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INTERNATTONAL ATOMIC WEIGHTS (1922)
O = 16.
Aluminium Al 27.0
Antimony Sb 120.2
Argon A 39.9
Arsenic As 74.96
Barium Ba 137.37
Bismuth Bi 209.0
Boron B 10.9
Bromine Br 79.92
Cadmium Cd 112.40
Caesium Cs 132.81
Calcium Ca 40.07
Carbon C 12.005
Cerium Ce 140.25
Chlorine CI 35.46
Chromium Cr 62.0
Cobalt Co 68.97
Columbium Cb 93.1
Cow)er Cu 63.57
D^prosium Dy 162.5
Erbium Er 167.7
Europium Eu 162.0
Fluorine F 19.0
Gadolinium Gd 157.3
Gallium Ga 70.1
Germanium Ge 72.5
Glucinum Gl 9.1
Gold Au 197.2
Helium He 4.00
Holmium Ho 163.5
Hydrogen H 1.008
Indium In 114.8
Iodine I 126.92
Iridium Ir 193.1
Iron ....Fe 65.84
Krypton Kr 82.92
Lanthanum La 139.0
Lead Pb 207.20
Lithium Li 6.94
Lutecium Lu 175.0
Magnesium Mg 24.32
Manganese Mn 64.93
Mercury Hg 200.6
o -le.
Molybdenum Mo 96.0
Neoaymium Nd 144.3
Neon Ne 20.2
Nickel Ni 68.68
Niton (radium emanation) Nt 222.4
Nitrogen N U.OOt
Osmium Os 190.9
Oxygen O 16.00
Palladium Pd 106.7
Phosphorus P 31.04
Platinum Pt 196.2
Potassium K 39. IC
Praseodymium Pr 140.9
Radium .Ra 226.0
Rhodium Rh 102.9
Rubidium Rb 85.46
Ruthenium Ru 101 .7
Samarium Sa 150.4
Scandium Sc 45.1
Selenium Se 79.2
Silicon Si 28.1
Silver Ag 107.88
Sodium Na 23 .00
Strontium Sr 87.63
Sulphur S 32.06
Tantalum Ta 181 .6
Tellurium Te 127.5
Terbium Tb 159.2
Thallium Tl 204.0
Thorium Th 232.16
Thulium Tm 169.9
Tin Sn 118.7
Titanium Ti 48.1
Tungsten W 184.0
Uranium U 238.2
Vanadium V 51 .0
Xenon Xe 130.2
Ytterbium (Neoytterbium) Yb 173.6
Yttrium Yt 89.33
Zinc Zn 66.37
Zirconium Zr 90.6
For a table of the Periodic System, see p. 278.
For a table of Atomic Numbers, see p. 649.
©&IIforn!a CcSl
C '
fM^«
BT THE SAME AUTHORS
SMITH'S LABORATORT OUTLINE
OF INTERMEDUTE CHEMISTRY
BT ALEXANDER SMITH
EXPERIMENTAL
INORGANIC CHEMISTRY
Fifth Edition, Revised
INORGANIC CHEMISTRY
Third Edition iigi?)* RewriUen
92s + xvi pp. With 166 Figures
A briefer Text on the same pUn
GENERAL CHEMISTRY FOR COLLEGES
Second Edition, Revised
662 + xi pp. With 138 Figures
LABORATORY OUTLINE OF COLLEGE
CHEMISTRY
266 + V pp. With 30 Figures
A Text-book of
ELEMENTARY CHEMISTRY
439 + viii pp. Wth 98 Figures and 6 Plates
A LABORATORY OUTLINE OF
ELEMENTARY CHEMISTRY
137 pp. With 18 Figures
New York, THE CENTURY CO.
London, GEO. BELL AND SONS
.ct.
SMITH'S
INTERMEDIATE CHEMISTRY
EEVISED AND REWRITTEN
JAMES KENDALL
Pbofebsor of Chemibtrt, Coluubia UNtvEBSiTT
EDWIN E. 8L0SS0N
Gditob of "Science SBBvicii," Author or "Cbkatitk
Chbmibtey," etc.
NEW YORK
THE CENTURY CO.
19??
Copyright, 1919, 1922,
BY
THE CENTURY CO
Printed in the U. S. A.
"c^.O?!
WHY STUDY CHEMISTRY?
CJhemistry began as a secret science. The early chemists con-
cealed their knowledge — and more often their ignorance — under
a cloak of symbols and ciphers of the most mysterious and awe-
inspiring sort. But now the Black Art has been opened to day-
light. The modem chemist is more anxious to tell people what
he knows than people are to listen to him. He still uses symbols
and has a fondness for long words, but these are designed to reveal,
not to conceal.
Still there lingers about chemistry something of the witchery
of its antiquity. It has the air of being much harder to under-
stand than it really is. The curious structural formulae of organic
compounds are no more difficult to work out than a Chinese puzzle
and quite as much fun.
Chemistry is especially fitted to give training in the scientific
method, for it is experimental from the start. Properly taught
— or rather properly learned — it inculcates self-reliance and
independence of thought. If the pupil will take the teacher's
word for the names of things and follow the advice of the book
as to what experiments to try, he can find out and think out the
most important part of the science for himself. He can work
out a system of analysis by testing known substances in a sys-
tematic way and then when he enters upon unknown mixtures
he can attack them with the courage of self-confidence. The
student of astronomy never gets a chance to handle a star, or
even an asteroid. But the substances that the young chemist
studies are always weighable, usually tangible, generally visible
and frequently smellable. The student of geology never has the
opportimity to make a mastodon and all he knows of a volcano or
a ge3rser is the picture of it. But the Freshman chemist makes
• ••
m
399130
iv smith's intekmediate chemistky
oxygen the first week and if he gets through the term without
making a volcano or geyser he is lucky.
This is not said to the disparagement of other studies. They
all have their peculiar merits which their professors are at liberty
to demonstrate. But chemistry has the advantage of most of
them in that it points toward the future and instigates to action.
The study of the history of man as written in books and of the
history of the earth as recorded in rocks gives one a sad sense of
irreparableness. What they deal with is dead and gone and all
that is left is their lesson which may or may not be applicable to
present problems. Meteorology is a discouraging science.
" People are always complaining about the weather but nobody
ever does anything about it." Astronomy reduces one to a state
of impotent awe. It is interesting of course to find out the struc-
ture of the solar system, but if you do not like it you cannot change
it. You cannot put Saturn in the place of Venus so as to get a
brighter evening star with two rings and eight sateUites running
around it. You cannot even alter the inclination of the axis of
the earth so as to give the polar regions a chance to grow bananas.
But when you find out the structure of a chemical molecule you
can alter it to suit yourself. You can substitute a bromine for a
chlorine atom and hook up a carbon chain into a ring.
In this field man is master of his material and the only limita-
tions to his power are his own ignorance and the innate intracta-
biUty of the elements. Berthelot calls chemistry the most crea-
tive of the sciences, because it penetrates most profoundly into
the nature of things and deals with the infinitesimal parts of
which all substances are composed. The chemist begins by taking
things apart (analysis); then he proceeds to put them together
again (synthesis). Sometimes he puts them together in a differ-
ent way and then he gets something quite different, perhaps some-
thing that never existed on the earth before. Some 300,000 dis-
tinct substances are described in chemical dictionaries and only
a small fraction of these are to be found in nature.
WHY STUDY CHEMISTBY? V
Chemistry is the science of power. All the energy of man and
beast and all the energy of machinery, except that furnished by
windmills and water wheels, is chemical energy and nearly all
of this is derived from two simple and similar chemical reactions,
the oxidation of hydrogen to form water and of carbon to form
carbon dioxide. Food and fuel plus the oxygen of the air pro-
vide the motive force of the modem world. The automobile
and the elephant, the airplane and the eagle, the whale and the
steamboat, man and his machines would alike be lifeless unless
empowered by the oxidation of hydrogen-carbon compounds.
Potential chemical energy is peculiarly effective in that it is so
compact, portable and permanent. A few bags can contain
enough of it to blow up a building or propel a projectile seventy-
five miles.
Chemistry has control of life and death. An animal or plant
deprived for a time of the proper chemical compounds will startre
or if it imbibes a minute amount of the wrong chemical com-
pounds it dies at once. The vital processes of the body are con-
trolled by chemicals, some of which are already known, and a
little more or less of them in the blood determines whether we
shall be tall or short, bright or dull, handsome or ugly, cheerful
or melancholy.
Chemistry is the science of terror and pity. Warfare has been
an intermittent branch of appUed chemistry ever since the inven-
tion of gunpowder, but it was not until 1914-18 that its terrible
power of destruction and devastation was realized. But chem-
istry may heal the wounds it causes, even with medicaments
drawn from the same unpromising protean stuff, coal tar, that
yields the high explosives and poison gas. Chemistry is the chief
weapon of man in his lifelong fight against the bacteria and pro-
tozoa that invade his body and ultimately destroy it. It is with
chemicals that he combats the insects that eat his crops and infect
him with disease. It is chemistry that brings relief to the sleep-
less and suffering and deprives surgery of its pain. Chemistry
vi bboth's intermediate chemistry
may, it is true, demolish a building, but it is chemistry that erects
it, even the steel-ribbed skyscrapers and mammoth monoliths
of concrete.
Chemistry is the democratic science. It bestows upon the
poorest what once were the gifts of kings; jewels and bright-
dyed garments, sugar and scents, vessels of porcelain and metal,
foreign fruits and out-of-season foods, books and pictures. Chem-
istry carries music into farmhouse and tenement on disks of resin
and reveals to multitudes from the celluloid film scenes from a
distance and past events; thus short-circuiting space and time.
Chemistry has brought light into dark dwellings by means of
windows and lamps. It supplies food to the hungry, for with
nitrogen extracted from the air and phosphorus from the rocks
it makes two blades of com grow where one grew before.
Chemistry is the joyous science. It contributes most to the
pleasures of life. Every scent and every savor is due to chem-
istry. The tints that delight the eye in flower or fruit, in gem
or painting, the greens and browns of the forest leaves, the reds
and yellows of the brick-built town, the black of print, the dyes of
cloth, the blush on the maiden's cheek — whether it be natural or
artificial — all colors are chemical. We may not gratify our ears
with music more than once a week, but we satisfy our chemical
senses three times a day and chemical colors delight our eyes all
the hours they are open. The students should realize that chem-
istry is not exclusively concerned with the stuff in the bottles be-
hind the prescription counter, but has just as much to say about
the soda-water side to which they are quite naturally more at-
tracted. They should understand that chemistry includes not
merely the process of cooking but the process of digestion. This
explains why we have introduced some unconventional topics and
experiments in the following pages.
Chemistry is a practical science. None touches everyday life
at more points except its sister science, physics, between which
and chemistry no clear boimdary can be drawn. None has more
WHY STUDY CHEMISTRY? vii
avenues for profitable employment. Some studies can be made
a profession only by teaching them. But the qualified chemist
has two strings to his bow. If he does not like teaching, he may
leave the campus for the industrial field. He may find employ-
ment in the big establishments devoted to the reduction of ores,
to steel or other manufactures or to the production of foods, fer-
tilizers, dyes, drugs, soaps, gas, rubber, oil, cements or explosives.
Here he may be engaged in routine analysis of the raw material
and finished products. Or he may be put in control of the fac-
tory processes. Or he may be engaged in research to prevent
waste, to utilize by-products, to invent new compounds or to
contrive new uses. In any case he will be near to the heart of
the industry and if he has inventive genius or managerial ability
he may rise high. In some fields of chemistry women have almost
as much of a chance as men.
On account of its varied opportunities and its expandmg field
chemistry is drawing an increasing number of college graduates.
Many more take the degree of Doctor of Philosophy in chemistry
than in any other branch. The American Chemical Society
with 15,500 members is the largest special scientific body in the
world. When the Government took a census of chemists to see
how many were available for war service, 25,000 names were
registered, and this did not include them all. No other science
is growing so fast through original research. In the files of the
National Research Council more than 2100 living Americans are
recorded as having made contributions to chemistry. Compared
with this, the research men in botany number 1400, in zoology
1255, in geology 750, and in mathematics 600, in psychology 450
and in astronomy 250.
But it is not suflScient that the ranks of research chemists and
technologists be kept filled. Chemistry has an interest and a
;isefulness for everyone. What would be the value of an art
gallery if nobody entered it but artists? And what would become
of the book business if nobody read but authors? Chemistry
viii smith's intermediate chemistry
is of value to those who never intend to become professional chem-
ists because it gives one an imderstanding of household arts and
modem manufactiu*es, of the principles of agriculture and bodily
processes, but more because it opens the eyes of the mind to the
molecular meaning of the marvelous metamorphoses of the visible
world.
• E. E. S.
CONTENT^
CHAPTER PAOB
I. SXTBSTANCES AND PROPERTIES 1
II. Chemical Change and the Methods op Studying It . . . 12
III. Air and Oxygen 26
IV. The Measurement op Gases. The Molecular Hy-
pothesis 44
V. Hydrogen , 50
VI. Water 58
VII. Chemical Units op Weight. Formula 72
VEIL Application op the Molecular Hypothesis in Chem-
istry 82
DC Making op Formula and Equations 98
X. Solution 108
XI. Hydrochloric Acid. Calculations 125
XII. Chlorine. Calculations 189
XIII. Energy and Chemical Change 153
XIV. Sodium and Sodium Hydroxide 164
XV. Acids, Bases, and Salts 170
XVI. Ionization 181
XVII. The Halogen Family 199
XVIII. Valence 211
XIX. Oxidizing Substances 219
XX. Chemical Equilibrium 230
XXI. Sulphur and Hydrogen Sulphide 248
XXII. Oxides and Oxygen Acids op Sulphur 257
XXIII. The Periodic System 273
XXIV. Nitrogen. The Atmosphere 286
XXV. Ammonia 299
XXVI. Nitric Acid 308
XXVII. Phosphorus, Arsenic, Antimony, Bismuth 319
XXVIII. Carbon and the Oxides op Carbon 328
XXIX. The Hydrocarbons and their Deriyattves. Flame . . 343
ix
X CONTENTS
•
CHAPTER PAGE
XXX. Silicon, Boron 359
XXXI. Compounds of Sodium and Potassium 365
XXXII. The Recognition of Substances, I. — A Review of
THE Non-Metallic Elements 374
XXXIII. Calcium and Its Compounds 383
XXXIV. Plant Life. Cellulose, Starch and Sugar 395
XXXV. Plant Life. Osmosis. Fertilizers 405
XXXVI. Plant Products. Fermentation and Fuels 417
XXXVII. Animal Life and Animal Products. Foods 428
XXXVIII. Magnesium and Zinc. Ionic Equilibria 446
XXXIX. Aluminium 466
XL. Synthetic Organic Products 474
XLI. Iron, Nickel, Cobalt 486
XLII. Lead and Tin. 502
XLIII. Copper and Mercury 510
XLIV. Silver, Gold, Platinum 519
XLV. Manganese and Chromium 527
XLVI. The Recognition of Substances, II. — A Review of
THE Metallic Elements 536
XLVII. Radium. Atomic Energy and Atomic Structure 542
Appendix 557
Index 559
LIST OF ILLUSTRATIONS
PAGE
Alexander Smith Frontispiece
Alexander Smith; B. Sc, Edinburgh, 1886; LL. D., 1919; Ph. D.,
Mimich, 1889. Fellow of the Royal Society of Edinburgh;
Member of the National Academy of Sciences; recently Head of the
Department of Chemistry, Colimibia University, New York City.
Joseph Priestley 28
Joseph Priestley (1733-1804) was a nonconformist minister of
Birmingham, England, but devoted his leisure and his means to
chemical research, the results of which were published in his six
volumes on "Experiments and Observations on Different Kinds of
Air.'* He was the first to prepare pure oxygen, wJiich he called
"dephlogisticated air.'* He also made nitrous oxide, nitric oxide,
ammonia, hydrochloric acid, sulphur dioxide, and silicon fluoride.
He foimd that water could be made to absorb carbon dioxide
under pressure, thus laying the foimdation of the "soda-water"
industry. On accoimt of his sympathy with the French Revolu-
tion and his heterodox theological views a mob burned his home
and laboratory, July 14, 1791. Priestley narrowly escaped with
his life and went to America, where he settled at Northumberland,
Pennsylvania. Here a group of chemists gathered on August 1,
1874, to celebrate the centenary of the discovery of oxygen, and
foimded the American Chemical Society, now the largest scientific
association in the world.
Antoine Laurent Lavoisier 29
Antoine Laurent^ Lavoisier (1743-1794) is called "the father of
modem chemistry" because he placed the science on the solid
basis of quantitative experiments and devised the system of naming
compoimds by their composition. He proved that the air con-
tained two gases; one which was inert and another which supported
combustion and united with metals. He called the former gas
"azote" because it would not support life and the latter "oxygen"
because it formed acids. He was guillotined during the Reign of
Terror. The Commission on Weights and Measures petitioned
that he might be allowed to complete his work on the metric system,
but Cbffinhal, vice-president of the revolutionary tribunal, refused,
saying: "The republic has no need for savants.'
li
yt
XU LIST OF ILLUSTRATIONS
PAGE
SvANTB Arrhbnius 182
Arrhenius put forward his theory of ionization while still a student
at Upsala. His own description of its reception is worth repeating:
"I came to my professor, Cleve, whom I admire very much, and I
said: 'I have a new theory of electrical conductivity as a cause of
chemical reactions.' He said: ^That is very interesting,' and
then said * Good-bye.' He explained to ine later, when he had to
pronounce the reason for my receiving the Nobel Prize for that work,
that he knew very well that there are so many different theories
formed, and that they are almost all certain to be wrong, for after
a short time they disappear; and therefore, by using the statistical
manner of forming his ideas, he concluded that my theory also would
not exist very long."
It scarcely needs to be added here that the theory is still very much
alive I
The Electric Current is Passed through Salt Water in These
Nelson Cells to Obtain Caustic Soda and Chlorine
Gas 183
In the U. S. Government Chlorine Plant at Edgewood Arsenal
during the Great War there were eight cell rooms similar to that
shown in the picture, with a total capacity of 100 tons chlorine per
day. Each cell room consists of six circuits — 74 cells per circuit,
or a total of 444 cells per room.
Sir William Ramsay .296
Sir William Ramsay (1852-1916) will always be famous for his
work on the inert gases of the atmosphere and for his discoveries in
the field of radioactivity. He first became interested in chemistry
in a rather imusual way. While a boy at school, he broke his leg
in a football game, and was kept on a couch for several weeks. To
kill time, he read Graham's Chemistry, hoping to find out how to
make fireworks. For the next few years his bedroom was full
of bottles and test-tubes, and often full of strange smells and star-
tling noises.
The First Helium-Fillbd Dirigible After Its Landing at Wash-
ington 297
When Ramsay was presented with the Longstaff Medal for his
work on the inert gases, the President of the London Chemical
Society remarked: ''If I may say a word of disparagement, it is
that these elements are hardly worthy of the position in which
they are placed. If other elements were of the same unsociable
cbairaqt^, Chemistry would not exist."
LIST OF ILLUSTRATIONS XIU
PAQE
This very unsociability of helium, however, renders it the ideal
gas for balloons and avoids all the risk of fire or explosion attendant
upon the use of hydrogen. Helium exists in small quantity in
many natural gases (see p. 345), and is obtained pure by liquefying
out the other components.
Chemical Smoke-Screenb Against Submarines. A Vessel Usmo
Silicon Tetrachloride 360
The sihcon tetrachloride is contained in a cylinder fixed on the deck
at the stem of the ship, and is forced out into a long funnel, in the
form of a fine spray, by the pressure of Uquid carbon dioxide. Am-
monia gas issues from a second cylinder, and rapid mixture is effected
in the presence of excess of moist air by means of an electric fan
placed at one end of the funnel. The current of air produced
drives the smoke out at the other end of the funnel.
Chemical Smoke-Screens Against Sxtbmarines. A Vessel Usmo
Oleum 361
Oleum (fuming sulphuric acid) was also employed to conceal ships
against submarines during the Great War. A finely-divided spray
of oleum was injected into the smokenstacks of the vessel and
carried by the hot gases issuing therefrom into the atmosphere in
the form of vapor. In contact with the cool moist air over the
ocean, this vapor condensed to give tiny droplets of sulphuric
acid, which hung over the surface as a very persistent mist (compare
pp. 263-4), extremely difi&cult to distinguish from a natural fog.
Eoppers Benzol Recovery Plant, River Furnace Co., Cleve-
land, Ohio 424
The recovery of valuable hydrocarbons of the aromatic series
(benzene, toluene, naphthalene, etc.) from coal gas is a feature of
every modem coke and gas plant. The bulk of the peace-time
demand for "benzols" — as the fight oils extracted by washing
the gas are called in industry — is for employment as motor fuels.
From one ton of coal approximately 3 gallons of " benzol " is
recovered, giving from 20 to 30 per cent more power and mileage
than high-grade gasoline, and causing less carbon trouble.
Oil Shale Cliff, Utah 425
Enormous deposits of oil shale (see p. 345) have recently been
opened up in the United States, notably in Colorado, Utah, and
Nevada. Though at first sight it would appear that the cost of
producing oil from shale would prohibit competition with well
petroleum, yet the value of the by-products offsets this disadvantage
XIV LIST OP ILLUSTRATIONS
PAGB
to a large extent. In any case, in view of the apprehended exhaus-
tion of the world's petroleum fields in the near future, these shale
deposits constitute a most important national asset.
Photomicrographs Showing the Structure op Steel Made by
Propessor E. G. M ahin op Purdue Universiit 490
1. Cold-worked steel showing ferrite and sorbite (enlarged 500
times).
2. Steel showing pearlite crystals (enlarged 500 times).
3. Structure characteristic of air-cooled steel (enlarged 50 times).
4. The triangular structure characteristic of cast steel showing
ferrite and pearhte (enlarged 50 times).
Repairing the Broken Stern-Post of the U. S. S. " Northern Pa-
cific ": THE Biggest Marine Weld in the World .491
On the right the fractured stem-post is shown. On the left it is
being mended by means of thermite (see page 468). Two crucibles
each containing 700 pounds of the thermite mixture are seen on the
sides of the vessel. From the bottom of these the melted steel
flowed down to fill the fracture.
Madame Curie in Her Laboratory. 642
Marie Sklodovska, a Polish refugee stranded in Paris, was first
engaged in the physical science department of the Sorbonne to
wash bottles and prepare the furnace. Professor Lippmann, the
pioneer in color photography, promoted her to setting up apparatus,
and put her to do work with Pierre Curie, one of his assistants. In
1895 she became Madame Curie.
In 1903 the results of Madame Curie's work on radium were pre-
sented to the faculty as a thesis for the doctor's degree. The thesis,
unlike that of Arrhenius, was favorably received, and shortly after-
wards the Nobel Prize in Physics was divided between the Curies
and Becquerel, whose previous work on uranium had suggested the
research.
In 1911 Madame Curie was awarded the Nobel Prize in Chemistry.
Fog-Tracks prom Radium 543
The positively -charged helium atoms (alpha-particles), thrown
off from radium, in passing through the air ionize the molecules
with which they collide, and these ionized molecules have the same
power that dust possesses (see p. 293) of affording nuclei on which
moisture may condense. Hence, when a particle of a radium
compound is supported in a flask containing air saturated with
LIST OF ILLUSTRATIONS XV
PAGE
moisture, and the air is suddenly cooled by expansion, the paths
of the particles become lines of fog. With poWd^ul illumination, the
fog-tracks can be photographed and the lengths of the paths can
be measured.
The negatively -charged electrons (beta-particles) give fog-tracks
which are mudi fainter and extremely tangled.
Oz^ViTcrr.lc:. Ccllozo cfPh^rmctc-:,
Oa!!fcrn!a ColJo^:© cf Pharmaey
SMITH'S
INTERMEDIATE CHEMISTRY
CHAPTER I
SUBSTANCES AND PROPERTIES
When exact information in regard to any sort of material is
required, we hand the material to a chemist. To learn something
about the nature of chemistry, let us watch the chemist at work
on a typical problem.
Properties. — Suppose that the material is a piece of cloth,
and we desire to know whether it is all wool, or partly cotton.
The chemist places a piece of the cloth in a test-tube (Fig. 1),
and pours in an amount of lye sufficient to cover
it. Lye is a solution in water of a white soUd,
named by chemists, conmionly, " caustic soda "
and, more formaUy, sodium hydroxide. The con-
tents of the test-tube are then heated over a flame
and are kept at the boiling point for ten minutes.
U the cloth dissolves entirely, leaving a Uquid,
clear like water, then it was composed of nothing
but wool. The chemist draws this conclusion
because wool, although not aflFected by the boiUng
water, has the property of turning into a soluble
substance when caustic soda is heated with it. If, on the other
hand, the piece of cloth becomes thinner, obviously losing a part
of its material, but leaving a part undissolved, then it contained
cotton. This conclusion depends on the fact that cotton has
the property of not being dissolved by caustic soda solution.
1
Fig. 1.
2 smith's INHIBMEDIATE CHEMISTRY
The same conclusion could have been reached in other ways.
For example, some threads could have been taken from the
e<^s of the sample and placed under a microscope. In this
case, they would have been seen to be made up of long, hollow
fibers or tubes (Fig. 2). But the cotton fibers (A) are smooth on
the surface, while the woolen fibers (B) are covered with scales.
By this difference in propertiea, the presence of both kinds, or of
only one of the kinds, could quickly be found out. Still again, the
chemist knows that wool will " take " almost any dye, while
cotton remains uncolored by
the greater number of dyes.
He could, therefore, boil a piece
of the cloth with a solution of a
soluble dye — such as red ink
(eosin) — for a few minutes, and
then wash the sample thor-
oughly in clean water. On
examining the cloth with a
microscope he could then ob-
serve whether any fibers were still white. Here wool has the
property of uniting with the dye, while cotton has not. The last
plan could be used only with white goods, or goods not already
strongly dyed.
The first method is the one which the chemist would probably
employ in practice, because by its means he can ascertain and
report, not only the presence of cotton, but the proportion of
cotton present. To do this, he weighs the dry piece of cloth
before placing it in the test-tube. Then, after the boiling with
caustic soda solution, he washes what is left of the sample veiy
thoroughly in running water, dries it, and weighs it ^ain. The
weight of this " residue " is that of the cotton. The difference
between this and the origmal weight is the weight of the wool.
The chemist is then able to state the percentE^es by weight of
cotton and of wool in the original material.
FiQ. 2
SUBSTANCES AND PROPERTIES 3
Substances.— Upon considering these operations, we can
discover a general plan which the chemist has devised and employs
in his work. Different samples of cloth are of diflFerent colors
and appearance,. even when they contain the same proportions of
cotton and wool. A casual inspection is, therefore, of little
value. But in certain respects all samples of wool are alike,
such as in dissolving in caustic soda, in " taking " certain dyes
and in possessing a scaly surface. Those respects in which aU
samples of wool are alike are called the specific properties of wool.
Similarly all samples of cotton are alike in certain respects, which
are called the spedfi^^ properties of cotton. And these properties
of cotton and of wool are different, many of them very different
indeed. For the pm^pose of stating what he means, the chemist
calls a kind of material, all specimens of which possess a certain
set of specific properties, a substance. Wool is one substance,*
and cotton another substance. Every part of a specimen of a
substance has the same specific properties as any other part. If
any portions can be foimd which have different specific properties,
these are portions of another substance, accidentally or inten-
tionally mixed with the first. The chemist calls the foreign mat-
ter an impurity y and the specimen an impure sample of the sub-
stance of which it is mainly composed.
A substancei then, is a species or kind of matteii and all speci-
mens of it show the same set of specific properties. Any par-
ticular substance is recognized by the specific properties which
it exhibits when exposed to various tests. Specific properties,
conversely, are those quaUties which are characteristic of any
particular substance.
The plan which the chemist uses in his experimental tests,
therefore, is that of ascertaining whether a given material exhibits
* In point of fact, wool contains several substances, but they are all alike
in respect to the three properties mentioned above. They differ slightly in
respect to other properties, and so can be distinguished from one an-
other.
4 smith's INTERBiEDIATE CHEMISTRY
the properties of one, or of more than one substance. He. then
describes it by naming the substances he finds in it.
Another Illustration. — This view-point is peculiar to the
chemist. Each art or science has its own view-point — its own
way of thinking about a given object. By the geologist, a piece
of granite is at once thought of as belonging to the older rocks of
the earth, and the geologist considers when it was formed and how
it was formed (namely, by soUdification of a molten mass).
To the builder, it is a very hard stone, expensive to cut and pol-
ish, but very ornamental, and very durable. How and when
it was formed does not make any diflFerence to the builder.
To the chemist, as a chemist, on the other hand, the expense
of cutting granite, and its ornamental character, are of no interest.
Instead, the chemist notices at once that it is spotted, and, upon
examining it closely, he observes that it appears to be a mixture.
He breaks up a portion, and studies the properties of the fragments.
Some are transparent like glass, are very hard, and in fact are in
all respects Uke qvxirtz or rock crystal (Fig. 3). All specimens
of rock crystal, from whatever source, are alike in properties,
and quartz is, therefore, a distinct
substance or species of matter. Again,
certain of the particles in granite are
dark, and with a penknife can be
easily spUt into transparent leaves,
thinner than paper. These fragments
are in all respects Uke mica (sheets of
which, under the name of " isinglass,'' are used to close the win-
dows of stoves), which is another substance (Fig. 4) well-known to
the chemist. Still others of the fragments are less transparent
than the quartz and less hard. They can be spUt, but with
much greater difficulty than the mica. They are crystals,*
* Crystals are natural forms of a geometrical nature, assumed by solid sub-
stances (p. 94).
A
XT
Fig. 3.
Fig. 4.
SUBSTANCES AND PROPERTIES 5
oblong in shape. These are pieces of a third substance, felspar
(Fig. 5). All the particles in the granite belong to one or other
of these three kinds. The chemist, then, studies the specific
properties, such as the hardnesses and the crystalline forms of
various parts of the specimen and seeks to state or describe the
nature of the specimen in terms of the substances he finds
in it.
A Third Illustration. — When flour is examined by the chem-
ist, it appears to the eye to be all aUke. Under the microscope,
even, all he can learn is that it consists largely of grains, which
have the characteristic appearance of grains of starch (see Fig.
99, p. 399). He places some flour on a square piece of cheese-
cloth and encloses it by tying with a thread (Fig. 6). On knead-
ing the little bag in a vessel of water, the water becomes milky.
When the milky water stands,
the white material settles to the
bottom, the water can be poured
oflf, and the deposit can be dried.
This white substance, when boiled
with water, gives an ahnost clear
hquid which jeUies on cooUng.
This is another property of starch.
A Uttle tincture of iodine (solu-
tion of iodine in alcohol), dropped on a part of the starch,
causes the latter to tmn blue. This is a very characteristic
property of (and therefore test for) starch. When the bag of
flour is kneaded persistently in water which is frequently changed,
the material finally ceases to render the water milky. The starch
has all been washed out. When the bag is now opened, a sticlqr
material is found in it. This is called gluten.
The chemist therefore finds that flour contains starch and
gluten. He learns this by separating these two different sub-
stances.
Fig. 5.
Fig. 6.
6 SMITHES INTERMEDIATE CHEMISTRY
ITie Law of Component- Substances. — Every material
consists of certain substances, each of wliich has a definite set
of specific properties. In terms of these properties the material
can be described. This is the first and simplest law of chemistry,
and at the same time the most fimdamental.
Mixtures and Impurities. — A material containing more
than one substance is called a mixture. The characteristic of a
mixture is that each of the substances present, although mixed
with the others, possesses exactly the same properties as if it were
present alone. No one of the ingredients affects any other
ingredient, or alters any of its properties. Granite and flour
are typical mixtures.
When a specimen is composed mainly of one substance, and
contains only minute amoimts of one or more other substances,
it is frequently spoken of as a specimen of the main substance
containing certain specified substances as impurities. To be
called an impurity, the foreign matter need not be dirty or offen-
sive. Thus, common salt usually contains a Uttle magnesium
chloride, a white crystaUine soUd, as an impurity, and it is this
impurity which becomes damp in wet weather. Again, com-
pounds of lime and magnesium are conmion impurities in drink-
ing water.
Component. — The ingredients of a mixture are called the
components (Latin, piU with), because they are simply placed
together, without change, and can be separated without change.
Other Specific Properties. — Beside the specific properties
which happen to have occurred in these illustrations, there are
others which are constantly foimd useful by the chemist. Thus,
in the case of solids, besides the hardness and crystaUine form, he
gives special attention to the temperature at which the substance
melts (the melting-point), the specific gravity or density (weight
SUBSTANCES AND PROPERTIES 7
of 1 C.C., see Appendix I), the color and the solubility or non-solu-
bility in water. In the case of Uquids, the temperature at which
the liquid boils under atmospheric pressure (the boiling-point) ,
the qiedfic gravity, the mobility, the odor and the color (if any)
are never exactly the same for any two different liquids. We shall
learn more about these properties as occasions for using them
arise.
Attributes and Conditions. — It should be noted that
mass is not a specific property. Different specimens of a given
substance may have different masses, but they all have the same
specific properties. The mass of each, although fixed, is a quality
of that particular specimen only. So is it with the volume and
the dimensions of a specimen. These are attributes of a specimen.
Again the temperature is not a specific property. There is no
particular temperature peculiar to quartz. Even the very same
specimen may be at different temperatures at different times.
Temperature and pressure are variable and are called conditions.
The temperature or the pressiu'e of a specimen may be changed
at will. But the specific properties of a substance under any
given conditions cannot be changed, so long as we have to do
with the same substance.
Law of Chemical Change. — When the wool was boiled
with caustic soda solution, it was in some way acted upon by
the caustic soda. It became soluble, and disappeared into the
Kquid. Wool wiU not dissolve in boiling water. But in the
fonner operation its material acquired at least one new property,
namely that of being soluble. Since we have defined a substance
as a species of matter, with a definite set of specific properties,
we are compelled to decide that when a piece of material changes
its properties, it has, in doing so, become a new substance. This
experiment, then, calls our attention to the second of the funda-
mental laws of chemistry, namely that material forming one or
8 smith's intermediate chemistry
more substances, without ceasing to exist, may be changed
into one or more new and entirely different substances. When
such an alteration occurs, it is called a chemical change or reac-
tion. To learn more about this most remarkable fact, we shall
take up in the next chapter some examples of a simple and long
familiar nature.
Definition of Law. — In science, a law, or generalization or
nile, is a statement describing some general fact or constant
mode of behavior. Its uses are to condense a great many similar
facts into one statement, and thus to make the whole set of facts
more easy to remember.
Constituents. — As we have seen, we speak of the substances
in a mixture as the components. When we wish to refer to the
forms of matter which are chemically united in a compound, we
call them the constituents (Lat., standing together) of the com-
pound substance. Thus, iron and oxygen are the constituents of
rust.
The chemist separates (p. 13) the components of a mixture,
for that is all that is necessary. He liberates the constituents of a
compound, however, because they are bound together in chemical
combination.
The names given to compounds are usually devised so as to
indicate the nature of the constituents. Thus, iron-rust is oxide
of iron (or ferric oxide, from Lat., ferrum, iron). The yellowish
powder obtained when lead is heated in air is lead oxide or oxide
of lead, and the white powder similarly derived from tin is oxide
of tin.
A Condensed Form of Statement. — We may represent a
chemical combination, or indeed any kind of chemical change, in a
condensed form, thus:
Iron + Oxygen -» Oxide of iron (ferric oxide)
SUBSTANCES AND PROPERTIES 9
Each name stands for a substance. Two substances in contact
with one another (mixed), but not united chemically, are con-
nected by the + sign. The arrow shows where the chemical
change comes in, and the direction of the change. We read the
statement thus: Iron and oxygen brought together under suit-
able conditions imdergo chemical change into oxide of iron, called
also ferric oxide. Similarly we may write:
Lead + Oxygen -^ Oxide of lead.
Tin + Oxygen — > Oxide of tin.
Explanation of Rusting. — Experiment shows that the
process of rusting is accompanied by a slow increase in the
weight of the soUd, due to the gradual addition of oxygen to the
metal. Now, this increase in weight ceases of its own accord,
when a certain maximum has been reached. This occurs when
the last particles exhibiting the properties of the metal have dis-
appeared. Thus, lead gains in weight until every 100 parts of the
metal have gained 7.72 parts of oxygen, and tin until every 100
parts have gained 26.9 parts of oxygen. When these increases
have occurred, the metal is foimd to have been all used up, and
prolonged heating and stirring cause no further imion with oxy-
gen and no further change in weight. This fact, that each sub-
stance limits itself of its own accord to combining with a fixed
proportion of the other substance, in forming a given compound,
is one of the most striking facts about chemical combination. In
mixtures, any proportions chosen by the experimenter may be
used. In chemical union, the experimenter has no choice; the
proportions are determined by the substances themselves. Thus,
100 parts of iron when turning into ordinary red rust take up 43
parts of oxygen, no more and no less.
This fact enables us to make our condensed statements more
specific and complete by including in them the proportions by
weight used in the chemical change:
10 smith's intermediate chemistry
Iron (100) + Oxygen (43) -> Ferric oxide (143).
Lead (100) + Oxygen (7.72) -> Oxide of 'ead (107.72).
The following numbers, which represent the same proportions
by weight, are the ones commonly used by chemists:
Iron (111.68) + Oxygen (48) -^Ferric oxide (159.68).
Summary. — Thus far, we have learned that chemistry deals
with substances and their specific properties, and with the changes
which substances undergo. We have discussed and defined a
number of important words expressing fundamental chemical
ideas. Finally, we have touched upon the weights of the ma-
terials used in chemical change, a subject of great iniportance
which will be more fully developed in a later chapter.
Exercises. — 1. Describe the following, by mentioning some of
their specific properties: (a) water, (b) wool, (c) cotton (pp. 1-3).
2. If any of the following are mixtures, mention the facts which
show them to contain more than one substance (p. 3) : (a) muddy
water, (b) an egg, (c) milk.
3. State and illustrate the first two laws of chemistry (pp. 6, 7).
4. Make definitions of " pure *' and " impure " as appUed to a
sample of a substance (p. 6).
5. Give a list of the specific properties mentioned in this chap-
ter.
6. In recognizing a specimen to be quartz, does the chemist
consider (a) the weight, (b) the temperature, (c) the length of the
specimen (p. 7 )? If not, why not?
7. Take one by one the words or phrases printed in black type
and the titles of the sections in this chapter, and endeavor to
recollect what you have read about each. In each case try,
(a) to recall the meaning and to state it in your own words; (6) to
recall the facts associated with, and the reasoning which lead up to
the point in question; (c) to recall examples illustrating the con-
ception and to apply the conception in detail to each example.
SUBSTANCES AND PROPERTIES 11
Whenever memory faik to give a perfectly clear report of the
matter in hand, the text must be read and re-read mitil the essen-
tial point can be repeated from memory.
Use the same method in all futm*e chapters. A useful prac-
tice is to employ a pencil as you read and to underline systematic-
ally all the important facts and statements, and then to go back
and apply to each marked place the process described above.
8. Define the following terms: Specific gravity, tenacity, melt-
ing-point, specific physical property, pure body, vacuum.
9. Is it logical to say " piu'e substance? "
10. Why do we decide that granite is a mixtiure and iron a
smgle substance?
11. What weight of oxygen would be required to convert 25
grams of lead into oxide of lead?
12. Make a list of the technical words we have defined, and
place the definition opposite to each.
CHAPTER II
CHEMICAL CHANGE AND THE METHODS OF STUDYING IT
We must now take up two new examples of chemical change.
They will aid us in introducing one or two additional conceptions
and laws. These are continually used by the chemist, and without
them we cannot begin the systematic study of the science.
Another CcLse of Combination: Iron and Sulphur. —
Since oxygen is an invisible gas, there is a slight difficulty in real-
izing that rusting consists in the union of two substances — this
gas and a metal. The present example is less interesting historic-
ally, but it is simpler because both substances are visible and are
easily handled. The case of iron and sulphur will enable us to
illustrate the same point of view and to
practice the appUcation of the same tech-
nical words. It will also introduce us to
two manipulations — filtration and evapn
oration — which are frequently used by
the chemist.
We begin by observing and contrasting
Fig. 7 the specific properties of the two sub-
stances. Sulphur is a pale-yellow sub-
stance of low specific gravity (sp. gr. 2). It is easily melted (m.-p.
114.5° C). It does not dissolve in water — that is, it does not
mix completely with and disappear in water, as sugar does on
stirring. It does dissolve readily in certain other Uquids, such as
carbon disulphide, however. It crystalUzes in rhombic forms
(Fig. 7). It is not attracted by a magnet. Iron is a lustrous
greyish substance of much higher specific gravity (sp. gr. 7.8).
12
CHEMICAL. CHANGE AND THE METHODS OF STUDYING IT 13
It is very difficult to melt it (m.-p. over 1500° C). It does
not dissolve in any common liquids at ordinary temperatures.
It crystal izes in cubes, and is attracted strongly by a magnet.
Study of the Mixture^ before Combination. — Now, if
some iron filings and pulverized sulphur are stirred together in a
mortar, the result is a mixture. True, the color is not that of
either substance, but with a lens particles of both substances can
be seen. Passing a magnet over the mixture will
easily remove a part of the iron, and with the
help of a lens and a needle the mixture could be
picked apart particle by particle, completely.
We can separate the components of the mixture
more expeditiously, however, by using manipula-
tions based upon other more suitable properties.
Thus, sulphur dissolves in carbon disulphide while
iron does not. If, therefore, a part of the mix-
ture is placed in a dry test-tube along with some
carbon disulphide (Fig. 8), and is shaken, the Uquid dissolves
the sulphur and leaves the iron. To complete the separation, the
iron must he removed from the liquid by
filtration, and the sulphur recovered by
evaporation of the carbon disulphide.
Filtration. — Iron, or any solid, when
it is mixed with a liquid or with a solu-
tion (Uke the solution of sulphur in car-
bon disulphide) is said to be suspended
in the Uquid. If the soUd is one that
settles rapidly, the liquid may be sep-
arated from the solid, in a rough way, by
pouring off as much of the clear, super-
iiatant liquid as possible. This is called decantation.
A complete separation is effected by pouring the mixtiure on to a
Fig. 8
Feo. 9
14 smith's interbiediate chemistry
cone of filter paper supported in a glass funnel (Fig. 9) . The liquid,
together with anything that may be dissolved in it, runs through
the pores of the paper and down the hollow stem of the funnel.
The liquid is then called the filtrate. The particles of the sus-
pended solid are too large to pass through the pores, and so col-
lect on the surface of the filter paper. This operation, like every-
thing the chemist does, takes advantage of differences in the
specific properties of the various materials.
The material remaining on the paper (the residue), when dry,
is wholly attracted by a magnet and shows all the other properties
of iron.
Evaporation. — To recover the sulphur, the solution in carbon
disulphide — the filtrate — is poiu'ed into a porcelain evaporat-
ing dish. (Carbon disulphide is veiy m-
flammable ! Keep flames away) . When
the vessel is set aside, the Uquid grad-
ually passes off in vapor (e-vapor-ates).
„ Sulphiu", however, does not evaporate
at room temperature and remains as a
residue, in the fonn of crystals of rhombic outline in the bottom
of the dish (Fig. 10). Here, again, differences in specific properties
have been utiUzed.
Since the physical properties of two substances are not changed
by mixing, we have thus used the properties of the iron and sul-
phiu" so as to separate them once more. The iron is on the paper;
the sulphur is in the dish.
Combination of Iron and Sulphur.-^ But iron and sulphur
are capable of combining to form a new substance, if we alter
the conditions by raising the temperatiu'e. When we place some
of the original mixture of iron and sulphur into a clean test-tube
and warm it, we soon notice a rather violent development of heat
taking place, the contents begin to glow, and what appears to
CHEMICAL CHANGE AND THE METHODS OF STUDYING IT 15
. be a form of combustion spreads through the mass. The heatmg
employed at the start falls far short of accounting for the much
greater heat produced. When these phenomena have ceased, and
the test-tube has been allowed to cool, we find that it now con-
tains a somewhat porous-looking, black soUd. This material is
brittle; it is not magnetic; it does not dissolve in carbon disul-
phide; and close examination, even under a microscope, does not
reveal the presence of different kinds of matter. This substance
is known to chemists as ferrous sulphide and, as we see, its prop-
erties are entirely different from those of its constituents.
In this connection we must not omit to notice that, as in rust-
ing, a certain fixed proportion wiU be used in forming the com-
pound.
Iron (55.84) + Sulphur (32.06) -> Ferrous sulphide (87.90).
If more iron is put into the original mixture, then some
unused iron will be found in the mass after the action. If too
much sulphur is employed, some may be driven off as vapor
by the heat and any that remains, beyond the correct propor-
tion, can be dissolved out of the ferrous sulphide with car-
bon disulphide. The sulphur which has combined with the iron,
however, is no longer present as sulphur — it has no longer the
properties of sulphur, and therefore cannot be dissolved out.
Another Illustration: Mercuric Oxide. — It has long been
known that air contains an active and an inactive gas. The
Chinese called them yin and yang, respectively. Mayow (1643-
1679) showed that the active gas caused rusting, that it was ab-
sorbed by paint (really by the linseed oil) in " drying," that it
supported combustion of wood and sulphur, and that it is
necessary to life, being absorbed by the blood from the air en-
tering the lungs. It was not until 1774, however, that a pure
specimen of this gas was obtained, by Bayen, and was recog-
nized to be a special kind of gas different from ordinary air.
16 smith's intermediate chemistry
The gas (later to be named oxygen) was made by Bayen from
mercuric oxide, a bright red, rather heavy powder. When
the oxide is heated, (Fig. 11), we find that
a gas is given off. This gas is easily shown
to be different from air, since a glowing splin-
ter of wood is instantly reUghted on being im-
mersed in it. The gas is pure oxygen. During
the heating, we notice also that a metalUc
coating appears on the sides of the tube, in the
form of a sort of mirror. Apparently the vapor
of some metal is coming off with the oxygen and
condensing on the cool parts of the tube. As this shining
substance accumulates it takes the form of globules, which may be
scraped together. It is, in fact, the metal mercury, or quicksilver.
If the heating continues long enough, the whole of the red powder
eventually disappears, and is converted into these two products.
Second Variety of Chemical Change: Decomposition* —
Bayen's experiment introduces to us a second, and very common
kind of chemical action. The first variety was combination or
imion (p. 14). The second is called decomposition. It consists
in starting with a single substance (here mercuric oxide) and
splitting it into two (or more) substances, which differ in proper-
ties from the substance taken and from one another. Here, the
red powder gave mercury, a liquid metal, and oxygen, a colorless
gas.
Simple and Compound Substances* — We have seen that
two (or more) substances, Uke lead and oxygen, can combine to
form a compound substance. Are all substances, then, com-
pounds? We find that some are not. We have never succeeded
in obtaining lead, or oxygen, or iron, or tin, or sulphur by com-
bining any two substances. We can decompose mercuric oxide
by heat, and we have other ways of decomposing compounds Uke
oxide of tin and ferrous sulphide, but we have never succeeded in
CHEMICAL CHANGE AND THE METHODS OP STUDYING IT 17
decomposing the mercury or the oxygen, the iron or the sulphur
themselves. Substances which we are not ablci by chemical
means, to decompose into, or to make by chemical miion from^
other substances are called simple or elementary substances.
The distinction between simple and compoimd substances was
first drawn by Boyle in 1678. Later, and independently, it was
stated very clearly by Lavoisier (1789).
Several substances, regarded in Lavoisier's time as elementary,
have since been shown to be compounds. Thus, quickhme was a
simple substance until Davy, in 1808, prepared the metal calcium
and showed that quicklime was the oxide of this metal. Hence,
we do not say that the substances regarded as simple cannot be
decomposed, but only that they are substances which we " are not
able '' (at present) to decompose.
The phrase " by chemical means " is also important. Although
by chemical methods we are not able to effect any decomposition
of the elements, yet we cannot regard them as absolutely unalter-
able and permanent. The element radium cannot be decomposed
by chemical means, but it undergoes continuous and spontaneous
" disintegration " into the elements heUum and lead (see p. 546).
It has recently been discovered by Rutherford that other elements,
such as nitrogen and aluminum, may also be disrupted to give
hydrogen, under the action of the tremendous forces of the
swiftly-moving particles ejected from radium in the course of
this disintegration. Such phenomena, however, do not affect
our conception of elements as appUed to ordinary chemical reac-
tions.
Elements. — The word element is used in two senses. It is
applied to the simple substance. Thus we speak of " the element
iron," meaning the metal iron. It is appUed also to the iron-
matter contained in ferrous sulphide or in ferric oxide. The
reader should note that it is correct usage to speak of the element
tttm and the element sulphiu- in ferrous sulphide, but a chemist
18 smith's intermediate chemistry
would never say that this compound contained the simple sub'
stances iron and sulphur. If he did, we should understand him to
mean that it was a mixture, and we should expect parts of the
material to be magnetic like iron, and other parts to be yellow and
soluble in carbon disulphide, which is not the case. In the same
way the name of an element (such as iron) is appUed both to
the material in combination and to the free substance. Thus
" u-on " may mean free, uncombined, metaUic iron, or iron-matter
m some compound. The sense in which the word is employed
must be inferred from the context or circumstances. When a
chemist speaks, as he sometimes does, colloquially, of " iron " in a
drinking water, for example, we know at once that he refers to
iron in the form of some compound, for metallic iron does not
dissolve in water.
The word element, then, means one of the simple forms of
matter, either free or in combination.
In formally describing a body or specimen, the chemist always
avoids the ambiguity just referred to by naming the components,
i.e,j the substance or substances it contains. He assumes that
the nature and constituents of these substances will be known to
anyone hearing or reading the description. If he says the body
contains zinc and sulphur, it is understood that the body is a
mixture of these simple substances. If it contained these ele-
ments in combination, the chemist would report that it was sul-
phide of zinc.
The Common Elem^ents. — Thousands of different com-
pound substances are known but, when they are decomposed, it is
found that the number of different elements contained in them is
not great. Dozens of substances contain iron, hundreds contain
sulphur, thousands contain oxygen. In fact, by combining a
limited number, two, three, or four, of simple substances together,
in varying proportions by weight, an almost unlimited number of
different compound substances could be produced.
CHEMICAL CHANGE AND THE METHODS OF STUDYING IT 19
A list of the elements appears on the inside of the cover, at the
end of this book, and contains about eighty names. Of these, a large
number are rare, and seldom encomitered. More than 99 per cent *
of terrestrial material is made up of eighteen or twenty elements and
their compounds. Only about twenty elements occur in nature
in their simple, uncombined condition. Three-fourths of the
whole number are found in combination exclusively, and must
be Kberated by some chemical action.
Taking the atmosphere, all terrestrial waters, and the earth's
crust, so far as it has been examined, F. W. Clarke has estimated
the plentifulness of the various elements. The first twelve, with
the quantity of each contained in one hundred parts of terrestrial
matter, and constituting together 99 per cent, are as follows:
Oxygen 49.85 Calcimn 3.18 Hydrogen 0.97
Silicon 26.03 Sodium 2.33 Titanium 0.41
Aluminium 7.28 Potassium 2.33 Chlorine 0.20
Iron 4.12 Magnesium 2.11 Carbon .0.19
The significance of these figures is more clearly shown in Fig.
12. It will be seen that oxygen accounts for nearly one-half of the
whole mass. SiUcon, the oxide of which when piu'e is quartz
and in less pure form constitutes ordinary sand, makes up half of
the remainder. Valuable and useful elements, like gold, silver,
sulphur, and mercury, are among the less plentiful which, all
tiJcen together, furnish the remaining one per cent.
Law of Definite Proportions. — In the decomposition of
mercuric oxide (p. 16)* we find that, for every 100 parts of
* References to previous pages are used in order to save needless repetition
in writing. The beginner requires endless repetition in his reading, however,
and must /arm the habit of examining, in conjunction with the current text, the
parts referred to. The passages cited are, by the reference, made part of the
current text, which will usually not be clear without them. The same remark
applies to topics referred to by name. Such topics must be sought in the
index.
All terms, and especially those borrowed from physics, if not perfectly
familiar, must be looked up in a work on physics or in a dictionary.
20
smith's intekmediate chemistry
mercury liberated, almost exactly 8 parts of oxygen by weight
are set free. Using the nimibers commonly employed in chemis-
' try, which represent the same proportion by weight :
Mercmic oxide (216.6) -^ Mercury (200.6) +
Oxygen (16).
We find also that mercury and oxygen can be
made to combine to form mercuric oxide, and
the proportions by weight required are the same.
Moreover, every sample of mercuric oxide, whether
made by combination, or in any of the other pos-
sible ways, always contains this proportion of the
two elements. We have already seen that the
oxides of lead and tin contam fixed proportions
(p. 9) of the metal and oxygen and that ferrous
sulphide has a constant composition by weight.
The same principle is found to apply to all
chemical compounds, and is stated in the law of
definite or constant proportions: In every sample
of any compound substance, formed or decom-
posed, the proportion by weight of the constituent
elements is always the same.
Certain elements have recently been shown to
exist in two or more forms (isotopes), which would,
if separable, give compounds possessing the same
specific properties yet differing in composition. Detailed discussion
of this point must be deferred to a later chapter (pp. 550-552).
^
1
•
6
^
^
d
•u
0
o
?
u
<0
>>
*»
X
>l
^^^
O
<0
1
i
—
Hyd.
T,T.
Othbh
Elts,
Fig. 12
Conservation of Mass. — The most painstaking chemical
work seems to show that, if all the substances concerned in a
chemical change are weighed before and after the change, there is
no evidence of any alteration in the qicantity of matter. The two
weights, representing the smns of the constituents and of the
products, respectively, are, indeed, never absolutely identical, but
chbhicaij chanqe and the methods of studying it 21
the more careful the work and the more delicate the instrument
used in weighir^, the more nearly do the values approach identity.
We are able to state, therefore, that the mass of a system is not
affected by any chemical change wifliln the system.
Fro. 13 Fig. 14
This statement simply means that the great law of the conserva-
tion of mass holds true in chemistry as it does in physics. Chemi-
c^ changes, thoroughgoing as they are in respect to aH other
quahtiea, do not affect the mass; an element carries with it its
weight, entirely unchanged, through the most complicated chemi-
cal transformations.
Superficial observation, as of a growing tree, might seem to give
evidence of the very opposite of conservation of matter. But
here the carbon dioxide gas in the air, the most important source
of nourishment for plants, is overlooked. Similarly, the gradual
disappearance of a candle by combustion seems to illustrate the
destruction of matter. But if we catch the gases which rise
through the flame (Fig. 13), we find that the gases weigh even
more than the part of the candle which has been sacrificed in
making them. When we take account of the weight of the oxy-
gen obtained from the air which sustains the combustion, we find
that there is really neither loss nor gain in weight. If we cany
22 smith's intermediate chemistry
out chemical changes in closed vessels (Fig. 14), which permit
neither escape nor access of material, we find that the weight does
not alter.
Physics in Chemistry. — It will be seen that one cannot
accomplish anything in chemistry without acquiring and using
some knowledge of physics. We measure quantities by means of
the physical attributes, weight and volume. We produce chemi-
cal change by arranging the physical conditions, for example, by
mixing, heating, or using an electric current. Physical means are
the only means we possess for producing, stopping, or modifying
chemical changes. Again, we ascertain whether a chemical
change has taken place or not by observing the physical properties
of the materials before and after the experiment. Thus, we noted
that the red, powdery oxide of mercury, when heated, gave a
liquid metal and a gas. AU the phenomena of chemistry are
physical. A phenomenon is literally something that is seen or,
more generally, something that aflfects any of the senses. Ob-
serving physical phenomena is, therefore, our sole means of study-
ing chemical changes. Chemical work is, in fact, entirely de-
pendent upon the skilful use of physical agencies, and upon the
close observation of physical phenomena for its success.
It is only the inference, foUowing the experiment and the obser-
vation, that is strictly chemical. If one substance gives two
different substances, or if two substances give one different sub-
stance, for example, we infer that a chemical change has occurred.
We then try to recognize the substances by their properties and
name them.
Changes like that of ice into water, or of water into steam, and
vice versa, are not regarded as chemical changes. These are
called changes of state (see p. 64).
Law: Explanation: Scientific Meth€fd. — There is a widely
spread impression that a sdence, like chemistry is a part of
CHEMICAL CHANGE AND THE METHODS OF STUDYING IT 23
the natural order of the universe. It is thought that we are try-
ing to find the boundaries of chemistry, as they have been pre-
determined by nature, and to discover the facts, relations of
facts, and laws which nature has provided as a means of classify-
ing the content of the science. Now, the situation is precisely
the reverse of this. Nature provides only the materials and the
phenomena, and man is attempting to classify them. He divides
the whole into groups, such as physics, chemistry, botany, etc.
Then he classifies the facts within each group, in order that he
may more easily remember them and perceive their relations.
He often finds that, when new facts are discovered, parts of the
classification have to be changed.
In the preceding pages, we have discussed some of the ways
that have been invented for classifying the materials and facts
assigned to chemistry. Thus, we pick out a number of facts of a
Uke nature and try to make a single statement which will cover all
these facts. For example, we find about one hundred thousand
different substances and, in the case of each svbstance, every speci-
men that we have examined contains the same proportions of
the constituent dements. So we formulate the law of constant
proportions.
A law or generalization in chemistry is a brief statement
describing some general fact or constant mode of behavior. We
must remember, however, that laws are only true so long as no
facts in conflict with them are known. There are no laws in
nature. Natiu'e presents materials and phenomena as she pleases.
The laws are parts of science, which is made by mxin, and is a de-
scription of natiu-al facts as man knows them.
One section (p. 9) was entitled: ''Explanation of rusting."
If that paragraph be now re-read, it will be found that, in the
ordinary (as distinct from the scientific) sense of the word, no
explanation was given! When we ask a man to " explain '' some
feature in his conduct, we recognize that he might have chosen to
act otherwise, and we wish to know why he acted precisely as he
24 smiths' intermediate chemistry
did. Nature, however, has no free will, and cannot tell why she
presents certain phenomena, and not others.
On examining the explanation, we find that it simply shows
that when iron rusts it combines with oxygen from the air. This
is an additional fact. It shows how iron rusts, namely, by taking
up oxygen, but not why it is able to imite with oxygen. We
simply do not know why iron can combine with oxygen gas and
platinum cannot.
Explanations in chemistry are of three kinds. (1) We usually
try to show that the phenomenon is not an isolated one. Thus,
we show that other metals rust. This reconciles us to some ex-
tent to the fact that iron rusts, and we feel some mental satisfac-
tion. This is the method of showing that the fact to be explained
is a member of a large class of similar facts. (2) Next, we try
to get more information about the fact to be explained. Thus,
when, to the acquaintance with the outward manifestations of
rusting, we add the further information that there is an increase
in weight, and that this is due to union of oxygen from the air
with the iron, we feel increased satisfaction, and say that the
fact has been " explained.'' (3) If we are still dissatisfied, and
can discover no further useful facts, we imagine a state of affairs
which, if true, would classify the fact or add to what we know
about it. This step we call explaining by means of an hjrpothesis.
We then devote our attention to trying to verify the hypothesis.
The formulation of laws and the making of attempts to explain
facts are part of what is called the scientific method. The purpose
of this method is to convert the subject matter into a science, that
is, into an organized body of knowledge.
Summary. — In this chapter we have learned: (1) that,
while there are many substances, there is a limited number of
entirely different kinds of matter (elements) ; (2) that, in addition
to definite specific properties, each substance has a constant
composition by weight. We have also learned that physical
CHEMICAL CHANGE AND THE METHODS OF STUDYING IT 26
properties are utilized in manipulations, like filtration and evapo-
ration, as well as for identifjdng substances, and that physical
attributes are used for measuring quantities in chemistry and
physical conditions for guiding chemical change. Finally, we
have seen that a science is not a natural, but a manufactured
product, and that the science of chemistry is still in the making.
Exercises.* — 1. What physical properties are used (a) in
filtration, (fc) in evaporation, (c) in the separation and identifica-
tion of the products from heating mercuric oxide (p. 15)?
2. Describe: (a) a red-hot rod of iron, 10 cm. long by 1 cm.
diameter, weighing 58.5 g.; (6) a solution of 5 g. of sulphur in
20 c.c. (26 g.) of carbon disulphide at 18° C. In doing so, divide
the description into attributes, conditions, and properties.
3. CJonsider the following materials and state whether, so far
as you can now judge, each is a single substance or a mixture: (a)
a candle, (&) a cake of soap, (c) an egg.
4. What are the two most direct ways of showing a substance
to be a compound? Illustrate each.
5. If we say that quickhme contains calcium, do we mean the
element or the simple substance calcium?
6. What explanation was given, (a) of the disappearance of
mercuric oxide when heated, (&) of the absence of iron and sulphur,
as substances, from ferrous sulphide? Which of the three kinds
of explanation was used in each case?
7. What weight of oxygen will be required to combine with 15
grams of lead (p. 10)?
8. If 5 grams of lead and 4 grams of oxygen were heated
together, which of the two would remain in part unused? How
much of this one would remain (p. 10)?
* The exercises should in all cases be studied with minute care. They not
only serve as tests to show that the chapter has been understood, but very
frequently (as in No. 4) also call attention to ideas which might not be ac-
quired from the text alone, or (as in Nos. 1, 2, 5) assist in elucidating ideas
giyen in the text which, without the exercises; might not be fully grasped.
CHAPTER III
AIR AND OXYGEN
We have seen that metals absorb a gas, called oxygen, from
the air, and turn into a rust or oxide. Let us now consider what
happens to the air during this process.
The Nature of Air. — Can a metal, like iron, when rusting,
absorb the whole of a sample of air, or does it select a part of the
air only? If we sprinkle some powdered iron in a test-tube,
having first moistened the interior to cause
the powder to adhere to the inside surface,
and then set the tube, mouth downwards,
m a dish of water (Fig. 15), we obtain be-
fore long an answer to this question. As
the iron slowly removes the oxygen, the pres-
sure of the atmosphere outside pushes the
water up the tube. But, after ascending
only about one-fifth of the total height of the
tube, the water comes to rest. Inspection shows reddened parti-
cles where rusting has taken place, but much of the iron is still
dark grey, and is as Uttle able to rust in the remaining gas as in a
vacuum. Four-fifths of the air, then, is composed of gases which
do not combine with iron, and only one-fifth is oxygen. The
four-fifths is in fact almost all (99 per cent) nitrogen, a substance
which combines with very few materials, while the balance (1 per
cent) is made up of argon and other gases which do not enter
into combination with any known substances. Oxygen, on the
contrary, combines with almost all simple substances, although
in many cases, as in those of lead and tin, heating is required to
hasten the process.
26
Fig. 15
AIR AND OXYGEN 27
Activity and Stability. — A substance which enters into
combination easily, is called active, so that oxygen is spoken of as
an active element, nitrogen as a relatively inactive or indifferent
element. An active element, since it combines greedily, holds
tenaciously to that with which it has combined. An active
element means, therefore, also, one which is in general difficult to
liberate from combination. Its compounds are in general rela-
tively stable, or inactive. The compounds of more indifferent
elements, on the other hand, are in general relatively easy
to decompose. In other words, they are unstable or active.
Law of the Influence of Heating. — Even oxygen, active
as it is, does not combine visibly with tin, when both are cold.
Lead rusts very slowly at the ordinary temperatiu'e. Iron rusts
very much faster when heated than when cold. In every chemical
change we find that raising the temperature hastens the process
very considerably. Other things being equal, it causes a greater
quantity of material to undergo the change in a given time. This
is the third law of chemistry.
As exceptions to this law, radioactive disintegrations (see p. 545)
must be noted. The rate at which reactions of this type proceed
is in every case absolutely independent of the temperature. The
speed of a photochemical reaction (a chemical change induced by
the action of light; see p. 498) is also, in general, very Uttle
affected by raising the temperature. For normal chemical
changes, however, the influence of temperature is very marked,
a rise of only ten degrees (approximately) causing most reactions
to proceed at twice their original rate.
Oxygen. — We cannot do better than begin the more systematic
study of chemistry with oxygen, for it is a most interesting as well
as useful substance. It is the active component of the air. We
depend upon it for life, since in its absence we suffocate, for heat,
28
smith's intermediate chemistry
since wood, coal, and gas will not bum without it, and even for
light where oil, gas, or a candle is used.
We wish to know with which substances it can combine, as well
as the substances on which it has no action. This information will
show us how to work, in future, without interference from the
oxygen in the air and whether oxygen has probably played a part
in some experiment or not.
Let us take up, then, (1) the history of the element, (2) what
materials contain oxygen (occurrence), (3) how we can obtain it
in a pure state (preparation), (4) what its specific physical proper-
ties as a substance are, and (5) what it does, and what it cannot do
in nature and in the laboratory (chemical properties). The
classification of the facts about this, and other substances under
five heads, is somewhat mechanical, but has the advantage of
enabling the reader quickly to find any required information.
History of Oxygen. — Leonardo da Vinci (1452-1519) seems
to be the first European to mention the presence of two gases in
the air. Mayow (1669) measured the proportion of oxygen in the
air and discussed fully its uses in combustion, rusting, vinegar-
making, and respiration, but did not make a pure sample. Hales
(1731) made it from saltpeter, and measured the amount obtain-
able, but did not see any connection between it and the air!
Bayen (Apr., 1774) was the first to make it by heating mercuric
oxide. Priestley (Aug. 1, 1774) made it
by heating the same substance and quite
purposelessly, as he admits, thrust a lighted
candle into it and was deUghted with the
extreme briUiance of the flame. Scheele, a
I Swedish apothecary, had made it in 1771-2
from no less than seven different sub-
stances and understood clearly that atmos-
pheric oxygen combined with metals, phosphorus, hydrogen,
linseed oil and many other substances. Finally, Lavoisier
Fig. 16
AIR AND OXYGEN
29
(1777) heated mercury in a retort (Fig. 16), the neck of
which projected into a jar standing in a larger dish of mercury.
The air, thus enclosed within the jar and the retort, during
twelve days lost one-fifth of its volume. Simultaneously, red
particles of mercuric oxide accumulated on the surface of the
mercury in the retort. The residual gas, nitrogen, no longer
supported life or combustion* ' The oxide, on being heated more
strongly, by itself, gave off a gas whose volume exactly corre-
sponded with the shrinkage undergone by the enclosed air, and
this gas posses3ed in an exaggerated degree the properties which
the air had lost. The proof that oxygen was a component of
the atmosphere was therefore complete. Later, Lavoisier, in
the mistaken bdief that the new element was an essential con-
stituent of aQ sour substances, named it oxygen (Greek, add-
producer).
Occurrence. — As we have seen, nearly 50 per cent of terres-
trial matter is oxygen. Water contains about 89 per cent, the
human body over 60 per cent, and common materials
like sandstone, limestone, brick, and mortar more than ly
50 per cent of this element. One-fifth by volume [p — »
(nearly one-fourth by weight) of the air is free oxy- ^"^
gen.
Preparation of Oxygen. — 1. Th® oxygen of com-
merce is now made chiefly from liquefied air. The
Kquid oxygen boils at — 182.5°, but the nitrogen boils
at an even lower temperatiu'e ( —194°). Since the
Kquid air has a temperature of about —190°, some-
what above that of boiling nitrogen, the latter
evaporates much more freely than does the oxygen.
After a time, when the remaining Uquid is almost pure
oxygen (96 per cent), the gas coming off is compressed by pumps
into the steel cylinders (Fig. 17) in which it is sold. In medicine,
Fig. 17
30
smith's intekmediate chemistry
patients suffering from pneumonia or suffocation obtain some
relief by inhaling it in this form. It is also used in feeding
flames, instead of air, when intense heat is required (see acetylene
torch, p. 352j and calcium light, p. 385).
2. Unfortunately, it is difficult to Uberate oxygen from natural
substances. Saltpeter (potassium nitrate), for example, which
is found in many soils and can be dissolved out with water, gives
off oxygen only when raised to a bright red heat by the Bunsen
flame or blast lamp. But, even at this temperature, it gives
up only one-third of the oxygen it contains.
3. In practice, we are compelled to use manufactured sub-
stances. Amongst the artificial substances are mercuric oxide,
expensive but historically interesting (p. 15), potassium chlorate,
perhaps the most convenient for laboratory use, and sodium per-
oxide. Potassium chlorate is a white crystalline substance
Fig. 18
used, on accoimt of the oxygen it contains, in large quantities in
the manufacture of matches and fireworks. When heated in a
tube similar to that in Fig. 11, it first melts (357®) and then, on
being more strongly heated, it effervesces and gives off a very
large volume of oxygen. Examination shows that the whole
of the oxygen it contains (39 per cent) can be driven out. The
white material which remains after the heating is identical with the
mineral sylvite. To the chemist it is known as potassium chloride.
AIE AND OXYGEN
31
The change, t(^ether with the weights of the materials, is as
follows:
Potassium chlorate (122.56) -+ Potassium chloride (74.56) + Cbtygen (48)
Potassium (39.1)
Chlorine (35.46)
Oxygen (48)
Potassium (39.1)
Chlorine (35.46)
A peculiarity of this action is that admixture of manganese
dioxide (the mineral pyrolusite) increases very markedly the
speed with which the decomposition of the potassimn chlorate
takes place. Hence powdered manganese dioxide is generally
mixed with the chlorate in laboratory experiments (Fig. 18), and
in its presence a sufficient stream of oxygen
is obtained at a relatively low temperature
(below 200^).
4. Oxygen can be obtained conveniently
from sodium peroxide and water by means
of generators (Fig. 19) similar to the acety-
lene generators used on automobiles. When
the metal sodium is burned in air, sodimn
peroxide is obtained as a powder. This
powder, after being melted, soUdifies in
compact, solid form, and is sold as oxone.
The oxone is bought in a small, sealed tin
can, the ends of which are perforated in
several places just before use. When the
valve (jB) is opened, so that the oxygen
escapes, the water, which fills the generator
almost to the top, enters the can (C) by the holes in the bottom
and interacts with the oxone. When the valve is shut, the gas
continues to be generated imtil it has driven the water down
again below the level of the bottom of the can.
Sodium peroxide (78) + Water (18) -> Sodium hydroxide (80; -r Oxygen (16)
Sodium
Oxygen
1
D
Fig. 19
Sodium (46)
(32)
Hydrogen (2.016) Sodium (46)
Oxygen (16) Oxygen (32)
Hydrogen (2.016)
This method is convenient because it works at room temperature
32
smith's intermediate chemistry
and can be started and stopped at will. The sodium hydroxide
produced is very soluble in water and remains dissolved.
Catalytic or Contact Action. — The influence of manganese
dioxide in causing the potassium chlorate to decompose more
easily (p. 31) well deserves notice. The effect is very striking if
some pure potassium chlorate is melted carefully, to avoid super-
heating, in a wide-mouth flask (Fig. 20). The flask is provided
with a wide exit tube, from which a rubber tube may lead to a
bottle inverted in a trough filled with water as in Fig. 18. A
little manganese dioxide is contained in the upper, closed tube.
No effervescence of the chlorate can be seen at its melting-point
(357°) — only a Uttle air, expanded by the heating, issues from
the tube. When, however, the closed tube containing the man-
ganese dioxide is rotated into a vertical position (see dotted lines),
and the black powder falls into the chlorate, the oxygen comes
off in torrents, in consequence of the enormous acceleration of the
decomposition. As a pre-
caution against injury from
an explosion of the flask, it is
advisable to wrap the latter in
a towel before turning the tube.
It must also be noted that
the manganese dioxide is not
itself permanently altered. If
the material left after the
action is shaken with water,
the potassium chloride dis-
solves, while the dioxide does
not. Filtration (p. 13) thus
enables us to recover the latter, and to ascertain that it has
been changed neither in quantity nor in properties.
The only effect of the dioxide is to hasten the decomposition of
the chlorate, which would otherwise be too slow at 200® (p. 31),
Fig. 20
AIR AND OXYGEN 33
or even at 357° (its m.-p.) to be of any practical value. Sub-
stances which hasten a chemical action without themselves under-
going any pennanent change are called contact agents, catalytic
agents, or catalysts. The process is called contact action or cataly-
sis (Greek, decomposition, not a very fortunate choice of words).
Such substances are frequently used in chemistry. The addition
of a suitable catalyst is one of the conditions (p. 7) for carrying
out actions in which a contact agent is necessary. Many sub-
stances of this class are secreted by animals and plants and play
an important part in digestion, fermentation, and other physio-
logical changes. Their presence often enables very complex
chemical actions to proceed rapidly at rather low temperatures.
The oxone, mentioned above, always contains a catalyst,
a trace of cuprous oxide, which hastens the action on water.
Specific Properties of Two Kinds^ Physical and ChemicaL
— We have learned that every substance has its own set of specific
properties. In describing a substance, it is convenient to divide
the properties into two classes. The Ust of substances with
which the given substance can enter into chemical combination,
for example, we place under specific chemical properties. Rela-
tions of the substance to any of the varieties of chemical change
belong to this class.
On the other hand, we do not consider melting or boiUng to be
chemical changes, so we place the temperatures at which the sub-
stance melts (m.-p.) and boils (b.-p.), its color, etc. (for list, see p.
6), under specific physical properties.
Properties of either class may be used for recognizing a substance.
Specific Physical Properties of Oxygen. — Oxygen resembles
air in having neither color, taste, nor odor. The density of a sub-
stance is, strictly speaking, the weight of 1 cubic centimeter (1 c.c).
In the case of a gas, we frequently prefer to give the weight of
1000 CO. (1 liter), at 0"* and 760 mm. (1 atmosphere) barometric
34 smith's intermediate chemistry
pressure. For oxygen this weight is 1.42900 grams (Morley).
The corresponding weight for air is 1.293, so that oxygen is slightly
heavier, bulk for bulk, than air' (in the ratio 1.105 : 1). Oxygen
can be liquefied by compression, provided its temperature is
first reduced below —118°, which is its critical temperature.*
The gas is only sUghtly soluble in water j the solubiUty at 20° being
3 volimies of gas in 100 volmnes of water (at 0°, 4.9 : 100).
The solubiUty of oxygen in water, although slight, is in some
respects its most important physical property. Fish obtain oxy-
gen for their blood from that dissolved in the water. With air-
breathing animals (Uke man), the oxygen could not be so readily
taken into the system, if it did not first dissolve in the moisture
contained in the walls of the air sacs of the lungs, and then pass
inwards in a dissolved state to the blood.
Liquid oxygen, first prepared by Wroblevski, has a pale-blue
color. At one atmosphere pressure, that is, in an open vessel, it
boils at — 182.5°. Its density (weight of 1 c.c.) is 1.13, so that it
is sUghtly denser than water. By cooUng with a jet of Uquid
hydrogen, Dewar froze the Uquid to a snow-Uke, pale-blue soUd.
A tube of liquid oxygen is noticeably attracted by a magnet.
Six Specific Physical Properties of Each Gas. — Although
every substance has many physical properties, we shaU mention
only those which are used in chemical work, with occasionaUy the
addition of any pecuUar or unexpected quality. It wiU aid the
memory to recall the physical properties of a gas, if we note that,
as a rule, only six such properties are mentioned: (1) color, (2)
taste, (3) odor, (4) density, (5) liquefiabiUty, defined by the
critical temperature, (6) solubiUty, usuaUy in water only.
* Each gas has an individual critical temperature (p. 91) above which no
pressure, however great, will produce liquefaction. The farther the tempera-
ture of a specimen of the gas is below the critical point, the less will be the
pressure required to liquefy it.
AIR AND OXYGEN 35
Specific Chemical Properties of Oxygen. — The chemical
properties of pure oxygen are like those of atmospheric air, only-
more pronounced.
Reactions vrith Non-metallic Elements. Sulphur, when raised in
advance to the temperature necessary to start the action, unites
vigorously with oxygen, giving out much heat and producing a
familiar gas having a pungent odor (sulphur dioxide). This
odor is frequently spoken of as the " smell of sulphur," but in
reality sulphur itself has no odor, and neither has oxygen. The
odor is a property of the compound of the two. The mode of
experimentation can be changed and the oxygen led into sulphiu*
vapor through a tube. The oxygen then appears to bum with
a bright flame, giving the same product as before.
Phosphorus, when set on fire, blazes in oxygen very vigorously,
forming a white, powdery, soUd oxide — phosphorus pentoxide.
Burning carbon, in the form of charcoal or hard coal, glows bril-
liantly and is soon burnt up. It leaves an invisible, odorless gas
— carbon dioxide. At high temperatures, oxygen combines
readily with one or two other non-metals {e.g., siUcon, boron, and
arsenic), and to a small extent (1 per cent at 1900®) with nitrogen.
It will not combine directly with chlorine, bromine, or iodine,
although oxides of the first and last can be prepared by using
other varieties of chemical change. With the six members of the
helium family of which no compounds are known, and with
fluorine, oxygen forms no compounds.
Sulphur (32.06) +Oxygen (32)->Sulphur dioxide (64.06).
Phosphorus (62.08)+ Oxygen (80)-»Phosphoruspentoxide(142.08).
Carbon (12)+0xygen (32)->Carbon dioxide (44).
Reactions with Metallic Elements. Iron, as we have seen, rusts
exceedingly slowly in air and, even when red-hot, gives hammer-
scale, the black soUd which is broken o£f on the anvil, rather
deliberately. In pure oxygen, a bundle of picture-wire, if once
ignited, will bum with surprising briUiancy, throwing off sparkUng
36 smith's intermediate chemistry
globules of the oxide, melted by the heat. This oxide is a blact
brittle substance, identical with hammer-scale, and different from
rust (ferric oxide). It contains, in fact, a smaller proportion of
oxygen than does the latter, and is called magnetic oxide of iron.
Iron (167.52) + Oxygen (64) -^ Magnetic oxide of iron (231.52).
All the familiar metals, excepting gold, silver, and platinum,
when heated, combine with oxygen, some more vigorously, others
less vigorously than does iron. Oxides of the three metals just
named can also be made, but only by varieties of chemical change
other than direct combination.
Com/pound substances, if they are composed largely or entirely
of elements which combine with oxygen, are able themselves to
interact with oxygen. Usually, they produce a mixture of the
same oxides which each element, separately, would give. Hence,
wood, which is composed of carbon and hydrogen with some
oxygen, when burnt in oxygen, produces carbon dioxide and water
(oxide of hydrogen) in the form of vapor. Again, carbon disul-
phide bums readily, giving carbon dioxide and sulphur dioxide,
just as do carbon and sulphur, separately. Ferrous sulphide
gives, similarly, sulphur dioxide and magnetic oxide of iron.
Tests. A Test for Oxygen. — A test is a property which,
because it is easily recognized (a strong color, for example), or
because it is especially distinctive, is commonly employed in
recognizing a substance.
Oxygen, when pure, is recognized by the fact that a spUnter
of wood, glowing at one end, bursts into flame when introduced
into the g«s. Only one other gas (see nitrous oxide) behaves
similarly.
The Measurement of Combining Proportions. — In a
number of condensed statements we have given the proportions
by weight of the materials combining. It is now desirable that
AIR AND OXYGEN Zf
we should know how the necessary measurements are made. The
most exact measurement of the proportions in which the elements
combine to form compoimds involves manipulations too elaborate
to be gone into here. One or two brief statements, diagrammatic
rather than accm*ate, will
show the principles, how- °^' ^' -■r..i..M,.r '^11='
ever.
If we take a weighed
quantity of iron in a test-
tube and heat it with more
than enough sulphur (an
excess of sulphur), we get
free sulphiur along with p^^ 2i
the ferrous sulphide (p.
15), and no free iron survives. We may remove the free sulphur
by washing the soUd with carbon disulphide. The difference be-
tween the weights of the ferrous sulphide and the iron gives the
amoimt of sulphm* combined with the known quantity of the
latter.
As an example of the study of the combination of a metal with
oxygen, we may weigh a small amoimt of copper in the form of
powder in a porcelain boat and pass oxygen over the heated metal
until it is completely converted into black cupric oxide (Fig.
21). The original weight of the copper, and the increase in
weight, representmg oxygen, give us then the data for determin-
ing the composition of this oxide. The data furnished by one
rough lecture-experiment, for example, were as follows:
Weight of boat + copper 4.278 g.
Weight of boat empty 3.428 g.
Difference = weight of copper 0.860 g.
Weight after addition of oxygen 4. 488 g.
Weight without oxygen 4.278 g.
Difference = weight of oxygen 0. 210 g.
38 smith's intermediate chemistry
The proportion of copper to oxygen, so far as this one measure-
ment goes, is therefore 85 : 21.
The results of quantitative experiments are often recorded in
the form of parts in one hundred. To find the percentage of each
constituent, we observe that the proportion of copper is 85 : 85 +
21, or -^j^ of the whole. That of the oxygen is -^ of the whole.
Thus the percentages are:
Copper, 106 : 85 :: 100 : a:. x'= 80.2.
Oxygen, 106 : 21 :: 100 : a;', x' = 19.8.
Naturally, the mean of the results of a number of more carefully
managed experiments will be nearer the true proportion. The
percentages at present accepted as most accurate are 79.9 and 20.1.
In the case of mercuric oxide, we may decompose a known
weight of the oxide (p. 16), collect the mercury and weigh it, and
ascertain the oxygen by difference.
The names of the constituent elements in a compound, together
with the proportion by weight in which they are present, are called
the composition of the substance. Thus, the composition of
cupric oxide is copper : oxygen :: 79.9 : 20.1. This is the per-
centage composition, but other numbers expressing the same pro-
portion (such as 63.57 : 16) will serve the purpose.
All experiments involving measurement, such as those used in
determining composition, are called quantitative experiments.
Another Quantitative Experiment. — The following will
show how the combining proportions may be measured when the
product is a gas, the weight of which must be ascertained. Sul-
phur bums in oxygen to form sulphur dioxide. A known weight
of sulphur is placed in a porcelain boat in a hard-glass combustion
tube (Fig. 22). The U-shaped tube to the right contains a solu-
tion of potassium hydroxide, which is capable of absorbing the
resulting gas. The oxygen enters from the left. When the sul-
phur is heated, it bums in the oxygen, and the gain in weight of
^^
AIR AND OXYGEN 39
the U-tube shows the weight of the compound produced. By
subtracting from this the weight of the sulphur taken, we get
the quantity of oxygen with which it has combined.
In one experiment, the weight of
sulphur was 1.21 g. The weight -^b ^ny^ ^
gained by the U-tube was 2.42 g. The
difference (= oxygen) is 1.21. The
proportion of sulphur to oxygen in ^iq 22
sulphur dioxide is therefore 1.21 : 1.21
or 1 : 1 or, in percentages, 50 : 50. This proportion is very close
to the accepted value, 32.06 : 32.
The same method could be used for carbon, for the carbon
dioxide produced would be absorbed in the solution of potassium
hydroxide.
Combustion. — Violent union with oxygen is called, in
popular language, combustion or burning. Yet, since similar
vigorous interactions with other gases are conunon, the term
has no scientific significance. Even the imion of iron and sul-
phur gives out Ught and heat, and is quite similar in the chemical
point of view to combustion.
A misleading term often used in this connection is kindliag
temperature. It gives the impression that there is a definite tem-
perature at which combustion will start. But the temperature
is only one of the conditions which produce combustion. Finely
powdered iron will start burning at a lower temperature than wiU
an iron wire, because it presents relatively more surface to the gas.
Again, if the oxygen is at less than one atmosphere pressure, the
wire will require to reach a higher temperature before combustion
will begin. Finally, the vapor of methyl alcohol and air requires
to be raised above a red heat before combustion starts, but a
pocket cigar-lighter sets fire to this very mixture by meaps of a con-
tact agent (a thin platinum wire) without any other means of
heating being required. Therefore, the conditions under which
40 smith's intebmediate chemistry
combustion begins involve the physical condition of the solid, the
pressure of the gas or vapor, the presence or absence of a contact
agent and the nature of the contact agent, as well as the temper-
ature. No definite kindUng temperatiu'e can be given, unless the
other conditions are specified also. Kindling conditions involve
several variables, of which the temperature is only one.
Oxidation. — The slower union with oxygen which occurs in
rusting is called oxidation. We shall see later, however, that it
has been found convenient to stretch this term so as to cover com-
binations of other elements than oxygen, and even to include
actions not involving combination. At this point we can dis-
cuss only oxidation by oxygen.
This process of slow oxidation by oxygen, although less con-
spicuous than combustion, is really of greater interest. Thus the
decay of wood is simply a process of oxidation whereby the same
products are formed as by the more rapid ordinary combustion.
Sewage is mixed with large volumes of river water, the object
being, not simply to dilute the sewage, but to mix it with water
containing oxygen in solution. This has an oxidizing power like
that of oxygen gas and, through the agency of bacteria, quickly
renders dissolved organic matters innocuous by converting them
for the most part into carbon dioxide and water. Thus, a few
miles further down the stream, the water may becoriie as suitable
for drinking as it was before the sewage entered. In our own
bodies we have Ukewise a familiar illustration of slow oxidation.
Avoiding details, it is sufficient to say that the oxygen, from the
air taken into the lungs, combines with the haemoglobin in the red
blood-corpuscles. In this form of loose combination, it is carried
by the blood throughout our tissues and there oxidizes the food-
stuffs which have been absorbed during digestion. The material
products are carbon dioxide and water, of which the former is
carried back to the lungs by the blood, and finally reaches the
air during exhalation. The important product, however, is
AIR AND OXYGEN 41
not material; but the heat, given out by the oxidation, which
keeps the body warm.
The opposite of oxidation, the removal of oxygen, is spoken of
in chemistry as reduction. But this term, also, has been stretched
to cover other similar kinds of chemical change.
Spontaneous Combustion. — Sometimes a mere slow oxi-
dation develops into a combustion, which is then known as spon-
taneous combustion. To understand this, we must note the fact
that a given weight of material, say, iron, in combining with
oxygen to form a given oxide, will hberate the same total amount
of heat whether the union proceeds rapidly or slowly. If the
action proceeds slowly, and the material being oxidized is freely
exposed to the air, the latter will become heated and will carry
o£F the heat as fast as it is produced. Thus, no particular rise
in temperature will occur. If, however, the material is a poor
conductor of heat, hke hay or rags, and there is suflScient air for
oxidation, but not enough to carry o£f the heated air, the heat may
accumulate and a temperature suflScient to start combustion may
be reached.
Such a situation sometimes arises in hay-stacks. It occurs
also when rags, saturated with oils used in making paints
(linseed oil and turpentine) are left in a heap. These oils,
in " drying," combine with oxygen from the air and turn into
a tough, resinous material. The rags, being poor conductors
of heat, may finally become hot enough to burst into flame, and
serious conflagrations often owe their origin to causes such as
this. Oily rags should always be disposed of by burning, or
should at least be placed in a closed can of metal. Fires in coal
bunkers of ships arise from the same cause — slow oxidation, with
accumulation of the resulting heat. That coal does undergo
slow oxidation, especially when freshly mined, is shown by the
fact that such coal, if left exposed to the air for months, may
lose from 2 to 5 per cent or more of its heating value.
42 smith's intermediate chemistry
Uses of Oxygen. — In the foregoing sections we have referred
to its use in breathing, its role in decay, which is a beneficient
process because it removes much useless matter which might
otherwise cause disease, and its value in the disposal of sewage.
Power and heat for conmiercial purposes are almost all obtained
by the burning of coal, in which oxygen from the air plays a large
part. If we had to purchase the oxygen as well as the coal, we
should require at least three tons of oxygen for every ton of coal.
Oxygen in cyhnders and oxygen generators are used to restore
the supply in the atmosphere of submarine boats, and to assist
the respiration of aviators at very great altitudes.
Substances Indifferent to Oxygen. — Finally, since the
atmosphere contains so large a proportion of oxygen, substances
which do not oxidize and, when heated, do not burn, have many
uses. Gold, silver, and platinum are of this kind (p. 36), and are
used for ornaments. The last is used for crucibles in which bodies
are heated in the laboratory. Although iron bums in pure oxygen,
it does not oxidize rapidly in the air even when heated, and so is
used for making vessels for cooking and in constructing fireproof
buildings.
Compounds, already fully oxidized, are naturally not com-
bustible. Of this nature are sandstone, granite, brick, porcelain,
glass, and water. All these are, therefore, fireproof. More-
over these substances do not give off oxygen when heated. It is
this inactivity which renders glass and porcelain suitable materials
for laboratory apparatus, since they experience no change in
weight when heated.
Exercises. — 1. What percentage by weight of free oxygen is
obtained by heating: (a) mercuric oxide, (6) potassium nitrate,
(c) potassium chlorate? At $1.50, $0.15, and $0.15 per kilogram,
respectively, which is the cheapest source of oxygen?
2. Using the data on p. 34, calculate the weight of oxygen dis-
solved by 100 c.c. of water at 0° under one atm. pressure.
AIR AND OXYGEN 43
3. Why does a forced draft make a fire bum more rapidly?
4. Why does a naked flame sometimes cause an explosion in a
mine, when the air of the mine is filled with coal dust?
5. The substances, like phosphorus and sulphur, which bum
rapidly in ordinary oxygen, combine very, very slowly with oxygen
which has been freed from moisture by careful drying. How is
this effect of water to be classified?
6. Air is 20 per cent oxygen. Why does iron bum brilliantly
m pure oxygen, but not in air?
7. What weight of copper would be required to combine with
the amount of oxygen contained in 100 grams of mercuric oxide?
\
\
V
CHAPTER IV
THE MEASUREMENT OF GASES. THE MOLECULAR
HYPOTHESIS
After the discussion in regard to the proportions of oxygen
in the air and the measurement of the volume of oxygen removed
(Fig. 15, p. 26), it will readily be imagined that measuring the
volume of a sample of gas is a common operation in chemistry.
Indeed, it is much easier to measure quantities of gas by noting
their volumes than by weighing them. Some facts have to be
taken account of, however, in order that the measurements of the
volume may be of value.
A sample of gas diminishes in volume when the pressure in-
creases, and it increases in volume when the temperatiwe rises.
Hence, when the volume of the gas is measured, the pressure
and the temperature must also be stated.
Measurement of Pressure. — In order that the pressure may
be easily measured, we arrange the sample of gas so that its pres-
sure is the same as that of the atmosphere at the
moment. To do this, if, for example, the gas is con-
tained in the narrow tube (Fig. 23), we bring the water
inside the tube to the same level as that outside by lower-
ing the tube. Then we read the volume by means of the
graduation (not shown) on the tube and at the same time
we ascertain the atmospheric pressure from the height of
the barometer.
Fig. 23 -^ simple form of barometer is shown in Fig. 33, p. 61
(tube on left of diagram only). A glass tube, about 1
meter long and closed at one end, is completely filled with mercury
and carefully inverted, with its open end dipping into a mercury
reservoir. The mercury falls inside the tube, leaving a vacuum
44
^\
1!HE MEASUKEMENT OF GASES. MOLECULAB HYPOTHESIS 45
at the top, until the pressure of the mercury column inside bal-
ances the pressure of the air outside. The atmospheric pressm-e
is therefore expressed by the barometer as equivalent to the
pressure exerted by so many millimeters of mercury. For accurate
work, all readings obtained must be reduced to a standard tem-
perature (0° C), by correcting for the expansion of mercury above
that temperature.
Correction of the Volume to 760 mm. Pressure. — Since
the barometer varies in height, and the volume of a sample of gas
therefore varies also, it is convenient to " correct " the volmne
further by " reducing " it to that which the gas would occupy at
" standard " pressure, namely 760 mm. Now, the volume of a
sample of gas varies inversely with the pressure (Boyle's law).
If, for example, the volume is 23 c.c. and the observed pressure
of the barometer is 745 mm., Boyle's law enables us to calculate
the volume the same sample of gas would occupy at 760 mm. The
volume changes in the ratio of these two pressures. If the pres-
sure of the gas were actually changed to 760 mm., — a greater
pressure — the gas would assume a smaller volume. Hence,
745
the new volmne = 23 X =^ = 22.5 c.c. That is to say, if the
new volume is to be less, we place the smaller pressure in the
numerator.
If a sample of gas occupies 15 c.c. at 850 mm., what volume will
it occupy at 500 mm.? Here the new pressure is smaller, and the
new volume therefore greater. New volume = 15 X ^7^ = 25.5 c.cl
oUU
Correction of the Volume to (f C — All gases at 0° are
found to gain 1/273 of their volume when heated 1 degree, 2/273
for 2 degrees and 273/273 for 273 degrees. Thus at 273° the
volume is doubled. When cooled below 0°, every gas similarly
loses 1/273 of its volume for each degree. At -273**, if the
46 smith's intermediate chemistry
regular contraction continued so far, the sample would, by cal-
culation, at least, lose all its volume. This temperature (the
lowest temperature that could possibly be attained) is called the
absolvie zero. In point of fact, however, all gases liquefy before
the temperature has fallen to —273°.
The rule contained in these statements is known as Charles'
law. By applying an arithmetical device, we can state the law in
a form which makes its use in calculations quite easy. The
device consists in adding 273 algebraically to all temperatures. The
temperature, when 273 has been added, is called the absolute
temperature. The rule then reads: The volume of a sample of
any gas is directly proportional to the absolute temperature.
Thus, a sample of gas occupies 45 c.c. at 15°, what would be its
volume at 10°? After we have appUed the device, this reads:
a sample of gas occupies 45 c.c. at 15 + 273 = 288° Abs., what
would be its volume at 10 + 273 = 283° Abs.? The volume
changes in the ratio of these absolute temperatures. Since the
new temperature (10° C. or 283° Abs.) is lower, the volume be-
comes smaller. Therefore, putting the smaller number in the
283
numerator, the volume at 283° = 45 X ^^qq = 44.2 c.c.
Again, a sample of gas occupies 125 c.c. at 25°, what will be its
volume at — 15°? The absolute temperatures are 25 + 273 =
298° Abs., and - 15 + 273 = 258° Abs. As the new temper-
ature is lower, the volume will be less. Hence the new volume
125 X |g = 108.2 c.c.
In practice, when a sample of gas is measured, we read the
existing temperature, and correct the volume to that which the
sample specimen would occupy at 0° C. For example, the
volume is 102 c.c. at 18°, what is it at 0°? The absolute temper-
atures are 18 + 273 = 291° Abs., and 0 + 273 = 273° Abs. The
new volume will be smaller. Hence, the new volume is 102 X
THE MEASUREMENT OF GASES. MOLECULAR HYPOTHESIS 47
273
sqt = 95 c.c. This correction enables us to compare all volumes
of gases as if they had been measured at the same standard tem-
perature, namely at 0°.
CtHrrections for Pressure and Temperature Combined. —
Since the volume changes, due to alterations in pressure and
in temperature, are independent of one another, the corrections
may be made either separately or together. The latter is more
convenient. Thus, a sample of gas occupies 190 c.c. at 17® and
750 mm., what will be the volume under standard conditionSi
namely 0® and 760 mm. pressure?
New volume = 190 X =^7^ X t^^?. = 176.6 c.c.
jyu 7ou
Correction for the Tension of Aqueous Vapor. — When
a gas is measiu-ed over mercury, the latter gives oflf practically
no vapor at room temperatiu'e, and the foregoing are the only
corrections required. If, however, the sample of gas is stand-
ing over water, then the volume is not that of the gas, but that
of the gas plus a certain amount of water vapor. The latter
must be subtracted. To do this we have to remember that,
in a gaseous mixture, each one of the several gases (or vapors)
exercises the same pressure as if it were present alone (Dalton's
law of partial i»:essures)« Now the pressure of the water vapor,
in .a gas standing over water, at each temperature is known (see
p. 62 and Appendix IV). It varies from 13.5 mm. at 16® to
23.5 mm. at 25**. When, therefore, the gas is measiu-ed over
water, the pressure of the water vapor is subtracted from the
barometric reading, before the above-mentioned corrections are
applied.
For example, a specimen of a gas, standing over water, occupies
175tC.c. at 19° and 752 mm., what is the volume of the same gas
at 0^ and 760 mm. when the gas is free from water? At 19®
48 smith's intermediate chemistry
the aqueous tension is 16 mm. The pressure of the gas alone is
therefore 752 — 16 = 736 mm. The fully corrected volume =
Molecular Hypothesis, — The relations between pressure and
volume (Boyle's law) and between either of these and temper-
ature (Charles' law) in gases may be explained by the molecular
hypothesis. According to this idea, all matter is composed of
minute particles called molecules, those of any given substance
being all aUke in nature and in mass.
In soUds and hquids these molecules are closely packed
together. In gases, however, they are widely scattered, with
much vacant space between them. A gas is in fact a vacuum,
with numerous relatively minute particles scattered through it.
When a gas is compressed, only the spaces between the molecules
are reduced. By assuming further that, in gases, the molecules
are in rapid motion, and produce pressure by striking the walls
of the vessel, and that this motion is increased by raising
the temperature, all the laws of gases can be completely explained
(see p. 88-93).
Exercises. — 1. Show that, if the levels of the water inside and
outside the tube (Fig. 15, p. 26) are equal, the pressure of the gas
inside must be equal to the atmospheric pressure.
2. Find the volume that 48 c.c. of gas at 732 mm. would occupy
at 760 mm.
3. Reduce 48 c.c. of gas at 780 mm. to standard pressure
(760 mm.).
4. Find the volume which 28 c.c. of gas at 775 mm. would
occupy if the pressure changed to 730 mm.
5. Find the volume which 320 c.c. of gas at 20® would occupy
at OP (pressure imaltered).
6. Reduce 600 c.c. at 25° and 760 mm. to standard conditions
(0^* and 760 mm.).
THE MEASUREMENT OF GASES. MOLECULAR HYPOTHESIS 49
7. What change in volume would occur if 1 liter of gas at 200°
were cooled to 0°.
8. Reduce 260 c.c. of gas at 10° and 742 mm. to standard
conditions.
9. Reduce 35 c.c. of gas standing over water at 21° and 732 mm.
to dry gas under standard conditions.
10. Give the featm-e of the molecular hypothesis which corre-
sponds to each of the following facts about gases:
(a) A specimen of gas can be compressed to l/2000th of the
volume it occupies at 760 mm.
(b) A gas, even if heavier than air, gradually leaves an open
cylinder.
(c) A specimen of a gas exercises pressure equally in all
directions.
(d) Crowding the specimen of gas into one-third of its
original volume triples its pressure.
(e) When the temperature is raised, the pressure of a speci-
men of gas increases (if the volume remains im-
changed).
11. When we drink lemonade through a straw, what causes the
liquid to flow up into our mouths?
CHAPTER V
HYDROGEN
After considering the atmosphere, and particulariy its most
active component, we naturally turn to water, which, Uke air, is
so closely associated with our daily Ufe. We find that water is a
compound of oxygen with hydrogen, and the latter element,
therefore, next claims our attention. Hydrogen is of interest
upon its own account. It is often used for filling balloons. Illimii-
nating gas, of the kind (water gas) used in most large cities,
consists of hydrogen to the extent of about 40 per cent.
Preparation by the Action of Metals on Water. — Hydro-
gen is not easily Uberated from water, since this oxide of hydrogen,
Uke many other oxides, is very stable. It is necessary to use some
element with which oxygen will combine even more
eagerly than with hydrogen, and to oflfer this element
in exchange for the hydrogen.
If a piece of one of the very active metals, such
as potassium, sodium, or calcium is thrown into watery
hydrogen is Uberated and comes oflf in bubbles.
The former two metals are Ughter than water, and
run about on the surface. The action with potassium
is so violent that the hydrogen usually catches fire,
and both with sodium and potassium much heat is
produced. The action often ends with a sUght explosion, so that
a glass plate should be held up to protect the eyes. This experi-
ment is too dangerous to be tried by a novice. With calcium
the action is rapid, but not violent, and there is no danger. The
metal sinks to the bottom of the vessel (Fig. 24), so that a test-
tube or bottle filled with water can be inverted over the metal to
catch the gas as it ascends.
50
<
»
.:<
•••
i"
/
^ ■
Fig. 24
HYDBOGEN 51
The metals, of coiirse, act upon a small part only of the whole
vesselful of water. In each case the metal displaces one-half of
the hydrogen from that part of the water upon which it acts:
Sodium (23) + Water (18) -^ Hydrogen (1) + Sodium hydroxide (40).
Hydrogen Sodium
Hydrogen £^drogen
Oxygen Oxygen
The products are hydrogen, along with potassiimi hydroxide,
sodium hydroxide, or calcium hydroxide. The first two hydrox-
ides are very soluble in water, but most of the calcimn hydroxide
is not dissolved, and may be seen suspended in the Uquid.
Magnesium will Uberate hydrogen from water, provided the
latter is hot. If steam be passed through a heated tube con-
taining iron filings, a mixture of hydrogen, with unused steam,
issues at the other end. The magnetic oxide of iron, not hydrox-
ide, remains in the tube:
Iron + Water —> Hydrogen + Magnetic oxide of iron.
Hydrogen Iron
Oxygen Oxygen
This is a method much used in making hydrogen for commercial
purposes.
Silver, gold and platiniun, ^^hich do not combine directly with
pure oxygen, and even copper and mercury, which do, are all
unable to form oxides and to Uberate hydrogen when heated in
steam.
Displacement. — The foregoing actions, in which hydro-
gen is Uberated, present us with a new — the third — variety of
chemical change. Here an element displaces one of the elements
from a compound, setting it free, and tmites with the rest of the
constituents of the compound. Thus, calcimn displaces part of the
hydrogen, and unites with the oxygen and the rest of the hydrogen.
Preparation by the Action of Metals upon Dilute Acids* —
All the metals which displace hydrogen from water or steam will also-
act upon cold dilute adds and displace the hydrogen they contain.
This action is the one most commonly employed in the laboratory
62
smith's intermediate chemistry
(Fig. 25). The gas, being much Ughter than air, is collected by
downward displacement of the air (Fig. 26b). Heavy gases are
collected by upward displacement of air (Fig. 26a).
r\ Extremely pure zinc is almost inactive, but com-
I mercial zinc, on account of the contact action of
the slight impurities it contains, gives a steady, not
too violent, evolution of hydrogen. Sulphuric acid
and hydrochloric acid, in each case diluted with
0 water, are convenient acids. Iron shows about the
same activity, but, on account of the impurities
Fig. 25 usually present in iron filings or wire, the hydrogen
contains other gases which exhibit a distinct odor.
Fig. 26a
Fig. 26b
Zinc (65.4) + Sulphuric acid (98) -^ Hydrogen (2) + Zinc sulphate (161.4).
Hydrogen (g^PJ^- Z-jS^.^e'^
Iron (56) + Hydrochloric acid (72.9) -* Hydrogen (2) + Ferroi^^chloride
Hydrogen
-Mi
Iron
Chlorine
Cnlorine
It will be seen that the action is of the form we have called
displacement (p. 51).
The Proportions hy Weight. — It may be well to remind
ourselves that the weights of the various materials (given in
brackets) are obtained by laboratory measurements. It is alwaj^
found that the total weight of the product is exactly equal to that
of the materials used (p. 20). Also, that a given weight of the
metal, say zinc, will always displace and liberate the same weight
HYDROGEN 53
of hydrogen. Also, that the proportions by weight of the con-
stituent elements in the compound produced are always the same
(p. 20). If we place a weighed piece of zinc in hydrochloric acid,
and wait imtil the zinc has all disappeared, we can then boil
away the water and unused acid, and weigh the white, soUd zinc
chloride. We find that 65.4 grams of zinc always leave 136.3
grams of zinc chloride. The difference, 70.9, is the chlorine, and
whatever weights we take, the proportion of zinc to chlorine in the
zinc chloride is always in the ratio 65.4 : 70.9.
Zinc (65.4) -h Hydrochloric acid (72.9) -► Hydrogen (2) + ^13^3)^^®
Hydrogen 2 Zinc 65.4
Chlorine 70.9 Chlorine 70.9
Chemically Equivalent Quantities, — It will be observed
that 65.4 parts of zinc displace 2 parts of hydrogen, whether
the acid used is sulphuric acid or hydrochloric acid. The propor-
tion is, in fact, the same with every acid. Hence 65.4 parts of
zinc and 2 parts of hydrogen are spoken of as chemically equiv-
alent quantities. The quantities of the displacing and of the
displaced element are in all cases referred to as chemically equiv-
alent.
The term equivalents is applied also to the quantities liberated
hy decomposition of a compound Uke the 100 parts of mercury
and the 8 parts of oxygen (p. 19). It is Ukewise used of the pro-
portions combining when chemical union takes place, as in the
case of phosphorus 31 parts and oxygen 40 parts (p. 35). The
proportions of the elements in zinc chloride (see above) are also
chemically equivalent.
Chemically equivalent quantities (or, simply, equivalents) of
two substances are exact quantities which enter into or result
from a chemical reaction.
The Order of Activity of the Metals. — It will greatly
aid us in remembering a number of the facts already given, as
well as many others, if we compare carefully with those facts the
64
smith's intermediate chemistbt
order in which the metals stand in the adjoining list. The most
active metals are at the top. All above hydrogen displace this
element from dilute acids (and, with more difficulty, from water) ;
those below it do not.
The first displaces the hydrogen from water so violently that
the gas catches fire, the second displaces it vigorously, the third
less vigorously. Magnesium requires hot water and
iron superheated steam. Copper and the metals
following it do not Uberate hydrogen from water.
Again, the upper metals act too violently on dilute
acids, and zinc is used to prepare the gas. Copper
and the metals following it do pot displace hydrogen
from dilute acids.
Still again, we recall the fact (p. 36) that, when we
heat metals in pure oxygen, the last three do not be-
come oxidized at all. Those preceding silver do com-
bine with pure oxygep — mercury with difficulty, and
the others more and more vigorously as we ascend
the Ust. On the other hand, if we start with the
oxides of all the metals, we find that those at the
foot of the Ust, up to and including mercuric oxide,
lose all their oxygen when heated,
leaving the metal.
Other facts of a sunilar nature wiU
be mentioned as we encounter them.
Meantime, it may be noted that the metals found
uncombined in nature are those following hy-
drogen. Again, the metals known to have been
first used by the human race were gold and silver.
In the ''bronze age" means of Uberating copper
from its ores had been discovered. Lead, tin, and
iron came later. The Ust, read from the bottom
up, gives, therefore, roughly, the historical order in which the
metals came into use.
Ordbb of
Activity.
Metalg
Potassium
Sodium
Calcium
Magnesium
Aluminium
Manganese
Zinc
Chromium
Iron
Nickel
Lead
Tin
Hydrogen
Copper
Bismuth
Antimony
Mercury
Silver
Platiniun
Gold
/^
T
4*
•••
m
Fig. 27
HYDROGEN 55
Preparation by Electrolysis of a Dilute Acid. — A
convenient way of obtaining pure hydrogen is by passing a
current of electricity through a dilvie add (Fig. 27). The gas
18 liberated at the negative wire (cathode) and collects in
the tube (also filled with the dilute acid). The direct, 110-volt
current, passing through a 16-c.-p. lamp placed in series with
the electrolytic apparatus, may Uberate 60 c.c. of hydrogen in
7-8 minutes. Every acid contains hydrogen, combined with
other elements. The other elements are carried to the positive
plate (anode) and therefore do not interfere with the collection
of pure hydrogen. What may be Uberated at the positive plate
depends upon the acid used. With hydrochloric acid, it is chlor-
me; with sulphuric acid, oxygen comes off and sulphuric acid is
regenerated.
Hydrochloric acid (in Aq. Soln.) —> Hydrogen (neg. plate) +
Chlorine (pos. plate).
The process is called electrolysis, from the Greek, meaning
decomposed by electricity.
Physical Properties of Hydrogen. — The gas is colorless,
odmless, and ta^steless. It is, bulk for bulk, the lightest known
gas, the density of air being about 14.5 times as great. It can
be liquefied by compression below —234° (its critical temperature).
It dissolves in water to the extent of 1.8 volmnes in 100 volmnes of
water at 15**.
The lightness of the gas may be. shown by pouring it upwards
from one jar to another, or by balancing an inverted beaker with
shot, and allowing hydrogen to flow in and displace the air.
Several metals can adsorb (" occlude ") hydrogen gas. Iron
takes up about 19 times its own volume, platinum 50 volumes,
and palladium from 500 to 800 volumes.
Liquid hydrogen, when allowed to evaporate rapidly, freezes to
a colorless solid, which melts again at —260°.
56
smith's intermediate chemistry
Chemical Properties of Hydrogen. — That hydrogen, when
it bums in the air, forms water, was first shown by Cavendish
(1781). A test-tube or beaker, filled with cold water and held
over a fiame of burning hydrogen, will condense the steam to
droplets of water (Fig. 28).
Hydrogen + Oxygen —> Water.
The union is very violent, so that when a mixture of hydrogen
with pure oxygen is set on fire (they do not unite when cold)
much heat is Uberated in the explosion. The gases can be made
¥
Fig. 28
Fig. 29
to bum quietly, but with an exceedingly hot flame, by the use of
an oxy-hydrogen burner (Fig. 29), which is constructed like a blast
lamp. Iron melts and bums in the flame. A piece of
quicklime, held in the flame, glows with a brilLant
white hght — the calcium light (Drummond light or
lime-Ught). For such uses the gases are obtained in
compressed form in iron cylinders (Fig. 30).
Hydrogen combines vigorously with chlorine, giving
hydrogen chloride, a gas of which hydrochloric acid is
a solution. It unites with the three most active
metals in the list (p. 54). Calcimn hydride is sold
under the name of hydroljrte, and is used, on ac-
count of its action on water, as a source of hydro-
gen.
Q* ^0 Hydrogen acts upon many compounds containing
oxygen, removing the latter to form water. Thus when the
HYDROGEN 67
oxide of iron or of copper is heated in a tube in a stream of
hydrogen, water is produced and the metal remains:
Magnetic oxide of iron + Hydrogen —> Iron + Water.
Iron Hydrogen
Oxygen Oxygen
Oxides of metals above iron in the " order of activity " (p. 54),
however, are very stable. Hydrogen is unable to remove the
oxygen from such oxides and leave the metal.
Reduction. — The removal of oxygen from a compound by its
union with some other substance is called reduction and the
substance (in the foregoing instance, hydrogen) is called a reduc-
ing agent. Carbon, in the form of coal or coke, is the agent of this
kind most commonly used in chemical industries. The term
reduction is appUed to some other chemical actions, in which oxy-
gen is not concerned. In aU cases, however, reduction is the
opposite of oxidation (p. 41).
Exercises. — 1. Name three varieties of chemical change (pp.
8, 16, 51) and explain the difference between them.
2. What do you infer as to the composition of a substance
when it is named: (a) an oxide, (b) an hydroxide (p. 51)?
3. What are the equivalent quantities of: (a) carbon and oxy-
gen (p. 35), (b) zinc and sulphiw?
4. What law shows that the ratio of chemically equivalent
quantities of any two substances must be constant?
5. Name the metals which: (a) do not liberate hydrogen from
water or dilute acids, (b) are found free in natiwe.
6. Why does hydrogen gas, when poured out, flow upwards?
Why is it an excellent gas for filling balloons?
CHAPTER VI
WATER
As oxygen is necessary to the life of plants and animals, so also
is water. The human body is saturated with it, and water to
make up for evaporation, as well as to aid in digestion, is a most
necessary part of our food. The ocean covers about three-
fourths of the surface of the earth, and the " dry '' land would be
uninhabitable if it were really dry. The air always contains more
or less water vapor.
Measurement of the Composition by Weight. — An ar-
rangement by which the proportion by weight of hydrogen and
Fig. 31
oxygen in water can be determined is shown in Fig. 31. The
bulb A contains cupric oxide, which is heated. Hydrogen from a
generator or cylinder enters on the left and reduces the oxide,
forming copper and water:
Hydrogen + Cupric oxide — ► Water + Copper.
58
WATER 59
The water is carried as vapor by the excess of hydrogen and
passes into the U-tube B. This tube contains calcium cliloride,
a substance which absorbs water greedily and is used therefore
for drying gas^. Here the water is all caught, while the hydro-
gen passes on. The tubes A and B with their contents are
weighed just before, and again just after, the experiment. The
loss of weight in A is the weight of the oxygen. The gain in
weight in B is the water. The difference between these numbers
is the hydrogen. It is foimd that the weights
of hydrogen and oxygen thus ascertained always
stand in a ratio close to 1 (Hyd.) : 7.94 (Ox.) or
1.008 : 8, the proportions accurately determined
by Morley.
Measurement qf the Compoaition by VeA-
ume. — The proportions by volume in which
hydrogen and oxygen combine may be shown by
intEoducing the two gases into a tube, filled with
mercury and inverted in a cyhnder of mercury
(Fig. 32). The volumes, at atmospheric pres-
sure, are read by lowering the tube, after the
introduction of each gas, until the levels of the
merciuy inside and outside are ahke. A spark
from an induction coil passed between the
platinum wires inserted at the top of the tube
causes the union of the gases. The water con-
denses to a slight dew and the volume of the
single gas which remains is measured. Thus,
if 19.5 c.c. of oxygen and 20 c.c. of hydrogen ^•'- 32
are taken, the volume of gas remaining is 9.5 c.c. and this gas is
afterwards found to be oxygen (test, p. 36). The volumes con-
sumed were therefore 19.5 - 9.5 = 10 c.c. of oxygen and 20 c.c.
of hydrogen. The ratio by volume is therefore 2 Hyd. : 1 Ox.
If these exact proportions are used, the mercury fills the tube
60 smith's intermediate chemistry
after the explosion, but is apt to break it by striking the top
violently.
By taking the gases in the exact ratio 1 : 2, and surround-
ing the tube by a wider one through which steam passes, the
condensation of the resulting steam is prevented. It is found
that, when all the gases are measiwed at the same temperature
(here about 100°), a shrinkage of one-third occurs. That is
to say:
1 vol. Oxygen + 2 vols. Hydrogen — ► 2 vols. Steam.
Gay-Lussac^s Law of Volumes. — When other chemical
actions between gases are studied in the same way, it is found
that, in every case, the volumes of the gases used and produced
in a chemical change can always be represented by the ratio of
small whole numbers. This fact is exceedingly interesting.
It was first discovered by Gay-Lussac in 1808. There is no
such simple relation amongst the proportions by weight, which
usually can be expressed only by large niunbers (see p. 36) or by
irregular fractions, so that this evidence of the existence of a
simple rule in regard to combining proportions is the first we
have encountered and is very welcome. The use made of it by
the chemist will be developed in the next chapter.
Physical Properties. — Water is without odor or taste. It is
very pale blue in eolor, as is shown when we look through a con-
siderable depth of water.
Melting ice and freezing water have the same temperatiwe.
The point at which the mercury column of a thermometer stands,
when the instnunent is immersed in such a mixture of ice and
water, is marked 0® on the centigrade scale. This is the freezing-
point of water. The density of ice is Uttle over nine-tenths that of
water.
When water is heated, it gives off vapor more and more freely
until finally it boils. This point is recognized by the fact that
WATER
61
bubbles of vapor form within the Uquid^ rise through it, and burst
on the surface. The temperature of the water (if it is pure), and
of the steam, are now found to be identical, and this thermome-
ter reading is marked 100° C, the boiUng-point. Water is thus
the standard substance used in the graduating of thermometers.
When so used for fixing the temperature of 100°, the atmospheric
pressure must be normal, that is 760 mm., for the temperature of
boiUng is lower when the pressure is lower.
When water has reached the boiling-point, the temperature
ceases to rise, and the heat supphed is used in changing the water
into steam. The evaporation of 1 gram of water consumes 537
heat units or calories (heat of vaporization). The melting of 1
gram of ice consumes 79 calories (heat of fusion). The calorie is
the average amount of heat required to raise 1 gram of water one
degree in temperature between 0° and 100°.
The amoimt of heat required to raise the temperatiwe of a
given mass of water one degree is greater than that
required for an equal mass of any other common
material. Hence, the temperatiwe of the sea changes
more slowly, and within a smaller range, than that of
the rocks which compose the land. For this reason
the climate of islands surrounded by much water is
less variable from season to season within the year
than is that of the continents.
Steam. — The tendency of water to evaporate
at various temperatures is best measured by the
gaseous pressure it exercises. A Uttle water intro-
duced into a barometric vacumn (B, Fig. 33) will
depress the mercury, and the difference in height is
a measure of the vapor pressure of the water. The
temperature may be changed by putting hot water into the tube
surrounding the barometer, and thus the increase in vapor pres-
sure with rising temperature may be shown. At 0° the colimm
B
Fig. 33
62 smiths' intermediate chemistry
of mercury is depressed 4.6 mm., at 10*^ the vapor pressure becomes
9.2 imn., at 20'' 17.4 imn. (see Appendix IV). At IW the level of
the mercury in the tube would be depressed 760 mm., and would
sink to the level of that in the trough. If a Uttle air is first placed
above dry mercury, causing it to fall, the additional depression
produced by adding water is the same as if the air had been absent
(p. 47).
Steam at 100°, subjected to a pressure over one atmosphere, is
condensed to water. Steam cannot be condensed by pressure,
however, if the temperature is over 374°, its critical temperature.
At 100° the steam occupies more than 1700 times the volume of an
equal weight of water.
Steam is a perfectly invisible gas. The visible cloud of fog,
issuing from a valve when steam escapes, is composed of mi-
nute drops of water formed by condensation.
Molecular Relations of Liquid and Vapor. — When the
water was introduced above the barometric column, the vapor, or
gaseous water, could have resulted only from the spontaneous
motion of the molecules in the Uquid. Some of the molecules,
moving near the surface, went off into the space above the water
and became gaseous. To be consistent, we must also conclude
that the vapor above the water is not composed of the same set of
molecules one minute as it was during the preceding minute.
Their motions must cause many of them to plunge into the Uquid,
while others emerge and take their places. When the water is
first introduced, there are no molecules of vapor in the space at all,
so that emission from the water predominates. The pressure of
the vapor increases as the concentration of the molecules of vapor
becomes greater, hence the mercury column falls steadily. At the
same time the nimiber of gaseous molecules plunging into the
water per second must increase in proportion to the degree to
which they are crowded in the vapor. The rate at which mole-
cules return to the water thus begins at zero, and increases
WATER 63
steadily; the rate at which molecules leave the water maintains a
constant value. Hence the rate at which vapor molecules enter
the water must eventually equal that at which other water mole-
cules leave the Uquid. At this point, occasion for visible changes
ceases and the mercury comes to rest. We are bound to think,
however, of the exchange as still going on, since nothing has oc-
curred to stop it. The condition is not one of rest but of rapid
and equal exchange. Such, described in terms of molecules, is
the state of affairs which is characteristic of a condition of equi-
librium. The condition is d3niamic, and not static.
Equilibrium. — This term is used so often in chemistry, and is
used in so unfamiUar a sense, that the reader should consider
attentively what it imphes. Three things are characteristic of a
state of equiUbrium:
1. There are always two opposing tendencies which, when
equiUbrium is reached, balance each other. In the foregoing
instance, one of these is the hail of molecules leaving the Uquid,
which is constant throughout the experiment. It represents the
vapor tension of the liquid. The other is the hail of returning
molecules, which, at first, increases steadily as the concentration
of the vapor becomes greater. This is the vapor pressure of the
■
vapor. These have the effect of opposing pressures and, when the
latter becomes equal to the former, equilibrium is estabUshed.
In all cases of equiUbrium we shaU symboUze the two opposing
tendencies by two arrows, thus:
Water (Uq.) ^ Water (vapor).
2. Although their effects thus neutraUze each other at equilib-
rium, both tendencies are still in full operation. In the case in
point, the opposing hails of molecules are stiU at work, but neither
can effect any visible change in the system. EquiUbrium is a
state, not of rest, but of balanced activities.
3. (and this is the chief mark of equiUbrium.) A slight change
64 smith's intermediate chemistry
in the conditions produces, never a great or sharp change, but
always, and instantly, a corresponding small change in the state of
the system. The change in the conditions accompUshes this by
favoring or disfavoring one of the two opposing tendencies. Thus,
for example, when the temperature of a Uquid is raised, the motion
of its molecules is increased, the rate at which they leave its sur-
face becomes greater, the vapor tension increases and, hence, a
greater concentration of vapor can be maintained. The system,
therefore, quickly reaches a new state of equiUbrium in which a
higher vapor pressure exists.
In the preceding illustration, the evaporating tendency was
favored by a rise in temperature. As an example of a change in
conditions disfavoring one tendency, take the case where the liquid
is placed in an open, shallow vessel. Here the condensing ten-
dency is markedly discouraged, for there is less chance of return
of the emitted molecules. Hence complete evaporation finally
takes place. Elevation of the temperature hastens the process.
A draft insures the diffusion of the vapor away from the surface
of the Uquid, and has therefore the same effect. The two methods
of assisting the displacement of an equiUbrium, and particularly
the second, in which the opposed process is weakened and the
forward process triumphs solely on this account, should be noted
carefuUy. They are appUed with surprising effectiveness in the
explanation of chemical phenomena.
The States of Matter. — Most substances are known in at
least three different states, namely, a crystalUne (or soUd), a
liquid, and a gaseous form. There is no magic g.bout the number
three, however. Very many substances are known in more
states than three. Thus sulphur has a vapor state, two Uquid
states, and several different crystalline forms. There are no
fewer than five forms of ice, different in physical but identical in
chemical properties.
When we wish to transform a substance from one state to
WATER 65
another, we change its conditions (see p. 7). Thus imder at-
mospheric pressure water is converted to ice by reducing the
temperature below 0°, and to steam by raising it above 100°.
Under a pressure of 100 atmospheres, however, water freezes
at — 1°, and boils only at 330°. With still higher pressures, the
change in the freezing-point is much more marked. Thus under
2000 atmospheres pressure water does not freeze above —20°.
If the pressm-e is increased further, different crystalline forms of
ice make their appearance, and the freezing-point rises again
imtil under 20,000 atmospheres pressure water is found to freeze
at +78°!
Water as a Solvent: Natural Waters, — Water has a re-
markable power of dissolving many other substances, and is said,
therefore, to have great power as a solvent. Rain is the purest
natm-al water. As it is formed by condensation of water vapor,
and has been in contact with the atmosphere only, it contains
only oxygen and other gases dissolved from the air, together
with a Uttle dust. Sea water contains the greatest amount of
dissolved rruiterial, namely about 3.5 per cent. River and, espe-
cially, well waters contain materials in solution which have been
dissolved from the soil and the rocks. These materials act chem-
ically upon soap so that well waters are more or less hard, while
rain water is soft. We have already learned (jp. 40) that natural
waters may also contain bacteria, which give rise to putrefaction
and disease. Many river waters contain large amounts of clay
and other insoluble substances suspended in them. The sus-
pended matter can be seen, as it renders the water turbid, but the
bacteria are invisible, so organic matter and bacteria may be pres-
ent in waters which look perfectly clear.
Purification of Water. — The suspended impurities, includ-
ing the bacteria, may be removed by filtration. City waters are
often filtered through extensive beds of gravel, but this treatment
66
SMITH 8 INTERMEDIATE CHEMISTRY
will not remove all bacteria. In many cases^ small amounts of
pjmn, or almn and lime, are added to the water, and the
suspended matter is then allowed to settle, which it does very
quickly, in large basins or reservoirs. By this coagulation
method (p. 470), all but a few of the bacteria are re-
moved. Sometimes the remaining organisms are
destroyed by adding a Uttle bleaching powder before
the water is distributed (p. 226).
In the household, the simplest appliance is the
Pasteur filter. It consists of a tube of unglazed
porcelain, closed at one end (Fig. 34), through the
pores of which the water is forced by its own pres-
sure. The cyUnder (" bougie ") should be cleaned
daily with a brush to remove the mud and the or-
ganisms which collect in its surface.
All forms of filters must be cleaned at short in-
tervals. If this is not done, the organisms multiply
and soon the filter pollutes the water instead of
purifying it.
Filtration does not remove dissolved matter, and
therefore does not soften hard water (p. 13). For this latter
purpose washing powders are used in the laimdry (see p. 390-1).
All organisms can be kiUed by boiling unfiltered water, but the
boiUng should continue for at least 10 to 16 minutes to be effective.
Pure water for chemical purposes is prepared by distUlation
(Fig. 35). Dissolved soUds remain in the flask (or boiler). The
steam is condensed by cold water circulating in the jacket of the
condenser. Freshly distilled water contains only gases dissolved
from the air. If kept in a vessel, however, such water quickly
dissolves traces of glass or porcelain. The purest water is made
by using a platinum tube in the condenser and a platinum bottle
as the receiver.
Chemical Properties: Stability. — As we should infer from
the vigor with which its constituents combine, water is a very
Fig. 34
WATER
67
stable substance. When steam is superheated, hardly a trace of
decomposition occurs. Even when the temperature reaches
2000** (far above a white heat), only 1.8 per cent of the vapor is
Fig. 35
broken up into oxygen and hydrogen. When the steam cools,
these elements recombine:
Water <=± Hydrogen + Oxygen.
\
Hydrates. — Many substances unite with water to give
compounds called hydrates. Thus if we take zinc sulphate (p. 52)
and dissolve it in water and allow the excess of the latter to evapo-
rate, the soUd appears in long transparent crystals. When these
are dried with blotting paper and heated in a test-tube, they give
off a large amount of steam. The hydrate of zinc sulphate decom-
poses and leaves anhydrous (Greek, deprived of water) zinc sul-
68 SMITHES INTERMEDIATE CHEMISTRY
phate. The latter, when once more moistened, changes back
into the hydrate.
Zinc sulphate + Water <=± Hydrate of Zinc sulphate.
Many conmion chemicals are in fact such hydrates. Thus com-
mon blue-stone, used in gravity batteries, is a hydrate of cupric
sulphate. When heated, it loses water and leaves the colorless,
anhydrous cupric sulphate. These are cases of simple com-
bination and decomposition.
Efflorescence. — Some hydrates are so imstable that the
water passes off, even at room temperature, when the hydrate
is left in an open vessel. Thus crystals of washing soda (hydrate
of sodium carbonate) crumble to powder (effloresce) when not
kept in a closed vessel.
Sodium carbonate + Water ^ Hydrate of sodium carbonate.
If a crystal of a hydrate like this is placed in the barometno
vacuum (Fig. 33, p. 61) a considerable vapor pressure of water
is indicated, so that the tendency of the hydrate to decompose,
when this vapor is allowed to escape, is easily understood. The
pressure of water vapor in equilibrium with such hydrates, when
partially dehydrated at ordinary temperatures, is found to be
greater than the average pressure of water vapor in the atmo-
sphere.
On the other hand, when anhydrous cupric sulphate and zinc
sulphate, obtainable from the hydrates by heating only, are
spread out in the air, they return slowly to the hydrated con-
dition. They combine with the moisture in the air. The vapor
pressure of water in the air is greater than the pressure of water
vapor in equilibrium with these hydrates and their anhydrous
products at ordinary temperatures. Calcium chloride absorbs
water vapor (p. 59) because of its tendency to form a hydrate.
Other Chemical Properties. — Water combines directly with
some oxides. Its union with quicklime yields slaked lime or
WATER 69
calcium hydroxide. Other oxides, for example sulphur dioxide or
phosphorus pentoxide, give acids (sulphurous acid or phosphoric
acid). Cases in which water acts chemically upon substances
dissolved in it will be noted later (p. 145).
Reversible Chemical Actions. — The contrary effect upon
an unstable hydrate of leaving the bottle open or closed, referred
to on p. 68, deserves a moment's notice. When understood, it
explains many things in chemistry. The hydration and dehy-
dration are opposite directions of the same chemical change, and
the condensed statements of the actions were written with the
double arrow to indicate this (p. 68). When the bottle is closed,
the water vapor is imable to escape and recombines with the
anhydrous particles as fast as other particles of the hydrate de-
compose. A reversible action therefore can not complete itself, if
the products of the action are kept together and not allowed to
separate. On the other hand, a reversible action will go to com-
pletion, if one of the products escapes, as the water vapor does
when the bottle is left open. This idea enables us to answer
several interesting questions.
For example, why does steam decompose to the extent of 1.8
per cent at 2000 degrees, but not any further? All its parts are
aUke and are therefore equally capable of decomposing. The
answer is, because neither the oxygen nor the hydrogen is removed,
or can easily be removed, from the steam, and so the completion
of the decomposition is prevented by continual recombination of
these gases.
When a reversible action has come to a standstill, we say that
equilibrium has been reached. This means that two opposing
tendencies are neutraUzing one another's effects.
When will reversible actions go to completion? The products
must be of such a nature that they separate easily. In practice
this happens when one is a gas or vapor, like the water vapor
70 smith's intermediate chemistry
coming from a hydrate, while the other is not. The settling of one
product as a precipitate, while the other stays in solution, is, as we
shall see (pp. 127, 129), another common way in which the separa-
tion occurs.
When iron and water are heated in a closed vessel, the hydrogen
and oxide of iron which are produced (p. 51) react with one another
to give back water and iron.
•
Iron + Steam ;:± Hydrogen + Magnetic oxide of iron.
Hydrogen Iron
Oxygen Oxygen
In a closed vessel we could never use this method for preparing
any quantity of hydrogen. To prepare hydrogen by this action
we must leave the tube open, and let the steam sweep the hydrogen
out. This separates it effectually from the oxide of iron, and
prevents the reversal of the action.
Devices depending on mechanical principles Uke this are con-
tinually used in chemistry for securing easy methods of prepar-
ing substances.
The reader can now answer for himself the question why we are
able to prepare oxygen by heating mercuric oxide (p. 15), in spite
of the fact that the action is reversible (p. 29).
Exercises. — 1. If 50 c.c. of hydrogen and 37 c.c. of oxygen are
exploded in a closed tube, which gas remains and what volume
of it is left?
2. Why do not bubbles of steam ordinarily form in water and
rise through it at temperatures below 100**?
3. How many calories of heat would be required to change 5
grams of ice at 0° into steam at 100°?
4. How could one find out how much soUd matter was dissolved
in a sample of water?
5. Define: filtration, distillation, efflorescence, chemical equilib-
rium.
WATER 71
6. Explain how water, even at room temperature, gradually
dries up (p. 64).
7. Why does a strong wind hasten the evaporation of water
(p. 64)?
8. Why does the lower surface of a sheet of ice on the pavement
melt in the sun before the upper one?
CHAPTER VII
CHEMICAL UNITS OF WEIGHT. FORMULA
As we have seen (p. 60), when the volumes occupied by sub-
stances in the gaseous condition, rather than the weights, are
taken as the basis of measurement, the combining proportions are
simple and are expressible by small whole numbers (Gay-Lussac's
law). This shows that there must be some relationship, connected
with chemical combination, between the amounts of different sub-
stances contained as gases in equal volumes. It suggests that
we might do well to take such amounts (weights of equal volumes)
as the standard or unit quantities for chemical purposes. Now
this is precisely what the chemist has found it in practice most
convenient to do, and the present chapter deals with the units of
material based upon comparing equal volumes.
Illustrations of Gay^Lussac^s Law. — Let us first famil-
iarize ourselves with the volimie-measuring point of view in
chemical actions. The following are a few observed facts, begin-
ning with the union of hydrogen and oxygen already discussed
(p. 59) :
(1) Hydrogen (2 vols.) + Oxygen (1 vol.) — > Steam (2 vols.).
(2) Hydrogen (1 vol.) + Chlorine (1 vol.) -^ Hydrogen chloride
(2 vols.).
(3) Chlorine monoxide (2 vols.) — > Chlorine (2 vols.) + Oxy-
gen (1 vol.).
(4) Mercuric oxide (not volatile) — > Mercury (2 vols.) + Oxy-
gen (1 vol.).
(5) Phosphorus (1 vol.) + Oxygen (5 vols.) — > Phosphorus
pentoxide (1 vol.).
(6) Zinc (at 1000^ 2 vols.) + Sulphur (at 1000°, 1 vol.) ->
Zinc sulphide (not volatile)
72
CHEMICAL UNITS OF WEIGHT. FORMULA 73
It will be noted that in some cases, like (2), there is no change in
the total volume. In others there is a shrinkage, as in (1) and (6).
In still others, like (3), where chlorine monoxide decomposes,
there is an increase in volume. In (5), in order that all the mate-
rials may be gaseous, the whole experiment must be done at a
very high temperature (and at some suitable pressure). In (4)
the mercuric oxide itself does not become gaseous, but decomposes,
so that its own relative volume cannot be given. In (6) the zinc
and sulphur can be combined as vapors at 1000°. The product
(zinc sulphide) will not remain gaseous at any temperature at
which its volume could be measured, however, and so its volume
is not recorded.
It must be kept constantly in mind that the law applies to
volumes in the state of gas or vapor only. There is no rule about
the proportions by volume required for the chemical combination
of Uquids and solids.
One can read these illustrations in different ways. For
example: (1) A given volume of steam is formed by union of the
same volume of hydrogen with half as great a volume of oxygen.
(4) Mercuric oxide, when decomposed by heating, gives two
volumes of mercury vapor and one volume of oxygen in every
three volumes of the escaping gases. (5) One volume of phos-
phorus vapor, together with an equivalent quantity of oxygen,
will give one volume of the vapor of phosphorus pentoxide, all
being measured at the same temperature. In fact, whenever two
vaporizable substances are amongst the factors and products of a
chemical change, their volumes thus ^are either equal, or are to
one another in the ratio of whole numbers.
The Standard or Unit Volume. — The volumes in the fore-
going paragraph are simply relative, and the statements are true
of gaseous volumes of any actual dimensions (large or small),
provided only they bear the proper relationship, such as 2 : 1,
1 : 1, or 1 : 5, in each case. An actual value has been chosen,
74
smith's intermediate chemistry
however, for the volume which is the standard or unit in chemistry.
This is the volume occupied by 32 grams of oxygen, which is 22.4
liters at 0° and 760 mm. This volume is equal to that of a cube
about 11 inches in height (Fig. 36). At other temperatures and
pressures this volume, in order to contam the same amount of
material, alters its value, in accordance with the laws of Boyle and
Charles (p. 45). The reason for selecting this particular volume
will be readily seen, so soon as we shall have
presented the actual weights of various
materials which fill it.
The Weights Filling the Unit Vol'
unie, 22.4 Liters. — The following table
contains a few sample substances, and gives
Fig. 36 ^^^ weight (in grams) of each which, in the
gaseous condition, at 0° and 760 mm.,
occupies the cube — 22.4 Uters. In the cases of compound sub-
stances, like water, we have given also the weights of the constituent
elements which together make up the total weight of the com-
pound filling the unit volume:
Weights of Gases in 22.4 Liters (at 0** and 760 mm.).
Substance
Total wt.
Wt. ox.
Wt. hyd.
Wt. chJor.
Oxyiren
32
18.016
2.016
36.468
70.92
86.92
32
16
"m"
* 2.016 "
2.016
1.008
•
^^ .^ o
Water
Hvdroeen
Hydrogen chloride
Chlorine
35.46
70.92
Chlorine monoxide
70.92
The first colxmm (" Total wt.") gives the nmnber of grams of each
substance occupying, as a gas, 22.4 liters at 0** and 760 mm. This
is the standard or unit weight of that substance. In the case of
water, which is a liquid at room temperatures and pressures, a
known volume of steam is weighed. The volume is reduced by
CHEMICAL UNITS OP WEIGHT. FORMULA 75
rule to 0® and 760 mm., and the weight of 22.4 Uters is calculated
from the reduced volume and the measured weight.
These standard or unit weights of substances are conunonly
called molecular weights (see p. 83).
Unit or Standard Weights of the Elements* — Let us now
examine the weights of the constituent elements making up a
cubeful of each substance, as shown in the last three columns of
the above table. We must first be sure we understand what these
numbers are. They are combining proportions, such as we have
given on previous occasions (e.g., pp. 15, 20, 31). They are equiv-
alents (p. 53). We can use them in oiu- condensed form for
representmg chemical changes:
Oxygen (16) + Hydrogen (2.016) -^ Water (18.016).
Hydrogen (1.008) + Chlorine (35.46) -> Hydrogen chloride
(36.468).
We observe at once that the weights in the oxygen column, or
the chlorine column, for example, are not identical. There was
no reafion to expect that they would be aUke, since different sub-
stances diifer in composition. But we do observe that the weights
of any one element are all exact multiples of the smallest number in its
column, either by unity or some other whole number. Thus, for
hydrogen, the weights are: 2.016, 2.016 and 1.008. The unit
weight of water contains exactly the same weight of hydrogen as
does the unit weight of hydrogen itself, and exaMy twice as much
as does the unit weight of hydrogen chloride. Similar relations
hold in the oxygen and chlorine columns. This is a very surpris-
ing, natural fact, and, better still, one for which we instantly per-
ceive a use. This fact greatly simplifies our task of finding some
way of expressing the compositions of substances in a simple
manner. The fact does not apply to a few compoimds only. If
our table had included all the hundreds of compounds of chlorine
(for example) which are capable of being converted into vapor, we
should have found, indeed, many multiples of 35.46 larger than
the two units (70.92) in chlorine monoxide, but no number smaller
76
smith's intebmbdiate chemistry
than 35.46 and none which was not a multiple of 35.46 by a whole
number. Clearly we shall find it convenient to accept 35.46 as the
unit or standard weight of chlorine for expressing the compositions
of its chemical compounds. Its use will enable us to state the
exact composition of any given compound by simply giving the
whole number (1, 2, 3, etc.) by which the hdsal weight 35.46 is to he
multiplied in the given case.
Similariy, we take 1.008 as the unit weight of hydrogen and 16
as the unit weight of oxygen. By including volatile compounds of
other elements in our investigation, we can similariy pick out the
most convenient unit for each element. Those unit weights for
elements are often called also combining or reacting weights, and
still more frequently atomic weights (see p. 86). A complete
list of their values for all the known elements is given in a table
inside the rear cover of this book. These numbers will hereafter
be in constant use.
The following table presents the results obtained from a larger
number of substances and gives their formulae (see p. 78) :
Substance
Hydrogen chloride
Chlorine dioxide
Phosphorus trichloride . .
Phosphorus oxychloride .
Phosphorus pentoxide . . .
Phosphine
Water
Methane
Acetylene
Ethylene
Formaldehyde
Acetic acid
Mercurous chloride
Mercuric chloride
Molsc-
ular
Weight
36.468
67.46
137.42
153.42
284.16
34.06
18.016
16.03
26.02
28.03
30.02
60.03
236.06
271.52
Weights of Constituents in Molecular Weight.
0
I
1.008
3.024
2.016
4.032
2.016
4.032
2.016
4.032
0
I
35.46
35.46
106.38
106.38
35.46
70.92
d
32
• • •
16
160
16
16
32
o
so
O
,4
31.04
31.04
124.16
31.04
a
o
12
24
24
12
24
2
200.6
200.6
Molecular
Formula
HCl
CIO2
PCla
POCl,
P4O10
PH,
H2O
CH4
C2H2
C2H4
CH2O
C2H40a
HgCl
HgCl,
CHEMICAL UNITS OF WEIGHT. FORMULA 77
The Case of Non-volatile Compounds. — The same ele-
ments enter also into many compounds which are not volatile.
But the unit weights of such elements, determined by the use of
volatile compounds, are found (by using multiples, when neces-
sary) to express the composition of the involatile compoimds also.
For example, the unit weight for oxygen (16) and that for mer-
cury (200.6), both formed by studying volatile compoimds, are
foimd correctly to express the composition of mercuric oxide
(p. 20) , which is not volatile.
In the cases of some elements no easily vaporizable compound
is known, and the unit weight cannot be determined by the present
method. In such instances, an entirely different way of obtaining
the value of the unit weight is employed (see p. 87).
The Law of Combining Weights. — The general fact which
we have developed in the preceding sections is known as
the law of combining weights: All the proportions in which the
elements combine with one another may be represented by a set of
numbers (one for each element) or by multiples of these numbers
Definition of Reacting or Atomic Weight, — Many differ^
ent values for each element would satisfy the law of combining
weights as stated above (see, for example, the values given for
oxygen on p. 10). The particular values chosen as units in this
chapter, however, fulfil an additional condition which fixes the
value in each case, absolutely. The chosen reacting or atomic
weight of an element is the smallest weight of the element found
in 22.4 liters (at 0° and 760 mm.) of the vapor of any volatile com-
pound of that element (p. 74). The amounts of the element in
22.4 hters of other compounds are all either the same amoimt or
multiples thereof by a whole number.
Simplification of Condensed Expressions for Chemical
Actions — Symbols. — Our condensed expressions for chemical
78 smith's intermediate chemistry
changes showed the name of each substance and also weights to
indicate the combining proportions. Having now foimd a imit
weight for each element, we can condense the statement still
further, by itsing a letter or pair of letters, called a symbol, to stand
for one chemical unit weight of each element. Thus O stands for one
unit or 16 parts of oxygen, H for one unit or 1.008 parts of hydro-
gen, CI for one unit or 35.46 parts of chlorine. A pair of letters is
required when the names of several elements begin with the same
letter. Thus, C stands for 12 parts of carbon, Ca for 40 parts of
calcium, Cr for 52 parts of chromium, each of these amounts being
one chemical unit of the element. These symbols are interna-
tional, and are alike in all languages. In some instances they are
based upon the Latin name for the element. Thus, Fe stands for
55.84 parts of iron (Jerrum), Sn for 119 parts of tin (stannum), Ag
for 107.88 parts of silver (argentum) and Hg for 200.6 parts of
mercury (hydrargyrum). Again K stands for 39.1 parts of potas-
sium (German, kalium) and Na for 23 parts of sodium (German,
natrium).
Formuke. — The composition* of a substance can be shown
briefly by putting together the symbols of the constituent elements,
and using numbers for the multiples of the unit weights, where
such are required. The resultmg expression is called a formula.
In the last colmnn of the table on p. 76 the formulae of the various
substances therein considered are indicated. Thus water contains
oxygen 16 parts (0) and hydrogen 2.016 parts (= 2 X 1.008 =
H2), and receives the formula H2O. Hydrogen chloride contains
hydrogen 1.008 parts (H) and chlorine 35.46 parts (CI), and its
formula is HCl.
It must be noted particularly, that the formtda, to be consistent,
must represent also the total weight of the unit quantity of the sub-
stance (wt. of 22.4 1.) in the gaseous state. Thus, HCl (1.008
* The composition means the names of the elements contained in the
substance, and also the proportion by weight of each element (see p. 38).
CHEMICAL UNITS OF WEIGHT. FORMULiE 79
+ 35.46 = 36.468) is correct, since 36.468 g. is the amount filling
the cube. Again, the formula for oxygen gas itself must represent
32 g. (= 2 X 16), the weight of 22.4 1., and is therefore O2. Half
as much as this may enter into a compound, H2O, ZnO, etc., but,
logically, the formula for free oxygen must record the double
weight required to fill the cube when the gas is present alone, in the
free condition (see pp. 86, 105). Similarly the fonnula for hydrogen
gas is H2 (2 X 1.008 = 2.016, the weight of 22.4 1.). The formula
of every volatile substance must thus he written so as to show the
weight of the chemical unit quantity (the molecular weight, p. 84).
When the substance is not easily volatiUzed, this unit cannot be
measured, and the simplest formula is employed.
Information Contained in Each Formula. — A fonnula
thus contams, in condensed form, several items of infonnation.
It shows:
1. The elements making up the substance,
2. The proportion by weight of those elements,
3. The total unit weight (molecular weight) of the substance.
Given the formula, we can read these facts in it.
Thus, if we are given the formula of carbon dioxide, CO2, we
consult the table of reacting or atomic weights (inside rear cover),
and leam that C = 12 parts by weight of carbon and O2 = 2 X 16
parts by weight of oxygen. The proportion by weight of the ele-
ments in this compound is, therefore, 12 of carbon to 32 of oxygen.
The total weight (molecular weight) is 12 + 32 = 44, and this
must be the weight fiUing 22.4 1. (at (f and 760 mm.).
Why was 22.4 Liters Chosen as the Unit Volume? — It is
clearly advantageous to employ a unit volume of such dimen-
sions that no element shall receive a unit weight smaller than
unity. At first sight, indeed, it would seem self-evident that we
ought to choose, for our own convenience, a unit volume based on
80 smith's intermediate chemistry
a smallest unit weight of exactly 1. Now hydrogen is the element
which enters in the smaUest proportions into chemical combination,
and we have seen that in 22.4 1. of some compomids there is as
Uttle as 1.008 g. of this element (see table, p. 76).
If a shghtly smaller miit volmne (22.21.) had been chosen,
hydrogen would possess a unit weight of exactly 1. Why was
not, it will be asked, this volume selected instead of 22.4 1. as our
standard? It was, in fact, at one time the standard, the unit
weight of oxygen being then about 15.9 and the unit weights for
all of the other elements being correspondingly reduced. The
transfer to the present scale O = 16 was made for a purely prac-
tical reason; because of the greater activity of oxygen (see p. 27),
the exact determinations of the combining proportions of most
other elements are generally given by experiment on an oxygen,
not on a hydrogen, basis. It so happens, also, that the majority
of atomic weights on the scale O = 16 are very close to whole
nimibers (see table), an aid to calculations being thereby afforded
which does not exist under the standard H =^ 1. Recent work
on the " complexity " of elements, it may be added, has given fur-
ther justification of our choice of the oxygen standard here adopted
(see p. 552).
The accepted scale is, therefore, that of 32 for the unit weight of
oxygen gas; 22.4 1. (the volume of 32 g. of free oxygen) for the
chemical unit of volume; and 16 for the unit weight of oxygen in
compounds containing that element. The unit weights of all the
other compounds and elements are based upon this scale.
Further Discussion. — At this point we have completed the
explanation of the derivation of the units of weight, and of the
sjrmbols and formula used in chemistry. We are now in a posi-
tion to proceed with the application of these conceptions, as they
are developed in Chapter IX. This would be possible, at least
so far as strict logic is concerned. But, in shaping our course
so as to reach the results by the shortest route, we have omitted
CHEMICAL UNITS OF WEIGHT. FORMUKB 81
a number of interesting and useful ideas. These were not indis-
pensable links in our reasoning, but, now that we have covered
the essential steps, a consideration of these ideas will be of great
value as a review of the subject and a means of mastering it
thoroughly. FormulaB and their uses come up so constantly
in every part of the science that we must not omit any means of
securing a perfect imderstanding of them. The following chapter
is therefore devoted to restating in other ways certain parts of
the same subject.
CHAPTER VIII
APPLICATION OF THE MOLECULAR HYPOTHESIS IN CHEMISTRY
Avogadro^s Hypothesis. — We have seen (p. 48) that the
physical behavior of matter, and particularly of gases, may be
explained by the conception that matter is composed of molecules.
The easy compressibility of gases is, under this view, a consequence
of the smallness of the molecules and the relative vastness of the
empty space between them. The pressure exerted by a gas is
regarded as due to the innumerable blows which the molecules
deliver when they strike the boundary walls of the space in which
the gas is confined. Now, when gases interact chemically, the
voliunes required for complete interaction are either equal or
stand in the ratio of small whole numbers (Gay-Lussac's law,
pu 72). Although that is a chemical, and npt merely a physical
fact, the molecular hypothesis can explain it also.
Since matter is composed of molecules, a chemical action
betwieen two kinds of matter must consist really in an interaction
of the molecules of each kind. Molecules of the two kinds must
meet and they may then either combine to form a compound
molecule, or they may exchange material in some fashion. Since
equal volunies are often the exact quantities required for the
action, it appears most Ukely that in equal volimies of different
gases (at the same temperature and pressure) the ntmibers of
molecules present are equal. This addition to the molecular
hypothesis was first suggested by an Italian physicist, Avogadro
(1811). When two gases interact in equal volumes (Uke hydrogen
and chlorine), one molecule of each is all that is required for a
small sample of the change in question. Since two volumes of
hydrogen are required to unite with one volmne of oxygen (p. 60),
clearly the interaction involves two molecules of hydrogen for
82
APPLICATION OF THE MOLECULAR HYPOTHESIS 83
every one of oxygen. Since two volumes of steam are produced,
evidently the two molecules of hydrogen and one of oxygen yield
two molecules of water.
This hypothesis of Avogadro helps us to a clearer notion of how
these chemical changes take place. The idea that all gases con-
tain equal numbers of molecules in equal volumes also explains
why aU gases behave aUke when subjected to equal pressures,
or to equal changes in temperature (Boyle's and Charles' laws, p.
45)*. No facts which conflict with this hypothesis are known, and
all the known facts confirm it. Hence, Avogadro's hypothesis has
been accepted by chemists, and since 1858 has been the keystone of
chemical theory.
Consequences of Avogadro* s Hypothesis — Molecular
Weights. — Equal volimies of the same gas (at the same tem^
perature and pressure) have equal weights. But equal volumes
of different gases have different weights. The differences are often
very great. Thus, bulk for bulk, oxygen is sixteen times as
heavy as hydrogen, and mercury vapor one hundred times as
heavy. Now, if equal volumes of different gases contain equal
nimibers of molecules, these differences must be due to the differ-
ing weights of the several kinds of molecules. Thus, measuring
the weights of equal volumes of different gases will give us the
relative weights of their molecules. For example, since 22.4 1. of
oxygen weigh 32 g. (p. 74), while the same volume of hydrogen
weighs 2.016 g. and of water vapor 18.016 g., and these are the
weights of equal numbers of molecules, the individual molecules
must differ in weight in the ratio 32 : 2.016 : 18.016. These are the
relative weights of the three kinds of molecules. In chemistry the
weights of 22.4 1. (at 0° and 760 mm.) of various gases are called
the molecular weights of those gases. The unit quantities of
various substances are therefore spoken of, technically, as the
molecular weights of those substances. The unit volmne, 22.4 L,
is called the gram-molecular volume (G.M.V.).
84 smith's intermediate chemistry
The numoer of molecules actually contained in the G.M.V.
has been detennined by several independent methods, the agree-
ment of the results obtained furnishing very strong suppbrt for
the vaUdity of the molecular hypothesis. The value at present
accepted as most accurate is that of Millikan, 6.06 X 1(F*. It
is of importance that the student should obtain some idea of the
significance of this stupendous number, in order to appreciate
the more detailed discussion of the properties of gases in the follow-
•
ing sections (p. 88-93). Imagine a G.M.V. of a gas at 0° and 760
mm. to be divided equally among the inhabitants of the United
States (say 110,000,000). Each person would receive about one-
fifth of a cubic millimeter as his share. Imagine, further, that
the market price of a million molecules of this gas was one cent.
Few of the recipients would think it worth while to cash in their
microscopic sample; those doing so, however, would benefit to
the extent of over 50 miUion dollars!
Molecular Weight, Measurement of. — The molecular
weight is measured by weighing any convenient volume of the gas
(say 200 c.c), and calculating by proportion the weight of 22.41.
If the gas or vapor was not measured at 0° and 760 mm., the
measured volume must be reduced by rule to standard conditions
before the weight of 22.4 1. is calculated, f
* This is the method used in physics and chemistry to express numbers
so large as to be cumbrous and incomprehensible if written in the ordinary
way. It means that the unit is to be followed by the number of zeros
indicated by the exponent. In this case, the quantity written out at length
would be 606,000,000,000,000,000,000,000.
t In practice, owing to the fact that Boyle's and Charles' laws do not
describe the behavior of any known gas exactly (they apply only to a "perfect"
gas), certain additional, small corrections have to be applied when very pre-
cise values are required. It may be noted that, in order to make the funda-
mental significance of the nimierical data presented in the preceding pages
immediately inteUigible to the beginner, all of the gases so far considered have
been assumed to be "perfect." The actual experimental values differ only
very slightly from those given.
APPLICATION OF THE MOLECULAR HYPOTHESIS 85
For example, 190 c.c. of a gas at 0° and 760 mm. weigh 1.23 g.
If X be the weight of 22.4 Uters (= 22,400 c.c.),
190 : 1.23 :: 22,400 : x (= 145 g.).
Again, 210 c.c. of a vapor at 100° and 743 mm. weigh 1.12 g.
This volume at QP and 760 mm. would become:
210X|^X^ = 150.3 c.c.
and
150.3 : 1.12 :: 22,400 : x (= 167 g.).
Molecular Weight and Density. — Density is the term used
in phjrsics for the weight of 1 cubic centimeter of a substance.
Thus, the density of water at 4° C. is 1, because at 4° C, 1 c.c. of
water weighs 1 g. The density of anmionia gas (wt. in g. of 1 c.c.
at 0° and 760 mm.) is 0.000759.
Since the molecular weight is the weight of 22.4 L, or 22,400 c.c,
the molecular weight of a gas is obtained by multiplying the
density (if known) by 22,400. Thus, the molecular weight of
ammonia is 0.000759 X 22,400 = 17.0.
Densities are often given on the scale, density-of-air = 1.
Now 22.4 Uters of air weigh 28.95 g. If a gas has a density twice
that of air, 22.4 Uters of this gas would weigh 28.95 X 2 g. For
example, the density of carbon dioxide (air = 1) is 1.52. The
molecular weight is therefore 28.95 X 1.52, or 44.0 (see p. 79).
Atomic Weights. — Since a compoimd substance can be
formed by union of the elementary substances, and decomposed
to give these substances, its molecule may be assumed to contain
those constituent elements as distinct parts of the mass. Those
elementary parts of a molecule are caUed atoms. When two ele-
mentary substances combine, the process involves the union of
the two kinds of atoms to form compound molecules. Now, in
the table (p. 76), we recorded the weights of the constituent ele-
ments which together made up the molecular weights of the com-
86 smith's intermediate chemistry
pounds. Then, upon examining the weights of any one element
contained in molecular weights of different compoimds, we saw
(p. 77) that one (the smallest) could be taken as the unit weight,
of which the others were multiples by some whole nimiber. Ob-
viously then, if the molecular weights are the relative weights of
different kinds of molecules,^ the unit weights of the elements
(35.46 for chlorine, etc.) are tne relative weights of the atoms of
the different elements. The relative weights of the different
kinds of atoms, such as 1.008 for hydrogen, 16 for oxygen, 35.46
for chlorine, are called the atomic weights of the respective ele-
ments. These weights, as we have seen (p. 80), are relative to
the weights of the atom of oxygen, when the weight of the latter is
taken as 16, and the weight of the molecule of oxygen is taken as 32.
A compound molecule may contain one, or more than one, atom
of each of the elements forming the compound. In the molecular
weight of compounds, weights of elements which are smaller than
the atomic weight (p. 77), are never found to occur. This
indicates that, in chemical change, fractions of atoms play no
part. The name atom (Greek, not cut, or not divided) records
this fact.
The Atomic Hypothesis and Definite Proportions. — The
idea of atoms furnishes at once an explanation of the law of definite
proportions. Evidently, every molecule of a given substance must
always contain the same numbers and the same kinds of atoms, so
that the proportions by weight of the constituents of the com-
pound as a whole must be aUke in all samples.
Dulong and PetiVs Law. — When an element gives no volatile
compoimds, as is the case with calcium, we can still find the atomic
weight. We first find for the element an equivalent weight (p.
53). Thus, the weight of calcium which combines with 35.46
parts of chlorine is 20 parts. This weight, or some multiple of it
by a whole number, does in fact always correctly represent the
APPLICATION OP THE MOLECULAR HYPOTHESIS
87
amount of calcium combined with atomic weights of elements
whose atomic weights are known. Hence 20, or some multiple of
20 is the required atomic weight. Dulong and Petit (1819) dis-
covered a rule by which the correctness of such an atomic weight
could be checked. When the atomic weight of an element is
multiplied by the specific heat of the element in the solid form,
the product is roughly 6.4. Thus, the known atomic weight of
magnesium is 24.3, its specific heat 0.245, and the product is
5.95. Again, the known atomic weight of mercury is 200.6, and
its specific heat 0.0335, and the product 6.7. Now the specific
heat of calciiun is 0.170, and the product 20 X 0.170 is equal to
3.4. But 2 X 20 X 0.170 = 6.8. The value 2 X 20, or 40 for
the atomic weight is therefore the correct one. The following
table contains additional illustrations of this law:
Element
Lithium . . .
Sodium
Magnesium
Silicon. . . .
Phosphorus
(Yellow).
Calcium. . .
Atomic
Wt.
Sp. Ht.
Prod-
uct
7
0.94
6.6
23
0.29
6.7
24.3
0.245
6.0
28.3
0.16
4.5
31
0.19
5.9
40
0.170
6.8
Element
Iron
Zinc
Bromine (Solid)
Gold
Mercury (Solid)
Uranium
Atomic
Wt.
56
65
80
197
200
238
Sp. Ht.
40
10.112
093
0.084
0.032
0.0335
0276
50
Prod-
uct
6.3
6.1
6.7
6.3
6.7
6.6
Another way of expressing this law will give it greater chemical
significance. The specific heats are the amounts of heat required
to raise one gram, that is one physical unit, of each element through
one degree. When we multiply this by the atomic weight, we
obtain the amount of heat required to raise one gram-atomic
weight of the element, that is, one chemical unit, through one
d^ree. The values of this product are approximately equal.
Since there are equal nmnbers of atoms in one gram-atomic weight
of each element, it follows that: Equal amounts of heat raise
equal number of atoms of all elements in the solid form through
equal intervals of temperature.
88 smith's intermediate chemistry
Relations between the Structure and Behavior op
Matter
We have seen that matter is composed of exceedingly minute
particles called molecules. Just as we can thoroughly understand
the behavior of a watch or an automobile engine only if we know
the details of its structure, and how the parts work, so we can
understand the physical and chemical behavior of matter in masses
only if we are familiar with its ultimate mechanism. Hence, we
must now take up the structure of matter in its three states, the
gaseous, the liquid, and the crystalline (or solid). In doing this,
we shall keep constantly in view the connection between the molec-
ular relations and the general behavior of the matter.
The Properties of Gases, — The most remarkable thing about
a gas, considering the looseness with which its material is packed,
is the total absence in it of any tendency to settling or subsidence.
Since the molecules cannot be at rest upon one another, as the
great compressibility shows, we are driven to conclude that they
are widely separated from one another, and that they occupy the
space, otherwise a complete vacuum, by constantly moving about
in all directions. But a moving aggregate of particles which
does not even finally settle must be in perpetual motion. We must,
therefore, believe the molecules to be wholly unlike larger particles
of matter in having perfect elasticily, in consequence of which
they undergo no loss of energy after a collision. They must
continually strike the walls of the vessel and one another and re-
bound, yet without loss of motion. The fact that each gas is
homogeneoKSy efforts to sift out lighter or heavier samples having
failed, requires the supposition that all the molecules of a pure
gas are closely alike.
The diffimUlity of gases is due to the motion of the moleculesi
and their permeability to the space available to receive molecules
of another gas. These two modes of behavior involve no addi-
tional molecular properties. The word " diffusion " is often
APPLICATION OF THE MOLECtJLAB HYPOTHESIS
89
thought to mean the property of a given mass of gas in virtue erf
which another gas can mix with the given mass. This property
is not diffusibiUty but permeabiUty. It is the other gas, which
makes its way into the given gas, which is diffusing. Diffusion
is spontaneous motion of the parts of a gas away from their original
location. Unless this motion is into an empty space, the diffusing
molecules must, of course, move into another body of gas. In
the case of the jars of hydrogen and air (p. 55), each gas moved in
part out of its original jar (diffused), and each received parts of
the other gas into its jar (was permeated).
Bayle^s Law and Charles^ Law. — Passing now to Boyle^s
law (p. 45), the thing to be accoimted for is that when a sample of
a gas diminishes in volume, its pressure increases in the same pro-
portion. Let the diagram (Fig. 37) represent a cyUnder with a
movable piston, upon which weights may be placed to resist the
pressing. Now the pressure exercised by the gas under the
piston cannot be like the pressure of the hand upon a table, since
we have just assmned that the particles are not even approxi-
mately at rest, and the spaces between them are enormous com-
pared with the size of the molecules themselves. The gaseous
pressure must therefore be attributed to the colossal
hailstorm which their innimierable impacts upon the
piston produce. If this is the case, the compressing
of a gas must consist simply in moving the partition
downwards, so that the particles as they fly about
are gradually restricted to a smaller and smaller space.
Their paths become on an average shorter and
shorter. Their impacts upon the walls become more
and more frequent. So the pressure which this causes
becomes greater and greater, and is proportional to
the degree of crowding (the concentration) of the molecules.
There are two other points to be added. When we diminish the
voliune to one-half, we find from experience that the pressiure
Fig. 37
90 smith's intermediate chemistry
becomes exactly, or almost exactly, twice as great. This must
mean that, although the particles are becoming crowded, they
do not interfere with one another's motion, excepting of course
where actual coUision causes a rebound. Only in the absence
of interference would doubUng the nimiber of molecules per unit
of volume give exactly double the number of impacts on the walls.
Hence the molecules must have practically no tendency to cohe-
sion. Finally, the molecules must be supposed to move in straight
lines between collisions.
Boyle's law therefore adds four more details concerning molec-
ular behavior, namely, that the impacts of the particles produce
the pressure f that the crowding of the molecules represents the con-
centration of the maierial and that the particles move in straight
lines and show almost no cohesion, since pressure and concert-
trdtion are very closely proportional to one another.
How, now, can we account for Charles' law (p. 46), according to
which an increase in voliune (or in pressure, if the volume is kept
constant) results from heating a mass of rapidly moving mole-
cules? The action of a particle coUiding with a surface is meas-
ured in physics in terms of its mass and its velocity. It is evident
that heating a cloud of molecules would not increase the mass
of each, and it must therefore increase the velocity.
Diffusion. — It is found by experiment that the Ughter a gas is,
the faster its particles move by diffusion (p. 89) in any direction.
Exact measurement shows that the rate is inversely proportional
to the square root of the density of the gas. From the results ob-
tained for any particular gas, the average speed at which the
molecules of which it is composed are moving can be calculated.
For the hydrogen molecule at room temperature, this speed is
1840 meters per second, or faster than a rifle bullet. The speed
of the oxygen molecule, which is almost exactly 16 times as heavy,
is consequently the square root of one-sixteenth (that is, one quar-
ter) of this value, or 460 m. per sec.
APPLICATION OF THE MOLECULAR HYPOTHESIS 91
Each one of the enormous number of molecules in any given
sample of gas is moving, therefore, with an enormous velocity,
and the pressure exerted by gases (explained above as due to the
impact of the molecules on the walls of the containing vessel)
becomes more intelUgible. In air, the number of molecules
striking a single square centimeter of surface per second would
fill no less than 12 Uters.
Liquefaction of Gases: Critical Temperature* — All gases
can be liqiiefied by sufficient cooling and compression. This fact
compels us to suppose that, after all, even gaseous molecules
have some tendency to cohesion. This cohesion is, in general,
scarcely perceptible so long as the gas is warm and is diffuse but
it is possible, by careful measurement, to establish the fact that it
does cause sUght deviations from Boyle's Law even imder such
conditions. Thus, 2 Uters of oxygen at 1 atmosphere pressure,
when subjected to 2 atmospheres pressure, give 0.9991 Uters
instead of 1 Uter. The additional contraction of 0.0009 Uters
(0.9 c.c.) is to be regarded as due to the effect of cohesion when
the molecules are thus crowded closer together. The gases which
are more easily Uquefied than is oxygen show greater
effects. Thus, 2 Uters of sulphur dioxide at 760 mm., f^
when subjected to 2 atmospheres pressure, give only 0.974
Uters, showing a contraction due to cohesion of 26 c.c.
These data refer to 0°. At lower temperatures the con-
tractions due to cohesion become rapidly greater.
We can readily understand, therefore, that when the
velocity of the molecules is sufficiently reduced by cool-
ing, and the molecules are brought sufficiently close
together, the tendency of the molecules to cohere causes yjq 33
the gas to condense and assume the Uquid form. In 1869
Andrews foimd that carbon dioxide could be Uquefied at 0° by
38 atmospheres pressure, and at 30° by 71 atmospheres, but
tiiat above 31.36° it could not be Uquefied by any pressure.
92 smith's intermediate chemistry
He discovered that each gas has a critical temperature, as he called
it. For carbon dioxide, this temperature can be observed by
placing a heavy-walled, glass tube (Fig. 38), half-filled with
liquid carbon dioxide, in a beaker of water, and gradually raising
the temperature of the latter. At 31.35°, the surface between
the Uquid and gas becomes hazy and vanishes. At this tem-
perature the Uquid state disappears, merging into the gaseous.
When the temperature falls once more, the surface marking the
boimdary between Uquid and gas reappears.
The critical temperature of oxygen is —118°, of hydrogen
-234°, of sulphur dioxide 156°, of water 358°.
Another Deviation from the Laws of Gases. A Perfect
Gas. — It is also found by experiment that when a gas is already
imder very high pressure, and therefore very closely packed,
an increase in the pressure does not prodiLce quite as greai a diminvr
Hon in volume as Boyle's law leads us to expect. This reminds us
that we are diminishing only the space between the molecules,
and not the volumes of the molecules themselves, and therefore
not the total volimie of the gas. When, on severe compression,
the volume occupied by the inolecules themselves has become
an appreciable fraction of the whole vrolume, additional compres-
sion does not affect the whole volmne, and the contraction is
smaUer than Boyle's law would indicate. Thus, 200 Uters of
hydrogen, at 16° and imder one atmosphere pressure, when sub-
jected to 200 atmospheres pressure, give 1.134 Uters, instead of
1 Uter.
The last two effects (namely, those due to the tendency to
cohesion of, and the space occupied by the molecules) are called
deviations from the laws of gases. In consequence of these in-
dividual deviations, there are not exactly equal numbers of mole-
cules in equal volumes of any two different gases, at the same
temperature and pressure. An imaginary gas, which exhibits
neither deviation, caUed a perfect gas, is often referred to in
discussing the behavior of gases (see footnote, p. 84).
APPLICATION OF THE MOUBCULAR HYPOTHESIS
93
Summary. — We may now summarize the principal facts
about gases in mass, with the corresponding f eatiures of the molec-
ular relations.
I^BotB About Gases in Mass
Compressibility
Diffusibility
Permeability
Non-«ettling
Homogeneity
Pressure
Boyle's law
Charles' law
Gay-Lussac's law
Law of diffusion
Gases can be liquefied,
and show lower com-
pressibility at high
pressures.
Corresponding Relations of Molecules
Vacuum + molecules widely separated.
Molecules in rapid motior.
Empty space relatively large.
Molecules perfectly elastic.
Molecules of any one substance closely alike.
Due to impacts of molecules.
Pressure proportional to concentration of the
molecules. Molecules move in straight lines
and, when widely scattered, show no tendency
to cohesion.
Rise in temperature increases the velocity of
the molecules.
Avogadro's law.
Speed of molecules inversely proportional to
square root of mass.
Molecules do possess some tendency to cohesion,
and do occupy an appreciable fraction of the
space when pressure is very great.
The Properties of Liquids. — The fact that even great pres-
sures produce Uttle diminution in the volume of a Uquid shows
that the free space, so predominant in gases, is relatively unim-
portant in Uquids. This is only natiu'al in view of the smaller
volume occupied by substances in the hquid state.* Thus we
have seen (p. 62) that steam at lOO"^ occupies more than 1700
times the volume of an equal weight of water. To reduce the
volume of water to one-half would require, not doubUng the
pressure as in the case of steam, but increasing it from 1 to 10,000
atmospheres. Evidently the molecules are already very closely
packed.
This close packing of the molecules causes cohesion to be much
more pronounced in liquids than in gases. A small amoimt
of Uquid, when poured into a vessel, does not fill the whole of
94
SMITH S INTERMEDIATE CHEMISTRY
the available space, but foxxDB coherent drops, the curvature
of the surface of which indicates that tremendous forces are
existent, pulling the outside molecules towards the interior of
the liquid. Nevertheless, molecules do escape outwards from a
liquid surface to form vapor (see p. 62^). We must therefore
conclude that the molecules of a hquid are still in rapid motion.
Similarly, when hquids which are capable of mixing (e.g,, alcohol
and water) are placed in separate layers in the same vessel, they
do mix, slowly, by diffvMon, The rate of dispersion of the mole-
cules, although much impeded by their close packing, has not
been annihilated.
A Uquid still possesses, therefore, in a modified degree, many
of the properties exhibited by a gas. For any particular substance,
the differences between the behavior of the hquid and the gaseous
states grow less and less as the temperature is raised, until at
the critical temperature (see p. 91) the two states become identical.
The Properties of Solids. — True solids are sharply distin-
guished from Uquids by their crystalline forms (Figs. 39-43, see
Fig. 39
Fig. 40
Fig. 41
Fig. 42
Fig. 43
Octahedron Square Pris- Rhombic
(Alum) matic (Niter)
Monosymmetric Asymmetric
(Gypsum) (Hydrated cupric
sulphate)
also pp. 4, 5 and 12), which possess definite planes of cleavage and
by their behavior towards Ught and X-rays give further evidence
of regular structure. Substances such as glass, which do not exhi-
bit any specific crystalline form, although commonly called solids,
are strictly speaking still in an extremely viscous Uquid state.
In such substances, as in hquids and gases, the cohesive forces
APPLICATION OF THE MOLECULAR HYPOTHESIS 95
between adjacent molecules are exercised equally in all directions,
with the result that the molecules are aU in haphazard, imordered
positions relative to one another and no particular arrangement
of particles in space can persist. Substances of this character
are therefore caUed amorphous (Greek, withoiU form) . In crystals,
on the other hand, since each substance shows an individual struc-
ture, the forces between adjacent particles must be exercised in
definite directions.
By using crystals of different substances as X-ray gratings
(see p. 548), W. H. and W. L. Bragg (1914) have been able to
show that crystals are composed of particles arranged in rows, the
spacing of these rows with respect to one another determining
the geometrical form of the crystal. They have also succeeded
in proving that the particles, which so arrange
themselves in definite patterns in crystals, are
not molecules, much less aggregates of molecules,
but atoms of the constituent elements of the sub-
stance. Thus when common salt (sodium chloride),
which crystallizes in cubes, is examined by reflect-
O «Na» • -CI.
ing X-rays from an appropriate plane, it is foimd fig. 44
to consist of alternate rows of sodium atoms and
chlorine atoms, the rows being arranged in space in such a
way as to build up a cubical structure. The framework of
a unit cube of sodiimi chloride is shown in Fig. 44. The actual
length of a side of this cube is approximately only one himdred
millionth of an inch! However, by imagining other cubes to
be packed all round this, with atoms of sodium and of chlorine
placed at their alternate comers, the reader will obtain for him-
self an idea of the ultimate structure of a crystal of common salt.
Most substances crystallize in less simple geometrical forms and
consequently possess a much more complex structure. The frame-
work of the carbon atoms in a diamond, for example, is indicated
in Fig. 45-
In a substance in the crystalline state, therefore, definite molec-
96
SMITH S INTEBMSDIATE CHEMISTRY
ular units no longer exist. Thus^ in sodium chloride, no sodium
atom can be said to be combined specifically with any one chlorine
atom. It is, instead, imprisoned by a number of chlorine atoms,
stationed at definite intervals around it, and among these chlorine
Fig. 45
atoms its combining forces must be regarded as impartially
dispersed. Very powerful forces, it will be evident, must be
called into play in order to constrain the separate atoms in a
crystal to retain their regular positions with respect to one another.
The nature of such forces will be discussed in a later chapter (see
p. 554-5).
The resistance offered by the atoms in a crystal to forcible
changes of position is shown by the very low compressibility of
matter in the crystalline state. Nevertheless, some degree of
motion of the particles must still persist, since many crystalline
substances show a measurable vapor pressure. Some vapors,
indeed (e.g., phosphorus, iodine), can be condensed directly
to crystals without passing through the intermediate Uquid state.
The results of X-ray work indicate that this motion of the atoms
of a crystal consists of rapid vibrations about a mean position.
Atoms vibrating violently at the surface are evidently exposed to
the risk of breaking away altogether, after which they combine
to form molecules of vapor. Similarly, molecules of vapor strik-
ing the crystal surface may stick thereto, their constituent atoms
APPLICATION OF THE MOLECULAR HYPOTHESIS 97
arranging themselves in a continuation of the crystal pattern.
Eqnihbrium relations between crystals and vapor resemble closely,
therefore, those between Uquid and vapor discussed in an earUer
chapter (p. 62-4), to which reference should here be made.
Exercises. — 1. State the change which takes place in the total
volume of the gases or vapors in each of the six actions mentioned,
p. 72.
2. If 1 Uter of oxygen at 0*^ and 760 nmi. weighs 1.429 g., what
is the molecular weight (pp. 74, 84)?
3. Calculate the molecular weight of a gas, 200 c.c. of which at
0° and 760 nmi. weigh 2.1 g. (p. 85).
4. Find the molecular weight of a gas, of which 250 c.c. at 18*^
and 752 nmi. weigh 2.5 g. (p. 85).
5. The foUowing are the weights of sulphur contained in the
molecular weights of several of its compounds: 32.06, 64.12,
96.18. What is the atomic weight of sulphur (pp. 74, 76)?
6. What information is contained in each of the formulae: CS2,
PCU, AI2O8 (p. 79)?
7. If two gases combine in the ratio 5 : 2 by volume, in what
relative numbers do their molecules interact (p. 82)?
8. The density of sulphur dioxide is 0.00286. What is its
molecular weight (p. 85)?
9. The density of a gas, air = 1, is 2.3. What is its molecular
weight (p. 85)?
10. An element combines with hydrogen in the proportion
10.03 : 1, and its specific heat is 0.2. What is its atomic weight
(p. 87)?
11. The molecular formula of a gas is CH4. What is the gram-
molecular weight? What volmne does this weight occupy?
What is the weight of 1 Uter of the gas?
CHAPTER IX
MAKING OF FORMULiE AND EQUATIONS
The formula (p. 78) is a condensed statement of the composi-
tion of a substance. Before we can make (i.e., calculate) the for-
mula for a substance, we must (1) measure the proportions by weight
of the constituent elements. Then, we must (2) express these propor-
tions in multiples of the known atomic weights of the elements.
Analysis and Synthesis. — In the case of water H2O we
saw (p. 58) how the weights of hydrogen and of oxygen required
to give a measured amount of water were determined. The
composition of the water was found out by putting the substance
together out of the elements. This method is called Sjrnthesis
(Greek, putting together).
In the case of mercuric oxide we can take a weighed amount
of the oxide, decompose it, and weigh the mercury formed. The
difference is the weight of the oxygen. This process, of decom-
posing a substance to learn its composition, is called analysiSi the
Greek word for decomposition.
One or other — sometimes both — of these plans can be used
with every compound.
Some of the results of such experiments have been given in the
earUer chapters. For example:
Tin (100) + Oxygen (26.9) -^ Tin oxide (p. 9).
Lead (100) + Oxygen (7.72) -^ Lead oxide (p. 10).
Iron (100) + Oxygen (43) -^ Ferric oxide (p. 10).
Zinc (2.04) + Sulphur (1) -^ Zinc sulphide.
Mercuric oxide (108) -► Mercury (100) + Oxygen (8) (pp. 15, 19).
98
MAKING OF FORMUK® AND EQUATIONS 99
Making Formuhe. — In the formnlse, these proportions are
to be replaced by multiples of the atomic weights by whole nimi-
bers. We therefore divide the quantity of each element by the
corresponding atomic weight. This operation gives us the factors
by which the atomic weights are to be multiplied. The atomic
weights we find in the table, where the values determined by the
most expert chemists are coUected.
For example, — In the case of tin oxide the proportion of tin to
100
oxygen is kjt^ • The atomic weights are 119 and 16, respectively.
100
100 ^ 119 = 0.84, and 26.89 -^ 16 = 1.68. The proportion ^^
, 119 X 0.84
now becomes jg^^-j^ .
Now this proportion — Uke all chemical proportions — must be
expressed in multiples of the atomic weight by whole numbers.
Hence, we next find the greatest common measure of the two
factors. It is 0.84. Indeed, in this simple instance, we can see
that the ratio of the factors is 1 : 2. Dividing above and below
by 0.84, we get ^^ .
Now, the symbols stand for the atomic weights. Substi-
tuting these symbols, the proportion becomes p. ^ • The
U X ^
formula is therefore Sn02.
Applying the same process to the lead oxide, we get
100 ^ 207.1 X 0.483 ^ 207.1 X 1 _ Pb X 1 p, ^
7.72 16 X 0.483 16 X 1 0 X 1 ' ^^
Treating the other data in the same manner, we find: ferric oxide
FejOa, zinc sulphide ZnS and mercuric oxide HgO.
If the composition of the substance is given in percentages,
the same process is used. Thus, the case of sodium sulphate
works out as follows:
100
smith's intermediatb cHEansrav
Elements
Percentages
At. wt.
Quotient -i-
Formula
SoHilim
32.43
22.55
45.02
23 X
32 X
16 X
1.41 0.705
0.705 0.705
2.814 0.705
Na X
S
0 X
2
Sulphur
Oxyeen
4
The formula is
therefore Na2S04.
1
1
Making Equations. — The condensed statements of chemical
changes which we have been using can now be stiff further sim-
plified by using the formulce in place of the names of the sub-
stances (p. 78). Thus
Sn + O2 -> SnOa
This is to be read: 119 parts (or 1 atomic weight) of tin, acting
chemically vrith 2 X 16 parts (or two atomic weights) of oxygen,
give 151 parts of stannic oxide. We may also read it thus: 1
atom of tin with 1 molecule of oxygen gives 1 molecule of stannic
oxide.
In making an equation there are four stages or steps:
1. Find out by experiment what the substances used and pro-
duced are.
2. Learn the molecular formula of each substance.
3. Set down the molecular formulae in the form of a skeleton
equation. Place the formulae of the initial substances on the left,
and those of the products on the right.
4. Adjust, or balance the equation.
For example:
(1) When hydrogen and oxygen combine, water is formed.
(2) The molecular formulae are H2, O2, and H2O.
(3) Skeleton equation : H2 + O2 — > H2O.
(4) In accordance with the law of conservation of mass, the
nimibers of atomic weights (or atoms) of each element must be
MAKING OF FORMULA AND EQUATIONS 101
the same after the action as before it. Now the skeleton equation
shows two atomic weights of oxygen before, and, thus far only one
after the action, whereas there ought to be two there also. With
O2 (2 X 16 parts) we have enough oxygen to give 2H2O, which
contains 2 X 16 parts of oxygen. But this will require us to take
2H2 to " balance " the equation. The final equation is, therefore:
Balanced Eqvxition: 2H2 + 02-^ 2H2O.
Observe, we could not adjust the difficulty by writing H2 + O
-^ H2O, because each substance must be represented by its moleo-
vlar formula, which stands for the weight of the substance in the
standard volume of 22.4 Uters, or one chemical unit weight, and
in the case of oxygen this is O2 (= 32 g.). Putting this in terms
of the hypothesis, each formula must represent 1 molecule, and
the molecules of oxygen contain 2 atoms. Hence we could not
divide the oxygen molecule. But we covM take more than one
molecule of hydrogen, so we took 2 molecules of this substance.
The coefficients in front of formulae multiply the whole formula.
2H2O is equivalent to 2(H20), or two whole molecules of water.
Balancing Equations. — Learning to balance equations
correctly comes only by practice. Take, again, the case of iron
rusting. The substances are iron (Fe), oxygen (O2) and ferric
oxide (FegOa"). The skeleton equation is
Fe + O2 -> FeaOa.
We are not permitted to alter these formulae themselves, but we
may put coefficients in front of any of them to make the number
of atomic weights alike on^both sides. A good rule is to pick out
the largest formula and reason hack from that Here, this is FeaOa.
To get oxygen atoms in threes y we must clearly take 3O2 (= 60).
That will give us 2Fe203. This, in turn, will require 4Fe:
Balanced: 4Fe + 3O2 -> 2Fe203.
102 (Smith's intermisdiate chemistry
Equntions for Actions Already Studied. — In the prepor-
ration of oxygen (p. 15) we used mercuric oxide and got mercury
and oxygen:
Skeleton: HgO-^ Hg + 0»-
Balanced: 2HgO -► 2Hg + O2.
Potassium chlorate has a composition shown by the formula
KClOs. It gives (p. 29) potassium chloride (KCl) and oxygen
Skeleton: KClOs-* KCl + 0»-
Balanced: 2KClOz -* 2KC1 + 3O2.
The variety of chemical change, where one substance gives two
(or more) substances, decomposition (p. 16), is readily recognized
in these equations.
SubstanceSi like the manganese dioxide (catalytic agent) used
here, and the water so often employed as a solvent, when they
undergo no chemical changei are omitted from the equation.
When the water takes part in the action, however, it must, of
course, be included. Thus sodimn peroxide (Na202) and water
(H2O) interact (p. 31) to give sodimn hydroxide and oxygen:
2Na202 + 2H2O -> 4NaOH + O2.
The Preparation of Hydrogen (p. 60) from sodium (Na) and
water gives sodium hydroxide (NaOH) and hydrogen (H2) :
2Na + 2H2O -* 2NaOH + H2.
When steam is passed over iron (p. 51), we get hydrogen and
magnetic oxide of iron (FezOi) :
3Fe + 4H2O ?± Fe804 + 4H2.
The Uberation of hydrogen by the action of zinc (Zn) upon
sulphuric acid (H2SO4), where the products (p. 52) are hydrogen
and zinc sulphate, is shown thus:
Zn + H1SO4 -> H2 + ZnS04.
MAKING OF FORMUKaS AND EQUATIONS 103
Again, iron and hydrochloric acid (HCl) give hydrogen and ferrous
chloride (FeCU) :
Fe + 2HC1 -> H2 + FeCl2.
In the last two equations the variety of chemical changes called
displacement (p. 51), where one .elementary substance displaces
another from a compound, is well illustrated.
The equation for the formation of water by union of hydrogen
and oxygen,
2H2 + O2 -> 2H2O,
has already (p. 59) been discussed. The reduction of an oxide,
such as magnetic oxide of iron or cupric oxide, by hydrogen (p. 57),
gives the metal and water:
Fe304 + 4H2 <=± 3Fe + 4H2O.
CuO + H2 -> Cu + H2O.
Upon examining these equations for reductions, we perceive that
they are illustrations of displacement also.
Reversible Actions (p. 69), Uke the decomposition of water by
heating, and the recombination of the elements on cooUng (p. 67),
are shown by using two arrows:
2H2O T± 2H2 + O2.
The equation may be read from either end. The decomposition
and formation of hydrates (p. 68) are also reversible actions. In
the case of zinc sulphate, the equation is
ZnS04 + 7H2O <i± ZnS04,7H20.
Reaction Formulae. — In the foregoing formula for the
hydrate of zinc sulphate, it will be seen that we do not add
together all the atoms of oxygen, and write ZnHi4S0ii. The
latter would show the composition of the substance correctly,
but it would show nothing more. Now chemists find it convenient,
frequently, to alter the formula so that it shall indicate also some
important chemical property or reaction of the substance. Hence
104 smith's intermediate chemistry
the fonnula ZnS04,7H20, which indicates at a glance the rela-
tionship of the substance to zinc sulphate (ZnSO*). The hydrate
is made from zinc sulphate by adding water, and is easily decom-
posed into these two substances again. The reaction formula
hints at this familiar reaction. Note, however, that the conmia
(,) does not indicate a mixture' oi the materials, such as ZnSOi
and H2O, but a single substance composed of both. The plus
(+) sign is used between the formulae of different, uncombined
substances in a mixture.
In accordance with this plan, washing soda, hydrate of sodium
carbonate (p. 68), is written Na2CO8,10H2O, and bluestone (p.
68), hydrate of cupric sulphate, CuS04,5H20.
Dissociation. — A decomposition which, like that in p. 67, is
reversible is called a dissociation. When heat is the agent pro-
ducing the change, it is sometimes called a thermal dissociation.
Not all decompositions are reversible. Thus potassimn chlorate
decomposes to give potassiimi chloride and oxygen, but these
products will not combine under any known conditions, directly,
to give potassium chlorate.
Molecular Formulae. — In this chapter, for the sake of
simpUcity, we have so far left in the background the fact that the
formula must represent a molecular weight of the substance, as
weU as its composition. The total weight, for which the symbols
in a formula stand, must be equal to the weight of the substance
occupying the gram-molecular volume. In other words, the
formvla-AJoeight must represent one cube-full (Fig. 36, p. 74) of the
substance. This is true, as we have seen, of the formulae H2O and
HCl (see p. 74).
In the cases of tin oxide (Sn02) and ferric oxide (Fe203), we have
substances which cannot be converted into vapor or dissolved, so
that their molecular weights are imknown. In sxich cases, we
use the simplest formula that will show the correct proportions.
MAKING OF FORMtJluffi AND EQUATIONS 106
Molecular FormuloB of Simple Substances. — With oxy-
gen (O2) and hydrogen (H2), however, the double fonnulae are
used, instead of the simpler 0 and H, because the weights 2 X 16
and 2 X 1.008 are the ones which fill the cube. The molecules
of all elements are not diatomic, however. Thus the cube-full of
mercury vapor weighs only 200.6, the same as the atomic weight,
and the correct molecular formula of the element is therefore
Hg. Similarly, the correct formulae are Na (sodimn), K (potas-
sium), and Zn (zinc). But the weight of 22.4 Uters requires us to
write CI2 for chlorine, N2 for nitrogen, P4 for phosphorus vapor
and S2 for sulphur vapor (above 1000°).
After all, there is nothing surprising in the fact that the mole-
cules even of elementary substances should, in some cases, contain
several atoms. All that it means, in the case of an element with
diatomic molecules, such as oxygen (O2 = 32), is that, when
oxygen combines with another element, each molecule of oxygen
will be divided between two molecules of the product if the latter
contain only 16 parts of oxygen each.
As an illustration, the imion of hydrogen and oxygen to form
steam (p. 61) may be considered.
Hydrogen Oxygen Steam
1000 1000 + 1000 -> 1000 1000
If each of the above rectangles represents a small volmne con-
tabling 1000 molecules of gas, then 2000 molecules of hydrogen
and 1000 molecules of oxygen give 2000 molecules of water vapor.
Since each molecule of water vapor must contain at least one
atom (see p. 86) of oxygen, at least 2000 atoms of oxygen were
required, and must have been furnished by the 1000 molecules
of oxygen. Each of these molecules must therefore have spUt
into at least two atoms. We have no reason, however, for sup-
posing that there are more than two atoms in the oxygen molecule.
Hence we accept the formula O2 as correct.
106 smith's intekmediate chemistry
Similar confirmation of the formulse H2 and CI2 will be foimd
on p. 148.
Molecular FormuloB of Compounds. — The need of atten-
tion to making our formulae molecular comes out also in the cases
of many compoimds. Thus, formaldehyde (a disinfectant) has
the composition CH2O, and its molecular weight is 30, so that
CH2O (= 12 + 2 + 16) is the correct formula. But acetic acid
(the sour substance in vinegar) has the same composition, CH2O,
only its molecular weight is 60, and the formula is therefore written
C2H402(= 24 + 4 + 32).
For gaseous and volatile substances the correct molecular for-
mulae are always used. Thus, for phosphorus pentoxide (p. 72),
P4O10 is preferred to P2O5 because the molecular weight of the
substance in the state of vapor is 284 and not 142.
The correct equation for reaction 5, p. 72, is therefore:
P4 + 5O2 -> P4O10.
Only by bearing in mind the true molecular formulae can we
include the volume and molecular proportions of the reacting sub-
stances in the vapor state, as well as their weight and atomic
proportions (p. 100), in our condensed statement of a reaction.
Often, however, when a reaction takes place wholly between
soUds and Uquids, we use for convenience the simplest possible
formulae throughout. Thus for the combination of iron and sul-
phur (p. 15) we write : Fe + S -^ FeS.
Warnings. — Always place the formulae of the products on the
right-hand side of the equation, and the formulae of the reacting
substances on the left.
Point the arrow in the direction of the reaction; that is, towards
the products.
Use the molecular formulae for elementary substances (O2, Hj,
N2, etc.). The molecular formulae, when they are known, are the
only ones given in the text. The symbols of the elements, as
MAKING OF FORMULAE AND EQUATIONS 107
given in the table on the rear cover, must be made into molecular
formulae before use in equations.
Exercises. — 1. Using the data given on p. 35, calculate the
formulae of sulphur dioxide, phosphorus pentoxide, and carbon
dioxide.
2. Using the results of 1, make and balance the equations for
the union of each of the three elements, sulphur, phosphorus, and
carbon, with oxygen.
3. Using the information on p. 36, make and balance the
equations for the interaction with oxygen of: (a) carbon disul-
phide (CS2); (6) zinc sulphide (zinc oxide is ZnO), and (c) wood
(assmning the formula of the latter to be that of cellulose, CeHioOs).
4. Make and balance the equations for the actions on water
of: (a) potassium (giving KOH), and (6) calcium (giving Ca02H2,
usually written Ca(0H)2).
5. Using the data in regard to the action of zinc on hydro-
chloric acid, given on p. 53, calculate the formula of zinc chloride
and make the equation.
6. Make equations for the action of magnesium and almninium
upon hydrochloric acid (giving MgCU and AICI3) .and upon sul-
phuric acid (giving MgS04 and Al2(S04)3).
7. Make a molecular equation for the decomposition of hydro-
chloric acid by electrolysis (p. 55).
8. Make equations for the formation of the hydrates of sodium
carbonate (Na2CO3,10H2O) and of cupric sulphate (CuS04,5H20),
by union of the anhydrous substances with water.
CHAPTER X
SOLUTION
The property that many substances have of dissolving in others
is a most interesting and valuable one. The value Ues chiefly in
the fact that some substances are easily soluble in a given Uquid
and others are, practically, not soluble in it at aU. These differ-
ences in solubiUty enable us to accompUsh, both in the laboratory
and in chemical industry, many things otherwise impossible.
Thus, we separated sulphur from iron (p. 13), by using carbon
disulphide (CS2) to dissolve the former. In the same way the
refining of silver (its separation from the lead in which it is con-
tained) is carried out on a large scale in actual practice by the use
of molten zinc as a solvent. We must first learn precisely what is
meant by a solution, and then we shall be ready to understand
the uses and properties of solvents and solutions.
Solution. — :We distinguish carefully between a solution and
a mere mixture, also between a solution and a compoimd.
A mechanical mixture, i§uch as that of iron and sulphur, can
never be perfect, as will be evident from what has been said in an
earher chapter (p. 84) regarding the size of molecules. However
finely we may powder up such a mixture, we cannot possibly
bring about a sufficiently intimate dispersion of the particles of
its components among one another to justify us in beUeving that
the whole mass has become homogeneous in its ultimate structure.
In any true solution, however, intermingUng of the particles of the
separate components down to molecular magnitudes has actually
been accompUshed. In a solution of salt in water, for example,
the dissolved substance is completely and permanently dissipated
throughout the liquid. However long the solution is allowed to
108
SOLUTION 109
stand, salt never settles out. Only by evaporating oflf all the
water can a complete separation be efifected.
Practically speaking, there is no limit to the amomit of dissipa-
tion which may thus be produced. Thus a single small crystal
of potassium permanganate, a conunon disinfectant which gives
a very deep purple solution in water, may be dissolved in a Uter
or even in a hxmdred Uters of water, and the purple tinge which it
imparts to the hquid will still be perfectly perceptible in every
portion of the solution. We may note here the distinguishing
characteristic of a solution as opposed to a compound, Com-
poimds contain definite proportions by weight (p. 19) and simple
atomic ratios (p. 86) of their constituent elements. The com-
position of a solviiorij on the other hand, can within certain limits
(see p. 110) be varied continuously.
Sometimes, when we shake up a finely-divided soUd with a
Uquid, the latter becomes dull, or cloudy, or muddy. The soUd
particles are here simply suspended in the Uquid, not dissolved,
and will eventuaUy settle out. Sand, shaken with water, settles at
once. Flour, mixed with water, settles more slowly. The parti-
cles of flom- can be readily separated from the water, however,
by filtration (p. 13), the flour remaining on the paper while the
water runs through. Such mixtures are called siispensions.
In exceptional cases, the subdivision of a suspended substance
in a Uquid, while not approaching molecular magnitudes, is so
minute as to make its retention by filter-paper impossible, or
even to prohibit it from settUng out in any reasonable time. So-
lutions of soap, starch and gelatine in water are of this nature.
Such suspensions are known as colloidal sicspensions. To the
unaided eye, they appear to be true solutions. Their main prop-
erties, however, are essentiaUy dififerent from those of true solu-
tions, as will be seen later (pp. 440-1).
Milk owes its cloudy, white appearance largely to droplets of
oily matter, which reflect much Ught from their surfaces. They
pass easily through filter paper. But when milk is aUowed to
110 smith's intermediate chemistry
stand they slowly rise to the top, being lighter than the water
in which they are not dissolved, but suspended. A mixture of two
liquids of this natiu*e is called an emulsion.
Chemists conmionly call the dissolved substance the solute and
the substance in which it is dissolved the solvent. In many cases,
however, (when we take two Uquids such as alcohol and water,
for example) the terms solvent and solute are interchangeable.
Gases, liquids, and soUds may all be solutes, and dissolve in
suitable gaseous, Uquid, or soUd solvents.
Solvents. — ; Water is by far the commonest and most useful
solvent. Very many inorganic substances dissolve in it easily.
The fact that many (Uke sulphur and sand) do not, enables us to
separate the components of a mixture containing a soluble and an
insoluble substance.
Many organic substances, such as fats, paraflin, petroleiun, tar,
rubber, cotton, paper, shellac, and so forth, do not dissolve to any
measiu-able extent in water. But fats dissolve readily in ether
(C4H10O), in carbon disulphide (CS2), in carbon tetrachloride
(CCI4), and in chloroform (CHCI3). For this reason these sub-
stances remove grease which has accidentally got into cloth.
Paraflin, petroleiun, and tar dissolve in gasoline (petrol), and in
benzene (CeHo). Cotton and pure paper (like filter paper) will
dissolve in strong sulphuric acid. Alcohol (CaHeO) dissolves
shellac (to make varnish).
Again, water dissolves Uttle carbon disulphide, chloroform,
carbon tetrachloride, gasoUne or benzene. But it dissolves alcohol
in any amoimt, and ether in limited quantity. Some organic
substances, like sugar, dissolve easily in water, but hardly at all
in the other solvents just mentioned. Hence candy or molasses
can be taken out of cloth by water, but not by solvents for
fats.
Saturation. — As a rule, not more than a certain amount
of a solute is dissolved by a given quantity of the solvent. By
SOLUTION 111
shaking the solute with the solvent for a sufficient length of time,
this maximum amoimt will finally be dissolved. The solvent is
then said to be saturated by the solute in question. Thus, 100 c.c.
of water at 18** will dissolve as much as 6.6 g. of potassium chlorate,
but not more. The same amount of water will dissolve 213.4 g.
of silver nitrate, however, before the solvent becomes satu-
rated. On the other hand, a satiu'ated solution of chalk
(calcium carbonate) in water will contain only 0.00013 g. in
100 c.c.
To describe these cases we should say that potassiiun chlorate
is only moderately soluble in water,' ^silver nitrate very soluble, and
chalk insoluble. But no substance is absolutely insoluble.
The number of grams of the solute required to saturate 100 c.c.
of the solvent we call the solubility of the substance (at the exist-
ing temperatiu*e). The solubilities at 18° of one hundred and
forty-two substances in water are given in a table printed inside
fte covery at the front of this book. A few additional examples
are given below (p. 113).
In some cases there is no limit to the solubiUty, and therefore no
possibility of the solution reaching satiu'ation. Thus alcohol or
glycerine and water will dissolve in one another in any proportion.
Such pairs of substances are said to be miscible in all proportions.
Dilute and Concentrated Solutions. — A dilute solution
is one containing Uttle dissolved matter, whether the matter is
naturally very soluble or not. A concentrated solution is one
containing much of the dissolved substance, and such a solution
can be made with very soluble solutes only.
Conditions affecting the Solubility of a Gas. — When the
dissolving substance is a gas, led through, or confined above the
liquid at a definite pressure, the gas dissolves until a state of
equilibrium between dissolving and emission is reached, for
example, Oxygen (gas) t=^ Oxygen (dissolved), and the liquid is
then saturated with the gas.
112 smith's intermediate chemistry
It is found, as the molecular theory would lead us to expect,
that the concentration of the saturated solution of a gas is pro-
portional to the pressure at which the gas is supplied (Henry's
law).
This equilibrium, Gas (gaseous) ^ Gas (dissolved), can be
reached, naturally, from the other direction, namely by starting
with a solution of the gas and a space above the solution contain-
ing, at first, none of the gas. The gas leaves the solution imtil the
rates of emission and return become equal. Hence, a gas may be
entirely removed from solution by bubbUng a foreign gas through
the Uquid. The bubbles furnish the space to receive the emitted
gas, and have a large surface, so that the process goes on rapidly.
The bubbles also escape, and carry with them the emitted gas,
so that, in this case, there is no re-solution. This is a case of
nuUifying one of the two opposed tendencies (p. 64).
When a mixture of two gases is shaken with a Uquid, the gases
behave independently of each other (Dalton's law, p. 47). Each
has the same pressure, and therefore the same solubiUty, as it
would possess if it alone occupied the whole space above the
Uquid.
Two Immiscible Solvents: Law of Partition. — An inter-
esting appUcation of the same ideas may be made to a case which
occurs very conmaonly in chemical work. If we shake up a small
particle of iodine with water, we find that it dissolves slowly,
giving eventuaUy a saturated but very dilute solution. K now
ether in sufficient quantity be shaken with the aqueous solution,
the greater part of the iodine wiU find its way into the ether, and
be contained in the brown layer which rises to the top. The proc-
ess of removing a substance practic^Uy from solution in one sol-
vent and securing it in another is caUed extraction. We find in
such cases that neither solvent can entirely deprive the other of
the whole of the dissolved substance, if the latter is soluble in
both independently. A state of equiUbrium is finaUy reached:
SOLUTION 113
I2 (in water) ^ I2 (in ether). The partition of the substance
takes place in proportion to its solubility in each solvent. It
is found that any amount of the solute, up to the maximum the
system cai^ contain, provided this does not involve too high a
concentration in either solvent, is divided so that the ratio of the
concentration in the two solvents is always the same. In the case
of iodine divided between water and ether, this ratio is about
1:200.
This principle is used in Parke's process for extracting silver
from molten lead, by means of melted zinc as the second solvent.
It is employed in separating interesting compounds from animal
secretions and vegetable extracts, and in purifying such compounds.
Nicotine from tobacco and cocaine from coca leaves are secured
in this way.
Temperature and Solubility. — The solubility of every
substance in any solvent varies more or less with the temperatiu'e.
The solubiUty of niter (potassimn nitrate KNO3) in water shows
great variation, namely 13 g. in 100 c.c. at 0°, 26 g. at 20°, 140 g.
at 70*^. On the other hand, the solubility of conmion salt (sodium
chloride NaCl) is nearly constant,^35.5 g. at 0°, 36.5 g. at 20°, 38 g.
at 70°, 40 g. at 100°.
Usually, as in these two cases, the solubility of soUds in Uquids
(and of Uqmds in Uquids) increases with rise in temperatiu-e, but
in a few cases it diminishes. Thus, the solubihty in water of
slaked lime (calcium hydroxide Ca(0H)2, used to make lime
water) is 0.175 g. at 20° and 0.079 g. at 100°, and that of an-
hydrous sodimn sulphate (Na2S04) is 55 g. at 32.5° and 42 g. at
100°.
The solubiUty of gases in Uquids diminishes with rising tem-
peratitfe. This may be iUustrated by heating cold tap-water in a
beaker. The dissolved gases, originaUy obtained from the air,
appear in bubbles on the bottom and sides as the temperature
rises.
114 smith's intermediate chemistry
Crystallizalion. — If the solvent has been saturated while
warm, and the substance is one that is less soluble at lower tem-
peratiu-es, then, when the temperature falls, the solute begins to
come out of solution. The amoimt appearing, of course, is only
the excess beyond what is needed to saturate the solvent at the
lower temperatiu-e.
If the solute is Uquid at the new temperature, it appears at
first as a cloud of drops, rendering the Uquid milky. This may be
shown by cooling a solution of phenol (carbohc acid) in hot water.
If the solute is soUd, then the particles, as they appear, take
the form of ciystals (p. 94). These grow by taking on more of
the separating solute. If the cooling goes on slowly, very large
crystals can finally be obtained. On the other hand, with rapid
cooling, new crystals are continually formed, and a fine crystal-
meal falls to the bottom of the solution. The crystals in this
meal, however, when viewed through a lens, are seen to be just
as perfect as the larger ones.
When a more dilute solution is used, instead of a saturated one,
crystals may still be obtained. A part of the solvent must first
be removedj however. This may be done, either by boiling the
solution for some time, or by leaving it to evaporate in a wide
dish in which a large surface is exposed.
When the dissolved substance can form a compound with the
solvent (e.g., a hydrate, see p. 67) which is stable at the tem-
perature of crystaUization, the crystals are composed of this com-
poimd.
The whole of the solvent may be boiled off. But in this case,
good crystals of the solute are never obtained — the residue is
usually a crust composed of imperfect crystals.
When the substance is more soluble in cold than in hot water,
then crystallization is produced by raising the temperature.
Crystallization from a Melted Mass. — In this connection,
it should be noted t^t there is another way of obtaining crystals.
i
SOLUTION
115
This is to meU the substance (without any solvent), and allow the
mass to cool slowly. When a part has soUdified, the rest of the
liquid is rapidly poiu-ed oflf. Metals and many other fusible sub-
stances give good crystals in this way. Water itself, when it
freezes, deposits radiating, hexagonal crystals of ice.
Superaaturation. — When a hot, saturated solution is cooled,
there is quite conmionly some delay before the crystals begin to
appear. The solution, pending the appearance of the crystals, is
then said to be supersaturated. In most cases the crystals soon
appear in due coiu^e, especially if the liquid is shaken or stirred.
But certain substances have a tendency to remain indefinitely in
the state of supersaturated solution. The hydrates of sodium
sulphate (Na2SO4,10H2O) and of sodiiun thiosulphate (photog-
raphers' " hypo '' Na2S208,5H20) give solutions in water of this
nature. The addition of a minute crystal of the
substance concerned (" inoculation ''), however,
always starts the crystaUization (Fig. 46).
Many pure Uquids, similarly, when cooled below
their freezing-point, do not always crystallize out
at once. Thus water can be taken down to — 10°
without the appearance of ice. In this condition
it is said to be supercooled. Shaking, or stirring,
or (better still) inoculating with a fragment of
ice, induces crystallization in this case also. It
may be noted that the opposite phenomenon has
never been observed; ice invariably melts sharply
at 0® under atmospheric pressure.
Fig. 46
Heat of Solution. — Most substances absorb heat when they
dissolve, making the solution cooler, and give out heat when they
cr3rstallize. Thus, in the cases of the two sodiiun salts last
mentioned, when the crystalhzation is brought about in the cool,
supersatmrated solutions, the rise in temperatiu-e is considerable.
116 smith's intermediate chemistry
This fact has been utiUzed in devising a sort of hot-water
bottle. The bottle is made of rubber and contains a super-
saturated solution of sodium acetate. Whenever the heat is
wanted, the stopper is taken out, rubbed with the finger, and
screwed back. The rubbing spreads on the inner surface of the
stopper, next the Uquid, some of the crystals adhering to the screw,
and so starts the crystaUization. The bottle then becomes warm
and remains so for a considerable time. After it has cooled, it is
placed, without being opened, in boiling water to redissolve the
crystals, and, when cold, is ready for use again.
Influence of the Solute upon the Solvent. — The dissolving
of a substance alters the properties of the solvent. The observed
changes may be divided into two classes.
In the first class, the amount of the change varies with the
substance dissolved. Very striking and difficult to explain, for
example, are the erratic changes in volume which occur when
solution takes place. Specific effects of this class show that
chemical changes often accompany solution. For example, 58.5
g. of sodium chloride (volume 27.5 c.c.) and 10,000 c.c. of water
have a volume, totalUng 10,027.5 c.c, but, when they are dis-
solved, the solution measures only 10,016.5 c.c. This is a very
dilute solution (about ^ per cent), so that the contraction of 11 c.c.
is relatively considerable. On the other hand, 214 g. of ammonium
chloride (volume 142.5 c.c.) and 843.5 c.c. of water have a total
volume of 986 c.c, but when dissolved give 1000 c.c of solution.
Here there is an expansion of 14 c.c In the case of table-sugar
and water, however, there is almost no change in volume.
Another important property of solutions in which the influence
of the solute is specific is conduction of electricity. Pure water is an
exceedingly poor conductor. A solution of table-sugar in water
is also practically non-conducting. But when acetic acid is
dissolved in water a solution is obtained which conducts the
current fairly well, while a solution of sodium chloride is an ex-
SOLUTION 117
ceedingly good conductor. The significance of these differences
in behavior will be taken up later (p. 118).
In the case of many properties of solutions, however, it has
been found that eqiuil numbers of dissolved molecules of different
substances proditce the same amount of change. The effect appears
here to be due essentially to physical causes, and is discussed in
the following sections in the Ught of the molecular hypothesis.
Before attacking these sections, the student is reconunended to
refer back to p. 62-4 and read these pages through again care-
fully, noting that the equiUbrium relationships between Uquid
water and water vapor, therein discussed, can obviously be ex-
tended to any volatile substance in contact with its own vapor.
Vapor Pressure of Solutions. — When we take equal quan-
tities of a volatile Uquid {e,g,, benzene, CeHe) and add to each
equal weights of different non-volatile solutes {e,g,, naphthalene,
anthracene, camphor; three organic soUds which are practically
non-volatile at ordinary temperatiu-es) we find that the vapor
pressures of all the resulting solutions are less than that of the
pure solvent, but the depression is different in each case. But if,
instead of adding equal weights of the different solutes, we add
equal numbers of molecules (as we can do by dissolving, for example,
1 g. molecular weight of each substance in 1000 g. of benzene),
we find that the depression is the same in every case. The depres-
sion is proportional^ moreover, to the fraction of soluie molecules
in the solution. This very striking fact is explained by the mo-
lecular hypothesis as follows.
Every molecule at the surface of a pure volatile liquid has an
equal chance to escape into the vapor above the Uquid. But
as soon as we add to such a Uquid a solute which is practicaUy
non-volatile, we have a Uquid in which some of the molecules
have no tendency to pass into the state of vapor, but are fixed
in the Uquid state. Suppose, for instance, we consider a solution
in which one molecule in every ten is non- volatile; the intensity
118 smith's INTERBfECIATE CHEMISTRY
of the hail of molecules leaving the liquid will evidently be reduced
by one-tenth. EquiUbriiun between liquid and vapor over such
a solution will be re-established only when the intensity of the
hail of vapor molecules returning to the Uquid is also reduced by
one-tenth, since otherwise more molecules will be returning than
leaving. This means that the vapor pressure of the solution must
be one-tenth less than that of the pure solvent.
It is important to note that the naiure of the solute is here im-
material, the essential factor is the number of molecules it furnishes
to the solution. We have here a method of determining the
molecular weights of nonvolatile substances. By dissolving a known
weight of such a substance in a known weight of a suitable sol-
vent and determining the relative depression of vapor pressure
thereby produced, we learn what fraction of the molecules in the
solution belong to the solute, and hence can calculate its molecular
weight.
AU aqueous solutions show a lower tension of water vapor than
does pure water. With conducting solutes (e.gf., sodium chloride),
indeed, the vapor pressure depressions obtained are abnormally
large, and do not agree with the accepted molecular weights.
This is a point to which we shall return later (p. 177).
If a substance is very soluble in water, the solution may give
a vapor pressure of water less even than that commonly present
in the atmosphere. Such a solution, placed in an open vessel,
will not evaporate. On the contrary, it will take up moisture
from the air and increase in bulk. For this reason very soluble
substances are commonly moist and, when exposed to the air,
extract water from the latter and dissolve in this water. This
behavior is called deliquescence^ and is shown, for example, by
the hydrate of calcium chloride CaCl2,6H20, used to dry gases
(p. 59).
Boiling'Points of SoluUona. — The boiling-point of a liquid
is that temperature at which the vapor pressure reaches 760 mm.
SOLUTION 119
(see p. 62). Since the addition of a non*volatile solute lowers the
vapor pressure of a pure Uquid^ it naturaDy raises the boUing-
paint to a higher temperature.
In dilute non-conducting solutions, equal numbers of mole-
cules of different solutes raise the boiling-point of a given sol-
vent to the same extent. Thus, one molecular weight of sugar
(Ci2H»0ii = 342 g.) or of glycerine (CsHgOs = 92 g.), dissolved
in 1000 g. of water, will each raise the boiling-point from lOO*'
to 100.52^.
Molecular weights of nonrvokuile, rumrconducting substances can
therefore be determined by finding out what weight of the sub-
stance, when dissolved in 1000 g., is required to raise the boiling-
point of water from 100*" to 100.52*".
Freessing'Points of Solutions. — The addition of a solute
similarly tends to prevent the freezing of the solution, for freezing
means the separation of a part of the pure solvent in the form of
ice. Hence solutions can be frozen only at temperatures below
those of the piu-e solvents. Thus, one molecular weight of a
substance, such as sugar (342 g.) or glycerine (92 g.), dissolved
in 1000 g. of water, will cause the water to freeze at —1.86° in-
stead of 0°. Molecular weights can be measiu*ed by this method
also.
This behavior explains why sea water is frozen in cold weather
much less often than fresh water.
It explains also why salt thraum on ice will cauM the latter to melt.
Saturated salt solution freezes only at —21°, to give a mixture
of pure ice and pure salt, both in soUd form. Hence, ice and
salt can not permanently exist together above that temperature.
When the outside temperature is below —21°, salt will no longer
melt the ice. But calcium chloride, which is more soluble, will
do so. A mixture of ice and salt, giving the temperature —21°,
is called a freezing mixture. Such a mixture is used in freezing
ioe cream and ices.
120 smfth's intebmediate chemistry
Definition of a Saturated Solution: A Warning. — To
avoid a common misconception, it must be noted that solution
knot a, process of filling the pores of the liquid. If that were true,
approximately equal weights of all substances would find accommo-
dation in equal volmnes of water. The fact is that, for example,
100 c.c. of water can dissolve 195 g. of silver fluoride, but only
0.00000035 g. of silver iodide, although the space available (if
there is any such space) is the same in both cases.
The same conclusion is reached when we consider that two forms
of the same salt may have different solubiUties. Thus, at 20°,
Na2SO4,10H2O can give about 18 g. of Na2S04 to 100 c.c. of water.
But anhydrous sodimn sulphate Na2S04 at 20° gives 59 g. to the
same amount of water.
The reader is also warned against the frequent definition of a
saturated solution as one containing all of the solute that it can
hold. A supersaturated solution evidently holds more. The
satiu'ated solution imder any given conditions is that solution
which, when placed in contact with excess of the solute, is found to
he in equilibrium.
The molecular hypothesis may again be called to our assistance
in this connection. When we have a solute (either crystalUne, or
liquid or gaseous) in contact with its saturated solution, and there-
fore in equiUbriimi with it, two opposing tendencies must balance
each other at the surface of contact. One of these is the tendency
of the solute particles to escape into solution, the other is the tend-
ency of the solute particles already in solution to return back to
the solute. The first of these tendencies (the intensity of the
hail of particles thrown oflf from the surface of the solute into a
given solvent, under given conditions of temperatiu-e and pressure)
we may regard as constant. The second tendency (the intensity
of the hail of particles returning from the solution to the surface
of the solute) will increase steadily as the concentration of the
solute particles in the solution increases. At one definite con-
centration only, therefore, can these two opposing tendencies
SOLUTION 121
counterbalance, namely that of the saturated soliUion. The rate
at which solute particles are returning from the solution just
equals, in this case, the rate at which they are entering. With
unsaturated solutions, containing less solute, the number of re-
turning particles will be deficient, and the solute will continue to
dissolve until it all disappears or saturation is reached. With
supersaturated solutions, on the other hand, containing mare
solute, the niunber returning will be in excess, and deposition of
home-coming particles on the solute surface will continue until
this excess is wiped out.
If, however, we have a supersaturated solution in which no
free solute is present, the solute particles in the solution have no
home to return to, no surface upon which they can deposit them-
selves. They are therefore compelled to continue wandering
around and aroimd in the solution, having lost their equiUbrium
completely. By violent shaking or stirring we may succeed in
inducing crystalhzation in such a solution, but the only certain
means of estabUshing equiUbrium conditions is inoculation with
a small fragment of the solute.
Units Used in Expressing Concentrations. — The concen-
trations of solutions, satiu'ated and otherwise, are sometimes
expressed in physical, and sometimes in chemical, units of weight.
When physical units are employed, we give the nmnber of grams
of the solute held in solution by one hundred grams of the solvent.
When chemical units of weight are employed, two different plans
are possible, and both are in use. Either the equivalent (p. 53) or
the atomic weights may be taken as a basis of measurement. In
the former case, the solutions are called normal solutions, and in
the latter, molar solutions.
A normal solution contains one gram-equivalent of the solute
in one liter of solution (not in 1 1. of solvent). The word " equiva-
lent '' has been used hitherto only of elements, and this appUcation
of the expression involves an extension of its meaning. An
122 smith's intermediate chemistry
equivalent weight of a compound is that amount of it which will
interact with one equivalent of an element Thus, a fonnula-
weight of hydrochloric acid HCl (36.5 g.) is also an equivalent
weight, for it contains 1 g. of hydrogen, and this amount of hydro-
gen is displaceable by one equivalent weight of a metal. A for-
mula-weight of sulphuric acid H2SO4 (98 g.), however, contains
two equivalents of the compound, and a formula-weight of alu-
minium chloride AlCls (133.5 g.) three equivalents. Hence
normal solutions of these three substances contain, respectively,
36.5 g. HCl, 49 g. H2SO4, and 44.5 g. AICI3 per Uter of solution.
The special property of normal solutions is, obviously, that equal
volmnes of two of them contain the exact proportions of the solutes
which are required for complete interaction. Solutions of this
kind are much used in quantitative analysis. We frequently use
also decinormal or one-tenth normal solutions (0.1 N or N/10), and
seminormal (0.5 N or N/2), and six times normal solutions (6 N),
and so forth.
A molar solution contains one mole (gram-molecular weight)
of the solute in one liter of solution (not in 1 1. of solvent). When
molecular formulae (p. 78) are used, this means one gram-formula
weight per Uter. In the cases cited above, the molar solution
contains 36.5 g. HCl, 98 g. H2SO4, and 133.5 g. AICI3 per Uter.
Is Dissolving a Physical or a Chemical Change? — This
is a question still much discussed amongst chemists. Prob-
ably in simple, typical cases, Uke dissolving paraflin in gasoUne
or benzene, the process may be considered piu-ely physical, and
the solution contains both components in imchanged chemical
condition.
On the other hand, when water is used, as it is more frequently
than any other solvent, chemieal changes undoubtedly take place.
The water itself, at least, is always changed. Water in the li-
quid stcUe is not shnply H2O. Its physical properties indicate that
it is an associated Uquid, extensive combination having taken
SOLUTION 123
place between simple H2O molecules to fonn more complex
molecules with the general fonnula (H20)n. Dissolving any sub-
stance in water must upset the equilibrium amongst these different
kinds of molecules:
(H20)n ^ nHgO
and produce more of one and less of the other kind. This is the
extent of the chemical change in the water.
The dissolved substance probably combines also, in the ma-
jority of cases, with part of the water. The nature of the com-
pounds is hard to determine, and no simple statement can as yet
be made about them. But the compounds, whatever they are,
are physically dissolved in the rest of the water.
Dissolving, therefore, is partly a chemical, and only partly
a pure physical process. The striking differences in solubility
already mentioned (p. 110) may consequently be accounted for
partly on a chemical, and partly on a physical basis. The main
chemical factor is compound formation between the components
of the solution. The more extensive this is, the greater, in general,
is the solubiUty. Thus substances which form definite hydrates
with water are mostly extremely soluble, while substances which
are only difficultly soluble in water invariably crystalUze out
from an aqueous solution in an anhydrous state. The main
physical factor is the relative magnitude of the cohesive forces
between the various types of molecules present in the solution.
Thus in the case of water and benzene, the water molecules at-
tract one another much more strongly than they do the benzene
molecules. Molecules of benzene endeavoring to intermingle
with water molecules encounter, therefore, very considerable
resistance, and are almost certain to be squeezed out. The two
liquids^ indeed, are found to be practically inmiiscible.
Exercises. — 1. Give two ways of separating a mixture, con-
sisting of a suspended solid and a liquid (p. 109).
2* If you had a spot on your clothing consisting of: (a) grease,
124 smith's intbrmbdlajte chemistry
(b) sugar, or (c) sugar and grease together, or (d) varnish, how
should you proceed in each case to remove the spot?
3. If chalk (5 g.) and potassium chlorate (5 g.) were mixed, how
should you separate them (p. 13)? Explain how you could secure
each substance.
4. Could you make (a) a concentrated, (b) a saturated solution
of chalk in water (p. 111)? Of alcohol in water?
5. If you saturated 200 c.c. of water at 70° with (a) salt, or
(b) potassium nitrate, and then cooled the clear liquid to 20°,
what weight of the soUd substance would separate out in each case
(p. 113)?
6. To make as concentrated a solution of lime water as possible,
should you use hot water or cold (p. 113)?
7. Explain why boiled water has a sUghtly different taste from
tap-water that has not been boiled (p. 66).
8. If 100 g. of a non-volatile substance, dissolved in 1000 g.
of benzene (CeHe), lower the vapor pressure from 74.8 to 68.0
mm., what is the molecular weight of the substance (pp. 117-8)?
9. Explain why potassium carbonate becomes wet, and finally
dissolves, when exposed to moist air. How must calcium chloride
be preserved from becoming moist?
10. If 52 g. of a substance dissolved in 1000 g. of water gives a
solution boiling at 100.26°, what is the molecular weight of the
substance (p. 119)?
11. If 68.5 g. of a substance, dissolved in 500 g. of water
gives a solution freezing at —1.86°, what is the molecular weight
of the substance (p. 119)?
12. How much glycerine (CsHsOa) could you dissolve in 1000 g.
of water, and still be able to freeze the water with ice and salt
(p. 119)?
13. Explain why a sodimn acetate hot-water bottle can be used
over and over again. What is the source of the heat it gives out
each time it is used?
y
CHAPTER XI
HYDROCHLORIC ACID- CALCULATIONS
Thus far, the substances we have studied have been mainly
air and its components and water and its constituents. Another
of the simpler, familiar substances, common salt or sodium chloride
(NaCl) may now be taken up. Large amounts of it are used in
the household, in cooking and in making freezing mixtures. StiU
larger quantities are consmned in manufacturing washing soda
and soap, for both of which it supplies the necessary sodium. It
is employed also to fimiish the chlorine for bleaching materials.
We shall consider it first as a means of making compounds of
chlorine.
Preparation of Hydrogen Chloride. — When a few drops
of commercial, concentrated sulphmic acid (H2SO4) are poured
upon common salt in an open dish, vigorous effervescence begins.
This indicates that a gas is forming bubbles upon the salt and
that the bubbles are rising through the layer of acid and burst-
ing. The gas is itself invisible, but when we breathe upon the
contents of the vessel, a heavy fog is produced. This is due to
condensation of water vapor (in the breath) to droplets of water,
in which the gas has dissolved. The fog is composed, in fact,
of drops of a solution of hydrogen chloride (HCl) in water, which
receives the name of hydrochloric acid (in conmierce, muriatic
acid).
In order to handle the gas more readily, the sulphuric acid
may be allowed to fall from a fimnel, drop by drop, upon salt con-
tained in a flask (Fig. 47). Soon the air in the flask is all dis-
placed by the gas, and the latter issues from the open delivery
tube. If a U-tube containing some water is attached to the
125
126
smith's intermediate chemistry
delivery tube, the gas dissolves in the water as fast as it is
formed.
If the correct proportions of the materials are used, then, when
the action is over, all that remains in the flask is a white soUd,
diflFerent from salt, and called sodium-hydrogen sulphate NaHS04.
A part of this may be in solution in a
Uttle water, contained originally in the
commercial sulphuric acid, of which
water commonly forms from 6 to 7 per
cent. The equation is easy to make
from the formulse given,
NaCH-H2S04?i^HCl t +NaHS04, (1)
and requires no further balancing. (An
arrow pointing upward is used in
equations to indicate that the sub-
stance to which it refers removes
itself from the reaction by escaping in the form of a gas.)
Other Sources of Hydrogen Chloride. — Chlorides of other
metals could be substituted for the sodium chloride in this action^
and all but the less soluble ones would give hydrogen chloride
freely. Common salt is employed because it is the cheapest of
the chlorides.
While theoretically any acid would, Uke sulphuric acid, furnish
the required hydrogen, and Uberate hydrogen chloride, yet in
practice no other acid works so well. Some, Uke phosphoric
acid H3PO4, act too slowly, because they do not dissolve sodium
chloride so readily. Others, Uke hydrofluoric acid HF, are too
volatile, and the heat of the action would send them over with
the hydrogen chloride in the form of vapor. Others, Uke nitric
acid HNO3, would react chemically with hydrogen chloride. Still
others, Uke hydriodic acid HI, being gases, could be used only in
aqueous solution, and the water would dissolve the hydrogen
chloride produced, and prevent its escape from the vessel. Aside
HYDROCHLORig ACID. CALCULATIONS 127
from these objections, all the other acids are more expensive than
sulphuric acid.
The Molecular View of the Interaction of Sulphuric Acid
and Salt. — One who has used the above-described methods for
making hydrogen chloride without reflection would not reaUze
the complexity of the machinery by which the result is achieved.
The means are apparently very simple. Yet the mechanical
features of this experiment, when laid bare, are extremely curious
and interesting. A single fact will show the possibiUties which
are concealed in it.
If we take a saturated solution of sodium-hydrogen sulphate in
water and add to it a concentrated solution of hydrogen chloride in
water (concentrated hydrochloric acid), we shall perceive at once
the formation of a copious precipitate. This is composed entirely
of minute cubes of sodium chloride:
NaHS04 + HCl -^ H2SO4 + NaCl l (2)
(An arrow pointing downwards is used in equations to indicate
that the substance to which it refers removes itself from the
reaction in the form of a precipitate.) Now this action is nothing
less than the precise reverse of (1), yet it proceeds with equal
success. In fact, this chemical interaction is not only reversible
(p. 103), but can be carried virtually to completion in either
direction. It is only in presence of a large amoimt of water,
sufficient to keep both the hydrogen chloride and the salt all in
solution, that it stops midway in its career and is valueless for
securing a complete transformation in either direction:
NaHS04 + HCl :^ H2SO4 + NaCl.
In an action which is reversible, if the products remain as per-
fectly mixed and accessible to each other as were the initial sub-
stances, their interaction will continually undo a part of the work
of the forward direction of the change. Hence, in such a case the
128 smith's intermediate chemistry
reaction must, and does, come to a standstill while as yet only
partly accomplished; but this was not the case with actions
(1) and (2). Let us examine the means by which the premature
cessation of each was avoided.
In equation (1) the salt dissolved to some extent in the sulphuric
acid, NaCl (soUd) ^^^ NaCl (dslvd.), and so, by intimate contact
of the two kinds of molecules in the resulting solution, the prod-
ucts HCl and NaHS04 were formed. On the other hand, the
hydrogen chloride, being practically insoluble in sulphuric acid,
escaped as fast as it was formed: HCl (dslvd.) 4=^ HCl (gas).
Hence, in that case, almost no reverse action was possible, and the
double decomposition went on virtually to completion. With
all the sodium-hydrogen sulphate in the bottom of the flask, and
most of the hydrogen chloride in the space above, the two products
might as well have been in separate vessels so far as any eflicient
re-interaction was concerned. This plan, in which water is pur-
posely excluded, forms therefore the method of making hydrogen
chloride.
In equation (2), on the other hand, the hydrogen chloride was
taken in aqueoits solviion, and was mixed with a concentrated
solution of sodium-hydrogen sulphate. The acid was, therefore,
kept permanently in full contact with the sodJtim-hydrogen sul-
phate. It had in this case, every opportunity to interact with
the latter and no chance of escape. Every molecule of each
ingredient could reach every molecule of the other with equal
ease. Furthermore, the sodium chloride, produced as a result
of their activity, is not very soluble in concentrated hydrochloric
acid (far less so than in water), and so it came out as a precipi-
tate: NaCl (dslvd.) ?=^ NaCl (solid). But this was almost the
same as if it had gone off as a gas. It meant that the greater part
of the salt was in the soUd form. In this form, it was no longer
able to interact effectively molecule to molecule with the other
product, the sulphuric acid. Hence, there was Uttle reverse
action to impede the progress of the primary one. Thus (2) is
HYDBOCHLORIG ACID. CALCULATIONS 129
nearly as perfect a way of liberating sulphuric acid as (1) is of
Hberating hydrogen chloride.
Precipitation. — When two soluble substances are dissolved
separately in water, and the solutions are mixed, chemical inter-
action frequently is evident between the dissolved materials. If
one of the products is not very soluble, then a supersaturated
solution (p. 115) of this product may he thus produced. As a rule,
this substance inamediately becomes visible as a fine powder,
called a predpitatCi suspended in the Uquid. More oi; less rap-
idly, according to its fineness of dispersion, this precipitate settles
out, leaving the solution clear. Equation (2) in the preceding
paragraph is an example of such a reaction.
Often the precipitated product can be recognized by the physi-
cal appearance of the precipitate, and so this sort of action is used
as a test for one of the original substances. Thus, precipitates
are classified according to their color. Again, precipitates of the
same color differ ux degree of dispersion, and may be described
as gelatinouSi curdyi pulverulenti or crystalline. In the two
former cases, at least, the precipitation is so sudden that there is
not time for crystals to be formed, and the product is amorphous
(see p. 95).
Physical Properties of Hydrogen Chloride. — Hydrogen
chloride is a colorless gas. It is sour in taste, and has a sharp odor.
It is irritating, but not poisonous in smaU amounts.
The gas is exceedingly soluble in water, one volume of which,
at 15*^, will dissolve no less than 455 volumes of the gas. The
saturated solution at 15° contains nearly 43 per cent of the gas
by weight. The concentrated hydrochloric acid of commerce
contains about 35 per cent.
The density of the gas (weight of 1 c.c.) is 0.001628. Of more
interest to the chemist is the weight of 22,400 c.c. or 22.4 liters
(the gram-molecular volume), namely 36.468 grams. This is the
130 SMITHES INTERMEDIATE CHEMISTRY
molecular weight of the substance. As we have seen (p. 74),
it is made up of 1.008 g. of hydrogen combined with 35.46 g. of
chlorine.
Is the gas heavier or Ughter than air? This question is an-
swered at once if we recall the fact that the 22.4-Uter cube-full of
air weighs 28.95 g. (p. 85). The gas is one-fourth heavier. It
may therefore be collected by upward displacement (Fig. 26a, p. 52).
The gas can be liquefied by pressure alone at any temperature
below 52° (its critical temperature). One atmosphere pressure
will Uquefy it at —84°, which is therefore the boiling-point of
Uquefied hydrogen chloride.
When the concentrated aqueous solution is heated, it is the hydro-
gen chloride and not the water which is vaporized, for the most part.
When the concentration has been reduced to 20.2 per cent, the
rest of the mixture distils unchanged at 110°. This occurs because,
at this concentration, the hydrogen chloride is carried off in the
bubbles of steam in the same proportion in which it is present
in the Uquid. If a dilute solution is used, water is the chief prod-
uct of distillation (about 100°), but gradually the boiling-point
rises and, when the concentration has reached 20.2 per cent once
more, the same hydrochloric acid of constant boiling-point (110°
at 760 nmi.), as it is called, forms the residue.
Chemical Properties of Hydrogen Chloride* — In the case
of a compound, the chemical property in regard to which we first
enquire is its stability (p. 27). Is it easy or difficult to decompose
by heating? Hydrogen chloride must be heated above 1500°
before even a trace of it is dissociated into hydrogen and chlorine.
Pure hydrogen chloride is therefore a very stable and, from a
chemical point of view, rather an inactive substance. It has no
action on non-metals, such as phosphorus, carbon, sulphur, etc.
However, many of the more active metals (see p. 64), such as
potassium, sodium, and magnesium, decompose it. Hydrogen
is set free, and the chloride of the metal is formed.
2K + 2HCI-^2KC1 + H2T.
HYDROCHLORIC ACID. CALCULATIONS 131
When hydrogen chloride is mixed with ammonia NH3 the gases
unite to form a cloud of fine, soUd particles of ammonium chloride.
HCl + NH3 -> NH4CI i .
Chemical Properties of Hydrochloric Acid. — The solution
of hydrogen chloride in water is an entirely different substance in
its behavior from hydrogen chloride. (1) The solution is sour
in taste, (2) It changes the color of litmits, a vegetable coloring
matter, from blue to red. (3) It is a conductor of electricity y and
is decomposed by the ciurent, hydrogen being liberated at the
negative wire (p. 55). (4) When the metals preceding hydrogen
in the order of activity (p. 54) are introduced into hydrochloric
acid, hydrogen is displaced and Uberated.
In a later chapter (ch. xv) we shall see that these four properties
of hydrogen chloride in aqueous solution are properties conmion to
all substances called acids. We may sum up the main proper-
ties of a solution of hydrogen chloride in water in one word, there-
fore, by saying that it is an acid.
Hydrochloric acid interacts with many other compounds in
solution. In some instances, one of the new substances produced
can be seen, because it appears as a precipitate. One such ex-
ample has already been discussed in detail (see equation 2, p. 127).
When hydrochloric acid is added to a solution of silver nitrate
(AgNOa limar caustic), a precipitate of silver chloride (AgCl)
is obtained,, which is white and curdy in appearance. The other
product, nitric acid (HNO3), remains dissolved and invisible:
HCl + AgNOa -^ AgCl i + HNOa.
Uses of Hydrochloric Acid. — This substance is used as a
source of chlorine. It is employed for cleaning metals. Al-
though present in very small proportions (about 1 part in 500)
in the gastric juice of the stomach, it is a most important compo-
nent of this fluid. It is sometimes given as a medicine, when the
natural supply is too small.
132 SMITHES INTERMEDIATE CHEMISTRY
Double Decomposition. — In this chapter we have met for
the first time with another variety of chemical change. If we ex-
amine the equation for the action of silver nitrate on hydrochloric
acid (p. 131), we shall see that the silver nitrate decomposed as if
it had been made up of two parts, namely (Ag) and (NO3). The
hydrochloric acid similarly separated into its two parts (H) and
(CI). The (Ag) then united with the (CI) and the (H) with the
(NOa).
(Ag)(N03) + (H)(C1) ~> (Ag)(Cl) + (H)(N03).
Since both original substances decomposed, this whole change
is called a double decomposition. A sort of exchange between
the halves of the decomposing substances took place.
The hydrogen chloride was prepared by an action (p. 126),
which, if we write it as follows, is seen to be of the same class:
(Na)(Cl) + (H)(HS04) -^ (H)(C1) + (Na)(HS04).
The Varieties of Chemical Change, — Almost all chemical
changes belong to one or other of the varieties we have already
met with and defined (pp. 14, 16, 51). These, along with one
example of each, are now placed together:
1. Combination: Zn + S -* ZnS.
2. Decomposition: 2KCIO3 -^ 2KC1 + 3O2.
3. Displacement: Zn + H2SO4 -* H2 + ZnS04.
4. Double Decomposition: AgNOs + HCl -* AgCl + HNO3.
In the first, 2 substances give 1 substance.
In the second, 1 substance gives 2 (or more) substances.
In the third, 1 element and 1 compound give 1 element and 1 comr
pound.
In the fourth, 2 compounds give 2 compounds.
This classification suffices for most purposes. But, for special
kinds of cases, some other names are used. Thus, a dissociation
HYDROCHLORIC ACID. CALCULATIONS 133
(p. 104) is an action which belongs to both of the first two classes,
because it is reversible. For example,
2H2O ?± 2H2 + O2.
Again oxidation (p. 40) and reduction (p. 57) are connected with
the particular substances, such as oxygen, which are concerned
in the action. The first classification (Nos. 1 to 4) paid no at-
tention to the kinds of elements which were present. Thus,
every decomposition is a decomposition. If it is reversible, then
it is also a dissociation. If oxygen is set free, then it is a reduction
as well.
Calculations
Calculations Connected with FormuloB. — In a previous
chapter (p. 99) we have seen that formulae represent the com-
position of substances, and we have seen how the formula of
each substance is worked out from the data obtained by experi-
ment. Some ways in which the information contained in formulae
can be used may now be taken up.
The Composition front the Formula. — Take, for example,
the formula for silver nitrate, AgNOs. To learn the composition
of this compound, we look up the atomic weights (see rear cover
of this book). We find Ag = 107.88 parts of silver, N = 14.008
parts of nitrogen, O3 = 3 X 16 or exactly 48 parts of oxygen.
The proportions of the constituents, in the same order, there-
fore, are 107.88 : 14.008 : 48.
What is the proportion of oxygen to nitrogen alone? It is
48 : 14.008, or 3.427 : 1.
Significant Figures. — The division of 14.008 into 48 really
gives the quotient 3.426613. But no atomic weights have been
measured so accurately that we know the values of the numbers
beyond the third place of decimals. For many elements, we do
134 smith's intermediate chemistry
not know even the first place accurately. Hence S.4^7 is just
as likely to be the exact value a« the longer number. In the most
exact calculation we round the number off, usually, at the second
decimal. For rougher purposes the first decimal place is suffi-
cient.
The Formula' Weight. — The sum of the weights of the
constituents indicated in the formula is^called the formula-weight.
For silver nitrate this has the value, 107.88 + 14.008 + 48, or
169.89 (Query: Why not 169.888?).
If the substance is a gas, or is easily volatile, the formula-weight
will be also the molecidar weight. Thus, acetylene gas (used in
Ughting) has the formula C2H2. The composition is C2 = 2 X
12.005 or 24.010 parts of carbon and H2 = 2 X 1.008 or 2.016
parts of hydrogen. The molecular weight is 24.010 + 2.016
= 26.026. Again, hydrogen peroxide (used in medicine), a li-
quid, is decomposed when boiled, but it dissolves in water and
depresses the freezing-point. Its molecular weight has therefore
been determined (p. 119) and is H2O2. The molecular weight on
which this formula is based is 2.016 + 32, or 34.016.
When the substance is neither volatile nor soluble, the simplest
formula is always used, and therefore only the formula-weight
can be ascertained.
To Find the Percentage Composition. — In silver nitrate
the proportions are 107.88 of silver, 14.008 of nitrogen, and 48 of
oxygen in a total of 169.89. In one hundred parts, the silver will
be j57|8 ^ 100, or 63.50; the nitrogen j|^ X 100, or 8.25,
48
and the oxygen ^^^ ^^ X 100 = 28.25.
loy . oy
The same results may be obtained by the rule of proportion.
Thus, for the silver, 169.89 : 107.88 :: 100 : a; where x is the per-
centage of silver.
HYDROCHLOKIC ACID. CALCULATIONS 135
*
Calculations by Use of Equations. — We frequently desire
to know what weight of a product can be obtained from a given
weight of the necessary materials. For example, what weight of
zinc sulphide can be made with 100 g. of sulphur? It is under-
stood, of course, that the necessary zinc is available.
In such calculations mistakes are easily made. The foUovh
ing rules must be strictly followed:
First, write down the equation:
Zn + S->ZnS.
Second, place beneath each formula the weight for which it
stands:
Zn + S -» ZnS.
65.37 32.06 97.43
Third, read the whole statement. In this case it reads: 65.37
parts by weight of zinc combine with 32.06 parts of sulphur to
give 97.43 parts of zinc sulphide.
Fourth, re-read the original problem. Then place the amount
given in the problem (100 g. of sulphur) under the formula of
the substance concerned. Then observe that the problem asks
" What weight of zinc sulphide?" and place an interrogation
point or an X under the formula of thai substance:
Zn + S -> ZnS
65.37 32.06 97.43
100 ? or x
Fifth, read the problem as it now appears in this expanded
equation: 32.06 g. of sulphur will give 97.43 g. of zinc sulphide,
therefore 100 g. of sulphur will give z g. of zinc sulphate.
SioUh. The answer may be now obtained by stating the pro-
portion in the same order:
32.06 : 97.43 :: 100 : a; (= 303.9).
136 smith's intermediate chemistry
If the expanded equation has been prepared correctly, this final
statement is purely mechanical. It will be seen that only two of
the three quantities in the equation were really used.
Alternative to the Sixth Step. — We may also say: If
32.06 g. of sulphur will give 97.43 g. of zinc sulphide, 1 g. of sulphur
97 43
will give oc^ r^a S- ( = 3.039 g.) of zinc sulphide. Then, if 1 g. of
sulphur gives 3.039 g. of zinc sulphide, the 100 g. of sulphur will
give 100 X 3.039 g. ( = 303.9 g.) of zinc sulphide.
Warnings. — In solving the exercises at the end of the chapter,
beware of three hinds of mistakes commonly made by beginners.
1. Conquer a tendency to say that the symbols Zn and S
stand for " 1 part " of zinc or of sulphur. They stand for 1 chemi-
cal unit, or atomic weight, or atom, in each case, — that is to
say, for 65.37 " parts " and 32.06 " parts," respectively.
2. Follow the rules laid down above. When one has once
become familiar with the art of solving such problems, running
through the rules takes only a few seconds. The chemist does it
almost unconsciously. The beginner always thinks he can ignore
these rules, and he fails in consequence. Writing the equation
in the expanded form, and then reading the problem into it are
absolutely essential steps.
3. Do not read the original problem carelessly and make the
equation backwards, that is, with the sides reversed. If there
seems to be confusion somewhere, when the last steps are reached,
this hint will probably show the cause of the difficulty.
Another Example. — What weight of hydrogen is required to
reduce 45 g. of magnetic oxide of iron to metallic iron?
Following the rules, as before, we reach the expanded equation:
Fe304 + 4H2 -* 3Fe + 4H2O.
3 X 55.84 + 4 X 16 8 X 1.008 3 X 55.84 4 (2 X 1.008 + 16)
167.52 + 64 8.064 167.52 4 X 18.016
231.52 8.064 167.52 72.064
45 g. ^
HYDROCHLOKIC ACID. CALCULATIONS 137
Observe that the atomic weights are multiplied by the sub-nmn-
bers, so that, for example, Fez = 3 X 55.84. Observe also that
the formula weights are multiplied by the coefficients, when such
occur, in front of the formula, so that, for example, 4H2O = 4
X 18.016.
The proportion 231.52 : 8.064 :: 45 : a; (= 1.57) supplies the
answer, 1.57 grams of hydrogen.
Using the alternative plan (p. 130) : If 231.52 g. of magnetic
oxide are reduced by 8.064 g. of hydrogen, 1 g. will be reduced
8 064
by ' p,^ g. ( = 0.035 g.) of hydrogen. Hence, if 1 g. of magnetic
oxide is reduced by 0.035 g. of hydrogen, 45 g. will be reduced by
45 X 0.035 g. (= 1.57 g.) of hydrogen.
Exercises. — 1. Complete the equation ZnCU + H2SO4 -^
ZnS04 + , and attach the name of the substance to each for-
mula in it.
2. Point out the differences in physical properties between
oxygen and hydrogen chloride.
3. Make equations for the displacement of hydrogen from
hydrochloric acid by zinc and by sodium (pp. 51 and 52).
4. Give additional examples of the four varieties of chemical
change (p. 132).
5. Classify (p. 132) the following actions: (a) the action of
steam on iron (p. 51); (b) the rusting of iron; (c) the electrolysis
of dilute hydrochloric acid (p. 55); (d) the effect of heating the
hydrate of cupric sulphate (p. 68).
6. Is the decomposition of mercuric oxide a case of: (a) dis-
sociation, or (b) oxidation or reduction?
^ 7. Wbatis the proportion of: (a) sodium to one part of chlorine
m salt; (b) one part of hydrogen to nitrogen in nitric acid (HNOs)?
fr- 8. Calculate the percentage composition of: (a) sulphuric
add; (b) acetylene C2H2.
•^ 9. What weight of hydrogen is displaced by the action of 100 g.
of zinc upon an excess of hydrochloric acid (ZnCl2 is formed)?
138 smith's intermediate chemistry
10. What weight of silver chloride is formed by the interaction
of 5 g. of sodium chloride with silver nitrate?
11. What should be the difference in cost of 1000 g. of oxygen
according as it is obtained by decomposing mercuric oxide at $3.00
per kilog. or potassiiun chlorate at 30 cents per kilog.?
CHAPTER XII
CHLORINB. CALCULATIONS
Chlorine was discovered by Scheele (1774). It was supposed
to be a compound containing oxygen, until the contrary was
proved by Davy (1809-1818). It is used in immense quantities
for making bleaching and disinfecting agents, explosives and
dyestuffs. For its use in gas warfare, see Chapter XL.
Occurrence. — Many compoimds of chlorine occur in nature,
but the most plentiful is common salt (NaCl). Most of the mat-
ter dissolved in sea water is sodimn chloride. Salt also occurs
underground, either in strata in almost pure form, or mixed with
rocky material. Near such deposits, wells and springs of salt
water are common.
Preparation by Electrolysis of a Chloride. — Chlorine is
liberated by passing a current of electricity through a concen-
trated aqueous solution of a chloride, such as hydrogen chloride
or sodium chloride (see Fig. 27, p. 54). Much of it is, in fact,
manufactured by electrolyzing natural brines. The chlorine is
liberated, first in solution and later as a gas, at the positive wire
(anode).
Other Products of Electrolysis of Chlorides. — In elec-
trolyzing hydrochloric acid, the hydrogen is set free at the nega-
tive wire (cathode). With a solution of sodium chloride we might
expect to get free sodiimi at this wire. It will be recalled, however,
that sodium is very much more active than is hydrogen, and
indeed displaces hydrogen from water. Hence the electrical
energy seta free the more easily liberated element — the hydrogen —
130
140 smith's intermediate chemistry
and the sodium remains in the solution as sodium hydroxide
(NaOH). The process is best shown by a diagram:
:2 Na: CI -^ CI2 T (pos. wire)
(neg. wire) t H2 ^ 2H j OH \
The chlorine CI2 and hydrogen H2 being liberated, leave behind
in the solution the constituents of 2NaOH:
2NaCl + 2H2O + Elect. -^ H2 + CI2 + 2NaOH.
Preparation from Oxygen and Hydrogen Chloride. —
When hydrogen chloride and oxygen gases are heated, they inter-
act very slowly to give water and chlorine. The action is greatly
hastened by contact with copper chloride. Lumps of pumice,
saturated with a solution of
this catalyst (see p. 32), are
Pjq 43 placed in a tube. When the
mixture of gases is passed over
the heated pumice (Fig. 48), steam, chlorine, and about 20 per
cent of unchanged oxygen and hydrogen chloride issue at the
other end:
Skeleton equation: HCl + 02^ H2O + CI2.
Balanced equaiion: 4HC1 + O2 ^ 2H2O + 2CI2.
Longer heating does not alter the proportion of the materials
successfully transformed. This is Deacon's process.
That 80 per cent is changed, and 20 per cent unchanged, is
due to the fact that the action is reversible. If we lead pure
chlorine and steam through the tube (read the equation back-
wards), 20 per cent of hydrogen chloride and oxygen are formed.
No more than 20 per cent is formed, because these products are
continually being used up again and reproduce steam and chlorine.
If one product could be separated (p. 128) from the other, to pre-
vent the backward action, the yield CQuld be r^^i^ed to 100 per
CHLORINE. CALCULATIONS
141
cent. But all the four substances are gases (in the hot tube),
and mix perfectly.
The results here noted are interesting, because they show that,
under the conditions of the experiment, oxygen is somewhat more
active than chlorine in combining with hydrogen. The precise
proportions of the four gases present at equiUbrium depend on
the temperature at which the experiment is carried out. In the
commercial appUcation of Deacon's process, this is near 345°.
Preparation from Hydrochloric Acid and an Oxidizing
Agent. — The best way to make a supply of chlorine in the
laboratory is to place potassium
permanganate crystals (KMn04)
in a flask (Fig. 49) and allow
concentrated hydrochloric acid,
previously mixed with an equal
volume of water, to fall upon
them drop by drop. The gas
is rather soluble in water, and is
best collected by displacing the
air from bottles. When one
bottle is full, it is stoppered
and a fresh one substituted. To
avoid the escape of the very irritating gas into the room the tube
from the collecting bottle dips beneath sodium hydroxide solution.
The essential feature of this reaction is that the oxygen of the
potassium permanganate unites with the hydrogen of the hydro-
chloric acid to give water. The potassiimi and manganese take
as much chlorine as they require to form their chlorides, KCl
and MnCl2. The rest of the chlorine is Uberated.
Skeleton: KMnO* + HCl -* H2O + KCl + MnCla + CI?.
To convert all the oxygen of the KMn04 to water, we re-
quire 8HC1. The formation of KCl and MnCU uses up 3 of the
Fig. 49
142 smith's intermediate chemistry
8 atoms of chlorine thereby made available, leaving 5C1 to be
Uberated. Since this chlorine is obtained as CI2, it is necessary
to double our quantities throughout.
Balanced: 2KMn04 + 16HC1 ~> 8H2O + 2KC1 + 2MnCl2 + SCU.
This action is an oxidation of the hydrogen chloride by the per-
manganate. The potassium pennanganate, which suppUed the
oxygen, is called the oxidizing agent Since the pennanganate
lost oxygen, it was itself reduced. In all oxidations one substance
is oxidized and another reduced.
Deacon's process (p. 140) is also an oxidation of hydrogen chlo-
ride (by free oxygen). The oxygen is reduced to water.
Chlorine may be prepared by using other substances to oxidize
hydrochloric acid. Amongst those which are suitable are man-
ganese dioxide Mn02, potassium chlorate KCIO3 and red lead
Pb804.
Manganese Dioxide and Hydrogen Chloride. — The action
of manganese dioxide upon hydrochloric acid is an instructive one.
It is a general rule, of which we shall meet many applications, that
when an acid interacts with an oxide of a metal, there are two
constant features in the result, namely: (1) The oxygen of the
oxide combines with the hydrogen of the acid to form water, and
(2) the metal of the oxide combines with the acid radical of the
acid. Here the skeleton equation should be Mn02 + HCl -^
H2O + MnCU. With O2, to form water, 4HC1 is required, and
the product is 2H2O. Hence the equation is
Balanced: Mn02 + 4HC1 -^ 2H2O + MnCU.
This is, undoubtedly, what happens in the first place. The prod-
ucts actually obtained on heating the mixture, however, are
water, manganous chloride MnCl2 and chlorine. We owe the
chlorine to the fact that the tetrachloride is imstable. At low
temperatures it decomposes into m£tng£tnese trichloride (MnCli)
CHLORINE. CALCULATIONS 143
and chlorine. When the mixture is warmed, the MnCU breaks
down further into MnClz and chlorine. The complete series of
reactions may be represented in one equation as follows:
Mn02 + 4HC1 -^ 2H2O + MnClz + CI2. (1)
If we had used manganous oxide MnO, we should have had a
double decomposition:
MnO + 2HC1 -* H2O + MnCl2, (2)
but we should have got no chlorine. The difference between
these two actions will be discussed in a later chapter (p. 215).
Physical Properties. — Chlorine is a greenish-yellow gas, and
takes its name from the Greek for this color. It has an exceedingly
disagreeable odor and irritates the lining of the nose and throat.
Alcohol vapor or anmionia, when breathed, reheves the irritation.
The density of the gas is recorded in the formula CI2. The
22.4-liter-cube-full weighs 70.92 g., against 28.95 g. for air, so
that chlorine is about 2^ times heavier. Two volumes of the
gas dissolve in one volume of water at 20°. The solution is called
chlorine-water.
The gas is liquefied by pressure below 146*^ (its critical tem-
perature), the pressure required at 20° being 6.6 atmospheres.
The Uquid boils at —33°, and sohdifies at —102° (its melting-
point).
Chemical Properties. — Chlorine is an element with about
the same degree of activity as oxygen (compare p. 141), and it
unites with very much the same hst of other elements. The
compoimds are called chlorides.
Unites with Metals. — When powdered antimony (cold) or
iron powder (warmed) are thrown into chlorine, they combine
with it, and red hot particles of the chlorides, SbCU or FeCU, fall
to the bottom. Copper leaf (Dutch metal, used in " gilding "),
144 smith's intermediate chemistry
or heated copper foil, bums in the gas, giving a fog of solid cupric
chloride CUCI2.
Skeleton: Sb + CI2 -^ SbCU.
Balanced: 2Sb + 3CI2 -^ 2SbCl3.
Sodium bums briUiantly in chlorine, giving sodium chloride.
That a shining metal and a poisonous irritant like chlorine, in
uniting, should 3deld a mild, household article like common
salt illustrates very well the extraordinary nature of chemical
change.
When thoroughly freed from moisture, chlorine no longer
combines with metals like copper and iron. Water seems to
be needed as a cordact agent, in these, as well as in hundreds of
other chemical actions. Hence, carefully dried chlorine in com-
pressed hquid form can be, and is, stored and sold in iron cylinders
(see detinning, p. 508).
Unites with Hydrogen. — A jet of burning hydrogen, lowered
into a bottle of chlorine, continues to bum, giving hydrogen
chloride HCl, the presence of which is shown by the fog produced
by allowing the gas to come in contact with moist air:
H2 + CI2 -^ 2HC1.
Hydrogen and chlorine, mixed, do not combine when cold,
provided strong Ught is excluded. But sunlight, or light from
burning magnesium (" flashUght powder "), starts the combina-
tion, which occurs with explosive violence. Plimging a Ughted
taper into the mixture has, of course, the same effect.
Acts upon Compounds Containing Hydrogen. — Because
of its activity toward hydrogen, chlorine removes hydrogen from
many compounds. Thus, if a Ughted wax taper be plunged into
chlorine, it continues to burn, though with a feebler flame. Dense
smoke, composed of particles of free carbon, rise from the flame.
CHLORINE. CALCULATIONS 146
Blowing the breath into the jar, afterwards, gives the fog due to
hydrogen chloride. Thus the presence of hydrogen and carbon
in the wax of the taper is proved. From this we learn, also, that
chlorine has a relatively small tendency to combine with carbon.
A few drops of warm turpentine (CioHw) upcm a slip of filter paper
will blaze up in chlorine, giving hydrogen chloride and fitn immense
cloud of soot:
Skeleton: C10H16 + CU-^ C + HCl.
Balanced: CioHw + 8CI2 -* IOC + 16HC1.
Acts upon Water. — We have seen that chlorine acts on
steam, reversing Deacon's reaction to the extent of 20 per cent.
It a^ also upon cold water, when dissolved in the latter, although
in a similarly incomplete way. With half-saturated chlorine-
water at 10*^, about one-third of the chlorine is transformed.
One of the products, hypochlorous acid HOCl, is of especial
interest, because it is an exceedingly active substance, much used
as an oxidizing agent (see p. 223) and for bleaching:
CI2 + H2O ^ HCl + HOCl.
Unites with Non-metals. — Phosphorus bums in the gas,
giving the vapor of phosphorus trichloride PCI3. This substance,
with excess of chlorine, forms the pentachloride PCls. When
phosphorus pentachloride is heated, it vaporizes and dissociates
again partially into the trichloride and chlorine, according to the
reversible reaction:
PCI5 ^ PCI3 + CI2.
Sulphur, when heated with chlorine, combines slowly, giving sul-
phur monochloride S2CI2, used in vulcanizing rubber.
Although chlorine does not combine very readily with carbon,
several compounds of carbon and chlorine are very important,
such as chloroform CHCI3, carbon tetrachloride CCI4 (p. 332)
and phosgene COCI2 (p. 484).
146 smith's INTERUED£A.T£ chemistbt
CompositUtn of Hydrogen Chloride. — Now that we are
familiar with the properties of chlorine, as well as with those
of hydrogen, we may return to the
question of the. proportion by vol-
ume in which they are produced by
decomposition of hydrogen chloride.
When we electrolyze hydrochloric acid
in the apparatus shown in Fig. 27
(p. 54), we find that the chlorine dis-
solves to a lai^e extent in the liquid,
and its true volume as gas is not easily
ascertained. The apparatus (Brown-
lee's) in Fig, 50 avoids the difficulty
by enabling us to saturate the Uquid
with chlorine before observing the col-
lected gases. The volumes of the two
gases are found to be equal.
A simpler apparatus (Fig. 51) may be
used to show the same fact. The gases
are generated, by electrolysis, in the test-
tube, pass through the straight 'tube,
driving the air before them, and finally bubble through sodium hy-
D*— ^
droxide solution. The whole apparatus must be covered with a
dark cloth to exclude light, and handled in diffused light. In fifteen
CHLORINE. CALCULATIONS
147
minutes, chlorine ceases to be dissolved in the test-tube, and
the gases come off in their natural proportions. In half-an-hour
more they have filled the tube. The stop cocks are now closed,
the tube is set in a tall cyUnder containing potas-
sium iodide (KI) solution (Fig. 52) and the lower
stop cock is opened. The potassium iodide acts
upon the chlorine, giving potassium chloride and
free iodine which is a sohd,
2KI + CI2 ~> 2KC1 + I2,
i
"ff
and the Uquid rises until it fills half the tube. The
remaining gas bums and is hydrogen.
The volume of the hydrogen chloride, in relation to
the volumes of the constituents, may be learned
by using a different apparatus (Fig. 53). A test-
tube of heavy glass is filled with dry hydrogen
chloride and closed with a rubber stopper greased
with vaseUne. A Httle sodium amalgam (solution
of sodium in mercury) is in-
troduced, and the stopper in-
stantly replaced. When the
contents are shaken for one
or two minutes, the sodium combines with
the chlorine and the hydrogen is Uberated
and remains. The mouth of the test-tube
is then immersed in a jar of water, and
the stopper withdrawn. The water rises
and fills about half the tube.
Conclusion, The hydrogen has half the
volume of the hydrogen chloride, and the volume of the chlorine
is equal to that of the hydrogen. Therefore:
1 vol. hydrogen + 1 vol. chlorine — > 2 vols, hydrogen chloride.
Fig. 52
Fig. 53
This result further illustrates Gay-Lussac's law (p. 60).
148 smith's intermediate chemistry
Confirmation of the FormuUe Ch o,nd fl2« — According to
Avogadro's hypothesis, there are equal numbers of molecules in
equal volumes of these gases. Let the rectangles represent small
volumes, containing 1000 molecules each:
Hydrogen Chloride Hydrogen Chlorine
1000 1000
1000 + 1000
It appears from this that 2000 molecules of hydrogen chloride
come from 1000 molecules of hydrogen and 1000 molecules of
chlorine. Now, each molecule of hydrogen chloride contains
at least one atom of hydrogen, so that the 1000 molecules of hydro-
gen must have given at least 2000 atoms of hydrogen, one for
each molecule of the compound. Hence each molecule of hydro-
gen contains at least two atoms. The same is true of each mole-
cule of chlorine. We have no reason, however, for supposing
that there are more than two atoms in either molecule; no sub-
stance is known which contains less than 1.008 g. hydrogen or
35.46 g. chlorine in its unit weight (p. 75). Hence the con-
clusion is confirmed which we reached before (p. 79), namely, that
the formulae of the free gases are H2 and CI2, and that single atoms
of the elements occur only in combination (as in NaCl, HCl, etc.).
Calculations
Calculations Involving Volumes of Gases. — For such cal-
culations the simplest method is always to use the gram-molecular
volume — the 22.4-liter cube.
Relative Densities of Gases. — What are the relative weights
of equal volumes of chlorine and hydrogen chloride? The for-
mulae record the weights of 22.4 hters: CI2 = 2 X 35.46 = 70.92,
and HCl = 1.008 + 35.46 = 36.468. The required relation is
70.92 : 36.468, so that chlorine is almost twice as heavy, bulk for
bulk, as hydrogen chloride.
CHLORINE. CALCULATIONS 149
Density Compared with Air. — We have seen (p. 85) that
22.4 liters of air weigh 28.95 g. Is water vapor heavier or lighter,
at the same temperature and pressure? The formula, H2O, shows
that 22.4 liters of steam weigh 2 X 1.008 + 16 = 18.016. Air is
heavier, in the ratio of 29 : 18, approximately.
The Volume of Gas from Given Weights of Material. —
What volume of oxygen at 0*^ and 760 mm. is obtained by heating
15 g. of potassium chlorate? Follow the rules given (p. 135) for
calculating weights:
2KC108 -> 2KC1 + 3O2.
Weights: 2 (39.1 + 35.46 + 3 X 16) 2 (39.1 + 35.46) 3 X 32
or 245.12 149.12 96
Now, insert the volumes in the case of substances which are gases.
Here oxygen is the only one. Remembering that O2 (=32 g.)
occupies 22.4 Kters, we see that the volume of the oxygen is 3 X
22.4 liters, or 67.2 liters. As the next step, add the data given
and the x in their proper places. The x here goes under the vol-
ume of oxygen:
2KC10s -> 2KC1 + 3O2.
245.12 g. 149.12 g. ' 96 g. (or 67.2 liters)
15 g. X
The problem now reads tnus: If 245.12 g. of potassium chlorate
give 67.2 liters of oxygen, 15 g. of potassium chlorate will give
X liters of oxygen:
245.12 : 67.2 :: 15 : a; (= 4.07),
where a?, the answer, is 4.07 Uters of oxygen.
Observe that, when we use the gram-molecular volume, one
proportion gives the answer. It is not necessary to make two
steps in the calculation, by finding first the weight of the oxygen,
and then its volume. The beginner always tends to do this, until
he learns by experience that it takes twice as long to solve the
problem in this way and that the chance of obtaining the wrong
answer by making arithmetical errors is greatly increased.
150 smith's intermediate chemistry
Another Example. — What weight of copper will combine with
15.2 liters of chlorine (at 0** and 760 mm.)?
Cu + CI2 ~> CuCla.
63.57 g. 70.92 g. or 22.4 L 134.49 g.
X 15.2 liters
The proportion is:
22.4 : 63.57 :: 15.2 : z (= 43.14 g.).
At Other than Standard Conditions. — If the problem
concerns a gas at some temperatm^ and pressure not ff^ and 760
nmi., then correction must be made as a separate calculation.
For example, if the 15.2 Uters of chlorine, in the foregoing illus-
tration, had been at 15° and 742 nmi., then the volume would
first have to be reduced to 0° and 760 mm. (See p. 47.)
Questions of Volume Alone. — When weights are not men-
tioned in the question, but volumes only, the calculation is very-
simple. For example: What are the relative volmnes of oxygen
and hydrogen chloride used in Deacon's process?
4HC1 + O2 -> 2H2O + 2CI2
Molecules: 4 12 2
Volumes: 4 X 22.4 1 X 22.4 2 X 22.4 2 X 22.4
Remembering that equal numbers of molecules occupy equal
volumes, and that 4HC1 = 4 molecules, O2 = 1 molecule, etc., we
perceive that 4 volmnes of HCl will be required for 1 volume of
oxygen.
Does the volume change during the process? Yes, 4+1
volumes become 2 + 2, or 5 volumes give 4.
Warnings. — The conmionest mistake made in these calcula-
tions is to neglect to use, in the equations, the molecular formulae.
We must use H2, O2, C2H2 (acetylene), etc., (and not H, 0, or
CH), because the 22.4 liters hold the weights represented by the
molecular formulae.
CHLORINE. CALCULATIONS 151
Before calculating the volume (gaseous) of a substance repre-
sented in an equation, consider whether it is a volatile substance.
Only volumes of gases can be calculated by the rules given above,
Exercises. — 1. What would be the results of electrolyzing
aqueous solutions of: (a) potassium chloride; (b) cupric chloride
(cf. p. 54).
2. How should you separate the chlorine and the steam pro-
duced by Deacon's process?
3. Make equations showing the interactions with hydrochloric
acid of: (a) manganese dioxide; (b) potassiimi chlorate; (c)
red lead. The metals form MnCL, KCl, and PbCk, respectively.
4. What would be the pressure in a cyUnder of liquid chlorine
at 20°?
5. Make equations for the union of chlorine with: (a) copper;
(b) sodiimi; (c) iron; (d) phosphorus; (e) sulphur.
6. When plunged into chlorine, a jet of illuminating gas con-
tinues to bum. A stream of soot rises from the flame, however,
and blowing the breath into the jar, afterwards, produces a fog.
What do you infer as to the constituents of illuminating gas?
7. What information is conveyed by the fact that the formula
of the chloride of sulphur is written S2CI2, and not SCI?
8. Make the molecular equation for the action of sodiimi upon
hydrogen chloride (p. 131). Why does not the mercury interact
with the latter (c/. p. 54) ?
9. What facts led us, in Chap. VIII, to the conclusion that the
molecular weight of chlorine was 70.92 while its atomic weight
was 35.46?
10. What are the relative densities (p. 85) of: (a) oxygen and
chlorine; (b) hydrogen and hydrogen chloride; (c) hydrogen and
air; (d) acetylene C2H2 and air?
11. What volume of hydrogen chloride at 0*^ and 760 mm. is
obtained by the interaction of 65 g. of sodium chloride, and an
excess of sulphuric acid (p. 126)?
152 smith's intermediate chemistry
12. What weight of zinc is required to make 100 liters of hydro-
gen, at 10*^ and 750 mm., by displacement from hydrochloric acid
(p. 53)?
13. What are the relative volumes of the factors and of the
products in the interaction between: (a) turpentine vapor and
chlorine: (b) oxygen and carbon disulphide vapor (giving SO2
and CO2)?
14. What are the relative volumes of the products in the de-
composition of: (a) mercuric oxide; (b) chlorine monoxide CI2O
(chlorine and oxygen are formed)?
15. What are the relative volumes of the volatile substances
concerned in the action of water vapor on iron (p. 102)?
16. Using the relative volumes in which oxygen and hydrogen
combine to form steam, prove that each molecule of free oxygen is
composed of at least two atoms (cf. p. 105).
17. (a) Calculate the weight of chlorine dissolved by 100 c.c.
of water at room temperature (p. 143). (b) In half-saturated
chlorine-water, what is the concentration of hypochlorous acid
(p. 145)?
CHAPTER XIII
ENERGY AND CHEMICAL CHANGE
In the description of chemical changes, the fact that heat was
evolved has frequently been mentioned. In several instances a
current of electricity has been used to produce chemical change.
It is now necessary to collect these scattered facts and classify
them for future use.
Physical Accompaniments of Chemical Change. — When
iron and sulphur combined (p. 14), and when iron burned in oxy-
gen or copper in chlorine, much heat was developed. On the
other hand, the decomposition of mercuric oxide, as was pointed
out (p. 15), owed its continuance to the persistent apphcation of
heat and ceased as soon as the source of heat was withdrawn.
Here, apparently, heat was consumed during the progress of the
change, and the chemical action was limited by the amoimt of
heat supplied. The production or consumption of heat may,
therefore, be a feature of chemical change.
In the burning of iron or magnesium in oxygen, and in the ac-
tions of chlorine on copper and tm^pentine, light was also pro-
duced. Conversely, silver chloride can be kept any length of
time in the dark, but in sunUght it becomes first bluish and then
brown, simultaneously giving off chlorine gas and finally leaving
only silver as a fine powder. Silver bromide or iodide, in pho-
tographic plates, films, and paper, is changed by Hght in a similar
way, hberating the bromine or iodine. It would appear, there-
fore, that light may be given out or consumed in connection with
chemical change.
We have seen (p. 55) that a cmrent of electricity may be em-
ployed to decompose hydrochloric acid and other chlorides, and
153
154 smith's intermediate chemistry
the battery, or other source of the current must be kept |[omg or
the chemical change stops. The inverae of this is hkewise familiar.
If we place in dilute sulphuric acid a stick
''''^<'y_^ '_, of the metal zinc, we find that hydrogen is
^\\ili-^'/^^-/^";i.|i,l)' given off (Fig. 54), that the zinc goes into
solution as zinc sulphate (p. 102), and that
a large amount of heat is developed. The
presence of small quantities of other metals
in the zinc catalyzes this reaction. If im-
pure zinc in fine particles, with much sur-
face, is used, the temperature of the liquid
may even rise spontaneously to the boiling-
point. This form of the action produces
heat. If, however, we attach the same
stick of zinc to a copper wire and, having
provided a plate of platinum also con-
Pj^ 54 nected with a wire, immerse the two simul-
taneously in the acid (Fig. 55), then a
galvanometer, with which the wires are connected, shows at once
the passage of a current of electricity round the circuit. Ejcactly
the same chemical chaise goes on as before. The sole difference
ENERGY AND CHEMICAL CHANGE 155
is that the gas appears to arise from the surface of the platinum.
It is easy to show, however, that the platinum by itself is not
acted upon by dilute acids and, in this case. Undergoes no change
whatever; it serves simply as a suitable conductor for the elec-
tricity. Here, then, in place of the heat which the first plan
produced, we get an electric current. The arrangement is, in
fact, a battery-cell, for a battery is a system in which a chemical
action which would otherwise give heat furnishes electricity in-
stead. Thus, electrical energy may be consumed or produced
m connection with a change in composition.
Even violent rubbing in a mortar, in the case of some substances^
can effect an appreciable amount of decomposition in a few min-
utes. In this way silver chloride can be separated into silver and
chlorine, just as by Ught. It is the mechanical energy which is the
agent, and part of it is consumed in producing the change, and
only the balance appears as heat. Conversely, the production of
mechanical energy, as the result of chemical change, is seen in
the behavior of explosives and in the working of our muscles.
Thus, mechanical energy may be used up or produced in chemica
changes.
Summing up our experience, we may state that no change in
composition occiu's without some accompaniments, such as the
production or consumption of heat, Ught, electrical energy, or, in
some cases, mechanical energy.
Classification of the Accompaniments of Change in
Composition: Energy. — The problem of classifying (t.e.,
placing in a suitable category) things Uke heat, light, and elec-
tricity has occupied much attention. In all changes in com-
position, one of these natiu'al accompaniments is given out or
absorbed, sometimes in great amount, yet in none is any alteration
in weight observed.* There are many things which are real,
* Electrons (see p. 195) do possess mass, but it is very small compared with
that of the materials oonoemed.
156 smith's intermediate chemistry
however, even if they are not affected by gravitation. In the
present instance we reason as follows:
A brick in motion' is different from a brick at rest. The former
can do some things that the latter cannot. Furthermore, we can
easily make a distinction in our minds. The brick can be deprived
of the motion and be endowed with it again. Thus, we can get
the idea of motion as a separate conception. Similarly, we ob-
serve that a piece of iron behaves differently when hot, and when
cold, when bearing a current of electricity, and when bearing
none. We conceive then of the brick or the iron as having a
certain amount and kind of matter which is unalterable, and as
having motion, heat, or electricity added to this or removed.
Thus, we describe our observations, by using two categories, one
of which includes the various kinds of mattery and the other,
various things whose association with matter seems to be in-
variable and is often so conspicuous^ The latter we call the forms
of energy.
The Practical Importance of Energy in Chemistry. —
The absorption or Uberation of energy accompanying a chemical
transformation of matter is often, of the two, the more important
feature. We do not bum coal in order to manufacture carbon
dioxide gas. We are glad to get rid of the material product
through the chimney. It is the heat we want. We do not buy
gasoUne (petrol) for an automobile in order to obtain various gases
to expel through the muflBier. We really pay for the mechanical
energy. It is the same with burning illuminating-gas or magnesium
powder when we want hght, and with eating food, which we do,
chiefly, to get energy to sustain our activity. We do not run
electricity for hours into a storage battery, in order to make a
particular compound (lead dioxide, for example), but in order to
save and store the energy for future use. In industry and life
fully half the total amoimt of chemical change involved is set in
motion by us, solely on account of the energy changes it involves.
BNBEGT AND CHEMICAL CHANGE 167
But the productioD of energy in cberaical change is not only thus
of practical importance; it is also of scientific interest, as wiU be
seen in the section on energy and chemical activity (see below).
InterconverttbiUty of Forma of Energy: Conservation. —
At first sight, these different forms of energy seem to be quite
unrelated. But a relation between them can be found. If the
heat of a Bunsen flame or of the sun is brought under a hot-air
motor (Fig. 56) violent motion results. Again, if the motor
is connected with a dynamo, electricity may be generated. Still
again, if the current from the dynamo flows through an incandes-
cent lamp, heat and light are evolved. Conversely, when motion
of the hotrair motor is impeded by a
brake, heat appears. When a current
of electricity is run through the dynamo,
the armature of the latter turns and
motion results. But the most ^gnifi-
cant facte are still to be mentioned.
The heat absorbed by the motor is
found to be greater when the machine is
permitted to move and do work, than
when it is not. Thus, it is found that
when work is done some heat disappears,
and this heat is, in fact, transformed
into work. Similarly, when the poles of
the dynamo are properly connected and Fig. 56
electricity is being produced, and only
then, motion is used up. This is shown by the effort required
to turn the armature under these circmnstances, and the ease
with which it is turned when the circuit is open. So, with a
conductor hke the filament in the lamp, tmless it offers resist
ance to the current and destroys a sufllcient amount of electrical
energy, it gives out neither Ught nor heat. Finally, motion gives
no heat unless the brake is set, and effort is then demanded to
158 smith's intermediate chemistry
maintain the motion. These experiences lead us to believe that
we have here a set of things which are fimdamentally of the
same kind, for each form can be made from any of the others.
We have, therefore, invented the conception of a single thing, of
which heat, Ught, electricity, and motion are forms, and to it we
give the name energy: energy is work and every other thing
which can arise from work and be converted into work.
Closer study shows that equal amounts of electrical or mechani-
cal energy always produce equal amounts of heat. No loss is
ever observed in the transformations of energy, any more than
in the transformations of matter. Hence we have been led to
the conclusion that in a limited system no gain or loss of energy
is ever observed. This brief statement of the results of many
experiments is called the law of the conservation of energy.
Application of the Conception of Energy in Chemistry. —
At first sight it looks as if the statement that energy is conserved
is not appUcable in chemistry. Heat and electricity, for example,
seem to be produced and consumed, in connection with changes
in composition, in a mysterious manner. We trace Ught in an in-
candescent lamp back to the electricity, and this in turn to the
mechanical energy, and this again to the heat in the engine. But
what form of energy gave the heat developed by the combustion
of the coal under the boiler, or by the union of iron and sulphur
in our first experiment? Since we do not perceive any electricity,
Ught, heat, or motion, in the original materials, and yet wish to
create an harmonious system, we are boimd to conceive of the
iron and the sulphur, and the coal and the air, as containing an-
other form of energy, which we caU internal energy. Similarly,
when heat is used up in decomposing mercuric oxide, or Ught in
decomposing silver chloride, we regard the energy as passing into,
and being stored in the products of decomposition in the form of
internal energy.
These conclusions compel us, for the sake of consistency, to
ENERGY AND CHEMICAL CHANGE 159
think of all our materials as repositories of energy as well as of
matter, each of these two constituents being equally real and
equally important. A piece of the substance known as " iron "
must thus be held to contain so much iron matter and so much
internal energy. So ferrous sulphide contains sulphur matter,
iron matter, and internal energy. Thus, by a svbstance we mean
a distinct species of matter, simple or compound, with its appro-
priate proportion of internal energy. During the progress of a
chemical change, like the union of iron and sulphur, the internal
energy of the system also changes. The total energy which can
thus be made available as the result of a chemical action, and
converted (through, say, heat or electrical energy) into work, is
called the free energy of the reaction.
In the course of this discussion it has become clear that it is
characteristic of chemical phenomena that, besides a change in the
nature of the matter, there is always an alteration in the amount
of internal energy in the system. This alteration involves the
production of internal energy from, or the transformation of
internal energy into some other form of energy.
Energy and Chemical Activity, — Other things being equal,
when the free energy of a reaction is large, the reaction proceeds
rapidly; that is to say, a large proportion of the reacting materials
are changed in the unit of time. Those reactions in which the
change of internal energy is small proceed more slowly. The
speed of a chemical change, and the quantity of energy available
because of it, are therefore closely related. Now, we are ac-
customed to speak of materials which, Uke iron and sulphur,
interact rapidly and with Uberation of much energy as "chemic-
ally active,'' or as possessing great " chemical affinity " for one
another. Thiis, relative chemical activity or affinity may be
estimated, (1) by observing the speed of a change or, in many
cases (2) by measuring the heat developed or (3) by ascertaining
the electromotive force of the current, when the materials are
160 smith's intermediate chemistry
arranged in the form of a battery-cell. The order of activity of (he
metalsj given on p. 54, is obtained by this third method.
It is evident that the chemical activity or aflSnity of a given
substance will not be the same towards all others. Thus, iron
unites much more vigorously with chlorine than with sulphur and,
with identical amounts of iron, more heat is Uberated in the former
case than in the latter. With silver, sodium, and many other
substances, iron does not unite at all. One of the tasks of the
chemist is to make such comparisons as this. He calls the results,
the specific chemical properties of the substances in question.
Care must be observed, however, in making comparisons by
the above methods. Although it is true that most chemical
changes that take place readily develop heat, yet at high tem-
peratures compounds can be formed by the direct union of their
elements with absorption of heat (see, for example, p. 312). Re-
versible reactions must also be accounted for. In such reactions,
one of the chemical changes taking place necessarily absorbs
just as much heat as the other develops, yet both have a definite
speed under any given conditions, as is shown by the fact that
a fixed point of equilibrium is reached. Deacon's process (p. 140)
is a good instance. As we shall see later (chapter XX), the relative
concentrations of the reacting substances, as well as their affinities,
must be taken into consideration.
The Cause of Chemical Activity or Affinity. — The
reader will undoubtedly be inclined to inquire whether we can
assign any cause for the tendency which substances have to under-
go chemical change. Why- do iron and sulphur unite to form
ferrous sulphide, while other pairs of elements taken at random
will frequently be found to have no effect upon one another under
any circumstances? A final answer to this question cannot, of
course, be given. As the facts regarding chemical activity or
afl^ty, however, become better known, we may arrive at a
stage where a logical eocplanaiion (see p. 23) can be advanced which.
ENERGY AND CHEMICAL CHANGE 161
basing itself on these facts, affords a means of classifying them
and suppUes us with a new and useful hypothesis regarding mat-
ter in its various forms (iron, sulphur, etc.) and the energy con-
tained therein. A brief discussion of such a hypothesis will be
found in the concluding chapter (pp. 562-5).
The Speed of Chemical Actions: a Means of Measuring
Activity. — One means of measuring the relative chemical activi-
ties of several substances is to observe the speed with which they
undergo the same chemical change. Thus we may compare the
activities of the various metals by allowing them separately to
interact with hydrochloric acid and collecting and measuring the
hydrogen Uberated per minute by each. It will be seen, even in
the roughest experiment, that magnesium is thus much more
active than zinc. The comparison must be made wit!} such pre-
cautions, however, as will make it certain that the conditions
under which the several metals act are all aUke. Thus, in spite
of the heat evolved by the action, means must be used, by suitable
cooUng, to keep the temperature at some fixed point during the
experiment, for all actions become more rapid when the tempera-
ture rises (p. 27). Again, the pieces of the various metals must
be arranged so that equal surfaces are exposed to the acid in each
case. It is found that the order in which this comparison places
the metals is much the same as that in which they are placed by
a study of other similar actions. A single table, showing the
order of activity (p. 54), suflSces, therefore, for all purposes.
Thermochemistry. — Chemical changes in which heat is
liberated are called exothermal. Those in which heat is continu-
ously absorbed (pp. 15, 30) are called endothermal changes. Since
the activities, or afl^ties of two substances (say, two metals)
may often be compared by observing the amounts of heat Uber-
ated when each combines with a third substance (say, oxygen),
it will be instructive now to consider some of the elementary facts
of thermochemistry.
162 smith's intermediate chemistry
The chemical interactions to be studied thermally are arranged
so that they may be carried out in a small vessel which can be
placed inside another containing water. The whole apparatus
is called a calorimeter (Greek, heat-measurer). The heat de-
veloped raises the temperature of this water. Where gases Uke
oxygen are concerned, a closed bulb of platinum forms the inner
vessel. The quantity of heat capable of raising one gram of water
one degree in temperature at 16° Centigrade is called a calorie.
Thermochemical Equations. — While in physics the unit of
quantity is the gram, in chemistry the unit which we select is
naturally a gram-atomic weight or a gram-molecular weight of the
substance. Thus, the heat of combustion of carbon means the
heat produced by combining twelve grams of carbon with thirty-
two grams of oxygen, and is suflScient to raise nearly 100,000
grams of water one degree. This is expressed as follows:
C + O2 -» CO2 + 96,820 cal.
In other words, the combustion of less than half an ounce of car-
bon will raise over two pounds of water from 0° to the boiUng-
point.
When the action is one which absorbs heat, this fact is indicated
by the negative sign preceding the number of calories. Thus for
the dissociation of water vapor into hydrogen and oxygen we
have the thermochemical equation:
2H2O -^ 2H2 + O2 - 116,200 cal.
If the action is reversible, as this one is, the heat absorbed when
it proceeds in one direction is equal to that liberated when it goes
in the other direction:
2H2 + O2 -> 2H2O + 116,200 cal.
Answers to Possible Questions. — It is always foimd that
the same quantities of any given chemical substances, undergoing
the same chemical change under the same conditions, produce
ENERGY AND CHEMICAL CHANGE 163
or absorb, according as the action is exothermal or endothermal,
amounts of heat which are equal.
The rate at which a given chemical action is allowed to take place
has no influence on the total amount of heat consumed or pro-
duced. It may not at first sight appear obvious that rusting
evolves heat, but a deUcate thermometer will show that a heap of
rusting nails is somewhat higher m temperature than surrounding
bodies. Poor conductors, hke oily rags and ill-dried hay, show a
tendency to spontaneous combustion owing to accumulation of the
slowly developing heat of oxidation (p. 41). The warmth of our
own bodies is due to the same cause.
It should be noted that production or absorption of heat is not,
in itself, an evidence of chemical action. Physical changes are all
likewise accompanied by the same phenomena. Thus, the evapo-
ration of water absorbs heat, and condensation of a vapor and the
crystaUization of a supercooled liquid liberate heat.
Exercises. — 1. Which form of energy is delivered as such,
and paid for as such, in most cities?
2. How many calories are required to raise 600 g. of a sub-
stance of specific heat 0.5 from 15° to 37°?
3. The combustion of 1 g. of sulphur to sulphur dioxide de-
velops 2220 calories. What is the heat of combustion of sulphur?
Write the thermochemical equation.
CHAPTER XIV
SODIUM AND SODIUM HYDROXIDE
In our study of common salt we have taken up one of its con-
stituents, namely, chlorine, and its commonest derivative hydro-
chloric acid. The latter is a good example of an acid. We now
turn to the other constituent, sodiimi, and one of its famihar
compoimds, namely sodiimi hydroxide (NaOH). The latter is an
example of a different kind of substances, called alkaUes or bases.
Salt, hydrochloric acid, and sodiimi hydroxide are examples of the
three largest and most important classes of substances known to
inorganic chemistry.
Preparation of the Metal Sodium, — Sodium cannot be
made by displacement (like hydrogen), because it is close to the
top of the activity list (p. 54) of the metals, and
no more active, and at the same time easily ob-
tained, metal is available to displace it. It was
first prepared by Davy (1807) by electrolyzing
melted sodiimi hydroxide, and is still manufac-
tured in this way. The aqueous solution of a
sodium compound, such as sodium chloride, can-
not be used, because, as we have seen (p. 140),
hydrogen from the water is Uberated in place of
the metal.
The dry sodium hydroxide is melted (about 318°) in an iron
vessel (Fig. 57), which is connected with the positive wire from the
dynamo, and the oxygen is Uberated on the iron (anode). The
negative wire is connected with rods of carbon which stand up
through the bottom of the vessel, and here (cathode) the sodium
and the hydrogen are set free. Being Ughter than the fused com-
164
SODIUM AND SODIUM HYDROXIDE 165
pound, the hydrogen rises in bubbles, and the sodium in melted
globules, to the surface. Here they collect under an iron cylinder.
Tha latter is made of wire-gauze at the lower part, to permit cir-
culation of the Uquid, but prevent the escape of the globules of
sodium. It is closed at the top, to prevent the heated sodium
from burning, as it would do if air could reach it. The melted
sodiima is ladeled into cyUndrical moulds, and the sticks of the
metal are preserved in air-tight tin boxes.
Physical Properties of Sodium. — Sodiimi is a silver-like
metal, of specific gravity 0.97. It is soft and can be cut with a
knife. It melts at 95.6° and boils at 742°. The gram-molecular
volume of sodium vapor weighs 23 g., the same as the atomic
weight, so that the molecular formula is Na.
Chemical Properties. — Sodium bums in chlorine, giving
sodium fchloride NaCl. It burns also in oxygen (or air) to form
sodimn peroxide Na202. It acts violently on water, as we have
seen (p. 50), displacing hydrogen and forming sodium hydroxide.
Skeleton: Na + H2O -^ NaOH + H2.
Balanced: 2Na + 2H2O -> 2NaOH + H2.
For this reason it tarnishes quickly in moist air. In the labora-
tory small amoimts are kept under kerosene, which contains no
compounds of oxygen.
Uses. — Sodiimi is used in the manufacture of many complex
organic compoimds. By contact action, it converts isoprene
(CbHs) into caoutchouc (CioHi6)n or raw rubber (see p. 480).
This is a method of making rubber artificially. It cannot yet be
carried out so cheaply as to compete with the natural product
under ordinary circumstances, but in Germany during the war,
when the supply of natural rubber was cut off entirely, consider-
able quantities were manufactured by this sjmthetic method.
166 smith's intermediate chemistry
Sodium peroxide (" oxone '')> niade by burning sodium, is used
as a source of oxygen (p. 31).
Preparation of Sodium Hydroxide. — As we have seen, this
compound is formed by the action of water on sodium, and may
•be obtained by evaporating the solution to dryness. But manu-
facturing it from an expensive substance Uke sodium is out of the
question.
Much sodium hydroxide is made by boiling an aqueous solution
of sodiimi carbonate Na^COs with slaked Ume (calcium hydroxide,
Ca(0H)2) :
Ca(0H)2 + NaaCOs ^ 2NaOH + CaCOs i-
The calcium carbonate, which is chemically the same substance
as Ume-stone and chalk, is much less soluble than any of the other
substances taking part in the reaction, and is consequently pre-
;cipitated out of the solution. This reaction is another example
of a double decomposition (p. 132). All double decompositions
are reversible, but here the precipitation of the calcium carbonate
prevents it from acting upon the sodium hydroxide and reversing
the action to any appreciable extent. After the precipitate has
settled, the solution of sodium hydroxide is drawn off and evap-
orated to dryness.
Electrolytic Process. — Sodium hydroxide is also manufactured
by electrolysis of sodium chloride solution (pp. 139-40), the other
product, chlorine, being of great commercial value also. The
Nelson cell (Fig. 58) is now most extensively used. A porous
diaphragm of asbestos separates the perforated steel cathode
from the carbon anode, which is immersed in a current of brine
flowing through the cell between the electrodes. Chlorine is
liberated at the anode and rises in bubbles to the surface of the
solution. It is drawn off, dried, and compressed to liquid form
in iron cylinders, or is naade directly into bleaching compounds
(p. 224). Sodium, as we have seen in the discussion of electroly-
SODIUM AND SODIUM HYDROXIDE 167
eis on pp. 139-40, remams in the solution as sodium hydroxide,
which collects around the cathode and flows out into a catch
basin. It is purified from residual sodium chloride by fractional
Copper Bus Bar
Copp«r ■ftrmuial
SUM.
cattiiMe
crystallization. Hydrogen is also liberated at the cathode, and
is a valuable by-product of the process.
Physical Properties of Sodium Hydroxide.— The substance
is a white crystalline solid. Generally it shows the form of the
iron drums, into which it is run when melted, or of the sticks into
which it is cast. It is exceedingly solitble in water. Its solution
gives to objects the smooth, soapy feeling which is characteristic
of alkalies. The solution is sometimes called soda-lye, and the
solid, caustic soda.
Chemical Properties. — Sodimn hydroxide is exceedingly
stable, being melted, but not decomposed, by beating.
The aqueous solution possesses the following important prop-
168 smith's intermediate chemistry
erties. (1) The solution has an acrid taste, like soap or borax.
(2) It changes the color of litrmiSj reddened by a trace of an acid,
back again from red to blue. (3) It is a condiLctor of electricity,
and is decomposed by the current, oxygen being liberated at the
positive wire.
In the following chapter we shall see that these three proper-
ties of sodiimi hydroxide in aqueous solution are properties com-
mon to all substances called alkalies. The reader should at this
point refer back to p. 131, and contrast these properties with
those exhibited by adds.
Sodium hydroxide in solution enters into double decomposition
with many substances. Frequently one of the products is in-
soluble, and appears as a precipitate. For example, with a solu-
tion of cupric chloride, sodium hydroxide gives a precipitate of
cupric hydroxide.
Skeleton: NaOH + CuCla -^ Cu(0H)2 i + NaCl.
Balanced: 2NaOH + CuCls -^ Cu(0H)2 i + 2NaCl.
As this equation shows, sodiimi hydroxide behaves in such actions
as if composed of two parts, namely (Na) and (OH) (compare p.
132). The reaction consists, essentially, of a transfer of (OH)
groups from (Na) to (Cu).
Alkalies and Bases. — It will be seen that the chemical proper-
ties of sodium hydroxide solution may be summed up by saying
that it is an alkali.
Solutions of the alkaUes also a^t upon animal matter, e,g, wool
(p. 1), especially when hot, converting it largely into soluble
substances. For this reason they are called caustic alkaUes.
They Hkewise act slowly upon the components of glass. For
this reason a precipitate is often visible in the caustic soda reagent
bottle, and the inner surface of the glass is always etched.
A very dehcate test for an alkaU is given by phenolphthalein,
a colorless organic substance. One drop of phenolphthalein
SODIUM AND SODIUM HYDKOXIDE 169
solution added to an alkali in water solution produces an in-
tense red (when dilute, pink) coloration. Addition of excess acid
renders the solution colorless again.
The alkalies, however, are simply the more active members
of a much larger class of substances called bases. Solutions of the
less soluble bases, of which cupric hydroxide is an extreme ex-
ample, do not show, distinctly, all the properties exhibited by
alkalies. Thus, those which are least soluble have, naturally,
no taste, do not visibly affect litmus, do not conduct the electric
current very well in solution, and are not soapy to the touch or
corrosive towards glass. But they all show the tendency to
double decomposition, in which the group (OH) is transferred, as it
was from NaOH to Cu(0H)2 in the foregoing example.
Uses. — Sodium hydroxide is used in immense quantities along
with fats, in the manufacture of soap. Some bleaching Uquids
are made by saturating it with chlorine. It is employed also in
making many other sodium compounds which are used in the arts.
Exercises. — 1. In the electrolysis of sodium hydroxide, why
is the metal not Uberated as a sohd or as a vapor, but as a liquid?
2. Make molecular equations for the burning of sodium in
chlorine and in oxygen.
172 smith's intermediate chemistry
we have met with. But there are well-known acids correspond-
ing to them, namely, chloric acid HCIO3, hydrogen sulphide H2S,
hydrogen peroxide H2O2, and carbonic acid*H2C03.
In general, then, all positive radicals combine with OH to
give bases, all negative radicals combine with H to give acids.
In general, also, each positive radical will combine with any
negative radical to give a salt. In a few exceptional cases only,
the compound cannot be formed, presumably because it is un-
stable under ordinary conditions.
Reaction Formulae. — The formulae of acids, bases, and salts
are written in a uniform manner to show the behavior of the
substances represented, when in solution. Thus, the radical
written first is usually the positive one which, when a solution is
electrolyzed, is attracted by the negative wire, and the negative
radical follows. Again, in a compound Uke calcium hydroxide,
the formula Ca02H2 would conceal the existence of the hydroxyl
group. So the radicals are written in brackets, with the coefficient
outside — Ca(0H)2. Thus we write also Cu(N03)2, and not
CUN2O6, and (NH4)2S04, not N2H8SO4. These formulae are all
reaction formulae. That is, they indicate, not simply the com-
position, but also the parts into which the compounds decompose,
and from which they are formed in double decompositions (p. 132,
and p. 173, below).
Properties Common to Acids^ Bases and Salts in Solution.
— There are four of these properties, all of which have come up
previously.
1. Displacement. — A simple radical belonging to an acid^
base, or salt in solution can be displaced by another element, and
is thereby obtained in the free state. Thus we have already seen
(p. 51) that hydrogen is Uberated from acids by the addition of
the more active metals:
Zn + 2HC1 -> H2 + ZnCla.
ACIDS, BASES, AND SALTS 173
Exactly the same type of reaction takes place when we add to
the solution of a salt any metal higher up in the activity series
than the positive radical of the salt. For example, zinc will
displace copper from a solution of cupric sulphate or any other
soluble cupric salt:
Zn + CUSO4 -> ZnS04 + Cu.
The copper is obtained as a red precipitate. This principle is
extensively used in the purification of the more valuable metals
at the foot of the activity series (p. 54). Thus copper will displace
silver, and silver will displace gold, from solutions of their re-
spective salts.
Similarly, a simple negative radical can be displaced by a more
active element. Thus the iodide radical in potassium iodide is
displaced by gaseous chlorine, iodine being Kberated:
2KI -f CI2 -> 2KC1 + I2.
2. Double Decomposition. — Several examples of this type
of reaction between acids, bases and salts in solution have already
been discussed (pp. 126, 132, 166 and 168). In fact, whenever
two solutions of such substances, which contain no radical in
conmion, are mixed, a double decomposition occurs. Any acid
or base will therefore react with any salt of a different acid or base.
Any acid will give a reaction with any base. Salts containing
no conmion radical will also react in pairs.
Very frequently we obtain instant evidence of such reactions
by the appearance of a precipitate, one of the products formed
being only sUghtly soluble in water. For example:
Salt and acid : AgNOs + HCl -> AgCl i + HNOa (1)
Salt and base: CuCk + 2NaOH -> Cu(0H)2 j + 2NaCl (2)
Acid and base: H2SO4 + Ca(0H)2 -> CaS04 i + 2H2O (3)
Salt and salt: AgNOa + NaCl -> AgCl i + NaNOs (4)
Such precipitations furnish us with very useful tests for estab-
lishing the presence or absence of certain radicals in an unknown
174 smith's intermediate chemistry
substance. For example, a curdy white precipitate of silver
chloride is obtained not only by adding to a solution of silver
nitrate a solution of hydrochloric acid (equation 1) or sodium
chloride (equation 4), but by addition of any solution containing
the radical Ag to any solution containing the radical CL If we
suspect that a solution contains Ag, we add, therefore, a solution
of any chloride. If we suspect that a solution contains CI, we
add a solution of any soluble silver salt. In either case a negative
result is conclusive evidence of the absence of the radical tested for.
If a positive result is obtained, the precipitate must be examined
further, to prove whether or not it is silver chloride.
Even when no precipitate appears, however, some interaction
takes place. Thus a solution obtained by mixing sodium nitrate
and potassium chloride solutions is identical in aU its properties
with a solution obtained by mixing sodium chloride and potassium
nitrate solutions. The appearance of a precipitate is notj there-
fore, an essential feature of a double decomposition. The most
important point for us to notice is that, in all such reactions, each
substance present behaves exactly as if it consisted of two dis-
tinct radicals. We may therefore regard double decompositions
as a simple result of the liberty of radicals to interchange partners.
Normally, this exchange will not be complete. Thus in the
instance cited immediately above we have the reversible reaction:
(Na)(N03) + (K)(C1) ^ (Na)(Cl) + (K)(N08).
Whichever pair of salts we start with, we reach the same result
on mixing their solutions. Every double decomposition is sim-
ilarly reversible in theory, and gives an equilibrium mixture. But
because, in many mixtures, one of the four possible compounds
withdraws from the exchange of partners by disappearing from the
solution either as a gas (p. 126) or as a precipitate (p. 127), the
reaction in such cases becomes practically complete in one direction.
Another significant point may now be noted. Not only does
every acid, base and salt in solution behave, in double decom-
ACmS, BASES, AND SALTS 175
positions, as if it consisted of two distinct radicals^ but it can be
shown to possess two independent sets of properties, one referring
to each radical. Thus a solution of cupric chloride in water
exhibits one set of properties which can be referred directly to the
cupric radical, and which is accordingly not pecuUar to cupric
chloride, but is common to all cupric salts in aqueous solution.
To mention only two of these; (1) the color of the solution when
diluted with water, is blue, and (2) the addition of a base gives a
pale blue, gelatinous precipitate of cupric hydroxide (see equation
3 above). The same solution exhibits a second set of properties
which can be referred directly to the chloride radical and which is
consequently conmion to all chlorides. To mention only two of
these again; (1) when the solution is heated with concentrated
sulphuric acid, HCl is evolved (see p. 126), and (2) the addition of
a silver salt gives a precipitate of silver chloride.
The properties of acids (p. 131) are properties of the hydrogen
radical. The properties of bases (p. 167) are properties of the
hydroxyl radical OH.
A radical, then, is an atom, or group of atoms, which behaves as
a distinct unit in double decompositions, and which confers a defi-
nite independent set of properties upon solutions of all acids, bases
or salts of which it forms one constituent. Simple radicals also
behave as separate imits in displacements.
3. Conductivity.-^ Solutions of acids, bases, and salts in
water are all conductors of electricity. Acids, bases and salts
are therefore called " electrolytes." In all cases the solution
is decomposed by the passage of the current. The positive and
negative radicals of which the solute is composed are attracted
to the opposite electrodes. There, unless special circimistances
prohibit (see for example, the electrolysis of sodiimi chloride,
p. 139), they are Kberated. Thus, all acids give hydrogen at the
negative pole, the other radical passing to the positive pole.
Of all the properties they have in conmion, this one of being
176
SMITH'S INTERMEDIATE CHEMISTRY
electrolytes is perhaps the most remarkable. It appears more
surprising when we consider that no substances, other than acids,
bases, and salts, undergo electrolysis in aqueous solution. This
is an exclusive property of these classes of bodies.
The conducting power of solutions can be examined, roughly,
by the apparatus in Fig. 59. The platinum electrodes are con-
nected with a direct-current circuit. The lamp, which is on one
of the wires, serves, by its resistance, to cut down the current.
Its glowing also shows when the Uquid is a conductor, and by
varying brightness indicates roughly the conducting power of the
solution.
A solution of sugar in water
shows no conductance, and
the lamp remains dark. Solu-
tions of acids, bases, and salts
in water enable the lamp to
glow.
We quickly find that differ-
ent solutions, when they con-
duct, do so in different degrees. Solutions of hydrochloric and
nitric acids conduct well. So do solutions of sodium and potas-
sium hydroxides. Salt solutions are practically all good conduc-
tors. But many acids, Uke acetic acid, conduct poorly in aqueous
solution. The same is true of some bases, Uke ammonium hy-
droxide.
Of course, acids, bases, and salts which are only very sKghtly
soluble in water give poorly conducting solutions. But in esti-
mating the conductivity for chemical purposes, we have to take
into consideration the amount dissolved. Thus silver chloride,
being a salt, is an excellent conductor, when we allow for the
extreme diluteness of the solution. A saturated soliUion of silver
chloride, in point of fact, shows a sUghtly higher conductance than
a solution of sodium chloride of the same concentration.
Substances which give solutions with high conducting power
Fia. 59
ACIDS, BASES, AND SALTS 177
are -called strong electrolytes. Substances which give poorly con-
ducting solutions are called weak electrolytes,
4. Vapor Pressures; Boiling'Points; Freezing-'Points, —
We have seen (pp. 11 7-9) that equal numbers of molecules of
different solutes dissolved in equal weights of water normally
depress the vapor pressure, raise the boiling-point, and lower
the freezing-point by constant amoimts. Thus, one molecular
weight of sugar (342 g.) or of glycerine (92 g.) dissolved in 1000
g. of water will raise the boiKng-point from 100° to 100.52°, and
will lower the freezing-point from 0° to —1.86°. But this is
uniformly true only of non-conducting solutions^ in other words,
solutions of substances which are not acids, bases, or salts. Gram-
molecular weights of substances of these three classes, dissolved
in 1000 g. of water, raise the boiUng-point more than 0.52 degrees
and lower the freezing-point by more than 1.86 degrees. We
say they give abnormal elevations of the boiUng-point and ab-
normal lowerings of the freezing-point. Vapor pressure depressions
in such solutions are also ahnormaL
Thus, a solution of 58.46 g. of sodiimi chloride in 1000 g. of water
boils at 100.97°, and freezes at —3.42°. The elevation in the
boiKng-point of the water is 0.97° instead of 0.52°. The depres-
sion of the freezing-point is 3.42° instead of 1.86°. The effect
in each case is nearly twice as great as the normal one. ' In the
same way, a gram-molecular weight of potassium chloride dis-
solved in 1000 g. of water at 20° depresses the vapor pressure
by 0.554 milUmeters, while a gram-molecular weight of mannite
(a normal sugar which gives a non-conducting solution) in the
same amount of water lowers the vapor pressure by only 0.313 mm.
Again the effect is nearly twice the normal.
The only conclusion we can draw from these results is that
nearly twice the normal numbers of solute molecules are present
in such solutions. In other words, not only do we have two inde-
pendent sets of properties exhibited by the constituent radicals
178 smith's intermediate chemistry
in solutions of electrolytes, but these same solutions actually
behave as if the radicals were, to a large extent, uncombined with
each other. It appears as if sodium chloride, for example, is
largely decomposed in aqueous solution into independent sodium
and chloride radicals. Additional evidence in this direction is
suppUed by the behavior of electrolytes of more complex types,
such as sodium sulphate Na2S04 or zinc chloride ZnC^. Both
of these substances, in dilute solution, give boiling-point and
freezing-point changes which are nearly three times the normal,
indicating that they are largely decomposed into their three con-
stituent radicals (2Na and SO4; Zn and 2C1). In the same way,
substances containing four radicals, Uke ferric chloride FeCls,
give effects approaching four times the normal in dilute solution.
One last point remains to be mentioned. It has been noted
above that some acids, Uke acetic acid, and some bases, Uke
ammonium hydroxide, are only poor conductors in solution. Just
these same two classes of substances, it has been found, give only
very sUghtly abnormal changes in the three physical properties
of solutions here under examination. The fuU significance of
this difference in behavior wiU appear in the foUowing chapter.
A Warning. — The reader is urged to keep in .mind the fact
that it is only in solution (and particularly in aqueous solution)
that the special properties of acids, bases, and salts become ap-
parent. Their behavior is often quite different in the absence
of a solvent. If, for example, we mix together dry ammonium
carbonate (NH4)2C03 and partiaUy dehydrated, soUd cupric
nitrate Cu(N03)2, and apply heat, a violent interaction begins.
This interaction is nothing so simple as a double decomposition,
however. An immense cloud of smoke and gas is thrown out of
the tube, and the soUd remaining is either black or reddish, in
parts, according to the proportions of the substances employed.
This residue contains black cupric oxide CuO, and sometimes red
cuprous oxide CU2O. The gas evolved is tinged red by the pres*
ACIDS, BASES, AND SALTS 179
ence of nitrogen peroxide NO2, while a careful analysis would
show that it contained also carbon dioxide CO2, nitrogen N2,
nitrous oxide N2O, water vapor H2O, and perhaps still other prod-
ucts.
The contrast, when the substances are dissolved in water before
being brought in contact, is very great. A pale green precipi-
tate is inunediately produced, which rapidly settles out and proves,
on examination, to be a carbonate of copper. Evaporation of
the solution gives us ammonium nitrate.
(NH4)2C08 + Cu(N03)2 -> CuCOs i + 2NH4NO8.
The reaction is essentially a double decomposition (although,
strictly speaking, the precipitate is not the normal carbonate of
copper, but a basic salt, see p. 192), similar in character to those
already discussed.
The diflferences in properties between dry hydrogen chloride
and hydrochloric add (p. 131) furnish another good example of
the fundamental changes involved in the addition of water as a
solvent. Carefully dried hydrogen chloride has practically none
of the properties of an acid. It does not affect htmus, it does
not conduct the electric current. It is true that it does react
with the more active metals (such as Na or K), hydrogen being
displaced and the chloride of the metal formed. But it is probable
that even this property would disappear if we could succeed in
eliminating the last traces of moistiu-e from oiu* reacting products.
A solution of sulphuric add acts on sodium so vigorously that the
mixtiu-e explodes, but pure hydrogen sulphate H2SO4, when very
carefully dehydrated, is so inert that metaUic sodium floats in it
without the slightest evidence of any interaction.
The rules that have been derived in the preceding sections re-
garding the behavior of acids, bases, and salts in solution will be
found to clear the ground wonderfully in the development of
subsequent chapters, for the substances which we shall meet with,
that are not included in the above groups, are very few indeed.
180 smith's intermediate chemistry
Furthermore, since we naturally prefer, when we carry out a
reaction either on a small scale in the laboratory or on a large
scale in chemical industry, to have it under our control as much
as possible, we contrive to perform most chemical changes with
our reacting materials in solution. By this means we are enabled,
in general, to forecast the course of the reaction accurately, to
arrange the conditions so that the reaction will be sure to proceed
smoothly, and to isolate without much difficulty those products
which we desire to obtain in a pure state. In the absence of a
solvent a reaction is much more Uable to get out of hand, perhaps
explosively, with the result that the proper conditions of the
process are not maintained, and we obtain either a poor yield or
an impure product.
Exercises. — 1. By what experiments should you determine
which were the radicals in substances of the following composition :
Cu(N03)2, CaCOa, NH4Br, NHJ, KCIO4 and KMn04.
There are always two ways, and usually three, of determining
the radicals — what are they (pp. 171-2)?
2. How should you determine whether a given substance were
an acid, base, or salt (electrolyte), or not?
3. Make equations showing the interactions of solutions of
aluminium chloride AICI3 and of cupric sulphate CUSO4 with
sodium hydroxide (p. 168). Name the products.
4. Make equations showing the interactions of solutions of
zinc chloride ZnCl2 and of ferric chloride FeCls.with silver sulphate
Ag2S04. Name the products.
5. What is the action of metaUic zinc on solutions of sodium
chloride, lead nitrate Pb(N08)2, silver sulphate Ag2S04, mag-
nesium sulphate MgS04, mercuric chloride HgCl2?
CHAPTER XVI
IONIZATION
We have learned in the preceding chapter that acids, bases,
and salts in solution exhibit two independent sets of properties,
one of which can be referred to the positive, the other to the
negative radicah We have seen that the passage of an electric
current through such a solution decomposes it, the positive radical
of the electrolyte proceeding to the negative electrode, the nega-
tive radical to the positive electrode. Finally, we have found that
the abnormaKties in certain physical properties of these solutions
seem to indicate that the radicals actually exist, to a large extent,
in an uncombined state in the solution.
Upon these facts a hypothesis has been based which has proved
of considerable assistance in explaining the special properties of
conducting solutions. The fundamental idea of this hypothesis,
which was first definitely advanced by the Swedish chemist Ar-
rhenius in 1887, is now accepted as an estabUshed fact. Our
knowledge of the field of conducting solutions is still, however,
very imperfect, and many important deductions from the hy-
pothesis of Arrhenius remain matters of conjecture and dispute.
The Ionic Hypothesis, — In this hypothesis of Arrhenius
it is assimied that the molecules of an electrolyte, such as hy-
drogen chloride, are largely broken up in solution into their con-
stituent radicals, each radical being electrically charged. These
charged radicals have been called ions, and the hypothesis is
hence known as the ionic hypothesis. A solution of hydrogen
chloride in water is supposed therefore to consist of two parts:
(1) an imdissociated part, made up of hydrogen chloride molecules
HCl; (2) a dissociated or ionized part, made up of equal numbers of
181
182 smith's intermediate chemistry
hydrogen atoms carrying a positive charge (H+) and chlorine
atoms carrying a negative charge (CI~). That part of the hydro-
gen chloride which remains undissociated in solution is under-
stood to be inactive. It plays no part in the conduction of the
current, it exerts a normal effect on the physical properties of the
solution (vapor pressure, boihng-point, freezing-point changes),
it does not convey to the solution any of the properties of an add,
it takes no direct share in displacements or double decompositions.
That part of the hydrogen chloride, on the other hand, which
breaks up into hydrogen ions H+ and chlorine ions Cl~ is active.
All of the special properties of the solution may therefore be
referred to these ions, as we shall see below.
Ionic Equations. — Of coimse, there is an equilibriimi between
the undissociated and dissociated parts of an electrolyte in solu-
tion, and this equihbrium may be represented by the ionic equa-
tion:
HCl ;=± H+ + C1-.
Ionization is, therefore, a kind of decomposition or rather
dissociation, and is in every sense a chemical change. In ionic
equations the charges upon the ions must be shovm, as they are
essential parts of the ionic substances.
NaCl ^ Na+ + CI" FeCls ^ Fe+++ + SCI"
Ca(N08)2 ^ Ca++ + 2NO3- Na^SO* ^ 2Na+ + S04=.
The number of positive charges must equal the number of negative
charges. This is proved most simply by the fact that the solution
of a substance like ferric chloride (FeCla) is electrically neutral,
as a whole. Thus some ions carry one charge, Uke Na+ and CI"
and H+, others two charges, Uke Ca++ and S04~, and so forth.
An ion is an atom, or group of atoms, carrying an electric charge,
or a number of such charges.
The positive ions are called the cations, since they move toward
s^L.^,
Q 'S
It
!
+
H
0©
IONIZATION 183
the negative electrode, or cathode. The negative ions are the
anions, and move toward the positive electrode, or anode.
We can usually tell which is the positive radical in a formula
because it generally consists of one atom of a metaUic element
(K+, Cu"'^", etc.) or of hydrogen. The negatiye radical may con-
tain a metal, hke the Mn in K^MnOi), but always along with a
non-metal hke oxygen.
Ions und Electrolysis. — Let us first attempt to imderstand
the phenomena of electrolysis in the Ught of the ionic hypothesis.
A solution of HCl in water contains three kinds of solute; im-
dissociated HCl molecules, posi-
tively charged hydrogen ions H+, - — ^^^""^ - ^
and negatively charged chlorine
ions CI"". These are scattered
indiscriminately throughout the
solution, as indicated in Fig. 60
(only the two ions are there
shown, the imdissociated part is
omitted since it plays no direct
part in the electrolysis). Equihbrium is kept up by continual
dissociation and recombination, according to the equation
HCl ^ H+* + CI-
As soon as the circuit is completed, all the ions in the solution
begin to migrate towards their proper electrodes. The positively-
charged hydrogen ions are attracted towards the negative elec-
trode, the negatively-charged chlorine ions are attracted towards
the positive electrode. Two orderly processions of ions, moving
in opposite directions, proceed therefore through the solution.
Ions (Greek, going) derive their name from this fact.
The rest is easily understood. When a positive ion reaches
and touches the negative electrode, its positive charge of elec-
tricity is neutraUzed, and the result is an ordinary atom of hydro-
gen. The atoms of free hydrogen unite to give molecules (H2)
III.
T
Fig. 60
184 smith's intermediate chemistry
and these form bubbles of the gas. Shnultaneously the negative
ions are discharged at the positive electrode, and the atoms of
free chlorine unite to give molecules (CI2).
Meanwhile, in the body of the solution, the departure of the
ions has disturbed the equiUbriimi HCl ;:± H"*" + Cl~. The undis-
sociated part HCl therefore continues to break up, attempting to
re-estabKsh equihbrium, until the electrolysis is complete.
By electrolysis of a solution of hydrochloric acid, therefore,
we obtain hydrogen and chlorine. With some electrolytes the
course of events is not so simple, secondary reactions taking place
at the electrodes. Thus when we pass a current through a solu-
tion of sodium chloride, we obtain chlorine at the positive elec-
trode, but hydrogen instead of sodium is hberated at the negative
electrode (see p. 139). In the same way, when we electrolyze a
solution of cupric sulphate, metaUic copper is deposited on the
negative electrode, but the radical SO4 cannot exist in the .free
state, and reacts with the water present to Uberate oxygen, ac-
cording to the equation:
2SO4 + 2H2O "^ 2H2SO4 + O2.
Even when we do not actually isolate the free radicals of an
electrolyte by electrolysis, however, we can show that they have
migrated with the current in the usual way by the fact that they
collect around the electrodes. Thus, in the electrolysis of so-
diimi chloride, the solution around the negative electrode becomes
alkaline^ owing to accumulation of sodium hydroxide. Similarly,
in the electrolysis of copper sulphate, the solution around the
positive electrode becomes add, owing to accumulation of sulphuric
acid.
Ions and Displacement. — When a metal acts upon a dilute
acid, and hydrogen is hberated, it is the ions alone that are di-
rectly concerned in the mechanism of the action. Thus the
equation for the action of zinc on hydrochloric acid: Zn + 2HC1 — >
IONIZATION 185
ZnCU + H2 1 iiiay be written:
Zn + 2H+ + 2Cl-->Zn++ + 2C1- + Ha t
or:
Zn + 2H+ -> Zn-H- + H2 1-
From this it appears (see however p. 195) that the action simply
consists of a transfer of positive charges from hydrogen ions to
atomic zinc, free hydrogen being liberated and zinc ion going into
solution. Similarly the action of zinc on a solution of a copper
salt may now be written:
Zn + Cu-H- -> Zn-H- + Cu j.
The activity series of the metals (p. 54) expresses, therefore,
the order of their preference for assuming the ionic state.
Ions and Double Decomposition, — The mechanism of reac-
tions of this type also becomes much clearer when we write the
equations in the ionic form. Thus, for the precipitation of silver
chloride by the action of hydrochloric acid on silver nitrate solu-
tion, we have: —
AgNOs ;=± Ag+ + NO3-
HCl ^ CI- + H+
11 11
AgCl i HNO3.
In the mixed solutions, we have four ions in quantity, Ag+
and NO3- from the AgNOs, H+ and CI" from the HCl. Before
mixing, these were in- equilibrium with xmdissociated AgNOs and
undissociated HCl respectively. But as soon as we bring all
four ions into the same solution, we furnish them with the op-
portimity of combining with other partners, Ag+ with CI" and H+
with NO3-. Undissociated AgCl and undissociated HNOs are
also present, therefore, in the mixed solutions, each substance in
equilibriimi with its respective pair of ions. (This is conveniently
indicated by the method of writing the equations which is em-
ployed above. The student should copy this method in analyzing
186 smith's intermediate chemistry
all double decompositions.) The mixture contains, then, no
fewer than eight different solutes, four ionic and four molecular,
and four reversible reactions control the equilibrium relation-
ships between them.
Now it so happens that one of the new solutes, AgCl, is prac-
tically insoluble in water. Unless the solutions are excessively
dilute, therefore, silver chloride separates out from the solution
as a predpitaie. This precipitation disturbs the existing equihbria
by withdrawing practically all silver ions Ag+ and chloride ions
Cl~ from the solution. Undissociated AgNOs and HCl continue
to break up, in an attempt to re-establish the equilibria, imtil
they also are practically eliminated, and only hydrogen ions H+
and nitrate ions NOa", in equiUbrium with imdissociated HNO3,
are left in quantity in the solution.
The reader may feel, at this point, that the previous method of
writing the reaction (p. 131) :
AgNOs + HCl -> AgCU + HNO3
expresses all this more simply, without the need of worrying about
ions. But he must note that any silver salt added to any chloride
in solution also gives a precipitate of silver chloride. This can
be predicted immediately by ionic equations; any two solutions
containing Ag"*" and Cl~ respectively must give a precipitate of
AgCl when mixed. Without the use of ionic equations, however,
we should have no guide to the course of any such reaction;
we should have to try it out for every mixture, memorize what
happened in each case, and then attempt to draw up some em-
pirical rules to assist us in retaining our tremendous mass of
isolated facts. The use of the ionic hypothesis obviates all this.
We do not require to remember all the properties of hundreds of
different acids, bases and salts in solution singly and mixed.
We know that each one of these substances exhibits two sets of
properties in solution, and that one set can be referred to the
positive ion, one set to the negative ion. If, therefore, we learn
IONIZATION 187
the properties of a few important positive and negative ions, we
leam, at the same time, the properties in solution of all the elec-
troljrtes of which these ions are constituents. This introduces a
great simplification into the study of chemistry.
To obtain the necessary practice in deahng with double de-
compositions, the student should now return to the section on
these reactions in the preceding chapter (p. 173) and rewrite all
of the equations there given or referred to in full ionic form.
It is of great importance that the principles involved should be
thoroughly understood in each case.
Ions and Conductivity. — Pure water is an exceedingly poor
conductor of electricity. The high conducting power of a solu-
tion of an electrolyte is due, therefore, to the ions present. The
actual conductance of any given solution will depend on the
number of ions between the electrodes and the rate at which they
move. The more numerous the ions are, and the more rapidly
they migrate towards the oppositely-charged electrodes, the
greater wiU be the number of discharges taking place per second
upon each electrode.
The rate at which the ions move, under given conditions,
has been carefully determined by methods which the student will
find discussed in text-books on physical chemistry. It will suffice
at this stage to give a few results. With electrodes one centi-
meter apart, and with a difference in electrical potential between
the positive and negative electrodes of one volt, the velocities
in centimeters per hour in dilute aqueous solution at 18*^ are as
foUows: H+ 10.8, Na+ 1.26, Ag+ 1.66, OH" 5.6, CI" 2.12, NO3-
1.91. The hydrogen ion is the fastest, the hydroxyl ion holds
the second place.
With the help of these figures we can calculate what is the
extent of ionization of the electrolyte in any solution containing
these ions, hydrogen chloride for example. We have already
learnt (p. 85) that a gram molecular weight of hydrogen chloride
188 smith's intermediate chemistry
contains 6.06 X KF molecules. If ionization in solution were
complete, therefore, a liter of a normal solution (p. 121) of hydro-
chloric acid would contain 6.06 X 10^ hydrogen ions, and the
same number of chlorine ions. These, moving at the rates given
above, would give a definite, calculable conductivity. The
conductivity actually obtained by experiment with a normal
solution of HCl at 18**, however, is only 78 per cent of this cal-
culated value. The conclusion has been drawn that, in normal
solution at 18*^, hydrochloric acid contains only 78 per cent of
free ions, the remaining 22 per cent of solute contributing nothing
towards the conductivity. We may express this conclusion by
writing the reversible ionic dissociation in the form:
(22%) HCl ^ H+ + CI- (78%).
Equihbrium is reached in a normal solution of hydrochloric acid
at 18**, therefore, when 78 per cent of the molecules are broken
up into free ions.
At other concentrations different degrees of dissociation would
be indicated. For example, a. 10 N solution gives only 17 per
cent of the calculated conductivity, a 0.1 iV solution 92 per cent,
a 0.01 N solution 97 per cent. At very high dilutions^ therefore,
the ionization becomes practically complete.
All electrolytes are not ionized to equal extents at the same
concentrations. Thus, a normal solution of acetic acid at 18°
shows only 0.4 per cent of the calculated conductivity for com-
plete ionization. Even a 0.001 N solution is only 12.5 per cent
ionized. A fxmdamental point in the ionic hypothesis of Ar-
rhenius, however, is that ionization always approaches com-
pletion as the solution becomes more and more dilute.
The following table shows the approximate degrees of ioniz-
ation of a number of typical electrolytes in tenth-normal solution
in water at 18*^. In the case of acids and bases containing more
than one displaceable unit of hydrogen or hydroxyl, the kind of
ionization on which the figure is based should be particularly noted.
IONIZATION 189
FRACTION IONIZED IN 0.1 N SOLUTIONS AT 18**
Acids
Nitric acid (H+,NO,-) 0.92
Sulphuric acid (2H+,S04-). .0.61
OxaUc acid (H+,HC04-) . . .0.50
Hydrofluoric acid (H+,F-) 0.085
Carbonic acid (H+jHCO,-) 0.0017
Hydrosulphuric acid (H+,HS-) .0.0007
Phosphoric acid (H+,Ha04-) 0 . 27 | Boric acid (H+,H2B0,-) 0 . 0001
Bases
Sodium hydroxide
(Na+,OH-) 0.91
Potassium hydroxide
(K+,OH-)... 0.91
Barium hydroxide (Ba++,20H-) 0.77
Ammonium hydroxide
(NH4+,0H-) 0.013
Salts
Potassium chloride (K+,C1-) 0.86
Sodium chloride (Na+,C1-) 0.85
Potassium fluoride (K+,F-) 0 . 85
Sodium nitrate (Na+,N08-) 0.83
Silver nitrate CAg+,N08-) . .0.81
Barium chloride (Ba-H-,2C1-) . . .0.76
Sodium sulphate (2Na+,S04") . 0 . 70
Zinc sulphate (Zn"H-,S04") 0.40
Copper sulphate (Cu-h-,S04-).. .0.40
Mercuric c&oride (Hg++,2C1-)<0.01
lonisation and Chemical Activity. — From a consideration
of the above figures and the results given by other acids, bases
and salts, the following important conclusions may be drawn:
1. Salts, with the exception of a few mercuric salts, are all
extensively ionized in 0.1 iV aqueous solution. The salts which
ionize most simply show the greatest degree of ionization. Com-
pare, for example, the series sodium chloride, sodiimi sulphate,
copper sulphate.
2. Acids show the most extreme differences in their degrees
of ionization. That is to say, equivalent solutions contain very
different concentrations of hydrogen ion. Since their activity
as acids depends upon this substance (p. 191), it follows that
acids will exhibit very marked differences in chemical activity
(for example, in their action on metals). In fact, they may be
divided roughly into three classes :
(a) Strong dcids, such as hydrochloric acid, nitric acid, sul-
phuric acid. These substances are highly ionized in 0.1 iV solu-
tion. Their solutions conduct the electric current excellently,
and are chemically most active.
190 smith's intermediate chemistry
(b) Transition ddds, such as phosphoric acid, hydrofluoric
acid. These substances are slightly ionized in 0.1 iV solution.
Their solutions conduct fairly well, and show moderate activity.
(c) Weak adds, such as acetic acid, carbonic acid, boric acid.
These substances are scarcely ionized at all in 0.1 iV solution.
Their solutions conduct the current very poorly, and exhibit
little activity as acids.
3. Bases also show very extreme differences in their degrees
of ionization. We have two main classes — strong hoses, such as
sodiimi hydroxide, bariimi hydroxide; and weak hoses, such as
ammonium hydroxide. Between these fall certain transition
hoses, such as silver hydroxide (which is only very slightly soluble,
but which is noticeably stronger than anmionium hydroxide)
and some organic derivatives of anmionium hydroxide.
4. Water itself is both an exceedingly weak acid and an ex-
ceedingly weak base. It breaks up into the two ions^H+ and 0H~
to a very minute extent indeed. At ordinary temperature the
fraction ionized is less than 0.000,000,002. Pure water conse-
quently conducts the electric current practically not at all. The
ionization of water, however, is a factor of vital importance in
the explanation of certain reactions, as we shall see later.
Ions and Vapor'Pressurcy Boiling'Point and Freezing'
Point Abnormalities, — The cause of the abnormal changes in
these three physical properties, exhibited by solutions of electro-
lytes, will now be evident. A solution of sodium chloride containing
one-gram molecular weight of salt in 1000 g. of water gives almost
twice the calcvlaied effect in each case, because almost all of the
solute is present as Na"*" and Cl~, instead of as undissociated NaCl.
The number of solute molecules present is almost doubled by
ionization. Similarly a dilute solution of ferric chloride shows
nearly four times the normal change in freezing-point lowering,
for example, owing to the fact that the majority of FeCU mole-
cules are broken up into four ions, Fe"*"*^ and 3C1~.
IONIZATION 191
Solutions of weak electrolytes, such as acetic acid and am-
monium hydroxide, exhibit practically normal results, because
only an insignificant fraction of the solute is broken up into ions.
Careful measurements show a close agreement in respect to
extents of ionization, as determined by the two independent meth-
ods, conductivity ratio and freezing-point depression, through-
out the whole list of hundreds of electrolytes. Minor divergences
exist in some cases which have not yet been entirely accoxmted for,
but the agreement in general is so remarkably close that it can-
not be a chance coincidence. We have here, in fact, very strong
confirmation of the truth of the ionic hypothesis.
The Properties of Acids, — The properties of adds (p. 131)
are now seen to be properties of hydrogen ion H+. An add is a
substance which contains hydrogen as a positive radical and,
in solution, gives hydrogen ion. Strictly speaking, only the
conducting solutions of such substances are acids, but for con-
venience we extend the term sometimes to include the pure sub-
stances. For example, HNO3 is usually called nitric acid, not
hydrogen nitrate.
Many substances, such as sugar C12H22O11, contain hydrogen,
but their solutions lack all of the properties of hydrogen ion.
They therefore do not contain hydrogen as a radical, and are not
acids.
A strong add is one which is highly ionized in solution, and
therefore shows the properties of hydrogen ion very markedly.
A weak add is one which is very httle ionized in solution, and
consequently exhibits the properties of hydrogen ion only feebly.
Solutions of very weak acids (such as boric acid) scarcely afifect
blue Utmus. Water, a still weaker acid, contains just as much
hydroxyl ion OH" as hydrogen ion £[+, and does not change the
color of either blue or reddened htmus. Its acid properties are
still evident, however, in its action on the most active metals,
such as sodium:
Na + 2H+->2Na+ + H2 T-
192 smith's intermediate chemistry
The Properties of Bases. — The properties of haseB (p. 167)
are properties of hydroxyl ion 0H~. A hose is a substance which
contains hydroxyl as a negative radical and, in solution, gives
hydroxyl ion.
Many substances, such as ethyl alcohol C2H5OH, contain the
hydroxyl group, but their solutions exhibit none of the charac-
teristic properties of a base. They are not bases, since they do
not contain hydroxyl as an ionizing radical.
A stroTig hose is one which is highly ionized in solution, and
therefore shows the properties of hydroxyl ion very markedly.
It will be observed that the most active bases (alkalies) are the
hydroxides of those metals (K, Na, Ba) which come first on the
activity Ust (p. 54). Weak hoses j like ammonium hydroxide and
copper hydroxide, are little ionized in solution, and exhibit the
properties of hydroxyl ion only feebly. Water is an exceedingly
weak base.
The Properties of Salts. — Salts are substances which contain
a positive ionizing radical other than hydrogen, combined with
a negative ionizing radical, other than hydroxyl. The properties
of a salt in solution are the properties of its two ions (see p. 175).
Some salts do indeed contain hydrogen or hydroxyl as a radical,
but always in addition to two other radicals. Thus sodiimi-hydro-
gen sulphate (p. 126) gives H+ in solution, as well as Na+ and
SO4". This is an a^ salt. The precipitate actually obtained
by the action of a soluble carbonate on a cupric salt (p. 179)
has the composition Cu2(OH)2C08. It is a ba^c saU.
Neutralization. — This is a special case of double decom-
position between electrol3rtes in solution, the reacting substances
being an acid and a base. The products of the reaction are a
salt and water. Thus the mixture of equivalent amounts of
hydrochloric acid and sodium hydroxide solutions gives a reaction
which is complete, although no substance concerned in the reaction
IONIZATION 193
has escaped either as a precipitate or as a gas. All of the peculiar
properties of the original components (such as their action on
Utmus) disappear, and we are left with a solution of common salt.
The explanation, according to the ionic hypothesis, is very-
simple. Before mixing the two solutions, we have our acid and
base almost entirely broken up into the four ions H+, Cl~, Na+
and 0H~. But as soon as mixture is effected, practically all
H+ and 0H~ ions withdraw from the solution, combining to form
water. The minute amoxmts of undissociated HCl and NaOH
continue to break up, attempting to regain equiUbrium with
their respective ions, until they are both eliminated, and only
chloride ions Cl~ and sodium ions Na+, in equihbrium with un-
dissociated NaCl, remain in the solution.
The reaction may be written in expanded ionic form as follows:
HCl ^ H+ + Cl-
NaOH ;=i OH- + Na+
1 11
H2O NaCl
The ionization of water is so minute that it may be neglected.
When either the acid or the base employed is weak, however,
the tendency of water to break up into H+ and OH" must be
taken into consideration (see hydrolysis, p. 369).
We may note here a third method of driving reversible reactions
between electrolytes to completion. Besides removing one
product as a precipitate or as a gas (see pp. 126-7), we may ar-
range the conditions so that one of the substances formed is prac-
tically non-ionized. This is a procedure which is very frequently
employed by chemists in carrying out reactions.
The Part Played by the Solvent in Ionization. — So far,
we have regarded water as playing merely a physical role in the
ionization of electrolytes. The ions in a solution of hydrogen
chloride, for example, have been regarded as H+ and Cl~, the
194 smith's intermediate chemistry
solvent breaking up the molecule HCl in some way, but not being
itself directly concerned in the reaction.
This view, although suflScient for many purposes, will not
stand strict investigation. Thus we know by experiment that
the ions, when they migrate with the electric current in elec-
trolysis, carry water with them. This indicates that ions are
hydrated. The solvent is therefore chemically active in ioniza-
tion.
Recent work, indeed, suggests that the distinction drawn
between solvent and solute in explaining ionization phenomena
is entirely misleading. Pure liquefied hydrogen chloride is prac-
tically a non-conductor, just like pure water. It is quite an
arbitrary procedure, therefore, to ascribe all of the conducting
power of a mixture of hydrogen chloride and water to the former
substance, and to regard the latter as quite inert. It would be
more logical to consider both components of the solution as equally
concerned in ionization. Now experiment shows that extensive
ionization in solution always accompanies extensive compound
formation on admixture. When no interaction at all between
the components occurs, the solution is non-conducting. All
strong acids, for example, give hydrates with water which are
suflSciently stable to be isolated in the solid state. No very weak
acids give isolable hydrates. Those mercuric salts which are
highly ionized, as mercuric nitrate Hg(N03)2, all yield hydrates,
such as Hg(N08)2,8H20. Those which are only sKghtly ionized
are aU non-hydrated. It is possible, therefore, to regard ion-
ization as due to the formation of solvent-solute complexes. The
attractive forces between the constituent groups in such complexes
would be considerably weaker than in the simpler molecules of
the two components, and disintegration into oppositely-charged
radicals could occur much more readily.
It should be added that ionization is not restricted to solu-
tions of electrolytes in water. Many other solvents, such as liquid
ammonia NHs, formic acid H.COOH, ethyl alcohol C2HfiOH.
IONIZATION 196
dissolve many electrolytes to give solutions of excellent con-
ducting power. Water, however, is the solvent most commonly-
used in chemical operations, and other ionizing solvents need
not be considered at this stage.
Ions and Electrons. — The question may be asked: Whence
do the ions obtain their electric charges? A brief answer to this
question may be attempted here, although for a clear reahzation
of its significance a knowledge of the subject-matter of the final
chapter (pp. 552-4) is necessary.
Matter is electrical in its ultimate nature, and the atoms of
all elements are more or less complex aggregates of positive and
negative electrical units. The positive units (protons) constitute
the main mass of the core or nucleus of the atom, and are fixed
therein, except in radioactive disintegrations. The outermost shell
of the atom consists of a nimiber of negative units (electrons), which
are less rigidly held. The atom as a whole, of course, is electrically
neutral. The hydrogen atom, to choose the simplest example, is
made up of a single proton and a single electron. The structure
of the atoms of other elements is, of course, more complex, but
all possess, in their outermost shell, a definite small number of
electrons, which are relatively loosely held. When two atoms of
different elements combine, it may happen that an electron (or
a nimiber of electrons) will pass from one atom to the other.
Atoms which lose electrons in this way become positive radicals,
the departure of an electron leaving the atom as a whole elec-
trically positive. Atoms which gain electrons become negative
radicals, arrival of an electron making the atom as a whole elec-
trically negative. Under normal circmnstances, the attractive
forces between such oppositely-charged radicals will be sufficient,
in most cases, to bind them firmly together as electrically neutral
molecules. If we weaken these forces, however, as we undoubtedly
do when we dissolve an electroljrte in a solvent such as water,
separation of the bound radicals into free positive and negative ions
196 smith's intermediate chemistry
can be effected much more readily, and extensive ionization may
result.
A positive ion, therefore, is a free atom, or group of atoms, which
has lost an electron, or a number of electrons, such as Na+,
(NH4)+, Zn++. A negative ion is a free atom, or group of atoms,
which has gained an electron, or a number of electrons, such as
C1-, (NOs)-, S-.
Some Possible Misunderstandings. — Before we close the
chapter, it will be profitable to anticipate some difficulties into
which the reader may fall. If the ionic hypothesis is not properly
understood, it appears to conflict so strongly with what the stu-
dent has learnt in previous chapters that he is apt to become
hopelessly confused. The following points of possible misunder-
standing and the explanations appended should therefore be read
through very carefully. For convenience of illustration, sodium
chloride is taken as a typical electroljrte in the questions and
answers Usted below. The student should test his knowledge of
the subject by substituting other electrolytes.
1. If sodium chloride is broken up in aqueous solution into
sodium and chlorine, why do we not find any of the properties of
sodium or of chlorine exhibited by the solution? — This has always
been a very common misapprehension of the ionic theory. Many
prominent chemists never could understand how sodium (a m£tal
which acts vigorously on water) and chlorine (an obnoxious gas)
could exist side by side in a solution of sodiiun chloride without
immediately notifying us of their presence by characteristic
reactions. The answer is that free sodium and free chlorine do
not exist in sodium chloride solution. The ionic hypothesis has
never stated that they do. What it does state is that sodium
ion Na"^ and chloride ion CI" are present in the solution. These
are entirely different substances from atomic sodium Na and mol-
ecular chlorine CI2. The electric charges on the ions change
their properties completely. There is no more reason why they
should behave like free sodium and free chlorine than there is
IONIZATION 197
for crystals of common salt to behave like a mixture of sodium and
of chlorine. Metallic sodiimi Na reacts with water to form a
solution of sodimn hydroxide. In sodium chloride solution,
however, the ionic sodium Na"*" is already in the same state as it
is in sodium hydroxide solution, and is in no need of trying to
enter that state.
2. Salt is a very stable substance. The imion of sodiimi and
chlorine evolves a great deal of heat. A great deal of work will
be required, therefore, to decompose sodiimi chloride. How can
mere addition of water break it up? — We have here the same
misunderstanding in another form. It is true that it would be
very difficult to decompose sodium chloride into free sodium and
free chlorine, but its dissociation into sodium ion and chloride ion
is an entirely different question. As a matter of fact, the heat
of ionization is extremely small. Sodium chloride is stable only
in the soUd state. In solution, it reacts with very great facility
witb many other electrolytes.
3. Why do not the ions Na"^ and Cl~ recombine at once, in re-
sponse to the attractions of their charges? — The answer is that
they do combine. The tendency towards combination is, how-
ever, opposed by the tendency of undissociated NaCl (the at-
tractive forces between the bound radicals of which are weakened
in solution) to decompose into free ions. An equihbrium be-
tween the two tendencies is, therefore, set up, which we may
express by the reversible reaction NaCl ^ Na+ + Cl~.
4. Why can we not separate sodium ions from chloride ions
in a solution of sodium chloride before we pass a current through
the solution? Does not this show that it is the electric current
which breaks up the sodium chloride? — The charges on the ions
are not derived from the electric current. Free sodium ions and
free chloride ions are present in the solution the instant the salt
is dissolved, whether a current is passing or not. Before we pass
a current through the solution, however, any portion of it of
sensible magnitude contains just as many sodium ions as chlo-
198 smith's intermediate chemistry
ride ions. By making use of the fact that the chloride ion diffuses
more rapidly than the sodium ion (into a layer of pure water, for
example, carefully poured over the solution) we can bring about
a slight separation of the two ions, the water layer becoming
negatively charged and the solution positively. The passage of
the current does not cause ionization, it merely makes its existence
more obvious, efifectually separating the ions by forcing them
to migrate in different directions towards the oppositely-charged
electrodes.
Exercises, — 1. Which are the anions and which the cations in
the substances whose formulae are given on p. 189?
2. Using the models given in p. 182, make the ionic equations
representing the ionization of all the acids, bases, and salts, the
formulae of which are given on p. 170.
3. Make an ionic equation (p. 185) for the displacement:
(a) of hydrogen from dilute hydrochloric acid by magnesium;
and (b) of copper from cupric sulphate solution by zinc.
4. Rewrite the double decompositions on pp. 166-168 in full
ionic form.
5. Why does a solution of 0.1 iV hydrochloric acid conduct
the current nearly twice as well as a solution of 0.1 iV sodium
hydroxide, and nearly four times as well as a solution of 0.1 iV
sodium chloride?
6. From the results given on p. 177, calculate the degree of
ionization of sodium chloride in a solution containing 1 gr. mol.
wt. NaCl to 1000 g. water (a) at the boiUng-point, and (b) at
the freezing-point.
7. From the results given on p. 177, calculate the degree of
ionization of potassium chloride in a solution containing 1 gr.
mol. wt. KCl to 1000 g. water at 18°.
CHAPTER XVII
THE HALOGEN FAMILY
The elements, if we may judge from those studied or mentioned
thus far, may be divided into two classes — the metallic or positive
elements, like sodium, zinc and magnesium, and the non-metaUic
or negative elements, hke oxygen, chlorine and sulphur. The
former give positive ions, such as Na+, Mg"'"^'. The latter give
negative ions, such as Cl~, S". Hydrogen constitutes the single
exception, giving the positive ion H+.
Natural Families of Elements, — We have a simple means
of subdividing within each of these two classes. We can place
together the elements of like chemical behavior. Thus sodium
and potassimn or zinc and magnesium resemble one another
very closely in their reactions. Also, oxygen and sulphur form
one group and chlorine, bromine, iodine and fluorine form an-
other. Groups of this kind are often spoken of as natural fami-
lies of elements. The last group is called the halogen family,
from the Greek for saU-producing, because these elements com-
bine with sodiimi to give substances all resembling conunon
salt. Usually, the elements of one family and their correspond-
ing compounds resemble one another in a number of ways, and
show at the same time a gradation in properties which it is in-
teresting to study.
Bromine Br2
The element was discovered by Balard in 1826 and derives its
name from its offensive odor (Greek, a stench).
Preparation. — The salt deposits and natural salt wells of
Cheshire, of Germany, and of Michigan, West Virginia, Ohio, and
199
200 smith's intermediate chemistry
Connecticut, contain some bromides, along with large quantities
of common salt. When the latter has been largely separated by
evaporation and crystallization, the bromides of sodium and
magnesimn, which are more soluble, collect in the mother Uquor.
The bromine can be Uberated at the positive electrode by elec-
trolysis. But usually a chemical process is employed.
In one process, chlorine gas is dissolved in the Uquor. This
displaces the bromine, and the latter can be distilled out by heat-
ing:
2Br- + CI2 -> 2C1- + Bra.
In another process, oxidation of a bromide by pulverized
manganese dioxide and sulphuric acid is employed, and this
method can be used in the laboratory (Fig. 81, p. 308).
The manganese dioxide is reduced to manganous sulphate
(compare p. 142), its oxygen combining with hydrogen from
H2SO4 to form water. The sodimn bromide used is converted
to sodium hydrogen sulphate (compare p. 126) and bromine is
libera l)ed.
Skeleton: NaBr + MnOj + H2SO4 -^ Brj + NaHSOi + MnSO* +
H2O.
Balanced: 2NaBr + Mn02 + 3H2SO4 -> Br2 + 2NaHS04 +MnS04
+ 2H2O.
Physical Properties. — Bromine is a liquid of a deep red-
brown color and the vapor, of the same color, has a suffocating
odor. It boiU at 69°. It is moderately soluble in water, giving
a 3.2 per cent solution (bromine- water), and is very soluble in
carbon disulphide. The density of the vapor gives it the formula
Br2. Great care must be used in handhng bromine, as, when
spilt upon the skin, it kills the tissues and the sore is very Uable
to become infected.
Treatment of Burns. — Bm-ns made by bromine or strong
acids should be washed instantly with water and then with bi-
THE HALOGEN FAMILT 201
carbonate of Boda solution, and covered thickly with vaseline, or a
salve of boric add in lanoline, to protect them from infection.
Chemical Properties. — A jet of burning hydrogen will con-
tinue to bum in bromine vapor, giving hydrogen bromide, a gas
which fumes in moist air like hydrogen chloride :
Hi + Brj -» 2HBr.
Many of the metals, when thrown in the form of powder, leaf,
or foil, into bromine vapor, combine directly, giving bromides.
The action is similar to that with chlorine, but less vigorous.
Hydrogen Bromide HBr, Preparation. — Hydrogen and
bromine vapor unite much less readily than hydrogen and chlo-
rine. A stream of pure hydrogen bromide is most easily made
by moistening red phosphorus with water, and allowing bromine
to fall drop by drop into the paste (Fig. 61). To absorb the
bromine vapor, carried by the gas, the latter is passed through
a U-tube containing dry red phos-
phorus mixed with broken glass or
beads:
2P + 3Br, -* 2PBr|.
PBn + 3H,0 -t 3HBr T + H,POa.
The bromine forms phosphorus tribro-
mide, which is immediately decom-
posed by the water. The phosphor-
ous acid HjPOj remains, dissolved in
the water, in the flask. The gas can
be collected by upward displacement Fia, 61
of air.
It might seem that a simpler action would be that of sulphuric
acid upon a bromide (compare p. 126) :
H,SO, -I- KBr -» KHSO. + HBr T •
202 smith's intermediate chemistry
This action does take place, but the hydrogen bromide formed,
being less stable than HCl, is oxidized rapidly by the concen-
trated sulphuric acid, so that, although some of the gas escapes
oxidation, it is mixed with much free bromine and sulphur dioxide:
H2SO4 + 2HBr -^ SO2 + 2H2O + Br2. This action, indeed, en-
ables us to recognize a bromide, by the color of the bromine vapor
and the fuming of the hydrogen bromide produced.
Properties of Hydrogen Bromide. — The gas, like hydrogen
chloride, is colorless, and has an irritating effect when breathed.
It is extremely soluble in water, and fmnes in moist air, giving a
fog of HBr dissolved in water.
Chemically, hydrogen bromide is stable, though not so much so
as hydrogen chloride. Its aqueous solution is an active add. As
such, it gives double decomposition with bases and salts. Thus,
with a salt of silver, we get a cream-colored precipitate of in-
soluble silver bromide:
AgNOa + HBr -^ AgBr i + HNOa.
Chlorine-water, added to a solution of any bromide, displaces
the bromine, which may be recognized by its brown color (test for
a bromide) :
CI2 + 2Br- -^ 2C1- + Br2.
A few drops of carbon disulphide, shaken with the mixture, will
settle to the bottom, carrying the brown bromine with it in a more
concentrated, easily recognizable form.
Uses of Compounds of Bromine. — Bromine is manu-
factured in large amounts in Germany and in the United States.
It is employed to make potassimn bromide, and other bromides.
These are utiUzed in medicine, and to precipitate silver bromide
in the manufacture of photographic films and plates.
THE HALOGEN FAMILY 203
Iodine I2
Sources, — Iodine was formerly all obtained from seaweed
(kelp), certain species of which use the traces of organic com-
pomids of iodine in sea water as part of their food. The dried
seaweed is carbonized in retorts, and sodium iodide remains in the
residue, along with much sodium carbonate and carbon. In an
improved process the iodine compounds are dissolved out of the
kelp, and from the latter a sort of gelatin, named algin, is extracted.
The greater part of our supply of iodine is at present obtained
from sodium iodate NalOa, which forms about 0.2 per cent of
crude Chile saltpeter.
Preparation. — The processes for obtaining iodine from an
iodide are precisely the same as those for bromine.
In France chlorine is used to displace the iodine:
Cl2 + 2I-->2Cl- + l2i.
The precipitate of iodine is pressed free from the solution.
In Great Britain the iodide is mixed with manganese dioxide
and sulphuric acid and heated:
3H2SO4 + MnOa + 2NaI -> MnS04 + 2NaHS04 + 2H2O + I2 T •
Iodine vapor condenses upon a cold surface, not to the Uquid,
but directly to the soUd, crystaUine form. Distillation which
gives a soUd product is called sublimation. The crude iodine is
purified by repetition of this process.
Physical Properties. — Iodine forms black, shining crystals.
The vapor, visible even at the ordinary temperature, is violet in
color (hence the name of the substance, from Greek, meaning like
a violet) and has a density corresponding to the formula I2. The
substance is very sUghtly soluble in water. It dissolves, however,
easily in carbon disulphide (violet solution), in alcohol or potas-
sium iodide solution (brown solution), and even in starch, upon
204 smith's intermediate chemistry
which a trace of it confers a strong blue color {test for free iodine,
p. 400). These colors are shown only by free iodine — the
iodides are colorless.
Chemical Properties. — Iodine unites very slowly and in-
completely with hydrogen, giving hydrogen iodide. It combines
readily with pfiosphorus (Pis) and with many of the metals, giving
iodides.
Hydrogen Iodide^ Preparation. — The gas is prepared by
the process used for hydrogen bromide. Red phosphorus and
iodine are mixed, and water is allowed to fall drop by drop upon
the mass (Fig. 26, p. 62) :
2P + 3I2 -> 2PI3.
Pla + 3H2O -► SHI + H3PO8.
The gas is very dense (HI = 1 + 127 = 128 g. per 22.4 Uters,
against 28.96 g. for air) and can be collected by upward displace-
ment of air.
The action of sulphuric acid upon an iodide does not give pure
hydrogen iodide> although the action Nal + H2SO4 -^ NaHS04
+ HI does take place (compare p. 202). Hydrogen iodide, being
much less stable than even hydrogen bromide, is a more active
reducing agent, and reduces the sulphuric acid to hydrogen sul-
phide. The odor of this gas is therefore very conspicuous when
an iodide is moistened with sulphuric acid:
H2SO4 + 8HI -> H2S + 4H2O + 4I2.
The violet vapor of iodine becomes visible if the test-tube is
warmed. A rough test for an iodide is afforded by this action.
Another method of making hydrogen iodide is frequently
employed when a solution of the gas in water is required, and not
the gas itself. Powdered iodine is suspended in water, and hydro-
gen sulphide gas is introduced through a tube in a continuous
stream. The iodine dissolves slowly in the water and acts by
THE HALOGEN FAMILY 205
displacement upon the sulphide ion S", derived from the solu-
tion of H2S in water. Sulphur separates in a fine powder, and
a solution of hydrogen iodide (hydriodic acid) is formed in ac-
cordance with the equation:
2H+ + S- + l2-^2H+ + 21- + S i.
The solution is freed from the deposit of sulphur by filtration, and
may be concentrated to 67 per cent of hydriodic acid by distilling
off the water.
Properties. — Hydrogen iodide is exceedingly soluble in water
and fumes strongly in moist air, giving a fog of HI solution. It
is colorless.
The aqueous solution is an active acid. The iodide-ion I",
which it contains, gives, with any soluble salt of silver, a precipi-
tate of insoluble yellow silver iodide Agl :
AgNOa + HI -► Agl i + HNOs.
Chlorine-water or bromine-water, added to a solution of this
or any other iodide, displaces the iodine:
CI2 + 21- -> 2C1- + I2.
The free iodine, even if present in minute amounts, may be recog-
nized by shaking the Uquid with a few drops of carbon disulphide.
The iodine gives a violet solution in the latter. A still more
deUcate test is the addition of a drop of very thin starch paste,
which gives a deep-blue tint with free iodine. Filter paper dipped
in starch paste and dried can also be used, by touching it with a
drop of the solution containing the free iodine.
Uses of Iodine and Its Compounds. — The alcoholic solution
(tincture of iodine), painted over the skin, reduces swellings and
inflammation. Iodoform CHIa is a soUd used for similar purposes,
lodothyrin is an organic compound found in the hmnan thyroid
gland, as well as that of other animals. An extract of sheeps'
206 smith's intbbmediatb chbmistbt
thyroids (thyroxiii) is admioiatered with remarkable succees in
cases of degeneration caused by abnormally small natural de-
velopment of this gland (cretinism). Potassium iodide ia alao
used in medicine, to cause absorption of blood-clots and effusions
of blood, for example in the eye. Silver iodide is contained in
the coatii^ on phott^aphic plates and films.
Fluorine Fs
Compounds of fiuorine are found in large quantities as minerals,
but the compounds are so stable that the element is very difficult
to liberate. The natural compounds, however, have many im-
portant uses.
Occurrence. — Calcium fluoride CaF* (fluoiite) occurs in
nature in beautiful cubical crystals. Cryolite AlFj,3NaF is used
in the modem manufacture of aluminium
(p. 466). Apatite CaB(P04)*F is a common
constituent of rocks and soils. When slowly
decomposed, by weathering, it furnishes soluble
phosphates. These are absorbed by plants, for
which they are a necessary food.
Preparation of Fluorine. — The element
is obtained by electrolysis of potassium-hydro-
gen fluoride KHFj dissolved in liquefied hydro-
gen fluoride. The electrodes are made of an
alloy of platinum and iridium, with which
fluorine has httle tendency to combine. The
vessel is a U-tube of copper (Fig. 62) and, to
prevent vaporization of the hydrogen fluoride (b.-p. 19.4°), is kept
at —23° to —40° during the operation. Hydrc^en is liberated
at one pole and fluorine at the other.
Properties. — Fluorine is a yellow gaa, with a density greater
than that of air (G.M.Y. weighs 38 g.). It is the most difficult
of the halogens to liquefy (b.-p. —187°).
FlQ. I
THE HALOGEN FAMILY 207
I
Fluorine is the most active of the non-metals. It combines with
all the metalsj but most slowly with platinmn and with gold. In
the preparation of the gas the copper is protected from serious
attack by the layer of fluoride first formed. Fluorine combines
also with hydrogen in the cold and, unUke chlorine, without the
assistance of Ught. It combines with most of the non-metalsj but
not with oxygen, chlorine, nitrogen, or the indifferent gases of the
atmosphere.
With water (vapor or Uquid) fluorine interacts, giving ozone
(see p. 219) and hydrogen fluoride H2F2:
Skeleton: F2 + H2O -^ H2F2 + O3.
Balanced: 3F2 + 3H2O -> 3H2F2 + O3.
Hydrogen Fluoride H2F29 Preparation. — When pulverized
calcium fluoride and concentrated sulphuric acid are placed in a
retort of platinmn or lead and the mixture is warmed, hydrogen
fluoride passes over. The vapor is usually led into water, in
which it is very soluble (hydrofluoric acid) :
CkFi + H2SO4 -^ CaS04 + H2F2.
The acid is kept in bottles of paraffin or rubber, as glass inter-
acts with it rapidly.
Physical Properties. — The vapor of hydrogen fluoride can
be condensed to a colorless liquid boiling at 19.4°. Being very
solvblcj it fiunes strongly in moist air. The vapor density below
40° corresponds with the formula H2F2, but at higher temperatures
gradual dissociation to HF occurs.
Chemical Properties.— The aqueous solution has all the
properties of a transition add (see p. 190). The substance has,
in addition, the remarkable property of acting upon silica Si02
(sand), and siUcates, to give silicon tetrafluoride SiFi (a gas).
Hence it attacks glass, which is a mixture of sodium siUcate
Na^SiOs and calcium siUcate CaSiOa:
Si02 + 2H2F2 -* SiF4 + 2H2O.
CaSiOa + 3H2F2 -^ SiF4 + CaF2 + 3H2O.
208 smith's intermediate chemistry
I
Thus, when glass is covered with melted parafl^ to protect the
surface, and marks or letters are made by removing the paraffin
with a sharp instrument, hydrogen fluoride will decompose the
glass at the parts thus exposed {test for fluorine). In this way the
graduation on thermometer stems and lettering on glass are
frequently made. The vapor gives rough, easily visible depres-
sions, the solution smooth, glossy ones.
On account of this property, hydrofluoric acid is used for re-
moving adhering sand from castings and for cleaning the outsides
of granite and sandstone buildmgs.
The Halogens as a Family
The reader is reconmiended to compare carefully the properties
of the several halogens and their compounds with hydrogen.
It wiU be found that, while very striking similarities exist
throughout the whole halogen family in the case of nearly
every physical and chemical property, there is always a
regular gradation in properties in the order of the atomic
weights, namely F, CI, Br, I. A few examples are noted
below; the student should tabulate the rest himself for his
own convenience.
Color in liquid state: F2 yellow; CI2 orange-yellow; Br2 brown;
I2 deep violet.
Bailing-paint: F2-187°; CI2 -34°; Br2 69°; I2 184°.
Action on hydrogen: F2 very rapid action in cold, without
Ught; CI2 action rapid in cold, only with strong Ught; Br2
action rapid only when heated; I2 action slow and incom-
plete even when heated.
As an aid to the memory, such a table is exceedingly valuable.
But it is really not necessary to attempt to memorize all the prop-
erties of each halogen and of each halogen compound separately .
The properties of bromine, for example, are aU intermediate
between those of chlorine and iodine. Indeed, when bromine
THE HALOGEN FAMILY 209
was first examined by liebig, he thought it was an xmstable com-
pound of chlorine and iodine, and so missed gaining the credit
of its discovery as an element.
The activity of the halogens, as is evident from their action on
hydrogen, decreases in the order of increasing atomic weight.
This is seen also in their displacement reactions in solutions.
Thus we have found that chlorine displaces bromine from bromides
and that bromine displaces iodine from iodides. Fluorine is
able to displace even chlorine from chlorides: 2C1+ + F2 -^ 2F~ +
CI2. When we note also that the halogens displace sulphur from
sulphides (compare p. 205) and that oxygen displaces iodine from
hydriodic acid, we are able to draw up an order of activity for
the non-metals, similar to the activity order for the metals. This
order of activity expresses the order of preference of the non-
metals for assuming the ionic state. Starting from the top, it is
F, CI, Br, O, I, S.
A last point worthy of mention is that some of the properties
of fluorides are pecuUar; Thus hydrogen fluoride at low tem-
peratures has the formula H2F2, and its solution in water is not
highly ionized. We shall see later that it is usual for the first
member of a family of elements or compounds to exhibit a few
peculiarities. In fact we have already noted, at the beginning of
this chapter, one pecuUarity of the first of all the elements, hydro-
gen. Although a non-metal, it gives a 'positive ion H+.
Exercises. — 1. Tabulate the properties of fluorine, chlorine,
bromine and iodine, and of their compounds with hydrogen.
2. How should you distinguish by chemical reactions the
chloride, bromide, iodide, and fluoride, (a) of hydrogen, (6) of
sodium from one another?
3. Write equations for the action, (a) of chlorine upon a solu-
tion of hydrogen sulphide, (6) of bromine upon a solution of
hydrogen iodide, (c) of oxygen upon a solution of hydrogen iodide,
(d) of fluorine upon a solution of hydrogen iodide.
210 SBilTH's INTERMEDIATE CHEMISTRY
4. Why does hydriodic acid, when left in the air, become brown
in color?
5. Make a list of all the acids we have encountered, and note
which are weak and which strong.
6. How should you make potassiimi bromide, starting with, (a)
potassium and bromine, (6) hydrogen bromide, (c) potassium
iodide?
7. Rewrite the equation for the action of hydrobromic acid
on silver nitrate solution in full ionic form.
CHAPTER XVIII
VALENCE
The differing number of charges on different ions has called
our attention vaguely to a subject which must now be explored
and set forth more clearly.
Valence. — The formulae of a number of common compounds,
including some that we have met with, are as follows:
NaCl ZnCk AlCls SnCU
NaBr ZnBr2 AlBrs SnBr4
Nal Znl2 Alia
We observe that one atomic weight of sodium appears to unite
with only one imit of another element, one unit weight of zinc
with only two imits of another element, a unit of aluminimn with
not more than three, and a unit of tin with only four imits.
It seems that an atomic weighl of each element has a fixed car
pacity for combining with not mxrre than a certain number of atomic
weights of other elements.
Other compounds of only two elements have the formulae:
HCl H2O NHa CH4 LisN CCI4.
So far as we may judge from this Umited Ust, CI combines with
only one atom of another element, O with two atoms, N with
three atoms, C with four atoms. Also an atom of hydrogen
combines with not more than one atom of another element,
although it may take more than one atom of hydrogen to satisfy
the atom of that other element (H2O, CH4, etc.).
This limited combining capacity of each kind of atomic weight
(or atom) is called its valence.
211
212 smith's intermediate chemistry
Mcwking the Valence. — UntU we are familiar with their
values in each case, it may be well to mark the valences thus:
Na^ Zn-^ Al°^ Sn'^ CV Bf I' O N"' O^
As we should expect, an atom with the double capacity can
combine with two of the single capacity, or with one of the double
capacity, and so forth. Thus we have compounds of oxygen:
Zn°0 Sn'^02'^ M2'^0^
Briefly stated, the quantities of the two elements which com-
bine must have equal total combining capacities. Thus Sn>^ has
the capacity four, and 02^^ has the total capacity of 2 X 2 or 4:
AI2™ has a total capacity of 2 X 3 (or 6) and so has O3" (3X2
= 6).
The imit of combining capacity of an atomic weight (or atom)
is called a valence. The atomic weights of H and CI are said to
be univalent; those of Zn and O, bivalent; those of Al and N,
trivalent; those of Sn and C, quadrivalent. The highest valence
known is eight.
Valence and Ionic Charges. — Comparison with the formulae
of the ions already given will now show that the valence is equal
to the number of charges on the corresponding ions: H'Gl' gives
H+ + CI- and Zn^CV gives Zn++ + 2C1-. Also, of course, the
total number of each kind of charges (positive and negative)
was equal, just as the total valences of each of two constituents
of each compound are equal.
Valence of Radicals. — What has been said applies to com-
pounds of not more than two elements — so called, binary com-
pounds. We cannot, by inspection, tell the valences in a com-
pound of three or more elements, Uke H2SO4. But, as we have
seen, all electrolytes behave like binary compounds, because they
divide into radicals, which move as wholes from one state of com-
bination to another. Hence we can assign a valence to the radical
VALENCE 213
SO4 as a whole. It is evidently bivalent, H2KSO4)", Zn"(S04)«.
Similarly, in K^NOa)', and in HKNOa)', the NOs is clearly uni-
valent. HaKPOi)^" shows PO4 to be trivalent.
Valence also by Displacement. — In the foregoing instances,
we have learned the valence of an element or radical by studying
its combinations. But, clearly, if an element is displaced from
combination, atoms of equal total valence must take its place.
Thus the action:
Zn + 2HC1 -> ZnCla + H2
shows Zn displacing 2H', and the valence of Zn must therefore be
two. We see that this is the case for, on displacing the 2H, it
combines with 2CP.
Summary. — We may now simi up all these facts by saying:
The valence of an element is a number representing the capacity
of one atomic weight of the element to combine with, or displace,
atomic weights of other elements, the unit of such capacity being
that of one atomic weight of hydrogen or chlorine. To make a
corresponding statement for the valence of a radical, we sub-
stitute, in the foregoing sentence, the word radical for element,
and the word formula-weight for atomic weight.
Application in Making Formuke and Equations. — We
can see at once that the rule of valence will be of great assistance
to us in making formulae and equations. Suppose, for example,
that we bum a piece of aluminium foil in chlorine, and get the
white aluminium chloride. What is its formula? Up to this
point, we should simply have looked for it in a book. And if,
subsequently, we had required the formulae of the oxide and sul-
phate of aluminimn, we should have looked these up separately
also.
But now, all we have to do is to find oiU the valence of aluminium.
Knowing already the valences of CI' and 0° and (804)"^, we have
214 smith's intermediate chemistry
then all the information we require for making the needed formulae.
Suppose we know that the atomic weight of aluminium is trivalent
Al"^^ (see next section). Making the total valences of each half
of the compound aUke, we get the formulae:
AliiiCV, Ala^Os", Al2iiKS04)3'^
When we know the valences of the elements and radicals, we can
make the formula of any required compound.
The reader mvst therefore make a special effort always to learn
the valences of each element and radical, and always to use them
in making formvlce.
The reader must also always check every formula he writes from
memoryy to make siu-e that it is correct. Thus, if he thinks the
formula of zinc nitrate is ZnNOs, he must coimt the valences,
Zn^'CNOa)'. Evidently, the correct formula is Zn(N03)2.
How to Learn the Valence of an Element. — To find out
the valence of an element, we must obtain the formula of one sim-
ple compound of the element, containing another element of known
valence. Thus, what is the valence of carbon? Its oxide is CO2.
The total valence of oxygen here is 2 X 2 = 4. Carbon C^^ is
therefore quadrivalent. Hence its chloride must be CivCU^ (car-
bon tetrachloride), and its compoimd with hydrogen Cr^H4^
(methane, composing a large part of natural gas). When car-
bon combines with a trivalent element, equi-valent amounts of
each element must be used, as in Al4™C3^ (aluminium carbide),
where AI4™ and Ca^ contain 3 X 4, or 12 units of valence each.
Again, when we know the formula of sodium iodide to be Na'I,
or that of hydrogen iodide to be H^I, we infer that iodine is univa-
lent. The formula of sihca (sand) Si02" shows siUcon to be
quadrivalenl, and indicates that the chloride must be SiCU. Simi-
larly the formula of calcium carbonate Ca^COs shows that the
radical CO3, which is common to all carbonates, must be bivalent.
The chemist does not memorize the valences themselves; he
VALENCE 215
recovers the valence of an element or radical, when needed, by
recalling the formula of a substance containing this element or
radical in combination with a more familiar element or radical,
such as CT or H^.
Elements with More than One Valence, — The rule of
valence is not so simple as it has thus far appeared to be. A
number of the elements have more than one valence. In other
words, the capacity of an atomic weight of such an element may
have two (or even more) values, according to the circumstances
under which it is combinmg with other elements.
Thus, antimony is usually trivalent, and gives compounds like
SbCla, Sb203, SbBrs. But it can also form compounds in which it
is quinquivalent, like SbCle. Similarly, iron forms two complete
series of compounds:
Bivalent: FeCk, FeO, Fe(0H)2, FeS04.
Trivalent: FeCU, FeaOs, Fe(0H)3, Fe2(S04)3.
Even the halogens, although uniformly imivalent in their com-
pounds with hydrogen and other positive radicals, show oxygen
compounds of higher valence, such as chlorine dioxide CIO2,
iodine pentoxide tOs. When an element does give more than one
series of compounds, however, we always make a strong point of
this fact, so that it may not be overlooked.
No simple rule, for telling, in advance, which valence will be
used in a given action, can be stated. But the ions Fe++ and
Fe+++, for example, have different properties, and are easily
recognized in practice.
As a rule, an element passes from one form of combination to
another without change of valence. But compounds of elements
like tin or manganese can undergo changes in the course of which
the valence alters. A case of this kind has already been en-
countered in the preparation of chlorine (p. 143).
MnOa + 4HC1 -> 2H2O + MnCk + CI2.
216
smith's intermediate chemistry
The valence of the atomic weight of manganese changes in the
course of this action from 4 to 2. When MnO acts on HCl,
however, manganese is bivalent throughout.
Exceptional Compounds. — A few compoimds will be met
with in which an element shows an exceptional valence. Thus,
nitrogen gives two series of compounds of N^^ and N^. But
there are three oxides, N2O, NO, and NO2, in which the valence of
nitrogen seems to be one, two, and four, respectively. However,
these are single compounds, not belonging to any series, and are the
only compounds of nitrogen showing any of those three valences.
Again, FeO and Fe208 belong to the two regular series of com-
poimds of iron. But there is the magnetic oxide, Fe804, where
the valence of iron appears not to be a whole number, but 8/3 or
2§. In this case the chemist makes the valence regular by sup-
posing the magnetic oxide to be a compoimd of the other two
oxides, and writing its formula, FeO,Fe208.
A List of Valences and Charges. — The following table
contains the valences of some famiUar ions and the commonest
Univalent
Bivalent
Trivalent
Quadrivalent
Na+
K+
H+
(NH4)+
Ca++
Ba++
Mg++
Zn++
Pb-H-
Ni-H-
Co++
Mn-H-
Cu-H- (cupric)
Fe-H- (ferrous)
Hg»+ (mercuric)
Sn-H- (stannous)
0-
r'-
(CO,)-
Oi" (peroxide)
Al-H-i-
Fe+++(ferric)
Cr+++
Sb-i-H-
Bi-H-h
(P04) =
As (AsH,)
B (B,0,)
N (NH„ Nrf),)
P (PH,)
Sn-H-H- (stannic)
(Si04)=-
C (CH4, CO,)
Br-
Quinquivalent
I-
F-
(OH)-
(NO,)-
(ClO,)-
N (N.0,)
P (Prf).)
Ar (AsiOt)
SedTalent
S(SO.)
VALENCE 217
valences of some elements. Many of these elements, however,
possess other regular valences, in addition to those shown, so that
the list does not pretend to be complete.
Where no charges are indicated, the element, by itself, does not
ordinarily form an ion.
Valence and Electrons. — We stated in om* discussion on
the source of the charges on ions (p. 195) that the outermost shell
of an atom of any element contained a definite small niunber of
negative electrical units (electrons), relatively loosely held. We
also assiuned that, when atoms of unlike elements combined or
reacted, a transfer of electrons from one atom, or group of atoms,
to another might take place.
On the basis of electrons, our idea of valence becomes some-
what more definite. The valence of an element is the number
of electrons that an atom of that element loses, or takes up, in entering
into combination with atoms of other elements. An atom of hydro-
gen has only one electron to lose, the hydrogen ion H+ consists
of nothing but the residual proton. Consequently hydrogen is
imivalent. Zn, however, can lose two electrons to form Zn"*~^",
and Al can lose three to give A1+++. CI can gain one electron to
form Cl~, S can gain two to give S".
The relation of valence to atomic structure in the case of these
and other elements will be taken up in detail in the final chapter.
A Suggestion. — Having just discussed the conception of
valence, we have now considered all the laws of chemical com-
position. At this point the reader should pause and review
thoroughly the subjects of the first eighteen chapters. The under-
standing of the fimdamental principles which this retrospect will
give will greatly lighten the task of understanding the new and
more complex substances we shall have to consider, and the new
kinds of reactions and new conceptions we shall encounter, in the
chapters inmiediately following.
218 smith's intermediate chemistry
Exercises. — 1. Mark the valences in the formnlse: InCU,
V2O6, OSO4, PtCU.
2. Mark the valences of the radicals in the formnlse: Zn(Se04),
Al2(Te04)8, H8(As04), HCSbOs).
r 3. If 26 g. of chromium displace 1 g. of hydrogen from hydro-
chloric acid, what is the valence of chromium in this displacement
(see Table of Atomic Weights)?
4. Correct the following formulse: CaNOs, CaP04, A1(P04)2,
liO, PbF, Bi(N08)2.
f 5. One gram of a quadrivalent element unites with 0.27 g.
of oxygen. What is the atomic weight of the element?
CHAPTER XIX
OXIDIZING SUBSTANCES
In the preceding chapters we have encountered several rather
confusmg oxidizing reactions. For example, in the preparation
of chlorine by the action of KMn04 on HCl (p. 142), it was stated
that the hydrogen chloride was oxidized to chlorine and the potas-
sium permanganate reduced to manganous chloride. The student
may have found it difficult to understand how we can regard
a substance as oxidized when no oxygen is added to it. In order
to explain oxidations more clearly, particularly fn connection with
the conception of valence, we must now learn more about oxida-
tion in general. We can do this best through the study of three
oxidizing substances which are of a simpler nature and are all
in conmion use. These are ozone, hydrogen peroxide, and h]rpo-
chlorous acid.
Ozone Os
When oxygen is blown from a small nozzle through the tip of a
Bunsen flame, a part of it is turned into ozone. The same thing
happens when a platinum wire, heated by an electric current, is
held under Uquefied oxygen. This shows us that, to get ozone, we
must add energy (for example, by strong heating) to oxygen. We
learn, also, that the ozone must be cooled at once and kept cold.
If it Ungers in the cooler (but not cold) region round the flame, it
decomposes again.
Ozone is also obtained by the action of fluorine on water (p.
207).
Preparation of Ozone. — In practice electrical energy, devel-
oped by passing a " silent discharge " through the oxygen, is
219
220
smith's intermediate CHEBnSTRY
employed. This discharge gives ozone very easily, because its
use involves no rise in temperature whatever.
3O2 -* 20,.
The poles of the induction coil are attached to the tinfoil upon
the outside of the outer tube and the inside of the inner tube
(Fig. 63). The " discharge " therefore passes through two layers
Fig. 63
of glass, as well as through the oxygen. The oxygen, from a
cylinder of the gas, flows slowly through the space between the
tubes. At best, only 6 to 7 per cent of the oxygen is usually
changed into ozone.
Physical Properties. — Ozone (Greek, to smell) is a gas of
deep blue color, with a fresh, highly individual odor. It is more
easily liquefied (b.-p. —119°) than is oxygen, and is also much
more solicble in water. Its density is one-half greater than that
of oxygen, and the formula Os records this fact. When a simple
substance shows more than one form in the same state, like oxygen
and ozone, we call them aUotropic modijications. Ice (p. 64)
exists in at least five allotropic modifications.
Chemical Properties. — Ozone is at rather low t^nperatm^s
{e.g,y from 10** to 500°) the less stable form of the element. Upon
standing, and more quickly when warmed, it changes into oxygen,
with liberation of the additional energy it contains.
Being possessed of more internal energy than oxygen, ozone
oxidizes the same substances as does oxygen, only more rapidly
and vigorously. For the same reason it oxidizes many substances
OXIDIZING SUBSTANCES 221
not affected by ordinary ojcygen. Thus, it rusts silver to black
silver peroxide Ag202.
2Ag + 20t -> Ag202 + 2O2.
Ozone also oxidizes a number of organic compounds which are
unchanged in atmospheric oxygen. For example, when ozon-
ized oxygen is bubbled through a dilute indigo solution, a yellow
substance, of much paler tint, isatin, is formed, and the indigo is
said to have been bleached. Indigo is taken, for illustration,
because it is a most widely used dye, employed in dying navy-blue
and blue-black goods, and is totally imaffected by Ught, and by
oxygen, soap, and other ordinary substances:
C16H10N2O2 + 208 -> 2O2 + 2C8H5NO2.
indigo isatin
litmus, and the traces of coloring matter in wax, starch, flour,
and ivory are all oxidized by ozone to colorless, or nearly colorless,
substances. For this reason it is used commercially in bleaching
the last-named materials.
Ozone is sometimes recommended for use, in connection with
ventilation, as a means of destroying minute organisms in the air.
Recent investigations have shown, however, conclusively, that
when thus diluted with air, it has Uttle value as a germicide. It
is employed by some cities for sterilizing the water supply.
Hydrogen Peroxide H2O2
Preparation. — Sodium peroxide Na202, produced by burning
sodimn in dry air, can be dissolved, a Uttle at a time, in ice-cold
water. When this solution is acidified with hydrochloric or sul-
phuric acid, a double decomposition takes place:
Na«02 + 2HC1 -^ 2NaCl + H2O2
and a dilute solution of hydrogen peroxide (mixed with common
salt) is obtained. The nature of the action shows the product to
be an acid, with the negative radical O2".
222 smith's intermediate chemistry
For manufacturing purposes it is more convenient to use barium
peroxide Ba02, suspended in water, and sulphuric acid:
BaOa + H2SO4 -> BaS04 i + H2O2
because the precipitation of insoluble bariimi sulphate carries
the reaction to completion. The precipitate is filtered off and a
clear solution of hydrogen peroxide obtained.
In pharmacy a 3 per cent solution is the one commonly sold.
As hydrogen peroxide decomposes rapidly at 100**, the pure sub-
stance can be obtained only by distiUing off the water under re-
duced pressure.
Properties. — Hydrogen peroxide is a colorless, sjnnipy
liquid of sp. gr. 1.5, miscible with water in all proportions.
Dilute solutions have a metallic taste.
Chemical Properties. — 1. In water solution, hydrogen peroxide
is a weak acid. It enters into double decomposition, particularly
with bases, giving salts containing the bivalent radical O2:
Ca(0H)2 + H2O2 -> CaOa i + 2H2O.
2. When the solution is heated, the compoimd decomposes, with
evolution of heat, giving water and oxygen:
2H2O2 -^ 2H2O + O2 T ^
Contact agents hasten the decomposition. Thus, it takes place
with frothing when the cold solution is appUed, as an antiseptic,
to cuts or sores, or when powders, such as manganese dioxide, are
thrown into the solution.
3. Since hydrogen peroxide, like ozone, gives off oxygen with
liberation of energy, it is an oxidizing agent also. In this respect
its behavior is very similar to that of ozone. It oxidizes colored
organic compounds to colorless ones, and is, therefore, used in
bleaching hair, feathers, silk, and ivory. It is also fatal to micro-
organisms, and is, therefore, employed in medicine to disinfect
wounds and as a throat wash.
In oil paintings the high lights are produced in part with white
lead (carbonate of lead). These disappear and the picture darkens
OXIDIZING SUBSTANCES
223
with age, because hydrogen sulphide in the air turns the white lead
into the black lead sulphide, PbS. In picture restoring the latter
is oxidized to lead sulphate, which is white, by treatment with
hydrogen peroxide solution.
PbS + 4H2O2 -> PbS04 + 4H2O.
4. The following reaction is used as a test for hydrogen peroxide.
When a solution of potassium dichromate K2Cr207 is acidified
with sulphuric acid and a drop of the mixture is added to aqueous
hydrogen peroxide, an unstable substance possessing a deep,
briUiant blue color is formed. By this test the presence of hydro-
gen peroxide in rain water can often be demonstrated.
Hypochlorous Acid HOCl
Pure Hypochlorous Acid. — A pure solution of the acid may
be made by dissolving chlorine monoxide CI2O in water. Chlorine
monoxide is a brownish-yellow, explosive gas, made by passing
chlorine gas over warmed mercuric oxide:
2CI2 + HgO -> HgCl2 + CI2O.
CI2O + H2O -> 2H0C1.
As an acid, hypochlorous acid is very weak, being very Uttle
decomposed into its ions, H+ and (0C1)~.
It is unstable, exposure to sunUght being sufficient to cause it
to give up oxygen, which rises in bubbles
through the solution (Fig. 64) :
2H0C1 -> 2HC1 + 02!.-
Heat is given out in the action, and the stable
hydrochloric acid remains.
It is a most a^Uive oxidizing agent, because
of this tendency to give up oxygen with Ubera-
tion of energy. Thus, its solution oxidizes
organic colored substances, producing color-
less or less strongly colored ones;
r\
Fig. 64
C16H10N2O2 + 2H0C1 -> 2C8H6NO2 + 2HC1.
indigo
iaatin
224 smith's intermediate chemistry
Used as a disinfectant (see p. 226), it oxidizes and destroys bac-
teria. Hypochlorous acid is more energetic as an oxidizing agent
than is ozone or hydrogen peroxide, and is used extensively in
bleaching.
Chlorine- Water. — It will be recalled that chlorine acts chemic-
ally upon water (p. 145) :
CI2 + H2O ^ HCl + HOCl
giving hydrochloric acid and hypochlorous acid. The action is
reversible (read the equation backwards), and in half-saturated
chlorine solution about one-third only of the chlorine has under-
gone the change shown in the equation. But, if a substance
which can be oxidized, such as a dye (attached, perhaps, to cloth),
is introduced into the solution, the HOCl which is present transfers
its oxygen to the dye-stuff. This leaves HCl alone in the solution,
and stops the backward reaction. Hence more of the chlorine
acts upon the water, and more hypochlorous acid is formed. This,
in turn, is used up. Thus, in a few moments, all the
free chlorine is gone, only dilute hydrochloric acid
remains, and a colorless organic compound is left
on the cloth or in the solution.
Chlorine itself is often, erroneously, spoken of as
Uthe bleaching agent. If a dry, colored cloth be
hung for a week in chlorine, dried by having sul-
phuric acid in the bottle (Fig. 65), little or no
change in color wiU occur. But a wet rag is bleached
as soon as the chlorine has time to dissolve in the
Fig. 65 water, and give the necessary hypochlorous acid.
Bleaching Powder. — CaCl(OCl) is made by the action of
chlorine on slaked lime:
Ca(0H)2 + CI2 -* CaCl(OCl) + H2O.
The action is not complete in practice, and the resultant product
is always a basic salt. In solution it gives the ions Ca++, CI", and
OXIDIZING SUBSTANCES 226
(OCl)-. In other words, a solution of bleaching powder in water
acts like mixed solutions of calcium chloride and calcimn hypo-
chlorite. Only the calcium hypochlorite is concerned in the
bleaching process. When the solution is exposed to the air, it
gradually absorbs carbon dioxide (see p. 336). This dissolves in
the solution to form carbonic acid H2CO8, which by double decom-
position liberates hypochlorous acid:
2CaCl(0Cl) + H2CO8 ^ CaCl2 + CaCOs + 2H0CL
If now materials which it is desired to bleach by oxidation
are introduced into the solution, the HOCl is used up, stopping
the backward reaction and carrying the decomposition finally
to completion.
Bleaching. — Cotton and linen, in their original states, are not
pure white. Bleaching is therefore an extensive and most impor-
tant industry. The yam or cloth must first be freed from cotton-
wax and tannin, since the former would hinder the action of the
bleaching agent, and both would also make the subsequent dyeing
imeven. The material is therefore first boiled with dilute caustic
soda-solution, and washed with water. The goods are then first
" chemicked " in cold bleaching powder solution; next " soured "
by immersion in very dilute sulphuric or hydrochloric acid; and
finally washed with extreme thoroughness.
The final washing, to remove all traces of chlorine and bleach-
ing powder, is absolutely necessary. If not removed, the hypo-
chlorous acid acts gradually upon the cotton or Unen, and " rots "
it. Bleaching agents, when used in the household, carelessly, are
Uable to cause extensive damage from this cause. A dilute solu-
tion of sodiiun thiosulphate (photographers' " hypo '') is often
used, as " antichlor,'' to interact with and remove the last traces
of chlorine.
Cotton and linen (CeHioOs)^ are rather indifferent chemical
substances (p. 398), and stand brief contact with dilute chlorine-
226 smith's intermediate chebhstry
water without much alteration. But wool and silk contain com-
pounds of nitrogen (proteins) and are acted upon as rapidly as are
the traces of colored matter themselves. Hence sulphiux)us acid
(see p. 259) is used for bleaching these materials.
Bleaching Powder as a Disinfectant. — A disinfectant is a
substance which destroys bacteria and other minute, and often
harmful organisms. Bleaching powder CaCl(OCl) has a distinct
odor. This is due to the slow action of the carbon dioxide and
moisture of the air upon the salt, Uberating hypochlorous acid.
Bleaching powder, when scattered around, will therefore dis-
infect the surrounding air, because the hypochlorous acid thus
liberated kills all bacteria present by oxidation.
When an epidemic of typhoid fever occurs, it is usually traced
to the presence of colon baciUi and typhoid organisms in the drink-
ing water. The most efiFective means of destroying these bacilU is
to add, at the distributing point, a small proportion of bleaching
powder (about 20 pounds per miUion gallons of water).
Recently, chlorine^water has in many cases taken the place of
bleaching powder for this purpose. Cylinders of liquid chlorine
(p. 144) were used in the Great War to kill germs as well as to
kill Germans, all water supplies being steriUzed, whenever possible,
by the addition of very minute amounts of chlorine.
Oxidations Previously Mentioned. — The simplest oxida-
tions are reactions in which free oxygen is actually used up, for
example in the union of oxygen with metals and with non-metals:
2Cu + O2 -> 2CuO.
S + 02->S02.
The displacement of another element from a compound by
oxygen is also oxidafionT^
4HCl + 02±=;2H20 + 2a2.
OXIDIZING SUBSTANCES 227
The transfer of combined oxygen from one substance to another
in a reaction is again oxidation:
2KMn04 + 16HC1 -^ 8H2O + 2KC1 + 2MnCl2 + 6CI2.
Mn02 + 4HC1 -> MnCl2 + CI2 +. 2H2O.
What is the substance oxidized in the last three reactions? It
is hydrochloric add, and the product of its oxidation is chlorine.
True, we have added no oxygen to chlorine itself in any of these
reactions, but we have done something which is exactly equivalent
to addition of oxygen, we have taken away hydrogen, and we have
not replaced that hydrogen by any other positive element.
Note that every oxidation is accompanied by reduction of the
oxidizing agent. Thus in the first of the three reactions just dis-
cussed, the free oxygen is reduced to water. In the other reactions,
KMn04 and Mn02 are reduced to MnCl2. In all three cases
HCl is the reducing agent.
The appearance of a product that could only be formed by
reduction is sometimes the first thing that calls our attention to
the fact that an oxidizing action has occurred. When concen-
trated sulphuric acid acts upon hydrogen iodide (p. 204), the
iodine vapor given ofif on warming shows that there was oxidation,
but the odor of the hydrogen sulphide is the first thing we notice
when doing the experiment:
H2SO4 + SHI -> H2S + 4H2O + 4I2.
Removal of the elements of water from a compound is neither
oxidation nor redu^Mon, for hydrogen and oxygen are both removed:
H2CO3 -^ CO2 + H2O.
NH4OH -^ NHs + H2O.
>
We can now see that oxidation, in the above cases, consists always
in adding oxygen or removing hydrogen.
Other Cdises of Oxidation. — But oxygen is only one of a class
of elements which we call non-metallic or negative elements, so
that we do not restrict the term " oxidation " to actions involving
228 smith's intermediate chemistry
oxygen. Thus, forming a sulphide, by heating a metal with
sulphur, is oxidation also :
Fe + S->FeS.
Similarly, changmg ferrous chloride FeClj to ferric chloride FeCU
is oxidation:
2FeCU + Cl2->2FeCU.
In every compoimd one of the eleipents is relatively positive
and the other relatively negative. Iron is positive, sulphur and
chlorine are negative.
Oxidation, then, is introducing, or increasing the proportion of
the negative element, or removing, or reducing the proportion of
the positive element. Reduction is the converse.
Oxidation and Valence. — Combining a negative element
with a metal raises the dctive valence of the latter from zero to
some finite value. Metallic copper has no valence in use. In
CuCU the copper is employing the valence II. The copper has
been oxidized. Similarly, changmg FeCl2 to FeCU increases the
active valence of the iron from II to III. Conversely, changing
HCl to CI2 alters the active valence of the chlorine from I to zero.
Hence, oxidation may be defined as increasing the active valence
of a positive element or decreasing that of a negative element.
Reduction is the converse.
Oxidation and Electrons. — Finally, since increasing the
valence of a negative atom means adding one or more electrons to
that atom, and increasing the positive valence of an atom means
removing one or more electrons, we reach the briefest definition by
saying: Oxidation is removing electrons and reduction is adding
electrons.
Exercises. — 1. Why does air containing ozone lose the latter
(by change into oxygen) quicker when warm than when cold?
2. Mark the valences of the radicals in bariiun peroxide and
hypochlorous acid.
OXIDIZING SUBSTANCES 229
/ 3. What volume (at 0** and 760 mm.) of oxygen would be
obtained by the decomposition of the hydrogen peroxide in 1 kilo-
gram of the 3 per cent solution?
4. Why does a given weight of chlorine in the form of hypo-
chlorous acid have twice as great a bleaching (or oxidizing)
capacity as has the same weight of chlorine in chlorine water?
5. Write down all the oxidizing actions mentioned in chapter
XVII, noting in each case the oxidizing agent and the reducing
agent.
CHAPTER XX
CHEMICAL EQUILIBRinM
In spite of its formidable title, this chapter will introduce noth-
ing novel. Its purpose is to collect together and organize more
definitely a nimiber of scattered facts and ideas which have already
come up in various connections. On this accoimt, however, it will
be all the more necessary for the reader to refresh his remembrance
of these facts and ideas by re-reading all pages to which reference
is made.
Reversible Actions. — In discussing Deacon's process (p. 140),
it was stated that the action 4HC1 + Oj ->2H20 + 2CI2 comes to
rest although a large amoimt ot both of the interacting substances
(20 per cent at 345®) still remains available. Now the materials
thus left imused are presumably no less capable of interacting
than were the parts which have already reacted. The solution
of this mystery lies in the fact that the products themselves interact
to reproduce the initial substances (read the equation backwards).
Thus two changes, one of which imdoes the work of the other, are
going on simultaneously. In consequence of this, neither action
can reach completion. As we should expect, experiment shows
that it makes no difference whether we start with pure chlorine
and steam, or with hydrogen chloride and oxygen; the proportions
of the four substances foimd in the tube, after it has been kept at
345® for a suflScient time, are in both cases the same. A general
statement may be foimded on facts like this, to the effect that a
chemical action must remain more or less incomplete when the
reverse action also takes place tmder the same conditions. Two
arrows pointing in opposite directions are used in equations
representing reversible changes.
230
CHEMICAL EQUILIBRIUM 231
The foregoing example of a reversible action, and the following
examples which very closely resemble it, should now be looked up
and studied attentively. The discussion in this and the following
sections, for which they furnish the basis, cannot otherwise be
understood: (1) the behavior of water vapor at the boiling-point
(pp. 63-4) and at 2000° (p. 67) ; (2) the depression of the vapor
pressure of a Uquid by a non-volatile solute (p. 117) ; (3) the action
of diliUe sulphuric acid on common salt (pp. 127-8) ; (4) the inter-
action of chlorine and water (pp. 145, 224) ; (5) the ionization of
electrolytes (p. 182).
Explanation in Terms of Molecules. — Restating these
reactions in terms of the molecules will enable us to reason more
clearly about this variety of chemical change. Suppose we start
with the materials represented on one side only of such an equation,
say the hydrogen chloride and oxygen in that on p. 230. The
molecules of these materials will encounter one another frequently
in the course of their movements. In a certain proportion of these
collisions the chemical change will take place. In the earliest
stages there will be few of the new kind of molecules (say, of
chlorine and steam), but, as the action goes on, these will increase
in number. There will be two consequences of this. In the first
place, the parent materials (in this case, hydrogen chloride and
oxygen) will diminish in amount, the coUisions between their
molecules will become fewer, and the speed of the forward action
will therefore become less and less. In the second place, the in-
crease in the number of molecules of the products will result in
more frequent coUisions between them, in more frequent occur-
rence of the chemical change which they can undergo, and thus in
an increase in the speed of the reverse action. The forward action
begins at its maximum and decreases in speed progressively; the
reverse action begins at zero and increases in speed. Finally the
two speeds mitst become equal, and at that point perceptible change
in the condition of the whole must cease.
232 smith's intermediate chemistry
The most immediate inference from this mode of viewing the
matter is, that the apparent halt in the progress of the action
does not indicate any cessation of either chemical change. Both
changes must go on, in consequence of the continued encounters
of the proper molecules. But since the two changes proceed
with eqiuil speeds they produce no alteration in the mass as a
whole. In fact, the final state is one of equiUbrium, and not
of rest, one of balanced activity and not of repose. Hence,
chemical changes which are reversible lead to that condition
of seemingly suspended action which we speak of as chemical
equilibrium.
Chemical Equilibrium and its Characteristics. — The de-
tailed discussion of the relations of liquid and vapor (pp. 62-64),
and of saturated solution and imdissolved soUd (pp. 120-121),
has already famiUarized us with the term equiUbrium and its
significance. We can, in fact, apply to the discussion of any kind
of reversible phenomena, the sets of ideas in regard to exchanges of
molecules there elaborated.
In particular, the reader will note that the three characteristics
of a state of equilibrium, developed and illustrated in the case of
the physical equiUbrium between a Uquid and its vapor (p. 63),
apply also to a typical case of chemical equiUbrium, such as that in
Deacon's process now before us. Thus:
1. There are the two opposing tendencies, which ultimately
balance one another. Here they are the tendency of the steam
and chlorine to produce hydrogen chloride and oxygen, and the
tendency of the hydrogen chloride and oxygen to reproduce steam
and chlorine by this interaction.
2. At equilibrium the two opposing tendencies or activities are
still in full operation, although their effects then neutraUze one
another.
3. (and this is the chief mark of chemical, as it is of physical
equiUbrium). The system is in a sensitive state, so that a change
CHEMICAL EQUILIBRIUM 233
in the conditions (temperature and pressing or concentration),
even if slight, produces a corresponding change in the state of the
system, and does this by favoring or disfavoring one of the two
opposing tendencies or activities. Such a change is called a dis-
placement of the equilibrium, for the system settles down in a new
state of equilibrium with new proportions of the two sets of sub-
stances, corresponding to the changed conditions. Thus, in the
present instance, a change from 345**, where there is 80 per cent
of the material in the form of steam and chlorine, to 384** results
in the diminution of this proportion to 75 per cent. The equiUb-
rium is affected by changes in concentration also, as we shaU
presently see.
Now, the foregoing facts show that the key to imderstanding
chemical activities, their magnitudes, their changes, and especially
their practical results, miLst lie in knowing how changes in the
conditions affect them. Hence, to the chemist, famiUarity with the
influence of conditions on chemical phenomena must be of the
greatest practical importance.
The " conditions " to be considered are famiUar, — temperature,
and concentration or, in the case of a gas, partial pressure. The
"activity" of an action is accurately measured by the speed with
which the action proceeds. Thus, if the foregoing section be
re-examined, it will be seen that we spoke throughout of the speed,
rather than of the tendency or activity.
Finally, temperature and other conditions influence also the
activities in, and therefore the speeds of, those actions which pro-
ceed to completion, and are not reversible. Hence, imless our
statements are expressly restricted to reversible actions arid to
states of equilibrium, they apply to all chemical changes.
The Influence of Concentration. — In the first place, let us
assume that the temperature is constant, and let us confine our
attention for the present to the influence of concentration upon a
chemical reaction. We have seen (p. 231) that the speed of a
234 smith's intermediate chemistry
chemical change is determined by the frequency with which the
molecules of the interacting substances encoimter one another.
The frequency of the encounters amongst a given set of molecules,
resulting in a definite chemical change, will in turn evidently
depend entirely upon the degree to which the molecules are con-
centrated in each other's neighborhood. Larger amounts of one
of the materials, for example, will not result in more rapid chemical
action, if the larger amoimt of material is also scattered through a
larger space. Chemical changes, therefore, are not accelerated by
increasing the mere quantity of any ingredient, but only by
increasing the concentratioii of its molecules. Thus, a large
amount of a 0.1 normal solution of hydrochloric acid with a piece
of zinc will generate hydrogen no faster than a smaller amoimt.
But substitution of a normal solution of hydrochloric acid, which
contains a higher concentration of hydrogen ions, will instantly
increase the speed of the action. In the second case, the number
of hydrogen ions reaching the zinc per second is greater, and
the displacement reaction Zn + 2H+ — > Zn++ + H2 t proceeds
more rapidly. So also, iron burns faster in oxygen (100 per cent)
than in air (20 per cent oxygen).
With a reversible action the effect on the speed is the same,
excepting that the continued activity of the reverse action pre-
vents the direct one from reaching completion.
Thus, if, in the action of hydrogen chloride upon oxygen, we
introduce into the same space an extra amoimt of oxygen, this
faciUtates the formation of steam and chlorine l)y increasing the
possibiUties of encounter between molecules of hydrogen chloride
and oxygen. At the same time it does not aflfect (c/. p. 230) the
number of encounters in a given time of steam and chlorine mole-
cules with one another which result in the reverse transformation.
The proportion of chlorine (and steam) formed, therefore, from a
given amount of hydrogen chloride will be greater, although the
total possible (by complete consumption of the materials) has not
been altered, since the quantity of one ingredient only has been
CHEMICAL EQUILIBRIUM 236
increased. The introduction of an excess of hydrogen chloride
would have had precisely the same effect.
An Experimental Illustration. — A reaction in which the
effects of different concentrations were carefully studied by Glad-
stone (1855) affords a good illustration. If ferric chloride and
ammonium thiocyanate are mixed in aqueous solution, a Uquid
containing the soluble, blood-red ferric thiocyanate is produced.
The compound radicals are (NH4) and (CNS), and the action is
a simple double decomposition:
FeCla + 3NH4CNS ^ Fe(CNS)8 + 3NH4CI.
The action is a reversible one, and the mixture is homogeneous,
i.e.y there is no precipitation. Now, if the two just-named salts
are mixed in very dilute solution in the proportions required by the
equation, say by adding 20 c.c. of a decinormal solution of each
salt to several liters of water, a pale-reddish solution is obtained.
When this is divided into four parts, and one is kept for reference,
the addition of a little of a concentrated solution of ferric chloride
to one jar, and of ammonium thiocyanate to another, will be
found to deepen the color by producing more of the ferric thio-
cyanate. On the other hand, mixing a few drops of concentrated
ammonium chloride solution with the fourth portion will be found
to remove the color almost entirely, on accoimt of its influence in
favoring the backward change.
The Law of Molecular Concentration. — The general prin-
ciple discussed and illustrated in this section may be called the
law of molecular concentration^ and may be stated as follows : In
every chemical change the activity, and therefore the speed
of the action, is proportional to the molecular concentration of each
interacting substance. This holds whether the action is reversible
or not.
The molecular concentration is expressed^ numerically j for each
substance, in terms of the number of moles (gram-molecular
236 smith's intbrmediate chemistry
weights, p. 122) of the substance contained in a Uter of the whole
mixture. There is the same number of molecules in a gram-
molecular weight of any substance (see p. 84). Hence the number
of moles per liter defines the concentration of each substance in
terms of this number of molecules in a Uter as the unit of concen-
tration.
As an example, the dissociation of phosphorus pentachloride
vapor into phosphorus trichloride and chlorine (p. 145) :
PCl6^PCl3 + Cl2
may be considered.
The law states that, in any mixture of the three substances, the
speed of the forward action, or decomposition, is proportional
to the concentration of PCU molecules. It is therefore equal
to the concentration of PCI5 molecules, which we may write
[PCU, multipUed by a constant, which we shall write Ki. Mathe-
matically expressed, if Si is the speed of decomposition:
Si = Ki [PCI5]
In the same way the law states that the speed of the reverse
action, or combination, is proportional to the concentration of
PCI3 molecules and also proportional to the concentration of CI2
molecules. It is therefore eqital to the concentration of PCI3
molecules [PCI3], multipUed by the concentration of CI2 mole-
cules [CI2], multipUed by another constant, which we shaU write
K2. If S2 is the speed of combination, we have now:
S2 = K2 [PCI3] • [CI2]
Each reaction, it will be noted, must slow up as the concen-
tration of the reacting substances diminishes. This is true of aU
reactions. The further they proceed, the less rapid their speed.
The burning of a candle in a confined space becomes less brilUant
as the available oxygen is used up. The evolution of hydrogen
in the action of an acid on a metal becomes very slow when the
acid is nearly exhausted, owing to the diminution in the concen-
tration of hydrogen ion.
CHEMICAL EQUILIBRIUM 237
The Condition for Chemical Equilibrium. — As we have
seen (p. 231), the characteristic of a system in chemical eqniUb-
rinm is that the speeds of the forward and reverse reactions have
become equal. Appljring this to the case of the dissociation of
phosphorus pentachloride discussed above, we see that, when we
have an equilibrium mixture of the three substances, where Si =
S2, then we also have the relationship:
Ki [PCld = K2 [PCI3] • [CI2]
This may be written in the form:
Ki _ [PCI3] - [CU]
K2 [PCI5]
The ratio K1/K2 is, of course, a constant, since Ki and K2 are
both constants. This ratio, which we may write K, is caUed the
equilibrium constant of the reversible reaction.
The equilibrium constant is a very important quantity. Once
we have determined it, by investigating one equiUbrium mixture,
we can calcukUe exactly what will happen to any mixture of the
three substances concerned at the same temperature. However
much we may vary the molecular concentrations of the three
substances, whether by changing the pressure or by adding an
excess amount of one of them, such as chlorine, the composition
of the mixture will adjust itself until, when equiUbrium is attained,
the ratio [PCla] • [Cy/IPCU] has again reached the value K. In
large-scale industrial processes, therefore, a knowledge of the
equiUbrium constant of the reaction involved is often of inesti-
mable value.
In many important reversible reactions, the concentrations
of the reacting substances on one side of the equation are, imder
equiUbrium conditions, very much greater than the concentra-
tions of those on the other side. Although the reaction is rever-
sible, it proceeds much farther towards completion in one direc-
tion than the other. This fact may often be indicated, very con-
238 smith's intermediate chemistry
veniently, by modifying the thickness of the arrows used to express
reversibiUty. Thus:
HCl ^ H+ + CI- H2O ^ H+ + OH-
In the case of phosphorus pentachloride vapor, 80 per cent of
the whole weight of material in the equiUbrium mixture at 250°
and 760 mm. pressure is dissociated into PCI3 and CI2.
Homogeneous and Inhontogeneous Systems. — While
there are all degrees of speed in chemical actions, yet in practice
we quickly distinguish two different classes. There is a class of
actions of which most examples are almost instantaneously accom-
pUshed, and a class in which, frequently, the operation takes
minutes or even hours. The classes overlap, but, in a general
way, the following distinction may be made.
To the former, speedy class belong the explosion of hydrogen
and oxygen or other gaseous mixtures, and the interactions when
solutions are mixed, as in precipitations. In view of the foregoing
explanations, we perceive that the rapid accomplishment of such
actions is due, not so much to any especially great intrinsic affin-
ity, as to the homogeneous state of mixture of the interacting
materials. This, of course, is a purely physical, and not a chem-
ical motive for speedy interaction. In intimate mixtures, every
molecule has an equal opportunity freely to encounter every other
molecule and there is therefore no mechanical impediment to
the operation of the affinities of the substances. Hence the ap-
parent activity is great.
To the second class, comprising the slower actions, belong cases
like the interaction of a piece of zinc with hydrochloric acid, or of
manganese dioxide (p. 142) with the same acid, whereby hydrogen
and chlorine, respectively, are slowly evolved, and the soUd is grad-
ually consumed. Here the hindrance is evidently the fact that
the interacting substances are not intimately mixed. In the slow
actions, the system is inhomogeneous. Pulverizing the soUd
before use will increase the speed, indeed, by providing more
CHEMICAL EQUILIBRIUM 239
surface and better mutual contact, but will not transfer the action
to the rapid class. It is chiefly the dissolved part of the substance
which interacts, for chemical action takes place between mole-
cules, and only the dissolved part is disintegrated in such a way
that the molecules are readily accessible. Thus, the action is
held back by continual waiting for the slow replenishment, from
the " insoluble " soUd, of the supply of dissolved molecules. In
the cases cited, the restraining influence of the dissolving process,
which is part of the whole phenomenon, may be formulated thus:
Zn (soUd) ^ Zn (dslvd.) + 2HC1 -^ ZnCla + H2.
MnOaCsoUd) ^ MnOaCdslvd.) + 4HC1 -^ MnClj + 2H2O + CI2.
Here, again, the mechanical details, depending on physical prop-
erties, have more to do with the progress of the action than has
the chemical aflSnity. In terms of the law of concentration, the
action is slow, and the apparent activity small, because the con-
centration of the acting molecules of one of the substances is very
small, and cannot be increased because of low solubiUty.
Displacement of Equilibria. — We have seen (pp. 234-5) that
one way in which a reversible action may be forced nearer to com-
pletion, in one direction or the other, is the introduction of an
excess of one of the ingredients contributing to the forward action.
This method of displacing the equiUbrium point, however, cannot
be very effective, unless it is possible to introduce an exceedingly
large excess of the selected ingredient in a high degree of molecular
concentration, since this operation does not in any way affect or,
in particidary restrain the reverse action which is continually imdoing
the work of the forward one. A much more effective means of
furthering the desired direction of such actions is found, therefore,
in the restraint or practical annulment of the reverse action. A
good way of accompUshing this is to allow the products of the
direct action to separate into an inhomogeneous mixture. Any
agency which could remove the water vapor as fast as it was
formed by the interaction of hydrogen chloride and oxygen, for
240 smith's intermediate chemistry
example, would entirely stop the reproduction of these substances,
and so would enable the forward action (4HC1 + 02-^ 2H2O
+ CI2) to run to completion.
This might be reaUzed by causing one end of a sealed tube
charged with the substances, after the contents had settled down
to a condition of equiUbrium, to project from the bath in which the
whole had been kept at 345° (Fig. 66, which is simply diagram-
matic). By cooling this end, a
large part of the steam would
quickly be condensed in it to the
Uquid form, while the other sub-
stances would remain gaseous.
^^^- ^^ In other words, the concentra-
tion of the water vapor would be greatly reduced. In fact, only
the trace of vapor which cold water gives would then be avail-
able to interact with the chlorine, and reproduce hydrogen chlo-
ride. Meanwhile the decomposition of the latter would go on, and
thus, eventually, almost all the water would be foimd in one end of
the tube, and the chlorine, all free, would occupy the rest. By
this purely mechanical adjustment the chemical change would
therefore be carried from 80 per cent completion to almost absolute
completion:
4HC1 + O2 ?=t 2CI2 + 2H2O (vapor) t± 2H2O (Uq.)
If, on the other hand, arrangements were made to have pow-
dered marble, in a sealed bulb of thin glass, enclosed in the tube,
we might imagine the very opposite of the above effect to be pro-
duced. The breaking of the bulb of marble, when equiUbrium
had been reached, would provide means for the removal of all the
hydrogen chloride,* while the other three substances would still be
* The hydrogen chloride wotdd be destroyed by interaction with the
maxble:
2HC1 + CaCOa -^ CaCl, + COj + H2O.
The calcium chloride is a solid. The gas, carbon dioxide, does not interact
with the other substances, and would not, therefore, interfere with the formar
tion of fresh hydrogen chloride.
CHEMICAL EQUILIBRIUM 241
gaseous. Thus, the compound (HCl) having been reduced in
concentration to the point of being removed entirely, there would
be no direct action to undo the work of the reverse action. The
whole chlorine would, therefore, soon have passed through the
form HCl. Hence, by another mechanical arrangement, an action
which ordinarily could progress to only 20 per cent would be
turned into a complete one:
2CI2 + 2H2O ^02 + 4HC1 (+ CaCOs -^ CaClg + H2O + CO2).
Reversibility Usually Avoided. — In every-day chemical
work, since our object is usually to prepare some one substance,
chemists either avoid chemical changes which are notably rever-
sible, or adjust the conditions, as is done in the foregoing illus-
trations, so that the reverse of the action which they desire is
prevented. In consequence of this, when carrying out the direc-
tions for making famiUar preparations, the fact that such actions
are reversible at all very readily escapes our notice. Arranging
the conditions so that the separation of a soUd body by precipita-
tion, or the Uberation of a gas, takes place, are the two commonest
ways of rendering a reversible action complete. Excellent exam-
ples of both of these are furnished by the chemical change used
in producing hydrogen chloride by the interaction of salt and
sulphuric acid, the full discussion of which (p. 127) should now be
studied attentively in the light of these explanations.
Double decompositions between electrolytes in solution may
also be carried to completion by arranging that one of the products
of the reaction shall be a non-ionized, or practically non-ionized,
substance. NetUrcdizaMon of an acid by a base is a case in point
(see p. 193).
The Influence of Temperature on the Speed of any ReaC"
tion. — The activity of chemical change, and therefore the
speed of all chemical changes, is increased by raising the temper-
ature and diminished by lowering it (qf. p. 27). Different actions
242 smith's intermediate chemistry
are affected in different degrees, and no simple rule accurately
defining the effect can be given. Roughly speaking, however,
a rise of 10° doubles the speed of every action. A rise of 100®
will therefore make the speed roughly 1024 times greater. Hence,
when the chemist finds that two substances show no evidence of
interaction he infers that there must be either slow action or
none, and he seeks to settle the question quickly by heating the
mixture.
The Influence of Temperature on a System in Equilib'
rium. — In a reversible change the two opposing reactions are
different actions and their speeds are therefore affected in different
degrees by the same alteration in temperature. Hence, when the
temperature is changed, the relative amount of the two sets of
material present is altered and the equiUbrium is displaced.
Thus, in Deacon's process, a rise of 40° in the temperature dis-
places the equiUbrium backwards (p. 233), and diminishes the
yield of chlorine by 5 per cent. In the vapor of phosphorus penta-
chloride (p, 236), the displacement is in the opposite direction.
At 250°, and 760 mm. pressure, 20 per cent of the material is
present as pentachloride and 80 per cent as trichloride and chlo-
rine. At 300° only 3 per cent of the pentachloride remains while
at 200° only a little more than 50 per cent is dissociated. Evi-
dently, here, raising the temperature favors the decomposition
of the pentachloride, and therefore increases the speed of its
dissociation more than it does the speed of the reunion of the
trichloride and chlorine.
Van'*t Hoff^s Law. — Now the facts mentioned above are con-
nected by a law which will answer many practical questions in
chemistry.
When phosphorus trichloride and chlorine combine (to form
PCIb), heat is given out. Conversely, when phosphorus penta-
chloride dissociates, heat is absorbed:
PCI5 + 30,000 cal. ^ PCI3 + CU.
CHEMICAL EQUILIBRIUM 243
Now, when the temperature is raised, the action proceeds in the
direction of decomposing more of the pentachloride. That is, the
equiUbrium is displaced in the direction which absorbs heat.
In Deacon's process, we find that the interaction of hydrogen
chloride and oxygen liberates heat:
4HC1 + O2 <^ 2H2O + 2CI2 + 28,000 cal.
and in this action raising the temperatm*e drives the equiUbrium
backwards, and a lowering in the temperature is required to
increase the yield of chlorine.
The rule is obvious, and apphes to all reversible reactions:
When the temperature of a system m equilibrium is raised, the
equilibrium point is displaced in the direction which absorbs heat.
In other words, a rise in temperature favors the interaction of that
one of the two sets of materials to which the heat is added (+ sign)
in the equation. If the equation happens to be written with a
negative heat of reaction (e.gr., p. 162), the heat can, of course, be
transferred to the other side with its sign changed. This law is
known as Van't HoflPs law of mobile equilibrium.
This law is of practical value. More than once, in chemical
factories, much time and money have been spent on trying to
arrange machinery to give a better jdeld of some substance at a
high temperature, when a reference to this law would have shown
that the chief change necessary was to use a lower temperature.
We shall frequently have occasion to refer to this law.
Application to Physical Equilibria. — Van't HoflPs law
applies also to physical equilibria. Thus, the vaporization of
a Uquid absorbs heat, and so an increase in temperature will
increase the pressure, and therefore the concentration of its vapor.
The special case of a saturated solution may be somewhat more
fully considered.
When we heat a saturated solution, with excess of solid present,
more soUd will dissolve as the temperature is raised, if solution
is attended with absorption of heat. This is the usual case, as is
244 smith's intermediate chemistry
shown by the way in which most solubilities increase with rising
temperature. The solution of a soUd in a Uquid, indeed, may be
regarded in normal cases as equivalent to the process of fusion,
which, Uke vaporization, always absorbs heat. The precipita-
tion of crystals from a supersaturated solution, similarly, may be
considered as analogous to the solidification of a pure substance,
the freezing-point being depressed below the normal va'ue, how-
ever, by the presence of the solvent (see p. 119). Freezing-point
lowering and solvbility phenomena are therefore identical. When
we determine the depression of the freezing-point of water on
addition of sugar by noting at what temperature ice begins to
separate out, we measure, simultaneously, the solubility of ice
in the solution at that temperature. In the same way, when we
determine the solubiUty of sugar in water at 25°, we establish
also how much water must be added to a given quantity of sugar
to depress the freezing-point of sugar to 25°. It is therefore
just as correct to say that sugar melts in tea as it is to say that
ice melts in iced tea.
Often, however, extensive chemical reactions occur in the pro-
cess of solution (see p. 123), and the heat effect of these reactions
may be considerable. Thus when water is added to concentrated
sulphuric acid, so much heat is evolved that the water boils.
In the case of a few salts, also, we find that solution in water takes
place with evolution of heat, and we draw the conclusion that the
heat of the chemical reactions involved more than counterbalances
the heat absorbed in the process of fusion.
When we heat a saturated solution, with excess of soUd present,
crystals will separate out as the temperature is raised, if predpitor
tion is attended by absorption of heat. This is the case with
anhydrous sodium sulphate Na2S04 and some calcium salts in
water (see p. 113).
Le Chatelier'^8 Law. — The above-mentioned law is really a
particular case of a more general one, the law of Le ChateUer.
CHEMICAL EQUILIBRIUM 245
K some stress (e,g,y by change of temperature, pressure, or con-
centration) is brought to bear on a system in equilibrium, a reac-
tion occurs, displacing the equilibrium in the direction which tends
to tmdo the effect of the stress. Thus, raising the temperature
furthers the change which absorbs heat — and therefore would
tend to lower the temperature. Increasing the concentration
of the molecules pushes the action in the direction which uses up
these very molecules (p. 234). Again pressure causes ice to melt,
because the water which is formed occupies a smaller volume, and
this change tends to reUeve the pressure. But pressure will not
cause most substances to melt, because usually the Uquid form
occupies a greater voliune and its production would tend to in-
crease pressure.
The student is cautioned against applying these laws to systems
not in equilibrium, for example, to unsaturated or supersaturated
solutions. To such cases they do not, necessarily, apply. Thus
the addition of a small quantity of cupric chloride to water is
attended by evolution of heat. It would be quite wrong to reason
from this, however, that the solubiUty must fall off as the temper-
ature is raised. The salt is extremely soluble in water, and if we
keep on adding it to the solution until the latter is saturated we
find that the last portions dissolve with absorption of heat. Now
it is only the saturated solution that is in equilibrium vrith the solid,
hence it is only with respect to this solution that we can apply
Van't Hoff's or Le Chatelier's law. In accordance with the behav-
ior of this solution, we find that the solubihty increases with rising
temperature.
The characteristics of systems in equiUbrium (pp. 63, 232)
should therefore be kept in mind carefully, in order to avoid mis-
takes.
Summary. — In this chapter we have answered three ques-
tions:
246 smith's intermediate chemistry
1. Why do some chemical actions cease, while still incomplete?
Answer: They are reversible.
2. What explains the position of the equiUbrium point? An-
swers: (a) Equal effects of opposed molecular actions; (6) Equal-
ity in speed of opposed reactions.
3. What will displace the equiUbrimn point? Answer: (a)
Change in concentration of one (or more) of the substances; (6)
Change in the temperature.
Exercises. — 1. Explain the completeness of the action by
which hydrogen chloride and water, respectively, are formed by
direct imion of the elements.
2. Explain the completeness of the action by which silver
chloride (p. 131) is formed.
3. Explain why the decomposition of potassium chlorate is
complete.
4. In view of the statement on p. 16, explain why mercuric
oxide is completely decomposed by heating.
' 5. Why can magnetic oxide of iron be reduced completely by
a stream of hydrogen (p. 57), and iron oxidized completely by a
current of steam (p. 51)?
6. With the phosphorus pentachloride system, say at 250°,
what effect would suddenly enlarging the space containing a given
amoimt of the vapor produce? What would be the effect of
diminishing the space? What would be the effect of introducing
additional chlorine into the same space (p. 234)?
"^ 7. Are the following systems in equihbrium or not: (a) a cap-
tive balloon puUing at its rope, (&) a balloon floating freely
in cooling air, (c) a man wearing a heavy suit in winter only,
(d) a woman changing to a velvet hat in August and to a straw
hat in February? What is the stress in each case? In which
systems does a reaction take place which opposes the effect of the
stress?
CHEMICAL EQUILIBRIUM 247
8. What inference should you draw from the fact that : (a) the
solubilities of potassium nitrate and of Glauber's salt increase
with rise in temperature; (b) those of calciiun hydroxide (p.
113) and triethylamine decrease with rise in temperature?
CHAPTER XXI
SULPHUR AND HYDROGEN SULPHIDE
Sulphur, the compounds of which have been so often men-
tioned, provides us, in sulphuric acid, with a substance which has
more extensive and more important appUcations in commerce
than any other chemical. The element sulphur, itself, enters,
with potassium nitrate and charcoal, into gunpowder. Vulcanite
is a compoimd of caoutchouc (rubber) and sulphur. Sulphur is
employed to destroy fimgi on grape-plants, and furnishes sidphur
dioxide for bleaching and disinfecting.
Sources. — The greater part of tne sulphur of conmierce come?
from Sicily, Louisiana and Texas. In Sicily, free sulphur is
mixed with pimiice and other rocks. When the lumps of rock,
obtained by mining or quanying, are heated by setting fire to the
sulphur (there is no coal in Italy), the sulphur melts and runs to
the bottom of the kiln. This product is far from pure, and is dis-
tilled from iron rfetorts. The vapor is condensed in chambers of
brick, and the Uquid is run into moulds, giving roll sulphur. The
first vapor condensed, while the chambers are cold, yields flowers
of sulphur.
In Louisiana the sulphur occurs in a deposit over half a mile in
diameter, below 900 feet of clay, quicksand, and rock. It is
obtained by means of borings, which permit foiu* pipes, one within
the other, to reach the deposit. Water, previously heated imder
pressure to a temperature of 170°, is pmnped down the two outside
pipes (6 and 8 inches in diameter). After time has been allowed
for the melting of a quantity of the sulphur (it melts at 114.5°),
compressed air is pumped down through an inner, one-inch pipe.
The melted sulphur, alone, has twice the specific gravity of the
248
SULPHUR AND HYDROGEN SULPHIDE 249
water in the outer pipes. But the air breaks up into small bubbles,
forming with the Uquid sulphur an emulsion which has a lower
specific gravity, and this flows freely up a three-inch pipe which
surrounds the air pipe. The sulphur runs into wooden enclosures,
measuring 150 by 250 feet, in which it quickly solidifies. The
product is so piu*e that, for most piu*poses, no other treatment
is required. The output of Louisiana and Texas — 500 tons a day
from each well and, in all, over 1,000,000 tons annually — suppUes
the whole demand of the United States, and could easily be
increased.
A nimiber of sulphates, such as gypsum (CaS04,2H20) and
barite (BaS04), and several sulphides, such as galena (PbS), zinc
blende (ZnS), and pyrite (FeS2), are foimd in large quantities as
minerals. The last two sulphides are used in the manufacture of
sulphuric acid.
Allotropic Forms of Sulphur. — Sulphur appears in two
different liquid forms, and in two famiUar and perfectly distinct
soUd varieties. The two latter are called, from their crystalline
forms, rhombic and monoclinic sulphur.
Physical Properties of Rhombic Sulphur. — This form is
yelloWy with specific gravity 2.06. Natural sulphur, roll sulphur,
and practically all of most specimens of flowers of sidphur are of
this variety, and are identical in all physical properties. Speci-
mens of natural sulphur often show the rhombic crystalline form
very clearly. All the forms of sidphur are insoluble in water, and
all the crystalline forms are soluble in carbon disulphide. Good
rhombic crystals are obtained from the solution (Fig. 10, p. 14).
The rhombic form is stable when not heated above 96°. K
kept above this temperature, it slowly changes into monoclinic
sidphur.
Monoclinic Sulphur. — This form is obtained most quickly
by first melting some sulphur (m.-p. 114.6®), and then allowing it
250 smith's INTBBMEDIATE CHBHI8TRT
slowly to cool. As the temperature is now above 96°, the crystals
which grow in the liquid are of the monoclinic variety. They are
loi^, transparent, pale-yellow needles (Fig. 67), almost rectangular
in section, and bevelled at the points. The specific gravity is 1.96.
This form can be kept indefinitely above 96°, but, when allowed
to cool below that temperature, it slowly becomes opaque, chang-
ing into particles of rhombic sulphur.
The temperature at which a substance
changes its crystalline form is called a tTonsition
•point. It is analogous to the fusion point in the
case of a solid and a liquid; only at this one point
can both forms exist together in equilibrium.
„ y The Tmo Liquid Forms: Amorphous Sul-
phur.— When sulphur is melted, and the Uquid
is heated, two fluid, mutually soluble forms of sulphur are pro-
duced. These are known as Sx and S^ or amorphous sulphur.
As the temperature rises, the second variety increases in quantity
at the expense of the first variety. When the temperature is
lowered, the reverse change occurs:
Suf^S^ (amorphous).
It the temperature is lowered gradually, therefore, only mono-
clinic sulphur (by crystallization of the Sx) is obtained, the reac-
tion proceeding to completion in the reverse direction owing
to the removal of Sx (compare p. 239). But the change from S^
to Sx takes place only very slowly, except at temperatures near
the boiling-point. Consequently, if the liquid is quickly chilled,
by pouring into a cold vessel or into cold water, the S^ is found
as a non-crystalline substance mixed with the crystalline form.
The crystalhne form can be dissolved out with carbon disulphide,
leaving the amorphous sulphur which is not soluble. The propor-
tion of S^ varies from 3.6 per cent at 120°, to 11 per cent at 160°
and about 34 per cent at 445° (the boiling-point of sulphur). S,,
SULPHUR AND HYDROGEN SULPHIDE 251
is very viscous, so that, as its quantity iacreases, the whole Uquid
becomes thick. At 120° molten sulphur is a limpid fluid, at 260°
a vessel containing it can be inverted without loss of material.
Amorphous sidphur is a super-cooled liquid, and not a true soUd,
for true sohds are all crystalline (see p. 94). At room tempera-
ture it changes into rhombic sulphur, but so slowly that the
transformation even of a small part of it can be detected (by
treating with carbon disulphide) only after the lapse of many
months. At 100° the change is complete in less than an
hour (compare p. 242).
Elastic sulphur, — When melted sidphur is chilled, the
amorphous sulphur does not at once become hard. Sidphur
which has been heated to a high temperature, therefore, and then
suddenly cooled, consists at first of a sticky, transparent, elastic
material; called elastic or plastic sulphur. In the course of forty-*
eight hours, however, this becomes opaque and hard, because of
the separation of the crystalline and the hardening of the amor-
phous varieties.
Melting and Freezing'Points. — Amorphous sulphur, like
glass and other amorphous substances, softens when heated, but
has no sharp melting temperature. The two crystaUine forms
have different melting-points, rhombic melting to form Sx at
112.8°, and monocUnic at 119.25°. But these are difficult to ob-
serve, as the rhombic begins to turn into monocUnic above 96.5°,
and gradual transformation of Sx to S^, to produce an equihbrium
mixture of the two, occurs in both cases in the Uquid state. Hence,
the only temperature which is easy to observe is that at which
both the soUd forms melt when heated very slowly, and that at
which the Uquid freezes if cooled very slowly, namely 114.5°.
This is the so-called natural freezing-point of sulphur.
Chemical Properties. — The vapor density of sulphur indi-
cates that the vapor is a mixture of the molecules Sg, Se and Ss,
252 smith's intermediate chemistry
the former diminishing and the latter increasing in number as the
temperature is raised.
All the metals, excepting gold and platinum, combine with sul-
phur to form sulphides, and in most cases much heat is given out
during the union. Sulphur imites with Morine to give sulphur
monochloride StCl2, used in vulcanizing rubber, and bums in
oxygen to give sulphur dioxide:*
S + 02-^802.
In these compoimds the valence of an atomic weight of sulphur
appears to be one (in S2CI2) or four (in SO2). These are excep-
tional values, however, the common valences being two (in H2S,
ZnS, etc.) and six (in SO3, SO2CI2, etc.).
Moist sulphur is slowly oxidized at ordinary temperatures to
sulphuric acid :
SkeUton: S+ H2O + 02-^H2S04.
Balanced: 2S + 2H2O + 3O2 -> 2H2SO4.
In the equations, the simple formula S is used in place of a
molecular formida. The latter is needed only when questions
about the volume of the vapor are asked, and sidphur is almost
always used only in solid or melted form. Then, too, the vapor
contains several kinds of molecules, and using Sg or Sg would
introduce large and inconvenient coefficients.
Hydrogen Sulphide H2S
Occurrence. — Sulphur is a constituent of albumen, of which,
for example, the white of an egg is composed. When decay takes
place within the shell, so that air is excluded and the oxidation
which accompanies ordinary decay is prevented, the sulphur gives
hydrogen sulphide. The latter can be recognized by its odor.
Some mmeral waters contain a smaU amount in solution.
* Traces of sulphur trioxide are found at the same time. They give minute
drops of sulphuric acid, which cause a haziness in the gas when it is formed
by this action.
SULPHUR AND HYDROGEN SULPHmE
253
Preparation. — Hydrogen and sulphur combine so slowly that
at 310° the completion of the imion requires seven days. A trace
may be obtained in a few minutes by leading hydrogen over
sulphiu*, melted in a bulb (Fig. 68). A strip of paper, dipped in
lead acetate solution and placed in the wide part of the tube, is
darkened by the formation of insoluble lead sulphide PbS (black),
while acetic acid is also formed:
1
Pb(C02CH8)2 + H2S -> PbS i + 2HCO2CH8.
Fig. 68
Laboratory Method. — The gas is conunonly
made by double decomposition, using a sulphide to get the S
radical, and an acid for the H radical. Fer-
rous sulphide, made by heating iron filings
and sulphur, is the cheapest sulphide, and it
interacts easily with hydrochloric acid or sul-
phuric acid:
FeS + 2HC1 <^ FeCfc + H2S t •
The action, like all double decompositions, is
reversible. But use of an excess of hydro-
chloric acid forces it forward, and the escape
of the gaseous hydrogen sulphide reduces the
backward action almost to zero. The gas can
be made in a flask fitted like that in Fig. 25
(p. 52), or in a Kipp's automatic generator
(Fig. 69). It can be collected by upward dis-
FiG. 69 placement.
Physical Properties. — Hydrogen sulphide is a colorless gas
with an odor recalling rotten eggs. It is rather easily liqvsfied,
and the hquid boils at about —60° and freezes at —83°. The
density, imphed in the formula H2S, shows that 22.4 Uters weigh
32 + 2 or 34 g., so that the density is only one-sixth greater than
that of air (of which 22.4 1. weigh 28.95 g.). The gas is moder-
254 smith's intermediate chemistry
ately soluble in water (290 vols, in 100 vols, water at 20®), a prop-
erty which enables us to carry out many reactions of the gas upon
substances in solution.
Physiological Properties. — Care must be taken to allow as
Uttle of the gas as possible to escape into the air, and all work with
it shoidd be done in a well-ventilated hood. The proportion must
reach 1 part in 200 of air, however, before fatal results follow
breathing the mixture. The best antidote is very dilute chlorine.
Chemical Properties. — 1. The gas bums in the air, giving
water and sulphur dioxide:
Skeleton: US +02-^ H2O + SO2.
Balanced: 2H2S + 3O2 -* 2H2O + 2SO2.
2. The compoimd is not very stable. When heated, for example,
in the interior of its own flame, it is partially decomposed into free
sulphur and hydrogen. A cold porcelain dish (Kg. 70) placed in
the flame will condense some of the sulphur
on its surface.
3. On accoimt of its instabiUty, and the
ease with which it gives up hydrogen, the gas
is a reducing agent. Thus, when jars of hydro-
gen sulphide and sulphur dioxide are placed
mouth to mouth, a deposit of sulphur gradu-
ally appears:
Fig. 70 SO2 + 2H2S -> 2H2O + 3S.
Part of the free sulphur found in nature seems to be Kberated by
the action of these gases, both of which are found in volcanic
regions. The gases must be moist, for, without water vapor as a
contact agent, no interaction occurs.
In this action the sulphur dioxide loses its oxygen. We say that
the H2S was oxidized by the SO2, or that the SO2 was reduced by
SULPHUR AND HYDROGEN SULPHIDE 255
the H2S. As we have abeady noted, every reduction involves
also an oxidation.
4. The metals, down to and including silver in the activity
series, quickly receive a coating of sulphide when exposed to the
gas:
2Ag + H2S-^Ag2S + H2T.
The tarnishing of silver in the household is due to the presence of
a trace of hydrogen sulphide in the illuminating gas which escapes
from sUght leaks in the pipes.
Chewnical Properties — An Acid. — The aqueous solution
is an acidy and hence the compound is frequently called hydro-
sulphuric acid. It turns faintly tinted Utmus paper distinctly
pink. The poor conductivity of the solution shows the substance
to be Uttle ionized, and therefore a weak acid.
Like all acids, it enters into double decomposition with bases
and salts. A number of these actions are used in anal3rtical
chemistry. Thus, with cupric sulphate solution, we get cupric
sulphide (black), and with antimony trichloride antimony tri-
sulphide (orange), both as precipitates:
CUSO4 + H2S <=^ CuS i + H2SO4.
2SbCl3 + 3H2S <± SbaSs i + 6HC1.
Sulphides. — Many sulphides of metals are foimd as minerals.
Most sulphides are insoluble, and can therefore be made by double
decomposition. They may also be prepared by reduction of
sulphates. Thus, when sodium sulphate is heated on a piece of
charcoal (such as a half-burnt match) the sulphide is formed:
NaaSO* + 4C -> Na2S + 4C0.
It will be observed that in sulphides, H2S, Na2S, ZnS, CuS, and
so forth, sulphur is invariably bivalent.
Carbon Bisulphide CS2. — This compound is an important
solvent for sulphur, caoutchouc (rubber), and other substances
256 smith's intermediate chemistry
which do not dissolve in water. It is manufactured by heating
coke and sulphur together to a very high temperature.. The
sulphur (vapor) mixed with the coke combines with the latter,
and carbon disulphide passes oflf as vapor and is condensed. The
Uquid boils at 46°, and is highly inflammable:
CS2 + 302->2S02 + C02.
Large quantities are employed in the destruction of prairie
dogs and for freeing grain elevators of rats and mice.
Exercises. — 1. Write equations for the imion of aluminium
and of zinc with sulphur.
2. What experiments should you use to recognize a piece of
sulphur?
3. In what proportions by volume do, (a) sulphur dioxide and
hydrogen sulphide, (&) oxygen and hydrogen sulphide interact?
4. Write full ionic equations for the precipitation of antimony
trisulphide, and for the other double decompositions given in
this chapter.
5. Is heat evolved, or absorbed, when monocUnic sulphur
changes over to rhombic sulphur? Is heat evolved, or absorbed,
when Sx changes over to S^? Apply van't Hoff's law (p. 242).
6. Would equal weights of rhombic and monocUnic sulphur
give out equal or different amounts of heat on burning? If differ-
ent, which would give the most and which the least?
7. What would be the effect of passing hydrogen sulphide
through a red-hot tube?
8. Why is chlorine an antidote for hydrogen sulphide poison-
ing (see p. 209)?
CHAPTER XXII
OXIDES AND OXYGEN ACIDS OF SULPHUR
There are two familiar oxides^ namely sulphur dioxide or sul-
phurous anhydride SO2, and sulphur fcrioxide or sulphuric anhy-
dride SO3. Each of these dissolves in water and combines with it
to form an acid. The former gives sulphurous acid H20,S0r
or H2SO8, and the latter sulphuric acid H20,S03 or H2SO4.
Acidic and Basic Oxides. — An oxide, like carbon dioxide
CO2 (p. 336) or sulphur dioxide SO2, which combines with water
to form an acid, is said to be the anhydride of the acid. The oxides
of the non-metaUic elements, when they combine with water, in
so doing invariably form adds. In the next chapters we shall
meet with other examples (e.gr., N2O6, P2O6, Si02, etc.). On the
other hand, the oxides of metallic elements, when they are able to
combine with water, generally give ba^es (e.g. Ca(0H)2 from CaO).
For convenience, therefore, we shall often speak of an oxide as an
acidic oxide or a basic oxide, as the case may be.
Nomenclature. — The acids and salts vnthin one group are
distinguished by the terminations of, and prefixes to, their names.
Thus we have:
Hydrochloric acid HCl Sodiiun chloricfe NaCl
Hypocblorous acid HOCl Sodimn hypochlorite NaOCl
Chlorot^s acid HCIO2 Sodiimi chlorite NaC102
Chloric acid HClOs Sodiiun chlorate NaClOs
Perchloric acid HCIO4 Sodiiun perchloro^c NaC104
The proportion of oxygen to the other elements is at the basis of
the system. The terminations oils and ite indicate less oxygen
than ic and ale. The prefix hypo (Greek, below) imphes still less
257
258 smith's intermediate chemistry,
oxygen, the prefix hydro implies none at all. The per-acid con-
tains the most oxygen.
The names of compomids containing only t\yo elements (the
binary compounds) end in ide: Zinc sulphide ZnS, magnesium
nitride Mg8N2, calcium carbide CaC2, sodiimi chloride NaCl, and
the oxides CaO, etc.
Sulphur Dioxide and Sulphurous Acid
Preparation of Sulphur Dioxide SO2. — In commercial prac-
tice sulphur dioxide is obtained in three ways:
1. By burning sidphur.
2. By burning natural sulphides, such as pyrite:
Skeleton: FeS2 + O2 -> FeaOg + SO2.
Balanced: 4FeS2 + 1 IO2 -* 2Fe208 + 8SO2.
With fairly pure pyrite the combustion has only to be started,
the heat evolved in the reaction being sufficient to offset loss of
heat by radiation and to keep it going of its own accord. But
with some sidphides, Uke zinc blende ZnS, which is used as a
source of sulphur dioxide as well as of zinc, the air must be strongly
heated throughout to maintain the combustion:
2ZnS + 3O2 -> 2ZnO + 2SO2.
Forced combustion of an ore, Uke this, is called roastingi or
calcining, and is the first stage towards obtaining the metal. The
oxide is subsequently reduced by heating with coke.
3. By dropping concentrated sidphuric acid into red-hot iron
retorts:
2H2SO4 -> 2H2O + 2SO2 + O2.
4. In the laboratory a steady stream of the gas is easily obtained
by dropping hydrochloric acid upon crystals of sodimn-hydrogen
sulphite (Fig. 25, p. 52) :
NaHSOs + HCl ^ NaCl + HjSOs, (1)
H2S08?=^H20 + S02T. (2)
OXIDES AND OXYGEN ACIDS OF SULPHUR 259
This method takes advantage of the facts that sulphurous acid
is only sUghtly ionized in solution, which renders reaction 1 prac-
tically complete (see p. 241), and that this acid is imstable and
decomposes (equation 2) when there is not a large excess of water
present.
Physical Properties. — The usual six physical properties may
be noted: The gas is colorless , but has a characteristic taste and
odor. It has a density considerably greater than that of air (SOj
= 64 against 28.95). It can be liquefied below 165° (the crit.
temp.) and the Uquid boils at —8°. As the pressure required
at 20° is only 3J atmospheres, the liquid can be kept in bottles
like syphons, or in sealed tin cans. It is extremely soluble in
water (about 40 vols, in 1 vol. water at 15°). The solution is
sulphurous acid.
Chemical Properties and Uses. — Sidphur dioxide is very
stable. It combines with water giving a solution of sulphurous
acid. The gas is used in bleaching straw, silk, and wool (compare
p. 225). The bleaching action seems largely to consist in combi-
nation with the coloring matter, to give a colorless compoimd.
Hence straw hats recover the yellow color of straw by exposure
to Ught, which slowly reverses the reaction and Uberates the
sulphur dioxide.
To prevent the growth of fimgi or other organisms, wine casks
are fiunigated with sulphur dioxide before being filled. Dried
peaches and apples are prepared by exposing sUces of the fruit
on trays to sidphur dioxide. The sulphurous acid produced
bleaches the fruit, keeps insects away, and prevents the formation
of dark-colored substances during the subsequent drying.
Enormous quantities of sulphur dioxide are employed in the
manufacture of sidphuric acid and of sulphites.
Properties of Sulphurous Acid H2SO3. — Sulphurous acid,
in aqueous solution, shows all the properties of a transition add.
260 smith's intermediate chemistry
As already noted, concentrated solutions are very unstable. A
solution of sulphurous acid therefore smells strongly of sulphur
dioxide.
Being rather easily convertible into sulphuric acid H2SO4,
sulphurous acid is a redicdng agent. Thus oxygen from the air
acts slowly upon the solution :
2H2SO8 + 02^ 2H2SO4
and iodine is turned into hydrogen iodide:
H2SO3 + H2O + I2 -^ H2SO4 + 2HL
Sulphites and Bisulphites. — Sulphites are formed by
neutraUzation of sulphurous acid with a base:
2NaOH + H2SO3 ^ Na^SOs + 2H2O.
With excess of sulphur dioxide passed into the solutions of the
bases, the acid sulphites are formed:
NaOH + H2SO3 -^ NaHSOs + H2O.
Ca(0H)2 + 2H2SO3 -^ Ca(HS08)2 + 2H2O.
Such acid salts are known in commerce as bisulphiteSi because,
the proportion of the metal being half that m a sulphite, the pro-
portion of the sulphite radical is, relatively, twice as great.
They are used extensively on paper manufacture (see p. 398).
Dibasic Acids. — Acids containing two atoms of hydrogen in
each molecule are called dibasic acids. H2CO3 (p. 336), H2S,
H2SO3 and H2SO4 are such acids. Each molecule is able to
react with two molecules of a base like sodium hydroxide, as
may be seen in the first of the equations in the preceding section.
When half the quantity of the base is used, an add salt (p. 192) is
produced, as the two other equations show. Phosphoric acid
H3PO4 is a tribasic acid, and forms two series of acid salts, for
example NaH2P04 and Na2HP04. Hydrochloric and nitric
acids HNO3 are monobasic.
OXIDES AND OXYGEN ACIDS OF SULPHUR 261
Sulphur Trioxide and Sulphuric Acid
Sulphur Trioxide. — Sulphur dioxide and oxygen, when
heated together to 400°, unite very slowly with evolution of heat
to give sulphur trioxide:
2SO2 + 02^ 2SO3 + 45,200 calories.
This reaction, however, cannot be utiUzed for the manufacture
of sulphur trioxide except under special conditions, for at 400°
the union takes place far too slowly for use in industrial work,
while at higher temperatures the reverse action becomes appre-
ciable and poor yields are obtained. If we apply van't HofiF's
law to the reversible reaction: 2SO2 + 02;=^2S08 we see that,
since the forward change is exothermic, raising the temperature
will favor the backward change. In actual practice it is found
that at 400°, 98-99 per cent of the materials unite; at 700°, only
60 per cent; at 900°, practically none.
Sulphur trioxide is a white soUd which exists in two allotropic
crystaUine forms. One melts at 15°, and is therefore fluid at
ordinary temperatures. The other vaporizes without melting
at 50°. Both forms react vigorously with water, causing a hissing
noise due to the steam produced by the heat of the union:
SO3 + H2O ^ H2SO4.
Sulphur trioxide combines also with sulphuric acid to give oleums
or fuming sulphuric add, H2S2O7:
S03 + H2S04-^H2S207.
The Contact Pnpcess for Sulphuric Acid. — The interac-
tion of sulphur dioxide and oxygen is hastened by many substances,
such as glass, porcelain, ferric oxide and, more especially, finely-
divided platinum, which remain themselves unchanged and
simply act as contact or catalytic agents. The contact process,
as this is called, is now very extensively used in the manufacture
of sulphuric acid.
262 smith's intermediate chemistry
The efficiency of the contact agent depends on the amount of
surface it presents to the gases. . The action may be illustrated by
dipping asbestos in a solution of chloroplatinic acid and then
heating the mineral in the Bunsen flame:
HjPtClc^Pt + 2HC1 1 + 2CI2 T .
The platinum is thus spread in a fine grey powder on the fibers
of the asbestos. The latter is placed in a tube (Fig. 48, p. 140),
where a mixture of oxygen (or air) and sulphur dioxide may be
passed over the heated material. The sulphur trioxide issues as
vapor at the other end of the tube, where its presence is recognized
by the dense fumes (droplets of sulphuric acid), produced when
it meets the moisture in the air. The vapor can be condensed
to liquid form in a cooled flask.
In practice the contact agent employed is platinum, dispersed
in a very finely-divided condition throughout a suitable carrying
material, or base. The Grillo process uses as a base magnesium
sulphate. This gives a catalytic mass just as active as platinized
asbestos, and requires only one-hundreth part the amoimt of
platinum. With silica gel (p. 360) as a base, the platinum con-
tent of the contact mass can be still further reduced. This is a
very important point in the economics of the process.
It is absolutely necessary, in employing the contact process,
to remove from the sulphur dioxide all traces of substances such
as arsenious oxide and similar impurities derived from the cal-
cimng of pyrite or some other mineral sulphide. The most
minute quantities of these substances act as poisons on the cataly-
tic agent, and soon render it quite inoperative. The sulphur
dioxide is therefore very carefully purified before reaching the
contact chamber. Excess oxygen is used in the reacting mixture,
in order to driv^ the reaction more readily towards complete
formation of SOs (compare p. 234). The temperature in the
contact chamber is kept between 380° and 450°. The system has
a tendency to get hotter during the reaction, owing to the heat
OXIDES AND OXYGEN ACIDS OF SULPHUB 263
evolved. This tendency, if unchecked, would lead to a decreased
yield of sulphur trioxide; the cold entering gases are therefore
first led over the outside of the pipes which contain the catalyst,
in order to keep the temperature constant inside.
The issuing gases, consisting mainly of sulphur trioxide vapor
mixed with excess oxygen, are condensed by being led into 97-99
per cent sulphuric acid, the concentration of the Kquid being
maintained at this point by a regulated influx of water. If
oleum, or fuming sulphuric acid, is required, the addition of water
is omitted.
It would seem to be simpler to dissolve the gaseous sulphur tri-
oxide in water, to give sulphuric acid H2O + SO3 — > H2SO4, rather
than in 98 per cent sulphuric acid, but this cannot be done. The
mixture O2 + 2S08 is very incompletely absorbed by water.
When a bubble of this mixture enters water, the latter evaporates
into the bubble in the attempt to saturate the space occupied by
the bubble with water vapor (p. 62). The water which so evap-
orates, however, combines immediately with the sulphur trioxide
to form a fog, consisting of droplets of Kquid sulphuric acid, and
so more and more water evaporates into the bubbles. Now the
molecules of SO3, so long as they remain gaseous, move with great
velocity, namely 292 meters per second at room temperature,
and still faster in this hot gaseous mixture (see p. 90). Hence,
all the molecules that escape combination with the water vapor,
strike the wall of the bubble, and combine with the water in a few
seconds. The droplets of sulphuric acid, forming the fog, how-
ever, are not molecules but large aggregates of molecules. They
do not therefore move like the molecules of a gas, but are rela-
tively stationary. The chance of their striking the wall of the
bubble is therefore reduced enormously. Hence, after the sulphur
trioxide that escapes combination has dissolved, the droplets of fog,
carried by the excess of oxygen, can be bubbled through a whole
series of vessels of water in succession without any appreciable num-
ber of the droplets being dissolved. The same fog can be shaken in
264 SMITHES INTERMEDIATE CHEMISTRY
a flask with water, violently and continuously, without any appre-
ciable solution. When the water is thrown, by the shaking,
through the oxygen, the oxygen is spUt up by the water, and driven
about, but the fog particles move with the oxygen, so that the
water never reaches them. On the other hand, when the mixture
of gases bubbles through 97-99 per cent sulphuric acid, as is done
in practice, there is practically no water available for evaporation,
the sulphur trioxide remains gaseous, and its rapidly moving
molecules in a few seconds have all plunged into the sulphuric
acid and combined with it, either uniting with the 1 to 3 per cent of
water present, or, when oleum is made, uniting with the sulphuric
acid to form H2S2O7.
This case afifords an admirable illustration of the importance of
physics in practical chemistry (p. 22). The chemical reaction
occurs with water, but the physical condition of the fog of sulphuric
acid prevents its dissolving and, if water were used in a factory, a
large proportion of the sulphuric acid would pass with the excess of
oxygen into the air and be lost. In fact, it would kill vegetation,
and make life unbearable in the neighborhood.
Chamber Process for Sulphuric Acid. — Although salts of
sulphuric acid, such as calcium sulphate CaS04, are exceedingly
plentiful in nature, the preparation of the acid by chemical action
upon the salts is not practicable. The sulphates, indeed, inter-
act with all acids, but the actions are reversible. The completion
of the action by the plan used in making hydrogen chloride (p. 126),
involving the removal of the sulphuric acid by distillation, would
be difficult on account of the involatiUty of this acid. It boils at
330°; and suitable acids, less volatile still, which might be used to
liberate it, do not exist. We are therefore compelled to build up
sulphuric acid from its elements.
The gases, the interactions of which result in the formation of
sulphuric acid, are: water vapor, sulphur dioxide, nitrous anhy-
OXIDES AND OXYGEN ACIDS OF SULPHUR 266
dride N2O8* (see p. 314), and oxygen. These are obtained, the
first by injection of steam, the second usually by the burning of
pyxite, the third from nitric acid HNO3, and the fourth by the
introduction of air. The gases are thoroughly mixed in large
leaden chambers, and the sulphuric acid forms droplets which fall
to the floors. In spite of elaborate investigations, instigated by
the extensive scale upon which the manufacture is carried on and
the immense financial interests involved, some imcertainty still
exists in regard to the precise nature of the chemical changes which
take place. According to Lunge, supporting the view first sug-
gested by Berzelius, the greater part of the product is formed by
two successive actions, the first of which yields a complex com-
pound that is decomposed by excess of water in the second:
0-H
H2O + 2SO2 + N2O8 + 02-^ 2SO2 :: (1)
^O-NO
The group —NO, nitrosyl, is found in many compounds. Here,
if it were displaced by hydrogen, sulphuric acid would result.
Hence this compound is called nitrosylsulphuric acid:
^0-H ^OH
2S02^ +H20^2S02^ +N2O3. (2)
^0-NO ^OH
The equations (1) and (2) are not partial equations for one inter-
action, but represent distinct actions which can be carried out
separately. In a properly operating plant, indeed, the nitrosyl-
sulphuric acid is not observed. But when the supply of water is
deficient, white " chamber crystals,'' consisting of this substance,
collect on the walls.
The explanation of the success of this seemingly roundabout
method of getting sulphuric acid is as follows: The direct union of
* This gas is unstable, breaking up in part into nitric oxide NO and nitro-
gen tetroxide NO2: N2O8 t^ NO + NO2. In this process, however, the mix-
tiure behaves as if it were all N2O8, and so only nitrous anhydride is named in
this connection.
266 smith's intermediate chemistry
sulphur dioxide and water to form sulphurous acid is rapid, but the
action of free oxygen upon the latter, 2H2SO3 + 02-^ 2H2SO4, is
exceedingly slow. Reaching sulphuric acid by the use of these two
changes, although they constitute a direct route to the result, is
not feasible in practice. On the other hand, both of the above
actions, (l) and (2), happen to be much more speedy, and so, by
their use, more rapid production of the desired substance is secured
at the expense of a sUght complexity.
The progress of the first action is marked by the disappearance
of the brown nitrous anhydride and, on the introduction of water,
the completion of the second stage results in the reproduction of
the same substance. The nitrous anhydride takes part a large
number of times in these changes, and so facilitates the conversion
of a great amount of sulphur dioxide, oxygen, and water into sul-
phuric acid, without much diminution of its quantity. Some
is imavoidably lost, however.
The loss of nitrous afihydride is made good by the introduction
of nitric acid vapor into the chamber. This acid is made from con-
centrated sulphuric acid and commercial sodium nitrate NaNOs:
NaNOs + H2SO4 T± HNOs t + NaHS04.
On account of the volatiUty of the nitric acid, a moderate heat is
suflScient to remove it from admixture with the other substances,
and its vapor is swept along with the other gases into the appara-
tus. The first reaction which this vapor undergoes may be
written:
H2O + 2SO2 + 2HN08 -^ 2H2SO4 + N2O8.
Details of the Chamber Process. — The sulphur dioxide is
produced in a row of furnaces A (Fig. 71). The gases from the
various furnaces pass into one long dust-flue, in which they are
mingled with the proper proportion of air, and deposit oxides of
iron and of arsenic, and other materials which they transport
mechanically. From this flue they enter the Glover tower G,
in which they acquire the oxides of nitrogen. Having secured
OXIDES AND OXYGEN ACIDS OP SULPHUK 267
all the necessary constituents, excepting water, the gases next
enter the first of the lead chambers, large structures lined com-
pletely with sheet lead. These measure as much as 100 X 40 X
40 feet, and have a total capacity of 150,000 to 200,000 cubic
feet. As the gases drift through these chambers they are thor-
FiQ. 71
ougMy mixed, and an amount of water considerably in excess
of that actually required is injected in the form of steam at various
points. The acid, along with the excess of water, condenses and
collects upon the floor of the chamber, while the unused gases,
chiefly nitrous anhydride and nitrogen, the latter derived from
the air originally admitted, find an exit into the Gay-Lussac
tower L.
This is a tower about fifty feet in height, filled with tiles, over
which concentrated sulphuric acid continually trickles. The
object of this tower, to catch the nitrous anhydride and enable it
to be reemployed in the proce^, is accomplished by a reversal of
action (2) above. The acid which accumulates in the vessel at
the bottom of this tower contains, therefore, nitrosylsulphuric
268 smith's intermediate chemistry
acid, and by means of compressed air this acid is forced through
a pipe up to a vessel at the top of the Glover tower G. A neigh-
boring vessel at the top of this tower is filled with dilute sulphuric
acid, and when the contents of both vessels are mixed by allow-
ing their contents to trickle down through the tower, nitrous
anhydride is once more set free by the interaction of the water in
the dilute acid (action (2)). The Glover tower is filled with
broken flint or tiles, in order that the descending hquid may ofifer
a large surface to the hot gases ascending the tower, and thereby
facilitate the acquisition by these gases of a sufficient quantity
of nitrous anhydride. Their high temperature also causes a con-
siderable concentration of the diluted sulphuric acid as it trickles
downward. This acid, after traversing this tower, is sufficiently
strong to be used once more for the absorption of nitrous anhy-
dride.
To replace the part of the nitrous anhydride which is inevitably
lost, fresh nitric acid is furnished by small open vessels N, contain-
ing sodium nitrate and sulphuric acid, placed in the flues of the
pyrite-bumers. About 4 kg. of the nitrate are consumed for every
100 kg. of sulphur.
The acid which accumulates upon the floors contains but 60 to
70 per cent of sulphuric acid, and has a specific gravity of 1.5-
1.62.
This crude sulphuric acid is appHcable directly in some
chemical manufactures, such as the preparation of superphos-
phates (p. 412) . For many purposes, however, a more concentrated
acid is required. Concentration is effected by evaporating off
water from the chamber acid in pans fined with lead, which are
frequently placed over the pyrite-bumers in order to econoniize
fuel. The evaporation in lead is carried on imtil a specific grav-
ity 1.7, corresponding to 77 per cent concentration, is reached.
Up to this point the sulphate of lead formed by the action
of the sulphuric acid produces a crust which protects the metal
from further action. When a still more concentrated acid is
OXIDES AND OXYGEN ACIDS OF SULPHUR 269
wanted other methods of driving off the water, such as the cas-
cade system or the Gaillard tower, must be employed.
The cascade system consists of a series of small siUca or silicon-
iron basins set over an inclined flue, and so placed that each basin
delivers by a spout into the one below. The flue is heated by
gas or coke firing at the lower end, and dilute acid is fed continu-
ously into the basin at the top end. As the acid passes from
basin to basin, it meets hotter and hotter conditions and becomes
more and more concentiated. ^
The Gaillard plant consists essentially of a large hollow tower
built of acid-resisting stone, and filled with small fragments of
similM- material. Dilute acid is sprayed in at the top, and meets
hot furnace gases injected in at the bottom. Most of the water
contained m the acid is carried off by these gases, and concen-
trated acid collects at the bottom of the tower.
A more convenient method of obtaining very concentrated acid,
which avoids the difliculties of evaporation entirely, is to add to
the chamber acid the requisite quantity of oleum, prepared by the
contact process already described. Commercial sulphuric acid,
oil of vitrioli has a specific gravity 1.83-1.84, and contains about
93.5 per cent H2SO4.
Physical Properties of Sulphuric Acid. — The pure acid is a
colorless, oily liquid of sp. gr. 1.84, which freezes to a sohd at 10°.
It mixes with water in all proportions, and much heat {heat of
solviioTij see p. 244) is given out when it dissolves. It boils at 330°,
but the vapor is largely decomposed into free water and sulphur
trioxide, which recombine when it cools.
Chemical Properties. — 1. The acid is more stable than sul-
phurous acid, but has a slight tendency to lose SO3 even at ordi-
nary temperatures and as already noted, decomposes largely at
the boiling-point.
2. In aqueous solution, sulphuric acid is much more active
270 smith's intermediate chemistry
as an add than is sulphurous acid, but is somewhat inferior in
this respect to hydrochloric HCl and nitric acids HNO3. Like
other active, soluble acids, its solution turns Htmus red, gives
hydrogen upon addition of active metals, and enters into double
decomposition with bases and salts. Thus, insoluble barium
sulphate is obtained as a white precipitate by the action of dilute
sulphuric acid on any soluble barium salt:
BaCl2 + H2SO4 ^ BaS04 i + 2HC1.
Any soluble sulphate will, V course, give the same precipitate
with barium chloride, and the action is used as a test for this ion.
Some other salts of barium are also insoluble in water, but the sul-
phate is recognized by the fact that it is too insoluble to be acted
upon by dilute pure hydrochloric acid or nitric acid. The other
insoluble salts of barium interact with these acids and dissolve.
The addition of one of these acids is therefore part of the test for
SO4' ion.
On accoimt of its high boiUng-point, the double decompositions
of the concentrated acid can be used for preparing more volatile
adds:
NaCl + H2SO4 ^ NaHS04 + HCl t (gas at room temp.j.
NaNOs + H2SO4 ^ NaHS04 + HNO3 T (volatile at 86°).
3. Concentrated sulphuric acid combines with water to form a
stable hydrate H2S04,H20. Hence it removes the elements of
water from many substances containing hydrogen and oxygen,
and is called a dehydrating agent. Thus, paper (cellulose),
moistened with the acid and warmed, turns black from the hbera-
tion of carbon. Sugar (C12H22O11) is decomposed even more easily:
C12H22O11 ^ 12c + IIH2O.
4. Finally, concentrated sulphuric acid acts as an oxidizing
agent. Sulphur and carbon, boiled in it, are oxidized:
2H2SO4 + S -> 3SO2 + 2H2O.
2H2SO4 + C ^ 2SO2 + 2H2O + CO2.
OXIDES AND OXYGEN ACIDS OF SULPHUR * 271
The reducing action of HBr and HI on sulphuric acid has abeady
been noted (pp. 202, 204). The more active metals, like zinc,
reduce it to hydrogen sulphide, the less active, like copper, give
sulphur dioxide. Hydrogen is not hberated, because practically
no hydrogen-ion is present in concentrated sulphuric acid. Grold
and platinum alone are not attacked.
4Zn + 5H2SO4 -^ 4ZnS04 + 4H2O + H2S.
Cu + 2H2SO4 ^ CUSO4 + 2H2O + SO2.
Uses of Sulphuric Acid. — The acid has innumerable applica-
tions, some of which will be taken up in detail in later chapters.
It is employed in almost every chemical industry, something
Hke 6,000,000 tons being produced yearly in the United States
alone. It is used in the manufacture of sulphates, hydrochloric
acid, nitric acid, sodium carbonate, etc., in making fertiUzers and
dyes, in bleaching, electroplating and so on. Its dehydrating
power is especially valuable in making explosives (pp. 481-2).
Other Oxygen Acids of Sulphur. — Many other oxygen
acids of sulphur exist, such as hyposulphurous add H2S2Q4 and
persidphuric acid H2S2O8. The acids themselves are very imstable,
and cannot be isolated in the pure state. Some of the salts,
however, are in common use, and will be dealt with later under
their several positive radicals.
When acid sulphates, such as NaHS04, are heated, water is
given oflf and a pyrosulphate (Greek prefix, fire) remains.
2NaHS04 -^ Na2S207 + H2O.
The pyrosulphates are salts of oleum, or fuming sulphuric acid,
which has already been mentioned. Oleum possess all of the
dehydrating and oxidizing powers of sulphuric acid in an accen-
tuated form, and is widely used in the industries on accoimt of
these properties.
272 smith's intermediate chemistry
Exercises. — 1. Which contains more oxygen: (a) a phosphate'
or a phosphite; (b) a nitrite or a nitrate; (c) a borate or a per-
borate? Name the acids corresponding to these six salts.
2. Make equations for: (a) the roasting of stannic sulphide
(SnSs) giving SnOa; (b) the action of concentrated sulphuric acid
on silver giving silver sulphate and SO2; (c) the dissociation of
sulphuric acid vapor.
3. What are the formulae of magnesium sulphite and bisulphite,
respectively?
4. Give two reasons why boiling sulphuric acid, when spilt upon
the flesh, causes most serious bums.
5. By what facts or tests could you recognize concentrated
sulphuric acid?
6. Why is the wooden laboratory shelf commonly " burned "
where the sulphuric acid bottle stands?
7. What would be the reaction between sodium sulphite and
sulphuric acid?
8. How could you distinguish between, and recognize, the sul-
phide, sulphite, and sulphate of potassium?
9. Why is it not desirable to make chamber sulphuric acid of a
concentration higher than 60-70 per cent?
10. Justify the nomenclature in the case of hyposulphurous
and persulphuric acids.
CHAPTER XXIII
THE PERIODIC SYSTEM
In an earKer chapter (p. 208) we saw that the elements fluorine,
chlorine, bromine and iodine exhibited striking similarities in
their chemical properties, and we grouped these four elements
together under the name of the halogen family. Now there are
two rather rare elements, selenium and tellurium, which resemble
sulphur very markedly in their chemical properties. Both give com-
poimds with hydrogen, hydrogen selenide H2Se and hydrogen
telluride H2Te, corresponding with hydrogen sulphide H2S, but
less stable. Both give compoimds with oxygen, selenium dioxide
Se02 and tellurium dioxide Te02, corresponding with sulphur
dioxide SO2. These dioxides dissolve in water to form weak
acids similar to sulphurous acid. These acids, again, can be oxi-
dized to yield selenic acid H2Se04 and telluric acid H2Te04, analo-
gous to sulphuric acid H2SO4. All these compounds show grad-
ations in properties, as we go upwards in atomic weights from
S to Te, which are very reminiscent of the gradations encountered
in a series of halogen compounds, such as HCl, HBr, HI.
Furthermore, just as we have in the halogen family a first mem-
ber with rather irregular habits, fluorine, so we note in the sulphur
family a corresponding light element showing decided peculiari-
ties, oxygen. Oxygen forms a compound with hydrogen H2O,
which is akin to H2S in being a weak acid and, as might be pre-
dicted, is much more stable. The metaUic oxides are very sim-
ilar in their properties to the metaUic sulphides. Ozone may be
regarded as oxygen dioxide OO2, analogous to sulphur dioxide SO2.
The family resemblance in other compounds, however, is more diffi-
cult to trace.
273
274 smith's intekmediate chemistry
The question naturally arises: can we group aM of the elements
into famiUes like the halogen family and the sulphur family?
K so, then we shall Ughten considerably the burden of chemical
facts that we need to remember, for the behavior of one element
in a family will suggest to us immediately how the other members
of the same group will act under similar conditions. Classifica-
tion of this kind is part of the method of science, and furnishes
a very useful guide in investigation.
Metallic and Non- Metallic Elements. — Thus far we have
found the division into metaUic and non-metaUic elements very
serviceable for classification in terms of chemical relations. The
metallic or positive elements (p. 54), form positive radicals
and ions containing no other element (c/. p. 171). Thus
the metals., give sulphates, nitrates, carbonates, and other salts,
which furnish a metaUic ion, such as Na+ or K+, together with
the ions SO4", NOs", and COa". Their hydroxides, KOH, Ca-
(OH2), etc., give the same metaUic ion, and the rest of the mole-
cule forms hydroxide-ion. That is to say, their hydroxides are
bases and their oxides are basic. The metaUic elements often
enter, but only with other elements, into the composition of a
negative ion, as is the case with manganese in K.Mn04, with chro-
mium in K2.Cr207, and with sUver in K.Ag(CN)2.
The non-metallic or negative elements are found chiefly in
negative radicals and ions. They form no nitrates, sulphates, car-
bonates, etc., for they could not do so without themselves alone
constituting the positive ion. We have no such salts of sulphur,
carbon, or phosphorus, for example. Their hydroxides, although
their formulae may be wriUen CIO2OH, P(0H)3, S02(OH)2, fur-
nish no hydroxyl ions, as this would involve the same conse-
quence. These hydroxides are divided by dissociation, in fact, so
that the non-metal forms part of a compound negative radical, and
the other ion is hydrogen-ion, CIO3.H, PO3H.H2, SO4.H2. Their
oxides are acidic. Their halogen compounds, Uke PBrg (p. 201)
THE PERIODIC SYSTEM 276
and S2CI2 (p. 252), are completely decomposed by water, and the
actions are not, in general, reversible. The hahdes of the typical
metals are not hydrolyzed. (see p. 369), and with those that are not
typical, the action is reversible.
The distinction is not perfectly sharp, however. Thus, zinc
gives both salts Uke the sulphate, Zn.S04, and chloride, Zn.CU, and
compoimds Uke sodium zincate Na.HZn02.
Classification by Atomic Weights. — Newlands (1863-4) dis-
covered a surprising regularity that became apparent when the
elements then known were placed in the order of ascending atomic
weight. Omitting hydrogen (at. wt. 1) the first seven were:
lithium (7), gluciniun (9), boron (11), carbon (12), nitrogen (14),
oxygen (16), fluorine (19). These are all of totally different
classes, and include first a metal forming a stongly basic hydroxide,
then a metaUic element of the less active sort, then five non-metals
of increasingly negative character, the last being the most active
non-metal known. The next element after fluorine (19) was
sodium (23), which brings us back sharply to the elements that
form strongly basic hydroxides. Omitting none, the next seven
elements were sodium, (23), magnesiiun (24.4), aluminiiun (27),
siUcon (28.4), phosphorus (31), sulphur (32), chlorine (35.5).
In this series there are three metals of diminishing positive char-
acter, followed by four non-metals of increasing negative activity,
the last being a halogen very Uke fluorine. On accoimt of the
fact that each element resembles most closely the eighth element
beyOnd or before it n the Ust, the relation was caUed the law of
octaves. After chlorine the octaves become less easy to trace.
That this periodicity in chemical nature is more than a coinci-
dence is shown by the fact that the valence and even many phys-
ical properties, such as the specific gravity, show a similar fluc-
tuation in each series. In the first two series the compounds with
other elements are of the types:
276 smith's intermediate chemistry
liCl GICI2 BCI3 CCI4; CH4 NH3 OH2 FH
U2O GIO B2O3 CO2 N2O5 — —
NaCl MgCls AICI3 SiCU; SiH4 PHs SH2 CIH
NaaO MgO AI2O8 Si02 P2O5 SOs CI2O7
Thus the valence towards chlorine and hydrogen ascends to four
and then reverts to one in each octave. The highest valence,
shown in oxygen compounds, ascends from hthium to nitrogen
with values one to five, and then fails because compounds are
lacking. In the second octave, however, it goes up continuously
from one to seven.
Again, the specific gravities of the elements in the second series,
using the data for red phosphorus and Uquid chlorine, are:
Na 0.97, Mg 1.75, Al 2.67, Si 2.49, P 2.14, S 2.06, CI 1.33.
Mendelejejps Scheme. — In 1869 Mendelejeff pubhshed an
important contribution towards adjusting the difficulty which the
elements following chlorine presented, and developed the whole
conception so completely that the resulting system of classifica-
tion has been connected with his name ever since. The table
on page 278, in which the atomic weights are expressed in round
nimibers, is a modification of one of Mendelejeff's, extended to
include elements more recently discovered.
The chief change made by Mendelejeff from the arrange-
ment in 'simple octaves is that the third series, beginning
with potassiimi, is made to furnish material for two octaves, potas-
sium to manganese and copper to bromine, and is caUed a long
series. The valences fall in with this plan fairly well. Copper,
while usually bivalent, forms also a series of compounds in which
it is univalent. Iron, cobalt, and nickel fall between the two
octaves, and cannot be acconmiodated in either.
Every long series contams three elements of this character,
closely resembUng one another. As will be seen from the table
these transition elements, as they are called, may be placed together
THE PERIODIC SYSTEM 277
in an eighth group. At the time Mendelejefif made the table,
three places in the third, long series had to be left blank, as a tri-
valent element [Sc] was lacking in the first octave of the series, and
a trivalent [Ga] and a quadrivalent one [Ge] in the second. These
places have since been filled, as we shall presently see.
The fourth series, which is a long series exactly similar to the
third, contained many blanks at the time of Mendelejeff, but is
now nearly complete. It begins with an active alkaU metal,
rubidiiun, and ends with iodine, a halogen. The rule of valence
is strictly preserved throughout the series, and in general the
elements fall below those which they most closely resemble.
The fifth series is still somewhat incomplete, but the order
of the atomic weights and the valence enable us satisfac-
torily to place most of those elements which are known. The
chemical relations to elements of the fourth series justify the
position assigned to each. Csesiiun, for example, is the most
active of the alkaU metals; bariinn has always been classed
with strontium, and bismuth with antimony.
The sixth and last series contains only a few radioactive ele-
ments. No element with an atomic weight greater than that of
uranium (238) has yet been discovered.
The most important change made in the table since the time of
Mendelejeff is the addition of another group, the family of the
inert gases of the atmosphere (see p. 296). These elements were
unknown before 1894, but fall logically into a new group at the
left hand side of the table as here given.
278
smith's intermediate chemistry
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THE PERIODIC SYSTEM 279
•
General Relations in the System. — In every octave the
valence towards oxygen ascends from one to seven, while that
towards chlorine and hydrogen ascends to four and then reverts
to one. The long series octaves therefore exhibit the same periodic
changes with respect to valence as do the short series octaves
already discussed (p. 276). Furthermore, the elements in the
new group on the left hand side of the table fall directly into line
with the rest by exhibiting zero valence. The inert gases, in other
words, form no compounds with other elements. The transition
elements on the right hand side of the table, similarly, justify
their position by forming a few compoimds in which a valence
of eight is shown, for example " osmic acid " OSO4. It must be
admitted, however, that lower valences are more frequently
displayed by these transition elements.
The physical properties, both of the elements themselves and of
corresponding compounds, fluctuate withm the Umits of each
series in the same way. Thus the meUing-points of the first
eight elements in the third series are as follows:
A - 188°, K 62°, Ca 810°, Sc - , Ti 1790°, V 1720°, Cr 1520°, Mn 1260°.
All of the elements in the same column do not show the same
degree of resemblance. We find, instead, that there are two
well-defined famiUes in each of the columns forming the octaves.
In each long series the element in the first octave falls into one of
these famiUes, the element in the second octave into the other.
In the table on p. 278 these two famiUes are differentiated in each
colimm by spacing one towards the left and the other towards
the right of the available space. Thus in the second colunm of
the table we have the family of the alkali metals (Li,Na,K,Rb,Cs)
and the copper family (Cu,Ag,Au). The members of the first
family, and their corresponding compounds, are all strikingly
similar in physical and chemical properties. The members of the
second family, on the other hand, have httle in common with those
of the first except in their valence, and even here abnormal values
280 smith's intermediate chemistry
are shown in well-known compounds such as cupric salts, which
contain the radical Cu".
The inert gases, on the left of the octaves, constitute a single
family. As for the transition elements on the right, while the
three members in any one series resemble one another in. many
respects, yet a closer relationship between elements in different
series, according to the vertical arrangement shown in the table,
dividing the group into three f amihes, is also evident.
If we examine the physical properties of successive elements,
or corresponding compounds of successive elements, of any one
family we find a uniform gradation observable, just as in the cases
of the halogens and their hydrogen compounds studied in chapter
XVII. Thus the melting-points of the alkaU metals are as follows :
U 186°, Na 96°, K 62°, Rb 38°, Cs 26°.
As yet no exact mathematical (quantitative) relation between
the values for any property and the values of the atomic weights
has been discovered; only a general (qualitative) relationship can
be traced. Anticipating the discovery of some more exact mode
of stating the relationship in each case, and remembering that
similar values of each property recur periodically, usually at inter-
vals corresponding to the length of an octave or series, the prin-
ciple which is assimied to underhe the whole, the periodic law,
is stated thus: All the properties of the elements are periodic
functions of their atomic weights.
That the chemical relations of the elements vary just as do the
physical properties of the simple substances is easily shown.
Thus, each series begins with an active metalUc (positive) element,
and ends with an active non-metaUic (negative) element, the inter-
vening elements showing a more or less continuous variation
between these Hmits. Again, the elements at the top are the least
metallic of their respective columns. As we descend, the members
of each group are more markedly metaUic (in the first columns), or,
what is the same thing, less markedly non-metaUic (in the later
colimms).
THE PERIODIC SYSTEM 281
Applications of the Periffdic System. — The system has
found application chiefly in four ways:
1. In the prediction of new elements. Mendelejefif (1871) drew
attention to the blank then existing between calcimn (40) and
titanium (48). He predicted that an element to fit this place
would have an atomic weight 44 and would be trivalent. From
the nature of the surrounding elements, he very cleverly deduced
many of the ph3^cal and chemical properties of the unknown
element and of its compounds. In 1879 Nilson discovered scan-
dium (44), and its behavior corresponded closely with that pre-
dicted. Mendelejeff described accurately two other elements, like-
wise unknown at the time. Inl875 Lecoque de Boisbaudran found
gaUium, and in 1888 Winkler discovered germanimn, and these
blanks were filled.
2. By enabling us to decide on the correct values for the atomic
weights of some elements, when the equivalent weights have been
measured, but no volatile compound is known (c/. pp. 77 and 86).
Thus, the equivalent weight of indium was 38 and, as the element
was supposed to be bivalent, it received the atomic weight 76. It
was quite out of place near arsenic (75), however, being decidedly
a metaUic element. As a trivalent element with the atomic weight
115, it fell between cadmium and tin. Later work fully justified
the change. More recently, when radiiun was discovered, it was
found to have the equivalent weight 113 and to resemble barimn.
Consequently we assiune that, Uke barium, it is bivalent, and
asgign it a vacant place under this element, in the last series.
3. By suggesting problems for investigation. The periodic
system has been of constant service in the course of inorganic
research, and has often furnished the original stimulus to such
work as well.
For example, the atomic weights of the platinum metals at first
placed them in the order, Ir (197), Pt (198), Os (199), although
the resemblance of osmium to iron and ruthenium would have
led us to expect that this element should come first. For similar
282 smith's intermediate chemistry
reasons platinum should have come last, mider palladium. A re-
investigation of the atomic weights, suggested by these consider-
ations, was undertaken by Seubert, and the old values were
found in fact to be very inaccurate. He obtained:
Os = 191, Ir = 193, Pt = 195.
Originally feod, although it fell in the fourth column, possessed
only one compound Pb02 in which it seemed to be undoubtedly
quadrivalent. Search for salts of the same form, however, speed-
ily yielded the tetrachloride PbCU, tetracetate, and many others.
The existence of osmic acid OSO4, and a corresponding compound
of ruthenium, suggests that other compounds of the elements of
the eighth group, displaying the valence eight, may be capable
of preparation. The collocation of copper, silver, and gold, in
the same column with the alkah metals, is not at present per-
fectly satisfactory, and suggests the advisabihty of strengthening
their position, if possible, by further investigation.
In the same way, incorrect values of many physical properties
have been detected, and have been rectified by more careful work.
4. By furnishing a comprehensive classification of the elements,
arranging them so as to exhibit the relationships among the physi-
cal and chemical properties of the elements themselves and of their
compounds. Constant use will be made of this property of the
table in the succeeding chapters. Having disposed of the halogen
and sulphur famiUes, situated, respectively, in the eighth and
seventh colimms of the table on p. 278, we shall presently take up
nitrogen and phosphorus from the right side of the sixth column.
Then from the fifth column, we shall select carbon and siUcon,
and from the fourth boron, leaving the other, more decidedly
metaUic elements for later treatment.
Defects in the Periodic System. — The periodic system is
often described as if it furnished a classification of tHe properties
THE PERIODIC SYSTEM 283
of chemical substances which was complete in its scope, and ideal
in its exactness. This, however, is far from being the case.
The order of activity of the metals (p. 54) and of the non-metals
(p. 209) smnmarizes many properties, and explains many featm'es
of the chemical behavior of the elements. This hst is scattered
through the periodic table (compare both), without any trace of
regularity.
The periodic system concentrates attention too largely on one
of the valences of each element. Thus, for manganese, it focuses
attention on the septivalent form in the permanganates. But
manganous salts are more hke the ferrous, the cobaltous, the
chromous, and other sets of salts, none of which are in the same
column of the table. Similarly, the manganic salts are hke the
ferric salts and the salts of aluminiiim. Again, copper is univalent
in one series of salts, but in its better known salts it is bivalent.
Silver, which belongs to the same periodic family, is always
univalent, while gold, also in the same family, is univalent or
trivalent, and in the latter case is almost wholly a non-metaUic
element. If it were possible to place each element in several
different columns, one for each of the valences that it shows, the
table would then include far more of the properties of the elements.
But this cannot be done, for, according to the order of magnitude
of the atomic weights, there is but one place for each element.
In other words, the periodic system largely ignores the variety
of different classes of chemical relations which an element with
several valences always shows.
The position of hydrogen in the system is still a matter of dis-
pute. It is more famiUar to us as a univalent positive radical
resembhng the alkah metals in forming compounds with negative
radicals such as chlorine, but it can also function as a univalent
negative radical, resembling the halogens in forming hydrides
with the alkah metals which are analogous to chlorides. Most
chemists shelve the difficulty by giving hydrogen a position all to
itself at the top of the table.
284 smith's intermediate chemistry
Between cerium (140) and tantalum (181.5) in the fifth series,
there occur twelve rare elements, called the elements of the rare
earths, which have been omitted entirely from the Mendelejefif
system. What is to be done with these elements is a point on
which agreement has not yet been reached.
Finally, reference to the table will show that in three cases a
sUght displacement of the order of the elements according to
atomic weights is necessary. Argon, an inert gas, is placed before
potassiimi, an alkah metal, although its atomic weight is 0.8
higher. Cobalt is put before nickel because it resembles iron
more closely. Tellurium and iodine are placed in that order to
bring them into the sulphur and halogen groups, respectively.
Their valence and other chemical relations both require this.
These three cases constitute undoubted exceptions to the Mendele-
jeff system of classification. The general agreement, however, is
obviously far too remarkable to be due entirely to chance.
In a later chapter it will be shown that recent work on atomic
stritcture throws considerable hght on the several abnormahties
discussed above, supplying us with a more logical basis for the
periodicity exhibited by the elements in respect to valence and
other properties than is furnished by the use of Mendelejefif's
system alone. Nevertheless the latter will be found to be
of valuable service to us throughout the remainder of the
book.
Exercises. — 1. There is a blank at the end of the fifth long
series, where we should expect to find another halogen (see p.
278). If the element that should fill this blank were to be dis-
covered, what would be its physical and chemical properties?
What would be the properties of its compound with hydrogen?
2. How should you attempt to obtain H2Te, and what physical
and chemical properties should you expect it to possess?
3. Make a Ust of bivalent elements and criticize this method of
gh)uping as a means of chemical classification.
THE PERIODIC SYSTEM 285
4. Write down the symbols of the elements in the fourth series
(that begmning with rubidiimi, and ending with iodine) on p.
278. Record the valence of each element toward oxygen, using
for reference the chapters in which the oxygen compounds are
described.
CHAPTER XXIV
NITROGEN. THE ATMOSPHERE
It is time now to return to the atmosphere, of which the most
active component, oxygen, has already been discussed. The
other chief component, nitrogen, will lead us to ammonia NHj
and nitric acid HNOs, both of which are of great commercial
importance, and have interesting derivatives.
Occurrence of Nitrogen. — Aside from the free nitrogen,
which forms nearly fom'-fifths of the bulk of atmospheric air,
much nitrogen is found in nature in combination.
Potassiimi nitrate KNOs is formed in the soil by the
action of bacteria upon animal matter, and sgdium
nitrate NaNOa is obtained from an immense deposit
in Peru and ChiU. Nitrogen is an essential constitu-
ent oTan important class of organic substances called
the vroteins, which are found in plants, particularly
in the fruit, and in the muscles and other tissues of
the animal body.
Preparation. — Nitrogen may be obtained from
the air by simply removing the oxygen. This
nitrogen is not pure, however, as it retains about one
per cent of other gases — the " inert gases " of the
atmosphere. The oxygen can be removed by allowing pieces
of moist phosphorus (Fig. 72) slowly to oxidize in an enclosed
specimen of air. The phosphoric acid H8PO4 and other products
of the oxidation of the phosphorus dissolve in the water.
Pure nitrogen can be obtained from pure compounds of nitrogen.
Thus, ammonia gas may te passed over heated cupric oxide
286
Fig. 72
NITROGEN. THE ATMOSPHERE 287
(Rg. 73), and the water removed by bubbling the gas through sul-
phuric acid.
Skeleton: NH3+ CuO-^ Cu+ H2O + N2.
Balanced: 2NH3 + 3CuO -^ 3Cu + 3H2O + N2.
A steady stream of nitrogen is most easily made by heating
sodium nitrite and ammoniimi chloride very gently along with a
Uttle water in a flask:
NaNOa + NH4CI ^ NaCl + NH4NO2 -^ 2H2O + N2 1 •
The double decomposition is reversible, and the first action might
be expected to be only partially com-
pleted. But the ammoniiun nitrite *====BCIIiSiMCiZZSS====»
NH4NO2 is imstable, and decomposes Fig. 73
as fast as it is formed, so that one of
the substances required to reverse the first reaction is removed, and
the reversing action does not occur.
Physical Properties. — Nitrogen is a colorless, tasteless, and
odorless gas. Its density is indicated in the formula N2 (mol. wt.
2 X 14 = 28). It is very litde soluble in water. When liquejied
it boils at - 194°.
Chemical Properties. — Nitrogen is chemically a rather
indifferent gas. It unites easily with a very few elements, notably
some of the most active metals, such as calciiun and magnesium.
When magnesium bums in the air, the white powder which is
formed contains some of the nitride of magnesiiun Mg3N2, along
with much of the oxide:
3Mg + N2 -^ Mg8N2.
The presence of the nitride may be shown by the odor of ammo-
nia, given off when the ash is moistened with water:
MgsNa + 6H2O -> 3Mg(OH)2 + 2NH, t •
288
SMITH'S INTERMEDIATE CHEMISTRY
The compounds with oxygen, such as NO and HNOs, and with
hydrogen such as NHs, are of immense commercial value, but,
not being very stable, they are formed only in traces by direct
union of the elements. The processes for utiUzing these tenden-
cies to union, feeble as they are, for manufacturing purposes, will
be described under the compounds themselves.
The Atmosphere
The components of the air may be conveniently divided into
regvlar components and accidental components. The regular
components, again, consist of three which are present in practi-
cally the same proportions in all samples, and three (namely
water, carbon dioxide and dust) which vary markedly in quan-
tity.
Components Present in Constant Proportions. — The
components whose proportions are practically invariable are
nitrogen, oxygen, and the group of inert gases. When the vari-
able components are removed, the proportions of the constant
ones are as follows:
Nitrogen
Oxygen..
Argon . . .
By Volume
78.06
21.00
0.94
By Weight
75.5
23.2
1.3
The inert gases, excepting argon, are present in traces only.
The Water Vapor. — The proportion of water vapor in the
air is exceedingly variable. When air becomes cool, the moisture
separates in cloud and fog, which are composed of minute drops of
Kquid water. When much moisture is condensed, the drops are
larger and fall as rain. When they fall through a cold region,
they freeze to hail. When condensation takes place in air already
NITROGEN. THE ATMOSPHERE 289
below 0®, the fog is composed of solid, and not of liquid particles.
The hexagonal crystalline structures of ice which are deposited
form snow.
On the other hand, when the weather becomes warm, evapo-
ration goes on rapidly, especially in the neighborhood of seas,
lakes, or moist country, and the proportion of water vapor in
the air may be considerably increased.
Humidity. — The moisture is usually defined in terms of
relative htunidity, the standard being the quantity required to
saturate the air at the existing temperature. A space filled with
air can take up aqueous vapor only until the partial pressure of
water vapor becomes equal to the vapor pressure of water (p. 61)
at the same temperature. The humidity is then said to be 100
per cent. If the partial pressure actually reached is only half as
great as the vapor pressure of water at the same temperature, the
humidity is 50 per cent. The average himiidity may be placed
very roughly at about 66 per cent.
At 18° (64.4° F.), the vapor pressure of water is 15.4 nmi.
(Appendix IV). If the total pressure of the atmosphere were
760 nmi., then the air would be saturated with moisture at 18°,
15 4
and have a humidity of 100 per cent, when -^^^ or about 2 per
cent of it by volume was water vapor. Upon cooHng to 0°, at
which temperatiu'e the vapor pressure of water is 4.6 nam., this
46
air would retain only =^, or about 0.6 per cent of moisture. At
18° there would be 16.3 grams of water in a cubic meter of air
and at 0° only 4.9 grams. The difference, 11.4 grams (11.4
c.c), would be precipitated as fog or rain from each cubic meter.
Test for Moisture in Air. — The presence of moisting in air
may be shown by placing any deliquescent (p. 118) salt, such as
calciimi chloride, in an open vessel. The quantity can be meas-
ured by driving a known volume of air slowly through a weighed
290 smith's intermediate chemistry
tube containing dry calcium chloride. It may be ascertained also
by noting the temperature to which the air has to be cooled before
it becomes saturated and deposits fog or dew. For example, if
air at 18° has to be cooled to 11° before it deposits dew, it contains
water vapor at a partial pressure of 9.8 nmi. If saturated at 18°,
it would have contained water vapor under a partial pressure
of 15.4 mm. Its relative humidity was therefore 9.8/15.4, or
63.6 per cent.
'Moisture and Comfort. — The chemical changes occurring
in oiu" bodies, and particularly the oxidation of waste and of
digested food by oxygen carried by the blood, are accompanied by
Uberation of heat. Yet oiu* bodies must remain at 98.4° F.
(37° C). A rise of a few tenths of a degree (F.) produces notice-
able discomfort. Much of the heat is lost by radiation from
the surface. The extent of this loss depends upon the surface,
which is invariable, and upon the surrounding temperature, which
we can not always control. Non-conducting clothes reduce the
radiation, and are increased in thickness in cold weather. The
real adjustment, however, is accompUshed, independently of
radiation, by evaporation of water at the surface of the skin.
The evaporation of 1 gram of water requires about 540 calories of
heat. Evaporation of a single half-ounce (14^ g.) of water will
therefore lower the temperatiu'e of 76 J kilograms (168 pounds)
of water (or of flesh, which is largely water) by one-tenth of
a degree C. (nearly 0.2° F.).
Our comfort, then, depends upon the possibiUty of continual,
moderate evaporation from the surface of oiu* bodies. " Much ''
moistiu'e in the air means, to us, not necessarily a great absolute
amount, but a near approach to the maximum possible at the
existing temperature. So the ratio of the amount present, to
the maximum — the humidity — is the significant fact for a
practical piu'pose, such as feeling comfortable (or drying the wash
quickly).
NITKOGEN. THE ATMOSPHERE 291
Ventilation. — In winter, cold air is brought into our rooms.
The amount of water vapor contained in this air, even if it is
saturated with moisture, is very small (see Appendix IV). When
this air has been heated, therefore, its relative humidity is too low,
discomfort is felt because there is too much evaporation, and mois-
ture has to be added artificially. Here the moisture afiforded by
evaporation from our bodies has little effect on the air. In smn-
mer, however, the outside air is often already nearly saturated at
the temperature of the room. At such times the speed of dis-
placement by the ventilating appUances may not be great enough
to keep the relative humidity down, and discomfort will arise
from the opposite cause. To reUeve it, the evaporation may
be promoted by electric fans. They do not remove or add any
air, but they stir it, and blow away the moist, nearly saturp.ted,
layers next to the skin.
The chief purposes of ventHation are, therefore, to supply fresh
air, to keep it in motion, and to maintain a humidity that is
neither too low nor too high.
The Carbon Dioxide. — The breathing of animals, the com-
bustion of coal and wood, and the decay (oxidation) of vegetable
and animal matter produce carbon dioxide CO2. The same gas
issues from volcanoes, and often in great quantities from the soil
in regions which are no longer volcanic. The proportion in the air
is therefore greatest in cities and in some volcanic regions, and
least in the country and over the sea. It varies from 3.5 parts in
10,000 in the country, to 1 per cent in crowded rooms.
Its presence may be proved in any air, and very quickly in
the breath itself, by bubbUng the air through calcium hydroxide
solution (Ume-water). Calcimn carbonate is precipitated (p.
336).
Carbon Dioxide and Respiration. — We draw about half a
liter of air into our lungs at each breath, or half a cubic meter per
292
smith's intermediate chemistry
hour. In the lungs some of the oxygen is removed, and som«
carbon dioxide is added.
Fresh Air.
Per Cent
Expired Air,
Per Cent
OxvKen
21.00
0.04
15.9
3.7
Carbon dioxide
A candle flame goes out when the proportion of oxygen has fallen
to 16.5 per cent. But air will sustain life until the proportion has
fallen to about 10 per cent.
Nearly all experts are now convinced that the unhealthiness
of over-crowded, " stuffy '' rooms is not due to the increase in the
proportion of carbon dioxide, which is seldom great enough to do
any damage. Nor is it due to " poisons " given off by the lungs or
skin. In spite of many experiments the presence of such sub-
stances has never been proved — they are imaginary. The
harm is caused by the stillness of the air, which, as we have seen
(p. 291), prevents the removal of the water vapor near the skin,
and therefore hinders evaporation.
Dust in the Air. — A beam of simlight crossing a dark room
can be seen by the hght reflected from the particles of dust which
all air contains. These are chiefly soUd bodies, and are composed
of salts, limestone, clay, and other rock materials, of soot and other
particles of unbumt fuel, of bits of hay or straw, and of fragments
of insects and other debris of plants and animals. They also
include living particles, such as bacteria, and spores of plants
such as moulds. The latter, when they settle upon food, germi-
nate and give rise to putrefaction. Some of the bacteria also pro-
duce disease, when they enter the body at a place where the skin
has been damaged by a cut or bum.
It is instructive to note that natural soil contains about 100,000
micro-organisms, and good, unfiltered river water from 6000 to
NITROGEN. THE ATMOSPHERE
293
20,000, in each cubic centimeter. Ordinary, pure air contains
only 4 to 5 micro-organisms per liter. Most of these bacteria
come from the drying of soil and the dispersion of the resulting
dust.
If dust were not present, we should soon notice its absence.
There would be no clouds or rain. It appears that moistiu'e
will not condense to fog or rain in air which has been filtered,
by being drawn through a wide tube containing a long (20 inches
or more) plug of cotton, and has so been freed from dust. The
particles act as nuclei, round which the Uquid grows at the expense
of the vapor. In the absence of dust, the condensation
would occur directly upon the surfaces of plants, houses, and
animals. Thus, in a dust-
less atmosphere, an open
shed or shelter, or an imi-
brella, would afford no pro-
tection whatever against a
wetting.
The formation of fog from
ordinary air, and its non-
formation in filtered air are
easily shown in a darkened
room (Fig. 74). The flask
contains water to saturate the air. When the tube leading to
the water pump is opened for an instant, the saturated air in
the flask expands and is cooled. In such circmnstances, ordinary
air gives a fog, briUiantly illuminated by the beam of Ught, while
filtered air (dustless) gives none.
Fig. 74
Air a Mixture. — The air does not contain in combined con-
dition the various substances we have named. Each of the
substances in air shows precisely the same properties which it
exhibits when free, separate, and pure. This behavior is charac-
teristic of a mixture.
294 smith's intermediate chemistry
Thus, the observed density of the air is precisely that which we
find by calculation from the known proportions and several densi-
ties of the components. The solubility of each gas is observed to
be the same as if it were alone present.
Again, when liquefied air is allowed to evaporate in a suitable
apparatus, the nitrogen, being more volatile, can be separated
almost completely from the oxygen. When the oxygen is, in turn,
allowed to evaporate, the carbon dioxide and water remain as
soUds, frozen by the low temperature.
Finally, the exa^ proportions can not be represented by a chemical
formvla. This shows that the law that, in chemical compounds,
the proportions can be represented by multiples of the atomic
weights by whole numbers (p. 77), does not apply to air.
In spite of the fact that air is a mixture, the composition of the
air is remarkably uniform and constant. The imiformity is due
to constant mixing by the winds. The steadiness of the com-
position from year to year is due to the fact that, although decay
and combustion continuaUy remove oxygen and add carbon diox-
ide, vegetation as continually consumes the latter and restores
the former (p. 396). The mass of carbon dioxide in the whole
atmosphere of the planet, about 2450 thousand million tons, is so
great that the amoimts added and removed by the agencies just
mentioned are small by comparison.
Liqueftiction of Gases. — The principle now used in h'quefy-
ing gases depends on the fact that, although a perfect gas, when
expanding into a vacuimi, should suffer no fall in temperature,
since it does no work, ordinary gases do become cooled very
sUghtly. The work which they do in expanding in such circum-
stances is done in overcoming the cohesion between their mole-
cules (p. 91), so that a tearing apart of the substance, which
consumes heat, has to take place. Since this cohesion becomes
more conspicuous the lower the temperature, the cooUng effect of
expansion becomes greater and greater as the temperature falls.
NITROGEN. THE ATMOSPHEHB 295
The most successful apparatus for use on a small or lat^e scale
is that devised by Hampson. In this apparatus {Fig. 75), two
concentric copper pipes, about 130 meters in length, are coiled
closely in a cyLndrical form, with non-conducting covering to
prevent access of heat from the outside. Air at 130-150 atmos-
Fm. 75 Fio. 76
pheres pressure is forced through the inner pipe (upper opening
Fig. 75), When it reaches the extremity of this pipe, it suddenly
escapes into a closed vessel. This expansion lowers its tempera-
ture. A spiral partition between the coils produces the outer
tube of which we have spoken. The gas in the tube A (Fig. 76)
is under a pressure of 130-150 atmospheres. "Hie distance of the
nozzle D from the plug C is adjusted so that the pressure of the
gas in the chamber and spiral outer tube is reduced to one atmos-
phere. The air can now escape only by traveling back through
the outer pipe to the final, wider esdt near the top. In doing so,
296 smith's intermediate chemistry
it cools the highly compressed air in the inner pipe. The cooler
air, on reaching the closed vessel, expands and becomes colder
than ever, and in passing backwards lowers the temperature of the
air in the inner pipe still further. Finally, the air in this pipe
Uquefies and drops of Uquid air are expelled into the closed vessel.
This is allowed to run out through a valve, from
time to time, as it accumulates.
Liquid air can be kept in Dewar flasks (Fig. 77).
The space between the inner and outer flasks is
evacuated, so that there is no gas to carry heat
Fig. 77 from the atmosphere in to the Uquid air. The
inner surface of the outer flask is often silvered, so
that radiant heat, from surrounding bodies, may be reflected and
not absorbed. Similar containers are in common use for keeping
hquids hot or cold for a long time (Thermos flasks).
Liquid Air. — Liquid air varies in composition, as the nitrogen
(b.-p. —194°) is less condensible than the oxygen (b.-p. — 181.4°).
When Uquid air evaporates, therefore, the first portions of gas
that come off consist almost entirely of nitrogen. Pure nitrogen
obtained in this way is used in the manufacture of ammonia by
the Haber process (p. 300), and in the formation of calcium cyana-
mide (p. 392). By allowing evaporation to continue a Uquid
containing 75 to 95 per cent of oxygen is obtained. This is
pumped into cyUnders and sold as compressed oxygen. It con-
tains about 3 per cent of argon, and is a convenient source of this
element. Cartridges made of granular charcoal and cotton waste,
when saturated with oxygen-rich Uquid air, are used as an eocplo^
sive in mining.
The Inert Gases
Argon. — Lord Rayleigh was the first to observe that a Uter of
pure nitrogen weighed 1.2505 g., while a Uter of atmospheric
"nitrogen" weighed 1.2572 g. The natural inference was that
SIR WILUAM RAUSAT
NITKOGEN. THE ATMOSPHERE 297
the latter contained a Kttle of some heavier gas. In 1894 Ramsay,
in consultation with Rayleigh, succeeded in separating this gas by
passing the " nitrogen " repeatedly over heated magnesium, and
so removing the real nitrogen as soUd magnesium nitride Mg3N2.
The remaining gas, about 1 per cent of the whole, was named
argon (Greek, lazy or inactive), because it would combine with no
other element.
Argon has a molecular weight of 39.9 (nitrogen only 28), and
when Uquefied hoils at — 186° and freezes at — 189.5°. It is used
in fiUing electric Kght bulbs.
Helium. — An indifferent gas, previously known to be given
off when uranium ores were heated in a vacuum, was found by
Ramsay (1895) to be neither nitrogen, nor yet argon. By its
spectrum it was recognized to be helium (Greek, the sun), a sub-
stance shown in 1868 to be present in the sim. Its molecular
weight is 4, so that it is only twice as dense as hydrogen. It was
the last gas to be hquefied (by Onnes), and the Uquid boils at
—268.7° (4.3° Abs.). Like argon, it does not enter into chemical
combination. HeUum is now being used to fill balloons, because
it is not combustible.
Other Inert Gases. — When liquefied argon was allowed to
evaporate, the first vapor coming off was found to contain another
gas, neon (Greek, new; Mol. wt. 20), along with heUiun. Care-
ful distillation of the remaining liquid gave two other gases
krypton (Greek, hidden; Mol: wt. 83) and xenon (Greek, stranger;
Mol. wt. 130). The total amount of these four gases, however,
was only 1 part in 80, the remaining 79 parts being pure argon.
None of these gases form any compounds. They do not combine
with themselves even, as do the more common gases such as H2,
O2, CI2. The molecule in each case is monatomic, for example
He, A. The valence throughout the group is therefore zero.
298 smith's intermediate chemistry
Exercises. — 1. Classify (p. 132) each of the reactions repre-
sented by equations in this chapter.
2. What are the radicals of sodium nitrite, and what their
valences? Justify the nomenclature.
3. At 77° F. the air of a room contains water vapor at a partial
pressure of 20 mm. What is the percentage of humidity?
4. What weight of water is contained in a cubic meter (1000
liters) of air saturated at 10® C?
5. What weight of carbon is contained in the total carbon
dioxide in the earth's atmosphere?
6. Air at 18® has to be cooled to 14® before it deposits dew or
fog. What is the percentage humidity at 18®?
7. Why is the air nearest the ground heated (by the sun) to a
higher temperatiu'e than the upper air?
CHAPTER XXV
AMMONIA
The interest in ammonia centers largely in the use of liquefied
ammonia for refrigeration, in the employment of the gas in making
carbonate of soda, and in the value of its compoimds as fertiUzers
and explosives.
Manufacture. — Ammonia is formed when m'trogenous or-
ganic matter is heated, in absence of air. It was formerly made
by distiUing scraps of hoofs, horns, and hides. The solution of the
gas thus obtained was called " spirit of hartshorn." The pungent
odor of smoldering feathers, leather, or fur is, therefore, partly due
to its presence in the escaping vapors. From the proteins of the
original plants, coal derives a considerable proportion of nitroge-
nous matter. Hence, when coal is distilled for the making of
coal gas, or, on a far larger scale, for the making of coke, much
ammonia can be separated, by washing with water, from the mix-
ture of gases produced. The aqueous solution is sejfeirated from
the tar, neutralized with sulphuric acid, and evaporated to give
the salt, ammonium sulphate (NH4)2S04.
NHs + H2O -> NH4OH (ammonium hydroxide).
2NH4OH + H2S04-^ (NH4)2S04 + 2H2O.
The distillation of coal is the chief source of commercial am-
monia. In the United States, prior to the war, nearly all the
coke was made in " beehive " ovens, in which the vapors issuing
from the coal are burned, uselessly, on the spot. Since the war,
about 75 per cent of coke is made in " by-product " coke ovens,
in which the ammonia and innumerable other by-products are
collected and utiUzed (see p. 424). In Scotland, oil-bearing shale
299
/
300 smith's intermediate chemistry
is distilled for the purpose of extracting the petroleum, and much
ammonia, hberated at the same time, is collected. Formerly
it was allowed to escape, but, in the absence of a protective tariff,
the competition of petroleum from American and Russian wells
compelled economy. Now, the profit on the sale of the
ammonimn sulphate pays the whole cost of mining and distilling
the shale.
Synthetic Ammonia. — The latest method of manufacturing
ammonia is by the direct union of hydrogen and nitrogen.
Exactly the same difficulties are encountered in the commercial
operation of this reaction (Haber's process) as in the manufacture
of sulphur trioxide by the contact process (p. 261), but in a greatly
accentuated form. The union of the gases, which is exothermic,
is exceedingly slow in the absence of a suitable catalyst:
N2 + 3H2 -^ 2NH8 + 24,000 calories.
In the presence of a contact agent — such as a specially prepared
mixture of iron and molybdenmn — combination is greatly
hastened. Traces of other gases, however, such as carbon monox-
ide or hydrogen sulphide, must be very carefully eUminated from
the reacting mixture, since they act as poisons on the catalyst,
that is, they destroy or impair its activity.
The reaction is reversible, and much more incomplete than is
the union of sulphur dioxide and oxygen under similar conditions.
Since the forward action evolves heat, the reverse action is favored
by raising the temperature (van't Hoff's law, p. 242), hence the
yield of anmionia in the equihbrimn mixture becomes less and
less the higher the temperature employed. Thus, under one
atmosphere pressure, the proportions of the gases that combine
in a mixtiu'e of one volume nitrogen and three voliunes hydrogen
are as follows: at 200°, 15.3 per cent; at 300°, 2.2 per cent; at
600°, 0.13 per cent; at 1000°, 0.004 per cent.
The preponderance of this reverse reaction, or in other words
the tendency of ammonia to decompose into its constituent
AMMONIA 301
elements as the temperature is raised, makes it impossible to
obtain high yields of anunonia by the Haber process at high
temperatures, while at lower temperatures the combination is too
tardy, even in the presence of a catalyst. Fortimately we are
able to make use of another of our general laws, the principle
of Le ChateUer (p. 244), to improve matters. It will be noted
from the equation given above that the union of hydrogen and
nitrogen to form ammonia is accompanied with diminution of
volume, 1 volume of nitrogen + 3 volumes of hydrogen = 2 vol-
imies of ammonia. Consequently the forward action will be
favored by increase of pressure. In fact, imder 200 atmospheres
pressure the yield of ammonia in the equihbrium mixture is as
follows: at 200°, 86 per cent; at 500°, 17.6 per cent; at 600°,
8.2 per cent; at 1000°, 0.9 per cent.
There still remains the question of speed of combination. This
decreases very rapidly as the temperature is lowered (compare
p. 241), and no catalyst has yet been prepared which is sufficiently
active to make the combination of nitrogen and hydrogen speedy
enough to permit the process to be operated on an mdustrial
scale much below 600°. A yield of about 8 per cent anmionia,
therefore, is the best that can be obtained.
Details of the Haber Process. — The hydrogen may be
obtained either as a by-product in an electrolytic process (p. 166),
or by the action of steam on iron (p. 51), or by careful purifica-
tion of water-gas (see p. 338). The preparation of pure hydrogen,
it may be noted, is the most costly feature of the whole process.
The nitrogen is obtained from liquid air. After removal of all
impurities injurious to the catalyst, the mixed gases are passed
under high pressure into the vessel containing the catalyst. This
consists of a steel " bomb " specially adapted to withstand liie
enormous pressure. Very serious disasters have taken place
owing to the explosion of such bombs. After passing over the
catalyst, the reaction mixture is cooled and its ammonia content
302 smith's intermediate chemistry
(6 to 8 per cent) removed either by refrigeration or by absorption
in water. The residual nitrogen and hydrogen are returned to
the plant for further treatment.
The ammonia obtained is used, in times of peace, mainly in
the manufacture of fertilizers, such as ammoniimi sulphate. In
war times, however, it is required more urgently for the produc-
tion of explosives. Nitric acid HNO3, which is necessary in the
preparation of most explosives, is obtained from ammonia by
oxidation (see p. 313). By the neutraUzation of nitric acid with
ammonia, ammonium nitrate NH4NO3 is formed. A mixture of this
substance with trinitrotoluene (p. 483) was used extensively as a
high explosive during the Great War, imder the name of amatoL
Sjmthetic ammonia may also be prepared by the calciimi
cyanamide process (see p. 392). But for these two processes,
Germany would never have been able to continue fighting in
the Great War beyond the first year. With all foreign suppUes
of nitrates (ChiU saltpeter) cut off, the only other available source
of ammonia was the by-product coke industry, and this was
already being utiUzed almost to its maximum. In the aUied
countries, unfortunately, the Haber process during the war was
not developed beyond the experimental stage.
The productive capacity of Haber process plants in 1920 was
no less than one and a half million tons (calculated as ammoniimi
sulphate).
Preparation in the Laboratory. — In the laboratory am-
monia is most readily made by heating a mixtiu'e of a salt of
ammonium, such as the chloride (NH4CI) or sulphate, and slaked
lime Ca(0H)2.
Ca(0H)2 + 2NH4CI -> CaCl2 + 2NH4OH -> 2NH3 + 2H2O.
The anmionium hydroxide, formed by the double decomposition,
immediately decomposes. To free the gas from water vapor, it is
passed through a tower filled with quickUme CaO (Fig. 78) :
CaO + H20-^Ca(OH)2.
AMMONIA
303
Sometimes^ a stream of the gas is generated by warming com-
mercial ammonium hydroxide solution (aqua ammonia), and dry-
ing the gas as above. Liquefied ammonia is obtainable in small
iron cylinders, and is a convenient
source when much of the gas is
required.
The liberation by hydrolysis of
nitrides, already noted (p. 287), is
interesting:
MgsNa+eHaO-^ 3Mg(OH)2+2NH3.
Composition of Ammonia. —
That ammonia contains nitrogen may
be shown by passing the gas over
cupric oxide heated in a tube, and
collecting the nitrogen over water:
2NH3 + 3CuO -^ 3Cu + 3H2O + N2.
Fig. 78
The hydrogen may be liberated by drying the ammonia, if
necessary, with soda-lime, and leading it through a tube containing
heated magnesium ribbon:
2NH8 + 3Mg -^ MgaNa + SHj.
Physical Properties. — Ammonia is a colorless
gas. It has a soapy taste, and a very pungent odor.
Its density, recorded in the formula JfHs, indicates
that it is only about half as heavy, bulk for bulk,
as air (14 + 3 = 17 against 28.95). It is easily
liqtiefied, boiling at —38.5°, and exerting a pressure
of 6 atmospheres at 10°. The gas is exceedingly
soluble in water (1 vol. water dissolves 1300 vol. of
NH3 at 0°). A 35 per cent solution is sold as " con-
centrated ammonia."
The extreme solubility in water may be shown by
fountain " experiment (Fig. 79). The flask is filled with am-
304 smith's intermediate chemistry
monia by downward displacement of air. The long tube is closed
by a short rubber tube and a clip at the bottom (not shown).
The " dropper " contains water, and is closed at the tip with
soft wax. A few drops of water, squirted into the flask by
pinching the " dropper," dissolve at once so much of the gas
that the water rushes in, Uke a fountain, through the longer
tube, when the cUp is opened.
Chemical Properties. — Anmionia, as we have seen, is
not very stable, and decomposes rapidly and almost completely
above 700*^. A discharge of sparks from an induction coil has
the same effect more gradually, and so a sample of the gas con-
fined over mercury in a closed tube may be shown to double in
volume when decomposed. Every two molecules give four:
2NH3 -^ 3H2 + N2.
The most characteristic property of ammonia is that it comr
bines directly with acidSy giving ammonium salts:
NHs (gas) + HCl (gas) -^ NH4CI (soUd particles).
It combines also with water at low temperatures to give ammo-
nium hydroxide NH4OH and ammonium oxide (NH4)20, white
soUds melting around —80°. These compounds are unstable
at ordinary temperatures, so that a solution of the gas, in a great
excess of water, is the only form of ammonium hydroxide con-
venient for use:
NH8 + H20^NH40H.
Ammonium Hydroxide. — This substance, as indicated by
the way in which we have written its formula, is a base. The
ions are (0H)~, given by all bases, and (NH4)"'", ammonium-ion,
which is found also in the salts mentioned above. The latter is a
compound positive radical, plajdng the part of a univalent metal-
lic element, such as Na or K.
As a base, ammonium hydroxide, although rather weak (little
AMMONIA 305
ionized), turns red litmus blue, possesses the characteristic soapy
taste and feeling, and enters into double decomposition with acids,
neutraliziAg them:
NH4OH + HCl -> H2O + NH4CI.
The salts, obtained by evaporation, are, of course, identical with
those formed by union of ammonia with the same acids.
Anmionium hydroxide used to be known as " volatile alkaU," in
reference to the fact that it decomposes into its constituents
(NH3 + H2O), both of which are volatile, while the other alkaUes
(NaOH, etc.) are not volatile (" fixed ")• This property was
utiUzed in the laboratory method of making ammonia (p. 302).
The Salts of Ammonium. — The salts are ionized in
aqueous solution, giving NH4 as the positive ion:
(NH4)2S04 -^ 2NH4+ + SO4".
When hecUedf dry, in a tube, they are decomposed. Most of them
give ammonia and an acid:
NH4Cl^NH3T+HClt.
If the acid is also volatile at, or below, a red heat, Uke sulphuric
acid H2SO4, the whole salt usually vaporizes. These actions are
reversible (read the equation backwards). Hence the acid and
the ammonia recombine, and the salt condenses again in a cold
part of the tube. This behavior helps us to recognize a salt
of ammonium, for the salts of mercury are the only others which
behave in this way.
Uses of Ammonia. — Some of the uses have already been men-
tioned. The anunonia process for making carbonate of soda
is described under the latter substance (p. 366).
Refrigeralion by Uquid anamonia depends upon the fact that
liquid ammonia, Uke any other Uquid, absorbs heat in evapo-
rating. It absorbs 260 cal. per gram. To freeze one gram of
water at O*', 80 calories have to be subtracted. Thus, 1 gram
of Uquid ammonia in evaporating will convert about 3 grams of
306
smith's intermediate chemistry
water to ice. The same principle is also largely used for cooling
air (in storage rooms for meat, etc.).
The machinery is represented diagrammatically in Fig. 80.
The ammonia, first admitted from a bomb of liquefied ammonia,
is driven by the pump F along the tube E and condenses to liquid
form in the tube coiled in the tank AB.
Cold water circulates through AB, and
removes the heat produced by the com-
pression and Uquefaction of the gas. The
liquid ammonia is allowed to drip through
the stopcock G into the lower coil. This
is kept exhausted by the compresser F,
and the Uquid ammonia evaporates. In
doing this, it takes heat from a 30 per
cent solution of calcium chloride in water,
which does not freeze even at —12°.
This cooled brine leaves the tank at D,
circulates through another tank (not
shown) in which the water-filled ice
moulds are suspended, and returns to C.
When used for cooUng refrigerating
chambers, the brine passes through a system of pipes suitably
placed in the cold-room. The whole machinery is made of iron,
as copper and brass are corroded by the ammonia.
Ammonium hydroxide solution is sold under the name of
household ammonia, and is used, in washing and cleaning, to
soften the watei;.
Exercises. — 1. Classify the reactions shown by equations in
this chapter.
2. How could you recognize the nitrogen and the hydrogen
obtained in the decomposition of ammonia (p. 304)?
3. Why can we not dry ammonia gas with concentrated sul-
phuric acid or with phosphorus pentoxide?
4. How could you separate a mixture of o^gen and ammonia?
Fig. 80
AMMONIA 307
5. In whaft relative volumes do ammonia and hydrogen chloride
imite?
6. Write in ionic form the equations for the interaction of
ammonium hydroxide: (a) with sulphuric acid; (6) with hydro-
chloric acid (p. 305).
CHAPTER XXVI
NITRIC ACID
Nitric acid HNOa is used in large quantities for making explo-
sives like guncotton, picric acid and TNT, and plastics like cellu-
loid, as well as innumerable drugs and dyes. Nitrates are largely
used as fertilizers (p. 410).
Manufacture. — Nitric acid is obtained in two ways, namely,
by the action of sulphuric acid upon natural sodium nitrate and
by oxidation of the nitrogen of the atmosphere. The processes
which utiUze the latter method will be referised to in a later section.
Sodium nitrate, Chile saltpeter, is found in an immense deposit
(2 by 220 miles) on the boundary of Chile and Peru. This salt is
mixed with concentrated sulphuric acid in iron retorts and gently
heated to drive off the nitric acid. The sodimn-hydrogen sulphate
remains in the retort:
NaNOa + H2SO4 ^ NaHS04 + HNO3 T •
The vapor is condensed in gla^ tubes (cooled with water) and the
acid collected in vessels of earthen-
ware. Sulphuric acid (b.-p. 330°) is
used because it is much less volatile
than nitric acid, and so only the latter
J, gj is vaporized. The acid boils at 86°
(760 mm.), but, to present loss by
decomposition, a lower boiling-point is secured by reducing the
pressure in the whole apparatus. -
In the laboratory the same action is employed, without, how-
ever, the reduction in the pressure (Fig. 81).
Physical Properties. — Pure nitric acid is a colorless liquid,
boiling at 86°. It is misdble in all proportions with water. The
308
NITRIC ACID 309
vapor, like hydrogen chloride gas, condenses moisture from the
air, giving a fog of droplets of the solution. The concentrated
nitric acid of commerce contains 68 per cent of the acid and boils
at 120.5°.
Chemical Properties — Decomposition of Nitric Acid.
— The acid is not very stable. It decomposes, in part, even when
simply distilled (86°), giving a red gas, nitrogen peroxide NO2,
water and oxygen:
Skeleton: HNO3 -^ NO2 + O2 + H2O.
Balanced: 4HNO3 -^ 4NO2 + O2 + 2H2O.
When a reducing agent is present, oxides containing less oxygen
than NO2 may be formed, such as nitric oxide NO.
Nitric Acid as an Acid. — An aqueous solution of nitric acid
turns blue litmus red, and is a very active^ highly ionized add.
With bases it gives nitrates, which can be obtained from the solu-
tion by evaporation:.
NaOH+ HNOa-^ H2O + NaNOs.
Ca(0H)2 + 2HNO3 -^ 2H2O + Ca(N03)2.
•» .
Nitric Acid as an Oxidizing Agent. — Since nitric acid gives
up oxygen with Uberation of energy, it is
also active as an oxidizing agent. Glowing
charcoal, as a powder or in the form of a
stick, will bum when pure nitric acid is ^^
poured upon it (Fig. 82). Carbon diaxide jwad red nitrogen per-
oxide NO2 are evolved.
Skeleton: HN03 + C-^ NO2 + CO2 + H2O.
Balanced: 4HNO3 + C -^ 4NO2 T + CO2 1 + 2H2O.
Nitric acid oxidizes indigo and other colored organic compounds,
in the same way as do the three oxidizing agents described in
Chapter XIX. It also oxidizes hydrochloric acid, upon which
310 smith's intermediate chemistry
hydrogen peroxide does not act; so that it is a more active oxi-
dizing agent than is that substance:
HNOs + 3HC1 -^ NOCl + 2H2O + CI2.
The mixture of concentrated hydrochloric acid, nitric acid, and
water is called aqua regia, and has strong oxidizing properties, due
to the presence of hypochlorous acid (from CI2 + H2O) as well as
nitric acid.
Action of Nitric Acid on Metals. — Magnesium, and metals
above it in the activity Ust, will displace hydrogen freely, especially
from diluted nitric acid:
Mg + 2HNO3 -^ Mg(N08)2 + H2 T .
But, with the less active metals, oxidaMon takes place, and instead
of hydrogen, we get water and, of course, a reduction product of
the nitric acid. The nitrate of the metal, however, is formed also.
Even metals, Uke copper and silver, which do not displace hydro-
gen, are acted upon by nitric acid in the same way (compare
sulphuric acid, p. 271). For example, diluted nitric acid acts
vigorously upon copper, giving nitric oxide NO as the reduction
product:
Skeleton: Cu + HNOj-^ Cu(N03)2 + NO + H2O.
Balanced: 3Cu + SHNOa -^ 3Cu(N03)2 + 4H2O + 2N0.
The nitric oxide is a colorless gas, but unites with oxygen in the
air to give NO2.
A test for a nitrate may be founded on this action. To the
nitrate sulphuric acid is added to Uberate nitric acid. Then
copper turnings are thrown in to give NO. A gas timiing red as
it meets the air shows that a nitrate was present.
Action Upon Organic Compounds. — Nitric acid stains the
skin and naUs yellcyw, by giving colored compounds with the pro-
teins. It gives similar compounds with wool (a protein), and
NITRIC ACID 311
therefore produces, on clothing, yellow stains which cannot be
removed.
The explosives made by the use of nitric acid are discussed in
Chapter XL.
Nitric Oxide NO. — This oxide is made by the action of diluted
nitric acid upon copper, by the action already discussed (p. 310).
It is a colorless gas, and almost insoluble in water.
Vigorously burning phosphorus continues to bum in it:
4P + lONO -^ 2P2O6 + 5N2.
Its most important property is that of uniting with oxygen to
give nitrogen peroxide:
2N0 + 02--^ 2NO2.
This reaction is reversible at high temperatures.
Nitrogen Peroxide NO2. — This oxide is formed by the union
of oxygen with nitric oxide. It is given off, also, when concerv-
trated nitric acid acts upon metals and other reducing substances.
It is further produced, along with oxygen, when nitrates, excepting
those of potassium, sodium and ammonium are heated, dry:
2Cu(N08)2 -^ 2CuO + 4NO2 + O2.
Potassium and sodium nitrates, when heated, give off only
oxygen, and leave the nitrites:
2KN08->02 + 2KN02.
Nitrogen peroxide is a red gas, with a choking odor. When
heated strongly, it becomes colorless, being dissociated into NO
and oxygen. When cooled, it becomes pale yellow, and its den-
sity becomes twice as great. This is due to the formation of mole-
cules of the formula N2O4:
2N02?:±N204.
312 smith's intermediate chemistry
The most interesting property of nitrogen peroxide is its action
upon water, whereby nitric acid is formed, and nitric oxide escapes:
3NO2 + H4O -» 2HN0, + NO.
When oxygen is present also, then the NO gives more NO2, and
this in turn gives more nitric acid. This action plays an important
part in the making of nitric acid from the nitrogen of the air (see
next section).
The gas is sometimes used for bleaching flour, but traces of the
oxide remain in the bread.
Fixation of Atmospheric Nitrogen. — Oxygen and nitrogen
have no natural tendency to combine at the ordinary temperature,
but rather the reverse — the compounds tend to decompose with
evolution of heat. But a high tem-
perature will supply the necessary
enei^y. Even so, however, the
union extends to only 1 per cent of
the mixture at 2000° and 5 per cent
at 3000°:
Ns + Oj F± 2N0.
Note that, since the formation of
NO is an endothermic reaction, the
yield of NO is ino'eased by raising
the temperature (van't Hoff's law,
p. 242). In spite of the poor yield
obtainable under the best condi-
tions, the supply of natural nitrates is so limited that machinery
has been devised, and is now in successful use, for carrying on
the combination on a commercial scale. Three devices are in use,
and all employ hydro-electric power.
In the Birkeland-Eyde process (Fig. 83), used at Notodden and
elsewhere in Norway, an arc discharge between rods of carbon is
spread, by the influence of powerful electromagnets, into a circular
NITBIG Acm 313
brush discharge several feet in diameter. The figure shows a
cross section of the space filled by the discharge. In the center is
a section of one of the carbon rods. Air is blown through the
flame, giving 1 per cent of NO, and is cooled to permit of union of
the nitric oxide with oxygen, to give the per- ^,^^ ^
oxide, NOi, The air containing NOg is then °*'°* !
passed through absorbing towers down which
water trickles. Here the action mentioned in
the last section takes place, and an aqueous
solution of nitric acid is produced. In peace
times, the nitric acid is mixed with calcium
hydroxide (slaked hme) :
Ca(OH)s + 2HN0» -* Ca(NO»)a + 2HiO
to give calcium nitrate, which, being very solu-
ble, is sold for use as a fertihzer. In war times,
the acid is concentrated for use in the manu-
facture of explosives.
The Schonherr process, used in the same
factories in Norway, employs a discharge
through a tube 22. feet long (Fig. 84). The
column of air rotates as it traverses the tube
and so every part is exposed to the discharge.
The Pauling process, used in Italy and
Austria, uses preheated air, and a different
arrangement of the discharge. The principles
employed are, however, the same.
The productive capacity of plants employ-
ing these arc processes in 1920 was three "^^
hundred thousand tons (calculated as nitrate of lime).
Nitric Acid from Ammonia. — The ammonia oxidation proc-
ess for the production of nitric acid was developed on a large
scale during the Great War, particularly in Germany, Gaseous
ammonia in the presence of air and a suitable catalyst undergoes
314 smith's intermediate chemistry
oxidation^ with the fonnation of oxides of nitrogen and water
vapor. The oxides of nitrogen can be recovered by absorption
in water, yielding dilute nitric acid. This can be concentrated
further if desired, or neutraUzed with a base for the production
of nitrates for use as explosives (ammonium nitrate) or f ertiUzers
(calcium nitrate).
Platinum gauze is almost universally employed as a catalyst.
A mixture of anmionia and oxygen-enriched air passes through
one or more layers of the gauze, which is heated electrically to
650-700° to start the reaction and which is maintained at that
temperature subsequently by the heat of combustion of the
ammonia. The reaction that takes place may be represented by
the equation:
4NH8 + 5O2 -> 4N0 + 6H2O.
The excess oxygen present converts the NO into NO2 as the issu-
ing gases cool. Unless the reaction is very nicely regulated,
however, the yield pf oxides of nitrogen is diminished either by
incomplete combustion of ammonia, or by the dissociation of
NO into nitrogen and oxygen (see p. 312). The time of contact
with the catalyst must not be too long, or this latter effect will be
appreciable. With proper precautions, a conversion efficiency of
90-95 per cent is obtained. Poisoning of the catalyst must be
guarded against by careful purification of the gases, especially
from non-volatile impurities such as dust particles, which choke
the surface of the gauze and render it inoperative.
Nitrous Acid HNO2 and Nitrous Anhydride VtOs, —
When an acid, such as sulphuric acid, is added to a solution of
a nitrite, Uke potassium nitrite (p. 311), nitrous acid HNO2 is
formed:
2KNO2 + H2SO4 -> K2SO4 + 2HNO2.
Nitrous acid, however, like sulphurous acid, is imstable and
nitrous anhydride N2O8 is at once liberated and escapes as a gas:
2HN02-^H20 + N20,t.
NITKIC ACID 315
This gas is used as a catalytic agent in the chamber process (p.
264) for making sulphuric acid. The acid and its anhydride are
employed in the manufacture of many dyes.
Nitrous Oxide N2O. — When ammoniuna nitrate, a white salt,
is heated, it decomposes into steam\and nitrous oxide: *
.NH4NO8 "^ 2H2O t + N2O T .
Nitrous oxide is somewhat more easily Uquefied (b.-p. —90°) than
is carbon dioxide. At 12° its vapor pressure is 41 atmospheres.
It is sold, as a Uquid, in steel cylinders, and used as an anaesthetic
for minor operations, chiefly in dentistry. The hysterical symp-
toms which accompany its use caused it to be named "laughing
gas."
Like oxygen, it reUghts a glowing spUnter of wood, and supports
combustion brilliantly. It does not interact with nitric oxide (p.
311), as does oxygen, however, to form nitrogen peroxide.
The Writing of Equations. — The reader will have discovered
that some of the equations in the present chapter are rather harder
than usual to balance correctly, even when the products of the
reaction are all known. The following sections should be read
through carefully, in order to obtain a thorough understanding
of the points under discussion. The hints given will be found of
general assistance in writing difficult equations.
Points About Oxygen Acids. — In dealing with an acid that
contains oxygen, Uke nitric acid, there are some things which we
must acquire the habit of keeping in mind.
Thus, an oxygen acid, as we have seen (p. 257), can be deprived
of water, leaving the anhydride. The chemist always thinks of
the one as soon as the other is mentioned. If the acid is named,
he instantly subtracts water from its formula to get the formula
of the anhydride:
HaCOa-^HjO + COi.
Skeletan: HOCl -^ H2O + ?
316 SMITHES INTERBiEDIATE CHEMISTRY
To balance the second equation, hydrogen atoms must be supplied
in pairs, so the HOCl must be multiplied by two:
Balanced: 2H0C1 -^ H2O + CUO.
With nitric acid the operation is similar:
Skeleton: HNO3 -> H2O + ?
Balanced: 2HNO3 -* H2O + N^Os.
This operation should be practiced:
Skeleton: HCIO4 -* H2O + ? H3PO4 -^ H2O + ?
Balanced: 2HCIO4 -> H2O + ? 2H3PO4 -* 3H2O + ?
The valence of the characteristic non-metal is ascertained by this
process. What is the valence of CI in HOCl? The anhydride is
CI2O. The valence is Cl^. What is the valence of S in H2SO4?
The anhydride is SOs. The valence is S^. What is the valence
of N in HNOs? The anhydride N2O5 shows the valence to be N^.
What are the valences of CI in HCIO4 and of P in H3PO4 and in
HPO3?
When we get SO2 from sulphuric acid, by a chemical action, how
do we know that the acid has been reduced (and something else
oxidized)? Because in sulphuric acid we have 8^3", and in the
product S'^02^. The valence of S has been lowered from VI to
IV. When from nitric acid we get N2O6, has there been reduction?
No, because N2O6 + H2O = 2HNO3 (nitric acid). If we get
NO2 or NO, has there been reduction of the acid? Yes, because
the valence has been reduced: N'^02, N"0. We then proceed
to pick out the other substance that has been oxidized.
Analyzing the formula of the acid, to get that of the anhydride,
also aids ils to balance equations.
Balancing Equations, — When dilute nitric acid acts upon
copper, the following skeleton equation represents the products
of the reaction.
Skeleton: Cu + HNOa -> Cu(N08)2 + NO + HjO.
NITRIC ACID 317
The valence of N in HNO3, from its anhydride N2O5, is N^.
In Cu(N03)2 the valence of N is unchanged; all that has happened
here is the displacement of 2Hi by Cu^. The valence of N in
NO, however, has been reduced to N^.
Let us see how we can utiUze these facts to balance our equation.
The formation of one molecule of Cu(N03)2 involves the displace-
ment of 2 atoms of hydrogen. The reduction of one molecule
of N2O5 to 2N0 involves the Uberation of 3 atoms of oxygen.
Both reactions are occuring simultaneously, so that neither hydro-
gen nor oxygen is actually obtained in the free state, they will of
course unite to form water. But 3 atoms of oxygen are equivalent
to 6 atoms of hydrogen, therefore 3 molecules of Cu(N03)2 must
be formed for every molecule of N2O6 that is reduced. To obtain
3Cu(N03)2, we must use 3Cu and GHNOs. To reduce IN2O6,
we must use 2 more HNO3, that is, 8HNO3 in all. The balanced
equation therefore becomes: I
Balanced: 3Cu + 8HNO3 -* 3Cu(N03)2 + 2N0 + 4H2O.
I'hese operations should be practiced on the following skeleton
equations:
HNO3 + C -* NO2 + CO2 + H2O (p. 309) .
H2SO4 + HI -^ H2S + H2O + I2 (p. 204).
Zn + H2SO4 -^ ZnS04 + H2O + H2S (p. 271).
These are typical examples of rections which it is very difficult
to write correctly without some such guide as is given above.
Exercises. — 1. What is the valence of carbon in carbonic
acid H2CO3?
2. What is the anhydride of nitrous acid HNO2, and what the
valence of nitrogen in this compound?
3. When nitric acid acts upon copper, which substance is oxi-
dized and which reduced?
4. How could you show experimentally that both nitrogen per-
oxide and oxygen are formed when cupric nitrate is heated?
318 smith's intermediate chemistry
5. How could you distinguish nitric oxide, (a) from hydn^n,
(b) from oxygen?
6. How could you distinguish nitrous oxide from oxygen?
7. Make a list of the names and formulae of the oxides of nitro-
gen, arranging them in the order of increasing proportions of
oxygen.
8. Write full ionic equations for the preparation of ammonium
nitrate (p. 309)?
CHAPTER XXVII
PHOSPHORUS, ARSENIC, ANTIMONY, BISMUTH
We now take up the elements which, with nitrogen, form the
nitrogen family. These elements all have two regular valences,
being trivalent and quinquivalent. Nitrogen, phosphorus and
arsenic are non-metallic elements, that is, they do not form posi-
tive radicals of salts. Antimony is non-metallic, although in
its trivalent condition it acts also as a metallic element. BismuHi
is metallic.
Phosphorus P
Occurrence. — Calcium phosphate Ca8(P04)2 forms about 25
to 27 per cent of the material of the bones and teeth of animals.
The same salt occm^ in deposits, as a mineral, and is foimd scat-
tered through all fertile soils. Complex organic compoimds of
phosphorus, such as lecithin, are essential constituents of the
muscles, nerves and brains of animals and are foimd also in plants.
The average man's skeleton contains 1400 g. of phosphorus, his
muscles 130 g., and his nerves and brain 12 g. Amongst foods,
egg-yolks and beans contain an imusually large proportion, nuts,
peas, and wheat (entire grain) coming next.
Phosphorus was discovered by Brand in 1669, and by Kunkel
in 1670, by distilling at a white heat the solid residue from evapo-
rated animal matter. They were both searching for the philoso-
phers' stone. Scheele in Sweden prepared it from bones in 1771.
The element is used chiefly in the manufactiu'e of matches and,
to a small extent, in roach paste and rat poison.
Manufacture of Phosphorus. — Phosphorus is now manu-
factured by mixing natural calcium phosphate with sand (SiQ^)
319
320 smith's intermediate chemistry
fuid coke, and heating the mixture in an electric furnace (Fig.
85). The mixture is admitted by moving the traps below the
hopper, and ia carried into the furnace by the worm conveyor.
The rmetance of the mass between the electrodes causes great
development of heat. The actions may be shown by partial
equations, which, when added together, give the complete equa-
tion:
Ca,(P04), -» 3CaO + P»a
3CaO + SSiOa -» SCaSiO. (calcium sOicate)
P.Ob + 5C-»2P + 5CO
. Ca,(PO0. + 3SiOt + 5C -> SCaSiO, ^ + 2P T + 5C0 T .
The calcium silicate is melted and runs out below as a slag. The
phosphorus (vapor) and carbon monoxide
(gas) pass off through the opening near the
top. The phosphorus vapor is condensed
under cold water.
White Phosphorus. — The product, after
purification, is a cohrhss, transparent waxy
solid (sp. gr. 1.83), which meUa at 44° and
' hoiU at 287°. It is insoluble in water, but
dissolves in carbon disulphide. It has a
Fia. 85 strong odor, resembhng ozone.
White phosphorus oxidizes in the air, giv-
ing, when moist, phosphorous acid and phosphlsric acid, and
emitting a faint light from which the element derives its name
(Greek, liglU bearer). It catches fire at a low temperature (about
35°), and in burning forms a cloud of the sohd phosphorus
pentoxide PiOs. It combines readily, even when cold, with the
halt^ens, and when heated it unites with sulphur and the more
active metals.
White phosphorus is a very active poison (fatal dose, 0.15 g.).
When traces of the vapor are breathed day after day, a disease,
PHOSPHORUS, ARSENIC, ANTMONY, BISMUTH 321
frequently shown by workers in match-factories and consisting in
ulceration of the bones of the jaw, makes its appearance. The
use of white phosphorus is forbidden by law in Sweden, France,
Great Britain, and Switzerland and is penalized by a special tax
in the United States.
Red Phosphorus. — When white phosphorus is heated at 230
to 300® in a tightly closed vessel (air excluded) it changes into red
phosphorus. This material is composed of small crystals, of dull
red color f and variable specific gravity 2.19 to 2.34. It is insoluble
in carbon disulphide, has no odor, and is not poisonous. On dis-
tillation the vapor condenses to white phosphorus.
This allotropic form of phosphorus is formed from the white
variety with Uberation of much heat. It thus contains less energy^
and is much less active. It bums to form the pentoxide, but
has to be heated to about 240® before it will catch fire in the air.
It combines also with elements other than oxygen much less
readily than does white •phosphorus.
Manufacture of Matches. — These are of two kinds, ordinary
matches, which strike on any rough surface, and " safety "
matches. Ordinary matches are still made in some coimtries by
dipping the splints of wood in melted paraflSn, and then in a paste
made of 4 to 7 per cent of white phosphorus, lead dioxide about
60 per cent, water, and dextrin (paste). The head, when dry, is
dipped in varnish to exclude air and moisture. Where the use of
white phosphorus is very properly forbidden, a sulphide P4S3 is sub-
stituted. When the match is struck, the friction explodes the
mixture of phosphorus trisulphide (combustible) and lead diox-
ide (or other oxidizing agent), and the resulting heat sets fire to
the paraffin and this, in turn, to the wood.
Safety matches carry no phosphorus, but only a mixture of
substances containing oxygen, such as potassimn chlorate or
potassium chromate, with a combustible, Uke antimony trisul-
322 smith's intermediate chemistry
phide, some dextrin and a filling (e.g., chalk). The box is coated
with a mixture of red phosphorus, antimony trisulphide, dextrin
and fiUing. The friction converts a trace of the red phosphorus
into the white variety, and the latter sets fire to the head.
Phosphorus Pentoxide P2O5. — When phosphorus is burned
in dry air or oxygen, imder a bell jar, the cloud of pentoxide slowly
settles as a white powder. If the pentoxide is thrown into cold
water, chemical union takes place. The heat developed pro-
duces a hissing sound, caused by the formation and condensation
of minute bubbles of steam. The solution, when concentrated at
a low temperature, gives crystals of orthophosphoric add H8PO4
(m.-p. 42^).
P206 + 3H20-^2H8P04.
On account of its tendency to unite with water, the pentoxide
is used for drying gases.
When phosphorus is burnt in moist air, the cloud of pentoxide
forms tiny droplets, consisting of a concentrated solution of phos»
phoric acid, which remain suspended in the atmosphere as a fog
(compare sulphur trioxide, p. 263) . Burning phosphorus was there-
fore used in the war for screening the movement of vessels. In
land warfare, shells containing white phosphorus were also em-
ployed for incendiary purposes. Such shells produced terrible, and
usually fatal, bums on any enemy within the radius of their explo-
sion.
As a triboMC acid, phosphoric acid H8PO4 gives three series of
salts by interaction with bases. In the oriJio or normal phosphates,
such as Na3P04, all of the hydrogen is replaced by a metaUic
radical. Two series of acid salts are also known (see p. 277).
The phosphates of calcium are of particular value as fertilizers
(see p. 411).
Test for a Phosphate. — Most phosphates (and phosphoric
acid), when mixed intimately with dry sodium carbonate, char-
PHOSPHOKUS, ARSENIC, ANTIMONY, BISMUTH 323
coaly and magnesium powder, and heated in a narrow tube closed
at one end, give a phosphide of the metal (NasP or Mg3P2). When
the mass is moistened, the odor of phosphine PHs can be recog-
nized:
NasP + 3H20^3NaOH + PH, T .
Shells containing metallic phosphides are of service in naval
warfare as markers in night engagements, the impure phosphine
produced by the action of the water on the phosphides burning
spontaneously on the surface and indicating the position of the
shot.
Phosphorus and Nitrogen Compared. — Although the
simple substances, phosphorus and nitrogen, have Uttle in com-
mon, they form compounds of similar composition, which are m
many ways aUke. Thus we have ammonia NHs and phosphine
PHs, both gases. The phosphides are analbgous to the nitrides
(p. 287). Then there are the oxides N2O8 and N2O6, to which the
oxides P2O3 and P2O6 correspond. All these oxides are anhydrides
of acids. Both nitrogen and phosphorus are typically non-
metaUic elements (p. 274), entering into negative radicals. Both
elements are trivalent in one series of compoimds (NH3, N2O8,
PHs, P2O8) and quinquivalent (N2O6, P2O6) in another series.
Aksenic As
Arsenic, the third member of this family, is metallic in appear^
ance and in physical properties, although in combination it behaves
as a non-metaUic element. The metal is used with lead in making
small shot, and the oxide As20s in preparing medicines and insec-
ticides.
Preparation. — Arsenical pyrites FeSAs, a mineral similar to
pyrite FeSj, but containing arsenic in place of half of the sulphur,
is one of the commonest natural forms of arsenic. When this
324 smith's intermediate chemistry
mineral is heated (air excluded), arsenic passes oflf as vapor and
condenses as a crystalline metallic powder:
FeSAs -> FeS + As t .
Most natural sulphides (PbS, FeS2, SnS2, CuFeS2, etc.) contain
more or less arsenic, which takes the place of a part of the sulphur.
When these ores are oxidized in a draft of air (roasted), as one
step towards the ultimate extraction of the metal, the metal,
sulphur, and arsenic are all converted into oxides. The sulphur
dioxide passes off as gas, but the arsenic trioxide AS2O3, a soUd,
settles in the flues. By distilUng the deposit with carbon, free
arsenic is obtained:
AS2O3 + 3C->2As t + SCO T .
Properties and Uses. — The element has a silvery luster, but
tarnishes quickly. When it is heated, its vapor reaches a pres-
sure of 760 nam. before the melting-point is attained, so that the
metal sublimes without melting.
The metal bums in air with a bluish-white flame, giving clouds
of white particles of the trioxide AS2O3 (poisonous).
In the making of small shoti about 0.5 per cent of arsenic is
added to the lead. The latter is then rim into a vessel, with a
perforated bottom, placed at the top of the shot tower. The
arsenic, Uke any dissolved substance (p. 119), lowers the freezing-
point of the solvent (lead), and delays the soUdification of the
lead until the drops have assumed perfect spherical form. At
the foot of the tower the drops fall into water and are cooled.
The arsenic also renders the metal harder than pure lead, and
less apt to be deformed during the explosion of the cartridge.
Arsenic Trioxide As203* — This oxide, formed when arsenic
bums, is a white crystalUne powder (white arsenic). It is dddic
and forms salts with bases. Paris green and Scheele's green are
made by dissolving the oxide in boiling water and adding a copper
PHOSPHOBUS, ARSENIC, ANTIMONY, BISMUTH 325
salt. They are green, insoluble compounds used, as insecticides,
for spra3ring plants. On account of their poisonous character,
they are no longer employed as ingredients in paints.
Arsenic as a Member of the Nitrogen Family. — While free
arsenic is as different physically from phosphorus as the latter is
from nitrogen, the compounds have much in common. Arsenic
forms arsine AsHa (corresponding to NH3 and PH3), arsenious
oxide AS2O3, arsenic oxide AS2O5, and the acids HsAsOs and
H3ASO4. In these compounds it is non-metallic, and shows the
valences three and five.
Antimony Sb
Preparation and Properties. — Antimony is found free in
nature. The sulphide, stibnite Sl^Sz, is also a well-known mineral.
When the latter is melted with iron, ferrous sulphide and free
antimony are formed:
3Fe + SbzSs -^ 2Sb + 3FeS.
The molten ferrous sulphide (sp. gr. 4.8) floats upon the molten
antimony (sp. gr. 6.5), and the products, being mutually insoluble,
are easily separated.
The metal is brilliantly silvery and non-tarnishing. It is brittle
and the black powder obtained by pulverizing it, " antimony
black," is rubbed on plaster casts to give them a dull, metallic
appearance. When heated, antimony volatilizes and hums in the
air with a brilliant white light to form the white trioxide Sb203.
The trisulphide Sb2S3 (orange-colored when precipitated) is used
in making matches and fireworks.
Alloys Containing Antimony. — The metal is chiefly used
as an ingredient in alloys. Lead, when solidifying, shrinks and
antimony coimteracts this tendency. Hence type metal, which,
when cold, must fill the mould completely, is made by melting
together 15 to 25 per cent of antimony, 10 to 20 per cent tin, and
326 smith's tNTERMEDIATE CHEMISTRY
the rest lead. The alloy is also harder than lead, and is less
quickly deformed by handling and by use in the printing press.
Babbitt*s metal (Sb 3, Zn 69, As 4, Pb 5, Sn 19), and other
anti-friction alloys, used in lining bearings, contain antimony
along with zinc, copper, and other metals. Molten mixtures of
metals (alloys), when solidifying, do not always form a homogene-
ous, solid mass. In an anti-friction alloy, what is wanted is a
mass, in general soft, but. containing hard particles. The latter
bear most of the pressure, yet, as the alloy wears, they are pressed
into the softer matrix so that a smooth surface is always presented.
An alloy which has the opposite composition, that is, which gives
a hard mass containing softer particles, develops heat by friction
much more rapidly.
Bismuth Bi
As the atomic weight increases, the members of this family
become more like metaUic elements in their chemical properties.
Thus bismuth is a true metaUic element. Its oxides are basic,
and its compounds give positive ions Bi+++ and include salts like
the carbonate, sulphate, chloride, and phosphate.
Metallic Bismuth. — The metal occurs free in nature. It is
a brittle metal, with a pink metaUic luster. It melts at 270** and
vaporizes at a high temperature. It does not tarnish. It is
used in preparing alloys with very low melting-points. Thus
Wood's metal contains bismuth (m.-p. 270**) 4 parts, lead (m.-p.
326**) 2 parts, tin (m.-p. 233**) 1 part, and cadmium (m.-p. 320**)
1 part. As is the case with other solutions, the melting-point is
lower than that of any of the components, namely 60**. Alloys of
this class are used as plugs in sprinkler systems and stops to hold
steel fire-doors open. When, in consequence of a fire, the tem-
perature rises, the alloy melts, the water exits are opened and the
fire-doors swing shut. Safety plugs in steam boilers, made of a
similar, but less fusible alloy, melt when, as the result of failure
PHOSPHOBU8, ARSENIC, ANTIMONY, BISMUTH 327
of the safety valve, the steam pressure, and therefore the tempera-
ture, exceeds a predetermined value. They behave in the same
way when the water is dangerously low and the metal above the
water becomes too hot.
Compounds of Bismuth. — When strongly heated, the metal
burns to form a brown trioxide Bi208. This oxide gives salts with
acids. Thus, with nitric acid it dissolves to form a solution of
bismuth nitrate :
BiaOa + 6HN08 -^ 2Bi(NOs)8 + SHaO.
When the solution is evaporated, the nitrate appears in color-
less crystals. If the crystals are placed in water, a white, insolu-
ble basic nitrate is formed:
Bi(N08)8 + H2O T± BiON08 i + 2HN08.
This is used in medicine under the name of bismuth subnitrate,
to cure stomach troubles.
Exercises. — 1. By what chemical experiments could you
recognize red phosphorus?
2. Explain the fact that solid Wood's metal floats upon melted
Wood's metal.
3. Why is white phosphorus always kept imder water?
CHAPTER XXVIII
CARBON Ain> THE OXIDES OF CARBON
The majority of the substances composing, or produced by,
living organisms, such as starch, fat, and sugar, are compounds
of carbon. Hence the chemistry of these compounds is known as
organic chemistry. It was at first supposed that the artificial
production of such compoimds, e.g,, without the intervention
of life, was impossible. But many natural organic products have
now been made from simpler ones or from the elements, and the
preparation of the others is delayed only in consequence of diflScul-
ties caused by their instabiUty and complexity. On the other
hand, himdreds of compoimds imknown to animal or vegetable
life, including many valuable drugs and dyes, have now been
added to the catalogue of chemical compounds. Hundreds of
thousands of different compounds containing carbon are known,
and thousands more are added every year.
The elements entering into carbon compounds are chiefly hydro-
gen and oxygen. After these, nitrogen, phosphorus, the halogens
and sulphur may be named.
Cakbon C
Occurrence. — Large quantities of carbon are found in the
free condition in nature. The diamond is the purest natural
carbon. Graphite, or plimabago, which is the next purest, is
found in limited amounts, and is a valuable mineral. Coal occurs
in numerous forms containing greatly varying proportions of
free carbon. Small quantities of the free element have been found
in meteorites.
In combination, carbon is found in marsh-gas^ or methane
CH4, which is the chief component of natural gas. The numerous
328
CARBON AND THE OXIDES OF CARBON 329
compounds found in plants and animals have already been men-
tioned. The mineral ails consist almost entirely of mixtures of
various compoimds of carbon and hydrogen (hydrocarbons).
Whole geological formations are composed of carbonates of common
metals, particularly calcium carbonate or limestone.
Diamond. — This allotropic modification of carbon is dis-
tinguished by its natural crystalline form, which often resembles
the octahedron (Fig. 39, p. 94). The ultimate struc-
ture of the diamond crystal is represented in Fig. 45
(p. 96). Its specific gravity is 3.5. For ornamental
purposes the diamond is " cut " by grinding new faces
so as to give artificial forms called " briUiants '' (Fig.
86) and " rosettes." It is the hardest of famiUar
substances, and can be scratched or poUshed only
by rubbing with diamond powder. The colorless
stones and those with special tints are valuable.
The black (" carbonado '') and badly colored specimens are less
valuable and are used for grinding, for glass-cutting, and on the
points of driUs.
Diamonds are foimd chiefly in South Africa and Brazil. They
are separated from the rock by weathering and washing. They
are sold by the carat (1 international carat = 200 mg.) and the
value increases with the size. The largest known specimen,
the CuUinan, weighed 3032 carats before being cut.
Small synthetic diamonds are obtained when molten iron con-
taining dissolved carbon is suddenly chilled. If the fused mass
is allowed to cool gradually, however, the carbon separates out
in the form of graphite.
Graphite is found in nature in Siberia, Cumberland, Brazil,
Ceylon and elsewhere. It forms dark grey or black hexagonal
tablets, and, when pulverized, it gives slippery scales of micro-
scopic size. Unlike the diamond, it is quite soft, has a specific
330 smith's intermediate chemistry
gravity of 2.3, and conducts electricity. Natural graphite is
usually mixed with foreign matter, and even the purest specimen
leaves, when burned, from 2 to 5 per cent of ash. It is called also
plumbago, or black lead.
Graphite (Greek, 7 write), moulded into blocks, is sawn into
rods for the cores of " lead " pencils (first used in the 16th cen-
tury). Clay is added in varying proportions to give different
degrees of hardness. Because of its infusibiUty, it is used to make
crucibles. Smeared on a plaster cast (non-conductor), it gives a
conducting surface on which metals (copper or silver) can be
deposited by electrolysis. A thin layer, used as stove-polish,
protects the iron from rusting. In electro-chemical industries it
is used for electrodes at which chlorine is to be Uberated; all
other conductors interact chemically with this element and
are destroyed. It is employed also as a lubricant, when wooden
beams sUde upon one another.
Large amounts of pure graphite are now manufactured by
heating coke with some pitch and a Uttle sand or ferric oxide
(Acheson's process). The mixture (3 to 3| tons) is piled (Fig.
87, p. 333) between the electrodes connected with a dynamo, and,
on account of its high resistance, becomes strongly heated. The
operation is complete in from 24 to 30 hours.
Other Forms of Carbon. — The apparently amorphous
varieties of carbon are numerous. They include wood-charcoaly
lampblack, animal charcoal, coal (e.g., bituminous coal and anthra-
cite) and coke. All of these substances will come up for dis-
cussion in later chapters. ' None of them, it may be noted here, is
composed of pure carbon, other elements being present, mostly
in combination with carbon, in very variable amoimts.
Examination of " amorphoiLS " charcoal by X-ray methods
indicates that it possesses a crystalline structure identical with
that of graphite. Charcoal is not to be regarded, therefore, as a
supercooled Uquid (see p. 94) Uke glass. It consists of tiny frag-
CARBON AND THE OXIDES OF CARBON 331
ments of graphite. The amorphous appearance is due to the
extreme minuteness of the crystals, which are interspersed, with
attendant impurities, through a highly porous mass.
Chemical Properties of Carbon. — 1. Carbon unites vigor-
ously vrith oxygen. With an excess of oxygen it forms carbon
dioxide, a gas:
C + Oa-^COa.
With a Umited supply of oxygen, it forms carbon monoxide,
also a gas:
2C + 02-^2CO.
2. In consequence of this tendency to imite with oxygen,
carbon is much used as a reducing agent. Thus, when oxide of
copper is heated with pulverized charcoal, carbon dioxide is
formed, and the metal is Uberated:
2CuO + C-^2Cu + C02.
In the same way the oxides of tin, of lead, and of many other
metals may be reduced. Copper, tin, and lead are manufactured
from the ores in this way.
3. Carbon unites directly mih some elements, particularly with
sulphm* to form carbon disulphide CS2 (p. 255) and with certain
of the metals. Thus, when dissolved in molten iron, it forms
iron carbide FesC.
The union with hydrogen is ordinarily too slow to be observed.
But when the carbon is mixed with pulverized nickel (contact
agent) and hydrogen is passed over the mixture at 250®, methane
CH4 is formed (99 per cent). The action is reversible and exo-
thermal, and is therefore, at higher temperatiu'es, less complete
(compare p. 242), at 850® reaching only 1.5 per cent. On the other
hand, an electric arc, between carbon poles in an atmosphere of
hydrogen, gives traces of acetylene C2H2, this action being endo-
thermal. The other compoimds of carbon and hydrogen are all
obtained by indirect reactions.
332 smith's intermediate chemistry
The valence of carbon is almost always four. This is clearly
seen in Civ02^^ and C^vH4^. In a few compounds, of which CO is
the commonest, carbon is bivalent.
Carbon Tetrachloride CCU* — This compomid is manu-
factured by leading dry chlorine into carbon disulphide, in which
a Uttle iodine (contact agent) is dissolved:
CS2 + 3CI2 -> CCI4 + S2CI2.
On distilUng the resulting mixture^ the carbon tetrachloride CCI4
(b.-p. 77°) passes oflf and is condensed, while the sulphur mono-
chloride S2CI2 (b.-p. 136°) remains.
Carbon tetrachloride is a colorless Uquid. It dissolves fats
and tars and other organic compounds, and has the advantage
over benzine and gasoline of being non-inflammable. It is there-
fore used in taking the grease out of wool, linen cloth, oil-bearing
seeds, and bones. " Carbona," sold for dry cleaning and remov-
ing stains from clothing, gloves, etc., is benzine (see p. 345)
to which sufficient carbon tetrachloride has been added to
render the mixture non-inflammable. " Pyrene " fire extin-
guishers contain a Uquid which is mainly carbon tetrachloride.
When the liquid is directed upon burning material, the carbon
tetrachloride is vaporized. The vaporization cools the mass by
using up the heat, and the vapor at the same tin^e displaces the
air and stops the combustion.
Carbides and the Electric Furnace. — Chemical actions
which proceed only at very high temperatures are most economic-
ally carried out by using electricity as the source of heat. In
such cases the electricity has no electrolytic or other chemical
action. There are two types of electric furnaces. In the making
of graphite (p. 330) and of carbon disulphide (p. 255), which
illustrates one of them, the resistance of the carbon furnishes the
occasion for the rise in temperature.
Of the same type is the furnace used for making carborundum
GABBON AND TBE OXIDES OF CABBON 333
(SiC, silicon carbide), mamifactured in large quantitieB at Niag<
ara Falls (Acheson's process). The coke and sand (silicon dioxide
SiOi) are piled between the ter-
minals, and the resistance of
the former causes the produc-
tion of the heat {Fig. 87) :
3C + SiOi -» SiC + 2CX) t •
Here the carbon reduces the
oxide, and combines with the _ „
element (Si) as well. The
product (SiC) is exceedingly hard, and, after pulverization and
mixing with other materials, is moulded into grinding wheels.
In the other type of furnace the air between the terminals
furnishes the resistance, and the arc (a dischai^e carried by the
badly conducting air and carbon vapor) furnishes the heat.
The arc is used in making caldum carbide (CaCe), by heating a
mixture of lime (CaO) and coke:
CaO + 3C^CO + CaCj.
Cold water acts vigorously with calcium carbide, giving acetylene
gas CjHi (see p. 351) and calcium hydroxide (slaked lime) :
CaC + 2H2O — Ca(OH), + C3H2 T ■
Carbon Dioxide
Occurrence. — Carbon dioxide ia found in nature issuing from
the ground, especially in volcanic neighborhoods, and dissolved
in effervescing natural waters, such as Saratoga and Vichy. It is
found in the air (3.5 Uters in every 10,000 liters of air) and in the
breath (37 hters per 1000 liters).
Preparation. — 1, Carbon dioxide is most easUy prepared in
the laboratory by the action of an add such as hydrochloric acid
upon a naiural carbonate hke calcium carbonate (marble or hme-
stone). The action occurs in two stages. The first is a double
334 smith's intermediate chemistry
decomposition, such as all acids and salts exhibit when brought
in contact with one another:
CaCOs + 2HC1 ?± CaCU + HjCOs.
The calcium chloride CaCU remains dissolved in the water con-
tained in the hydrochloric acid. The carbonic acid HjCOs is
unstable, however, and immediately dissociates into water, which
lemains, and carbon dioxide gas, which escapes:
HaCOs^HzO + COit.
The apparatus used is similar to that employed in making chlorine
(Kg. 49, p. 141).
2. For commercial purposes the carbon dioxide is either used
as it is produced, or else it is compressed into wrought-iron cylin-
ders and shipped in the form of a Uquid. Three sources of such
commercial carbon dioxide are in use:
When carbon, for example, in the form of coke, is burned with
a plentiful supply of air, all the carbon is converted into carbon
dioxide:
C + 02-*C02.
Since, however, there are four volumes of nitrogen to one of
oxygen in the air, this carbon dioxide is diluted with nitrogen
in the same proportion. The gases must therefore be separated
by leading them through potassium carbonate solution, which
absorbs the carbon dioxide:
CO2 + H2O :^ H2CO8 + K2CO8 ^ 2KHC0,.
The bicarbonate of potassium KHCOj thus produced is subse-
quently decomposed by heating.
3. Calcium carbonate (limestone) CaCOs, or, more easily, mag-
nesium carbonate (magnesite) MgCOj, may be decomposed by
heating in a kiln:
CaCOs -^CaO + COat.
MgCOs-^MgO + COat.
CARBON AND THE OXIDES OF CARBON 336
In the former ease the quicklime (CaO), which is fonned at the
same time, is a valuable product also.
4. Carbon dioxide is fonned in fermentationy and so is collected
from the vats in which beer is brewed (see p. 417).
Physical Properties. — Carbon dioxide is a colorless, odorless,
almost tasteless gas. As the molecular weight (CO2 = 44) shows,
it is one-half heavier than air. Its greater specific gravity may
easily be shown by pouring it from one jar into another, or into a
beaker placed on one pan of a balance with an equipoise of shot
on the other pan. It is much more soluble in water than is air.
One volimie of water at 15^ dissolves an equal volume of the gas.
At two atmospheres pressure, two volumes are dissolved, at three
atmospheres, three volumes. Pure water, charged at 3 or 4
atmospheres pressure is known as soda water. EflFervescent
waters, such as Selters and Vichy, contain dissolved salts in addi-
tion.
The gas can be liqitefied at any temperature below 31.35° (see p.
91). At 20° the pressure required is 60 atmospheres and this is
therefore the pressure in a cylinder of liquid carbon dioxide at that
temperature. To withstand the pressure very massive cyUnders
are required, and they weigh, when empty, about twice as much as
does the Uquid they will hold when full.
When liquid carbon dioxide is allowed to run from a cyhnder
nto a cloth bag (non-conductor of heat), the rapid evapora-
tion of a part of the Uquid consumes so much heat that the rest of
the Uquid freezes to a snow-Uke mass of solid carbon dioxide. In
the laboratory this soUd is used as a cooUng agent, being mixed
with ether or alcohol to secure closer contact with the object to
be cooled.
Chemical Properties. — Stability. — 1. Carbon dioxide is
very stable (only 7.5 per cent dissociated at 2000°), and so, although
it contains much oxygen, it wiU not support combustion. Mag-
336 smith's INTERBfEDIATE CHEMISTRY
nesium or aluminium powder, however, will bum when placed on a
cake of soUd carbon dioxide and set on fire with burning magne-
sium ribbon :
2Mg + CO2 -* 2MgO + C.
The gas extinguishes burning wood, oil, or candles, and 15 to
16 per cent of it in air is sufficient to extinguish ordinary com-
bustibles. For this reason some fire extinguishers contain a dilute
solution of bicarbonate of sodium (NaHCOa, p. 366) and sulphuric
acid:
2NaHC08 + H2SO4 ^ Na2S04 + 2H2CO8 -^ 2H2O + 200, T-
When the instrument is inverted, these materials are mixed,
and water and carbon dioxide are forced out by the pressure of the
gas.
Chemical Properties — Carbonic Acid. — 2. Carbon dioxide,
when dissolved in water, combines in part to farm carbonic add:
H20 + C02;=iH2C08.
The gas is therefore often called the anhydride (Greek, without
water) of carbonic acid. The solution has all the properties of an
acid, although, as the acid is very little ionized (p. 189), it exhibits
them rather feebly. It tastes slightly sour, turns blue Utmus
faintly red, and neutraUzes bases. The last action is easily shown
by shaking the gas with limewater (solution of calcium hydroxide,
a base) :
Ca(0H)2 + H2CO8 ^ CaC08 i + 2H2O.
The carbonate of calcium is precipitated and the liquid becomes
milky in appearance. This action is used by sugar refiners for
removing the Ume employed in piuifying the sugar. Manu-
facturers of white lead (carbonate of lead) also employ carbon
dioxide, because of its entering into double decomposition to give
carbonates. The same property is utiUzed in making bicarbonate
of sodium and washing soda (carbonate of sodiiun, p. 367).
CAKBON AND THE OXIDES OF CARBON 337
Since the molecule of carbonic acid (H2CO8) contains two atoms
of hydrogen, either one or both of these atoms may be replaced by
a metal — the acid is dibasic. Thus we may have sodium car-
bonate Na2C03, or its hydrate, washing soda Na2CO3,10H2G, and
also sodium bicarbonate (baking soda, sodium-hydrogen carbon-
ate) NaHCOa (see pp. 366-7).
3. The most marvellous chemical action into which carbon
dioxide enters is the action by which plants vse the gas as food.
This important action is discussed in detail in a later chapter.
Carbon Monoxide
Preparation. — Carbon monoxide CO is most easily prepared
in the laboratory by heating formic acid (or sodiimi formate, a
white crystaUine soUd) with concentrated sulphuric acid. The
latter combines with the water, but is not otherwise changed:
HCO2H -^ CO + H2O.
When coke, or any form of carbon bums with a limited supply of
airy or oxygen, the same gas is produced:
2C + 02->2CO.
The gas therefore rises from the surface of a coal fire, sometimes
escaping unbumed, but often burning with a blue flame above
the coal.
Prtpducer Gas and Water Gas. — When air is led through
burning coke, the mixture of carbon monoxide (39 per cent) with
nitrogen (60 per cent) obtained is called producer gas. It is
combustible, and is used in industrial estabUshments for heating
and to drive gas engines for power.
Conmiercially, large amounts of carbon monoxide mixed with
hydrogen (water gas), are manufactured by blowing steam over
white hot coke or anthracite :
C + H2O -> CO + H2 - 28,300 calories.
338 smith's intermediate chemistry
The coke is first set on fire in a brick-lined cylindrical structure
and brought to vigorous combustion by blowing in air for ten
minutes/ Then steam is substituted for the air.
The interaction, as the equation shows, takes place with absorp-
tion of heat. Hence, at the end of a few minutes, the coke be-
comes too cool. It is then necessary to turn the steam oflF and to
turn the air on again, and so on alternately. The mixture of
carbon monoxide (40 to 50 per cent) and hydrogen (45 to 50
per cent), containing also some carbon dioxide (4 to 7 per cent),
nitrogen (4 to 5 per cent), and oxygen (1 per cent), is known as
water gas. It is almost wholly combustible, burning with a blue
flame, and is used as a source of heat and, by driving internal com-
bustion engines, to furnish power. It is used also in manufactur-
ing illuminating gas (see p. 356).
If both air and steam are driven together over the burning coke,
the air enables the coke to biun continuously, and a fuel gas
which is a cross between producer gas and water gas is obtained.
Fuel gases are employed on a large scale in steel works, and
other industrial plants. They give a uniform and easily regulated
heat, they leave no ash, and their use involves no labor for stoking.
Industrial Hydrogen from Water Ga^s. — Hydrogen is
required in large quantities in chemical industry for the manu-
facture of ammonia (p. 300) and for hydrogenating oils (p. 433).
It is essential that this hydrogen should be carefully purified
from traces of other gases, such as carbon monoxide and sulphur-
etted hydrogen, which act as poisons on the catalysts employed in
the above processes. The cheapest source of industrial hydrogen
is water-gaSy and much work has been done to devise a method for
eliminating the imdesirable carbon monoxide from this.
When a mixture of water-gas (substantially H2 + CO) and
superheated steam is passed over a suitable catalyst, such as
iron oxide, a reaction occurs as follows.
CO + H20;;±C0i + Ha + 10,000 calories.
CABBON AND THE OXIDES OF CARBON 339
The reaction being reversible, and exothennie in the forward
direction, its equilibrium point is displaced towards the left,
favoring the backward reaction, as the temperature is raised.
It is therefore desirable to work the process at as low a temper-
ature as possible. With iron oxide alone as a catalyst, however,
the interaction between the gases becomes too slow to be eflFective
at temperatures much below 600°. The addition of small quan-
tities of other oxides, such as nickel oxide and chromium oxide,
has been found to increase the activity of the catalyst very con-
siderably. Substances which act in this manner (catalyzing a
catalyst, so to speak) are termed promoters.
In actual practice the reaction, carried out at 450-600®, gives
a mixture of gases containing only about 2 per cent residual
CO. Excess of steam is employed to drive the equiUbrium as
far as possible towards the right, but the excess of hydrogen
present in the original water-gas favors, of course, the opposite
reaction (compare p. 234). It should be noted that nearly twice
as much hydrogen as was contained in this water-gas is obtained
by the process, the second half being derived from the decomposi-
tion of the steam.
The bulk of the carbon dioxide present in the final mixture
(approximately 30 per cent by volume) is removed by washing
the gas with water imder pressure. The last traces of carbon
dioxide are absorbed by means of lime or alkaUes. The removal
of the 2 per cent residual CO presents difficulties. Absorption
of CO by hot caustic soda solutions and by anmioniacal solutions
of cuprous salts has been employed. The most efficient method,
however, consists of preferential combustion of CO to CO2 with the
requisite quantity of air or oxygen in the presence of a second
oxide catalyst. If due precautions are taken, CO bums almost
quantitatively to CO2 without any Hj present burning to H2O.
The small amoimt of CO2 formed is then removed as already
described.
340 smith's intermediate chemistry
Physical Properties of Carbon Monoxide. — Carbon monox-
ide is a colorless, odorless, and tasteless gas. It is a little lighter
than air (mol. wt. 28), and is very slightly soluble in water. It is
difficult to liqicefy. Its boiling-point, when liquid, is —190°,
close to that of liquid air.
Chemical Properties. — When set on fire, the gas burns in
air or oxygen with a blue flame. Carbon dioxide is formed, and
the presence of the latter may be shown with Ume-water (p. 336) :
2C0 + O2 -> 2CO2.
On account of this property, carbon monoxide reduces the oxides
of the less active metals, such as those of iron and of the metals
below iron in the order of activity. Commercially, the ores of
iron are reduced by this gas (essentially producer gas) in the blast
furnace. The oxides of the metals above iron are not reduced.
Physiological Properties. — The gas is an active poison, and
1 volume in 100,000 volumes of air produces symptoms of poison
ing, while one volume in 760 to 800 voliunes produces death in
about thirty minutes. The gas combines with the haemoglobin
of the blood corpuscles, forming a stable compound, and thus
preventing the absorption of oxygen by the blood (p. 34). This
gas is the chief poisonous substance in illuminating gas. The
poisonous effect of tobacco smoke, particularly when inhaled, is
due mainly to the carbon monoxide produced by the necessarily
incomplete combustion.
Combustions or explosions in confined spaces (such as in a mine-
shaft, or in the interior of a warship during an engagement) may
cause many deaths through CO poisoning. Gas masks for use
in rescue work in such cases are fitted with canisters containing
a mixture of metaUic oxides, as Mn02, CuO, C02O3 and Ag20 (hop-
calite). A mixture of this kind acts catalytically, any carbon
monoxide passing into the canister being oxidized to carbon
dioxide by the oxygen of the air.
CARBON AND THE OXIDES OF CARBON 341
Exercises. — 1. (a) What physical property of graphite enables
it to cover the surface of a stove so eflFectively? (b) How does
" poUshing " with a brush contribute to the result? (c) Why not
use paint on a stove? (d) Explain why graphite can be used as a
lubricant.
2. If a metal formed the positive electrode (anode) in electrolyz-
ing sodium chloride solution, what chemical change might it
undergo (p. 54), and which metals would be least rapidly attacked?
What objection is there to using the latter metals in practice?
3. When one cubic meter of oxygen acts upon carbon, what
volumes (at the same temperature and pressure): (a) of carbon
dioxide; (b) of carbon monoxide can be obtained?
4. Make the equation: (a) for the formation of methane by
union of carbon and hydrogen; (b) for the reduction of stannic
oxide (Sn02) by carbon.
5. Make equations for: (a) the action of sulphuric acid upon
calciimi carbonate; (b) carbon dioxide on sodium hydroxide solu-
tion (p. 336) ; (c) the burning of aluminium in carbon dioxide.
6. From the fact that the molecular weight of carbon dioxide is
44, how do we infer that it is one-half heavier than air?
7. Why does soda water remain quiescent in the closed bottle,
and why does it eflFervesce when the bottle is opened?
8. Rewrite the equati(&is on p. 336 in full ionic form.
9. Name the variety of chemical change (p. 132) to which
belongs the reaction shown in each equation in this chapter.
10. Assuming that air contains oxygen and nitrogen in the pro-
portion of 1 : 4 by volume, what are the theoretical proportions of
carbon monoxide and nitrogen in producer gas?
11. (a) What volume of water gas is produced from each liter of
steam, and (b) what is the proportion of the component gases in
the product? (c) What impurities should you expect to find in
water gas? (d) How should you attempt to separate the com-
ponents of water gas?
342 smith's intermediate chemistry
12. Why is water gas an especially valuable source of heat when
high temperatures are required?
13. Make a list of metals the oxides of which would be reduced
by carbon monoxide.
CHAPTER XXIX
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME
The compounds of carbon and hydrogen are called the hydro-
carbons. Hundreds of different hydrocarbons, containing differ-
ent proportions of the two elements, are known. The natural
oil petroleum is a mixture of many substances of this class.
The hydrocarbons fall into several distinct series, the chief
one of which contains methane CH4 as its simplest member.
On account of the fact that certain members of this set are found
in parajBSn, it is conmionly known as the paraffin series. For
the reason that in this series the carbon has all its four valences
employed, the members are also called the saturated hydrocar-
bons.
Paraffin or Saturated Series of Hydrocarbons. — The
foUowing is a Ust of the names, formulae, and boiling-points of
seven of the simplest hydrocarbons of this series, and of two of the
higher members of the series:
Methane CH4
b.-p.
- 164°
Hexane CeHu b.-p. 71°
Ethane CaHa
- 89.5°
Heptane CyHw 99°
Propane CsHs
-37°
Hexadecane Ci«H84 287 . 5°
Butane C4H10
+ r
m.-p. 18°
Pentane CeHu
35°
Pentatriacontane CiJ3.n m.-p. 74.7°
After the fiffst four, the names are based on the Greek numerals
corresponding to the number of carbon atoms in the molecule.
On comparing the formulae, we observe that in each, the number of
units of hydrogen is equal to twice the number of carbon units
plus two. The general formula is therefore CnH2n+2. The series
affords a striking illustration of the law of multiple proportions.
We note, further, that the first four are gases at the ordinary tem-
343
344 smith's intermediate chemistry
perature. The members of the series from pentane to penta-
decane (C16H32) are liquid mider ordinary conditions. From
hexadecane onwards they are solids, with higher and higher melt-
ing-points.
In these compoimds the carbon is quadrivalent, and each sub-
stance is related to the preceding one by containing the additional
units CHa. The formulae of the fib-st three members may be writ-
ten graphically to illustrate these two facts:
H
H H
H H H
1
1 1
1 1 1
H-C-H
H-C-C-H
H-C-C-C-H
1
1 1
1 1 1
H
H H
H H H
Petroleum. — Petroleum is a thick, greenish-brown oil. When
borings are made into the oil-bearing strata, the oil either gushes
up, or is pumped to the surface. In the United States many thous-
ands of miles of pipe-lines are used to transport the oil, with the
aid of force pimips, to the refineries, and in 1920 nearly 450 mil-
Uon barrels (42 gal. each) were produced. The world's produc-
tion in 1920 was 700 million barrels.
After the United States, Mexico and Russia are the chief
producers of petroleum.
Oil Refining. — The natural oil is a complex mixture, and is
partially separated by distillation (p. 67) into products which
are still mixtures, but are suited to special purposes. The com-
ponents of lower boiUng-point come off first and the temperature
rises steadily as these components are eliminated and those of
higher and higher boiling-point enter the vapor. As certain
temperatures are reached (or as the sp. gr. of the distillate attains
certain values) the condensed hquid is diverted into different
vessels, so as to collect together the " fractions '' of the same kind.
This is called fractional distillation.
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 345
At some suitable stage, the residual oil is chilled, and a quan-
tity of the soUd members of the series (C22H46 to C28H58) crystal-
lizes in flakes (solid paraffin) and is separated by filtration in
presses. The final residue is used for lubricants and for fzcel.
The fractions are still mixtures, but contain mainly colnpoimds
lying close together in the series. Some of the products are as
follows:
Name
Main Components
B.-P.
Uses
Petroleum ether
Gasoline
Pentane, hexane ....
Hexane, heptane
Heptane, octane ....
Octane, nonane
Decane-hexadecane. .
40°- 70°
70°- 90°
80°-120°
120°-150°
150°-300°
Solvent, gas-making
Solvent, fuel
Naphtha
Solvent, fuel
Benzine
Solvent
Kerosene
Illuminating oil
Vaselinei C22H46 to C^H48, is separated in some refineries.
Solid paraffin is employed for waterproofing paper, as an ingredi-
ent in candles, and in making chewing gum.
Asphalt, a natural mixture of the sohd hydrocarbons, found
particularly in Trinidad, is used in road-making.
Oil Shale. — In Scotland, petroleum is also obtained by heating
shale. The shale is a clay deposit, which contains no oil as such,
but which when heated gives off fuel gas, ammonia (see p. 299),
petroleum oils, and many valuable hydrocarbon derivatives.
The richer shales yield from 30 to 40 gallons of oil per ton. The
oil usually contains a much larger percentage of unsaturated
hydrocarbons than well petroleum. The yield can be somewhat
increased, and the proportion of undesirable unsaturated hydro-
carbons diminished, by blowing superheated steam into the
retorts during the distillation.
Natural Gas: Methane CH4. — Natural gas is obtained from
wells, tapping strata close to those which contain petroleum, and
in the same locahties. It often issues under very high pressure.
346 smith's intermediate chemistry
It owes its combustibility to its chief component (over 90 per cent),
methane CH4. It is largely used as a fuel in the regions in which
it is found and, in the United States, the annual value of the gas
so consiuned is nearly $160,000,000 (1919). The same gas issues
from many coal seams (" fire-damp ")> ^^^ forms explosive
mixtures with the air of mines. It rises to the surface when stag-
nant pools containing decomposing vegetable matter are stirred
(" marsh-gas ")•
The formation of methane by direct union of carbon and hydro-
gen has already been discussed (p. 331).
Chemical Properties of the Hydrocarbons. — The hydro-
carbons, whether pure or in solution, show no conductivity for
electricity. They have none of the chemical properties of acids,
bases, or salts, and therefore do not enter into double decom-
positions with substances of these classes. The satiu*ated hydro-
carbons are in fact quite indifferent to the presence of most chem-
ical reagents.
All the hydrocarbons hum with oxygen or air to form carbon
dioxide and water:
CH4 + 2O2 -^ CO2 + 2H2O.
C7H16 + IIO2 ^ 7CO2 + 8H2O.
The water can be shown by its condensation on a cold vessel
held over the flame. The carbon dioxide gives a precipitate
of calciimi carbonate (p. 336) when the gases rising from the flame
are drawn through lime-water.
All the hydrocarbons, when heated strongly (air excluded),
decompose or crack. They usually lose a part of their hydrogen
and become unsaturated. These of high molecular weight break
up to give a mixture of hydrocarbons of low molecular weight.
Ethylene C2H4, for example, is produced in large amoimts by
heating the higher members of the series to a red heat. On the
other hand, the lower members of the series, when heated, often
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 347
give compounds of higher molecular weight. Thus, methane
gives ethylene and acetylene, along with hydrogen:
2Cri4 — > C2H4 -|- 2xi2.
2Cxi4 — > C2H2 "f" 3H2.
At a white heat all the hydrocarbons decompose into hydrogen
and free carbon.
The latter is deposited in a dense form called gas-carbon, which
is used in making carbon rods for arc lights and electric furnaces,
and carbon plates for batteries, and for the electrodes employed
in electrolysis. The carbon is ground up, moistened with petro-
leum residues, subjected to hydraulic pressure and finally heated
strongly to expel volatile matter.
Derivatives of the Hydrtfcarhons. — Although the hydro-
carbons are themselves almost inert chemically, yet many impor-
tant classes of organic substances may be regarded as their deriv-
atives, one or more atoms of hydrogen in the graphic formula
(p. 344) being replaced by other elements. The following tab-
ulation should be carefully studied, and the graphic formula
of each compoimd mentioned should be written down by the
student.
1. Halogen Derivatives. When a mixture of methane and
chlorine is exposed to sunUght several successive changes occur:
CH4 + CI2 ^ HCl + CH3CI (methyl chloride)
CHsQ + CI2 -^ HCl + CH2CI2 (methylene chloride)
CH2CI2 + CI2 -^ HCl + CHCI3 (chloroform)
CHCla + CI2 ^ HCl + CCI4 (carbon tetrachloride).
Chloroform CHCI3, used as an ansesthetic, and carbon tetra-
chloride CCI4 (p. 332) are familiar substances. Iodoform CHIs
is employed in surgical dressing. These substances are not salts,
and are not ionized in solution.
348 smith's interbiediate chemistry
2. Hydroxy I Derivatives (Alcohols). Water acts dotdy
upon methyl chloride, according to the equation:
CHaQ + H2O ;=±HC1 + CH3OH (methyl alcohol).
The reaction is reversible and incomplete, but can be accelerated
and carried to completion by addition of a base, which removes
the HCl as fast as it is formed. Ethyl chloride CaHgCl gives
ethyl alcohol C2H6OH.
Although containing the radical OH, the alcohols are not ion-
ized in solution, and are therefore not bases. They are extensively
used as solvents for other organic substances (see also p. 418).
3. Oxygen Derivatives (Ethers). When ethyl alcohol and
concentrated sulphuric acid are heated to 140*^, water and ether
(C2H6)20 distil oflf. The action occurs in two stages:
(1) C2H5OH + H2SO4 -> H2O + C2H6.HSO4
(2) C2H6.HSO4 + C2H5OH -> H2SO4 + (C2H5)20.
Ether (C2H6)20 is a very volatile and inflammable liquid, used
as an anaesthetic and as a solvent for resins, fats and oils.
4. Aldehydes. By fractional combustion of methyl alcohol
(passage of a heated mixture of alcohol vapor and air over a metal
catalyst), formaldehyde H.CHO is obtained:
2CH8.0H + O2 -> 2CH2O + 2H2O.
Formaldehyde is a gas. Its solution in water (formalin)
is employed as an antiseptic and disinfectant. Its property of
hardening gelatins makes it valuable in the leather industry
and in the manufacture of artificial silk. It is also used in making
dyes and in the production of bakelite (p. 481). The correspond-
ing derivative of ethyl alcohol is acetaldehyde CHs.CHO. The
group .CHO is characteristic of aldehydes.
5. Acids. By fiu1;her partial oxidation, alcohols or aldehydes
give members of the fatty acid series. Thus ethyl alcohol can
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 349
be oxidized directly to acetic acid CH8.C00H by passing a mixture
of the vapor, with air, over specially prepared platinum as a
catalyst:
C2H6.OH + 02-^ CHs-COOH + H2O.
The first acid of this series is formic acid H.COOH, a corrosive
liquid secreted by red ants and present in stinging nettles. Acetic
acid (see p. 419) has many industrial uses.
The lower members of the fatty acid series are perfectly mis-
cible with water, and are slightly ionized in aqueous solution.
By neutraUzation with bases we obtain salts, such as the formates
and the acetates. Only the hydrogen of the characteristic .COOH
group, it must be noted, is replaceable by metals.
6. Ketones. When calciiun acetate Ca(CH8.COO)2 is heated,
acetone (CH8)2.CO distils off:
Ca(CHs.COD)2 -> CaCOa + (CHa)2.C0
The ketones resemble 'ohe aldehydes in many respects, but their
characteristic group :C0 is not directly combined with hydrogen.
Acetone is a Uquid boiling at 56^, used in large quantities in the
industries as a solvent.
7. Esters. Alc^Dhols and acids interact slowly and incom-
pletely to form esters. Thus when ethyl alcohol and acetic acid
are used, we obtain ethyl acetate C2H5.COO.CH8:
C2H5.OH + CHs-COOH ^ H2O + C2H5.COO.CH8
The action may be catalyzed by the addition of a little sulphuric
acid.
The equation, as given above, bears certain resemblances to a
neiUralization. It differs sharply, however, from a true neutraH-
zation in several respects. An alcohol is not a base, neither is an
ester a salt. Both classes of substances are non-ionized in solu-
tion. True neutralization takes place instantaneously, while
the foregoing action, and all like it, proceed very slowly.
350 smith's intermediate chemistry
The esters form the sweet-smelling constituents of plants.
Many are now produced synthetically as substitutes for natural
and fruit essences.
A few additional classes of hydrocarbon derivatives will be
given later (p. 353).
Unsaturated Hydrocarbons. — In addition to the saturated
series of hydrocarbons, several other series are known in which
smaller proportions of hydrogen are present. Thus, ethylene
C2H4, to which illuminating gas largely owes the luminosity of its
flame, belongs to a series CnH2», all the members of which contain
two atoms of hydrogen less than the corresponding compoimds
of the first series. Again, acetylene C2H2 is the first member of a
series CnH2»-2, and benzene CeHe begins a series CnHan-e. These
are all unsaturated because the full valence of the carbon is not
in use, and these compoimds, therefore, i:nite more or less readily
with hydrogen, chlorine, bromine, and - ^ncentrated sulphuric
acid. The hydrocarbons of all the series.^ re mutually soluble,
but none of them dissolve in water.
Members of the ethylene and acetylene vseries are found in
petroleum, and are formed also to some exten . by decomposition
during the distillation. As oil containing thtm acquires dark-'
colored products by chemical change, the oils t^re always refined
before being sold. They are agitated with conce^itrated sulphuric
acid, which unites with the unsaturated substa*nces and, being
insoluble in the oil, collects in a layer below it. Vhe oil is finally
washed free from the. acid with dilute alkali and ^vith water.
Ethylene €2^^. — The formation of one molecu^^e of ethylene
from two molecules of methane with elimination of fi wo molecules
of hydrogen suggests its graphic formula:
H H H H iH H
I I II III
H-C-H + H-C-H -^ H-C-C-H or H-lC = C-H
I I II
H H
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 351
Ethylene is a gas which bums in the air with a highly luminous
flame (see p. 356). It combines directly with bromine to form
ethylene bromide C2H4Br2.
The hydrocarbons of the ethylene series are known as olefines.
They are of value as illuminants. Their derivatives are similar
in character to those of the paraflin hydrocarbons, but are more
active chemically in view of their imsaturation. When heated
with hydrogen in the presence of a catalyst, such as finely-divided
nickel, they give the corresponding saturated derivatives (com-
pare pp. 433-4).
Acetylene C2H2--^ A mixture, containing acetylene, is formed
when any hydrocarbons is heated strongly (p. 346), air being
excluded. As in the case of ethylene, the formation from methane
by loss of hydrogen (p. 347) suggests the graphic formula:
H . H
I I II
H-C-H-fH-C-H->H-C-C-H or H-C=C-H
I I II
H H
Pure acetylene is prepared by the action of water on calcium
carbide (p. 333) :
CaC2> 2H2O ->,Ca(0H)2 + C2H2 T •
Calcium hydroxide (slaked lime) remains. The gas bums with a
flame even more luminous than that of ethylene. It is there-
fore made in generators by the foregoing action for use on auto-
mobiles and for hghting buildings remote from a pubhc supply of
illuminating gas. Acetylene tanks, which are also in use, contain
acetylene dissolved, under high pressure, in acetone.
The Acetylene Blowpipe or Torch. — Acetylene decomposes,
when heated, with liberation of heat:
C2H2 -> 2C + H2 + 58,100 cal.
When acetylene bums with oxygen, therefore,
2C2H2 + 5O2 -^ 4CO2 + 2H2O,
352 smith's intermediate chemistry
we obtain not only the heat due to the combustion (p. 162) of
the carbon to carbon dioxide (4 X 96,820 cal.) and of the hydrogen
to water (116,200 cal.) but also the heat due to the decomposition
of the gas (2 X 58,100 cal.). The temperature of the flame is,
therefore, the highest that can be reached by the combustion of
any easily obtainable gaseous mixture. The oxy-acetylene flame,
produced by means of a suitable burner (Fig. 29, p. 56), the gases
being furnished from small, portable tanks, is now used for cutting
metals. Such a flame will melt its way through a 6-inch shaft
of steel, or a heavy steel plate several feet wide, in less than one
minute, cutting the object in two. Steel buildings have been taken
apart rapidly by this device.
Blau gas and oil gas, mixtures of hydrocarbons made by " crack-
ing " (see p. 346) heavy oils, are now largely displacing acetylene
for uses hke those just mentioned. They give flames which are
almost as effective, and are more easily controlled. • Even the
oxy-hydrogen torch is remarkably eflicient, when applied to the
same purposes.
Benzene CeHe* — This is the first member of the aromatic
hydrocarbon series. It may by synthesized by heating acetylene
in a closed vessel at a moderately low temperature:
3C2H2 — ^ Celie.
In practice it is obtained, with many of its valuable derivatives,
as a by-product in the production of coke (p. 424).
The graphic formula of benzene is represented as a closed
ring structure:
H
I
/^%
H-C C-H
II I
H-C C-H
\C^
I
H
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 353
The hydrocarbons of the benzene series exhibit, in fact, many of
the properties of olefines, combining directly with hydrogen and
with the halogens to form satm'ated compoimds such as hexa-
hydrobenzene C6H12 and benzene hexabromide CeHcBre.
The second member of the aromatic series is toluene CeHs-CHs.
The third is xylene C6H4.(CH8)2. Naphthalene CioHs and anthra-
cene C14H10 are members of more complex series, containing n^ore
than one ring.
Of all hydrocarbons, those of the aromatic series, with their
derivatives, are the most important. They are of particular
significance in the dye-stuff and explosive industries, and in the
manufacture of synthetic perfumes and drugs. They give alde-
hydes, acids, esters, etc., similar to those Usted under the paraflin
hydrocarbons (pp. 347-50). In addition to these, the following
extremely valuable classes of derivatives should be noted:
(1). Phenols. The substitution of hydroxyl for a hydrogen
atom in benzene gives phenol CeHeOH, a substance quite different
in many of its properties from an alcohol, although it resembles
the alcohols in forming esters with acids (see p. 349). Phenol,
when pure, is a colorless soUd melting aroimd 40°, with a charac-
teristic odor. It is strongly antiseptic, corrosive and poisonous.
In solution it is a weak add.
(2). Nitro'compounds. Nitrobenzene C6H6NO2 is obtained
by the action of a mixture of concentrated nitric and sulphuric
acids upon benzene in the cold:
CeHe + HNOs -^ CeHe.NOj + H2O.
It is a pale yellow hquid, with a smell resembUng bitter almonds,
and is used in scenting cheap soap.
(3). Amino'Compounds. Reduction of nitrobenzene by a
metal in acid solution gives aniUne C6H5.NH2. The amino-
compoimds are derivatives of ammonia, and their solutions
accordingly are weakly basic. Aniline, when pure, is a colorless,
oily hquid, boiUng at 185°. It is the parent substance of the
countless aniline dyes.
354 sboth's intermediate chemistry
In subsequent chapters we shall return to the different classes
of organic substances, tabulated here and on pp. 347-50, and dis-
cuss their properties and industrial uses in greater detail.
Flame
We have encountered a variety of flames, from the simple one
of hydrogen burning in air to the more compUcated case of the
luminous flame of ethylene or acetylene. The subject will now
repay a somewhat closer study.
The Simple Flame. — The flame of hydrogen (giving water),
or of carbon monoxide (forming carbon dioxide) is very simple in
structure (Fig. 88). We find that there is a tapering colimm of
imbumt gas in the interior, surrounded by a layer of hot
gas — the flame itself. The flame is therefore a hollow
cone. That the flame is hollow is easily shown by hold-
ing a wooden match across it. The match is charred at
the two points at which it crosses the flame, and remains
unheated in the middle. These flames are simple, be-
cause only one chemical change occurs in them. The
flames are rather large, because suflScient oxygen to
bum all the gas does not reach the latter at once, and the gas
travels upwards and diffuses outwards a certain distance before
being all consmned.
If oxygen is substituted for air, by lowering the jet into a jar
of that gas, the flame becomes much smaller. In the absence
of atmospheric nitrogen, there is now five times as much oxygen
within a given range of the center of the jet as before. This
chemical union, Uke any other, proceeds more rapidly with an
increase in the concentration of the interacting substances (p. 233).
It is therefore completed before the gas has time to diffuse very
far from the opening of the jet.
The Candle Flame. — A candle is made of a mixture of paraffin
and stearic acid (one of the higher members of the fatty acid
THE HYDROCARBONS AND THEIB DERIVATIVES. FLAME 355
series (p. 348), with the formula C17HM.COOH, made from fat).
When it bums, the whole phenomenon is vastly more comphcated
than the burning of hydrogen. The following are some of the
stages in the process, which is operated by the flame's own heat.
.To start with, the wax is melted and ascends the wick by capillary
action. This is merely a physical phenomenon. Then the
chemical changes begin. (1) The melted compounds of carbon
are decomposed by the heat (cracked, p. 346), being turned into
more volatile compounds and gases which occupy the central
hollow of the Same. (2) The compounds forming the gases and
vapors are further decomposed at a white heat, giving free carbon
and hydrogen (p. 347). (3) All the materials finally reach
a sufficient supply of oxygen and are burned to water and
carbon dioxide. There are thus three chemical chaises, each of
which takes place in a definite region that can be
observed by the eye {Fig, 89), The formation of
the gases from the melted wax (without gas, there
would be no flame) takes place in the dark central
region where there is no oxygen. The carbon is set
free and glows brilliantly in the liuninous cone that
surrounds the gas and extends far above it. The
final combustion occurs in a fainter cone of flame
covering the whole exterior.
That there is unbumt gas (produced by decomposition of the
wax) in the center is easily shown by inserting a narrow tube,
throi^h which some of the gas will ascend. The free carbon in
the luminous zone will show its presence by blackening a cold
dish placed across the flame.
Lampblack. — When an iron vessel, cooled by a stream of
water circulating through it, is suspended in the luminous flame of
natural gas or burning petroleum, the carbon (soot) is deposited
on the vessel. By rotating the latter, the soot can be continu-
ously scraped off by a stationary piece of metal. The product,
366
smith's intermediate chemistry*
lampblack, mainly very finely divided carbon, is used in making
printers' ink, India ink, and black varnish.
Carburetted Water Gas. — To fit water gas, essentially
H2 + CO (p. 337), which bums with a pale blue flame, for pubhc
service as an illuminating gas, unsaturated hydrocarbons, and
particularly ethylene C2H4, which bum with a highly luminous
flame, must be added. The water gas is passed through a tower,
filled with strongly heated brick on which oil is continually sprayed.
Mixed with the vapor of the oil, the gas goes into the " super-
heater " where, at a higher temperature, the decomposition into
unsaturated hydrocarbons (cracking) takes place. The gas is
then cooled and washed to remove the condensible hydrocarbons,
which would otherwise collect in the service pipes with resulting
waste of combustible material as well as obstruction in the deUvery
of the gas. A typical carburetted water gas has the composi-
tion: Illuminants (largely ethylene) 16.6 per cent; heating
gases — methane 19.8 per cent, hydrogen 32.1
per cent, carbon monoxide 26.1 per cent; im-
purities (nitrogen and carbon dioxide) 5.4 per
cent.
Carburetted water gas has now largely sup-
planted coal gas (p. 423) for Ughting and heating
purposes.
NoTt'luminous Gas Flames. — When gas is to
be used for heating, the complete combustion of
the gas, without any intermediate Uberation of free
carbon, is desirable. This is achieved in the
Bimsen burner (Fig. 90) by admitting air at the
bottom of the burner, in such a way that the air
mixes with the gas before the latter reaches the
flame. The air cools the middle zone of the flame, so that at
this point the temperature required for dissociating the ethy-
FiG. 90
THE HYDROCARBONS AND THEIR DERIVATIVES. FLAME 357
lene, and liberating carbon, is not reached. The oxygen in
the air plays no part — mixing carbon dioxide or pure nitrogen
with the gas has exactly the same effect. A flame of this kind
is non-luminous.
Although the middle zone of the non-luminous flame is cooler
than that of the luminous flame, the average temperature of the
flame as a whole is higher. This is the case because the same
total amoimt of heat is Uberated in both cases, but the non-
luminous flame as a whole is smaller in size.
The Bunsen type of burner, placed in a horizontal position
(Fig. 91) is used in the ordinary gas cooking range. As with the.
Bimsen burner, some care is required to get good results. The
holes which admit the air to the mixer must be kept clear of
obstructions, as otherwise luminous flames are produced, smoke
and soot are formed, and less heat is generated. The size of
the openings must also be adjusted so that
the admission of too much air will not cause
the flame to flash down the burner, and set
fire to the gas within the mixer.
Flames with Incandescent Mantles.
— When gas is burned in a Bunsen burner,
a bright Ught may still be obtained from
the flame. This is managed by suspending
in the flame a structure (" mantle ") made
of the oxides of thorium (99 per cent) and
of cerium (1 per cent). These oxides act as
a cordad agent, hastening the combustion
and hberation of heat close to their sur-
face, which thereby becomes incandescent.
The hght has about ten times the illuminating power of a flat
flame burner using the same amount of gas.
Exercises. — 1. When vegetable matter decays in the air
the carbon it contains is finally all turned into carbon dioxide.
Fig. 91
358 smith's intermediate chemistry
When the same matter decays imder water, it gives methane (p.
346). Explain the difference in the result.
2. What is the density (air = 1) of (a) methane, (b) ethylene?
3. Write the graphic fonnulae for propane, propyl chloride,
propyl alcohol, dipropyl ether, propionaldehyde, propionic acid,
dipropyl ketone, propyl propionate.
4. Write the graphic formulae for toluene, phenol, nitrobenzene,
aniUne.
5. (a) Given a flame of hydrogen burning in a jar of air, what
would be the effect on the flame of lowering the pressure of the
air by means of an air pump? (b) What would be the effect on
the average temperature of the flame? (c) How about the heat
produced by burning 1 g. of hydrogen in each case? (d) What
differences would be observed in using an alcohol lamp at the
bottom and on the top of a high moimtain?
6. (a) In the candle or gas flame, what is the source of the
light? (b) Why does such a flame become smoky when placed
in a draft?
CHAPTER XXX
SILICON; BORON
Silicon belongs to the carbon family, being, like carbon,
quadrivalent and non-metallic.
Although silicon does not occur free in nature, yet its com-
pounds are so plentiful that about 26 per cent of the terrestrial
globe is sihcon. Instead of naming all the rocks which contain
it, such as sandstone, basalt, granite, and so forth, it is easier
to say that hmestone is the only common rock which is not siliceous.
Silicon Si. — The element is now manufactured at Niagara
Falls and elsewhere, by heating sand (Si02) with coke in an elec-
tric furnace. The process closely resembles that for making
carborundum (p. 333), except that less coke is used:
2C + SiOa -> 2C0 t + Si.
The element, as prepared in this way, is a grey, crystalhne
material.
Silicon Dioxide Si02 (Silica), Physical Properties. — Color-
less rock-crystal, often showing large hexagonal crystals, is pure
sihcon dioxide, deposited from natural solutions. When im-
purities enter into it, smoky quartz, rose quartz (pink), and
amethyst (violet) are formed. Often the impurity changes dur-
ing the growth of the deposit, and beautifully variegated spe-
cimens, Uke jasper, catseye, and agate are produced. Chalcedony,
opal, and flint contain a small amount of water in combination.
The nodules (rounded masses) of flint break in spKnters, when
struck, and our prehistoric ancestors dexterously fashioned their
implements and weapons from this material. The soUd struc-
ture of sponges and diatoms is also hydrated sihcon dioxide.
359
360 smith's intermediate chemistry
Diatomaceous earth (Tripoli powder) is used in making polishing
powders and for removing coloring matters from oils.
Chemical Properties. — Silicon dioxide (sand) is not acted
upon by acids, with the exception of hydrofluoric acid, which
gives silicon tetrafluoride SiF4 and water:
SiOa + 4HF -^ SiF4 T + 2H2O.
When siUcon dioxide is fused with sodium carbonate, carbon
dioxide gas is hberated, and soditun silicate Na2Si03 is formed:
Si02 + Na^COs -^ CO2 T + Na^SiOs.
The resulting salt is very soluble in water and a strong solu-
tion is sold imder the name of water-glass or soluble glass. This
material is used as a filler in cheap soaps, as an ingredient in
artificial stone, a coating to render wood or cloth fireproof, and a
cement for uniting glass or porcelain. Eggs are preserved by
being submerged in a solution of this salt.
When an acid is added to sodium aiUcate solution, silicic acid
H4Si04 or H2Si08 (gelatinous, the degree of hydration varies with
the conditions) is precipitated. If this acid is heated strongly,
sihca Si02 remains as a powder.
Incompletely dehydrated sihcic acid, containing 5-7 per cent
of water {silica gel), is employed as an adsorbent material for
recovering valuable vapors (such as sulphur dioxide, oxides of
nitrogen, and volatile organic solvents) from the issuing gases
in many large-scale industrial processes. The adsorbed vapors
are given up on heating, and the gel is ready for renewed use.
SiKca gel is also of service in deodorizing petrolemn oils, and as a
catalyst (see p. 262).
Silicon Tetrachloride SiCU. — This compound is made by
direct imion of the free elements. It is more conveniently pre-
pared by passing chlorine over a strongly heated mixture of siUcon
I
I
s 1
is
silicon; boron 361
dioxide and carbon. The gaseous products enter a condenser in
which the tetrachloride assumes the liquid form:
2CI2 + Si02 + 2C -^ SiCU + 2C0.
Silicon tetrachloride is a coloriess liquid (b.-p. 59°) which fumes
strongly in moist air, giving siUcic acid and HCl. Mixed vapors
of SiCU, NH3 and H2O produce a very dense white smoke, con-
sisting of minute particles of NH4CI and silicic acid. This smoke
was utilized during the war for screening vessels from submarines.
Glass. — Calcium carbonate (limestone) interacts at a high
temperature with sand in the same way as does sodium car-
bonate:
Si02 + CaCOs -^ CO2 T + CaSiOs
giving calcium silicate. Now sodium silicate, when alone, is
soluble in water. Calcium siUcate is insoluble, but forms a
brittle, crystaUine mass. By using both sodium and calcium
carbonates, and employing a larger proportion of sand than
that shown in the equation, a material is obtained which has
the qualities required in glass. When cooled, the molten mass
becomes viscous and finally, for all practical purposes, soKd»
Yet it does not crystaUize — it is amorphous. It is also prac-
tically insoluble in water.
By pouring the viscous material into moulds, or stamping
it with dies, articles of pressed glass are obtained. Bottles are
blown, by taking up a suflScient mass of the hot, thick Uquid on
the end of an iron tube, inserting it in a mould, and blowing until
the outline of the mould is filled. Window glass is made by
blowing an immense, elongated bubble (6 by 1^ ft.), ripping it
while still hot and soft, and flattening it out. Plate glass for
windows and mirrors is manufactured by pouring out the material
upon a cast-iron table, with a raised rim, and passing a large,
heated iron roller over it. The plate is subsequently ground flat
on both sides and poKshed with rouge (Fe203).
362 smith's intermediate chemistry
Soda-calcium glass is called soft glass, because it is easily soft-
ened by heating. When potassium carbonate is substituted for
sodium carbonate, a less fusible substance, used in making some
chemical apparatus, and called hard glass, is obtained. When
lead oxide is employed in place of the limestone, a potassium-lead
siUcate K2Si03,PbSi03,xSi02 is formed which, on being cooled,
gives flint glass. This glass has a higher density and greater
briUiancy than soft glass, and is used in making vessels of cut glass
and lamp chimneys. The cutting is done with a revolving grind-
ing wheel.
When glass is allowed to cool quickly, the product is very brittle
and apt to crumble to pieces on receiving a shock or scratch.
Glassware is therefore all annealed, by being passed on a slowly
moving frame through a long furnace, which is very hot at the
entrance and much cooler at the exit.
Colored glass is made by adding oxides of metals which, with
the siUca, give colored siUcates. Oxide of chromium gives green
siUcates, oxides of copper and of cobalt blue siUcates, and oxide of
manganese violet. Gold oxide is reduced to the metal, which
goes into colloidal solution and gives ruby glass. Milky glass is
made by adding calcium fluoride, or stannic oxide. The green
color of bottle glass is due to iron (ferrous siUcate) derived from
impure sand or Umestone.
The rough surface of grotmd glass is produced with a sand blast.
For engraved glass, the surface is covered by a stencil to pro-
tect it from the sand blast, and only the pattern is left exposed.
In granite iron ware the surface is covered with a thin layer of
easily fusible glass (enamel, see borax).
Pure quartz can be melted in the oxy-hydrogen blowpipe,
and recently chemical apparatus (silica ware) has been made
out of it. It has the advantage of being less soluble than glass,
and of not breaking even when it is heated white hot and quenched
in cold water. Glass breaks when chilled, because the parts first
cooled shrink considerably and a great strain is produced. Quartz
silicon; boron 363
suffers very little change in volume with change in temperature,
and so unequal cooUng causes almost no strain.
Pyrex glass, a borosiUcate, has also come into extensive use
lately for laboratory ware and for cooking vessels. Its low co-
efficient of expansion renders it much less Uable to crack under
sudden temperature changes. Like pure siUca, it is very re-
sistant to chemicals. It possesses the further advantage of
withstanding much greater mechanical shocks.
Boron B
The element boron resembles silicon and graphite in appearance.
It has no appUcations.
Boric Acid H3BO3. — This acid is contained in the steam which
issues from the ground in certain parts of Tuscany. It is caught
in water, placed in basins built over the " soffioni," and separated
by evaporation. Much of it is also made from borax.
Boric acid crystallizes in white, sUppery scales. It dissolves
somewhat in water (4 : 100 at 18°), and the saturated solution,
mixed with an equal volume of water, is used as an eye-wash.
Boric acid is an exceedingly weak acid; its solution scarcely
affects htmus. It is a mild antiseptic, and preserves foods by
preventing the development of moulds and bacteria. It is often
added to talcum powder to prevent infection of irritated skin.
When heated, it loses water and gives tetraboric acid:
Skeleton: H3BO3 -^ H2B4O7 + H2O
Balanced: ' 4H3BO3 -^ H2B4O7 + 5H2O
and eventually boric anhydride B2O3.
Borax. — This salt is the decahydrate of sodium tetraborate
Na2B4O7,10H2O.
It is made by adding calcium borate, foimd in California,
to sodium carbonate solution. The precipitate of calciimi car-
364 smith's intermediate chemistry
bonate is separated by filtration, and the solution is concentrated
until crystals appear upon cooling.
It is a white crystalline salt. It is added to the glass, used for
enameUing and glazing, to make it more fusible and easier to
spread in a thin layer. It is a preservative. Since it contains
but a small proportion of the metallic oxide (Na20,2B203), it
combines with other metallic oxides when fused with them. For
this reason the powdered salt is sometimes sprinkled on tarnished
metallic surfaces which are to be soldered or brazed. The heat
of the bolt or blowpipe melts the borax, and the latter removes
the oxides and permits perfect running of the solder over the
surface. The borates thus formed are often colored, and the
colors afford a means of recognizing the metaUic compoimd which
produced them. In chemical analyses a bead of borax, produced
by fusion on a platiniun wire, is heated with a particle of the im-
known compoimd and its color then examined. The colors are
similar to those already described imder colored glass (p. 362).
When hydrochloric or nitric acid is added to a hot, concentrated
solution of borax, boric acid crystaUizes out:
Na2B407 + 2HC1 + 5H2O -^ 2NaCl + 4H3BO3 i .
Boric acid may be recognized by the green tint which it confers
on the Bunsen or alcohol flame.
Exercises. — 1. Make an equation for the preparation of, (a)
lead silicate PbSiOa by fusion of Utharge PbO and sand, (b)
potassimn silicate K2Si03.
2. The inside surface of the bottle of sodium hydroxide solu-
tion becomes etched and dull. To what is this diie?
3. What is the valence of boron?
4. Why is not all of the boric acid deposited from a hot solu-
tion containing it?
5. Write an equation for the effect of heat upon borax.
6. Why does the addition of borax render a glass more easily
fusible?
CHAPTER XXXI
COMPOUNDS OF SODroM AND POTASSIUM
We have already considered sodium and sodium hydroxide
(Chap. XIV). In this chapter we take up the other important
compoimds of sodium and their uses, and we devote some space
also to potassium and its more useful compounds.
In general, we shall find that these metals and their correspond-
ing compounds are very much aUke in properties. The chief
differences are that the sodium compounds are usually cheaper,
and that, on account of the difference in the atomic weights of
the two elements (sodium 23, potassiimi 39), smaller weights of the
sodium compounds suflSce for a given use involving chemical in-
teraction. For these reasons the sodium compounds find, in
most cases, more applications.
Sodium and potassium are unwdlent^ and both are very active
as metaUic elements. Their hydroxides being strongly alkaUne,
the elements are often called the metals of the alkalies.
Sodium Na
Sodium derives its symbol Na from its German name, natrium.
All compounds of sodium, when heated with a Bunsen burner,
confer a strong yellow tint upon the flame.
Sodium Chloride NaCl. — Sea water contains about 2.5 per
cent of sodium chloride NaCl. The same compound is found in
extensive deposits at Stassfurt in Germany, in Cheshire (Eng-
land), at Syracuse (New York), at Sahna (Kansas), in Utah,
California and many other parts of the United States.
The pure salt is obtained from these deposits by re-crystal-
365
366 smith's intermediate chemistry
lization from water. A certain amount is made, by evaporation,
from sea water.
The salt crystallizes in white cubes. Its chief application is in
the manufactm^ of other compoimds of sodiimi.
Sodium Bicarbonate NaHCOs* Manufacture. — This salt
is manufactiu'ed by the interaction of sodium chloride and am-
monium-hydrogen carbonate in the Solvay or ammonia-soda proc-
ess. Very concentrated solutions are used, and from them a
great part of the sodium-hydrogen carbonate, which is a much
less soluble salt, is precipitated:
NaCl + NH4HCO3 ^ NaHCOa i + NH4CI. (1)
In practice salt is dissolved in water and the solution is saturated
with ammonia gas. The mixture is placed in an iron tower
filled with perforated shelves. Carbon dioxide, made by heating
limestone in special kilns, is forced in at the bottom. The per-
forations spht up the gas into small bubbles, and faciUtate its
solution to form carbonic acid H2CO3. With the ammoniiun
hydroxide NH4OH in the Uquid it gives ammonium-hydrogen
carbonate:
NH4OH + H2CO3 -> NH4HCO3 + H2O. (2)
This product interacts as in equation (1). The sodium bicar-
bonate is precipitated and is freed from the Uquor in filter-presses.
The ammonia is recovered for use by treating the residual liquor,
containing ammonium chloride, with the quicklime CaO from the
kilns. The quicklime, with the water, gives slaked hme Ca(0H)2
and the latter Uberates the ammonia (p. 302) :
2NH4CI + Ca(0H)2 -> CaCl2 + 2NH4OH -> 2NH3 T + 2H2O.
Properties of Sodium Bicarbonate NaHCOs* — This salt is
a fine, whitey not obviously crystaUine powder, which is only
slightly soluble in water. It is commonly known as baking soda.
It decomposes slowly in an open vessel, even when cold. When
COMPOUNDS OF SODIUM AND POTASSIUM 367
heated, it rapidly gives off carbon dioxide and water, and leaves
sodium carbonate:
SkeUton: NaHCOs ^ Na«COa + HaCO, ^ H2O + CQi.
BaUinced: 2NaHC08 ^ Na«COa + H2O + COj.
like all carbonates, when treated with an add, it gives carbonic
acid and this, in turn, gives water and carbon dioxide.
NaHCOs + HCl ^ NaCl + HgCOs ?=^ H2O + CQi T •
This property leads to its use in fire extinguishers.
Sodium Carbonate Na2C03* — This salt is manufactured by
heating sodium bicarbonate. It is made also by the Le Blanc
process. Sodium chloride is treated at a red heat with an equiv-
alent amoimt of sulphuric acid, giving sodiimi sulphate Na2S04.
The latter is roasted with powdered coal and limestone. The
coal reduces the sulphate to sodium sulphide NaaS, and the latter
interacts with the limestone giving sodium carbonate and calcium
sulphide CaS. The carbonate is separated *by solution in water.
The solution, when concentrated, gives crystals of the decahy-
drate, washing soda NaaCOsjlOHaO. The latter, when heated,
leaves " soda-ash " or " calcined soda '' NaaCOs.
NaCl + H2SO4 -^ Na2S04 + 2HC1 T
Na^S04 +2C -►Na^S + 2CO2 1
Na«S + CaCOa -> Na^COa + CaS.
Sodiimi carbonate Na2COs is used in making sodium hydroxide
(p. 166) and glass (p. 361), and to soften water (see p. 389). As
about two-thirds of washing soda is water, it is cheaper to ship
the anhydrous form, except when the hydrate is wanted, as for
washing.
Properties of Sodium Carbonate Solution. — The solution
is not neutral to Utmus, as we might expect, but distinctly alkaline.
The explanation by the ionization theory is of general interest.
368 smith's INTERBfEDIATE CHEMISTBY
Water, it will be remembered, is ionized to a very small extent.
When any saU is dissolved in water, therefore, there exists the
possibility of a double decomposition taking place. Thus with
sodiiun chloride :
NaCl ;=i Na+ + Q-
H2O ^OH- + H+
II 11
NaOH HCl
In this case, however, the extent of formation of NaOH and HCl
is altogether negUgible. NaOH and HCl are both very highly
ionized in aqueous solution, their existence in appreciable amount
woidd involve the presence of both 0H~ and H+ in quantity, and
these ions would inunediately combine to form water.
When a salt Uke sodimn carbonate is dissolved in water, the
state of affairs is different, as may be seen by studying the arrows
in the ionic equations:
fc^ SVM C..^^ -* ' Na«COa ;=i 2Na+ + CO,"
^ ,^ 2H2O ^20H- + 2H+
11 u
2Na0H H2CO,
Double decomposition with the ions of water here involves the
formation of some NaOH and some HaCOs. Now the latter sub-
stance is an exceedingly weak add. While, therefore, any NaOH
formed is almost entirely ionized again, furnishing OH" to the
solution, any HaCOs formed remains, on the contrary, almost
in the non-ionized state. Now we have learnt (p. 241) that one
way of driving a reversible reaction to completion is to remove
one product as a nonr-ionized substance. This reaction cannot
be driven to completion in this way, since we have the formation
of another practically non-ionized substance, water, tending to
drive the reverse reaction to completion. In such circumstances,
obviously, a balance must be struck between the two conflict-
ing tendencies, and we will be left with a solution in which partial
COMPOUNDS OF SODIUM AND POTASSIUM 369
double decomposition has occurred. But such a solution will
contain a much higher concentration of OH" from the highly
ionized NaOH than of H+ from the practically non-ionized H2C0$.
It will, therefore, react Uke an alkaU.
Hydrolysis of Salts. — This interaction of a salt with water is
called hydrolysis (Greek, decomposition by water). All salts are
hydrolyzed, at least to a sUght extent. The action is the reverse
of neutraUzation (p. 192), water and a salt giving, by double
decomposition, an acid and a base. The effect is noticeable^
however, only when the add and base are of very unequal activity.
A salt which, by hydrolysis, gives an active base and a weak add,
furnishes a solution the reaction of which is basic.
Ck)nversely, if the salt gives, by hydrolysis, a weak base and an
active acid, then the solution is acid in reaction. Thus, the solu-
tion of cupric sulphate is acidic because cupric hydroxide is a
feeble base.
The extent of hydrolysis, even in cases where it is distinctly
observable, such as sodium carbonate and cupric sulphate men-
tioned above, is in general only small. Borax, a salt of an ex-
tremely weak acid (p. 363), reacts distinctly alkaUne in water,
but the degree of hydrolysis in a 0.1 iV solution is only one-half
of one per cent. The reason, of coimse, Ues in the fact that water is
very much less ionized even than exceedingly weak acids Uke
boric acid or exceedingly weak bases Uke cupric hydroxide. The
tendency towards the completion of neviralizationy therefore,
preponderates considerably over the tendency towards the com-
pletion of hydrolysis.
Sodium Nitrate NaNOs. — This salt is prepared by recrjrstal-
Uzing Chile saltpeter (p. 308). When heated, it gives off oxygen,
leaving sodium nitrite NaN02, a compoimd much used as a source
of nitrous acid in the manufacture of organic dyestuffs.
Sodium nitrate is at present the chief source of nitric add (p.
370 smith's intermediate chemistry
308). On account of its solubility and its nitrogen content, it
is valuable as a fertilizer (p. 410). It is used also in the manu-
facture of cheap grades of gunpowder. Much of it is converted
into potassium nitrate, as this salt is less deliquescent, and the
gunpowder made from it keeps better.
Other Salt8 of Sodium. — Several of them, such as the
peroxide (p. 221), the silicate (p. 360), and the tetraborate (p.
363), have been described ahready. Sodium sulphate Na2S04
is used as a substitute for sodiiun carbonate in making cheap
glass. Sodium sulphite Na2S08 (p. 260), made by the action
of sulphur dioxide on an aqueous solution of sodium hydroxide,
is a convenient source of sulphur dioxide and is also used as a
preservative. Sodimn thiosulphate Na2S208 (" hypo ") can be
obtained readily by boiling sodiiun sulphite solution with sidphur:
NajSOs + S -^ Na2S20,
and is used in fixing photographs. Sodium cyanide NaNC
(preparation, see p. 393) is used in extracting gold from its ores.
Potassium K
Potassium receives its symbol from the initial of the German
word for it, kalium (related to the word alkali). All compounds
of potassium confer a violet color upon the Bunsen flame.
The metal itself may be prepared from its hydroxide, as in the
case of sodium (p. 164), and exhibits similar properties. It has
no uses. Certain of its salts, however, are of the greatest value,
chiefly as fertilizers (p. 412) and explosives.
Occurrence. — Silicates containing potassiimi, such as felspar
and mica, are common constituents of volcanic rocks. These
minerals have not yet been brought into commercial use as sources
of potassium compounds. Many salt deposits (p. 412) contain
potassium chloride, alone (sylvite) and in combination with
Other salts, and most of the compoimds of potassium are manu-
COMPOUNDS OF SODIUM AND POTASSIUM 371
factured from this- material. Potassiimi sulphate occurs also
in the salt layers, and is used directly as a fertilizer.
•
Potassium Carbonate K2CO89 Preparation. — The water
used in scouring wool leaves, when evaporated, the potassium
salts of organic acids. When the residue is roasted, potassiimi
carbonate remains. Wood ashes contain considerable amounts
of potassium carbonate, and were indeed originally the only
source of this compound. The sugar beet takes up exceptional
quantities of potassiiun and, after all the sugar has been removed
from the extract, potassiimi carbonate is obtained by evaporating
the Uquid and calcining the residue. Much of the salt is also
manufactured direct from potassimn chloride.
Potassiiun carbonate is used in making soft soap and difficultly
fusible glass (p. 362).
Potassium Bicarbonate EHCO3. — When carbon dioxide is
led into potassium carbonate solution potassium bicarbonate
KHCOs is formed:
CQi + H2O ^ HtCOa + K2CO8 T^ 2KHC0a.
The bicarbonate decomposes easily, especially when warmed,
reversing the above action (read the equation backwards). This
occurs even in the solution. Hence carbon dioxide can be forced
by pressure in large amounts into a warm solution of potassiiun
carbonate, and Uberated again by pumping so as to create a
vacuum. This plan is used as a means of purifying carbon
dioxide (p. 334). The same quantity of potassium carbonate
can be used over and over again.
Potassium Hydroxide KOH. — This alkali is made like
sodium hydroxide. Either potassium carbonate is treated with
slaked lime, or potassium chloride solution is electrolyzed in a
Nelson cell (see p. 166). The solution is evaporated, and the
substance cast in slender sticks.
372 smith's intermediate chemistry
Its properties are much like those of sodium hydroxide. It is
used in making soft soap and other compounds of potassium.
*
Potfissium Nitrate ENOs. — The supply of the natural nitrate
being insufficient, the salt is made by double decomposition from
the Chile saltpeter NaNOs:
NaNOa + KCl -► KNOa + NaCl i .
Sodium chloride is not much more soluble in hot water than in
cold. The thriee other salts, however, become very soluble as the
temperature rises. Hence, when sodium nitrate and potassium
chloride are heated with very little water, they dissolve, sodium
chloride is precipitated, and potassium nitrate remains in solution.
The mass is filtered quickly through canvas to separate the pre-
cipitate, and potassium nitrate crystallizes from the filtrate as it
cools.
The salt is used in making gunpowder and fireworks. It is
employed also in preserving ham and corned beef.
Gunpowder » — Gunpowder is composed of potassiimi nitrate
(75 per cent), charcoal (15 per cent) and sulphur (10 per cent).
The ingredients are moistened with water, and intimately mixed
by grinding utnder the heavy rollers of a mill. The " mill cake "
is then broken up and granulated to the required size.
The explosion results largely from the union of the charcoal with
the oxygen from the nitrate and of the sulphur with the potassium.
One gram of powder jdelds 264 c.c. of gas (CO2, CO, and N2)
measured at 0*^ and 760 mm., and a much larger volume at the
temperature of the explosion. One gram produces about 660
calories of heat. The explosion is due to the suddenness with
which the gases are generated and the heat is developed. The
smoke is composed of particles of solid compoimds of potassium
and is therefore very slow in dissipating itself. Smokeless powder
(p. 483) produces no soUds when it explodes.
COMPOUNDS OF SODIUM AND POTASSIUM 373
Gunpowder is still used in mining, and for detonating charges
of smokeless powder.
Other Compounds of Potassium* — Potassiiun bromide KBr
(p. 202), used in medicine to relieve nervousness and to produce
sleep, potassium iodide KI (p. 206), used also in medicine, and
potassium chlorate KClOs, employed in making matches and fire-
works (p. 321), have already been mentioned. Potassium sul-
phate K2SO4, made by the action of sulphuric acid upon potassiimi
chloride at a red heat, is employed in making alimi (p. 469).
Exercises. — 1. In the Solvay process why is the anmionia
dissolved in the salt solution, and not separately in water? Make
the equation for the action of heat on limestone.
2. What is the exact percentage of water in washing soda?
3. What will be the reactions to Utmus of aqueous solutions of:
(a) sodium phosphate, (b) sodium bromide, (c) sodiiun siUcate^
(d) sodium peroxide, (e) sodium nitrate?
4. Write full ionic equations to show why cupric sulphate solu-
tion reacts acid to Utmus.
5. How will a solution of anmioniimi acetate react towards
litmus?
6. Write full ionic equations for the Solvay process (p. 366).
7. Why does a piece of glass, when strongly heated, confer a
yellow color on the Bimsen flame?
CHAPTER XXXII
THE RECOGNITION OF SUBSTANCES, I. — A REVIEW OF THE
NON-METALLIC ELEMENTS
In the chapters preceding we have classified our substances
under one of the chemical elements they contained. Thus,
chloroform, alcohol, and ether were put under carbon. Hydrogen
sulphide and sulphuric acid were considered as compounds of
sulphur. Now this classification is of a theoretical nature. That
chloroform, alcohol, and ether all contain carbon can not be told
by mere inspection. We have to make experiments, and to reason
about the results, before we reach this inference. Thus we put
our inference as the title of the chapter, and distributed the ob-
servations and data through it. There is, however, another way
of classifying the facts, which is just the opposite of this one.
It is the practical classification. When we obtain a specimen, or
when a substance appears in the course of an experiment, we must
be able to tell what particular substance it is. If it is a white
powder, it may contain almost any of the whole Ust of elements.
It may be any one of several thousands of substances. We can
recognize it only by its physical properties (p. 6) and by the
physical properties of other substances that we can get from it
by interaction with known chemical compounds. We need,
therefore, a plan of operation, and this plan must be based upon
a classification by physical properties, not by constituents.
One benefit of the discussion of such a plan is that it will afford
us a review of some of the facts already mentioned, by presenting
the same facts from a different view-point, and by showing the
uses to which they may be put. To avoid unnecessary repeti-
374
THE RECOGNITION OF SUBSTANCES, ETC. 375
tion, we shall refer frequently to the previous statements of the
facts. To understand what is here said, therefore, the reader
should look up every reference and reread the statement.
Is the Material a Mixture? — The first question that occurs
to us is, whether the material is a single substance or a mixture.
When it is a mixture, we can often, though not always, very
quickly find this out.
If it is a solid mass or powder, we examine it with the naked
eye and with the help of a lens. If we see two or more kinds of
particles, as in granite (p. 4), in a mixture of sand and sugar, or
in a piece of rusty iron, the kinds differing in form or color or both,
then it probably is a mixture.
Whether it appears to the eye to be a mixture, or not, we can
next try a solvent, such as water, ether, or carbon tetrachloride.
If a part of a small sample refuses to dissolve {e.g, sand), while
the rest dissolves (e.g. sugar), we have shown that there are two
different sets of physical properties, and therefore two different
substances (p. 13) present. As there may be very Uttle of the
soluble substance in the mixture, we may not perceive at once that
anything has dissolved. So we allow a drop of the liquid to
evaporate on a clean watch crystal, and observe whether any resi-
due remains.
If the material is a liquid) we depend largely on differences in
the volatility of different substances to find out whether it is
a mixture. If a drop evaporates on a watdfc crystal, leaving a
residue (solid or liquid) which does not evaporate, then it is a
mixture. If this test fails, because all, or none, evaporates, then
we must distil the liquid, with the bulb and most of the stem of a
thermometer in the vapor (Pig. 35, p. 67), and note whether the
whole comes over at one temperature (single substance, in most
cases) or whether the temperature changes as the distillation
proceeds (mixture, such as petroleum, p. 344).
These are simply a few examples intended to show how, when
376 smith's intermediate chemistry
fifc practical problem is before us for solution we use phy steal prop'
ertiea as the basis of our reasoning and classification.
The Recognition of a Single Substance. — The majority of
the substances we have met with are acids, bases, or salts. In
identifying a substance of one of these classes, it is convenient to
attempt the recognition of the positive and of the negative radicals
(or ions) as two almost separate problems. In other words, we
investigate one radical at a time. On the other hand, substances
which do not belong to any of these classes, such as simple sub-
stances (sulphur, carbon, chlorine, etc.), oxides (sulphur dioxide,
carbon monoxide, etc.), and many organic substances (e.g. car-
bon disulphide, ethyl alcohol), are investigated as a single prob-
lem.
Scope of this Chapter. — In discussing the recognition of a
single substance we shall, for the present, limit ourselves to the
non-metalhc elements. We shall consider the elementary sub-
stances themselves (sulphur, oxygen, etc.), the oxides of such
elements, the few organic compounds described, the non-metallic
negative ions, and ammonium-ion. We shall leave out of con-
sideration imtil later (Chap. XLVI) the metalhc elements (in-
cluding As, Sb, Bi, Na and K). We shall also ignore the metallic
positive radicals, although one of these (or hydrogen-ion) must
inevitably be combined with the negative radical under consider-
ation.
External Examination (Solids). — The specimen may be a
solid, a liquid, or a gas. We should note in which of these states
it exists under room conditions. What follows appUes only to
SOLIDS — the Uquids and gases will be taken up later (pp. 380,
381).
Without training in crystallography we can tell little about the
crystalline fonn (p. 94) of the specimen. But anything con-
THE EECOGNITION OF SUBSTANCES, ETC. 377
spicuous, such as needle-shape or cubical f onnation of the particles,
must be noted.
The COLOR, if any, is significant. If yellow, the specimen
may be sulphur (p. 249), if black, ^carbon (p. 330), if black and
crystalline, with violet vapor, iodine (p. 203). Most substances
are colorless.
The odor, if any, must be noted. Many salts of ammonium
(carbonate, sulphide, etc.) smell of ammonia (p. 303). Some sul-
phides (of Bodium, potassiiun, ammonium, etc.) smell of hydrogen
sulphide (p. 253) . Some acetates smell of acetic a4nd and hypochlo-
rites of hypoddorous acid (p. 223).* Some chlorides {e,g. PCU
and AlCU), in moist air, give hydrogen chloride^ with the odor and
fumes characteristic of that substance.
Effect of Heating (Solids). — A great deal may be learned by
heating as much of the specimen as will fill the rounded bottom
of a small, dry test-tube (Pig. 92).
The substance may fuse. Many substances, such
as some chlorides and most hydrates, do so. Continue
heating.
A sublimate (soUd deposit on the cold part of the
tube) may appear. Black crystals (from violet vapor)
are iodine: white crystals, with the Umitations we have
set, show the substance to be a salt of ammonium (p. 305).
Confirm by smelling, and by adding sodium hydroxide to the
original substance (p. 302). (Salts of mercury sublime also.)
A reddish-brown liquid condenses, becoming, when cold, a yel-
low sohd. The substance was sulphur from a sulphide (such as
FeS2, p. 486).
Water condenses. Hydrates, add-saUs, and some organic com-
pounds. Test the water with Utmus paper. An add reaction in-
* Many experimental details, essential for the successful performance
by a beginner of the tests described in this chapter are here omitted. They
will be found in the Authors' Intermediate OvUine of Elementary Chemistry.
U
378 smith's intermediate chemistry
dicates an add-saU (p. 192), an easily hydrolyzed salt (e.g. FeCU,
6H2O) or organic add. Continue heating, inclining the tube
mouth downwards, removing condensed water with filter paper
imtil no more comes ofif, and heat the residue.
A gas is given off. The gas may be violet (some iodides) or
brown (some nitrates, p. 311, and some bromides). If brown,
lower a glass rod dipped in silver nitrate solution into the gas
in the tube. Bromine will give a white precipitate (AgBr, p.
202). In case of a negative result try the test for nitrates.
The gas may be colorless. If it has an odor, it may be am-
monia from a salt of ammonium, sulphur dioxide from a bisulphite
(p. 258) or from oxidation of a sulphide (p. 258). A stifling odor
with fumes may be sulphur trioxide from some sulphates or de-
composition products from some organic matters.
The gas is colorless and odorless. It may be oxygen (test with
long, glowing splinter of wood) from a peroxide, chlorate, or nin
trate (of K or Na). To learn which of these it is, dissolve or sus-
pend a Uttle of the substance in water, add dilute sulphuric acid,
and test for hydrogen peroxide (p. 223). In case of a negative
result examine the residue (as in p. 379).
A colorless, odorless gas may be carbon dioxide, coming from a
bicarbonate or a carbonate (except of K or Na). Lower a glass rod
dipped in Ume-water into the gas in the test-tube (white precipi-
tate, CaCOs [).
The substance carbonizes or chars and gives an odor of smolder-
ing wood or burning flesh. The compoimd is organic. Identify,
by properties (pp. 346-50, 353).
Heating may produce no effect. On the other hand, more than
one of these effects {e.g. both water and sulphur dioxide from a
bisulphite) may be given by the same specimens.
If heating produces any effect, continue heating imtil all change
ceases, and preserve the residue for use in p. 379.
In most cases other distmctive properties and tests wiU be
found on the p^es referred to.
THE RECOGNITION OP SUBSTANCES, ETC. 379
Effect of Sulphuric Acid on Solids. — Fill the rounded
bottom (only) of a test-tube with the substance, add just enough
concentrated sulphuric acid to moisten the sample, and warm
sUghtly.
A gas (effervescence) which fumes in the breath may be given
off. If the gas is brown or yellow, it may be bromine (bleaches
litmus paper) mixed with hydrogen bromide from a bromide
(p. 202) . It may be nitrogen peroxide (odor) from a nitrate (p. 3 11) .
If violety with brown deposit, accompanied by an odor of hydro-
gen sulphide, it is iodine from an iodide (p. 204).
If the gas fumes, but is colorless, it may be hydrogen chloride
from a chloride (add manganese dioxide to get chlorine, p. 142).
It may be hydrogen fluoride from a fluoride (a moistened glass rod
acquires white precipitate of silicic acid produced by decompo-
sition of silicon fluoride, p. 207).
The gas does not fume. If yellow it may be chlorine from
bleaching powder (p. 224) or nitrous anhydride from a nitrite
(p. 314). A chlorate also gives a yellow gas (chlorine dioxide
CIO2) when heated with concentrated sulphuric acid. In this
case oxygen will have obtained in the test on page 378. Heat
very carefuQy, since chlorine dioxide is explosive (stop heating
when material in tube begins to crackle!).
The gas does not fume and is colorless. An odor of sulphur
dioxide indicates a sulphite (p. 258) . An odor of hydrogen sulphide
indicates a sidphide. If the gas is odorless, it may be carbon
monoxide (bums, leaving carbon dioxide) from a formate (p. 337),
or oxygen from some oxides or a peroxide, or nitrous oxide from
ammonium nitrate (p. 315), or carbon dioxide from a carbon-
ate.
No gas evolved indicates a silicate (p. 360), sulphate (p. 270),
phosphate (p. 322), or a basic oxide.
Sulphuric Acid on the Residue from p. 378. — If the sub-
stance gave off oxygen when heated alone (p. 378), add a drop
380 smith's intermediate chemistry
or two of concentrated sulphuric acid to the residue. If the
specimen now gives a yellow gas (nitrous anhydride), the original
substance was a nitrate (of K or Na), from which the nitrite was
formed by heating (p. 311). If it gives a colorlesSi fuming gas
(HCl), the original substance was a chlorate (p. 30).
Examination of a Liquid. — Test the specimen with litmus
paper. A marked acid reaction may be due to an dddy such as
concentrated or dilute sulphuric acid (p. 270), concentrated or
dilute nitric acid (p. 309), concentrated or dilute hydrochloric
acid (p. 131), hydrobromic acid (p. 202), hydriodic acid (p. 205),
phosphoric acid (p. 322), sulphurous acid (p. 259), or an organic
acid (p. 348); also an acid-salt (p. 192), or a hydrolyzed salt (p.
369).
If it bleaches litmus paper, it may be chlorine-^ater or bromine^
water (odor).
If it is markedly alkaline in reaction, it may be a solution
of a base (NaOH, KOH, NH4OH, etc.) or a hydrolyzed salt (p.
355).
Note the odor. Ammonium hydroxide, hydrogen sulphide solu-
tion, sulphurous acid, concentrated nitric acid and concentrated
haUde acids all have odors. Alcohol, acetic acid, carbon disul-
phide, carbon tetrachloride, and hydrocarbons (e.g. gasoline)
have odors easily distinguished from those of the foregoing.
Evaporate a few drops to dryness on a watch crystal. A soUd
residue shows that the original substance was a solviion in
water (or possibly alcohol or some other solvent, if the vapor has an
odor indicating this). If there is a soUd residue, a quantity of it
may be obtained by evaporating a larger amoimt of the Uquid,
and may then be treated as a soUd (pp. 377-9).
If the specimen leaves no residue, and is not acid or alkaline
but has an odor, it may be one of the volatile organic compoimds
named above. If it is odorless, it may be a solution of hydrogen
peroxide (p. 222) or simply pure water.
THE RECOGNITION OF SUBSTANCES, ETC. 381
Examiwiation of a Gets. — The gas has a color. A brown
gas may be bromine or nitrogen peroxide. The former liberates
iodine from potassium iodide solution (p. 205), but not bromine
from a bromide (insert rods moistened with solution of an iodide
and a bromide). The latter becomes deeper brown on warming
(p. 311). A greenish-yellow gas is chlorine. It bleaches, and
displaces bromine from a solution of a bromide (p. 202).
The gas may become colored (yellow or brown) on admitting
air. It is nitric oxide (p. 311).
The gas may have a distinctive odor. Sulphur dioxide, hy-
drogen sidphide, nitrogen peroxide, and ammonia are of this
kind.
The gas may fume in the breath. The chloride, bromide, and
iodide of hydrogen do so. Distinguish by dissolving in Kttle
waxer and adding manganese dioxide.
The gas may be combustible. Burning with a blue flame in-
dicates hydrogen (vessel bedewed with moisture), or carbon m,on-
oxide (leaving carbon dioxide, test, p. 336). Bummg with a
sUghtly Itmiinous flame indicates methane (p. 345) and a very
luminous flame (often depositing carbon) indicates ethylene (p.
350) or acetylene (p. 351).
The gas may relight a glowing splinter of wood. This is ooi^-
gen, or nitrous oxide (p. 315). The former, with nitric oxide,
gives a brown gas (p. 311), the latter does not.
The gas may give a white precipitate (CaCOj) with lime-
water. This is carbon dioa^ide.
The gas having none of these properties is nitrogen (p. 287).
In most cases other distmctive properties will be found on the
pages referred to.
Exercises. — 1. Look up the references, and give the proper-
ties (physical as well as chemical), other than those mentioned in
p. 377, by which you should recognize: (a) sulphur, (b) carbon,
and (c) iodine.
382 smith's intebmediate chemistry
2. Same question in regard to: (a) ammonia, (b) hydrogen
sulphide, and (c) hypochlorous acid.
3. Same question in regard to: (a) an acid salt, (b) a hydrate,
(c) nitrogen peroxide, (d) bromine, (e) sulphur trioxide, (f) hy-
drogen peroxide, (g) anmionium nitrate, (h) carbon dioxide.
4. Same question in regard to (p. 379) : (a) a bromide, (b) an
iodide, (c) a nitrate, (d) bleaching powder, (e) sulphur dioxide,
(f) carbon monoxide, (g) a silicate, (h) a sulphate, and (i) a phos-
phate.
5. Same question in regard to (p. 380) : (a) cone, and (b) dil.
sulphuric acid, (c) cone, and (d) dil. nitric acid, (e) cone, and (f)
dil. hydrochloric acid, (g) hydrobromic acid, (h) hydriodic acid,
(i) phosphoric acid, (j) sulphurous acid.
6. Same question in regard to: (a) chlorine-water, (b) alcohol,
(c) acetic acid, (d) carbon disulphide, (e) carbon tetrachloride,
(f) a hydrocarbon.
7. Same question in regard to: (a) nitric oxide (b) hydrogen,
(c) methane, (d) ethylene, (e) acetylene, (f) nitrous oxide, (g)
ozone, (h) nitrogen.
8. Why does sodium sulphide smell of HjS (see pp. 118, 369)?
9. When anmioniimi nitrate is heated in a test-tube, the gases
evolved do not relight a glowing splinter. How can this be
reconciled with p. 3157
CHAPTER XXXIII
CALCIUM AND ITS COMPOUNDS
Calcium belongs to a family of metallic elements which includes
also strontium, barimn, and radium. This family resembles that
to which sodium and potassium belong in so far that the metals
are second in activity only to the two last named and that the
hydroxides are active bases. The chief differences are that the
metals of the present group are bivalent and that all the carbonates
and many other single compounds are insoluble.
Compoimds of calcium confer a brick-red color upon the Bimsen
flame.
Calcium. — The metal is made by electrolyzing melted cal-
cium chloride in a graphite crucible, which forms the anode.
The cathode is a rod of iron, one end of which dips into the fused
salt. The calcium, Uberated at this point, adheres to the rod.
The latter is slowly raised, in such a way that the calcium always
remains in contact with the Uquid. In this fashion a long " cab-
bage-stalk " of calcium is finally produced.
The metal is slightly harder than lead and has a sHver-white
luster. It decomposes cold water, Uberating hydrogen (p. 50).
Calcium Carbonate CaCOs* — The carbonate is the com-
monest compound of calciiun. White marble is a pure variety,
composed of crystals compactly wedged together. Limestone
does not show much crystalline structure and usually contains
clay and other impurities. Chalk is made up of shells of minute
marine organisms. Shells, coral, and pearls are likewise mainly
calcium carbonate. Well-formed crystals (calcite, or Iceland
spar — Fig. 93 — and aragonite — Fig. 94) are common.
383
384 sutth's intermediate chemistbt
Limestone is used for building and in road-making. Much of it
is employed in making quicklime, cement (p. 471), and glass
(p. 361), and as a flux in metallui^cal operations (see p. 486).
Marble, often variegated by the presence of impurities, is used
in building and sculpture.
Fia. 93 FiO. 94
Aa we have seen (p. 333), calcium carbonate reacts wUh adds to
give carbonic acid :
CaCO. + 2HC1 -» CaClj + H,CO, ^ H^ + CO, t •
When heated, all forms of calcium carbonate give off carbon
dioxide, and leave calcium oxide:
CaCO,?iCaO + CO,.
Calcium Oxide CaO, Manufacture. — The manufacture of
calcium oxide or quicklime {i.e.
live lime) ia one of the moat ancient
chemical industries. The lime-
stone is thrown into a kiln lined
with brickwork (Fig. 95). The
flames and heated gases from the
fire pass through the limestone and
the carbon dioxide is liberated and
carried off by the draft. When
this gas is to be used, as in the
Solvay process or in the refining
of sugar, coke (smokeless) is chosen
as the fuel. When no use is to be
made of the escaping gas, coal may be employed.
The use of as low a temperature as possible is important. A
CALCIUM AND ITS COMPOUNDS 385
high temperature causes the unpurities in the Umestone (the clay,
etc.) to interact with the quickhme and form fusible silicates,
which fill the pores and retard the subsequent action of water in
slaking the lime. Calcium carbonate gives a pressure of only
25 mm. of carbon dioxide at 700° and one atmosphere at about
900°. The action is reversible (see equation), and if the gas ac-
ciunulates in the kiln, the carbon dioxide recombines with the
quickUme as fast as it is Uberated — unless a temperature above
900° is used. When, however, the gas is continually removed,
the backward action is prevented, and a lower temperatiu^ suf-
fices to complete the dissociation of the compoimd. Hence a
low temperature and a good draft of air through the kiln are
essential:
Properties and Uses of Quicklime. — Calcium oxide is a
white, amorphoiLS material. It melts only in the electric arc.
When heated strongly, it glows with an imusually briUiant and
white Ught. The Drummondi or oxy-hydrogen light, more com-
monly called the lime light or calcium light, is produced by al-
lowing a flame of burning oxygen and hydrogen or illuminating
gas to play upon a cylinder (rf calciiun
oxide (Fig. 96). The gases are contained
in iron cylinders, under pressure, and so
the apparatus can be used for iQimiination ^^ j 'QP-
or projection where neither electricity nor y^^ qq
a local supply of gas is available.
When vHiter is poured upon quicklime, it is at first absorbed,
and then enters iato combraation to form calciiun hydroxide
(slaked lime) :
CaO + H2O -> Ca(0H)2.
So much heat is given out that the excess of water is converted
into steam. The quicklime swells and falls to powder.
Quicklime is used most largely in making slaked lime for mor-
386 smith's intermediate chemistry
tar, and also in the manufacture of bleaching powder. Other uses
are mentioned under the hydroxide.
Quicklime deteriorates when exposed to the air. It combines
both with the moisture and carbon dioxide in the atmosphere and
becomes air-slaked.
Calcium Hydroxide Ca(0H)2. — The hydroxide is a white,
amorpfioiLS powder. It is sUghtly soluble in water (about 1 : 600
at 18°), giving lime-water. The solution has a strong alkaline
reaction, however, showing that so much as is dissolved is very
largely ionized. Milk of lime, a satiu*ated solution with a large
excess of calcium hydroxide suspended in it, is employed in many
operations (see e.g. pp. 166, 371). As the dissolved part under-
goes chemical change, more goes into solution. Being cheap,
it is used whenever an alkaU is needed, provided a dilute alkali
will serve the purpose. It interacts with acids giving salts of
calcium, and shows the other properties of a base (pp. 167, 192).
Slaked lime is used in making mortar (see below) and alkaUes
(p. 166) and in piuifying sugar (p. 402). It is employed to re-
move the hair from hides, before tanning, an action which recalls
the solubiUty of wool (sheep's hair) in an alkaU (p. 1). It finds
appUcation, also, in softening water (p. 389) and as whitewash.
Mortar. — Mortar is made by mixing slaked lime with three or
four times its bulk of sand, and making the whole into a paste
with water. When the water evaporates, a porous, rather crumbly
material remains. This, however, at once begins to harden,
owing to the action of the carbon dioxide in the air upon the hme:
Ca(0H)2 + CO2 -> CaCOs + H2O t .
The crystalUne calcite (CaCOa) adheres to, and is interlaced with
the sand, and gives a rigid, though porous, structure attached
firmly to the brick or stone. The pores facihtate the penetration
of the air into the deeper parts and thus provide, both for the
fresh suppUes of carbon dioxide required for the continuance of
CALCIUM AND ITS COMPOUNDS 387
the action shown in the equation, and for a considerable amount of
useful ventilation through walls of the building.
Calcium Sulphate CaS049 Various Forms. — Calcium sul-
phate is a very common mineral. It occurs, as anhydrite CaS04,
in salt deposits. Gypsum CaS04,2H20 is foimd in masses, and
also in single crystals (selenite. Fig. 42, p. 94). Alabaster is
highly crystalline gypsum, tmted by impurities.
Gypsum CaS04,2H20 is the commonest form, and is the one
produced when calcium sulphate is precipitated. It is white
and much softer than calcite. It is only sUghtly soluble in water
(1 : 500 at 18°). It is used as a fertiUzer and in making plaster
of Paris and is the chief component of blackboard crayon or
" chalk."
When gypsum is heated, the vapor pressure of the water it
gives ofif soon exceeds that of the moisture in the atmosphere,
and the compound begins to decompose:
2[CaS04,2H20] ^ (CaS04)2,H20 + SHjO t .
The hemi-hydrate which remains (plaster of Paris) gives a much
lower pressure of water vapor and is more stable. Plaster of
Paris is manufactured in large quantities by heating gypsum in
kilns. When moistened with water, it sets in about half an hour
to a soUd mass of gypsmn. The temperature used in making it
must not exceed 125°, otherwise the hemi-hydrate is itself de-
composed, the plaster is ** dead bmut," and it no longer sets
readily. The setting involves, simply, the reversal of the equation
given above.
Plaster of Paris swells somewhat, in setting, and so fills out
completely every detail of a mould and appUes itself closely to
the outUne of an object on which it is spread. It is used in mak-
ing casts, and in surgical bandages where movable parts are to be
held rigidly in place. Stucco is made with sizing or glue instead
of pure water.
388 smith's intebmediate chemistry
Casts are made smooth and non-porous (" ivory " smface)
by a coating of paraffin which fills the pores. Excellent imita-
tions of bronze or other castings are produced by rubbing with
pulverized metab.
WTiat Makes Water Hard. — All natural waters except rain
water, which is " soft," contain salts of calciiun and magnesium
in solution and are more or less " hard." These salts are dissolved
by the water in its passage over and through the soil.
Although limestone is very insoluble in pure water (0.013 g.
per Uter), yet it interacts with the carbonic acid contained in all
natural waters, giving calcium bicarbonate which is about thirty
times more soluble under atmospheric conditions:
C0» + H2O ^ H2CO3 + CaCO, ^ Ca(HCO,),.
When the water is boiled, these actions are all reversed. The
carbon dioxide is driven out of solution, the carbonic acid is
decomposed, and the calcium bicarbonate gives calcium carbon-
ate, most of which is at once precipitated. Iron carbonate is
also held m solution as bicarbonate Fe(HC08)2 and is precipi-
tated as FeCOs by boiling. These two bicarbonates constitute
temporary hardness. Their decomposition causes the " fur " in
a kettle.
The sulphates of calcium (solubility 2 g. per Uter) and of mag-
nesiiun (very soluble) are also commonly foimd in natural waters.
These salts are not altered by boiling and, along with magnesium
carbonate (soPty 1 g. per 1.) and calcium carbonate (soPty 0.013 g.
per 1.), give permanent hardness to the water.
Consequences of Hardness in Water. — When hard water
is used in a steam boiler, the salts, of course, are not carried off
with the steam, but accumulate amazingly as fresh water is in-
jected and steam alone is drawn off. In time, heavy deposits
of bailer crust settle on the tubes of the boiler, and interfere with
the transference of heat from the metal to the water. One-fourth
CALCIUM AND ITS COMPOUNDS 389
of an inch of crust will increase the bill for fuel by 50 per cent.
In addition to this the iron is heated to a higher temperature
and may even become red hot. In consequence, it combines
more rapidly with oxygen on the outside and displaces hydrogen
from the water (p. 51) on the inside, giving in both cases Fe804.
Thus the Ufe of the boiler is shortened. If the formation of the
crust is not prevented, or if the crust is not removed, the boiler
may explode and great damage may be done.
When hard water is used for washing , in the hoiLsehold or laund-
ry, much soap has to be dissolved before the necessary lather
can be secured. Soap, which consists of a mixture of the sodium
salts of several organic acids, such as palmitic acid H.CO2C15H31
(see p. 438), interacts by double decomposition with the salts of
calciimi and magnesium giving palmitates etc. of these metals.
These salts are insoluble and form a " curd.'* With sodium
palmitate Na(C02Ci5H8i), for example, the action is
CaS04 + 2Na(C02Ci5H8i) ^ Ca(C02Ci6H3i)2 i + NbSOa.
Not imtil all the salts causing the hardness have been decom-
posed, does the permanent solution of soap which is required for
washing begin to be formed. The waste thus involved is often
very great and expensive.
Treatment of Hard Water. — Temporary hardness is com-
monly removed, on a large scale, by adding slaked lime (made into
milk of lime) in exacdy the qiumtity shown by an analysis of the
water to be required, and stirring for a considerable time:
Ca(HC03)2 + Ca(0H)2 -> 2CaC08 i + 2H2O. (1)
The bicarbonate is neutraUzed and all the lime precipitated. The
latter is reioaoved by filtration.
Permanent hardness is not aflfected by slaked lime, but is re-
moved by adding sodimn carbonate in the necessary proportion:
CaS04 + NaaCOs -> CaCOa i + NaaSO*. (2)
390 smith's intermediate chemistry
When both kinds of hardness are present, crude caustic soda
(sodium hydroxide) may be employed. It neutralizes the bi-
carbonate, precipitating CaCOa:
Ca(HC03)2 + 2NaOH -> CaCOa i + NagCOa + 2H2O (3)
and giving sodium carbonate. The latter then acts as in equation
(2).
Instead of this, the treatments indicated in equations (1) and
(2) may be appUed in combination (Porter-Clark process).
In the new pei^nutite process the water is simply filtered through
an artificial sodium sihco-aluminate (permutite) which is sup-
pUed in the form of a coarse sand. The calcium, etc., in the
water is exchanged for sodium, which does no harm. If we use
for permutite the abbreviated formula NaP, we may write the re-
action thus:
Ca(HC03)2 + 2NaP -^ 2NaHC03 + CaPj.
After twelve hours' use, the permutite is covered with 10 per cent
salt solution and allowed to remain for the other twelve hours of
the day, when it is ready for employment once more:
2NaCl + CaPa -> CaCk + 2NaP.
Only salt, which is inexpensive, is consumed, and calcium chloride
solution is thrown away. Permutite removes magnesium, iron,
manganese, and other elements in the same way. The life of a
charge is said to be over twenty years.
•
Hard Water in the Laundry. — As we have Been (p. 389),
soap will soften water, but the calcium and magnesium salts of
the soap acids, which are precipitated, are sticky, and soil the
goods being washed. Other substances that soften water not
only give non-adhesive precipitates, but are also much cheaper,
and an attempt is generally made to utiUze them. The use of
slaked lime is impracticable on a small scale.
CALCIUM AND ITS COMPOUNDS 391
Washing soda Na2CO8yl0H2O is added to precipitate both kinds
of hardness:
Ca(HC0,)2 + NasCOa -► CaCO, i + 2NaHC0,
CaSOi + NajCOs -^ CaCOs i + NajSOi.
The small amounts of salts of sodimn which remain in the water
have no action on soap.
Household Ammonia NH4OH acts like sodium hydroxide
(p. 390) :
Ca(HC03)2 + 2NH4OH -► CaCOa i + (NH^iCOa + 2H,0
CaS04 + (NH4)2C03-> CaCOa i + (NH4)iS04
except that it will not precipitate magnesium-ion (see p. 539).
When borax Na2B4O7,10H2O (p. 363) is added, it is hydrolyzed
and the sodium hydroxide contained in its solution acts as akeady
described.
The supposed bleaching or whitening action of borax or soda
is a myth; these salts prevent staining by (he iron in the water.
They simply precipitate the iron (present as Fe(^COz)i), which
almost all waters contain, as FeCOs before the goods are put in.
This precipitate is easily washed out in rinsing. The palmitate,
etc., of iron, however, which the soap itself would throw down, is
sticky and adheres to the cloth. The air subsequently oxidizes
it and gives hydrated ferric oxide (rust), which is brownish-red.
It is evident that, properly to achieve their purpose, the soda
and borax must be added, must be completely dissolved, and
must be allowed to produce the precipitation of FeCOs, CaCOs,
etc., all before the soap, or the goods, is introduced. If the soap
is dissolved before or with the soda, it will take part in the pre-
cipitation, and give sticky particles containing the iron and cal-
cium salts of the soap acids.
Washing ^wders are, or ought to be, mainly sodium carbonate,
mixed with more or less pulverized soap.
392 SMTTH S INTEBMEDIATB CHEMISTRT
Calcium Chloride, CaCla. — Chloride of calcium is obtained
as a by-product in the Solvay process (p. 366) and in other indus-
tries. It crystallizes from water as the white hexahydrate,
CaCU,6HiO, and Is very soluble. The porous, granular variety,
used for drying gases, is made by driving moat of the wator out of
the hexahydrate by heat. The granular form is used in lai^e
amounts for sprinkling on dusty roads. The salt, being deliques-
cent (p. US), attracts water from the air and moistens the dust
with calcium chloride solution. The saturated solution does not
freeze until —48° is reached, so that chilled calcium chloride
brina is used in refrigerating appliances (p. 306).
Calcium Cyanamide CaCNj. — Calcium carbide, when
strongly heated, absorbs nitrogen, giving a mixture of calcium
cyanamide and carbon (aitro-Iime);
CaCa-|-N,^CaCN»"+C.
The carbide is pulverized and placed in a cylin-
drical furnace (Fig. 97), holding 300 to 450 kg. The
heat (800 to 1000°) is furnished by the passage of
a current of electricity through a thin carbon rod,
which passes through the axis. The tube sur-
rounding the rod and the other partitions are of
cardboard, which bums up and leaves opening
for the circulation of the nitrogen. The latter is
made by the fractionation of liquid au- and is in-
troduced under pressure. In thirty-five hours
nitrogen ceases to be absorbed, and the product is pulverized
when cold.
Calciiun cyanamide is now manufactured in large quantities
at Niagara Falls (Ontario) and Odda (Norway) for use as a
fertilizer. It is also a valuable source of ammonia, and was utilized
very extenuvely as such for the manufacture of explosives diuing
CALCIUM AND ITS COMPOUNDS 393
the Great War. When treated with hot water, calcium cyana-
mide is decomposed as follows:
CaCN2 + 3H2O -^ CaCOa + 2NHa.
In practice, pulverized nitro-lime is fed into an autoclave charged
with water. Impurities such as free calcium carbide and cal-
cium phosphide are immediately decomposed, and the gases
evolved allowed to escape. Small amoimts of alkali are then
added to facilitate the evolution of ammonia and to prevent the
formation of complex nitrogen compounds. The autoclave is
closed and steam admitted until the pressure rises to 3-4 atmo-
spheres, the heat evolved by the decomposition of the cyanamide
being sufficient to carry the reaction to rapid completion. The
ammonia given off is absorbed in water, or converted directly
to ammonium sulphate. When the gas ceases to be given off,
steam is blown through the liquor in the autoclave until all residual
dissolved ammonia is expelled.
The productive capacity of cyanamide plants in 1920 exceeded
1,750,000 tons, calculated as nitro-lime.
Nitro-lime, when fused with sodium carbonate, gives sodium
cyanide NaNC, used in the extraction of gold:
CaCNi + C + Na^COa -> CaCOa + 2NaNC.
Other Compounds of Calcium. — Calcium fluoride CaF2
occurs as a mineral (fluorite). It is our source of hydrofluoric
acid (p. 207), and is used in metallurgy to lower the melting-point
of slags. The phosphates (p. 411) and bisulphite (p. 398) of
calcium, and bleaching powder CaCl(OCl) (p. 224) are elsewhere
discussed.
Strontium^ Sr ^nd Barium^ Ba. — The compounds of these
elements closely resemble those of calcium in physical properties
and chemical behavior. Strontium salts confer a carmine-red
color to the Bimsen flame, barium salts a green color. Both are
394 smith's intermediate chemistry
used as ingredients in fireworks and signal lights. Barium
sulphate is used in making white paint (" permanent white ")•
Exercises. — 1. Why can not lime-water be kept in an open
bottle?
2. Why does whitewash become so firmly attached to the wall?
3. What is the percentage of nitrogen in nitro-lime, assuming
100 per cent efficiency of conversion?
4. Make equations for the action of sodium palmitate: (a)
upon calcium bicarbonate; (b) upon magnesium sulphate.
5. In softening water: (a) what would be the objection to
using an excess of milk of lime; (b) why is prolonged stirring
required (p. 389); (c) why must the precipitate be removed by
filtration?
6. Explain why wood ashes are sometimes used to soften water,
and how they act.
7. Make equations for the action of chlorine upon quicklime.
8. Why does fluorite lower the melting-point of a slag?
CHAPTER XXXIV
FLAirr LIFB. OSLLULOSB, STARCH AND SUGAR
Plants and animak are similar in composition. They contain
much the same efomente, and these are present in the form of
similar compounds. They differ sharply, however, in the foods
they use in constructing these compounds. Plants use simple,
inorganic m^aterials; animals absolutely require complex, organic
substances as food. The main domical processes, therefore, are
very different in the two groups.
Hotr the Plant Feeds. — The walls of the cells which form the
frame-work of a plant are made of cellulose (C«Hio06)y. In the
cells, especialliy in certain parts of the plant, granules of starch
(CeHioOs)* are foimd. These complex substances differ in prop-
erties, although they have the same composition. The plant
juice (sap) contains sugars, such as cane-sugar or sucrose, C12H22O11,
in variable amounts, and also esters (vegetable oils, p. 432) and
alkaloids (vegetable bases, p. 479) in much smaller quantities.
The plant cells also contain still more complex substances, known
as proteins. Gluten, the sticky portion of wheat flour (p. 5), is
a typical protein. These proteins are the chief components of
the protoplasm, a semi-liquid substance lining each active plant-
cell, and the real seat of life of the plant. Now all these substances
contain carbon, hydrogen, and oxygen, and plant food must
furnish these elements, which constitute over 95 per cent, on the
average, of all plants. Hence, in addition to large quantities of
water ascending from the soil through the roots and stem, and
sufficient amounts of compounds of nitrogen, potassium, phos-
phorus, and other elements, all plants require an abundant supply
of carbon in absorbable form. This carbon is practically all
395 ♦
1
396
sbhth's intermediate chemistry
taken up by plants in the form of atmospheric carbon dioxide.
It is admitted through minute openings (stomata), situated
mainly in the surface of the leaves.
The Refwtion Involved. — Comparison of the formulae of
carbon dioxide CO2 and of any plant substance, like starch
(C«Hio06)x, shows at once that the latter contains a far smaller
proportion of oxygen, relatively to the amoimt of carbon, than
does the former. Hence, during the
digestion or assimilation of the carbon
dioxide by the plant, this compound
must be reduced. In point of fact, the
chlorophyll (green coloring matter) and
protoplasm in the leaves act upon the
carbon dioxide, causing oxygen to be
Uberated:
r\
1
Fia. 98
6CO2 + 5H2O -^ CftHioOfi + 6O2 1-
This action goes on only in the sim-
Ught. The steps by which sugar,
starch, and cellulose are manufactured
by the plant out of water and carbon
dioxide, are not yet perfectly xmder-
stood. But the Uberation of the oxygen is easily shown by placing
a green plant imder water in a jar, and setting the jar in the sim-
light (Pig. 98). Bubbles of gas appear on the leaves, grow larger,
and then detach themselves and rise to the top. The gas re-
lights a glowing sphnter of wood, and is pure oxygen.
The results of recent investigations suggest the following stages
of the reaction:
(1) Carbonic acid, formed by the union of water and carbon
dioxide, is reduced to formaldehyde (p. 348), oxygen being liber-
ated:
aCOa -> H.CHO + O2.
PLANT LIFE. CELLULOSE, STABCH AND SUGAB 397
(2) Fonnaldehyde molecules quickly combine together, or
polymerize, to give simple sugars with the formula C6H12O6:
6H.CH0 -> CeHiaOe.
(3) These sugars lose a molecule of water, and polymerize
further to fonn starch and cellulose:
nCeHiaOe -^ nHgO + (C6Hio06)n.
Reactions (1) and (2) have been carried out in the laboratory
with ultra-violet light as a catalyst. Chlorophyll thus appears
to act in the r61e of a promoter (compare p. 339), in the presence
of which the reactions are able to proceed in visible light.
Reverse reactions also take place in the plant. Thus starch,
which first accumulates in the leaves, is later turned back into a
sugar soluble in the sap, and is thus able to pass to parts of the
plant requiring new material. Some carbon dioxide is also
liberated from plant surfaces by oxidation of sugars.
The Thermochemistry of the Reaction. — In the combina-
tion of carbon and oxygen, during combustion of wood or coal,
much heat is liberated. Hence, when oxygen is taken out of
carbon dioxide agam, heat or energy in some form must be sup-
plied. When this takes place in a plant, the energy is evidently
furnished by the sunUght, for the action proceeds more slowly m
the shade, and ceases in the dark.
The energy required can be measured, and may be expressed in
calories. The energy required to produce one simple formula-
weight of cellulose (CeHioOg = 6 X 12 + 10 X 1 + 5 X 16 = 162
g.) is 671,000 calories. The whole may be represented in a rough
equation, in which the unknown intermediate steps are left out,
and only the starting substances and the final products are shown:
6CO2 + 5H2O + 671,000 cal. -> CeHioOs + 6O2.
Cellulose (CeHioOs)^. — This substance, named cellulose be-
cause it forms the walls of the cells, composes much of the frame-
308 smith's intermediate chemistry
work and intricate structure of plants. We are familiar with
pure cellulose in the forms of filter paper and cotton. The latter
consists of fine, hollow tubes of cellulose (see Fig. 2, p. 2), large
tufts of which surroimd the seed of the cotton plant. Linen is
almost pure cellulose, wood is largely cellulose, and paper pulp
is practically all cellulose.
Cellulose interdcts with very few chemical substances. It is
because it thus remains unchanged, by most substances that come
in contact with it, that it can be used as a filter paper. When it
does undergo chemical change, it acts as if it contained hydroxyl
(OH) groups, and behaves therefore chemically like an alcohol
(see p. 480) forming valuable derivatives (explosives and plastics)
which will be discussed later (Chap. XL).
Paper Manufacture. — Paper is composed of cellulose
(C6Hio05)y and is made from a mixture of cotton or linen pulp and
wood pulp — the cheapest varieties from the latter alone. The
wood is cut into chips and heated (" cooked ") with a solution
of calcium bisulphite Ca(HSOs)s. This dissolves out the lignin,
which, together with cellulose, makes up the solid part of its
structure. The pulpy material is then washed, beaten with water
to reduce it to minute shreds, and bleached with very dilute
chlorine-water. The pure cellulose, now paper pulp, suspended in
water, is spread on screens, drained, pressed, and dried. During
the process other substances are usually added. Thus size (glue
or gelatine) prevents the ink from running; pulverized calcium
sulphate (gypsum), and other white solids (" loading ") give body
to the paper and make possible the subsequent production of a
smooth smface by rolling (" calendering ")• Ultramarine (blue)
and other colored powders are added to the pulp when special
tints are required.
Other Uses of Cellulose. — Cellulose dissolves in hot, concen-
trated zinc chloride solution. When the liquid is pressed throug}i
PLANT LIFE. CELLULOSE, STARCH AND SUGAR 399
a small orifice into alcohol, the cellulose is reprecipitated in the
form of a thread. By carbonizing, this is made into filaments
for incandescent electric lamps.
Cellulose is solubl?. also in a solution of cupric hydroxide in
excess of ammoniiun hydroxide, and is reprecipitated by dilute
sulphuric acid. Paper or cotton goods can be passed through
first one and then the other of these liquids, and so receive a tough,
waterproof surface. Artificial silk is made by pressing the solu-
tion through dies into the precipitant. It can be dyed to any
desired tint, and is at least as brilliant in appearance as the natural
article. Its strength, however, does not equal that of the natural
silk fibre, especially when wet.
Cotton, when dipped in strong sodium hydroxide solution,
and then stretched while drying to prevent the shrinkage which
otherwise occm^, acquires a brilliant luster and is used in enor-
mous quantities under the n*ame of mercerized cotton.
Finally, mercerized cotton, or wood pulp treated with caustic
soda, combines with carbon disulphide to give viscose. Viscose
dissolves in water, and decomposes in solution giving a plastic
form of cellulose. This can be rolled into transparent sheets,
made into caps for bottles, moulded into any form, or pressed
through dies into solutions of salts of ammoniimi to give an-
other form of artificial silk.
Starch (CeHioOs)*. — Starch is found in plants in little colorless
granules of various roimded shapes (Fig. 99) which may readily
be seen under the microscope. These
granules are massed in large numbers in
the ears of wheat and oats, in the tubers
of potatoes, in the grains of maize \qM W^ (^
(com) and in peas and beans. Even vgr ^-g.
in the leaves they can be seen, im- 'Br
mediately after the plant has been ex-
posed to sunlight. They gradually disappear from the leaves
400 smith's intermediate chemistry
in the dark. They can be recognized, not only by their appear-
ance, but, without a microscope, by the iodine test. When a
drop of a potassium iodide solution, rendered brown by the ad-
dition of a little free iodine, is placed on the leaf or other part of
the plant, the granules of starch become blv^ while the other parts
are not affected.
Preparation of Starch. — If flour, which is made by grinding
wheat, and is three-fourths starch, is placed in a muslin bag and
kneaded under water, the granules of starch are washed out and
render the water milky (p. 5). After a time the granules settle
and the water can be poured off. Starch is manufactured by
washing disintegrated potatoes (in Europe) or maize (in America)
on sieves, and collecting and drying the white powder deposited
in the water used for the washing.
Starch is not soluble in water. If it be boiled with water, how-
ever, the granules swell and break, and the starch becomes finely
diffused through the water, forming a clear liquid. With little
water, a sort of transparent jeUy is produced. When the liquid
is poured through a filter, a large part of the starch goes through
the paper as if it were truly dissolved. Such a liquid is caUed a
colloidal suspension (p. 109). Imitation solutions like this are
constantly met with in using complex organic compoxmds such as
enter into jeUies, glues, soaps, and the juices of the bodies of an-
imals. Even inorganic substances, of the insoluble class, give
such suspensions. A description of their peculiarities must be
noticed under soap (p. 440).
The colloidal suspension of starch is v^ed in the laundry, for
stiffening white goods. Glucose is manufactured from it.
Since neither cellulose nor starch can be vaporized without
decomposition, and since they do not form true solutions in any
of the common solvents, we have no means of determining their
molecular weights, and are therefore forced to write their formulse
in the indeterminate forms (CeHioOs)^ and (CeHioOe)* respectively.
PLANT LIFE. CELLULOSE, STARCH AND SUOAR 401
The molecule in each case is, in all probability, exceedingly com-
plex.
Glucose C6H12O6 from Starch. — When starch is boiled with
water, to which a few drops of an acid (catalyst) such as hydro-
chloric acid have been added, the liquid, after neutralization of
the acid, is found to be sweet in taste. A kind of sugar, glucose
C6H12O6, can be obtained in crystals by evaporation. In com-
merce the evaporation is stopped before crystallization begins,
and the syrup (" corn-syrup,'' if maize is the source of the starch)
is sold for making candy and for preserving fruits.
(CeHioOe), + a;H20 -^ aKDeHiaOe-
Glucose is known also as dextrose, and as grape sugar. Brown-
ish crystalline granules foimd in dried grapes (raisins) are mainly
composed of it. When pure, it is almost colorless. It reduces
cupric hydroxide, in Fehling's solution (p. 513), to cuprous oxide.
The Sugars. — The common sugars are divided into two
classes. There are several sugars, having the same formula,
C6H12O6, but different properties, which are called monosacchar-
ides. Other sugars, having twice as many carbon units m the
formula C12H22O11, are caUed disaccharides. The sugars we have
occasion to mention here are the following:
Monosaccharides: Glucose (dextrose or grape sugar) C6H12O6.
Fructose (fruit sugar) C6H12O6.
Disaccharides: Sucrose (cane-sugar, beetnsugar, saccha-
rose) C12H22O11.
Maltose (formed by action of malt on
starch) C12H22O11.
Lactose (milk-sugar, found only in animals)
C11H22O11.
Carbohydrates. — Since cellulose, starch and the sugars are
freely changed, one into another, they are grouped together in one
402 smith's intermediate chemistry
class, the carbohydrates. The word refers to the fact that they
contain hydrogen and oxygen in the proportions required to form
water, and are, therefore, in a sense, hydrates of carbon. When
dehydrating agents Uke concentrated sulphuric acid (p. 270)
act on the carbohydrates, a black mass of carbon is left.
Sucrose or CanC'Sugar C12H22OU. — The sugar-cane and the
beet produce exceptionally large amoxmts of this sugar, which is
the one commonly used as table sugar. Maple sugar, obtained
by evaporating the sap of the tree, is composed mainly of the same
substance.
The sugar-cane forms stalks from ten to twelve feet high.
The juices are extracted by crushing the plants between rollers.
The liquid is evaporated in closed pans. A vacuum maintained
in the pans permits the boiling of the solution at a low tempera-
ture (about 65 degrees) and prevents the decomposition of a part
of the sugar which would otherwise occur. When the syrup cools,
the sugar crystallizes and the crystals are freed from the liquid in
centrifugal machines. The crystals are brown in color. At the
sugar refinery they are dissolved, and the solution is passed through
a column of bone charcoal. This adsorbs the coloring matter,
and the filtrate is once more evaporated and allowed to crystallize.
Refined cane-sugar has a faint yellow tint, and a small amount
of ultramarine is added to cover up this tint, and give the
white appearance which is popularly connected with purity in
sugar.
The sugar beets, which contain 16 per cent or more of cane-
sugar, are sliced and steeped in water to extract the sugar. The
liquid contains gummy material in colloidal suspension. This is
coagulated and precipitated by adding " milk of lime '' (calcium
hydroxide Ca(0H)2 suspended in water) and boiling. Carbon
dioxide is then passed through the solution to precipitate the
excess of lime:
Ca(0H)2 + CO2 -^ CaCOa i + H2O.
PLANT LIFE. CELLULOSE, STARCH AND SUGAR 403
The solution is decolorized with charcoal and evaporated to
crystallization in the same way as is the extract from the sugar-
cane.
Properties of Sucrose. — Sucrose crystallizes in four-sided
prisms, the form of which is seen in " rock-candy." It melts
at 160**. It does not reduce Fehling's solution (p. 401). When
heated to 200 to 210** it begins to decompose, slowly losing water
and leaving a brown, soluble mass called caramel, used in color-
ing whiskey and soups.
When boiled with water, to which a trace of an acid catalyst
has been added, it is hydrolyzed, giving a mixture of the two
monosaccharides, glucose and fructose:
C12H22O11 -|- H2O — > C6H12O6 "h C«Hi206.
This mixture of glucose and fructose is caUed invert sugar
and is found in many sweet fruits and in honey. Each sugar
interferes with the crystallization of the other, by lowering the
freezing-point (p. 119), and so invert sugar is added in making
" fondant " candy and candy that is to be " pulled," both of
which are intended to remain soft for some time. With the same
object in view, vinegar, lemon juice, or cream of tartar is added
to a syrup made from cane-sugar, in order that the acid contained
in them may produce some invert sugar and so give a less crystal-
lizable mixture (icing for cakes). Prolonged heating has the
same effect.
Exercises.—- 1. What inference do you draw as to the composi-
tion of tapioca, sago, and rice from the facts that they are plant
products and when boiled with water and cooled give a jelly-like
mass? How should you confirm your inference?
2. (a) Why does a concentrated solution of sugar boil at a
temperature far above that of boiUng water? (b) In evaporation
why is the boiling-point lower in a vacuum than in air?
404 SBflTH's INTERBfEDIATE CHEMISTRY
3. In what relative volumes are carbon dioxide used and
oxygen produced by a plant?
4. What products must be formed when paper is burned?
Make the equation.
CHAPTER XXXV
PLAirr LIFE. OSMOSIS. FERTILIZERS
For successful growth, plants require carbon, hydrogen, oxy-
gen, phosphorus, potassium, nitrogen, sulphur, calcium, iron, and
magnesium. The carbon is chiefly supplied by the carbon dioxide
in the air, as we saw in the preceding chapter. Water supplies
hydrogen and oxygen, entering through the leaves, and also
through the roots and stems. Oxygen is also supplied directly
by the au-. The oxygen produces oxidation of substances in the
plants, and gives heat. Phosphorus, nitrogen, and sulphur are
absorbed from the soil by the roots as soluble phosphates, nitrates,
and sulphates, and nitrogen sometimes as ammonia. Potassium
commonly enters as carbonate or bicarbonate. Calcium and
magnesium are absorbed as phosphate, nitrate, sulphate, or
bicarbonate; and iron as ferric hydroxide (Fe(0H)8). Man-
ganese, chlorine and silicon are also present in many plants.
In some species, sodium salts can take the place of potassium
salts.
The last named elements are used in such small amoxmts,
relative to the available supply in the soil, that they rarely need
attention. Nitrogen^ phosphorus^ potassium and calcium, how-
ever, the presence of which in quantity is essential to the life and
development of plants, often need to be added to soils which
are deficient in these elements. Suitable compounds of nitro-
gen, phosphorus, potassium and calcimn are therefore used in
enormous quantities as fertilizers.
Before discussing individual fertilizers, we may profitably take
up the general question of how the plant derives its food from the
soil.
405
406 smith's intermediate chemistry
Haw the Phmt Feeds. — We have mentioned (p. 395) that the
plant juice or sap contains soluble sugars, which travel to those
parts of the plant which require new material for their growth.
The moist walls of the root-hairs of the plant are freely permeable
to water, the entrance of which into the plant from the soil is
necessary to offset evaporation from the leaves and stems. Sol-
uble salts present in the soil water are also able to pass, although
less freely, through these walls, and are incorporated in the sap.
If they are of nutritive value to the plant, they react, as they
circulate through the plant from cell to cell, with the organic
constituents there present, forming more complex compounds
which are imable to permeate the boimdary walls of the cells,
and thus become permanently fixed in the growing parts. If
they are not of nutritive value, they complete the circuit im-
changed.
Now the passage of salts from the soil water into the plant
continues only imtil the sap inside contains the same concen-
tration of each salt as the solution outside. Hence while nutritive
constituents, which are removed from the sap dming its circula-
tion through the plant, can continuously enter to keep up the
supply necessary for growth, non-nutritive constituents soon reach
their equilibrium concentration and are thereafter rejected.
Osmosis. — The membranes lining the cell walls and the root-
hairs of plants exercise, as we saw above, a selective action with
regard to the passage of different substances through them. Water
is able to permeate the membranes quite freely, dissolved salts
pass through less readily, while the movement of complex organic
substances, such as proteins, is completely blocked. Upon this
selective flow of materials through the plant membranes, it must
be noted, the life of a plant is absolutely dependent. If the
membranes were freely permeable to the organic materials con-
tained in the plant sap, the plant would soon lose these materials
to the outside soil and die of exhaustion.
PLANT LIFE. OSMOSIS. FERTILIZERS
407
The selective flow of certain components of a solution through
a membrane is known as osmosis. Osmosis and its consequence,
osmotic pressure, are phenomena which may be very clearly
illustrated in the laboratory by means of a solution of cane sugar
and an artificial membrane of precipitated cupric ferrocyanide
Cu2.Fe(CN)6.
Osmotic Pressure. — A suitable semi-permeable membrane
may be obtained by soaking a clean porous pot in a solution of
potassium ferrocyanide K4.Fe(CN)6 (p. 497),
rinsing in water, and then aUowing to stand
in a solution of cupric sulphate. The diffusion
of the latter substance into the cell produces,
by double decomposition, a film of insoluble
copper ferrocyanide within its walls. This
film is freely permeable to water molecules,
but not to molecules of sugar dissolved in the
water.
A simpler method is to use a diffrision shell
of specially treated parchment, of test-tube
form. This, however, is not entirely im-
^fiili ^ fl-^^j*^ permeable to sugar jnolecules.
The porous pot or diffusion shell, filled
with sugar solution, is securely attached to
a long glass tube, and suspended in pure
water (Fig. 100). It is found that the level of
the Uquid in the tube gradually rises until, if
the membrane remains intact and is truly im-
permeable to sugar, a definite hydrostatic
pressure is established, the magnitude of this
osmotic pressure depending only upon the
temperature and upon the fraction of sugar molecules in the
solution within the cell.
FiQ. 100
408 smith's intermediate chemistry
Explanation of Osmotic Pressure. — For a complete dis-
cussion of osmotic pressure, the reader is referred to a modem
text-book of physical chemistry. A brief explanation by means
of the molecular hypothesis, however, may be given here.
The molecules of water id the pure water outside the cell, and
the molecules of water and of sugar in the solution inside the cell,
are all in rapid motion (p. 94). When, in consequence of this
motion, they strike the membrane, water molecules have a chance
of passing through, but sugar molecules are all turned back.
Now the concentration of water molecules id the pure water otd-
sidcj striking the membrane and attempting to enter the cell, is
greater than the concentration of water molecules id the solution
inside, striking the membrane and attempting to leave the cell
(compare vapor pressures, p. 117). Hence more water molecules
will be entering than leaving, and the level of the Uquid inside the
tube must rise in consequence.
Why does the level of the Uquid stop rising when a definite
hydrostatic head has been established? Because now, although
there is still a greater concentration of water molecules in the
pure water outside than in the solution iDside, water molecules
attempting to enter the cell through the membrane are opposed
by the hydrostatic pressure, while water molecules attempting
to leave are assisted in their passage. We have on one side of the
membrane more water molecules with a smaller chance of getting
through, on the other side fewer water molecules with a greater
chance of gettiDg through. Equilibrium is reached, evidently,
v/hen these two factors coimterbalance.
The student should note very carefully the fact that the sugar
molecules are not directly concerned in the phenomenon of osmotic
pressure. Their function is merely to reduce the concentration
of water molecules in the solution inside the cell. Any solute, with
respect to which the membrane is similarly impermeable, will
give the same effect, the osmotic pressure at any given tempera-
ture being dependent only upon Oie fraction of solute molecules in
PLANT LIFE. OSMOSIS. FERTILIZEBS 409
the aolutionj that is, upon the extent to which the concentration
of waier molecules has been reduced (compare, again, vapor pres-
sure depression of solutions, p. 117).
A solution containing 1 g. molecular weight of an inert solute
dissolved in 1000 g. water should give at 20**, with a perfectly
semi-permeable membrane, an osmotic pressure of 23.6 atmos-
pheres.
Osmotic Pressure in Plant Life. — Osmotic phenomena
were first studied by Pfeflfer (1877), a botanist, who used certain
plant cells for the purpose. The cell content included a liquid
containing various salts in solution, and a protoplasmic layer
lining, but not firmly attached to, the cell wall. This protoplas-
mic layer behaved like an imperfect semi-permeable membrane.
When such cells were immersed in a concentrated solution of
any substance, the water passed from the interior of the cell to
the solution, and by means of a microscope a shrinkage of the
protoplasmic layer away from the cell wall could be observed.
Conversely, when such cells were placed in pure water, or a solu-
tion of a very dilute nature, water passed from the outside into
the interior, and the protoplasmic layer was distended so aS' to
cause the cell to become tm^d.
Osmotic pressure is, therefore, a subject of great interest in
connection with the physiology of plants. It aids in explain-
ing why a withered flower, containing a solution in its cells, re-
vives when placed in pure water. The latter enters through the
walls of the cells, and the pressure thus produced distends the
structure and stiffens it. Similarly, the wilting of plants, when
too high a concentration of salts as fertilizers is added to the soil,
is explained. In the animal body also, osmosis plays a large part.
Fertilizers, — Many soils are either naturally deficient in one
or more of the necessary plant foods, or the supply may have
been exhausted by repeated cropping. Every crop removes
410 smith's intermediate chemistry
permanently a certain part of the supply. Thus, m the case
of nitrogen, an average crop of maize or com (45 bushels) removes
63 poimds per acre, a crop of cabbage (15 tons) removes 100 poimds
per acre, clover hay (2 tons) 82 poimds, and wheat (15 bushels)
31 pounds. When the supply becomes reduced, the crops be-
come poor. Moreover, the necessary elements must be present
in soluble form, or they cannot enter into the plant system.
Felspar KAlSisOg is a common constituent Df many rocks, such
as granite (p. 4). When such rock material, contained in the soil,
is decomposed by weathering j through the action of carbonic acid
from the atmosphere, the felspar gives day HAlSi04 and soluble
compounds of potassium. There are immense quantities of
felspar available, but the process of weathering is very slow, and
in many agricultural regions the soil is therefore deficient in sol-
uble salts of potassium.
It is just as necessary to feed crops as to feed cattle, and equally
foolish to starve either of them. Fertilizers are used to make
good the original, or acquired deficiences of the soil in the most
important elements, nitrogen, phosphorus, potassium and calcium.
It is absolutely necessary, in addition^ to keep up ia sufficient
supply of fresh organic material in the soil, or the use of fertiUz-
ers may result in more harm than benefit.
The value of the systematic use of fertilizers is indicated by
comparison of the average crop of wheat per acre in different
countries. The average of ten successive years is: Denmark 40
bushels. Great Britain 33, Germany 29, United States 14.
Nitrogen. — The nitrogen is suppUed as sodium nitrate or
guano (p. 369), calcium nitrate (p. 313), anmionium sulphate
(p. 299), calcium cyanamide (p. 392), manure or the offal (" tank-
age ") and ground bones from slaughter houses.
Over every acre of the earth's surface there are 34,000 tons of
free atmospheric nitrogen. Plants in general are incapable of
drawing upon this inmiense store for the nitrogen necessary for
PLANT LIFE. OSMOSIS. FERTILIZERS 411
their growth, but peas, beans, clover, alfalfa and other leguminous
plants bear round their roots colonies of a special kind of bacteria
which has the power to bring free nitrogen into combination.
In these root nodules (Fig. 101) the bacteria first produce pro-
teins, which later decompose, and ultimately, by bacterial action,
yield nitric acid. In this way a crop of clover will
fertilize the soil, not only for itself, but also for the
following crop. The advantage of rotation of crops
is therefore explained.
The beneficial action of bacteria upon plant growth
is not limited to this special case. The soil is crowded
with bacteria which assist in the decomposition or
rotting of vegetable and animal matter. The product,
the black, gummy stuff of a fertile soil, is called yiq loi
humus. Certain varieties of bacteria convert insoluble
carbohydrates such as cellulose into simpler soluble mate-
rials, immediately available for plant use. Other species lib-
erate the nitrogen from proteins in the form of ammonia and
amino-compounds (p. 353). Others, again, oxidize these com-
poimds to nitrous add si,nd nitrites. Still others, finally, oxidize
these products to nitric add and nitrates.
The actual assimilation of nitrogen by plants, in whatever
form that element is originally present in or added to the soil,
takes place almost exclusively through soluble nitrates. The
co-ordinated team-work of all of the above classes of bacteria is
therefore an essential point in the efficient utilization of other
nitrogen-containing fertilizers.
Phosphorus. — Natural calcium phosphate Ca8(P04)2 is the
orthophosphate of calcium. It is foimd in considerable deposits
in S. Carolina, Florida, Tennessee, and several western states,
and in Algeria and Tunis. This insoluble salt, however, affords
only an exceedingly dilute phosphorus diet for plants. Hence,
a more soluble compoimd is to be preferred. This is found in
412 smith's intermediate chemistry
calcium acid-phosphate (" superphosphate ") CaHiCPOi)^, which
is made by heating pulverized natiuul calcium phosphate with
sulphuric acid, containing the requisite proportion of water:
Ca3(P04)2 + 2H2SO4 + 4H2O -^ CaH4(P04)2 + 2CaS04,2H20.
The whole turns into a dry mixture, consistiag of the superphos-
phate and gypsum (hydrated calcium sulphate). The latter does
not interfere with the fertilizing power of the soluble superphos-
phate, so the mixture is placed directly in sacks and sold as " super-
phosphate of lime."
In slaughter houses the bones, after being deprived of fat and
gelatin, give a residue containing much calcium phosphate. This
residue is treated with sulphiuic acid and made into fertilizer.
Potassium. — Wood ashes contain much potassium carbonate,
and are used as fertilizers for this reason. The giant sea-weeds
{Kdp) of the Pacific coast have also been foimd to contain an
imusually large proportion of salts of potassium. By far the
most important source of this element, however, is potassium
chloride, obtained from natural salt deposits.
The average production of the Grerman deposits at Stassfurt
in the ten years preceding the war exceeded 1,000,000 tons (cal-
culated as K2O). These deposits constituted practically a world
monopoly, and the shortage of potassium salts for fertilizer pur-
poses during the war was consequently extreme. The most
strenuous efforts to develop the potassium resources of the United
States (natural brines. Kelp, recoverable by-products from the
cement and molasses industries, etc.) culminated in a production
of only 50,000 tons K2O in 1918. Fortunately, the cession of
Alsace to France has destroyed the Grerman monopoly, since very
extensive deposits exist near Mulhouse. The working of these
deposits has been greatly hampered by the damage done to shafts
and machinery by the Germans before their evacuation, but
production in 1920 already exceeded 200,000 tons K»0.
PLANT LIFE. OSMOSIS. FERTILIZEBS 413
Potassium chloride occurs in the Stassfurt deposits as syhine
KCl, but is chiefly associated with magnesium chloride as car-
naUite KCl,MgCl2,6H20. When water is added to camaUite, a
large part of the potassium chloride, which is much less soluble,
separates out. Complete extraction and purification of the salt
involves a series of recrystallizations. The Alsace deposits con-
tain very little magnesium salts, and the separation of the potas-
sium chloride from the sodium chloride with which it is mixed is
comparatively simple.
Potassium sulphate is also obtained from salt deposits, and is
substituted for potassium chloride as a fertilizer for certain crops,
such as tobacco. Chlorides, in general, melt at lower tempera-
tures than sulphates, and the presence of a chloride in tobacco
results in an ash that fuses on burning. This is, obviously, an
undesirable property, especially for cigars.
Calcium. — Calcium is naturally present in many soils as
calcium carbonate CaCOs. In contact with water containing car-
bonic acid (see p. 400), this gives a solution of the more soluble
bicarbonate Ca(HC08)2. Other compoimds of calcium which are
used as fertilizers include lime CaO or Ca(0H)2, calcium phos-
phate and acid phosphate, calcium sulphate or gypsum (p. 387),
and calcium cyanamide.
Manure. — One ton of farm manure contains about 10 pounds
of nitrogen (chiefly as urea CO(NH2)2 and proteins), 5 poimds of
phosphoric acid and 10 poimds of potash. The manure is mixed
with the soil just before planting seed, or is used as top dress-
ing.
The bacteria in the air and in the soil assist materially in the
changes in the manure. Thus urea is hydrolyzed to anmioniiun
carbonate (NH4)2C08:
CO(NH2)2 + 2H2O -^ (NH4)2CO,.
414 smith's intermediate chemistry
The proteins are changed by air bacteria into ammonia. The
formation of nitrates seldom happens in manure piles, but scat-
tered manure, with the help of soil bacteria, develops nitrates.
The potassium compoimds form potassium hydrogen carbonate
KHCOa. The phosphorus and sulphur compoimds become sol-
uble phosphates and sulphates.
Indirect Fertilizers. — Not all substances which are added
to the soil are employed with the direct object of their assimilation
for plant growth. Often indirect effects induced by their pres-
ence are of greater importance. A few cases where calcium salts
are of service as indirect fertilizei-s may be briefly presented as
illustrations of this point.
(1) Gypsum CaS04,2H20 is added to manure at the rate
of 100 pounds per ton. This slightly soluble salt reacts in the
soil solution with the ammonium carbonate produced by the hy-
drolysis of urea, precipitating the much less soluble calcium car-
bonate and leaving ammonium sulphate in solution:
(NH4)2C08 + CaS04 -^ (NH4)2S04 + CaCOs i .
Now ammonium sulphate, being a salt of a strong acid with a
weak base, is only very slightly hydrolyzed in solution (see p.
369). Ammonium carbonate, however, is a salt of a weak acid
and a weak base, and is extensively hydrolyzed. This hydrolysis,
unless a large excess of water is present, would lead to rapid loss
of ammonia from the manure. The smell of free ammonia, in-
deed, is often very noticeable in manure piles. The addition of
gypsum fixes this valuable constituent in the fertilizer for plant
use. Lime, on the other hand, would assist in the liberation of
ammonia.
(2) Gypsum is often added, also, to clay soils with the object
of converting insoluble compoimds of potassium into more sol-
uble compounds. Lime and calcium carbonate are employed to
PLANT LIFE. OSMOSIS. FERTILIZERS 415
change insoluble phosphates of iron and aluminium into more
soluble calcium phosphates.
(3) Some soils are either naturally acidic or sour, or have
become so by excessive use of sulphate fertilizers, and are hence
imfavorable to plant growth. This acidity is corrected by addi-
tion of lime or calcium carbonate, but not gypsum.
(4) Salts which are ivQurious to plants, such as soluble magne-
sium compoimds or *' black alkaU " (sodium carbonate), are
converted into less soluble compounds, such as magnesium hy-
droxide or carbonate, by addition of hme or calcium carbonate,
or into non-poisonous compoimds, such as sodium sulphate, by
gypsum.
Adsorption in Soils, — As we have seen, fertilizers must
contain the elements necessary for plant growth in soluble form.
It is imdesirable, however, to use salts which are exceedingly
soluble in water as fertilizers, since they will obviously be rapidly
washed away from the surface soil. A large proportion of all
fertiUzers is unavoidably wasted in this manner. The minute
particles of the soil, however, possess the power of conserving
dissolved substances in the soil solution by concentrating and
holding them upon their moist surfaces. This phenomenon,
adsorption, is of considerable interest and importance in other
connections, and will be taken up in detail in a later chapter
(p. 421).
Exercises. — 1. Given an osmotic pressure cell and an im-
known substance, soluble in water, to which the membrane is
perfectly impermeable, how could you determine the molecular
weight of the substance?
2. What would be the eflfect of putting fresh flowers in a vase
containing a concentrated salt solution?
3. At 20** 100 c.c. of water, shaken up with excess NaCl and
KCl, dissolve 30 g. NaCl and 15 g. KCl. At lOO*" the same
416 smith's intermediate chemistry
amount of water, in equilibrimn with the same two solids, con-
tains 20 g. NaCl and 40 g. KCl. On the basis of these figures^
devise a method for separating KCl from NaCl in the Alsace
deposits.
4. The use of ammonimn sulphate as a fertilizer is apt to
result in an acid soiZ. Explam why.
5. What are the three elements most needed in fertilizers?
Of which does every country have a free and unlimited supply?
How can it be made available?
6. What valuable soil ingredients are lost to a locality when a
carload of wheat is shipped away? A carload of pure sugar?
Of cotton, (a) ginned or (b) unginned? Of peanut oil?
CHAPTER XXXVI
PLANT PRODUCTS. FERMENTATION AND FUELS
Having described the chemistry of plant life, we may now pro-
ceed to the chemistry of substances resulting from plant life and
growth. Foods will be taken up separately in the following chap-
ter. In the present chapter we shall restrict ourselves to two
other main branches, fermentation products and fuels.
Enzymes. — All fermentations are brought about either
directly or indirectly by the activities of animal or vegetable
organisms. The most familiar ferment, of course, is yeast.
Yeast belongs to a low order of plants and consists of minute
cells. Its value lies in the fact that, while growing and multiply-
ing, it secretes within each cell small amounts of two very active
chemical substances which are dissolved in the Cell contents.
These substances are known as zymase and invertase (or sucrase)i
and belong to the class of organic materials called enzymes.
Enzymes produce remarkable chemical changes in organic ma-
terials by their mere presence (contact actions). These changes
are specific^ each enzyme acting only on certain carbohydrates,
for example, and being quite inert towards others.
Fermentation of Sugars. — When a cake of yeast is broken in-
to an aqueous solution of glucose or grape-sugar (p. 401), the small
amount of zymase present causes the gradual decomposition of
the sugar. The most favorable temperature is about 30°. Bub-
bles of carbon dioxide soon begin to rise to the surface, and the
gas can be led oflF (Fig. 102) to exhibit its characteristic action
(p. 336) on limewater. At the same time alcohol C2H6OH ac-
cumulates in the liquid as the sugar disappears:
CoHwOe -* 2CO2 1 + 2C8H60H.
417
418
smith's intebmediate chemistry
a
\
The liquid extracted from the yeast cells works as well as does the
plant itself.
Yeast will ferment fructose (fruit sugar, p. 401), with the same
result, but more slowly.
Zymase does not act upon cane-
sugar (sucrose C12H22O11). But the in-
vertase (sucrase), which is also con-
tained in the yeast, hydrolyzes the
sucrose in the same way as does a dilute
acid, giving invert sugar (p. 403). The
latter is then decomposed by the zy-
mase. Hence cane-sugar in solution is
decomposed by yeast into alcohol and
carbon dioxide, just as is grape-sugar,
only more slowly.
\^ In the manufacture of wines the glu-
Fia. 102 cose contained in the grape juice is fer-
mented by a species of yeast always foimd on the skins.
Fermentation of Starch. — Barley, which has been allowed
to sprout, and is then dried, is called malt. This contains an
enzyme, diastase (or amylase), which is able to hydrolyze starch
into maltose C12H22O11 (p. 401). Maltose is further hydrolyzed
by another enzyme, maUasey to form glucose, and the latter is
then decomposed by zymase into alcohol and carbon dioxide.
Whisky is made by treating the starch of rye, maize, or barley
in the above way, with subsequent distillation. Beer is made
similarly from various kinds of grain, particularly barley, except
that the fermented liquid is not distilled.
Industrial Alcohol, — Alcohol has very extensive uses, apart
from its historic value as a beverage. It is employed as a solvent
in making varnishes for wood and lacquers for metal, as well
as for plastics like celluloid, collodion and artificial silk (p. 399),
PLANT PRODUCTS. FERMENTATION AND FUELS 419
It is of service in the purification of many natural organic products,
such as turpentiney and in the preparation of many synthetic
organic products, such as dyes. It is also rapidly coming into
use as a fuel, its smokeless flame and efficiency of combustion
making it of special importance for aeroplane and motor engines.
Solidified alcoholj obtained by the addition of cellulose esters
(p. 480), is now largely employed for cooking purposes. Another
new development is the catalytic production of ethylene for cid-
ting and welding purposes (p. 352).
Alcohol, as used in the industries, is denatured^ or rendered
unsuitable for drinking purposes, by addition of small quantities
of benzine (p. 345), pyridine bases, or other disagreeable and
non-removable organic liquids. The exact formula of denatured
alcohol depends upon the use for which it is intended.
The cheap production of industrial alcohol is rendered possible
by the utilization of certain waste materials rich in carbohydrates.
When the price of food is high, grains are employed in the
manufacture of alcohol only when a crop has been damaged in
some manner so that it cannot be sold as a food material.
When sawdust or wood refuse is heated with dilute sulphuric
acid imder pressure, the cellulose is converted into fermentable
sugars by hydrolysis (compare p. 401). At the present time,
however, the most important source of industrial alcohol in the
United States is molasses. Only a few years ago the dis-
posal of molasses furnished a very troublesome problem to the
sugar mills, but in 1918 nearly 120,000,000 gallons of industrial
alcohol were obtained from this " waste product " in the United
States alone.
Acetic Add CH3COOH. — This acid is formed by the partial
oxidation of alcohol (p. 348). Vinegar (crude acetic acid) is
manufactured by oxidizing alcohol with atmospheric oxygen,
using a bacterium (jB. Aceti, " mother of vinegar ")> or more
probably an enzyme which it secretes, as a contact agent. The
420 BMITH S IimaiMEDIATB CHEHIBTRT
dilute alcohol, in the form, for example, of "hard " cider (fep-
mented apple juice), is allowed to trickle over shavings in a barrel.
The shavings are inoculated with the B. aceti by preliminary
wetting with vinegar. Holes in the sides admit a plentiful sup-
ply of air, to the action of the oxygen of which the hquid is ex-
posed by being spread over the surface of the shavings:
CiHsOH + 0» -> CH,.COOH + HjO.
The liquid (vinegar), which issues at the bottom, contains from
5 to 15 per cent of acetic acid, besides coloring and flavoring
matters derived from the fruit juices.
Pure acetic acid may be prepared by distilling the vin^ar
repeatedly. It is derived more cheaply, however, from the liquid
distUlate obtained by heating wood in the manufacture of char-
coal. Lai^e quantities are used in the manufacture of various
synthetic organic products (see, for example, p. 480).
Fuels
Destructive Distillation of Wood. — When dry wood is
heated in u-on retorts in absence of air, the compounds which
it contains are decom-
posed. Much of the
carbon remains in the
form of charcoal. The
vapors which pass off
(through the pipe on
the right, Fig. 103)
^'^- 1"^ deposit, when cooled,
much liquid material. The uncondensed gases are combustible
and are used for heating the retorts or other similar purposes.
Hard wood furnishes, approximately, 25 per cent of its weight
of charcoal, 25 per cent of gases, and 50 per cent of liquids.
TTie liquid contains acetic acid (10 per cent), methyl alcohol
CHjOH or wood sphit (3 per cent), a complex, tany mix-
PLANT PRODUCTS. FERMENTATION AND PUEtS 421
ture, used in road-making (10 per cent), water (77 per cent),
and a little acetone (p. 349). The distillate from resinous wood
also contains valuable quantities of turpentine, CioHu, an un-
saturated hydrocarbon used extensively as a solvent. The gases
evolved contain a lai^e part of the nitrogen of the original proteins
in the form of ammonia, which ia dissolved out with water.
When charcoal only is desired, the wood is stacked, covered
with turf (Fig. 104), and set on fire. A part ia burned, the rest
is converted into charcoal, and all
the valuable volatUe products are lost.
Properties of Wood Charcoal.
Adsorption. — The charcoal retains
the structure — a complex network of
minute cells — of the original wood, and therefore has a surface
which is vast in proportion to the amount of material it contains.
Upon this surface it is capable of taking up or of adsorbing many
times its own volume of gases, especially of the more condensible
ones. Thus, boxwood charcoal takes up ammonia (90 volumes),
hydrogen sulphide (55 volumes), and oxygen (9 volumes).
TTie adsorption is extremely rapid and, in the case of a conden-
sible gas contained in small quantity in air, practically complete.
For this reason, charcoal and other substances with very finely
divided surfaces are used as adsorbent materials for industrial
gases and vapors (compare silica gel, p. 360).
The toxic gaaes employed in the Gre^ War are also readily
adsorbed by charcoal. Hence the canisters of gas masks con-
tain layers of porous charcoal, together with granulated sodar
lime and potassium permanganate, which react chemically with
certain of the noxious gases liable to be present. During the war
vast strides were made in increasing the adsorptive power of
various kinds of charcoal by modifications in methods of carboniz-
ation, the most efficient of all charcoals being dense varieties
derived from cocoaniU shells and fruit pits. Canisters packeC
422 smith's intermediate chebhstry
with such charcoal reduced the concentration of all toxic gases
employed in the war (see Chapter XL) below the danger limit.
Toxic smokes, however, were not satisfactorily adsorbed. The
explanation is similar to that advanced in a preceding chapter
(p. 263) for the persistence of the fog obtained when a mixture of
sulphur trioxide and oxygen is bubbled through water. The
molecules of a gas are in such rapid motion that they are prac-
tically certain to strike the surface of the charcoal while passing
through the canister, and to be adsorbed on this surface if the gas
is easily condensible. The dimensions of the solid smoke par-
ticles, however, are much larger than molecular, and the particles
are relatively stationary. Most of them, in consequence, are
able to pass through the air channels between the charcoal granules
without touching.
Pulverized charcoal, when shaken with a liquid, is also able to
extract from it any dissolved substances, and to concentrate them
upon its surface. Other finely-divided materials, such as soil
particles (p. 415), possess the same power of adsorbing substances
from solution. Salts are, in general, only partially taken up;
organic solutes are removed more completely. This property
of charcoal is made use of in water purification and in sugar re-
fining (p. 402). Charcoal is used, also, in making gunpowder, in
reducing ores, and as a fuel (smokeless).
CoaL — When wood bums with a plentiful supply of oxygen,
it gives nothing but carbon dioxide, water, free nitrogen, and a
certain amount of ash (oxides and carbonates of the metals).
What happens when it is heated in absence of oxygen, we have
just seen. In natiu^, however, the intermediate case of slow
decomposition of vegetable matter, without much heating and
without access of oxygen, takes place on a large scale. Clay and
sand, or even simply water, cover the vegetation and exclude the
air, and the products are anthracite coal, bituminous coal, or peat.
Little is known of the actual compounds contained in coal. We
PLANT PRODTTCTS. FERMENTATION AND FUELS 423
are concerned mainly with the products obtained by heating it
in the absence of air, and with its use aa a fuel.
Bituminous coals give much, and widely varying amounts of
volatile matter; anthracite coals give very little. The ash is
the mineral matter of the original plants, with additional rock
materials m some specimens. The coal is selected according to
the purpose for which it is to be used. For coal gas, and even for
coke, a variety h^h in volatUe matter is chosen. For water gas
(p. 203) anthracite or coke itself is employed.
(Ami Gas. — The gas plant (Fig, 105) includes (1) the fire-brick
retorts in which the coal is heated (externally) to 1300°, (2) the
hydraulic main (a wide iron pipe) immediately above them in
Fia. 106
which most of the tar collects, (3) the condenser and wash box for
cooling, condensing, and removing oils, (4) the scrubbers (ver-
tical and rotary) where the ammonia is taken out by water drip-
ping over strips of wood and by stirring the gas with water, (5)
the purifier where hydrogen sulphide is taken up by hydrated
ferric oxide, and (6) the holder in which the gas collects.
The yield of gas varies considerably with the type of coal used.
One ton of good bituminous coal should produce approximately
10,000 cubic feet of coal gas, 1300 pounds of coke, 5 pounds of
ammonia and 12 gallons of tar. The average composition of coal
gas is: Illuminants 4 per cent, carbon monoxide 8 per cent, hydrogen
50 per cent, methane 29 per cent, ethane 3 per cent, carbon dioxide
2 per cent, oxygen and nitrogen 4 per cent.
424 smith's interbcediate chemistry
The ammonia is made into ammonium sulphate. The tar may
be used for road-making, as a waterproof material in building,
and wherever pitch is applicable. More frequently it is sepa-
rated by distillation, and other forms of treatment, and jdelds
benzene CbH^, naphthalene CioHg, anthracene C14H10, phenol or
carbolic acid CeHsOH and innxmierable other valuable substances.
Coke Ovens. — The by-product coke oven is very much like
the plant used for making coal gas. The difference is thac the
heating is arranged so as to decompose the volatile matter and
cause it to leave as much as possible of its carbon behind. The
resulting gas is consequently poor in illmninants, but excellent
as a fuel. The ammonia and tar are also diminished in amount,
but are still produced in paying quan-
tities.
The beehive coke oven (Fig. 106),
now largely discarded, is a primitive
device of fire-brick, shaped like a bee-
hive. It is simply filled with coal, part
of which is allowed to biun with a
limited supply of air. It yields 66 per
cent coke, against 73 per cent from the
by-product oven. All the volatile matter, with its gas, ammonia,
and tar, escapes through an opening at the top, where it bums in
a large flame and is wasted.
Properties and Uses of Coke. — Coke is a grey-black, hard
material of spongy texture. It bums without flame, and gives a
higher temperature than does coal, because no heat is used in
vaporizing moisture and volatile matter. On accoimt of these
and other properties, it is used in immense quantities in reducing
ores of iron and other metals, and in smaller amounts in electric
furnace work and in making electric light carbons.
OIL SHALE CLIFF, UTAH
PLANT PRODUCTS. FERMENTATION AND FUELS 425
Coal CIS Fuel. — The quality of a fuel coaJ, and whether* it is
worth its price, is learned by measuring its calorific (heating)
power. A sample (about 1 g.) is burned in a bomb calorimeter.
This is a closed, metal vessel, filled with oxygen and submerged
in a known weight of water. The coal is set on fire by a wire
heated electrically and, after it has burned, the increase in tem-
perature of the water is read off. Hence the heat in calories
(p. 162) evolved by the burning of 1 g. of coal is obtained. In
engineering practice they use the nxmaber of British Thermal
Units (1 B.T.U. = heat required to raise 1 poimd of water 1® F.)
developed by 1 poimd of coal, and call the result the calorific
power.
Hydrogen
Charcoal (to CO2)
Wood (seasoned) .
Calories
per 1 g.
28,800
8,080
4,750
B.T.U.
per 1 lb.
61,840
14,544
8,550
Bituminous coal
Anthracite
Petroleum
CalcH'ies
per 1 g.
7,800
8,000
11,000
B.T.U.
per 1 lb.
14,040
14,400
19,800
Knowing that 100 cal. will raise 1 g. of water from 0® C. to
100® C, and 539 cal. more will convert it into steam, it is possible
to calculate how much steam should be furnished by 100 kilog.
of coal of known heat of combustion. If the quantity falls short,
then the furnace, draft, or method of firing may be defective.
Too much draft, for example, merely introduces additional, useless
air to be heated. Thus, if the flue gas, upon analysis, is foimd
to contain, not 12 per cent carbon dioxide (normal), but only 3
per cent, then for every ton of coal burned, 62 tons of imnecessary
air have been raised to the temperature of the furnace. By
chemical tests, made in ways like this, the efficiency of every
device in the modem factory is (or ought to be) controlled. If
the coal is bought without heed to its calorific value, and used
without experimental checks, the boiler house alone may easily
waste the whole profit earned by the rest of the plant.
426 smith's intermediate chemistry
Source of the WorWs Energy. — The energy that does the
world's work comes mamly from two sources, namely, water power
and the combustion of wood, or of coal (which is fossil wood).
The water comes from vapor, generated by the sun^s heaij con-
densed as rain, and collected in lakes or reservoirs. The source
of the energy of coal or wood is a little less obvious. When wood
(which is largely cellulose) bums, it gives carbon dioxide, water,
and heat. In fact, its combustion is represented by the equa-
tion given on p. 397, when the equation is read backwards.' Thus
the sunlight, working through the machinery of the plant, takes
the carbon dioxide and water, furnishes the energy (as Ught),
and gives us wood and oxygen. And the wood and oxygen, when
burned, give us back the original substances, and the equivalent
of the original energy in the form of heat. Hence, our other
main source of energy turns out to be the same as the first — the
sun's rays — although the route by which the energy comes to us
is a little less direct.
If, instead of burning the starch of the plant, we consume it as
food, it goes through a series of changes instead of only one. But
the end products are the same, namely, carbon dioxide and mois-
ture issuing from our limgs, and heat and other forms of energy
such as are developed in living organisms. Thus, whether we
use our muscles, a steam engine, or a waterwheel to do work,
sunlight is in each case the ultimate soiu-ce of energy employed.
Exercises. — 1. In fermentation, why does not carbon dioxide
appear in bubbles at oncef
2. How do we ascertain that acetic acid in aqueous solution is
only slightly ionized? Give as many methods as possible.
3. (a) Why are charcoal and coke smokeless fuels? (b) Ex-
plain why bitiuninous coal bums with flames while anthracite
does not.
4. Point out the analogies between the processes used in making
coke and charcoal, and between their properties and uses.
PLANT PRODUCTS. FERMENTATION AND FUELS 427
5. A gas of sp. gr. 0.43 (air = 1) gives, on burning, 610 B.T.U.
per cu. ft. How many B.T.U. is this per pound?
6. How many kilog. of steam, from water at 20®, can be made
by bmning 100 kilog. of coal, the heat of combustion of 1 g. of
which is 8500 caL?
7. At 6 atmos. pressure and 152® C, how many cubic meters
will 190 kilog. of steam occupy?
8. What is " conservation?" What foiu* industries or opera-
tions (or ways of performinjg operations) that are wasteful have
been mentioned in this chapter (compare p. 424)?
CHAPTER XXXVII
ANIMAL UFB AND ANIMAL PRODUCTS. FOODS
Only a few of the more important points in the chemistry
of animal life and growth can be touched upon here. In the same
way, the chemistry of foods is presented merely in outline. Some
food products, derived from plants, have been described in earlier
chapters. The single animal product dealt with in any detail in
the present chapter is soap. In connection with soap, the subject
of colloids is also briefly discussed.
Composition of the Human Bfpdy. — The following gives,
roughly, the percentage of each element in the himian body.
We have already learned that the calcium and phosphorus are
chiefly in the bones (p. 412). The nitrogen, sulphur, and iron are
in the proteins. The sodiimi is largely present as salts, in the
fluids of the body. The potassimn is in the soft tissues and in
special secretions like milk. As in the plant, the carbon, hydrogen,
and oxygen are in the form of carbohydrates, proteins, and fats,
and there is also much water.
Certain amounts of all these elements leave the system daUy.
Water evaporates from the lungs and skin. The carbon leaves
in large amounts, chiefly from the lungs as carbon dioxide, and
also as excreted fats, proteins and carbohydrates. Much of the
nitrogen is eliminated, chiefly as urea CO(NH2)2. The salts
are removed in the same way.
428
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 429
Animal Nutrition. — Since there is continual loss, there must
be continual replacement. The animal resembles the plant, in
the fact that it can take up into its system only dissolved material.
It differs from the plant, however, in the fact that it is provided
with a wonderful laboratory in which insoluble substances are
changed into soluble ones. This is the digestive trad, consisting of
the mouth, stomach, and intestine. The production of soluble
substances of suitable composition is called digestion.
The processes are too complex for detailed treatment here. Only
a few typical illitstrations can be given. The principles concerned
have all been used and illustrated already, and many of the facts
are contained in previous chapters.
FfH^ds. — First, let us examine the table showing the percentage
composition of the edible portion of several articles of food:
Food material
Beef (\esi,n)
Cod
Eggs
MiE*
Butter
Cheese (cheddar) .
Oatmeal
Wheat flftur
Beans (dried)
Almonds
Maize (green corn)
Potatoes
Lettuce
Sugar
Water
Protein
Fat
Carbo-
hydrate
73.8
22.1
2.9
82.6
16.8
0.4
73.7
14.8
10.5
87.0
3.3
4.0
5.0
11.0
1.0
85.0
27.4
27.7
36.8
4.1
7.3
16.1
7.2
67.5
11.9
13.3
1.5
72.7
12.6
22.5
1.8
59.6
4.8
21.0
54.9
17.3
75.4
3.1
1.1
19.7
78.3
2.2
0.1
18.4
94.7
1.2
0.3
2.9
« • • «
• • • •
• • • •
100.0
Aflh
1.2
1.2
1.0
0.7
3.0
4.0
1.9
0.6
3.5
2.0
0.7
1.0
0.9
* The emulsified fat separates slowly as the cream; the protein (casein, oolloidally suspended
in the sldm milk) is coagulated by rennet and constitutes cheese; the carbohydrates (lactose, a
•usar) are then left in the water, along with inorganic salts.
We note, at once, that there is little more water in milk than in
cod; that the animal foods, except milk (carrying lactose, p. 401),
contain no carbohydrates; that potatoes and com when dried
430 smith's intermediate chemistry
are nearly all carbohydrate (starch) ; that lean beef when dry is
nearly all protein; that some seeds (wheat and beans) contain
almost no fat, some (oats) much more, and some (almonds and
nuts) a very large amount; and that lettuce and other leaves are
mainly water, with dissolved inorganic salts (valuable), contained
in a Ught framework of cellulose (non-digestible).
Digestion of Starch. — The carbohydrates, in most foods
which contain a large proportion of them, are mainly in the form
of starch. The exceptions are milk, sweet fruits, and sugar itself.
Starch is insoluble in water, and can not be directly absorbed.
But we have seen (p. 401) that, when boiled with a dilute acid, it is
hydrolyzed, giving glucose. When bread and potatoes are masti-
cated, an enzyme (p. 417), named ptyalin, contained in the saliva
(alkaline) turns a part of it, by hydrolysis, into a soluble sugar,
maltose. Later, in the small intestine, amylopsin completes this
process. Here also another enzyme, maltase, splits the mal-
tose into glucose. The glucose then passes through the intestinal
wall and so goes into the circulation, where most of it is oxi-
dized.
The cooking of starch (baked bread, boiled potatoes, etc.) breaks
up the grains and makes the mixing with the enzyme more perfect
and the digestion more rapid and complete.
Baking Powders. — The purpose of the powder is to generate
carbon dioxide in the dough. The bubbles of the gas are re-
tained by the sticky gluten of the flour. They expand when the
dough is baked, and give to it the open texture which, when the
bread is eaten, faciUtates access of the saliva to every particle.
Baking soda NaHCOs, if used alone, will give ofif, when heated,
half the carbon dioxide it contains (p. 367). The sodimn car-
bonate which remains in the bread, however, has an acrid taste.
By its action on the gluten in the flour, it gives also a yellow color
and an unpleasant odor. Finally, the carbonate of soda tends to
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 431
neutralize the gastric juice (acid) of the stomach and so to inter-
fere with digestion.
To obviate these difficulties sour milk (containing lactic acid)
is sometimes used in making the dough. Occasionally vinegar
(p. 420) is added. Most frequently a baking powder, containing
an acid substance along with the soda, is employed. The acid
substances contained in baking powders are almn (p. 469), acid-
phosphate of calcium or sodiiun, or potassiiun-hydrogen tartrate.
The last, known commonly as cream of tartar HKC4H4O6, acts as
follows:
HKC4H4O6 + NaHCOs -> NaKC4H406 + HgCOs -> H2O + CO2.
The sodium-potassium tartrate which remains is better knowi^
imder the name of Rochelle Salt.
It is important that the soda and the acid substance should be
used in the correct proportions, which can be calculated from the
equation. In commercial baking powders a little corn-flour is
added, to keep the particles of the other compounds apart and
prevent that gradual interaction which otherwise would be bound
to occur. The acid substance should also be somewhat insoluble,
so that, even when wet, it will not act upon the soda imtil ample
time has been allowed for complete mixing with the dough.
Bakers^ Bread. — The '^raising" of bakers' bread is effected by
adding yeast. The batch is " set " in a warm place for some hours
to permit the yeast to propagate and to act upon the sugar in the
flour. In this action, as we have seen (p. 417), carbon dioxide and
alcohol are produced. A little sugar, molasses, or malt extract
is added to the dough, to afford a larger supply of the sugar re-
quired for the production of the carbon dioxide.
The whites of eggs " raise '' cake, without the presence of any
soda, because of the expansion under heat of the bubbles of air
entangled with the albumen when the eggs are " whipped."
432 smith's intermediate chemistry
Fats and Oils. — The fats and oils found in the bodies of
animals, or pressed from the seeds of plants, are composed mainly
of various esters (p. 349). As such, they are formed by the inter-
action of a tribasic alcohol, glycerine C3H6(OH)3, with higher
acids of the paraflSn and olefine series of hydrocarbons, such as
the saturated acids CwHsi.COOH (palmitic acid) and C17H35.COOH
(stearic acid), and the unsaturated acid C17H33.COOH (oleic acid).
For example:
C8Hb(OH)8 + 3H(COO.CibH3i);=^ C8H6(COO.CibH8i)3 + 3H2O
glycerine palmitic acid glyceryl palmitate
The glycerine esters of the saturated acids are solid at ordinary
temperatures (fats), while those of the unsaturated acids are
liquid (oils). Beef suet is a mixture of about three-fom1;hs
glyceryl palmitate (palmitin) C3H5(C02CibH3i)3 and glyceryl
stearate (stearin) C3H6(C02Ci7H35)3, and one-fourth glyceryl
oleate (olein) C3Hb(C02Ci7H83)3. Hog lard contains about 40
per cent of the former and 60 per cent of the latter, and is
therefore softer. Butter includes the same esters, with about
14 per cent of water. When butter is dried, the remaining fat
contains about 8 per cent of glyceryl butyrate (butyrin)
C3Hb(C02C3H7)3. Olive oil contains 75 per cent of olein. Cot-
ton seed oil is similar in nature.
All these fats and oils contain, also, a small amount of the
free adds. They must not be confused with mineral oils like
petroleum, which are mixtures of hydrocarbons.
Oleomargarine is an artificial butter. It is made by straining
melted beef fat and allowing it to stand at 24°. Much of the
stearin crystallizes out and the remaining liquid (the " oleomar-
garine ") is pressed out and allowed to solidify. The solid is
finally mixed with a little of some oil, to render it softer, and is
churned with milk to impart the proper flavor. Although it
lacks the butyrin, the product is similar in chemical nature to
butter, and just as nutritious and wholesome.
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 433
Hydrolysis of Fats and Oils. — The chief chemical property
of the fats and oils, and in fact of all esters, is that each can be
decomposed^ or hydrolyzed, to give back the alcohol and acid from
which it is derived. Thus, when ethyl acetate is boiled with
water, it is slowly decomposed into ethyl alcohol and acetic acid:
C2H5(C02CH3) + H2O -> C2H5OH + HCO2CH3.
In the case of the fats and oils, if water alone is used, they must
be heated in a closed vessel, or autoclave, imder pressure so as
to secure a high temperature (about 200°) :
C8H5(C02CibH3i)3 + 3H2O -> CeHsCOH), + 3HCO2C16H81.
palmitin glycerine palmitic acid
When the mixture has cooled, the acid, which is insoluble in water,
forms a solid cake, while the glycerine is dissolved in the water.
With water alone as a hydrolyzing agent, however, the reaction
is slow and incomplete and, at the high temperatures which it is
necessary to employ, some destruction of fatty matter is apt to
occur. In the presence of dilute sulphuric acid as a catalyst, the
hydrolysis can be carried out much more satisfactorily and rap-
idly, even at 100°. Sulpho-derivatives of the fatty acids are still
more efifective as catalysts, since they are freely soluble both in
fats and oils and in water, and hence promote the miscibiUty of
the two layers (Twitchell process).
When tallow (beef fat) is treated in this way, the solid is a
mixture of palmitic, stearic, and oleic acids. The latter, being
liquid, is pressed out, and the solid material is used with paraffin in
making candles. The glycerine is separated from the water and
used in making nitroglycerine (p. 481) and in medicine.
Hydrogenation of Oils. — The market value of solid fats
is much higher than that of liquid oils. Many natural oils, in-
deed, possess disagreeable characteristics (taste, odor, etc.) which
render them totally unsuitable for edible purposes. They may
be converted into more appetizing edible fats, however, by
434 smith's intermediate chemistry
hydrogenation. At 200°, in the presence of finely divided
nickel as a catalyst, these unsaturated compounds take up hydro-
gen and become saturated, the oil hardening to a fat in the
process.
Many substances which were formerly waste products, such as
cotton-seed oil, are now treated in enormous quantites in this
way. The hydrogenation of low grade oils (for example, fish oU)
is also of great industrial importance in the manuf actm^ of soap
(see p. 438) and candles.
Digestion of Fats. — At body temperature, the fats and oils
present in foods are all insoluble in water, and therefore can-
not be directly absorbed into the system. But fats, if already
emulsified (p. 110), as in milk, are hydrolyzed by a lipase (enzyme
for fat) in the gastric juice of the stomach, and are decomposed
into the acid and glycerine (p. 433). Fat in larger masses is
hydrolyzed by lipases in the bile and here the acid (insoluble in
water) is dissolved. The acid and the glycerine then diffuse
through the intestinal wall and finally recombine to form fat in
the blood. Some of this fat is deposited in the tissues and some is
oxidized (giving muscular energy and heat).
Cooking (application of heat) does not affect the digestibility of
fat. However, when fat is heated too strongly, the beginning of
destructive distillation produces unsaturated compounds. These
are intensely irritating to the digestive organs — as the way their
vapors bring tears to the eyes would lead us to expect.
Digestion of Proteins. — The proteins contained in foods,
of which the white of an egg (albiunen) is a typical example, are
not affected by saliva, but, when mixed with the gastric juice of
the stomach, they are changed by the free hydrochloric acid it
contains into syntonin. This in turn is hydrolyzed by the pep-
sin (enzyme), also contained in the gastric juice, into peptones
which are soluble in water. These changes, only partly carried
ANIMAL LIFE AND ANIMAL PRODUCTS. POODS 435
out in the stomach, are completed in the small intestine by the
trypsin of the pancreatic juice, and the peptones (or amino-acids
into which they are spUt) pass through the intestinal wall into the
circulation. The casein of milk, being in colloidal suspension,
is completely hydrolyzed to peptones in the stomach.
When heated, as in cooking, the proteins do not behave alike.
Some, like albumen (white of egg) become coagulated, though
probably not less digestible. The same is true of the blood pro-
teins (haemoglobin, etc.) of beef. On the other hand, the connec-
tive tissue of meat (chiefly collagen) is insoluble in cold water, but
in hot water goes into colloidal suspension as gelatine. It is
therefore softened by judicious roasting (under-done meat), pro-
vided the operation is not carried so far (over-done meat) that the
water in the meat is largely evaporated.
Fuel Value. — While food is needed primarily to replace the
material which is continually eliminated from the system, the
organism also requires energy to maintain the routine motions of
the heart, intestines, limgs, and other organs, and the normal
muscular tension, as well as the movements of the muscles in
walking and working. If the heat derived from routine changes
is not sufficient to maintain the temperature (37° C.) of the body,
then additional food material is oxidized by the system for this
specific purpose (compare p. 426). The values of foods are
therefore conveniently estimated in terms of the heat they pro-
duce when burned — their fuel values.
The average fuel values, as measured in the calorimeter, with
certain necessary corrections, and expressed, as is usual in this
work, in large calories* per gram, are: Carbohydrates 4 Cal., fats
9 Cal., proteins 4 Cal. The fuel values per pound (= 453.6 g.)
are 453.6 times greater; Carbohydrates 1800 Cal., fats 4080 Cal.,
proteins 1800 Cal.
* One large calorie (1 Cal.) is equal to one thousand small calories (1000
cal.)i as hitherto defined (p. 162) and used.
436
smith's intermediate chemistry
Normal Diet. — There is much uncertainty, as yet, in regard
to the best choice of foods, in respect to the exact distribution in
kind and quantity. We know, however, that life cannot be
maintained on one kind (say, sugar or gelatine) alone. A mixed
diet is necessary. In general, it appears that 100 g. of proteins
(giving 4 X 100 Cal.) per day, and a suflScient amount of other
foods to bring the total fuel value up to 2200 Cal. per day, is
sufficient for a person leading a strictly sedentary life. For work
involving physical exercise, larger values, up to about 3800 Cal.,
are required.
From the data given in the table (p. 429) the fuel value of 100
g. of each kind of food can easily be calculated.
Fuel Values and Prices of Foods, — If the current prices are
considered, one can also readily calculate the fuel value obtainable
for a given sum of money invested in each kind of food. Thus:
lean beefsteak contains 22.1 per cent of protein, or 0.221 poimds
per pound of meat. The fuel value of this protein is 0.221 X 1800,
or 398 Cal. per pound.
%
Price* per
1 Pound,
Cents
Fuel Values per 1 Pound
Cal. per
Protein
Fat
Carbohyd.
Total
10 Cents
Steak
25
24
5
5
40
2.5
25
5
398
266
290
249
378
40
500
> • •
118
428
294
61
2240
4
1500
• • • •
i2i5
1309
311
331
74
1800
516
694
1799
1619
2929
375
2074
1800
206
Ekits
290
Oatmeal
3600
Flour
3240
Almonds
732
Potatoes
1500
Cheese
830
Suear
3600
* The prices vary gr atly with the quality, the season of the year, the demand, the 8iu>ply, ^c
Vitamins, — A diet may be carefully balanced with respect
to carbohydrates, fats and proteins, and yet lead to malnutrition
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 437
through deficiency in vitamins. These are substances present
in minute quantities in most fresh or unsteriUzed foods, particu-
larly in milk and green vegetables, without the aid of which cer-
tain parts of the animal mechanism either cease to develop or
fail to perform their functions entirely. The three known types
are:
(1) Fat-soluble vitamin A, plentiful in milk, butter, the yolks
of eggs, cod-liver oil and the leaves of green plants, but not found
in grains, sugars, or refined vegetable fats and oils. This vitamin
is needed to promote the growth of children; its absence leads to
rickets and a disease of the eye called xerophthalmia.
(2) Water-soluble vitamin B, plentiful in the outer hull of
grains, beans, green leaves and fruit and yeast, but not in the
kernels of grains, such as polished or milled rice. Deficiency of
this vitamin in the diet leads to boils and skin eruptions, and in
extreme instances to beriberi.
(3) Water-soluble vitamin C, plentiful in citrous fruits, to-
matoes, cabbage, lettuce and other fresh fruits and vegetables.
The value of this vitamin lies in the prevention and cure of
scurvy. It is easily destroyed, except in the case of acid foods,
by heating, drying or ageing. Since milk is nevtral, infants fed
entirely on pasteurized milk are almost sure to develop a mild
case of scurvy unless the diet also includes a little orange juice
or some other source of water-soluble vitamin C. Owing to the
acid nature of tomatoes, even the canned product is rich in this
vitamin. A knowledge of this fact was put to good use by the
authorities of the British armies in Mesopotamia and Palestine
during the war, when fresh fruits and vegetables were unobtain-
able.
Animal Products.^ Many valuable products, apart from
foods, are derived from animal life. The use of animal products
in fertilizers (calcium phosphate from bones, nitrogen compounds
from manure, etc.) has already been discussed (Chapter XXXV).
I
438 smith's intermediate chemistry
When bones or dried blood are subjected to destructive dis-
tillation (compare p. 420), the residue consists of animal charcoal.
The charcoal from bones (Jbone black) contains 90 per cent of
mineral matter, largely calcium phosphate, and only 10 per cent
of carbon. Animal charcoal, being a very active adsorbent
(p. 421), is used in stigar refining.
The basis of woolj and of hair fibers in general, is a protein
called keratin. The silk fiber is also of animal origin, but differs
very widely from wool in its structure and properties. As spim
by the silkworm in the preparation of its cocoon, it consists of
two filaments composed of a protein called fibroin surroimded and
cemented together by a gluey substance known as seridn or silk-
gum.
Gins, an impure form of gelatine (p. 435), is obtained from the
skin and bones of animals by extraction with water imder pressure.
Leather is prepared from the hides by tanning (see p. 533).
Casein, a proteia contained in milk and precipitated therefrom
by dilute acids, has recently found many interesting applications.
In the modified form of cheese, of course, it has long been of value
as a food. Mixed with various alkalies it gives glices, cements
and putties. It is used as an ingredient in paints, and for sizing
or enameling paper. Paper bottles are made waterproof by im-
pregnatiag them with caseia and then exposing them to the vapor
of formaldehyde (p. 348). Galalith (milk-stone) is a plastic,
harder than celluloid (p. 480) and non-inflammable, made by
the same hardenuig action of formaldehyde on caseia.
Soap. — When fats are hydrolyzed by heating with a solution
of caustic soda NaOH, instead of water, the sodium salts of the
adds are obtained. These sodiimi salts are known as soaps and
the operation is called saponification:
C8H5(C02CisH3i)3 + 3NaOH -> C3H5(OH)8 + 3Na(C02Ci6Ha)
palmitin glycerine sodium palmitate
(a soap)
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 439
The sodium palmitate or other soap is soluble in the water. When
common salt is added, however, the soap coagulates and separates
into a floating layer which solidifies on cooling.
Soft soap is made with potassiiun hydroxide, and is composed
of the potassium salts of the organic acids.
Manufacture of Sitap. — Soap is made in large iron caldrons.
These contain closed steam coils for heating, and pipes delivering
live steam, when needed, for stirring the mass. The fat is mixed
with caustic soda solution containing, at first, only about one-
fourth of the total amoimt of the base that the above equation
requires. When, after heating and stirring, a uniform mixtiu'e
has been produced, the rest of the alkali is added gradually and
the heating is continued imtil the reaction appears to be complete.
Salt is now dissolved in the mixtiu'e and the soap separates as a
curd. The curd floats, leaving the " spent lye," containing the
salt solution and much of the glycerine, in the lower layer. This
process is called salting out.
When the process stops at this point, the upper layer is known
as curd soap, and may be dipped out and allowed to cool and
solidify. Most " Marseilles " soaps are ciu'd soaps made in this
way. A large part of the imported " Castile " soaps are of the
same kind.
Cuid soaps contain salt, glycerine, adhering lye, and other
impurities. To prepare a purer soap, the spent lye is run oflF,
dilute brine is added to the soap, and the curd and brine are stirred
up. When separation has again occurred, the brine is run oflf
and the process repeated. Finally, some water is added, and
steam is run in imtil the curd mixes completely with the water.
When the solution stands, it " settles," that is, resolves itself once
more into two distinct layers. The upper layer is called settled
soap. The washing with brine and temporary dissolving in
water remove the impurities, and hence settled soap is the pm-est
variety. The greater part of the soap made in the United States
440 smith's INTERBiEDIATE CHEMISTRY
and in Great Britain, and much of that made in other eoimtries,
is of this kind.
The qualities of soaps are varied by adding " fillers," such as
sodium carbonate, borax, or sodium silicate. Soap powders are
often made of ground soap mixed with sodium carbonate. Dyes
and perfumes are sometimes added to soaps. Air bubbles are
mixed with the soap, by beating, to give the floating varieties.
Soap for scoimng contains fine sand. Transparent soaps are
made by dissolving the soap in alcohol, or by the addition of
glycerine or sugar.
Chemical Properties of Soaps. — The soaps, being sodium
salts, dissolve in water and have the icsiuil properties of salts. Thus,
when an acid is added, double decomposition takes place:
NaCCOaCigHsi) + >HC1 -> NaCl + HCCaCisHai) i .
sodium palmitate palmitic acid
The adds, being insoluble, are thrown down as white precipitates.
When other salts are added, similar actions occur. Thus with
calcium chloride solution, calciimi palmitate is formed, and being
insoluble, is precipitated:
2Na(C02Ci5H3i) + CaCk -> 2NaCl + CaCCOjCiBHsOs i.
This action is important in connection with " hardness " in water
(p. 389).
Colloidal Suspension. — We have seen (pp. 109, 400) that
starch can be suspended in water in such a fine state of division
that the liquid is transparent, and runs through a filter without
leaving any solid on the paper. Yet this " suspension " lacks
some of the characteristic properties of a solution.
The simplest proof, that this is a case of imitation solution, and
not of true solution (molecular dispersion, p. 108), is obtained
by examining the liquid with the ultra-microscope. A converg-
ANIMAL LIFE AND ANIMAL PRODXTCTS. FOODS 441
ing beam of strong light is sent through the liqmd horizontally
(Fig. 107), and the illuminated place is viewed
from above through a microscope. When the
room is dark, and the only light comes from
the horizontal beam, a colloidal solution
shows minute points of light. A true solu-
tion — such as one of common salt or of
alcohol — remains perfectly dark. The points
of light are produced by myriads of suspended
partickSy which although extremely minute, are of far larger than
molecular dimensions. Solutions of soap, gelatine, and many
dyes, blood serum, and innumerable other liquids contain such
suspended particles.
The particles of a colloid (like starch), when viewed in this way,
show a continual unordered, zig-zag movement (Brownian
movement), which is more rapid the smaller the particles, and is
due to the impacts of the molecules of the solvent.
Other Properties of Colloidal Suspensions. — The freezing-
point, boiling-point and vapor pressure of a liquid containing a
colloid in suspension are all practically identical with those of the
pure solvent. This is quite different from what we have seen
to be the case with true solutions (p. 117), and indicates that the
fraction of the " solute " particles present is substantially zero.
The unit particles of suspended material, indeed, are complex
aggregates so much larger than the ultimate molecules into which
true solutes are broken up that the number of them present, com-
pared with the nimiber of solvent molecules, is entirely negligible.
When a " solution " of a colloid is placed in a " diffusion-
shell " (test-tube shaped) of parchment, surrounded by pure
water, none of the colloid escapes through the minute pores of the
shell. Ordinary, non-colloidal solutes do escape, more or less
rapidly (salt rapidly, sugar slowly, see p. 407). In this way a
mixture of coUoid and nonrCoUoid (say, starch and salt) can he
442 smith's intermediate chemistry
separated, if the water surrounding the shell is replaced by pure
water at intervals until all the non-colloid has b^n removed.
This method of separation is called dialysis.
Finally, matter in colloidal suspension may be coagulated (or
flocculated) by addition of electrolytes or other colloids, and is
then precipitated.
Colloidal Matter in Stmp Solutions. Explanation of
Salting Out. — Soaps, being salts of weak acids, are some-
what hydrolyzed in solution. Letting R stand for the hydrocar-
bon part of the acid radical:
Na(C02R) + H2O ^ H(C02R) + NaOH.
The free acid HCO2R thus produced combines with the original
salt NaC02R to form an acid salt NaH(C02R)2. This add salt
is a colloidal suhstance, and exists in colloidal suspension in the
soap solution, in equiUbriimi with the ions and molecules of the
original salt and the NaOH.
The capacity for being coagulated and precipitated, which is
characteristic of colloidal matter, is shown very clearly by soap
solutions. Most sodium salts will coagidaie a soap solution and
precipitate the soap as a curd. The acid salt NaH(C02R)2 seems
to collect (adsorb) and carry down with it the most of the sodium
hydroxide. As both of the substances on the right of the equa-
tion (above) are thus precipitated, the equilibrium is displaced
to the right, and the precipitation becomes complete. This
explains the process of " salting out " (p. 439) which plays so large
a part in the manufacture of soap.
Causes of the Cleansing Action of Soap. — The chief use of
soap solution is in removing grease and dirt from yam, cloth, or
clothing, and from woodwork and kitchen utensils. Soap solu-
tion has two more or less distinct properties, one of which enables
it to remove oil or grease (viscous, insoluble liquids), and the
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 443
other of which enables it to remove dust and dirt (largely minute,
solid particles of carbon — soot). The former is its emulsifying
power J the latter is probably connected with its nature as a col-
loidal suspension.
The Emulsifying Action of Soap. How Soap Removes
Crease. — When an insoluble oil, such as kerosene or lubricating
oil, is shaken with water it is divided into minute droplets sepa-
rated by water from one another. When the shaking ceases, how-
ever, the droplets begin to run together and soon the oil and
water have separated once more into two layers of transparent
liquid. When very dilute soap solution and oil are shaken to-
gether, however, the droplets do not run together, but remain
permanently suspended. The mixture is opaque and more or
less viscous. Such a permanent mixture of two insoluble liquids
is called an emulsion (compare p. 110).
Soap solution, when rubbed on oily or greasy goods, emulsifies
the grease, converts it into droplets, surrounded by soap solution
and separated from the cloth, and so permits it to be washed off.
In mayonnaise dressing, which is a thick, almost solid, per-
manent emulsion, the oUve oil is emulsified by the colloidal matter
of the yolks of the eggs which have been dissolved in the vinegar.
How Soap Removes DirU — When a solution containing
colloidal substances, such as many dyestuffs and organic coloring
matters, is shaken with finely pulverized charcoal, the colloidal
substance adheres to the surface of the powder (is adsorbed) and
the Uquid is consequently decolorized (see p. 422).
When dilute soap solution is shaken with infusorial earth and
the mixture is filtered, the clear liquid is found to have been de-
prived of the soap. The soap is evidently precipitated (adsorbed)
on the surface of the particles of the soUd.
Clearly there is a tendency to cohesion between the colloid in
a soap solution and the particles of a fine powder. When there is
444 smith's intermediate chemistry
much of the powder, and little of the soap m solution, the powder
takes the soap out of the solution. When, however, there is
much of the colloid in the form of soap solution, and little of the
solid, and that very finely divided, the same tendency to adsorp-
tion exists, ordyy in this case, the colloidal particles carry off the
powder. In short, the dirt is removed by adsorption into the
solution.
Possible Objections to the Foregoing Explanation. — For-
merly soap solution was supposed to remove grease (and soot?)
because of its slight alkaline reaction, due to hydrolysis. This
explanation must be given up because: (1) an alkali so dilute that
it exists in equiUbrium with the free fatty acid, can not possibly
saponify the ester contained in a grease spot. (2) Pure alkali of
the same concentration (or stronger) has no more emulsifying
power than water. Such an alkaline solution will indeed emulsify
an animal or vegetable oil (cod-liver oil, cotton oil, castor oil), but
it does so by interacting with the free fatty acid always present
in such oil (p. 432) B,nd forming therefrom a soap. Such an alkaline
solution does not emulsify kerosene, although soap solution does.
The emulsifying agency is always a soap. (3) Very dilute alkali
has no more effect upon soot than has water — but soap solution
takes clean (greaseless) soot instantly into permanent suspension.
The power of forming an emulsion depends, theoretically, upon
the abnormally low surface tension of dilute soap solution. Very
dilute alkali has the same high surface tension as has pure water.
*
Exercises. — 1. Make equations for the formation: (a) of malt-
ose from starch; (b) of glucose from maltose.
2. Make a connected statement showing the stages in the
digestion of mUk.
3. Why does fat appropriately form a larger proportion of the
diet in the Arctic regions than elsewhere?
4. Give the weights of carbohydrate, protein, and fat which
ANIMAL LIFE AND ANIMAL PRODUCTS. FOODS 445
would supply a menu, such that the total food value was 3000
Cal., and that 100 g. of protein was included, and the remaining
fuel value was divided equally between carbohydrates and fats.
5. Calculate the calorific value of 1 kilog. of: (a) wheat flour,
(b) oatmeal.
6. (a) In what proportions by weight should baking soda and
cream of tartar be used in raising bread? (b) What is the objec-
tion to using too large a proportion of baking soda? (c) Why
must baking powder be kept in a dry, cool place?
7. Why does vinegar liberate carbon dio;dde from baking soda?
Make the equation for the action.
8. Make equations for the formation of: (a) ethyl formate;
(b) glyceryl formate; (c) ethyl stearate.
9. Make equations for the action of superheated water on:
(a) stearin, and (b) olein; and for the action of caustic soda on:
(c) stearin, and (d) ethyl acetate.
10. Siunmarize the facts which show soaps to be salts.
11. (a) Write the full equation for the hydrolysis of sodium
palmitate. (b) What reaction (acid or alkaline) should soap
solution possess, and why (p. 369)?
12. Why does " French dressing " (vinegar, salt, and olive oil)
give an emulsion, which is much less durable than mayonnaise
dressing?
CHAPTER XXXVIII
MAGNESIUM AND ZINC. IONIC EQUILIBRIA
We shall now return to a consideration of the metallic elements.
We can most easily remember magnesium and zinc by the facts
that they are silver-white metals with a markedly crystalline
structure, and that they displace hydrogen from dilute acids. In
these respects they resemble aluminium, but the latter is trivalent
in all its compounds, while the present two elements are bivalent
exclusively (see Periodic System).
Magnesium Mg
Occurrence. — Magnesium carbonate is foimd in dolomite
CaCOsjMgCOa, a common rock, and in small amounts as mag-
nesite MgCOs. The sulphate and chloride are foimd at Stassfurt.
Several natural silicates of magnesiimi, such as meerschaum,
asbestos, talc or soapstone, and olivine, are familiar minerals.
Asbestos, a fibrous material, is used in making fireproof cloth
and cardboard. Soapstone is made into sinks and table tops for
use, for example, in laboratories.
The Metal Magnesium. — The metal is made by electrolyzing
a molten mixture of magnesiiun, potassium, and sodium chlorides.
A carbon rod forms the anode, and the iron crucible the cathode,
on which the metal collects in globules. The mass, when cold,
is broken up and the metal is recast in bars. The metal can be
drawn, through a die, into ribbon or wire.
Magnesium ritsts in the air, gradually crumbling to a white
powder of a basic magnesium carbonate. It burns in air, with a
brilliant white light, producing a mixture of the oxide MgO and
nitride Mg3N2 (see argon, p. 297). Magnesium filings, mixed
446
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 447
with potassium chlorate, give flash-light powder. Signal lights
are made of shellac, barium nitrate and magnesimn powder.
Oxide MgO and Hydroxide Mg(0H)2. — The oxide is made
by heating magnesimn carbonate, and is therefore called calcined
magnesia. Being very infusible (" refractory '') it is used in
lining electric furnaces. On accoimt of its poor heat conduc-
tance, it is also employed very extensively for insulating pipes
and boilers, and so reducing heat losses (86% Magnesia). It
combines slowly with water, giving the hydroxide Mg(0H)2,
which with water gives a mortar that hardens under the action
of the carbon dioxide of the air (see mortar, p. 386). The oxide
is basic, and with acids gives salts by double decomposition.
Magnesium hydroxide Mg(0H)2, being insoluble, is easily pre-
cipitated by adding sodimn hydroxide to a solution of a salt of
magnesium. When moistened and mixed with a little magnesimn
chloride, it sets to a h^rd basic chloride of variable composition.
The mixture, to which sawdust is sometimes added, is used as a
plaster in house decoration.
Salts of Magnesium. — Magnesium carbonate MgCOs is
foimd in natiu^. That made by precipitation is a basic carbonate
3MgC08,Mg(OH)2, magnesia alba, which is used in tooth powder
and for polishing silver.
Magnesium sulphate MgS04 is commonly sold as the hepta-
hydrate, MgS04,7H20, Epsom salt. It is found in the salt depos-
its and in many aperient mineral waters. Thus Hunyadi water
contains Uttle beside 47 g. Epsom salt and 52 g. sodimn sulphate
(Na2SO4,10H2O) and 1 g. sodium bicarbonate per liter. The salt
is used for loading cotton goods and as a purgative,.
Magnesium chloride MgCl2 is found in sea water and in some
natural waters. It is very deUquescent (p. 118) and, being present
in impure table salt, causes the latter to cake or even become moist
in damp weather. Addition of a very little sodiiun bicarbonate
448 smith's intermediate chemistry
to the salt remedies this difficulty. Magnesium cnloride is a very
objectionable form of hardness in water, because hot water par-
tially hydrolyzes the salt and liberates hydrochloric acid, which
attacks and corrodes the iron of the boiler and tubes. Hence
sea water can not be used in marine boilers.
Zinc Zn
Occurrence and Manufacture. — Zinc is found as zinc
blende ZnS (in large amounts in Missouri) and smithsonite ZnCOs
(Spain and U. S.).
In the case of the carbonate ore, the oxide ZnO is first obtained
by heating. When zinc blende ZnS is the ore, it is crushed and
pulverized, and then roasted (p. 258) to remove the sulphur and
leave the oxide:
2ZnS + 3O2 -> 2ZnO + 2SO2.
The ore is fed in at the top of a huge, box*like furnace (Fig. 108,
diagrammatic) through which rush the flames and heated gases
from fuel gas burning with an excess of air.
The ore is turned over and gradually displaced
forward by moving rakes imtil, at the end, it
drops to the next level. Here it is raked in the
opposite direction, imtil it falls to the third
Fig. 108 j^^^j .pj^^ ^j.^ collects at the bottom fully
oxidized, while the sulphur dioxide in the gases is made into
sulphuric acid. The oxide from either ore is then reduced by
heating with powdered coal:
ZnO + C->Zn + CO.
This treatment of zinc ores should be carefully considered,
Since orefe of most metals consist of the carbonate, sulphide, or
oxide of the metal, these steps are common to most metallurgical
processes. In the case of other metals, only the forms of the
furnaces and other details vary.
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 449
In the case of zinc, because it is a volatile metal, the heating of
the mixture of oxide and coal is conducted in retorts, from which
the metal issues as vapor. The mixture is placed in fire-clay
cylinders (4 to 5 ft. long), which are arranged in several tiers in an
oblong, gas-heated furnace (Fig. 109). A fire-clay receiver is
Fig. 109
luted on with clay. The carbon monoxide bums with a blue flame
at the nozzle of each receiver, while the zinc condenses to liquid
within it. From time to time the liquid metal is raked out into
a traveling iron pot, from which it is poured into moulds.
Properties and Uses* — The metal is silvery and crystalline.
At 120 to 150° it can be rolled into sheets between hot rollers, at
200 to 300° it becomes brittle, at 419° it melts, and at 950° it boils.
The density of the vapor shows it to be monatomic. Zinc vapor
bums with a bluish flame, giving ZnO. In air the metal does not
rust, being protected by a non-porous coating of a basic carbonate
which adheres closely to the surface.
Sheet zinc is used for gutters and cornices. Iron is coated
(galvanized) with zinc by thorough cleaning with dilute sulphuric
acid or the sand blast, and dipping in melted zinc. Iron netting,
corrugated iron for sheds and roofing, and iron gutters, tanks, and
pipes are coated with zinc, either in this way, or by electroplating
(see p. 515). Sherardized iron is made by covering the article
450 smith's intermediate chemistry
with zinc dust and baking it. The zinc protects the iron, primar-
ily because it excludes the air from the surface, and secondarily
because, even when the coating is broken, the zinc, being the
more active metal of the two (p. 54), is rusted first. Zinc is used
also in extracting silver from crude lead (p. 519), as the active
metal (anode) in batteries, and in several alloys {e.g. Babbitt's
metal, p. 326, brass, and German silver). In the laboratory
granulated zinc, made by pouring the melted metal in a thin
stream into water, and zinc dust (impure, contains ZnO), are the
forms commonly employed.
Zinc Oxide ZnO and Hydroxide Zn(0H)2. — The oxide is
made by burning zinc vapor in air. It is yellow while hot, and
white when cold. Mixed with oil, it is used as a paint (Chinese
white). It has less covering power than has white lead paint (4
coats of the former equal 3 of the latter), but it does not darken
from exposure to hydrogen sulphide in the air (ZnS is white, PbS
black).
Zinc hydroxide Zn(0H)2 is formed by precipitation. Both the
oxide and hydroxide are weakly basic, and give salts with acids.
But with respect to strong bases they are weakly acidic, dissolving
for example in excess of sodium hydroxide, giving sodium zincate
Na.HZn02.
Zinc Chloride ZnCl2. — The chloride is formed in the action
of zinc or zinc oxide on hydrochloric acid. It is a white, deliques-
cent solid. Its aqueous solution gelatinizes cellulose and dissolves
it (p. 398), and thus is used in parchmentizing paper and in im-
pregnating wood to prevent decay. The aqueous solution is acid
(hydrolysis) and is used for cleaning the surface of metals before
soldering. Solder " runs " on a hot brass, copper, or lead surface,
provided the latter is clean, and adheres perfectly when cold. But
it does not dissolve oxides, or melt them, and therefore cannot
even reach the surface of the metal, much less adhere to it, if the
slightest tarnish is present.
MAGNESIXm AND ZINC. IONIC EQUILIBRIA 451
Other Compounds. — Zinc sulphate, ZnS04,7H20, is made
by the action of sulphuric acid on zinc or zinc oxide. It is used
in preserving hides and as a mordant in cotton printing (see
dyeing, p. 475). Zinc sulphide ZnS (white) is precipitated prac-
tically completely when a solution of ammonium sulphide (NH4)2S
is added to a salt of zinc:
ZnS04 + (NH4)2S -> ZnS i + (NH4)2S04.
When hydrogen sulphide is employed instead
ZnS04 + H2S <± ZnS + H2SO4
only a small part of the zinc is precipitated. The cause of this
difference will come up for discussion in a subsequent section (p.
469).
A mixture of zinc sulphide and barium sulphate BaS04, pre-
pared in a special way, is called lithopone. Used as a white
pigment, it has greater covering power than has white lead and
is, besides, non-poisonous, .
Ckidmium. — Aside from the rare mineral greenockite CdS,
this element is found only in small amounts (about 0.5 per cent),
as carbonate and sulphide, in the corresponding ores of zinc.
During the reduction, being more volatile than zinc, it distils
over first (b.-p. 778°). The metal is white, and is much more
malleable than is zinc. It melts at 320°.
It displaces hydrogen from dilute acids, but is itself displaced
from solutions of its compounds by zinc, since it is less electro-
positive. It is used in making fimble alloys.
Ionic Equilibria
In view of the importance of ionic actions in the chemistry of
the metals, we must now consider more closely the subject of
ionic equilibria. The whole basis for this exact consideration
has already been supplied, and only more specific application of the
452 smith's intermediate chemistry
principles is demanded. The basis referred to, which should now
be re-read as a preliminary to what follows, is contained in the
discussion of chemical equilibriimi in general (pp. 230-241).
One important point may be briefly recapitulated here, as an
aid to the reader.
Equilibrium in Reversible Reactions. — We saw (p. 237)
that the point of equilibrium in the dissociation of phosphorus
pentachloride:
PCl6 ;=i PCI, + CI,
is dependent upon the molecular concentrations of the reacting
substances, and may be represented by the equation:
[pq] ■ [CI2] ^^
[PCI5]
In this equation the quantities within the brackets represent
the equilibrium concentrations of the respective substances, and
K is the equilibrium constant of the reversible reaction. How-
ever much we may vary the molecular concentration of any one
substance at any fixed temperature (say, by adding excess of
CI2 to the gaseous mixture), the only effect will be a re-adjustment
of the equilibrium point such that the final ratio [PCls] • [CU] /
[PClfi] is unchanged. In other words, the equilibriiun constant
K is a fixed quantity.
Application to Ionization. — Exactly the same principles
can be applied to the ionic dissociation of an electrolyte. The*
behavior of acetic acid, a weak, slightly ionized acid, will serve
as an illustration.
In normal solution (60 g. in 1 1.) acetic acid is only 0.004 ionized
(p. 190), so that, in the equation for the equilibrium,
(0.996) HC2H802i=5H+ (0.004) + CjHsOr (0.004),
the relative proportions are as shown by the nimibers in paren-
theses. If the whole of the acid (60 g.) were ionized, there would
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 453
be 1 g. of hydrogen-ion per liter. Yet, even in the much smaller
concentration actually present (0.004 g. per liter), the acid taste of
the H"*" and its efifect upon indicators can be distinctly recognized.
If, noWy solid sodium acetate is dissolved in the solution, the liquid
no longer gives an add reaction with one of the less delicate indica-
tors, like methyl orange (an organic dye which turns pink in the
presence of a fair amoimt of hydrogen-ion). The explanation is
simple. Sodium acetate is highly ionized. It gives, therefore,
a large concentration of acetate-ion to a liquid formerly contain-
ing very little. This causes a greatly increased union of the H"*"
ions and C2H802~ ions to occur, and the former, being already
very few in number, disappear almost entirely. Hence the solu-
tion becomes, to all intents and purposes, neutral. There is no
less acetic acid present than before, but the concentration of
hydrogen-ion is very much smaUer.
Formulation and Quantitative Treatment of the Case of
Excess of One Ion. — If [H"*"] and [C2H3O2""] represent the moUc-
vlar concentrations of hydrogen-ion and acetate-ion, respectively,
and [HC2H8O2] that of the acetic acid molecules at equilibrium,
then :
[H^] X [C2H302i ^ ^
[HC2H8O2]
The value of the ionization constant K is unchanged, whether
the concentration of the solution of acetic acid is great or small,
and even when another substance with a common ion is present.
In the latter case, [C2H302~] and [H+] stand for the whole concen-
trations of each of these ionic substances from both sources.
Now, in normal acetic acid [Hi = 0.004, [C2H3O2-] = 0.004 (for
the number of each kind of ions is the same), and [HC2H3O2] =
0.996, practically 1. Substituting in the formula:
aooi^MQ4 ^ K (^ 0,046).
454 smith's intermediate chemistry
When, however, sodium acetate is dissolved in the liquid until the
solution is normal in respect to this substance obo, which is 0.528
ionized in normal solution, the following additional equilibrium
has to be considered:
(0.472) NaC2H302 <=± Na+ (0.528) + C2H8O2- (0.528).
The concentration of acetate-ion from this source is 0.528, so that,
in the mixture of acid and salt, the concentration of acetate-ion
[C2H3O2-] will be 0.528 + 0.004 = 0.532, or 133 times larger
than in the acid alone. Hence, in order that the ionization con-
stant K may recover, as it must, its original value, [H+] must
be diminished to something like xir of its former magnitude.
That is, [H+] will become equal to about 0.00003,
0.00003 X 0.532 ^. aaiax
z = K (= O.O4I0),
the rest of the hydrogen-ion uniting with a corresponding amount
of the acetate-ion to form molecular acetic acid. The concen-
tration of this increases only from 0.996 to 0.99997; in other words,
so much is already present that its concentration is practically
unchanged. The chief effect of adding this amount of sodiiun
acetate therefore is, as we have seen, to reduce the concentration
of the hydrogen-ion below the amount which can be detected by
use of an indicator like methyl-orange.
This efifect is of course reciprocal, and the ionization of the
sodium acetate will be reduced also. But the acetate-ion fur-
nished by the acetic acid is relatively so small in amount that the
efifect it produces on the ionization of the salt is imperceptible.
Although, therefore, acetate-ion and hydrogen-ion must dis-
appear in equivalent quantities, for they unite, there is, however, so
much of the former that the loss it sustains goes unremarked, while
there is so little of the latter that almost none of it remains.
When substances of more nearly* equal degrees of ionization are
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 455
used, both effects are relatively inconspicuoiLS. It is altogether
impossible, for example, to reduce the concentration of the hydro-
gen-ion given by an a^ive acid like hydrochloric acid, by addition
of a salt containing a common ion, like sodimn chloride, below
the limit at which indicators are afifected, for there is no way of
introducing the enormous concentration of the other ion which
the theory demands.
With more crude means of observation than indicators afiford,
effects like this last, however, may sometimes be rendered visible.
This is the case with cupric bromide solution, to which potassium
bromide is added. The blue of the cupric-ion disappears from
view, while much cupric-ion is still present, because the brown
color of the molecular cupric bromide covers it up completely.
The color changes discussed on p. 235 in connection with
the reaction FeCla + 3NH4.CHS ^ Fe(CNS)3 + 3NH4CI also
illustrate this point. The red color is due entirely to undis-
sociated ferric thiocyanate. The reader should re-examine this
equiUbrium carefully, writing the equation in full ionic form.
He will then be in a position to understand why the addition of
any ferric salt or of any thiocyanate to a given mixture will deepen
the red color, while the addition of any chloride (except ferric
chloride) or any anmionium salt (except ammonium thiocyanate)
will tend to lighten it.
Special Case of Saturated Solutions. — The commonest as
well as the most interesting application of the conceptions devel-
oped above is met with in connection with saturated solutions,
especially those of relatively insoluble substances.
The situation in a system consisting of the saturated solution
and excess of the solute has been discussed already (read p. 120).
In the case of potassium chlorate, for example, we have the follow-
ing scheme of equilibria:
KClOa (solid) <± KCIO3 (dslvd.) ^K+ + ClOa"".
456 smith's intermediate chemistry
If we apply to the two reversible reactions represented in this
scheme the principles which have been developed above, we obtain
the two equations:
[KClOa (dissolved)] _
[KCIO3 (solid)] ^' ^^^
[^-^J ' t^^-J ^ K. (2)
[KClOa (dissolved)]
where Ki and K2 are the equilibrium ^constants of the two reac-
tions. By multiplying equation (1) by equation (2), we obtain a
third relationship:
[K+] [CIO3I
[KCIO3 (soUd)]
= K1.K2 (3)
Now [KCIO3 (solid)], the concentration of KCIO3 in the solid state,
is invariable at any fixed temperature. Equation (3) conse-
quently reduces to the form:
[K-^] • [CIO3-] = K1K2 • [KClOs (soUd)] = a constant quantity.
In other words, in a saturated solution, the product of the molar
concentrations of the ions is a constant. This product is called the
solubility product of the substance. The law of the constancy
of the ion-product in a saturated solution is one of the most useful
of the principles of chemistry. It enables us to explain all the
varied phenomena of precipitation and of the solution of pre-
cipitates in a consistent manner, as will be seen below. Upon
the accuracy of the law of the constancy of the solubility product
our whole system of quantitative analysis is based. It is of
importance therefore to mention the fact that, although experi-
mental test has shown that both equation (1) and eqvaiion (2),
by means of which the law was derived, are grossly inaccuraie for
all strong electrolytes (neither Ki nor K2 remaining by any means
constant as conditions are varied), yet eqvxition (3), the law itself,
always approximates very closely to the truth.
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 457
niMistration of the Principle of Solubility'PriHiuct Cort'
stancy. — When, to a saturated solution of one of the less soluble
salts, a concentrated solution of a very soluble salt having one
ion in common with the first salt is added, partial precipitation
of the first salt is found to take place. This happens, for example,
with a saturated solution of potassium chlorate, which is not very
soluble (molar solubiUty 0.62, see Table). The concentrations
[K"*"] and [ClOa"] being small, one may easily increase the value
for one of the ions, say [ClOa*], fivefold, by adding a chlorate
which is sufficiently soluble. To preserve the value of the product
[K+] X [ClOs"], the value of [K+] will then have to be diminished
at once to one-fifth of its former value. This can occur only
by union of the ionic material it represents with an equivalent
amount of that for which [ClOa"] stands. The molecular material
so produced will thus tend at first to swell the value of [KClOa].
But the value of [KCIO3] cannot be increased, for the solution
18 already saturated with molecules j so that the new supply of mole-
culeSy or others in equal ninnbers, will be precipitated. Hence
the ionic part of the dissolved substance may be diminished, the
equilibria (p. 455) may be partially reversed, and we may actually
precipitate a part of the dissolved material without introducing
any substance, which, in the ordinary sense, can mteract with it.
In point of fact, when, to a saturated solution of potassiiun
chlorate there is added a saturated solution of potassium chloride
KCl (molar solubiUty, 3.9) or of sodium chlorate NaClOa (molar
solubility, 6.4), a precipitate of potassium chlorate is thrown down.
The precipitation of sodium chloride from a saturated solution,
by the introduction of gaseous hydrogen chloride (p. 127), is to be
explained in the same manner. The equilibria:
NaQ (soUd) <± NaCl (dslvd ) <=> Na+ + Cl-
are reversed by the introduction of additional CI" from the very
soluble, and highly ionized HCl.
458 smith's intermediate chemistry
Precipitation hy Double Decomposition, — The mechanism
of this type of reaction has abeady been discussed in some detail
in an earUer chapter (p. 185), to which reference should here be
made by the reader. The solubility-product law, however, throws
much additional light on the subject, and enables us to forecast
the completeness of any given precipitation under given conditions.
The first thing to be remembered is that the precipitate which
we observe, however insoluble its material may be, does not
include all of the substance, but only the excess beyond what is
required to saturate the water. The liquid surrounding the pre-
cipitate is always a saturated solution of the substance precipitated.
If it were not so, some of the precipitate would dissolve until the
liquid became saturated. Thus, for example, when we add ammo-
nium sulphide solution to zinc chloride solution (p. 451) :
(NH4)2S T± 2NH4+ + S=
ZnCl2 <=^ 2C1- + Zn++
2NH4CI ZnS T=± ZnS
(dissolved) (solid)
the liquid is a saturated solution of zinc sulphide, with the excess
of this salt suspended in it as a precipitate.
Looking at the matter from this viewpoint, we perceive the
application of the rule of solubility-product constancy. In this
saturated solution the product of the ion-concentrations, [Zn++] X
[S=], is constant. If the original solutions had been so very
dilute that, when they were mixed, the product of the concentra-
tions of these two ions had not reached the value of this constant,
no precipitation would have occurred. As a matter of fact, the
original salts are so extensively ionized in solution, and the solu-
bility-product of zinc sulphide is so small, that in all ordinary
mixtures the product [Zn"*"^] X [S"] considerably exceeds the
requisite value, and hence the salt is thrown down until the bal-
ance remaining gives the value in question. The rule for pre-
cipitation, then, is as follows: Whenever the product of the con-
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 459
centrations of any two ions in a mixture exceeds the value of the
ion-product in a saturated solution of the compound formed by
their union, this compound will be precipitated. Naturally the
substances with small solubilities, and therefore small solubility-
product constants, are the ones most frequently formed as pre-
cipitates.
Completeness of Precipitation. — In the above case, we
precipitate zinc sulphide practically completely from solution
by adding excess of ammoniiun sulphide. This substance, like
all salts, is highly ionized in solution, and therefore a solution
which contains it in excess contains a high concentration of sul-
phide ion S". The amount of zinc ion Zn++ that can exist simul-
taneously in such a solution is negUgible, since otherwise the solu-
biUty-product of zinc sulphide would be exceeded. Zinc is there-
fore precipitated quantitatively as zinc sulphide by addition of excess
of ammonium sulphide to a solution of any zinc salt. This is a
fact which is turned to practical use in quantitative analysis.
If we attempt to carry>ut the precipitation with hydrogen sul-
phide instead of anmioniiun sulphide, however, we find that only
partial precipitation of the zinc as sulphide occurs (p. 451). How-
ever great an excess of H2S we add, precipitation is incomplete.
The reason for this is immediately apparent from a study of the
full ionic equations:
H2S 4^2H+ + S-
ZnCl2?^2Cl- + Zn++
tl tl
2HC1 ZnS ?=fc ZnS
(dissolved) (solid)
Hydrogen sulphide is so slightly ionized in solution that the
concentration of sulphide ion S" that it supplies is extremely
small. A large amount of zinc ion Zn++ can therefore remain in
solution without the solubility-product of zinc sulphide being
exceeded.
460 sbhth's intermediate chemistry
Rule for Solution of Substances. — The rule for solution of
any electrolyte follows at once from the foregomg considerations,
and may be formulated by changing a few of the words in the rule
just given: Whenever the product of the concentrations of any
two ions in a mixture is less than the value of the ion-product
in a saturated solution of the compound formed by their union,
this compound, if present in the solid form, will be dissolved.
When applied to the simplest case, this rule means that a sub-
stance will dissolve in a Uquid not yet saturated with it, but
will not dissolve in a Uquid already saturated with the same
material. The main value of the rule lies, however, in its appU-
cation to the less simple, but equally conmion cases, such as
when an insoluble body is dissolved by interaction with another
electrolyte.
Applications of the Rule for Solution to the Solution of
Insoluble Substances. — So long as a substance remains in pure
water its solubility is fixed. Thus, with calciimi hydroxide, the
system comes to equiUbrium at 18° when 0.17 g. per 100 c.c. of
water (0.02 moles per liter) have gone into solution:
Ca(0H)2 (soUd) ^ Ca(0H)2 (dslvd.) <=^ Ca++ + 20H-.
But if an additional reagent which can combine with either one of
the ions is added, the concentration of this ion at once becomes less,
the actual numerical value of the ion-product therefore begins to
diminish, and further solution must take place to restore its
value. Thus, if a little of an acid (giving H"*") be added to the
solution of calcium hydroxide, the union of 0H~ and H"*" to form
water removes almost all the 0H~ (see p. 368), and solution of the
hydroxide proceeds until the acid is used up. There are now
more Ca++ than 0H~ ions present, but the iorir^rodiLCt reaches
the same value as before, and then the change ceases. If a further
supply of acid is added, the removal of 0H~ to form H2O begins
again. Hence, with excess of acid, the calcium hydroxide finally
all dissolves.
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 461
This particular action is a neutralization of an insoluble base.
But the other kinds of actions by which insoluble electrolytes pass
mto solution all resemble it closely, and differ only m details. The
general outlines of the explanation are the same in every case.
We proceed now to apply it to the common phenomenon of the
solution of an insoluble salt by an acid.
Interaction of Insoluble Salts with Acids, Resulting in
Solution of the Salt. — Zinc sulphide passes into solution
when in contact with acids, especially active acids. Thus, with
hydrochloric acid, it gives zinc chloride and hydrosulphuric acid,
both of which are soluble:
ZnS T + 2HC1 1^ ZnCl2 + H2S. (1)
The action of acids upon insoluble salts is so frequently mentioned
m chemistry and is so important a factor in analytical operations
that it demands separate discussion. This example is a typical
one and may be used as an illustration.
According to the rules already explained (p. 457), zinc sulphide
(or any other salt) is precipitated when the nimierical value
of the product of the concentrations of the two requisite ions
[Zn++] X [S"] exceeds the value of the ion-product for a saturated
solution of zinc sulphide in pure water. When, on the contrary,
the product of the concentrations of the two ions falls below the
limiting value, a condition which may arise from the removal in
some way either of the Zn"*"^ or of the S" ions, the solid will
dissolve to replace them imtil the ionic concentrations necessary
for equilibrium with molecules have been restored or until the
whole of the solid present is consumed. Here the sulphide-ion
from the zinc sulphide combines with the hydrogen-ion of the
acid (usually an active one) which has been added, and forms
molecular H2S:
S- + 2H+ 1^ H2S. (2)
462 smith's intermediate chemistry
It will be seen that the removal of the ions in this fashion can
result in considerable solution of the salt only when the acid pro-
duced is a feebly ionized one. Here, to be specific, the concentra-
tion of the S" in the hydrosulphuric acid equilibrium (2) above
must be less than that of the same ion in a saturated zinc sulphide
solution. Now hydrosulphuric acid belongs to the least active
class of acids, and its pure solution contains only a very small con-
centration of S= (p. 189). There is, however, another factor in the
situation which we have not yet taken into account. The hydro-
chloric acid which we used for dissolving the precipitate is a very
highly ionized acid and gives an enormously greater concentration
of hydrogen-ion than does hydrosulphuric acid. Hence the
hydrogen-ion is in large excess in equation (2), and the condition
for equilibrium, - — rrToT — " = K, will be satisfied by a corre-
1-H.2toj
spondingly much smaller concentration of S"". In this particular
case, therefore., the [S""] of the hydrosulphuric acid is far less than
that given by the zinc sulphide. The whole change, therefore,
depends for its accomplishment, not only on the mere presence
of hydrogen-ion, but on the repression of the ionization of the hydro-
sulphuric add by the great excess of hydrogen-ion furnished by
the active add that has been used. The whole scheme of the equi-
libria is as follows:
ZnS (solid) ^ ZnS (dissolved) t^' Zn++ + S-
2HC1 t=^2Cl- + 2H+
ti u
ZnCla H2S ?=fc H2S
(dissolved) (gas)
A generalization may be based on these considerations: an
insoluble salt of a given acid will in general interact and dissolve
whexi treated with a solution containing another add, provided
that the latter acid is a much more highly ionized (more active)
one than the former.
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 463
Precipitates Insoluble in Acids. — But even active acids
frequently fail to bring salts of weak acids into solution. Here
the cause lies in the fact that such salts are even less soluble
than those of the zinc sulphide type. Thus, even hydrochloric
acid (normal) will not appreciably dissolve cupric sulphide. The
solubility product [Cu++] X [S=] for this salt is so small that, after
an infinitesimal amount has gone into solution, the sulphide-ion
concentration is sufficient, in spite of the repressive action of the
large hydrogen-ion concentration furnished by the HCl, to bring
the product [Cu"*^] X [S"] up to its maximum possible value.
In this case the first link in the chain of equilibria:
CuS (soUd) 1=5 CuS (dissolved) t=^ Cu++ + S"
2HC1 t^2Cl- +2H+
tl . u
CuCk H2S T± H2S
(dissolved) (gas)
tends so decidedly backward that only the use of very concen-
trated acid will increase the concentration of the H"*" to an extent
sufficient to secure even a slight advance of the whole action.
We must add, therefore, to the above rule: provided also that the
salt is not one of extreme insolubility.
Illustrations of the application of these generalizations are
countless. Carbonic acid is made from marble (p. 333), hydrogen
sulphide from ferrous sulphide (p. 253), hydrogen peroxide from
sodium peroxide (p. 221), and phosphoric acid from calcium phos-
phate (p. 412). In each case the acid employed to decompose the
salt is more active than the acid to be liberated. On the other
hand, calcium phosphate is insoluble in acetic acid because this
acid is weaker than is phosphoric acid. We have thus only to
examine the list of acids showing their degrees of ionization (p.
189) in order to be able to tell which salts, if insoluble in water,
will be dissolved by acids and, in general, what acids will be
sufficiently active in each case for the purpose. In chemical
464 smith's intermediate chemistry
analysis we discriminate between salts soluble in water, those
soluble in acetic acid (the insoluble carbonates and some sulphides,
FeS and MnS, for example), those requiring active mineral acids
for their solution (calciiun phosphate, zinc sulphide and the more
insoluble sulphides, for example), and those insoluble in all acids
(bariiun sulphate and other insoluble salts of active acids).
Precipitation of Insoluble Salts in Presence cf Adds. —
The converse of solutioni namely, precipitationi depends upon the
same conditions: an insoluble salt which is dissolved by a given
acid cannot be formed by precipitation in the presence of this
acid. Thus zinc sulphide can be precipitated in presence of acetic
acid, but not in presence of active mineral acids in ordinary con-
centrations. Cupric sulphide or barium sulphate can be precipi-
tated in presence of any acid, but ferrous sulphide and calcium
carbonate only in the absence of acids.
From this it does not follow that zinc sulphide, for example,
cannot be precipitated if once an active acid has been added to the
mkture. To secure precipitation, aU that is necessary is to
remove the excess of hydrogen-ion which is repressing the ioniza-
tion of the hydrosulphiuic acid. This can be done by adding a
base, or by addmg ammonium sulphide.
Exercises. — 1. (a) How do the electrolytic methods of mak-
ing calcium and magnesium differ? (b) Why not electrolyze an
aqueous solution of magnesiiun chloride in making magnesium?
(c) Why use both potassium and sodium chlorides in making mag-
nesiiun? (d) Why is magnesiiun, but not potassium or sodium,
liberated?
2. Why are magnesium and zinc not found free in nature?
3. Why does magnesium rust completely (in time), while zinc
does not?
4. Make equations for: (a) the action of magnesium on hydro-
chloric acid; (b) the burning of magnesium in air; (c) the heating
MAGNESIUM AND ZINC. IONIC EQUILIBRIA 465
of magnesium carbonate; (d) the precipitation of magnesium
hydroxide from the sulphate; (e) the hydrolysis of magnesium
chloride.
5. Why is salt containing magnesium chloride, after mixing
with sodium bicarbonate, no longer deliquescent?
6. What is the density of zinc vapor (air = 1)?
7. Make equations for: (a) the action of hydrogen sulphide on
zinc oxide; (b) the precipitation of zinc hydroxide; (c) the action
of sodium hydroxide on zinc hydroxide.
8. What will be the effect of adding a concentrated solution
of silver nitrate to a saturated solution of silver sulphate (see
table of solubilities)?
9. Write full ionic equations for the reactions mentioned on p.
468.
CHAPTER XXXIX
ALnHinnrH
The family of trivalent elements to which the metal aluminium
belongs mcludes the non-metal boron (p. 363), and several rare
metals.
Occurrence. — Aluminiiun, although it does not occur free, is
the third element in orderof quantity (seep. 19). Its sUicates, such
as clay (kaolin) HAISiO,, mica K\lSi04, and felspar KAISisOg, are
amongst the most plentiful minerals. The oxide AljOj occurs as
corundum, sapphire, ruby, and emery (impure form). Bauxite is
a valuable hydrated oxide. Garnets Ca4Al2(Si04)3 are mined and
pulverized to make " sand " paper. Cryolite 3NaF,AlFs (Greek,
ice-stone) is imported from Greenland.
Manufacture. — The making of aluminium on a large scale
originated in C. M. Hall's discovery (1886) that the oxide could be
± „ electrolyzed in solution in molten cryo-
lite. Iron boxes (Fig. 110), about 5
by 3 feet and 6 inches deep, are lined
n with a compressed mixture of coke and
tar which is afterwards baked. The
lining forms the cathode, while the
oxygen is liberated at the anodes —
a series of rods of carbon about 3
"'^' ' " inches in diameter which are attached
to copper rods. The cryolite is melted (1000°) by the arcs
struck by the carbon rods. The latter are then raised some-
what, the alimiinium oxide is added, and some coal (which floate)
ie thrown in to cover the surface and obscure the blinding glow.
466
ALUMINIUM 467
From time to time more of the oxide is added and the melted
alimiinimn (m. p. 650 to 700°) is tapped off. The oxide must be
made from carefully purified bauxite, as the metal itself can not
be purified conmiercially. In 1866 it cost $250 to $750 per kilo-
gram but now sells at about 50 cents.
Properties. — The metal has a lower specific gravity (2.6)
than any other metal that could be put to the same uses (sp. gr.
iron 7.8, copper 8.8). It has malleability and the foil is taking
the place of tinfoil to some extent for wrapping foods. It has
considerable tensile strength, and is a good conductor of elec-
tricity. When heated, the metal bums brilliantly.
In the air it acquires only a slight film of closely adhering oxide.
This film prevents it from actmg upon water (hot or cold). When
the surface is cleaned and amalgamated with mercury, by dipping
in mercuric chloride solution, however, this metal acts as a contact
agent, and hydrogen is rapidly displaced :
2A1 + 6H2O ^ 3H2 T + 2A1(0H)3 i .
Uses. — The largest quantity of alimiinium is consumed by
steel-makers. When added in small amount (less than 1 : 1000)
to molten steel, it combines with the gases, and gives sound ingots
free from blow holes. Next to this comes its use for long distance
transmission of electricity. A cable of the requisite capacity
is larger than one of copper for the same current, but is lighter
and puts less strain on the supports. Cooking vessels of alumin-
ium are not corroded and are largely used. Cameras and opera
glasses are made of it. On account of its lightness, it is used
extensively in the metal parts of dirigibles and aeroplanes. Pul-
verized aluminium, mixed with oil, gives a paint which protects
iron admirably.
Aluminium bronze (copper, with 5 to 12 per cent almninium)
has a brilliant golden yellow color and is stable in air and easily
worked. Magnalium (containing 1 to 2 per cent of magnesium)
468 smith's intermediate chemistry
can be filed, turned, or polished like a mirror, and is better for
many purposes than the pure metal.
On account of its great chemical activity (p. 64), aluminiimi
displaces many other metals {e.g. iron, manganese, and chromium)
from combination. Thus powdered alimiinium and oxide of
iron when mixed (thermite) in a crucible, and started by a burning
magnesium ribbon, interact with great violence:
2A1 + Fe208 -> AI2O8 + 2Fe.
A temperature of 3000 to 3500° is reached, the molten iron (m.-p.
1500°) collects at the bottom, and the molten alimiinium oxide
(m.-p. 2050°) floats to the top. Steel rails are welded together,
and large objects of steel like broken propeller shafts are mended,
by enclosing a mass of thermite round the joint and firing it.
Aluminium Hydroxide A1(0H)3. — The hydroxide is precipi-
tated when anunonium hydroxide, or other alkaline hydroxide,
is added to a solution of a salt of aluminium:
Al2(S04)8 + 6NH4OH -> 2A1(0H)3 i + 3(NH4)2S04 (dslvd.).
It tends to remain in colloidal suspension (p. 441), and forms a
white gelatinous precipitate. It is both weakly boMC and feebly
accidie in chemical properties. In acids it dissolves forming salts
of aluminium, such as the chloride AICI3 or sulphate Al2(S04)3.
Solutions of these salts in water give an acid reaction, owing to
hydrolysis (p. 369).
Aluminium hydroxide dissolves also in sodium hydroxide solu-
tion, to form sodium aluminate NaaAlOs :
3NaOH + HsAlOs -> NasAlOs + 3H2O.
The aluminates are hydrolyzed by water, and their solutions
have an alkaline reaction.
Aluminium hydroxide, precipitated from the sulphate, is used
in sizing paper (to fill the capillaries and pores), in purifying
water (see p. 470), and as a mordant (see p. 475) in dyeing. Deli-
ALUMINIUM 469
cate fabrics (cloth) are rendered waterproof by saturating them
with aluminium acetate solution and boiling to promote the
hydrolysis. The alimiinium hydroxide is precipitated in the
capillaries of the cotton or linen, rendering them non-absorbent.
Aluminium Oxide (Alumina) A1203. — Corundumi and the
impure variety emery, are next to the diamond in the scale of
hardness, and are used as abrasives. Ruby and sapphire are also
crystaUized aluminium oxide, containing traces of impurities
(iron and titanium in the one case and chromium in the other) to
which they owe their colors. By ingenious methods of fusing the
oxide, " synthetic " sapphires to the extent of six million carats
and rubies to the extent of ten million carats are now made annu-
ally. The artificial gems are chemically identical with the natural
ones, and can be distinguished only by the fact that they are free
from microscopic bubbles and other defects. Alundum, an arti-
ficial abrasive, and refractory material for crucibles and muffles,
is made by barely melting the oxide in the electric furnace.
Aluminium Sulphate AI2 (804)39 I8H2O. — The sulphate is
manufactured by the action of sulphuric acid on bauxite. It
crystallizes in leaflets, which usually have a faint yellow tinge due
to the presence of iron (Fe(0H)3) derived from the mineral. The
salt is used in fireproofing cloth, since, when heated, it melts in its
water of hydration. It is used as a source for precipitated alu-
minium hydroxide in paper-making, water purification, and
dyeing.
Alums. — When almninium sulphate and potassiiun sulphate
are dissolved together in molecular proportions, the solution
deposits transparent octahedral (Fig. 39, p. 94) crystals of
potash-alum K2S04,Al2(S04)8,24H20. This salt is more easily
freed from impurities (e,g, compounds of iron) by recrystaUization
than is aluminium sulphate, and is therefore used instead of the
470 smith's intermediate chemistry
latter in medicine, in dyeing (delicate shades), and to replace the
cream of tartar (p. 431) in making baking powder.
Sodium - alumi Na2S04,Al2(S04)3;24H20, ammonium - alum
(NH4)2S04, Al2(S04)3, 24H2O and chrome-alum K2SO4, ^2(804)3,
24H2O are made in the same way, and crystallize in the same form.
The first two are used as sources of aluminium hydroxide, and the
last in the " fixing bath '' to harden the gelatine on photographic
films and plates.
Water Purification — Coagulation Prtfcess. — The sus-
pended matter in water to be used for a domestic supply can be
coagulated into larger particles by introducing a small amount of
the gelatinous precipitate of aluminium hydroxide. These larger
particles, which adsorb also the greater part of the bacteria, settle
rapidly and the process therefore permits the use of relatively
small settling ponds. Aluminium sulphate, made from crude
bauxite, and lime are added to the water:
3Ca(OH)2 + Al2(S04)3-^3CaS04 + 2A1(0H)3.
K the water has much temporary hardness, lime is not required:
»
3Ca(HC08)2 + Al2(S04)8 -> BCaSO* + 2A1(HC03)8.
The carbonate of aluminium, being a salt of both a very weak
acid and a very weak base, if formed, would be instantly hydro-
lyzed:
A1(HC03)8 + 3H2O -^ A1(0H)8 i + 3H2CO3
so that aluminium hydroxide is precipitated.
The few remaining bacteria are destroyed by the addition of
bleaching powder (p. 224).
Crude ferrous sulphate FeS04 (copperas), being in many places
cheaper than aluminium sulphate, is often used instead of the
latter. The lime precipitates ferrous hydroxide Fe(0H)2. This
is quickly oxidized to the red ferric hydroxide Fe(0H)3, which
coagulates the suspended matter.
ALUMINIUM 471
Clay and Pottery. — Pure clay (kaolin) is white. It is hydro-
gen-aluminum siUcate HAlSi04, derived from the weathering
of felspar (p. 410). Common clay contains impurities such as
sand (silica), limestone, and compounds of iron. Both kinds are
plastic when wet and can be moulded. When heated strongly
the material shrinks (so that the products are porous) and becomes
hard. Bricks, and tiling for roofs and drains, are made of common
clay and, when red, owe their color to oxide of iron (Fe203).
The j&ring is done with fuel gas in ovens or kilns of brickwork.
To glaze drain pipes and some bricks, salt is thrown into the kiln.
The vapor of the salt produces a more fusible sodium-almninium
silicate, which fills the surface pores. Clay for fire-brick (infusible)
must contain free silica, but no lime.
China and porcelain are white, translucent and non-porous.
They are made of pure clay to which a little of the more fusible
felspar is added. After firing, the articles are dipped in water,
in which the materals for the " glaze,'' namely finely ground
felspar and silica, are suspended. Having thus acquired a thin
coating of these substances, they are fired again at a higher tem-
perature and for a longer time! Colored decoration is done with
materials which melt (third firing) to colored enamels (p. 362).
Ultramarine. — This material is made by heating together
kaolin, sodium carbonate, sulphur, and charcoal, pulverizing the
green mass, and heating it again with more sulphur. The product
is used as laundry blueing, and in making blue-tinted paper. It
is added also to correct the yellow shade of linen, starch, sugar
(p. 402), and paper stock.
Cement. — Portland cement is made by heating a pulverized
mixture of a material rich in lime, such as limestone CaCOs,
with one in which silica, iron oxide and alumina are the main
constituents, such as common clay. Some natural rocks contain
all of the necessary elements in suitable proportions. The finely
472 smith's INTERBfEDIATB CHEMISTRY
powdered material is first burnt or raised to a temperature of
1400-1600**, at which temperature it fuses partially and forms
lumps or dinkers. When these have cooled, they are mixed with
2-3 per cent of gypsum and pulverized again. The resulting
product is Portland cement, the manufacture of which in the
United States alone now exceeds 90,000,000 barrels (of 380 poimds)
yearly. •
Portland cement is essentially a mixture of calcium silicate
and calcium alimiinate, with excess lime. The calcium silicate is
simply a filler. The calciiun aluminate is hydrolyzed on addition
of water, according to the equation:
Ca3(A103)2+ 6H2O -^ 3Ca(OH)2 + 2H3AIO3.
The calcium hydroxide thus formed slowly crystallizes, connecting
the particles of the calcium silicate. The aluminium hydroxide
fills the interstices and renders the whole compact and impervious.
The small amoimt of gypsum added regulates the setting time
of the cement. The iron oxide is necessary to assist in the burning
and to lower the temperature at which the mixture begins to fuse.
K too small a quantity of excess lime is present, the cement will be
unsound and crack on drying. Too little excess lime gives a
cement which sets too quickly and is lacking in strength.
Concrete is a mixture of cement with sand and crushed stone or
gravel, all made into a paste with water. It sets to a solid mass,
suitable for walks, and for the foimdations, walls, and floors of
buildings. Since no carbon dioxide from the air is required in the
hardening process (contrast with p. 386), it sets equally well
under water (hence hydraulic cement), and is employed in con-
structing dams, levees, and the foundations of bridges. Rein-
forced concrete contains twisted rods of iron, embedded in the
mass, and is much used in building construction.
Blast-furnace slag, when pulverized and heated with limestone,
has been found to yield an excellent quality of cement, and a
valuable use has thus been found for what was formerly an annoy-
ing encumbrance.
ALUMINIUM 473
Exercises. — 1. Make equations for the following actions: (a)
aluminium on hydrochloric acid; (b) aluminium on mercuric
chloride HgCU; (c) displacement of manganese from manganese
dioxide by aluminium.
2. How must alimiinium hydroxide be ionized so as to behave
both as an acid and a base?
3. Make the equation for heating calcium carbonate with:
(a) aluminiiun hydroxide; (b) aluminium oxide.
4. Explain why the reaction of solutions of aluminates is alka-
line (p. 369) and that of alums is* acid.
6. Make the equation: (a) for the action of sulphuric acid on
bauxite, assuming the formula of the latter to be Al20(OH)4; (b)
for the formation of potash-alum.
6. Why is the tarnish on aluminium the oxide, and not the
carbonate (as on Zn and Mg)? What qualities in a tarnish enable
it to protect the metal from further oxidation (p. 467)?
7. Make equations for the action of bicarbonate of soda and
aluminium sulphate (alum baking powder) when heated.
Explain what raises the bread.
CHAPTER XL
SYNTHETIC ORGANIC PRODUCTS
Mention was made several times in the preceding chapter of
the use of salts of almninium in dyeing, A brief smnmary of the
chemistry of dyeing and dyestuffs is all that can be presented
here. A still more restricted r^sum^ of some other important
synthetic organic products is also given. The products included
are, on the constructive side, perfumes, drugs and plastics (includ-
ing rubber); on the destructive side, explosives and toxic gases.
Some dyes, perfiunes and drugs are still, of course, obtained
from natural sources, and synthetic rubber continues to show
but little promise of superseding plantation rubber in the near
future. Nevertheless, the interest of the chemist in the fields
here under discussion is almost entirely synthetic, seeking to
duplicate natural products by laboratory methods and to dis-
cover new products of even greater use to humanity.
Dyeing. — The problem of the dyer is to confer the desired
color upon a fabric made, usually, of cotton, linen, wool, or silk,
and to do this in such a way that the dye is fast to {i.e., is not
removed or destroyed by) rubbing and light, and often, also,
to washing with soap. To understand the means by which this
is achieved, it must be noted that cotton and linen consist of
smooth hollow fibers (Fig. 2, p. 2) of cellulose. Wool is made
of hollow fibers with a scaley surface, and silk of solid filaments,
but these are composed of proteins (p. 438). Now, the proteins
are much more active chemically than is cellulose, and also, as
colloidal materials, seem to have a much greater tendency to
adsorb other substances (see pp. 422, 443) than has cellulose.
Hence, accidental stains on wool or silk are much less often remov-
474
SYNTHETIC ORGANIC PRODUCTS 475
able than are those on cotton, and when samples of the three
materials are dipped in a solution of a dye, the first two are per-
manently dyed, while from the last most dyes can be completely
washed out with water.
Three modes of dyeing may be mentioned:
1. Insoluble Dyes. If the colored body can be produced by
precipitation, after the solution has filled the capillary and wall
of every fiber of the goods, then, if the dye is sufficiently insolu-
ble, it is mechanically imprisoned in every fiber and cannot be
washed out. This plan may be applied to any kind of goods.
For example, if cotton, silk, or wool is first boiled in a solution of
lead acetate, and is then soaked in a boiling solution of potassium
chromate K2Cr04, it is dyed a brilliant, permanent yellow. Lead
chromate is the colored body:
PbCCOaCHs) + K2Cr04±^2K(C02CH3) + PbCr04 i .
The part precipitated on the outside of the goods can be, and is,
at once washed off by rubbing in water, but the particles inside
the fibers can come out only by being dissolved, and they are
insoluble in water. Indigo C16H10N2O2, which is used in larger
amounts than any other dye, belongs to this class. The cloth
is saturated with an alkaline solution of indigo white C16H12N2O2,
a soluble, slightly acid substance, and the oxygen of the air subse-
quently oxidizes this and deposits the insoluble indigo blue within
the fibers:
2C16H12N2O2 + O2 -> 2C16H10N2O2 i + 2H2O.
2. Mordant or Adjective Dyes. Since cotton is inactive chem-
ically and has but a slight tendency to adsorb dyes, it is usually
necessary first to introduce into the fibers of cotton some colloidal
substance with greater adsorptive powers. Substances of this
kind are tannic add (p. 499) for basic dyes, and gelatinous col-
loidal hydroxideSy such as those of alimiinium, tin, iron and chro-
mium, for non-basic (including acid) dyes. They are called mor-
dants (Lat. mordere, to bite). Thus, if in three jars we place
476 smith's intermediate chemistry
very dilute solutions of aluminium sulphate, ferric chloride FeCl3
and chromous acetate Cr(C02CH8)2, then add a few drops of a
solution of a dye to each, and finally introduce a little of a base
(like sodium hydroxide) to precipitate the hydroxide of the metal,
this hydroxide will adsorb the dye and carry it into the precipitate.
Such a precipitate of mordant and dye is called a lake. With
the same dye, the three lakes have different colors. Thus, in
the above-mentioned experiment, if alizarin (madder) is used as
the dye, the colors are red (Turkey red), violet, and maroon,
respectivdy. This is probably due to the different degrees of
dispersion in the three colloidal materials. If aluminium hydrox-
ide is to be used, by first saturating the cloth with hot aluminium
acetate solution (p. 468), or by using first aluminium sulphate
and then ammonium hydroxide, the aluminium hydroxide is
precipitated within the fibers of the goods. When the material
is then dyed, the coloring matter is adsorbed by the mordant, with
which it forms an insoluble lake, within the fibers. Basic dyes,
like Malachite green and Methylene bltte, behave similarly with
tannic acid, or an insoluble salt of tannic acid, as mordant. It
will be seen that, so far as the fabric is concerned, this process,
like the first, is a mechanical one, and is independent of the chem-
ical nature of the goods.
3. Direct or Substantive Dyes. Most organic dyes are direct
dyes on silk or wool, and require no mordant with these materials.
The actions seem to be sometimes chemical, but more often cases
of adsorption by the silk or wool (both colloids) themselves. A
few dyes are also fast on cotton. Congo-red is fast both on cotton
and wool, but is no longer much used. Chrysophenin is now one
of the conunonest dyes of this class. These dyes, which are
sodium salts of complex organic acids, are colloids like soap (p.
442), and are salted out within the fibers of the goods by adding
sodium sulphate to coagulate them and assist the adsorption by
the cotton. Once adsorbed in this way, unlike soap, they cannot
be washed out.
SYNTHETIC ORGANIC PRODUCTS 477
Dye'Stmffs. — Natural dye-stufifs have now been almost
entirely superseded by manufactured products, which can be
prepared more cheaply and are of superior quality. Logwood,
still used as a black mordant dye for silk, is the only important
exception. The total value of natural dyes imported into the
United States in 1919 was only $1,250,000, while the artificial
dyes made in the country in the same year were worth over $70,-
000,000.
The vast bulk of sjmthetic dyes are built up from ring hydro-
carbons (p. 352) and their derivatives, extracted from coal tar
(p. 424). By substituting suitable groups into the molecule
of the simpler colorless products, more complex derivatives with
great brilliancy of color are obtained. Thus indigOi formerly
the most extensively used of all natural dyes, is now manufactured
most conveniently with aniline (p. 353) as a starting-point. Ali-
zarin (Tm-key red), once extracted from madder root, is obtained
from anthracene.
By varying the position of the substituted groups in the mole-
cule, the most delicate variations in color can be effected. The
most precious of all dyes in ancient times was Tyrian purple,
obtained from certain species of sea snails (Murex). The secret
of preparing this substance was lost for centuries, but in 1909
Friedlaender gathered 12,000 of these mollusks and succeeded in
isolating 1.5 grammes of the coloring material for analysis. He
showed it to be a derivative of indigo, containing two bromine
atoms in place of two of the hydrogens. This identical substance
had been prepared synthetically five years earlier, but found to
be inferior to another dye containing the bromine atoms in diflfer-
ent positions in the molecule!
Preparation of Typical Dyes, — The student is recommended
to attempt the following preparations, which are at the same
time easy and instructive.
(a) Phenolphihalein. Take 0.1 gram of phthaUc anhydride (the
478 smith's intermediate chemistry
anhydride of phthalic acid, C6H4.(COOH)2) and 0.1 gram of phenol
C6H5.OH in a test-tube, add 2 drops of concentrated sulphuric
acid, and heat carefully for a minute over a small flame. The
mixture, which should be well shaken during the heating, will
turn dark-red in color. Allow to cool, add a few c.c. of water,
and then add drop by drop a dilute solution of sodium hydroxide
until a pink color persists on shaking.
Take a portion of this solution, and test the action of phenol-
phthalein as an indicator by adding first dilute hydrochloric acid,
then dilute sodium hydroxide (see p. 168).
(6) Fliurrescein, To 0.1 gram of phthalic anhydride and 0.1
gram of resorcinol C6H4(OH)2 add 3 drops of concentrated sul-
phuric acid, and heat carefully for a minute. Allow to cool, add
a few c.c. of water, then add sodium hydroxide until alkaline
(use litmus paper as a test).
Shake up a few drops of this solution with a test-tube of water.
The dye imparts to the solution a brilliant green fluorescence,
hence its name.
Perfumes. — Many natural perfumes and fruit flavors can
also be produced synthetically. The basis of most of these con-
sists of esters (p. 349). The fragrance of ripe apples is due to
minute amounts of the alnyl esters of formic, acetic and caproic
acids. In bananas the characteristic ester is amyl acetate; in
grapes it is methyl anthranilate. Almond flavor is due to benz-
aldehyde, CeHs.CHO; the smell of geraniums to diphenyl ether,
(C6H5)20. The chief ingredient of otto of roses is geraniol; of
the vanilla bean, vanillin; of the perfume sold as " new mx/wn
hay" coumarin; all complex hydrocarbon derivatives which
are now commercial products. Some of the synthetic flavors
and perfumes on the market are identical with the substances
that give odor to fruits and flowers; others are merely more or
less adequate imitations. Thus, natural oil of wintergreen is
essentially the same as synthetic methyl salicylate. On the other
SYNTHETIC OBGANIC PRODUCTS 479
hand amyl valerate, which is sold as apple essence, is not contained
in the fruit, though it smells like it. Most perfumery nowadays,
including the most expensive, consists of mixtures of natural and
synthetic products.
Preparation of Typical Perfumes. — Take 0.1 gram of
beta-naphthol C10H7.OH and 10 drops of methyl alcohol in a
test-tube, add 2 drops of concentrated sulphuric acid, and warm
gently for a few minutes. The methyl ether of beta-naphthol,
which is formed, has a most powerful odor, reminiscent of acacia
blossoms.
Repeat the experiment with ethyl alcohol instead of methyl
alcohol. The ethyl ether of beta-naphthol is produced; its odor
recalls the perfiune of orange flowers.
Drugs. — In the same way, many drugs formerly obtained from
natural sources are now built up in the laboratory, and many new
compounds have been made which possess as valuable medicinal
properties as any found in nature. The first stage in this work
consisted in determining the constitution of the active ingredients
of plant products. Quinine was isolated from cinchona bark,
morphine from the seed capsules of the opium poppy, strychnine
from the seeds of nicx vomica. These and other substances of
similar character are now classed together as aZfcaZoicfe,^ complex
nitrogeneous substances possessing basic properties. The struc-
ture of many of the alkaloids is now completely worked out.
It is not necessary, however, to construct the entire complicated
molecule if the same results can be secured'with simpler substances,
and this has been demonstrated already in many instances.
Aspirin is an ester of salicylic acid C6H4.OH.COOH; acetavr
Hide is a derivative of aniline C6H5.NH2. Acetophenone
(hypnone), CeHs.CO.CHa, is* used as a hypnotic. Novocains, pre-
pared from diethylaniline, has practically replaced the natural
alkaloid cocaine in dentistry. Saccharin is obtained from toluene;
480 smith's interbiediate chemistry
its solution in water has an intensely sweet taste^ hence it is used
as a substitute for sugar in war times and in cases of diabetes.
It has no food value, however, and hence cannot replace sugar
m nutrition.
Plastics. — These are substances, like celluloid, which can
be moulded or shaped into any desired form. Among natural
plastics may be mentioned resins, gums, and rubber. Synthetic
rubber equal to the natural product in durability and cheapness
has still to be prepared (see p. 165), but many other plastics are
now of considerable industrial importance.
Certain cellulose plastics have already been described (p. 399).
Cellulose behaves chemically like an alcoholy and as such forms
esters with acids (see p. 349). When cotton is treated with acetic
acid (in the form of acetic anhydride (CH8.CO)20, see p. 315) it
gives cellulose acetate. The viscous liquid dries to a tenacious film.
On account of its waterproof character, non-inflammability and
non-conductance of electricity, it is now used for coating the wings
of aeroplanes, for making moving picture films, and for insulating
electric wires. Artificial horsehair {e.g. for making women's hats)
and bristles for hair brushes are made of it.
By the action of nitric acid upon cellulose, various cellidose
nitrcUes may be formed, according to the niunber of OH groups
replaced by NOs (see gimcotton, below). An mcompletely
nitrated'ester, when worked between rollers with camphor* and a
little alcohol, forms a viscous solution. When the alcohol evap-
orates, a transparent colorless solid, celluloid, remains. Photo-
graphic films are made by rolling the dough into sheets. Fillers
and dyes can be added to the dough and the latter can be moulded
to any form. In this way ivory-like or black combs and brush
handles, opaque white knife handles, articles of " artificial amber "
and so forth can be made.
* A white solid with the formula CioHwO, obtained commercially by distill-
ing with steam the wood of the camphor tree, but recently also prepared syn-
thetically.
SYNTHETIC ORGANIC PRODUCTS 481
The same sort of guncotton dissolves in a mixture of alcohol and
ether, giving a solution called collodion, used in photography and
in medicine.
When collodion is forced under great pressure through minute
holes in a steel die, the threads dry as they issue from the open-
ings and can be wound on spools. The product is treated with
an alkah, which decomposes the ester, leaving a material of the
composition of the original cotton. The product is another
form of artificial silk (p. 399).
Another plastic, not chemically related to the preceding, is
bakelite, prepared from formaldehyde CH2O and phenol CoHbOH
(carbolic acid). Under suitable heat treatment the mixture grad-
ually sets to a solid, hard, infusible, resinous mass, which is insol-
uble in all known solvents. Before it sets, it can be dyed or
" filled," and it can be applied as lacquer, or moulded to any form.
Switchboards, dolls, ornamental buttons, artificial jewels, phono-
graph records, billiard balls, and stereotyping matrices are
amongst the objects into which it is now fashioned.
Explosives. Nitroglycerine. — Some inorganic explosives,
gunpowder (p. 372) and ammoniimi nitrate (p. 302), have been
discussed in earlier chapters. The main organic explosives are
also compounds of nitrogen.
As already mentioned, the alcohols interact with inorganic
acids, as well as with organic ones, to produce esters. A familiar
illustration is met with in the manufacture of nitroglycerine
(glyceryl trinitrate) by the action of glycerine and nitric acid:
C8H6(OH)8 + 3HN08 -> C8H6(N08)8 + 3H2O.
To assist in the Uberation of the water, the nitric acid is mixed
with a dehydrating agent. The glycerine then is added slowly
to the cooled reagents. The nitroglycerine is an almost color-
less oil which fioats to the surface of the acid mixture. It is
shaken repeatedly with water, in which it is insoluble, and then
482 smith's intermediate chemistry
with sodium carbonate solution, in order to free it from all traces
of the acids.
Nitroglycerine explodes violently, often from the slightest shock.
It owes this power to the fact that its carbon and hydrogen can
combine with the oxygen it contains to form carbon dioxide and
water:
2C8H5(N08)8 -> 6CO2 + 5H2O + 3N2(+0).
The latter are very stable substances and much heat is liberated
in forming them. They are both produced as gases and, at the
high temperature of the action, they and the nitrogen tend to
occupy a great volume — or to exert an enormous pressure in the
effort to do so.
The explosion is also so siuiden, compared with that of gun-
powder, that nitroglycerine would shatter the breech of a rifle
before the bullet had time to move. It also pulverizes rock,
instead of breaking it into fragments of usable size. For these
reasons, as well as on account of the danger in handling, and
impossibility of safely transporting the substance, it is made into
blasting gelatine (see below). The old form of dynamite was
made by soaking a porous earth (infusorial earth, kieselguhr)
with nitroglycerine.
Guncotton. — When cotton is steeped for half an hoiu* in
a cooled mixture of nitric and sulphuric acids, it is converted into
cellulose trinitrate or guncotton:
C6H702(OH)8 + 3HNO3 -> C6H702(N03)8 + 3H2O.
The equation, as written above, shows that three hydroxyl groups
OH in the empirical cellulose formula CeHioOs are replaced by
three nitrate groups NO3, with simultaneous formation of three
molecules of water. The sulphuric acid hastens the reaction and
carries it to completion by acting as a dehydrating agent and re-
moving this water (see p. 270), The fibers have the same appear-
SYNTHETIC ORGANIC PRODUCTS 483
ance as before, but are crisper to the touch. The guncotton is
washed thoroughly with water to remove the acids, which would
cause slow decomposition and perhaps accidental explosion.
Dried guncotton bums briskly (deflagrates) when set on fire.
While wet, it can be moulded and cut without danger. It
explodes only when " set off " by a small amount of another
explosive. Fulminate of mercury Hg(0NC)2, used in percussion
caps, is commonly employed. By such means the explosion is
brought about in wet guncotton as easily as in dry.
In pure form guncotton is used only in torpedoes and submarine
mines. It explodes too rapidly to be used in fire-arms or for
blasting.
Smokeless Powfier and Dynamite. — The violence of gun-
cotton is reduced by compressing it, and still more by dissolving
it and allowing the solvent to evaporate. Thus, cordite is made
by dissolving guncotton (65 parts), nitroglycerine (30 parts) and
vaseline (5 parts) in acetone. The resulting paste is rolled and
cut into pieces of different dimensions, according to the rate of
explosion desired. When the acetone evaporates, the homy
cordite remains. These explosives are smokeless because they
differ from gunpowder (p. 372) in yielding no solids when they
decompose.
Blasting gelatine, giant powder, and other forms of djoiamite
are made by dissolving guncotton in nitroglycerine. Substances
like nitrate of sodium or of ammonium and sawdust or flour are
added to adjust the rate of explosion so that, for example, coal
may be split up, but not shattered.
High Explosives. — These are substances which develop their
explosive affect at an extremely rapid rate, and are used therefore
when a shattering effect of great violence is required, for example
in bursting shells or in anti-svbmarine mines, Trinitro-toluene
(TNT) is made by nitrating toluene (p. 353), picric acid (trinitro-
484 smith's intebmediate chemistry
phenol) by nitrating phenol. Picric acid* is a strong acid and forms
salts which, with the exception of ammonium picrate, are more
sensitive to shock and friction than the acid itself. The action
of this explosive on the metals with which it comes in contact
must therefore be guarded against. TNT, however, is inert
towards metals, and is insensitive towards ordinary shocks. It
melts at a low temperatiu-e (80°), and so can be readily lique-
fied and poured into shells. Mixed with ammonium nitrate, it
gives amatol (p. 302).
Toxic Gases. — The poisonous substances employed in the
Great War were mainly sjmthetic organic products. Most of
these, it is true, were not gases, but Uquids or solids of low volatil-
ity. Their vaporization was, however, favored by the explosion
of the shell in which they were contained, part of the material
being converted to gas by the heat of the explosion and the
remainder being scattered aroimd in a finely-divided condition.
Some of the substances used were so powerful in their action upon
the human system that a concentration of 1 part in 10,000,000
in the air was sufficient to incapacitate anyone unprotected with a
gas-mask.
The first method employed in gas warfare was to release a
highly-volatile poisonous substance {chlorine^ or a mixture of
chlorine and phosgene COCI2) from cylinders under pressure. A
cloud of gas was thus evolved, which under favorable wind condi-
tions was carried over the enemy's lines. This cloiid method was
soon abandoned in favor of the sheU method described above.
Among the substances thus employed may be mentioned ddor-
pterin CCI3.NO2 (made by the action of bleaching powder and
lime upon picric acid) and mustard gas (CH2C1.CH2)2S. The
* Picric acid is an interesting example of the close interrelation of synthetic
organic compounds. It is used in dyeing to give a yellow color, in warfare
as an explosive and in medicine as an antiseptic. It may happen that picric
acid is used in a base hospital to cure the woimd that picric acid caused.
SYNTHETIC ORGANIC PRODUCTS 485
latter substance, the chemical name of which is dichlordiethyl
sulphide, may be regarded as diethylether (p. 348) in which oxy-
gen is replaced by sulphiu- and two hydrogen atoms are replaced
by chlorine. It was produced by the action of ethylene C2H4
upon sulphur monochloride S2CI2.
Both of the above substances are actually toxic, that is, a
sufficient concentration will induce death. Many other sub-
stances were used, however, which merely put the victim tempo-
rarily out of action. Among these were lachrymatories (tear-
producing substances) such as benzyl bromide C6H6.CH2Br, and
8temiUatories (sneeze-producing substances) such as diphenyl-
chlorarsine (C6H6)2AsCl. The latter substance is a derivative
of arsine AsHs, the three hydrogen atoms being replaced by two
phenyl groups CeHs and one chlorine. Mustard gas in very low
concentration also acted as a skin irritant.
The methods employed for obtaining protection against toxic
gases have already been discussed (p. 421).
Exercises. — 1. Write graphic formulae (see Chapter XXIX)
for the following synthetic essences: ethyl formate, ethyl
butyrate, benzaldehyde, diphenyl ether.
2. What chemical change would occur after mixing nitro-
glycerine with sodium hydroxide solution? Name the kind of
reaction and give the equation.
3. When nitroglycerine explodes, in what relative volumes are
steam, carbon djoxide, and nitrogen produced? What principle
is used in answering this question?
4. Make an equation for the decomposition of guncotton,
similar to that given for nitroglycerine (p. 482).
5. Make an equation for the denitration of guncotton by an
alkali.
6. Write the graphic formulae for the following substances:
acetic anhydride, trinitrotoluene, picric acid, phosgene, chlor-
picrin, mustard gas, benzyl bromide, diphenylchlorarsine.
CHAPTER XLI
mON, NICKEL, COBALT
Iron Fe
Occurrence. — Iron as a free metal is rarer than gold (see p. 54).
Masses of iron are found in meteorites, lately arrived on earth
from airless space. Minute particles of the metal may be detected
in igneous rocks, freshly broken open. Compounds of iron are
very widely distributed. Pyrite FeS2 (fools' gold) is used mainly
as a source of sulphur for sulphuric acid. The ores which yield
iron itself are:
Fe203 (ferric oxide), red haematite. Red when pulverized.
2Fe203,3H20 (hydrated ferric oxide), brown haematite.
Fe304 (magnetic oxide of iron), magnetite. Black when pul-
verized.
FeCOs (ferrous carbonate), spathic iron ore.
The carbonate, mixed with clay (clay iron-stone), furnishes most
of the iron in Great Britain, but less than one per cent of it in the
United States. The ore is first calcined to produce the oxide.
The Blast Furnace. — Coke is used to reduce the oxides and
as fuel. The carbon monoxide CO, produced by the burning of
the coke and air, is the actual reducing agent:
Fe804 + CO ?=^ 3FeO + CO2.
FeO +CO?=^Fe + CO2.
Since the ores contain rocky material (gangue), such as silica
Si02 and silicates of aluminium, limestone is added in the
proportion required to give a fusible slag.
486
IRON, NICKEL, COBALT
487
The blast furnace (Fig. Ill) is an iron structure 40 to 100 feet
high, lined with fire-brick. A circular pipe delivers a blast of
pre-heated compressed air to several nozzles
(tuyeres) near the foot. The ore, coke and lime-
stone are admitted at the top. Reduction of the
ore occurs continuously as the solid materials, q
passing downwards, become hotter and hotter
through combustion of the coke. The melted
iron and slag (immiscible) finally collect in two
layers in the hearth or crucible at the bottom.
From time to time the slag is allowed to flow
from an opening near the top of the crucible,
and the iron from a similar opening at the bottom.
Plugs of wet clay close the openings and are in-
stantly baked hard. The iron is taken in ladles
to other parts of the plant* or is cast into " pigs "
in steel moulds and chilled in water.
Fig. Ill
Reactions in the Blast Furnace. — The actions (see equa-
tions, above) are both reversibk and the carbon dioxide formed
tends to react to reproduce the original materials. At any partic-
ular temperature it is necessary therefore to keep the proportion
of CO2 to CO in the furnace gases below a certain value in order
to prevent the reversal of the reactions. This proportion of CO2
to CO is regulated, however, by a third reversible reaction:
C + CO2 ^ 2C0.
At high temperatures (above 1000*^) this reaction is almost com-
plete in the forward direction. As the temperature falls, the pro-
portion of CO2 in the equilibrium mixture increases rapidly.
If all three reactions had time to attain equilibriiun conditions
in the blast furnace, reduction of Fe304 to FeO, and of FeO to Fe,
would occur at about 650*^ and 690° respectively, and the residual
gases would contain very little CO. In practice, however, the
488 smith's interbiediate chemistry
slowness of the reactions necessitates the use of much higher
temperatures and a large excess of CO. The gases that escape
are therefore combustible, and are led ofif through openings near
the top of the furnace and used for pre-heating the air-blast.
In 1920, nearly 40,000,000 tons of pig iron were produced in
the United States. The production in Great Britain exceeded
8,000,000 tons.
CmsI Iron. — Pig iron contains 4 to 5 per cent of carbon and
varying amoimts of silicon (as silicide of iron), phosphorus (as
phosphide) and sulphur (as sulphide). These impurities lower
the melting-point from 1510° to about 1100°. The material is
hard and brittle. Most of it is made into wrought iron or steel,
but some is used in making objects of cast iron, such as ranges,
stoves, pipes, and radiators, which are not to be subjected to
shocks or strains. Cast iron expands on solidifying and forces
itself into the details of the mould.
By adding pyrolusite Mn02 in the blast furnace, cast iron con-
taining from 20 per cent of manganese (spiegeleisen) up to 80 per
cent (ferromanganese), and carbon up to 6 per cent, is made for
use in steel manufacture.
Wrought Iron. — Wrought iron is commercially pure iron.
The broken pigs are placed in a reverberatory furnace (Fig. 112),
the hearth of which is covered with a bed of hsematite ore Fe208
and silicates. The flames and heated gases, deflected by the low
roof, play upon the iron and melt it. The oxygen in the hsema-
tite combines with the carbon, phosphorus, sulphur, and silicon,
giving the oxides. The mass is worked vigorously with iron rods
upon the bed of haematite (puddled), carbon monoxide escapes,
and the iron becomes more viscous as its melting-point rises on
account of the removal of the impurities. Finally, it is collected
in balls (blooms) on the iron rods. The treatment occupies an
hour and a half. To press out the slag, the blooms are first passed
IRON, NICKEL, COBALT
489
through a squeezer and then put through the rolls. The result-
ing bars are repeatedly cut, " piled " in a bundle, reheated, and
rolled. These treatments, and the presence of a Uttle slag dis-
tributed through the mass, give wrought ironthe valuable prop-
erties which distinguish it from all other iron products, namely
Vtt
m
nnmninffl
iiiiiiiii
1 , 1
I . I
i.'i.'i.^
i!i!i!i!i!i'i!:!i!i!i
■:■:■:■:■:■
masK^^^^^^^^^^
'.I '.I I
Fig. 112
its fibrous structure and its extreme toughness. On account of
these properties it is used for anchors, chains, and bolts. It is
drawn into wire, and, when heated, can be cut, shaped, and
welded under the hammer. The impurities having been greatly
reduced (to 0.1 or 0.2 per cent), this iron is much less fusible than
cast iron, and is used for fire bars.
Crucible Steel. — Steel contains 0.75 to 1.5 per cent of carbon,
and is freed as far as possible from other impurities. Small lots,
for special purposes, are made in clay (or graphite and clay)
crucibles in melts of 60 to 100 pounds. The charge in Shefiield
consists of bUster steel, i.e., carburized Swedish wrought iron of
varying carbon content. The modem method is to melt " melt-
ing bar,'* a very pure open hearth steel with charcoal, or even
pure pig iron. Crucible steel is used in making razors (1.5 per
cent C), tools (1 per cent C), dies (0.75 per cent C), pens, needles,
and cutlery.
Electric heating (e.g, in the H^roult furnace), recently intro-
duced, permits the steel maker, first to wash the molten iron with
490 smith's intermediate chemistry
basic slags of high oxidizing power until perfectly pure, and then
by suitable additions to give it any required final composition.
Properties of Steel. — Cast iron can be melted and cast, is
hard when chilled, but can not be forged or rolled. Wrought
iron is slag-bearing and malleable, and is not hardened by quench-
ing from a high temperature. It is never cast. Steel is free
from slag, being cast from an originally liquid condition. If its
carbon content is high enough, it can be hardened by quenching.
Steel has also greater tensile strength* than wrought iron, and it
can be permanently magnetized. In addition, high carbon steel
can also be tempered to the required degree of hardness.
Tempering. — To understand the last fact, it must be noted
that carbon dissolves readily in molten iron, and is partly
converted to a carbide of iron (FeaC, 6.6 per cent C by weight)
in solution. As the temperature falls, the solubility of carbon in
iron diminishes. When white hot steel (up to 2 per cent C) is
suddenly chilled, there is no time for any changes to occur during
the cooling, and a supersaturated solid soliUion is obtained which
is very hard and brittle. When, however, the cooling is slow,
some of the carbon separates in minute crystals of cementite,
carbide of iron FeaC, imtil at about 700° there remains only 0.9
per cent of carbon in solution. At this temperature the solid
solution breaks down into a mechanical mixture of pure iron
which is soft, and carbide of iron which is hard. Steel is thus
a mixture, and not homogeneous, when slowly cooled. When
therefore hard, chiUed steel is heated once more for the purpose
of tempering, the extent to which the softer material is formed
depends on the temperature reached, and on the rate and dura-
tion of the cooling permitted. By varying these the degree of
hardness allowed to remain can be adjusted.
* Tensile strength or tenacity is measured by the weight (in kilograms)
required to break a wire of the metal 1 sq. nmi. in section. Lead 2.6, copper
51, iron 71, steel 91.
Ill
Ill
i!
l!if
ill
si
Si
ii
|i
B» 3
ill
IRON, NICKEL, COBALT 491
Phosphide of iron makes steel brittle when cold (" cold short ")•
Sulphide of iron makes it brittle when hot (" red short ")> and
imsuitable for forging. Hence phosphorus and sulphur are
reduced to the lowest possible amoimts.
«
Bessemer Process. — Pig iron is melted and nm into the
converter (Fig. 113), which is lined with fire-brick, measures about
15 by 8 feet, and holds 15 tons. An air-blast, entering through
one axle, blows through the metal >^s^^j^
from tuyeres at the bottom. The J ^^^
oxidation of the carbon and silicon, | .^ 1
which raises the temperature above I T" | I
the melting-point of pure iron, is ^W^[0^
over in 20 minutes. Spiegeleisen ^
. Fig. 113
is then added to give the desired
percentage of carbon and manganese, and the liquids, first the
slag and then the iron, are poured into ladles, and the metal is
cast.
Sulphur and phosphorus are not removed by the air. If present
in too great amounts, they are removed by lining the converter
with basic material such as magnesium and calcium carbonates
(Thomas-Gilchrist process). The slag then contains phosphates,
and is valuable as a fertilizer.
Bessemer steel is used for heavy and light machinery castings,
and is rolled into bridge and structiu^l iron. It contains from
0.1 per cent (soft) to 1 per cent (hard) of carbon.
Operi'hearth {Siemens- Martin) Process. — In this process
pig iron and scrap iron (up to 75 tons) are melted on a hearth lined
with fire-brick and sand (Fig. 114). At a later stage haematite is
added to furnish oxygen (as in puddling). To secure economically
the temperature necessary to keep the pure product (iron) fused,
Siemens contributed the idea of preheating the fuel gas and air by
a regenerative device. The spent air and gas pass out through a
492 smith's intermediate chemibtry
cheelrerwork of brick. When this has become hot, the valves
are reversed, the gas and air now enter through the hot brickwork
and pass out through the checkerwork on the opfxisite side, raising
ite temperature. The direction of the gases is changed every
twenty minutes or so, and the whole operation lasts 8 to 12 hours.
Fia 114
Towarcfe the end some alimiinium is added to combine with
oxygen (present Ets CO), and give sounder ingots. Recently, iron
containing 10 to 15 per cent of titanium has been added instead.
The titanium combmes with both nitrogen and oxygen, and
passes into the slag. Rails made of steel purified with this
element are less liable to breakage and are 40 per cent more dur-
able than ordinary open-hearth rails.
The advantage of this process over that of Bessemer is that
it is not hurried and is therefore under better control. The
material can be tested by sample at intervals, and defects cor-
rected. The product is of better and more uniform quality.
As in the Bessemer process, phosphorus and sulphur are re-
moved by using a basic lining.
IBON, NICKEL, COBALT 493
Open-hearth steel is used for the better class of rails, for railway
bridges, for shafts, armor-plate, and heavy guns, and whereVer
the steel is subject to much vibration.
Steel Alloys. — We must distinguish between manganese,
almninium, silicon, or titanium added in small amoimts (" med-
icine ") to purify the iron, and passing (in combination) into
the slag, as described in preceding sections, and the present
subject, which concerns metals added so as to produce regular
alloys.
Manganese steel (7 to 20 per cent Mn) is exceedingly hard even
when cooled slowly. It therefore does not lose its temper readily
when heated by friction. It is used for the jaws of rock-crushing
machinery and for burglar-proof safes.
Chromium-vanadium steel (1 per cent Cr, 0.15 per cent Va) has
great tensile strength, can be bent double while cold, and ofiFers
great resistance to changes of stress, and to torsion. It is used for
frames and axles of automobiles, and for connecting rods.
Tungsten steel (timgsten 8 to 20 per cent, and chromimn 3 to
5 per cent) is used for tools in high-speed metal turning, as it can
become red hot (from friction) without loss of temper.
Nickel steel (containing 2 to 4 per cent of nickel) resists cor-
rosion, and has a very high limit of elasticity and great hardness.
It is used for armor-plate, wire cables, and propeller shafts. The
alloy with 36 per cent nickel, called invar, is practically non-
expansive when heated and is valuable for meter-scales and pen-
dulum rods.
Duriron or tantiron (15 per cent Si) is rustproof and is not
attacked by sulphuric, nitric or acetic acid, hot or cold, dilute or
concentrated. Vessels made of this alloy are therefore used
industrially in acid concentrations.
Properties of Pure Iron.— Pure iron may be made by
electrolysis, or by reduction of a pure salt by hydrogen. It has
494 smith's intermediate chemistry
a white luster, is very tough, and melts at about 1510®. The
purest iron does not nest in cold water.
Ordinary iron rusts in moist air or imder water, forming a
hydrated ferric oxide 3Fe208,H20. The impurities act as contact
agents. The rust is a brittle, porous, non-adherent coating, which
does not protect the metal below. Oil protects iron from rusting
because, although oxygen penetrates the oil, being soluble in it,
moisture does not. Iron displaces hydrogen from hydrochloric
and sulphuric acids, giving ferrous salts:
Fe + 2H+->Fe+++*H2T.
The impurities — carbide, sulphide, and phosphide — produce
hydrocarbons, hydrogen sulphide, and phosphine PHs, and the
last two confer an odor on the escaping gas.
Iron bums in oxygen, and acts when heated upon steam, in both
cases producing magnetic oxide of iron Fe804 (p. 51). A thin film
of this oxide is adherent, and protects the iron (" Russia " iron).
The articles to be treated are put into a closed retort and exposed
first to a current of superheated steam and then to a current of
producer gas (p. 337) to reduce any higher oxides that may have
been formed. Watch hands, buckles and the like may be given a
protective coating by dipping them in an oxidizing bath such as
melted saltpeter. Another method of rust proofing is to immerse
iron articles in a hot solution of ferrous phosphate. This salt is
appreciably hydrolyzed in solution, and the free acid acting upon
the iron converts its surfaces into an adherent film of basic phos-
phate.
Iron hAis Two Valences. — One atomic weight of iron may
hold two or three atomic weights of a imivalent element in com-
bination. Thus, we have ferrous chloride FeCl2 and ferric chlo-
ride FeCla, the bromides FeBr2 and FeBrs, the oxides FeO and
Fe208, ferrous sulphate FeS04 and ferric sulphate Fe2 (804)3, and
so forth. It may therefore be bivalent or trivalent, according to
the chemical conditions.
IBON, NICKEL, COBALT 495
The ferrous salts are pale green and give colorless solutions,
containing ferrous-ion Fe++. The ferric salts, containing the ion
Fe"*"^^, are usually yellow in solution, on account of ferric hydrox-
ide produced by hydrolysis.
Ferrous Sulphate FeS04. — When the bath of dilute sul-
phuric acid, used in cleaning iron for making tin-plate (p. 508),
and galvanized iron (p. 449), is becoming exhausted, scrap iron is
thrown in to use up the rest of the acid. The solution gives, on
evaporation, pale green crystals of ferrous sulphate, FeS04,7H20
(copperas or green vitriol). The salt is used in making ink (see p.
499) and rouge (see p. 496), and in purifying water (p. 470).
Chlorides of Iron. — Ferrous chloride FeCl2 is obtained in
solution when iron displaces hydrogen from hydrochloric acid (p.
52), and is isolated by evaporation. The hydrate FeCl2,4H20 is
pale green, the anhydrous salt colorless. When chlorine is dis-
solved in the solution, or when the latter, acidified with hydro-
chloric acid, is exposed to the air, ferric chloride FeCls is produced :
2FeCl2 + CI2 -> 2FeCl3 or 2Fe++ + CI2 -> 2Fe+++ + 2C1-
4FeCl2 + O2 + 4HC1 -^ 4FeCl8 + 2H2O
and is familiar in the form of a yeUow hydrate FeCl3,6H20 ob-
tained by evaporation. Other oxidizing agents, such as nitric
acid, produce the same change.
Ferric chloride, in solution, has an add reaction, due to hydroly-
sis. It is reduced to ferrous chloride by shaking the solution, or
more quickly by boiling it, with iron fiUngs:
2FeCl3 + Fe -> SFeCU or 2Fe+++ + Fe -^ 3Fe++.
Hydroxides of Iron. — Ferric hydroxide Fe(0H)3 appears as
a brown precipitate when an equivalent amoimt of sodium hydrox-
ide is added to a solution of a ferric salt:
FeCl, + 3NaOH ^ Fe(0H)3 i + 3NaCl.
496 smith's intermediate chemistry
When only a little sodium hydroxide is added, the brown pre-
cipitate, formed locally, redissolves to give a deep reddish solution.
This contains ferric hydroxide in colloidal suspension. The
sodium chloride and unused ferric chloride can be separated by
dialysis (p. 442), and a pure colloidal suspension of the hydroxide
obtained.
Ferrous hydroxide Fe(0H)2 produced by precipitation, is
white when pure, but becomes quickly green and then brown by
oxidation.
Oxides of Iron. — When ferric hydroxide is heated, ferric
oxide Fe208 remains as a red mass:
2Fe(OH)8 -^ FeaOa + ZR2O t .
It is made by calcining (roasting) ferrous carbonate, ferrous sul-
phide, or ferrous sulphate:
4FeS04 + 02-^ 2Fe203 + 4SO3.
4FeS + 7O2 -> 2Fe203 + 4SO2.
The oxide, when pulverized in a ball mill, gives a powder of
more or less briUiant red color conmionly used in paints (Venetian
red and Indian red). That from ferrous sulphate is rouge, used
in polishing plate glass and lenses, and as a pigment. Yellow
ochre is a natural hydrated ferric oxide 2Fe203,3H20, which
acquires various depths of color during calcination, and constitutes
the sienna and umber used for paints.
Magnetic Oxide of Iron Fe304 is regarded as a compoimd dt
ferrous and ferric oxides FeO,Fe203. It is formed by strongly
heating ferric oxide:
GFejOs -^ 4Fe304 + O2
and is formed when iron is oxidized at a high temperature by oxy-
gen, air (blacksmith's scale), or steam. It can be magnetized, and
natural specimens are often strongly magnetic (lodestone).
IRON, NICKEL, COBALT 497
Ferrous oxide FeO is a black substance made by cautious
reduction of ferric oxide by a stream of hydrogen.
Ferrous Carbonate FeCOs. — The carbonate occurs in nature
as an impurity in clay, in clay iron stone, and pure as siderite.
Water containing carbonic acid dissolves it, giving the bicar-
FeCOs + H2CO3 -^ Fe(HC08)2.
Thus well and river waters all contain at least traces of ferrous
bicarbonate as a part of their hardness. Exposure to the air
causes oxidation and, as ferric carbonate is not stable, rust (ferric
hydroxide or a hydrated ferric oxide) is deposited:
4Fe(HC08)2 + O2 + IOH2O -^ 4Fe(OH)3 + 8H2CO8.
This red deposit is seen in white vessels in which such water drips
or stands. It also " yellows " goods washed in such water, if the
carbonate of iron is not previously precipitated by soda or some
other softening agent (see p. 391).
Ferro' and Ferri-cyanides. — Potassium ferrocyanide K4Fe-
(CN)6, or yellow prussiate of potash, is a pale yellow, soluble salt.
The iron is contained in the negative radical and ion [Fe(CN)6]",
and the solution therefore gives the reactions of this ion, and not
of ferrous- or ferric-ion. One of the double decompositions of this
«alt — namely, that with ferric salts — is important because it
gives a gelatinous precipitate of Prussian blue "(ferric ferrocy-
anide) :
4FeCl3 + 3K4Fe(CN)6 -^ Fe^^ (Fe(CN)6)3^ i + 12KC1.
Prussian blue is employed in making paints, and is the usual
pigment in laundry blueing. Although insoluble, it is such a fine
powder that it appears to dissolve in the water. It is used in the
laimdry to correct the yellowish tint derived from the ferrous
bicarbonate in the water (p. 391). If the goods are not freed
by rinsing from soap and soda, however, the alkali liberated by
498 smith's intermediate chemistry
hydrolysis of the latter enters mto double decomposition with the
Prussian blue and produces more rust:
Fe4(Fe(CN)e)a + 12NaOH ^ 4Fe(OH)8 i + 3Na4Fe(CN)«.
Potassium Fenicyanide K3i[Fe(CN)e]™.— The difference
between this and the preceding salt can be seen by writing the
formulse thus: 4KCN,Fe(CN)2 and 3KCN,Fe(CN)8. In the
ferricyanide the iron is trivalent and the negative ion Fe(CN)e^
is also trivalent as a whole. It is a soluble salt, of red color,
made by oxidizing the ferrocyanide. With ferric salts it gives
only a brown solution, but with ferrous salts it gives a deep blue
precipitate of ferrous ferricyanide — Tumbull's blue:
3FeCl2 + 2K8Fe(CN)e -^ Fe8(Fe(CN)6)2 i + 6KC1.
We can distinguish ferrous-don Fe"*^ from ferrioion Fe"'"'"'" by
this reaction. An equally sharp distinction is obtained by adding
potassium thiocyanate, for, although the ferrous and ferric thio-
cyanates are both soluble, the latter is blood red in color (see p.
235).
Blue^Prints. — Some ferric salts are reduced to ferrous salts
when exposed to light. Thus ferric oxalate Fe2(C204)8 will keep
in the dark, but in light gives ferrous oxalate FeC204:
Fe2(C204)8 -^ 2FeC204 + 2CO2 T •
When paper is -dipped in ferric oxalate solution and dried, and a
fern (or ink drawing on transparent paper) is placed over the
prepared sheet, sunhght will reduce the iron to the ferrous condi-
tion excepting where the fern or ink lines protect it from the light.
When the sheet is now dipped in potassium ferricyanide solution
(developer), the ferric oxalate gives only the brown substance
which can be washed out. But the parts exposed to the light turn
deep blue from the precipitation of ferrous ferricyanide in the
paper. The pattern is white on a blue groimd. In regular blue-
print paper ammonium-ferric citrate takes the place of the oxa-
IRON, NICKEL, COBALT 499
late, and the ferricyanide has aheady been appUed to the paper,
so that only exposure and washing remain to be done.
The student may make blue-prints for his own use as follows.
Dissolve 10 g. of potassiimi ferricyanide in 100 c.c. of water, and
13 g. of anmionium-ferric citrate in a second 100 c.c. Mix equal
volumes of the two solutions and filter if there is any precipitate.
Paint evenly over the paper with a clean camePs-hair brush,
dry, and keep in a dark place imtil required.
Ink. — Writing ink is conmionly made by adding ferrous
sulphate to an extract of nut-galls. The active constituent of
this extract is tannic acid, useful also in dyeing (p. 475)
and tanning leather (p. 533). Tannic acid is not, strictly speaking,
a single substance, but a mixture of complex phenolic acids (see
p. 353). Ferrous tannate is soluble and almost colorless, but
is slowly oxidized, when exposed to air, to the insoluble, black
ferric tannate. To make the writing visible from the first, a blue
or black dye is added to the ink.
Stains of fresh writing ink, being soluble, can usually be washed
out with water, if the latter is used at once. After the oxidation
has occurred, the ferric tannate must be reduced again, by soak-
ing the part for 12 hours or longer in anmionium oxalate solution,
and the ink can then be washed out. Rust stains are often ren-
dered soluble by ammonium oxalate also.
Cobalt Co and Nickel Ni
These metals, like iron, are attracted by a magnet. Cobalt,
like iron, has two series of compounds, in which it is bivalent and
trivalent, respectively. In its salts, nickel is bivalent only, but
the oxide and hydroxide Ni203 and Ni(0H)8 are also known.
In olden times ores containing cobalt and nickel were frequently
mistaken for iron and copper ores, and were treated accordingly.
Failure to isolate the expected metals was regarded as due to
supernatural influences, hence the word cobalt is derived from the
500 smith's intermediate chemistry
German Kobald, an evil spirit (akin to the English " goblin ")»
while the connection between nickel and the chief of evil spirits is
obvious. Cobalt continued to justify its name until very
recently; only within the last few years have any applications of
the metal been discovered. Nickel, on the other hand, has long
ago found many uses.
Cobalt. — The metal has a silvery luster, tinged faintly with
pink, and does not tarnish. It displaces hydrogen very slowly
from dilute acids, but is acted upon rapidly by nitric acid.
An alloy of cobalt, chromium and tungsten (stellite) is used for
high-speed tools. When heated, it keeps its temper better than
the steel alloys (p. 493).
The oxide is used as a pigment in the glass and china industries
(p. 362).
Cobaltous chloride CoCl2,6H20 is red in color, and when par-
tially or wholly dehydrated becomes deep blue. Writing made
with a dilute solution of this salt leaves pink traces so faint as to
be invisible. But, when the paper is warmed, the hexahydrate is
decomposed, and the writing appears blue. When the breath
is now blown on the writing, it disappears once more (sympathetic
ink).
Metallurgy of Nickel. — Nickel occurs in all iron meteorites.
The chief source of nickel is pentlandite, a mixture of nickel,
copper and iron sulphides, from Sudbury, Ontario. The ore is
roasted and smelted and finally bessemerized (p. 491). The
result is an aUoy of nickel and copper which is much used for
sheet metal work under the name of monel metal. Pure nickel
is separated from the copper by an electrolytic process (see copper,
p. 515), or by the Mond process (see below). In 1918, more than
1,500,000 tons of nickel ore (approximating 3 per cent Ni) were
smelted.
Properties and Uses of the Metal. — Nickel is a white, hard,
malleable metal which takes, and keeps, a high polish even in
IRON, NICKEL, COBAIiT 601
moist air. Nickel-plating, deposited electrolytically on iron, has
the same qualities. The metal is used also in alloys such as monel
metal (copper, nickel, approximately 1:1), German silver (cop-
per, zinc, nickel, 2:1:1), and nickel coinage (copper, nickel,
3:1). Nickel steel is used for armor-plate.
Compounds of Nickel. — These salts are green, and give
green solutions (Ni"*^). The sulphate NiS04,7H20 is familiar,
as is also the double salt, ammonium-nickel sulphate (NH4)2S04,
NiS04,6H20, used in nickel-plating. Nickel carbonyl Ni(C0)4 is
a volatile, colorless liquid (b.-p. 43°). It is formed by passing
carbon monoxide over warm, finely divided nickel, and is decom-
posed again, yielding nickel 99.6 per cent pure, by heating at
180°. These actions are the basis of the Mond process for
separating nickel.
Exercises. — 1. Why does paint protect iron from rusting?
2. Why does iron, in time, turn completely into rust, while
zinc and aluminimn receive only a slight film of tarnish?
3. How could it be ascertained that ferric hydroxide is in col-
loidal suspension, and not dissolved?
4. Make equations for: (a) the interaction of potassium ferro-
cyanide and cupric sulphate CUSO4, giving a brown precipitate of
cupric ferrocyanide; (b) the action of hydrochloric acid on ferric
hydroxide (or rust).
CHAPTER XLII
LEAD AND TIN
Lead and tin are the best-known metallic members of the fam-
ily to which the non-metals silicon and carbon also belong. In
their compoimds they are bivalent or quadrivalent.
Lead Pb
Metallurgy of Lead. — The chief ore of lead is galenite PbS.
The ore, if rich, is roasted in a reverberatory furnace (Fig. 112, p.
489) until a part has been converted into the oxide PbO and sul-
phate PbSOi. The air is then shut oflf, and the temperature
raised, so that these products may oxidize the remaining galenite:
PbS + 2PbO -> 3Pb + SO2 T
PbS + PbS04 -^ 2Pb + 2SO2 T .
The melted lead flows out.
Ores poorer in lead are sometimes reduced by heating with
scrap iron, or with a mixture of iron ore and coke.
The lead usually contains silver and some gold, which are re-
covered by Parke's process (see silver, p. 519).
Properties of Lead. — Lead is a soft malleable metal, with a
bluish-white or grey luster when freshly cut. It is quickly, but
only superficially, oxidized by the air. It acts very slowly upon
hydrochloric acid, but not upon sulphuric acid (see p. 268). With
nitric acid it gives lead nitrate Pb(N03)2 and oxides of nitrogen.
Soft water, in presence of air, dissolves it in appreciable amoimts
as hydroxide Pb(0H)2, and carbonic acid assists the process.
Hard water, however, produces a skin of carbonate and sulphate
(insoluble) which protects the surface. Hence, lead pipes may
502
LEAD AND TIN 503
be used for hard water, but not for rain water. In presence of air,
acids (even feeble vegetable acids) interact with the metal, which is
therefore unsuitable for kitchen utensils.
Uses of Lead. — The metal is rolled into sheets, which are used
for lining t^nks. Lead pipes are made by pressing the metal,
while hot, through dies. Their pliability, and the ease with which
they can be cut and soldered, fits them for use in plumbing and for
covering electric cables.
Type metal) containing 20 per cent of antimony, is harder than
lead, and expands on solidifying. Small shot (p. 340) contain 0.5
per cent of arsenic. Solder (lead, tin 1:1) remains melted at a
lower temperature than pure lead (m.-p. 326°) and so can be
applied to a lead joint without danger of melting the lead itself.
Oxides. — Lead monoxide PbO is made by oxidizing melted
lead in a current of air. At low temperatures a buff colored
powder, massicot, is obtained. When the oxide is allowed to melt,
it solidifies to a reddish-yellow, scaley mass — litharge. The
oxide is predominantly basic, absorbing carbon dioxide from the
air and, with acids, giving salts. It is used m making glass and
enamels. Stone and glass can be cemented with a mixture of
massicot and glycerine.
Minium or red lead Pb304 is a bright-red powder formed by
oxidizing the monoxide in air at 470 to 480°:
6PbO + 02^ 2Pb804.
Red lead is used in making flmt glass and m paints.
Lead dioxide Pb02, is a brown powder, made by treating red
lead with diluted nitric acid.
Paints. — A paint usually contains three ingredients: 1. The
oil, which hardens (" dries ") to a tough resin, being oxidized by
the air, and adheres firmly to the surface being painted.
604 smith's intermediate chemistry
2. The body, a fine powder which makes the paint opaque.
Since the powder does not shrink, it also " fills " the paint and
prevents the fonnation of minute pores which otherwise would
appear in the oil after drying. White lead (see below) is the most
common material for the body, but zinc oxide and other sub-
stances are also used.
3. Except in the case of white paint, a pigment is added.
Various oxides, such as minium, colored salts, and lakes (p. 476)
are used as coloring matters.
The oil does not " dry " by evaporation but gives a resin by
oxidation (see p. 41). Linseed oil and hemp oil are conunonly
used. They contain glyceryl esters (p. 432) of unsaturated acids,
such as that of linoleic acid, (C8H5(C02Ci7H3i)8) . The unsaturated
part of the molecule takes up the oxygen. By previously boiling
the oil with manganese dioxide and other oxides, it is rendered
more active, and " dries " more quickly.
Pliunbers use a cement made of minium and linseed oil, in which
the former oxidizes the latter, without access of air being neces-
sary, to make joints tight.
White Lead. — White lead is a basic carbonate 2PbC03, Pb-
(0H)2. It is a heavy, white, insoluble, amorphous substance.
Mixed with linseed oil, it forms a white paint valued for
its " body " or covering power (opacity). Its disadvantage is
the darkening, due to formation of the blade lead sulphide PbS,
which is produced by the hydrogen sulphide in the air (see p. 223).
Its poisonous character is also objecti6nable.
The old Dutch process for making white lead is still used exten-
sively. Gratings (" buckles ") of lead are placed above a little
vinegar in small pots. The pots are covered with boards, on
which manure or spent tan bark is spread. Other tiers of pots,
boards and bark are placed on the first, until the shed is full. Car-
bon dioxide, warmth and moisture are furnished by the decajdng
bark. The gratings, by the end of three months, are converted
LEAB AND TIN 606
into cakes of white lead. The vapor of acetic acid arising from
the vinegar may be regarded as a catalytic agent.
In Mild's process melted lead is atomized by a Jet of steam, and
the lead dust is beaten with vinegar, ak, and carbon dioxide in a
vat for about seven days. In the French process white lead is
precipitated by a stream of carbon dioxide from a solution of the
basic acetate.
Other Compounds of Lead. — Lead chromate PbCr04 is
precipitated by adding potassiimi chromate solution to a solution
of a salt of lead. It is used as a pigment (chrome-yellow). Lead
chloride PbCU (white) is very Uttle soluble in cold water and the
iodide Pbl2 (yeUow) is insoluble. Both are formed by precipita-
tion. Lead sulphide PbS (black) is pryecipitated by hydrogen
sulphide, even from acid solutions (see p. 253). The sulphate
PbS04 is a very insoluble salt. On this accoimt, the workmen
in white lead works add a little sulphuric acid to the water they
drink.
Zinc, or any of the metals more active than lead, when placed
in a solution of a soluble salt of lead, will displace the metal, and
deposit it in a mossy form (" lead-tree ") :
Zn + Pb-H--^Zn"H- + Pb j .
The Storage Battery. — In the ordinary lead accumulator the
plates consist of leaden gratings. The openings in these gratings
are filled with finely divided lead in one plate and with lead dioxide
in the other. These, and the dilute sulphuric acid in the cell, are
the active substances when the cell is charged. When the battery
is used, the SO4" ions migrate towards the plate filled with the
lead (Fig. 115), and convert this lead into a mass of the insoluble
lead sulphate: SO4" + Pb -» PbS04 + 2 0- This plate there-
fore becomes negatively charged. Simultaneously, the H"*" ions
move towards the other plate and there reduce to monoxide the
lead dioxide with which it is filled.
PbOa + 2H+-^H20 -h PbO -h 2 ®.
506
smith's INTERBfEDIATE CHEBHSTRY
This plate consequently becomes positively charged and, by inter-
action of the lead monoxide with the sulphuric acid, becomes filled,
like the negative-plate, with lead sulphate. During the dischai^,
much sulphuric acid is thus removed from the cell fluid, and the
approaching exhaustion of the cells can therefore be ascertained
PbOi
PbSO^
Fig. 115
Fig. 116
by measuring the specific gravity of the fluid. The E.M.F. of the
current is a Uttle over 2 volts.
The cell may be recharged by passing a high-voltage current
through the cell, in the opposite direction (Fig. 116). The H+
ions are attracted to the negative plate and an equivalent munber
of SO4" ions are formed, so that only lead remains:
PbS04 + 2H+ + 2 0 ~> Pb -f- 2H+ + SOr.
Simultaneously, the SO4" is attracted by the positive plate and,
with the lead sulphate there present, forms lead disulphate:
SO4" + PbS04 + 2 ® ~> Pb(S04)2. The disulphate is at once
hydrolyzed and the filling of this plate is thus changed into lead
dioxide : Pb(Sai)2 + 2H2O -> PbOa + 2H2SO4. Both plates are
thus brought back to the condition in which they were before the
discharge.
LEAD AND TIN 507
The lajst set of charges consumes energy, whUe the first set
liberates energy. Both may be stated m a smgle equation:
charge — »
2PbS04 + 2H2O ^Fh + 2H2SO4 + PbOa.
^ discharge
In the Edison cell, when charged, one plate is of iron and the
other contains nickelic oxide Ni208. The cell liquid is a solution
of potassium hydroxide. When the cell operates, the nickelic
oxide is reduced to Ni(0H)2 and the iron is oxidized to Fe(0H)2,
an action which delivers energy:
>
Fe + 3H3O + NizOa ^ Fe(0H)2 + 2Ni(OH)2.
When the cell is recharged, the nickel is reoxidized and the iron
reduced.
Tin Sn
Metallurgy. — Tin is obtained from cassiterite or tin-stone
SnOg (stannic oxide). The world's production averages 120,000
tons antiually. Formerly the mines in Cornwall (England) con-
stituted the chief source of the metal, but now the largest supply
comes from the East Indies, the next largest from Bolivia. The
ore is roasted to expel sulphur and arsenic, and reduced with coal
in the reverberatory furnace. The melted metal is cast ui blocks
(" block tin ")• The metal was well known to the ancients (found
in Egyptian tombs).
Properties. — The metal is white, and markedly crystalline.
It is soft and maUeahle (tinfoil), and melts at 232°.
Tin does not tarnish in the air. With concentrated adds it
acts rapidly. Hydrochloric acid gives stannous chloride SnCl2 and
hydrogen. Concentrated sulphuric acid gives stannous sul-
phate SnS04, sulphur dioxide (p. 271) and water:
Sn + 2H2SO4 ^ SnS04 + SO2 1 + 2H2O.
508 smith's INTERBfEDIATE CHEMISTRY
Nitric acid gives the insoluble, white meta-stannic acid H^Qt:
4HN0, + Sn-^HaSnO, i + 4NO2 + HjO.
Uses. — Tin plate, used in making "tin" cans, is produced by
dipping cleaned sheets of mild steel in melted tin. So long as the
layer of tin remains intact, the iron is protected from rusting.
But, if the tin layer is damaged, the iron rusts. The iron being
the more active metal of the two, the tin acts as a contact agent
and actually hastens the rusting.
Tin is suflSciently valuable to render the detinning of scrap
tin plate from can factories, bearing 3 to 5 per cent of tin, a pay-
ing process. In the Goldschmidt process the scrap is cleaned,
dried, and exposed to dry chlorme, which converts the tin into the
liquid stannic chloride SnCU, but leaves the iron imaflfected. The
chloride is used in mordantmg.
Copper vessels for cooking and brass pins are also coated with
tin, to preserve them from the action of air and moisture. Tin
pipes are used where lead would be unsafe, as, for example, for
beverages.
Tin enters into many useful alloys, such as bronze (with copper)
and solder.
Compounds of Tin. — Stannous chloride SnCl2,2H20, "tin
crystals," and a hydrate of stannic chloride SnCl4,5H20 are used
in mordanting (p. 475). With soda the former gives stannous
hydroxide Sn(0H)2 and the latter stannic acid Sn(0H)4, both
gelatinous substances which are precipitated in the goods to be
dyed. Stannic acid is also sometimes precipitated in flannelette
(a material made of cotton) to render it non-inflammable, and
always in silk to " weight " it (the increase may be from 25 per
cent to 300 per cent or more).
Stannous sulphide SnS (brown) and stannic sulphide SnSs
(yellow) are precipitated by hydrogen sulphide, even from acid
solutions of stannous and stannic salts, respectively.
LEAD AND TIN 609
Exercises. — 1. Make equations for: (a) the action of air and
water on lead to form the hydroxide; (b) the precipitation of lead
chromate by adding K2Cr04 to lead nitrate solution; (c) the action
of hydrogen sulphide on white lead; (d) the precipitation of
PbCl2 and of Pbt.
2. Tm melts at 232^ and lead at 326°. Solder (1 : 1) melts
(and solidifies) at 210°. Why is this? What is the advantage of
using solder?
3. Make equations for: (a) the action of soda on solutions of
stannous and stannic chlorides, respectively: (Jb) the action of
hydrogen sulphide on the same salts.
CHAPTER XLIII
COPPER AND IfBRCURY
In this chapter we encounter the first metals (if we except As
and Sb) which are below hydrogen in the activity list, and do not
displace that element from dilute acids. Copper and mercury
both have two valenceSy so that we have cupric-ion Cu++ and cu-
prous-ion Cu"*" and mercimc-ion Hg+^" and mercurous-ion Hg+.
All the soluble compoimds of both are poisonous.
Copper Cu
Occurrence. — Copper occurs free in considerable amounts,
particularly on the Michigan shore of Lake Superior. Cuprous
oxide CU2O and basic carbonates, like malachite CuC08,Cu(0H)»,
are less common. The latter is often used as an ornamental stone.
A large proportion of commercial copper is obtained from chal-
copyrite CuaSjFcgSa.
Metallurgy. — The free copper, after being concentrated
(freed from gangue) by washing, is smelted with a flux. The
carbonate is roasted, leaving the oxide. The oxides are reduced
with coal.
The sulphide ores are more difficult to reduce, and the presence
of so much iron complicates the process. They are first roasted.
This removes much of the sulphur as sulphm* dioxide, leaving
CU2O and Fe208. Next the roasted material is treated in a blast
furnace, along with " green " (imroasted) ore, sand (if silica is not
present in the ore), and coke. Some of the iron is oxidized and
removed in the slag (as silicate). The product, known as cppper
" matte,'' is a mixture of cuprous sulphide CuS with ferrous sul-
phide FeS. The third stage is to bessemerize the melted matte
510
COPPER AND MERCURY 611
with sand in a converter. Here the rest of the iron is oxidized and
eliminated as silicate in the slag, and the sulphur escapes as SO2.
The slag and metallic copper are poured separately. The latter
gives oflf some dissolved sulphur dioxide in bubbles as it solidifies
and, from its appearance, is named blister copper. Finally y since
the copper now contains dissolved cuprous oxide CU2O, the blister
copper is melted and " poled," by stirring with green wood. The
gases (hydrocarbons, etc., p. 435) given oflf by the heated wood
reduce the oxide to copper. If the copper is to be refined electro-
lytically (p. 485), it is then cast in plates 3 feet square and f inch
thick.
The old methods of concentrating copper sulphide ores
by simple washing left large amoimts of copper in the rejected
gangue. Recent flotation processes prevent this loss. The ore
is finely ground and beaten up with water containing a small
quantity of oil. The particles of copper sulphide become wetted
by the oil, the particles of gangue are preferentially wetted by the
water. When air is forced through the pulp, the copper sulphide
particles float with the froth to the top and are scraped oflf, while,
the gangue sinks to the bottom. Sixty million tons of ore per
year are now treated by this method.
Properties. — Copper has a characteristic bright yellow-pink
luster, quickly darkened by oxidation. It is second only to silver
in electrical conductivity and to iron in tenacity. It is third in order
of malleability. It melts at 1057°.
In moist air copper acquires a green coating of basic carbonate,
which protects the metal. It is not aflfected by dilute hydro-
chloric or sulphuric adds, when air is excluded. Hot concentrated
sulphuric acid gives cupric sulphate and sulphur dioxide (p. 271),
and nitric acid gives cupric nitrate and oxides of nitrogen (p. 310).
Uses. — Pure copper is used for electric wires and cables.
Traces of other metals greatly reduce the conductivity. Kettles,
512 smith's intermediate chemistry
stills, and evaporating pans are made of copper. It is used for
sheathing ships and for bolts, because it resists corrosion by sea
water.
It enters into important alloys, such as brass (18 to 40 per cent
of zinc), and bronze for coins (4 per cent tin and 1 per cent zinc),
for gun-metal (10 per cent tin), and for bell-metal (20 to 24 per
cent tin). Aluminium bronze (5 to 10 per cent aluminium) is used
for the hulls of yachts. All of these are composed in part of com-
pounds, such as CuaSn and Cu2Zn8.
Cupric Sulphate CuS04}6H20. — The hydrated sulphate,
bluestone or blue vitrioli is made by continuously agitating granu-
lated copper with air and warm dilute sulphuric acid:
2Cu + O2 + 2H2SO4 -> 2CUSO4 + 2H2O.
The blue crystals form on strips of lead suspended in the warm,
saturated solution.
Because of slight hydrolysis, giving an active acid and a weak
base Cu(0H)2, the aqueous solution is add in reaction.
Cupric sulphate is used in battery solutions. The salt is em-
ployed in minute proportions to destroy algae, which otherwise
confer a disagreeable taste and odor on water that has been stand-
ing in reservoirs. Seed for cereals is moistened with a dilute solu-
tion to prevent " smuts."
Cupric Hydroxide Cu(0H)2. — The hydroxide is a ftZtie,
gelatinous precipitate , formed when an alkali is added to cupric
sulphate solution. It is used as a mordant. A mixture of cupric
sulphate solution and milk of lime (Ca(0H)2), — Bordeaux mix-
ture — containing this precipitate, is used extensively as a spray
on grape vines and other plants, to prevent the growth of fungi.
Cupric hydroxide dissolves in ammonium hydroxide solution,
giving a liquid of deep blue color. The explanation of this, accord-
ing to the ionic hypothesis, is as follows. The solubility product
COPPER AND MERCURY 613
of copper hydroxide in water is very small; in other words, we
cannot have a high concentration of Cu"^ and 0H~ simultane-
ously in solution. But when ammonia is added to a solution
containing the very small amount of Cu"^ and 0H~ requisite to
be in equilibrium with precipitated Cu(0H)2, it combines with
the Cu"+^" to form complex ions of the formula [CuCNHa)^"'^. To
replace Cu"'"'", more Cu(0H)2 goes into solution, and the process
continues, if enough ammonia is added, until all Cu(0H)2 has
been dissolved. We shall meet with other cases of a similar char-
acter later; in fact the principle involved — the formation of a
complex ion — is used very extensively in analysis (see p. 538)
and in industrial processes (see p. 523).
An ammoniacal solution of cupric hydroxide is employed as a
solvent for cellulose. It dissolves also in a solution of sodium-
potassium tartrate NaKC4H406 (p. 354), giving Fehling's solu-
tion, a reagent used in testing for glucose and similar reducing
agents. When this reagent is added to a liquid containing glu-
cose (p. 401), red cuprous oxide CU2O is precipitated.
Cuprous Oxide CU2O. — This oxide, mixed with CuO, is
formed by gentle heating of copper in air, and is best prepared by
use of Fehling's solution. It is employed in making ruby glass
and in coloring porcelain.
Cupric Oxide CuO. — When the liquid containing the blue
precipitate bf cupric hydroxide is boiled, the blue color changes
to black and cupric oxide is thrown down:
Cu(0H)2 -> CuO + H2O.
This oxide is used in the laboratory to ascertain the composi-
tions (and formulffi) of organic compoimds (determination of car-
bon and hydrogen). A weighed amount of the organic compoimd
is placed in a horizontal tube, between heated masses of the oxide.
A stream of oxygen or air carries the vapor of the organic com-
514 smith's intermediate chemistry
poiind over the cupric oxide, which oxidizes it to water and car-
bon dioxide. The first is absorbed in a weighed U-tube filled with
calcium chloride, and the second is caught in a weighed vessel con-
taining potassium hydroxide. From the increase in weight in
each case, the corresponding weights of hydrogen and carbon
(derived from the weighed portion of the organic compoimd) are
calculated.
Tests for Copper. — The blue color of cupric salts in dilute
solution is distinctive (p. 175). Hydrogen sulphide precipitates
cupric sulphide CuS (brownish-black) even from acid solutions
of cupric salts (see p. 463).
More active metals, such as zinc or iron, displace copper from
solutions of its salts, so that a blade of a knife, for example, re-
ceives instantly a red coating of copper when immersed in such a
solution:
Fe + Cu++->Fe++-hCui.
Copper Plating. — When platiniun or carbon plates, connected
with a battery, are immersed in a solution of cupric sulphate,
copper is deposited on the negative plate (cathode). The SO4"
migrates towards the positive wire (anode) and there produces
oxygen and sulphuric acid:
2SO4 + 2H2O -> 2H2SO4 + 02!-
If the anode is made of copper itself,
however, the SO4" migrates, but is
not discharged. Instead, copper goes
into solution (Fig. 117) as Cu++, in
amount equal to that deposited on the
other plate. Thus the quantity of
cupric sulphate in solution remains
unchanged, and the effect is, virtu-
ally, to transfer copper from the copper anode to the cath-
ode.
COPPER AND MERCURY
515
Electrotypes. — A copper electrotype of an object like a
medal is made by first preparing a cast of the medal in plaster
of Paris or wax. The surface of this is rubbed with graphite,
to render it a conductor, and the cast is used as the cathode
in a cell like that just described. The deposit of copper, when
stripped oflf, is found to show an exact reproduction of the
engraving, etc., on the object.
Book plates are made by taking a cast of each page of type,
preparing the copper electrotype, and then strengthening and
thickening it by filling the back with melted lead. The printing
is then done from the electrotype. For newspapers this process
is too slow, and the plate is made from the cast by means of melted
stereotype-metal (lead, antimony, tin; 82 : 15 : 3).
fjjijujjjjjfjjjjjijijjijjjfjjTTm
Copper Refining. — The copper, as obtained from the ores by
the treatment already described (p. 510), contains a certain
amoimt of silver, gold, and baser metals. The former pay for the
cost of refining, and the simultaneous ^ ^
removal of the latter gives pure copper
suitable for electrical purposes.
The principle is the same as that used
in electroplating. The heavy plates of
poled copper (p. 511) are himg at intervals
in large, lead-lined vats of copper sulphate
solution and form the anodes (Fig. 118,
diagrammatic, view from above). The
metal is deposited on thin sheets of copper,
which are coated with graphite to permit
the deposit to be easily stripped ofif . These
sheets hang in the vat between the anodes,
and are connected with the negative wire. The copper, along
with such traces of more active metals, like zinc, as are present, is
ionized and goes into solution, until the anode is reduced to a
skeleton and is exchanged for a fresh one. The less active metals,
[fjJJJJJJJJJJJJfJJJJJJJJJJ^WTgJ
Fig. 118
616 smith's intermediate chemistry
such as silver sCnd gold, as well as traces of sulphides, are not
ionized. They fall to the bottom of the vat, as a sort of heavy-
mud. At the cathode copper ions alone are discharged, and
deposited, because copper is the least active of the metals present
in ionic form. In this way copper, 99.8 per cent pure, can be
obtained, and gold and silver are recovered from the mud.
Nickel Plating. — Here the bath contains an ammoniaeal
solution of ammonium-nickel sulphate (p. 501), and a plate of
nickel forms the anode. The article to be plated is carefully
cleaned, to secure a uniform deposit, and is suspended as cathode
in the vat. The surface of the deposit is afterwards burnished.
Mercury Hg
Metallurgy. — Mercury occurs both free and as mercuric sul-
phide HgS, cinnabar. Most of the ore comes from Califomia and
Spain.
The ore is roasted, sulphur dioxide escapes, and the vapor of
mercury is condensed in long, tortuous flues.
Properties. — Mercury is a liquid at ordinary temperatures,
hence its name, quicksilver (i.e. live silver). It freezes at —40**,
and boilsy giving an invisible, non-conducting vapor, at 357°. The
vapor density shows the molecules to be monotomic, as indeed are
the molecules of all metals of which the densities have been meas-
ured.
The metal has a silvery metallic luster , which is not aflfected by
the air, and a high specific gravity.
Mercury dissolves other metals, forming alloys or amalgams.
Mercury, when moderately heated, combines with oxygen, form-
ing mercuric oxide (red), but the action is reversible, and the
oxide is decomposed by stronger heating (p. 15). It combines
readily with sulphur and the halogens. Dilute hydrochloric and
sulphuric adds are not attacked by mercury. C!oncentrated
nitric acid attacks and dissolves it readily.
COPPER AND MERCURY 617
Uses. — Mercury is used in filling thermometers and barom-
eters. Sodium amalgam is used in the laboratory and the zinc
plates of batteries are amalgamated superficially to protect them
when the battery is not in use. Dentists fill teeth with mixtures
of mercury with silver, copper, cadmium and other metals, which
quickly set to a solid amalgam. The pulverized ores of gold and
silver, mixed with water, are allowed to trickle over thin layers of
mercury. The latter dissolve the particles of the precious metals,
while the sand passes on.
Compounds of Mercury. — As in the case of copper, there are
two sets of compoimds — mercurous Hg' and mercuric Hg". All
the common salts are completely volatile, with or without decom-
position, when strongly heated (vapor poisonous!).
Mercuric chloride or corrosive sublimate HgCl2 is made by
subliming a mixtiu'e of mercuric sulphate and sodium chloride.
It is a white substance, soluble in water. Hydrogen sulphide
precipitates mercuric sulphide HgS (black) from the solution.
Mercuric chloride is a violent poison. Albumen (white of eggs)
forms msoluble compounds with it, and is used as an antidote.
Mercurous chloride or calomel HgCl is precipitated as an
amorphous white powder when a chloride is added to a solution of
a merciu'ous salt. It is used in medicine to stimulate all organs
producing secretions. Mercuric fulminate Hg(0NC)2 is used in
percussion caps (p. 483).
Baser metals precipitate mercury from solutions of its salts.
The grey deposit is best seen on a clean strip of copper foil:
Cu + Hg++ -> Cu++ + Hg i .
Exercises. — 1. Make an equation: (a) for the oxidation of
ethyl alcohol by heated cupric oxide; (b) for the precipitation of
cupric sulphide from cupric sulphate solution.
2. When we electrolyzed sodixun chloride solution (p. 139),
hydrogen was liberated at the cathode. What principle, used in
518 smith's INTERBfEDIATE CHEMISTRY
the electrolytic refining of copper, does this phenomenon illus-
trate?
3. How could you recognize cupric sulphate solution by show-
ing that it contained, (a) cupric-ion, (b) sulphate-ion (p. 270)?
4. Make an equation for the liberation of mercury from cinna-
bar.
5. What gases should you collect over mercury?
6. What other salts, beside those of mercury, are volatile?
7. How should you recognize a salt of mercury in solution?
CHAPTER XLIV
SILVER, GOLD, PLAHNUM
Descending the activity list, we now reach the "noble'' metals,
which are the least active. They do not displace hydrogen from
dilute acids. They do not combine with oxygen, even when
heated.
Silveb Ag
Occurrence. — Native silver is found in many localities, usually
in small amounts. The chief supply of the metal is obtained
from the ores of lead and copper, which contain silver sulphide
Ag2S. The chief localities are in California, Australia, and Mexico.
Metallurgy. — After the lead (bearing silver) has been ex-
tracted from the ore (p. 502), it is melted in large caldrons, a
small proportion of zinc is added, and the whole is vigorously
stirred (Parke's process). Zinc is only sUghtly soluble in lead,
but it combines with silver in several proportions. The zinc-
silver alloy rises to the surface, solidifies (while the lead is still
molten), partly as alloy and partly as a compound (usually Ag2Zn6),
and is skinmied off. The most of the adhering lead is pressed out,
and the compound (or mixture) is placed in graphite retorts, in
which the zinc is removed by distillation. The silver and lead
which remain are heated in a blast of air (cupelled) to oxidize
the lead. The melted litharge flows oflf and the silver is then cast.
The gold, which accompanies the silver through this treat-
ment, is separated electrolytically (see copper, p. 515). The
silver-gold alloy forms the anode and silver nitrate the vat-liquid.
The silver, being the more active metal, is ionized and deposited
519
520 smith's intermediate chemistry
on the cathode, while the gold collects as a powder in a bag which
surrounds the cathode.
An older method, still used, is to heat the silver with concen-
trated sulphiuric or nitric acid, either of which will dissolve the
silver and leave the gold. From the solution the silver is dis-
placed by the action of plates of copper:
Cu + 2Ag+-^ Cu++ + 2Ag i.
The world's production of silver in 1^20 was 170,000,000 ounces
(troy) , of which the mines of Mexico were responsible for 65,000,000,
and those of the United States for 55,000,000.
Properties. — Silver is fairly hardy considering its great
ductility and maUeability, It is the best conductor of electricity.
When an electric discharge passes between the ends of two
silver wires, held under water, silver is dispersed at the points
and forms a colloidal solution. The color of the solution varies
from brownish to pink, according to the conditions. Colloidal
solutions of gold and platinmn can be made in the same way.
The metal is oxidized by ozone (p. 221), although not by oxy-
gen. Sulphur compounds in the air tarnish the surface (Ag2S,
see p. 255), as do also eggs, secretions from the skin and rubber
which contains sulphur.
Cold nitric add and hot concentrated sulphuric acid are at-
tacked by it, giving the nitrate and oxides of nitrogen, and the
sulphate and sulphur dioxide, respectively:
3Ag + 4HN08 (dU.) -> SAgNOa + 2H2O + NO f .
Uses. — For silver ware and coins the metal is alloyed with
copper. American coins contain 90 per cent of silver (" 900
fine ")• British coins formerly contained 92.5 per cent, which is
the proportion in " sterling silver,'' but the rise in price of the metal
has recently necessitated a reduction to 50 per cent. Articles
of baser metal are plated with silver. The bath contains potassium
argenticyanide KAg(CN)2, made by adding potassium cyanide
SILVER, GOLD, PLATINUM 521
in excess to sUvet nitrate solution. This solution gives a compact
deposit on electrolysis. The anode is of silver, so that the silver
in the solution is replenished as quickly as it is deposited.
Mirrors are silvered by cleaning the surface and pouring over
it a solution containing silver nitrate, ammonium hydroxide, and
a reducing agent like formaldehyde CH2O, or grape sugar:
4AgOH + CH2O -> 3H2O + 4Ag i + CO2.
The film of silver adheres to the glass and is washed, dried, and
varnished.
Silver Nitrate AgNOs* — This exceedingly soluble salt is
deposited from hot solution (p. 520) in colorless crystals. Its
solution is nevtral in reaction, which shows that silver hydroxide
is not a weak base. It melts easily and is cast in sticks (lunar
caustic), which are used in cauterizing sores. It is the chief source
of the other compoimds of silver. It is used in some hair dyes,
and in indelible ink. In the latter case the organic matter in
the goods reduces it, with the help of light, to metallic silver.
Silver Halides. — Silver chloride AgCl is precipitated (white)
when a soluble chloride is added to a solution of a salt of silver:
AgNOs + KCl ^ AgCl i + KNO3.
Silver bromide AgBr is precipitated with bromide-ion and silver
iodide Agl with iodide-ion. These compoimds have a yellowish
tinge. The chloride and bromide are easily dissolved by am-
monium hydroxide solution, giving the complex ion Ag(NH8)2''"
(see p. 513), and also by sodium thiosulphate solution Na2S203
C' hypo ").
All the halides of silver are decomposed by hght, which liber-
ates the halogen, and finally leaves metaUic silver.
Photography. — The taking of a photograph involves four
processes — preparation of the plate, exposure, development, and
fixing.
522 smith's intermediate chemistry
In preparing the plate, silver bromide is first precipitated in
water containing gelatine. The mixture is kept warm, to permit
the precipitate to become more sensitive to light by acquiring a
coarser grain (" ripen "). The " emulsion " is applied to plates of
glass or strips of transparent celluloid (films).
The brief exposure of the plate to the image of the object,
well-focussed in the camera, produces no visible effect. But the
bromide is thereafter more easily reduced to metaUic silver, in
proportion to the intensity of the light that fell upon each part.
Development consists in appl3ring a reducing agent, of such
sUght activity that its efifect during the process on non-illumi-
nated parts of the bromide is practically zero. Ferrous oxalate,
or an alkaline solution of pyrogallol C«H3(OH)8 or of quinol
CftH4(OH)2 (two substances belonging to the class of phenols,
see p. 353) may be used. The reduction goes fastest and de-
posits most silver where the illumination was most intense. Thus,
the plate becomes most opaque where the object was brightest,
and vice versa. On accoimt of this reversal, the plate is called
a negative. With the potassium salt of quinol, quinone C6H4QS
is formed:
2AgBr + C«H4(OK)2 -^ 2Ag + 2KBr + C»H402.
The foregoing processes are all carried out in a faint red light,
which is almost without action on silver bromide. To prevent
the gradual reduction of the remainmg, unchanged bromide to
silver by daylight, it is dissolved out by soaking the plate in so-
diimi thiosulphate (fixing). The plate is now dear where no silver
was deposited. The negative is finally washed thoroughly to re-
move all except the gelatine and the silver image, and is then dried.
In printing, the prepared paper is illmninated through the nega-
tive, and light and dark are again reversed. The denser parts of
the negative protect the paper below them, and leave these parts
white. .On printing paper, silver chloride suspended in egg al-
bumen is the sensitive substance, and the silver is liberated in a
SILVER, GOLD, PLATINUM 523
reddish, colloidal condition. The color is improved by toning
with a solution containing gold chloride, as part of the silver goes
into solution and gold (purplish) is deposited in its place. The
print is fixed with hypo, washed, and dried. Papers like velox
(invented by Baekeland) are essentially like plates (silver bromide
in gelatine), and are exposed, developed, and fixed in the same
way.
Gold Au
Occurrence and Extraction. — Gold is found in the free
condition in veins of quartz and in alluvial deposits, resulting
from the breaking up of such rock. It is foimd also in cohibination
with tellm-imn (p. 273).
In vein mining (e,g. in the Transvaal) the rock is pulverized
with iron stamps working in an iron trough. The powder is
washed in the form of mud over plates of copper amalgamated
with mercury, in which 55 per cent of the gold dissolves. The
amalgam is afterwards scraped ofiF, the mercury removed by dis-
tillation, and the gold residue refined. The tailings still contain
46 per cent of the gold, adhering to the particles of rock. They
are covered with sodiimi cyanide (p. 393) solution, and exposed to
the air, imti! the gold has been dissolved as sodiimi am-ocyanide
NaAu(CN)2. From this solution the gold is deposited by elec-
trolysis, or displaced by zinc.
The alluvial deposits are washed, on a small scale, in " cradles "
(shallow pans) and, on a large scale, by being carried by water
down a long trough (Placer Mining). The gold, having a much
higher specific gravity than the rock, sinks to the bottom, while
the rock particles are carried away. In the trough the gold settles
between elects nailed across the bottom. In hydraulic mining, a
modification of placer mining, very heavy streams of water are
thrown against the deposit.
In 1920, the gold production of the world exceeded $330,000,000
in value. $165,000,000 of this came from the Transvaal, and
524 smith's intermediate chemistry
$50,000,000 from the United States. In 1912 the world yield
was $475,000,000, but the increase in the cost of working has
caused many mines to be shut down.
Properties. — Gold is yellow in color. It is the most
malleable and ductile of metals. It melts at 1075°. To enable it
to resist wear, it is alloyed with copper. Pure gold is " 24 karat "
fine. British gold coins are 22 karat, and American coins 21.6
karat (90 per cent gold).
Gold is not affected by oxygen or by hydrogen sulphide. It
does not interact with any single add. It combines directly,
however, with chlorine and bromine. It dissolves in aqua regia
(hydrochloric and nitric acids, mixed). This happens, not be-
cause aqua regia is more active as an oxidizing agent than the sub-
stances it contains, but because it oxidizes, and also furnishes the
chloride-ion Cl~ required to produce the exceedingly stable negor
tive ion of chlorauric acid HAuCU, namely AuCh".
Uses. — Most of the metal is used in coins and bars as a medium
of exchange. It is beaten into gold leaf. It is employed in
making potassium chloraurate KA^uCU for toning photographs.
Gold plating on silver and other metals is carried out by using
a gold anode and a bath of sodium aurocyanide NaAu(CN)s
solution.
Platinum Pt
Occurrence and Extraction. — Platinmn is found in the free
condition in alluvial sand, chiefly in the streams of the Ural
Mountains and the Caucasus and in Colombia. The separation
from osmium, iridium, and other metals, which accompany it, is
a complex operation.
Properties. — Platinum is a malleable and dudtHe metal with
a greyish-white luster. It meite in the oxy-hydrogen,flame (m.-p,
about 1780°), but not in the Bimspn flame,
SILVER, GOLD, PLATINUM 525
The metal is not affected by aivy water, or acidSj excepting
aqua regia. In the latter instance, chloroplatinie acid H2ptCl« is
formed.
Platinum adsorbs hydrogen and oxygen. When finely divided,
so as to present a large surface, it catalyzes powerfully many
chemical actions (pp. 262, 314). The forms used are platinum
black, a powder precipitated from chloroplatinie acid by an active
metal {e,g. zinc) ; platinum sponge, a porous mass made by decom-
posing ammonium chloroplatinate by heat; and platinized as-
bestos, made by dipping the asbestos in chloroplatinie acid
solution and heating.
It unites, slowly, with chlorine and bromine. It combines also
with carbon, phosphorus, and silicon, and alloys itself with many
metals, so that reactions which liberate any of these elements can
not be carried out in vessels of platinum. It also acts upon fused
caustic alkalies giving platinates.
Uses. — Since the metal does not melt in the Bunsen fiame, and
is not affected by many substances, it is employed in laboratory
operations in the forms of wire, foil, and crucibles. Much of the
metal is used by dentists, and in photography. Having the same
expansibility as glass, it is fused through the bottoms of incan-
descent lamps, to connect the filament with the exterior. Since
the metal is not oxidized by the air, even when heated, it is used
for electrical contacts. The metal is employed most extensively
for jewelry.
On accoimt of the increasing demand, and the failure of the
Russian supply, platinum is now (1922) worth $93.00 per ounce
troy, or more than four times the value of gold (nominally $20.67
per ounce).
Other Metals of the Platinum Group. — Associated in
small quantities with platinum in nature are five other metals
— ruthenimn, rhodium, palladium, osmium and iridium. These
526 smith's intermediate chemistry
metals all fall in the last group of the periodic system (see pp.
^78, 282). Iridium is harder than platinum, and is alloyed with
it for special purposes (pen-points and vessels to resist fluorine).
A platinum-iridium alloy is used for the international standards of
length and mass, kept in Paris. A palladiumrgold alloy (palau)
has recently been devised as a substitute for platinum in labo-
ratory crucibles.
Exercises. — 1. Make an equation for : (a) the action of con-
centrated sulphuric acid on silver; (b) the decomposition of
silver chloride by light.
2. Make an equation for: (a) the displacement of gold from
potassium chloraurate by silver; (b) the decomposition of chloro-
platinic acid, and of (c) ammonium chloroplatinate by heat; (d)
the action of zinc on chloroplatinic acid.
3. What would be the advantages and disadvantages of using
gold instead of platinum for crucibles?
CHAPTER XLV
MANGANESE AND CHROBOUM
The first metallic elements we considered form simple positive
ions {e.g. Na"*", Ca""^) only. The last two, on the other hand,
appear almost exclusively in complex, negative ions, just as do
the non-metallic elements (AuCU", Au(CN)2'~, PtCl«"). Some of
the intermediate metals can give negative ions {e.g. Zn02", AlO^),
although in most of their compoimds they are positive (Zn"*^,
Al"*"^). The two elements taken up in the present chapter are
equally famiUar in both r61es.
Manganese Mn
Ore and Preparation. — The conunonest ore of manganese is
pyrolusite Mn02, a soft, black mineral. The metal is obtained in
piu'e form by mixing the pulverized dioxide with aluminiimi
(Goldschmidt's process, p, 468) in a clay crucible and starting the
reaction with magnesiiun:
3Mn02 + 4A1 -> 3Mn + 2AI2O3.
At the high temperature (over 3000°) the molten manganese
sinks to the bottom, and the alumina floats above it.
Properties and Uses. — Manganese is hard and crystalline,
with a greyish-white luster. It is tarnished superficially by moist
air. In fine powder it slowly displaces hydrogen from boiling
water. It acts vigorously on dilute hydrochloric and sulphuric
adds, giving manganous salts:
Mn + 2H+->Mn++ + H2 1 •
The preparation and uses of spiegeleisen and ferro-manganese
(p. 488), and of manganese steel (p. 493) have already been de-
627
528 smith's intermediate chemistry
scribed. Traces of manganese seem to be essential to the growth
of plants.
The Chemical Relations of the Element. — Manganese
stands, at present, alone on the left side of the eighth coluinn of
the periodic table. The right side is occupied by the halogens.
It is never univalent, as are the halogens, but its heptoxide Mn207
and the corresponding acid, permanganic acid HMn04, are in
many ways closely related to the heptoxide of chlorine and per-
chloric acid HCIO4. Of the lower oxides of manganese, MnO
is basic, and Mn203 feebly basic. Mn02 is feebly acidic, MnOj
more strongly so, and permanganic acid (from Mn207) is a very
active acid. Contrary to the habit of feebly acidic and feebly
basic oxides, such as those of zinc, aluminium, and tin, the basic
oxides of manganese are not at all acidic, and the acidic oxides,
with the exception of Mn02, are not also basic. There are thus the
five following, rather well-defined sets of compoimds, showing
five different valences of the element. Of these the first, fourth,
and fifth are the most stable and the most important.
1. Manganous compounds, MnO, Mn(0H)2, MnS04 etc. These
compounds resemble those of the magnesiiun family (and those of
Fe++). The salts of weak acids, such as the carbonate and sul-
phide, are easily made, and there is Uttle hydrolysis of the halides.
The salts are pale-pink in color.
2. Manganic compounds, Mn203, Mn(0H)8, Mn2(S04)3, [MnCU].
The salts resemble aluminium salts in behavior, but are very
unstable (see p. 142) and are completely hydrolyzed by water.
The salts are violet in color.
3. Manganites, Mn02, H2Mn08, CaMnOa. The alkali man-
ganites are strongly hydrolyzed, Uke the siUcates.
4. Manganates, MnOa, H2Mn04, K2Mn04. The salts resemble
the sulphates and chromates, but are much more easily hy-
drolyzed. They are green in color.
MANGANESE AND CHROMIUM 529
6. Permanganates, Mn2C)7, HMn04 (hydrated), KMn04. The
salts resemble the perchlorates, and are not hydrolyzed by water.
They are reddish-purple m color.
It will be seen that the element manganese changes its char-
acter totally with change in valence, and in each form of combi-
nation resembles some set of elements of valence identical with
that which it has itself assimaed. Since the valence represents
the nmnber of electrons gained or lost by each atom (p. 217), it
is thus evident that the chemical properties of an element depend
more upon the electrical constitution of its atom than upon the
atomic weight (see p. 553).
Compounds of Manganese* — Manganese dioxide Mn02
is an oxidizing agent, and is mixed with black paints to make
them " dry " rapidly. For the same reason, and because it is
a conductor, it is used as a depolarizer in dry batteries and in the
Leclanch^ cell. It is added to molten glass to remove the green
color due to compounds of iron (p. 362).
Potassium permanganate KMn04 crystaUizes in needles which
have a deep purple color with a greenish luster by reflected light.
The purple color of the solution is the color of the ion Mn04".
The solution of potassium permanganate, especially when, by
addition of an acid, permanganic acid HMn04 has been formed,
is an active oxidizing agent. We employed it in oxidizing hydro-
chloric acid to make chlorinB (p. 141). For the same reason a
dilute solution of potassium p3rmanganate is used as an antiseptic
and disinfectant.
Chbomium Cr
Ore and Preparation. — The chief ore of chromium is
chromite FeO,Cr203, which is mined chiefly in California, Rho-
desia and New Caledonia. The U. S. production in 1918, stim-
ulated by war requirements, exceeded 80,000 tons. In 1921 it
sank to less than 300 tons.
530 smith's intekmediate chemistry
The metal is easily obtained, by reducing chromic oxide OjOj
with aluminimn (Goldschmidt process, p. 468).
Properties. — Chromium is a lustrous crystaUine metal. It
does not tarnish in air. It displaces hydrogen, however, from
dilute adds, giving chromous salts CrCU, CrS04, etc. Chromium
(3 per cent) gives a hard steel, and with nickel is used in armor-
piercing shells and in armor plate.
Alloys which, although composed entirely of active metals,
are hardly affected even by boiling acids (including nitric acid),
usually contain chromimn {e.g. 60 per cent Cr, 36 per cent Fe,
4 per cent Mo, no C). ** Stainless " cutlery contains 12 to 14
per cent of chromiimi.
The Chemical Relations of the Element. — Chromium
gives four classes of compoimds, and most of them are colored
substances (Greek, chroma, color). The chromates are derived
from chromic acid H2Cr04, which is analogous to sulphuric acid
H2SO4. The free acid, however, is itself unstable, and leaves the
anhydride CrOa when its solution is evaporated. The oxide and
hydroxide in which the element is trivalent, namely Cr203 and
Cr(0H)3, are weakly basic and still more weakly acidic. Hence
we have chromic salts such as CrCla and Cr2(S04)3 which are
somewhat hydrolyzed, but no car'jonate, and no sulphide which is
stable in water. The compounds in which the same hydroxide
acts as an acid are the chromites, and are derived from the less
completely hydrated form of the oxide CrO(OH). Potassium
chromite K.Cr02 is more easily hydrolyzed, however, than is
potassimn zincate or potassimn aluminate. Finally, the chro-
mous salts, such as CrCU and CrS04, correspond to chromous
hydroxide Cr(0H)2, in which the element is bivalent. This
hydroxide is more distinctly basic than is chromic hydroxide, and
forms a carbonate and sulphide which can be precipitated in
aqueous solution. The chromous salts resemble the ferrous salts
in being easily oxidized by the air.
MANGANESE AND CHBOMIUM 631
Potassium Chromate K2Cr04* — Powdered chromite is
roasted with potash and lime:
4FeO,Cr203 + 8K2CO3 + 7O2 -> 2Fe208 + SKzCrO* + 8CO2 T
and the potassium chromate dissolved out of the residue. It is
a yellow, soluble salt (the ion 0104*" is yellow), with which in-
soluble chromates of other metals, such as lead chromate PbCr04,
can be precipitated.
The Dichromates. — When a solution of potassium sulphate
is mixed with an equivalent amount of sulphuric acid, potassium
bisulphate is obtainable by evaporation: K2SO4 + H2SO4 — >
2KHSO4. The dry acid salt, when heated, loses water (p. 271),
giving the pyrosulphate (or disulphate) : 2KHSO4 ^ K2S2O7 +H2O,
but the latter, when redissolved, returns to the condition of acid
sulphate. The second action is instantly reversed in presence of
water. Now, when an acid is added to a chromate we should
expect the chromic acid H2Cr04, thus liberated, to interact, giving
an acid chromate (say, KHCr04). No acid chromates are known,
however, and instead of them, pyrochromates or dichromates
are produced, with elimination of water. In other words, the
second of the above actions is not appreciably reversible in pres-
ence of water when chromates ye in question :
K2Cr04 + H2SO4 -> (H2Cr04) + K2SO4.
K2Cr04 (+ H2Cr04) -^ K2Cr207 + H2O.
2K2Cr04 +H2SO4 -^K2Cr207 + H2O + K2SO4. (1)
In terms of the ionic hypothesis, S2O7" is unstable in water, and
interacts with the 0H~ ion it contains, giving water and sulphate-
ion, while Cr207' is stable in water and is formed from the inter-
action of water and chromate-ion :
S2O7- + 20H- 1=? H2O + 2SO4-,
Cr207- + 20H- ±=? H2O + 2004-. (2)
532 smith's intermediate chemistry
The dichromates of potassium and sodium are made by adding
sulphuric acid to the crude solution of the chiiomate obtained from
chromite (p. 531). They crystaUize when the liquid cools, and
the mother-Uquor, containing the potassium sulphate and im-
deposited dichromate, is used for extracting a fresh portion of
cinder. As the dichromates are much less soluble than the chro-
mates, they crystaUize from less concentrated solutions, and can
therefore be obtained in purer condition. For this reason the
extract is always treated for dichromate.
Chemical Properties of the Dichromates. — 1. When
concentrated sulphuric acid is added to a dichromate, chromic
anhydride CrOs separates in red needles:
Na^CrgOy + H2SO4 -> Na^SO* + H2O + 2Cr03 j .
2. Although a dichromate lacks the hydrogen, it is essentially
of the nature of an acid salt. Hence, when potassium hydroxide
is added to a solution of potassium dichromate, potassium chro-
mate is formed:
KaCraOT + 2K0H -> 2K2Cr04 + H2O.
The solution changes from red to yellow, and the chromate is
obtained by evaporation. In this way the pure alkali chromates
are made.
3. With free acids the dichromates give powerful oxidizing
mixtures, in consequence of their tendency to form chromic salts
(see below).
Uses of Dichromates. — When paper is coated with gelatine
containing a soluble chromate or dichromate and, after being
dried, is exposed to Ught, chromic oxide is formed by reduction,
and combines with the gelatine. This product will not swell
up or dissolve in tepid water, as does pure gelatine. .This action
is used in many ways for purposes of artistic reproduction. Thus,
MANGANESE AND CHROMIUM 533
if the gelatine mixture is made up with lampblack and, after the
coatLDg has dried, is covered with a negative and exposed to Ught,
the parts which were protected from illumination may afterwards
be washed away, while the carbon print remains. The gelatiae
layer can be transferred to wood or copper before washing. When
materials of different colors are substituted for the lampblack,
prints of any desired tint may be made by the same process.
Sodiimi dichromate is used, instead of tan-bark, in tanning
kid and glove leathers. A reducing agent is employed to precipi-
tate chromic hydroxide Cr(0H)3 in the leather. Its use dimin-
ishes the time required for the process from 8 or 10 months to a
few hours. The hide is a mixture of colloidal materials, and the
hydroxide is adsorbed.
Chromic Salts: Chrome' Alum. — The chromic salts are
made by reduction from chromates or dichromates. These latter
are derived from the anhydride CrOa, in which chromiimi is
sexivalent, while the chromic salts are obtained from the oxide
CraOa, in which the element is trivalent.
When concentrated HCl is heated with a dichromate, it is
oxidized to chlorine (compare p. 142) :
KgCrzOT + 14HC1 -> 2KC1 + 2CrCl3 + THjO + SCla.
The most important chromic salt is chrome-alum, a double sul-
phate of potassimn and chromimn, which crystallizes from solu-
tion as beautiful piu-ple octahedra K2S04,Cr2(S04)3,24H20 (see
p. 470).
Chromic Hydroxide Cr(0H)8 and Oxide Ct20z* — When an
alkali is added to chrome-almn solution, chromic hydroxide
Cr(0H)3 (green) is precipitated:
Cr2(S04)8 + 6NaOH -> 2Cr(0H)s i + 3Na«S04.
Chromic hydroxide is the substance precipitated in the leather in
chrome tanning. It is used also as a mordant in calico printing.
534 smith's intermediate chemistry
When the hydroxide is heated, chromic oxide Cr208 remains as a
green powder.
Other Elements of the Chromium Family. — Reference
to the periodic system of the elements (p. 278) shows three ele-
ments — molyhdenvmy tungsten and uranium — which should
resemble chromium in their properties and derivatives. Molyb-
denmn and tungsten, indeed, give acid anhydrides MoOa and
WO3, and the salts derived therefrom correspond to the chromates
and dichromates. The metals are used, like chromium, in making
special steel alloys (p. 493). Tungsten has a higher melting-
point (3640°) than any other metal and, on this accoimt, and
because it is less volatile than carbon, is now used for filaments
in electric lamps. A carbon filament also requires 3.25 watts per
candle power while a timgsten filament uses only 1.25 watts per
1 c.p. The powdered metal obtained by reduction can be pressed
into wire form and then rolled while strongly heated by an elec-
tric current until a compact wire is obtained.
Uraniimi, besides giving uranates and diuranates, also ex-
hibits base-forming properties. Pitchblende, which contains
the oxide UaOs along with smaller amounts of many other elements,
is found mainly in Joachimsthal (Bohemia) and in Cornwall.
Camotite, a uranate and vanadate of potassium, K20,2UQj,
V205,3H20, occurs in Colorado. Pitchblende is roasted with
lime, the calcium uranate CaU04 thus formed is decomposed with
sulphuric acid, giving uranyl sulphate UO2SO4. When excess of
sodium carbonate is added to the solution of the latter, the foreign
metals are precipitated and sodium diuranate Na2U207,7H20,
which is also thrown down, dissolves in the excess as Na2U04.
After filtration, the diuranate of sodium is reprecipitated by
neutralizing with sulphuric acid and boiling. This salt is used in
making uranium glass, which shows a yellowish-green fluorescence.
The most striking property exhibited by uranium and its
compounds, radioactivity, will be dealt with in the final chapter.
MANGANESE AND CHROMIUM 535
Exercises. — 1, Give a list of the classes of manganese com-
pounds and, in each class, the formulse of several compounds, the
prevailing color, and the valence of manganese in that class.
2. Make equations for the formation of: (a) manganous sul-
phate; (b) manganous carbonate (insoL); (c) sodimn manganate;
(d) the oxidation of hydrochloric acid by potassium permanganate.
3. What action did we find to be catalyzed by manganese
dioxide?
CHAPTER XLVI
THE RECOGNITION OF SUBSTANCES, IL — A REVIEW OF THE
IfETALUC ELEMENTS
As in Chapter XXXII, so here, we assume that the specimen to
be identified contains a single substance. We consider first the
metallic elements, and limit ourselves to those that have been
described in the context. Our review will cover, mainly, the
properties of each simple metallic positive radical.
Although arsenic has been stated to be a non-metaUic element,
and antimony to be partially non-metallic, it is more convenient
in the problem of recognition to classify them with the metallic
elements.
External Examination* — The color is often significant.
Most of the conmiori compoimds of iron, nickel, cobalt, copper,
gold, manganese and chromium are colored (see text). A metallic
luster (scrape off the tarnish) usuaUy, though not always, indi-
cates a free metal or an alloy. The crystalline form should be
noted. The odor usually gives information about non-metallic
constituents (p. 377) only. As regards state, the vast majority
of the metals and their compounds are solids. When a liquid
presents itself, therefore, it is usually an aqueous solution of
some compoimd. Obtain the soUd by evaporation.
Solubility and Reaction of the Solution. — Ascertain
whether the substance is soluble in water (p. HI). Note whether
the solution is acid, alkaline, or neutral in reaction (p. 369). No
substance can be identified by the preceding observations alone,
but the final conclusion as to the natiu^ of the specimen must be
in harmony with them.
536
THE RECOGNITION OP SUBSTANCES, II, ETC. 537
A salt which gives an acid reaction must be an acid salt of a
polybasic acid (p. 260), or a derivative of a strong acid and a weak
base. Similariy a salt which reacts alkaline must be a basic salt,
or a derivative of a weak acid and a strong base.
Recognition by ReacV.ons in Solution. — Starting with the
substance in solution, its identity can be ascertained by using
reactions involving mainly precipitations and oxidations or reduc-
tions, which separate the metals into five distinct groups.
The following plan, taken in conjunction with the statements in
the context, shows how a single cation may be identified. What
will be said applies only to the case of a solution containing salts
like the chlorides, nitrates, or sulphates of one or more cations,
and leaves the oxalates, phosphates, cyanides, and some other
salts, out of consideration.
Before attempting to understand this plan, the student should
turn to the discussion on ionic equilibria (Chapter XXXVIII)
and read it carefully through. The sections on the solvbility of
precipitates (pp. 460-4) should be particularly studied, since upon
the principles therein formulated the whole plan is directly based.*
Group /. — Add to the solution hydrochloric acid. A precipi-
tate indicates that cations giving insoluble chlorides are present.
Silver, mercurous and lead salts .give the white AgQ, HgCl, and
PbCl2 respectively. The last-named salt, being appreciably
soluble, will be only incompletely precipitated.
These three chlorides can be distinguished from one another
very easily. When excess of ammonimn hydroxide is added to
the precipitate, silver chloride dissolves (p. 521). Mercurous
chloride turns black, owing to the formation of a finely divided
mixture of free mercury and mercuric amidochloride Hg(NH2)Cl.
Lead chloride remains apparently unchanged.
* Many experimental details, essential for the successful performance of
the tests described in this chapter, are here omitted. They will be found
in the Authors' Laboratory OvUine of Intermediate Chemistry.
538 smith's intermediate chemistry
Group //. — If no precipitate appears on addition of hydro-
chloric acid, hydrogen sulphide is led into the solution. The sul-
phides insoluble in active acids, namely, HgS, CuS, PbS, BisSs,
CdS, AS2S3, Sb2S3, SnS, are therefore thrown down. The first
four are black or dark brown, the next two are yellow, and the
last two are orange and brown respectively. K too much HCl
is present, the precipitation of several of these sulphides will be
incomplete. On the other hand, if too little HCl is used, zinc
sulphide may be partially precipitated (see pp. 461-3).
This group is readily subnlivided. The last three sulphides
pass into solution when warmed with yellow ammonium sul-
phide, for they give soluble complex svlphides. The first five
sulphides will be unaffected.
Group Ha. — HgS, CuS, PbS, Bi2S8, CdS. A yellow precipitate
insoluble in ammonium sulphide mdicates cadmium. To dis-
tinguish between the remaining four sulphides, boil with HNOs;
HgS alone does not go into solution. Dilute the solution, and
add H2SO4. Lead gives a white precipitate of PbS04. If no
precipitate is obtained, add NH4OH till alkaline. BismiUh gives
a white predpitaie of Bi(0H)3, copper a blii£ solviion.
Group lib. — AS2S3, Sb2S3, SnS. The color distinction is not
always certain. Reprecipitate the sulphides by adding HCl,
and boil the precipitate with concentrated HCl. AS2S8 does not
dissolve. If the precipitate does dissolve, cool the solution and
place in it a piece of bright tin. A black deposit forming on the
surface (by displacement, see p. 54) indicates Sb.
Group Ilia. — If no precipitate is obtained with H2S, the solu-
tion is boiled, and a few drops of concentrated HNOs added to
oxidize any jerrouB salt to the jerric state. Ammonium chloride
is added, and then ammonium hydroxide in excess, A white,
gelatinous precipitate of A1(0H)8 indicates aluminium, a bluish-
THE RECOGNITION OP SUBSTANCES, II, ETC. 539
green precipitate of Cr(0H)3 indicates chromium, a reddish-brovm
precipitate of Fe(0H)3 indicates iron. The presence of NH4CI is
necessary to reduce the concentration of 0H~ furnished by the
NH4OH below the point at which other less insoluble hydroxides
(such as Mn(0H)2) would be precipitated. To tell whether iron
was originally present in the ferrous or the ferric condition, the
ferricyanide test (p. 498) should be applied.
Group Illh. — If, still, no precipitate is obtained, hydrogen
sulphide is led into the alkaline solution. Sulphides which are
insoluble in water, but soluble in active acids (see pp. 460-4),
now appear. They are CoS and NiS (both black), MnS (flesh-
colored) and ZnS (white). To distinguish between Co and Ni,
add NaOH to the original solution. CobaU gives a blii£ precipi-
tate of a basic salt, changing to pink Co(OH)2 on boiling; nickel
gives a lightrgreen precipitate of Ni(0H)2.
Group IV. — If negative results are still obtained, add
(NH4)2C03. A white precipitate is given by three of the remain-
ing metals whose carbonates are insoluble, calcium, strontium
and barium. Distinction between these three may be made by
the flame test, A small portion of the precipitate is taken up on
a platinum wire and held in the Bunsen flame. A brick-red
coloration signifies Ca, a crimson-red Sr, a green Ba.
Group V. — The only other common positive radicals are
Mg, NH4, K and Na. On addition of anmioniiun phosphate to
the solution from Group IV, magnesium, if present, is precipitated
in the form NH4MgP04 (white). An anmionium salt may be
recognized by boiling some of the original solution with NaOH,
when ammonia is evolved. Potassium salts confer a violet color-
ation to the Bunsen flame; sodium salts a bright yellow.
Confirmatory Tests. — From the context in earlier chapters,
the student will be able to pick out for each particular metal other
540 smith's intermediate chemistry
tests which may serve to confirm the conclusions arrived at in
the course of the above analysis. Often the metallic radical can
be recognized by a displacement reaction (compare pp. 514, 517).
Often, again, color changes which occur during the course of the
scheme of operations outlined above give indications of value.
Thus a yellow solution changing to green during the passage of
H2S in Group II signifies a chromate (p. 533), a purple solution
becoming colorless signifies a permanganate (p. 529).
Tests for Negative Radicals. — Precipitation reactions sim-
ilar in nature to those outlined above may be utilized as additional
tests for the negative radical of an imknown substance (see chap-
ter XXXII). Thus a chloride gives with silver nitrate a white
precipitate, soluble in ammonia. A svlphate gives a white pre-
cipitate with barium chloride^ insoluble in hydrochloric acid.
From the context, the student will be able to discover for Imnself
confirmatory tests of this kind for most of the common negative
radicals discussed in this book. For a complete scheme, how-
ever, a manual of quaUtative analysis should be consulted.
•
Insoluble Substances, — If the unknown substance is insol-
uble in water, try to bring it into solution by boiling successively
with dilute HNO3, concentrated HNOa, and agiLa regia (p. 310).
In case it dissolves, evaporate off the excess of acid and proceed
with the analysis as above. The only common substances which
are still insoluble are: the sulphates of Pb, Sr, Ba; certain mineral
oxides such as AI2O8, Cr208, Fe208, Sn02; some silicates; CaFj;
AgCl (soluble in NH4OH). Fuse with Na2C08 in a crucible, cool,
extract with water, and filter.
The residue contains the positive radical as carbonate, and may
be analyzed for this after dissolving in HN08. The fiUrate con-
tains the negative radical as sodium salt, and may be examined
as in Chapter XXXII.
THE RECOGNITION OP SUBSTANCES, II, ETC. 541
Exercises. — 1. Name some substances that have a metallic
luster, but are not metals or alloys.
2. . Name the metals whose salts with active acids will give : (a)
neutral aqueous solutions; (b) acid aqueous solutions. What
classes of salts will give alkaline solutions?
3. Write full ionic equations for the chloride precipitations
mentioned in Group I. Why is silver chloride soluble in am-
monia?
4. Write full ionic equations for the sulphide precipitations
mentioned in Group II. Which is more soluble in water, mer-
curic sulphide or cadmium sulphide?
5. Write full ionic equations for the hydroxide precipitations
mentioned in Group Ilia. Which is the more soluble in water,
aluminium hydroxide or magnesium hydroxide?
6. Show, by means of ionic equations, why it is possible to
prevent the precipitation of magnesium hydroxide when am-
moniimi hydroxide is added to a soluble magnesium salt, by
prior addition of anmionium chloride.
7. Why is it not possible, in the same way, to prevent the
precipitation of magnesium hydroxide when sodium hydroxide
is added to a soluble magnesium salt, by prior addition of sodium
chloride?
8. Write full ionic equations for the sulphide precipitations
mentioned in Group Illb. Which is the more soluble n water,
nickel sulphide or zinc sulphide?
9. Write full ionic equations for the carbonate precipitations
mentioned in Group IV.
10. Write fuU ionic equations for the reactions mentioned in
Group V.
11. Give a precipitation test for the negative radical in each
of the following substances: sodiimi hydroxide, potassiimi car-
bonate, sodium sulphide, sodium phosphate. Write fuU ionic
equations in each case.
CHAPTER XLVII
RADIUM, ATOMIC EHERGT, Ain> ATOMIC STRUCTURE
The Discovery of Radium. — In 1896 Henri Becquerel dis-
covered that a crystal of a salt of uranium could, in the dark,
reduce the silver bromide on a photo-
graphic plate, even when a sheet of black
paper (impervioua to light) was placed be-
tween. Evidently a radiation, different
from light, was given out by the salt.
Next he discovered that an electrometer
(Fig. 119), in which the go.d leaves had
been caused to separate by charging with
*'■ electricity, lost its charge rapidly when a
salt of uranium was brought near to the knob connected with the
leaves. Evidently the salt rendered the air a condudor {" ionized "
the air), and this permitted the escape of the electricity. These
discoveries, in the hands of a multitude of observers, have led to
the development of an entirely new branch of our science, namely
radio-chemistry.
The radioactivity of every pxu-e uranium compound is pro-
portional to its uranium content. The ores are, however, rela-
tively four times as active. This fact led M. and Mme. Curie,
just after 1896, to the discovery that the pitchblende residues, from
which practically all of the uranium had been extracted, were
nevertheless quite active. About a ton of the very complex resi-
dues having been separated laboriously into the components, it
was found that a large part of the radioactivity remained with the
sulphate of barium. From this a product free from barium, and at
least one million times more active than uranium, was Qnally
secured in the form of the bromide. The nature of the spectrum
and the chemical relations of the element, now named radium,
542
FCXi-THACKS FROM RADIUM (C. T. R. WILSON)
r. Paths of helium atoms. 3. Pait ol >, enlwged. 4. Paths of dectnO'
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 543
placed it with the metals of the alkaline cjarths The ratio by
weight of chlorine to radium in the chloride is 35.46 : il3, so that,
on the assumption that the element is bivalent, its chloride is
IlaCl2 and its atomic weight is 226. With this value it occupies
a place formerly vacant in the periodic table.
In 1910 Mme. Curie obtained metallic radium by electro-
lyzing a solution of radium chloride, using a mercury cathode,
and expelling the mercury by distillation. It was a white metal
(m.-p. 700°) which, like calcium, quickly tarnished in the air and
displaced hydrogen from water.
The Nature of the ** Bays." — Many properties show that
the " rays " emitted by compounds of uranium and of radium are
of three kinds. They are most sharply distinguished from one
another when allowed to pass through a powerful magnetic field..
The alpha-rays are positively charged and are bent in one direction
while the beta-rays are negative and are bent in the other. The
gammorrays are not affected.
The alpha-rays are atoms of helium (p. 297) thrown olBf in
straight lines with varying initial velocities, averaging about one-
tenth that of Kght. Each such atom bears a double positive
charge (the imit being the charge on a imivalent positive ion),
and a delicate electrometer readily indicates the impact of a
single atom. These alpha-particles, being each four times as
heavy as an atom of hydrogen, plough their way through tens of
thousands of air-molecules and usually go about 3-8 cm. before
being stopped. The emission of atoms of helium can be de-
tected by means of Crookes' spintharoscope (Fig. 120). The
particle of radium bromide is at B, and some of
the charged helium atoms strike a surface C
covered with zinc sulphide, producing faint
flashes of Ught. The lens A magnifies the ^^ ^^q
flashes, which can be seen in a dark room after
the eye has become thoroughly rested (15-20 minutes). The
644 smith's intermediate chemistry
hdiiun gas given oflf by radium compounds was collected by
Soddy working with Ramsay and identified, and its rate of pro-
duction was measured. The amount was equal to 158 cubic
mm. per 1 g. of radium per year.
The beta-particles are electrons (p. 195), or unit charges of
negative electricity, and are shot out with a velocity approach-
ing that of light (300,000 kiloms. per sec). Their apparent mass
is very small (about t^tt ^^* of ^^ atom of hydrogen). Owing
to collisions with the relatively ponderous air-molecules, half of
them are lost after going about 4 cm.
The gamma-rays are identical with X-rays (see p. 548), and are
presumably produced like the latter by the impacts of the elec-
trons on the surrounding matter.
The helium atoms are almost all stopped by a sheet of paper or
by aluminium foil 0.1 mm. thick. The electrons have greater
penetrating power, many passing through gold-leaf, but being
practically all blocked by a sheet of aluminium 1 cm. thick.
The ganmia-rays (X-rays), however, are able to penetrate rela-
tively thick layers of metals and other materials of low atomic
weight.
One of the most striking facts is that the stoppage by the air
of so many rapidly moving particles results in the production of
much heat. One gram of radium would produce about 120 cal.
per hour.
Disintegration. — The emission of atoms of heUum and of
electrons was first explained by Rutherford (1902-3), then of
McGill University, Montreal, as being due to the spontaneous
disintegration of the atoms of uranium, radium, and other radio-
active elements. Thus, Rutherford was the first to show that
radium compounds produced a gaseous substance called the
radium emanation (niton), which was the residue left after the
emission of one atom of heUum from an atom of radium. This
gas was itself radioactive and underwent further disintegration,
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 545
depositing a solid radioactive residue on bodies in contact with it.
Furthermore, every known uranium ore contains radium (McCoy)
and radium emanation (Boltwood) in amounts proportional to
the uranium content. Also, after the radium has been removed,
the pure uranium compound gives olBf at first only a-particles,
but gradually recovers its whole radioactivity and is then found
to contain radium emanation once more (Soddy). It thus ap-
pears that uraniimi is the starting point, and that the disinte-
gration proceeds by steps, producing a number of different prod-
ucts. Each of these is formed from one such product and by
disintegration furnishes another.
Unlike ordinary chemical change, the rate of disintegration is
not affected by conditions. It can neither be started nor stopped
at will. It is no more vigorous at 2000° than at —200°. Other
changes occur between atoms, these within each atom.
The law, due also to Rutherford, describing the rate at which
any one radioactive element disintegrates is simple. Only a
certain fraction of the whole of any one specimen undergoes the
change in unit time. Thus, as the total amount diminishes be-
cause of the change, the amount changing during the next unit
of time, being a constant fraction of the whole, must be less.
Hence an infinite time would be required for the complete dis-
integration of any one specimen. For convenience in expressing
the rate of disintegration, however, we calculate and tabulate the
average life of the element.
Radium emits helium atoms at the rate of 3.4 X 10^® per gram
per second. From this fact, we can calculate its average life to
be about 2400 years. Hence, if it were not continuously being
produced (from uraniimi), the whole supply would have been
exhausted long before the earth reached a habitable condition.
The Uranium Group of Radioactive Elements. — The
following shows the various elements produced from uranium by
successive disintegrations. When a heliiun atom or an electron
546
SMITH'S INTERMEDIATE CHEMISTRY
is expelled, the fact is shown by the symbols He and €, respectively.
The first number below each element is the average life of that
member of the series (y = year, d = day, h = hour, m = min-
ute, s = second). The second niunber is the atomic weight,
obtained by subtracting from the at. wt. of uranium (238) the
weight (4) of each heUiim atom emitted.
Ui -.
He + U-Xi -^ € + U-X,
-> €+Vt
— ► He + Ionium
8XlO»y.
35.5 d. 1.65 m.
3 X 10»y.
2XlO*y.
238
234 234
234
230
He -Mia
— > He 4- Niton -n
> He 4- Ra-A
-» He + Ra-B
2440 y.
5.55 d.
4.3 m.
38.5 m.
226
222
218
214
€ + Ra-C
-► €+RarCi -H
► H +Ra-D
-»€ + Ra-E
28.1m
10-«s.
24 y.
7.2 d.
214
214
210
210
€ + Ra-F
-» He + Pb (end)
196 d.
210
206
The radium emanation was shown by Ramsay to be one of the
inert gases (p. 296), and was renamed niton. Its density was
determined experimentally with a small sample, using a micro-
balance capable of weighing 1/500,000 mgm., and found to be
222.4 (density of oxygen = 32).
The end-product of the disintegration is lead, and all uraniimi
ores contain lead (see p. 550).
Thorium, found as phosphate in monazite sand, is also radio-
active and furnishes a similar series of disintegration products.
The final material is again lead.
Actinium and polonium are other radioactive elements, which
have not yet been fully investigated.
Transmutation of Elements; Atomic Energy. — The
phenomena of radioactivity establish the transmutation of ele-
ments, long regarded as a delusion of the alchemists, as an in-
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 547
disputable fact. It is true that we have not yet discovered any
simple means of disintegrating the more common elements (see,
however, p. 17). We cannot even control in any way the rate of
disintegration of radioactive elements (see p. 545). If, however,
some method of inducing or hastening radioactive changes on
a large scale is devised in the future, a wonderful new source of
power will be put into our hands, namely, atomic energy.
The energy change in radioactive disintegrations is enormously
greater than in ordinary chemical reactions. One gram of radium,
as already mentioned, would evolve about 120 cal. per hour, and
would continue to evolve this heat, at a gradually decreasing rate,
for centuries. The total heat available would be over 2,000,000,000
cals. per gram, whereas a gram of carbon burning to CO2 gives
only 8040 cals. The disintegration of a pound of uranium salts
would furnish enough power to drive an ocean liner across the
Atlantic, but 8,000,000,000 years is entirely too long to wait
for the completion of the trip. Chemists are already looking
forward, however, to the possibUity of using the enormous stores
of energy here available so soon as a catalyst for the reaction is
obtained.
Another interesting by-product of this subject is the calcula-
tion that the heat given off by the disintegration of the radium
known to exist in the earth (niton is foimd in the soil and in well
waters) is sufl&cient aione amply to account for the maintenance of
its temperature. A globe the size of the earth, possessing orig-
inally only heat energy, and cooling from a white-hot condition
to the temperature of interstellar space, would have passed
through the stage of habitable temperatures in a much shorter
time than that which geological deposits and fossils show to have
been actually available. The discovery of the enormous, but
gradually released, disintegration energy of radium, enables us
now to explain the prolonged period during which life has existed
ill the earth.
This energy is derived from within the atom itself, by re-arrange-
648 smith's intermediate chemistry
ments of the protons and electrons of which it is constructed (see
p. 195). Building up a more complex atom from its disintegration
products would require just as much energy as is evolved in the
disintegration. This is another step which remains for the
future.
We may now proceed to examine more intimately the question
of atomic structurey already discussed in brief in earlier chapters
(pp. 195, 217).
Atomic Numbers. — Visible light, X-rays, and wireless elec-
tric waves are all vibrations of the same nature in the ether.
They differ only in wave-length, the order of the wave-lengths
being 10^ cm., 1(H cm., and 10* cm. (10 kilometers), respectively.
Now, just as the spectrum of visible light is obtained by using a
grating, on which the rulings are separated by distances of the
order of the wave-length of such Ught, so ordinary crystals give
spectra of X-rays, because they are composed of particles arranged
in rows about one thousand times closer and so form a suitable
grating for X-raj^. This fact
was first discovered by Dr.
Laue of the University of
Zurich (1912). The X-rays
are produced in an evacuated
tube by cathode rays, which
are streams of electrons emanating from the cathode (C, Fig.
121), when they strike the anticathode (A).
With different elements on the anti-cathode. X-rays of sUghtly
different wave-lengths, and therefore giving different X-ray
spectra, are produced. The greater the number of free protons
(unit positive charges) in the nucleus of the atom, the shorter
should be the wave-length of the characteristic X-rays. It was
shown by Moseley (a brilliant young English physicist, killed at
Gallipoli) that when the elements are arranged in the order of
these wave-lengths, whole numbers can be assigned to each which
Fig. 121
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 649
are inversely proportional to the wave-lengths of corresponding
lines in their X-ray spectra. These atomic numbers have been
determined for most of the elements, the atomic weights of which
lie between those of aluminium and uranium. In the following
table, the atomic numbers for these elements are given and, for
the sake of greater completeness, numbers for the twelve elements
preceding Al have been inserted also.
ATOMIC NUMBERS (Mosblby)
H 1
He 2
Ne 10
A 18
U
Na
K
Cu
Rb
A«
Cs
Au
3
11
19
29
37
47
55
79
87
Gl 4
Mc 12
Ca 20
Zn 30
Sr 38
Cd 48
Ba 56
Hg 80
Ra 88
B 5
Al 13
So 21
Ga 31
Y 39
In 49
La 57
Tl 81
Ac 89
C 6
Si 14
Ti 22
Ge 32
Zr 40
Sn 50
Ce 58
Pb 82
Th 90
N 7
P 15
V 23
Ab 33
Cb 41
Sb 51
Ta 73»
Bi 83
U-X*91
0 8
S 16
Cr 24
Se 34
Mo 42
Te 52
W 74
Po 84
U 92
F 9
CI 17
Mn 25
Br 35
- 43
I 53
- 75
- 85
Fe 26
Co 27
Ni 28
Kr 36
Ru 44
Rh45
Pd 46
Xe 54
Os 76
It 77
Pt 78
Nt 86
* The atomio nixmbers 59-72 are those of the metals of the rare earths: Pr 59, Nd 60, — 61,
Sa 62, Eu 63, Gd 64, Tb 66, Dy 66, Ho 67, Er 68, Tm 69, Yb 70, Lu 71, - 72.
It will be seen that there is a whole number available for every
known element, up to and including uranium, and not omitting
the rare earths which have no satisfactory place in the periodic
system. There are four blank numbers in the table, which cor-
respond to three spaces below Mn in the periodic system and one
xmder gold, and two more amongst the rare earths, indicating
only six elements with atomic weights less than that of uranium
yet to be discovered. The atomic numbers of argon and potas-
sium place them in the chemically correct order, while the atomic
weights do not. The same is true of cobalt and nickel and of
teUurium and iodine.
The atomic numbers represent the number of free positive
charges of electricity in the nucleus of the atom of each element.
It must be noted that the nucleus also contains, in all cases except
hydrogen, a number of bound positive charges associated with an
550 smith's intermediate chemistry
equal number of electrons, indicated by the difference between
the atomic weight and the atomic number.
The atomic numbers apparently determine all the properties
of each element, and are more fundamental than the atomic
weights. The latter are secondary properties, in most cases
modified by other factors, and in a few cases actually thrown out
of order by such factors.
Atomic Numbers of Radioactive Elements; Isotopes. —
When an atom of a radioactive element loses an atom of helimn,
it also loses two free positive charges from its nitdeus. Its atomic
number is consequently reduced by two (for example. Radium
= 88, Niton = 86). When, on the other hand, a radioactive
change takes place involving the loss of an electron^ a positive
charge in the nucleus, previously bound, becomes free, and the
atomic number is found to be increased by one.
With these facts in mind, an examination of the uranimn
disintegration series discloses that several elements (for example,
Radiiun-B, Radium -D, Radiiun-G and Lead) must ex-
ist which possess the same atomic numbers^ bui different atomic
weights. Such elements are known as isotopes. Isotopes are
identical in all of their chemical properties, although they differ
in atomic mass (see p. 20). This shows conclusively that atomic
weight is not a fundamental property, but atomic number.
Ordinary lead chloride contains the elements lead and chlorine
combined in the following proportions by weight:
Lead (207.2) — Chlorine (70.9) Lead chloride (278.1).
Richards, however, has found that the lead contained in uranium
ores gives a chloride in which as little as 206.1 parts by weight of
lead may be combined with 70.9 parts of chlorine, while Soddy
has shown that the lead extracted from thorium ores gives a
chloride which contains as much as 208.4 parts by weight of lead
to 70.9 parts of chlorine. We have, therefore, three lead chlorides,
all possessing the same specific properties, and being therefore the
,
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 551
same substance^ yet differing in composition. Other cases of a
similar nature undoubtedly exist, although not yet encountered
in actual practice.
Isotopes of Common Elem^ents. — It has recently been
shown by Aston that many conunon elements also are isotopic,
or contain chemically identical atoms of dififerent weight. The
method employed by Aston was that of positive ray analysis
(Fig. 122). The positive rays from a discharge tube are sorted
Photographic
Plate
Fig. 122
out into a thin ribbon by means of the two parallel slits Si and S2,
and are then passed between the oppositely charged plates Pi and
P2. The rays are deflected towards the negative plate P2, and
are spread out into an electric spectrum. A portion of this
spectrum deflected through a given angle is selected by the dia-
phragm D and passed between the circular poles of a powerful
electromagnet O, the field of which is such as to bend the rays
back again to fall on a photographic plate placed as shown. If
all the rays with a single charge have the same mass, they will
converge to a focus at F. If, however, the rays are derived from
an element which consists of a mixture of isotopes, each isotope
is distinguished by a separate band on the photographic plate,
and from the relative position of each band the mass of the atom
to which it corresponds can be obtained.
Chlorine (at. wt., 35.46), examined in this way, showed itself
to be a mixture of two isotopes with atomic weights exactly 35
and exactly 37. Bromine (at. wt., 79.92) gives isotopes with
552 SMITHES INTERMEDIATE CHEMISTRY
atomic weights exacUy 79 and exactly 81. Mercury (at. wt.,
200.6) appears to exist in as many as six forms, with atomic
weights ranging from 197 to 204. Other elements, however, such
as oxygen, nitrogen, and iodine, give no indications of isotopes.
The fundamental atomic weights obtained by Aston are with
one exception whole numbers, within the limits of experimental
error. The single exception is hydrogen (at. wt., 1.008).
Atomic Structure. — On the basis of the above resxilts, general
theories of atomic structure have been built up by Harkins and
Rutherford, postulating hydrogen and helium atoms as the " bricks "
building up the atoms of all elements. The elements with atomic
weights divisible by 4 are considered as constructed entirely of
charged helium nuclei, with surroimding electrons; thus C =
SHe""^ + 6e, O = 4He"'~^ + 8c, etc. Other elements must be
assiuned to contain hydrogen atoms also in their structure (see
p. 17). The decrease in the mass of the hydrogen atom from
1.008 in hydrogen itself to exactly 1 in all other atomic types
has been ascribed to a " packing effect."
Valence and Atomic Structure. — The electrons surrounding
the nucleus are arranged, according to a theory recently developed
by Lewis and Langmuir, in successive concentric shells. The
total number of electrons in these shells must be equal, since the
atom as a whole is electrically neutral, to the number of free "pro-
tons in the nucleus, in other words, to the atomic number.
The case of the hydrogen atom (atomic number = 1) has already
been considered (p. 195). The helium atom (atomic number =
2) has two electrons, which are supposed to be situated on opposite
sides of the nucleus — a very stable arrangement.
No more electrons can be contained in the first shell, hence in
succeeding elements the additional electrons begin to build up a
second outer shell. Only the electrons in this outer shell can be
added to or lost in interactions with other atoms (see p. 217),
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 553
and it is found that, for all of the next 8 elements from lithium to
neon, the tendency is either to lose electrons until none are left,
or to gain electrons until a stable ring of 8 is formed. The ar-
rangement of the electrons in this second shell has been pictured
by Lewis as shown in Fig. 123. The electrons are assmned to
occupy the comers of an imaginary cube, in the center of which is
the atomic nucleus.
7T A
7{ Av\
V)^
y
Li
Be
B
7\
7\
V h^j/' icnr l^Jr
€) QH 5)
N
0 F
Fig. 123
Ne
The valence of any one of these elements is therefore repre-
sented either by the nmnber of electrons that can be lost (positive
valence) or by the niunber required to form a stable ring of 8
(negative valence). The next eight elements (sodium to argon)
exhibit the same behavior. We have not the space here to dis-
cuss in detail the electron arrangement in these and later elements.
It will suffice to mention that not only are all of the relationships
suggested by the Periodic System (Chapter XXIII) confirmed, but
many of the points of difficulty in MendelejefiF's tabulation are
satisfactorily explained.
Co'Valence. — When combination takes place between two
atoms {e.g., Li and F) by loss and gain of electrons, we are left
564
smith's intermediate chemistry
with a compound J such as LiF (Fig. 124), in which the constitu*
ent atoms are apparently separate, being held together only by the
attraction of their opposite charges. An atom may also com-
plete its stable ring of eight, however, by sharing electrons with
another atom, as in the case of the fluorine molecule F3 (Fig. 125).
Li+ F-
FiG. 124
Similarly in carbon tetrachloride CCI4, we may assume that all
five atoms have completed their " octets," each chlorine atom
holding a pair of electrons in common with the central carbon
atom.
The number of pairs of electrons which an atom of an element
can thus share with other atoms is called its co-valence.
Polar and Nori'Polar Compounds. — Compounds like LiF
or NaCl, in which an electron has already passed from one atom
to another, are evidently potentially ionized, and if we can diminish
the attractive forces between the two charged atoms suflSciently
to enable them to break away from one another, we obtain im-
mediately the free ions (as in aqueous solution). Such com-
pounds are termed polar. In substances like CCI4, on the other
hand, where electrons are held in conunon, the molecule will not
tend to break up in this way. Such substances do not, conse-
quently, ionize in solution, and are termed non-polar.
Strictly speaking, however, the distinction between polar
and non-polar substances is not fundamental, but one of degree
only. In no non-polar compound, probably, are pairs of eleo-
RADIUM, ATOMIC ENERGY, AND ATOMIC STRUCTURE 655
trons held in common so impartially that they will not tend, to
some (albeit very small) extent, to pass over to one atom rather
than to the other. In the same way, in no polar compound has
the electron passed from one atom to another completely; the at-
tractive forces between th^ two atoms, tending to restore it to its
original position, must induce some distortion of the cubical ar-
rangement.
Atomic Structure and Chemical Affinity.— In cases where
combination between two elements produces a molecule (e.gr.,
LiF) in which the arrangement of electrons is much more stable
than in the original substances, we shall clearly obtain a con-
siderable diminution in the internal energy of the system (see p.
160) as the result of the interaction. Chemical activity or aflBnity
hence appears to depend, finally, upon atomic structure. Where
the electron arrangement is extremely stable, as with the inert
gases, the element will be inert. Where electrons are readily
gained or lost, as with the halogens or the alkali metals, the ele-
ment will be active, and will form very stable compoimds with
such elements as assist most readily in the interchange.
Chemistry " within the atom " is still in its infancy, but it
cannot be doubted that its development will lead to results of
the greatest importance in the near future. The facts presented
in this volume were almost all derived on a piu^ly experimental
basis, and the construction of hypotheses to correlate and explain
these facts has been a long and painful process. On the basis of
atomic structure, however, the next generation of chemists will
be able to predict physical and chemical properties in advance.
A multitude of new facts will thus be brought to light, and many
new applicat'ons of chemistry to industry will become evident.
Exercises. — 1. What justification has been obtained in this
chapter for the use of 0 = 16 rather than H = 1 as a basis for
atomic weights?
556 smith's intermediate chemistry
2. If there are two isotopes of chlorine, and six of mercury,
how many varieties of mercuric chloride HgCU are possible?
In what respects would these varieties be different? In what
respects woxild they be identical?
3. Draw diagrams showing the electron arrangements in the
following molecules: hydrogen fluoride, water, oxygen (O2),
carbon dioxide, sodium chloride.
APPENDIX
I. The Metric System
Length. 1 meter = 10 decimeters = 100 centimeters (100 cm.)
= 1000 millimeters (1000 mm.).
1 decimeter = 10 centimeters = nearly 4 inches.
Volume. 1 Uter = 1000 cubic centimeters (1000 c.c.) =» a cube
10 cm. X 10 cm. X 10 cm.
1 liter = 1.057 quarts (U. S.) or 1.136 quarts (Brit.).
1 fl. ounce (U. S.) = 29.57 c.c. 1 ounce (Brit.) = 28.4 c.c.
Weight. 1 gram (g.) = wt. of 1 c.c. of water at 4° C. 1 kilo-
gram = 1000 g.
1 kilogram = 2.2 lbs. avoird. (U. S. and Brit.).
1 ounce avoird. (U. S. and Brit.) = 28.35 g.
1 nickel (U. S.) weighs 5 g. 1 halfpenny (Brit.) weighs 5 to
5.7 g.
n. Scale of Hardness
Each of the following minerals will scratch the surface of a
specimen of any one preceding it in the list.
1. Talc 6. Felspar
2. Gypsum (or NaCl) 7. Quartz
3. Calcite (or Cu) 8. Topaz
4. Fluorite 9. Corundum
5. Apatite 10. Diamond
Glass is sUghtly scratched by 5, and easily by those following.
Glass will not scratch 5 distinctly, but will scratch those pre-
ceding 5.
A good knife scratches 6 slightly, but not those following.
A file will scratch 7, but not those following.
667
568
SMITH'S INTERMEDIATE CHEMISTRY
in. Temperatures Centigrade and Fahrenheit
Upon the centigrade scale, the freezing-point of water is 0** C.
and the boiling-point 100*^ C. Upon the Fahrenheit scale, the
same points are 32® F. and 212*^ F., respectively. The same in-
terval is thus 100® on the one scale and 180° on the other. The
degree Fahrenheit is therefore — or - of 1° Centigrade. Any
180 9
temperatures can be converted by using the formulsB:
C.® = I (F.® - 32), F.° = I (C.°) + 32.
The following table (IV) contains the temperatures from 0® C.
to 35° C, with the corresponding values on the Fahrenheit scale
(32° F. to 95° F.).
IV. Vapor Pressures of Water
Both t^ie Fahrenheit (F) or ordinary and the Centigrade (C) temperatures are given.
Temperature
•
Temperature
^ramiiTA TnTn
Pressure, mm.
. F.
C.
X AvDOUAvf m 1 1 M «
F.
C.
32**
0**
4.6
71. 6**
22 *»
19.7
41
5
6.5
73.4
23
20.9
46.4
8
8.0
75.2
24
22.2
48.2
9
8.6
77.0
25
23.6
50.0
10
9.2
78.8
26
25.1
61.8
11
9.8
80.6
27
26.5
63.6
12
10.5
82.4
28
28.1
66.4
13
11.2
84.2
29
29.8
67.2
14
11.9
86.0
30
31.5
59.0
15
12.7
87.8
31
33.4
60.8
16
13.5
89.6
32
35.4
62.6
17
14.4
91.4
33
37.4
64.4
18
15.4
93.2
34
39.6
66.2
19
16.3
95.0
35
41.8
68.0
20
17.4
60.8
21
18.5
212.0
100
760.0
INDEX
(''
*** Adds are all listed under " acid," and salts under the positive radical.
ACETYLENTC, 333, 351
Acid, acetic, 419
boric, 363
carbonic, 336
hydrochloric, 125-31
composition of, 146-7
hydrosulphuric, 255
hypochlorous, 223-6
nitric, 308-11, 313-4
nitrosylsulphiiric, 265
nitrous, 314
phosphoric, 322
picric, 484
siUcic, 360
sulphuric, 261-71
chamber process, 264
concentration, 269
contact process, 261
fimiing, 271
sulphurous, 259
Acids, properties of, 131, 170, 191
organic, 348
Activity, 27, 159-61
order of metals, 54
order of non-metals, 209
Adsorption, 421
AflSnity, 159-60
Air, 26, 288-96
a mixture, 293
carbon dioxide in, 291-2
dust in, 292-3
liquid, 29, 294-6
water vapor in, 288-90
Alcohol, industrial, 418
Alcohols, 348
Aldehydes, 348
AlkaUes, 168
Allotropic modifications, 220, 249
Alloys, anti-friction, 326
fusible, 326
steel, 493
Alimiinates, 468
Aluminium, 466-8
salts of, 468-70
Alimis, 469
Amalgams, 516-7
Ammonia, 299-^06
-oxidation process, 313
Anmioniimi salts, 304-5
Anhydrides, 257
Animal life and nutrition, 428-9
Anthracene, 353
Antimony, 325-6
Appendix, 557-8
Aqueous tension, correction for, 47
Argon, 296
Arsenic, 323-5
Atmosphere, 288-96
Atomic disintegration, 544
energy, 546-8
numbers, 548-50
structure, 552-3
weights, 76, 77, 85, 281
Atoms, 85
Attributes, 7
Babbitt's metal, 326
Bakelite, 481
Baking powders, 430
soda, 366
559
660
INDEX
Barium, 393
Barometer, 44, 61
Bases, properties of, 167, 192
Battery, storage, 505-7
Benzene, 352
Bessemer process, 491
Birkeland-Eyde process, 312
Bismuth, 326-7
Blast furnace, 486-7
Bleaching, 221, 225, 259
powder, 224-6
Blue-prints, 498
BoiUng-point, of liquids, 61
of solutions, 118, 177, 190
Borax, 363-4
Boron, 363
Brass, 512
Bromine, 199-201
Bronze, 512
Bunsen burner, 356
Cadmium, 451
Calcium, 383
bisulphite, 398
carbide, 333
c^bonate, 383-4
chloride, 392
cyanamide, 392
fluoride, 207
hydroxide, 386
oxide, 384
phosphates, 411
salts as fertilizers, 413-5
sulphate, 387
Calculations, 133, 148
Calorie, 61, 162
large, 435
Candle flame, 355
Cane-sugar, 401-3
Carbides, 332-3
Carbohydrates, 401, 429, 435
Carbon, 328-32
allotropic forms, 328-30
dioxide, 333-7
Carbon dioxide, as plant food, 396
in air, 291-2
disulphide, 255
monoxide, 337-40
tetrachloride, 332
Carborundum, 332
Casern, 438
Catalysis, 32-3
Celluloid, 480
Cellulose, 397-9, 480, 482
Cement, 471
Chamber process, 264
Charcoal, 330, 420-2, 438
Chemical change, and energy, 155
efifect of temperature, 27, 241
physical accompaniments, 153
varieties of, 132
Chemical equilibrium, 230-46
properties, 33
units of weight, 74, 78
Chlorine, 139-45
molecular formula of, 148
-water, 224
Chromium, 529-30
salts, 530-4
Clay, 471
Coal, 422-5
gas, 423
Cobalt, 500
Coke ovens, 424
Collodion, 385
Colloidal suspensions, 109, 440-2
Combination, 14, 132
Combining proportions, 36
weights, 76-7
Combustion, preferential, 339
spontaneous, 41
Complex ions, 513
Components, 6
Composition, 38, 134
Compoimds, 16
Conditions, 7
Conductivity, electrical, 176, 187
Constituents, 8
INDEX
561
Contact action, 33
process, 261
Copper, 510-12
plating, 514
refining, 515
salts, 512-4
Cotton, 1, 398
Cracking of oils, 346, 352, 35^
Crystallization, 114
Crjrstal structure, 94-6
Deacon's process, 140, 230
Decantation, 13
Decomposition, 16, 132
Deliquescence, 118
Density, 6, 85, 148
Dextrose, 401
Dialysis, 442
Diamond, 95, 329
Diet, normal, 436
Diffusion, 90
Digestion of foods, 430, 434
Displacement, 51, 132, 172
valence by, 213
Dissociation, 104, 133
Distillation, 66
fractional, 344
Double decomposition, 132, 173, 185
Drugs, 479
Dust in air, 292-3
Dyeing, 474-6
Dyestuffs, 477-8
D3niamite, 483
Edison cell, 507
Efflorescence, 68
Electrons, 195
and oxidation, 228
and valence, 217
Elements, 17, Inside hack cover
common, 19
periodic system, 273-84
transmutation of, 546
Electric furnace, 332
%
Electrolysis, 55, 139, 183
Electrolytes, conductivity of, 176
Electrotypes, 515
Emulsions, 110
Endothermal changes, 161
Energy, 156, 158, 426
and chemical activity, 159
atomic, 546-8
conservation of, 158
internal, 158
Enz3rmes, 417
Equations, 81, 135
balancing of, 101, 316-7
ionic, 182
making of, 100
thermochemical, 162
Equilibrium, 63, 69, 230-46
displacement of, 239
Equivalent weights, 53, 122
Esters, 349, 432
Ethers, 348
Ethylene, 360
Evaporation, 14
Exothermal changes, 161
Explosives, 481-4
Extraction, 112
Fats and oils, 432
hydrolysis of, 433
Fehling's solution, 401, 513
Felspar, 5, 410
Fermentation, 417
Fertilizers, 409-15
Filtration, 13, 66
Fixation of nitrogen, 300, 312
Flame, 354-7
Flotation process, 511
Hour, 5
Fluorine, 206-7
molecular structure, 664
Foods, 429
fuel value of, 435-^
Formaldehyde, 348
Formuke, 78
562
INDEX
Formulae, in calculations, 133
making of, 99
molecular, 104-6
Freezing mixtures, 119
point of liquids, 60
of solutions, 119, 177, 190
Fructose, 401
Fuels, 435
Gaillard tower, 269
Gases, facts about, 93
inert, 29&-7
liquefaction of, 91, 294-6
measurement of, 44
perfect, 92
properties of, 88-93
Gasoline, 213
Gay-Lussac tower, 267
Gelatine, 435
Glass, 361-3
Glover tower, 266
Glucose, 401
Gluten, 5
Gold, 623-4
Gram-molecular volume, 83
Granite, 4
Graphite, 329
Guncotton, 482
Gunpowder, 372
Gypsum, 387, 414
Haber process, 300-2
Halogen derivatives of hydrocarbons,
347
family, 199, 208-9
Hard water, 388-91
Hardness, scale of, 557
HeUum, 297, 543
Homogeneous systems, 238
Humidity, 289
Hydrates, 67
Hydrocarbons, 343-53
saturated, 343-7
Hydrocarbons, imsaturated, 350
Hydrogen, 50-6
from water gas, 338
molecular formula, 105, 148
position in periodic system, 283
structure of atom, 195
Hydrogen bromide, 201-2
chloride, 129-31
fluoride, 207-8
iodide, 204-5
peroxide, 221-3
sulphide, 252-5
Hydrolyte, 56
Hypothesis, Avogadro's, 82
ionic, 181
molecular, 48, 88-93
Ice, 64
Impurities, 6
Inert gases, 296-7
Ink, 499
Iodine, 203-4
Iodoform, 205
Ion-product constancy, 456
Ionic equations, 182
equilibria, 451-64
Ionization, 181-98
and chemical activity, 189
Ions, 182
and conductivity, 187
and displacement, 184
and double decomposition, 186
and electrolysis, 183
and electrons, 195
and valence, 212
complex, 513
Iron, 486-94
Bessemer process, 491
cast, 488
galvanized, 449
open-hearth process, 492
salts, 495-8
Isotopes, 550-2
INDEX
563
Ketones, 349
Kindling conditions, 39
Krypton, 297
Lactose, 401
Lampblack, 355
Laundry, hard water in, 390-1
Law, Avogadro's, 82
Boyle's, 45, 89
Charles^ 46, 90
Dalton's, 47, 112
definition of, 8, 23
Dulong and Petit's, 86
Gay-Lussac's, 60, 72
Henry's, 112
Le Chatelier's, 244
Newland's, 275
of chemical change, 7
combining weights, 77
component substances, 6
definite proportions, 19, 86
effect of heating, 27, 241
molecular concentration, 235
partition, 112
solubility product, 456
periodic, 280
van't Hoff's, 242
Lead, 502-3
accumulator, 505-7
salts, 503-5
white, 504
Le Blanc process, 367
Lewis-Langmuir theory, 552
Lime, 384-6
Liquids, properties of, 93
Lithopone, 451
Magnesium, 446
salts of, 447-8
Maltose, 401, 418
Manganese, 527
dioxide, 31, 142, 529
salts of, 528-9
Manure, 413
Mass, conservation of, 20
Matches, 321
Matter, states of, 64
structure and behavior of, 88-97
Mendelejeff's system of elements, 276
Mercury, 516-7
salts, (Af 517
Metals, order of activity, 64
Methane, 346
Metric system, 557
Mica, 4, 466
Mixtures, 6, 12
Molar solution, 122
Molecular concentration, law of, 235
formuke, 104-6, 148
hypothesis, 48, 88-93
weights, 75, 83, 84
Monel metal, 500
Mortar, 386
Mustard gas, 484
Naphthalene, 353
Natural gas, 345
Neon, 297
Neutralization, 192
Nickel, 500
plating, 516
salts of, 501
Nitrogen, 286-7
compoimds as fertilizers, 410
peroxide, 309, 311
Nitroglycerine, 481
Nitrous anhydride, 265, 314
oxide, 315
Non-metals, order of activity, 209
Normal solution, 121
Octaves, law of, 275
Oil refining, 344
shale, 345
Olefines, 351
Oleum, 261, 271
Osmosis, 406
Osmotic pressure, 407-0
564
INDEX
Oxidation, 40, 226-8
Oxides, acidic and basic, 257
Oxone, 31
Oxygen, 27-42
member oi sulphur family, 273
molecular formula, 105
Ozone, 21^21
Painto, 305-4
Paper manufacture, 398
Parafl^ series, 343-7
Parke's process, 519
Pasteur filter, 66
Pauling process, 313
Perfumes, 478-0
Periodic system, 273-80
app^cations of, 281-2
defects of, 282-4
Permutite process, 390
Petroleum, 344
Phenols, 353
Phosphine, 323
Phosphorus, 319-23
compounds as fertilizers, 411
pentachloride, 236
pentoxide, 322
Photography, 521-3
Physical properties, 33
Physics in chemistry, 22
Plant hfe, 395-7, 405-9
Plaster of Paris, 387
Plastics, 480
Platinimi, 524-5
group of metals, 281, 525-6
Polar and non-polar compounds, 554
Portland cement, 471
Potassium, 370
bicarbonate, 371
carbonate, 371
chlorate, 30
chloride, 413
chromate, 531
dichromate, 531
ferricyanide, 498
Potassium ferrocyanide, 497
hydroxide, 371
nitrate, 30, 372
permanganate, 141, 529
salts as fertilizers, 412
Precipitation, 129
theory of, 458-9
Preferential combustion, 339
Pressinre, atmospheric, 44
osmotic, 407-9
Producer gas, 337
Promoters, 339
Proteins, 434
Protons, 195
PjT^x, 363
Quartz, 4
Radicals, 170
positive and negative, 171
valence of, 212
Radioactivity, 542-8
Radium, 542-3
Recognition of substances, metallic.
536
non-metaUic, 374
Reduction, 57, 228
Refrigeration, 306
Respiration, 291-2
Reverberatory furnace, 489
Reversible reactions, 69, 127, 230
Rouge, 496
Rusting, 9, 26, 494
Salts, 170, 192
Saponification, 438
Saturated solutions, 120, 243, 455
Saturation, 110, 120
Schonherr process, 313
Scientific method, 24
Selenitun, 273
Shale, 345
Silica gel, 360
SiHcon, 359
INDEX
565
Silicon dioxide, 359
tetrachloride, 360
tetrafluoride, 207, 360
Silk, 438
artificial, 399, 481
SUver, 519-21
salts, 521
Soap, 438-44
Sodium, 164r-5
bicarbonate, 366
carbonate, 367
chloride, 95, 125, 365
cyanide, 393
hydroxide, 166-8
nitrate, 369
peroxide, 31, 221
silicate, 360
Solder, 503
Solids, properties of, 94-7
Solubilities, table of. Inside front
cover
Solubility and temperature, 113, 243
of a gas, 111
-product law, 456
Solution, 108
heat of, 115
of insoluble precipitates, 460-2
saturated, 110, 120
supersaturated, 115, 121
volume change in, 116
Solutions, boiling-points of, 119, 177
freezing-points of, 119, 177, 190
molar, 122
normal, 121
osmotic pressure of, 407-9
vapor pressure of, 117, 177, 190
Solvay process, 366
Solvents, 110
influence of solute on, 116
part in ionization, 193
Specific properties, 3, 33
Speed of chemical reactions, 161,
238
effect of temperature, 27, 241
Spintharoscope, 543
StabiUty, 27
Starch, 5, 399-401
digestion of, 430
fermentation of, 418
iodine test for, 400
Steam, 61
Steel, 489-93
Stellite, 500
Strontium, 393
Sublimation, 204
Substance, 3
Substances, amorphous, 95
compoimd and simple, 16
Sucrose, 401-3
Sugar refining, 402
Sugars, fermentation of, 417
list of conmion, 401
Sulphur, 248-52
dioxide, 258-«
trioxide, 261
Supersaturation, 115, 121
Suspensions, 109
coUoidal, 400, 440-2
Symbols, 78
Synthesis, 98, 300
Tellurium, 273
Temperature, absolute, 46
critical, 92
effect on solubility, 113
Temperatures, Centigrade and Fahr
renheit, 558
Test, 36
Thermite, 468
Thermochemistry, 161
Tin, 507-8
salts of, 508
Toluene, 353
Toxic gases, 421, 484-5
smokes, 422
Transition point, 250
Trinitrotoluene, 483
Twitchell process, 433
566
INDEX
Ultramarine, 471
Ultrarmicroscope, 440
Uranium, 634, 645-6
Valence, 211-7, 228
and atomic structiire, 662
Vapor pressure, of solutions, 117, 177,
190
of water, 47, 63, 668
Vaporization, 61
Ventilation, 291
Vinegar, 419
Viscose, 399
Vitamins, 436-7
Volume, standard 73, 79, 83
Washing soda, 367, 391
Water, 68-70
as a solvent, 65-122
association of, 122
coagulation process, 470
composition by volume, 59, 106
by weight, 68
Water, hardness, 388-0
ionization of, 190
purification, 66
softening, 389-91
vapor in air, 288-90
pressure of, 47, 63, 658
Water-gas, 337-8
Waters, earburetted, 356
natural, 66
Weights, atomic, 76-7, 85, 281
combining, 76-7
equivalent, 53, 122
formula, 134
molecular, 75, 83-4
Wood, distillation of, 420
Wood's metal, 326
Wool, 1, 438
Xenon, 297
Xylene, 353
Zinc, 448-6
salts of, 460-1
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1923
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Ll.
Smith, Alexander. 39950
Smith's intermediatelf
chemistry, rev. by James
Kendall and E.E.Slosson.
S9950
V
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INTERNATIONAL ATOMIC WEIGHTS (1922)
O » 16.
Aluminium Al 27.0
Antimony Sb 120.2
Argon A 39.9
Arsenic As 74.96
Barium Ba 137.37
Bismuth Bi 209.0
Boron... B 10.9
Bromine Br 79.92
Cadmium Cd 112.40
Caesium Cs 132.81
Calcium Ca 40.07
Carbon C 12.005
Cerium Ce 140.25
Chlorine CI 35.46
Chromium Cr 52.0
Cob&H- Co 58.97
ColumSium Cb 93.1
CJopper Cu 63.57
Dysprosium Dy 162.5
Erbium Er 167.7
Europium Eu 152.0
Fluorine F 19.0
Gadolinium Gd 157 .3
Gallium Ga 70.1
Germanium Ge 72.5
Glucinum Gl 9.1
Gold Au 197.2
Helium ,.He 4.00
Hohnium Ho 163.5
Hydrogen H 1 .008
Indium In 114.8
Iodine I 126.92
Iridium Ir 193.1
Iron Fe 55.84
Krypton Kr 82.92
Lanthanum La 139 .0
Lead Pb 207.20
Lithium Li 6.94
Lutecium Lu 175.0
Magnesium Mg 24.32
Manganese Mn 54.93
Mercury Hg 200.6
O » 16.
Molybdenum Mo 96.0
Neodymium Nd 144.3
Neon .Ne 20.2
Nickel Ni 58.
Niton (radium emanation) Ni 222 . 4
Nitrogen . . .N 14.00fe
Osmium Os 190.9
Oxygen O 16.00
Palladium Pd 106.7
Phosphorus P 31.04
Platinum Pt 195.2
Potassium K 59. IC
Praseod3nnium Pr 140.9
Radium Ra 226.0
Rhodium Rh 102.9
Rubidium Rb 85.46
Ruthenium Ru 101 .7
Samarium Sa 150.4
Scandium Sc 45.1
Selenium. Se 79.2
Silicon Si 28.1
Silver Ag 107.88
Sodium Na 23 .00
Strontium Sr 87.63
Sulphur S 32.06
Tantalum Ta 181.5
Tellurium Te 127.5
Terbium Tb 159 .2
ThaUium Tl 204.0
Thorium Th 232.16
Thulium Tm 169.9
Tin... Sn 118.7
Titanium Ti 48.1
Tungsten W 184.0
Uranium U 238.2
Vanadium V 51 .0
Xenon Xe 130.2
Ytterbium (Neoytterbium) Yb 173.5
Yttrium Yt 89.32
Zinc Zn 65.37
Zirconium Zr 90.6
For a table of the Periodic System, see p. 278.
For a table of Atomic Numbers, see p. 549.
-/ »•%» *■
'Si-
% > '. V H ••