OF
CHEMISTRY
SHEPARD
ELEMENTS
INORGANIC CHEMISTRY,
DESCRIPTIVE AND QUALITATIVE.
BY
JAMES H. SHEPARD,
INSTBUCTOR IN CHEMISTRY, YPSILANTI HIGH SCHOOL.
BOSTON:
PUBLISHED BY D. C. HEATH & CO.
1886.
Entered according to Act of Congress, in the year 1885, by
JAMES II. SHEPARD,
in the Office of the Librarian of Congress, at Washington.
J. 8. GUSHING & Co., PRINTERS, 115 HIGH STREET, BOSTON.
PREFACE,
THIS elementary treatise is based upon plans and
methods which have been employed in the author's
laboratory throughout a series of years, and no work has
been incorporated in the text or in the exercises that
has not there been proven practicable.
A love for the science of chemistry would have for-
bidden any attempt to add another text-book to the
already too extended list of Elementary Chemistries had
not the hearty commendations of teachers of national
reputation and undoubted ability encouraged both the
author and the publisher to put this work in permanent
form.
During the correspondence which grew out of the
issue of this work, it has become evident that many
of the best teachers in all sections of the country are
pursuing independently a plan essentially the same ; and
the deepest regret which the author feels in seeing the
work go to press arises from the fact that it signals
for the close of his correspondence and labors with such
an enthusiastic corps of fellow-workers. If it shall be
IV PREFACE.
found that this work, towards which they have contrib-
uted so freely, meets with their hearty commendations,
he will rest satisfied with his labor of love.
It now only remains to return thanks to those who, so
patiently and ofttimes so laboriously, have assisted the
author in completing this work.
Dr. Ira Remsen, Professor of Chemistry in Johns
Hopkins University, has critically read the work in man-
uscript and in proof, and has contributed much toward
the accuracy and the arrangement of the topics treated,
particularly those which pertain to chemical theories.
The following well-known and enthusiastic teachers of
chemistry have read the work in proof, and have given the
author constant advice as, from time to time, the sheets
appeared : —
Otis Coe Johnson, Assistant Professor of Applied Chem-
istry, University of Mich. ; Robt. B. Warder, Professor of
Chemistry, Purdue University, and State Chemist of Ind. ;
W. W. Daniells, Professor of Chemistry, University of
Wis. ; Jas. A. Dodge, Professor of Chemistry, University
of Minn.; E. J. Bartlett, Professor of Chemistry, Dart-
mouth College ; Delos Fall, Professor of Natural Science,
Albion College ; Albert C. Hale, Instructor in Chemis-
try, Central Grammar School, Brooklyn, N.Y. ; George
Weitbrecht, Chemist and Instructor Natural Science,
High School, St. Paul, Minn. ; Leroy Griffin, Professor
of Natural Sciences, Lake Forest University, 111. ; Herbert
C. Foote, Chemist and Instructor Natural Science, High
PREFACE. V
School, Cleveland, O. ; Louis McLouth, Professor of
Natural Science, Michigan State Normal School, Ypsi-
lanti (now of Michigan Agricultural College, Lansing) ;
H. N. Chute, Instructor Natural Science, High School,
Ann Arbor, Mich. ; W. G. Rappleye, Teacher of Physics
and Chemistry, Normal School, Oswego, N.Y. ; Adolf T.
Bechdolt, Supt. Schools, Mankato, Minn.; J. C. Crawford,
Supt. Schools, Green Bay, Wis. ; and many other teach-
ers of Chemistry in preparatory, normal, and collegiate
departments.
Article 234 on the Natural Classification of the Elements
is due to the kindness of Professor Warder.
The author is aware that many data, not usually given
in works for beginners, appear in the text; but, in the
laboratory, these will be found to be useful and valuable
additions.
It has been the constant aim, in preparing this book, to
make the labors of the teacher as light as possible, and to
place the laboratory work where it would do the most
good, in the hands of the students.
J. H. S.
OCT. 15, 1885.
TO THE TEACHER.
I.
METHODS.
IT is with no little diffidence that the author approaches the
subject of Methods. He is fully aware that every teacher has
his own method, and that all successful methods are entitled to
respectful consideration. There are, however, some principles
upon which all are agreed, and a classification and a brief dis-
cussion of the different methods which have been employed may
at least prove suggestive.
The problems before us are these : —
1. If we teach chemistry at all, what advantages has this
science to offer as factors in developing the youthful mind, and
what good results will follow its study ?
2. If the study of chemistry be positively desirable, what
method of presentation will best accomplish the desired results ?
1. Neglecting for the present the claims of those who would
become chemists by profession, let us consider chemistry as
a means of education. In this capacity, when properly taught,
chemistry awakens and cultivates a spirit of investigation ; it
encourages the student to ask Nature questions, and it is unex-
celled by any other branch of learning in the clearness and con-
clusiveness of the answers received ; it insists upon the strictest
habits of observation ; it leads to the concentration of thought
and of energy ; it educates the senses ; it trains the hand to
delicate manipulation ; it exercises the faculty of reason and
the power of judging ; it affords useful information peculiarly
Viii TO THE TEACHEE.
its own, and thus forms an important part of a good, general
education.
Backed by such advantages as these, it really seems that
chemistry should be deemed worthy of a place in all liberal as
well as in purely scientific education.
2. When it comes to methods of instruction, the teacher has
many from which to choose. These methods may be arranged
approximately under four general divisions : —
(1) The Classical or Didactic method.
(2) The Laboratory method, in which the teacher does all the
experimentation in the presence of the class, and accompanies
the experimentation by didactic instruction.
(3) The Working-Laboratory method, in which the student
does his own experimentation, and receives little or no
didactic instruction. This method varies somewhat, in its
application : —
(a) The student may be required to work with no aid from
text-books, etc., relying apon his work alone for the benefits
to be obtained, the instructor in this case acting really as a
demonstrator.
(b) The student may have a text-book as a guide, the in-
structor acting as before.
(4) A method which the author begs leave to christen THE
SCIENTIFIC METHOD ; this embodies all the good features of the
preceding methods.
The experience of many careful instructors would warrant the
following estimate of the relative value of these methods : —
The first method affords some special information ; otherwise,
chemistry, when thus taught, is equalled as an educational factor,
by history, and by kindred subjects ; and is excelled by mathe-
matics and the classics.
The second method accomplishes as much as the first, and to
a very limited extent cultivates observation ; farther than this,
no advantages are to be gained by its use.
The third method incites to investigation ; trains the senses to
TO THE TEACHER. IX
observe ; trains the hand to careful manipulation ; and encour-
ages the student to originate. But (a) is too slow ; it requires
more time than can be devoted to this study ; and although the
student may " know well what little he does know,'' his rea-
soning powers are not developed, and his fund of information is
not sufficiently increased. (6) accomplishes its ends somewhat
more rapidly than (a) , and consequently yields more informa-
tion in a given length of time ; otherwise, it is not better than
(a). As a rule, students taught by the third method are very
weak in chemical theory.
An insight into the fourth method may best be obtained by
a description of the manner of its application. This method
contemplates : didactic instruction by the teacher ; a good text-
book, and as many books of reference as possible ; much work
by the student, who should keep a careful record of all work
done, and who should recite frequently ; and work by the
teacher, either in the presence of the class where the class is
large, or personal directions to the student when the class is
small.
The use of this method is extremely simple. The teacher
assigns a lesson from the text, indicating such parallel reading
as the time at the student's disposal may permit ; he then goes
over the lesson, and gives such working directions and cautions
as the subject and the student's capabilities may demand, thus,
in most cases at least, saving the student from wasting his time
in repeating the useless blunders of those who worked centuries
ago ; if the experiments be dangerous, or if the line of work be
new, the teacher either makes the experiment for the class with
little or no explanation, or he explains the general principles,
leaving the student, when safety permits, to work out the
details ; after this, the student is sent to his desk, where he
works, reads, and makes his notes for the next recitation.
The following day the student is questioned concerning his
work, and is encouraged to tell truthfully and exactly how he
succeeded, if he has succeeded, or why he failed, if he has failed.
X T° THE TEACHEB.
In case the student has failed, and does not know the reason,
or gives the wrong reason, the teacher, meanwhile explaining
nothing himself, calls upon other members of the class until the
point in question is elucidated. If all have failed, which rarely
happens, the teacher gives directions anew, and the students try
again. In general, the teacher aims to do as little work for
the class as possible, and to tell the student nothing that he can
find out for himself in a reasonable length of time.
Reviews and those topics which are necessary to the science
as a whole, and which are not covered by the student's work,
are faithfully taught by didactic methods.
Variety is introduced and practical results are obtained in
several ways, 'and thus the student's interest is never permitted
to flag. Students are assigned essays upon various topics ;
are given unknown substances to analyze ; are required to make
analyses of -substances with which they are familiar, such as
coins, worn-out articles of jewelry, alloys, common salt, baking-
powder, samples of drinking-water, crude drugs from the drug
store, etc., etc.
In keeping his notes, the student constantly recognizes the
fact that the knowledge he is seeking is to be drawn from
phenomena observed while working with known factors. A
good form for the headings of a note-book is as follows : —
1. Required Conditions.
2. Known Conditions.
3. Operations.
4. Conclusions.
Under 1 the student enumerates what he wishes to know ;
under 2 he enumerates his working materials ; under 3 he tells
what he does ; and under 4 what conclusions he has reached.
In making his notes, the student is warned that he may err :
(1) by taking a trivial required condition; (2) by assuming a
required condition that will not follow from the premises ;
(3) by an indefinite or obscure description of his operations ;
(4) by reaching a conclusion more general than the premises
TO THE TEACHER. XI
warrant ; and (5) by employing bad English in any of the pre-
ceding divisions.
In this book, written with special reference to the fourth
method, the student's work, so far as practicable, is not fore-
stalled by telling him what phenomena are to occur, and many
queries are left to be answered by an experiment which the
student may devise. In the closing portions of the "book, all
experiments (as such) are purposely omitted with the sugges-
tion that, as the student is no longer, in the strict sense of the
word, a beginner, he should be thrown still farther upon his
own resources. He is asked to prepare various salts and com-
pounds of the metals, and to describe their preparation as
experiments ; this is work well adapted to afford an exercise
more exacting than anything previously attempted. Another
good exercise for the student is to prepare working solutions for
himself and his classmates, starting with the crude materials.
By this method, the student will not only secure a lasting
benefit from chemistry, as an educator of hand and rnind, but
in case he so desires, he will find himself amply prepared for
further pursuing this delightful study.
II.
WHAT SHOULD THE STUDENT MEMORIZE?
As in all other studies, this question is frequently asked con-
cerning chemistry. In the curriculum of all schools in which
chemistry is taught to beginners other studies are found, or
should be found, which are peculiarly adapted to cultivate the
faculty of memory ; the amount of memorizing required in
chemistry should be quite limited, depending more or less upon
the curriculum itself.
In general, it is safe to say that much valuable time has been
frittered away by requiring the student to memorize unimportant
details which not even an expert retains. Because certain facts
xii TO THE TEACHER.
or numerical data are given in a text, it does not follow that the
student's memory must be burdened with them ; there are other
uses for such data, and especially so in a working text. Thus,
for example, the weight of one litre of a gas, atomic heats,
specific gravities, densities, etc., etc., may be utilized in
solving problems.
It is riot even necessary to memorize the atomic weights or
such units as the weight of one litre of hydrogen, since the stu-
dent will learn these data by frequently using them, just as we
all have learned the multiplication table. It is, however, a
positive advantage to have these data given in the body of the
text, since a frequent reference to them serves in a certain way
as a review.
Again, the author has never required his classes to memorize
tests and separations, and still his students, by way of final
work, have been able correctly to analyze complex unknown solu-
tions without the aid of reference books or of text-books ; this
was accomplished by simply giving the student much work to
do, and then by asking him to explain his work. And again, it
would be manifestly absurd to require the student to memorize
the language of the text in experiments. And finally it should
suffice to bear in mind that to be able to do, to reason, to origi-
nate, is far better than to be able to repeat from memory things
not half understood.
III.
A BRIEFER COURSE.
FOR various reasons some teachers may wish to use certain
portions of the text and to omit the rest. There is no reason
why this may not be done. Experience has shown that, in
a working text, even of the most elementary character, it is
desirable to have the book quite complete, thus lightening the
labor of the teacher, and providing for emergencies which often
TO THE TEACHER. xiii
and unexpectedly arise. For example, one piece of apparatus
may be broken, or it may be wanting, while another, which
may be made to answer the same purpose, is available ; or, a
student in his work may come upon something which not even
the teacher could foresee ; one chemical may have been entirely
consumed, while another, which will answer, may still be plenti-
ful, etc., etc. In view of all these considerations, it is evident
that a somewhat full text will be more satisfactory to both
student and teacher, even though certain portions of it are
omitted, or dwelt upon quite lightly. There is no truth in the
tradition that " to omit certain parts of a book causes the
student to be less thorough " ; on the contrary, such a process
should teach him to select what he really wants from what he
does not want, — a lesson he must learn sooner or later. There
is one thing, at least, that a full text certainly does do, and that
is, it forever banishes from the student's mind the idea that he
has learned all there is to know of chemistry.
The following hints may serve to show how the work ma}7 be
lessened or how the course may be shortened : —
1. Omit the experiments marked OP.
2. When two or more experiments tend toward the same
general result, omit as many as desirable, selecting those most
readily performed by the apparatus and working material
available.
3. Omit the rarer elements and their compounds.
4. In the compounds of the common elements, dwell at length
upon the most useful ones, e.g., in the compounds of nitrogen,
place the stronger work upon ammonia, nitrogen monoxide, and
nitric acid, omitting or dwelling but briefly upon the remaining
compounds.
5. The qualitative work may be curtailed by omitting some
of the separations, etc.
6. Sometimes, also, the teacher may prefer to modify the
order of presenting the various topics ; for example, he may
wish to discuss molecules more thoroughly at the outset, or he
XIV .TO THE TEACHER.
wish the class to experiment with the oxides of nitrogen
before discussing them in their bearing upon the law of multiple
proportions, etc., etc. In this way he may conform to his own
ideas of presentation.
CO^TETTTS.
HISTORICAL SKETCH.
PAGE.
THE ANCIENTS. - — The Arabs. — Alchemy of the Middle Ages.
— Medical Chemistry. — Pneumatic Chemistry. — Modern
Chemistry 1-7
INTRODUCTION.
EXPERIMENTATION. — Elements. — Compounds. — Chemistry
Defined. — Three Forms of Matter. — Chemism. — Laws of
Definite and Multiple Proportions. — Combining Number.
— Atomic Theory. — Atomic Weight. — Determination of
Atomic Weight. — Names of the Elements. — Symbols. —
A Table of the Elements 8-22
CHAPTER I.
OXYGEN: its occurrence, preparation, properties, and tests. —
The Bunsen Burner and the Blow-pipe. — Ozone: prepara-
tion, properties, and tests 23-33
CHAPTER II.
HYDROGEN : its occurrence, etc. — Water : its occurrence, etc.
— Composition of Water. — The Oxy-hydrogen Blow-pipe. —
Impurities in Drinking-water, and Tests for. — Hydrogen
Dioxide : its preparation, etc 34-49
CHAPTER III.
NITROGEN : its occurrence, etc. — Ammonia : its occurrence,
etc. — Nitrogen Monoxide: its occurrence, etc. — Nitrogen
XVI CONTENTS.
PAGE.
Dioxide. — Nitrogen Trioxide. — Nitrogen Tetroxide. —
Nitrogen Pentoxide. — The Nitrogen Acids : Hyponitrous,
Nitrous, and Nitric Acids. — Hydroxylamine. — Estimation
of Ammonia in Drinking-water 50-72
CHAPTER IV.
BINARY COMPOUNDS. — Acids. — Bases. — Salts. — Acid and
Normal Salts. — Writing Equations 73-81
CHAPTER V.
THE ATMOSPHERE. — Atmospheric Pressure. — Measurement of
the Temperature of the Atmosphere. — Impurities in the
Atmosphere. — Determination of the Volumes of Oxygen and
Nitrogen in the Atmosphere. — Effects of Heat and Pressure
on the Volume of a Gas. — Weight and Density of Gases. —
Useful Problems 82-91
CHAPTER VI.
CHLORINE : its occurrence, etc. — Hydrochloric Acid : its
preparation, etc. — Oxides of Chlorine: Monoxide, Trioxide,
Tetroxide ; their preparation, etc. — The Chlorine Oxacids :
Hypochlorous, Chlorous, Chloric, and Perchloric Acids, and
their preparation, etc. — Estimation of Chlorine in Drinking-
water % ...... 92-107
CHAPTER VII.
BROMINE : its occurrence, etc. — Hydrobromic Acid : its prepara-
tion, etc. — Hypobromous, Bromous, and Perbromic Acids,
and their preparation, etc 108-114
CHAPTER VIII.
IODINE and FLUORINE. — Occurrence, etc., of Iodine. — Hydri-
odic Acid : its preparation, etc. — lodic an4 Periodic
Acids, and their preparation, etc. — Fluorine. — Hydrofluoric
Acid . 115-124
CONTENTS. XV11
CHAPTER IX.
PAGE.
CARBON : its occurrence, etc. — Methane. — Ethylene. — Acety-
lene. — Illuminating Gas. — Carbon Monoxide. — Carbon
Dioxide. — The Carbonates. — Cyanogen. — Prussic Acid. —
Estimation of Carbon Dioxide in Living-rooms . . . 125-148
CHAPTER X.
MOLECULES. — Avogadro's Hypothesis and the Computation of
Molecular Weights. — Determination of Atomic Weights by
Means of Avogadro's Hypothesis. — Valence. — Substituting
Power and Valence 149-156
CHAPTER XI.
SULPHUR: its occurrence, etc. — Hydrogen Sulphide. — Hydro-
• gen Per-sulphide. — Sulphur Dioxide. — Sulphur Trioxide.
— Sulphurous Acid. — Sulphuric Acid. — Nordhausen, or
Fuming Sulphuric Acid. — Thiosulphuric Acid and the
Thiosulphates. — Carbon Bisulphide. — SELENIUM : its occur-
rence, etc. — TELLURIUM : its occurrence, etc 157-183
CHAPTER XII.
SILICON and BORON. — Occurrence, etc., of Silicon. — Silica. —
The Silicon Oxacids and the Silicates. — Other Compounds
of Silicon. — Occurrence, etc., of Boron. — Boron Com-
pounds, etc , 184-192
CHAPTER XIII.
PHOSPHORUS : its occurrence, etc. — Phosphorus and Hydrogen.
— Phosphorus Oxides. — Phosphorus Oxacids : Hypophos-
phorus Acid, Phosphorous Acid, Phosphoric Acid, Meta-
phosphoric Acid, Pyrophosphoric Acid. — Examination of
Unknown Substances for the Acids previously given . . 193-207
CHAPTER XIV.
INTRODUCTION TO THE METALS. — Properties. — Alloys. —
Analytical Classification of the Metals. — Salts of the Metals.
— A Natural Classification of the Elements . . 208-223
XV111 CONTENTS.
CHAPTER XV.
THE FIRST GROUP METALS. — Lead, and its occurrence, prep-
aration, properties, uses, compounds, and tests. — Silver: its
occurrence, preparation, etc. — Mercury: its occurrence, etc.
— Separation and Identification of Lead, Silver, and
Mercury • 224-240
CHAPTER XVI.
THE SECOND GROUP METALS. — Arsenic, and its occurrence,
preparation, properties, uses, compounds, and tests. — Anti-
mony : its occurrence, etc. — Tin : its occurrence, etc. — Sepa-
ration and Identification of Arsenic, Antimony, and Tin. —
Bismuth : its occurrence, etc. — Copper : its occurrence, etc. —
Cadmium: its occurrence, etc. — Separation and Identification
of Bismuth, Copper, and Cadmium. • — Separation, etc., of
the Metals of the Second Group. — Separation of the Metals
of Groups I. and II. — The Rare Metals of the Second
Group: Gold, Platinum, Palladium, Ruthenium, Iridium,
Rhodium, Osmium, Tungsten, Molybdenum 241-272
CHAPTER XVII.
THE THIRD GROUP METALS. — Iron : its occurrence. — Iron
Ore. — Preparation of Iron. — The Iron Furnace. — Wrought
Iron. — Steel. — Properties, Uses, and Compounds of Iron. —
Tests for Iron. — Chromium : its occurrence, etc. — Tests for
Chromium. — Aluminum : its occurrence, etc. — Tests for Alu-
minum. — Separation and Identification of Iron, Chromium,
and Aluminum. — Nickel : its occurrence, etc. — Tests for
Nickel. — Cobalt : its occurrence, etc. — Tests for Cobalt. —
Separation and Identification of Nickel and Cobalt. — Man-
ganese : its occurrence, etc. — Tests for Manganese. — Zinc :
its occurrence, etc. — Separation and Identification of Nickel,
Cobalt, Manganese, and Zinc. — The Rare Metals of the
Third Group : Beryllium, Indium, Gallium, Yttrium, Lan-
thanum, Cerium, Didymium, Terbium, Erbium, Thorium,
Titanium, Zirconium, Uranium, Tantalum, Niobium, and
Vanadium . 273-308
CONTENTS. XIX
CHAPTER XVIII.
PAGE.
THE FOURTH GROUP METALS. — Barium : its occurrence, etc. —
Tests for Barium. — Strontium : its occurrence, etc. — Tests
for Strontium. — Calcium : its occurrence, etc. — Tests for
Calcium. — Separation and Identification of Barium, Stron-
tium, and Magnesium. — Magnesium : its occurrence, etc. —
Tests for Magnesium 309-319
CHAPTER XIX.
THE FIFTH GROUP METALS. — Potassium: its occurrence, etc. —
Tests for Potassium. — Sodium : its occurrence, etc. — Soda
Preparation by the Black Ash and Ammonia Processes. —
Tests for Sodium. — Ammonium. — The Ammonium Salts.
— The Analysis of Unknown Substances 320-340
APPENDIX.
DEVOTED to The Laboratory, Apparatus, Working Material,
Reagents, etc., etc 341-366
HISTOEIOAL SKETCH.
1. The word Chemistry is probably derived from Che-
mia, which is an old name for Egypt. The word signifies
simply the Egyptian art ; and it was so called since chem-
istry was first practised by the Egyptians.
Like all sciences which have to deal with Nature,
chemistry has been developed by a k>ng and tedious
series of experiments. Since the art of experimenting
is a comparatively modern one, the Ancients, as one
would naturally infer, were not deeply versed in this
science. The principal obstacle in the way of their prog-
ress is apparent when we know that they made great use
of the speculative method ; that is, when they wanted an
explanation of any fact in Nature, they simply thought
about it, without seeking to verify their conclusions by
the test of rigid experiment.
The Egyptian priests were the learned class of their
time ; and their researches were carried on with such an
air of mystery, and at such uncanny times, and in such
secret places, that Chemistry was spoken of as the Black,
or Secret Art. We find, however, that the Egyptians
possessed a considerable knowledge of the arts of dye-
ing, painting, and glass-making ; and that they were quite
skilled in metallurgy and the manufacture of pottery.
About the time of Aristotle (fourth century B.C.) it was
2 HISTORICAL SKETCH.
believed by some that all bodies are only modifications of
one fundamental substance ; by others, that all substances
are but the dwelling-places of four properties, — viz., heat,
cold, moisture, and dryness, — and that these four prop-
erties of matter are best represented in the four sub-
stances, fire, air, water, and earth. It was further believed
that these properties could be transferred from one body
to another, and, as a consequence, that the ordinary metals,
such as iron, could be transformed into the noble metal,
gold. It will be readily understood that this thought fur-
nished a powerful incentive to work, which incidentally
contributed something towards the advancement of chem-
istry. Considering the object he had in view, it is not
surprising that the chemist practised his art in caverns
and at night, where no prying eyes could see his opera-
tions, nor that he recorded his transactions in ambiguous
terms and in mysterious characters.
We thus find the ancients making but little progress in
true chemical science. Moreover, we now know that their
pernicious methods and theories were detrimental for many
centuries afterwards, notwithstanding the fact that chem-
istry originated in these self-same theories and methods.
2. The Arabs, in the year 640 A.D., invaded Egypt and
became acquainted with the Egyptian sciences.
Geber, an Arabian alchemist of the eighth century (the
Arabs gave chemistry the name Al-Chemia), wrote the first
book on chemistry. He understood many chemical manip-
ulations, discovered a solvent for gold, a mixture of nitric
and hydrochloric acids or aqua regia, and proposed the first
theory of the chemical composition of the metals, viz., that
sulphur and mercury were the simple or primary sub-
stances from which all the different metals are derived.
HISTORICAL SKETCH. 3
In this period, then, we find an encouraging advance ;
chemical processes are becoming more generally known,
and a suggestive though erroneous theory is announced,
which is destined to develop, through many modifications,
from error into truth. As an instance of the manner in
which this theory was afterwards modified and extended,
we may here mention the fact that Basil Valentine of the
fourteenth century, accepting sulphur and mercury as
the primal elements, extended the conception to all sub-
stances ; and that Boyle, three centuries later, doubtlessly
influenced by this same theory to investigate this problem,
announced the true solution.
3. During the Middle Ages the Arabians fostered the
sciences. Their academies in Spain were sought by stu-
dents from all parts of the civilized world ; these philoso-
phers, returning to their native countries, taught chemistry
there. Thus we find, in the thirteenth century, Raymond
Lully in Spain, Albertus Magnus in Germany, Arnold
Villanovanus in France, and Roger Bacon in England.
All these believed in the transmutation of the metals, and
the philosophy of their time teemed with mysticism, and
nonsense. We must here note that the all-absorbing theme
was the Philosopher's Stone, a substance which should
transform the baser metals into precious gold. The
writings of this period are extravagant, confused, and
purposely so written that they are nearly unintelligible.
Bacon, however, to clear himself of the charge of sorcery
(chemistry was still the Black Art), wrote a treatise in
which he showed that many things supposed to be caused
by supernatural agencies are produced by natural causes.
The search for the Philosopher's Stone during this
period brought to light many facts in inorganic chem-
4 HISTORICAL SKETCH.
istry; and thus do we find alchemy slowly but surely
paving the way for genuine chemistry.
4. In the era of Medical Chemistry, chemists directed
their investigations into a different channel. They then
sought the Elixir Vitae, or Elixir of Life, — a cordial
which should cure all the ills of mankind, and give per-
petual youth. By a strange misinterpretation of Aristotle,
some chemists also conceived the idea that the Philoso-
pher's Stone, when found, would achieve the same results.
Paracelsus (1493-1541) was the most noted of these
investigators. By his great achievements he earned the
title, The Father of Medicine.
Agricola (1490-1555) wrote the first treatise on Metal-
lurgy and Mining.
Libaviiis wrote the first Hand-Book of Chemistry, his
Alchemia, which was published in 1595.
Van Helmont (1577-1644) deserves special mention,
since he was the first to emancipate himself from the
theories of the Aristotelian school. He also discovered
various gases, and showed that metals are not destroyed
when dissolved in acids. But he, too, had his delusion: it
was liis Alkahest, a universal solvent as well as a universal
medicine.
Robert Boyle (1627-1691) advanced still further: he
claimed that the exact number of the elements was not
known, and he clearly stated the difference between the
elements and the compound substances. He also raised
chemistry to the dignity of a true science, which was not
to be studied as a part of any other, but as one of the
great Natural Sciences.
During this period many useful and potent medicines
were discovered, and, although error was by no means
HISTOKICAL SKETCH. 5
completely banished, the fundamental principles of chem-
istry were well grounded in truth. Hereafter, the history
of chemistry is a history of improvements, discoveries,
and researches extending to all the different branches
into which this science has developed.
5. Pneumatic Chemistry was the next phase in the
development of our science. This period was remarkable
for the investigation of the properties of gases, and the
phenomena of combustion.
Stahl sought to explain combustion by assuming the
existence of a combustible principle, or element, which he
termed Phlogiston. According to his views, this element
must be taken away from combustible bodies to render
them incombustible.
Among the believers in Phlogiston were three remarka-
ble men : —
1. Joseph Priestley, who discovered oxygen gas in 1774,
and afterwards other and important gases.
2. Henry Cavendish (1731-1810), who experimented
with inflammable air (hydrogen gas), determined the
density of the gases, and discovered the unvarying com-
position of the atmosphere.
3. Charles William Scheele (1742-1786), a Swedish
chemist, who discovered chlorine gas, prussic acid, gly-
cerine, and the pigment, Scheele's green. He also made
such other researches that he is entitled to be placed
among the founders of Quantitative Analysis.
None of these three ever discovered the true explana-
tion of combustion. The Phlogiston theory, however,
could not stand the test of rigid experiment ; and Lavoi-
sier, by exposing its fallacies, ushered in the new era of
chemistry, or
6 HISTORICAL SKETCH.
6. The Modern Era. — From his own experiments and
those of his predecessors, Lavoisier determined that a burn-
ing body unites with, or takes up a combustible element,
oxygen. By the use of the balance he discovered the
great fundamental truth, that, however great the changes
matter .may undergo, no loss in weight occurs, or, in other
words, that matter is indestructible. He also introduced
a system of chemical nomenclature, which has been of
inestimable value, as chemists not only disagreed as to the
names of the substances with which they were acquainted,
but often and purposely called one substance by so many
names that their meaning was not at all certain.
Dalton, next to Lavoisier, gave a great impetus to the
study of chemical phenomena by the discovery of the laws
of combination, known as the laws of " definite and mul-
tiple proportions," and by the propounding of the atomic
theory.
Gay Lussac discovered the law of combination of gases
by volume.
In 1808 Sir Humphry Davy discovered, by means of
electrolysis, the compound nature of the alkalies.
In 1828 Wohler prepared urea from inorganic sub-
stances, thus crossing out the division line between or-
ganic and mineral chemist^.
Spectrum analysis, dating back scarcely farther than
1860, has not only revealed the existence of many new
terrestrial elements, — such as caesium, thallium, rubi-
dium, indium, etc., — but has enabled us to determine
the composition of the sun and stars themselves.
Chemistry is no longer the Black Art, nor the handmaid
of astrology, but a legitimate science, exact in its methods,
and beneficent in its results. While, as a pure science, its
HISTOEICAL SKETCH. 7
aim is the investigation of truth, it has in its practical
application formed an important factor in the industries
of all civilized countries.
SUGGESTION. Read Rodwell's Birth of Chemistry; Roscoe's Spectrum
Analysis ; Whewell's History of the Inductive Sciences, pp. 261-310; Roscoe
and Schorlemmer's Treatise, pp. 1-40. Write short biographical sketches
of the chemists mentioned (consult an Encyclopedia).
rNTBODTJCTION,
DEFINITIONS. — LAWS OF COMBINATION IN DEFINITE AND
MULTIPLE PROPORTIONS. — ATOMIC THEORY. — ATOMIC
WEIGHTS. — NAMES OF ELEMENTS. — SYMBOLS. — TABLE
OF THE ELEMENTS.
7. To Experiment with a substance is to place it
under certain conditions or with certain substances to
ascertain its properties and behavior.
An experiment is a question intelligently put to
Nature.
EXPERIMENT 1 p. (To the student.) Since this is your
first experiment in chemistry, you may feel uncertain as to
what you are expected to do, or how }TOU are to derive the
most benefit from your work. In general, it is a safe policy
always to work carefully, and to note all phenomena that
occur ; from these phenomena you are then expected to derive
certain desired conclusions. It is true, that, for various rea-
sons, you may sometimes need assistance in reaching these
conclusions ; in such cases you must necessarily rely upon the
experience of others. Although this latter method is a legiti-
mate and often an indispensable way of obtaining knowledge,
we may safely say that he has the most truly scientific spirit
and methods, who, so far as possible, works and observes for
himself.
In the experiments }'ou are about to make, you may watch
for any changes that take place in the substances experimented
upon. Some of these changes may be perceptible to the sense
INTRODUCTION. 9
of sight, and some to the sense of smell ; others may be per-
ceptible to the sense of touch ; and still others to general sensi-
bility ; but, as a usual thing, the chemist depends mainly upon
sight and smell to detect any changes in the substances upon
which he is working. Now let us ask of Nature a few ques-
tions.
Steadily and persistently hold a platinum wire in a Bunsen
flame (Art. 28). What occurs? Now cut off a very short
piece (say 2mm) of the wire, place it upon a piece of charcoal,
and heat it by means of the blow-pipe flame (Art. 28) . What
takes place? Then cover the bit of wire with a mixture of
sodium carbonate (Na2CO3) and potassium nitrate (KNO3),
and slightly moisten the whole. Again heat in the blow-pipe
flame as before. What results? Now wash the piece of wire
clean, and place it in a test-tube ; then add nitric acid (HNO3) ,
and warm gently in the Bunsen flame. What occurs? Again
wash the wire, add hydrochloric acid (HC1), and warm as
before. What have you observed? You may possibly be
inclined to answer, u Nothing of importance." But let us
see. Did you succeed in separating the platinum into two or
more different substances ? Assuredly not ; nor could you have
so separated it by any process known to man. Now that is
important, since there are, besides platinum, about sixty- eight
other substances that have not been separated into simpler
ones : and these should have a class name. Hence the fol-
lowing name and definition : —
8. An Element is a substance that has not been divided
into two or more simpler substances.
EXAMPLES. Gold, Iron, Silver, Tin, Oxygen, Potassium.
NOTE (to the student). You are not to infer that all these sixty-eight
elements would behave precisely like platinum : such, indeed, is not the
case. Very few of them could have withstood the above treatment with-
out undergoing marked changes. None of them, however, would have
yielded two different substances, in which respect alone do they all agree
with platinum.
10 INTRODUCTION.
QUERY. What is a definition 1
SUGGESTION. Try, as above, bits of lead, copper, iron, zinc, etc. Com-
pare the results with those obtained from platinum.
EXP. 2 p. Place in a test-tube a short piece of thoroughly
dried pine wood as thick as a lead-pencil. Heat it over a
Bunsen flame, or a spirit-lamp. What collects on the sides of
the tube, what escapes, and what remains behind? Burn this
remainder on platinum foil, and what will then remain ?
QUERIES. What did you obtain from the wood 1 What became of the
charcoal when burned 1 Did any tar escape with the smoke q How do
you know ? Any water ? Prove it. (SuG. Hold a piece of cold glass
in the escaping vapors.) Will a piece of brick give the same results 1
Try it.
EXP. 3 P. Place in a hard glass tube, open at both ends, a
small piece of galena (PbS). Hold the tube somewhat slant-
ing in the Bunsen flame, so that the greatest heat shall strike
underneath the galena. Notice the odor of the fumes which
soon issue from the tube. These are the fumes of burning sul-
phur. Now place the residue in a shallow, cup-shaped cavity,
which you are to make in a piece of charcoal. Cover the resi-
due with sodium carbonate (Na2CO3) , and slightly moisten the
whole. Heat it before the blow-pipe flame and you will obtain
a metallic bead. What metal is it?
It is evident that wood and galena are not elements ;
and, as the student's experience increases, he will learn
that there is a very large class of substances which can
thus be separated into simpler ones, and that these simpler
substances are united in definite proportions by weight.
Hence the following name and definition : —
9. A Compound (chemical) consists of two or more ele-
ments chemically combined in definite proportions.
(Art. 17.)
Ex. Salt (NaCl) ; Water (H2O) ; Sugar (C,2H23OU).
INTRODUCTION. 11
EXP. 4 p. Mix thoroughly 0.56g of very fine iron-filiugH
and 0.32s powdered sulphur. Although the mixture resembles
neither iron nor sulphur, this is only a mechanical mixture,
and the microscope reveals the particles of iron and sulphur
lying side by side : moreover, they may be separated by
mechanical means. Now heat one-half the mixture to red-
ness in an iron spoon ; a glow diffuses itself throughout the
mass, and the iron combines with the sulphur in definite pro-
portions. No microscope can now distinguish the iron and
sulphur particles, nor can they be separated except by chemi-
cal means. The iron and sulphur have exactly entered into
chemical union.
QUERIES. Can you, with a magnet, separate the iron from the sulphur
before heating7 Try it. Will bisulphide of carbon (CS2) dissolve out the
sulphur from the iron particles before heating 1 Try it. Should the sul-
phur dissolve, evaporate the solution to dryness on a watch crystal, and
see if the sulphur will remain as a residue.
After heating, pulverize the mass and try as above. What difference
do you find "*
From the above we derive the two following defini-
tions : —
10. A Mechanical Mixture is formed when substances
are put together in no definite proportions, and the result-
ing substance retains the properties of its constituents.
11. A Chemical Combination or Reaction takes place
when two or more substances unite in definite proportions
to form one or more substances entirely different from
the original ones.
12. Chemistry is that science which treats of the ele-
ments found in nature, their properties, compounds, and
actions and reactions upon one another.
Matter exists in three forms ; viz., Solids, Liquids, and
Gases.
12 . INTRODUCTION.
13. Solids do not readily change their forms, since in
them the attractive (inter-molecular) forces exceed the
repellent forces.
14. Liquids do readily change their forms, since their
attractive and repellent (inter-molecular) forces are equal,
or nearly so.
15. In Gases, the repellent forces are greater than the
attractive forces, consequently gases always tend to occupy
a larger space.
SUG. Name several solids. Liquids. Gases. Show, by heating a
piece of ice till it vaporizes, that water exists in all three conditions.
16. Chemism is an attractive force which is exerted
between the elements, causing them to enter into com-
bination with one another.
NOTE. Cohesion and chemism tend to draw particles together. In all
solid and liquid compound bodies, both chemism and cohesion operate :
the former holds the elements together, and determines the composition
of the body ; the latter holds the particles of the compound together, and
gives us the mass. Heat is a repellent force, and tends to separate the
small particles of all bodies, as is shown by the expansion of bodies when
heated.
17. Law of Definite Proportions. — If we examine
any chemical compound, — such, for example, as water,
which consists of the elements hydrogen and oxygen ;
common salt, which consists of the elements sodium and
chlorine, — we find that the compound always contains
exactly the same proportions of its constituents. Water
always contains 88.89 per cent of oxygen and 11.11 per
cent of hydrogen ; common salt always contains 39.32 per
cent of sodium and 60.68 per cent of chlorine. As a re-
sult of the careful analysis of a very large number of
INTRODUCTION. 13
chemical compounds, the law of definite proportions was
propounded. The law may be stated in this form : —
Any given chemical compound alivays contains the same
elements in the same proportions by weight.
KEM. It is, of course, impossible for the beginner to prove the cor-
rectness of this law, for the reason that the proof cannot be furnished
without the employment of some of the most delicate and difficult chemi-
cal processes.
18. Law of Multiple Proportions. — Some elements
form more than one compound with each other. Thus
hydrogen and oxygen form not only water but hydrogen
dioxide ; iron and sulphur form three compounds ; nitro-
gen and oxygen form five compounds. If we examine the
proportions by weight in which the elements unite, we find
very curious and interesting relations. Thus, in water we
find : hydrogen 1 part, oxygen 8 parts ; in hydrogen diox-
ide, hydrogen 1 part, oxygen 16 parts. (See Art. 38.)
In the compounds of iron and sulphur (Art. 293), there
are :
Compound 1, 32 parts of sulphur and 56 parts of iron.
Compound 2, 64 " " " 56 " "
Compound 3, 96 " " 112 " "
In the compounds of nitrogen and oxygen (Art. 56),
there are:
Compound 1, 28 parts of nitrogen and 16 parts of oxygen.
Compound 2, 28 " k< " 32 " "
Compound 3, 28 " " 48 " "
Compound 4, 28 '-' " " 64 " "
Compound 5, 28 " " " 80 " "
The amount of oxygen in the second compound of
hydrogen and oxygen is just twice as great, — not one
14 . INTRODUCTION.
and one-half, nor any fractional number of times, as great,
as in the first.
The amounts of sulphur in the three compounds of iron
and sulphur bear to each other the relation of 1 : 2 : 3 ; and
the amounts of iron are to each other as 1:1:2.
Finally, in the compounds of oxygen and nitrogen, the
amounts of oxygen are to each other as 1:2:3:4:5; the
amount of nitrogen remaining constant.
These cases illustrate what is known as the law of multi-
ple proportions, which may be stated thus :
If two elements, A and B, form several compounds with
each other, and we consider any fixed amount of A, then the
different amounts of B which combine with this fixed amount
of A bear a simple ratio to each other.
19. Combining- Number. — For each element we can
select a certain number which will enable us always to
express the proportion by weight in which this element
enters into combination.
Thus, we .may select the number 16 for oxygen, and we
find that no matter what the compound may be in which
we find the oxygen, its proportion may be expressed by 16
or some simple multiple of 16. In the same way we find
that 32 may be selected for sulphur ; 14 for nitrogen ; 56
for iron, etc., etc. The figures thus selected are known as
the combining numbers. Elements always combine with
each other in the proportions expressed by their combin-
ing numbers, or by simple multiples of these numbers.
Thus, according to this, if sulphur and oxygen unite, we
would expect to find them in their compounds in the
proportions of 32 parts of sulphur to 16 parts of oxy-
gen ; 32 parts of sulphur to 32 parts of oxygen ; 32 parts
of sulphur to 48 parts of oxygen, etc. Compounds cor-
INTRODUCTION. 15
responding to the last two proportions are known. (See
Art. 164.)
20. Atomic Theory. — To account for the fact that ele-
ments unite in fixed proportions, it is assumed that all
matter is made up of indivisible particles called atoms, and
that each different kind of atom has its own particular
weight. When chemical combination takes place, it is
supposed that this consists of a union of the atoms of the
elements which take part in the action. Thus, when iron
and sulphur are brought together, at first no action takes
place ; but when they are very intimately mixed, and the
mixture heated, it is believed that each atom of iron seizes
upon an atom of sulphur, uniting with it. Now, as these
atoms have definite weights, it follows that, no matter how
many unite, the compound formed must always contain
the elements in the proportion of the weights of the atoms.
The simplest kind of combination is that in which the
elements unite in the proportion of one atom of one ele-
ment to one of the other. But the elements may unite in
the proportion of one atom of one to two, or three, or even
four of the other, etc. Or, two atoms of one may unite
with three of another, etc. Hence, it follows that the
amounts of any element found in different compounds
must bear simple relations to each other.
21. Atomic Weights. — The numbers called combining
numbers are believed to express the relative weights of the
atoms of the elements, and are now called atomic iveights.
The numbers now in use are intended to express the
weights of the atoms of the elements as compared with the
weight of the atom of hydrogen taken as unity. Thus,
when we say that the atomic weight of oxygen is 16, and
that of nitrogen 14, we mean that the weight of the atom
16 INTRODUCTION.
of oxygen is 16 times as great as that of the atom of
hydrogen ; and that the weight of the atom of nitrogen is
14 times as great as that of hydrogen.
22. Determination of Atomic Weights. — To deter-
mine the atomic weight of an element is by no means a
simple matter; indeed, it is extremely difficult. If all
the elements united with each other in only one propor-
tion it would not be difficult to agree upon atomic weights.
Thus chlorine and hydrogen unite with each other in the
proportion of 35.5 parts of chlorine to 1 of hydrogen ;
bromine and hydrogen in the proportion of 80 parts of
bromine to 1 of hydrogen ; iodine and hydrogen in the
proportion of 127 parts of iodine to 1 of hydrogen ; and
these elements do not unite with hydrogen in any other
proportions. Hence, we may assume that in the com-
pounds formed we have, in the first place, one atom of
chlorine united with one atom of hydrogen ; in the second,
one atom of bromine with one of hydrogen ; and in the
third, one of iodine with one of hydrogen. We are thus
led to the conclusion that the atom of chlorine weighs
35.5 times as much as the atom of hydrogen, or that the
atomic weight of chlorine is 35.5 ; and, in the same way,
that the atomic weight of bromine is 80, and that of iodine
127.
When, however, two elements unite in more than one
proportion, — and this is the rule rather than the excep-
tion,— it is clear that we must be left in doubt as to the
number to select as the atomic weight. Thus, hydrogen and
oxygen, as was remarked above, unite in two different pro-
portions. In the -first there are 8 parts of oxygen to 1 of
hydrogen ; in the second, 16 parts of oxygen to 1 of hydro-
gen. From this we might conclude that 8 is the atomic
INTRODUCTION. 17
weight of oxygen. But we may just as well express the
proportions by saying that in the first there are 16 parts
of oxygen to 2 of hydrogen ; and in the second, 16 parts
of oxygen to 1 of hydrogen. And we might, with equal
justice, conclude that 16 is the atomic weight of oxygen.
We shall find that two methods are in general use for
the determination of atomic weights. The first is based
upon a consideration of the specific gravity of elements
and compounds in the form of gas or vapor ; the second,
upon the specific heat of elements and compounds. These
methods will be described after some of the elements and
their compounds have been considered. (Art. 157.)
23. Names of the Elements. - - The ancients were
acquainted with only seven elements ; viz., gold, silver,
copper, iron, mercury, lead, and tin. They dedicated
these to the heavenly bodies ; e.g., silver was dedicated
to the moon or luna. In this fanciful way some of the
names of chemical compounds originated ; e.g., nitrate of
silver is yet called lunar caustic.
The elements have received their names in different
ways : —
1. Some retain their ancient names.
2. Some are named from some marked characteristic ;
e.g., phosphorus, light-bearer ; bromine, a stench.
3. The names of some end in " ine " or " on," to indi-
cate a similarity of properties in those so terminating/
4. Some are named from the place of their discovery.
5. The names of recently discovered substances pos-
sessing metallic properties end in " um " or " ium."
SUG. Student find illustrations to above from Art. 25.
24. Symbols. — In expressing the composition of chemi-
cal compounds, it is desirable to have a system of symbols.
1 8 INTRODUCTION.
Those now in use consist of letters which stand for the
names of the different elements. Thus, O stands for Oxy-
gen, H for Hydrogen, N for Nitrogen, etc.
When only one element is known, whose name begins
with a certain letter of the alphabet, that letter is used
as the symbol.
When two or more are known, the names of which begin
with the same letter, that one best known or first discov-
ered is generally designated by the letter, while the others
are designated by this letter and some other letter occur-
ring in the name, e.g., Carbon, C ; Chlorine, Cl ; Calcium,
Ca; Caesium, Cs ; Cadmium, Cd; Cobalt, Co; etc.
Some elements have symbols derived from their Latin
names. This is perplexing to the student, but this list
will explain : —
Antimony, Sb, from Stibium.
Copper, Cu, " Cuprum.
Gold, Au, " Aurum.
Iron, Fe, " Ferrum.
Lead, Pb, " Plumbum.
Mercury, Hg, " Hydrargyrum.
Potassium, K, from Kalium.
Silver, Ag, " Argentum.
Sodium, Na, " Natrium.
Tin, Sn, " Stannum.
Tungsten, W, " Wolframium.
The symbol stands not only for the name of the element,
but for its atom. Thus, O means not only oxygen, but an
atom of oxygen ; 2 O or O2 means two atoms of oxygen,
etc. In expressing the composition of bodies by means of
these symbols, we simply place the latter side by side.
Thus, HC1 stands for a body which consists of hydrogen
arid chlorine in the proportions, 1 part by weight of hydro-
gen to 35.5 of chlorine ; or, in terms of the Atomic Theory,
it stands for a body which is formed by the union of hydro-
gen and chlorine in the proportion of 1 atom of hydrogen
to 1 of chlorine. An expression like HC1 is called a for-
mula.
INTRODUCTION. 19
In expressing the composition of a body in which more
than one atom of the same kind is present, a small figure
is added below the line to the right of its symbol. Thus,
potassium nitrate, which consists of potassium, nitrogen,
and oxygen, in the proportion of 1 atom potassium, 1
nitrogen, alid 3 oxygen, is written KNO3. A large figure
placed before a formula affects every symbol in the for-
mula. Thus, if we want to express two parts of potassium
nitrate, we usually write 2 KNO3, and not (KNO3)2. We
repeat a group of atoms (NO3, NH4, etc.) which we wish
to keep together as a whole (Art. 159), thus : Pb(NO3)2,
(NH4)2S.
Following is a list of the elements which have thus far
been discovered. The table includes not only the names
of the elements, but their atomic symbols, atomic weights,
— as determined by every available method, — and their
specific gravities.
The small Roman numerals or indices added to the sym-
bols are intended to indicate the valence (see Art. 158)
of the elements. Usually the symbol is written without
these.
20
INTRODUCTION,
25. A Table of the Elements.
Names.
Symbols.
Atomic
Weights.
Physical
condition at
ordinary
temperature.
Specific Gravity.
Aluminum
Al""
27.
Solid
* 2.60
Antimony
Sb'"'v
120.
K
6.71
Arsenic
As'"'v
75.
«
5.73
Barium
Ba"
137.
i t
3.75
Beryllium
Be"
9.
a
2.07
Bismuth
Bi'"'v
208.
a
9.80
Boron
Bo'"
11.
a
2.5?
Bromine
Br''v
80.
Liquid
3.187
Cadmium
Cd"
112.
Solid
8.60
Caesium
Cs'
133.
a
1.88
Calcium
Ca"
40.
a
1.57
Carbon
C""
12.
n
3.5-.6
Cerium
Ce'"'""
141.
it
6.68
Chlorine
Cl'-v
35.5
Gas
2.450
Chromium
O"">vi
52.
Solid
6.50
Cobalt
Co"-""
59.
tt
8.5-.7
Copper
Cu"
63.3
(i
8.95
Didymium
D'"
142.3
u
6.54
Erbium
E'"
166.
u
—
Fluorine
F'
19.
1.313
Gallium
Qfflf
. 69.
Solid
5.95
Gold
An''"'
196.5
u
19.32
Hydrogen
H'
1.
Gas
0.069
Indium
In""
113.6
Solid
7.42
Iodine
I''v
127.
it
4.948
Ividium
Jj.ff.fW.vi
193.
u
22.42
Iron
Feff,f"f,vi
56.
((
7.86
Lanthanum
La'"
138.2
«
6.10
Lead
Pb"' ""
207.
u
11.37
Lithium
Li'
7.
(i
0.59
Magnesium
Mg"« ""'vi
24.
t i
1.74
Manganese
Mn"
55.
it
8.03
Mercury
Hg"
200.
Liquid
13.55
Molybdenum
Mo"'""'vi
96.
Solid
8.60
Nickel
Ni»."«
58.
n
8.90
INTRODUCTION.
21
Names.
Symbols.
Atomic
Weights.
Physical
condition at
ordinary
temperature.
Specific Gravity.
Niobium
Nbv
94.
Solid
7.06
Nitrogen
Osmium
Os".',",vi
14.
199.
Gas
Solid
0.971
22.48
Oxygen
Palladium
0"
Pd"'""
16.
106.
Gas
Solid
1.105
11.40
Phosphorus
•pi, tll,v
31.
" {
Colorless 1.83
Red 2.20
Platinum
pt", ""
195.
"
21.50
Potassium
K'
39.
"
0.87
Rhodium
Ro"'""'vi
104.
"
12.10
Rubidium
Rb'
85.
<-<-
1.52
Ruthenium
Ru". ""»vi
103.5
1;
12.26
Samarium
Sm
150.
"
—
Scandium
Sc
44.
"
—
Selenium
Se"i '"'. vi
79. •
k i
4.50
Silicon
Si""
28.
U
2.39
Silver
Sodium
Ag'
Na'
108.
23.
k k
10.53
0.978
Strontium
Sr"
87.5
"
2.54
Sulphur
Tantalum
Tav
32.
182.
kt
2.05
10.40
Tellurium
«Pe"'""'Vi
125.?
kk
6.40
Terbium
Tb
148.5?
kk
—
Thallium
'PI'.'"
204.
"
11.85
Thorium
Th""
232.
"
11.00
Tin
Sn"'""
118.
"
7.29
Titanium
Ti"'""
48.
(I
—
Tungsten
Uranium
"117""", vi
184.
239.8
It
U
19.12
18.70
Vanadium
V'"'v
51.5
«
5.50
Ytterbium
Yb
173.
"
— .
Yttrium
Y'"
89.
it
—
Zinc
Zn"
65.
i«
7.15
Zirconium
Zr""
90.
kk
4.15
'2*2 INTRODUCTION.
HEM. 1. Many elements occurring in the earth have also been dis-
covered in the sun and stars.
REM. 2. Some elements occur in such very small quantities that their
properties are not accurately known ; while others have been discovered
so recently that they have not been fully investigated. ( See Chem. News,
Nov. 7, 1883, for List of Elements.)
REM. 3. More elements will be discovered, undoubtedly ; and some sub-
stances now known as elements may prove to be chemical compounds, as
our chemical researches advance.
SUG. Student, learn to spell the names of the elements. Learn to give
the symbol when the element is named, and vice versa.
REM. 4. In estimating the specific gravity of the elements, water is
taken as the standard for solids and liquids, while air is taken for gases.
REM. 5. The chemist also uses hydrogen as a standard for estimating
the density of gases, as will be explained later.
SUMMARY OF STUDENT'S WORK IN INTRODUCTION.
1. Make those experiments whose numbers are followed by the letter
"p."
GENERAL NOTE. When "P" follows the number of an experiment, the
student should be able to do the work : if, however, the student cannot
do the work, owing to various causes for which no text can provide, or if
the teacher wishes the work done differently, a few simple oral directions
from the teacher to the class will assist greatly.
Experiments marked " T " are to be made by the teacher before the
class. Let the pupils assist as much as possible.
In experiments marked " TP," it is advisable for the teacher to make the
experiment for the class before requiring the student to do it.
Experiments marked " OP " are optional.
Encourage the student to exert his ingenuity in overcoming obstacles,
and he will soon become quite independent in manipulation.
CHAPTER I.
OXYGEN : ITS OCCURRENCE, PREPARATION, PROPERTIES,
AND TESTS. — OZONE.
THE ELEMENT OXYGEN.
SYMBOL O". — ATOMIC WEIGHT, 1(5; SPECIFIC
GRAVITY, 1.1056.
26. Occurrence. — Oxygen occurs well-nigh everywhere
in nature. It constitutes 44 to 48 per cent of the weight
of the earth's crust, 88.89 per cent of water, arid about
23 per cent of the atmosphere.
Oxygen occurs in combination with every known ele-
ment except fluorine.
27. Preparation. — EXP. 5 p. Heat one gram mercuric
oxide, HgO, in a hard glass test-tube. The oxygen is driven
off, while the mercury is condensed on the sides of the tube.
Test the presence of the gas with a glowing match. (HgO =
Hg + O.)
QUERY. Aug. 1, 1774, Joseph Priestley made this experiment for the
first time. What gas did he discover ?
EXP. 6 P. Minium, or red oxide of lead, Pb3O4, is to be
heated as above. A part of the oxygen is driven from the
red oxide of lead with great difficulty. (Pb3O4 = 3 PbO + O.)
Test as before.
SUG. Try KC103 with and without Mn02, as above. Also heat, as
above, KC103 with a pine splinter. What occurs ? Explain.
24 THE ELEMENT OXYGEN.
QUERIES. Why does not the red oxide of lead, Pb304, part with its
oxygen as readily as mercuric oxide, HgO 1 Ans. The lead has a stronger
chemism for oxygen than mercury has. It is upon the principle of variable
degrees of chemism existing between different substances, that double
chemical reactions are always based. Do you obtain metallic lead in this
experiment ? Heat some red oxide of lead on charcoal, with sodium car-
bonate (Na2CO3), before the blow-pipe. Do you now obtain metallic lead ?
What, effect do the sodium carbonate and charcoal have on substances
treated thus ? Ans. The charcoal abstracts oxygen from the oxide, or
acts as a strong reducing agent. The sodium carbonate serves as a " flux,"
preventing the lead from again taking up atmospheric oxygen.
Oxygen can be prepared most easily from the com-
pounds which it forms with other elements ; as, mercuric
oxide, HgO ; manganese dioxide, MnO2 ', potassium chlo-
rate, KC1O3, etc.
Potassium chlorate, KC1O3, is the most available sub-
stance for preparing moderately large quantities in small
laboratories ; but if very large quantities are required, it
may be prepared more cheaply from manganese'dioxide,
although special apparatus is necessary.
Potassium chlorate gives up its oxygen more readily
and at a lower temperature when mixed with manganese
dioxide (KC1O3 = KC1 + 3 O ). The manganese dioxide is
unchanged. This method is best for laboratory use.
EXP. TT. Pulverize WO8 potassium chlorate, KC1O3, and mix
thoroughly with 25g manganese dioxide, MnO2. Place the mix-
ture in an iron or copper retort, and arrange to wash the gas
through two Woulff bottles : the first containing water, the sec-
ond sodium hydroxide, NaOH. Now heat strongly but care-
fully, and, when the air is expelled from the apparatus (test
with a match) , connect with the gas receiver. Notice that at
a certain point the gas is given off with great rapidity. The
heat must be moderated immediately to avoid accident. You
will thus obtain about 301 of pure oxygen gas.
THE ELEMENT OXYGEN. 25
CAUTION. Organic matter or carbon, when present, may produce a
serious explosion. It is best, therefore, to try a little of this mixture in a
test-tube before heating the retort. Use C. P. materials.
NOTE. It is always best to have the class present when preparing
such experiments as this last. Arrange the pneumatic trough, bell jars,
wires, etc., and make the following experiments in a dark room.
28. Properties of Oxygen. — EXP. ST. Plunge into ajar
of oxygen a glowing pencil of thoroughly charred bark charcoal.
It will burn with brilliant scintillations. (C -f- 2 O= C02.)
NOTE. This illustrates the combustion of fuel.
EXP. 9 T. Place a bundle of very fine iron wires, tipped
with sulphur and ignited, in a jar of ox}rgen. The wires will
burn with a reddish light, and at times with beautiful scintilla-
tions. (3 Fe + 4 O = Fe3O4.)
NOTE. This illustrates the great chemical activity of pure oxygen.
EXP. 10 T. File the end of a watch-spring till very thin.
Draw the temper in a spirit-lamp, and uncoil it. Make a
hook on the thin end, tip with sulphur, and ignite. Place
in a jar of oxygen. The spring will burn with great energy.
(3Fe-t-40 = Fe304.)
EXP. 11 T. Place a piece of phosphorus in a jar of oxygen.
Ignite. It burns with a brilliant white light. (2 P + 5 O = P2O5. )
See Phosphorus.
EXP. 12 T. Treat a piece of sulphur as in last experiment.
It burns with a violet light. (S -f- 2 O = SO2.)
NOTE. Do not allow the fumes from the burning of phosphorus and
sulphur to escape in the room, as they are very disagreeable.
EXP. 13 T. Cut zinc foil into fine strips ; make into a bundle ;
tip with sulphur ; ignite. White light in oxygen. (Zn + O =
ZnO.)
NOTE. The product formed is called " Philosopher's Wool."
26 THE ELEMENT OXYGEN.
Now that you have prepared and experimented with
oxygen, you will be ready to appreciate several of its
physical and chemical peculiarities which we term prop-
erties. Oxygen is an invisible, odorless, tasteless gas. Its
specific gravity is 1.10563 ; and I1 at 0° and 760mm pressure
weighs 1.430g.
It has been liquefied by a pressure of 25.85 atmospheres
at a temperature of - 131.6°. (Read R. and S., p. 516,
Vol. II., Pt. II.)
EXP. 14 OP. Place a live mouse upon a cork raft, under a
bell jar filled with air, over the pneumatic trough. Secure the
jar so that no communication with the outside air is possible.
Does the water rise in the jar? What does this indicate?
QUERIED How does the oxygen come in contact with the blood?
What harm ensues from persons living in a room without ventilation 7
Is the blood purified by a physical or chemical process ?
Oxygen is that constituent of air which is essential to
breathing, and all animals consume it. When inhaled, it
enters into combination with some of the tissues of the
body, actually burning them out, and thus liberating heat
and energy. Air that has been breathed over too many
times loses its vitality, the oxygen having been consumed.
As oxygen occurs in the atmosphere, it is largely diluted
with nitrogen.
EXP. 15 OP. Place a live fish in a sealed jar of water. What
follows? Why?
Water absorbs free oxygen, and fishes consume this
oxygen by means of their gills, which serve as lungs.
QUERY. How does a jet fountain render water fit for preserving the
life of fishes?
THE ELEMENT OXYGEN.
27
EXP. 16 OP. Place a burning taper in a closed jar of air.
When the oxygen of the air is consumed, what occurs?
QUERIES. Why does blowing the fire cause it to burn more briskly ?
Why does blowing a candle extinguish it 7 (See next Exp.)
Fire or Combustion is produced by the union of tho
fuel with atmospheric oxygen. Before a substance can
unite with oxygen, it must be heated to what is called its
burning temperature or kindling point ; and to produce
flame, it must be converted into a gas. A flame is a burn-
ing gas.
EXP. 17 P. Carefully place a bent glass tube very near the
wick of a lighted candle, within the flame zone. The gas
escaping from the wick will be forced up through the tube, and
may be lighted at the other end of the tube.
Since the gas which escapes from the wick burns only
when mixed with air, the flame of a candle has but a
thin outer zone, in which the gas is entirely consumed.
(Fig. 1.)
EXPLANATION OF FIG. 1.
T, bent glass tube.
c, centre of unconsumed gas.
p, zone of incomplete combus-
tion.
v, light zone, or zone of com-
plete combustion.
F', unconsumed gas burning at
end of glass tube.
QUERIES. What does this ex-
periment prove 1 Why can you
not ignite a lump of anthracite
coal with a match 1 Why does
a blow-pipe give such a hot
flame ?
28 THE ELEMENT OXYGEN.
BUNSEN BURNER. — This burner is almost exclusively
used in laboratories provided with gas for heating purposes.
It gives a very hot, clean flame, owing to the fact that it
is so arranged that the gas, before ignition, is thoroughly
mixed with air, which insures its complete combustion.
The tube e, shown in Fig. 2, is pierced with holes at its
base, and the gas is discharged at about the
height of these holes. Now, as the gas ascends
the tube e, it draws a current of air along with
it ; the air and gas mix in their ascent through
6, and burn with a hot, non-luminous flame
when ignited at the top of e. A ring a pierced
with holes surrounds e ; by turning this ring,
the holes dd may be closed, when the gas burns
with an ordinary luminous flame, which is the
flame used for the blow-pipe. And here the
student may learn the meaning of the terms Oxidizing
Flame and Reducing Flame, for which he will hereafter
find frequent application. The way in which these flames
are produced is as follows : —
1. The Oxidizing Flame. — First close the openings dd,
and make a moderately small luminous flame. Now place
the tip of the blow-pipe in the centre of the luminous
flame, and blow gently, using the cheeks like a bellows.
The oxidizing flame should be non-luminous. In case
you do not succeed in making it so, do not try to remedy
the evil by blowing harder, which will end only in ex-
hausting you, but moderate the flow of gas, and try again.
After a little practice you should be able to keep the flame
steady for half an hour, without becoming much fatigued.
This flame tends to oxidize substances when they are
placed in it, since it contains an excess of oxygen at a
very high temperature.
THE ELEMENT OXYGEN. 29
QUERIES. Whence comes this excess 1 Should air from the lungs be
used in blow-piping ? Why ?
SUG. Examine a blow-pipe, and give a short description.
2. The Reducing Flame is made by placing the jet of
the blow-pipe just outside of the luminous flame from the
burner, with the openings closed. This flame is slightly
luminous, and reduces or takes oxygen away from bodies
placed in it, since it contains an excess of hydrogen and
carbon (illuminating gas is a hydrogen-carbon compound)
at a high temperature ; both hydrogen and carbon have a
strong affinity for oxygen.
Moreover, the bead, or assay, is to be kept within the
zone of complete combustion (Fig. 1) when you are using
the oxidizing flame ; when using the reducing flame, the
proper position of the assay is within the zone of incom-
plete combustion.
QUERY. If you use your lungs for bellows, can you keep the blow-pipe
flames steady ?
EXP. 18 P. Make a borax bead by fusing borax on a loop of
platinum wire. Slightly moisten this bead in ferrous sulphate,
FeSO4, and heat a short time in the oxidizing flame. The bead
thus treated should be of a reddish color when hot, fading to a
light yellow when cold. Now heat the same bead persistently
in the reducing flame. It should become colorless unless too
strongly saturated with the ferrous sulphate, when it becomes
pale green. Unless the proper flames are used, these results
cannot be obtained.
Spontaneous Combustion. — The combination of oxygen
and other substances always produces a definite amount of
heat depending upon the nature of the substance. When
iron rusts slowly, the heat is imperceptible ; but when
greasy rags or waste are thrown in a heap, the heat pro-
30 THE ELEMENT OXYGEN.
ducecl by the oxidation of the oils may, in time, be suffi-
cient to raise the mass to the temperature of ignition.
This kind of action, known as spontaneous combustion, is
not unfrequently the cause of disastrous fires.
EXP. 19 P. Sift very fine iron-filings over the flame of an
ordinary lamp. What results?
QUERY. Why is this ? What does it illustrate ?
SUG. Try fine dust from a malt house, flour mill, wood-working shop,
etc., as above. The best place to collect the dust is from rafters or high
beams. Why ?
Fine dust collecting in the attics of large mills and malt
houses has sometimes exploded when ignited, causing great
destruction of life and propert}^. Again, the sun's rays,
when brought to a focus on inflammable substances, or
steam pipes coming in too close proximity to inflammable
substances, have produced unlooked-for conflagrations.
EXP. 20 P. Place green plant-leaves in the sunlight, under
a bell jar filled with water. Bubbles of oxygen will collect at
the top of the jar.
Oxygen is given off by plants growing in the sunlight.
Enough oxygen is returned to the air in this way to keep
its composition nearly uniform.
29. Tests for Free Oxygen. — 1. Char a small pine
stick, as a match, and, with one end glowing, place it in
a jar or current of free oxygen, when it will burst into
flame.
2. Fill a flask with oxygen gas. Pour in a small quan-
tity of potassium hydroxide, KOH. Shake, and no change
in the liquid takes place. Then add a small quantity of
pyrogallic acid, C6H3(OH)3. Shake again, and the liquid
turns brown, oxygen being absorbed.
OZONE. 31
N.B. In testing an unknown gas in this way, it is absolutely necessary
to exclude all air, as the free oxygen of the air gives this reaction. It is
best, therefore, to fill the flask over mercury. (See App.)
OZONE.
30. Ozone is a peculiar or allotropic form of oxygen
found in the atmosphere, and produced by electrical dis-
charges, or by evaporation, or by both. When an element
occurs in more than one form, the unusual one is called
an allofcopic form. It is easily prepared by several
methods.
31. Preparation. — EXP. 21 p. Place a small quantity of
a solution of potassium permanganate, K2Mn2O8, in a flask or
test-tube. Add a few drops strong sulphuric acid, H2SO4.
Notice the odor of the gas given off. It is ozone. Apply
Test 1 for ozone.
Ozone may be prepared by suspending a clean stick of
phosphorus in a closed jar containing a little water and
atmospheric air at a temperature of 15° to 20°. Ozone is
formed very rapidly.
When an electrical machine, in good working order, is
in action, a peculiar odor is observed which is due to
ozone. (Use Test 1, Art. 33, for ozone.)
Ozone may also be obtained by passing a silent electri-
cal discharge, carefully avoiding sparks, through a closed
jar of oxygen.
SUG. Produce ozone by one or all of the above methods.
32. Properties. — Ozone is three volumes of oxygen
condensed to two volumes, the condensation being proba-
bly accompanied by some deep-seated change in the relation
of the atoms.
32 OZONE.
There are good reasons for believing that the molecule
(Art. 155) of ordinary oxygen consists of two atoms, as
indicated by the formula O2, and that the molecule of
ozone should be represented by the formula O3. Ozone
is readily changed into ordinary oxygen. It is an active
oxidizing agent. When brought in contact with mercury
and some other substances in the dry state and at ordi-
nary temperatures, it converts them into oxides, and itself
becomes ordinary oxygen.
Ozone readily acts upon organic substances, and is sup-
posed to destroy the germs of contagious diseases. When
present in large quantities, ozone has an irritating effect
on the lining membranes of the throat and nostrils, where-
fore it should be dilute if inhaled.
Atmospheric ozone is more plentiful in the open country
than in cities, and more is found out of doors than in
dwellings. (Why ?)
NOTE. It is extremely difficult to determine whether the substance in
the atmosphere which is commonly called ozone, is really ozone or not.
There are certainly other substances present which in some of their
properties closely resemble it; such, for example, as hydrogen dioxide
(Art. 44).
33. Tests for Ozone. — 1. A paper strip saturated with
a solution of starch paste and potassium iodide, KI, turns
blue when exposed to the action of ozone.
REM. This test is the one employed to determine the presence and
amount of ozone in the atmosphere ; but it is not reliable, since some of
the oxides of nitrogen, which also exist in the atmosphere, affect the paper
similarly.
2. Its odor, resembling dilute chlorine, betrays ozone
when present in considerable quantities.
3. Metallic mercury, Hg, when dropped into a flask
containing ozone, immediately tarnishes.
OZCXNE. 33
SUMMARY OF STUDENT'S WORK IN O. AND OZONE.
1. Make the experiments as indicated.
2. Make Tests 1 and 2, Art. 29 ; also test a flask of common air by 2.
3. Make Tests 1, 2, and 3, Art. 33 ; also fit a delivery tube to the
florence flask used in Exp. 21 r, and direct the jet of ozone against a
globule of Hg in the bottom of a test-tube. What result ?
4. Allow the jet of ozone to pass into a test-tube containing a solution
of starch paste and KI. What occurs ?
5. Read R. and S., Vol. 1., p. 194, et seq., for a more complete discus-
sion of ozone.
6. Read Huxley's Elementary Lessons in Physiology on the arterialization
of the blood.
7. The manganese bead (Art. 316) will furnish the student excellent
practice in the use of the oxidizing and reducing flames.
8. Will ozone give the oxygen test with the glowing match ? How
can you distinguish between ozone and oxygen ?
CHAPTER II.
HYDROGEN. — ITS OCCURRENCE, ETC. — WATER. —
HYDROGEN DIOXIDE.
HYDROGEN.
SYMBOL H'. — ATOMIC WEIGHT, 1; SPECIFIC GRAVITY, 0.0692.
34. Occurrence. — Hydrogen is found, nearly always,
combined with other substances. It occurs free, however,
in very tsraall quantities in certain volcanic gases, and
absorbed in meteorites.
It occurs combined with oxygen in the form of water,
of which it constitutes 11.11 per cent by
weight.
It is a constituent of ammonia, coal gas,
marsh gas, and of nearly all organic sub-
stances.
35. Preparation. — EXP. 22 T. Use the
apparatus shown in Fig. 3. Add 1 part by
weight of pure sulphuric acid, H2SO4, to 20
parts distilled water ; then open the stop-cocks
S and S'. Pour the acidulated water into the
tube B until it issues from the tubes O and H.
Then close the stop-cocks, and fill B up to the
bulb. Connect the platinum wire Z, which is
melted through the tube H, and terminates in
a platinum strip, with the zinc pole of a Grove's
FIG. 3. battery consisting of five or six cells. Also
HYDROGEN.
35
connect the platinum wire P (which is like Z in every respect)
to the platinum pole of the battery. Hydrogen collects in tube
H, and oxygen in tube O. The hydrogen in tube H may be
tested by slightly opening the stop-cock S', and igniting.
Hydrogen burns with a very hot flame, although it emits but
little light.
QUERIES.
oxygen ?
In which tube is the volume of gas greater 7 How test the
EXP. 23 T. Make an amalgam by rubbing, in a porcelain
mortar, one-half gram metallic sodium, or potassium, together
with 5g mercury.
Fill a jar with water, and
arrange as in Fig. 4. Place
the amalgam in a wire gauze
cage, and insert under the
mouth of the jar. Hydrogen
is liberated, which rises, and
fills the jar.
Test by carefully raising
the jar, mouth downwards, FIG. 4.
and plunging a lighted taper
upward into the jar. The taper is extinguished, but the hydro-
gen burns at the mouth of the jar. The taper may be relighted
in this flame.
NOTE. This experiment usually ends with a slight but not dangerous
explosion.
QUERIES. What becomes of the mercury of the amalgam after being
dipped into the water ? Drop a piece, not larger than a pea, of metallic
sodium or potassium into a dish of warm water. What results 1 Do
you now see the reason for amalgamating the K or Na 1 What is the
reason 7 Is the water alkaline ?
NOTE. Alkalies turn a strip of red litmus paper, blue. Acids turn blue
litmus paper, red.
36 HYDROGEN.
Hydrogen is best prepared in a pure state by the de-
composition of water. It may be prepared in many other
ways ; but it then contains impurities from which it is
difficult to free it.
The action of potassium on water is expressed by the
equation
K + H2O == KOH + H.
Now let us inquire particularly as to the meaning of an
equation. Primarily, it means that potassium and water
give a substance called potassium hydroxide (KOH) and
hydrogen. It will be seen that the sign + is read and, and
the sign = is read give. But the equation means more
than this. It tells us the exact proportions in which the
substances act. For each one of the symbols stands for
a certain proportion of the element corresponding to its
atomic weight. In the above, 39 parts of potassium act
upon 18 parts of water (made up of 2 x 1 parts of hydro-
gen and 16 parts of oxygen), and give the compound
potassium hydroxide (made up of 39 parts of potassium,
16 parts of oxygen, and 1 part of hydrogen), and 1 part
of hydrogen. These relations are maintained whenever
potassium acts upon water. From the use of a given
amount of potassium, provided there be enough water,
we get a definite amount of hydrogen.
PROBLEM. How much hydrogen will be formed if 100s of potassium
were allowed to act upon water in such a way as to prevent the burning
of the hydrogen1? How much potassium hydroxide will be formed"'
How much water will be decomposed ?
SUG. Teacher will give a number of other similar problems, calling
attention to the fact that instead of saying parts we may say grams,
ounces, pounds, tons, or whatever unit of weight we may choose to take.
Student review the equations inclosed in parentheses, and explain.
HYDKOGEN. 37
EXP. 24 T. Place a quantity of granulated zinc in the gener-
ating flask A, Fig. 5. Through the funnel tube B introduce a
liberal quantity of dilute sulphuric acid, H2SO4, consisting of
one part acid by weight to four of water. Allow the gas to
escape through the delivery tube D for some time, to free the
apparatus from air. Then collect in gas bags, or in the gas
receiver, or in jars over the pneumatic troifgh. The reaction is
represented by the equation
Zn + H2SO4 = ZnSOt -f 2 H.
We take no account of the water added, as it serves merely
as a solvent for the zinc sulphate, ZnSO4, as fast as formed.
FIG. 5.
Hydrogen is prepared in large quantities, when absolute
purity is not especially requisite^ by allowing dilute acids
(HC1, H2SO4) to act on certain metals, such as iron and
zinc.
NOTE. Hydrogen made in this way may contain sulphuretted hydro-
gen and other impurities, which, for the most part, are destroyed by pass-
ing the gas through a solution of potassium permanganate. The gas may
be dried by passing it through sulphuric acid or calcium chloride, or both.
(Student should arrange an apparatus for making and purifying hydro-
gen.)
Having a quantity of hydrogen stored, teacher and
students make the following experiments: —
38 HYDROGEK.
36. Properties. — EXP. 25 r. Fill collodion balloons to
illustrate the lightness of hydrogen. Allow one or two of
them to rise to the ceiling, and remain as long as they will.
Even though they do not leak, they will, nevertheless, sink to
the floor after a time. Why?
EXP. 26 TP. Make hydrogen soap-bubbles, which will burn
when touched with the flame of a taper.
QUERIES. Are these bubbles heavier or lighter than air ? How can
you tell the same of other gases 1
EXP. 27 TP. Discharge the Irydrogen pistol, illustrating the
explosiveness of hydrogen and oxygen.
QUERY. Should you fill the pistol full of H, could you discharge it 1
Why?
EXP. 28 TP. Produce singing flame. To succeed well with
this, fit a- long, straight jet into a generating flask containing
metallic zinc and dilute sulphuric acid. When the gas is coming
off freely, light the jet. Hold glass tubes of various lengths and
bores, down over the burning jet. In this way different tones
may be produced.
QUERY. What produces the tones ?
EXP. 29 TP. Fill a bell jar with hydrogen, by holding the
mouth of the jar downward, and allowing the hydrogen to flow
up into the jar. Now reach up into the jar with an inverted
dipper (ordinary) . Keeping the dipper bottom side up, draw it
slowly downward out of the jar, and remove it some distance
away ; then bring a lighted taper under the dipper. What
ensues ? Explain.
Pure hydrogen is an odorless, tasteless, invisible gas,
which was discovered and described by Cavendish in 1766.
Its specific gravity (air — 1) is 0.0692. Hydrogen is
the lightest substance known: I1 at 0° C. and 760mm
pressure, weighs 0.0896g.
HYD110GEX. 89
PROB. How many grams II, at 0° and 760mm, will a bell jar of 201
capacity hold 1
M. Pictet claims to have liquefied hydrogen at —140° C.,
and 650 atmospheres ; but the so-claimed liquid gave no
meniscus in the tube in which he was endeavoring to
condense the gas.
SUG. Student, half fill a test-tube with water, and note the meniscus at
the upper level of the liquid.
Hydrogen is highly combustible, burning with a very
hot but slightly luminous flame, and, when mixed with
considerable quantities of air or free oxygen, explodes
with violence.
The metal palladium absorbs hydrogen in large quan-
tities at moderate temperatures. Platinum and iron also
absorb it, but in much smaller proportion than palladium.
It seems as if the hydrogen forms an alloy with them,
acting very much like a metal itself. When a jet of
hydrogen is directed against a piece of spongy platinum,
at ordinary temperatures, so much heat is evolved as to
cause the jet to ignite. Hydrogen is slightly soluble in
water. It is not directly poisonous, but produces a weak-
ening and sharpening effect on the voice, when inhaled.
It is very diffusible, and is apt to contain atmospheric
air. The extreme lightness of hydrogen caused it to be
used for filling balloons ; but, owing to its great diffu-
sibility and the expense of its manufacture, it has been
superseded by coal-gas. One gram of hydrogen, when
burned, produces enough heat to raise the temperature
of 34,462g of water through one degree. Hence its ca-
lorific power is said to be equal to 34,462 thermal units,
— the thermal unit or Calorie being the amount of heat
necessary to raise the temperature of one gram of water
one degree Centigrade.
40 HYDROGEN AND OXYGEN COMPOUNDS.
37. Test. — Hydrogen may be recognized by its flame
and behavior, as in the preceding experiments.
HYDROGEN AND OXYGEN COMPOUNDS.
38. Hydrogen and Oxygen form two chemical com-
pounds ; viz. : —
1. Water, H2O ; and
2. Hydrogen dioxide, H2O£.
WATER, H2O.
39. Occurrence. — With water we are all well acquainted.
It occurs everywhere, — in streams, lakes, and the bound-
less ocean. It exists in the
atmosphere as vapors, fogs,
and clouds, and is precipi-
tated upon the earth as dew,
rain, hail, and snow. It is
absorbed by the soil and
rocks, while in crystalline
structures it enters into
closer combination as water
__- of crystallization.
40. Preparation. — It is
FIG. 6.
not necessary to prepare
water chemically, owing to its great abundance every-
where, but for the sake of illustration use the apparatus
shown in Fig. 6.
G is a hydrogen generator.
B is a drying bulb, containing granulated calcium chloride.
II is a bell jar. The hydrogen jet burning in this jar unites with the
oxygen of the air, producing water, which soon collects, and falls down
in drops.
HYDROGEN AND OXYGEN COMPOUNDS. 41
SCG. Student, write the equation. Also write a description of the
whole apparatus and manipulations.
41. Question. — What is the chemical composition of water,
and what its formula?
We may determine the composition of water, first by
analysis, and then, if possible, by synthesis. The experi-
ment described on page 34 showed that when water is
decomposed by the electric current, it yields only hydro-
gen and oxygen, and these in the proportion of 2 vol-
umes of hydrogen to 1 of oxygen. Knowing the relative
weights of the gases, we see that they are obtained from
water in the proportion of 1 part by weight of hydrogen
to 8 of oxygen, or 2 of hydrogen to 16 of oxygen.
SUG. Student, show that this statement is correct.
To prove that hydrogen and oxygen alone are necessary
to form water, and that they are present in the proportions
found by analysis, we may cause the two gases to unite as
follows : —
EXP. 30 TP. The apparatus shown in Fig. 7 is called Ure's
Eudiometer. The graduated limb and part of the plain limb
are filled with mercury ; then, by means of a
curved tube, 10 divisions of the graduated
limb are filled with pure oxygen ; then fill
say 25 more with pure hydrogen. An elec-
tric spark is now passed through the wires
attached to the graduated limb, while the
thumb is held firmly over the plain limb.
20 divisions of hydrogen will unite with 10
divisions of oxygen; i.e., 2 of Irydrogen to
1 of ox}rgen.
QUERY. After passing the spark, where is the water to be seen ?
N.B. Before passing the spark, see that the plain limb is not entire';
full of mercury, and hold the thumb as firmly as possible.
42
HYDROGEN AND OXYGEN COMPOUNDS.
SUG. Student, examine this apparatus, and write a full description of
it and the experiment.
We see, thus, that the analysis and synthesis of water
both lead us to the conclusion that in it hydrogen and
oxygen are united in the proportions above stated, and
these proportions are expressed in the formula H2O, the
full significance of which cannot be explained at this
stage. For the present, suffice it to say that formula)
express primarily the composition of bodies by weight.
Hereafter, we shall see that they also have to deal with
the volumes of bodies when in the form of a gas.
PROB. How many grams of O in 100s of H20 ? How many of H 1
42. The Oxy-Hydrogen Blow-Pipe. — Small laborato-
ries will not be likely to contain this apparatus; but,
owing to its great value and the frequent references
made to it, the student should become acquainted with
it.
The oxygen and hydrogen holders are not shown in this
cut (see App.). They may be provided 'with safety-
valves, to prevent the flow of the gas from one to the
other.
J is a jet containing a jet within,
a space being left between the inner
jet and the outer one for hydrogen to
pass through.
H is a stop-cock to admit hydro-
gen into this space. 0 is a stop-cock
to admit oxygen into the inner jet,
which is not quite so long as the outer
jet. By this arrangement the two gases
are thoroughly mixed upon issuing into
the air.
C is an adjustable cup for holding
a piece of chalk in the flame, when
the design is to produce the brilliant
FIG. 8. calcium light
HYDEOGEN AND OXYGEN COMPOUNDS. 43
The heat of the flame of this blow-pipe is intense
enough to melt most of the refractory metals.
The calcium light is equaled only by the electric light.
43. Properties of Water. — Water is an almost univer-
sal solvent; consequently, pure water does not occur in
nature. Snow and ice waters are nearly pure, but they
still contain dust, and various gases found in the air.
Lake Superior water is also very nearly pure, since the
bed of the lake is composed of the old Azoic rocks which
are but slightly soluble, and the lake is fed with ice, snow,
and rain. Sea water contains nearly every known sub-
stance in solution.
Water is at its maximum density at +4° Centigrade.
When the temperature passes either above or below this
point, water expands. This is a most fortunate provision,
as otherwise, ice would be heavier than water and would
sink to the bottom ; thus, many of our lakes and rivers
might be frozen solid to their beds, and the summer sun.
would not suffice to thaw them. Aquatic plants and
animals could not exist, and our temperate zones would
become uninhabitable.
QUERY. Why does the pail burst when the water freezes in it 7
EXP. 31 OP. Place a thermometer through an opening in the
ice of a frozen lake. At any depth it will read nearly +4° C.
QUERY. What deductions may be derived from this experiment ?
THE LATENT HEAT OF WATER is 79 CALORIES OR THERMAL
k UNITS.
Illustrate this statement, thus : —
EXP. 32 OP. Mix lk of ice at 0° C. with lk of water at 79°
C. The ice will melt, and the temperature of the 2k of water
44 HYDROGEN AND OXYGEN COMPOUNDS.
will be 0° C. Hence we see that the 79 thermal units contained
in the kilogram of water have disappeared while melting the ice,
or, in other words, have become latent. When water freezes,
it gives off its latent heat.
QUERY. What effect, upon the temperature of a room, would be pro-
duced by a tank of freezing water.
THE LATENT HEAT OF STEAM is 536 THERMAL UNITS.
To illustrate this, proceed thus : —
EXP. 33 OP. Into 5.36k of water at 0° C. pass steam at 100°
C. until the water boils. You will then have 6.36k of water at
100° C. Now, since lk of steam has parted with sufficient latent
heat, while condensing to water (of the same temperature, i.e.,
100°), to raise 5.36k of water 100°, or 536k 1°, we have measured
its latent heat, which is 536 thermal units.
NOTE. Experiments 32 p and 33 p involve quite large experimental
errors.
When steam condenses to water it gives off all its latent
heat ; hence its great usefulness for heating dwellings, etc.
DRINKING -WATER.
Drinking-water is apt to contain many impurities,
organic and inorganic, some of which are believed to be
very deleterious to health, frequently leading to various
forms of disease, such as typhoid fever, etc.
QUERY. How does drinking-water become contaminated with impuri-
ties ?
Let the student make the following tests upon drink-
ing-water obtained from his own well, or from the usual
source of water for drinking purposes.
HYDROGEN AND OXYGEN COMPOUNDS. 45
TESTS FOR IMPURITIES IN DRINKING-WATER.
EXP. 34 P. For Organic Impurities. — Fill a tall glass jar
with the water to be tested. Add a few drops of sulphuric
acid, H2SO4 ; then add a solution of potassium permanganate,
K2Mn2O8, until the whole assumes a deep purplish tint. Stand
in a warm place for one hour. If organic impurities are
present, the solution will be decolorized.
Another Test. — When much organic matter is present. —
Fill a tightly-stoppered bottle nearly full of the water to be
tested. Set in a warm place for several da}~s. An offensive
odor indicates organic impurities. Such impure water, it is dan-
gerous to drink. A good charcoal and gravel filter will remove
organic impurities if only a small amount be present.
SUG. Teacher explain the construction of a filter, and how to take
care of it.
TEST FOR AMMONIA.
EXP. 35 P. Distil the water in perfectly clean glass apparatus
(after dissolving a small quantity of sodium carbonate, Na2CO3,
in the water to be tested) .
Collect the distillate in tall glass jars in volumes of 50CC
each, numbering them successively 1, 2, 3, 4, etc.
Add about 2CC of Nessler's Test Solution (see App.) to
each of these jars. If ammonia be present in any or all of
them, such as contain it will be tinged brownish-yellow.
N.B. Drinking -water containing much ammonia is unfit to drink,
since the presence of ammonia indicates that the water of the well has
percolated through decaying vegetable or animal substances.
TEST FOR CHLORINE OR CHLORIDES.
EXP. 36 P. Concentrate 50CC of water to be tested to 25.
Acidulate with nitric acid ; then add a few drops of a solution
of silver nitrate, AgNO3. If a white precipitate is made which
46 HYDROGEN AND OXYGEN COMPOUNDS.
is soluble in ammonia, NH4OH, and insoluble in nitric acid,
HNO3, chlorine is present.
The presence of much chlorine is to be looked upon
with suspicion (as sewage water always contains chlorine
in considerable quantities), unless in the vicinity of salt
wells or of the ocean.
TEST FOR NITRITES.
EXP. 37 P. Acidify the water to be tested with acetic acid,
H(C2H3O2) ; then distil in a clean glass retort, allowing the first
part of the distillate to drop into a solution of starch paste and
potassium iodide, KI, to which a few drops of sulphuric acid,
H2SO4, have been added. Nitrites, if present, will turn this
solution blue, owing to the liberation of iodine.
NOTE. Nitrites will also bleach a solution of potassium permanganate,
KMn04, when acidulated with sulphuric acid, H2S04 ; but this test is not
reliable, since organic matter acts in the same way.
The presence of nitrites is an indication of sewage,
especially when chlorides and ammonia are present.
Too much stress cannot be laid on the danger of drink-
ing water contaminated with sewage. Fevers and pesti-
lence may follow its use.
TEST FOR HYDROGEN SULPHIDE, H2S.
EXP. 38 P. Acidify, with sulphuric acid, H2SO4, about I1 of
the water to be tested. Place it in a stoppered flask holding
say 21. Suspend above the liquid a strip of bibulous paper
moistened with lead acetate, Pb (C2H3O2)2. Cork tightly, and
set in a warm place for several hours. Hydrogen sulphide, if
present, will blacken the paper.
HYDROGEN AND OXYGEN COMPOUNDS. 47
TEST FOR HARDNESS.
EXP. 39 P. Employ Clark's Soap Test thus : Place 70CC of
the water to be tested in a stoppered glass flask. Add lcc
of Clark's Soap Solution ; then shake thoroughly. If a perma-
nent lather be not formed, again add lcc of the soap solution,
and shake as before, and thus proceed until a permanent
lather remains, for three minutes, unbroken over the surface of
the water. (See App. for Clark's soap solution.)
The number of centimetres soap solution added will be
equal to the number of degrees of hardness, or to the
number of grains plus 'one of hardness, per imperial gallon.
Hardness is usually caused by the presence of calcium
and magnesium carbonates.
NOTE. Hardness and hydrogen sulphide do not necessarily impair the
qualities of drinking-water ; on the contrary, they often serve useful pur-
poses.
HYDROGEN DIOXIDE, H2O2.
44. Preparation. — Hydrogen dioxide does not occur
in nature in quantity, though it is present in small
amounts in the air, and in rain and snow. It may be
prepared chemically in several ways, of which we give only
one, the best way.
EXP. 40 P. Treat pulverized barium dioxide with dilute sul-
phuric acid (5 parts water to 1 part acid) in a beaker. Stir
thoroughly to bring all the barium dioxide in contact with the
acid. A white precipitate, barium sulphate, BaSO4, will settle
to the bottom upon standing, and the clear fluid will contain
the hydrogen dioxide. This separation can be effected more
quickly by filtering. The clear fluid which comes through con-
tains the hydrogen dioxide in dilute solution. The reaction
may be expressed thus : BaO, + H2S04 = BaS.O, + H202-
48 HYDROGEN AND OXYGEN COMPOUNDS.
PROS. How many grams H202 may be obtained from 10s Ba02?
This dilute solution of hydrogen dioxide may be con-
centrated by allowing it to stand in a beaker placed over
strong sulphuric acid in the vacuum of an air-pump ; but,
after and during concentration, it should be kept at a low
temperature.
45. Properties. — Hydrogen dioxide is a very unstable
liquid, slowly separating into water and oxygen at low
temperatures (student, write the equation), rapidly decom-
posing at +20°, and exploding with violence at 100°.
It is syrupy, transparent, and colorless, possessing a very
nauseating and stringent taste.
Its specific gravity is 1.452, and it has not been frozen.
Aqueous hydrogen dioxide is sold commercially for
bleaching old engravings arid paintings. It is also used
to change dark hair to lighter shades, which is a danger-
ous practice, since it is an active poison when brought
upon the skin, often producing white blisters which finally
become very painful.
46. Test for Hydrogen Dioxide. — Acidulate a small
quantity of a solution of hydrogen dioxide (or the liquid
to be tested) with two or three drops of sulphuric acid,
H2SO4, in a test-tube. Add a small quantity of ether,
(C2H5)2O, also five or six drops of potassium chromate,
K2CrO4; shake well. Hydrogen dioxide, when present,
turns the whole to a splendid blue color. On standing,
the ether absorbs this color, and separates out in a blue
layer.
HYDROGEN AND OXYGEN COMPOUNDS. 49
SUMMARY OF STUDENT'S WORK IN H, H2O, AND H2O2.
1. Make the experiments as indicated.
2. If the laboratory contain an oxy-hydrogen blow-pipe, teacher and
students should use it in making the calcium light, fusing bits of metals,
as Fe, Au, Pt, etc.
3. Art. 43. Draw up reports giving results of the Exp. 34 p- 39 p.
This is work sufficient for a whole week.
4. Albuminoids may be detected thus : Add solid KOH to the water
until strongly alkaline, and boil a short time. Now pour into a retort and
add K2Mn208, and distil, collecting and testing the first portions of dis-
tillate as in Exp. 35 P, since albuminoids thus treated yield ammonia.
5. Read Wanklyn's Water Analysis. In case it is desirable to de-
termine the amounts of ammonia, etc., present in drinking-water, complete
directions are to be found in this work. One should hesitate to pronounce
upon the potableness of drinking-water without first making quantitative
determinations.
6. PROB. The imperial gallon contains 70,000 gr. of distilled water ;
the U. S. gallon contains 53,328.88 gr. How many grains of hardness
per U. S. gallon does the sample of water that you have analyzed contain 1
How many milligrams per litre does it contain (lcc Clark's soap solution
precipitates lms of hardness)?
7. Try to remove, by boiling, the hardness from a sample of water. In
case you succeed, the hardness is said to be temporary ; and it is due to the
presence of calcium carbonate, CaC03, and perhaps magnesium carbonate,
MgCO3. Should you not succeed in thus removing it, the hardness is
called permanent, and probably consists of the sulphates of calcium and
magnesium.
QUERY. How can you determine if both permanent and temporary
hardness be present ?
8. PROB. How many grams 0 can be obtained by decomposing 100CC
of water (lcc=r Is)? How many grams H ?
9. PROB. In the equations enclosed by parentheses, assume 10s of
the first substance, and ascertain how many grams will be required of the
remaining substances.
10. Art. 44. It is not necessary to condense H202 in vacuo, unless a
concentrated solution is required.
Student test H202, as in Art. 46. A dilute solution of H202 will answer
well for this purpose. Use K2Cr207, also, in place of K2Cr04. Do you
obtain the same color as before 1
CHAPTER III.
NITROGEN. — ITS OCCURRENCE, ETC. — AMMONIA. — OXIDES
OF NITROGEN. — THE NITROGEN ACIDS. — HYDROXYLA-
MINE.
NITROG-EN.
SYMBOL, N'". — ATOMIC WEIGHT, 14; SP. GRAY., 0.9713.
47. Occurrence. — Nitrogen occurs free in the atmos-
phere, of which it constitutes nearly four-fifths by volume,
or 77 per cent by weight. It also occurs in many chemical
compounds, such as potassium nitrate, KNO3; sodium
nitrate, NaNO3; ammonia, NH3; and in many organic
substances, particularly those of animal origin.
48. Preparation. — EXP. 41 T. Place in an iron sand-
bath about 2g of phosphorus ; ignite the phosphorus, and float
the sand-bath on the water in a pneumatic trough. Immedi-
ately place over the burning phosphorus a bell-jar of about
41 capacity, allowing the mouth of the jar to be under water,
so that no outside air can enter. The phosphorus enters into
combination with the oxygen of the air contained within the
jar, forming dense white fumes of phosphorus pentoxide,
P2O5, and perhaps of the trioxide, P2O3. In a short time
these fumes settle, and are dissolved by the water, leaving the
nitrogen nearly pure.
Nitrogen may be prepared in many ways, but the
method above indicated is a cheap and convenient one
for laboratory use. -
NITROGEN. 51
49. Properties of Nitrogen. — Exr. 42 p. Bend a small
glass tube in the shape of a letter V, and draw out one extrem-
ity into a jet. Now insert the plain end of the tube into the
jar of nitrogen obtained in Exp. 41, and press the jar down
into the water of the pneumatic trough until the nitrogen issues
through the jet. Try to ignite the gas. Does it burn?
EXP. 43 P. Fill with nitrogen a large test-tube. Insert a
glowing match ; a burning match ; a lighted taper. What results ?
EXP. 44 p. Try to fire the hydrogen pistol when filled with
a mixture of nitrogen and common air. What occurs?
Nitrogen is a gaseous element, and, like hydrogen and
oxygen, none of its physical properties render it percep-
tible to sight, taste, or smell. Its specific gravity is 0.971 ;
and I1 at 0° C. and 760mm pressure weighs 1.256^.
Chemically considered, it is not an active element, as
shown by the apathy which it exhibits in entering into
combination with other elements.
Indirectly, however, it unites with hydrogen, oxygen,
and carbon to form important chemical compounds.
Its greatest value in nature is due to its mildness, and
the remarkable persistency with which it remains in a free
state. It thus serves to dilute the oxygen of the atmos-
phere, which is simply a mechanical mixture of these gases,
consisting of 23.1 parts oxygen and 76.9 parts nitrogen,
by weight. As one would infer, it has no poisonous prop-
erties, neither will it burn nor support combustion.
Nitrogen is but slightly soluble in water, and has been
condensed to a liquid at — 146° C. and under a pressure
of 33 atmospheres.
50. Test for Nitrogen. — Owing to its passive nature,
nitrogen does not give any reaction whereby it may be
52 NITROGEN AND HYDliOGEN.
readily detected when present in small quantities. Larger
amounts are indirectly tested by the negative results
obtained.
NITROGEN AND HYDROGEN.
51. Ammonia. — Nitrogen and hydrogen unite to form
an important compound, viz. : —
AMMONIA, NH3.
52. Occurrence. — Ammonia occurs free in the atmos-
phere, being produced by the decay of organic matter
containing nitrogen. It is also found dissolved in small
quantities in rain water and many surface waters. Its
compounds, such as ammonium chloride, NH4C1, and am-
monium carbonate occur but sparingly in nature, although
they are common articles of commerce, obtained by arti-
ficial processes. Ammonia solution, or aqua ammoniae, is
also a staple article of commerce.
53. Preparation. — EXP. 45 p. Place in one hand a small
quantity of dry quicklime, CaO, and in the other an equal
bulk of pulverized ammonium chloride, NH4C1. Note that
neither substance emits an odor. Now rub them together
between the palms of the hands, and carefully smell the invisi-
ble gas given off. It is ammonia.
EXP. 46 P. To a solution of ammonium chloride in a test-
tube add a few drops of potassium hydroxide, KOH. Warm
gently, and note the fumes. Do you again obtain ammonia?
Also try in the same way a solution of ammonium nitrate,
NH4NO3. What result? Moisten a glass stirring-rod with
hydrochloric acid, HC1, and hold it in the escaping vapors ;
notice the white fumes that are formed. Try the same with
NITROGEN AND HYDROGEN.
53
nitric acid, HNO3. What takes place is indicated in the two
following equations : —
1. NH3+HC1 =NH4C1.
2. NH3+ HN03= NH4N03.
Sue. Explain these equations. Do you obtain the same substances
with which you commenced ?
EXP. 47 T. Thoroughly mix two parts, by weight, of finely-
pulverized, dry ammonium chloride, and one part of dry quick-
lime. Quickly place the mix-
ture in the generating-flask F
(Fig. 9), and then add a thick
layer of dry quicklime, which
will serve to dry the ammonia
as it rises through it. Insert
the cork containing the bent
tube, and gently heat the flask.
Ammonia will collect in the
bottle B. 5
This method is used when
dry ammonia gas is required.
What occurs in the flask is
indicated by the equation,—
CaO + 2 NH4C1 = 2 NH3
Fn 9.
H2O + CaCl2.
QUERIES. CaCl2 is a substance called calcium chloride, and is a solid.
Where is it to be found after the reaction ? What becomes of the water 1
Would passing the gas through a long tube filled with quicklime tend to
insure the dry ness of the ammonia ? From the position in which the bottle
B is held, should you judge ammonia to be lighter, or heavier, than air 1
:
EXP. 48 T. Prepare thick pastes (with water) of the same
substances used in the last experiment, employing the same
Toportions of the dry substances before moistening. Arrange
an apparatus similar to that shown in Fig. 10. The first wash-
bottle, A, acts as a safety-valve to prevent water from return-
54
NITHOGEN AND HYDKOGEN.
ing into F ; also to prevent explosions. Notice that the centre
tube alone dips beneath the water. The bottle B contains
cold water, and serves as a condenser. In this bottle the entry
and centre tubes extend below the surface of the water. C is
also a bottle containing cold water. It is best to place B
and C in vessels, and to surround them with a freezing mixture
of snow, or pounded ice and salt. Now pour these pastes
as rapidly as possible into F, shake quickly, and connect with
FIG. 10.
the wash-bottles. Apply heat to F, and boil for some time. An
aqueous solution of ammonia will be found in B and C.
SUG. Explain the reasons for arranging A in the manner described.
After the experiment, note that the contents of B and C differ in no way
from ordinary aqua ammoniae.
The last experiment indicates the general process em-
ployed in manufacturing commercial aqua ammoniae.
Other methods of preparing ammonia are as follows : —
1. Ammonia is obtained in small quantities by mixing
nitrogen and hydrogen in a eudiometer, and by passing
for some time a silent electric discharge.
QUERY. How does this explain the production, at certain times, of
ammonia in the atmosphere ?
NITROGEN AND HYDROGEN. 55
2. It is produced by allowing heaps of compost and
urine to decompose.
QUERY. How do you explain the presence of ammonia in stables ?
EXP. 49 P. Heat in a test-tube a small ball of hair, or
wool, or a few hoof-clippings. What is given off?
3. Ammonia may be had by the dry distillation of such
nitrogenous bodies as hair, hoofs, hides, and horns. It was
formerly prepared in this way, and thus received the name
spirits of hartshorn.
4. In the distillation of coal to make illuminating gas,
ammonia is formed as a by-product. In this case the
nitrogen and part of the hydrogen contained in the coal
are driven off, combined as free ammonia and ammonia
compounds. These arnmoniacal products are led into
water containing hydrochloric or sulphuric acids; from
the compounds thus formed we obtain the ammonia of
commerce.
54. Properties. — EXP. 50 p. Place in a generating-flask
a concentrated solution of ammonia ; pass through this solution
a current of oxygen gas. The escaping mixture of ammonia
and oxygen will burn, at the mouth of the flask, with a yellowish
flame. Under ordinary conditions, ammonia does not burn.
SUG. Student, try to light a jet of NH3.
EXP. 51 P. Heat to bright redness a long spiral coil of
platinum wire, and quickly introduce it into the mouth of a
common reagent bottle containing a strong solution of am-
monia. The wire will continue to glow while the ammonia is
decomposed, thus : —
2NH3 + 3 O = NH4NO2 + H_9O.
QUERIES. How do you account for this phenomenon ? Whence comes
the O indicated in the above equation ?
56 NITROGEN AND HYDROGEN.
The composition of ammonia gas may be determined by
introducing the dry vapor into the graduated limb of Ure's
eudiometer, and passing a succession of electric sparks,
when the volume of the enclosed gas is doubled. This
may now be proven, by introducing oxygen and exploding,
to consist of one volume of nitrogen and three volumes of
hydrogen.
We may here learn a useful fact ; viz., that the formula
NHs, in addition to its other significations, also represents
two volumes of ammonia in the form of a gas; and the same
is true of all formulas representing gases.
The name ammonia originated from the fact that the
gas was first prepared from sal-ammoniac, NH4C1, a sub-
stance formerly confounded with the salt, NaCl, produced
near the ruins of the temple of Jupiter Ammon, in Lybia.
Ammonia is an invisible gas possessing a powerful, irri-
tating odor, and intensely alkaline properties.
SUG. Try the effect of ammonia upon a strip of moist, red litmus paper.
This gas is very soluble in water, lcc of water at 0° C.
absorbing 1148CC of ammonia ; and I1 at 0° and 760mm pres-
sure weighs 0.762g. It can easily be condensed to a liquid
under a pressure of 7 atmospheres at 4- 15.5° C. ; and this
liquid, on being cooled to — 75° C., becomes a transparent
solid.
In passing from a liquid to a gaseous state, gases always
absorb a large amount of heat. M. Carre* has taken advan-
tage of this fact in constructing an ice machine. (Fig. 11.)
B is a boiler containing a strong solution of ammonia. C is
a condenser with an air-tight space between its double walls,
the whole being surrounded by the non-conducting covering, H.
B is gradually warmed over a slow fire, while C is placed in
a vessel of cold water. The ammonia of the aqueous solution
NITROGEN AND HYDROGEN.
57
FIG. 11.
in B is driven into the air-tight space in C, where it is con-
densed by its own pressure. Water is now placed in C, and
B is subjected to a cold bath,
when the liquid ammonia in the
walls of C quickly evaporates,
and is absorbed by the cold
water in B. This evaporation
abstracts so much heat from
the water in C that it is soon
frozen.
EXP. 52 p. Place in a beaker
glass a dilute solution of nitric
acid, HNO3 ; now carefully add
ammonia until the solution is
neutralized so that it does not
affect litmus paper. Gently evaporate this solution to dryness,
when a crystalline salt is obtained.
Ammonia unites with acids to form salts, and is known
as the volatile alkali. This action with acids may be
illustrated by the equations: —
NH3 -f HC1 = NH4C1.
NH3 + HNO3 = NH4NO3.
It will be seen that the ammonia, NH8, is added directly
to the acid. The compounds thus formed are called
ammonium compounds, the group NH4 contained in them
being known as ammonium.
Since ammonia neutralizes an acid, it is used in cases of
accidents when acids are spilled upon the clothes or flesh.
But should the acid be received in the face, it is best to
wash it off quickly with much water, then with a weak
solution of ammonia, and finally, without rubbing, to cover
the injured parts with sweet oil.
58 NITROGEN AND OXYGEN.
Ammonia produces a stimulating effect upon the human
system when inhaled, and is often employed in cases of
fainting, or where over-doses of chloroform, laughing gas,
ether, etc., have been taken. It also neutralizes the effects
of such poisonous or irritating gases as chlorine, sulphur
dioxide, and nitrogen tetroxide.
55. Tests for Ammonia, NH3. — 1. When present in
very small quantities, as in drinking-water, ammonia is
best detected by means of Nessler's test solution. (See
App.)
2. When present in considerable quantities, add to the
solution to be tested potassium hydroxide, KOH, and
warm gently. Ammonia, if present, is driven off, and may
be recognized as follows : —
(a) By its pungent smell.
(6) By turning moistened red litmus paper blue.
(c) A warm glass rod previously moistened in hydro-
chloric acid, HC1, is coated white by ammonia gas. Char-
acteristic white fumes (NH4C1) are also produced when
much ammonia is present.
NITROGEN AND OXYGEN.
56. Nitrogen indirectly unites with oxygen to form five
oxides or compounds, viz. : —
N2O, Nitrogen Monoxide or Nitrous Oxide.
NO, (or N2O2) Nitrogen Dioxide or Nitric Oxide.
N2O3, Nitrogen Trioxide or Nitrous Anhydride.
NO2 (or N2O4), Nitrogen Tetroxide.
N2O5, Nitrogen Pentoxide or Nitric Anhydride.
NITROGEN AND OXYGEN.
59
NITROGEN MONOXIDE, N2O.
57. Occurrence. — This substance is a gas, and never
occurs free in nature. It is often known by the common
name of nitrous oxide.
58. Preparation. — EXP. 53 P. Place in a test-tube a
small quantity of ammonium nitrate, NH4NO3, and heat
gently in the Bunsen flame. Note the sweetish odor. The
gas thus obtained is nitrogen monoxide. Insert a glowing
match, as in testing for oxygen ; also try a burning match.
QUERIES. How does this gas behave, in comparison with oxygen ?
How can you distinguish it from oxygen ?
FeS04
FIG. 12.
H20
EXP. 54 T. Place 20g ammonium nitrate, NH4NO3, in a
generating flask, and connect with three wash-bottles, as
shown in Fig. 12. The thistle-top tube contains a small quan-
tity of mercury, and will serve as a very efficient safety-valve.
A moderately strong heat will serve to decompose the contents
of the flask, thus : —
NH4NO3 = N2O + 2 H,O.
But the nitrogen monoxide may contain impurities such as
nitric oxide, NO, and chlorine. It is accordingly washed through
60 NITROGEN AND OXYGEN.
a solution of ferrous sulphate, FeSO4, which is placed in the
first wash-bottle, to remove the nitric oxide. The second wash-
bottle contains a solution of potassium hydroxide, to remove the
chlorine ; while the third bottle contains water. The contents
of the bottles must be warm, since nitrous oxide is somewhat
soluble in cold water, and but slightly so in warm solutions.
The gas thus prepared is best collected in rubber gas-bags,
where it may be kept for experimental purposes in studying its
properties.
Nitrogen monoxide may be prepared by other methods ;
but the one given above is always used in its practical
preparation.
59. Properties. — EXP. 55 OP. Inhale a small quantity of
pure nitrogen monoxide, as prepared above, and note its odor,
and taste.
SUG. Student, make the same experiments with nitrogen monoxide as
with oxygen.
Nitrogen monoxide is a colorless gas possessing a pleas-
ant smell and sweetish taste, and when mixed with air, and
inhaled, produces a peculiar intoxication, while conscious-
ness remains, whence it derived its name " Laughing Gas."
When inhaled in a pure state, it affects the system thus: —
1. Intoxication and singing in the ears are experienced.
2. Insensibility follows.
8. If continued long enough, death ensues.
This gas is chiefly used for anaesthetic purposes, by den-
tists and physicians, who keep it stored under pressure in
tanks or cylinders.
It is soluble in water, 100CC of water at 0° dissolving 130CC
of nitrogen monoxide, while alcohol dissolves still greater
quantities.
This gas can be liquefied at 0° by a pressure of 30 atmos-
NITROGEN AND OXYGEN. 61
pheres, or at ordinary pressures by reducing its tempera-
ture to — 88° C. By mixing this liquid with carbon
bisulphide, CS2, and by placing the mixture in a receiver
from which the air and vapors are afterwards rapidly
exhausted, the remarkably low temperature of — 140° C.
has been reached.
Nitrogen monoxide will support combustion ; but in order
to initiate the process, some substances, as sulphur, must
be freely burning. Ignited sodium, potassium, and phos-
phorus, however, burn in it quite as briskly as in oxygen.
The specific gravity of this gas is 1.527; and I1 at 0° and
760mm weighs 1.972&.
60. Tests for Nitrogen Monoxide, N2O. — This gas
closely resembles oxygen, from which it is easily distin-
guished, first by its odor and taste, and second by its great
solubility in cold water.
NITKOGEN DIOXIDE, NO.
61. Preparation. — This oxide of nitrogen, also called
nitric oxide, does not occur free in nature.
EXP. 56 P. Place copper filings in the generating-flask A ;
then adjust the cork with the tubes B and C, as shown in
Fig. 13. Now, through the tube B, introduce into A dilute
nitric acid (sp. grav., 1.2).
At first A will be filled
with reddish - brown fumes ;
but these disappear as soon
as the air is expelled from
the apparatus, and a color-
less gas, NO, collects in G.
Note the disagreeable odor.
Allow some of the gas to
escape into the air. What
do you observe? FIG. 13.
62 NITROGEN AND OXYGEN.
Nitrogen dioxide can be prepared from nitric acid by
the action of other metals than copper, such as iron, zinc,
silver, and mercury. It might be well to know, however,
that this gas thus prepared contains impurities such as
nitrogen and nitrogen monoxide ; but these impurities are
insignificant in qualitative work. The reaction with cop-
per is expressed thus : —
3 Cu + 8HNO3 = 3 Cu (NO3)2 + 4H2O + 2NO.
62. Properties. — Nitrogen dioxide is a colorless gas,
but when brought in contact with the air, it unites with
atmospheric oxygen to form the reddish-brown fumes,
NO2, seen at the beginning of the last experiment.
Sodium, potassium, and phosphorus, when very strongly
ignited, will burn in this gas, but not so readily as in nitro-
gen monoxide, since it does not decompose as readily as
the latter gas, to supply oxygen for the purpose of com-
bustion. What takes place when bodies burn in these
oxides of nitrogen may be seen from the following equa-
tions : —
1. N2O-f-2Na = Na2O + 2N.
2. NO + 2K =K2O + N.
SUG. Try NO with a glowing match. What result ?
Nitrogen dioxide was formerly considered as an incon-
densible gas, but it became a liquid at — 11° under
104 atmospheres.
The specific gravity of this gas is 1.038 ; and I1 under
standard conditions weighs 1.343g.
QUERY. At what temperature and pressure have we given the weights
of the gases up to this time ?
63. Tests for Nitrogen Dioxide, NO. — 1. We can
distinguish this gas by the brownish-red fumes which it
gives upon escaping into the air.
NITROGEN AND OXYGEN. 63
2. When passed into a solution of ferrous sulphate, FeSO4,
the solution turns brown.
NOTE. By heating this solution, chemically pure NO may be obtained.
NITROGEN TRIOXIDE, N2O3.
64. Preparation. — This gas also does not occur in
nature. The following is the best method of preparing
it: —
EXP. 57 P. To a few grains of starch in a test-tube add
reagent nitric acid, HNO3, and gently heat in the Bunsen flame.
Dark-reddish fumes of the trioxide are given off.
EXP. 58 P. Place in a generating-flask 10s of starch, and
cover with nitric acid. Cork the flask tightly with a rubber
stopper, carrying a bent deliveiy-tube, which projects into
another flask filled with cold water and surrounded by a mix-
ture of ice and salt. Gently heat the generating-flask con-
taining the starch and nitric acid, when the trioxide is plen-
tifully produced, and absorbed by the cold water with which
it unites, thus : —
N203+H2O = 2HN02,
nitrous acid being formed by this union. Also, pass a portion
of the nitrogen trioxide into a cold solution of potassium
hydroxide, when potassium nitrite will be formed, thus : —
2 KOH + N2O3 = 2 KNO2 + H2O.
Preserve the above for work under nitrous acid.
Nitrogen trioxide is of itself unimportant, except as being
e starting-point from which nitrous acid and its com-
pounds are formed. Consequently we will again refer to
it, omitting its tests, etc., for the present, since they are
the same as for nitrous acid. (Art. 72.)
64 NITROGEN AND OXYGEN.
NITROGEN TETROXIDE, NO2 (or N2O4).
65. Preparation, etc. — This oxide of nitrogen is unim-
portant, and is easily obtained by artificial processes, e.g.,
when lead nitrate, Pb(NO3)2, is heated in a hard glass
retort, dense reddish fumes of the tetroxide are evolved,
thus : —
Pb(N03)2 = PbO + 2 NO2 + O.
These fumes can be condensed by passing them into a
U-tube surrounded by a freezing mixture. When passed
into water, the following reaction occurs : —
2 N02+ H20 = HN02+ HN03.
NITROGEN PENTOXIDE, N2O5.
66. Preparation, etc. — Nitrogen pentoxide is a white
crystalline solid assuming the form of rhombic crystals or
six-sided prisms. Although from a scientific standpoint it is
an important compound, being the anhydride of nitric acid
(by anhydride of an acid we mean a certain oxide that,
uniting with water, produces that acid), it is, nevertheless,
so unstable, and difficult of preparation, that it is not
advisable to attempt its production in small laboratories.
There are several methods of obtaining the pentoxide, one
of which is by passing dry chlorine gas through a glass
tube containing silver nitrate. The reaction occurs in two
stages, thus : —
1 . AgNO3 + 2 Cl = NO2C1 -f- AgCl + O.
2. NO2C1 4- AgNO8 = N2O5 4- AgCl.
The pentoxide unites with water, thus : —
NITROGEN AND OXYGEN. 65
THE COMPOUNDS OF NITROGEN, OXYGEN, AND HYDRO-
GEN; OK, THE NITROGEN ACIDS.
67. There are three acids in this series, viz. : —
1. Hyponitrous Acid (theoretical), HNO.
2. Nitrous Acid, HNO2.
3. Nitric Acid, HNO3.
None of these acids occur free in quantity, and only the
last two have been prepared in the free state. It is in-
teresting to note the manner in which these acids may be
supposed to originate from the union of their anhydrides
with water, thus : —
N2O + H2O = 2 HNO.
N2O3 4- H2O = 2 HNO2.
N A + H2O = 2 HNO3.
QUERY. Can they all thus be produced ?
HYPONITROUS ACID, HNO, AND HYPONITRITES.
68. Preparation. — Hyponitrous acid has not been iso-
lated, but its compounds (called hyponitrites) with certain
metals are known.
EXP. 59 P. Add sodium amalgam to a strong solution of
potassium nitrate in a beaker, until hydrogen gas escapes.
Potassium hyponitrite will be formed, thus : —
KN03 + 4H = KNO + 2H2O.
QUERIES. Is the above solution alkaline ? Whence comes the H of
the above reaction ?
Retain the solution thus prepared to make the follow-
ing : —
69. Tests for Hyponitrites. — 1. Hyponitrites in alka-
line solutions precipitate lead hyponitrite, Pb(NO)2, upon
66 NITROGEN AND OXYGEN.
addition of lead acetate, Pb(C2H3O2)2. This precipitate is
white, changing to yellow.
Pb(C2H3O2)2 + 2KNO = Pb(NO)2 + 2K(C2H3O2).
2. Hyponitrites in alkaline solutions do not turn a solu-
tion of starch paste and potassium iodide, KI, blue, while
acid solutions (use acetic acid to acidulate) do effect this
change.
3. In solutions acidulated with acetic acid they bleach
a solution of potassium permanganate, K2Mn2O8.
4. Upon adding silver nitrate to a nearly neutral hypo-
nitrite solution, silver hyponitrite, AgNO, a yellow pre-
cipitate, is thrown down.
KNO + AgNO3 = AgNO + KNO3.
NITROUS ACID, HNO2, AND NITRITES.
70. Preparation. — Nitrous acid is a very unstable com-
pound ; bat its salts, called the nitrites, are stable and well
known. It may be prepared by the action of nitrogen tri-
oxide upon water. (Art. 64.)
QUERY. How can you obtain a nitrite ?
71. Properties. — Nitrous acid, upon standing or upon
being heated, undergoes decomposition, thus : —
3 HNO2 = HNO3 + 2 NO + H2O.
The nitrites are all soluble in water ; and since they are
produced upon the surface of the earth by the transforma-
tions of decaying nitrogenous substances, they will be
found in drinking water contaminated with sewage. (See
Exp. 37.)
All nitrites deflagrate when thrown upon hot charcoal,
and they are decomposed by the action of stronger acids,
giving off fumes of nitrogen trioxide, N2O3.
NITROGEN AND OXYGEN. 67
EXPERIMENTAL PKOBLEM. Given : Starch, nitric acid, and potassium
hydroxide. Prepare, and test as you proceed, lst> N2O3; 2d, HN02; 3d,
K:x'O2; 4th, N2O3; using in each case the last substance produced to obtain
the next succeeding compound.
72. Tests for Nitrous Acid ami Nitrites. — 1. Free
nitrous acid turns a solution of starch paste and potassium
iodide blue.
2. It bleaches a solution of potassium permanganate.
3. In solutions acidified with acetic acid, the nitrites
bleach a solution of potassium permanganate. When
acidulated with acetic acid, they give a white precipi-
tate, AgNO2, with silver nitrate.
EXP. PROB. Let the student have two unlabelled solutions, one a nitrite
and one a hyponitrite ; then let him determine which is the nitrite.
NITRIC ACID, HNO3, AND THE NITRATES.
73. Occurrence and Preparation. — Free nitric acid
barely occurs in nature ; but its compounds, as potassium
nitrate, KNO3, or saltpetre, and sodium nitrate, NaNO3, or
Chili saltpetre, are found in large quantities.
EXP. 60 T. Place in the retort A (Fig. 14) equal parts, by
weight, of strong sulphuric acid and pulverized potassium
nitrate. Surround the receiver R with snow or ice, or allow
a stream of cold water continually to flow over it. Apply heat
to A, which rests upon a piece of wire gauze, when nitric acid
will be given off aud condensed in R. As soon as the opera-
tion is finished, pour the acid into a glass-stoppered bottle,
and reserve for a few experiments which will be given under
"Properties."
Nitric acid thus prepared is apt to be colored, owing to
the presence of some of the lower oxides of nitrogen ; but
chemically pure nitric acid is colorless. The commercial
68 NITROGEN AND OXYGEN.
acid is prepared on the large scale by treating Chili salt-
petre, NaNO3, in iron retorts with sulphuric acid, the vapors
being condensed in stoneware condensers. The acid is
afterwards purified by distillation. If much water be
present, a weak acid is at first obtained ; if little water be
present, a stronger acid distils over ; but in either case an
acid of the specific gravity of 1.4 is finally obtained.
74. Properties. — EXP. 61 T. Place in an evaporating dish
2cc or gcc of the strong acid obtained above. With the aid of
a long-handled deflagrating spoon, drop in a small piece of
FIG. 14.
phosphorus. It usually takes fire, and that, perhaps, with
explosive violence. There is some danger attendant upon this
experiment.
EXP. 62 P. Heat to redness some finely-powdered charcoal
in an iron sand-bath. A few drops of strong nitric acid will
cause the charcoal to deflagrate.
Nitric acid is an exceedingly powerful oxidizing agent
owing to the ease with which it gives up a part of its oxy-
gen. This oxygen, when in a nascent condition, that is, at
the moment it is liberated, is by far more active chemi-
NITROGEN AND OXYGEN. 69
cally than when in a free condition. We may here note
that the same is true of all elements when in a nascent
state.
EXP. 63 P. Ignite a small quantity of spirits of turpentine
in an evaporating dish by carefully adding a few drops of a
mixture of equal parts nitric and sulphuric acids.
Nitro-glycerine is prepared by treating common glycerine
with these acids, at low temperatures. What is dynamite ?
EXP. 64 p. Sprinkle upon red-hot charcoal finely-powdered
potassium nitrate. What occurs?
Common gunpowder is a mechanical mixture of potas-
sium nitrate, sulphur, and charcoal. Gun-cotton and
wood-powder are made by treating vegetable fibres with
nitric acid.
EXP. 65 P. To about 20CC pure water in an evaporating dish
add one or two drops of nitric acid. Now drop in some goose-
quill clippings or the parings of the finger-nails. Evaporate the
solution to dryness, when the cuttings will turn yellow. Also
try the same upon white silk thread. What occurs?
QUERY. How does nitric acid act upon the skin and similar organic
substances ?
EXP. 66 P. Drop a few drops of nitric acid upon copper
filings in a test-tube, and note the brownish-red fumes evolved.
What are these fumes ?
When rnetals react with nitric acid, substances called
nitrates are obtained, some of which are very useful, as we
shall hereafter see.
Nitric acid is one of the most important acids known in
chemistry. In addition to the uses above indicated, we
may add that it is used in making coal-tar colors and
various other economical products, while in the laboratory
70 NITROGEN AND OXYGEN.
it is used as an indispensable reagent, serving as a solvent
for most metals, — since the nitrates are all soluble in
water, — and as a point of departure in the preparation of
all the other oxides and acids of nitrogen.
75. Tests for Nitric Acid, HNO3, and the Nitrates. —
1. Make in a test-tube a solution of ferrous sulphate,
FeS04, and add sulphuric acid, H2SO4. Shake well, and
allow to stand till cool ; then, without mixing, carefully
pour in the solution to be tested. Now lightly tap with
the finger on the side of the test-tube. If nitric acid or a
nitrate be present, a brown ring will be formed where the
liquids meet. Upon shaking, the ring disappears.
EXPLANATION. Nitric oxide, NO, is liberated, which, uniting with the
ferrous sulphate, forms the brown substance of the ring, thus : —
2 KN03 + 4 H2S04 + 10 FeSO4
K2S04 + 3Fe2(SOJ3 + 4 H2O + 2(FeS04)2NO.
2. A solution of a nitrate with sulphuric acid and a few
bits of copper will give off reddish fumes.
EXPL. 2 KNO3 + 4 H2SO4 + 3 Cu = K2S04 + 3 CuS04 + 4 H2O + 2NO.
The NO coming in contact with the air absorbs atmospheric oxygen, thus :
2 NO + 2 O = 2 N02. What is NO2 ?
3. Nitrates are distinguished from nitrites thus : Add
acetic acid to ferrous sulphate, then add the solution to be
tested. Nitrates produce no change : nitrites turn the
solution brown.
HYDROXYLAMINE, NH3O.
76. Nitrogen, hydrogen, and oxygen, form another com-
pound of some scientific interest, hydroxylamine, NH3O,
which may be regarded as a compound of ammonia and
NITROGEN AND OXYGEN. 71
oxygen. If we represent ammonia thus, N -s H, we may
in us, IN -s n
H (-H
represent hydroxylarnine thus, N -j H , showing that the
(.OH
group OH, or hydroxyl, which the student must have
noticed as occurring in the hydroxides, has displaced one
atom of hydrogen in ammonia.
This substance in an aqueous solution possesses strong
reducing powers, being capable of throwing down, in a
finely-divided state, some of the metals from their solu-
tions. It is produced by the action of nascent hydrogen
on nitric oxide, thus : —
It turns a solution of ctipric sulphate orange-yellow,
forming cuprous oxide, Cu2O. This reaction serves as a
test.
EXERCISES IN NITROGEN.
1. PROB. How many litres of air would be required in preparing 51 of
nitrogen ? How much phosphorus 1
2. Make a cork boat, and place thereon a small quantity of iron filings
moistened with ammonium chloride ; float the boat on water, and place
over it a tall glass jar, the mouth of which is to dip under water ; note the
volume of air, and in two or three days again examine. What change has
occurred in the iron ? What alteration in the volume of air ? How can
you determine the volume of 0 and N in air, provided the iron has united
with all the oxygen ? Test the residual gas. Is it nitrogen ?
3. PROB. How many grams of NH3 can you obtain from 80s of NH4C1?
4. PROB. How many grams of nitrous oxide are to be had from 400s of
NH4N03 ?
5. PROB. How many pounds of nitric acid may be obtained from one
ton of Chili saltpetre ? How much H2S04 will be required to produce it ?
6. If you had KNO3, H2SO4, Na, Hg, starch, and NH4OH, and no other
reagents, show how you could prepare all the oxides and acids of nitrogen
excepting N205.
72 NITROGEN AND OXYGEN.
7. Which compound treated in this chapter is the most valuable to
commerce ? Which next ?
8. Write a short sketch of the chemist Rutherford, who discovered
nitrogen in 1772.
9. What is the derivation of the words nitre and nitrogen ? (Consult a
dictionary.)
10. NH3 represents two volumes of ammonia. In its production, one
volume of N has united with three volumes of H. How much condensa-
tion has occurred? What is the density (Art. 87) of NH3 as compared
with H ? (Sue. 14 + 3 = 17 and 17 -j- 2 = 8.5.) What is the density of
N2O ? Of N2O3 ?
11. In practice one determines the amount of ammonia present in
drinking-water, thus : Proceed as in Exp. 35, using I1 of the water to be
tested. The first jar contains three-fourths of all the ammonia in the
sample (Wanklyn). In a similar tall jar is placed 50CC pure water and
about 4CC Nessler's solution. To this, from a burette graduated to tenths
of a cubic centimetre, is added a standard solution of ammonium chloride,
NH4C1, drop by drop, with constant stirring until the same color is reached
as in the first jar (Exp. 35). The number of cubic centimetres standard
solution added equals the number of milligrams of ammonia per litre.
The standard solution of ammonium chloride is prepared by dissolving
3.15s of the dry salt in I1 distilled water.
Good drinking-water should not contain over 0.08 parts per 1,000,000,
of free ammonia.
QUERY. Should the qualitative tests fail to detect organic matter,
ammonia, nitrites, etc., can there be a question as to the potableness of
the water under examination ?
12. Nitric acid containing oxides of nitrogen may be freed from the
latter by passing through it for some time a current of pure air.
GENERAL NOTE. Recent investigations throw doubt upon the existence
of free nitrogen trioxide in a gaseous condition. Some authors also give
ammonium hydroxide, NH4OH ; but there are grave doubts as to its exist-
ence. It is, at least, decomposed by boiling.
CHAPTER IV.
BINARY COMPOUNDS. — ACIDS. — BASES. — SALTS. -
CHEMICAL NOMENCLATURE.
77. Binary Compounds are those which consist of but
two elements. Oxygen unites with all other elements
except fluorine, and the compounds thus formed are known
as oxides. Similarly, the binary compounds of sulphur are
known as sulphides ; those of chlorine, bromine, and iodine,
as chlorides, bromides, and iodides.
The principal elements whose binary compounds are
named in this way are bromine, chlorine, fluorine, iodine,
oxygen, selenium, sulphur, and tellurium.
To distinguish between the different oxides, chlorides,
etc., the name of the element in combination with oxygen,
chlorine, etc., is prefixed. Thus sodium chloride is the
compound of sodium and chlorine ; magnesium chloride is
the compound of magnesium and chlorine ; barium oxide
is the compound of barium and oxygen ; potassium iodide,
the compound of potassium and iodine, etc.
SUG. Student, name the compounds, the formulae of which are here
CaO, BaCl2, KF, Nal, MgO.
It sometimes occurs that oxygen, chlorine, bromine, etc.,
unite with other elements in more than one proportion, as
illustrated by the formulae, HgO and Hg2O, CuO and
Cu2O, FeCl2 and Fe2Cl6, etc. In these cases the simple
prefixing of the name of the element which is in combina-
tion with oxygen, chlorine, etc., will not suffice. Hence
74 BINARY COMPOUNDS.
the name is modified by the suffixes -ic and -ous. We
have not simply mercury compounds, but mercuric and
mercurous compounds, etc. That compound which con-
tains the smaller proportion of oxygen, chlorine, etc., is
designated by the suffix -ous, and that which contains
the larger proportion is designated by the suffix -ic.
Thus, of the two compounds of mercury and oxygen, that
which has the formula Hg2O is called mercurous oxide,
because it contains less oxygen in proportion to the mer-
cury than the other compound, HgO. The latter is called
mercuric oxide. In naming compounds of copper, iron, tin,
lead, and some other elements, when the syllables -ic
and -ous are necessary, the Latin names of the elements
are used. Instead of speaking of copperous and copperic,
or of ironous and ironic compounds, we use the words
cuprous and cupric, ferrous and ferric, compounds, etc.
The compounds CuO and Cu2O are known respectively as
cupric arid cuprous oxides : FeCl2 and Fe2Cl6 are called
ferrous and ferric chlorides.
There are cases in which a given element unites with
oxygen, chlorine, etc., in more than two proportions. It
is then necessary to use other methods in naming the com-
pounds. Manganese forms four compounds with oxygen.
These have respectively the compositions expressed by the
formulas MnO, Mn2O3, Mn3O4, and MnO2. To these are
sometimes given the names manganous oxide, MnO ; man-
ganic oxide, Mn2O3; manganoso-manganic oxide, Mn3O4,
the name signifying that the compound is made up of
manganous and manganic oxides ; and manganese dioxide,
MnO2.
It is not uncommon to indicate by the name the number
of oxygen atoms represented in the formula, as in the
case of oxides.
i
ACIDS. 75
Those containing one atom of oxygen are called monoxides ;
" two atoms *' " dioxides;
" three " " " trioxides;
" four " " tk tetroxides;
" j#y« " " " pentoxides, etc.
The relation 2 to 3 is sometimes expressed by the word
" sesqui," e.g., Fe2O3, sesquioxide of iron, which is the old
name for what is now called ferric oxide.
78. Acids. — Among the compounds thus far considered
are nitric and nitrous acids ; and frequent reference has
been made to sulphuric acid and hydrochloric acid. Indeed,
it would be difficult to write a page on any chemical sub-
ject without the use of the word "acid." What is an
acid? An exact definition cannot well be given. By the
term " acid " we mean a body with certain physical and
chemical properties, the chief of which are the following :
a sour taste ; the power to turn certain vegetable colors,
as to turn blue litmus red ; the power of giving up hydro-
gen, and taking up metals (bases) in its place.
EXP. 67 P. Student, test with blue litmus every acid to be
found in the laboratory. Do all the acids have the same effect
on the color? In testing, take a few drops of the acid in a
test-tube half full of water. Try substances which are not
acids, as common salt. What effect is produced?
According to the above statement regarding the proper-
ties of acids, all acids must contain hydrogen. It does not
follow that all bodies which contain hydrogen are acids.
Ammonia, NH3, for example, has properties quite the
opposite of those possessed by acids; and many other
examples might be cited. In order to have acid proper-
ties, we must have the hydrogen in combination with cer-
tain elements, or gixnips of elements.
76 ACIDS.
The elements whose hydrogen compounds are markedly
acid are chlorine, bromine, iodine, and fluorine, which give
hydrochloric acid, HC1 ; hydrobromic acid, HBr ; hydriodic
acid, HI ; and hydrofluoric acid, HF. The hydrogen com-
pounds of sulphur, selenium, and tellurium, are weak acids.
Most acids consist of hydrogen in combination with
oxygen and some other element, as nitric acid, HNO3 ;
nitrous acid, HNO2; sulphuric acid, H2SO4, etc. They
are commonly called oxygen acids to distinguish them from
those which contain no oxygen. There are a great many
acids belonging to this class, but only a few of them are
in common use.
In naming the oxygen acids, the same suffixes -ous and
-ic are used, as in the case of binary compounds, and
with the same significance .
If an element forms only one acid with oxygen and
hydrogen, the suffix -ic is used. If it forms two acids,
that which contains the smaller proportion of oxygen is
designated by the suffix -ous, and that which contains
the larger proportion of oxygen, by the suffix -ic. Thus
we have
Nitrous acid, HNO2,
and Nitric acid, HNO3 ;
Sulphurous acid, H2SO3,
and Sulphuric acid, H2SO4, etc.
In those cases in which more than two acids are formed
by the same elements, prefixes are used in addition to the
suffixes. A good illustration of the use of these prefixes is
furnished by the acids of chlorine. This element forms
four acids with oxygen and hydrogen. They are repre-
sented by the formulae HC1O, HC1O2, HC1O3, and HC1O4.
Of these the second and third are known as chlorous and
BASES AND SALTS. 77
chloric acids. The first is called hypochlorom acid, which
signifies that it is below chlorous acid as regards the
amount of oxygen it contains. The fourth is called per-
chloric acid, which signifies that it is beyond chloric acid
in the series. These prefixes hypo- and per- are frequently
used in this sense.
79. Bases. — There are certain compounds which have
properties almost exactly the opposite of those of acids.
They are called bases. The name "base" has been applied to
various bodies, and with different meanings. In general, we
mean by a base a substance which has the power of neutral-
izing acids, that is, destroying their acid properties. The
bases, like the acids, consist of certain elements in combi-
nation with oxygen and hydrogen. Some elements unite
with oxygen and hydrogen to form acids; and others
unite with oxygen and hydrogen to form bases. Nearly
all the compounds which the metals form with hydrogen
and oxygen are bases. Examples are : potassium hydrox-
ide, KOH ; calcium hydroxide, Ca(OH)2, etc. The stronger
bases are known as alkalies, among which are the hydrox-
ides of potassium and sodium, formerly called caustic potash
and caustic soda.
80. Salts. — When an acid and a base react, they tend
to neutralize each other. The acid properties and the
basic properties are usually both destroyed, and a new
body is formed which is neither acid nor base. This new
body is called a salt. The relation between an acid and
the salts derived from it will readily be seen by examining
the following formulae : —
r NaCl,
Hydrochloric acid, HC1, yields the salts •< KC1,
(CaCl2, etc.
78 SALTS.
fKN03,
Nitric acid, HNO3, yields the salts . •] NaNO3,
(Ba(NO3)2, etc.
rK2S04,
Sulphuric acid, H2SO4, yields the salts -< BaSO4,
(Na2SO4, etc.
On comparing the salts with the acid from which they
are derived, we see that the difference between them is
simply this, that the acid contains hydrogen while the
salts contain something in the place of the hydrogen.
We shall see later that this something which takes the
place of hydrogen is called a metal. Thus, in the exam-
ples given above, the metals sodium, Na, potassium, K,
calcium, Ca, and barium, Ba, take the place of hydrogen
in the acids.
Each acid can yield at least one salt with every metal,
and in some cases more than one. The salts of each acid
receive a general name, and we distinguish between the
different salts of the same acid by prefixing the name of
the metal.
The salts of the simplest acids, such as hydrochloric,
hydrobromic, and hydriodic acids, are named, as described
above, under the head "Binary Compounds" (see p. 73).
Salts of the oxygen acids are named thus : when the
name of the acid ends in »<?, the name of its salts ends in
ate ; and, when the name of the acid ends in ous, the name
of its salts ends in ite. Thus, a salt of nitric acid is called
a nitrate ; of nitrous acid, a nitrite ; of sulphuric acid, a
sulphate ; of sulphurous acid, a sulphite, etc. From nitric
acid we thus have a series of nitrates corresponding to
the different metals. We distinguish between them by
using the names of the metals as adjectives, as in the case
of binary compounds. The potassium, sodium, and cal-
SALTS. 79
cium salts of nitric acid, for example, are called potassium
nitrate, KNO3, sodium nitrate, NaNO3, and calcium nitrate,
Ca(N03)2.
The metals mercury, iron, copper, and some others yield
two different classes of salts, corresponding to the lower
and higher oxides alread}7 mentioned. Just as we have
mercurous and mercuric oxides and chlorides, ferrous and
ferric chlorides, etc., so also we have mercurous and mer-
curic nitrates, sulphates, etc., and ferrous and ferric ni-
trates, sulphates, etc. The two nitrates of mercury will
serve as examples. We have
Mercurous nitrate, Hg2(NO3)2,
and Mercuric nitrate, Hg(NO3)2.
The principle of nomenclature adopted for these salts is
the same as that described in connection with the oxides,
chlorides, etc. The name of that salt which contains the
smaller proportion of the acid constituent ends in ous,
while the name of that one which contains the larger pro-
portion of the acid constituent ends in ic.
The action of metals upon acids may be illustrated by
the following equations : —
Zn + H2SO4 = ZnSO4 + 2 H ;
Zn + 2HCl = ZuCl2 +2 H.
In these cases the metal simply replaces the hydrogen
which is set free. This action takes place only in the case
of the stronger acids.
When an acid acts upon a base, the action is as repre-
sented below : —
KOH + HNO3 = KNO3 + H2O ;
NaOH + HNO3=NaNO3-f H2O ;
2 KOH + H2SO4 = K2SO4 + 2H2O;
+2H2O.
80 ACID AND NORMAL SALTS.
This kind of action takes place between all acids and
all bases.
81. Acid and Normal Salts. — The simplest acids, such
as hydrochloric and nitric acids, yield only one salt each
with most of the metals. Thus hydrochloric acid and
potassium yield only one potassium chloride, KC1, which
is a neutral body ; nitric acid and sodium yield only one
sodium nitrate, NaNO3, which is also neutral.
There are some acids, like sulphuric acid, H2SO4, which
have the power of yielding two or more salts with the
same metal. Thus sulphuric acid yields with potassium
riot only the salt, K2SO4, potassium sulphate, but another
salt, of the formula KHSO4, which contains only half as
much potassium as the first. In the first case all the
hydrogen 'of the acid has been replaced, and the resulting
compound has no acid properties. It is a normal salt. In
the second case a part of the hydrogen is left, and the
compound still has acid properties. It is both acid and
salt, and is called an acid salt.
A normal salt is one which is formed by replacing all
the hydrogen of an acid with a metal.
An acid salt is one which is formed by replacing only a
part of the hydrogen of an acid with a metal.
In naming the acid salts it is customary to indicate the
number of atoms of the metal which are represented in
the formula. Thus the salt KHSO4 is called mono-potas-
sium sulphate ; the salt Na2HPO4 is called disodium phos-
phate. Sometimes they are referred to as acid salts,
mono-potassium sulphate being called acid potassium sul-
phate.
Applications of these principles of nomenclature will be
met with when the salts are considered. Meanwhile the
ACID AND NOKMAL SALTS. 81
student should familiarize himself with the main points
by means of examples furnished by the teacher. A few
examples are here given.
Student, name the compounds, the formulae of which
are given below : —
Cu2Cl2, KNOS, Ca(NO3)2, Fe(NO3)2, HgO, NaHSO4,
CuCl2, NaN02, Ba(N03)2, Ve2(KOs)<» HgA K2SO4.
WRITING EQUATIONS.
It is important to know how to write chemical equations, and thus avoid
the necessity of committing them to 1116111017. In the first place we know
what substances we put together or experiment upon, and these are placed
in the first member, and connected by the sign +. The substances formed
are determined by experiment, or, when our knowledge is sufficient, by
analogy or by induction. They are then placed in the second member,
and connected by the + sign. It now remains to balance the equation.
The fundamental principle to be remembered here is, that " matter is in-
destructible "; that is, just as many atoms of a given dement as appear in one
member, just so many must also appear in the other. Let us, for example,
write the equations for the reaction of NaCl and H2S04 : we first write
NaCl + H2S04 = • • • ; by experiment, we know that under certain conditions
(Art. 94) HC1 and HNaS04 are formed, and we proceed to the next step,
thus : NaCl + H2SO4 = HNaS04 + HC1. By inspection, we see that the
equation balances and is complete. But let us suppose that the conditions
are different, and that Na2SO4 and HC1 are produced ; the second step
gives us NaCl + H2SO4 = Na2SO4 + HC1. By inspection we here see that
the equation is not true, since two atoms of Na appear in the second mem-
ber, and but one in the first ; also one atom of H in the second, and two
in the first. We may obtain the required amount of Na by doubling the
NaCl ; and, when this is done, the necessity for doubling the HC1 becomes
apparent, and the equation balances, thus : —
2 NaCl + H2SO4 = Na2SO4 + 2 HC1.
Water, which is almost always present, must sometimes be taken
into consideration. The equations previously given will afford good prac-
tice, also those to follow, especially those relating to the metals. Before
we reach that point, however, molecular equations will be explained.
CHAPTER V.
THE ATMOSPHERE. — LAWS OF PRESSURE, TEMPERATURE,
DENSITY, AND VOLUMES OF GASES. — PROBLEMS.
THE ATMOSPHERE.
82. The earth is everywhere surrounded by an ocean of
gaseous vapor, called the atmosphere, which varies from
fifty to one hundred miles in height. This variation at
any one point is never ceasing, for just as in the oceans of
water, so 'in this ocean of air, do huge waves continually
surge to and fro, — waves so vast that their altitudes are
measured in miles.
Every object upon the surface of the earth is subjected
to the pressure exerted by the weight of air above. This
pressure varies constantly, and, owing to the great mobility
of the particles of air, it is exerted in all
directions, — downwards, upwards, and side-
wise.
This pressure is measured by an instru-
ment called a barometer (Fig. 15). A is a
glass tube about 800mm long, sealed at the
upper end, open at the lower, and provided
with a scale. This tube is filled Avith mer-
cury, and inverted in a cup of mercury,
C. Now, since the tube itself sustains
the pressure which the atmosphere would
FIG. 15. exert on this column of mercury within the
THE ATMOSPHERE. 83
tube, in every direction except upwards, it follows that
the column will remain at a higher altitude than the level
of the mercury in the cup. The height of this column
of mercury will depend upon how hard the atmosphere
presses it upward.
At the level of the sea, in the latitude of Paris, and at
0° C., the average height of this column is 760mra; hence
760mm is taken as the standard pressure of the air.
As you ascend from the sea-level the column falls
(why?), and as you descend it rises (why?).
As the density (Art. 88) of the mercuiy and the at-
mosphere varies, owing to changes of temperature, the
height of the barometer varies ; hence the necessity of
taking a standard temperature, which is 0° C.
83. Measurement of the Temperature of the Atmos-
phere.— This is accomplished by means of instruments
called thermometers. There are three scales in use, —
Centigrade, Fahrenheit, and Reau-
mur (Fig. 16). Thermometers are
made by blowing bulbs on capillary
tubes. The bulbs and tubes are filled
with mercury, and then heated till
the mercury issues in vapor, when
the ends are suddenly sealed by the
blow-pipe flame. They are graduated
by first plunging them into melting
ice, the height of the column of mer-
cury being marked 0° C., 0° R., or
32° F. The instruments are next
placed in the steam of boiling water,
and the height of the column of mercury marked 100° C.,
80° R., or 212° F. The distances between these points are
84 THE ATMOSPHERE.
then divided into spaces (or degrees), there being 100
divisions C., 80 R., or 180 F. ; divisions of the same length
are also made above and below these points. From the
manner of laying off these scales, it follows that
5° C. = 4° R. = 9° F.
The following formulae will assist in changing from the
reading of one scale to another : -
(a)C. = (F.-32)f
(6) F. =f C.. + 32.
(c) R. = tC.
P-ROB. Change 98° C. to F. ; 87° R. to F. ; 91° F. to R. ;
- 18° C. toF. ; -40°F. to C.
Sue. It would be advisable for the student to learn the points of a
good thermometer from his text-book in physics, and to review the metric
system in Iris arithmetic.
NOTE. The centigrade thermometer and the metric system of weights
and measurements are used throughout this work, as they best answer its
purposes, and are the ones used by scientists in general.
84. Impurities in the Atmosphere. — As we have
already learned, air is a mechanical mixture of nitrogen
and oxygen. By this we mean pure air. But atmos-
pheric air is never pure. It contains, —
(a) Moisture, as invisible vapors, clouds, and fogs.
These, being lighter than the atmosphere, cause a lower
barometer, especially when they are present in large
quantities.
(6) Carbon dioxide, CO2, produced by combustion, by
the respiration of all air-breathing animals, and by the
decomposition of animal and vegetable tissues.
(c) Ammonia (Art. 51).
(d) Ozone (Art. 30), or other substances having marked
oxidizing power.
THE ATMOSPHERE. 85
(e) Dust and smoke.
(/) Other gases in small quantities, which are liberated
in various ways.
85. Determination of the Volumes of Nitrogen and
Oxygen in the Atmosphere. — This determination is
made by means of Ure's eudiometer. A measured quan-
tity of pure air is introduced into the graduated limb,
and then a volume of hydrogen, more than sufficient to
combine with the oxygen of the air, is added. The whole
volume is now carefully noted, the spark passed, and the
diminution of volume carefully ascertained.
One-third of this diminution equals the volume of the
oxygen contained in the air.
The volume of oxygen is subtracted from the volume of
air introduced at the beginning, and this gives the volume
of the nitrogen.
In this way we learn that the air consists of oxygen 21
volumes, and nitrogen 79 volumes, in 100 parts. These
proportions vary but slightly in any locality or season.
QUERIES. What chemical action takes place when the spark is passed ?
How do you know one-third of the volume of diminution to be the volume
of 0 ? Through what substances would you pass air to remove its im-
purities ?
86. Effect of Pressure on the Volume of a Gas. — If
a mass of gas be confined in an air-tight cylinder, and a
perfectly-fitted piston be pressed down into the cylinder,
the gas will be compressed into a smaller volume.
The law for the volume of a gas under such conditions
is : —
LAW I. The volume of any gas, its temperature remaining
constant, varies inversely as the pressure.
86 THE ATMOSPHEKE.
We mean by this that volume 1 under pressure 1
becomes volume \ under pressure 2, volume J under
pressure 3, or volume J under pressure 4, etc. ; and the
reverse of this is also true when the temperature remains
the same in both cases. (How could this law be discov-
ered if it were unknown ?)
NOTE. On l<Jcm, at the standard pressure of 760mm, the atmosphere
exerts a pressure of 1033.3s (nearly 15 Ibs. per sq. in.), which is called a
pressure of 1 atmosphere. A pressure of 2 atmospheres is 2 X 1033.3s, etc.
Now, since gases are subject to the pressure of the atmosphere, their
volume varies with every change of the barometer.
Sue. The student should consult some work on physics for the experi-
mental demonstration of the above law, as well as for that of the succeed-
ing law, since we use them as an application of physics to chemistry.
PROBLEMS. 1. What volume will 101 of gas at 762mm occupy
when the barometer stands at 758mm?
SOLUTION. Since the volume varies inversely as the pressure, we have
the proportion,
758 : 762 : : 10 : x == 10.0527 + litres. Ans.
2. 190CC of gas at 760mm pressure becomes how many cubic
centimetres at 765"""?
3. A mass of gas, 1001 under 755mm, is subjected to a
pressure of 4.5 atmospheres; what volume will it occupy?
Ans. 22. 076+ litres.
NOTE. In these problems the temperature is considered constant.
87. Effect of Heat on the Volume of Gases. — It has
been found by experiment that 273 volumes of any gas
at 0° become 274 volumes when its temperature is raised 1°,
275 when raised 2°, etc., increasing one volume for each
degree of increase in its temperature ; also that 273 volumes
at 0° become 272 volumes when its temperature is lowered
1°, 271 volumes when lowered 2°, 270 when lowered 3°,
THE ATMOSPHEKE. 87
etc., decreasing one volume for each degree of decrease in
its temperature. According to this, the volume of a gas at
— 273° C. would be 0 ; and this point is designated as the
absolute 0 of temperature. Hence the absolute tempera-
ture of any body is the temperature above the ordinary
0 -f 273, or t + 273. Taking these ideas into account, we
have : —
LAW II. The volume of any gas, its pressure remaining con-
stant, varies as its absolute temperature, i.e, in the ratio of
273 + t to 273 + t1.
REM. 1. t is the observed temperature of the gas, and *' the required
temperature.
REM. 2. Since any gas surrounded by the atmosphere will usually be
of the same temperature as the atmosphere itself, it follows that the
volume of that gas will vary as the thermometer varies.
PROBLEMS. 1. At +15° the volume of a gas is 841 ; what
will be its volume at + 85° ?
SOLUTION. 273 + 15 : 273 + 85 : : 84 : x = 104.4166 + litres. Am.
2. A gas at —15° has a volume of 181 ; what will be its
volume at 100°?
SOLUTION. 273 — 15 : 273 + 100 : : 18 : x = 26^ litres. Am.
3. 981 of gas at - 4° become how many at - 24°?
4. 1761 of gas at + 100° become how many at — 140°?
5. 801 of gas at 0° become how many at — 18°?
6. 1441 of gas at — 15° become how many at 0°?
NOTE. In these problems the pressure is considered constant.
PROBLEMS in which both pressure and temperature vary : —
1. A mass of gas at -f- 15° and 762mm pressure occupies 941 ;
what will be its volume at + 25° and 758mm pressure ?
SOLUTION. We here have a combination of the principles of Arts. 86
and 87, pressure and temperature both affecting the volume of the 941 in
88 THE ATMOSPHBKE.
question. We will consider them separately ; hence the compound pro-
portion : —
• (1) Temperature, 273 + 16 : 273 +
I (2) Pressure . . . 758 : 762
. . _ „._„„,
2. 90CC of gas at 0° and 760mm occupy what volume at - 140°
and 40 atmospheres? Ans. 1.09+ cubic centimetre.
NOTE. From this the student may judge of the effect of pressure and
reduction of temperature.
3. 721 at — 12° and 4 atmospheres pressure become how
many litres at 100° and 760mm?
4. In the evening a quantity of oxygen gas was generated
in a laboratory. The oxygen receiver, holding 1121, was filled
while the barometer read 760mm and the thermometer + 15° C.
The next morning the barometer fell to 758mm, and the janitor
allowed the temperature of the room to reach -f- 40°. At 4 P.M.
that day 'the barometer read 760mm and the thermometer -j- 18°,
when the master ascertained that he had but 941 of the gas
remaining ; upon which he charged a student, who alone had
access to the laborator}', with having used some of the oxygen.
Allowing half a litre to have been absorbed by the water during
the night, the temperature and pressure remaining constant mean-
while, how much gas, measured at + 15° and 760mm, was lost
owing to the fall of the barometer and negligence of the janitor ?
Was the master justifiable in making the charge against the
student? How much (if any) of the oxygen did the student
use?
88. Relation of Weight to Density. — By the density
of a substance we mean the amount of that substance
contained in a given volume. We have seen how the
volume of a gas varies under differences of pressure and
of temperature. Now, it is evident that its density varies
also ; i.e.< whatever tends to make the volume less makes
the density greater, and whatever tends to make the vol-
THE ATMOSPHERE. 89
ume greater makes the density less. Again : it is evident
that the denser a given amount of gas, the greater will be
its weight, and the less dense the gas, the less its weight ;
or, —
The weight of a given volume of gas varies directly as its
density.
PROBLEMS. 1. How much will 101 of oxygen weigh at + 15°
and 765mm?
SOLUTION. We know that I1 of oxygen at 0° and 760mm weighs 1.430s;
therefore we will find how many litres this gas will be at 0° and 760mm, as
in Art. 87, and then multiply that result by 1.430, thus : —
( 288 • 273 )
I 760 : 765 / : : 10 : * " ®M ' ' ^ 9>M ' * M^' ^S'
2. How much will 201 of hydrogen weigh at 755mm and +20°?
3. How much will 151 of nitrogen weigh at —112° and 29
atmospheres pressure ?
89. Useful Problems. — I. To find the percentage com-
position of a compound. We will explain this by solving
a problem : What per cent of N and H in NH3 ?
SOLUTION. Now, it is evident that T37 of NH3 is hydrogen, and ^
is nitrogen. T3T expressed in the form of per cent equals
3H = 3 300 _:. 17 = 17.65 % of H. One can also readily understand
NH3 = 17 that 100 % - 17.65 % = 82.35 % of N.
These percentages are valuable in that they enable us
to make computations more rapidly. For example, if we
wish to know how much hydrogen there is in 10g of am-
monia, we have simply to multiply 10g by the per cent of
hydrogen, and divide the result by 100, when we have the
weight of the hydrogen in grams, thus : —
(17.65 x 10) -H 100 = 1.765*.
1 . What per cent of oxygen in HgO ? KC1O3?
2. What per cent of chlorine in NaCl? Of sodium?
90 THE ATMOSPHERE.
II. To find what volume ivill be occupied by a gas
obtained from a certain weight of chemicals. We can
also best understand this by a problem : How many litres
of oxygen can be obtained from 10g of potassium chlorate,
KC1O3, when the barometer reads 750mm and the thermom-
eter 25°?
SOLUTION. We will first ascertain what weight of oxygen 10s of KC10,
will yield. We can best do this by multiplying 10 by the per cent of O in
KC103. Thus we find the weight of oxygen to be 3.918s. We will now
ascertain how many litres 3.918s of oxygen will occupy at 760mm and
(P. One litre of oxygen under these conditions weighs 1.430s ; hence,
3.918 + 1.430 = 2.73981, or the number of litres at 0° and 760mm. We
can now finish the problem by Arts. 86 and 87, thus : —
1. How many litres of oxygen gas may be had from 100s
HgO when the barometer stands at 755mm, the thermometer
reading 20° ?
SUG. HgO (heated) = Hg + O.
2. How many litres of ox}Tgen gas may be had from 5008
MnO2, at 20° and 4 atmospheres pressure?
SUG. 3 MnO2 (heated) = Mn304 + 2 O.
3. How many litres of nitrous oxide may be obtained from
lk of NH4NO3 when the barometer reads 750mm and the ther-
mometer + 22"?
SUG. See Nitrogen Monoxide.
III. To find the weight of chemicals required to yield a
certain volume of gas. Let us again have recourse to a
problem : How many grams of KC1O3 will be required to
fill with oxygen a receiver of 321 capacity at 20° and
750mm?
THE ATMOSPHERE. 91
SOLUTION. We will first find what volume 321 of oxygen at 20° and
750nnn WOuld become when reduced to 0° and 760mm, in order to find the
required weight of the oxygen, thus : —
Now, 29.423 X 1.430 = 42.074^, or the required weight of oxygen. We
may now obtain the desired weight of the potassium chlorate by dividing
the weight of the oxygen by the percentage of O in KC1O3, or by the
proportion: 48 : 42.074 :: 122.5 :*.
1. How many grams KC1O3 will be required to yield 20l of
oxygen at — 20° and 760mm ?
2. How many grams of zinc and sulphuric acid are needed to
yield 401 of hydrogen at •+• 24° and 765mm ?
SUG. Zn + H2SO4 = ZnS04 + 2 H.
3. What weights of CaO and NH4C1 are required to make
25l of NH3 at 15° and 749mm?
SUG. CaO + 2 NH4C1 = CaCl2 + H2O + 2 NH3.
EXERCISES.
1. The following equations may be of service in making calculations
upon gases : —
til VH V'HI
1~'273 + * 273 + *'
V, H, and t represent respectively the volume, height of barometer, and
temperature of a gas under observed conditions, while V, H' ', and t' rep-
resent the same under required conditions, one of which will be unknown.
When t = t' we have,
(2) VH= VH' (Art. 86).
When H=H' we have,
2. Make a table showing the relations between the acids and their salts.
CHAPTER VI.
CHLORINE. — ITS OCCURRENCE, ETC. — HYDROCHLORIC
ACID. — AQUA REGIA. — CHLORINE OXIDES. — CHLO-
BINE OXACIDS.
CHLORINE.
SrifeoL CL'. — ATOMIC WEIGHT, 35.5; SPECIFIC
GRAVITY, 2.450.
90. Occurrence. — Chlorine does not occur free in
nature, owing to its great chemical activity ; in combi-
nation with certain metals, however, it occurs in large
quantities, as in sodium chloride, NaCl, or common salt.
Silver, chloride, AgCl, potassium chloride, KC1, calcium
chloride, CaCl2, and magnesium chloride, MgCl2, occur in
smaller quantities.
91. Preparation. — EXP. 68 p. In a test-tube place a
small quantity of manganese dioxide, MnO2, and add hydro-
chloric acid, HC1. Upon gently warming, chlorine is evolved
as a heavy, yellowish, suffocating gas, thus : —
MnO2 + 4 HC1 = MnCl2 + 2 H2O + 2 CL
Hold in the escaping gas strips of moistened litmus paper and
calico printed in organic colors ; they will be bleached. Also
note the fumes.
This method is sometimes employed in making chlorine
gas for the manufacture of bleaching-powder. (Art. 349.)
EXP. 69 P. Drop three or four small crystals of potassium
chlorate into a test-tube, and add hydrochloric acid. Warm
CHLORINE. 93
gently, and, when the chlorine fumes begin freely to appear,
immediately add 3cm or 4cm of cold water. What occurs may
be indicated thus : —
4 HC1 + 2 KC103 = 2 KC1 + 2 H2O + C12O4 + 2 Cl.
Try the effect of this solution upon vegetable colors as before.
Also add a few drops of the solution to tinctures of litmus,
carmine, and indigo ; they will be bleached.
The above method is one often employed by the
chemist in preparing chlorine water for such purposes as
testing iodine and bromine. Hereafter the student will
find frequent use for chlorine water thus prepared. It
might be well to say that the C12O4 and KC1 are in no
wise detrimental to the solution.
Another method of preparing chlorine in the manufac-
ture of bleaching-powder is of interest, since it is con-
tinuous and quite inexpensive. Hydrochloric acid gas,
mixed with air, is' passed over heated cupric sulphate,
CuSO4 ; the cupric sulphate undergoes no change, while
the oxygen of the air and ttie hydrochloric acid react,
thus : —
Another and common method of preparing chlorine is
as follows : —
EXP. 70 T. In the generating-flask A (Fig. 17) place equal
weight of common salt, NaCl, and manganese dioxide, MnO2,
which have been thoroughly pulverized and mixed. Then add
to this mixture twice its weight of dilute sulphuric acid (con-
sisting of equal weights of water and acid) . Apply a gentle
heat, and chlorine gas is plentifully given off. The Woulff
bottle B (Fig. 17) contains a little warm water to absorb any
hydrochloric acid gas that may be produced, while C contains
strong sulphuric acid to dry the gas. The thistle-top tube con-
94
CHLORINE.
tains a little sulphuric acid. Collect the chlorine in tall jars.
This may be accomplished by delivering the gas by means of
a long glass tube extending to the bottom of the upright jar.
The air will be pushed up and out of the jar.
NOTE. This method of collecting a gas is called displacement, and is
employed with those gases heavier than air.
If a pure, aqueous solution of chlorine be desired, it
may be obtained by attaching two or three Woulff
bottles, nearly filled with
cold water, and surrounded
with a cooling or freezing
mixture. Should the tem-
perature of any bottle con-
tained in the series nearly
reach 0°, a crystalline hy-
drate of chlorine is formed,
whose composition is
Cl + 5 H2O. In thus pre-
paring chlorine we may rep-
17. resent the reaction by, —
2 NaCl + MnO2+ 3 H2SO4 = 2 NaHSO4+ MnSO4 + 2 H2O + 2 Cl.
In reality, however, two distinct processes are involved.
In the first place the sulphuric acid, H2SO4, acts upon the
sodium chloride, NaCl, giving hydrochloric acid, HC1, and
mono-sodium sulphate, NaHSO4. Then the manganese
djoxide, MnO2, acts upon the hydrochloric acid, HC1,
giving manganous chloride, MnCl2, free chlorine, and
water, H2O. If an excess of sulphuric acid is present, it
decomposes the manganous chloride, MnCl2, giving man-
ganous sulphate, MnSO4, and hydrochloric acid ; and
the latter again acts upon manganese dioxide, yielding
chlorine.
CHLORINE.
95
The equations which give the best insight into the
reactions are the following : —
2 NaCl + H28O4 = Na2SO4 -f- 2 HC1,
and 4HC1 + MnO2=MnCL> + 2 H2O + 2 Cl.
Manganese dioxide readily gives up one part of its oxygen,
and it is this which, uniting with the hydrogen of hydro-
chloric acid, sets the chlorine free.
QUERIES. Why can you not collect chlorine over water or mercury ?
How can you collect hydrogen by displacement ?
92. Properties. — Chlorine is a heavy, greenish-yellow
gas having a strong and suffocating odor, and producing
great irritation to the lining membranes of the throat
and nostrils; and, when inhaled in sufficient quantities, it
is capable even of producing suffocation and death.
EXP. 71 P. Write with an organic (carmine) ink upon a
slip of printed paper ; moisten, and hold it in a large test-tube
full of chlorine gas. The writing disappears and the printing
remains. Printer's ink is made of lampblack (carbon), and is
not bleached.
QUERIES. How can you distinguish between organic and mineral
colors ? Try wall paper. Would chlorine water answer as well ?
Chlorine in the presence of moisture is an invaluable
bleaching reagent, acting upon vegetable coloring-matters
thus : —
2C1 + H2O=2HC1
Now this oxygen (liberated, as it is, within the fibres of
the substance to be bleached), while in a nascent condi-
tion, seizes upon the coloring-matters, and destroys them,
or changes them into colorless compounds.
EXP. 72 T. Saturate with hot turpentine, C10H16, a strip of
blotting-paper, and plunge it into a jar of dry chlorine gas.
96 CHLORINE.
*
The turpentine takes fire, the chlorine and hydrogen uniting,
while carbon is deposited as soot.
SUG. Student, write the equation.
EXP. 73 P. Plunge a lighted taper into a large test-tube of
chlorine. It continues to burn with a dull, red, smoky flame,
the chlorine again uniting with the hydrogen contained in the
substance of which the taper is composed, while the carbon is
set free.
REM. Oils, resins, gums, waxes, tallows, etc., are compounds containing
C, H, and 0, in varying proportions.
We thus see that chlorine possesses a powerful chemism
for hydrogen, even decomposing compounds to obtain it.
We shall hereafter see that the great chemism of chlorine
enables it to displace from their binary compounds the
nearly-allied elements, bromine and iodine. The sulphides
are also dissociated thus : —
Chlorine is extensively used as a deodorizer and disin-
fectant, owing its efficiency to its power of liberating from
water oxygen, which, as already explained, while in a nas-
cent state, oxidizes putrefactive vapors and disease germs
to their destruction.
Chlorine is soluble in water, lcc of water absorbing
nearly 3ee of this gas. It may be condensed to a liquid at
0° by a pressure of 6 atmospheres, or by 1 atmosphere at
- 34°. I1 at 0° and 760mm weighs 3.173^, and its specific
gravity is 2.450.
93. Tests for Chlorine. — Free chlorine gas or its
aqueous solution may be recognized by its color, odor,
or behavior, as in the preceding experiments.
CHLORINE AND HYDROGEN. 97
CHLORINE AND HYDROGEN.
HYDROCHLORIC ACID, HC1.
94. Occurrence and Preparation. — We now come to
an important and useful acid, the only compound formed
by hydrogen and chlorine, — hydrochloric acid, HC1. This
acid rarely occurs in nature, although it is a staple ar-
ticle of commerce. The following is the general method
of its preparation : —
EXP. 74 T. Heat to redness, in a crucible, 5g of common
salt, NaCl ; pulverize, and place in a generating-flask. Now
.add 10g strong sulphuric acid, H2SO4, and heat gently. Hydro-
chloric acid, in the form of a gas, is freely given off, and can be
collected by displacement, or over mercury.
By passing through two or three wash bottles, it may be
obtained in aqueous solution, the form in which it is used and
found for sale. The reaction is : —
NaCl 4- H2SO4 = HNaSO4 + HC1.
If a larger proportion of salt be used, the reaction may be
represented by this equation : —
2 NaCl + H2S04 = Na2SO4 + 2 HC1.
As we shall hereafter see, commercial hydrochloric acid
is almost exclusively obtained as a by-product of the
alkali works where common " soda " is prepared.
95. Properties. — Hydrochloric acid gas is extremely
soluble, lcc water at 0° dissolving no less than 505CC of this
gas. Its specific gravity is 1.247, it condenses at — 4°
under 25 atmospheres pressure, and I1 weighs 1.632s.
The aqueous solution of hydrochloric acid is one of the
most useful chemicals. It acts upon bases to form chlo-
98 CHLOK1NE AND HYDROGEN.
rides; the principal one of these, common salt or sodium
chloride, NaCl, occurs in nature in large quantities. The
gas has a pungent odor. In contact with the air it forms
dense white fumes, in consequence of its attraction for
water. The strong water solutions give off the gas read-
ily; weak ones may be concentrated by boiling.
EXP. 75 P. Take three test-tubes. In the first, place a solu-
tion of silver nitrate, AgNO3 ; in the second, a solution of mer-
curons nitrate, Hg2(NO3)2; and in the third, a solution of
plumbic acetate, Pb^H^Og^- To all three now add hydro-
chloric acid. What takes place ?
As the silver, lead, and mercurous chlorides are insolu-
ble in water, it precipitates these metals from solutions
in which they are contained. Other chlorides, as a rule,
are soluble.
This fact is taken advantage of in analyzing unknown
substances. Suppose, for example, we have a solution
which may contain any or all known metals. If we add
hydrochloric acid to it, and get a precipitate, we know
that one or more of the metals whose chlorides are insolu-
ble in water must be present. We know, in other words,
that one or more of the three metals, silver, lead, and
mercury, must be present ; and further, as their chlorides
are insoluble, we know that the addition of hydrochloric
acid to the solution removes these metals. By the
use of other chemical substances, other groups may be
precipitated in a similar way; and thus the problem of
determining what is in the substance under examination
is more and more narrowed down, until we know exactly
what is present. Substances which are used for the pur-
pose of precipitating groups of metals in analysis are
called GROUP-REAGENTS.
CHLORINE AND OX Y GEN. 99
When hydrochloric acid is mixed with one-third its
volume of nitric acid, AQUA REGIA or nitro-hydrochloric
acid is produced, which is the strongest solvent known ;
even gold and platinum are dissolved in it. The great
power of aqua regia lies in the fact that it readily gives
up chlorine, which, in a nascent condition, is very active.
The salts formed by aqua regia are chlorides. In using
this solvent it should be but slightly warmed ; a stronger
heat drives off chlorine to waste.
96. Test for Hydrochloric Acid, or Chlorides. — Their
solutions, even when acidulated with nitric acid, give a
white precipitate of silver chloride, AgCl, with silver
nitrate, AgNO3. This precipitate is insoluble in nitric
acid, and soluble in ammonia.
Sue. Student, try a solution of NaCl. Write the equation.
CHLORINE AND OXYGEN.
97. Chlorine and oxygen unite to form three compounds,
which have been isolated, viz., —
Chlorine monoxide C12O,
Chlorine trioxide ( ? ) • . . . C12O3,
and Chlorine tetroxide C12O4.
These oxides never occur free in nature, nor can they
be produced by the direct union of chlorine and oxygen ;
they may, however, be obtained by indirect processes.
Since they are unimportant, and dangerous to prepare,
owing to the ease with which they decompose, we shall
treat each but briefly.
98. Chlorine Monoxide, C12O. — This substance is a
100 CHLOKINE AND OXYGEN.
yellow-colored gas, prepared by passing chlorine gas over
dry mercuric oxide in the cold, thus : —
HgO + 4 Cl = HgCl2 4- C12O.
This gas may be pressed into a U-tube, surrounded with a
freezing mixture, arid condensed to a yellow liquid ; but
if the tube be suddenly jarred or scratched, as with a file,
it explodes with great violence. If exposed to direct sun-
light, it is also decomposed, but without explosion. It
unites with water, thus: —
C12O + H2O = 2 HC1O. (See Hypochlorous Acid.)
99. Chlorine Trioxide, C12O3. — This is a greenish-yel-
low gas, of great instability and explosive power. It can
be prepared in different ways, one of which is as follows :
Make a 'thin paste of 4 parts potassium chlorate, KC1O3,
and 3 parts of arsenious oxide, As2O3, with water ; place
in a generating-flask, and add a solution of 12 parts nitric
acid and 4 parts water; warm gently. This gas may also
be condensed to a liquid, but, owing to its extremely
uncertain and explosive propensities, the student should
not attempt its preparation.
It unites with water, forming chlorous acid, thus : —
C12O3 4- H2O = 2 HC1O2. (See Chlorous Acid.)
100. Chlorine Tetroxide, C12O4 or C1O2. — This is a
dark-yellow gas of small importance, as it forms no acids,
and consequently no distinct series of salts (it is also dan-
gerously explosive) ; but some idea of its deportment, as
well as that of the other chlorine oxides, may be gained
by the following experiment, which may be safely made if
care be used : —
THE CHLOKINE OXACIDS.
101
EXP. 76 P. Drop into a test-tube three or four small crys-
tals of potassium chlorate, KC1O3 ; then, holding the tube with
a pair of tongs, its mouth turned away from all persons pres-
ent, add a few drops of strong sulphuric acid. Warm gently,
when chlorine tetroxide gas will appear ; but a sharp and
vicious explosion soon terminates the experiment. The con-
tents of the tube are' thrown violently out, but the tube itself is
seldom broken. Note the odor of the gas.
THE CHLORINE OXACIDS.
101. This series contains four acids, none of which are
of commercial importance, nor are they of special value as
reagents ; and they all decompose upon standing. Their
salts, however, are stable, well known, and of great util-
ity. These acids are : —
Hypochlorous acid .... HC1O,
Chlorous acid HC1O2,
Chloric acid HC1O3,
and Perchloric acid HC1O4.
SUG. Student, name the salts these acids form with potassium.
HYPOCHLOKOUS ACID, HC1O.
102. Preparation. — This acid has been prepared only
in dilute aqueous solution. Owing to its instability, the
student must prepare it freshly for the purpose of studying
its properties. It is obtained by treating freshly-precipi-
tated mercuric oxide, HgO, with chlorine water, thus : —
3 HgO + 4 C1 + H2O = 2 HgO, HgCl + 2 HC1O.
EXP. 77 T. Dissolve as much mercuric chloride, HgCl2, as
possible in 250CC hot water ; then add KOH as long as a pre-
cipitate (yellowish-red) is formed. You thus obtain the fresh
mercuric oxide : —
102 THE CHLOKINE OXAC1DS.
2 KOH + HgCl2 = HgO + 2 KC1 + H2O.
Filter out tins precipitate, and wash it by adding much water
to it as it lies upon the filter-paper. Finally make a hole
in the point of the filter-paper, and wash the precipitate through
into a half-litre flask by means of 250CC cold water. Then
gradually add chlorine water, thoroughly shaking meanwhile,
until the remaining brownish-red precipitate ceases to dissolve
(i.e., be careful to keep an excess of HgO. If the chlorine
water be fairly-well saturated, you. will require less than 200CC) .
The remaining precipitate is the compound represented by the
formula HgO, HgCl. Allow the flask to stand in a cool place
until this precipitate settles, when you will be able to pour off
the slightly-colored aqueous solution of hypochlorous acid,
which may be used for experimental purposes. Note the odor
of the acid differing from chlorine.
EXP. 78 P. In a florence flask fitted with a bent deliver}'
tube, generate chlorine gas from sodium chloride, manganese
dioxide, and sulphuric acid. Pass this gas into a cold, dilute
solution of potassium hydroxide, stopping short of saturation.
You will thus obtain for experimental purposes a solution of
potassium hypochlorite, KC1O, thus : —
2 KOH + 2 Cl = KC1O + KC1 + H2O.
103. Properties. — Hypochlorous acid, when in dilute
aqueous solution, is a yellowish liquid, possessing a char-
acteristic odor and strong bleaching properties. A con-
centrated solution cannot be distilled without undergoing
decomposition ; indeed, it soon decomposes at ordinary
temperatures, of its own accord, giving off chlorine and
oxygen gases.
EXP. 79 P. Moisten in dilute hydrochloric acid pieces of
unbleached cotton cloth and suspend them for a moment in the
solutions of hypochlorous acid and potassium liypochlorite, as
prepared above. Finally wash them in pure water, allow them
THE CHLORINE OXACIDS. 103
to dry, and note that they are bleached. Also make this
experiment with a solution of bleaching-powder.
The hypochlorites are of great importance, especially
the calcium compound, which is used in bleaching-factor-
ies under the name of bleaching-powder. 'Enormous quan-
tities of this powder are prepared by passing chlorine gas
into chambers containing slaked lime, Ca(OH)2, thus : —
2 Ca (OH)a + 4 01 = 2 H2O + (CaCl2 + Ca(ClO)2).
It thus appears that bleaching-powder is a mixture of cal-
cium hypochlorite with calcium chloride.
The cloth to be bleached, after a thorough cleansing, is
drawn through a solution of bleaching-powder, and then
through very dilute sulphuric acid, which decomposes
the powder, liberating free chlorine in the fibres of the
cloth. By this means, as previously explained, the color-
ing-matters are destroyed. The effect, upon hypochlorous
acid or the hypochlorites, of stronger acids may be seen,
thus : —
HC1O + HC1 = H2O + 2 Cl,
and KC1O + 2 HC1 = KC1 + H2O + 2 Cl.
104. Tests for Hypochlorous Acid, or the Hypochlo-
rites.— 1. An aqueous solution of the free acid bleaches
litmus paper or solution.
2. The odor of the free acid identifies it.
3. Hypochlorites in solution require acidulating with an
acid, as acetic or hydrochloric acid, before they produce
their bleaching effects.
QUERY. Will a hypochlorite bleach when acidified with HN03? H2SO4?
Try it.
SUG. Carefully distinguish between bleaching a substance and changing
its color, as from blue to red. When it has been bleached, an alkali will
not restore the original color ; when simply changed, the color may thus
be restored.
104 THE CHLORINE OXACIDS.
CHLOROUS ACID, HC1O2.
105. This acid and the salts it forms are unimportant.
As already explained, it may be obtained by dissolving
chlorine trioxide in cold water, but it does not bear con-
centration. It is readily decomposed by heat, as likewise
are its salts, the chlorites. It also possesses bleaching
properties. Its action upon the alkaline bases is very
slow and feeble.
106. Tests for Chlorous Acid and the Chlorites.—
Test as for a hypochlorite, when the same results are
obtained. Then to a fresh portion add a small quantity
of arsenious oxide, As2O3, and a drop or two of nitric acid.
If the solution be that of a hypochlorite, its bleaching pow-
er is destroyed. If that of a chlorite, it will still bleach.
NOTE. This acid and its salts may well be dismissed with simply a
reading of the two preceding paragraphs.
CHLORIC ACID, HC1O3.
107. This acid is also unimportant, and, moreover,
somewhat dangerous to experiment upon ; its prepara-
tion, therefore, should be omitted.
Potassium chlorate, KC1O3, the most important salt of
chloric acid, is made by passing chlorine into a concen-
trated, warm solution of potassium hydroxide, KOH : —
6 C1 + 6 KOH = 5 KC1 + 3 H2O + KC1O3.
QUERY. What takes place when the solution of potassium hydroxide
is cold and dilute ?
In order to get the free acid from this potassium salt,
the latter is treated with a solution of hydrofluo-silicic
acid, H2SiFG:-
2 KC1O3 + H2SiF6 = KgSiFg + 2 HC1O3.
THE CHLOIUNE OXACIDS. 105
The potassium salt thus formed is insoluble ; conse-
quently, after it has subsided, the dilute solution of
chloric acid may be poured off, and afterwards concen-
trated in a vacuum over sulphuric acid.
Concentrated chloric acid is, indeed, a powerful oxidiz-
ing agent, uniting so eagerly with vegetable tissue, as
paper and wood, that it ignites them.
SUG. Student, name the uses of KC103 as suggested by the experi-
ments up to this point.
108. Tests for Chloric Acid and the Chlorates. —
1. Free concentrated chloric acid may be recognized by
its odor and by its charring a slip of paper.
2. The dry chlorates, when treated with strong sul-
phuric acid, }deld a yellowish, explosive gas, C12O4 (see
Exp. 76); with hydrochloric acid, they yield free chlorine
gas. (Exp. 69.)
PERCHLORIC ACID, HC1O4.
109. This acid and its salts are also of but small impor-
tance, and the free acid should not be prepared. It is to
be had by distilling dry potassium perchlorate, KC1O4,
with strong, boiled sulphuric acid. Perchloric acid is
one of the most powerful oxidizing agents known. When
dropped upon charcoal, it explodes with violence, while
dry wood and paper are instantly ignited. Upon the skin
it produces deep and dangerous wounds.
One of its salts, potassium perchlorate, may be prepared
as follows : —
EXP. 80 P. Heat in a generating-flask 5g potassium chlorate,
carefully noting when the oxygen ceases readily to be evolved,
and the mass becomes pasty or semi-solid, —
2 KC1O3 = KC1 + KC1O4 + 2O.
106 EXERCISES IN CHLORINE.
Remove the heat, allow the flask to cool, and dissolve its con-
tents in much hot water. Upon cooling, the potassium per-
chlorate separates out in crystals, while the potassium chloride
remains in solution. These crystals may be removed, dried,
and used for experimental purposes.
110. Tests for Perchlorates. — 1. Dry per chlorates
yield no yellow explosive gas with sulphuric acid, and
with hydrochloric acid yield no free chlorine.
2. They require for their decomposition a higher tem-
perature than the chlorates.
CHLORINE AND NITROGEN.
ill. Chlorine and nitrogen unite to form a dangerous
explosive, which rivals nitro-glycerine, and whose composi-
tion is not definitely known. It is prepared by passing a
current of chlorine through a moderately warm solution of
ammonium chloride. Under no circumstances should the
student thus bring these chemicals together. The eminent
chemists, Dulong, Davy, and Faraday, were seriously
maimed while experimenting with this capricious com-
pound.
EXERCISES IN CHLORINE.
1. Given: NaCl, H2S04, Mn02, HgO, As203, and KOH. From these
chemicals show how you could prepare chlorine and all the compounds
treated in this chapter.
2. PROB. How many tons of salt, NaCl, would it require to prepare
10 tons of hydrochloric acid ?
3. PROB. How many litres of chlorine gas can be obtained from 75k
of NaCl when the barometer reads 755mm and the thermometer 18° C. ?
4. How are acids formed from their anhydrides ?
5. Given: The formula of an acid to determine the formula of its
anhydride. Proceed thus : 2 HN03 — H20 = N205. In a like manner
EXERCISES IN CHLOK1N E. 107
determine the anhydrides of HC1O4, HNO2, H2SO4, HC1O, HC103, and
HI03.
6. What per cent of HC1 is hydrogen ? Chlorine ?
7. Determine the percentages of H, N, and 0 in the nitrogen oxacids.
8. PROB. 201 of Cl, measured at standard temperature and pressure,
increased to 20.51 owing to a fall in the barometer. How many millimetres
did the barometer fall ?
9. Chlorine gas was discovered in 1774. Who was its discoverer? He
used the chemicals HC1 and MnO2. Describe the process, and write the
equation.
10. An aqueous solution of chlorine changes, upon standing, to an
aqueous solution of IIC1. What gas is liberated ? Write the equation.
11. How can you prepare chlorine gas from bleaching-powder ?
12. The water analyst, in determining by titration the amount of
chlorine in drinking-water, proceeds thus: He first prepares a standard
solution of silver nitrate, by dissolving 4.79s AgN03 in I1 of distilled water ;
he then measures out 70CC of the drinking-water, and adds sufficient potas-
sium chromate, K2Cr04, to tinge the water light-yellow. Now, from a
burette graduated to tenths of a cubic centimetre, he adds to the water
thus prepared the standard silver solution, drop by drop, with constant
stirring, until the red color at first formed in the liquid becomes perma-
nent. The number of cubic centimetres silver solution added is equal to
the number of milligrams of chlorine per imperial gallon. How much
silver nitrate does lcc of the standard solution contain ? How much silver ?
Show how this amount of silver will precipitate lm& of chlorine. (Sua.
108m«8 of Ag precipitate 35.5ras* of Cl ; therefore, to precipitate lin« of Cl
requires 108 -=- 35.5 = 3.03mss Ag.) The permanent red color is due to the
formation of silver chromate, Ag.2CrO4; this formation does not occur
until the chlorine is all precipitated. The potassium chromate thus serves
as an indicator, showing when the right amount of AgN03 has been added.
Why does the silver unite with the chlorine first ?
GENERAL NOTK. Recent investigators doubt the existence in a free
state of chlorine trioxide.
CHAPTER VII.
BROMINE, ITS OCCURRENCE, ETC. — THE BROMINE ACIDS.
BROMINE.
SYMBOL BR'. — ATOMIC WEIGHT, 80 ; SPECIFIC GRAVITY, 3.1872.
112. Occurrence. — Bromine does not occur in a free
condition, but is found combined with magnesium, sodium,
potassium, and perhaps with some organic compounds, as
bromides in sea water, certain mineral waters, and in most
saline deposits. It also occurs combined with silver in
the silver mines of Mexico and South America.
Balard, in 1826, discovered bromine in sea water. He
obtained it from the concentrated solution or "mother
liquor" from which the crystals of common salt, NaCl,
had been removed.
Bromine, although by no means a plentiful element, is,
nevertheless, an article of commerce, considerable quanti-
ties of it being produced from the concentrated "mother
liquors " of salt wells in various parts of the world. The
United States produces the greater part of the commercial
article.
113. Preparation. — EXP. 81 p. Dissolve in a test-tube a
crystal of potassium bromide, KBr ; add a small quantity of
chlorine water. Notice that the liquid turns somewhat darker
than the chlorine water added ; this is due to free bromine.
Now add three or four drops of carbon bisulphide, CS2, and
BROMINE. 109
shake thoroughly. What co.lor does the carbon bisulphide
assume ?
EXP. 82 T. Thoroughly mix 30g granulated manganese di-
oxide, MnO2, with 40g potassium bromide, KBr, and place in a
retort. Use the same apparatus as for nitric acid, excepting
that the condenser must contain 0.41 cold water, and the neck
of the retort must dip below the water in the condenser ; or a
rubber cork with a bent tube dipping below the water may be
fitted into the neck of the retort. Now pour into the retort
105g sulphuric acid, H2SO4, previously diluted with 70CC water,
and warm gently. Bromine will distil over in reddish-brown
fumes and condense under the water in the condenser. A part
of the bromine will also be dissolved in the water, thus giving
bromine and bromine water at one operation. Save them both,
each in separate bottles accurately fitted with ground glass
stoppers, and keep in a cool place.
QUERY. The specific gravity of the H2SO4 is 1.84. How many cubic
centimetres equals 10.5s ?
Bromine is best prepared for class purposes by treating
potassium bromide, KBr, with manganese dioxide and
sulphuric acid, thus: —
3 H28O4 + 2 KBr + MnO2 = MnSO4 + 2 HKSO4 + 2 H2O + 2 Br.
SUG. Student compare this equation with that given in Preparation of
Chlorine, Art. 90.
Bromine is liberated when occurring in saline waters,
by adding a small quantity of manganese dioxide, and
then just enough sulphuric acid to liberate sufficient
chlorine to free the bromine. This process depends upon
the fact that free chlorine gas liberates bromine from its
compounds.
QUERY. Which possesses the greater chemism, chlorine or bromine?
110 BROMINE AND HYDROGEN.
114. Properties of Bromine. — Bromine is a dark-red
colored liquid, at ordinary temperatures always giving off
pungent, irritating fumes. It bleaches organic coloring-
matter, but not so powerfully as chlorine. Its principal
use is as a disinfectant.
Bromine has a specific gravity of 3.1872 at 0°, freezes at
- .7.5°, and boils at + 59.3°.
NOTE. When removing the stopper of a bottle containing bromine or
its aqueous solution, always turn your face away. Why ?
115. Tests for Free Bromine. — 1. Free Bromine, even
in dilute solutions, when shaken (in a test-tube) with
carbon disulphide, CS2, colors the latter brownish-red.
2. Colors ether yellowish-red, which color is destroyed
by shaking with potassium hydroxide, KOH.
3. Colors starch-paste solution orange-yellow.
SUG. Student make these tests upon a dilute solution of the bromine
water prepared as above. Also try the bleaching effect as with chlorine.
BROMINE AND HYDROGEN.
116. Hydrobromic Acid, HBr, is the only acid formed
by bromine and hydrogen. It is 'unimportant to the be-
ginner, and he should not attempt its preparation. It is
usually obtained by allowing liquid bromine to act upon
amorphous phosphorus and water. This is accomplished
by placing 10 parts of liquid bromine in a stoppered
funnel provided with a stop-cock to allow the bromine to
fall drop, by drop into a generating-flask containing one
part phosphorus and two parts water : —
p + 5 Br + 4 H2O = H3PO4 + 5 HBr.
What takes place in this reaction can best be understood
BROMINE AND HYDROGEN. Ill
by considering it in two phases. When bromine acts
upon phosphorus the two unite directly, forming either
phosphorus tribromide, PBr3, or the pentabromide, PBr5,
according to the relative quantity of bromine present.
Now each of these compounds is decomposed by water, as
represented in the following equations : —
PBr3 + 3 H2O = H8PO3 + 3 HBr ;
PBr5 + 4 H2O = H3PO4 + 5 HBr.
Thus, in each case, all the bromine appears finally in
combination with hydrogen in the form of hydrobromic
acid.
We should naturally expect that the simplest method
for making hydrobromic acid would be like that used for
making hydrochloric acid, but strong sulphuric acid decom-
poses hydrobromic acid, and hence, although the reaction,
2 KBr + H2SO4 = K2SO4 + 2 HBr,
actually does take place, a further reaction also takes
place, as follows : —
2 HBr + H2SO4 = 2 H2O -f- SO2 + 2 Br,
giving the gas, sulphur dioxide, SO2, and free bromine in
the form of vapor, and from these it is very difficult to
separate the hydrobromic acid.
This acid is a colorless, irritating gas whose chief inter-
est to us lies in the fact that it yields the salts called
bromides, some of which are applied to useful purposes,
thus : silver bromide, AgBr, is used in photography ;
potassium bromide is used in medicine; while others, as
magnesium bromide, MgBr2, are much esteemed ingredi-
ents of certain mineral springs.
Hydrobromic acid is used in organic laboratories, and it
is now an article of commerce.
112 BROMINE AND HYDROGEN.
117. Tests for the Bromides. — 1. Place the solution
in a test-tube, and liberate the bromine by means of
chlorine water; then add a few drops of carbon bisul-
phide, CS2, arid shake thoroughly. The carbon bisulphide
is colored brownish-yellow.
NOTE. An excess of Cl must be avoided, otherwise a chloride of
bromine is formed which does not color the bisulphide.
2. With silver nitrate, AgNO3, this solution gives a
yellowish-ivhite precipitate, AgBr, insoluble in nitric acid,
difficultly soluble in ammonia, and easily soluble in potas-
sium cyanide, KCy.
QUERY. How do HC1 and the chlorides deport themselves with AgN03,
etc.?
NOTE. When bromides and nitrates occur in the same solution, the
tests interfere. The bromine may be readily detected, but not so the nitrates,
since the H2S04 and FeS04 liberate free bromine which obscures the ring.
SUG. Student try a bromide as if testing a nitrate.
118. Bromine, Oxygen, and Hydrogen. — No com-
pounds of bromine and oxygen have been isolated, but
two, and possibly three, acids have been prepared, viz : —
HBrO, Hypobrornons acid,
HBrO3, Bromic acid,
and HBrO4, Perbromic acid.
In regard to the existence of the last there is much doubt.
These acids never occur in nature, and are of small
importance to the beginner, so we shall here notice them
but briefly.
119. Hypobromous Acid, HBrO, may be prepared in
the same way as hypochlorous acid, HC1O, thus : —
HgO -f 4 Br + H,O = HgBr2 + 2 HBrO. (See HC1O.)
BROMINE AND HYDKOGEN. 113
It possesses bleaching powers, is of a straw-yellow color,
and easily breaks up into water, bromine, and oxygen.
QUERY. What salts does this acid form 1
120. Bromic Acid, HBrO3, is formed by treating silver
bromate, AgBrO3, with bromine water, thus: —
5 AgBrO3 + 6 Br + 3 H.O = 5 AgBr + 6 HBrO3.
The salts (bromates) formed by this acid somewhat re-
semble the chlorates in their properties, but are of little
importance commercially.
EXP. 83 P. The easiest way to get a bromate is to dissolve
bromine in a strong solution of potassium hydroxide, when a
mixture of potassium bromate and bromide is formed : —
6 Br + 6 KOH = 5 KBr + KBrO3 + 3 H2O.
The bromate will soon separate out in crystals, which the
student may try as he did the chlorates.
121. Tests for the Bromates. — 1. They are decom-
posed by hydrochloric acid, giving free bromine (which
may be detected as in Art. 115).
QUERY. What effect does HC1 have upon KC103 ?
2. The bromates yield no explosive gas with sulphuric
acid, but they are decomposed, affording free bromine and
oxygen. .
QUERY. How does H2SO4 affect the chlorates ?
EXERCISES IN BROMINE.
1. What chemicals are needed to prepare bromine and its compounds?
2. What per cent of KBr is potassium ? Bromine ?
3. How many grams of NaBr would be required to prepare 10s Br 1
4. Compare the bromine and chlorine acids from a commercial stand-
point.
114 BROMINE AND HYDROGEN.
5. What does the word Bromine signify ?
6. In analyzing a sample of mineral water, a chemist found 0.1678s
bromine per litre. In combining his bases and acids he united this bro-
mine with magnesium. How many grams per litre of magnesium bromide,
MgBr2, did he report ? Ans. 0.1929s.
7. How many grams of chlorine gas would be required to free the
bromine of one gram KBr ? ( KBr + Cl = KC1 + Br. )
8. To precipitate all the bromide in 50CC of a solution, required 0.215s
silver nitrate. How much bromine per litre did the solution contain ?
SUG. 108 parts Ag precipitate 80 parts Br.
9. Try to prepare HBr by passing H2S through bromine water.
Explain the equation, —
H2S + 2Br=2HBr-f S.
Filter the solution and test for HBr. Can you prepare HC1 by passing
H2S through chlorine water ? Write the equation.
10. Boil in an evaporating dish a mixture of solid K2Cr207 and H2S04
until the mixture turns bright red. When cool, place a portion of the
substance thus formed in a test-tube fitted with a bent delivery tube, and
add a solid chloride, as NaCl. Note the bright brownish-red gas evolved : —
4 NaCl + K2Cr207 + 3 H2S04 = 2CrO2Cl, + 2 Na2S04 + K2S04 + 3 H20.
Now lead this gas, which becomes plentiful by applying a gentle heat, into
a test-tube containing a dilute solution of ammonia. Note the yellow
liquid formed : —
2 NH3 + 2 H2O + Cr02Cl2 = (NH4),Cr04 + 2 HC1.
Acidify this solution with acetic acid, and add lead acetate : —
Pb(C2H302)2 + (NH4)2Cr04 = PbCr04 + 2 NH4C2H302.
Note the yellow precipitate thus obtained.
Thus try with KBr instead of NaCl. What results ? Try the same
with a mixture of KBr and NaCl. Do you obtain the same reaction as
with NaCl alone '? How can you distinguish a chloride in presence of a
bromide 1 See Douglas and Prescott, Qual. Anal., p. 159.
CHAPTER VIII.
IODINE. — THE IODINE ACIDS. — SEPARATION OF CHLO-
RIDES, BROMIDES, AND IODIDES. — FLUORINE. — HYDRO-
FLUORIC ACID.
IODINE.
SYMBOL 1'. — ATOMIC WEIGHT, 127; SPECIFIC
GRAVITY, 4.948.
122. Occurrence. — Iodine, like bromine and chlorine,
does not occur free. It is chiefly obtained from sea water,
from which it is taken up by seaweeds. These weeds,
especially on the coasts of Ireland and Scotland, are
washed ashore during storms ; then they are collected,
placed in shallow trenches, dried and burned in thin
layers so that the temperature may not rise high enough
to vaporize the iodides of sodium, potassium, etc., con-
tained in the ashes or kelp, as it is popularly termed.
These iodides are soluble in water, and are removed, by
washing, from the ashes. Small quantities of bromides
are also obtained in this process. Plantations of this sea-
weed are cultivated in some parts of the ocean, and at
the proper times vessels are sent to collect the weed.
Iodine also occurs together with Chili saltpetre in the
form of sodium iodide, Nal, and of late a considerable
quantity of that which comes into the market has been
obtained from this source. It also occurs as silver iodide,
Agl, in certain American silver mines.
116 IODINE.
123. Preparation. — EXP. 84 p. Treat a crystal of potas-
sium iodide, KI, as in Exp. 81. What results do you obtain?
EXP. 85 T. It is not necessary to prepare iodine for class
purposes, since it is an article of commerce, procurable at any
drug store. It may be readily obtained, however, by treating
potassium iodide, KI, with manganese dioxide and sulphuric
acid, as in preparing bromine and chlorine. The iodine vapors
may be condensed in a suitable flask surrounded by a cooling
mixture.
Commercial iodine is prepared from the iodides by
treating them, as above, in iron retorts, when it is
liberated in violet vapors and condensed in black, shining
crystals upon the sides of suitable condensers : —
2 KI + MnO2 + 3 H2SO4 = MnSO4 + 2 HKSO4 + 2 H2O + 2 I.
SUG. Compare this equation with that for bromine and chlorine.
124. Properties. — EXP. 86 p. Heat a small crystal of
iodine in a test-tube. What is the color of the vapor? Note
the odor.
Iodine at ordinary temperatures is a black, shining
solid, possessing a decidedly metallic appearance, and
always giving off fumes of a peculiar odor. When heated,
iodine is easily converted into vapor of a splendid violet
color and characteristic odor. The specific gravity of this
vapor is 8.72.
QUERY. Courtois discovered iodine in the year 1811. He named it
from !oct>8r]s, violet-colored. Why did he thus name it?
Iodine is much used in medicine for various purposes,
especially in reducing swellings, such as goitre and weep-
ing sinews. It is also used in checking the spread of
eruptive diseases, like erysipelas. When thus applied it
is used in the form of a solution prepared by taking, by
IODINE AND HYDROGEN. 117
weight : iodine, 20 parts ; potassium iodide, 30 parts ;
water, 900 parts. Free iodine when brought in contact
with the skin turns it brown.
Iodine is only slightly soluble in water, but easily
soluble in alcohol, carbon bisulphide, chloroform, and in
an aqueous solution of potassium iodide.
125. Tests for Free Iodine. — 1. Free iodine colors
carbon bisulphide, CS2, violet.
2. Colors starch paste blue.
NOTE. Those substances heretofore mentioned as coloring a solution
of starch paste and potassium iodide blue, produce this effect by liberat-
ing iodine which unites with the starch to form a blue substance.
QUERY. What substances act in this way ?
IODINE AND HYDROGEN.
126. Hydriodic Acid, HI, is the only compound of
hydrogen and iodine. This is a colorless gas resembling
hydrochloric acid. It is of no commercial importance, owing
to its instability. Its principal use is for organic work and
as a blow-pipe reagent.
Exp. 87 P. Suspend in a test-tube half full of cold water a
few crystals of iodine. Pass through this solution sufficient
sulphuretted hydrogen, H2S (Art. 165), to decolorize it. Hy-
driodic acid will be formed and sulphur deposited : —
H8S + 2 I = S + 2 HI.
The sulphur will soon subside, and the clear solution of the acid
maybe poured off. Reserve this solution for the next experiment.
Hydriodic acid may also be made by the method
which was described under hydrobromic acid ; that is, by
gradually adding iodine to amorphous phosphorus under
118 IODINE AND HYDROGEN.
water. The reactions are the same as in the case of
bromine, the iodides of phosphorus being first formed,
but afterwards decomposed by the water.
EXP. 88 P. In one test-tube place a solution of mercuric
chloride, HgCU ; in another, a solution of silver nitrate, AgNO3 ;
in a third, a solution of lead acetate, Pb(C2H3O9)2. To each of
these now add a portion of the hydriodic acid solution prepared
as above. Note the brilliantly colored precipitates which are
respectively the iodides of mercury, silver, and lead. Repeat
the experiment, using a solution of potassium iodide, KI, in
place of the acid. Do you obtain the same results?
Hydriodic acid unites with bases to form the iodides,
many of which are valuable. Some of these iodides
possess very bright and distinctive colors which are of
service in identifying some of the metals whose salts are
in the solution to be analyzed. Since the acid itself is
unstable and somewhat troublesome to prepare, the chem-
ist preferably uses a solution of potassium iodide for this
purpose.
Sue. Explain these equations : —
Pb(C2H302)2 + 2 HI = PbI2 + 2 HC2H302 ;
AgNO8 + HI = Agl + HN03.
Also write the same equations with KI in place of HI.
127. Tests for Hydriodic Acid or the Iodides. — 1. To
the solution add chlorine water. Then add a few drops
of carbon bisulphide and shake. Iodine is freed and colors
the bisulphide violet.
NOTE. If the iodide is not readily soluble, the iodine may be freed by
warming the insoluble iodide in a test-tube with a crystal of potassium
chlorate and hydrochloric acid; the bisulphide may then be directly
added.
IODINE AND HYDROGEN. 119
2. With silver nitrate, AgNO8, a yellow precipitate is
given, insoluble in nitric acid; sparingly soluble in am-
monia; soluble in potassium cyanide, KCy.
SUG. Compare this test with the similar ones for chlorine and bromine.
QUERY. Which test in the case of chlorine is distinctive ?' Of bro-
mine ? Of iodine ?
NOTE. Solutions of iodides and nitrates do not readily yield a test for
nitric acid for the same reasons as those given under bromine. The test
for the iodide is readily obtained. Try a solution of KI as for a nitrate.
Try a solution of KI and KBr with CS2, etc. Which test do you obtain ?
128. Detection of Chlorides, Bromides, and Iodides
in the same Solution. — The student is probably aware
that the precipitates obtained with silver nitrate do not
afford sufficiently marked characteristics to distinguish
these compounds, and that the carbon bisulphide tests also
fail, especially in the case of bromides in presence of
iodides. To separate and distinguish these substances is
not an easy task, and of the many ways proposed, none
are entirely satisfactory and at the same time simple and
convenient. The following method requires careful ma-
nipulation.
f
EXP. 89 P. Let us suppose the solution to contain NaCl,
KBr, and KI. Divide it in three portions and add to numbers
1 and 2 an excess of silver nitrate, when the precipitates ob-
tained in each will consist of AgCl, AgBr, and Agi. Filter
out these precipitates and wash them thoroughly with hot water,
then wash them through a hole in the point of the filter-paper
into separate beakers. To the first beaker now cautiously add
hut two or three drops of potassium bromide, and to the second
carefully add three or four drops of potassium iodide, and boil
for a short time. Again filter the contents of the first beaker,
and test the clear liquid which runs through for chlorides,
120 IODINE, OXYGEN, AND HYDROGEN.
Art. 96. Filter the contents of the second beaker, and test the
solution for bromides (Art. 117, 1).
Try a part of the third portion directly for iodides by Art.
127, 1. In case you do not succeed, proceed thus : To the re-
mainder of the third add a few drops of ferrous sulphate,
FeSO4, ^nd copper sulphate, CuSO4, when a light-green pre-
cipitate of cuprous iodide, Cu2T2, will be thrown down. Test
this insoluble iodide by Art. 127, 1, Note.
EXPLANATION. What occurs in the three cases may thus be ex-
plained : —
1. In number 1, AgCl + AgBr + Agl + KBr = Agl + 2 AgBr -f KC1.
The KBr and AgCl react, yielding KC1, which is soluble and in the solution
tested for chlorides.
2. AgCl + AgBr + Agl + 2 KI = 3 Agl + KBr + KCL
The KC1 and KBr are soluble and readily yield the test for bromides.
3. This last is readily understood when we remember that the iodine is
partially precipitated in the Cu2Ir
NOTE. The foregoing method is not sufficiently accurate for quanti-
tative determinations where an excess of potassium bromide or iodide
would necessarily be employed. Care must be used to avoid an excess
of either reagent when employed for qualitative work.
IODINE A!*D OXYGEN.
129. There is but one known oxide of iodine, IODINE
PENTOXIDE, I2O5. This oxide may be obtained by heat-
ing iodic acid, HIO3 as described in the next article.
IODINE, OXYGEN, AND HYDROGEN.
130. There are but two oxygen acids of iodine, viz : —
IODIC ACID, HIO3,
Periodic Acid, HIO4.
IODINE, OXYGEN, AND HYDROGEN. 121
These acids and their salts are unimportant; we shall
therefore notice only the first, and that but briefly.
EXP. 90 OP. Heat one part, by weight, of free iodine with
ten parts strong nitric acid (sp. grav. 1.5) until red fumes cease
to come off and the iodine is dissolved. Evaporate the solution
to dryness and heat in the air-bath to 200°. The resulting
white powder is iodine pentoxide.
The first product formed in this process is iodic acid,
HIO3. When this acid is heated to 200° it breaks up into
water and iodine pentoxide : —
By again dissolving the pentoxide in water, pure iodic
acid may be obtained.
SUG. Write the equation for the action of iodine on nitric acid,
remembering that NO, HI03, and H20 are formed. Also show the action
of H2O on I2O5.
Iodic acid rapidly oxidizes organic substances. When
this acid or the pentoxide is heated with powdered char-
coal, phosphorus, sulphur, etc., it oxidizes them so rapidly
that the action is accompanied by flame.
It forms normal salts, the iodates, as KIO3. Acid salts,
as KIO3HIO3 or HK(IO3)2, are also known.
131. Tests for Iodic Acid or the Iodates. - - To a
solution containing either the free acid or its salts add
starch paste and chlorine water ; no change in color
occurs. Now add a solution of sodium sulphite, Na2SO3,
when iodine is liberated and the solution turns blue.
122 FLUORINE.
FLUORINE.
SYMBOL, F'. — ATOMIC WEIGHT, 19; SPECIFIC
GRAVITY, UNKNOWN.
132. Occurrence. — Free fluorine is unknown. It occurs
combined with calcium as calcium fluoride, CaF2, or fluor
spar, in cubical crystals which are usually somewhat trans-
lucent and often quite transparent. It also occurs in
the mineral cryolite, which is a fluoride of sodium and
aluminium. Other sources of fluorine are unimportant.
Fluorine has resisted all attempts to isolate it, and they
have been many. This fact appears to be due to its great
chemism when nascent, at which time, it invariably
attacks and combines with the vessel in which it is gener-
ated. Nothing is known of its physical properties, and
but little of its chemical deportment other than its great
chemism.
Fluorine forms no oxides, no oxygen acids, and but one
hydrogen acid, viz. : —
HYDROFLUORIC ACID, HF.
133. Preparation. — This acid is also a gas correspond-
ing to hydrochloric, hydrobromic, or hydriodic acid. It
is best prepared by treating calcium fluoride in a leaden
evaporating-dish, with sulphuric acid : —
CaF2 + H2SO4 = CaSO4 + 2 HF.
This gas is a dangerous poison, and great care must be
exercised in its preparation.
134. Properties. — EXP. 9 IT. Pulverize 4g calcium fluor-
ide, and place in a leaden dish, which can be made by cutting
FLUOttlNE. 123
off a piece of lead pipe, splitting it open lengthwise, and then
placing it in an iron mortar where it can, by the aid of an iron
pestle, be hammered out into the shape of an evaporating-dish.
Next prepare a sheet of glass by coating both sides with
beeswax or paraffin. Upon one side of this glass engrave, by
means of a pin or sharp, soft wire, some design. Now put
the evaporating-dish, supported by a ring-stand, in a gas-
chamber or where there is a current of air to carry off all
fumes, and support the plate a short distance above the dish.
Add strong sulphuric acid to the calcium fluoride, when hydro-
fluoric acid will be quickly liberated, especially if a gentle heat
be cautiously applied. In a few minutes the design will be
neatly etched into the glass. Be very careful not to inhale any
hydrofluoric acid fumes, as they are exceedingly poisonous.
Hydrofluoric acid is often employed as above in etching
thermometer scales.
This acid seems to have great chemism for such sub-
stances as calcium, silicon, and potassium, in consequence
of which glass is immediately attacked and can not be
used to store the gas or its aqueous solution. Leaden or
vulcanite bottles are employed for this purpose.
The action of hydrofluoric acid upon sand and glass,
which is a compound of sand with bases, is largely due to
the action represented by the equation , —
SiO2 + 4 HF = 2 H2O + SiF4.
The silicon tetrafluoride, SiF4, thus formed, escapes as a
gas.
135. Tests for Hydrofluoric Acid in Fluorides. — The
best is the etching test, but care must be taken not to
scratch the glass with the graver used in cutting through
the wax.
124 EXERCISES IN IODINE AND FLUORINE.
EXERCISES IN IODINE AND FLUORINE.
1. How many grams of silver nitrate would be required exactly to
combine with 10« of potassium iodide ?
2. How many pounds of iodine can be obtained from one-half ton of
sodium iodide ?
3. What chemicals are necessary to prepare from potassium iodide
iodine and its compounds 1
4. Make a comparison between the commercial values of the acids of
chlorine, bromine, and iodine.
5. Compare the same three elements according to their physical con-
ditions at ordinary temperatures ; also according to their atomic weights,
specific gravities, and chemism. Make a table comprising the acids they
form.
6. Will nitro-hydrochloric acid liberate bromine and iodine from their
compounds ? Try it.
7. Class prepare a sheet of glass as directed in Exp. 91, writing the
names of the class through the wax. Under the teacher's direction etch
with HF, This will be a good memento to leave in the Laboratory.
8. Test a solution of NaCl and KN03 for the different acids combined
with bases in these salts.
9. Under potassium and sodium learn the tests for these metals, and
try for them in the above solution.
10. It would now be well for the student to practise daily upon
unknown solutions, as in 8 and 9. These solutions should not contain acids
that interfere, and the bases with which the acids are combined should
preferably be potassium, sodium, and ammonium.
11. In working upon an unknown solution a student obtained tests
for K, Na, NH3 and H2S04, HC1, and HNO3. What salts may have been
dissolved in the solution ? In case the laboratory contains only NH4N03,
NH4C1, KN03, KC1, NaCl, and K2S04, what salts may the teacher have
employed in preparing this solution ?
12. See Trans. Roy. Soc. Canada, 1883, sec. 3, pp. 65 et seq., for " Hy-
driodic Acid as a Blow-pipe Reagent." Dr. Haanel's paper on this topic
is accompanied by very fine plates.
13. To a solution containing an iodide and a bromide add CS2 ; now
by the addition of sufficient chlorine water try to obtain first the color for
iodine, and second the color for bromine. Explain.
CHAPTER IX.
CAEBON. — CARBON AND HYDROGEN. — OXIDES OF
CARBON. — CARBONIC ACID. — CYANOGEN. — PRTJSSIC
ACID.
CARBON.
SYMBOL Civ. — ATOMIC WEIGHT, 12. — SPECIFIC GRAVITY: DIA-
MOND, 3.5-.G ; GRAPHITE, 2.25; CHARCOAL, 1.57.
136. Occurrence. — Carbon is a very widely distributed
element, occuring chiefly in an impure state or in chemical
compounds. It is an important constituent of all organic
substances, mineral carbonates, carbonic acid gas, and the
cyanides. In a free condition, it exists in three widely
differing forms.
1. In pure, transparent, glittering, octahedral crystals,
as Diamonds, which are found in earthy detritus or clayey
shales in Africa, South America, Australia, and other
localities.
SUG. Write an essay on diamond-mining, diamond-cutting, and famous
diamonds.
2. In dark, shining, six-sided slabs as Graphite, Plum-
bago, or Black Lead, which occurs in England, Ceylon,
the United States, and other countries.
3. In impure forms as Coal, Soot, and Lamp-black. Of
coal we find a number of varieties, as Charcoal, Anthracite
coal, Bituminous coal, etc.
SUG. Write a short paper on coal-mining.
126 CARBON.
137. Preparation. — Tt is not necessary to prepare car-
bon for class illustration, since any of the above-named
modifications are easily to be obtained.
Small diamonds are said to have been made artificially
by a somewhat complicated process, which cannot be
profitably described at this stage. The method of their
formation in nature is not understood.
Crraphite has been frequently observed in iron-smelting
furnaces, having been artificially produced at high tem-
peratures.
In Exp. 2, Charcoal was obtained by heating wood in a
test-tube. The principles therein involved are made use
of in preparing charcoal for commerce. In practice the
wood is heated in closed iron cylinders, or burned in
large pits or kilns with a limited supply of air. In the
latter case a part of the wood thus treated is completely
consumed in order to furnish the heat requisite for charring
the remainder.
Lamp-black or soot is prepared by burning a carbonaceous
substance, such as oil, resin, etc., in a limited supply of air.
The lamp-black appears as a black smoke which is easily
collected upon a cold surface.
QUERIES. What makes a lamp smoke ? Why is the lamp-chimney
blackened 1 Explain the deposition of soot in stovepipes and chimneys.
Can soot be obtained from the Bunsen flame ? Luminous flame ? Alcohol
flame ? Ordinary candle flame ? Why does pitch pine give such a smoky
flame 1 If one wishes to know a fact which comes within the province of
Experiment, how should he proceed ?
138. Properties. — Carbon is absolutely indispensable
to all organic structures. With other elements, such as
hydrogen, oxygen, and nitrogen, it is capable of forming
an almost endless number of chemical compounds. As a
matter of convenience these are generally considered
CARBON. 127
under the head of the Chemistry of the Compounds of
Carbon, or ORGANIC CHEMISTRY.
Carbon has many industrial uses. It is chiefly used in
reducing metals from their ores and for heating and
illuminating purposes.
The colorless diamond is highly prized as a jewel. A
colored variety is used in glass-cutting, while its dust is
employed for polishing hard and refractory substances.
Drills armed with diamond points are used by miners and
others ; these drills will quickly cut through the hardest
rocks. Smoky or black diamonds and carbonado, an
impure massive form, are principally used for this latter
purpose.
The diamond is the hardest substance known, its value
in the " scale of hardness," by which mineralogists estimate
the hardness of minerals, being 10°. This scale, in which
each substance is able to scratch all that are below it in
the scale, is as follows : —
Diamond 10 5
Sapphire 9°
Topaz 8°
Quartz 7°
Feldspar 6°
Apatite oc
Fluorspar 4C
Calcspar 3C
Gypsum 2C
Talc . lc
The primary form of a diamond crystal is octahedral ;
but it occurs in many different forms derived from this
primary crystal. When first removed from its matrix, the
diamond is often rough and lustreless, and afterwards
requires cutting and polishing ; this latter is accomplished
by means of its own dust. Like all hard substances it is
brittle and quite easily broken. In acids and alkalies the
diamond is completely insoluble. When heated to a high
temperature in a current of oxygen it burns, the product
being carbon-dioxide gas, CO2, with a small amount of
128 CARBON.
residual ash. Upon light the diamond exerts a very high
refractive influence, to which property it owes its great
brilliancy.
QUERY. What properties cause the diamond to be so highly esteemed
as a jewel ?
Graphite is greasy to the touch. It is largely used for
polishing purposes, such as for coating shot and powder,
and, owing to its great permanence in the air, is largely
employed in the manufacture of stove-polish. Its particles,
however, are very hard, arid the saws used in cutting it are
quickly worn out, and a knife, when employed for the same
purpose, soon loses its edge.
Graphite, owing to its great infusibility, is now mixed
with clay and extensively used in making crucibles which
are employed by metallurgists, while its employment in
the manufacture of leads for the common lead-pencil is
a well-known application.
SUG. Prepare a paper on the manufacture of lead-pencils.
Coal is probably the remains of a magnificent vegeta-
tion which flourished during the carboniferous age. It
has been brought into its present condition by heat and
pressure. The heat is thought to have been supplied by
the heated interior of the earth, while the pressure was
due to the influence of water and the rocks which subse-
quently formed above the coal. This explanation con-
templates the idea that during some post-carboniferous
convulsion which swept over the globe, the land sank
down, and the vegetation was overwhelmed by the inrush
of water, while the rocks were afterward deposited. The
ashes and " clinkers " of burned coal are the mineral
sediments which were entangled by the vegetation, as
well as the mineral constituents of the plants themselves.
CARBON. 129
Anthracite coal is used for heating purposes, and for
reducing metals from their ores. Its reducing power
depends upon the chemism of carbon for oxygen.
QUERY. What is meant by reduction ?
Bituminous coal differs from anthracite in that the
former contains more hydrogen-carbon compounds, and
evidently has not been subjected to so high a temperature
or to so great a pressure by natural agencies. This variety
of coal burns with a very hot arid sooty flame, and needs a
large supply of air for its combustion.
Coke is a form of carbon obtained by driving off, at a
high temperature, the volatile constituents of coking-
coal. It is left behind in the retorts when coal is distilled
for the purpose of making illuminating gas.
Gas Carbon is also produced in distilling coal. This
form of carbon is much used in making negative plates
for batteries and for the terminals of electric lamps.
Peat is a form of fuel nearly akin to bituminous coal,
and is formed from the roots and stems of certain plants
growing in bogs or marshes.
Lignite is a peculiar form of coal formed from such
sources as our present deciduous trees, and often exhibits
a distinctly woody structure.
Jet is a black variety of lignite, much used in jewelry.
Jet readily takes a high polish.
Lampblack is much used as a paint, and in making
printers' ink.
Charcoal is employed as a reducing agent in preparing
iron from its ores.
QUERIES. For what purposes do you use charcoal in the laboratory ?
What class of artisans employ charcoal ? What other common uses does
it have ?
130 CARBON.
Charcoal possesses some remarkable properties : —
EXP. 92 p. Place a filter-paper in a funnel ; then fill the
paper nearly full of bone-black or freshly-burned charcoal
powder. With a filter thus arranged see if you can produce
any changes in the following solutions by filtering them several
times: 1. Vinegar; 2. Syrup of brown sugar; 3. Dilute black
molasses ; 4. Indigo solution ; 5. Carmine solution ; 6. Beer ;
7. Potassium dichromate solution.
QUERIES. What changes occurred ? Does 7 behave like the others ?
Why ? Explain the changes.
We thus see that charcoal is capable of decolorizing
and purifying such organic liquids as were mentioned. The
reason why it is employed in filtering drinking-water is
now apparent. It is supposed that this action of charcoal
is partially due to the fact that it absorbs oxygen, and
possesses the power of causing certain organic substances
to combine with this oxygen. However this may be,
the charcoal soon loses its efficacy unless it be frequently
washed and exposed to the air. For the same reason
charcoal will destroy the gases from putrefying sub-
stances.
QUERIES. Why should a filter be frequently cleaned ? Is it best con-
tinuously to keep a filter full of water ? What is the use of gravel in
filters ? Why should a rapid river flowing over stones and with numerous
falls be purer than one with a sluggish current and a sandy or muddy
bottom ?
EXP. 93 P. Place in an evaporating-dish a few grains of
common sugar ; add a few drops of strong sulphuric acid. Do
you obtain carbon ? Also thus try starch. What results?
From the above experiment and from previous work
the student may learn that many substances, such as
sugar, oils, resins, fats, waxes, tallow, and alcohol are
CAEBON. 131
compounds of carbon. We may add to this list nearly
every substance used as food by man and by animals, and
all the vegetable drugs known to chemistry and com-
merce. We should not forget also that it is to the
compounds of carbon that we are indebted for our rai-
ment, and even for a portion of our dwellings.
QUERY. How could we obtain light without the aid of carbon ?
Kerosene, gasoline, naphtha, benzine, and paraffin are
all derived from PETROLEUM, or rock oil, which is a
mixture of many compounds of carbon and hydrogen
found in company with coal deposits. The limits of our
work forbid a further notice of these interesting sub-
stances.
139. Tests for Carbon. — 1. Free carbon, as soot, coal,
lampblack, etc., may be recognized by its physical proper-
ties and by its insolubility in all acids and alkalies ; also
by the manner in which it burns when heated on platinum
foil.
2. Graphite may be recognized by its properties, and by
the black, insoluble streak which it leaves when drawn
across paper.
SUG. Write with a lead-pencil on white paper. Try to bleach it. What
results ?
3. The diamond is recognized by its brilliancy and
hardness, being able to produce a scratch upon the hard-
est substance.
QUERY. The hardness of glass is less than 6°. Is the fact that a
given substance makes a scratch upon glass sufficient evidence that it is a
diamond 1
CARBON AND HYDKOGEN.
CARBON AND HYDROGEN.
140. Carbon and Hydrogen form many compounds,
but three of which we shall notice here : —
1. Methane, or Marsh Gas, CH4 ;
2. Ethylene, or Olefiant Gas, C2H, ;
3. Acetylene, C2H2.
METHANE, CH4.
141. Methane, or Marsh Gas, may thus be prepared for
illustration : —
EXP. 94 p. 2g sodium acetate, NaC2H3O2, are heated in a
hard glass test-tube fitted with a jet, with 8g sodium hydroxide,
NaOH, and 2g finely-powdered quick-lime, CaO. As soon as
the gas issues freely from the jet it may be ignited, when it
burns with a bluish-yellow, non-luminous flame. The reaction
NaC2H3O2 + NaOH = Na2CO3 + CH4.
QUERY. What purpose does the CaO serve 1 (Compare the use of
Mn02 in producing oxygen from KC103.)
This gas occurs free in nature, and is formed in stagnant
pools by the decay of leaves and other vegetable material,
whence it derives its name, Marsh Gas. It also occurs in
coal seams and in coal mines, where it is known as Fire
Damp. Methane condenses at —10° under 50 atmospheres,
and boils at — 160° under 1 atmosphere.
EXP. 95 P. Discharge the hydrogen pistol by means of a
mixture of marsh gas and air.
When mixed with air or oxygen, methane is often the
cause of most violent explosions. To prevent these ex-
plosions, Sir Humphrey Davy invented his Safety Lamp,
which consists of an ordinary lamp, the flame of which is
CARBON AND HYDROGEN. 133
surrounded with a wire-gauze cage. This cage prevents the
temperature of the surrounding mixture of methane and
air from rising to the point of ignition. The specific
gravity of methane is 0.558.
QUERY. Why does the wire gauze placed between the Bunsen flame
and chemical vessels prevent them from breaking ?
SUG. Student ascertain the particulars of several noted colliery ex-
plosions.
EXP. 96 p. Hold moistened strips of red and blue litmus
paper in a jet of methane. The gas does not affect them.
We are thus led to the conclusion that methane does
not resemble either the acid or the alkaline gases already
studied. These compounds of carbon and hydrogen differ
in many respects from the compounds of other elements
with hydrogen. A very large number of the hydrogen-
carbon compounds is known and new ones are being con-
stantly discovered. We may regard as derived from these
the compounds treated in organic chemistry.
Since methane is not readily acted upon by reagents, the
color of its flame, and its explosiveness when mixed with
air, will answer our purposes as tests.
ETHYLENE, OR OLEFIANT GAS, C2H4.
142. Ethylene is formed in distilling coal, and is, there-
fore, a constituent of coal gas. It is prepared most readily
by the following method, which may be shown for class
illustration : —
EXP. 97T. Heat in a generating-flask fitted with a jet
delivery-tube 10g of ethyl alcohol, C2H6O, with 50g strong
sulphuric acid. Note the odor and taste of the gas issuing
from the jet, and then ignite it. It burns with the ordinary
134 CARBON AND HYDROGEN.
gas-flame. The sulphuric acid simply abstracts one molecule
of water from the alcohol, thus : —
C2H60 = H20 + C2H4.
QUERY. How many cubic centimeters of alcohol and acid are required
above, the specific gravity of H2SO4 being 1.843 and that of ordinary
alcohol being 0.815 1
Ethylene is explosive when mixed with three times its
volume of oxygen.
QUERY. What substances are formed ? Student write the equation.
When equal volumes of ethylene and chlorine gases are
brought together, an oily liquid, called " Dutch Liquid,"
C2H4C12, the odor of which resembles chloroform, is formed.
The specific gravity of ethylene is 0.9784 ; it can be
condensed to a liquid at 10° by a pressure of 51 atmos-
pheres, and boiling under 1 atmosphere at — 100°.
143. Test for Ethylene. — Fill a jar with the gas sup-
posed to contain ethylene ; then pass a current of chlorine
gas into the jar. If the oily Dutch Liquid mentioned
above be formed, ethylene is present.
NOTE. This liquid is insoluble in water.
ACETYLENE, C2H2,
144. Acetylene is also a gas, and possesses a powerful
and disagreeable odor, which is particularly noticeable
when an ordinary Bunsen burner strikes back and con-
tinues to burn at the base.
It has been prepared by passing sparks from a powerful
battery through an atmosphere of hydrogen, the termi-
nals of the electrodes being carbon. No other hydro-
carbon compound has been thus directly produced. It
CARBON AND HYDROGEN. 135
burns with a bright, luminous flame, and has a specific
gravity of 0.92. The odor of acetylene betrays its
presence.
145. Illuminating- Gas is obtained, together with many
bye-products, *by distilling coal in retorts. It contains
hydrogen, methane, and ethylene, and many other hydro-
carbon compounds. It also contains in small quantities
the impurities: ammonia; hydrogen-sulphide, H2S; carbon
dioxide, CO2; carbon monoxide, CO; atmospheric oxygen;
and nitrogen. These impurities are mostly removed by
passing the gas through a series of washing and absorbing
reagents.
SUG. Student visit the gas works. Write a description of the process
of gas manufacture. Consult R. and S.
The student may test for these impurities thus : —
1. Ammonia is detected by holding a strip of moistened
faintly-red litmus paper in a stream of the illuminating gas.
The paper turns blue if ammonia be present.
2. Hydrogen sulphide will blacken a strip of bibulous paper
moistened with lead acetate, Pb(C2H3O2)2, when the paper is
held in a current of the gas.
3. Carbon dioxide may be detected by shaking lime-water,
Ca(OH)2, in a flask of the gas (see test for CO2).
4. Oxygen may be detected as directed under tests for
oxygen. Art. 29, 2.
5. The nitrogen and carbon monoxide cannot be detected with
certainty by an}' means likely to be at the beginner's disposal.
Coal Tar. — It has been mentioned that there are many
bye-products formed in distilling coal in the manufacture
of illuminating gas: of these coal tar is, from a chemical
standpoint, the most remarkable. It is used directly for
136 CARBON AND OXYGEN.
various industrial purposes which are so well known as
to need no description. The attention of many chemists
has been given to this substance, and from it they have
produced a large number of articles which are in daily
use in the arts and manufactures. The beautiful aniline
anthracine and naphthaline dyes are obtained from this
source, and their production has revolutionized not only
the art of dyeing, but also the industries of whole coun-
tries, and made it possible for even the laborer to em-
bellish his home with colors which before were only
accessible to the opulent. From coal tar, then, we may
see the artificial production of substances which formerly
were only obtained from natural sources ; and thus is the
distinction between the so-called organic and inorganic
substances rapidly passing away.
SUG. The student who reads German may obtain valuable information
upon this topic by consulting Schultz's Chemie des SteinJcohlentheers.
QUERY. What are the natural sources of indigo 1 Cochineal ?
CARBON AND OXYGEN.
146. There are two oxides of carbon, viz : —
1. Carbon Monoxide, CO ;
2. Carbon Dioxide, CO2.
Of these two oxides the latter is to us of the greater im-
portance. Both are gases under ordinary conditions.
CARBON MONOXIDE, CO.
147. Preparation, etc. — This gas is a product of com-
bustion, and is formed when carbon is burned in a limited
supply of oxygen : —
C + O = CO.
CARBON AND OXYGEN. 137
It is also formed at high temperatures by the action of
carbon on carbon dioxide : —
CO2 + C = 2 CO.
QUERY. Of what kind of action is this an example ?
EXP. 98 P. Carefully heat in a generating-flask with a de-
livery-tube, 2g potassium ferrocyanide, K4FeCy6, with 20g strong
sulphuric acid. Ignite the stream of escaping gas, CO ; care-
fully note its odor, if any, and the color of the flame.
Carbon monoxide burns with a lambent blue flame, as
seen in coal stoves when the supply of air is limited, and
at the upper surface of the coal in grate fires. The com-
bustion at the bottom of the coal first produces carbon
dioxide ; this substance coming in contact with the heated
coal near the upper surface is reduced to carbon monoxide;
and when this latter meets the air above the coal, it again
burns, forming carbon dioxide, the combustion now being
complete : —
CO + O = C02.
Carbon monoxide is colorless and tasteless, and has a
faint and peculiar odor. It acts upon the animal economy
as a deadly poison, producing headache, giddiness, and
insensibility. It seems to produce its effects upon the
system by combining with the haemoglobin of the blood,
leaving traces which betray its action even after death.
Great care should be taken not to allow this poisonous
gas to accumulate in rooms warmed by coal fires. One
per cent is a sufficient quantity to prove fatal. The joints
of the stove should be tight, the draft strong, and, above
all, the ventilation should be perfect. Death has been
produced from warming poorly-ventilated rooms by means
of charcoal fires in open vessels from which carbon mon-
138 CARBON AND OXYGEN.
oxide is given off; and people have perished by going to
sleep beside a lime or brick kiln or a charcoal pit, being
suffocated and poisoned by the gaseous oxides of carbon.
Tobacco smoke contains more or less carbon monoxide.
Hence, the inhalation of the air of a room in which many
persons are smoking may produce pernicious effects upon
the system.
Carbon monoxide has a specific gravity of 0.968, and
condenses at — 139.5°, under a pressure of 35.5 atmos-
pheres ; under 1 atmosphere it boils at — 190°.
148. Test for Carbon Monoxide. — This gas may be
recognized, when present in sufficient quantity, by its
bluish flame.
CARBON DIOXIDE, CO2.
149. Occurrence. — This gas, commonly known as car-
bonic acid gas, occurs widely distributed in nature. It
occurs free in the atmosphere in small but persistent
quantities, and combined in all the carbonates, from which
it is readily liberated by the stronger acids. Calcium
carbonate, or limestone, CaCO3, is a very plentiful sub-
stance. Whole geological formations consist of this
material. It also is the chief constituent of shells and
most corals. Whole islands are being constantly built
up by the corals in the tropical regions.
SUG. Write a paper on coral formations.
150. Preparation. — EXP. 99 p. In a wide test-tube or
a small beaker place about 5CC calcium hydroxide solution,
Ca (OH) 2. By means of a small glass tube force air from the lungs
through the solution, when a white precipitate will be formed.
Continue to breathe some minutes through the liquid ; the pre-
cipitate dissolves.
CARBON AND OXYGEN. 139
This white precipitate is calcium carbonate, CaCO3,
and was produced by the action of the carbon dioxide
which is thrown out of the lungs as a waste product at
every respiration : —
Ca(OH)2 + CO2 = CaCO3 + H.O.
Large quantities of carbon dioxide must thus neces-
sarily be liberated in the air, since it is produced in the
same way by all air-breathing animals.
QUERY. Why does the air in poorly-ventilated living-rooms contain
more carbon dioxide than those that have good ventilation ?
EXP. 100 P. Carefully lower into a wide-mouth bottle a
burning taper. When the taper is extinguished, add a small
quantity of calcium hydroxide ; cork the bottle, and shake.
Do you again obtain the white precipitate ?
All carbon compounds when burning in the air produce
carbon dioxide. This gas is also emitted during volcanic
action.
QUERY. In what ways may C02 be liberated in living-rooms ?
EXP. 101 P. To a dilute solution of sugar or molasses in
water add a little bakers' yeast. Place in an evaporating-dish
a small quantity of this solution ; also fill a test-tube with the
solution, invert the tube, and place its mouth below the solution
in the evaporating-dish. The whole is now to be left standing
in a warm place. Fermentation soon begins, bubbles of gas
rise in the tube, and the liquid is forced down. When the tube
is full of gas, pour the latter out into another tube (as if it were
water), add calcium hydroxide, and shake as before. Is the
gas carbon dioxide ?
Carbon dioxide is also produced in fermentation, and
iu the spontaneous decomposition of animal and vegetable
substances.
QUERY. In what ways is carbon dioxide liberated in the atmosphere 1
140 CARBON AND OXYGEN.
EXP. 102 P. Break into pieces about 10g calcium carbonate,
or marble, CaCO3. Place in a generating-flask, and cover with
water. Fit the flask with a V-shaped delivery-tube, and collect
the materials mentioned in the following experiments. Upon
adding hydrochloric acid to the contents of the flask, carbon
dioxide will be plentifully given off, although a gentle heat
may sometimes be required. The equation is : —
CaCO3 + 2 HC1 = CaCl2 + H2O + CO2.
NOTE. The CaCl2 solution should be evaporated to dryness, fused in
a sand crucible, and kept in a tightly-corked bottle. It is useful for
drying gases and for other purposes.
Carbon dioxide may be readily obtained in larger quan-
tities by treating the carbonates with strong acids. With
the gas which the student is now ready to prepare he may
proceed to study the
151. Properties. — EXP. 103 p. Fill a wide test-tube with
carbon dioxide. Note the odor and color, if any, and try the
effect upon a glowing match ; a burning match ; a lighted
taper. What results? Tr}' to ignite a jet of this gas.
Carbon dioxide is a colorless, odorless gas which does
not support combustion. Advantage has been taken of
this fact in making an engine to extinguish fires. The
gas is generated from sodium carbonate, Na2CO3, and sul-
phuric acid, and allowed to escape through a hose.
SUG. Explain the construction of a Babcock fire extinguisher. Write the
equation for the reaction of Na2C03 and H2S04. For HNaCO2 and H2S04.
EXP. 104 P. In the centre of a pine ruler 2cm wide and
100cm long drive two needles. This ruler will serve as the
beam of a balance, while the needle-points will serve instead of
a knife-edge bearing. These points are to be placed upon a
flat metallic surface. Now from one end of the beam suspend,
by means of a thread, a small paper sack, and from the other
CARBON AND OXYGEN. 141
end a larger paper sack. Into the smaller sack carefully drop
small pieces of iron, chalk, sand, or any heavy substance, until
the beam is in equilibrium. Into the larger sack now deliver a
jet of carbon dioxide. The larger sack will soon become
heavier and sink.
QUERIES. With what gas was the larger sack filled before introducing
the C02 ? Is C02 lighter or heavier than air ? Suppose the sack be sus-
pended mouth downwards, what would occur if a jet of hydrogen were
allowed to flow up into it ? In what other way have you compared the
weight of air with that of gases ? How should a jar be placed when filling
it with C02, mouth down or up ?
EXP. 105 P. Place a lighted taper in an open jar of air.
Now fill a second jar with carbon dioxide, and then pour the
contents of this jar into the first. As soon as the taper is
immersed in carbon dioxide it is extinguished. Try to transfer
by means of a siphon the contents of a jar of carbon dioxide
into another arranged with a taper like the first, treating the
gas as if it were a liquid.
Carbon dioxide is heavier than air, its specific gravity
being 1.529. I1 at 0° and 760mm weighs 1.9658. It can be
condensed to a liquid by pressure or by reduction of its
temperature. Under one atmosphere it liquefies at —78°;
a still further reduction, which may be accomplished by
allowing the liquid to escape into a box with a sieve-like
bottom, freezes the liquid to a snow-like solid.
EXP. 106 OP. Place any small animal in a jar of carbonic
acid gas ; note the symptoms and time of death. Also thus
proceed with a jar of carbon monoxide. How do the symptoms
compare ? The time of death ?
Pure carbon dioxide seems to produce its deadly effects
by asphyxiation, the lungs being unable to effect the
decomposition of the gas, and thus to appropriate the
needed oxygen, which it certainly contains, but holds with
142 CAEBON AND OXYGEN.
an exceedingly tenacious grasp. As one would infer, this
gas is very stable; but its decomposition can, nevertheless,
be accomplished.
EXP. 107 P. Into a jar of carbon dioxide place a brightly
burning magnesium ribbon. It continues to burn. Is carbon
set free? Also try a piece of burning sodium. Is the gas
again decomposed? Are other products than carbon formed?
If these products are MgO and Na2O, write the equations.
Since carbon dioxide is liberated in so many different
ways, it is present in the atmosphere in considerable
quantities. It varies from 2.7 to 3.5 volumes in 10,000
volumes of air. This gas is more plentiful in living-rooms
than out of doors, but the amount present should never
be allowed to exceed 7 or 8 parts per 10,000. It is not
so much that carbon dioxide is itself very poisonous, as
that other and more dangerous animal impurities are
thrown off by the lungs together with the carbon dioxide.
We may therefore practically employ the amount of car-
bon dioxide present in a living-room as an index to meas-
ure the purity of the air, as will hereafter be explained.
PROB. Calculate the number of cubic metres of CO2 in the
atmosphere, assuming the extent of the air to be as stated
under Atmosphere. Compute its weight.
Carbon dioxide gas is often found in mines, caves, old
wells, and vats. When so occurring it is termed Choke
Damp, and many persons yearly lose their lives through
a lack of caution in entering such places. Before ventur-
ing into a place where choke damp is likely to occur, it is
best to lower a lighted candle ; should the candle be ex-
tinguished, it is unsafe to go in. A well may sometimes
be freed from choke damp by dashing in much water, the
CARBON AND OXYGEN. 143
gas being thus absorbed ; and, again, a vat may be made
safe by making an opening in the bottom. Why?
Carbon dioxide is indispensable to plant life. It can be
shown that in sunlight the leaves, roots, and green parts
of plants absorb carbon dioxide and give off oxygen ; also
on moonlight nights and under the influence of the electric
light the same processes go on more slowly ; but, in the
dark, carbon dioxide is given off, and oxygen is quite
freely absorbed.
SUG. Devise an experiment to show the effect of a growing plant, in
sunlight, upon carbon dioxide.
QUERIES. Are plants in a living-room conducive to health ? In a
sleeping-room ? How are plants and animals interdependent through
carbon dioxide and oxygen ? What prevents the excessive accumulation
of carbon dioxide in the atmosphere'? What would result if all the
oxygen of the air were consumed ? All the carbon dioxide ? If there
were an excess of the latter gas ?
EXP. 108 P. Fill a bottle of about I1 capacity with water,
and invert it over the pneumatic trough, or better, over a basin
of pure water. Now fill the bottle three-fourths full of carbon
dioxide. Cork the bottle with its mouth under water ; remove,
and shake it thoroughly. Again place the mouth of the bottle
under water, and uncork. Does the water rise in the bottle?
Is carbon dioxide soluble in water? Reduce the temperature
of the bottle by means of a freezing mixture, and shake as
before. Again remove the cork under water. Does a greater
diminution of the volume of the gas take place? Boil a portion
of the water in the bottle, and test by adding calcium hydroxide
(Art. 151). Does the boiled water give a reaction? Testa
portion of the water in the bottle, with blue litmus paper. Is it
acid? Taste of the water in the bottle, or drink of it if you
wish. How does it taste?
NOTE. The gas for this experiment should be washed through a
solution of sodium carbonate. Why ?
144 CARBON AND OXYGEN.
Although carbon dioxide is injurious when inhaled, it
is, nevertheless, when taken into the stomach, sometimes
an aid to digestion. Certain springs and artesian wells
owe their excellent properties to the carbon dioxide
absorbed in their waters, while soda water is simply pure
water highly charged (under pressure) by this gas.
SUG. Examine and describe a soda-water fountain.
QUERIES. What causes the effervescence of champagne? Beer?
Cider ? What effect does vinegar produce upon common baking-soda 1
Can you thus generate carbon dioxide ? What causes dough or " empty-
ings " to rise ? By what process is the gas furnished in this latter case ?
What is meant by heavy bread ?
Carbon dioxide is soluble in cold water, lcc of water
at 0° dissolving about 1.8CC of the gas ; if the pressure be
increased, the solubility is also increased. An increase
of the temperature of the water drives off the gas, the
process being complete at 100°.
QUERIES. What Exp. shows that limestone is soluble in water contain-
ing free carbon dioxide, but insoluble in water containing none of this gas ?
How is the deposition of limestone formations to be explained ? How
the crust formed in the tea-kettle ? The formation of caves ?
When carbon dioxide is passed into water, the solution
is slightly acid, and it is believed that an acid of the
formula H2CO3 is thus formed : —
C02 + H20 = H2C03.
We also consider that the carbonates, such as calcium
carbonate, are derived from this acid. The acid itself, if
it exist at all, is very unstable, thus breaking up when
liberated : —
H2C03 = H20 + C02.
The carbonates, however, are very stable and of great
CARBON AND NITROGEN. 145
importance, and they occur, as previously noted, in im-
mense quantities.
152. Tests for Carbon Dioxide and the Carbonates.
-1. The free gas is detected by conducting it through a
solution of calcium hydroxide, Ca(OH)2, with which it
forms the white precipitate, calcium carbonate, CaCO3.
2. The free gas in water solution may be detected by
adding the same solution as before.
3. The carbonates will effervesce with any strong acid,
preferably nitric or hydrochloric acids, yielding free carbon
dioxide, which may be tested as in 1.
CARBON AND NITROGEN.
CYANOGEN.
153. Cyanogen, CN or Cy, is the only known compound
of carbon and nitrogen. It has been isolated; but its'
constituents do not directly unite to produce it. The
cyanogen compounds, as potassium cyanide, KCy, prussic
acid, HCy, and other substances containing the group of
atoms, CN, are of importance.
Cyanogen gas is prepared by heating mercuric cyanide,
HgCy2, in a hard glass test-tube provided with a delivery-
tube so arranged that the gas may be collected over
mercury. It is soluble in water, and can be condensed at
moderate temperature under a pressure of four atmos-
pheres. It possesses an agreeable odor resembling peach
blossoms, and burns with a purple flame. This gas is so
poisonous that the student should hesitate to experiment
with it.
The specific gravity of cyanogen gas is 1.806.
146 CARBON AND NITROGEN.
HYDROCYANIC OR PRUSSIC ACID, HCN OR HOY.
154. Prussic Acid is one of the most deadly poisons
known. It acts so quickly that antidotes are of little use,
though in some cases ammonia and chlorine have been of
service in counteracting its effects. It is formed by the de-
composition of amygdalin, a complicated substance which
occurs in the leaves of some plants, and in the kernels of
peach pits, bitter almonds, and other fruits. It can be
prepared in a pure, liquid state by passing hydrogen sul-
phide gas, H2S, over mercuric cyanide, HgCy2 : —
HgCy2 + HoS = 2 HCy + HgS.
It should be remembered, however, that this acid is a
volatile liquid, and that its vapors are a deadly poison and
instantaneously fatal if inhaled in any considerable quanti-
ties. The deadly effects of even dilute hydrocyanic acid
may be illustrated by the following experiment which
would better by far be omitted : —
EXP. 109 OP. Dissolve 9E tartaric acid, C4H6OG, in 60CC of
water ; place ID a 70CC flask, and add 4g potassium cyanide, KCy.
Shake, and allow to settle, when a dilute solution, containing
about 3.6 per cent prussic acid, will be obtained (R. and 8.).
Administer to a cat about a teaspoonful, and note effects.
The specific gravity of hydrocyanic acid at 18° is 0.6969.
NOTE. See larger manuals for the remaining numerous compounds of
cyanogen. The more important cyanides of the metals will be noted
under the metals in question ; but the student is not to forget that many
of them are extremely poisonous.
155. Tests for Hydrocyanic Acid and the Cyanides.
— 1. Prussic acid, HCy, when in dilute solution, maybe
thus detected : To the solution add ammonium sulphide,
EXERCISES IN CARBON. 147
NH4HS, and evaporate nearly to dryness on the water-
bath. Ammonium sulphocyanate, NH4SCy, is formed ; this
substance, when dissolved in water and treated with ferric
chloride, Fe2Cl6, turns to a deep-red color.
2. To detect a cyanide in solution, add a few drops of
potassium hydroxide, KOH, and then add ferrous sul-
phate ; shake well, and acidify with hydrochloric acid, when
prussiari blue will be formed if a cyanide be present.
Sue. Use a solution of KCy for these tests.
EXERCISES IN CARBON.
1. Prepare carbon from 10 different articles of food.
2. Write a short description of the carboniferous age in respect to the
condition of the atmosphere and vegetation. (Consult some text-book
on Geology.)
3. Collect snail shells, clam shells, oyster shells, and a few specimens
of limestone, and test for carbonates.
4. PROB. The temperature of the laboratory is 72° F., and the barom-
eter reads 752mm. How many litres of C02 gas may be generated from
25? CaC03 ? How many grams of HC1 are necessary ? How many grams
of CaCl2 will be produced *
5. Fill a common clay pipe with walnut, hickory-nut, or butternut
meats. Seal the bowl by means of a thick paste of plaster of paris and
water. Allow the paste to dry, then heat the bowl in the Bunsen flame.
Ignite the gas which soon issues from the stem, and prove that it contains
hydrogen and carbon.
SUG. Hold a cold glass tube over the flame. Also hold a piece of
cold porcelain against the flame.
6. Produce carbon from marble, snail shells, etc.
7. The value of a sample of coal for reducing iron from its ores is
ascertained by making the following quantitative determinations : 1. Moist-
ure; 2. Volatile matter; 3. Fixed carbon; 4. Ash; 5. Phosphorus; 6. Sul-
phur. The first four determinations may be made thus : Place in a
weighed porcelain crucible about 5s of the coarsely-powdered sample, and
heat at 100° for several hours. Weigh, and note the loss of weight as
" Moisture." Lute on the oover of the crucible by means of a paste of
wood ashes, leaving a very small opening in one side. Allow the luting
148 EXERCISES IN CARBON.
to dry, and weigh the- whole. Now heat to redness for one hour; weigh,
and the loss in weight equals the " Volatile matter." The last weight
minus the weight of crucible and luting equals the weight of "Coke."
Now remove the cover, carefully clean off the luting, and weigh again ;
then burn the residue in the crucible, and weigh, noting the loss of weight
as " Fixed carbon." The last weight minus the weight of crucible equals
the " Ash." The value of a coal partly depends upon the amount of fixed
carbon it contains. (See Sulphur and Phosphorus.)
8. For valuable information concerning the varieties of coal, coal
analysis, etc., see Dana's System of Mineralogy, pp. 751-760.
9. As previously stated, it is customary to measure the amount of
carbon dioxide as an index to the purity of the atmosphere of a room.
This is accomplished by litration ; and a litre-flask and two reagent solu-
tions are required.
The first solution consists of 58 barium hydroxide, Ba(OH)2, dissolved
in I1 of distilled water ; the second, 2.863s pure freshly-crystallized oxalic
acid, H2C2O4(H20)2, in the same amount of water. From the manner of
using this latter solution, lcc corresponds to lms carbon dioxide.
The litre-flask is filled, by several puffs of a hand-bellows, with the air
to be tested, and the temperature of the room carefully noted. A quantity
of the first solution, equal to the space above the litre-mark on the neck of
the flask, is now added, and the flask vigorously shaken. A portion of the
solution in the flask is neutralized, —
Ba(OH)2 + C02 = BaCO3 + H20,
and a portion is unchanged. The remainder is now carefully neutralized
by means of the second solution, —
Ba(OH)2 + H2C204,(H20)2 = BaC204 + 4 H20,
and the number of cubic centimetres is carefully noted; a phenol phthalein
solution is employed as an indicator. An amount of the first solution
equal to that placed in the flask is now directly titrated with the second
solution, and the number of cubic centimetres of the latter noted. It is
evident that the difference between the two numbers thus obtained equals
the number of milligrams of C02 per litre.
QUERIES. Why do we take 2.863s oxalic acid ? Having the number
of milligrams C02 per litre, multiply the result by 10, and then calculate
the number of cubic centimetres per 10,000. What principles apply ?
What is titration ? An indicator ?
10. The student who wishes to obtain a clearer insight into the
processes employed in the Chemistry of the Carbon Compounds, will do
well to consult Dr. llemsen's work on that subject.
CHAPTER X.
MOLECULES. — MOLECULAR FORMULAE. — VALENCE.
156. Molecules. — What is meant by the word atom
has already been explained. The chemical atom is the
smallest particle of an element that can take part in
chemical reactions. Now, if we consider any chemical
compound as, for example, hydrochloric acid, it is clear
that the smallest particle of this compound which can be
imagined must contain both hydrogen and chlorine, and
must contain at least one atom of each of these elements.
Such a smallest particle of a compound is called a
molecule.
The molecules of compound bodies are made up of
atoms of different kinds. The molecules of the elements
are made of atoms of the same kind. The theory com-
monly held is that when the elements exist in the free
state their atoms unite to form molecules.
The formulae which we use to represent compounds are
intended to represent molecules, just as the symbols of
the elements are intended to represent atoms. Thus the
formulae H2O, NH3, HC1, HNO3, etc., represent the mole-
cules of water, ammonia, hydrochloric and nitric acids ;
and we see from them that the molecule of water is made
up of 2 atoms of hydrogen and 1 of oxygen; that the
molecule of ammonia consists of 1 atom of nitrogen and 3
atoms of hydrogen, etc. Knowing the weights of the
150 MOLECULES.
atoms which mak'e up a molecule, we know the weight of
the molecule. It is the sum of the weights of the atoms
contained in it. The molecular weight of water is 18,
which is the sum of the weight of 2 atoms of hydrogen
(2 x 1) and of 1 atom of oxygen, 16. The molecular
weight of ammonia is 14^(the atomic weight of nitrogen)
+ 3 (the weight of three atoms of hydrogen) = 17.
QUERY. What is the molecular weight of hydrochloric acid ? of nitric
acid ?
157. Avogadro's Hypothesis. — If the atomic weights
of all the elements were known to us there would be little
difficulty in determining the molecular formulae of com-
pounds. Thus, if we knew that the atomic weight of
oxygen is 16, and on analysis found that water consists of
hydrogen and oxygen in the proportion of 1 part of hydro-
gen to 8 of oxygen, the simplest formula which we could
give to the compound would be H2O, and we might
assume that this represents the molecule. A molecular
formula, according to this, would be nothing more than
the simplest formula which could be used to express the
composition of a body, assuming the correctness of the
commonly accepted atomic weights. In reality, the molec-
ular formulae mean more than this ; they are dependent
upon a very ingenious and valuable hypothesis, known as
the hypothesis of Avogadro.
On comparing the specific gravities of a number of
gaseous compounds with the molecular weights of the
same compounds, it is found that the two sets of figures
bear the same relation to each other. In other words, the
specific gravity of any compound gas is to the molecular
weight of the compound, as the specific gravity of any
other gas is to its molecular weight. This leads to the
MOLECULES. 151
conclusion that equal volumes of bodies in the form of gas
or vapor contain the same number of molecules, and this is
Avogadro's hypothesis. According to the hypothesis, if
a cubic inch of hydrochloric acid gas contains (say) 1000
molecules, a cubic inch of any other gas or vapor,
measured under the same conditions of pressure and
temperature, also contains 1000 molecules. We can not
determine the absolute number of molecules present in
any given volume, and hence, of course, can not determine
the absolute weight of the molecules; but accepting the
hypothesis we can easily determine the relative weights
of molecules of all substances which are gaseous or can be
converted into vapor. These relative weights compared
.to some standard are what we know as the molecular
weights.
We may take any simple molecule, as hydrochloric
acid, as a standard. The simplest formula which can be
assigned to this substance to express its composition is
HCl, in which the atomic weight of chlorine is assumed
to be 35.5. The molecular weight of a compound of this
formula is 36.5. Let this be the standard molecule. The
problem now is to determine the weights of the molecules
of other bodies in terms of this standard, and in accordance
with the principle laid down in Avogadro's hypothesis.
We simply determine the relative weights of equal vol-
umes of hydrochloric acid and the other gases or vapors,
and, knowing that the molecular weights bear to one
another the same relation as these relative weights, the
molecular weights can easily be deduced.
The figures which express the relative weights of equal
volumes of bodies are called the specific gravities. We
have then only to compare the specific gravities of gases
152 MOLECULES.
with that of hydrochloric acid to know the molecular
weights of these bodies.
If S' is the specific gravity of hydrochloric acid, and
36.5 its molecular weight ; S the specific gravity of some
other gas, and M its molecular weight, we have : —
S':36.5::S:M,
but S' is known. It is 1.247. Hence we have: —
1.247 : 36.5 : : S (the sp. gr. of any gas) : M (its molecular wt.) .
In other words, the relation between the specific gravity
of any gas and its molecular weight is represented by a
constant quantity which is about 28.8, i.e., —
M = 28.8, or M = 28.8 X S.
b
The molecular weights of all bodies which can be con-
verted into the form of vapor have been determined by
means of this rule, and the molecular formulae are based
upon these determinations.
158. Determination of Atomic Weights by means
of Avog-adro's Hypothesis. — In order to determine
atomic weights by means of the hypothesis of Avogadro,
we first determine the molecular weights of all compounds
which are gaseous or can be converted into vapor. We
then analyze these same compounds. On now examining
the results of the analysis, we select the smallest quantity
of an element which occurs in any of its compounds, as
its atomic weight.
The method may be illustrated by taking some of the
compounds of carbon as examples.
MOLECULES. 153
Molecular ru«n*u,
Wt. Found.
Carbon monoxide . 27.96 12 parts C ; 16 parts O.
Carbon dioxide . . 44.16 12 " C ; 32 " O.
Marsh gas . . . 16.1 12 " C; 4 " H.
Ethylene .... 28.0 24 " C ; 4 " H.
Acetylene . . . 26.0 24 " C • 2 " H.
The smallest quantity of carbon contained in any of
these compounds is represented by the figure 12, and
consequently this is accepted as the atomic weight, unless
there is some other compound the molecular weight and
analysis of which lead us to a smaller figure.
159. Valence. — Having determined the molecular for-
mulae of chemical compounds, we see that they differ
markedly from one another. Take, for example, the
hydrogen compounds of some of the elements thus far
considered. We have hydrochloric acid represented by
HC1, water by H2O, ammonia by H3N, and marsh gas by
H4C. A fundamental difference between these compounds
is noticed in the number of hydrogen atoms contained in
each one. In HC1 we have 1 H ; in H2O, 2 H ; in H3N,
3 H; and in H4C, 4 H. The atoms of chlorine, oxygen,
nitrogen, and carbon are thus seen to differ from one
another in regard to the number of hydrogen atoms which
they can hold in combination. The power of any atom to
hold a certain number of the simplest atoms in combina-
tion is called its valence. This term is also applied to the
elements. We speak of a univalent element meaning an
element the atom of which has the power of holding one
of the simplest atoms in combination. Thus chlorine and
hydrogen are univalent elements.
We may measure the valence of any element by any
154 MOLECULES.
univalent element with which it will unite. Thus we
measure the valence of oxygen by hydrogen. It is
bivalent because its atom unites with two atoms of hydro-
gen. In the same way we regard nitrogen as trivalent
because its atom unites with three atoms of hydrogen ; and
carbon as quadrivalent because its atom unites with four
atoms of hydrogen.
Some elements do not unite with hydrogen. In these
cases we may measure the valence by means of any other
univalent element, as chlorine. Thus potassium does not
unite with hydrogen, but it does unite with chlorine,
forming the compound KC1, which shows that potassium
is univalent; calcium forms the compound CaCl2, which
shows that calcium is bivalent. The valences of all the
elements have thus been determined by a study of the
formulae of their compounds. In many cases one and
the same element has more than one valence, as shown in
the two chlorides of phosphorus, PC13 and PC15, in the
first of which phosphorus appears as a trivalent and in the
second as a quinquivalent element.
160. Substituting1 Power and Valence. — We have
seen that in the formation of salts the hydrogen of the
acids is replaced by metals. The number of atoms of
hydrogen which the atom of any metal can replace is
determined by the valence of the metal. The atom of a
univalent metal replaces 1 atom of hydrogen, as is shown
in the formation of potassium nitrate, KNO3, from HNO3 ;
the atom of a bivalent metal replaces 2 atoms of hydrogen,
as in the formation of calcium nitrate, Ca(NO3)2, from
HNO3, in which case the calcium atom is represented as
taking the place of two atoms of hydrogen in two mole-
cules of nitric acid. In barium sulphate, BaSO4, one
EXERCISES IX EQUATIONS. 155
atom of the bivalent metal barium takes the place of the
two hydrogen atoms in sulphuric acid H2SO4. In making
hydrogen by treating sulphuric acid with zinc, we had
another illustration of the replacement of the two hydro-
gen atoms of sulphuric acid by one atom of the bivalent
metal zinc. Numerous illustrations of the different sub-
stituting powers of the metals will present themselves
when the salts come up for consideration.
NOTE. It is customary to consider the part of an acid which remains
in combination with a metal after the hydrogen has been displaced as a
group of atoms, and when we wish to take this group more than once, as
above, we write Ca(NO3)2 and not CaN2O6. By so doing the formula
shows at a glance what acid took part in forming the compound.
EXERCISES IN EQUATIONS.- USEFUL PROBLEMS.
1. The equations previously given might with propriety be termed
" Atomic Equations," since they show what we believe takes place at the
instant dissociation of a compound occurs.
We may also write " Molecular Equations," showing the state of
affairs after all reactions are complete. In order to do this we only need,
in addition to what we have already practised, to represent the molecules
of the free elements in some appropriate manner, so that the formula for
the molecule shall show the number of atoms it contains. It is now
becoming customary to do this by the use of subscript figures ; thus, O2,
H2, N2, P4, S2, etc., represent the molecule of oxygen, nitrogen, phosphorus,
etc. Let us now again take up some of the atomic equations already
given, and rewrite them to represent molecular conditions : —
K + H2O = KOH + H, when rewritten gives 2 K + 2 H2O = 2 KOH + H2 ;
Zn + H2S04 = ZnS04 + 2 H becomes Zn + H,SO4 = ZnS04 + H2 ;
2 P + 5 0 = P2O5 becomes 2 P4 + 10 02 = 4 P~05 ;
S + 2 O = SO2 becomes S2 + 2 02 = 2 S02.
By inspecting the equations thus rewritten it becomes apparent that
molecular equations are somewhat the more complex of the two, and that
to write them properly requires a knowledge of the molecular formulae of
the elements. In the compounds, as previously stated, the formula also
represents the molecule ; not so however with the symbols of the elements ;
and since it is first necessary to determine the vapor density of an element
156 EXERCISES IN EQUATIONS.
before we can determine its molecular formula, it is evident that when we
come to solids not readily volatilized it is manifestly absurd to write such
a formula as Au2, Pt3, etc., especially if we agree to represent the mole-
cules of elements by subscript figures.
"Write in molecular formulae : —
C + 20 = C02;
3Fe + 4O = Fe304;
Zn + 0 = ZnO.
2. To calculate the weight of a given volume of any gas from its
molecular weight : —
PROB. 1. How much does I1 of HC1 gas weigh at 0° ?
SOLUTION. The molecular weight equals 35.5 + 1 = 36.5, and the density
(with reference to H) equals 36.5^2 = 18.25. Now I1 of H at 0° and
760mm weighs 0.0896s, and it is evident that the required weight equals
18.25 X 0.0896. In case the temperature and pressure vary from standard
conditions the problem may be finished by Art. 87.
NOTE'. Note that The density of a gas (H = 1) equals one-half its molec-
ular weight. This follows from the fact that we take the hydrogen
molecule, H2, as 2 ; or the half molecule, H, as unity.
PROB. 2. How much do 61 of chlorine weigh at 15° and 750mm ?
SUG. The molecular formula of chlorine is C12, and its density equals
2x35.5-^2 = 35.5 or the atomic weight of Cl. We may here note that
the density and atomic weights of the gaseous elements are numerically
equal.
PROB. 3. Compute the weights of I1 of the following gases : O, N,
N20, N203, NH3, H2S, S02, C02, CO.
3. To compute the specific gravity (air= 1) of a gas from its molecular
weight. Divide the weight of I1 of that gas by the weight of I1 of air, or
1.293.
PROB. 4. What is the specific gravity of C02 ? H2S ? CO ? NH3 ?
4. Show that one needs simply to remember the atomic weights of
the elements to compute: 1. The molecular weight of any gas; 2. Its
density ; 3. The weight of I1.
CHAPTER XL
SULPHUR. — SELENIUM AND TELLURIUM. — THEIR OCCUR-
RENCE, PREPARATION, TESTS, ETC.
SULPHUR.
SYMBOL, S". — ATOMIC WEIGHT, 32. — SPECIFIC GRAVITY
(CRYSTALS), 2.05.
161. Occurrence. — Sulphur occurs native in volcanic
regions, and in its compounds with other elements it is
widely distributed. The most plentiful of these com-
pounds are the sulphides, iron pyrites, FeS2, or Fool's
Gold ; galena, PbS ; cinnabar, HgS ; and the sulphates,
gypsum, CaSO4 + 2 H2O ; heavy spar, BaSO4; green
vitriol or ferrous sulphate, FeSO4 + 7 H2O, etc.
Native sulphur occurs in regular, yellowish, transparent,
octahedral crystals, and in other forms derived from this
primary crystal. It is also found in a massive state
being then known as volcanic sulphur.
162. Preparation. — Since sulphur in its various forms
is a common article of commerce it may readily be pro-
cured for class purposes. The common roll sulphur or
brimstone is prepared by distilling the crude ore in large
earthen-ware retorts, and condensing the vapors in stone-
ware condensers. More frequently, however, it is ob-
tained by building up the crude ore in the form of a kiln
158 SULPHUR.
or charcoal pit, where the ore is roasted by burning a
portion of the sulphur as a fuel. The sulphur is melted
from its accompanying impurities, and runs down into a
receptacle prepared to receive it at the bottom of the pit.
It is afterwards purified by distillation, and cast into the
ordinary rolls or sticks.
Flowers of Sulphur, also an article of commerce, are
obtained by vaporizing a quantity of sulphur and bringing
the vapor into a cold condenser, where this variety is pro-
duced in a manner analogous to snow.
EXP. 1 10 P. Dissolve 2s flowers of sulphur in 13CC of water,
to which has been added 1s slacked lime (prepared by treating
1 part quicklime with 3 parts water). The product calcium
pentasnlphide, CaS5, is formed. Write the equation. Now
add to the solution hydrochloric acid, when the liquid turns
white, very finely divided sulphur being obtained.
The substance thus prepared is an article of commerce
known as lac sulphuris or milk of sulphur.
163. Properties. — EXP. Ill p. Dissolve lg sulphur in 3g
carbon bisulphide, CS2. Place the solution in a beaker glass,
and allow it to evaporate, without heat, in the atmosphere.
Octahedral sulphur crystals will be obtained. Allow these
crystals to stand for several days, noting from time to time
any changes that may occur.
Sulphur crystals occur in no less than thirty different
forms all derived from the primary octahedron. The
specific gravity of these primary crystals at 0° is 2.05.
EXP. 112x. Melt in an evaporating dish 100s sulphur and
heat to 230°, when the molten mass will turn black. Now
pour into a basin of cold water, and when cold remove and
examine the product obtained. Leave for several days in the
water, and occasionally observe what changes occur.
SULPHUK. 159
The modification of sulphur thus obtained is known as
plastic sulphur, and at first strongly resembles caoutchouc,
in that it is elastic ; it soon becomes brittle, however,
upon standing. The specific gravity of this form is 1.96.
EXP. 113T. Melt in a sand crucible a quantity of sulphur
and allow it to cool slowly. When a crust forms over the
surface of the molten sulphur make an opening through the
crust and pour off the liquid portion. Note the peculiar needle-
shaped crystals attached to the solid crust.
QUERIES. How many different forms or modifications of sulphur have
you observed1 What changes take place in the crystals last obtained
when they are allowed to stand ?
. Sulphur is extensively used in making sulphuric acid
and in the manufacture of rubber goods. When heated
at moderate temperatures with crude rubber gurn, 2 to 3
per cent of sulphur is absorbed, and the product obtained
is firmer and better adapted to some industrial require-
ments than the pure gum itself. When the temperature is
raised to a higher degree the substance called vulcanite
or ebonite is obtained.
QUERY. What developments in the rubber industry are due to Samuel
Goodyear ?
EXP. 114 p. Dip into powdered sulphur a pine splinter and
ignite ; note the flame and the odor emitted. What does the
odor resemble ? The fumes have the formula SO2. Write the
equation.
Sulphur is used in the manufacture of matches and is
burned for bleaching straw goods. Some forms are also
employed in medicine.
It is capable of uniting directly with most metals to
form sulphides.
160 SULPHUR AND HYDROGEN.
164. Tests for Free Sulphur. — 1. Free sulphur is dis-
tinguished, if in considerable quantities, by its physical
properties, and by its flame and the odor of its fumes.
2. If the quantity be too small to test as in 1, fuse it on
platinum foil with sodium carbonate, Na2CO3; then place
the fused mass, which is sodium sulphide, Na2S, on a bright
piece of silver, and moisten with a drop of water. If
free sulphur be present, a black spot of silver sulphide
will be obtained.
CAUTION. The Na2C03 and charcoal must be free from sulphur; like-
wise the illuminating gas used for the blow-pipe flame. The alcohol lamp
is best to use for this test.
NOTE. Since sulphur blackens silver, egg spoons, mustard spoons, etc.
are gilt to prevent their tarnishing. Silver ware blackened by sulphur is
easily brightened by washing in a solution of potassium cyanide, KCy ;
this is better than scouring, since the cyanide does not attack the pure
silver, bow may the black spot obtained in 2 be removed ?
SULPHUR AND HYDROGEN.
165. Sulphur and hydrogen form two compounds, viz. : —
Hydrogen Sulphide, H2S,
Hydrogen Persulphide, H2S2(?).
Of these the first alone is of importance to the beginner.
HYDROGEN SULPHIDE.
166. Occurrence. — Hydrogen sulphide, commonly
known as sulphuretted hydrogen, is of wide occurrence,
both free and combined. The waters of many famous
"sulphur springs" contain this gas in large quantities.
It is a product of volcanic action and of the decomposition
of albuminous substances; thus the peculiar odor of
SULPHUR AND HYDROGEN. 161
rotten eggs is partly due to the hydrogen sulphide
evolved.
The sulphides, which may be regarded as derived from
this acid, are found in great abundance, as already
mentioned.
167. Preparation. — EXP. HDP. Place in a test-tube n
small quantity of water, say 10CC, and add a small piece of
ferrous sulphide, FeS ; now add lcc of sulphuric acid, and close
the tube quickly with a perforated cork containing a glass
U-shaped jet delivery-tube. The gas will soon issue through
the jet, when it may be ignited. Note the odor, but do not
allow more gas than is necessary to escape, since it is some-
what poisonous. The contents of the tube should be poured
into the sink as soon as a sufficient amount of gas has been
obtained, but in case a considerable piece of the sulphide re-
mains this may be saved for further use.
This is the general method and the one almost exclus-
ively employed in laboratory practice for the production
of hydrogen sulphide. The chemist thus produces it for
analytical purposes, as will subsequently be explained.
It is well to have a gas chamber wherein this gas may be
produced and wherein the whole contents of the test-tube
may be retained, since another reagent, ferrous sulphate,
is thus produced : —
FeS + H2S04 = H2S + FeSO4.
This latter compound may be separated by crystallization.
In case large quantities of sulphuretted hydrogen are
required, a generating flask may be employed instead of a
test-tube, and the gas may be washed through warm
water. An aqueous solution in cold water is to be had,
but the gas itself, freshly generated, is preferable for
qualitative work.
162 SULPHUR AND HYDROGEN.
Hydrogen sulphide is also formed by the action of some
of the other acids on the sulphides ; by burning sulphur
in an atmosphere of hydrogen; by, passing hydrogen
through boiling sulphur, and by heating paraffine with
sulphur. All these methods are, for various reasons, not
well adapted for obtaining the gas in practice.
168. Properties. — Hydrogen sulphide is a colorless,
inflammable gas, possessing a disagreeable odor somewhat
resembling rotten eggs. It is condensed, at ordinary tem-
peratures under a pressure of 17 atmospheres, to a color-
less liquid which boils at — 61.8° and freezes at — 85°.
Its specific gravity at 0° is 1.191, and I1 weighs 1.522g.
lcc of water at 0° absorbs about 4.4CC hydrogen sulphide,
forming a slightly acid solution.
EXP. 116 P. Place in several different test-tubes solutions
of metallic salts, such as copper sulphate, CuSO4 ; mercuric
chloride, HgCl2 ; lead acetate, Pb(C2H3O2)2, and silver nitrate,
AgNO3. Generate hydrogen sulphide as in Exp. 115, and
successive^' place the jet into these solutions, allowing the gas
to bubble up through them. Precipitates which are respectively
the sulphides of the different metals will be formed.
It is thus that the chemist employs hydrogen sulphide
in analytical operations, and the great utility of this gas
becomes apparent when it is known that by its aid the
metals may be separated into groups. In short, it is
another group reagent (p. 98). The same is true of one
of its compounds, ammonium sulphide, (NH4)2S.
SUG. Try the effect of H2S upon solutions of arsenic, antimony,
cadmium, copper, and tin. Note the colors of the precipitates.
EXP. 117 P. Pass sulphuretted hydrogen through nitric acid ;
aqua regia ; strong hydrochloric acid ; sulphuric acid. Do you
SULPHUK AND HYDROGEN. 163
obtain precipitates ? If so, collect and burn on a pine splinter.
Note the odor of the fumes. What is the sediment obtained ?
What effect do stronger acids have upon hydrogen sulphide?
Make a solution of lead nitrate, Pb(NO3)2, and strongly acidify
with nitro- hydrochloric acid. Now pass hydrogen sulphide.
Do you obtain lead sulphide ? Wiry ?
169. Tests for the Sulphides. — 1. Free hydrogen sul-
phide in quantity is distinguished by its odor and by its
blackening effect upon paper moistened with lead acetate,
Pb(C2H3O2)2. Also see Exp. 38.
2. A sulphide, when fused on platinum foil or a bit of
porcelain, — as a piece of broken evaporating dish, — with
sodium carbonate, and moistened, produces a black spot
when placed on a clean piece of silver.
QUERIES. How do the sulphides, as FeS, behave with sulphuric acid 1
What is meant by a test ?
NOTE. The salts of easily reducible metals, such as those of lead and
mercury, must not be fused on platinum, since these metals form with the
platinum alloys which are fusible at high temperatures. The platinum
may thus be ruined. In such cases it is necessary to fuse on charcoal or
porcelain. What disadvantage does this latter process involve ?
HYDROGEN PERSULPHIDE, H2S2(?).
170. Hydrogen persulphide may be prepared by boiling
(say) lg slacked lime with 16CC water and 2g flowers of
sulphur. The cold clear solution is then poured into
dilute hydrochloric acid, when the persulphide falls to the
bottom of the vessel as an oily liquid.
It has a very disagreeable odor, more pungent than that
of hydrogen sulphide. It is not important for the be-
ginner.
164 SULPHUK AND OXYGEN.
SULPHUR AND OXYGEN.
171. There are two oxides of sulphur deserving special
mention, viz. : —
Sulphur Dioxide, SO2,
and Sulphur Trioxide, SO3.
These oxides are respectively the anhydrides of sulphurous
and sulphuric acids. The manner in which they combine
with a molecule of water is worthy of notice : —
1. H2O + SO, = H2SO3 ;
2. H
It will be seen that in either case one molecule of water
and one molecule of oxide form but one molecule of acid.
In the case of the oxacids of nitrogen, bromine, chlorine,
and iodine two molecules of acid were thus formed.
Two other oxides corresponding to the formulae, S2O3
and S2O7, are known.
SULPHUR DIOXIDE, SO2.
172. Occurrence. — This oxide is the gas formed when
sulphur is burned in the atmosphere. It occurs free in
volcanic gases, and combined with other elements, as in
the sulphites or salts of sulphurous acid.
173. Preparation. — EXP. 118 p. Place in a generating
flask fitted with a delivery-tube lg very fine copper filings and
6CC strong sulphuric acid. Heat until a gas begins to escape.
Note the odor, and collect by displacement in a large test-tube,
or small, tall jar.
SUG. Some other metals when thus treated also yield sulphur dioxide.
Try several, such as iron, mercury, and lead.
SULPHUK AND OXYGEN. 165
When sulphur dioxide is thus prepared the reaction
may be indicated by the equation : —
Cu + 2 H2SO4 = CuSO4 + 2 H2O + SO,.
Notice the difference between this reaction and that which
takes place when sulphuric acid and zinc are brought
together. In the latter case the reaction is represented
thus : —
Zn + H2SO4 = ZnSO4 + H2.
Whenever a metal reacts with an acid the first action con-
sists in the replacement of the hydrogen of the acid by
the metal. The hydrogen is liberated and a salt is formed.
In the case of copper and sulphuric acid, however, the
reaction does not take place at ordinary temperatures,
and at higher temperatures the hydrogen which is first
liberated acts upon the sulphuric acid reducing it to
sulphur dioxide : —
H2SO4 + H2 = 2 H2O + SO2.
A common method of preparing this gas is to burn
sulphur in the air. Many other methods are also known,
such as heating sulphur and carbon with sulphuric acid,
roasting pyrites, etc.
174. Properties. — Sulphur dioxide gas is easily con-
densed by passing it through a spiral glass tube surrounded
by a freezing mixture. It is very soluble in water, lcc of
which dissolves, at 0°, about 79.8CC of this gas. Its specific
gravity is 2.211, and I1 weighs 2.862g. It condenses at
4- 59° under 79 atmospheres and boils at — 8° under 1
atmosphere.
Sulphur dioxide is used in great quantities for preparing
sulphuric acid, in which case it is prepared by burning
sulphur or iron pyrites in a current of air.
1G6 SULPHUH AND OXYGEN.
EXP. 119 p. Suspend in a jar of sulphur dioxide a strip of
moistened unbleached silk ; a moist wheat straw ; a piece of
white woollen yarn.
Sulphur dioxide is used for bleaching such goods as
chlorine would injure. It produces its effects by reduc-
tion instead of oxidation, as in the case of those bleaching
reagents previously noticed. It unites with the oxygen
of water, liberating hydrogen, and this latter gas enters
into combination with the coloring matter to form color-
less compounds.
QUERIES. How do milliners prepare the sulphur dioxide which they
use in bleaching straw goods ? What other substances have been men-
tioned as reducing agents 1 What is meant by reduction ?
A solution of sulphur dioxide in water becomes oxidized
if it cornes in contact with air, sulphuric acid being
formed. It is probable that in the solution sulphurous
acid, H2SO3, is present, and that this takes up oxygen, thus
passing into sulphuric acid. Write the equations.
Sulphur dioxide is also a good disinfectant, and will
prevent the decay of meats and vegetables when applied
for that purpose. It also prevents fermentation.
175. Tests for Sulphur Dioxide. — 1. Its odor is marked
and well known, resembling that of burning matches.
2. Suspend in this gas a strip of paper which has been
dipped into a solution of starch paste and potassium
iodate, KIO3. Iodine is liberated, and the paper becomes
blue : —
2 KIO3 + 5 SO2 + 4 H2O = 2 HKSO4 + 3 H2SO4 + 12.
NOTE. The sulphur dioxide must not be present in excess, or the paper
will be bleached, hydriodic acid being produced. Write the equation.
THE SULPHUR OXACIDS. 167
SULPHUR TRTOXIDE, SO3.
176. Sulphur trioxide is somewhat difficult of prepara-
tion and very unstable owing to the eagerness with
which it unites with water.
It is prepared for commerce by passing sulphur dioxide
together with oxygen over finely divided platinum in a
highly-heated porcelain tube. It may also be prepared by
heating strong sulphuric acid with phosphorus pentoxide :
H2SO4 + PA = 2 HPO3 + SO8.
Sulphur trioxide was formerly supposed to be the true
sulphuric acid, but as soon as it was separated it proved to
be a white crystalline solid without action upon the metals
in absence of moisture. The discovery of this substance
brought about a marked change in the views held in
regard to salts and acids, and was one of the many causes
which have led up to our present conceptions concerning
chemical reactions and chemical formulae.
THE SULPHUR OXACIDS.
177. In this series eight different acids are known, the
names and formulae of which are shown by the subjoined
list : -~
Hyposulphurous acid . . . H2SO2 ;
Sulphurous acid H2SO3 ;
Sulphuric acid H2SO4 ;
Thiosulphiiric acid .... H2S2O3 ;
Dithionic acid H2S2O6 ;
Trithionic acid H2S3O6 ;
Tetrathionic acid . . . . H2S4O6 ;
Pentathionic acid .... H2S5O6?
NOTE. The root "thion" is of Greek derivation, signifying sulphur.
168 THE SULPHUR OXACIDS.
By inspection, it will be seen that all these acids are
dibasic, possessing two atoms of replaceable hydrogen ;
hence, they yield both acid and normal salts, e.g., mono-
sodium sulphite, HNaSO3; sodium sulphite, Na2SOc, etc.
Acids which contain tut one replaceable hydrogen atom
are called monobasic acids; those which contain two re-
placeable hydrogens are called bibasic acids; those with
three are called tribasic acids, and those which contain
four are called tetrabasic acids. Most common acids
belong to the first two classes. All the acids previously
considered, excepting carbonic acid, are monobasic ; the
latter and the sulphur acids are bibasic. The principal
tribasic acid is phosphoric acid, H3PO4. There is no
common tetrabasic acid.
SUG. Name the salts formed by the sulphur acids and potassium.
Write their formulae.
NOTE. The student who has thus far followed these pages will have
noted that the rarer acids are chiefly of interest to the scientist, and that
they are all unstable and somewhat difficult of preparation. His experi-
ence, moreover, with these unimportant compounds will have served to
give him a sufficient conception as to the characteristics of the class of
substances to which they belong. We shall therefore note but three acids
of this series ; viz., sulphurous, sulphuric, and thiosulphuric acids.
SULPHUROUS ACID, H2SO3.
178. This acid, as previously noted, is formed when
sulphur dioxide is passed into water. It is an unstable
acid constantly giving off sulphur dioxide fumes ; but the
sulphites are a well known class of salts.
EXP. 120 P. Pass sulphur dioxide gas into a test-tube of
cold water ; also into a cold solution of sodium or potassium
hydroxide. What does each tube contain after passing the
THE SULPHUR OXACIDS. 169
gas? Gently evaporate to dryness the contents of the second
tube, and a salt is obtained. Complete this equation, —
Use the contents of these tubes for the following : —
179. Tests for Sulphurous Acid and the Sulphites.
— 1. Free sulphurous acid in quantity is recognizable by
its odor.
2. In traces it may be detected by a solution of starch
paste and potassium iodate, owing to the blue tinge pro-
duced. It will also blacken a strip of paper moistened
with silver nitrate.
3. The sulphites in solution upon addition of a stronger
acid (HC1, H2SO4) remain clear, yielding sulphur dioxide
fumes. (See Thiosulphuric Acid.)
4. When barium chloride is added to a solution of a
sulphite, the white precipitate barium sulphite, BaSO3, is
thrown down. Divide this precipitate in two parts : to
the first add hydrochloric acid ; it is soluble. To the second
add nitric acid ; the sulphite is oxidized to barium sulphate,
BaSO4, a white precipitate insoluble in acids.
SUG. Complete and balance the following equations, and explain the
principles they illustrate : —
5H2S03 +2KI03 = I+HKSO4+H2S04+ . . .;
Na2SO3 + HC1 = S02 + . . . + NaCl ;
Na2SO3 + BaCl2 =NaCl+ • • •;
BaS03 +HN03 = BaS04+H2O + . . .
SULPHUIUC ACID, H2SO4
180. Occurrence. — Although sulphuric acid does not
occur in nature except in volcanic waters, it is the most
important acid known to the chemist and to commerce,.
170
THE SULPHUR OX ACIDS.
It has even been stated that the prosperity of a country
may be estimated by the amount of sulphuric acid which
that country consumes.
Its salts are very stable and of great value, as, for ex-
ample, blue vitriol, CuSO4 + 5 H2O, a salt of copper used
for galvanic batteries and many other purposes; gypsum
or land plaster, CaSO4 + 2H2O, used by farmers as a manure ;
green vitriol or ferrous sulphate, FeSO4 + 7 H2O, a well-
known salt used in the laboratory as a reagent, and also
used for purifying water-closets, sewers, etc. ; Glauber
salts, Na2SO4 + 10 H2O ; Epsom salts, MgSO4 + 7 H2O, and
the sulphates of the alkaloids used in medicine.
181. Preparation. — EXP. 121 T. Although the student
will have this acid upon his table, where he may study its
properties at his leisure, it might be well to illustrate the inter-
FIG. 18.
esting process of its manufacture. The formation of sulphuric
acid may be beautifully shown by employing the apparatus
illustrated in Fig. 18. G is a glass globe used as a condensing
chamber. B is a generator containing copper filings and
THE SULPHUR OXACIDS. 171
sulphuric acid for the purpose of producing sulphur dioxide.
C is a flask containing water for generating steam. A contains
copper filings and nitric acid for generating nitrogen dioxide
and nitrogen trioxide. D is used to convey air into the con-
densing chamber, and is attached to a hand-bellows. E is an
escape-pipe to allow the waste gases, nitrogen and nitrogen
dioxide, etc., to be forced out of the chamber. In practice most
of these gases are utilized, but in this experiment E must be
placed in a good ventilating draft. When the products of A,
B, and C begin to fill the condenser, a steady, but gentle,
current of air from the bellows must be forced through G until
the close of the experiment. Sulphuric acid is thus produced,
and falls to the bottom of the condenser.
In preparing commercial sulphuric acid the materials
and principles used vary but slightly from those illustrated
in the foregoing experiment. The sulphur dioxide is
prepared by burning sulphur or roasting iron pyrites,
FeS2, in a current of air. The fumes are conducted
into immense lead-lined chambers where they are mixed
with air and steam and the higher oxides of nitrogen, as
N2O3 and NO2 ; or at first a little nitric acid formed from
sodium nitrate and sulphuric acid is used. The steam is
obtained from a boiler arid is blown into the chamber
through jets stationed at different points.
The chemical processes involved in the manufacture of
sulphuric acid are quite complicated. The essential
features will appear from the following brief description :
When the sulphur dioxide and nitric acid first come to-
gether in the presence of steam this reaction takes place : —
2 HNO3 + 3 SO2 + 2 H2O = 3 H2SO4 + 2 NO.
As will be seen, the nitric acid is reduced to nitric oxide,
NO, and this is incapable of oxidizing any more sulphur
172 THE SULPHUR OXACIDS.
dioxide ; but the oxygen of the air which is present im-
mediately transforms the nitric oxide into nitrogen tetrox-
ide, NO2 (NO + O = NOa), and this, in the presence of
steam, converts a further quantity of sulphur dioxide into
sulphuric acid, as indicated in this equation : —
SO2 + H2O 4- NO2 = H2SO4 + NO,
and is itself again reduced to nitric oxide. This NO again
takes up oxygen to form nitrogen tetroxide, which in turn
oxidizes sulphur dioxide, and so on, indefinitely. Thus,
theoretically, starting with a small quantity of nitric
acid, an infinite quantity of sulphur dioxide could be
converted into sulphuric acid, as, after the nitric oxide,
NO, is once formed, it simply serves the purpose of trans-
ferring oxygen from the air to the sulphur dioxide. Prac-
tically, of course, there is always some loss of the oxides
of nitrogen, and this loss must be made good by a fresh
supply in order to make the operation continuous.
The acid formed in the leaden chambers is a weak acid
having a specific gravity of 1.55. It is withdrawn into
large leaden pans, and concentrated until its specific
gravity reaches 1.71, when it is quickly removed, since
any further concentration would result in the destruction
of the pan.
It is further concentrated and purified in glass or plati-
num stills until its specific gravity becomes 1.84, when it
is ready for the market.
182. Properties. — Commercial sulphuric acid has an
oily appearance, and was formerly prepared by distilling
green vitriol or ferrous sulphate : owing to these facts it
received the name oil of vitriol.
When exposed to the atmosphere it soon absorbs mois-
THE SULPHUR OXAC1DS. 173
ture, thereby becoming dilute. In consequence of its
great hygroscopic power, it is employed under the receiver
of the air-pump to aid in concentrating aqueous solutions
of such substances as would not bear heating without
undergoing decomposition. Pumice stone moistened with
sulphuric acid is used to dry those gases upon which the
acid has no action. The pure acid may also be used in a
wash-bottle.
QUERIES. Eor which gases already considered may it be used ? For
which ones should it not be used ?
When sulphuric acid is brought together with water in
quantities proportional to their molecular weights, the
hydrate of sulphuric acid, H2SO4 + H2O, is formed.
When this acid mixes with water much heat is evolved.
In diluting it with water it is best slowly to add the acid
to the water, and not the water to the acid, otherwise the
vessel containing the acid may be broken and a serious
accident ensue.
EXP. 122. Try the effect of strong sulphuric acid upon a
splinter of wood ; a bit of cloth ; a lump of sugar. What
occurs ?
Sulphuric acid chars vegetable substances by abstracting
water, or the elements of water, hydrogen and oxygen.
In its industrial uses, sulphuric acid is employed very
extensively in the manufacture of soda (sodium carbonate,
Na2CO3), artificial fertilizers, nitroglycerine, etc., and in
the refining of petroleum.
QUERY. Eor what purposes has sulphuric acid thus far been employed
in the laboratory ?
183. Tests for Sulphuric Acid and the Sulphates. —
1. Sulphuric acid or a soluble sulphate may be detected by
174 THE SULPHUR OX ACIDS.
adding to the solution barium chloride, BaCl2, when the
white precipitate, barium sulphate, BaSO4, is obtained.
This precipitate is insoluble in acids.
2. An insoluble sulphate may be fused on platinum foil
or a bit of porcelain with sodium carbonate ; the moist-
ened residue produces no spot on silver. If fused in the
same way on charcoal a spot will be produced.
QUERIES. If a sulphate, when treated on charcoal with sodium carbon-
ate, yields sodium sulphide, Na2S, what action upon the acid has occurred ?
If the black spot on silver be Ag2S, what other compound is probably
formed in the reaction : —
2 Ag + Na2S = Ag2S + . . . ?
(Suo. H and 0 are present in H20 to unite with Na.) Write this equation
in full, and balance. How can you distinguish a sulphate from a sulphide,
by fusing, etc. ?
NORDHAUSEN, OR FUMING SULPHURIC ACID, H2S2Or.
184. This acid is made by heating dried ferrous sulphate
which still contains a little moisture. The reaction is
represented thus : —
4 FeSO4 + H2O = 2 Fe2O3 -f 2 SO2 + H2S2O7.
It may also be made by passing sulphur trioxide, SO3,
into strong sulphuric acid : —
H2SO4 + SO3=H2S2O7.
It breaks up readily, forming sulphur trioxide and sul-
phuric acid. When a vessel containing it is opened, fumes
of the trioxide escape ; hence it is called fuming sulphuric
"Water acts violently upon it, converting it into ordinary
sulphuric acid : —
H2S2O7 + H2O = 2 H2SO4.
THE SULPHUK OX ACIDS. 175
The principal uses of this acid are for dissolving indigo
in the process of dyeing Saxony blue and for manufactur-
ing the coal-tar colors.
QUERY. Since H2S207= H2S04+ S03, should this acid be regarded as
a distinct acid or as a solution of S03 in H2S04 1
THIOSULPHURIC ACID, H2S2O3.
185. This acid, in a free state, is so unstable that its
existence is somewhat problematical; but its salts, the
thiosulphates, are well-known articles of commerce. The
principal one, sodium thiosulphate, Na2S2O3, is used by
photographers as a solvent for the unchanged silver salts
in their prints, which are thus " fixed," as the process is
termed. This salt is formed by fusing sodium sulphite
with flowers of sulphur, thus : —
Na2SO3 + S = Na2S2O3.
When a thiosulphate in a hot solution is treated with
hydrochloric acid or sulphuric acid, free sulphur is de-
posited, and sulphur dioxide fumes evolved, thus : —
Na2S2O3 + 2 HC1 = 2 NaCl + S + SO2 + H2O.
QUERY. How does a sulphite behave with hydrochloric acid ?
NOTE. This sulphur acid was formerly known as hyposulphurous
acid, and its salts as hyposulphites ; while the stcid of the formula H2S(X
was called hydrosulphurous acid, and its salts hydrosulphites. Sodium
thiosulphate is still commonly known to druggists as hyposulphite of
sodium.
186. Tests for the Thiosulphates. — 1. With hydro-
chloric acid their solutions yield a precipitate of sulphur,
and give off sulphur dioxide fumes.
2. Barium chloride, when added to a solution of a thio-
sulphate, yields a white precipitate soluble in hydrochloric
acid, but leaving a residue of sulphur.
176 THE SULPHUR OXACIDS.
187. To distinguish between the Soluble Salts of
the Sulphur Acids. — The solution may contain a sul-
phide, a sulphite, a sulphate, or a thiosulphate. There
are many ways of making this distinction, one of which is
as follows : —
1. Evaporate a portion of the solution to dryness, and
fuse on charcoal with sodium carbonate, etc. A black
spot on silver indicates any of these acids. Then fuse on
porcelain, etc. ; no spot indicates a sulphate.
2. To a portion of the solution add silver nitrate,
AgN03:-
(a) A black precipitate formed at once indicates a sul-
phide.
(6) No precipitate indicates a sulphate.
(<?) A white precipitate, obtained by adding a single
drop of the silver nitrate, and which does not dissolve upon
shaking, indicates a sulphite. This precipitate, upon stand-
ing, or upon being heated, turns black, metallic silver
being the final product obtained.
(d) A white precipitate from a single drop of the nitrate,
which dissolves upon shaking, indicates a thiosulphate.
Add an excess of nitrate, and boil. A black precipitate,
Ag2S, is finally obtained.
(e) If the student is still in doubt as to whether the
solution contains a sulphite or a thiosulphate, add hydro-
chloric acid to a fresh portion of the solution; sulphur
dioxide fumes from a clear solution indicate a sulphite;
the same fumes from a clouded solution indicate a thio-
sulphate.
SUG. Try to distinguish these acids by means of barium chloride,
BaCl2, etc.
SULPHUR AND CARBON. 177
SULPHUR AND CARBON.
188. Carbon Bisulphide, CS2, is the only known com-
pound of sulphur and carbon. This is a colorless, inflam-
mable, highly refracting liquid, boiling at +46°, and
possessing a specific gravity of 1.292. It has a powerful
odor, in its impure commercial forms, and its fumes are
poisonous ; when pure it has a pleasant, ethereal odor.
It is prepared by passing the vapor of sulphur through
a cylinder heated to redness and containing charcoal.
Carbon bisulphide is employed for a variety of purposes.
In the laboratory it is used as a solvent for bromine
and iodine, as we have previously seen ; in the manufac-
tures it is employed as a solvent for various gums, such as
rubber gum ; shoemakers mend shoes with a cement made
by dissolving crude rubber in carbon bisulphide ; in woollen
manufacture it is used to regain the oils with which the
wool is treated during some of the necessary processes ;
in optics the hollow prisms used for decomposing light,
and for spectrum analysis, are filled with carbon bisul-
phide ; in agriculture it is employed as an insecticide,
and (in the form of salts) in combating the phylloxera.
It is also said to be of value in exterminating woodchucks
and other burrowing animals, for which purpose it is placed
in their burrows, which are then tightly closed with earth.
The odor of carbon bisulphide betrays its presence, and
serves as a test.
SELENIUM.
SYMBOL, SE". — ATOMIC WEIGHT, 79. — SPECIFIC GRAVITY
(CRYSTALLINE), 4.3.
189. Occurrence. — Selenium is a rare element closely
resembling sulphur. It was discovered in 1817 by Berzelius
178 SELENIUM.
while examining the deposits of the sulphuric acid chambers at
Gripsholm. It does not occur native, but is found in the
selenides, such as lead selenide, PbSe, and the double selenides
of mercury, lead, silver, and copper.
190. Preparation. — Owing to the rarity of this element,
the student will probably do no work with it, therefore general
processes alone will be briefly given.
The residue of the sulphuric acid chambers is mixed with
potassium nitrate and then thrown into a red-hot crucible, where
it deflagrates, forming potassium selenate, K2SeO4, which is
now contaminated with many impurities contained in the
chamber residue. This impure mass is now digested with
hydrochloric acid, and the solution filtered and evaporated
nearly to dryness, selenious acid, H2SeO3, being formed. This
acid is then treated with sulphurous acid, thus : —
H2SeO3 + 2 H2SO3 = 2 H2SO4 + H2O + Se.
The finely divided selenium thus produced is separated by
filtration.
191. Properties. — Finely divided selenium when viewed
by transmitted light has a reddish color. In its properties and
compounds it resembles sulphur. It is known in three modifi-
cations ; viz., amorphous, vitreous, and crystalline. Flowers
of selenium, a scarlet powder, is obtained in a manner similar
to flowers of sulphur.
The specific gravity of selenium varies from 4.5 to 4.8.
192. Selenium Compounds. — 1 . Selenium and hydrogen
form hydrogen selenide, H2Se, a poisonous gas obtained by the
direct union of the vapor of selenium with hydrogen, or by
treating potassium selenide with hydrochloric acid.
2. Selenium and oxygen form selenium dioxide, SeO2, when
the former is burned in a current of the latter, or by treating
the former with strong nitric acid.
TELLURIUM. 179
Selenium dioxide and water form selejiious acid, H2SeO3,
from which the selenites may be derived. Selenium dioxide
has the odor of rotten cabbage or horseradish.
3. Selenic acid, H2SeO4, is obtained by passing a stream of
chlorine gas through water in which finely divided selenium is
suspended, thus : —
Se + 3 C12 + 4 H2O = 6 HC1 + H2SeO4.
This acid forms salts called selenates.
193. Tests for Selenium and its Compounds. — 1. Free
selenium burned in the air gives the odor of the dioxide.
2. Hydrogen selenide is distinguished by its very offensive
odor. It causes inflammation of the eyes and seriously affects
the lining membranes of the nose.
3. The selenides when heated on charcoal give the dioxide
fumes ; when fused with potassium nitrate, and when the solu-
tion of the residue in hydrochloric acid is treated with sulphur
dioxide, they yield free selenium.
4. The seleuites when heated on charcoal also give the fumes
of burning selenium ; their solutions with sulphur dioxide yield
free selenium.
5. The selenates, with sulphur dioxide, yield free selenium
when acidulated with hydrochloric acid. The fumes of a
selenate heated on charcoal are also those of the dioxide.
TELLURIUM.
SYMBOL, TE". — ATOMIC WEIGHT, 128.? — SPECIFIC
GRAVITY, 6.24.
194. Occurrence. — Tellurium is a rare element which
occurs native in small quantities and in combination with
certain metals, as tellurides, particularly with gold, silver, lead,
and bismuth.
180 TELLURIUM.
195. Preparation. — Tellurium is prepared by mixing bis-
muth telluride (which also contains some sulphur as an im-
purity) with sodium carbonate and oil ; this mixture is rubbed
to a paste, placed in a closed crucible, and strongly heated.
The mass is then lixiviated with water, when a solution of
sodium telluride and sulphide is obtained. Upon exposure to
light and air tellurium, in the form of gray powder, is deposited
in this solution ; this powder is purified by distilling it in an
atmosphere of hydrogen.
196. Properties. — Tellurium is a very brittle, bluish-
white solid, possessing a metallic lustre, and a specific gravity
of 6.24. It burns in the air with a bluish flame, giving white
fumes of tellurium dioxide.
197. Compounds. — 1. Hydrogen telluride, H2Te, is a very
poisonous gas resembling hydrogen sulphide. It is prepared
thus : —
ZnTe + 2 HC1 = ZnCl2 + H2Te.
It burns with a blue flame, is soluble in water, and forms the
tellurides.
2. Tellurium dioxide, TeO2, is obtained by burning the metal
in the air or in oxygen. It also occurs native in tellurite. When
melted it forms a light-yellow liquid.
3. Tellurous acid, H2TeO3, is formed by dissolving the metal
in dilute nitric acid and pouring the liquid into water.
4. Tellurium trioxide, TeO3, is prepared by strongly heating
telluric acid, thus : —
H2TeO4 = H2O + TeO3.
This oxide is an orange-}7ellow crystalline solid.
5. Telluric acid, H2TeO4, is produced by oxidizing tellurium
with potassium nitrate.
198. Tests for Tellurium and its Compounds. —
1. Free tellurium, when dissolved in strong sulphuric acid,
EXEKCISES. 181
forms a purplish-red solution, from which tellurium may be pre-
cipitated by adding water.
2. Tellurium in any compound may be detected by mixing
with sodium carbonate and a little charcoal dust, after which it
is placed in' a sealed tube and heated to redness. When cool
the tube is broken and the contents dissolved in hot water.
Sodium telluride, Na2Te, is dissolved out, coloring the water
purple. Upon standing, free tellurium is deposited.
3. Tellurates are first heated to redness, whereby they are
reduced to tellurites. The tellurites when dissolved in hydro-
chloric acid and afterwards treated with sulphurous acid yield
tellurium.
EXERCISES.
1. In what experiment did sulphur unite directly with a metal to form a
sulphide ? In how many ways may a salt be formed ?
2. What varieties of sulphur may be purchased at the drug store ?
(Suo. Ask your druggist what varieties he has for sale, and by what
names they are known.)
3. Try to obtain sulphur from a piece of vulcanized rubber.
4. Try to prepare H2S from various sulphides that you may find in
the laboratory. Use H2S04, HC1, and HN03. Try "Fool's Gold" or iron
pyrites. If the acids do not give the desired results, fuse the pyrites on
charcoal with sodium carbonate and again apply the acids. Do you thus
obtain H2S ? Why ?
5. The amount of hydrogen sulphide in a solution, as in mineral water,
may easily be determined by titration. For this purpose a standard solu-
tion of iodine and a fresh solution of starch paste (an indicator) are
required. The standard solution is prepared thus : —
Weigh out in a small corked vial (weighing flask) about Is of pure
iodine ; then dissolve about 5s potassium iodide in 20CC distilled water ;
uncork the vial and immerse it in the iodide solution. When the iodine
is dissolved, dilute with water so that lcc of the standard solution shall
contain lms of free iodine ; preserve this in a perfectly corked bottle in a
dark place.
The titration is made thus : To 100CC of the water to be tested add
about 2CC starch paste, and then, in the usual manner, add the standard
solution of iodine, until a permanent light-blue color is reached. The
182 EXERCISES.
number of cubic centimetres standard solution required (N) equals the
number of milligrams of iodine required to decompose the hydrogen
sulphide : —
It is usually safe to deduct from N 1 or 2mss to allow for the iodine
required to color the starch paste, although this is best determined by
trial. (As soon as the H2S is decomposed, upon what does the I act?
What causes the blue color?)
The computation is made thus : —
254 : N : : 34 : x = wt. of H2S in 100CC.
In case the amount of H2S per litre is required, it = 10 x. Why ? How
obtain the number of cubic centimetres of H2S gas per litre ? Whence
come the numbers, 254 and 34 ? How compute the number of cubic inches
of H2S per U.S. gallon *
6. Coal containing much sulphides is not adapted to reducing iron
from its ores. Why ? Sometimes the sulphides are oxidized to sulphates,
which are not so objectionable, by piling coal in heaps exposed to the air : —
7. Try to obtain a sulphate by treating sulphur or a sulphide with a
mixture of KC1O3 and HN03. Test for the sulphate with BaCl2.
8. Sulphuric acid or a sulphate is determined quantitatively, thus : To
(say) 50CC of the solution containing a sulphate (e.g., K2SOJ add hydrochloric
acid and boil ; while hot add an excess of barium chloride and thoroughly
agitate : —
BaCl2 + K2SO4 = BaS04 + 2 KC1.
Now filter out the BaS04 and thoroughly wash with much hot water; the
ash of the filter-paper used should be known ; the precipitate and filter-
paper are now carefully dried and the precipitate carefully transferred (as
completely as possible) to a weighed porcelain crucible; the filter-paper is
now burned and its ash placed within the crucible, which is then heated to
redness; when the crucible is cool its weight (W) is determined: —
W — wt. of cruc. — wt. of filter-ash = wt. of BaSO4.
Sometimes the chemist estimates the anhydride of an oxacid. How
much S03 in 17.241s BaS04? With how much potassa, K2O, will this
amount of S03 unite, and how much K2S04 will it yield ?
9. The salts of many acids are decomposed and their acids set free by
sulphuric acid. Why ?
EXEKCISES. 183
Complete and balance these equations: —
KN03 +H2S04= . . .
NaCl + H2S04 = . . .
Na2C03+H2SO4 = . . .
10. Make a table showing the similarity of the formulae of the oxides
and acids of S, Se, and Te.
11. AN EXERCISE IN VALENCE. If to the number representing the
valence of an element we assign a positive or negative sign, we shall find
that the algebraic sum of these numbers in any stable chemical compound
always equals zero, — provided we take : —
1. The number for H = + 1.
2. The number for 0 = — 2.
3. The number for any metal in combination as + (except As, Sb,
etc., with H).
We may utilize these data to determine the valence of an element in
combination ; e.g., what is the valence of I in HI03 ?
SOLUTION. O3 = 3 X — 2 = — 6. H — + 1. Now the question simply
is, what number must be added to the + 1 to make +6 (or a number which
added to —6 will give 0). The number required is evidently + 5. Accord-
ingly we may conclude that I in HIO3 is a pentad.
QUERIES. What is the valence of S in the following compounds :
H2S; H2S02? H2S03; H2SO4; H2S2O3? Of Cl in: HC1; HC10; HC102;
HC103; KC1(V Of Brin: HBr; HBrO; HBr03; HBr(V Of P, Si, B,
and N in their compounds ?
SUG. Read Johnson on Oxidation in Douglas and Prescott's Qualitative
Analysis, pp. 251-253.
CHAPTER XII.
SILICON AND BORON.
SILICON.
SYMBOL, Siiv. — ATOMIC WEIGHT, 28. — SPECIFIC
GRAVITY, 2.49.
199. Occurrence. — Silicon is a very abundant element,
occurring in combination with oxygen, or with oxygen and
other elements. Silica, SiO2, known under the names
quartz, sand, agate, etc., is a very widely distributed sub-
stance, found in every geological formation.
The silicates, such as feldspar, mica, and certain clays,
are well-known compounds. Silicon constitutes from 22.8
to 36.2 per cent of the earth's crust.
In a free state, it may be prepared in three modifications,
— amorphous, graphitoidal, and crystalline.
200. Preparation. — EXP. 123T. Silicon may be obtained
by heating in an iron tube potassium hydrofluosilicate, K2SiF6,
with metallic sodium or potassium : —
K2SiF6 + 4 K = 6 KF + Si.
A violent reaction occurs. When cool the fused mass is
treated with water to dissolve the potassium fluoride, while the
silicon remains as a brown amorphous powder. (See Art. 208
for K2SiF6.)
EXP. 124 T. Place in a porcelain crucible a small quantity
of amorphous silicon. Carefully lute on the cover with a paste
SILICON AND OXYGEN. 185
of wood ashes, and after thoroughly drying heat the crucible,
gently at first, and finally to redness. The amorphous mass
contracts, becoming denser, and assuming the form of plates of
graphite.
EXP. 125T. Crystalline silicon is best obtained by the fol-
lowing method : A mixture of 3 parts dry sodium Irydrofluo-
silicate, Na2SiF6, and 1 part sodium cut in pieces, is rapidly
introduced into a hessian crucible heated to bright redness.
Then 9 parts well dried granulated zinc are rapidly added ; and
finally, the whole covered with a layer of dried sodium chloride.
The crucible is then closed, the fire allowed to go down, and
the crucible allowed to cool in the furnace. The silicon under
these circumstances crystallizes from its solution in molten
zinc, and the zinc afterward solidifies, enclosing the crystals of
silicon. By treating the mass with hydrochloric acid the zinc
is dissolved and the crystals left behind.
201. Properties. — Amorphous silicon, as obtained above,
is inflammable in the air, when strongly heated, producing
silicon dioxide. The graphitoidal form is not so readily
inflammable. At a high temperature, and in absence of
oxygen, silicon can be fused. Hydrochloric acid does not
dissolve it, but in strong alkalies it is soluble, thus r —
Si + 2 KOH 4- H2O = K2SiO3 + 2 H2.
NOTE. The student will not meet with free silicon in his work unless
lie prepares it or buys it as such, when he can examine it, ignite it, and
test for silicon dioxide. Art. 205.
SILICON AND OXYGEN.
SILICON DIOXIDE, OK SILICA, SiO2.
202. Occurrence. — Silicon and oxygen form one well-
known compound, which occurs in many modifications, as :
186 SILICON AND OXYGEN.
1. Quartz crystals, glassy hexagonal prisms terminating in
hexagonal pyramids.
2. Amethyst, smoky quartz, rose quartz, and chrysoprase,
colored varieties of quartz.
3. Quartzite, a sedimentary rock.
4. Sand and sandstone, fine fragments of quartz more or
less cemented together.
5. Honestone or novaculite, a fine-grained quartz rock.
6. Chalcedony, a mixture of crystalline and non-crystalline
quartz.
7. Agate, consisting of layers of crystallized and amorphous
quartz of various colors.
8. Flint and chert, a coarse variety of chalcedony.
9. Opal, a hydrated form of silica.
10. Various modifications of the above in which one form is
passing into another.
203. Preparation. — Silica may be artificially obtained
in two forms : as the so-called " soluble silica," and as an
insoluble powder.
EXP. 126 P. Melt in a crucible, 6g each, potassium carbonate
and sodium carbonate ; then add 3g pulverized quartz or white
sand, and heat till the whole is melted. The molten mass is
now to be poured out and dissolved in dilute hydrochloric acid.
The solution thus obtained is now placed in a tray (dialyzing) ,
which may be prepared by stretching parchment paper over a
wooden hoop, s% 10cm in diameter. This tray is now floated
on a tub of pure water, when the hydrochloric acid and saline
substances of the solution pass through the parchment into the
water of the tub, while the soluble silica remains in the tray.
It will take about four days to effect this separation, and there
must be much water in the vessel on which the tray is floated,
or it must be often changed.
SILICON AND OXYGEN. 187
NOTE. This method of separation is called Dialysis, and depends upon
the fact that crystallizable substances will pass through the parchment,
while colloid or non-crystallizable substances will not pass through.
In this manner a colorless, tasteless, limpid solution is
obtained, which may be concentrated in a generating-flask ;
but if the concentration be carried too far, the solution
becomes of a jelly-like consistency.
Though we here speak of having silica in solution, the
substance dissolved is really a form of silicic acid, probably
ortho-silicic acid, H4SiO4. This loses water very readily,
and is converted into meta-silicic acid, H2SiO3, and this,
when dried, loses more water, and passes into silicon di-
oxide, SiO2.
NOTE. The relations between silicon dioxide and silicic acid, H2Si03,
are similar to those existing between carbon dioxide and carbonic acid.
Student will indicate the points of resemblance.
EXP. 127 P. Evaporate strictly to dryness (in an evaporat-
ing-dish) a portion of the solution obtained in the last experi-
ment. The powder thus obtained is pure silica. Is it now
soluble in acids? In alkalies?
204. Properties. — Natural crystals of silicon dioxide or
quartz are of a glassy lustre, and rank 7° in the scale of
hardness. They present no cleavage, and a conchoidal
fracture. The specific gravity of quartz is 2.6 ; of tridy-
mite, another form, it is 2.3.
All forms of silica are somewhat soluble in alkalies,
especially when digested under pressure ; consequently
many waters, such as those of the Hot Springs in Arkan-
sas, and the geysers of Iceland, contain, in solution, silica,
which is deposited upon standing. This explains the
existence of siliceous sedimentary rocks, like quartzite,
etc., and of the siliceous petrifactions which so frequently
occur, especially in the rocks of the Cretaceous Period.
188 THE SILICON OXACIDS.
Tripoli is the siliceous remains of the shells or valves of
microscopic plants, — the Diatoms.
Sandstone is composed of fragments of quartz cemented
together by deposited silica ; while
Conglomerates are larger pebbles similarly joined. Arti-
ficial conglomerate is now used as a building stone.
The many different colors which quartz assumes are due
to the fact that soluble silicon compounds readily absorb
coloring matters. These colors are either destroyed or
changed upon application of heat.
Some forms of quartz, owing to their hardness, and sus-
ceptibility to a high polish, are prized as ornaments.
Agates are somewhat porous; when soaked in honey,
then treated with sulphuric acid, and afterwards polished,
they exhibit curious and beautiful markings.
205. Tests for Silicon Dioxide. — The student will soon
learn to recognize any of the natural forms of silica by
their appearance when crystallized, and by their hardness
and fracture. (Also see tests for Silicates.)
THE SILICON OXACIDS.
206. The silicon acids are hardly known in the free
state, being very unstable like carbonic acid. Notwith-
standing the instability of the acids of this series, there
are three well-marked classes of salts which we may fairly
suppose to be derived from these acids : —
1. The mono-silicates.
2. The bi-silicates.
3. The tri-silicates.
Wollastonite, CaSiO3, and steatite, Mg3H2(SiO3)4, are
THE SILICON OXACIDS. 189
examples of the first; serpentine, Mg3Si2O7, and ortho-
clase, Al2K2(Si3O8)2» are examples of the second and third.
Besides these there are known many polymeric forms of
each of these classes.
SUG. Read R. and S., Vol. I., p. 573.
The various forms of silicic acid may be regarded as
derived from the acid H4SiO4 by abstraction of water in
different proportions. The simplest case is represented
thus : —
H4Si04 - H20 = H2Si08,
the salts of the acid thus formed being the monosilicates.
Then we have : —
2 H4SiO4 - H2O = HsSi2O7,
from which the bisilicates are derived ; and, finally,
3 H4SiO4 - 4 H2O = H4Si3O8,
from which the trisilicates are derived.
207. Tests for the Silicates. — Fuse the solid substance
with sodium carbonate on charcoal; dissolve the fused
mass in hydrochloric acid, and evaporate the solution to
dryness. If a white powder (SiO2), insoluble in hydro-
chloric acid, and soluble in potassium hydroxide, be
obtained, silicic acid, or some of its derived forms, is
present.
OTHER COMPOUNDS OF SILICON.
208. Silicon may be made to unite with nearly all the
elements previously considered, but their compounds are
unimportant. We may mention here that ; —
1. Silicon hydride, SiH4, is a gas prepared by acting
upon an alloy of magnesium and silicon with very dilute
hydrochloric acid, in the absence of air.
190 BORON.
If this gas be allowed to escape through water in
bubbles, each bubble, upon coming in contact with the
oxygen of the air, ignites spontaneously, forming ring-
shaped clouds of silicon dioxide.
2. Silicon fluoride, as we have previously seen, is ob-
tained by acting upon glass or silicon with hydrofluoric
acid (Art. 134).
3. Hydrofluo silicic acid, H2SiF6, is prepared when silicon
fluoride is dissolved in water : —
3 SiF4 + 4 H2O = H4SiO4 + 2 H2SiF6.
The sodium or potassium salts of this acid may thus be
prepared : —
EXP. 128T. Silicon fluoride is first prepared by treating in
a generating-flask sand and fluorspar, CaF2, with sulphuric
acid. .This gas is led into water, thus forming a solution of
hydrofluosilicic acid. When potassium or sodium carbonate is
added to this solution, a precipitate of the sodium or potassium
salt is obtained. Care must be taken to avoid an excess of the
alkaline carbonate, as the salts of hydrofluosilicic acid are de-
composed by alkalies.
BORON.
SYMBOL, B'". — ATOMIC WEIGHT, 11. — SPECIFIC GRAVITY
(CRYSTALS), 2.5.
209. Occurrence. — Boron occurs only in combination
with other elements. The chief compounds are boric
acid, H3BO3; borax, Na2B4O7 + 10 H2O ; and boracite,
2 Mg3B8015, MgCl2.
210. Preparation. — Boron may be prepared in two
modifications, viz., amorphous and crystalline.
BORON. 191
EXP. 129 T. Amorphous boron, a dark-brown, odorless,
tasteless powder, may be obtained by heating boron trioxide,
B2O3, with metallic potassium in an iron tube.
EXP. 130 T. Crystalline or adamantine boron is obtained by
fusing amorphous boron, in the absence of air, with metallic
aluminium.
This modification of boron ranks 9° in the scale of hard-
ness, and its crystals are prisms or monocliiiic octahedra.
211. Boron Compounds. — 1. Boron trioxide, B2O3, is
the only known oxide of boron, and may be obtained by
heating to redness boric acid, H3BO3. It is a brittle, glassy
solid, readily uniting with water to form boric acid.
SUG. Write the equation.
2. Boric acid, H3BO3, occurs dissolved in the waters of
certain lagoons in Tuscany, and the market is mostly sup-
plied from that source. In the vicinity of these lagoons
are volcanic jets of steam, whose heat is used to evaporate
the water containing the acid, which is thus obtained in
crystals; its purification is effected by re crystallization
from a water solution.
There are in California several dried up lake beds
containing massive borax, said to be sufficient to supply
our wants. Here the acid is obtained by treating the
borax with hydrochloric acid, and dissolving in hot water.
From this solution boric acid is also obtained by crystal-
lization.
Boric acid is soluble in water and in alcohol. It forms
the borates.
212. Tests for Boric Acid and the Borates. — 1 . When
in solution, the free acid turns a strip of turmeric paper
192 EXERCISES.
brown, and this color is not changed by dilute hydrochloric
acid, as is the case with the alkalies.
NOTE. It is best for the beginner to compare the action of an alkali
on this paper with the action of boric acid, noting how the hydrochloric
acid affects the colors. Also dip a piece of turmeric paper in boric acid ;
then moisten with Na2C03 and note the greenish-black color produced.
2. When a solid borate is heated on a platinum loop
in the reducing flame, the flame is tinged green.
NOTE. This test is most striking when the solid has first been calcined,
then dipped in sulphuric acid and heated to expel the acid, and finally
moistened with glycerine and treated as in 2.
EXERCISES.
1. Silicates in solution are estimated quantitatively as follows : The
solution is acidulated with hydrochloric acid and evaporated strictly to
dryness, without allowing the temperature to rise sufficiently high to cause
the silica, Si02, obtained again to unite with any bases present. The
residue is again treated with hydrochloric acid ; the white insoluble powder
Si02 is next removed by filtration, and in a manner similar to that
employed in estimating the sulphates directly determined as silica.
2. Soak bits of agate in honey ; treat with sulphuric acid, and polish
on a grindstone or emery wheel. The peculiar markings of the agates are
thus brought out.
3. Unite the edges of broken bits of glass with the so-called " soluble
silica " ; allow the mended articles to dry for two days ; then test the
strength of the silica as a cement.
4. What is glass ? Write an essay on the manufacture of glass.
5. Ask a blacksmith for what purposes he uses borax. Ask him if a
mixture of salt and sand will answer as well. What is a flux ?
6. For what use does the barber employ borax ?
7. Does borax soften "hard " water 1 Try it.
8. What is the anhydride of boric acid, H3BO3 ?
9. Dissolve a little borax in HC1 ; then to the solution add alcohol.
Warm and ignite the alcoholic solution of boric acid thus obtained and
note the characteristic green flame.
10. The waters of all our streams abound in diatoms. Examine some
under the microscope.
CHAPTER XIII.
PHOSPHORUS. — ITS OCCURRENCE, COMPOUNDS, ETC. —
GENERAL EXAMINATION OF UNKNOWN SUBSTANCES
FOR ACIDS.
PHOSPHORUS.
SYMBOL, P. — ATOMIC WEIGHT, 31. — SPECIFIC
GRAVITY, 1.83.
213. Occurrence. — Owing to its great affinity for
oxygen, phosphorus, although widely distributed, never
occurs in the free state. Its principal compounds are
with calcium ; as, phosphorite, Ca3(PO4)2, and apatite,
8 Ca3(PO4)2 + CaClF. It also unites with iron to form
vivianite, Fe3(PO4)2 + 8 H2O. It is also found in the
igneous rocks, from whose disintegration our alluvial soils'
have been produced ; hence every fertile soil must contain
phosphates. These phosphates are taken up from the soil
by growing plants, of whose ripened seeds they form an
essential constituent. Again, animals consume the plant
and its seeds, and appropriate the phosphates for building
up the solid or inorganic portion of their bones ; and it is
from bones that the greater part of our commercial phos-
phorus is now obtained. Sombrerite, an impure form of
calcium phosphate, found in the island of Sombrero, is
another source of commercial phosphorus.
214. Preparation. — Phosphorus is obtained from the
ashes of burned bones. As a matter of economy, the bones
194 PHOSPHORUS.
are not directly turned, but are subjected to a preliminary
treatment, in order to save some of their other constitu-
ents. Thus they are either first digested with water,
under pressure, in closed vessels, in order to extract the
gelatine ; or they are distilled in closed retorts, the vola-
tile products (bone oil) being utilized to some extent;
while the remaining solid substance, or "bone black" is
used for clarifying sugar until worthless for that purpose.
In either case the remaining solid residue of the bones is
reduced to ashes by burning in the open air.
Bone ash, which consists largely of calcium phosphate,
Ca3(PO4)2, is first treated with sulphuric acid, when an
acid calcium phosphate, soluble in water, is obtained : —
Ca3(PO4)2 + 2 H2SO4 = CaH4(PO4)2 + 2 CaSO4.
This solution of " super-phosphate of lime," as it is usually
called, is then evaporated to dryness, and afterward heated
nearly to redness, when calcium meta-phosphate is ob-
tained : —
CaH4(P04)2 = Ca(P03)2 + 2 H2O.
The meta-phosphate is then intimately mixed with fine
charcoal dust, and heated to redness in earthen crucibles
placed in tiers inside of a furnace, their necks extending
outside of the furnace, and dipping under water in a con-
denser. Only one-half of the phosphorus is thus liber-
ated and condensed under the water. The phosphorus is
now removed, melted under water, and purified by strain-
ing through chamois leather under water, when it is cast
into the ordinary sticks of commerce. Before it is cast
into sticks, the phosphorus may be purified by treating
it with sulphuric acid and potassium dichromate, K2Cr2O7.
All the phosphorus contained by the bone ash may be
PHOSPHORUS. 195
liberated by mixing the m eta-phosphate with sand and
charcoal dust, after which it is treated as before. The
reactions are : —
1. 2 Ca(PO3)2 + 5C = Ca2P2O7 + 5 CO + 2 P.
2. 2 Ca(PO3)2 + 2 SiO2 4-100 = 2 CaSiO8 4- 10 CO 4- 4 P.
215. Properties. — Phosphorus is a highly inflammable
substance, taking fire at low temperatures. When exposed
to the air it slowly oxidizes, emitting a phosphorescent
glow, or luminous and evanescent flashes of light. A
slight blow or scratch is often sufficient to ignite it. It
burns with great heat, and when in contact with the flesh
it produces deep and painful wounds; hence great care
should be exercised in handling it. It should not be taken
in the hands nor cut in the air, but should be held by a
pair of forceps, and cut under water.
Phosphorus should always be stored, for safe keeping, in
a bottle of water fitted with a good cork to prevent the
water from evaporating ; the bottle should then be kept in
a tightly-covered can, and the whole placed in a cool
place.
Owing to the low temperature of its ignition, phosphorus
is employed in tipping the common lucifer match. The
composition of match-tips varies ; but nearly all the com-
pounds employed for making tips contain phosphorus,
sulphur, and potassium nitrate.
Phosphorus is also used as an ingredient of many ver-
min "exterminators," but about five-sixths of all the
phosphorus produced is consumed for making matches.
The fumes of phosphorus are characteristic, possessing
poisonous properties, and an odor with a faint resemblance
to garlic. When taken internally, phosphorus is a virulent
196 PHOSPHORUS.
poison; one decigram may produce fatal results. Severe
pains in the stomach, vomiting of substances with an odor
of garlic, and even the characteristic fumes emitted with
the breath, are the symptoms of phosphorus poisoning.
Turpentine is a proposed antidote.
Phosphorus is known in three different modifications,
viz. : —
1. Ordinary, or waxy phosphorus, the form usually seen
in sticks.
2. Crystalline phosphorus, obtained by dissolving the
common form in carbon bisulphide, and allowing the solu-
tion to evaporate.
3. Red, or amorphous phosphorus, obtained when either
of the other two modifications is heated to 240° in the
absence of the air. This variety is not so inflammable as the
ordinary phosphorus, nor does it give off poisonous fumes ;
hence it is sometimes used by the matchmakers, who thus
avoid the dreaded effects of phosphorus poisoning. The
specific gravity of this variety is 2.106.
All three varieties of phosphorus burn in the air with a
bright, luminous flame, forming dense white fumes of phos-
phorus pentoxide.
QUERY. Should an excess of phosphorus be employed in experiment
41, of what variety would the remainder be 1
216. Tests for Free Phosphorus. — 1. Phosphorus, in
considerable quantity, may be detected by its physical
properties and odor.
2. In minute quantity, as in cases of phosphorus pois-
oning, phosphorus is detected by dissolving in water the
substance to be tested, after which it is boiled in a gen-
erating-flask, and the steam is led through a glass con-
densing-tube into another flask containing cold water.
PHOSPHORUS AND HYDROGEN. 197
Now, if the room be dark, and if phosphorus be present, a
phosphorescent glow is noticeable at the point where the
steam condenses.
PHOSPHORUS AND HYDROGEN.
217. Phosphorus and hydrogen form three compounds : —
1 . Gaseous phosphoretted-hydrogen or hydrogen-phosphide, PH3.
2. Liquid phosplioretted-hydrogen or hydrogen-phosphide, PH2.
3. Solid phosplioretted-hydrogen or Irydrogen-phosphide, (P2H?).
Of these we shall consider only the first.
218. Gaseous Hydrogen Phosphide, or Phosphine,
PH3, is a gas which ignites spontaneously upon coming in
contact with the oxygen of the air, owing to the presence
of traces of the liquid compound PH2, this latter substance
being obtained by the same process that yields the former.
If the tube from which the phosphine escapes be bent
upward under water, each bubble upon reaching the air
ignites, forming beautiful ring-shaped clouds of phosphorus
pentoxide, P2O5. In a still atmosphere these clouds have
a peculiar rotary motion, illustrating what is known as
vortex motion. This striking experiment may be exhibited
thus : —
EXP. 131 T. In a generating-flask place a strong solution of
potassium hydroxide, KOH, and add several small pieces of
stick phosphorus. Now gently warm, and as soon as flames
begin to appear at the mouth of the flask, insert a cork carry-
ing a bent delivery-tube. The lower end of this tube is to dip
under water placed in an open vessel. As each bubble of the
gas comes into the air, it ignites with a slight report : —
4 P + 3 KOH + 3 H2O = 3 KH2PO2 + PH3.
198 PHOSPHORUS AND OXYGEN.
It is somewhat safer to put the apparatus together, and then
to pass hydrogen through it long enough completely to displace
the air ; or the air may be expelled by pouring a little ether
over the solution before warming. During the experiment cur-
rents of air in the room are to be avoided. Save the contents
of the generating-flask for work under hypophosphorous acid.
QUERIES. What is the object of these last precautions ? What other
gas "behaves like PH3? Show how PH4Br and PH4I are obtained from
PH3, HI, and HBr. Does PH3 form salts similar to NHS ?
In this experiment liquid hydrogen-phosphide may be
obtained by passing the gas through a suitable condensing-
tube, but both this and the solid form are of no impor-
tance to the beginner. None of the hydrogen phosphides
possess acid properties.
Sue. Make a list of the binary acids. Also make a list of the non-
acid hydrogen compounds of the elements previously considered. Which
one is alkaline ?
PHOSPHORUS AND OXYGEN.
219. There are two known oxides of phosphorus, viz. : —
1. Phosphorus Trioxide, P2O3.
2. Phosphorus Pentoxide, P2O5.
1. Phosphorus trioxide is formed when phosphorus is
burned in a limited supply of air. It is a white powder,
which possesses a garlic odor, and unites with water to
form phosphorous acid : —
3 H2O + P2O3 = 2 H3PO3.
2. Phosphorus pentoxide is obtained by burning phos-
phorus in the open air or in oxygen. It is also a white
powder, which eagerly unites with hot water to form phos-
phoric acid : —
THE PHOSPHORUS OXACIDS. 199
THE PHOSPHORUS OXACIDS.
220. There are three acids in this series : —
1. Hypophosphorous acid . H3PO2,
2. Phosphorous acid . . . H3PO3,
3. Phosphoric acid . . . H3PO4,
from which are derived : —
a. Metaphosphoric acid . . HPO3,
b. Pyrophosphoric acid . . H4P2O7.
Since the last two acids may be derived from phosphoric
acid, all three will be treated tinder one article, after
the consideration of the first two acids in the series.
HYPOPHOSPHOKOUS ACID, H3PO2.
221. EXP. 132 p. In a generating-flask place 10CC of a
solution of barium hydroxide, Ba(OH)2, and add two or three
small pieces of phosphorus. Add a little ether and boil until
the following reaction is completed : —
3 Ba(OH)2 + 2 P4 + 6 H2O = 3 Ba(H2PO2)2 -f 2 PH3.
The remaining solution is now to be filtered, when the barium
hypophosphite is obtained in clear solution. To this solution
carefully add dilute sulphuric acid to precipitate the barium,
when hypophosphorous acid is obtained, thus : —
Ba(H2PO2)2 + H2SO4 = BaSO4 + 2 H3PO2.
This acid is a colorless liquid, oxidizing to phosphorous
and phosphoric acids, when standing exposed to the air.
It is mono-basic, only one atom of its hydrogen being
displaceable. If we represent by M; any univalent metal,
the general formula for a hypophosphite may be repre-
sented thus: M'(H2PO2).
The hypophosphites may be prepared as in Exp. 131 T,
by boiling phosphorus with an alkali. The principal use
200 THE PHOSPHORUS OXACIDS.
of these salts is for medicinal purposes. The acid and its
salts are strong reducing agents.
222. Tests for Hypophosphorous Acid and the Hypo-
phosphites. — 1. The acid or its salts when heated in a
test-tube yield phosphine, PH3.
2. With silver nitrate a solution of the acid or its salts
gives a white precipitate, which soon changes to brownish-
black : —
4 AgN03 + H3P02 + 2 H20 = 4 HNO3 + H3PO4 + 4 Ag.
3. To the solution of this acid or of its salts add an ex-
cess of cupric sulphate, CuSO4; an insoluble hydride of
copper, CuH, is formed. Boil a short time ; hydrogen is
liberated and metallic copper is obtained.
NOTE. No. 3 distinguishes H3P02 from H3PO3. Thus test the latter.
SUG. Try hypophosphorous acid, or a hypophosphite, with mercuric
chloride, HgCl2. Do you obtain metallic mercury ^ 'In which tests do
you find examples of reduction ? Write the equations for HgCl2 and
CuS04 with KH2P02.
PHOSPHOROUS ACID, H3PO3.
223. This acid may be obtained by passing chlorine gas
through a layer of melted phosphorus under water.
Phosphorus trichloride, PC13, is at first formed, and im-
mediately reacts upon the water, thus : —
PC13 + 3 H2O = H3PO3 + 3 HC1.
The hydrochloric acid is expelled by heat. If the addition
of the chlorine gas does not stop short of saturation, i.e.,
before the phosphorus has all disappeared, phosphoric acid
is produced. Indeed, it is difficult thus to obtain phos-
phorous acid free from traces of phosphoric acid.
THE PHOSPHORUS OXACIDS. 201
Phosphorous acid is generally dibasic, and M'2(HPO3)
may be taken as a general formula for the phosphites,
although there are some phosphites known in which the
acid is tribasic, all the hydrogen being displaced.
224. Tests for Phosphorous Acid or a Phosphite.—
1. To the solution add a few drops of sulphuric acid, and
then add potassium permanganate until a purplish tint is
reached. This color fades slowly in a cold solution, but
rapidly when heat is applied.
SUG. Thus try H3P02. How does it behave ? Also try H3P03 with
CuS04, as you tried H3P02. What results ?
2. To the solution add calcium hydroxide, Ca(OH)2;
a white precipitate is thrown down.
SUG. Thus try a hypophosphite. Do you obtain a precipitate ?
QUERY. How can you distinguish between a phosphite and a hypo-
phosphite 7
PHOSPHORIC ACID, H3PO4.
225. This acid is also known as orthophosphoric acid,
and its salts as the orthophosphates. It may be obtained
thus : —
EXP. 133 P. In an evaporating-dish place a small quantity
of red phosphorus, and add reagent nitric acid (sp. grav. 1.2) ;
now heat gently, adding more nitric acid, until the phosphorus
disappears and red fumes cease to come off. The evaporation
is to be continued until the excess of nitric acid is expelled.
The acid thus obtained is a thick, syrupy mass, free from
odor and readily soluble in water ; when allowed to stand,
rhombic, six-sided crystals are obtained.
Phosphoric acid is a tribasic acid, forming acid and normal
THE PHOSPHORUS OXACIDS.
salts, the phosphates. M/3PO4 is a general formula for the
phosphates.
Phosphoric acid is used in medicine, arid its salts are of
common occurrence and much used as fertilizers. The
phosphates are found in the blood and fluids of animals;
they are excreted from the kidneys as acid phosphate of
sodium and phosphates of calcium and magnesium. When
urea in urine decomposes a double salt of ammonium and
sodium, NaNH4HPO4, or microcosmic salt is formed. It was
from this source that, in 1669, Brandt first prepared phos-
phorus.
METAPHOSPHORIC ACID, PIPO3.
This acid is formed when orthophosphoric acid is heated
to 400°. It is the form in which phosphoric acid is com-
monly met with in the market (glacial phosphoric acid).
Its fortnation is illustrated thus : —
H3PO4 - H2O = HPO3.
At ordinary temperatures, in solution in w^ater, it is slowly
changed to orthophosphoric acid ; the change takes place
rapidly in boiling water.
Salts of metaphosphoric acid are formed by igniting
phosphates belonging to the class represented by the
formula M'H2PO4 as, for example : —
KH2PO4 - H2O = KPO3.
QUERY. In what process already considered does a transformation
from an orthophosphate to a metaphosphate take place 1
PYROPHOSPHORIC ACID, H4P2O7,
Is formed when orthophosphoric acid is heated at 200—
300°, until a small specimen neutralized with ammonia
gives a pure white precipitate with silver nitrate. The
change is : —
2 H3PO4 - H2O = H4P2O7.
THE PHOSPHORUS OXACIDS. 208
Its salts are formed by igniting phosphates of the order
M'2HPO4, thus: —
2 K2HPO4 - H2O = K4P2O7.
QUERY. In what connection have pyrophosphates been mentioned in
this book ?
226. Tests for the Phosphates or their Correspond-
ing- Acids. — 1. To the solution add a few drops of silver
nitrate, AgNO3.
(a) A light-yellow precipitate, soluble in ammonia, nitric
acid, and acetic acid, H(C2H3O2), indicates phosphoric acid
or its salts.
(£) A white precipitate, soluble in nitric acid (without
effervescence) and in ammonia, indicates pyrophosphoric
acid or its salts.
(ci) A .gelatinous white precipitate, soluble in nitric acid,
indicates metaphosphoric acid or its salts.
2. We may also distinguish metaphosphoric acid or its
salts by acidulating its solution with acetic acid and add-
ing the white of an egg, which immediately coagulates.
SUG. Try H3P04 and H4P207 with the white of an egg. What results ?
3. The most delicate test for orthophosphoric acid or its
salts is made by adding to the acid or to one of its salts
dissolved in nitric acid an excess of ammonium molybdate,
(NH4)2MoO4; upon heating, a yellow precipitate of am-
monium phospho-molybdate is obtained. See App. for
reagent ammonium molybdate.
NOTE. This test is sufficiently delicate to detect even very minute
traces of phosphoric acid or of the phosphates.
4. An orthophosphate with ammonium chloride, am-
monia and magnesium sulphate, gives a crystalline precipi-
tate of magnesium-ammonium phosphate, MgNH4PO4.
SUG. Try the phosphorus oxacids with salts of lead, calcium, barium,
and mercury. What results 1
204 EXAMINATION OF UNKNOWN SUBSTANCES.
EXAMINATION OF UNKNOWN SUBSTANCES FOR
ACIDS.
227. We have now learned something about the principal
inorganic acids. As we have already seen, Art. 79, some
elements are acid formers, others form bases ; and we may
now mention that there are still other elements — as, for
example, chromium and manganese — that are indifferent,
acting in certain compounds as acids, in- other compounds
as bases. The consideration of the acids of the indifferent
elements will be deferred for a time.
It frequently occurs that the chemist, while working,
comes upon substances entirely unknown to him ; and
among other things that he is called upon to determine are
the acids, which form essential constituents of all salts.
It is true that the substance may not be acid, but, as we
have previously seen, the salt of any acid yields the test for
that acid. Thus, KNO3 gives the test for nitric acid, and
NaCl the test for hydrochloric acid, etc. Now since there
are many acids, it is neither best nor profitable to test at
random for first one acid and then another ; some methodi-
cal plan should be followed. One method of procedure is
as follows : —
If the substance be in liquid form and neutral, evaporate
it to dryness or nearly so, carefully avoiding a high heat,
which might decompose certain unstable compounds and
drive off their acids in vapors. If the substance under
examination be a solid, no preliminary treatment is neces-
sary. If the substance in solution be acid, it is either a
free acid or an acid salt : in this case the solution must be
directly tested. Thus two cases naturally arise.
I. Let us suppose that the substance is neutral and a
EXAMINATION OP UNKNOWN SUBSTANCES. 205
solid, or, if a neutral solution, that we have evaporated it
to dryness. Proceed thus : —
Place a small portion of the substance in a test-tube ;
add sulphuric acid ; heat it gently, and note the results as
follows : —
1. If a rapid effervescence of an odorless, colorless gas
occur, the substance is probably a carbonate or an oxalate.
Now turn to the test for Carbonates or Carbonic Acid,
Art. 152, and try a fresh portion of the substance by all
the tests there given.
In case it prove not to be a carbonate, it is, very likely,
an Oxalate, a salt of the organic oxalic acid, H2C2O4.
This acid (in this connection) may be recognized by its
giving with calcium chloride, CaCl2, a white precipitate
of calcium oxalate, CaC2O4, soluble in hydrochloric acid,
but insoluble in acetic acid.
2. Slower effervescence of a colorless gas possessing odor,
(a) The odor of rotten eggs indicates a sulphide. Test
by Art. 169.
(&) The odor of burning matches; try for H2SO3, Art. 179,
or H2S2O3, Art. 186.
(c) Odor of peach blossoms; try for HCy, Art. 155.
(d) Odor of vinegar ; try for acetates, which are the
salts of acetic acid, HC2H3O2, thus : Dissolve the original
substance in water, add ferric chloride, Fe2Cl6, and boil. A
red solution of ferric acetate, Fe2(C2H3O2)6, is formed ; the
color is destroyed by adding hydrochloric acid.
(e) An irritating odor indicates HC1, Art. 96 ; HNO3,
Art. 75 ; or HF, Art. 135.
3. If a gas having a color and an irritating odor be
liberated, try for HI, Art. 127; HNO2, Art. 72; or HC1O,
Art. 104.
206 EXAMINATION OF UNKNOWN SUBSTANCES.
4. If a sudden explosion occur, try for HC1O3, Art. 108.
5. If none of the preceding phenomena occur, try for
H2SO4, Art, 183; H3P04, Art. 226; HPO2, Art. 224;
H4Si04, Art. 207; H3BO3, Art. 212; HIO3, Art. 131 ; or
HBrO3, Art. 121.
The student should remember that the foregoing data
are. valuable as indications only, arid that these indications
point toward certain acids to which he should refer, and
which he should try until he is satisfied that he has f6und
the right one.
II. If the solution be an acid one, proceed thus : —
1. To a portion of the solution add HC1; then add
BaCl2. If a white precipitate be obtained, the acid present
is H2SO4, since barium sulphate, BaSO4, is the only
barium . salt (except the salt formed with the rare acid
H2SiF6) which is insoluble in hydrochloric acid.
2. To a fresh portion of the solution add HNOs, and
then AgNO3.
The following acids give a precipitate insoluble in nitric
acid: HC1; HIj HBr ; H2S ; HCy; HC10 ; and the rarer
acids, hydro-ferro-cyanic acid, H4FeCy6, and hydro-ferri-
cyanic acid, H FeCy6. For these last two acids, see Iron.
3. Test in order for the following acids, using each time
afresh portion of the solution: HNO3; H2CO3; H3PO4;
H4Si04; H3BO3; H2S2O3 ; H2SO3 ; HNO2; H2C2O4 ;
H(C2H302); HC103.
If the acid is not discovered by working carefully up to
this point, it is a rare acid, and the student will be obliged
to try for all those previously mentioned in the text and
not mentioned above. It is true that the acid may be
quite a common one, belonging to the acids of the indiffer-
ent acid-forming elements, such as chromium, arsenic, or
manganese.
EXERCISES. 207
In such a case the student needs farther experience to
determine the acid. He will find directions under the
elements just named.
EXERCISES.
1. Phosphorus in iron ores, or in coal used in reducing iron ores, makes
the iron brittle. The presence of phosphorus in coal may be determined
by testing the ash for phosphates.
2. Make a list of the commonly occurring acids; also a list of the
rarer acids previously mentioned. In testing for acids a substance that
occurs native, would you expect to find rare acids ?
3. Dissolve the salt of an acid, and test with litmus paper; some salts
are acid, some are neutral, and some are alkaline. By trying many salts
and tabulating the results, the student may learn that normal salts may
belong to any of the three classes. Do any of the acid salts that you
have tried belong to the last two classes ?
4 If in an unknown solution NH3 and HNO3 be found, what salt is
present ? If Na and HC1, what salt ?
5, The student should be assigned many unknown (to him) salts and,
by reference to the text, he should determine the acids present. In this
way he will soon know the tests for the common acids. More than one acid
may be assigned in one solution, provided the acids do not decompose one
another, or their tests do not interfere. The metals of many metallic
salts obscure the test for the acids of the salts ; in this case the metals must
first be removed, as will hereafter be explained. Na, K, NH4, Ca, Mg, Sr,
and Ba do not thus interfere.
6. For an improved method of obtaining phosphorus, see Chemical
News, Apr. 4, 1879, p. 147.
CHAPTER XIV.
THE METALS.
INTRODUCTION.
228. The elements have been divided arbitrarily into
Metals and Non-metals, but the dividing line is nowhere
distinctly drawn. Certain elements, such as arsenic, anti-
mony, and bismuth, stand midway, in regard to their phys-
ical and chemical properties, between the two proposed
classes, and may be fairly placed in either ; consequently
we may justly consider the elements as forming but
one class with a regular gradation of properties. In view
of these facts it is impossible to give a strict and valid
definition of a metal ; but, in general, we may say : —
Definition. — A metal is an element which possesses a
peculiar lustre, known as a metallic lustre, especially when
in a solid or coherent condition, and the higher oxides of
which only, arid then in very few instances, are acid-forming
compounds.
SUG. All, or nearly all, of the oxides of the non-metals form acids.
State a few exceptions.
NOTE. Opacity, high specific gravity, and great atomic weight are
not exclusively characteristic of the metals.
229. Properties of the Metals. — Some of the metals
are barely known to exist, while others have been known
since the highest antiquity, and their properties have been
thoroughly investigated.
THE METALS.
209
Of the properties of metals we may note the following: —
(a) Specific Gravity. — As a rule the specific gravity
of a metal is greater than unity; only three — sodium,
potassium, and lithium — are less than 1.000. Osmium
(sp. grav. 22.48) is the heaviest metal, while lithium (sp.
grav. 0.59) is the lightest. (See Art. 25.)
QUERIES. With what are solids and liquids compared to determine
their specific gravities ? Gases 1 How is specific gravity determined 1
(5) Specific Heat. — The specific heat of an element is
equal to the number of thermal units required to raise
one kilogram of that element through 1° C.
The specific heat of any metal is less than unity, and
varies somewhat according to the temperature at which
the observation is made. The following observations,
which were made at 55°, will serve as an illustration: —
Cd 0.0567
Zn 0.0955
Ag 0.0570
Mn . . 0.1220
Co 0.1070
Ni 0.1080
Au 0.0324
Ft . . 0.0324
(f) Atomic Seat. — When the specific heat of any ele-
ment is multiplied by its atomic weight, a nearly constant
quantity (about 6.4) is obtained. This product, in the
case of any element, is termed the atomic heat of that ele-
ment. Take, for example, gold and zinc : —
0.0324 (sp.ht. of Au) X 196.5 (at.wt.of Au) = 6.4— (at. lit. of Au).
0.0955 (sp. ht. of Zn) x 65.0 (at. wt. of Zn) = 6.4 — (at. lit, of Zn) .
From an inspection of the results thus obtained was
deduced Dulong and Petit's law, viz. : —
The specific heat of an element varies inversely as the
atomic weight of that element.
This law is but approximately true, but so nearly true
210 THE METALS.
that it is the best method known for selecting the atomic
weights of some of the rarer metals. The accepted atomic
weights of indium, cerium, didymium, and lanthanum
were thus selected. For example, the atomic weight of
cerium as determined by analysis of its compounds was at
first assumed to be about 92 or 94, and the formulae of
its principal oxides were taken as CeO and Ce3O4. The
specific heat of cerium, however, was found to be 0.04479;
this would make the atomic heat about 4.2 instead of 6.4.
Accordingly the atomic weight was increased in the ratio
|, and the same oxides were assigned the formulae Ce2O3
and CeO.,. Both hypotheses agree equally well with the
percentage composition of these oxides.
QUERIES. In what different ways are the atomic weights of elements
determined ? If Ce = 94, what per cent of 0 is found in CeO and Ce304 ?
If Ce = 141, what per cent is found in Ce203 and Ce02 ?
(c?) The Conductivity of the metals for heat and elec-
tricity is greater than that of the non-metals or any of
the compounds of either.
(e) The Melting-points of the metals, so far as determined,
vary from — 39° to + 1090° C. Iron and cobalt fuse at
a white heat, platinum and iridium require the intense
heat of the oxy-hydrogen blow-pipe, while osmium has
not been fused at all. The determined melting-points of
a few metals are : —
Hg - 40°
G +30°
K + 62.5°
Na + 95.6°
Li +180°
Sn +235°
Bi . . . , +270°
Tl 4- 294°
Cd + 316°
Pb + 334°
Zn .....+ 423°
Sb + 425°
Ag +1000°
Cu . , +1090°
It is difficult to measure the temperature required to
THE METALS. 211
melt a metal whose fu sing-point is higher than that of
copper.
(/) The Molecular Heat of the Salts. — The molecular
heat of a metallic salt usually equals the sum of the atomic
heats of its constituent elements : e.g., the atomic heat of
potassium is 6.5, that of bromine is 6.7, while the molecu-
lar heat of potassium bromide equals 13.2 or 6.5 + 6.7.
The observed molecular heat of a salt agrees veiy closely
with the theoretical results thus obtained.
230. Alloys. — Metals mix in definite and in indefinite
proportions to form alloys, which possess properties both
like and unlike the properties of the metals composing
these alloys. These compounds are of great utility. The
following list gives the composition of some of the princi-
pal alloys : —
1. Gold Coin (U.S. ), 90 parts gold, 1 part silver and 9 parts copper.
2. Silver Coin (U.S.), 90 parts silver, 10 parts copper.
3. Brass, varying proportions of copper and zinc.
4. Britannia, varying proportions of brass, tin, antimony, and bismuth.
5. Pewter, 4 parts tin, 1 part lead.
6. Queen's Metal, 9 parts tin and 1 each of antimony, bismuth, and lead.
7. Solder, lead and tin in varying proportions.
8. Speculum Metal, 1 part tin, 2 parts copper.
9. Bell Metal, 18 parts tin, 22 parts copper.
10. Bronze, tin, copper, and zinc in varying proportions.
Bronze Coin, 95 parts copper, 4 parts tin, and 1 part zinc.
11. Type Metal, 1 part tin, 2 parts lead, and 1 part antimony.
12. German Silver, 5 parts copper, 2 parts nickel, and 2 parts zinc.
L3. Fusible Metal (melting at 93.75°), 1 part each of tin and lead, and 2 of
bismuth. This is called " Rose's Metal."
Fusible Metal (melting at 65°), 8 parts lead, 5 parts bismuth, 4 parts
tin, and 3 parts cadmium. This is "Wood's Alloy."
SUG. Name the uses of the alloys.
231. Amalgams. — Certain metals, such as silver, gold,
zinc, tin, copper, etc., unite with mercury to form amal-
212 THE METALS.
gams. Some of these amalgams are of great value in the
arts : battery zincs are amalgamated to prevent local cur-
rents and the needless waste of the zincs ; mirrors are
made by coating glass with a silver amalgam ; articles to
be electroplated are first slightly amalgamated to prevent
the plating from peeling ; gold and silver are extracted
from their ores by amalgamation ; etc.
QUERY. What uses does the chemist make of sodium amalgam 1
Amalgams are made in different ways : —
(a) By the direct union of the metal with mercury.
(Exp. 23.)
(7>) By adding metallic mercury to the solution of a
metallic salt : —
EXP. 134 P. To a solution of silver nitrate in a test-tube
add a drop of metallic mercury. Allow the tube to stand some
time. The splendid crystals formed are silver amalgam. These
cn'stals often assume an arborescent form, whence the name,
arbor Diance.
(c) By placing a metal in a solution of a salt of
mercury : —
EXP. 135 P. Into a solution of a salt of mercury succes-
sively place bits of different metals, such as copper, iron, and
zinc. Also try a nickel coin or a two-cent piece. What ones
are amalgamated? Is the coating permanent?
SUG. Mercurous nitrate, Hg2(NO3)2, is a good salt to use for this pur-
pose.
GENERAL CAUTION. Do not bring mercury in contact with valuable
articles, consisting of such metals as gold, silver, etc. Why ?
232. Classification of the Metals. — Various methods
of classification have been proposed and followed, such as
a classification according to those properties which are
THE METALS. 213
made use of in the analysis of substances. Prominent
among these properties are the solubilities of the metallic
chlorides, sulphides, hydroxides, carbonates, and phosphates
in various reagents. This method of classification is
well adapted to the analytical separation and recognition
of the various metals, while it interferes in no way
with their proper consideration in other respects. To
effect this separation in practice various group reagents are
employed, and five groups are obtained : —
A. THE FIRST GROUP METALS.
In the separation of this group hydrochloric acid is the
reagent employed, and all the metals belonging to the
group may be precipitated as chlorides. We mean by this
that any soluble salt containing a first group metal as a
base gives, upon the addition of hydrochloric acid to
a solution of that salt, an insoluble chloride. This group
contains three metals : —
Lead ..... Pb,
Silver .... Ag-,
Mercury . . . Hg ( in mercurous salts only).
NOTE. Mercury, as we have already mentioned, gives two series of
salts, which will be described under Mercury. The mercurous salts alone
are precipitated by hydrochloric acid. The mercuric salts belong to the
second group. It is necessary to state here that lead is not completely
precipitated by hydrochloric acid.
B. THE SECOND GKOTJP METALS.
The metals of this group are characterized by yielding
with hydrogen sulphide, H2S, metallic sulphides which are
insoluble in dilute acids. It is customary in analytical
operations first to acidulate the solution with hydrochloric
acid, and then to pass the hydrogen sulphide through the
214 THE METALS.
solution in question. If any or all the metals of this group
are present, the precipitate obtained consists entirely of
the sulphides of those metals. This group embraces the
common metals : —
Arsenic As,
Antimony .... Sb,
Tin Sn,
Bismuth .... Bi,
Copper Cu,
Cadmium .... Cd,
Mercury .... Hg" (in mercuric salts) ;
and the rarer metals : —
Gold . . . . Au
Platinum . . Pt
Palladium . Pd
Ruthenium . Ru
Iridium . . . Ir
Rhodium . Rh
Osmium . . Os
Tungsten . . W
Molybdenum . Mo
NOTE. The sulphides of the first group are also insoluble in dilute
acids an'd might be obtained in this group ; but in the course of analysis it
is best first to remove with hydrochloric acid the first group metals.
C. THE THIRD GROUP METALS.
The metals of this group are those whose hydroxides
and sulphides are soluble in dilute acids but insoluble in
alkaline solutions. Ammonia and ammonium sulphide,
(NH4)2S, are the third group precipitants, and it is cus-
tomary first to add to the solution under consideration
ammonium chloride before adding the group reagents.
This group includes the common metals : —
Iron Fe,
Chromium ... Cr,
Aluminum . . . Al,
Nickel Ni,
Cobalt Co,
Manganese . . . Mn,
Zinc Zn;
THE METALS. 215
and the rarer metals : —
Beryllium .
. Be
Cerium . .
. Ce
Titanium .
. Ti
Indium . .
. In
Didymium .
. D
Zirconium .
. Zr
Gallium
. Ga
Terbium
. Tb
Uranium .
. Ur
Yttrium
. Yt
Erbium . .
. E
Tantallum .
. Ta
Lanthanum
. La
Thorium .
. Th
Niobium .
. Nb
Vanadium .
V
NOTE. Many metals of this group form no sulphides in the wet way.
NiS and CoS are very sparingly soluble in cold dilute HC1.
D. THE FOURTH GROUP METALS.
We cannot isolate this group by means of their sul-
phides, etc., since these salts are soluble in acid and alkaline
solutions. The metals of this group are separated by
means of their carbonates which are thrown down by ammo-
nium carboTiate, (NH4)2CO3, in solutions made alkaline
with ammonia. The metals belonging to this group
are: —
Barium Ba,
Strontium . . . Sr,
Calcium .... Ca,
Magnesium . . . Mg.
NOTE. The carbonate of magnesium is somewhat soluble in ammonia,
and completely so in the presence of ammonium chloride; hence in practice
it is customary first to add ammonia, ammonium chloride, and then am-
monium carbonate. The magnesium salts are thus retained in solution and
afterwards precipitated as a phosphate.
E. THE FIFTH GROUP METALS.
These rnetals give no precipitates with common reagents,
since their salts are all soluble. This group includes : —
Potassium . . . K,
Sodium Na,
Ammonium . . . NH4
(known only in salts; see Ammonia),
Lithium .... Li;
216 THE METALS.
and the rarer metals : —
Rubidium Rb,
Caesium Cs.
In A, B, C, D, and E are outlined the general principles
which, with a few details to be explained further on,
enable us to separate the metals into groups. These
groups may again be taken up and each metal separated
and identified. It is thus that we may analyze a solu-
tion containing any or all the metals. It might be well
here to give a definition of " analysis " as applied in
chemistry.
1. Qualitative Analysis is the separation and detection
of the individual substances in a given compound.
2. Quantitative Analysis is the determination of the
weight or amount of each substance present in a given
compound.
QUERY. Which analysis must be made first ? Why ? .
233. Salts of the Metals. — The metals may be said to
react with all the acids previously mentioned to form salts.
The relations existing between the acids arid the salts are,
as a rule, simple and easily understood. A metal replaces
a certain number of hydrogen atoms depending on its
valence, a univalent metal replacing one hydrogen atom,
a bivalent metal, two, etc.
The simplest salts are those which are derived from
monobasic acids and univalent metals ; as,
Potassium nitrate . . . KNO3,
Sodium nitrite .... NaNO2,
Potassium chlorate . . KC1O3,
Lithium perchlorate . . LiClO4, etc.,
in each of which one atom of the metal replaces one atom
THE METALS. 217
of hydrogen, forming a normal salt, or one that contains
no more replaceable hydrogen.
In the case of bivalent metals and monobasic acids the
relations are also simple enough, one atom of the metal
replacing two atoms of hydrogen in two molecules of the
acid; as, for example: —
Calcium hypochlorite . . Ca(ClO)2,
Barium nitrate . . . . Ba (NO3)2,
Copper nitrate .... Cu(NO3)2,
Magnesium chlorate . . Mg (C1O3)2, etc.
A monobasic acid generally yields but one salt with any
given metal. A few curious exceptions to this rule will
be mentioned further on.
Taking now a bibasic acid, its two hydrogen atoms may
be replaced, (1) By two univalent atoms of the same
kind, as in
Potassium sulphate . . K2SO4,
Sodium carbonate . • . Na,CO3, etc.
(2) By two univalent atoms of different kinds as in
Sodium potassium carbonate, NaKCO3,
Sodium ammonium sulphate, Na(NH4)SO4, etc.
(3) By one bivalent metal, as in
Barium sulphate . . . BaSO4,
Zinc carbonate .... ZnCO3,
Copper sulphate . . . CuS04, etc.,
or (4) Only one of the hydrogen atoms may be replaced,
thus giving rise to the formation of a substance which is
called an acid salt, as in
Mono-potassium carbonate, KHCO3,
Mono-sodium sulphate . NaHSO4, etc.
218 THE METALS.
The matter becomes more complicated when we have
tribasic and tetrabasic acids, and trivalent and quadri-
valent metals to deal with; but still the student should
carefully trace the relation between the most complex
acids and their salts. Most acids are either monobasic
or bibasic, and only a few of those which we commonly
have to deal with are tribasic.
STJG. Let the student classify according to their basicity all the acids
thus far considered.
We shall learn that most metals which we commonly
have to deal with are either univalent, bivalent, or
trivalent.
A normal salt frequently unites with a hydroxide to
form a basic salt, e.g., —
Pb(N03)2 + Pb(OH), = 2 Pb
Again, water may thus act on a normal salt, e.g., —
Bi(N03)3 + 2 H20 = Bi (O)* + 2 HNO3.
Basic mercuric sulphate may be supposed to originate
thus : —
HgS04 + 2 HgO = Hg3S06.
The structures of some basic salts are exceedingly com-
plex.
EXERCISE. Taking as examples of univalent metals, potassium, sodium,
and ammonium (NHJ ; of bivalent metals, calcium, barium, and stron-
tium ; and of trivalent metals, aluminium and chromium, let the student
write the formulae of the following named salts: barium hypochlorite,
calcium nitrate, mono-potassium phosphate, tri-silver phosphate, tri-calcium
phosphate, aluminium meta-phosphate, barium iodate, chromium sulphate,
potassium aluminium sulphate, magnesium ammonium phosphate.
SUG. The teacher should add to this list, practicing the student until
it is evident that the principles involved are thoroughly understood.
A NATURAL CLASSIFICATION OF THE ELEMENTS. 219
A NATURAL CLASSIFICATION OF THE ELEMENTS.
234. As previously explained, the elements may be
roughly divided into metals and non-metals. There are
other characteristics affording methods of classification,
such as valence, in which the elements may be classed as
monads, diads, triads, etc. Again, as we have seen, there
are elements bearing a close resemblance to one another in
their chemical compounds, properties, etc., such as chlorine,
bromine, iodine, and fluorine ; or sulphur, selenium, tellu-
rium, etc. But a true understanding of natural relation-
ships requires a careful study of all the available properties
of the elements and their compounds, and cannot be based
upon any one characteristic alone.
That property of the elements which can be expressed
with the greatest certainty and definiteness is the atomic
weight. The specific gravity, although varying within
certain limits, may assist us to compare those elements
which are solid at ordinary temperatures. Again, we may
use the atomic volume, which is found by dividing the
atomic weight of an element by its specific gravity ; the
number thus obtained shows how many cubic centimetres
of an element are required to weigh as many grams as
there are units in the atomic weight of that element.
The following table presents a number of facts in regard
to the best known and most distinctly characterized ele-
ments including all those whose atomic weights are less
than 88. The lists of compounds are made as full as the
limits of the table allow, no facts being suppressed in the
interests of any theory. The student of nature will feel
best satisfied with that arrangement or classification which
most fully expresses the natural harmonies. Many dis-
crepancies are still to be expected through our lack of
knowledge, or our imperfect appreciation of chemical facts.
220 A NATURAL CLASSIFICATION OF THE ELEMENTS.
Atomic
Weight.
Atomic
Difference.
'3 °>
gU
Atomic
Volume.1
Highest
Hydrogen
Compound.
ill
BOO
Lowest
Oxygen
Compound.
REMARKS.
H
1.0
jj
HA
H20
Li
7.0
6.0
0.59
11.9
Li202(?)
Li20
Be
9.1
2.1
2.07
4.4
BeO
BeO
There is some reason
B
11.0
1.9
2.5
4.1
BA
BA
for assuming for Be
C
12.0
1.0
3.5
3.4
CH4
do,
CO
the at. wt. of 13.65
N'
14.0
2.0
Gas.
NH3
N205
N.O
with the oxideBe2O3.
0
16.0
2.0
Gas.
OH2
0,
02
Compare Ozone,p.31.
F
19.1
3.1
i
FH
Na
23.0
3.9
0.97
23.7
Na4H2
Na2O2
Na20
Mg
23.9
0.9
1.74
13.8
MgO
MgO
Al
Si
27.3
28.0
6.4
0.7
2.60
2.39
10.6
10.7
SiH4
A1203
Si02
AJA
Si02
P
31.0
3.0
2.20
12.8
PH3
P205
P203
P20 is suspecte'd to
s
32.0
1.0
2.05
16.
SH2
S03
so.
exist.
Cl
35.4
3.4
Gas.
26.
C1H
CIO,
C120
Perchloric acid,
K
Ca
39.0
39.9
3.6
0.9
0.87
1.58
45.4
25.3
K4H2(?)
K202
Ca02
K20
CaO
HC104, suggests a
hypothetical anhy-
Sc
44.0
4.1
1
?
dride, C1207.
Ti
48:0
4.0
4.1
12.5?
Ti02
Ti 0 C1}
V
51.2
3.2
5.5
9.2
V2o
Cr
52.4
1.2
6.8
7.65
CrOj
CrO
Mn
54.8
2.4
7.14
7.6
Mn207
MnO
Fe
55.9
1.1
7.86
7.09
FeH2
FeA
FeO
Ni
58.0
2.1
8.90
6.31
MO
Co
59.0
1.0
8.5
6.82
CoA
CoO
Cu
63.3
4.3
8.9
7.13
Cu2H2
CuO2
Cu40
Zn
64.9
1.7
7.2
9.37
ZnO
ZnO
Ga
68.0
3.1
5.9
11.5
Ga203
As
74.6
5.6
13.1
AsH3
AsA
As20
[Se03.
Se
79.0
4.4
4.8
16.5
SeH2
Se02
Se02(?)
H2Se04 suggests
Br
79.7
0.7
3.19
25.
BrH
Br2O, Br2O5, and
Rb
85.2
5.5
1.52
56.3
Rb20
Rb20
Br207 are hypothet-
Sr
87.2
2.0
2.5
34.4
SrO2
SrO
ical anhydrides.
I g
107.7
126.5
10.6
4.95
10.2
25.6
IH
Aigo°2
Ag40
JA
HI04 suggests the
Te
128.0
1.5
6.25
20.4
TeH2
Te053
Te02
hypothetical anhy-
dride I207.
The following table is based on that of a Russian chemist
named Mendelejeff, but modified in view of suggestions
from L. Meyer, Huth, and Muir : —
1 Prom Huth's Das periodische Gesetz der Atomgeiviclite, Frankfurt a. Oder, 1884.
A NATURAL CLASSIFICATION OF THE ELEMENTS. 221
I.
II.
III.
IV.
V.
VI. VII. VIII.
R20
RO
R,O3
RO2
RA
RO;j R2O7
HI
RH4
RH3
RH2 RH
1
y
Be
3
9
11
12
N)
® W
N,a
23
Mg
Al
14
16 19
K
39
24
Ca
40
27
Sc
44
28
Ti
AQ
P
31
-v
1 ci
35.5
Cr
i
Cu
63
Zn
Ga
4o
51
52 "y.n Fe, Ni, Co
56 58.6 59
rs
O)
n
'B
Rb
85
Ag
108
65
Sr
87
Cd
112
69
Y
89
In
114
72(?)
Zr
90
Sn
118
As
75
- Nb
94
'Sb
Se Br
79 M
Mo
% — Rh, Ru, Pd
('} 104 104.5 106
Te 1
NOTE. Spread gum on
Cs
133
Ba
137
La
139
Ce .
120
r Di
128 127
141
144
-
166
... .:
1A7C ?\
169 (?)
170 (?)
IRCQ
.Yb
173
f Ta
w
Au
176
Hg
200
Tl
204
Pb
207
182
•Bi
184 — |r, Os, Pt
1 ' 192.5 193 194
209
_
Th
_
!
232 (?)
237(?)
OytfA
222 A NATURAL CLASSIFICATION OF THE ELEMENTS.
In this table the elements are arranged in eight vertical columns, rep-
resenting eight groups ; while successive series are presented in nearly
horizontal lines. These are made to incline slightly, so that on rolling the
table Na will immediately succeed F, K will succeed Cl, and so on in a spiral
line. The first eight or twelve elements present very marked individuality
of character; some of these are typical of natural groups which follow.
Elements of most distinct basic character are found towards the left;
non-metals predominate in the upper and middle parts of Groups V., VI.,
and VII. ; while the lower part of the table is marked by the more indif-
ferent elements. A double spiral will be traced beyond Si (beginning with
P and V respectively) and distinguished by heavy-face and light-face type.
Many familiar relationships can now be traced out ; thus, K, lib, and Cs
are more closely related to each other than they are to Li and Na ; Ca, Sr,
and Ba are very closely related in their properties, while Mg resembles these
elements in some respects and Zn and Cd in others. Very many facts in
regard to the properties and compounds of the several elements may be
fixed in the mind by the law of association when studied with the aid of
this table, while they could only be retained by a severe effort of memory,
if viewed independently.
It wiU be noticed that the first series has but one member ; group VIII.
is represented in the even series only, beginning with the fourth ; and the
element of highest atomic weight yet discovered is in the twelfth series,
group VI. It is necessary to transpose I and Te in the table, in view of
their properties. The blanks represent the probable position and approxi-
mate atomic weights of elements not yet discovered or investigated.
When Mendelejeff published his table (in 1869) lie left two blanks which
have since been filled by Sc and Ga ; and the properties of these elements
agree very closely with those expressly predicted from the analogies indi-
cated in the table. The true position of some of the rarer metals (espe-
cially those of the cerium group) is still uncertain; these are here
arranged as in Muir's Principles of Chemistry.
The harmony of nature here exhibited is most impressive. Is it possi-
ble that the so-called elements are really compounds ? Did the various
" elements " of the earth and sun once exist as hydrogen, when our solar
system was a nebula ? l And will modern chemists ever revive the famed
problem of the alchemists, and seek to turn the base metals into gold ?
Far more precious than gold is the search for truth ; and the more we
learn of science, the broader becomes our conception of what we know in
part, and the deeper should be our reverence for the infinite thought of
the Creator.
1 See a paper by F. W. Clarke in Popular Science Monthly for Feb. 1876, p. 463.
A NATURAL CLASSIFICATION OF THE ELEMENTS. 223
QUERIES.
1. In column 2, p. 220, how many numbers differ from whole numbers
by less than 0.1 ? If the numbers are calculated by comparing with O = 16,
they approximate still more closely to whole numbers.
2. In column 3, what differences are greater than 3 "? Which are less
than 2 ?
3. In column 4, where do the numbers increase ? Where do they
diminish ? Note the same in column 5.
4. What monads are indicated by the compounds of column 6 ? What
diads ? Triads ? Tetrads ?
5. Note the valence indicated by the oxides in the table, or by any
other compounds that you may know.
6. Imagine it possible to begin with an atom of hydrogen, and to build
up an atom of each of the elements by successive additions of matter; can
you show that the specific gravity of the product would alternately increase
and diminish ? What kind of variation is observed in the atomic volumes 1
In valence ?
7. At what points of the series are elements of strongly marked non-
metallic (or electro-negative) character brought into juxtaposition with
those of strongly marked metallic or electro-positive character ? Is this
transition marked by a relatively large or small increase of weight ?
CHAPTER XV.
THE FIRST GROUP METALS.
The metals of this group are, as previously explained,
Lead, Silver, and Mercury. They are of great importance,
and are utilized in manifold ways.
LEAD.
SYMBOL, PB". — ATOMIC WEIGHT, 207. — SPECIFIC HEAT, 0.0315.
— MELTING-POINT, 334°.
235. Occurrence. — Native, or free, metallic lead occurs
in very small quantities in certain lead-bearing ores and in
volcanic tufa. The principal source of lead is its sulphide,
Galena, PbS. This ore is distributed throughout nearly
every geological period, but the largest deposits in the
United States are in the Lower Silurian. Nearly every
ore of lead is argentiferous, i.e., silver-bearing ; and it is
not uncommon to find lead associated with other metals,
as copper, tin, zinc, arsenic, antimony, molybdenum, tung-
sten, etc.
236. Preparation of Metallic Lead. — EXP. 136 P. In
a test-tube containing a solution of lead acetate, Pb(C2H3O2)2,
place a clean strip of metallic zinc. A dark deposit of lead
soon forms on the zinc. Complete the equation
Pb(CsH8Og)2'+Zn= . . .
Collect this deposit of lead, place it on charcoal, and cover
LEAD. 225
with sodium carbonate ; now heat the mass before the blow-
pipe (reducing-flame) , when a bead of metallic lead is easily
obtained. Place this bead on an anvil and strike it a light
blow with a hammer. Is it easily malleable ? Cut the flattened
bead with a knife, and scratch it with the finger-nail, carefully
noting the hardness, lustre, tarnish, etc.
EXP. 137 P. Take any lead compound or lead ore, such as
red lead, Pb3O4, or galena, PbS ; place it on charcoal, and heat
it before the reducing-flame. Do you again obtain a bead?
Try this bead as before.
EXP. 138 P. Solder to each terminal of a Grove, Grenet, or
Bunsen battery (2 cells) a narrow ribbon of platinum foil.
Place the platinum strips about lcm apart in a beaker-glass
containing a strong solution of lead acetate. Treat the deposit
obtained as before. Is this deposit lead?
QUERY. What is Electrolysis 1
Lead is easily reduced from its ores, in consequence of
which it has been known since the highest antiquity. In
the smelting works, where it is prepared from its ores,
three distinct processes are employed : —
1. The air reduction process. In this process the ore
employed is the sulphide, PbS, which is simply roasted in
a reverberatory furnace until one portion of the sulphide is
changed to the sulphate, PbSO4, and another to the oxide,
PbO ; the heat is then increased, when the unaltered sul-
phide reacts with the oxide and sulphate thus : —
(a) 2 PbO + PbS = 3 Pb + SO2.
(b) PbSO4 + PbS = 2 Pb + 2 SO2.
2. The carbon reduction process. In this process the
sulphide is mixed with peat, or other carbonaceous mate-
rial, and reduced in a blast-furnace. The first and second
processes are adapted to very pure ores only.
226 LEAD.
QUEKY. In what experiment did you obtain lead by a combination of
these two processes ?
3. The precipitation process is adapted to the prepara-
tion of lead from impure ores. In this process the ore is
melted with cast iron or iron slag. A portion of the lead
is obtained pure, while the remainder, contaminated with
other metals, is afterwards so treated that all are saved.
QUERIES. In how many different ways can you prepare metallic lead ?
Why is not the zinc process an economical one 1
237. Properties, Uses, and Compounds of Lead. —
Lead is a soft, heavy, malleable metal possessing a high
lustre, which is best seen on a freshly-cut surface. On
exposure to the air this surface soon oxidizes, thus: —
4Pb + O2 = 2Pb2O.
The oxide Pb2O is a bluish-gray substance which soon
forms a coating over the exposed surface, and prevents
further oxidation.
EXP. 139 p. Draw the dull surface of a bit of lead over
a clean white paper. Note that the surface of the lead becomes
bright, and that a black streak is made on the paper.
Lead is insoluble in pure cold water free from air, but
water is seldom or never pure ; hence water flowing
through, or standing in, leaden pipes or vessels is almost
certain to contain lead salts in solution. Now, since lead
and its salts affect the system as a virulent, cumulative
poison, such waters should never be used for drinking or
cooking purposes.
SUG. Explain the action of a cumulative poison.
EXERCISE. Let the student name all the uses of metallic lead that he~
can call to mind. Also let him name the alloys of lead and their uses.
A good solvent for lead is dilute nitric acid, since the
nitrate is a very soluble lead salt. The nitrate and
LEAD. 227
•*
acetate solutions in water are the best ones to use as work-
ing solutions.
THE PRINCIPAL COMPOUNDS OF LEAD ARE : —
(a) Lead Oxide or Massicot, PbO, a yellow powder. Lith-
arge is an impure form of lead oxide, containing oxides of
other metals, as of copper, iron, etc. Both forms are obtained
by heating lead in the air. Litharge is used in glazing earthen-
ware, in preparing flint glass, and in the manufacture of Red
Lead or Minium, Pb3O4. This last is much used as a pigment,
and in steam-pipe fitting.
(b) Lead Acetate, or Sugar of Lead, Pb(C2H3O2)2, is used
in medicine ; in the laboratory it is a valuable reagent ; in the
arts it is used with potassium bichromate, K2Cr2O7, for dyeing.
EXP. 140 P. Moisten a strip of white cotton cloth in a solu-
tion of lead acetate, and then moisten it in a solution of potas-
sium bichromate. What color is the strip dyed ?
(c) White Lead, 2 to 3 PbCO3+Pb(OH)2, the principal and
best white paint known. It is prepared ("Dutch method ") by
placing rolls of sheet lead in earthenware vessels containing
vinegar or crude acetic acid. These vessels are then piled in
tiers, layers of manure or spent tan-bark being placed between.
The whole is then covered with manure, which, by its decompo-
sition, furnishes sufficient heat to cause the sheet lead and acetic
acid to react and to form the compound Pb(C2H3O2)2.2PbO.
This compound is next decomposed by the carbon-dioxide
which escapes from the fermenting mass around the vessels.
In four or five weeks the process is completed.
In the " French method " white lead is prepared bypassing
carbon-dioxide through an aqueous solution of litharge in lead
acetate.
(d) Galena, PbS, is a dark, shining solid, crystallizing in
cubes and in other forms belonging to the regular system. It
is the principal ore of lead.
228 SILVEK.
(e) Lead Chromate, or Chrome Yellow, PbCrO4, is a pig-
ment obtained by treating a soluble lead salt with potassium
bichromate, K2Cr2O7.
(/) Lead Chloride, PbCl2, is important as a precipitate
met with in the regular course of analysis. It is a crystalline,
white solid, soluble in hot water.
EXERCISE. Precipitate a dilute solution of lead acetate or nitrate with
K2Cr2Or Note the color of the precipitate, and try its solubility in HNO3
and ammonia. Thus try other precipitants, as HC1, H2SO4, (NHJ2S,
KOH, and KI, and test the solubility of the precipitates as before.
Tabulate the results and keep them for future reference.
238. Tests for Lead. — 1. Metallic lead is recognized by
its lustre, tarnish, streak, and malleability. (See Exp. 136.)
REM. If the student is not sure, he may dissolve a bit of the metal in
dilute nitric acid, and test by 2.
2. Lead, in a solution of its salts, is detected by the
colors of its precipitates : —
(a) H2S gives PbS (black),
(b) K,Cr2O7 gives PbCrO4 (yellow) ,
(c) (NH4)2CO3 gives PbCO3, Pb(OH)2 (white),
(d) KI gives PbI2 (yellow scales) ,
(e) H2SO4 gives PbSO4 (white).
3. Lead in an unknown solid is detected by reduction
on charcoal with sodium carbonate, and, if farther identi-
fication be necessary, by 2, after dissolving the bead in
nitric acid.
SILVER.
SYMBOL, AG'. — ATOMIC WEIGHT, 108. — SPECIFIC HEAT, 0.0570.
— MELTING-POINT, 1000°.
239. Occurrence. — Silver occurs native in considerable
quantities with native copper deposits, but its chief sources
SILVER. 229
are from the lead furnaces, mentioned in Art. 236, and
from the following ores : Argentite, Ag2S ; Ruby Silver,
Ag3SbS3; Silver-Copper Glance, Ag2Cu2S2; Horn Silver,
AgCl; and other compounds containing silver, copper,
antimony, arsenic, and sulphur in varying proportions.
240. Preparation of Metallic Silver. — EXP. 141 p. In
a test-tube containing a solution of silver nitrate place a strip
of zinc. Fuse the dark-colored deposit on charcoal before the
blow-pipe. Try the bead as you tried lead. What differences
do you find ? Thus reduce silver by means of strips of copper
and Iron. "Write the equations and explain the reactions.
Will mercury thus yield metallic silver ?
EXP. 142 p. Try to prepare silver by the electrolysis of
silver nitrate (see Exp. 138). Fuse to a bead on charcoal
the substance obtained. Does silver tarnish like lead ?
EXP. 143 P. Heat any silver salt, as AgCl, on charcoal be-
fore the reducing-flame. Try to oxidize the bead thus obtained,
by using the oxidizing flame. Can you thus oxidize lead? Can
you separate silver from lead by the blow-pipe ?
EXP. 144 p. To a solution of silver nitrate add tartaric acid,
H2C4H4O6, and heat. Note the silver mirror deposited on the
sides of the test-tube. Is this an instance of oxidation or
reduction ? What then occurs to the acid ?
EXP. 145 p. To a solution of silver nitrate add a little of a
solution of chloral hydrate, C2HC13O, H2O. Make the mixture
faintly alkaline with ammonia, and heat. Is a silver mirror
again formed on the sides of the test-tube ?
QUERIES. In how many ways have you prepared silver ? Would the
process of reducing silver with zinc be an economical one, provided there
were no better processes 1 Heat the crystals of the Arbor Dianae (Exp.
134) in an iron spoon. Can you regain the silver, thus separating silver
from mercury ?
Silver is reduced from its ores in three different ways.
280 SILVER.
(1) by Cupellation, or Oxidation of lead; (2) by Amal-
gamation; (3) by Solution and Precipitation.
1. In the lead furnaces metallic silver is obtained to-
gether with the lead. ^ This alloy of lead and silver, when
molten, is allowed to cool slowty; when the temperature
reaches a certain point, most of the lead separates out in
crystals, which are removed by means of perforated dip-
pers. In this way an alloy of lead rich in silver is
obtained.
Again, in certain localities, where some of the purer
ores of silver occur, the crude ore is melted with pure
lead ; thus a similar alloy is obtained.
The alloy obtained in either case is freed from lead by
cupellation, i.e., it is strongly heated in bone-ash vessels,
called cupels, over which a current of air is flowing. At a
high temperature the lead is oxidized, while the silver is
not changed; upon completion of the process, metallic
silver remains in the bottom of the cupel.
QUERY. In what experiment were principles of this process employed 1
2. In the amalgamation process the ore is ground fine
(sometimes first roasted) and mixed with sodium chloride
and mercury; copper sulphate is also frequently added.
In this way a silver amalgam is obtained. The silver is sep-
arated by distilling off the mercury in iron retorts. The
mercury is condensed in cool receivers and again employed
for the same purpose.
QUERY. What experiment foreshadows this process ?
3. Iii the third process the silver ore is first roasted.
If it contains the sulphides of iron and copper, which is fre-
quently the case, the silver is oxidized to the sulphate,
Ag2SO4, which, by means of water, may be dissolved out
from the insoluble oxides of copper and iron formed by
SILVER. 231
the roasting. From the solution of silver sulphate thus ob-
tained metallic silver is precipitated by introducing metallic
iron.
If the ore does not contain the sulphides of iron and
copper, sodium chloride is mixed with the ore before
roasting. Silver chloride is obtained, which is then dis-
solved in sodium thiosulphate. Silver sulphide is next
precipitated by adding to the thiosulphate solution sodium
sulphide. Finally the sulphur is driven off by heating the
silver sulphide in a muffle furnace.
QUERY. What experiment illustrates these principles ?
EXP. 146 P. To a solution of silver nitrate add ferrous sul-
phate. Do you obtain a precipitate ? Is it silver?
. QUERIES. Is silver easily reduced from its salts ? What substances
have been mentioned which are capable of thus reducing silver ? Why
does silver nitrate blacken the skin or other organic materials 1 Will a
solution of sugar reduce silver ?
241. Properties, Uses, and Compounds of Silver. —
Silver is one of the precious metals, and has been known
and valued since the highest antiquity. It is white, bril-
liant, and very ductile and malleable. It does not oxidize
in the air at any temperature, hence its use in coinage and
jewelry.
Silver is readily attacked by ozone, chlorine, bromine,
iodine, phosphorus, sulphur, and sulphuretted hydrogen.
QUERIES. Why does silver coin blacken when carried in the pocket
with matches 1 Why do egg and mustard spoons blacken ? Why drink-
ing cups used with sulphur waters ? What substance is a good solvent
for silver sulphide 1 Explain its action on blackened silverware. What
gases from soft-coal grates and from burning illuminating gas blacken
silver ?
EXERCISE. Name all the uses for which silver is employed.
The best solvent for silver is nitric acid, and silver nitrate
is the best salt to use in working solutions.
232 SILVER.
THE PRINCIPAL COMPOUNDS OF SILVER ARE:
(a) Silver Nitrate, AgNO3, or Lunar Caustic. This salt is
prepared by dissolving silver in nitric acid. It is extensively
used in medicine as an escharotic agent ; in photography ; in
the laboratory as a reagent ; as an indelible ink, etc. Sticks
of lunar caustic are prepared by fusing the ordinary crystals of
silver nitrate and casting the fused mass in moulds.
EXP. 147 P. Moisten a sheet of paper with silver nitrate;
dry the paper in the dark ; lay upon the prepared paper a fern
leaf, a skeleton leaf, or a bit of lace. Cover with a sheet of
glass. Expose the whole to sunlight until the sheet is black-
ened. Now in the dark treat the paper with a solution of
sodium thiosulphate, and then wash perfectly clean with pure
water. Explain the formation of the " print" obtained. Thus
prepare a print from a lantern- slide or a photographer's nega-
tive. •
(b) Silver- Plating /Solution may be obtained by dissolving
silver chloride in an excess of potassium cyanide. The prepa-
ration of this substance is shown by the two equations : —
1. AgNO3 + NaCl = AgCl + NaNO3.
2. AgCl + 2 KCy = AgCyKCy + KC1.
This is used as an electro silver-plating solution. See Gore's
Electro-Metallurgy for full directions for electro-plating. Also
read R. and S., Vol. II., Pt. I., pp. 361-65.
NOTE. The AgCl is freed from the NaNO., by filtration and washing.
The silver potassium cyanide solution with the potassium chloride may be*"
used as a silver electro-plating solution.
(c) Silver Chloride, AgCl, is important, in that it is a group
precipitate, obtained by adding hydrochloric acid to a silver salt
solution. This precipitate is soluble in ammonia : —
2 AgCl + 3 NH3 = (NH3)3(AgCl)2 ( ?) .
MERCURY. 233
(d) /Silver Bromide, AgBr, is used in photography, and may
be thus prepared : —
AgNO3 + KBr = AgBr + KNO3.
In an impure form silver bromide occurs native.
EXERCISE. Precipitate a silver nitrate solution with HC1, and test the
solubility of the precipitate with NH3, KCy, Na2S203, and HN03. Thus
obtain and try the precipitates with KBr, KI, and KOH or NaOH. Tabu-
late the results (with colors of precipitates) and preserve them for
future reference.
242. Tests for Silver. — 1. Metallic silver is recognized
by its lustre and other physical properties.
2. If the student is not sure, he may dissolve a bit of
the metal in HNO3 and add HC1. A white precipitate,
insoluble in HNO3, and soluble in ammonia, indicates
silver.
3. Unknown solids may be tested on charcoal with the
blow-pipe. The bead may be examined as in 1 and 2.
4. Unknown solutions are tested by adding : —
(a) HC1, etc., as in 2 ;
(6) FeSO4, as in Exp. 146 ;
(c) H2S gas, which gives a black precipitate, Ag2S, solu-
ble in KCy and strong HNO3.
5. Silver may be separated from lead by using the oxid-
izing flame of the blow-pipe as in Exp. 143.
MERCURY.
SYMBOL, Hg'1". — ATOMIC WEIGHT, 200. — SPECIFIC HEAT, 0.0319.
— MELTING-POINT, —40°. — BOILING-POINT, 357.25°.
:
3243. Occurrence. — Metallic mercury occurs only in
ry minute globules disseminated through its chief ore,
234 MERCUKY.
Cinnabar, HgS. Cinnabar occurs in Mexico, California,
Spain, Bavaria, China, Japan, and other countries.
244. Preparation. — EXP. 148 p. In a hard glass tube,
open at both ends, place a small quantity of vermillion or cinna-
bar, HgS. Hold the tube somewhat slanting in the Bunsen
flame, and heat strongly. Sulphur dioxide fumes escape from
the upper end of the tube, while mercury is deposited in the
tube in the form of a mirror. Write the equation.
EXP. 149 P. Heat red oxide of mercury, HgO, in a test-
tube, and explain what takes place. In what connection have
you thus treated HgO ?
EXP. 150 P. In a solution of a mercury salt suspend a strip
of zinc. In what form do you thus obtain mercury? Thus
proceed with a piece of clean copper wire ; an iron wire. Do
you obtain mercury in both cases? Compare the precipitate
obtained by zinc in the mercury salt solution with those pre-
cipitates obtained in silver and lead salt solutions.
EXERCISE. Prepare a table showing the action of copper, zinc, and
iron upon the salts of the first group metals.
QUERY. Can you obtain metallic mercury from its salts by means of
reducing agents, such as sugar, chloral hydrate, FeS04, SnCl2, etc. ?
Compare by means of a table the results obtained with those obtained
with lead and silver salts.
The commercial preparation of mercury is a very simple
process. Cinnabar is simply heated in a furnace so con-
structed that a current of air is passed through the highly
heated ore. The sulphur is oxidized to sulphur dioxide ;
the mercury is vaporized, and afterwards condensed under
water in a cooling chamber.
QUERIES. Why was the tube in Exp. 148 open at both ends 1 In the
next Exp. why could one end of the tube be closed ? Which of these
experiments illustrates the process for manufacturing mercury ?
MERCURY. 235
245. Properties, Uses, and Compounds of Mercury. —
Metallic mercury is a silver-white liquid, vaporizing slowly
at all temperatures between its freezing-point and boiling-
point. Its properties were discovered and discussed by
the alchemists, and some of its compounds were found to
possess great medicinal properties.
Mercury acts as a poison upon the human system, especi-
ally when in the form of vapor.
Metallic mercury is used in constructing thermometers,
barometers, and other instruments used in physical meas-
urements. Its amalgams are of great value.
The best solvent for mercury is nitric acid. Solutions of
mercurous nitrate, Hga(NO3)2, and mercuric chloride are good
working solutions.
THE PRINCIPAL COMPOUNDS OF MERCURY ARE : —
(a) Cinnabar, HgS, an ore of mercury ; the artificial sul-
phide is used as a paint (vermillion) .
(b) Red Oxide of Mercury, HgO, also called Red Precipitate;
it is used in medicine. This compound is obtained by heating
a very intimate mixture of mercury and mercuric nitrate until
no red fumes are given off. It may also be obtained as an
orange-yellow powder by adding an excess of sodium or potas-
sium hydroxides to the solution of a mercuric salt.
(c) Mercurous Chloride, or Calomel, Hg2Cl2, is used in medi-
cine. This substance is prepared by subliming an intimate
mixture of mercuric chloride and mercury. It is also obtained
when an excess of hydrochloric acid is added to a solution of
mercurous nitrate. When thus obtained, it is a Group Pre-
cipitate which turns black with ammonia : —
Hg2Cl2 + 2 NH3 = NH2Hg.2Cl + NH4C1.
Mercurous chloride is soluble in nitro-hydrochloric acid.
(d) Mercuric Chloride, or Corrosive Sublimate, HgCl2, is a
236 MERCURY.
deadly poison ; it is used in medicine, and in the laboratory as a
reagent. This substance is prepared by subliming a mixture of
mercuric sulphate and common salt.
(e) Mercurous Nitrate, Hg2(NO3)2, is often sold fraudulently
as a silver-plating solution. This is to be had by treating an
excess of metallic mercury with cold, dilute nitric acid. If the
acid be in excess, mercuric nitrate, Hg(NO3)2, is obtained.
It will be seen that mercury forms two compounds with
chlorine and two with nitric acid. Were we to examine the
entire list of the salts of mercury, we should find numer-
ous other illustrations of this tendency on the part of the
metal to form two distinct series of derivatives, of which
the two chlorides and two nitrates mentioned are good
representatives. The simplest formulae which can be
assigned to the two chlorides are HgCl and HgCl2, and to
the twe nitrates, HgNO3 and Hg(NO3)2. It would appear
from these formulae that in the simpler compounds, HgCl
and HgNO3, mercury acts as a univalent element ; whereas,
in the more complicated compounds, HgCl2 and Hg(NO3)2,
it acts as a bivalent element. It appears more probable,
however, that the formulae of the simpler compounds
should be doubled, becoming Hg2Cl2 and Hg2(NO3)2; and
perhaps in these compounds the two mercury atoms are
Hg-
united with each other, as indicated thus, I , forming
a bivalent group. Hg—
However this may be, it is more common nowadays to
write the doubled formulae, and we thus have the two
series of mercury compounds corresponding to Hg2Cl2 and
HgCl2. The former is called mercurows chloride, and the
latter, mercunc chloride. The chloride containing the
smaller proportion of the acid constituent is designated by
the terminal syllable -ous, while that chloride which con-
MERCURY. 237
tains the larger proportion of the acid constituent is desig-
nated by the syllable -ic. The mercurous salts correspond
in composition to mercurous chloride. The mercuric salts
correspond to mercuric chloride.
Similar series of salts are known in connection with :
iron, whicli gives ferrous and ferric salts ; copper, which
gives cuprous and cupric salts; and many other metals.
The most marked cases are those of mercury, iron, and
copper.
Only the mercurous salts are precipitated in the first
group. The mercuric salts are thrown down in the second
group.
EXERCISE. With a solution of a mercurous salt try the precipitants
HC1, H2S, KI, and KOH. Try the solubility of each precipitate in HN03,
and in nitre-hydrochloric acid. Thus proceed with a mercuric salt; make
a table comparing the results.
246. Tests for Mercury and the Mercury Compounds.
-1. Metallic mercury is readily recognized by its physi-
cal properties.
2. An unknown solid is tested for mercury by heating
it in a test-tube with anhydrous sodium carbonate, Na2CO3.
A mirror of metallic mercury is formed on the sides of the
test-tube.
3. An unknown solution is tested for mercury by add-
ing : —
(#) HC1. If a white precipitate be formed, filter it out
and moisten it on the filter-paper with ammonia. If the
precipitate turns black, mercury in the mercurous condi-
tion is present.
(£>) Through the filtrate from (a) pass hydrogen sul-
phide. Mercury, in the mercuric condition, gives a black
precipitate, which is to be tested farther by dissolving it
238 MERCURY.
in nitre-hydrochloric acid, and evaporate to expel the excess
of acid; to this solution add stannous chloride, SnCi2: —
HgCl2 + SnCl2 = Hg + SnCl4.
The mercury thus obtained appears (usually after some
time) as a finely-divided, black precipitate.
4. A copper wire in a solution containing a mercury
salt is soon coated with a silver-white deposit.
5. A solid may be dissolved in nitro-hydrochloric acid,
and the solution directly tested by SnCl2, or by 4.
247. To separate and identify Lead, Silver, and Mer-
cury (Hg2). — 1. To a solution containing salts of these
three metals add HC1 ; the compounds PbCl2, AgCl, and
Hg2Cl2 are thus obtained together in the form of a pre-
cipitate. Filter and wash this precipitate with a little cold
water. •
NOTE. Lead is not completely precipitated by HC1 ; consequently some
lead usually passes over into the second group.
2. Add much hot water to the precipitate as it lies on
the filter-paper. The lead chloride, PbCl2, is thus dis-
solved, and will now run through the filter-paper. Collect
this solution in a beaker, and test for lead by Art. 238, 2.
3. The undissolved precipitate on the filter-paper now
consists of AgCl and Hg2Cl2. The silver chloride, AgCl,
may now be dissolved out by adding a little ammonia.
Collect the solution, (NH3)3(AgCl)2, as it runs through,
and test for silver by acidulating it with nitric acid ; AgCl
is again precipitated : —
(NH3)3(AgCl)2 + 3 HN03 = 2 AgCl + 3 NH4NO3.
NOTE. The formation, in this connection, of the white precipitate,
AgCl, upon adding nitric acid, is sufficient to identify silver ; but in case
the precipitate be plentiful, a bead of metallic silver may be had, as in
Art. 242, 3.
MERCUKY. 239
4. At the same time that the silver chloride dissolves
in ammonia, the mercurous chloride, Hg2Cl2, turns black
(245, 6*) and remains on the filter-paper. This blackening,
in this connection, is a sufficient indication that mercurous
compounds are present. A farther test, by 246, 5, may be
employed, however, if desirable.
248. To separate Mercury, Lead, and Silver by the
Blow-Pipe. — EXP. 151 P. Make an amalgam of silver, as
in Exp. 134. Also make an amalgam of lead by warming and
rubbing bits of lead and mercury in an evaporating-dish ; mix
the two amalgams ; you thus have a compound of metallic
mercury, lead, and silver.
1 . Carefully heat a bit of this compound on charcoal in the
r 'educing Aflame until you think you have driven off the mercury.
Dissolve the residue in dilute nitric acid, and test by Art. 247.
If the separation of the mercury was complete, you will obtain
tests for lead and silver only.
2. Heat another bit of the compound on charcoal in the
oxidizing flame. You may thus drive off the mercury and
oxidize the lead, leaving a bead of metallic silver. Dissolve
this bead in nitric acid, and try for silver, lead, and mercury by
Art. 247. If the separation of mercury and lead was complete,
you will obtain a test for silver only.
NOTE. In case the amalgam, when heated as in 1, spits out and is
lost, this step may be accomplished by carefully heating and shaking a
fresh portion in an iron spoon.
249. Reactions in Group I. — Balance these equations.
Iifhat principles do they illustrate ?
(1) Pb + HN03 = Pb(N03)2
(2)
(3) Hg1+HN08=Hg8(NOs), + NO
(dilute)
(4) Pb(NQ3)2 + HC1 = PbCl, + HNO3
240 MEKCUKY.
(5) AgN03 + HC1 = AgCl + HNO3.
(6) Hg2(N03)2 + HC1 == Hg2Cl2 + HN
(7) PbCl2 + H20 = Sol. of PbCl2.
(8) PbCl2 + K2Cr04 = PbCrQ4 + KC1.
(9) AgCl + NH8=(NH,)8(AgCl)a.
(10) (NH3)3(AgCl)2 + HN03 = AgCl
(11) Hg2Cl2 + NH3 = NH2Hg,Cl + NHAC1.
(12) NH2Hg2Cl + NOC12 + Cl + H2O = HgCl2 + N02.
(13) HgCl2 + SnCl2 = Hg + SnCl4.
SUG. The precipitates are underscored. Let the student determine
which of the substances on the right of the sign = are gases.
MODEL RECITATION. Equations 1, 4, 7, and 8 illustrate the reactions
previously described in lead. No. 1 shows how lead is dissolved; Pb(NO3)2
is the substance in solution; NO is a gas. No. 4 shows the precipitation
of lead with HC1. PbCl2 is the precipitate. No. 8 shows the distinctive
test for lead (PbCr04 being yellow) which was separated from silver and
mercury by No. 7. Give the name and the formula of each compound.
EXERCISES.
1. Observe that different products are obtained when some substances
react, depending upon which substance is in excess. Thus, if Hg be in
excess, silver amalgam is obtained in Exp. 134, while an excess of AgN03
gives pure silver. What other instances have been given ?
2. Read R. and S., Vol. II., Pt. I., pp. 388 to 392, for different processes
of preparing mercury.
3. Expose some freshly prepared AgCl to the action of sunlight. What
changes in color occur ?
4. Compute the atomic heats and atomic volumes of lead, silver, and
mercury.
NOTE. The expression "group precipitate," as used in the text in connec-
tion with a single metallic salt, signifies the precipitate of that metal
obtained by the group reagent.
CHAPTER XVI.
THE SECOND GROUP METALS.
250. The second group metals are those the sulphides
of which are insoluble in dilute acids. These metals are
separated from all others by removing the first group
metals with hydrochloric acid, after which hydrogen sul-
phide, H2S, is passed through the acidulated solution.
NOTE. Tellurium and Selenium, which are precipitated with these
metals, have already been described. Lead which has not been fully
removed from the first group also appears in this group.
The metals of this group exhibit many kindred proper-
ties. Their oxides, excepting those of arsenic, are nearly
insoluble in water ; they do not decompose water except
at high temperatures, and then but four, viz., bismuth,
antimony, tin, and molybdenum, give this reaction to any
considerable extent ; all the commonly occurring metals
form soluble nitrates (excepting antimony and tin, which
form oxides) when treated with nitric acid ; all are
readily reduced to the metallic state when heated on
charcoal in the reducing-flame.
In deference to some requirements in analysis, we may
divide the common metals of this group into two divisions.
The metals of the first division yield sulphides which are
soluble in yellow ammonium sulphide, (NH4)2S2, while the
sulphides of the second division are insoluble in that
reagent. Let us distinguish these divisions by the letters
A and B ; then each division is as follows : —
242 AESENIC.
f Mercury (in mercuric salts),
f Arsenic. I Lead.
DIVISION A 1 Antimony. DIVISION" B -{ Bismuth.
^ Tin. j Copper.
[ Cadmium.
NOTE. The yellow ammonium sulphide for the purposes mentioned in
this chapter may be prepared by gently warming in a test-tube a little
reagent ammonium sulphide, (NHJ2S, with a small quantity of flowers of
sulphur. The reagent sulphide, upon standing, also changes to the yellow
variety.
DIVISION A.
ARSENIC.
SYMBOL, Asiii)V. — ATOMIC WEIGHT, 75. — SPECIFIC HEAT,
0.0822. — MELTING-POINT, 356°.
251. Occurrence. — In nature arsenic occurs free in
kidney-shaped masses, which usually may be split up into
thin laminse or leaves ; but commercial arsenic is obtained
chiefly from some of the following ores : Iron Arsenide,
FeAs2; Nickel Arsenide, NiAs; Mispickel, (FeS)2As; Re-
algar, As2S2 ; Orpiment, As2S3; and from Arsenic Trioxide,
As2O3, combined with lead, calcium, and cobalt as arsenites.
252. Preparation. — EXP. 152 p. Make a pellet of arsenic
trioxide, As2O3 (commonly known as "arsenic"), with pow-
dered charcoal and a drop or two of water. Place the pellet in
the bottom of a hard glass test-tube, and heat gently to expel
the water. Now insert a loosely-fitting stopper (made of chalk)
nearly down to the pellet, which is then to be heated to redness.
Arsenic is freed and vaporized ; the vapors condense, above
the chalk, on the sides of the test-tube, forming a metallic
mirror.
EXP. 155 P. Heat any arsenic compound, as As2O3, on
charcoal before the reducing-flame. Arsenic is freed in form
of a vapor which has an odor somewhat resembling garlic.
ARSENIC. 243
Similarly treat a bit of metallic arsenic. Do you obtain the
same odor ? How do you now know that arsenic was freed by
heating As2O3?
Commercial arsenic is prepared by heating its ores,
especially mispickel and orpimeut, in earthen vessels or
tubes. The arsenic is driven off in vapors, which are
condensed in sheet iron tubes or condensers.
To purify the arsenic thus obtained it is sublimed with
charcoal, when it condenses in rhombohedral crystals pos-
sessing a bright metallic lustre.
253. Properties, Uses, and Compounds of Arsenic. —
Arsenic is a solid substance possessing a steel-gray color
and a metallic lustre. When heated under ordinary
pressure, it seems to vaporize without melting, at 356°;
under greater pressure, however, it may be obtained in a
liquid state.
As previously noted, arsenic stands midway between
the metals and non-metals ; in its chemical compounds
and chemical deportment it is closely allied to phosphorus
on the one hand, while, on the other hand, the physical
properties of arsenic and its compounds bear a close re-
semblance to those of antimony.
Arsenic oxidizes quite readily in warm, moist air, form-
ing a dark substance (probably a low oxide) known as
fly powder. When strongly heated in oxygen, arsenic
burns with a white light, forming arsenic trioxide, As2O3.
With oxygen and hydrogen, arsenic forms the acids
arsenious acid, H3AsO3(?), and arsenic acid, H3AsO4, which
closely resemble the corresponding acids of phosphorus.
The vapors of arsenic possess a strong odor resembling
garlic. Both arsenic and its soluble salts act as deadly
poisons when taken into the system, and even arsenical
244 ARSENIC.
vapors produce the symptoms of arsenic poisoning when
inhaled or absorbed through the pores of the skin.
The best antidote for arsenic is freshly prepared ferric
hydroxide, Fe2(OH)6, made by adding ammonia to a solu-
tion of ferric chloride, Fe2Cl6. The ferric hydroxide is
filtered out and washed, when it is ready for use. Mag-
nesia, MgO, is also an antidote ; both these substances
form insoluble compounds with the arsenic, thus prevent-
ing its absorption by the system. An emetic, such as a
teaspoonful of mustard in a cup of warm water, should
soon follow the antidote, and that in turn should be fol-
lowed by castor oil.
Arsenic is dissolved by nitro-JiydrocJiloric acid or by chlorine
water : —
2 As + 5C12 + 8H2O = 2H3AsO4+ 10HC1.
A good working solution of an arsenite can be made
thus : —
As2O3 + 6 NaOH = 2 Nu;,AsO3 + 3 H2O ( + an excess of NaOH) .
THE PRINCIPAL COMPOUNDS OF ARSENIC ARE: —
(a) Arsenic Trioxide, As2O3. This oxide is sometimes called
arsenious anhydride, and is usually sold in drug stores as
" arsenic." This a white crystalline powder, used for destroy-
ing vermin, as a medicine, and in taxidermy as a dryer and
antiseptic. Arsenious Acid, H3AsO3, has not been isolated ;
from it are derived the arsenites. All the soluble arsenites
are poisonous.
(6) Arseniuretted Hydrogen or Hydrogen Arsenide, AsHv
which is an exceedingly poisonous, inflammable gas evolved by
treating any compound of arsenic with nascent hydrogen. The
same apparatus used for hydrogen sulphide may be employed
for this purpose. This gas is to be had by placing in the test-
tube any arsenic salt together with metallic zinc and dilute
ARSENIC. 245
sulphuric acid. This gas is generated in making the u spot
test" (Art. 254) for arsenic, and great care must be used not
to inhale any of it. Allow the acid and zinc to work until the
apparatus is free from air before adding the arsenic compound ;
the gas escaping from the jet should be immediately ignited.
(c) Scheele's Green or Copper Arsenite, CuHAsO3. This com-
pound is to be had by adding an aqueous solution of arsenic
trioxide to an ammonia-copper sulphate solution ; this latter
solution is prepared by adding ammonia to a solution of copper
sulphate until the precipitate, which is at first formed, dis-
solves.
Schweinfurth's Green is a copper aceto-arsenite. (CuOAs2O3)3-
Cu(C2H3O2)2. Both of these compounds are used as pigments,
and are sold under the name of Paris Green. Gardeners use
them as anti-insect powders. Wall paper frequently owes its
green tints to the pre^ance of one of these compounds ; such
paper is dangerous, sometimes giving rise to aggravated cases
of arsenic poisoning. It seems that such papers give off
arsenical vapors or dust, which are disseminated through the
air and absorbed by the pores of the skin and by the lungs.
(d) Arsenic Pentoxide, As2O-;, and Arsenic Acid, II3AsO4.
The first is prepared by dissolving arsenic in strong hot nitric
acid, after which the solution is first evaporated and then fused
at a dark -red heat. Arsenic acid is obtained by treating arsenic
with chlorine as previously explained. From this acid we
obtain the ar senates.
(e) Arsenious Sulphide, As2S3, is the group-reagent precipi-
tate, and may be had by treating any soluble arsenic salt with
hydrogen sulphide. This is a yellow powder soluble in yellow
ammonium sulphide, (NH4)2S2.
(/) /Sodium Arsenate is used to remove the mordant in
calico printing. The impure form thus employed is made by
dissolving arsenic trioxide in sodium hydroxide, after which
sodium nitrate is added ; the solution is then evaporated to
dryuess. (Read R. and S., Vol. II., Pt. I., p. 125.)
246 ARSENIC.
254. Tests for Arsenic and its Compounds. — 1. Metal-
lic arsenic is to be distinguished by its physical properties
and by its giving a garlic odor when heated in the reduc-
ing-flame on charcoal.
2. Solutions or solids are best tested by the " spot or
mirror test." The solid or solution is first treated with
a crystal of potassium chlorate and hydrochloric acid to
oxidize the arsenic (if any be present) to arsenic acid.
The excess of chlorine is expelled, and the prepared solu-
tion is now treated with arsenic-free zinc and dilute
sulphuric acid. Hydrogen arsenide is thus evolved. The
escaping gas is delivered through a jet and is ignited.
Now hold a piece of cold porcelain in the flame. Arsenic
if present is deposited on the porcelain as a bright steel-
gray spot or mirror.
Mate several spots, and make sure that they are arsenic,
thus : —
(a) Try one spot with a drop of yellow ammonium
sulphide ; it turns yellow.
(&) Try another with a drop of hydrochloric acid ; it
does not dissolve.
(<?) Add to another a drop of a solution of bromine or
chlorine in potassium hydroxide ; it dissolves.
(cT) Try another with hot nitric acid ; it dissolves clear.
Then to this clear solution add a drop of silver nitrate ;
no change in color occurs. Now treat the solution with
ammonia vapor, which may be forced against the solution
by blowing through a blow-pipe across the mouth of an
uncorked ammonia bottle ; the solution turns brick-red or
yellow. You may now be assured that arsenic in some
form is present.
3. To distinguish an ar senate from an arsenite. Make a
ANTIMONY. 247
clear solution of magnesium sulphate, MgSO4, ammonia,
and ammonium chloride, NH4C1. To this clear solution
add the unknown solution, a portion of which has been
found to contain arsenic by 2. A white precipitate (in
the absence of phosphates) indicates an arsenate. An
arsenite gives a white precipitate with MgSO4 which is
soluble in ammonia and NH4C1. A solution of arsenic
trioxide in an excess of sodium hydroxide, when treated
with copper sulphate, gives a blue solution from which a
red precipitate of Cu2O is thrown down on boiling.
NOTE. Any arsenic compound in solution gives a yellow precipitate,
As2S3, with H2S.
ANTIMONY.
SYMBOL, SB'">V. — ATOMIC WEIGHT, 120. — SPECIFIC HEAT,
0.0523. — MELTING-POINT, 425°.
255. Occurrence. — Native antimony occurs in small
quantities as scaly masses which are contaminated with
iron, silver, etc.; but its chief source is Stibnite, Sb2S3.
Other ores of less importance also occur.
256. Preparation. — EXP. 154 p. Make a pellet of a
thoroughly pulverized antimony compound, as stibnite, Sb2S3,
with potassium cyanide or with sodium carbonate and a drop of
water. Heat on charcoal in the reducing-flame ; a bright
metallic bead of antimony is obtained. Try the malleability,
etc., of this bead as you did of the lead or silver bead. In com-
parison , how does it behave ?
EXP. 155 P. Pulverulent antimony, or antimony black, may
be prepared by placing a zinc strip in a solution of antimony
chloride, SbCl3 (see Art. 257, (c)). How does the precipitate
compare with those thus obtained in the first group metals?
Preserve this powder for future use.
248 ANTIMONY.
Commercial antimony is prepared from stibnite. The
crude ore is first melted in vessels the bottoms of which
are perforated by small openings. The sulphide is melted
and runs through these openings nearly pure. The sul-
phide is next melted with metallic iron, which combines
with the sulphur, leaving the antimony free and ready to
be drawn off in a molten condition.
By another process the sulphide is converted into an
oxide in a reverberatory furnace. The oxide is then re-
duced by heating it with charcoal or some other reducing
agent.
The antimony of commerce often exhibits a stellated
surface, which is obtained by allowing the purified molten
metal to cool slowly.
257. • Properties, Uses, and Compounds of Antimony.
— Antimony is a bluish-white metal, so brittle that it may
be finely pulverized. It tarnishes slowly in warm, moist
air and burns with a white light when heated to redness
in the air, forming the trioxide, Sb2O3. It vaporizes at a
white heat in the absence of oxygen.
Metallic antimony is used principally in making alloys,
to which it imparts the property of hardness and that of
expansion when cooling from a molten state. Hence it is
extensively employed in manufacturing type-metal.
Antimony is also used in many pharmaceutical prepara-
tions. That form of antimony which is obtained in Exp.
155, and which is an article of commerce, is employed to
impart a metallic surface to plaster casts. It is also used
as a medicine for horses.
EXP. 156 P. Coat a small plaster of paris image with anti-
mony black, and polish until the surface assumes a metallic
lustre.
ANTIMONY. 249
With hydrogen and oxygen, antimony forms both acids
and bases. With acids it forms salts, in which.it plays
the part of a trivalent metal, as in antimony sulphate,
Sb2(SO4)3. It also forms basic salts, in which the group
SbO, which is univalent, takes the place of one atom of
hydrogen. These are called antimonyl salts. Antimonyl
sulphate, (SbO)2SO4, may serve as an example. The prin-
cipal acid of antimony is antimonic acid, H3SbO4, which
closely resembles phosphoric and arsenic acids.
The best solvent for antimony is hot nitro-hydro chloric
acid, and the salt thus obtained (SbCl3) is a good solution
for working purposes.
THE PRINCIPAL COMPOUNDS OF ANTIMONY ARE : -
(a) The oxides, Sb2O3, Sb2O4, and Sb2O5, which give rise to
a series of acids similar to those of phosphorus. (Art. 220.)
Antimonic Acid, H3SbO4, is obtained by oxidizing antimony
in nitric acid. None of these acids are employed for industrial
purposes, although antimonic acid was formerly used as a
medicine.
(b) Tartar Emetic, C4H4KSbO7, which is used in medicine.
It is prepared by dissolving antimony trioxide, Sb2O3, in cream
of tartar or potassium tartrate, KHC4H4O6.
(c) Butter of Antimony, Antimony Trichloride, SbCl3. This
is prepared by dissolving antimony trisulphide, Sb2S3, in hydro-
chloric acid. It is used in staining iron or steel utensils, such
as gun-barrels.
(d) Stibnite, Antimony Trisulphide, Sb2S3, which is one of
the antimony ores and is of a dark-gray color. That which is
obtained by precipitating an antimony salt with hydrogen
sulphide is an orange-colored powder. It is a group-reagent pre-
(.•ifiifate insoluble in dilute acids, soluble in ammonium sulphide.
The pentasulphide, Sb2S5, resembles the trisulphide.
250 ANTIMONY.
(e) Hydrogen Stibide, SbH3. This is an inflammable gas
used in the " spot test" for antimony. This is obtained from
an antimony salt by treating it with zinc and sulphuric acid, as
in preparing AsH3, Art. 254, 2.
GENERAL NOTE. Antimony and its salts are poisonous when taken
internally, but they are neither so dangerous nor so active as arsenic and
its compounds.
258. Tests for Antimony. — 1. Solids containing anti-
mony may be tested in the reducing-flame with sodium
carbonate on charcoal. A silver-white, brittle bead is
obtained.
2. A very delicate test for antimony, free or combined,
is the " spot test." Make several spots by directly treat-
ing the substance with zinc and dilute sulphuric acid.
These spots are distinguished from arsenic spots by the
color. Those of antimony are black or velvety-brown.
More certain distinctions are as follows: —
(«) The antimony spot with, yellow ammonium sul-
phide turns orange.
(6) With hot nitric acid turns white.
(<?) In a solution of bromine or chlorine in potassium
hydroxide it is insoluble.
(d) The white spot, formed in (6), treated with silver
nitrate and ammonia fumes gives no color; but when a drop
of ammonia solution is added, the spot turns black.
3. Upon addition of water to the solution of an anti-
mony salt acidulated with hydrochloric acid, a portion of
the antimony is precipitated as a basic salt soluble in
tartaric acid. (See Art. 267, 2.)
NOTE. H2S forms an oraw^e-colored precipitate, Sb2S3 or Sb2S5, with
any antimony compound in solution.
SUG. See Chemical News, June 5, 1885, p. 267, and June 19, 1885, p.
292, for some delicate tests for antimony.
TIN. 251
TIN.
SYMBOL, SNU'IV. — ATOMIC WEIGHT, 118. — SPECIFIC HEAT,
0.0548. — MELTING-POINT, 230°.
259. Occurrence. — Small quantities of tin occur native.
Its chief ore is Tin Stone, Cassiterite or stannic oxide, SnO2.
This ore occurs in veins in the older schistose arid crystal-
line rocks, and also as nodules or "stream tin " in the beds
of rivers traversing the above-mentioned rocks.
The principal tin mines of the world are in Cornwall
(England), Australia, Bolivia, and Peru. The mines of
Cornwall are the oldest tin-mines known ; they were
probably worked as far back as during the Bronze Age.
260. Preparation. — EXP. 157 P. Into a solution of a tin
salt place a strip of zinc. What results ?
EXP. 158 P. Make a paste of a tin salt with solid potassium
cyanide, KCy, and a drop of water. Heat this paste on char-
coal before the reducing-flaifle. Small beads of tin are thus
obtained with great difficulty. (See Art. 262, 2, for test for tin.)
The first step in its preparation for commerce is to crush
the ore and to remove as many impurities as possible by
washing. The ore is then roasted in revolving, inclined
cylinders through which a continuous blast of air and
flame are passing. In this way volatile substances, such
as arsenic and sulphur, are driven off, while other impuri-
ties are oxidized. The roasted ore is now washed again,
and is thus obtained quite pure. It is now reduced to
metallic tin by mixing it with anthracite and heating it in
a blast-furnace.
The metal is next drawn off and further purified by
liquation, i.e., it is gradually melted in a reverberatory
252 TIN.
furnace ; the pure tin is more fusible than its alloys,
which are present, and melts first. It is then drawn off
and stirred with poles of green wrood ; a dross separates
out and is removed. In this way the tin is brought to ;>
state of great purity.
PROPERTIES, USES, AND COMPOUNDS OF TIN.
261. Tin is a white, lustrous metal which is quite per-
manent in the air at ordinary temperatures and which., in
the absence of oxygen, can be vaporized at a white heat.
It is very malleable, and is extensively used in the form of
thin sheets as tinfoil. When bent or bitten, bar-tin emits
a crackling sound, supposed to be due to the motion of its
particles over one another ; this goes to show that solid
masses of tin probably assume a granular structure.
Tin' can be obtained in a crystalline form in different
ways : (1) Melt it, and allow the molten mass partially to
cool ; pour off the liquid portion, when prismatic crystals
of tin remain. (2) Decompose « chloride of tin, as SnC]2
or SnCl4, by means of a weak galvanic current. (3) Make
a solution of a chloride of tin alkaline, and insert a bright
strip of zinc.
Since tin is quite readily reduced from its ores, it has
been known from an early time. Its uses are many and
its alloys are important.
QUERIES. What is "Block Tin"? What is tin plate, and how is it
made'? What are the uses of metallic tin 1
Tin as a base yields two series of salts, — the stannous
and the stannic salts. These are well typified by the
chlorides SnCl2 and SnCl4.
The tin acids yield two series of salts of small importance,
— the stannates and the met as tan nates.
TIN. 253
The best solvent for tin is hydrochloric acid, stannous
chloride, SnCl2, being the salt produced. Nitro-hydro chloric
acid (with excess of HC1) dissolves tin, forming stannic
chloride, SnCl4. These are good working solutions.
THE PRINCIPAL TIN COMPOUNDS ARE : -
(a) Tin Stone, or Cassiterite, SnO2 ; this is the principal ore
of tin.
Stannic Add, H2Sn<>, may be supposed to originate thus : —
SnO2 + H2O = H2SnO3.
In practice this acid is obtained when calcium carbonate is
treated with an excess of stannic chloride. One of the salts of
this acid, sodium stannate, Na2SnO3, is largely used (as " pre-
paring salts ") in calico printing.
Metastannic Acid probably has the formula H10Sn5015. Both
these acids form salts chiefly with the metals of the fifth group.
(6) Stannous Chloride, SnCl2, and Stannic Chloride, SnCl4.
These salts are used as reagents in the laboratory. How are
they made ?
(c) Stannous Sulphide, SnS, which is a brown powder, while
Stannic Sulphide, SnS2, is a yellow one. These are the group-
reagent precipitates, thrown down by hydrogen sulphide in acid
solutions ; they are soluble in yellow ammonium sulphide.
QUERY. If the solution be a stannic salt, which sulphide is obtained ?
A stannous salt ?
262. Tests for Tin. — 1. Metallic tin is recognized by
its lustre and by the crackling sound when bent or bitten.
2. An unknown solid is tested by the blow-pipe, Exp.
158. If the oxidizing flame be used, a coating of stannic
oxide, SnO2, is formed upon the charcoal around the assay.
This coating is pale yellow when hot, white when cold.
3. A solid insoluble in water is dissolved in hydrochloric
acid, and mercuric chloride, HgCl2, is added (see Art. 246).
254 TIN.
NOTE. At first a white precipitate is obtained, if a stannous salt be
present; this soon turns gray and then (usually after some time) black,
when metallic mercury is found to have been precipitated. The white
precipitate is probably Hg2Cl2. Write the equation. This reaction is of
importance, since by it we may identify both tin and mercury.
4. Aii unknown solution is tested by adding : —
(a) HgCl2 (see 3).
(ft) H2S (see Art. 261 (c)).
(cf) Ammonia and a zinc strip (Art. 261).
QUERY. How can you distinguish between a stannous and a stannic
salt?
263. To separate and identify Arsenic, Antimony,
and Tin. — There is no simple method which is at the
same time very accurate. A fairly good one is the fol-
lowing : Bring the precipitate which contains the sul-
phides of arsenic, antimony, and tin into a small flask, and
boil with concentrated hydrochloric acid as long as the
odor of hydrogen sulphide can be detected. The sul-
phides of antimony and tin are dissolved, while the
sulphide of arsenic remains undissolved. Filter and wash,
and then treat the undissolved substance with hydrochloric
acid and potassium chlorate. It is thus converted into
arsenic acid, which may be detected by means of the
reactions given in Art. 254, 3. Test also for arsenic by
(a) The spot test (Art. 254, 2),
(6) Hydrogen sulphide (a yellow precipitate) .
The solution containing antimony and tin is treated
with zinc, which reduces the compounds to the metallic
state. After a time pour off the solution, wash the
residue with water, and treat with hydrochloric acid.
Only the tin is dissolved. It may be detected by means
of mercuric chloride (see Art. 262, 3). Examine the
residue and convince yourself that it is antimony by
BISMUTH. 255
(a) The spot test,
(b) Hydrogen sulphide (an orange-colored precipitate) .
SUG. Write the equations for the steps involved.
DIVISION B.
The metals of this division of the second group are
those whose sulphides are not soluble in yellow ammonium
sulphide.
BISMUTH.
SYMBOL, Bi'". — ATOMIC WEIGHT, 210. — SPECIFIC HEAT.
0.0305.— MELTING-POINT, 270°.
264. Occurrence. — Bismuth is a comparatively rare
metal. It usually occurs native, but it is always contami-
nated with a small percentage of other metals, such as
iron, copper, lead, silver, etc.
Of its ores Bismuth Ochre, Bi2O3, is the principal one.
Bismuthite, Bi2S3, ranks next in importance. Most of the
bismuth of commerce comes from Saxony.
265. Preparation. — EXP. 159 p. Make a pellet of any
bismuth compound with sodium carbonate and a drop of water.
Heat it in the reducing-flame on charcoal. Try the bead as you
did those of lead, silver, etc. What difference do you find?
Treat the bead with the oxidizing-flame. Note the coating on
the charcoal. This coating, Bi2O3, is characteristic.
EXP. 160 P. Into a solution of bismuth chloride, BiCl3, place
a zinc strip, and proceed as usual. Try the same salt with the
galvanic current.
Bismuth can be extracted incompletely from its ores by
fusion ; the extraction can be made complete by roasting
256 BISMUTH.
them first and afterward fusing them with iron, slag, and
charcoal. The crude bismuth thus obtained is purified by
melting it at the lowest possible temperature on an inclined
plane ; the molten metal runs slowly down the plane while
the impurities remain behind.
Commercial bismuth is also prepared as in Exp. 160.
266. Properties, Uses, and Compounds of Bismuth.
— Metallic bismuth is not employed in a pure state in
any of the arts. It is chiefly used in alloys and in making
pharmaceutical preparations; nearly 25,000 kilograms are
thus consumed annually.
Bismuth is a hard, brittle metal of a grayish-white color
with a distinct tinge of red. It oxidizes but slowly in
the atmosphere, but the gases of the laboratory cause it
quickly to tarnish.
It expands during solidification, and it imparts this
property to its alloys, which are, on this account, used in
making delicate castings. Many of the alloys of bismuth,
especially those with tin, lead, and cadmium, melt at very
low temperatures (see Art. 230). These "fusible metals"
or alloys are used in stereotyping and electrotyping ; they
are also used as solders and for making safety plugs for
steam boilers.
SUG. Explain the use of the safety plug.
Bismuth, like antimony, forms two kinds of salts, those in
which its atom takes the place of three atoms of hydrogen,
as in bismuth nitrate, Bi(NO3)3, and those in which the
group BiO, called bismuthyl, takes the place of one atom
of hydrogen, as in the salt (BiO)(NO3). Salts of the
former class are decomposed by water and transformed
into salts of the latter class, which are known usually as
basic salts.
BISMUTH. 257
Aii acid called bismuthic acid, and supposed to have the
formula HBiO3, lias been described ; but very little is
known regarding it or its salts.
The best solvent for bismuth is nitric acid. Hydrochloric
acid also reacts feebly with this metal. The solutions thus
obtained are good working solutions.
THE PRINCIPAL BISMUTH COMPOUNDS ARE AS FOLLOWS :
(a) Of the Bismuth Oxides, Bi2O2 and Bi2O3 are the principal
ones. Of these two the trioxide Bi2O3 is the 'more important.
It is the chief ore of bismuth, and is used as a pigment.
(b) Bismuth Nitrate, Bi(NO3)3 + 3H2O, is obtained by dis-
solving the metal in nitric acid.
. The Sub-nitrate of Bismuth, BiO . NO3, H2O (of the pharma-
copoeia), is prepared by precipitating bismuth nitrate by
the addition of water to the solution. The sub-nitrate is used
in medicine as a remedy for cholera and dysentery. It is also
used as a cosmetic, under the names of Blanc d'Espange and
Blanc de Fard. It is further used in glazing porcelain, to
which it imparts an iridescent surface. This salt is a white
powder, now known as Bismuth Basic Nitrate.
(c) Bismuthite, Bi2S3, is an ore of bismuth and the group-
reagent precipitate. It is obtained from an acid solution of a
bismuth salt by passing through it hydrogen sulphide. It is
soluble in hot nitric acid.
267. Tests for Bismuth. — 1. Unknown solids are tested
for bismuth by the blow-pipe. When the bead is treated
with the oxidizing-flame, Bi2O3 is formed, and the charcoal
is coated orange-yellow while hot, lemon-yellow when
cold. The edges of the coat are bluish-white when cold.
2. A solution is tested by adding : —
(a) Water, which yields a basic salt, as a white pre-
cipitate insoluble in tartar tc acid. (See Art. 258, 3.)
258 COPPER.
(5) H2S, a black precipitate, Bi2S3, soluble in HNO3.
(<?) Ammonia, a white precipitate, Bi(OH)3.
(d) K2Cr2O7, a yellow precipitate (BiO)2Cr2O7, which is
insoluble in KOH, a distinction from lead.
(e) KI in acid solution gives BiI3, a brown, unstable
precipitate soluble in an excess of HC1.
COPPER.
SYMBOL, Cu'1". — ATOMIC WEIGHT, 63. — SPECIFIC HEAT,
0.0952.— MELTING-POINT, 1090°.
268. Occurrence. — Copper occurs native in large
quantities, and the commercial metal is obtained princi-
pally from this source. The most plentiful deposits are
found in upper Michigan, where masses of the pure metal
weighing many tons have been found. It occurs in sheets
or veins, intersecting red sandstone and trap rocks, but
the largest deposits are found as granular masses mixed
through a rocky matrix. Native copper also occurs in
many other localities, and nearly every deposit is silver
bearing.
The ores of copper occur plentifully and are widely dis-
tributed. The principal ores are : Cuprite, Cu2O ; Copper
Glance, Cu2S ; Malachite, CuCO3 + Cu(OH)2; Azurite,
2CuCO3 + Cu(OH)2; and Copper Pyrites, CuFeS*
The argentiferous copper ores of the Rocky Moun-
tains, especially those of Montana, have of late years
furnished a large amount of the copper in the market, at
times so reducing the price of the metal as to necessitate a
temporary suspension of the mines till a higher price
would render the mining and smelting of the ores more
profitable.
COPPER. 259
269. Preparation. — Native copper usually requires
little treatment except smelting ; but the reduction of its
ores to obtain commercial copper is a somewhat complicated
process of minor interest, at present, to the American
student. From its soluble salts copper may be obtained
by precipitation and by electrolysis.
EXP. 161 P. Place a bright strip of iron in a solution of
copper sulphate, CuSO4. It is soon coated with a film of metallic
copper. Thus try a strip of zinc. What result? Try two
strips at once, one of zinc and one of iron. What result?
Does the iron increase in weight owing to the deposit of cop-
per? In what ratio?
We have many familiar examples of the reduction of
copper from the solution of its salts. In gravity batteries
the copper plates are soon covered with a deposit of
copper ; in electro typing, a metallic film is deposited upon
a wax mould of the type, and this film is afterward strength-
ened by a fusible metal (it is thus that the plates were
prepared from which these pages were printed) ; the hypo-
phosphites, when heated with the solution of a copper salt,
reduce metallic copper ; and the following metals will give
metallic copper with a solution of a copper salt: iron,
zinc, cobalt, nickel, lead, cadmium, bismuth, and tin.
270. Properties, Uses, and Compounds of Copper. —
Owing to its abundance in the native state, copper was
probably the first metal used by man. The prehistoric
copper miners of Lake Superior employed the rudest
methods imaginable for mining and working copper.
They confined their operations chiefly to the sheet-like
veins which were visible at the surface. Owing to the
dip of the rocks only the edge of the sheet was within their
reach. They built wood fires upon the rocks until the
260 COPPER.
stone would crumble and leave a narrow ribbon of copper
exposed ; then, by means of a stone from the lake shore,
which served as a hammer, the ribbon was hammered off
into strips, which were afterward rudely fashioned, by
means of two stones — one a hammer, the other an anvil —
into knives, spearheads, arrow-points, fish-hooks, needles,
and other utensils. The relics of the ancient copper miners
are found in all parts of America, and some of the richest
mines in the world are located upon the sites of prehistoric
mines.
Copper is a tough, malleable metal of a reddish color
which tarnishes quickly in air containing moisture and
carbon dioxide. In the native state it sometimes occurs
in regular octahedral crystals, which are also obtainable
by the electric current. Copper forms two series of salts
and no* acids.
SUG. Examine the copper plates of a gravity battery which has been
in operation several weeks.
EXERCISE. Write an essay, giving the uses of copper and describing
the process of electrotyping.
The best solvent for copper is nitric acid, and a solution of
the salt thus obtained is a good one for practice. Copper
sulphate solution answers the same purpose.
THE PRINCIPAL COMPOUNDS OF COPPER ARE: —
(a) Copper Sulphate, or Blue Vitriol, CuSO4 -f- 5 H2O ; this
salt is used in electrotyping, in calico printing, in the prepara-
tion of Paris green, and for galvanic batteries. How is it
prepared ?
(b) Copper Nitrate, Cu(NO3)2, which is used in calico print-
ing. How is it prepared ?
(c) Cupric Sulphide, CuS, and Cuprous Sulphide, CuJS,
which are the group-reagent precipitates. These are obtained
COPPEE. 261
by passing hydrogen sulphide through the solutions of the cor-
responding copper salts.
(d) The principal Oxides of copper, Cuprous Oxide, Cu2O,
and Cupric Oxide, CuO. The former is used to impart a red
color to glass ; it occurs native. Cupric oxide is used m color-
ing glass green in imitation of the emerald. This oxide also
occurs native as Melaconite. Both oxides may be prepared
artificially. (See R. and S., Vol. II., Pt. I., pp. 329 and 330.)
GENERAL NOTE. The copper salts act as poisons when taken internally.
271. Tests for Copper. — 1. Any compound of copper
may be reduced on charcoal, by the usual method, to
minute red metallic beads.
2. Solutions are tested thus: (a) Make a borax bead
upon a platinum wire ; moisten with the solution and heat
in the oxidizing-flame. If copper be present, the bead will
be green while hot, blue when cold. (5) To the solution
add:-
(1) An excess of ammonia, a blue solution ;
(2) H2S, a black precipitate.
3. Potassium ferrocyanide, K4FeCy6, in dilute solutions
gives a reddish-brown solution ; in concentrated solutions,
a precipitate, Cu2FeCy6, of the same color.
4. Copper chloride colors the Bunsen flame blue.
NOTE. The tests by the blow-pipe for copper are as unsatisfactory as
those for tin ; the reduction occurs, but the beads are of microscopical
dimensions. If the fused mass be rubbed in a mortar spots of copper
become visible when the flux and charcoal are removed by washing.
272. To separate and identify Bismuth and Copper.
-If the substance is a solid, dissolve it in nitric acid;
then add an excess of ammonia. Bismuth hydroxide,
Bi(OH)3, is obtained as a white, flocculent precipitate
powder, while the copper remains in the blue solution as a
cupro-ammoniuin salt. Filter out the precipitate, dissolve
262 CADMIUM.
it in hydrochloric acid and expel the excess of acid ; again
add water, when bismuthyl chloride, BiOCl, is precipitated
as a white powder. This identifies the bismuth; the blue
solution identifies the copper.
CADMIUM.
SYMBOL, CD". — ATOMIC WEIGHT, 112. — SPECIFIC HEAT,
0.0567.— MELTING-POINT, 315°.
273. Occurrence. — Cadmium is a somewhat rare metal,
which is found in nature associated with zinc. Its sul-
phide, CdS, or Greenockite, also occurs in small quantities.
274. Preparation. — In smelting zinc, the cadmium is
oxidized to form the compound CdO, which readily passes
off in dark-yellow vapors. These vapors are condensed in
suitable chambers, and afterward reduced to a metallic con-
dition by heating in closed tubes with charcoal. The impure
metal thus obtained is purified in the wet way, as in : —
EXP. 162 p. Dissolve a bit of cadmium in hydrochloric acid.
After expelling an}T excess of acid, suspend a strip of zinc in
the solution. The cadmium is precipitated as a spongy, gray
precipitate. Collect, fuse to a bead, then oxidize strongly.
What occurs ?
275. Properties, Uses, and Compounds of Cadmium.
— Cadmium was discovered in 1817 by Stromeyer in zinc
carbonate from Salzgitter. It is a tin-white metal, which
vaporizes at 860°. It oxidizes slowly in the air, arid the
surface of the metal is apt to present a yellowish tint,
owing to the formation of the oxide, CdO. It takes fire
if vaporized in the air.
Cadmium closely resembles tin in its physical properties ;
but, unlike tin, it has but few uses in the arts. Cadmium
CADMIUM. 263
amalgam is used in filling teeth, since it is pasty at first,
but afterwards hardens.
The best solvent for- Cadmium is nitric acid. In hydro-
chloric acid and sulphuric acid it dissolves less readily.
Employ the nitrate or the chloride as a working solution.
THE PRINCIPAL COMPOUNDS OF CADMIUM ARE : -
(a) Cadmium Oxide, CdO. How is this compound formed ?
(b) Cadmium Iodide, CdI2, which is used in photography ;
it is prepared by boiling metallic cadmium and iodine in water.
(c) Cadmium Sulphate, 3 CdSO4 -f- 8 H2O, used in medicine
in diseases of the eye.
(d) Cadmium Sulphide, CdS, used as a yellow pigment. It
occurs native, as Greenockite, and is the group-reagent precipi-
tate thrown down by hydrogen sulphide in acid solutions.
276. Tests for Cadmium. — 1. A solid heated on char-
coal in the oxidizing-flame gives brownish-yellow fumes of
CdO, also a coating of the same on the charcoal, if
cadmium be present in sufficient quantity.
2. An acidulated cadmium solution with hydrogen sul-
phide gives a yellow precipitate insoluble in yellow am-
monium sulphide.
3. A cadmium salt colors the borax bead yellow while
hot, colorless when cold.
277. To separate and identify Bismuth, Copper, and
Cadmium. — 1. To the solution containing salts of these
three metals add ammonia in excess. The bismuth is
precipitated, and identified as in Art. 272. The copper
and cadmium remain in solution. The copper is identified
by the blue solution.
2. Separate the copper and cadmium remaining in
264 CADMIUM.
solution thus: 'precipitate these two metals by hydrogen
sulphide ; filter out and wash the precipitate, then add
hot dilute sulphuric acid to the precipitate on the filter-
paper; the copper sulphide is unaltered, while the cad-
mium sulphide is dissolved and runs through, thus
effecting the separation. Or, to the ammoniacal solution
containing copper and cadmium add potassium cyanide
until the blue color is destroyed; then pass hydrogen
sulphide into it, and the cadmium is precipitated as the
yellow sulphide, CdS, to be further identified by Art. 276.
NOTE. Copper sulphide is soluble in potassium cyanide, forming the
double cyanide 6 KCy . Cu2Cy. Cadmium forms a similar cyanide, but it
is decomposed by hydrogen sulphide.
278. To separate and identify the Metals of the
Second Group. — Acidulate the solution containing the
salts of one or all of these metals with hydrochloric acid,
and precipitate by hydrogen sulphide. The precipitate
may be a sulphide of arsenic, antimony, tin, bismuth, cop-
per, or cadmium ; or sulphides of them all.1
Wash the precipitate, and wash it through into an
evaporating dish ; add yellow ammonium sulphide, and
digest for some time, when the sulphides of division A
dissolve, while the sulphides of division B remain un-
altered. Filter and treat the filtrate as in 1, the remaining
precipitate as in 2.
1. Add hydrochloric acid to the filtrate. This decom-
poses the compounds present, and precipitates the sulphides
of tin, arsenic, and antimony. Filter out and wash the
precipitate, and proceed according to Art. 263.
2. Dissolve the precipitate while on the filter-paper in
hot nitric acid, and expel the excess of acid by evaporat-
1 Sulphides of lead and mercury may also be present.
REACTIONS IN GROUP II. 265
ing the solution to dryness. Dissolve in water, and pro-
ceed by Art. 277, for bismuth, copper, and cadmium. (See
Note 2.)
NOTE 1. Should the precipitate fail to dissolve completely in HN03, the
residue is probably mercury, which was present in the original solution as
mercuric salts. Therefore dissolve this residue in nitro-hydrochloric acid,
and test by adding SnCl2. (See Art. 246.)
NOTE 2. Before trying for bismuth, copper, and cadmium, be sure
there is no lead salt in the solution. It is best to try a small portion of
the solution with H2S04 for lead ; should a precipitate occur, add H2S04
to the whole, which will remove the lead as a precipitate. Filter, and
proceed with the solution by 277.
279. Separation of the Metals of Groups I. and II. —
To a cold solution containing one or more metals of both
groups add hydrochloric acid ; the first group is precipi-
tated, but not completely. (See Notes 1 and 2, Art. 278.)
280. Reactions in Group II. — Balance the following
equations, and ascertain what operations they indicate,
and what principles they illustrate : —
(1) As + Cl+H2O=:H3As04+HCl.
(2) As203 + HC1 = AsCl3 + H20,
and AsCl3 + H20 = H3As03 + HC1.
(3) Sb+Cl=SbCl3.
(4) Sn + HC1 = SnCl2 + H.
(5) Bi+HN03 = Bi(N03)3 + N02+H20.
(6) Cu + HN03 - Cu(N03)2 + NO + H20.
(7) Cd + HN03 = Cd(N03)2 + NO + H20.
(8) H3As04 + H2S = As2S3 + H2O + S2.
(9) SbCl3 + H2S = Sb,S3 + HC1.
(10) SnCl2 + H2S = SnS + HC1.
(11) Bi(N03)3 + H2S = Bi2S3 + HN03.
(12) Cu(N03)2 + H2S = CuS + HNO3.
(13) Cd(N03)2+H2S = CdS+HN03.
266 THE RARE METALS OF GROUP II.
(14) As2S5 + Cl + H20 - II3AsO, + HC1
(15) Sb2S5 + Cl + I-I20 = H4Sb207 + HC1
(16) SnS + Cl= SnCl4+ S.
(17) H3As04 + H = AsH3 + H20.
(18) AsII3 + AgNO3 + H20 = H3As03 + Ag + HN03.
(19) H4Sb207 + H = SbH3 + H20.
(20) SbH3 + AgN03 = Ag3Sb + HNO3.
(21) SnCl2 + Zn =_Sn + ZnCl2.
(22) PbS + HN03 = Pb(N03)2 + S + NO + H20.
(23) Pb(NO.,), + H2S04 = PbS04 + HN03.
(24) Bi2S3 + HNO3 = Bi(NOs)3 + S + NO + H2O.
(25) Bi(N03)3 + NH4HO = Bi(QH)3 + NH4N03.
(26) Bi(OH)3+HCl-BiCl3+H20.
(27) BiCl, + H,0 = BiQCl + HC1.
(28) CuS + HN03 = Cu(N03)2 +. S + NO + H2O.
(29) Cu(N03)2 + NH.HO = Cu(NH3)20, NH4N03 + H20.
• (30) CdS + HN03 = Cd(N03)2 + S + NO + H20.
(31) Cd(N08)2 + H2S = CdS+HN03.
(32) Sn + HNO3 = SnO2 + H2O + NO.
(33) Sn + HNO3 = Sn(N03)2 + H2O + NH4N03.
(34) Sb + HN03 = Sb203 + H20 + NO.
(35) Sb + HN03 = Sb02 + H20 + NO.
QUERIES. Which equations show the precipitation of Group II.
Which show special reactions or tests ?
THE RARE METALS OF GROUP II.
GOLD.
SYMBOL, Au'". — ATOMIC WEIGHT, 196.5.
281. Gold always occurs native owing to its feeble chemical
affinity ; it has been known since the earliest times, therefore,
and ever highly prized. It occurs very widely distributed in
the older sedimentary and igneous rocks, and rivers running
GOLD. 267
through these rocks wash down fine particles of gold and sand.
From these sands the miner separates the precious metal by
washing in a shallow pan or cradle. Nuggets of gold of great
value have been found in the various gold-bearing districts,
especially in Australia. The largest deposits of gold are in the
western United States and Australia.
Gold-mining by hydraulic power has been conducted on
an enormous scale in the West. The auriferous deposits are
loosened by powerful streams of water which are directed
against them. Thus the detritus is loosened, and afterward
carried down the mountain slopes in sluices in which are placed
pockets containing mercury. The fine particles of gold are
caught in the pockets, as they readily form an amalgam with
mercuiy.
In quartz-mining, the coarse rocks are crushed b}r machinery,
and the gold likewise extracted by amalgamation.
Gold may be obtained in a pure state by first dissolving any
of its alloys in nitro-hydrochloric acid, when its principal salt,
AuCl3, is formed, and by then adding ferrous sulphate, FeSO4,
thus : —
2 AuCl3 + 6 FeSO4 = 2 Au -f- 2 Fe2(SO4)3 + Fe2Cl6.
The finely divided gold thus obtained may be fused to a yellow
bead on charcoal.
The uses of gold are many, but in a pure state it is too soft
to wear well, hence it is alloyed with silver and copper. Gold
is very malleable and ductile, and does not tarnish in the air.
Its salts are few and of little interest to the beginner. Name
the uses of gold.
The test for gold is the formation of the purple of Cassius,
thus : —
Dissolve the substance in aqua-regia, and expel the excess
of acid. Fill a test-tube half full of water, then add one drop
each of stannous chloride, SnCl2, and stannic chloride, SnCl4 ;
then add a few drops of the solution first made. If gold be
268 PLATINUM.
present, a purple tint appears ; but should the reagents be in
excess, metallic gold is reduced, and the color is brownish
or reddish-black.
NOTE. A gold electroplating solution is made by adding potassium
cyanide to gold chloride.
PLATINUM.
SYMBOL, PT"". — ATOMIC WEIGHT, 195.
282. Platinum is a silver-white, lustrous metal of great value
to the chemist ; without it, many chemical processes would be
impossible, and our knowledge of the rare elements would be
less extended.
It was probably discovered about three hundred years ago,
but its properties and preparation have not been understood for
more than one century.
It occurs only native, and always together with several kin-
dred metals ; it occurs in many localities, but in small quantities.
Platinum is obtained from its ores by first dissolving them in
aqua regia, when platinic chloride, PtCl4, is formed, together with
the chlorides of the accompanying metals ; ammonium chloride is
then added, when a double chloride of platinum and ammonium
is precipitated, viz., (NH4)2PtCl6. This precipitate when heated
yields spongy platinum, which is afterwards fused in lime
crucibles by means of the oxy-hydrogen blow-pipe.
The high melting-point of platinum and its power of resisting
the action of solvents, together with its malleability and duc-
tility, render it most useful in chemical laboratories, where it is
used as wire, foil, battery-plates, crucibles, weights, etc.
Platinum may be welded like iron. It possesses a remarkable
property of condensing gases upon its surface, especially when
in a spongy state, as exhibited when a stream of hydrogen gas
is directed against it ; the hydrogen is thus ignited. Platinum
wire acts similarly, and maty even decompose a gas, as shown
PALLADIUM. 269
when a heated spiral is held in the fumes of ammonia or ether :
the wire continues to glow, so great is the action upon its sur-
face.
The salts of platinum are very numerous, but as the}' are used
very little, they need not be considered here.
Nitro-hydrochloric acid is the best solvent for this metal.
The test for platinum is made thus : —
Dissolve the substance in aqua regia, expel any excess of
acid, add NH4C1, when a yellow crystalline precipitate is thrown
down, thus : —
PtCU + 2 NH4C1 = (NH4)2PtCl6.
Spongy platinum is to be had from this precipitate upon heating
strongly.
GENERAL NOTE. Do not fuse any of the salts of the easily reducible
metals on platinum, since they form fusible alloys with it, as will the
silicon in charcoal, when platinum is heated in contact with the glowing
coals. Do not heat a platinum crucible in a smoky flame, which will
cause it to blister.
Try the purity of platinum by boiling in HC1, then in pure
HNO3. It should not dissolve.
To clean a platinum crucible, fuse acid potassium sulphate,
KHSO4, in it.
PALLADIUM.
SYMBOL, PD". — ATOMIC WEIGHT, 106.
283. Palladium, a silver-white metal resembling platinum,
occurs in connection with gold and platinum ores. It was first
prepared by Wollaston, in 1804. Its uses in the arts are few,
the chief ones being to make graduated scales for astronomical
instruments, to plate silverware, and to take the place of gold
in dentistry. Its best solvent is nitric acid, but spongy palladium
is readily soluble in hydrochloric acid. The nitrate and chloride
270 RUTHENIUM. — IRIDIUM.
of palladium are used to separate chlorine, bromine, and iodine ;
but these reagents are too expensive for general use.
Palladium is detected by adding potassium cyanide to its
solution in hydrochloric acid. A yellowish-white precipitate,
soluble in hydrochloric acid and ammonia, is thrown down. It
also gives a black precipitate with potassium iodide insoluble in
Irydrochloric acid.
RUTHENIUM.
SYMBOL, Ru". — ATOMIC WEIGHT, 103.5.
284. Ruthenium also resembles platinum. It was discovered,
in 1845, by Glaus, in the Ural platinum ores. It is very little
used in the arts, nor are its salts particularly valuable.
Ruthenium is detected by passing hydrogen sulphide through
its solution : the solution turns blue, afterwards brown. Also,
when water is added to its chlorides, an inky, soluble oxy-
chloride is formed.
IRIDIUM.
SYMBOL, In". — ATOMIC WEIGHT, 193.
285. Iridium occurs, as the preceding metals, in small grains,
alloyed with platinum or osmium. It forms a valuable alloy
with platinum, consisting of 1 part iridium and 9 parts plati-
num, which is very hard, elastic, insoluble, and unchanging in
the air ; it also takes a splendid polish and has a small coeffi-
cient of expansion, hence its use in making the standard measure
of the metric system. Iridium alloys dissolve in aqua regia.
Iridium is very refractory ; but it may be fused (as a phos-
phide) at a high temperature by the addition of phosphorus.
The phosphide is adapted to many purposes where hardness
and the property of resisting chemical action are requisite.
The phosphorus can be withdrawn, however, by repeatedly
heating the phosphide in contact with lime. Iridium or the
compounds mentioned are used for the tips of gold pens, for
&HOD1 UM. OSM LUM. T UNGSTEN. 271
" stylographic pens," and sometimes for the bearings of chem-
ical balances. For valuable information concerning Iridium
see The Chemical News for Jan. 1, 1885, p. 1, and Jan. 9,
1885, p. 19.
Iridium is detected \>\ the dark-red crystalline precipitate
formed by adding ammonium chloride to its concentrated solu-
tions.
RHODIUM.
SYMBOL, RH". — ATOMIC WEIGHT, 104.
286. Rhodium also occurs with platinum, and was discovered
in 1804 by Wollaston. It remains in solution after precipi-
tating the platinum with ammonium chloride, and is to be
obtained from this liquor. It is but slightly soluble in an}-
solvent, but its alloys are soluble in aqua regia, when the pro-
portion of rhodium is very small.
Rhodium is reduced from acid solutions by metallic zinc.
OSMIUM.
SYMBOL, Os". — ATOMIC WEIGHT, 199.
287. Osmium likewise occurs with platinum, and is remark-
able as forming a volatile oxide, OsO4. This metal is the
heaviest substance known, its specific gravity being 22.477.
It has never been fused. An alloy of osmium and iridium is
used to tip gold pens ; also to make the bearings of mariner's
compasses.
Osmium is detected by the odor of its volatile compound,
OsO4, obtained by treating its soluble compounds with nitric
acid.
TUNGSTEN.
SYMBOL, Wiv. — ATOMIC WEIGHT, 184.
288. Tungsten occurs most plentifully in Wolfram , a tungstate
of iron and manganese, and has not been prepared in a coherent
state. Its proposed industrial use is to improve tool steel.
272 MOLYBDENUM.
Tungsten is detected by first fusing an}' of its compounds
with potassium hydroxide, and afterwards dissolving the fused
mass in hydrochloric acid ; into this solution a strip of zinc is
immersed ; the solution turns blue if tungsten be present.
MOLYBDENUM.
SYMBOL, Mo". — ATOMIC WEIGHT, 96.
289. Molybdenum is a silver-white metal, occurring in Molyb-
denite, MoS2 ; this sulphide was mistaken, in ancient times,
for plumbago, which substance it closely resembles. Ammo-
nium molybdate is an important test for phosphoric acid (see
App.) . Ammonium phospho-molybdate, 2 (NH4)3PO4 + 22 MoO3
+ 12 H2O, is used in chemistry as a reagent for detecting alka-
loids.
Molybdenum is detected in the same way as tungsten, the
solution turning successively blue, green, and brown.
EXERCISES.
1. Dissolve in nitric acid a small silver coin, and see what metals you
can detect in the solution.
2. Allow a drop of molten bismuth to fall upon the floor, and note what
occurs.
8. What metals have atomic weights of about 104 ? About 195 ? Find
their positions in the table on p. 221. By what similarity of properties are
they marked 1
4. How can you show by the same table that iron, cobalt, nickel, chro-
mium, manganese, and copper are closely related ?
5. What elements are closely related to phosphorus ?
6. Dissolve a bit of worn-out " gold " jewelry in aqua regia and
determine what metals are present.
7. Compute the atomic heat and the atomic volume of the common
metals of the second group.
8. Analyze a sample of "antimony black" obtained from the drug
store.
" Pulverized Sb.2S3 is known as ' antimony black ' although a mixture of
anthracite and marble is sometimes sold under that name." — WARDER.
CHAPTER XVII.
THE THIRD GROUP METALS.
290. The metallic hydroxides and sulphides of this
group are soluble in dilute acids, but insoluble in alkaline
solutions. There are different methods of separating this
group from the other groups ; and, moreover, the individ-
ual metals of the third group may be separated and
identified by different processes. We shall here pursue
that plan which is as simple as possible, and which is in
most cases preferable. Just as in the preceding group,
we may likewise divide the commonly occurring metals of
this group into two classes. We may obtain the precipi-
tates of this group as follows: Suppose the solution con-
tains any or all the metallic salts of the five groups. The
first and second group metals are removed by hydrochloric
acid and hydrogen s^phide ; in any chromium compound
that may be present chromium is now combined as a base
through the agency of the reagents employed, while iron
salts by the same means are reduced to the ferrous con-
dition. In case no first and second group metals are
present, it is still necessary to use hydrochloric acid and
hydrogen sulphide to insure that chromium may be present
as a base (Art. 303, 4, Note). The solution is now boiled
to expel any excess of hydrogen sulphide ; then nitric acid
is added and the whole boiled a moment to oxidize iron
salts to the ferric condition ; the solution is now ready for
the application of the group reagents.
274 THE THIRD GROUP METALS.
Ammonia arid ammonium chloride are immediately
added ; thus the hydroxides Fe2(OH)6, Cr2(OH)6, and
A12(OH)6 are precipitated. This precipitate is now re-
moved by filtering. To the filtrate ammonium sulphide,
(NH4)2S, is added ; this gives the precipitates NiS, CoS,
MnS, ZnS. Let us indicate these divisions as in the
previous group : —
r Nickel.
f -ron- C b It
DIVISION A Chromium. D™«w B J
I Aluminum. [
Sue. Try to determine why it is necessary to oxidize ferrous to ferric
salts, and why chromium, if present, must be a base.
The most strongly marked characteristics of the metals
of this group are as follows : —
(a) Their surfaces gradually oxidize in the air, forming
oxides most of which are to be reduced to the metallic state
only at a white heat in the presence of reducing agents.
(6) On charcoal many of these metals cannot be readily
reduced from their compounds by means of the blow-pipe.
(<?) Their oxides and hydroxides ^ie insoluble in water;
in certain cases, however, the hydroxides are soluble in an
excess of the alkali used as a precipitant.
(c?) None of the common metals of this group give
spectra or color the flame unless the temperature be
higher than that of the Bunsen flame.
(e) Most of these metals give a characteristic color to
the borax or microcosmic bead, when heated on a platinum
loop before the blow-pipe.
GENERAL NOTE. Most of the common metals of the third group and
those of the fourth and fifth cannot be reduced to the metallic state by
any means likely to be at the command of workers in small laboratories.
On this account, and by reason of the fact that the general principles un-
IKON. 275
deriving the reduction of the metals have been illustrated previously, the
attention of the student may now most profitably be given to the analyti-
cal reactions of the various metallic salts. In case time permits, it would
also be well to encourage the student to prepare such of the salts as the ap-
paratus, chemicals, etc., at his command will permit. Owing to his previous
training, the student will now be able, in his work, to devise methods and
to keep his notes accurately and intelligently. In furtherance of this object
many topics, by aid of the descriptions given, may be rewritten by him in
the form of experiments. Let him also make tables for each metal show-
ing the effect of precipitants and solvents upon the salts of that metal.
IRON.
SYMBOL, FE". — ATOMIC WEIGHT, 56. — SPECIFIC HEAT,
0.1140. — MELTING-POINT, A WHITE HEAT.
291. Occurrence. — Native metallic iron occurs in in-
significant quantities, and as meteoric iron ; it is also said
to fall everywhere and constantly upon the earth as a fine
dust. Meteorites frequently contain iron ; the largest
mass on record weighs 32,000 pounds, while others of less
weight are frequently found.
The iron compounds are present in most rocks arid soils,
and play an important part in the animal and the vegetable
economy ; the color of the blood and of all vegetation is
due to the presence of iron.
The ores of iron, from which the metal is reduced,
usually occur in somewhat limited areas, but in many
localities. In the United States, the Lake Superior region
supplies very valuable ores; the Southern, Eastern, and
Western States likewise produce immense quantities. In
the Old World, Scandinavia, Elba, Great Britain, the
Ural Mountains, and some other parts of Europe are
famous iron-producing regions.
Among the many ores of iron, space permits a descrip-
276 IRON.
tion of only those varieties which are valuable to com-
merce.
1. Haematite, or /Specular Ore, Fe2O3, occurs in veins,
beds, and pockets. This ore is frequently of great purity,
and is the ore chiefly employed in the reducing furnaces
of the United States.
It assumes different modifications, as : —
.(#) Amorphous Ore, a reddish, massive variety which
resembles, in appearance, " iron rust."
(5) Micaceous Ore, occurring in glittering scales.
(c) G-ranular Ore, of a crystalline structure.
(d) Grape Ore, occurring in masses resembling bunches
of grapes, and therefore frequently called botryoidal ore.
2. Brown Haematite, Fe2O3 + Fe2(OH)6, is frequently
known as bog ore, and is the ore chiefly employed for
reduction in France and Germany.
3. Magnetite, or Lodestone, Fe3O4, is the well-known
natural magnet ; it occurs in all the previously mentioned
localities. A good quality of iron is obtained from this
ore, but its reduction is somewhat difficult.
4. Siderite, or Spathic Ore, is a carbonate of iron, FeCO3,
containing also the carbonates of calcium, magnesium,
and manganese. The renowned Styrian steel is manufac-
tured from this ore at Erzberg.
5. Argillaceous Ore, or Clay Iron-Stone, occurs in connec-
tion with coal, and is the ore chiefly employed in England.
This ore also occurs in Maryland, Pennsylvania, and Ohio.
292. Preparation of Iron. — The reduction of iron as
now carried on is one of the greatest industries of the age.
In the primitive stages of this industry the process was a
simple one which, however, permitted the employment of
the purest ores only.
IKON.
277
A simple hearth was built with an opening at the
bottom to admit a blast of air from a rude bellows, while
another opening allowed the exit of the slag. After heat-
ing the hearth, the ore and fuel were arranged in layers,
and a continuous blast was maintained. In from four to
six hours a porous " bloom " of iron weighing from 5 to
30 Ibs. was obtained. This "bloom" was then brought
into the requisite shape by alternately heating and ham-
mering it. Since the limits of our work prevent the
giving of a consecutive history of the improvements on
this method, it must suffice to say that the developments
have been such that the present processes are applicable
to any ore.
The process now employed consists of two steps : —
(a) The production of cast iron from
the ore ;
(6) The manufacture of wrought iron
from cast iron.
(a) The ore is first crushed and then
mixed with crushed limestone and coal,
when it is ready for the furnace.
The best form of blast-furnace is
shown in Fig. 19. This furnace, which
is from 50 to 90 feet high, and from 14
to 20 feet broad in its widest part, is
constructed of masonry and has a lining
of fire-brick. The whole stack is en- FJG- 19-
closed down to the point A in riveted iron boiler-plates.
The masonry of the stack does not extend below A before
the furnace is prepared to go into blast, but strong iron
pillars (not shown in the cut) extend from that point to
the ground. The hearth H consists of fire clay. It is
278 IRON.
here that the molten metal collects. This hearth has two
openings, the upper one for removing the slag, the lower
one for drawing off the molten metal. The top of the
stack D is funnel-shaped and is closed by an inverted
cone, E, which lowers to admit the ore, fuel, etc., when
the furnace is in operation, and which can be quickly
raised again to close the opening.
.When a furnace is about to go into blast, the spaces H
and B are filled with cord-wood, after which the whole
portion below A, excepting the egress openings for the
molten metal and slag, is enclosed by masonry ; a number
of blow-pipes, or " Tuyeres," C, are also inserted, through
which powerful blasts of air are to be driven.
The wood is now ignited, the blast is started, and
coal is introduced through the funnel D. When the
interior of the furnace is sufficiently heated, the ore, mixed
with coal and limestone, is admitted through D at regular
intervals.
Barring accidents, a furnace, after going into blast, runs
night and day, shutting down but once or twice a year to
renew the fire-brick linings and to make other needed
repairs.
The chemical changes which take place in the stack are
not thoroughly understood, but the products obtained are
cast iron, a glassy slag, carbon dioxide gas, carbon mon-
oxide gas, hydrogen, and perhaps certain hydro-carbons,
graphite, and cyanogen. The gases are not wasted, but
are led through the pipe G — whose opening is seen at
F — into burners situated under the boilers which supply
steam to the engines used to drive the crushers, hoisters,
air-condensers, etc., etc.
The cast iron of the blast-furnace is not pure iron, but
contains carbon and silicon, with traces of phosphorus,
IRON. 279
arsenic, and sulphur, besides other metals in small quan-
tities.
(6) The second step is changing the cast iron into
wrought iron. This is accomplished by the processes
termed "refining'' and "puddling," which consist in
burning out the impurities and hammering the metal into
coherence. The metal is then rolled into bars and sent to
market.
Another important branch of the iron industry is the
manufacture of steel.
Steel, in its chemical composition, stands midway
between cast and wrought iron in the amount of carbon
and silicon which it contains. It was formerly prepared
at great expense by heating wrought iron in contact with
carbon. The Bessemer process, however, has cheapened
the production of steel, so that its use is becoming
general.
This process is, briefly, as follows : —
Cast iron is melted in a cupola furnace and run into
an egg-shaped vessel called a " converter," which is so
arranged that a blast of air may be driven up through
the molten iron.
In this way the carbon and silicon of the cast iron are
burned out; then a sufficient quantity of pure molten cast
iron is added to convert the whole into steel. In about
half an hour from five to twelve tons of steel are thus
obtained from one converter.
293. Properties, Uses, and Salts of Iron. — Pure iron
is a nearly silver-white metal, which rusts easily when
exposed to dampness, ferric hydroxide and oxide being
formed upon its surface. Iron is a very tenacious metal,
and possesses the peculiar property of softening before it
melts, thus allowing different pieces to be " welded."
280 IRON.
The uses of iron and steel are so numerous and various
that this age has well been termed the IRON AGE.
EXERCISES. Name the uses of cast iron; wrought iron ; steel. What
is meant by the " temper " of steel ? How is steel tempered ? How is cast
iron chilled ? What is malleable iron 7 How are twist gun-barrels made ?
Laminated steel gun-barrels ? Damascus steel gun-barrels? (Ask your
gunsmith about the value and manufacture of gun-barrels.) How do cast
iron, wrought iron, and steel differ in their properties ? Can you "clinch"
a 'common cut iron nail 1 Try it. Heat it red hot, cool it slowly, and then
try. How can you draw the temper of a steel tool 1 State the effects of
magnets upon iron and steel. In what acids will iron dissolve ? Which
acid is the* best solvent 7 (Try, at least, H2S04, HN03, HC1, and aqua
regia.) What differences are there between the composition of cast iron,
of wrought iron, and of steel 1
Iron, like mercury (which see), forms two series of
salts. Its compounds are numerous and important. Only
a few of the most important ones can be given here.
THE PRINCIPAL COMPOUNDS OF IRON ARE : -
(a) The Oxides, FeO, Fe2O3, and Fe3O4. The last two,
which are ores of iron, have already been noticed. Ferric Hy-
droxide, Fe2(OH)6, is a brownish precipitate obtained by adding
ammonia to a cold solution of ferric chloride, Fe2Cl6 ; this
hydroxide is used iu medicine. It is the group-reagent precipi-
tate. The first oxide, FeO, is unimportant.
(b) Ferric Chloride, Fe2Cl6. This salt is prepared by dis-
solving iron wire in Ivydrochloric acid, after which the solution
is thoroughly saturated with chlorine gas. It is used in medi-
cine ; in the laboratory it is employed as a reagent.
Ferrous Chloride, FeCl2, is prepared by dissolving iron wire
in hydrochloric acid ; it is also used in medicine.
(c) Ferrous Sulphate, Green Vitriol or Copperas, FeSO4 +
7 H2O. This salt is obtained when iron or ferrous sulphide is
dissolved in sulphuric acid. It is used as a reagent and for
preparing inks, dyes, and Prussian blue. Its uses as a deodor-
ant and a disinfectant have alreadv been mentioned.
IRON. 281
(d) Ferrous Sulphide, FeS. This useful compound is made
by stirring a portion of molten sulphur with a white-hot rod of
wrought iron until the sulphur disappears. It is used in the
laboratory for obtaining hydrogen sulphide.
(e) Iron Pyrites, or Fool's Gold, FeS2, occurs native as yel-
low, shining cubical crystals. It is found in various geological
formations. Its principal use is for manufacturing sulphuric
acid.
Fe2S3 probably occurs in magnetic pyrites. It can be pro-
duced artificially, but is of small importance.
(/) Potassium Ferrocyanide, or Yellow Prussiate of Potash,
K4FeCy6. This salt is obtained by heating scrap iron in closed
iron retorts with potash and animal matter such as hoofs, horns,
hides, feathers, etc.
This salt is of great importance, since it serves as the point
of departure in the preparation of all the cyanogen compounds.
In the laboratory it is used as a reagent for detecting iron. In
the arts it is used for preparing Prussian blue, (Fe2)2(FeCy6)3
or Fe7Cy]8. This pigment is obtained when ferrous sulphate
or ferric chloride is added to the ferrocyanide, K4FeCy6.
(g) Potassium Ferri-cyanide, or Red Prussiate of Potash,
K3FeCy6, is obtained by oxidizing K4FeCy6 by the action of
chlorine. It is used to some extent as a reagent in the labora-
tory.
294. Tests for Iron. — 1. It is best to dissolve solids,
and to test by 2. The blow-pipe tests are not satisfactory
to beginners.
2. Any solution is tested for iron by the reagents,
potassium sulpho-cyanide, KCyS, and the ferro and ferri-
cyanides, K4FeCyc.and K3FeCy6. The change produced
upon any solution depends upon whether the substance
under examination contains a ferrous or a ferric salt.
These changes are exhibited by the following table : —
282
CHROMIUM.
REAGENT.
F.KKKIC SALT.
FERROUS SALT.
KCyS
K4FeCy6
K3FeCy6
lied sol. [Fe2(CyS)6]
Deep blue prec. [Fe4 ( FeCy6 ) 3]
No prec. Reddish brown sol.1
No change.
Pale blue prec. [K2Fe(FeCy6)]
Deep blue prec. [Fe3(FeCy6)2]
3. By employing a ferric salt (Fe2Cl6) and a ferrous
salt (FeSO4) as reagents, it is evident that the table given
in 2 affords tests for Ferrocyanic and Ferricyanic acids or
their derived salts.
EXERCISES.
1. Ignite in the Bunsen flame any ferrous salt, as FeSO4, on platinum
foil. Try the residue with a magnet. Is it magnetic ? What oxide of iron
is thus obtained ?
2. Heat any ferric salt, as Fe2(N03)6, on charcoal in the reducing-flame.
Do you obtain the same oxide as before ?
3. Prepare, solid, Fe2Cl6, and test for ferric salts. Write out a descrip-
tion of the process, etc., in the form of an experiment.
4. In what experiment is the action of iron in the borax bead described ?
When have you prepared FeSO4 ?
5. Try to precipitate a solution of FeSO4 with NH3 and NH4C1. Acidu-
late a fresh portion of the solution with HN03 ; boil, and try as before.
Explain.
CHROMIUM.
SYMBOL, CR'". — ATOMIC WEIGHT, 52. — SPECIFIC HEAT, 0.09975.
— MELTING-POINT (HIGHER THAN THAT OF PLATINUM).
295. Occurrence. — Chromium is a somewhat rare
metal which never occurs free in nature. Its chief ores
are Crocoisite, or chrome yellow, PbCrO4, and chrome iron-
stone, Cr2O3(FeO). The color of many minerals is due to
the presence of traces of the chromium compounds.
296. Preparation. — Chromium is not employed in the
metallic state. It is obtained for scientific purposes by
1 If the color is very dark, dilute the solution until you can see whether a blue pre-
cipitate is not also formed. In such a case, you have both the ferric and ferrous salts
present.
CHROMIUM. 283
mixing its oxide with sugar, after which the mixture is
strongly heated in a lime crucible. Thus obtained, it is a
gray, crystalline powder.
297. Properties, Uses, and Compounds of Chromium.
— Metallic chromium presents a crystalline, silvery ap-
pearance under the microscope. Its principal use is to
harden steel, to which it imparts a superior hardness.
The best working solutions for chromium are solutions of
chromous chloride, CrCl2 ; chrome alum ; potassium chro-
mates ; or the acid chromate. Metallic chromium is solu-
ble in hydrochloric acid, CrCl2 being obtained.
THE PRINCIPAL CHROMIUM COMPOUNDS ARE : -
(a) The Oxides, Cr2O3 and CrO3. Chromic oxide, Cr2O3, is
used in coloring glass and enamel green. It is obtained by
fusing potassium bichromate, K2Cr2O7, with sulphur or with
ammonium chloride, after which the fused mass is treated with
water. The oxide CrO3 may be regarded as the anhydride of the
hypothetical chromic acid, H2CrO4 : H2O + CrO3 = H2CrO4.
Chromic Hydroxide, Cr2(OH)6, is the group-reagent precipi-
tate obtained by adding ammonia and ammonium chloride to
the solution of a salt in which chromium is combined as a base.
Guignet's Green, Cr2O(OH)4, is now largely used as a pig-
ment ; it is sold in drug stores under the name chrome green.
This pigment is prepared by fusing potassium bichromate mixed
with crystallized boric acid in quantities proportional to the
molecular weights of the substances employed. The fused
mass is then ground to a fine powder.
(b) Potassium Chromium Sulphate, or Chrome Alum,
K2Cr2(SO4)4 + 24 H2O. This salt is obtained as a by-product
in the manufacture of alizarine and in many other oxidations.
It is used in dyeing, tanning, and in calico printing.
(c) Potassium Chromate, K2CrO4. This salt is obtained by
284 CHROMIUM.
adding potassium hydroxide to a solution of potassium bichro-
mate. It is used in the laboratory as a reagent.
(d) Potassium Bichromate, K2Cr2O7, or Acid Potassium
Chromate. This is an important salt obtained from chrome
iron ore by three steps : (1) Roasting the ore to oxidize it;
(2) fusing the roasted ore with lime and potassium carbonate ;
(3) lixivating the fused mass with as little water as possible,
and then treating the liquor with sulphuric acid. This salt is
used to prepare chrome yellow, PbCrO4, to dye wool, and to
prepare other chromium compounds. It is also used in the auto-
type process and as a reagent.
The relation between potassium chromate and bichromate
may be understood best by considering the acids from which
they are derived. Hypothetical chromic acid probably has the
composition H2CrO4. Its normal potassium salt is K2CrO4. If we
imagine chromic acid to lose water according to this equation, —
2 H2Cr04 - H20 = H2Cr2O7,
we have left the acid from which potassium bichromate, K2Cr2O7,
is derived, bichromic or pyrochromic acid.
QUERY. Are there any similar relations met with in connection with
sulphur compounds ? Compare carefully sulphuric and chromic acids.
(e) Lead Chromate, or Chrome Yellow, PbCrO4. This com-
pound occurs in nature as crocoisite, and is also prepared
artificially by precipitating a lead salt with potassium bichro-
mate ; used in calico printing and as a pigment. Chrome Red,
Pb2CrO6, and Chrome Orange, a mixture of chrome red and
chrome yellow, are much used as paints.
The following peculiarities will be noticed concerning
the chromium compounds : As a base it forms the chro-
mous and chromic salts, of which chromous chloride, CrCl2,
and chromic chloride, Cr2Cl6, may be taken as examples.
As an acid-forming element, it gives rise to three series
of salts. — the chromites, chromates, and the bichromates ;
CHROMIUM. 285
of these, ferrous chromite, FeCr2O4, lead chromate, PbCrO4,
and potassium bichromate, K2Cr2O7, may be taken as
representatives.
NOTE. As the chromates do not yield precipitates with NH3 and
XH4C1, it is advisable to use a salt like chrome alum, in which chromium
is a base, for the solution which the beginner is to analyze for the third
group metals. The use of HC1 and H2S may then be dispensed with.
298. Tests for Chromium. — 1. Chromium, free or in
compounds, gives the borax or microcosmic bead an
emerald-green color.
2. Fuse the substance to be tested on platinum foil or
porcelain with KNO3 and Na2CO3. These reagents will
oxidize any chromium present to a chromate. Now dis-
solve the yellow mass in water, acidify with acetic acid,
and add lead acetate ; a dense yellow precipitate, PbCrO4,
indicates chromium.
3. A chromate or a bichromate may be recognized by
adding : —
(a) H2S, when the color changes to green.
(5) Lead acetate, which gives yellow lead chromate,
PbCr04.
(V) Silver nitrate, which gives brownish-red silver
chromate,
NOTE. Potassium chromate is a yellow crystalline solid, while the
bichromate is of a red color.
EXERCISES.
1. Prepare (and describe its preparation as an experiment) Cr203.
2. Similarly prepare and describe Guignet's green.
3. Likewise prepare and describe K2Cr04 and chrome yellow.
4. Try to precipitate a chromate with NH3 and NH4C1. Acidulate with
HC1 a fresh portion of chromate, pass H.2S, and try the same precipitants.
What results ? Warm a chromate with (NH4).2S. What results?
286 ALUMINUM.
ALUMINUM.
SYMBOL, AL'". — ATOMIC WEIGHT, 27. — SPECIFIC HEAT,
0.2140. — MELTING-POINT, 700° C.
299. Occurrence. — Next to oxygen and silicon, alu-
minum is the most plentiful and widely-occurring ele-
ment. It is the basis of all clayey soils, and occurs as
feldspar or K2Al2Si6Oi6 in granite, gneiss, syenite, trachite,
porphyry, etc.
Albite, a sodium feldspar, also occurs in large quan-
tities. Kaolin, or porcelain clay and china clay, is feld-
spar which has been disintegrated and decomposed by
exposure to the atmosphere.
The different varieties of garnet, mica, and slate stones
are iiriportant silicates of aluminum and other metals.
Alumina or aluminum oxide, A12O3, is known as corun-
dum or emery when coarse, but when crystallized it con-
stitutes the important jewels sapphire, ruby, oriental
emerald, oriental topaz, and oriental amethyst.
300. Preparation. — Metallic aluminum is prepared
from bauxite, Al2Fe2O8H4.
From this substance the oxide of aluminum, A12O3, is
first prepared thus : The bauxite is heated with soda in a
reverberatory furnace, when a soluble compound of sodium
and aluminum is formed; this compound is dissolved in
water, and a current of carbon dioxide passed through the
solution, precipitating the required alumina.
This oxide is then mixed with charcoal and sodium
chloride and heated to a white heat ; then chlorine gas is
passed through the mixture, thus forming a volatile double
ALUMINUM. 287
chloride of aluminum and sodium, from which the metal
is obtained by fusing with metallic sodium and cryolite.
Aluminum (alloyed with copper)- has lately been ob-
tained directly from corundum, A12O3, by reduction with
carbon, in the absence of air, in an electric circuit.
301. Properties, Uses, and Salts of Aluminum.—
Aluminum is a white, malleable metal which does not tar-
nish or oxidize under ordinary circumstances. Its extreme
lightness (sp. grav. = 2.67), elasticity, tenacity, and the
fact that it is not easily oxidized fit it for many uses which
the cost of its production alone prevents. It is chiefly
employed at present in making philosophical instruments.
Aluminum forms valuable alloys with copper and silver.
. The lest solvent for aluminum is hydrochloric acid.
THE PRINCIPAL COMPOUNDS OF ALUMINUM NOT PREVI-
OUSLY MENTIONED ARE : —
(a) Sodium Aluminate, Na6Al2O6(f)> obtained by fusing
bauxite with sodium sulphate and carbon ; it is used as a mor-
dant in dyeing and calico printing, for preparing colored lakes,
and for sizing paper.
(b) Aluminum Sulphate, A12(SO4)3. This is used in immense
quantities as a mordant and for weighting paper, and is obtained
by roasting kaolin, which is then dissolved in sulphuric acid, and
the solution evaporated till it will solidify when cool.
(c) The Alums, of which there are many. We give the
formulae of the most important: K2A12(SO4)4 -f 24 H2O ;
Ag2Al2(S04)4 + 24 H20 ; and (NH4)2A12(SO4)4 + 24 H2O.
(d) Aluminum Hydroxide, A12(OH)6, obtained by adding
ammonia to a soluble salt of aluminum ; it is the white, gelati-
nous group-reagent precipitate.
(e) Phosphates of Aluminum. The principal one is tur-
quois, a well-known jewel, which owes its bluish or greenish
288 ALUMINUM.
color to the presence of copper. Its formula is A12PO4(OII)3 4-
ILiO ; the ancient gem was cut from odontolite, a fossil tooth or
bone. s
(/) Silicates of Aluminum. The principal ones are : —
Topaz, Al2F2SiO4. A yellowish-colored jewel whose coarser
forms are frequently nsed instead of emery for polishing purposes.
'Beryl, Al2Be3Si6O18. A green variety, the true emerald, is
used in jewelry. What is the " oriental emerald?"
Lapis Lazuli, a bluish mineral of unknown chemical con-
stitution containing silicates of aluminum and sodium, besides
sulphur. It is used for ornamental purposes, and when pow-
dered is known as ultramarine, a valuable paint. The best
ultramarine is now manufactured in large quantities by fusing
together a very fine variety of clay, sand, sulphur, and resin.
302. Tests for Aluminum. — 1. Solids are fused with
Na2CO3or HKSO4 and then dissolved in hydrochloric acid.
This solution is treated as in
2. Add an excess of NH4C1 and ammonia to the solu-
tion. Aluminum gives a white, gelatinous precipitate,
A12(OH>.
3. The solution may be tested further by adding to
another portion : —
(CL) Na2CO3, — the precipitate, A12(OH)C;
(7>) Na2HPO4, — a white precipitate, A12(PO4)2, soluble
in KOH, insoluble in acetic acid.
EXERCISES.
1. Obtain at your druggist's different kinds of alum, and determine what
bases are present.
2. For a valuable paper on alum as a purifier for drinking-water, see
the "Chemical News," May 22, 1885, p. 241.
3. Does aluminum act both as a base and an acid-forming element ?
4. Ask your jeweller to show you specimens of turquois, beryl, topaz,
ruby, sapphire, emerald, and lapis lazuli. How do their values compare ?
5. Examine baking powder for alum.
ALUMINUM. 289
303. To separate and identify Iron, Chromium, and
Aluminum. — From the solution remove the first and
second group metals (if any be present) by means of HC1
and H2S. Boil the nitrate (if HC1 and H2S were em-
ployed) to expel all the hydrogen sulphide. If any iron
salts be present, they are now in the ferrous condition;
therefore add a little nitric acid, and boil a short time to
oxidize the ferrous to ferric salts. The solution is now
prepared for the following treatment : —
1. Add ammonia till the solution is alkaline, and then
add ammonium chloride ; the precipitate obtained may be
any or all of the hydroxides, Fe2(OH)6, Cr2(OH)6,
A12(OH)6. Filter and wash the precipitate.
2. Pierce the point of the filter-paper, and wash the pre-
cipitate through into a beaker-glass; add potassium or
sodium hydroxide, and boil for several minutes. The
hydroxides of iron and chromium remain unchanged,
while the aluminum is dissolved. Now filter the contents
of the beaker, and treat the precipitate for iron and
chromium as in 3 and 4. Treat the filtrate for aluminum
as in 5.
8. Dissolve a portion of the precipitate in HC1, and test
for iron as in 294, 2. If iron be present, test a portion of
the original solution to determine if the salt be in the fer
rous or ferric condition.
4. Fuse on platinum foil a second portion of the same
precipitate with sodium carbonate and potassium nitrate
Any chromium present is thus oxidized to a chromate.
Dissolve the fused mass in water, and test by 298, 3.
Test the original solution for chromates.
NOTE. If the original solution contained chromates, they would be re'
duced by H2S to salts in which chromium would give a precipitate with
NH4C1 and NH3. If no first and second group metals were present, and
290 NICKEL.
HC1 and H2S were not employed, it would be necessary to test the original
solution (unless it is colorless) directly for chromates, since chromates
do not give Cr2(OH)6 with NH3 and NH4C1. In case the solution con-
tains the salts of the metals of division B or those of the fourth group, it
is necessary to use HC1 and H2S.
5. To the filtrate from 2 add sufficient hydrochloric
acid barely to acidify it ; then add ammonia ; a white
precipitate, A12(OH)6, identifies aluminum.
SUG. Explain the significance" of these equations : —
(1)2 H2Cr04 + 6 HC1 + 3 H2S = Cr2Cl6 + 3 S + 8 H2O.
(2) 2 Fe" + 8 HN03 = Fe"'2(NO8)6 + 2 NO + 4 H20.
hot
(3) Fe'"2(N03)6 + H2S - 2 Fe"(N03)2 + 2 HN03 + S.
Which equation shows the transformation of chromium in a chromate
to chromium as a base ? Which equation shows the change of iron from
the ferric to the ferrous condition ?
NICKEL.
SYMBOL, Ni". — ATOMIC WEIGHT, 58. — SPECIFIC HEAT,
0.1080. — MELTING-POINT, NEARLY A WHITE HEAT.
304. Occurrence. — Nickel never occurs native, but its
ores are usually found in connection with cobalt ores.
Meteoric iron always contains nickel. Its chief ores are
Kupfer-nickel, NiAs, — which is the most important ore,
and which is found in Saxony, Styria, and the United
States, — Nickel-glance, Ni(AsS)2; Breithauptite, NiSb ;
Nickel-blende, NiS. (Also see Cobalt.)
305. Preparation, — Metallic nickel is now obtained
mostly in the wet way. The ores are roasted, and then
dissolved in hydrochloric acid ; this solution usually con-
tains iron, cobalt, and copper, which accompany the nickel
in its ores. In such cases the iron salts are oxidized by
chlorine and then precipitated by adding limestone ; the
NICKEL. 291
copper is precipitated by hydrogen sulphide and the cobalt
by bleaching-powder. The remaining clear solution con-
tains the nickel, which is now precipitated by the addition
of an alkali, usually sodium hydroxide. From the nickel
hydroxide thus obtained the nickel is reduced by fusion
with charcoal.
306. Properties, Uses, and Salts of Nickel. — Nickel is
a white, hard metal, scarcely tarnishing in the air, and
susceptible of a brilliant polish. It is accordingly used
extensively in coinage and in plating other metals. The
salt used in electro-plating is a double sulphate of nickel
and ammonium. Its chief alloy is German silver.
SUG. For a valuable paper on Electro-Nickel Plating as an Industry,
which gives the history and development of the details of the processes
employed in electro-nickel plating, see " Scientific American Supplement "
for May 10, 1884, p. 6957.
Nickel, like iron, can be welded, and is likewise at-
tracted by the magnet.
The best solvent for nickel is dilute nitric acid.
THE PRINCIPAL COMPOUNDS OF NICKEL ABE: —
(a) The Oxides, NiO and Ni2O3. We must also note the
apple-green hydroxide, Ni(OH)2, which may be obtained b}T the
action of an alkali, as NaOH, on a solution of a nickel salt.
(b) Nickel Sulphate, NiSO4 -f 7 H2O. Prepared by dissolv-
ing the metal or its hydroxide in sulphuric acid.
(c) Nickel Ammonium Sulphate, Ni(NH4)2(SO4)2 + 6 H2O.
Obtained by adding ammonium sulphate to nickel sulphate.
(d) Nickel Sulphide, NiS. Obtained as a black powder by
adding ammonium sulphide to an alkaline solution of a nickel
salt. It is the group-reagent precipitate; it also occurs in na-
ture as Millerite in rhombohedral or capillary crystals.
307. Tests for Nickel. — 1. Nickel compounds in the
292 COBALT.
oxidizing-f\.'<ime ,give the borax bead a brownish-red color
when hot, yellow when cold. In the reducing-ft&me the
bead assumes a grayish color owing to the reduction of
metallic nickel. The presence of cobalt may obscure this
test. See
2. In the wet way the solid is dissolved in water or aqua
regia. This solution is then tested for nickel thus : —
. (a) With a drop or two of ammonia the apple-green
hydroxide, Ni(OH)2, is thrown down ; but if ammonia be
added to excess, the hydroxide dissolves, forming a blue
solution. Again add potassium hydroxide to this blue so-
lution, and the apple-green precipitate again appears.
(5) Add potassium or sodium hydroxide to the original
solution, — an apple-green precipitate. In general, all the
salts of nickel are greenish.
EXERCISES.
1. Dissolve a nickel three-cent piece in nitric acid, and determine what
metals the coin contains.
2. Heat a nickel salt with Na2C03 on charcoal. Does a magnet attract
the powder obtained ? What other substance (metallic oxide) is thus
attracted ?
3. Read R. & S., Vol. II., Pt. II., pp. 146-149, on the alloys of nickel.
Prepare M(N03)2 and NiS04, and describe the processes, etc., as ex-
periments.
COBALT.
SYMBOL, Co". — ATOMIC WEIGHT, 59. — SPECIFIC HEAT,
0.10674. — MELTING-POINT, A WHITE HEAT.
308. Occurrence. — Cobalt does not occur free, and its
ores are neither plentiful nor widely distributed. Some of
its ores are Speiss Cobalt, Co(Ni,Fe)As2 ; Skutterrudite,
CoAs3; and Cobalt-glance, CoFeAs2S2.
COBALT. 293
309. Preparation. — Metallic cobalt is of little or no
use in the arts, and is prepared in small quantities only
for scientific purposes. The metal is reduced by strongly
heating the oxide or chloride of cobalt in an atmosphere
of hydrogen, when the cobalt is obtained as a gray, metal-
lic powder.
The metal can also be obtained in a coherent state by
fusing its oxalate under a layer of powdered glass and
afterwards fusing the metal in a graphite crucible.
310. Properties and Compounds of Cobalt. — Coherent
metallic cobalt resembles iron in its whitish color and in
being attracted by the magnet. It oxidizes but slowly
when in the coherent condition, but when in the form of
a powder it oxidizes quickly at ordinary temperatures.
The compounds of cobalt are valuable, and are prepared
directly from the arsenical ores, preferably speiss cobalt,
Co(Ni,Fe)As2.
The ore is roasted to vaporize the arsenic, then fused
with lime and sand to remove the iron. The residue is
now dissolved in hydrochloric acid, and any remaining
impurities are precipitated by adding successively chlorine
calcium carbonate, and hydrogen sulphide. The remaining
solution is thus freed from all the first and second group
metals, the cobalt only remaining in solution.
The oxide of cobalt is obtained from this solution by
the aid of bleaching-powder. The oxide is an article of
commerce, and is used for coloring glass blue, and for pre-
paring the salts and compounds of cobalt, which are of
great value as pigments.
There are two series of cobalt salts, — the cobaltous and
cobaltic. In the latter series, the salts are of a varying
formula, and not of sufficient importance to warrant a
294 COBALT.
notice here. The cobalt salts are violet when anhydrous,
and pink in color when hydrated ; hence they are used
to make "sympathetic ink," which becomes visible upon
warming.
Cobalt or its oxides, CoO and Co20s, are soluble in dilute
THE PRINCIPAL COBALT SALTS AKE : —
(a) Cobaltous Chloride, CoCl2. Prepared by dissolving the
metal or its carbonate in hydrochloric acid ; it is used as a
sympathetic ink.
(b) Cobaltous Nitrate, Co(NO3)2. Prepared by dissolving
the metal or its carbonate in dilute nitric acid ^ it is used as a
reagent.
t(c) Cobaltous Sulphide, CoS, the group-reagent precipitate.
A black precipitate obtained by adding ammonium sulphide to
a solution of a cobalt salt.
(d) /Silicates of Cobalt. These are prepared artificially, and
are known as " Smalt." The ore is first roasted sufficiently to
oxidize the cobalt, and then fused with quartz and potash,
when a dark-blue glass is formed, which is crushed into dust,
under water, by granite millstones, and is sold as a pigment.
(e) Rinmann's Green. A pigment prepared by precipitating
a solution of zinc and cobalt sulphates by sodium hydroxide.
This precipitate is then heated and reduced to an impalpable
powder. Its formula is unknown.
(/) Thenard's Blue, or Cobalt Ultramarine. A valuable
pigment prepared by heating alumina with a cobalt salt. Its
formula varies.
311. Tests for Cobalt. — 1. Any cobalt compound
colors the borax or microcosmic bead blue, — often appear-
ing black when the cobalt is in excess ; when powdered,
the dust obtained from the bead is blue in all cases.
MANGANESE. 295
NOTE. Should iron or nickel compounds be present, they may be
reduced to a colorless condition (metallic) by the continued application of
the reducing-flame, so that they will not interfere with this test.
312. To separate and identify Nickel and Cobalt. —
To a somewhat concentrated solution of the salts of these
two metals add acetic acid and potassium nitrite (KNO2).
Warm the solution gently for some time, and allow it to
stand for about twenty-four hours. At the end of this
time a yellow, crystalline precipitate of potassium-cobaltic
nitrite will settle.
(a) Obtain the nickel test from the solution. (Art. 307.)
(6) Apply the cobalt test (Art. 311) to the precipi-
tate.
SUG. Read the " Chemical News " for April 10, 1885, p. 170, for a new
method of separating nickel and cobalt.
MANGANESE.
SYMBOL, MN". — ATOMIC WEIGHT, 55. — SPECIFIC HEAT,
0.1217. — MELTING-POINT, A WHITE HEAT.
313. Occurrence. — Manganese never occurs free. Its
chief ore is Pyrolnsite, MnO2; it also occurs in Braun-
ite, Mn2O3, Hausmannite, Mn3O4, Rhodocrosite, MnCO3,
and Manganite, Mn2O2(OH)2.
314. Preparation. — Metallic manganese is not em-
ployed for practical purposes. It is obtained by fusing,
at a white heat, a mixture of any one of its oxides and
charcoal in a closed crucible lined with graphite.
315. Properties and Compounds of Manganese. —
Manganese is a reddish-white, brittle metal, oxidizing so
easily in the air that it must be kept under naphtha or
coal-oil.
296 MANGANESE.
THE IMPORTANT COMPOUNDS OF MANGANESE ARE : —
(a) The Oxides. — Manganous Oxide, MnO. A grayish-
green powder.
Mangano so -Manganic Oxide, or Red, Oxide of Manganese,
Mn304. This substance crystallizes in acute, quadratic pyra-
mids.
Manganic Oxide, Mn2O3. This oxide is brownish-black, and
crystallizes in obtuse quadratic pyramids. As it occurs in
nature it is known as Braunite.
Manganese Dioxide, or Black Oxide of Manganese, MnO2.
This is the most important of the manganese oxides. It is used
in the laboratory in many ways, but its principal use is for pre-
paring chlorine gas, thus : —
Mn02 + 4 HC1 = MnCl2 + 2 H2O + C12.
The chloride, MnCl2, is not a waste product, since it may be
again converted into the dioxide : —
(1) MnCl2 + CaCO3 = MnCO3 + CaCl, ;
(2) MnCO3 + O (heated in a blast of hot air) = MnO2 + CO2.
This illustrates a process employed in generating chlorine for
the manufacture of bleaching-powder.
Manganese Heptoxide, Mn2O7, is a dark, reddish-brown liquid,
which yields, with water, permanganic acid : Mn2O7 + H2O =
2 HMnO4.
The other oxides of manganese are basic oxides ; of these
MnO is the strongest base.
(6) Manganic Acid, H2MnO4, and Permanganic Acid,
HMnO4. The first is a very unstable acid not known in the
free state ; its salts, the manganates, are green in color, and
very unstable, except in the presence of an excess of alkali.
Permanganic acid is prepared thus : —
Ba(MnO4)2 + H2SO4 = 2 HMnO4 + BaSO4.
This acid in aqueous solution is a deep-red liquid possessing
MANGANESE. 297
a bitter, metallic taste ; it is readily decomposed by heat or
exposure to light. Of its salts the principal one is
Potassium Permanganate, KMnO4 or K2Mn2O8. This is a
crystalline substance, the color of which varies through green,
black, and steel-blue, depending upon the age and exposure of
the crystals. The uses of the permanganate are numerous,
with many of which the student is already acquainted. Organic
substances, as in drinking-water, reduce permanganates to lower
compounds, partially manganates.
Condy's Disinfecting Liquid is obtained by dissolving the
permanganate in water. The commercial article, however, is
not a pure permanganate, and is prepared on the large scale
by heating to redness for several hours caustic soda with man-
ganese dioxide ; the fused substance is then lixiviated with
water, and the solution is afterwards concentrated, when it is
ready for the market.
EXP. 163 T. Chameleon Mineral is a remarkable compound
which may be readily prepared as follows : Fuse in a crucible
equal weights of solid potassium hydroxide and finely levi-
gated manganese dioxide. Fill a tall jar with pure water, and
slowly drop in the powdered and cooled mass formed by fusion.
Note the colors obtained as the fine particles find their way to
the bottom of the vessel.
NOTE. The chemical changes which take place in this experiment are,
first, the formation of a salt of the composition K2MnO4, potassium man-
ganate,
3 MnO2 + 2 KOH = K2MnO4 + Mn203 + H20,
the solution of which is green. This salt is unstable unless free alkali is
present. "When poured into water it is converted into the permanganate,
K2Mn208 (or KMn04), the solution of which has a beautiful purplish-red
color. Hence, in the above experiments, the color changes from green to
purplish-red, and various intermediate colors are observed.
(c) Manganese Sulphide, MnS, is the flesh-colored group-re-
agent precipitate obtained by ammonium sulphide in an alkaline
solution of any salt of manganese. It also occurs as the min-
298 ZINC.
eral Alabandite, or Manganese Blende, in cubical or octahedral,
steel-gray crystals.
316. Tests for Manganese. — 1. To the borax and
microcosmic beads in the oxidizing-flame manganese gives
a violet color when hot, amethyst-red when cold.
In the reducing-flame the bead becomes colorless.
2. When fused on platinum foil with Na2CO3 and
KNO3, manganese compounds give a bright-green mass,
(what?). Dissolve this mass in water and add HNO3; a
red solution is formed.
3. The manganese acids may be distinguished by the
color of their salts in solution, and by further yielding
the reactions in 1 and 2.
What substances bleach a permanganate ?
ZINC.
SYMBOL, ZN". — ATOMIC WEIGHT, 65. — SPECIFIC HEAT,
0.0955.— MELTING-POINT, 423° C.
317. Occurrence. — Zinc seldom or never occurs native.
Its chief ore is Smithsonite, or ZnCO3. Franklinite,
(Zn,Fe)O + Fe2O3; Zinc blende, ZnS; Willemite, Zn2SiO4,
and a red oxide which owes its color to a reddish oxide
of manganese, are the ores chiefly employed in the reduc-
tion of zinc in the United States.
318. Preparation. — The ores of zinc are first roasted
and ground fine, then mixed with coal-dust to the amount
of one-half their weight. This mixture is then placed in
clay retorts, and heated until the zinc is reduced and
vaporized, when the escaping vapors are condensed in iron
ZINC. 299
condensers. Zinc thus prepared is the commercial article,
and is seldom pure, since it contains small quantities of
carbon and iron, lead, arsenic, antimony, and other metals.
319. Properties, Uses, and Salts of Zinc. — Zinc is a
malleable, ductile, bluish-white metal which is used for
many purposes. Its uses in our dwellings are familiar to
all. It is also used, when alloyed with copper to form
brass, in ways innumerable. In the laboratory, zinc is
used in batteries, in preparing hydrogen, in desilvering
lead, and in reducing other metals from their solutions.
Sheet iron, when covered with a coating of zinc, is said to
be galvanized.
SUG. Let the student name the metals which he can obtain frem solu-
tions of their salts by suspending a strip of zinc therein.
In the form of dust, zinc is used in chemistry as a reduc-
ing agent. Zinc-dust burns in the Bunsen flame with a
white bluish light.
Pare zinc dissolves very slowly in sulphuric acid, hence
it is well to add a small quantity of platinum chloride to
the granulated zinc employed by the student. Cover the
zinc with water, pour in the platinum chloride, and a coat-
ing of platinum black is soon deposited on the zinc. Now,
when the sulphuric acid is poured in, a galvanic current is
established, and the zinc readily dissolves. A solution of
pure copper sulphate answers the same purpose.
Zinc salts, when taken internally, are poisonous. Canned
goods may become poisonous when the tin cans are sol-
dered by the aid of zinc chloride, which is a soldering fluid
often employed by tinners.
Sue. Eead the " Chemical News/' June 5, 1885, p. 268, for valuable
information concerning poisoned canned goods.
300 ZINC.
THE PRINCIPAL COMPOUNDS OF ZINC ARE: —
(a) Zinc White, ZnO. Used as a paint.
(6) Zinc Chloride, ZnCl2. Used as a caustic in surgery,
and in organic chemistry for removing the elements of water
from many substances. It is also used in " weighting" cotton
goods.
(c) Zinc Sulphate, or White Vitriol, ZnSO4 + 7 H2O. Used
in medicine and in dyeing.
(d) Zinc Sulphide, ZnS, the group-reagent precipitate (white).
320. Tests for Zinc. — 1. Unknown solids are tested
for zinc by the blowpipe. When heated in the oxidizing
flame on charcoal, zinc compounds with Na2CO3 give a
coating around the assay, which is yellow when hot, white
when cold.
2. If -after being treated as in 1, the mass be moistened
with cobaltous nitrate and heated again, it turns green.
This color is a beautiful one, known as Rin man's green.
3. Solutions are first made alkaline, and ZnS (white) is
precipitated with (NH4)2S. This sulphide is insoluble in
dilute acetic acid, and is further tested by 1 and 2.
321. To separate and identify Nickel, Cobalt, Man-
ganese, and Zinc. — Obtain the precipitates NiS, CoS, MnS,
and ZnS, as directed in Art. 290. Warm the test-tube
containing the precipitate until the sulphides settle ; then
filter out and wash this precipitate and wash it through
into a test-tube ; dissolve as much of it as possible with
cold dilute HC1. Any residue may be NiS or CoS, or
both; filter and test as in (a). The filtrate is tested by
(5) for manganese and zinc.
(a) This residue is always black when containing Ni or
REACTIONS IN GROUP III. 301
Co, or both metals. Test it by the borax bead. (See
Arts. 307, 1 and 2, and 311, 1, note.) It is well, also, to
proceed by 312 to make sure whether both are present.
NOTE. Some free sulphur usually remains on the paper with this
residue. Whence came it ?
(6) Boil to expel H2S, and add a decided excess of KOH
to the hydrochloric acid solution in a test-tube. Allow it
to stand some time, and shake it frequently. If manga-
nese be present, it will be precipitated as Mn(OH)2, when
it must be filtered out, and tested by Art. 316, 2.
NOTE. Test this precipitate also for Ni and Co, since the HC1 is apt to
dissolve small portions of their sulphides.
For detecting the zinc, acidify with acetic acid the fil-
trate just obtained, and add (NH4)2S, — a white precipitate,
ZnS. Also test by Art. 320, 1 and 2.
GENERAL NOTE. Phosphoric acid or phosphates, when present in third
and fourth group solutions, cause them both to be precipitated with
(NH4)2S. For separation of third and fourth group metals in presence of
phosphates, see Douglas and Prescott's " Qual. Anal./' p. 241.
SOME REACTIONS IN GROUP III.
Balance and explain these equations : —
(1) Fe + H2SO4 = FeS04 + H.
(2) Fe + (cold dil.) HN03 = Fe(N03)2 + NH4N03 + H2O.
(3) Fe + (hot dil.) HN03 = Fe2(NO3)6 + NO + H20.
(4) Fe + HC1 = FeCl2 + H.
(5) FeCl2 + Cl = Fe2Cl6.
(6) Fe +HN03 (cold, very dil.) = Fe(N03)2 + H.
(7) Fe2Cl6 + H2S = FeCl2+HCl + S.
(8) FeS04 + (NH4)2S = FeS + (NH4)2S04.
(9) Fe2Cl6 + KOH = Fe,(OH)fi + KC1.
(10) A1+H2SO4=A12(SO4)3+H.
302 THE RARE METALS OF GROUP III.
(11) Al + HC1 = A12C16 + H.
(12) A12C16 + (NH4)2S + H20 = Al2(OH)fi + NH4C1 + H2S.
(13) K2Cr207 + (NH4)2S2 + H2O = Cr2(OH)6 + K2S03 + NH3 + S.
(14) Mn02 + HC1 = MnCl2 + Cl + H20.
(15) Mn02 + H2S04 + NaCl = MnS04 + Na2S04 + C12 + H2O.
(16) MnS04 + (NH4)2S = MnS + (NH4)2SO4.
(17) Co+HN03 = Co(N03)2+H20 + NO.
(18) CoN03 + (NH4)2S = CoS + (NH4)N03.
(19) CoCl2 + KNQ2 + HC2H302 + H2O = (KNO2)fi, Co20(N02)4>
H20 + KC1 + NO + KC2H302.
(20) Ni + HC1 = NiCl2 + H.
(21) M+HN03 = Ni(N03)2+H20 + NO.
(22) NiCl2 + (NH4)2S = NiS + NH4C1.
(23) NiCl2 + KOH = Ni(OH)2 + KC1.
(24) Zn+H2S04=ZnS04+H.
(25) Zn + HC1 = ZnCl2 + H.
(26) ZnS04 + (NH4)2S = ZnS + (NH4)2S04
(27) ZnCl2 + KOH = Zn(OH)2 + KCl.
QUERY. Are all of the metals of the third group precipitated with
ammonium sulphide ?
Sue. Separate and identify Ni and Co thus : To a solution of their salts
add KCy until the precipitate at first formed dissolves ; slightly acidulate
with HC1, and boil some time ; a little HC1 now precipitates NiCy2.
THE RARE METALS OF GROUP III.
BERYLLIUM.
SYMBOL, BE. — ATOMIC WEIGHT, 9.
322. Beryllium, also known as Glucinum, is a silver-white
metal occurring in Beryl, Be3Al2Si6O18.
It is prepared by fusing BeCl2 with metallic sodium or
potassium.
The salts of beryllium have a sweetish taste, from which
fact the metal first received the name glucinum.
INDIUM. — GALLIUM. 308
In the regular course of analysis beryllium is obtained along
with aluminum, from which metal it can be separated by pre-
cipitation with (NH4)2CO3.
INDIUM.
SYMBOL, IN. — ATOMIC WEIGHT, 113.6.
323. Indium occurs in zinc blende, and was discovered by
means of the spectroscope. It is a soft, white metal which
scarcely undergoes any change in the air. It is prepared from
its ores in the wet way.
Indium is detected by moistening its compounds with hydro-
chloric acid ; then it is placed in the non-luminous Bunsen flame
by means of a looped platinum wire. It colors the flame blue.
GALLIUM.
SYMBOL, GA. — ATOMIC WEIGHT, 69.
324. This metal also occurs in zinc blende, and was dis-
covered by the spectroscope. Its more prominent properties
were predicted previous to its discovery by Mendelejeff, under
the name of " Ekaluminum." (See p. 222.) It is prepared in
the wet way, and is a bluish-white metal, oxidizing readily in
the air, and melting at the extremely low temperature of 30/L° C.
It is detected by the spectroscope. Its luminous spectrum
contains two violet lines.
NOTE. The luminous spectrum is obtained by igniting a substance on
platinum wire in the Bunsen flame, or by means of a powerful cur-
rent of electricity, and exposing the flame directly to the spectroscope.
The absorption spectrum is obtained by igniting as above, and placing a
luminous gas-flame or other absorbing material between the burning metal
and the spectroscope. The spark spectrum is obtained by moistening the
carbon terminals of a dynamo or other powerful electric machine with a
solution of the substance to be tested, after which sparks are allowed to
pass.
304 YTTRIUM. LANTHANUM. CERIUM. DJDYMIUM.
YTTRIUM.
SYMBOL, YT. — ATOMIC WEIGHT, 89.
325. Yttrium occurs along with erbium.
It is detected by the spark spectrum of its chloride, which
gives many bright lines, of which the most marked are two
groups near the sodium line.
LANTHANUM.
SYMBOL, LA. — ATOMIC WEIGHT, 138.2.
326. Lanthanum occurs in the mineral Lanthanite as La2(CO3)3
+ 8 H2O. It is best prepared by the electrolysis of its chloride,
and is a soft grayish metal which readily tarnishes in the air,
assuming a steel-blue tint. It is detected by its spark spectrum
containing many characteristic lines.
•
CERIUM.
SYMBOL, CE. — ATOMIC WEIGHT, 141.
327. Cerium occurs along with Lanthanum, and is similarly
prepared. It is a soft, gray metal which tarnishes in damp air,
assuming, successively, the colors yellow, blue, and green.
It burns with great brilliancy when heated in the air, and
is detected by its spark spectrum which contains three bright
lines in the green.
DIDYMIUM.
SYMBOL, Di. — ATOMIC WEIGHT, 142.3.
328. This metal occurs along with the rare metals previously
mentioned, and is prepared similarly to Lanthanum. It has a
yellowish lustre, and burns brightly when heated in the air.
It is detected by its absorption spectrum. Its salts have a
rosy tint, and it colors the microcosmic bead rose-red.
TEKBIUM. EKBIUM. — THOIIIUM. — TITANIUM. 305
TERBIUM.
SYMBOL, TB. — ATOMIC A^EIGHT, 148.5.
329. This metal has not been prepared, but its oxide, Tb2O3,
is an orange-yellowish powder. It is difficult to separate ter-
bium from the preceding kindred metals, and no sure means of
detection is known, since it gives no absorption spectrum.
ERBIUM.
SYMBOL, ER. — ATOMIC WEIGHT, 166.
330. This metal occurs with the foregoing, and has not been
obtained pure. It is detected by its continuous luminous spec-
trum, which is crossed by bright lines which are darkened in
the same position in the absorption spectrum.
THORIUM.
SYMBOL, TH. — ATOMIC WEIGHT, 232.
331. Thorium occurs in Thorite and other complex minerals,
and is prepared by heating its chloride with potassium. This
metal as thus prepared is a gray powder which burns brightly
in the air.
Thorium is detected by the precipitation of its carbonate or
hydroxide ; these are soluble in an excess of the precipitant.
TITANIUM.
SYMBOL, Ti. — ATOMIC WEIGHT, 48.
332. This metal occurs in Rutile and in Titanite, TiCaSiO5,
and other minerals. It forms a considerable per cent of some
of the Lake Superior iron ores. Titanium is prepared by heating
a double fluoride of potassium and titanium in a closed crucible
with metallic potassium; the fused mass is then lixiviated with
water, when the titanium remains as a dark-gray powder.
306 ZIRCONIUM. — URANIUM.
At a high temperature this metal unites directly with nitro-
gen,— a marked peculiarity; it also burns when heated in
the air. In blast furnaces, when reducing iron ore containing
titanium, a peculiar compound, Titanium Cyano-nitride, TiCy2
-f- 3 Ti3N2, is obtained.
Titanium is detected by imparting to the microcosmic bead in
the reducing-flame a yellow color when hot, violet when cold ;
when iron is present the bead is red. The oxidizing-flame
gives no color.
ZIRCONIUM.
SYMBOL, ZR. — ATOMIC WEIGHT, 90.
333. Zirconium occurs in the mineral Zircon, ZrSiO4, and is
prepared in the same wa}~ as titanium, which metal it strongly
resembles. The amorphous form burns easily, but a crystalline
variety takes fire in the air only at the highest temperatures.
Zirconium is detected by precipitating its sulphate by K2SO4,
which gives a basic salt insoluble in water and hydrochloric
acid. Its spectrum is characteristic.
URANIUM.
SYMBOL, U. — ATOMIC WEIGHT, 239.8.
334. Uranium occurs in pitch blende, U3O8, and is prepared
in the wet way, or by fusing its chloride with potassium.
This is a hard, grayish-white metal, which also burns in the
air.
The black oxide, U2O5, is used for painting on porcelain.
The uranium salts are fluorescent, and impart this property to
" canary" glass.
Uranium is detected by its giving to the microcosmic bead
in the oxidizing-flame a yellow color when hot, green when
cold ; when farther heated the color is darkened.
The spectrum of uranium is distinctive.
TANTALUM. NIOBIUM. VANADIUM. 807
TANTALUM.
SYMBOL, TA. — ATOMIC WEIGHT, 182.
335. Tantalum occurs together with many of the rare metals
previously noticed, but more especially with niobium, from
which metal it has not been separated.
Tantalite, Columbite, Pyrochlor, Yttrotantalite, Pitch Blende,
and many other minerals contain small quantities of this metal.
Tantalum has not been obtained pure.
It is detected by converting the compound into tantalic acid,
and adding potassium f errocyanide to its solution ; this yields a
yellow precipitate. The conversion is effected by heating the
compound with carbon in a current of chlorine to obtain the
chloride TaCl5 ; this chloride, when mixed with water, yields
the acid HTaO3. A solution of nut-galls gives a yellow precipi-
tate with solutions of this acid.
NIOBIUM.
SYMBOL, Nb. — ATOMIC WEIGHT, 94.
336. Niobium occurs with Tantalum, and is prepared by
passing the vapor of its chloride and hydrogen through a red-
hot porcelain tube. It is a steel-gray metal, burning easily in
the air.
Niobium is detected similarly to tantalum, the precipitate with
K4FeCy6 being brown ; with nut-galls solution, orange-red.
.
VANADIUM.
SYMBOL, V. — ATOMIC WEIGHT, 51.5.
337. This metal occurs in Vanadanite, 3 Pb3(VO4)2+ PbCl2,
and is prepared as a grayish powder by heating its chloride in
hydrogen.
308 VANADIUM.
Vanadium brdnze, or metavanadic acid, is now used in place
of gold bronze for gilding.
Vanadium is detected by placing a strip of zinc in a solution
of vanadium chloride ; the solution turns blue. When hydrogen
dioxide and ether are added to the solution of a vanadate, the
solution turns red.
GENERAL NOTE. Observe those formulae like Co(Ni, Fe)As.2; these
do. not signify that both Ni and Fe are present, but that one or the other
is found in such a compound.
CHAPTER XVIII.
THE FOURTH GROUP METALS.
338. The fourth group metals are commonly known
as the Metals of the Alkaline Earths.
Their chlorides, hydroxides, and sulphides are soluble in
water, acids, and alkalies. In the course of analysis they
are precipitated as carbonates by ammonium carbonate,
(NH4)2CO3, in the presence of ammonia and ammonium
chloride. We must except magnesium, however, from
the above statement, since its carbonate is soluble in
ammonium compounds. It is best to filter out the precipi-
tates obtained by ammonium carbonate, and to precipitate
the magnesium from the filtrate by means of di-sodiuin
phosphate, Na2HPO4.
THE FOURTH GROUP METALS ARE: —
f Barium, Ba.
DIVISION A 5 Strontium, Sr. DIVISION B \ Magnesium, Mg.
' Calcium, Ca.
These metals oxidize easily in the air, and consequently
never occur free ; they are strongly basic, hence they are
not easily reduced to a metallic state ; they form no acids ;
they decompose water to form alkaline hydroxides.
810 BARIUM.
BARIUM.
SYMBOL, BA". — ATOMIC WEIGHT, 137. — SPECIFIC HEAT, .
MELTING-POINT, HIGHER THAN CAST IRON.
339. Occurrence. — The most abundant ore of this
metal is Heavy Spar, BaSO4. Barium also occurs in small
quantities in Witherite, or BaCO3, in certain silicates in
feldspathic rocks, in seaweeds, and in mineral waters.
340. Preparation. — Barium amalgam is prepared by
electrolyzing a thick paste of BaCl2 and dilute HC1 in the
presence of mercury. This amalgam is then heated to
vaporize the mercury, thus leaving a porous mass of
metallic barium. Barium oxidizes rapidly in the air, and
burns with great brilliancy.
341. Compounds and Uses of Barium. — Metallic
barium is not used in the arts.
ITS PKINCIPAL COMPOUNDS ARE: —
(a) Barium Monoxide, or Baryta, BaO, which is prepared by
heating the nitrate until nitrous fumes cease escaping.
Barium Hydroxide, or Caustic Baryta, Ba(OH)2, is obtained
by moistening BaO with water ; a solution of this hydroxide
is used as a reagent known as Baryta Water. Caustic baryta
is now largely used in refining cane sugar, which it precipitates
from its impure solutions as C12H22OnBaO. The barium is after-
wards removed by treatment with carbon dioxide gas, which
precipitates the insoluble compound, BaCO3, while the sugar
dissolves.
Barium hydroxide is now prepared in large quantities by
passing moist carbon dioxide gas through heated barium sul-
BARIUM. 311
phicle, which gives BaCO3 ; this carbonate is then treated with
superheated steam, when this reaction occurs : —
BaCO3 + H2O = Ba(OH)2 + CO2.
(6) Barium Chloride, BaCL. This salt is used as a reagent
to detect and estimate sulphuric acid ; it is prepared by dissolv-
ing barium carbonate, BaCO3, in hydrochloric acid. Write the
equation.
(c) Barium lodate, Ba(IO3)2, which is used to prepare iodic
acid, HIO3 ; this iodate is prepared thus : —
BaCl2 + 2 KIO3 = Ba(IO3)2 + 2 KC1.
(d) Barium Sulphate, or Heavy /Spar, BaSO4. This mineral
is an important barium ore, used for weighting paper and as a
paint. It is prepared for commerce thus : —
BaCl2 + H2SO4 = BaSO4 + 2 HC1.
(e) Barium Nitrate, Ba(NO3)2. This is prepared thus : —
BaCO3 + 2 ELNO3 = Ba(NO8)2 + H2O + CO2.
It is used in making green fires for tableaux and pyrotechnics.
(/) Barium Carbonate, BaCO3, which occurs in nature as
Witherite; it is also the group-reagent precipitate, prepared by
precipitating a barium salt by means of an alkaline carbonate.
It is largel}' used to prepare soluble barium salts.
342. Tests for Barium. — 1. Solids are fused with
sodium carbonate, if necessary, and then dissolved in
hydrochloric or nitric acid ; this solution gives these pre-
cipitates : —
(a) With K2Cr2O7 and ammonia, a yellow precipitate,
BaCrO4, insoluble in acetic acid.
(5) With H2SO4, a white precipitate, BaSO4, insoluble
in acids.
(c) CaSO4 gives an immediate precipitate of BaSO4
even in dilute solutions.
012 STRONTIUM.
2. Barium sarlts tinge the non-luminous flame green.
3. The barium spectrum, although complicated, is
readily distinguished by the green lines Baa and Ba/3.
STRONTIUM.
SYMBOL, SR". — ATOMIC WEIGHT, 87.2. — MELTING-POINT,
A RED HEAT.
343. Occurrence. — Strontium occurs most plentifully
in the two ores, Celestine, SrSO4, and Strontianite,
SrCO3. It also occurs in a few mineral waters and in
sea-water.
344. Preparation. — This metal is prepared by electro-
lyzing its chloride, or by heating this compound with a
sodium amalgam ; the strontium amalgam thus formed is
then washed, dried, and, finally, ignited in a current of
hydrogen.
345. Properties, Compounds, and Uses of Strontium.
— Strontium is a yellow, malleable metal, oxidizing in the
air, and burning brightly when heated.
THE PRINCIPAL STRONTIUM COMPOUNDS ARE : -
(a) Strontium Carbonate, SrCO3. This precipitate is ob-
tained b}r precipitating a strontium salt solution with an alkaline
carbonate.
(b) Strontium Nitrate, Sr(NO3)2. This is prepared thus : —
SrCO3 + 2 HNO3 = Sr(NO3)2 + H2O + CO2.
It is used in producing red fire for tableaux, etc. Material for
red fire is best produced by mixing about equal parts of finely
pulverized and thoroughly dried Sr(NO3)2 and KC1O3 with an
CALCIUM. 313
equal bulk of powdered shellac, or with one-fourth part flowers
of sulphur ; the shellac is preferable, as it gives off no suffo-
cating fumes of sulphur dioxide. Green fire is obtained simi-
larly, by using barium nitrate, Ba(NO3)2, in place of strontium
nitrate.
CAUTION. These ingredients must be powdered separately, and after-
wards mixed with a bone knife on paper, since any concussion may pro-
duce an explosion.
346. Tests for Strontium. — 1. Most strontium com-
pounds, when moistened with hydrochloric acid, impart a
beautiful crimson tint to the non-luminous flame. Sul-
phates should, be reduced to sulphides in the reducing-
flame and then moistened with HC1 before ignition.
NOTE. When both barium and strontium are present, the strontium
color appears when the substance is first brought into the flame. A cau-
tion, also, is needed here lest the student mistake the pale yellowish-red
flame of calcium for that of strontium. Compare the colors yielded by
the pure salts of these two metals.
2. The spectrum of strontium contains the prominent
lines : Sm, orange ; Sr/3, red ; and SrS, blue.
3. In the wet way, strontium when precipitated with
carbonates, phosphates, and oxalates, resembles barium.
It may be separated from barium by precipitating the latter
with ammonia and K2Cr2O7. It may be separated from
calcium by precipitating strontium with CaSO4.
CALCIUM.
SYMBOL, CA". — ATOMIC WEIGHT, 40. — SPECIFIC HEAT,
0.1804. — MELTING-POINT, A RED HEAT.
347. Occurrence. — The most abundant compound of
calcium is the carbonate, CaCO3. This mineral occurs in
enormous quantities and widely distributed ; uncrystal-
314 CALCIUM.
lized CaCO3 oc'curs as limestone and chalk ; the crystal-
lized forms are many, such as marble, Iceland Spar, Calc
Spar, and Dog-tooth Spar. Shells and corals are chiefly
carbonates of calcium, while bones and teeth are princi-
pally phosphates of this metal. Calcium Sulphate, CaSO4,
occurs in Gypsum, Anhydrite, and Selenite. Some moun-
tain ranges and geological formations are chiefly composed
of these calcium compounds.
348. Preparation. — This metal is prepared by electro-
lyzing its chloride, or by fusing calcium iodide with
metallic sodium in closed iron retorts.
349. Properties, Compounds, and Uses of Calcium. —
Calcium is a malleable metal, which oxidizes most rapidly
in moist air, and burns with an orange-yellow light.
THE 'MOST USEFUL COMPOUNDS OF CALCIUM AEE:-
(a) Quick-lime, CaO, prepared by heating the carbonate,
CaCO3. Give the equation.
Calcium Hydroxide, Ca(OH)2, which is prepared by treating
CaO with water. When this substance is in a dry powder or of
the consistency of paste, it is called "slaked lime." Why? A
saturated water solution of calcium hydroxide, called lime-
water, is used as a reagent for detecting free carbon dioxide
gas.
Slaked lime is used for many purposes, such as for making-
mortar, purifying illuminating gas, whitewashing, etc. Mortal-
consists of sand, three to four parts, and lime, one part,
thoroughly mixed with water.
SUG. Describe the method of making mortar. (Ask a mason or
plasterer, if you do not know.) What is "putty coat" or "hard finish" ?
Lime containing about ten per cent of silica is known as
hydraulic cement or water-lime, and possesses the peculiar
CALCIUM. 315
property of hardening under water. This cement is artificially
prepared by mixing finely pulverized burnt clay and limestone.
Calcium hydroxide absorbs carbonic acid gas from the air,
which fact explains the hardening of the mortar. It may also
combine with the silica.
QUERY. Does age improve the hardness of cement or mortar 1 Does
the cement of the ancient Roman masonry owe its stone-like character to
its age or to the process of manufacture *
(b) Gypsum, CaSO4 + 2 H2O. This occurs native, and
when ground is used as land plaster ; when calcined, it is
known as "Plaster of Paris," which is used in making casts
and for filling writing-paper.
QUERY. What is the object of the calcining 1 Explain the setting of
the plaster.
(c) Calcium Chloride, CaCl2. This substance is prepared by
dissolving Iceland spar or pure marble in hydrochloric acid.
When fused, it is used as a dryer for gases, owing to its great
absorptive power for moisture.
(d) Fluor Spar, CaF2, a well-known mineral used in prepar-
ing fluorine compounds.
(e) Bleaching Powder. This is an article of commerce, and
one of the most useful substances known to the arts. It is
made by passing chlorine gas into large chambers, on the floors
of which slaked lime is spread. It is used in bleaching paper,
rags, cotton goods, etc. This powder affords a convenient
source of chlorine, which is liberated by the addition of an acid,
as sulphuric or hydrochloric acid.
QUERY. Upon what does the bleaching power of chlorine depend ?
(/) Superphosphate of Lime is a substance obtained by treat-
ing bones with sulphuric acid; iJt is used in preparing phos-
phorus, and also as a fertilizer. The superphosphate is a mix-
ture of calcium sulphate and acid phosphate.
(g) Calcium Carbonate, CaCO3, previously mentioned under
316 MAGNESIUM.
the carbonates. .This substance forms one of the constituents
called " hardness" in drinking-water (see p. 49).
When a soap is brought into a hard water, insoluble calcium
salts are formed with the organic acids contained in the soap ;
hence the peculiar, unpleasant feeling experienced on attempt-
ing to wash the hands with soap in hard water. All the cal-
cium carbonate in solution must be precipitated before the
soap will act in the desired way and form a lather.
Iceland Spar, a beautiful crystalline variety, possesses the
property of " double refraction."
350. Tests for Calcium. — 1. The volatile calcium salts
tinge the flame orange-red.
2. The spectrum shows the green line Ca/3 and the
orange line Caa, which are distinctive.
3. In solutions, calcium may be separated from barium
and strontium by precipitating the latter metals with
K2S04;' to the filtrate ammonia and ammonium oxalate,
(NH4)2C2O4, are added ; the oxalate gives a white precipi-
tate, CaC2O4, which under the circumstances is distinctive.
QUERY. Is calcium sulphate easily soluble in water ? Try it.
MAGNESIUM.
SYMBOL, MG.'' — ATOMIC WEIGHT, 24. — SPECIFIC HEAT, 0.245.
MELTING-POINT, 750°.
351. Occurrence. — Magnesium ores are found plenti-
fully in many localities, among which we notice : Magne-
site, MgCO3; Dolomite, (Mg,Ca)CO3; Kieserite, MgSO4+
H2O; Carnallite,(Mg,K)Cl2 + 6H2O; Spinelle, MgOAl2O3;
Asbestos, (Mg,Ca)SiO3; Talc, Mg3H2(SiO3)4 ; and Meer-
schaum, Mg2H2(SiO3)3.
Magnesium sulphate also occurs in certain medicinal
springs, while the chloride is a constituent of sea-water.
MAGNESIUM. 817
Magnesium limestone is a double carbonate of calcium
and magnesium.
352. Preparation. — Magnesium, like calcium, may be
prepared by the electrolysis of its chloride, but the com-
mercial article is obtained by fusing a mixture of the dry
chloride, fluor spar, and metallic sodium in a closed cruci-
ble. The metal is afterward purified by distillation, and,
when in a semi-molten condition, it is pressed into wires,
which are flattened finally into ribbons.
353. Properties, Uses, and Compounds of Magnesium.
— Magnesium is a silver-white metal, quite permanent in
dry air; in damp air, however, its surface becomes coated
with oxide. It takes fire readily in any ordinary luminous
flame, and burns with a painfully bright and dazzling
light, which is very rich in chemical rays. Owing to this
important property, magnesium ribbon is now employed in
photographing caverns and other objects inaccessible to
the sun's rays. This metal is also employed in pyrotechny
and signaling. It is further employed in chemical analy-
sis, especially in cases of arsenic poisoning, in place of
zinc, since magnesium contains no traces of arsenic.
THE MOST IMPORTANT COMPOUNDS OF MAGNESIUM AKE
THE FOLLOWING: —
(a) Magnesia, MgO, which is prepared by igniting the car-
bonate, MgCO3. It is used in medicine.
(b) Magnesium Chloride, MgCl2, is obtained from sea-water
and salt springs. It is used in dressing cotton goods.
(c) Epsom Salts, MgSO4 -f- 7 H2O, are prepared from Kie-
serite, or by treating MgCO3 with sulphuric acid. It is used in
medicine as a cathartic, and is also used in dressing cotton
goods.
318 REACTIONS IN GROUP IV.
(d) Magnesium Carbonate, or Magnesite, MgCO3, an ore of
magnesium. This is artificially prepared by roasting dolomite,
and treating the moistened residue with carbon dioxide gas under
pressure ; a bicarbonate is thus formed, which is decomposed by
means of superheated steam. This compound as thus formed
is a white powder, which is an important article of commerce.
It is used in medicine ; also used as a face-powder.
354. Tests for Magnesium. — 1. After removing the
metals of the fourth group by ammonium carbonate, etc.,
di-sodium phosphate, Na2HPO4, when added to the filtrate,
throws down a white precipitate, MgNH4PO4; this forms
in a dilute solution after stirring the solution with a glass
rod for a few minutes. This precipitate, under the circum-
stances, is distinctive.
NOTE. The spectrum of magnesium is not a practical test, as it is not
very marked at the temperature of the Bunsen flame.
355. Separation and Identification of the Fourth
Group Metals. — 1. Make the solution to be tested neu-
tral or slightly alkaline, and then remove the metals of
Groups I., II., and III. by the usual methods.
Save the filtrate, and boil for some time to expel free
H2S; filter.
2. Add ammonia, NH4C1, and (NH4)2CO3 to precipitate
barium, strontium, and calcium. Filter out this precipi-
tate, and save it to test by 3 ; also save the filtrate, and
test it by 4 for magnesium.
3. Dissolve this precipitate in acetic acid.
(«) Test a small portion of the solution for barium
by adding K2Cr2O7 and ammonia; a yellow precipitate,
BaCrO4, indicates barium. If barium be present, thus re-
move it from the whole solution. This precipitate may
be filtered out and dissolved in hydrochloric acid ; then,
BE ACTIONS IN GROUP IV. 319
upon addition of H2SO4, the insoluble sulphate, BaSO4,
will confirm the test.
(&) Test a portion of the nitrate from (a) for calcium
by Art. 350, 3.
(<?) Precipitate the calcium and strontium from the ni-
trate not used in (b) by means of ammonia and ammonium
carbonate. Filter out the precipitate, and dissolve it in
HC1, and expel excess of acid; then add CaSO*. A white
precipitate, SrSO4, formed after a few minutes, indicates
strontium. Further test this precipitate by 346, 1.
4. To the nitrate from 2 add Na2HPO4, and stir for
some time with a clean glass rod, if necessary ; a white
precipitate, MgNH4PO4, indicates magnesium.
REACTIONS IN GROUP IV.
(1) CaCl2 + (NHJ2C03 = CaCOs + NH4C1.
(2) Sr(N03)2 + (NH4)2C03=_SrC05 + NH4N03.
(3) BaCl2 + (NHJ2C03 = BaC03 + NH4C1.
(4) MgS04 + Na2HPO4 = MgHP04 + Na2S04.
(5) CaC03 + H(C2H3O2) = Ca(C2H302)2 + H20 + C02
(6) SrC03 + H(C2H302) = Sr(C2H302)2 +
(7) BaC03 + H(C2H302) =
(8) Ca(C2H302)2 + (NHJ2C204 - CaC2O4 + (NHJ(C2H302).
(9) Ba(C2H3O2)2 + K2Cr207 + H20 + NH3 = BaCr04 + KC2H3O2
+ (NH4)2CrO4.
(10) Ba(02H302)2 + H2S04 = BaS04 + H(C2H3O2).
(11) BaCl2 + K2C03 = BaC03 + KC1.
(12) MgS04 + Na2C03 = MgC03 +
(13) CaC03 + HCl= +
SUG. The student should do much work with the preceding groups ;
the quickest way to become acquainted with a substance is to work with
it. Unknown solutions give an added zest to the student's desire for
mastering processes.
CHAPTER XIX.
THE FIFTH GROUP METALS.
356. The metals of the fifth group are known as the
" Metals of the Alkalies." They do not yield precipitates
with the usual reagents, since the compounds thus formed
are soluble ; but they are detected by the color which
their compounds impart to the non-luminous flames, or by
their spectra.
These metals are POTASSIUM and SODIUM, also the com-
pound AMMONIUM, NH4; the rare metals are LITHIUM,
RUBIDIUM, and CESIUM.
Of course ammonium is not to be considered a true
metal, but its compounds are alkaline, and it behaves
much like metals of this group. In distinction from the
other or " Fixed Alkalies," ammonium is termed the "Vol-
atile Alkali," since most of its salts are volatile.
The metals of this group form a natural series ; they
are all acted upon by the moisture of the air, and hence
they must be kept under naphtha ; all decompose water at
ordinary temperatures to form strongly alkaline hydrox-
ides; each one forms but one series of salts, many of
which are exceedingly stable and useful.
QUERIES. To what group do these metals belong in MendelejefPs
Table ? Which belong to the < ' series ? Does Na or K show the
even
more intense action when thrown upon the water ?
POTASSIUM. 321
POTASSIUM.
SYMBOL, K'. — ATOMIC WEIGHT, 39. — SPECIFIC HEAT,
0. 1655 (?). — MELTING-POINT, 62.5°.
357. Occurrence. — The potassium-bearing compounds
are widely distributed; they occur in mineral waters, sea-
waters, and all fruitful soils, and are utilized by plants and
animals. Sheep excrete, through the skin, potassium and
other compounds, termed " Fat " and Suint. These com-
pounds are of considerable commercial value ; they are
retained by the wool, of which, before washing, they con-
stitute nearly one-third part by weight.
Some potassium compounds are the following minerals :
Sylvite, KC1; Saltpetre, KNO8; Orthoclase, K2Al2(Si3O8)2;
Carnallite, (KMg)Cl3; and Alum, K2A12(SO4)4 + 24 H2O.
358. Preparation. — Acid potassium tartrate is first
heated in closed iron retorts ; in this way, a very intimate
mixture of potassium carbonate and carbon is obtained.
This mixture is then placed in iron tubes covered with
clay, which are afterwards placed in a furnace, and heated
to a white heat. Metallic potassium is given off in the
form of vapors, which are passed into shallow, box-like
condensers placed outside the furnace; in these con-
densers they are quickly cooled to a liquid state ; the
liquid potassium then flows out into vessels containing
rock oil. (See Fig. 20.)
Formerly frequent explosions occurred, owing to the
formation of a black substance, KCO ; but this trouble
is now obviated by the shallow condensers.
Sir Humphry Davy first prepared potassium by electro-
lyzing the moistened hydroxide. This marked a new era
822 POTASSIUM.
in chemistry, as the alkalies were previously supposed to
be elements ; and, moreover, with the discovery of potas-
sium, the discovery of other rare metals became possible.
QUERY. What rare metals are now prepared by the aid of metallic
potassium or sodium "?
359. Properties, Uses, and Compounds of Potassium.
— Potassium is a silver-white metal when first cut, but
soon afterward exposes a bluish surface. It is brittle at
0° C., and waxy at ordinary temperatures.
It ignites at a low heat, — often while being cut, —
and requires the utmost care while being handled; it
must be kept under rock oil or naphtha. It quickly
decomposes water, liberating hydrogen with such violence
that it frequently takes fire and explodes.
It dissolves in ammonia, forming a blue solution, from
which it may be again obtained unchanged. The princi-
pal use of metallic potassium, other than for class demon-
stration, is in preparing the rare metals, as previously
noticed.
THE PRINCIPAL POTASSIUM COMPOUNDS NOT HERETO-
FORE NOTICED ARE: —
(a) Potassium Hydroxide, or Caustic Potash, KOH. This is
prepared by treating potassium carbonate with slaked lime,
thus : —
K2CO3 -f-Ca(OH)2 = 2 KOH + CaCO3.
The aqueous solution thus prepared is evaporated to dryness,
fused, and marketed. In this condition, it is extensively used
as a lye. It is purified for reagent purposes by dissolving the
crude salt in alcohol, and, after evaporation, again fusing and
casting it into sticks. It is kept in air-tight bottles, since it has
a powerful attraction for carbon dioxide and moisture, and soon
POTASSIUM. 828
deliquesces ; neither must it be handled with the hands, since it
destroys the skin.
Sue. Leech some common wood ashes by passing water through them.
Examine the filtrate obtained.
(b) Potassium Chloride, KC1, occurs naturally as Sylvite,
and in many brines. It is used as a fertilizer and in preparing
other potassium salts.
(c) Potassium Bromide, KBr. This salt is obtained together
with bromate of potassium by dissolving bromine in potassium
hydroxide ; the bromate is afterwards decomposed by a gentle
heat. It is used in medicine as a sedative, and in the labora-
tory as a source of bromine for demonstration.
QUERY. How is Br prepared ?
(d) Potassium Iodide, KI, may be prepared in the same way
as the bromide. It is used extensively in medicine and for
other purposes ; in the laboratory it is a source of iodine for
purposes of demonstration and is a reagent.
SUG. The potassium salts will be found in the laboratory; let tho
student examine them, note the forms of the crystals, etc., and write a
description.
(e) Potassium Chlorate, KC1O3, is obtained by passing a cur-
rent of chlorine gas through a solution of caustic lime until
calcium chlorate, Ca(ClO3)2, is formed; potassium chloride is
then added with the following results : —
Ca(ClO3)2 + 2 KC1 = 2 KC1O3 + CaCl2.
The chlorate of potassium is obtained from this solution by
crystallization. This salt is used in medicine for inflammation
of the throat, and in the laboratory as a source of oxygen.
QUERY. How is oxygen obtained from KC1O3 ? How may potassium
chlorate be prepared from chlorine and potassium hydroxide ? How is
KC104 prepared ? (See Perchloric Acid.)
(/) Potassium Sulphate, K2SO4, occurs native, and is pre-
pared as a by-product in the manufacture of other potassium
324 POTASSIUM.
compounds, as the bichromate, etc. It is used in medicine as
a purgative ; it is further used in the manufacture of alum, and
in the laboratory as a reagent.
An acid sulphate, KHSO4, is obtained in manufacturing-
nitric acid.
SUG. Write the equation.
(g) /Saltpetre, or Nitre, KNO3, occurs as an incrustation on
the soil of some hot, dry climates, as in India and in Egypt, where
it is produced by the oxidation of nitrogenous organic substances
in contact with the potassium compounds contained in the soil.
It has recently been shown that the formation of nitrates which
takes place in the soil is caused by minute organisms or fer-
ments. The process is similar to the familiar fermentation of
sugar, which causes the formation of alcohol and carbon
dioxide.
It is artificially prepared by treating sodium nitrate, which
occurs native in immense deposits, with potassium chloride,
thus : -
NaNO3 + KC1 = KNO3 + NaCl ;
and also in the so-called u saltpetre plantations." These are
constructed by piling up refuse animal matter, mixed with wood
ashes and lime, and moistening with urine or stable drainings.
At intervals the outer layer is removed, and extracted with
water.
The term " saltpetre" is derived from the fact that this salt
was and is still obtained from certain oily or feldspathic rocks
by boiling the weathered rock with slaked lime and potash.
Saltpetre is used in the laboratory as a source of nitric acid
for demonstration, as an oxidizing agent (substances are fused
with KNO3 for this purpose), and in preparing cooling mix-
tures.
QUERY. How are freezing mixtures prepared ? Explain the philoso-
phy of the process.
POTASSIUM. 325
In domestic economy, it is used as a preservative of meat ;
but the most important purpose for which nitre is used is in
manufacturing gunpowder.
Gunpowder consists of an intimate mixture of nitre, sulphur,
and charcoal, in somewhat varying proportions. Sporting pow-
der consists of nitre, 78.99, sulphur, 9.84, and charcoal, 11.17
parts. The explosive force of gunpowder depends upon the
fact that it contains within itself the necessary amount of oxy-
gen for its own combustion, whereby large volumes of heated
gases (principally carbon dioxide and nitrogen) are liberated.
QUERY. What effect has the invention of gunpowder had on civiliza-
tion ? Give the philosophy of explosions in general.
(h) Potassium Carbonate, or Potash, K2CO3, is usually ob-
tained from wood ashes. The ashes are lixiviated or "leached,"
and the lye thus obtained is evaporated till the solution is satu-
rated, when impure crystals of the carbonate are deposited.
These crystals are purified by roasting in a reverberatory
furnace.
Other sources of potash are potassium sulphate, beet-root
ashes, and suint.
Potassium carbonate is used in preparing other salts, as
potassium cyanide, chromate, acetate, etc., and as a reagent.
An acid salt, KHCO3, is prepared by passing a current of
carbon dioxide gas through a solution of the normal carbonate.
(i) Potassium Cyanide, KCN or KCy, is an important com-
pound, used in the laboratory as a reducing agent ; also used in
photography, and as a solvent for silver sulphide or oxide.
It is prepared by heating the ferro-cyanide with the carbo-
nate to a red heat in iron crucibles, thus : —
K4Fe(CN)6 + K2CO3 = 5 KCN + KCNO + CO2 + Fe.
The chemically pure cyanide is prepared by passing hydrocyanic
acid gas into an alcoholic solution of potassium hydroxide.
(j) There are other potassium salts in which the metal is
326 SODIUM.
combined with organic acids, and some of which are used in the
laboratory.
The student will notice the tartrate, oxalate, and acetate.
360. Tests for Potassium. — 1. Potassium compounds,
on the platinum loop, color the Bunsen flame violet ; but
the presence of sodium obscures this test, hence it is neces-
sary to observe the flame through thick cobalt-blue glass,
which shuts off the sodium rays but transmits the potas-
sium color.
NOTE. Always thoroughly clean the wire before testing.
2. The spectrum furnishes two easily distinguished
lines, — Ka in the extreme red, arid K/9 in the violet.
3. Potassium salts, in concentrated solutions, and in the
absence of all non-alkaline bases, yield, with tartar ic acid,
a white, distinctive precipitate, KHC4H4O6, this is granular-
crystalline, and may be tested further by 1.
SODIUM.
SYMBOL, NA'. — ATOMIC WEIGHT, 23. — SPECIFIC HEAT, 0.2394.
— MELTING-POINT, 95.6°.
361. Occurrence. — The chief and most plentiful sodium
compound is common salt, sodium chloride, NaCl. Salt
occurs in sea-water, most mineral waters, and drinking
water, while traces of it are to be found in nearly all river
waters. In some localities in the United States — as at
Syracuse, N.Y., and the Saginaw Valley, Mich. — salt
water or brine is found in vast reservoirs at a considerable
depth below the surface of the earth. Wells are sunk in
such localities, and the brine is raised to the surface by
pumps, and utilized as a source of the salt used in com-
merce. Again, large beds of native salt or rock salt occur
in various localities.
SODIUM.
327
Another source of sodium is the native nitrate, NaNO3,
or Chili saltpetre, which occurs in beds in Chili and Peru.
A large tract of territory in the western United States
is known as the Alkali Plains, owing to the occurrence of
sodium compounds : the water and the very earth itself
are saturated with alkali to such an extent that but scant
vegetation grows, and, with the exception of one or two
species of worms, the waters of the lakes, although clear
as crystal, are uninhabited.
FIG. 20.
A is the iron tube retort coated with clay.
C is the condenser.
D is the cup containing rock oil.
In its distribution, sodium is the most persistent and
universal of all the metals ; indeed, it is nearly impossible
to find a compound that will not yield the sodium test.
362. Preparation. — Sodium is prepared precisely like
potassium, excepting that the carbonate and charcoal,
instead of the tartrate, are employed. It is somewhat
328 SODIUM.
more easily obtained, however, and no explosive compound
is formed.
Fig. 20 will give a good idea of the furnace employed
in obtaining metallic sodium and potassium. After the
condenser is filled with the metal, it is taken off and put
under rock oil, after which the metal is scratched off.
363. Properties, Uses, and Compounds of Sodium. —
Sodium is a light, silver-white metal which oxidizes readily
in damp air.
It does not act upon water with as much violence as
potassium, but it will take fire when thrown upon hot
water, starch paste, or wet paper.
QUERIES. "What purpose does the starch paste serve ? Explain the
phenomenon of sodium burning on hot water. What metals are obtained
by the aid of metallic sodium ?
SODIUM FORMS MANY USEFUL SALTS, OF WHICH WE
NOTICE THE FOLLOWING : —
(a) Sodium Hydroxide, or Caustic Soda, NaOH, is prepared
on the large scale by decomposing sodium carbonate with
slaked lime, thus : —
Ca(OH)2 + Na2CO3 = 2 NaOH + CaCO3.
The aqueous solution is then treated precisely in the same man-
ner as caustic potash. Caustic soda is also prepared in large
quantities from the red liquors from which the black crystals
obtained in the soda-ash process are deposited.
QUERIES. When metallic sodium acts on water, is NaOH obtained ?
Try it. How can you decide what this substance is ?
The principal use of caustic soda is in soap making. In the
laboratory it is a useful reagent.
(b) Sodium Chloride, or Common Salt, NaCl, is obtained
from various sources, as previously indicated. The strong
brine of the salt wells is evaporated in shallow tanks by the aid
SODIUM. 329
of steam until the salt crystals are deposited. Salt is obtained
from sea-water by allowing it to flow into large, shallow pans or
vats called " salterns," where it is evaporated through the agency
of the wind and sun.
SUG. Student mention the many uses of common salt.
(c) Sodium Nitrate or Chili Saltpetre, NaNO8, occurs in vast
deposits in Peru and Bolivia, and is now used as a source of
nitric acid and as a fertilizer.
QUERIES. What other use of NaN03 was mentioned above under
potassium ? What element is obtained from Chili saltpetre ?
(d) Acid Sodium Hyposulphite, NaHSO2, is obtained by
treating a solution of sodium hydrogen sulphite, NaHSO3, with
granulated zinc.
It is used by dyers and calico printers to reduce indigo, and
in the laboratory for estimating free oxygen quantitatively.
(e) Sodium Sulphate, Na2SO4, with some admixture of the
acid sulphate, NaHSO4, is prepared in the first stage in the
manufacture of soda or sodium carbonate. It is known as
u salt-cake."
(/) Sodium Thiosulphate, Na2S2O3 -f 5 H2O, is used as an
antichlor by paper manufacturers, and in the photographic proc-
ess for dissolving out the unaltered silver salts. It is prepared
by boiling caustic soda with sulphur, and then passing sulphur
dioxide gas until the yellow solution obtained is decolorized.
Its solvent action on silver salts is due to the formation of a
double salt of sodium and silver, NaAgS2O3 : —
Na^SA + AgCl = NaAgSA + NaCl.
(g) Sodium Hypophosphite, NaH2PO2, is prepared by adding
calcium hypophosphite to a solution of sodium carbonate. The
filtered solution is then evaporated in vacuo. It is used in
medicine.
(h) Disodium Phosphate, Na2HPO4, is used in medicine as a
mild cathartic, and in the laboratory as a reagent. It is pre-
pared by treating phosphoric acid with sodium carbonate.
330 SODIUM.
(i) Sodium Carbonate, Na2CO3, is the chief product of soda-
ash manufacture. Soda-ash is a mixture of the carbonate and
hydroxide. The normal carbonate is used as an indispensable
reagent in dry reactions in the laboratory.
The manufacture of soda-ash is a great industry by itself.
The English process is thus described by Roscoe : —
" This substance, known in commerce as soda-ash, is manu-
factured in England on an enormous scale, and used for glass
making, soap making, bleaching, and various other purposes in
the arts. Formerly it was prepared from barilla or the ashes of
sea-plants, but now it is wholly obtained from sea-salt by a
series of chemical decompositions and processes, which may be
divided into two stages : —
"1. Manufacture of sodium sulphate, or salt-cake, from
sodium chloride (common salt) ; called salt-cake process.
"2. Manufacture of sodium carbonate, or soda- ash, from
salt-cake ; called soda-ash process .
FIG. 21.
"1. Salt -Cake Process. — This process consists in the
decomposition of salt by means of sulphuric acid. This is
effected in a furnace called the Salt-Cake Furnace. Fig. 21
shows the section of such a furnace. This is drawn to a scale
from one actually in use. It consists of (1) a large covered
iron pan, a, placed in the centre of the furnace, and heated by
fire placed underneath ; and (2) two roasters or reverberatory
furnaces, dd, placed one at each end, and on the hearths of
which the salt is completely decomposed. The charge of half a
ton of salt is first placed in the iron pan, and then the requisite
SODIUM. 331
quantity of sulphuric acid allowed to run in upon it. Hydro-
chloric acid gas is evolved, and escapes through a flue, e, with
the products of combustion into towers or scrubbers filled with
coke or bricks moistened with a stream of water. The whole
of the acid vapors are thus condensed, and the smoke and
heated air pass up the chimney. By recent act of Parliament,
the alkali makers are compelled to condense at least 95 per
cent of the hydrochloric acid gas they produce ; and so perfectly
is this condensation as a rule carried out, that the escaping
gases do not cause a turbidity in a solution of silver nitrate,
proving the absence of even a trace of the acid gas. After the
mixture of salt and acid has been heated for some time in
the iron pan, and has become solid, it is raked on to the hearths
of the furnaces at each side of the decomposing pan, where the
flame and heated air of the fire complete the decomposition into
sodium sulphate and hydrochloric acid.
FIG. 22.
"2. /Soda-Ash Process. — This process consists (1) in the
preparation of sodium carbonate, and (2) in the separation and
purification of the same. The first chemical change which the
salt-cake undergoes in its passage to soda-ash is its reduction
to sulphide, by heating it with powdered coal or slack : —
Na2SO4 + C4 = Na2S + 4 CO.
The second decomposition is the conversion of the sodium
sulphide into sodium carbonate, by heating it with chalk or
limestone (calcium carbonate) : —
Na2S + CaCOs = Na*COg + CaS.
These two reactions are in practice carried on at once, a mixture
332 SODIUM.
often parts of salt-cake, ten parts of limestone, and seven and
a half parts of coal being heated in a reverberatory furnace
called the Balling Furnace (shown in section in Fig. 22) until
it fuses and the above decomposition is complete, when it is
raked out into iron wheelbarrows to cool. This process is gen-
erally termed the black-ash process, from the color of the fused
mass.
' ' The next operation consists in the separation of the sodium
carbonate from the insoluble calcium sulphide and other impuri-
ties. This is easily accomplished by lixiviation, or dissolving
the former salt out in water. On evaporating down the solu-
tion, for which the waste heat of the black-ash furnace is used,
the heated air passes over an iron pan (see 6, Fig. 22) contain-
ing the liquid. On calcining the residue, the soda-ash of com-
merce is obtained."
Ammonia Process. — Another process for converting sodium
chloride into sodium carbonate is now used extensively. It
consists in treating a solution of sodium chloride with ammonia
and carbon dioxide : —
NaCl + NH3 + H2O + CO2 = NaCl + NH4HCO3.
The acid ammonium carbonate acts upon the sodium chloride,
forming acid sodium carbonate, NaHCO3, which is difficultly
soluble and is deposited : —
NaCl + NH4HCO3 = NH4C1 + NaHCO3.
The acid carbonate is heated and thus converted into the neu-
tral salt : —
2 NaHCO3 = CO2 + H2O + Na2CO3 ;
and the carbon dioxide given off is used for the purpose of satu-
rating the ammonia contained in the original solution. The
ammonium chloride obtained in the second stage of the process
is decomposed either by lime, CaO, or magnesia, MgO, and the
ammonia thus recovered. This process is also known as
SODIUM. 383
the Solvay process, as its introduction is due to the exertions
of M. Solvay.
Soda Crystals, or /Sal Sodae, much used in softening hard
water, are obtained by dissolving soda-ash in water, and allow-
ing the crystals to deposit from a saturated solution. These
crystals possess the formula Na2CO3 +10 H2O.
Acid Sodium Carbonate, NaHCO3, can be obtained from
soda crystals by allowing them to be acted upon by CO2 gas.
This substance is known as Bicarbonate of Soda, and is
employed in medicine and for preparing effervescing drinks. In
domestic economy it is used as Saleratus and as an ingredient
of Baking Powder.
(/) Silicates. Glass is a silicate of calcium and either
sodium or potassium. Ordinary glass contains sodium. The
difficultly fusible Bohemian glass contains potassium. For
some purposes, lead is introduced instead of calcium. Glass
made in this way, having a high refractive power, is very use-
ful for optical purposes. Ordinary glass is made by melting
together quartz and quicklime or calcium carbonate and sodium
carbonate.
(k) Man}r other salts of sodium may be obtained in the
shops, and are very useful in preparing test solutions, especially
when the student is working for acids in the non-metals.
364. Tests for Sodium. — 1. Sodium compounds color
the noil-luminous flame intensely yellow, and this color is
obscured by the blue glass.
NOTE. Any substance, as dirt on the platinum wire, will give this test
for sodium. Therefore, clean the wire carefully, and convince yourself
that the color is not caused by the ordinary impurities. Try some known
sodium compound till you recognize the flame.
2. The sodium spectrum gives two intense lines in the
yellow which lie so close that they often seem but one.
They coincide with Fraunhofer's D lines in the solar
spectrum.
334 AMMONIUM.
AMMONIUM.
SYMBOL, NH4. — MOLECULAR WEIGHT, 18.
365. When sodium amalgam containing one to three
per cent of sodium is thrown into a strong solution of
ammonium chloride, a curious spongy substance is formed,
which gradually rises in the vessel, filling a large amount
of space. It is very unstable, giving off ammonia and
hydrogen, and leaving metallic mercury. This substance,
according to the most careful examinations, contains nitro-
gen and hydrogen in the proportions indicated in the for-
mula NH4, and this is simply in combination with mercury.
As this group plays the part of a metal in the salts
obtained from ammonia and the acids, — as in (NH4)C1,
(NH4)NO3, (NH4)2SO4, etc., — it is called ammonium, and
the compound with mercury, ammonium amalgam ; hence,
further, the salts obtained with ammonia are called ammo-
,riium salts. The metal ammonium, NH4, is, however,
hypothetical.
OF THE AMMONIUM SALTS WE NOTICE I —
(a) Ammonium Chloride, or /Sal Ammoniac, NH4C1, which
occurs as a natural deposit, but is now prepared from the
ammoniacal liquors of gas works. The ammonia gas is liber-
ated from the gas liquors by adding slaked lime, and is led into
a dilute solution of hydrochloric acid, from which this salt is
obtained by evaporation ; the chloride is afterwards purified by
sublimation. This salt is used as a reagent and as a source of
ammonia in the laboratory, and as an important aid in solder-
ing, welding, etc.
(b) Ammonium Nitrate, NH4NO3, is used as a source of
Laughing Gas or Nitrous Oxide, and can be prepared by neu-
tralizing nitric acid with ammonia.
THE RARER METALS OF THE FIFTH GROUP. 335
(c) Sodium - Ammonium Phosphate, or Microcosmic Salt,
HNaNHJPO4 + 4 HX), is much used in blow-pipe work, since
it forms a colorless bead on the platinum wire, and receives a
color by adding certain substances. It is formed by the decom-
position of urine, and is artificially prepared by dissolving five
parts of sodium phosphate with two parts of ammonium phos-
phate in hot water, and allowing the solution to cool.
(d) Ammonium Carbonate, (NH4)2CO3, is used as a group
reagent, and is now prepared by subliming CaCO3 with ammo-
nium sulphate, and digesting the product formed with strong
aqua ammoniac.
(e) Ammonium /Sulphide, (NH4)2S, is used as a group
reagent, and is very unstable, passing into (NH4)2Sx upon
exposure. This reagent is readily prepared in the laboratory
when needed by passing a current of hydrogen sulphide gas
into aqua ammoniae until the solution will not precipitate mag-
nesium sulphate.
366. Tests for Ammonium. — The tests for ammonium
have already been given (Art. 55), and it only remains
to add that, in the course of analysis, although the ammo-
nium salts remain in the fifth group, it is necessary to
apply these tests directly to the original solution.
THE RARER METALS OF THE FIFTH GROUP.
LITHIUM.
SYMBOL, Li'. — ATOMIC WEIGHT, 7.
367. Lithium is a rare metal which is found in Lepido-
lite, Triphylline, (Li,Na)8PO4 + (Fe,Mn)3PO4, and some
other minerals. This metal occurs in most surface waters
and in many mineral waters, and easily finds its way into
the animal and vegetable kingdoms.
336 RUBIDIUM. — CAESIUM.
It is prepared* by electrolyzing its chloride, and is a
silver-white metal, readily oxidizing in the air.
The principal salt is the carbonate, which is used in
medicine. The chloride, nitrate, sulphate, etc., can be
prepared by treating the carbonate with the proper acid.
QUERY. Why are the carbonates of the metals chiefly employed in
preparing the rarer salts ?
368. Tests for Lithium. — 1. Lithium compounds color
the flame intensely crimson ; this color is obscured only by
very thick blue glass.
2. The spectrum of lithium affords a certain test, yielding
the bright-red line, Lia, and the weak yellow line, Li/3.
RUBIDIUM.
SYMBOL, RB'. — ATOMIC WEIGHT, 85.
•
369. Rubidium is prepared like potassium, which metal
it closely resembles. It is widely distributed, but occurs
only in very minute quantities. It is found in Lepidolite,
Triphylline, Mica, Orthoclase, and other minerals, as well
as in various waters and soils.
Rubidium is detected by its coloring the flame some-
what more red than potassium, but more certainly by its
spectrum, which yields two violet lines, Rbo. and Rb/3.
CJESIUM.
SYMBOL, Cs'. — ATOMIC WEIGHT, 133.
370. Caesium is the first metal discovered by the spectro-
scope, and occurs with the other alkali metals. It has not
been prepared, but its salts are known.
Ccesium is detected by its spectrum, which yields the
bright-blue lines, Csa and Cs/3.
DETECTION OF THE FIFTH GROUP METALS. 337
371. Detection of the Fifth Group Metals. — 1. Test
the original solution for ammonium.
2. Free the solution from the first four groups (magne-
sium exeepted) by adding NH3, NH4C1, and NH4CO3; the
filtrate is to be tested for Na, K, and Li ; accordingly,
evaporate the solution nearly to dryness, and proceed
thus : —
(a) The sodium flame is to be observed by the naked
eye, and is intensely yellow.
NOTE. Remember that traces of sodium are usually present.
(6) Sodium obscures the violet potassium flame, but the
potassium flame becomes visible when observed through
the blue glass which shuts off the sodium color.
(V) The lithium flame is readily determined by its crim-
son color. It is obscured only by very thick blue glass.
The lithium flame is visible even when Na and K are
present.
GENERAL NOTE. The student is not to infer that the analytical grouping
of the metals or the numbering of the groups is otherwise than purely
arbitrary. Many different groupings can be made, depending upon the
reagents employed in the course of analysis. The following table will
enable the student to compare the grouping and numbering used in this
book with those used by Fresenius : —
I K, Na, NH4, Li V.
II Ba, Sr, Ca, Mg IV.
Ill Al, Cr
IV.... Zn, Mn, Ni, Co, Fe...
y ( Ag, Hg, Pb I.
( Bi, Cu, Cd ) n
VI As, Sb, Sn )
The Roman numerals in the first column indicate the groups given in
Fresenius.
372. To Analyze an Unknown Solution. — In making
a complete qualitative analysis of an unknown solution, it
338
TO .ANALYZE AX UNKNOWN SOLUTION.
is desirable to proceed by a methodical plan. From what
has preceded, it is evident that the first step should be to
determine the bases ; this may be accomplished as indi-
cated in the following table. When we know what bases
are present, we are then prepared to determine the acids.
In case we obtain arsenic, chromium, manganese, etc., we
know that these elements are apt to be present as acids.
Accordingly we first try for the acids formed by those
elements. In case these elements are not present, we
remove the bases by E (as explained farther on), and then
test for acids as in Art. 227.
A.
The solution may contain a salt of : —
1.
2.
;*.
4.
5,
Pb, Ag, or Hg'
Hg, Cd
Bi, As,
, Pb, Cu,
•Sb, Sn
Fe, Cr, Al, Zn,
Mn, Ni, Co....
Ba, Sr,
Ca, Mg
K, Na,
NH4, Li
The precipitates : —
Hd = ^>C1«, AgC1, HS«.C1« |
white white white
Solutions of 2, 3, 4C
and 5.
Filter out the precipitate, and proceed by Art. 247.
Treat the filtrate by B.
Filtrate from A : —
Hg, Cd, Pb, Cu,
Bi, As, Sb, Sn
Fe, Cr, Al, Zn,
Mn, Ni, Co...
Ba, Sr, Ca, Mg
Na, K, NH4, Li
B.
The precipitates : —
As2S3 Sb2S3>
SnS SnS PbS
yellow orange
BiS CuS CdS
brown yellow black
HgS ( Solutions of
black' black' yellow' black I 3, 4, and 6.
Filter out the precipitate, and proceed by Art. 278.
Boil the filtrate to expel H2S, and add a little HN03, and boil a short
time to oxidize ferrous to ferric salts, and then proceed by C.
TO ANALYZE AX UNKNOWN SOLUTION. 339
c.
Filtrate from B : — The precipitates : —
Mn, Ni, Co....
Fe Cr Al Zn + NH3 + NH4C1 = JC2l^;6, ^"-"Je, ^2l"^o. +
reddish brown bluish green whitegelatinous
Solutions of Zn, Mn, Ni, Co, 4, and 5.
Ba, Sr, Ca, Mg [ Filter out these precipitates, and proceed by Art. 303.
K NH I (To the filtrate) + NH4S =
, IV, JN±14,.L1 \
flesh col. black black white
Solutions containing 4 and 5.
Filter out this precipitate, and proceed by Art. 321.
Boil the filtrate to expel H2S, and proceed by D.
D.
Filtrate from C : — The precipitates : —
5.
_ \. . XTTT i "Mtr r*t i /ATTT \ rir\ BaCOo SrCOo CaCOo
+ NH3+NH(C1+(NH4)2C03=- , -
Na,K.NH<,Li
Mg and 5.
Filter out these precipitates, and proceed by Art. 355.
Divide the filtrate in two parts ; to one of these parts add Na2HP04 :
MgNH4P04
precipitate, white
Test the second part by Art. 371 for 5.
Test for acids by E.
E.
1. If the solution contains arsenic, chromium, or manganese, etc., test
the solution for the acids formed by these elements.
2. When the solution contains only the metals of the fifth group, test
the original solution directly for acids, following the directions under
Art. 227, and as given under each acid in the non-metals.
3. When other metals, not acid forming, are found, it is best to make
the solution neutral with KOH, and then to add K2C03 to precipitate them.
Filter out the precipitate, and test the filtrate. In case calcium super-
phosphate be present, the phosphate will be found in the precipitate.
Now, since we have added a carbonate, the filtrate contains the added
carbonate. In consequence of this, we must test the original solution for
carbonates. Before proceeding as in Art. 227, it is best to remove the
added carbonate by means of HC1 ; in this way we get a solution which
340 TO ANALYZE AN UNKNOWN SOLUTION.
may be tested for all. the non-metallic acids excepting HC1. We may pre-
pare another portion of the filtrate containing the added carbonate by
adding HNO3 ; this solution is to be tested for HC1.
Test for some Organic Acids given under F.
F.
1. Tartaric Acid, H2(C4H406), is detected by adding AgN03 to the nor-
mal solution ; a white precipitate is thrown down, which turns black on
boiling. And further, when tartaric acid is ignited, it gives off the odor of
burnt sugar. CaCl2 gives a white precipitate, Ca(C4H406), soluble in cold
solution of KOH.
2. Acetic Acid, H(C2H302), forms a red solution with Fe2Cl6, which is
not decolored by adding HgCl2, while red KCyS solutions with Fe2Cl6 are
thus decolored.
Also, when warmed with sulphuric acid and a little alcohol, acetic acid
gives off the odor of acetic ether.
3. Citric Acid, H3(C6H507), gives a white precipitate with AgN03, which
does not blacken on boiling; also it gives a white precipitate with lead
acetate. Further, concentrated nitric acid produces from it acetic and
oxalic acids.
4. Oxalic Acid, H2C204, is decomposed into COj and CO by H2S04.
When treated with CaCl2, the oxalates give a white precipitate, soluble
in HC1, insoluble in acetic acid. (See Art. 227.)
APPENDIX.
APPENDIX.
THE LABORATORY.
1. The Room selected for the chemical laboratory should be dry,
well lighted, and well ventilated. Generally an upper room is preferable
to a basement ; basements are apt to be damp, and poorly lighted, and
the laboratory fumes are not so easily restrained from diffusing them-
selves through the building ; with proper precautions, however, little or no
inconvenience will arise from the use of a dry, well ventilated basement
room.
It is desirable that the rooms devoted to chemistry and physics should
be adjacent to each other, as many pieces of apparatus will illustrate
portions of both studies. If communication between the two rooms can
be secured by sliding doors, so much the better ; this arrangement offers
many advantages in those schools where chemistry and physics are taught
by the same teacher. In case the rooms cannot be adjacent, they should
be as near together as possible.
GENERAL FIXTURES.
In case the building is heated by steam, and lighted by gas, many of
the general fixtures are easily provided.
2. The Condenser for procuring distilled water may be connected
directly to the steam-pipes used in heating the building. A plain sheet
copper cylinder 30cm in diameter, and 135cm high, will afford all the
distilled water thirty students will require : this cylinder simply needs
a faucet at the bottom, through which the water may be drawn when
needed, and a small pet-cock at the top, through which the air is to be
blown out when the steam is first turned on. The steam is admitted
at the top of the cylinder which stands upright, and which needs no
internal coil nor external jacket. The cylinder should be able to carry
all the pressure that the boilers are likely to put upon it, and it may stand
in any convenient part of the room, as no hissing or other disagreeable
noise is heard.
344 APPENDIX.
In case the building does not contain steam, permission may be
obtained from some factory or mill to connect such a condenser to the
boilers used there. The connecting pipe should be as small as possible,
and the steam should be allowed merely to leak through the valve, by
means of which the condenser is shut off from the boiler.
Many other devices are to be had, some of which are applicable, under
one condition, while under other conditions another device may succeed
more satisfactorily, e.g.
Small quantities of distilled water are to be had by means of a Liebig
condenser, in connection with a still heated by a gasoline stove, or by an
ordinary stove ; a coil may be passed through a cask containing cold
water, etc., etc.
One fact should be noted here ; ordinary rain-water, and water as
usually prepared by distillation, usually contain free ammonia. Water free
from ammonia may be obtained as explained in App. 77.
3. The Tank for Wash- Water may be placed in a corner of the room,
and its bottom should be four or five feet higher than the faucets from
which the water is drawn. Pipes leading from the tank may carry the
water to a sink and to each student's desk.
Pure cistern water is best for ordinary washing purposes in a labor-
atory ; the water may be raised to the tank by a force-pump, or a cistern
may be constructed under the roof of the building.
4. A Gas Chamber is useful for many purposes. It may be built of
sash with glass, and it may stand in any convenient place, so that it may
be connected to a good ventilating shaft. By means of such an arrange-
ment, the operator can observe what is taking place, and the unwhole-
some gases generated are carried out of the building. It is convenient to
have two or three separate apartments not in communication with one
another, and each one with a separate door. The size of such a chamber
will depend upon the requirements of the school, but one 3 ft. square X 6 ft.
high will answer for most small laboratories.
5. Cases for chemicals, apparatus, etc., are convenient and inexpen-
sive. It is desirable to have a portion of the case provided with sasli
doors, and the remainder is to be cased with panel doors, thus providing
dark closets in which stock chemicals and reagents may be kept to better
advantage.
6. Working Tables may be placed against the walls of the room, or
through its centre. A table 15 ft. long, 3 ft. 1 in. high, and 3 ft. 4 in.
wide, and standing from the walls, will afford ample room for eight
APPARATUS FOE STUDENT'S DESK. 345
students to work at a time. If the class be divided into two working
divisions, such a table will accommodate sixteen students, while the
apparatus per student will thereby be materially lessened. In the centre
of the table are placed four desks, while sink-bowls are placed between.
One side of such a desk is shown in the Frontispiece ; this cut is taken
from the photograph of a desk in Ypsilanti High School Laboratory. In
the table just under the desk is a drawer, used by the student to keep his
apron and other personal property which he requires in his work.
The tables may be supported by legs or by square posts ; in the latter
case, cupboards may be constructed under the tables ; but in case cup-
boards are made, a bottom or an extra floor should be put in, so that the
base-board under the doors may not form an obstruction in sweeping out
any dust, etc., that may collect in the cupboard.
The dimensions of the desk shown are as follows : Height, 2 ft. 4 in. ;
length, 2 ft. 6 in. ; breadth at bottom, 14 in. ; at top, 12 in. ; space under-
neath first shelf, 11 in. ; second space, 8 in., and third space 6 in. The
top of the desk may be utilized as a shelf. A partition through the desk
divides it into halves, thus forming two working cupboards, one on each
side of the desk.
The gas chamber, tables, desks, and cases, can be made by any car-
penter.
APPARATUS AND REAGENTS.
In considering the materials under this heading, it will be convenient to
follow the order : —
(a) Apparatus for the student's desk ;
(6) Reagents for the student's desk;
(c) Reagents for the side table ;
(d) Working material;
(e) General apparatus for the laboratory.
APPARATUS FOR THE STUDENT'S DESK.
Perishable apparatus, such as glass and porcelain ware, should be kept
in stock in order to supply quickly any loss by breakage, etc.
7. Test-Tubes. — At the start the student should have twelve 4-in. test-
tubes, and two 8-in. test-tubes of a larger diameter. The latter are to be
fitted with rubber stoppers pierced with one hole, through which is inserted
a bent delivery-tube ; they are used as generators.
Test-tubes are perishable, but they are not expensive. A liquid may
346 APPENDIX.
be heated in a test-tube by placing the tube directly in the Bunsen or
alcohol flarne, provided the flame does not strike the tube at the upper
level of the liquid.
When heating a substance in a test-tube, the student should never hold the
mouth of the tube towards himself nor towards others, since any explosion, as
of steam o'r other gases, might result seriously ; it is best to move the test-
tube gently through the flame when heating any substance.
With a little practice the student may mend a test-tube, the bottom of
which has been broken. To accomplish this, the tube is first to be cleaned
and dried ; the broken end is then strongly heated in the Bunsen flame
until the glass becomes soft ; the broken edges of the tube are now forced
together by means of a bit of glass tubing ; when the bottom is closed,
the end of the tube is freed from unnecessary material by carefully draw-
ing out the highly heated end of the tube with the glass rod ; the end of
the tube is now strongly heated until it becomes somewhat thicker than
the walls of the tube ; now the mending is to be finished by blowing gently
into the tube, in order to give the end a rounded form. When heating the
tube, it should be rolled over constantly in the flame, so that all sides may
be heated alike. An alcohol or gas blast-lamp may be used to good advan-
tage for this work, and for such other glass-work as usually must be done
in the laboratory.
8. Hard Glass Tubing. — Each student should have a tube 8 in. long,
and with a bore of about j in. ; this is used for heating solids as in
Exp. 3 P.
A hard glass test-tube has been mentioned in the text. These are more
expensive than ordinary test-tubes ; for most purposes mentioned a com-
mon tube may be used, but it is almost invariably ruined ; this is of no
great moment, however, if a tube that has been mended is employed.
9. A Test-Tube Rack for holding test-tubes is shown in the Frontis-
piece. The student can make this for himself by taking a suitable block
of wood and setting in one edge of it a row of wooden pins 3 in. high ; in
the other edge holes are bored, which will serve to hold tubes containing
liquids.
10. A Test-Tube Swab for washing out test-tubes is also to be made
by the student. It is simply a wooden stick as large as a lead-pencil, upon
the end of which a bit of sponge is fastened.
Test-tube brushes of various designs are also to be had in the market, but
the swab will answer for nearly every purpose.
11. A Glass Stirring-Rod may be made from suitable solid glass rods
APPARATUS FOR STUDENT'S DESK. 347
which are to be kept in the laboratory. This rod should be about as large
and long as a common slate-pencil. The ends of the rod must be melted
smooth and round in the Bunsen flame.
12. Platinum Wire and Platinum Foil are much used. The wire
should be about 3 in. long, and one end of the wire should be fused into a
short glass tube ; the other end of the tube should be closed. The plat-
inum foil may be about 1 in. X f in. The uses of these articles are
described in the text in the appropriate places.
13. A Blow-Pipe of the form shown in the Frontispiece (Bp), known
as Black's, is the best of the cheaper forms. A blow-pipe should last
many years.
14. Steel Tongs (T in Frontispiece) are useful to handle hot evap-
orating dishes, hot crucibles, etc. The student may readily hold a test-
tube while boiling solutions, etc., by putting a narrow strip of cloth around
the upper end of the tube and clasping the ends of the strip in these tongs.
These tongs should last five or six years.
15. Funnels are shown in FN ; these are of glass. The student
should have two, — one 2 in. and one 3 in. or 4 in. in diameter. The fun-
nels should have their stems ground off at an acute angle to facilitate the
process of filtration. Funnels are seldom broken.
16. Filter-Papers should be cut round, and should be furnished the
student in packages. The proper size papers for the funnels are 4 in. and
6 in. in diameter. These papers should be kept in a tin box of proper
form and size.
The filter-papers are placed in the funnel as follows : First, they are
folded through the centre ; then another fold, at right angles to the first,
is made, which leaves the paper in the form of a sector of a circle ; now,
by inserting the apex of the sector into the funnel, the paper may be
opened out in form of a cone that will fit the funnel. It will be seen that
two pockets are formed in the paper, either of which will serve as a recep-
tacle for the fluid to be filtered. It is best to wet the paper with distilled
water before filtering a solution containing a precipitate, as this tends to
prevent the precipitate from adhering so closely to the paper.
Beginners are often at a loss as to how they may divide small precipi-
tates into several parts ; this may be accomplished in different ways, of
which these two are as convenient as any : First, the point of the filter-
paper containing the well-washed precipitate may be pierced, and the damp
precipitate may be washed through into a beaker glass by means of dis-
tilled water; the precipitate may now be agitated with a stirring-rod
348 APPENDIX.
until it is suspended in the water, when portions of it may be poured out ;
Second, the precipitate may be left on the filter-paper, and whether damp
or dry may be separated into portions by tearing the filter-paper into the
requisite number of parts. If the precipitate be damp, it may be washed
off each part as needed, by means of water. If the precipitate be dry, and
the student wishes to dissolve the dried precipitate, he may put the paper
and all in a test-tube, and after dissolving may remove the particles of the
filter-paper by passing the solution through a new filter.
For filter ing acids a little spun glass is best; this may be crowded down
into -the stem of the funnel, and after passing the acid through it may be
washed and preserved for further use.
17. Generating Flasks, one each of 2-oz. and 4-oz. capacity, will
answer for the student's needs. These flasks are used for generating gases,
etc., and are fitted with delivery-tubes as shown in the Frontispiece, F.
These flasks are sometimei broken.
18. Two Beaker Glasses (see BK in Frontispiece), one of 2-oz. and
the other of 4-oz. capacity, are needed. In them solutions are boiled,
crystals are allowed to form, and solutions for working purposes are kept
temporarily, etc., etc. Neither these beakers nor the Florence flasks men-
tioned in If should be heated in the naked Bunsen flame. They should
always be placed on wire gauze or on a sand-bath.
19. Evaporating Dishes (see E in Frontispiece), one each of about
3-oz. and 4-oz. capacity, are needed. These should be heated on the sand-
bath or on wire gauze. They are seldom broken. Prof. Weitbrecht's stu-
dents frequently use saucers as evaporating dishes.
20. A Bunsen Burner is shown at B. The use of this has been ex-
plained. In laboratories not containing gas for heating purposes, alcohol
lamps are the best substitute. In nearly every place in the text where
" Bunsen flame " has been used, " alcohol flame " may be substituted.
21. A Wash-Bottle, or " Blow-Bottle " as it is familiarly termed by
students, is shown at W in Frontispiece, as made by a student from whose
desk this cut was taken. Each student can make his own bottle ; the de-
livery-tube should be drawn out into quite a fine jet, so that the stream of
water issuing from it, upon blowing into the mouth-piece, shall be quite
small.
22. Each Student should provide himself with a toy magnet, a clay
tobacco pipe for blowing soap-bubbles, a sponge, a towel, a bundle of soft
white rags, a box of matches, a watch-crystal, an oil-cloth apron, and a
pair of rubber sleeves. The uses of these are too evident to need men-
tioning.
APPARATUS FOR STUDENT^ DESK. 349
23. A Ring Stand is shown at A. This is used to support funnels
while filtering, and sand-baths, retorts, generating flasks, etc., while heat-
ing. The rings may be removed or clamped in any position upon the
upright standard.
24. A Blue-Glass is shown at G. This is a frame containing two thick-
nesses of glass. One blue-glass will answer for two desks.
25. A Sand-Bath is shown at S, resting on a ring. This is a saucer-
shaped sheet-iron dish, which may be hammered out by any tinsmith. It
must be large enough to rest on the largest ring. The dish is filled with
clean white sand, and in this sand beakers, evaporating dishes, etc., are
set ; the heat is applied to the sand-bath. There is one objection to a
sand-bath, — the sand is apt to get scattered on the student's desk and find
its way into the waste pipes leading from the sink-bowls. It is safer, how-
ever, to heat glass ware, etc., in a sand-bath than it is on a
Wire Gauze. This gauze is of fine brass wire, and is placed between
the flame and the evaporating dish. It will be found to be neater and less
objectionable in several respects than the sand-bath, but it is not quite so
safe to heat fragile ware upon it.
Professor Foote recommends asbestos paper in place of the sand-bath
and wire gauze.
26. A Match-Safe should be furnished to each desk, and the student
should not be allowed to put matches in his drawer. Employ sulphur
matches ; parlor matches are too dangerous.
27. If the student is to do a little quantitative work, he will need, in
addition to the foregoing, a porcelain crucible with cover, a feather, a
sheet of glazed paper, and a triangle made by joining three common clay
tobacco-pipe stems by means of iron wire.
28. Litmus Papers. These papers may be purchased ready for use,
or they may be prepared in the laboratory by dipping sheets of bibulous
paper in litmus solution ; the papers thus prepared are blue. Red and blue
papers are needed ; the red papers may be prepared by moistening the blue
papers in dilute acetic acid. The papers should be cut into strips 4cm
long and 4mm wide ; they may be kept in a bottle or cardboard box.
29. Charcoal. A fine variety of charcoal is to be purchased of chem-
ical dealers, but selected pieces may be obtained from ordinary charcoal
that will answer all purposes. Charcoal should not be kept in the drawers
or on the desk. Separate pans with legs should be provided to avoid dan-
ger from fires.
350 APPENDIX.
The reagents for the student's desk should be kept in stock in the lab-
oratory, i.e., a sufficient quantity of the dry salts and of the liquid reagents
should be purchased at the beginning of the year to last throughout that
year. Some of these reagents are more convenient in a dry form ; but
most of these are used in the form of solutions.
The solutions should be kept in good glass-stoppered bottles, holding
I-1 or 4 oz., similar to those shown in the Frontispiece. It is desirable that
these bottles have permanent acid-proof names and symbols.
The dry salts should be kept in small 2-oz. salt-mouth bottles, and
these are best when provided with glass stoppers.
A few words of caution concerning the care of reagent bottles are in
place here. A good reagent bottle must have its stopper ground to fit it,
and this stopper will not fit any other bottle in the set. Consequently the
stoppers should never be interchanged. Again, the stoppers of all re-
agent bottles, excepting sulphuric acid, should be paraffined with gum-
stock paraffin, otherwise they are quite apt to stick ; often the bottles are
ruined or cracked by trying to remove the stoppers. There is no excuse
for breaking a reagent bottle. The solutions should not be allowed to
freeze, as the bottles may thus be broken.
The student should not lay down the cork of a reagent bottle while
pouring out a solution, since he may thus change stoppers with his bottles
or contaminate his reagents. Again, no solution but the one correspond-
ing to the name on the bottle should ever be placed in a reagent bottle.
Another important item is that each bottle have a place on its shelf,
and always be put in its place ; thus the student comes to know where to
find a reagent, just as a printer knows where to find the letters in his case.
Since some order must be followed, that in which the reagents are
described below may be insisted on. Commencing with the first name in
the list on the upper shelf, left-hand side, arrange the bottles toward the
right ; and, when the shelf is full, begin again on the left-hand of the next.
Since systems of nomenclature vary somewhat, and since labels and
names are apt to vary decidedly, all the names are given in connection
with each reagent, the most preferable coming first, the symbol next, and
thereafter the various other names, in order of their preference, excepting
the name given in italics, which is that of the United States Pharmaco-
poeia ; its position has no reference to its preferment.
In naming the acids, the common names are given first, for the reason
that these names are good ones, and in spite of all attempts to do away
with them, they still persist in remaining ; and it is perhaps but wise to
submit to the inevitable. Thus, that acid whose formula is H2S04, is
called sulphuric acid ; hydrogen sulphate, for some reasons, would be better,
LIQUID REAGENTS. 351
but the change is not universally accepted. Again, hi/dric sulphate has
been proposed, but this is still less favorably received ; while the oldest
name of all, oil of vitriol, is scarcely used or known by the last generation
of chemists, though still retained by manufacturers.
The reagents enumerated below (with a few exceptions, which are
noted) should be chemically pure. Of all persons, a beginner should have
the best materials to work with ; moreover, good material is now so cheap
that there is neither profit nor sense in using goods of a poor quality.
LIQUID REAGENTS.
30. Sulphuric Acid, H2S04; Hydrogen Sulphate; Hydric Sulphate ;
Dihydric Sulphate ; Oil of Vitriol ; Acidum Sulphuricum.
This acid should be bought in a concentrated form, sp. grav. 1.843, and
should be dealt out to students in this form ; it should evaporate on plati-
num foil without leaving any residue, and it should be colorless.
The commercial acid may be contaminated with arsenic, antimony, iron,
aluminum, calcium, potassium, sodium, lead, magnesium, hydrochloric
acid, nitrous acid, nitric acid.
31. Nitric Acid, HN03 ; Hydrogen Nitrate; Hydric Nitrate; Aqua
Fortis ; Acidum Nitricum.
This acid may be bought in a concentrated form, and afterward re-
duced with water to reagent strength, which is 32 per cent acid, sp. grav.
1.32. (See 34 for computation.)
Pure nitric acid is colorless, but, on standing exposed to the light, it
may become colored by the lower oxides of nitrogen, which, as a usual
thing, are not harmful. They may be removed by passing a current of air
through the acid by means of a glass tube attached to a hand-bellows.
The commercial acid may contain calcium, sodium, iron, oxides of nitro-
gen, hydrochloric acid, sulphuric acid.
32. Hydrochloric Acid, HC1 ; Hydrogen Chloride ; Hydric Chloride;
Muriatic Acid ; Chlorhydric Acid ; Chlorhydrate ; Spirit of Salt ; Acidum
Hydrochloricum.
This acid may likewise be purchased in a concentrated form, and after-
wards reduced to the reagent strength, 24 per cent acid, sp. grav. 1.12.
The pure acid is colorless, and leaves no residue upon evaporation ; upon
standing, it may become colored by free chlorine.
The commercial acid may contain iron, sodium, aluminum, arsenic, sul-
phuric acid, sulphurous acid.
33. Acetic Acid, H(C2H30.2); Hydrogen Acetate ; Hydric Acetate;
Acidum Aceticum.
352 APPENDIX.
Since acetic acid is not so extensively used as the preceding acids, it
may be purchased of a reagent strength, 30 per cent acid, sp. grav. 1.04.
The pure acid is colorless, and leaves no residue upon evaporation.
The commercial acid may contain sodium chloride, lead, copper, iron,
empyreumatic substances, sulphuric acid, sulphurous acid, nitric acid.
34. Ammonia, NH3. The reagent solution contains 10 per cent of the
gas NH3, and has a sp. grav. 0.96. It is prepared from the "Stronger
Water of Ammonia," or Aqua Ammonia (28 per cent gas; sp. grav. 0.90 ;
U. S.-P.), by the addition of distilled water. The concentrated form is
more convenient to keep in stock, as it requires less space for storage. In
the case of ammonia, and of the concentrated acids previously mentioned,
the label of the original package should state the per cent and sp. grav.
Commercial aqua ammonia may contain ammonium chloride, ammonium
carbonate, calcium sulphate, empyreumatic material.
The amount of water to be added to a given volume of a stronger solu-
tion may be determined by calculation. Thus, in the case of ammonia :
We know that I1 of the strong solution weighs 900s, and that 28 per cent
of that weight, or 252s, is NH3. It is evident that this 252s is to form
10 per cent of the weight of the reagent solution ; hence, the whole weight
of the reagent solution will be 252 -^- .1 — 2520s. Now, we already have
taken 900s of the strong solution ; consequently 2520 — 900 = 1620s, or
the weight of distilled water to be added to I1 or 1000CC of the strong
solution. It is further evident that one part, by volume, of the strong
solution requires 1.62 parts, by volume, of distilled water.
35. Ammonium Carbonate, (NHJ2C03; Carbonate of Ammonia; Am-
monic Carbonate ; Volatile Salt ; Ammonii Carbonas.
This solution is prepared by dissolving 1 part by weight of the dry salt
in 4 parts by weight of water, after which one part of reagent ammonia
solution is added.
The commercial salt may contain calcium, iron, lead, chlorides, iodides,
sulphates.
It is not necessary to weigh the water, since lcc of water weighs is.
The graduated ware used in measuring solutions is graduated at a certain
temperature, usually 15° C. When accuracy is required, the temperature
of the water or of the solution to be measured should be that at which the
apparatus is graduated.
36. Ammonium Sulphide, (NHJ2S ; Sulphide of Ammonium ; Am-
monic Sulphide.
This solution may be purchased ready for use, or it may be prepared in
LIQUID REAGENTS. 353
the laboratory by passing hydrogen sulphide gas through a reagent solu-
tion of ammonia until the solution no longer precipitates magnesium
sulphate. This reagent changes, upon standing, to the yellow variety.
Although the formula of the yellow ammonium sulphide has been given
as (NH4)2S2, its composition varies greatly.
37. Ammonium Chloride, NH4C1 ; Chloride of Ammonium ; Am-
monic Chloride ; Muriate of Ammonia; Sal Ammoniac; Ammonii Chloridum.
To prepare this reagent solution, dissolve 1 part of the crystallized salt
in 8 parts of water.
The commercial salt may contain iron, sulphates, organic matter.
38. Ammonium Oxalate, (NH4)2C2O4 ; Oxalate of Ammonium ; Am-
monic Oxalate.
This is prepared by dissolving the crystallized salt, (NH4)2C204+ H2O,
in 24 parts of water.
The commercial salt mat) contain sodium, potassium, calcium, aluminum,
lead, sulphates, nitrates. %
39. Potassium Hydroxide, KOH ; Potassium Hydrate; Potassic
Hydrate; Caustic Potash.
This solution is prepared by dissolving 1 part of the dry sticks in 20
parts water. It is not absolutely essential that this salt be strictly C. P. ;
there is a good white article ("rein weiss") containing a little silica, and
perhaps a trace of chlorine, that will answer most purposes, and it is much
cheaper than the C. P. article.
The commercial article may contain iron, aluminum, sodium, calcium, or-
ganic matter, silica, chlorides, sulphates, carbonates.
44'. Sodium Hydroxide, NaOH, is preferred by many chemists to
potassium hydroxide, since the former is much cheaper. This solution is
made by adding 1 part of the fused substance to 9 parts water. The im-
purities are much the same as in the potassium compound.
40. Potassium Carbonate, K2C03 ; Carbonate of Potassium; Potas-
sic Carbonate ; Carbonate of Potash (potassa) ; Potassii Carbonas.
Make this solution by dissolving 1 part of the dry salt, K2C03 + 3H20,
in 10 parts water.
The commercial article may contain iron, aluminum, silica, sodium, chlo-
rides, sulphates, sulphides.
41. Potassium Iodide, KI; Iodide of Potassium; Potassic Iodide;
Potassii lodidum.
Dissolve 1 part of the salt in 20 parts of water.
354 APPENDIX.
The commercial artide may contain sodium, iodates, sulphates, chlorides,
carbonates.
42. Potassium Bichromate, K2Cr207 ; Bichromate of Potassium; Po-
tassium Bichromate ; Potassic Bichromate ; Bichromate of Potash ; Red
Chromate of Potash; Potassic Acid Chromate ; Potassii Bichromas.
1 part of the salt is dissolved in 10 parts of water.
The commercial salt may contain iron, calcium, aluminum, sulphates,
chlorides.
43. Potassium Sulpho-Cyanide, KCyS ; Sulpho-Cyanide of Potas-
sium; Potassic Sulpho-Cyanide; Potassium Sulpho-Cyanate.
This solution is made by dissolving 1 part of the salt in 25 parts of
water.
The commercial article may contain iron, sulphates, chlorides.
44. Potassium Ferro-Cyanide, K4FeCy6; Ferro-Cyanide of Potas-
sium ; Potassic Ferro-Cyanide ; Yellow Prussiate of Potash ; Potassii Ferro-
cyanidum.
This solution is made by dissolving 1 part of the crystallized salt,
K4FeCy6,3H2O, in 12 parts of water.
45. Disodium Phosphate, Na2HPO4; Sodium Phosphate ; Phosphate
of Sodium ; Disodium-Hydrogen Phosphate ; Disodic-Hydric Phosphate ;
Sodii Phosphas.
This solution is prepared by dissolving 1 part of the crystallized salt,
Na2HPO4+ H20, in 10 parts of water.
The commercial salt may contain arsenic, iron, lead, sulphates, chlorides.
46. Barium Chloride, BaCl2 ; Chloride of Barium ; Baric Chloride ;
Bar ii Cldoridum.
Dissolve 1 part of the crystallized salt, BaCl2 + 2H2O, in 10 parts of
water.
The commercial article may contain calcium, strontium, iron, aluminum,
silica.
47. Calcium Hydroxide, Ca(OH)2; Calcic Hydrate ; Lime Water;
Liquor Calcis.
This solution is best prepared in the laboratory. " Slake the lime by
the gradual addition of 6 parts of water, then add 30 parts of water, and
stir occasionally during half an hour. Allow the mixture to settle, decant
the liquid and throw this away. Now add to the residue 300 parts of
distilled water, stir well, and wait a short time for the coarser particles to
subside, and then pour the liquid, holding the undissolved lime in suspen-
sion, into a glass-stoppered bottle. When wanted for use, pour off the
clear liquid." — U. S. P.
LIQUID REAGENTS. 355
48. Magnesium Sulphate, MgS04 ; Sulphate of Magnesium ; Mag-
netic Sulphate ; Sulphate of Magnesia ; Epsom Salt ; Magnesii Sulphas.
Dissolve 1 part of the crystallized salt, MgSO4 + 7H2O, in 10 parts of
water.
The commercial salt may contain calcium, iron, silica, zinc, manganese,
chlorides.
49. Mercuric Chloride, HgCL2 ; Bichloride of Mercury ; Perehloride
of Mercury ; Corrosive Sublimate ; Corrosive Chloride of Mercury ;
Hydrargyri Cldoridum Corrosivum.
Dissolve 1 part of the crystallized salt in 70 parts of water.
The commercial salt may contain iron, lead, calcium, antimony, tin.
50. Silver Nitrate, AgN03; Nitrate of Silver; Argentic Nitrate;
Lunar Caustic ; Argenti Nitras.
Dissolve 1 part of salt in 70 parts of water.
The commercial salt may contain iron, lead, copper.
51. Lead Acetate, Pb(C2H3O2)2; Acetate of Lead; Plumbic Acetate ;
Sugar of Lead ; Plumbi Acetas.
Dissolve 1 part of the crystallized salt,Pb(C2H302)2 + 3 H20, in 10 parts
of water. If the solution is not clear, filter it.
The commercial salt may contain sodium, calcium, iron, lead, copper,
chlorides, nitrates.
52. Ferric Chloride, Fe2ClG ; Perehloride of Iron ; Sesquichloride
of Iron ; Ferri Chloridum.
Dissolve 1 part of the solid salt, Fe2Cl6 + 6 H2O, in 15 parts of water.
The commercial article may contain feurous chloride, aluminum, nitrates,
sulphates.
53. Alcohol, C2H6O ; Ethyl Alcohol ; Spirits of Wine.
The alcohol used should be the " Spirits of Wine," having a specific
gravity of .815, and containing about 95 per cent of the spirit. This
should be purchased ready for use.
54. Cobaltous Nitrate, Co(N03)2.
This solution is prepared by dissolving 1 part of the crystalline salt,
Co(NO3)2 + 5H2O, in 20 parts of water.
This solution is used merely for moistening the bead on the platinum
wire, and should be kept in a small half-ounce bottle, as this amount will
last a long time.
356 APPENDIX.
DRY REAGENTS.
55. Ferrous Sulphate, FeSO4 + 7H20 ; Sulphate of Iron ; Green
Vitriol ; Ferri Sulphas.
This reagent is used in solution, 1 part of the salt to 10 parts of water;
but the solution oxidizes rapidly to a ferric condition, in consequence of
which, it is best to make the solution in a test-tube, as required from time
to time ; the proportions need not be exact.
The dry salt also oxidizes by standing ; hence, in practice, a crystal of
the salt is dropped into the test-tube, and a little water added ; the crystal
is now shaken until the white coating of the ferric salt disappears, and the
crystal is of a clear green color ; this water is now thrown out, and a fresh
portion added; heat is then applied to hasten the solution.
56. Sodium Carbonate, Na2C03 ; Carbonate of Sodium ; Sodic Car-
bonate ; Sodii Carbonas.
This reagent is used in the form of the dry, powdered salt ; the bottle
containing it should be kept well corked to prevent the reagent from
absorbing the gases of the laboratory.
The commercial salt may contain iron, aluminum, silica, calcium, lead,
chlorides, sulphates, sulphides.
57. Sodium Borate, Na20(B203)2; Borate of Sodr::n; Borax; Sodti
Boras.
This reagent is used in a dry, powdered form.
The commercial article may contain iron, sodium, aluminum, silica, cal-
cium, chlorides, sulphates.
58. Sodium-Ammonium Phosphate, NaNH4HP04 . 4H20 ; Microcos-
mic Salt ; Sodii et Ammonii Phosphas.
This is used in a dry state.
59. Ferrous Sulphide, FeS.
The method of using this sulphide is explained in the text, Art. 167.
60. Potassium Chlorate, KC103 ; Chlorate of Potassium ; Potassic
Chlorate; Chlorate of Potash; Potassii Chloras.
The crystallized salt is used.
61. Metallic Zinc, Zn.
The granulated metal is employed. This form is obtained by pouring
molten zinc into water. It must be absolutely free from arsenic. (See
Art. 319.)
REAGENTS FOR THE SIDE-TABLE.
REAGENTS FOR THE SIDE-TABLE.
These reagents are those required occasionally by the student. One
set should be prepared and placed on a side-table, or in a c&pboard conven-
iently located, so that it is accessible to all the students in the laboratory.
The solutions may be kept in 4-oz. bottles similar to those on the student's
desk. The eorks of all these bottles, excepting those for ether and carbon
bisulphide, should be paraffined. The dry salts are to be kept in convenient
broad-mouth bottles.
62. Carbon Bisulphide, CS2 ; Carbon Bisulphide ; Bisulphide of Car-
bon; Carbonei Bisulphidum.
This reagent is purchased ready for use. It is very volatile, and the
bottle should be closed with a good chemical cork stopper.
63. Ether, (C2H5)2O; Aether; Sulphuric Ether.
This reagent is purchased ready for use, and the bottle should be closed
with a chemical cork stopper.
64. Potassium Sulphate, K2S04 ; Sulphate of Potassium ; Potassic
Sulphate ; Sulphate of Potash ; Potassii Sulphas.
Dissolve 1 part of the crystallized salt in 12 parts of water.
65. Potassium Ferri-Cyanide, K3FeCy6 ; Ferricyanide of Potas-
sium ; Red Prussiate of Potash.
Dissolve 1 part of the salt in 12 parts of water. This solution will not
keep long without undergoing decomposition.
66. Potassium Chromate, K2Cr04; Chromate of Potassium; Potas-
sic Chroinate.
Dissolve 1 part of the salt in 10 parts of water.
67. Potassium Cyanide, KCy; Cyanide of Potassium; Potassic
Cyanide ; Potassii C yanidum.
1 part of the solid is dissolved in 4 parts of water. The poisonous
nature of this reagent should not be forgotten.
68. Potassium Permanganate, K2Mn2O8; Permanganate of Potas-
sium ; Permanganate of Potash ; Potassii Per many anas.
Dissolve 1 part of the crystallized salt in about 500 parts of water.
69. Sodium Sulphite, Na2S03; Sulphite of Sodium ; Sodic Sulphite;
Sodii Sulphis.
Dissolve 1 part of the crystallized salt, Na2S03+7H20, in 5 parts of
water.
358 APPENDIX.
70. Calcium Sulphate, CaSO4; Sulphate of Calcium; Calcic Sul-
phate ; Calcii Sulphas.
This solution is made by dissolving all the salt, CaSO4+ 2H20, that the
water will take up ; or, in other words, it is a saturated solution.
71. Calcium Chloride, CaCl2; Chloride of Calcium; Calcic Chloride ;
Culcii Chloridum.
Dissolve 1 part of the salt, CaCl2 + 6 H20, in 8 parts of water.
72. Stannous Chloride, SnCl2 ; Protochloride of Tin.
To 6 parts of water add 1 part of the crystallized salt, SnCl2 -f 2 H20 ;
then add hydrochloric acid, drop by drop, until the solution turns clear.
73. Copper Sulphate, CuS04; Sulphate of Copper; Cupric Sulphate;
Blue Vitriol ; Blue Stone ; Cupri Sulphas.
Dissolve 1 part of the crystallized salt, CuS04+5H20, in 8 parts of
water.
74. Starch Paste. This solution is made by dissolving 1 part of
starch in 500 parts of water. In case the student desires a solution of
starch paste and potassium iodide, he may place a little of the starch paste
solution in a test-tube, and add a drop or two of the reagent potassium io-
dide solution.
75. Ammonium Molybdate, (NHJ2Mo4. Dissolve 60s of the dry
salt in 400CC of reagent ammonia solution; add 400CC of distilled water;
then cautiously add 500CC nitric. acid (sp. grav. 1.4).
GRADUATED SOLUTIONS, ETC.
76. Clark's Soap Solution is prepared by dissolving IQs of good
castile soap in I1 of dilute alcohol containing about 35 per cent of the
spirit. The dilute alcohol may be prepared from the reagent alcohol
by mixing 368.5CC alcohol with 631.5CC distilled water.
To test the soap solution a reagent solution of calcium chloride is required.
This solution is prepared by dissolving is of Iceland spar in hydrochloric
acid ; the solution is then evaporated to dryness to expel any excess of
acid, after which the residue is dissolved in I1 of distilled water. Now if
12CC of the solution just formed be diluted to 70CC and brought into a flask,
it will require just 13CC of the soap solution to make a permanent lather,
provided the soap solution be of the right strength. In case the soap so-
lution is not of the right strength, it must be made so, or allowances must
be made when calculating the degrees of hardness of a sample of water.
The soap solution deteriorates by standing.
INDICATORS. 359
77. Nessler's Solution is prepared by dissolving 13" mercuric chlo-
ride, HgCl2, in about 400CC of distilled water ; now 358 of potassium iodide,
KI, are dissolved in (say) 200CC of water, and these two solutions are then
mixed. To this solution add 100s of solid potassium hydroxide, KOH, and
when it is dissolved and the solution cool, dilute the whole with water to
I1. Keep this solution in a dark, cool place, and take a portion of it in a
small bottle for immediate use.
Before using the solution it is necessary to " sensitize " it ; this is
accomplished by adding slowly a saturated solution of mercuric chloride,
with constant stirring, until the red precipitate first formed ceases to dis-
solve. Either filter the solution or allow it to stand till the solids have all
subsided. It is now ready for use, and should be of a light, straw-yellow
color. This solution loses its sensitiveness by standing.
78. A Few Graduated Solutions have been mentioned in the text ;
as, for example, Barium Hydroxide Solution and Oxalic Acid Solution,
p. 148 ; Silver Nitrate Solution, p. 107 ; Iodine Solution, p. 181 ; Ammo-
nium Chloride Solution, p. 72. These have been sufficiently described, so
that there is nothing to add, unless it be to note that in case these solu-
tions prove too strong that they may be diluted to some other standard of
strength ; for example, it is evident that if lcc of the ammonium chloride
solution be added to 99CC of distilled water, lcc of the solution thus formed
will correspond to .Olms of ammonia. It is usually necessary to work,
when estimating the ammonia of drinking-water, with this dilute solution.
Now, if the burette used be graduated to .lcc, it is evident that by this
means the ammonia in drinking-water, etc., may be determined to .001Ing.
It might be well, in this connection, to call attention to the extreme accu-
racy obtainable in titration.
N.B. A few words of caution concerning the estimation of chlorine may
be in place here. It is evident the chromate used for an indicator must
be free from chlorine ; also, in order to have the end reaction sharp, the
solution must be exactly neutral.
In estimating ammonia, the water used in connection with the standard
solution of NH4C1 must be free from ammonia. This may be obtained
by taking (say) 21 of distilled water, and distilling until the distillate gives
no reaction for ammonia. The water remaining in the retort is evidently
free from ammonia.
INDICATORS.
Solutions of various substances are employed to indicate what is called
"End Reactions." The method of using these indicators has been ex-
360 APPENDIX.
plained in the text. Lt now remains to show how a few of these solutions
are made.
79. Litmus Solution is prepared by digesting for several hours 10s of
solid litmus with 500CC of distilled water ; allow the liquid to become clear,
or filter it when it is ready for use, when the end reaction is to be aeid ;
one portion of it may be prepared for solutions, when the end reaction is
to be alkaline, by adding to it a few drops of acetic acid.
80. Cochineal Solution is obtained by digesting 3s of the powder in
250CC of 20 per cent alcohol. This is very sensitive ; acids bleach it, alka-
lies redden the bleached solution.
81. Phenol-Phthalein Solution is made by dissolving 1 part of the
solid in 100 parts of 60 per cent alcohol ; this gives a colorless solution
which is reddened by alkalies. This red solution is bleached by acids. It
may be used as a qualitative test for carbon dioxide. See "American Chemi-
cal Journal," 3, 55, 232. For a paper on Lakmoid, Phenol-Phthalein, and
other indicators, see " The Chemical News " of July 10, 1885, p. 18, and
July 17, 1885, p. 29.
82. A Soap-Bubble Solution is prepared thus: To about 100s of
finely-cut best castile soap in a litre flask add nearly a litre of distilled
water ; shake until the solution is saturated with soap ; then allow it to
settle clear ; to two volumes of soap solution add one volume of glycerine.
GENERAL NOTE. In order to lessen the first cost of equipping the
laboratory, many of the reagents, enumerated as belonging to the student's
desk, may be placed on the side-table. Many good laboratories are thus
arranged.
WORKING MATERIAL.
The substances enumerated under this heading are arranged in the same
order as the Elements and their compounds in the text, and none are
repeated. It is not necessary in every case that the chemicals which fol-
low should be chemically pure. The reagents, etc., already named are not
given.
83. Introduction. Galena; iron filings ; flowers of sulphur.
84. Oxygen. Mercuric oxide; red lead; manganese dioxide (C. P.) ;
bark charcoal ; iron wire ; broken watch-springs ; phosphorus ; zinc foil ;
pyrogallic acid.
85. Hydrogen. Metallic sodium and potassium ; mercury ; well-water ;
barium dioxide.
WORKING MATERIAL. 361
NOTE. For generating large quantities of hydrogen when purity is not
especially requisite, sheet zinc may be employed; this is cut into bits, and
to help the action along a few nails may be thrown into the generator.
86. Nitrogen. Quicklime ; ammonium chloride ; ammonium nitrate
(C. P.); copper filings; potassium nitrate; spirits of turpentine.
87. Chlorine. Indigo solution ; sodium chloride.
88. Bromine. Potassium bromide ; bromine.
89. Iodine. Iodine.
90. Fluorine. Calcium fluoride ; beeswax, or paraffin.
91. Carbon. Lampblack ; graphite ; various kinds of coal ; bone-
black ; sugar; sodium acetate; yeast; calcium carbonate; magnesium
ribbon ; clam shells, snail shells, corals, and other carbonates.
92. Sulphur. Roll sulphur; iron pyrites.
93. Silicon. As many varieties of silicon dioxide as possible.
94. Boron. Boric acid.
95. Phosphorus Stick phosphorus ; red phosphorus.
When working with the metals, it is desirable to have as many ores of
each metal as possible ; not that these ores are absolutely indispensable to
the work in the text, but because of the advantage the student may derive
from their examination or from working with them.
96. The First Group Metals. Metallic silver and ores of silver;
metallic mercury and ores of mercury; metallic lead in its commercial
forms, and ores of lead.
97. Second Group Metals. Arsenic and arsenic trioxide ; antimony,
antimony sulphide, and ores of antimony ; metallic tin in its commercial
forms, and ores of tin ; metallic bismuth and ores of bismuth ; sheet copper,
native copper, and ores of copper ; metallic cadmium, ores of cadmium.
98. The Third Group Metals. Iron in its commercial form and ores
of iron ; chrome alum or other chromium salts ; metallic aluminum and
as many commonly occurring aluminum compounds as possible ; metallic
nickel and ores of nickel ; cobalt ores ; manganese ores ; commercial forms
of metallic zinc.
99. The Fourth Group Metals. Barium dioxide, hydroxide, and as
many barium-bearing minerals as possible ; strontium nitrate ; many cal-
cium bearing minerals ; metallic magnesium ribbon, and many magnesium-
bearing minerals.
The Fifth Group Metals are already provided for.
362
APPENDIX.
GENERAL APPARATUS.
Under this heading is included that apparatus which is of general
utility. The teacher may need some of it for special purposes, while some
of it is so placed that the students may have access to it at any time.
Much of this apparatus may be used in physics also.
100. A Becker or Troemner Balance, Fig. 23, is to be recommended
on account of its cheapness, neatness, accuracy (sensitive to 2ms), and
durability. By placing a small shelf or table over one pan, so that the
FIG. 23.
balance may play freely, it will answer well for specific gravity. The
author's students have used this balance for three years, and it is still as
good as new. Accompanying it is a set of weights in a polished, velvet-
lined box, with forceps, and a tray divided into compartments for the
small weights, and covered with a glass slide. These weights were im-
ported at a cost of $3.50 ; they run from 50s to Is in brass and 500mss to
lm& in platinum.
1O1. A Pair of Counter-Poised Watch-Crystals are useful in
weighing those substances which would attack the pans of the balance.
GENERAL APPARATUS. 363
102. A Weighing Flask for iodine and other volatile substances is
desirable.
103. A Specific Gravity Bottle of 50CC capacity is useful in deter-
mining the specific gravity of fluids.
104. A Pair of Hydrometers. One for fluids lighter than water, and
one for fluids heavier than water.
105. A Pair of Good Centigrade Chemical Thermometers. One
graduated from —20° to + 240°, and one from — 10° to + 360°.
106. Graduated Flasks. One I1, one £i , and one J1.
These are fitted with glass stoppers, and bear only one mark around the
neck. These are useful when I1, etc., is wanted quickly.
107. Litre Cylinder for mixing reagent solutions. These are gradu-
ated into cc's to read up and down.
108. Two Burettes, capacity 50CC each ; graduated to O.lcc. These
are used in titration.
109. A Pipette, capacity 5CC, graduated to O.lcc. Used for taking out
small quantities of liquids from bottles, etc.
110. A Lipped Graduated Jar, capacity 100CC, graduated to lcc.
Used in measuring out liquids.
111. Ure's Eudiometer. This is shown and explained in Fig. 7.
112. Hofinann's Apparatus, as shown and explained in Fig. 3.
113. Spectroscope. Spectroscopes are now to be had quite reason-
ably. The needs of the school should determine the expense of the instru-
ment purchased.
114. Bell Jars are used in experimenting with gases. Those used in
connection with the air-pump may be employed, or large bottles may be
cut off at the bottom. This may be accomplished by cutting a crease
around the bottle with a three-cornered file ; this crease is then followed
up with a minute blow-pipe flame until the bottom cracks off. The edges
may then be ground smooth on a sheet of emery-paper stretched on a flat
board.
115. Large Beakers, Funnels, Evaporating Dishes, and Ring
Stands similar to those shown in the Frontispiece, only larger, are found
useful in preparing solutions, reagents, etc.
116. Retorts and Receivers, similar to those shown in Fig. 14, are
used in distillation, etc.
364 APPENDIX.
117. A Ldebig's Condenser is often used in connection with the
retorts.
118. Tall Jars are useful in experimenting with gases.
119. An Iron Mortar and a Porcelain or a Wedgewood-Ware
Mortar, with pestles
120. Assorted Glass Tubing of various sizes suitable for " hydrogen
tones," connections, etc.
121. Funnel Tubes for Generators, as shown in Fig. 5.
122. Blast-Lamp, for alcohol or gas, is useful in working glass.
123. A Copper Oxygen Retort, for generating oxygen. An iron
retort may be used, or a common glass generating flask will serve the
same purpose.
124. Mercury Trough of Porcelain.
125. A Hydrogen Pistol may be made from a gas-pipe 1| in. in
diameter, and 6 in. long. One end is closed with a cap ; a small opening
is drilled in for a vent, and the mouth is closed with a common cork.
126. A Pneumatic Trough. There are many designs in use. As
a general rule, the simpler the trough, the better.
127. Gas Holders. Any tinsmith can make very satisfactory gas
holders. Or they can be made from a barrel, and a cask that will go
inside the barrel. The heads are removed; the barrel is filled with water,
and the cask is inserted in the barrel and suitably weighted ; a stop-cock,
for attaching rubber-hose, is inserted in the head of the cask.
128. Chemical Corks and Rubber Stoppers of assorted sizes.
129. Rubber Tubing of assorted sizes, for connections, etc.
130. Rubber Gas-Bags. One of 2 gals., and one of 1 gal. capacity.
131. Oxyhydrogen Blow-Pipe. One form of this apparatus is
shown in Fig. 8. Prof. Weitbrecht has constructed a cheap instrument
from J-in. gas fixtures. The instrument is T "shaped ; into the stem of the
T is screwed a Springfield musket cap-nipple which serves as a jet; in
each arm of the T is a stop-cock. The hydrogen is admitted into one arm
and the oxygen into the other. Illuminating gas may be used in place of
hydrogen.
132. A Furnace, known as the Fletcher Furnace, and provided with
bellows and a blast-jet for illuminating gas, is not expensive, and will fuse
such metals as gold, silver, etc.
THE LIBRARY. 365
133. Crucibles. Hessian crucibles and plumbago crucibles are used.
The sand, or Hessian crucible, is inexpensive, and may be bought in nests.
THE LIBRARY.
A reference library should be kept in the laboratory. It should be
easy of access, and the students should be permitted to make use of any
book at any time. Books should not be taken out of the laboratory. In
the following list no attempt at completeness is made ; a few good books
that are within the reach of all schools are named. Roscoe and Schor-
lemmer's " General Treatise " will be found useful for general descriptive
work.
Douglas and Prescott's " Qualitative Analysis," or a standard edition
of Fresenius's " Qualitative Analysis," will be useful in qualitative work.
Sutton's "Volumetric Analysis" is recommended for methods of
titration.
Fresenius's " Quantitative Analysis " is useful, if quantitative work is
attempted.
Elderhorst's " Blow-Pipe Analysis " is to be used in expanding any
work with the blow-pipe.
Wanklyn's " Analysis of Water, Milk, and Air," may be used in case it
is desired to do work in that direction. These books are published in
separate volumes.
Dana's " Mineralogy " is valuable as affording information concerning
ores, coal, etc.
Gore's " Electro-Metallurgy " will afford information in that direction.
Some good work on Spectrum Analysis is desirable. Schellen, though
popularly written, is good. Roscoe's work is more technical.
The " U. S. Dispensatory," and the " Pharmacopoeia" are often useful.
Bailey's " Chemist's Pocket-Book " contains many valuable data for com-
putations, conversions, etc., etc.
One or two chemical journals, as " The Chemical News " and the
" American Journal of Chemistry," will serve to create an interest, by
calling the student's attention to the present tendencies of the science.
In response to numerous inquiries from teachers, concerning apparatus,
etc., the author would take this occasion to say that he will gladly give
any information in his power concerning the same ; and, in case any
school wishes aid in purchasing, that he has, made arrangements with
Messrs. Eberbach and Son, Ann Arbor, Mich., whereby any apparatus or
chemicals necessary for this text can be supplied promptly, and at the
366
APPENDIX.
lowest market price for the high grade of goods recommended. All cor-
respondence on this subject should be addressed to the author. A priced
list will be sent on application.
DATA FOR CONVERTING METRIC AND ENGLISH WEIGHTS AND
MEASURES.
1mm = 0.0394 in.
1cm _ 0.3937 in.
1 in. = 2.539954cm.
I cu. in.= 16.386176<*.
lcc = 0.06103 cu. in.
II = 61.02709 cu. in.
18= 15.43235 grains.
1 grain — 0.0648s.
1 Ib. avoirdupois = 453.59s.
1 oz. avoirdupois = 28.34954s.
1 gal. U. S. = 231. cu. in.
1 gal. Imp. — 277j cu. in.
INDEX.
[The numbers refer to pages.]
Acetic acid 340
Acetylene 134
Acid, Antimonic 249
Boric 191
Bromic 113
Chloric 104
Chlorous 104
Citric 340
Fuming sulphuric 174
Hydriodic 117
Hydrobromic 110
Hydrochloric 97
Hydrocyanic 146
Hydrofluoric 122
Hydrofluosilicic 190
Hypobromous 112
Hypochlorous 101
Hyponitrous 65
Hypophosphorous 199
lodic 120
Manganic 296
Meta-phosphoric 202
Meta-stannic 253
Nitric 67
Nitrous 66
Nordhausen 174
Orthophosphoric 201
Oxalic 340
Perchloric 105
Permanganic 296
Phosphoric 201
Acid, Phosphorous 200
Prussic 146
Pyrophosphoric 202
Acids, Basicity of 168
defined 75
General examination for,
204, 339
Acid salt 217
Acid, Selenic 179
Selenious 179
Sodium carbonate 333
Stannic 253
Sulphuric 169
Sulphurous 168
Tartaric 340
Telluric 180
Tellurous 180
Thiosulphuric 175
Agate 186
Agricola 4
Albite 286
Alchemy .... 2
Alkali plains 327
Alloys 211
Alum 321
Alumina 286
Aluminum hydroxide 287
Aluminum, Occurrence 286
Preparation 286
Properties 287
Compounds 287
368
INDEX.
Aluminum, Tests . . . : 288
sulphate 287
Alums 287
Amalgams 211
Amethyst 18G
Amorphous iron ore 276
Ammonia, Albuminoid, Estima-
tion of 49
Ammonia, Estimation of 72
in drinking-water 45
Occurrence 52
Preparation 52
Process 332
Properties 55
Tests 58
Ammonium 334
carbonate 335
chloride 334
molybdate 272
nitrate 334
phospho-molybdate 272
sulphide 335
Analysis denned 216
of unknown substances, 337-340
Ancients, Chemistry of 1
Ancient copper-miners 259
Anhydrite 314
Anthracite 129
Antimonic acid ... 249
Antimony black 247, 272
Antimony, Butter of 249
Occurrence 247
Preparation 247
Properties 248
Compounds 249
Tests 250
oxides 249
trichloride 249
trisulphide 249
Apatite 193
Arabs, Chemistry of 2
Argentite 229
Argillaceous iron ore 276
Aristotle, Doctrines of 2
Arsenic, Occurrence 242
Preparation 242
Properties 243
Compounds 244
Tests 246
pentoxide 245
trioxide 242, 244
Arsenious sulphide 245
Arseniuretted hydrogen 244
Asbestos 316
Atomic heat 209
theory 15
weights, Determination of,
15, 16, 152
Atmosphere 82
Estimation of oxygen and
nitrogen of 85
Impurities of 84
Temperature of 83
Avogadro's hypothesis 150
Azurite 258
Boracite 190
Barium carbonate 311
chloride 311
hydroxide 310
iodate 311
monoxide 310
nitrate 311
Barium, Occurrence 310
Preparation 310
Compounds 310
Tests 311
sulphate 311
Barometer 82
Baryta, Caustic 310
water 310
Bases denned . . 77
IXDKX.
369
Basic salt
218
Bronze
. 211
Bauxite
28(5
Brown haematite
. . . 276
Bell metal
211
Bunsen hurner
... 28
Beryl
.. 288
B'crylliuin
302
Cadmium iodide
263
Bessamer process
279
Cadmium Occurrence
262
Bismuthite
255 257
Preparation
262
Bismuth nitrate ...
257
Properties
262
oclire
255
Compounds
263
Bismuth Occurrence
255
Tests
263
Preparation
255
Sulphate
... 263
256
Sulphide
263
257
Caesium
336
Tests
257
Calcium chloride
315
257
carbonate
. . . 315
subnitrate
. ... 257
hydroxide
. ... 314
Binary compounds
73
Occurrence
313
129
Preparation
. ... 314
276
Properties »
... 314
Borax
.. 190
Compounds
.... 314
Boron Occurrence
190
Tests
315
Preparation
190
sulphate
313, 314
Tests
191
Calc spar
. ..313
276
Carbonado
127
Boyle
4
Carbon bisulphide
178
Black lead
125
dioxide, Estimation of .
148
Blanc de fard . ...
257
dioxide, Occurrence . . .
. ... 138
257
Preparation
138
Blast furnace
277
dioxide, Properties . . . .
140
315
Tests
145
Brass
211
hydrides
.... 132
Braunite
295
monoxide, Preparation.
.... 136
290
Tests
138
157
Occurrence
125
Britannia
211
Preparation ....
.... 126
Bromic acid
113
Properties
.... 126
112
Tests
.... 131
108
oxides
136
108
Carnallite
316, 321
110
Carre ice machine
.... 57
Tests .
. 110
Cassiterite
. 251
370
INDEX.
Cast iron , 278
Caustic potash 322
soda . 328
Cavendish 5
Celestine 312
Cerium 304
Chalcedony 186
Chalk 314
Chameleon mineral 297
Charcoal 126
Chemical reaction . 11
Chemism 12
Chemistry denned 11
Chemistry, Derivation of 1
Chert 186
Chili saltpetre 327
China clay 286
Chloric acid 104, 105
Chlorine, Estimation of 107
in drinking-water 45
Occurrence 92
Preparation 92
Properties 95
Tests 96
oxacids 101
oxides 99
monoxide 99
trioxide 100
tetroxide 100
Chlorous acid 104
Choke damp 142
Chrome alum 283
Iron stone 282
yellow 228, 284
Chromium hydroxide 283
Occurrence 282
Preparation 282
Properties 283
Compounds 283
Tests 285
oxides.. . 283
Chrysophrase 186
Cinnabar 234, 235
Citric acid 340
Claus 270
Coal 125, 128
Coal analysis 147
Cobalt, Occurrence 292
Preparation 293
Properties 293
Compounds 293
Tests 294
glance 292
Cobaltous chloride 294
nitrate 294
sulphate 294
Cobalt ultramarine 294
Coin, gold, silver, bronze 211
Coke 129
Columbite 307
Combining number 14
Combustion 27
Spontaneous 29
Compounds 10
Conductivity 210
Condy's disinfecting liquid 297
Conglomerates 188
Copperas 280
Copper glance 258
nitrate 260
Copper, Occurrence 258
Preparation 259
Properties 259
Compounds 260
Tests 261
oxides 261
pyrites 258
sulphate 260
sulphides 260
Corundum 286
Crocoisite . . . 282
INDEX.
871
Cuprite . .
Cyanogen
258
145
Dalton 6
Davy 6
Dialysis 187
Diamonds 125, 126, 127
Didymium 304
Disodium phosphate 329
Dog-tooth spar 314
Dolomite 316
Dulong and Petit's law 209
Dutch liquid 134
Ekaluminum 303
Egyptians, Chemistry of 1
Elements defined . . 9
Elements, Names of 17
Classification of 219-222
Table of 20
Elixir Vitae 4
Epsom salts 317
Equations, Atomic and molecular, 155
Meaning of 36
Writing of 81
Emery 286
Erbium 305
Etching on glass 122
Ethylene 133
Test for 134
Experiment defined 8
Fat 321
Feldspar 286
Fermentation 139
Ferric chloride 280
hydroxide 280
Ferrous chloride 280
sulphate 280
sulphide 281
Fire ., 27
Fire damp 132
Fixed alkalis 320
Flint 186
Flowers of sulphur 158
Fluorine 122
Fluorspar 315
Fly-powder 243
Fool's gold 281
Formula . . 18
Franklinite 298
Fresenius's analytical classifica-
tion of the metals 337
Fusible metal 211
Galena 224, 227
Gallium 303
Gas carbon 129
Geber 2
German silver 211
Glass 333
Glucinum 302
Gold 266
Granular iron ore 276
Grape iron ore 276
Graphite 125, 126, 128
Greenockite 263
Green vitriol 280
Guignet's green 283
Gunpowder 325
Gypsum 314, 315
Hardness of water 47, 316
Estimation of 49
Haematite 276
Hausmannite 295
Heavy spar 310, 311
Hone stone 186
Horn silver 229
Hydriodic acid 117, 118
Hydrobromic acid 110-112
Hydrochloric acid, Occurrence. . 97
372
INDEX.
Hydrochloric acid, Preparation . . 97
Properties 97
Tests 99
Hydrofluoric acid, Preparation . . 122
Properties 122
Tests 123
Hydrofluosilicic acid 190
Hydrogen arsenide 244
Hydrogen dioxide 47, 48
Hydrogen, Occurrence 34
Preparation 34
Properties 38
Test 40
persulphide 163
phosphides 197
selenide 178
stibide 250
sulphide, Estimation of .... 181
sulphide, Occurrence 160
Preparation 161
Properties 162
Tests 163
telluride 180
Hydroxyl 71
Hydroxylamine 70
Hypobromous acid 112
Hypochlorous acid 101-103
Hyponitrites 65
Hyponitrous acid 65
Hypophoshorous acid 199, 200
Iceland spar 314, 316
Illuminating gas 135
Indium 303
lodic acid 121
Iodine, Occurrence 115
Preparation 116
Properties 116
Tests 117
oxacids 120
oxides . 120
Iridium 270
Iron arsenide 242
Iron, Occurrence 275
Preparation 276
Properties 279
Compounds 280
Tests 281
oxides 280
pyrites 281
Jet
129
Kaolin 286
Kelp 115
Kieserite 316
Kupfer-nickel 290
Lac sulphuris 158
Lampblack 126
Lanthanite 304
Lanthanum 304
Lapis lazuli 288
Laughing gas 60
Law of definite proportions 12
of multiple proportions .... 13
Lavoisier 6
Lead chloride 228
chromate 228, 284
Lead, Occurrence 224
Preparation 224
Properties 226
Compounds 227
Tests 228
Lepidolite 335
Libavius 4
Lignite 129
Limestone 314
Lithium 335, 336
Liquation 251
Lodestone 276
Lunar caustic. . . 232
INDEX.
373
Magnesia 317
Magnesite 316, 317
Magnesium carbonate 317
chloride 92, 317
limestone 316
Magnesium, Occurrence 316
Preparation 316
Properties 316
Compounds 317
Tests 317
Magnetite 276
Manganese acids 296
Manganite 295
Manganese, Occurrence 295
Preparation , 295
Properties 295
Compounds 296, 297
Tests 298
oxides 296
sulphides 297
Marble 314
Massicot 227
Matter 12
Mechanical mixture 11
Meerschaum ... 316
Melting-points 210
Mendelejeffs classification . .220, 221
Mercuric chloride 235
Mercurous chloride 235
nitrate 236
Mercury, Occurrence 233
Preparation 234
Properties 235
Compounds 235
Tests 237
Red oxide of 235
Metal, Analytical classification
of 212, 216
Metal defined 208
Metals of the alkalies 320
Metals, Salts of 216
Metaphosphoric acid 202
Metastannic acid 253
Methane 132
Meteorites .... 275
Micaceous iron ore 276
Microcosmic salt 335
Mispickel 242
Molecules 149
Molecular heat 211
Molecular weight, Determination
of 151, 152
Molybdenite 272
Molybdenum 272
Nickel ammonium sulphate 291
arsenide 242
blende 290
glance 290
Nickel, Occurence 290
Preparation 290
Properties 291
Compounds 291
Tests 291
Nickel oxides 291, 293
sulphate 291
sulphide 291
Niobium 307
Nitre 324
Nitric acid 67-69
Nitrites 66
in drinking-water 46
Nitrogen chloride 106
dioxide 61, 62
monoxide 59-61
Nitrogen, Occurrence, etc 50, 51
Nitrogen oxacids 65
oxides 58
pentoxide 64
tetroxide 64
trioxide 63
Nitrous acid . . 66, 67
INDEX.
Nitrous oxide «. 59
Novalculite 186
Odontolite 288
Oil of vitriol 172
Opal 186
Oriental amethyst 286
emerald 286
topaz 286
Orpiment 242
Orthoclase 189, 321
Orthophosphoric acid 201
Osmium 271
Oxalic acid 340
Oxidizing-flame 28
Oxygen, Occurrence 23
Preparation 23, 24
Properties 25
Tests 30
Oxy-hydrogen blow-pipe 42
Ozone '. 31, 32
Palladium 269
Paracelsus 4
Peat 129
Perchloric acid 105, 106
Pewter 211
Philosopher's stone 3
Phlogiston 5
Phosphate of aluminum 287
Phosphates, Tests for 203
Phosphoric acid 201
Phosphorite 193
Phosphorous acid 200, 201
Phosphorus, Occurrence 193
Preparation 193
Properties 195
Tests 196
Phorphorus oxacids 199
oxides 198
pentoxide 198
Phorphorous trioxide 108
Pitch blende 306, 307
Plastic sulphur 159
Platinum 268
Plumbago 125
Pneumatic chemistry 5
Potash 325
Potassium bichromate 284
bromide 323
carbonate .... 325
chlorate 323
chloride 323
chromate 283
chromium sulphate 283
cyanide 325
ferrocyanide 281
hydroxide 322
iodide 323
Potassium, Occurrence 321
Preparation 321
Properties 322
Compounds 322
Tests 326
Potassium permanganate 297
sulphate 323
Priestley 5
Prussic acid 146
Pyrochlor 307
Pyrolusite 295
Pyrophosphoric acid 202
Quartz 186
Quartzite 186
Queen's metal 211
Quicklime 314
Realgar 242
Red precipitate 235
Reducing-flame 29
Rhodium 271
Rhodocrosite ., .295
INDEX.
375
llinmann's green L".) I
Rose's metal 211
Rubidium 336
Ruby 286
silver 229
Ruthenium 270
Rutile 305
Safety-lamp 132
Salt-cake process 330
Saltpetre 321, 324
Sal sodae 333
Salts, Acid and normal 80
defined 77
Sand 186
Sandstone 188
Sapphire 286
Scale of hardness 127
Scheele 5
Schweinfurth's green 245
Scheele's green 5, 245
Selenite 314
Selenic acid 179
Selenious acid 179
Selenium dioxide 178
Selenium, Occurrence 177
Preparation 178
Properties 178
Tests 179
Separation of arsenic, antimony,
and tin 254
of bismuth, copper, and cad-
mium 263
of chlorides and bromides. . 114
of chlorides, bromides, and
iodides 119
of copper and bismuth 261
of cobalt, manganese, nickel,
and zinc 300
of first group metals 238
Separation of first and second
group metals 265
of fourth group metals 318
of iron, chromium, and alu-
minum 289
of nickel and cobalt . . . 295, 302
of second group metals .... 264
Serpentine 189
Siderite 276
Silica 184, 188
Silicates 188, 189
of cobalt 294
Siliceous springs 187
Silicon fluoride 190
hydride , 189
Silicon, Occurrence, etc. . . . 184, 185
Silver bromide 233
chloride 92, 232
copper glance 229
iodide 115
nitrate 232
Silver, Occurrence 228
Preparation 229
Properties 231
Compounds 232
Tests 233
plating solution 232
Skutterrudite 292
Slaked lime 314
Smalt 294
Soda-ash 330
process 331
Soda crystals 333
Sodium aluminate 287
Sodium ammonium phosphate. 335
arsenate 245
carbonate 330
chloride 92, 328
hydroxide ... 328
hyposulphite 329
hypophosphite 329
376
INDEX.
Sodium nitrate ..............
Sodium, Occurrence ..........
Preparation ..............
Properties ........ ____
Compounds ...............
Tests ....................
silicates ..................
thiosulphate ..............
Solder .......................
Sombrerite ...................
Soot .........................
Spathic iron ore ..............
Specific heat ..................
Spectra ......................
Specular ore ..................
Speculum metal ..............
Speiss cobalt ............. 292,
Spinelle .....................
Spirits of hartshorn ...........
Stannic acid .................
sulphide. .". ...............
Stannous sulphide .............
chloride ..................
Steam, latent heat of .........
Steatite ......................
Steel .........................
Stibnite .................. 247,
Stream tin ....................
Strontianite ...................
Strontium carbonate ...........
nitrate ...................
Strontium, Occurrence, etc. 312,
Substituting power and valence.
Suint ................... 321,
Sulphur acids, Test* for ........
Sulphur dioxide, Occurrence,
etc ................. 164,
Sulphuretted hydrogen ........
Sulphuric acid, Estimation of . . .
fuming ...................
Hydrate of ...............
^ulphuric acid, Occurrence 169
327
32S
333
333
329
211
193
126
276
209
303
276
211
293
316
55
253
253
253
253
44
188
279
249
251
312
312
312
313
154
325
176
166
160
182
174
173
Preparation ............... 170
Properties ............... 172
Tests .................... 173
Sulphur, Occurrence ........... 107
Preparation ............... 107
Properties ................ 108
Tests .................. 160
Sulphur oxacids ............... 167
oxides ................... 164
Sulphurous acid ........... 168, 169
Sulphur trioxide .............. 167
Superphosphate of lime ........ 315
Sylrite ....................... 321
Symbols, Chemical ............ 17
Talc
316
Tantalite .................... 307
Tantalum .................... 307
Tartar emetic ................. 249
Tartaric acid .................. 340
Tellurium acids ............... 180
dioxide ................... 180
Tellurium, Occurrence ......... 179
Preparation, etc ........... 180
Tellurium trioxide ............. 180
Terbium ..................... 305
Thenard's blue ................ 294
Thermometers ................ 83
Thiosulphuric ............... 170
Thorite ...................... 305
Thorium . .................... 305
Tin foil .....................
Tin, Occurrence, etc ....... 251-253
Tin stone ..................... 201
Titanite .....................
Titanium .....................
Titanium cyano-nitride ......... 306
Topaz ........................ 288
Triphylline ............... 335, 336
Tripoli ...................... 188
INDEX.
S77
Tungsten 271
Turquois 287
Type metal 211
Ultra marine 288
Uranium 300
Useful problems 89
Valence 153
Vanadium 307
bronze 308
Van Helmont 4
Vivianite 193
Volatile alkali 320
Volume of a gas affected by heat, 86
by pressure 85
Formulae for computing the, 91
Water 40-43
drinking, Impurities of . . . 45-47
Water-lime 314
Weight and density 88
White lead *. 227
Willemite 298
Witherite 310, 311
Wiihler 6
Wood's alloy 211
Wolfram 271
Wollaston 271
Wollastonite 188
Wrought iron 279
Yttrium 304
Yttrotantalito 307
Zinc blende 298
Zinc, Occurrence, etc 298-300
Zircon 306
Zirconium . . .306
BERWICK A SMITH, PRINTERS, BOSTON.
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Guides for Science -Teaching.
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