Substituting these values.

-V—

1000000 X 1000

1.4

we find the speed per second to be about 46000 cm.
(over a quarter of a mile), or about as fast as a rifle bullet.
Hydrogen, in accordance with (7), should move four
times as fast.

404 APPENDIX.

These equations do not take into consideration the size
of the molecules or the actual space they occupy. A study
of gases under high pressure shows that the diameter of
the hydrogen molecule is less than one ten-millionth of a
centimeter, and that the molecules occupy about one
thousandth of the volume.

In the case of liquids, the molecules have considerable
freedom of motion, as is evident in the diffusion of liquids.
The molecules are probably closer together than in gases,
thus resisting pressure, and their mutual attraction is suffi-
cient to prevent their rapid separation. When a solid^ is
dissolved in a liquid, it acts physically very much as if it
were a gas occupying the volume of the solvent.

In solids the motion is much more circumscribed and
the molecules are still closer together. Diffusion and
evaporation are less in the case of solids than in liquids.
The application of conceptions derived from the kinetic
theory gives one mathematical expression applicable to
solids :

Specific heat x atomic weight = 6.3 (approximately),
which is Du Long and Petit's law.

INDEX

References are to pages.

Heavy-face nnmerals Indicate the principal
reference.

Abrasives ....... 260, 320

Absolute temperature ... 9
change of Centigrade to . . 10

zero of 10

Absorbent cotton 390

Acetaldehyde 386

Acetic acid 385

fermentation 386

prodaction of 375

Acetylene 236,379

Acids 116

definition 116

general method for prepara-
tion of 66

strong and weak 117

Agrate 255

Air Chap. XIX, 170

a mixture, proofs 171

composition of 170

Alabaster 274

Alcoholic beveragres . ... 383

Alcohols 881

denatured 384

ethyl 382

grain 382

methyl 381

wood 381

Aldehydes 385

Aluminates 319, 321

Aluminum . . . Chap. XXX, 817

acetate 322

alloys 320

bronze 299, 320

chemical properties .... 318

compounds 320

test for 320

Aluminum — Continued

double sulphates 321

foil 319

hydroxide 321

occurrence 317

oxide 320

paint 319

physical properties .... 318

preparation 317

silicates 317,323

uses 319

Alums 821

chrome 362

iron 343

Amalgramatlon process . . 312

Amalgrams 285

Amethyst 255

Ammonia 179

chemical properties .... 182

commercial production . . . 179

formation in nature .... 179

fountain 182

physical properties .... 182

preparation 181

uses 184

Ammonium alum 321

hydroxide 181

nitrate 185

nitrite 168

radical 183

sulphate 181

Analysis, as type action ... 94
Anesthetics

chloroform 381

ether 389

methyl chloride 380

nitrous oxide 185

406

INDEX,

References are to pages.

Anhydride, acid, defined . . 156
Anti-friction metals .... 351
Antimony 204

alloys 205

Aquafortis 188

Aqua regria 1^2

Aqueous tension, defined . . 15

table of 18,400

Argron 174

Arsenic 203

compounds 204

Artificial ice 184

Asbestos 278

platinized 159

Asphyxiation 251

Atmosphere . . Chap. XIX, 170
Atomic hypothesis .... 51
Atomic weifiThts . .Chap. XI, 79

definition 79

table of 396

Atoms Chap, y II, 51

definition 55

in molecule of gaseous ele-
ments 74

graphic demonstration . . 76

number in molecule .... 82

Auric chloride 313

Avosradro's hypothesis . . 74

exercises involving .... 78

explained by kinetic theory . 403

Babbitt Metal 205

Bacteria, nitrifying 195

Bakingr powder 248

Baklnsr soda 134

Ballastite 391

Barium chloride 163

peroxide 47

Barometer 12

Base 116

definition 116

sodium hydroxide as a typical 105

Base bullion 353

Basic lining process, for iron 334

Bauxite 317

Beer 383

Benzene 378

series 379

Benzine 376

Benzol 378

Bessemer process, for copper . 294

for iron 333

Bismuth 206

Blast-furnace, for copper . . 293

for iron 331

Bleaching, by chlorine ... 62

by hydrogen peroxide ... 48

by sulphur dioxide .... 157

powder 63, 276

Blueing 343

Blueprints 311,844

Boiler scale 268

Boiling-point, effect of dis-
solved substances on . . 113

of elements, table 396

Bone ash 200

Boneblack 232

Bone oil 376

Borate 262

beads 263

uses 263

Bordeaux mixture .... 301

Boric acid 261

properties 262

uses 262

Boron Chap. XXV, 261

Boyle's law 11

explained by kinetic theory . 402

Brand. 200

Brandy 383

Brass 283

Bread, raising of 248

Brickmaking 323

Brimstone 143

Brine, electrolysis of ... . 57
Britannia metal .... 205, 351

Bromine 208

bleaching action 211

chemical properties .... 210

ions, test for 214

occurrence 208

INDEX.

407

References are to pages.

Bromine — Continued

physical properties .... 210

preparation 209

replacement of 213

uses 211

water 210

Bronze 283

alnminum ........ 320

BtimerB, self-lighting .... 314

Bumlnsr, chemical nature of . 6

Butane 378

Calmerorm stone 255

Calclte 268

Calcium . . . Chap. XXVI, 265

bicarboniEite 247

carbide ....... 236

carbonate 267

occurrence ...... 267

properties 268

uses 270

varieties 267

chemical properties .... 266

hydroxide 272

properties 272

uses 273

Ught 32

burner '. 272

oxide 270

manufacture 270

propeirties 271

uses ........ 271

phosphates 203, 275

physical properties .... 266

preparation 265

sulphate ........ 274

uses 267

Calico printiner . . . . . . 322

Calomel 287

Caramel 394

Carat 313

Carbides 236

Carbohydrates 390

Carbolic acid 376

Carbon . . . Chap. XXIII, 227

allotropic forms 233

amorphous 2^33

chemical properties .... 235

Carbon — Continued

compounds, organic . . .

Chap. XXXV, 874

dioxide 244

chemical properties . . . 246
cycle in nature .... 173

. mair 173

S^hysical properties . . . 245

^fceparation 244

test for 247

uses 245,248

disulphide 235

importance in nature . . . 227

monoxide 249

chemical properties ... 251
physical properties . . . 250

preparation 249

uses 251

occurrence 228

oxides of . . Chap. XXIV, 244

uses 237, 242

Carbonates 247

Carbonic acid 247

Carborundum .... 237, 260

furnace 261, 262

products 260

Camallite 278

Camelian 256

Cast iron 829

composition 332

gray 332

manufacture 329

properties . ' 332

uses 3:^2

white ......... 332

Caustic potash ...... 131

soda .......... 131

Cave formation 269

Celluloid 192, 391

Cellulose 390

Cement, hydraulic . . . 273,325
hardening of .... 273, 325

Chalcedony 255

Chalcopyrite 292

Charcoal 232

Charles' laTV 9, 10

explained by kinetic theory . 402

408

INDEX.

S^erences are to pages.

Chemical ohanfire 1

definition 6

Chile saltpeter 137

China 324

Chloride of lime 276

Chlorides 68

insoluble .68

test for '. « 68

Chlorine Chap. VUI, 57

chemical properties .... 59

occurrence 57

physical properties .... 59

preparation 57

equations 96

uses 61

Chloroform 381

Chlorplatinic acid 314

Choke damp 244

Chromates 362

conversion into dicliromates . 363

Chrome alum 362

yellow 358

Chromic compounds . . . 868

acid 362

anhydride 362

chloride 363

oxidation of 363

oxide 362

Chromite 361

Chromium . . Chap. XXXIII, 861

compounds 363

occurrence 361

oxides 362

preparation 361

properties 361

sulphate 362

uses 361

valence 363

Chromous compounds . . . 363

Cinnabar 285

Clarke, T.W., quoted. ... 56

Clay 323

CoafiTTilum in water purifica-
tion 302,323

Coal 228

anthracite 230

Coal — Continued

bituminous ....... 230

cannel 230

composition of 228

formation in nature .... 230
gas (illuminating) . . . 179, 375

tar 375

Cobalt Chap. XXXI. 844

chloride 345

cyanides, potassium .... 346

extraction 345

glance 345

nitrate 345

test for aluminum . . . 320

for magnesium . . . 279

for zinc 282

ores 344

properties 345

speiss 345

sulphide 346

Coins, copper (bronze) .... 351

gold 313

nickel 347

silver 306

Coke 282,375,377

Collodion 390

Combininsr weierhts . Chap. V, 86

method of determining ... 38

Combustion - 20

definition 25

relation of air to 21

Compound, definition .... 6

Concrete 826

Condenser 41

Conservation of matter, law

of 51

Contact process for sulphu-
ric acid 160

Converter, for copper .... 294

for iron . 333

Copper . . . Chap. XXVm. 291

alloys 299

blister 296

compounds 299

matte 294

metallurgy 292

native 291

INDEX.

409

R^erences are to pages.

Copper— Continued

occurrence 291

oxides 299

poling of 296

properties 298

refining, electrolytic .... 297

sulphate 900

preparation 900

properties 301

uses 301

uses (of copper) 299

valence 299

Copperas 342

Coquina 267

Cordite 391

Correction of srases . Chap, n, 8

for difference in level ... 13

for pressure 12

for pressure of water vapor . 14

for temperature 10

Corrosive sublimate. ... 287

Corundum 317

Courtois 214

Cream of tartar 248

Crockery 324

Cryolite 222,317

Crystals 46

Crystallization, water of . . 45

Cupellation 305

Cyanide process 312

Dalton,Jolm 52

Davy, Sir Humphry . 103, 185, 265
Decomposition, as type action . 94
Decrepitation defined ... 133
Definite proportions, law of . 35
explained by atomic hypothe-
sis 53

Deliquescence 47

Density, of gases 79

Depolarizer, manganese dioxide

as 360

Destructive distillation . . 874

of bones 375

of coal 180 375

of wood 374

Dextrine 392

Diamond 234

artificial 234

Diastase 392

Diatomaceous earth .... 255

Dichromates 362

conversion into chromates . . 363

reduction of 363

Disinfectants

bleaching powder 63

chlorine 63

formaldehyde 386

hydrogen peroxide .... 48

sulphur dioxide 158

Dissociation of electrolytes . 113

Distillation 41

destructive, defined .... 232

fractional 383

of petroleum

Ddbereiner

Dolomite

DTilongr and Petit, law of . .
Dutch process for white lead

Dyeingr

Dynamite

Ecuiihenware . .
Eflaorescence . .
Electric furnace
Electrolysis, defined

376
365
278
404
356
322
197

324

237

110

explanation of 114

of water 26

Electrolytes, defined .... 109
chemical activity of ... . 113

dissociation of 113

Electrolytic copper, produc-
tion of 297

Electroplatiner 807

gold 313

nickel 347

silver 307

Elements, definition .... 6
physical constants of . . . 396

table of 396

Emery 320

Epsom salts 279

410

INDEX.

Se/erenees

Bqustions, chemical, Chap. XIII, 92

balancing of 93

calculation of relative weight!)

from 97

significance 92

Bquivalent weifiThts .... 85

method of determining ... 38

Esters 2

Btchinflr of fflass 223

Ethane 378

substitution products ... 382

Ether 389

Ethereal salts 386

Explosive erelatlne .... 391

Explosives 196

high 391

Families of Elements ... 369

Fats 387

FehllnsT's solution 300

Felspar 257,323

Fermentation 382

acetic 386

Ferric salts 840

ammonium citrate .... 344

chloride 342

ferrocyanide 343

hydroxide 329,341

oxide 340

reduction of 124

sulphate 342

tannate 342

Ferricyanides 343

Ferrocyanides 343

Ferro-maneranese 360

Ferrous salts 840

carbonates 328

chloride 341

ferricyanide 344

hydroxide 341

oxidation of 125

oxide 340

sulphate 342

tannate 342

Fertilizer 139,276

Films, photographic 308

Filter-paper 391

are to pages.

Firebricks 334

Fire-damp 377

Fire extinsruishers .... 248

Flame, candle 238

gas 241

Flashllsrht powder . . 270, 326

Flint 255

Fluorine 220

peculiarities 220

preparation 221

properties 222

Fluor-spar 222

Flux 270,330

Fool's srold 329

Formaldehyde 385

Formalin 385

Formic acid 385

Formulas .... Chap. XII, 86
calculated from percentage

comi)osition 87

meaning of 86

problems in determination of 91
Fractional distillation ... 383
Freezingr mixtures .... 43
FreezinfiT point, effect of dis-
solved solids on Ill

Fructose 392

Funsricide, copper sulphate as . 301

Furnace, blast, for copper . . 293

blast, for iron 329

electric 237

open hearth 334

poling 296

reverberatory 363

roasting 281

Fuse wire 355

Fusible metals .205

Galena 352

Qas carbon 232

Qaseous elements 78

number of atoms in molecule . 75

size of molecule 404

velocity of molecule .... 403

Gases, definition 8

causes of change in volume . 8

INDEX.

411

References are to pages.

Oases — Continued

kinetic theory of 401

standard conditions for meas-
urement 11, 12

weight of liter, table . . . 399
Gas, illuminating, coal . . . 179, 375

natural 377

water 251

Gasolene 376

Gas volumes, corrections of .

Chap. II, 8

for difiference in level . . . l.'i

for pressure 11

for pressure of water vapor . 14

for temperature 10

problems in 16, 17

Gay-Lussac 214

law of 73

German silver 2^^

Gin 384

Glass 257

Bohemian 257

coloring of 259

crown 258

cut 258

etching of 223

flint 2.59

furnace 257

plate 258

window 258

Gluclnum 366

Glucose 382,392

test for 300

Glycerine 387

Gneiss 257

Gold Chap. XXIX, 311

chloride 313

fineness 313

leaf 312

metallurgy 312

mining 311

occurrence 311

properties 312

uses 313

Gram-molecular volume . . 81

Granite 257

Graphite 233

Guano 275

Gimcotton 197

Gunpowder 138, 148

Gypsum 274

Halogens . . . Chap. XXII, 208

compared 218

heats of formation 218

relative replacement of . . . 219

tabular comparison .... 218

Hard waters 247

Hare, Dr. Robert 265

Heat of formation, defined . 219

influence on chemical action . 220

Heat of neutralization ... 120

Helium 175

Hematite 329

Hornblende 267, 278

Horn silver 304

Humidity of air 173

indicators 345

Hydraulic cement .... 325

Hydraulic mininfir 311

Hydriodic acid 217

Hydrobromic acid . . . .211

preparation 211

properties 212

Hydrochloric acid . . Chap. IX, 65

chemical properties .... 67

composition by volume ... 69

physical properties .... 66

preparation 65

equations {¥>

uses 69

Hydrofluoric acid 222

Hydrofluosilicic acid ... 261

Hydrofiren Chap. IV, 26

chemical properties .... 30

number of atoms in molecule 75

peroxide 47

preparation 47

properties 47

uses • . 48

physical properties .... 29

preparation 26

equations 94

sulphide . . . Chap. XVII, 148

412

INDEX,

R^erences are to pages.

Hydrofiren — • Continued

sulphide, chemical properties 150

occurrence 149

preparation 149

physical properties . . . 149

uses 151

uses (of hydrogen) 31

Hydrolysis 135

Hygrroscopic substances . . 47

Hypo 308

Ice

artificial, manufacture . . .

Iceland spar

IlluminatiniT SraSt coal . . .

water gas

Inert grases in atmosphere .

position in periodic classifica-
tion

Infusorial earth

Ink

Insolubility, in relation to
double decompositions .

Invertase

Iodides

Iodine

chemical properties ....

physical properties ....

preparation

starch test

tincture of

uses

Iodoform

Ionization, hypothesis ....

effect of dilution on ... .
Ions 114

and atoms distinguished . .

charges carried by ....

common, table of

Iridium

Iron Chap. XXXI,

Bessemer

carbide 332,336

cast * . .

chlorides

classification of

compounds

42
184
268
179
251
174

371
255
342

121
3i>4
217

214
216
215
215
216
215
216
381

114
116

,115
115
122
122
313

328
333

,337
330
342
338
340

Iron — Continued

ferricyanides 343

ferrocyanides 343

galvanized 283

hydroxides Ml

occurrence 328

ores 328

oxides 310

properties of pure 339

Russia 341

rust 341

sulphates 335

tabular comparison with steel 348

tannates 342

valence 340

wrought 332

Jasper

255

Kaolin 323

Kelp 216

Kerosene 376

Kiln, lime 270

pottery 324

Kindling: temperature ... 22

definition 25

Kinetic theory of grases . . 401

Avogadro's hypothesis . . . 403

Boyle's law 402

Charles' law 402

Dulong and Petit's law . . . 404

law of diffusion 403

statement 401

states of matter 404

Krypton 175

Lakes , ... 322

Lampblack 2;W

Laugrhingr gras 185

Lavoisier, experiments of, on

heating metals in air . . 4, 5

Lead .... Chap. XXXII, 862

acetate 357

basic carbonate 356

burning 32

chromate 358

compounds 355

effect on water 351

INDEX.

413

<^'

B^erences are to pages.

Iiead — Continued
hydroxide ....
metallurgy ....
oxides

pig

poisoning

properties

red

reduction, electrolytic

sulphide

tree

uses

white

Ligrnite

Lime

air-slaked ....

chloride of ... .

kiln

light

manufacture of. . .

milk of

properties ....

slaked

uses

water

limestone

caves

Limonite

Liquids, definition . .

Litharge

Lithium

Lockyer

Lodestone

Lubricatingr oils . .
Lunar caustic . . .

32,

367
352
355
353
354
353
356
353
352
354
354
356
230
270
272
275
270
272
270
272
271
272
271
272
267
268
329
8
355
107
175
340
377
307

Magrnallum 320

Magrnesite 278

Magnesium . .Chap. XXVII, 278

compounds 279

test for 279

Malachite 292

Malt 383,392

Maltose 392

Mangranates 361

Mangranese . Chap. XXXIII, 360

dioxide 360

Mangranese — Continued

dioxide as catalytic agent . 19
dioxide as oxidizing agent . 58

Marble 267

Marl 325

Marsh-gra^ 377

Massicot 355

Matches 202

safety 204

Matte, production of .... 293

Matter, definition 8

Melting: points of elements,

table 396

Mendelejeff 366

Mercuric chloride 287

Mercurous chloride .... 287

Mercury . . . Chap. XXVII, 285

chemical properties .... 286

compounds 287

occurrence 285

physical properties .... 285

separation 285

uses 286

Metals heated in air .... 2

Meteorites 328

Methane 377

Methyl alcohol ...... 381

chloride 380

Meyer, Lothar 366

Mica 257

Minium 356

Mirrors, making of . . . 286, 30T

Moissan 221

Molecular composition . . .

Chap. X, 73
Molecular weigrhts Chap. XI, 79

definition 81

determination of . . . . 80, 81

problems 84

Molecules .... Chap. VII, 61

definition 55

Molybdenum 337

Monochlormethane .... 380

Mordants 322

aluminum hydroxide as . . 322

414

INDEX,

Refer encen are to pages.

Mordants— Continued

copper sulphate as ... . 302

stannous chloride as . . . . 351

Mortar 273

hardening of 273

Multiple proportions, law of . 48

explained by atom ic hypothesis 53

Muriatic acid 66

Naphtha 376

Naphthalene 376

Nascent state 63

Natural gas 377

Neon 175

Neutralization, defined ... 105
explained by ionization hy-
pothesis 118

heat of 120

products of 119

Newlands 365

Nickel . . . Chap. XXXI, 346

alloys 347

ammonium sulphate .... 347

carbonyl 347

coins 347

extraction 346

ores 346

plating 347

properties 346

salts, color of 346

steel 347

sulphate 347

uses 346

Nitrates 193

properties li)4

test for 194

Nitre beds 138

Nitric acid 188

action with metals .... 191

chemical properties .... 190

physical properties .... 189

preparation 188

reduction products .... 190

uses 192

Nitric anhydride 188

Nitric oxide 186

Nitrides 170

Nitrification 195

Nitrites . .' 138

Nitrocelluloses 390

Nitrogen . . . Chap. XIX, 167

chemical properties .... 168

compounds . . Chap. XX, 179

cycle in nature 173

group, elements of, Chap. XXI, 200

tabular comparison . . . 206

occurrence 167

peroxide 186

physical properties .... 168

preparation 167

Nitroglycerine .... 196, 387

Nitrous anhydride 188

oxide 185

Non-electrolytes 109

Occlusion 313

defined ........ 29

Ochre, yellow . 340

Octaves, law of 366

Oil of wintergreen .... 387

Oils 387

Oleic acid 387

Onyx 256

Opal 255

Open-hearth process . . . 334

Organic acids 885

acetic 385

formic 385

oleic 387

palmitic 387

properties 385

stearic 387

vinegar . 386.

Organic chemistry, defined . 374
Organic compounds . . .

Chap. XXXV, 874

Orpiment 204

Osmium 313

Oxalic acid 250

Oxidation 20

and reduction 124

slow 21

definition 25

INDEX.

415

References are to pages*

Oxides 20

aud acids of salpbur . . .

Chap. XVIIl, 154

definition 25

of nitrogen 185

Oxygren Chap. Ill, 19

chemical properties .... 20

number of atoms in molecule 75

occurrence 23

physical properties .... 20

preparation 19

equation 94

relation to life 23

Oxy-hydrogren blowpipe . . 31

Ozone 24

Painter's colic 354

Palmitic acid 387

Paper making: 391

Paraffin . . . * 377

series 378

Paris grreen 204

Parkes' process for silver

extraction 304

Peat 230

Percentagre composition, de-
termined from formula ... 87
Periodic Law . Chap. XXXIV, 365
groups of elements .... 369
history of development . . . 365

long periods 369

position of hydrogen .... 371

short periods 367

statement of law 367

table 368

vacant spaces 371

value 371

Permangranates 361

Petrified wood '2m

Petroleum 376

distillation 376

products 376

refining 377

Pewter 351

Phosphates 203

Phosphoric acid 203

anhydride 203

oxide 203

Phosphorite

Phosphorus anhydride . . .

oxide

Phosphorus . . Chap. XXI,

alio tropic forms

chemical properties ....

compounds

occurrence

physical properties ....

preparation

red

yellow

Photogrraphy

blue-priuts

developing

fixing

negative

plates

positive

tiu-types

toning

Physical changre

Physical constants, table of .

Pigr-iron

Plaster of Paris

Platinized asbestos . . . 160,
Platinum . . . Chap. XXIX

black

compounds

occurrence

properties

spongy

uses

Plugrs, fusible

Poling: of copper

Porcelain

Potassium . . . .Chap. XIII,

alum

carbonate

chemical properties ....

chlorate

chloride

chromate

compounds . . Chap. XVI,
test for

cyanide 307,

dichromate

275

203
203

200
201
202
203
200
201
200
201
201

808%
311-

309-
309-
308-
310*
310-
310 •

1
396
330
274
313

SIS
313
314
313
313
313
314
206
296
324

107
321
134
107
19
133
362

128
315
312

416

INDEX.

lieiferences are to pages.

PotasBlum — Continued

ferricyanide 343

ferrocyanide 343

hydroxide 128

preparation 128

properties 130

uses 131

manganate 361

nitrate 137,193

occurrence 137

preparation 138

uses 138

permanganate 361

physical properties .... 107

preparation 107

silicate 257

Pottery 323

Powder, smokeless 197

Producer gras 252

Propane 378

Protelds 167,228

Prussian blue 343

Prussiates of potash . . . 343

Puddlingr process 332

Pyrolusite 360

Quartz 255

Quicklime 270

Quicksilver 285

Quinine 374

Radium 176

Rayleigrh 174

Beactingr quantities, calcu-
lated from equation . . 97

problems involving .... 101
Beactingr volumes, calculated

from equations ... 98

problems involving .... 101

Beactingr weigrhts 35

and volume weights of gases,

relation between ... 73

definition 38

method of determining ... 38

Beactions in solution . . . 120

Bealgrar 204

Bed compound, Lavoisier's

experiment with ... 4

Bed pigrments 340

Beducingr agrents, definition . 31

Beduction, definition .... 31

Befrigreratingr process . . . 184

Beplacement, as type action . 94

double, as type action ... 96

of bromine by chlorine . . . 213

Bochelle salt 249, 300

Bock crystal 255

Bose's metal 205

Bougre 340

Buby 317

Bum 384

Bust, iron 341

Saccharose 392

Sal ammoniac 181

Salt, common 131

Salts, definition 68

solubility table 398

Saltpeter 137,193

Chile . 137,193

Sand 256

Sandstone 255

Sapphire 317

Scheele 200

Schweitzer's reagrent ... 390

Segrgrar 324

Seltzer 248

Shale 325

Shot 355

Siderite 328

Sienna, raw and burnt . . . 340

Silica 255

Silicates 257

Silicic acid 261

Silicon .... Chap. XXV, 256

carbide 260

dioxide 255

properties 256

uses 256

varieties 255

fluoride 261

INDEX.

417

B Terences are to pages.

Silicon— Conrtntwei

hydride

occurrence

properties

Silver .... Chap. XXIX,

alloys

amalgams

bromide

chemical properties . . . .

compounds

cupellation of

halogen salts of

horn

metallurgy

nitrate

occurrence

plating

refining, electrolytic . . .

sterling

uses

Slagr

Slakingr of lime

Slate

Smeltinsr, of copper ....

of iron

Smithsonite

Smokeless powders . . . .
Soap 131,

composition

hard

hard water and

manufacture

soft

Soda

Soda water

Sodium .... Chap. XIV,

bicarbonate

preparation

carbonate

preparation

uses

chemical properties ....
chloride

crystalline form ....

extraction from sea water

occurrence

properties

261
255
255
804

306
306
212
306
307
306
308
304
304
307
304
307
305
306
306
330
272
257
293
329
279
391

887
387
387
389
388
387
lU
245

108

184
134
135

134
134

i;«

104
131

1:53
132
131
132

Sodium — Continued

purification 132

uses. 133

compounds . . Chap. XVI, 128

dichromate 362

hydroxide 128

as a typical base .... 105

preparation by electrolysis 129

preparation by lye process 130

properties 130

uses 131

nitrate 137, 193

physical properties .... 104

preparation 103

silicate 267

stearate 387

thiosulphate 309

Solder 351

Solids, definition 8

Solubilities, table of .... 398

Solubility, factors 43

in water, conditions deter-
mining 43

rules 399

Solute, definition 42

Solution Chap. XV, 109

definition 42

saturated 42

supersaturated 45

Solvay process 134

Solvent, definition 42

Specific gravity, gases . . 79, 80

of elements, table 396

Spectra, frontispiece

Spectroscope 106

Spectrum analysis .... 106

Spelter 281

Spiegreleisen 334

Spontaneous combustion 22

Stalactites 270

Stalagrmites 270

Stannic chloride 352

sulphide 352

Stannous chloride 351

sulphide a52

Starch 892

formation in nature .... 246

418

INDEX.

References are to pages.

Stassfurt deposits .
States of matter

Steam

Btearates . . . .

134,208
. . 8
. . 41
. . 387

Stearic acid 387

Stearine 387

Steel 388

Bessemer {>r<>ce8s 333

basic lining pr<»ce8s ... 331

cementatiou process .... 337

chrome 361

crucible process 336

manganese 337

nickel 347

open-hearth process .... 334

properties 348

self-hardening 337

tabular comparison with iron 348

tempering of 337

uses 338

Storafire batteries 356

Stucco 275

Styptic, alum as 321

Sublimation, defined .... 216

Substitution products ... 379

Sufirar 892

barley 3^

beet 393

cane 3J)3

fruit 392

grape 392

maple 393

refining 393

Sulphates, insoluble . . . . 16^^

test for 163

Sulphides 147

preparation of 150

Sulphites 156

Sulphur .... Chap. XVII, 141

allotropic forms 143

amorphous 145

chemical properties .... 146

commercial forms 142

dioxide 154

bleaching by 157

chemical properties ... 156

physical properties . . . 165

preparation 154

Sulphur — Continued

flowers of 142

milk of 146

native 141

occurrence 141

orthorhombic 144

oxides and acids of Chap. XVIII, 154

plastic 145

preparation 141

prismatic 144

roll 142

springs 151

trioxide 158

preparation 159

uses 148

Sulphuric acid 160

chemical properties .... 161

manufacture 160

physical properties .... 161

uses 163

Sulphuric anhydride ... 159

Sulphurous acid 156

Sulphurous anhydride . . . 156

Superphosphate of lime . . 275

Symbols Chap. XII, 85

meaning of 85

of elements, table 396

Sympathetic ink 345

Synthesis, as type action . . 94

Talc 257

Tallow 387

Tannic acid 342

Tellurium 311

Tempering: of steel .... 337

Thomas-Gilchj-ist process . 334

Thomas slagr 334

Tiles . 323

Tin Chap. XXXH, 850

alloys 351

block 350

chlorides 351, 352

cry 360

crystals 351

foil 351

metallurgy 360

oxide 350

properties 350

INDEX.

419

References are to pages.

Tin— Continued

sulphides 352

uses 351

valence 352

Tin-types 310

Tinware 351

Triculs of elements .... 365

Trichlormethane 381

Tri-iodomethane 381

Tungsten 337

Tumbuirs blue 344

Turpentine 374

Type metal 205

Ultramarine 394

Umber, raw and burnt .... 340

Valence 123

definition 123

of elements, table 396

relation to charges carried by

ions 124

yariations in 124

Vaseline 377

Vichy 248

Vinegrar 386

Vitriol, blue 300

green 342

oil of 161

Volatile compoimds, table of 399
Volatility, in relation to double

decomposition . . . .121

table 399

Washingr soda 134

Water . . . Chaps. V, VI, 34. 40

decomposition by iron ... 28

decomposition by potassium . 27

decomposition by sodium . . 27

determination composition of 34

distillation of 41

electrolysis of 26

formation of, equation ... 94

gas 251

hard, permanent 248

hard, temporary 247

action with soap .... 389

of crystallization 45

Water— Continued

physical properties .... 40
purification of

by aluminum hydroxide . 323

by copper sulphate . . . 302

by distillation 41

synthesis, gravimetric ... 35

volumetric 34

vapor, in air 172

pressure of, table . . .18, 400
Weigrlit changre on heatingr

metals in air .... 3

Welding: 333

Welsbcbch burner 241

Whiskey 384

White metal 351

Wine 383

Wood . , 390

pulp 391

Wood alcohol 381

production 375

Wood's metal 205

Wrougrht iron 882

composition 333

manufacture 332

properties 333

uses 333

Xenon

175

Zinc .... Chap. XXVII, 279

blende 279

chemical properties .... 281

compounds 283

test for 282

forms of 281

hydroxide 283

metallurgy 279

occurrence 279

oxide 283

physical properties .... 281

purification 281

sulphate 284

sulphide 284

precipitation of ... . 284

uses 283

Zincite 279

Zymase 382, 394

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THIS has been for many years a standard class-room text-
book on the subject for secondary schools, and has been
thoroughly tested by its long service. Its aim is to formulate
clear statements of laws and principles ; to illustrate them amply
by simple experiments and appropriate problems ; and to observe
a logical sequence of topics, so that the pupil may pass from sub-
ject to subject with the aid of what he has already acquired. The
experiments have been carefully selected with a view to their
availability for high school work as well as for their scientific
value.

Professor C. F. Brackett, Princeton University: I have examined this
work with care and with pleasure, for it presents the fundamental prin-
ciples of phjrsics with exactness and with clearness.
63

SCIENCE.

Practical Physiography.

By Dr. Harold Wellman Fairbanks, of Berkeley, California.
8vo, cloth, 570 pages. 403 Illustrations. Price, $1.60,

IN this volume the author has tried to work out a practical, con-
crete treatment of the subject of Physical Geography. The
book is intended as an aid to study, not as a compendium of
information ; consequently a description of the world as a whole
is omitted. Attention is devoted specifically to the region of the
United States, and the typical examples afforded by it are studied
as representatives of universal processes.

The purely descriptive method has been discarded as far as
practicable, the object being to lead the student to investigate
and find out for himself. Instead of being told everything, he is
asked to use his observing and reasoning powers.

No separate chapters are devoted to the relation between
physical nature and life, but, instead, this relation is brought
out in its appropriate place in connection with each topic through-
out the book. Such an arrangement, it is believed, will make the
whole matter much more vital.

Another feature which the author trusts will meet with favor
from the practical teacher is the distribution of the questions and
exercises throughout each chapter, in close connection with the
descriptive portions to which they refer. The placing of ques-
tions and exercises by themselves at the close of each chapter, as
is done in many text-books on the subject, puts a premium upon
mere memorizing of the text and the omission of all practical
work. It is not expected, however, that the pupils will be able
to answer all the questions without aid and direction from the
teacher.

The illustrations are a marked feature of the book. Photo-
graphs have been used wherever possible, as they appeal with
much more force to pupils of high school age than do diagrams
or sketches. Most of the views are fi^om the author's own
negatives.

SCIENCE.

Herbarium and Plant Descriptions.

THIS is an herbarium and plant record combined, enabling
the student to preserve the specimens together with a
^cord of their characteristics.

A sheet of four pages is devoted to each specimen. The first
page contains a blank form, with ample space for a full descrip-
tion of the plant, and for notes of the circumstances under
which it was collected. The pressed specimen is to be mounted
on the third page, and the entire sheet then serves as a species-
cover. Each portfolio contains fifty sheets, which are separate,
so as to permit of scientific rearrangement after mounting the
specimens.

The preliminary matter gives full directions for collecting,
pressing, and mounting plants, as well as a synopsis of botanical
terms.

The portfolio is strong, durable, and attractive in appearance.

In the class-room and in the field this work has been found
helpful and stimulating. It encourages observation and research,
and leads to an exact knowledge of classification.

Professor D. P. Penhallow, McGill University, Montreal, Can,: The idea
is a good one, and well carried out. I am sure it will prove most useful
in the botanical work of schools and academies, for which I would
strongly recommend it.

Professor G. H. Perkins, University of Vermont, Burlington^ Vt, .• It is the
best thing of the sort I have seen ; very attractive and very helpful to
beginners in calling attention to points that would be overlooked.

Professor B. P. Colton, Normal University, III, .• It is a very ingenious ar-
rangement, and neatly gotten up. It speaks well for the publishers, as
well as the designer. It is the neatest scheme of the kind I have seen.

0. D. Robinson, Principal of High School, Albany, N, Y, : It appears to me
to be a very complete arrangement, admirable in every respect, and vex^
moderate in price.

P. S. Hotalingi Formerly Principal of High School, Framingham, Mass. :
Last year's work in botany was made so much more interesting and valu»
ble by the use of the Herbarium that we find it now a necessity.
69

SCIENCE,

First Principles of Chemistry.

By Raymond B. Brownlee, Far Rockaway High School; Robert
W. Fuller, Stuyvesant High School; William J. Hancock, Eras-
mus Hall High School ; Michael D. Sohon, Morris High School;
and Jesse E. Whitsit, DeWitt Clinton High School; all of New
York City.

THIS manual includes a treatment of the common elements
and their important compounds, together with a full dis-
cussion of the fundamental theories of chemistry. These theo-
ries are simply, clearly, and accurately stated without being buried
in a confusing mass of detail. The recent developments in chem-
ical theory are given due recognition.

In the development of laws and hypotheses, the historical order
is followed, the experimental facts in each case being described
before stating the discovery that resulted from them. Theoret-
ical topics are introduced as soon as the pupil is able to take in
their full significance; this, on the ground that to delay their
presentation is to deprive the pupil of a very useful tool in his
acquisition of chemical ideas.

This book marks a step in advance by^ts omission of much
material that has^found a place in other elementary manuals
through tradition rather than for its real value. On the other
hand, a number of the metallic elements, sometimes neglected, are
given the thorough treatment that their industrial importance
deserves. One of the chief aims of the descriptive matter is to
show the student the many points of contact between the life
about him and the science he is studying.

An important feature of the book is the brief summary and the
test exercises given at the end of each chapter. The summary
is a series of pithy statements emphasizing the essentials and
aflfording systematic review.

This book has been prepared by five teachers of long experi-
ence in both college and secondary school work. Three years of
careful discussion and revision by the authors have produced a
unified text-book especially planned for the beginner in chemistry.

'i\ «•

This book should be returned to
the Library on or before the last date
stamped below.

A fine of five cents a day is incurred
by retaining it beyond the specified
time.

Please return promptly*

OUEJKC-^^^

T'l-r

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