■I
J
COLLEGE OF LIBERAL ARTS.
Boston University*
From
^- -^^
T{eceived
4o\P.
EXPERIMENTAL
CHEMISTRY
BY
LYMAN C. NEWELL, Ph.D. (Johns Hopkins)
PROFESSOR OF CHEMISTRY, BOSTON UNIVERSITY,
BOSTON, MASS.
'"^ Library
College of Liberal Arts
Boston University
D. C. HEATH & CO., PUBLISHERS
1909
Copyright, 1900
By D. C. Heath & Co.
Poo
PREFACE.
The purpose of this book is to promote the more efficient
teaching of chemistry by modern methods. The choice and
arrangement of subject-matter is based on the author's extended
experience with students of varied abihty. The book as a whole
is the outcome of a desire to provide a course in chemistry which
shall be a judicious combination of the mductive and deductive
methods.
The concise directions and specific questions in the experiments
aim to obviate the difficulties and perplexities which hinder the
progress of beginners in chemistry. The experiments are repre-
sentative. No important ones have been omitted, and no senseless
or dangerous ones have been included. Every experiment need
not be done by each pupil. Indeed, many experiments may be
advantageously performed by the teacher either on the lecture
table or in the laboratory ; others may be done by the more
skilful students and the results utilized in the class-room. Abun-
dant material has been purposely provided to meet varied condi-
tions. The Teacher's Supplement suggests different uses of many
experiments, and the author invites correspondence regarding the
adaptation of the book to special demands.
Two supplementary sets of questions arise from most completed
experiments.' One class may and often must be answered in the
laboratory where necessary data are available. The Laboratory
Exercises aim to meet this demand. The second class of ques-
tions concerns the vast amount of descriptive and historical
matter indirectly connected with the experiments ; this material
is needed for a broad knowledge of the subject, but its presence
iv Preface.
would often destroy the real value of the experiments by pre-
cluding original thought on the part of the student. Hence, at
appropriate points, Class-room Exercises are inserted. They
are not to be considered in the laboratory, but may be used as
the basis of lectures, recitations, examinations, quizzes, or reviews.
Both kinds of exercises may be omitted in whole or in part, if the
teacher so desires. The Bibliography contains titles of books
helpful for a discussion of these exercises, and additional infor-
mation may be found in the Teacher's Supplement. Despite the
variety and amount of matter presented in these exercises, the
ordinary text-books will meet the student's needs in most cases
except those involving the recent discoveries and appUcations.
A judicious selection and distribution of the points actually
needed will obviate the immediate necessity of an extensive
library. Facts not accessible to students may be easily supplied
by the teacher. It should not be forgotten, however, that refer-
ence books are indispensable adjuncts to laboratory work.
Opinions differ widely regarding the kind of experiments which
should constitute a course in chemistry. There is a growing
belief, however, that selected experiments involving accurate
measurements possess fundamental disciplinary value. This book
contains several quantitative experiments, most of which have
been repeatedly performed in the author's classes under a variety
of conditions. The details have been worked out with the
greatest care. It will be noticed that each quantitative experi-
ment illustrates one or more of the fundamental principles of
chemistry, hence the results permit conclusions which furnish a
sound and attractive basis for the treatment of chemical theory.
Teachers are earnesdy urged to give these experiments a trial and
to utilize the results in class work.
An unusually large number of problems has been intentionally
incorporated to allow a generous distribution among several classes
or sections. It is not intended that each student shall solve all
the problems. The author will be grateful for information of
errors or ambiguity detected in these problems. Only those who
Preface. v
have prepared similar work for publication can realize the count-
less sources of error. The Teacher's Supplement contains answers
to all problems.
The Appendixes contain directions for manipulation which could
not be conveniently incorporated in the text, a complete treat-
ment of the subject of gases, and miscellaneous data not easily
found elsewhere.
Material intended exclusively for teachers has been put in a
Teacher's Supplement. It is believed that the contents and the
manner of its presentation will not only lessen their labors, but
guide them in the choice of experiments, apparatus, etc. Copies
can be obtained only by ordering directly of the publishers.
The entire manuscript has been read by —
Mr. Fred L. Bardwell, Assistant Professor of Chemistry, Insti-
tute of Technology, Boston, Mass.
Dr. William B. Shober, Instructor in Organic Chemistry, Lehigh
University, Bethlehem, Pa.
Mr. Irving O. Palmer, Instructor in Chemistry and Physics, High
School, Newton, Mass.
The proof has been read by —
Dr. Edward C. Franklin, Professor of Chemistry, University
of Kansas, Lawrence, Kansas.
Mr. F. F. Coburn, Principal of the State Normal School, Lowell,
Mass.
Dr. Charles A. Pitkin, Professor of Chemistry, Thayer Academy,
South Braintree, Mass., and Professor of General Chemistry, Tufts
College Medical School, Boston, Mass.
Mr. William H. Snyder, Master of Science, Worcester Academy,
Worcester, Mass.
Dr. A. M. Muckenfuss, Professor of Chemistry, Millsaps College,
Jackson, Miss.
Dr. H. G. Shaw, Instructor in Chemistry, High School, Melrose,
Mass.
Dr. A. J- Hopkins, Assistant Professor of Chemistry, Amherst
College, Amherst, Mass.
vi Preface.
Dr. J. E. Bucher, Professor of Chemistry, College of Agriculture
and Mechanic Arts, Kingston, R. I.
Dr. C. E. Boynton, Instructor in Chemistry, North Division
High School, Chicago, 111.
Dr. M. D. Sohon, Instructor in Physics, Mixed High School,
New York.
Dr. Milo S. Walker, Instructor in Chemistry, West Division
High School, Chicago, 111.
Mr. Ernest A. Congdon, Professor of Chemistry, Drexel Insti-
tute, Philadelphia, Pa.
Mr. George W. Benton, Instructor in Chemistry, High School,
Indianapolis, Ind.
Dr. J. Elliott Gilpin, Associate in Chemistry, Johns Hopkins
University, Baltimore, Md.
The chapter on "Some Common Organic Compounds" was
read in manuscript by Dr. James F. Norris, Instructor in Organic
Chemistry, Institute of Technology, Boston, Mass.
The author is profoundly grateful to these teachers for their
helpful and intelligent criticism, and he gladly takes this oppor-
tunity to express his appreciation of their kindness.
Numerous courtesies have been received from the L. E. Knott
Apparatus Co., and for these the author hereby tenders his thanks.
To his former pupils and colleagues in the English High School
at Somerville, Mass., the author is under obligation for generous
assistance.
Suggestions for a short course may be found on page 398.
It is believed that the intelligent and careful examination of
both manuscript and proof by the above teachers has produced a
book relatively free from errors ; nevertheless, for such errors as
may be found the author willingly assumes responsibility.
L. C. N,
Lowell, Mass.,
June, 1900.
CONTENTS.
Introduction *
Bunsen Burner. — Cutting Glass. — Bending Glass.—
Heating. — Filtration. — Smelling and Tasting. — Cleanli-
ness.—Metric System. — Balance. — Form of Record of
Experiments.
Experiments i-8. — Class-room Exercise I. — Problems
I, II.
CHAPTER I.
Physical and Chemical Changes 15
Introduction. — Chemistry. — Preliminary Definitions. —
Changes in Metals when Heated. — Products from the heated
Red Powder. — Identification of the Gas. — Definitions.
Experiments 9-13- — Laboratory Exercise I. —Class-room
Exercise V.
CHAPTER II.
Oxygen .24
Introduction. — Preparation and Properties. — Chemical
Compounds. — Combination of Oxygen with Magnesium. —
Discussion of Experiment 15. — Relation of Oxygen to Po-
tassium Chlorate. — Discussion of Experiment 16. — Weight
of a Liter of Oxygen. — Definitions. — Oxidation. — Law of
Definite Proportions by Weight. — Compounds and Mixtures.
Experiments 14-17. — Laboratory Exercise 11. — Class-
room Exercise VI. — Problems VII.
vii
viii Contents.
CHAPTER III.
PAGE
Hydrogen 38
Introduction. — Preparation and Properties. — Chemical
Reaction. — Tests. — Crystallization. — Crystallization of
Zinc S'j'phate. — Test for Zinc Sulphate. — Quantitative
Examination of the Reaction between Zinc and Sulphuric
Acid. — Discussion of Experiment 22. — Burning Hydrogen.
— Discussion of Experiment 23. — Weight of a Liter of
Hydrogen.
Experiments 18-23. — Laboratory Exercise III —Class-
room Exercise VII. — Problems VIII.
CHAPTER IV.
Symbols — Formulas— Conservation of Matter — Chemi-
cal Equations . ... -53
Symbols. — Formulas. — Conservation of Matter. — Chemi-
cal Equations — Quantitative Interpretation of Equations. —
Other Chemical Equations Studied. — Problems based on
Foregoing Equations.
Class-room Exercise VIII. — Problems IX.
CHAPTER V.
General Properties of Water 64
General Distribution. — Water of Crystallization. — Detec-
tion of Water of Crystallization. — Determination of Water
of Crystallization in Barium Chloride. — Definitions. — Efitlo-
rescence. — Deliquescence. — Impure Water. — Simple Tests
for Impurities in Water. — Purification of Water. — Distilla-
tion. — Relation of Water to Temperature. — Freezing and
Boiling Points of Water. — Solution. — Solutions of Gases.
— Solutions of Liquids. — Solubility of Liquids — Solutions
of Solids. — Solubility of Solids. — Saturation and Super-
saturation. — Supersaturated Solutions. — Thermal Phe-
nomena of Solution. — Solution and Chemical Action.
Experiments 24-37. — Class-room Exercises IX-XII. —
Problems X.
Contents. ix
CHAPTER VI.
PAGK
Composition of Water . 83
Introduction. —Electrolysis of Water. — Hydrogen and
Water. — Decomposition of Water by Iron. — Oxygen and
Water. — Decomposition of Water by Chlorine. — Interac-
tion of Sodium and Water. — Quantitative Composition of
Water. — Volumetric Composition of Water. — Gravimetric
Composition of Water. — Density of Steam. — Summary. —
Volumetric Composition of Steam. — Reduction.
Experiments 38-44. — Class-room Exercises XIII, XIV.
— Problems XI.
CHAPTER VII.
The Atmosphere i°5
Introduction. — Composition of Air. — Quantitative Exami-
nation of Air. — Other Constituents of the Atmosphere. —
Weight of a Liter of Air. — Weight of a Liter of Nitrogen.
Experiments 45-47- — Laboratory Exercises V, VI. —
Class-room Exercises XV-XVII. — Problems XII.
CHAPTER VIII.
Acids, Bases, and Salts 116
Introduction. — General Properties of Acids. — Nomen-
clature of Acids. — General Properties of Bases. — Nomen-
clature of Bases. — A General Property of Salts. — No-
menclature of Salts. — Testing the Nature of Substances.
— Nature of Common Substances. — Nature of Salts. — Neu-
tralization. — Quantitative Examination of the Interaction of
Acids and Bases. —Discussion of Experiments 52 and 53. —
Definitions.
Experiments 48-53. — Laboratory Exercise VII. — Class-
room Exercise XV 111.
CHAPTER IX.
Atoms, Molecules, and Related Subjects . . . . 132
Equivalents. — Equivalent of Magnesium and of Alu-
minium.—The Atomic Theory. —Atomic Weights. — Law
Contents.
PAGE
of Specific Heats. —Various Figures for Atomic Weiglits. —
Molecular Weights. — Avogadro's Law. — Calculation of
Atomic and Molecular Weights from Experimental Data. —
Valence. Percentage Composition. — Formula. — Elevation
of Boiling Point and Depression of Freezing Point. — Ions
and Ionization. — Applications of the Theory of Electrolytic
Dissociation.
Experiments 54-55. —Class-room Exercise XIX. — Prob-
lems XIII.
CHAPTER X.
Chlorine and Hydrochloric Acid i57
Chlorine. — Prepai'ation and Properties. — Bleaching by
Chlorine. — Weight of a Liter of Chlorine. — Hydrochloric
Acid. — Characteristic Properties. — Synthesis of Hydro-
chloric Acid Gas. — Preparation and Properties of Hydro-
chloric Acid. — Interaction of Sodium Chloride and Sulphuric
Acid. — Composition of Hydrochloric Acid Gas.— Discus-
sion of Experiment 63. — Formula of Hydrochloric Acid Gas.
Experiments 56-63. — Class-room Exercises XX, XXI. —
Problems XIV.
CHAPTER XL
Compounds of Nitrogen 172
Ammonia. — Preparation and Properties of Ammonia Gas
and Ammonium Hydroxide. — Composition of Ammonia
Gas. — Qualitative Composition of Ammonia Gas. — Quanti-
tative Examination of the Composition of Ammonia Gas. —
Discussion of Experiments 66 and 67. — Equations for the
Chemical Action in Experiments 66 and 67. — Volumetric
Composition of Ammonia Gas. — Discussion of Experiment
68. _ Weight of a Liter of Ammonia Gas. — Formula of
Ammonia Gas. - Law of Gay-Lussac. — Ammonium Com-
pounds. — Interaction of Ammonium Hydroxide and Hydro-
chloric Acid. — Preparation of Nitric Acid. — A Special
Property of Nitric Acid. — Test for Nitric Acid. — Interaction
of Sodium Nitrate and Sulphuric Acid. —Composition of
Nitric Acid. — Synthesis of Nitric Acid. —Analysis of Nitric
Contents. xi
PAGE
Acid. — Interaction of Nitric Acid and Metals. — General
Character of this Interaction. — Interaction of Nitric Acid
and Copper. — Discussion of Experiment ']']. — Nitrates. —
Action of Nitrates with Heat. —Action of Ammonium Nitrate
with Heat. — Special Test for Nitrates. — Law of Multiple
Proportions. —Aqua Regia. — Its Characteristic Property.
Experiments 64-81. — Laboratory Exercise VIII. — Class-
room Exercises XXII-XXV. — Problems XV-XVII.
CHAPTER XII.
Carbon and its Simplest Compounds 204
Introduction. — Distribution of Carbon. — Decolorizing
Action and Deodorizing Action of Charcoal. — Reduction by
Carbon. — Preparation and Properties of Carbon Dioxide.—
Synthesis of Carbon Dioxide. — Weight of a Liter of Carbon
Dioxide. — Carbon Dioxide and Combustion. — Carbonic
Acid. — Formation and Properties of Carbonates. — Acid
Calcium Carbonate. — Determination of Carbon Dioxide in a
Carbonate. —Carbon Monoxide. — Action of Hot Charcoal
on Carbon Dioxide. — Weight of a Liter of Carbon Monoxide.
— Preparation and Properties of Carbon Monoxide.
Experiments 82-94. — Laboratory Exercises IX-XII. —
Class-room Exercises XXVI-XXVHI. — Problems XVIII,
XIX.
CHAPTER XIII.
Compounds of Carbon and Hydrogen — Illuminating Gas
— Flames — Oxidation and Reduction . . . .221
Hydrocarbons. — Preparation and Properties of Methane.
— Preparation and Properties of Ethylene. — Preparation
and Properties of Acetylene. — Illuminating Gas. — Prepara-
tion and Properties of Illuminating Gas. — Combustion of
Illuminating Gas. — Bunsen Burner. — Construction of a
, Bunsen Burner. — Bunsen Burner Flame. — Candle Flame.
— Oxidizing and Reducing Flames. — Borax Bead.— Tests
with Borax Beads. — Blowpipe and its Use. — Reduction and
Oxidation with the Blowpipe.
xii Contents.
Experiments 95-106. — Laboratory Exercises XIII, XIV.
— Class-room Exercises XXIX-XXXIl. — Problems XX,
XXI.
CHAPTER XIV.
Some Common Organic Compounds 236
Introduction. — Composition of Organic Compounds.—
Hydrocarbons. — Alcohols. — Preparation of Alcohol. —
Properties of Alcohol. — Fermentation. — Formula of Alco-
hol. — Ether. — Properties of Ether. — Aldehyde. — Prepa-
ration and Properties of Aldehydes. — Acetic Acid.—
Properties of Acetic Acid. — Test for Acetic Acid. — Prop-
erties of Vinegar. — Preparation of Acetates. — Formula of
Acetic Acid. — Soap. — Preparation and Properties of Soap.
— Hardness of Water — Carbohydrates — Fehling's Test
for Sugar.
Experiments 107-118 — Laboratory Exercises XV-XX. —
Class-room Exercises XXXIll-XXXIX. — Problems XXII,
XXIII.
CHAPTER XV.
Bromine, Iodine, and Hydrofluoric Acid . . • .250
Introduction. — Preparation and Properties of Bromine. —
Preparation and Properties of Hydrobromic Acid. — Proper-
ties of Potassium Bromide. — Preparation and Properties of
Iodine. — Carbon Bisulphide Test for Iodine. — Starch Test
for Iodine. — Detection of Starch. — Preparation and Prop-
erties of Hydriodic Acid. — Properties of Potassium Iodide.
— Preparation and Properties of Hydrofluoric Acid.
Experiments 119-128. — Class-room Exercises XL, XLI.
— Problems XXIV, XXV.
CHAPTER XVI.
Sulphur and its Compounds 260
Physical Properties of Sulphur. — Amorphous Sulphur.—
Crystallized Sulphur. — Combining Power of Sulphur. —
Hydrogen Sulphide. — Preparation and Properties of Hydro-
Contents. xlii
gen Sulphide. — Sulphides. — Preparation and Properties of
Some Sulphides. — Separation of Sulphides. — Preparation
and Properties of Sulphur Dioxide. — Sulphurous Acid. —
Sulphur Trioxide. — Action of Sulphuric Acid with Water
and with Organic Matter. — Test for Sulphuric Acid.
Experiments 129-139. — Laboratory Exercises XXI-
XXIII. — Class-room Exercises XLII-XLV. — Problems
XXVI, XXVII.
CHAPTER XVII.
Silicon and Phosphorus and their Compounds — Review 273
Silicon. — Relation of Silicon Dioxide to Other Compounds
of Silicon. — Preparation and Properties of Silicic Acid. —
Discussion of Experiment 140. — Phosphorus. — Review.
Experiments 140-141. — Laboratory Exercise XXIV. — ■
Class-room Exercises XLVI-XLVIII. — Problems XXVIII,
XXIX.
CHAPTER XVIII.
Sodium — Potassium — Copper — Silver — Magnesium —
Calcium — Zinc — Mercury 279
Metals. — Properties of Sodium. — Compounds of Sodium
and their Properties. — Preparation of Pure Sodium Chloride.
— Preparation of Sodium Hydroxide. — Properties of Potas-
sium.— Compounds of Potassium and their Properties. —
Preparation of Potassium Hydroxide. — Preparation and
Properties of Potassium Carbonate. — General Properties of
Copper. — Tests for Copper. — Interaction of Copper and
Metals. — Important Compounds of Copper. — Oxides of
Copper. — Preparation and Properties of Cuprous Oxide.
— Alloys. — Preparation and Properties of Silver. — Tests
for Silver. — General Properties of Magnesium. — Tests for
Magnesium. — Tests for Calcium. — Compounds of Calcium.
— General Properties of Zinc. — Tests for Zinc. — Interaction
of Zinc and Metals. — General Properties of Mercury. —
Preparation of Mercury. — Definitions. — Tests for Mercury.
— Properties of Mercurous and Mercuric Compounds.
XIV
Contents.
PAGE
Experiments 142-166. — Laboratory Exercises XXV-
XXXIV. -- Class-room Exercises XLIX-LVI. — Problems
XXX-XXXVI.
CHAPTER XIX.
Aluminium Tin — Lead — Chromium — Manganese — Iron
— Separation of Metals .... -303
General Properties of Aluminium. — Action of Aluminmm
with Acids and Alkalies. — Preparation and Properties of
Aluminium Hydroxide. —Discussion of Experiment 169.—
Tests for Aluminium. — Preparation and Properties of Com-
mon Alum. — General Properties of Tin. — Action of Tin
with Acids. — Tests for Tin. — Discussion of Experiment
174. — Deposition of Metallic Tin. — General Properties of
Lead. — Tests for Lead. — Soluble Lead Compounds. —
Action of Lead on Water. — Properties of the Oxides of
Lead. — Tests for Chromium. —Chromium in its Compounds.
— Properties of Chromates. — Definition. — Chromic Com-
pounds.— Reduction of Chromates to Chromic Compounds.
— Properties of Chrome Alum. — Preparation and Properties
of Chromic Hydroxide. — Tests for Manganese. —Oxidation
with Potassium Permanganate. — General Properties of Iron.
— Ferrous and Ferric Compounds. — Behavior of Ferrous
Compounds. — Behavior of Ferric Compounds. — Reduction
of Ferric Compounds. — Oxidation of Ferrous Compounds. —
Separation of Metals. — Separation of Lead and Silver. —
Separation of Lead, Silver, and Mercury (-ous). — Analysis
of a Silver Coin. — Analysis of Brass. — Analysis of Solder.
— Separation of Copper, Iron, and Sodium. — Separation of
Zinc and Iron. — Separation of Iron and Aluminium. — Sepa-
ration of Zinc and Aluminium.
Experiments 167-201. — Laboratory Exercises XXXV-
XLI. — Class-room Exercises LVII-LXIll. — Problems
XXXVII-XLII.
APPENDIX A.
Manipulation — Weighing and Measuring . . • . • 3^9
Hard Glass Tubing. — Making Ignition Tubes. — Heating
Contents. xv
'AGE
Glass and Porcelain. — Evaporation. — Filtration. — Stop-
pers, Corks, Joints, and Safety Tubes. — To cut off the
Bottom of a Bottle. — To insert a Glass Tube into Rubber
Tubing. — To tit a Glass Tube to a Stopper. — To dry the
Inside of a Tube or Bottle. — To clean the Inside of a Bottle.
— To introduce a Powder into a Tube. — Pouring Liquids
and transferring Solids. — To make a Platinum Tip. —
To seal a Platinum Wire into a Glass Rod. — Collecting
Gases. — Gas Holder. — Aspirators. — Metric System. —
To counterpoise a Balance. — Weights. — Weighing. —
Measuring Liquids and Gases. — Thermometer. — Barometer.
Experiment 6.
APPENDIX B.
Gas Laws — Manipulation of Gases 354
Law of Charles. — Correction for Temperature. — Absolute
Temperature. — Law of Boyle. — Correction for Pressure. —
Aqueous Tension. — Formula for the Reduction of Gas Vol-
umes to Standard Conditions. — Manipulation of Gases.
Class-room Exercises II-IV. — Problems III-VI.
APPENDIX C.
Miscellaneous Data 370
Glossary of Chemical Terms. — Bibliography. — Apparatus
and Chei;nicals. — Reagents. — Important Elements and their
Atomic Weights. — Colors of Coatings on Charcoal. — Colors
of Borax Beads. — Colors of Residues moistened with Co-
baltous Nitrate. — Colors of Flames. — Conversion of Ther-
mometric Readings. — Specific Gravity and Melting Point
(Approximate) of Metals. — Solubility of Salts in Water. —
Weight (in Grams) of a Liter of Dry Gases at 0° C. and
760 mm. — Formula for preparing an Alcoholic Solution of
Desired Strength. — Composition of Typical Coals. — Peri-
odic Arrangement of the Elements. — Emergency Set.
Index 399
CHEMISTRY.
INTRODUCTION.
Much of the apparatus used in this book can be con-
structed in part by the student. The directions for such
work here given are exphcit, and the student is earnestly
advised not only to acquire at the outset reasonable skill
in glass working and other operations performed in the
laboratory, but also to make constant use of the hints and
directions concerning general manipulation. More detailed
directions may be found in Appendix A.
BUNSEN BURNER.
The Bunsen Burner is used as a source of heat in most
chemical laboratories. A common form of this burner is
shown in Fig. i. It is attached to the
gas cock by a piece of rubber tubing.
When the gas is turned on, the current
of gas draws air through the holes at
the bottom of the tube, and this mixture,
when lighted, burns with an almost color-
less, i.e. non-luminous, flame. It is a hot
flame and deposits no soot. The burner
should be lighted by turning on a full
current of gas and holding a lighted
match in the gas about 5 centimeters (two inches) above
I
Fig. t.— Bunsen
burner.
2 Experimental Chemistry.
the top of the burner; the height of the flame should
then be regulated to meet the special demand.
The air supply may be reduced or entirely cut off by
turning the ring at the bottom of the burner so that the
holes in the tube are partly or completely closed. As
the air supply is lessened, the flame gradually becomes
yellow, and finally is luminous, Uke an ordinary gas flame.
The luminous flame deposits soot and is not hot enough
for most laboratory work. The non-luminous flame, which
is often called ''the Bunsen flame," or simply "the flame,"
should be used in all experiments, unless the directions
otherwise state.
The parts of a Bunsen burner are shown in Fig. 6S.
The theory of the burner and structure of the flame are
considered in Chapter XIII.
A wing-top burner is shown in Fig. 2. This
movable attachment slips over the top of the
tube (see Figs. 4, 5, 7) and thereby produces a
flat flame, which is used to heat through a
longer space than the Bunsen flame permits, c.o^.
in heating glass tubing which is to be bent.
GLASS WORKING.
Glass Tubing. — Soft glass tubing is used for all pur-
poses except those in which intense heat is employed. It
melts easily, and can be bent and blown into numerous
shapes.
Cutting. — Glass tubes and rods are cut as follows :
Determine the length needed, lay the glass on the table,
and with a forward stroke of a triangular file make a short
but deep scratch where the glass is to be cut. Grasp the
tube in both hands, one on each side of the mark, and hold
Introduction.
Fig. 3. — Cutting a glass tube.
the thumbs together behind the scratch. Now push gently
with the thumbs, pull at the same time with the hands,
and the tube will break
at the desired point.
Hard glass tubing is cut
in the same manner as
soft tubing, though the
scratch must be deeper.
See Fig. 3.
The sharp ends should
be rounded by rotating the end of the glass slowly in the
flame until a yellow color is distinctly seen. See Fig. 4.
Certain precautions must
be observed in cutting
glass.
(i) If the glass does not
break with slight pressure,
do not push hard, but make
a deeper scratch and try
again.
(2) If the tube is mod-
erately large or of hard
glass, wrap it in a piece of
cloth before attempting to break it. Never try to break
large tubes by this method. Ask the teacher for directions.
Bending — General Directions. —( O Use a wing-top
burner, which produces a flat flame.
(2) Use a slightly yellow flat Bunsen flame for thin-walled
tubes, and the colorless flame for thick-walled, or hard
tubes.
(3) Never put a glass tube into the flame suddenly, nor
a hot tube on a cold surface, lest it crack.
Fig. 4. — Rounding the sharp ends of a
glass tube — " fire-polishing."
Experimental Chemistry.
Pig. 5. — Bending a tube into a right
angle — I.
(4) Heat evenly, and cool gradually. All bends when
made should be held at the top of the large yellow flame
until covered with soot. This deposit protects the outside
from the cooler air, and thus prevents sudden or unequal
cooling. This operation
is called ** annealing."
(5) If two or more
bends are made in the
same tube, they should
all lie in the same plane,
i.e. all parts of the tube
should touch a level
surface.
Operation i. To bend a
tube into a right angle.—
Hold the tube in the flame as shown in Fig. 5, and slowly rotate it
between the thumbs and forefingers in order to heat it evenly. When
it feels soft and ready to
yield, take it out of the flame
and slowly bend it into a
right angle, as shown in
Fig. 6. Compare the angle
with the corner of a square
block of wood, or with any Fig. 6. -Bending a tube into a right
other right angled object, ange —
and if not correct, repeat the operation with another tube.
Operation 2. To bend a tube into an oblique angle. — Hold the
tube in the flame, as shown
HI Fig. 7. Rotate it, and at
the same time move it slowly
to the right and left so that
from five to eight centimeters
are evenly heated. When
soft, remove it and bend
into the desired angle, as
shown in Fig. 8. It is help-
^'"'"'^S^e - 1! '"'° """ '''''^'' ^^^ t^ ^^^^^ ^^^^^ ^^Sle sketched
Fig. ^.
Introduction. 5
roughly on a piece of paper near by, so that the tube may be held over
it while being bent.
Precautions. — (i) Tubes to be bent at an oblique angle should
be heated through about twice the space required for a right angle.
A very slight bend, however, is often made by heating a short space.
v^^
Fig. 8. — Bending a tube into an oblique
angle — II.
Fig. 9. — A worthless bend.
(2) The heat should be applied continuously, and the tube never
bent at a low temperature, otherwise the curves may be flattened or
creased, or may even collapse, as shown in Fig. 9. Such bends are
brittle, and therefore worthless.
CUTTING AND BENDING GLASS.
Experiment i. — Cut into halves a glass tube about 30 centimeters
long. Bend one piece into a right angle and the other into an oblique
angle. Submit each to the teacher
for criticism. ^ — >.
Experiment 2. — Bend a glass
tube, about 35 centimeters long and
5 millimeters in external diameter,
into a double right angle, so that
each arm is about 10 centimeters
long. See Fig. 10. Preserve this
tube, as it is used in many experi-
ments.
Experiment 3. — Bend a glass tube, 40 to 45 centimeters long, into
the shape shown in Fig. 11. The numbers indicate the approximate
length of the different portions. The shorter arm is made by holding
Fig. 10. — A double right angle bend.
Experimental Chemistry.
the tube at that point (7 centimeters from the end) at
almost right angles to the flame, and then bending slowly,
regulating the pressure by
the cooling of the tube. Pre-
serve this delivery tube for
future use.
MAKING STIRRING RODS.
Experiment 4. — A piece
of glass rod about 25 centi-
meters long and 5 millime-
ters in diameter is
needed. Heat the
rod in the middle in
the ordinary — not
flat — Bunsen flame, and when soft draw it out slightly
into the shape shown in Fig. 12. Cut it .into two rods by
making a slight scratch at the desired point of the narrow
Fig. II. — A convenient
delivery tube.
5
I
I
Fig. 12. — Stirring rod partially made.
portion. It is convenient to have one rod rather blunt
for stirring hot metals and the other somewhat pointed for
general use. All rough points should be melted off in the
flame. A completed rod is shown in Fig. 13.
CLOSING TUBES.
Experiment 5. — Small tubes are closed, or sealed, by
heating one end of the tube in the Bunsen flame ; the tube
should be constantly rotated to insure even heating. Large
tubes are closed as described in Exp. 6. Close a small tube
about 10 centimeters long, and submit it to the teacher for
criticism.
Directions for working with hard glass tubing
may be found in Appendix A.
Fig. 13.—
A stirring
rod — ex-
act size.
Introduction.
MAKING IGNITION TUBES. OPTIONAL. ,
Experiment 6. — See Appendix A, § i.
CLASS-ROOM EXERCISE. I.
1. Why does thick glass crack when suddenly heated?
2. Why is glass brittle, if it has been suddenly cooled?
3. Why does a hot tube crack, if laid on a cold surface or if wet
with w^ater?
HEATING.
Effective use of the Bunsen burner flame is learned only
by experience, but a few preliminary hints may prove
serviceable. The burner should always be lighted before
any piece of apparatus is held over it, or before it is placed
under the piece of wire gauze which supports a dish (see
App. A, § 2) — sudden heat may crack the apparatus.
A small flame is more easily con-
trolled than a large one. Usually
a flame from 5 to 10 centimeters
(2 to 4 inches) high gives sufficient
heat, and is not easily blown about
by sudden drafts.
Chemical glassware is made
of uniformly thin glass, and if
heated with certain precautions
will not crack. A wet test tube
should never be put in or over the
bare flame. The temperature of
dry test tubes, or of those contain-
ing only solids, should be raised
gradually by moving them in and
out of the flame or by holding them in the flame and roll-
ing them slightly between the thumb and forefinger. If
the test tube contains a solid, special care must be taken
to distribute the heat evenly ; if the test tube contains a
Fig. 14. — Test tube and
holder.
8 Experimental Chemistry.
liquid, it should be slightly incHned so that the greatest
heat is not on the bottom, which is often thinner than
the sides ; the flame should never come above the sur-
face of the liquid, or it will heat the tube and then the
liquid touching this spot will cause the tube to crack.
In most experiments test tubes can be held between the
thumb and forefinger without discomfort If they are
too hot to handle, a test tube holder should be used.
See Fig. 14.
FILTRATION.
A finely divided solid suspended in a liquid may be
separated from the liquid by filtration or filtering. A
Fio. 16.— Folded paper Fio. 17.— Folded Fig. 18. — A fil-
— first stage. paper — second ter paper folded
Fig. 15. — Piece of stage. ready for the
filter paper. funnel.
circular piece of porous paper (Fig. 15) is folded so that
it exactly fits a funnel, and when the mixture is poured
upon this paper, the solid — the residue or precipitate — is
retained, while the liquid — the filtrate — passes through
the paper. The paper is prepared for the funnel by
folding it successively into the shapes shown in Figs. 16
and 17, and then opening as shown in Fig. 18, so that
three thicknesses are on one side and one on the other.
(See App. A, § 4.)
MISCELLANEOUS SUGGESTIONS.
Smelling and Tasting. — A safe rule to follow is never
to smell or taste unfamiliar substances, unless so directed,
Introduction. 9
and even then with the utmost caution. Never inhale a gas
vigorously, but waft it gently toward the nose. Taste by
touching a minute portion of the substance to the tip of
the tongue, and as soon as the sensation is detected, reject
the solution at once — never swallow it.
Cleanliness. — Successful laboratory work is largely meas-
ured by general cleanliness. All apparatus should be clean
before use, and should be washed as soon as it has been
used. In quantitative work it is absolutely necessary to
have balance, weights, crucibles, forceps, and every other
piece of apparatus perfectly clean.
WEIGHING AND MEASURING.
Exact relations by weight or volume are found by weigh-
ing or measuring, or by both. Such operations require a
working knowledge of the system of weights and measures
used in science and a familiarity with the chemical balance
and other instruments of precision.
The Metric System of weights and measures is used in
science. For a brief outline of the system, with tables, see
Appendix A, § i8.
The relation of the metric system to the weights and
measures in common use is shown by the following: —
Table of Equivalents.
* I meter = 39-37 inches
I liter
=
1.056 quarts (liquid)
I liter
=
0.908 quarts (dry)
I gram
=
15.432 grains
* I kilogram
=
2.2 pounds (avoir.)
* I inch
=
2.54 centimeters
I mile
=
1.6 kilometers
I cubic inch
=
16.39 cubic centimeters
I quart (liquid)
=
0.9465 liters
I pound (avoir.)
=
0.4536 kilograms
lo Experimental Chemistry.
The equivalents marked with a * should be learned.
Other equivalents can be easily deduced from the above.
(See App. A, § i8, Table of Transformation.)
A convenient relation (true only in the case of water)
to remember is i 1. = i kg. = i cu. dm. = looo cc. = looo gm.
= 2.2 lb.
The customary abbreviations of the most common de~
nominations are as follows : —
Meter, m. Liter, 1. Centigram, eg.
Decimeter, dm. Kilogram, kg. or Kg. Milligram, mg.
Centimeter, cm. Decigram, dg. Cubic centimeter, cc.
The preferable abbreviation for gram is gm. Occasion-
ally cu. cm. is used for cc. The same abbreviation may be
used for the plural and singular, e.g. i gm., 4 gm., .04 gm.
PROBLEMS. I.
1. Express: —
(«) I km. in millimeters. {c) 127.5 ^^^' i^ inches.
(J)) I cm. in decimeters.
2. Add I km., 2 m., 4 dm., 6 cm., 9 mm., and express the result in
centimeters.
3. Subtract i mm. from 4 cm., and express the answer in centi-
meters.
4. How many tubes 1.5 cm. long can be made from 249 m. of
tubing, allowing 1.5 m. for loss ?
5. Express in cubic centimeters : —
{a) I 1. {b) 27 cu. dm. (c) \'j'2\ 1. (rt?) 24.8 1.
6. How many liters in a tank 75 m. long, 27 cm. deep, and 3 dm.
wide ?
7. What is the capacity in cubic centimeters of a box 2 m. long,
8 dm. wide, and 7 cm. deep ?
8. How many grams in 1.647 kg. ?
9. How many cubic centimeters in 721 1. ?
10. Add 2 gm., 9 eg., 14 dg., and 237 mg., and express the sum in
grams.
11. How many pounds in 25 kg.?
12. How many milligrams in 5 cc. of water at 4°C.?
Introduction. II
13. The standard pressure at which a gas is measured is 760 mm.
Express the same in inches.
14. How many millimeters in 1.575 m.? In 0.5 m.? In 0.005 m.?
15. How many meters in 100.056 dm. ? In 1.5 dm. ? In 50 dm. ?
16. A box is 0.5 m. deep, 25 dm. long, and 1.2 m. wide. What is
the volume in cubic centimeters ?
17. If a cylinder holds 141. 3 gm. of water, and its height is 20 cm.,
what is its diameter ?
18. How many grams in 2000 dg.? In 1800 eg.?
19. How many centigrams in 21 gm.? In 1900 mg.? In 1565 dg. ?
20. How many cubic centimeters in 100 1.? In 0.5 1.? In 15 cu.
dm.? In 0.5 cu. m.?
21. How many grams of water in a cylinder 9 dm. high and 5 cm.
in diameter ?
22. How many millimeters, centimeters, and meters are respectively
contained in 0.437 dm. ?
23. How many grams in 1.725 kg.?
24. How many centigrams in 2.567 kg.?
25. A piece of platinum foil measuring 10.5 cm. by 1.5 cm. weighs
0.723 gm. Into how many pieces each weighing one decigram may it
be divided ?
26. A circular piece of filter paper is 10 cm. in diameter. What is
its area ?
27. A cistern is 2 m. long, 1.5 m. wide, and i m. deep. How many
liters of water will it contain ?
28. A cyHndrical gas holder is i m. in diameter and 1.5 m. high.
How many liters of gas will it hold ?
29. A wire 255 mm. long weighs 0.172 gm. What length of this
wire is necessary to make a rider weighing i eg.?
30. Sulphuric acid is 1.8 times heavier than water. How many
grams of acid will a liter flask hold ?
31 . Alcohol is 0.8 as heavy as watqr. What is the weight of 1200 cc
of alcohol ?
32. Express in meters and in millimeters : —
{a) 6 in. {b) 1.25 ft. {c) 3.6 yd.
33. How many pints in 2.5 liters ?
34. How many liters in 2 gal. ?
35. How many grams in 150 grains?
36. How many pounds in 7 kg.?
12
Experimental Chemistry,
A balance sufficiently accurate for the exact experiments
of this book is shown in Fig. 19. It is often called a
» "horn pan balance," because
the pans are made of that
material. Some of the ex-
periments give satisfactory
results, if the ordinary trip
or platform scales are used.
Fig. 19. — Horn pan balance.
Fig. 20. — Trip or platform scales.
See Fig. 20. All rough or approximate weighing should
be done on the trip scales.
Both of the balances above mentioned must be coun-
terpoised before use. Directions for counterpoising and
for weighing may be found in Appendix A, §§ 19, 21.
WEIGHING.
Experiment 7. — {a) Weigh a small object such as a bottle, stone,
or piece of lead on the trip scales. Weigh to a decigram, and express
the result in grams and a decimal fraction of a gram. Record the weight
in the proper place below.
{b) Verify the result by weighing the same object on an accurate
balance, if one is available ; if one is not available, repeat the weighing
on the same scales. Record the weight in the proper place below.
(c^ Results : —
^ ^ GrAMo.
Weight of object on trip scales
Weight of object on balance ^
Error in weighing
Introduction. 13
Form of Record of Experiments. — No one form of record
is adapted to every experiment, but all forms should include
at least (^?) a short description of the operation, {b) all
weights and measurements, and {c) answers to all questions.
The following will serve as a preliminary model : —
WEIGHING.
Experiment 7. — A small bottle was weighed on the trip scales and
then on a balance. The following results were obtained : —
Grams.
Weight of bottle on trip scales
Weight of bottle on balance
Error
Every record must, of course, include the title and
number of the experiment. It is often helpful, also, to
know the date when a particular experiment was performed.
The notes, or record, should be written, as far as possible,
during the performance of the experiment.
WEIGHING AND MEASURING.
Experiment 8. — {a) Weigh a small, dry, clean, empty bottle on
the trip scales. Record the weight at once in the note-book.
{b) Fill the bottle to the neck with water and weigh again. Be sure
that the outside of the bottle, including the bottom, is dry, before stand-
ing it on the scale pan. Record the result as shown below.
{c) Transfer all the water to an empty graduated cylinder. Pour the
water into the cylinder without spilling a drop. This can be done by
pouring the water down a glass rod (see App. A, § 12), or by a
preli4ninary movement quick enough to prevent the water from running
down the outside of the bottle. Read the exact volume of water in
the graduate. Record the result as shown below. What is the rela-
tion between the weight and volume of water?
Form of Record of Experiments. — The record of Exp. 8
should be somewhat as follows : —
14 Experimental Chemistry.
WEIGHING AND MEASURING.
Experiment 8. — A small bottle was weighed empty and then nearly
full of water, and the volume of water measured.
Results : —
Grams.
Weight of bottle and water . . . 115.2
Weight of bottle, empty .... 65.4
Weight of water 49.8
Volume of water ... 50 cc.
Weight of volume of water . . . 50.0
Error . . . 0.2
The weight and volume are nearly equal.
(date)
PROBLEMS. II.
Foi'mnlas for the co7iversion of tJiermometric readings
will be found in Appendix A, § 23.
1. Convert into Fahrenheit readings the following readings on the
centigrade scale : —
(a) 60.5 (d) -5 C^) 720 (/) 1000
(fi) 40 (e) -36.6 (//) 860 (k) o
(c) 70 (/) 643 (0 973 (O 100
2. Convert into centigrade readings the following readings on the
Fahrenheit scale : —
(a) 207
(d) 0
(g) 100
0) 2941
(d) 180
(0 -30
(//) 463
(^) -31
(0 114
(/) 674
(0 3478
(/) -76
3. What is the difference in centigrade degrees between 72° C. and
72° F.?
4. What temperature is expressed by the same number on each scale?
CHAPTER I.
PHYSICAL AND CHEMICAL CHANGES.
The matter of the universe is constantly changing.
Sometimes the change temporarily modifies the special
properties of the matter under examination, but often the
change is permanent and another substance or kind of
matter is the result. When the properties of a given por-
tion of matter are so changed that a different kind of mat-
ter is formed, then the change is called a chemical change.
If the properties are temporarily changed, then the sub-
stance has undergone a pJiysical cJiange. Most chemical
changes are accompanied by physical changes, sometimes
as causes, sometimes as effects. So closely related are
they that it is often impossible to distinguish them, even
for study.
Chemistry is largely a study of these chemical and
physical changes, and, especially, their causes, products,
and attendant phenomena.
Preliminary Definitions. — Many of the substances used
in experimental chemistry have exceedingly simple proper-
ties, so simple that it is believed that these substances can-
not be subdivided into simpler ones, and they are therefore
placed in a class by themselves, called elements. Some are
familiar and others will become familiar as the experiments
proceed. The common metals are elements. Thus copper
and zinc have properties which are so simple that chemists
regard these metals as elements. Platinum and magnesium
'5
1 6 Experimental Chemistry.
may not be familiar, but they resemble other metals in their
luster, hardness, power to conduct heat and electricity, as
well as in their chemical behavior toward other elements.
All metals do not have the same properties, but the best
known have an undeniable resemblance to the familiar
shining substances called metals.
Elements contain no other substances. Zinc is only zinc,
hence, if zinc changes into a substance having different prop-
erties, the change is due to something else than the zinc.
This fact must be remembered in interpreting the experi-
ments in this chapter. A broad knowledge of elements,
whether metals or not, is obtained only by extensive study ;
hence the ideas gained in this preliminary investigation
will be enlarged by additional experiments.
Changes in Metals when Heated. — A knowledge of
some of the conditions accompanying chemical and phys-
ical changes is necessary as a foundation for the study
of chemistry, especially of experimental chemistry. An
introductory study should be simple and limited, and the
present chapter will be confined to an examination of
the changes which metals undergo when they are heated
in the air.
GENERAL CHANGES.
Experiment 9. — Supplies: Platinum wire, copper wire, magnesium
ribbon, sheet zinc, forceps.
Examine successively a piece of platinum, copper, zinc, and magne-
sium carefully enough so that they could be detected with certainty,
if seen again under similar conditions. Then take each successively
in the forceps and hold it in the upper part of the Bunsen burner flame
long enough to produce a definite change. If the change occurs sud-
denly, remove the metal and observe the character of the change. Try
each metal until there is a definite result. Look for simple things,
such as change of color, new substances, etc. In examining the zinc
Physical and Chemical Changes. 17
hold the burner at an angle, so that the melted zinc will not drop inside.
When definite results have been obtained, answer the following : —
(i) Are all the metals changed?
(2) Are all changed permanently? Give reasons for the conclusion.
(3) Examine each product and state briefly how they differ, if they
do, from the original metal.
(4) Since nothing in the metal assisted the heat in causing the
change, what did ?
Further study is necessary to verify any theory of the
auxiliary cause of the changes in Exp. 9.
EFFECT OF HEATING A METAL WHEN UNCOVERED AND
WHEN COVERED.
Experiment 10. — Supplies: Porcelain crucible, pronged tripod or
similar support, pointed glass rod (blunt form), lead, magnesium oxide.
{a) Put a small piece of lead in a porcelain crucible and stand the
crucible on a tripod as shown in Fig. 21. If such a support is not
available, use a triangle and ring of an iron stand.
Heat the crucible for a few minutes with a low flame,
— about 5 cm. high, — then gradually increase the
heat until the bottom of the crucible is hot. Occa-
sionally scrape aside the greenish product with the
blunt glass stirring rod. The heat must not be suffi-
cient to melt this product.
Continue to heat until there is enough of the new
substance to examine. While the lead is still molten, ^^^^^^'""^^'''^'^^^
firmly grasp the crucible near the edge with the for- j^^po^!^^ ^ ^
ceps, and pour out the contents of the crucible upon
an iron pan, or a piece of asbestos board, or a block of wood. Com-
pare the product with the unchanged lead, stating briefly the differences.
How does the change in the lead resemble the changes in Exp. 9 ? Is
it a similar chemical change? (If the crucible is not clean, heat it
where the lead sticks, and scrape off the softened lead with a file.)
{F) In the same crucible put another piece of lead and cover it com-
pletely with magnesium oxide, pressing down the white, infusible sub-
stance with the finger. Heat the crucible as in {a), but do not disturb
the contents. The lead should be completely covered during the whole
1 8 Experimental Chemistry.
operation. Heat until it is certain that the lead has been melted several
minutes — long enough at least to produce the greenish powder, if any
can form. Pour out the contents as before. Has any greenish powder
formed? Why.-' What, then, assists heat in the chemical changes in
Exp. 9 and Exp. lo (a) ? Does this conclusion verify your theory
drawn from Exp. 9? If not, modify your notes accordingly.
The result of Exp. 10 gives a clew to the nature of the
permanent changes which metals undergo when heated in
the air, but it does not prove conclusively whether by the
change the metals undergo a loss or a gain. This question
is considered in the next experiment.
RESULT OF HEATING A KNOWN WEIGHT OF A METAL IN
THE AIR.
Experiment 11. — Supplies: Powdered iron, zinc dust, crucible and
support, glass rod used in Exp. 10, crucible block.
{a) Clean and dry a porcelain crucible and weigh it to a decigram on
the trip scales. Slide powdered iron into the crucible from a narrow
strip of smooth paper creased in the middle, until about 3 gm. have
been added. (See App. A, § 12.) Do not spill the iron on the scale
pans. The iron should not coat the inside of the crucible, but be
in a heap on the bottom. Wipe off any iron which is on the outside
oi" the crucible or on the inside above the heap, before weighing the
crucible and contents. The iron need not weigh exactly 3 gm., but
the exact weight, whatever it is, must be known. Record the weights
n the note-book, as soon as made, thus: —
Grams.
Weight of crucible 12.2
Weight of iron and crucible . . c 15.4
Weight of iron 3.2
In carrying a crucible to and from the scales, it should not be carried
in the hand nor by the forceps, but placed in the crucible block which
is shown in Fig. 22. Support the crucible as before and heat it for
about ten minutes. Heat with a low flame at first and gradually in-
crease the heat. Finally stand the burner under the crucible, taking
Physical and Chemical Changes. 19
care, however, that the flame is below the top of the crucible. Touch
the metal occasionally with the pointed glass rod
and if it is hard, pierce the crust cautiously in
several places to expose a fresh surface. If iron
clings to the rod, hold the pointed end over the
crucible and tap the rod gently against the crucible
so that the particles will fall back. Avoid losing
anything from the crucible. After the crucible is
cool enough to bear the hand without discomfort, Fig. 22. — Crucible
weigh again. If the weight has changed, record block.
thus : —
Weight of crucible and contents before heating
Weight of crucible and contents after heating
Gain or loss in weight
Grams.
15.4
If there is no change in weight, heat again, and then weigh ; if still
no change, consult the teacher. What is proved by the change in
weight?
{b) The general result obtained in {a) may be verified by heating
zinc dust. Weigh the crucible and weigh in it about 3 gm. of zinc
dust, observing the same precautions as in {a) . Record thus : —
Grams.
Weight of crucible 13-0
Weight of zinc and crucible 15 -9
Weight of zinc 2.9
The zinc must be heated with the utmost care to avoid loss. Use a
low flame, not more than 5 cm. high, during the entire heating. Do
not stand the burner under the crucible, and if the zinc glows or gives
off" a white smoke, remove the burner at once. This smoke weighs
something, and its loss means a loss in weight which cannot be esti-
mated. If a crust forms on the zinc, break it as in (a). Avoid losing
any zinc or any of the white product. Heat the crucible for about 10
minutes, then cool and weigh. If there is a change in weight, record
thus : —
Weight of crucible and contents before heating
Weight of crucible and contents after heating
Gain or loss in weight
Grams.
15.9
20 Experimental Chemistry.
If there is no change, heat again and then weigh ; if still no change,
consult the teacher. What is proved by the change in weight ?
If magnesium, lead, copper, or mercury, in fact, almost
any metal, is heated as in Exps. lO and ii, similar results
are produced. The nature of the interacting substance,
as far as the above experiments show, is unknown, but
whatever it is, it has combined with them in such a way
as to change their properties. It forms chemical com-
pounds of which the metal is one part and this " unknown
substance" is another part. If the process could be re-
versed by decomposing one of these compounds and get-
ting this ''unknown substance," it could be studied, and
its nature would throw some light on the real cause of the
changes which metals undergo when heated in the air.
None of the compounds — the products of Exps. lo and
II — is suited to such an experiment. Fortunately we
have available the very compound used by the chemist
who discovered this ''unknown substance." It is a red
powder and is formed by heating the metal mercury in the
air, just as the white powder was formed from zinc, the
black film from copper, and the greenish powder from
lead. When heated, this red powder decomposes, into a
well-known metal and a gas which is the "unknown
substance."
PRODUCTS OBTAINED BY HEATING THE RED POWDER.
Experiment 12. — Supplies: Apparatus shown in Fig. 23, pneu-
matic trough, small bottle, iron stand and clamp, the red powder, splin-
ter of soft wood, ignition tube.
The delivery tube for this experiment was made in Exp. 3. If the
ignition tube is not available, it may be made according to the direc-
tions in Appendix A, § i. If the ignition tube is not large, it may be
attached to the delivery tube as shown in Fig. 23 ; if it is rather large,
Physical and Chemical Changes.
21
then it must be attached to the delivery tube by putting the rubber con-
nector inside the ignition tube, the upper end of the delivery tube itsell
4=, r-^
Fig. 23.
Apparatus for heating the red powder and collecting the gaseous
product.
lemaining, as before, within the connector
i
fU
v^
A cross section of such a
connection is shown in Fig. 24.
Shp a Httle of the red powder from a paper into the
ignition tube and connect with the deHvery tube. (See
App. A, § II.) Clamp the tube near the top so that the
apparatus shall have about the position shown in Fig. 23.
Tap the tube gently to spread out the powder in the tube.
The outer end of the delivery tube reaches just below the
shelf of a pneumatic trough arranged to collect the gas
over water. Fill the bottle full of water, invert it, and
stand it on the shelf near the hole. (See App. A, § 15.)
Heat the whole ignition tube with a low flame at first,
gradually increasing the heat where the red powder is
located. As the heat increases, bubbles pass up through
the water. Reject the first few, which are air, then slip
the bottle over the hole in the shelf. If the powder
" crawls ''"' up the ignition tube, tap the tube gently until
the contents slips back. If the evolution of gas ceases,
heat more strongly. Do not wave the burner back and
forth, but apply the heat steadily by a low flame which
envelopes that part of the tube containing the powder.
Fig. 24.—
Section of
an inside
connector.
22 Experimental Chemistry.
Avoid heating the end of the tube, as it is often so thin that jt
melts.
Collect the bottle at least half full of gas and remove it thus : Slip
the bottle from the shelf, still keeping the mouth under water, and
cover the mouth with a piece of moistened filter paper. Keep hold of
the paper with one hand and with the other invert the bottle and stand
it, still covered, on the desk. Remove the delivery tube at once from
the trough to prevent the water from being drawn up into the hot
apparatus.
The gas from the red powder is the " unknown substance " which
was taken from the air by a certain metal when heated in the air. Ex-
amine the deposit on the inside of the ignition tube. If its nature is
doubtful, pour it out on a block of wood. What is it? Test the gas
thus : Light a splinter of wood, let it burn a few seconds, then blow it out
and plunge the glowing end into the gas in the bottle. What happens ?
Remove the splinter, if there is any decided change, and once more
plunge it, while glowing, into the gas. Answer the following : —
(i) Is the gas air? Why not?
(2) Where must the gas have come from originally?
A substance is usually identified by comparing its prop-
erties with those of some known substance with which it is
suspected to be identical.
IDENTIFICATION OF THE GAS OBTAINED FROM THE RED
POWDER.
Experiment 13. — Obtain a bottle of oxygen from a gas holder
and plunge a glowing splinter into it. Has oxygen the same character-
istic property as the " gas obtained from the red powder " ? What one
conclusion can be drawn from this experiment?
Definitions. — An Oxide is a compound of oxygen and
one other element, such as iron, zinc, or lead. Chemical
action is a term applied to all classes of chemical changes,
such as the addition of oxygen to iron, or the decomposi-
tion of the red powder into oxygen and mercury. The
Physical and Chemical Changes. 23
cause, accompaniments, ' and products of chemical action
will be constantly studied.
LABORATORY EXERCISE. I.
1. Summarize the results of Exps. 9 to 13 into a concise argument.
2. Name all the oxides studied or used in this chapter.
3. What is the chemical name of the red powder?
CLASS-ROOM EXERCISE. V.
1. Discuss the states and properties of matter.
2. Physical and chemical changes.
(a) Give three illustrations of both occurring in every-day life.
(d) What is supposed to cause many chemical changes? Is
change constant or intermittent?
(c) What aids and what retards chemical change ?
3. Review the Metric System.
4. Manipulation.
(a) Precautions in heating hard glass and porcelain.
(d) Most effective use of Bunsen flame.
(c) Give the successive steps in collecting gases.
5. Give evidences of chemical action observed in Exps. 9 to 13.
CHAPTER II.
OXYGEN
The experiments in Chapter I. teach that one important
constituent of the air is oxygen. Subsequently, air will be
examined with special reference to its composition. We
shall now study oxygen, which is the most important of
the few simple substances to which the matter of the uni-
verse may be reduced.
PREPARATION AND PROPERTIES OF OXYGEN.
Experiment 14. — Supplies: Retort, ring and stand, pneumatic
trough, six bottles (about 250 cc. capacity), filter paper, splinter of soft
wood, deflagrating spoon, roll sulphur, piece of charcoal fastened to a
wire, short piece of wire picture cord unwound at one end, magne-
sium ribbon, 15 gm. of crystallized potassium chlorate, and 15 gm.
of coarsely powdered manganese dioxide.
Fig. 25. — Apparatus for the preparation of oxygen.
Procure all the supplies before attempting to prepare the oxygen.
Arrange the apparatus as shown in Fig. 25, and collect the gas ovei
24
Oxygen.
25
water. (See App. A, § 15.) The potassium chlorate and manganese
dioxide should be free from pieces of cork, paper, straw, or wood ;
the retort must be dry and clean and provided with
a tightly fitting glass stopper.
Mix the potassium chlorate and manganese dioxide
on a piece of smooth paper, and pour the mixture into
the retort as shown in Fig. loi. Invert the bottles full
of water in the trough, and have the filter paper where
it can be conveniently reached. Adjust the apparatus,
and heat the bare retort carefully with a low flame.
Direct the heat upon the part of the retort containing
the mixture. Do not collect the gas before it bubbles
freely through the water. If the gas is evolved too
rapidly, lessen the heat; if not at all, then examine
the stopper for a leak. Collect six bottles of gas,
covering each when full with a piece of wet filter
paper, as soon as taken from the trough.
When the gas has been collected, remove the neck
of the retort immediately from the water, lest the cold
water be drawn up into the hot bulb of the retort, as
the gas contracts.
Proceed at once to study the properties of oxygen
by the following experiments : —
{a) Thrust a glowing splinter of wood into one
bottle. Remove the splinter, if there is any change,
and repeat as many times as possible. Describe the
result. What property of oxygen does this experiment
show? What is the essential difference observed be-
tween burning in air and in oxygen?
{b) Put a small piece of sulphur in the deflagrating
spoon (Figs. 26 and 27), hold the spoon in the flame
until the faint blue flame of the bufning sulphur can
be seen, then lower the spoon into a bottle of oxygen.
Describe and explain any change in the flame. Smell
the vapor cautiously by wafting a little toward the
nose, and state of what the odor reminds you. To
what class of compounds, in all probability, does this vapor belong?
(If the odor of the burning sulphur is disagreeable, plunge the spoon
into water, or remove it to the hood.)
Fig. 26. — Defla-
grating spoons.
i6
Experimental Chemistry.
{c) Hold the charcoal in the flame long enough to produce a glow,
then lower it into a bottle of oxygen. Describe and explain any change
in the appearance of the charcoal. How does the
chemical change resemble that in {b) ?
{d) Dip the unwound end of the picture cord into
the deflagrating spoon containing the melted sulphur
used above. If the sulphur on the end is not burning,
light it, and quickly lower it into a bottle of oxygen.
The picture cord should take fire and burn brilliantly.
A picture cord is chemically the same as iron. Can the
change be explained similarly to that in {b) and (^) ?
Is there any visible, material proof of change besides
the vapor from the burned sulphur ? Does the product
of the chemical change belong to the same class as the
products in (/;) and {c) ?
{e) Hold one end of
a piece of straightened mag-
FlG. 27.— Com-
bustion cup
(or deflagrat-
ing spoon)
made from a
piece of cop-
per wire and
a piece of
crayon hav-
ing a cavity
at the upper
end.
nesium ribbon about 10 cm. long in the forceps, light
the other end, and instantly lower it into a bottle of
oxygen. In what general ways does the result resemble
those obtained in the other experiments with oxygen?
What conclusive evidence is there that the magnesium
was chemically changed ? How does the product resem-
ble that formed by burning magnesium in the air? Are
the products identical ?
Write a short account of this experiment in your
notebook, including a sketch of the apparatus used for
the preparation of oxygen. Arrange the notes some-
what in the order of the above directions, using the same
letters, and answering all questions.
LABORATORY EXERCISE. II.
1. Summarize the properties of oxygen.
2. What is its most characteristic property?
3. What oxides were formed in Exp. 14?
4. From what class of substances might oxygen be
prepared ?
5. From what substances have you previously prepared it?
6. What chemical change probably takes place when a substance
burns ?
Oxygen. 27
7. If air contains something besides oxygen, what must be the
general properties of this other constituent?
It was shown in Exp. 1 1 that when a metal is heated
in the air it gains in weight, and in Exps. 12 and 13 that
the gain is actually due to the formation of a new substance,
composed of the metal and oxygen. The experiments with
oxygen, also, show that this gas unites easily and vigorously
with metals, like iron and magnesium, and with other sub-
stances, like sulphur and charcoal (chemically known as
carbon), forming oxides. These oxides belong to a large
class of substances called Chemical Compounds. The experi-
ments thus far performed show that chemical compounds
are produced under special conditions, and that they possess
specific properties. Additional facts may be learned about
chemical compounds by studying the mathematical relation
of oxygen to two of its compounds.
THE COMBINATION OF OXYGEN WITH MAGNESIUM.
Experiment 15. — Method: Heat a weighed quantity of magnesium
in a covered crucible at a low temperature, weigh the product, and the
gain will be the weight of oxygen which has combined with the known
weight of magnesium.
Supplies : Porcelain crucible and cover, powdered magnesium, for-
ceps, pronged tripod or similar support, crucible block.
Process: Clean and dry the crucible and cover, and weigh both
together. Record the weight as shown below. Weigh in the crucible
from 0.4 to 0.5 gm. of magnesium, taking care to weigh to a centigram
the exact amount. Record the weights thus : —
Grams.
Weight of crucible, cover, and magnesium . . 13.32 ^
Weight of crucible and cover 12.81
Weight of magnesium 0.51
1 These and similar results which will be given are solely for illustration.
28 Experimental Chemistry.
Stand the crucible on the tripod, as shown in Fig. 28, and heat for
five minutes with a flame which just touches the bottom of the crucible.
Grasp the cover firmly by the ring with the clean forceps, cautiously
lift it, and if the magnesium glows, cover the crucible instantly. Repeat
this operation at frequent intervals, until the glow
is dull red and ceases to spread through the mass ;
then adjust the cover so that a small opening is
left between the cover and the crucible, and heat
strongly for ten or fifteen minutes. If the contents
has ceased to glow, heat the crucible, uncovered,
for five minutes. Take care not to upset the cover
by accident or insecure handling with the forceps.
Fig, 28. — Covered At no time should the flame touch the cover of the
crucible supported crucible ; roughly speaking, the flame should reach
by a tripod. , , ' ^ , , . 1 • • 1
as high outside as tlie magnesium does inside.
The operation is complete when the product is gray or white. Cool
the crucible gradually. When cool enough to handle, it is cool enough
to weigh. Weigh, and record thus : —
Grams.
Weight of crucible, cover, and contents, after heating . . 13.66
Weight of crucible, cover, and contents, before heating . . 13.32
Weight of oxygen which has combined with the magnesium 0.34
Heat the uncovered crucible again strongly for five minutes, cool, and
weigh again as before. If the weight is not the same, continue until
the last two weights are approximately the same.
Calculation: Since 0.34 gm. of oxygen combined with 0.51 gm. of
magnesium, the ratio in which they combined may be expressed by the
proportion —
0.34 :o.5i .w.x
.'. x= 1.5
Class average =
Discussion of Experiment 15. — This result means that
for every gram of oxygen one and one-half grams (approxi-
mately) of magnesium are necessary to form a chemical
compound. The truth of this statement is more vivid
when the class average is known ; for if the work has
Oxygen. 29
been done accurately, the class average will agree closely
with the theoretical result. More important conclusions
will be drawn later from this experiment, but it is sufficient
at present to grasp the idea of definite relations by iveigJit
which exist between the components of a chemical com-
pound.
THE RELATION OF OXYGEN TO POTASSIUM CHLORATE.
Experiment 16. — Method: Heat a weighed quantity of potassium
chlorate in a covered crucible until all the oxygen is evolved, and the
loss will be the oxygen which was in chemical combination in the
potassium chlorate.
Supplies : Crucible and cover, forceps, support (as in Exp. 15), dry,
powdered potassium chlorate, crucible block.
Process: Weigh the crucible and cover as in Exp. 15, and weigh
in the crucible about 1.5 gm. of dry, powdered potassium chlorate.
Weigh to a centigram the exact amount taken, and record the weights
thus : —
Grams.
Weight of crucible, cover, and potassium chlorate . . 15-34
Weight of crucible and cover 14.00
Weight of potassium chlorate 1.34
Arrange the apparatus as shown in Fig. 28, and proceed with the
heating somewhat as in Exp. 15. The potassium chlorate first melts,
and as the oxygen, which is liberated, bubbles through the molten mass,
it may spatter a little potassium chlorate up on the inside of the cover.
Regulate the heat, if possible, to avoid this difficulty. Remove the
cover with the clean forceps occasionally, and if it is coated with potas-
sium chlorate, lay it (ring side down) on a block of wood, deftly loosen
the thin layer with a pin or sharp rod, then hold the cover over the
crucible and gently scrape the pieces off into the crucible with a sharp
rod. If any fall upon the desk and are touched with the moistened
finger, they will cling to the finger and may be returned thereby to the
crucible. Regulate the flame so that the bubbling will be slow. As
the temperature increases, the mass solidifies, and the action apparently
stops. Increased heat, however, drives off more oxygen. As soon as
the second evolution of oxygen is passed, remove the cover, lay it (ring
30 Experimental Chemistry.
side down) on a block of wood where it will be safe, and then heat
the crucible strongly for five or ten minutes. The decomposition is
complete when the product is perfectly white and non-crystalline, and
not changed in appearance by a flame nearly as high as the top of the
crucible. Cool gradually, and when cool, weigh the crucible, cover, and
contents, as before. Record the weights thus : —
Grams.
Weight of crucible, cover, and contents before heating . 15.34
Weight of crucible, cover, and contents after heating . 14.82
Weight of oxygen lost 0.52
Heat again strongly for ten minutes, then cool and weigh as before.
If the weight is the same as after the first heating, proceed with the
calculation ; if not the same, heat and weigh until two successive weights
are the same, i.e. heat to constant weight.
Calcidatwn : From 1.34 gm. of potassium chlorate 0.52 gm. of
oxygen was obtained, or —
1.34 : 0.52 : : 100 : x
X = 39.3 per cent.
Class average = per cent.
Discussion of Experiment 16. — This result means that a
definite proportion of oxygen by weight can always be ob-
tained from potassium chlorate. Individual experiments
may produce slightly varying results, but an average of
several results will be substantially identical with the theory.
As in Exp. 15, the full significance of the result is not yet
apparent, and later the experiment will be discussed in its
widest application. This experiment is intended to teach,
at this stage of the experimental work, that definite zveigJit
relations are essential to chemical compounds.
The Weight of a Liter of Oxygen will be serviceable, later,
in explaining fundamental facts. Before it can be deter-
mined, however, the contents of Appendix B should be
mastered.
Oxygen.
31
PROBLEMS. III., IV., v., VI.
See Appendix B, §§ 2, 5, 7.
CLASS-ROOM EXERCISE. II., III., IV.
See Appendix B, §§ 5, 6, II.
WEIGHT OF A LITER OF OXYGEN.
Experiment 17. — Method: Weigh the volume of water displaced
by generating a known weight of the gas.
Apparatus : The apparatus is shown in Fig. 29. A is a test tube
containing a mixture of potassium chlorate and manganese dioxide ; it
is attached to the bent tube F
by a small rubber stopper. B
is a bottle to be filled with
water; it is provided with a
two-hole rubber stopper,
through which pass F and C,
the latter being the glass por-
tion of the whole delivery tube,
CC, while C is rubber. A
Hofmann screw is attached at
the point E. A large bottle,
D^ serves to catch the w^ater
forced over from B through
CC by the oxygen generated
in ^. The bent tube /^ retains
any moisture driven over from
A^ and is weighed with A. All joints must be air tight.
Process: (i) Fill A two-thirds full of a mixture of manganese
dioxide (2 parts) and potassium chlorate (3 parts). Each substance
must be powdered and free from organic matter {e.g. paper, cork, straw).
The mixture should be dried by heating it in an oven to about 110° C,
on a radiator or on some convenient heated object. Push a little
glass wool, or shredded asbestos (previously ignited to a red heat), into
the tube to hold the contents in place. The tube must be free, inside
and out, from loose particles of manganese dioxide, or other matter
which might be rubbed off after the weighing — clean, in other words.
Weigh AF to a centigram. Weigh the empty, dry, clean bottle, D,
to a decigram on the trip scales.
Fig. 29. — Apparatus for determining the
weight of a liter of oxygen.
32 Experimental Chemistry.
(2) Fill B with water nearly to the neck. Fill CC with water and
tighten the Hofmann screw to prevent the water from running out.
Insert AJ*^ into the stopper of B. Push the stopper into the bottle,
slowly at first, then hard ; if water rises in F^ loosen the screw at E
slightly, remove A, and blow gently into F io force the water back into
B. When properly adjusted, the water should be in B and CC, but not
in B\ Replace A, taking care not to crush the thin glass by pushing it
too hard upon its stopper ; open the screw at E. If the apparatus is
tight, the water will not flow out. It should be adjusted until air tight.
Leave the screw open.
(3) Heat A gently with a low flame, beginning at the closed end
and keeping the flame back of any moisture which might be driven out of
the mixture. The liberated oxygen will force the water from B into D.
Heat A just hot enough to cause a gentle flow of water into D. When
D is about three-fourths full, decrease the heat gradually. If moisture
appears near the top of the mixture, heat behind it, and either drive it
over into F or cool A so gradually that the moisture will not shatter the
tube. While A is cooling sufficiently to weigh, stand a thermometer
in D ; also read the barometer. When A is cold, raise B until the
water is at the same level in B and D, pinch C tight and remove it
from D. Read and remove the thermometer. Dry D on the outside,
if necessary, and then weigh it, using the same large weights as before ;
the gain in weight (in grams) of D gives the volume (since i gm. of
water =1 cc.) of oxygen liberated. Weigh AF\ its loss in weight
is the weight of the oxygen which passed into B.
Calculation : The observed volume must be reduced to the volume
it would occupy, if it were at 0° C, 760 mm., and in the dry state, i.e.
reduced to standard conditions. (See App. B, § 7.) This is done by
the formula —
^,^ V-(P--^)
760 (i + .00366/)
Substitute the proper values in this formula, and solve for F— the cor-
rected volume of oxygen liberated.
F= cc.
Since i 1. contains 1000 cc, then
= 1. of oxygen liberated.
1000 ^°
Oxygen. 3;^
The weight of oxygen liberated is found thus : —
Grams.
Weight of A/"' before liberation of oxygen
Weight of A/^ after liberation of oxygen
Weight of oxygen liberated . .
Therefore the weight of i 1. of oxygen equals
Wt. of oxygen liberated _
Vol. of oxygen liberated
gm.
Substitute in this formula the proper values. Tabulate the total result
of the experiment thus : —
WEIGHT OF 1 L. OF OXYGEN.
Definitions. — Combination may mean a process or its
product. Thus the process employed in Exp. 1 5 is called
combination, and the product, magnesium oxide, is also
called a combination of magnesium and oxygen. The
context, however, usually eliminates any ambiguity. A
synonym of the process, combination, is Synthesis, which
literally means " a putting together." Decomposition means
the separation of a chemical compound into its parts.
These parts are called decomposition products, or, better,
Components. And as decomposition may be partial or
complete, these components may be elements or com-
pounds. A synonym of decomposition is Analysis, which
literally means "a breaking up." Thus in Exp. 16, the
34 Experimental Chemistry.
potassium chlorate was actually broken up into the ele-
ment, oxygen, and the compound, potassium chloride. A
Determination is an accurate chemical examination. It
may be synthetical, analytical, gravimetric — involving
weights, or volumetric — involving volume. Experiment i6
was a gravimetric determination of the per cent of oxygen
in potassium chlorate. A Residue is usually the product
remaining after something has been removed from a com-
pound or a mixture. Thus in Exp. i6 the residue is
the potassium chloride remaining in the crucible. A Chlo-
ride belongs to a class of compounds, each member of
which consists of the element chlorine and one other ele-
ment. Chlorides bear the same relation to chlorine as
oxides bear to oxygen.
Oxidation. — Oxides were mentioned in Chapter I.
The process of forming an oxide, or in general of adding
oxygen to an element or compound, is called Oxidation.
Thus Exp. 15 was an oxidation of magnesium. The
metals heated in the experiments performed in Chap-
ter I. were also oxidized. Compounds which yield oxygen
easily are called oxidizing agents ; thus nitric acid, a
compound to be studied later, decomposes easily, and
one of the decomposition products is oxygen, which
often immediately combines with some other element or
some compound. Oxidation is an important chemical
operation.
Law of Definite Proportions by Weight. — Experiment 1 5
showed that when magnesium is heated so that it has
access to oxygen the two elements combine in the ratio of
I to 1.5. This also means that the product, magnesium
oxide, always contains 1.5 times as much magnesium as
oxygen by weight. Experiment 16 shows that potassiucn
Oxygen. 35
chlorate always yields about 39 per cent of oxygen. A vast
number of experiments has verified these, and similar
results, in the case of other compounds. The fact of
definite proportions by weight revealed by these two
experiments (15 and 16), and extended to cover many
others, is the result of a law known as the " Law of
Definite Proportions by Weight." It was estabHshed as
the outcome of a controversy between two chemists, Proust
and Berthollet, which lasted from 1799 to 1806, though
isolated suggestions of the law had previously appeared.
The law may be stated thus : —
A chemical compound ahvays contains the same elements
in the same proportions by weight.
Or,
A chemical compoimd has a definite composition by weight.
This law is one of the foundation stones of chemistry.
Compounds and Mixtures. — Proust, in his discussion
with Berthollet, showed conclusively the distinction be-
tween mixtures and compounds. The "gradual changes "
found in the substances analyzed by Berthollet were shown
by Proust to be due to the fact that Berthollet analyzed
mixtures — masses contai7iing ingredients in varying pro-
portions. Mixtures are the opposite of compounds ;
they may have any composition ; their ingredients are
not held together by that fornr of energy called chemical
affinity, or attraction. A mixture of sulphur and iron
can be easily separated by mechanical means, but in a
compound of iron and sulphur the components are held
together by a force which does not yield to the influence
of a solvent or a magnet, nor to any other purely physi-
cal agent.
^6 Experimental Chemistry,
CLASS-ROOM EXERCISE. VI.
1. Additional study of oxygen.
(a) Occurrence and distribution.
{d) Chemical relation to plants and animals.
(c) Liquid oxygen.
(^) Uses.
(e) Preparation on a large scale.
2. Historical.
(a) Discovery of oxygen.
(d) Essential facts of Priestley's life.
{c) Scheele's and Lavoisier's actual contribution to the dis
CO very.
(d) Names of oxygen, by whom so called, and why.
3. Combustion.
(a) Theories of Becher and Stahl.
(d) Lavoisier's work.
(c) Present views.
4 Manipulation.
(a) Rules for weighing.
(d) Heating and cooling porcelain.
(c) General precautions in accurate experiments.
(<^) Sources of error in Exps. 15 and 16, how detected,
avoided, and remedied.
5. Elements and compounds.
(a) Definition of each.
(d) Enumerate all the elements and compounds thus far used
or studied.
(c) What compounds have been made in the previous experi-
ments? What compounds have been decomposed?
(d) Define and illustrate : a mixture, combination, decomposi-
tion, determination, residue, oxide, oxidation, chloride,
component.
6. Law of Definite Proportions by Weight.
(a) Historical development.
(d) Statement in several different forms.
(c) Illustrations.
Oxygen. 37
PROBLEMS. VII.
1. How many grams of magnesium will exactly combine with 16
gm. of oxygen? With 32 gm. of oxygen? With 79 gm. of oxygen?
2. How many grams of oxygen will exactly combine with 24 gm.
of magnesium? With 48 gm. of magnesium? With 90 gm. of
magnesium ?
3. What weight of oxygen gas could be obtained by decomposing
200 gm. of potassium chlorate? What weight of oxygen, if the potas-
sium chlorate had contained 10 per cent of an impurity?
4. What weight of oxygen gas could be obtained from 100 gm. of
pure potassium chlorate? If the chlorate used contained 12 per cent
of its weight of potassium chloride, what would be the difference in the
amount of oxygen obtained ?
5. If a liter of oxygen weighs 1.43 gm., what will be the weight of
oxygen in a room 25 m. long, 17 m. wide, and 15 m. high?
6. How many liters of oxygen can be obtained from 122.5 S^- ^^
potassium chlorate? From 255 gm. of potassium chlorate, allowing 2
per cent for impurities in the potassium chlorate ?
7. A quantity of potassium chlorate weighing 721 gm. is decom-
posed, the oxygen collected, and magnesium is burned in the oxygen
until no gas remains.
(a) What weight of oxygen is produced ?
(d) What volume of oxygen is produced?
{c) What weight of magnesium is required?
(d) What weight of magnesium oxide is produced?
CHAPTER III.
HYDROGEN.
The element hydrogen is a gas under ordinary condi-
tions. It was recognized as a ''peculiar and distinct sub-
stance " in 1766 by the English chemist Cavendish. It is
a component of many compounds, and is the standard of
several scientific measurements.
Hydrogen was first prepared by the interaction of an
acid and a metal.
PREPARATION OF HYDROGEN BY METALS AND AN ACID.
Experiment 18. — Supplies: Test tubes, granulated zinc, iron
filings, dilute sulphuric acid, dilute hydrochloric acid.
{a) Put a few pieces of granulated zinc in a test tube and cover
them with dilute sulphuric acid. A gas will bubble through the liquid.
After a minute or two test the gas by holding a lighted match at the
mouth of the test tube. Describe the result.
{b^ Repeat («), using dilute hydrochloric acid. Is the final result
the same ? If not, repeat both {a) and (Jb).
(^) Put a little iron filings in a test tube and cover them with
dilute sulphuric acid. Test the gas evolved, as in {a). Is it the same
gas?
(^d) Repeat (^), using dilute hydrochloric acid, and test the gas, as
before. Is it the same gas ?
What is the source of the hydrogen? What are its most character-
istic properties ?
The slight explosion heard in some, perhaps all, of the above experi-
ments is due to impure hydrogen. What, in all probability, is the
impurity? Does the observation suggest any danger in working with
hydrogen? What is the simplest way to avoid the danger?
38
Hydrogen.
39
The specific nature of hydrogen may be learned from its
preparation on a larger scale than in Exp. i8 and a study
of its properties.
A is a. 500 cc flask
a safety tube. C is
Q
B
PROPERTIES OF HYDROGEN.
Experiment 19. — Supplies: Hydrogen generator, 25 gm. of granu-
lated zinc, dilute sulphuric acid, pneumatic trough, five bottles (each
250 cc), taper.
Apparatus: The generator is shown in Fig. 30
provided with a two-hole rubber stopper. B is
the double right an-
gle bend which was
made in Exp. 2. D
is a rubber connector
about 4 cm. long. E
is a delivery tube of
such length that its
end is level with the
bottom of the flask.
The flask stands on
the desk, and E dips
into the pneumatic
trough.
Precaidions : All
joints must be tight.
The stopper must fit
perfectly, and both B
and C must turn with
difficulty after the
stopper has been
pushed well into the
flask. Acid must be
left in the bend of
the safety tube to
about the point indicated by the dotted line.
II
Fig. 30. — Hydrogen generator.
(See App. A, §5.) All
flames, large and stnall, should be extinguished in the vicinity during
this experiment, because mixtures of air and hydrogen, when ignited,
explode violcjitly. Save the contents of the generator for Exp. 20.
40 Experimental Chemistry.
Process: Incline the flask and slide the zinc into it from a paper.
Do not drop it in — the bottom is thin and may crack. Insert the
stopper with its tubes by holding the flask by the neck in one hand
and with the other working the stopper slowly" into the flask. Adjust
the flask and delivery tube so that the end of E is on the bottom of the
pneumatic trough and under the hole in the shelf. Have the bottles
filled with water and inverted in the trough, and filter paper ready to
cover the bottles of gas.
Pour enough dilute sulphuric acid through the safety tube into the
flask to cover the zinc, taking care to leave some acid in the lower bend
of the safety tube. This latter precaution prevents the gas from flowing
out the back of the apparatus ; if at any time the gas should flow back-
ward, pour a little more acid into the bend. If the acid will not run down
the safety tube, loosen the stopper for an instant. Remember, how-
ever, that this operation may admit air, and the gas evolved for the next
few minutes will be impure. As soon as the acid and zinc begin to
interact, hydrogen will be rapidly evolved. Allow the action to proceed
several minutes to free the apparatus from air, then collect five bottles
of gas over water, as in the case of oxygen. As fast as the bottles are
full, remove and cover them with wet filter paper. When the gas has
been collected, remove the generator to the hood. Proceed at once to
study the properties of hydrogen as follows : —
{a) Uncover a bottle for an instant and then drop a lighted match
into the bottle. Describe and explain the result.
{b) Remove the paper from a bottle of hydrogen and allow it to
remain uncovered for three minutes — by the clock. Then show the
presence' or absence of hydrogen by dropping a lighted match into the
bottle. Explain the result.
{c) Verify the result in (J?)
thus : Hold a bottle of air
over a bottle of hydrogen,
remove the paper from the
hydrogen bottle, then bring
the bottles into the relative
positions shown ia Fig. 31.
Hold them there for a minute
or two, then stand the lower
Fig. 31. — Pouring hydrogen.
bottle on the desk, cover the upper one with wet filter paper and stand
it on the desk. Drop a lighted match into each bottle. What has
Hydrogen. 41
become of the hydrogen? What property of hydrogen is shown by
{b) and (0 ?
{d) Read the directions carefully and watch all parts of the appa-
ratus for definite residts. I nvert a bottle of hydrogen, remove the paper,
and thrust a hghted taper up into it. Withdraw the taper slowly and
then insert it again. Does the hydrogen burn ? If so, where ? Does
the taper burn when in the bottle? When out of the bottle? Feel of
the neck of the bottle, and describe and explain. Repeat with the
remaining bottle of hydrogen. What three properties of hydrogen are
shown in {d) ?
LABORATORY EXERCISE. III.
1. Sketch in your note-book the generator used to prepare hydrogen
in Exp. 19.
2. Summarize the properties of hydrogen.
3. What is its most characteristic property?
4. Why is there danger of an explosion in generating hydrogen?
How may the danger be averted?
5. From what class of substances may hydrogen be prepared?
Chemical Reaction. — The mutual chemical action which
results when two or more elements or compounds are added
to each other, is called a Reaction. Thus when zinc and
sulphuric acid come together there is chemical action, ap-
parent so far only by the formation of hydrogen. Inas-
much as every experiment involves one or more reactions,
an examination of the reaction which takes place in the
preparation of hydrogen will serve as a preliminary study
of reactions. Not every fact about a reaction can be proved
at the outset. Certain fundamental facts are always util-
ized by chemists. For our present purpose, we must re-
gard zinc as an element and sulphuric acid as a compound
of hydrogen, sulphur, and oxygen. Now in this reaction
the zinc and hydrogen exchange places, and the result is
the formation of hydrogen — already shown — and zinc
sulphate. Furthermore, we can prove that a definite rela-
42 Experimental Chemistry.
tion by weight exists between the zinc used and hydrogen
produced. The present problem, then, has two steps: (i)
the identification of zinc sulphate as the second product of
the reaction, and (2) the actual determination of the re-
lations by weight between the interchanging zinc and
hydrogen.
Tests. — The identity of an element or a compound is
established by a Test, i.e. by subjecting it to the action
of various chemicals and comparing the results with the
recorded properties of the substance under examination.
Sometimes the test is a single operation, sometimes several
different examinations are necessary. The test for zinc
sulphate involves two distinct steps: (i) the identification
of the group to which it belongs — the sulphate group, and
(2) the determination of the kind of sulphate, i.e. the
establishment of those properties which distinguish zinc
sulphate from all other members of the sulphate group.
If the substance in question does not show these properties,
then, of course, it is not zinc sulphate, and other tests must
be applied to determine its nature.
Crystallization. — Most substances to be tested must first
be purified. A convenient method of purification is crystal-
lization. Most solid substances are more soluble in hot
than in cold water. Hence a hot liquid on cooHng will
often deposit some of the dissolved solid, especially if the
liquid contains much solid in solution. If the hot liquid
is cooled slowly, the solid is deposited in masses having
a more or less definite geometrical form, called crystals.
This operation is called Crystallization, and the substance
is said to Crystallize. If a solution evaporates slowly in
the air, crystals will form when the solution reaches the
Hydrogen. 43
same condition as a cooled solution, i.e. a point where
the liquid cannot hold all the solid in solution. This point
must be found by trial, hence no rules can be given to
obtain well-shaped crystals, except perhaps (i) slow evapo-
ration in the air, if time permits, or (2) frequent cooling,
if the evaporation is caused by heating. When a substance
crystallizes, most of the impurities remain behind in solu-
tion, hence the crystallized solid is usually pure. Repeated
crystallization — recrystallization — is a common method
of purification, whenever it is applicable. Usually every
chemical substance has one or more distinct geometrical
forms in which it crystallizes. This crystal form is also a
means of identification, but the chemical tests are more
decisive.
CRYSTALLIZATION OF ZINC SULPHATE.
Experiment 20. — Fold a filter paper according to the directions
given on page 8, fit it into a glass funnel, and moisten the paper with
water to hold it in place. Support the funnel by a ring or filter stand,
and pour the contents of the hydrogen generator gradually upon the
filter, taking care that no liquid comes above the edge of the paper.
(See App. A, § 4.) If the contents of the generator consists partly of
a white solid, add a little hot water and shake the mass until the solid
is dissolved. The filtrate should be caught in a casserole, or evaporating
dish. Wash any remaining zinc with water to remove the acid, and
preserve for future use. When the filtrate measures about 100 cc. place
the casserole on a gauze-covered tripod, or similar support, and heat the
liquid for ten or fifteen minutes, but do not allow it to boil violently.
(See App. A, §§ 2, 3.) Then cool it slowly, and crystals of zinc sul-
phate will be deposited. If no crystals separate after a few hours,
evaporate again until a thin film appears on the surface when the solution
has cooled. Then if none are deposited from the cold solution, con-
sult the teacher. If for any reason the filtrate cannot be immediately
evaporated, let it remain undisturbed, and crystals will be deposited
as the water evaporates. This liquid, of course, should not be boiled.
In any case remove the crystals and dry them between filter paper.
Describe them, giving color, luster, approximate shape and size, and
44 Experimental Chemistry.
any other striking property. Put any remaining crystals into the stock
bottle marked Zinc Sulphate.
TEST FOR ZINC SULPHATE.
Experiment 21. — Dissolve some of the crystals of zinc sulphate
obtained in Exp. 20, in a test tube half full of water, divide the solution
into halves, and proceed as follows : —
(a) Test for zinc: To one part add a little sodium hydroxide solu-
tion. A white jellylike solid will be formed, but it will disappear if
considerable sodium hydroxide solution is added and the mixture well
shaken. The insoluble jellylike compound is zinc hydroxide, which is
changed by the excess of sodium hydroxide into soluble sodium zincate.
Now add a little hydrogen sulphide solution, and a white solid is formed.
It is zinc sulphide. Compare its color and texture with those of zinc
hydroxide. Shake the test tube containing the zinc sulphide to sus-
pend the solid in the Hquid, and divide into halves. To one add dilute
hydrochloric acid, and to the other add acetic acid. Shake each tube
and observe the results. Zinc sulphide is decomposed by hydrochloric
acid, but not by acetic acid. These facts are characteristic of compounds
containing zinc, but not true of compounds of other metals. Hence
zinc must be one of the components of the solid product of the reaction
between zinc and sulphuric acid.
{b) Test for a sulphate: To the other half of the original solution
add a little barium chloride solution. A white solid is formed, called
barium sulphate, which will be found to be insoluble, in both hydro-
chloric and acetic acids. Try it. Barium sulphate is always formed when
barium chloride is added to the solution of a sulphate, or to sulphuric
acid, which is closely related to all sulphates. Hence the zinc com-
pound formed in the hydrogen generator must be a sulphate.
{c) Draw a final conclusion from the results obtained in («) and {b).
Definitions. — The test made in Exp. 21, which showed
that zinc sulphate is the compound formed by the reaction
which liberates hydrogen, is called a Qualitative test, since
it determined only the quality of the compound, not the
quantity by weight or volume of each component. The
latter operation is called a Quantitative test. The replace-
ment of one element or group of elements by another ele-
Hydrogen.
45
ment or group of elements is called Substitution. Thus,
the formation of zinc sulphate is the substitution of zinc
for hydrogen in sulphuric acid. So also the formation of
barium sulphate consists in the substitution of barium for
zinc in the zinc sulphate ; the zinc, as could be easily
proved, takes the place left by barium in the barium chlo-
ride. Substitution is a broad, complex subject, which has
a deeper meaning than mere replacement. It will be
clearer as the experiments proceed. A Sulphate is a com-
pound derived from sulphuric acid by substituting a metal
for the hydrogen in the acid. A Sulphide is a compound
of sulphur and ojie other element or group acting like an
element. Sulphides correspond to oxides and chlorides —
all being binary compounds, i.e. compounds of
two elements or groups acting like elements.
QUANTITATIVE EXAMINATION OF THE REACTION
BETWEEN ZINC AND SULPHURIC ACID.
Experiment 22. — MetJwd:
Measure the hydrogen liberated
from sulphuric acid by a known
weight of zinc.
Apparatus : The construc-
tion and arrangement of the
apparatus is shown in Fig. 32.
^ is a large test tube provided
with a two-hole rubber stopper.
The tube B passes through one
hole to the bottom of the test
tube and is connected by a
short piece of pressure tubing
with the cup D. A Mohr's
pinchcock, C, controls the flow
of acid from D into A. If the
pinchcock is not strong enough
Fig. 32. — Apparatus for determining the
quantitative relation between zinc and
hydrogen.
46 Experimental Chemistry.
to close the pressure tubing, then ordinary rubber tubing will answer,
if fresh. The part DB is made by cutting a thistle tube into the
requisite lengths. The ends of B and D are as
close together as the pinchcock will permit. F is
n a shallow dish, and 6^ is a 100 cc. graduated tube.
The modification shown in Fig. 33 is more conven-
ient, since it is not easily broken and requires no
support.
Process : Select a single piece of zinc weighing
not less than .22 nor more than .24 gm., weigh it
exactly, wind a short piece of platinum wire around
it and drop it into A. Fill A full of water and
insert the stopper with its tubes. Fill the w/iok
apparatus with water thus : Fill /J and admit water
repeatedly until all air is forced out of A, B, and
FIG. 33.-Alternate ^^^^^ ^^ j^^ ^^^ ^^^^^^ j^ ^ ^^j^ ^^_
apparatus. ' =>
low the point indicated by the dotted line. Fill F
nearly full, fill G full and invert it in F over the end of F. All water
used in this experiment must have stood in the laboratory at least a day.
Fill the cup D with hot dilute sulphuric acid, and let the acid run
slowly into A, but never below the dotted line. Introduce in this way
and at this time about 50 cc. of acid. Do not add acid after the action
begins unless absolutely necessary, as the zinc may be forced over
into F. Hydrogen will be liberated and will pass through F up into G.
Let the gas accumulate in A at the top, and there will be little danger
of losing the zinc. Let the action continue until the zinc disappears,
then force over into G any gas in the apparatus by admitting water,
observing the precautions previously given.
Unclamp G, close the end with the second finger, remove it to a tall
jar of water which has been standing long enough to assume the tem-
perature of the room, and clamp it so that the water is about the same
height within and without the tube. Stand a thermometer in the water
and allow the whole to remain undisturbed for at least fifteen minutes,
and longer if possible. Then, without touching the tube where it con-
tains gas, adjust the height so that the meniscus just touches the sur-
face of the water in the jar. (See App. A, § 22, (i), (2).)
Read the gas volume, the thermometer, and the barometer, observing
all the precautions mentioned in App. B, II., (3) and (4). Record
numerical results as shown below. Perform the calculation as directed
Hydrogen. 47
ft
below, and, if possible, before the tube G is used again. The latter
precaution may obviate a repetition of the experiment.
Calculations : (The figures recorded here are only for illustration.
Those actually obtained will probably be different.)
Data : —
Weight of zinc, 0.24 gm.
Volume of hydrogen, 88.5 cc.
Temperature, 14° C.
Barometer, 758 mm.
Aqueous tension, 11.9.
Reduce the observed volume of hydrogen to the volume it would
occupy at 0°, 760 mm., and in the dry state by the formula —
„_ V{P-a^
760 {I +.00366^)
If the proper values are substituted in this formula, the corrected vol-
ume of hydrogen is found to be 82.64 cc. A liter (1000 cc.) of dry
hydrogen under standard conditions weighs .0896 gm. Hence the
weight of 82.64 cc. is found by solving the proportion —
1000 : 82.64 '• ' -0896 : X
X = .0074 gm.
Weight of corrected gas volume = .0074 gm.
Since .0074 gm. of hydrogen is replaced by .24 gm. of zinc, the
weight of zinc which will replace i gm. of hydrogen is found by the
proportion
.0074 : .24 II I : X
:r= 32.43.
Therefore 32.43 gm. of zinc are chemically equivalent to i gm. of
hydrogen.
Class average =
Final Directions : Wash the whole apparatus free from acid. Pre-
serve the platinum wire for future use. If the final result is incorrect,
first examine the arithmetical work, then repeat the readings (gas volume,
pressure, and temperature), and recalculate, if blunders were made.
48 Experimental Chemistry.
Discussion of Experiment 22. — The result found in this
experiment should be approximately 32.5. The number,
32.5, is called the equivalent of zinc. The Equivalent of
an element is the weight of that element which will replace
one part by weight of hydrogen. The equivalent of zinc
is the same whatever the acid used. The fact that an
average of independent results, such as would be obtained
by a class, agrees within the limits of experimental error
shows that definite relations by weight exist between
hydrogen and zinc, when zinc and sulphuric acid interact.
This definite and fixed relation has a deeper meaning than
can be considered at this stage of the subject. Its signifi-
cance will be explained after more facts have been accu-
mulated. Just now, it serves as an introduction to the
quantitative study of chemical equations.
Burning Hydrogen. — It was shown in Exp. 19 that
hydrogen burns, but the product was not examined. It
was shown in Chapter I. that when metals are burned in
air, they combine with oxygen.
PRODUCT OF BURNING HYDROGEN.
Perform this cxperimciit with the utmost precaution,
since a mixture of air and Jiydrogen explodes violeiitly, if
ignited. Proceed exactly according to the directions.
Experiment 23. — Method: Generate hydrogen, and examine the
product, from a burning jet.
Supplies : Hydrogen generator used in Exp. 19, U-tube with two one-
hole stoppers, platinum tip, small bottle, calcium chloride.
Apparatus: The construction and arrangement is shown in Fig. 34.
The lower portion of the delivery tube is removed and the generator is
then connected directly with the U-tube, which is filled about two-thirds
full of lumps of calcium chloride separated from the stoppers by plugs
Hydrogen.
49
Fig. 34.-
K:^
Apparatus for burning
hydrogen.
of cotton. The delivery tube is bent and attached to the U-tube as
shown in the figure. The platinum tip, which is made by the method
given in App. A, § 13, is attached to the delivery tube by a short,
gas-tight^ rubber connector. All joints and connections must be gas
tight.
Process: Pour slowly but con-
tinuously through the safety tube
enough (about 50 cc.) dilute sul-
phuric acid upon at least 25 gm.
of granulated zinc to produce a
steady current of hydrogen gas for
about eight minutes. It is advis-
able to use considerable zinc and
a moderate amount of acid. Acid
must not be added after the evolu-
tion of gas begins, unless, of course,
the experiment is begun anew. Let
the gas bubble through the acid for at least two inimiies by actual
observation, then attach the platinum tip to the rubber connector at
the end of the delivery tube, leaving a short space between the ends
of the two glass tubes so that the rubber tube may be compressed
suddenly, if necessary. Let the gas run for another full viinute.
This latter precaution is to drive all air out of the tip. Light the
hydrogen, and observe at once the nature of the flame, its color,
heat (by holding a match, or copper, or platinum wire over it), and
any other striking property Then hold a small dry bottle over the
flame in such a position that the flame is just inside the bottle.
When conclusive evidence of the product of burning hydrogen is seen
inside the bottle, remove the bottle, and extinguish the flame at once by
pinching the rubber connector. Remove the generator to the hood,
and if the evolution of hydrogen is still brisk, dilute the acid by pour-
ing water through the safety tube. Examine the inside of the bottle.
What in ail probability is the deposit ? Explain its formation. What
previous experiments does this one resemble ? What experiment would
have enabled you* to predict the result in this experiment ?
Sketch in your note-book the apparatus used in this experiment.
Final Directions : When the apparatus is taken apart, replace the
stoppers of the U-tube with corks to preserve the contents for further
use. Insert the glass end of the platinum tip into the hole at the smaller
50 Experimental Chemistry.
end of a rubber stopper or cork and insert it into a test tube. This
device will protect the platinum end of the tip. Pour off the acid
from the zinc, wash the zinc with water to remove all acid, and preserve
for further use.
Discussion of Experiment 23. — The product of burning
hydrogen is water, though the final proof cannot be made
until the composition of water is studied.
LABORATORY EXERCISE. IV.
1. Calcium chloride absorbs water. Explain the use of the U-tube
in Exp. 23.
2. Is water an oxide ? Why ?
3. How does the heat of the hydrogen flame compare with its
luminosity ?
Weight of a Liter of Hydrogen. — Hydrogen is the lightest
known substance. A liter at 0° C. and 760 mm. weighs
only .0896 gm., while a liter of oxygen under the same
conditions weighs 1.43 gm., as was found by Exp. 17.
The number .0896 must be remembered. Hydrogen is
usually taken as the standard of density in the case of
gaseous elements and compounds, though air is some-
times so regarded. Thus the density of oxygen on the
air standard is 1.105, and on the hydrogen standard is
approximately 16, —
.0896: 1.43 :: I : 16.
The significance of this relation will subsequently appear.
CLASS-ROOM EXERCISE. VII.
I. Additional study of hydrogen.
(a) Distribution.
(d) Compounds of hydrogen thus far studied or used.
(^) Uses.
(d) Liquid hydrogen.
(e) Solid hydrogen.
Hydrogen. ^i
2. History of hydrogen.
{a) Discovery.
(d) Essential facts of Cavendish's life.
(c) Names of the gas, by whom so called, and why ?
3. Som-ces of error in Exp. 22 ; how detected, avoided, and remedied.
(a) Weighing the zinc.
(d) Reading the graduated instruments.
(c) Blunders in the calculation.
4. Compound blow-pipe.
(a) Structure (with outline sketch).
{d) Action.
(c) Properties of its flame.
(d) Precautions in using it.
(e) Uses.
5. Lavoisier's experiment.
Lavoisier passed steam through a red-hot tube containing iron
filings. A combustible gas and a black solid were obtained.
(a) What is the name of the gas ?
(d) What in all probability are the components of the black
solid ?
(c) Devise an apparatus for this experiment.
(d) Read the historical value of this experiment in a history
of chemistry.
6. Impurities are often removed from hydrogen gas by passing it
through a solution of potassium permanganate. This compound con-
tains a relatively large proportion of oxygen, which under certain con-
ditions is easily liberated. What must be the general character of the
change in the removal of the impurities ?
7. Crystallization and precipitation.
(a) Favorable conditions for crystallization.
(d) Why are perfect crystals rare ?
(c) How does precipitation differ from crystallization ?
(d) The distinguishing characteristics of the six crystal
systems.
(<?) Define "crystallized," "amorphous," "dimorphous,"
"pseudomorph.'"
(/) Define "precipitate " and "precipitation."
r2 Experimental Chemistry.
8. Miscellaneous.
(a) What is effervescence, and how does it differ from boil-
ing ? What experiments in this chapter illustrate
effervescence ?
(d) What is meant by the "diffusion of gases'' ? What
part of Exp. 19 illustrates diffusion ?
(c) Describe a generator, a safety tube, a delivery tube, a
tripod.
(d) What synthesis was made in one or more of the ex-
periments with hydrogen ? What oxidation ? What
determination ? What chemical compounds addi-
tional to those in previous experiments were made or
used ? What new elements (free or combined) were
prepared or used ?
PROBLEMS. VIII.
1 . Calculate the equivalent of zinc from the following data actually
obtained by a student : —
Weight of zinc .12 gm.
Corrected volume of hydrogen 40.89 cc.
2. Calculate the equivalent of zinc from —
Weight of zinc .50 gm.
Observed volume of hydrogen 183 cc.
Pressure 748 mm.
Temperature 9° C. (a = 8.57.)
3. What is the weight of 500 cc. of dry hydrogen at o" C. and
760 mm.? Of 1800 cc? OfQl.?
4. How many grams of zinc will liberate 100 gm. of hydrogen from
sulphuric acid ? How many grams of zinc will liberate 100 1. ?
5. What is the weight of a liter of hydrogen measured over water
at 50" C. and 790 mm. ? {a — c)\ .98.)
6. How much zinc (in grams) is necessary to liberate from sulphuric
acid 100 1. of dry hydrogen at 9i°C. and 800 mm.?
7. How many times heavier than a liter of hydrogen is a liter of
oxygen, both being dry and under standard conditions ?
CHAPTER IV.
SYMBOLS — FORMULAS — CONSERVATION 01 MATTER —
CHEMICAL EQUATIONS— QUANTITATIVE INTERPRE-
TATION OF EQUATIONS — PROBLEMS BASED ON EQUA-
TIONS—PROBLEMS.
Symbols. — A symbol is an abbreviation of the name of
a chemical element. It is usually the first letter of the
name of the element. Thus O is the symbol for oxygen,
H for hydrogen, N for nitrogen ; as more than one element
has the same initial letter, another letter in some cases is
added. Thus B is the symbol of boron, but Ba of barium,
Bi of bismuth, Br of bromine, etc. The symbol of several
metals is derived from their Latin name. Thus, the sym-
bol of iron is Fe {fei'nnn), of lead is Pb {plinnbum), of
sodium is Na {natrinui), of potassium is K {kaliuni), etc.
These symbols denote one atom of the element, i.e. H
means the smallest particle of hydrogen which can take
part in a chemical change. They also represent a certain
number, called the atomic weight, which belongs to that
atom. They are sometimes loosely used to mean any
convenient amount of the element. A list of symbols
is given in Appendix C, Table I. If more than one
atom is to be designated, the proper numeral is placed
before the symbol. Thus : —
2 O means 2 atoms of oxygen.
3 H means 3 atoms of hydrogen. YA\5\ ^^ ^ \ kt*^
4 P means 4 atoms of phosphorus, r \ \V)6^
rk<%c
\(\ft
j4 Experimental Chemistry.
But if the atoms are in chemical combination, either
with themselves or other atoms, then a small numeral is
placed after and a little below the symbol. Thus : —
H2 means 2 atoms of hydrogen in combination.
Ng means 3 atoms of nitrogen in combination.
P4 means 4 atoms of phosphorus in combination.
Formulas. — A formula is a group of symbols which
expresses the composition of a compound. Thus, KCIO3
is the formula of potassium chlorate, and MnOg of
manganese dioxide. Formulas represent single molecules.
Thus KCIO3 means one molecule of potassium chlorate
containing one atom each of potassium and chlorine and
three atoms of oxygen. They also represent the molecular
weight of the compound. More than one molecule is
designated by the proper numeral placed before the formula,
thus : —
2 KCIO3 means 2 molecules of potassium chlorate.
3 HgO means 3 molecules of water.
4 H2SO4 means 4 molecules of sulphuric acid.
Groups of atoms acting like a single atom in reactions
are enclosed in a parenthesis, or separated by a period.
Thus (NH4)N03 represents a molecule of ammonium
nitrate, in which the group of atoms (NH4) bears such
relations to the NO3 that the group is represented as a
unit; and CgH^ • OH similarly represents a molecule of
alcohol. A group of atoms to be multipHed is also
enclosed in a parenthesis. Thus, Pb(N03)2 is the formula
of lead nitrate, and means that the (NO3) group bears
to the atom of lead (Pb) such relations that the group
must be multipHed by two. Many facts similar to the
last will be clearer as the experiments proceed.
Conservation of Matter — Chemical Equations. 55
Symbols and formulas have a deeper significance, which
will be considered as the necessary facts accumulate.
CLASS-ROOM EXERCISE. VIII.
1 . Give the symbol of each of the following elements, all of which
have been used or studied in the preceding experiments : —
Oxygen, Hydrogen, Sulphur, Carbon, Iron, Magnesium, Zinc, Mer-
cury, Platinum, Lead.
2. Name the elements which correspond to the following symbols : —
Fe, Pb, Pt, Zn, Hg, Na, K, Cu.
3. Give the formula of each of the following compounds : —
Magnesium Oxide, Mercuric Oxide, Potassium Chlorate, Potassium
Chloride, Sulphuric Acid, Hydrochloric Acid, Zinc Sulphate, Sulphur
Dioxide, Carbon Dioxide.
Conservation of Matter. — Most experiments transform
matter, but they never destroy or create it ; no weight is
gained or lost. The total weight of matter involved in a
reaction is the same before and after. This vast conception
was first demonstrated by Lavoisier, and, largely as a result
of his initial work, it has become a fundamental law of
chemistry. It is called the Law of the Conservation of
Matter, and is often stated thus: —
No weight is lost or gained in a cJiemical reaction.
Chemical Equations. — The mathematical significance of
symbols and formulas, together with the conception of the
conservation of matter, allows reactions to be represented
quantitatively by equations. Chemical equations differ
from algebraic equations in one vital respect: they are
the result of experiment. Thus in Exp. 15 it was shown
that magnesium unites with oxygen to form magnesium
^6 Experimental Chemistry.
oxide. These facts are represented in the simplest way,
thus : —
Mg + O = MgO (i)
Magnesium Oxygen Magnesium
Oxide
This equation means that one atom of magnesium unites
with one atom of oxygen, and forms one molecule of
magnesium oxide In Exp. i6 it was shown that potassium
chlorate when heated yields oxygen and a residue (called
potassium chloride). This reaction is represented in the
simplest way by the equation : —
KCIO3 = 3O 4- KCl (2)
Potassium Oxygen Potassium
Chlorate Chloride
This equation means that one molecule of potassium chlo-
rate yields by decomposition three atoms of oxygen and
one molecule of potassium chloride. It has been shown as
the result of several experiments that the interaction of zinc
and sulphuric acid produces hydrogen and zinc sulphate.
This fact is represented by the equation : — -
(3)
Zn
+
H2SO4
= H2 +
ZnS04
Zinc
Sulphuric
Acid
Hydrogen
Zinc
Sulphate
This equation means that one atom of zinc interacting
with one molecule of sulphuric acid produces two (com-
bined) atoms of hydrogen and one molecule of zinc sulphate.
These equations have not been completely proved by
the particular experiments, but the work of others sup-
plements the results obtained and permits important
conclusions.
Quantitative Interpretation of Equations. 57
It should be noted —
(i) That these equations are expressions between atoms
and molecules, not between the quantities actually seen in
action.
(2) That some substances are not included in the equa-
tion. Thus in equation (3) water is not represented because
it takes no chemical part in the reaction ; it simply serves
to dissolve the zinc sulphate from the surface of the zinc,
and thereby allow the reaction to proceed. Likewise in
equation (i) no nitrogen appears. Nitrogen combines with
magnesium only when the latter is at a red heat in the
absence of oxygen — obviously impossible in Exp, 15.
(3) That only the beginning and end of reactions are
represented. Thus, in Exp. 9, certain compounds are
formed and decomposed again as the temperature in-
creases ; but these facts do not appear in the equation,
because here we are concerned not with phases, but only
with the actual ultimate decomposition of potassium chlo-
rate into oxygen and potassium chloride.
(4) That chemical equations /;'<?7'^ nothing. They simply
help us interpret an experiment. They are not a sufficient
record of an experiment, and they should never be used
unless they have been proved in whole or in part.
The atoms or molecules entering into the initial stage
of a reaction are called factors, those present in the final
stage are called prodjicts. Thus in equation (3) those atoms
and molecules at the left of the equality sign are factors,
those at the right are products. Chemical equations are
read from left to right, and are very rarely reversible.
Quantitative Interpretation of Equations. — Since matter
cannot be destroyed, all weight entering a reaction can be
accounted for. Much more information, however, lurks
58 Experimental Chemistry.
in an equation. It was found by Exp. 15 that when
magnesium is heated in air, the ratio of the weights in
which magnesium and oxygen combine is 1.5:1. Equa-
tion (i) might, therefore, be written —
Mg + O = MgO
1.5 + I =2.5
Any number might be substituted for i, provided, of
course, the ratio 1.5 : i is not destroyed. Hence the above
equation might be written —
Mg + O = MgO
24 + 16 = 40
This is the simplest kind of an equation, but the facts
similar to those revealed here may be discovered about
every equation. Equation (3) might be written —
Zn + H2SO4 = H2 + ZnS04
65+98 =2 + 161
Experiment 22 proved that 32.5 (approximately) gm.
of zinc correspond to i gm. of hydrogen. If the work
had been extended, it would have been found that 65 gm.
of zinc correspond to 161 gm. of zinc sulphate. The
equation would then become, as the result of experiment, —
Zn + H2SO4 = H.3 + ZnS04
65 + 98 = 2 + 161
163 163
Finally, in Exp. 16 it was shown that potassium chlorate
yields about 39 per cent of oxygen. If we select for 3 O
Other Chemical Equations Studied. ^9
the number 48 (equal to 3 x 16), then we can form the
proportion
48 : 39 : : .r : 1 00
X = 123
If the experiment should be conducted with the utmost
care and precaution, the proportion would be slightly modi-
fied, so that instead of equalling 123, x would have equalled
122.5, a number nearer the true relation existing between
potassium chlorate and the oxygen it yields. Equation (2)
may now be written —
KCIO3 = 3 O + KCl
122.5 = 48 + 74.5
From the above discussions it is evident that equations,
accurately evolved and correctly interpreted, reveal the
exact proportions by weight in which compounds react in
the particular chemical change expressed by the equation.
Other facts of fundamental importance are contained in
chemical equations, and these will be considered in the
proper place.
Other Chemical Equations Studied. — Experiment 12 was
a qualitative experiment, but if the factors and products
had been weighed, the final result might have been ex-
pressed thus : —
HgO = Hg + O (4)
216 = 200 + 16
This means not only that mercuric oxide decomposes
into mercury and oxygen, but that 216 parts by weight of
mercuric oxide produce 200 parts of mercury and 16 parts
of oxygen, since careful experiments show that these num-
bers represent the correct proportions.
64
64
c
12
+ O2
+ (2 X
16) =
12
CO2
+ (2 X
16)'
60 Experimental Chemistry.
Again, when sulphur and charcoal were burned in
oxygen in Exp. 14, the final results might be expressed
thus : —
3 + 0^= SO, (5)
32 + (2 X 16) = 32 + (2 X 16)
(6)
44 44
Equation (5) is correct, because many experiments prove
that when 32 parts of sulphur are heated in oxygen, they
combine with 32 parts by weight of oxygen, and, further-
more, the compound formed, called sulphur dioxide, yields
by decomposition the same proportion of sulphur and
oxygen by weight. Equation (6) is true for similar
reasons. The experiment has been tried repeatedly, and
the results show that carbon burning in an excess of oxy-
gen always combines with 32 parts of oxygen for every
12 parts of carbon. The product of the reaction, called
carbon dioxide, has been analyzed many times, and when
analyzed always yields carbon and oxygen in the ratio of 12
to 32 parts by weight.
Problems Based on the Foregoing Equations. — In the
experiments which led to the above equations, no special
weight of matter was used. The equation simply states
the proportions which participate in the reaction. Obvi-
ously any convenient weight of magnesium might be
heated in the air, or any weight of potassium chlorate
might be decomposed, if necessity or choice prevailed, but
the proportions of the factors and products are always the
same as the proportions which accurate experiments show
Problems Based on Foregoing Equations. 6 1
are represented by the numbers given in the equation.
Thus in equation ( i ),
Mg + O = MgO,
24 + i6 = 40
if 12 gm. of magnesium are taken, then 8 gm. of oxygen
will be necessary to convert all the magnesium into mag-
nesium oxide, which will weigh 20 gm. — the sum of the
weights of the two factors. Hence from the weight of one
substance participating in a reaction, we can calculate by
proportions the actual weights of all other participants, if
we know the equation corresponding to the reaction.
Suppose 17 gm. of magnesium are to be converted into
magnesium oxide, and it is desired to know (a) how much
oxygen is necessary and (l?) how much magnesium oxide
will be formed. The questions are answered by solving
the following proportions : -
24 : 16 : : ly : X.
X = 1 1.3 gm. of oxygen.
24 : 40 : : 17 : x.
X — 28.3 gm. of magnesium oxide.
Similarly, when 45 gm. of zinc interact with sulphuric
acid, the weights of {a) acid required, {b) hydrogen formed,
and {c) zinc sulphate produced are found by the propor-
tions: —
65 : 98 : : 45 : ^•
X = 6'].'^ gm. of sulphuric acid.
65 : 2 : : 45 : .r.
X = 1.38 gm. of hydrogen.
65 : 161 :: 45 '. x.
X = 1 1 1.4 gm. of zinc sulphate.
62 Experimental Chemistry.
PROBLEMS. IX.
1. How many grams of oxygen can be prepared from 122.5 S^- ^^
potassium chlorate? From 245 gm.? From 421 gm.?
2. How much hydrogen can be prepared from 65 gm. of zinc?
From 130 gm. ? From 297 gm. ?
3. How much zinc is needed to prepare 2 gm. of hydrogen ? 4 gm. ?
17 gm.?
4. How much zinc sulphate can be prepared from 98 gm. of sul-
phuric acid? From 196 gm.? From 427 gm. ?
5. How much mercury and how much oxygen is obtained by heat-
ing 10 gm. of mercuric oxide?
6. How much oxygen can be prepared from 50 gm. of potassium
chlorate?
7. A certain weight of potassium chlorate was heated and 298 gm.
of potassium chloride remained. What weight of potassium chlorate
was heated and what weight of oxygen was formed ?
8. A certain weight of potassium chlorate was heated until com-
pletely decomposed, and the residue weighed 20.246 gm. What
weight of potassium chlorate was heated, and how much oxygen was
evolved?
9. If 60 gm. of mercuric oxide are completely decomposed, what
volume of oxygen is obtained at 91" C. and 380 mm. ?
10. If 400 gm. of potassium chlorate are completely decomposed,
what volume of oxygen is obtained at 27° C. and 760 mm.?
11. If 100 gm. of zinc interact with sulphuric acid, what volume of
hydrogen is evolved ?
12. A balloon holds 132.74 kg. of hydrogen. How much zinc and
sulphuric acid are needed to produce the gas ?
13. How much potassium chlorate is needed to prepare 36.48 1. of
oxygen at 15° C. and 750 mm.?
14. How many liters of oxygen can be obtained from a kilogram of
potassium chlorate, if the gas is measured at 10° C. and 755 mm. ?
15. How much zinc and sulphuric acid (in kg.) will yield hydrogen
enough to fill a 350 cc. flask at 15° C. and 735 mm.?
16. If 12 gm. of carbon are burned in the oxygen obtained by de^
composing 122.5 g"^- of potassium chlorate, what weight of carbon
dioxide is formed, and what weight of oxygen, if any, remains?
Problems. 6;^
17. A lump of carbon weighing 24 gm. is burned in air.
(a) What weight of carbon dioxide is formed ?
(d) What weight of oxygen is needed?
{c) If a liter of oxygen weighs 1.43 gm., what volume of oxygen
is needed?
18. What weight of carbon dioxide may be obtained by burning
112 lbs. of coal containing 15 per cent of impurities?
19. A sulphuric acid maker accidentally set fire to a storehouse con-
taining 8794 kg. of sulphur. A rough estimate showed that 67 per
cent had burned. What weight of sulphur dioxide was formed?
20. Calculate the weight of oxygen necessary to burn 731 gm. of
sulphur containing 15 per cent of impurities.
21. A lump of sulphur weighing 32 gm. is burned in air. Calculate —
(a) The weight of oxygen required.
(/^) The weight of sulphur dioxide formed.
CHAPTER V.
GENERAL PROPERTIES OF WATER.
Water is the most abundant of all chemical compounds,
and on account of its remarkable and varied properties it
demands extensive study.
GENERAL DISTRIBUTION.
Experiment 24. — Supplies: Test tubes, wood, meat, potato.
Heat successively in dry test tubes a small piece of wood, of meat,
and of potato (or any fresh vegetable). Hold the test tube at such an
angle that any moisture which may be liberated will not flow down
upon the hot glass. Is there conclusive evidence of water? Is the
amount unexpectedly large in any case?
These substances are types of animal and vegetable matter. Draw
a general conclusion.
Fig. 35. — Potato showing (by shaded
part) the proportion of water. Re-
mainder is mainly starch and fiber.
Fig. 36. — Carrot showing (by shaded
part) the proportion of water. Re-
mainder is mainly starch, sugar,
and fiber.
The water driven off by heating many organic substances
may be simply mixed with other ingredients, or it may be
64
General Properties of Water. 65
due to the presence in them of hydrogen and oxygen in
such proportions that these two gases combine to form
water at the instant the compound is decomposed. Many
examples of this latter kind are known and will be sub-
sequently examined.
The wide distribution of water in vegetables is strikingly
shown in Figs. 35 and 36. The human body is 70 per
cent water. Water is essential to life.
Water of Crystallization. — Dry crystals deposited from
the water solution of many solids often contain water,
which seems to be an essential part of the chemical com-
pound. This water is called water of crystallization. In
some crystals the water passes off at the ordinary temper-
ature, and in all crystals when they are heated.
DETECTION OF WATER OF CRYSTALLIZATION.
Experiment 25. — {d) Heat a few small crystals of sodium carbon-
ate or of alum in a dry test tube, inclining the test tube so that any water
liberated may run out. What is the evidence that they contained water
of crystallization ? If there is any marked change in the appearance of
the crystals, describe and explain it.
{])) Repeat, using a crystal of gypsum. Answer the question asked
in {a).
{c) If time permits, examine successively crystals of the following
compounds for water of crystallization, and describe the result in each
case : Potassium nitrate, potassium dichromate, and magnesium sul-
phate.
{d^ Heat two or three small crystals of copper sulphate in an evapo-
rating dish which stands on a gauze-covered tripod. As the action
proceeds, hold a dry funnel or glass plate over the dish. Is there con-
clusive evidence of escaping water of crystallization? Do the crystals
change in color? In shape? Can the form of the crystals be changed
by gently touching the mass with a glass rod? Continue to heat until
the resulting mass is a bluish gray. Let the dish cool. Meanwhile
heat a test tube two-thirds full of water. When the dish has cooled
66 Experimental Chemistry.
somewhat, pour tlie hot water slowly into the dish upon the copper
sulphate. Explain the change in color, if any. If there are any lumps,
crush them with a glass rod and heat the solution over the gauze until
it is clear. Let the solution evaporate for several hours. Are crystals
deposited? If not, heat a few minutes, and cool again; if so, why?
Have they water of crystallization, and, if so, when did they get it ?
The amount of water of crystallization in crystals is not
arbitrary. It is constant in the same compound when
crystallized under uniform conditions, but the amount varies
between wide limits in different substances. Its necessity
in some crystals and not in others, as well as its varying
quantity, has never been explained.
DETERMINATION OF WATER OF CRYSTALLIZATION IN BARIUM
CHLORIDE.
Experiment 26. — MetJwd: Heat to constant weight a weighed
quantity of crystallized barium chloride in an uncovered crucible, and
the loss will be the water of crystallization.
Supplies: Porcelain crucible and support, crystallized barium chlo-
ride, forceps, crucible block.
Process: Weigh accurately a clean, dry crucible. Weigh in it from
1.5 to 2 gm. of barium chloride, weighing exactly the amount taken.
Record the results thus : —
Grams.
Weight of crucible and barium chloride . .
Weight of crucible
Weight of barium chloride
Stand the crucible on a pronged tripod or similar support, and heat
with a low flame for about fifteen minutes, then gradually increase the
heat, as in previous experiments, until the flame is just below the edge
of the crucible. Continue the intense heat for ten minutes, then gradu-
ally lower the flame until it just touches the bottom of the crucible.
Move the flame to and fro under the crucible, until all danger of crack-
ing disappears. Stand the crucible in the crucible block, and weigh
when cool enough to bear the hand. Record the results thus : —
General Properties of Water. 67
Grams.
Weight of crucible and crystallized barium chloride . .
Weight of crucible and barium chloride
Weight of water of crystallization
To be sure that all water has been driven off, heat again strongly for
five minutes, then cool and weigh, as before. If there is an appreciable
loss, repeat until the last two weights are constant. If there is no
appreciable loss upon the second heating, calculate the per cent of
water lost.
Per cent of water lost =
Class average =
PROBLEMS. X.
1. Calculate the per cent of water of crystallization in a sample of
barium chloride, 1.5 gm. of which lost, on heating, 0.22 gm.
2. A student found that 2 gm. of barium chloride lost, on heating,
0.295 g'""- What per cent of the sample was water of crystallization ?
Definitions. — Crystals which have lost water of crystal-
lization are said to be Dehydrated or Anhydrous. Thus the
bluish white product obtained by heating the blue crystal-
lized copper sulphate is usually called " dehydrated copper
sulphate." The words dehydrated and anhydrous have
been extended to mean any substance from which water
has been removed, as anhydrous alcohol or ether. The
opposite term Hydrated is sometimes applied to a compound
to emphasize the fact that it contains water of crystallization.
The term Dissociation is applied to that special kind of
decomposition in which the components may reunite under
favorable conditions. If crystallized copper sulphate is
heated, water, which is driven off by heat, readily re-
unites with the copper sulphate to form the blue crystal-
lized compound, which, strictly speaking, is copper sulphate
united with water, though it is loosely called copper sul-
68 Experimental Chemistry.
phate. Crystallized copper sulphate, therefore, is said to
dissociate when heated. There are many different varieties
of this kind of chemical action.
Efaorescence. — ^Nlany crystals contain water of crys-
tallization so loosely combined that the water passes off
readily when the crystal is exposed to the air. This is
called efflorescence, and the crystals are said to effloresce,
or to be efflorescent.
EFFLORESCENCE.
Experiment 27. — Put a fresh crystal of sodium carbonate and of
sodium sulphate on a piece of filter paper and leave them exposed to the
air for an hour or more. Describe any marked change. What does
this change show about the air ?
It is believed that the water in the crystals called
efflorescent tends to pass off into the atmosphere. This
tendency resembles the tendency of water to evaporate.
And like water, this water of crystallization passes off
only when the vapor tension of the crystal exceeds the
tension of the water vapor in the atmosphere. Thus sodium
sulphate at ordinary temperatures has a greater vapor ten-
sion than the atmospheric water vapor, but gypsum (crys-
tallized calcium sulphate) has a lower vapor tension than
the atmospheric water vapor. Hence sodium sulphate
effloresces, but calcium sulphate does not. This is an easy
way to distinguish the two compounds.
Deliquescence. — Many substances, crystallized and un-
crystallized, absorb water when exposed to moist air, or often
even to ordinary air, which always contains a little water
vapor. This is called deliquescence, and such substances
are said to deliquesce, or to be deliquescent
General Properties of Water. 69
DELIQUESCENCE.
Experiment 28. — Place on a piece of glass or in an evaporating
dish a small piece of calcium chloride, potassium hydroxide, sodium
hydroxide, and potassium carbonate, and leave them exposed to the air
for an hour or more. Describe any marked change which takes place.
How does the action differ from tliat in Exp. 27 ? Where and why was
this property of calcium chloride utilized ?
Deliquescence is a property of substances very soluble
in water. When such substances, potassium carbonate or
calcium chloride for example, are exposed to the air, the
water vapor forms with the substance a small quantity of
a saturated solution. This saturated solution has a lower
vapor pressure than that of the atmosphere, that is, the
water is held by the substance, it does not tend to escape,
hence more water vapor is added from the air, and finally
the substance is entirely dissolved in this condensed vapor.
Common salt or sodium chloride often appears to deliquesce,
but the deliquescence is due to the very soluble magnesium
and calcium chlorides which are usually mixed with com-
mercial sodium chloride. Sodium nitrate is very soluble
in water at the ordinary temperature, but potassium nitrate
is only slightly soluble. Hence potassium nitrate, and not
sodium nitrate, is used in the manufacture of gunpowder.
Impure Water. — Water, if impure, is usually contami-
nated with two kinds of impurities, — -organic and inorganic.
The former is, in general, decomposing animal and vege-
table matter ; the latter consists of various mineral sub-
stances dissolved from the earthy matter through which
the water percolates. The term impure must be inter-
preted broadly, since many mineral zvatei's are whole-
some. " The rule generally adopted by chemists in dealing
70 Experimental Chemistry.
with water is to pronounce any water dangerous which is
contaminated by sewage." (Remsen.)
SIMPLE TESTS FOR IMPURITIES IN WATER.
Experiment 29. — (a) Organic Matter : Fill a clean test tube two-
thirds full of distilled water and another with water containing a little
dirt or a bit of paper. Add to each test tube a drop or two of concen-
trated sulphuric acid and sufficient potassium permanganate solution
(made from distilled water) to color each liquid a light purple, as nearly
alike as possible. Label one tube, and then heat gently nearly to the
boiling point the tube containing the impure water. As soon as a
definite change is seen, heat the other cautiously, as too sudden heat
may cause the liquid to ''bump out.''' Organic matter decolorizes
potassium permanganate solution. Which tube shows the more organic
matter?
(J)) Chlorides : To a test tube half full of distilled water add a few
drops of nitric acid, and then a few drops of silver nitrate solution. Do
the same with faucet water or water known to contain a chloride in solu-
tion. What is the difference between the results? The cloudiness,
or solid, is due to the precipitation of silver chloride, which is always
formed when silver nitrate is added to hydrochloric acid or a chloride
in solution. Silver chloride is soluble in ammonium hydroxide. Try
it. This is the usual test for chlorides (and conversely for soluble sil-
ver compounds), and will hereafter be used without further description.
(c) Sulphates: Apply the usual test for sulphates to distilled water
and to impure water known to contain sulphates (see Exp. 21 (<^))-
{d) Calcium Coiupounds : Add a few drops of a fresh solution of
ammonium oxalate to a test tube half full of clear lime water. Lime
water is a solution of calcium hydroxide, and the white precipitate
formed is calcium oxalate, which is soluble in hydrochloric acid but not
in acetic acid. Try it. This is the test for calcium compounds, often
called " lime " compounds, because lime, which is calcium oxide, is so
well known. Apply this test to distilled water and to water known to
contain calcium compounds, and compare the two results.
Purification of Water. — Water is purified by exposure
to the air, by filtration, or by distillation. The last opera-
tion is convenient only with relatively small quantities, and
General Properties of Water.
71
is performed by means of a condenser, which is shown in
Fig. 37 arranged for use. The condenser consists of an
outer tube, AA', provided with an inlet and outlet for a
Pig. 37, — Condenser arranged for the purification of water by distillation.
current of cold water, which surrounds an inner tube, BB\
These tubes are shown in Fig. 38. The vapor from the
liquid boihng in the flask, C, condenses .in the inner tube,
J
Pig. 38. — Inner and outer tubes of a condenser.
owing to the decrease in temperature, and drops off the
lower end of this tube, as the distillate, into the receiver, D.
The condenser and receiver should, obviously, be scrupu-
lously clean.
72 Experimental Chemistry,
DISTILLATION,
Experiment 30. — Fill the 500 cc. flask, C, half full of watei known
to contain the impurities mentioned in Exp. 29, add a few crystals
(3 or 4) of potassium permanganate, and connect with the condenser
as shown in Fig. 37. Attach the inlet tube to the faucet, fill the con-
denser slowly, and regulate the current so that a small stream flows
continuously from the outlet tube into the sink or waste pipe. Heat
the liquid in C gradually, and when it boils, regulate the heat so that
the ebullition is not too violent. Collect about 25 cc. in the receiver, D.
Test separate portions of this distillate for organic matter, chlorides,
sulphates, and calcium compounds. Is organic matter found ? Is
mineral matter found ? If the liquid in C had contained some volatile
substance like alcohol, ammonia, or hydrochloric acid, where would this
volatile matter have been at the end of the operation ? Can water be
separated from any solid by distillation ? Sketch the apparatus used in
this experiment.
CLASS-ROOM EXERCISE. IX„
1. Additional study of water.
{a) Color, odor, and taste of pure water.
{b) Distribution of water in animal and vegetable matter.
{c) Rain water.
{d) Mineral waters : kinds, general character, medicinal value.
2. Water of crystallization.
{a) Relation to color and form of crystals.
{b) Additional examples.
(<:) Theories.
3. Efflorescence and Deliquescence.
{a) Additional examples.
4.. Drinking water.
(a) Organic matter : source, danger of its presence, how detected
other than chemically, how removed.
{b) Significance of presence of sulphates and chlorides.
{c) Hard water : meaning of term, meaning of '' temporary " and
" permanent," how softened, how estimated, how detected
other than chemically, why objectionable. What causes
boiler scale ? What causes the '' furring " of a kettle ?
General Properties of Water.
73
{d) How is city water filtered ? Of what value are large fil-
ters ? Why should drinking water be exposed to the air ?
{e] Why are lead compounds often found in water ? How
detected ? Why dangerous ?
(/) What is a " water analysis " ?
Distillation.
{u) Literal meaning.
{b) Is it a recent process of purification ?
{c) Is there any dilTerence between water vapor and steam ?
(^) industrial applications.
Experimental.
Some substances have a moist surface and some have water mechani-
cally enclosed within them. How would you distinguish this moisture
from water of crystallization ?
5-
6.
Relation of Water to Temperature. — The freezing point
of water (or the melting point of ice) and the boiUng point
of water are the fixed points on a
thermometer. They are easily
found. They vary, however, under
certain conditions.
THE FREEZING AND BOILING POINTS
OF WATER,
Experiment 31. — Supplies: Ther-
mometer fitted loosely to a cork, iron stand
and clamp, funnel, 250 cc. Erlenmeyer flask
or similar vessel, tripod or iron ring, ice
(or sLLw), salt.
(a) Crush the ice, if snow is not avail-
able, and put it in a funnel which stands
in the flask. Clamp the thermometer
around the cork, as shown in Fig. 39, and
adjust the height so that the bulb is buried
in the ice. After a short time, read the
thermometer. Repeat the reading at brief
intervals until the mercury is constant.
Fig. 39. — Apparatus for deter-
mining the melting point
of ice.
74
Experimental Chemistry.
Record the final reading. When the whole experiment is finished,
sketch this apparatus in the note-book.
{b) Remove the thermometer, carefully transfer the ice to the flask,
and fill the flask half full of water. Stand the flask on a gauze-covered
tripod or ring, and clamp it loosely around the neck. Adjust the
thermometer so that the bulb dips into the liquid. The mouth of the
flask should be open. Read the thermometer. Heat gently and watch
the thermometer. What happens ? Repeat the readings until the
mercury is constant. Record the final reading.
(<;) Remove the thermometer and add from lo to 20 gm. of sodium
chloride — common salt. Adjust the thermometer as before, and find
the boiling point of this solution. If not appreciably different, add
more salt. Record the boiling point.
{d) Make a mixture of equal parts by weight of salt and cmshed
ice, and find, as in {a), the melting point of this mixture. Record the
melting point.
(e) Summarize the results as follows : —
Table of Freezing and Boiling Points.
Substance.
Freezing Point.
Boiling Point.
Found.
Standard.
Found.
Standard.
Water
C.
F.
c.
F.
C.
F.
C.
F.
0
32
100
212
Class Average.
Class Average.
Water
C
F.
C.
F.
and
Salt
CLASS-ROOM EXERCISE. X.
I. Boiling.
{a) Exact definition.
(b) Effect of pressure on the boiling point.
General Properties of Water. 75
(c) Effect of a dissolved substance on the boiling point. Cause
of this effect.
((/) Other conditions affecting the boiling point.
2. Freezing.
(a) Exact definition.
{&) Effect of dissolved substance on the freezing point. Cause
of this effect.
3. Miscellaneous.
(a) Explain the term maxivuun density of water.
ip) What is the scientific reason for using water to extinguish fires ?
{c) What scientific measurements are based on water as a
standard ?
{d) State the changes in volume which occur when (i) ice
melts, (2) water freezes, (3) water is heated from o^ C. to
15° C, (4) water is cooled from 15° C. to o" C.
Solution. — Many solids, liquids, and gases disappear
when put into water. This operation is called dissolving
or putting into solution. The resulting liquid is called a
solution of the substance used. The liquid in which the
substance dissolves is called the Solvent, and the substance
dissolved is called the Solute. ^ If the solute is not volatile,
it may be recovered by evaporation or distillation of the
solution. The degree of solubility is usually expressed by
the terms slightly soluble, soluble, and very soluble. It is
desirable, however, to state in all possible cases the exact
proportions of solvent and solute. Substances which do
not dissolve are called insoluble. A solution which con-
tains a relatively small proportion of the solute is called a
Dilute solution; one containing a relatively large proportion
is called a Concentrated solution. Thus, dilute sulphuric
acid contains usually one volume of sulphuric acid to three
or more volumes of water, while concentrated sulphuric
* See Glossary, Appendix C.
.76
Experimental Chemistry.
acid is nearly 98 per cent acid. The terms dilute and con-
centrated are loosely used in several senses, which will be
clear as the experiments proceed. Other descriptive terms
are applied to solutions, and these will be discussed in the
proper place. Usually a solution means a solution in water,
and the term is so used in this book, unless otherwise
stated.
Solutions of Gases. — There is no general relation be-
tween the solubility of gases and their chemical composi-
tion. Their solubility varies between wide limits.
SOLUBILITY OF GASES.
Experiment 32. — (a) Warm a little faucet water in a test tube. Is
there immediate evidence of a previously dissolved gas? Is there evi-
dence of much gas? What effect has increased heat?
(^) Warm slightly a few cubic centimeters of ammonium hydroxide
in a test tube. Do the results resemble the observations in {a) ? As
soon as the result is obtained, pour the remaining liquid down the sink
and flush well with water.
{c) Repeat {b), using a little concentrated hydrochloric acid. Do
the results resemble those of {a) and {b) ?
The volume of a moderately soluble gas which is dis-
solved by water is directly proportional to the pressure,
if the temperature is constant. This relation is illustrated
by the following: —
Table of Solubility of Carbon Dioxide Gas.
Vol. of Water.
Vol. of Cakbon Dioxide Measured
UNDER Normal Conditions.
Pressure.
I 1. at 0° C.
900 CC.
1800 CC.
3600 CC.
7200 CC.
.5 atmosphere
1 atmosphere
2 atmospheres
4 atmospheres
General Properties of Water.
77
CLASS-ROOM EXERCISE. XI.
1. What is "soda '' water? Why should it be kept cold? Why kept
in a strong vessel? Why so called?
2. Is oxygen soluble in water? Is hydrogen? Is air? (Consult a
text-book regarding the solubility of hydrogen and of oxygen in water.;
3. What is carbonated water? Sulphur water?
4. Why are subterranean gases often easily dissolved by water?
5. Henry's Law of the solubility of gases.
Solutions of Liquids. — The solubility of liquids in water
varies between wide limits. Some liquids, especially oils
and organic liquids, are practically insoluble in water, and
their insolubility is seen by the formation, after agitation,
of two distinct layers of liquid.
SOLUBILITY OF LIQUIDS.
Alcohol, ether, carbon dis nip hide, and kerosene, which are
used in tins experiment, take fire easily. No flames should
be 7iear.
Experiment 33. — {a) Measure into a test tube 10 cc. of water from
a graduate or a pipette. If a pipette is used, proceed as follows : Rinse
out the pipette with water, dip the
pointed end into water and suck up
the liquid nearly to the top, then
remove the pipette from the mouth
and quickly close the top with the
forefinger, as shown in Fig. 40 ; now
cautiously lessen the pressure of the
finger enough to allow the liquid to
sink slowly to the mark around the
stem, then press down tightly, lower
the end of the pipette into a dry
test tube, and remove the finger.
If it is a ID cc. pipette, let all the
water run out ; if not, then let 10 cc.
run into the test tube and reject the
Fig. 40. — Using a pipette. This
illustration shows the correct
position of the hand after the
pipette has been filled with liquid.
78
Experimental Chemistry.
rest. Add to the water about half as much alcohol, and shake. Is
there evidence of solution? Add a little more, and shake. Then
add a third portion. Is there still evidence of solution? Draw a
general conclusion regarding the solubility of alcohol in water.
{b) To 10 cc. of water add a few drops of ether. Observe the result
and draw a conclusion. Add to the water about half its volume of
ether, and shake. Observe the result.
From these two results make a general statement about the solubility
of ether in water.
{c) Repeat {b), using successively kerosene, carbon disulphide, and
glycerine. Observe the results and conclude accordingly.
Summarize the results as follows : —
Table of Solubility of Some Liquids.
Liquid.
Solvent.
1. Alcohol
2. Ether
3. Kerosene
4. Carb. disulph.
5. Glycerine
10 cc.
of Water
at Temperature
of
Laboratory
I.
2.
3-
4-
5-
A rise of temperature in many cases increases the
solubiUty of Uquids in water.
Solutions of Solids. — The sokibility of solids in water is
a subject of vast practical importance, and recent investi-
gations are giving it fundamental scientific value. Its
principles can be mastered only after extensive study, but
the following experiments will illustrate many of the funda-
mental facts.
SOLUBILITY OF SOLIDS.
Experiment 7,^. — Supplies : About 20 gm. of powdered copper sul-
phate, 6 gm. of powdered potassium chlorate, i gm. of calcium sulphate.
{a) Label three test tubes I., II., III. Measure 10 cc. of water into
each. To I. add i gm. of powdered copper sulphate, to II. add i gm. of
General Properties of Water, 79
powdered potassium chlorate, to III. add i gm. of calcium sulphate.
Shake each test tube, and then allow them to stand undisturbed for a
few minutes. Is there evidence of solubility in each case? Is there
evidence of a varying degree of solubility? If III. is doubtful, carefully
transfer a portion of the clear liquid to an evaporating dish by means
of a pipette, or glass tube, and evaporate to dryness. Is there now
conclusive evidence of solution? Draw a general conclusion from this
experiment. Save solutions I. and II. for {b).
Tabulate the results of {a) as follows, using the customary terms to
express the degree of solubility : —
Table of Solubility of Typical Solids.
Solute.
Solvent.
Results.
1 . Copper sulphate
2. Potassium chlorate
3. Calcium sulphate
10 cc. of Water at
Temperature of
Laboratory
I.
2.
3-
(J)) Heat I. and add gradually 4 more grams of powdered copper
sulphate. Does it all dissolve? Heat II. and add 4 more grams of
powdered potassium chlorate. Does it all, or most all, dissolve? What
general effect has increased heat on the solubility of solids ? What is
the difference between this general result and that in Exp. 32? Save
the solutions for {c).
(^) Heat I. and II. nearly to boiling, and as the temperature increases
add the respective solids. Do not boil the liquid away. If the volume
of liquid becomes appreciably less, add a little water. Is there a limit to
their solubility? Draw a general conclusion from these typical results.
Saturation and Supersaturation. — Some of the solutions
obtained in Exp. 34 are Saturated. That is, the given vol-
ume of the liquid (or weight, since i cc. = i gm.) will dis-
solve no more solid at that temperature, although some of
the solid remains undissolved. A saturated solution repre-
sents a state of stable equilibrium. There is no simple
general relation between the temperature and the amount
8o Experimental Chemistry.
of solid dissolved. Since the maximum amount of solute
usually depends upon the temperature, it is customary to
say that a solution is saturated at such and such a tempera-
ture. Furthermore, the weight of the solvent must be
given. Usually, solubility is stated as the number of
grams soluble in lOO gm. of water at a certain tempera-
ture. Thus the number of grams of potassium nitrate
which lOO gm. of water dissolves at —
0°C. is 13.3
20°C. is 31.2
50°C. is 85.0
ioo°C. is 246.0
In Exp. 34 an excess of undissolved solid was present in
all cases when the solution was saturated. If, however,
solutions can be brought to the point of saturation, either
by cooling or by evaporation, frequently no solid will
separate from the Hquid. That is, solutions can exist
which contain a larger amount of dissolved solid than is
required for saturation at a given temperature. These
solutions are called Supersaturated.
SUPERSATURATED SOLUTIONS.
Experiment 35. — {a) Fill a test tube nearly full of crystallized
sodium sulphate, and add not more than 5 cc. of water. Warm slowly.
As solution occurs, heat gradually to boiling. Add sodium sulphate
until no more will dissolve. Pour the solution into a warm, clean, dry
test tube and let it stand until cool. Then drop in a small crystal of
sodium sulphate and watch for any simple but definite change. What
happens? Is the excess of solid large?
{b) Repeat with sodium thiosulphate, if time permits.
A supersaturated solution seems to be a case of unstable
equilibrium.
General Properties of Water.
8i
THERMAL PHENOMENA OF SOLUTION.
Experiment ^Jb. — Supplies : Test tubes, thermometer, 5 gm. each
of fused calcium chloride, potassium nitrate, ammonium nitrate, dehy-
drated copper sulphate, i gm. each of sodium hydroxide and potassium
hydroxide, and a few cubic centimeters of concentrated sulphuric acid.
Measure 10 cc. of water into a test tube, take the temperature, add at
one time 5 gm. of fused calcium chloride. As it dissolves, stir with the
thermometer and observe the highest reading. Record in the proper
place in the table below. Repeat successively with separate portions of
water and the ammonium nitrate, potassium nitrate, -dehydrated copper
sulphate, potassium hydroxide, sodium hydroxide, and concentrated
sulphuric acid (add the acid to the water). Tabulate the results as
follows : —
Summary of Thermal Phenomena of Solution.
Substances.
Solvent.
Temperature.
Thermal
Change.
Name.
Amt.
Name.
Amt.
Initial.
Maximum
or
Minimum.
Maxi-
mum.
Kind.
1. Calcium Chloride
2. Ammonium Nitrate
3. Potassium Nitrate
4. Copper Sulphate (dehy.)
5. Potassium Hydroxide
6. Sodium Hydroxide
7. Sulphuric Acid
In the act of dissolving, heat is necessary, hence solution
is usually accompanied by a fall of temperature. If the
dissolving body, however, coml3ines with some of the Uquid,
heat is produced, as is the case in chemical combination.
Sometimes the amount of heat produced is so large that it
more than balances the slight fall of temperature. Thus
when dehydrated copper sulphate is added to water, the
dehydrated salt combines with water to form the blue com-
pound, which is known to be copper sulphate combined with
82 Experimental Chemistry.
its water of crystallization. No doubt there was a slight fall
of temperature due to the mere act of solution, but the heat
due to the chemical combination is so much in excess that
heat alone is observed as the result of solution. It has
been proved that when sulphuric acid is dissolved in
water, compounds of water and sulphuric acid — hydrates of
sulphuric acid — are actually formed. So great is the heat
produced with large quantities that the solution often boils,
and sometimes so suddenly that the hot acid is spattered.
Hence, the acid should always be poured slowly into water
with constant stirring.
SOLUTION AND CHEMICAL ACTION.
Experiment -^. — Supplies: A gram each of powdered tartaric
acid, sodium bicarbonate, lead nitrate, potassium dichromate ; mortar
or evaporating dish, large vessel of water.
{a) Mix in a dry mortar or evaporating dish i gm. of powdered tar-
taric acid and an equal weight of sodium bicarbonate. Is tliere any
decided evidence of chemical action ? Pour the mixture into a large
vessel of water. Is there conclusive evidence of chemical action ?
{b^ Repeat, using powdered lead nitrate and powdered potassium
dichromate.
Describe the results in {a) and {b^ and interpret the whole experi-
ment from the standpoint of solution.
CLASS-ROOM EXERCISE. XIL
1. Solvent powder of water.
{a) Cleansing agent. {d) Solvent of chemicals.
{b) Erosive agent. {e) Composition of sea water.
{c) Solvent of drugs and medicines.
2. Relation of taste and solution.
3. Why are so many solutions used in the laboratory ?
4. A liter of sea water was evaporated to dryness, and the residue
weighed 36.4 gm. What per cent of the sea water was " salt " ?
5. How is a cold saturated solution of sodium chloride influenced by
rise of temperature ?
CHAPTER VI.
COMPOSITION OF WATER.
Water was thought to be an elenient until about the end
of the eighteenth century. At that time its composition
was shown by the efforts of Cavendish, Priestley, and Lavoi^
sier, and, later, by Humboldt, Gay-Lussac, Nicholson and
Carlisle, Dumas, and Davy. Their labors were so inter-
mingled it will be impossible to follow their experiments
chronologically. We shall, therefore, consider only the main
facts which contributed to the discovery of the composition
of water, together with some later work.
DECOMPOSITION OF WATER BY ELECTRICITY.
Water was first decomposed by electricity in 1800 by
Nicholson and Carlisle, and confirmed by Davy by a series
of brilliant experiments extending through a period of six
years. The decomposition of water by electricity is called
Electrolysis.
ELECTROLYSIS OF WATER.
Experiment 38. — Fill the Hofmann apparatus, Fig. 41, with water
containing 10 per cent of sulphuric acid so that the water in the
reservoir tube stands a short distance above the gas tubes after the
stop-cock in each has been closed. Connect the platinum terminal wires
with a battery. As the action proceeds, small bubbles of gas rise and
collect at the top of each tube. Allow the current to operate until the
height of the gas in the smaller tube is from 8 to 10 cm. Measure the
63
84
Experimental Chemistry,
height of each gas column. Assuming that the tubes have the same
diameter, the vokunes are in approximately the same ratio as their
heights. How do the volumes compare ?
Test the gases as follows : (a) Hold a glowir>g
taper over the tube containing the smaller quan-
tity of gas, cautiously open the stop-cock to allow
the water (or air) to run out of the glass tip, and
JL JL II then let out a little gas upon the glowing taper.
^ '^ ™ What is the gas ? Repeat until the gas is ex-
hausted. Care must be taken not to lose the gas
by clumsy manipulation. It is advisable to have
at hand several partially burned tapers or thin
splints, in case the escaping water extinguishes
the first one. (^) Open the other stop-cock long
enough to force out the water in the glass tip;
close the stop-cock, and slip a platinum tip (see
I if V i stop-cock again, let out the gas slowly, and hold
^^^^^fcg^ at the same time a lighted match at the end of
^«*i^^^^^ the tip, then immediately thrust a taper into the
small and almost colorless flame. What is the
gas ? Repeat until the gas is exhausted.
Draw a general conclusion from this experiment. Describe the
whole experiment and sketch the apparatus.
Fig. 41. — Hofmann
apparatus.
ELECTROLYSIS OF WATER. — Optional Method.
Experiment 39. — The Hofmann apparatus may be replaced by the
one shown in Fig. 42. A large bottle. A, is cut into two parts by the
method described in App. A, § 6. The upper part is used in this
experiment and it is about 15 cm. deep. A bottle with a small mouth is
preferable, though the one shown in the figure gives satisfactory results.
Provide A with a stopper or cork having two holes. Through these
holes pass two glass tubes into the upper ends of which are sealed
platinum wires. 'The upper end of each wire is attached to a narrow
strip of platinum foil by passing the wire in and out through several holes
pricked by a pin along the longer diameter of the strip. The lower ends
of these wires are attached to the battery by double connectors. When
the stopper is fitted with its tubes, press it into the neck of the bottle so
Composition of Water.
8s
firmly that there is no crack between the neck and stopper. This is
easily accomplished, if the neck is small and the stopper is rubber.
Corks are difficult to fit, unless both cork
and neck are perfectly circular. Place the
bottle in the ring of an iron stand and
pour ill warm water to heat the glass.
Meanwhile melt some paraffine and when
the bottle is warm enough not to be
cracked by the hot paraffine, pour out
the water and pour in the paraffine slowly
until it about fills the neck of the bottle.
As it cools, press it around the tubes
and where the glass and paraffine meet.
When cold, pour water into the bottle
to detect any leak. If the apparatus is
not tight, add more paraffine. It. may be
necessary to adjust the position of the
platinum strips when the test tubes, C
and D, are lowered into the bottle. The
electrolysis is accomplished the same as
in Exp. 38. A IS nearly filled with water
containing 10 per cent of sulphuric acid,
C and Z^ are filled with the same solu-
tion and clamped over the platinum strips
as shown in Fig. 42. When one tube is
full of gas, the current is stopped, the
volume measured as in Exp. 38, and each
gas is tested with a lighted taper or blaz-
ing stick of wood.
What is the ratio of the volumes ?
What are the gases ?
Fig. 42. — Apparatus for the elec-
trolysis of water. — Optiotial
Method.
It has been shown by many accurate trials of Experi-
ment 38 (i) that only two gases are produced, viz.,
hydrogen and oxygen, (2) that the ratio of their vol-
umes is 2 to I, (3) that the sum of the weights of the
products equals the weight of the water decomposed. It
therefore follows that water is composed of hydrogen and
86 Experimental Chemistry.
oxygen combined in the proportion of 2 to i by volume.
Additional experiments are necessary to prove the exact
composition of water.
Hydrogen and Water.— The fact that hydrogen is a
component of water, as shown in Exp. 38, may be verified
by an experiment first performed by Lavoisier. He passed
steam through a red hot gun barrel containing bits of iron.
The solid product was, he said, '' in the state of the black
oxide precisely Uke that which had been [formed from iron
when] burnt in oxygen." This experiment is historically
interesting because it led to the name " hydrogen."
Lavoisier says in his notes, "No name has appeared to
us more suitable than that of hydrogen, that is to say,
* generative principle of water ' — from hudor water, and
gehiomai^ I produce."
DECOMPOSITION OF WATER BY IRON.
Experiment 40. — Method: Pass steam over heated iron filings
and collect the gaseous product over water.
Apparatus: An iron tube about 30 cm. long and 2 cm. in diameter,
iron filings, steam generator, glass trap, pneumatic trough, three bottles
(250 CO.), and stoppers and connectors.
Process: Fill the iron tube about half full of clean iron filings, hold
it horizontal, and tap it gently on the table to provide a free channel
for the gases. Support the tube in a furnace, or lay it on the
rings of two iron stands. Provide each end with a one-hole rubber
stopper. To one end attach a delivery tube which passes into a
pneumatic trough. To the other end attach the steam generator. This
is a 500 cc. flask half full of water and provided with a two-hole rubber
stopper; through one hole passes a straight glass tube about 40 cm.
long, open at both ends, and reaching to the bottom of the flask — to
relieve any excessive back pressure ; through the other hole passes a
short glass tube bent at right angles and connected by rubber tubing
with a glass trap,i which in turn is connected by rubber tubing with a
1 See Teachers' Suppleinent.
Composition of Water. 87
short glass tube projecting from the stopper of the iron tube. The
trap collects any water driven over from the flask. The distance be-
tween the flask and iron tube must be as short as possible to prevent
condensation of the steam before it reaches the iron tube. It is advis-
able to put in the flask several pieces of granulated zinc, pipe stem, or
glass tubing, to insure steady boiling.
Fill the bottles with w^ater and invert in the trough, but not over the
hole in the shelf, since the first bubbles will be air. Heat the whole
iron tube by means of one or more wing-top burners. Meanwhile
heat the water in the steam generator, and by the time the steam has
been formed, the iron tube will be sufficiently hot. If drafts cool the
tube, protect it at the ends with asbestos boards. It is advisable, also,
to keep the ends of the tube cool by hanging a piece of asbestos board
from each end just inside the stoppers. Excessive heat may melt the
stoppers. If they should melt and stop up the tube, the pressure gauge
in the steam generator will indicate the difficulty. Reject the first por-
tions of the gas. As soon as the hydrogen begins to be delivered, a
few minutes will suffice to collect two bottles nearly full. Test the gas
for hydrogen and for oxygen, applying as many different tests as the
circumstances permit.
When the iron tube is cool, examine the contents. Compare it with
the original iron and with magnetic oxide of iron.
Describe this experiment, stating exactly what it proves.
Other metals, zinc and magnesium, for example, will also
decompose water, and hence furnish additional evidence
of the fact that hydrogen is one component of water.
CLASS-ROOM EXERCISE. XIII.
1. State exactly what Exp. 40 proves and suggests alDOut the com-
position of water.
2. What does Exp. 23 prove and suggest about the composition of
water ?
3. If coal instead of iron had been used in Exp. 40, what industry
would be illustrated ? (Hint : look up "water gas.'')
4. Summarize the evidence thus far presented about the composition
of water.
5. Actual facts contributed by Davy to our knowledge of the com-
position of water.
88 Experimental Chemistry.
Oxygen and Water. — The fact that oxygen is a constitu-
ent of water, shown in Exp. 38 and suggested in Exp. 23,
may be verified by allowing chlorine and water to interact
in the sunlight. Chlorine is a gas and an element, and
later will be fully studied.
DECOMPOSITION OF WATER BY CHLORINE.
Experiment 41. — Construct a chlorine generator and prepare
chlorine as directed in Exp. 57. Pass the gas into a deep vessel of
water until a sample of the liquid smells strongly of chlorine. The
delivery tube should reach to the bottom of the vessel of water. Com-
pletely fill a flask with a slender neck with this saturated solution of
chlorine., cork tightly, and stand in the sunlight. After several hours a
small quantity of gas will collect at the top. Test the gas with a glow-
ing match. Repeat the experiment, if the result is not satisfactory.
As the amount of gas is usually small, the apparatus shown in Fig.
43 gives more satisfactory results than the flask. It is a glass tube
about 2 cm. in diameter and i m. long,
closed at one end, either by sealing or
] i\ by inserting a cork or solid stopper, and
drawn out at the other end so that the
Fig. 43. — Tube for the decompo- smaller portion is about 4 cm. long and
sition of water by chlorine. . ,. ^ r^, . -, : nu a
^ I cm. m diameter. The tube is filled
with chlorine water, and the smaller end
immersed in a vessel containing the same solution. When sufficient
gas has been collected, cover the smaller end of the tube with the
thumb, invert the tube, and test the gas, which will nearly fill the
smaller portion of the tube.
Describe this experiment, stating what it shows about the composi-
tion of water. Sketch the apparatus used.
Interaction of Sodium and Water. — Additional tacts con-
cerning the composition of water may be obtained by a
study of the interaction of sodium and water.
Sodium is an element. It is also a metal, though it is
so soft it can be cut with a knife, and is much lighter in
Composition of Water. 89
weight than those metals which are familiar. Nevertheless
it has the luster characteristic of the metals, and chemically
Its behavior is markedly like the other metals. Sodium is
kept beneath oil to protect it from the moisture of the
atmosphere. // should be tised cautiously and strictly
iccording to directions. Small fragments sJiould not be
left about or thrown into the refuse jar, but into a large
vessel of zvater especially provided for the purpose,
INTERACTION OF SODIUM AND WATER.
Remember tJie precautions to be observed in iising .^odiuin.
Experiment 42. — -Supplies: Sodium, pneumatic trough filled with
water as usual, tea lead, test tube, forceps, litmus paper, iron stand and
clamp.
{a) Scrape the brown coating from a piece of sodium, press it
between filter paper to remove the excess of oil, cut off a piece not
larger than a small pea, and drop it upon the water in the trough.
Stand far enough away so that you can just see the action. Wait until
you are sure the action has stopped, and then describe all you have
seen.
{b) The action in {a) may be further studied as follows : Fill a test
tube with water, invert it and clamp it in the trough so that the mouth
is over the hole in the shelf of the trough. Wrap a small piece of
sodium loosely in a piece of tea lead about 5 cm. square, make two or
three small holes in the tea lead, and then thrust it under the shelf of
the trough with the forceps. A gas will rise into the test tube. Proceed
similarly with additional small pieces of sodium and dry tea lead until
the test tube is nearly full of gas ; then unclamp and remove, still keep-
ing the tube inverted. Hold a ligKted match, for an instant, at the
mouth cf the tube. Observe the result, watching especially the mouth
of the tube. What is the gas ? Why ? Remembering that sodium is
an element, where must the gas have come from ? If there is any doubt
about the nature of the gas, collect more, and subject it to those tests
which will prove its nature.
(c) Put a piece of filter paper on the water in the trough, and before
it sinks drop a small piece of sodium upon it. Stand back and observe
^o Experimental Chemistry.
the result. Wait for the slight explosion which usually occurs soon
ifter the action stops. Describe all you have seen. What burned ?
What caused it to burn ? To what is the vivid color probably due ?
(In answering these questions, utilize your knowledge (i) of the prop-
erties of the gases previously studied, and (2) of the usual accompani-
ment of chemical action, suggested here by the melting of the sodium.)
{d) Test the water in the trough with red litmus paper. Push the
paper to the bottom or to the place where it is certain that chemical
action between water and sodium has taken place. Test, until the red
litmus paper has undergone a decided change in color. Describe this
final result. With another piece of red litmus paper test a solution
made by dissolving a small piece of sodium hydroxide in a test tube
half full of water. Is the result similar ? Dip the platinum test wire
(see App. A, § 14) into this solution and hold it in the Bunsen flame.
Describe the result. Is the color of this flame and that noticed in {c)
the same ?
Experiment 42 verifies the previous observation regard-
ing the presence of hydrogen in water. It shows also
that when water and sodium interact another substance is
formed. This is sodium hydroxide. Its name suggests its
components, for it is a compound of sodium, hydrogen,
and oxygen. It belongs to ■ a class of substances called
alkaUes, which will be considered later. These alkalies turn
red Htmus paper blue; and sodium hydroxide, like all
sodium compounds, colors a Bunsen flame an intense yellow.
The color is due to the volatiUzed sodium, and is the test
not for hydroxides, but for sodium compounds. The sodium
hydroxide is in solution in the trough, and by evaporating
the water could be obtained, as a white solid, having all the
properties of the sodium hydroxide actually used to make
the solution in(<^). Since sodium liberates hydrogen from
water, and forms at the same time a compound also contain-
ing hydrogen, then the hydrogen in water must be divisible
into two parts. If in Exp. 42 \b) the sodium had been
weighed and its weight compared with the volume of
Composition of Water. 91
hydrogen liberated, it would have been found that .1 gm.
of sodium liberates 48.22 cc. of hydrogen. And if the
sodium hydroxide thus formed had been dried and then
heated with sodium, 48.22 cc. more of hydrogen would
have been obtained. This shows that the hydrogen in
water is divisible into two equal parts.
The facts revealed by the interaction of sodium and
water, together with others which have been proved, permit
us to express the reaction in its simplest form by the
following equation : —
Na + H.,0 = H + NaOH
Sodium Water Hydrogen Sodium Hydroxide (7)
23 18 I 40
This equation means that 23 gm. of sodium are needed to
liberate i gm. of hydrogen from 18 gm. of water, and to
form, at the same time, 40 gm. of sodium hydroxide.
The Quantitative Composition of Water. — The preceding
experiments on the composition of water have been mainly
qualitative. They have shown by analysis and synthesis
that water is composed of hydrogen and oxygen, and that
the ratio of their volumes is approximately 2 to i. The
most decisive evidence of the quantitative composition
of water is obtained by the determination of (i) its exact
volumetric composition, (2) its exact gravimetric composi-
tion, and (3) the density of steam. Volumetric means
"by volume," and gravimetric ^eans "by weight."
VOLUMETRIC COMPOSITION OF WATER.
Experiment 43. — Method: Explode measured volumes of hydrogen
and oxygen and measure the residual gas, using an excess of either gas.
Apparatus : The apparatus is shown in Fig. 44. ^ is a level tube
about 90 cm. long, provided with a base, or fitted at one end with a
one-hole stopper, to which is connected a piece of pressure tubing, I^
92
Experimental Chemistry.
about 50 cm. long. When not in use, A is clamped in any convenient
position. F IS 2. eudiometer so arranged that the open end can be
moved about in the reservoir, E. This reservoir is about 15 cm. deep
and is made from a five-pint acid (or similar) bottle according to the
directions given in App. A, § 6. It is supported by a ring or clamp,
and is provided with a one-hole stopper, through which passes a short
^\.
Fig. 44. — Apparatus for determining the volumetric composition of water.
glass tube, G, the ends of which have a flange, made by melting, each
end in the Bunsen flame and pressing it quickly upon a hard surface.
Before the flange is made, a rubber stopper, C, just large enough to
fit the eudiometer, is pushed down firmly upon the other stopper.
After the flange is made, the upper end of G is pulled down hard upon
C, the tubing, B, is pulled over the lower end of G until it touches
Composition of Water. 93
the large stopper, and tied securely near the stopper and near the
lower flange by shoemaker's waxed thread. A Hofmann screw is
attached near the point D.
Process: Loosen the Hofmann screw, pour into A water that has
been standing in the laboratory for at least twenty-four hours, until the
reservoir, E, is two-thirds full ; then tighten the screw and add more
water until the level tube, A, is three-fourths full ; clamp A at any con-
venient height to the iron stand which supports the apparatus. Pinch
B to remove any imprisoned air. Fill the eudiometer, F, with some of
the same water, invert in the reservoir, and clamp it so that the open
end is just below the surface and near the wall of E.
Introduce into /^ about 10 cc. of oxygen free from air. The oxygen
may be taken from a gas holder, or it may be prepared from the usual
mixture in a test tube, since only a small volume is needed. Unclamp
/^ and press it down upon the stopper, C\ clamp F again, loosen the
Hofmann screw, stand a thermometer in E^ and allow the whole appara-
tus to remain undisturbed for at least ten minutes. Meanwhile, read
the barometer, and write all necessary notes. Hold the top of the
level tube near the eudiometer and raise or lower the level tube, until
the water is at the same height in both A and F, then read the volume
of oxygen in the eudiometer and the temperature of the water in E.
Clamp A again in any convenient position. Record all readings as
illustrated below.
Unclamp the eudiometer and lift it from C, and introduce into it
from 25 to 30 cc. of hydrogen free from air. The hydrogen may be
prepared from zinc and sulphuric acid in a test tube. Proceed as in the
reading of the oxygen, observing the same precautions, especially the
interval before reading the gas volume. This time, however, the eudi-
ometer must be pressed down hard upon the stopper, C, taking care, of
course, not to force the larger stopper out of place. Record the total
volume and other readings as illustrated below.
Make sure that all stoppers and jaints are tight, lower A to expand
the gases in the eudiometer, clamp the Hofmann screw as tight as pos-
sible just below the glass tube, G, and clamp A as before. Explode the
mixture of hydrogen and oxygen by passing an electric spark through
the eudiometer. The spark may be obtained from a RuhmkorfF coil or
plate machine. The explosion is indicated by a slight click, often
accompanied by an instantaneous flash passing down the tube from the
platinum wires to which the coil is connected. The ends of the plati-
94 Experimental Chemistry.
num wires must be about 5 mm. apart in order to insure an explosion.
After the explosion, loosen the screw and allow the water to seek its
own level in the two tubes. Let the whole apparatus remain undis-
turbed for ten minutes or more, then read as before the gas volume,
thermometer, and barometer.
A minute quantity of water is formed when the two gases unite, but
so minute that its volume is not measurable. The residual gas can be
proved by the usual tests to be hydrogen. The data obtained enable
us to compute at once the volumetric composition of water.
Calculation : The recording of the data and calculation of the final
result may be illustrated by an experiment actually performed by a ^
student with this apparatus. Other determinations will vary somewhat
from these figures, though the final result may be the same.
Data : —
I. Oxygen. II. Hydrogen and oxygen.
V = 10.7 cc. y - 51.5 cc.
/ =i7"C. / =i7°C.
P' = 767 mm. P' =■ 767 mm.
a = 14.42 mm. a = 14.42 mm.
III. Residue after the explosion.
y = 20.4 cc.
/ =i7°C.
P' = 767 mm.
a = 14.42 mm.
These volumes, corrected (see App. B, formula 5) for temperature,
pressure, and aqueous tension, become —
I. Oxygen added 9.97 cc.
II. Hydrogen and oxygen added 48.01 cc.
III. Hydrogen left 19.02 cc.
Since 19.02 cc. of hydrogen remain, then the volume of hydrogen
and oxvsren which combined must be
48.01 - 19.02 = 28.99 cc.
and volume of hydrogen which united with th^xygen must be
28.99 ~ 9-97 = ^9-°2 cc.
Composition of Water.
95
If 19.02 cc. of hydrogen united with 9.97 cc. of oxygen to form
water, then the ratio in which they combined was
19.02 : 9.97 or 1.9 : i.
Accurate repetitions of this experiment have shown that
the ratio in which the pure gases, hydrogen and oxygen,
combine is 2 to i, that is, the vohuiietric composition of
water is two parts hydrogen and one part oxygen. The
same result was obtained in Exp. 38, to be sure, but the
result of Exp. 43, ^^btained by a more accurate method,
establishes the fact beyond dispute.
GRAVIMETRIC COMPOSITION OF WATER.
Experiment 44. — Method: Pass dry hydrogen over a weighed
amount of copper oxide and collect the water formed in a weighed tube.
Apparatiis : The construction and arrangement of the apparatus,
except the generator, is shown in Fig. 45. The hydrogen may be
generated as in Exp. 19, but it is difficult with this apparatus to obtain
a slow evolution of gas as well as to be sure the generator will not be
exhausted before the experiment is completed. If a gas holder or a
in,
r-
A
^f^
vcy
Fig. 45. — Apparatus for determining the gravimetric composition of water.
reliable Kipp's apparatus is available, either one may be used. A
simple, compact, and reliable generator is shown in Fig. 46. Two tubes,
A and B^ about 30 cm. long, 2 cm. in diameter, and open at both ends,
are connected, as shown in the figure, by a piece of rubber tubing, 6",
about 40 cm. long. A Hofmann screw is attached at E. The generator
is connected at D with the rest of the apparatus. The generator is
operated as follows : Fill B two-thirds full of granulated zinc, separat-
#
96
Experimental Chemistry,
ing it from the bent tube at the bottom by a plug of cotton. Tighten
the screw and fill A with dilute sulphuric acid. When the screw is
loosened, acid will flow into B, and hydrogen will be evolved. It is
advisable to clamp B into a permanent position,
/>- — ' open the passage a little, and regulate the evo-
r 1 1. II. lution of hydrogen by raising or lowering A by
another clamp. A height will be found by trial
which will produce the desired rate of evolution.
In Fig. 45 ^ is a Drechsel drying bottle
which should be half full of concentrated sul-
phuric acid, and B is aK^-tube to be filled with
calcium chloride — both serving to dry the hy-
drogen. CC is a hard glass tube from 15 cm.
to 20 cm. long, and from 1.5 cm. to 2 cm. in
diameter. It is attached directly by rubber stop-
pers to B and Z), the latter being a small (10 cm.)
iMarchand tube to be filled with fused calcium
chloride. E is a U-tube to be filled with calcium
chloride to prevent the moisture in the air from
entering D. Only the tubes CC and V are
weighed before and after the experiment. At
the points C and C a copper wire (No. 20) is
wound once or twice loosely around the com-
bustion tube, and allowed to project at several
points, like the spokes of a wheel, to prevent the
tube from cracking and the stopper from melting
at these points. This wire is fto^ weighed with
FIG. 46. -A generator ^j^^ ^^^^^ ^jj ■ -^^ ^^^t ^^ ^j l^t^ The appa-
for producing a steady , •' 1,1 ^ ^1
current of hydrogen ratus may be supported by clamps at the neces-
(and other gases to be sary points, but not, of course, where they
subsequently studied) . ^yjn interfere with the application of heat to
CC. Clean and dry the tubes B and E, and
provide them with one-hole rubber stoppers, the holes of which are
temporarily plugged with a short glass rod ; fill each three-fourths full
with lumps (not powder) of fused calcium chloride, push a plug of
cotton down loosely upon the calcium chloride, wipe oflf any small
particles of calcium chloride, and close the tubes with the plugged
stoppers — removing the plugs, of course, when the tubes are in use.
Clean and dry D and prepare it as follows : Push enough absorbent
Composition of Water. 97
cotton down the straight limb and up the other limb to fill loosely the
lower bulb, and then fill the tube with fused calcium chloride ; nothing
is put in the upper bulb ; the straight limb is provided with a one-hole
rubber stopper, which has previously been cleaned by boiling it succes-
sively in dilute sodium hydroxide, hydrochloric acid, and water. This
stopper should have a short glass connector, bent at a right angle,
and each end of the tube
should be protected from { -Z^^^^--- -ll' ^
the air by a short rubber
, , 1 . , , Fig. 47. — Glass plug.
tube plugged with a glass
rod (see Fig. 47). The water which collects in the empty bulb during
the experiment may be removed by a bit of rolled filter paper, and the
tube thereby used several times without refilling. The stoppers fitted
to CC should be cleaned, before use, by boiling as above described.
Process: Clean and dry the combustion tube CC and put in it
about 20 gm. of granulated copper oxide which has been heated in an
iron or porcelain dish to remove moisture and organic matter. Intro-
duce the copper oxide by the method illustrated in App. A, § ii, and
hold it in place by two loose plugs of shredded asbestos which has been
heated red-hot by the blast lamp. It may be heated in a porcelain
dish, if the dish is heated and cooled gradually. Weigh the combustion
tube and contents. Weigh the Marchand tube, D, without the pro-
tectors. Record the weights as shown below. Connect the apparatus
as shown in Fig. 45, joining CC directly to B and Z>, after the copper
wires have been slipped over the ends of the combustion tube. Attach
the generator, and pass a slow current of gas through the whole appara-
tus to drive out the air and incidentally to detect any leak. If there is
no leak, then the gas will not bubble through A, when the open end of
E is covered for an instant with the finger. If there is a leak, stop it
before proceeding. Regulate the flow of hydrogen so that about three
bubbles a second pass through A. An ordinary Bunsen burner and a
wing-top burner are needed, the latT;er for constant use, the former for
an emergency. Heat the whole combustion tube slowly at first with
the wing-top burner, gradually increasing the heat where the copper
oxide is located. As the heat increases, moisture will collect in CC .
Heat the combustion tube from C toward C so that the moisture will
not collect at C, but be driven along beyond the copper oxide toward
C and over into the bulb oi D. The point C must be kept hot. just
as hot as possible without melting the stopper ; if the moisture collects
^8 Experimental Chemistry.
here, it is difficult to remove. The moisture often collects in the
oxide, hence this must be kept hot enough from the outset to volatilize
the water. If drops gather at C, heat this point cautiously with a low
Bunsen flame. The wing-top burner may be held at various angles
to CC, thereby allowing heat to be easily directed upon short spaces.
When the copper oxide ceases to glow, lessen the heat somewhat and
drive over into D any moisture seen, or thought to be, in the combustion
tube. Allow the hydrogen to run while the apparatus is cooling.
When CC is cool enough to handle, disconnect the generator and draw
air through the apparatus by applying the lips at the open end of E.
Disconnect CC, wipe out any organic matter from the inside at the-
ends, and then weigh. Disconnect D and weigh as before. Record
the weights as shown below.
Calculation: The form of recording the data and calculating the
final result may be illustrated by figures taken directly from a studenfs
note book. Other determinations of the gravimetric composition of
water may, of course, give entirely different figures, though the final
results may agree.
Data : —
I. Combustion tube.
Grams.
Weight of combustion tul)e and copper oxide before heating . 65.688
Weight of combustion tube and copper oxide after heating . 64.783
Weight of oxygen lost - . . 0.905
II. Marchand tube.
Weight of Marchand tube after experiment 76-741
Weight of Marchand tube before experiment .... 75-720
Weight of water formed 1021
But 0.905 and 1.021 are in the same ratio as 8 and 9 (approximately).
Thus 0.905 : 1.021 ;: I ;;r
x—\ J28
But I : 1.128 : :8 ;9.024
That is, oxygen makes up | of water, and the remaining \ is hydrogen.
Composition of Water. 99
Final dh-ections : Protect from the air the whole drying apparatus
by connecting it "tandem"' and plugging the ends with the rod or
rubber plug, as the case demands. Protect B with its rubber plugs.
Pour the acid from the generator, wash the zinc and rubber tube free
from acid, and recharge, or preserve empty, for future use.
Numerous chemists, whose work is reliable, have found
this proportion (|) to be approximately correct.
Density of Steam. — The density of steam is the ratio of
the weight of a given volume of steam to an equal volume
of hydrogen under the same conditions. It could easily be
calculated if we knew the weight of a liter of steam under
given conditions of pressure and temperature, because
the weight could then be compared directly with the weight
of a liter of hydrogen. It has been found that .8063 gm.
is the weight which a liter of steam would have, if it could
exist as steam at the normal temperature and pressure.
The determination is carried on, of course, at a tempera-
ture above 100° C, and the volume reduced to standard
conditions, which are simply arbitrary points, selected
irrespective of the physical states of matter. Dividing
the weight of a liter of steam by the weight of a Hter of
hydrogen, we have —
:8o63^
.0896 ^
The significance of this fact will appear later, when the
subject of molecular weights is studied. Just now all that
can be said is that the molecular weight of a gas is twice
its density. Hence 18 is the molecular weight of steam.
Summary. — The following facts have been revealed,
partly by our experiments, and partly by suggestion, in
the study of the composition of water : —
lOO Experimental Chemistry.
1. Water is a chemical compound of hydrogen and
oxygen.
2. It is formed by burning hydrogen in the air, or by
exploding a mixture of hydrogen and oxygen.
3. It can be decomposed by electricity into hydrogen
and oxygen in the proportion of two volumes of the
former to one volume of the latter.
4. Sodium liberates hydrogen from water, and forms at
the same time a solid containing hydrogen equivalent to
the hydrogen liberated. Iron also liberates hydrogen from
water.
5. Chlorine liberates oxygen from water.
6. Two volumes of hydrogen and one volume of oxygen
combine to form water, and the weight of the water
formed equals the weight of the gases.
7. Water is formed by the union of two parts by weight
of hydrogen with sixteen parts by weight of oxygen.
8. The molecular weight of water is 18.
The foregoing facts permit us to write the equation
(8)
as the simplest expression of the reaction for the forma-
tion of water from hydrogen and oxygen.
The significance of the density of oxygen (see page 50)
is now apparent, at least in part, since 16 is the smallest
number which can be used in the above equation. The
meaning of the result of the determination of the gravi-
metric composition of water (see page 98) is likewise
clearer. All facts point to the truth of equation (8).
H^
+ 0 = H2O
Hydrogen
+ Oxygen = Water
2
-1- 16 = 18
Composition of Water. loi
CLASS-ROOM EXERCISE. XIV.
I, Actual facts contributed to our knowledge of the composition of
water by —
{i^^ Cavendisho
(^) Dumas.
(^) Gay-Lussac.
2 The formula of water.
(a) Why HP?
(/^) Why riot HO, as formerly?
{c) Why not U.,OJ
3. Which of the following are oxidation and which reduction : —
{a) Burning of hydrogen.
(<^) Formation of iron oxide from steam and iron.
{c) Formation of magnesium oxide by heating magnesium
in air.
4. Symbol of Sodium, Iron, Copper, Hydrogen, Oxygen, Chlorine,
Platinum.
5. Formula of water, steam, ice, sodium hydroxide, copper oxide,
magnetic oxide of iron.
6. Define gravimetric, volumetric, eudiometer, ratio, density.
7. Sources of error in Exps 43 and 44, and how avoided.
8. Litmus.
(a) Source.
(d) Preparation of commercial form.
(c) Uses.
Volumetric Composition of Steam. — If water is decom-
posed, and the mixed gases ""are exploded when the sur-
rounding temperature is above 100° C, then the product is
gaseous water, that is, steam. If the experiment is per-
formed quantitatively, it is found that three volumes of the
mixed gases form, two volumes of steam. The impor-
tance of this fact will appear later, when similar facts
regarding the two gases, hydrochloric acid and ammonia,
2H2 +
O2 = 2 H^O.
2 Volumes of
Hydrogen
I Volume of _ 2 Volumes of
Oxygen ~ Steam
4 -H
32 = 36
102 Experimenta! Chemistry.
are revealed. The simple relation existing between the
volumes of the components and resulting gas is expressed
by the equation —
(9)
This equation means that 4 gm. of hydrogen require 32
gm. of oxygen to form 36 gm. of steam.
Reduction. — The chemical change in the combustion
tube in Exp. 44 is called Reduction. It usually means the
removal of oxygen, and is, therefore, the opposite opera-
tion from oxidation. In Exp. 44, most, perhaps all, of the
copper oxide was reduced to metallic copper, that is,
oxygen was removed. Hydrogen is therefore a reducing
agent, just as oxygen is an oxidizing agent. Reduction is
an important chemical operation, especially industrially,
since many ores are reduced from the oxide to the metal.
PROBLEMS. XI.
1. What weight of hydrogen and of oxygen can be obtained from
125 gm. of water?
2. If 10 gm. of water are decomposed by electricity, what weight
and what volume of hydrogen and of oxygen are formed?
3. If 216 gm. of water are decomposed by electricity, how many
liters of oxygen are formed?
4. How many grams of water are formed by burning 5 1. of hydrogen ?
5. 10 1. of hydrogen are to be burned in oxygen.
(a) What volume of oxygen is necessary?
{b) What weight of water will be formed?
6. What volume of oxygen is necessary to unite with 100 gm. of
hydrogen?
Composition of Water. 103
7. A piece of sodium weighing 150 gm. is allowed to interact with
water. Calculate —
(a) The weight and volume of hydrogen liberated.
(d) The weight of sodium hydroxide formed.
8. A gram of sodium amalgam (a mixture of sodium and mercury)
liberated from water 200 cc. of hydrogen at 13° C. "What per cent of
sodium did the amalgam contain?
9. How many centigrams of hydrogen will be liberated when 23 eg.
of sodium interact with water? How many centigrams, if the sodium
had contained 8 per cent of impurities?
10. How many grams of hydrogen can be obtained by using 3 gm.
of sodium ? How many cubic centimeters ?
11. Hydrogen is passed over 2.48 gm. of copper oxide, which at the
end of the experiment weighed only 2.24 gm. ; the water formed weighed
0.27 gm. In what ratio did the hydrogen and oxygen combine?
12. Calculate the ratio in which hydrogen and oxygen combine to
form water from the following data : —
Grams.
Weight of copper oxide tube before the experiment . 105.6
Weight of copper oxide tube after the experiment . . 10 1.6
Weight of drying tube after the experiment .... 84.8
Weight of drying tube before the experiment . . . 80.3
13. Calculate the ratio in which hydrogen and oxygen combine to
form water from the following data, which were obtained by Dumas : —
Grams.
Weight of copper oxide tube before the experiment . 334-598
Weight of copper oxide tube after the experiment . 314.236
Weight of drying apparatus after the experiment . . 449.263
Weight of drying apparatus before the experiment . 426.358
14. What weight of copper remains after 159 gm. of copper oxide
have been reduced by hydrogen ?
15. If 20 gm. of hydrogen are passed without loss through a tube
containing hot copper oxide, what weight of w^ater is formed ?
16. Berzelius and Dulong, in 1820, obtained the following data in
three determinations of the gravimetric composition of water : —
Loss of copper oxide tube in grams, respectively, 8.051, 10.832,
and 8.246.
I04 Experimental Chemistry.
Weight of water formed in grams, respectively, 9.032, 12.197, 9.27.
Calculate in each case the ratio of combination. What is the average
result ?
17. Dumas and Stas, in 1843, repeated the work of Berzelius and
Dulong, and as an average of nineteen determinations found that
840.161 gm. of oxygen formed 945-439 g^'^- of water. Calculate the
ratio of combination.
CHAPTER VII.
THE ATMOSPHERE.
The atmosphere is the gaseous envelope surrounding
the earth and extending into space. It is often called the
air. The word air was used by the earUer chemists in the
sense in which the word gas is now employed. The words
atmosphere and air are often used interchangeably. The
study of the atmosphere has occupied the time of many
skilful chemists, and the progress of science has been has-
tened by their labors, especially by their investigations of
the action of air on living things, its relation to combustion,
the influence of its weight, and the proportion and inter-
action of its constituents. Boyle (i 626-1 691) began these
investigations, and they are still in progress. A new
impetus has recently been given to the work by the dis-
covery of argon in the atmosphere by Rayleigh and
Ramsay. This gaseous element, which had hitherto es-
caped detection, forms about ^Jq part of the atmosphere.
Composition of Air. — It has already been shown that air
contains oxygen and at least one other gas. It is now a
favorable time to study this other constituent and the pro-
portions in which these two^ases exist in the air. These
two objects may be simultaneously accomplished by a
QUANTITATIVE EXAMINATION OF AIR.
Phosphorus is used in this experiment, before touching
the phosphorus, read carefully the precautions stated belozv,
and observe them.
105
io6
Experimental Chemistry,
Experiment 45. — Method: Remove the oxygen from a measured
volume of air by phosphorus, and measure the residual gas.
Apparatus : The arrangement of the apparatus is shown in Fig. 48.
^ is a tall jar, B is a 100 cc. graduated tube, C is a piece of worm-
shaped phosphorus attached to a long piece (40-60 cm.) of copper wire,
Z^ is a clamp to support B.
Process : Fill A two-thirds full of water which has been exposed to
the air in the laboratory for at least twenty-four hours. Fill the tube, B,
with the same water, cover
the end with the thumb or
finger, hold it over the jar,
and allow enough water to
run out so that when the
tube is lowered into the jar
the 100 cc. mark is a short
distance below the surface,
if the water is at the same
height within and without
the tube. Clamp the tube
in this position, stand a
thermometer in the water
in the jar, and allow the
whole to remain undisturbed
for about fifteen minutes.
While the air is coming to
the temperature of the water
over which it is enclosed,
prepare the phosphorus.
Phosphorus is an element
which oxidizes so quickly in the air that it often burns unexpectedly.
To prevent its oxidation it is kept under water. // /s dangerous to
handle, since the burns made by it are deep and painful. Transfer a
piece of phosphorus about 5 cm. long with the forceps from the bottle
to a large porcelain mortar, battery jar, or metal dish nearly full
of water. Wind the copper wire around it in the manner shown in
Fig. 48. Quickly, but deliberately, transfer the phosphorus to the jar
of water, A.
Without touching the tube with the hand (see App. B, II., (3) and
(4) ), adjust it so that the water is at the same level within and without
Fig.
Apparatus for the quantitative exam-
ination of air.
The Atmosphere. 107
the tube. Read the volume of air in the tube, the thermometer, and
barometer. Record as shown below.
Push the phosphorus up into the tube just above the surface of the
water. White fumes indicate immediate action. In about ten minutes
push the phosphorus half-way up the tube, attach the wire to the
jar as shown in the figure, and allow the whole to remain undis-
turbed for an hour or more. Then pull the phosphorus down into the
jar, still keeping it under water, and gently rub off with the fingers any
deposit on the surface of the phosphoms. Insert the cleaned phos-
phorus into the tube as before, and if the white fumes appear, let the
phosphorus remain a half-hour longer. If no fumes appear, then lower
the phosphorus into the jar and let it remain there, taking care to keep
the copper wire below the surface (to avoid accidental removal of the
phosphorus from the jar). Lower the tube, without touching it with the
hand, until the water is about the same level within and without, stand
the thermometer in the water, and let the whole remain undisturbed
for fifteen minutes. Then readjust the tube, if necessary, so that the
inner and outer levels are the same, read the volume of gas, the ther-
mometer, and barometer. Record as shown below. Unclamp the tube,
grasp it near the end with the hand, close the open end with the thumb
or finger, lift it from the water, invert, and test the residual gas several
times with a glowing stick. Observe carefully its action with a flame,
and record the observations. The name of this gas is nitrogen.
Calculations :
Data : —
I. Before the removal of oxygen.
Volume of air taken = V =
Temperature = / =
Pressure = P' =
Aqueous tension = a =
Reduce this volume to the volume it would occupy under standard
conditions by the formula
^ 760(1 -f- .00366 X t)
Corrected volume — V-^=
io8
Experimental Chemistry,
II. After the removal of oxygen.
Volume of gas remaining = F' =
Temperature = / =
Pressure — P' ■=.
Aqueous tension —a =
Reduce to standard conditions as in I.
Corrected volume = Kj =
From V^ and K, find the percentage of oxygen and nitrogen in air.
Insert the answer in the proper place in the
Table of the Composition of Air.
Constituent.
Per Cent by Volume.
Found.
Class Average.
Correct.
Oxygen
Nitrogen (and Argon) . .
20.96
79.04
LABORATORY EXERCISE. V.
1. Upon vv'hat property of phosphorus does the chemical action in
Exp. 45 depend ?
2. What became of the white fumes seen in Exp. 45?
3. Summarize the most obvious properties of nitrogen.
4. Compare the properties of nitrogen with those predicted in
Laboratory Exercise II.
CLASS-ROOM EXERCISE. XV.
I. Atmospheric nitrogen.
{a) Discovery and distribution.
{b) Physical functions in the atmosphere.
{c) Chemical relation to plants and animals
(^) Its direct assimilation by plants.
{e) How prepared in large quantities.
The Atmosphere.
109
2. Atmospheric oxygen.
(a) Physical and chemical functions in the atmosphere.
3. Experimental.
(a) Give several methods for showing that air contains oxygen.
(d) Suppose air is passed over red-hot copper, what are the
products? If the experiment were conducted quantita-
tively, what would it prove ?
4. Historical.
(a) Who first proved the gravimetric composition of the atmo-
sphere ? Describe the apparatus and state the method.
(d) Bunsen's work on the volumetric composition of the air.
(^) Lavoisier's work on the composition of the air.
(d) Aristotle's conception of the word ai'r.
5. Miscellaneous.
(a) Sources of error in Exp. 45.
(d) Precautions in handling phosphorus.
The atmosphere is not entirely composed of oxygen and
OTHER CONSTITUENTS OF THE ATMOSPHERE.
Experiment 46. — (a) Prove by an experiment
that water vapor is present in the atmosphere.
(d) Construct an apparatus like that shown in
Fig. 49. Fill the test tube half full of lime water
and draw air through the apparatus by means of an
aspirator bottle or filter pump. (See App. A, § 17.)
If no pump or aspirator is available, expose lime water
in a bottle or beaker to the air. Allow the operation
to continue, in either case, until a definite change is
produced in the appearance of the lime water. In the
aspirated lime water the change will appear through-
out the liquid ; in the exposed lime water look upon
the surface for the change. If neither operation is
satisfactory, use a concentrated solution of barium
hydroxide in place of lime water.
change is due to the formation of
The
calcium carbonate. Lime water is the pop-
Fig. 49. — Appara-
tus for detecting
the presence of
parbon dioxide
in the air.
Ca(OH),
+ CO2 -
CaCOg + H2O
Calcium
Carbon
Calcium Water
Hydroxide
Dioxide
Carbonate
1 10 Experimental Chemistry.
ular name of a solution of calcium hydroxide, and when
it is exposed to carbon dioxide, the two interact. The
chemical change is represented in its simplest form by
the equation —
(10)
The formation of this easily detected calcium carbonate
is the conclusive test for carbon dioxide. Usually a glass
rod which has been dipped into hme water is held in the
escaping gas, and if the lime water turns milky, this is a
proof of the presence of carbon dioxide. Barium hydroxide
acts the same as calcium hydroxide, and is sometimes used
instead of the latter.
CLASS-ROOM EXERCISE. XVI.
1. Atmospheric water.
(a) Source.
. (d) Relation to plants and animals.
(c) What are clouds? What is dew, rain, frost, snow, hail?
(d) Amount present in air on a pleasant day.
(e) What is meant by the "dew point" of air?
2. Atmospheric carbon dioxide.
(a) Source.
(d) Relation to plants and animals.
{c) Carbon dioxide is heavier than air. Why, then, is the
former gas not always at or near the earth's surface?
(^) Proportion in the air?
(e) How determined quantitatively?
3. Miscellaneous.
(a) Name the gases occasionally found in the atmosphere;
source of each.
(^) Name those rarely found.
The Atmosphere.
Ill
(c) "The air is a mixture/'' Give in your own language tliree
reasons for this statement,
(rt') If the nitrogen were entirely removed from a vessel of air.
would the physical properties of the oxygen be changed ?
(e) If air is a mixture, why do we speak of its '* composition"?
(/■) Chemical effect of lightning on the atmosphere.
WEIGHT OF A LITER OF AIR.
Experiment 47. — Method: Find the weight and volume of air ex-
hausted by a pump from a bottle.
Apparatus : The apparatus is shown in Fig. 50. A is a large bottle
(2 liters or more) provided with a one-hole rubber stopper, through
which passes a glass tube, B^ bent as shown in the figure. C is a
piece of pressure tubing (40-50 cm.
long). A Hofmann screw is at-
tached at the point D. It is advis-
able to make a flange on the outer
end of the glass tube by melting it
in the Bunsen flame and pressing
it quickly upon a hard surface.
The stopper must fit air tight, and
a little vaseline should be rubbed
on that part of the stopper which
enters the neck of the bottle. If
a filter pump is used to exhaust the
bottle, a straight calcium chloride
tube, filled as usual, should be
placed between the pump and the
bottle.
Process : Slip C over the end of
B^ and unless it fits exceedingly
tight, tie it just above the flange
with waxed thread. Attach the
screw loosely at Z>, push in the stopper as far as possible, wipe the
bottle with a piece of cheese cloth or soft paper, and weigh the whole
apparatus on the trip scales. Stand the bottle in the center of the
left-hand scale pan and arrange the weights symmetrically around the
center of the other. Weigh to a decigram, and leave the weights on
Fig. 50. — Apparatus for determining
the weight of a hter of air.
112 Experimental Chemistry.
the scales ready for the second weighing, so that the same large weights
may be used each time. This precaution will eliminate many of the
errors due to weighing. Call this weight «, and record as shown below.
Slip the end of the rubber tube over the exhaust nozzle of an air
pump, and exhaust the air by rapid, short strokes of the handle.
Tighten the Hofmann screw so that it will be air tight. The screw
should not be more than five centimeters from the end. Detach from
the pump, and weigh again — on the same scales. If less than i gm. of
air has been removed, exhaust again. Call this weight b, and record
as shown below.
Lower the rubber tube into a jar of water, open the clamp cautiously
under water ^ and allow the water to run in slowly. The sudden rush
of water into the bottle is easily controlled by pinching C near the end
between the thumb and forefinger. The jar can be kept full by allow-
ing water to flow continuously through a tube reaching to the bottom
of the jar. Care must be taken to keep the jar free from air bubbles by
adapting the inflow to the outflow. When no. more water runs into
the bottle, raise or lower the bottle until the water is at the same
level in both jar and bottle, pinch the tube and remove it from the jar.
Tighten the screw, loosen the stopper, wipe the whole apparatus dry,
and weigh again, observing the previous precautions. Record this
weight as c, as shown below.
Calculation :
Data : —
Grams.
a Weight of bottle before exhaustion
d Weight of bottle after exhaustion
c Weight of bottle and water
Therefore
Therefore
a — b = Weight of air removed = gm.
c — b = Volume of air removed = cc.
c — b: a — b\\ looo : x
x= gm.
The Atmosphere.
113
Summary of results : -^
Weight of a Liter of Air.
Found.
Class Average.
Weight under Standard Conditions.
I.
2.
3-
Average,
1.293
LABORATORY EXERCISE, VI.
1. Draw a general conclusion from the result of Exp. 47.
2. Tabulate the weights of a liter of hydrogen, oxygen, and air,
stating the conditions under which the weights were obtained.
Weight of a Liter of Nitrogen. — A liter of nitrogen
under standard conditions weighs 1.256 gm.
CLASS-ROOM EXERCISE. XVIL
Atmospheric pressure.
{a) Meaning.
{p) How measured.
{c) What is "an atmosphere"?
(^) Why does it vary ?
{e) Why is it less on a stormy than on a pleasant day ?
Liquid air.
(rt) Early experiments.
{b) Dewar^s work.
{c) Tripler's process.
(<^) Properties.
{e) Industrial uses.
Historical.
{a) Discovery and properties of argon.
{b) Discovery and properties of helium.
{c) Discovery and properties of "recently announced con-
stituents of the air "
114
Experimental Chemistry.
PROBLEMS. XII.
1. A liter of dry air is passed over heated copper. The copper in-
creases in weight 0.297 gm. What per cent of oxygen by weight is
contained in this sample of air ?
2. A mixture of 25 cc. of air and 50 cc. of hydrogen is exploded, and
the residue measures 60.3 cc. What per cent of oxygen was contained
in this sample of air ?
3. Dumas, in determining the composition of air by passing air over
heated copper and measuring the residual nitrogen, obtained the follow-
ing data : —
Grams.
Weight of tube and copper before experiment . 120.00
Weight of tube and copper after experiment . . 121. 15
Weight of globe exhausted 852.00
Weight of globe and nitrogen 855.85
Calculate the per cent of each constituent by weight and volume.
4. Bunsen analyzed the air at Marburg in 1846 by exploding a
measured volume with a measured volume of hydrogen. Calculate the
per cent of each constituent from the following analyses : —
II.
Air used ....
Air and hydrogen used
Residue ....
Air used . . .
Air and hydrogen
Residue . . .
428 93 cc.
74977 cc.
480.09 cc.
448.00 cc.
743.01 cc.
461.72 cc.
5. A student placed a piece of phosphorus in a volume of air which,
when corrected, was 87.9 cc. The residue when corrected was 70.58 cc.
Calculate the per cent of each constituent. Is the result high or low?
How much?
6. If the composition of normal air by weight is 76.995 per cent
of nitrogen and 23.005 per cent of oxygen, what is the per cent of each
constituent by volume?
7. Dumas and Boussingault in 1841 found in a sample of air
12-373 gm. of nitrogen and 3.68 gm. of ox3'gen. What per cent of
each constituent did they find?
8. What weight of hydrogen is necessary to unite with the oxygen
in 100 gm. of air to form water?
The Atmosphere. iic
9. Assuming that air contains exactly 23 per cent of oxygen by
weight, how many grams of carbpn are necessary to unite with the
oxygen in loogm. of air to form carbon dioxide? (Hint: see equa-
tion (6).)
10. How many kilograms of air are necessary to yield 100 kg. of
oxygen ?
11. 20 cc. of air are mixed with 30 cc. of hydrogen and the
mixture is exploded. What volumes of what gases remain?
12. The simplest equation for the reaction between phosphorus and
oxygen is —
2 P + 5 O = P,05
How much phosphorus is nev^essary to remove the oxygen from a
liter of air?
13. Calculate in the following cases the weight of a liter of air •
(a) Weight of air exhausted 3.4 am.
Volume of air exhausted 2571.6 cc.
(d) Weight of air exhausted 1.3 gm.
Volume of air exhausted 990.6 cc.
If the results are high or low, state how much.
14. Calculate the weight of the following volumes of air: 5 1.,
8.5 1., 10 1.
CHAPTER VIII.
ACIDS, BASES, AND SALTS.
Extensive study of the properties of chemical com
pounds has led to a general, though incomplete, classifica-
tion of many of them into three grand divisions, long known
as Acids, Bases, and Salts. Each division has its distinctive
properties, though the divisions are closely related and
occasionally overlap. The old limits of the classification
and its distinguishing terms have not been modified to meet
the new compounds. The nomenclature of these com-
pounds is antiquated. " Even to-day each of these words
(acid, base, and salt) is encumbered with a residue of
ancient usage from which it is almost impossible to set it
free." The present chapter is, therefore, a preliminary
study, not from choice, but from necessity.
GENERAL PROPERTIES OF ACIDS.
Experiment 48. Supplies: Test tubes, dilute sulphuric, nitric, and
hydrochloric acids, acetic acid, glass rod, litmus paper (both colors),
zinc, calcium carbonate (marble), lime water.
Fill a test tube one-third full of one of the. following dilute acids :
sulphuric, hydrochloric, and nitric : fill another test tube one-third full
of concentrated acetic acid. Label the tubes in some distinguishing
manner.
' (^) Dip a clean glass rod into each acid and cautiously taste it.
Describe the taste by a single word.
{b) Dip a clean glass rod into each acid and put a drop on both kinds
of litmus paper. The striking change is characteristic of acids ; draw
a general conclusion from it.
116
Acids, Bases, and Salts. 117
(c) Slip a small piece of zinc into each test tube successively. If no
chemical action results, warm gently. Test the most obvious product
by holding a lighted match at the mouth of each tube. If no decisive
action results, provide the test tube with a stopper and simple delivery
tube, and collect any product in a test tube over water. This latter
method will probably be unnecessary except with the acetic acid.
Draw a general conclusion from the result in the case of all except the
nitric acid.
((^) Wash the test tubes, add fresh acid, and slip a piece of calcium
carbonate (marble) into each tube successively. Proceed as in (c),
only in this case the gas must be tested not only with a burning match,
but as follows : Dip a clean glass rod into lime water, and hold it in
the escaping gas. What is the gas? (Hint: recall a similar experi-
ment under air). Draw a general conclusion from these results.
The results in Exp. 48 are characteristic of all acids, and
substances acting thus are said to be acid, or to have acid
properties, or to have an acid reaction. The action in {b)
is true only for litmus, but it is a striking, simple test, and
should be remembered. In {c) nitric acid seems excep-
tional. It is not, however, and its action with metals will
be explained under Nitric Acid (see. page 197). Any
metal could be used in place of zinc. The result in {d) is
true of all carbonates, though the action, as in the case
of metals, varies in vigor.
An acid may be a solid, Hquid, or gas, or an aqueous
solution of each. It may be dilute or concentrated, and
the general properties vary somewhat with the strength
of the acid. Each acid has special properties, which will
be considered under the acids^ themselves.
Nomenclature of Acids. — Oxygen is a constituent of
most acids, and the names of these acids correspond to
the proportion of oxygen which they contain. The best
known acid of an element usually has the suffix -ic, e.g.
sulphuric, nitric, phosphoric, etc. If the element forms
ii8 Experimental Chemistryo
another acid containing less oxygen, this acid has the suffix
-oils, e.g. sulphurous, nitrous, chlorous, etc. Some elements
foim an acid containing less oxygen than the -ous acid;
these acids retain the suffix -oiis, and have, also, the prefix
/lypo-, e.g. hyposulphurous, hyponitrous, hypochlorous, etc.
If an element forms an acid containing more oxygen than
the -ic acid, such acids retain the suffix -ic, and have, also,
the prefix /^r-, e.g. persuli^huric, perchloric, periodic, etc.
Acids which do not contain oxygen have the prefix
hydro and the suffix -ie, e.g. hydrochloric, hydrobromic,
etc. This last class is small.
The nomenclature of acids is well illustrated by the
series of chlorine acids : —
ACIDS OF THE ELEMENT CHLORINE.
Name. Formula.
Hydrochloric . HCl
Hypochlorous , . . HCIO
Chlorous HCIO2
Chloric , , . HCIO3
Perchloric . . , , HCIO^
The series is incomplete in the case of many elements,
but the names are formed according to the above principles.
GENERAL PROPERTIES OF BASES.
Experiment /^^.— Supplies: Test tubes, litmus paper (both kinds),
glass rod, short glass tube, solid sodium hydroxide and potassium
hydroxide; ammonium hydroxide and calcium hydroxide.
Make dilute solutions of sodium hydroxide and potassiun, hydroxide
by adding a very small piece to a test tube half full of water, provide,
also, an equal quantity of dilute ammonium hydroxide and of calcium
hydroxide.
{a) Rub a little of e^ch liquid between the fingers, and describe
the feeling. Cautiously taste each liquid by touching to the tip of
the tongue a rod moistened in '-ach, and describe the result.
Acids, Bases, and Salts, 119
{b) Test each liquid with both kinds of litmus paper. Describe the
most striking result. Compare with the action of acids on litmus paper.
{c) Blow through a glass tube into each liquid, except ^the ammo-
nium hydroxide, for several minutes. Describe the change, if any. Then
add to each test tube in succession a few drops of dilute hydrochloric
acid and test the gas evolved for carbon dioxide by holding in each
tube a glass rod with a drop of lime water on the end. The amount of
gas may be small, hence the rod must be held near the surface of the
liquid as soon as possible after the acid has been added. Explain the
chemical change.
The results in Exp. 49 are characteristic of bases, and
substances acting thus are said to be basic, to have basic
properties, or to have a basic react io?i. The word alkaline
is often used instead of basic. The result in {a) is typical,
though it varies in intensity between wide limits from
mere unpleasantness to violent caustic action. The results
in {b) are generally true, though there are important excep-
tions. This is the customary test for a base. The result
in {c) is not usually significant, though it serves to empha-
size the relation which acids and bases bear to carbonates.
Nomenclature of Bases. — The term base is exceed-
ingly broad in meaning and loose in appUcation. Most
bases are Jiydroxides, since they consist of hydrogen and
oxygen combined with a metal. Hydrate is sometimes used
as a synonym of hydroxide, both words suggesting the re-
lation of these compounds to water. Hydroxides are usually
called alkalies, and their chemical action is called alkaline.
There is no general rule covering the nomenclature of
bases. Hydroxides are distinguished from each other by
placing the name of the metal before the word hydroxide,
e.g. sodium hydroxide, potassiunv hydroxide, calcium hy-
droxide, barium hydroxide. The common hydroxides have
long been known by several different names. Thus sodium
I20 Experimental Chemistrv,
hydroxide is called sodium hydrate, or caustic soda; am-
monium hydroxide is called ammonia, or ammonia water.
A GENERAL PROPERTY OF SALTS.
Experiment 50. — Supplies: Test tubes, litmus paper (both colors)^
glass rod, sodium chloride, potassium nitrate, potassium sulphate, and
barium chloride.
JVlake, as in the preceding experiment, dilute solutions of sodium
chloride (common salt), potassium nitrate, potassium sulphate, and
barium chloride. Test each solution with litmus paper and describe the
result. Compare with the action of acids and of bases on litmus paper.
This result is true only of salts which are called normal
salts (see below). It is also characteristic of bodies which
are, chemically speaking, not salts, such as water. Sub-
stances, whether salts or not, which act toward litmus
Hke common salt, are said to be neutral, or to have a
neutral reaction. A salty or saline taste is characteristic
of most salts, and this property doubtless led to their
classification with sodium chloride as a type, since this
compound has been known for ages. Some compounds
which are salts as far as their chemical relations are con-
cerned do not have a neutral reaction. Thus sodium
carbonate, which is the sodium salt of carbonic acid, is one
of the most marked alkahes, being, in fact, known in com-
merce simply as "alkali."
Nomenclature of Salts. — Salts containing only two
elements follow the rule for binary compounds, aiid hence
end in -ide. This suffix is added to an abbreviated form
of the name of the non-metal, e.g. chloride, bromide, fluor-
ide, etc. Notice that the prefix hydro- is omitted, and that
the name of the metal precedes, e.g. sodium chloride, potas-
sium bromide, calcium fluoride, etc. It will soon be shown
experimentally that salts may be regarded as derived from
Acids, Bases, and Salts. 121
acids by replacing the hydrogen of the acid by a metal.
Only the ending of the acid is changed to indicate this
relation. Thus —
-ic becomes -ate.
-oiis becomes -ite.
Hence —
Sulphuric acid forms sulphates.
Sulphurous acid forms sulphites.
Nitric acid forms nitrates.
Nitrous acid forms nitrites.
Chloric acid forms cJilorates.
Hypochlorous acid forms JiypocJilorites.
Permanganic acid forms permanganates.
The name of the replacing metal is retained, e.g. potas-
sium sulphate, sodium nitrite, calcium hypochlorite, potas-
sium permanganate. Notice that the prefixes, hypo- and
per-, are not changed.
Salts in which all the hydrogen is replaced by a metal
are called normal salts. Sometimes only a part of the
hydrogen is replaced, and the salt still has acid properties;
such salts are called acid salts, e.g. acid sodium sulphate is
the salt derived from sulphuric acid by replacing half of
the hydrogen by sodium; its formula is HNaS04; the
formula of normal sodium sulphate is Na2S04. On the
other hand, if a salt is so made ^that it still retains a portion
of the base, it is called a basic salt, e.g. basic nitrate of
bismuth may be regarded as the salt derived from bis-
muth hydroxide, in which only a portion of the base is
replaced by the essential part of nitric acid ; its formula is
Bi(OH)2N03. Basic lead nitrate, Pb(0H)N03, is formed
by boiling solutions of lead nitrate and lead oxide. Basic
122
Experimental Chemistry.
salts bear the same relation to hydroxides as the acid
salts to acids. They are prepared in numerous ways, and
are often complex compounds. Thus basic lead carbon-
ate (white lead) has the formula 2PbC03 . PKOH)^, and
a basic copper carbonate (malachite) has the formula
CuCOg . Cu(0H)2.
Testing the Nature of Substances, — The presence or
absence of an acid or base may be found by the litmus
test, but such a test reveals only the properties of the
substance examined. Other tests must be made to deter-
mine whether or not a particular substance is an acid, base,
or salt.
THE NATURE OF COMMON SUBSTANCES.
Experiment 51. — Determine by the litmus test the nature of lemon
juice, vinegar, soap, sweet and sour milk, washing soda, borax, wood
ashes, faucet water, baking soda, sugar, cream of tartar, the juice of any
ripe fruit and any green fruit.
Make a solution of each of the solids before testing. Tabulate the
results as follows : —
Nature of Common Substances.
Acid.
Neutral.
Acids, Bases, and Salts. 123
LABORATORY EXPZRCISE. VII.
1. Predict the reaction of ''acid phosphate," "^soda water," sour
bread, basic lead acetate, sour wine, tart preserves, pickles.
2. Determine experimentally and then explain the reaction of (a) cop-
per sulphate, (/?) ammonium chioride, (c) alum, (d) potassium carbon-
ate. (If necessary, consult a text-book for the explanation.)
The Nature of Salts. — The relation of salts to acids and
bases is found experimentally by a study of —
THE INTERACTION OF ACIDS AND BASES.
Experiment 52. — Dissolve a small piece of sodium hydroxide in
an evaporating dish half full of water. Slowly add dilute hydrochloric
acid, until a drop taken from the dish by means of a glass rod reddens
blue litmus paper. Then evaporate to dryness by heating over a piece
of wire gauze. Since the residue mechanically holds traces of the excess
of hydrochloric acid added, it is necessary to remove this acid before
applying any test. Heat the dish until all the yellow color disappears,
then moisten the residue carefully with a few drops of warm water and
heat again to remove the last traces of acid. This precaution is essen-
tial to the success of the experiment. Test a portion of the residue
with litmus paper to find whether it has acid, basic, or neutral proper-
ties. T?ste a little. Test (a) a solution of the residue for a chloride,
and {d) a portion of the solid residue for sodium. Draw a definite con-
clusion from the total evidence.
Neutralization. — The operation in Exp. 52 is called
neutralizatioit. Other acids and bases produce similar
results. It shows qualitatively the relation existing between
acids and bases, and the relation of salts to each of the
other two classes. Further information on the subject of
neutralization may be obtained by determining the quanti-
tative relation between acids and bases when representatives
cf the two classes interact.
124
Experimental Chemistry.
A QUANTITATIVE EXAMINATION OF THE INTERACTION
OF ACIDS AND BASES.
Experiment 53. — Method: Neutralize a measured quantity of an
acid with a base.
Apparatus : Two burettes and a support, two beakers of the same
size, two glass rods (blunt form).
Solutions: (a) Add 8 cc. of dilute hydrochloric acid (i vol. of acid
to 3 vols, of water) to 200 cc. of filtered water. Keep the solution in a
clean, stoppered flask or bottle which
is labelled Hydrochloric Acid.
(d) Make a solution of dilute
sulphuric acid (i vol. of acid to
5 vols, of water) as in (a). Label
it Sulphuric Acid.
(6) Dissolve about 2 gm. of so-
dium hydroxide in 200 cc. of filtered
water. Preserve as described above,
and label it Sodium Hydroxide.
(d) Make a solution of potas-
sium hydroxide as in (c). Label
it Potassium Hydroxide.
(e) Dissolve i gm. of solid phe-
nolphthalein in 100 cc. of alcohol (50
per cent — see Table X., App. C).
Keep in a glass stoppered bottle.
This solution is to be used instead
of litmus. It colors alkaline solu-
tions magenta, and undergoes no
appreciable change in acid solutions.
This solution or any other used
similarly is called an md/cator,
because it indicates the nature of a
solution, whether acid or alkaline.
Process : Wash two burettes with
distilled or filtered water, and dry
the inside with a stick having a piece of soft paper attached to one end.
Arrange them as shown in Fig. 51. Label one clamp "acid'' and the
other ''alkali." Fill the burette marked "acid" with the hydrochlui:^
Fig. 51. — Burettes arranged for use.
Acids, Bases, and Salts. 125
acid already prepared, taking care that the liquid is free from air bubbles
at the lower end, and finally that the surface of the acid is level with
the zero mark. Remember that the lowest point of the meniscus is the
correct position. The burette holds 50 cc. from the o mark to the
50 mark, but it should be filled to the glass tip, and drawn off from o to
50 — but no lower. Any number of cubic centimeters may be obtained
by opening the pinchcock and allowing the liquid to flow out of the
glass tip until the level has fallen to the desired point. Fill the burette
marked " alkali '^ with the sodium hydroxide solution in the same way.
When the operation is over, temporarily or finally, the burettes should
be washed and left standing full of clean water.
See that each burette reads o, then place a beaker on a sheet of white
paper under the "acid" burette, open the pinchcock, and allow 15 cc.
of acid to run into the beaker. Add two or three drops of phenol-
phthalein solution. Then allow the sodium hydroxide solution to run
slowly, drop by drop, into the acid. Stir constantly with the glass rod.
As the action proceeds, the magenta color will be more permanent ; add
the alkali very slowly, as this point is approached. The aim is to add
just enough alkali to neutralize the acid. It is customary to pass slightly
beyond this point, just far enough to be sure the color is permanent.
It requires skill and judgment to stop the action at the right time.
Record as shown below the exact number of cubic centimeters of sodium
hydroxide necessary to neutralize the 15 cc. of hydrochloric acid. Draw
into another beaker 15 cc. of acid and find the exact amount of alkali
necessary to neutralize it. Do the same with 10 cc. of acid, refilling
the " alkali " burette, if necessary. The tint should be the same in each
case. Tabulate the results thus : —
I. Hydrochloric acid and sodium hydroxide.
15 cc. of acid neutralize cc. of alkali.
15 cc. of acid neutralize cc. of alkali.
10 cc. of acid neutralize cc. of alkali.
Ratio : 1 Acid, 3-3-2. — Alkali, - - .
If the work has been performed with reasonable accuracy, the results
in I. will show : —
(i) Definite quantities of acid neutralize definite quantities of alkali.
(2) Neutralizing power is independent of the amount of solution
used.
iSame as 15-15-10.
126 Experimental Chemistry.
Wash the burette used for the alkali, and, after drying it, fil\ with the
potassium hydroxide solution. Refill the " acid " burette with hydro-
chloric acid, and proceed with these solutions as before. Tabulate the
results thus : —
II. Hydrochloric acid and potassium hydroxide.
15 cc. of acid neutralize cc. of alkali.
15 cc. of acid neutralize cc. of alkali.
ID cc. of acid neutralize cc. of alkali.
Ratio: Acid, 3-3-2. — Alkali, - - .
These results likewise reveal the facts stated in I. (i) and (2;. They
also give (3) the ratio of the neutralizing power of the two alkalies,
viz., — to — . ((3) is found by comparing the quantity of each alkali
necessary to neutralize a definite quantity — 15 cc. for example — of
acid.)
Wash and dry the "acid" burette, and fill with the sulphuric acid.
Refill the " alkali " burette with the potassium hydroxide solution. Pro-
ceed as before, using successively 15, 15, and 10 cc. of acid and the
necessary amount of alkali. Tabulate the results thus : — •
III. Sulphuric acid and potassium hydroxide.
15 cc. of acid neutralize cc. of alkali.
15 cc. of acid neutralize cc. of alkali.
10 cc. of acid neutralize cc. of alkali.
Ratio: Acid, 3-3-2. — Alkali, - - .
Refill the "acid" burette, but replace the potassium hydroxide bj
sodium hydroxide. Proceed as above, and tabulate the results thus : —
IV. Sulphuric acid and sodium hydroxide.
15 cc. of acid neutralize cc. of alkali.
15 cc. of acid neutralize cc. of alkali.
ID cc. of acid neutralize cc. of alkali.
Ratio: Acid, 3-3-2. — Alkali, - - .
If the whole experiment has been performed with reasonable accu^-
racy, it will also be found that (4) the ratio of the neutralizing power oi
the two alkalies is the same for both acids. An actual example will
make this point clear. A student found that the ratio of the alkalies in
the case of one acid was 11. 7: 19.5, and in the case of the other acid
Acids, Bases, and Salts.
12'
was 12.9: 21.6. Now, by calculation the second ratio is 12.9: 21.5, the
slight difference being readily explained by the fact that the experi-
mental errors are combined in this one quantity, 21.6.
Tabulate the results so that it will be seen that the ratio of the two
alkalies is the same for both acids. Sketch the apparatus.
Discussion of Experiments 52 and 53. — These experi-
ments show that when a base and an acid interact a salt
is formed. Accurate repetitions of these and similar experi-
ments permit the chemical changes involved in the above
experiments on neutralization to be written thus : —
(II)
HCl
+
NaOH --
= NaCI
+
H^O
Hydrochloric
Sodium
Sodium
Water
Acid
Hydroxide
Chloride
36.5
+
40
= 58.5
+
18
HCI
+
KOH =
= KCl
+
u,o
Hydrochloric
Potassium
Potassium
Water
Acid
Hydroxide
Chloride
36.5
+
56
74-5
+
18
H2SO4
+
2 KOH =
= K,S04
+
2H2O
Sulphuric
Potassium
Potassium
Water
Acid
Hydroxide
Sulphate
98
+
I 12
174
+
36
HjSO^
+
2 NaOH --
= Na.2S04
+
2H2O
Sulphuric
Sodium
Sodium
Water
Acid
Hydroxide
Sulphate
98
+
80
142
+
36
(12)
(13)
(14)
These equations show, besides the usual facts, that in
neutralization the hydrogen of the acid and the metal of
the base exchange places, forming thereby a salt and
water.
128 Experimental Chemistry.
Definitions. — The experiments in this chapter show that,
in addition to certain general properties, acids, bases, and
salts are chemically related to a marked degree. No
adequate definition can be given of each. The following,
however, are generally accepted. An Acid is a substance
containing hydrogen which is easily .replaced by a metal.^
The replacing metal may be obtained from a compound of
the metal, r.^'-. the oxide, carbonate, or hydroxide. Acids
are, as a rule, sour, and turn blue litmus red. A monobasic
acid contains in a molecule only one atom of replaceable
hydrogen, e.g. nitric acid, HNO3. Dibasic, Tribasic, etc.,
acids contain, respectively, two, three, etc., replaceable
atoms, e.g. sulphuric acid, H2SO4, and phosphoric acid,
H3PO4. The basicity of acids explains the existence of
normal and acid salts. A base is a substance containing
a metal combined with oxygen and hydrogen, the metal of
which easily replaces the hydrogen of acids when the base
interacts with an acid. Bases, as a rule, turn red litmus
paper blue. A monacid base is one in which the metal
can be replaced by one atom of hydrogen, e.g. sodium
hydroxide, NaOH. Diacid and Triacid bases contain a
metal replaceable by respectively two and three atoms of
hydrogen, e.g. Ca(OH).^and A1(0H)3. A Salt is the main
product of the interaction of an acid and a base. Salts
may be formed in several ways, but they may always be
theoretically viewed as if they were formed directly from
a fundamental acid or base.
A Metal can now be regarded not merely as a hard,
lustrous substance, but as a base-forming element. A
Non-metal likewise may be defined as an acid-forming
element. All non-metallic elements, except hydrogen.
1 Water is an exception.
Acids, Bases, and Salts. 129
and a few others, form oxides which unite with water and
thereby produce an acid ; thus : —
SO3 + H.p = H2SO4
Sulphur Water Sulphuric
Trioxide Acid
These oxides are often called Anhydrides. Similarly, many
metallic elements form oxides which unite with water and
thereby form hydroxides; thus : —
CaO + H2O =
Ca(0H)2
Calcium Water
Calcium
Oxide
Hydroxide
Such oxides are called Basic Oxides. There are some
oxides which are neither acidic nor basic. Hydroxyl is a
name applied to the group OH, which may be regarded
as a factor in the molecules of many compounds. It does
not exist as elements and compounds exist, but in many
reactions this group acts as a unit. It is often called a
radical, owing to its fundamental character, — -j/ being
the specific suffix of many radicals.
CLASS-ROOM EXERCISE. XVIII.
1. Summarize the properties of acids, bases, and salts.
(a) Physical properties.
(d) Chemical properties.
2. Neutralization.
(a) Definition.
(d) Illustration.
(c) Exactly what does Exp. 53 prove?
(d) Application of the principle in quantitative analysis.
(^) Sources of error in Exp. 53, and how avoided.
ijo Experimental Chemistry.
3. Explain and illustrate the expressions: —
(a) Acidity of bases.
{d) Basicity of acids.
(c) "Every oxygen acid has a corresponding anhydride."
(^) " Most normal salts are neutral/'
{e) " Sulphuric acid is the starting-point of sulphates."
(/) "Metals form basic compounds."
(g) " Hydroxyl is a radical."
4. Historical.
(a) Historical development of the terms ac/(/, alkali^ and salt.
(J?) Lavoisier's contribution to the development of the subject of
acids.
{c) Davy's work on the theory of Lavoisier.
5. Miscellaneous.
(i) Explain the terms volatile alkali, caustic alkali, and Jixed
alkali. What are the alkali metals ?
(2) Why are there no acid nitrates ?
(3) Are there any exceptions to the statement, " all substances
from which a metal liberates hydrogen are acids"'?
(4) Literal meaning of acid (adj.), per-, hypo-, alkali.
(5) What is the etymological relation of hydrogen to the terms
hydrate, hydroxide, hydro- (in hydrochloric), anhydride?
(6) The operation in Exp. 53 is called titration. What is
meant by ''titration to color"?
6. Nomenclature.
(a) Name the —
Sodium salt of hydrochloric acid.
Potassium salt of hydrochloric acid.
Lead salt of hydrochloric acid.
Calcium salt of hydrochloric acid.
Barium salt of hydrochloric acid.
Zinc salt of hydrochloric acid.
Silver salt of hydrochloric acid.
{b) Name the same salts of nitric acid and of sulphuric acid-
{c) Name the same salts of hypochlorous acid, of chloric acid,
and of nitrous acid
{d) Name the base corresponding to sodium, potassium, ammo-
nium, calcium, barium, zinc, lead, copper.
Acids, Bases, and Salts. 131
(e) Name the —
Sodium salt of sulphurous acid.
Potassium salt of manganic acid.
Lead salt of nitric acid.
Potassium salt of perchloric acid.
Calcium salt of hydrofluoric acid.
Sodium salt of nitrous acid.
Calcium salt of hypophosphorous acid.
Sodium salt of carbonic acid.
Potassium salt of tartaric acid.
Lead salt of chromic acid.
(/) Significance of the terms hydroxide, hydrate, anhydride, and
hydroxyl.
{g) Select the hydroxyl groups from —
(1) NaOH (3) Ca(OH), (5) A1(0H)3
(2) KOH (4) Ba(0H)2 (6) C^H^COH)
CHAPTER IX.
ATOMS, MOLECULES, AND RELATED SUBJECTS.
Equivalents. — The equivalent of zinc was found in
Exp. 22, and the result, together with similar results,
which may be easily obtained, furnish fundamental data
for the subjects discussed in this chapter.
EQUIVALENT OF MAGNESIUM.
Experiment 54. — Determine the equivalent of magnesium by the
method of Exp. 22. Use the same apparatus. The reaction proceeds
so easily that only about 25 cubic centimeters of warm dilute hydro-
chloric acid (and no platinum wire) are needed. Use from 0.06 to
0.08 gm. of clean magnesium ribbon, which should be slipped into a
short glass tube so that it will sink. Collect the gas in a 100 cc. tube.
Proceed in other respects exactly as directed in Exp. 22. Record and
calculate as follows : —
. Data: —
Weight of magnesium taken
=
gm.
Observed volume of hydrogen
=
cc.
Observed temperature
=
0
Observed pressure
= .
mm.
Tension of aqueous vapor
=
mm.
II. Calculation : —
Formula for reduction to standard conditions —
V {F - a)
V=-
760 (i +. .00366 /j
132
Atoms, Molecules, and Related Subjects. 133
Corrected volume of hydrogen = cc.
Weight of hydrogen = gm.
Since gm. of magnesium is equivalent to gm. of hydrogen,
then gm. of magnesium are equivalent to i gm. of hydrogen.
Equivalent of magnesium =
EQUIVALENT OF ALUMINIUM.
Experiment 55. — Determine the equivalent of aluminium by the
method used for zinc and magnesium. Use from 0.06 to 0.07 gm. of
aluminium for a 100 cc. tube. Use warm concentrated hydrochloric
acid instead of sulphuric acid ; no platinum wire is needed. The action
may be vigorous and must be watched to prevent the metal from being
forced over into the dish. Otherwise, proceed as in the other deter-
minations. Record and calculate as in Exp. 54.
The equivalent or equivalent weight of an element was
defined in Chapter III. as that weight of an element chemi-
cally equivalent to one part by weight of hydrogen. More
specifically it is the number of grams of an element which
combines with or replaces one gram of hydrogen. Thus
in Exp. 22, 32.5 gm. (approximately) of zinc were found to
be equivalent to i gm. of hydrogen. In Exps. 54 and 55,
I gm. of hydrogen was found to be equivalent respectively
to 12 gm. of magnesium and 9 gm. of aluminium (approx-
imately), and these numbers, 12 and 9, are the equivalents
of these two elements. Some metals do not liberate
hydrogen easily from acids, and their equivalents are
indirectly found. Sodium, for example, liberates hydrogen
from water, and careful experiments show that 23 gm. of
sodium are equivalent to i gm. of hydrogen. Oxygen is
withdrawn from water by passing steam over powdered
iron (see Exp. 40), water itself is formed by passing
134 Experimental Chemistry.
hydrogen over heated copper oxide (see Exp. 44), and both
of these experiments, if performed accurately, show that
8 gm. (approximately) of oxygen are equivalent to i gm. of
hydrogen. Exp. 1 5 also shows that 8 gm. of oxygen always
combine with 12 gm. of magnesium, a number which has
been found by an independent experiment (Exp. 54) to
be the equivalent of magnesium.
Many metals precipitate another when one is suspended
in the solution of the other, and if the metals are weighed
the equivalent of one is easily found, thus : —
Grams.
I gm. of zinc precipitates .97 of copper.
I gm. of zinc precipitates 333 of silver.
I gm. of magnesium precipitates 2.64 of copper.
I gm. of magnesium precipitates 9.00 of silver.
I gm. of copper precipitates 3.43 of silver.
Now if we recalculate these values and express the results
in terms of the equivalents already known, the table reveals
other equivalents, thus : —
G-RAMS. Grams.
32.5 of zinc are equivalent to 317 of copper.
32.5 of zinc are equivalent to 108.2 of silver.
12. of magnesium are equivalent to 31.7 of copper.
12. of magnesium are equivalent to 108. of silver.
31.7 of copper are equivalent to 108. of silver.
Actual analysis of chlorides reveals the following facts : —
I gm. of hydrogsn combines with 35.5 gm. of chlorine.
23 gm. of sodium combine with 35.5 gm. of chlorine.
16 gm. of sulphur combine with 35.5 gm. of chlorine.
39 gm. of potassium combine with 35.5 gm. of chlorine.
108 gm. of silver combine with 35.5 gm. of chlorine.
Atoms, Molecules, and Related Subjects. 135
Analysis of bromides gives the following results : —
I gm. of hydrogen combines with 80 gm. of bromine.
23 gm. of sodium combine with 80 gm. of bromine.
39 gm. of potassium combine with 80 gm. of bromine.
108 gm. of silver combine with 80 gm. of bromine.
Analysis of some oxides shows the following results in
addition to those given above : —
23 gm. of sodium combine with 8 gm. of oxygen,
39 gm. of potassium combine with 8 gm. of oxygen.
108 gm. of sih-er combine with 8 oti. of oxygen.
-J &"
Some sulphides show similar results : —
I gm. of hydrogen combines with 16 gm. of sulphur.
23 gm. of sodium combine with 16 gm. of sulphur.
39 gm. of potassium combine with 16 gm. of sulphur.
108 gm. of silver combine with 16 gm. of sulphur.
A summary of the above data may be tabulated thus : —
Element. Equivalent.
Hydrogen i (by definition)
Oxygen 8
Magnesium 12
Aluminium 9
Zinc , . . . 32. 5
Copper 31.7
Silver 108
Sodium 23
Potassium 39
Chlorine 35.5
Bromine 80
Sulphur 16
136
Experimental Chemistry.
A comparison of some equivalent weights and the approx-
imate atomic weights of the same elements accepted by
chemists reveals a remarkable coincidence.
Comparative Table of Equivalents and Atomic Weights.
Hydrogen .
Oxygen
Chlorine .
Bromine .
Sulphur
Zinc . .
Copper
Magnesium
Sodium
Potassium
Silver . .
Aluminium
In many cases they are the same, but in some cases the
atomic weights are a simple multiple of the equivalent.
The first weights called atomic weights were, in many in-
stances, identical with equivalents, and it has taken much
labor and thought to select the multiple which is the proper
atomic weight of the element in question.
Atomic weights cannot be interpreted without a prelimi-
nary consideration of the atomic theory.
The Atomic Theory. — Up to the time of Dalton (1766-
1844) there were two views regarding the constitution of
matter. The first of these is known as the dynamic. This
view advocated infinite subdivision ; it claimed that there
Atoms, Molecules, and Related Subjects. 137
is no limit to the division of matter. A piece of iron, for
example, according to this theory, could be divided forever.
The other view is known as the atomic. This view advo-
cated finite divisibility. It said that if subdivision goes on,
particles will finally be reached which cannot be further
subdivided These particles were called atoms by the
Greeks, and this latter theory is substantially the atomic
theory of the Greeks.
Both these views are old ; they date from the time of
Lucretius (95-52 b.c). They are opposed to each other,
they are entirely speculative, they are metaphysical, un-
scientific, and finally have no relation to chemistry as we
know it.
Dalton shared the general views of his contemporaries
regarding the constitution of matter. He believed in the
solid particles of Newton, i.e. in the corpuscular theory.
He also believed in the existence of individual gases, par-
ticularly those which were discovered by Priestley, Caven-
dish, Black, and others. He believed in the law of definite
proportions by weight, and in the law of multiple propor-
tions (see page 201), discovered by Dalton himself, both
of which emphasized the general idea of combination by
units.
Dalton's meteorological observations led him to believe
that aqueous vapor exists separately from the other con-
stituents of the air, also that gases themselves are made up
of distinct particles. His work on the solubility of gases
in water led him to undertake to determine "the relative
weights of ultimate particles of bodies."
Now the above work was largely physical. His views
on all matter seem to be the result of shrewd generaliza-
tions from his general and vaguely supported ideas of
matter, mainly in its physical relations. These ideas were
138 Experimental Chemistry.
expressed at several times and more particularly in a course
of lectures delivered in 1809- 18 10 at the Royal Institution
in London. Regarding atoms, he said in substance in this
lecture : " Matter is composed of particles of definite size
and weight. These particles are called atoms. All atoms •
of the same kind of matter are alike. Atoms of different
kinds of matter are different. All atoms combine by
wholes to form compounds ; their relative weight only can
be determined." This is the atomic theory of Dalton.
One of the essential properties of matter is weight, and
Dalton's first efforts were to determine some of the weights
of atoms. To do this he formulated rules and actually
found several weights, which he called atomic weights.
This work, mingled with crude speculation on the shapes
of atoms, seems to have been done at different times and
probably preceded to some extent Dalton's final views and
work on atoms, though the public expression of his views
undoubtedly gave an impetus toward the gravimetric con-
ception of atoms. Dalton's contemporaries did not receive
his theory favorably. It is true, no facts, then known, con-
flicted with the theory, but it is equally true that no set
of facts actually supported it. The work of determining
atomic weights spread, however, and was given a strong
impetus by the Swedish chemist, Berzelius. It has never
ceased to attract skilful, persistent workers.
Analysis of chemical compounds reveals the proportions
of the components by weight. Thus, one part by weight
of hydrogen combines with 35.5 parts by weight of chlorine,
and with 80 of bromine. These parts by weight of chlorine
and bromine combine in each instance with 39 parts by
weight of potassium. The early experimenters soon found
that some elements enter into combination with other
elements in several proportions. Thus, 8 parts by weight
Atoms, Molecules, and Related Subjects. 139
of oxygen combine with one part by weight of hydrogen,
but in the larger number of instances 16 parts by weight
of oxygen combine with various parts by weight of other
elements, e.g. zinc, sulphur, and magnesium. Further-
more, sulphur combines with hydrogen in the proportion
of 16 parts to I, but in the vast majority of cases the num-
ber of parts by weight of sulphur which combines with
other elements is found to be 32. In a word, multiples of
combining weights were found, and the problem then
reduced itself to selecting the particular multiple to be
called the atomic weight. That number which represents
the combining proportion of an element in the greatest
number of cases, its favorite number, so to speak, was se-
lected as the atomic weight of that element. The number,
then, which we call the atomic weight represents not an
absolute weight, but a relative weight, i.e. how many times
heavier the atom in question is than the atom of hydrogen.
The absolute weight of an atom of hydrogen is not known.
Its weight is arbitrarily chosen as one.
The atomic theory is viewed at the present time as a
satisfactory explanation of the laws of combination of
matter revealed by experiment. It is generally adopted
as a working hypothesis, not as a fact.
An atom according to the present views is that particle
of matter which has not yet been subdivided into smaller
particles. It combines with itself or with other atoms and
thereby forms molecules.
Atomic Weights. — The atomic weight of an element,
therefore, is that number which represents how many times
heavier the atom of an element is than the atom of hydro-
gen (see " Various Figures for Atomic Weights," page 142).
Hence, the numbers hitherto called the equivalents of
140 Experimental Chemistry.
the elements are also, in many cases, the atomic weights of
the respective elements. Thus,
23 is the atomic weight of sodium.
108 is the atomic weight of silver.
80 is the atomic weight of bromine.
127 is the atomic weight of iodine.
35.5 is the atomic weight of chlorine.
The determination of exact atomic weights is a difficult
operation, and no single method furnishes sufficient data
for a final selection of the numbers. If atoms combined
in only one proportion, atom for atom, analysis would re^
veal the proportions, and if one were the standard, or if its
weight in terms of the standard were known, the relative
weight of the other could be easily calculated. For exam-
ple, Stas, a Belgian chemist who made masterly determina^
tions of atomic weights, found that 121.4993 gm. of silver
chloride were formed by burning 91.462 gm, of silver in
chlorine. It was known to him that an atom of silver
unites with one atom of chlorine to form silver chloride,
and that the atomic weight of chlorine is 35.453. Hence,
he calculated the atomic weight of silver thus: —
121.4993
91.4620
30.0373 weight of chlorine used.
91.462: 30.0373 ::,r: 35.453
X = 107.95, atomic weight of silver.
Since many atoms combine in several proportions, the
number selected as the atomic weight is based on consid-
Atoms, Molecules, and Related Subjects 141
erations such as vapor density, agreement with the periodic
law, specific heat, etc. Only one of these approximate or
check methods, however, will be considered.
Law of Specific Heats. — Dulong and Petit in 1819 an-
nounced the Law of Specific Heats, viz. : —
TJic product of the atomic zveigJit and the specific heat of
the solid elements is a constant quantity.
This law is not perfectly accurate, since in many cases
this product is only approximately 6.4 — the number se-
lected by Dulong and Petit as the constant — as appears
from the
Table of Specific Heats. ^
Element.
Specific Heat.
Atomic Weight.
Product.
Aluminium .
Calcium . .
Copper . .
Iron . . .
Lead . . .
Magnesium .
Potassium .
Silver . . .
Sodium . .
Sulphur . .
Tin . . .
Zinc . . .
0.214
0.170
0.095
0.114
0.031
0.250
0.166
0.057
0.293
0.178
0.055
0.094
27.0
39-91
63.18
55.88
206.4
24-3
39-03
107.66
23.00
31.98
118.8
65.1
5.8
6.8
6.0
6.4
6.4
6.1
6.5
6.1
6.7
5-7
6.5
6.1
An example will make this law more intelligible. The
specific heat of silver is .057; if 6.4 is divided by this num-
iFrom Meyer's "Outlines of Theoretical Chemistry," p. 22.
142 Experimental Chemistry.
ber, the quotient is 1 12.2. This is not the accepted atomic
weight of silver, but it checks the other determinations,
since 112.2 points to 108 and not to any of its multiples.
On the other hand the specific heat of mercury is .032 ; if
6.4 is divided by this number, the quotient, 200, is the atomic
weight of mercury, and this number is confirmed by other
methods.
Various Figures for Atomic Weights. — Obviously the
figures for the atomic weight of the same element will
vary with the standard. Hence different atomic weight
tables will be found, some on one standard, some on
another. Hydrogen was, until recently, the generally
adopted standard. Since many atomic weights" are now
directly determined with reference to oxygen, the hydro-
gen standard is being replaced by the oxygen standard, i.e.
16 and not i is the standard atomic weight. But as the
atomic weight of hydrogen becomes only 1.008, if oxygen
is taken as exactly 16, the variation in the usual table of
atomic weights is very slight, especially when whole num-
bers are used. Appendix C contains a table of atomic
weights on both standards. The approximate atomic
weight (see App. C, Table I.) of the most common ele-
ments should be learned, not slavishly, but by constant
association.
Molecular Weights. — Since atoms combine to form
molecules, a molecular weight is the sum of the weights of
the atoms in a molecule. If the molecule is the simple
combination of two atoms, then the molecular weight is
the sum of the respective atomic weights. Thus experi-
mental evidence points conclusively to the fact that sodium
chloride contains one atom of sodium and o.ne atom of
chlorine, hence its formula is NaCl and its molecular
Atoms, Molecules, and Related Subjects. 143
weight is 58.5 (23 + 35.5). Since molecules may theoreti-
cally contain any number of atoms, the molecular weight
and formula are rigidly connected. It is customary to
assume the simplest formula until evidence reveals an
error. Thus HO was formerly regarded as the formula
of water ; but when the density of steam showed that the
molecular weight must be 18, this fact, together with other
data of a similar value, led to a change from HO to the
present formula, H2O.
Avogadro^s Law. — Various generalizations have con-
tributed to the solution of the problem of molecular weight
determination. The oldest of these is Avogadro's Law,
which is —
There is an equal number of molecules in equal volumes
of all gases at the same temperature and pressure.
This law was announced by Avogadro, an Italian physi-
cist, in 181 1, but was not recognized .by chemists and
physicists as a fundamental law of science until about
1868. It has been repeatedly verified experimentally and
deduced mathematically (from the kinetic theory of gases).
This law means that the weights of equal volumes of gases
are in the same proportion as the molecular weights of
the gases themselves. Thus the weight of a liter of oxy-
gen was found by Exp. 17 to be approximately 1.43 gm.
A liter of hydrogen under the same conditions weighs
.0896 gm. But .0896 and 1.43 are in the same ratio as
2 to 32, the accepted molecular weights of hydrogen and
oxygen, since it is believed that these gases, like most ele-
mentary gases, have two atoms in a molecule. The den-
sity of all compound gases has been calculated, and the
relations of some appear in the following table : —
144
Experimental Chemistry.
Gas.
Weight of a Liter
in Grams under
Standard Con-
ditions.
Weight of a Mole-
cule compared
with the Weight of
One Molecule of
Hydrogen (H2 = 2).
Weight of a Mole-
cule compared
with the Weight
of One Atom of
Hydrogen (H = i).
Carbon dioxide
Hydrochloric acid
Chlorine . . .
Ammonia . . .
Nitrogen . . .
Steam ....
1.97
I.61
3.18
0.77
1.256
.806
22
18.25
35-5
8.5
14
9.0
44
36.5
71
17.0
28
18.0
The numbers in the last column are the molecular
weights of the respective gases. They are equal to the
sum of the weights of the combination of atoms which
represents a molecule of the different gases. The molecu-
lar weight of a gas is, therefore, twice its specific gravity
referred to hydrogen.
Calculation of Atomic and Molecular Weights from Ex-
perimental Data. — Many of the experiments in this book
furnish data for the calculation of approximate atomic and
molecular weights.
The atomic weight of magnesium may be calculated by
utilizing the results of Exps. 1 5 and 54. A student found
that .37 gm. of magnesium combined with .25 gm. of oxygen
in forming magnesium oxide. Accepting 16 as the atomic
weight of oxygen, we may write the proportion —
Weight of . Weight of ^ ^ Atomic weight
oxygen " magnesium " " of oxygen
.25
37
16
Atomic weight
of magnesium
237
Atoms, Molecules, and Related Subjects. 145
The same student verified this result by finding the
equivalent of magnesium to be 11.95. The atomic weight
of magnesium is twice the equivalent, hence the atomic
weight becomes 23.9 {i.e. 2 x 11.95), a result agreeing
fairly well with 23.7 previously found. The approximate
atomic weight of magnesium is 24.
Again, from the results of the first experiment the mo-
lecular weight of magnesium oxide may be found. Thus
.25 gm. of oxygen formed .62 gm. {i.e. .2.5 -f- .37) of mag-
nesium oxide. Hence the proportion : —
Weight of . Weight of . . Atomic weight . Molecular weight of
oxygen * magnesium oxide ' ' of oxygen ' magnesium oxide
.25 : .62 :: 16 \ x
^r = 397
The molecular weight of magnesium oxide is 40 {i.e.
24 + 16), which is in fair agreement with the above experi-
mental determination. It is generally accepted that a
molecule of magnesium oxide contains one atom each of
oxygen and magnesium.
The atomic weight of zinc is similarly found. A student
obtained these results : —
1.05 gm. of zinc united with .26 gm. of oxygen to form
1.3 1 gm. of zinc oxide. Using the same form of proportion
as in the case of magnesium, we may write
.26 : 1.05 : : 16 \x
x = 64.61
An average of the determinations of a large class work-
ing in the same laboratory on the equivalent of zinc was
32.45. Since the atomic weight of zinc is twice the equiva-
lent, the second result gives 64.9 as the atomic weight of
zinc. The approximate atomic weight of zinc is 65.
146 Experimental Chemistry.
A student in performing Exp. 16 obtained these results :
Grams.
Weight of potassium clilorate used 1.500
Weight of oxygen lost .... .585
Weight of potassium chloride left .915
Now chemists believe that a molecule of potassium
chlorate contains three atoms of oxygen, hence we may
write the proportion : —
Weight of three Molecular weight of Weight of Weight of
atoms of oxygen ' potassium chlorate ' ' oxygen ' potassium chlorate
or 48 : X :: .585 : 1.5
X = 123.07
The molecular weight of potassium chlorate, found by
adding the approximate atomic weights, is 122.5. Again,
the molecular weight of potassium chloride may be found
by the proportion : —
Weight of three Molecular weight of Weight of Weight of
atoms of oxygen * potassium chloride ' " oxygen ' potassium chloride
or 48 : X :: .585 : .915
,r= 75.07
The molecular weight of potassium chloride, found by
adding its atomic weights, is 74.5, it being assumed that
the molecule contains one atom each of potassium and
chlorine.
The molecular weight of water may be calculated from
the result of Exp. 44. A student found that .905 gm. of
oxygen formed 1.021 gm. of water. Hence: —
Weight of
oxygen
or ,905
;r-= 18.05
Weight of
Atomic weight
Molecular weight
water
of oxygen
of water
1.02 1
:: 16
: X
Atoms, Molecules, and Related Subjects. 147
This result agrees closely with 18, the molecular weight
of water. These results are ouly approximate, but they
serve to illustrate some of the methods employed in the
determination of atomic and molecular weights.
Valence. — Investigation shows that each element has
power to hold in combination a definite number of atonis
of other elements, and that the power varies with the
elements which interact. Thus chlorine, oxygen, nitrogen,
and carbon, under certain conditions, combine respectively
with I, 2, 3, and 4 atoms of hydrogen. This capacity is
called valence. It is measured by the number of hydrogen
atoms with which one atom of the element combines to
form a molecule. Hydrogen and all elements combining
with one atom of hydrogen, or its equivalent, are called
univalent elements, e.g. sodium and potassium ; those com-
bining with two atoms are called bivalent ; with three, tri-
valent ; with four, tetravalent ; with five, pentavalent, or
quinquivalent. Elements of the same valence replace each
other atom for atom. If the valence is different, such a
number of atoms or groups of atoms must be taken that
the valence of the interchanging parts will be equal. Thus
Cu(N03)2 is the formula of copper nitrate, because in this
compound copper has the valence two and is equivalent to
two NO3 groups which are univalent, since they combine
directly with one atom of hydrogen. Similarly, Na20' is
the formula of sodium oxide, but CaO is the formula of
calcium oxide, the valence of oxygen and calcium being
two, while the valence of sodium is one. The valence
of an element is the same in the same compound, but
often varies in different compounds, according to circum-
stances, such as temperature and character of the other
element. Valence must not be assumed, but learned by
148 Experimental Chemistry.
experience, since the whole theory is not a sharply defined
conception.
Percentage Composition. — The formula of a compound
is not only an abbreviation of its name, but also an expres-
sion of its composition. Thus H2SO4 means that in every
molecule of sulphuric acid 2 parts are hydrogen, 32 are
sulphur, and 64 are oxygen, if the parts are taken by
weight. As there are 98 parts (2 + 32 + 64) in the mole-
cule, then —
— are hydrogen.
98
^ are sulphur.
98
-A are oxygen.
98
Should the composition in per cent be desired, it can
be calculated thus : —
2:
198:
: X :
: 100
x =
2.04% of
hydrogen.
32:
198:
: X
: 100
X =
32.65% of
sulphur.
64:
:98:
: X
: 100
X =
Total,
65.31% of
oxygen.
100.00%
By this method it is possible to calculate the percentage
composition of any compound, if its formula is known.
Formula. — If the percentage of each element in a com-
pound is divided by the atomic weight of the element, the
quotients will be in the same ratio as the number of atoms
Atoms, Molecules, and Related Subjects. 149
of each clement in the compound. Thus the percentage
composition of sulphuric acid is —
H = 2.04.
S = 32.65.
0=65.31.
Now if each percentage is divided by the proper atomic
weight, we may write (approximately) : —
-^ = 2 = number of hydrogen atoms.
■52.6c
= I = number of sulphur atoms.
■^'^ = 4 = number of oxygen atoms.
Hence the simplest formula of sulphuric acid is H2SO4.
The simplest formula calculated from the percentage com-
position is not always the molecular formula, that is, it
does not always express the composition and number of
atoms which are in a molecule of the compound when
in a gaseous state. Since every formula is designed to be
a molecular formula, it is usually necessary to determine
it from independent experimental data. The molecular
weight of a compound is twice its vapor density, hence
if the density is known or can be determined, then the
molecular formula can easily^be calculated from the sim-
plest formula. For example, the simplest formula of
water based on gravimetric analysis is HO, but since its
vapor density is 9, its molecular weight is 18, a number
which requires the formula H2O. Again, the formula
H2O is supported by numerous facts independently dis-
covered, such as the atomic weight of oxygen, the divisi-
i^o Experimental Chemistry.
bility of hydrogen into two equal parts, and the ratio of
the volumes of the components. Another illustration will
make the subject still clearer. The simplest formula of a
compound of carbon and hydrogen was found to be CH2.
Its vapor density was found to be 81.4. Therefore its
molecular weight must be 162.8, or nearly twelve times
that corresponding to CH2, hence the molecular formula
is C12H24. Molecular formulas of other compounds may
be similarly found.
Until recently the vapor density was the only reliable
check on molecular weight determinations, but, as will
presently be seen, newer and more comprehensive methods
are now applicable.
Elevation of Boiling Point and Depression of Freezing
Point. — It has long been known that a solution of common
salt boils at a higher temperature and freezes at a lower
temperature than pure water. (See Exp. 31.) The elevation
of the boihng point and depression of the freezing point are
controlled by laws, which are easily determined by methods
appHcable to many compounds, and molecular weights are
now studied by these methods, especially by the depres-
sion of the freezing point. Investigation of molecular
weights determined by the latter method shows that the
depression of the freezing point of a solution is propor-
tional to the weight of the substance dissolved in a given
volume, and that the same number of molecules of any
dissolved substance produces the same fall in temperature
of the freezing point of the solution; i.e. for solutions
of different substances but of equal concentration, the
depression of the freezing point is inversely proportional
to the molecular weights. Thus if X is the depression
found experimentally to be produced by a one per cent
Atoms, Molecules, and Related Subjects. 151
solution of sugar, and V the depression produced by a
one per cent solution of urea, the following proportion may
be written : —
V: X:: mol. wt. of sugar : mol. wt. of urea.
The molecular weight of sugar is known to be 342, and
the molecular weight of urea calculated from the above
proportion is 60 — a number agreeing closely with the
molecular weight found by other methods. This method
is used to determine whether a molecular weight is a given
number or its multiple, a question often arising in the case
of different organic compounds having the same percentage
composition. The method is applicable, however, only in
those cases in which (i) there is no chemical action between
the solvent and the solute, and (2) the solute does not
decompose.
Ions and Ionization. -^ The molecular weights of many
substances which had been determined by the older
methods agree with the values obtained by the newer
methods. There is a marked disagreement, however, in the
case of solutions of strong acids, strong bases, and salts,
or in general, in the case of ehxtrolytes, i.e. all substances the
aqueous solutions of which conduct electricity. In some
instances the molecular weight is only half of the molecular
weight found by other methods, i.e. the solution acts as if
twice as many molecules were in solution. Thus the molec-
ular weight of sodium chloride was found by two experi-
ments to be 29.4 and 30.2, instead of 58.5. An extensive
experimental study of solutions, supplemented by keen
conclusions, derived from facts of the most varied kind,
has led to the general adoption of the theory — known as
the Theory of Electrolytic Dissociation — that —
152 Experimental Chemistry.
When acidsy bases, and salts dissolve in water their mole-
cules partly split up into iojis.
Ions are atoms, or groups of atoms, charged with elec>
tricity ; they are carriers of electricity, which they give up
when they reach the electrodes of the battery, or source of
electric current. There are always two kinds in every
solution, — negative ions or anions, and positive ions or
cations. Hydrogen, metals, and metallic radicals (Hke
ammonium) form positive ions ; salt radicals, like SO4, and
the hydroxyl group OH, form negative ions. Potassium
chloride in solution is largely split up into the ions, K and
- . + -
CI ; potassium nitrate becomes largely K and NO3.
This dissociation, or ionization, increases as the dilution
increases, and in the case of strong acids, bases, and their
salts the dissociation is nearly complete even in moderately
concentrated solutions. The percentage of dissociation of
some familiar substances (normal solutions at 18°) is shown
in the following —
Table of Ionization.
Substance.
Per Cent of Ionization.
Hydrochloric acid
Nitric acid
78
82
Sulphuric acid
Potas^'uin chloride
51
75
64
53
77
77>
Potassium nitrate ....
Potnssiiim siilnhatp .........
Potassium hvdroxide
Sodium hydroxide
Applications of the Theory of Electrolytic Dissociation. —
Many reactions, already studied, are more intelligible when
Atoms, Molecules, and Related Subjects. 153
interpreted in the light of this theory. Thus, when solu-
tions of silver nitrate and sodium chloride are mixed, in-
soluble silver chloride is formed. The two solutions before
mixing contain the ions of sodium, chlorine, silver, and the
group NO3; when mixed, the ions of silver and chlorine
unite to form silver chloride, a compound insoluble in water,
and hsnce not decomposed or ionized ; the solution still
contains ions of sodium and of NO3, together with the ions
of the salt added in excess. If solutions of silver nitrate
and potassium chlorate are mixed, no silver chloride is
formed, because no chlorine ions are available. Po-
+ -
tassium chlorate dissociates into the ions • K and CIO3.
Ordinary tests are tests for ions. This theory has been
widely applied by the German chemist, Ostwald, to the
explanation of many fundamental facts of analytical
chemistry.
CLASS-ROOM EXERCISE. XIX.
1. Explain and illustrate the following: —
(«) Atom, molecule, mass.
{b) Equivalent, atomic weight, molecular weight.
{c) Specific heat.
(^) Ion, anion, cation, ionization, electrolytic dissociation, de-
pression of freezing point, elevation of boiling point.
2. Historical.
{a) Essential facts in the life of Dalton.
(J?) The contributions of Berzelius to the subject of atomic
weights.
{c) The work of Stas on atomic weights and its value at the
present time.
{d) Dulong and Petit and their joint labors.
(e) Avogadro, Ampere, and Cannizzaro, and their contributions
to Avogadro's Law.
1^4 Experimental Chemistry.
3. The modern theory of solution.
(a) Historical development.
(d) Brief but exact statement of the theory.
(c) The hydrate theory.
(d) Application to simple reactions.
4. Valence.
(a) Give the valence of —
Sodium in sodium chloride.
Calcium in calcium chloride.
Oxygen in water.
Oxygen in calcium oxide.
Nitrogen in ammonia gas (NH..).
Nitrogen in nitric acid (HNO,).
Sulphur in sulphur dioxide (SO2).
Sulphur in sulphur trioxide (SOo).
Sulphur in sulphuric acid.
Carbon in carbon dioxide.
Carbon in carbon tetrachloride (CCl^).
Zinc in zinc sulphate.
Magnesium in magnesium sulphate.
PROBLEMS. XIII.
1. If 0.34 gm. of sodium yields 165 cc of hydrogen (corrected
volume), what is the equivalent of sodium ?
2. If 0.48 gm. of magnesium yields 443 cc. of hydrogen (corrected
volume), what is the equivalent of magnesium ?
3. A piece of zinc weighing 1.95 gm. evolved 670 cc. of hydrogeii
from sulphuric acid. Calculate the equivalent of zinc.
4. If 0.48 gm. of zinc replaces 1.6 gm. of silver, calculate the
equivalent of silver.
5. A gram of aluminium yields 1271 cc. of hydrogen at lo^C. by
interaction with hydrochloric acid. What is the equivalent and the
atomic weight of aluminium ?
6. When 0.5 gm. of aluminium was warmed with potassium hydrox-
ide, 660 cc. of hydrogen at 13' C. were liberated. What is the atomic
weight of aluminium ?
Atoms, Molecules, and Related Subjects. 155
7. If 0.3697 gm. of aluminium liberates 0.04106 crm. of hydroiijen
from sodium h3^droxide, what is the atomic weight of aluminium ?
8. A chemist found that 8.2144 gm. of alum, (NH,)._,SO,Al.,(SO^)...
24 H^,0, dried at 2i''-25"C., yielded 0.9258 gm. of aluminium oxide,
AUO;.. What is the atomic weight of aluminium ? (Assume O- 15.961,
S = 31.966, N = 14.01.)
9. The same chemist found that 6.9617 gm. of aluminium bromide,
AlBi'y, required 8.4429 gm. of silver to precipitate all the bromine.
What is the atomic weight of aluminium ? (Assume Ag ~ T07.649,
and Br = 79 754.)
10. Ii 0.2 gpi. of magnesium liberates 195.5 cc. of hycu^ogen at 13^ C,
what is the atomic weight of magnesium ?
11. A compound of silver and chlorine weighing 121.4993 gm. was
produced by heating 91.462 gm. of silver in a current of chlorine gas.
What is the atomic weight of silver ? (Assume CI = 35-5, and that
one atom of silver unites with one atom of chlorine.)
12. A molecule of calcium oxide contains one atom of calcium and
one atom of oxygen. The percentage composition is 71.43 per cent of
calcium and 28.57 per cent of oxygen (O = 16). What is the atomic
weight of calcium ?
13. Calculate the percentage composition of —
(a) Water.
(d) Magnetic oxide of iron, Fe.jO^.
{c) Crystallized sodium carbonate, Na^^CO;,. io^H.>0.
(rt') Ca,(PO,),.Ca,,(PO,)F.
(e) Alum, Al.K,(SO,)4 . 24 H,0.
(/) Zinc blende, ZnS.
(g) Calamine, zinc carbonate, ZnCOo.
(/i) Feldspar, K.^Al, S{^p^^..
(/) Cu,(P0,),.3H,0.
(/•) Ca,,,(P0,),.CaC03.
14 Calculate the percentage composition of —
(a) Potassium chlorate.
(d) Sugar, Ci.H.sO^.
(c) Calcium sulphate, CaSO^.
{d) Zinc sulphate.
{e) Magnesium oxide.
(/) Copper oxide, CuO.
156 Experimental Chemistry.
15. Calculate the simplest formula of the compounds which have
the indicated percentage composition, and give the name of the
compound : —
(a) Hydrogen = 11. 11 (c) Carbon = 27.27
Oxygen = 88.89 Oxygen = 72.72
(^) Sodium — 32.39
Sulphur = 22.54
Oxygen = 45.07
16. Calculate the simplest formula of the compounds which have
the indicated percentage composition : —
(a) Nitrogen = 82.353 (c) Hydrogen = i.o
Hydrogen = 17.647 Carbon = 11.99
(d) Oxygen = 30 Potassium = 39.06
Iron = 70 Oxygen = 47.95
ly. Calculate the molecular weight of the following compounds by
finding the sum of the atomic weights : —
(a) Copper sufphate. (e) Sodium hydroxide.
(d) Barium chloride. (/) Potassium hydroxide.
(^) Manganese dioxide. (^) Sodium carbonate.
(d) Calcium oxide. (//) Potassium nitrate.
18. A liter of sulphurous oxide gas (SO^,) weighs 2.8672 gm. What
is its molecular weight ?
19. If 1500 cc. of carbon monoxide gas (CO) weigh 1.8816 gm.,
what is its molecular weight ?
20. A volume of gas measures 5.40675 1. at 15° C. and 740 mm.
and weighs 7.616 gm. What is its molecular weight ?
21. Calculate the molecular formula of the compouads having the
Indicated percentage composition and vapor density : —
(a) Carbon = 73.8 (c) Carbon = 39.9
Hydrogen = 8.7 Hydrogen = 6.7
Nitrogen = 17.1 Oxygen = 53.4.
Vapor density = 80.2 Vapor dens'tv - 3^-S
(b) Carbon = 92.3
Hydrogen = j.j
Vapor density = 38.8
CHAPTER X.
CHLORINE AND HYDROCHLORIC ACID.
Chlorine. — It has been shown in previous experiments
that hydrochloric acid contains hydrogen, but as hydro-
chloric acid is not an element, it must contain something
else besides hydrogen. Chemists have shown that hydro-
chloric acid is a compound of two elements, and confirma-
tory evidence of this fact will be obtained later. Obviously,
then, to obtain this other constituent of hydrochloric acid,
we must treat hydrochloric acid with something which will
liberate this constituent and at the same time combine with
the hydrogen. Several experiments have shown that
oxygen combines readily with hydrogen, hence if hydro-
chloric acid is heated with some substance which yields
oxygen easily, this other element in the acid will be liberated,
provided it does not combine easily with oxygen. It has
also been shown that some oxides when heated yield oxygen,
and, as a matter of fact, manganese dioxide was used by
early chemists to prepare oxygen. Furthermore, Scheele,
the discoverer of this other constituent of hydrochloric acid,
actually prepared it by heating hydrochloric acid with ores
(mainly oxides) of manganese. He did not recognize the
new product as an element and it was not until about 1810
that Davy proved its elementary nature and gave it the
name Chlorine, by which it is now known. Chlorine should
not be inhaled, as it irritates the sensitive lining of the throat
and nostrils.
157
158
Experimental Chemistry.
THE PREPARATION OF CHLORINE FROM HYDROCHLORIC
ACID AND MANGANESE DIOXIDE.
Perform this Experiment in the Hood.
Experiment 56. — Put a dozen small pieces of manganese dioxide
in a test tube, cover with concentrated hydrochloric acid, and heat gently.
Describe the action and enumerate the characteristic properties of the
most apparent product.
PREPARATION AND PROPERTIES OF CHLORINE.
Perform this Experiment in the Hood.
Experiment 57. — Apparatus : Construct an apparatus like that
52. The flask A holds 500 cc. and stands on a sand
bath ; the connecting tube a is about 7
cm. long and the ends of E and B are
about 5 cm. apart; C is a wooden block
with a hole in the center large enough to
admit B, which reaches to the bottom of
the bottle D. To remove D, hold C in
one hand and with the other pull D aside,
bending at the same time the whole de-
livery tube at the point a.
Supplies: When the apparatus is ar-
ranged, provide 30 gm. of granulated
manganese dioxide, six bottles, a little
powdered antimony, a little powdered
iron, bundle of fine brass wire, a taper, a
strip of calico, paper with writing in black
and in red ink, litmus paper, short glass
tube drawn out to a fine point, deflagrating
spoon, turpentine.
Fig. 52. — Apparatus for the
preparation of chlorine gas.
Caution — - Turpentine is inflammalde.
Process: Put the manganese dioxide in the flask, insert the stopper,
and slowly pour through the safety tube enough concentrated hydro-
chloric acid to cover the manganese dioxide. Heat gently and finally
stand the burner under the sand bath so that the flame just touches the
Chlorine and Hydrochloric Acid.
59
iron pan. Chlorine is rapidly evolved and passes into D, which should
be removed when full (as seen by the color) and covered with a glass
plate. Collect six bottles and proceed at once with the following
experiments : —
{a) Sprinkle a pinch or two of powdered antimony into a bottle of
chlorine. Heat some powdered iron in a deflagrating spoon and pour
it into another bottle. Into a third bottle thrust a loose bundle of very
fine brass wire which has been slightly heated. Describe the action,
giving especially in each case the evidence of chemical action and of
new products.
(^) Into a bottle of dry chlorine put a piece of calico, litmus paper
(both colors), and paper containing writing in black and in red ink.
Allow the whole to remain undisturbed for a few minutes and then
describe the change, if any. Add several drops of
water, and describe the change. Draw^ a general con-
clusion from the whole experiment.
{c) Hold a burning taper in a bottle of chlorine
long enough to observe the result. Draw a conclusion
and verify it thus : Attach the pointed glass tube to
a rubber tube connected with the gas cock, light the
gas and lower a s>nall jet into a bottle of chlorine.
Does the result verify the above conclusion ? The
same result may be shown as follows : Fold a- strip
of filter paper (about lo cm. wid^) into the shape
shown in Fig. 53; cautiously heat ^ about 10 cc. of
turpentine in a large test tube ; saturate the paper with
the hot turpentine and drop it into a bottle of chlorine.
Describe the result. When the action is over, exam-
ine the paper, and draw a conclusion regarding the
action between hot turpentine and chlorine.
Wax (in the taper), illuminating gas, and turpentine are mainly
compounds of hydrogen and carbon.. Explain the result in {c).
BLEACHING BY CHLORINEo
Experiment 58. — Put a little bleaching powder into a test tube
and add enough water to make a thin paste. Add a few drops of dilute
Fig. 53.— Fluted
paper fo r
Exp. 57 (c).
1 Hold the test tube with the holder. Remember that turpentine ignites
easily. If the turpentine catches fire, press a damp towel over it.
i6o Experimental Chemistry.
sulphuric acid to the paste, and then dip a strip of bright colored calico
into the mixture. Remove the calico and wash it with water. Describe
the change.
If time permits, repeat with a piece of unbleached cotton cloth and
describe the change.
A liter of chlorine under standard conditions weighs
3.18 grams.
CLASS-ROOM EXERCISE. XX.
1 . Bleaching and bleaching powder.
{a) What is bleaching powder ? How is it made ? For what
kind of bleaching is it commonly used ?
(J)) The phenomenon in Exp. 57 {b) is due to what general
chemical change ? Will chlorine bleach a lead pencil mark
or newspaper ink ? Why ? Predict the result when red
hot charcoal is thrust into a bottle of chlorine.
{c) Why was the turpentine heated in Exp. 57 (<:)?
{d) What is the per cent of "available chlorine" in bleaching
powder?
2. Additional study of chlorine.
{a) Summarize the physical properties of chlorine. What is its
most striking property? Is it heavier or lighter than air?
How do you know? The method in Exp. 57 is called
" collecting a gas by dozvnward displacement." Why so
called?
{b^ What class of chemical compounds is formed in Exp. 57 {a) ?
Name others in this class already used or studied. Could
chlorine be collected over mercury? Why?
{c) What does Exp. 57, as a whole, show regarding the chemical
activity of chlorine? In what respects does its activity
resemble and differ from the activity of oxygen? Is
chlorine ever found uncombined? Why? Does it form
many compounds ? Are they in general stable or unstable ?
(^) For what chemical element has chlorine a strong attraction?
What experiment shows this?
{e) How may chlorine be quickly distinguished from all other
gases previously studied?
(/) The industrial preparation of chlorine. (^) Uses.
Chlorine and Hydrochloric Acid. i6i
3. The equation for the interaction of hydrochloric acid and man-
ganese dioxide is —
4HCI + MnO^ = Clo + MnCl. + 2H2O
146 + 87 =: 71 + 126 + 36-
Name all the products. How much chlorine can be made from 247 gm.
of manganese dioxide ?
4. A bottle of chlorine water was exposed to the sunlight until all
the chlorine disappeared.
(a) What two products were formed?
{b) Write the equation for the reaction.
{c) What weight of chlorine is necessary to form 20 gm. of the
gaseous product?
{d) What volume of chlorine is necessary to form 20 gm. of
the other product?
5. If dry steam and chlorine are passed through a red-hot porcelain
tube, two gases are formed.
{a) Name the gases.
{b) Write the equation for the reaction.
(^) How may the gases be separated?
6. Historical.
{a) The discovery of the elementary nature of chlorine.
(J)) Essential facts in the life of Davy ?
{c) A brief account of Faraday's contributions to chemistry.
7. Liquid chlorine.
{a) Faraday's work.
(J?) Present method of liquefaction.
{c) Transportation.
{d) Uses.
8. Nascent state.
{a) Definition.
{b) Illustration.
(^) Theoretical value of the conception.
{d) Chlorine is sometimes prepared by heating a mixture of
sodium chloride, sulphuric acid, and manganese dioxide.
Interpret the reaction from the standpoint of the nascent
state. Write the equation for the reaction.
1 62 Experimental Chemistry.
9. Miscellaneous.
(a) Atomic weight of manganese.
(^) Molecular weight of manganese dioxide, and of manganese
chloride.
(<:) Symbol of antimony, carbon, and manganese.
Hydrochloric Acid. — It was shown in Exp. 19 that
hydrochloric acid contains hydrogen, and in Exp. 56 that
it contains chlorine. These results should be verified by a
synthesis of the acid. It will first be necessary, however,
to know the characteristic properties of hydrochloric acid,
so that the product of the synthesis may be readily recog-
nized. It should be remembered that hydrochloric acid,
as we have used it, is a water solution of the gas, and that
the same term, hydrochloric acid, includes both the gas
and its solution.
CHARACTERISTIC PROPERTIES OF HYDROCHLORIC ACID GAS.
Experiment 59. — Heat a little concentrated hydrochloric acid in
a test tube, and observe the white fumes and pungent odor. The gas
is so soluble in water that it draws the moisture from the air, forming
thereby dense white fumes. Try the action of the gas on moist-
ened litmus paper and on a burning match. Moisten a glass rod with
ammonium hydroxide, hold it near the mouth of the test tube, and note
the striking result. This is a confirmatory test for hydrochloric acid
gas, and should always be used to verify its presence. Sunnnarize the
characteristic properties of hydrochloric acid gas.
SYNTHESIS OF HYDROCHLORIC ACID GAS.
Remember that a mixtiire of hydrogen and air expiodes
if broiigJit 7iear aflame.
Experiment 60. — Construct a hydrogen generator (see Exp. 19) and
attach about 15 cm. of rubber tubing to the delivery tube. Before
generating the hydrogen, collect a bottle of chlorine from a generator
If heat is necessary, clamp A \.o tx. stand and
Chlorine and Hydrochloric Acid. 163
h
iengthen B (see Fig. 82). Generate a rather rapid current of hydrogen
and let it run about three minutes ; attach the platinum tip (see App. A,
§ 13) to the rubber tube, and let the gas run two
minutes longer; then light <-he hydrogen and lower
the flame well into the bottle of chlorine. Let the
action continue long enough to observe the color
and size of the flame, and the color and quantity of
the product;, tiien extinguish the flame by pinching
the rubber tube. Cover the bottle with a glass plate
and. after the action in the hydrogen generator has
been stopped or the generator removed to a safe
place, test the contents of the bottle as the gas was \^
tested in Exp. 59. Do the results agree in every ^^^- 54- — Simple
particular ? If not, repeat until the disagreement is fworine^^ras °^
removed or explained.
Summarize the evidence of the qualitative composition of hydro-
chloric acid gas.
It was shown in Exp. 57 (<^) that chlorine unites easily
with metals, forming chlorides, just as oxygen forms oxides.
There are many compounds of chlorine and metals, the
most common being sodium chloride, which would have
been formed in Exp. 57 (<^) if sodium, instead of iron and
antimony, had been burned in chlorine. It is expensive
and inconvenient to prepare large quantities of hydrochloric
acid by a synthesis, so the chlorides are allowed to interact
with a compound which yields hydrogen easily. Sulphuric
acid and sodium chloride are usually used.
PREPARATION AND PROPERTIES OF HYDROCHLORIC ACID.
Experiment 61. — Construct and arrange an apparatus like Fig. 52
to collect the gas. The apparatus shown in Fig. 55 is used to prepare
a solution of the gas, and is the same as shown in Fig. 52, except that
B passes into the absorption apparatus, /^G. Instead of /% the appa-
ratus //, shown in Fig. 56. may be used. The flask A stands on a sand
bath, and the gas is collected in D by downward displacement, exactly
as in the collection of chlorine.
164
Experimental Chemistry.
I. To 15 cc. of cold water add slowly and with constant stirring
40 gm. of concentrated sulphuric ackl. While the mixture is cooling
put 40 gm. of sodium chloride in the flask A, arrange the apparatus to
collect the gas, and then pour the
cold acid through the safety tube
upon the salt. Heat gently with
a low flame. Collect four bottles,
as in the preparation of chlorine,
cover with glass plates, and set
aside until needed. Disconnect
Fig. 55. — Apparatus for preparing
hydrochloric acid.
Fig. 56. — Apparatus //, which may
replace F in Fig. 55.
at a, and replace D by FG (or HG). The flask F (or H) contains
about 50 cc. of water, and the tube reaches just to the surface of the
water, but in G the tube dips into the water. Pass the gas into the
absorption apparatus, until the mixture in A seems exhausted, or for
a reasonable time.
II. Meanwhile proceed to study the hydrochloric acid gas already
collected.
{a) Invert a bottle, and hold its mouth beneath the water in a pneur
matic trough. Describe the result. What property of the gas does this
result illustrate ? How would a rise of temperature aff"ect the result ?
Compare this property with that of oxygen, hydrogen, nitrogen, and
air under similar conditions. Apply the litmus test to the liquid in
the bottle, or to the liquid beneath the mouth of the bottle.
{b) Insert a burning stick into another bottle of the gas. Describe
the result. How do oxygen, hydrogen, nitrogen, chlorine, carbon diox-
ide, and air act under similar conditions ?
Chlorine and Hydrochloric Acid. 165
(c) Hold a piece of wet filter paper near the mouth of another bottle
of the gas. Describe the result. What is the cause ?
(d) Drop in the remaining bottle of hydrochloric acid gas a piece
of filter paper saturated with ammonium hydroxide. The dense white
fumes are ammonium chloride, and are formed by direct combination
of ammonia gas and hydrochloric acid gas. The equation is —
HCl + NH3 == NH.Cl
Ammonia gas Ammonium
Chloride
Determine the solubility of the ammonium chloride in water.
(e) Enumerate other properties of hydrochloric acid gas which you
have observed.
III. Disconnect J^G and continue to heat the flask (in the hood) until
no more gas is evolved. Save the contents of the flask for Exp. 62.
Meanwhile study the general properties of the solution of hydrochloric
acid gas in /^according to Exp. 48. Record the results.
IV. Every acid has properties peculiar to itself. To a solution of
silver nitrate add a few drops of nitric acid and then a little of the
hydrochloric acid made in /^. A white, curdy precipitate, called silver
chloride, is formed. Expose a little of this precipitate to the sunlight.
Determine its solubility by adding separate small portions to cold
water, hot water, and warm ammonium hydroxide. Tabulate the prop-
erties of silver chloride. This whole experiment (IV.), solubility
included, constitutes the test for hydrochloric acid. The same test
answers for chlorides. Try it. (Compare Exp. 29 (<^).)
It has been shown that when zinc and sulphuric acid
interact (i) a new compound is formed, and (2) a definite
weight relation exists between. the zinc and hydrogen. It
is possible, though not so convenient, to prove that similar
relations exist between the factors and products when
sodium chloride and sulphuric acid interact. Hydro-
chloric acid gas is not so easily collected as hydrogen ;
furthermore, the chemical action between sodium chloride
and sulphuric acid is not always complete. However,
(66 Experimental Chemistry.
extensive work by chemists has shown that only one sub-
stance besides hydrochloric acid is formed, if the chemical
action is complete. It is only necessary, therefore, in a
qualitative test, to confine our examination to this substance
in a study of
THE INTERACTION OF SODIUM CHLORIDE AND SULPHURIC
ACID.
Experiment 62. — I. Add enough warm water to the flask A to
loosen the sohd, and pour a portion of the contents into a casserole or
evaporating dish. Evaporate to complete dryness slowly over a piece
of wire gauze and then heat strongly (in the hood) until all fumes of
sulphuric and hydrochloric acids have been driven off.
II. {a) Dissolve a little of the residue from I. in distilled water.
Test part of the solution for a chloride and the remainder for a sul-
phate. Draw a general conclusion from the result.
{b) Test another portion thus : Moisten the platinum test wire
(see App. A, § 14) with water, dip the loop into the solid, and hold it
in the Bunsen flame. Note the intense color. Clean the wire with
hot water or sandpaper, and repeat with any compound known to
contain sodium, or which you have made from sodium. (Compare
Exp. 42 (<f).) What does this result prove qualitatively about the
specific nature of the substance under examination ? Give the name
of the substance. State the chemical changes which led to its
forifiation.
COMPOSITION OF HYDROCHLORIC ACID GAS.
This experiment, including the preparatio:: ^f the sodium
amalgam, should be done in the hood.
Experiment 63. — Method: The composition ot hydrochloric acid
gas may be found by decomposing a known volume of the gas with
sodium amalgam. Hydrogen is the only volatile product, and is meas-
ured directly in the tube used in the experiment. The mercury in the
amalgam takes no part in the reaction.
Supplies: The apparatus shown in Fig. 57; a tightly fitting cork
for the tube £", tall jar, crucible (Hessian), 100 gm. of mercury, 5 gm.
Chlorine and Hydrochloric Acid. 167
of sodium, stiff iron wire, iron pan, glass-stoppered bottle. In Fig. 57,
A is a 500 cc. flask, ^ is a small dropping funnel, (7 is a U-tube to be
filled with fragments of glass drenched with concentrated sulphuric
Fig. 57. — Apparatus for determining the composition of hydrochloric acid gas.
acid, Z? is a delivery tube reaching to the bottom' of die tube E. The
tube E is graduated and its total capacity should be known.
I. Preparation of sodiitni amalgam: Put 100 gm. of mercury in a
rather large Hessian or iron crucible, slip the crucible into a ring, clamped
to an iron stand, so that about two-thirds of the crucible is below the
ring. The crucible should have an iron, porcelain, or sand cover. Cut
about 5 gm. of sodium into pieces the size of a bean and slide all
the sodium at once from a piece of paper into the mercury. Cover the
crucible immediately and stand a lighted burner under it. Hold the
rubber tubing between the thumb and forefinger and stand aside from
the apparatus. As soon as the heat is sufficient, the reaction takes
place suddenly, often accompanied by a flash of light. The instant a
flash or noise occurs, pinch the tube to extinguish the burner flame,
and stand back until the fumes of mercury disappear. Extinguish the
gas, remove the cover, and stir the amalgam vigorously with a stiff
iron wire. Before the amalgam cools, pour it into a shallow iron pan.
1 68 Experimental Chemistry.
Allow it to solidify, then break it quickly into small pieces and put it
into a glass-stoppered bottle. Preserve carefully, as sodium amalgam
is decomposed by moisture.
II. Process: {a) Arrange the apparatus so that the hydrochloric
acid gas will pass from the generator through the drying apparatus to
the bottom of the graduated tube, as shown in Fig. 57. Generate the
gas by allowing concentrated sulphuric acid to drop slowly from the
funnel into concentrated hydrochloric acid.
(/;) Fill the tube E completely with dry gas free from air. Allow the
generator to run. moderately fast, and occasionally test the escaping gas
at the open end of the tube with moist blue litmus paper. As the opera-
tion proceeds, gradually draw the graduated tube away from the delivery
tube, keeping it as nearly vertical as possible. Do not touch the tube
with the hands, but use a test tube holder. Have ready about 10 gm.
of powdered sodium amalgam, wrapped in a small piece of thin paper,
and when the tube is full of gas, withdraw it from the delivery tube,
thrust the sodium amalgam into the tube and instantly insert the stop-
per. Invert the tube repeatedly to allow all the gas to come in contact
with the sodium amalgam. As the action proceeds, a white deposit
coats the inside of the tube and the sodium amalgam slowly liquefies.
When the action seems complete, uncork the tube under water. If the
preceding part has been correctly performed, the water will rush into
the tube. Clamp the tube so that the water is at the same level within
and without, but not so that the tube is over the stream of hydrogen
arising from the excess of amalgam at the bottom of the jar. After
it has stood five minutes, adjust the levels (if necessary) without
touching the tube with the hand, and read the volume. What is
the relation between the original and the final volume?
{c) Unclamp the tube, place the thumb over the open end, lift it from
the water, invert, remove the thumb, and instantly hold a lighted
match at the mouth. What gas was in the tube?
(^) Summarize the evidence of the composition of hydrochloric
acid gas. Sketch the apparatus.
Discussion of Experiment 63. — The result of Exp. 63
shows approximately that a simple relation exists between
the volumes of hydrogen and chlorine derived from hydro-
chloric acid gas. It has been proved by many experiment-
Chlorine and Hydrochloric Acid. 169
ers that these volumes are equal. This result has been
confirmed by decomposing hydrochloric acid by an electric
current, just as water was decomposed ; furthermore it has
been shown that when equal volumes of hydrogen and
chlorine are allowed to combine, hydrochloric acid is
formed without a residue of either component. When
the gases so combine, the equation for the reaction is : —
H2 + Cl2= 2HCI.
This simple relation, together with the equally simple
relation revealed by the volumetric composition of steam,
will be subsequently explained.
The Formula of Hydrochloric Acid Gas is HCl. — No de-
rivative of hydrochloric acid is known which contains less
hydrogen or less chlorine in a molecule. Hence HCl is
the simplest expression of the facts. Again, the molecular
weight of hydrochloric acid gas is 36.5, since its density is
18.25.
CLASS-ROOM EXERCISE. XXI.
I. Theoretical.
The equation for the complete interaction of sodium chloride and
sulphuric acid is —
2 NaCl + H.SO, = Na,SO^ + 2 HCl
Sodium •j \
Sulphate
117 + 98 = 142 + 73
(«) Why is this equation correct ?
{d) Why does it contain 2 NaCl instead of NaCl ?
(c) How would you prove this equation ?
n. Further study of hydrochloric acid.
(a) How made industrially ? (^) Its usual impurities.
(c) Cause of its yellow color.
(d) Commercial name, and why so called?
lyo Experimental Chemistry.
(e) Explain the statement, "The coefficient of absorption ol
hydrochloric acid gas at o° C. is 503.'"
(/) What happens when a solution of hydrochloric acid is boiled?
(g) The liquefaction of hydrochloric acid gas.
3. Historical.
(a) The history of hydrochloric acid up to Priestley's time.
(d) Priestley's work. (c) Davy's work.
(<'/) The various names of chlorine and hydrochloric acid, and
their significance.
4. Miscellaneous.
(a) Name and describe any compounds of chlorine not already
considered.
I
PROBLEMS. XIV.
1. How many grams of chlorine can be obtained from the complete
action of hydrochloric acid on one kilogram of manganese dioxide ?
How many liters of chlorine ?
2. How much chlorine (by weight) will 3.6 1. of water absorb, pro-
vided it absorbs twice its volume ?
3. A manufacturer needs 10 1. of commercial hydrochloric acid,
sp. gr. 1.174. How much sodium chloride and sulphuric acid are
necessary ?
4. What volume of chlorine would be obtained at 13° C. and 740 mm.
by heating 10 gm. of manganese dioxide with an excess of hydrochloric
acid ?
5. How much manganese dioxide will yield a liter of chlorine ?
6. A liter of commercial hydrochloric acid, sp. gr. 1.16, is required.
How m.'jch sodium chloride, sulphuric acid, and waiter (in the receiver)
must be used ?
7. How many liters of hydrogen gas can be obtained from i 1. of
hydrochloric acid gas ?
8. Assume (i) 100 parts of potassium chlorate yield 60.85 parts of
potassium chloride, (2) 100 parts of potassium chloride yield 192.4 parts
of silver chloride, (3) the formulae KCIO.,, KCl, and AgCl, (4) O = 16,
and (5) 100 parts of silver combine with 32.87 parts of chlorine. From
these assumptions, find the atomic weights of chlorine, potassium^ and
silver.
Chlorine and Hydrochloric Acid. 171
9. A chemist found that 91.462 gm. of silver, when heated in a cur-
rent of chlorine, yielded 121.4993 gm. of silver chloride. What is the
atomic weight of chlorine ? (Assume Ag = 108.09.)
10. If a volume of dry chlorine measares 22.32 1. at o C. and 760
mm., what will it measure at 17° C. and 730 mm.?
1 1 . The simplest form of the equation for the interaction of chlorine
and water is —
H,0 + CI, = 2 HCl + O.
(a) How many liters of hydrochloric acid gas can be made
from 875 gm. of water ?
(d) How many liters of oxygen will be produced ?
{c) How many liters of chlorine will be required ?
12. If 20 I. of hydrogen are mixed with 10 I. of chlorine, (a) how
many liters of hydrochloric acid gas are produced, and (d) which gas
and how much remains in excess ?
13. Hydrogen and chlorine were mixed and exploded, and the prod-
uct weighed 73 gm.
(a) What weight of each gas was used?
(I?) What volume of each gas was used?
14. What volume of hydrogen is necessary to unite with one gram
of chlorine ?
15. Calculate the percentage composition of —
(a) Hydrochloric acid gas.
(^) Sodium chloride.
(^) Sodium sulphate.
(c/) Silver chloride.
(t') Potassium chloride.
(/) Potassium chlorate.
(g) Acid sodium sulphate, HNaSO^.
16. How much sodium chloride is necessary to make a kilogram of
hydrochloric acid gas ?
17. How many pounds of sodiu;n chloride are necessary to make 150
lb. of hydrochloric acid solution containing 31.5 per cent of available
acid ?
18. How many grams of manganese dioxide are necessary to make
100 gm. of chlorine from hydrochloric acid?
19. How many grams of manganese dioxide are necessary to make
40 1. of chlorine at 37*^ C?
CHAPTER XL
COMPOUNDS OF NITROGEN.
Ammonia. — The word aviuiojiia is loosely used to include
both the gas and its solution, just as hydrochloric acid
means either the gas or its solution. The solution of
ammonia gas is, more strictly, ammonium hydroxide, as
will appear from subsequent experiments.
Ammonia is best prepared from its compounds.
PREPARATION OF AMMONIA.
Experiment 64. — {a) Mix and grind together in a mortar a little
ammonium chloride and lime (calcium oxide). What product is re-
vealed by the odor?
(Jj) Dissolve a little ammonium chloride in water, add a little sodium
hydroxide solution, and warm gently. What is the most ajDparent
product?
PREPARATION AND PROPERTIES OF AMMONIA GAS AND
AMMONIUM HYDROXIDE.
Experiment 65. — Method: Prepare the gas and the solution sub-
stantially as in the case of hydrochloric acid.
Apparatus : Prepare the gas in a modified form of the apparatus
shown in Fig. 52. The gas is collected by upward displacement in the
apparatus shown in Fig. 58. It is an inverted form of the '' B-C-D part "
of Fig. 52; 79 is connected with E by «, as before. The solution —
ammonium hydroxide — is prepared in the apparatus shown in Fig. 55.
This experiment must be performed without interruption, and all parts
of the apparatus must be constructed and ready for immediate use
before the experiment proper begins.
172
Compounds of Nitrogen.
173
Supplies : 30 gm. of ammonium chloride, 30 gm. of quicklime which
has been exposed to the air long enough to make it crumble easily, 5
bottles, 4 glass plates, litmus paper, stick or splint of wood, pneumatic
trough or dish filled (as usual) with cold water,
filter paper.
Process: I. Mix thoroughly the ammonium
[HO=i
t
:c
Fig. 58. — Apparatus for
collecting ammonia gas.
chloride and lime. Slip the mixture into the
flask A^ insert the stopper and its tubes, and
add a little water through the safety tube to
moisten the mass. Connect the collection ap-
paratus, and see that all joints are tight. Heat
A gently with a low flame. Collect four bottles
of gas (in dry bottles), removing each when
full, and covering with a glass plate. It can
easily be told when the bottle is full by hold-
ing a piece of moist red litmus paper near the
mouth. Do not smell at the mouth of the bottle.
Disconnect at a, attach the absorption ap-
paratus FG^ and allow the ammonia gas to be
absorbed by the water in both F and G, while
the collected gas is being studied, remembering
to shake F occasionally.
II. The general properties of ammonia gas .are found by the follow-
ing experiments : —
{a) Try the action of a bottle of the gas on moist litmus paper ; try
the action on a burning stick or taper (having a small flame) ; notice,
also, the color, and very cautiously the odor by wafting the gas gently
toward the nose. Is it heavier or lighter than air, chlorine, and hydro-
chloric acid gas? Has it acid or alkaline properties? Will it burn or
support combustion ?
{p) Hold a bottle with its mouth just below the surface of the water
in a pneumatic trough or other vessel. Watch the result. Describe it.
W^hat property of the gas does it reveal? Is it a marked property?
Test the liquid in (or just below) the bottle with litmus paper (both
colors). What has become of the gas?
{c) Pour a few^ drops of concentrated hydrochloric acid into an
empty, warm, dry bottle. Cover it with a glass plate, invert it, and stand
it upon a covered bottle of ammonia gas. Remove both plates at once,
and hold the bottles together by grasping them firmly about their necks.
1*74 Experimental Chemistry.
Describe the action. Give all the evidence of chemical action. If heaV
was evolved, state where and why. Allow the white product to settle,
then scrape out as much as possible upon a paper, transfer it to a test
tube, and warm gently with a little sodium hydroxide solution. What
is formed? What was the white product? (If the last question cannot
be answered from the evidence here presented, consult Exp. 6i {d)
and Exp. 64 {b).)
III. Disconnect the absorption apparatus, and remove the generator
to a hood. Study the properties of the ammonium hydroxide in F as
follows : —
{a) Determine the general properties by repeating Exp. 49 («)
and {b).
{b) Boil a little of the liquid gently in a test tube. What gas is
evolved ?
{c) Try the effect of ammonium hydroxide on a grease spot. What
practical application is made of this property ?
The equation for the interaction of ammonium chloride
and lime — calcium oxide — is —
2NH4CI + CaO - 2NH3 + CaCl2+H20
Ammonium Calcium _ Ammonia Calcium ^^.
Chloride Oxide Gas Chloride ^
107 + 56 = 34 + III + 18
Composition of Ammonia Gas. — Ammonia gas is a com-
pound of two elements. A qualitative and a quantitative
examination of the gas will reveal its components and
their relations.
QUALITATIVE EXAMINATION OF THE COMPOSITION OF
AMMONIA GAS.
Experiment (id. — Method: Pass dry ammonia gas over heated
magnesium powder and test the two products.
Apparatus: The apparatus is shown in Fig. 59. A is a 500 cc.
flask, provided with a safety gauge tube and a delivery tube connected
Compounds of Nitrogen.
175
directly with a U-tube B. The left-hand limb of B is to be empty, the
bend loosely plugged with cotton, and the other limb filled with lumps
of quicklime. CC is a glass tube about 25 cm. long and 2 cm. in diam-
eter; it is loosely filled with soda lime, except at the ends, where the
soda lime is replaced by loose plugs of cotton. Before the final ad-
justment, this tube should be tapped along its entire length to insure
an unobstructed channel for the ammonia gas. These tubes, B and
CC, serve to dry the ammonia gas — a precaution absolutely essen-
tial to the success of the experiment. If the ammonia gas is evolved
slowly, these tubes are sufficient ; but if haste is necessary, then a
calcium chloride drying jar, filled with lime, should be placed between
B and CC . The tube CC is attached to B and DD' , as shown in the
ffl=H]
»f
Fig. 59. — Apparatus for determining the qualitative composition of ammonia gas.
figure. DD' is a combustion tube from 12 to 15 cm. long. Two cop-
per wires at D and D'^ arranged as usual, serve to prevent accidental
fracture of the tube or melting of the stoppers. A delivery tube,
partly shown as £", passes from the combustion tube into a glass dish,
arranged, as usual, to collect a gas over water. The apparatus is sup-
ported by iron stands provided with clamps which may be attached at
convenient points, such as C and-Z?'. Various modifications may be
made to economize space, but such changes should not shorten the
drying tubes. The flask A stands on a sand bath or asbestos board,
and the ammonia gas is evolved by heating concentrated ammonium
hydroxide over a low flame, which is controlled by a Hofmann screw
attached to the gas tube (lengthensd, if necessary) near the operator.
The combustion tube is heated with a wing-top burner. Two or three
empty bottles will also be needed.
1-76 Experimental Chemistry.
Process: Fill A half full of concentrated ammonium hydroxide and
connect securely with the tubes B and CC . Put enough magnesium
powder in DD' to make a thin layer from 5 to 7 cm. long. (See App. A,
§ II.) Connect DD' with the drying and delivery tubes. Fill the
glass dish with enough dilute sulphuric acid to submerge the end of
the (lengthened) delivery tube E. Fill the bottles (preferably of 125 cc.
capacity) with cold water, and stand them conveniently near the dish,
so that they may be instantly inverted, if necessary. Heat the flask A
gently, and begin to heat the combustion tube. Search for leaks, which
may be detected by the odor, as soon as any ammonia gas is known to
have been evolved. Watch three points — (i) the safety gauge tube,
which will reveal any excessive back pressure ; (2) the magnesium,
which will change in color as the experiment proceeds ; and (3) the end
of the delivery tube, where the gas which is to be collected rises, and
where the escape of the ammonia gas indicates, approximately, the rate
of evolution of the latter gas. When it is judged that the apparatus is
free from air, heat the combustion tube throughout its entire length,
and, finally, heat the magnesium strongly. A slow current of ammonia
gas should be kept passing all the time. As the temperature rises, the
magnesium changes to a greenish yellow color, and a gas is evolved
which does not dissolve in water. When this color appears, slip a
bottle over the end of the delivery tube and collect the gas. Continue
to heat the magnesium until the evolution of this new gas slackens.
Collect three bottles of the gas evolved and set them aside temporarily.
Cool the combustion tube slowly, and, when cool enough to handle,
disconnect the combustion tube from the delivery tube and the drying
apparatus, and quickly pour the contents of the combustion tube into a
small dry bottle provided with a glass stopper. Preserve for subse-
quent examination.
Test each bottle of gas with a lighted match. What is the gas?
What is its source? What, then, is one component of ammonia gas?
Test the greenish powder formed in the combustion tube thus :
Pour a little into a test tube half full of water. What gas is evolved?
Repeat, if the observation is not specific. Record this observation, and
reserve it for subsequent explanation.
This experiment reveals one component, of ammonia
gas, and the product of the interaction of the greenish
yellow powder and water suggests the other.
Compounds of Nitrogen.
177
The other component of ammonia gas may be found as
"the residual gas" by
THE QUANTITATIVE EXAMINATION OF THE COMPOSITION OF
AMMONIA GAS.
Experiment 67. — Method: Decompose ammonia gas in a measured
volume of chlorine. The method depends mainly upon the fact, al-
ready shown, that chlorine will withdraw hydrogen from compounds of
that element.
Apparatus: In addition to the special apparatus shown in Fig. 60,
there is needed a chlorine generator, arranged as shown in Fig. 52, ex-
cept that the " B-C-D part " is replaced by a tube shaped
like E in Fig. 30, and long enough to reach to the bottom A/
of a pneumatic trough ; a tall jar filled with water at the
temperature of the room ; a little (about 25-30 cc.) con-
centrated ammonium hydroxide, and dilute sulphuric acid.
The apparatus shown in Fig. 60 consists of the graduated
tube A, of known capacity ; it is provided with a gas-
tight, one-hole rubber stopper, carrying the small drop-
ping funnel B, which is never removed from the stopper.
The latter may be made by sealing a piece of tubing
(5 cm. long and 2 cm. in diameter) to a stop-cock tube
near the stop-cock ; the other end is then cut off at any
desired point. It is desirable, though not necessary,
that the tube A be graduated. The dropping funnel
shown in Fig. 32 may be used instead of B, but it
is not so satisfactory.
Process : Generate the chlorine in the hood as directed
in Exp. 57. Meanwhile ascertain the total capacity of the
tube A (see App. A, § 22). Then fill it with water, at
the temperature of the room, invert, and clamp it in the
pneumatic trough so that the open end is over the hole in
the shelf. By this time the chlorine generator should be
free from air ; this condition, however, can be told only by experience,
but since it is absolutely essential to an accurate result, the generator must,
in case of doubt, be allowed to run at least fifteen minutes after the time
the chlorine bubbles from the end of the delivery tube. When the
chlorine is believed to be free from air, slip the end of the delivery tube
Fig. 60. — Ap-
paratus for
determining
the quanti-
tative com-
position of
ammonia
gas.
lyS Experimental Chemistry.
under the shelf, and pass a rapid stVeam of gas up into the tube ; the
tube will soon be filled, since the gas will not dissolve in the water to
any appreciable extent. When the tube is full, let it stand a few min-
utes until the water runs down ; be sure the tube is completely full
before removing it from the trough.
Meanwhile, fill the dropping funnel half full of concentrated ammo-
nium hydroxide, open the stop-cock, and allow the liquid to fill the lower
tube. Close the stop-cock and stand the funnel in a beaker or bottle
in another hood, or in some convenient place not far from Ihe chlorine
apparatus. Unclamp the tube, cover the open end with the thumb or
second finger, invert, and quickly insert the stopper with its funnel.
Push in the stopper gas-tight. Fill the funnel with concentrated am-
monium hydroxide, cautiously open the stop-cock, and allow a drop or
two of the ammonium hydroxide to enter the tube. Considerable heat,
and sometimes a faint flash of light, together with dense, white fumes,
indicate chemical action. Add ammonium hydroxide, drop by drop,
until the reaction seems complete. Take care not to let any gases out of
or any air into the tube when the stop-cock is opened. Now add dilute
sulphuric acid through the funnel until at least one-third of the tube is
full. Fill the funnel with water, cover with a small piece of filter paper,
and invert the whole apparatus in a tall jar of water. Remove the paper,
open the stop-cock, allow the water to run in, and let the whole appara-
tus remain undisturbed for at least fifteen minutes. Then, without
touching the upper part of the tube, raise or lower it until the water is
at the same height within and without ; clamp the tube in this position
and read the volume of "the residual gas." What proportion of the
original volume is it?
Close the stop-cock, lift the tube from the water, remove the stopper,
lower a lighted match attached to a wire into the tube, and observe the
result. Test the gas also with lime water, as previously directed.
What is the gas? Why? Why not any of the other gases studied?
Record the two definite results of this experiment.
Sketch the apparatus used for the quantitative examina':-cn of the
composition of ammonia gas.
Discussion of Experiments 66 and 67. — When ammonium
hydroxide, which has been shown to be identical with a
solution of ammonia gas, is added to chlorine, the chlorine
withdraws the hydrogen from the ammonia gas, forming
Compounds of Nitrogen. 179
hydrochloric acid gas and ''the residual gas." But the hy-
drochloric acid gas instantly combines with the ammonia
gas to form ammonium chloride, which is seen at first as the
dense white fumes, but which finally dissolves in the water
added. These changes continue until all the chlorine is
removed. The excess of ammonia present, both as gas
and in solution, is neutraHzed with sulphuric acid, and the
solid product — ammonium sulphate — also dissolves in
the water. Hence at the end of the experiment only "the
residual gas" is present as agaSy and the volume and nature
may be found as directed. It was shown in Exp. 63 that
chlorine combines with its own volume of hydrogen. There-
fore there must have been three times as much hydrogen
as nitrogen in the ammonia gas, since there was three
times as much chlorine at the beginning as there was nitro-
gen at the end of the experiment.
Since it has been shown that ammonia is a compound
of hydrogen and nitrogen, the yellowish green compound
formed in Exp. 66 must be a compound of magnesium and
nitrogen. It is, in fact, magnesium nitride, Mg3N2. When
added to water the interaction yields ammonia gas as one
product, thus confirming the evidence of the composition
of ammonia gas furnished by Exp. 6^.
Equations for the Chemical Action in Experiments 66
and 67. — The following equations represent in the simplest
form the chemical action in these experiments. In Exp.
66-^
2NH3 + 3Mg = MggN^ + 3H2
Magnesium
Nitride
MgsNa + 3H2O - 2NH3 + 3MgO
t
1 80 Experimental Chemistry.
In Exp. 6^ —
NH3
+
3 CI
= 3 HCl
HCl
+
NH3
= NH.Cl
H2SO4
#
+
2NH8
= (NH.XSO,
Ammonium
Sulphate
+ N
The Volumetric Composition of Ammonia Gas. — Addi-
tional evidence of the composition of ammonia gas is fur-
nished by a determination of
THE VOLUMETRIC COMPOSITION OF AMMONIA GAS.
Experiment 68. — Method: Pass electric sparks through a meas-
ured volume of dry ammonia gas and measure the final volume.
Apparatus : Eudiometer, mercury, mercury trough, ammonia gas
generator and drying apparatus, induction coil, battery, meter stick,
barometer, and thermometer.
Process : Fill the eudiometer with clean, dry mercury, invert it in a
trough of mercury and clamp it in a perpendicular position. Generate
ammonia gas and dry it as in Exp. 66. Pass the current through the
drying apparatus long enough to expel all the air, and then introduce
about 20 cc. of the dry ammonia gas into the eudiometer. Arrange a
thermometer so that the bulb dips into the mercury, and allow the
whole apparatus to stand undisturbed for at least fifteen minutes. Then,
without touching that part of the tube which contains the gas, read and
record the following : —
Volume of the gas = V = cc.
Height of mercury in the tube = 1/ = mm.
Thermometer = t =
Barometer = P' = mm.
Connect the eudiometer with the induction coil and pass sparks
through the gas for about half an hour. Meanwhile reduce the
observed gas volume to standard conditions by the formula —
^,_ V'(P'-N)
760(1 +(.00366 X ^))
Compounds of Nitrogen. i8i
The factor H must be subtracted from the barometer reading because
the gas is supporting only that column of mercury which is the dif-
ference between the barometer column and the column extending from
the surface of the mercury in the trough to the upper surface of the
mercury in the tube — found by actual measurement with a meter stick.
Read the volume again roughly, noting also the barometer and
thermometer. Continue to pass sparks through the gas. Read again
soon, and if there is any marked change in volume, continue to pass
the sparks. The sparking must continue until the volume is constant.
When there is no further change in volume, disconnect the wires, allow
the whole apparatus to stand undisturbed for ten or fifteen minutes, then
read accurately as before. Reduce the observed volume to standard
conditions. What is the relation between the original and final
volumes ?
Discussion of Experiment 68. — The result of this experi-
ment simply shows that when ammonia gas is subjected to
the action of electric sparks the volume is doubled. Further
consideration is necessary to determine the nature of the
chemical change, which may be best explained by an
illustration. A volume of ammonia gas measured lOO cc,
and after sparking, the residual volume was 200 cc. Then
90 cc. of dry oxygen were added, the mixture was exploded,
and 65 cc. of gas remained. It is known, of course, from
previous experiments, that ammonia gas yields by decom-
position the two gases, hydrogen and nitrogen. Hence
the total volume of hydrogen and oxygen which dis-
appeared as a result of the explosion must have been
225 cc. {i.e. 290-65). But two-thirds of this volume must
have been hydrogen, and the remainder oxygen, since
these gases unite in the proportion of two to one by
volume to form water. Therefore, 150 cc. (f of 225 cc.)
of hydrogen must have come from the decomposed am-
monia gas ; and since only 200 cc. of mixed gases were
formed by the sparking, the remaining 50 cc. must be the
nitrogen. These two gases — hydrogen and nitrogen —
i82 Experimental Chemistry.
therefore, arc combined in ammonia gas in the proportion
of three parts to one by vohmie.
The facts revealed by Exp. 68 are most simply expressed
by the equation : —
N2 + 3H2 = 2NH3 (17)
One volume Three volumes _ Two volumes of
of nitrogen of hydrogen " ammonia gas
A liter of ammonia gas under standard conditions xveighs
77 gi"i^-
Formula of Ammonia Gas. — The gravimetric composi-
tion of ammonia gas reveals the fact that fourteen parts
of nitrogen combine with three parts of hydrogen. The
vapor density of ammonia gas is 8.5 (approximately on the
hydrogen standard), and hence its molecular weight is 17.
The simplest formula which meets all these requirements
is NH3.
Law of Gay-Lussac. — A comparison of the volumetric
composition of water vapor (steam), hydrochloric acid gas,
and ammonia gas (see pages loi and 169) reveals the
fact that in each case the volume relations can be
expressed by whole numbers. Thus —
Two volumes of hydrogen '
unite with" ■ to form two volumes of water vapor.
one volume of oxygen
One volume of hydrogen 1 ^ ^ . , ^ ^ i 1 1 ;„
. ; '^ to form two volumes or uydrochlonc
unites with r ■ 1 <r -
one volume of chlorine J ' ^'
Three volumes of hydrogen 1
unite with to form two volumes of ammonia gas.
one volume of nitrogen J
The simple relation between the volumes of the com-
ponents, and between the components and their compounds
in the three cases examined, is just as simple in all cases.
Compounds of Nitrogen. 183
The law was first pointed out by Gay-Lussac, who stated
the relation substantially as follows : —
Gases combine in volumes zvhicJi bear a simple relation to
each other and to the product.
The law may be expanded thus — '' When two or more
elementary gases unite to form a gaseous compound, the
volume of a fixed quantity of the compound stands to the
volumes of the combining elements in a ratio which can be
expressed by whole numbers."
Ammonium Compounds. — It has been shown in several
experiments (Exp. 64 {b) and Exp. 65 {c)) that ammonium
chloride is formed by direct combination of ammonia gas
and hydrochloric acid gas. This compound may be formed
in the same way as any other salt. ^
INTERACTION OF AMMONIUM HYDROXIDE AND
HYDROCHLORIC ACID.
Experiment 69. — Put about 50 cc. of ammonium
hydroxide in an evaporating dish or casserole, and add
dilute hydrochloric acid, drop by drop, until the solution
is just neutral or faintly alkaline. A convenient drop-
ping apparatus is shown in Fig. 61. Evaporate to dry-
ness on a water bath or very slowly over a piece of
wire gauze. Test the residue as follows : —
{a) Is it an acid, base, or salt ?
{b) Warm a little with sodium hydroxide soaition.
What is formed? Draw a conclusion as to the nature of
the residue.
{c) Support the dish on a piece cf wire gauze and
warm gently until a decided change occurs. Describe Fig.6i.— Con-
the result. What compound do the fumes suggest ? venkmt droj)-
{d) Verify the observations and conclusions by re- {"^"1 pointed
peating {b) and {c) with ammonium chloride from the ^ube open ai
laboratory bottle. both ends.
V
184 Experimental Chemistry.
{e) Draw a conclusion from these observations regarding the main
product of the interaction of ammonium hydroxide and hydrochloric
acid.
The compound formed in Exp. 69 was ammonium chlo-
ride, and it belongs to a series of compounds known as the
Ammonium compounds. If sulphuric acid or nitric acid
had been used in Exp. 69, then ammonium sulphate or
ammonium nitrate would have been formed. The
ammonium series of compounds is strictly analogous to
the sodium and other metallic series, i.e. ammonium and
sodium enter into similar chemical reactions. Thus, these
two series of compounds are well known : —
NaCl NH4CI
NaOH NH4OH
NaNOg NH4NO3
Na2S04 (NH4)2S04
NagPO^ (NH4)3P04
Ammonium, however, is not an element, like sodium,
but a compound; and not a compound, strictly speaking,
but a group of atoms — one atom of nitrogen and four
atoms of hydrogen — so related to each other and to other
atoms, that they act as one atom of a metal. Such a
group of atoms, which is incapable of independent ex-
istence, is called a radical. The chemical relations of
radicals are exactly the same as elements. Ammonium
has never been isolated, or if it has, it is so unstable that
it immediately decomposes into ammonia and hydrogen.
It is believed that ammonium hydroxide is a solution of
an extremely unstable compound. The other compounds
of ammonium mentioned above are well defined.
The symbol of ammonium is NH4, and its valence is one.
The formulas of its compounds have already been given.
Compounds of Nitrogen.
CLASS-ROOM EXERCISE. XXIL
Further study of ammonia gas.
(a) Source of atmospheric ammonia gas.
(d) Specific gravity.
(c) Its relation to the subject of combustion.
(d) Sokibility in water.
(e) Quantitative test for ammonia gas.
Liquid ammonia.
(a) Faraday's work.
(d) Present method.
(r) Properties.
(d) Use in manufacture of ice.
(e) Formula. Why ?
Historical.
(a) Early knowledge of ammonia gas and its compounds.
(d) Discovery of the gas by Priestley.
Ammonium hydroxide.
(a) Exact meaning of the name as usually used.
(d) Source of commercial substance.
{c) Specific gravity and strength of commercial substance.
(d) What is the difference between liquid ammonia and am-
monium hydroxide ?
(^) Uses for domestic purposes, in medicine, and in agriculture.
(/) What is anhydrous ammonia ?
Names.
(^) Why is ammonia gas sometimes called alkaline air, volatile
alkali ?
(b) Why should ammonium chloride be called sal-ammoniac,
muriate of ammonia, sal-animoniaciun, chloride of am-
monia ?
{c) What different meanings may the term ammonia have ?
(^) Ammonium carbonate is sometimes called sal-volatile. Why ?
{e) What is liqnor ammoniac ?
if) What is ammoniacal liquor ?
(^) What is gas liquor ?
(/i) Is there any difference between lime, quicklime, and calcium
oxide ? If so, what is it ?
1 86 Experimental Chemistry.
(f) What is (igua aruvionia ? Ammonia water ? Spirits ot
hartshorn ?
6. Ammonium sahs.
{a) Give the formula, method of preparation, properties, and uses
of ammonium chloride, ammonium nitrate, and ammonium
sulphate.
7. Composition of ammonia gas.
{a) Summarize the experimental evidence of the qualitative com-
position of ammonia gas.
{b) Summarize the evidence of the volumetric composition of
ammonia gas.
{c) Why is NH3 the simplest formula of ammonia gas ?
(^) When ammonia gas is passed over red-hot copper oxide,
what are the products? What would the experiment prove
in regard to the composition of ammonia gas, if the copper
oxide and the products were weighed ?
8. Miscellaneous.
{a) What is the valence of nitrogen in ammonia gas? In
ammonium? In ammonium hydroxide?
(J)) Calculate the atomic weight of nitrogen from data already
given.
(<:) Why is the formation of white fumes from the combination
of ammonia gas and another gas not necessarily evidence of
ammonium chloride?
9. Gay-Lussac's Law.
{a) Exact statement.
(^) Illustrations.
{c) Theoretical significance.
(^) Essential facts in the life of Gay-Lussac.
{e) Humboldt's contribution to the law.
PROBLEMS. XV.
1. How many grams of ammonia gas can be obtained from 2140
gm. of ammonium chloride by heating with lime? (See Equation 16.)
2. What volume of ammonia gas can be obtained by heating 31.47
gm. of ammonium chloride with lime?
Compounds of* Nitrogen. 187
3. How many grams of ammonium chloride must be heated with
lime to make 80 1. of ammonia gas ?
4. Water absorbs about 50 per cent of its weight of ammonia gas
at the ordinary temperature and pressure. How much ammonium
chloride and calcium oxide are needed to produce 10 kg. of ammonium
hydroxide ?
5. If 22.4 1. of hydrogen at o" C. and 760 mm. weigh 2 gm., what
will an equal volume of ammonia gas weigh at 100' C. and 380 mm. ?
6. If 250 cc. of chlorine are measured at 10'' C. and 750 mm. and
then allowed to interact with ammonia gas, what volume of nitrogen
will be liberated at the same temperature and pressure ?
7. How many cubic centimeters of the component gases can be
obtained by the decomposition of a liter of ammonia gas ?
8. A volume of dry ammonia gas equal to 29 cc. is decomposed by
electric sparks, and the resulting mixture is exploded after 30 cc. of
oxygen have been added. What gases remain, and what are their
volumes ?
9. 15 cc. of ammonia gas are decomposed by electricity, and 40 cc.
of oxygen are added; the mixture is then exploded. What are the
gases and their volumes before and after the explosion ?
10. What weight of ammonia gas and of chlorine is necessary to
produce a liter of nitrogen ?
11. A kilogram of ammonium chloride is heated with lime. What
volume of ammonia gas at 12° C. is produced ?
12. 180 cc. of dry ammonia gas are decomposed by electric sparks.
(a) What is the volume of each product ?
{d) If 130 cc. of oxygen are added and the mixture exploded,
what gases and what volumes remain ?
13. 100 cc. of dry ammonia gas are decomposed by electric sparks;
then 100 cc. of oxygen are added and the mixture exploded.
(a) What gases and what volumes remain after the decomposi-
tion of the ammonia gas ?
(If) After the explosion ?
14. What is the weight of 30 1. of dry ammonia gas at the normal
temperature and pressure ?
15. What volume of nitrogen and of hydrogen can be obtained by
passing electric sparks through one hter of dry ammonia gas, and what
volume of oxygen will be necessary to unite with all the hydrogen thus
obtained to form water ?
1 88 Experimental Chemistry.
16. Calculate the percentage composition of —
{a) Ammonium chloride, NH^Cl.
(J?) Ammonium hydroxide, NH^OH.
{c) Ammonium sulphate, (NH4)^S04.
{d) Ammonium nitrate, NH4NO3.
(^) Magnesium nitride, MgoN^.
17. Calculate the formula of a substance having the percentage
composition : —
{a) Nitrogen = 82.35
Hydrogen = 17-64
(J?) Nitrogen = 26.17
Hydrogen = 7.48 1
Chlorine = 66.35
18. How many pints of a solution of ammonium hydroxide (sp. gr.
0.88) weigh 4 lb. ?
Nitric Acid. — Another useful compound of nitrogen is
nitric acid.
PREPARATION OF NITRIC ACID.
Precaution — Do not get concentrated sulpJmric or nitric
acids on the skin or clotJiing.
Experiment 70. — Prepare the acid in the apparatus shown in
Fig. 62. A\%2i tubulated retort supported by a clamp and resting on a
sand bath. The neck of the retort passes into the receiver B^ which
catches the nitric acid as it distils from the retort. Condensation of
the acid fumes is facilitated by wrapping wet filter paper around the
neck of the retort where it enters the receiver, and upon the flask itself,
if it becomes hot. The neck is clamped just tightly enough to hold the
retort in position. Turn B occasionally to expose a fresh surface to
the fumes ; it is advisable, though not always necessary, to place a block
of wood against the bottom of the receiver to keep it in the desired
position.
Slip about 35 gm. of sodium nitrate from a paper into the retort
(see Fig. loi, App. A, § 12), and pour upon it, through a funnel, stand-
ing in the tubulure, about 25 cc. of concentrated sulphuric acid. Adjust
the apparatus as shown in Fig. 62. Heat gently, and nitric acid will
pass into the receiver Distil at as low a temperature as possible, as
Compounds of Nitrogen.
189
long as any nitric acid runs down the neck of the retort. Pour the
nitric acid into a small bottle for use in the next experiment. Allow the
retort to cool, add just enough warm water to loosen the solid mass,
pour the contents into any convenient vessel, and preserve for Exp. y^-
Fig. 62.— Apparatus for preparing nitric acid.
The general properties of nitric acid have been shown
in a preceding experiment (see Exp. 48). They should
be recalled or the work repeated.
A SPECIAL PROPERTY OF NITRIC ACID.
Experiment 71. — Add twice its volume of water to the nitric acid
made in Exp. 70, and proceed as follows : —
{a) Boil a piece of a quill toothpick in a portion of this diluted nitric
acid. How is the quill changed at first ? What is the effect of contin-
ued heating ? Pour off the acid, and wash the quill with water. Is the
color permanent ?
i^o Experimental Chemistry.
(J?) Add a dozen or more drops of nitric acid to a dilute solution of
indigo. Describe the change. Will ammonium hydroxide restore the
original color ? Is the change temporary or permanent ? What, in all
probability, is the general character of the change — combination or
decomposition ?
Draw a general conclusion from the whole experiment regarding the
action of nitric acid on organic matter, which is typified by the quill and
indigo.
LABORATORY EXERCISE. VIII.
1. Examine a bottle of pure nitric acid and of the commercial acid,
and record —
{a) Color of each acid.
(b) Presence or absence of a gas above the liquid, and color of
gas, if any.
2. Examine a bottle of nitric acid which has been exposed to the
sunlight. What does this observation show about the stability of nitric
acid ?
3. Remove the stopper from a bottle of strong nitric acid. Is it
very hygroscopic ? What other acid acts similarly ?
4. Smell of the acid cautiously, and record the result.
A TEST FOR NITRIC ACID.
Experiment 72. — To a test tube one-fourth full of water add five
or six drops of concentrated nitric acid ; add also an equal volume of
concentrated sulphuric acid. Shake until the acids are well mixed, then
cool by holding the test tube in nmning water. Make a cold, dilute
solution of fresh ferrous sulphate and pour this solution carefully down
the side of test tube upon the nitric acid mixture. Where the two solu-
tions meet a brown or black layer will appear, consisting of a compound
formed by the interaction of the nitric acid and the ferrous sulphate.
It is an unstable compound and will often decompose, if the test tube is
shaken. Record the observation.
This test is also used for a nitrate. Try it with a solution of sodium
nitrate. Record the result.
The Interaction of Sodium Nitrate and Sulphuric Acid. —
It has already been shown that nitric acid is one product
of the interaction of sodium nitrate and sulphuric acid, and
Compounds of Nitrogen. 191
since chemists have shown that if tJie action is complete
only one other substance is formed, it remains to determine
its nature.
THE INTERACTION OF SODIUM NITRATE AND SULPHURIC
ACID.
Experiment 73. — Pour the solid residue obtained in Exp. 70 into
a casserole or an evaporating dish, and evaporate to dryness over a
piece of wire gauze in the hood. As the mass approaches pasty con-
sistency lessen the heat to avoid spattering. When the mass is dry,
heat strongly as long as white, choking fumes are evolved. This last
operation is done to remove all traces of sulphuric acid, and to complete
the chemical change. Allow the dish to cool gradually, and when cool,
dissolve some of the white solid in distilled water and test separate
portions for a sulphate and nitrate. Which is it ? Test another por-
tion for sodium, as previously directed. WMiat is the name of the white
substance ?
Draw a general conclusion regarding the chemical action which occurs
in the preparation of nitric acid by the interaction of sulphuric acid and
sodium nitrate.
There are two equations expressing the chemical changes
revealed by the preceding study of the preparation of
nitric acid. If the temperature is low, then the reaction is
represented thus : —
NaNOg + H2SO4 = HNO3 -f HNaSO^
Sodium Acid Sodium
Nitrate Sulphate (18)
85 H- 98 =- 63 + 120
But if the temperature is high, then the reaction is
represented thus : —
2 NaNOg + H.SO^ = 2 HNO3 + Na2S04
170 -f 98 = 126 + 142
192 Experimental Chemistry.
Composition of Nitric Acid. — Numerous facts pointing to
the quaUtative composition of nitric acid are found by
simple experiments.
SYNTHESIS OF NITRIC ACID.
Experiment 74. — Apparatus: The apparatus consists ot a small
bottle provided with a two-hole rubber stopper ; a platinum wire passes
through each hole and is held firmly in place by a short piece of glass
rod ; each wire terminates in a narrow strip of platinum foil. There
will be needed, also, a Ruhmkorff coil and a battery, or a static machine,
and the usual connections for the battery.
Process : Put in the bottle a few drops of neutral water and a piece
of blue litmus paper. Adjust the platinum terminals so that they are near
the bottom of the bottle and from 3 mm. to 6 mm. apart. Connect the
platinum wires with the coil, and pass sparks through the apparatus for
about twenty minutes, or until there is definite evidence of a chemical
change revealed by the litmus paper.
What is the nature of the change ? Independent evidence has shown
that the compound formed under these conditions is nitric acid. Since
only air and water are present, what elements combine to form nitric
acid ?
ANALYSIS OF NITRIC ACID.
Experiment 75. — (a) Generate hydrogen from zinc and dilute
sulphuric acid in a large test tube or small flask, and when the evolu-
tion is brisk drop dilute nitric acid slowly into the vessel. The nitric
acid may be dropped in from a burette or pipette, or from a glass tube
(15 cm. long) shaped like Fig. 61. A brown gas may appear at inter-
vals. If so, it indicates a secondary and undesirable action. The diffi-
culty may be prevented by (i) regulating the evolution of hydrogen so
that it is continuous but not too rapid, and (2) by adding the nitric
acid slowly. The essential point of the experiment is to allow the nitric
acid to interact with the hydrogen, but not to be decomposed by inter-
action with the zinc. Add about 10 cc. of nitric acid, then pour off
the liquid from any excess of zinc, and evaporate the liquid to dryness ;
heat at first over the free flame in a casserole or in a dish which stands
on a piece of wire gauze, and finally heat over a low flame which is kept
conr,tantly moving to prevent spattering.
Compounds of Nitrogen. 193
Put a portion of the residue in a test tube and warm gently with
sodium hydroxide sokition. A thick, white precipitate will be formed,
hence the tube must be shaken vigorously and the heat applied with
care. A gas will be liberated. Test it by determining the odor and
by holding a piece of moist red litmus paper in the escaping vapor —
not against the side of the test tube which may have been covered with
the sodium hydroxide. What is the gas ? What is its source ? What
evidence does it furnish regarding the elements contained '*^ nitric acid .'*
(^) Perform this experi)neni in the hood.
Boil a few grams of coarsely powdered ferrous sulphide in a test tube
with a litde concentrated nitric acid. Allow the actLn to proceed a few
minutes, and after the tube is cool, dilute with two or three times its
volume of distilled water. Filter, and test the filtrate for a sulphate.
Remembering the essential diflference between a sulphide and a sul-
phate, explain the general chemical change which has taken place.
What evidence of the qualitative composition of nitric acid is con-
tributed by this experiment ?
CLASS-ROOM EXERCISE. XXIII.
1. Historical.
(a) The alchemists' knowledge of nitric acid.
{b) Lavoisier's contribution to our knowledge of the composition
of nitric acid.
{c) Cavendish's determination of the exact composition of nitric
acid.
{d^ Derivation and significance of the term aqua fortis.
2. Commercial nitric acid.
{a) Preparation.
(Jj) Specific gravity.
(c) Impurities.
\d) Uses.
3. Composition of nitric acid.
{a) Summarize the experimental evidence of the composition of
nitric acid.
{b) If nitric acid gas is passed over heated copper, two of the
three products are water and nitrogen. Of what elemen+s
must the nitric acid have been composed ?
194 Experimental Chemistry.
{c) If nitric acid is allowed to trickle through a red-hot porce-
lain tube, a gas is formed which vigorously supports
combustion. What gas (probably) is it, and what does
its formation show about the composition of nitric acid ?
About the stability of nitric acid ?
(d) If a mixture of dry ammonia gas and air is passed over
heated platinum sponge {i.e. over a mass of porous plati-
num), nitric acid is formed. What does this fact prove
about the composition of nitric acid ? (The platinum takes
no chemical part in the reaction.)
(f; Salts of nitric acid are often found in the water which falls
at the beginning of a thunder-storm. Why ?
4. Miscellaneous.
((2) Why are there no acid nitrates ?
(^) What is the valence of nitrogen in hitric acid ?
(c) Explain and illustrate the statement " nitric acid is a power-
ful oxidizing agent."
PROBLEMS. XVI.
1. {a) How much sulphuric acid is needed for the complete trans-
formation of 606 gm. of potassium nitrate into nitric acid ?
{b) How much sulphuric acid, if the reaction takes place at a low
temperature ?
2. Five hundred grams of potassium nitrate are transformed into
nitric acid by heating with sulphuric acid at a low temperature. Cal-
culate —
(a) The amount of nitric acid produced.
(^) The amount of sulphuric acid required.
(c) The amount of acid potassium sulphate formed.
3. What is the least quantity of sulphuric acid w^hich can be used to
decompose 500 gm. of potassium nitrate in the preparation of nitric acid?
4. How many pounds of nitric acid can be obtained by distilling
400 lb. of sodium nitrate with sulphuric acid ?
5. How many grams of nitric acid can be obtained by heating a
kilogram of sodium nitrate with sulphuric acid at a low temperature ?
6. How much {a) sodium nitrate and {b) potassium nitrate are
necessary to produce the maximum weight of nitric acid by heating
each with 140 kg. of 97 per cent sulphuric acid ?
Compounds of Nitrogen. 195
7. Calculate the percentage composition of —
(a) Nitric acid.
(d) Potassium nitrate.
(c) Sodium nitrate.
(^) Acid sodium sulphate.
(e) Acid potassium sulphate.
The Interaction of Nitric Acid and Metals. — The interac-
tion of nitric acid and metals is complicated. It depends
upon (i) the particular metal, (2) the strength of the acid,
(3) the temperature, and (4) the presence of salts resulting
from the chemical action.
THE GENERAL CHARACTER OF THE INTERACTION OF NITRIC
ACID AND METALS.
Perform this experiment in the hood.
Experiment 76. — Stand four test tubes in the test tube rack and
slip into each a few small pieces of one of the following metals : zinc,
copper, tin, and iron. Add to each test tube in succession just enough
concentrated nitric acid to cover the metal. Observe the changes in
each case, particularly (i) the vigor of the action, (2) the nature and
properties of all the products, especially color and solubility, and (3)
evidence of presence or absence of hydrogen.
Tabulate these observations.
The interaction of nitric acid and copper typifies the
chemical action in many cases and demands additional
examination.
THE INTERACTION OF NITRIC ACID AND COPPER.
Perform this experiment in the hood.
Experiment 77. — I. Construct an apparatus like Fig. 30 and
arrange it to collect a gas over water as in the preparation of hydrogen
(Exp. 19). Put from 10 to 15 gm. of copper borings in the flask, insert
196 Experimental Chemistry.
the stopper tightly, adjust the delivery tube, fill three bottles with water,
and invert them in the trough. Pour just enough concentrated nitric
acid through the safety tube int: the flask to cover the copper, taking
care to seal the bend of the safety tube with acid. Dense brown fumes
are evolved. If the action is too vigorous, add a little water through the
safety tube. Collect three bottles of the gas which bubbles from the
delivery tube. Cover them with glass plates and stand them aside
tempoi^rily.
Pour the contents of the flask into a casserole or evaporating dish,
and evaporate slowly to crystallization (not to dryness) on a water bath.
If thc-re is much of the original copper remaining, add more nitric acid
at intervals. The crystals, after being dried between filter paper, should
be preserved in a glass-stoppered bottle.
II. Meanwhile study the gas as follows : —
(a) Observe its general properties while covered.
(d) Uncover a bottle. Describe the result. Is the brown gas
identical with the one observed in the generator at the beginning of the
experiment ?
(c) Uncover a bottle, pour in about 25 cc. of water, cover with the
hand and shake vigorously, still keeping the bottle covered. Why has
the brown gas disappeared? Uncover the bottle for an instant, then
cover and shake again. Is the result the same ? Repeat, if the result
is not definite, or does not agree with previous observations.
(rtf) With the third bottle determine whether the two gases will burn
or support combustion. A convenient flame is a burning match fastened
to a stiff wire. Plunge it to the bottom at first and gradually raise it
into the brown gas.
Answer the following: —
(i) What is the source of the colorless gas?
(2) What is the general chemical relation of the two gases to each
other? To the air?
(3) Why is not the brown gas collected in the bottles by displace-
ment of water?
(4) Will either gas burn, or support combustion?
(5) Which gas has been observed before? In what experiment
(other than Exp. 76) ?
(6) What is the general relatioa of these gases to nitric acid?
Compounds of Nitrogen. 197
III. Study the properties of the crystals by determining: —
(a) Solubility in water (cold and hot).
(b) Action of heat.
(<:) Action of their solution upon an iron nail.
{d) Action of their solution when added to ammonium hydroxide.
Compare the observed properties with those of copper nitrate
obtained from the laboratory bottle. Are the two substances id3ntical ?
Discussion of Experiment 77. — The colorless gas is nitric
oxide, the brown gas is nitrogen peroxide, and the blue
'solid is copper nitrate. The following equation is usually
given as the simplest expression of the interaction of copper
and nitric acid : —
3Cu + 8HNO3 = 3Cu(N03)2 + 4H2O + 2NO
Copper Copper Nitrate Nitric
Oxide
When the nitric oxide is exposed to the air, it combines with
oxygen. This reaction is most simply represented by the
equation —
NO + O = NO2
Nitric Nitrogen
Oxide Peroxide
Nitric oxide is not always the only gas produced by the
interaction of nitric acid and copper, but it is most abun-
dantly produced and is loosely regarded as the only oxide of
nitrogen evolved. In the case of most metals, one or more
of the numerous oxides of nitrogen are produced. ''Concen-
trated nitric acid and zinc yield ammonia, as well as nitric
oxide, among the gaseous products, though the ammonia
usually combines at once with the excess of acid to form
ammonium nitrate. Some metals, such as tin and antimony,
form an oxide and not a nitrate. Nitric acid is such a
powerful oxidizing agent that hydrogen which is displaced
198 Experimental Chemistry.
from it by metals never appears as hydrogen, but is oxidized
to water. (See page 117.)
Nitrates. — The salts of nitric acid have one common
property, viz. extreme solubility in water, and hence many
laboratory solutions are nitrates of the metallic elements.
Nitrates behave in various ways when heated.
ACTION OF NITRATES WITH HEAT.
Jl/Xperiment 78. — I. Heat a little sodium nitrate in an ignition
tube. A high temperature is necessary to produce any extensive
chemical change. Insert a glowing match into the ignition tube. What
gas is detected .'* Dissolve the residue in water and add a few drops of
dilute sulphuric acid. Describe the result. Could the residue have
been a nitrate? Why? If the experimental evidence is not definite
enough for a final conclusion, proceed as follows : —
Heat a mixture of 10 gm. of sodium nitrate (or potassium nitrate)
and 20 gm. of lead in a sand bath pan, which stands on a tripod, or
similar support. Stir the melted mass with a stiff iron wire or blunt
glass rod. Some of the lead will disappear and a yellowish brown
powder will be seen in the molten mass. The action should proceed
until most of the lead has disappeared. Allow the mass to cool, transfer
to a mortar, pulverize, add hot water, and filter the clearer portion ; add
more hot water to the residue and filter this portion. Add to the com-
bined filtrates a few drops of concentrated sulphuric acid. Describe
the result. Compare the result with the action of concentrated sulphuric
acid on a solution of sodium nitrite. Is the white residue (produced
in the ignition tube) a nitrate? Why?
The yellowish product is lead oxide. What general chemical change
led to its formation? How must the nitrate have been changed?
Draw a general conclusion from the observations made above re-
garding the action of heat on nitrates of the alkali metals.
II. Pulverize 8 or 10 gm. of lead nitrate, and heat the powder in an
ignition tube or evaporating dish. Describe the result. How does it
differ essentially from the action of heat on the alkali nitrates?
Two gaseous products are formed, though one — the oxygen — is not
easily detected in the presence of the other. More definite results are
obtained as follows : Heat the lead nitrate in an ignition tube provided
Compounds of Nitrogen.
199
with a delivery tube passing to the bottom of a test tube half full of
sodium hydroxide solution and connected by a delivery tube with an-
other test tube, or small bottle, arranged to collect a gas over water.
The gas will be found to be oxygen. Add a few drops of concentrated
sulphuric acid to the test tube which contained the sodium hydroxide
solution. What compound is present? What chemical change pro-
duced it? The residue in the ignition tube is lead oxide. Compare
these results with those obtained in Exp. 76, III. {b).
Draw a general conclusion regarding the action of heat on the
nitrates of the heavier metals.
The exceptional behavior of ammonium nitrate when
heated places it in a class by itself. An examination of
its behavior involves also a study of the main product,
which is another oxide of nitrogen.
THE ACTION OF AMMONIUM NITRATE WITH HEAT.
Experiment 79. — The apparatus is shown in Fig. 63. The am-
monium nitrate is heated in A, which is connected with an empty bottle,
D ; the latter is pro-
vided with a delivery
tube, E, which dips into
a pneumatic trough,
filled, as previously
described, with mod-
erately warm water.
Fill ^one-third full
of crystals of ammo-
nium nitrate, adjust
the apparatus, and
heat A gently with a
low flame. , The am-
monium nitrate melts
and appears to boil.
Regulate the heat so
that the evolution of
gas will be slow. Notice the continuous formation of another product
besides the gas. Collect three bottles of the gas,//-^^ from air, covei
Fig. 63. — Apparatus for decomposing ammonium
nitrate.
200 Experimental Chemistry.
each with a glass plate, and stand them aside until needed. As soon
as the last bottle has been removed from the trough, disconnect the
apparatus at a, and stand the generator in the hood to cool.
Test the gas as follows : —
{a) Allow a jar to remain uncovered for a few seconds. How does
this oxide of nitrogen differ from nitric oxide?
{h') Thrust a glowing stick of wood into the same bottle of gas.
Describe the result. Is the gas combustible? Does it support com-
bustion?
{c) The observations in {b) suggest that the gas is oxygen, but it is
not, though this fact is not easily proved by a single experiment. Put
a small piece of sulphur in a deflagrating spoon, light it, and lower the
burning sulphur at once into another bottle of gas. If the experiment
is conducted properly, the sulphur will not burn so brightly as it would
in a bottle of oxygen.
(^) Stand the other bottle mouth downward in the pneumatic
trough, or better, in a vessel of cold water. Describe the result. If
the result is not conclusive, fill the bottle half full of water, cover with
the hand, and shake. Would this observation help distinguish the gas
from oxygen?
What in all probability is the other product of the chemical change in
this experiment? Could it have been an impurity in the ammonium
nitrate?
How would you distinguish ammonium nitrate from all other nitrates ?
How would you distinguish the third oxide of nitrogen — nitrous oxide
— from {a) the other oxides of nitrogen, {b') air, (<:) oxygen, (^) hy-
drogen, {e) nitrogen, (/") carbon dioxide?
CLASS-ROOM EXERCISE. XXIV.
1. Summarize the properties of the three oxides of nitrogen studied.
2. State briefly the effect of heat on (rt) potassium nitrate, (J>) copper
nitrate, (<:) ammonium nitrate.
3. Predict the simplest equation for the reaction in the case of
{a) sodium nitrate, (J)) copper nitrate, {c) lead nitrate, (^d) zinc
nitrate, {e) ammonium nitrate, when heated. Verify the equations.
4. Select from Exp. ']'j^ III. a test for copper.
The fact that all nitrates, except ammonium nitrate, yield
oxygen when heated furnishes —
Compounds of Nitrogen,
20I
A SPECIAL TEST FOR NITRATES.
Experiment So. — Heat a piece of charcoal in the Bunsen flame, lay
it on a board, or iron pan, and cautiously sprinkle powdered potassium
nitrate upon the hot surface. Stand back when the action begins.
Observe and describe the action, especially its violence and rapidity,
also the color of the flame, the effect on the charcoal, and any other
characteristic result.
This kind of chemical action is called deflag}'ation. What causes it?
Law of Multiple Proportions. — There are two oxides of
nitrogen besides the three already studied, and these five
oxides aptly illustrate the Law of Multiple Proportions.
This law is usually stated thus : —
A fixed iveigJit of one element so combines with different
iveights of another element that the relations betiveen these
different iveights are expressed by small whole numbers.
The composition of compounds is usually expressed in
per cent. If, however, a definite weight is adopted as a
unit for one component, and the composition is expressed
in terms of this unit, the simple integral relation existing
between the different proportions of the other element
is clearly seen. The following table illustrates the law of
multiple proportions : —
Name.
Composition in
Per Cent.
Unit
Weight.
Ratio.
. N - 0.
N.
N - 0.
Nitrous oxide
Nitric oxide
Nitrogen trioxide ....
Nitrogen peroxide . . .
Nitrogen pentoxide . . .
63.6 - 36.4
46.6 - 53.4
36.8 - 63.2
30.4 - 69.6
25.9-74.1
7
7
7
7
7
7 -4
7-8
7-12
7-16
7-20
202 Experimental Chemistr
From this table it is clear that the proportions of oxygen
in combination with a fixed weight of nitrogen are as
1:2:3:4:5. This law together with the law of definite
proportions has profoundly influenced the development of
the atomic theory of Dalton.
CLASS-ROOM EXERCISE. XXV.
1. Gunpowder.
{a) Composition.
(<^) General methods of manufacture.
{c) Kinds.
\d) Uses.
{e) Products of action.
2. Composition, manufacture, and uses of gun-cotton, nitroglycerine,
and dynamite.
3. Miscellaneous.
« {a) Define and illustrate nitrification.
{b) What are etchings and how are they made?
4. Law of Multiple Proportions.
{a) History.
{b) Exact meaning.
{c) Other illustrations.
{d) Theoretical significance.
5. What is the valence of nitrogen in HNO^? In Cu(NO ).,? In
N.O, NO, N2O3, NO2, and N2O5?
Aqua Regia is an old term still applied to a mixture of
concentrated nitric and hydrochloric acids (one volume of
the former to three of the latter).
THE CHARACTERISTIC PROPERTY OF AQUA REGIA.
Experiment 81. — Touch a small piece of gold leaf with the end of
a moist glass rod, and wash the gold leaf into a test tube by pouring a
few cubic centimeters of concentrated hydrochloric acid down the rod.
Heat gently until the acid just begins to boil. Does the gold dissolve?
Compounds of Nitrogen. 203
Wash another piece of gold leaf into another test tube v;ith concen-
trated nitric acid, and heat as before. Does the gold dissolve ?
Pour the contents of one tube into the other, and warm gently. Does
the gold dissolve? Draw a conclusion.
Answer the following : —
(i) What is the literal meaning and significance of the term aqua
regia ?
(2) What other metals does aqua 7'egia dissolve?
(3) What is the chemical action oi aqua regia on gold?
(4) Upon what property of nitric acid does the action of aqua regia
depend'
PROBLEMS. XVII.
1 . Find the simplest formula of the substances having the indicated
composition : —
{a) H = 1.58
N = 22.22
O = 76.19
{b~) 0 = 47-52
N = 13.86
K = 38.61
2. What weight of pure nitric acid would yield 100 gm. of oxygen
if completely decomposed ?
3. The specific gravity of nitric acid is 1.522. {ji) What will
100 cc. weigh? {b) What volume must be taken to weigh 100 gm.?
4. What volume is occupied by 10,000 gm. of nitric acid? (Assume
density = 1.5.)
5. One gram of gunpowder yielded 280 cc. of gas at the normal tem-
perature and pressure. At the instant of explosion the temperature
rose to 2000° C. Calculate the volume occupied at this temperature
(assuming an unchanged pressure).
chaptp:r XII.
CARBON AND ITS SIMPLEST COMPOUNDS.
Carbon is a constituent of every living thing. It forms
a vast number of compounds. It is often the essential,
and sometimes the only, constituent of the remains of
animals and vegetables.
DISTRIBUTION OF CARBON.
Experiment 82. — (a) Cover the bottom of a Hessian crucible with
a thin layer of sand. Put on the sand a small piece of wood, a small,
compact wad of cotton, a small bone, and a lump of starch. Fill the
crucible loosely with dry sand, and slip it into the ring of an iron stand.
Heat with a flame which extends just above the bottom of the crucible
until the smoking ceases (approximately a half hour)-. After the crucible
has cooled sufficiently to handle, pour the contents out upon a block of
wood or an iron pan. Examine the contents. What is the residue?
What is hereby shown about the distribution of carbon ?
(d) Heat about i gm. of sugar in an old test tube until the vapors
cease to appear. What is the most obvious product ?
(c) Close the holes at the bottom of a lighted Bunsen Lirner, and
hold a piece of crayon or glass tubing in the upper part of the flame
long enough for a thin deposit to form. Examine it, name it, and state
its source.
(d) Hold a piece of clean crayon in the flame of a candle which
stands on a block of wood, and compare the result with that in (c).
(e) Light a wax taper ana note the most obvious product of the flame.
Explain. If time permits, heat in separate test tubes, or on an iron
pan, a piece of meat, a little flour, a bit of albumen, a piece of bread.
What is the most obvious product in each case?
Draw a general conclusion regarding the distribution of carbon.
204
Carbon and Its Simplest Compounds. 205
Carbon in the form of wood charcoal and of animal
charcoal (often called bone black) possesses remarkable
properties.
DECOLORIZING ACTION OF CHARCOAL.
Experiment 83. — Fill a test tube one-fourth full of animal char-
coal (see App. A, § 11), add 10 cc. of indigo solution, shake thoroughly
for a m^'nute, and then warm gently. Filter through a wet filter paper
into a clean test tube. Compare the color of the filtrate with that of
the indigo solution. Explain the change in color.
Other organic substances besides indigo are similarly changed. Draw
a general conclusion regarding the decolorizing power of charcoal.
DEODORIZING ACTION OF CHARCOAL.
Experiment 84. — Smell of a weak solution of hydrogen sulphide
gas. Fill a test tube half full of powdered wood charcoal, add 5 cc. of
hydrogen sulphide solution, and cork securely. If the tube leaks, make
the opening gas-tight with vaseline. Shake thoroughly. After fifteen
or twenty minutes, remove the stopper and smell of the contents. Is
the odor much less oflfensive? Repeat, unless a definite result is ob-
tained. Explain the change.
The decolorizing and deodorizing power of charcoal is
largely a physical operation, and is mainly due to its
porosity.
The attraction of carbon for oxygen — already shown by
experiment and often observed in various forms of combus-
tion — permits
REDUCTION BY CARBON.
Experiment 85. — Prepare an apparatus like that shown in Fig. 23.
Fill the ignition tube half full of a mixture of copper oxide (6 parts)
and powdered wood charcoal (i part) ; introduce the mixture by the
method described in App. A, § 11. Arrange the apparatus as directed
in Exp. 12. The delivery tube iii this case, however, should dip into
a small dish of lime water. Heat the whole ignition tube gently
at first ; increase the heat gradually, and finally heat strongly that part
!2o6 Experimental Chemistry.
of the ignition tube containing the mixture. If the mixture "crawls"
up the ignition tube, tap the tube gently near the top. A gas will be
evolved. Heat until a decided change is produced in the lime water,
and then immediately remove the end of the delivery tube from the
liquid. Describe the change in the lime water. What caused the
change ? Draw a conclusion regarding chemical change which oc-
curred in the ignition tube. Examine the contents of the tube to
verify the conclusion. If the verification is not decisive, proceed as
follows : —
Fill a porcelain crucible half full of the original mixture, and cover
it with a thin layer of charcoal. Cover the crucible and stand it on a
pronged tripod, or a triangle, and heat for five or ten minutes. Let
the crucible cool, still covered, and, when cool enough to handle com-
fortably, turn out and examine the contents. Does the observation
verify the previous conclusion? If the chemical change in the crucible
had been complete, and the proportions absolutely correct, what would
the residue have been?
LABORATORY EXERCISE. IX.
1. Complete the equation —
CuO + C = CO., +
2. Define reduction in terms of Exp. 85.
3. What other element will reduce copper oxide ?
CLASS-ROOM EXERCISE. XXVI.
1. Allotropism or Allotropy.
(a) Definition.
(d) Illustration by means of the allotropic modifications of carbon
(c) Reason for existence of allotropes.
({i) Literal meaning and significance of the word allotropy.
2. Diamond.
{a) Occurrence. (r) Properties.
(Z") Artificial preparation. (^d) Famous diamonds.
3. Graphite.
(rt) Occurrence. (c) Properties.
{b) Preparation. {d) Uses.
(<?) Literal meaning of the word.
(/) What is "black lead" ? Plumbago ?
Carbon and Its Simplest Compounds. 207
4. Animal charcoal.
(a) Preparation. " (,) Average composition.
(d) Properties. (,^) Uses.
(e) To what property is its efficiency due ?
(/) What is bone black ? Why so called ?
5. Charcoal.
(a) Various kinds, and how obtained.
(d) General properties.
(c) Uses.
6. Coal.
(a) Kinds.
(^) Relation of kinds to each other.
(c) General mode of formation.
(d) What are peat and lignite ?
(e) Composition of typical coals (see App. C, Table XL).
7. Give the method of preparation (or source), properties, and use.«
of the following forms of carbon : —
(a) Lampblack. (c) Gas carbon.
(^) Coke. (^) Soot.
8. Experimental.
(a) Examine a section of coal through a microscope.
9. Miscellaneous.
(a) How would you prove the chemical identity of the three
allotropic modifications of carbon ?
(b) What is black smoke ?
{c) What industries utilize the decolorizing po\ver of animal
charcoal .?
{d) Use of charcoal as a disinfectant.
{e) What immense industry utilizes the principle illustrated
in Exp. 85 }
Carbon Dioxide is the best^ known of the simpler com-
pounds of carbon.
PREPARATION AND PROPERTIES OF CARBON DIOXIDE.
Experiment d>6.- Method: Prepare carbon dioxide from hydro-
chloric acid and calcium carbonate, and study the gas collected over
water.
2o8 Experimental Chemistry.
Apparatus: Use the same apparatus as in the preparation ol
hydrogen (see Exp. 19 and Fig. 30). Other materials needed are
lumps of marble, sand, concentrated hydrochloric acid, stick or splinter
of wood, candle fastened to a wire, lime water, five bottles.
Process: Cover the bottom of the flask with sand, add a litile water,
and carefully slip into it a dozen or more small lumps ot marble.
Arrange the apparatus to collect the gas over water, as previously
directed. Add through the safety tube just enough concentrated
hydrochloric acid to cover the marble. Reject the first portion of
the gas evolved. Collect five bottles, cover with glass plates or wet
filter paper, and stand aside till needed.
Allow the action in the flask to continue, and preserve tna contents
for subsequent examination.
I. Study the properties of carbon dioxide gas as follows : —
(«) Plunge a burning stick into one bottle. Describe the result.
\b) Lower a lighted candle into a bottle of air, and invert a bottle
of carbon dioxide over it, holding the bottles mouth to mouth.
Describe the result. What does this result show about the specific
gravity of carbon dioxide ? Devise a simple experiment to verify this
conclusion, using for that purpose one of the remaining bottles of
carbon dioxide.
{c) Pour a little lime water into a bottle of carbon dioxide, cover
with the hand, and shake vigorously. Describe and explain the result.
{d) Fill a bottle of carbon dioxide one-third full of water, cover
with the hand, and shake vigorously. Invert, still covered, in a
vessel of water. Does the result reveal any facts about the solubility
of carbon dioxide ?
II. Filter the contents of the flask into a casserole or evaporating
dish, adding a little warm water beforehand, if the contents is solid.
Evaporate to dryness in the hood over a free flame as long as much
liquid remains. As the residue approaches pasty consistency, add a
little water and continue the evaporation. If a casserole is used, move
it about rapidly to avoid spattering. If an evaporating dish is used,
stand it on a gauze-covered support and move the lighted burner under-
neath. Heat the residue until no fumes of hydrochloric acid are
evolved. Dissolve some of the residue in distilled water and test
portions for {a) a carbonate, {b) a chloride, {c) a calcium compound
(see Exp. 29). If a calcium compound is found, confirm the observa-
tion thus : —
Carbon and Its Simplest Compounds. 209
Dip a clean, moist platinum test wire into the solid residue, and hold
Jt in the Bunsen flame. If calciam is present, the flame will be colored
a yellowish red.
What is the residue? Verify the conclusion by a simple experiment.
The simplest equation for the interaction of hydrochloric
acid and calcium carbonate is : —
2HCI + CaC03 = CO2 + CaCU +
Calcium Carbon Calcium
Carbonate Dioxide Chloride
73 + 100 = 44 + III +
(20)
18
The qualitative composition of carbon dioxide may be
shown by the
SYNTHESIS OF CARBON DIOXIDE.
Exjwriment 87. — Method: Draw, simultaneously, purified air over
hot charcoal and the product of the reaction through barium hydroxide
solution or lime water.
Apparatus: The essential part of the apparatus is shown in Fig. 64.
Short pieces of sodium hydroxide are placed in the left-hand limb of
ir?
Fig. 64. — Apparatus for the synthesis
of carbon dioxide.
A and calcium chloride in the other, the two substances being separated
by a wad of cotton ; small wads of cotton are also placed just below the
stopper in each limb. This U-tube removes the carbon dioxide and
water vapor from the air. The tube BB^ is hard glass and from 15 to
20 cm. long. A spiral of copper wire is attached at the point B' . The
test tube C is connected directly with the combustion tube by a glass
2IO Experimental Chemistry.
tube passing to the bottom of the test tube ; the tube D serves as an
outlet for the excess of air and is attached directly to a filter pump or
aspirator (see Fig. 107). The apparatus may be supported by clamps
at the points B and B\
Process: Place three or four lumps of dry wood charcoal in the com-
bustion tube. Fill the absorption tube half full of barium hydroxide
solution. Connect as shown in the figure. Draw- a slow current of
air through the apparatus to detect the leaks, if any. Readjust the
connections, if necessary. Regulate the current of air so that a bubble
a second, approximately, is drawn through the apparatus. Heat the
whole combustion tube gently at first, and finally concentrate the heat
at the place where the charcoal is located. A wing-top burner is well
adapted for this operation. In a few minutes (approximately 10) a
marked change is seen in the absorption tube. Describe it. Since
the nitrogen of the air takes no part in the chemical change, explain
the change. What does the experiment prove about the composition
of carbon dioxide.
The simplest equation for the synthesis of carbon
dioxide is : —
C + O2 = CO, ^ ^
(21)
12 + 32 = 44
A liter of carbon dioxide under standard conditions weighs
1.977 gm-
Carbon dioxide is vitally connected with the subject of
combustion.
CARBON DIOXIDE AND COMBUSTION.
Experiment 88. — (^d) Exhale through a glass tube into a test tube
half full of lime water. Describe and explain the result.
{b) Lower a lighted candle into a bottle and allow it to burn for
a few minutes. Remove the candle, pour a little lime water into the
bottle, and shake vigorously. Describe and explain the result.
{c) Allow a stick of wood to burn for a short time in a bottle, and
then proceed as in {[)). Describe the result. Does it confirm the
results in {a) and {b)^
Carbon and [ts Simplest Compounds. 211
Answer the following : —
(1) What is the main product of the combustion of substances
containing carbon ?
(2) What gas was formed in Exp. 14 (c)?
(3) How does this experiment verify previous work on "burning in
the air"?
(4) How does (a) show the relation between carbon dioxide and
combustion?
CLASS-ROOM EXERCISE. XXVII.
1. Additional study of carbon dioxide.
(a) Occurrence. (d) Modes of formation.
(c) Properties not shown experimentally.
(d) Liquid carbon dioxide. (e) Solid carbon dioxide.
2. Review.
(a) Combustion.
(d) Carbonated mineral waters, (c) Respiration and decay.
3. Historical.
(a) Van Helmont's work on carbon dioxide.
(d) Black's work. (c) Lavoisier's work.
(d) Why was carbon dioxide called^^j- sylvestre a.ndyixed air ?
4. Miscellaneous.
{a) How would you distinguish carbon dioxide from all other
gases ?
{b) Summarize the most important properties of carbon dioxide.
(<:) What is the decisive test for carbon dioxide ?
{d) What happens when a burning stick is plunged into
oxygen ?
Into nitrogen ? Into hydrogen? Into carbon dioxide?
{e) What happens when lime water is poured into oxygen:
Into hydrogen? Into nitrogen? Into carbon dioxide?
5. Deduce the molecular weight and simplest formula of carbon
dioxide from the following data : —
{a) A liter of carbon dio^fide weighs 1.977 gm.
{b) Carbon dioxide contains its own volume of oxygen.
6. Dumas and Stas found that 80 parts of oxygen by weight com-
bined with 30 parts of carbon. Deduce the atomic weight of carbor4
from this fact together with the results obtained in 5 {supra).
212 Experimental Chemistry.
PROBLEMS. XVIII.
1. How many grams of calcium carbonate are necessary to produce
15 1. of carbon dioxide?
2. How much calcium carbonate and hydrochloric acid are needed
to form 132 gm. of carbon dioxide?
3. What weight of carbon burned in air will produce 11 gm. of
carbon dioxide?
4. Twelve grams of carbon were burned in the oxygen liberated
from 122.5 g"^- o^ potassium chlorate. How much carbon dioxide was
formed, and what w^as the excess of oxygen?
5. An excess of air was passed over red-hot charcoal and thereby
formed 21 1. of carbon dioxide. What volume of air was deprived of
its oxygen?
6. A piece of pure graphite weighing 7 gm. is completely burned
in oxygen? What volume of carbon dioxide is formed?
7. What volume of oxygen is necessary to burn a kilogram of
carbon ?
8. Eighteen grams of carbon are to be burned in air. How many
liters of air are needed, and how many liters of carbon dioxide will be
formed ?
9. What volume of carbon dioxide at 12° C. and 750 mm. will be
produced by the action of hydrochloric acid on 10 gm. of marble?
10. What weight of water must be decomposed to furnish sufficient
oxygen to form, with pure carbon, 44 gm. of carbon dioxide ?
11. How much oxygen by weight and by volume is required to
unite with pure carbon to form 132 gm. of carbon dioxide?
12. How much carbon by weight is there in a liter of carbon
dioxide?
13. How much carbon dioxide by weight and volume is in the air
of a room 6 m. long, 4 m. wide, and. 3 m. high, if there is one volume
of carbon dioxide in 1000 volumes of air?
Carbonic Acid. — Carbon dioxide gas is often called
carbonic acid gas, or simply carbonic acid. It is believed
that when carbon dioxide is passed into water it com-
bines with the water and forms a weak, unstable acid.
The case is analogous to the formation of the base
ammonium hydroxide by the solution of ammonia gas in
Carbon and Its Simplest Compounds. 213
water. Carbonic acid has never been isolated, but its
salts, the carbonates, are numerous and well-defined com-
pounds. The formula H^COg has been given to carbonic
acid.
CARBONIC ACID.
Experiment 89. — Construct a carbon dioxide generator like that
shown in Fig. 46. B is tilled nearly full of lumps of marble. Concen-
trated hydrochloric acid (i vol. to i vol.) is put in A. The generator
is operated as previously described. (See Exp. 44.) Attach Z^ to a
wash bottle (see Figs. 56,65, and 114) containing water or concentrated
sulphuric acid to free the gas from any hydrochloric acid carried over
mechanically; connect the wash bottle with an absorption apparatus
consisting of a 1 irge test tube or bottle ; the tube from the wash bottle
should reach to the bottom of the absorption apparatus. A simple
combined generator and wash bottle is shown in Fig. 65.
Fill the absorption apparatus nearly full of water, add a few drops of
a solution of phenolphthalein and just enough sodium hydroxide solu-
tion to color the liquid a faint pink. Allow a slow current of carbon
dioxide to bubble through the apparatus until a definite change is
produced in the absorbing liquid. Describe and explain it.
Sketch the essential part of the apparatus.
Carbonates are salts of the hypothetical carbonic acid.
Many carbonates have already been studied. Additional
experiments, however, are needed to illustrate the —
FORMATION AND PROPERTIES OF CARBONATES.
Experiment 90. — {d) Pass carbon dioxide free from acid into a
bottle full of lime water until considerable precipitate is formed. The
generator shown in Fig. 46 may l^e used, if a wash bottle containing
water or concentrated sulphuric acid is placed between the generator
and the bottle of lime water. The combined generator and wash bottle
shown in Fig. 65 also gives satisfactory results. The gas 's generated
in the large test tube, A^ and washed in B^ and any convenient delivery
tube is attached to C; the apparatus may be stood in a test tube rack.
When sufficient precipitate has formed, disconnect the generator, and
allow the precipitate to settle. Decant the supernatant liquid, add a
214
Experimental Chemistry.
few drops of hydrochloric acid to its residue, ;ind test the gas evolved
with a burning match and with a rod moistened with lime water. What
is the gas? What is the precipitate?
' Complete the equations : —
CO, + Ca(OH), = H.O +
^
:^
CaCO,
HCl
CaCl., +
{b) Pass carbon ^\ox\d& free f?'om
acid through a weak solution of so-
dium hydroxide (or potassium hydrox-
ide) for several minutes. Add a few
drops of hydrochloric acid to a por-
tion of flie liquid. Describe and
explain the result, testing the main
product of the action, if necessary.
In what essential respect does the
compound formed from the sodium
hydroxide differ from the one formed
from calcium hydroxide?
{c) Heat a litde powdered magnesium carbonate in a test tube or
ignition tube, and test (as in {a)) the gas evolved. Explain the result
Complete the equation : —
MgCOg = MgO +
Magnesium
Carbonate
Fig. 65. — Apparatus for generatin
and washing carbon dioxide.
LABORATORY EXERCISE. X.
1. What is the decisive test for a carbonate?
2. In what two ways may carbonates be decomposed?
3. How are carbonates formed?
4. Complete the equations : —
{a) KOH + CO, = KXO, +
(b) K.fO,
HCl = CO.,
5. How may lime water be easily distinguished from solutions of
sodium and potassium hydroxide?
Carbonic acid is dibasic, hence, besides ordinary calcium
carbonate, there is
Carbon and Its Simplest Compounds. 215
ACID CALCIUM CARBONATE.
Experiment 91. — Pass carbon dioxide free from acid into lime
water until the precipitate disappears. Filter, if the liquid is not per-
fectly clear, and then heat. Describe the change. Why should the
precipitate disappear before heating ? Why reappear after heating ?
(If the latter question cannot be answered from evidence already given,
consult a text-book.)
LABORATORY EXERCISE. XI.
1. If Na2C03 is the formula of sodium carbonate, what is the formula
of acid sodium carbonate ?
2. What is the formula of acid calcium carbonate?
DETERMINATION OF CARBON DIOXIDE IN A CARBONATE.
Experiment 92. — Method: Decompose a known weight of a car-
bonate (preferably calcite, which is a pure variety of calcium carbonate)
and calculate the loss in weight as carbon dioxide.
Apparatus: A 125 cc. Erlenmeyer flask is provided
with a one-hole rubber stopper fitted with a small bulb
trap as show in Fig. 66. The trap is about 45 mm.
high, and is made by sealing a short tube into a bulb
about 25 mm. in diameter. It is filled with fused cal-
cium chloride held in place by absorbent cotton. Its
object is to prevent the loss of water vapor which might
be carried off mechanically by the escaping carbon diox-
ide. The whole apparatus should not w^igh more than
60 gm. before adding the acid
Process: Clean and dry the flask. Fill the bulb trap
as previously directed. Put in the flask 25 cc. of dilute
hydrochloric acid (one volume of acid to two volumes
of water). Slip into the flask a lump of calcite (crystal-
lized calcium carbonate) weighing about one gram —
its exact weight need not be known. Insert the stopper and allow
the action to proceed without interruption. The object of this pre-
liminary operation is (i) to fill the apparatus with carbon dioxide,
since it is to be weighed finally in that condition, and (2) to allow
the liquid to absorb carbon dioxide, so that only a very small
amount will be retained in subsequent operations. Meanwhile weigh
Fig. 66. — Ap-
paratus for
determining
the propor-
tion of car-
bon dioxide
in a car-
bonate.
2i6 Experimental Chemistry.
exactly to a centigram from 1.3 to 1.5 gm. of calcite. When the
evokition of carbon dioxide has ceased, open the flask for an instant,
close it, and weigh to a centigram. Slip in the weighed piece of calcite
and insert the stopper. Allow the action to proceed until no inore gas
is evolved, then open the flask for an instant, close it, and weigh again.
Record the results as follows : —
Grams.
Weight of apparatus
Weight of calcite
Total
Final weight of apparatus
Weight of carbon dioxide lost
Per cent of carbon dioxide in the calcite . . . .
Summary.
Per Cent of Carbon Dioxide in Calcite.
Found.
Theory.
Class Average.
2.
3-
Average.
•
Several determinations may be made with the same apparatus with-
out replacing the acid, if time permits a continuous operation. While
the calcite is being decomposed, another piece may be weighed ready
for use as soon as the ''Final weight"' is found, which then becomes,
of course, Ihe "Weight of apparatus" in the second determination.
Carbon Monoxide is another compound of carbon and
oxygen. It may be produced by the
ACTION OF HOT CHARCOAL ON CARBON DIOXIDE.
T/ie product of this experiment is a poisonous gas and
should not be allowed to escape into the air.
Carbon and Its Simplest Compounds. 217
Experiment 93. — MdJiod Pass carbon dioxide over hot charcoal,
collect over water the gaseous product, and test it.
Apparatus: The essential part of the apparatus is shown in Fig. 64.
A carbon dioxide generator (see Fig. 46) is connected with the U-tube A,
which is to be filled with calcium chloride. The combustion tube BB' is
the same as that used in Exp. 87. The test tube C contains a solution
of sodium hydroxide to absorb the excess of carbon dioxide. A deliv-
ery tube is connected with D and passes into a pneumatic trough
arranged to collect a gas over water. The apparatus may be supported
by clamps placed at B and B' or any other convenient points. It is
advisable, though not absolutely necessary, to add a little sodium
hydroxide solution to the water in the trough. Four or five bottles and
several lumps of charcoal are also needed.
Process: Put three or four lumps of charcoal in the combus ion tube
and connect the diiferent parts of the apparatus. Regulate the flow of
carbon dioxide so that the rate is (approximately) a bubble a second.
Stop the leaks, if any are detected. Heat the whole combustion tube
at first, and finally concentrate the heat where the charcoal is located.
Collect all the gas evolved, but reject the first two bottles, as they con-
tain air (and possibly carbon dioxide). Collect two or three more bot-
tles, cover with glass plates, and set them aside temporarily. Stop the
generator, remove the delivery tube from the trough, and gradually cool
the combustion tube. Test the gas thus : —
{a) Notice that it is colorless.
{b) Hold a lighted match at the mouth of a bottle for an instant.
Note the flame, especially its color and how it burns. After the flame
has disappeared, drop a lighted match into the bottle. Describe the
resu'". Draw a conclusion and verify it by {c).
(c) Burn another bottle of gas, and after the flame has disappeared,
pour a little lime water into the bottle and shake. Does the result
verify the conclusion made in (d) ? If not, repeat with another bottle
of gas.
LABORATORY EXERCISE. XII.
1. Summarize the observed properties of carbon monoxide.
2. What is the chemical li^-^tion of carbon monoxide to carbon
dioxide ?
3. How can each be changed into the other? What two general
processes do the changes illustrate?
21 8 Experimental Chemistry.
The simplest equation for the action of carbon dioxide
on hot charcoal is
CO2 + C = 2CO
Carbon , .
f 22)
Monoxide
44 + 12 = 56
A liter of carbon monoxide under standard conditions
weighs 1.25 gm.
If the properties of carbon monoxide cannot be conven-
iently studied by Exp. 93, they may be studied from the
gas prepared by the usual method, as follows : —
PREPARATION AND PROPERTIES OF CARBON MONOXIDE.
{Optional >j
Experiment 94. — Construct an apparatus similar to that used for
the preparation of chlorine (see Fig. 52). The flask should be smaller,
though this modification is not necessary, and the tube U should pass to
the bottom of a wash bottle fitted with a two-hole rubber stopper. The
wash bottle is filled two-thirds full of sodium hydroxide solution. A
delivery tube passes from the wash bottle into a pneumatic trough
arranged to collect a gas over water.
Put 10 gm. of crystallized oxalic acid {oxalic acid is poisonous) in the
flask, insert the stopper with its tubes, see that all joints are tight, and
then pour 30 cc of concentrated sulphuric acid through the safety tube
into the flask. Heat the flask gently, and carbon monoxide will be
evolved. A small flame must be used, because the gas is rapidly
evolved as the heat increases. It is advisable to remove or lower the
flame as bubbles appear in the flask — re<^ulate the lieat by the efferves-
cence. Collect all the gas, but do not use the first two bottlc.i, covering
the bottles with glass plates as tliey are filled, and setting them aside
temporarily. Disconnect the generator at a, and stand it in the hood
to cool. Test the gas as directed in Exp. 93.
Add a little acid to the wash bottle. What gas besides carbon
monoxide was produced ?
Sketch the essential part of the apparatus.
Carbon and Its Simplest Compounds. 219
CLASS-ROOM EXERCISE. XXVIII.
1. Additional study of carbon monoxide.
(a) Various methods of formation.
(d) Properties besides those revealed by experiment.
(c) Its reducing power.
{d) What is water gas ?
(e) Danger from carbon monoxide.
2. Combustion and the oxides of carbon.
(a) Which oxide is formed in an excess of air ? Which in an
excess of carbon ?
(d) State and explain the various chemical changes which occur
from :he entrance of oxygen (in the air) below the grate
of a red-hot coal fire to the end of the burning of the
carbon monoxide at the top of the coal.
3. Miscellaneous.
(a) How may a mixture of the oxides of carbon be separated
into its constituents ?
(^) What is the valence of carbon in carbon monoxide and in
carbon dioxide ? Is there any satisfactory explanation of
this fact ?
(0 Illustrate the Law of Multiple Proportions by the two oxides
of carbon.
(d) Carbon dioxide is often called carbonic anhydride. Why ?
What experiment proves this ?
(e) Calculate the molecular weight of calcium carbonate from
the result obtained in Exp. 92. Compare with the theo-
retical weight.
(/) Devise an experiment to verify the result obtained in Exp. 92.
PROBLEM^. XIX.
^ I. If 20 gm. of carbon are heated in the presence of 44 gm. of carbon
dioxide, what weight of carbon monoxide is formed, and what weight,
if any, of carbon remains ?
2. What volume of carbon dioxide must be passed over red-hot
charcoal to yield 159 1. of carbon monoxide ?
3. How many liters of carbon dioxide must be passed over red-hot
charcoal to yield 84 gm. of carbon monoxide ?
'i20 Experimental Chemistry.
4. What volume of carbon dioxide must be passed over what weight
of carbon to produce 10 1. of carbon monoxide at 14'' C. and 760 mm.?
5. What per cent of carbon by weight is contained in carbon
monoxide ? In carbon dioxide ?
6. Carbon dioxide is passed into Hme water and forms a precipitate
weighing 5 gm. How much carbon dioxide can be obtained from this
precipitate ?
7. Calculate the percentage composition of —
(a) Calcium carbonate.
(d) Carbon monoxide.
(c) Carbon dioxide.
(^) Magnesium carbonate.
CHAPTER XIII.
COMPOUNDS OF CARBON AND HYDROGEN — ILLUMINATING
GAS — FLAMES — OXIDATION AND REDUCTION.
Hydrocarbons. — Three important compounds of carbon
and hydrogen are methane, ethylene, and acetylene. They
are called hydrocarbons, and are the first members of three
large classes of similar compounds.
PREPARATION AND PROPERTIES OF METHANE, OR MARSH GAS.
Precaution. — A mixture of methane and air explodes
violently when ignited. Guard against flames and leaks.
Experiment gc,.^,-Ge7ieral Directions: Use the same apparatus as
in Exp. 12. Fill the ignition tube nearly full of the mixture described
below and insert a plug of glass wool, or asbestos between top of the
mixture and the inner end of the delivery tube. Clamp the apparatus
in the position shown in Fig. 23, and collect the gas over water.
Supplies: Heat a few grams of sodium acetate in a porcelain dish
or iron pan until the water of crystallization has been expelled. Use a
small flame and stir constantly toward the end of the operation. Pre-
pare one of the following mixtures : —
{a) Three parts of dehydrated sodium acetate, and two parts soda
lime.
(p ^ Equal parts of dehydrated sodium acetate, dry sodium hydroxide,
and quicklime.
Process: Fill the tube with either one of the above mixtures and
arrange the apparatus as previously described. Heat the whole ignition
tube gently at iirst, increasing the heat gradually until a gas is evolved.
Reject the first bubbles. Collect two small bottles of the marsh gas,
and then remove the end of the delivery tube from the water. Hold a
lighted match at the mouth of one of the bottles, and observe the nature
221
+
2H20
Water
+
2 vols.
+
36
221 Experimental Chemistry.
of the flame, especially its color, luminosity, visible products (if any;,
temperature, etc. Repeat with the other bottle and carefully observe
any properties needing confirmation. Pour a little lime water into one
of the bottles and shake. Explain the result. What evidence of the
composition of methane is presented by this experiment?
When methane burns in air, the combustion is represented
by the equation : —
CH4 + 2O2 = CO2
Methane Oxygen Carbon
Dioxide (23)
I vol. + 2 vols. = I vol.
16 + 64 = 44
PREPARATION AND PROPERTIES OF ETHYLENE.
Precautions. — (a) A mixture of etJiylcne and air explodes
violently, if ignited. Guard against flames a7id leaks.
(b) Hot sulpJiuric acid bums f-ightfully. Gnard agairist
breakage.
Experiment 96. — Apparatus: Use the apparatus shown in Fig.
63. Arrange it to collect a gas over water.
Supplies : Put 5 cc of water in a test tube and slowly pour upon it
15 cc. of concentrated sulphuric acid. Cool the acid by holding the
test tube in a stream of cold water. Put 5 to 7 cc. of alcohol in the
test tube A, add a little clean sand, or a few glass beads, and then
slowly pour in the cotc^ acid. Two or three small bottles will be needed.
A dish should stand under A to catch the contents, in case of accident.
Process : Adjust the apparatus as shown in the figure, taking care not
to crush the test tube. Heat the test tube gently between the bottom
and the surface of the contents to detect any leaks in the apparatus.
Readjust, if necessary. Heat gently to drive out the air, and when it is
judged that the gas which is being evolved is ethylene, collect two or
tliree bottles. As the heat increases the mixture is apt to froth or
'*bump"; sometimes the gas is evolved suddenly. Hence, the heat
must be so regulated that the evolution of gas is slow. Especial care
must be taken not to heat the test tube above the surface of the con-
tents, otherwise a sudden movement of the hot liquid might crack the
Hydrocarbons. 223
tube. As soon as the gas has been collected, disconnect at a^ and
stand the generator in the hood to cool. When cool enough to handle,
pour the contents down the sink or into a receptacle especially provided
for dangerous mixtures.
Test the gas by holding a lighted match at the mouth of a bottle.
Observe and record the color and temperature of the flame, its luminos-
ity, rapidity of combustion, visible products, and any other characteristic
properties. Add a little lime water to one of the bottles in which the
gas was burned, shake, and explain the result.
What evidence does this experiment present regarding the composi-
tion of ethylene?
When ethylene burns, the combustion is represented by
the equation : —
(24)
PREPARATION AND PROPERTIES OF ACETYLENE.
Experiment 97. — Put about 10 cc. of water in a test tube, stand the
test tube in a rack, and drop two or three very small pieces of calcium
carbide into the test tube. Acetylene is evolved. After the action has
proceeded long enough to expel the air, light the gas by holding a
lighted match at the mouth of the tube. Observe and record the nature
of the flame, especially its color, intensity, visible products (if any), tem-
perature, etc. Hold a cold glass plate over the flame. Wliat does the
result suggest about the composition of acetylene ? What other evidence
of its composition is revealed by the properties previously observed?
CLASS-ROOM EXERCISE. XXIX.
I. Hydrocarbons.
{a) Exact meaning of term.
{b) Reason for the vast number.
(^) Meaning of the term " homologous series."
{d) Illustrate by hydrocarbons the terms : saturated, unsaturated,
isoinerisin, polyinerisni, substitution product.
C2H4 +
3O2 =
2C02 +
2H20
Ethylene
I vol. +
3 vols. =
2 vols, -h
2 vols.
28 +
96 =
%8 +
36
224 Experimental Chemistry.
2. Methane.
(«) NameSy and why so called.
{b) Occurrence.
{c) Properties (additional).
{d) What is fire damp ? choke damp ?
3. Ethylene.
{a) Names, and why so called.
(J?) Equation for the reaction in Exp. 96.
{c) Properties (additional).
(^) Liquid ethylene and its uses.
(^) One essential difference between ethylene and methane.
4. Acetylene.
(a) Synthesis, and historical significance of this method of
formation.
(b) Industrial preparation from calcium carbide.
(c) Properties.
Id) Uses.
{e) Essential difference between methane, ethylene, and acetylene.
5. Miscellaneous.
(a) Illustrate the Law of Gay-Lussac by the facts revealed by the
combustion of methane and ethylene.
PROBLEMS. XX.
1. What volume of oxygen will be needed and what volume of
carbon dioxide will be formed in the combustion of {a) 2 1. of methane,
and {b) 2 1. of ethylene ?
2. How many volumes of oxygen are required and how many volumes
of carbon dioxide will be formed in burning 500 cc. of {a) ethylene and
{b) methane ?
3. What weight of air is necessary for the complete combuscion of
1000 cc. of marsh gas at 10° C. and 752 mm. ? What are the volumes
of the products measured at 300° C. and 752 mm. ?
4. What weight of air is needed for the complete combustion of
4 gm. of ethylene ?
5. Calculate the percentage composition of —
{a) Marsh gas, CH^.
{b) Ethylene, CgH^.
{c) Acetylene, C^^-
Illuminating Gas
225
Illuminating Gas is formed by the destructive distillation
of coal. It is a mixture of many gases, and its composition
varies with the coal employed. The constituents are divided
into three classes, viz. illuminants, diluents, and impurities.
The first class contains, among other gases, ethylene and
acetylene. Marsh gas, hydrogen, and carbon monoxide,
all of which burn with a feeble — non-yellow — flame, are
the most common diluents. The common impurities are
carbon dioxide, hydrogen sulphide, and ammonia; some of
these, however, are entirely removed from the purified gas.
Illuminating gas prepared as above is often called coal
gas.
PREPARATION AND PROPERTIES OF ILLUMINATING GAS.
Experiment 98. — Apparatus i The apparatus is shown in Fig. 67.
A A' is an ignition tube from 10 to 15 cm. long. A spiral of copper
wire is placed near A% and the tube is supported by a clamp as near this
end of the tube as convenient. An empty test tube or bottle is con-
nected with the combustion tube by a bent tube passing to the bottom
C
'W
a
B C
C
Fig. 67. — Apparatus for preparing illuminating gas.
of j5; this vessel retains tarry matter which comes from the ignition
tube. The U-tube contains moistened pink Htmus paper in the limb
C and a narrow strip of filter paper moistened with a lead compound
(nitrate or acetate) in the limb C, the latter serving to detect hydrogen
sulphide. The bottle D is connected as shown in the figure, and is to be
226 Experimental Chemistry.
one-third full of lime water. The tube E is to be connected with a
delivery tube passing into a pneumatic trough arranged to collect a gas
over water. Three or four bottles will be needed.
Process: Fill A A' two-thirds full of coarsely powdered soft coal
which should be held in place with a loose plug of glass wool, or
shredded asbestos. See that all connections are gas tight by heating the
ignition tube gently; if the apparatus is tight, the expanded air will
bubble through the bottle D. Readjust, if necessary.
Heat the whole ignition tube gently at first, and gradually increase
the heat, but avoid heating either end very hot, otherwise the closed
end may soften and burst or the rubber stopper may melt. As the heat
increases, watch for marked changes in B, CC, and D. As soon as the
slow bubbling shows that all air has been driven out of the apparatus,
collect, as previously directed, two or three bottles of the gas evolved.
Cover the bottles with wet filter paper or glass plates as fast as they are
removed from the trough. When the last bottle has been removed,
disconnect the apparatus at any convenient point between A' and C
Let the ignition tube cool. Test the gas by bringing a lighted match
near the mouth of a bottle. Observe and record the color and heat of
the flame, the rapidity of combustion of the gas, whether or not smoke
is formed, and any other characteristic property. Repeat with the
remaining gas and observe more closely any facts suggested but not
clearly revealed by the first observations.
Examine the contents of the ignition tube. Does it suggest coke or
some form of carbon ? Examine the bottle B for tarry matter. Does
the paper in C reveal the presence of any ammonia ? If the paper in
C is brown or black, it is due to lead sulphide which is formed by the
interaction of hydrogen sulphide and lead compounds. Did the gas
contain hydrogen sulphide ? Did it contain carbon dioxide ? Record the
result of each observation. Sketch the essential parts of the apparatus.
COMBUSTION OF ILLUMINATING GAS.
Experiment 99. — Attach a pointed glass tube (see Exp. 57 (c)) to
the rubber tube connected with the gas jet and lower a small flame of
illuminating gas into a cold, dry bottle. Observe the most definite
result. Remove and extinguish the flame, add a little lime water, and
shake. Describe the result. What do these two observations show
about the composition of the main constituents of illuminating gas .''
Illuminating Gas. 227
CLASS-ROOM EXERCISE, XXX.
1. Illuminating gas (coal gas).
(a) Industrial preparation, including a description of the appa-
ratus for purification and storage.
(^d) Constituents.
(c) Removal of impurities.
(d) By-products — use and value.
2. Water gas.
(«) Industrial preparation.
(d) Two main constituents.
(c) Why is it often '^ enriched " ?
(d) Dangers from its use.
3. Miscellaneous.
(a) Upon what does the value of an illuminating gas depend ?
(d) How is its value determined ?
(<:) Explain the statement, "This is a 20 candle-power gas."
(d) Describe the Welsbach light.
4. Review.
(«) Combustion.
(d) Distribution of carbon.
(c) Sources of ammonia gas and ammonium compounds.
(d) Source and use of (a) gas carbon, (<^)'coke.
(<?) Law of Multiple Proportions in the light of the facts revealed
by the composition of hydrocarbons.
PROBLEMS. XXI.
1. How much illuminating gas at 10° C. and 530 mm. is required to
fill a tank having a capacity of 800 cu. m. ? (Specific gravity of illumi-
nating gas is 0.5 referred to air, and 14.43 referred to hydrogen.)
2. A bottle contains 53.2 cc. of moist illuminating gas at 760 mm.
and i8.5°C. What is the volume of the dry gas under standard conditions.^
3. When steam is passed over red-hot coal in the preparation of
water gas, the simplest equation for the reaction is^
C + H.O = CO + H...
How many liters of hydrogen and of carbon monoxide at Io° C. and
750 mm. will be formed from 100 gm. of steam .''
228
Experimental Chemistry.
Bunsen Burner. — The Bunsen burner which is used
so constantly in chemical operations was devised by the
German chemist, Bunsen. The
form varies, but the principle is
the same in all burners. The
parts of a common form are shown
in Fig. 6S.
O
O
I
CONSTRUCTION OF A BUNSEN
BURNER.
Experiment lOO. —Take apart a Bun-
sen burner and study the construction.
Sketch the essential parts. Write a short
description of the burner.
The theory and practical advan-
tages of a Bunsen burner are only
partly shown by its structure. It
is also necessary to study the
BUNSEN BURNER FLAME.
Experiment loi. — I. (a) Close the
holes at the bottom of a Bunsen burner
and hold a piece of crayon in the upper
part of the flame. Note the black deposit.
What is it ? Where did it come from ? Open the holes and hold the
blackened crayon in the colorless flame. What becomes of the de-
posit ? How is the flame changed, if at all ? Does the experiment
suggest a cause of the luminosity of a flame ? What is it ?
(d) Dip a glass tube a short distance into powdered wood charcoal,
place the end containing the charcoal in one of the holes at the bottom
of the burner, and blow gently two or three times into the other end.
Describe and explain the result. I3oes it verify the answer to the last
question in (a) ?
(c) Open and close the holes of a lighted burner several times.
Describe the result. Pinch the rubber tube to extinguish the flame,
then light the gas at the holes. What change is produced in the flame.''
What causes the change ?
Fig. 68.
Parts of a Runscn
burner.
Flames.
229
j^
Answer the following : —
(i) What is the object of the holes ?
(2) Why does the gas burn at the top and not inside of the burner?
(3) Why does it sometimes '' strike back " and burn inside ?
(4) Why is the Bunsen flame non-luminous ?
li {^a) Hold a match across the top of the tube of a lighted Bunsen
burner. When it begins to burn, remove and extinguish it. Note
./here it is charred, and explain the result. The
same fact may be shown by sticking a pin through
a (sulphur) match 15 mm. from the tip, suspending
it across the burner, and then lighting the gas. The
position of the match is shown in Fig. 69. Turn on
a full current of gas before lighting it. What does
this experiment show about the structure of the Bun-
sen flame? Verify your answer by (^).
(d) Bend a glass tube about 15 cm. long into the
shape shown in Fig. 70. Hold the shorter arm in
the flame about 2 cm. from the top of the burner
tube. Hold a lighted match for an instant at the
upper end of the tube. What does the result show
about the structure of the Bunsen flame ? Does it
verify (a) ?
{c) Find the hottest part of the flame, when a full current of gas is
burning, by holding a platinum or cop-
per wire in the flame. Measure its dis-
tance, approximately, from the top of
the burner tube.
(d) Examine a typical Bunsen flame
— one w^hich shows clearly the outlines
of the inner part. What is the general
shapQ of each main part? Draw a ver-
tical and a cross section of the flame.
CANDLE FLAME.
Experiment 102. — Attach a candle
to a block of wood by means of a little
melted candle wax, and proceed as fol-
lows : —
l^G. 69. —A (sul-
phur) match sus-
pended across
the top of a Bun-
sen burner.
Fig. 70. — Bent tube for examining
the structure of a Bunsen flame.
230 Experimental Chemistry.
(a) Hold a cold, dry bottle over the lighted candle. Describe the
result. What is the product ? What is its source ? . Remove the
bottle, pour a little lime water into it, and shake. Describe and explain
the result. What are the two main products formed by a burning candle ?
(d) Blow out the candle flame, and quickly hold a lighted match in the
escaping smoke. Does the candle relight ? Why ? What is the general
nature of this smoke? How is it related to the candle wax? How does
(d) contribute to the explanation of (a) ?
(c) Stand a lamp chimney over the lighted candle. How is the flame
affected? Prop up the chimney on two blocks of wood. Does the
candle continue to burn? Why? Now cover the top of the chimney
with a piece of tin or similar covering. What is the result? Why?
LABORATORY EXERCISE. XIII.
1. Sketch a candle flame.
2. Examine a lamp burner. Is it constructed according to the prin-
ciples revealed in Exp. 102 (c) ?
3. What is the essential difference between a Bunsen and a cand'e
flame ?
4. Is there any essential difference between a candle flame and a gas
or lamp flame ?
5. Explain the statement: "Large lamps have a powerful central
draft."
CLASS-ROOM EXERCISE. XXXI.
1 . Essential facts in the life of Bunsen.
2. What is meant by the ignition point or kindling temperature of a
flame?
3. Describe a miner's safety lamp, and state the exact principle on
which its use depends.
4. Give additional facts about the Bunsen burner flame.
5. Give additional facts about the luminosity of flames.
6. Why does a draught of cool air often cause a lamp flame to smoke ?
Oxidizing and Reducing Flames. — It is convenient to
consider the Bunsen flame from two standpoints, viz. its
power to give or to take oxygen, cr in other words its power
to oxidize or to reduce. That portion of the flame giving
Oxidation and Reduction.
231
oxygen freely is called the oxidizing flame, and the part tak-
ing oxygen is called the reducing flame. A diagram of the
general relation of these two flames is shown in Fig. 71
A is the most effective part of the oxidiz-
ing flame, and B of the reducing flame.
At A metals are oxidized, and at B oxides
are reduced.
LABORATORY EXERCISE. XIV.
I, Sketch the oxidizing and reducing flames.
' 2. Why do oxidation and reduction occur respec-
tively at A and B (Fig. 71) ?
Borax Bead. — A practical application
of the oxidizing and reducing flames is
often made by '* testing with a borax bead."
Borax, when heated, melts and forms a
clear, transparent globule, which resem-
bles a glass bead. Many metallic sub-
stances dissolve in melted borax and form
colored beads. These beads differ in
color and thus serve to. identify the sub-
stance or at least its essential elementary
constituent.
TESTS WITH BORAX BEADS,
Fig. 71. — The oxi-
dizing {A) and re-
ducing {B) flames.
Experiment 103. — Make a small loop on the end of the platinum
test wire (see App. A, § 14), moisten it, and dip it into powdered borax.
Heat it in the flame, rotating it slowly ; at first the borax swells, but
finally shrinks to a small, transparent bead. If the bead is too small,
add more borax and heat again. After use, the bead may be removed
by dipping it, while hot, into water; the sudden cooling shatters the
bead, which may then be easily rubbed or scraped from the wire.
{a) Cobalt compounds. Touch a transparent borax bead with a glass
rod which has been moistened with cobalt nitrate or chloride solution.
Heat the bead in the oxidizing flame- Notice the color when cold. If
2^2 Experimental Chemistry.
it is black, melt a little more borax into the bead ; if faintly colored,
moisten again with the cobalt solution. The color is readily detected
by looking at the bead against a white object in a strong light, or by
examining it with a lens. When the color has been definitely deter-
mined, heat again in the reducing flame. Compare the color of the
cold bead with the previous observation.
(d) Copper coiiipowids. Make another transparent bead, moisten it
with copper sulphate solution, and heat it first in the oxidizing flame,
and then in the reducing flame. Compare the colors of the cold beads,
and draw a conclusion.
{c) Manganese cojupounds. Make another transparent bead, moisten
it with manganese sulphate solution, or touch it with a minute quantity
of manganese dioxide, and proceed as in {b). Compare the colors of
the cold beads, and draw a conclusion.
(^) Tabulate the results of this experiment.
Differently colored hot and cold beads are characteristic
of the compounds of different elements, and the bead test
is often used to confirm other observations or to suggest
a further special examination.
The Blowpipe and its Use. — Sometimes it is not con-
venient to apply the bead test, nor to hold a substance
directly in the flame. It is then customary to study the
action of a substance when heated on charcoal in a small
flame, much Hke the Bunsen flame and produced by
Fig. 72. — A blowpipe.
a blowpipe (Fig. 72). A blowpipe tube is put inside
the burner tube (Fig. 73) to produce a luminous flame.
The tip of the blowpipe rests on this tube (Fig. 74), and
when air is gently blown through the blowpipe, a long,
slender flame is produced (Fig. 75), exactly like a Bunsen
Oxidation and Reduction.
'^33
flame as far as its oxidizing and reducing powers are con-
cerned. In Fig. 75, A is the oxidizing and B is the
Fig. 73. — a blowpipe Fig. 74. — The blowpipe
tube in position. tip in position.
Fig. 75. —The blowpipe
flame.
reducing part, i.e. these points are the most effective for
these respective operations.
USE OF THE BLOWPIPE.
Experiment 104. — Slip the blowpipe tube into the burner, light
the gas and lower the flame until it is about 4 cm. high. Rest the tip
of the blowpipe on the top of the tube, as shown in Fig. 74, placing
the tip just within the flame. Put the other end of the blowpipe be-
tween the lips, puflf out the cheeks, inhale through the nose, and exhale
into the tube, using the cheeks somewhat as a bellows. Do not blow
in puffs, but produce a continuous flow of air by steady and easy in-
haling and exhaling. The operation is natural and simple, and, if
properly performed, will not make one out of breath. The flame
should be an inner blue cone surrounded by an outer and almost in-
visible cone, somewhat as shown in Fig. 75, though its shape varies
with the method of production. Practice until the flame is produced
234 Experimental Chemistry.
voluntarily and without exhaustion. Watch the flame and learn to dis-
tinguish the two parts, so that they may be intelligently utilized.
Sketch a blowpipe and a blowpipe flame.
REDUCTION WITH THE BLOWPIPE.
Experiment 105. — (a) Lead compoimds. Make a shallow hole at
one end of the flat side of a piece of charcoal. Fill the hole with a
mixture of equal parts of powdered sodium carbonate and lead oxide,
and heat the mixture in the reducing flame. The sodium carbonate
melts and assists the fusion of the oxide, but the former is not changed
chemically. In a short time bright, silvery globules will appear on the
charcoal. Let the mass cool, and pick out the largest globules. Put
one or two in a mortar, and strike with a pestle. Are they soft and
malleable, or brittle and hard? State the result when a globule is
drawn across or rubbed upon a white paper. How do the properties
compare with those of metallic lead ? What has become of the oxygen ?
Of what chemical use is the charcoal ?
(J)) Grind together in a mortar a little sodium sulphate and wooO
charcoal, adding at intervals just enough water to hold the mass to-
gether. Heat this paste for a few minutes in the reducing flame as in
{a). Scrape the fused mass into a test tube, boil in a little water, and
put a drop of the solution on a bright silver coin. If a dark brown
stain is produced, it is evidence of the formation of silver sulphide.
Repeat, if no such stain is produced. State all tli^ chemical changes
which led to the production of the silver sulphide, explaining at the
same time how the experiment illustrates reduction.
OXIDATION WITH THE BLOWPIPE.
Experiment 106. — {a) Heat a small piece of zinc on charcoal in
the oxidizing flame. What is the product? Observe and record the
color of the product, and the color of the coating on the charcoal when
both hot and cold. (See OO-)
(J)) Heat a piece of lead as in (<^), and note and record the presence
or absence of fumes, as well as the color of the coating wlien hot and
cold.
{c) Heat a small piece of tin in the oxidizing flame, and observe
and record the presence or absence of fumes, and the color of the coat-
ing, if any.
Oxidation and Reduction. 23c
{d) Tabulate the above results, stating in each case (i) color of the
coating on the hot charcoal, (2) color of the coating on the cold char-
coal? (3) presence or absence of fumes, (4) name of product.
CLASS-ROOM EXERCISE. XXXII.
1. What industrial use is made of the principle shown in Exp. 105 {b) .?
2. Does the chemical action in Exp. 106 differ essentially from com-
bustion in the air? Then why use a blowpipe?
3. Review.
{a) Oxidation and reduction.
{b) Compound blowpipe.
CHAPTER XIV.
SOME COMMON ORGANIC COMPOUNDS.
Carbon forms a vast number of compounds. Many of
them were first obtained directly or indirectly from living
things, and they naturally came to be considered under the
comprehensive title of '' Organic Chemistry." Such a
subdivision is now made for convenience, not for any
intrinsic chemical properties of these carbon compounds, or
organic compounds, as they are often called.
Several organic compounds have already been studied ;
these will be reviewed from a broader standpoint, and a
few others which are fa.niliar will be examined.
The common o- ganic compounds are members, or deriva-
tives, of the classes : —
1. Hydrocarbons.
2. Alcohols.
3. Acids.
4. Carbohydrates.
The Composition of Organic Compounds has already been
partially shown. They usually contain carbon and hydro-
gen, with or without oxygen, or nitrogen, or both ; a few
contain sulphur, phosphorus, and other elements. When
burned or heated they yield carbon, or carbon dioxide,
water vapor, ammonia gas, sulphur compounds, etc., ac-
cording to their composition. The general character of
typical organic compounds may be shown by a study of the
236
Some Common Organic Compounds. 237
COMPOSITION OF ORGANIC COMPOUNDS,
Experiment 107. — {a) Carbon. (O Recall or repeat the experi-
ments which showed that carbon is a constituent of wood, cotton,
bone, starch, sugar, illuminating gas, candle wax, meat, flour, bread,
albumen. (2) Heat 2 or 3 cc. of turpentine in a porcelain or iron
dish, and then set fire to it. Does it contain carbon? Hold a bottle
over the flame long enough to collect any product, and then test the
contents for carbon dioxide ; does the observation verify the previous
conclusion? (3) Repeat with alcohol. Does it contain carbon?
Burn a small lump of camphor in a dish or on a block of wood. Does it
contain carbon? (4) Hold a bottle over a burning kerosene lamp long
enough to collect any product, and test as in (2). Does kerosene
contain carbon?
{b) Hydrogen, (i) Set fire to i or 2 cc. of the following liquids ir
a porcelain dish (or crucible), and hold over the flame a cold dry bottle
long enough to allow the condensation of the water vapor which is
always one product of the combustion of organic compounds which
contain hydrogen: alcohol, turpentine, kerosene. (2) Heat in separate
test tubes the following dry solids, and if they contain hydrogen, a little
water vapor will condense on the upper part of the test tube : sugar,
starch, flour, wood, paper, hair. (3) Hold a cold, dry bottle, for a few
seconds over a burning kerosene lamp, a Bunsen flame, an ordinary
gas flame, a burning candle, a burning taper. Is hydrogen a component
of kerosene, illuminating gas, and wax?
The oxygen, which unites with the hydrogen to form the water, may
come from the substance, as in the case of sugar, starch, wax, wood,
paper, or it may come from the air. No simple experiment will deter-
mine the source of the oxygen.
(^) Nitrogen. Mix a little granulated gelatine (one part) with dry
soda lime (two parts) and heat the mixture in a test tube. Hold a
piece of moist red litmus paper In the escaping vapor. It will be
turned blue by escaping ammonia gas. Gelatine (also horn, glue, and
leather) contains nitrogen which is liberated in combination whh hydro-
gen as ammonia gas.
(rt') Sulphur, (i) Put a little mustard paste on a clean silver coin.
The brown stain is silver sulphide. Explain. (2) Why is a silver
spoon tarnished by a cooked egg?
Draw a general conclusion regarding the composition of organic
compounds.
238 Experimental Chemistry.
Hydrocarbons. — These compounds of hydrogen and
carbon have already been considered in Chapter XIII.
LABORATORY EXERCISE. XV.
1. Recall or repeat the experiments illustrating the preparation and
properties of methane, ethylene, and acetylene.
2. What are the products of the combustion of these hydrocarbons?
3. Write the equations for the combustion of methane and ethylene.
CLASS-ROOM EXERCISE. XXXIII.
1. Review Class-Room Exercise xxix., Chapter XIII., page 223.
2. Give the components (if a compound) or the main constituents
(if a. mixture) of kerosene, petroleum, natural gas, illuminating gas,
naphtha, gasoHne, turpentine, benzine, benzene.
PROBLEMS. XXII.
I. Review Problems XX., Chapter XIII., page 224.
Alcohols. — These are compounds of carbon, hydrogen,
and oxygen. Ethyl alcohol is the best known member of
this class, and is therefore called alcohol, just as sodium
chloride is called salt. Alcohol is formed by the fer-
mentation of glucose, or grape sugar.
PREPARATION OF ALCOHOL.
Experiment 108. — Prepare alcohol from one of the following
mixtures : —
{a) 300 gm. of grape sugar. {b) 50 gm. of grape sugar.
2 1. of water. i 1- of water.
500 cc. of yeast. 60 cc. of yeast.
{c) 25 gm. of grape sugar.
150 cc. of water.
A of a compressed yeast cake.
I. Dissolve the grape sugar in the water and add the yeast. Put
the mixture in a large bottle or flask provided with a one-hole rubber
Some Common Organic Compounds. 239
stopper fitted with a delivery tube which reaches to the bottom of a
small bottle half full of lime water.
Pour enough kerosene down a rod upon the lime water to prevent it
from interacting with the carbon dioxide of the air. Stand the appa-
ratus in a dark, or moderately dark place, where the temperature is
25°-30°C.
Fermentation begins at once, and carbon dioxide, one of the prod-
ucts, bubbles through the lime water. Examine the stopper for a
leak, if no change occurs in the lime water. The operation should be
allowed to conf'nue at least a day, and longer if possible. The flask
will then contain mainly water, unchanged grape sugar, alcohol, and
some products of minor importance. Decant the liquid, agitate it with
a httle bone black to remove the odor and color, and filter. The alco-
hol, which varies in quantity with the conditions, is dissolved in a large
excess of water and must be separated by distillation.
II. The distillation is performed with the apparatus used in Exp. 30.
Fill the flask half full of the liquid from I., add a few pieces of pipe-
stem (or granulated zinc, or glass tubing) to prevent "bumping." and
distil about 50 cc. Save the distillate. Replace the residue in the
flask by more liquid from I., distil again, and repeat this operation
until all the liquid has been used. Replace the one-hole stopper with
a two-hole stopper, insert a thermometer in one hole so that the bulb
just touches the surface of the combined distillates which should now be
distilled. Heat gently, and collect in a separate receiver the distillate
which is formed when the liquid boils between 80° and 93° C. This
distillate contains most of the alcohol.
Note the odor. Drop a little into a warm dish, and hold a lighted
match over it. If it does not burn, it shows that the alcohol is too
dilute. Put a little in a dish, warm gently, and light the vapor. De-
scribe the result.
If this experiment is indifferently successful, the general
character of alcohol may be learned by a study of
THE PROPERTIES OF ALCOHOL.
Experiment 109. — (a) Determine cautiously the odor and taste of
alcohol. Drop a little on a glass plate or on a piece of paper, and watch
it evaporate. Is its rate of evaporation more rapid than that of water ?
240 Experimental Chemistry.
(b) Weigh a measured quantity (about 25 cc.) of 95 per cent alcohol
and calculate its specific gravity.
{c) Alcohol dissolves many organic substances. Try camphor, pow-
dered shellac, or rosin.
{d) Burn a little alcohol in a dish and observe the nature of the
flame. What are the products of combustion ?
{e) Cautiously add a few drops of concentrated sulphuric acid to
equal volumes (about 5 cc. each) of acetic acid and alcohol. Shake
carefully, and then warm gently. The pleasant fruit-like odor is due to
the vapor of ethyl acetate, a volatile liquid which is always formed under
these circumstances. Its formation is a simple test for either alcohol or
acetic acid. Its nature will be subsequently considered.
LABORATORY EXERCISE. XVI.
1. Summarize the properties of alcohol.
2. Complete the equation —
C.H,p + = CO2 + H.O.
Fermentation is a destructive process caused, probably,
by various organisms, and the products vary with the organ-
ism. The equation for the reaction when sugar ferments is
QHi^Oe = 2 CsHgO + 2 CO2.
Glucose Alcohol
Cane sugar (ordinary sugar) does not ferment. If boiled
with acid, it changes thus : —
C12H22O11 + H.p = C.HiPe + C.HiPe,
Cane Sugar Glucose Fructose
and fermentation may then occur as usual.
The Formula of Alcohol is often written C2H5.OH, be-
cause conclusive evidence shows that this general grouping
of atoms exists in the molecule. The radical C2H- is called
ethyl. Like ammonium, it exists only in combination. The
C2H5.OH+ HCl
Ethyl Chloride
H2O.
C2H5.OH + C2H4O2
= C2H5.C2H3O2
+
Acetic Acid
Ethyl Acetate
Some Common Organic Compounds. Q,^i
ethyl group is present in many organic compounds, and its
presence in alcohol gives rise to the term et/ij// alcohol.
Alcohols are analogous to metallic hydroxides. Thus
C2H5.OH corresponds to KOH and NaOH. Both form
salts. Hence we may write, as in the case of the experi-
ment on neutralization : —
H,0.
Ethyl acetate is analogous to sodium acetate. It is an
organic salt, and it is often called an etJiereal salt, or ester.
CLASS-ROOM EXERCISE. XXXIV.
1. Additional study of alcohol.
{a) Occurrence.
{h) Various methods of preparation.
{c) Properties.
(^) Purification. (/) Constitution.
{e) Test. {g^) Uses.
(Ji) Determination of alcohol in liquors.
2. Fermentation.
{a) General nature. {c) Applications.
{b) Alcoholic fermentation. {d) Nature of ferments.
3. Composition of brandy, whiskey, wines, and beer.
4. Methyl alcohol.
{a) Preparation. {c) Uses.
{b) Properties. (^) What is methylated spirit }
5. Miscellaneous.
{a) What is "proof spirit" ?
{b) History of alcohol.
(^) Derivation of the word alcohol.
{d) Absolute alcohol.
(^) Tinctures.
242 Experimental Chemistry.
6. Ethereal salts.
(a) Occurrence in flowers and fruits.
(3) General properties.
(c) Artificial preparation. (d) Uses.
Ether is a compound of carbon, hydrogen, and oxygen,
and is made from alcohol by heating with sulphuric acid.
In its narrow sense the term ether means the best known
member of an homologous series of organic compounds
which is analogous to the metallic oxides. Ordinary ether
is ethyl oxide, (C2H5).p or C2H5 . O . C2H5.
PROPERTIES OF ETHER.
Ether vapor is easily ignited, and should never be brought
near a flame, unless special directions are so given.
Experiment no. — (a) Pour a little ether into a dish or test tube
and observe the odor and volatility. Taste cautiously. Pour a drop
upon a glass plate or a block of wood. How does its rate of evapora-
tion compare with that of alcohol ?
(d) Recall or repeat the experiment (see Exp. 33 (d)) illustrating
the solubility relations of ether and water.
(c) Add a bit of wax to a few cubic centimeters of ether. The result
is typical ; draw a conclusion.
LABORATORY EXERCISE. XVII.
1. From what has already been shown about ether, predict its
approximate boiling point.
2. Explain the statement, ''ether is miscible with alcohol in all
proportions.^'
3. Complete the equation for the combustion of ether : -
(C.H^),© + 6O0 = CO., + H,,0.
CLASS-ROOM EXERCISE. XXXV.
I. Additional study of ether.
(a) Industrial preparation. (3) Properties. (c) Uses.
(t/) Formula, constitution, and equations for its preparation.
Some Common Organic Compounds. 243
Miscellaneous.
(a) Ethers.
(d) Significance of the terms e^/iyl ether and sulphuric ether.
{c) Explain the expression, " ether is prepared by a continu-
ous process."
(^) Anhydrous ether.
PROBLEMS. XXIII.
i. Ether boils at 34.9^0. and alcohol at 78.4'^ C. What are the
corresponding points on the Fahrenheit scale?
2. Calculate the weight of alcohol which may be obtained from a
kilogram of cane sugar.
3. Calculate the percentage composition of —
{a) Alcohol, C,H,;0.
{b) Acetic acid, C^H^o^
(c) Cane sugar, CjoHoaOi^.
4. If 10 gm. of pure alcohol are burned, what weight of each product
is formed?
5. One and a half kilograms of absolute alcohol are burned.
{a) HowMiiany liters of oxygen are needed?
{b) How many liters (at standard conditions) of carbon diox-
ide and of water vapor are formed?
6. Calculate the simplest formula of the substances possessing the
indicated composition : —
{a) Carbon
40.00
{c) Carbon
49.05
Hydrogen
6.67
Hydrogen
5.14
Oxygen
53-33
Oxygen
17.20
{b) Carbon
J5.80
Nitrogen
28.61
Hydrogen
5.26
(y) Carbon
74.07
Nitrogen
36.84
Hydrogen
8.64
Sulphur
42.10
Nitrogen
17.29
if) Carbon
54-55
{g) Carbon
35-71
Hydrogen
9,09
Hydrogen
2.38
Oxygen
36.36
Nitrogen
33-33
{d^ Carbon
91.30
Oxygen
28.58
Hydrogen
8.69
244 Experimental Chemistry.
Aldehyde is a compound of carbon, hydrogen, and
oxygen, which is formed by the oxidation of alcohol.
Aldehyde, like alcohol and ether, is the name both of
an homologous series and of its best known member.
PREPARATION AND PROPERTIES OF ALDEHYDES,
Experiment iii. — («) Acetic Aldehyde. Add a little concentrated
hydrochloric acid and a few drops of alcohol to a few cubic centi-
meters of potassium dichromate solution. Warm gently, and observe
the peculiar-smelling gaseous product. It is aldehyde vapor, aldehyde
itself being a colorless, extremely volatile liquid which boils at 20.8° C.
See Exp. 183.
{b) Formic Aldehyde or Formaldehyde. Put a few cubic centi-
meters of methyl alcohol in a test tube and stand the test tube in a
rack. Wind a piece of copper wire into a spiral around a glass rod or
lead pencil. Slip the spiral from the rod, grasp one end in the forceps,
and heat the wire red-hot in the flame. Then quickly drop it into the
methyl alcohol. The pungent vapor which is suddenly produced is the
vapor of formaldehyde.
The preparation of formaldehyde is represented by the
equation —
CH3.OH 4- O = H.CHO + H2O
Methyl Alcohol Formaldehyde
The formula of acetic aldehyde is CHg . CHO.
LABORATORY EXERCISE. XVIII.
1. Complete the equation for the formation of acetic aldehyde : ■
C^H^.OH + O =
2. Derivation of the term ^Z^.?/^/^.?.
3. Why may acetic aldehyde be called ethaldehyde?
CLASS-ROOM EXERCISE. XXXVL
1. Formaldehyde.
{a) Preparation. {b) Properties.
{c) Uses as disinfectant and germicide.
2. Use of aldehydes in formation of mirrors.
Some Common Organic Compounds. 245
Acetic Acid is an exceedingly important organic com-
pound. The acid and its derivatives find numerous appli-
cations.
PROPERTIES OF ACETIC ACID.
Experiment 112. — Recall or determine again the general proper-
ties of acetic acid.
Warm a little in a test tube and compare the odor with that of other
acids.
TEST FOR ACETIC ACID.
Experiment 113. — Repeat Exp. 109 (^), and describe the result.
What is the name of the sodium compound which corresponds to the
main product of the reaction ?
Vinegar is dilute, impure acetic acid.
PROPERTIES OF VINEGAR.
Experiment 114. — Show experimentally that vinegar is (i) an
acid, and (2) contains acetic acid.
Acetates are salts of acetic acid.
PREPARATION OF ACETATES.
Most acetates are poisonous.
Experiment 115. — Prepare one or all of the following acetates: —
{a) Sodium acetate. Dissolve 20 gm. of sodium carbonate in 10 cc.
of water in a large porcelain or agate dish, and slowly add 30 cc. of
commercial acetic acid, with constant stirring. If the solution is not
acid, add a little more acetic acid. Filter the solution, if not clear.
Evaporate to crystallization. When the crystals have formed, remove
and dry them. Describe the crystals. Prove that they contain water
of crystallization. Test the acetate as follows: (i) Dissolve a little
in water, add a few drops of concentrated sulphuric acid, and boil.
What does the odor show is present? What other acids have been
similarly prepared? (2) Dissolve as in (i), add a few drops of alcohol
and of sulphuric acid, and boil. What does the odor conclusively
prove ? Preserve the crystals, finally, in a glass-stoppered bottle, or
in one having a cork covered with paraffine.
246 Experimental Chemistry.
(/;) Lead acetate {poisonous^. To 10 gm. of litharge add 18 cc. of
commercial acetic acid in small portions. Stir the mixture constantly
during the addition of acid. After all the acid has been added, heat
gently until the action ceases. If the solution is green or bluish, it is
due to a copper compound. The copper may be precipitated and re-
moved mechanically by standing a strip of lead in the solution for an
hour or more. Decant the clearer liquid and then filter. Evaporate to
crystallization on a water bath. Remove the crystals from the liquid,
and dry at a moderate temperature. Preserve the crystals finally as in
{a). Describe the crystals. Test them for lead (see Exp. 98, last para-
graph, and Exp. 134 {/) (i)), and for an acetate.
The Formula of Acetic Acid is C.^H^Og or CHg. COOH.
This acid is formed by the oxidation of ethyl alcohol,
thus : —
CsHeO + 02 = C2H4O2 + H2O.
The change is due to fermentation. Acetic acid is mono-
basic.
CLASS-ROOM EXERCISE. XXXVII.
lo Additional study of acetic acid.
{a) Industrial preparation. (^) UseSo
{b) Glacial acetic acid. (^) Test.
ic) Properties. (/") Constitution.
2. Acetates.
(d;) General method of preparation.
(^) Chemical name and formula of iron liquor, red liquor, verdi-
gris, sugar of lead.
{c) Formula of sodium, lead, silver, chromium, ammonium,
aluminium, and calcium acetates.
(^) Uses of acetates.
3. Vinegar.
{a) What is it ? {c) Industrial preparation
{Jj) Chemistry of its formation. (^) Uses.
4. Oxalic acid.
{a) Preparation. {b^ Properties.
(^) Formula and constitution. -
(<'/) What experimental use has been made of oxalic acid? What
do the results show about its composition ?
Some Common Organic Compounds. 247
5. Show by formulas the relation between ethyl alcohol, acetic
aldehyde, and acetic acid, and by equations the possibility of passing
from one to the other.
Soap is a mixture of alkaline salts of organic acids,
mainly palmitic and stearic acids. It is made by boiling
caustic soda or potash with fats. The process is called
saponification, and consists simply in decomposing an or-
ganic salt (fat) into the corresponding alcohol (glycerine)
and an alkahne salt (soap). Fats are analogous to ethyl
acetate, ordinary fats being mixtures of glyceryl palmitate,
stearate, and oleate.
PREPARATION OF SOAP.
Experiment 116. — Prepare soap in an iron or a tin dish by one
of the following methods : —
(a) Dissolve 10 gm. of sodium hydroxide in 75 cc. of water, add 30
gm. of lard, and boil until the mixture begins to solidify. Then add
20 gm. of fine salt in small portions. Stir constantly during the addi-
tion of the salt. Boil a few minutes. Let the mass cool, and then
remove the soap, which will form in a cake at .the surface.
(d) Dissolve 13 to 15 gm. of sodium hydroxide in 100 cc. of water,
add 100 cc. of castor oil, and boil for about half an hour. Add 20 gm.
of salt, and then proceed as in (a).
(c) Dissolve 8 gm. of potassium hydroxide in 150 cc. of alcohol,
add 10 gm. of lard, and stir constantly while the mixture is being heated
on a water bath to syrupy consistency. Allow the contents to cool.
The jelly-like product is soap.
Preserve a sample.
PROPERTIES OF SOAP.
Experiment 117. — (a) Leave soap shavings exposed to the air for
several days. What does the result show about the presence of water
in the soap ?
(d) Test a yellow soap for free alkali. Is such a soap carefully pre-
pared ?
(c) Add considerable dilute sulphuric acid to a soap solution. The
greasy precipitate, which finally rises to the top, is a mixture mainly of
palmitic and stearic acids.
248 Experimental Chemistry.
(d) To a little soap solution in separate test tubes add calcium sul-
phate and magnesium sulphate solutions. Describe the result. Boil
for a few minutes and describe the result. Prepare a solution of acid
calcium carbonate by passing carbon dioxide into lime water until the
precipitate is redissolved. Add some of the solution to a soap solution,
and describe the result. Boil, as above, and describe the result.
Hardness of Water is illustrated by Exp. 1 17 (^)- Mag-
nesium and calcium sulphates and calcium carbonate (in
solution as the acid carbonate) form insoluble salts with
soap. As long as water contains these calcium and mag-
nesium compounds, the soap is useless as a cleansing
agent, since it is immediately transformed into a sticky
solid. If the hardness cannot be removed by boiling, as in
the case of water containing magnesium and calcium sul-
phates, the hardness is called permanent. Heat decom-
poses acid calcium carbonate, the calcium carbonate is
precipitated and carbon dioxide is liberated; hence hard-
ness due to calcium carbonate can be removed by boiling,
and is therefore called temporary. Soft water contains no
magnesium or calcium salts.
CLASS-ROOM EXERCISE. XXXVIII.
1. Additional study of soap.
{a) Industrial preparation. (0 Kinds.
\b) Chemistry of its preparation. (rtf) Uses.
2. Glycerine (or Glycerol).
{a) Industrial preparation.
{b) Properties. (rtT) Formula and composition.
{c) Relation to soap industry. ie) Uses.
3. Miscellaneous.
(«) Stearin, — preparation and uses.
{b) What is saponification?
{c) Organic acids related to stearic acid.
Carbohydrates are compounds of carbon, hydrogen, and
oxygen. Sugar and starch are familiar representatives.
Some Common Organic Compounds. 249
LABORATORY EXERCISE. XIX.
1. Examine sugar and summarize its general properties.
2. What happens when it is heated? When treated with concen-
trated sulphuric acid ? What do these experiments show about its
composition?
3. Examine sugar through a lens. What is its crystal form?
FEHLING's test for SUGAR.
Experiment 118. —Perform Exp. 153.
LABORATORY EXERCISE. XX.
1. Examine starch through a microscope, and describe the result.
How does it differ from sugar?
2. What is the effect of heat on starch ? Of concentrated sulphuric
acid? What does it contain, according to these experiments?
3. Summarize the general properties of starch.
4. Is starch soluble in cold water?
CLASS-ROOM EXERCISE. XXXIX.
1 . Sugars.
(a) Distribution of sugars.
(d) Industrial preparation and properties of glucose.
(c) Cane sugar, — distribution, source, industrial preparation,
properties.
(</) Review " Fermentation."
(<?) Formula and composition of several sugars.
2. Miscellaneous.
(a) What is molasses? Sugar of milk? Fructose? Glucose?
(d) Beet root sugar industry.
3. Starch.
(a) Distribution and source. (d) Uses.
(d) Industrial preparation. (e) Formula.
(c) Properties. (/) Kinds.
4. Chemistry of bread-making.
5. Dextrine — preparation, properties, uses, and relation to starch
and dextrose.
6. Cellulose — distribution, properties, and uses.
CHAPTER XV.
BROMINE, IODINE, AND HYDROFLUORIC ACID.
The elements hitherto studied suggest Httle or no rela-
tion to each other. Chlorine, however, bears a strong
resemblance to bromine and iodine, and the experiments
in this chapter are intended to reveal, among other things,
this relation as well as to present the special facts about
bromine and iodine and their compounds.
PREPARATION AND PROPERTIES OF BROMINE.
Precaution. — Bromine is a corrosive liquid ivhicJi forms,
at the ordinary temperature, a suffocating vapor. Perform
in the hood all experimejits ivJiicli Jtse or evolve bromine.
Experiment 119 — Bend a glass tube about 30 cm. long into the
shape shown in Fig. 76. The lower end passes through the hole of a
rubber stopper which should tightly fit a large test tube.
Put five or six crystals of potassium bromide in the test
tube, add an equal quantity of manganese dioxide and
10 cc. of dilute sulphuric acid. Insert the stopper and
its tube securely, and boil gently. Do not hold the test
tube in the band, but use the test tube holder. Brown
fumes soon appear in the test tube and pass out of the
delivery tube. Regulate the heating so that this vapor
will condense and collect in the lower bend of the deliv-
ery tube. Both vapor and liquid are bromine. When no
further boiling produces bromine vapor in the test tube,
pour the bromine into a bottle of water. Observe and
record the physical properties of this bromine, especially
the color, solubility in water, specific gravity, volatility,
and physical state. Try the action of the contents of
250
Fig. 76. — Ap-
paratus for
preparing
bromine.
Bromine, Iodine, and Hydrofluoric Acid. 251
the bottle on litmus paper ; if the action is not marked, push the paper
down near the bromine. Determine the odor by smelling cautiously
of the water in the bottle. As soon as these observations have been
made, pour the contents of the bottle into the sink and flush with
water, or pour into a jar in the hood. Wash the test tube free from
all traces of bromine, taking care to get none on the hands.
Answer the following : —
1. In what ways does bromine physically resemble chlorins ? In
what ways does it differ from chlorine ?
2. How is it essentially different from all other elements previously
studied ?
The simplest equation for the reaction in the preparation
of bromine is —
2 KBr + 2 H2SO4+ Mn02= Br^ + MnS04 + K2SO4 + 2 H2O
Potassium Bromine Manganese Potassium
Bromide Sulphate Sulphate ^^^^
238 + 196 + ^7 =160+ 151 + 174 + 36
PREPARATION AND PROPERTIES OF HYDROBROMIC ACID.
Perform this experiment in the hood.
Experiment 120. — Put three or four crystals of potassium bromide
in a test tube, and add a few drops of concentrated sulphuric acid.
The white product is hydrobromic acid. Observe its .most apparent
properties — cautiously. Test the gas with both kinds of litmus
paper. Does this product resemble hydrochloric acid ? Are they simi-
lar enough to be closely related ? Give any evidence of a secondary
reaction.
Since 2 NaCl + H.SO^ = 2 HCl + Na.SO^ is the equation for the
preparation of hydrochloric acid, what is the equation for the prepara-
tion of hydrobromic acid ?
Hydrobromic acid forms bromides, just as hydrochloric
acid forms chlorides.
252
Experimental Chemistry.
PROPERTIES OF POTASSIUM BROMIDE.
Experiment 121. — Examine a crystal of potassium bromide. Com-
pare it with potassium or sodium chloride in regard to crystal form,
color, and solubility in water.
Add a few drops of silver nitrate solution to potassium bromide so-
lution. Is the solid product soluble in dilute ammonium hydroxide ?
Compare the interaction with the interaction of silver nitrate and chlo-
rides. Do the properties of bromides, typified by potassium bromide,
suggest any marked relation to chlorides ?
PREPARATION AND PROPERTIES OF IODINE.
Experiment 122. — Mix and grind together in a mortar three or
four crystals of potassium iodide and an equal quantity of manganese
dioxide. Put the mixture in a 250 cc. Erlenmeyer flask,
add about 8 cc. of water, and about 5 cc. of concentrated
sulphuric acid, drop by drop. Invert a dry funnel upon
the flask, as shown in Fig. 'j'js and stand the flask on a
sand bath. Heat with a low, steady flame. Too much
heat will vaporize the water, or crack the flask — perhaps
both. The vapor of iodine will fill the flask, and crystals
of iodine will collect on the sides of the funnel. Regu-
late the heat so that the flask is filled with vapor, but
so that little or none escapes from the funnel. If vapor
escapes from the funnel, plug the stem with cotton or
soft paper. If crystals of iodine collect on the upper
portions of the flask, a gentle heat will remove them.
Continue to heat until no more iodine is evolved, or until
sufficient iodine has been collected for the subsequent
experiments. Scrape the crystals into a dish. Observe
and record the physical properties of iodine, especially the color of the
solid and of the vapor, volatility, and odor (cautiously). Heat a
crystal in a dry test tube, and when the tube is half full of vapor invert
it. What does the result show about the specific gravity of iodine
vapor ? Touch a crystal with the finger. What color is the stain ?
Will water remove it ? Will alcohol ? Will a solution of potassium
iodide ? What do these results show about the solubility of iodine ?
Preserve the crystals.
Fig. 77. — Ap-
paratus for
preparing
iodine.
Bromine, Iodine, and Hydrofluoric Acid. 253
Answer the following : —
1. In what ways does iodine resemble chlorine and bromine? In
what ways does it differ from them ?
2. How does it differ essentially from all other elements previously
studied ?
The simplest equation for the chemical action in the
preparation of iodine is —
2KI + 2H2S04+Mn02=l2+MnS04+K.^S04-r2H20
Potassium Iodine
Iodide (26)
332 + 196 + 8y =254-1-151 + 174 + 36
THE CARBON BISULPHIDE TEST FOR IODINE.
Precaution. — Carbon bisulphide is iiiflamuiable. It
should not be nsed near flames.
Experiment 123. — {a) Free iodine. Add a few drops of carbon
bisulphide to a very dilute solution of iodine, made by dissolving a crys-
tal of iodine in a solution of potassium iodide or in alcohol, and observe
the color of the carbon bisulphide, which, being much heavier than
water, will sink to the bottom of the test tube. How does it resemble
the color of iodine vapor ?
{b) Combined iodine. Add a few drops of carbon bisulphide to a
very dilute solution of potassium iodide. Is there positive evidence
of iodine ? Now add a few drops of chlorine water (see App. C,
Reagents), and shake. How does this result compare with the final
result in (a) ? The result is due to the fact that chlorine liberates iodine
from its compounds, and the iodine, being free, forms the characteristic
color.
THE STARCH TEST FOR IODINE.
Experiment 124. — Grind a gram of starch in a mortar with a little
water to creamy consistency. Pour this into about 100 cc. of boiling
water, and stir the hot liquid. Allow it to cool, or cool it by holding
the vessel in a stream of cold water, and then decant the clear liquid.
Use this starch solution to test for iodine.
254 Experimental Chemistry.
(a) F7'ee iodine. Add a few cubic centimeters of the starch solution
to a test tube nearly full of water, and then add a few drops of a solu-
tion of iodine. The deep blue color is due to the presence of a com-
pound which is always formed under these circumstances, but the
composition of which is unknown. If the color is black, pour out half
of the liquid and add more water, or pour some of the liquid into a
vessel of water.
{b) Co7nbi7ied iodine. Add a few cubic centimeters of the starch
solution to a very dilute solution of potassium iodide. Is the blue com-
pound formed.? Add a few drops of chlorine water and shake. Com-
pare with theJinal result in Exp. 123 (]>).
The facts revealed by Exp. 124 are practically utilized
in — •
THE DETECTION OF STARCH,
Experiment 125. — Test the following for starch by grinding a
small portion of each with water in a mortar, and then adding a few
drops of the extract to an exceedingly dilute solution of iodine : {a)
wheat flour, {]?) oatmeal, {c) potato, {d^ rice, (^) beans, (/) bread,
{g) cracker, (//) pepper, (/) baking powder.
Use, if possible, about the same amount in each case, and the results
will then be an approximation of the relative amount of starch. Tabu-
late the results. If the starch is detected in {Ji) and (/), give a reason
for its presence.
PREPARATION AND PROPERTIES OF HVDRIODIC ACID.
Experiment 126. — Put three or four crystals of potassium iodide
in a test^ tube and add a few drops of concentrated sulphuric acid.
Fumes of hydriodic acid are formed among other products. Observe
the most obvious properties of these fumes as in Exp. 120, and com-
pare with the fumes of hydrochloric and hydrobromic acids. Are the
acids similar?
Record the evidence of secondary reactions in this experiment. If
nothing but hydriodic acid were formed, what would be the simplest
equation? Why would this method be impracticable for the prepara*
tion of hydriodic agid on a large scalQ ?
Bromine, Iodine, and Hydrofluoric Acid. 255
Hydriodic acid forms iodides, which correspond to
chlorides and bromides.
PROPERTIES OF POTASSIUM IODIDE.
Experiment 127. — Examine a crystal of potassium iodide. Com-
pare it with potassium cliloride and potassium bromide in regard to
crystal form, color, and solubility in water.
Add a few drops of silver nitrate solution to potassium iodide solution.
Compare the precipitate with the one obtained in the case of chlorides
and of bromides. Test its solubility in dilute ammonium hydroxide.
Does potassium iodide exhibit any marked resemblance to polassiun:
chloride and potassium bromide?
CLASS-ROOM EXERCISE. XL.
1. Additional study of bromine.
(a) Occurrence.
(d) Sources of commercial bromine.
(<:) Industrial preparation.
(d) Summary of properties experimentally observed.
(e) Additional properties.
(/) Preparation and properties of bromine water.
(g) Uses.
2. Additional study of iodine.
(a) Occurrence.
(d) Industrial jDreparation from seaweed and from Chili saltpeter.
(c) Summary of properties experimentally observed.
(d) Additional properties. (e) Vapor density.
(/) Solubility and color of various solutions. (g) Uses.
3. Historical.
(a) Discovery of bromine by Balard. Liebig's blunder.
(d) Contributions of Courtois, Davy, and Gay-Lussac to the
discovery of iodine.
4. Prepare a table showing the relation between chlorine, bromine,
and iodine. The table should include the atomic weight, physical state,
color, typical method ofpreparation.properties.and analogous compounds.
5. Tabulate the properties o' hydrochloric, hydrobromic, and hydri-
odic acids.
2^6 Experimental Chemistry.
6. What chemical properties are possessed in common by chlorine,
bromine, and iodine?
7. Miscellaneous.
(«) What is " drug store iodine "?
(b) How is Exp. 124 utilized in chemical analysis?
{c) What is a " bromide '"' print?
(^ ) What is iodoform, and for what is it used ?
(e) Meaning and significance of the terms " halogen," " haloid,"
and "hahde."
8. Prepare a short essay on " The Stassfurt Deposits."
PROBLEMS. XXIV.
1. Bromine boils at 138.7° F. What is its boiling point on the
centigrade scale ?
2. The specific gravity of bromine is 3. What volume does a pound
occupy ?
3. Calculate the percentage composition of —
{a) Potassium bromide. {d) Hydrobromic acid.
(^b) Potassium iodide. (<?) Hydriodic acid.
{c) Silver bromide. (/) Iodoform, CHIo.
4. Marignac found that 3.946 gm. of silver dissolved in pure nitric
acid were precipitated by 4.353 gm. of potassium bromide. What is
the atomic weight of bromine? (Assume K = 39.)
5. Iodine vapor is 8.716 times heavier than air. What is its molec-
ular weight ?
6. How much potassium iodide is necessary to prepare 63.5 gm. of
iodine?
7. How much iodine is liberated from potassium iodide by all the
chlorine obtained by boiling 6 gm. of pure manganese dioxide with
hydrochloric acid? (Equation is 2 KI + MnO^ + 4 HCl = Ig + MnClg
+ 2 KCl + 2 H2O.)
8. How much potassium bromide is necessary to prepare 10 gm. of
bromine ?
The Element Fluorine is not adapted to examination in
the laboratory on account of its corrosive properties. Its
acid, however, may be studied, if it is prepared with caution.
Bromine, Iodine, and Hydrofluoric Acid. 257
Fig. 78. — a lead dish.
PREPARATION AND PROPERTIES OF HYDROFLUORIC ACID.
Precaution. — Hydrofluoric acid gas is a corrosive poison.
An aqueous solution of the gas — conwtercial hydrofluoric
acid — burns the flesh frigh t fully.
Experiment 128. —Warm a glass plate about 10 cm. square by
dipping it into hot water or by standing it near a warm object, such as
a radiator or drying oven.
If it is held over a flame
it is liable to crack. Coat
one surface with beeswax
or paraffine. The surface
should be uniformly cov-
ered with a thin layer.
Scratch letters, figures, or
a diagram through the wax
with a pointed instrument. Be sure the instrument removes the wax
through to the glass, and that the lines are not too fine.
Put 5 gm. of calcium fluoride in a lead dish (Fig. 78), and add just
enough concentrated sulphuric acid to form a thin paste. Stir the
mixture with a file or nail. Place the glass plate, wax side down, upon
the lead dish and stand the whole
apparatus in the hood for several hours,
or until some convenient time. Re-
move the plate. Scrape the contents
of the dish, immediately, into a waste
jar in the hood, and wash the dish free
from acid. Most of the wax can be
scraped from the glass plate with a
knife. The last portions can be re-
moved by rubbing with a cloth moist-
ened with alcohol or turpentine. If
time permits, the wax can be removed by supporting the plate on nails
in a dish of water and boiling the water for a few minutes. Do not
attempt to melt off the wax over the flame. If the experiment has
been properly performed, the plate will be etched where the glass
was exposed to the hydrofluoric acid gas- Figure 79 shows an
etched plate.
Fluorine.
M o i s s a n.
18 8 6.
Fig. 79. — Glass plate etched by
hydrofluoric acid gas.
258
Experimental Chemistry.
The simplest equations for the two reactions in Exp. 128
are —
r^F. -4- R.^n ^ o RF _L. ^Q^r
(27)
7^ + 98 =40+136
and —
CaF2
+
H2SO4
= 2 HF
+
CaSO,
Calcium
Hydrofluoric
Calcium
Fluoride
Acid (gas)
Sulphate
7^
+
98
40
+
136
SiOs
+
4HF
= SiF^
+
2H,0
Silicon
Silicon
Dioxide
Tetrafluoride
60
+
80
104
+
36
{2S)
It will be learned later that silicon dioxide is the essential
constituent of the mixture which is made into glass. Silicon
tetrafluoride is a gas.
CLASS-ROOM EXERCISE. XLI.
Fluorine.
{a) Occurrence. {c) Properties.
{b) Isolation by Moissan. {d) Liquefaction of fluorine.
Further study of hydrofluoric acid.
{a) Industrial preparation of the commercial acid.
{b) Properties. (d) Vapor density,
(t) Preservation and transportation. {e) Uses.
Miscellaneous.
(jo) Compare the properties of fluorine with those of chlorine,
bromine, and iodine.
{b) Compare the properties of hydrofluoric acid with the
corresponding acid of the other halogens.
The periodic classification of the elements.
{a) Natural tendency of elements to assume groups.
{b) The Periodic Law stated, expanded, and illustrated by the
halogen groups of elements.
{c) Value of the law in correcting atomic weights and in predict-
ing new elements.
(d) Explain the statement : "Argon may cause a rearrangement
of the periodic classification.''
Bromine, Iodine, and Hydrofluoric Acid. 259
PROBLEMS. XXV.
. I . What is the percentage composition of —
(a) Fkiorspar, CaF2 ?
(^) Cryolite, Na^AlFg ?
2. If 100 gm. of fluorspar are heated with sulphuric acid, how much
calcium sulphate and hydrofluoric acid are formed ?
3. If 100 parts of calcium fluoride yield 174.5 parts of calcium
sulphate, what is the atomic weight of fluorine ? (Assume CaFg and
Ca = 40.)
CHAPTER XVI.
SULPHUR AND ITS COMPOUNDS.
Sulphur is one of the most important of the non-metallic
elements.
PHYSICAL PROPERTIES OF SULPHUR.
Experiment I2g. — (a) Examine a piece of roll sulphur and write
a short account of its most obvious physical properties.
(d) Weigh a lump of roll sulphur to a decigram. Slip it carefully
into a graduate previously filled with water to a known point — about
half full — and note the increase in the volume of water. This increase
in volume is equal to the volume of the sulphur. Calculate the specific
gravity of sulphur from the observed data.
AMORPHOUS SULPHUR.
Experiment 130. — Put a few pieces of roll sulphur in an old test
tube. Heat carefully until the sulphur boils, and then quickly pour the
contents of the test tube into a dish of cold water. This is amorphous
sulphur. Note its properties. Preserve, and examine it after twenty-
four hours. Describe the change, if any.
Define amorphous, and illustrate it by this experiment.
Sulphur exists not only as the amorphous variet}^, but
it crystallizes under favorable conditions in two distinct
crystal systems.
CRYSTALLIZED SULPHUR.
Experiment 131. — (a) Monoclinic. Fill a small Hessian crucible
nearly full of roll sulphur. Support the crucible in a ring, and heat until
all the sulphur is melted. Let it cool, and the instant crystals shoot
out from the walls just below the surface, pour the remaining melted
sulphur into a dish of cold water. When the crucible can be handled
without discomfort, crack it open lengthwise. The contents should
resemble the conventional sketch of a typical mass of crystals shown in
Fig. 80. Observe and record the properties of the crystals, especially
260
Sulphur and Its Compounds.
261
the shape, size, color, kister, l)rittlcncss or elasticity, and any other
characteristic property. Allow the best crystals to remain undisturbed
for a day or two ; then reexamine, and record
any marked changes.
{b) Orthorho7nbic. Put 3 gm. of flowers
of sulphur in a test tube and add about 5 cc.
of carbon disulphide — remember the precau-
tion to be observed in using this liquid (see
Exp. 123). wShake until all the sulphur is
dissolved, then pour the clear solution into
a small dish to crystalHze. It is advisable,
though not absolutely necessary, to stand
the dish in the hood, where there is no
flame and where the offensive vapor will be
quickly removed. Watch the crystalliza-
tion toward the end, and, if perfect crystals
form, remove them with the forceps (see
Fig. 81) . Allow the liquid to evaporate almost entirely, then remove and
dry the crystals. Examine them as in {a) and record their properties.
Fig. 80. — Section of a cruci-
ble showing monoclinic
crystals of sulphur.
Fig. 81. — Orthorhombic crystals of sulphur.
{c) Tabulate the essential result$"in {a) and {U).
{d) Make an outline sketch of a crystal of orthorhombic sulphur.
Sulphur combines with both non-metals and metals, and
forms sidphidcs in most cases.
COMBINING POWER OF SULPHUR.
Experiment 132. — {d) Set fire to a little sulphur in a deflagrating
spoon, and lower the spoon into a bottle. Caiitioiisly waft the fumes
262 Experimental Chemistry.
toward the nose, and observe and dcscril)e the odor. The product is a
compound of sulphur. From what has already been shown experi-
mentally about comi)ustion, state the components of the fumes.
(d) Mix 3 gm. of flowers of sulphur with 3 gm. of iron powder or
clean iron filings, and heat half of the mixture in an old test tube. As
soon as there is decided evidence of chemical action, remove the test
tube from the flame. When the tube has cooled, break tlie end by
rapping it sharply on a hard surface. Examine the contents. What
is the evidence of the formation of a new compound ? V^erify the
conclusion by adding a little dilute hydrochloric acid to the product
and then to the remainder of the original mixture, testing the gaseous
product in each case by the odor.
Summarize the evidence regarding the power of sulpliur to combine
with iron.
Similar results are obtained with copper and other metals. Draw a
general conclusion regarding the combining power of sulphur.
CLASS-ROOM EXERCISE. XLII.
1. Additicnal study of sulphur.
(a) Occurrence of free sulphur.
(d) Occurrence of combined sulphur.
(c) Theories of the volcanic and the organic formation of
sulphur.
(^) Extraction of sulphur from native sulphur.
(<?) Other sources of sulphur.
(/) Purification.
(^) What is roll sulphur ? Brimstone ? Flowers of sulphur ?
(//) Additional properties, e.g. solubility.
(/■) Illustrate allotropy by means of sulphur.
(/) What is milk of sulphur ? (k) Uses.
2. Miscellaneous.
(a) Review of the orthorhombic and monoclinic systems of
crystallization. (See class-koom kxercise. vii., 7 (^/).)
(d) Illustrate the term amorphous by sulphur.
ic) What part has sulphur played in the history of chemistry ?
(^/) In what crystal form does free crystallized sulphur occur ?
Hydrogen Sulphide. — An important compound of sul-
phur is the gas, hydrogen sulphide. It is formed in many
Sulphur and Its Compounds.
263
chemical changes, it possesses characteristic properties, it
forms many compounds, and is an indispensable laboratory
reagent. It is a poisonous gas, and should never be lib-
erated in the laboratory.
PREPARATION AND PROPERTIES OF HYDROGEN SULPHIDE.
Perform this experiment in the hood, and do not inJiale
the gas.
Experiment 133.— Construct a generator like that shown in Fig. 82.
The test tube A is clamped to an iron stand at such a height that
the end of C reaches to the bottom of a bottle
in which the gas is collected by downward dis-
placement and the solution subsequently prepared.
The essential part of the apparatus is also shown
in Fig. 54, the tube C being joined to B by the
rubber connector a. Fill the test tube one-third
full of coarsely powdered ferrous sulphide, cover
\^
B
Fig. 83. — Perforated card to prevent the
diffusion of hydrogen sulphide gas.
with dilute hydrochloric acid, and quickly insert ^^^ 82. — Apparatus
the stopper. Hydrogen sulphide gas is rapidly for generating hy-
evolved. DiiTusion of the gas maybe prevented drogen sulphide gas.
by covering the bottle with the block of wood used
in the collection of chlorine, or with a piece of stiff paper shaped as
shown in Fig 83, and cut along the line ab to allow it to be slipped
264 Experimental Chemistry.
easily upon the delivery tube. Collect three small bottles, removing
each, and covering with a glass plate as soon as full. If the evolution
of gas stops or slackens, warm gently or add more hydrochloric acid.
After the bottles of gas have been collected, allow the gas from the
generator to bubble through a bottle half full of water. Hydrogen
sulphide water will be formed, which should be preserved in a stop-
pered bottle for future use ; or, if more convenient, it may be poured
into the stock bottle.
Meanwhile test the hydrogen sulphide gas as follows : —
(a) Waft a //Ule of the gas cautiously toward the nose, and then
describe the odor. This is characteristic of hydrogen sulphide, and is
a decisive test. Has the gas color ?
(d) Test the gas from the same bottle with both kinds of litmus
paper. Is it acid or alkaline ?
(c) Bring a lighted match to the mouth of a bottle. Observe the
properties of the flame as in previous experiments. Observe cautiously
the odor of the product of the burned gas ; to what compound is the
odor due .'* What, then, is one component of hydrogen sulphide ?
(d) Burn another bottle of hydrogen sulphide and hold a cold bot-
tle over the burning gas. What additional experimental evidence does
this result give regarding the composition of hydrogen sulphide ?
LABORATORY EXERCISE. XXI.
1 . Summarize the properties of hydrogen sulphide and the experi-
mental evidence of its composition.
2. Complete the equation for the combustion of hydrogen sul-
phide:— 2H,S + O = 2 SO. +
The simplest equation for the preparation of hydrogen
sulphide gas is —
FeS + 2HCI = H^S + FeCl^
Ferrous Hydrogen Ferrous
Sulphide Sulphide Chloride
88 -I- 73 = 34 + 127
A liter of dry hydrogen sulphide gas under standard
conditions weighs 1.542 gm.
Sulphur and Its Compounds. 26^
Sulphides. — Hydrogen sulphide is a weak acid and
forms salts called sulphides. Many of these sulphides
have characteristic colors and they vary in solubility, hence
they are easily detected and separated. They are readily
formed by adding to solutions of compounds of metals a
solution of hydrogen sulphide gas, — usually called hydro-
gen sulphide water, — or by passing the gas through the
solution of the metal ; the moist gas itself acts immediately
on some metals.
PREPARATION AND PROPERTIES OF SOME SULPHIDES.
Experiment 134. — (a) Shake the bottle of hydrogen sulphide
water prepared in Exp. 133 (or a similar solution), and Hold succes-
sively at the mouth or in the neck of the bottle (i) a clean copper wire
or piece of sheet copper, (2) a bright strip of lead, and (3) an un-
tarnished silver coin. Describe the result in each case. These com-
pounds are sulphides of the respective metals.
(d) Put a gram of litharge — the brownish-yellow oxide of lead — in
a test tube, cover it with hydrogen sulphide water, and warm gently.
Describe and explain the change. The formula of pure litharge is PbO.
Complete the equation : —
PbO + H.S = H.O +
(c) Prepare the following solutions in separate test tubes : —
(i) Dissolve a small crystal of lead nitrate in water.
(2) Boil a minute quantity of arsenic trioxide in dilute hydro-
chloric acid.
(3) Dissolve a minute quantity of tartar emetic — a compoimd of
the metal antimony — in water.
(4) Dissolve a small crystal of zinc sulphate in water.
All the above solutions are poisonous. The test tubes should be
about half full. Add about one-fourth its bulk of hydrogen sulphide
water to each solution, and shake thoroughly. Observe the color of
each sulphide, and tabulate the results.
The varying solubility of sulphides permits —
266 Experimental Chemistry.
THE SEPARATION OF SULPHIDES.
Experiment 135. — Mix equal quantities of dilute solutions of
(i) copper sulphate, (2) ferrous sulphate, and (3) sodium chloride.
Add dilute hydrochloric acid. Pass hydrogen sulphide gas through the
mixture (use the same generator as in Exp. 133), or add considerable
hydrogen sulphide water. In either case, warm gently, allow the pre-
cipitate to settle, and then add a few drops of hydrogen sulphide water
to ensure complete precipitation. If precipitation is not complete,
add more hydrogen sulphide water, and proceed as above. If com-
plete, decant the clear liquid into a clean vessel by pouring it down a
glass rod, and save it for subsequent examination. Add dilute hydro-
chloric acid to the precipitate, warm gently, and filter. There are now
three parts to test. Proceed as follows : —
(i) Precipitate. Stand the funnel in a clean test tube, make a hole
with a pointed glass rod in the apex of the filter paper, and wash the
precipitate into the test tube with dilute nitric acid by pouring the acid
down a glass rod directly upon the solid. Use as little acid as possible.
Boil until the black precipitate is dissolved, and then add considerable
ammonium hydroxide. The solution should become a distinct blue,
which is evidence that a copper compound is present. Had an iron
compound been present, the result would be entirely different.
(2) Filtrate. Add ammonium hydroxide to the filtrate, and a black
precipitate will be formed. This is iron sulphide. Dissolve it in hy-
drochloric acid, add a few drops of nitric acid, boil, and then add
ammonium hydroxide. A rusty-red precipitate of ferric h3droxide is
formed, showing that an iron compound was the main constituent of
the filtrate.
(3) Clear liquid. Filter, if this liquid is not free from particles.
Apply the flame test for sodium to this liquid.
Summarize the results of the whole experiment.
The above experiment illustrates briefly the general
method of separation and detection of metals employed
in Qualitative Analysis. Sulphides are not the only com-
pounds, however, which permit systematic separation. A
more explicit treatment of this subject will be found in
Chapter XIX., page 322.
Sulphur and Its Compounds. 267
CLASS-ROOM EXERCISE. XLIII.
1 . Additional study of hydrogen sulphide.
(a) History.
(i>) Names and their significance.
(c) Occurrence.
(d) Properties not experimentally shown.
(e) Uses.
2. Sulphides.
(a) Natural sulphides.
(d) Exact name and color of sulphides of arsenic, antimony,
zinc, lead.
(c) Why does a silver spoon blacken in mustard? In a
cooked egg?
{d) Why does white paint often turn brown or black ?
{e) Why does silverware blacken?
(/) Name and formula of cinnabar, stibnite, copper pyrites,
and pyrite.
3. Composition of hydrogen sulphide gas.
(i) If a spiral of platinum wire is heated in a bottle of hydrogen sul-
phide gas, a yellow powder and an invisible gas are formed ; the powder
burns with a pale blue flame, producing a choking vapor, and the in-
visible gas burns with a colorless, hot flame. W^iat evidence do these
facts present about the composition of hydrogen sulphide?
(2) Sulphur is often deposited in bottles of hydrogen sulphide
water. Chlorine passed into hydrogen sulphide forms sulphur and
hydrochloric acid. What do these facts show about the composition of
hydrogen sulphide ?
(3) Give additional evidence of the composition of hydrogen
sulphide.
PROBLEMS. XXVI.
1. How much does a stick of brimstone i m. long and 4 cm. in
diameter weigh ?
2. Dumas heated 10.000 grains of silver in sulphur vapor and obtained
1^481.5 grams of silver sulphide. What is the atomic weight of sul-
phur?- (Assume Ag = 108, and Ag,, + S = Ag.,S.)
3. How many grams of sulphur can be obtained from 600 cc. of
hydrogen sulphide gas ?
268 Experimental Chemistry.
4. How many grams of ferrous sulphide are necessary to yield a liter
of hydrogen sulphide gas?
5. Calculate the percentage composition of —
(a) Galena, PbS. (c) Iron pyrites, FeSg.
(d) Zinc blende, ZnS. (d) Ferrous sulphide, FeS.
6. What weight and what volume of hydrogen can be obtained from
102 gm. of hydrogen sulphide?
The two oxides of Sulphur, sulphur dioxide ana sulphur
trioxide, and their related acids, are the most important
compounds of sulphur.
PREPARATION AND PROPERTIES OF SULPHUR DIOXIDE.
Perform this experiment in the hood.
Experiment 136. — Construct a generator like that shown in Fig.
82. The test tube should be clamped high enough to be conveniently
heated with a free flame and the delivery tube must be long enough
to reach to the bottom of an empty bottle. Put a few grams of cop-
per borings in the test tube, cover with concentrated sulphuric acid,
and arrange the apparatus to collect a gas as in Exp. 133. Heat gently
with a low flame, and sulphur dioxide gas is evolved. As the reaction
proceeds, the evolution of gas may become violent. Remove the flame
at intervals, regulating the heat so that the evolution of gas is slow and
steady. Remember that hot sulphuric acid is a dangerous liquid.
Collect four small bottles of gas, covering each with, a glass plate as
soon as full. Then remove the stand with its attached generator to
the hood or some place where it will not be disturbed. Study the sul-
phur dioxide gas as follows : —
{a) Observe the most obvious physical properties, e.g. co or, odor
(cautiously), specific gravity.
{b) Prove that it is the same gas as was formed in Exp. 132 (<?).
{c) Hold a blazing stick in a bottle of the gas. Will the gas burn
or support combustion? What previously acquired facts would have
enabled you to predict this result?
{d) Pour 10 or 15 cc. of water into a bottle of sulphur dioxide, cover
with the hand, and shake. What is the evidence of solution? Is the
resulting liquid acid, alkaline, or neutral?
Sulphur and Its Compounds. 269
{e) Moisten a pink or purple flower with a few drops of water, hang
it in the remaining bottle of sulphur dioxide, holding it in place by putting
the stem between the glass and a cork. Observe and describe any
change in the color of the flower. If the action is slow or indefinite,
stand a deflagrating spoon containing burning sulphur in the bottle for
a minute or two. What is this operation called?
LABORATORY EXERCISE. XXII.
1. Light a sulphur match and observe and describe the entire action,
as far as the sulphur is concerned.
2. What is the function of the sulphur in a burning match?
Sulphurous Acid is formed when sulphur dioxide dis-
solves in water. Sulphur dioxide is, therefore, sulphurous
anhydride. The simplest equation expressing this fact is —
SO2 + H2O = H2SO3
Sulphurous
Acid
It is an unstable acid, and gradually forms sulphuric
acid by combining with oxygen from the air. Its salts —
the sulphites — yield sulphur dioxide when treated with
acids.
CLASS-ROOM EXERCISE. XLIV.
I. Additional study of sulphur dioxide.
{a) Occurrence.
(J?) Additional properties, e.g. solubility, specific gravity.
{c) Relation to combustion.
{d) Explanation of bleaching power.
{e) Liquid sulphur dioxide."'
{/) Composition and formula.
Sulphur Trioxide, SO3, is formed when a mixture of sul-
phur dioxide and oxygen is passed over heated platinized
asbestos. It is the anhydride of sulphuric acid. The
equation expressing the latter fact is —
SO3 + H2O = H2SO4
270 Experimental Chemistry.
Sulphuric acid has not, until recently, been industrially
prepared in this way. Its preparation is not well adapted
to the laboratory, but its properties may be profitably
studied.
ACTION OF SULPHURIC ACID WITH WATER.
Experiment 137. — Recall or repeat that part of Exp. 36 which
shows the action of sulphuric acid with water. Record the result.
This experiment illustrates the necessity of always add-
ing the acid to the zvatcr. The most obvious result is due
to the formation of a compound of sulphuric acid and
water. The experiment also proves the avidity with which
sulphuric acid unites with water. This affinity often decom-
poses compounds which contg.in hydrogen and oxygen in
the proportion to form water.
ACTION OF SULPHURIC ACID WITH ORGANIC MATTER.
Experiment 138. — {a) Write some letters or figures with dilute
sulphuric acid on a sheet of white paper, and move the paper back and
forth over a low flame, taking care not to set fire to the paper. As the
water evaporates the dilute acid becomes concentrated. Observe and
describe the result. Paper is largely a compound of carbon, hydrogen,
and oxygen, and the hydrogen and oxygen are present in the propor-
tion to form water. Explain the general chemical change in this
experiment.
{b) Dissolve 3 or 4 gm. of sugar in a test tube half full of warm
water, stand the test tube in the rack, and add cautiously several drops
of concentrated sulphuric acid. If there is no decided result, add more
acid. What is the black product? Compare the final result with that
obtained in Exp. 82 {b). Is the chemical action the same in each
experiment? Are the statements made in {a) about paper also true of
sugar ?
{c) Repeat {b)^ using powdered starch instead of sugar. Describe
the result. How does the result resemble that in (Jb) and in Exp. 82 {a) ?
Predict the components of starch. In what simple way may the pre-
diction be verified?
Sulphur and Its Compounds. 271
(d) Stand a stick of wood in a test tube one-fourth full of concen-
trated sulphuric acid. Allow it to remain in the acid for fifteen minutes,
then remove the stick and wash off the acid. Describe the change.
Does it resemble that in (a), {b), and {c), and in Exp. 82 {a) ?
LABORATORY EXERCISE. XXIII.
1. Can sulphuric acid be used to dry gases? Can it be used to dry
ammonia gas? Why?
2. Is sulphuric acid volatile? What experiment proves this?
3. W^hat happens when concentrated sulphuric acid is exposed to
moist air?
4. Determine the specific gravity of sulphuric acid by weighing a
measured volume (25 cc.) in a graduate. Compare the result with the
specific gravity of the commercial acid.
TEST FOR SULPHURIC ACID.
Experiment 139. — Devise a test for sulphuric acid. Verify it. Is
the same test applicable to soluble sulphates ? Devise, or find in a text-
book, a test for sulphates almost insoluble in w^ater, e.g. calcium sul-
phate. Verify it. How can a sulphate be distinguished from a
sulphite? Is barium sulphate absolutely insoluble in all reagents?
CLASS-ROOM EXERCISE. XLV.
1 . Additional study of sulphuric acid.
(«) Complete account of its industrial preparation.
(b) Properties not experimentally found.
{c) Uses. {d) History.
{e) Significance of the term oil of vitriol.
2. Miscellaneous.
{a) Explain and illustrate the^ statement, "sulphurous and sul-
phuric acid are dibasic."
{b) Valence of sulphur in H,S, SO,, SO,, H.SO,, H,SO,.
(c) Give the chemical name and formula of gypsum, heavy spar,
plaster of Paris, white vitriol, blue stone, green vitriol,
Glauber's salt, kieserite.
{d) Illustrate the Periodic Law by sulphur, selenium, and tel-
lurium.
{e) What is ''hypo"? What is sodium thiosulphate? Use of
the latter.
272 Experimental Chemistry.
PROBLEMS. XXVII.
1. What volume of air is necessary to burn 1000 gm. of sulphur?
2. How many grams of sulphur can be obtained from a liter of
sulphuric acid? (Assume the commercial density.)
3. Calculate the percentage composition of —
(a) Barium sulphate.
{b) Zinc sulphate.
(c) Copper sulphate.
(d) Sodium sulphate. 1
{e) Chamber crystals, SO..NO0.HO.
(/*) Nordhausen acid, H^S^O;.
CHAPTER XVII.
SILICON AND PHOSPHORUS AND THEIR COMPOUNDS.—
REVIEW.
Silicon does not occur in the free state, but its com-
pounds are abundant and widely distributed. Approxi-
mately one-fourth of the earth's crust is silicon. Its
most important compound is silicon dioxide, Si02, which
is familiar as sand and quartz.
The properties of some compounds of silicon are shown
by the
RELATION OF SILICON DIOXIDE TO OTHER COMPOUNDS
OF SILICON,
Experiment 140. — Fill a porcelain or platinum crucible half full of
a fusion mixture — sodium carbonate (4 parts) and potassium carbonate
(5 parts). Support the crucible on a tripod, or a triangle, and heat
gently, until the fusion mixture is melted. Then add a pinch of clean,
exceedingly fine sand, and heat with the hottest available flame — a
blast lamp flame, if possible. Add three or four pinches of sand from
time to time, and finally heat until no particles of sand can be seen
suspended in the molten mass. The heat must be intense and steadily
applied, but so regulated that there is- no undue frothing. Cool the cru-
cible gradually, especially if it is porcelain, and then place it in a beaker
or casserole ; cover with water and boil gently until the mass is detached.
Remove the crucible, and heat until the residue is dissolved. Filter, it
the solution contains any sand. Add dilute hydrochloric acid slowly to
the clear liquid until the carbonates are decomposed. The gelatinous
precipitate is siHcic acid. Rub some between the fingers and describe
the result. Evaporate the precipitate to dryness upon a water bath in a
porcelain dish, or casserole, in the hood. As the mass hardens, stir it
273
274 Experimental Chemistry.
with a glass rod. Toward the end, add more liydrochloric acid and
evaporate to complete dryness. Then stand the dish on a gauze-covered
tripod and heat strongly for live minutes. The residue is silicon dioxide
mixed with chlorides of sodium and potassium. Rub some between the
fingers or across a glass plate. Is any grit detected?
If the apparatus required for Exp. 140 is not available,
the same results may be shown by the
PREPARATION AND PROPERTIES OF SILICIC ACID.
Experiment 141. — Add hydrochloric acid to a solution of sodium
silicate — " water glass." The gelatinous precipitate is silicic acid.
Proceed from this point as in Exp. 140.
Discussion of Experiment 140. — When silicon dioxide is
fused with alkaline carbonates, an alkaline sihcate is formed.
The alkaline silicates are soluble in water, and when hydro-
chloric acid is added to such a solution, the gelatinous pre-
cipitate formed is a siHcic acid having the formula H2Si03
(probably). This acid is decomposed by heat into silicon
dioxide and water, thus : —
HaSiOg = SiO^ + H2O
Silicic Silicon
Acid Dioxide
The silicic acids are numerous and complex, and the
silicates which they form constitute a large class of im-
portant minerals and rocks.
CLASS-ROOM EXERCISE. AL.VI.
I. Additional study of silicon.
(a) Occurrence and distribution.
{d) Preparation and properties of amorp.ious and of crystallized
silicon.
{c) Resemblance to carbon.
Silicon and Phosphorus and Compounds. 275
2. Silicon dioxide.
(a) Occurrence and distribution.
{d) Properties of various forms.
(c) Geological significance.
(d) Uses.
{e) Reason for abundance.
{/) Meaning of the term silica.
3. Other compounds of silicon.
{a^ Silicon tetrafliioride. {c) Silicates.
ib) Silicic acids. id) Carborundum.
4. Miscellaneous.
{a) What is diatomaceous earth ? Examine it through a micro-
scope.
{p) What is " electro-silicon " ?
(c) What is opal ?
{d) Crystallization of quartz,
ie) Significance of the term rock crystal.
(/) What is the essential chemical constituent of many rocks?
{g) Write an essay on " Glass."
(//) Prepare a short account of the manufacture of pottery and
porcelain.
PROBLEMS. XXVIII.
1. Scheele found that 0.6738 gm. of silicon tetrachloride gave 2.277
gm. of silver chloride. Calculate the atomic weight of silicon. (Assume
the equation, SiCl^ + 4 AgNOg + 2 HgO = SiO., + 4 AgCl + 4 NH^NOg.)
2. How much silicon can be obtained from 119 gm. of potassium
silicofluoride? (Assume the equation, K^,SiFg + 4 K = Si + 6 KF.)
3. Calculate the percentage composition of —
{a) Willemite, Zn^,SiO^.
{b) Steatite, Mg.Sip,, . H.O.
Phosphorus itself is not adapted to extended laboratory
examination on account of its properties. It takes fire
suddenly and produces dangerous burns. The element has
some uses, and its compounds have many industrial appli-
cations.
276 Experimental Chemistry.
LABORATORY EXERCISE. XXIV.
1. Recall or read again the notes made about the properties of phos-
phorus observed in Exp. 45.
2. Smell of the end of a phosphorus-tipped match, then smell of the,
fumes arising from a bottle containing phosphorus under water. Do
they suggest each other?
3. If convenient, rub the end of a phosphorus-tipped match in the
dark, and observe and describe the result.
CLASS-ROOM EXERCISE. XLVII.
1. A study of yellow phosphorus.
{a) History.
(^) Significance of name.
(c) Occurrence and sources.
(d) Manufacture.
(^) Properties.
(/) Uses.
(g) Danger in its use.
(//) How it should be kept.
2. A study of red phosphorus.
(a) Preparation.
((^) Properties.
(c) Uses.
3. Compounds of phosphorus.
(a) Preparation and properties of phosphorus pentoxide.
(d) Orthophosphoric acid, H0PO4, and its salts.
(c) Metaphosphoric acid, HPO^, and its salts.
4. Miscellaneous.
(a) Use of phosphates as food of plants.
(d) Fertilizer — manufacture, constituents, use.
(c) Test for phosphorus.
(d) Write an essay on " Matches."
(e) What is bone black? Bone ash? A cupel? Apatite?
Phosphate rock ? Microcosmic salt ? Disodium phosphate ?
(/) Illustrate allotropy by phosphorus.
(^) To what group in the periodic classification does phosphorus
belong?
Silicon and Phosphorus and Compounds. 277
S. A study of arsenic.
{a) Occurrence. {b) Preparation. (<:) Properties.
{d) Preparation, properties, and uses of arsenic trioxide.
(6') What is Marsh's test for arsenic?
(y) What is " arsenic,^' white arsenic, arsenious oxide ?
{g) Antidote for arsenic poisoning.
Review. — The foregoing part of this book contains most
of the fundamental principles of chemistry illustrated by
experiments as far as such a method permits. It is advis-
able at this stage to glance backward and review these
principles in the light of the facts gleaned from all sources.
CLASS-ROOM EXERCISE. XLVIII.
1 . Discuss the atomic theory as an expression of the Laws of Multiple
and Definite Proportions.
2. Discuss molecules and molecular weight in the light of the Laws
of Avogadro and Gay-Lussac, vapor density, vapor tension, elevation of
the boiling point, and depression of the freezing point.
3. Illustrate Gay-Lussac's Law by the volumetric formation of water,
ammonia gas, hydrochloric acid gas, hydrogen sulphide, sulphur dioxide,
and by the oxides of carbon.
4. Illustrate the Law of Multiple Proportions by the oxides of nitro-
gen, carbon, and sulphur, and by the hydrocarbons, methane, ethylene,
and acetylene.
5. Illustrate the Law of Definite Proportions by Exps. 15, 16, 22,
26, 44, 54, 55, 92.
6. Discuss the causes, accompaniments, and results of chemical
action.
7. Discuss : —
{a) Heat of combustion. {c) Allotropism.
{b) Solution. {d) Valence.
8. Discuss with illustrations : —
(«) The Periodic Law.
(J?) Reduction and oxidation.
{c) Equations.
{d) Elements and compounds.
{^e) Symbols and formulas.
278 Experimental Chemistry.
9. State the test for a chloride, sulphate, sulphide, sulphite, nitrate,
carbonate, and for Zn, Na, Ca, Cu, Pb, Ag, C. O, H, Sb, As, CO,, N,
CI, Br, I, S, H^S, SOo, SiO^,, P, an alkali, an acid, and a neutral salt.
10. State the chemical name and formula of lime, lime water, lime-
stone, chloride of lime, bleaching powder, marble, muriatic acid, oil of
vitriol, blue vitriol, ammonia, sal ammoniac, salt, soda, potash, caustic
soda, quicklime, slaked lime, cream of tartar.
11. Prepare a chronological table of the following chemists who
may be regarded as the founders of modern chemistry : Boyle, Becher,
Stahl, Black, Cavendish, Priestley, Dalton, Davy, Faraday, Graham,
Liebig, Lavoisier, Berthollet, Proust, Gay-Lussac, Dumas, Scheele,
Stas, Berzelius.
12. Discuss the principles underlying the nomenclature of acids,
bases, and salts.
PROBLEMS. XXIX. (Review.)
1. How many grams of hydrogen may be obtained by the inter-
action of zinc and 196 gm. of pure sulphuric acid?
2. How much mercury will remain after obtaining 48 gm. of oxygen
by heating mercuric oxide?
3. What weight of potassium chlorate is necessary to produce
144 gm. of oxygen?
4. What weight of potassium chloride remains after obtaining
8 gm. of oxygen from potassium chlorate?
5. How much zinc is necessary to obtain 2 1. of hydrogen?
6. How many grams of oxygen can be obtained from 490 gm. ot
potassium chlorate?
7. If the density of ice is 0.92, what volume will a liter of water at
A° C. occupy when frozen?
8- The formula of crystallized gypsum is CaSO^ . 2 H^,0. What
volume of steam at 300° C. is liberated by heating 100 gm. of gypsum?
9. Ten liters of hydrogen at 30° C. are passed over hot copper
oxide. How much copper is set free, and what volume of steam at
300° C. is formed? (The equation is CuO + Hg = Cu -f H^.)
10. A quantity of carbon monoxide was passed through a red-hot
tube containing ferric oxide (Fe^,0;5). The resulting gas was absorbed
by caustic potash. The gain in weight of the caustic potash was 0.86
gm. What was the volume of the carbon monoxide at the standard
temperature and pressure?
CHAPTER XVIIL
SODIUM — POTASSIUM — COPPER — SILVER — MAGNESIUM
— CALCIUM — ZINC — MERCURY.
Metals. — Most of the elements hitherto studied are gases.
The few soHds — carbon, sulphur, phosphorus, and iodine —
bear almost no resemblance to the elements called metals.
The popular conception of a metal is that of a lustrous
solid, possessing certain characteristics, e.g., malleability or
ductility, tenacity, varying degrees of hardness, etc. It is
more satisfactory to make no sharp division into non-metals
and metals, but to regard all the elements as a series,
broken into related groups the members of which have a
marked gradation of propierties, some elements emphasizing
basic, some acid, and some both properties. If an element
emphasizes basic properties at any or all times, it is then a
metal. (See page 128.)
PROPERTIES OF SODIUM.
Reniefnber tJie precautions to be observed in using sodium.
(See page 89.)
Experiment 142. — {a) Examine a small piece of sodium, and
record its most obvious physical properties, e.g. color, luster, whether
hard or soft, etc.
{b) Repeat Exp. 42, and read its accompanying text.
(<;) Answer the following : —
(i) Is sodium heavier or lighter than water ?
(2) What properties suggest that it is a metal ?
(3) Is it harder or softer than most metals ?
279
28o Experimental Chemistry.
The Principal Compounds of Sodium are the chloride,
sulphate, carbonate, nitrate, and hydroxide.
GENERAL PROPERTIES OF SODIUM COMPOUNDS.
Experiment 143. — Recall (or determine again) the flame test,
solubility in water, and color of the most common sodium compounds.
Draw legitimate conclusions.
Sodium Chloride is the most abundant compound of
sodium. It is familiar to all under the name salt, or com-
mon salt.
PREPARATION OF PURE SODIUM CHLORIDE.
Experiment 144. — Grind 50 gm. of salt with 150 cc. of water in a
mortar, adding the salt in small portions. Fiher this saturated solution
into a dish or beaker, and conduct hydrochloric acid gas into it. Gen-
erate the gas by the method employed in Exp. 61 or Exp. 63. A funnel,
or calcium chloride tube, or a thistle tube, should be attached to the end of
the delivery tube so that it wall dip just below, the surface of the salt solu-
tion. As the operation proceeds pure sodium chloride separates, but the
magnesium chloride and other impurities remain in solution. When
considerable sodium chloride has separated, remove the generator, allow
the solid to settle, decant the clear liquid, add 15 cc. of cold water, and
decant again. If a filter pump is available, remove the last traces of
water with it ; if not, wasli and decant once or twice as before. Finally
dry the salt by heating in a porcelain dish ; stir constantly with a glass
rod during the heating. Describe it.
Preserve a specimen in a small tube or bottle.
Sodium Hydroxide is a useful compound of sodium.
Its properties have already been studied.
PREPARATION OF SODIUM HYDROXIDE.
Experiment 145. — Dissolve 25 gm. of sodium carbonate in 150 cc.
of water and heat gently in an iron dish (an ordinary iron spider is well
adapted for this work). Meanwhile slake 10 gm. of lime and add just
enough water td make a milky liquid — " milk of lime." Add the milk
Sodium. 281
of lime to the sodium carbonate solution and boil for several minutes,
stirring constantly with an iron wire or file. Let the precipitate subside,
remove a little liquid with a pipette or small tube, and if it effervesces
with hydrochloric acid, add more milk o^ lime and boil ; if not, decant
the clear liquid into a convenient vessel, let it stand for a few minutes
or until the solid settles, and siphon off the liquid through a glass
siphon. This solution of sodium hydroxide may be evaporated to
dryness and the solid product tested and the remainder preserved, or the
solution may be tested at once as follows : —
{a) Rub a httle between the fingers and describe the feeling.
{b) Add to a small portion a solution of zinc sulphate. Add an
excess of sodium hydroxide. These changes have already been
explained (Exp. 21). Reconsider them, and state what they prove
about this solution.
{c) Try the flame test.
{d) Determine the specific gravity by weighing a measured volume.
CLASS-ROOM EXERCISE. XLIX.
1. Additional study of sodium.
{a) Occurrence. {b) Discovery.
(c) Industrial preparation.
(^) Properties not experimentally shown. (<?) Uses.
(/) Derivation and significance of the words natriiini and sodium.
2. Oxides of sodium.
(^) Monoxide.
{b) Peroxide — preparation, properties, and uses.
3. Sodium hydroxide.
{a) Names and formula. {c) Additional properties.
{b^ Industrial preparation (<:/) Uses.
4. Sodium chloride.
{a) Sources. - {b) Purification.
{c) Properties not experimentally found.
{d) Uses. ie) Physiological value.
5. Sodium carbonate.
{a) Names.
{b^ Industrial preparation, including the most important equations
in each process. {c) Uses.
(^) Chemical name and formula of soda, washing soda, soda
crystals, alkali, soda ash.
282 Experimental Chemistry.
6. Sodium bicarbonate.
(a) Formula.
(d) Names and their significance.
(c) Properties and uses.
7. Other sodium compounds.
(a) Sodium sulphate — source, preparation, properties, uses,
historical significance.
(0) Sodium nitrate — names, source, properties, and uses.
8. Miscellaneous.
(a) Valence of sodium in NaOH, Na,SO^, HNaSO,, Na.O, Na..PO^
(sodium phosphate), H^,NaPO^ (dihydrogen sodium phos-
phate), HNa.PO^ (diso'dium phosphate), HNaCNHJPO^
(microcosmic salt) .
PROPERTIES OF POTASSIUM.
Observe the same preeautions as in the use of sodium.
Experiment 146. — (<?) Examine a very small piece of freshly cut
potassium, and record its most obvious physical properties. Touch it
slightly. Does it suggest caustic potash and soda?
{b^ Scrape the coating from a piece of potassium, press it between
filter paper to remove the oil, cut off a very small piece, and drop it on
the water in a pneumatic trough or similar .vessel. Stand just near
enough to see the action. Describe the action. How does it differ
from the action of sodium? Test the water as in Exp. 142.
From what has already been learned about sodium and potassium,
predict the main chemical change observed in {b) and write the equation.
{c) Answer the following : —
(i) Is potassium heavier or lighter than water?
(2) What properties suggest that it is a metal ?
(3) How does it reseml)le and differ from sodium?
(4) What is the usual test for potassium?
(rtf) If possible, examine the flame of a potassium compound {e.g.
potassium chloride) through a spectroscope, and describe the result.
The Principal Compounds of Potassium are the chlorate,
chloride, hydroxide, nitrate, carbonate, bromide, and iodide.
Potassium. 283
GENERAL PROPERTIES OF POTASSIUM COMPOUNDS.
Experiment 147. — Recall (or determine again) the solubility of
these compounds. Subject them to the flame test. How does each
color the flame? Draw a legitimate conclusion.
LABORATORY EXERCISE. XXV.
1. What experimental use has been made of potassium chlorate,
potassium bromide, and potassium iodide ?
2. What happens to potassium hydroxide and potassium carbonate
air
when exposed to the
3. What is the thermal effect produced by dissolving potassium
nitrate and potassium hydroxide separately in water?
Potassium Hydroxide has properties which are analogous
to those of sodium hydroxide. It is similarly prepared.
PREPARATION OF POTASSIUM HYDROXIDE.
Experiment 148. — Proceed as in Exp. 145, but use potassium car-
bonate instead of sodium carbonate. Test as in the case of sodium
hydroxide.
Potassium Carbonate is the historical potash, though that
term is now usually applied to potassium hydroxide. Some
of its properties have already been considered.
PREPARATION AND PROPERTIES OF POTASSIUM CARBONATE.
Experiment 149. — {a) Heat strongly 5 gm. of cream of tartar —
acid potassium tartrate — in an iron dish in the hood until the residue
is white. Grind this solid with water in a mortar, and filter. Test the
filtrate (i) with both kinds of litmus paper, (2) for potassium, and
(3) for a carbonate. Record the results.
ib) Fill a test tube half full of wood ashes, add half the volume of
water, shake, and warm gently. Filter, and test the filtrate as in {a).
If test (3) is not decisive, repeat the experiment on a larger scale.
Record the results.
284 Experimental Chemistry.
Answer the following : —
(i) What is the source of cream of tartar?
(2) What do (a) and (d) show about the distribution of potas-
sium? Of its assimilation by plants?
(3) What is the literal meaning of the word potash ?
{c) Mix 5 gm. of powdered potassium nitrate and i gm. of powdered
wood charcoal. Spread the mixture on an iron dish or pan, place the
pan on a tripod or ring, stand a lighted burner under it, and step aside.
Wait for a decided result. As soon as the mass is cool, grind It with
a little water in a mortar, filter, and test the filtrate as in {a). How has
the potassium nitrate been changed ?
CLASS-ROOM EXERCISE. L.
1. Additional study of potassium.
{a) Occurrence and sources.
{b) Discovery. {c) Industrial preparation.
{d) Properties not experimentally shown. {e) Uses.
(/) Derivation and significance of the words kaliion and
potassium.
2. Potassium hydroxide.
(«) Names and formula. {b) Industrial preparation.
{c) Explain the statement, '-the reaction in the preparation of
potassium hydroxide is reversible."
{d') Additional properties. {c^ Uses.
3. Halides of potassium.
{a) Sources, preparation, properties, and uses of each.
(J)) Their similarity.
4. Potassium chlorate.
{a) Industrial preparation. {b) Properties. {c) Uses.
5. Potassium carbonate.
(«) Sources. {c) Properties.
{b) Preparation (several methods). (^/) Uses.
(J) What is pearlash ? Potash ?
6. Potassium nitrate.
(«) Names and their significance.
(^) Sources. {d) Properties,
(r) Preparation. (e) Uses.
(/) Review the subject of gunpowder.
Potassium. 285
7. Soduim and potassium.
(a) How do they illustrate the Periodic Law ?
((^) "Their molecules are monatomic." Explain this statement.
(c) Why are they called " alkali metals '' ?
(d) Significance of their discovery by Davy.
8. Miscellaneous.
(a) What is "lye"?
(d) Valence of potassium in KOH, K^SO^, HKSO^, Kp, KgCOg,
KBr.
(c) Elements related to sodium and potassium — especially their
properties.
(d) The terms volatile, fixed, mineral, and vegetable, have all
b.^en applied to alkali. Explain each term.
{e) The spectroscope — its construction and use.
(/) Review " The Stassfurt Deposits " with special reference to
potassium compounds.
{g) Review "Fertilizer" from the standpoint of potash.
9. Ammonium compounds.
{a) Review the compounds previously mentioned or used.
{b) Why may they be considered with the "alkali" compounds?
PROBLEMS. XXX.
1. Stas found that 100 gm. of pure silver (dissolved in nitric acid)
required 54.2075 gm. of sodiur.v chloride for complete precipitation.
Calculate the atomic weight of sodium. (Assume atomic weights of
silver and chlorine and AgNOg + NaCl = AgCl + NaNOg.)
2. One hundred gm. of potassium chloride give 52.42 gm. of potas-
sium and 47.58 gm. of chlorine. Calculate the atomic weight of
potassium — making two assumptions.
3. How much sodium can be obtained from 2 kg. of sodium car-
bonate, if the latter is heated with carbon? The equation is —
Na^COg + C. = Na^ + 3 CO.
4. How much potassium carbonate is necessary to prepare a kilogram
of potassium hydroxide? The equation is —
K,C03 + Ca(OH),, = 2KOH + CaCO,.
5. How much sodium carbonate is necessary to prepare a kilogram
of sodium hydroxide ?
286 Experimental Chemistry.
6. What per cent of Glauber's salt, Na^,SO^.ioHp is sodium
sulphate ?
7. How many pounds of salt are needed to make one-fourth of a ton
of salt cake ? The equation is —
2 NaCl + H,,SO^ = NaoSO, + 2 HCl.
8. How much sodium carbonate can be made from 500 kg. of salt?
The first equation is given in problem 7, the other is —
Na._,S04 -f CaCO.5 + 2 C2 = Na,,CO.. + CaS + 4 CO.
9. How much sodium sulphide is formed by heating 100 gm. of
sodium sulphate with charcoal? The equation is —
Na^SO^ + 2 C2 = Na,S + 4 CO.
Copper is an abundant element, and its compounds are
numerous. Its properties, which are familiar, adapt the
metal to various uses.
GENERAL PROPERTIES OF COPPER.
Experiment 150. — (a) Examine several forms of copper — wire,
sheet, filings, etc. — and observe and state the most obvious physical
properties.
(d) Determine the specific gravity by weighing copper in air and then
in water, or by weighing it in air and comparing this weight with the
volume of water it displaces when immersed in a known volume in a
graduate.
LABORATORY EXERCISE. XXVL
1. Is copper a good conductor of heat? Of electricity? On what
evidence is your answer based?
2. Is copper ductile? Malleable? Brittle? Tough? Hard or soft?
3. What happens to copper when heated? When exposed to the air?
TESTS FOR COPPER.
Experiment 151.— (^?) Heat a copper wire in the Bunsen flame.
The color is characteristic of copper and its compounds, though not a
conclusive test, since the same color is produced by other substances.
(^) Add a few drops of ammonium hydroxide to copper sulphate
solution, and observe the result; now add an excess of ammonium
Copper. 287
hydroxide. The bhie solution is a characteristic and decisive test for
copper.
(c) Add to a test tube half full of water a drop of copper sulphate
solution and shake ; then add a few drops of acetic acid and of
potassium ferrocyanide solution. The brown precipitate is copper
ferrocyanide.
Copper replaces some metals, if suspended in solutions
of their compounds, and is in turn replaced by other metals.
Both operations precipitate or deposit a metal.
INTERACTION OF METALS COPPER.
Experiment 152. — (a) Put a clean copper wire into a solution
of any compound of mercury. After a short time, remove the wire and
wipe it with a soft cloth or paper. Describe the change. What has
become of some of the copper?
(d) Put in separate test tubes half full of copper sulphate solution a
bright nail and a strip of clean zinc. After a short time remove the
metals and examine them. What is the deposit? What has become
of some of the zinc and iron? Does the final color of the solution
indicate any chemical change ? How would you prove the answer to the
last question?
The Important Compounds of Copper are the copper oxides,
copper sulphate, and copper nitrate.
LABORATORY EXERCISE. XXVIL
1. What experimental use has been made of one copper oxide?
2. Examine a crystal of copper sulphate and observe its most
obvious properties. What happens to it when heated?
3. How has copper nitrate been prepared? Summarize its proper-
ties. What is its formula?
4. Prove that each of the above important compounds contains copper.
5. What is the color of many copper compounds?
Oxides of Copper. — There are two oxides of copper, the
black oxide, already studied (see Exps. 9 and 151 (a)),
and the red oxide. The formulas are respectively CuO
and CugO, and their exact names are cupric and cuprous
288 Experimental Chemistry.
oxides. These oxides are the first members of two classes
of copper compounds — the -ic and -ous. In the former
the metal is combined with a larger proportion of oxygen
than in the latter, or is in a higher state of oxidation —
** farther on in the series." Other metals form similar
classes.
PREPARATION AND PROPERTIES OF CUPROUS OXIDE.
Experiment 153. — Dissolve i gm. of copper sulphate in 15 cc. of
water, and 5 gm. of Rochelle salt (sodium potassium tartrate) in 10 cc.
of water; mix these solutions, warm, and filter. Add to this filtrate
about 5 cc. of a dilute solution of grape sugar (or of cane suga; which
has been boiled at least five minutes with dilute sulphuric acid), and
enough sodium hydroxide solution to dissolve any precipitate. The
solution must be strongly alkaline. Boil until a decided change is
produced. The j^roduct is cuprous oxide. Describe it. Show by a
comparative table the general properties of the two oxides of copper.
The mixture of copper sulphate, Rochelle salt, and sodium
hydroxide is called Fehhng's solution, and is widely used to
detect the presence of sugar.
Alloys are mixtures or compounds of two or more metals.
The properties of the alloy vary with the constituents and
their proportions. If mercury is one constituent, the al-
loy is called an amalgam. The important alloys in which
copper is the larger constituent are brass, bronze, German
silver, and their allied alloys.
CLASS-ROOM EXERCISE. LI.
1. Additional study of copper.
{a) Occurrence of native and combined copper.
{p) Preparation from its ores. {d') Uses.
(c) Additional properties. (<?) History.
(/) What is electrolytic copper and how is it prepared.'*
2. Oxides.
{a) Review the preparation and properties of cuprous and cupric
, oxides.
(Jy) Significance of the terminations -ous and -ic.
Copper. 289
3. Copper sulphate.
(rt) Names. (c) Properties.
(d) Preparation. (d) Uses.
4.. Copper sulphides.
(a) Occurrence. (^) Properties. (c) Names and formulas.
j. Copper compounds.
(a) Give the chemical name and symbol of ruby ore, copper
pyrites, copper glance.
(d) Give formula and color of azurite and malachite.
6. Miscellaneous.
(a) What is the valence of copper in CuO, Cup, CuSO^,
Cu(N03)2, Cu.S?
(d) Derivation of the word ciiprinn.
7. Alloys of copper.
{a) Composition, preparation, properties, and uses of common
brass.
(J)) Alloys allied to brass.
{c) Composition, properties, and uses of (i) gun metal, (2) bell
metal, (3) speculum metal, (4) antique bronze, (5)
aluminium bronze.
(^/) Composition, properties, and uses of (i) ordinary German
silver, (2) Chinese Pakfong, (3) white German silver,
(4) casting German silver.
8. Experimental.
{a) Starting with copper, how would you prepare successively
copper nitrate, cupric oxide, and cuprous oxide ?
(Jj) Give the steps for the transformation of copper into copper
sulphate and cuprous oxide?
PROBLEIVIS. XXXI.
1. When 1.3305 gm. of copper were heated in oxygen, 1.6675 gm. of
cupric oxide were formed. Calculate the atomic w^eight of copper.
(The equation is Cu + O = CuO.)
2. How much cupric oxide is formed by heating 1467 gm. of copper
in air?
3. How much cuprous oxide is formed by boiling 100 gm. of copper
sulphate with potassium hydroxide and grape sugar? (The equation is
2CuSO,.5H.O + 4KOH - Cup + O + I2H,0 + 2 K,SO,.)
290 Experimental Chemistry. •
4. Calculate the percentage composition of —
(a) Malachite, CuC03.Cu(OH)2.
(d) Azurite, 2CuCO..Cu(OH)2.
Silver is familiar as the alloy of silver in coins. Two
of its compounds, silver nitrate and silver chloride, have
already been considered in testing for chlorine in the form
of chlorides.
PREPARATION OF SILVER.
Experiment 154. — Prepare silver by one or both of the following
methods : —
(a) Fill a small dish (e.^i^. a porcelain crucible) half full of silver
nitrate solution, and add a few drops of mercury. Allow the whole to
stand undisturbed for a day or more, and then examine. The delicate
crystals attached to the mercury are silver. Pick them out with the
forceps, wash well with water, and preserve them.
(d) Dissolve a ten-cent piece in 10 cc. of concentrated nitric acid,
dilute with a little water, and add hydrochloric acid until the precipita-
tion is complete. Let the precipitate settle, and wash several times by
decantation. Then filter and wash until the filtrate is neutral. If con-
venient, let the precipitate dry ; if not, scrape half from the opened
paper with a knife or spatula, put it in a shallow dish, cover with dilute
sulphuric acid, and add a piece of zinc; put the other half in a cavity
at the end of a piece of charcoal, cover with sodium carbonate, and re-
duce it with a blowpipe flame. In the first case, the silver will collect
as a grayish powder ; remove any excess of zinc, filter, wash with water
and dry the residue. It may be preserved as a powder, or fused into a
bead with a blowpipe flame. In the second case, minute globules of
silver will appear on the charcoal ; scrape them together and fuse into
a single bead.
PROPERTIES OF SILVER.
Experiment 155. — Examine the silver formed in Exp. 154, and
state briefly its most obvious properties.
TEST FOR SILVER.
Experiment 156. — Devise a test for combined silver, based upon
previous experiments. Verify it.
Silver. 291
LABORATORY EXERCISE. XXVIII.
1. What caused the blue color of the filtrate in Exp. 154 (d) ?
2. Name the haloid compounds of silver. State their resemblances
and differences.
3. Explain the chemical change in the formation of silver in the first
casein Exp. 154 (d). Does it differ fundamentally from the change in
the second case?
4. How may silver sulphide be formed? What is its color? What
happens chemically when silver "tarnishes"?
CLASS-ROOM EXERCISE. LIL
1 . Additional study of silver.
(a) Occurrence of native and combined silver. What is argen-
tiferous lead? '
(d) The various industrial methods of preparation.
(c) Properties.
(d) Uses.
(e) History.
(/) Derivation of the word argentum.
2. Halides of silver.
{a) Occurrence, properties, and uses of silver chloride.
(Jf) Preparation and uses of silver bromide and iodide.
(6) Properties common to this group.
3. Silver nitrate.
{a) Names.
{b) Preparation.
(6-) Properties.
\d) Uses.
4. Miscellaneous.
(«) Describe the process of silver plating.
{b^ What is silver glance ? Horn silver ? Lunar caustic ?
{c) Compare copper and silver in the light of the Periodic Law.
(rt') In what respect does gold resemble copper and silver ?
5. Alloys of silver.
{a) Composition of standard silver coin.
((^) What is sterling silver ?
{c) Why are coins stamped and not cast ?
292 Experimental Chemistry.
PROBLEMS. XXXII.
1. Calculate the percentage composition of —
(a) Silver chloride, AgCl. (d) Silver sulphide, AggS.
2. If 53.1958 gm. of silver form 92.6042 gm. of silver bromide, AgBr,
what is the atomic weight of silver ?
3. If I gm. of silver chloride yields 0.7526 gm. of silver, what is the
atomic weight of silver ?
4. If 480 gm. of silver interact with nitric acid, how much silver
nitrate is formed .''
Magnesium. — The conduct of magnesium has been
studied when the element was heated in air, in oxygen,
and in ammonia, and when it was treated with hydro-
chloric acid.
LABORATORY EXERCISE. XXIX.
1. What happens to magnesium when it is heated in air, in oxygen,
and in ammonia gas ?
2. What is the equivalent of magnesium ? The atomic weight ?
Summarize the experimental evidence on which the answers are based,
GENERAL PROPERTIES OF MAGNESIUM.
Experiment 157. — (a) Examine a piece of magnesium, and write
a short account of its most obvious physical properties.
(d) If a sufficient quantity is available, determine its specific gravity.
Weigh it first in air, then in water, and compute the specific gravity
from these data. Compare the result with the specific gravity ot other
metals.
Magnesium Oxide is the only important compound of
magnesium previously considered.
LABORATORY EXERCISE. XXX.
1 . When magnesium is burned in air, what compound is formed ?
What is its formula ?
2. What is magnesia ? State its observed properties. Predict an
important application of magnesia.
Magnesium. 293
TESTS FOR MAGNESIUM.
Experiment 158. — (a) To a solution of magnesium sulphate or
chloride add successively solutions of ammonium chloride, ammonium
hydroxide, and disodium phosphate. A precipitate of ammonium mag-
nesium phosphate is formed, NH^MgPO^. It is voluminous at first,
but finally crystalline. It is soluble in acids. Try it.
(d) Put a litde powdered magnesium carbonate or oxide in a cavity
at the end ot a piece of charcoal, moisten with water, and heat intensely
in a blowpipe flame. Cool, and moisten with a drop of cobaltous nitrate
solution. Heat again, and when cool observe the color. If the experi-
ment has been conducted properly, a pink or pale flesh-colored residue
coats the charcoal.
CLASS-ROOM EXERCISE. LIII.
I. Additional study of magnesium.
(a) Occurrence of its compounds.
(d) Preparation.
(c) Additional properties.
(d) Uses.
(e) How was magnesium utilized in the discovery of argon ?
2- Magnesium oxide.
(a) Names. (c) Properties.
(d) Preparation. ((i) Uses.
(e) What is magnesia mixture ? Calcined magnesia ? Fluid
magnesia ? Magnesium citrate ?
3. Magnesium sulphate.
Ca) Names. -^ (c) Properties.
(d) Preparation. (^) Use.
4. Miscellaneous.
(a) Chemical name and formula of inagnesia alba, magnesite,
brucite, periclase, dolomite.
(J?) Valence of magnesium in MgO, MgoCO.., MgClo, Mg(OH)..
{c) Review " The Stassfurt Deposits '' with reference to magne-
sium compounds.
5. Black's work on Magnesia alba.
294 Experimental Chemistry.
PROBLEMS. XXXIII.
1. What volume of nitrogen at 20° C. and 780 mm. will combine
with 50 gm. of magnesium ? (Equation is 3 Mg + N^ = Mg.jN2.)
2. How many grams of magnesium will be formed by heating 100
gm. of potassium with magnesium chloride .-* (Equation is Kg +
MgCl2=Mg + 2 KCl.)
3. How much sodium is necessary to decompose a kilogram of mag-
nesium chloride, and how much magnesium will be formed ? (Equation
is similar to 2.)
4. How much magnesium can be obtained from 100 gm. of (a)
magnesite, MgCO^, (d) dolomite, MgCa^^COa)^, (c) Epsom salts,
MgSO, . 7 H,0 ?
Calcium is an uncommon metallic element, but its com-
pounds are exceedingly widespread and important.
TESTS FOR CALCIUM.
Experiment 159. — (a) Subject calcium chloride to the flame test.
Record the result.
(d) Repeat Exp. 29 (d).
(c) If possible, examine a calcium flame through a spectroscope.
Compare the result with the position of the potassium lines.
The Principal Compounds of Calcium are the hydroxide,
oxide, carbonate, chloride, sulphate, fluoride, and hypo-
chlorite (bleaching powder).
LABORATORY EXERCISE. XXXL
1. What experimental use has been made of calcium hydroxide,
calcium chloride, calcium fluoride, and calcium hypochlorite .''
2. What happens to calcium oxide and calcium chloride when
exposed to the air ?. What thermal eff"ect is produced by dissolving
each in water ?
3. Give the common name of each of the above compounds (except
calcium chloride).
4. How is lime water prepared ? How should it be kept ?
5. What happens to crystallized calcium sulphate (gypsum) when
heated ? To calcium carbonate when treated with acids ?
Calcium. 295
6. Examine specimens of selenite, gypsum, and plaster of Paris.
Describe them. Recall or repeat the experimental study of the solu-
bility of calcium sulphate. Mix a little plaster of Paris with enough
water on a block of wood to form a thin paste. Let it stand undisturbed
for ten or fifteen minutes, and then examine. Describe the change.
What is meant by the "set" of plaster of Paris?
7. Starting with limestone, how would you prepare lime, lime water,
milk of lime, and chloride of lime?
8. Calcium sulphate is nearly insoluble in water ; how can it be
proved experimentally that it is a sulphate?
CLASS-ROOM EXERCISE. LIV.
1. Additional study of calcium.
{a) Occurrence of combined calcium.
{h) Discovery. {c^ Properties.
2. Calcium oxide.
{a) Names. {c) Properties.
{b) Industrial preparation. {d) Uses.
{e) Explain and illustrate the statement, "the term hme is
ambiguous."
3. Calcium hydroxide.
{a) Names.
{b) Properties of actual substance.
{c) Uses.
(^) What is "milk of lime"? Mortar? Cement? Plaster?
Slaked lime? Explain the hardening of mortar, cement,
and plaster.
4. Calcium carbonate.
{a) Names. {c) Crystallization.
{b) Occurrence. ^ {d) Properties.
(e) Uses of various varieties.
if) What are stalactites and stalagmites? How are they formed?
5. Calcium chloride.
{a) Natural sources. {c) Properties.
{b) Preparation, {d) Uses.
6. Calcium fluoride.
(a) Occurrence. {c) Properties.
{b) Crystallization. {d) Uses.
296 Experimental Chemistry.
7. Calcium sulphate.
(a) Names. (c) Properties.
(d) Occurrence. (^) Uses.
(e) Plaster of Paris — source, properties, and uses. Why so
called ?
8. Calcium hypochlorite.
(a) Review the subject as considered under chlorine
((^) Formula.
(c) What is " bleach " ? '' Chloride of lime " ?
9. Calcium phosphates.
(rt) Review the subject as considered under phosphorus.
10. Miscellaneous.
(a) Elements related to magnesium and calcium.
(d) Review "Hardness of water." (See Class-Room Exercise
IX, 4 (c) and page 248.)
(c) What is the chemical name of Iceland spar? Of satin spar?
Of selenite?
(^) What is the "lime" light?
(e) Valence of calcium in CaO, CaCL, CaCOo, Ca(0H)2,
CaSO^.
(/) Discuss magnesium and calcium in the light of the Periodic
Law.
PROBLEMS. XXXIV.
1. How much calcium may be obtained from 100 gm. of (a) marble,
(d) gypsum, CaS0^.2Hp, (0 Auor spar, CaFg, (d) superphosphate
oflime, CaH,(PO,),?
2. How many tons of limestone must be heated to form 100 tons
of quicklime? The equation for the reaction is CaCOg = CO., + CaO.
3. Find the simplest formula of a compound having the percentage
composition Ca — 40, C = 12, O = 48.
4. If 100 gm. of calcium carbonate yield 56 gm. of calcium oxide,
what is the atomic weight of calcium ?
5. How many grams does a pound of selenite lose when heated to
constant weight? The equation is CaSO^ . 2 H,0 = CaSO^ + 2 H^O.
Zinc is a familiar metal, and the compounds, zinc sul-
phide and zinc sulphate, have been experimentally studied.
Zi
nc.
297
GENERAL PROPERTIES OF ZINC.
Experiment 160. — (a) Examine a piece of zinc and record its
most obvious properties.
(3) Determine the specific gravity by the method previously de-
scribed. (See Exp. 150 (d).) Compare the result with the specific
gravity of other metals.
LABORATORY EXERCISE. XXXII.
1. What happens to zinc when heated? Describe and name the
product. When it is exposed to air?
2. Is zinc hard or soft? Malleable? Ductile? Brittle? Tough?
Does it melt easily?
TESTS FOR ZINC.
Experiment 161. — (a) Recall or devise a simple test for combined
zinc.
(d) Recall or repeat the action of zinc when heated in the oxidizing
flame.
(c) Fill a small cavity at one end of a piece of charcoal with zinc
oxide, moisten with water, and heat strongly in the blowpipe flame.
Cool, and moisten with a drop of cobaltous nitrate solution, then heat
again. Cool and examine. A green incrustation is caused by zinc
compounds.
Zinc interacts with other metals similarly to copper.
INTERACTION OF METALS ZINC.
Experiment 162. — (a) Put a strip of clean zinc in copper sul-
phate solution, as in Exp. 152. Describe and explain the result which
soon occurs.
((^) Put a strip of clean zinc in lead nitrate or lead acetate solution,
and describe and explain the resuli
(c) Clean a strip of zinc and put it in a solution of any salt of mer-
cury. Examine after a short time. What is amalgamated zinc, and
for what is it used ?
LABORATORY EXERCISE. XXXIII.
I . What happens when zinc is added to sulphuric acid ? Describe
the purified solid product. Illustrate by an equation the chemical
change.
298 Experimental Chemistry.
2. What happens when a little sodium hydroxide solution is added
to a zinc compound (in solution)? When an excess is added? Name
all the compounds of zinc formed in these changes.
3. What happens when a sulphide, e.g. hydrogen sulphide or am-
monium sulphide, is added to a zinc compound (in solution) ? What
characteristic property has the solid product? If the zinc solution were
acid (instead of neutral), how would the result be changed?
4. How may zinc be transformed into zinc nitrate and then into
zinc oxide?
CLASS-ROOM EXERCISE. LV.
1. Additional study of zinc.
{a) Occurrence of combined zinc. {c) Properties.
{b) Industrial preparation. {d) Uses.
2. Compounds.
{a) Chemical name and formula of calamine, zinc blende,
gahnite, zinc white, white vitriol, sphalerite.
3. Alloys.
{a) What alloys contain zinc?
4. Miscellaneous.
{a) What is galvanized iron?
{b) Valence of zinc in ZnSO^, ZnCU, ZnO, ZnS.
PROBLEMS. XXXV.
1. Review Problems i, 2, 4, and 6 in Chapter III., page 52.
2. Review Problems 2, 3, 4, 1 1, 12, and 15 in Chapter IV., page 62.
3. Review Problem 3 in Chapter IX., page 154.
Mercury is the only metallic element which is liquid at
the ordinary temperature, and it is more or less famiUar
from its extensive use in thermometers and barometers.
GENERAL PROPERTIES OF MERCURY.
Merauy and its compounds arc poisonous.
Experiment 163. — {a) Examine some mercury, and state its
characteristic physical properties. Pour a drop or two into a large
shallow dish, and agitate the dish ; does the result suggest the reason
for the name " quicksilver " ?
Mercury. 299
(fi) Lift a bottle of mercury. Estimate its specific gravity. Verify
the estimate by consulting a book, or by weighing a measured small
quantity, as in the case of sulphuric acid (Exp. 4, Lab. Ex. xxiii.,
page 271).
Mercury is prepared from its compounds by roasting
them alone, or with sodium carbonate or Hme. Cinnabar,
mercuric sulphide, is the natural compound used in the
industrial preparation of mercury.
PREPARATION OF MERCURY.
Perform in the Hood.
Experiment 164. — {a) Recall or repeat the preparation of mercury
from mercuric oxide.
{b) Put a little vermilion or powdered cinnabar near one end of
a glass tube open at both ends, and heat slowly at first, but finally
strongly where the powder is located. The tube should be held at a
slight angle to the flame and rotated. Fumes of sulphur dioxide (and
possibly of mercury) will escape from the upper end of the tube, and
mercury will be deposited in the colder portion somewhat as in {a).
((f) Mix any dry salt of mercury, e.g. mercuric chloride, with three
times its weight of dry sodium carbonate, and heat strongly in an
ignition tube. A test tube may be used, but it is liable to crack or
melt. Mercury will be deposited on the upper part of the tube.
Definitions. — The layer of mercury obtained in Exp. 164
is a sublimate. The operation is called sublimation, or
subliming. The preparation of iodine illustrated sublima-
tion. A sublimate is usually purer than its source.
TESTS FOR MERCURY.
Experiment 165. — {a) What is the simplest test for free mercury ?
{b) Recall or devise a test for combined mercury. Verify it.
Mercuric Oxide is the only compound of mercury pre-
viously studied.
jOO Experimental Chemistry.
LABORATORY EXERCISE. XXXIV.
1. Describe briefly but accurately the effect of heat on mercuric
oxide. Write the equation for the change.
2. What historical interest has this compound ?
Other Important Compounds of Mercury belong to the two
series ?nalogous to the copper compounds previously de-
scribed, viz., mercurous and mercuric nitrate and chloride.
PROPERTIES OF MERCUROUS AND MERCURIC COMPOUNDS.
Experiment i66. — (a) Merairoits. Add a few drops of hydro-
chloric acid to a little mercurous nitrate solution. The white pre-
cipitate is mercurous chloride. Note its insolubility in water and in
dilute hydrochloric acid. Add a few drops of ammonium hydroxide.
The black precipitate is mainly mercurous ammonium chloride. Its
formation is a delicate test for mercury in mercurous compounds.
ib) Mercuric. Add a few drops of hydrochloric acid to a little
mercuric nitrate solution. Compare the result with that in {a). Add
a few drops of ammonium hydroxide, or enough to produce a decided
change. Compare with {a). The precipitate is mercuric ammonium
chloride.
The formula of mercurous nitrate is Hg./N03)2, and of
mercurous chloride is Hg2Cl2; some authorities write these
formulas HgNOg and HgCl. The formula of mercuric
nitrate is Hg(N03)2, and of mercuric chloride is HgCl2.
These compounds are analogous to the corresponding -ohs
and -ic oxides. Thus,
Mercuric oxide is HgO
Mercuric chloride is HgCla
Mercuric nitrate is Hg(N03)2
Mercurous oxide is HggO
Mercurous chloride is ... . Hg2Cl2
Mercurous nitrate is , . . . Hg2(N03)2.
Mercury. 301
CLASS-ROOM EXERCISE. LVI.
1. Additional study of mercury.
(a) Occurrence of free and combined mercury.
(d) Industrial preparation.
{c) Impurities, and how removed.
(d) Properties.
(e) Uses.
(/) Significance of the common name.
(£■) History.
2. Mercuric oxide.
(a) Preparation of the varieties.
(<^) Names.
(c) Historical significance.
3. Mercurous compounds.
(rt) Preparation, properties, uses, formulas^ and common name
of mercurous chloride.
{3) Other mercurous compounds.
(^) Valence of mercury in mercurous compounds.
4. Mercuric compounds.
(a) Preparation, properties, uses, formula, and common name
of mercuric chloride.
(d) Other mercuric compounds.
(c) Valence of mercury in mercuric compounds.
5. Miscellaneous.
(a) Explain the statement, " the vapor of mercury and of zinc
consists of monatomic molecules."
(d) Discuss the relation of zinc to mercury, and of both metals
to magnesium and calcium in the light of the Periodic
Law.
(c) Literal meaning of word " hydrargyrum."
6. Amalgams.
(a) Definition of an amalgam.
(d) Preparation, properties, and uses of sodium amalgam.
(c) What is amalgamated zinc, and for what is it used ?
(d) For what do dentists use amalgams ?
(e) How is gold often extracted from its ores ?
302 Experimental Chemistry.
PROBLEMS. XXXVI.
1. If 195 cc. of oxygen at 37° C. and 620 mm. are obtained by heat-
ing 2.7 gm. of mercuric oxide, what is the atomic weight of mercury ?
2. How much mercury is formed by heating 400 gm. of pure cinna-
bar ? (Equation is HgS + O^ = Hg + SO^.)
3. If 70 gm. of mercuric oxide are heated, what volume of oxygen
at 21'' C. and 740 mm. is evolved ?
4. If 177.1664 gm. of mercuric sulphide yield 152.745 gm. of mer-
cury, what is the atomic weight of mercury ?
CHAPTER XIX.
ALUMINIUM — TIN — LEAD — CHROMIUM— MANGANESE -
IRON — SEPARATION OF METALS.
Aluminium, owing to its extensive preparation by elec-
trolytic methods, is a familiar metal. Alum has long been
known.
GENERAL PROPERTIES OF ALUMINIUM.
Experiment 167. — (a) Examine a piece of aluminium (sheet or
wire), and observe its physical properties. Has it any "spring 'Mike
brass ? Is it ductile, malleable, soft, hard, tough, brittle ?
(d) Compare roughly the weight of a piece of aluminium with a
piece of zinc, pasteboard, and glass having approximately the same
volume.
If sufficient aluminium is available, determine the specific gravity.
(See Exp. 150 ((^).) Compare the result with' the specific gravity of
other metals.
It has been asserted that aluminium is not acted upon
by acids and alkalies.
ACTION OF ALUMINIUM WITH ACIDS AND ALKALIES.
Experiment 168. — (a) Add a small piece of aluminium to separate
test tubes containing dilute sulphuric acid and dilute hydrochloric acid.
Warm, if necessary. Describe the action. Test the gas evolved. What
compound is formed in each case ?
(d) Add a small piece of aluminium to a test tube half full of a
dilute solution of sodium hydroxide, and boil. Test any gas evolved.
If only a little gas is liberated, attach a simple delivery tube and collect
the gas over water.
Other acids and alkalies act similarly ; draw a general conclusion
from this experiment. Is the aluminium examined in Exp. 167 tar-
303
304 Experimental Chemistry.
nished ? Scrape the surface, and examine again. Does aluminium
differ from most metals in this respect ?
Aluminium Hydroxide is an important compound of
aluminium.
PREPARATION AND PROPERTIES OF ALUMINIUM HYDROXIDE.
Experiment 169. — {a) Add slowly a little potassium hydroxide or
sodium hydroxide solution to a test tube half full of alum solution.
The gelatinous precipitate is aluminium hydroxide. Now add an ex-
cess of the alkali to one half, and dilute, hydrochloric acid to the other.
Describe the results.
{b) Recall or devise another simple method of preparing aluminium
hydroxide.
{c) Add a little solution of ammonium sulphide to a solution of
alum. The precipitate is not a sulphide, but aluminium hydroxide,
because aluminium forms no sulphide in the wet way.
{d) Add a little alum solution to a dilute solution of cochineal, then
add ammonium hydroxide. The colored product is called carmine
lake. It belongs to a class of dyes formed by the combination of a
vegetable dye and a metallic hydroxide, usually aluminium hydroxide.
Discussion of Experiment 169. — When sodium hydroxide
or potassium hydroxide is first added to a solution of a
compound of aluminium, aluminium hydroxide is formed.
The simplest equation for this reaction is —
Al2(S04)3 + 6K0H = 2Al(OH)3 + 3 K2SO4.
Aluminium Aluminium
Sulphate Hydroxide
An excess of alkali produces a chemical change repre-
sented by the equation : —
A1(0H)3 -f 3KOH = A1(0K)3 + 3H2O
Potassium
Aluminate
Aluminium. 305
The chemical changes are similar to the changes of zinc
compounds under the same conditions. (See Exp. 21.)
TESTS FOR ALUMINIUM.
Experiment i*jo.— (a) What is a simple test for metallic alu-
minium?
(/;) Recall or devise a test for combined aluminium. Verify it.
How can aluminium compounds be distinguished from those of zinc?
(c) Heat a little aluminium sulphate or aluminium hydroxide on
charcoal in the blowpipe flame. Cool and moisten with a drop of co-
baltous nitrate solution. Heat again, and if the operation has been
conducted properly, a blue residue will coat the charcoal. This color
is characteristic of aluminium compounds. Compare this result with
the action of other metallic compounds under similar circumstances.
Alum is by far the most useful compound of aluminium.
PREPARATION AND PROPERTIES OF COMMON ALUM.
Experiment 171. — (a) Dissolve about 10 gm. of aluminium sul-
phate in the least possible amount of hot water. Dissolve 3 gm. of
potassium sulphate in the same way. Mix the clear, hot, saturated
solutions in a small shallow dish, and allow the solution to cool undis-
turbed. Crystals of potassium alum will be deposited. Remove the
best ones ; dry, and examine. Describe them, giving color, luster, size,
and crystal form.
(d) Prove by actual tests that (i) they are a sulphate of potassium
and of aluminium, and that (2) they contain water of crystallization.
CLASS-ROOM EXERCISE. LVIL
1. Additional study of aluminium. "^
(a) Occurrence.
(d) Proportion in the earth's crust. (^/) Properties.
(c) Industrial preparation. (e) Uses.
(/) History.
(g) Derivation of the words aluminiDn and al/n/iniium.
2. Alums.
{a) General formula. {c) Properties of alums.
{b) Varieties. {d) Potassium alum. {e) Uses.
3o6 Experimental Chemistry.
Compounds.
(^?) Chemical name and formula of corundum, ruby, sapphire,
cryolite, turquoise, bauxite, alumina, feldspar.
Miscellaneous.
(a) Valence of aluminium in AljO.^ and Al._, (804)3.
(d) What is burnt alum, a mordant, red liquor?
(c) What elements are closely related to aluminium?
(d) What is aluminium bronze? Uses.
Complete the following equations (note that one side is already
complete) : —
(a) Al + O = AloO,
(d) Al + H,SO, = A1„(S0,), + 3H2
(0 Al,(SO,), + KOH = 2A1(0H)3 + 3 KgSO^
PROBLEMS. XXXVII.
How much aluminium can be obtained theoretically from 100 gm.
of-
(a) ALOo-aSiOo?
(I?) Cryolite, AlNagF^j?
(c) Turquoise, Al^.p^ . H^AiPe . 2 H,,0 ?
2. How much aluminium can be obtained by the interaction of a
kilogram of sodium aluminium chloride and sodium? (Equation is
AlCL.NaCl + 3Na = Al + 4 NaCl.)
3. What volume of oxygen at I5°C. is needed to burn 5 gm. of
aluminium to Al^Og?
4. Half a gram of aluminium when warmed with potassium hydroxide,
liberated 660 cc of hydrogen at 13° C. (over water). Calculate the atomic
weight of aluminium. (Equation is Al, + 2 KOII + 2 Hp =
2 KAIO., + 3 H,.)
5. Review Problems 5, 6, 7, 8, 9 in Chapter IX, page 154.
Tin is a familiar metal owing to its extensive use as a
protective coating for iron. Its compounds are not nu-
merous, though one, stannous chloride, is widely used in
dyeing.
Tin. 307
GENERAL PROPERTIES OF TIN.
Experiment 172. — {a) Examine a stick of tin or a piece of pure
tinfoil, and state the most obvious physical properties. Bend a stick
c^ tin, and note the crackling sound.
{b) If a stick of tin or a piece of block tin pipe is available, deter-
mine the specific gravity by either of the methods described in Exp.
150 (b). Compare the result with the specific gravity of other metals.
ACTION OF TIN WITH ACIDS.
Experiment 173. — {(i) Put a small piece of tin — about .5 gm. —
in a test tube, cover with concentrated hydrochloric acid, add a little
water, and heat — in the hood. Heat gently at first, and when action
begins regulate the heat accordingly. Most of the tin disappears,
soluble stannous chloride being formed. Save this solution for Exp.
174 (^).
{b) Treat a small piece of tin with concentrated nitric acid — in the
hood. It is advisable to stand the test tube in the rack or in a bottle as
soon as the action begins. The white, amorphous product is metastan-
nic acid. How does the action of nitric acid on tin differ from and
resemble its action on other metals, zinc, for example?
Stannous chloride has the formula SnCl2, but the formula
of metastannic acid varies with the temperature at which
it is dried.
TESTS FOR TIN.
Experiment 174. — {ci) What is a simple test for metallic tin ?
{b) Recall or repeat the action of tin when heated in a blowpipe
flame.
{c) Add a few drops of mercuric chloride solution {poison) to a
small portion of the stannous chloride solution prepared in Exp. 173 {a).
The white precipitate is mercurous chloride. Add a little more stannous
chloride solution and heat gently. The mercurous chloride is reduced
finally to mercury, which appears as a grayish powder.
Discussion of Experiment 174. — Stannous compounds
readily form stannic compounds, i.e. by an extension of
the conception of oxidation and reduction, they reduce
3o8 Experimental Chemistry.
other compounds and become oxidized, though no oxygen
may be involved. Thus the simplest equation for the
interaction of stannous and mercuric chlorides is —
2 HgCl2 + SnCla = Hg2Cl2 + SnCl^
Mercuric Stannous Mercurous Stannic
Chloride Chloride Cliloride Chloride
The mercuric chloride is reduced to mercurous chloride,
while the stannous chloride is oxidized to stannic chloride.
Reduction in its broadest sense means " going lower," and
oxidation "going higher" in a series of compounds. Stan-
nous chloride is often used as a reducing agent.
DEPOSITION OF METALLIC TIN.
Experiment 175. — Put a strip of zinc in a slightly acid solution of
stannous chloride. Examine after a short time, and the tin will be
found adhering to the zinc as a grayish black deposit; sometimes it
appears as minute scales. What becomes of the zinc ?
CLASS-ROOM EXERCISE. LVIII.
1. Additional study of tin.
(a) Occurrence of combined tin.
.{d) Industrial preparation. (^/) Uses.
(c) Properties. (e) History.
2. Compounds.
(a) Chemical name and formula of cassiterite, tin crystals, strean:
tin, tin dioxide, oxymuriate of tin.
3. Miscellaneous.
(«) What is tin plate ? "Tin"? Block tin ?
(d) Valence of tin in SnCU, SnCl^, SnOg, SnS.
4. Name the important alloys of tin.
5. Complete the equations : —
(a) Sn + = SnO..
(^) SnCl^ + = SnCl^.
(0 SnCl^ + = SnCL + 2 HCl
Lead. 309
PROBLEMS. XXXVIII.
1. Calculate the percentage composition of —
(a) Tinstone, SnO^,.
{d) Stannous chloride, SnCl^.
(c) Stannic chloride, SnCl^.
2. If 100 gm. of tin are heated with nitric acid, and the stannic
oxide formed on heating the product weighs 127. i gm., what is the
atomic weight of tin ?
Lead, both free and combined, is well known, and has
numerous applications. Its principal compounds are the
oxides, sulphide, sulphate, chromate, nitrate, and carbonate.
GENERAL PROPERTIES OF LEAD.
Experiment 176. — (a) Examine a piece of freshly cut lead and
state its most obvious physical properties.
(d) Determine its specific gravity by either method described in
Exp. 150 (d). Compare the result with the specific gravity of other
metals.
LABORATORY EXERCISE. XXXV.
1. What happens to lead when heated ? When exposed to the air?
Is it easily melted and tarnished ?
2. What physical properties adapt it for its extensive use ?
3. Draw a piece of lead across a sheet of white paper and describe
the result. What is erroneously called "black lead" ? Is there any
lead in a lead pencil ?
TESTS FOR LEAD.
Experiment 177. — (a) Recall or repeat the reduction of lead oxide
in the blowpipe flame.
(3) Recall or repeat the action of hydrogen sulphide with the solu-
tion of any lead compound.
(c) Add dilute hydrochloric acid to a little lead nitrate solution until
precipitation ceases. Note the insolubility of the lead chloride which
is formed. Warm gently as long as any decided change occurs. De-
scribe the action. This is characteristic of lead chloride and permits
its separation from the chloride of silver and of mercury (in the -ous
condition).
jio Experimental Chemistry.
{d) Add dilute sulphuric acid to a little lead nitrate solution until
precipitation ceases. The precipitate is lead sulphate. Observe its
properties. Is it soluble in hot water ? Try it.
{e) Repeat {d), using potassium chromate or dichromate instead of
sulphuric acid. The precipitate is lead chromate. Describe it, espe-
cially the color.
Discussion of Experiment 177. — The test in {c) is usually
employed as the preliminary test, and it is confirmed by
{b), {d), and (e\ or any of these three.
Lead, like other metals, interacts with zinc.
INTERACTION OF METALS LEAD.
Experiment 178. — Repeat Exp. 162 (d), and describe the result.
Soluble Lead Compounds are often formed when lead is
subjected to the action of water containing much gas, e.g:
air and carbon dioxide.
ACTION OF WATER ON LEAD.
Experiment 179. — Fill a bottle half full of water, pass carbon
dioxide into it from the lungs, and shake vigorously. Put a piece of
clean lead in the water, cork the bottle, and let the whole stand undis-
turbed for a day or two. Remove the lead and test separate small por-
tions of the water for lead by Exp. 177 (c), (d), (e). If no lead is
detected, let the action continue another day, or evaporate the liquid
to a small bulk and test as above.
Ordinarily, drinking water has no action on lead, but if
the solvent power is increased by the presence of an excess
of gases or of organic matter containing sodium or potassium
nitrate, enough lead compound may be taken up by the
water to cause lead poisoning when the water is drunk.
Oxides of Lead. — There are several oxides of lead ; the
most important are lead monoxide, PbO, lead dioxide, PbOg,
and lead tetroxide, PbgO^.
Lead,
PROPERTIES OF THE OXIDES OF LEAD.
311
Experiment 180. — (a) Examine the three oxides and tabulate
their most obvious physical properties, stating the exact chemical name
and formula and the popular name of each oxide.
(d) Repeat or recall the experiment in which lead was heated in the
oxidizing flame, especially the color of the coating. What oxide of
lead is thereby formed .'*
(c) Warm a little lead tetroxide with dilute nitric acid. The solid
product is lead dioxide. Describe it.
LABORATORY EXERCISE. XXXVL
1. How might lead tetroxide be prepared ?
2. If lead tetroxide is heated strongly, lead monoxide is formed.
What does this fact reveal about the stability of lead tetroxide ?
3. When lead dioxide and concentrated hydrochloric acid are mixed
and heated, chlorine is evolved. Complete the equation —
PbO, + HCl = PbCl., + 2H2O +
How does this interaction resemble that of manganese dioxide and
hydrochloric acid ?
4. How may lead nitrate be formed ? What happens when it is
heated strongly? Prove that it is a nitrate and a lead compound.
5. Examine lead carbonate and state its most obvious properties.
Prove that it is a carbonate and contains lead.
6. Examine a lump of galena and state its most obvious properties.
Prove that it is lead sulphide. Complete the equation : —
PbS + O = 4- SO.,.
CLASS-ROOM EXERCISE. LIX.
Additional study of lead.
(a) Occurrence. (c) Properties.
(d) Metallurgy. (^/) Uses. (e) History.
Oxides of lead. Give the preparation, properties, and uses of —
(a) Lead monoxide. (c) Lead tetroxide.
(d) Lead dioxide.
312 Experimental Chemistry.
3. Lead carbonate.
(a) Occurrence.
(d) Industrial preparation — old and new methods.
(c) Properties. (^/) Uses.
4. Lead sulphide.
(a) Properties of galena.
(d) Properties of the artificial sulphide.
5. Compounds of lead.
Give the chemical name and formula of galena, cerussite, anglesite,
sugar of lead, white lead, red lead, litharge, lead peroxide, chrome
yellow.
6. Miscellaneous.
(a) Valence of lead in PbO, PbO.,, PbS, PbCO,, PbCla, PbCl^.
(i>) Illustrate the Periodic Law by the compounds of tin and
lead.
7. Alloys of lead.
(a) Approximate composition of hard and soft solder, and of
pewter; uses of each.
{d) Approximate composition of Britannia metal and of type
metal.
(c) Components of a fusible alloy. Uses.
(^) What is shot ?
PROBLEMS. XXXIX.
1. What per cent of lead is contained in —
(a) Galena, PbS.
(d) Cerussite, PbCOg.
(c) Anglesite, PbSO^.
(d) Lead acetate, PbCC.H.^O.Oa • 3 HoO
2. How much litharge can be made from 40.5 gm. of lead ? (Equa-
tion is Pb + O = PbO.)
3. An analysis of lead monoxide showed that 100 gm. contained
7.1724 gm. of oxygen. Calculate the atomic weight of lead-
4. If 5 gm. of lead chloride give 5.16 gm. of silver chlonde, what
is the atomic weight of lead ? (Equation is PbCh, + 2 AgNOg =
2 AgCl + Pb(N03)2-)
5. If 100 gm. of lead form 159.9703 gm. of lead nitrate, what is the
atomic weight of lead ? (Assume Pb + 2 HNO3 = Pb(NO,)^ + H3.)
Chromium.
3^3
Chromium is an uncommon metal, but several of its com-
pounds, especially potassium chromate, potassium dichro-
mate, and chrome alum, are familiar, and have numerous
industrial appHcations.
TESTS FOR CHROMIUM.
Experiment i8i. — (a) Prepare a borax bead (see page 231), touch
it with a minute quantity of any chromium compound, ^.^^. chrome alum,
and heat in bodi the oxidizing and reducing flame. The green „olor is
a characteristic and delicate test for chromium.
(d) Mix equal small quantities of potassium carbonate, potassium
nitrate, and powdered chrome alum,, place the mixture on a platinum
foil, and hold it with the forceps in the upper Bunsen flame so that the
mixture will fuse. If a platinum foil is not available, use a porcelain
crucible. A yellow mass, due to the presence of potassium chromate,
results. If the color is not decided, dissolve the mass in water, add
acetic acid, slowly at first, and boil to expel the carbon dioxide. Add
a few drops of lead nitrate solution to a portion, and yellow lead chro-
mate is precipitated. (Compare Exp. 177 (e).) If the precipitate is
white, it is lead carbonate, and shows that not all the potassium car-
bonate was decomposed, as intended.
(c) Add lead nitrate solution to potassium, dichromate solution.
Name and describe the precipitate. Try the solubility of the precipitate
in acetic acid, dilute nitric acid, and sodium hydroxide.
Chromium in its Compounds is sometimes metallic and
sometimes non-metallic toward the other components.
Thus, in chromium trioxide, CrOg, the anhydride of the
hypothetical chromic acid, H2Cr04, chromium acts as a
non-metal, just like sulphur in sulphuric acid; hence in
chromates the chromium is acidic, or non-metallic, in its
chemical relations with the other elements. In chromic
compounds, however, chromium acts as a metal. Thus
chromium hydroxide, Cr(0H)3, is analogous to aluminium
hydroxide ; chrome alum is potassium chromium sulphate,
and is analogous to aluminium alum. The two classes pass
into eagh other by appropriate operations,
314 Experimenta' Chemistry.
Chromates. — The two important chromates are potassium
chromate, K2CrO^, and potassium dichromate, K2Cr207.
PROPERTIES OF CHROMATES.
Experiment 182. — {a) Examine crystals of potassium chromate and
dichromate, and state their most obvious physical properties. From
previously determined facts, predict their solubility or insolubility in
water. Verify the prediction.
(J?) Recall the properties of lead chromate. Complete the equation : —
K^CrO^ + Pb(N0,,)2 = PbCrO^ +
Lead
Chromate
(c) Add a few drops of concentrated hydrochloric acid to a dilute
solution of potassium chromate and observe the change of color. This
is due to the presence of potassium dichromate, which was formed from
the chromate. The equation is —
2KXr04 + 2HCI = K,Cr,0, + 2 KCl + Hp.
Potassium
Dichromate
{d) Add potassium hydroxide solution to 10 cc. of potassium dichro-
mate solution until the color just changes. The yellow color is due to
the presence of potassium chromate, which was formed from the
dichromate. Thus —
K,Cr,0. + 2KOH = 2K,Cr04 + HoO.
(e) The chromates are oxidizing agents. Add a few drops of con-
centrated hydrochloric acid to powdered potassium chromate and dichro-
mate in separate test tubes. Chlorine is evolved, owing to the oxidation
of hydrochloric acid. The simplest equation in the case of the chro-
mate is —
KaCrO^ + 8 HCl = 3 CI + CrCl, + 2 KCl + 4 H,0.
Chromic
Chloride
Definition. — Oxidation sometimes means the withdrawal
of hydrogen from a compound, as well as the addition of
Chromium. 315
oxygen. If the hydrogen can be removed, it is ultimately
oxidized to water, hence oxidation actually occurs.
Chromic Compounds. — The chromic chloride formed in
Exp. 182 {e) illustrates the possibility of passing from a
chromate to a chromic salt. Here the chromate was
reduced. Conversely, Exp. 181 (d) illustrates the forma-
tion of a chromate by the oxidation of a chromic compound.
REDUCTION OF CHROMATES TO CHROMIC COMPOUNDS.
Experiment 183. — Add to a few cubic centimeters of potassium
dichromate solution a little concentrated hydrochloric acid and a few
drops of alcohol. Warm gently. Two important changes occur. The
chromate is reduced to chromic chloride which colors the solution
green ; the alcohol is oxidized to aldehyde, which is detected by its
peculiar odor. (See Exp. in (^).) The equation is —
K.^Cr.Pj + 8 HCl + 3 C,>H,0 = 2 CrClg + 3 C.Hp + 2 KCl + 7 Hp.
Alcohol Aldehyde
PROPERTIES OF CHROME ALUM.
Experiment 184. — (a) Examine chrome alum and state its most
obvious physical properties.
(<^) Recrystallize a little and describe the crystals.
(c) Prove that chrome alum is a sulphate, and that it contains
chromium and water of crystallization.
PREPARATION AND PROPERTIES OF CHROMIC HYDROXIDE.
Experiment 185. — (a) Add a little sodium hydroxide solution to
a solution of chrome alum. The precipitate is chromic hydroxide.
Describe it. Add an excess of sodium hydroxide solution and shake.
Describe the result Boil, and state the result.
(6) Add a little, and then an excess, of ammonium sulphide to a
solution of chrome alum. Compare the result with that in (a). Docs
chromium form a sulphide? Do aluminium and zinc?
ji6 Experimental Chemistry.
LABORATORY EXERCISE. XXXVII.
1. Compare the action of sodium hydroxide on ordinary alum and
on chrome alum.
2. riow can aluminium hydroxide be distinguished from chromic
hydroxide ?
3. Read the discussion of Exp. 169.
CLASS-ROOM EXERCISE. LX.
1. Additional study of chromium.
(«) Occurrence and source. (c) Properties.
(d) Preparation. (d) Uses.
2. Chromates.
Give the preparation, properties, and sources of —
(a) Potassium chromate.
(d) Potassium dichromate.
3. Chromic compounds.
(a) How do these compounds differ essentially from chromates ?
(d) Names and formulas of those experimentally studied.
(c) Preparation and properties of chrome alum.
4. Miscellaneous.
(a) Valence of chromium in Cr(OH)o, Cr^,(S0^)3, K^,CrO^,
PbCrO,, K,Cr,0..
(^d) Chemical name and formula of chromite, crocoite, chrome
yellow, Guignet's green, chrome green, chiome alum,
yellow chromate of potash.
PROBLEMS. XL.
1. Find the percentage composition of —
{a) Lead chromate, PbCrO^.
(d) Chromic oxide, Cr^Og.
(c) Chrome ironstone, Cr^,0.5 . FeO.
(J) Potassium chlorochromate, KClCrOg.
(e) Chromium monoxide, CrO.
(/) Chromium trioxide, CrOg.
2. If 100 gm. of lead nitrate produce 97.576 gm. of lead chromate,
what is the atomic weight of chromium? (Equation is KXrO^ +
Pb(N03)2 = PbCrO^ + 2 KNO3.)
Manganese. 317
3. If 100 gm. of chromous chloride, CrCl^, yield 57.5 gm. of chlorine,
what is the atomic weight of chromium ?
4. If 36.865 gm. of chromic chloride form 100 gm. of silver chloride,
what is the atomic weight of chromium ? (Equation is 2 CrClg + 6 AgNOo
= 2Cr(N03)3 + 6AgCl.)
Manganese, like chromium, is a rare metal. Two of its
compounds, manganese dioxide and potassium permangan-
ate, have already been used.
TESTS FOR MANGANESE.
Experiment 186. — (n) Subject a minute quantity of manganese
dioxide to the borax bead test, and note the color of the bead after
heating in each flame.
(d) Fuse on a platinum foil or in a crucible a litde manganese dioxide
mixed with potassium carbonate and potassium nitrate. (See Exp. 181
(d).) The green mass is a characteristic test for manganese. It is
due to the presence of potassium manganate.
(c) Add. ammonium sulphide to manganese sulphate or chloride
solution. The flesh-colored precipitate is manganese sulphide. Divide
it into two parts. Add hydrochloric acid to one and acetic acid to the
other, then add an excess of ammonium hydroxide to each. Draw a
conclusion regarding the solubility of manganese sulphide.
LABORATORY EXERCISE. XXXVIII.
1. For what has manganese dioxide been used in the laboratory?
(s it an oxidizing agent?
2. For what has potassium permanganate been used in the labora-
tory ?
3. Describe potassium permanganate. What can be said of its solu-
bility in water?
Potassium Permanganate is a powerful oxidizing agent.
OXIDATION WITH POTASSIUM PERMANGANATE.
Experiment 187. — (a) Add a few drops of sulphuric acid to a
weak solution of fresh ferrous sulphate ; then add, drop by drop, a dilute
solution of potassium permanganate. Us color is changed^ owing to
31 8 Experimental Chemistry.
the loss of oxygen which changes the ferrous to the ferric sulphate ;
the decomposition of the permanganate also allows the formation of
potassium and manganese sulphates.
{d) Pour a solution of potassium permanganate upon a piece of filter
paper. Describe and explain the result.
LABORATORY EXERCISE. XXXIX.
1. What is the formula of potassium permanganate? Does the
formula give a clue to the oxidizing power?
2. Potassium permanganate solution is often used as a disinfectant.
Upon what property is this application based ?
CLASS-ROOM EXERCISE. LXL
1. Additional study of manganese.
(a) Occurrence. (c) Properties.
(6) Preparation. (d) Use. ,
2. Manganese dioxide.
(a) Names. (c) Properties.
(d) Occurrence in native state. (d) Uses.
3. Potassium permanganate.
(a) Preparation. (d) Properties. (c) Uses.
4. Miscellaneous.
(a) Valence of manganese in MnO, MnO^„ Mn^Og, MugO^,
KMnO^, MnS, MnClg.
PROBLEMS. XLL
1. If manganous sulphate, MnSO^, yields 42.392 per cent of oxygen,
what is the atomic weight of manganese ?
2. Calculate how much manganese can be obtained from 100 gm.
of—
(a) Manganese dioxide, MnOo-
(d) Manganese trioxide, Mn^,0.;.
(c) Manganese sulphide, MnS.
(d) Manganese alum, K^,Mn^,(SO^)^ . 24 H^,0.
3. Calculate the weight of oxygen liberated from 100 gm. ol potas'
slum permanganate when heated with sulphuric acid. (Equation is
? KMnO^ + 3 H3SO4 = 50 + 2 MnSO^ + K2SO4 + 3 H3O.)
Iron. 319
4. How much manganese ore containing 85 per cent of manganese
dioxide is needed to prepare 300 lb. of chlorine? (Equation is
MnO. + 4 HCl = MnClg + 2 H.O + Cl^.)
Iron and many of its compounds are familiar and have
numerous applications.
GENERAL PROPERTIES OF IRON.
Experiment 188. — (^) Examine cast iron, wrought iron, and steel,
and state their most obvious physical properties. Try the action of a
magnet on each. Drop a pinch of iron powder into the Bunsen flame.
Hold a piece of fine iron wire in the Bunsen flame. Describe the results
and draw conclusions.
{b) Determine the specific gravity of a piece of iron. Compare the
result with the specific gravity of other metals.
LABORATORY EXERCISE. XL.
1. In previous experiments iron has been (i) heated in air, (2) heated
in steam, (3) treated with acids, and (4) heated with sulphur. Recall
and record briefly the essential result of each experiment.
2. What is " iron by hydrogen,'' and " alcoholized iron" (or " iron by
alcohol ■") ?
3. What use has been made of ferrous sulphide in the laboratory.
Ferrous and Ferric Compounds. — Iron forms two series of
compounds, the ferrous and ferric. They are analogous to
cuprous and cupric compounds. The ferrous compounds
in the presence of free acid pass into the corresponding
ferric compound by the action of oxidizing agents, e.g.
oxygen, nitric acid, potassium chlorate, chlorine, and
bromine. Conversely the ferric compounds are reduced to
the ferrous by reducing agents, e.g. hydrogen, hydrogen
sulphide, sulphur dioxide, and stannous chloride. The
passage from one state to the other occurs easily, especially
from ferrous to ferric.
320 Experimental Chemistry.
Ferrous Compounds. — The preparation, properties, and
tests of ferrous compounds are shown by the
BEHAVIOR OF FERROUS COMPOUNDS.
Experiment 189. — (a) Put a few grams (3 to 5) of iron filings in
a test tube, add about 10 cc. of dilute hydrochloric acid, and warm
gently. Ferrous chloride is formed (in solution) . ( i ) Pour a little into
a test tube one-third full of sodium hydroxide solution. The precipi-
tate is ferrous hydroxide. Watch the changes in color. To what are
the changes due? (2) Add a second portion to potassium ferricyanide
solution. The precipitate is ferrous ferricyanide. Describe it. (3) Add
a third portion to potassium thiocyanate solution. If ferric salts are
absent, no change results. (4) Add a fourth portion to potassium
ferrocyanide solution. The precipitate is ferrous ferrocyanide. De-
scribe it.
The above tests, especially (2), serve to distinguish
ferrous from ferric compounds, as will be readily seen by a
study of the
BEHAVIOR OF FERRIC COMPOUNDS.
Experiment 190. — To a little ferric chloride solution add (i)
sodium hydroxide solution. The precipitate is ferric hydroxide. De-
scribe it. Add to ferric chloride solution (2) a little solution of potas-
sium ferricyanide. Compare the negative result with (2) in Exp. 189,
Add as above (3) a little solution of potassium thiocyanate. The rich
wine-red coloration is caused by the soluble ferric thiocyanate. This
test distinguishes ferric from ferrous compounds. Add as above
(4) a little solution of potassium ferrocyanide. The precipitate is ferric
ferrocyanide. Describe it.
Tabulate the results of Exps. 189 and 190.
REDUCTION OF FERRIC COMPOUNDS.
Experiment 191.— Put a piece of zinc in ferric chloride solution
made slightly acid by hydrochloric acid. The nascent hydrogen re-
duces the ferric to ferrous chloride. After the operation has proceeded
for about fifteen minutes, test a portion of the liquid for a ferrous
and a ferric compound. If the tests are not conclusive, continue the
Describe the result.
Iron. 321
OXIDATION OF FERROUS COMPOUNDS.
Experiment 192. — (a) To a solution of fresii or freshly washed
ferrous sulphate add a little hydrochloric acid, warm gently, and then
add a few crystals of potassium chlorate. After heating a short time,
test portions of the liquid for a ferric and a ferrous compoun.L
{b) Add 10 cc. of concentrated nitric acid, drop by drop, to a hot
solution of ferrous sulphate to which a little sulphuric acid 'las been
added, and boil. Test portions of the liquid for a ferric and a ferrous
compound.
{c) Recall a third illustration of the oxidation of a ferrous to a ferric
compound.
LABORATORY EXERCISE. XLI.
1. Read the discussion of Exp. 174.
2. Examine ferrous sulphate, and state its physical properties. Test
a crystal, which has been exposed to the air, for both ferric and ferrous
compounds. Explain.
3. Which seems the more stable form, ferrous or ferric?
4. Examine specimens of hematite, limonite, and magnetite. De-
scribe each. Draw the first two across a sheet of rough paper or a
piece of ground glass, and describe the "streak-' made by each. What
is the formula of each (pure) compound? Significance of each name?
5. Examine iron pyrites and state its most obvious physical proper-
ties. It is iron disulphide. What is the formula? For what is it used?
6. Examine specimens of siderite. Describe them. Siderite is
ferrous carbonate. What is the formula? How is it formed in the
earth's crust? Try the action of warm hydrochloric acid on a little
powdered siderite. How does iron get into clay banks and :Xo the soil ?
CLASS-ROOM EXERCISE. LXIL
I. Additional study of iron.
(rt) Occurrence of free and of combined iron.
(^b) Ores of iron.
(c) Metallurgy of iron and steel.
{d) Properties of cast iron, wrought iron, and steel.
{e) Composition of the three common kinds of iron.
(/") What is passive iron?
{g) Uses of the various kinds.
322 Experimental Chemistry.
2. Explain the general relation of ferrous to ferric compounds, and
illustrate by equations the method of passing from one to the other.
3. Compounds of iron.
Give the chemical name and formula of red hematite, iron pyrites,
magnetite, loadstone, copperas, green vitriol, iron liquor.
4. Miscellaneous.
(a) What elements are related to iron ? What metals are closely
related to iron, and how is this relation emphasized.''
(/; Valence of iron in Fe.fi.^, Fe^O^, FeCOs, Fe(OH)3, FeO,
FeSO^, Fe,,(SOJ,, FeS, FeS,.
(c) The group CN, known as cyanogen, has the valence one.
What is the valence of the iron atoms in ferrous ferricyanide,
Fes (Fe(CN ),;).? In ferric ferrocyanide Fe^(Fe(CN)^)3?
(d) What is ink? Common bluing? Galvanized iron? Iron
rust?
PROBLEMS. XLII.
1. Berzelius found that 1.586 gm. of iron formed 2.265 g'""- of ferric
oxide. Calculate the atomic weight of iron. (Equation is 2 Fe + 3 O
= Fe,03.)
2. A gram of iron liberated 389.74 cc. of hydrogen at 14° C. and
820mm. What is the equivalent of iron?
3. What volume of oxygen is used in the oxidation of 100 gm. of
ferrous oxide to ferric oxide? (Equation is 2 FeO + O = Fe^,03.)
4. What volume of hydrogen at 13° C. and 780 mm. is needed to
reduce 63 gm. of ferric oxide to metallic iron? (Equation is Fe.,03 +
6 H = 3 H,0 + 2 Fe.)
Separation of Metals. — Many of the foregoing experi-
ments illustrate the fundamental principles of quahtative
analysis. The experiments, however, emphasize only the
detection of elements or groups, e.£: zinc and the sulphate
group when testing zinc sulphate. In many instances the
various tests would be interfered with by the presence of
other substances. Hence it is necessary to separate the
essential compounds before applying tests. It is custom-
ary to speak of the separation of metals, though usually
a compound of the metal is meant. Thus, if a mixture
Separation of Metals. ^'^3
of lead chloride and silver chloride is boiled with water,
the lead chloride dissolves, and filtration separates the lead
chloride from the silver chloride, or, loosely, the lead from
the silver. Once separated, the special tests may be used
to detect the various metals. Some separations are simple,
but others are complex. A common method of separation
is illustrated in Exp. 135, since sulphides are usually solids
and behave in various definite ways.
The following experiments illustrate simple separations,
but are in no respects a substitute for a course in quali-
tative analysis. Such a course is beyond the scope of this
book.
SEPARATION OF LEAD AND SILVER.
Experiment 193. — Mix 5 cc. each of lead nitrate and silver nitrate
solutions, and add dilute hydrochloric acid drop by drop, until precipita-
tion ceases. Allow the mixed precipitates of lead chloride and silver
chloride to settle, decant the supernatant liquid down a glass rod, add
a little water to the precipitate, and boil. Filter. Test portions of the
filtrate for lead (see Exp. 177 (d) and (e)). Test the precipitate for
silver by (a) fusing a portion in the blowpipe flame, or (d) determining
the solubility in ammonium hydroxide, or (c) exposing it to the sunlight
(a change to purple indicates silver chloride).
SEPARATION OF LEAD, SILVER, AND MERCURY (-OUS).
Experiment 194. — Mix 5 cc. each of lead nitrate, silver nitrate, and
mercurous nitrate solutions, and add dilute hydrochloric acid drop by
drop, until precipitation ceases. Separate and test the lead as in
Exp. 193. Pierce a hole in the point of the filter paper with a glass
rod, and wash the mixed precipitates of silver and mercurous chlorides
into a test tube with dilute ammonium hydroxide. Warm gently and
shake. Filter, and test the filtrate for silver as in Exp. 193. The black
residue (see Exp. 166 (a)).h a sufficient test for mercury. Its presence
may be confirmed thus : Dissolve the black precipitate in a very little
agua regi'a, dilute with water, and add a clean copper wire ; remove the
wire in a few minutes, wipe gently, and mercury will be seen on the
wire as a bright silvery coating.
324 Experimental Chemistry.
SEPARATION OF SILVER AND COPPER, OR ANALYSIS
OF A SILVER COIN. , ♦
Experiment 195. — Repeat the first part of Exp. 154 {b). Test the
precipitate for silver, and the filtrate for copper (see Exp. 151 {b)
and {c)).
SEPARATION OF COPPER AND ZINC, OR ANALYSIS
OF BRASS.
Experiment 196. — Dissolve a few grams of fine brass wire in a
little dilute nitric acid, evaporate nearly to dryness, dissolve the residue
in water, add 10 or 15 cc. of hydrochloric acid, and pass hydrogen
sulphide gas into the solution for about twenty minutes, or add con-
siderable hydrogen sulphide water. Filter a little, and add hydrogen
sulphide water to see if precipitation is complete ; if not, proceed as
before, but if complete, continue the filtration.
The filtrate contains the zinc as zinc chloride ; the precipitate is
copper sulphide. Test the filtrate for zinc (see Exp. 21 (a)). Dissolve
the copper sulphide in warm dilute nitric acid, and filter, if the solution
is not clear. Test the filtrate for copper (see Exp. 151 {b) and (<:)).
SEPARATION OF LEAD AND TIN, OR ANALYSIS
OF SOLDER.
Experiment 197. — Dissolve a gram of solder filings in as small a
quantity of hot aqua regia as possible, evaporate nearly to dryness,
dissolve the residue in water, add 10 to 15 cc. of hydrochloric acid, and
precipitate the metals as sulphides as in Exp. 196. Filter, wash -with
hot water, pierce a hole in the filter paper, and wash the precipitate into
a test tube with yellow ammonium sulphide. Add more ammonium
sulphide, and shake. Filter.
The filtrate contains the tin as ammonium sulphostannate ; add to it
dilute hydrochloric acid, and yellow stannic sulphide appears.
The precipitate is lead sulphide. Dissolve it in hot dilute nitric acid,
filter, and test the filtrate for lead.
SEPARATION OF COPPER. IRON, AND SODIUM.
Experiment 198. — Repeat Exp. 135
Separation of Metals. 325
SEPARATION OF ZINC AND IRON.
Experiment 199. — Mix 5 cc. each of zinc sulphate and ferric
chloride solutions, add an excess of sodium hydroxide drop by drop,
and shake vigorously. Filter a little, and add sodium hydroxide to see
if precipitation is complete ; if not, proceed as before, but if complete,
continue the filtration. Test the filtrate for zinc. Scrape the piecipi-
tate from the paper into a dish, dissolve in dilute hydrochloric ac'd, and
test portions for ferric iron (see Exp. 190).
SEPARATION OF ALUMINIUM AND IRON.
Experiment 200. — Mix 5 cc. each of alum (common) and ferric
chloride solutions, precipitate, and separate as in Exp. 199. Test the
filtrate for aluminium (see Exp. 169 (a)) and the precipitate for iron as
in Exp. 190.
SEPARATION OF ZINC AND ALUMINIUM.
Experiment 2GI. — Mix 5 cc. each of zinc sulphate and alum (com-
mon) solutions. Precipitate and redissolve the hydroxides of zinc and
aluminium with sodium hydroxide solution. Pass hydrogen sulphide
gas into the solution, or add considerable hydrogen sulphide water.
Filter. The precipitate is zinc sulphide, which may be further tested,
if desired, though the formation of zinc sulphide under these condi-
tions indicates the presence of zinc. The aluminium is in the filtrate as
sodium aluminate. Add to the solution of sodium aluminate hydro-
chloric acid to acid reaction, and boil ; filter, if necessary, and test the
filtrate for aluminium.
CLASS-ROOM EXERCISE. LXIII,
I. Devise and describe a method for the separation and detection of
the metals in a mixture of — ■
(a) Manganese sulphate and chrome alum.
(^d) Silver nitrate and alum.
(c) Copper and aluminium (aluminium bronze).
(d) Lead sulphide and silver sulphide.
APPENDIXES,
APPENDIX A.
MANIPULATION — WEIGHING AND MEASURING.
I, Hard Glass Tubing is infusible in ordinary flames, and
is used in all operations which require intense heat. It is
cut in the same manner as soft tubing, though the scratch
must be deeper. It can be bent b}^ using the hottest pos-
sible flat flame of a Bunsen burner. Ignition tubes and
combustion tubes are made of hard glass tubing and unless
they are heated and cooled slowly they will surely crack.
Constant care should be taken in using such glass. It is
advisable to heat the whole tube gently at first, before
directing the flame upon any particular part. Combustion
tubes may be safely heated by either (i) winding a piece
of fine wire gauze (iron or copper) around the part of the
tube to be heated, or (2) winding a piece of copper wire
(No. 20) once or twice around each end of the tube so that
pieces of the wire project like spokes of a wheel. See
Figs. 45 and 59. The latter device has been repeatedly
used Avith success in all the experiments in this book
demanding such a precaution.
MAKING IGNITION TUBES.
Experiment 6. — Ignition tubes have thick walls and are made of
hard glass. Read again the precautions to be observed in heating hard
glass.
Select a piece of hard tubing about 20 centimeters long and of the
desired internal diameter (from 7 to 10 millimeters is a convenient size).
Warm it by holding it over the flat Bunsen flame, and gradually lower
329
330
Experimental Chemistry.
B
Fi(^.
the tube into the flame so that the middle is in the hottest part.
Rotate it slowly and evenly. When soft, remove it from the flame,
and pull it apart a short distance, as shown in Fig. 84. Then heat A
so that the flame will melt that part
along the dotted line ; as soon as it
yields to a gentle pull, remove from
the flame, and draw it out a little more-
Continue this operation until A and B
are separated. Do not pull the two
portions apart too quickly, or the glass
will be thin at the end. The part A
may have a little thread or bit of glass on the end. Heat this end and
pifich it off" with the forceps ; if large, it may be pulled off, but this opera-
tion is liable to remove too much glass and make the end thin. If the
closed end is not symmetrical, heat it, and then blow gently into the
tube, heating and blowing until the desired shape is produced. Heat
B at the narrower part and proceed as with ^i.
If a blast lamp is available, better results may be obtained by the
same procedure. Hard glass tubes, after heating, should always be
coated with soot from the yellow flame and thus allowed to cool slowly.
84. — Ignition tube, partially
made.
2. Heating. — Flasks, beakers, retorts, and all glass
vessels should never be heated when empty, nor over a
Fk;
direct or free flame even if they
contain somethin^:, unless the
f wire "auze.
Fig. 86. — Porcelain dish.
directions so indicate. They should be placed on a piece of
iron or brass wire gauze (Fig. 85) supported by a tripod or
the ring of an iron stand, and heated gradually from beneath.
Porcelain dishes (Fig. S6) should be heated with even
more precaution than large glass vessels. They should
Appendix A.
33
always be placed on a piece of gauze, or similar protection,
and heated and cooled gradually. They should never be
laid on a cold surface when they are hot, but on a block
of unfinished wood, a piece of asbestos board, or a straw
Fig. 87. — Covered porcelain crucible.
Fig.
Triangle.
ring. Porcelain crucibles (Fig. Sy), however, owing to
their thin walls and small size are usually supported on a
triangle (Fig. 88), or a pronged tripod (Fig. 21), and heated
safely with a low, free flame. The covers of crucibles
often snap, and should be heated with more precaution
than the crucibles themselves, especially if the cover has
been weighed, since the loss of a weighed cover means a
repetition of the experiment.
Fig. 89. — Casserole.
An exceedingly convenient and inexpensive porcelain
vessel for general use is a casserole (Fig. 89). It is a deep,
thin-walled dish with a handle and a flat bottom. It is
made in several sizes, and the one containing 125 cubic
33^
Experimental Chemistry.
centimeters is best adapted to ordinary use. It may be
heated over a free flame, if partly filled with liquid and
kept in motion, though it is safer to employ a piece of
gauze.
3. Evaporation is the slow conversion of a liquid into
vapor. It may be accomplished in several ways. A free
flame is best suited for rapid evaporation, but it should
never be used for concentrated solutions or semi-fluid
masses, as loss or accident may occur by spattering, de-
composition, or breakage. As the concentration increases,
it is advisable to transfer the vessel to a water bath (Fig. 90),
which is intended to ac-
^^^__ complish slow evapora-
tion. The containing ves-
sel, which is copper or
iron, has a top consist-
ing of movable concentric
rings to adapt it to dishes
of various sizes. It is filled
two-thirds full of water,
supported on a tripod or
ring of an iron stand, and heated underneath, the vessel
meanwhile having been so placed on the bath that it is
almost entirely in the steam. The temperature of the
solution to be evaporated is thus slowly raised to nearly
100° C. The water in the bath must never be alloweU
to evaporate entirely or * run dry."
Equable heat at a higher temperature than 100° C.
may be conveniently obtained by using a sand batho
This is simply a shallow iron pan (Fig. 91) filled with
dry, clean sand and supported on the ring of an iron
stand or on a tripod. The dish to be heated is partly
Vir.. 90. — \\'aterl)atn
Appendix A
333
imbedded in the sand, and heat is suppHed by a burner
under the bath. A small tin pan serves the purpose,
though iron
m
ay
be obtained from
Fig. 91. — Sand bath pan.
pans of all sizes
dealers.
Evaporation is also performed
by placing the vessel containing
the solution on a piece of as-
bestos board about 15 centi-
meters square. The asbestos
is supported in the same way as the sand bath. An air
bath is sometimes used for evaporation (Fig. 92). It is
also called a dry-
ing oven, for it
is really nothing
but a small copper
oven supported
on four legs. A
hole in the top
contains a cork
carrying a ther-
mometer which
allows the tem-
perature of the
bath to be noted
without opening
the door. An oil
stove and its ac-
companying bak-
ing oven make
a convenient air
bath.
4. Filtration. — The folded paper should rest firmly on the
walls of the funnel so that the weight of the added liquid
Fig. 92. — Air bath
334
Experimental Chemistry.
will not tear the filter paper. If moistened with water or
the liquid to be filtered after being fitted to the funnel, the
paper will keep its place more securely and filter more rap-
idly. The paper should never project above the edge of
funnel. The stem of the funnel should be near the wall of
the vessel and just above the surface of the liquid to pre-
vent spattering. It is better, also, to pour the liquid to be
filtered down a rod upon the side of the paper, as shown
Fkj. 93. — Filtering —correct
Fig. 94. — Filtering - mcorrect.
in Fig. 93, otherwise the liquid may run down the outside
of the containing vessel, as shown in Fig. 94. If a pump
is used for rapid filtration, the apex of the filter paper
should be protected by a platinum cone, by a small cone
of parchment paper pricked with holes, or by a small
square of cheese cloth folded into the point of the filter
paper.
5. Stoppers, Corks, Joints, and Safety Tubes. — Rubber
stoppers with one and with two holes are recommended.
They are decidedly superior to corks. If corks are used.
Appendix A.
335
they should be free from cracks and be softened by rolHng
or pressing. A convenient way is to wrap the cork in
paper to protect it from dirt, and then roll it under the
foot on the floor. Holes of any size are made in corks
by a cork borer. A set of cork borers is shown in Fig.
95. The stout wire which accompanies each set of cork
borers serves as a handle when passed through the holes
in the cap of the borer, and also as a piston to remove the
cyUnder of cork which often remains in the borer after the
operation.
Fig. 95. — Set of cork borers.
To bore a hole in a cork proceed as follows : Select a
cork free from cracks or channels and use a borer which
is one size smaller than the desired hole. Hold the cork
between the thumb and forefinger, as in Fig. 96, press the
larger end against a firm
but soft board, and slowly
push the borer by a ro-
tary movement through
the cork, taking care to
keep the borer perpen-
dicular to the cork. If the hole is too small, enlarge it
with a round file. If corks are used instead of rubber
stoppers, the apparatus should always be tested before use
by blowing into it, stopping of course all legitimate out-
lets. A poor cork often means a failure, to say nothing
of wasted time.
Rubber stoppers to be used in quantitative experiments
should be boiled in dilute sodium hydroxide solution, rinsed
Fig. 96. — Boring a hole in a cork.
jj6 Experimental Chemistry.
with water, then boiled in dilute hydrochloric acid, and
finally washed with water. This operation removes all
particles of matter which might adhere to tubes and thereby
introduce into the final result a needless error.
Glass tubes are joined by short pieces of rubber tubing
called rubber connectors. Such joints should be gas-tight.
Before experiments begin leaks should be discovered by
testing the apparatus at the necessary points. Dj not stop
leaks by wax, vaseline, or any temporary makeshift. Select
pieces of apparatus which fit. Rubber connectors may be
tightened by tying them to the glass tube with a waxed
thread. Joints are often made tighter by having on
the end of the glass tube a flange, made by heating the
end of the tube and pressing it while hot upon a hard
surface.
Safety tubes are recommended in place of the ordinary
straight thistle tube. See Fig. 30. They should fit per-
fectly and always have enough liquid in the lower bend to
prevent a backward escape of gas. If the liquid to be
introduced will not run down, loosen the stopper slightly.
If the straight thistle tube is used, it must dip into the
liquid in the flask or bottle.
6. To Cut off the Bottom of a Bottle. — Select a bottle with
walls of uniform thickness. Tie a piece of cotton string
loosely around the bottle at the point where it is to be cut.
The string should be just tight enough to stay in place
and the knot very small. Pour a little kerosene on the
string, turning the bottle slowly at the same time, until
the string is saturated. The kerosene should follow
the string and in no case be allowed to run down tne
side of the bottle. Light the kerosene at one point and
rotate the bottle slowly so that the flame will heat the glass
Appendix A. 337
beneath the string. When the circle is complete, quickly
plunge the bottle into a pail of water, and it will crack
evenly at the desired point. The edges should be smoothed
with emery paper. Both parts of the bottle may be utilized
in many experiments. The ordinary five-pint acid bottle,
or Hthia bottle, is well suited to this method.
7. To Insert a Glass Tube into Rubber Tubing. -- Cut
the rubber tubing at an angle, as shown in Fig. 97, moisten
the smoothed end of the glass tube with water, place the end
of the glass tube in the angular shaped cavity so that both
Fig. 97. — Rubber tube cut at an angle (exact size).
tubes are at about a right angle, and then slip the rubber tube
slowly up and over the end of the glass tube. If the glass
tube is large or the rubber tube stiff, the rubber tube
must be held firmly between the thumb and forefinger to
keep it from slipping off until it is securely adjusted.
8. To Fit a Glass Tube to a Stopper. — First round the
edges of the tube in the flame. When it is cold, moisten
the end with a little water, glycerine, or vaseline (preferably
the first), grasp it firmly about an inch from the end, hold
the stopper between the thumb and forefinger of the other
hand, and work the tube into the hole by a gradual rotary
motion. Proceed in the same manner, if the tube is to be
pushed through the stopper. Never point the tube toward
the palm of the hand which holds the stopper. Never
grasp a safety tube or any bent tube at the bend when
inserting it into a stopper — it may break.
33^
Experimental Chemistry.
9. To Dry the Inside of a Tube or Bottle.— Moisten the
inside with alcohol, and then with a bellows force in air
through a glass tube reaching to the bottom of the vessel ;
hold the bottle mouth downward and withdraw the glass
tube as the alcohol evaporates. Never stand a wet bottle
on a hot support, for a drop of water may be jostled down
upon the heated bottom and shatter the bottle. The inside
will dry if the bottle is exposed to the sun, but the operation
is tedious.
10. To Clean the Inside of a Bottle. — Partly fill the bottle
with water, drop in wads of soft paper, shot, or sand, and
shake the bottle vigorously. This device is also applicable
to the large tube of a condenser, and to large flasks,
11. To Introduce a
Powder into a Tube. —
First fold a narrow strip
of smooth paper so that
it will slip into the tube
easily. Place the pow-
der at one end of the
troughUke holder, and
slowly push the paper into the tube, as far as necessary,
moving the tube rather
than the paper. This
operation is shown in
Fig. 98. Rotate the tube
or turn the paper, and
the powder will be de-
posited at the desired
point (Fig. 99). Care-
fully withdraw the paper. Introduce all powders in this
way, whether the tube is open or closed, large or small.
Fig. 98. — Introducing a powder into a tube
first stage.
Fig. 99.
Introducing a powder into a tube —
second stage.
Appendix A
339
Fig. ioo. — Pouring a liquid down a rod.
12. Pouring Liquids and Transferring Solids. — Liquids
may be poured from a vessel without spilling, by moisten-
ing a glass rod with the
liquid and then pouring
it down the rod as is
shown in Fig. lOO. The
angle at which the rod
is held varies with cir-
cumstances. This is a
convenient way to pour
a liquid from a vessel containing a solid without disturbing
the solid. Solids should never be poured directly from a
large bottle into a test
tube, retort, or similar
vessel. A convenient
method is as follows :
Rotate the bottle slowly
so that the solid will roll
out in small quantities ;
catch this solid on a nar-
row strip of paper folded along the middle, and slide the
solid from the paper into the desired vessel. ' The last
part of the operation
is shown in Fig. lOi.
Liquids can often be
poured from a bottle
by holding the bottle
as shown in Fig. 102.
Notice that the stopper
and bottle are held in pjg. 102. —The way in which a glass stopper
the same hand. This should be held while a liquid is being
poured from a bottle.
is accomplished by
holding the palm of the hand upward and removing the
Fig. ioi. — Pouring a solid into a vessel with
a small opening.
340
Experimental Chemistry.
stopper by grasping it between the fingers before the bot-
tle is lifted. All stoppers should be removed this way,
when possible, and never be laid down, because the im-
purities adhering to the stopper may run down into the
bottle and contaminate the solution. Never return any
reagent to a bottle ; if too much has been taken, throw it
away. Stoppers and bottles will not become mixed if
each bottle and stopper has a similar label.
13. To Make a Platinum Tip. — The steps arc shown in
Fig. 103. A is a. piece of glass tubing about 12 centime-
1 A
D B
:>c
Fig. 103. — Stages in making a platinum tip.
ters long and 5 millimeters in internal diameter. Heat this
in the middle in the Bunsen flame and draw it out into the
shape shown in B ; draw it slowly at first, then rapidly.
Cut B at the point indicated by the dotted Hne. Roll a
piece of thin platinum foil, about 25 millimeters square
into a conical tube, as shown in D ; begin at one corner and
roll it around the small end of a round file, gradually shap-
ing the platinum tube so that it will slip into the larger end
of C but not out of the smaller end. Heat B where the
dotted line is drawn, and the glass and platinum will firmly
unite. The completed tip is like B, though it may vary
Appendix A. J41
with circumstances. It is used in burning hydrogen and
other gases and is attached to the generator by a rubber con-
nector. This tip is more expensive than the customary
form, but it is safe and durable. A short piece of pipe-
stem or capillary tubing may be used instead of the plati-
num tip, but these are not so satisfactory.
14. To Seal a Platinum Wire into a Glass Rod. — Rotate
one end of a piece of glass rod, about 10 centimeters long,
in the flame until it softens. At the same time grasp a piece
of platinum wire from 5 to 7 centimeters long firmly in the
forceps about i centimeter from the end and hold it in the
flame. When the rod is soft enough, gently push the hot
wire into the rod. Cool the rod gradually by rotating it
-o
Fig. 104. — Platinum test wire.
in the flame. The completed wire is shown in Fig. 104.
If a glass tube is used, instead of a rod, it should be
drawn out to a very small diameter (see Fig. 103, B)
before inserting the platinum wire, but in other respects
the two operations are practically identical.
15. Collecting Gases. — Gases are usually collected over
water by means of a pneumatic trough, a common form of
which is shown in Figs. 23 and 25. The vessel to be filled
with gas is first filled with water, covered with a piece of filter
paper, inverted, and placed mouth downward on the shelf of
the trough, which is previously filled with water just above
the shelf. The paper is then removed and the vessel
placed over the hole in the shelf of the trough. Glass plates
instead of filter paper may be used to cover the bottles
34^
Experimental Chemistry.
Figure 105 is a conventional sketch of a cylinder in position
in a pneumatic trough. It is represented partly filled with
gas which comes through a delivery tube and bubbles up
through the water into the cyl-
inder. As the gas rises, water
is forced down out of the bottle
into the trough. Bottles are
usually used in place of cyl-
inders, and the delivery tubes
have various shapes. All gases
insoluble in water may be col-
lected over water in this way.
Some heavy gases, like chlo-
rine, hydrochloric acid, and
sulphur dioxide, are collected
Fig. 105. — Cylinder partly filled with by allowing the gaS tO flow •
gas in position in a pneumatic ^^^^^^^^^^ ^^,^0 an empty bottlc
trough. ^ J
and displace the air in the
bottle, i.e. by dowmvard displacement. See Fig. 52. Am-
monia gas, being a light gas, is collected by allowing the
gas to flow upward into a bottle, i.e. by upward displace-
ment. See Fig. 58.
16. A Gas Holder. — A small gas holder is constructed
as follows : A bottle, as large as available, is provided with
a two-hole rubber stopper through which pass two tubes,
each bent at a right angle ; one tube, B, reaches to the
bottom, the other. A, is just even with the stopper. Attach
about 20 centimeters of rubber tubing to each glass tube
and provide each rubber tube with a Hofmann screw (Fig.
106) near the end of the glass tube. All joints must be
air tight. The holder is filled as follows : Fill the bottle
and tubes with water, close A with the screw about 5 cen-
Appendix A.
343
Fig. io6. — Hofmann screws.
timeters from the end, and put the outer end of B m a
vessel of water. When the apparatus which generates or
contains the gas to be introduced is ready, pinch B with the
thumb and forefinger, loosen
the screw on A, connect the
delivery tube with A, and im-
mediately remove the press-
ure from B. The gas will
flow in through A and the
water out through B. The
initial pressure of the en-
tering gas must be strong
enough to start the flow through B, and once started
the operation proceeds smoothly. When the holder is
full, pinch A, instantly slip off the delivery tube, and com-
press A firmly with the screw a few centimeters from the
end. Attach the end of B to the lower opening of an
aspirator bottle or to a large funnel supported above the
holder to force out the gas. B is loosely clamped or left
open, but if clamped, the flow from the reservoir to the
holder is easily controlled. The gas is driven from the
holder by loosening the screw attached to A. Do not
attach B to the faucet, to force out the gas, as air from the
water pipe is apt to pass into the holder along with the water.
17. Aspirators. — It is often necessary to draw (or force)
a gas through a tube. This may be done by a filter
pump, or by an aspirator. The former may be obtained
from the dealer, the latter is easily constructed. A com-
mon form is shown in Fig. 107. A five-pint acid bottle is
provided with a two-hole rubber stopper carrying two open
tubes. The shorter is bent at a single right angle, it ex-
tends just within the bottle, and is called the inlet tube.
344
Experimental Chemistry.
The longer tube is bent at a double right angle, the inner
arm of which extends nearly to the bottom of the bottle,
while the other arm terminates at any
convenient point below the bottom
of the bottle. A Hofmann screw,
attached to a piece of rubber tubing,
A, near the end of this outer tube,
serves to regulate the flow of water.
If more convenient, the rubber tube
may replace the greater portion of
the outer arm of the outlet tube.
An aspirator bottle may be obtained
from the dealer, but one made as
above serves the purpose. To draw
— or aspirate — a gas, fill the bottle
with water and insert the stopper
with its tubes, attach the shorter
arm to the apparatus, loosen the
screw, start the water flowing, and regulate the flow by
the screw.
Fig. 107. — Aspirator.
18. The Metric System originated in France during the
French Revolution, and its use has since been required or
legalized in most civilized countries. The fundamental
unit of the system is the meter, which is approximately
equal to the ten-millionth part of the distance from the
equator to the north pole. This distance was ascertained
by actual measurement of an arc of a meridian passing
through Barcelona in Spain and Dunkirk in France.
The legal equivalent of the meter in the United States is
39.37 inches.
The superior advantage of the metric system is its deci-
mal character, which allows rapid transformations from
Appendix A. 345
volume into weight and vice versa, and between denomina-
tions of the same unit. Each unit has multiples and sub-
multiples, which are designated by prefixes attached to the
particular unit. The prefixes denoting multiples are deca-^
hccto-, and kilo-, equivalent respectively to 10, 100, and
1000. The submultiple prefixes are deci-, ccnti-, and milli-,
which correspond respectively to .1, .01, and .001.
The meter — the unit of length — is seldom used in
chemistry, but the centimeter is often employed to express
the length of a tube, or the linear dimensions of a large
vessel. The height of the barometer is stated in centi-
meters or millimeters, and the length of the smallest pieces
of apparatus is often expressed in the latter denomination.
The unit of weight is the gram, which, like the meter,
has its multiples and submultiples. The gram itself is
actually derived from the kilogram, which is the weight in
a vacuum of a cubic decimeter, or 1000 cubic centimeters,
of pure water at its maximum density. Y{Q.xiZQ^ 07ie gram
is the zveigJit of one cubic centimeter of zvater. The weights
of small masses are expressed in terms of the gram. If
an object weighs, for example, 2 grams, 2 centigrams, and
5 milligrams, the weight is written 2.025 gi'ams, though the
two small weights may read 20 and 5 milligrams. Two
milligrams is more often written as .002 gram than as
2 milHgrams, though both forms are used. The decimal
form of the fraction is always used in the metric system.
Thus, 4 decigrams is not written -^^ grams, but .4 gram, or
often, 0.4 gm.
The unit of volume is the liter, which is used for both
dry and liquid measure. It is equal to the capacity of the
vessel containing the standard kilogram. It therefore con-
tains 1000 cubic centimeters, i.e. it is a cubic decimeter.
A liter of water weighs looo grams. Volume is usually
346
Experimental Chemistry.
expressed in cubic centimeters, and most graduated volu-
metric apparatus is described by this denomination. Thus
a liter flask is often marked looo cubic centimeters;
burettes are made to deliver 50 cubic centimeters, 01 100
cubic centimeters, and pipettes from i to 100 cubic centi-
meters. The following tabular view will make clear the rela-
tions between the units, their multiples, and submultiples : —
Length.
Weight.
V Volume.
Notation.
Kilometer
Kilogram
Kiloliter
1000.
Hectometer
Hectogram
Hectoliter
100.
Decameter
Decagram
Decaliter
10.
METER
GRAM
LITER
I.
Decimeter
Decigram
Deciliter
O.I
Centimeter
Centigram
Centiliter
O.OI
Millimeter
Milligram
Milliliter
O.OOI
From this table it is evident that 10 milligrams equal i
centigram, 10 centigrams equal i decigram, 10 decigrams
equal i gram, and so on. The numerical significance of
the prefixes is also apparent.
The passage from the English to the metric system may
be accomplished by utilizing the following: —
TABLE OF TRANSFORMATION.
To Changi
Inches to centimeters ....
Centimeters to inches ....
Cubic inches to cubic centimeters
Cubic centimeters to cubic inches
Ounces to grams
Grams to ounces
Multiply by
2.54
3937
16.387
.061
28.35
0-0353
Appendix A.
347
19. To counterpoise a balance proceed as follows : Clean
the pans with soft paper or cheese cloth, and allow them
to swing freely to ascertain the exact condition of the
balance. If the pointer does not make equal excursions
on each side of the middle or zero point, then add to the
lighter side weight sufficient to restore equilibrium. Do
not wait for the pointer to come to rest, but estimate dis-
tances to the right and left. Balls of paper, bits of wire
or of match, may be added to the proper pan of the hern
pan balance to secure equilibrium. Often a piece of wire
bent in the form of a spring is attached to the beam, if
the balance is persist-
ently ''off." Weight is
added to the lighter pan
of the trip scales by
properly adjusting the
screws near the pointer.
The horn pan balance
may be protected from
drafts by hanging the
balance from the top
of a box shaped like an
ordinary balance case.
The box is open in front,
but the rear is covered
with cheese cloth, held firmly in place by half-round
moulding. The box and its enclosed balance are shown
in Fig. io8.
Fig, io8. — Horn pan balance in a case.
20. Weights from the kilogram to the gram are made
of iron or brass, and the smaller weights of platinum or
some other durable metal. They may be obtained in sets
possessing any degree of accuracy. A set from 50 gm. to
348
Experimental Chemistry,
I eg. is large enough for the exact work in this book.
See Fig. 109.
21. Weighing. — The object to be weighed is placed in
the center of the left-hand pan of the counterpoised
balance, and in the
other pan a weight
is placed, which is
assumed to be the
approximate weight
of the object. If
the weight is too
heavy, return it to
the box and place
on the pan a weight
or several weights
smaller in amount.
Proceed thus, add-
ing and removing
weights, until the
balance is judged to be in equilibrium ; then allow the
pointer to swing freely several times, and note the dis-
tances. If they are unequal, adjust the weights accord-
ingly; if they are equal, the balance is in equilibrium and
the sum of the weights is the weight of the object. Record
the weight immediately in the proper place in a note-book
— not on a scrap of paper. The most satisfactory method
of determining the correct sum of the individual weights is
as follows : Add the weights missing from the box and
record their sum ; then add the weights on the 'oalance
and compare with the amount recorded ; finally, as the
weights are replaced, beginning with the largest, add
again and check the first result This operation takes
Fig, 109. — Set of weights.
Appendix A. 349
time, but assuredly less time than that consumed by a
repetition of the weighing or of an experiment.
Certain precautions must be observed in weighing.
( 1 ) Substances should not be weighed on the bare pan,
but on a piece of smooth paper creased on the sides or in
the middle, or in some counterpoised vessel, e.g. a watch
glass, beaker, crucible, tube, or flask.
(2) Never touch the weights with the fingers. Use
clean forceps.
(3) Arrange the weights symmetrically about the center
of the pan. This is especially necessary with large weights
on the trip scales.
(4) Never attempt to weigh the exact amount specified,
e.g. 1.49 gm., but weigh accurately an approximate amount,
e.g. 1.47 gm., or 1.50 gm., unless, of course, the directions
state exactly 1.49 gm. The expression "from 2.1 to
2.9 gm." means any weight between these two, but the
weight, whatever it is, must be exactly known; the two
amounts, 2.1 and 2.9 gm., are simply limits adapted to
the experiment. So also the expression, ''weigh about
2.4 gm.," means an exact amount which approximates
2.4 gm.
(5) If the same object is to be weighed more than once,
it should be weighed, if possible, with the same balance
and weights. If the balances are carefully counterpoised,
this precaution is, however, not absolutely necessary.
(6) The balance and weights must always be left in a
usable condition — and they will always be so found.
22. Measuring Liquids and Gases. — Liquids are meas-
ured in graduated cylinders, or graduates, tubes, burettes,
and pipettes. Vessels of various capacities are used, de-
pending upon the volume of liquid to be measured. A
350
Experimental Chemistry.
liter graduate is shown in Fig. lie. A graduate holding
lOO cc. is best adapted to general use. Small volumes are
more conveniently measured in a burette, or in a pipette
(Fig. 40). Burettes and pipettes also hold
various volumes, though the burettes most
often used hold 50 cc., and pipettes ic cc.^
25 cc, and occasionally 50 cc. The method
of using a burette is explained in the ex-
periment on neutralization (Exp. 53). The
pipette is used when definite small volumes
are wanted, such as 10 cc, or from i to 5 cc
The pointed end is dipped into the liquid,
which is then gently sucked up into the tube
to a point just above the mark on the stem,
and the top of the tube is quickly closed
with the forefinger. The relative position of
the hand and pipette at this stage is shown
in Fig. 40. If the pressure of the finger is
Fig. no. — Liter
graduate.
lessened, the liquid will slowly fall to the
mark on the stem, and farther if desired.
Pipettes hold the indicated volume between this mark and
the extreme lower end. Occasionally flasks holding a liter
or its fractions are used to obtain exact volumes.
Gases are measured roughly by collecting them over
water in a vessel the capacity of which is known or subse-
quently determined by filling the space occupied by the gas
with water from a graduate or burette. In accurate experi-
ments, the gas is collected in a graduated tube called a gas
measuring tube (see Figs. 32 and 48), or in a graduated
tube having two platinum wires sealed into the walls near
the closed end, and called a eudiometer (see Figs. 111
and 44). The latter is used only in those cases in which
electric sparks are to be passed through the contents of
Appendix A.
35^
the tube. Often the capacity of a plain tube may be found
by a burette and used instead of the more expensive
graduated tube.
Certain essential precautions must be ob-
served in reading volumes of liquids or gases,
(i) Notice the exact value of a unit space
on the graduated scale, since single spaces
often vary in value with vessels of different
capacities. Thus one space may, and often
does, equal o.i cc, 0.2 cc, i cc, 2 cc, and
even 5 cc. on different vessels.
(2) The surface of most liquids in a tube,
especially in a small tube, is not horizontal
but concave. This curved surface is called
the meniscus. It is customary to call the low-
est point of the meniscus the correct height
of the column of water, or depth of the column
of gas, as the case may be.. Thus in Fig. 112
the correct reading of
the meniscus is ii./cc,
though it may be re-
garded erroneously as
more or less, if the eye
were at II., or III., in-
stead of at I., the cor-
rect position. The eye
must be in a line tangent to the low-
est point of the meniscus in order to
determine the correct reading. The
reading 11.7 cc. is the depth of the
gas column in Fig. 112. If the
height of the water column is desired, as is often the case,
then 1 1.7 must be subtracted from the graduated capacity
Fig. III. — Eu
diometer.
11-
12-
,,11
--III
Fig. 112. — Meniscus. Cor-
rect reading is along the
line I.
352
Experimental Chemistry.
of the vessel, if the graduation is from the
top downward. Usually the graduation of
a cylinder or graduate begins at the bottom,
and the height of the meniscus is then the
actual height of the water column. Some-
times tubes are graduated in both directions,
and the reading is then simple.
23. A Thermometer (Fig. 1 13) is an instru-
ment for measuring temperatures. There
are two kinds in use ; the one in scientific
use is the Centigrade, the one in general
use in this country is the Fahrenheit. The
change in temperature is measured by ex-
pansion and contraction of the mercury, or
popularly speaking, by " rise and fall of the
thermometer." The abbreviation for centi-
grade is C. and for Fahrenheit is F.
The point where the mercury stands when
held for a time in boihng water is called the
boiling point of water, and the point where
the mercury stands when the thermometer
is held in melting ice is called the freezing
point of water.
On the centigrade thermometer the boil-
ing point is 100, and the freezing point is
zero ; on the Fahrenheit thermometer the
boiling point is 212, and the freezing point
is 32. The space between these two points
is divided into equal portions called degrees.
The abbreviation for degrees is °, s,g: 212°.
Between these two points on the centigrade
scale, as the space is called, there are 100
■^
Appendix A. ^^^
equal portions; but on the Fahrenheit scale there are i8o
equal portions. Hence lOO degrees centigrade equal i8o
degrees Fahrenheit. Since
.
I go' C. = 180'' F.
then
5^C. = 9^F.
or
i^C. = f of i"F.
Buc as all readings are made from 0°, 32 must be added
when the change is made i/ito Fahrenheit, and subtracted
when the change is made fnwi Fahrenheit. Hence
F. =f C. + 32
C. = |(F.-32)
Notice that in the second formula 32 is to be subtracted
before multiplying by |.
Observe the following precautions in using a ther-
mometer : —
(i) A thermometer should never be thrust suddenly into
a hot liquid, but held in the vapor until the temperature
registered is nearly that of the liquid.
(2) Always allow the thermometer to stand in a liquid
long enough to assume the temperature of the Hquid, be-
fore reading the thermometer.
(3) Verify all readings by a second or independent
observation.
(4) Record all readings in the note-book as soon as made.
24. The Barometer is used to measure the pressure of
the atmosphere.
The normal height of the barometer is 760 mm. It is
assumed in all problems in this book that the pressure is
760 mm., unless otherwise stated.
Precautions (3), (4) under thermometers must also be ob-
served in reading the barometer.
APPENDIX B.
I. GAS LAWS.
1. Law of Charles. — It has been found that all gases
under constant pressure expand or contract uniformly for
the same change of temperature. This law, known as
the Lazv of Charles, may be stated more explicitly thus : —
A given volume of gas under constant pressure expands or
co}itracts 273 of its volume at zero degrees centigrade for
every degree through wJiich it is heated or cooled.
This law means that —
273 cc. at 0° become 273 + i = 274 cc. at 1° C.
273 cc. at o"^ become 273 + 2 = 275 cc. at 2° C.
273 cc. at 0° become 273 + 3 = 276 cc. at 3° C.
or 273 volumes at t° become 273 + t volumes at f" C.
Let V= the volume at 0° C.
and F' = the volume at /° C.
Then V\ V ::273 : 273 + t.
This proportion may be solved for either V or V ;
thus —
273 + / ^ ^
273
354
Appendix B. 355
From (i) we can easily calculate the volume a gas would
occupy if at 0° C. For example, suppose a volume of gas
at 17° C. measures 245 cc, its volume at 0° C. would be
found thus : —
r, 245x273
y= _Ho ^ ^ 230.6 cc.
273 + 17
And from (2) the volume at any temperature may be
found, if the volume at 0° C. is known. Thus if the
volume is 230.6 cc. at 0° C, the volume at 17° C. would
be found by substituting the proper values in (2). Thus —
^,^230.6(273+17)^2 ^^
273
2. Correction for Temperature. — Since the volume varies
with the temperature, gas volumes to be comparable must
be at the same temperature. This situation is not always
possible, hence it is customary to reduce the observed
volume by formula (i) to the volume it would occupy if it
were to exist at o^ C, which is the normal or standard
temperature. The operation is called ''reducing to standard
temperature," or "correcting for temperature." The sig-
nificance of the law and the correction for temperature will
be clearer after the solution of some typical problems.
PROBLEMS. III.
/?i tJie fcllowing- problems f/ic pressure is to be regarded
as co7istant.
1. Find the volume which 173 cc. of gas at 12'' C. would occupy at
o°C.
2. If a gas volume measures 466 cc. at 14'' C, what would it meas-
ure at 0° C. ?
3. What would be the volume at o'^ C. of a volume of gas which
measures 706 cc. at 15° C..''
2^6 Experimental Chemistry.
4. A certain volume of a gas measures 25 cc. at 27^ C. What would
be its volume at o" C. ?
5. Two gases are equal in volume ; one is at — 20" C, the other is at
+ 20° C. What would each volume be at 0° C. ?
6. A gas volume measures 3 1. at 0° C. At what temperature
would its volume be 4 1. ?
7. A volume of gas at 13° C. measures 100 cc. What would be its
volume at - 130^ C. ? At - 13° C. ? At + 130° C. ?
8. If 1000 cc. of gas are heated from o^ C. to 39^^ C , what is the new
volume ?
9. If 1000 cc. of gas are heated from —39' C. to 52"" C, what is the
new volume?
10. Suppose 10 cc. of gas are measured at 15' C. What volume will
the gas occupy at 150° C. ?
11. If 500 cc. of nitrogen at 27'' C. arc cooled to —5" C, what is the
new volume?
12. A volume of gas measured 120 cc. at 15° C. What will be the
volumes at 30" C, o" C, and — 10° C. ?
13. If 743 cc. of gas are heated from 47 C. to 83^0., what is the
new volume ?
14. One thousand cubic centimeters of air at 13" C. w^ill occupy what
volume at 65° C?
15. If 155 1. of air are cooled from 150° C. to 0° C, what will be the
new volume?
16. A liter of oxygen is heated from 14'' C. to 42" C. Find the new
volume.
17. A volume of carbon dioxide measures 1500 cc. at o" C. What
will be its volume at each of the following temperatures : (a) I5°C.,
(d) 50"^ C, (c) 100^ C, (<'/) 30o''C. ? At what temperature will it
measure a liter?
18. If a volume of hydrogen measures 100 cc. at 100" C, what will
be the volume at — 100'' C. ?
19. Twenty-five volumes of air measured at — 10"' C. are heated until
they occupy 30 volumes. To what temperature must they have been
heated ?
3. Absolute Temperature. — Since a gas contracts 273 of
its volume at zero C. for every degree through which it is
cooled, then at —273° C. the gas would have no volume!
Appendix B. j^y
This condition is not possible, and as a matter of fact as
gases approach this point (which has not yet been reached),
they deviate from the law. However, for convenience the
poinc —273" C. is regarded as a starting or zero point, and
IS called absolute zero. Absolute temperature is reckoned
from this point, just as centigrade temperature is reckoned
from zero on the centigrade scale. Degrees on the abso-
lute scale are found by adding 273 to the readings of the
centigrade thermometer. Thus —
273° absolute =0° C.
274^ absolute = + 1° C.
272° absolute = — 1° C.
Hence the Law of Charles is often stated thus : —
TJie volume of a gas undei' constant pressmr varies
directly as its absolute temperature.
4. Law of Boyle. — The effect of pressure on gases was
early discovered, but it was not until about 1662 that
Boyle announced his Law: —
The volume of a gas at a c oust ant tempei^ature varies
inversely as tJie pressure.
This law means that 200 cc. of gas under a pressure of
125 mm. become —
100 cc. under a pressure of 250 mm.
50 cc. under a pressure of 500 mm.
25 cc. under a pressure of 1000 mm.
or, 200 X 125 = 25 X 1 000.
358 E>
[penmental Chemistry.
If
200 cc.= V
X
125 mm.=P
25 cc.= V
1000 mm. = P',
then
VxP= V'xP'
or,
V: V ::P' -.P
This proportion
may be solved for Vox V'\ thus —
VP'
^ P
VP
V' —
By agreement among chemists, P is equal to 760 mm.
Hence the above equations become respectively: —
760
V'J^ (4)
Suppose a gas at 740 mm. measures 420 cc, its volume
at 760 mm. would be found thus : —
^^ 420x740 _
Formula (4) is not often used in actual work, but it serves
to extend the conception of the general application of the
law.
5. Correction for Pressure. — In order to compare gas
volumes they must be measured when under the same
pressure. This, however, is impracticable, or, at least, in-
convenient. Hence it is customary, as in the case of the
temperature correction, to measure each volume under the
Appendix B. 359
existing pressure, which is read directly from the barometer,
and then to reduce the observed volume to the volume it
would occupy if at 760 mm., which is the standard
or normal pressure. This operation is called ''reducing to
standard pressure," or ''correcting for pressure."
PROBLEMS. IV.
/;/ the folloiving problems the temperature is to be re-
gardeei eis constant.
1. The volume of a gas at 740 mm. is 200 cc. What is the volume
at 760 mm . ?
2. If a gas volume at 780 mm. is 25 cc, what would the volume be
at 760 mm.?
3. A gas holder contained 4.5 1. of oxygen when the barometer stood
at 755 mm. What would the volume be at 762 mm. ?
4. A tube contained 137 cc. of air when the barometer stood at
766 mm. The next day the barometer reading was 757 mm. What
was the new volume?
5. A gas measures 1000 cc. at 770 mm. .What is its volume at
530 mm- ?
6. A volume of hydrogen measures 467 cc at 756 mm. What would
its volume be at the normal pressure ?
7. A volume of gas was measured and found to be 195 cc when
the barometer stood at 740 mm. If the volume should be increased
to 200 cc, what would be the barometer reading?
8. A gas volume was 467 cc. at 756 mm. What would the volume
have been at 29.92 in. by the barometer?
CLASS-ROOM EXERCISE. II.
1, Historical.
{a) Essential facts in the life of Charles.
{b) Essential facts in the life of Boyle.
(J) Mariotte's share in the discovery of Boyle's Law.
2. Experimental.
{a) Proof of Boyle's Law.
{b) Proof of Charles's Law.
360 Experimental Chemistry.
6. Aqueous Tension. — When different gases are confined
in the same vessel, each gas exerts its pressure against the
atmospheric presstire. Hence, when a gas is collected over
water, as many gases are, it is saturated with water vapor
and does not bear the full pressure of the atmosphere,
since part is borne by the water vapor. Hence the volume
of a gas is increased by the pressure of water vapor. The
pressure of water vapor, or aqueous tension, as it is usu-
ally called, has been calculated. in terms of miUimeters of
mercury, and may be subtracted directly from the barom-
eter reading. Hence the pressure under which a dry gas
is to be regarded as existing is really ?' — a, a being the
symbol commonly employed to designate aqueous tension.
The amount of water vapor present in a volume of
saturated gas varies with the temperature, hence the pres-
sure exerted by this confined water vapor in its effort to
escape varies with the temperature.
The conception and significance of the tension of aqueous
vapor is not always grasped by students. The fact may
be illustrated as follows : —
Take a glass tube about 90 cm. long and closed at
one end, and dry and warm it. Fill it with dry, warm
mercury and invert in a mortar containing the same liquid.
The mercury will fall a little in the tube. Introduce a
Httle water by means of a small pipette or by a medicine
dropper, and the mercury will soon fall sHghtly, owing to
the pressure of the water vapor — tension of aqueous
vapor — in the space above the mercury. Warm the
upper end of the tube with the hand or a lighted candle,
and the mercury falls still lower. If it were possible to
boil water in the tube, the pressure of the vapor would be
equal to the pressure of the atmosphere at that instant,
since water boils when it just overcomes atmospheric pres-
Appendix B.
J6i
sure. The correction is often larger than might be sus-
pected and must be made in all accurate work.
Gases are assumed to be saturated when measured
over water, but their volumes are compared when dry.
Hence a ''correction for aqueous tension" as well as for
temperature and pressure must be made before the vol-
umes are comparable.
The following is a —
TABLE OF THE TENSION OF AQUEOUS VAPOR (REGNAULT).
i.
mm.
i.
mm.
t.
mm.
i.
mm.
lO
9.17
16
13-54
22
19.66
28
28.10
947
13-97
•5
20.27
•5
28.93
II
9-79
17
14.42
23
20.89
29
29.78
.5
10.12
14.88
•5
21.53
•5
30-65
12
10.46
18
15-36
24
22.18
30
31-55
.5
10.80
15.85
•5
22.86
•5
32.46
13
II. 16
19
16.35
25
23-55
31
33-41
.5
"•53
16.86
•5
24.26
•5
34-37
14
11.91
20
17-39
26
24.99
32
35-36
.5
12.30
17.94
•5
25-74
•5
36-37
15
12.70
21
18.50
27
26.51
Z3
37-41
13. II
19.07
•5
27.29
•5
38.47
CLASS-ROOM EXERCISE. IIL
1. Would all liquids have the same vapor tension at a givan temDer-
ature 1
2. Why should a gas, the volume of which is to be measured over
water, be saturated before its volume can be accurately determined ?
3. Does the amount of water vapor alter the pressure foi a given
temperature ?
7. Formula for the Reduction of Gas Volumes to Standard
Conditions. — Since gas volumes to be comparable must be
362 Experimental Chemistry.
corrected for temperature, pressure, and aqueous tension
{i.e. reduced to standard conditions), it is convenient to
make the corrections simultaneously by the formula : —
V^ V\P'--) (5)
760(1 +.00366/)
In tnis formula —
V — the corrected volume.
V — the observed volume.
P' = the observed pressure.
/ = the observed temperature.
a = the aqueous tension at /° C.
" Corrected " means " under standard conditions," and
"observed" means "under any other conditions."
This formula is so important, its derivation should be
mastered.
(a) Temperature Correction. — Since all gases expand
; 3 of their
grade, then
oj 3 of their volume at zero for each rise of a degree centi-
V' = F+ ( -^ X F
V273
But — .00366 X A Hence the above equation becomes —
273
V' = V+ (.00366 X 0 X F
= V X (i + .00366 t)
.•.F= "^
(l -f .00366 f)
(b) Pressure Correction. — Since volumes of gases are
inversely proportional to pressure, the volume relations of
Appendix B. 2^3
the same gas under two different pressures may be expressed
by the proportion —
V: V ::P' :P.
Therefore VP = V'P'
or ^"^ "7^
VP'
and since P = 760, .'. V = — ^ —
760
(c) Combined Corrections for Temperature and Pressure. —
To reduce a gas volume to 0° and 760 mm. it must first
I P'
be multiplied by ; — — and then by -— -
^ (i + .00366 t) 760
T P'
Hence V= V x ^ — ^ — — x — -
(i + .00366 t) 760
I r>l
VP
— (6)
760 (i + .00366/)
Formula (6) is used when there is 710 correction for
aqueous tension.
(d) Combined Corrections. — Since the correction for
aqueous tension is made by subtracting from the observed
pressure the aqueous tension in milHmeters of mercury
(found in the Table in § 6), formula (6) becomes-
V =
V'(P' - a)
760 (i + .00366 /)
An example will make the formula more intelHgible.
A student actually found that a mixture of potassium
chlorate and manganese dioxide yielded 2192.4 cc. of
oxygen when the barometer read 757.8 mm. and the ther-
364 Experimental Chemistry.
mometer 21.5° C. The aqueous tension was taken as
19.097. Substituting these vakies, the formula becomes —
^^^2192.4(757.8- 19-097)
760 ( I 4- .00366 X 21.5)
2192.4 X 738.703 „ ^
= ^ ^ ^^ ^ ^ = 1980.6+ CO.
760(1 +.07869)
Therefore, 1980.6 cc. is the volume which would be
occupied at standard conditions by 2192.4 cc. at the ob-
served temperature and pressure.
In this solution notice (i) that the aqueous tension is
subtracted from the barometer reading and tJiat remainder
is multiplied by the observed gas volume, and (2) that the
temperature is multipHed by .00366 and i is added to that
result before multiplying by 760.
PROBLEMS V.
/;/ tJiese problems the gas is to be regarded as dry, unless
otherwise stated.
1. Reduce to standard conditions the following: —
{a) 147 cc. measured at 570 mm. and 136.5° C.
{b) 320 cc. measured at 950 mm. and 91'^ C.
(c) 480 cc. measured at 380 mm. and 68.25° C.
{d) 25 cc. measured at 780 mm. and 27° C.
(<?) 14 cc. measured at 763 mm. and 1 1'' C.
(/) 18 cc. measured at 742 mm. and 14° C.
(g) 0.18 cc. measured at 764 mm. and 20° C.
(//) 375 cc. measured at 745 mm. and 17^ C.
2. A vessel 10 cm. long, 5 cm. wide, and 3.5 cm. deep, is filled with
gas at 770 mm. and 100° C. What is its volume under standard
conditions?
3. A gas under standard conditions has its pressure doubled and its
temperature raised until the volume is the same as the original volume.
What is the final temperature?
Appendix B. 365
4. A gas volume measures 12 liters at o'^C. and 760 mm. What
would its volume be at 800 mm. and 17^ C?
5. A gas measures 150 cc. at 10'' C. and 500 mm. What would be
its volume at 16.4° C. and 520 mm. ?
6. If 852 cc. of hydrogen are measured at 1 1° C. and 760 mm., what
would be the volume at 27^ C. and 900 mm. ?
7. A certain volume of hydrogen measures 250 liters at 745 mm. and
— 15" C. What would be its volume at 20'^ C. and 765 mm. ?
8. What volume under standard conditions would be occupied by
200 cc. of gas at 14° C. and 756 mm. ?
9. Reduce to standard conditions 1328 cc. of gas saturated with
water vapor and measured under the following conditions : —
Pressure, 765 mm.
Temperature, 18" C.
Aqueous tension, 15.357 mm.
10. Correct 99.5 cc. of air for pressure, temperature, and aqueous
tension from the following data: —
Pressure, 756.3 mm.
Temperature, 20.5° C.
Aqueous tension, 17.943 mm.
11. What would be the corrected volume of ']'].'] cc. of saturated
air, if measured at 17.5" C. when the barometer stood at 755.5 mm.,
allowing 14.889 mm. l"or aqueous tension?
12. Reduce to standard conditions in the following cases the gas
volume measured under the given conditions : —
{a) 97.2 cc. {c) 81.2 cc. {/) 100 cc.
20.3° C. 746^ mm. 756 mm
756 mm. 19.5'' C. 24° C.
17.65 a. 16.87 ^- 22.18 a.
{b) ^7 cc. {d) 100 cc. (/) 99.8 cc.
17.5^0. 755.3 mm. 753.2 mm.
755 mm. 18.5° C. 20.3° C.
14.89 a. 15.85 a. 17.65 a.
^66 Experimental Chemistry.
(g) 79 cc. (/) 786CC. (k) 77 a
754 mm. 758.4 mm. '/58.4 mm.
I9.5°C. 22.5^C. i9.5°C.
16.35 rt. 20.29^. 16.87 «.
(/i) 98.8 cc. 0) 97-8 cc.
756.5 mm. 756-3 mm.
20.5° C. 19.5° C.
17.94 a. 16.87 a.
PROBLEMS. VI. (REVIEW.)
1. A certain volume of air measures 150 cc. when the barometeT
stands at 760 mm. On the following clay its volume had decreased
1.52 cc. Calculate the change in the height of the barometer.
2. 146 cc. of air at 10^ C. and 730 mm. are measured over water.
What would the dry air measure under standard conditions?
3. A certain quantity of nitrogen measures 155 cc. at lo""' C, and
under a pressure of 530 mm. What will the volume become at 18.7° C.
and under a pressure of 590 mm. ?
4. A certain weight of air measures a liter at 0° C. How much will
the air expand if heated to 100° C. ?
5. A liter of air at 39° C. is cooled to — 26° C. Find the new
volume.
6. 134 1. of air are heated from - 30" C. to 60° C. Find the new
volume.
7. 50 cc. of air at 10'' C. occupy what volumeat 24' C. ?
8. 100 cc. of air at 12'' C. are heated until they occupy 145 cc.
Find the new temperature.
9. yoo cc. of air at 77^ C. and 1000 mm. measure what at 163° C.
and 872 mm. ?
10. A quantity of air which measures 230 1. at 14° C. and 740 mm.
will measure what at 0° C. and 760 mm. ?
II. MANIPULATION OF GASES.
Certain facts, simple in themselves but easily overlooked,
render work with gases somewhat difficult.
(i) Most gases are colorless and their escape is not
easily discovered.
Appendix B. 367
(2) They are difficult to dry and purify.
(3) Their volume changes quickly and considerably with
changes of temperature and pressure.
(4) They cannot be weighed directly without delicate
apparatus and equable surroundings, hence their weight
must be indirectly determined.
(5) Several exact readings must be made.
Nevertheless, gases present a most fruitful field of study,
and if certain precautions are heeded, fairly accurate results
are possible. Inasmuch as many of the exact experiments
in this book deal with gases, the difficulties above mentioned
warrant further discussion.
(i) Since most gases are collected over water, the whole
system should be watched — height of water, bottle, or
tube, and space occupied by the gas. The water level is
an indicator of the gas level, and the eye should lead the
mind to regard this water level as the essential indicator of
the conditions. Again, all apparatus used in gas experi-
ments should be rigorously tested before the experiment
begins, or, better still, as the arrangement of the apparatus
proceeds. If the gas is to be under increased or dimin-
ished pressure at any time, the connections should be
perfect. Stoppers 7niist fit and tubing must be the exact
size. Do not use wax to stop leaky joints. Never use
vaseline unless it is prescribed in the directions.
(2) Each gas has its pai:ticular drying or purifying
agents, and these agents vary in efficiency. Water vapor
is the most common ingredient to be removed. It is
customary to pass a gas through a bottle or U-tube contain-
ing the purifying agent. The two substances used for the
removal of water vapor are concentrated sulphuric acid and
calcium chloride. Either may be put in a U-tube. The
sulphuric acid drier is more efficient if both limbs are
368
Experimental Chemistry.
nearly but loosely filled with glass beads or fragments of
glass tubing and then enough acid poured into each limb
to drench the glass but not seal the bend. This device
exposes a large surface of acid, permits a steady current of
gas, and relieves the pressure on the generator, since the
gas does not have to overcome
the pressure of the heavy acid in
the bend of the U-tube. If the last
precaution is unnecessary, then a
Drechsel drying bottle (Fig. 114)
is a convenient substitute for the
U-tube. The Drechsel bottle may
be replaced by the apparatus shown
in Fig. 56. If calcium chloride is
used, it should be the fused variety
and be free from powder, as the
latter hardens when moistened and
clogs the tube. Often the U-tube
may be successfully replaced by a
calcium chloride tube or jar, or
even an ordinary glass tube. All
drying apparatus must be protected
from the air by tight corks or glass
plugs (Fig. 47). Ammonia is dried by passing it over lime
or soda lime.
(3) and (4) Since gases cannot be easily weighed, their
weight is calculated from their volume. If the volume
changes and the change in conditions is ignored or not
revealed by the thermometer or barometer, the weight will
apparently change, while the mass is constant. The heat
of the hand or even of the body will often alter a gas
volume, yet the thermometer shows no appreciable change.
Hence, tubes contaiimig gas to be measured must never be
Fig. 114. — Drechsel drying
(or washing) bottle.
Appendix B. 369
touched zvitJi tJu hand. If it is necessary to move them, do
so by the clamp, or with a slip of paper wound around the
tube, or a piece of straw matting, or a test tube holder.
The temperature of a gas is taken either by hanging a
thermometer beside the gas tube, or finding the temperature
of the water over which the gas stands. The latter way is
more reliable. Nevertheless, errors lurk in this method.
The gas must stand over the water at least fifteen viiniites
before the temperature is taken so that the whole system
may be at the same temperature. It is convenient to
measure gases over the water in tall jars which have
been standing in the laboratory for several days before
use. The thermometers may remain in the jars all the
time and the tubes containing the gas may stand in the
water over night, readings being taken at the close of
the session or on the following morning before the regular
laboratory period. Again, the reading of the gas volume
and the temperature must be made simultaneously, and
the thermometer must be read while the bulb is in the
water. A lens is often necessary to read the exact height
of the mercury, especially if the reading is estimated to
the fraction of a degree.
A barometer reading taken at the beginning of the
laboratory period and posted in plain sight is sufficiently
accurate for several hours, unless violent atmospheric
changes are in operation. All readijigs should be verified,
either by the teacher or a student, and any disagreement
remedied before the calculations are made.
CLASS-ROOM EXERCISE. IV.
1. Define law, hypothesis, and theory, as the terms are used in
science. Of what value is each ?
2. Why should gas volumes be reduced to standard conditions ?
APPENDIX C.
I. GLOSSARY OF CHEMICAL TERMS.
A WORKING familiarity with the language of Chemistry is indispensable to
intelligent reading, speaking, and note-taking. Many of the following defini-
tions are only suggestive; all should be interpreted by the larger and more
recent text-books.
Acid, adj. Sour; reddens blue litmus,
n. A substance which neu-
tralizes a base, decomposes a car-
bonate, and reddens blue litmus.
Acidic, adj. Same as Acid (adj.).
Affinity, n. See Attraction.
Agitate, v. To shake; to mix by
violent shaking.
Alkali, n. Popularly a substance
which neutralizes an acid, and turns
red litmus blue. A strong hydroxide.
A caustic substance, A narrow use
of the term Base,
Alkaline, adj. Having the properties
of an alkali, Basic.
Allotrope, n. An allotropic form,
Allotr5p'-ic, adj. Illustrating Allot-
ropy,
Allotropism, A115t'-ropy, n. Vari-
ation in the physical properties
assumed by some elements, espe-
cially sulphur, phosphorus, and
carljon.
Alloy', n. A compound or mixture
of metals having metallic properties,
e.g. brass.
Amalgam, n. Alloy in which mer-
cury is a constituent, e.g. sodium
amalgam.
Amorphous, ailj. Non-crystalline, eg.
amorphous sulphur. Often applied
to powders, as amorphous carbon.
Analysis, n. A separation ; a partial
or complete decomposition. The
opposite of synthesis. The deter-
mination by a chemical operation
of the presence of a substance, as
qualitative analysis, or of the amount
of a sulistance, as quantitative,
gravimi;tric, or volumetric analysis.
Anhydride, n. An t)xide which unites
with water to form an acid, e.g.
sulphurous anhydride, carbonic an-
hydride. Often called an acid
anhydride.
Anhydrous, adj. Applied to a com-
pound from which water of crystal-
lization has been removed, e.g.
anhydrous copper sulphate. Any
" water-free " substance, as anhy-
drous alcohol or ether.
Anion, n. A negative or acidic ion.
370
Appendix C.
371
Anneal, v. To render less brittle by
slow cooling.
Aqua, n, Latin for water. Aqua
?-t'gia means royal water.
Aqueous, adj. Relating to water, as
aqueous vapor, aqueous tension, or
an aqueous solution as distinguished
from an alcoholic solution.
Aspirate, v. To draw or force air
or any other gas through an ap-
paratus.
Aspirator, n. A filter pump. A
bottle to force or draw a gas
through an apparatus.
Atmospheric, adj. Relating to the
atmosphere.
Attraction, n. Affinity. Force which
causes and maintains combination
of elements.
Bare, adj. Same as free in " free
flame." Bare flask is one not pro-
tected by a gauze or a similar cover-
ing.
Base, n. A substance which neu-
tralizes an acid. A loose general
term including metallic oxides and
hydroxides. See Alkali and Acid.
Basic, adj. A synonym of alkaline.
The opposite of Acid (adj.) and
Acidic. Mono-, di-.tribasic, etc., in-
dicate the number of replaceable
hydrogen atoms in a molecule of
an acid. Applied to a class of
salts.
Bi-. A prefix meaning two. Same
as Di-.
Binary, adj. Applied to a compound
of two elements.
Bump, V. Loose term meaning sud-
den or violent boiling.
Cation, n. A positive or metallic
ion.
Caustic, adj. Burning, corroding,
e.g. caustic potash, lunar caustic
(silver nitrate).
Chemical, adj. Relating to chemis-
try, e.g. chemical change. Implies
deep-seated action.
n. Chemical compound ; re-
agent ; any substance used in the
laboratory.
Combustible, adj. Inflammable, burns
with flame, ignites easily.
Combustion, n. Chemical action
accompanied by light and heat.
Strictly, a union with certain ele-
ment.s, <?.^. oxygen; broadly, a burn-
ing.
Commercial, adj. Not chemically
pure; as ordinarily furnished by
dealer.
Compound, n. Product of the com-
bination of two or more elements
in definite proportions. Often
means the opposite of mixture.
Concentrated, adj. Strong. Oppo-
site of Dilute.
C. P. Abbreviation of " chemically
pure." Free from impurities.
Crystal, n. A solid, bounded by geo-
metricaUy arranged surfaces and de-
posited usually from a vapor or
saturated solution.
Crystalline, adj. Consisting of crys-
tals or resembling a .crystal. Oppo-
site of Amorphous.
Crystallize, v. To change into or to
form crystals.
Crystallized, adj. Having the form
of crystals. Not powdered, or dis-
solved, or amorphous.
372
Experimental Chemistry.
Decant, v. To pour a supernatant
liquid from a vessel without disturb-
ing the sediment.
Decantation, n. Act of decanting,
as " to wash by decantation."
Decompose, v. To separate into
simpler parts.
Decomposition, n. Act of decom-
posing.
Deflagrate, v. To burn suddenly.
Deflagrating, adj. Applied to a spoon
in which deflagration occurs. Some-
times applied to a mixture, or to a
single substance.
Deflagration, n. Rapid and sudden
burning.
Deliquesce, v. To absorb moisture
rapidly from the air.
Deliquescence, n. Liquefaction or
dissolving due to absorption of mois-
ture from the air.
Deliquescent, adj. Applied to sub-
stances which deliquesce.
Diffusion, n. Mixing, intermingling,
especially of gases.
Di-. A prefix. See Bi-.
Digest', V. To soften by heat and
moisture ; to heat in a closed
vessel preparatory to chemical
operations.
Dilute, adj. Weak. See Concentrated.
V. To weaken, usually by addi-
tion of water.
Displacement, n. Collection of a gas
by alltjwing it to fall or rise into a
vessel. The former method is down-
ward displacement, the latter is up-
ward displacement. Substitution of
one element for another in a com-
pound ; replacement.
Distil, V. To evaporate a liquid and
recover it by subsequent condensa-
tion.
Distillate, n. The condensed vapor.
Distillation, n. Act of distilling.
Destructive distillation is often ap-
plied to the preparation of coal gas,
or to the manufacture of various
liquids from wood.
Ductile, adj. May be drawn into wire.
Dyad, n. Element with valence of
two. A bivalent element.
Ebullition, n. Boiling or bubbling of
a lifiuid due to escaping gas.
Effervescence, n. The bubbling of a
gas through a liquid. The result of
chemical action, and usually occur-
ring without the application of heat.
A^ot hoiliug. E.g. the escape of car-
bon dioxide from the decomposition
of a carbonate by an acid.
Effloresce, v. To lose easily water of
crystallization.
Efflorescence, n. A turning to powder
or loss of luster due to escape of
water of crystallization at the ordi-
nary temperature.
Efflorescent, adj. Applied to sub-
stances which effloresce.
Electrode, n. Pole or terminal of a
wire conducting an electric current.
Element, n. A substance which has
not yet been decomposed by any
means known to science.
Equivalence, n. See Valence.
Evaporate, v. To convert into vapor.
To heat until all liquid has been
turned into vapor; eg. "evaporate
to dryness," means to heat until no
liquid remains.
Evolve, v. See Liberate.
Apperxdix C.
373
Evolution, n. Liberation; escape as
a result of chemical action, as evolu-
tion of oxygen from potassium
chlorate.
Excess, n. A large quantity. More
than a given quantity. "A slight
excess " is a little more than is
necessr.ry to produce certain results.
Extract, v. To remove the essential
part of a mixture, usually by dissolv-
ing it in hot water, or alcohol, or a
special solvent, and then filtering.
Filter, v. To separate a solid from a
liquid; to purify.
n. Loosely appHed to the filter
paper, or sometimes to the whole
apparatus by which filtration is ac-
complished.
Filtrate, n. The liquid which passes
through the paper. Opposite of
Precipitate (n,).
Fixed, adj. Not volatile.
Formula, n. A group of symbols
expressing the composition of a
compound.
Formulae, )
Formulas, i
Free, adj. Not opposed. Free flame
touches the object. Free channel
for gases has no obstruction.
Fuse, v. To melt.
Fusible, adj. Capable of being
melted.
Gas, n. The aeriform state of matter.
Popularly, a mixture of gases, eg.
illuminating gas. See Vapor.
Gelatinous, adj. Like jelly. Applied
to precipitates like aluminium hy-
droxide.
n. Plural of formula.
Generate, v. See Liberate and Evolve.
Generator, n. Apparatus in which
gases are generated.
Gram, n. Unit of weight in the
Metric System. Equal in weight to
looo mg. and to i cc. of water at
4°C.
Granulated, adj. In sma'l pieces
varying in size from a single crystal
of sugar to a hazel nut. Applied
to zinc usually used to generate hy-
drogen.
Gravimetric, adj. By weight.
Halogen, n. and adj. A group of ele-
ments consisting of chlorine, bro-
mine, iodine, and sometimes fluorine.
Halogen means a " sea salt pro-
ducer." It is applied to the mem-
bers of this group because their
sodium salts resemble sea sal.t.
Chlorides, bromides, etc., are some-
times • called halogen salts, or
halides.
Hard, adj. Applied to water v/hich
contains salts of calcium or mag-
nesium, or both.
Hydrate, n. Used inaccurately as a
synonym of hydroxide.
Hydrated, adj. Combined with water,
opposite of dehydrated and anhy-
drous. "Opal is hydrated silica."
Hydro-. Prefix meaning («) an acid
containing no oxygen, or {b') a com-
pound containing hydrogen, as in
the word hydrocarbon.
Hydroxyl, n. The radical OH.
Hygroscopic, adj. Applied to bodies
which readily absorb moisture from
the atmosphere, e.g. potassium car-
bonate.
374
Experimental Chemistry
Hypo-. A prefix meaning below.
Applied tJ a compound in which
the distinctive element has a low,
and usually the lowest, valence, 6 g.
hypochlorous acid.
-ide. Ending of binary compounds.
Inflammable, adj. Applied to a vapor
or gas which burns readily. Ihe
vapor of such liquids as carbon bi-
sulphide, ether, alcohol, naphtha,
and benzine is inflammable.
Infusible, adj. Not fusible; not easily
melted.
Ignite, V. To set on fire, to burn, to
kindle.
Ignition, n. The act of igniting. Used
with tube, e.g. ignition tube, a tube
which will not melt at a high tem-
perature.
Inorganic, adj. Opposite of organic;
relating to mineral substances, as in-
organic or mineral acids. Applied
to all compounds except the com-
plex derivatives of carbon.
Insoluble, adj. Not soluble.
Interact, v. To react, to undergo
chemical action, to act.
Interaction, n. Reaction. Especially
that chemical action in which there
is double decomposition. Tlie re-
action consisting of both decom-
position and combination.
Ion, n. An electrically charged atom
or atomic group. See Chapter IX.
See Anion and Cation.
Ionization, n. The act of breaking
up into ions.
Liberate, v. To set free, to evolve,
eg. to liberate hydrogen from an acid.
Limb, n. An arm or portion of a
tube, e.g. limb of a U-tul)e.
Liter, n. Unit of volume in the Metric
System. A cubic decimeter, equiva-
lent to looo cc, a kilogram, 2.2 lb.,
or 1.06 quarts of water.
Malleable, adj. May be rolled into
sheets. Compare Germ, hdiiimer-
bar.
Manipulation, n. Experimenting.
" Handling " apparatus. The manual
part of an experiment.
Metal, n. A substance, generally an
element, possessing characteristic
properties. They are usually hard,
their smooth surfaces reflect light,
they are good conductors of heat
and electricity. All, except mer-
cury, are solid at the ordinary
temperature. A base-forming ele-
ment. An alloy.
Molecular, adj. Applying to a mole-
cule, as molecular weight. A
weight proportional or equal to
the molecular weight, as molecular
proportions.
Molecule, n. A group of atoms
chemically combined. The small-
est particle of a compound pos-
sessing the properties of the com-
pound.
Molten, adj. Melted, fluid or semi-
fluid, thick, viscous. Applied to a
hot mass.
Monad, n. Element with a valence
of one. A univalent element.
Monatomic, adj. One atom to the
molecule.
Mono-. A prefix meaning one. Same
as Uni-.
Appendix C.
375
Nascent, adj. Applied to an ele-
ment the instant it is evolved or
released from a compound. An
active state.
Neck, n. The slender part of a vessel,
e.g. the long stem of a retort, the
part of a bottle or flask to which a
stopper is fitted.
Neutral, adj. Indifferent to litmus,
as, "v^'ater is neutral." Neither
acid nor alkaline, as, "sodium
chloride is neutral."
Neutralization, n. The interaction
of an acid and base resulting in the
formation of a salt.
Neutralize, v. To make neutral.
Nomenclature, n. The language of
chemistry. A system of terms and
terminology. The naming of ele-
ments and compounds.
Non-metal, n. An element not a
metal. An acid-forming element,
such as sulphur or nitrogen.
Normal, adj. Standard. Normal
Pressure = 760 mm. Normal Tem-
perature = o" C. Normal volume
is one measured at or reduced to
these conditions. Salt containing
no replaceable hydrogen, such as
K2SO4.
Notation, n. The written language of
chemistry, especially the symbols
and formulas.
Organic, adj. Applied to natural or
artificial products from animal or
vegetable life. More specifically,
complex compounds of carbon.
Opposite of Inorganic.
Oxidation, n. Act of causing c> com-
bination with oxygen. j
Oxide, n. Compound of oxygen with
one other element.
Oxidize, v. To change to an oxide.
To add oxygen. To increase the
valence. To withdraw hydrogen.
Oxidizing Agent. A compound
yielding oxygen easily, as nitric acid
or potassium chlorate.
Per-. A prefix meaning beyond.
Originally denoting the highest
valence, as persulphuric acid — the
acitl in which sulphur has the highest
valence. Often denoting that an
element has a higher valence than
in similar compounds.
Physical, adj. Relating to Physics, e.g.
physical change. Implies a change
of properties which can be readily
seen or detected by physical means.
Pneumatic, adj. Relating to gases, e.g.
pneumatic trough, a vessel for col-
lecting gases. The Pneumatic Period
of chemistry was characterized by
the discovery of several gases.
Precipitate, v. To produce an insol-
uble solid by mixing solutions.
n. The solid so produced. The
solid retained by the paper in filter-
ing. Literally, a substance thrown
down, because the insoluble solid
often falls to the bottom of the vessel.
Proto-. ) A prefix meaning first or
Prot-. i lowest. Denoting the first
of a series of compounds, as pro-
tochloride, protoxide. Sometimes
used synonymously with Mono.
Qualitative, adj. Referring only to
the kind of matter.
Quantitative, adj. Referring to exact
weight or volume.
376
Experimental Chemistry.
Radical, n. A group of different
atoms acting as a single element in
a series of compounds and incapa-
ble of independent existence, as
NH4, OH, CN.
React, V. To act chemically, to
interact.
Reaction, n. A chemical change,
usually applied to a complete change.
Behavior toward litmus, as acid
reaction. Loosely used as syngny-
mous with equation.
Reagent, n. A mixture or compound,
often in solution, used to produce a
desired chemical change, e.g. a
group reagent. Often applied to
any *' chemical."
Receiver, n. Any vessel which catches
or receives a distillate.
Reduce, v. To remove oxygen. To
lower the valence. To add hydro-
gen. See Oxidize.
Reducing Agent. A compound ab-
stracting oxygen easily, such as
hydrogen or carbon.
Reduction, n. Act of reducing, or
removing oxygen.
Replace, v. See Substitute.
Residual, adj. Whatever remains, as
"residual gas."
Residue, n. A solid left after filtra-
tion or evaporation. A final product.
A remainder.
Salt, n. A compound formed by the
replacement of the hydrogen of an
acid by the metal of a base. Popu-
lar name of sodium chloride.
Saturated, adj. Containing a maxi-
mum amount under given conditions.
A saturated solution remains un-
changed on the addition of more
solid. State in which all chemical
attraction is satisfied. A compound
to which nothing can be chemically
added, such as water.
Seal, V. To close. To seal a tube is
to melt the open end (or ends)
together. To seal a flask is to close
it — usually air tight.
Sesqui-. A prefix meaning one and
one-half. Applied to compounds
containing three atoms of one ele-
ment and two of another, as Fe^O.s.
Soluble, adj. Applied to all sub-
stances which dissolve (usually in
water). Opposite of Insoluble.
Solute, n. The solid which the liquid
dissolves.
Solvent, n. The liquid in which the
solid dissolves.
Stable, adj. Not easily decomposed.
Standard. See Normal.
Sublimate, n. A substance sublimed.
Sublimation, n. Act of distilling a
solid. Usually, converting a soUd
into a vapor which immediately
solidifies in a purer state higher up
in the containing vessel.
Substitute, v. To replace one element
or group by another.
Substitution, n. Replacement in a
compound.
Sulpho-. Prefix denoting a compound
of sulphur.
Supernatant, adj. Applied to a liquid
standing over a solid, especially a
precipitate. See Decant.
Supersaturated, adj. Applied to a
solution containing more solid than
it would contain if an undissolved
portion were present.
Appendix C.
377
Suspension, n. Used in the phrase " in
suspension" to describe a solid distrib-
uted in fine particles through a liquid.
Symbol, n. The abbreviation of the
name of an element.
Synthesis, n. Putting together. Op-
posite of Analysis.
Tetra-. A prefix meaning four.
Tetrad, n. Element with the valence
of four. A tetravalent or quadriv-
alent element.
Thio-. Same as Sulpho-.
Treat, v. To add to, to mix with, to
allow to react with; e.g. "treat with
hydrochloric acid " means to add
hydrochloric acid.
Tri-. A prefix meaning three.
Triad, n. Element with the valence
of three. A trivalent element.
Tubulated, adj. Provided with a tu-
bulure. Applied to a glass-stop-
pered vessel, e.g. a retort.
Tubulure, n. A short tube-like open-
ing in the bulb of a retort or side of
a bottle.
Uni-. Same as Mono-.
Unstable, adj. Easily decomposed.
Valence, n. The replacing or com-
bining power of an element in terms
of hydrogen.
Vapor, n. Gaseous form of a substance
normally liquid or solid. Any gas
near its condensing point. Popu-
larly, moisture in the air. Loosely,
any cloud-like substance.
Volatile, adj. Applied to all sub-
stances which easily change into a
vapor.
Volumetric, adj. By volume. See
Gravimetric.
-yl. Suffix of some radicals, e.g.
hydroxyl.
II. BIBLIOGRAPHY.
Successful experimental work is enlivened and fixed in mind by intelligent
use of the best books. A chemical library should be in every laboratory ;
pupils should have free access to books, and should be encouraged to consult
them frequently at all stages of the work. The following list contains only a
portion of the books which the author regards as useful to both teacher and
pupil in elementary work. The starred (*) titles indicate books which are
intended for the teacher's use. In most cases the book which is accurate and
simple in statement, full but not ponderous, and new, is first in its own list;
such a discrimination is, of course, largely personal. The prices in some cases
are subject to discount. Special books, monographs, and pamphlets are noted
under separate chapters in the Teacher's Supplement.
I. DICTIONARIES.
* I. Dictionary of Chemistry, Watts, Morley, and Muir. Longmans,
Green, & Co. 4 vols. $65.
* 2. A Dictionary of Chemical Solubilities, Comey. The Macmillan
Co. 515 pp. $5, net.
378 Experimental Chemistry.
♦3. Scientific American Cyclopedia OF Receipts. Munn & Co., New
York. $^. (Contains 12,500 useful receipts.)
*4. CilEMiKER Kalendar, BiedenTiann. Springer, Berlin (annually).
$1. (Contains numerous tables.)
II. LARGE DESCRIPTIVE.
I. Textbook of Inorganic Chemistry, Newth. Longmans, Green, &
Co. 682 pp., 146 illus. ^1.75..
* 2. Treatise on Chem'STRy, Roscoe and Schorlemmer. D. Appleton &
Co. 2 vols. ^8. (Inorganic.)
3. Chemistry (Advanced Course), Remsen. Henry Holt & Co.
850 pp. ^2.80. (Includes experiments.)
♦4. Principles of Chemistry, Mendeleeff. Edited by Greenaway.
Longmans, Green, & Co. 2 vols. ^10.
5. General Inorganic Chemistry, Freer. Allyn & Bacon, Boston.
559 PP- ^3- (Includes experiments.)
III. SMALL DESCRIPTIVE.
1. Elements of Chemistry, Shepard. D. C. Heath & Co. 366 pp. $1.40.
2. Chemistry (Briefer Course), Remsen. Henry Holt & Co. 435 PP-
;^i.i2.
3. Elementary Manual of Chemistry, Storer and Lindsay. American
Book Co. 453 pp. $1.20.
4. Elements of Chemistry, Williams. Ginn & Co. 397 pp. ^i.io.
IV. THEORETICAL AND PHYSICAL.
I. Physical Chemistry for Beginners, Van Deventer. Translated
by Boltwood. John Wiley & Sons, New York. 154 PP- $^-S^-
*2. Outlines of General Chemistry, Ostwald. Translated by Walker.
The Macmillan Co. 386 pp. ^3.50.
* 3. Outlines of Theoretical Chemistry, Meyer. Translated by
Bedson and Williams. Longmans, Green, & Co. 232 pp. $2.50.
4. Principles of Theoretical Chemistry, Remsen. Lea Bros., Phila-
delphia. 361 pp. $2.50.
* 5. Outline of the Theory of SolutioNj Morgan. John Wiley &
Sons. ^i.
*6. Elements of Physical Ciikmisiry, Morgan. John Wiley & Sons.
299 pp. $2.
* 7. Scientific Foundations of Analytical Chemistry, Ostwald,
The Macmillan Co. 216 pp. ^2.
Appendix C. 379
*8. Solutions, Ostwald. Translated [ry Aluir. Longmans, Green, & Co,
310 pp. ^2.
9. Chemical Theory for Beglnners, Dobbin and Walker. The Mac-
millan Co. 236 pp. $.70.
10. Scientific Memoirs, J. S. Ames, Editor. American Book Co.
(a) Modern Theory of Solution, Jones, 134 pp. ^i.
{/>) Fl'ndamental Laws of Electrolytic Conduction, Good-
wm. 96 pp. $1.
*ii. Atomic Theory, Wurtz. D, Appleton & Co. 344 PP- $i-50.
*i2, Lntroduction to Physical Chemistry, Walker. Ths Macmillan
Co. 332 pp. $3-
Books in this class are constantly appearing. Of the above, 2, 3, and 8
are excellent, though not the latest; i is good, though brief; 7 and 10 are
special, but exact; 12 is the latest and from many standpoints the best for a
first study of the whole subject.
V. HISTORICAL.
I. History of Chemistry, Venable. D. C. Heath & Co. 172 pp. $1.
* 2. History of Chemistry, von Meyer. Translated by M'Gowan. The
Macmillan Co. 544 pp. $4-50-
* 3. Essays in Historical Chemistry, T. E. Thorpe. The Macmillan
Co. 381 pp. $2.20. (Biographies of eleven famous chemists.)
4. Gases of the Atmosphere, Ramsay. The Macmillan Co. 240 pp.
$2. (Covers pneumatic period.)
5. Alembic Club Reprints, Chemical Bub. Co., Easton, Pa. ^.50 each.
1. Experiments on Magnesia Alba.
2. Foundations of the Atomic Theory.
3. Experiments on Air.
4. Foundations of the Molecular Theory.
5. Extracts from Micrographia.
6. Decomposition of the Fixed Alkalies.
7. Discovery of Oxygen. Part i.
8. Discovery of Oxygen. Part 2.
9. Elementary Nature of Chlorine.
10. Researches on the Arseniates, Phosphates, and Modifications of
Phosphoric Acid.
11. Increase in Weight of Tin and Lead on Calcination.
12. Liquefaction of Gases.
13. Early History of Chlorine.
14. Researches on Molecular Asymmetry.
jSo Experimental Chemistry.
6. Century Science Series. The Macmillan Co. $1.25 each.
1. Humphry Davy, T, E. Thorpe.
2. John Dalton, Roscoe.
3. Michael Faraday, Thompson.
4. Justus von Liebig, Shenstone.
7. Heroes of Science — Chemists, Muir. E. & J. B. Young &. Co.,
N.Y. 350 pp. $1.50. Popular but accurate.
8. Heroes of Science — Physicists, Garnet. (Same as 7.)
* 9. History of the Warfare of Science with Theology, White.
D. Appleton & Co. 2 vols. ^5. Vol. i. Chap. 12, contains a valuable account
of the growth of chemistry.
10. Faraday as a Discoverer, Tyndall. D. Appleton & Co. 171 pp.
VI. ORGANIC.
1. Organic Chemistry, Perkin & Kipping. J. B. Lippincott Co., Phila-
delphia. 552 pp. (2 vols.). $2.
2. Organic Chemistry, Remsen. D. C. Heath & Co. 364 pp. $1.30.
3. Laboratory Manual of Organic Chemistry, Orndorff. D. C.
Heath & Co. 78 experiments. $.40. Designed to accompany 2 (supra).
* 4. The Practical Methods of Organic Chp:mistry, Gattermann.
Translated by Shober. The Macmillan Co. 330 pp. $1.60.
* 5. Organic Chemistry, Richter. Translated by E. F. Smith. P. Blak-
iston's Son & Co. Vol. I., 625 pp. $3. Vol. II., 671 pp. $3.
VII. MISCELLANEOUS.
1. Outlines of Industrial Chemistry, F. H. Thorp. The Macmillan
Co. 528 pp. ^3.50.
2. Methods of Glass Blowing, Shenstone. Longmans, Green, & Co.
96 pp. ^.50.
* 3. Physical Properties of Gases, Kimball. Houghton, Mifflin, & Co.
238 pp. $1.25.
4. Minerals and how to study Them, E. S. Dana. John Wiley &
Sons. 368 pp. ^1.25.
♦5. Questions ON Chemistry, Jones. The Macmillan Co. ^.75. (About
1500 unanswered questions.)
6. Fermentation, Schiitzenlierger. D. Appleton & Co. 331 pp. ^1.50.
7. Chemistry of Daily Life, Lassar-Cohn. J. B. Lippincott Co. ^1.75.
Appendix C. 381
8. Aluminum and Aluminum Alloys. The Pittsburg (Pa.) Reducing
Co. 266 pp. ^1.50.
*g. Development of the Periodic Law, Venable. Chemical Pub. Co,
321 pp. ^2.50.
10. Chemistry of Common Life, Johnston. D. Appleton & Co.
592 pp. $2.
* II. Manual of Chemical Technology, Wagner. D. Appleton & Co.
568 pp. ^7.50.
12. The New Chemistry, Cooke. D. Appleton & Co. 393 pp. $2.
13. The Chemistry of Cookery, Williams. D. Appleton & Co. ^1.50.
14. Story of a Piece of Coal, Martin. D. Appleton & Co. 165 pp.
$.40.
15. Coal and the Coal Mines, Green. Houghton, Mifflin, & Co.
240 pp. $.75.
16. Inorganic Chemical Preparations, F. H. Thorp. Ginn & Co.
238 pp. $1.50.
17. Laboratory Practice, Cook.e. D. Appleton & Co. 192 pp. $1.
18. Liquefaction of Gases, Hardin. The Macmillan Co. 244 pp. $1.50-
III. APPARATUS AND CHEMICALS.
Various specific directions regarding laboratory supplies are given in the
Teacher's Supplement, in the list of Reagents, and in the directions for the ex-
periments. The subjoined lists are inserted simply to assist in the preparation
of an order. Quantities and prices have been intentionally omitted, since it is
believed that both dealer and purchaser will make more fehcitous arrange-
ments, if both are unrestricted by a list unavoidably incomplete. No teacher
can judge fcr another without knowing the working conditions, and no author
has a right to insist that a dealer shall always furnish an article for a fixed sum.
The market price fluctuates, and prices vary with quantity and quality as well
as with the method of ordering. The author, at his own suggestion, has lodged
with the L. E. Knott Apparatus Co., 16 Ashburton Place, Boston, Mass., in-
formation regarding the quantities of apparatus and chemicals used by his own
classes. It is hoped that teachers will avail themselves of the opportunity to
correspond with both author and dealer in the preparation of order lists.
I. INDIVIDUAL APPARATUS.
This list includes the apparatus constantly used by a single student, who
should be provided with each piece, and be held responsible for its return in
382
usable condition.
than $6.75.
Experimental Chemistry.
The apparatus, as JtcmizeJ below, should not cost more
I
100
5
Test tubes, 6xf.
Test tubes, 8x1.
Test tube holder.
Test tube rack.
Test tube brush.
Bunsen burner.
Feet rubber tubing, \ in. int. diam.
Wing-top burner.
Blowpipe.
Blowpipe tube.
Horn spoon, 5 in.
Filter papers, 4 in.
Bottles, salt mouth, 250 cc. (8 oz.)
Funnel, 2| in.
Safety tube.
Flask, 500 cc.
Two-hole rubber stopper, to fit
above.
One-hole rubber stopper, to fit
above.
Evaporating dish, 3 in.
Crucil)le with cover. No. o.
Crucible block.
Casserole, 125 cc.
Pair iron forceps.
Triangular file.
Mortar and pestle, 3 in.
Ignition tube, 4 in. x ^ in.
Deflagrating spoon.
Towel.
II. SPECIAL APPARATUS.
This list includes the apparatus used occasionally, but of such a nature that
it should be given only when demanded, and returned as soon as used. The
only items costing over $2.00 in this list are the burettes (if with glass
stop-cock) and the platinum crucible. The latter need not cost over $8.00
(see Exp. 140). In some cases cheaper forms than those called for may answer,
but in the author's judgment it is imprudent to economize on permanent
equipment.
Beads.
Beakers, 2.
Bottle, Drechsel.
Burettes.
Calcium chloride tube.
Chlorine decomposition tube.
Condenser.
Crucible, Hessian.
Dish, lead.
glass (shallow).
Eudiometer, 100 cc.
Flask, Erlenmeyer, 250 cc and 125 cc.
Funnel, dropping-
Gas tube, 100 cc. and 50 cc.
Marchard tube.
Pinchcock, Mohr.
Ilofmann
Pipe, iron.
Pipette.
Platinum crucible.
foil.
wire.
Thermometer.
Thistle tube top.
Trap, bulb (for Exp. 92).
steam.
U-tube.
Appendix C.
3^3
HI. GENERAL
This list includes all apparatus inte
$io or over arc marked with a *.
*Air bath.
* pump.
Asbestos boards, 6x6 in.
Aspirator bottle.
Balance, horn pan.
Barometer.
Battery.
Blast lamp.
Bottles, wide mouth, 2 oz. and 4 oz.
Candles.
Chimneys, student lamp.
♦Combustion furnace, 10 burner.
Corks, assorted.
Cork borers, set.
Cylinders, graduated (set) 7.
Dowel rod.
Electrolysis apparatus, io| in.
Emery paper.
File, round.
P'ilter pump.
*Gas holder.
Gauze, iron, 6x 6 in.
Glass plates, 6x6 in.
tubing, soft.
tubing, hard.
rod.
Graduates. See Cylinders.
Induction coil.
Iron stands, with 3 rings and 2 clamps
APPARATUS.
nded for general use. Items costing
Jars, battery, 6 x 8 in.
tall, 3 x 15 in.
waste.
Labels (Dennison, 223, 205, 219, 201).
Magnets.
Magnifying glass.
Meter stick.
Pneumatic troughs.
Rule, 30 cm.
Sand bath pans, 4 in.
Scales, trip.
Scissors, 6 in.
Shears, 12 in.
* Spectroscope.
Stoppers, rubber, with i and 2 holes.
Tapers, box.
Thread, cotton.
Triangle, porcelain.
Tripods.
Tubing, rubber (ordinary and press-
ure), i, t\, i in. (int. diam.).
Water baths, 5 in.
Weights (for horn pan and trip scales).
Wire cutter.
Wood (splints).
Wooden blocks, 6x6x1 in.
6x6x f in.
4x4x1 in.
I Woulff bottle, 2-neck.
IV. CHEMICALS.
This list includes all chemicals needed for this book.
Acid, acetic. Acid, oxalic.
citric. sulphuric,
hydrochloric tartaric,
nitric.
384
Experimental Chemistry.
Alcohol, ethyl.
methyl.
Alum, chrome.
potassium.
Aluminium, metal.
sulphate.
Ammonium, chloride.
hydroxide.
nitrate.
oxalate.
sulphide.
Antimony, metal.
Arsenious 'oxide.
Asbestos, shredded.
Barium chloride.
hydroxide.
Beeswax.
Bleaching powder.
Borax (powd.).
Brass wire. No. 24.
Calcite (lump).
Calcium carbide.
carbonate (marble),
chloride,
fluoride,
oxide (lime),
sulphate.
Carbon bisulphide.
Charcoal, animal (powd.).
lump.
wood (powd.).
Cogl, scft.
Cobalt nitrate.
Cochineal.
Coin (silver).
Copper borings,
nitrate,
oxide (granulated),
sheet,
sulphate (cryst.).
sulphate (dehydrated).
wire.
Cotton (absorbent).
Cream of tartar.
Ether.
Galena.
Gelatine.
Glass wool.
Glycerine.
Gold leaf (book).
Hematite.
Indigo
Iodine.
Iron, by hydrogen.
chloride (2V).
filings.
pyrites.
sulphate (ous),
sulphide (ous).
Kerosene.
Lead acetate.
carbonate.
dioxide (peroxide).
nitrate.
monoxide (litharge).
sheet,
tea.
tetroxide.
Limonite.
Litmus, cubes,
paper.
Magnesium carbonate (magnesite),
oxide.
powder.
ribbon.
sulphate.
Manganese dioxide.
sulphate.
Appendix C.
385
Mercury.
Mercuric chloride.
nitrate.
oxide.
sulphide (cinnabar).
Mercurous nitrate.
Mustard.
Parafifine.
Phenolphthalein.
Phosphorus, yellow (worm-shaped),
Picture cord (iron).
Potassium, metal.
bromide.
carbonate,
chlorate (cryst,).
chlorate (powd.).
chloride.
chromate.
dichromate.
ferricyanide.
ferrocyanide.
hydroxide.
iodide.
nitrate.
permanganate.
sulphate.
sulphocyanide.
Sand.
Selenite (g}'psum, cryst.).
Shellac.
Siderite.
Silver nitrate.
Soap.
Solder.
Soda lime.
Sodium, metal.
acetate.
bicarbonate.
carbonate.
chloride.
hydroxide.
hyposulphite (thiosulphate).
nitrate.
nitrite.
phosphate (disodium phos-
phate).
silicate.
sulphate.
Stannous chloride (tin crystals).
Starch.
Sugar, cane,
grape.
Sulphur, flowers.
roll.
Tartar emetic.
Tin, granulated.
Turpentine.
Vaseline.
Vinegar.
Water, distilled.
Wood ashes.
Zinc, dust.
granulated.
oxide.
pure.
sheet.
sulphate.
IV. REAGENTS.
The following apparatus is helpful in the preparation of reagents : —
Porcelain dish, 14 in.
Funnel, 7 in.
Graduate, 1000 cc.
I Graduate, 500 cc.
I Iron dish (frying dish).
I Tin dish (8 in. sauce pan).
386
Experimental Chemistry.
I Mortar, 7 in.
I Set of hydrometers.
Glass rod (i cm. in diam.).
Glass wool.
Asbestos, shredded.
Filter paper, 12 in.
Filter pump, flask and Witt plate.
Lithia bottles.
Glass stoppered bottles (5 pint acid),
The preparation of large quantities of a solution is hastened by using a
dissolver. One is described in the American Chemical Journal, Vol. 22,
No. 5 (November, 1899), p. 407.
The following list of reagents includes all those required for this book. A
few others have been inserted for completeness. Reagents not mentioned
are either unimportant or are fully explained in the experiments requiring
their use. If a reagent is not mentioned in this list, consult the index.
Acetic Acid, C2H4O0.— (i) Use the
commercial acid (sp. gr. 1.04).
(2) Add 1000 cc. of glacial acetic
acid to 2500 cc. of water.
Alcohol, C2HCO. — Use the commer-
cial liquid (sp. gr. .815 ; contains
95 per cent of alcohol). See Ap-
pendix C, Table X.
Alum. — Solid. Use the commercial
crystals.
Solution. Dissolve 100 gm. of the
commercial alum in looo cc. of
water. This proportion answers for
any alum.
Aluminium Chloride, AICI3. — (i)
Dissolve 100 gm. of the commer-
cial salt in 1000 cc. of water.
(2) Precipitate aluminium hy-
droxide by adding ammonium hy-
droxide to a solution of alum. Wash
the precipitate free from acid, and
add a slight excess of concentrated
hydrochloric acid.
Ammonium Carbonate, (NH4)oC03.
— Dissolve 250 gm. of the commer-
cial salt in icxx) cc. of water, with-
out heating, and add 100 cc. of
commercial ammonium hydroxide.
Filter. Keep tightly stoppered.
Ammonium Chloride, NH4CI. — Solid.
Use the commercial salt.
Solution. Dissolve 100 gm. of
the commercial salt in 1000 cc. of
water. Filter.
Ammonium Hydroxide, NH4OH. —
Concentrated. Use the commercial
liquid (26° or sp. gr. = ,96).
Dilute. Add three volumes of
water to one volume of commercial
ammonium hydroxide.
Ammonium Molybdate,(NH4)2Mo04.
— (i) Add 100 gm. of molybdenum
oxide (M0O3) to 400 cc. of com-
mercial ammonium hydroxide, and
stir until dissolved. Cool, and add
all at once 1 000 cc. of nitric acid
(sp. gr. 1.2), using no heat.
(2} Dissolve 150 gm. of the salt
in I OCX) cc. of water and potir this
solution into 1 000 cc. of nitric acid
.(i vol. acid to I vol. water).
(3) Dissolve 60 gm. of the dry
salt in 400 cc. of commercial am-
monium hydroxide, add 400 cc. of
Appendix C.
387
water, and then add slowly 500 cc.
of commercial nitric acid.
Ammonium Oxalate, (NH4)2C204. —
Dissolve 40 gm. of the crystallized
commercial salt in 1000 cc. of water.
Filter. Should be made as needed.
Ammonium Sulphide, (NH4)oS. —
(i) Add 500 cc. of the commercial
solution to 500 cc. of water.
(2) Pass purified hydrogen sul-
phide gas into 600 cc. of dilute
ammonium hydroxide until the so-
lution is saturated, and then add
400 cc. of dilute ammonium hy-
droxide. The saturation point may
be found either by the intense odor
of the solution, or by the fact that
the saturated solution gives no pre-
cipitate when heated gently with
magnesium sulphate. Ammonium
sulphide when exposed to the air
soon forms the yellow aiiujiouiuin
sulphide. The latter may be pre-
pared by dissolving a little flowers
of sulphur in ammonium sulphide.
Ammonium Sulphide (Yellow),
(NH4)2Sx. — Saturate 1500 cc. of
ammonium hydroxide (sp. gr. .90)
with hydrogen sulphide gas. Add
1 100 cc. of ammonium hydroxide
(sp. gr. .90) and 2500 cc. of water.
Dissolve in the above from 50 to
75 gm. of flowers of sulphur. Filter."
Ammonium Sulphocyanide (or Thi-
ocyanate), NH4SCN. — Dissolve
10 gm. of the commercial salt in
1000 cc. of water.
Barium Chloride, BaClo. — Solid.
Use the crystallized commercial
salt.
Solution. Dissolve 50 gm. of the
commercial salt in 1000 cc. of water
free from sulphuric acid and sul-
phates. Filter if not clear.
Barium Hydroxide, Ba (0H)2. —
Dissolve 50 gm. of the commercial
compound in looo cc. of hot water
free from carbon dioxide and car-
bonates. Let the solution stand a
day, and then filter if not clear.
Keep tightly stoppered.
Bromine Water. — Shake 500 cc. of
water with a few drops of bromine.
Keep tightly stoppered and stand
in the hood, where the bottle can-
not be accidentally upset. A satu-
rated solution contains about 40 gm.
to the liter.
Calcium Chloride, CaClo. — Solid.
For ordinary work use the commer- .
cial salt; for quantitative work use
lumps of the fused variety. Keep
tightly stoppered, or preserve in a
quick- sealing jar.
Solution. Dissolve 100 gm. of
the commercial salt in 1000 cc. of
water. Let the solution stand a
day, and then filter if not clear.
Calcium Sulphate, CaS04. — Solid.
For water of crystallization use
crystallized gypsum (selenite). For
other work use the powdered va-
riety.
Solution. Make a saturated solu-
tion by allowing lOOO cc. of cold
water to stand over 10 gm. of the
powdered salt. Shake often, and
then siphon off the clear supernatant
liquid.
Chlorine Water. — (i) Pass chlorine
388
Experimental Chemistry.
into water until the solution smells
strongly of the gas.
(2) Slowly add concentrated hy-
drochloric acid to a few grams of
crysfallized potassium chlorate, and
then add 500 cc. of water. Keep
tightly stoppered and in a dark
place. This solution does not keep
well, and should be made as needed.
Chrome Alum. — See Alum.
Cobalt Nitrate, Co(N03)-2. — Dis-
solve 10 gm. of the commercial salt in
1000 cc. of water. Filter if not clear.
Cobalt Chloride, C0CI2. — Dissolve
50 gm. of the commercial salt in
1000 cc. of water. Philter if not
clear.
Cochineal. — (i) Grind a few pieces
of cochineal, or a little powder, in
water, and dilute with water to the
desired tint.
(2) Digest 5 gm. of cochineal
powder with 50 cc. of alcohol (25
per cent). (See Table X., Appen-
dix C.)
Copper Sulphate, CUSO4. — Dissolve
100 gm. of the crystallized commer-
cial salt in 1000 cc. of hot water.
Disodium Phosphate, HNa2P04.—
Dissolve 100 gm. of the crystallized
commercial salt in 1000 cc. of water.
Ferric Chloride, FeCls-— (0 Dis-
solve 50 gm. of the commercial salt
in 1000 cc. of water.
(2) Boil ferrous sulphate solu-
tion with nitric acid, add am-
monium hydroxide, filter, wash the
precipitated ferric hydroxide free
from acid and alkali, and dissolve
in concenr.rated hydrochloric acid.
Ferrous Sulphate, FeS04. — Dissolve
100 gm. of the clean, fresh, or
freshly washed, commercial salt in
locK) cc. of cold water. The solu-
tion should be made as needed. If
necessary to preserve it, keep iron
wire in the bottle, and add occa-
sionally (once a day) a drop or
two of concentrated sulphuric acid.
Fusion Mixture. — Mix four parts of
anhydrous sodium carbonate and
five parts of dry potassium car-
bonate.
Hydrochloric Acid, HCl. — Concen-
trated. Use the commercial acid
(sp.gr. 1.2).
Dilute. Add one volume of
commercial acid to four volumes of
water.
Hydrogen Peroxide, H2O2. — Use the
commercial solution. Keep in a
cool place, and occasionally loosen
the cork for an instant.
Indigo. — ^i^//^. Use the commer-
cial substance.
Solution, (i) Dissolve a little
commercial indigo paste in looo cc.
of water.
(2) Slowly add 10 gm. of pow-
dered commercial indigo to 25 cc.
of concentrated sulphuric acid. Let
the mixture stand a day, then add
it slowly, with constant stirring, to
1000 cc. of water.
Iodine, I. — Solid. Use the commer-
cial crystals.
Solution. — (i) Grind to com-
plete solution in a mortar 10 cc. of
water, 12 gm. of commercial iodme,
and 20 gm. of potassium iodide,
Appendix C.
389
and then add the solution to 1000 cc.
of water.
(2) Use the commercial solution.
Lead Acetate, Pb(C2H302) 2. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water. Let the
solution stand a day, then filter if
not clear.
Lead Nitrate, PbCNOs) 2. — Dissolve
100 gm. of the commercial salt in
1000 cc. of water. Let the solution
stand until clear, then decant or
filter.
Lime Water, Ca(0H)2.— Slake quick-
lime in an iron or tin dish, add
water enough to make " milk of
lime," pour into a stoppered bottle,
shake occasionally for several days,
then let stand, and siphon off the
clear supernatant liquid. Several
portions of water may be added to
the same solid.
Litmus. — Solid. Use the commer-
cial blocks.
Solution. (i) Grind a few
blocks of litmus, with a little water,
in a mortar, and add the paste to
enough water to produce the de-
sired color.
(2) Grind logm. of litinus blocks,
with hot distilled water, in a mortar,
and, after the mass has stood for fif-
teen minutes, pour off the Hquid.
Extract once or twice more with
hot distilled water, dilute the liquid
with about 150 cc. of distilled water,
and keep in a loosely stoppered
bottle. A few drops of chloroform
occasionally added will prevent the
formation of mould.
Magnesium Sulphate, MgS04.—
Dissolve 100 gm. of the commercial
salt in 1000 cc. of water.
Magnesium Chloride, MgCl2. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water.
Manganous Chloride, MnCl.. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water.
Mercuric Chloride, HgCl2. — Dissolve
50 gm. of the commercial salt in
1000 cc. of warm water. Poison.
Mercurous Nitrate, Hg2(N03)2.—
(i) Dissolve 50 gm. of the crystal-
lized salt in 1000 cc. of distilled
water, add 75 cc. of concentrated
nitric acid and a little mercury.
(2) Dissolve metallic mercury
moderately warm concentrated ni-
tric acid, taking care to leave a
little mercury undissolved.
Nessler's Reagent. — Dissolve 3.5
gm. of potassium iodide in 10 cc.
of water; dissolve 1.6 gm. of mer-
curic chloride in 30 cc. of water.
Add the latter to the former gradu-
ally and with constant stirring until
the precipitate ceases to be redis,
solved. Then add a solution of 6.^
gm. of potassium hydroxide in 60
cc. of water. Filter. Keep tightly
stoppered.
Nitric Acid, HXO3. — Concentrated.
Use the commercial acid (sp. gr.
1.42).
Dilute. Add one volume of
commercial acid to four volumes
of water.
Phenolphthalein. — (i) Dissolve i
gm. of the solid in 100 cc. of alco-
390
Experimental Chemistry,
hoi (50 per cent.) (See Table X.,
Appendix C.)
(2) Dissolve I gm. of the solid in
100 cc. of alcohol (95 per cent).
Dilute with distilled water until a
precipitate forms; then add just
enough alcohol to redissolve the
precipitate.
(3) Use the commercial solution.
Platinic Chloride, PtCU. — (0 Buy
a solution of desired or known
strength.
(2) Dissolve scrap platinum in
aqua regia, evaporate to pasty dry-
ness, and dissolve the residue in
distilled water.
Potassium Bromide, KBr. — Dissolve
50 gm. of the commercial crystals
in 1000 cc. of water.
Potassium Carbonate, K2CO3. —
Solid. Use the dry commercial
salt, which should always be kept
tightly stoppered, or preserved in a
quick-sealing jar.
Solution. Dissolve 100 gm. of
the dry commercial salt in looo cc.
of water.
Potassium Chloride, KG. — Dissolve
50 gm. of the commercial salt in lOOO
cc. of water. Fiker if not clear.
Potassium Chromate, K..C'r04.—
Dissolve 100 gm. of the commercial
salt in icxx) cc. of water. Filter if
not clear.
Potassium Cyanide, KCN. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water. This so-
lution is a violent poison, and should
be prepared in small quantities,
plainly labelled, and kept where it
cannot be promiscuously handled.
Potassium Dichromate (or Bichro-
mate), KoCtoOt. — Dissolve 50 gm.
of the powdered commercial salt in
1000 cc. of water.
Potassium Ferricyanide,K3Fe(CN) 6.
— Dissolve 75 gm. of the com-
mercial salt in 1000 cc. of water.
Keep tightly stoppered and in a
dark place.
Potassium Ferrocyanide,K4Fe(CN)6.
— Dissolve 75 gm. of the commer-
cial salt in 1000 cc. of water.
Potassium Hydroxide, KOH.—
Solid. Use the stick form.
Solution. Dissolve 100 gm. of
the commercial lump potash in 1000
cc. of water. Use an iron dish, and
stir constantly. If the solution is not
clear, filter through glass wool or
shredded asbestos. Keep tightly
stoppered. For absorption of carbon
dioxide in quantitative work, use a
solution containing 500 gm. to the
liter (sp. gr. = 1.27 approximately).
Potassium Iodide, KI. — Solid. Use
the commercial crystals.
Solution. Dissolve 50 gm. of the
commercial salt in 1000 cc. of water.
Potassium Nitrate, KNO3. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water. Filter.
Potassium Permanganate, KMn04. —
Dissolve 5 gm. t)f the commercial
crystals in 1000 cc. of water.
Potassium Sulphate, K2SO4. — Dis-
solve 100 gm. of the commercial
salt in 1000 cc. of water.
Potassium Sulphocyanide (or Thio-
cyanate), KSCN. — Dissolve 10
gm. of the commercial salt in lOOO
cc. of water.
Appendix C.
391
Silver Nitrate, AgNOs- — Dissolve
50 gm. of the commercial crystals in
1000 cc. of distilled water. Let the
solution stand a day, and if not clear,
filter through glass wool, or better,
siphon off the clear supernatant
Uquid. This solution should be kept
free from dust and organic matter.
Soap Bubble Solution. — Dissolve 10
gm. of thin shavings of castile soap
in 100 cc. of distilled water. Shake
until solution is complete, and filter
if .lot clear. Add 50 cc. of glycerine.
Sodium Amalgam. — See Exp. 63.
Sodium Carbonate, Na2C03. — Solid.
Use the commercial powder or crys-
tals as directions require.
Solution. Dissolve 100 gm. of the
crystallized commercial compound
in 1000 cc. of water. A saturated
solution is made by dissolving about
430 gm. in 1000 cc. of water.
Sodium Chloride, NaCl. — Dissolve
100 gm. of the commercial salt in
1000 cc. of water. Filter.
Sodium Hydroxide, NaOH. — Solid.
Use the stick form.
Solution. Dissolve 100 gm. of the
commercial lump caustic soda in
1000 cc. of water. Use an iron
dish, and add the solid slowly to the
liquid with constant stirring. Filter
through glass wool or shredde^l
asbestos. Keep tightly stoppered.
Sodium Sulphite, NaoSOs. — Dis-
solve 200 gm. of the crystallized
commercial salt in 1000 cc. of water.
Stannous Chloride, SnClo. — ( i ) Use
the commeicial solution.
(2) Dissolve granulated tin in an
excess of hot concentrated hydro-
chloric acid until the evolution of
hydrogen ceases. Add 5 volumes
of water to the solution.
(3) Dissolve 500 gm. of the com-
mercial salt in 1000 cc. of hot
concentrated hydrochloric acid.
Keep a few pieces of tin in the
solution.
(4) Dissolve 112 gm. of the com-
mercial salt in 200 cc. of hydrochloric
acid (i to i) and add 800 cc. of
water. Keep a few pieces of tin in
the solution.
Stannous chloride solution does
not keep unchanged, and should be
made as needed.
Starch Paste. — Grind 10 gm. of
starch to a paste with a little cold
water in a mortar. Add the paste
to 200 cc. of water. Boil. Use
when cold.
Sulphuric Acid, USO^.- — Concen
trated. .Use the commercial acid
(sp.gr. 1.84).
Dilute. Slowly pour one volume
of commercial acid into four volumes
of water. Stir constantly during the
operation. Cool and filter, if the
commercial acid is not colorless.
Tartar Emetic, K(SbO)C4H406.—
Dissolve 100 gm. of the commercial
salt in 1000 cc. of water.
Zinc Chloride, ZnCL.. — Dissolve 50
gm. of the coimiiercial salt in looo
cc. of water.
Zinc Sulphate, ZnS04. — Dissolve
100 gm. of the commercial salt in
1000 cc. of water. Filter if the
solution is not clear.
39'^
Experimental Chemistry.
V. TABLES.
TABLE I. — IMPORTANT ELEMENTS AND THEIR ATOMK
WEIGHTS.
Symbol.
Atomic Weight.
Name.
Clarke.
Richards.
German.
Approxi-
mate.
H = i.ooo
0 = 16
0 = 16
0 = 16
0 = 16
Aluminium . .
Al
26.9
27.1
27.1
27.1
27
Antimony
Sb
1 19-5
120.4
120.0
120
120
Arsenic .
As
7445
75
75-0
75
75
Barium .
Ba
136.4
137-4
13743
1374
137
Bismuth .
Bi
206.5
208.1
208
208.5
208
Boron . .
B
10.9
"
II.O
II
11
Bromine .
Br
79-34
79-95
79-955
79.96
80
Cadmium .
Cd
"1-55
II 2.4
1 1 2.3
112
112
Calcium .
Ca
39.8
40.1
40.1
40
40
Carbon
C
11.9
12
12.001
12
12
Chlorine .
CI
35-18
35-45
35-455
3545
35-5
Chromium
Cr
517
52.1
52.14
52.1
52
Cobalt . .
Co
58.55
59
59.00
59
59
Copper
Cu
63.1
63.6
63.60
63.6
63.5
Fluorine .
F
18.9
19.05
19.05
19
19
Gold . .
Au
195-7
197.2
197-3
197.2
197
Hydrogen
H
I
1.008
1.0075
I.OI
I
Iodine . .
I
125.89
126.85
126.85
126.85
127
Iron . .
Fe
55.6
56
55-9
56
56
Lead . .
Pb
205.36
206.92
206.92
206.9
207
Lithium .
Li
6.97
7-03
7-03
I'O?,
7
Magnesium
Mg
24.1
'24-3
24.36
24.36
24
Manganese
Mn
54-6
55
55-02
55
55
Mercury .
Hg
198.50
200
200.0
200.3
200
Nickel. .
Ni
58.25
58.70
58.70
58.7
58.5
Appendix C.
393
TABLE L — IMPORTANT ELEMENTS AND THEIR ATOMIC
WEIGHTS — {Continued) .
Symbol.
Atomic Weig
^T.
Name.
Clarke.
Richards.
German.
Approxi-
mate.
H = i.ooo
0 = 16
0=16
0 = 16
0= 16
Nitrogen . , .
N
13-93
14.04
14.045
14.04
14
Oxygen . . ,
0
15.88
16
16.000
16
16
Phosphorus . .
P
3075
31
31.0
31
31
Platinum . .
Pt
193-4
194.9
195.2
194.8
195
Potassium . .
K
38.82
39"
39-H
39-15
39
Selenium . . .
Se
78.6
79.2
79.2
79.1
79
Silicon . . .
Si
28.2
28.4
28.4
28.4
28
Silver ....
Ag
107. II
107.92
107-93
107.93
108
Sodium . . .
Na
22.88
23-05
23-05
23-05
23
Strontium . .
Sr
86.95
87.60
87.68
87.6
87.5
Sulphur . . .
S
31-83
32.07
32.065
32.06
32
Tellurium . .
Te
126.5
127.5?
127.5?
127
127
Tin
Sn
118.1
119
II 9.0
118.5
119
Zinc ....
Zn
64.9
65.4
65.40
65-4
65
The above list is taken from \h^ Journal of the American Chemical Society,
Vol. XXII., No. 2 (February, 1900). Richards's list was revised by a private
communication in the final proof. The approximate list is chosen to facili-
tate calculation, ami should be used in solving all problems in this book.
TABLE IL — COLORS OF COATINGS ON CHARCOAL.
Metal.
Hot.
Cold.
Lead
Tin
Zinc
Lemon-yellow
Yellow (faint)
Yellow
Lemon-yellow
White
White
394
Experimental Chemistry.
TABLE III. — COLORS OF BORAX BEADS.
Oxidizing Flame.
Reducing Flame.
Metal.
Hot.
Cold.
Hot.
Cold.
Chromium
Cobalt
Copper
Manganese
Reddish-
Yellow
Blue
Green
Violet
Yellowish-
Green
Blue
Greenish-
Blue
Violet
Green
Blue
Colorless
Colorless
Green
Blue
Red
Colorless
TABLE IV. — COLORS OF RESIDUES MOISTENED WITH
COBALT NITRATE.
Metal.
Color (best seen when cold).
Aluminium
Magnesium
Zinc
Blue
Pink
Green
TABLE v. — COLORS OF FLAMES.
Metal.
Color of Bunskn Flame.
Barium
Green
Calcium
Yellowish-red
Potassium
Violet
Sodium
Yellow
Strontium
Crimson
TABLE VI. — CONVERSION OF THERMOMETRIC READINGS.
Fahrenheit to Centigrade.
Centigrade to Fahrenheit.
^' - 1.8
or
C _ 5 (F. - 32)
9
F.= (C. X 1.8) -1-32
or
Appendix C.
395
TABLE VII. — SPECIFIC GRAVITY AND MELTING POINT
(APPROXIMATE) OF METALS.
Metal,
Aluminium . .
Copper . . .
Iron (wrought)
Lead ....
Mercury
Platinum . .
Silver . . .
Tin ... .
Zinc ....
Specific
Gravity.
2.6
8.9
7.8
"•35
13-59
21.5
IO-53
7-3
6.86-
7.2
Melting Point.
Centigrade.
625
1054
1600
326
-38.5
1775
954
232.7
433
Fahrenheit.
"57
1929
2912
619
- 37-3
3227
1749
450-9
811. 4
TABLE VIII. — SOLUBILITY OF SALTS IN WATER.
Salt.
Copper sulphate (cryst.)
Potassium chlorate . .
Potassium chloride . .
Potassium dichromate .
Potassium nitrate . .
Potassium sulphate . .
Sodium chloride . . .
Sodium sulphate (anhyd.)
Sodium sulphate (cryst.)
Grams soluble in 100 gm. at
20° c
100" C.
42.31
203.32
7.2
59-5
35
57
13
102
3^-7
246
10.6
26
36
39-7
20
43
58-35
212.47
TABLE IX. —WEIGHT (IN GRAMS) OF A LITER OF DRV
GASES AT 0° C. AND 760 MM.
Air
Ammonia
Carbon dioxide .
Carbon monoxide
Chlorine .
Hydrochloric acid
Hydrogen . . .
Hydrogen sulphide
Oxygen ....
Nitrogen . . .
Steam ....
Sulphur dioxide .
.0896
1-542
1.429
1.256
.806
2.869
?^9^
Experimental Chemistry.
TABLE X. — FORMULA FOR PREPARING AN ALCOHOLIC
SOLUTION OF DESIRED STRENGTH.
_ ac
X = vol, required.
a = % of solution used.
d := %of solution desired.
c = vol. of stock solution used.
TABLE XL —COMPOSITION OF TYPICAL COALS.
Kind.
C.
H.
0.
N.
Wood
50
59
69
82
95
6
7
5-2
5
2.5
43
34
25
12.2
2.5
I
Peat
0
Lignite
Bituminous .
Anthracite
.8
.8
trace
TABLE XIL — PERIODIC ARRANGEMENT OF THE ELEMENTS,
ACCORDING TO MENDELEEFF.
Revised by F. W. Clarke. On the Basis of O = 16.
Groip I.
Group II.
Group III.
Group IV.
Series.
R2O.
RO.
R.P3.
RH4, RO^.
.T
H = 1.008
2
Li -. 7.03
Gl = 9.1
B= II
G= 12
3
Na = 23.05
Mg = 24.3
Al = 27.1
Si = 28.4
4
K = 39.11
Ca = 40
Sc = 44.i
Ti = 48.1
5
(Cu = 63.6)
Zn = 65.4
Ga = 69.9
Gi = 72.5
6
Rb = 85.4
Sr = 87.6
Yt = 89
Zr = 90.4
• 7
(Ag = 107.92)
Cd = 1 12.4
In= 1 13.8
Sn = 119
8
Cs= 132.9
Ba= 137.4
La = 138.6
Ce = 139.4
9
( )
10
Er = 166.3
II
(Au = 197.2)
Hg = 200
Tl = 204.1
Pb = 206.9
12
Th = 232.6
Appendix C.
TABLE y.\\.— {Conliimed).
397
Group V.
Gkoup VI.
Group VII.
Group VIII.
Series.
RH3, R2O5.
RH2, RO3.
RK, R0O7.
RO4.
I
2
N = 14.04
0=16
F=I9
3
P=3i
S = 32.07
CI = 3545
4
V = 5i.4
Cr = 52.i
Mn = 55
Fe = 56 Ni = 58.07
Co = 59 Cu = 63.6
5
As = 75
Sc = 79.2
Br = 79.95
6
Cb = 93.7
Mo = 96
= 100.
Ru=ioi.7 Rh=i03
Pd= 106.4 Ag= 107.92
7
Sb= 120.4
Te = 127.5
1= 126.85
8
9
lO
Pr = 140.5
Nd = 143.6.
— — — —
Ta = 182.8
W = 184.8
Os = 191 Ir = 193. 1
Pt= 194.9 Au = 197.2
II
Bi = 208.1
—
12
U = 239.6
— — — —
Note. — This classification does not include argon, helium, other recently
announced constituents of the atmosphere, and the rare earth metals.
XIII.— EMERGENCY SET.
A box or cabinet provided with the following articles should be kept in a
convenient place : —
Absorbent cotton, \ lb.
Linen iclls, I in. and 2 in.
Court-plaster, thin and thick, cut in
strips.
Emulsion of lime water and oil, 500 cc.
Sodium bicarbonate.
Cuts and burns should receive prompt attention. Fires may be extinguished
by wet sand, if ca'ised by burning liquids, or by a damp towel or blanket, if
due to burning clothing. Acids and alkalies should be removed from the flesh
with warm water, and the injured spot immediately covered with a paste made
by mixing sodium bicarbonate and the emulsion of lime water and oil ; the
paste may be held in place by cotton and a linen bandage.
Vaseline, i lb. box.
Mortar or shallow dish.
Smelling salts.
Ammonium hydroxide.
Sand.
Blanket.
j^S Experimental Chemistry.
VI. SUGGESTIONS FOR A SHORT COURSE.
The experiments incorporated in this book constitute a liberal course
in chemistry. In view of the fact, however, that conditions, often tem-
porary or unavoidable, may prevent the pupil from performing all the
experiments, this course may be shortened without detriment to its
integrity, if the following selection and apportionment of experiments
be made : —
A.
PUPIL'S
EXPERIMENTS
I,
24,
46,
69,
86,
133,
176,
2,
25,
47,
70,
88,
138,
177 (^)
3>
27,
48,
71,
go (a)
139.
(0,
4,
28,
49,
72,
and (d),
150,
w,
8,
29,
50,
73'
102,
151,
w,
9j
32,
51,
76,
107,
152,
188,
ID,
33,
52,
77,
112,
160,
189,
11(a)
34,
56,
81,
114,
161 (a),
190,
(orii(^)):
, 35,
58,
82,
116,
162,
191,
14,
36,
59,
83,
117,
167,
192.
i8,
37,
64,
84,
129,
168,
B. TEACHER'S EXPERIMENTS.
12,
19'
23^ 42,
61,
75,
108 (or 109
1), 128,
134,
13,
20,
30' 57,
62,
98,
124,
130,
136,
17,
21,
38 (or 39), 60,
65,
99,
125,
131,
179.
The course provided by the above experiments is short, representa-
tive, and continuous. Suggestions regarding these and other experi-
ments may be found in the Teacher's Supplement, which will be sent
to teachers on application to the publishers. The author takes advan-
tage of a second opportunity to invite correspondence concerning the
adaptation of the book to meet special demands.
INDEX
ADbreviaiious of metric denominations,
lO.
Acetates, 215.
Preparation, Exp. 115.
Acetic acid, 245.
Formula, 246.
Properties, Exp. 112.
Test for, Exp. 1 13.
Acetic aldehyde, Exp. iii (a).
Acetylene, preparation and properties,
Exp. 97.
Acid, definition, 117, 128.
Acids and bases, interaction of, Exp. 52.
Quantitative examination of, Exp. 53.
Acids, bases, and salts, 116.
Acids, general properties, 117, Exp. 48.
Nomenclature, 117.
Air bath, 333.
Air, composition, 105.
Quantitative examination, Exp. 45.
See Atmosphere.
Weight of liter of, Exp. 47.
Alcohol, 238, 240.
Formula, 240.
Preparation, Exp. 108.
Properties, Exp. 109.
Alcoholic solution of desired strength,
396.
Aldehyde, 244.
Aldehydes, preparation and properties,
Exp. III.
Alloys, 283.
Alum, common, preparation and proper-
ties, Exp. 171.
Chrome, Exp. 184.
Aluminium, action with acids and alka-
lies, Exp. 168.
Equivalent, Exp. 55.
General properties, Exp. 167.
Tests for, Exp. 170.
Aluminium hydroxide, preparation and
properties, 304, Exp. 169.
Amalgam, 288.
Sodium, Exp. 63.
Ammonia, 172.
Preparation, Exp. 64.
Ammonia gas, composition, 174.
Formula, 182.
Preparation and properties, Exp. 65.
Qualitative examination, 178, Exp. 66.
Quantitative examination, 178, Exp. 67.
Volumetric composition, 181, Exp. 68.
Weight of liter of, 182.
Ammonium, 184.
Compounds, 183, 184.
Ammonium hydroxide, 172.
And hydrochloric acid, Exp. 69.
Preparation and properties, Exp. 65.
Ammonium nitrate, action with heat,
Exp. 79.
Analysis, definition, 33.
Anhydride, definition, 129.
Anhydrous, definition, 67.
Annealing, 4.
Apparatus, lists, 381-383.
Aqua regia, 202.
Characteristic property, Exp. 81.
Aqueous tension, 360.
Table of, 361.
Argon, 105, 397.
Aspirators, 343.
Atmosphere, 105.
Other constituents of, Exp. 46.
See Air,
Atom, 53, 140, 141.
Atomic and molecular weights, calcula-
tion of, 144.
Atomic theory, 136.
399
400
Index.
Atomic weights, 139.
Various figures lor, 142.
Table of, 392.
Avogadro's law, 143.
Balance, counterpoising a, 347.
Hornpan, 12, 347.
Barium chloride, water of crystallization
in, Exp. 26.
Barometer, 353.
Using a, 353, 369.
Base, definition, 128.
Bases, general properties, 118, Exp. 49.
Nomenclature, 119.
See Acids.
Basic oxides, 129.
Bending glass, 2-6, Exps. 2, 3.
Bertholiet, 35.
Berzelius, 138.
Bibliography, 377-381.
Black, 137.
Bleaching by chlorine, Exp. 58.
By sulphur dioxide, Exp. 136 {e).
Blowpipe and its use, 232.
Blowpipe flame, 233.
Oxidation with, Exp. 106.
Reduction with, Exp. 105.
Blowpipe tube, 232.
Blowpipe, use of, Exp. 104.
Boiling point, elevation of, 150.
Of water, Exp. 31.
Bone black, 205.
Borax bead, 231, 232.
Tests with, Exp. 103.
Bottle, Drechsel, 368.
Drying (or washing), 368.
To clean the inside of a, 338.
To cut off the bottom of a, 336.
Boyle, 105.
Law of, 357.
Brass, analysis of, Exp. 196.
Bromides, 251.
Bromine, preparation and properties,
Exp. 119.
Bunsen burner, 1-2, 228.
Construction of, Exp. 100.
Flame, Exp. loi.
Burette, use of, Exp. 53.
Burner. See Bunsen.
Wing-top, 2.
Calcium, compounds of, 294.
Tests for, 221, 294, Exp. 159.
Candle flame, Exp. 102.
Carbohydrates, 248.
Carbon, 204, 236. See Charcoal.
Distribution of, Exp. 82.
In organic compounds, Exp. 107 (a).
Reduction by, Exp. 85.
Carbonate, acid calcium, Exp. 91.
Determination of carbon dioxide in a,
Exp. 92.
Carbonates, 213.
Formation and properties, Exp. 90.
Carbon bisulphide test for iodine, Exp.
123.
Carbon dioxide, 207.
Action of hot charcoal on, Exp. 93.
And combustion, Exp. 88.
Determination of, in a carbonate, Exp.
92.
Preparation and properties, Exp. 86.
Synthesis of, Exp. 87.
Weight of liter of, 210.
Carbon monoxide, 216, Exp. 93.
Preparation and properties, Exp. 94.
Weight of liter of, 218.
Carbonic acid, 212, 214, Exp. 89.
Casserole, 331.
Cavendish, 38, 51, 83, 137.
Changes, physical and chemical, 15.
Charcoal, action of carbon dioxide on
hot, Exp. 93.
Animal, 205.
Decolorizing action of, 205, Exp. 83.
Deodorizmg action of, 205, Exp. 84.
Charles, law of, 354.
Chemical action, 22.
Chemical compounds, 27, 29, 30, 3^"
Chemical reaction. See Reaction.
Chemical terms, glossary of, 370-377-
Chemicals, list of, 383-385.
Chemistry, definition, 15-16.
Organic, 206.
Chloride, definition, 34.
Test for, FIxps. 29 {d), 61 Iv.
Chlorine, 157, 250.
Bleaching by, Exp. 58.
Decomposition of water by, Exp. 41.
Preparation and properties, E.xps. 56,
57-
Index.
401
Weight of liter of, 160.
Chromates, 314.
Properties, Exp. 182.
Reduction of, to chromic compounds,
Exp. 183.
Chrome akim, properties, Exp, 184.
Chromic compounds, reduction of chro-
mates to, Exp. 183.
Chromic hydroxide, Exp. 185.
Chromium, 313.
Tests for, Exp. 181.
Cinnabar, Exp. 164 {d).
Class-room Exercises.
Acetates, 246.
Acetic acid, 246.
Acetylene, 224.
Acids, bases, and salts, 129-131.
Air, 109-111, 113.
Alcohol, 241.
Alkali, 130.
Allotropy, 206, 276, 277.
Alloys, 289, 291, 298, 308, 312.
Aluminium, 305.
Alums, 305.
Amalgams, 301.
Ammonia, 185, 186.
Ammonium compounds, 185, 186, 285.
Ampere, 153.
Animal charcoal, 207.
Apparatus, 52.
Aristotle, 109.
Arsenic, 277.
Atmosphere. See Air.
Atomic theory, 277.
Atomic weights, 162.
Atoms, 153.
Avogadro, 153.
Balard, 255.
Becher, 36.
Berzelius, 153.
Black, 211, 293.
Bleaching by chlorine, 160.
Bleaching powder, 160.
Boiling, 74.
Boyle, 359.
Bread-making, 249.
Bromine, 255.
Bunsen, 109, 230.
Bunsen fiame, 23.
Class-room Exercises,
Calcium and its compounds, 295-296.
Cannizzaio, 153.
Carbon, 206-207.
Carbon dioxide, no, 211.
Cavendish, 193.
Cellulose, 249.
Charcoal, 207.
Charles, 359.
Chemical action, 23, 26, 277.
Chemists, 278.
Chlorides, 72.
Chlorine, 160-161.
Chromium and its compoundi, 316.
Coal, 207.
Collecting gases, 23.
Combustion, 36.
Compound blowpipe, 51.
Compounds, 36, 55, 277.
Copper and its compounds, 288-289.
Courtois, 255.
Crystallization, 51.
Dalton, 153.
Davy, 87, 130, 161, 170, 255.
Deliquescence, 72.
Dextrine, 249.
Diamond, 206.
Diffusion of gases, 52.
Distillation, 73.
Dulong and Petit, 153.
Effervescence, 52.
Efflorescence, 72.
Electrolytic dissociation, 153.
Elements, 36, 55, 277.
Equations, 277, 306, 308.
Ether, 242-243.
Ethereal salts, 242.
Ethylene, 224.
Faraday, 161, 185.
Fermentation, 241.
Fertilizer, 276, 285.
Filtration, 73.
Flames, 230.
Fluorine, 258.
Formaldehyde, 244.
Formulas, 55, loi, 246, 271, 277, 278,
312, 316.
Freezing, 75.
Gas, illuminating and water, 227.
Gases in the atmosphere, iii.
402
Index.
Class-room Exercises.
Gases, solubility in water, jj.
Gay-Lussac, 255.
Law of, 186, 224.
Glass, 7, 275.
Glycerine, 248,
Graphite, 206.
Gunpowder, 202.
Heating glass and porcelain, 23, 36.
History of air, 109.
Ammonia, 185.
Bromine, 255.
Carbon dioxide, 211.
Chlorine, 161.
Hydrochloric acid, 170.
Hydrogen, 51.
New gases in air, 113.
Nitric acid, 193.
Oxygen, 36.
Water, 87, loi.
Hydrocarbons, 223, 238.
Hydrochloric acid, 169-170.
Hydrofluoric acid, 258.
Hydrogen, 50, 130.
Sulphide, 267.
Hydroxyl, 131.
Hypothesis, 369.
Illuminating gas, 227.
Iodine, 255-256.
Iron and its compounds, 321-322.
Lavoisier, 36, 51, 109, 130, 193, 211.
Law, 369.
Of definite proportions, 36.
Of Gay-Lussac, 186, 224.
Of multiple proportions, 202, 227.
Laws, 277.
Lead and its compounds, 311-312.
Lead in water, 73.
Liebig, 255.
Litmus, loi.
Magnesium and its compounds, 293.
Manganese, 162, 318.
Mariotte, 359.
Matches, 270,
Matter, 23, 72.
Mercury and its compounds, 301.
Metals, separation of, 325.
Methane, 224.
Mixture, iii.
Moissan, 258.
Class-room Exercises.
Molecular weight, 162, 277.
Molecules, 153, 277.
Multiple proportions, law of, 202, 227,
Nascent state, 161.
Neutialization, 129, 130.
Nitrates, 200, 202,
Nitric acid, 193, 194.
Nitrogen, 108.
Oxides of, 200.
Nomenclature, 130, 278.
Organic acids, 248.
Organic matter, 72.
Oxalic acid, 246.
Oxidation, loi, 235, 277.
Oxygen, 36, 109.
Periodic classification, 258, 271, 276,
277, 285, 291, 296, 301, 312.
Phosphorus, 109, 276.
Physical and chemical changes, 23.
Potassium and its compounds, 284.
Potassium permanganate, 51.
Precipitation, 51.
Priestley, 36, 170.
Reduction, loi, 235, 277.
Relation between chlorine, bromine,
and iodine, 255, 256.
Review, 277.
Rocks, 275.
Salts, ethereal, 242.
Scheele, 36.
Separation of metals, 325.
Silicon and its compounds, 274-275.
Silver and its compounds, 291.
Soap, 248.
Sodium and its compounds, 281-282.
Sodium and potassium, 285.
Solids, solubility of, in water, 82.
Solution, 277.
And taste, 82.
Modern theory of, 154.
Specific heat, 153.
Stahl, 36.
Standard conditions, 369.
Starch, 249.
Stas, 153.
Stassfurt deposits, 256, 285, 293.
Sugars, 249.
Sulphates, 72.
Sulphide, hydrogen, 267,
Index.
403
Class-room Exercises.
Sulphides, 267.
Sulphur, 262.
Sulphur dioxide, 269.
Sulpliuric acid, 271.
Symbols, 55, loi, 162, 277.
Tests, 278.
Theory, 369.
Tin and its compounds, 308.
Valence, 154, 186, 194, 202, 271, 277,
282, 285, 289, 293, 296, 298, 30X,
306, 308, 309, 312, 316, 318, 322.
Van Helmont, 211.
Vapor tension, 361.
Vinegar, 246.
Volumetric composition, 277.
Water, 72.
As a standard, 75.
Composition of, 87, loi.
Drinking, 72.
Formula of, loi.
Gas, 227.
Hard, 72.
History of, loi.
In atmosphere, no.
Lead in, 73.
Maximum density of, 75.
Of crystallization, 72, 73.
Relation to temperature, 75.
Solubility of gases in, 77.
Solvent power of, 82.
Weighing, 36.
Zinc and its compounds, 29S.
Cleanliness, 9.
Closing tubes, Exps. 5, 6.
Coals, composition of, 396.
Combination, definition, 33.
Components, definition, 33.
Composition of organic compounds, 236,
Exp. 107.
Composition of water, summary of, 99.
Composition, percentage, 148.
Concentrated, definition, 75.
Condenser, 71.
Connector, 338.
Section of inner, 21.
Cork, boring a, 335.
Corks, 334.
(-opper, 286. See Cuprous-
Compounds of, 287.
General properties, Exp. 150.
Interaction with metals, Exp. 152,
Oxides, 287.
Tests for, Exp. 151.
Corrected, definition, 362.
Correction, combined, for temperature
and pressure, 363.
For aqueous tension, 361, 363.
For pressure, formula for. 358, 363.
For temperature, formula for, 354, 362.
Counterpoising a balance, 347.
Crucible block, 19.
Crystallization, 42.
Of zinc sulphate, Exp. 20.
Rules for, 43.
Cuprous oxide, preparation and proper-
ties, Exp. 153.
Cutting glass, 2, Exp. i.
Dalton, 137, 138, 202.
Davy, 83, 157.
Decomposition, definition, 33.
Deflagrating spoon, 25, 26.
Deflagration, 211.
Dehydrated, definition, 67.
Deliquescence, 68-69, Exp. 28.
Determination, definition, 34.
Dilute, definition, 75.
Displacement, downward and upward,
342.
Dissociation, definition, 67.
Distillate, definition, 71.
Distillation, Exp. 30.
Dulong and Petit, 141.
Dumas, 83.
Efflorescence, 68, Exp. 27.
Electrolytes, 151.
Electrolytic dissociation, theory of, 151,
152.
Element, definition, 15.
Elements, table of, 392.
Emergency set, 397.
Equations.
Combination of calcium oxide and
water, 131.
Hydrochloric acid gas and am-
monia gas, 182.
404
Index.
Equations.
Hydrogen and chlorine, 171.
Hydrogen and nitrogen, 184.
Oxygen and magnesium, 56, 58, 61.
Oxygen and phosphorus, 115.
Sulphur trioxide and water, 131.
Sulphuric acid and ammonia, 182.
Combustion of ethylene, 223.
Methane, 222.
Decomposition of ammonia gas by
chlorine, 182.
Mercuric oxide, 59.
Potassium chlorate, 59.
Silicic acid, 274.
Water by carbon, 227.
Water by chlorine, 173.
Fermentation of glucose, 240.
Formation of steam from hydrogen
and oxygen, 102.
Water from hydrogen and oxygen,
100.
Interaction of alcohol and acids, 241.
Ammonia gas and magnesium, 181.
Copper and nitric acid, 207.
Hydrofluoric acid and silicon diox-
ide, 258.
Lime water and carbon dioxide, no.
Magnesium nitride and water, 181.
Sodium and water, 91.
Sodium chloride and sulphuric acid.
Stannous and mercuric chlorides,
308.
Sulphuric acid and potassium per-
manganate, 318.
Neutralization, 129.
Oxidation of carbon to carbon dioxide,
60, 220.
Ferrous oxide, 322.
Iron, 322.
Sulphur to sulphur dioxide, 60.
Preparation of aldehyde, 315.
Aluminium, 306.
Aluminium hydroxide, 304.
Ammonia gas, 177.
Ammonium chloride, 167.
Bromine, 251.
Carbon dioxide, 219.
Carbon monoxide, 228.
Chlorine, 163, 319.
Equations.
Chromic chloride, 314.
Cuprous oxide, 289.
Formaldehyde, 244.
Hydrofluoric acid, 258,
Hydrogen, 56.
Hydrogen sulphide, 264.
Iodine, 253.
Lead chromate, 316.
Nitric acid, 201.
Nitric oxide, 207.
Nitrogen peroxide, 207.
Potassium aluminate, 304, 306.
Potassium chromate, 314.
Potassium dichroinate, 314.
Potassium hydroxide, 285.
Sodium, 285.
Sodium carbonate, 286.
Sodium sulphate, 286.
Sodium sulphide, 286.
Sulphuric acid, 269,
Sulphurous acid, 269.
Reduction of ferric oxide, 322.
Equations, chemical, 55-57. 59-6o.
Problems based on, 60-61.
Quantitative interpretation of, 57-59-
Equivalent, definition, 48, 132.
Of aluminium, Exp. 55.
Of magnesium, Exp. 54.
Of zinc, Exp. 22.
Ester, 241.
Etching, Exp. 128.
Ether. 242.
And water, Exp. 33 {b).
Properties, Exp. no.
Ethyl, 240.
Acetate, 241, Exp, 109 {e).
Alcohol. See Alcohol.
Chloride, 241.
Ethylene, combustion of, 223.
Preparation and properties, Exp. 96.
Eudiometer, 351, Exps. 43, 68.
Evaporation, 332.
Experiments, form of record of, 13.
Fats, 247.
Fermentation, 240, 246, Exp. 109.
Ferric compounds, behavior of, Exp. 190
, Reduction of, Exp. 191.
index.
405
Ferrous and ferric compounds, 319.
Ferrous compounds, behavior of, Exp
189.
Oxidation of, Exp. 192.
Filter, folding a, 8.
Pump, 334, 343.
Filtrate, definition, 8.
. Filtration, definition, 8, 333.
Flame, Bunsen, 2.
Bunsen burner, Exp. loi.
Candle, Exp. 102.
Flat, 2.
Luminous and non-luminous, i, 2.
Oxidizing and reducing, 230-231.
Fluorine, 256.
Formaldehyde, Exp. iii (d).
Formulas, 54.
Calculation of, 149.
For preparing an alcoholic solution,
396.
Freezing point, depression of, 150.
Of water, Exp. 31.
Gas from the red powder, identification
of, Exp. 13.
Gas holder, 342.
Gas, how to light, 1-2.
Gases, collecting, 341,
Drying agents for, 367.
Manipulation of, 366-369.
Measuring, 349.
Solubility of, 76, Exp. 32.
Gay-Lussac, 83.
Law of, 182.
Generator for producing steady current
of gas, 96.
Glass stopper, holding a, 339.
Glass tube, to fit a, to a stopper, 337,
To insert a, into a rubber tube, 337.
Glass working, 2-5.
Glossary of chemical terms, 370-377.
Glycerine, 247.
Gram, 345.
Hard glass tubing, 329.
Hardness of water, 248, Exp. 117 (d).
Heating glass and porcelain, 330.
Heating glassware, 7.
Heating hard glass, 329.
Hofmann apparatus, Exp, 38.
Hofmann screw, 343.
Humboldt, 83.
Hydrate, definition, 119.
Hydrated, definition, 67.
Hydriodic acid, preparation and proper-
ties, Exp. 126.
Hydrobromic acid, preparation and
properties, Exp. 120.
Hydrocarbons, 221, 236, 238.
Hydrochloric acid, 162.
Preparation and properties, Exp. 61.
Hydrochloric acid gas, composition of,
168, Exp. 63.
Formula of, 169.
Properties, Exp. 59.
Synthesis of, Exp. 60.
Hydrofluoric acid, preparation and prop-
erties, Exp. 128.
Hydrogen, 38.
And water, 86, 90, Exp. 40,
Flame, Exp. 23.
Generator, 39.
In organic compounds, Exp. 107 (d).
Preparation and properties, Exps. 18,
19.
Product of burning, Exp. 23.
Weight of liter of, 50.
Hydrogen sulphide, 262.
Preparation and properties, Exp. 133.
Weight of liter of, 264.
Hydroxides, 119.
Hydroxy!, 129.
Ignition tubes, making, 329.
Illuminating gas, 225.
Combustion of, Exp. 99.
Preparation and properties, Exp. 98.
Iodide, potassium, properties of, Exp.
127.
Iodides, 255.
Iodine, 250.
Carbon bisulphide test for, Exp. 123.
Preparation and properties, Exp. 122.
Starch test for, Exp. 124.
Ions and ionization, 151.
Iron, 319. See Ferrous and Ferric.
General properties, Exp. 188.
Joints, 336.
4o6
Index.
Laboratory Exercises.
Acids, bases, and salts, 123.
Air, 27, 108.
Calcium, compounds cf, 294.
Chemical action, 108.
Chromium, compounds of, 316.
Copper, 286, 287.
Ether, 242.
Flames, 230, 231.
Hydrocarbons, 238.
Hydrogen, 41.
Sulphide, 264.
Iron, 319, 321.
Lead, compounds of, 311,
Magnesium and its compounds, 292.
Manganese, compounds of, 317.
Mercury, compounds of, 300.
Nitric acid, 190.
Nitrogen, 108.
Oxides, 23, 26.
Oxygen, 26.
Phosphorus, 108, 276.
Potassium, compounds of, 283, 318.
Reduction, 206.
Silver, 291.
Starch, 249.
Sugar, 249.
Sulphide, hydrogen, 264.
Sulphur dioxide, 269.
Sulphuric acid, 271.
Zinc and its compounds, 297.
Lavoisier, 51, 55, 83, 86.
Law of Avogadro, 143.
Boyle, 357.
Charles, 354. 357.
Conservation of matter, 55.
Definite proportions by weight, 34.
Gay-Lussac, 182.
Multiple proportions, 201.
Specific heats, 141.
Lead, 309.
Action of water on, 310, Exp. 179.
General properties, Exp. 176.
Interaction with metals, Exp. 178.
Tests for, Exps. 98, 177.
Lead acetate, preparation, Exp. 115 {b).
Lead oxides, properties of, 310, Exp. 180.
liquids, measuring, 349.
Pouring, 339.
Solubility of, Exp. 33.
Liter, 9, 10, 345.
Lucretius, 137.
Magnesium, 292.
Atomic weight of, 144.
Combination of, with oxygen, Exp. 15.
Equivalent of, Exp. 54.
General properties, Exps. 9, 14 {e),
157-
Tests for, Exp. 158.
Magnesium nitride, 179.
Magnesium oxide, molecular weight of,
145-
Manganese, 317.
Tests for, Exps. 103 (t), 186.
Marsh gas, Exp. 95.
Matter, 15.
Conservation of, 55.
Organic, 69, 310, Exps. 29 (a), 138.
Meniscus, 351.
Mercurous and mercuric compounds
300, Exp. 166.
Mercury, 298.
Compounds of, 300.
General properties, Exp. 163.
Preparation, 299, Exp. 164.
Tests for, Exp. 165.
Metal, definition, 15, 16, 128, 279.
Change in, when heated, 16, Exp. 9.
Effect of heating a, covered and un-
covered. Exp. 10.
Result of heating known weight of,
Exp. II.
Metals, separation of, 322.
Methane, combustion of, 222.
Preparation and properties, Exp. 95.
Meter, 344.
Metric system, 9-10, 344.
Mixture, definition, 35.
Molecular weights, 142.
Multiple proportions, law of, 201.
Neutralization, 125, 127, Exps. 52, 53.
Newton, 137.
Nicholson and Carlisle, 83.
Nitrates, 198.
Action of, with heat, Exps. 78, 79.
Special test for, Exp. 80.
Test for, Exp. 73,
Index.
407
Nitric acid, analysis, Exp. 75.
Interaction with metals, 195, 197, Exps.
76, n-
Preparation, 198, 201, Exps. 70, 73.
Special property, Exp. 71.
Synthesis, Exp. 74.
Test for, Exp. 72.
Nitrogen in the air, Exp. 45.
In organic compounds, Exp. 107 (c).
Weight of liter of, 113.
Nitrogen, oxides of, 197, Exps. 'j'j, 79.
Non-metal, 128.
Observed, definition, 362.
Organic chemistry, 206.
Organic compounds, composition of, 236,
Exp. 107.
Organic matter, 69, 310, Exps. 29 {a) , 138.
Ostwald, 153.
Oxidation, 34, 308, 314.
With blowpipe, Exp. 106.
With potassium permanganate, Exp.
187.
Oxide, definition, 22.
Oxygen, 24.
And water, 88.
Combination of, with magnesium, Exp.
15-
Density of, 102.
In the air, 105, Exp. 45.
Per cent in potassium chlorate, Exp.
16.
Preparation and properties, Exp. 14.
Weight of liter of, 30, Exp. 17.
Palmitic and stearic acids, 247.
Percentage composition, 148.
Phosphorus, 275.
Precaution in using, 106.
Pipette, use of, -jt, 352, Exp. 33.
Platinum test wire, 341.
Platinum tip, 340, Exps. 23, 60.
Platinum wire, to seal a, into a glass rod,
341-
Pneumatic trough, 21, 24, 344.
Potassium, properties, Exp. 146.
Properties of compounds, Exp. 147.
Potassium bromide, properties, Exp. 121.
Potassium carbonate, preparation and
properties, 283, Exp. 149.
Potassium chlorate, molecular weight of,
146.
Per cent of oxygen in, Exp. 16.
Potassium chloride, molecular weight of,
146.
Potassium hydroxide, preparation, 283^
Exp. 148.
Potassium iodide, properties, Exp. 127.
Potassium permanganate, oxidation with,
Exps. 29 (a), 187.
Powder, to introduce a, into a tube, 338.
Precautions in bending glass, 5.
In cutting glass, 3.
In generating hydrogen, 39, 48.
In preparation of methane, Exp. 95.
In preparation of ethylene, Exp. 96.
In reading the barometer, 353.
In reading volumes, 351.
In using carbon bisulphide, 253.
In using ether, Exp. no.
In using phosphorus, 106.
In using sodium, 89.
In using thermometers, 353.
In using turpentine, 158, 159.
In weighing, 349.
Precipitate, definition, 8.
Pressure, correction for, 358, 363.
Priestley, 83, 137.
Problems.
Air, 114-115, 212, 224.
Alcohol, 243.
Aluminium, 154, 155, 306.
Ammonia gas, 186-187.
Ammonium compounds, 187.
Atomic weights, 154, 155. 170, 171, 256,
259, 275, 285, 289, 292, 296, 306,
312, 316, 318, 322.
Barium chloride, 67.
Boyle, law of, 359.
Bromine, 256.
Calcium, 155, 296.
Calcium carbonate, 212.
Carbon, 62, 115, 211, 212.
Carbonates, 214, 215.
Carbon dioxide, 62, 63, 211, 212.
Charles, law of, 355-356.
Chlorine, 170-171.
Chromium, 316-317.
Composition of water, 102-104.
4o8
Index.
Problems.
Copper, 103, 289.
Ether, 243.
Equivalents, 154.
Fluorine, 259.
Formula, 156, 188, 203, 243.
Gas volumes, reduction of, 364-366,
Hydrocarbons, 224.
Hydrochloric acid, 170-171, 212.
Hydrogen, 52, 62, 102, 103, 114, 171.
Sulphide, 268.
Illuminating gas, 227.
Iodine, 256.
Iron, 322.
Lavk' of Boyle, 359.
Of Charles, 355-356-
Lead, 312.
Magnesium, 37, 154, 155, 294.
Manganese, 318.
Manganese dioxide, 171, 172.
Mercuric oxide, 62.
Mercury, 302.
Metric system, lo-ii.
Molecular weight, 156.
Nitrates, 194.
Nitric acid, 194, 203.
Nitrogen, 114, 187.
Oxygen, 37, 62, 63, 102, 103, 114,1115, 224.
Percentage composition, 155, 171, 188,
195, 224, 243, 256, 259, 268, 272,
275, 290, 291, 309, 316.
Phosphorus, 115.
Potassium, 170, 285.
Potassium chlorate, 37, 62.
Review, 278, 366.
Silicon, 275,
Silver, 155, 170, 292.
Sodium, 103, 154, 285.
Sulphide, hydrogen, 268.
Sulphur, 267, 272.
Sulphur dioxide, 63.
Sulphuric acid, 62, 194.
Thermometers, 14, 243, 256.
Tin, 309.
Water, 102-104, 212.
Water of crystallization, 67.
Zinc, 52, 62, 154.
Problems based on equations, 60-61.
Proust, 35.
Qualitative analysis, 266, 342.
Radical, 184, 240.
Rayleigh and Ramsay, 105.
Reaction between zinc and sulphuric
acid, Exp. 22.
Reaction, chemical, 41.
Reagents, list of, 385-391.
Red powder, products obtained by heat-
ing the, Exp. 12.
Reducing agents, 102.
Reduction, 102, 308.
By carbon, Exp. 85.
Of gas volumes to standard conditions,
361-364.
With blowpipe, Exp. 105.
Residue, definition, 8, 34.
Safety tubes, 336.
Salt, ethereal, 241.
Salts, 121, 128.
General property of, 120, Exp. 50.
Nature of, 125.
Nomenclature of, 120-121.
Sand bath, 332.
Saturation, 79.
Scales, trip, 12, 347.
Scheele, 157.
Separation of aluminium and iron, Exp.
200,
Copper, iron, and sodium, Exp. 198.
Lead and silver, Exp. 193.
Lead, silver, and mercury, Exp. 194.
Zinc and aluminium, Exp. 201.
Zinc and iron, Exp. 199.
Silicic acid, preparation and properties,
Exp. 141.
Silicon, 273.
Silicon dioxide, relation of, to other com*
pounds of silicon, 274, Exp. 140.
Silicon tetrafluoride, 258.
Silver, 290.
Analysis of, coin, Exp. 195.
Preparation, Exp. 154.
Properties, Exp. 155.
Tests for, Exp. 156. •
Smelling and tasting, 8.
Soap, 247, 248.
Preparation, Exp. 116.
Properties, Exp. 117.
Index.
409
Sodium, 88.
And water, interaction of, 88, 90, 91,
Exp. 42.
General properties of compounds of,
Exp. 143.
Precautions in using, 89.
Properties, Exp. 142.
Sodium acetate, preparation, Exp.
115 {a).
Sodium amalgam, Exp. 63.
Sodium chloride, preparation of pure,
Exp. 144.
Sodium hydroxide, preparation, 280, Exp.
145-
Sodium nitrate and sulphuric acid, inter-
action of, 190, 191, Exp. 73.
Solder, analysis of, Exp. 197.
Solids, solubility of, 78, Exp. 34.
Transferring, 339.
Solubility, terms for expressing, 75.
Solute, definition, 75.
Solution, 75.
And chemical action, Exp. 37.
Thermal phenomena of, Exp. 36.
Solutions, list of, 385-391,
Of gases, 76, Exp. 32.
Of liquids, 77, 78, Exp. 33.
Of solids, 78, Exp. 34.
Solvent, definition, 75.
Specific heats, law of, 141.
Standard temperature, reducing to, 355.
Standard pressure, reducing to, 359.
Starch, detection of, Exp. 125.
Test for iodine, Exp. 124.
Stirring rods, Exp. 4.
Steam, density of, 99.
Volumetric composition of, loi.
Sublimate, definition, 299.
Subliming, definition, 299.
Substitution, definition, 45.
Sugar, Fehling's test for, Exp. 118.
Sulphate, definition, 45.
Test for, Exp. 2i{l>).
Sulphide, definition, 45.
Sulphide, hydrogen, 262.
Preparation and properties, Exp. 133.
Weight of liter of, 264.
Sulphides, 261, 265.
Preparation and properties of some,
Exp. 134.
Separation of, 266, Exp. 135.
Sulphites, 269.
Sulphur, amorphous, Exp. 130.
Combinmg power of. Exp. 132.
Crystallized, Exp. 131.
In organic compounds, Exp. 107 (d).
Oxides of, 268.
Physical properties of, Exp. 129.
Sulphur dioxide, preparation and prop-
erties, Exp. 136.
Sulphur trioxide, 269.
Sulphuric acid, 270.
Action with organic matter, Exp. 138.
Action with water, 270, Exp. 137.
Test for, Exps. 21 {d), 139.
Sulphurous acid, 269.
Supersaturation, 79, Exp. 35.
Symbols, 53.
List of, 392.
Synthesis, definition, 33.
Tad/es.
Colors of borax beads, 394.
Of coatings on charcoal, 393.
Of flames, 394.
Of residues moistened with cobalt
nitrate, 394.
Composition of air, 108.
Of typical coals, 396.
Conversion of thermometric readings,
394-
Densities and molecular weights of
gases, 144.
Equivalents and atomic weights, 136.
Freezing and boiling points, 74.
Important elements, symbols, and
atomic weights, 392.
Ionization, 152.
Metric equivalents, 9.
System, 346.
Transformation, 346.
Periodic arrangement of the elements,
396.
Solubility of carbon dioxide, 76.
Of salts in water, 395.
Of some liquids, 78.
Of typical solids, 79.
Specific gravity and melting point of
metals, 395.
Specific heats, 141.
4TO
Index.
Tables.
Tension of aqueous vapor, 361.
Weight of a liter of gases, 395.
Temperature, 352.
Absolute, 356.
Correction for, 355, 362.
Tension, aqueous, 360.
Vapor, 68, 69.
Test, definition, 42, 44.
Tests.
Acetic acid, Exp. 109 {e).
Alcohol, Exp. 109 {c).
Alkali, 90, Exps. 42 {d), 49.
Aluminium, Exp. 170.
Base, Exp. 49.
Calcium, Exps. 29 {d), 86 II, 159.
Carbon dioxide, Exps. 46, 86 1.
Chlorides, Exps. 29 (<^), 61 IV.
Chromium, Exp, 181.
Cobalt, Exp. 103 {a).
Copper, Exps. 103 {b), 151.
Hydrochloric acid, Exp. 61 IV.
Iodine, Exps. 123, 124.
Iron, Exps. 189, 190.
Lead, 310, Exps. 105 (a), 106 (/'). I77-
Magnesium, Exp. 158.
Manganese, Exps. 103 {c) , 186.
Mercury, Exps. 165, 166, 174.
Nature of common substances, Exp. 51.
Nitrate, Exps. 72, 78, 79, 80.
Nitric acid, Exp. 72.
Organic matter, Exps. 29 (a), 72, 138.
Potassium, Exp. 146.
Silver. See Exps. 29 {b), 61 IV.
Sodium, 90, Exp. 42.
Starch, Exp. 125.
Sugar, 153.
Sulphate, Exps. 21 {b), 105 {b).
Sulphides, Exp. 124.
Sulphuric acid, Exp. 21 {b).
Tin, Exps. 106 {c), 174.
Zinc, Exps. 21 {a), 106 {a), 161.
Thermal phenomena of solution, 81,
Exp. 36.
Thermometer, 352.
Using a, 353, 369.
Tin, 306.
Action with acids, Exp. 173.
Deposition of metallic, Exp. 175.
General properties, Exp. 172.
Tests for, 307, 308, Exps. 106 (a), 174,
Triangle, 331.
Tube, gas measuring, 350, 369.
Marchand, Exp. 44.
To dry the inside of a, 338.
U-, 367, 368, Exps. 23, 44.
U-tube, 367, 368, Exps. 23, 44.
Valence, 147.
Vinegar, 215.
Properties, Exp. 114.
Water, action of, on lead, 310, Exp. 179.
And hydrogen, 86, 90, Exps. 40, 42.
And oxygen, 88, Exp. 41.
And sodium, 88, Exp. 42.
Bath, 332.
Decomposition of, by chlorine, P2xp. 41.
By iron, Exp. 40.
Electrolysis of, 83, Exps. 38, 39,
Freezing and boiling points of, Exp. 31.
General distribution, Exp. 24.
Gravimetric composition of, Exp. 44.
Hardness of, 248, Exp. 117 {d).
Impure, 69.
Mineral, 69.
Molecular weight of, 146.
Purification of, 70, Exp. 30.
Quantitative composition of, 91.
Relation to temperature, 73, Exp. 31.
Tests for impure, Exp. 29.
Volumetric composition of, 95, Exps.
38, 39. 43-
Water of crystallization, 65-67.
Detection of, Exp. 25.
Determination of, Exp. 26.
Weighing, Exp. 7.
And measuring, Exp. 8.
Rules for, 348.
Weights, 347.
Zinc, 296.
Atomic weight of, 145.
Equivalent of, 45-48, 145, Exp. 22,
General properties, Exp. 160.
Interaction of, with metals, Ex;). 162.
Tests for, Exps. 21 {a), 161.
Zinc sulphate, crystallization of, Exp. 20.
Test for, Elxp. 21,
ADVERTISEMENTS
Science
Austin's Clinical Chemistry. A manual of applied physiological chemistry. $1.75.
Benton's Guide to General Chemistry. A manual for the laboratory. 35 cents.
Boyer's Laboratory Manual in Biology. Treats of both animals and plants. 80 cts.
Boynton, Morse and Watson's Laboratory Manual in Chemistry. 50 cents.
Burrage and Bailey's School Sanitation and Decoration. Illustrated. ;?i.5o.
Cheston, Gibson and Timmerman's Physics. Theoretical and descriptive. ^11.25.
Chute's Physical Laboratory Manual. Revised edition. Illustrated. 80 cents.
Chute's Practical Physics. For schools and colleges. $1.12.
Clark's Methods in Microscopy. Detailed descriptions of successful methods. $1.60,
Coit's Chemical Arithmetic. With a short system of analysis. 50 cents.
Coleman's Elements of Physics. For secondary schools. $1.25.
Colton's Physiology : Practical and Descriptive. Illustrated. $1.40.
Colton's Physiology : Briefer Course. For earlier years in high schools. Illus-
trated. 90 cents.
Colton's Practical Physiology. A laboratory course. 60 cents.
Colton's Zoology : Descriptive and Experimental. Illustrated. $1.50. Parti,
Descriptive, $1.00. Part 11, Experimental, 60 cents.
Fisher and Patterson's Elements of Physics. Experimental and descriptive. 60
cents.
Grabfield and Bums's Chemical Problems. For review and drill. Paper, 25 cts.
Hyatt's Insecta. A practical manual for students and teachers. Illustrated. $1.25.
Newell's Descriptive Chemistry. A full exposition of modern inorganic chem-
istry. Illustrated. #1.20. Part I, Without experiments. $1.00. Part II,
Experiments. 40 cents.
Newell's Experimental Chemistry. A modern course for high schools and col-
leges. $1.10.
Orndorff's Laboratory Manual in Organic Chemistry. Boards. 35 cents.
Palmer's Questions and Problems in Chemistry. 20 cents.
Pepoon, Mitchell and Maxwell's Plant Life. A laboratory guide. 50 cents.
Remsen's Organic Chemistry. $1.20.
Roberts's Stereo-Chemistry. Its development and present aspects. $1.00.
Sanford's Experimental Psychology. Part I. Sensation and perception. $1.50.
Schoch's Experiments and Discussions in Chemistry. 50 cents.
Shaler's First Book in Geology. Cloth, 60 cents. Boards, 45 cents.
Shepard's Inorganic Chemistry. Descriptive and qualitative. $1.12.
Shepard's Briefer Course in Chemistry, with chapter on Organic Chemistry. 80 cts.
Shepard's Laboratory Note-Book. Boards. 35 cents.
Spalding's Botany. Practical exercises in the study of plants. 80 cents.
Stevens's Introduction to Botany. Illustrated. $1.25. Key and Flora, 40 cents.
Botany, with Key and Flora, $1.50.
Stevens's Chemistry Note-Book. Laboratory sheets and covers. 50 cents.
Venable's Short History of Chemistry. For students and the general reader. $1.00.
Weed and Crossman's Laboratory Guide in Zoology. Emphasises essentials. 60 cts.
Whiting's Physical Measurement. Parts 1-1 V, in one volume. $3.75.
Whiting's Mathematical and Physical Tables. Paper. 50 cents.
J^or eloiietttary ivorks see onr list 0/ books in Elementary Science.
D. C. HEATH & CO., Publishers, Boston, New York, Chicago
Mathematics
Barton's Plane Surveying. With tables and a chapter on levelling. Jii.50.
Barton's Theory of Equations. A treatise for college classes. $1.50.
Bauer and Brooke's Plane and Spherical Trigonometry, ;fi.5o.
Bowser's Academic Algebra. For secondary schools. #1.12.
Bowser's College Algebra. A full treatment of elementary and advanced topics. $1.50.
Bowser's Plane and Solid Geometry. $1.25, Plane bound separately. 75015.
Bowser's Elements of Plane and Spherical Trigonometry, ^octs. ; with tables, $1.40.
Bowser's Treatise on Plane and Spherical Trigonometry. $1.50.
Bowser's Five-Place Logarithmic Tables, sects.
Candy's Analytic Geometry. Plane and Solid, $1.50.
Cohen's Differential Equations. An elementary treatise. JI2.00.
Fine's Number System of Algebra. Theoretical and historical. $1.00.
Gilbert's Algebra Lessons. Three numbers : No. i, to Fractional Equations; No. 2,
through Quadratic Equations ; No. 3, Higher Algebra. Each number, per dozen, ^11.44.
Hopkins's Plane Geometry. Follows the inductive method. 75 cts.
Rowland's Elements of the Conic Sections. 75 cts.
Lefevre'S Number and its Algebra. Introductory to college courses in Algebra. $1.25.
Lyman's Geometry Exercises. Supplementary work for drill. Per dozen, |i.6o.
McCurdy's Exercise Book in Algebra. A thorough drill book. 60 cts.
Nichols's Analytic Geometry. A treatise for college courses. $1.25.
Nichols's Calculus. Differential and Integral. #2.00.
Osborne's Differential and Integral Calculus. Revised edition. $2.00.
Peterson and Baldwin's Problems in Algebra. For texts and reviews. 30 cts.
Robbins's Surveying and Navigation. A brief and practical treatise, sects.
Waldo's Descriptive Geometry. With problems and suggestions. 80 cts.
Wells's Academic Arithmetic. With or without answers. $1.00.
Wells's Algebra for Secondary Schools. $1.20. Pocket er^ition, $1.20.
Wells's Essentials of Algebra. For secondary schools. $1.10.
Wells's Academic Algebra. With or without answers. $1.08.
Wells's Text Book in Algebra. $1.40.
Wells's New Higher Algebra. For schools and colleges. J 1.32.
Wells's Higher Algebra. S1.32.
Wells's University Algebra. Octavo. $1.50.
Wells's College Algebra. $1.50. Part II, beginning with quadratics, ^1.32.
Wells's Advanced Course in Algebra. A college algebra. $1.50.
Wells's Essentials of Geometry. $1.25. Plane, 75 cts. Solid, 75 cts.
Wells's New Plane and Spherical Trigonometry. For colleges and technical schnof.«
$i.oo. With six-place tables, $1.25. With Robbins's Surveying and Navigation, f '. so
Wells's Complete Trigonometry. Plane and Spherical. Qo cts. With tables, #1.08.
Plane, bound separately, 75 cts.
Wells's New Six-Place Logarithmic Tables. 60 cts. Pocket edition, small type 36 crs
Wells's Four-Place Tables. 2s cts.
Wright's E::ercises in Concrete Geometry. 30 cts.
For Arithmetics see our list of books in Elementary Mathematics.
D. C. HEATH & CO., Publishers, Boston, New York, Chicago
History
Allen's History Topics. Covers Ancient, Modem, and American history and gives an
excellent list ot books of reference. 121 pages. Paper, 25 cents.
Allen's Topical Outline of English History. Including references for literature. Boards,
25 cents ; cloth, 40 cents.
Boutwell's The Constitution of the United States at the End of the First Century.
Presents the Constitution as it has been interpreted by decisions of the United States Su-
preme Court from 1789 to 1889. 430 pages. Buckram. J->.5o ; law sheep, $3.5°.
Fisher's Select Bibliography of Ecclesiastical History. An annotated list of the most
essential books for a theological student's library. 15 cents.
Flickinger's Civil Government: as Developed in the States and the United States.
An historical and analytic study of civil institutions, for schools and colleges. 374 pages-
Cloth, $1.00.
Hall's Method of Teaching History. " Its excellence and helpfulness ought to secure it
many readers."— T/ie Nation. 405 pages. $1.50.
Pratt's America's Story for America's Children. A series of history readers for ele-
mentary schools.
1 The Beginner's Book. Cloth. 60 illustrations. 132 pages. 35 cents.
II. Discoverers and Explorers: 1000 to 1609. Cloth. 152 pages. 52 illus, 40 cents.
III. The Earlier Colonies: 160 1 to 1733. Cloth. 160 pages. Illus. 40 cents.
IV. 'The Later Colonies. Cloth. Illus. 160 pages. 40 cents.
V. The Revolution and the Republic. Cloth. Illus. 160 pages. 40 cents.
Sheldon's American History. Follows the "seminary" or laboratory plan. " By it the
pupil is not robbed of the right to do his own thinking." Half leather. ^1.12.
Teacher's Manual to Sheldon's American History. 60 cents.
Sheldon's General History. For high schools and colleges. The only general history
following the " seminary " or laboratory plan. Half leather. 572 pages. #1.60.
Teacher's Manual to Sheldon's History. Puts into the instructor's hand the key to the
above system. 172 pages. 85 cents.
Sheldon's Greek and Roman History. Contains the first 250 pages of the General
History, ^i.oo.
Sheldon-Barnes's Studies in Historical Method. Suggestive studies for teachers and
students. Cloth. 160 pages, go cents.
Shumway's A Day in Ancient Rome. With 59 illustrations. Should find a place as a
supplemefitary reader in every high-school class studying Cicero, Horace, Tacitus, etc.
96 pages. Paper, 30 cents ; cloth, 75 cents.
Thomas's Elementary History of the United States. For younger grades. Maps and
illustrations. Cloth. 357 pages. 60 cents.
Thomas's History of the United States. Revised and rewritten. Edition of 1901. For
schools, academies, and the general reader. A narrative historv \vith copious references
to sources and authorities. Fully illustrated. 592 pages. Half leather, $1.00.
English History Readers. English history for grammar grades.
SVilson's Compendium of United States and Contemporary History. For schools and
the general reader. 1 14 pages. 40 cents.
57il80n'S The State. Elements of Historical and Practical PoUtics. A book on the
organization and functions of government. Revised edition, largely rewritten. 69;.
pages. 552.00.
Sent postpaid on receipt of price by the publishers.
D. C. HEATH & CO., Publishers,Boston, NewYork,Chicago
Heath's English Classics.
V.
Addison's Sir Roger de Coverley Papers. Edited by W. H. Hudson. Cloth. 232 pages.
Nine fuU-^iage illustrations and two maps. 35 cents.
Burke's Sp ech on Conciliation with America. Edited by A. J. George, Master in
the Newton Olass.) High School. Cloth. 119 pages. 20 cents.
Carlyle's Essay on Burns. Edited, with introduction and notes, by Andrkw J. George.
Cloth. 15.J pa5,'es. Illustrated. 25 cents.
Coleridge's Rime of the Ancient Mariner. Edited by Andrew J. George. Cloth.
96 pages. Illustrated. 20 cents.
Cooper's Last of the Mohicans. Edited by J. G. Wight, Principal Girls' High School,
New York City. Cloth. Illustrated. 659 pages. 50 cents.
De Ouincey's Flight of a Tartar Tribe. Edited by G. A. Wauchope, Professor in the
University of South Carolina. Cloth. 112 pages. 25 cents.
Dryden's Palamon and Arcite. Edited by William H. Cravvshaw, Professor in Col-
gate University. Cloth. 158 pages. Illustrated. 25 cents.
George Eliot's Silas Marner. Edited by G. A. Waitchpoe, Professor in the University
of South Carolina. Cloth. 288 pages. Illustrated. 35 cents.
Goldsmith's Vicar of Wakefield. With introduction and notes by W. H. Hudson.
Cloth. 300 pages. Seventeen illustrations by C. E. Brock. 50 cents.
Irving's Life of Goldsmith. Edited bv H. E. Coblentz, South Division High School,
Milwaukee. Cloth. 328 pages. Maps and illustrations. 35 cents.
Macaulav's Essay on Milton. Edited by Albert Perry Walker, Master in the
English High Scliool, Boston. Cloth. 146 pages. Illustrated. 25 cents.
Macaulay's Essay on Addison. Edited by Albert Perry Walker. Cloth. 192 pages.
Illustrated. 25 cents.
Macaulay's Life of Johnson. Edited by Albert Perry Walker. Cloth. 122 pages.
Illustrated. 25 cents.
Milton's Paradise Lost. Books i and ii. Edited by Albert Perry Walker. Cloth.
188 pages. Illustrated. 25 cents.
Milton's Minor Poems. Edited by Aluert Perry Walker. Cloth. 190 pages.
Illustrated. 25 cents.
Pope's Translation of the Iliad. Books i, vi, xxii, and xxiv. Edited by Paul Shorey,
Professor in the Univ. of Chicago. Cloth. 174 pages. lUus. 25 cents.
Scott's Ivanhoe. Edited by Porter Lander MacCltntock. Cloth. 556 pages. Seven-
teen full-page illustrations by C E. Brock. 50 cents.
Scott's Lady of the Lake. Edited bv L. Dupont Syle, Professor in the University of
Cahfornia. Cloth. 216 pages. Illus. and map. 35 cents.
Shakespeare. ?,Q&Xhe: Arden Shnkes/>enre. Per vol., 25 cents.
Tennyson's Enoch Arden, and the two Locksley Halls. Edited by Calvin S. Brown,
University of Colorado. Cloth. 168 pages. 25 cents.
Tennyson's Idylls of the King. Four idylls, edited by Arthur Beatty, University of
Wisconsin. Cloth. 190 pages. Illus. and map. 25 cents.
Tennyson's The Princess. With introduction and notes by Andrew J. George. Cloth.
148 pages. Illustrated. 25 cents.
Webster's First Bunker Hill Oration. With introduction and notes by Andrew J.
George. Cloth. 55 pages. 20 cents.
D. C. HEATH & CO., Boston, New York, Chicago
BOSTON UNIVERSITY
1 17n D17M7 5353
} 30
DIRECT ^
Jn to
I
<ET
^m ooK
MEMBER
.♦i??^^
BOSTON UNIVERSITY LIBRARIES
771 Commonwealth Ave,
Boston, Mass. 02215
Live Music Archive
Librivox Free Audio
Metropolitan Museum
Cleveland Museum of Art
Internet Arcade
Console Living Room
Books to Borrow
Open Library
TV News
Understanding 9/11