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Full text of "Laboratory exercises in inorganic chemistry"

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INTERNATIONAL CHEMTCAL SERIES 
H. P. TALBOT, Ph.D., Sc.D., Consulting Editor 



LABORATORY EXERCISES 

IN 

INORGANIC CHEMISTRY 



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LABORATORY EXERCISES 

IN 

INORGANIC CHEMISTRY 



BY 

JAMES F. NORRIS 

Professor of Organic Chemistry, Massachusetts Institute 

OF Technology; Author of "A Textbook of Inorganic 

Chemistry for Colleges," "The Principles op 

Organic Chemistry" and "Experimental 

Organic Chemistry" 



KENNETH L. MARK 

Professor of Chemistry, Simmons College 



First Edition 



McGRAW-HILL BOOK COMPANY, Inc, 

NEW YORK: 370 SEVENTH AVENUE 
LONDON: 6 & 8 BOUVERIE ST., E. C. 4 

1922 



Copyright 1922, by the 
McGraw-Hill Book Company, Inc. 

I 6 3 ^ fi- 






PREFACE 

This laboratory manual contains, in addition to the experi- 
ments ordinarily found in books of this class, a large number 
of new experiments that have been designed to illustrate the 
general principles which are being more and more emphasized 
in elementary instruction in chemistry. About one-fourth of 
the book is devoted to the study of gases, liquids, and solids, 
solutions, thermo- and electro-chemistry, and chemical equili- 
brium. Directions for a number of inorganic preparations are 
also given. In designing these experiments the authors have 
had in mind the student who has had a good training in 
chemistry in the high school, and who should have his interest 
stimulated in his college course and should be given the oppor- 
tunity to extend his experience. Many of the experiments are 
quantitative ; but care has been taken to have the manipula- 
tions and apparatus of such a simple character that the exer- 
cises can be performed successfully by the average student in 
a laboratory for beginners. 

The experiments to illustrate the properties of metals and 
their compounds are followed by an outline of a general 
method for testing a simple salt for the metal and acid radical 
present. The short time available in the first year for qualita- 
tive analysis makes it impossible for the student to master 
the methods used in the separation of the metallic elements. 
He can, however, by the simple procedure outlined learn to 
identify pure salts. The method described has been found 
to be an excellent introduction to qualitative analysis. 

Each experiment is preceded by a ''discussion" which em- 
phasizes the principles to be illustrated. The questions incor- 
porated into the directions for the experiments have been 
selected in such a way that it is necessary for the student to 
read the discussion in order to answer them. As a result, it 
is impossible for him to complete the record of his work with- 
out understanding what is to be learned from the experiment. 

Questions in connection with the laboratory directions serve 



VI PREFACE 

another purpose. It is the opinion of the authors that the 
student can not be left to himself in his first experimental 
work; he must be guided if he is to learn to observe closely 
and to differentiate the essential from the non-essential. A 
question in the right place furnishes this guidance and avoids, 
in most cases, the repetition of the experiment. 

The notes of the student consist, in the main, of answers 
to these numbered questions. The plan makes it possible for 
the instructor to see at a glance whether the student has under- 
stood his work. 

The authors are greatly indebted to Raymond E. Neal, In- 
structor in Chemistry, Simmons College, for valuable sugges- 
tions and for the preparation of the drawings, from which the 
cuts were made. 

James F. Norris 
Kenneth L. Mark 

Boston, Mass. 
Aj/ril, 1922. 



CONTENTS 

PAGE 

Preface v 

General Directions 1 

Physical and Chemical Changes i^O 

Classification of Changes, 10. 

Mixtures and Pure Substances 14 

Mixtures and Pure Substances, 14 — Formation of a Compound, 16 — 
Decomposition of a Compound, 20. 

Oxygen 24 

Preparation of Oxygen, 24 — Properties of Oxygen, 30 — Kindling 
Temperature, 32. 

Hydrogen 36 

Preparation of Hydrogen by Various Means, 36 — Preparation of 
Hydrogen in Quantity, 40 — Properties of Hydrogen, ^A. 

Quantitative Studies of Chemical Reactions 48 

Percentage of Oxygen in Potassium Chlorate, 48 — Percentage of 
Chlorine in Silver Chloride, 50 — Vapor Pressure, 52 — Boyle's 
and Charles' Laws, 56 — Weight of a Liter of Oxygen, 60 — 
Atomic Weight of Zinc or of Aluminium, 66. 

Water 74 

"occurrence of Water, 74 — Purification of Water, 76 — Hydrates, 80- - 
Efflorescence, Dehquescence, 82. 

Chlorine 88 

Preparation of Chlorine, 88 — Properties of Chlorine, 92 — Other 
Means of Preparing Chlorine, 96. 

Hydrochloric Acid 102 

Preparation of Hydrochloric Acid, 102 — Properties of Hydrochloric 
Acid, 104 — Preparation of Lead Chloride, 110 — Quantitative 
Determination of the Hydrogen and Chlorine in Hydrochloric 
Acid, 112. 

Hydrogen Peroxide 118 

Preparation and Properties of Hydrogen Peroxide, 118 — Quantitative 
Analysis of Hydrogen Peroxide, 120. 

Properties of Liquids and Solids 124 

Determination of Density, 124 — Determination of Specific Heat, 
126— Effect of Pressure on the Boiling Point of Water, ISO- 
Heat of Vaporization, 136 — Distillation of a Mixture of Two 
Liquids, 138 — Determination of a Freezing Point, 142 — Super- 
cooling, 146. 

vii 



viii CONTENTS 

PAGE 

Carbon and its Compounds 152 

Properties of Carbon, 152 — Preparation of Carbon Dioxide, 156 — 
Properties of Carbon Dioxide, 158 — Formation of Carbon 
Dioxide by Various Means, 162 — Formation of Carbonates, 166 
— Coal, 168 — Flame and the Bansen Burner, 170. 

Acids, Bases, and Salts 178 

Properties of Acids and Bases, 178 — Quantitative Neutralization. 
Titration, 184 — Percentage of Acetic Acid in Vinegar. Rela- 
tive Cost of Ammonium Hydroxide Solutions, 190. 

Solutions 194 

Formation of Ions, 194 — Electrical Conductivity of Solutions, 198 — 
Determination of the Heat of Neutralization, 202 — Electro- 
motive Series, 206 — Determination of Molecular Weights from 
the Freezing Points of Solutions, 210 — Determination of the 
Degree of Dissociation of Electrolytes from the Freezing Point 
of Solutions, 214. 

Chemical Equilibrium 220 

Reversible Reactions, 220 — Determination of the Effect of Increase 
in Temperature on the Rate of a Reaction, 224. 

Sulphur and Sulphides 232 

AUotropic Forms of Sulphur, 232 — Preparation of Hydrogen Sul- 
phide, 236 — Properties of Hydrogen Sulphide, 238 — Sulphides, 
240. 

Oxides and Acids of Sulphur 244 

Formation of Sulphur Dioxide, 244 — Preparation of Sulphur Dioxide, 
246 — Properties of Sulphur Dioxide, Sulphurous Acid, and 
Sulphites, 250— Sulphuric Acid, 256. 

Nitrogen and the Atmosphere 262 

Preparation and Properties of Nitrogen, 262 — The Atmosphere. 
Percentage of Oxygen, 266 — Quantitative Determination of 
Carbon Dioxide in Exhaled Air, 268 — Dew Point. Relative 
Humidity, 272. 

Ammonia 280 

Sources of Ammonia, 280 — Preparation and Properties of Am- 
monia, 282 — Determination of the Percentage of Nitrogen in 
Ammonia, 288. 

Acids and Oxides of Nitrogen 294 

Preparation and Properties of Nitric Acid, 294 — Nitrates, 298 — 
Nitrous Acid and Nitrites, 304 — Nitrous Oxide, 306 — Nitric 
Oxide, Nitrogen Dioxide, and Nitrogen Tetroxide, 308. 

Atomic and Molecular Weights 316 

Molecular Weight of a Gas, 316 — Determination of the Molecular 
Weight of Carbon Tetrachloride, 318 — Determination of the 
Atomic Weight of Manganese, 322 — Determination of the • 
Atomic Weight of Copper, 324. 



CONTENTS ix 

PAGE 

The Halogens 330 

Preparation and Properties of Bromine, 330 — Hydrobromic Acid, 
332— Iodine, 334— Ilydriodic Acid, 338— Hydrofluoric Acid, 
340 — Comparisons of the Halogens and of the Halides, 342 — 
Hypochlorites and Chlorates, 346. 

Phosphorus, Arsenic, Antimony, and Bismuth 350 

Phosphoric Acids and Phosphates, 350 — Halides of Phosphorus, 
354 — Preparation of Arsenic, 356 — Comparison of Phosphorus, 
Arsenic, Antimony, and Bismuth, 356. 

Organic Compounds 366 

Properties of Cellulose, 366 — Preparation of Soap, 368. 

Silicon and Boron 372 

Silicic Acid and Silica, 372 — Compounds of Boron, 374. 

Physical Properties of the Metals. Alloys 378 

Density of Metals, 378 — Hardness of Metals, 378— Alloys of Lead 
and Tin, 380. 

Chemical Properties of the Metals 388 

Behavior of Metals with Water, 388 — Action of Non-oxidizing 
Acids on Metals, 390 — Action of Oxidizing Acids on Metals, 
392 — Effect of the Presence of Oxj^gen on the Action of Acids 
on Metals, 396 — Action of Metals on Alkalies, 396. 

Electrochemistry 402 

Production of an Electric Current from Chemical Energy, 402 — 
Metallic Couples, 406 — Faraday's Law, 408 — The Lead Storage 
Battery, 414. 

Properties of Hydroxides and Salts 418 

Effect of Ammonium Salts on the Precipitation of the Hydroxides of 
the Bivalent Metals by Ammonium Hydroxide, 418 — Isomor- 
phism of Salts, 420— Heat of Solution of Salts, 422— Effect of 
Temperature on the Solubility of Salts, 424 — Solubility of 
Salts in Acids, 428 — Solubility of Salts in Solutions of Ammonia, 
432 — Preparation of a Double Salt: Ferrous Ammonium Sul- 
phate, 434 — Properties of Complex Salts, 436 — Hydrolysis of 
Salts, 438. 

Sodium and Potassium 444 

Action of Air upon Sodium, 444 — Baking Powder, 446 — Preparation 
of Sodium Carbonate, 450 — Flame Tests, 452 — Preparation 
of Potassium Nitrate, 456. 

Calcium, Strontium, and Barium 460 

Hard Water, 460 — Plaster of Paris, 464 — Comparison of Com- 
pounds of Calcium, Strontium, and Barium, 466. 

Magnesium, Zinc, and Mercury 472 

Magnesium Nitride, 472 — Magnesium Chloride, 472 — Zinc Hy- 
droxide, 474— Zinc Chloride, 476— Zinc Sulphide. A Test 
for Zinc, 478 — Compounds of Mercury, 480. 



X CONTENTS 

PAGE 

Aluminium 488 

Ahiminium Hydroxide, 488 — Method of Cleaning Silver, 490. 

Tin and Lead 494 

Hydroxides of Tin, 494— Chlorides of Tin, 496— Sulphides of Tin, 
498 — Oxides of Lead, 500 — Preparation of Lead Dioxide, 504 — 
Salts of Lead, 506. 

Silver and Copper 508 

Preparation of Metallic Silver, 508 — Cupric Compounds, 510 — 
Cuprous Compounds, 512 — Preparation ^ of Cupric-ammonia 
Sulphate, 514. 

Iron 518 

Tests for Iron, 518 — Reduction and Oxidation of Iron Salts, 520 — 
Preparation of Ferric Ammonium Alum, 522. 

Chromium 526 

Chromic Anhydride, Chromates, and Bichromates, 526 — Reduction 
of Potassium Bichromate, 528. 

Manganese 532 

Potassium Permanganate as an Oxidizing Agent, 532. 

Identification of Compounds 534 

Appendix 542 

Vapor Pressure of Water, 542 — Electromotive Series of the Ele- 
ments, 542 — Table of Solubilities, 543 — Lists of Apparatus, 544 
— List of Chemicals, 546. 



LABORATORY EXERCISES IN 
INORGANIC CHEMISTRY 



GENERAL DIRECTIONS 

The Object of Laboratory Work: Laboratory work is a 
form of study. A thorough knowledge of elementary chem- 
istry can not be acquired from books alone. Actual first-hand 
observation of the properties of substances must be made by 
the student himself, but even these observations will be of 
small value unless he appreciates the object of the experiment 
and the principle which it illustrates. Merely ''doing" the 
experiments, that is, mechanically following directions, is as 
valueless as reading a book without comprehending its mean- 
ing. No experiment should be begun until the discussion is 
thoroughly understood. 

Notes: Observations must be recorded immediately. To 
facilitate clearness and brevity they are to be entered on 
the blank pages facing the text, as answers to the numbered 
questions. 

The data obtained in quantitative experiments should be 
carefully labeled and the result of each measurement should 
be placed on a separate line, as illustrated by the following 
example in the determination of the percentage of oxygen in 
potassium chlorate, as carried out in Experiment 11. 

Wt. crucible empty 20.30 gms. 

Wt. crucible and KCIO, 22.03 

Wt. KCIO3 1.73 

Wt. crucible and KCIO3 after first heating 21.42 



2 LABORATORY EXERCISES 

Wt. crucible and KCIO3 after second heating 21.35 

Wt. crucible and KCIO3 after third heating 21.35 

Loss of weight = weight of oxygen 0.68 

Percentage of oxygen in KCIO3 39.3 

The data must be entered at once directly in the book, never 
on loose pieces of paper. All calculations should be made in 
the book. 

Individual Work: All experiments should be done entirely 
independently. A student can get no more out of his labora- 
tory exercises than he puts into them. If he allows another 
student to perform the experiments for him or copies his 
notes, he is deriving no benefit from his work, he is wasting 
his time and, moreover, he is dishonest. 

Use of the Text Book : The numbers given in the references 
are those of the paragraphs in **A Text Book of Inorganic 
Chemistry," by James F. Norris. The text book should be 
brought to the laboratory at each exercise, the paragraphs 
pertaining to the exercise of the day should be read, and 
reference to the book should be made whenever the student 
is in doubt. 

Repetition of Experiments ; If the results obtained are not 
those which you have been led to expect, search for the cause, 
making sure particularly that you have followed the direc- 
tions precisely. Not until a possible cause of error has been 
detected sliould the experiment be repeated. 

Cleanliness: Apparatus must be clean in order to obtain 
trustworthy results. Always have a clean towel at hand. 

The desk top should be frequently wiped with a wet 
sponge and any material which is spilled upon it should be 
washed off immediately with plenty of water. 

General apparatus must be cleaned and returned to its place 
at the close of each exercise. 

Materials: All the materials required for an experiment 
and listed in the text should be obtained before beginning 
work. It is assumed that each desk is supplied with solutions 
of dilute hydrochloric, sulphuric, and nitric acids, of sodium 
and ammonium hydroxides, and with concentrated sulphuric 



LABORATORY PROCESSES 3 

acid. These substances therefore are not included in the lists. 
It is also assumed that each student is provided with the set 
of apparatus given in the appendix and only additional ap- 
paratus is included in the list of materials. 

Only such amounts of substances as are actually required 
should be obtained. Stock bottles should never be carried to 
the student's desk. Solids may be measured out on a watch 
glass or on a piece of filter paper and liquids may be carried 
in a test tube or beaker. 

Always read the label on a stock bottle before taking ma- 
terial from it, and verify your reading by rereading after you 
have taken the substance. 

Unused material must never be returned to the stock bottle 
but thrown away into the crock or sink. 

Waste Material: Solids, particularly matches, should be 
disposed of in the crock, not in the sink. Liquids should be 
poured into the sink, and in the case of concentrated acids, 
the sink should be flushed immediately with a large amount 
of water. 

Accidents: Any corrosive liquid on the skin should be 
washed off immediately with plenty of water. 

Burns and cuts, even when slight, should be reported to an 
instructor for treatment. 

Acid spilled upon the clothing should be neutralized with 
ammonia. 

Fires should be extinguished by throwing a wet towel over 
the blaze. 

LABORATORY PROCESSES 

1. Heating: The Bunsen burner is lighted by first turning 
on the gas and then applying a burning match. This order of 
procedure allows time for the air in the tube and burner to 
be expelled while the match is being '* struck." The burner 
should always be lighted before it is placed under the appar- 
atus to be heated, never after it has been put in position. 

The character of the flame is regulated by adjusting the 
quantity of air which enters through the holes at the bottom 



4 LABORATORY EXERCISES 

of the burner. Ordinarily a clear bluish flame should be used. 
Sometimes the flame "strikes back," that is, begins to burn 
at the base where the air enters. The occurrence is usually 
due to the admission of too much air. Turn off the gas, and 
after the burner has cooled, diminish the amount of air by 
adjusting the movable ring and relight the burner. 

Thick vessels, like a bottle or a mortar, should never be 
heated because the uneven expansion of the material causes 
them to break. 

Porcelain evaporating dishes and crucibles can be heated 
directly in the flame, but beakers should rest on a piece of 
wire gauze. 

Test tubes should be held in the test tube holder. AVhen 
a liquid is being heated, the test tube should be inclmed and 
held in such a position that the flame strikes the glass oppo- 
site the upper part of the liquid. The tube should be slightly 
shaken constantly. If the tube is held still and the bottom 
heated, a large amount of steam may be formed suddenly 
and throw the contents out of the tube. 

All vessels must be dry on the outside when heated; other- 
wise they may crack. 

2. Filtering: Fold a disc of filter paper just in halves and 
then again in quarters. Open one of the segments, leaving 
three thicknesses of paper on one side and one on the other. 
A paper cone will be formed Avhich should be placed in a glass 
funnel. If water solutions are to be filtered, wet the paper 
and press it firmly to the glass with your finger. When a 
filtration is to be made into a beaker or an evaporating dish, 
the funnel should be supported by a ring or by a filter arm 
and the apparatus adjusted so that the stem of the funnel 
touches the side of the vessel which is to catch the liquid which 
comes through; otherwise the liquid will spatter out as it 
falls. 

3. Drying Test Tubes: Slip the washed test tube over the 
end of a glass tube about 1 foot in length and. inclining it 
mouth downward, wave it back and forth through the flame 
of a burner. When the test tube is hot blow through the 
glass tube for an instant. Repeat the procedure if necessa^^^ 



LABORATORY PROCESSES 6 

4. Preparing Glass Tubing: (a) Cutting: To cut a piece 
of glass tubing, first make a slight scratch on the tube with the 
edge of a triangular file. Then, holding the tube in both hands, 
place the thumbs together against the tube on the side opposite 
the scratch. Press the tube as though bending it away from 
the scratch. A clean break should result. The broken edge 
of the tube will be very sharp and must alivays be smoothed by 
holding the end of the tube in the flame until it becomes a dull 
red color. This process, which is called ''fire polishing," causes 
the sharp edges of the glass to melt and thereby become smooth. 

{!)) Bending: A good bend is a smooth curve rather than a 
sharp angle. Several inches of the tube, therefore, must be 
involved in the bend, and it becomes necessary to soften the glass 
by heating it in a wide flame. Put a ''flame spreader" on the 
top of the Bunsen burner and close the holes at the bottom of 
the burner so as to give a luminous flame. Hold the tube in 
the yellow part of the flame, letting one end rest lightly be- 
tween the fingers and thumb of the left hand and turning the 
tube slowly, but constantly, in one direction with the fingers and 
thumb of the right hand. When the tube feels pliable, remove 
it from the flame and make a bend of the desired angle. 

The following bent tubes should be prepared as they will be 
used frequently in these experiments : two right-angled bends each 
leg of which is 10 cm. long; one right-angled bend one leg of 
which is 10 cm., the other 20 cm. long; one 45° bend one leg of 
which is 10 cm., the other 15 cm. long; and one 135° bend one 
leg of which is 10 cm., the other 15 cm. long. 

5. Preparing Corks: Before being used corks should be 
softened. This can be done by means of a press, which is made 
for this purpose, or the cork can be rolled on the desk while it 
is being pressed firmly by means of a block of wood. Sharp 
cork borers should be used to make the holes of such a size 
that the tubes to pass through fit snugly. In boring corks it 
is advisable to push the borer with a rotary motion half way 
through the cork, taking care that the hole is bored through 
the center of the cork; the borer is then removed and a hole 
made from the center of the other end of the cork to meet that 
first made. By proceeding in this way the edges of the holes 



6 LABORATORY EXERCISES 

on the two sides of the cork will be clean cut, and thus make' 
a tight joint with the tube to be passed through the hole ; and 
the latter will run evenly through the axis of the cork. 

6. Setting Up Apparatus : Place the front of the ring stand 
parallel to the edge of the desk and so that the rod is away 
from you, not toward you. Arrange the clamps so the main 
weight of the apparatus is over the base of the stand. The 
movable jaw of the clamp should be on top. Be careful not 
to screw the clamp so tight as to crush thin apparatus. Keep 
vertical lines vertical and horizontal lines horizontal. 

In putting a glass tube through a rubber stopper, wet the 
outside of the tube and then push it in with a twisting motion. 
Give three distinct pushes, pause long enough to count three 
and then give three more pushes. If you do not pause, you 
involuntarily keep pushing harder and harder, often with the 
result that the tube breaks and the jagged end cuts the hand. 
Be sure that all stoppers fit tightly and that there are no 
leaks around the holes of stoppers through which glass tubes 
pass. 

7. Weighing: Objects and material to be weighed fall into 
two general classes, those in which only an approximate 
weight, to the nearest gram, is required and those in which an 
accurate weight, to the nearest centigram, is demanded. Cases 
in which a convenient amount of material for experimentation 
is to be weighed out belong to the first class ; the quantitative 
determination of the composition of substances belongs to the 
second class. Platform scales are used in the first instance; 
beam balances are necessary in the second. 

(a) The platform scales must always be tested before using 
them because they are exposed to the fumes of the laboratory, 
and easily become rusted. Be sure that the rider on the front 
of the scales is pushed to the extreme left, and then gently tap 
one of the platforms. The pointer should swing an equal num- 
ber of divisions on each side of the middle of the pointer scale 
and should not come to rest until at least two swings on each 
side have been completed. If the scales are found to be out of 
adjustment, an assistant should be called to fix them. 

In weighing out material for an experiment the empty con- 



LABORATORY PROCESSES 7 

tainer should be placed on the left-hand platform and a brass 
weight which is presumably too heavy to counterpoise it should 
be placed on the right-hand platform. When this weight has been 
tried and found too great, it should be removed and the next 
smaller one substituted. If this is not heavy enough the next 
smaller one should be added to it. If this second weight is found 
to be too great, the next smaller should be substituted for it ; or if 
it is found insufficient, the next smaller should be placed on the 
platform in addition. Thus all the weights in turn should be tried 
and rejected or keT)t as the case demands down to and including 
the 5-gram weight. The rider on the front of the scales should next 
be pushed to the right until the scales just balance. The weights 
on the platform should then be counted up, together with the 
weight indicated by the position to the rider and the number 
of grams immediately recorded in the notebook. The weights 
should then be returned to their proper holes in the weight box, 
beginning with the largest weight; and the number of grams 
again counted as they are removed from the platform. By this 
second counting errors are often discovered. 

The material to be weighed is now placed in the container 
and the weight of it and the container together should be deter- 
mined and recorded in the same way that the weight of the con- 
tainer alone has just been found. 

Sometimes objects, such as a piece of wire, which do not re- 
quire a container, can be weighed directly by placing a disc 
of filter paper on each platform, testing the scales, putting the 
material on the paper on the left and counterbalancing with 
weights on the right, as just described. 

Never place any substance directly on the platform without 
a container or paper to protect it. 

(h) The beam balances are delicate and should be handled 
with care. When not in use the beam is supported and held at 
rest by a mechanism that is controlled by a thumb screw in the 
front of the balance case. Free the beam by turning the screw, 
raise the door of the case, and set the balance in motion by 
waving your hand near one of the pans. The disturbance of 
air should be enough 'to set the beam swinging. Draw down 
the door and observe the swing of the pointer. When the 



8 LABORATORY EXERCISES 

balance is perfectly adjusted, the pointer should swing through 
an equal number of divisions on each side of the middle of the 
scale. The balance can be used satisfactorily, however, even 
when not in perfect adjustment, by taking as the ''zero" of the 
scale that divison mark which is midway between the extremes 
of the swing. For example, if the pointer swings through two 
divisions on the left of the middle of the scale and through 
four divisions on the right of the middle of the scale, the "zero" 
for this particular occasion is the first mark to the right of 
the middle. In other words, in making a weighing, the balance 
is in equilibrium if the pointer swings in the same way it did 
when nothing was on the pans. "When the swing of the pointer 
on one side of the middle exceeds the swing on the other side 
by more than four divisions, it is best to call an assistant to 
make a readjustment. 

Always test a balance before using it. After testing turn 
the thumb screw so that the beam is again supported. Now 
the object to be weighed may be placed on the left-hand pan and 
a weight, presumably too heavy, on the other pan. Never place 
on the pans or remove from them objects or weights while the 
beam of the balance is free to swing. If the beam is not sup- 
ported the sudden addition or subtraction of weight will make 
it lurch and may completely change the adjustment. The 
weights must always be handled with pincers, as the perspira- 
tion on your fingers will tarnish them. 

After the object and a weight have been placed on the pans, 
as just directed, slowly turn the thumb screw and free the beam. 
Observe whether the weight chosen is too great or too small, 
again support the beam and make the proper substitution or ad- 
dition of weights in the same way as directed in the use of 
the platform scales. Continue the systematic trial of the weights 
down to the 1-centigram weight. If a certain number of centi- 
grams appears to be insufficient and one more to be too much, 
choose the one which makes the balance nearer to equilibrium as 
judged by the swing of the pointer. Count the weights as they 
lie in the pan and immediately make a record in your notebook. 
The number of centigrams must be recorded even when it is 



LABORATORY PROCESSES 9 

zero ; otherwise it will appear that you did not weigh the objects 
to centigrams. 

Never attempt to weigh a warm object. The upward current 
of warm air arising from it will make it appear lighter than it 
really is. Never place any material directly on the scale pan. 
Never add anything to a container or take anything out of it 
while it rests on the pan. 



PHYSICAL AND CHEMICAL CHANGES 

Experiment 1. Classification of Changes. 

Reference : 5-11. 

Discussion : A substance is recognized by its characteristic 
properties, such as its color, odor, melting temperature, and be- 
havior when brought in contact with certain other substances, 
as distinguished from its variable properties, such as size and 
shape. When a substance is so treated that its characteristic 
properties disappear and a different set of properties appears, 
the change is summarized by giving to the substance a new name. 
Such a change is called a chemical change. On the other hand, 
when a substance is so treated that only its variable properties 
are affected, the change is termed physical. 

Materials : Magnesium ribbon, paper, asbestos, iodine, sugar, 
copper wire, iron nail. 

Directions: {a) Hold a piece of magnesium ribbon in the 
flame of your burner by means of a pair of pincers. (1) Have 
the characteristic properties of the substance changed? (2) 
Has a new substance been formed? (3) How should this 
change be classified? Repeat the experiment using a piece of 
platinum wire instead of the magnesium ribbon. (4) (5) (6) 
Answer the same questions as in (1) (2) (3). (7) Tabulate 
the answers to these questions and to similar questions for the 
following experiments by heading the columns of the table re- 
spectively — Substance, Cause of change. Effect on character- 
istic properties. Classification — and writing in the list of sub- 
stances — Magnesium, Platinum, Paper, Asbestos, Iodine, Su- 
gar, Copper, Blue solution. 

(h) Drop a crystal of iodine into a dry test tube, and heat 
gently the bottom of the tube until a violet vapor arises and 
a black deposit appears in the upper part of the tube. Ob- 
serve the crystals of which this black deposit is composed and 
then heat them by holding the upper part of the test tube in 

10 



12 LABORATORY EXERCISES 

the flame. (1) Are the crystals iodine? (2) Give reasons for 
your conclusion. (3) Are the size and shape of the crystals the 
same as those of the original iodine? (4) What class of prop- 
erties underwent change in the transformation of the iodine to 
the black deposit? Record your results in the table referred 
to above. 

(c) Put enough sugar into a test tube to make a layer about 
0.5 cm. in thickness and hold the tube in the flame of a burner 
until no further change takes place. When the tube has cooled, 
break it in your mortar and examine what has been formed. 
(1) Whal is its color? (2) Is it sweet? (3) Will it dissolve 
in water? (4) Is it sugar? 

(d) Int© a test tube one-fourth full of dilute nitric acid slip 
a piece of copper wire, warm the liquid gently and allow the 
action to proceed for 5 or 10 minutes. (1) Has the wire dimin- 
ished in size? (2) What is the color of the substance into 
which the copper has been transformed by the action of the 
acid? Dilute the contents of the test tube by adding an equal 
volume of water and cool the mixture by holding the tube 
under the tap. Carefully slip a bright iron nail into the liquid 
while holding the tube inclined so that the nail will not plunge 
down and break the bottom of the tube. After 5 minutes re- 
move the nail. (3) With what does the nail appear to be 
coated? (4) How many changes has the copper undergone? 
(5) Is the reappearance of a substance after a series of treat- 
ments a reason for classifying the changes involved as phys- 
ical? 



MIXTURES AND PURE SUBSTANCES— COMPOUNDS 
AND ELEMENTS 

Experiment 2. Mixtures and Pure Substances. 

Reference : 15-16. 

Discussion: Some substances are perfectly homogeneous, 
that is all particles of the substance have identical character- 
istic properties. In the case of other substances, different par- 
ticles have different characteristic properties ; these are mixtures. 
Often the particles of a mixture are so minute and so evenly 
distributed throughout the whole mass that a substance appears 
to be homogeneous when in reality it is not. In such cases prop- 
erties other than appearance must be depended upon to dis- 
tinguish a pure substance from a mixture. Differences in den- 
sity, in solubility, in action upon heating, and in behavior in 
many other tests may prove an apparently pure substance to 
be in reality a mixture. 

Materials : Salt, ' ' Old Dutch Cleanser, ' ' three numbered un- 
known substances. 

Directions: (a) Examine carefully a pinch of salt. (1) 
Do all particles appear to have the same characteristic prop- 
erties? Add a small pinch of salt to half a test tube of water, 
holding the tube quietly without shaking it. (2) Do all par- 
ticles appear of the same density as judged by their sinking 
in or floating upon the water? Shake the test tube vigorously. 
(3) Does the test of solubility show the salt to be made up of 
different or of similar particles? (4) Is salt a pure substance 
or a mixture? (5) State the grounds for your conclusion. 

(h) Examine some ''Old Dutch Cleanser.'' (1) Does it ap- 
pear to be homogeneous? Put a spatulaful of it into a test 
tube, fill the test tube half full of water and shake it. (2) 
Does it all dissolve? (3) Does it all settle to the bottom of the 
tube? Pour the material into a filter prepared as directed in 
paragraph 2 of the Laboratory Processes, page 4, and placed 

14 



16 LABORATORY EXERCISES 

in a clean test tube. When the liquid has run through the 
paper, shake the filtrate vigorously. (4) What is formed? Rub 
a few drops of the liquid between your fingers. (5) What 
does the substance in the water appear to be? Remove the 
filter paper from the funnel and spread it out flat. Feel of the 
material collected on it. (6) Does it feel like the substance in 
the filtrate ? Scrape some of the material from the paper with 
your spatula, put it in a test tube, add some water and shake 
it. (7) Is the result the same as with the filtrate? (8) Is ''Old 
Dutch Cleanser" a pure substance or a mixture? (9) What is 
the evidence? 

(c) Determine by trying various tests whether the materials 
marked ''No. 1," ''No. 2," and "No. 3" are mixtures or pure 
substances. (1) (2) (3) State in each case what you did and 
what you concluded. 

Experiment 3. Formation of a Compound. 

Reference: 16-18. 

Discussions : The ingredients of some mixtures may be made 
to enter into chemical action with each other and produce a new 
substance, the properties of the substances in the mixture dis- 
appearing and a single pure substance with different properties 
taking the place of the mixture. Substances thus formed are 
called compound substances or compounds. If all of both con- 
stituents of the mixture are to be made to disappear and only 
a pure compound substance is to be formed, the constituents 
of the mixture must be in a certain definite proportion to each 
other. Otherwise all of the substance in the mixture which was 
present in too large an amount will not be used up and a mixture 
of this remainder and of the compound will result. 

Materials: Powdered iron, powdered sulphur, carbon di- 
sulphide. 

Directions: (a) Mix one spatulaful each of powdered sul- 
phur and of fine iron filings by grinding them thoroughly in a 
mortar. Determine whether a mixture or a pure substance has 
been obtained by dividing half of the contents of the mortar 
(the other half should be saved for (6)) into three portions in 
three dry test tubes. 



18 LABORATORY EXERCISES 

Fill one of the test tubes half full of water, shake it, and 
allow the solid material to settle. (1) In what respect, if any, 
do different particles of the material differ from one another? 
(2) Is the material a mixture or a pure substance? 

To the second test tube add about 10 c.c. of dilute hydro- 
chloric acid and warm it until a gas begins to be given off. 
(Smell the gas in order to compare the odor with that of a 
gas given off in a later experiment under (b)). (3) Describe 
any observation of the action of the acid on the material which 
will aid you to decide by this means whether the material is 
a mixture or a pure substance. 

To the third test tube add about 5 c.c. of carbon disulphide 
and shake it. Precaution : Carbon disidphide is very inflam- 
mcible. It must never he brought near a flame. All experiments 
involving its use must he done under a hood where there is 
no burner. Any carbon disulphide left in a test tube or other 
apparatus after performing an experiment must not he thrown 
into a sink or crock hut emptied into a bottle provided for the 
purpose and kept in a hood. Filter the contents of the test tube, 
collect the filtrate in an evaporating dish, and allow the carbon 
disulphide to evaporate spontaneously. Examine what is left 
in the dish. (4) What is the substance? (5) What is left on 
the filter paper? (6) Does the action of carbon disulphide 
on the material show it to be a mixture or a pure substance? 
(7) How does it show this? 

(&) Place the remaining half of the material in the mortar 
in a dry test tube and hold the lower end of the tube in the 
flame until it begins to glow. At once remove the tube from 
the flame and observe the progress of the action through the 
material. (1) What is the evidence that action is taking place? 
(2) What kind of a change does this phenomenon indicate? 
When the action has ceased, again heat the lower end of the 
tube strongly for 2 minutes. This is done to drive into the upper 
part of the tube any sulphur which may not have acted. When 
the tube has cooled, wrap it in cloth and break it by a sharp 
blow with a pestle. Grind the black mass which was formed 
in the bottom of the tube in a mortar and place a third of it 
in each of three dry test tubes. Perform with this material 



20 LABORATORY EXERCISES 

three experiments similar to those which you did in (a) with 
water, with hydrochloric acid, and with carbon disulphide. 
(3) What does the action of the water and of the acid indi- 
cate as to whether the material is a pure substance or a 
mixture? (4) How does it show this? (5) Does the odor of 
the gas given off upon treatment with acid show that the 
material is a new substance? (6) How? (7) What does the 
action of carbon disulphide prove about the material? (8) 
What is your final conclusion as to the purity of the material? 
(9) As to the class of substances to which it belongs? (10) 
What kind of a change is involved when a mixture is changed 
to a compound? (11) What evidence was there in this experi- 
ment of such a change taking place? 

Experiment 4. Decomposition of a Compound. 

Kefercnce : 19-20. 

Discussion : The reverse of the process illustrated in Expeii- 
ment 3 can be brought about, that is, certain pure substances 
may undergo chemical changes by which a mixture of two or 
more substances are produced. Some pure substances, however, 
can not be decomposed. These are called elementary substances. 

Materials: Oxide of mercury, a small splinter (a match 
stick is not satisfactory, as the wood has usually been treated 
chemically). 

Directions: Fill an ignition tube about one-quarter full of 
oxide of mercury. Grasping the tube with the tongs in the left 
hand, hold the lower end of the tube in the flame, keeping the 
tube inclined. When a deposit has begun to appear in the upper 
part of the tube ignite a splinter, held in the right hand, by 
thrusting it into the flame, extinguish the blaze by blowing on 
it, and immediately thrust the glowing end into the mouth of the 
tube. A marked action should be produced upon the glowing 
splinter. If nothing happens, repeat the test with the splinter, 
after the oxide of mercury has been heated for a couple of 
minutes more. (1) What happened to the glowing splinter? 
(2) What is the evidence that some gas other than air was in 
the tube while it was being heated? (The name of this gas is 
oxygen.) (3) What familiar substance do you find deposited 



22 LABORATORY EXERCISES 

in the upper part of the tube? (4) Did you heat the tube long 
enough to decompose all the oxide of mercury? (5) Of what 
substances have you found oxide of mercury to be composed? 
(6) Name one characteristic property in which oxide of mer- 
cury differs from mercury and one in which it differs from 
oxygen. (7) Are the properties of oxide of mercury the sum 
of the properties of mercury and oxygen? (8) Is oxide of 
mercury a pure substance or a mixture? (9) Is it a compound 
substance or an elementary substance? (10) How does the 
experiment verify your answer to (9) ? 



OXYGEN 

Experiment 5. Preparation of Oxygen. 

Reference : 23-29. 

Discussion: Oxygen can most conveniently be prepared in 
small amounts, suitable for studying its properties, by the de- 
composition of some of its compounds. The temperature neces- 
sary to produce the decomposition and the expense of the ma- 
terials are of prime importance in making a choice. Oxide of 
m.ercury, for example, which was shown in Experiment 4 to 
be easily decomposed, so far as temperature is concerned, is too 
expensive for general use. Sodium nitrate, on the other hand, 
which is inexpensive, requires a very high temperature for de- 
composition. Potassium chlorate, however, can be decomposed 
at a lower temperature than sodium nitrate and is at the same 
time much less expensive than oxide of mercury. Potassium 
chlorate, therefore, is usually selected for the preparation of 
oxygen in the laboratory. Moreover the temperature at which 
oxygen can be evolved rapidly from potassium chlorate is con- 
siderably lowered by the addition of manganese dioxide, which 
acts as a ^'catalytic agent." 

Materials: Barium dioxide, sodium peroxide, potassium 
chlorate, manganese dioxide, splinters. 

Directions: (a) Sources. (1) In what experiment did you 
obtain oxygen from the oxide of mercury? (2) By what 
means? 

Heat some barium dioxide in an ignition tube and test 
with a glowing splinter, as you did in the experiment with 
oxide of mercury. The residue left in the tube is barium oxide. 
(3) What is given off upon heating? (4) How did you recog- 
nize it? 

Place in a small, dry test tube enough sodium peroxide to 
cover the bottom of the tube completely and add about 5 c.c. 

24 



26 LABORATORY EXERCISES 

of water. Test for oxygen at the mouth of the tube. (5) What 
test did you use? (6) What was the outcome of the test? 

Put enough potassium chlorate in a 6-inch test tube to 
make a tayer about 1 cm. deep. Heat this in the flame and 
test with a glowing splinter. It is useless to attempt a test 
until you can see that gas is being formed. (1) What change 
in appearance occurs when the potassium chlorate is first 
heated? (2) How can you tell when a gas begins to be 
formed? (3) What happened finally to the glowing splinter? 
(4) What does this action prove as to the identity of the gas 
evolved ? 

( h ) Catalysis. Into another test tube put as much potassium 
chlorate as you have just used, and have ready at hand on a 
spatula some manganese dioxide. Heat the potassium chlorate 
slowly and as soon as it has melted, hold a glowing splinter at 
the mouth of the tube. Eemove the tube from the flame, im- 
mediately add the manganese dioxide and test again with a 
glowing splinter. (1) Was the temperature of the pure melted 
potassium chlorate high enough to decompose it? (2) How do 
you know? (3) What was given off upon adding the man- 
ganese dioxide? (4) What effect does the presence of man- 
ganese dioxide have upon the decomposition temperature of 
potassium chlorate? Careful experiments have proved that no 
manganese dioxide is used up in this experiment. (5) What 
technical name is used for a substance which acts as the man- 
ganese dioxide does in this case? 

(c) Preparation in Quantity. Set up an apparatus as shown 
in Fig. 1, page 28, and so placed that the outlet of the trough 
will be over the sink. 

A is an 8-inch test tube provided with a one-hole rubber 
stopper through which passes a piece of glass tubing B which 
in turn is attached to the longer piece of glass tubing D by 
means of the rubber connector C. J5/ is a pneumatic trough 
upon the shelf of which rests a 250 c.c. bottle F. 

Weigh out, as directed in the general directions on page 
6, on pieces of filter paper, 10 grams of potassium chlorate 
and 5 grams of manganese dioxide. (1) Why is each of these 
particular substances used? Mix them thoroughly in a beaker 



28 



LABORATORY EXERCISES 



and then pour them into the tube A, which should be removed 
from the ring stand for the purpose. Holding the tube in a 
horizontal position, shake and roll it until the contents lie evenly 
distributed along two thirds of its length, and again place it 
in position. Be careful not to crush the tube by screwing the 




Fig. 1. 



clamp too tight, and be sure that the stopper, tubes, and con- 
nectors all fit so that there are no leaks. 

Pour enough water into the trough to cover the shelf, over 
the hole in which is to be placed an inverted bottle full of water. 
Fill a 250 c.c. bottle brimful of water at the tap and slide a 
glass plate over the mouth of the bottle so that no bubble of 
air is included. Invert the bottle in the water of the trough, 
remove the glass plate, and lift the bottle to its position on the 
shelf of the trough. Fill three other bottles in a similar way, but 
leave them on the bottom of the trough ready to be substituted 
for the one on the shelf, when that has been filled with gas. 
Adjust the end of the delivery tube under the hole in the shelf. 
Heat the tube gently by holding the burner in your hand and 
moving the flame back and forth under that part of the tube 
which contains the mixture. (The burner should not be held 
directly under the tube, but to one side at an angle, so that if the 
tube should break it would not fall upon your hand and burn it.) 
Bubbles of gas should immediately begin to issue from the 
delivery tube. (2) What is this gas that appears even at the 



30 LABORATORY EXERCISES 

first warming of the tube? (3) Why does it issue from the 
tube? (4) If no gas appears, what is wrong with your ap- 
paratus? Continue heating the tube until gas is evolved and 
rises into the bottle at the rate of one or two bubbles per 
second. When the evolution becomes more rapid than this, 
remove the flame until the gas comes over more slowly 
again. 

When the first bottle is full of gas, slip a glass plate under 
the mouth of the bottle while it is still under water and then 
remove it, keeping it covered with a glass plate or a wet filter 
paper. Substitute for it on the shelf one of the bottles from 
the bottom of the trough. In the same way fill all four. bottles 
with gas. 

Whenever you stop heating the tube so long that gas ceases 
to be evolved, disconnect the delivery tube at C. If this is not 
done water will rise in the tube into the test tube, thus ruining 
the mixture and often cracking the tube. (5) AVhy will the 
water run back into the test tube? (6) What is the gas you 
have collected? (7) From the decomposition of what com- 
pound did it come? (8) Why Avas the manganese dioxide 
added? 

Experiment 6. Properties of Oxygen. 

'Reference : 31-33. 

Discussion : Of the many properties which are associated 
with each substance, only those, as a rule, are studied which 
are important in showing its characteristic chemical or physical 
behavior. For example in the case of oxygen, the fact that it 
will completely fill any space in which it is confined is not 
emphasized, since this fact is true of all gases and is implied in 
the statement that oxygen is a gas. On the other hand the 
action of oxygen in the presence of heated sulphur is different 
from tliat of other gases and is therefore chosen as worth 
studying. 

Some particular and striking property, the presence or ab- 
sence of which may be easily proved, is selected as a ''test'' 
or means of identification of the substance among all other sub- 
stances commonly met. 



32 LABORATORY EXERCISES 

Materials: Splinter, sulphur, red phosphorus, steel wool, 
deflagrating spoon. 

Directions: Referring to the evidence from the last ex- 
periment, (1) state whether or not oxygen is freely soluble 
in water, and (2) the reason for your conclusion. 

Thrust a splinter of wood into one of the bottles of oxygen. 
(3) Does any action take place? 

Thrust a glowing (not blazing) splinter into the bottle. 
Remove the splinter and repeat the action several times. (4) 
What occurred? (5) What condition is necessary for this 
action to take place? (6) Did the oxygen burn? 

Heat in the flame of the burner a small piece of sulphur 
in a deflagrating spoon, and when it has begun to burn with 
a small blue flame, lower it into a bottle of oxygen. (7) What 
effect has oxygen on burning sulphur? (8) What is formed? 

Repeat substituting a little red phosphorus for the sulphur. 

(9) What is formed? 

With the tongs hold a piece of steel wool in the flame 
and when it is glowing quickly thrust it into a bottle of oxygen. 

(10) What is formed? (11) What properties of oxygen do 
the foregoing experiments illustrate? (12) What is a test for 
oxygen? 

Experiment 7. Kindling Temperature. 

Reference : 33. 

Discussion : It Avas evident from Experiment 6 that the sub- 
stances which were burned in oxygen liad to be heated before 
the combustion started, and that a higher temperature was 
necessary in the case of some substances, like the steel wool, 
than in the case of others, like the red phosphorus. All chem- 
ical changes take place more rapidly as the temperature is 
raised. In the case of reactions which evolve heat, either the 
rate of action may be so slow that the heat is lost to the 
surroundings as fast as it is produced and thus causes no rise 
in temperature and speeding up of the action, or the rate may 
be so rapid that more heat is evolved than is lost, which will 
cause a rise in temperature and consequent increase in rate 
of action and again the production of still more heat. When 



34 LABORATORY EXERCISES 

a combustible substance has been heated from an outside 
source, as for example by a burner, until the rate of action 
is so great that the heat evolved from the action is sufficient 
to maintain the temperature after the burner is removed, then 
the kindling temperature has been attained. The substance 
has been kindled and the burning will furnish enough of its 
own heat to continue the action. Whether or not a given 
substance always has the same kindling temperature is deter- 
mined in this experiment. 

Materials : Splinter of wood, steel wool, potassium chlorate.- 

Directions: (a) Hold the end of a splinter in the flame of 
a burner until it blazes. Withdraw the splinter and hold the 
burning end downwards, then upwards. Finally Avave the 
splinter through the air rapidly till it stops burning. (1) What 
is the technical term used to denote the temperature to which 
the splinter was heated when it began to burn? (2) Explain 
in a statement employing this term why the splinter burned 
more rapidly in one position than in the other, (3) and why 
the flame was extinguished when the splinter was waved 
through the air. 

(&) Hold in the flame, first a mass of steel wool with the 
tongs and then the end of a steel file. (1) Which one became 
heated to its kindling temperature? (2) Has steel a definite 
kindling temperature? (3) Upon what physical condition does 
the kindling temperature of a substance depend? 

(c) Heat a little potassium chlorate in a small test tube till 
you can get a distinct test for oxygen, and then withdraw 
the tube from the flame. (1) Does oxygen continue to be given 
off? (2) Does the evolution of oxygen produce heat or does 
the action appear to require the constant supply of heat? (3) 
How does the cessation of the action here differ in cause from 
the cessation of the burning of the splinter in (a) ? (4) Does 
the decomposition of potassium chlorate have a kindling tem- 
perature? (5) What kind of reactions have a kindling tem- 
perature? (6) What kind do not? 



HYDROGEN 

Experiment 8. Preparation of Hydrogen by Various Means. 

Reference : 41-48. 

Discussion : Hydrogen is prepared from its compounds, but 
the result is not brought about as in the case of oxygen 
by the decomposition of the compounds by heat. The 
hydrogen compound is treated with a metal which will unite 
with the other element or elements of the compound, and thus 
set free the hydrogen. The readiness with Avhich these actions 
take place depends upon the particular compounds and metals 
chosen and upon the temperature and the presence or absence 
of catalytic agents. Water, acids, and bases are compounds 
of hydrogen available for use ; and many different metals, but 
not all, may be employed. Certain metals which will not them- 
selves serve to liberate hydrogen from one of its compounds 
will act as catalytic agents when they are in contact with a 
metal which is being used for this purpose. 

Materials: Sodium, strips of zinc, iron nail, iron filings, 
aluminium, dilute acetic acid, copper sulphate solution. 

Directions: (a) Fill a small test tube with water and invert 
it in a pneumatic trough nearly full of water. Obtain from an 
instructor a small piece of sodium in a dry evaporating dish. 
Caution : Never handle sodium with your fingers, as enough 
moisture is present on them to cause action. Always use pincers. 
With the left hand raise the inverted test tube till its mouth 
is about an inch below the surface of the water in the trough. 
With the right hand quickly insert under the mouth of the 
test tube the piece of sodium held in the pincers. When the 
action of the sodium and the water has ceased, place your 
thumb over the mouth of the tube, carry the tube to the burner 
and uncover it in an inclined position with the mouth down- 
ward and close to the flame. (1) What occurred? (2) What 

36 



38 



LABORATORY EXERCISES 



was the gas collected? (3) What was the source, the water 
or the sodium? (4) How can you recognize hydrogen? 

(h) Slip a strip of zinc into a test tube half full of water. 
(1) Is hydrogen evolved? Repeat the experiment using an 
iron nail instead of zinc. (2) Is hydrogen evolved? (3) What 
results would have been obtained at much higher temperatures? 

(c) Set up an apparatus as shown in Fig. 2. A is an 8-inch 
test tube provided with a two-hole rubber stopper, B a thistle 
tube, C and D bent glass tubes, E a rubber connector, G a 250 
c.c. bottle filled with water, H a pneumatic trough. 

Slip into the test tube two strips of zinc, replace th^ stopper 




and pour down the thistle tube enough dilute hydrochloric 
acid to fill the test tube a quarter full. When the bottle has 
become a third full of gas, substitute in its place a small 
inverted test tube filled with water. When the test tube has 
become filled with gas, carry it to the fiame, with your thumb 
over its mouth, and test for hydrogen. (1) What was the gas 
first collected in the bottle? (2) Where did it come from? (3) 
What was the gas in the test tube? (4) Was it produced from 
the hydrochloric acid or from the zinc? (5) Of what element 
is hydrochloric acid a compound? 

Empty the large test tube and repeat the experiment, first 
substituting dilute sulphuric acid for the hydrochloric acid 
and then again perform the experiment using acetic acid and 



40 LABORATORY EXERCISES 

finally using dilute nitric acid. (6) In which eases did you 
find hydrogen to be evolved? (7) Which acid or acids are 
best adapted for preparing hydrogen under these conditions? 

(cZ) Using the same apparatus as in (c), perform experi- 
ments which will enable you to answer the following questions. 
(1) Is hydrogen evolved by the action of hydrochloric acid 
on iron (use iron filings) ? (2) on aluminium? (3) on copper? 
(4) Which metal or metals are best adapted for preparing 
hydrogen? 

(e) Into each of two test tubes standing in the test tube 
rack put two strips of zinc of the same size and just cover 
them with dilute sulphuric acid. To one add five drops of 
copper sulphate solution. The black deposit which appears 
on the zinc is copper from the copper sulphate. It is black 
because it is so finely divided. (1) In which tube is hydrogen 
evolved more rapidly? (2) Is the copper acting on the acid? 
(3) How do you know? (4) What technical term is used to 
express the action of the copper? 

Experiment 9. Preparation of Hydrogen in Quantity. 

Kef erence : 45. 

Discussion : As has l)een discovered in Experiment 8, dif- 
ferent metals and acids may be used in preparing hydrogen. 
Relative expense and the ease of procuring fairly pure sub- 
stances may therefore determine the selection. Sulphuric acid 
is cheaper than hydrochloric acid. Although iron is cheaper 
than zinc, commercial samples usually contain considerable 
iron carbide, which forms with acids other gases than hydro- 
gen. Sheet zinc contains a slight impurity of carbon, but this 
fact is an advantage, rather than the reverse, because the 
carbon acts as a catalytic agent, in the same way that copper 
did in Experiment 8 (e). The size and form of the apparatus 
will depend upon the amount of hydrogen to be prepared. In 
general a small generator is better than a large one, since it 
will contain less air, which must be driven out before pure 
hydrogen can be collected. 

Materials: Strips of zinc, copper sulphate solution. 

Directions: Caution: Keep all flames at least five feet 



42 LABORATORY EXERCISES 

away from a Ifiydrogen generatoVy as mixtures of hydrogen and 
air are very explosive. 

Set up an apparatus like the one used in Experiment 8 (c). 
Fill three more bottles with water and invert them on the 
bottom of the trough. Put in the test tube 4 strips of zinc, 
about 5 grams, and insert the stopper. If the presence of the 
zinc prevents the stem of the thistle tube from entering the 
test tube as it should, unclamp the test tube, disconnect the 
delivery tube, hold the test tube horizontal, and by gently tap- 
ping it make room for the stem and insert the stopper. Then 
restore the apparatus to its original position. Pour into your 
graduate 2 c.c. of copper sulphate solution, fill it with dilute 
sulphuric acid to the 20 c.c. mark and pour the contents of 
the graduate into the test tube through the thistle tube. Upon 
pouring in the liquid, gas should be forced over into the 
bottle. If the gas does not appear, what is wrong? If this is 
the case, locate the trouble and remedy it at once. 

Gas should begin to come over regularly into the bottle. 
When the bottle is a third full of gas, replace it with another 
bottle from the bottom of the trough, refill the first bottle 
with water and invert it with the others in the bottom of the 
trough. (1) What is the gas collected in the first bottle and 
rejected? After the action has proceeded for a few minutes, 
the gas will begin to come over faster. Feel of the outside 
of the test tube. (2) Why does the action increase in speed? 
If the action becomes so vigorous that froth rises in the test 
tube, wet the outside of the tube with cold water. If this is 
not effective, pour a little cold water down the thistle tube. 
(3) Why should this treatment reduce the speed of the action? 
When the bottle on the shelf of the trough is full, remove it 
as you did the bottles of oxygen and set it on the desk still 
covered and still inverted. Immediately substitute for it on 
the shelf one of the bottles from the bottom of the trough. 
Continue collecting the hydrogen till all of the bottles have 
been filled. After a time the hydrogen will come over more 
slowly. (4) To what is this due? When the rate is less than 
two bubbles per second, measure into your graduate 5 c.c. of 
concentrated sulphuric acid and pour 1 c.c. of this down the 



44 LABORATORY EXERCISES 

thistle tube. When the rate of action again diminishes, pour 
in again 1 c.c. of acid from the graduate. Do not add more 
than 5 c.c. of the concentrated acid altogether, as there is 
danger that it will collect in the bottom of the test tube and 
unexpectedly becoming mixed with the water will cause a 
violent action. Leave the apparatus and at once study the 
properties of the hydrogen as directed in Experiment 10. 

Experiment 10. Properties of Hydrogen. 

Reference : 49-54. 

Discussion : When a mass of gas is released into another 
gas, as into the air, it will rise or sink as a whole, depending on 
whether or not it is lighter or heavier than the gas into which it 
passes, just as an object thrown into water will float or sink 
according to its density compared to water. In addition to this 
movement, however, each particle of a gas has individual inde- 
pendent motion, which causes the gas to move against the 
force of gravity and to become equally distributed throughout 
the container which holds it. This property is called diffusion. 
These facts are true of all gases but they are most easily demon- 
strated in the case of hydrogen. 

In testing for hydrogen, it will be noticed that the loudness 
of the explosion varies greatly on different occasions. The 
noise of an explosion is due to the vibrations of the air caused 
by the sudden production of gas or by the sudden expansion 
of air by heat. When hydrogen and air have become thoroughly 
mixed before ignition in such proportions that the oxygen 
present is just the proper amount to unite with all of the 
hydrogen, conditions are right for simultaneous action through- 
out the whole mass, with the instantaneous production of much 
heat and a consequent loud report. If, however, the amount 
of either the hydrogen or the air present in the mixture is too 
great, the amount of the other gas must be correspondingly 
less, and less action and less noise will result on igniting the 
mixture. If the gases are prevented from mixing before they 
are ignited, action can take place only where they meet in 
the process of diffusing into each other, and slow union with 
the phenomenon of a steady flame rather than explosion will 
result. 



46 LABORATORY EXERCISES 

Materials : Splinters. 

Directions : Stand one of the bottles of hydrogen, collected 
in Experiment 9, upright, remove the cover and instantly drop 
in a lighted match. (1) What happened? Eepeat the experi- 
ment but wait fully half a minute before dropping in the match. 
(2) What happened? (3) Why? (4) What property of hydro- 
gen does this experiment illustrate? (5) Why were the bot- 
tles of hydrogen prepared in Experiment 9 kept inverted ? 

Place a covered bottle of hydrogen mouth downward over 
the mouth of a bottle of air and remove the glass plate between 
them. After they have stood for fully five minutes, test for 
the presence of hydrogen in each bottle. (6) What was the 
result? (7) What property of hydrogen is shown by this ex- 
periment? (8) Does air possess this property? (9) Is it a 
property of all gases? 

Thrust a lighted splinter up into an inverted bottle of 
hydrogen. Slowly withdraw the splinter so that it will relight 
at the mouth of the bottle. Look for a deposit on the inside 
of the bottle and feel of the neck of the bottle. (10) What 
took place in the bottle? (11) What was the evidence? (12) 
Why did no loud explosion, like those in the preceding experi- 
ments, occur? (13) What is formed when hydrogen burns in 
the air? (14) Does hydrogen '* support combustion"? (15) 
What was the evidence? 



QUANTITATIVE STUDIES OF CHEMICAL REACTIONS 

Experiment 11. Percentage of Oxygen in Potassium 
Chlorate. 

Reference : 61, 72. 

Discussion : One of the most fundamental laws of chemis- 
try is that the proportion by weight of the elements in a given 
compound is always the same. A knowledge of what these 
proportions are is, therefore, of prime importance in the study 
of any substance. In the case of potassium chlorate, the per- 
centage of oxygen can be determined by heating a known 
w^eight of the compound until all the oxygen is driven off and 
then weighing what remains. The difference is the weight of 
the oxygen. In the actual experiment the weights involved 
are small, so that they must be determined as accurately as 
possible and great care must be used to lose none of the solid 
material. As there is no simple means of knowing w^hen all 
of the oxygen has been driven off by heating, the substance 
must be heated, weighed, reheated and again weighed. If there 
is no loss of weight after the second heating, all of the oxygen 
was expelled on the first heating ; if, however, a loss of weight 
appeared after the second heating, it shows that the action 
was not completed the first time, and successive heating and 
weighing must be continued until there is no further loss in 
weight. 

Materials: Potassium chlorate, pin. 

Directions : Read carefully the directions for making weigh- 
ings on pages 7 to 9. 

Weigh carefully a clean, dry porcelain crucible with its 
cover, and (1) record its weight to centigrams in the appro- 
priate place in the tabulation given below. From a bottle of 
purest potassium chlorate, labeled for this experiment, transfer 
to the crucible with a clean, dry spatula from 1.5 to 2.0 grams 

48 



50 LABORATORY EXERCISES 

of the salt. This is about two spatulafuls. Then weigh the 
crucible and contents to centigrams and record the weight. 

Support the covered crucible and contents on a clay tri- 
angle, resting on the ring of the ring-stand and arranged at 
such a height that the tip of the flame of the Bunsen burner 
placed underneath it will be about one inch below the bottom 
of the crucible. Heat the crucible in this position for about 
twenty minutes, then lower the ring about 5n inch and a half, 
so that the flame plays upon the bottom of the crucible, and 
heat it for half an hour or more. The evolution of oxygen 
probably will have spattered some of the potassium chlorate 
on the bottom of the cover. This must be removed and returned 
to the crucible in the following way: When the crucible has 
cooled, place the cover, bottom side up, on a piece of filter 
paper and very carefully loosen the adhering solid with a pin. 
Put this solid back in the crucible together with any particles 
which may have lodged on the paper. Now heat the covered 
crucible for half an hour, allow it to cool for at least half an 
hour, weigh it, and record the weight. Heat it again for 
half an hour, cool for half an hour, and again weigh it. Record 
all weighings. 

If it has lost weight during the last heating, repeat the 
process until two successive weights do not disagree by more 
than one centigram. (2) Calculate the percentage of oxygen 
in potassium chlorate. (3) What is the percentage as cal- 
culated from the formula? (4) How great Avas your error? 

Wt. crucible empty. 

Wt. crucible and KCIO3. 

Wt. KCIO3. 

Wt. crucible and KCIO3 after first heating. 

Wt. crucible and KCIO3 after second heating. 

Wt. crucible and KCIO3 after third heating. 

Loss of weight = weight of oxygen. 

Percentage of oxygen in potassium chlorate. 

Experiment 12. The Determination of the Percentage of 
Chlorine in Silver Chloride. 

Reference : 74. 



52 LABORATORY EXERCISES 

Discussion : When silver chloride is heated in a stream 
of hydrogen, hydrochloric acid and silver are formed. By 
weighing the chloride in the beginning and the silver left 
after the reaction, the percentage of chlorine in the compound 
can be calculated. In the experiment described below, illumi- 
nating gas, which contains hydrogen, is used as a convenient 
source of hydrogen. 

Materials: Silver chloride. 

Directions: Weigh a clean test tube to centigrams and 
(1) record the weight. Place in it about 3 grams of silver 
chloride and weigh again. Support the tube in a horizontal 
position in a clamp placed at about one-fouvth of the length 
of the tube from the open end. Place in the end of a rubber 
tube, connected with the gas cock, a piece of glass tubing about 
twice the length of a test tube. Turn on the gas and light it 
as it comes from the glass tube. Adjust the flame so that 
it is about 1 cm. long. Blow out the gas and insert the tube 
in the test tube so that the end of the former reaches almost 
to the silver chloride. Light the gas at the mouth of the 
test tube. Heat the test tube cautiously, keeping the flame in 
motion under the silver chloride. The test tube should not 
be heated hot enough to give a yellow color to the flame. As 
the hydrochloric acid issues from the tube it imparts a slight 
green color to the blue part of the flame of burning gas. It 
will also produce a cloud Avhen a solution of ammonia is 
brought near the flame. When the reaction is completed these 
tAvo tests for the gas are no longer given. When this occurs, 
turn off the gas, let the tube cool and weigh it. Place the 
silver in the bottle provided for it. What was the weight (2) 
of the silver chloride? (3) of the chlorine? (4) Calculate the 
percentage of chlorine in silver chloride. (5) What is the true 
percentage as calculated from the formula AgCl? 

Experiment 13. Vapor Pressure. 

Reference : 83, 89. 

Discussion : It is a matter of common knoAvledge that water 
left in an open dish will disappear, that is will form a gas 
or vapor which passes off into the atmosphere, will in short 



54 LABORATORY EXERCISES 

evaporate. When the space above the water is enclosed, how- 
ever, the vapor can not escape and as it collects in larger and 
larger amounts more and more returns to the liquid state, that 
is condenses. Eventually the rate of condensation equals the 
rate of evaporation and apparently no more water evaporates ; 
vapor and liquid are said to be in equilibrium. The vapor in 
the enclosed space exerts a pressure on the confining walls, 
which pressure reaches a maximum when equilibrium between 
liquid and vapor, as just explained, has been attained. This 
maximum vapor pressure is what is referred to as the vapor 
pressure of the water. 

Since water evaporates faster when hot than when cold, it 
follows that when equilibrium has been attained in an enclosed 
space condensation also must be taking place faster above hot 
water than above cold. But to produce more rapid condensation 
more vapor must be present, which in turn will cause greater 
pressure on the confining walls. The (maximum) vapor pres- 
sure of water depends, therefore, on the temperature. 

The relations which exist between water and water vapor 
hold equally for the relation between all other liquids and their 
vapors. The values for these vapor pressures are different for 
different liquids, as would be expected from their well-known 
differences in rate of evaporation in the open air. 

Directions: (1) Make a drawing of the apparatus set up 
in the laboratory, which was prepared as follows. Each of the 
tubes, sealed at one end, was completely filled with mercury 
and then inverted in the trough of mercury. Under tube B 
was introduced a few drops of water, which rose to the surface 
of the mercury in the top of the tube. In a similar way ether 
was introduced into tube C. (2) What is in the space above 
the mercury in tube Af (3) What prevents the mercury from 
running out of the tube? (4) For what purpose is this kind 
of an apparatus used? (5) By what name is it known? (6) 
What is the length of the mercury column in tube A? (7) 
What would happen if a very little gas were let into the space 
above the mercury? (8) What is in the space above the mer- 
cury column in tube Bf (9) What is the length of the mercury 
column in tube B? (10) Why is it shorter than that in tube 



56 LABORATORY EXERCISES 

A? (11) What is the temperature as shown by the thermometer 
beside the apparatus? (12) What do you find to be the vapor 
pressure of water at this temperature, as expressed in terms 
of the height of a column of mercury? (13) What is the vapor 
pressure of ether at this temperature, as indicated by tube Cf 
Grasp the tube C with your fingers and thumb opposite 
the top of the column of mercury. (14) In which direction 
does the mercury move? (15) What is the effect of increase 
of temperature on the vapor pressure of a liquid? 

Experiment 14. Boyle's and Charles' Laws. 

Reference: 88-90. 

Discussion : The volume of a gas increases wdth rise in tem- 
perature according to Charles' law, which states that the 
volumes of a given sample of gas at two temperatures are pro- 
portional to these temperatures expressed on the absolute scale. 
The absolute temperature, T, is 273 degrees plus the tempera- 
ture measured on the centigrade scale, r° = 273 + ^°. This 
law applies not only to gases like oxygen, hj^drogen, and air, 
but also to the vapor formed from liquids such as water. In 
the exi)eriment described below, the volume of a sami)le of air 
contained in a bottle over water at room temperature is first 
measured. It is then heated. The temperature and volume 
under the new conditions are recorded. A calculation is then 
made to determine what the volume should be according to 
Boyle's and Charles' laws, and this value is compared with that 
obtained as the result of the experiment. 

The air contained in the bottle both at room temperature 
and at the higher temperature is mixed with water vapor, since 
the air is present in a bottle which also contains water. The 
pressure of water vapor changes with the temperature and this 
factor must be taken into account in the calculations. The 
formula which expresses tlie two laws is as follows: 

pv T 



p'v'~ r 

The value of p is determined by subtracting the pressure 
of water vapor at tlie room temperature, T, from the pressure 



58 



LABORATORY EXERCISES 



of the atmosphere; the value of p' is the pressure of the at- 
mosphere minus the pressure of the water vapor at the higher 
temperature T'. The volume of the gas at room temperature 
is V, and at the higher temperature is v\ In the experiment 
all the factors are noted. In the calculations v' is determined 
by substituting the other values. If the observed value of v' 
agrees with the value calculated from the equation the results 
are a confirmation of the laws expressed by the equation. 

Materials: 150 c.c. bottle, large beaker. 

Directions: Arrange the apparatus as represented in Fig. 
3. Place in the beaker just enough water to cover the bottle 
when it is completely immersed. Remove the bottle, place the 



Fig. 3. 



beaker on a wire gauze and heat the water to 80°-85°. While 
the water is heating, fill the bottle to over-flowing with water at 
room temperature and record this temper^ure in the tabula- 
tion given below. Insert the stopper carrying the glass tube 
and press it down into place. Place the finger over the end 
of the tube at a and remove the stopper. Pour from the bottle 
about 80 c.c. of water into a graduate and note the volume of 
the water, which is the volume of the air in the bottle. 

When the water in the beaker has reached 80° to 85°, fill 
the tube attached to the cork with water at room temperature, 
place the finger over the end a and replace the cork in the 
bottle. Remove the flame from under the beaker, put the 
bottle in the hot water and place it so that the water that 
leaves the tube at a falls into a graduate. Place a thermometer 



60 LABORATORY EXERCISES 

in the hot water and stir occasionally. When water no longer 
falls into the cylinder note the temperature and the barometric 
pressure. Place the graduated cylinder under running water 
to cool it and note the volume of the water it contains. The 
sum of this volume and the original volume of the gas in the 
bottle is the volume of the gas at the higher temperature. 
Record this volume. Record the vapor pressures of water at 
the two temperatures involved in the experiment, as they are 
stated in the table of vapor pressures posted in the laboratory. 

(1) Calculate in the way outlined in the discussion above what 
the volume of the gas at the higher temperature should be. 

(2) By how much does this value differ from that which you 
found experimentally? 

Temperature of room. 

Temperature when water stops flowing. 

Volume of air in the bottle. 

Volume of water collected. 

Volume of air at higher temperature. 

Barometric pressure. 

Vapor pressure of water at room temperature. 

Vapor pressure of water at the higher temperature. 

Experiment 15. The Weight of a Liter of Oxygen. 
Reference : 80, 82, 84-88, 90. 

Discussion : In finding the weight of a specified volume of 
gas, the direct method of weighing the gas in a flask or other 
container is subject to very considerable errors, unless unusual 
precautions are taken. These errors are due to the fact that 
gases are so light that the weight of a gas is very small in 
comparison to the weight of its container. Hence a small error 
in the weight of the container will make an error in the weight 
of the gas very large in proportion to its total weight. To 
avoid these difficulties the weight of the gas is often obtained 
indirectly, as in this experiment, by weighing the generator 
and contents before and after the gas has been evolved. The 
difference, that is the loss of weight, is the weight of the gas 
which has been formed and collected. 

The volume of the f?as is usually measured by observing 



62 LABORATORY EXERCISES 

the volume of water displaced upon collecting the gas. The 
mere volume of a gas has no significance, however, unless the 
existing conditions of temperature and pressure are stated, 
because the volume of the same mass of gas may be greater 
or smaller according as the temperature is higher or lower and 
the pressure upon it is diminished or increased. Hence in this 
experiment, the temperature of the gas and the pressure under 
which it is confined must be determined as well as the volume 
which it occupies. The temperature can be read directly on 
a thermometer and the experiment may be so arranged that 
the pressure of the gas will be the same as that of the outside 
air. The pressure of the air, that is atmospheric pressure, is 
measured on a barometer. When, however, the gas is collected 
over water, the space above the water contains not only the 
gas in question but in addition water vapor, which also exerts 
a pressure. In other words, if the gas appears to be at 
atmospheric pressure, its true pressure is atmospheric pressure 
minus the pressure of the water vapor. 

The weights of gases as usually stated are those under 
standard conditions, 0°C. and 760 mm. pressure, and the neces- 
sary calculation must be made by the application of the laws 
of Boyle and of Charles to convert the results obtained in the 
experiment to those which would obtain under these conditions. 

Materials: Potassium chlorate, manganese dioxide, copper 
wire. 

Directions: Thoroughly heat one spatulaful of manganese 
dioxide in a porcelain dish to drive off all the moisture. (1) 
What error would result if the moisture remained? 

While this is heating, set up an apparatus as shown in 
Fig. 4. A is a 1 or 2 liter bottle fitted with a two-hole rubber 
stopper. B is a glass bend extending nearly to the bottom of 
A. This can be the exit tube used in Experiment 8. C is a right- 
angled bend. Z> is a dry 750 c.c. flask. JE' is a glass jet. F 
is a pinch cock. G is an 8-inch test tube fitted with a one hole 
rubber stopper. ^ is a loop of copper wire. 

Mix in a mortar half a spatulaful, about 1 gram, of the dried 
manganese dioxide and one spatulaful, about 2 grams of potas- 
sium chlorate. Remove the test tube from the apparatus, put 



64 



LABORATORY EXERCISES 



into it the mixture from the mortar, suspend the test tube from 
the hook over the pan of a balance and determine its weight 
to centigrams. Before doing this read carefully the directions 
for weighing on pages 7 to 9. (1) Record all the data in 
tabular form. 

Fill the jet E and connecting tubing completely with water 
by opening the pinch cock F and blowing with the mouth through 
the tube C. At the same time raise the jet and attached rubber 
tubing above the level of F. Replace the test tube in the position 
shown in the figure. 

With the jet E dipped into a beaker half full of water, raise 




Fig. 4. 



the beaker until the level of the water in the beaker is on the 
same level with that in the bottle A, then open the pinch cock 
F for several minutes to allow the pressures to become equal. 
(2) Why is it necessary to equalize the pressures inside and 
outside the bottle? Close the pinch cock and dip the jet into 
a clean, dry, 750 c.c. flask, as shown in the figure, which has 
previously been weighed to grams on the platform scales. 

Now start heating the mixture in the test tube with a low 
flame and at the same time loosen the pinch cock. Water 
should flow out of the jet into the flask. (3) If it does not, what 
is wrong? Stop heating an instant, the water should cease 
flowing. (4) If it continues to flow, what is wrong with the 
arrangement of the apparatus? 



66 LABORATORY EXERCISES 

When you are sure that the apparatus is correctly set up, 
heat the test tube more vigorously, just so that a steady stream 
flows from the bottle to the flask. Care must be taken at this 
point, 7iot to heat the test tube to redness. If the heating is 
interrupted for any reason, be sure to keep the tip of the jet 
below the surface of the water in the flask. (5) Why? (6) 
What is replacing the water in the bottle? 

Continue heating until about 600 c.c. of water has been 
collected in the flask. Then, with the jet still in the water, 
allow the test tube to cool 15 minutes to room temperature 
again. (7) What is happening to the water in the flask? (8) 
What is causing this action? Raise the flask until the level 
of the water in the flask corresponds with that of the water 
in the bottle to equalize the pressures again. Close the pinch 
cock. 

Weigh the flask and water on the platform scales. (9) 
What is the weight of the water? (10) What is the volume 
of the gas collected? Weigh the test tube again to centigrams. 
(11) What does the loss in the weight of the tube represent? 
Record the temperature of the water in the flask and also the 
barometric pressure. 

To calculate the volume of the oxygen formed, subtract 
from the barometric pressure the vapor pressure of water at 
the observed temperature of the water. (See table on page 
542.) (12) Why? By means of Boyle's law and Charles' law 
calculate from the corrected pressure and from the observed 
temperature the volume of oxygen under standard conditions 
and from this value and the weight of the oxygen, the weight 
of a liter of oxygen. (13) Name all sources of error in this 
experiment. 

Experiment 16. Atomic Weight of Zinc or of Aluminium. 

Reference: 73-74. 

Discussion : When chemical reactions take place, the pro- 
portions by weight in which given substances act with each 
other are always the same, as is also the proportion of the 
substances formed compared with the original substances. For 
example, the ratio between the weight of zinc used and the 



68 LABORATORY EXERCISES 

weight of hydrogen evolved when the zinc reacts with acid 
is always the same. Furthermore it has been found that 
definite numbers may be assigned to all of the elements which 
express the proportion by weight in which they react with one 
another. These numbers represent the relative weights of the 
atoms of different elements. 

In order to assign to a given element its significant number, 
or atomic Aveight, it must be experimentally determined what 
weight of the element is involved in chemical action Avith an 
atomic weight of some element of knoAvn value. For example 
if the atomic weight of element **A" is 10, and if experiment 
shows that 10 grams of '*A" react Avith 100 grams of element 
^'B," then 100 is assigned as the atomic Aveight of "B. " In 
this example the assumption is made that an equal number 
of atoms of '*A" and of ''B" are involved. If, hoAvever, it 
is knoAvn that tAvice as many atoms of '*B" react as atoms of 
**A," then the number 100 Avill represent tAvice as many atoms 
as the number 10 and the atomic Aveight of half 100 or 50 
must be assigned to ''B." The experiment about to be per- 
formed is an illustration of this point, for it is knoAvn that 
for every atom of zinc Avhich reacts, tAvo atoms of hydrogen 
are evolved; or for every atom of aluminium, three atoms of 
hydrogen. 

In this experiment the reacting weights of zinc and hydro- 
gen are to be found. The zinc is Aveighed directly. The 
hydrogen is too light to Avcigh accurately Avith the available 
apparatus, so its volume is determined under observed condi- 
tions of temperature and pressure and, after the proper cor- 
rections have been made, its Aveight is calculated from the 
knoAvn Aveight of a liter of hydrogen. 

The atomic Aveight of aluminium may be determined as an 
alternative if desirable. 

Materials: Pure zinc or aluminium; concentrated, 20 per 
cent or 10 per cent hydrochloric acid. 

Directions: Weigh to centigrams a strip of zinc (about 
2.0 grams) or a strip of aluminium (about 0.8 gram) and record 
the weight Avith the data under 1. 

Set up an apparatus as in Experiment 8 for the prepara- 



70 LABORATORY EXERCISES 

tion of hydrogen, but substitute for the gas bottle a 750 c.c. 
tlask filled with water. The flask can be supported by slipping 
the neck of the flask through a ring, after the flask has been 
filled and while it is inverted in the trough, and by clamping 
the ring to a stand. Be sure that the delivery tube is so placed 
that no gas will be lost. The water in the trough should be 
at about room temperature. 

Put the weighed metal into the test tube, replace the stopper, 
and pour through the thistle tube 30 c.c. of hydrochloric acid 
which has been specially prepared for this experiment (20 per 
cent acid for the zinc, or 10 per cent for the aluminium) and 
to which a little copper sulphate has been added. Collect all 
the gas which comes over from the generator. If the evolution 
of the hydrogen becomes very slow before the last of the metal 
disappears, add 5 c.c. more hydrochloric acid, '^concentrated" 
for the zinc or 20 per cent acid for the aluminium. 

When all the metal has reacted, slip a glass plate over 
the mouth of the flask and quickly remove it from the trough, 
turning it right side up and being careful to lose none of the 
water which it still contains. Determine the volume of the 
water which remains in the flask, by measuring it with your 
graduate. Determine the total volume of the flask by filling 
it with water from your graduate. 

(1) Record these volumes and also the temperature of the 
water in the trough and the barometric pressure. (2) What 
was the volume of the gas collected? This volume must be 
corrected by subtracting the volume of the acid added to the 
generator. (3) Why? (4) What is the corrected volume? The 
hydrogen left in the generator at the end of the experiment 
was not collected and measured. (5) Why is no correction 
for this amount necessary? (6) What must be subtracted from 
the barometric pressure in order to get the true pressure of 
the hydrogen? (Refer to the table on page 542 for the amount 
of this correction.) (7) What is the corrected pressure? (8) 
Calculate what the volume of the hydrogen would have been 
under standard conditions. One liter of hydrogen at 0° and 
760 mm. pressure weights 0.090 gram. (9) What was the 
weight of the hydrogen collected? (10) Write in words the 



72 LABORATORY EXERCISES 

proportion which you are going to use in calculating the atomic 
weight of the metal from your data. (11) Why do you use 
twice the atomic weight of hydrogen in the proportion for 
zinc or three times the atomic weight of hydrogen in the pro- 
portion for aluminium? (12) Substitute for the words in the 
above proportion the values you have obtained in this experi- 
ment and make the necessary calculations. (13) What do you 
find to be the atomic weight of zinc or aluminium? (14) What 
is the correct value? (15) How great was your error? (16) 
What is the ratio of your error to the correct value? (17) 
What do you consider the most probable source of error in 
your work? (18) Why? 



WATER 

Experiment 17. Occurrence of Water. 

Reference : 92. 

Discussion : In addition to the obvious occurrence of Avater 
in the liquid form and in the solid form, as ice and snow, water 
is present in many materials which on casual examination seem 
to be dry. In these cases the water is so intimately mixed 
with the other substances present that it must be separated 
from them before it can be recognized. Often pressure alone 
is sufficient to do this, that is Avater can be squeezed , out of 
apparently dry objects; but in other instances the object must 
be heated so as first to vaporize the water, which later con- 
denses, before its presence can be detected. There is danger, 
however, of being misled in these latter cases, for heating 
may have produced water by chemical changes in the sub- 
stances of which the material is composed, instead of merely 
separating water which already existed as a constituent of 
a mixture. 

Water may be present in gases as water vapor, which is 
invisible. (Do not confuse water vapor with fog which is 
made up of minute drops of liquid.) Water vapor may often 
be detected by cooling the gas in which it is present to such 
a temperature that the vapor condenses to liquid and is then 
recognized. 

Materials: Potato, wood, meat, sugar, ice, salt. 

Directions: (a) Cut a piece of potato Avith a knife and 
examine the blade. Repeat the experiment using some dry 
wood instead of the potato. (1) What is the evidence as to 
the presence of water in potato? (2) In the wood? 

Heat a piece of wood gently in the bottom of a test tube, 
but do not let it char. (3) How does this experiment confirm 
or refute ansAver 2? 

74 



76 LABORATORY EXERCISES 

Determine experimentally whether or not water is present 
in meat. (4) What did you do? (5) What do you conclude? 

Heat a spatulaful of sugar in a test tube until it blackens. 
(6) What collects in the upper part of the test tube? (7) Does 
this prove that sugar contains water? (8) Why? 

(?)) Fill a small beaker with cracked ice, pour in water 
till it rises to half the height of the ice in the beaker, sprinkle 
into the beaker 3 spatulafuls of salt. (The presence of the 
salt lowers the temperature of the ice and water.) (1) What 
is found on the outside of the beaker after a few minutes? (2) 
Where did it come from? (3) Why did it appear? 

Experiment 18. Purification of Water. 

Reference: 105-106. 

Discussion : The impurities in water may be present as 
particles of solid, or they may be in solution. To remove solid 
impurities, the water can be filtered, the nature of the filter 
depending on the size of the particles and the amount of water 
to be filtered. Some suspended particles are so small that 
they pass through all ordinary filtering materials, but often 
the addition of another insoluble substance will cause the fine 
particles to coagulate or to adhere to the substance added and 
then the suspended matter will be held back on the filter. 

When the impurity is in solution, filtration can not be used 
and the water must be removed from the impurity by boiling 
the water and collecting the condensed vapor, that is by dis- 
tillation. If the soluble impui^ity also distills upon boiling the 
water, purification becomes very difficult and special means 
must be resorted to for individual cases. 

Materials : Powdered sulphur, clay, alumina cream, copper 
sulphate, litmus paper. 

Directions: (a) Prepare samples of impure water as fol- 
lows : 

Sample 1. Put half a spatulaful of powdered sulphur into 
a test tube half full of water and shake it. 

Sample 2. Grind together in a mortar half a spatulaful of 
clay and 30 c.c. of water. 

Sample 3. Completely dissolve in an 8-inch test tube full of 



78 



LABORATORY EXERCISES 



water enough copper sulphate to give a distinct blue color 
to the solution. 

Sample 4. To 30 c.c. of water in an 8-inch test tube add 
5 c.c. of ammonia. 

(b) Pour some of sample 1 into a funnel fitted with a filter 
paper, collecting the filtrate in a test tube. In this and the 
following filtration the process need be carried only far enough 
to show whether or not the impurity has been removed. (1) 
What has been accomplished in this case by filtration? 

(c) Repeat experiment (h), using some of sample 2 instead 
of sample 1. (1) What is the appearance of the filtrate? (2) 




Fig. 5. 



Why is the result different from that in experiment (6) ? Pour 
10 c.c. of sample 2 into an 8-inch test tube, add 5 c.c. of 
^'alumina cream" (a suspension of aluminium hydroxide) and 
shake it thoroughly. Filter. (3) What is the appearance of 
the filtrate? (4) How did alumina cream accomplish the result? 

(d) Repeat experiment (5) using some of sample 3. (1) 
Was the removal of copper sulphate accomplished? (2) Why? 
(3) From what class of substances can water be purified by 
filtration? (4) From what class is it impossible to do this? 

(e) Set up an apparatus for distillation as shown in Fig. 5. 
A is an 8-inch test tube provided with a one-hole stopper, B 
and C are glass tubes and D a rubber connector, E is a 6-inch 
test tube, and F a 250 c.c. bottle partly filled with cold water. 



80 LABORATORY EXERCISES 

Put 15 c.c. of sample 3 in tube A and boil the solution gently, 
taking care that none of it spatters up through tube B. When 
a few c.c. of liquid appear in tube E observe the color of what 
has collected. (1) What is it? (2) In what physical state 
did it pass from tube A to tube Ef (3) What is the purpose 
of the cold water in the bottle? (4) What is the process called 
which takes place in this experiment? (5) How can pure water 
be obtained from water containing a soluble impurity? (6) 
How could water fit to drink be prepared from sea water? 
(7) Could water be purified from an insoluble impurity by 
this method? 

(/) Wet a piece of pink litmus paper with water from the 
tap. (1) How does pure water affect pink litmus? (2) How 
does solution 4 afi'ect pink litmus paper? Repeat experiment 
(e) substituting solution 4 for solution 3. Test the distillate 
in tube E with litmus paper. (3) How does it act? (4) What 
does the distillate contain? (5) In Avhat physical state did it 
pass from tube A to tube Ef (6) From what class of sub- 
stances does distillation fail to purify water? 

Experiment 19. Hydrates. 

Reference: 98-99. 

Discussion : Some substances which have crystallized from 
water solution give off Avater when heated, and the original 
crystalline substance is again formed upon adding Avater to 
what remains after heating. If this water is present in the 
crystals as water, the material is obviously a mixture, a mixture 
of the substance and water. If such a mixture were heated, 
the water would be given off and the pure substance would be 
left unchanged. If, however, the crystals themselves were a 
pure substance, which Avas a compound of hydrogen, oxygen, 
and other elements, heat might decompose the compound caus- 
ing the hydrogen and oxygen to form water and leaAdng behind 
some new compound, composed of the other elements. In this 
case heat would have brought about a chemical change. More- 
over, if the hydrogen and oxygen were present as part of a 
compound, the law of definite proportions would apply, that 
is the same weight of water would always be formed from the 



82 LABORATORY EXERCISES 

same weight of crystal, but if the crystals were merely a mix- 
ture of a substance with water no such definite relation would 
be found. Crystalline salts which behave as described above 
and prove to hold the hydrogen and oxygen in chemical com- 
bination are called hydrates. 

Materials: Rock salt, copper sulphate, potassium chlorate. 

Directions ; (a) Grind in a dry mortar several large crystals 
of rock salt. (1) What evidence is there of the presence of 
water? Place a large crystal of the salt on a wire gauze over 
your burner and heat it. (2) What happened? (3) If the 
crystal contained within it a cavity filled with water, what 
would happen to this water on heating? (4) What effect 
would this have on the crystal? (5) Do you conclude crystals 
of rock salt to be pure or a mixture? (6) Why? 

(&) Grind in a dry mortar a spatulaful of blue crystals of 
copper sulphate. (1) Is there evidence of water? Put the 
copper sulphate into an evaporating dish, on a wire gauze, 
supported on a ring, attached to a ring stand. Cover the dish 
with an inverted funnel and heat. (2) What formed on the 
inside of the funnel? You must watch carefully to observe 
this deposit, as it soon disappears again. (3) Why? (4) What 
change in appearance has taken place in the material in the 
dish? (5) Has a chemical change taken place? (6) What is 
the evidence? When the dish has become cool enough to 
handle, let a few drops of water fall upon the material which 
has been heated. (7) What happens? (8) Does water react 
chemically with the substance? (9) Do you conclude blue 
copper sulphate to be a hydrate or not? (10) Why? 

(c) Find out experimentally whether or not water is given 
of¥ upon heating potassium chlorate. (1) What did you do? 
(2) What did you find as a result? (3) Are all crystalline 
substances hydrates? 

Experiment 20. Efflorescence. Deliquescence. 

Reference: 100. 

Discussion : As was shown in Experiment 19 some hydrates 
like copper sulphate may be decomposed by heating; Avith 
others, ordinary room temperature is sufficiently hifih to bring 



84 LABORATORY EXERCISES 

about this action. Another factor besides temperature is in- 
volved, however, and must not be overlooked, namely that the 
anhydrous compound will take up water and change to the 
hydrate, if water vapor is present in sufficient concentration. 
Thus when a hydrate is placed in a sufficiently dry atmosphere 
it decomposes, giving off water. This process is called 
efflorescence. On the other hand, if the amount of water vapor 
in the air is so great that the anhydrous compound, were it 
formed, would immediately take up water again, no appre- 
ciable efflorescence will occur. 

Another phenomenon, deliquescence, Avhich at first sight 
appears to be connected with efflorescence, may be studied at 
this time, although it is really dependent on another principle. 
It will be recalled that in Experiment 15 the water vapor in 
the air was condensed to liquid water on the outside of a 
beaker containing ice. This was because the vapor pres- 
sure of water was lowered below that of the vapor in the 
air by cooling. When a substance is dissolved in water, 
the vapor pressure of the water is also diminished, and in a 
concentrated solution its value may become less than that of 
the vapor in the air, just as in the case of cooling. So on the 
surface of the solution condensation from the air will take 
place, as it did on the cold surface of the beaker. Thus a 
substance which is very soluble will make of itself a solution 
by causing condensation of water from the air and may even 
gather enough to become completely dissolved. Substances 
which cause water to be condensed upon them in this way are 
said to be deliquescent. 

Materials: Sodium carbonate, sodium sulphate, potassium 
chlorate, calcium chloride, sodium hydroxide. 

Directions: (a) Put on a glass plate a lump of crystallized 
sodium carbonate, upon another plate a lump of crystallized 
sodium sulphate and upon a third plate a few crystals of potas- 
sium chlorate. Put into a watch glass a lump of calcium 
chloride and into another a piece of sodium hydroxide. Label 
these substances and set them away on the shelf in your desk 
until the next exercise. Then answer the folloAving questions. 
(1) Describe any changes in appearance of each. (2) Which 



86 LABORATORY EXERCISES 

were efflorescent? (3) Which were deliquescent? (4) Are all 
substances either efflorescent or deliquescent? (5) What is the 
evidence? (6) How can the rate of efflorescence be increased? 
(7) How can efflorescence be prevented? (8) What was the 
source of the water which gathered on the deliquescent sub- 
stances? (9) By decreasing what property of water is its 
vapor caused to condense? (10) What is the effect of solution 
on this property? (11) Do deliquescent substances form solu- 
tions? (12) What property of a substance causes it to be 
deliquescent ? 

(6) Put a spatulaful of dehydrated calcium chloride into a 
dry test tube and drop on top of it a crystal of hydrated copper 
sulphate. Into another dry test tube put a crystal of the copper 
sulphate alone. Cork the tubes and set them aside in your desk 
until the next exercise. Then answer the following questions. 

(1) What is the appearance of the copper sulphate in each tube? 

(2) Is the salt efflorescent in ordinary air? (3) What is the 
evidence? (4) What is the cause of its change in the tube 
containing the calcium chloride? 



CHLORINE 

Experiment 21. Preparation of Chlorine. 

Reference : 112. 

Discussion: The methods of preparing chlorine fall into 
four general classes, the electrolysis of a chloride, which is the 
chief industrial method, the oxidation of hydrochloric acid in 
which the oxygen from the oxidizing agent unites with the 
hydrogen of the acid and liberates the chlorine, the decom- 
position of a hypochlorite, and the formation of a chloride 
which immediately breaks up into chlorine and another chloride 
containing less chlorine. The last of these methods is the one 
commonly used for the laboratory preparation of the element. 

Manganese dioxide reacts with hydrochloric acid to form 
manganese tetrachloride, Avhich immediately decomposes to 
give chlorine and manganese dichloride. These reactions take 
place readily and it is usually not necessary to heat the sub- 
stances. If the generation of the gas becomes very slow to- 
ward the end of the preparation, the generator may be slightly 
heated with a flame. Care must be taken, however, not to heat 
it too much, for hydrochloric acid, which is a gas dissolved 
in water, will be driven off and become mixed with the chlorine, 
if the generator becomes too hot. 

A mixture of sodium chloride and sulphuric acid may be 
substituted for the hydrochloric acid in the preparation of 
chlorine. The sulphui'ic acid reacts Avith the sodium chloride 
and produces hydrochloric acid. This modification of the 
process, therefore, consists merely in making the hydrochloric 
acid in the generator instead of adding it already made to the 
generator. 

Materials: Manganese dioxide, concentrated hydrochloric 
acid. 

Directions: Precaution: Bo not inhale chlorine. All ex- 
perimeyiis with chlorine must he performed in the hood. Set 

88 



90 



LABORATORY EXERCISES 



up an apparatus as shown in Fig. 6. A is an 8-inch test tube, 
B a 250 c.c. bottle, C a piece of filter paper, through the 
hole in which passes the delivery tube D. The test tube is tipped 
to facilitate heating it, if that becomes necessary. 

Weigh out 8 grams of manganese dioxide on a filter paper, 
transfer it to the test tube, replace the stopper and pour in 
10 c.c. of concentrated hj^drochloric acid through the thistle tube. 
The evolution of gas should begin immediately. Place a sheet 
of white paper behind the bottle, so th^t you may see when it 
has become full. Precaution : Do not hcnd down, while watch- 
ing the bottle fill, so that your head is on a level with it, because. 




Fig. 6. 



if any chlorine has escaped from the bottle, it will lie chiefly 
along the floor of the hood and you may inhale it. When the 
bottle is full, replace it with another and collect in all four 
bottles full. Heat the test tube slightly if the generation of 
the gas becomes too slow toward the close of the experiment. 
(1) What may result if the generator gets too hot? As soon 
as the last bottle of gas has been collected substitute for it a 
test tube, held in the test-tube rack, about one-third full of 
water. Arrange the delivery tube so that it dips into the water 
and continue the generation of the chlorine for three minutes. 
At once disconnect the apparatus and empty the contents of 
the generator into the large bottle in the hood provided for 
the purpose. (2) Write the equation for the reaction which 



% 



92 LABORATORY EXERCISES 

has taken place. (3) Did the manganese dioxide act as a 
catalytic agent as it did in the preparation of oxygen? (4) 
Give the reasons for your answer. (5) To which of the four 
general classes of preparation of chlorine named in the dis- 
cussion does this method belong? 

Experiment 22. Properties of Chlorine. 

Reference: 117-130. 

Discussion : Chlorine can be recognized by its odor, by its 
color, and by its bleaching action on certain dyes, when it is 
present in fairly large amounts. Small amounts of chlorine can 
be detected by the action on a mixture of potassium iodide 
and starch, but as a similar action is produced by several other 
substances, such as bromine, ozone, and hydrogen peroxide, 
this test for chlorine can be used only when these substances 
are known to be absent. 

Chlorine will unite with elements of widely different chem- 
ical nature, such as phosphorus, copper, and hydrogen, but the 
difficulties of demonstrating all of these reactions are too great 
to make the experiments suitable for a large elementary class. 
The reactions of some of the metals, hoAvever, can easily be 
shown, as well as the behavior of chlorine toward the com- 
pounds of carbon. 

The reaction of chlorine with water illustrates a new type 
of reaction, one in which part of the chlorine forms one com- 
pound and part forms another, hydrochloric acid and hypo- 
chlorous acid. A new principle is also exemplified here, namely 
that of an incomplete reaction, where there exist together in 
equilibrium both some of the original substances and some of 
the products of the reaction. The reaction between chlorine 
and water is important because it is involved in the process 
of bleaching. 

Materials : Iron wool, copper wire, wax taper, colored cot- 
ton cloth, newspaper, fountain pen ink, litmus paper, starch, 
potassium iodide solution. 

Directions: Perform these experiments in the hood, {a) 
(1) What is the color of chlorine? (2) Describe its odor. (3) 
Is it heavier or lighter than air? (4) What is the evidence? 



94 LABORATORY EXERCISES 

(5) Why was it not collected by displacement of water, as was 
hydrogen? 

(6) Ignite a wad of iron wool in the burner and quickly 
thrust it into a bottle of chlorine. (1) Describe the result. 
(2) What is the name of the substance formed? (3) Write the 
equation. (4) Assuming that other metals behave like iron, 
make a statement concerning the action of chlorine with metals 
in general. 

(c) Thrust a blazing wax taper into a bottle of chlorine. 
The taper is composed of compounds of carbon and hydrogen. 

(1) What is deposited on the sides of the bottle? (2) What 
element does not unite directly with chlorine? Blow across 
the mouth of a bottle of concentrated hydrochloric acid. (3) 
What did you observe? (4) Blow across, not into, the mouth 
of the bottle into which the burning taper was thrust. Pre- 
caution : Be careful not to inhale the gas from the bottle. (5) 
What does this test show was formed by the action of the 
chlorine on the taper? 

(d) Fasten together on the ends of two copper wires two 
sets of strips each consisting of a piece of colored cotton cloth, 
a piece of litmus paper, and a piece of printed newspaper, upon 
which has been made a cross with fountain pen ink and one 
with lead pencil. (The ''lead" of a lead pencil and printers' 
ink are composed of forms of carbon.) Thoroughly wet one 
of the sets, and suspend each set in a bottle of chlorine. (1) 
State what happened to the cloth, the litmus i^aper, the printers' 
ink, the fountain pen ink, and lead pencil mark in each case. 

(2) Did the method of preparing the chlorine produce dry or 
moist chlorine? (3) AVhy? (4) Judging by the relative action 
on the two sets of strips, Avhat do you believe Avould have 
been the action on the dry set, had the chlorine been dried 
before using it? (5) Is dry chlorine a bleaching agent? (6) 
Is water? (7) What was the bleaching agent in this experi- 
ment? (8) What elementary substance is not bleached by 
chlorine in the presence of water? 

(e) Grind in your mortar a lump of starch the size of a 
small pea, transfer the starch to a small test tube, fill the tube 
half full of water, and heat it to boiling. Add 5 drops of 



96 LABORATORY EXERCISES 

potassium iodide solution and then wet a strip of filter paper 
with the liquid. Hold the piece of filter paper in the bottle 
of chlorine in which the dry cloth and paper were suspended 
in {d). (1) What happens to the starch iodide mixture on the 
paper? (2) What is a delicate test for chlorine? (3) What 
common substances must be absent in order to make this test 
trustworthy ? 

(/) Put a drop of dilute hydrochloric acid, of dilute sul- 
phuric acid, and of dilute nitric acid on each of three pieces 
of blue litmus paper. (1) What is the action of acids on blue 
litmus paper? (2) What is a test for an acid? Dip the end 
of a glass rod into the water into which you passed chlorine 
in Experiment 21 and touch a piece of blue litmus paper with 
it. (3) What change in color takes place immediately? (4) 
What kind of a compound or compounds does the water con- 
tain? (5) What are the names of the substances formed by 
the action of chlorine on water? (6) Write the equation for 
the reaction. Cautiously smell the solution. (7) Does it 
still contain chlorine? (8) Did the reaction, the equation of 
which you have just written, proceed to completion? (9) 
What is the evidence? (10) What substances are present in' 
** chlorine water"? (11) Which of the substances is the bleach- 
ing agent in ''chlorine Avater"? (12) What are your reasons 
for ascribing the action to this compound rather than to the 
other substances present? 

Experiment 23. Other Means of Preparing Chlorine. 

Bef erence : 116. 

Discussion : When hydrochloric acid is treated with a suit- 
able oxidizing agent, such as potassium permanganate, the 
oxygen from the oxidizing agent unites with the hydrogen of 
the hydrochloric acid and chlorine is liberated. These methods 
are often very couA^enient but are more expensive than the 
method involving the use of manganese dioxide. 

When a hypochlorite is treated with hydrochloric acid, 
hypochlorous acid is formed from the hypochlorite. It Avas 
shown in Experiment 22 / that the reaction between water and 
chlorine to form hydrochloric acid and hypochlorous acid was 



98 LABORATORY EXERCISES 

incomplete. Incomplete reactions are reversible, that is, the 
products of the reaction partially interact in turn and produce 
again some of the original substances. Thus chlorine and 
water form hydrochloric and hypochlorous acids and these will 
react again to form chlorine and water. Both reactions pro- 
ceed simultaneously, so that at any given moment all four 
substances are present. When now a hypochlorite is treated 
with an excess of hydrochloric acid both hypochlorous acid 
and hydrochloric acid will be present, and, as just stated, they 
will react to give chlorine. 

Bleaching powder is a complex salt from which both hydro- 
chloric and hypochlorous acids are produced when it is treated 
with many different acids, such for example as sulphuric acid. 
It is thus a convenient substance from which to prepare 
chlorine. 

Mati:rials: Potassium permanganate, concentrated hydro- 
chloric acid, sodium hypochlorite solution, bleaching powder, 
cotton and silk cloths. 

Directions: (a) Put five crystals of potassium permanganate 
in a test tube, add not more than 5 c.c. of concentrated hydro- 
chloric acid and heat the tube slightly. Cautiously smell at 
the mouth of the tube. (1) What was formed? (2) To what 
class of substances does potassium permanganate belong? (3) 
State the general method of preparing chlorine of which this 
experiment is an example. 

(5) To 5 c.c. of sodium hypochlorite solution in a test tube 
add 5 c.c. of concentrated hydrochloric acid and heat the tube 
slightly. (1) Was chlorine evolved? (2) How did you test 
for it? (3) What was formed in the first place by the action 
of the acid on the hypochlorite? (4) Write the equation for 
the reaction between this product and hydrochloric acid. (5) 
To what reaction studied in Experiment 22 is this related and 
how? (6) What term is applied to a pair of such reactions? 

(c) Put a spatulaful of bleaching powder into an evaporating 
dish and add enough water to make it into a thick paste. 
Add 10 drops of dilute sulphuric acid and cover one end of 
a strip of colored cloth with the mixture. After it has stood 
for 5 minutes, remove the cloth and wash it under the tap. (1) 



100 LABORATORY EXERCISES 

What has happened to the cloth? (2) What produced the 
action? (3) What was the source of this substance? 

{d) On a piece of white cotton cloth make a spot about the 
size of a dime with fountain pen ink. Spread the spotted part 
of the cloth over an evaporating dish and slowly pour drop 
by drop a few cubic centimeters of sodium hypochlorite on the 
spot. Next pour on in the same way a few cubic centimeters 
of dilute hydrochloric acid. Wash the cloth in plenty of water. 
(1) What was the action on the ink? (2) What practical 
application can be made of these solutions? (3) What would 
be the result of removing ink from colored cloth by this 
method? 

Repeat the experiment just described, using a piece of 
thin silk cloth instead of cotton cloth. (4) What Avas the 
result? (5) Draw a conclusion as to the range of usefulness 
of this means of removing ink stains. 



HYDROCHLORIC ACID 

Experiment 24. Preparation of Hydrochloric Acid. 

Eef erence : 138. 

Discussion : The study of the action of sulphuric acid upon 
sodium chloride is important, since this is the reaction which is 
commonly made use of in preparing hydrochloric acid, since it is 
an example of a very general method for the preparation of 
acids (namely, the action of sulphuric acid on a salt), and since 
it affords an excellent illustration of double decomposition. 

It will be recalled that in the preparation of hydrogen 
by the action of an acid on a metal in Experiment 8 the metal 
took the place of the hydrogen of the acid and the hydrogen 
thus liberated, having nothing to unite with, escaped as a gas. 
If, however, a compound of a metal had been used instead 
of the uncombined metal, the other element or elements of the 
compound would have been left when the metal replaced the 
hydrogen of the acid. This element or these elements thus 
set free would become available to unite with the hydrogen 
just displaced and to form a compound. Thus the metal and 
hydrogen would have changed places, that is a reaction of 
*'doul)le decomposition" would have taken place. In this case, 
therefore, the sodium of the sodium chloride will displace the 
hydrogen of the sulphuric acid, which in turn Avill imite with 
the available chlorine and produce hydrochloric acid. 

If this reaction is to be used to prepare hydrochloric acid 
economically, the reverse action, namely hydi'ochloric acid and 
sodium sulphate forming sodium chloride and sulphuric acid 
must be prevented. Since, however, hydrochloric acid is a gas, 
it may be allowed to escape as fast as formed and not enough 
will remain to allow the reversal. 

As sulphuric acid contains two atoms of hydrogen in each 
molecule, the sodium may replace one or both of the hydrogen 

102 



104 LABORATORY EXERCISES 

atoms forming either sodium hydrogen sulphate or disodium 
sulphate. At the temperature of the laboratory preparation 
the former reaction takes place, at higher temperatures the 
latter. 

Materials: Sodium chloride. 

Directions: Perforin this experiment in the hood. Measure 
out into an evaporating dish 8 c.c. of water and add to it 
10 c.c. of concentrated sulphuric acid. (Do not add the water 
to the acid.) While the solution is cooling set up an apparatus 
like the one for the preparation of chlorine, Experiment 19. 
Put 10 grams of sodium chloride into the test tube and pour 
the acid down the thistle tube. When the salt has become 
thoroughly moistened with the acid, heat the test tube with a 
low flame, holding the burner in your hand. To determine 
when the bottle has been filled, hold a piece of wet blue litmus 
paper at its mouth. Do not consider the bottle full when 
the first test is obtained but continue generating the gas until 
it overflows in some quantity. Colleot four bottles full and 
cover them with glass plates. 

Immediately empty the generator into the bottle provided 
for the purpose. (1) Write the equation for the reaction which 
took place in the generator. (2) For the reaction which would 
have taken place at a higher temperature. (3) To what class 
of reactions does this preparation of hydrochloric acid belong? 
(4) What elements changed places? (5) What property of 
hydrochloric acid made the completion of the reaction possible? 

Experiment 2,5. Properties of Hydrochloric Acid. 

Reference: 139-146. 

Discussion : A cloud is always composed of minute particles 
of liquid or solid. A gas or vapor is always clear and when 
colorless is invisible. In common parlance a cloud of water, 
like the clouds in the sky or like fog, is sometimes spoken of 
as vapor. This is, strictly speaking, incorrect, as the cloud is 
really made up of minute drops of liquid water, water vapor 
being invisible. 

When a gas ''fumes" in the air, that is when a cloud 
appears upon the gas coming in contact with the air, small 



106 LABORATORY EXERCISES 

drops of solution arc formed. This occurs when the gas is 
so soluble that a solution can be produced the vapoi' pressure 
of which is less than the pressure of the Avater vapor existing 
in the air. Condensation then takes place and drops of liquid 
are formed. This condensation depends upon the same prin- 
ciple as does deliquescence which was discussed in Experiment 
20 and should be here reviewed. 

For the recognition of hydrochloric acid in solution two 
independent tests are necessary. The solution must be shown 
to contain an acid and also a chloride. Many substances are 
acids, many other substances are chlorides, hydrochloric acid 
alone is at the same time both an acid and a chloride. 

In testing for a chloride the principle of elimination is made 
use of. The formation of a white precipitate upon the addition 
of silver nitrate indicates the presence of any of a certain 
group of substances ; the action of nitric acid upon the precipi- 
tate reduces the possibilities to a certain few; and the action 
of ammonia indicates the individual among these. 

Materials: Splinter, litmus paper, concentrated hydro- 
chloric acid, solutions of silver nitrate, sodium chloride, potas- 
sium chloride, sodium carbonate, potassium bromide. 

Directions: (a) (1) What is the color of hydrochloric acid? 
(2) Describe its odor. (3) Is it heavier or lighter than air? 
(4) What is the evidence? 

(b) Into a bottle of the gas collected in Experiment 24 thrust 
a lighted splinter. (1) What happened? (2) From what other 
gases which you have studied can hydrochloric acid be dis- 
tinguished by this experiment? 

(c) Invert in a pneumatic trough one of the bottles of 
hydrochloric acid. Keep the bottle covered until its mouth 
is under the water. Shake the bottle slightly and hold it so 
that its mouth is just below the surface of the Avater in the 
trough. Observe the height of the water inside the bottle, 
slip a glass plate over the mouth of the bottle and remove it 
together with the water which it contains. Test the water in 
the bottle with blue litmus paper. (1) What does the bottle 
contain? (2) What is the evidence? (3) When the bottle was 
inverted in the trough, what was the level of the water inside 



108 LABORATORY EXERCISES 

compared to that outside? (4) Explain why this was so. (5) 
Is hydrochloric acid soluble or insoluble in water? (6) What 
Vv'as the gas remaining in the bottle? (7) How did it get there? 
(8) Can hydrochloric acid be collected by displacement of 
water? (9) Why? 

(cZ) Blow across, not into, the mouth of a bottle of gas. (1) 
What appeared? (2) What was it composed of? (3) What 
was the source of the Avater? (4) Why did it condense? (5) 
Would an insoluble gas ''fume" in the moist air? (6) Why? 
(7) To substantiate your answer, name an insoluble gas you 
have studied and state whether or not it ''fumes." 

(e) Wet a piece of filter paper with a solution of ammonia 
and drop the paper into a bottle of hydrochloric acid gas. (1) 
What appeared? (2) Is it a solid, liquid, or gas? (3) What 
is the name of the substance? (4) Write the equation for the 
reaction by which it was formed. 

(/) (1) What does the reagent bottle marked "hydrochloric 
acid" contain? Blow across the mouth of a bottle of concen- 
trated hydrochloric acid and also across the mouth of a bottle 
of dilute hydrochloric acid. (2) What happened in each case? 
(3) Explain the reason for the difference in behavior. 

{g) (1) Hydrochloric acid may be regarded as the chloride 
of what element ? To three test tubes add 1 c.c. of dilute 
hydrochloric acid, 1 c.c. of sodium chloride solution, and 1 c.c. 
of potassium chloi'ide solution, respectively, and then fill each 
a third full of water. Dip into each a piece of blue litmus 
paper. (2) Record the results. (3) How can a solution of 
hydrochloric acid be distinguished from solutions of other 
chlorides? To each of the three solutions add 10 drops of 
silver nitrate solution, then 2 c.c. of nitric acid and finally 
5 c.c. of ammonium hydroxide. Shake the tubes and be sure 
that enough ammonium hydroxide has been added to give a 
distinct odor of ammonia at the mouth of each tube. (4) 
Record what occurred on the addition of each reagent to each 
tube. To two test tubes add 1 c.c. of sodium carbonate solution 
and 1 c.c. of potassium bromide solution respectively. To each 
add 10 drops of silver nih'ate solution. (5) What happened 
in each case? (6) Is the appearance of a Avhite precipitate 



no LABORATOHY EXERCi;SES 

upon the addition of silver nitrate a conclusive test for a 
chloride? (7) Why? Add 2 c.c. of nitric acid to each test 
tube. (8) What happened in each case? (9) How does the 
addition of nitric acid distinguish the precipitate produced 
from a chloride by silver nitrate from that produced from a 
carbonate by silver nitrate? To the test tube which contained 
the potassium bromide add 5 c.c. of ammonium hydroxide or 
enough to give a distinct odor of ammonia. (10) What hap- 
pened? (11) In what respect does a bromide behave differently 
from a chloride? (12) State the complete test for a chloride. 
(13) Why is it necessary to add nitric acid? (14) Why am- 
monium hydroxide? 

(/i) (1) In what experiment have you studied the action of 
hydrochloric acid on a metal? (2) What was formed? (3) 
Write the equation for the reaction. 

Experiment 26. Preparation of Lead Chloride from Lead 
Oxide. 

Reference : 142. 

Discussion: Hydrochloric acid reacts with oxides and 
forms chlorides. The preparation of lead chloride in this way 
is easy to carry out because the compound is only slightly 
soluble in cold water. The equation for the reaction is as 
follows : 

PbO + 2HC1 = PbCL + H.O 

Materials: Lead oxide, concentrated hydrochloric acid. 

Directions: Place 4 grams of lead oxide in a beaker and 
add to it 300 c.c. of water and 30 c.c. of concentrated hydro- 
chloric acid. Boil the solution for about 5 minutes and then 
pour it off from the undissolved solid into a second beaker. 
Let the solution cool to room temperature. Filter off the 
crystals and set the paper aside until it is dry. Weigh the 
lead chloride on the paper, using a piece of filter paper of 
the same size in the other pan to counter-balance the weight 
of the original paper. (1) Record the weight of the compound. 
(2) Describe its appearance. 



112 LABORATORY EXERCISES 

Experiment 27. The Quantitative Determination of the 
Hydrogen and Chlorine in Hydrochloric Acid. 

Reference: 142-143. 

Discussion : When zinc reacts with hydrochloric acid each 
atom of the metal replaces two atoms of hydrogen. This fact 
can be used to determine the amount of hydrogen in a given 
wei^fht of the acid by treating the latter with a weighed amount 
of zinc. When reaction is complete the metal is weighed again. 
The loss in weight is a measure of the hydrogen, since 65.4 
grams of zinc dissolve when 2.016 grams of hydrogen are set 
free. 

The chlorine in a known weight of hydrochloric acid can 
be determined by adding to the solution of the acid a solution 
of known concentration of silver nitrate, as long as silver 
chloride is precipitated. From the volume of the solution of 
silver nitrate used and from the concentration of the latter 
(number of grams AgN03 per cubic centimeter) can be cal- 
culated the weight of silver nitrate and the weight of chlorine 
in the acid. 

Materials: Sheet zinc, silver nitrate, concentrated hydro- 
chloric acid, 10 c.c. graduate. 

Directions: (a) Measure accurately from a 10 c.c. graduate 
5 c.c. of pure concentrated hydrochloric acid. Pour the acid 
into a graduated cylinder and dilute with water to 50 c.c. In 
order to avoid adding too much water, it is advisable to fill 
the cylinder up to between 45 and 50 c.c. and add the last 
portion from a glass tube used as a pipeTte. Stir the solution 
Avith a glass rod until it is thoroughly mixed. 

Exactly 10 c.c. of this solution is to be used to determine 
the chlorine and 40 c.c. to determine the hydrogen. Measure 
accurately from the graduate 40 c.c. of the solution into a 
beaker, and place in it a clean piece of sheet zinc about 4 cm. 
square that has been weighed to centigrams. (1) Record the 
weight. Warm the acid until reaction begins and then set it 
aside to use in part c. 

(&) Weigh to centigrams about 2.5 grams of silver nitrate, 
transfer it to a graduated cylinder, and add water to the 50 c.c. 
mark. Stir until the solution is uniform. 



.liiii 




114 LABORATORY EXERCISES 

Add to exactly 10 c.c. of the hydrochloric acid already 
prepared about 50 c.c. of water and then a dilute solution of 
sodium hydroxide until the solution is alkaline to litmus paper. 
From the cylinder add to the solution about 25 c.c. of the silver 
nitrate and stir. Add more of the latter cautiously until the 
colorless milky solution changes to a light yellow-brown color. 
(1) Record the volume used. As long as the solution contains 
a chloride white silver chloride is formed. When the chloride 
has been precipitated, silver oxide, which is brown, is formed 
as the result of the action of silver nitrate on the sodium 
hydroxide present. Calculate (2) from the weight of silver 
nitrate dissolved in 50 c.c. the weight of silver nitrate in 1 c.c. ; 
(3) from the number of cubic centimeters of silver nitrate 
used and from the weight in 1 c.c, the weight used to react 
with the hydrochloric acid; (4) from the fact that 1 AgNOg 
precipitates 1 CI as AgCl, calculate the weight of chlorine 
precipitated. This weight of chlorine was contained in 10 c.c. 
Calculate (5) the weight in the 50 c.c. which contained the 
5 c.c. of concentrated hydrochloric acid. 

(c) Returning to the zinc and hydrochloric acid solution in 
part a, warm the beaker if the reaction has not ceased. When 
the evolution of hydrogen is very slow, the solution can be 
boiled gently. Cool, and if the gas is no longer evolved, 
remove the metal, wash it in running water, and dry it on 
a clean towel. In about 5 minutes weigh it. Record the weight 
of zinc under a 1. (1) From the loss of weight of the zinc and 
from the fact that 65.4 grams of zinc liberate 2.016 grams of 
hydrogen, calculate the weight of hydrogen evolved. This 
amount of hydrogen was contained in 40 c.c. of the acid 
analyzed. (2) Calculate the amount in 50 c.c. The chlorine 
in this amount of acid was calculated in 6 5 above. (3) From 
these figures calculate the weight of hydrochloric acid contained 
in the 50 c.c. (4) From the weights of the acid, hydrogen, and 
chlorine calculate the percentages of chlorine and of hydrogen 
in hydrochloric acid. 

(d) Weigh a graduated cylinder, fill it to the 50 c.c. mark 
with concentrated hydrochloric acid and weigh again. (1) 
Calculate the weight of 1 c.c. of concentrated hydrochloric 



116 LABORATORY EXERCISES 

acid. (2) Calculate the weight of 5 c.c. of the acid. From this 
weight and the amount of pure acid found in c 3 above, cal- 
culate (3) the percentage of hydrochloric acid in concentrated 
hydrochloric acid. 



HYDROGEN PEROXIDE 

Experiment 28. Preparation and Properties of Hydrogen 
Peroxide. 

Reference: 165-167. 

Discussion : Hydrogen peroxide can be prepared by the 
double decomposition of a metallic peroxide and an acid. It 
is a rather unstable compound, easily giving np oxygen with 
the liberation of energy. It acts thus as an oxidizing agent 
and exhibits those properties which would be predicted from 
this fact, such as a bleaching action on certain colored sub- 
stances, and the ability to set free iodine from potassium iodide. 

In the test for hj^drogen peroxide with potassium dichromate 
in acid solution the colored substance formed is more soluble 
in ether than in water. Hence in making the test, if the 
solution is shaken with a little ether, the colored compound 
is extracted by it and concentrated into smaller volume, where 
it is more easily detected. The rate of decomposition of hydro- 
gen peroxide with the evolution of oxygen gas is greatly in- 
creased by the presence of any fine powder, like manganese 
dioxide, Avhich acts as a catalytic agent. 

Materials: Solution of hydrogen peroxide, potassium 
dichromate, ether, barium peroxide, manganese dioxide, cochi- 
neal, potassium iodide, starch. 

Directions: {a) To become familiar with the test for hydro- 
gen peroxide, add to a test tube half full of water 10 drops of 
dilute sulphuric acid, 2 c.c. of potassium dichromate solution, 
a layer of ether 2 cm. thick and finally 3 drops of hydrogen 
peroxide solution. Shake the mixture. (1) What color is the 
ether layer? (2) What is the function of the ether in the test? 

(7)) To 10 c.c. of water in a small beaker add 5 c.c. of dilute 
sulphuric acid and enough ice so that pieces will rest on the 
bottom of the beaker. Slowly sprinkle in a spatulaful of 

118 



120 LABORATORY EXERCISES 

barium peroxide, stir the mixture and filter it. Test the filtrate 
for hydrogen peroxide. (1) What did you do in making the 
test? (2) What did you conclude from the test? (3) Write 
the equation for the formation of the hydrogen peroxide. (4) 
Why was the ice used? 

(c) Into 10 c.c. of hydrogen peroxide solution in a test tube 
sprinkle a quarter of a spatulaful of manganese dioxide. If 
a vigorous evolution of gas does not immediately result, heat 
the mixture. (1) What in all probability is this gas? Verify 
your prediction by making a test. (2) How did you test the 
gas? (3) What did you find it to be? (4) Write the equation 
for its formation. (5) What was the function of the manganese 
dioxide? (6) In what experiment was it used for a similar 
purpose? 

{d) Prepare some starch iodide paper as you did in Experi^ 
ment 22 e and put a few drops of hydrogen peroxide solution 
on it. (1) What occurred? (2) The presence of what element 
is indicated by the color? (3) Write the equation for the 
reaction by which this element was formed. (4) What other 
substance which you have studied produces this action on 
starch iodide paper? (5) What property of hydrogen peroxide 
is illustrated by this action? 

(e) To 5 c.c. of dilute cochineal solution in a test tube add 
5 c.c. of hydrogen peroxide solution and heat. (1) What took 
place? (2) Because of what chemical property of hydrogen 
peroxide did this action take place? 

Experiment 29. The Quantitative Analysis of Hydrogen 
Peroxide. 

Reference: 166-167. 

Discussion : When a solution of hydrogen peroxide is 
treated with potassium permanganate, KMnO^, oxygen is set 
free, one-half of which comes from the peroxide and one-half 
from the permanganate. The volume of the gas liberated in 
this way from a known volume of a solution of hydrogen 
peroxide is a measilre of the strength of the latter. 

Materials : Potassium permanganate, solution of hydrogen 
peroxide. 



122 LABORATORY EXERCISES 

Directions : Arrange an 8-inch test tube with a thistle tube 
and a delivery tube to collect a gas over water. Add to the 
tube a pinch of potassium permanganate and 10 c.c. of water. 
Insert the stopper and place the end of the delivery tube under 
a gas bottle filled with water. Pour into the test tube 10 c.c. 
of a solution of hydrogen peroxide and then 10 c.c. of water. 
When gas ceases to be evolved, cover the mouth of the bottle, 
remove it from the trough and measure (1) the volume of the 
water it contains. Fill the bottle with water and measure (2} 
its volume. The difference (3) is the volume of the gas that 
was collected. To get the volume of the oxygen (4) subtract 
20 c.c. from (3). (5) Calculate what volume of gas is generated 
when 1 c.c. of hydrogen peroxide is decomposed by potassium 
permanganate. 



PROPERTIES OF LIQUIDS AND SOLIDS 

Experiment 30. Determination of Density. 

Reference: 175-176. 

Discussion : The density of a solid or a liquid is defined as 
the weight in grams of 1 c.c. of the substance. In order to 
determine the density of a substance it is necessary to know 
the weight and volume of a sample of it. If the substance 
is a liquid its density can be determined readily by weighing 
a measuring vessel, such as a graduated cylinder, placing in 
it the liquid, the volume of which is noted, and then Aveighing 
the two. The weight of 1 c.c. of the liquid can then be 
determined by calculation. In order to obtain a fair degree 
of accuracy, it is advisable to Aveigh as much as 50 c.c, if the 
platform scales are to be used. 

If the density of a solid is to be determined, its volume is 
found out by immersing it in a known volume of a liquid in 
which it does not dissolve, and noting the sum of the volumes 
of the liquid and the solid. This can be done by placing a 
known weight of the solid in a known volume of the liquid 
contained in a graduated cylinder. The increase in volume 
after the solid has been added is the volume of the latter. 
Water is used most conveniently in the case of metals and other 
substances which are insoluble in water. In the case of 
salts that dissolve in water, benzene, kerosene, or carbon 
tetrachloride can be used. 

Materials: Concentrated hydrochloric acid, iron nails, 
sodium chloride, benzene. 

Directions: (a) Weigh on the platform scales to decigrams 
your graduated cylinder, which should be dry. Fill it exactly 
to the 50 c.c. mark. In order to do this pour in water until 
it is nearly up to the mark and add the rest drop by drop 
from a piece of glass tubing full of water, over one end of 

124 



126 LABORATORY EXERCISES 

which the finger is held. If the pressure of the finger is reduced 
and a little air is allowed to enter the upper part of the tube 
the water can be made to run out of the tube drop by drop. 
(1) Weigh the cylinder and the water. (2) What is the error, 
if any, in the graduation of the cylinder? 

Pour out the water, dry the cylinder, fill it exactly to the 
50 c.c. mark with concentrated hydrochloric acid, and weigh 
again. (3) Calculate the density of the acid. 

(&) Place about 40 c.c. of water in the graduated cylinder 
and note accurately to tenths of a cubic centimeter its 
volume. Slide into the cylinder about 20 grams of nails, which 
have been weighed to decigrams on the platform scales. Read 
to 0.1 c.c. the volume of the water and iron. (1) Record all 
data. (2) What does the difference between this volume and 
that of the water i-epresent? (3) Calcuhite the density of iron. 

(c) Carry out a similar experiment using 20 c.c. of benzene 
and 20 grams of sodium chloride. (1) Record the results. (2) 
Why is benzene used instead of water? (3) What is the rela- 
tion between density and specific gravity? 

Experiment 31. Determination of Specific Heat. 

Reference : 177. 

Discussion : The specific heat of a substance is the amount 
of heat required to raise the temperature of 1 gram of the 
substance 1 degree centigrade. In determining specific heat 
use is made of the following facts: The specific heat of water 
is 1, that is, 1 caloi-ie raises the tempei*ature of 1 gram of water 
1 degree. The heat given off by a hot body Avhen its tempera- 
ture falls 1 degree is equal in amount to that required to raise 
its temperature 1 degree. 

Directions are given below for the determination of the 
specific heat of a solid. A known weight of brass is heated 
in water to a definite temperature. It is then transferred to 
a known weight of water at room temperature. The rise in 
temperature is noted. The product of the number of degrees 
rise in temperature and the weight of the water is evidently 
the number of calories given off by the brass. This figure 
divided by the weight of the brass gives the number of calories 



128 LABORATORY EXERCISE/S 

given off by each gram of brass. The figure obtained in this 
way divided by the number of degrees through which the 
temperature of the brass fell is the number of calories given 
off when 1 gram of brass drops 1 degree in temperature, that 
is, the specific heat of brass. 

In actual practice some of the heat given up by the brass 
is lost to the surrounding air and to the glass of the beaker 
and the thermometer. As a consequence, the rise in the tem- 
perature of the water will not be an exact measure of the 
heat lost by the brass. A correction is made, accordingly, in 
a way which can be understood from the following considera- 
tions. When two equal weights of water at different tempera- 
tures are mixed, the temperature of the mixture is just half 
way between the temperatures of the two portions before 
they were mixed, provided no heat is lost to the vessels con- 
taining the water. Thus, if 50 c.c. of water at 20° is mixed 
with 50 c.c. of water at 30°, the temperature after mixing would 
be 25^ if no heat were lost. If heat is lost, however, the 
difference between the observed temperature and the mean of 
the two temperatures is a measure of this heat. 

Materials: One 200 gram brass weight attached to a string. 

Directions: (a) Fill a 500 c.c. beaker about two-thirds full 
of water, place in it the brass weight, and heat the water to 
about 75°. Measure into a 200 c.c. beaker 100 c.c. of water 
at room temperature. (1) Record the temperature of the water 
estimating the temperature to 0.1 degree. Remove the beaker 
containing the brass from the burner, stir the water slowly 
with the thermometer, read the temperature (preferably when 
the column of mercury stands at a degree mark) and imme- 
diately transfer the brass by means of the string to the small 
beaker, and (2) note the time. Cool the thermometer to room 
temperature by placing it in water at this temperature, and 
insert it into the small beaker containing the brass. Observe 
the thermometer as the temperature rises slowly. (3) Record 
the highest temperature reached and the time when this occurs. 
From the two time observations it will be seen how many minutes 
are necessary for all the heat to be transferred from the brass. 
This time (3 to 5 minutes) will be required in the second part 



130 LABORATORY EXERCISES 

of the experiment. For convenience it will be referred to later 
as T. 

(h) To determine to what extent the loss of heat to the 
surroundings affects the rise in temperature proceed as follows: 
Measure into the small beaker used before 50 c.c. of water at 
rt)om temperature. (1) Record the temperature of the water to 
0.1 degree. Place 50 c.c. of water in a small flask and heat it 
until its temperature is approximately twice as many degrees 
above the temperature of the cold water as the observed rise in 
part (a) above. (2) Read the temperature of the warm water to 
0.1 degree, pour it at once into the 50 c.c. of cold w^ater in the 
beaker and (3) note the time. Cool the thermometer to room 
temperature by immersing it in water at this temperature, and 
put it into the beaker in which the two portions of water were 
mixed. (4) Record the temperature after the time T (see (a) 
above) has elapsed. If there had been no loss of heat from the 
water to the air, the beaker, and the thermometer, the final 
temperature would have been just half way between the tem- 
peratures of the cold and the hot water (the average of the 
two). The difference between this temperature and that ac- 
tually observed is the error introduced as the result of the loss. 
This difference (5) should be added to the rise in temperature 
observed in {a) above. 

(c) (1) From the corrected rise in temperature found in {a) 
calculate the specific heat of ])i'ass in the way outlined in the 
discussion. 

(2) What is meant by the specific heat of a substance? (3) 
Where did most of the heat of the brass go? (4) Where did 
the remainder go? (5) Why was all possible speed used in 
transferring the brass from the beaker? (6) Why was the time 
observed? (8) How was the error of heat loss corrected for? 

Experiment 32. The Effect of Pressure on the Boiling- Point 
of Water. 

Reference : 178. 

Discussion : Boiling is the formation of vapor within a 
liquid, as distinguished from evaporation from the surface 
only. A liquid boils when the pressure of its vapor just exceeds 



132 



LABORATORY EXERCISES 



the pressure exerted upon its surface. Under these conditions 
a bubble of vapor can form under the surface of the liquid. 
Thus, when a liquid is heated in a vessel open to the air, the 
pressure of the latter must be overcome before it boils. If a 
part of the air is withdrawn from the vessel, in any way, the 
air exerts less pressure and, as a consequence, the liquid boils 
at a lovv er temperature. This is due to the fact that a smaller 
vapor pressure, attained at a lower temperature, is now enough 
to exceed the diminished pressure of the air upon the liquid. 
The experiment described below illustrates this fact in a 




Fig. 7. 



simple way. Water is boiled in a vessel connected to a tube 
which dips under water in a second vessel (P'ig. 7). As the 
liquid boils the water vapor formed drives out most of the 
air which escapes through the tube. When the air has been 
driven out, the tube is closed and the source of heat removed 
from the boiling water. Since the apparatus is now closed, 
the only pressure upon the surface of the water is that of 
the small amount of air present, plus the vapor already there. 
As this latter condenses upon cooling, boiling continues for 
some time at the reduced pressure, while the temperature is 
falling. 



134 LABORATORY EXERCISES 

Materials: Rubber tubing. 

Directions : Arrange an apparatus as represented in Fig. 7. 
Use an 8-inch test tube and a glass tube about 18 inches 
long. The bulb of the thermometer should be under the surface 
of the water so that the temperature of the latter can be 
observed. (In determining the boiling point of a liquid the 
thermometer is ordinarily placed in the vapor of the boiling 
liquid, when heat is constantly supplied from an outside 
source.) 

In order to aid boiling place a few pieces of broken glass 
or tile in the test tube and add about 20 c.c. of water. Insert 
the stopper carrying the delivery tube which is fitted with a 
short piece of rubber tubing and a pinch cock. Open the latter 
and heat the water to boiling. Observe closely the change in 
size of the bubbles that escape through the water in the beaker 
as the boiling continues. (1) State what occurs and give a 
reason for the observed facts. 

When the steam makes a hissing sound as it condenses in 
the beaker (2) note the temperature of the boiling water. 
Cautiously heat the test tube with a burner held in one hand 
and close the pinch cock with the other hand. Remove the 
flame immediately, and raise the beaker so that the end of 
the glass tube is well covered by the cold water. Observe the 
thermometer as the water continues to boil. (3) Note the 
lowest temperature observed when the water is boiling. Tap 
the tube lightly to assist in the formation of bubbles of vapor. 

When boiling ceases open the pinch cock under the water. 
(4) What happened? (5) Why? Was there any pressure 
exerted by air on the boiling liquid when the system was 
closed? How could you calculate the value of this pressure 
from the observations made? (6) Did the apparatus contain 
any air when closed? (7) What is the evidence? (8) How 
does boiling differ from surface evaporation? (9) What con- 
ditions must exist for boiling to take place? (10) Ordinarily 
boiling is produced by the increase of what pressure? (11) 
How is this accomplished? (12) By the decrease of what pres- 
sure may boiling be brought about? (13) How can water be 
made to boil below 100°? 



136 LABORATORY EXERCISES 

Experiment 33. Heat of Vaporization. 

Reference : 180. 

Discussion : When a liquid passes into a vapor, heat is 
required to effect the change. The heat required to convert 
1 gram of a substance from the liquid to the vapor state is 
called the heat of vaporization. 

When a liquid is boiled over a flame the heat required is 
obtained from the burning gas. When a liquid evaporates of 
itself the heat is supplied by the surrounding air and by the 
liquid; as a consequence the temperature of the latter falls. 
If the conditions are such that the liquid evaporates rapidly, 
a relatively large amount of heat must be supplied in a short 
time, and the temperature of the liquid falls to a greater extent 
than when the evaporation takes place slowly and heat can be 
taken up from the surrounding air. 

The extent to which the change in temperature takes place 
is determined by the rapidity of evaporation, the heat of 
vaporization of the liquid, and the rate at which the material 
can take heat from the surroundings. 

Materials: Piece of string. 

Directions: (a) Cut from a piece of filter paper a strip 
about 2 cm. wide by 10 cm. long. Wrap the paper around 
the bulb of the thermometer and tie it in place with a piece 
of string. (1) Note the temperature of the air. Dip the end 
of the thermometer in ether, hang it up and observe the tem- 
perature. (2) Record the lowest temperature reached. 

{!)) Dry the paper or replace it by a new piece and repeat 
the experiment using carbon disulphide. (1) Record the tem- 
perature. 

(c) Repeat the experiment using alcohol. (1) Record the 
temperature. 

(d) Repeat the experiment using water. (1) Record the 
temperature. 

(2) What property of a liquid determines the rate at which 
it evaporates at a given temperature, provided it receives a 
constant supply of heat sufficient for the evaporation? (3) 
If two liquids are evaporating at the same rate what determines 
the relative amounts of heat absorbed in the two cases? (4) 



138 



LABORATORY EXERCISES 



The heat of vaporization of ether is 90 calories and that of 
alcohol 206 calories. What conclusion as to their relative vapor 
pressures can you draw from these facts and the results ob- 
tained in the experiment? (5) Upon heating a liquid why 
does the temperature cease to rise when it begins to boil, al- 
though the heating is continued? 

Experiment 34. Distillation of a Mixture of Two Liquids. 
Reference : 181. 

Discussion : A liquid possesses a definite boiling point at 
a fixed pressure, and a definite density at a fixed pressure and 




Fig. 8. 



tenipei-alLire. The density of a homogeneous mixture of two 
liquids varies with the proportions of the two substances 
present. When such a mixture is heated to boiling the lower 
boiling liquid does not distill off first in the pure condition at 
its own boiling point, but a mixture is obtained, the composition 
of which varies as the distillation proceeds. The distillate 
always contains a larger percentage of the more volatile liquid 
than is present in the boiling mixture. As the distillation pi'o- 
ceeds the proportion of the higher boiling liquid thus increases. 
The experiment described below is designed to illustrate the 



140 LABORATORY EXERCISES 

behavior of such a mixture when it is slowly distilled. The 
liquids used are methyl alcohol (wood alcohol), which boils 
at 66°, and water, which boils at 100°. 

Materials: 250 c.c. flask, methyl alcohol, rubber band. 

Directions: Set up the apparatus represented in Fig. 8. 
A test tube. A, placed in cold water in a gas bottle serves to 
collect the distillate. (1) Weigh a graduated cylinder to deci- 
grams. Place in the graduate 15 c.c. of methyl alcohol, note 
the volume to 0.1 c.c. and weigh again. (2) Calculate the 
density of the alcohol. Pour the alcohol into the flask. Wash 
out the cylinder with water and then add 30 c.c. of water 
to the alcohol. Shake the mixture so that it is uniform. Dry 
the cylinder. Pour into it the mixture. Note the volume 
accurately and weigh. (3) Calculate the density of the mixture. 

Pour 2 or 3 drops of the mixture on a glass plate and apply 
a lighted match. (4) Does it burn? Pour the mixture into 
the flask and place a rubber band around the tube which is 
to receive the distillate in such a position that it will serve to 
show when approximately 15 c.c. of liquid have been collected. 
To do this put 15 c.c. of water into the test tube, place the 
tube in position in the bottle and mark the level of the water 
by the position of the band. Empty and dry the test tube 
and place it again in the bottle which contains cold water. 
Arrange the apparatus as indicated in the diagram. As the 
flask is heated in the way described below, (5) note the tem- 
perature recorded on the thermometer when the liquid fii'st 
drops back from the tube at C and also when the distillation 
is stopped. Heat the flask cautiously with a flame. When the 
liquid begins to boil remove the flame, count 5 seconds, heat 
again till the liquid boils (about 1 second), and remove the 
flame. Continue heating cautiously in this way until 15 c.c. 
of the distillate have been collected. (6) Determine the density 
of the distillate, and return it to the test tube. 

Cool the contents of the flask under running water and (7) 
determine its density. 

Pour a few drops of the distillate on a glass plate and 
determine whether or not it will burn. Test the residue in the 
flask in the same way. (8) What were the results? (9) What 



142 



LABORATORY EXERCISES 



do you conclude about the relative amounts of alcohol in the 
distillate and in the residue? 

Pour the distillate into the flask, replace the stopper and 
thermometer and (10) determine the boiling point of the liquid. 
(11) Arrange in tabular form the boiling points, the densities, 
and the percentages of alcohol of the following : methyl alcohol, 
water, mixture, distillate, and residue. 

From the values of the densities obtained, the percentage 
of methyl alcohol in the mixtures can be found from the fol- 
lowing table : 

Methyl Alcohol and Water 



Density 


983 


.972 


.961 


.946 


.929 


.909 


.887 


.863 


.838 


810 








Per cent of alcohol 


10 


20 


30 


40 


50 


60 


70 


80 


90 


100 



(12) What is the density of a mixture of 1 volume of methyl 
alcohol and 2 volumes of water as calculated by finding the 
sum of one-third the density of the alcohol and two-thirds 
of the density of the water? (13) What was the observed 
density? (14) What occurred when the liquids were mixed? 
(15) Can the composition be calculated directly from the den- 
sity? (16) How did you find the relation of composition to 
density? (17) What is the relation of the composition of the 
distillate to that of the boiling mixture? (18) How does the 
composition of the mixture change during distillation? (19) 
Why does the boiling temperature rise? 



Experiment 35. The Determination of a Freezing Point. 

Reference : 187. 

Discussion : The freezing point of a pure liquid is the 
temperature at which it changes to a solid. The melting point 
of a pure solid is the temperature at which it changes to a liquid. 
These temperatures are the same, for any one substance. If 
heat is constantly supplied to a solid it rises in temperature until 
the melting point is reached. The temperature then remains 
constant as heat is supplied, since the added heat is used in con- 



144 



LABORATORY EXERCISES 



1- 



a) 



IW 








































































































































































65 
















































































































































































































^0 
















































































































































































































55 
















































































































































































































50 
















































































































































































































A5 
















































































































































































































AD 














L_ 




















1 









10 
Time in Mmu+e& 
Fig. 9. 



15 



20 



146 LABORATORY EXERCISES 

verting the solid into the liquid. When the solid has melted, 
the temperature of the liquid again begins to rise as the heat 
is supplied. Conversely, when heat is withdrawn from a liquid 
it falls in temperature uiitil the freezing point is reached. The 
temperature remains constant at this point as crystallization 
takes place, because the formation of crystals from the liquid 
occurs with the evolution of heat. 

The change which takes place when a liquid solidifies is 
illustrated by the experiment described below. The freezing 
point can also be determined in the way illustrated in the next 
experiment. 

Materials : Sodium thiosulphate, NasSsOajSH^O. 

Directions: Put approximately 10 grams of sodium thio- 
sulphate in a dry test tube and place the latter in hot water 
(about 70°) in a beaker until the solid has melted. Remove 
the test tube from the water, place the thermometer in it, and 
stir slowly. Read the temperature at intervals of a minute. 
When the thermometer falls to about 55° drop into the tube 
a small crystal of sodium thiosulphate. A little of the solid 
is added to prevent supercooling (see next experiment). Con- 
tinue the stirring and record the temperatures at intervals of 
a minute for 10 minutes. (1) Plot the results on Fig. 9. 

(2) What occurred in the liquid during the time that it 
did not change in temperature? (3) Why is heat lost to the 
surroundings continually throughout the experiment? (4) 
What is the evidence that heat is produced during part of the 
experiment? (5) Which part was this? (6) What was the 
source of the heat? (7) What is the temperature called at 
which this action occurs? (8) What other change will occur 
at the same temperature? (9) What is this temperature called? 

Experiment 36. Supercooling. 

Reference : 189. 

Discussion : Many liquids can be cooled below the tem- 
perature at which they freeze, provided the conditions exist 
which prevent the formation of the solid form. When a trace 
of the solid is added to such a supercooled liquid, crystalliza- 
tion takes place. As the change of a liquid to a solid is aecom- 



148 



LABORATORY EXERCISES 



panied by the evolution of heat, the temperature of the liquid 
rises to that of the melting point and remains constant as the 
solidification proceeds. 



V^vJ 
































































































































50 






























































































S 
































CO 
<L> 






























































































































^ 
































































40 






























































































































































35 

































10 



15 



Time in Minutes 
Fig. 10. 



Materials: Sodium thiosulphate. 

Directions: Melt approximately 10 grams of sodium thio- 



150 LABORATORY EXERCISES 

sulphate in a dry test tube placed in hot water in a beaker 
at about 70°. Care should be taken that the solid particles 
of the substance do not adhere to the walls of the tube above 
the liquid. Remove the test tube from the water, insert a 
thermometer, and cool the tube under running water until 
the temperature of the liquid is approximately 40°. Drop a 
small crystal of sodium thiosulphate into the tube, stir, and 
note the temperature each minute for 10 minutes. 

(1) Plot the results on Fig. 10. (2) At what degree 
did the temperature remain constant for a period? (3) 
What was taking place during this period? (4) How does this 
temperature compare with that of constancy in Experiment 35? 
(5) What is this temperature called? (6) In what condition 
is the liquid said to be when it is below this temperature? (7) 
What precaution must be observed to maintain this condition? 
(8) What was the source of heat when the temperature rose? 



CARBON AND ITS COMPOUNDS 

Experiment 37. Properties of Carbon. 

Eeference: 191-197. 

Discussion : Carbon does not possess a single set of proper- 
ties as do hydrogen and chlorine, but three very different sets. 
A given sample is called diamond, graphite, or amorphous car- 
bon, according to which set of properties it possesses. The speci- 
mens of amorphous carbon which are familiar have all been formed 
by the decomposition of compounds of carbon, often mixed 
with other substances, and so are seldom pure. The differences 
in the methods of formation and the presence of the impurities 
make the amorphous carbon from different sources appear to 
have dift'erent properties, as in the case of coal, charcoal, and 
soot ; but when the pure carbon is extracted from these sub- 
stances, it is all found to be the same. 

The physical properties of these allotropic forms vary more 
widely than do the chemical properties and it is the peculiar 
physical properties of diamond and of graphite which make 
these substances of commercial value. 

Materials: Powdered graphite, lumps of charcoal, powd- 
ered wood charcoal, animal charcoal, vinegar, hydrogen sul- 
phide solution, lead oxide. 

Directions: (a) Rub a little graphite between your thumb 
and forefinger. (1) Describe the feeling. (2) What use is 
made of graphite dependent on the property which causes this 
feeling? 

Rub a bit of graphite across a white paper. (3) What was 
produced? (4) What common use of graphite depends upon 
this property? 

Make a small hole in a lump of charcoal and fill it with 
some powdered graphite. Holding the burner in your hand, 
let the flame play directly on the hole. (5) AVhich ignites 
more readily, the graphite or the charcoal? 

152 



154 LABORATORY EXERCISES 

(6) Put a small lump of charcoal, not more than a cen- 
timeter cube, into water. (1) Does it sink or float? Hold the 
lump of charcoal with the tongs in the flame until it is red 
hot and quickly put it in the water again. (2) Does it now 
sink or float? (3) What is the cause of the difference in 
behavior? (4) What does this experiment suggest as to the 
structure of a piece of charcoal ? 

(c) Fill a test tube a quarter full of animal charcoal, add 
10 c.c. of vinegar, shake the test tube until the charcoal has 
become moistened throughout, and then heat it to boiling. 
Filter into a test tube and compare the color of the filtrate 
with some of the original vinegar. (1) What change in ap- 
pearance has been produced? (2) Was this a physical or a 
chemical change? (3) To what physical property of the char- 
coal was the action due? 

Repeat the experiment with the substitution of 10 c.c. of 
hydrogen sulphide solution for the vinegar. Compare the odor 
of the filtrate with that of the original solution. (4) What 
change has the charcoal produced? (5) What has become of 
the hydrogen sulphide? (6) What uses of charcoal do these 
experiments suggest? 

(d) Mix in your moi'tar one-quarter of a spatulaful of lead 
oxide and the same amount of powdered wood charcoal. Put 
some of the mixture in a hole in a lump of charcoal and heat 
it with the blow pipe. A little preliminary practice with the 
blow pipe may be necessary, which may be carried out as 
follows: First, with your mouth closed, puff out your cheeks 
and breathe through your nose several times, till you can do 
this freely. Now put the blow pipe in your mouth and repeat 
the performance, keeping your cheeks puffed out. The air in 
your mouth will become exhausted only slowly and may be 
replenished by exhaling into your mouth fi'om time to time, 
while still breathing through your nose. After a little practice 
you will be able to blow a continuous stream of air through 
the blow pipe. Slip the blow-pipe tube down the tube of your 
burner, light the gas and turn it down till the flame is about 
half an inch high. Hold the bloAv-pipe in your mouth, steady- 
ing it with the right hand and letting the tip rest in the crotch 



156 LABORATORY EXERCISES 

of the blow-pipe tube. Hold the piece of charcoal in the left 
hand and direct against it the flame produced by blowing into 
the tube. Heat the mixture of lead oxide and charcoal for 
two or three minutes, then shake it out on your desk and 
examine it. (1) What substance do you find? (2) Write the 
equation for the reaction by which it was formed. (3) As 
what kind of an agent has the charcoal acted? (4) In what 
commercial processes are similar reactions made use of? 

Experiment 38. Preparation of Carbon Dioxide. 

Reference : 200. 

Discussion : Of the various methods for the formation of 
carbon dioxide, which will be taken up later in Experiment 40, 
the most convenient for laboratory preparation is the action 
of an acid on a carbonate. The progress of the reaction de- 
pends upon the instability of carbonic acid, which breaks up 
into water and carbon dioxide. Substances are chosen, there- 
fore, which will react to form carbonic acid by double decom- 
position and these are, as just stated, a carbonate and an acid. 
If a slow steady stream of gas is desired, an insoluble car- 
bonate, usually calcium carbonate, is selected, since a soluble 
carbonate reacts too rapidly. The acid employed must be one 
which will form a soluble compound with the metal of the 
carbonate, for if an insoluble compound were formed, it would 
coat over the solid carbonate and prevent further access of 
the acid. 

Materials: Marble, concentrated hydrochloric acid. 

Directions: (a) Use the same apparatus as for the prepara- 
tion of hydrogen in Experiment 8. Fill the generator about 
one-third full of lumps of marble, slipping the pieces in while 
the tube is inclined, so that they will not crack it. To 20 c.c. 
of water in a graduate add 20 c.c. of concentrated hydrochloric 
acid and pour through the thistle tube enough of this mixture 
to cover completely the marble. Collect three bottles of carbon 
dioxide as you have collected gases over water before. (1) 
Which side up should the filled bottles be kept? (2) Why? 
(3) By the spontaneous decomposition of what substance is 
the carbon dioxide formed? (4) Write the equations for the 



158 LABORATORY EXERCISES 

two reactions taking place in this preparation. (5) To what 
class of reactions does the first belong? (6) The second? (7) 
"Why does the first take place to completion? (8) Why does 
the second? (9) Write the equations for the reactions which 
would have taken place if sodium carbonate had been used 
instead of calcium carbonate. (10) If sulphuric acid had been 
used instead of hydrochloric acid. (11) Why is calcium car- 
bonate preferred over sodium carbonate? (12) Why is a hydro- 
chloric acid used instead of sulphuric acid? (See table of 
solubilities, page 543.) 

(h) Pour the liquid out of the generator, put in more marble 
if necessary, disconnect the delivery tube at the rubber con- 
nector and substitute for it a tube passing to the bottom of 
an 8-inch test tube through a two-hole rubber stopper. Into 
the other hole of the stopper fit a glass tube, bent at a right 
angle and barely passing through the stopper. Connect this 
tube by means of a rubber connector to a second right-angled 
tube passing into a second test tube. Fill both test tubes one 
third full of water, pour acid into the generator as in (a), 
and allow the action to proceed for at least 10 minutes. Save 
the contents of both test tubes for Experiment 39. 

Expi']RiMENT 39. Properties of Carbon Dioxide. 

Reference: 201-205. 

Discussion : Carbon dioxide may be used to illustrate the 
behavior of a gas heavier than air, because, being odorless, it 
produces no disagreeable results if it escapes. Many other 
gases, as chlorine and sulphur dioxide, are much heavier but 
are not suitable for experiments on density. Likewise the 
action of carbon dioxide on burning matter is not a unique 
property of this gas but this phenomenon can be most con- 
veniently shown with carbon dioxide because it is so easily 
formed and entirely inoffensive. 

In the test for carbon dioxide calcium carbonate is produced. 
It is obvious that in using this test care must be taken to 
prevent acid from coming in contact with the reagent, for the 
acid would act on the calcium carbonate. A fine spray from 
the acid used in the production of carbon dioxide often acts 
this way unless precaution is taken. 



160 LABORATORY EXERCISES 

Materials: Splinter, lime water, silver nitrate solution, 
litmus paper. 

Directions: (^0 (1) What is the color of carbon dioxide? 

(2) What is its odor? 

(h) Thrust a blazing splinter into a bottle of carbon dioxide. 
(1) What occurred? (2) What other gas which you have 
prepared acted similarly? (3) Why is carbon dioxide to be 
preferred to this gas as a fire extinguisher? 

(c) Bend a splinter at a right angle, light an end of it and 
lower it into a bottle of air. Immediately pour into the bottle 
a bottleful of carbon dioxide, holding the bottles mouth to 
mouth, and pouring as if the carbon dioxide were a liquid. 
(1) What occurred? (2) Is carbon dioxide heavier or lighter 
than air? (3) What is the evidence? (4) What other gases 
which you have prepared have a similar density compared 
with air? 

{d) Into a bottle of carbon dioxide pour 5 c.c. of lime water 
(a solution of calcium hydroxide) and shake the bottle. (1) 
What appeared? (2) What is the name of the compound? 

(3) Write the equation for the reaction for its formation. (4) 
What is the test for carbon dioxide? Add 5 c.c. of hydro- 
chloric acid to the bottle. (5) What occurred? (6) What 
precaution must be taken in testing for carbon dioxide? 

(e) Test with litmus paper the solution contained in the 
second of the test tubes through which carbon dioxide was 
bubbled in Experiment 38 h. Add 5 c.c. of lime water to 
the solution. (1) What was the action on the paper? (2) 
With the lime Avater? (3) What compound does the solution 
contain? (4) Write the equation for its formation. (5) Does 
the reverse of this reaction take place? (6) What is the evi- 
dence? (7) Explain how it is that carbon dioxide reacts with 
water and yet can be collected over water. 

Test for a chloride in the water contained in the first test 
tube in Experiment 38 h. (8) What did you do to make 
the test? (9) What was the result? (10) What do you con- 
clude the water contains? (11) How did it get there? (12) 
Why did it not pass on into the second test tube? (13) Why 
was it necessary to have this tube as part of the apparatus? 



162 LABORATORY EXERCISES 

Experiment 40. Formation of Carbon Dioxide by Various 
Means. 

Reference: 200. 

Discussion : As was shown in Experiment 38, carbon dioxide 
is formed from the spontaneous decomposition of carbonic acid 
(hydrogen carbonate) which in turn is formed by the action 
of an acid on a carbonate. Carbonates other than hydrogen 
carbonate decompose giving carbon dioxide and an oxide, but 
they must be heated above room temperature. The heating of 
limestone to make lime is a notable example of this action. 

The burning of carbon in air is an obvious method of forma- 
tion, as is also the complete oxidation of compounds of carbon, 
either by combustion or natural processes of slow oxidation. 
Some compounds of carbon also form carbon dioxide by decom- 
position in the process of fermentation, which is brought about 
by the presence of micro-organisms. 

Materials : Sodium carbonate, lime water, magnesium car- 
bonate, marble, charcoal, paper, candle, ''Karo Syrup," yeast 
cake, copper wire. 

Directions: (a) To half a spatulaful of sodium carbonate 
in a test tube add 5 c.c. of dilute sulphuric acid. Hold in the 
mouth of the test tube, Avith a steady hand, a glass rod which 
has been dipped into lime water and upon the end of which 
a drop of the lime water is hanging. (1) How did the drop 
change in appearance? (2) What was formed? (3) What 
gas was found to l)e present? (4) If the drop shows no change, 
what is the probable cause for the failure of the experiment? 
(5) The action between Avhat other carbonate and acid have 
you studied? (6) Describe the test for a carbonate. 

(h) Heat in a small test tube half a spatulaful of manganese 
carbonate and then test for carbon dioxide as you did in ex- 
periment (a). (1) What is the evidence as to the decomposi- 
tion of the manganese carbonate? (2) Write the reaction. 
Repeat the experiment substituting a lump of calcium car- 
bonate for the manganese carbonate. (3) Is there evidence of 
decomposition? (4) In what commercial process is calcium 
carbonate decomposed by heat? (5) Why is this not used as 
a laboratory method for preparation? 



164 LABORATORY EXERCISES 

(c) Hold a piece of charcoal with the tongs in the flame 
until it glows and then quickly thrust it into a bottle of air 
for a minute. Test the gas in the bottle for carbon dioxide. 

(I) How did you make the test? (2) What was the result? 
(3) Write the equation for the burning of the charcoal. 

Drop a piece of burning paper into a bottle of air and test 
for carbon dioxide. (4) Of what element is paper a compound? 
(5) What is formed when it burns? (6) What is the evidence? 

Fasten a candle to a stout copper wire and after lighting 
it thrust it into a bottle of air for a minute. Test for carbon 
dioxide in the bottle. Seek evidence of the formation of two 
compounds. (7) What do you conclude they are? (8) By 
what observation or experiment did you recognize their 
presence? (9) Of what elements do you conclude a candle to 
be a compound? 

BloAV through a tube into a test tube half full of lime water. 
(10) What is one of the constituents of the exhaled breath? 

(II) What is the evidence of this? (12) By what kind of a 
reaction was this constituent of the breath formed? (13) What 
was the source of the necessary oxygen? (14) What is the 
evidence of the production of heat in the formation of the 
compound? (15) Name two kinds of action taking place in 
nature in which carbon dioxide is formed without the 
phenomenon of combustion. 

{d) Mix in a 300 c.c. Erlenmeyer flask 20 c.c. of "Karo 
Syrup'* and 150 c.c. of water. Grind in a mortar a quarter of 
a yeast cake with 10 c.c. of Avater and pour the mixture into 
the flask. Fit into the flask a one-hole rubl)er stopper through 
Avhich passes a tube bent at right angles and connect this 
tube by means of a rubber connector with another right-angled 
tube which dips into 50 c.c. of lime water in a 250 c.c. bottle. 
Set the whole apparatus away in your desk until the next 
exercise, when you should answer the following questions. (1) 
What has been formed in the bottle of lime water? (2) What 
is shown to be one of the products of fermentation? (3) Is 
this a reaction of oxidation? (4) What brought about the 
fermentation? 



166 LABORATORY EXERCISES 

Experiment 41. Formation of Carbonates. 

Reference : 204. 

Discussion : As appears from the preceding experiments, 
calcium carbonate is formed when carbon dioxide is passed 
into a solution of calcium hydroxide. Similar reactions occur 
between carbon dioxide and the hydroxides of other metals. 
The reaction may be regarded as taking place in two steps, 
first the production of carbonic acid from the carbon dioxide 
and the water of the solution and then a double decomposition 
between the acid and the hydroxide forming the carbonate and 
water. When a solution of some other compound of the metal 
is substituted for the hydroxide, as for example the chloride, 
an acid will be formed by the double decomposition instead 
of water. But since, as has been shown, an acid always reacts 
with a carbonate, these two could not be formed and exist 
together. In other words a carbonate can not be prepared by 
passing carbon dioxide into a solution, if the other product 
of the reaction is an acid. 

Acid carbonates are compounds of both a metal and hydro- 
gen in combination with the carbonate radical. They are 
formed by the action of an excess of carbonic acid on a normal 
carbonate. This reaction, however, can be reversed by boiling 
off the carbonic acid as carbon dioxide and water. 

Materials: Solutions of calcium chloride and of calcium 
hydroxide, marble. 

Directions: (a) Bubble carbon dioxide from the generator 
you used in Expei'imciit 38 h into 10 c.c. of calcium chloride 
solution. (1) How does the result differ from that obtained 
when carbon dioxide is passed into calcium hydroxide? (2) 
Explain the cause of this difference. (3) How can calcium 
carbonate be prepared from calcium chloride? Try the method 
which you suggest. (4) What leads you to believe that the 
substance which you have obtained is calcium carbonate? 

(h) Pass carbon dioxide which has been purified from hydro- 
chloric acid, as in Experiment 38 b, into a test tube one-third 
full of lime water, until the precipitate first formed disappears. 
(1) What was the precipitate? (2) Into what compound was 



168 LABORATORY EXERCISES 

it converted? (3) Write equations for all the reactions taking 
place in the test tube. 

Boil the solution in the test tube until a distinct change 
is noticeable. (4) What appeared? (5) Write the equation 
for the reaction which took place. (6) What is the relation 
of this reaction to the one in which the precipitate disappeared? 
(7) AVhat is such a pair of reactions called? (8) Which is 
more soluble, calcium carbonate or calcium acid carbonate? 

Experiment 42. Coal. 

Eeference: 220. 

Discussion : Coal is the product of the slow decomposition 
of vegetable material in the absence of oxygen. Since this 
decomposition is never complete, coal is found to be a mixture, 
different samples varying in their composition. Upon heating 
coal many of these undecomposed compounds, which are hydro- 
carbons, are given off, some as permanent gases and others as 
easily condensed substances, which appear as a tar. Much 
of the nitrogen and sulphur in combination in the original 
vegetable material are converted into ammonia and hydrogen 
sulphide, which are evolved with the other gases when the 
coal is heated. The residue, which is called coke, consists of 
uncombined carbon and inorganic material, the latter being 
left as ash when the coke is burned in the air. 

Materials: Bituminous coal, litmus paper. 

Directions: Half fill a test tube with powdered bituminous 
coal and heat it vigorously. Smell the fumes, and when they 
are coming off vigorously, try to ignite them. Test the liquid 
which collects near the mouth of the test tube with pink litmus 
paper. (1) Describe the appearance and odor of the fumes. 

(2) Of what class of compounds are they chiefly composed? 

(3) Are they combustible? (4) Describe the appearance of 
the deposit at the mouth of the tube. (5) What is it called? 
(6) What Avas the action upon litmus paper? (7) What is 
thus proved to be given off from the coal? (8) What is the 
residue called which is left upon heating the coal? (9) What 
industrial process is illustrated by this experiment? 



170 LABORATORY EXERCISES 

Experiment 43. Flame and the Bunsen Bxirner. 

Reference: 230-231. 

Discussion : The distinction between amount of heat and 
intensity of heat, that is, temperature, must be kept clearly 
in mind. The same amount of heat is produced when the same 
amount of gas is burned no matter what kind of flame is pro- 
duced, but the temperature of the flame will be higher when 
the combustion takes place in a smaller space and when there 
is less inert gas to be heated. Thus a given amount of gas 
burning as a luminous flame produces the same amount of heat 
as it does burning as a Bunsen flame, but the latter is hotter, 
as the space where the burning takes place is smaller. Like- 
wise the same amount of heat is produced when gas burns 
in the air as is produced when it burns in pure oxygen, but 
the latter flame is hotter since no heat is consumed, as 
in the former case, in raising the temperature of the inert 
nitrogen. 

The preliminary admixture of air to the gas, affected by 
the Bunsen burner, serves the double purpose of increasing 
the temperature of the flame and of preventing the deposit of 
soot on objects heated. Soot is the carbon formed by the 
decomposition of the hydrocarbons of the gas where there is 
not enough oxygen present to consume it. In the luminous 
flame the only source of oxygen is the surrounding air, as no 
oxygen exists in the center of the flame to consume the carbon; 
but in the Bunsen flame the air drawn in at the bottom of 
the burner furnishes oxygen within the flame for this purpose 
and so no toot is formed. 

Although the preliminary addition of air to the gas is 
advantageous, the amount must be such that the mixture will 
not ignite in the tube of the burner itself. In a properly 
regulated burner a combustible mixture of gas and air is rising 
through the tube, but it does not reach its kindling temperature 
because the flow of the cold mixture upward offsets the con- 
duction of heat from the flame downward. When the latter 
is more rapid the mixture becomes ignited and the gas burns 
at the base of the burner. The rate of flow of the mixture, 
however, is not the only factor involved, for the addition of 



172 LABORATORY EXERCISES 

too much air may cause the flame to ''strike back," since the 
additional amount of oxygen produces a mixture in the tube 
which is explosive and which is ignited by the gas burning 
at the top of the burner. If something, like a metal grid for 
example, is placed in the top of the tube of the burner to 
conduct away the heat of the flame and thus prevent the 
mixture of air and gas from being raised in temperature, a 
more combustible mixture, one with more air in it, can be 
used. This will enhance the effect produced by an ordinary 
Bunsen burner and a flame of still higher temperature will be 
obtained. The Meker burner is constructed on this principle. 

In the Bunsen flame three distinct regions exist. Directly 
al)ove the bui-ncr is a cone-shaped space fllled with the mixtui'o 
which has just emerged from the burner but which has not yet 
been heated to the kindling temperature and is not burning. 
This is surrounded by gas burning in the air drawn in through 
the burner, but here the combustion is incomplete, because 
there is an insufficiency of oxygen and an excess of gas. This 
part of the flame is known as the reducing cone, because a 
decomposable oxide introduced into it will have its oxygen 
taken aAvay to aid in the combustion. This reducing cone is 
in turn surrounded by a region where an excess of oxygen 
derived from the air exists and Avhere an oxidizable sul)stance 
may take on oxygen. This part of the flame is the oxidizing 
cone. 

Materials: Fine copper gauze, pin, copper wire, Meker 
burner. 

Directions: (a) Take apart a Bunsen ])urner. Three 
separate pieces are obtainable. (1) Make drawings of the 
parts of the burner. 

(h) Put the burner together again, close the holes in the 
bottom of it, light the gas, and hold a glass rod in the upper 
part of the flame. (1) What is deposited upon it? (2) Was 
this substance present in the unburned gas as it issued from 
the cock? (3) Where did it come from? (4) What becomes 
of it when no object upon which it may be deposited is held 
in the flame? (5) Why does this action not take place within 
the flame? 



174 LABORATORY EXERCISES 

Admit air into the burner and again hold the rod in the 
flame. (^6) How does the result differ from that previously 
obtained? (7) Why? 

(c) Turn the gas on fully, close the air inlet and measure 
the height of the flame by holding a ruler beside it. Open the 
air inlet and again measure the height of the flame. (1) Record 
the heights. Hold a platinum wire at the tip of the green 
cone of the Bunsen flame and then, after closing the air inlet, 
in the luminous flame. (2) Which flame was the hotter? (3) 
What is the evidence? (4) What is the cause? (5) Was more 
heat produced in one flame than in the other? (6) Why? (7) 
What are the advantages of the Bunsen flame over the luminous 
flame? 

(d) (1) Make a drawing of the Bunsen flame. 

Hold a piece of glass tubing at an angle of 45 degrees with 
the lower end in the cone immediately above the top of the 
burner and apply a lighted match to the other end of the tube. 
(2) What occurred? (3) What does this prove about the com- 
position of this part of the flame? 

Stick a pin through a match close to the head. Turn off 
the gas and hang the match head upwards in the tube of the 
burner by means of the pin. Turn the gas on fully and light 
it. (4) What happened to the match? (5) What does this 
prove about the temperature of this part of the flame? 

Hold a bright stout piece of copper wire in the tip of the 
flame. (6) What color does it become? (7) What is the sub- 
stance which produces the color? (8) What does this prove 
as to the composition of this upper part of the flame? (9) What 
is this part of the flame called? 

(10) What is the color of the part of the flame directly 
outside the innermost cone? Hold the end of the copper wire, 
which you have just heated, in this part of the flame. (11) 
What change in color do you observe? (12) What reaction 
is taking place on the surface of the wire? (13) AVhat does 
this prove about the composition of this part of the flame? 
(14) What is this part called? 

By holding a platinum wire in various parts of the flame 
determine where the temperature is highest and indicate the 



176 LABORATORY EXERCISES 

place on your drawing. Label your drawing to show the com- 
position of the various parts of the flame and the names by 
which they are known. 

(e) (1) Make a drawing of the Meker burner. (2) What 
is in the top of the tube of the burner? (3) What is its use? 
(4) Does a larger or smaller proportion of air enter the Meker 
burner or the Bunsen burner? (5) Why does the construction 
allow a different proportion? Light the burner. (6) What is 
the size of the flame compared with that of the Bunsen 
flame? (7) Which flame would you expect to be hotter? (8) 
Why? (9) What is the fact, as shown by testing with a 
platinum ware? 

(/) Bring down upon the Bunsen flame a piece of fine 
copper gauze held in the tongs. (1) Where is the flame now? 
Hold a lighted match above the gauze. (2) What happened? 

(3) Was unburned gas passing through the gauze originally? 

(4) Why did it not burn? Turn off the gas and then turn 
it on again without lighting it. Hold the gauze about two 
inches above the top of the burner and apply a lighted match 
above the gauze. (5) Where is the flame? (6) What is below 
the gauze? Try some experiment to verify your answer. (7) 
What did you do, and what occurred? (8) Explain how the 
gauze acts to produce the results obtained in these experiments. 

Light the gas with the air inlet of the burner open. (9) 
Why is air drawn into the burner? 

Very slowly diminish the supply of gas until it takes fire 
at the base of the burner. (10) What is said to have occurred 
when this takes place ? (11) Explain why it happened? (12) 
How does the appearance of the flame issuing from the burner 
differ from that of the luminous flame? After the gas has 
burned at the base of the burner for two or three minutes, 
turn it off, immediately turn it on again about half Avay and 
attempt to relight it in the ordinary manner. (13) Why did 
it become ignited at the base? (14) When a burner accidentally 
*' strikes back," what should be done before attempting to 
relight the gas? 



ACIDS, BASES, AND SALTS 

Experiment 44. Properties of Acids and Bases. 

Reference: 234-241. 

Discussion : Solutions of a large number of substances fall 
into two very distinct classes, the members of each class having 
many identical properties but quite different properties from 
the members of the other class. The substances forming one 
class of solutions are termed acids and in all cases are com- 
pounds of hydrogen with an element the physical properties 
of which place it in the category of non-metals. Many com- 
pounds of this class contain oxygen also. The substances 
forming the other class of solutions are the hydroxides of the 
metals and are called bases. Among the bases, however, are 
included the hydroxides of all the metals, whether they form 
solutions or not. Those particular bases which are soluble are 
the alkalis. A solution of ammonium hydroxide, although not 
the hydroxide of a metal but the hydroxide of a radical, pos- 
sesses the characteristic properties of solutions of the alkalis. 
A member of one class of solutions will react w^ith a member 
of the other class to form water and a substance which belongs 
to a third class called salts. This type of reaction is termed 
neutralization. 

Materials: Dilute solution of acetic acid; of calcium, 
sodium, and potassium hydroxides; phenolphthalein, methyl 
orange, 10 per cent sodium hydroxide solution, woolen cloth, 
cotton cloth, silk, vinegar, sugar, litmus paper. 

Directions: (a) Put 5 c.c. of dilute hydrochloric, of dilute 
sulphuric, and of dilute acetic acids into each of three test 
tubes respectively, fill the test tubes with water, and shake 
them. Taste the solutions by putting a drop of each on the 
tip of your tongue with a glass rod. (1) Describe the taste 
of each. (2) In what experiments have you studied the action 
of acids upon metals? (3) What substance was formed in 

178 



180 LABORATORY EXERCISES 

3very case? (4) In Avhat experiments have you studied the 
action of acids upon carbonates? (5) What Avas formed? (6) 
Name three properties common to acid solutions. (7) What 
is the chemical composition of acids? 

{h) Taste a solution of calcium hydroxide. Dilute solu- 
tions of sodium hydroxide, of potassium hydroxide, and of 
ammonium hydroxide as you diluted the acid solutions in (a) 
and taste them. (1) Describe the taste of each. (2) In what 
experiments have you studied the action of carbon dioxide on 
this class of compounds? (3) What was formed? (4) To what 
large class of substances do these compounds belong? (5) To 
what subdivision of this class do they also belong? (6) What 
is the chemical composition of bases? 

(c) Add five drops of methyl orange to each of three test 
tubes one-third filled with water. Put five drops of sodium 
hydroxide solution into one test tube, of potassium hydroxide 
into the second, and of ammonium hydroxide into the third. 
(1) What colors did you obtain? (2) What is a test for an 
alkali? 

Add dilute hydrochloric acid to one of the test tubes, dilute 
sulphuric to another, and dilute nitric to the third until a change 
of color occurs in each case. (3) What colors Avere produced? 
(4) What is a test for an acid? 

Repeat these experiments with the substitution of five drops 
of phcnolphthalein for the methyl orange. (5) What color is 
phenolphthaloin in alkaline solution? (6) In acid solution? 

(7) What else have you used in testing for acids and alkalis? 

(8) Make a table consisting of the names of three indicators 
in the first column, their respective colors in acid solution in 
the second column, and their colors in alkaline solution in the 
third column. 

(d) Measure out in your graduate 12 c.c. of 10 per cent 
sodium hydroxide solution, pour 10 c.c. of it into an evaporating 
dish and the remaining 2 c.c. into a test tube nearly full of 
water. To the solution in the evaporating dish add dilute 
hydrochloric acid little by little Avith constant stirring. Test 
the solution AA^ith litmus from time to time by touching the 
end of your stirring rod to a piece of litmus paper held in 



182 LABORATORY EXERCISES 

your hand. Do not dip the paper into the solution. Do not 
lay the paper on the desk while making the test, as acid on 
the desk may contaminate the paper. Continue adding the 
acid until the solution in the dish gives an acid test. Now 
add some of the dilute solution of sodium hydroxide, which 
you have prepared in the test tube and which should be well 
shaken, until the contents of the dish is alkaline. Prepare a 
very dilute solution of hydrochloric acid by pouring 5 c.c. of 
the dilute acid into a test tube full of water. Pour this diluted 
solution into the dish little by little with constant stirring and 
frequent testing until the last small addition makes the con- 
tents of the dish acid. 

Support the dish on a wire gauze and evaporate the solu- 
tion to dryness. Toward the end of the evaporation when a 
solid begins to appear in the dish, the burner should be held 
in the hand and the flame Avaved under the dish, as otherwise 
the solid will spatter. When all the water has been driven 
off and the dish allowed to cool for several minutes, add 5 c.c. 
of water and again evaporate to dryness. (1) Describe the 
appearance of the substance left in the dish. Taste a little 
of it. (2) What familiar substance should you judge it to be? 

(3) Write the equation for the reaction by w^hich it was formed. 

(4) Was the evaporation necessary to bring about this reaction 
or had it already taken place before the evaporation? (5) 
What is this kind of reaction termed? (6) What classes of 
substances enter into this kind of a reaction? (7) What sub- 
stance is always one of the products? (8) To what class of 
substances does the other product belong? 

(e) (1) What is the effect of vinegar upon litmus paper? 
(2) What kind of a substance does it contain? 

To 10 c.c. of vinegar in a test tube add successive small 
amounts of sugar until it ceases to dissolve upon shaking and 
test the solution after each addition with litmus paper. (3) 
What was the result of the last test? (4) Is a sour taste still 
to be detected in the solution? (5) Did the sugar neutralize 
the acid? (6) What class of substances neutralizes acids? (7) 
Does sugar belong to this class? Verify your answer by an 
experiment. (8) What did you do and what was the result? 



184 LABORATORY EXERCISES 

Experiment 45. Quantitative Neutralization — Titration. 

Reference : 240. 

Discussion: The object of this experiment is to find the 
concentration, namely the number of grams of acid per liter, of 
a solution of hydrochloric acid by determining how many cubic 
centimeters of the solution are required to neutralize a definite 
volume of sodium hydroxide solution, the concentration of 
which is known. For example if the sodium hydroxide solu- 
tion was known to contain 20.00 grams of the compound per 
liter and if it were found that 10.00 c.c. of this solution was 
neutralized by 12.00 of hydrochloric acid solution, it is 
obvious that in the 10.00 c.c. of alkaline solution there was 

20.00 X Yfvw^ ^ 0-200 gram of sodium hydroxide and that in 

the 12.00 c.c. of acid solution there was present the weight 
of hydrochloric acid which would react with 0.200 gram of 
sodium hydroxide. This weight can be calculated by the 
familiar method of making a proportion between the molecular 
weights and the actual weights of the compounds. Having 
thus found the weight of hydrochloric acid in 12.00 c.c. of 
solution, the weight in 1000 c.c. is calculated by the aid of 
simple proportion. 

In the actual experiment the relative volumes of alkali and, 
acid required for neutralization are determined several times 
and the average of the results is taken as being more reliable 
than the value obtained fi'om a single experiment. 

Certain precautions must be taken to avoid errors in using 
a burette, {a) The burette must be washed clean, but enough 
water may adhere to the inside of it to dilute the solution. 
As drying the burette would be too tedious, it is rinsed out 
with some of the solution to be used. What adheres to the 
inside now is some of the solution itself, so no dilution occurs 
upon filling the burette. 

(&) A bubble of air may be caught in the tip of the burette 
or in the rubber connector above it. If this bubble of air comes 
out during the neutralization, some solution from the burette 
will flow down to take its place and it will appear that this 
amount of solution has been used, whereas in reality only air 



186 LABORATORY EXERCISES 

will have issued from the tip. All bubbles of air must there- 
fore be removed from the burette after it has been filled and 
before it is used. 

(c) In looking through the burette, the surface of the liquid 
appears as a curved line, which is called the meniscus. In 
reading the burette the position of the lowest part of the 
meniscus is taken as the height of the liquid. In observing 
its position on the scale on the burette, the eye must be on 
a level with the surface of the liquid. If the eye Avere too 
low the line of vision to the center of the meniscus would pass 
through a scale division on the burette beloAV the true level 
and if the eye were too high the reverse would be true. 

Materials: Special solutions of sodium hydroxide and hy- 
drochloric acid, phenolphthalein, two burettes. 

Directions: (a) Thoroughly wash out two burettes and 
clamp them to a ring stand in a strictly vertical position. Fill 
them with water and make sure that the tips do not leak. 
Remove the air bubbles in the tips by bending the tips upward 
and opening the pinch cocks slightly so that the water forces 
the air upward and out. (1) In actual use what error may 
be caused by allowing bubbles of air to remain in the tip? 
Practice the manipulation of the pinch cock, letting the water 
run out rapidly and then a single drop at a time. Read the 
height of the water in the burette. (2) What volume is in- 
dicated by the smallest division marks? With your eye five 
or six inches below the meniscus make a reading. Again make 
a reading with your eye an equal distance above the meniscus. 
(3) How great was the difference in readings? (4) Make a 
draAving to show the cause of this difference. (5) At what 
level should your eye always be in making a reading? (6) 
Record the proper reading. Run out 10 drops of Avater and 
(7) read again. (8) What is the volume of 10 drops? Deter- 
mine again the volume of 10 drops, and (9) record all the 
readings. (10) What do you find to be the average volume 
of a drop? (11) To what fraction of a cubic centimeter, there- 
fore, should you attempt to read a burette? (12) In making 
your record how many figures beyond the decimal point should 
appear? (13) If the reading chances to be a whole number 



188 LABORATORY EXERCISES 

of cubic centimeters, should ciphers be placed after the decimal 
point? (14) Why? 

{h) Let all the water run out of the burettes and after 
they have stood for three minutes run out the remaining drops 
which will have collected. 

Procure in clean dry flasks 75 c.c. of the special solution 
of hydrochloric acid and of sodium hydroxide prepared for 
this experiment. Pour 5 c.c. of the acid into one burette and 
5 c.c. of the alkali into the other. Take down in turn each 
burette, thoroughly rinse it with the 5 c.c. of solution by turn- 
ing it up and down, replace it in position and run out the 
solution. (1) What is the object of this rinsing? Fill the 
burettes with their respective solutions, remove air from the 
tips and (2) record the initial reading of each. Run into a 
clean beaker about 15 c.c. of sodium hydroxide solution from 
the burette, add three drops of phenolphthalein and then slowly 
run in acid from the other burette with constant stirring until 
the indicator loses its color. Now add enough alkali to restore 
the color and again enough acid to destroy it. Continue adding 
small amounts of one and then of the other solution until one 
drop only will produce the change in the indicator. Finish 
by the addition of one drop of alkali to give a faint pink. 
Read the burettes, (3) record the readings, and subtract the 
initial readings from these to obtain the amounts of solution 
used. (4) Divide the number of cubic centimeters of acid by 
the number of cubic centimeters of alkali. (5) What does this 
quotient mean? 

Repeat the experiment which you have just performed (6) 
recording the initial and final readings and making the same 
calculations as before. (7) Should you expect to get the same 
quotient as before? (8) Why? If your results differ from 
each other by more than 1 per cent continue making deter- 
minations until you can get two successive results which agree 
with each other to this accuracy. (9) What is the average of 
these accepted results? 

One thousand cubic centimeters of the alkali solution con- 
tains 20.0 grams of sodium hydroxide. (10) 1 c.c. of this solu- 
tion contains how many grams? (11) Calculate how many 



190 LABORATORY EXERCIiSES 

grams of hydrochloric acid are required to react with this 
weight of sodium hydroxide. (12) In how many cubic centi- 
meters of your solution is this weight of hydrochloric acid 
contained? (13) How do you know? (14) What weight of 
hydrochloric acid is contained in 1000 c.c. of your solution? 
(15) What was the object of the experiment? (16) What is 
the final answer? 

Experiment 46. Percentage of Acetic Acid in Vinegar. 
Relative Cost of Ammonium Hydroxide Solutions. 

Reference : 240. 

Discussion : The object of this experiment is the determina- 
tion of the percentage of acetic acid in vinegar. The experi- 
ment may be carried out as was Experiment 45 with the sub- 
stitution of vinegar for hydrochloric acid, as described in the 
text. It is more usual in such cases, however, to measure out 
in a pipette a definite volume of the solution to be analyzed, 
to add to it an excess of a base, and to complete the neutraliza- 
tion with a standard solution of some acid. The latter method 
is employed in this experiment since it gives practice in addi- 
tional manipulation and calculation. 

M.ATERiALS: Standai'd solutions of hydrochloric acid and of 
sodium hydroxide, vinegar, phenolphthalein, methyl orange, 
*' household ammonia," pipette, two burettes. 

Directions: (a) Set up and fill burettes with sodium hy- 
droxide and hydrochloric acid solutions, as in Experiment 45 
and (1) record the initial readings. 

Rinse out a 10 c.c. pipette with a little vinegar, fill it Avith 
vinegar to the mark and let the vinegar run out into a clean 
beaker. Allow the pipette to drain for half a minute and blow 
out the drop, which will have collected in the tip, into the 
beaker. Add 25 c.c. of water and three drops of phenol- 
phthalein. Run in slowly sodium hydroxide from the burette 
with constant stirring until the pink color of the indicator 
appears. Add enough hydrochloric acid from the other burette 
to destroy the color of the indicator and complete the neu- 
tralization as in Experiment 45. (2) Record the final readings 
of the burettes and calculate the volumes of each solution used. 



192 LABORATORY EXERCISES 

(3) Referring to the ratio of alkali to acid determined in 
Experiment 45, calculate how many cubic centimeters of alkali 
is equivalent to 1.00 c.c. of acid. (4) How many cubic centi- 
meters of alkali are equivalent to the volume of acid which you 
used in this experiment? (5) Subtract this volume from the 
total volume of alkali used. (6) Why is this subtraction made? 
(7) What does the remainder represent? Repeat the experi- 
ment, starting with another 10 c.c. of vinegar, until you obtain 
results which agree to within 5 per cent of one another. (8) 
What is the average of the accepted results? (9) Using this 
average and assuming that the sodium hydroxide solution con- 
tains 20.0 grams per liter, calculate the weight of the alkali 
used to neutralize the vinegar. (10) What weight of acetic 
acid (HC2H3O2) will react with this alkali? (11) What was 
the percentage of acetic acid in the vinegar? 

(&) Determine by a method similar to the one just used in 
(a) the relative concentrations of solutions of ** household am- 
monia" and the ammonium hydroxide used as a laboratory 
reagent. Use methyl orange instead of phenolphthalein as an 
indicator. (1) Make records of your burette readings and 
such calculations as are necessary in order to state the average 
volumes of acid required to neutralize 100 c.c. of each solution 
of ammonium hydroxide. Inquire from the instructor the cost 
per liter of each of the ammonium hydroxide solutions. (2) 
Which is cheaper for equivalent amounts of ammonia? (3) In 
what ratio? 



SOLUTIONS 

Experiment 47. Formation of Ions. 

Reference: 243. 

Discussion: When two samples of material have the same 
characteristic properties, the fact is summarized by saying that 
they are the same substance. Likewise if two mixtures exhibit 
a number of characteristic properties in common, it is con- 
cluded that they each contain some of the same substance. 
Such mixtures, with similar properties, are common in certain 
classes of solutions, as for example in solutions of acids, where 
similarity in taste, action with indicators, with metals, etc., 
have been noted. It appears, then, that all acid solutions are 
mixtures containing some of the same substance, but since 
these solutions were made from different acids, dift'erent sub- 
stances, they must have reacted to form some of the same new 
substance in every case. By no type of reaction previously 
studied could the facts be so well explained as by the assump- 
tion that every acid in solution breaks up into two electrically 
charged substances one of which is always hydrogen and that 
this is the substance which gives the similar properties to all 
acid solutions. This particular kind of reaction which produces 
electrically charged substances in solution is termed electro- 
lytic dissociation and the charged substances are called ions. 

In the same way the common properties of all alkaline solu- 
tions are accounted for by the formation of hydroxyl ions, 
of all solutions of chlorides by the formation of chloride ions, 
of all sulphates by the formation of sulphate ions, etc. A 
solution, thus, is a mixture of two or more ionic substances 
together with the original substance put into the solution, pro- 
vided the reaction of dissociation has not been complete. It 
does, not follow, however, that all substances react this way 
when dissolved, or that all compounds of hydrogen form hy- 
drogen ions, of chlorine, chloride ions, etc. 

194 



196 LABORATORY EXERCISES 

The reaction of dissociation is caused by the presence of 
the water; but when the water is again removed from a solu- 
tion, the reaction tends to become reversed and when finally 
all of the water has been taken away, only the original sub- 
stance remains. 

Materials: Litmus paper, glacial acetic acid, aluminium, 
solutions of sodium chloride, potassium chloride, magnesium 
chloride, silver nitrate, and silver sulphate ; potassium chlorate, 
sugar, alcohol, chloroform. 

Directions: (a) (1) What was the action on litmus paper 
of tlie solutions of acids studied in Experiment 44? Fill a 
perfectly dry test tube one-quarter full of glacial acetic acid 
(this is a pure substance and contains no water) and transfer 
a drop of it with a rod to a piece of blue litmus paper. (2) 
How does the result differ from that obtained with the acid 
solutions in Experiment 44? Add a drop of water to the acid 
on the litmus paper. (3) What happened? (4) What property 
is exhibited by a solution of acetic acid which is not possessed 
by the pure substance itself? (5) When acetic acid is dis- 
solved in water is any new substance formed? (6) What is 
the evidence? (7) Do other acids when dissolved in water 
form the same substance? (8) What is the evidence? (9) 
What is the name of the substance? (10) What is its chemical 
composition? (11) How is it essentially different from other 
substances previously studied? (12) Can it be obtained in a 
pure state out of solution? (13) What happens to it when a 
solution containing it is evaporated? (14) What is a test 
for it? 

Drop a small piece of aluminium into the glacial acetic 
acid in the test tube and warm it. (15) What happened? 
Add an equal volume of water and again warm it. (16) What 
difference in result is occasioned by the presence of the water? 
(17) Explain how this experiment substantiates the conclusion 
reached from the one with the litmus paper. 

{h) To 5 c.c. each of solutions of sodium chloride, of potas- 
sium chloride, and of magnesium chloride in separate test 
tubes add a few drops of silver nitrate solution. (1) What 
was formed in each case? (2) What common property was 



198 LABORATORY EXERCISES 

possessed by the three original solutions? (3) What substance 
in solution do they possess in common? 

Repeat the experiment with the substitution of silver sul- 
phate solution for silver nitrate solution. (4) Record the results 
of the experiments. (5) What substance besides water exists 
in both silver nitrate and silver sulphate solutions? (6) Write 
equations in the ionic foi,'m for the reaction which took place 
in all six experiments. Cross out in each equation the formulas 
which occur on both sides of the equation. (7) What remains 
in each case? (8) Did the same reaction or six different ones 
take place? 

Dissolve a few crystals of potassium chlorate and add some 
silver nitrate to the solution. (9) What occurred? (10) Do 
all salts containing chlorine form chloride ions? (11) What 
is the evidence? 

(c) (1) Referring to Experiment 44, what is the evidence 
that all alkaline solutions contain a common substance? (2) 
What is its name? (3) What is a test for it? 

{d) Test for the presence of the hydrogen ion in a solution 
of sugar (C12H22O11) ; of hydroxide ion in alcohol (C2H.5OH) ; 
and of chloride ion in chloroform (CHCl,). (1) (2) (3) State 
what you did, what you observed, and what you concluded in 
each case. 

Experiment 48. Electrical Conductivity of Solutions. 

Reference: 242-244. 

Discussion: Whether or not a given liquid is a good con- 
ductor of electricity can be determined roughly by attempting 
to pass through the liquid a current whicn has a lamp included 
in the circuit. If the lamp gloAvs a current is passing and 
the liquid is proved to be a good conductor; if the lamp does 
not glow the liquid is not a good conductor. 

In this experiment the conductivity of some pure liquids 
is tried, then that of some solutions. It will be found that all 
solutions do not behave alike in this respect, only those of 
certain classes of substances being good conductors. Upon 
this fact and certain other peculiarities of these same substances, 
which were studied in Experiment 47, the conception of elec- 
trolytic dissociation is based. 



200 LABORATORY EXERCISES 

Materials: Glacial acetic acid, distilled water, alcohol, 
solutions of acetic, of sulphuric, and of hydrochloric acids, of 
sodium and potassium hydroxides, of sodium chloride, of potas- 
sium nitrate, and of sugar, conductivity apparatus. 

Directions: The apparatus consists of two carbon rods 
about 5 mm. in diameter and 10 cm. long held in one-hole 
stoppers, Avhich are fitted into holes 2 cm. apart in a thin piece 
of wood 15 cm. long and 5 cm. wide. The upper ends of the 
rods are connected by binding posts with lead wires which are 
attached to terminals of a board carrying a socket and lamp 
which in turn is connected through a switch with a lighting 
circuit or other source of current. The piece of wood which 
carries the carbon rods is held in the clamp of a ring stand, 
whereby the rods may be lowered into a beaker containing the 
liquid to be tested. 

Caution : Be sure that the switch is left open, that is, so no 
current can pass through it, except when actually making the 
tests. 

Select the smallest beaker into which the rods of the ap- 
paratus can be lowered. Wash off the rods and wipe them dry 
(Handle them carefully as they are fragile). Fill the beaker 
one-third full of glacial acetic acid, (this is a pure liquid which 
contains no water) lower the rods into the acid, turn on the 
switch and observe whetlier or not the lamp glows. Tui-n off 
the switch. (1) Record the result of this and the following 
experiments by making a column of the names of the sub- 
stances tested and noting at the right of each the behavior of 
the lamp and the conclusion as to the conductivity of the liquid. 
Pour the acid into the bottle marked ** Acetic Acid Returned,'' 
wash the rods and the beaker, and then determine the con- 
ductivity of distilled water in the same way that you tried 
that of glacial acetic acid. Try the conductivity of alcohol 
and of the following solutions in turn: acetic acid, dilute 
sulphuric acid, dilute hydrochloric acid, sodium hydroxide, 
potassium hydroxide, sodium chloride, potassium nitrate, al- 
cohol, sugar. Ten cubic centimeters of each solution should 
be put into the beaker and enough water added to fill the 
beaker a third full. The carbon rods must be carefully rinsed 



202 LABORATORY EXERCISES 

after each trial. (1) Record the results of the experiments. 
(2) Are pure liquids conductors? (3) Are all solutions con- 
ductors? (4) What classes of substances form conductive 
solutions? (5) Is anything new formed when an acid, like 
acetic in this experiment, is put into water? (6) What is the 
evidence? (7) What is this kind of reaction called? (8) What 
is the general name of the substance formed? (9) What classes 
of substances act thus? (10) Does the reaction take place 
when the solution is made, or when the current passes? (11) 
What carries the current through the solution? 

Experiment 49. The Determination of the Heat of Neu- 
tralization. 

Reference : 245. 

Discussion : Neutralization is the interaction of an acid and 
a base, as the result of which a salt and water are formed. 
From the standpoint of the theory of elctrolytic dissociation, 
in neutralization the hydrogen ions furnished by the acid unite 
with the hydroxyl ions furnished by the base. If the solutions 
are sufficiently dilute, both the acid and the base are completely 
dissociated, and the only change that occurs when they are 
mixed is the formation of undissociated water from its ions. 
It follows that equal volumes of equivalent solutions of acid^ 
and of bases should produce the same amount of heat when 
neutralization takes place. The experiment described below 
is designed to test this conclusion. Equal quantities of 
normal solutions of several acids and bases are mixed and the 
h^at developed in each case is measured. This is done by 
determining the rise in temperature that occurs when the 
solutions are mixed. Since dilute solutions are used, it is 
assumed in the calculations that the specific heat of the result- 
ing salt solutions is equal to that of water, which is 1. If the 
volume of the solution is multiplied by the rise in temperature, 
the product is the number of calories set free. 

A part of the heat, however, will be used in raising the 
temperature of the surrounding air, the thermometer, and the 
vessel containing the solution. As a consequence, a correction 
must be made, as was done in Experiment 31. 



204 LABORATORY EXERCISES 

Materials: Normal solutions of potassium hydroxide, 
sodium hydroxide, hydrochloric acid, and sulphuric acid. 

(If the solutions are not furnished, they can be prepared 
by dissolving separately in 100 e.c. of water in each case 8.0 
grams of sodium hydroxide, 11.2 grams of potassium hydroxide, 
5.6 c.c. of concentrated sulphuric acid, and 16 c.c. of concen- 
trated hydrochloric acid, and then adding water to make the 
volume of each solution 200 c.c. The solutions should be cooled 
to room temperature before being used.) 

Directions: {a) Measure 50 c.c. of a normal solution of 
sodium hydroxide in a graduated cylinder and pour it into 
a 200 c.c. beaker. Place a thermometer in the solution and (1) 
record the temperature to a tenth of a degree. Measure 50 c.c. 
of the solution of sulphuric acid in the graduate and pour it 
into the beaker, and at once observe the thermometer. Stir 
and record to tenths of a degree the highest point reached. The 
temperature changes rapidly and the reading must be made 
promptly. Repeat the neutralization in the same way to obtain 
a second observation and average the results. 

(h) To determine the correction that must be added to the 
observed rise in the thermometer pour out the solution used 
in {a) above, wash the beaker with water at room temperature, 
and add to it 50 c.c. of Avater at this temperature. Insert the 
thermometer and (1) reeoi'd the temperature to tenths of a 
degree. Next heat in a small flask 50 c.c. of water to a tem- 
perature about twice as many degrees above the temperature 
of the cold water as the rise in temperature observed in (a) 
above. (The wai'm water will be from 12 to 14 degrees above 
room temperature.) Read the temperature of the warm water 
to tenths of a degree, place the thermometer in the beaker and 
at once pour the warm water into it and read the temperature 
of the mixture. (2) Calculate the temperature half-way be- 
tween the two extremes used and subtract from it the tem- 
perature observed Avhen the two portions of water were mixed. 
Repeat this part of the experiment and add the average of the 
results to the average temperature determined in (a). 

(3) Calculate the number of calories set free in the reaction 



206 LABORATORY EXERCISES 

(rise in temperature X volume) and from this value the number 
set free when 1 gram-molecular-weight of the base is used. 

(c) Determine in the way described in (a) above the rise 
in temperature when 50 c.c. of the hydrochloric acid solution 
is mixed with 50 c.c. of sodium hydroxide. (1) Calculate the 
molecular heat of neutralization. 

(d) Repeat using 50 c.c. of potassium hydroxide and 50 c.c. 
of sulphuric acid. (1) Calculate the results as before. 

(c) Repeat using 50 c.c. of potassium hydroxide and 50 c.c. 
of hydrochloric acid. (1) Calculate the results as before, and 
(2) obtain the average of all the experiments. (3) By what 
percentage does your result differ from the true value which 
is 13,600 calories? 

(4) Does the same reaction, or do different reactions take 
place between the different acids and bases? (5) How do the 
results of this experiment confirm your answer? (6) What is 
the reaction involved? 

Experiment 50. Electromotive Series. 

Reference: 248-252. 

Discussion : All metals tend to pass from the condition of 
the free element to that of ions in solution; they differ 
markedly in the strength of this tendency, which is called the 
solution pressure of the element. In order to pass into tho 
ionic condition, however, a metal must acquire positive elec- 
trical charges. These charges can be gained from the ions 
already in a solution of a salt into which the metal is put, 
provided the solution tension of the metal is greater than the 
solution tension of the metal which would be formed from the 
existing ions by the loss of their charges. In other words if 
the metals are arranged in the order of the magnitude of their 
solution pressures, any metal may form ions at the expense 
of any metal below it in the series and existing as ions in 
solution. Or conversely, by observing which metals will replace 
which others from solutions, the order of their relative solution 
tensions may be determined. Since hydrogen as well as the 
metals forms positive ions and may be replaced by certain 
metals, it too has a place in the series. 



208 LABORATORY EXERCISES 

Not only do the metals tend to form positive ions, but like- 
wise the non-metals tend to form negative ions and may be 
arranged in a series in accordance with the value of this ten- 
dency. These series are known as the electromotive series. 

Materials: Zinc, copper, solutions of copper sulphate, 
sodium chloride, mercuric chloride, potassium bromide, potas- 
sium iodide, iodine, bromine water, chlorine water, carbon di- 
sulphide. 

Directions: (a) Slip strips of clean zinc into test tubes 
containing solutions of copper sulphate and of sodium chloride. 
(1) What happened in each case? Slip a strip of clean copper, 
or a piece of clean heavy copper wire, into a solution of mer- 
curic chloride. (2) What occurred? (3) Arrange zinc, sodium, 
copper, and mercury in the order of their tendency to form 
ions. (4) What is this tendency termed? 

Try the action of dilute hydrochloric acid on zinc, and on 
copper. (5) What happened? Insert hydrogen in its proper 
place in the list of metals in (3). (6) What is the name of the 
series of which this list is a part? (7) Would you expect 
hydrochloric acid to act upon mercury? (8) Why? 

(h) As a preliminary to this experiment make yourself 
familiar Avith the color of iodine in a solution of carbon disul- 
phidc, and of bromine in the same solvent. Since both of these 
elements are more soluble in carbon disulphide than in water, 
the former will extract them from water solutions. Precaution : 
Do not use carbon disulphide near a flame, since it is inflammable. 
To 10 c.c. of water in a test tube add 1 c.c. of bromine water, 
2 c.c. of carbon disulphide and shake the test tube. (1) 
What color is bromine dissolved in carbon disulphide ? Repeat 
the experiment substituting a solution of potassium bromide 
for the bromine water. (2) What is the appearance of the 
carbon disulphide? (3) Will the bromide ion color carbon 
disulphide? (4) What does give it a color? Repeat the experi- 
ment using 1 c.c. of iodine solution instead of bromine water. 
(5) What color is iodine dissolved in carbon disulphide? 

Having now become familiar with the colors of bromine 
and of iodine in carbon disulphide, add 10 c.c. of water, 1 c.c. 
of chlorine water and 2 c.c. of carbon disulphide to 2 c.c. of 



210 LABORATORY EXERCISES 

potassium bromide in one test tube and to 2 c.c. of potas- 
sium iodide in another test tube and shake them. (6) 
What was formed in each case respectively? Repeat the ex- 
periment with the potassium iodide solution substituting for 
the chlorine water bromine water, but adding it only a few 
drops at a time and shaking after each addition. (7) What 
was formed? (8) Arrange bromine, chlorine, and iodine in the 
order of their solution pressures. 

Experiment 51. The Determination of Molecular Weights 
from the Freezing Points of Solutions. 

Reference: 254 (Read also 253). 

Discussion : The freezing point of a solution is always lower 
than that of the pure solvent. It has been shown as the results 
of many experiments that in the case of a substance which is 
not ionized (a non-electrolyte) the lowering of the freezing 
point is proportional to the concentration of the solution. The 
concentration is usually expressed as the number of grams of 
the dissolved substance contained either in 1 liter of the solu- 
tion or in 1000 grams of the solvent. The freezing point of 
a solution made by dissolving 1 gram-molecular-weight of any 
non-electrolyte in 1000 grams of water is —1.86° C. This fact 
makes it possible to determine the molecular weight of non- 
electrolytes which are soluble in water. The experiment 
described below shows how this can be done. 

The freezing point of a solution can be determined in two 
ways. Either the solution can be cooled luitil the temperature 
is reached at which ice begins to form, or ice can be added to 
the solution and the temperature noted at which the thermom- 
eter after falling, begins to rise. In the latter method the 
temperature observed is not the freezing point of the solution 
originally used, for as the temperature falls, some ice melts 
and the solution becomes diluted. The observed temperature 
is the freezing point of the solution present when the thermom- 
eter is read. In order to determine the concentration of the 
solution, the ice is separated from it by pouring the mixture 
through a funnel into a graduated cylinder, as soon as the 
reading of the thermometer has been made. The volume of the 



212 LABORATORY EXERCISES 

solution is noted and from this and the weight of the dissolved 
substance used, the concentration of the solution Avhich had 
the observed freezing point can be calculated. 

This last method is used in the experiment described below, 
because it is more readily carried out than the apparently 
simpler method of determining directly the freezing point of 
a solution of known concentration. 

The data needed for the calculation of the molecular weight 
are the weight of substance used, the volume of the solution 
when the temperature was read, and the freezing point 
observed. 

Materials: Sugar, ice, large funnel. 

Directions: (a) Before making the experiment the ther- 
mometer to be used must be tested to see if it registers correctly 
at the freezing point of water. Fill a 200 c.c. beaker full 
of ice, cracked in pieces approximately the size of a walnut. 
Wash the ice by adding water until it is covered; stir and 
pour off the water. Next just cover the ice Avith water and 
insert the thermometer so that the zero point is just above 
the surface of the liquid. Stir slowly and when the temperature 
is constant read the thermometer. When making the reading 
have your eye at the same level as the top of the column of 
mercury. Estimate the position of the mercury to tenths of a 
degree. (1) Record the freezing point and use this value in 
calculating later the lowering of the temperature in the follow- 
ing experiment. Pour off the water and keep the ice for the 
next part of the experiment. 

(h) Weigh to centigrams on a filter paper approximately 
30 grams of sugar. Pour the sugar into a beaker, add 50 c.c. 
of water, and stir until the sugar has dissolved. Add next 
enough cracked ice so that some of the pieces rest on the bottom 
of the beaker. The pieces should be of such a size that after 
melting somewhat they will not pass through the funnel to 
be used later. Stir slowly w^ith the thermometer. At first the 
temperature will fall and then it will slowly rise. At this 
point read the thermometer to the nearest degree and estimate 
to tenths of a degree. In reading the thermometer be sure 
that your eye is at the same level as the top of the column 



214 LABORATORY EXERCISES 

of mercury. Immediately pour the contents of the beaker into 
a funnel, large enough to hold all the ice, supported in a 
100 e.e. graduated cylinder. (If only a 50 c.c. graduated 
cylinder is available, the funnel may be supported in a flask, 
and the volume of the solution subsequently measured.) Re- 
move the funnel, and read the volume of the solution. 

(1) How many grams of sugar were in your solution? (2) 
Since the depression of the freezing point is proportional to 
the concentration, how many grams would be necessary to 
make the solution freeze at —1.86° C. (3) How many grams 
of sugar would be contained in 1000 c.c. of such a solution? 
(4) Why is this number the molecular weight of sugar? (5) 
Calculate the theoretical molecular weight of sugar from its 
formula which is C12H22O11. (6) What was the percentage 
error in your determination? 

Experiment 52. The Determination of the Degree of Dis- 
sociation of Electrolytes from the Freezing Point of Solutions. 

Reference: 253-254. 

Discussion : When an electrolyte dissolves in water the 
solution contains, in addition to the undissociated compound, 
ions, the number of which depends upon the concentration of 
the solution. The number of particles present which affect 
the freezing point of the solution is, consequently, greater than 
in the case of a non-electrolyte. If a salt, sodium chloride 
for example, yields two ions when it dissociates, its effect upon 
the freezing point of water will be twice that of a non-electro- 
lyte at the same concentration, provided it is completely dis- 
sociated ; if it is but partly dissociated the effect will be less. 
From this effect it is possi])lo to calculate by the following 
method the extent of dissociation of the salt. 

If we consider 100 molecules of sodium chloride, and let a 
represent the number of these molecules which are ionized, then 
evidently a is the per cent ionized and 100 — ^ is the per cent 
not ionized. Since sodium chloride yields two ions, the number 
of ions is 2a. The total number of particles in the solution is 
the sum of the un-ionized molecules and the number of ions, 
which in this case is (100 — a) -\- 2n or 100 + a. The lowering 



216 LABORATORY EXERCISES 

of the freezing points of two solutions having the same volume 
are proportional to the number of particles present in the solu- 
tions. Consequently, 

100 lowering of freezing point if no ionization 
100 -f- a lowering of freezing point observed 

To calculate the extent of dissociation it is only necessary to 
substitute in the formula above the values of the expressions 
written out in words and solve for a. 

If a solution of sodium chloride is studied, the lowering of 
the freezing point observed is the quantity experimentally 
determined. The lowering of the freezing point if no ioniza- 
tion occurred is the value calculated from the freezing point 
of a solution of a non-electrolyte of the same molar concentra- 
tion as that used in the experiment. From the weight of the 
salt used and from the volume of the solution the molar con- 
centration is known. 

Materials: Pure sodium chloride, ice, large funnel. 

Directions: Determine the reading of your thermometer 
in pure ice and water. Weigh to centigrams on a piece of 
filter paper approximately 15 grams of sodium chloride. Pour 
the salt into a 200 c.c. beaker and dissolve it in 50 c.c. of 
water. Add cracked ice until some of it rests on the bottom 
of the beaker and stir slowly with a thei^mometer. At fii'st 
the temperature of the mixture will fall; when it begins to 
rise (1) record the temperature to tenths of a degree and 
immediately pour the mixture into a large funnel supported 
in a 100 c.c. graduated cylinder. Remove the funnel, and (2) 
record the volume of the solution. 

(3) How many grams of salt did your solution contain and 
in what volume? (4) What weight of salt would be contained 
in 1000 c.c. of a solution of this concentration? (5) What is 
the molecular weight of sodium chloride? (6) How many 
times molar was the concentration of your solution? (7) What 
is the freezing point of a molar solution? (8) What would 
have been the freezing point of your solution had no dissocia- 



218 LABORATORY EXERCISES 

tion occurred? (9) What was the observed freezing point? 
(10) Substitute these last two values in the equation given in 
the discussion and solve for a. (11) What does a represent? 
(12) What is the final result of this experiment? 



CHEMICAL EQUILIBRIUM 

Experiment 53. Reversible Reactions. 

Reference: 260-262. 

Discussion: If two substances, A and B, react and form 
the substances C and D, and if the latter can interact to form 
the former, the reaction involved is said to be reversible. This 
fact is represented by writing the equation for the reaction in 
a general form as follows: 

A^B^C -\-D 

IC any one of the substances involved in the reaction is con- 
stantl}^ removed as the reaction proceeds, the reaction runs to 
completion in one direction. In the general case given above if 
A or B, for example, is removed all of C and D are transformed 
into A and B; likewise if or P is removed A and B are com- 
pletely converted to C and D. 

The experiment described below is designed to illustrate 
the effect of the removal of one of the products of a reversible 
reaction. The reaction studied is that represented by the fol- 
lowing equatiodi : 

NaCl + H,S(), .^ NaHSO, + HCl 

To force the reaction in the direction indicated from left to 
right in the equation, diy sodium chloride is treated with con- 
centrated sulphuric acid. Under these conditions hydrogen 
chloride, which is a gas, escapes and all the salt is changed 
to sodium hydrogen sulphate. To force the reaction in the 
opposite direction an aqueous solution of sodium hydrogen 
sulphate is treated with concentrated hydrochloric acid. Since 
sodium chloride is only sparingly solubk^ in strong hydrochloric 

220 



222 LABORATORY EXERCISES 

acid, it is formed and separates out as a precipitate. As a 
result, the sulphate is changed to chloride. 

Materials: Sodium chloride, concentrated hydrochloric 
acid, solution of ferric chloride, ammonium chloride, ammonium 
thiocyanate. 

Directions: (a) (iJoocZ) To 10 grams of dry sodium chloride 
contained in a small beaker add 13 c.c. of concentrated sul- 
phuric acid. In 2 or 3 minutes place the beaker on a wire 
gauze, cover it with a watch glass, and heat it with a small 
flame that just touches the gauze. Blow across the top of the 
beaker. (1) What happened? (2) What is thus shown to be 
escaping from the beaker? 

When all the salt has dissolved and the beaker contains 
a clear liquid remove the flame, let the substance cool until it 
crystallizes, and observe the appearance of the crystals. Next 
add 20 c.c. of water and heat cautiously until the salt has 
dissolved. Cool the beaker to room temperature, and after 
pouring about half of the solution into a test tube, for future 
use in (h), place the beaker in cold water. When crystals 
have formed in the beaker, filter them off and (3) describe 
their form. Scrape the crystals into a watch glass and add 
to them a few drops of concentrated sulphuric acid. (4) Was 
any hydrochloric acid given off? (5) How do you know? (6) 
AVere the crystals sodium chloride? (7) What is the evidence? 
(8) What then Avere the crystals? (9) Write an equation for 
the original reaction. (10) Why did it take place? 

(h) Add to the solution in the test tube saved from (a) 
an equal volume of concentrated hydrochloric acid. (1) What 
happened? Filter off the precipitate and press out as much 
of the liquid as you can with a spatula. Remove the filter 
paper, open it out flat, and press it and the salt between filter 
papers. Remove the salt, place it on a dry piece of paper, 
and set it aside. Dissolve a pinch of the salt in a few drops 
of water on a watch glass and set it aside. (2) Describe the 
appearance of the crystals formed when the water has 
evaporated. 

Using the main portion of the solid on the filter paper, 
determine whether or not it is sodium chloride. (3) What 



224 LABORATORY EXERCISES 

did you do? (4) What was the result? (5) What do you 
conclude? (6) Write an equation for the formation of the pre- 
cipitate? (7) What is the relation of this reaction to that in 
(a)? (8) Why did this reaction take place? (9) Make a 
general statement as to the factors which cause a change in 
the equilibrium in a reversible reaction. 

(c) To 20 c.c. of water in a test tube add exactly from a 
glass tube 5 drops each of solutions of ferric chloride and of 
ammonium thiocyanate. Pour a quarter of the solution into 
each of three other test tubes and to one add 2 c.c. of ferric 
chloride, to another 2 c.c. of ammonium thiocyanate and to a 
third 5 c.c. of ammonium chloride solution. Compare the colors 
of these three solutions with that of the fourth which was 
kept as a standard of reference. (1) How did each of the three 
solutions change in color? (2) Write an equation for the 
equilibrium existing upon adding ferric chloride to ammonium 
thiocyanate. (3) What is the colored compound? (4) Explain 
the causes of the changes in color upon adding the ferric 
chloride, the thiocyanate, and the ammonium chloride respec- 
tively. 

Experiment 54. The Determination of the Effect of In- 
crease in Temperature on the Rate of a Reaction. 

Reference : 265. 

Discussion : The rate at which most reactions take place 
is approximately doubled for an increase in temperature of 
ten degrees. In the experiment described below determinations 
at two temperatures are made of the time required to reduce 
a solution of potassium permanganate by means of iron and 
sulphuric acid. Potassium permanganate has a purple color 
and is converted by reduction into a colorless compound. All 
the conditions except the temperature are kept constant by 
using equal volumes of the same solution in test tubes of the 
same size ; the iron is used in the form of nails so that the 
surface of the metal in each case is the same. 

The results are to be plotted on Fig. 11, expressing 
time on the vertical lines (abscissae) and the tempera- 
ture on the horizontal lines (ordinates). The theoretical 



226 



LABORATORY EXERCISES 



curve for the relation of time to temperature is also drawn, 
as based on the assumption that the speed of the reaction is 
..A 



^<J 
































































































































30 






























































































































































20 






























































































































































10 
































































































































































































B C D E 

Tempera+ure in Degrees 
Fig. 11. 

doubled for an increase of temperature of ten degrees. The 
observed values may then easily be compared with the 
theoretical. 



228 LABORATORY EXERCISES 

Materials: A solution is furnished which is made by dis- 
solving 1 gram potassium permanganate in 1 liter of water 
and adding 200 c.c. concentrated sulphuric acid. Two wire 
nails 21^ inches long, watch with second hand. 

Directions: (a) Measure accurately in a graduate 10 c.c. 
of the solution of sulphuric acid and potassium permanganate 
and pour it into a dry test tube. Place this tube in a large 
beaker filled with water at room temperature. (1) Note the 
temperature of the water to tenths of a degree, Avhich should 
be the same as that of the solution in the test tube. (2) Note 
the time to seconds as a nail is carefully dropped into the test 
tube. 

While the reaction is taking place, measure accurately 10 
c.c. of the acid solution into a dry test tube and place the 
latter in a beaker full of water. Heat the w^ater until its 
temperature is about 25 degrees above room temperature, re- 
move the beaker from the source of heat, (3) note to tenths 
of a degree the temperature of the acid solution, and (4) record 
the time when a nail is placed in the test tube. 

(5) Record the time when each of the acid solutions loses its 
color. (6) Note the temperature of the warm solution when its 
color disappears and use the mean of this temperature and 
that observed at the beginning of the experiment as the tem- 
perature of this solution. 

(5) To prepare the theoretical curve for the relation of 
speed of the reaction to time, proceed as follows using the 
diagram on Fig. 11. (1) Label the point B the tempera- 
ture of the first tube (room temperature), and label the 
points 0, D, and E, which are respectively 10, 20, and 30 degrees 
above room temperature. !Make a cross on the line rising from 
B (the line BA) to indicate the time of the reaction at room 
temperature. In choosing the number of divisions which you 
will let represent one minute, select such a number as will bring 
the cross near the upper end of the line. On the verti^^al line 
rising from C make a cross half as high above the base line 
as the cross on the line BA, since at temperature C the reaction 
will require only half as much time as before. On the line 
rising from D make a cross at half the height of that above C, 



230 LABORATORY EXERCISES 

and on the line rising from E a cross at half the height of that 
above D. Draw a smooth curve through the crosses. (2) What 
does this curve represent? 

Mark with the letter F the position on the base line cor- 
responding to the higher temperature at which you carried out 
the reaction. On the line rising from F make a cross at the 
height corresponding to the time which elapsed in the experi- 
ment. (3) If the speed of this reaction had increased exactly 
at the rate of doubling for every ten degrees rise, where would 
this cross above F have fallen in relation to the curve? 



SULPHUR AND SULPHIDES 

Experiment 55. AUotropic Forms of Sulphur. 

Reference: 273. 

Discussion : Sulphur can exist in two different crystalline 
forms and two different liquid forms, one of which when sud- 
denly cooled becomes so stiff that it closely resembles a solid. 
The variety of sulphur which is rhombic in crystalline form 
is stable up to 96° and the other variety, the monoclinic, 
from this temperature to its melting temperature 119°. The 
rhombic sulphur can be heated temporarily above 96°, and 
the monoclinic sulphur can be cooled below this temperature, 
because the rearrangement of the material of a solid takes 
place slowly; eventually, however, all sulphur below 96°, 
the ''transition temperature," becomes rhombic and above this 
temperatui'e monoclinic. 

In prepai'in^' ci'ystals of any substance, one of two methods 
is used; either the substance is melted and then slowly cooled, 
when crystals are formed, as in the case of ice, or the substance 
is dissolved and then the solvent evaporated, when the matei'ial 
may come out of solution as ci'vstals, as in the case of crystalliz- 
ing sugar from water solution. In the case of rhombic sulphur 
the fii'st metliod can not be used, because crystallization must 
tak(* place below the ti'ansition tempei'atui'c, Avhich in turn is 
below the solidification temperature of liquid sulphur. The 
second method must therefore be resorted to and evaporation 
cari'ied on below 96°. Foi' preparing monoclinic crystals the 
first method is suitable, since this form is stable directly below 
the solidification temperature of the liquid. 

All liquids become more viscous upon cooling. If some 
liquids are suddenly cooled below their solidification tempei-a- 
ture they become so stiff an^ rigid that they appear solid and 
may remain in this condition for a long time, since the 

232 



234 LABORATORY EXERCISES 

material does not easily arrange itself in a crystalline or true 
solid form because of the great viscosity. Glass is such a super- 
cooled liquid. The dark colored liquid sulphur, Avhen suddenly 
cooled, affords another example of this phenomenon and is thus 
often referred to as a third solid amorphous form. 

Materials: Roll sulphur, flowers of sulphur, carbon disul- 
phide. 

Directions: (a) Heat a test tube one-third full of sulphur 
slowly until it boils, observing changes in color and viscosity 
during the process, and then pour the boiling sulphur into an 
evaporating dish filled with cold water (1) What was the 
color of the sulphur just as it melted? (2) What was 
the color upon further heating? (3) What changes in viscosity 
occurred? (4) How many varieties of liquid sulphur exist? 

Examine the sulphur which was poured into the water. (5) 
What is its most striking physical property? (6) AVhat is this 
form called? Preserve it untir the next exercise and then 
answer the following questions. (7) How has it changed in 
appearance? (8) As what form does it now chiefly exist? (9) 
Why did it not assume this form immediately? 

(h) Fill a test tube nearly full of roll sulphur and melt it 
very slowly at as low a temperatui-e as possible, not allowing 
it to become dark. Pour the melted sulphur into a filter paper 
folded as for filtering, but held by its edge in your hand instead 
of in a funnel as usual. As soon as crystals begin to form in 
the filter paper, pour the remaining melted sulphur quickly 
into a dish of cold water. Unfold the filter paper and examine 
the crystals. (1) Record the properties of the crystals, es- 
pecially their shape, color, luster, and brittleness. (2) What 
is this form of sulphur termed? (3) Why is it so named? (4) 
Between what temperatures is it stable? (5) At about what 
temperature is it now? (6) How do you account for its 
existence at this temperature? Preserve it until the next 
exercise and then answer the following questions. (7) How 
has it changed in appearance? (8) As what form does it now 
exist? (9) Why did it change? 

(c) To 3 grams of flowers of sulphur (powder) in a test 
tube add about 5 c.c. of carbon disulphide. (Remember the 



236 LABORATORY EXERCISER 

precaution to be observed in using this compound, Experi- 
ment 50). Shake the tube until all the sulphur is dissolved, 
then pour the clear solution into an evaporating dish to 
crystallize in the hood where there is no flame. Cover the dish 
with a piece of filter paper to retard evaporation. Watch the 
crystallization toward the end, and, if perfect crystals form, 
remove them with the forceps. Allow the liquid to evaporate 
almost entirely, then remove and dry the crystals. (1) Ex- 
amine them as in (?;) and record their properties. (2) What 
is the name of this kind of sulphur? (3) Can it be formed 
directly from melted sulphur? (4) Why? (5) Name all the 
different forms of sulphur. (6) Which is the stable form at 
room temperature? 

Experiment 56. Preparation of Hydrogen Sulphide. 

Reference: 277. 

Discussion : As was pointed out in Expei-iment 53, hydro- 
chloric acid, or as it may be called hydrogen chloride, was 
formed by the double decomposition between a chloride and 
an acid, under such conditions that the hydrogen chloride could 
escape. In a similar way hydrogen sulphide may be prepared 
by double decomposition between a sulphide and an acid. Since 
hydrogen sulphide is only slightly soluble in water, the reaction 
may take place in <liliit(> solution, so that a solution of hydro- 
chloric acici may ])e used as the necessaiy acid. The sulphides 
of many metals are so ins()lul)le that this double decomposition 
takes place chiefly in the reverse dii'ection and so th(>se sul- 
phides can not be us(m1 for the preparation of hydi'ogen 
sulphide. 

Materials: Iron sulphide, lead acetate. 

Directions: {All experiments inili hydrogen sulphide gas 
should he performed in the hood.) 

Set up an apparatus similar to the one used for the prepara- 
tion of chlorine in Experiment 21. Put 15 grams of iron sul- 
phide in the test tube and pour in enough dilute hydrochloric 
acid to seal the end of the thistle tube. Cas should begin to 
be evolved at once and when the evolution slackens a little 
concentrated hydrochloric acid may be added, a few cubic 



238 LABORATORY EXERCISES 

centimeters at a time. Collect the gas, which is heavier than 
air, in dry bottles and determine when each bottle is full by 
testing with a piece of filter paper which has been dipped into 
a solution of lead acetate. Hydrogen sulphide turns lead com- 
pounds black. When two bottles of gas have been collected, 
allow the hydrogen sulphide to bubble into half a bottleful of 
v/ater for ten minutes and then empty the contents of the 
generator into the bottle in the hood provided for the purpose. 
(1) Write the equation for the reaction by Avhich hydrogen 
sulphide was prepared. (2) To what class of reactions does 
it belong? (3) Why does it take place to completion? (4) 
Why do not all sulphides act like iron sulphide? (5) Write 
the equation for the reaction taking place in testing for hydro- 
gen sulphide. 

Experiment 57. Properties of Hydrogen Sulphide. 

Reference: 280-283. 

Discussion : In the l)urning of hydrogen sulphide appear 
analogies between the hydrogen compounds of sulphur and of 
carbon in that both form water and the dioxide of the non- 
metal, and that both are unstable at the temperature of the 
flame, and deposit the uncombined element when burning in 
an insufficient supply of oxygen. 

In solution the acidic i)roperties of hydrogen sulphide appear 
and under the name of hydrosulphuric acid the compound may 
be compared with hydrochloric acid. 

Materials: Splinter, concentrated nitric acid, litmus paper, 
lead acetate, copper sulphide, zinc. 

Directions: (a) Thrust a burning splinter into a bottle of 
hydrogen sulphide and at once invert a dry bottle of air over 
it. (1) What happened? (2) What was deposited on the 
sides of the bottles? (3) Is hydrogen sulphide stal)le at the 
temperature of the flame? (4) What is the evidence? (5) 
Write the equation for the complete oxidation of hydrogen 
sulphide. 

{!)) Into another bottle of the gas pom* a few di'ops of 
concentrated nitric acid, cover \ho liottlo and shake it gently. 
(1) AVhat insoluble substance is formed? (2) By what kind 



240 LABORATORY EXERCISES 

of reaction was it formed? (3) What kind of an agent is 
nitric acid? (4) What kind of a reaction did the nitric acid 
undergo? (5) What kind of an agent is hydrogen sulphide? 

(c) (1) Is hydrogen sulphide soluble? (2) What is the 
evidence? Examine the bottle on the table which was filled 
with a solution of hydrogen sulphide but has been exposed 
for some days to sunlight and air. (3) What insoluble sub- 
stance is in the bottle? (4) Does the solution smell of hydrogen 
sulphide? (5) Is hydrogen sulphide stable in solution? 

(d) Test the solution of hydrogen sulphide, which you pre- 
pared in Experiment 56, with blue litmus paper. (1) What 
does the test show? (2) Why is the action less marked than 
it would be in a solution of hydrochloric acid of equal concen- 
tration? (3) Is hydrogen sulphide a weak or a strong acid? 

To 10 c.c. of hydrogen sulphide solution in a test tube add 
a few drops of lead acetate. (4) What did you observe? (5) 
For what ion is this experiment a test? (6) Why can not 
this procedure be used directly as a test for an insoluble 
sulphide? (7) How may an insoluble sulphide which will react 
with an acid be tested for? (8) In what recent experiment 
have you done this? 

Put into a test tube a quarter of a spatulaful of copper 
sulphide, a strip of zinc, and 10 c.c. of hydrochloric acid. Heat 
the test tube and test for hydrogen sulphide at its mouth. (9) 
How did you test? (10) What was the result? (11) Write 
equations for the reactions which took place in the test tube. 
(12) State the test for a sulphide insoluble in acids. 

Experiment 58. Sulphides. 

Reference : 282, 283. 

Discussion: The sulphides of the metals fall into three 
classes, those soluble in water, those which are insoluble but 
react with dilute acids and are thus said to be soluble in acid, 
and those which do not react with dilute acid. When hydrogen 
sulphide is passed into the solution of a salt of the first class, 
of course no precipitate is formed; when it is passed into a 
solution of a salt of the second class only part of the metal 
is precipitated as sulphide, because the reaction is reversible. 



242 LABORATORY EXERCISES 

since sulphides of this class react with acids formed by the 
double decomposition; and when it is passed into the third 
class, complete precipitation of the sulphide takes place. For 
complete precipitation of members of the second class a soluble 
salt of hydrogen sulphide must be used in order that no acid 
shall be formed. These facts are made use of in separating 
the metals in the course of analysis of salts. 

Materials: Solutions of sodium chloride, calcium chloride, 
zinc sulphate, ferrous ammonium sulphate, copper sulphate, 
arsenic chloride, hydrogen sulphide, ammonium sulphide. 

Directions: Put 5 c.c. of solutions of sodium chloride, 
calcium chloride, zinc sulphate, ferrous ammonium sulphate, 
copper sulphate, and arsenic chloride into separate test tubes 
and add 10 c.c. of hydrogen sulphide solution, (either some 
that you made in Experiment 56, or some of the laboratory 
reagent) into each. 

(1) Tabulate the results of the experiment by writing the 
names of the compounds in the first column and recording the 
appearance or non-appearance of precipitates in the second 
column. (2) AVi'ite equations for the formation of precipi- 
tates in those cases where they were formed. 

Add hydrochloric acid to the tubes containing precipitates 
until the precipitate disappeai-s or the test tube is full. Record 
the results in the table in (1). (3) What is the distinguishing 
characteristic of each of the three classes of sulphides. (4) 
To which class does each of those just studied belong? 

To 5 c.c. of zinc sulphate solution add 10 c.c. of hydrogen 
sulphide solution, filter, and add 10 c.c. of ammonium sulphide 
to the filtrate. (5) What Avas the precipitate finally obtained? 
(6) Did the hydrogen sulphide completely precipitate the zinc 
in the first place? (7) Why? (8) How may zinc be completely 
precipitated as sulphide? (9) How may precipitation of zinc 
sulphide be entirely prevented when hydrogen sulphide is 
added? 



OXIDES AND ACIDS OF SULPHUR 

Experiment 59. FormatioD of Sulphur Dioxide. 

Reference: 287-288. 

Discussion : Sulphur dioxide in many respects resembles 
carbon dioxide, and like carbon dioxide it may be prepared by 
the spontaneous decomposition of the acid of which it is the 
anhydride, and by the oxidation of the element itself. But 
unlike carbon dioxide it can also be prepared by the reduction 
of sulphuric acid, to Avhich there is no analogous compound 
of carbon. The reduction may be considered as the removal 
of oxygen from the sulphuric acid by the reducing agent and 
the subsequent decomposition of the resulting sulphurous acid. 

Materials: Powdered sulphur, sodium sulphite, copper, 
powdered charcoal. 

Directions: (a) Touch a bit of sulphur with a hot platinum 
w4re and bring the sulphur thereby into the flame. (1) Describe 
the odor of the substance formed. (2) Write the equation for 
the reaction Avhich takes place when sulphur burns in the air. 
(3) What other element reacts similarly to give an analogous 
compound? 

(h) To a few crystals of sodium sulphite in a test tube 
add dilute sulphuric acid and cautiously smell the «>as which 
is evolved. (1) What is it? (2) Write two equations to ex- 
press how the gas was formed. (3) Why does this reaction 
take place? (4) What analogous reaction takes place for the 
same reason? 

(c) Heat in a test tube half a spatulaful of copper turnings 
with 5 c.c. of concentrated sulphuric acid and cautiously smell 
the gas which is formed. (1) What is it? (2) AVhat element 
did the copper remove from the sulphuric acid? (3) AVhat 
compound would this leave? (4) Into what substances does 
this compound spontaneously break up? (5) How would you 

244 



246 LABORATORY EXERCISES 

expect the copper compound to react with more sulphuric acid? 

(6) Why do you predict this action? (7) Write equations 
for these various reactions and then combine them into one 
equation. (8) As what kind of an agent did the copper act? 

(9) The sulphuric acid? (10) What Avas the product of reduc- 
tion? (11) What was the product of oxidation? (12) What 
reaction of double decomposition took place? 

(d) Repeat experiment (c) with the substitution of 
powdered wood charcoal for the copper. (1) What gas did 
you recognize as being formed? (2) How did you identify it? 
(3) From what compound was it formed? (4) By what kind 
of a reaction? (5) What was oxidized? (6) What was the 
product of the oxidation? (7) Write equations for the reac- 
tions taking place and then combine them into a single equa- 
tion. (8) What striking difference exists between the final 
products in this case and those in the action between copper 
and sulphuric acid? (9) In what respect are the two reactions 
similar ? 

(e) Repeat experiment (c) with the substitution of powdered 
sulphur for the copper. (1) What gas Avas formed? (2) What 
was the final product of reduction of the sulphuric acid? (3) 
What product did the sulphur form? (4) Write equations for 
the separate reactions and then combine them into a single 
equation. (5) Which Avas oxidized and Avhich reduced? (6) 
Which was the reducing agent, Avhich the oxidizing agent? 

(7) What always happens to a reducing agent? (8) To an 
oxidizing agent? (9) Which of the five methods of formation 
of sulphur dioxide, Avhich you have just tried, do you consider 
best adapted for preparing the gas in the laboratory for study? 

(10) Why? 

Experiment 60. Preparation of Sulphur Dioxide. 

Reference: 287. 

Discussion: As Avas shoAvn in Experiment 59, sulphur 
dioxide can be prepared by the action of an acid on a sulphite. 
A concentrated solution of sodium acid sulphite, however, is 
much cheaper than sodium sulphite and is usually used in its 
place. Since a liquid is employed instead of a solid, a slight 



248 



LABORATORY EXERCISES 



change in the form of the gas generator becomes necessary. 
The presence of the water in the generator requires the use 
of concentrated sulphuric acid, which is much heavier than 
the solution. If it were poured down the thistle tube of an 
apparatus like the one used for preparing carbon dioxide, it 
would lie as a nearly distinct layer in the bottom of the test 
tube and not become mixed with the solution of sodium acid 
sulphite. It is therefore necessary to arrange the apparatus 
so that the acid will fall through the solution drop by di-op. 



^ 



r 



I 



'7K 



1^3 



^D 



Fig. 12. 



Materials: Concentrated solution of sodium acid sulphite, 
litmus paper. 

Directions: Set up an apparatus as shown in Fig. 12. A 
is a funnel supported by a ring and attached by a rubber con- 
nector to a tube passing through the stopper in the test tube C. 
The rubber connector must be securely bound with copper wire 
to the funnel and to the tube and must be long enough to allow 
the pinch cock B to close it completely. D, E and F are the 
delivery tube, bottle, and filter paper arranged as in former 
experiments for collecting gases. 

Put 15 c.c. of concentrated solution of sodium acid sulphite 
into the test tube and 10 c.c. of concentrated sulphuric acid 
into the funnel. By opening the pinch cock slightly allow the 
acid to flow into the test tube a few drops at a time, so that 



250 LABORATORY EXERCISES 

a continuous evolution of gas takes place. Moist litmus paper 
held at the mouth of the bottle will show when it has been 
filled. After the second bottle has been collected, replace it 
by a 300 c.c. Erlenmeyer flask one-third full of water, arrange 
the delivery tube so that the end of it is just above the surface 
of the water, and continue the evolution of the gas for ten 
minutes. Cork the flask and keep the solution for Experiments 
61 and 62. (1) "Write the equations for the reactions taking 
place in the generator and then combine them into a single 
equation. (2) Why was the funnel and pinch cock substituted 
for the thistle tube previously used in the gas generator? (3) 
Why was concentrated sulphuric acid used instead of dilute 
acid? 

Experiment 61. Properties of Sulphur Dioxide, Sulphurous 
Acid, and Sulphites. 

Reference: 289-294. 

Discussion : The reactions involving sulphur dioxide fall 
under two heads, namely those in which the valence of the 
sulphur does not alter and those in which it changes. In the 
first class the analogy to carbon dioxide is striking, but in the 
second class no similarity exists, because sulphur unlike carbon 
can exist in compounds in which it has a valence of six. When 
sulphur dioxide is converted into a compound in Avhich sulphur 
exists with the valence of six, the sulphur has been oxidized 
and hence the sulphur dioxide, by causing oxygen or its equiva- 
lent to be given up to the sulphur by the other substance 
involved in the reaction, has acted as a reducing agent. In 
contrast to the action of sulphur dioxide as a reducing agent, 
this substance may also act as an oxidizing agent, when it 
loses its oxygen and is thereby reduced to uncombined sulphur 
or still further to hydrogen sulphide. 

The bleaching brought about by sulphur dioxide differs 
chemically from that produced by chlorine, or, more strictly 
speaking, by hypochlorous acid. The chlorine bleaching 
consists of the oxidation of the colored substances and this 
action usually so far destroys the substance that no reversal 
of the action takes place ; that is the bleaching is permanent. 



252 LABORATORY EXERCISES 

This very completeness of the oxidation, however, precludes 
this method from use with many fibers, such as silk and wool, 
because the fibers themselves are easily oxidized and thereby 
injured by the bleaching agent. On the other hand, sulphur 
dioxide reduces or unites directly with many colored compounds, 
producing substances quite similar to the original ones in 
chemical composition but colorless. Many of these reactions 
are easily reversible, particularly in the presence of an oxidiz- 
ing agent, and therefore the bleaching produced by sulphur 
dioxide is not permanent. Nevertheless the process has the 
great advantage of not injuring the fiber. 

Sulphurous acid and its anhydride sulphur dioxide bear the 
same relation to each other as was found to exist between 
carbonic acid and its anhydride, carbon dioxide. No change 
in valence occurs in the formation of sulphurous acid from 
sulphur dioxide, so the acid is capable of undergoing the same 
sort of reactions of oxidation and reduction as its anhydride 
does. 

Materials: Sulphur dioxide prepared in Experiment 60, 
splinter, copper wire, dyed silk, carnation, potassium chlorate, 
chlorine water, litmus paper, barium chloride solution, hydro- 
gen sulphide, samples for analysis. 

Directions: {a) (1) What is the color and odor of sulphur 
dioxide, and its density compared with air? 

(h) Tnto one of the bottles of sulphur dioxide thrust a 
blazing splinter. (1) What happened? (2) What oxide which 
you have prepared behaves similarly? (3) Would sulphur 
dioxide be a suitable gas to use as a fire extinguisher? (4) 
Why? 

(c) By m_eans of a copper wire suspend a piece of wet 
pink silk (dyed with rosaniline) and a petal of a carnation 
in a bottle of sulphur dioxide until a distinct change takes 
place. (1) What occurred? Rinse the piece of silk under the 
tap and then boil it in a solution containing a spatulaful of 
potassium chlorate dissolved in a test tube half full of water. 
(2) How did the cloth chanfre in appearance? (3) What kind 
of agent is potassium chlorate? (4) By what kind of reaction 
was the color restored? (5) Bv what kind of reaction was it 



254 LABORATORY EXERCISES 

destroyed? (6) As what kind of an agent did the sulphur 
dioxide act? 

Boil a second piece of dyed silk in a solution of chlorine 
water. (7) What effect was produced on the color? (8) Was 
the strength of the fiber of the cloth effected? Rinse the 
cloth under the tap and boil it in the solution of potassium 
chlorate. (9) Was the color restored? (10) To what class 
of reactions does the bleaching process belong? (11) What 
advantage does it possess over bleaching with sulphur dioxide ? 
(12) What disadvantage? 

(c^) Test the solution into which you passed sulphur dioxide 
in Experiment 60 with litmus paper. (1) What was the imme- 
diate action on the paper? (2) The final action? (3) Write 
the equation for the reaction between sulphur dioxide and 
water. (4) Is the reaction reversible? (5) What is the evi- 
dence? (6) To what reaction is this analogous? 

(e) As a preliminary to this experiment make yourself 
familiar with the test for a sulphate by adding to 5 c.c. of 
dilute sulphuric acid in a test tube, 5 drops of barium chloride 
solution and then 5 c.c. of hydrochloric acid. (1) What was 
the result of adding the barium chloride? (2) Of adding the 
hydrochloric acid? (3) State the test for a sulphate. 

To 5 c.c. of the solution of sulphurous acid prepared in 
Experiment 60 add 5 c.c. of chlorine water, shake the solution 
and then test for a sulphate. (4) What was the result of the 
test? (5) Write the equation for the reaction. (6) What was 
oxidized? (7) What was the oxidizing agent? (8) As what 
kind of an agent did the sulphurous acid act? 

(/) Into a test tube half full of the solution prepared in 
Experiment 60 bubble hydrogen sulphide from the generator 
set up in the hood. (1) What insoluble substance was formed? 
(2) Write the equation for its formation. (3) What was 
oxidized? (4) What was reduced? (5) What was the oxidiz- 
ing agent? (6) What was the reducing agent? (7) In sum- 
mary, write three equations, one in which sulphur dioxide is 
oxidized, one in which it is reduced, and one in which neither 
action occurs. 

(g) (!) State the test for a sulphite based on the facts 



256 LABORATORY EXERCISES 

brought out in the foregoing experinienly. Test the unknown 
substances given out for the purpose and (.:,; state whether 
or not each is a sulphite. 

Experiment 62. Sulphuric Acid. 

Reference: 305-306. 

Discussion : Sulphuric acid may be formed by the oxidation 
of sulphurous acid, as was shown in Experiment 61 e, or by 
the oxidation of sulphur dioxide and the subsequent addition 
of water. In the commercial processes the oxidation is brought 
about by the oxygen of the air acting directly in the presence 
of a contact catalytic agent or acting indirectly by the oxida- 
tion of an oxide of nitrogen which in turn oxidizes the sulphur 
diox,ide. In the latter method, the higher oxide of nitrogen, 
having served its purpose as an oxidizing agent and having 
become thereby reduced, is again oxidized and thus used over 
and over again. In actual practice the preparation of the 
sulphuric acid takes place in steps with the preliminary forma- 
tion of nitrosyl sulphui'ic acid which is later decomposed. Part 
{a) of this experiment does not reproduce this method of 
preparing sulphuric acid but it is introduced to show that 
oxidizing compounds of nitrogen, as nitric acid, will oxidize 
sulphurous acid to sulphuric acid. 

Sulphuric acid has a high boiling temperature ; it forms 
addition products readily, particularly with Avater to form hy- 
drates ; it is an oxidizing agent ; and in solution it is highly 
ionized. These four general properties account for the im- 
portant chemical reactions which sulphuric acid exhibits. 

Materials : Solution of sulphurous acid, concentrated nitric 
acid, sugar, splinter, cotton cloth, zinc. 

Directions: (a) Test 5 c.c. of the solution prepared in 
Experiment 60 for sulphuric acid as you tested in Experiment 
61 e. (1) How did you make the test? (2) What was the 
result? (3) Is sulphurous acid oxidized rapidly by the air? 

To a test tube one-quarter full of the solution of sulphurous 
acid prepared in Experiment 60 add 10 drops of concentrated 
nitric acid and boil for two minutes. Dilute with an equal 
volume of water and test for sulphuric acid. (4) What was 



258 LABORATORY EXERCISES 

the result? (5) What kind of reaction did the sulphurous acid 
undergo? (6) What was the agent? (7) Why are special 
methods of oxidation required in the preparation of sulphuric 
acid? (8) What are these methods? (9) What is the ultimate 
source of the oxygen in these methods? 

(h) (1) Is concentrated sulphuric acid a solution like con- 
centrated hydrochloric acid or is it a pure liquid itself? By 
means of a glass rod put a drop of concentrated sulphuric 
acid on a glass plate and nearby put a drop of water. (2) 
Which spreads out more on the glass? (3) What is the name 
for sulphuric acid used in the chemical trade and dependent 
for its origin on this property? 

Weigh your empty graduate on the platform scales, pour 
into it about 10 c.c. of concentrated sulphuric acid, and weigh 
it again. (4) Record the w^eights and the volume of the acid. 
(5) Calculate the specific gravity of the acid. 

(c) Into 5 c.c. of water in a small test tube pour 5 c.c. 
of concentrated sulphuric acid and feel the outside of the tube 
or take the temperature of the mixture with a thermometer. 
(1) What was your observation? (2) What is the evidence 
that chemical action takes place between the acid and the 
water? (3) What is the nature of this action? 

To two spatulafuls of sugar (CjoHoaOn) in a test tube add 
an equal volume of water. Stand the test tube in the rack, 
add to it 10 di'ops of concentrated sulphuric acid and if no 
visible action takes place within one minute add 10 drops more 
of acid. Cautiously smell at the mouth of the test tube. (4) 
What is the insoluble substance formed? (5) What has prob- 
ably happened to the hydrogen and oxygen of the sugar? (6) 
What property of sulphuric acid caused the action to take 
place? (7) What substance did you recognize by its odor 
as being formed? (8) How do you account for its forma- 
tion? 

Put into a test tube a small splinter of wood, a scrap of 
filter paper, and a small piece of cotton cloth and pour over 
them a. few cubic centimeters of concentrated sulphuric acid. 

(9) What happenes to each in the course of a few minutes? 

(10) What do you conclude as to the chemical composition of 



260 LABORATORY EXERCISES 

these substances? (11) As to the nature of the action of sul- 
phuric acid upon them? 

By means of a glass rod make a cross on a piece of filter 
paper with dilute sulphuric acid. (12) Does any action take 
place? (13) Why? Dry the paper by moving it back and 
forth above the flame, taking care that it does not scorch. 
(14) What happened to the paper? (15) Why did it occur? 

(d) Heat in a test tube a strip of zinc and 5 c.c. of con- 
centrated sulphuric acid. (1) What insoluble substance was 
formed? (2) What gas or gases were given off? (3) How 
did you recognize it or them? (4) What kind of an agent 
was the acid? (5) What leads you to this conclusion? (6) 
In Avhat other experiments has sulphuric acid acted as this 
kind of an agent ? (7) What is the relation between the degree 
of reduction of sulphuric acid by a metal and the position of 
the metal in the electromotive series? (8) The action of the 
acid with what two metals illustrates this? 

(e) (1) What was formed in an earlier experiment by the 
action of dilute sulphuric acid with zinc? (2) How do you 
account for the difference between this action and that with 
concentrated sulphuric acid? (3) What evidence have you 
already acquired to show that dilute solutions of sulphuric 
acid contain hydrogen ions? (4) Sulphate ions? 



NITROGEN AND THE ATMOSPHERE 

Experiment 63. Preparation and Properties of Nitrogen. 

Reference: 316-318. 

Discussion : Although nitrogen is the most abundant of 
the uncombined elements, the removal of the other substances 
which are mixed with it in the air is so difficult that pure 
nitrogen is most easily prepared by the decomposition of one 
of its compounds. If an oxide of nitrogen is chosen as a source, 
the oxygen is removed by treatment with a reducing agent, 
as in the reaction between nitric oxide and copper; or if a 
hydrogen compound of the element is selected, it is treated 
with an oxidizing agent, as in the reaction between ammonia 
and copper oxide. The most convenient method, however, 
should be the preparation from a compound of nitrogen which 
contains both hydrogen and oxygen in the proportion to satisfy 
each other, that is in the proportion to form water. Ammonium 
nitrite, NH^NOo, is such a compound and it is easily decom- 
posed by heating. In fact it is so inistable a compound that 
it is usually prepared only at the instant at which it is to be 
decomposed. This series of reactions is brought about by heat- 
ing a mixture of sodium niti'ite and ammonium chloride, when 
the double decomposition between the two substances is imme- 
diately followed by the breaking up of the ammonium nitrite 
into nitrogen and water. 

The remarkable lack of chemical activity of nitrogen is 
easily understood from the fact that nitrogen forms with itself 
a very stable compound. Although the molecules of a com- 
pound are usually thought of as consistnig of atoms of different 
elements held in chemical combination, molecules of elementary 
nitrogen likewise consist of atoms held in chemical combination, 
and may be regarded as a compound of nitrogen with more 
nitrogen. Since the combination is a very difficult one to 

262 



264 LABORATORY EXERCISES 

decompose, elementary nitrogen does not readily undergo 
chemical reaction. 

Materials: Ammonium chloride, sodium nitrite, substances 
necessary to make the tests required in {h). 

Directions: (a) Mix thoroughly 5 grams of ammonium 
chloride with 7 grams of sodium nitrite and place the mixture 
in a test tube fitted with a thistle tube and delivery tube, as 
arranged for the preparation and collection of hydrogen in 
Experiment 8. Have in readiness also a beaker of cold water, 
which may be held so as to immerse the bottom of the test 
tube if necessary. Add to the generator 10 c.c. of water and 
warm the test tube gently. As soon as gas begins to be formed, 
remove the burner and if the action becomes so vigorous that 
the contents of the tube begins to froth, hold the beaker of 
cold water so that it Avill cool the tube until the action becomes 
less violent. A few cubic centimeters of water may also be 
poured down the thistle tube to decrease the frothing if it 
continues. Collect two bottles of nitrogen. (1) Write equa- 
tions for the two reactions taking- place in the genei'ator. (2) 
Why is not ammonium nitrite itself used for the preparation 
of nitrogen? (3) Of Avhat radicals is ammonium niti'ite com- 
posed? (4) In the decomposition of the compound, what ele- 
ment does the nitrogen of the ammonium radical lose? (5) 
What is the removal of this element from a compound termed? 
(6) What element did the nitrooen of the nitrite radical lose 
in the decomposition? (7) What is the loss of this element 
termed? (8) Why should the composition of ammonium nitrite 
make it particularly suitable for the preparation of nitrogen? 
(9) Was the decomposition exothermic or endothermic? (10) 
What was the experimental evidence? 

(h) (1) State the properties of nitrogen as to odor, color, 
and solubility in water. (2) From what common knowledge 
might you have predicted these properties without preparing 
pure nitrogen? 

Try experiments when necessary with the nitrogen which 
you have collected, such as will enable you to answer the 
following questions. Tn each case state what you did to prove 
your answer correct. (3) How can nitrogen be distinguished 



266 LABORATORY EXERCISES 

from oxygen? (4) From hydrogen? (5) From chlorine? (6) 
From hydrochloric acid? (7) From sulphur dioxide ? (8) What 
do you consider the most striking chemical property of nitro- 
gen? (9) What conclusion do you draw as to the stability 
of the nitrogen molecule? (10) Why do you draw this con- 
clusion? 

Experiment 64. The Atmosphere. Percentage of Oxygen. 

Reference : 321. 

Discussion : Air is a mixture, not a pure compound and, 
therefore, the proportion of its constituents is not constant. 
The proportion of oxygen, however, in samples taken from 
different localities varies only slightly and is usually expressed 
in per cent by volume. This percentage is determined by 
removing the oxygen from a measured volume of air and ob- 
serving the volume of the gas which remains, or the volume 
of oxygen absorbed. Since the volumes of gases change con- 
siderably with changes of temperature and changes of pressure, 
care must be taken that the initial and final measurements 
of volumes are taken under the same conditions. 

In the following determination, air confined in a flask of 
known volume is shaken with a solution of sodium pyrogallate, 
which absorbs the oxygen, thus causing a partial vacuum in 
the flask. The flask is now opened under water and a volume 
of water will flow in equal to the volume of the oxygen which 
has been absorbed. The percentage of oxygen in the air is 
then calculated from the volume of oxygen and the original 
volume of air in the flask. 

Matp:rials: Solutions of pyrogallic acid (10 per cent), and 
of sodium hydroxide (50 per cent) ; screw pinchcock. 

Directions: The apparatus consists of a 300 c.c. Erlen- 
meyer flask fitted with a one-hole rubber stopper into which 
is thrust a short piece of glass tubing flush with the bottom 
of the stopper and attached at the other end to a stout rubber 
connector provided with a screw pinchcock. Care must be 
taken that the stopper and tubes fit tightly. 

Fill the flask completely with water, loosen the pinchcock, 
push the stopper into the flask as far as it will readily go, 



268 LABORATORY EXERCISES 

close the pinchcock, remove the stopper, and measure the 
volume of water in the flask by means of your graduate. (1) 
Record the volume. 

Fill a pneumatic trough, adding enough hot water to bring 
the temperature to that of the air in the room. (2) Why? 

Pour into the flask 15 c.c. of the special solution of pyrogallic 
acid and 30 c.c. of the special solution of sodium hydroxide, 
each measured accurately in the graduate, and immediately 
insert the stopper in the flask and close the pinchcock. At 
once put the neck of the flask under Avater in the pneumatic 
trough and keep it there until the experiment is finished. (3) 
Why is the neck of the flask kept under water? Shake the 
flask at frequent intervals for fifteen minutes. (4) AVhat is 
taking place in the flask dui'ing this period? 

Hold tlie flask on its side in the water and loosen the screw 
of the pinchcock, taking care to keep the end of the rubber tube 
constantly below tlie surface. After the water has stopped run- 
ning in, remove the flask from the trough, measure carefully 
in your graduate the volume of the liquid in the flask, and 
(5) record the value. 

The original volume of air is equal to the volume of the 
flask as obtained in the beginning of the experiment minusi 
45 c.c, and the volume of tlie oxygen is ecjual to the final volume 
of liquid in the flask minus 45 c.c. (6) Why must 45 c.c. be 
subtracted in the fii'st case? (7) In the second case? (8) 
In calculating the pei'centage of oxygen, state why it is, or 
is not, necessary to take into account the bai'ometi-ic pressure; 
(9) the temperature of the air. (10) Calculate from your data 
the percentage of oxygen in the air. (11) What is the actual 
percentage? (12) What is the ratio of your error to the actual 
percentage? 

Experiment 65. The Quantitative Determination of Carbon 
Dioxide in Exhaled Air. 

Reference: 322. 

Discussion : The amount of carbon dioxide present in a 
sample of air can be determined quantitatively by shaking a 
known volume of the sample with a solution of barium hy- 



270 LABORATORY EXEHCISES 

droxide of known strength, and observing how much of the 
latter is required to convert the carbon dioxide present into 
barium carbonate. The reaction is carried out in the presence 
of an indicator in order to determine Avhen an excess of the 
hydroxide has been added. Exhaled air contains 4 per cent 
by volume of carbon dioxide. 

Materials: One 2-liter bottle, a 10 c.c. graduated pipette, 
solid Ba(OH)2,8H20, phenolphthalein. 

Directions: To obtain a solution of barium hydroxide of 
known concentration place in a test tube 2 grams of crystalline 
barium hydroxide and 20 c.c. of water, the temperature of 
which is approximately 20°. Close the tube Avith a cork and 
shake vigorously for 3 minutes. At the end of this time there 
should be present some crystals of the hydroxide which have 
not dissolved. Set the tube aside to allow the solid to settle. 

Practice in the following way +he use of a 10 c.c. pipette 
until you can deliver from it any volume desired. Moisten 
the index finger of the right hand and place the upper end 
of the pipette in the mouth and the lower end below the 
surface of some water. Draw up the water by suction until 
the level is above the zero point of the graduation. Place 
the index finger over the upper end. Slightly release the 
pressure of the finger until the liquid falls, drop by drop, from 
the pipette. To stop the flow increase the pressure of the 
finger. Practice delivering 1 c.c. of water. 

Determine the exact volume of a 2-liter bottle by filling 
it with water from a graduate. (1) Record this volume and 
also the temperature of the water and the barometric pressure. 
Insert the bottle filled with water in a pneumatic trough, and 
then completely replace the Avater in the bottle with air from 
the lungs by blowing into it through a rubber tube. Cork the 
bottle and remove it from the trough. Remove the cork, and 
place in the bottle 10 c.c. of water and about 5 drops of a 
solution of phenolphthalein. Replace the cork. 

(2) Record the temperature of the solution of barium hy- 
droxide. Without disturbing the solid on the bottom of the 
tube fill the pipette with the solution to the line slightly above 
the zero point. (The presence of a small amount of suspended 



272 LABORATORY EXERCISES 

barium carbonate in the solution will not affect the results.) 
Let the solution run back into the tube until the level of the 
liquid is just at the zero point, and then let 10 c.c. of it run 
into the bottle of air. Cork the latter tightly and shake it 
vigorously until the pink color of the indicator has disappeared 
(about 2 minutes). Add 1 c.c. of the barium hydroxide solu- 
tion and shake again. Continue the addition of the barium 
hydroxide, 1 c.c. at a time, until a deep pink color remains 
after shaking the bottle vigorously for 3 minutes. (3) Record 
the total number of cubic centimeters added. (If in the addi- 
tion of the base slightly more or less of it has been added 
as the result of inability to stop the flow from the pipette at 
the exact point, record the exact volume added.) 

The weights of Ba(OH). contained in 1 c.c. of a saturated 
solution at different temperatures are as follows: 18°, 0.035 g. ; 
20°, 0.037 g. ; 22°, 0.040 g. Calculate from the number of cubic 
centimeters used and the concentration of the solution at the 
temperature of the experiment (4) the weight of Ba(OH)o 
required to react Avith the CO2 present in the air ; (5) the weight 
of CO2 equivalent to this (Ba(OH),, + CO^ = BaC03 + H2O) ; 
(6) the volume of this weight of COo when mea.sured at the 
temperature and pressure of the air in the bottle; and (7) 
the total volume of air in the bottle, after correcting for the 
liquid added. (8) What is the percentage by volume of COo 
in the exhaled air examined? (9) By how much does this value 
dift'er from what has been determined by exact experimen- 
tation? 

A man exhales about 12 cubic meters of air per day. Cal- 
culate what (10) volume and (11) weight of CO, is exhaled 
per day. (12) What is the weight of carbon in this weight 
of COo? The carbon required is obtained fi'om the food. If 
sugar, which contains 42 per cent of carbon were the food, 
what weight would be required to furnish the necessary 
carbon? 

Experiment 66. Dew Point. Relative Humidity. 

Reference : 94-96, 323. 

Discussion : For every temperature there exists a definite 



274 LABORATORY EXERCISES 

maximum concentration of water vapor beyond which con- 
densation will take place if an excess of vapor is present. The 
concentration is expressed as the pressure of this vapor in 
terms of millimeters of mercury. This maximum concentration 
or pressure is the same as the vapor pressure of water at the 
particular temperature in question. For example, the maxi- 
mum vapor pressure of water vapor at 10° is 9.2 mm. and at 
20° is 17.4 mm. 

The pressure of water vapor in the air may be determined 
by cooling the air and observing the temperature at which 
condensation takes place. Thus, if condensation begins at 
10°, the pressure of the water vapor in the air is 9.2 mm. 
but if condensation begins at 20°, the pressure must be 17.4 
mm. The temperature at which condensation takes place is 
called the dew point. By observing the dew point and referring 
to a table of vapor pressures at different temperatures, the 
concentration or pressure of the water vapor in the air may 
be found. 

If the air has to be cooled in order to cause condensation, 
it is obvious that the air could have contained more water 
vapor at the original temperature. The maximum pressure 
of water vapor which could exist at this original temperature 
may be found by again referring to the table. The relation 
between this maximum pressure and the actual pressure as 
shown by the dew point is spoken of as the relative humidity 
of the air. To illustrate : the air On a certain day was at a 
temperature of 25° and it was found necessary to cool it to 
16° before the dew point was reached. By reference to the 
table, on page 542, it appears that the vapor pressure at 16° 
is 13.5 mm. and at 25° is 23.6 mm. Thus the actual vapor 
pressure was 13.5 mm., the maximum vapor pressure which 
could have existed at 25° was 23.6 mm. and the relative 
humidity on that day was 13.5/23.6 = .57 -|- or 57 per cent. 

To find the dew point experimentally, ice may be added 
slowly to water in a beaker, the outside of which has been 
carefully cleaned, and the temperature at which condensation 
begins on the outside of the beaker may be noted. Since it 
is very difficult to see the very first film of moisture, too low 



276 LABORATORY EXERCISES 

a temperature is usually recorded. If now the ice is removed 
from the beaker and the water inside allowed to rise in tem- 
perature, the condensed water will evaporate and should dis- 
appear when the temperature rises above the dew point. But 
here too there is usually a *4ag" and too high a temperature 
is found. The average between the temperature of appearance 
of the film and the temperature of its disappearance will be 
very close to the true dew point. 

Materials : Ice. 

Directions: Test the accuracy of the graduations on your 
thermometer as directed in Experiment 51. (1) What correc- 
tion do you find it necessary to make? 

Clean and polish the outside of a 200 c.c. beaker, fill it 
about one-third full of tap water and place it so the light 
from the window comes to your eye by reflection from the 
side of the beaker. 

Put into the beaker a lump of ice, about 10 c.c. in volume, 
stir the water with your thermometer, and watch the outside 
of the beaker for the appearance of a film of moisture. If 
the ice melts before the film appears add more ice, lump by 
lump, as it melts, until you see the condensed water. (1) 
Record the temperature at which the film appears. 

Remove the ice, continue the stii-i'ing, and (2) note the 
temperature at which the film disappears. Again determine 
and record the temperature at which the film appears by add- 
ing ice, and at which it disappears when the ice is removed. 
(3) Correct the recorded temperatures by the amount of the 
error of your thermometer and obtain the average of these 
four values. (4) What is this tempcT'ature called? 

Wipe the thermometer and hang it on a ring stand on your 
desk. When the mercury has ceased to rise, (5) record the 
observed temperature and the corrected temperature of the air. 

Referring to the table of vapor pressures of water on page 
542, (6) record the vapor pressure for the dew point which 
you have found and for the temperature of the air. (7) What 
relation has the first value to the pressure of the water vapor 
actually existing in the air? (8) AVhat is the maximum pres- 
sure of water vapor possible in the air at the existing tern- 



278 LABORATORY EXERCISES 

perature? (9) What is the ratio of the pressure of the actual 
vapor present to the pressure of the possible maximum? (10) 
What is this value expressed in per cent? (11) What is this 
percentage called? 



AMMONIA 

Experiment 67. Sources of Ammonia. 

Reference: 336-343. 

Discussion : Formerly all ammonia was derived from or- 
ganic nitrogen compounds which were decomposed by heating 
them. Those compounds occurring in coal were the chief source 
and still are. The ammonia formed in this way is always mixed 
with large quantities of other substances formed during the 
decomposition of the original material. When the gaseous 
mixture is passed into the solution of an acid, the ammonia 
forms a compound with the acid, which may be obtained as a 
solid upon evaporation of the solution. This ammonium com- 
pound will again yield ammonia in a fair state of purity, when 
treated with a base. 

Ammonia is the only gas of common occurrence which forms 
an alkaline solution with water and its presence may thus 
easily be detected with wet litmus paper. 

Materials: Clelatine, ammonium sulphate, concentrated 
solution of sodium hydroxide, powdered sodium hydroxide, 
powdei'ed sodium nitrate, powdei'cd iron, calcium cyanamide, 
litmus paper. 

Directions: (a) Heat in a dry test tube a spatulaful of 
gelatine and try the action of the gas formed upon a piece of 
wTt pink Htmus paper held at the mouth of the tube. (1) 
What was the action on the paper? (2) For what gas is this 
a test? (3) Of what kind of substances is this action of gela- 
tine typical? 

(&) Kefer to your notes on Experiment 42. (1) What is 
the evidence that ammonia is given off upon heating coal? (2) 
Why must ammonia from this source be further treated before 
being used? 

(c) To half a spatulaful of ammonium sulphate on a Avatch 

280 



282 LABORATORY EXERCISES 

glass add a few drops of concentrated sodium hydroxide and 
observe the odor of the gas given off. (1) What was the gas? 
(2) What is the evidence that it is in a purer state than the 
ammonia obtained directly from coal? (3) How can ammonium 
sulphate be prepared from ammonia? (4) By what series of 
steps can pure ammonia be obtained from coal? 

(d) Mix thoroughly half a spatulaful of powdered sodium 
hydroxide (handle this with care as it is very corrosive) with 
an equal amount of powdered sodium nitrate and 2 spatulafuls 
of powdered iron. Heat the mixture in a test tube and test 
for the evolution of ammonia. (1) How did you test and with 
what result? (2) By what sort of reaction was ammonia 
formed in this experiment? 

(e) Boil in a test tube a mixture of half a spatulaful of 
calcium cyanamide and 10 c.c. of water. (1) Was ammonia 
formed? (2) What did you do in order to answer question 
(1) ? (3) Write the equation for the reaction which took place. 
(4) Write two equations to show how calcium cyanamide is 
made. 

Experiment 68. Preparation and Properties of Ammonia. 

Reference : 336, 344-346. 

Discussion : The preparation of ammonia from ammonium 
chloride and calcium hydroxide resembles the preparation of 
carbon dioxide from an acid and a carbonate in that a double 
decomposition is immediately followed by the spontaneous 
breaking up of one of the products into water and a gas. In 
both cases also the secondary reaction is reversible and with 
ammonia this reversal, together with the great solubility of 
ammonia gas itself, is so marked as to require the use of solid 
material instead of solutions, in order to reduce as much as 
possible the amount of water present. 

Ammonia adds directly to the acids and forms a series of 
compounds resembling the salts. Just as oxygen and a non- 
metal may form the negative part of a salt, which exists only 
as an ion in solution and can not be obtained uncombined, 
so four atoms of hydrogen and one of nitrogen may form the 
positive part of a salt, which likewise exists only as an ion and 



284 



LABORATORY EXERCISES 



has not been obtained nncombined. This radical or ion is 
formed by the direct union of ammonia with the hydrogen 
of an acid, but upon attempting to decompose the compound 
to obtain the uncombined radical, the reaction of formation 
is reversed and the original ammonia and acid are obtained. 

Materials: Ammonium chloride, calcium hydroxide, splinter, 
litmus paper, concentrated hydrochloric, and nitric acids. 

Directions: {a) Set up an apparatus as shown in Fig. 13. 
A is a dry 250 c.c. bottle resting in a clamp B. C is a glass tube 
extending nearly to the top of the bottle and projecting about 
1 cm. below the bottom of the stopper in the 8-inch test tube I). 



c+. 



<4 



^ c 



3^ 



:^ 



Fig. 13. 



Mix 10 grams of ammonium chloride with 10 grams of 
calcium hydroxide in a mortal' and put the mixture into the 
test tube. Holding the burner in your hand, heat the test 
tube gently l)eginning at tlie i)oint marked D in the figure 
and slowly work forward. Test for ammonia at tlie mouth 
of the bottle with wet litmus paper and continue the evolution 
of the gas at least one minute after a test is obtained. Remove 
the bottle, and cover it with a glass plate, keeping it inverted 
all the time. Collect two more bottles of ammonia in the 
same way. 

(1) Write two equations for the reactions taking place in 
the generator and then combine them into a single equation. 
(2) Which of these reactions is easily reversible? (3) Why are 



286 LABORATORY EXERCISES 

the solid materials used instead of solutions of them? (4) 
Name substitutes which might have been used instead of am- 
monium chloride and calcium hydroxide. (5) \\"rite equations 
for the reactions which would have taken place had these sub- 
stitutes been used. (6) Why was the bottle held in an inverted 
position? (7) Why should the delivery tube extend nearly to 
the top of the bottle? 

(&) Thrust a burning splinter into one of the bottles of 
ammonia. (1) What was the result? (2) How does this com- 
pare with the similar experiment that you tried with the hydro- 
gen compound of sulphur? 

(c) Hold one of the bottles of ammonia mouth downward 
in the water in the pneumatic trough, shake the bottle keeping 
the mouth under water, replace the glass plate, and remove the 
bottle from the trough. (1) What occurred when the bottle 
was opened under water? (2) What property of ammonia 
caused this action? (3) What was the gas remaining in the 
bottle? (4) How did it get there? 

Drop a i)ieee of pink litmus paper into the water in the 
bottle. (5) What was the action on the litmus? (6) For what 
ion is this action on litmus a test? (7) Does either water or 
ammonia contain this ion? (8) Does ammonia react chemically 
with water? (9) State the logical steps upon which you base 
your ans\N'er. (10) Name the four substances beside water 
Avhich exist in a solution of ammonia. 

(d) Into the third bottle of ammonia thrust a glass rod which 
has been wet with a concentrated solution of hydrochloric acid. 
(1) What did you observe? (2) Write an equation for the 
reaction which occurred. 

Repeat the experiment using concentrated nitric acid on 
the rod instead of hydrochloric acid and again using dilute 
hydrochloric acid and concentrated sulphuric acid. (3) What 
did you observe in each case? (4) What is the reason for the 
difference in the results? 

(e) Thrust the end of a piece of glass tubing about 20 cm. 
in length into a one-hole rubber stopper. Put into the tube 
enough ammonium chloride to occupy about 1 cm. of its length 
and by tapping and rolling the tube get the material near the 



288 LABORATORY EXERCISES 

middle of it. Clamp the stopper to a ring stand so that the 
tube is inclined as much as is possible without allowing the 
ammonium chloride to slip out. Heat the substance and hold 
pieces of both kinds of wet litmus paper at each end of the 
tube. (1) What happened, and (2) what was the action due 
to in each case? (3) Write an equation for the reaction which 
took place upon heating the ammonium chloride. (4) Why 
was the experiment carried out in an inclined instead of hori- 
zontal tube? (5) What was deposited near the ends of the 
tube? (6) Why was it deposited instead of passing out? (7) 
By what reaction was it formed? (8) What will happen if 
ammonium chloride is heated in an open dish? Try the experi- 
ment by heating a little in a porcelain dish. (9) Was your 
prediction verified? 

Experiment 69. The Determination of the Percentage of 
Nitrogen in Ammonia. 

Reference: 345-346. 

Discussion: When ammonia is treated with a solution of 
bleaching powder, CaOCl,, (116), the chlorine of the latter 
unites with the hydrogen of the ammonia and nitrogen is set 
free. In the experiment described below, the amount of 
ammonia in a solution is determined by finding out how much 
hydrochloric acid of known strength is required to neutralize 
20 c.c. of the solution. The amount of nitrogen in 20 c.c. 
of the solution is obtained in the way indicated above. From 
the two results the percentage of nitrogen in ammonia can be 
calculated. 

The required solution of hydrochloric acid of known 
strength is made by diluting concentrated hydrochloric acid, 
the composition of which must first be determined. 

Materials : Concentrated hydrochloric acid, methyl orange, 
bleaching powder. 

Directions: (a) To determine the composition of the con- 
centrated hydrochloric acid, weigh your graduated cylinder to 
decigrams, pour into it about 10 c.c. of the acid, weigh again 
and (1) record the weights and also the volume to 0.1 c.c. (2) 
Calculate the density of the acid. 



290 LABORATORY EXERCISES 

The percentages by weight of hydrochloric acid in solutions 
of different densities are as follows: density 1.17, 33.5 per cent; 
1.18, 35.4 per cent; 1.19, 37.2 per cent. (3) What percentage 
of hydrochloric acid did your solution contain? (4) How many 
grams of the pure compound were contained in the solution 
you weighed out? 

Add water to the acid in the graduate up to the 50 c.e. 
mark, pour the solution into a dry beaker and add to it 150 c.c. 
more of water. (5) Calculate from the weight of acid obtained 
in (4) the weight in 1 c.c. of this diluted solution. 

(h) To exactly 20.0 c.c. of dilute ammonia (5 normal) add 
exactly 80.0 c.c. of water in a flask, shake the solution, pour 
exactly 20.0 c.c. of it into a beaker, and add to it 3 drops of a 
solution of methyl orange, to serve as an indicator. Put exactly 
50.0 c.c. of the dilute hydrochloric acid solution prepared in 
(a) into a graduated cylinder and add it little by little to the 
ammonia until the indicator just changes from yellow to red. 
(1) Record the number of cubic centimeters used. This is only 
a preliminary trial. Repeat the neutralization, adding nearly 
the required amount of acid quite rapidly and making the 
final addition drop by drop. (1) Record the volume of acid 
used to 0.1 c.c. (2) From the value obtained in (a) 5, calculate 
the weight of acid in this volume; (3) the w^ei^ht of ammonia 
required to react with this Aveight of acid (NH.. + HOI = 
.NH4CI). (4) What is the weight of ammonia in 20 c.c. of 
your solution? 

(c) Fit a 250 c.c. flask with a stopper carrying a thistle 
tube and a delivery tube arranged to collect a gas over water. 
Put in the flask 10 grams of fresh bleaching powder and 100 
c.c. of water. Insert the stopper and place the delivery tube 
under an inverted 250 c.c. bottle filled with water in a pneu- 
matic trough. Add to the flask exactly 20 c.c. of your ammonia 
solution and then 10 c.c. of water. Shake the flask occasionally. 
When gas is no longer evolved, cover the gas bottle with a 
plate and remove it from the trough. Pour in additional water 
from your graduate to fill the bottle completely and (1) record 
the volume added, the temperature of the water, and the 
barometric pressure. (2) Subtract from this volume 30 c.c. 



292 LABORATORY EXERCISES 

and state what this value represents. (3) Why was the 
subtraction made? (4) Calculate the volume of the nitrogen 
under standard conditions, not forgetting to take into account 
the vapor pressure of the water. One cubic centimeter of 
nitrogen under standard conditions weighs 0.00125 gm. (5) 
Calculate the weight of nitrogen contained in the 20 c.c. of 
ammonia solution which you used. (6) From this value and 
from the weight of ammonia obtained in (h) 4 calculate the 
percentage of nitrogen in ammonia. (7) Calculate from the 
formula the true percentage. 



ACIDS AND OXIDES OF NITROGEN 

Experiment 70. Preparation and Properties of Nitric Acid. 
Reference : 357, 360-366. 

Discussion: Since nitric acid boils at a much lower tem- 
perature than sulphuric acid, the double decomposition between 
a nitrate and sulphuric acid will proceed, if the temperature 
is high enough to make the nitric acid pass olf as fast as it 
is formed. Nitric acid, however, readily decomposes with the 
formation of a brown gas, when it is heated; for this reason 
the acid is prepared at the lowest temperature Avhich can be 
used to effect the reaction. As a consequence only the first 
step in the double decomposition, which involves the forma- 
tion of an acid sulphate, takes place ; the second step, producing 
a normal sulphate, requires a higher temperature. 

Nitric acid is so corrosive that it attacks rubber or cork 
stoppers and it thus becomes necessary to conduct the prepara- 
tion of the acid in an apparatus composed entirely of glass. 

It has already hecn shown, in Experiment ()2 on the oxida- 
tion of sulphurous acid to sulphuric acid, that niti'ic acid is 
a strong oxidizing agent, and the study of the action of nitric 
acid with a variety of substances fui-ther emphasizes this im- 
portant property. Since the reduction products of nitric acid 
usually are brown gases or gases which turn brown in the air, 
it is easy to recognize, as such, a reaction in which nitric acid 
is acting as an oxidizing agent. 

Materials: Sodium nitrate, sulphur, solution of barium 
chloride, retort, sand bath. 

Directions: Precaution: Nitric acid is very corrosive. Use 
great care not to get a drop of it on your skin or clothing, (a) 
Set up an apparatus as shown in Fig. 14. A is a retoi-t 
clamped to a ring-stand and resting on an iron ])an, /), filled 

294 



296 



LABORATORY EXERCISES 



with sand. The neck of the retort extends into an 8-inch test 
tube, B, which stands in a beaker, C, partly filled with cold 
water. 

Slip into the retort by means of a creased filter paper 12 
grams of sodium nitrate. Place a funnel in the opening of 
the retort and through it pour in 10 c.c. of concentrated sul- 
phuric acid. Insert the stopper of the retort and slowly heat 
the sand in the pan. When about 5 c.c. of nitric acid have 
distilled into the test tube, remove the flame, and raise the 
clamp which holds the retort several inches on the ring stand, 
so that the test tube may be slipped off and the retort allowed 




i^-c 



Fig. 14. 



to cool, a^\'ay fi'om the sand. Save the nitric acid for subse- 
quent experiments. 

When the retort has become quite cool, fill it about half 
full of warm water and let it stand until the solid which it 
contains is somewhat dissolved and becomes loosened from the 
glass. The contents of the retort may be poured into the sink 
and washed down the waste pipe with a large amount of water 
from the tap. 

(1) Write an equation for the reaction by which the nitric 
acid was formed. (2) Write the equation for another reaction 
which may take place between sodium niti-ate and sulphuric 
acid. (3) How could the second reaction be caused to take 
place? (4) Why is this not done? (5) How does your nitric 



298 LABORATORY EXERCISES 

acid differ in appearance from the concentrated nitric acid 
in the laboratory? (6) What is the cause of this difference? 
(7) What substances could be used instead of sodium nitrate 
in this preparation? (8) Could hydrochloric acid be substituted 
for sulphuric acid? Why? (9) Why is this form of generator 
used instead of the simpler one used in preparing hydrochloric 
acid ? 

{b) Pour about 1 c.c. of the nitric acid which j-ou have 
prepared into a 6-inch test tube standing in the test tube rack. 
Be careful not to let the acid run down the outside of either 
test tube. Hold the smaller test tube in a holder and heat the 
acid cautiously. (1) What was formed? (2) Write the equa- 
tion for the reaction. 

(c) Dilute tlie remainder of the acid which you prepared 
with an equal volume of water. If you do not have enough 
acid to perform the following experiments, the concentrated 
nitric acid in the laboratory may be used instead. 

To about 3 c.c. of your acid in a small test tube add a 
small pinch of sulphur and boil until the sulphur disappears. 
Add enough water to half fill the test tube and then test the 
solution for a sulphate. (1) How did you make the test? (2) 
AVhat was the evidence that sulphuric acid was present? (3) 
What property of niti'ic acid does this experiment show? (4) 
State how it proves it? (5) Write three partial equations to 
shoAv how the nitric acid decomposes, how the sulphur is 
oxidized, how sulphuric acid is foi-med from the product, and 
then combine these equations into a single one. (6) Why 
should a different reaction of decomposition of nitric acid be 
written in this case from that written for the decomposition 
in (h)1 

Experiment 71. Nitrates. 

Reference: 367-368. 

Disci^ssTON : It has already been shown in the preceding 
experiment that nitric acid is decomposed by heat. Nitric 
acid, however, is the nitrate of hydrogen and it is important 
to know whether the nitrates of other elements also decompose 
in the same way. This experiment is designed to answer this 



300 LABORATORY EXERCISES 

question. Sodium nitrate is chosen as an example of the 
nitrates of very active metals, of which group potassium is 
the only other common metal, and copper nitrate as an example 
of the nitrates of less active metals. It might be expected 
that ammonium nitrate, being the nitrate of a radical instead 
of a metal, might behave differently from all other nitrates. 
The action of ammonium nitrate will be studied more fully 
in Experiment 73. 

In testing for a nitrate advantage is taken of the fact that 
a very small amount of nitric oxide will form a dark 
brown compound with ferrous sulphate, FeS04,N0. As has 
already been shown nitric acid is reduced to nitric oxide by 
certain reducing agents, of which ferrous sulphate is one. In 
other words, a dark brown compound is formed when nitric 
acid is added to ferrous sulphate, because the latter first re- 
duces the nitric acid and then unites with the nitric oxide 
produced. Thus the formation of this compound may serve 
as a test for nitric acid. The test may be made applicable 
not only to nitric acid but also to any nitrate by the addition 
of concentrated sulphuric acid, since the sulphuric acid will 
convert the nitrate, if present, to nitric acid, which will in 
turn react with the ferrous sulphate as just desci'ibed. In 
actual pi'actice the sulphui'ic acid is added last and in such 
a way that it will form a layer in the bottom of the test tube, 
because it is heavier than the other liquids. The reactions 
will then take place where the layers meet and the brown 
compound will appear as a I'ing, if a nitrate is present. It 
must be remembered, however, that nitrous acid and the nitrites 
also produce niti^ic oxide, and consequently the brown com- 
pound, under these circumstances. The test is of value, there- 
fore, only when nitrites are known to be absent. 

Materivls: Sodium nitrate, copper nitrate, solutions of 
sodium nitrate, sodium nitrite, and ferrous sulphate. 

Directions: (a) Heat a few crystals of sodium nitrate iH 
an ignition tube. It is necessary to heat the substance to as 
high a temperature as possible. When a gas is seen to be 
vigorously given off, hold a glowing splinter at the* mouth 
of the tube. (1) What was the action? (2) What did this 



302 LABORATORY EXERCISES 

prove the gas to be? (3) Was any nitrogen dioxide given offT 
(4) What is the evidence? (5) Does sodium nitrate decompose 
in the same way as hydrogen nitrate? (6) If part, but not 
all, of the element which you identified Avere given off upon 
heating the sodium nitrate, what compound would remain? 
Test the truth of your answer by letting a drop of concentrated 
sulphuric acid run down into the ignition tube after it has 
cooled. (7) What happened? Try the action of a drop of 
concentrated sulphuric acid on a crystal of the substance 
named in the answer to question 6. (8) Was your answer to 
question 6 proved to be correct or not? (9) How? (10) 
Write an equation for the reaction taking place upon heat- 
sodium nitrate. (11) What other nitrate behaves in the same 
way? 

(6) Heat a few crystals of copper nitrate in a small test 
tube and continue the heating until no further change takes 
place. (1) What gas was given off? (2) How did you recog- 
nize it? (3) What remained in the test tube? (4) By what 
property did you recognize it? (5) Write an equation for the 
reaction of decomposition. (6) How does the decomposition 
resemble that of nitric acid? (7) Why was an ignition tube 
required for the heating of sodium nitrate, whereas a test tube 
could be used for copper nitrate? 

(c) To 5 c.c. of sodium nitrate solution in a test tube add 
an equal volume of ferrous sulphate solution. Holding the test 
tube inclined, carefully pour down its side 3 c.c. of concen- 
trated sulphuric acid, which should form a distinct layer in 
the bottom of the tube. (1) What happened between the two 
layers? (2) What is the formula of the compound? (3) What 
two functions does the ferrous sulphate serve? (4) Why was 
the sulphuric acid added? (5) Why was the test tube inclined 
when the acid was added? (6) State the test for a nitrate. 

Repeat the experiment with the substitution of a solution 
of sodium nitrite for sodium nitrate. (7) What is the observed 
result? (8) Can a nitrate be tested for in the presence of a 
nitrite? (9) Why? 



304 LABORATORY EXERCISES 

Experiment 72. Nitrous Acid and Nitrites. 

Reference: 369-373. 

Discussion : Nitrous acid is even less stable than nitric acid. 
This is shown from the fact that when the salts of the acids 
are treated with sulphuric acid at room temperature, the nitrate 
produces nitric acid which is only slightly decomposed, whereas 
the nitrite forms nitric oxide, nitric dioxide, and water instead 
of nitrous acid. 

Nitrous acid, like sulphurous acid, may be reduced and also 
oxidized, or in other words may act as an oxidizing agent or 
as a reducing agent. 

Materials: Solution of sodium nitrite, iodine, sodium 
iodide, potassium permanganate, starch. 

Directions: (a) To 5 c.c. of sodium nitrite solution in a 
test tube add 5 c.c. of dilute sulphuric acid. (1) What oc- 
curred? (2) What compound did you recognize as being 
formed? (3) What nitrogen compound would you expect to 
be formed by double decomposition? (4) What do you con- 
clude as to the stability of nitrous acid? (5) Formulate a test 
for a nitrite. (6) How can a nitrite be distinguished from 
a nitrate? 

(h) As a preliminary to this experiment it is necessary to 
become familiar with a test for iodine. Grind in your mortar 
with 5 c.c. of water a lump of starch the size of a small pea, 
put the paste into a large test tube half full of water and boil 
until the liquid no longer looks cloudy. Dilute 5 c.c. of 
iodine solution with 10 c.c. of water and add a few drops 
of this diluted iodine to 5 c.c. of the starch emulsion. (1) 
What color appeared? This is a test for iodine. (2) State the 
test for iodine. Add a few drops of sodium iodide solution 
to 5 c.c. of starch emulsion. (3) What occurred? (4) Can 
starch be used as a test for uncombined iodine as well as for 
iodine alone? 

To 5 c.c. of starch emulsion add a single drop of sodium 
iodide solution, 5 c.c. of dilute sulphuric acid, and finally 5 c.c. 
of sodium nitrite solution. (5) What happened? (6) What 
was thus proved to have been formed in the reaction? (7) 
Write an equation for the reaction by which it was formed, 



206 LABORATORY EXERCISES 

assuming that nitrous acid and hydrogen iodide had previously 
been formed by the action of sulphuric acid on the sodium 
nitrite and sodium iodide respectively. (8) What kind of 
reaction did the hydrogen iodide undergo? (9) What was the 
agent? (10) What sort of reaction did the nitrous acid 
undergo ? 

(c) When potassium permanganate is reduced the character- 
istic purple color disappears. 

To 5 c.c. of water in a test tube add 3 drops of potassium 
permanganate solution, 5 c.c. of dilute sulphuric acid and u c.c. 
of sodium nitrite solution. (1) What change in color occurred? 
(2) What happened to the nitrous acid present? (3) State in 
logical sequence the steps by which you arrived at this con- 
clusion? (4) What compound which you have already studied 
resembles nitrous acid in that it can be both reduced and 
oxidized? 

Experiment 73. Nitrous Oxide. 

Reference: 376-377. 

Discussion : Just as the hydrogen and oxygen of ammonium 
nitrite were found in Experiment 63 to unite to form water 
Avhen the compound was heated, so the hydrogen and oxygen 
of ammonium nitrate act when this compound is heated. But 
a molecule of ammonium nitrate contains one more atom of 
oxygen than the number required to react with the hydrogen 
present, so the resulting residue is nitrous oxide, NoO, instead 
of elementary nitrogen, Ng, as in the case of the nitrite. 

Nitrous oxide decomposes into a mixture of nitrogen and 
oxygen at a temperature below that at which most substances 
burn. Hence when a glowing splinter is thrust into nitrous 
oxide a mixture is formed which is richer in oxygen than is 
the air, and combustion therefore proceeds more vigorously 
than in the air. Moreover nitrous oxide is an endothermic 
compound, that is it evolves heat upon decomposition, and the 
additional heat thus furnished still further increases the rate 
of combustion. 

Materials: Ammonium nitrate, splinters. 

Directions: Set ut) an a])i)aratus consisting of an R-incli 



308 LABORATORY EXERCISES 

test tube provided with a one-hole rubber stopper through 
which passes a delivery tube leading to a pneumatic trough. 
Have in readiness a beaker of cold water with which to check 
the reaction if it becomes too rapid, as you did in Experiment 
63. Fill the pneumatic trough and gas bottles with water at 
about 40°. Put 10 grams of ammonium nitrate in the test 
tube, insert the stopper, and heat with a low flame until the 
ammonium nitrate melts and the nitrous oxide is slowly 
evolved. Collect two bottles full of the gas and immediately 
disconnect the delivery tube from the generator. (1) Why? 
(2) Write an equation for the reaction taking place in pre- 
paring the nitrous oxide. (3) What very similar reaction have 
you recently studied? 

(6) Thrust a glowing splinter into one of the bottles of 
nitrous oxide. (1) What occurred? (2) With what gas does 
a similar reaction take place? (3) What does this experiment 
suggest about the stability of nitrous oxide upon heating? (4) 
How does it indicate it? (5) Write an equation for the reaction 
of decomposition of nitrous oxide. (6) Would you expect 
combustion to take place more or less vigorously than in air? 
(7) Why? (8) What did you actually find to be the case? 
(9) Was the total heat evolved in burning the splinter in 
nitrous oxide derived from the oxidation of the wood? (10) 
Upon what property of nitrous oxide, stated in the discussion, 
do you base your answer? 

Open the second bottle of nitrous oxide inverted under cold 
Avater in the pneumatic trough. (11) What happened? (12) 
What property of nitrous oxide is shown by this action? (13) 
Why was the gas collected over warm rather than cold water? 

Experiment 74. Nitric Oxide, Nitrogen Dioxide, and Nitro- 
gen Tetr oxide. 

Reference: 379-382. 

Discussion : Nitric oxide and nitrogen dioxide have already 
been met as reduction products of nitric acid. In preparing 
nitric oxide an agent of the proper reducing power and nitric 
acid of the proper concentration must be used or else other 
products of the reduction of nitric acid will be formed instead. 



310 LABORATORY EXERCISES 

Copper and concentrated nitric acid to which has been added 
an equal volume of water are found to be the most convenient 
substances to producer nitric oxide in a fair degree of purity. 
It must be collected out of contact with the air because it 
unites directly with oxygen. Since it is much more stable than 
nitrous oxide, it will not support combustion, as does the latter. 

In the relation of nitrogen dioxide, NOg, to nitrogen 
tetroxide, N2O4, there is exemplified a kind of equilibrium not 
heretofore emphasized in these experiments, namely that which 
may exist in some cases between single molecules and double 
or poly molecules of the same percentage composition. Nitro- 
gen tetroxide, which is colorless, is converted into nitrogen 
dioxide, which is brown, as the temperature is raised. At room 
temperature both gases exist together in equilibrium. 

In the reactions of nitrogen tetroxide and of nitrogen 
dioxide with water is illustrated the oxidation of part of a 
compound at the expense of the remainder, which is thus 
reduced. Nitrogen exhibits the valence of four toward oxygen 
in nitrogen dioxide and tetroxide, but when the latter reacts 
with water nitric acid and nitrous acid are both formed, in 
which compounds nitrogen shows the valences of five and three 
respectively. That is, some of the nitrogen has been oxidized 
from valence four to valence five and some reduced from four 
to three. At a higher temperature more is oxidized to nitric 
acid; consequently that which is reduced must be fui'ther re- 
duced than in the first case; that is, nitric oxide, in which 
nitrogen has the valence of only two, is formed instead of 
nitrous acid. 

Materials: Copper tui'nings, concentrated nitrie aeid, lead 
nitrate, ice. 

DiRKCTiONS: (a) Arrange an apparatus like the one used 
in Experiment 8 for the prepai'ation and collection of hydrogen. 
Put 10 grams of copper turnings in the test tube. Dilute 10 
c.c. of concentrated nitric acid with 10 c.c. of water and pour 
the mixture down the thistle tube. If the reaction becomes 
too rapid, a little cold water may be added and the outside 
of the generator cooled as in previous experiments. Collect 
two bottles of gas. (1) What was the color of the gas in the 



312 LABORATORY EXERCISES 

generator immediately after the action began? (2) What was 
the color of the gas collected? (3) What is its name? (4) 
Write an equation for the reaction by which it was formed. 
(5) Why was it necessary to use diluted nitric acid? (6) Why 
was copper rather than a much more active metal chosen for 
this preparation? 

(5) Thrust a burning splinter into a bottle of nitric oxide. 
(1) What occurred? (2) How is the difference between this 
action and that of nitrous oxide in respect to combustion ac- 
counted for? (3) What did you observe upon opening the 
bottle of nitric oxide? (4) Write an equation for the reaction. 
(5) How do you account for the changes in color in the gases 
in the generator during the preparation of nitric oxide? (6) 
Can nitric oxide be collected by displacement of air? (7) Why? 

(c) Uncover the second bottle of nitric oxide, and pour into 
it 25 c.c. of water, immediately cover it tightly with the palm 
of your hand, and shake it. (1) What two occurrences show 
that the nitrogen dioxide formed upon opening the bottle is 
absorbed by the water? Let a little more air into the bottle 
and again shake it. After doing this several times, test the 
water with litmus paper. (2) What was the action? (3) Write 
an equation for the reaction between the gas and the water. 
(4) In forming nitric acid from the gas, what sort of a process 
has the nitrogen undergone? (5) What was the agent? (6) 
What sort of a change did the agent undergo? (7) What did 
it form? (8) What changes in valence of nitrogen are involved 
in this reaction? (9) What would have been formed at a 
considerably higher temperature? 

{d) Put a si)atulai'ul of lead nitrate into an 8-inch test tube 
and clamp the tube in a horizontal position. Fit it with a 
one-hole rubber stopper through which passes a delivery tube, 
bent at right angles and extending to the bottom of a small 
test tube supported in a gas bottle. Heat the lead nitrate 
until two test tubes have been filled with the gas. Cork the 
tubes and place one tube in a gas bottle full of hot water and 
the other in a bottle full of ice and water. (1) What difference 
in color is observable in the two samples of gas? (2) To what 
is this difference due? (3) What compound is chiefly present 



■X 



314 LABORATORY EXERCISES 

in the cold tube? (4) In the hot tube? (5) Write an equation 
to show the relation of these gases to each other. 

Allow both tubes to come to room temperature by standing 
them in the test tube rack. (6) Is any difference of color finally 
observable? (7) What is now present in the tubes? 



ATOMIC AND MOLECULAR WEIGHTS 

Experiment 75. The Molecular Weight of a Gas. 

Reference: 388-390. 

Discussion : The volume occupied by one gram-molecular- 
weight of a gas under standard conditions is 22.4 liters. Hence 
if the weight of 22.4 liters is known, the molecular weight 
is known. On account of the difficulty of handling so large 
a volume of gas as 22.4 liters, the weight of a measured volume 
is found and the weight of 22.4 liters is computed. 

Materials: Kipp carbon dioxide generator with apparatus 
for drying the gas. 

Directions : Clean and dry a 500 c.c. flask, select a tightly 
fitting cork, and weigh the flask and cork to centigrams. (1) 
Record the weight. Pass carbon dioxide, prepared in a Kipp 
generator and dried by passing through sulphuric acid, into the 
flask until it overflows as shown by the action on a burning 
match held near the mouth of the flask. The carbon dioxide 
delivery tube should extend to the bottom of the flask. (2) 
Why? When the flask has been filled, cork it and weigh it 
again. Record the weight under 1. Again pass carbon dioxide 
into the flask and weigh it. (3) If there is a gain in weight upon 
the second addition of gas, what do you conclude about your 
first attempt to fill the flask? Continue to successively pass 
in carbon dioxide and weigh the flask, until you are sure that 
you have the weight when it is full. (4) How will you know 
when you have obtained this weight? 

Weigh on the platform balance and (5) record the weight 
of the flask full of water and corked. (6) What is the volume 
of the flask? (7) Record the temperature and the barometric 
pressure. One liter of air under standard conditions weighs 
1.293 grams. (8) What is the weight of air in the flask? (9) 
Of the flask empty? (10) Of the carbon dioxide? (11) What 

316 



318 LABORATORY EXERCISES 

would be the volume of the carbon dioxide under standard 
conditions? (12) What would be the weight of 22.4 liters? 

(13) What do you find the molecular weight of the gas to be? 

(14) What is the true molecular weight? (15) What error 
in the molecular weight would be caused by an error of a 
centigram in weighing the flask when filled with carbon 
dioxide? (16) How great an error in weighing the flask full 
of water would have caused the same error in your final result ? 

Experiment 76. Determination of the Molecular Weight of 
Carbon Tetrachloride. 

Reference: 388-389. 

Discussion : The molecular weight of a gas is determined 
by weighing a sample of the gas at any temperature and pres- 
sure and calculating with the aid of the gas laws what weight 
of the gas would occupy 22.4 liters at 0° and 760 mm. pressure, 
provided the behavior of the gas with change in temperature 
and pressure was that indicated by the laws of gases. 

The experiment given below illustrates a method that can 
be used when only approximate results are desired. The sub- 
stance used is carbon tetrachloride, CCI4. It -is a liquid at 
ordinary temperatures (boiling point, 76.6°), but exists as a 
gas at the temperature of boiling water. The molecular weight 
of carbon tetrachloride in the gaseous condition is determined 
by heating some of the liquid in a weighed flask of known 
volume placed in boiling water until the flask contains only 
the vapor of the compound. The flask is next cooled to room 
temperature. The vapor liquefies and air enters the flask to 
take its place. The flask is weighed; the increase in Aveight 
is the v/eight of the known volume of the vapor of carbon 
tetrachloride at the temperature observed and at the pressure 
of the atmosphere. From the weight of the vapor and its 
volume reduced to 0° and 760 mm. the weight of 22.4 liters 
under these conditions can be calculated. This value is the 
molecular weight. 

Materials: A flask which holds approximately 300 c.c. 
when completely filled, large beaker as high as the flask used, 
copper Avire, carbon tetrachloride. 



320 LABORATORY EXERCISES 

Directions: Fit the dry clean flask with a cork stopper 
containing two holes. In one place a thermometer so that the 
bulb of the latter wdll be about 3 to 4 cm. below the stopper. 
Pass through the second hole a piece of glass tubing about 
3 to 4 cm. long, one end of which is drawn out so that the 
diameter of the opening is about 1 mm. The wide end of the 
tube should project below the cork about 0.5 cm. Place the 
stopper carrying the thermometer and glass tube in the flask 
and (1) weigh to decigrams on the platform scales. 

Measure about 5 c.c. of carbon tetrachloride in a dry 
cylinder and pour it into the flask. Place the flask in the 
beaker and wire it down so that it wdll stay in place when 
water is added to the beaker. This can be done conveniently 
by placing a piece of copper wire around the beaker just below 
the top, having one long end and one short end. Wind the 
short end around the wire itself and bring the long end over 
the stopper between the thermometer and the glass tube. Next 
bend the wire drawn tight over the edge of the beaker. 

Place the beaker on a stand, and pour in boiling water 
until it is three-fourths full. Heat the water to boiling. Wind 
a towel around the thermometer and let it rest on the beaker. 
Take care that the ends of the towel do not hang over the 
beaker and come in contact with the flame. The towel serves 
to retain the steam in the upper part of the beaker and thus 
secures a constant temperature in the flask. 

The carbon tetrachloride boils and escapes from the flask 
through the glass tube. A small amount condenses at first, 
however, in the latter and drops back into the flask. Continue 
the heating for about 10 minutes after the liquid ceases to 
drop. (The total time for the heating is about half an hour.) 
(2) Note the temperature recorded on the thermometer. Re- 
move the flask from the water. Dry it carefully with a towel 
and let it stand 15 minutes. (3) Weigh the flask with the 
stopper, thermometer and glass tube to decigrams. The dif- 
ference between this weight and that of the empty flask is 
the weight of the vapor at the temperature observed. To 
determine the volume of the flask pour out the carbon tetra- 
chloride, fill the flask completely with water, and insert the 



322 LABORATORY EXERCISES 

stopper carrying the thermometer and glass tube. Remove 
the stopper and (4) measure with a graduated cylinder the 
volume of the water that remains in the flask. The volume 
of the water is equal to that of the vapor contained in the 
flask. (5) By the use of the gas law formula reduce this 
volume to 0° and 760 mm. This volume of vapor has the 
weight found in the experiment. (6) From these two figures 
calculate the volume of 22.4 liters. (7) What is the molecular 
weight of carbon tetrachloride as calculated from its formula? 
(8) How great was your error? (9) What was your percentage 
error? (10) Why was no correction made for the weight of 
air in the flask? 

Experiment 77. Determination of the Atomic Weight of 
Manganese. 

Reference : 392. 

Discussion : The determinations of the atomic weights of 
metallic elements are usually made by converting one com- 
pound of the metal into another compound. If the atomic 
weights of the other elements in the compounds are known, 
it is possible to calculate the atomic weight of the metal. In 
the experiment described below the atomic weight of man- 
ganese is detei'mined by converting manganese cai'bonate, 
MnCOa, into manganous oxide, MnO : 

MnCO, = MnO + CO^ 

A weighed quantity of manganese carbonate is heated. The 
loss in weight is the weight of the carbon dioxide given off. 
The difference between this and the weight of the carbonate 
used is the weight of the manganous oxide. The calculation 
of the atomic weight is made by solving the following propor- 
tion: Molecular weight of manganous oxide (= atomic weight 
of manganese + 16) : molecular weight of carbon dioxide 
(= 44) = observed weight of manganous oxide: observed 
weight of carbon dioxide, or 

Mn + 16 : 44 = wt. MnO : wt. CO. 



324 LABORATORY EXERCISES 

Materials: Manganese carbonate. 

Directions: (1) Weigh a crucible and cover to centigrams. 
Place in it from 2.5 to 3 grams of manganese carbonate and 
weigh to centigrams. Support the crucible and cover on a 
triangle and heat cautiously by moving the flame of a burner 
under the crucible. At the end of about one minute place the 
burner under the crucible and have the flame of such a size 
that it reaches to the cover. Do not remove the cover until 
the crucible is cold. The carbon dioxide given off from the 
carbonate prevents the oxidation of the manganous oxide. If 
the latter is heated to redness in air it changes to Mn304. Heat 
for 10 minutes. Allow the crucible to cool for 10 minutes and 
(1) weigh. (2) Calculate the atomic weight of manganese 
in the way indicated in the discussion above. (3) What is the 
true atomic Aveight? (4) How great was your error? (5) 
What per cent of the true atomic weight was your error? (6) 
By what other method have you determined the atomic weight 
of a metal? 

Experiment 78. Determination of the Atomic Weight of 
Copper. 

Reference: 392-393. 

Discussion : In determining the atomic weight of a metal 
a rough approximation of the value can be calculated from the 
specific heat of the metal by applying Dulong and Petit 's law, 
which states that the atomic weight multiplied by the specific 
heat = 6.4. The accurate value is obtained by the analysis of 
a compound of the metal. In the experiment described below 
the atomic weight of copper is calculated from the weight ofi 
copper obtained from a known weight of copper sulphate, 
CuS04,5HoO. A solution of copper sulphate is treated with 
finely divided zinc. The reaction which takes place is indicated 
by the following equation : 

CuSO, + Zn = ZnSO, + Cu 

The excess of zinc is dissolved from the product by means of 
dilute hydrochloric acid, which does not affect the copper. 
Taking 32 as the atomic weight of sulphur, 16 as that of oxygen, 



326 LABORATORY EXERCISES 

and 1 for hydrogen, the atomic weight of copper can be cal- 
culated by solving the following proportion : Weight of CuSO^, 
5H2O taken : Weight of Cu formed = At. wt. Cu + 32 + 
4(16) +5 (18) :at.wt. Cu. 

Materials: Copper sulphate, zinc dust, alcohol. 

Directions: (1) Weigh to decigrams on the platform scales 
approximately 20 grams of copper sulphate, CuS04,5HoO. The 
crystals should be clear blue and not covered with a white 
deposit. Place the salt in a 250 c.c. beaker, add 100 c.c. of 
water, heat, and stir until the salt has dissolved. Cool the 
beaker in running water, and add cautiously in small portions 
7 grams of zinc dust, which has been rubbed with a glass rod 
so that it contains no lumps. Stir during the addition of the 
zinc. In about 5 minutes add in small portions 10 c.c. of con- 
centrated hydrochloric acid. Place the beaker on a wire gauze 
and boil gently for 30 minutes. Remove the beaker and stir 
to determine if hydrogen is still evolved from the excess of 
zinc used. > — -> 

When the gas is no longer given off, add lOt) c.c. of water, 
heat to boiling, remove the beaker from the gauze, let the 
copper settle and pour off most of the liquid yjough an 11 cm. 
filter paper. Add 100 c.c. of water to flii^^opper, heat to 
boiling, let the copper settle, and decant oft* the liquid as before. 
RemT)ve by means of a spatula as much of the copper as possible 
to the filter paper. With the aid of a little water transfer 
the rest of the copper to the filter paper. Heat about 50 c.c. 
of water to boiling and pour it through the filter to complete 
the washing of the copper. If alcohol is available the dryint? 
of the copper can be effected more rapidly. Pour alcohol on 
the copper up to the edge of the paper. When no liquid drops 
from the funnel wash again with alcohol. Remove carefully 
the paper containing the copper from the funnel, and place the 
former in a dry evaporating dish. Place the dish in a drying 
closet or on the steam bath. When the paper is thoroughly 
dry, weigh it and the copper to centigrams. Place on the other 
pan of the balance a piece of filter paper of the same size. The 
weight observed is the weight of the copper. Record this 
weight under 1. 



»i ^*<K 




a-^ 




\ 



^ 




328 LABORATORY EXERCISES 

(2) Calculate the atomic weight of the metal in the way 
indicated in the discussion above. (3) Calculate an ap- 
proximate atomic weight of copper from the fact that its 
specific heat is 0.091. (4) Which method of determination 
gives the more accurate result? (5) Why? (6) Of what value, 
then, is the other method? (7) In this experiment is it neces- 
sary to weigh the zinc accurately? (8) Why? 



THE HALOGENS 

Experiment 79. Preparation and Properties of Bromine. 

Reference: 401-403. 

Discussion : Pluorine, chlorine, bromine, and iodine show 
striking similarity in their chemical properties, and they all 
fall into the same column of the periodic classification of the 
elements. In studying such a group, or so-called ''family, " 
of elements, one may be studied in detail and then the differ- 
ences and resemblances between the members of the family may 
be emphasized. Thus bromine should be constantly compared 
with chlorine, and iodine in turn with these two. 

Materials: Potassium bromide, manganese dioxide. 

Directions: Precaution : Bromine is very corrosive and its 
vapor attacks the membrane of the nose and throat. Perform 
all experiments involving bromine in the hood. Grind in a 
mortar about 2 grams of potassium bromide, add an equal 
volume of manganese dioxide and transfer the mixture to an 
8-inch test tube, clami)ed to a ring stand and fitted with a one- 
hole stopper through which passes a delivery tube leading to 
the bottom of a small test tube placed in a bottle of cold water. 
To 5 c.c. of water add 3 c.c. of concentrated sulphuric acid 
(never add the water to the acid) and when the mixture has 
cooled, pour it into the test tube Avhieh contains the potassium 
bromide mixture and at once replace the stopper in the test 
tube. Heat the mixture until a few drops of liquid bromine 
have been collected in the receiver. When the preparation 
of the bromine has been completed, empty the generator into 
the bottle provided for the purpose. 

(1) Wrile equations for the two reactions involved in the 
preparation of bromine and then combine the equations into a 
single one. (2) What other member of this group of elements 
have you studied? (3) What elements are included in the 

330 



332 LABORATORY EXERCISES 

group? (4) In what respects does the preparation of bromine 
resemble or differ from the preparation of chlorine? (5) What 
is the color of bromine vapor? (6) Is the vapor heavier or 
lighter than air? (7) What is the evidence? (8) Calculate 
from the atomic weight of bromine, which is 80, the density 
of bromine vapor compared with air, the density of which is 
the same as that of a gas having the molecular weight of 28.8. 
(9) What is the color of liquid bromine? (10) Is it a com- 
paratively volatile liquid? (11) What is the evidence? (12) 
How does it compare with chlorine in this respect ? Pour 5 c.c. 
of water into the test tube containing the bromine. (13) What 
is the density of liquid bromine compared with water? (14) 
What is the evidence as to its solubility? Drop a small piece 
of litmus paper into the bromine and water. (15) What was 
the action on the litmus? (16) What other element acts in 
the same way? 

Pour the bromine and water into the bottle provided to 
receive it. 

Experiment 80. Hydrobromic Acid. 

Reference: 405^07. 

Discussion : When a bromide is treated with concentrated 
sulphuric acid, hydrobromic acid and an acid sulphate are 
formed, as would be expected ; but then a further reaction takes 
place. Just as hydrochloric acid can be oxidized by a suffi- 
ciently strong oxidizing agent, as was shown in Experiment 23, 
so hydrobromic acid also may be oxidized. But this latter 
compound is more readily oxidized than hydroehloi'ic acid and 
even sulphuric acid, which does not react with hydrochloric acid, 
serves to bring about this oxidation. Hence the hydrobromic 
acid expected from the double decomposition of a bromide and 
sulphuric acid Avill contain bromine and the reduction products 
of the sulphuric acid. Pure hydrobromic acid may be prepared 
by the substitution of phosphoric for sulphuric acid or by some 
entirely different method, such as the reaction between phos- 
phorus pentabromide and water. 

Materials: Potassium bromide. 

Directions: Place two or three crystals of potassium 



334 LABORATORY EXERCISES 

bromide in an evaporating dish and pour on them a few drops 
of concentrated sulphuric acid. (1) What color does the liquid 
become? (2) To what is the color due? (3) Write equations 
to show by what reactions it was formed. (4) Does a chloride 
behave in a similar Avay when treated with sulphuric acid? 
(5) Which is the stronger reducing agent, hydrochloric or 
hydrobromic acid? (6) What is the evidence? 

Blow across the top of the evaporating dish. (7) What 
was the result ? (8) With what compound did you once observe 
a similar action? (9) The presence of what compound is in- 
dicated by the phenomenon in this ease? (10) Explain how 
this phenomenon proves that hydrobromic acid is very soluble 
in water. (See Experiment 25.) (11) How can pure hydro- 
bromic acid be prepared? 

Experiment 81. Iodine. 

Reference : 411-413. 

Discussion : The preparation of iodine is similar to the 
preparation of chlorine and bromine. 

At room temperature iodine is a very volatile solid, that 
is it has a high vapor pressure. Upon heating it, the vapor 
pressure increases and reaches atmospheric pressure at a tem- 
perature below the melting point. Hence it sublimes instead 
of melting, as was explained in 186. The chemical activity of 
iodine is less than that of the other halogens but yet it will 
oxidize many compounds and is extensively used for this pur- 
pose, particularly in quantitative chemical analysis. An im- 
portant reaction in these processes is that between iodine and 
sodium thiosulphate. 

2Na.S203 + I2 = 2NaI + Na,S,Oe 

Materials: Potassium iodide, manganese dioxide, alcohol, 
ear])on disulphide, chloroform, solutions of potassium iodide, 
of sodium acid sulphite, of sodium thiosulphate, starch. 

Directions: (o) Grind in a mortar about 2 grams of potas- 
sium iodide, add an equal volume of manganese dioxide and 
put the mixture in a crucible on a wire gauze supported on 



I 



336 LABORATORY EXERCISES 

a ring of the ring stand. Moisten the material with water 
and add 3 drops of concentrated sulphuric acid. Cover the 
crucible with a watch glass, into which pour a little water to 
keep it cool. Heat the crucible gently until the bottom of the 
watch glass becomes coated with crystals of iodine. When 
the crucible has cooled, empty it into the bottle provided for 
the purpose. 

(1) Write equations for the two reactions involved in the 
preparation of iodine and then combine the equations into a 
single one. (2) To what group of elements does iodine belong? 

(3) AVhy did the iodine vapor condense to a solid as it was 
formed instead of first passing thi-ough the liquid state? (4) 
Describe the appearance of solid iodine. 

(h) Transfer a few crystals from the watch glass to a test 
tube and heat them. When the tube is full of vapor incline it 
mouth downward. (1) What is the color of iodine vapor? (2) 
Is it heavier or lighter than air? (3) What is the evidence? 

(4) Calculate the density of iodine vapor compared with air. 
(The molecular weight of iodine is 2 X 127 = 254.) 

(c) Put a small crystal of iodine into each of five test tubes 
and add to them respectively 5 c.c. of water, of alcohol, of 
potassium iodide solution, of carbon disulphide, and of chloro- 
form. (1) What were the colors of the solutions? (2) Is 
iodine as soluble in water as in the other liquids? (3) What 
is the evidence? (4) What is the evidence that iodine reacts 
Avith potassium iodide? (5) Write the equation for the 
reaction. 

(d) Make some starch emulsion as directed in Experiment 
28 d and to 10 c.c. of it add a drop of the water solution of 
iodine. (1) What color appears? Now add sodium acid sul- 
phite solution until the color disappears. (2) What does the 
disappearance of the color show concerning the iodine? (3) 
Write an equation which shows what has happened to it. (4) 
As what kind of an agent has it acted? Repeat the experiment 
using a solution of sodium thiosulphate instead of sodium acid 
sulphite. (5) Write an equation for the reaction which took 
place. (6) Under what circumstances is this reaction used? 



338 LABORATORY EXERCISES 

Experiment 82. Hydriodic Acid. 

Reference: 415-417. 

Discussion : Just as hydrobromic acid can be prepared by 
the action of water on phosphorus bromide, so hydriodic acid 
can be prepared from water and phosphorus iodide. The phos- 
phorus iodide can be made by the direct union of phosphorus 
and iodine. 

As was pointed out in Experiment 80, hydrobromic acid is 
a more powerful reducing agent than hydrochloric acid, being 
able to reduce sulphuric acid to sulphurous acid. From the 
position of iodine in the periodic classification, it would be 
predicted that hydriodic acid would be a still more powerful 
reducing agent than hydrobromic acid, and that further reduc- 
tion products of sulphuric acid might be expected, when an 
iodide is treated with sulphuric acid. Whether or not this 
prediction is verified is shown in part {h) of this experiment. 

Materials: Iodine, red phosphorus, potassium iodide. 

Directions: (a) Place one crystal of iodine, 2 or 3 mm. 
in diameter, on not more than twice its volume of red phos- 
phorus in an evaporating dish. Let one drop of water from 
the end of a glass rod fall upon the mixture. Marked action 
will usually take place within a minute. If it does not, the 
bottom of the dish may be warmed very slightly. Phosphorus 
tri-iodide is formed. (1) Write an equation for the reaction. 
After the reaction is completed, add another drop of water 
and blow across the top of the dish. (2) What appeared? 
(3) To what compound is the action presumably due? (4) 
What is the reason for this belief? (5) Write an equation 
for the reaction between the water and the phosphorus iodide. 

(h) Pour a few drops of concentrated sulphuric acid on a 
crystal of potassium iodide on a watch glass. Observe closely 
the substances formed in the dish and smell the gases given 
off by wafting them toward you with your hand. (1) What 
reduction products of sulphuric acid did you recognize? (Three 
are usually detectable.) (2) What further evidence was there 
of the oxidation of hydriodic acid? (3) Was any hydriodic 
acid given off? (4) How did you test for it? (5) Write an 
equation for the first reaction taking place between potassium 



340 LABORATORY EXERCISES 

iodide and sulphuric acid, and three equations for the forma- 
tion of the reduction products of sulphuric acid. (6) Is hy- 
driodic acid or hydrobromic acid the more powerful reducing 
agent? (7) What is the evidence? (8) From what fact could 
you have predicted their relative reducing power? 

Experiment 83. Hydrofluoric Acid. 

Reference : 424, 426-428. 

Discussion : Hydrofluoric acid is too corrosive and poison- 
ous a substance to be prepared by an elementary class in suffi- 
cient quantities to observe its physical properties. Enough can 
be made, however, to show its action on glass, which is one of 
its most important chemical reactions. 

Fluorine is the most electronegative, or non-metallic, of the 
elements, as would be anticipated from its position at the top 
of group VII of the periodic classification. Consequently it 
forms stable compounds, not only with the metals, but even 
with such non-metals as silicon. Thus, a reaction of double 
decomposition takes place between hydrofluoric acid and the 
oxide of silicon, just as between hydrochloric and the oxide 
of a metal. 

It is convenient to regard calcium silicate, CaSiO;^, as a 
compound of calcium oxide and silicon dioxide (CaO,SiO._,), 
for thus it is evident why hydrofluoric acid acts on calcium 
silicate to form the fluorides of both calcium and silicon, to- 
gether with the second product of double decomposition, water. 
Glass is a mixture of silicates and silicon dioxide, and when 
it is exposed to hydrofluoric acid it is acted upon in accordance 
with the reactions just described. As the action takes place 
unevenly on a polished surface of glass, the surface becomes 
pitted and produces a frosted or etched appearance. 

Materials: Calcium fluoride, paraffin. 

Directions: Precaution: Exercise great care not to inhale 
any hydrofluoric acid or allow it to touch your skin. Work in 
the hood. Warm a glass plate by holding it high above the flame, 
where it will not get so hot as to crack. Wlion the glass is 
warm enough to melt paraffin, rub a lump of this wax over one 
side of the plate so as to form a thin even coating. Scratch 



342 LABORATORY EXERCISES 

some design through the wax with the end of a file, being 
careful that the paraffin is entirely removed and the glass 
exposed along the lines. 

To 5 grams of calcium fluoride in a lead dish add just 
enough concentrated sulphuric acid to form a paste when 
thoroughly stirred with a file. Place the glass plate, wax side 
down, over the dish. At the end of half an hour remove the 
plate and at once scrape the contents of the dish into the bottle 
prepared to receive it and then wash the dish thoroughly. 

Scrape the paraffin from the glass plate with a knife and 
observe the action of the hydrofluoric acid on the glass. (1) 
What has happened to the glass? (2) Why was it coated with 
paraffin? (3) Write an equation for the reaction between the 
calcium fluoride and sulphuric acid. (4) Was there any evi- 
dence of the reduction of sulphuric acid, as in the case of 
treating sodium iodide with sulphuric acid? (5) Why could 
you have predicted this result? (6) Write equations for thd 
reactions taking place between the hydrofluoric acid and (a) 
the silicon dioxide, (&) the calcium silicate, and (c) the sodium 
silicate in the glass. (7) What do these reactions show con- 
cerning the chemical nature of fluorine? (8) Explain why the 
glass appears etched instead of smooth where it was exposed 
to the hydrofluoric acid. (9) Describe a test for a fluoride. 

Experiment 84. Comparisons of the Halogens and of the 
Halides. 

Reference : 441. 

Discussion : In the following comparisons the fluorides are 
omitted because fluorine itself is so difficult to prepare. 

The relative reducing power of the halogen acids has been 
discussed in Experiment 82 and the relative oxidizing power 
of the halogens themselves necessarily follows this property 
in the reverse order. Closely associated with the relative 
oxidizing power is the order in which one halogen may replace 
another from a solution of a compound. 

Valuable tests for the indentification of the halogens depend 
upon the insolubility of the silver halides and their relative 
tendencies to form soluble complex compounds with ammonia. 



344 LABORATORY EXERCISES 

Materials : Solutions of sulphurous acid, chlorine, bromine, 
barium chloride, sodium chloride, potassium bromide, potas- 
sium iodide, sodium carbonate, silver nitrate. 

Directions: (a) (1) Name the halogen acids, except hydro- 
fluoric acid, in the order of their reducing power. (2) Upon 
what experimental evidence do you base this answer? (3) 
At which end of the series would you expect hydrofluoric acid 
to stand? (4) Why? (5) What experiment accords with this 
prediction? 

(h) (1) In what experiment did you determine the relative 
replacing power of the halogens? (2) What did you find the 
order to be? (3) What was this series then called? (4) To 
what property of the elements was the replacing poAver stated 
to be due? (5) In respect to what other property do the 
halogens follow the same order? 

(c) To 5 c.c. each of solutions of sodium chloride, potassium 
bromide and potassium iodide add 3 c.c. of dilute niti'ic acid 
and then a few drops of silver nitrate solution. (1) State 
what happened in each case, including the color produced. (2) 
Write equations for the reactions, both in the molecular and 
ionic forms. 

Add to each test tube ammonium hydroxide, about 5 c.c. at 
a time, shaking after each addition, until the precipitate dis- 
appears or the solution contains a large excess of ammonia as 
indicated by the odor. (3) What happened in each case? (4) 
How can the three halides be distinguished from one another? 
(5) In the case of the disappearance of a silver halide upon 
adding ammonia, what kind of a compound was formed? 

To 5 c.c. of sodium carbonate solution add a few drops of 
silver nitrate and then an excess of nitric acid. (6) What 
happened after the addition of each reagent? The action of 
nitric acid upon the precipitate in this case may be taken as 
typical of the action on all insoluble silver salts other than 
the silver halides. (7) Formulate a test for the halides in 
general. (8) Why is it necessary to use nitric acid in this 
test? (9) State a test for a chloride, (10) two tests for a 
bromide and (11) three tests for an iodide. 



346 LABORATORY EXERCISES 

Experiment 85. Hypochlorites and Chlorates, 

Reference: 431-433. 

Discussion: The action of chlorine on water v/ith the 
formation of hydrochloric and hypochlorous acids was dis- 
cussed under chlorine in Experiment 22. If the water contains 
a base in dilute solution, the acids react to form their salts 
and a mixture of chloride and hypochlorite is formed. 

Bleaching powder, which in the solid state is a mixed salt 
(calcium chloride-hypochlorite), is commonly used for bleach- 
ing. A solution containing sodium hypochlorite, sometimes 
called Javelle water, is often more convenient and can be 
prepared by treating bleaching powder with sodium carbonate 
and filtering. 

As was pointed out in the discussion of the oxides of nitro- 
gen, Experiment 74, a compound which is capable of both 
reduction and oxidation may sometimes decompose of itself, 
part being oxidized at the expense of the remainder. Hypo- 
chlorites offer another example of this kind of reaction, where 
part of the hypochlorite is oxidized to chlorate and part re- 
duced to chloride. If chlorine is passed into a hot concentrated 
solution of a base, a chlorate and a chloride are formed instead 
of hypochlorite and chloride as in the case of a cold dilute 
solution. 

Materials: Bleaching powder, sodium carbonate, colored 
cloth, potassium hydroxide, chlorine, splinter. 

Directions: (a) (1) In what experiment which you have 
performed was hypochlorous acid formed? (2) Write an equa- 
tion for the reaction Avhich took place. (3) What action can 
be used as a test for hypochlorous acid? 

(h) To 10 grams of bleaching powder in an evaporating 
dish add enough water to make a thin paste. Place two 
strips of colored cloth in the paste and keep a third for a 
standard of comparison. Remove one of the pieces of cloth 
from the paste and dip it into a beaker of dilute sulphuric 
acid, again cover it with the bleaching powder, and again dip 
it in the acid. Repeat this four or five times and then wash 
both pieces of cloth under the tap and compare them with 
the original and with each other. (1) What changes have 



348 LABORATORY EXERCISES 

taken place and to what extent? (2) How can bleaching 
powder be treated to make it a more rapid bleaching agent? 

(c) Mix thoroughly in your mortar a spatulaful of bleach- 
ing powder and one of sodium carbonate. Put the mixture 
into a large test tube, add 20 c.c. of water, shake the tube for 
several minutes, and filter. Put a small piece of colored cloth 
into 5 c.c. of the filtrate. (1) What happened to the piece 
of cloth? (2) Write an equation for the reaction between 
the sodium carbonate and the bleaching powder. (3) What 
is the commercial name of the filtrate? 

{d) Dissolve 3 grams of solid potassium hydroxide in 7 c.c. 
of water in a test tube and while the solution is still warm 
bubble into it chlorine from the generator in the hood until 
crystals are precipitated. Cool the tube under the tap and 
filter out the crystals, keeping them in the tip of the cone of 
filter paper. Remove the filter from the funnel, spread it out, 
and carefully transfer the crystals to a sheet of dry filter 
paper. When the substance is thoroughly dry put some of it 
into an ignition tube and heat it. Be careful that no shreds 
of filter paper get into the tube. Test at the mouth of the 
tube for oxygen. (1) How did you test? (2) With what 
result? (3) What were the crystals? (4) Write equations 
in five steps and a combined equation to show how the sub- 
stance was formed. 



PHOSPHORUS, ARSENIC, ANTIMONY AND BISMUTH 

Experiment 86. Phosphoric Acids and Phosphates. 

Reference: 458-463. 

Discussion : When the oxide of a non-metal reacts Avith 
water, the resulting hydroxide is an acid. It is theoretically 
possible to form a large number of different acids from the 
same oxide and water by causing reactions to take place be- 
tween the two in different proportions. In the cases which 
have already been studied of the action of carbon dioxide, 
sulphur dioxide, and sulphur trioxide upon water only one 
of the possible acids has been formed in each case, because 
the others are unstable or do not exist at all. With phosphorus, 
however, three well-defined acids are known which may be 
regarded as being formed from one molecule of phosphorus 
pentoxide and one, two, and three molecules of water, respec- 
tively. The formulas for these acids are therefore P205,H^,0 
or HPO3, P20,,2H,0 or H.P.O^, and Po05,3H20 or "H3P6,. 
They are all called phosphoric acid, but are distinguished by 
the prefixes meta-, pyro-, and ortho- respectively. These three 
acids form corresponding salts which differ in their properties 
and can thus be made use of in tests for the three different 
acid radicals. Orthophosphoric acid introduces a further com- 
plexity, in that three sets of salts may be derived from it, 
namely those in which one, two, or three hydrogen atoms have 
been replaced by the same metal. These are called primary, 
secondary, and tertiary orthophosphates, respectively. 

When phosphorus pentoxide, or as it is commonly called 
phosphoric anhydride, is put into cold water, the meta acid 
is formed. Upon boiling this slowly it reacts with more water 
and produces the ortho acid. The pyro acid Avhich might be 
expected as an intermediate product is not formed. This acid, 
or more usually one of its salts, can be prepared, however, by 

350 



352 LABORATORY EXERCISES 

the reverse process, namely by the dehydration of a secondary 
ortho compound. The dehydration of a primary orthophosphate 
gives a metaphosphate. 

Materials: Disodium orthophosphate, microcosmic salt, 
phosphoric anhydride, solutions of ammonium molybdate, 
silver nitrate, calcium chloride. 

Directions: (a) Preparation of a pyrophosphate. Heat 
gently in a porcelain crucible 2 grams of disodium orthophos- 
phate. Gradually increase the heat to the highest temperature 
obtainable and continue the heating for 10 minutes. When 
the crucible has cooled, dissolve the contents in cold water 
and keep the solution for the tests in part (c). (1) Write an 
equation for the reaction. (2) Of what acid is the product 
a salt? (3) What is the relation of this acid, to phosphoric 
anhydride 1 

(6) Preparation of a metaphosphate. Kepeat experiment 
(a) with the substitution of microcosmic salt, NaNH^HP04, 
for the disodium phosphate, saving the solution for the tests 
in (c). Test the gas given off with red litmus paper. (1) 
What substance do you recognize as a product of decomposi- 
tion? (2) Write an equation for the reaction. (3) A salt of 
what acid is formed? (4) What is the relation of this acid 
to phosphoric anhydride? (5) Does microcosmic salt decom- 
pose like a primary or like a secondary orthophosphate? (6) 
Upon what do you base your answer? 

(c) Tests. (1) Prepare a table for recording the results 
of the following tests by writing in the first column the names, 
ortho, pyro, meta, by heading the second column molybdate, 
the third silver nitrate, and the fourth calcium chloi'ide. Record 
in this table the formation and color of precipitates when they 
appear. To 1 c.c. of a solution of disodium orthophosphate 
add 2 c.c. of dilute nitric acid, 3 c.c. of ammonium molybdate 
and warm the solution until a precipitate appears. 

Into each of two test tubes put 5 c.c. of a solution of 
disodium orthophosphate and add to one a few drops of silver 
nitrate solution and to the other a few drops of calcium chloride 
solution. Repeat the foregoing tests with the pyrophosphate 
and metaphosphate solutions prepared in (a) and (h). (2) 



354 LABORATORY EXERCISES 

What is a general test for all phosphates? (3) How can each 
of the three phosphates be distinguished from the other two ? 

(d) Drop a quarter of a spatulaful of phosphoric anhydride 
into 15 c.c. of cold water in a beaker and stir the solution. 
(1) Describe the appearance of phosphoric anhydride. (2) 
By what other name is it known? (3) What was the evidence 
of vigorous reaction betwen the oxide and water? Determine 
by testing a small amount of the solution which phosphoric 
acid was formed. (4) What test did you use and with what 
result? (5) What acid was present? (6) Write an equation 
for its formation. 

To the remainder of the solution add 5 c.c. of dilute nitric 
acid, which acts only as a catalyst, and boil it for 10 minutes. 
(7) Which acid is now present? (8) Upon what tests do you 
base this answer? (9) Write an equation for the formation 
of the acid. (10) Name in order the three compounds which 
may be regarded as successive stages of hydration of phos- 
phoric anhydride. 

Experiment 87. Halides of Phosphorus. 

Reference : 467. 

Discussion : Phosphorus tri-iodide has already been pre- 
pared by the direct union of phosphorus and iodine in Experi- 
ment 82. Two sets of halides would be expected corresponding 
to the two valences in which phosphoi'us appears in combina- 
tion. All of these compounds are known with the exception 
of phosphorus penta-iodide. Since phosphorus is a distinctly 
aeid-forming rathci* than ])ase-forming element, its halides may 
be completely hydrolyzed in water. 

Materials: Phosphorus trichloride, phosphorus penta- 
chloride. 

Directions: Perform these experiments in the hood, (a) 
Pour a few drops of phosphorus trichloride into a dry test tube. 
(1) Describe the appearance of the compound. Blow across 
the mouth of the test tube. (2) What occurred? (3) What 
do you conclude is being formed? 

Half fill the test tube Avith water. (4) What phenomenon 
do you observe? (5) Explain its cause. (6) Write an equation 
for the reaction between phosphorus trichloride and water. 



356 LABORATORY EXERCISES 

{h) Examine a small amount of phosphorus pentachloride, 
about 1 gram. (1) Describe its appearance. Determine 
whether or not it will hydrolyze. (2) What did you do and 
with what results? (3) What do you conclude? (4) What 
property of phosphorus might lead you to expect this action? 
(5) Write an equation for the reaction. 

(c) (1) In what previous experiment was phosphorus tri- 
iodide formed and by what means? (2) How did it react with 
water? (3) Write an equation for the reaction. 

Experiment 88. Preparation of Arsenic. 

Reference: 470-471. 

Discussion : Many elements which do not occur uncombined 
in nature may be prepared by heating a mixture of the oxide 
of the element with carbon. If the elementary substance 
formed in this way by the reduction of the oxide is volatile 
at the temperature of the reaction, it will pass off along with 
the oxide of carbon and be deposited again as a solid in the 
cold part of the apparatus. 

Materials : Arsenious oxide, powdered charcoal, lime water. 

Directions: Mix thoroughly about 1 gram of arsenious 
oxide with three times its volume of powdered wood charcoal 
and place the mixture in the bottom of a test tube, clamped 
to the ring-stand in a horizontal position, and fitted with a 
one-hole stopper and delivery tube, which dips into lime water 
in a test tube. Heat the mixture until a dark deposit is ob- 
served near the mouth of the tube and then at once remove 
the delivery tube from the lime water. (1) Why? (2) What 
is the deposit? (3) Was any carbon dioxide formed? (4) 
What is the evidence? (5) Write an equation for the reaction 
upon heating. (6) What kind of reaction did the arsenious 
oxide undergo? (7) What was the agent? (8) What process 
is illustrated by the appearance of the deposit at a distance 
from the place at which it was formed? 

Experiment 89. Comparison of Phosphorus, Arsenic, An- 
timony, and Bismuth. 

Reference : 449, 474, 478, 480, 485, 488, 490, 494, 495. 

Discussion : Comparison of the compounds of elements in 



358 LABORATORY EXERCISES 

the same group of the periodic classification shows that with 
increasing atomic weights their non-metallic or electro-negative 
properties decrease and their metallic or electro-positive prop- 
erties increase. This relation is very clearly exemplified in 
Group V in the case of phosphorus, arsenic, antimony, and bis- 
muth. 

Positively charged ions of phosphorus are not known. Posi- 
tive ions of arsenic presumably exist in small amounts in 
strongly acid solutions, since sulphides of arsenic can be pre- 
cipitated by hydrogen sulphide. Upon passing down the series 
to bismuth, well-defined metallic salts are found to be its most 
stable compounds. 

The hydroxides of phosphorus are entirely acidic, those of 
arsenic markedly so, whereas the hydroxide of antimony must 
be treated with an excess of a strong base to convert it into 
a salt, and the hydroxide of bismuth remains insoluble even 
in a large excess of a strong base. 

The action of the chlorides of these elements still further 
illustrates the progressive change in their chemical properties. 
The complete hydrolysis of the chlorides of phosphorus has 
already been observed. Arsenious chloride becomes hydrolyzed 
to a very great extent although some arsenious ions still exist 
in solution, as was stated above. Antimony chloride forms 
an oxychloride and bismuth chloride a basic chloride when 
these compounds are added to water. These facts indicate 
partial hydrolysis in both cases. 

The sulphides of arsenic and antimony are sufficiently acidic 
to form soluble compounds with ammonium sulphide ; but bis- 
muth is so electro-positive that its sulphide does not act in this 
way. 

Materials: Arsenious oxide, antimony trichloride, bismuth 
chloride, hydrogen sulphide, solutions of sodium hydroxide, 
concentrated hydrochloric acid, ammonium sulphide, litmus 
paper. 

Directions: (a) (1) In what experiment did you obtain 
a hydroxide of phosphorus? (2) Write an equation for the 
reaction by which it was formed. 

To a pinch of arsenious oxide in a test tube add 5 c.c. of 



360 LABORATORY EXERCISES 

water. (3) How does the vigor of the action compare with 
that of an oxide of phosphorus with water? 

Add 5 c.c. of sodium hydroxide and boil the solution until 
it is clear. (4) Write an equation for the reaction taking 
place. (5) What is the name of the arsenic compound? (6) 
Of what acid is it a salt? 

Add hydrochloric acid in excess, as shown by the action 
on litmus paper. (7) Write an equation to show the reaction 
of double decomposition. 

Bubble hydrogen sulphide into the solution. (8) Describe 
and name the compound which appears. (9) What ion must 
have been present in the solution, in order that hydrogen sul- 
phide could form the compound precipitated? (10) Write an 
equation to show how arsenious sulphide was formed. 

Filter out the precipitate and wash it on the filter paper 
with hot water until the filtrate is only faintly acid. Rest 
the funnel in a clean test tube, poke a hole in the tip of the 
filter paper, and wash the precipitate through into the test 
tube with 5 c.c. of ammonium sulphide. (11) What has hap- 
pened to the precipitate? (12) Write an equation for a hypo- 
thetical reaction between one molecule of arsenious oxide and 
three molecules of sodium oxide. (13) Write beneath this 
equation a similar one in which sulphur is substituted for 
oxygen and the ammonium radical for sodium. (14) For what 
reaction Avhieh you have carried out is this the equation? (15) 
What chemical characteristic of arsenic is the cause of this 
reaction? 

Add an excess of hydrochloric acid to the solution in the 
test tube. (16) What appeared? (17) Write an equation for 
the reaction. 

(h) Dissolve a small lump of antimony trichloride, about 
the size of a pea, in 5 c.c. of concentrated hydrochloric acid 
in a test tube. Add water to this solution, a few cubic centi- 
meters at a time, until a precipitate appears. (1) What is the 
precipitate? (2) Write an equation to show how it was formed. 

Add concentrated hydrochloric acid drop by drop until the 
precipitate disappears. (3) What relation does this reaction 
bear to the preceding one? 



362 LABORATORY EXERCISES 

Bubble hydrogen sulphide into the solution. (4) Describe 
and name the compound formed. (5) Write an equation for 
the reaction. , 

Filter out the precipitate, wash it, and treat it with am- 
monium sulphide as you did under corresponding circumstances 
with the arsenic compound. (6) What happened? (7) Write 
an equation for the reaction. 

Try the effect of acidifying the solution. (8) What ap- 
peared? (9) Write an equation for the reaction. 

To 5 c.c. of a solution of antimony trichloride add very 
slowly sodium hydroxide solution until finally a considerable 
excess is present, as shown by litmus paper. (10) What changes 
did you observe? (11) Write equations for the reactions. 

Repeat the last experiment with the substitution of am- 
monium hydroxide for sodium hydroxide. (12) What hap- 
pened? (13) How do you account for the difference in behavior 
in these cases? 

(c) Dissolve a small lump of bismuth chloride in 5 c.c. of 
concentrated hydrochloric acid and add water slowly as in the 
experiment with antimony chloride. (1) Record what hap- 
pened. (2) Write an equation for the reaction. (3) What is 
the relation of the compound formed in this case to that formed 
with antimony chloride? 

Add concentrated hydrochloric acid slowly until the pre- 
cipitate disappears and then bubble in hydrogen sulphide. (4) 
Name and describe the compound precipitated. (5) Write an 
equation for its formation. 

Filter and wash the precipitate and try the action of am- 
monium sulphide with it. (6) What happened? 

To 5 c.c. of a solution of bismuth chloride add sodium 
hydroxide slowly and finally in excess. (7) What was formed? 
(8) How did the action differ from the corresponding experi- 
ment with antimony chloride? (9) Write an equation for the 
reaction. 

(d) (1) In which group of the periodic classification do 
the elements studied in this experiment fall? (2) What other 
elements which you have studied are in this group ? (3) Arrange 
the elements of this group in the order of their increasing 



364 LABORATORY EXERCISES 

atomic weights. (4) Which of these five differs most widely 
from the others in its chemical properties? (5) In what other 
group have you found the element occupying the similar posi- 
tion differing from the other members? In making the follow- 
ing comparisons omit nitrogen from consideration. (6) What 
is the order of the elements of the fifth group in respect to their 
decreasing electro-negative nature? (7) What experiments 
show the relative positions of arsenic and phosphorus? (8) 
Of arsenic and antimony? (9) Of bismuth and antimony? 
(10) Formulate a test by which you could recognize a com- 
pound of arsenic. (11) A compound of antimony. (12) A 
compound of bismuth. 



ORGANIC COMPOUNDS 

Experiment 90. Properties of Cellulose. 

Keferences: 499, 502. 

Discussion : Cellulose is produced in the growth of plants 
and is the chief constituent of wood and cotton. The purified 
material is used in making paper and cotton and linen cloth. 
Silk and wool are animal products, containing nitrogen, which 
are called proteins; wool also contains sulphur. The behavior 
of cellulose with acids and alkalies is quite different from that 
of proteins; as a consequence it is possible to distinguish by- 
chemical means cotton from silk and avooI. The presence of 
sulphur in wool serves to distinguish it from silk, which does 
not contain this element. 

Cellulose is an alcohol, and like other compounds of this 
class it forms a nitrate when treated with strong nitric acid. 
A number of different nitrates of cellulose can be made in this 
way. The one containing the highest percentage of nitrogen 
is called gun-cotton, and is used in smokeless powder ; the one 
containing a smaller percentage of nitrogen is used in making 
collodion and celluloid. In the experiment described below the 
conditions are given that yield the product which can be used 
to make colloidon. 

Materials : Pieces, about 1 inch square, of cotton, silk, and 
wool; raw cotton; concentrated hydrochloric and nitric acids; 
10 per cent solution of sodium hydroxide; solution of lead 
acetate ; alcohol ; ether. 

Directions: (a) In separate test tubes shake pieces of 
cotton, silk, and wool with about 10 c.c. of cold concentrated 
hydrochloric acid, until one of them dissolves. (1) Record the 
result in each case. 

(6) Place a piece of cotton and a piece of wool in a mixture 
of 1 c.c. of concentrated hydrochloric acid and 5 c.c. of water. 

306 



368 LABORATORY EXERCISES 

Press out the solution, and dry the pieces on the steam bath 
or set them aside until the next exercise. Try to tear the 
material. (1) Record the results. 

(c) Place in a test tube a piece of cotton and one of wool 
and add 10 c.c. of a 10 per cent solution of sodium hydroxide. 
Heat the tube to such a temperature that it can just be held 
in the hand and shake for some time. (1) Record the result. 

(d) Add a few drops of a solution of lead acetate to 20 c.c. 
of a 10 per cent solution of sodium hydroxide. Divide the 
solution into two portions. In one dissolve by heating a piece 
of wool and in the other a piece of silk. (1) Record and 
explain the results. (2) How can cotton, wool, and silk be 
distinguished from one another? 

(e) Pour 10 c.c. of concentrated sulphuric acid into 10 c.c. 
of concentrated nitric acid. To the hot solution add 0.5 gram 
of raw cotton. At the end of 3 minutes withdraw the cotton, 
and drop it into a large amount of cold water. Wash the 
cotton in running water for about 1 minute, squeezing out 
the w^ater from time to time. Set it aside to dry. Hold a small 
bit of the dry nitrate with tongs and place it in a flame. (1) 
Record the result. (2) What is the chemical composition of 
the substance? 

Pour on to some of the nitrated cotton in a test tube 5 c.c. 
of a mixture of equal volumes of alcohol and ether. After a 
few minutes, pour off the clear liquid into a watch glass and 
let it evaporate sloAvly until all the liquid has disappeared. 
Place some water on the watch glass and after a few minutes 
lift the edge of the film and remove it from the glass. Dry 
it with filter paper. Hold a bit of the film Avith tongs and 
put it in a flame. (3) Record the result. 

Experiment 91. The Preparation of Soap. 

Reference : 502. 

Discussion : The animal and vegetable fats are composed 
of the glycerine esters of acids of high molecular weight. When 
these esters are heated with a solution of sodium hydroxide 
they are converted into glycerine and the sodium salts of the 
acids, Avhich are called soap. Soap is somewhat soluble in 



370 LABORATORY EXERCISES 

water but is insoluble in a solution of sodium chloride. This 
fact is used in separating soap from water and glycerine, and 
from the excess of alkali used in the saponification of fats. 

Materials: Lard, sodium hydroxide, toilet soap, sodium 
chloride, solution of calcium chloride. 

Directions: (a) Weigh into a small beaker 20 grams of 
lard, and add to it a solution of 5 grams of sodium hydroxide 
dissolved in 5 c.c. of water. Rest the beaker in a second beaker 
containing boiling water and heat the mixture in this way 
for from 45 to 60 minutes. Stir the mixture occasionally until 
the contents hardens. 

Remove the soap from the beaker. Add to a piece of it a 
drop of a solution of phenolphthalein. Test in the same way 
a piece of pure toilet soap. (1) Give a reason for the difference 
observed. (2) Write an equation for a reaction involved in 
the preparation of soap from lard, which contains the glycerine 
ester of stearic acid. 

(h) Cut about one-half of the soap into thin shavings, place 
it in 150 c.c. of water, and gently boil the solution until the 
soap dissolves. Prepare 100 c.c. of a filtered saturated solution 
of sodium chloride and add it to the soap solution. Set the 
beaker away until the next exercise. (1) What was the action 
of the salt when it was added to the soap solution? (2) Why 
is this procedure used? 

Remove the soap and let a part of it dry in a filter paper 
and use the rest for the following experiments. Dissolve a 
little of the soap in water and shake some of the solution in a 
test tube. (3) Record the result. Add calcium chloride to 
some of the solution. (4) What happened? Filter off the 
precipitate and shake some of it with water. (5) Is the pre- 
cipitate soap? Give a reason for your answer. (6) What is 
the precipitate? Stearic acid is insoluble in water. (7) How 
could it be obtained from soap? 



SILICON AND BORON 

Experiment 92. Silicic Acid and Silica. 

Reference: 508. 

Discussion: Various silicic acids are known, which may 
be regarded as hydrates of silicon dioxide. tJnlike the cases 
which have already been studied, these acids can not be pre- 
pared by the action of the anhydride on water, because silicic 
anhydride is extremely insoluble. They may be formed, how- 
ever, by the double decomposition between a silicate and an 
acid, the particular silicic acid produced being dependent on 
the substances used, and on their concentration and tempera- 
ture. The fully hydrated orthosilicic acid, H4Si04, is a jelly, 
which can be completely dehydrated upon heating. 

Materials : Solution of sodium silicate, concentrated hydro- 
chloric acid. 

Directions: (a) To 10 c.c. of sodium silicate in an 
evaporating dish add slowly with constant stirring 10 c.c. of 
dilute hydrochloric acid. (1) State how the substance formed 
appears and how it feels when rubbed between the fingers and 
thumb. (2) Assuming the original salt to be the meta silicate, 
Na^SiOg, write an equation for the reaction. 

Evaporate the contents of the dish to dryness in the hood, 
and when the dish has cooled, thoroughly moisten the i-esidue 
with concentrated hydrochloric acid. Again evaporate to dry- 
ness and then heat strongly for five minutes. When the dish 
is cool add 20 c.c. of water and stir to dissolve the sodium 
chloride. Carefully pour off the solution leaving the insoluble 
residue in the dish. (3) Describe its appearance and state how 
it feels when rubbed. (4) What is the substance? (5) Write 
an equation for the reaction by which it was formed. 

Transfer the substance from the dish to a test tube and 
boil it with 10 c.c. of water. (6) What happened? (7) State 

372 



374 LABORATORY EXERCISES 

a striking difference between the chemical behavior of this 
compound and other acid anhydrides you have studied. 

(6) (1) In what experiment did you study the action of an 
acid on silicon dioxide? (2) Write an equation for the reac- 
tion. (3) Do other acids act this way Avith silicon dioxide? 
(4) Upon what do you base your answer? 

Experiment 93. Compounds of Boron. 

Keference: 519-523. 

Discussion : The comparative insolubility of boric acid in 
cold water makes it possible to prepare the acid from one of 
its soluble salts by double decomposition with another acid. 

Borax, Na2B407, may be regarded as a compound of one 
molecule of sodium oxide with two of boric anhydride, Na^O, 
2B2O3. Thus, enough boron exists in borax to form sodium 
meta borate, NasO, B.fi.^, and also to unite with a molecule of 
some other basic oxide to form another borate. When borax 
is fused with the oxide of a metal, this reaction takes place 
and the glassy mass of sodium meta borate becomes mixed with 
the borate of the other metal. Many of these borates possess 
characteristic colors, which may serve as tests for identifying 
the metals. 

In making the special tests for compounds of boron it is 
necessary first to convert the compound to boric acid by adding 
an acid to the substance to be tested. 

Materials: Borax, alcohol, manganese dioxide, turmeric 
paper, concentrated hydrochloric acid, solutions of cobalt 
nitrate, and copper sulphate. 

Directions: (a) Dissolve 10 grams of borax in 15 c.c. of 
boiling water in a small beaker and add 5 c.c. of con- 
centrated hydrochloric acid. Set the solution aside to cool 
and then filter out the crystals of boric acid which will have 
precipitated. (1) Describe the appearance of boric acid. (2) 
Write an equation to show how it was formed. (3) Why does 
the reaction take place? 

(&) Put part of the crystals obtained in (a) into an 
evaporating dish, add 1 c.c. of concentrated hydrochloric acid, 
!10 c.c. of alcohol, and ignite the solution by directing the flame 



376 LABORATORY EXERCISES 

of the burner for a moment into the dish. (1) What color is 
the flame? (2) What use may be made of this coloration of 
the alcohol flame? 

(c) Dissolve about 1 gram of borax in a test tube half full of 
hot water, add 5 c.c. of dilute hydrochloric acid and wet a 
strip of turmeric paper with the solution. Dry the paper at 
100°. This can be done conveniently by wrapping the wet 
paper around the upper part of a test tube containing about 
5 c.c. of water and by then boiling the water until the paper 
is dry. (1) What color is the dried paper? Touch the paper 
with a drop of ammonium hydroxide. (2) What color is pro- 
duced? (3) For what is this action on turmeric a test? (4) 
Why was the hydrochloric acid added? 

(tZ) Make a loop about 3 mm. in diameter in the end of a 
platinum wire by winding it around the tip of a pencil. Al- 
ternately heat the wire in the Bunsen flame and dip it into 
powdered borax, until a clear glassy bead is formed in the loop 
of the wire. Touch the bead with a glass rod which has been 
dipped into a solution of cobalt nitrate, taking care that only 
a very small drop of the solution reaches the bead. Heat the 
bead in the flame until clear. (1) What color is it? (2) Write 
an equation to show what happens to cobalt nitrate \v^hen 
heated; (3) an equation for the reaction taking place between 
this compound and borax when fused. (4) Of what does the 
bead consist? (5) What gives it the color? After use, the 
bead may be removed by dipping it, while hot, into water. 
The sudden cooling shatters the bead, which may then easily 
be scraped from the wire. 

Make another borax bead as before, touch it with a solution 
of copper sulphate, and fuse it in the oxidizing part of the 
flame. (6) What is the color of the bead? Heat the bead 
again in the reducing flame, that is, just at the tip of the blue 
inner cone. (7) What is the color of the bead? 

Repeat the experiment, using a minute quantity of man- 
ganese dioxide instead of the copper sulphate, and heating in 
both parts of the flame. (8) What are the colors? (9) What 
use is made of the formation of these colored beads? 



PHYSICAL PROPERTIES OF THE METALS. ALLOYS 

Experiment 94. Density of Metals. 

Reference : 528. 

Discussion : It is easy to distinguish samples of the common 
metals from one another by a determination of their densities, 
because the values of the latter vary to such an extent that 
very accurate determinations are not necessary. The method 
described in Experiment 30 can be conveniently used. If the 
metals in the form of sheets are available, samples about 6 
inches wide can be rolled so they will slip into a graduated 
cylinder. If sheet metal is not available any form of such 
a size that it can be placed in the cylinder can be used. 

Materials : Samples of several metals weighing from 50 to 
7e5 grams each. 

Directions: Identify the samples of the metals furnished 
by a determination of the density of each in the way described 
in Experiment 30 h. (1) Tabulate your results giving in each 
case in separate columns the name of the metal, its Aveight, its 
volume, its density as found by you, and the recorded density. 

Experiment 95. The Hardness of Metals. 

Eef erence : 529. 

Discussion : It is possible to discover which of two sub- 
stances is the harder by determining which one scratches the 
other. Tests of the hardness of metals can be most readily 
made by attempting to scratch the surface of one metal with 
the sharp edge of a piece of the other metal in the form of 
a sheet. 

Materials: Pieces of sheet aluminium, copper, lead, tin, 
nickel, iron, zinc. 

Directions: Attempt to scratch the surface of one metal 
with the sharp edge of another. Place the harder to the left 
of the softer metal. Attempt to scratch the surface of each 

378 



380 LABORATORY EXERCISES 

of these two metals with a third metal and place it in its 
proper place, the harder metal to the left of the next softer. 
Continue with the rest of the metals. (1) Write down the 
names of the metals in the order of hardness found and place 
after each the hardness as given in the table on page 443 of 
the text. 

Experiment 96. Alloys of Lead and Tin. 

Reference: 538. 

Discussion : The melting points of the alloys of lead and 
tin, which are used as solders; vary with the percentage of 
the two metals in the alloy. Pure tin melts at 232°. When 
the molten metal is allowed to cool, it behaves in a manner 
which is characteristic of other pure substances. The tem- 
perature falls rapidly uJitil the freezing point, which is the 
same as the melting point, is reached. If the metal is not 
allowed to lose heat too rapidly, the temperature remains 
constant during solidification because, during the change in 
state, heat is given off at such a rate that it overcomes the 
cooling effect of the surrounding air. When the metal has 
solidified it again falls rapidly in temperature. By noting at 
equal intervals of time the temperature of the metal when it 
cools and changes from the liquid to the solid condition, it 
is possible to determine its freezing point. If the temperatures 
are plotted on a diagram in Avhich time and temperatures are 
the coordinates, the curve obtained by connecting the points 
made for each observation will be found to have a section 
parallel to the axis of time. This temperature is the freezing 
point of the metal. 

Pure lead, which melts at 327°, exhibits the same behavior 
as tin when it changes from the liquid to the solid state. 

When a small amount of lead is added to tin, the molten 
mixture begins to solidify at a temperature below the freezing 
point of tin. As the amount of lead in the mixture is increased 
the point at which solidification begins continues to fall until 
the mixture is reached which contains 69 per cent of tin and 
31 per cent of lead. As more lead is added the temperature 
at which solidification begins rises. The mixture of 69 per cent 



382 LABORATORY EXERCISES 

of tin and 31 per cent of lead freezes at a definite temperature 
and in this respect resembles a pure compound; it is called 
the eutectic. 

Other mixtures do not freeze at a definite temperature; as 
the molten alloy cools, the constituent which is in excess begins 
to separate at a definite temperature ; the temperature con- 
tinues to fall because the composition of the liquid changes as 
the result of the separation of the pure metal. When the 
composition of the liquid becomes that of the eutectic the 
temperature at which freezing takes place becomes constant, 
namely, that at which the eutectic freezes. 

Materials: Tin, lead, high boiling oil ('^Nujol"), 360° 
thermometer. 

Directions: (a) In making determinations of the melting 
points of tin and its alloys with lead it is necessary to have 
enough of the metal to cover, when liquid, the bulb of the 
thermometer used. To determine this, place in a test tube 
just enough water to cover the bulb of the thermometer when 
it rests on the bottom of the tube. Measure the volume of 
the water. This number multiplied by 7.3 (the density of tin) 
will give the weight of tin required. (1) Weigh on a piece of 
paper on the platform scales approximately this amount of tin ; 
the weighing should be to 0.1 gram. The pieces of the metal 
should be of such a size that they will readily slip into a test 
tube. Put into a perfectly dry test tube about 5 c.c. of a high 
boiling oil (^'Nujol") and support the tube in a vertical posi- 
tion by means of a clamp and ring-stand at such a height that 
it can be heated by a burner. 

Place a few pieces of tin in the oil and heat the latter 
until the tin melts. Drop into the heated tube pieces of the 
metal as rapidly as they melt. When all the metal has been 
melted, cautiously insert a perfectly dry thermometer into the 
liquid. Do this by first touching the surface of the liquid with 
the tip of the thermometer and removing the latter. Insert the 
thermometer again a little deeper and remove it immediately. 
The third time the thermometer can be left in the liquid. Rest 
the thermometer on the bottom of the test tube and heat 
the oil and metal until the thermometer registers about 250°. 



384 



LABORATORY EXERCISES 



Read and (2) record rapidly the temperature, count four 
slowly, record the temperature, count four, record, etc. Con- 
tinue the readings in this way until the temperature has 
fallen to 150°. In recording the temperatures it is well to 
write on a piece of paper and to transfer the readings later 
to the note book. The eyes should be kept on the thermometer 



dbU 












' 










r 


















































































































































































9?i;^ 


















































ceo 








































































































































































































m" 
























































































































































































































































175" 
























































































































































































































































i^n" 



















































60 



r20 
Time in Seconds 
Fig. 15. 



180 



240 



as the temperature falls, and the record made on the fifth 
second. 

(3) Plot the temperatures against the time on Fig. 15, and 
draw a curve through the points. 

(b) (1) Calculate what weight of lead must be added to 

the weighed amount of tin in the test tube to produce the 

eutectic, which contains 69 per cent tin and 31 per cent lead. 

wt Tin 
(Lead to be added = — — — X 31). Weigh this amount to 0.1 

69 



386 LABORATORY EXERCISES 

gram. Heat the test tube containing the tin until the metal 
melts and add the lead slowly. Heat the tube to about 250° 
and (2) observe the temperatures at intervals of 5 seconds as 
it cools to 150°. (3) Plot the temperature on Fig. 15, and 
draw a curve through the points. 

(4) Record the correct values of the melting point of tin and 
that of the eutectic, and those you obtained. Add to your observed 
value 6° in the case of the melting point of tin and 3° in the 
case of the eutectic. These values are the so-called ''stem 
corrections," that is, the amounts that must be added to the 
observed temperatures to correct for the error due to the fact 
that the mercury in the stem of the thermometer was not heated 
to the temperature of the mercury in the bulb. 

(c) Weigh out twice as much lead as the amount used in 
{h) above and add it slowly to the melted eutectic. When all 
the metal is liquid, heat the mixture to 250° and (1) record 
the temperature at intervals of 5 seconds as before until it 
falls to 150°. (2) Plot the temperatures and connect the point 
to make the so-called cooling curve. (3) Give a reason for 
the fact that the slant of the part of the curve to the melting 
point of the eutectic is different in this case from that of the 
other two curves. (4) Calculate the percentage composition 
of the alloy which gave this cooling curve. (5) At what tem- 
perature did solid begin to separate from the molten alloy? 
(6) Are the results in accord with the diagram in Section 538 
of the text? 



CHEMICAL PROPERTIES OF THE METALS 

Experiment 97. Behavior of Metals with Water. 

Reference : 545. 

Discussion : The metals vary widely in their activity with 
water. Some, like sodium, decompose water rapidly at room 
temperature; others do not react even with steam at a high 
temperature. When a metal reacts with water, hydrogen and 
a hydroxide of the metal are formed. If these adhere firmly 
to the surface of the metal, reaction soon ceases, as they prevent 
contact between the metal and the water. If a metal is brought 
in contact with a second metal below it in the electromotive 
series, a so-called '^couple" is formed. One of the reasons 
why a metal is more reactive when coupled with a second one 
is as follows : As the more active metal reacts, the hydrogen is 
evolved from the less active metal; it does not prevent, there- 
fore, contact between the water and the active metal. A couple 
can be made by bringing into contact the two metals. It is 
usually made,. however, by adding a solution of a salt of the 
less active metal to the more active metal. The latter liberates 
the former and thus its entire surface becomes coated with 
the less active metal. Couples made in this way are more 
active than those prepared by simple contact of the two metals, 
because larger areas of the two metals are in contact. Mercury 
or copper is commonly used as the less active metal in making 
couples. 

Materials: Magnesium ribbon; small piece of aluminium, 
tin, zinc, copper, and lead ; solutions of mercuric chloride and 
of copper sulphate. 

Directions: In order to avoid mistaking for hydrogen the 
dissolved air liberated when the water used in the experiment 
is heated, boil in a beaker for about 1 minute about 250 c.c. 
of water. Pour about 10 c.c. of the boiled water into each 
of six test tubes and put them in the hot water in the beaker 

388 



390 LABORATORY EXERCISES 

which has been removed from the flame. Place a bit of magne- 
sium ribbon and small pieces of aluminium, tin, zinc, copper, 
and lead in each of the tubes respectively. (1) Is hydrogen 
given oft' in any case? 

Add to each tube, except the one containing zinc, 2 or 3 
drops of a solution of mercuric chloride. To the zinc add 
2 or 3 drops of a dilute solution of copper sulphate. (2) Which 
metals evolve hydrogen? (3) Arrange in a column the metals 
in the order of decreasing activity, as judged by the rate of 
evolution of gas. Place side by side with this arrangement 
the symbols of the metals in the order of their position in 
the electromotive series. (4) Is the order in the two arrange- 
ments the same? (5) Explain the lack of action with water 
alone. (6) Why did the addition of solutions of mercuric 
chloride and of copper sulphate activate the metals? 

Experiment 98. Action of Non-oxidizing Acids on Metals. 

References : 546, 547. 

Discussion : Aqueous solutions of non-oxidizing acids react 
with metals down to hydrogen in the electromotive series. 
Whether the metal reacts completely or not is determined by 
the properties of the salt formed. If it is insoluble in water 
and adheres to the surface of the metal in such a way that 
the latter can not come into contact with the acid, the action 
soon ceases. The chlorides of all the metals above hydrogen 
in the electromotive series, except lead chloride, are freely 
soluble in water ; all these metals except lead react in cold dilute 
hydrochloric acid. Lead chloride is soluble in hot water; as 
a consequence the metal reacts with hot hydrochloric acid. 

In a similar way the reaction of metals in dilute sulphuric 
acid is determined by the solubility of the sulphates of the 
metals. Lead sulphate is insoluble in water but is soluble in 
concentrated sulphuric acid ; the metal does not react with the 
dilute acid but does react with concentrated sulphuric acid. 

The behavior of aluminium appears to be abnormal as it 
is only slightly attacked by sulphuric or nitric acids. This 
is probably due to the formation of a protective coating of 
insoluble basic sulphate or nitrate. 



392 LABORATORY EXERCISES 

When a metal reacts with an acid to set free hydrogen, the 
salt formed is always that derived from the metal in its lower 
valence, since hydrogen is a reducing agent. 

Materials: Lead, iron filings, concentrated hydrochloric 
acid, solutions of ferrous sulphate and of sodium carbonate. 

Directions: (1) Make a table to indicate the behavior of 
the following metals with dilute hydrochloric acid and with 
dilute sulphuric acid: magnesium- aluminium, zinc, iron, tin, 
lead, and copper. If you are not sure of the action in any 
case, make the test. (2) In the cases where there is no action, 
state the reasons for the inactivity. 

Place a bit of lead in a few cubic centimeters of concen- 
trated sulphuric acid. (3) Does the metal react? (4) Why? 

To determine which chloride of iron (FeCL or FeClg) is 
formed when the metal dissolves in hydrochloric acid, add a 
few pieces of iron filings, or a small nail, to about 3 c.c. of 
concentrated hydrochloric acid. While the metal is reacting 
add a few drops of a solution of a ferrous salt to a solution 
of sodium carbonate. (5) Record the color of the precipitate. 
Repeat, using a ferric salt. (6) Record the color of the pre- 
cipitate. Next pour a few drops of the solution prepared by 
the action of iron in hydrochloric acid, into a solution of sodium 
carbonate. (7) What is the color of the precipitate? (8) Is 
the compound formed from the metal and hydrochloric acid a 
ferrous ov a ferric salt ? (9) From what fact would you have 
predicted this result ? 

Experiment 99. The Action of Oxidizing Acids on Metals. 

Reference : 548. 

Discussion : The behavior of a metal with an oxidizing acid 
is determined by the activity of the two substances involved. 
The higher the metal is in the electromotive series the more 
readily it is oxidized. The concentration of the acid is an 
important factor in its activity as an oxidizing agent. Since 
the ionization of the acid increases Avith dilution and since 
the undissociated acid alone brings about oxidation, the ac- 
tivity of the acid as an oxidizing agent increases as the concen- 
tration of the acid inci*eases. 



394 LABORATORY EXERCISES 

Materials: Tin, aluminium, iron nail, concentrated nitric 
acid, solutions of stannous chloride, stannic chloride, ferric 
chloride, ferrous sulphate, ammonium sulphide. 

Directions: (a) Put a small piece of tin in a test tube 
and add to it dilute nitric acid made by mixing 1 c.c. of concen- 
trated nitric acid with 10 c.c. of water. AIIoav the tube to 
stand for about one-half hour and meanwhile try the following 
test to distinguish a stannous from a stannic salt. 

Add to a solution of stannous chloride, SnCls, a few drops 
of ammonium sulphide. (1) Record the color of the precipitate, 
which is stannous sulphide, SnS. Repeat with a solution of 
stannic chloride, SnCl^. (2) Record the color of the precipitated 
stannic sulphide, SnSg. 

At the end of the half hour determine whether a stannous 
salt or a stannic salt is formed as the result of the action of 
dilute nitric acid on tin. (3) State what you did, the result, 
and the conclusion. 

(6) Place about 1 gram of tin in a test tube and add 5 c.c. 
of concentrated nitric acid. After reaction has ceased, add 
water to the tube until it is full and shake. The precipitate 
is an oxide of tin. Filter about 5 c.c. of the liquid into another 
tube and add a few drops of ammonium sulphide to the filtrate. 
(1) Has any tin passed into solution? (2) What is the evidence? 

Mix about 5 c.c. of the original milky liquid with an equal 
volume of a solution of sodium hydroxide and boil the mixture 
for about 1 minute. The oxide dissolves very slowly in a hot 
solution of sodium hydroxide, but it is not necessary to wait 
until it has all dissolved before proceding. Cool the test tube, 
add dilute hydrochloric acid until the solution is acid and then 
a few drops of ammonium sulphide. (3) Record the color of 
the precipitate. (4) What is it? (5) What is the valence of 
tin in this compound? (6) What do (a) and {h) in this experi- 
ment show in regard to the action of dilute and of concentrated 
nitric acid on tin? 

(c) Place a little steel wool or an iron nail free from rust in 
a test tube and cover it with about 10 c.c. of a solution made 
by mixing 1 c.c. of concentrated nitric acid with 20 c.c. of 
water. Set the tube aside for about 10 minutes and then add 



396 LABORATORY EXERCISES 

some of the solution to a little sodium carbonate solution. (1) 
Record the result. (2) Is a ferrous or a ferric salt formed? 
(3) What is the evidence? 

(d) Treat a little iron with a mixture of one part of nitric 
acid and five of water. If the metal does not appear to react 
heat the tube gently. Add a little of the solution to a solution 
of sodium hydroxide. (1) What was formed? Make a state- 
ment as to the action of concentrated and moderately dilute 
nitric acid on iron. 

(e) Heat a piece of aluminium with dilute and with con- 
centrated nitric acid. (1) Is there evidence of any action in 
either case? 

Experiment 100. The Effect of the Presence of Oxygen on 
the Action of Acids on Metals. 

Reference: 549, 550. 

Discussion : Metals that either do not displace hydrogen 
from acids or react with acids vei-y slowly are more or less 
rapidly attacked by acids in the presence of oxygen. It is- 
probable that the oxygen converts the metal at the surface 
into an oxide, which then dissolves in the acid. 

Materials: Strips of copper and of lead, dilute acetic acid. 

Directions: Place in a small beaker about 20 c.c. of dilute 
acetic acid and stand in the solution a strip of sheet copper 
so that a part of the metal is above the surface of the liquid. 
Arrange a similai' experiment using lead instead of copper. 
Set the beakei's in your desk and examine them at the next 
exercise. Test the two solutions to determine whether the 
metals have reacted by adding to each a few drops of am- 
monium sulphide. (1) What do you observe? (2) What does 
this prove? (3) Describe the appearance of the metals above 
and below the line at the surface of the liquid. (4) Give an 
explanation of the appearance. 

Experiment 101. The Action of Metals on Alkalies. 
Reference : 554. 

Discussion : A number of metals form hydroxides that dis- 
solve in both acids and alkalies. When the hvdroxide dis- 



398 LABORATORY EXERCISES 

solves in an acid it is playing the part of a base. When the 
hydroxide dissolves in an alkali it is playing the part of an 
acid. Most metals that form hydroxides which dissolve in 
alkalies are attacked more or less by alkalies. Aluminium, for 
example, reacts rapidly with a solution of sodium hydroxide, 
and the metal passes into solution as a salt which is called 
sodium aluminate. The same substance is formed when 
aluminium hydroxide dissolves in sodium hydroxide : 

A1(0H)3 + NaOH = NaAlO^ + H^O 

Other metals of this class react very slowly with alkalies. 
Some react only when the metal is heated to a high temperature 
with the solid alkali. 

The metals can be activated toward alkalies by coupling 
them with other metals, just as they are activated toward 
water and acids. When zinc is in contact with iron the former 
reacts rapidly with a warm solution of sodium hydroxide. 
Hydrogen is evolved from the iron and the zinc passes into 
solution as sodium zincate, Na.Zn02. Tin behaves in a similar 
way and sodium stannite, Na2Sn02, is formed. It should be 
noted that in these cases the metal that dissolves is the one 
which forms a hydroxide soluble in sodium hydroxide. When 
two metals are coupled in neutral or acid solution hydrogen 
is evolved from the metal having the lower solution pressure — 
that is, the one lower in the electromotive series. When the 
solution contains an alkali, this is not always the case. Zinc 
is above and tin is below iron in the electromotive scries, but 
when activated by iron in the presence of an alkali they each 
dissolve. The solubility of the hydroxide in alkalies is the 
determining factor in the reaction. The zinc that passes into 
solution is not presnt in the form of zinc ions, but as zincate 
ions, ZnOo~"" which result from the ionization of the salt 
formed. 

Materials : Aluminium ; zinc ; tin ; solutions of zinc sul- 
phate, ferrous sulphate, and ammonium sulphide. 

Directions: In separate test tubes add a small amount of 
the following metals to a solution of sodium hydroxide: 



400 LABORATORY EXERCISES 

Aluminium, zinc, and tin. (1) What is the action in each case? 
Heat the solutions. Drop into the tubes containing zinc and 
tin a clean small iron nail. (2) From which metal is the hy- 
drogen set free? 

To determine which metal has dissolved proceed as follows : 
After the reaction has taken place for about two minutes pour 
off the liquid from the tube containing the zinc and add to 
it a few drops of a solution of ammonium sulphide. (3) What 
is the color of the precipitate? To determine whether the sul- 
phide precipitated is that of iron or zinc add ammonium sul- 
phide to solutions of ferrous sulphate and of zinc sulphate. 
(4) What is the color of the precipitate in each case? (5) Did 
the iron or the zinc dissolve? (6) As what ion does the dis- 
solved metal exist? 

Pour off the liquid from the tin, make the solution just 
acid with dilute hydrochloric acid and add a few drops of 
ammonium sulphide. (7) What is the color of the precipitate? 
(8) What is it? (9) How do you know? (10) Which metal 
dissolved? (11) Why did it dissolve? 



ELECTROCHEMISTRY 

Experiment 102. The Production of an Electric Current 
from Chemical Energy. 

Reference : 566. 

Discussion: When zinc is placed in a solution of hydro- 
chloric acid the metal dissolves, hydrogen is set free, and heat 
is developed. When copper is placed in dilute hydrochloric 
acid no reaction takes place. If zinc that is dissolving in 
hydrochloric acid is brought into contact with copper, the 
metal continues to pass into solution, and hydrogen is evolved 
from the copper. A part of the liberated energy that appeared 
as heat in the first case appears a« electrical energy when the 
metals are in contact. 

The explanation of the reactions in the two cases, using 
the conception of electrons is as follows: When zinc dissolves 
in an acid the atoms that pass into solution give up their 
electrons to the hydrogen ions and hydrogen is given off : 

Zn+2H+ = Zn+++H2+Heat 

When zinc is in contact with copper and the former passes into 
solution, the electrons that were on the zinc atoms before solution 
occurred pass from the zinc to the copper. It is from the 
latter metal that the hydrogen ions take up the electrons 
necessary to convert them into hydrogen gas. 

In the experiment described below two cells made of zinc, 
copper, and sulphuric acid are set up and joined in series, 
that is, the copper of one cell is joined to the zinc of the 
next. In this way it is possible to obtain a current that will 
decompose solutions of acids, bases, and salts. 

Materials: Two copper and two zinc electrodes of the 
dimensions represented in Fig. 16, Four pinch cocks. A 

402 



404 



LABORATORY EXERCISES 



solution of mercuric chloride containing 1 gram of the salt to 
100 c.c. of water, of potassium iodide, sodium chloride, phenol- 
phthalein, litmus paper. 

Directions: Bend over to make a right angle the narrow 
end of one zinc and one copper electrode; slide a pinch cock 
over the end of one electrode, slip the other over the first, and 
release the pinch cock. Place one electrode in a small beaker 
and the second in another (see Fig. 16) ; place the second 
zinc electrode in a solution of 1 gram of mercuric chloride 
in 100 c.c. of water and let it stand for 3 minutes. Mercury 
is deposited evenly over the surface of the zinc and as a result 
the metal reacts very slowly with dilute acids. This procedure 



\3cm. 




Fig. 16. 



is used when zinc is employed in batteries to prevent the metal 
from dissolving as the result of the setting up of couples be- 
tween the zinc and the impurities in it. By the ** amalgama- 
tion" of the zinc a uniform amalgam of zinc and mercury is 
formed (567). 

Place the second copper electrode in the beaker containing 
the zinc, and the zinc electrode in the beaker containing the 
copper. Join to each of these electrodes by means of pinch 
cocks a copper wire which passes through a cork as indicated 
in the figure. 

(a) Add 10 c.c. of concentrated sulphuric acid to 200 c.c. 
of water, cool the solution, pour about 20 c.c. of it into a small 
beaker and the rest into the two beakers of the apparatus. 
(1) Record the action of the acid on the four electrodes. 

Bring together the two terminals that pass through the 



406 LABORATORY EXERCISES 

cork. (2) What happens on the electrodes? Separate the two 
terminals and hold the small beaker containing the sulphuric 
acid so that they dip under the surface of the liquid. (3) 
What occurs at the terminals? (4) Is the terminal from which 
the gas is given off connected with zinc or copper? (5) What 
is the gas? (6) Why is gas not given off from the other 
electrode? (7) What is the terminal called from which the 
gas is given off? (8) State the direction in which the electrons 
in the circuit are flowing, (9) In which way is the positive 
current flowing? 

(&) Place a few cubic centimeters of a solution of potassium 
iodide in a small beaker and place the latter so the terminals 
dip into the solution. (1) State what happens at the two 
electrodes. 

(c) Repeat the experiment using a solution of sodium 
chloride to which has been added a few drops of a solution 
of phenolphthalein. (1) State what is observed, and explain 
the result. (2) How could you determine which of the ter- 
minals of an electric circuit is the positive pole? 

(d) Moisten a piece of litmus paper with a solution of 
sodium chloride, and press the paper against the two terminals. 
(1) Explain what happens. 

(e) Cut off about 1 inch from the anode terminal and wind 
around it a piece of platinum wire. Leave the end of the wire 
free to use as the electrode. Dip the terminals into a solution 
of sulphui'ic acid. (1) Is oxygen evolved? (2) Why in this 
case and not when copper was used as the electrode? 

Experiment 103. Metallic Couples. 

Reference: 567. 

Discussion: The activation of metals toward water by 
means of couples was illustrated in Experiment 97. The student 
should reread the discussion of that experiment in the light of 
the experiment on the production of an electric current (Ex- 
periment 102). In the following experiment additional couples 
are studied. 

Materials: Nail, strip of zinc, tin, lead, copper and 
aluminium. 



408 LABORATORY EXERCISES 

Directions: (a) Wrap tightly a small piece of zinc around 
the head and another piece around the point of a wire nail 
free from rust. Attach two pieces of tin in a similar way 
to a nail. Place the nails side by side, but not touching each 
other, on the bottom of a beaker. Cover the nails with a layer 
of water about 3 cm. deep, and set them aside. At the next 
exercise examine the two nails, the zinc, and tin carefully. (1) 
Record the appearance of the iron in the two cases. (2) Give 
an explanation of what happened. (3) Is the appearance of 
the zinc and the tin in accord with the explanation? (4) What 
is the deposit on the zinc? (5) What practical conclusion can 
be drawn from the experiment? 

{h) Devise and carry out an experiment to determine if 
contact with (1) lead, (2) copper, or (3) aluminium affects the 
rusting of iron. (1) Record the results in the three cases. 

Experiment 104. Faraday's Law. 

Reference: 569. 

Discussion: Faraday's law states (1) that the quantity of 
any one substance decomposed by the electric current is pro- 
portional to the quantity of electricity passed through its solu- 
tion, and (2) that the quantities of two or more substances 
liberated by equal quantities of electricity are proportional to 
the chemical equivalents of these substances. The first part 
of the law can be tested by passing a constant current through 
a solution of a salt, such as coppei* sulphate, and determining 
the weight of copper deposited during different time intervals. 
When a constant cui'rent flows the quantity of electricity that 
passes is proportional to the time during which the current 
flows. 

The second part of the law can be tested in the manner 
outlined in the following experiment. The direct current which 
is ordinarily furnished is one at 110 to 115 volts. By placing 
in the circuit a 60-watt lamp the current is reduced, as the 
result of the resistance of the lamp, so that it serves to deposit 
copper evenly on an electrode. In order to test Faraday's 
law such a current is passed first through a solution of copper 
sulphate in which are placed copper electrodes and then 



410 LABORATORY EXERCISES 

through a solution of tin chloride, SnClg, in which are placcvi 
tin electrodes. As the current passes through the copper sul- 
phate solution, the metal dissolves at one electrode and is 
deposited at the other. An analogous change takes place in 
the case of tin. The amounts deposited in the two cases should 
be in the proportion of the atomic weights of the two metals, 
because in the solutions used the metals have the same valence. 
On account of the fact that it is difficult to deposit tin in a 
form that adheres to the electrode, and can, as a consequence, 
be easily weighed, the amount of tin that dissolves is determined 
by noting the change in weight of the electrode from which 
the metal is dissolved. If pure tin is used this loss is equal 
to the gain in weight of the other electrode. 

It should be remembered (563) in answering the questions 
given below that watts = volts X amperes, and that coulombs X 
seconds = amperes. 

Materials: Two electrodes made of thin sheet tin of the 
size indicated in Fig. 16. Two electrodes made of thin sheet 
copper of the same size. Copper wire. Four pinch cocks. 
Solutions of copper sulphate and of stannous chloride, each 
containing approximately 20 grams of the salt to 100 c.c. of 
water. 

Directions: Bend the smaller end of one of the copper 
electrodes over the edge of a small beaker, so that the 
electrode will hang in the beaker. The end outside the beaker 
should be bent into a horizontal position. Arrange the other 
copper electrode opposite the first. Place the two tin electrodes 
in a second beaker in the same way. Connect a 60-watt lamp 
with the source of electricity and with a copper wire which 
is joined to one of the copper electrodes in the following man- 
ner: Slip a pinch cock over the end of the wire, and rest the 
latter on the horizontal part of the electrode outside the beaker ; 
open the pinch cock, and move it until it is over the joint 
and then close it. By means of a short piece of wire and 
two pinch cocks connect the second copper electrode with one 
of the tin electrodes. Connect the second tin electrode with 
the second terminal from which the current is taken. Remove the 
electrodes and scratch upon them with the point of a file the 



412 LABORATORY EXERCISES 

numbers 1, 2, 3, and 4 respectively. Weigh each electrode to 
centigrams. Assemble the apparatus as before and cover the 
copper electrodes with the solution of copper sulphate described 
under ' ' Materials, ' ' above, and the tin electrodes with the solution 
of stannous chloride. Switch on the current and note the time. 
Observe closely what happens in the two beakers. The solution 
of copper sulphate should be stirred gently from time to time. 
If this is not done all the copper in the part of the solution 
touching the electrode on which the metal is depositing will 
pass out of solution and hydrogen will be evolved by the 
current. 

As the tin is deposited the crystals will grow out into the 
solution. Care should be taken to prevent contact between 
them and the other electrode. If the crystals grow out very 
far into the solution gently press them back to the electrode" 
from which they have grown. Be careful not to break them 
away from the electrode. 

At the end of from 45 minutes to 1 hour, break the current 
and note the time. Disconnect the electrodes. Hold the copper 
electrodes under running water for a few seconds and then 
dry them by pressing them against a clean towel. Wash and 
dry in the same way the tin electrode from which the tin was 
dissolved. Hang the tin electrode to which the crystals are 
attached in a beaker of water for a few minutes; remove the 
electrode and hang it a second time in water. Remove it and 
hang it on an empty beaker until it is dry. Weigh the four 
electrodes to centigrams. 

(1) Tabulate on the opposite page the original weights of 
the electrodes, their weights after the metals have been de- 
posited, and the gain or loss in weight of each electrode. (2) 
Is the weight of the copper dissolved equal to the weight of 
the metal deposited? (3) If not, give a reason for the dif- 
ference? (4) Answer the same questions in regard to the tin. 
(5) Which is probably the more accurate measure in the two 
cases of the current that passed through the solutions? (6) 
Calculate from the weight of the copper deposited what weight 
of tin should have been deposited according to Faraday's law. 
Write under this the weight found in the experiment. (7) 



414 LABORATORY EXERCISES 

Calculate from the weight of copper dissolved the weight of 
tin that should have been dissolved. Write under this the 
weight found in the experiment. (8) The weight of copper 
deposited should be the most accurate of the measurements 
made. Is the weight of the tin deposited or dissolved more 
in accord with the amount required by Faraday's law? (9) 
Calculate from the fact that a 60-watt lamp was placed in 
the circuit the current (amperes) that were used in the experi- 
ment. (The resistance of the solution is so small that it can 
be neglected.) (10) Calculate from the weight of copper 
deposited the number of coulombs that were used. (11) From 
this last result and the time required for the deposition, cal- 
culate the current in amperes. (12) Compare the result in 
(11) with that obtained in (9) above, and state which is the 
more accurate measurement. 

Experiment 105. The Lead Storage Battery. 

Reference : 578. 

Discussion : If an electric current is passed through a solu- 
tion of sulphuric acid between lead electrodes, hydrogen is 
evolved at one electrode and on the other a deposit of lead 
dioxide is formed. If the circuit is broken and the electrodes 
are connected by a wire, a positive current flows through the 
wire from the electrode on which the lead dioxide was de- 
posited. The dioxide is converted into lead sulphate. At 
the other electrode metallic lead is also converted into lead 
sulphate. The equation for the reactions involved is as follows : 
PbO^ + 2H2SO, + Pb = PbSO, + 2H2O + PbSO^ + electricity. 
If next a current from an outside source is passed between the 
electrodes, reactions take place which are the reverse of those 
just given; the equation for the reaction is that given above 
when read from right to left. To force the reaction in this 
direction energy is necessary and this is supplied by the cur- 
rent. When the electrodes are joined the reaction reverses 
itself and energy in the form of electricity is set free. The 
equation read from left to right represents the changes in- 
volved when a lead storage battery is furnishing a current; 
read from right to left it represents the changes involved when 
the battery is being charged. 



416 LABORATORY EXERCISES 

Materials: Two lead electrodes of the form and size de- 
scribed in Experiment 102, 2 pinch cocks, copper wire, volt- 
meter. 

Directions: Place the electrodes in a small beaker con- 
taining dilute sulphuric acid made by adding 10 c.c. of con- 
centrated sulphuric acid to 90 c.c. of water. Connect the elec- 
trodes by means of pinch cocks with two copper wires that 
are joined to a 110-volt circuit through a 60-watt lamp. Put 
on the current and observe at once what happens. Break the 
circuit and hold in contact with the two electrodes a piece of 
copper wire for about 5 seconds. Remove the wire, put on 
the current for about 5 seconds; break it, and connect the 
electrodes as before. Repeat charging and discharging in this 
way about 5 times. 

Examine the charged electrodes and (1) describe their 
appearance. Test the voltage of the charged cell by means 
of a voltmeter such as those used in testing dry cells. (2) 
What voltage do you find? (3) What is the voltage of the 
lead storage battery? Charge the cell, remove the electrodes 
and pour on the surface of each a few drops of dilute hydro- 
chloric acid and at once note the odor. (4) What do you 
detect? (5) From which electrode was the gas given oft'? (6) 
Write an equation for the reaction. (7) Write an equation 
for the reaction taking place when the cell is being charged. 



PROPERTIES OF HYDROXIDES AND SALTS 

Experiment 106. The Effect of Ammonium Salts on the 
Precipitation of the Hydroxides of the Bivalent Metals by 
Ammonium Hydroxide. 

Reference : 589. 

Discussion : The hydroxides of most of the bivalent metals 
which are precipitated by a solution of ammonium hydroxide 
are not precipitated by this reagent if an ammonium salt is 
present in the solution. The extent of the ionization of am- 
monium hydroxide is small and when an ammonium salt is 
added to the solution the ionization of the hydroxide is reduced 
as the result of the presence of the added ammonium ions. 
Such an addition forces the equilibrium represented by the 
equation 

NH40Ht=>NH4++OH- 

in the direction which results in the formation of undissociated 
ammonium hydroxide. As a consequence, the concentration 
of the hydroxyl ions is reduced to such an extent that the 
solution no longer causes the precipitation of the hydroxides 
of the bivalent metals, on account of the fact that the latter 
compounds are slightly soluble in water. 

The solubility of hydroxides of the trivalent metals is ex- 
ceedingly small ; as a consequence the hydroxides of the triva- 
lent elements are precipitated by ammonium hydroxide in the 
presence of ammonium salts. 

Materials: Solutions of calcium chloride, magnesium 
chloride, ammonium chloride. 

Directions: (a) In a test tube add ammonium hydroxide 
to a solution of calcium chloride. Add a solution of sodium 
hydroxide to one of calcium chloride. (1) What happened in 
each case? (2) What is thus proved about the relative ioniza- 

418 



420 LABORATORY EXERCISES 

tion of ammonium hydroxide and of sodium hydroxide? (3) 
How did you arrive at this conclusion? 

(h) Repeat (a) above using magnesium chloride instead of 
calcium chloride. (1) What happened? (2) Is the hydroxide 
of calcium or of magnesium the more soluble? (3) What 
is the evidence? 

Add a solution of ammonium chloride to the tube to which 
ammonium hydroxide was added. (4) What occurred? 

Add to a solution of magnesium chloride an equal volume 
of ammonium chloride solution and then ammonium hydroxide. 
(5) What was the result? (6) What is the effect of ammonium 
chloride upon ammonium hydroxide? (7) What is the reason 
for the result in this experiment? (8) What is the reason 
for (4) ? 

(c) Fit a filter paper in a funnel ready for use. Add to 
a solution of ferrous sulphate an equal volume of a solution 
of ammonium chloride. Add ammonium hydroxide to the mix- 
ture and filter at once. Ferrous salts are oxidized in solution 
by the air; as a consequence it is necessary to filter off the 
ferric hydroxide precipitated. (1) What is the appearance of 
the filtrate as it drops from the funnel? Shake the filtrate. 
(2) What occurred? The ferrous salt is oxidized by the air 
and the ammonium hydroxide present causes the precipitation 
of ferric hydroxide. (3) Will ammonium chloride prevent the 
precipitation by ammonium hydroxide of ferrous hydroxide? 
(4) Of ferric hydroxide? (5) What is the evidence? 

Experiment 107. Isomorphism of Salts. 

Reference : 591. 

Discussion : When two salts that do not react with each 
other are allowed to crystallize from the same solution, crystals 
of the two are usually deposited side by side. If, however, 
the two salts resemble each other closely in crystalline form 
they form mixed crystals, the composition of which is deter- 
mined by the proportions of the two salts present in the solu- 
tion. When two salts form such mixed crystals, they are 
said to be isomorphous. 

Ordinary alum has the composition represented by the 



422 LABORATORY EXERCISES 

formula K2S04,Al2(S04)3,24H20. Chrome alum has a similar 
formula, but contains chromium instead of aluminium, 
K2S04,Cr2(S04)3,24H20. Both salts crystallize in octahedra. 
The aluminium salt is colorless and the chromium salt is purple. 
When mixtures of the two are allowed to crystallize together 
uniform mixed crystals are obtained, the color of which is 
determined by the proportions of the two salts in the solution. 

Materials : Potassium aluminium alum, potassium chromium 
alum. 

Directions : (a) Dissolve 16 grams of potassium aluminium 
alum in 100 c.c. of hot water, filter if necessary, and cool the 
solution to room temperature in running water. Shake occa- 
sionally for five minutes 8 grams of finely powdered potassium 
chromium alum with 100 c.c. of water at about 30°. The 
temperature should not be allowed to go higher than 35° in 
order to prevent the change of the chromium salt to the green 
modification. Filter the solution. In four beakers place the 
following: (1) 40 c.c. of the solution of aluminium alum and 
10 c.c. of the chromium alum, (2) 20 c.c. of the former and 
10 c.c. of the latter, (3) the rest of the solution of the former 
and (4) the rest of the solution of the latter. 

(h) Set aside the solutions and examine them at the next 
exercise. Remove some of the crystals from each of the beakers 
and place them side by side on a piece of filter paper. (1) 
Are the crystals from each beaker uniform in appearance? 
Examine closely the crystalline form of the crystals. (2) How 
many faces are there on a single crystal? (3) State what form 
they have. (4) Compare the colors of the crystals. (5) What 
inference can you draw from the colors? (6) What is meant 
by isomorphous salts? (7) Can isomorphous salts be separated 
from each other by crystallization? 

Experiment 108. The Heat of Solution of Salts. 

Reference: 592. 

Discussion: When a substance is dissolved in water there 
is a change in temperature of the solvent. The number of 
calories set free or absorbed when 1 gram-molocular-weight of 
a salt dissolves in such an amount of water that further addi- 



424 LABORATORY EXERCISES 

tion of the solvent produces no heat change, is called the 
molar heat of solution of the salt. In the case of most salts 
the heat of solution is negative, that is, heat is absorbed and 
the temperature falls. The molar heats of solutions of salts 
vary widely. The experiment described below is designed to 
illustrate the variation in this quantity. 

Materials: Potassium nitrate, sodium chloride. 

Directions: (a) Grind in a mortar to a fine powder about 
11 grams of potassium nitrate. (1) Weigh on a piece of paper 
on the platform scales 10.1 grams of the powdered salt. 
Measure into a 500 c.c. beaker 250 c.c. of water at room tem- 
perature and (2) read the temperature of the water to 0.1 
degree. Pour the potassium nitrate into the water and stir with 
the thermometer. The salt should dissolve in less than 30 
seconds. (3) Read the temperature when the salt has dissolved. 
(4) From the fall in temperature and the volume of the solution 
calculate the number of calories absorbed. Since the fall in 
temperature is small, the error due to the absorption of heat 
by the beaker and thermometer can be neglected. (5) What 
is meant by the molar heat of solution of a salt? (6) Calculate 
the molar heat of solution of potassium nitrate. 

(h) Repeat the experiment described in {a) above replacing 
potassium nitrate by sodium chloride. Use 5.8 grams of the 
salt and 250 c.c. of water. (1) Record the results. (2) Cal- 
culate the molar heat of solution of sodium chloride. 

Experiment 109. Effect of Temperature on the Solubility 
of Salts. 

Reference : 592. 

Discussion : The solubility of most substances increases 
with rise in temperature, but the rate of change varies greatly 
with different substances. In general, if the heat of solution 
of a salt is negative, that is, if the temperature falls when it 
dissolves in water, the salt is more soluble in hot than in cold 
water. Further, the change in solubility with change in tem- 
perature is greater in the case of a salt with a high heat of 
solution than in the case of one which produced a small heat 
effect when dissolved. In the previous experiment (number 



42G LABORATORY EXERCISES 

108) it was found that there was a marked difference between 
the molar heats of solution of potassium nitrate and sodium 
chloride. The following experiment serves to illustrate the 
relation between the heat of solution of a salt and the change 
in the solubility of the salt with rise in temperature. 

Materials: Potassium nitrate, sodium chloride. 

Directions: (a) Grind to a powder 35 grams of potassium 
nitrate. Weigh the salt to decigrams on a piece of paper on 
the platform scales. Measure into a small flask or beaker 
35 c.c. of water and set it aside. Fit a stopper to a 100 c.c. 
graduated cylinder so that when the latter is shaken the water 
will not leak out. Dry the graduate carefully and pour into 
it the potassium nitrate. Tap the cylinder on a book until 
the upper surface of the salt is level, and read the volume 
of the salt to 0.1 c.c. Tap the cylinder smartly again and 
read the volume again. Repeat the tapping until the volume 
remains constant. (1) Record the volume of the potassium 
nitrate. 

Add to the graduate the 35 c.c. of water previously 
measured, insert the stopper, and shake vigorously for one-half 
minute. Let the salt settle, and read its volume. Shake again 
for about one-half minute and read the volume a second time. 
(2) Record the volume of the salt and the temperature of the 
solution. (3) What is the volume of the potassium nitrate 
that dissolved in the 35 c.c. of water used? (4) From the 
original volume of the dry salt and its weight calculate the 
weight of 1 c.c. of it. (5) From this value and the volume 
of the salt that dissolved calculate the weight of potassium 
nitrate that dissolved in the 35 c.c. of water at the temperature 
used. (6) What would dissolve in 100 c.c. of water? 

(h) Place the cylinder in a beaker of water and heat until 
the temperature of the solution in the graduate is about 40°. 
Remove the graduate and shake it vigorously for one-half 
minute, let the salt settle, (1) measure its volume and imme- 
diately observe and record the temperature of the solution. 
(2) Calculate as before the solubility of potassium nitrate in 
100 c.c. of water at the recorded temperature. 

(c) Replace the graduate in the hot water and heat it until 



428 LABORATORY EXERCISES 

the temperature is about 55°. Proceed as before, (1) record- 
ing the volume of the salt and the temperature, and (2) cal- 
culating the solubility. 

{d) Plot the solubilities against temperatures in the way 
given in the diagram in 592 of the text. Mark also on your 
plot the correct values for the solubilities at these temperatures 
as obtained by inspection of the diagram in the text. 

(e) In the way described above determine the solubility of 
sodium chloride at room temperature and at about 55°. Use 
in this case 25 grams of salt. Plot the curve as before and 
also the values obtained from the diagram. (1) What relation 
exists between the relative heats of solution of potassium nitrate 
and sodium chloride, as determined in Experiment 108 and the 
relative effects of rise in temperature upon solubility as found 
in this experiment? 

Experiment 110. The Solubility of Salts in Acids. 

Reference : 593. 

Discussion : When an insoluble salt is treated with an acid, 
double decomposition may or may not take place depending 
upon the solubilities of the original salt and of that which may 
be formed, and upon the relative degrees of ionization and of 
solubility of the acid added and of the acid produced. For 
example, if the salt AB is treated with the acid HD, the reaction 
expressed by the equation 

AB + HD-^AD + HB 

may or may not take place, depending upon the solubilities of 
the salts AB and AD and upon the ionization and solubilities 
of the acids HD and HB. The properties of all four substances 
are involved. 

Or viewed wholly from the ionic standpoint, if the concen- 
trations of ions of B, depending upon the solubility of AB, and 
of ions of 2), depending upon the degree of ionization of HD, 
are great enough to cause the continued formation of HB, the 
double decomposition will take place. 

In the cases studied in this experiment, the salts represented 
by AD are to be highly soluble. The effect of variation in tlie 



430 LABORATORY EXERCISES 

solubility of the salt AB, the effect of the ionization of the acid 
HD, and the effect of the solubility of the acid HB will be 
exemplified. 

The action of a highly ionized acid on the salt of a poly- 
basic acid, with the formation of a hydrogen salt of the latter, 
depends upon the same general principle as that stated above, 
although the case may appear slightly more complex. The 
action of nitric acid upon calcium phosphate serves as an illus- 
tration. 

Calcium phosphate is slightly soluble in water and dis- 
sociates into its ions : 

(1) Ca3(P04)2 ^ 3Ca+++ 2PO4— 

The above equation indicates an equilibrium that exists be- 
tween the ions of the dissolved salt and the undissolved solid. 
When nitric acid is added, the ions of the acid are intro- 
duced into the solution. 

(2) HN03^H++N03- 

(3) 2H++PO4— ;=^H2P04- 

The equilibrium in equation (1) above is disturbed as the 

result of the removal of PO4 ions. As a consequence more 

of the solid passes into solution and finally it all dissolves. 

The solution now contains the following ions: Ca"'""'", H+, 
H2PO4", NO3-. If a base is now added to the solution, the H-^ 
and the H2PO4" ions disappear as the result of neutralization. 
The H2PO4" ion reacts with the base as follows : 

2Na++ 20H-+ H2PO4- = 2Na++ 2H2O + PO4 

The calcium ions and the PO4 ions unite and Ca3(P04)2 

precipitates. 

This explanation makes it evident why bases precipitate 
from the solutions in acids of slightly soluble salts, the salt 
itself and not the hydroxide of the element. 



432 LABORATORY EXERCISES 

Materials: Solutions of zinc chloride, copper sulphate, am- 
monium sulphide, sodium carbonate, and acetic acid. 

Directions: (a) Prepare in a test tube some zinc sulphide 
by diluting 5 c.c. of a solution of zinc chloride with 5 c.c. of 
water and adding 5 c.c. of ammonium sulphide. Pour half the 
suspended precipitate into another test tube. Add dilute hydro- 
chloric acid to one tube and acetic acid to the other. (1) What 
happened in each case? (2) Upon Avhat difference in the two 
acids used does the difference in their action depend? 

{h) Prepare some copper sulphide from copper sulphate in 
the same way you made zinc sulphide. Try the action of 
hydrochloric acid upon it. (1) What happened? (2) Why 
did the result differ from that with zinc sulphide and hydro- 
chloric acid? 

(c) Prepare some zinc carbonate from zinc chloride by add- 
ing sodium carbonate to the solution. Treat the zinc carbonate 
with acetic acid. (1) What happened? (2) What is the reason 
for the difference between this result and that obtained when 
zinc sulphide was treated with acetic acid in part a? 

(d) Add a solution of sodium phosphate to a solution of 
calcium chloride. (1) Write an equation for the reaction. 
Add to the solution dilute nitric acid until the precipitate dis- 
solves. (2) Write an equation for the reaction using ionic 
symbols. 

Add sodium hydroxide to the solution until precipitation 
is complete. Determine as follows whether the precipitate is 
calcium hydroxide or calcium phosphate. Filter the solution, 
wash the precipitate with water until the filtrate no longer 
gives a test for a phosphate. Dissolve a part of the precipitate 
in dilute nitric acid and test for a phosphate. (3) How did you 
test for a phosphate? (4) What was the result with the nitric 
acid solution of the precipitate? (5) What was the precipitate? 

Experiment 111. The Solubility of Salts in a Solution of 
Ammonia. 

Reference : 595. 

Discussion: The hydroxides and insoluble salts of certain 
metals dissolve in a solution of ammonia as the result of the 



434 LABORATORY EXERCISES 

formation of soluble compounds which yield complex ions of 
the metal. Copper hydroxide, Cu(0H)2, for example, is in- 
soluble in water; it dissolves in a solution of ammonia as the 
result of the formation of a soluble compound which has the 
composition Cu(NH3) 4(011)2. The ion Cu(NH3)^ + + gives a 
deep blue color to the solution. Similarly silver chloride forms 
the soluble compound Ag(NH3)2Cl. 

Materials: Solutions of copper and of zinc sulphates, of 
silver and of nickel nitrates. 

Directions: Add a few drops of sodium hydroxide in 
separate test tubes to solutions of a copper, silver, nickel, and 
zinc salt. (1) Write equations for the reactions that take place. 

Add a solution of ammonia to each of the tubes. (2) Record 
the results. (3) Write equations for the reactions. 

Experiment 112. Preparation of a Double Salt: Ferrous 
Ammonium Sulphate. 

Reference: 598, 750. 

Discussion : The so-called double salts are compounds 
which are formed as the result of the chemical combination 
of two salts. They are usually prepared by dissolving the two 
constituents in water and allowing the crystals to deposit from 
the solution. The composition of a double salt is definite; it 
does not vary with the proportions of the constituents present 
in the solution as is the case with the composition of the 
crystals obtained when mixtures of isomorphous substances 
are crystallized (see Experiment 107). 

Double salts break down more or less completely into their 
constituents when dissolved in water. The presence of the 
separate ions of the constituents of the salt can be shown by 
appropriate tests. 

Materials : Ferrous sulphate, ammonium sulphate, solution 
of sodium carbonate, materials needed to test for an ammonium 
salt and for a sulphate. 

Directions: Add 28 grams of crystalline ferrous sulphate, 
FeS04,7H,0 to 20 c.c. of water; heat just to boiling and filter 
hot if necessary. Add 13 grams of ammonium sulphate to 
18 c.c. of water; heat and filter the hot solution if necessary. 



436 LABORATORY EXERCISES 

Mix the two solutions and add 5 drops of concentrated sul- 
phuric acid. If a salt crystallizes out, warm the solution until 
it dissolves. Cover the beaker containing the solution with a 
piece of filter paper and set it aside until the next exercise. 
Filter off the crystals and spread them on a paper to dry. (1) 
Describe the crystals. (2) What are they called? (3) To what 
class of substances do they belong? (4) Are they a mixture 
or a pure substance? (5) How could this be proved? 

Dissolve a few of the crystals in water and test separate 
portions of the solution for a sulphate and for an ammonium salt. 
(6) State how the test for a sulphate w^as made and the result 
obtained. (7) State how the test for an ammonium salt was 
made and the result obtained. 

Add to a solution of sodium carbonate a few drops of a 
solution of ferrous sulphate. (8) Record the result. Add to 
a solution of sodium carbonate a few drops of the solution of 
the double salt. (9) Record the result. (10) Does the latter 
give the test for a ferrous salt? (11) What conclusion can 
you draw as to the behavior of double salts when they are 
dissolved in water? 

Experiment 113. Properties of Complex Salts. 

Reference : 598. 

Discussion : Complex salts are formed as the result of the 
union of two salts. They differ from double salts in that they 
do not break down in solution into the ions of the salts fi'om 
which they are formed. Potassium feri-ocyanide is an example 
of a complex salt. It can be prepared by dissolving ferrous 
cyanide in a solution of potassium cyanide : 

Fe(CN)2 + 4KCN = K,Fe(CN)e 

In Experiment 112 it was shown that the double salt formed 
from ammonium sulphate and ferrous sulphate broke down 
in solution into its constituents; the solution gave the test for 
a ferrous salt. In the experiment described below the complex 
cyanide of iron and potassium is studied. 

Materials: Potassium ferrocyanide, solutions of ferrous sul- 
phate and of sodium carbonate. 



438 LABORATORY EXERCISES 

Directions: Dissolve a small bit of a crystal of potassium 
ferrocyanide in water and add a few drops of the solution to 
a solution of sodium carbonate. (1) Is a precipitate formed? 
Add a few drops of a solution of ferrous sulphate to a solution 
of sodium carbonate. (2) What is the appearance of the 
precipitate? (3) Does the solution of the complex salt show 
the presence of a ferrous ion? 

Heat in a crucible a small crystal of potassium ferrocyanide 
until it turns black. When cold remove the solid, place it in 
a test tube, cover it with water and add a few drops of hydro- 
chloric acid and warm the solution. Add a little of the solution 
to a solution of sodium carbonate. (4) What was formed? 
(5) What conclusion can you draw as to the presence of iron 
in the complex salt? (6) What did heating do to the complex 
salt? (7) How does a complex salt differ from a double salt? 

Experiment 114. Hydrolysis of Salts. 

Reference: 600. 

Discussion : Neutralization is the name given to the reac- 
tion between an acid and a base as the result of which a salt 
is formed. The reverse of this, the reaction of water with 
a salt to form an acid and a base, is called hydrolysis. A salt 
is hydrolyzed by water if the acid from which it is prepared 
is weak, that is, if it is but slightly ionized in water. The hydrol- 
ysis is brought about as the result of the union of the hydrogen 
ions furnished by the water and the negative ions of the acid 
to form the undissociated acid. P^or example, the salts of acetic 
acid, which is a relatively weak acid, are hydrolyzed. 

The reactions which take place between the ions in a solu- 
tion of sodium acetate may be indicated as follows : 

H20^H+-fOH- 

+ 
NaC2H302 ^ C2H30-+Na-^ 

it 
HC.H.O 

The removal of hydrogen ions in this way leaves hydroxyl ions 
in the solution, which, accordingly, shows a basic reaction. 



440 LABORATORY EXERCISES 

The salts prepared from weak bases are hydrolyzed in a 
similar way ; in this case, however, as the result of the forma- 
tion of the undissociated base, the solution shows an acidic 
reaction. 

The extent to which hydrolysis of a salt takes place is 
determined (1) by the strength of the acid and of the base 
from which it is prepared, (2) the concentration of the solil- 
tion, and (3) the temperature. The weaker the acid or base 
the greater the amount of hydrolysis. Hydrolysis also in- 
creases with increased dilution and with rise in temperature. 

Whether an acid is weak or strong can be told by testing 
its aqueous solution with certain indicators. Litmus is very 
sensitive to hydrogen ions; such a weak acid as carbonic acid 
will change blue litmus paper to red. Congo red, on the other 
hand, requires a higher concentration of hydrogen ions to affect 
it; strong acids like hydrochloric acid change the dj'^e to a 
deep blue color; it is not affected by carbonic acid. 

Litmus is very sensitive to hydroxyl ions, and v/ill show 
when hydrolysis takes place to only a very slight degree. When 
phenolphthalein changes from a colorless to a red compound a 
higher concentration of hydroxyl ions is required than in the 
case of litmus. 

Matp:rials: Solutions of sodium chloride, sulphate, sulphide, 
carbonate, and acetate, of carbonic, acetic and concentrated 
hydrochloric acids, of calcium and ferric chlorides, of copper 
sulphate, and of phenolphthalein; litmus paper; Congo paper. 

Directions: (a) Determine by testing with litmus paper 
whether solutions of the following salts show a neutral, acidic, 
or basic reaction : sodium chloride, potassium nitrate, ^ sodium 
sulphate, calcium chloride, sodium sulphide, copper sulphate, 
ferric chloride, sodium carbonate, and sodium acetate. (1) 
Tabulate the results and state in the case of each salt the 
strength (weak or strong) of the acid and the base from 
which the salt is derived. 

(&) Test a dilute solution of hydrochloric acid with Congo 
red paper. (1) Tabulate the results of this and the following 
tests. Test a solution of carbon dioxide in water (carbonic 
acid) with blue litmus paper and with Congo red paper. (2) 



442 LABORATORY EXERCISES 

Explain the result. Test a solution of hydrogen sulphide with 
the two indicators. (3) Explain the result. (4) Are these 
results in accord with the conclusions drawn from the experi- 
ments in (a) above? 

Add 5 drops of glacial acetic acid to 10 c.c. of water, and 
5 drops of concentrated hydrochloric acid to 10 c.c. of water. 
Test the two solutions with blue litmus paper and Congo red 
paper. (5) What are the relative strengths of hydrochloric 
acid and acetic acid? 

(c) In a clean beaker boil for 15 minutes a filtered solution 
of 2 grams of copper sulphate dissolved in 50 c.c. of water. 

(1) Note carefully whether a small amount of solid separates. 

(2) Is the result in accord with the behavior of copper sulphate 
in (a) above? (3) What is the deposit? (4) What does the 
experiment show about the effect of rise in temperature on 
hydrolysis? (5) Give a reason for your answer. 

(d) Test 10 c.c. of a solution of sodium acetate with pink 
litmus paper and with 3 drops of a solution of phenolphthalein. 
(1) Explain the results. Heat the solution just to boiling. (2) 
Explain the result. Cool the solution. (3) Explain the result. 

(e) Add to 10 c.c. of a solution of ferric chloride a drop 
of a dilute solution of sodium hydroxide and shake the tube. 
Continue the addition of the alkali until after shaking a very 
slight precipitate remains. Filter the solution. Test it with 
litmus paper. (1) Is ferric chloride hydrolyzed? Add to the 
solution an equal volume of a solution of sodium acetate and 
heat to boiling. (2) What happened? Ferric acetate is first 
formed and then hydrolyzes. (3) Explain why in this case 
the hydrolysis is complete. 



SODIUM AND POTASSIUM 

Experiment 115. The Action of Air Upon Sodium. 

Reference: Experiments 8 a, 20, 41. 

Discussion : A study of the action of air upon sodium brings 
together in review a number of facts which have already been 
taken up in previous experiments. These are the composition 
of the air, the action of certain metals with water, deliquescence, 
the action of carbon dioxide with bases, hydrates, efflorescence, 
and the formation of acid salts. 

The great chemical activity of sodium is shown by the 
formation of the peroxide instead of the oxide when the metal 
is burned in the air. 

Materials: Sodium, starch iodide paper. 

Directions: Precaution: Handle sodium with the pincers. 
Bo not allow water to touch it. (a) Remove the oil from a 
small freshly cut piece of sodium by touching each side of the 
piece to filter paper. Scrape the coating from one side of the 
piece. (1) What is the appearance? (2) Is it permanent? (3) 
What is the hardness of sodium compared to common metals? 

Place the piece of sodium on a watch glass and observe 
closely the changes taking place on its surface. (4) Recalling 
the composition of the air and a previous experiment in which 
you used sodium, Avhat substance do you expect to see formed? 

(5) What evidence is there of the formation of this substance? 

(6) Write an equation for the reaction. (7) In what connection 
have you previously studied this reaction? 

Put the sodium away on a shelf in your desk until the 
next exercise. (8) What is the appearance now? (9) What 
is the compound on the watch glass? (10) How do you account 
for its present state? (11) What is the technical name of the 
property which caused the substance to be in this condition? 

Again set the watch glass and contents aside in your desk 
for one or more exercises, until the material appears as well 

444 



446 LABORATORY EXERCISES 

defined crystals. (12) Recalling once more the composition of 
the air, state what these crystals probably are. Transfer a 
small amount of the material to another watch glass and apply 
a simple test which would tend to corroborate your supposition, 
if true. (13) What did you do? (14) What was the result? 
(15) Write an equation for the reaction by which the crystalline 
substance was formed. 

Put the remaining crystals back in your desk and examine 
them at each exercise until another distinct change has taken 
place. (16) What is now the appearance of the substance? 
(17) In what experiment have you observed a substance 
undergo a similar change in appearance? (18) What was the 
phenomenon called? (19) What further chemical reaction may 
this substance undergo if exposed to the air for a sufficiently 
long time? 

{h) Place a piece of sodium on the inverted cover of a 
crucible and direct the flame of a burner down upon it until 
it ignites. When it has cooled, add a few drops of water and 
touch the solution with a piece of starch iodide paper, prepared 
as in Expei-iment 28. (1) What Avas the action on the paper? 
(2) What is thus proved to have been formed when the sodium 
burned? (3) Write an equation for the reaction. (4) What 
property of sodium does this reaction illustrate? 

Experiment 116. Baking Powder. 

Reference : 612. 

Discussion : leaking powders are mixtures of sodium bicar- 
bonate and some substance which, by producing an acid solu- 
tion on the addition of water, Avill evolve carbon dioxide from 
the carbonate. Starch is also added to prevent reaction during 
storage and to dilute the baking powder, so that it may be 
evenly mixed with the flour when used. The choice of the 
acid-forming ingredient is dependent upon several considera- 
tions. It must be soluble enough to act quickly, but not too 
quickly; it must leave a final product of reaction which is 
not harmful; and the cost must be taken into account. The 
efficiency of the powder will depend also upon the thoroughness 
with Avhich the ingredients are powdered and mixed. 



448 LABORATORY EXERCISES 

Materials: Sodium bicarbonate, potassium acid tartrate, 
sodium aluminium alum, primary calcium phosphate, starch, 
commercial baking powder. 

Directions: (a) Dissolve a quarter of a spatulaful of cream 
of tartar, KH(C4H40y), in a test tube half full of warm water 
and test the solution with litmus paper. (1) What was the 
result? (2) To what class of compounds does it belong? 

Repeat the experiment using primary calcium phosphate. 
(3) What was the action of the solution with litmus? 

Likewise try the action of a solution of the alum of the 
composition Na,S04,Al2(SOj3,24H20 with litmus. (4) What 
was the action? (5) What kind of process did the alum 
undergo in solution? (6) Write an equation for the reaction. 
(7) What properties make these substances suitable for making 
baking powders? 

{h) (1) Write an equation for the reaction between cream 
of tartar and sodium bicarbonate. (2) Calculate the weight of 
of the former which will react with 5 grams of the latter, and 
(3) the weight of the carbon dioxide produced. (4) Calculate 
the weight of starch Avhich must be added in order that the 
carbon dioxide shall be 12 per cent of the combined weights 
of the three ingredients. 

Prepare the baking powder by thoroughly mixing in a 
mortar the ingredients in the amounts calculated above. 

(5) Write an equation for the reaction taking place when 
alum baking powder acts. (6) Calculate the weight of alum 
necessary to use with 5 grams of sodium bicarbonate and the 
Aveight of starch to be added to make a mixture Avhich will 
produce 12 per cent by weight of carbon dioxide. (7) Write 
an equation for the reaction in a phosphate baking powder 
and (8) calculate as above the weights of the ingredients to 
be used. (9) Calculate from the prices of the ingredients the 
cost of one pound of each kind of baking powder and record 
the prices of commercial baking powders. 

(c) Add a pinch of your baking powder to a test tube 
half full of cold water and observe the rate of action. Repeat 
the experiment using hot water instead of cold. Try both of 
these experiments, using commercial baking powder instead of 



450 LABORATORY EXERCISES 

your mixture. (1) What differences in rate of action did you 
observe? (2) How do you account for these differences? 

Experiment 117. Preparation of Sodium Carbonate. 

Reference : 610. 

Discussion: As was recently stated in the experiment on 
the preparation of boric acid from borax, a compound may be 
obtained by double decomposition when it is sufficiently in- 
soluble to be precipitated under the conditions of the reaction. 
Thus sodium acid carbonate, which is much less soluble than 
sodium chloride, will be precipitated from a saturated solution 
of sodium chloride upon the addition of ammonium acid car- 
bonate. The ammonium acid carbonate, however, can not be 
added in solution, since the additional water would serve to 
keep the sodium acid carbonate also in solution. Therefore 
carbon dioxide is passed into a saturated solution of sodium 
chloride which contains ammonia. 

Materials: Sodium chloride, marble, lime water, phenol- 
phthalein. 

Directions: (a) Saturate 10 c.c. of ammonium hydroxide 
solution with sodium chloride in a large test tube by adding 
an excess of the salt, shaking the mixture vigorously, and then 
filtering out the undissolved remainder. Set up a generator 
for carbon dioxide as in Experiment 38 h, and pass the gas 
into the solution until a heavy precipitate is obtained. Filter 
out the precipitate and dry it by pressing it between sheets 
of filter paper. (1) What is this substance? (2) Write two 
equations for the reactions which takes place and cause 
the formation of the precipitated substance. (3) What 
caused the second reaction to take place? (4) Why is not a 
solution of ammonium acid carbonate used to produce this 
reaction with a solution of sodium chloride? 

(h) Treat a small portion of the precipitate in a test tube 
with 3 c.c. of water and add three drops of phenolphthalein. 
(1) What was the result? 

Put into a crucible the remainder of the precipitate, or 
enough of it to one-quarter fill the crucible, and heat vigor- 
ously. While the substance is being heated, hold a drop of 



452 LABORATORY EXERCISES 

lime water on a glass rod just above the crucible. (2) What 
appeared in the lime water? (3) What was being given off 
from the heated substance? 

When bubbles no longer appear in the fused substance, 
allow the crucible to cool and add water, a cubic centimeter 
at a time, until the substance dissolves. (4) How much water 
was required? (5) How does the solubility compare with that 
of the unheated substance? 

Add three drops of phenolphthalein to the solution. (6) 
What was the result? (7) Name three facts Avhich show that 
reaction took place upon heating the substance in the crucible. 
(8) Write an equation for the reaction. (9) What industrial 
process is illustrated by this experiment as a whole? 

Experiment 118. Flame Tests. 

Reference : 615. 

Discussion : When gases are heated to a sufficiently high 
temperature they emit colored light. The Bunsen flame is hot 
enough to volatilize the alkaline metals and alkaline earths 
and some of their compounds and thus to impart characteristic 
colors to the flame. All compounds of a given metal produce 
the same color, so that the appearance of the flame serves as 
a delicate test for that metal. Since the chlorides are the 
most volatile compounds, the material to be tested is usually 
treated with hydrochloric acid before it is put in the flame. 
The flame may be observed directly, or observed through a 
prism, which spreads out the light into separate colored bands. 

Sodium compounds are so commonly present that the bright 
yellow color they produce in the flame often masks the delicate 
shades caused by potassium. Since sodium light, however, does 
not pass through blue cobalt glass, the flame may be viewed 
through a piece of this glass, and the color imparted by potas- 
sium may be detected, if it is present. 

Colored fires may be produced by mixing with some easily 
combustible matter compounds which will give the desired color 
to the flame. 

Materials: Chlorides of sodium, potassium, calcium, stron- 
tium, and barium, calcium sulphate, powdered shellac, stron- 
tium nitrate, potassium chlorate, barium nitrate. 



454 LABORATORY EXERCISES 

Directions : (a) Clean a platinum wire by alternately dip- 
ping it into concentrated hydrochloric acid in a test tube and 
holding it in the Bimsen flame, until no color is imparted to 
the flame by the wire. 

Place a few crystals each of the chlorides of sodium, potas- 
sium, calcium, strontium, and barium on watch glasses or glass 
plates and moisten them with a few drops of concentrated 
hydrochloric acid. 

Touch the clean wire to the sodium chloride and then hold 
it in the flame. (1) Kecord the color of the flame after this 
and each following trial. Clean the wire as before and repeat 
the experiment using potassium chloride instead of sodium 
chloride. Try the action again while viewing the flame through 
a piece of cobalt glass. 

Wash your hands, draw the clean wire between your thumb 
and finger and then hold it in the flame. (2) What do you 
conclude as to the delicacy of the flame test for sodium com- 
pounds? 

Touch the clean wire to both the sodium chloride and the 
potassium chloride and hold it in the flame. (3) Is the color 
caused by potassium observable? Repeat the experiment view- 
ing the flame through cobalt glass. (4) How may potassium 
compounds be detected in the presence of sodium compounds? 

Determine the colors imparted to the flame by the chlorides 
of calcium, strontium, and barium, respectively. 

Moisten a little calcium sulphate with water and try the 
flame coloration produced by it. Repeat the trial with the 
calcium sulphate moistened with a little concentrated hydro- 
chloric acid. (5) What difference in the flame did you observe? 
(6) What is the cause of this difference? (7) Why are chlorides 
to be preferred in making flame tests? 

(h) Put into a small beaker half a spatulaful each of 
powdered shellac, strontium nitrate, and potassium chlorate. 
Pour the material back and forth frOm this into a second beaker 
several times until the substances are thoroughly mixed. Heap 
the mixture in an iron pan or on a brick in the hood and 
holding the burner at arms length let the flame play upon 
the mixture. (1) What happened? (2) What was the purpose 



456 LABORATORY EXERCISES 

of the shellac? (3) Of the potassium chlorate? (4) What 
caused the color? (5) Why was strontium nitrate used instead 
of the chloride? 

Repeat the experiment using barium nitrate instead of 
strontium nitrate. (6) What occurred? 

Experiment 119. Preparation of Potassium Nitrate. 

Reference: 619, 592. 

Discussion : The source of the nitrates of commerce is 
Chili saltpeter, NaNOg. Although potassium nitrate is a soluble 
salt it can be prepared from sodium nitrate by a double decom- 
position which takes place according to the reaction repre- 
sented by the following equation : 

NaN03 + KCl = KNO3 + NaCl 

The compound can be prepared in this way because potassium 
nitrate is but slightly soluble in water at low temperatures, 
whereas the other salts involved are much more soluble. In 
making potassium nitrate in this way equivalent quantities of 
sodium nitrate and potassium chloride are dissolved in such 
an amount of hot water that the sodium chloride formed it\ 
the reaction will remain in solution when the latter is cooled. 
The marked change in solubility of potassium nitrate with 
decrease in temperature and the relatively great solubility of 
sodium chloride in cold water assist materially in bringing 
about the metathesis. 

Materials: Sodium nitrate, potassium chloride. 

Directions: (a) To 22 grams of sodium nitrate and 19 
grams of potassium chloride contained in a small beaker add 
50 c.c. of water. Heat the mixture until the solution begins 
to boil. Filter, cover the solution with a piece of filter paper, 
and set it in your desk until the next exercise. (1) Describe 
the appearance of the crystals that separate. Filter off the 
crystals, and press them against the paper in the funnel to 
remove most of the mother liquor. When no more liquid drops 
from the funnel remove the beaker and set aside the filtrate 
and treat it as desci'ibed in (h) below. Add 5 c.c. of cold 



458 LABORATORY EXERCISES 

water to the crystals of potassium nitrate in the funnel and 
let them drain. Press the crystals between dry filter papers, 
and when they are dry, weigh them. (2) Record the weight. 

(3) Calculate the weight of potassium nitrate equivalent 
to the sodium nitrate used. (4) What percentage of the 
theoretical amount was obtained in the experiment? (5) Why 
is the yield not the theoretical? (6) How could you calculate 
the largest percentage of potassium nitrate that could be ob- 
tained in the preparation as carried out? (7) If the solution 
were cooled to 0° would the yield of potassium nitrate be mu<*-h 
greater? (8) Would the salt obtained be as pure? (9) Why? 

(h) Heat the filtrate from the potassium nitrate crystals 
to boiling and evaporate it to about two-thirds its original 
volume. (1) What is the compound that separates during the 
boiling? Filter the hot solution rapidly. Save the filtrate. 
Press the crystals between filter papers. Examine them with 
a magnifying glass. (2) Do they resemble the crystals of 
potassium nitrate in appearance? (3) Upon what fact does 
the separation of potassium nitrate from the other three sub- 
stances in solution depend? Examine the crystals that are 
deposited from the filtrate when it cools. (4) What are they? 



CALCIUM, STRONTIUM, AND BARIUM 

Experiment 120. Hard Water. 

Keference: 628-631. 

Discussion: Water which contains soluble compounds of 
calcium, magnesium or iron is called ''hard water," because 
the addition of soap produces a precipitate by reaction with 
the compound in the water, instead of forming suds, as with 
pure water. The precipitate is usually a stearate formed by 
double decomposition between the compound in solution and 
sodium stearate, NaCigHg^Oa, which is the chief constituent of 
ordinary soap. 

The most common compound found in hard water is calcium 
acid carbonate, formed from limestone by the action of water 
containing carbonic acid. As was shown in Experiment 41 h, 
on the formation of calcium acid carbonate, this compound is 
decomposed by boiling the solution, when the calcium is pre- 
cipitated as carbonate. The acid carbonates of magnesium and 
iron act similarly. Water containing any of these compounds 
is therefore called ** temporary" hard water, in distinction from 
water containing other soluble salts of these metals, which is 
known as ''permanent" hard water. 

Temporary hard water may be made **soft" by removing 
the excess of carbonic acid, which preserves the acid carbonate 
from decomposition, by boiling the water, as stated above, or 
by neutralizing the acid with a base, such as ammonium hy- 
droxide or calcium hydroxide. Both kinds of hard water may 
be "softened" by the addition of a reagent which will form 
an insoluble compound by reaction with the salt which is caus- 
ing the hardness. Borax is often used for this purpose. 

Materials: Dilute soap solution, marble, lime water, solu- 
tions of calcium chloride, magnesium chloride, ferric chloride, 
calcium sulphate, sodium carbonate, and borax. 

460 



462 LABORATORY EXERCISES 

Directions: (a) Add 5 c.c. of soap solution to a test tube 
half full of distilled water and shake it vigorously. (1) Describe 
the appearance. (2) What is the soap commonly said to have 
produced? 

(h) Prepare 60 c.c. of calcium acid carbonate solution as 
directed in Experiment 41 h. (1) Write equations for the 
reactions by which it was formed. (2) How is calcium acid 
carbonate formed in nature? 

To 20 c.c. of the solution add 5 c.c. of soap solution and 
shake it. (4) How does the appearance of the liquid differ 
from that obtained in {a) after 10 seconds? (5) What name 
is given to any water which acts this way with soap? (6) 
Write an equation for the reaction which took place. 

Add successive 5 c.c. portions of soap solution, shaking 
after each addition, until suds are formed which persist after 
10 seconds. (7) How many cubic centimeters of soap did you 
add in all? (8) How may the degree of hardness of water be 
determined? 

Boil vigorously a fresh 20 c.c. portion of the solution of 
calcium acid carbonate for 5 minutes. (9) What was formed? 
(10) Write an equation for the reaction. Filter and shake the 
filtrate with 5 c.c. of soap solution. (11) What was the result? 
(12) Why did it differ from that obtained before the solution 
was boiled? (13) What is water called which behaves thus? 

To the remainder of the original solution add ammonium 
hydroxide until an excess is present, as indicated by the odor. 
(14) What was precipitated? (15) Write an equation for the 
reaction. Filter, add 5 c.c. of soap solution and shake. (16) 
Was the water softened by the addition of the ammonium 
hydroxide? (17) What is the evidence? (18) What other 
compounds act as the calcium acid carbonate has in these 
experiments? (19) Write an equation to' show why lime is 
used in the industries to soften hard water. 

(c) To 20 c.c. each of solutions of calcium sulphate, calcium 
chloride, magnesium chloride, ferric chloride, and sodium 
chloride add 5 c.c. of soap solution and shake them. (1) State 
for each case whether or not suds were formed which remain 
permanent after 10 seconds. (2) Soluble salts of what elements 



464 LABORATORY EXERCISES 

cause water to be hard? (3) How do these compounds make 
the water hard? 

Boil 20 c.c. of a solution of calcium chloride. (4) Did a 
precipitate appear? Test the solution for hardness. (5) What 
did you do and observe? (6) Does boiling soften water which 
is hard on account of the presence of calcium chloride? (7) 
What is such hard water called? 

To 20 c.c. of a Golution of calcium sulphate add 10 c.c. 
of a solution of sodium carbonate. (8) Write an equation for 
the reaction taking place. Filter and test the filtrate for hard- 
ness. (9) What was the result? 

Repeat the experiment using first borax solution instead of 
sodium carbonate and then again substituting ammonium hy- 
droxide. (10) State for each case what happened upon the 
addtion of the reagent and (11) whether or not the water was 
softened by it. (12) How can permanent hard water be 
softened? 

Experiment 121. Plaster of Paris. 

Reference : 636. 

Discussion : Plaster of Paris is the hemihydrate of calcium 
sulphate, (CaS04)2,H20, and is made by heating gypsum, 
CaS04,2H20, to about 145°. When the proper amount of 
water is added to plaster of Paris, the reverse action takes 
place and the dihydrate which is produced forms a rigid mass 
of interlocking crystals. The material expands slightly during 
this change and will thus completely fill all the small indenta- 
tions in a mold in Avhich it is put before it ^^sets." This 
expansion makes the material suitable for making casts. 

Materials: Plaster of Paris, block of wood, vaseline, coin. 

Directions: Mix five spatulafuls of plaster of Paris on a 
block of wood with enough water to make a soft dough. Press 
into the plaster a coin which has been covered with a thin 
coating of vaseline, and after 15 or 20 minutes, remove the 
coin and examine the cast. (1) How has the physical condi- 
tion of the plaster changed? (2) What is this change called? 

(3) Are the indentations in the coin well marked in the cast? 

(4) To what property of the plaster is this due? (5) Write 



466 LABORATORY EXERCISES 

an eqaation for the reaction taking place when plaster of Paris 
is prepared ; (6) when it sets. 

Experiment 122. Comparison of Compounds of Calcium, 
Strontium and Barium. 

Reference : 589, 593, 642, 644, 651. 

Discussion: The following study of compounds of the 
'* alkaline earths" is made from the point of view of solubility 
and affords, in addition to a knowledge of the appearance of 
certain specific substances, a number of examples of the relation 
of ionization to precipitation. The underlying principle is that 
a reaction of double decomposition in solution tends to run 
to the formation of the least ionized substance. A small con- 
centration of ions may be due to one of two causes, either to 
a small degree of dissociation, as in the case of ammonium 
hydroxide and of acetic acid, or to the insolubility of the 
compound, as in the case of barium sulphate. 

If the equilibria involved are expressed in the following 
way, these relations become apparent : 



AB 


^A-+ 


-\-B- 




+ 


+ 


CD 


^D- 


■+C+ 




Ti 


U 



AD CB 

When, for example, AB represents calcium chloride and CD 
sodium hydroxide, both highly ionized substanese, AD (calcium 
hydroxide) is precipitated, as it is the least ionized combination, 
due to its insolubility; but when CD represents ammonium 
hydroxide, no AD is precipitated, because the slight dissociation 
of CD in this case makes this compound the least ionized com- 
bination. 

Materials: Solutions of calcium, strontium, and barium 
chlorides, ammonium carbonate, ammonium oxalate, acetic acid, 
and sodium chromate, calcium sulphate, strontium sulphate. 

Directions: (1) Prepare a table for recording the results 
of these experiments, indicating the divisions into separate ex- 
periments by drawing horizontal lines across the table below 



468 LABORATORY EXERCISES 

the reagents italicized in the following list, which constitutes 
the first column of the table: Reagents, sodium hydroxide, am- 
monium hydroxide, ammonium carbonate, ammonium oxalate, 
hydrochloric acid, acetic acid, sodium chromate, hydrochloric 
acid, acetic acid, sulphuric acid, hydrochloric acid, calcium sul- 
phate, strontimn sulphate. Head the second, third, and fourth 
columns calcium, strontium, and barium respectively and record 
in their proper places in the table your experimental results. 
When a precipitate appears, write '^P" and name the color; 
when an insoluble substance reacts with a reagent write **Il"; 
and when nothing is observable make a dash. 

(a) Add 5 c.c. of sodium hydroxide solution to each of three 
test tubes containing 5 c.c. of solutions of calcium chloride, 
strontium chloride, and barium chloride respectively. Repeat 
the experiments substituting ammonium hydroxide for the 
sodium hydroxide. (2) Make a diagram, of the kind given in 
the discussion, to show the equilibria involved in the addition 
of sodium hydroxide to calcium chloride and one involved in 
the addition of ammonium hydroxide to calcium chloride. 
Underline in each, the least ionized substance and state the 
causes of the small concentration of ions. (3) Why does sodium 
hydroxide produce a different result from ammonium hy- 
droxide? 

{h) To 5 c.c. each of solutions of calcium, strontium, and 
barium chlorides add a few drops of ammonium hydroxide, 
to insure the alkalinity of the solution, and then 5 c.c. of 
ammonimun carbonate solution. (2) Construct a diagram for 
the equilibria in the case of calcium chloride and ammonium 
carbonate. (3) Why does the addition of ammonium cai'bonate 
give a different result from the addition of ammonium hy- 
droxide? 

(c) To 5 c.c. each of solutions of calcium, strontium, and 
barium chlorides, made alkaline with a few drops of ammonium 
hydroxide, add 5 c.c. of ammonium oxalate. Pour half the 
contents of each test tube, including a share of the precipitates, 
into three other test tubes. Add an excess of hydrochloric 
acid to each of these test tubes and an excess of acetic acid 
to each of the original tubes. (2) Make equilibria diagrams 



470 LABORATORY EXERCISES 

for the eases of calcium oxalate and hydrochloric acid, ealeium 
oxalate and acetic acid and barium oxalate and acetic acid. 

(3) Why does hydrochloric acid differ from acetic acid in its 
action with calcium oxalate? (4) Why does acetic acid differ 
in the action with calcium oxalate and barium oxalate? (5) 
How may a soluble calcium salt be distinguished from a barium 
or strontium salt? 

{d) To slightly alkaline solutions of the three chlorides, as 
in (c), add 5 c.c. portions of sodium chromate solution. Divide 
the contents of the test tubes and try the action of hydrochloric 
acid and of acetic acid on each precipitate, as before in (c). 
(2) Make equilibria diagrams for the cases of calcium chromate 
and acetic acid and of barium chromate and acetic acid. (3) 
Why does acetic acid act differently with these two chromates ? 

(4) How may a soluble barium salt be distinguished from a 
soluble calcium or strontium salt? 

(e) To 5 c.c. each of solutions of calcium, strontium, and 
barium chlorides add 1 c.c. of dilute sulphuric acid. De- 
termine whether or not the sulphates formed will react Avith 
hydrochloric acid. 

To a test tube half full of water add about 1 gram of solid 
calcium sulphate, shake the mixture vigorously and then filter 
it. To half of the filtrate add 5 c.c. of strontium chloride 
solution and to the other half 5 c.c. of barium chloride solution. 
(2) Explain why the reaction took place in each case. 

Shake a mixture of 5 c.c. of water and 1 gram of stron- 
tium sulphate and filter. To the filtrate add 5 c.c. of bai'ium 
chloride solution. (3) What are the relative solubilities of the 
three sulphates studied? (4) How do these experiments es- 
tablish this order? (5) Where in the periodic classification 
do the elements studied in this experiment occur? (6) What 
is this group called? 

(/) Determine as in Experiment 118 the colors imparted to 
the flame by the chlorides of these elements. (1) Record the 
results. 



MAGNESIUM, ZINC, AND MERCURY 

Experiment 123. Magnesium Nitride. 

Reference : 318. 

Discussion : When magnesium burns in the air, the chief 
product is magnesium oxide, which is formed with the evolu- 
tion of much energy, as light and heat. Under these conditions 
the usually inert nitrogen of the air becomes sufficiently active 
to unite with some of the magnesium. The resulting magnesium 
nitride is decomposed by water. 

Materials : Magnesium powder, copper wire, litmus paper. 

Directions: Put half a spatulaful of powdered magnesium 
into a crucible supported on a clay triangle. Heat the crucible 
vigorously and then ignite the magnesium by directing the 
flame of the burner into the crucible. Continue to heat the 
crucible and stir the burning magnesium with a piece of copper 
wire until no further ignition takes place. When the crucible 
has cooled, allow a few drops of water to fall upon the contents 
and at the same time lay a piece of pink litmus paper across 
the top of the crucible. (1) What was the action on the litmus 
paper? (2) To what substance is the action due? (3) Write 
equations for the reactions which took place when the magne- 
sium burned. (4) Which was the chief reaction? (5) How did 
this facilitate the other reaction? (6) Write equations for the 
two reactions which took place upon the addition of water. 

Experiment 124. Magnesium Chloride. 

Reference: 656. 

Discussion : The great solubility of magnesium chloride, 
which accounts for its deliquescence, and the extensive hydro- 
lysis which takes place upon heating a solution of this salt, 
are shown in the following experiment. 

Materials : Magnesium chloride, litmus paper. 

Directions: Put half a spatulaful of magnesium chloride 

472 



474 LABORATORY EXERCISES 

into a 6-inch test tube and add 2 c.c. of water. (1) What is 
shown as to the solubility of the substance? 

Boil away all of the water in the test tube and test the 
vapor with litmus paper as it issues from the tube. (2) What 
was the result? (3) What property of magnesium chloride is 
thus shown? Heat the residue in the tube strongly for two 
or three minutes and when it is cold, add about 3 c.c. of water 
to the residue. (4) Did it dissolve completely? (5) How do 
you account for this? 

Experiment 125. Zinc Hydroxide. 

Reference : 595, 663. 

Discussion : Zinc hydroxide forms both hydroxide ions and 
hydrogen ions in solution, that is, it may act either as a base 
or as an acid. These facts are shown in the following ex- 
pression : 

2H++Zn02""^Zn(OH)2;=^Zn"'-^+20H- 

Zn(OH)2(solid) 

If an acid, as hydrochloric acid, is added to zinc hydroxide, 
the hydrogen ions of the acid unite with the hydroxyl ions 
from the zinc hydroxide and cause the reaction to proceed to 
the right in the expression just given. This action causes solid 
zinc hydroxide to dissolve to replace that originally in solution, 
and a solution of zinc chloride is the final product. If, on 
the other hand, a base, as sodium hydroxide, is added to the 
zinc hydroxide, the hydroxyl ions of the base unite with the 
hydrogen ions from the zinc hydroxide, the reaction proceeds 
to the left, the solid zinc hydroxide dissolves as before, but 
the final product is sodium zincate. 

The solubility of zinc hydroxide in ammonium hydroxide is 
due to a third cause. Ammonia forms by direct union with the 
zinc ion a complex ion of the composition Zn(NH3)4 + + . The 
formation of these ions removes zinc ions from solution. This 
causes the reaction again to proceed to the right as indicated 
in the expression above, with the consequent solution of the 
solid zinc hydroxide. 



476 LABORATORY EXERCISES 

In summary, any reagent which will remove one of the ions 
formed by zinc hydroxide will cause the solid to disappear. 
If the reagent is an acid, the hydroxyl ions are removed, if it 
is a base, the hydrogen ions, and if it is ammonia, the zinc ions. 

Materials: Solution of zinc sulphate. 

Directions: (a) To 10 c.c. of zinc sulphate solution in a 
test tube add sodium hydroxide solution drop by drop until 
the liquid contains a considerable precipitate. (1) What is 
the precipitate? (2) Write an equation for its formation. 

Shake the liquid and pour half of it into another test tube. 
To one portion add hydrochloric acid and to the other sodium 
hydroxide, until a distinct change takes place. (3) What 
occurred in each case? (4) Write equations for the reactions. 
(5) As what kind of a compound did the zinc hydroxide act 
when treated with the acid? (6) When treated with the base? 

(7) Write an expression to show how zinc hydroxide ionizes. 

(8) Explain with reference to this expression why the reaction 
with the acid and with the base took place. 

{h) To 10 c.c. of zinc sulphate solution add a few" drops 
of ammonium hydroxide. (1) What is formed? (2) Write an 
equation for the reaction. 

Add to the test tube ammonium hydroxide until a distinct 
change takes place. (3) What occurred? (4) What compound 
is present in a solution of ammonium hydroxide in addition 
to the ammonium hydroxide? (5) Write an equation for the 
reaction which occurred. (6) Explain with reference to the 
expression for the ionization of zinc hydroxide why the reaction 
took place. (7) Is the solution of the zinc hydroxide in this 
case due to the same cause as the solution of magnesium hy- 
droxide on the addition of ammonium salts, as observed in 
Experiment 106 hi (8) Wherein lies the resemblance or dif- 
ference? (9) Upon what general principle does the reaction 
of zinc hydroxide with acids, bases, and ammonia depend in 
all three cases? 

Experiment 126. Zinc Chloride. 

Reference : 664. 

Discussion : It was shown in Experiment 125 that the basic 



478 LABORATORY EXERCISES 

properties of zinc hydroxides are very weak and it therefore 
follows that soluble zinc salts will become hydrolyzed in solu- 
tion. If the conditions are such that the acid formed by this 
hydrolysis can pass off, the hydrolysis will proceed, leaving 
a basic salt. 

Materials: Zinc chloride. 

Directions: To 5 c.c. of water in a small test tube add 
several crystals of zinc chloride. Test the solution with litmus 
paper. (1) What was the result? (2) What does this prove 
concerning the action of zinc chloride with water? (3) Did 
any precipitate appear? (4) What does this fact show as to 
the extent of the action? 

Boil the solution, and when most of the water has 
evaporated, test the steam at the mouth of the tube with blue 
litmus paper. (5) What was the result? (6) What caused 
the action? (7) What was its source? 

Continue the evaporation to dryness and heat the residue 
in the test tube to as high a temperature as the glass will 
permit for several minutes. When the tube has cooled add 
5 c.c. of water. (8) Is the substance in the tube soluble? 
Add 5 c.c. of nitric acid; heat to boiling; filter, if necessary; 
and test the solution for a chloride? (9) What was the result 
of the test? (10) What was formed by heating magnesium 
chloride ? 

Experiment 127. Zinc Sulphide. A Test for Zinc. 

Reference: 666. 

Discussion : Zinc sulphide is characterized by being the 
only common white sulphide and by being insoluble in acetic 
acid but soluble in more active acids, such as hydrochloric or 
sulphuric acids. These characteristics may be made use of in 
recognizing compounds of zinc. 

Materials: Solutions of zinc sulphate, of ammonium sul- 
phide, of acetic acid, and hydrogen sulphide. 

Directions: (a) To 10 c.c. of zinc sulphate solution in a 
test tube add 10 c.c. of ammonium sulphide. (1) What is the 
appearance and the solubility of zinc sulphide? (2) Write an 
equation for the reaction of its formation. 



480 LABORATORY EXERCISES 

Filter out the precipitate, wash it on the paper with 5 c.c. 
of water, transfer the funnel to a clean test tube, poke a hole 
in the tip of the filter paper, and wash the precipitate into the 
test tube with 15 c.c. of water. Pour a third of the suspended 
zinc sulphide into each of two other test tubes. To one add 
10 c.c. of hydrochloric acid, to the other 10 c.c. of sulphuric 
acid and to the remainder of the zinc sulphide 10 c.c. of acetic 
acid. (3) What happened in each case? (4) Write expressions 
similar to those used in Experiment 122 to shoAv the equilibria 
involved between zinc sulphide and acetic acid. (5) Explain 
with reference to these expressions the difference in action of 
the acids with zinc sulphide. 

{h) Bubble hydrogen sulphide into 10 c.c. of zinc sulphate 
solution in a test tube until the solution smells strongly of 
hydrogen sulphide. (1) What insoluble substance was formed? 
(2) Write an equation for its formation. 

Filter out the precipitate and add a few drops of ammonium 
hydroxide to the filtrate. (3) What appeared? (4) Was all 
the zinc precipitated by the hydrogen sulphide? (5) What 
is the evidence? (6) Why was the reaction incomplete? (7) 
Why did it take place at all? (8) What reagent should be 
used to precipitate all the zinc as sulphide? 

(c) Prepare 10 c.c. of sodium zincate solution as in Experi- 
ment 125. (1) What did you do? To this add 5 c.c. of am- 
monium sulphide. (2) What was formed? (3) Write an equa- 
tion for the reaction. 

(d) (1) State what you would do in testing for zinc, and 
what would happen if it were present. 

Experiment 128. Compounds of Mercury. 

Reference : 672, 674, 676, 677, 679, 681. 

Discussion : The important reactions of the compounds of 
mercury depend upon the following facts: (1) Two sets of 
compounds exist, in which mercury appears Avith the valences 
of one and of two respectively. Some reagents may reduce 
mercuric compounds with the valence of two, to mercurous 
compounds with the valence of one, or may even reduce them 
to metallic mercury, as in the case of the chlorides. Some 



482 LABORATORY EXERCISES 

mercurons compounds, on the other hand, undergo spontaneous 
oxidation to mercuric compounds, with the necessary separation 
of metallic mercury, as in the case of the sulphide. 

(2) Hydroxides of mercury are unstable and consequently 
an oxide is precipitated where a hydroxide would be expected 
as a product of double decomposition. Moreover, the oxides 
of mercury do not react with an excess of a base to form 
salts, as does zinc hydroxide. 

(3) Mercuric salts react with ammonia to form ammono 
basic compounds. This is a type of reaction not heretofore 
studied. It w^ill be recalled that in the hydrolysis of a salt 
the hydroxide radical of water unites with the metal of the 
salt and the remaining hydrogen unites with the negative 
radical of the salt ; but that in the case of many salts, this 
action is incomplete, because the substitution of only part of the 
negative radicals by hydroxide takes place. Thus, mercuric 
chloride can be hydrolyzed in accordance with the expression: 
HgCl^ + H^O = Hg(OH)Cl + HCl. Just as water thus reacts 
as a compound of hydrogen with the hydroxide radical, so 
ammonia may react as a compound of hydrogen with a radical 
of the composition NHg ; and in the same way that many salts 
react with water in hydrolysis, so mercuric salts react with 
ammonia in ammonolysis. The resulting compound is analogous 
in composition to the basic salt, with the substitution of the 
radical NIL, for the radical OH. These compounds must not 
be confused with the complex salts formed by the direct union 
of ammonia to salts of zinc, copper, and other metals. 

Materials: Solutions of mercuric chloride, of mercurous 
nitrate, and of stannous chloride, and hydrogen sulphide. 

Directions: (a) Dilute 5 c.c. of mercuric chloride solution 
with an equal volume of water in a test tube and add to it 
ten drops of stannous chloride solution. (1) State the chemical 
name, the commercial name, the color, and the solubility of 
the mercury compound formed. (2) Write an equation for 
the reaction. (3) What kind of a process has the mercuric 
chloride undergone? (4) What was the agent? 

To the contents of the test tube add more stannous chloride 
until another distinct change takes place. (5) Describe the 



484 LABORATORY EXERCISES 

change. (6) Write an equation for the reaction. (7) What 
kind of a change has the mercury compound undergone in 
this case? 

(h) Bubble hydrogen sulphide into 5 c.c. of mercuric 
chloride solution. (1) Name and describe the appearance of 
the mercury compound formed. (2) Write an equation for 
its formation. Add to it 5 c.c. of concentrated nitric acid 
and heat it to boiling. (3) What do you conclude as to the 
solubility of mercuric sulphide? (4) Upon what do you base 
your conclusion? 

Bubble hydrogen sulphide into 5 c.c. of mercurous nitrate 
solution. (5) What appeared? This is a mixture of mercuric 
sulphide and mercury. (6) What would have been formed 
by double decomposition? (7) What does this experiment show 
concerning the stability of mercurous sulphide? (8) What 
process has the mercury undergone in changing from mer- 
curous nitrate to mercuric sulphide? (9) What was reduced? 

(c) To 5 c.c. of mercuric chloride solution add ten drops 
of sodium hydroxide solution. (1) Name and describe the 
precipitate. (2) Write an equation for the reaction. Add 
sodium hydroxide in excess. (3) What happened? 

Repeat the experiment using mercurous nitrate instead of 
mercuric chloride. (4) Name and describe the precipitate. 
(5) Write an equation for its formation. (6) Did it react with 
an excess of sodium hydroxide? (7) Does mercury show 
properties of an acid-forming element? (8) Upon what do 
you base your answer? (9) What other element in this group 
of the periodic classification differs from mercury in this 
respect^ 

{d) To 5 c.c. of mercuric chloride solution add 5 c.c. of 
ammonium hydroxide solution. (1) What is the name, color, 
and solubility of the mercury compound formed? (2) With 
what substance in the ammonium hydroxide solution did the 
mercuric chloride react? (3) To what type of reaction 
previously studied is this type of reaction similar? (4) Explain 
with the use of two equations where the similarity exists. 

Prepare some mercurous chloride by adding 5 c.c. of dilute 
hydrochloric acid to 5 c.c. of mercurous nitrate, filter, and add 



486 LABORATORY EXERCISES 

a few drops of ammonmm hydroxide solution to the precipitate 
on the paper. (5) What happened? The substance on the 
paper is a mixture containing ammono-basic mercuric chloride 
and mercury. (6) To what is the color due? (7) In what 
other reaction studied in this experiment did a similar spon- 
taneous oxidation and reduction take place? (8) Write an 
equation for the reaction. 

(e) (1) What characteristic action of a mercuric salt would 
you choose as suitable for the basis of a test? (2) Describe 
the procedure in making the test and the results if a mercuric 
salt were present. (3) What reaction would be suitable as a 
test for a mercurous salt? (4) Describe the procedure and the 
result of making a test if a mercurous salt were present. 



ALUMINIUM 

Experiment 129. Aluminium Hydroxide. 

Reference : 685. 

Discussion: In some of the preceding experiments, the im- 
portance has been emphasized of studying the behavior of the 
hydroxides of the elements with acids, with sodium hydroxide, 
and with ammonium hydroxide, as showing the base- or acid- 
forming properties of the elements, their ability to form com- 
plex or substituted compounds with ammonia, and the means 
of distinguishing compounds of one element from those of 
others. The following experiments on aluminium hydroxide 
are designed to show the chemical properties of this element 
in these same respects. 

Materials: Solution of aluminium sulphate. 

Directions: To 10 c.c. of aluminium sulphate solution in 
a test tube add 10 drops of sodium hydroxide solution. (1) 
Describe the appearance of the precipitate. (2) Write the 
equation for its formation. 

Pour one-third of the contents of the test tube into each 
of two other test tubes, to one of which add an excess of 
hydrochloi'ic acid. (3) What occurred? (4) Write an equa- 
tion for the reaction. (5) As what kind of a compound does 
aluminium hydroxide act? 

To the second of the three test tubes add an excess of 
sodium hydroxide. (6) What occurred? (The salt formed 
is that derived from a dehydi'ated aluminium hydroxide of the 
composition Al(OH)., — H.O = HAIO,.) (7) Write an equa- 
tion for the reaction. (8) As what kind of a compound does 
aluminium hydroxide act in this case? (9) Write an expression 
of the kind used in Experiment 125 to show the equilibria 
involved in a suspension of aluminium hydroxide. (10) Ex- 
plain with reference to this expression the actions of hydro- 

488 



490 LABORATORY EXERCISES 

chloric acid and of sodium hydroxide with aluminium hy- 
droxide. 

To the third test tube add an excess of ammonium hy- 
droxide. (11) What occurred? (12) Does aluminium hy- 
droxide form either complex compounds or substituted com- 
pounds with ammonia? 

(13) By the use of what reagent or reagents can a soluble 
aluminium salt be distinguished from a soluble calcium salt? 
State what happens with the reagent in each case? (14) Like- 
wise how can aluminium salts be distinguished from magne- 
sium salts? (15) From zinc salts? (16) Prom mercury salts? 

Experiment 130. Method for Cleaning Silver. 

Reference: 567. 

Discussion : It was shown in Experiment 97 that when a 
metal is put into a solution with which it will react with the 
evolution of hydrogen, and when also this metal is in contact 
with another metal beloAv it in tile electromotive series, the 
hydrogen appears on the surface of the less active metal and 
the more active metal goes into solution as a compound. Thus, 
when aluminium in contact with silver is treated with sodium 
hydroxide, hydi'ogen appears upon the silver and sodium 
aluminate is formed. If the silver is coated with the sulphide, 
the usual tarnish on silver ware, the hydrogen will reduce this 
sulphide. Use is made of these reactions in cleaning tarnished 
silver. 

Materials: Solution of sodium carbonate; strips of alumin- 
ium, and of tarnished silver to be returned after using. 

Directions: Heat a strip of aluminium, about one by three 
centimeters in size, Avith enough sodium carbonate solution to 
cover it in a test tube. (1) What is the evidence of chemical 
action? (2) How do you account for this action of sodium 
carbonate with aluminium? 

Slip into the test tube a strip of tarnished silver so that it 
will touch the piece of aluminium and allow only enough solu- 
tion to remain in the tube so that but half the strip of silver 
is submerged. Heat the solution to boiling, remove the test 
tube from the flame and observe the piece of silver. (3) What 



492 LABORATORY EXERCISES 

is the gas set free on its surface? (4) What relation has 
this phenomenon with the relative positions of aluminium 
and silver in the electromotive series? Take the silver from 
the tube and compare the appearance of the two ends of the 
strip. (5) How has the part which was in the solution changed? 
(6) Of what does silver tarnish consist? (7) Write an equa- 
tion for the reaction which took place when it disappeared. 
(8) How can this action be made of practical household use? 



TIN AND LEAD 

Experiment 131. Hydroxides of Tin. 

Reference: 698-700. 

Discussion : As in the case of other elements recently 
studied, a knowledge of the reaction of its hydroxides is im- 
portant in summarizing the properties of tin in respect to its 
electropositive or electronegative characteristics. The position 
of tin in the fourth group of the periodic classification would 
lead to the prediction of distinctly acidic properties for stannic 
hydroxide, yet the fact that tin occurs so far doAvn as the sixth 
member of the group would indicate the possibility of some 
basic characteristics. 

Stannic hydroxide can be prepared from the chloride by 
precipitation by a base or it may be formed by oxidation of the 
metal in presence of water, when a hydrated oxide results. The 
former compound is called a-stannic acid, tlie latter /?-, or meta 
stannic acid. 

Materials: Solution of stannic chloride, concentrated nitric 
acid. 

Directions: (a) Dilute 5 c.c. of stannic cliloride solution 
witli an equal volume of water in a test tube and add sodium 
hydroxide solution, a few drops at a time, until a precipitate 
is formed. (1) Describe the appearance of this precipitate. 
(2) What is it called? (3) Write an equation for its formation. 

Pour half of the suspended precipitate into another test 
tube and add an excess of hydrochloric acid to one tube and 
an excess of sodium hydroxide to the other. (4) What hap- 
pened in each case? (5) Write equations for the reactions. 
(6) How might these reactions have been predicted from the 
position of tin in the periodic classification? 

{h) Add 5 c.c. of concentrated nitric acid to a small piece 
of tin in a test tube in the hood. Warm the tube, if necessary, 
to start the reaction and then stand it in a gas bottle to 

494 



496 LABORATORY EXERCISES 

catch any liquid which may overflow. When the action has 
ceased, dilute the solution with three times its volume of water 
and filter. Wash the precipitate with three 5 c.c. portions of 
water and then transfer it to a clean test tube by puncturing 
the filter paper. (1) What gas was evolved in the reaction? 
(2) What do you conclude as to the role of nitric acid? (3) 
What was oxidized? (4) What is the composition of the 
precipitate? (5) What is it called? 

Pour half of the suspended precipitate into a clean test 
tube and treat one portion with hydrochloric acid and the other 
with sodium hydroxide. (6) What happened in each case? 
(7) The oxide of what other element in the same group of the 
periodic classification as tin behaves in the same way with 
solutions of acids and bases? 

(c) To 5 c.c. of stannous chloride solution add a solution of 
sodium hydroxide very slowly, and finally in excess. (1) What 
were the results obtained? (2) Write equations for the reac- 
tions. 

Boil the solution until a precipitate appears. Metallic tin 
and sodium stannate are formed. (3) Write an equation for 
the reaction. 

Experiment 132. Chlorides of Tin. 

Reference: 701. 

Discussion : As would be predicted from the position of 
tin in the electromotive series, non-oxidizing acids Avill react 
with this metal to form salts and evolve hydrogen. If hydro- 
chloric acid is used, mercuric chloride can be used to determine 
whether stannous chloride, or stannic chloride is produced. 
This is the same reaction which was employed in the reduction 
of mercuric chloride to mercurous chloride, studied in Experi- 
ment 128 a. 

Suitable reagents can oxidize or reduce stannous chloride 
or stannic chloride, respectively, so that conversely these com- 
pounds can be used for reduction and oxidation. Since the 
oxygen of the air will slowly oxidize stannous chloride, special 
means must be used to preserve solutions of this compound in 
the laboratory. 



498 LABORATORY EXERCISES 

Materials: Tin, solutions of mercuric chloride, of ferric 
chloride, concentrated hydrochloric acid, concentrated nitric 
acid. 

Directions: Ponr upon a small piece of tin in a test tube 
5 c.c. of concentrated hydrochloric acid and 5 c.c. of water. 
Heat the solution if necessary and keep it at such a temperature 
that vigorous action continues until the tin has disappeared. 
(1) What was the gas evolved? (2) Why might you expect 
that it would be formed? 

To 5 c.c. of mercuric chloride solution add a few drops of 
this solution. (3) What is the substance which appears? (4) 
What does this prove as to the tin compound formed by the 
original action of the hydrochloric acid? (5) Write an equa- 
tion for the reaction. 

(Ferric salts are reddish brown and ferrous salts colorless 
in dilute solutions.) Dilute 1 c.c. of ferric chloride solution 
with 3 c.c. of water and add some of your original solution 
drop by drop till a change of color occurs. (6) What was the 
change in color? (7) Write an equation for the reaction. (8) 
What kind of change has the tin compound undergone? (9) 
As w^hat kind of an agent has it acted? 

To 5 c.c. of the original solution add 10 drops of concen- 
trated nitric acid and heat to boiling. Cool the solution and 
test for a stannous salt. (10) How did you test? (11) With 
what result? (12) What has happened to the stannous 
chloride? (13) Will this reaction be brought al)out by the 
air? (14) What is used in the laboratory to preserve solutions 
of stannous chloride from change? 

Experiment 133. Sulphides of Tin. 

Reference : 478, 702. 

Discussion : The sulphides of tin are chiefly of interest in 
relation to their use in separating tin from other metals and 
in identifying tin in the course of qualitative analysis. Their 
important properties are the insohi])ility in dilute acid, and 
the solubility in ammonium polysulphide and concentrated 
hydrochloric acid. 

Materials: Solutions of stannic chloride, of stannous 



500 LABORATORY EXERCISES 

chloride, of ammonium polysulphide, concentrated hydrochloric 
acid, and hydrogen sulphide. 

Directions: (a) Bubble hydrogen sulphide into 10 c.c. of 
a solution of stannic chloride. (1) Describe the precipitate. 

(2) Write an equation for its formation. 

Filter out the precipitate and wash it with three portions 
of water of 5 c.c. each. After puncturing the tip of the filter 
paper, wash the precipitate into a clean test tube with as little 
water as possible. Transfer half of the precipitate to another 
test tube and treat it with 5 c.c. of ammonium polysulphide. 

(3) What happened? (4) Write an equation for the reaction. 
(5) The sulphides of what other elements react in a similar 
way? 

To the remainder of the precipitate add 5 c.c. of concen- 
trated hydrochloric acid and heat to boiling. (6) What hap- 
pened? (7) Referring to the properties of the sulphides of 
arsenic studied in Experiment 89, state how stannic sulphide 
may be distinguished from them. 

(b) Bubble hydrogen sulphide through 5 c.c. of stannous 
chloride solution. (1) What is the appearance of the pre- 
cipitate? (2) Write an equation for its formation. 

Filter, wash, and treat the precipitate with ammonium poly- 
sulphide as in part (a) above. (3) State what happened. (4) 
Write an equation for the reaction. 

Treat the solution with dilute hydrochloric acid. (5) What 
is the precipitate? (6) How do you recognize it? (7) What 
kind of change has the tin undergone during these reactions? 
(8) What was the agent? 

Experiment 134. Oxides of Lead. 

Reference : 706. 

Discussion : Since lead exhibits two valences in its com- 
pounds, two different oxides, lead monoxide and lead dioxide, 
would be anticipated. In addition to these, however, a third 
oxide, red lead, is known, the constitution of which is to be 
determined in a part of this experiment. 

Materials: Lead, oxides of lead, concentrated hydrochloric 
acid, starch iodide paper, bromine water, charcoal. 



502 LABORATORY EXERCISES 

Directions: {a) Place about one gram of lead in a shallow 
hole in a piece of charcoal and heat it in the oxidizing flame of 
a blow pipe until a deposit is formed about the hole. (1) What 
is the color of this deposit? Compare it with the samples of 
oxides of lead in the laboratory. (2) What is it? 

(b) To 5 c.c. of lead nitrate solution add sodium hydroxide 
solution until a considerable precipitate is formed. To this 
add bromine water drop by drop until the precipitate under- 
goes a distinct change. (1) What is the appearance of the 
new compound? Compare it with the samples of the oxides 
of lead. (2) What is it? (3) Write an equation for its forma- 
tion. (4) What kind of a change has the lead undergone? 
(5) What was the agent? 

(c) Warm half a spatulaful of lead dioxide with 5 c.c. of 
concentrated hydrochloric acid in a test tube and smell 
cautiously at the mouth of the tube. (1) What do you con- 
clude was formed in the test tube? 

Add 10 c.c. of water and filter. When about 5 c.c. of filtrate 
have been collected, add a few drops of dilute sulphuric acid. 
(2) What appeared? (3) What is the substance? (4) Write 
an equation for the reaction taking place in making the test. 
(5) What kind of a change has the lead of the lead dioxide 
undergone? (6) As what kinxl of agent has it acted? (7) Upon 
what did it thus act? (8) Write an equation for this reaction. 

(d) Warm a quarter of a spatulaful of red lead with 10 c.c. 
of dilute nitric acid in a test tube, and when no further action 
takes place, filter the liquid. (1) What is the appearance of 
the residue? (2) What do you take it to be? (3) Judging 
by the inactivity of this substance Avith dilute acid, does lead 
here exhibit acid-forming or base-forming properties? 

To the filtrate add a few drops of sulphuric acid. (4) What 
did the filtrate contain? (5) State in logical order the steps 
by which you came to this conclusion. (6) To what class of 
elements does lead appear to belong in this compound? 

(7) What is the name and formula of the acid of which 
lead dioxide is the anhydride? (8) Of the lead salt of this 
acid ? (9) What is the evidence that red lead is this compound? 



504 LABORATORY EXERCISES 

(10) Write an equation for the reaction between red lead and 
nitric acid. 

Experiment 135. Preparation of Lead Dioxide. 

Reference : 706. 

Discussion: When a lead salt in which the metal has the 
valence 2 is treated with an alkaline solution of a hypochlorite 
or a hypobromite, oxidation takes place and lead dioxide is 
precipitated. In the preparation described below the hypo- 
bromite is prepared by dissolving bromine in a cold solution 
of sodium hydroxide: 

2NaOH + Br, = NaOBr + NaBr + H,0 

To this solution which contains an excess of sodium hydroxide 
is added a solution of lead acetate. The equation for the 
reaction which takes place is as follows : 

Vh (0,11,0,), + NaOBr + 2NaOH = 

PbO^ +2NaC2H302 + NaBr + H^O 

Materials: Bromine, crystalline lead acetate, sodium hy- 
droxide. 

Directions: Dissolve 17 grams of sodium hydroxide in 200 
c.c. of water, and add to the cooled solution 5 c.c. of bromine. 
Dissolve 34 grams of lead acetate, Pb(C,H302)o,3H,0, in 200 
c.c. of hot water and cool the solution. (If a small amount 
of substance remains insoluble it will not affect the result.) 
Add the solution of sodium hypobromite to the solution of 
lead acetate in a large beaker, and heat the mixture to boiling 
for 5 minutes. Let the solid settle and then pour off carefully 
most of the liquid. Fill the beaker with hot water nearly to 
the top, heat to boiling, let the solid settle and decant the 
liquid as before. Repeat the washing twice in this way. Filter 
off the lead dioxide, wash it twice with about 100 c.c. of boiling 
water, and set it aside to dry. 

Report (1) the weight of the dry dioxide, (2) the theoretical 
weight calculated from the weight of lead acetate used, and 
(3) the percentage of this obtained in the preparation. 



506 LABORATORY EXERCISES 

Experiment 136. Salts of Lead. 
Reference : 707, 709, 710. 

Discussion : Many lead salts are insoluble and thus can be 
prepared by precipatition from solution. A knowledge of their 
appearance is important in identifying compounds of lead. 

Materials: Solutions of lead nitrate, and of potassium or 
sodium iodide, hydrogen sulphide. 

Directions: (a) Dilute 3 c.c. of a solution of lead nitrate 
with 5 c.c. of water and bubble hydrogen sulphide into the 
solution. (1) Name and describe the precipitate. (2) Write 
an equation for its formation. Add 5 c.c. of dilute nitric acid. 
(3) Does the precipitate dissolve in nitric acid? 

(h) To 3 c.c. of lead nitrate diluted with 5 c.c. of water 
add 10 drops of dilute hydrochloric acid. (1) Name and 
describe the precipitate. (2) Write an equation for its forma- 
tion. Heat the solution. (3) What change occurs? Cool it 
again. (4) What appears? (5) What do you conclude about 
the solubility of lead chloride? (6) What other insoluble 
chlorides have you made? (7) How does lead chloride differ 
from them? 

(c) To 3 c.c. of lead nitrate diluted as before add 3 c.c. 
of potassium or sodium iodide solution. (1) Write an equation 
for the reaction. Heat the solution to boiling. (2) What hap- 
pened? Cool the solution again. (3) Name and describe the 
precipitate. 

(d) To 3 c.c. of lead nitrate diluted as before add 10 drops 
of dilute sulphuric acid. (1) Name and describe the pre- 
cipitate. (2) Compounds of what other metals show a similar 
result upon the addition of sulphuric acid? (3) From the 
foregoing experiments select two reactions which in conjunc- 
tion can be used to identify compounds of lead in solution. (4) 
State the procedure in testing for lead. 



SILVER AND COPPER 

Experiment 137. Preparation of Metallic Silver. 

Reference : 714. 

Discussion : Silver may be prepared from its compounds 
by displacement by a metal above it in the electrochemical 
series, or by a process of reduction. Numerous reducing agents 
may be used in the latter case in solution, or the reaction may 
be accomplished by heating the silver compound in a reducing 
flame. 

Materials: Ten cent pieces cut in quarters, concentrated 
nitric acid, mercury, anhydrous sodium carbonate, charcoal. 

Directions: (a) Place a quarter of a 10 cent piece in a 
100 c.c. beaker, pour upon it 5 c.c. of water and 5 c.c. of 
concentrated nitric acid. When the metal has disappeared, 
pour a few drops of the solution into a test tube and add 
ammonium hydroxide very slowly until an excess is present. 
(1) What is the color of the solution? (2) For what is this 
a test? (3) With what metal is coinage silver alloyed? 

(h) Put 1 c.c. of mercury in an evaporating dish and add 
about a third of the solution in the beaker. Set the dish aside 
in your desk until the next exercise, when the following ques- 
tions should be answered. (1) Describe the substance deposited 
on the mercury. (2) What is it? (3) Write an equation for 
its formation. (4) State why this reaction takes place. 

Remove the crystals with pincers and pour the mercury into 
the bottle prepared to receive it. Take great care that no 
mercury gets into the sink. (5) Why? 

(c) Heat the solution remaining in the beaker and add 
dilute hydrochloric acid to it slowly until no more precipitation 
takes place. Add hot water until the beaker is nearly full, 
stir the precipitate, and when it has settled, pour off as much 
of the water as possible without losing any of the precipitate. 

508 



510 LABORATORY EXERCISES 

Wash the precipitate three times with hot water in this way 
and finally transfer it to a filter. Spread out the filter paper, 
scrape off the silver chloride, put some of it in a cavity in a 
piece of charcoal, and cover it with an equal volume of an- 
hydrous sodium carbonate. Heat the mixture in a reducing 
flame of a blow pipe and when there is space in the cavity add 
the remainder of the silver chloride and more sodium carbonate. 
P^inally heat the mixture until metallic globules appear. Poke 
these together and fuse them into a single mass. (1) What is 
the metal? (2) By what kind of process was it foi-med? (3) 
What was the agent? (4) Was any oxide deposited about the 
sides of the cavity, as in the case of lead? (5) Why? 

Experiment 138. Cupric Compounds. 

Reference: 724, 729. 

Discussion : When sodium hydroxide is added to a solution 
of a cupric salt, cupric hydroxide is precipitated. By trying 
the action of this compound with an excess of sodium hy- 
droxide, evidence is gained as to whether or not copper exhibits 
acid-forming properties, as do many of the metals already 
studied. Cupric hydroxide is rather unstable, so that even 
below 100° it is partially dehydrated. 

As has been noted in numerous cases already, ammonium 
hydroxide often differs from sodium hydroxide in its action 
upon salts, particularly when present in excess. In the case 
of cupric salts the ammonia present in the solution forms a 
complex ion of the formula Cu(NH.5)4++. 

Cupric ferrocyanide, Cu2Fe(CN)8, is sufficiently character- 
istic in appearance to be serviceable in testing foi' cupric salts. 

Materials: Solutions of cupric sulphate, and of potassium 
ferrocyanide. 

Directions: (a) To 5 c.c. of cupric sulphate solution add 
sodium hydroxide solution a few drops at a time and, finally, 
in excess. (1) Name and describe the precipitate. (2) Did 
it react with an excess of sodium hydroxide? (3) What do 
you conclude as to the acid-forming properties of copper? (4) 
Why should this be anticipated from the position of copper 
in the periodic classification? 



512 LABORATORY EXERCISES 

Heat the contents of the test tube to boiling. (5) How did 
the precipitate change? (6) What is the nature of the reaction? 

{h) To 5 c.c. of cupric sulphate solution add a few drops 
of ammonium hydroxide and then an excess of this reagent. 
(1) What was the precipitate first formed? (2) What was 
the final appearance of the solution? (3) To what ion is the 
color due? (4) Write equations for the reaction caused by 
the first addition of ammonium hydroxide and for the one with 
the excess of the reagent. (5) With compounds of what metals 
do similar reactions take place? 

(c) To 5 c.c. of cupric sulphate solution add a few drops 
of potassium ferrocyanide solution. (1) What happened? (2) 
Write an equation for the reaction. (3) In what procedure 
is this reaction employed? (4) What other reaction is used for 
this purpose? 

Experiment 139. Cuprous Compounds. 

Reference : 723, 725, 726. 

Discussion : When soluble cupric compounds are reduced 
in alkaline solution, cuprous oxide is precipitated. Since 
ordinary cupric salts can not exist in an alkaline solution, 
because of the precipitation of cupric hydroxide, complex com- 
pounds must be used. Such a compound is formed by adding 
a solution containing sodium potassium tartrate and sodium 
hydroxide to one of cupric sulphate. When this mixture, 
known as Fehling's solution, is boiled with a reducing agent, 
such as glucose, cuprous oxide is precipitated. 

Cuprous chloride can be prepared by the reduction of 
cupric chloride by copper in the presence of concentrated 
hydrochloric acid. The soluble compound, HoCuClg, at first 
formed under these conditions, decomposes upon dilution, with 
the precipitation of the simple chloride. 

The oxidizing action of cupric compounds is exemplified 
in the action between a soluble cupric salt and potassium 
iodide, where cuprous iodide and iodine are formed instead of 
cupric iodide. 

Materials: Fehling's solution, solutions of glucose, of 
cupric chloride, of cupric sulphate, of potassium iodide, con- 
centrated hydrochloric acid, copper filings, starch. 



514 LABORATORY EXERCISES 

Directions: (a) Mix 5 c.c. of Fehling's solution number 1 
with 5 c.c. of Fehling's solution number 2. (1) What does each 
of these solutions contain? (2) What is the appearance of 
the copper compound produced? (3) What is the nature of 
its composition? 

Add 1 c.c. of glucose solution and boil the mixture. (4) 
Name and describe the precipitate. (5) What kind of process 
has the copper undergone? (6) What was the agent? 

(5) Dissolve 2 grams of cupric chloride in 15 c.c. of water 
in an Erlenmeyer flask, add 3 c.c. of concentrated hydrochloric 
acid and 5 grams of copper turnings and boil the solution in 
the hood until the green tint is no longer perceptible in the 
liquid. (1) What is the formula of the copper compound in 
solution? (2) What kind of process has the copper undergone? 

(3) What was the agent? 

Pour the liquid into a large test tube nearly full of water. 

(4) Name and describe the precipitate. (5) Write two equa- 
tions to show the steps by which it was formed from cupric 
chloride. (6) If the water assumes a slightly bluish color, 
what is the cause? Allow the sunlight to fall upon the pre- 
cipitate. (7) What change in color is observable? (8) The 
chloride of what other metal resembles this compound in solu- 
bility and in change in sunlight? (9) Why might this similarity 
have been anticipated? 

(c) Dilute 3 c.c. of cupric sulphate solution with 3 c.c. of 
water, add 3 c.c. of potassium iodide solution and filter. (1) 
Name and describe the precipitate. 

Add a few drops of the filtrate to a test tube half full of 
starch solution. (2) What was the result? (3) What was in 
the filtrate? (4) By what kind of process was it formed? (5) 
What was the agent? (6) Write an equation for the reaction. 
(7) In what previous experiment was the precipitate formed 
by a similar process? 

Experiment 140. Preparation of Cupric-ammonia Sulphate. 

Reference : See Experiment 111. 

Discussion : A number of the soluble salts formed as the 
result of the addition of ammonia to the salts of copper crystal- 



516 LABORATORY EXERCISES 

lize well and can be prepared readily. The compound formed 
from copper sulphate has the composition represented by the 
formula Cu(NH3),S0„H20. 

Materials: Copper sulphate. 

Directions : Place 20 grams of finely powdered copper sul- 
phate, CuSO^jSHgO in a small flask and, under the hood, pour 
upon it 25 c.c. of concentrated ammonia. Shake the flask for 
about 1 minute and then heat it gently until all the solid has 
dissolved. Place the flask in a large beaker containing cold 
water. When the solution has cooled, hold the flask under 
running water. Filter off the crystals. Press them with the 
end of a rod and let them drain. Remove the filter paper, open 
it, and place it on two or three sheets of filter paper. Press 
the crystals with a piece of paper and set them aside to dry. 
(1) Weigh the crystals. (2) Calculate the weight of the com- 
pound that should be formed from the amount of copper sul- 
phate used. (3) Calculate the percentage of this amount ob- 
tained in the experiment. 



IRON 

Experiment 141. Tests for Iron. 

Reference : 751, 753, 756. 

Discussion : Iron forms complex cyanides with potassium 
cyanide, both in the ferrous and ferric conditions, in which 
it is part of the negative complex radicals. These radicals 
form salts, which are highly colored, in one case with ferric 
iron as the positive part of the salt and in another with ferrous 
iron. Since most ferrous salts are easily oxidized by the air, 
they are seldom free from ferric compounds. Ferrous am- 
monium sulphate, however, is less easily oxidized than the 
simple ferrous sulphate and is therefore used in this experi- 
ment. Ferric thiocyanate is also highly colored and is thus 
serviceable as a test for ferric salts. 

Materials: Solutions of ferric chloride, of ferrous am- 
monium sulphate, of potassium ferrocyanide, of potassium 
ferricyanide, of ammonium thiocyanate. 

Directions : Put 5 c.c. of ferric chloride solution into a 
small test tube, fill the tube with water, shake the solution 
and then pour a third of it into each of two other test tubes. 
Dilute and divide into thirds in the same way 5 c.c. of ferrous 
ammonium sulphate solution. Add a few drops of solutions of 
potassium ferrocyanide, of potassium ferricyanide, and of am- 
monium thiocyanate to the three solutions of ferric chloride 
respectively and likewise to the three solutions of ferrous am- 
monium sulphate. (1) Tabulate the results obtained in these 
six tests. (2) Do the results indicate the presence of any 
ferric salt in the ferrous salt? Would you expect the pro- 
duction of any color in the absence of ferric salts? (3) State 
one test for a ferrous salt and two tests for a ferric salt. 
(4) Write equations for the reactions taking place in the six 
test tubes. 

518 



520 LABORATORY EXERCISES 

Experiment 142. Reduction and Oxidation of Iron Salts. 
Reference : 127, 128, 752. 

Discussion : Soluble ferric salts may be reduced by various 
reagents such as stannous chloride, by metallic iron in the 
presence of acids, and by nascent hydrogen. In this connection 
the difference in activity between hydrogen gas and nascent 
hydrogen can be well illustrated. 

Materials : Solutions of ferric chloride, of stannous chloride, 
of potassium ferrocyanide, of potassium ferricyanide, iron, zinc. 

Directions: {a) To 5 c.c. of ferric chloride solution add 
5 c.c. of stannous chloride solution. Test a portion of the 
solution for a ferrous salt. (1) How did you test? (2) With 
what result? (3) Write an equation for the reaction. 

{h} To 5 c.c. of ferric chloride solution add 5 c.c. of dilute 
hydrochloric acid and 5 grams of iron filings. After five 
minutes, test a portion of the solution for a ferrous salt. (1) 
What was the result of the test? (2) Write an equation for 
the original reaction. (3) With what chloride have you carried 
out a similar experiment and with what result ? (4) Which 
chloride would you expect to be formed when iron is treated 
with hydrochloric acid? (5) Why? 

Vei'ify your prediction by treating 5 grams of iron filings 
with 10 c.c. of dilute hydrochloric acid and after five minutes 
testing portions of the solution for a ferrous salt and for a 
ferric salt. (6) What were the results of the tests? (7) Did 
your prediction pi'ove correct? (8) Was there any evidence 
to the conti'ary? (9) If so, how do you account for it? 

(c) Set up an apparatus for generating hydrogen from 
hydrochloi'ic acid with zinc and bubble the hydrogen into 5 c.c. 
of ferric chloride solution for three minutes. Test some of 
the solution for a feri'ous salt. (1) What was the result of 
the test? (2) Will hydrogen gas reduce a ferric salt in solu- 
tion? 

(d) To 5 c.c. of ferric chloride add 5 c.c. of dilute hydro- 
chloric acid and a strip of zinc. After five minutes, test a 
poi-tion of the solution for a ferrous salt. (1) What was the 
result of the test? (2) Had the ferric chloride been reduced? 
(3) What was the agent? (4) Wi'ite an equation for the 



522 LABORATORY EXERCISES 

reaction. (5) How do you account for the difference between 
the result in this case and in part (c) ? 

(e) To 5 c.c. of a solution of ferrous ammonium sulphate 
add 5 c.c. of dilute nitric acid and heat to boiling. Test the 
solution for the presence of a ferrous salt. (1) What did you 
use and with what result? (2) What happened to the ferrous 
salt? (3) What was the agent? 

Try to bring about this same kind of change by using some 
other reagent than nitric acid. (4) What did you use? (5) 
How did you prove that the desired reaction had been accom- 
plished? 

Experiment 143. The Preparation of Ferric Ammonium 
Alum. 

Reference: 752, 688. 

Discussion : The sulphates of the trivalcnt metals form 
double salts with the sulphates of ammonium and the alkali 
metals. All these salts have analogous compositions which are 
-represented by the general formula: 

I III 

M2S0„M2(S0J3,24H20 

They are isomorphous with one another, crystallize in octahedra, 
and are called alums. In the experiment described below 
ferric ammonium alum is prepared from ferrous sulphate and 
ammonium sulphate. The ferrous sulphate is first oxidized by 
means of nitric acid to ferric sulphate. 

Materials: Ferrous sulphate, ammonium sulphate, concen- 
trated nitric acid, materials used in testing for ferric and for 
ammonium compounds. 

Directions : Place 28 grams of crystalline ferrous sulphate, 
FeS04,7H20 in a beaker and add 25 c.c. of water and 5 c.c. of 
concentrated sulphuric acid. Heat the mixture until the salt 
dissolves. Support the beaker on a tripod under the hood 
and add cautiously in small portions 10 c.c. of concentrated 
nitric acid. Heat the solution cautiously and boil it for 5 
minutes. (1) What does the change in color of the solution 
indicate? (2) Wi'ite the equation for the reaction involved. 



524 LABORATORY EXERCISES 

Dissolve 13 grams of ammonium sulphate in 25 c.c. of boil- 
ing water and filter if necessary. Mix the two solutions and 
set aside the beaker until the next exercise. (3) Describe the 
form of the crystals. 

Filter off the crystals, spread them on a filter paper and 
when they are dry weigh them. (4) Record the weight. 
Determine if the salt contains water of crystallization and 
(5) state what you did to find out. Apply a test to determine 
if the salt contains ferric iron and (6) state what you did and 
your conclusion. Test the crystals for an ammonium com- 
pound. (7) What did you do and with what result? (8) Give 
the name and formula of the crystals. (9) To what class of 
substances do they belong? (10) Write the formulas for two 
other substances Avhich are isomorphous with this compound. 



CHROMIUM 

Experiment 144. Chromic Anhydride, Chromates, and 
Dichromates. 

Reference: 777. 

Discussion : As was stated in Experiment 93, it is often 
convenient to regard a salt as composed of the oxide of a 
metal and of the oxide of a non-metal, that is, of an acid anhy- 
dride. Thus one molecule of chromic anhydride with one of 
potassium oxide may be regarded as forming the combination 
K20,Cr03 or K2Cr04, potassium chromate, and two of chromic 
anhydride with one of potassium oxide as forming K20,2Cr03 
or K2Cro07, potassium dichromate. 

When potassium chromate is treated with concentrated sul- 
phuric acid, the chromic acid, which would be formed by double 
decomposition, becomes dehydrated by the concentrated sul- 
phuric acid and chromic anhydride is produced. A similar 
reaction takes place when potassium dichromate is treated with 
concentrated sulphuric acid. 

Materials: Solutions of potassium dichromate, of potas- 
sium chromate, of concentrated sodium hydroxide, potassium 
dichromate, glass wool. 

Directions: (a) Powder in a mortar 1 gram of potassium 
dichromate, transfer it to a test tube, add 5 c.c. of water and 
heat it to boiling. Cool the tube under the tap, allow the 
crystals to settle and pour off the saturated solution into a 
small beaker. To this add 10 c.c. of concentrated sulphuric 
acid and when the liquid has again become cold, filter out 
the precipitate through a funnel into the tip of which has been 
forced some glass wool. (1) Name and describe the precipitate. 
(2) Write an equation for its formation. (3) Why was paper 
not used for filtering? 

Rest the funnel in a clean test tube and pour over the 

526 



528 LABORATORY EXERCISES 

precipitate drop by drop 2 c.c. of concentrated sodium hydroxide 
solution. (4) What color is the liquid? (5) What does this 
lead you to suppose the compound to be? (6) Write an equa- 
tion for its formation. 

(b) To 5 c.c. of potassium chromate solution add 5 c.c. of 
sulphuric acid. (1) What change in color occurred? (2) How 
do you account for it? 

(c) To 5 c.c. of potassium dichromate solution add sodium 
hydroxide slowly until a change of color appears. (1) What 
Avas the change? (2) Write an equation for the reaction. (3) 
What are the relations in composition among chromic anhy- 
dride, a chromate, and a dichromate? 

Experiment 145. Reduction of Potassium Dichromate. 

Reference: 775, 779. 

Discussion : Potassium dichromate can be reduced by fer- 
rous salts in acid solution. This reaction is frequently used 
for the oxidation of ferrous salts in the course of quantitative 
analysis. In writing equations for the reduction of a dichro- 
mate, it is convenient again to regard the compound from the 
point of view of chromic anhydride and to write hypothetical 
steps in the reaction. The anhydride loses oxygen and the 
resulting chromic oxide and potassium oxide are converted into 
salts by the acid. The oxygen furnished in this way accom- 
plishes the oxidation. 

Potassium dichi-omate will oxidize oxalic acid, C2H402,2K,0, 
in sulphuric acid solution to water and carbon dioxide. Upon 
evaporating the resulting solution of potassium and chromic 
sulphates, a double compound of these salts with twenty-four 
molecules of water of crystallization, which belongs to the class 
of alums, separates out in well-defined crystals. If the solution 
is heated, however, a green compound of very dilferent proper- 
ties is formed and no crystals can be obtained. 

Materials: Solutions of ferrous ammonium sulphate, of 
potassium dichromate, of potassium ferrocyanide, of potassium 
f erricyanide ; potassium dichromate, oxalic acid. 

Directions: {a) To 10 c.c. of ferrous ammonium sulphate 
solution add 10 c.c. of water, 5 drops of concentrated sulphuric 



530 LABORATORY EXERCISES 

acid, and 10 c.c. of potassium dichromate solution. Test a 
small amount of the solution for a ferrous salt, and if any 
is present, add successive portions of the dichromate solution 
of 5 c.c. each until no ferrous salt remains. Test the solution for 
a ferric salt. (1) What test did you make and with what result ? 
(2) Write equations for the reaction in three hypothetical steps 
and then combine them into a single equation. (3) In what 
analytical process is this reaction employed ? (4) Explain how 
it is used. 

(h) (1) Write equations in steps for the reduction of potas- 
sium dichromate by oxalic acid in the presence of sulphuric 
acid and combine the equations into a single one. (2) Calculate 
the weight of crystallized oxalic acid required to react with 
5 grams of potassium dichromate and (3) the volume of sul- 
phuric acid (Sp. gr. = 1.84) required. 

Add 5 grams of potassium dichromate to 35 c.c. of hot water 
in a beaker, and when it has dissolved, cool the solution, and 
add the calculated amount of sulphuric acid. Place the beaker 
in cold water in a pneumatic trough and add two grams more 
than the calculated weight of oxalic acid, a few crystals at a 
time, with constant stirring and with care that the solution 
remains cold. (4) What color is the solution? 

Pour 5 c.c. of the solution into a test tube and heat it to 
boiling. (5) What color is the solution? (6) Why was the 
solution cooled during the reaction? Transfer the remainder 
of the solution to an evaporating dish and set it aside in your 
desk until the next exercise, when crystals should have begun 
to form. Pick out with the pincers four of the largest crystals, 
put them in a small beaker, and pour over them the solution, 
taking care that no other crystals get into the beaker. (7) 
Give the name and write the formula of these crystals. Again 
set the solution a.side to allow the crystals to grow and at each 
exercise remove the small crystals. In this way very large 
well-formed crystals may be produced. 



MANGANESE 

Experiment 146. Potassium Permanganate as an Oxidizing 
Agent. 

Reference: 794. 

Discussion : Potassium permanganate is a useful oxidizing 
agent in quantitative analysis. In neutral or alkaline solution 
the manganese is reduced from the valence of seven to the 
valence of four, thus furnishing oxidation to the extent of three 
valences, but in acid solution it is reduced to the valence of 
two, furnishing five valences for oxidation. 

Materials: Solutions of ferrous ammonium sulphate and 
of potassium permanganate. 

Directions: (a) To 5 c.c. of ferrous ammonium sulphate 
solution add potassium permanganate solution very slowly, a 
few drops at a time toward the end, until a slight excess is 
present, as indicated by the color of the solution. (1) What 
is the precipitate? (2) Write hypothetical equations, similar 
to those written for potassium dichromate in Experiment 145, 
to show the composition of potassium permanganate as being 
potassium oxide and an acid anhydride, and to show the reduc- 
tion of the anhydride as it occurs in alkaline or neutral solu- 
tion. (3) Write equations for the remainder of the hypothetical 
steps in the oxidation of the ferrous salt and combine them 
into a single equation for the reaction. 

(h) To 5 c.c. of ferrous ammonium sulphate solution add 
10 drops of concentrated sulphuric acid and then potassium 
permanganate solution as in part (a). (1) What is the appear- 
ance of the solution? (2) Write an equation, similar to the 
one you wrote in part (a), to show the amount of oxidation 
furnished by potassium permanganate in acid solution. (3) 
Complete the equations for the reaction with the ferrous 
sulphate. 

532 



INDIVIDUAL PROBLEMS 

Individual assignments for the identification and prepara- 
tion of compounds are made during the latter part of the year. 
Work on these problems should be done only after the class 
work for the day has been finished. 

IDENTIFICATION OF COMPOUNDS 

Discussion : Application of the tests studied in many of 
the preceding experiments and of tests now to be used for 
the first time is made by applying these tests to compounds 
given out individually for identification. Single compounds 
only will be issued, since the analysis of mixtures often requires 
a process of separation, involving a systematic scheme of 
qualitative analysis, before tests can be made. 

The procedure in general, after noting the appearance of 
the substance, is to get the material into solution, first trying 
water alone and subsequently, if necessary, acids. The metallic 
element of the compound is identified first. An indication of 
what this may be is obtained by adding ammonium hydroxide 
to the solution, and then a confirmatory test is tried to prove 
definitely the presence of the element suspected. From the 
knowledge of the solubility of the compound and of the metallic 
element which it contains, a conclusion may be drawn by 
reference to the table of solubilities, as to what acid radical 
may be present. The addition of silver nitrate and of barium 
chloride to portions of the solution causes the precipitation 
of certain groups of acid radicals if present, and thus offers 
additional basis for judging what radicals must be tested for. 
Finally a few confirmatory tests will identify the acid radical 
and consequently the compound. 

Often the acid radical can be recognized when the com- 

534 



IDENTIFICATION OF COMPOUNDS 535 

pound is first dissolved, particularly when a gas is evolved in 
the process. 

Many opportunities for thoughtless misinterpretation exist 
in making these tests, as for example in testing for a chloride 
when the compound has been dissolved in hydrochloric acid. 
A student can frequently save himself needless work by 
thoughtful interpretation of his results. 

Complete notes must be kept, since a negative result in 
making a test is as important as a positive one. Abbreviations 
and symbols may be used, as illustrated in the following sample 
report, which is given as a suggestion for the form of notes. 

No. 6 April 6, 1922. 

White powder, insol. water, sol. cold dil. HNO3 with effervescence. 
Metal. 

NH40H = wh. ppt. insol. excess, =Pb, Hg, Bi, Sb, Sn, Al or Mg. 

HCl and H2S = black ppt., =Pb, Hg, or Bi. 

HC1= . . . H2S04= . . . NotPb. 

SnCl2= .... NotHg. 

NH4OH ppt. dissolved in 2 drops cone. HCl, water gave wh. ppt. = Bi 
Acid Radical. 

Table of sol. HF, H2SO4, HCl, HBr, HCIO3, HNO3 impossible. 

HCl, to solid salt, gas, colorless, odorless. 

Ca(0H)2 on rod milky = CO2 . 
Report, bismuth carbonate. 

Directions: Procure at the storeroom the substance 
assigned for identification by giving your name and the number 
of the salt you want, e.g., first, second, etc. A little of the 
salt should always be reserved until after the report on the 
analysis has been made, so that the instructor may test the 
original sample, if a test is necessary. A second sample of any 
salt may be got if the first has been used up before a satisfactory 
result has been obtained. 

Examine and record the appearance of the substance. If it 
is not already powdered, grind it in the mortar. Try a pinch 
of the substance with each of the following solvents until one 
is found which will dissolve it : cold water ; boiling water ; cold, 
dilute hydrochloric acid; boiling, dilute nitric acid; hot aqua 
regia (three parts of concentrated hydrochloric acid and one 
part of concentrated nitric acid) . When using concentrated acids 



536 LABORATORY EXERCISES 

avoid large quantities. Sulphides which dissolve in boiling, 
dilute nitric acid will leave a ball of spongy sulphur on the 
surface, which, in itself, serves as a test for a sulphide. When 
a suitable solvent has been found, get a considerable quantity 
of the substance into solution, using about one spatulai'ul to 
50 c.c. of solvent. 

If acid has been used to dissolve the salt, neutralize as nearly 
as possible with sodium hydroxide, since the presence of am- 
monium salts prevents the precipitation of certain metals b}^ 
ammonium hydroxide. Do not allow a permanent precipitate 
to form, even if the solution has to be left slightly acid. If 
a precipitate forms, redissolve it with a few drops of the acid. 

IDENTIFICATION OF THE METAL 

To 5 C.C. of the solution add ammonium hydroxide drop 
by drop and finally in excess. Record the result and the con- 
clusion to be drawn in accordance with the facts given below 
as to the action with ammonium hydroxide. 

If no precipitate appears, try the flame test by moistening 
the platinum wire in hydrochloric acid, dipping it into the salt 
to be tested and then into the flame. This will identify barium 
and potassium. Calcium and strontium, however, give flames 
so nearly alike that the confirmatory tests must always be 
applied. A yellow flame does not indicate sodium unless it 
persists ten or fifteen seconds, as sodium compounds are present 
as impurities in nearly all chemicals. Never report sodium 
unless the test for potassium has first been applied. Absence 
of flame test indicates arsenic or ammonium compounds. 

If ammonium hydroxide gives a white precipitate insoluble 
in excess, acidify a fresh portion of the solution with hydro- 
chloric acid and pass in hydrogen sulphide. A black precipitate 
indicates lead, mercuric mercury, or bismuth, which should be 
tested for only in the order given. An orange precipitate in- 
dicates antimony; a yellow or brown one, tin; and no precipi- 
tate, aluminium or magnesium. A dirty white precipitate of 
sulphur sometimes is formed, which has no significance. 



IDENTIFICATION OF COMPOUNDS 537 

ACTION WITH AMMONIUM HYDROXIDE 

No Precipitate. Arsenic, barium, strontium, calcium, sodium, 
potassium, or ammonium. 

White Precipitate, soluble in excess. Cadmium or zinc. 

White Precipitate, insoluble in excess. Lead, mercury (mer- 
curic), bismuth, antimony, tin, aluminium, magnesium. 

Blue-Green Precipitate. Chromium. 

Blue Precipitate, soluble in excess to a brown solution. 
Cobalt. 

Blue Precipitate, soluble in excess to a dark blue solution. 
Copper. 

Light, Dirty Green Precipitate. Iron (ferrous). 

Green Precipitate, soluble in excess to a light blue solution. 
Nickel. 

Brownish White Precipitate, turning darker on exposure. 
Manganese. 

Red-Brown Precipitate. Iron (ferric). 

Brown Precipitate, soluble in excess to a colorless solution. 
Silver. 

Black Precipitate. Mercury (mercurous). Some iron salts 
also give a black precipitate. 

CONFIRMATORY TESTS FOR METALS 

Make confirmatory tests for the element or elements indi- 
cated as probably present. 

Aluminium. Ammonium chloride and ammonium hydroxide 
give a white precipitate of aluminium hydroxide, insoluble in 
excess ; soluble in sodium hydroxide. 

Ammonium. Heated with sodium hydroxide, ammonia is 
formed, which can be detected by the odor or by moist litmus 
paper held in the gas. 

Antimony. Hydrogen sulphide precipitates orange antimony 
sulphide in the presence of dilute hydrochloric acid. The sul- 
phide is soluble in yellow ammonium sulphide, and soluble in 
hot, concentrated hydrochloric acid. 

Arsenic. Hydrogen sulphide precipitates yellow arsenic 



538 LABORATORY EXERCISES 

sulphide in the presence of dilute hydrochloric acid. The sul- 
phide is soluble in yellow ammonium sulphide, insoluble in 
hot, concentrated hydrochloric acid. 

Barium. Imparts green color to the flame. Sulphuric acid 
precipitates white barium sulphate. 

Bismuth. Hydrogen sulphide precipitates black sulphide 
in presence of dilute hydrochloric acid. 

Ammonium hydroxide precipitates white hydroxide which, 
when filtered, washed, and dissolved on the filter with two or 
three drops of concentrated hydrochloric acid, and poured into 
water, gives a white precipitate of bismuth oxychloride. 

Cadmium. Hydrogen sulphide precipitates yellow cadmium 
sulphide in presence of dilute hydrochloric acid, which is in- 
soluble in yellow ammonium sulphide. 

Calcium. Gives a dull red color to the flame. Gives no 
precipitate with calcium sulphate (distinction from strontium). 

Chromium. Mixed with sodium carbonate and potassium 
nitrate and fused in a platinum wire loop gives yellow potas- 
sium chromate. A solution of the bead in acetic acid, gives 
yellow precipitate with lead acetate. 

Cobalt. Borax bead colored deep blue. 

Copper. Hydrogen sulphide precipitates black copper sul- 
phide in presence of dilute hydrochloric acid. Salts give green 
flame. Potassium ferrocyanide gives brown precipitate of cop- 
per ferrocyanide in presence of acetic acid. 

Iron. Potassium ferricyanide gives blue color with ferrous 
salts. Potassium feri-ocyanide gives blue color with ferric salts 
and potassium thiocyanate gives red color with ferric salts. 

Lead. Hydrochloric acid precipitates from solutions not 
too dilute, white lead chloride, soluble in hot Avater. 

Sulphuric acid precipitates white lead sulphate. 

Magnesium. No precipitate with ammonium chloride and 
ammonium hydroxide. Disodium phosphate added to the mix- 
ture precipitates white, crystalline magnesium ammonium 
phosphate. 

Manganese. Ammonium sulphide precipitates flesh-colored 
manganese sulphide. 

Compounds when fused with sodium carbonate and potas- 



IDENTIFICATION OF COMPOUNDS 539 

slum nitrate in platinum wire loop, give green potassium man- 
ganate. 

Mercury. Mercuric. Stannous chloride precipitates white 
mercurous chloride or gray mercury. 

Mercurous. Hydrochloric acid precipitates white mercurous 
chloride, which is converted into a black compound by am- 
monium hydroxide. 

Nickel. Salts are usually green. Compounds color borax 
beads violet when hot, red-brown when cold. 

Potassium. Violet flame. If flame is yellow, look through 
two pieces of cobalt glass to detect the reddish color in the 
flame. Do not be misled by the color of the heated wire. 

Silver. Hydrochloric acid precipitates white silver chloride, 
soluble in ammonium hydroxide. 

Sodium. Yellow flame persisting for ten or twenty seconds. 
Before reporting sodium, prove the absence of potassium and 
ammonium by proper tests, since an impurity of sodium gives 
a strong flame. 

Strontium. Bright red flame. White precipitate with cal- 
cium sulphate, forming slowly in concentrated solution. (Com- 
pare with calcium.) 

Tin. Stannic. Hydrogen sulphide precipitates yellow stan- 
nic sulphide in presence of dilute hydrochloric acid. Pre- 
cipitate soluble in yellow ammonium sulphide, soluble in hot, 
concentrated hydrochloric acid (distinction from arsenic). 

Stannous. Hydrogen sulphide precipitates brown stannous 
sulphide in the presence of dilute hydrochloric acid. Precipi- 
tate soluble in yellow ammonium sulphide. 

Zinc. In acetic acid solution, i.e., containing no strong acid 
or alkali, hydrogen sulphide precipitates white zinc sulphide. 

IDENTIFICATION OF THE ACID-RADICAL 

When the metallic element has been identified, determine 
by reference to the table of solubilities on page 543 what acid 
radicals may be present in accordance with the solubility and 
color of the compound. 

To a small portion of the solid substance add a few cubic 



540 LABORATORY EXERCISES 

centimeters of dilute hydrochloric acid and heat. If a gas is 
evolved, identify it by its color, its odor, its action with lime 
water, or with lead acetate paper. If the compound is insoluble 
in the acid and if the previous results indicate the possibility 
of a sulphide, add a strip of zinc to the hydrochloric acid mix- 
ture and test for hydrogen sulphide. 

If no gas is evolved, the presence or absence of a halide 
may be ascertained by adding silver nitrate to a nitric acid 
solution of the compound. 

If ammonium hydroxide gave no permanent precipitate Avith 
a solution of the compound, barium chloride may be added 
to such an alkaline solution and the presence or absence of a 
sulphate, phosphate, borate, or fluoride is indicated by whether 
or not a precipitate appears. When, however, the metallic 
element of the compound is one which forms an insoluble 
chloride, barium nitrate must be substituted for barium 
chloride. 

The procedure described above will have identified the 
acid radical present or have eliminated a large number of 
possibilities. In the former case a confirmatory test Avith a 
fresh portion of the compound should be made, and in the 
latter case confirmatory tests for all the remaining possibilities 
should be tried. 

TESTS FOR THE ACID-RADICALS 

Acetate, (a) The salt heated with a few drops of concen- 
trated sulphuric acid gives the characteristic odor of acetic 
acid. 

(6) A pinch of the salt heated with five drops of concen- 
trated sulphuric acid and 1 c.c. of alcohol gives a fruity odor, 
due to the formation of ethyl acetate. 

Borate. Turmeric paper, after having been dipped in a 
hydrochloric acid solution of the salt, turns bright pink upon 
drying at 100° and dark green upon subsequent treatment 
with ammonium hydroxide. 

Bromide. One c.c. of carbon disulphide added to 5 c.c. of 
the solution is turned cherry rod after 3 c.c. of chlorine water 
has been added and the mixture shaken. 



IDENTIFICATION OF COMPOUNDS 541 

Carbonate. Dilute hydrochloric acid causes the evolution 
of an odorless gas, which gives a precipitate with lime water. 

Chlorate. Silver nitrate gives no precipitate in a nitric acid 
solution ; but a pinch of the salt heated strongly in a dry test 
tube, cooled and dissolved gives a precipitate with silver nitrate. 

Chloride. Silver nitrate gives in nitric acid solution a white 
precipitate, which dissolves readily in ammonium hydroxide. 

Fluoride. To a pinch of the salt in a dry test tube is added 
an equal amount of sand and the mixture is moistened with 
concentrated sulphuric acid and gently heated. A platinum 
wire with a drop of water adhering to a loop in the end, is 
carefully thrust into the tube above the mixture. A white 
precipitate in the drop of water shows the presence of a 
fluoride. 

Hydroxide or Oxide. All acid radicals having been proved 
absent, the compound may be reported as an oxide or hy- 
droxide. The latter may sometimes be distinguished by heat- 
ing a little of the compound in a dry test tube, when a small 
deposit of moisture will collect in the upper part of the tube. 

Iodide. A deep blue color is produced upon adding to the 
solution a few drops of starch solution and a few drops of 
chlorine water. 

Nitrate. When to 5 c.c. of solution is added 5 c.c. of ferrous 
sulphate solution and when 5 c.c. of concentrated sulphuric 
acid is carefully poured dovs-n the side of the test tube, a brown 
ring appears. This test is of value only in the absence of a 
nitrite. 

Nitrite. Dilute sulphuric acid evolves a brown gas from 
the solid when it is heated. 

Phosphate. Three c.c. of ammonium molybdate, heated 
with 1 c.c. of the solution to which has been added 2 c.c. of 
dilute nitric acid, produces a yellow precipitate. 

Silicate. A few drops of concentrated hydrochloric acid 
produce a gelatinous precipitate when added to a few cubic 
centimeters of a concentrated solution of the salt. 

Sulphate. Barium chloride gives a white precipitate, in- 
soluble in hydrochloric acid. 

Sulphide. Hydrochloric acid, or hydrochloric acid and zinc, 



542 



LABORATORY EXERCISES 



evolves hydrogen sulphide, recognizable by its odor and by 
its action on lead acetate paper. 

Sulphite. Hydrochloric acid evolves sulphur dioxide, recog- 
nizable by its odor. 

Thiosulphate. Hydrochloric acid evolves sulphur dioxide 
and at the same time produces a precipitate of sulphur. 

APPENDIX 

Vapor Pressures of Water 

The temperatures are given in degrees centigrade, and the pressures in milH- 
meters of mercury. 



Temp. 


Press. 


Temp. 


Press. 


Temp. 


Press. 


Temp. 


Press. 


0° 


5 


14° 


12 


20° 


17 


26° 


25 


4 


7 


15 


13 


21 


19 


27 


27 


8 


8 


16 


13 


22 


20 


28 


28 


10 


9 


17 


14 


23 


21 


29 


30 


12 


10 


18 


15 


24 


22 


30 


32 


13 


11 


19 


16 


25 


24 


100 


760 



Electromotive Series of the Elements 



Potassium 


Iron 


Silver 


Sodium 


Cobalt 


Palladium 


Barium 


Nickel 


Platinum 


Strontium 


Tin 


Gold 


Calcium 


Lead 




Magnesium 


Hydrogen 


Iodine 


Aluminium 


Copper 


Bromine 


Manganese 


Arsenic 


Oxygen (O ) 


Zinc 


Bismuth 


Chlorine 


Chromium 


Antimony 


Oxygen (OH") 


Cadmium 


Mercury 


- 



IDENTIFICATION OF COMPOUNDS 



543 



-P«oNX 


:^ : :g :^^ :^ :^ :<^^^ : :^^^^^ : :^ 


ac^iqding 


^^ : :<^<«'§:<«^««^<^« :^ 


apiqding 


<<^<1<^<<<<^5<^-<<<<^<1<1<1<J<1<^^<J^^<1'^<: 


o^i3qding 


^^^ :-^^g^^^^^5^^^^^^^^-^ :^ 


a^BoinS 


5 : : :<i :<i<i<i<5-<<<<<:-<^ : :<i^ ■.^< : :<i 


8:^^qdsoqj 


<i^^<i<<<i<^<<i<<-<<<i<i<j<<<i<^<:^^^-<<i-<<5 


9^1-i^TN 


:^ : :^ :^^ : : : : :^^ :*:^^^g^^ : :^ 


a^Bj^i^ 


^^ : :^^^^^^^^^^^^^^^^^^^ : :^ 


apipoj 


^^^^^<^^^^^^^^^^^^<i^^-^^^^^ 


apixojpiCjj 


<^<<^<<^<«<«<< : :<i5^ :^^<i5"^ 


opijoni^ 


^^^^g<g^^g<<g^-i5^ :gg^^^g^^g 


apuoiqo 


^^^^^^^^g^^^^g^^^^^^-^^^^^ 


a^^joiqo 


^^ : :^^^^^^^^^^^^^^^^^^^^ :^ 


a^BuoqjtJQ 


:^ : :-<^<j<: :<i<<<< :<<<<«^<^< : :< 


apiraoig 


^^^^^^^^g^^^^g^^^^^^-^^ : :^ 


a-^BJog 


"^^ : :<i<i^^<i<:<^-5i<i<i-^-< : :<^<^« :< 


a^^-^aoy 


^^ : :^^^^^^^^^^^^^^^^^^^^^^ 

























:::::::::::::::::::::::::: 


:::::::::::::::::::::::::: 


♦Aluminium . . . 

Ammonium . . . 
♦Antimony .... 

Arsenic 

Barium 

♦Bismuth 

Cadmium. . . . 

Calcium 

♦Chromium 

Cobalt 

♦Copper 

Iron++ 

♦Iron+++ 

Lead 

Magnesium. . . 

Manganese . . . 
♦Mercury+ .... 
♦Mercury ++ . . . 

Nickel 

Potassium. . . . 

Silver 

Sodium 

Strontium .... 

Tin++ 

♦Tin++++ 

Zinc 





LABORATORY 


EQUIPMENT 


Balance, beam 




Generator, Kipp 


Balance, platform 




Litmus paper 


Barometer 




Rings 


Clamps 


, 


Ringstand 


Copper wire 




Splinters 


Corks 




Spoon, deflagrating 


Cork borers 




Thread 


Cork press 




Trough, pneumatic 


Crock for waste 




Voltmeter 


Funnel support 








INDIVroUAL 


EQUIPMENT 


5 Beakers, Nos. 1-5 




Pincers, 1 pair 


1 Blowpipe 




1 Pinehcock 


1 Blowpipe tube 




1 Platinum wire, 5 cm. 


4 Bottles, gas, 250 c.c. 




Rubber tubing, 2 ft. 


1 Bunsen burner 




2 Rubber connectors _ 


1 Cloth 




3 Rublier stoppers. No. 5 


1 Crucible and cover 




2, 2-hole 


1 Cylinder, graduated 




1, 1-hole 


1 Evaporating dish, 8 


cm. 


1 Ruler 


4 Flasks 




1 Spoon, horn 


1 Erlenmeyer, 300 c.c. 


1 Sponge 


1 Florence, 100 c.c. 




18 Test tubes 


1 Florence, 500 c.c. 




2, 6-in. 


1 Florence, 750 c.c. 




6, 8-in. 


1 Flame spreader 




1 Test tube brush 


1 File, triangular 




1 Test tube holder 


Filter paper, 1 pkg., 


10 cm. 


1 Test tube rack 


2 Funnels, 6 cm. 




1 Thermometer, -10° to 110' 


2 Glass plates, white. 


5 cm. 


1 Thistle tube 


2 Glass plates, blue, 5 


cm. 


Tongs, 1 pair 


Glass rod, 2] ft. 




1 Triangle, clay 


Glass tubing, 5 ft. 




3 \\'atch glasses 


3 Ignition tubes 




2, 4 cm. 


1 Key 




1, 7 cm. 


Matches 




Weights, 1 box 


1 Mortar and pestle, ^ 


\ cm. 


1 Wire gauze 



544 



APPARATUS 



545 



SPECIAL EQUIPMENT 



Beaker, 1000 c.c. 
Bottle, 150 c.c. 
Bottle, 2000 c.c. 
Burette, 50 c.c. 
Burner, Meker 
Cylinder, graduated, 10 c.c. 
Dish, lead 
Electrodes, carbon 
Flask, 250 c.c. 



Funnel, 10 cm. 

Glass wool 

Pinchcock 

Pipette, graduated, 10 c.c. 

Retort, 250 c.c. 

Sand bath 

Thermometer, 360° 

Weight, brass, 200 gr. 



LIST OF CHEMICALS 



SOLIDS 



Aluminium, sheet 
Ammonium chloride 
Ammonium nitrate 
Ammonium nitrite 
Ammonium sulphate 
Antimony trichloride 
Arsenious oxide 
Asbestos 
Barium chloride 
Barium hydroxide 
Barium nitrate 
Barium oxide 
Barium peroxide 
Bismuth chloride 
Bleaching powder 
Calcium carbonate, marble 
Calcium chloride 
Calcium cyanamide 
Calcium fluoride 
Calcium hydroxide 
Calcium phosphate, primary 
Calcium sulphate 
Charcoal, boneblack 
Charcoal, lumps 
Charcoal, powdered 
Copper, gauze 
Copper, sheet 
Copper, turnings 
Copper nitrate 
Copper sulphate 
Copper sulphide 



Ferrous sulphate 

Ferrous sulphide 

Graphite, powder 

Iodine 

Iron, nails 

Iron, powder 

Iron, steel wool 

Lead, sheet 

Lead acetate 

Lead dioxide 

Lead monoxide 

Lead nitrate 

Lead oxide, red 

Magnesium, powder 

Magnesium, ribbon 

Magnesium carbonate 

Manganese carbonate 

Manganese dioxide 

Mercuric oxide 

Oxalic acid 

ParaflSn 

Phosphorus, red 

Phosphorus pentachloride 

Potassium aluminium sulphate 

Potassium bromide 

Potassium chlorate 

Potassium chloride 

Potassium chromium sulphate 

Potassium dichromate 

Potassium ferrocyanide 

Potassium hydrogen tartrate 



546 LABORATORY EXERCISES 

Potassium hydroxide Sodium nitrate 

Potassium iodide Sodium peroxide 

Potassium nitrate Sodium sulphate, crystalHzed 

Potassium permanganate Sodium sulphite 

Silica, sand Sodium tetraborate, borax 

Silver, tarnished Sodium thiosulphate 

Silver chloride Starch 

Sodium Strontium chloride 

Sodium aluminium sulphate Strontium nitrate 

Sodium ammonium phosphate Sugar 

Sodium carbonate, crystallized Sulphur, powder 

Sodium chloride Tin, granulated 

Sodium chloride, pure Tin, sheet 

Sodium chloride, rock salt Zinc, dust 

Sodium hydrogen carbonate Zinc, pure 

Sodium hydrogen sulphate Zinc, sheet 

Sodium hydroxide Zinc chloride 

LIQUIDS AND SOLUTIONS 

The numbers indicate the weights in grams (jr the volumes in cubic 
centimeters per liter of solution of the respective substances as commonly 
obtained, often in the hydratcd form. 

Acid, acetic, glacial 

Acid, acetic (dilute), IN, 60 

Acid hydrochloric (cone), sp. gr. 1.20, 13N 

Acid, hydrochloric, 20 per cent, 469 c.c. of HCl, sp. gr. 1.20 

Acid, hydnifliloric, 10 per cent, 224 c.c. of HCl, sp. gr. 1.20 

Acid, hydrochloric, IN, 77.6 c.c. of HCl, sp. gr. 1.20 

Acid, hydrochloric, hN, 38.8 c.c. of HCl, sp. gr. 1.20 

Acid nitric (cone), 16N, sp. gr. 1.42 

Acid, sulphuric, IN, 27.8 c.c. of H2SO4, sp. gr. 1.84 

Alc(J^l, ethyl, 95 per cent 

Alcohol, methyl 

Ammonium carbonate, IN, 67 + 150 c.c. NH4OH 

Ammonium chloride, IN, 53.5 

Ammonium molybdate, 75 + 180 c.c. HNO3 

Ammonium oxalate, iN, 35 

Ammonium sulphide (dilute ammonia sp. gr. 0.90 with equiv. vol. of water, 

sat. with H-iS) 
Ammonium sulphide, yellow (NH4)2S+S 
Ammonium thiocyanate, \N, 19 

Arsenious chloride, \N, 8 g. AS2O3 + IO g. NaOH, acidify with HCl 
Barium chloride, ^N, 61 
Benzene 
Bromine 



LIST OF CHEMICALS 547 

Bromine water, sat. 

Calcium chloride, IN, 56 

Calcium hydroxide, 0.04N, sat. 

Calcium sulphate, sat. 

Carbon disulphide 

Carbon tetrachloride 

Chlorine water, sat. 

Chloroform 

Cobalt nitrate, iN, 36 

Copper sulphate, |N, 62 

Ether 

Fehling's solution No. 1, CuS04,5H20, 69 

Fehling's solution No. 2, KNaC4H406,4H20, 246+NaOH, 100 

Ferric chloride, f N, 27 

Ferrous sulphate, jN, <35 

Ferrous ammonium sulphate, jN, 98 

Glucose, 50 

Hydrogen peroxide, 3 per cent 

Hydrogen sulphide, sat. 

Lead acetate, iN, 47 

Lead nitrate, jN, 41 

Magnesium chloride, ^N, 51 

Mercuric chloride, |N, 34 

Mercurous nitrate, |N, 75+4 c.c. HNOa+Hg 

Mercury 

Methyl orange, 0.5 

Nickel nitrate, iN, 36 

Phenolphthalein (1 g. in 500 c.c. alcohol) 

Phosphorus trichloride 

Potassium bromide, |N, 24 

Potassium chloride, ^N, 38 

Potassium chromate, ^N, 49 

Potassium dichromate, |N, 77 

Potassium dichromate, jN, 38 ,: 

Potassium ferricyanide, |N, 27 ■^'' 

Potassium ferrocyanide, jN, 26 i 

Potassium hydroxide, IN, 56 0^ 

Potassium hydroxide, ^N, 28 

Potassium iodide, |N, 33 

Potassium nitrate, ^N, 50 

Potassium permanganate, IN, 32 

Potassium permanganate, 1 gr.+200 c.c. H2SO4 '^ 

Potassium thiocyanate, jN, 29 4^ 

PyrogalHc acid, 10 per cent, 111 

Silver nitrate, IN, 34 

Soap, 30 # 



548 LABORATORY EXERCISES 

Sodium carbonate (anhyd.), ^N, 26 

Sodium chloride, 5.5N, sat. 

Sodium chloride, ^N, 29 

Sodium chromate, jN, 43 

Sodium hydrogen sulphite, sp. gr. 1.26 

Sodium hydroxide, cone, 50 per cent, 765 

Sodium hydroxide, IN, 40 

Silver sulphate, iN, 31 

Sodium hydroxide, 10 per cent, 112 

Sodum hydroxide, |N, 20 

Sodium nitrate, ^N, 43 

Sodium nitrite, |N, 37 

Sodium phosphate, secondary, ^N, 55 

Sodium siHcate, sp. gr. 1.38 

Sodium thiosulphate, jN, 31 , 

Stannic chloride, 50 c.c.+500 c.c. HCl 

Stannous chloride, IN, 112+200 c.c. HCl+Sn 

Strontium chloride, jN, 33 

Sulphurous acid, sat. 

Zinc chloride, ^N, 17 

Zinc sulphate, ^N, 36 

DESK REAGENTS 

Ammonium hydroxide (dilute), 5N, 685 c.c. of NH4OH sp. gr. 0.90 
Hydrochloric acid (dilute), 5N, 388 c.c. of HCl, sp. gr. 1.20 
Nitric acid (dilute), 5N, 317 c.c. of HNO3, sp. gr. 1.42 
Sulphuric acid (cone), sp. gr. 1.84, 36N 
Sulphuric acid (dilute), 5N, 139 c.c. of HjSO^, sp. gr. 1.84 
Sodium hydroxide, 3N, 120 



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Library Bu'- au Ci! No 1137 



QD45,N65 




3 5002 00126 8247 



Norris, James Flack 

Laboratory exercises In inorganic chemis 



QHEMlSTFvY LfBRARt: 




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