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273 filD

XPERIMENTAL
CHEMISTRY

BY WALTER MONROE

J&N 61906

LIBRARY

OF THE

UNIVERSITY OF CALIFORNIA.

GIFT OF"

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EXPERIMENTAL CHEMISTRY

FOR

HIGH SCHOOL STUDENTS

WALTER MONROE,

Instructor in Chemistry in the Monrovia High School

MONROVIA, CALIFORNIA,

1899

COPYRIGHTED 1899

WALTER MONROE

Commercial Printing House
Los Angeles

PREFACE.

(HIS course of experiments has been arranged for
students in high schools, and is intended to lead
directly to the courses in the State University.
Since many of the principles of chemistry are founded
upon quantitative relations, a number of such experiments
has been introduced; but no exercise has been incorpor-
ated into the course that has not been successfully per-
formed. Special care has been taken in clearly describing
the experiments, so that pupils may not fail to obtain
good results. Considerable attention has also been given
by questions and suggestions to cause the student to think
carefully over observed phenomena.

The use of symbols has been avoided in the earlier
stages of the work, and qualitative analysis has been
entirely omitted. Very little of real value can be
accomplished in the way of qualitative analysis in so
short a time.

A number of the most important elements and their
compounds has been considered, and the experiments out-
lined are only those which may be performed by the
student during the year. The course may be modified,
whenever desirable, by supplementary directions prepared
by the instructor. The writer is of the opinion that
a few elements and their most important compounds
studied carefully, are of much more value than a great
many studied hurriedly.

The exercises are intended to develop in the pupil a

140475

4 PREFACE

* 'scientific habit of thought' ' and lead him to discover for
himself by a purely inductive method many of the funda-
mental principles of the science.

A course of recitations, and an occasional lecture
illustrated by experiments should accompany the labora
tory work. The note-book should be used in the labora-
tory only, so that it may always represent the pupil's own
thoughts.

Large "folders" with perforated paper that may be
removed make the best laboratory note-books. The
written sheets may be passed in for examination by the
instructor, and afterwards fastened in the folders and kept
for future reference.

Very few directions have been given to pupils, for
these are best left to the discretion of the teacher.

The writer is indebted to Prof. Rising of the State
University for many valuable suggestions. It is likewise
a pleasure to acknowledge the assistance received from
Mr. Sharwood, formerly of the University, who has so
kindly aided in the presentation of the book. Much of
the material has been adapted from his course on the
"General Laws of Chemical Action."

It was not the intention to have this course put into
print until next year; but the demand for the work by a
number of instructors has induced the author to have a
limited number printed for use this year. Next year the
book will appear in better form, and with it a text book
which is now in preparation and especially adapted to the
course. Further suggestions from the instructors in
chemistry at the University, and from teachers who may
use the book the coming year will be thankfully received.

WALTER MONROE.
Monrovia, Cal.y June 75,

CONTENTS.
PART I.

THE NON-METALS AND THEIR COMPOUNDS.

EXP. NO.

1 Physical and Chemical Changes.

2 Mechanical Mixture and Chemical Compound.
3-5 Oxygen, Preparation and Properties.

6-9 Hydrogen, Preparation and Properties.

10-13 Water.

14 Estimation of Water in Copper Sulphate.
( Volumetric Measurement of Gases.
{ Equivalent Weight of Zinc and Magnesium.

1 6 Nitrogen, Preparation and Properties.

17 Composition of the Atmosphere.

1 8 Ammonia, Preparation and Properties.

19 Nitrous Oxide, "
2 . Nitric Oxide,

21 Nitric Acid,

22 Hydrochloric Acid, Preparation and Properties.

23 Chlorine,
24-25 Carbon,

26 Carbon Dioxide,

27 Carbon Monoxide, " " "

28 Flame, Combustion, Heat, Light, etc.

29 Sulphur and Its Properties.

30 Sulphur Dioxide, Preparation and Properties.

31 Hydrogen Sulphide, " " "

32 Sulphuric Acid, " " "

33 Acids, Bases, Salts, etc.

34 Neutralization.

CONTENTS.

PART II.

THK METALS AND THEIR COMPOUNDS.

EXP. NO.

35 Potassium.

36 Potassium Compounds.

37 Sodium.

38 Sodium Compounds.

39 Ammonium Compounds.

40 Calcium Compounds.

41 Zinc.

42 Zinc Compounds.

43-44 Oxidation and Reduction.

45 Iron.

46 Iron Compounds.

47 Copper.

48 Equivalent Weight of Copper.

49 Lead and Its Compounds.
50—51 Silver and Its Compounds.

52 Mercury and Its Compounds.

53 Action of Acids on the Metals.

54 Chlorides.

55 Oxides.

56 Hydroxides.

57 Sulphides.

58 Nitrates.

59 Sulphates.

60 Carbonates.

6 1 Crystallization.

62 Determination of Atomic Weights.

CONTENTS

APPENDIX.

Table of Principal Elements, Atomic Weights and

Symbols.
Complete List of Chemicals and Apparatus Required

for the Course.
Individual Apparatus.
Preparation of Solutions.
Addresses of Supply Houses.

GENERAL DIRECTIONS.

Students should provide themselves with a towel,
some pieces of cloths, and an apron — one with sleeves is
the best protection.

Keep your desk in as good order as possible; and
before leaving the laboratory, see that your apparatus is
clean and the table dry. Never use dirty apparatus.
Glass vessels should be cleaned as soon after use as possi-
ble, rinsed and left to drain and dry — not wiped unless
required for immediate use. No one can do good work
in chemistry who neglects these essential points.

Do not put down the stopper when using a re-agent
bottle, but hold it between the fingers. Any excess of a
re-agent must not be poured back into the re-agent bottle.
Always see that your flasks and tubes are dry on the out-
side before heating, and apply the heat gently at first.

Use great care in handling the strong acids and alka-
lies. Forceps should be used in handling phosphorus,
potassium, and sodium. Do not perform experiments
outside of the regular work without permission as serious
accidents sometimes occur

Before commencing an experiment, read the directions
carefully, and ascertain the object of it. Always ask for
information when you are in doubt. The quality of the
work done by each student and his ability to draw cor-
rect conclusions from observed facts, is of far more value
than the bare number of experiments performed.

Watch carefully for changes that may take place, and
write in your note-book in the laboratory an account of
the experiment as you have performed it, (a sketch of the

10 GENERAL DIRECTIONS

apparatus will quite often save time in description) and
state any conclusions you have been able to draw. Be
careful to note everything the experiment shows, and
answer all questions fully.

Write the description neatly, accurately, and in good
English. Do not use abbreviations for chemicals, but
write the name in full. When your instructor performs
an experiment, write an account of it just as if you had
performed it yourself.

PRELIMINARY WORK.

Make yourself thoroughly familiar with the practical
use of the metrical standards, and learn to estimate vol-
umes in cubic centimeters.

Graduate a test tube to measure 5 and 10CC by pour-
ing water from a graduate into the test tube. Mark the
level of the water by a small file scratch. For most
experiments, not quantitative, the amounts to be taken
are merely approximate and may be weighed on paper on
a rough balance. Volumes of liquids can be measured
in graduated vessels.

Under the direction of your teacher, learn to do the
following well: — Bend glass tubing and round the ends,
draw out tubing, close the ends of a glass tube, make
bulbs, bore corks and fit tubes in them. Always have
your apparatus neatly arranged. Ask your teacher to
assist you in fitting up apparatus until you become
familiar with the work.

PART I.

EXPERIMENTS.

THE NON-METALS AND THEIR COMPOUNDS.

PHYSICAL CHANGE AND CHEMICAL CHANGE.

1. Changes which do not alter the nature or compo-
sition of substances are known as physical changes. Ice
may change to water and water to steam, but the compo-
sition of the substance is unchanged and the changes
are physical.

Changes in the nature and composition of substances
as well as a change in properties are chemical changes.
The burning of wood and the rusting of iron are exam-
ples of chemical change.

(a) Dissolve 2 grams of sugar in about 5CC of water.
Taste the solution. What change has taken place?
Always give reasons for your answers. Add slowly 5CC
of concentrated sulphuric acid. Note any change of
color, or temperature of the tube. Is there any sugar
left? What kind of a change has taken place?

(b) Dissolve a gram of salt in 10CC of water in an
evaporating dish. Evaporate todryness and examine the
residue. What is its color and taste? Name the change
that took place.

(c) Hold a piece of platinum wire in the flame of
your lamp, remove and allow to cool. Repeat with a
piece of magnesium ribbon; hold the magnesium with a
pair of pincers. What change takes place in each
case? Give reasons. Mention other changes similar to
these with which you are familiar.

12 EXPERIMENTAL CHEMISTRY

MECHANICAL MIXTURE AND CHEMICAL COMPOUND.

2. (a) Mix intimately about one gram of flowers of
sulphur with an equal weight of fine iron filings. Exam-
ine the mixture. Can you distinguish the particles of
iron and those of sulphur? Examine again with a lens.

(b) Spread some of the mixture on a piece of paper,
and pass a magnet over it. What is the effect? The oper-
ation may be repeated a number of times if necessary.

(c) Put a small quantity of the mixture in a test
tube and cover well with carbon disulphide. (Carbon
disulphide is a very volatile liquid and should never be
heated.) Shake vigorously several times, and notice
whether any change takes place. Filter the substance
(ask for directions); catch the filtrate upon a watch
glass, and allow it to evaporate. What is left on the filter
paper? Allow it to become dry and test with a magnet.
What is the color of the substance on the watch glass?
What is it? Has the separation of the iron and sulphur
been accomplished by physical or chemical processes?

(d) Take all but a small quantity of the mixture
and place it in a small ignition tube. Heat strongly for
some time and observe closely the changes that take
place. After the action is over and the tube is cooled,
break it and catch the contents on a piece of paper.
Compare the substance with the original mixture.
Examine it with the lens, magnet, and carbon disulphide.
What kind of a change has taken place?

EXPERIMENTAL CHEMISTRY 13

•

OXYGEN.

3. (a) Mix intimately, without pulverizing, 25
grams of coarsely powdered potassium chlorate with an
equal weight of pure coarsely powdered manganese diox-
ide. Explosions sometimes occur by using impure man-
ganese dioxide. Test the mixture by heating a little in
a test tube. If no violent action takes place, put the
mixture in a retort (copper, if obtainable), or a flask, and
insert a delivery tube. Heat the retort and fill several
receivers (cylinders or wide-mouthed bottles) with the
gas by the displacement of water.

(b) Put a lighted splinter into a small receiver;
remove it and put it in again while glowing. Does oxy-
gen support combustion? Inhale a little of the gas.
What can you say of its odor, taste and appearance?

(c) Put some sulphur into a deflagrating-spoon (a
piece of crayon hollowed out and attached to a copper
wire will answer ) and hold it in one of the receivers.
Into another receiver introduce a piece of charcoal; and
into a third, a small piece of phosphorus. L,et them
stand a few minutes, and note any changes that take
place.

(d) Repeat with the same elements, but ignite them
before introducing into the receivers. How does oxygen
act upon substances at ordinary temperatures and at high
temperatures?

(e) Fill three test tubes with oxygen and put a piece
of dry phosphorus into one, a piece of roll sulphur into
another, and a piece of charcoal into the third. Cork
the tubes loosely, and place them in a beaker of cold
water. Heat the water slowly, or pour boiling water
slowly into the cold. Observe the changes that take

14 EXPERIMENTAL CHEMISTRY

place. Which element has the lowest kindling tempera-
ture, and which the highest?

4. It has been found by numerous experiments that
whether a substance burns in pure oxygen or in the air,
the result is generally the same

(a) Ask for directions for using the delicate balance,
and weigh accurately a porcelain crucible which has been
previously heated and allowed to cool, or which has been
left in a desiccator for some time. Why? Ask if you do
not know.

(b) Take about one-half gram of magnesium rib-
bon, weigh accurately and place it in the crucible. Place
the crucible on a triangle over a flame and heat it until
the magnesium is converted into a powder, and then place
it in a desiccator and allow ,it to cool.

(c) Weigh the crucible and contents accurately.
How has the magnesium changed in weight? Note any
change in appearance, conduct, etc., of the magnesium,

1 and state what the experiment shows.

Lead may be used in this experiment instead of mag-
nesium. Use about a gram; but before weighing, it
should be brightened with sandpaper, or have the dark
part cut off with a sharp knife. After the lead is melted,
it should be stirred continually with a stout iron wire
until it ceases to be in a liquid form.

5. (a) Put about one gram of mercuric oxide in a hard
glass tube, and insert a delivery tube. Heat the tube to
redness if necessary and collect the gas in a test tube or
a wide-mouthed bottle over water. Test the gas col-
lected. How do you know what it is? Note any change
that has taken place in the mercuric oxide.

(b) Repeat (a) using about four grams of manganese
dioxide instead of the mercuric oxide.

EXPERIMENTAL CHEMISTRY 15

HYDROGEN.

6. (a) Place 15 to 20 grams of granulated zinc in
a flask or Wolff's bottle; attach a delivery tube, and
through a second opening insert a thistle or funnel tube.
Add about 100CC of dilute hydrochloric acid or sulphuric
acid. What is the effect? Allow time for the air to be
expelled from the generator and tube, and then fill several
cylinders or bottles with gas by the displacement of water.
Fill a test tube with the gas, invert it and bring a lighted
match to its mouth. What happens? I/ight a candle or
splinter and introduce it into one of the vessels of gas,
having the mouth of the receiver downward. Does
hydrogen behave like oxygen?

(b) Take two vessels containing the gas; place one '
with its mouth upward, the other with its mouth down-
ward. Uncover both vessels and after a few moments
introduce a lighted taper into each. Bxplain the differ-
ence.

(c) Take a cylinder or bottle containing air and pour
the gas from one of the receivers into it. How must the
vessels be held? How can you tell that the gas has been
poured from one vessel to the other? What does the
experiment show?

(d) Collect some of the gas from the generator after
passing it through a bottle containing a solution of potas-
aium permanganate. Has hydrogen any odor, taste or
color? What is the use of the potassium permanganate?

7. (a) Arrange an apparatus as in the previous
experiment, but instead of the delivery tube, attach a U-
shaped tube filled with calcium chloride. Roll up a piece
of platinum foil so as to make a small tube and seal it to
one end of a glass tube about 15cm long. Connect the

16 EXPERIMENTAL CHEMISTRY

other end of the glass tube to the calcium chloride tube
so that the platinum point will be almost in a vertical
position.

(b) Add dilute hydrochloric acid and after action
has been going on long enough to expel all air, light the
hydrogen escaping from the jet. (It is well to test the
gas before lighting. How? Ask if you do not know.
Wrap a towel around the flask to prevent any serious
accident in case of an explosion.) Place a bell-jar or
other wide-mouthed vessel over the flame at an angle, and
examine after a few minutes. Explain the chemical
change that takes place when hydrogen burns in air.
What is the use of the calcium chloride tube? Why is
the platinum tube used in the end of the glass tube?
What is the appearance of the hydrogen flame?

8. (a) Take an iron gas or water pipe about 50cm
long and 20mm internal diameter, and fill the middle por-
tion with iron turnings. Fasten the pipe in a horizontal
position by means of a clamp attached near one end.
Partly fill a flask with water, place it on a ring stand,
and connect with one end of the iron pipe. Insert a
delivery tube in the other end of the pipe and arrange to
collect gas over water. Heat the iron pipe to redness in
the center, and boil the water in the flask.

(b) Collect several samples of the gas and test them.
What is the gas? Give reasons. After the tube has
cooled, examine the iron turnings. Explain the changes
that have taken place during the experiment.

9. (a) Fit a porous earthen cup to a glass tube about
25 or 30cm long, having a small internal diameter. (In-
stead of the porous cup, a plaster of Paris plug may be
inserted in the end of the glass tube which may, in this
case, be shorter and about 2cm in internal diameter.)

EXPERIMENTAL CHEMISTRY 17

Put the other end of the tube through a cork in one neck
of a Wolff's bottle containing some water colored with
litmus, or any other coloring matter. Do not pass the
tube below the surface of the water. Through the other
neck of the bottle pass a small glass tube below the sur-
face of the water. The upper end of this tube should
extend above the bottle about 10cm, be bent outward
slightly, and drawn out at the end to a small opening.

(b) Fill a bell-jar with dry hydrogen by the dis-
placement of air, and bring it over the porous cup. Note
what happens and explain all that you have seen. Does
the bell-jar filled with air produce the same effect?

WATER.

10. Procure the apparatus for the electrolysis of
water, and pour water into the glass vessel until the plat-
inum electrodes are below the surface. Take two large
test tubes of the same size, and fill them with water con-
taining a small amount of sulphuric acid. " Invert the
test tubes and place them over the platinum electrodes;
then pour into the glass vessel about one-twelfth as much
strong sulphuric acid as it contains water. Water alone
will not conduct a current, but when sulphuric acid is
added it acquires the power to convey the current. Con-
nect a battery of two or three cells in series to the copper
wires in connection with the platinum electrodes. Notice
the electrodes. When the water has been completely
forced out of one tube, how full is the other? Test the
gas thus collected. How do you know what the gasses
are? Which gas collects the faster? In what proportion
do the two gasses unite to form water?

11. From the knowledge you already have of the
behavior of water, state its boiling and freezing points.
At what temperature does it weigh most? What is the

18 EXPERIMENTAL CHEMISTRY

weight of one cubic centimeter at that temperature?
Is the temperature at which water boils affected by the
pressure under which it boils?

(a) Fill a flask of about one liter capacity one-third
full of water from the hydrant and connect it with a con-
denser. Distil about 50 to 100CC of the water and com-
pare it with the hydrant water as to taste, odor and color.

(b) Evaporate to dry ness — in a watch glass set over
a beaker of water, or in an evaporating dish — about 5CC of
the distilled water.

(c) Treat in a similar manner an equal amount of
water from the hydrant. Compare results. Could sea
water be made fit to drink by distillation?

12. (a) Heat about one gram of copper sulphate
crystals in a small test tube. Observe what takes place.

(b) Repeat (a) using one gram of alum.

(d) Place a few crystals of clear sodium sulphate
(Glauber's salt) on a watch glass and leave exposed to the
air until the next day.

(e) Treat similarly a few pieces of calcium chloride.
Compare with (d).

13. Nearly all substances dissolve more or less in
water. Platinum, however, is insoluble. Some liquids
are soluble in water in all proportions, some are only par-
tially soluble, while liquids like oils are very slightly sol-
uble. The solubility of a substance depends upon the
temperature, but it is always definite for a given tempera-
ture. The solubility of solids generally increases with
the temperature; but the solubility of gases decreases
with an increase of temperature.

(a) Put one gram of powdered potassium chlorate

EXPERIMENTAL CHEMISTRY 19

(weighed on a balance) in a test tube, add three or four
cubic centimeters of water and heat to boiling. Exam-
ine the solution; allow it to cool and examine again.

(b) Add about 15CC of water and heat till completely
clear; cool and compare with (a).

(c) Put 5CC of cold water in a test tube and add a
few drops of carbon bisulphide. Shake well and let it
stand a few minutes. Examine.

(d) Pour 5CC of water in a test tube and add a few
drops of ether. Shake, let stand and examine; add a
little more ether, and again examine; finally add 4 or 5CC
more and examine again.

(e) Repeat (d) using alcohol instead of ether.

ESTIMATION OF WATER IN COPPER SULPHATE.

14. Dry and weigh a small porcelain crucible. (See
directions, Exp. 4.) Introduce into it about a gram of
bright crystals of pure copper sulphate, slightly powdered
and accurately weighed. Heat in an oven (or on asbestos
cloth) for at least an hour at a temperature between 100°
and 110° C. Cool in a desiccator and weigh. Heat the
crucible and contents again under similar conditions for
about half an hour, cool and weigh as before. Continue
the heating until you find the weight constant. Now heat
over a flame almost to a low red heat for about fifteen
minutes; cool in a desiccator and weigh. Calculate the
amount of water expelled below 110° C.; also the amount
expelled between 110° and low red heat. What fraction is
this last quantity of the whole amount of water expelled?
Calculate the percentage of water in the blue crystals
taken.

(b) Put the dry copper sulphate in a test tube and
add a few drops of water. Note any change in appear-
ance and temperature.

20 EXPERIMENTAL CHEMISTRY

VOLUMETRIC MEASUREMENT OF GASES.

If has been found by experiment that the volume of a
gas varies with the temperature and pressure as follows:

(a) The volume of a given quantity of dry gas
increases ^ part of its volume at zero Centigrade for
each rise in temperature of one degree Centigrade. (Law
of Charles, also law of Gay-Lussac.) Taking — 273 as
absolute zero and 273 plus the temperature Centigrade
the absolute temperature, the volume varies as the abso-
lute temperature divided by the pressure.

(b) The volume of a confined mass of gas is inversely
proportional to the pressure to which it is exposed. The
volume times the pressure is equal to a constant (Law of
Boyle, also Marriotte's Law). The atmospheric pressure
varies with the height above sea-level, but it approxi-
mates to that of a column of mercury 760mm high which
is taken as the standard.

From the above laws we get the following formula
which may be used to determine \vhat volume a gas,
measured at one temperature and pressure, would occupy
at another:

Let V represent the volume of a given mass of gas at
0° C. and 760mm pressure.

Its volume at 1° C. and 760mm press ure=V-f^-3V.

Its volume at 2° C. and760mm pressure=V-j-~V.

Its volume at t° C. and 760mm pressure=V+ ~ V. =:

273+tx
Vl 273 )'

Its volume at t° C. andpmmpressure=V(^)-^.

The volume of a gas in contact with water is increased
owing to pressure of water vapor. Correction may be
made by subtracting from the given barometric pressure
the tension of aqueous vapor at the given temperature;

EXPERIMENTAL CHEMISTRY 21

that is, substitute for p. in the formula the observed
barometric reading ^ninus the tension of water vapor at
the temperature / centigrade.

TENSION OF WATER VAPOR FOR DIFFERENT TEMPER-
ATURES.
Temp. C. Tension of vapor Temp. C. Tension of vapor

0° 4.6mm 21° t, 18.49mm

4° r.... 6.0 " 22° ....19.69"

8° 8.0 " 23° 20.88"

10° 9.16" 24° 22.18"

11° 9.8 " 25° 23.55"

12° 10.5 " 26° 24.98"

13° 11.6 " 27° 26.5 "

14° 12.0 " 28° 28.1 "

15° 12.7 " 29° 29.78"

16° 13.5 " 30° ....31.54"

17° 14.4 " 31° 33.4 "

18° 15.35" 32° 35.35"

19° 16.3 " 33° 37.4 "

20° 17.3 "

PROBLEMS.

(1) 18.2 liters of gas are measured at 0° C. and
760mm pressure; what would be the volume if measured
at 17° C. and 950mm pressure?

(2) A quantity of oxygen occupies a volume of
121 6CC at 15° C. and 750mm pressure; what is its volume
under standard conditions?

(3) 100 cubic centimeters of hydrogen are measured
off at 27° C. and 950mm pressure; what volume would it
occupy at 17° C. and 750mm pressure?

(4) 200CC of a gas are measured over water at 20° C'

22 EXPERIMENTAL CHEMISTRY

and 760mm pressure; what volume would the gas occupy
measured dry at 0° C. and 760mm pressure?

15. (a) Clean a small piece of zinc, approximately
.05 grams, and weigh accurately. Place it in a beaker
on a ring stand, and introduce a funnel with the mouth
covering the zinc — stem upward, but not long enough to
reach the mouth of the beaker. Add water until the
funnel is completely covered, and insert a test tube, filled
with water, over the stem; fasten securely with a wire
holder or with a clamp. Then with a pipette or a funnel
pour 10CC of concentrated sulphuric acid to the bottom of
the beaker, and add a few drops of copper sulphate solu-
tion. If the action is very slow, warm slightly. When
the zinc has all been dissolved remove the clamp, lower
the tube into the beaker, and set the beaker into a basin
of cold water and leave until cold. Raise or lower the
tube till the level of the liquid is the same inside of the
tube as outside. Why? Cover with the thumb while so
adjusted, remove and invert. Let the gas escape, and
measure the volume of the gas. It is equal to the vol-
ume of liquid now required to fill the tube. Fill a Bur-
ette with water; take the reading, and then run water
from it into the test tube until it is exactly as full as in
the beginning when inverted over the funnel. The dif-
ference in reading gives the volume of gas. (The
Burette is read by noting the level of the lower meniscus
of surface.) Read the thermometer and the barometer to
get the temperature and pressure, and also find the water
vapor tension at this temperature. From these data cal-
culate the volume dry at 0° C. and 760mm pressure.
Hence calculate the volume of gas liberated by one gram

EXPERIMENTAL CHEMISTRY 23

of zinc. One liter (1000CC) of dry hydrogen under stand-
ard conditions has been found by careful experiment to
weigh 0.0896 grams nearly. Therefore, calculate the
equivalent of zinc; that is, the number of grams of zinc
required to yield one gram of hydrogen from sulphuric
acid. Compare your results with those of other students.

(b) Repeat the above experiment using hydrochloric
acid and a different weight of zinc. Is the final result
the same?

(c) Repeat the experiment using .02 grams of mag-
nesium, and sulphuric acid. What is the equivalent
weight of magnesium?

NITROGEN.

16. Hollow out one end of a piece of crayon about
three cm long and attach it to a piece of wire. Fasten
the wire to the shelf of a pneumatic trough, having water
3 or 4cm above the trough. Have the crayon in an up-
right position and of such a height that it will reach
about half way to the bottom of a large wide-mouthed
bottle when the bottle is inverted over it. With a pair of
forceps, put a piece of phoshorus, the size of a pea, into
the crayon cup. Ignite the phosphorus, and quickly
invert the wide-mouthed bottle over it. Allow it to stand
until the white fumes disappear. What are these fumes
and what becomes of them? Cover the mouth of the
bottle with a glass plate, and turn it mouth upward.
Note the odor and color of the gas. Test the effect of
the gas on a burning splinter, phosphorus, and sulphur,
introduced one after the other. Does nitrogen support
combustion? Where did it come from?

THE ATMOSPHERE.

The air is a mechanical mixture, of which over

24 EXPERIMENTAL CHEMISTRY

ninety-nine per cent, is oxygen and nitrogen.* Water
vapor and carbon dioxide are always found present in
varying proportions. Besides these, ammonia and some
other gases are usually found in small quantities.

17. (a) Take a piece of glass tubing about 20cm
long, closed at one end, and from 10 to 15mm in diameter.
Bend the closed end to a right angle making the bend
about 5cm long. Slip two small rubber rings over the
open end, and drop a piece of phosphorus the size of a
pea into the bent portion. Dip the open end into a
beaker of water, and suck air out with a rubber tube
until the wrater rises about 3cm in the tube. Adjust the
tube until the water is at the same level inside and out-
side. Why? Mark the mean level by the lower rubber ring,
and fasten the tube securely with a clamp. Warm the
closed end containing the phosphorus, but not enough to
drive the expanded air out of the tube. I^et the tube
cool after the phosphorus has entirely ceased to act. Pour
water over the tube; and after the white fumes have dis-
appeared, adjust the upper rubber ring to indicate the
volume of the gas left. Is the volume of the gas greater
or less than that of the original air? Why?

(b) Fill the tube, first to one rubber ring and then
to the other, with water from a Burette, and note the vol-
ume in each case. Calculate the percentage composition
of air by volume, assuming it to be composed entirely of
oxygen and nitrogen. What previous experiment showed
the presence of moisture in the atmosphere?

AMMONIA.

18. (a) Add a few drops of caustic soda solution

* In the year 1894; Raleigh and Ramsay discovered that about 1 per cent,
of what was supposed to be nitrogen was really a new element which they
called Argon, meaning inert substance. Since then other elements have been
found in very small quantities.

EXPERIMENTAL CHEMISTRY 25

to a little ammonium chloride in a test tube, warm and
notice the odor of the gas given off. Repeat the opera-
tion using caustic potash instead of caustic soda.

(b) Mix in a mortar about a gram of quicklime and
a gram of ammonium chloride, and note the odor as before.
What kind of an odor has ammonia? By what other name
is it commonly known? Write equations representing the
reactions that took place in each case. (The student is
supposed to have had, by this time, some practice in
writing equations.)

(c) Mix thoroughly about 25 grams of slaked lime
with an equal weight of ammonium chloride. Place the
mixture in a flask and attach a delivery tube. Heat the
flask gently in a sand bath or over asbestos, and collect
three cylinders of the gas by the displacement of air;
cover securely with glass plates. The mouths of the
receivers should be kept downward throughout the
experiment unless otherwise directed, as the gas is very
much lighter than air. Try to light the gas as it escapes
from the generator. Will ammonia burn? Allow the
current of gas from the generator to pass into a beaker of
water while the cylinders of gas are being examined.

(d) Note the odor and color of the gas. Into one of
the receivers put a piece each of wet blue, and red litmus
paper. What change takes place? Introduce a burning
candle or splinter into the same receiver. Does the gas
support combustion? Invert one of the cylinders with its
mouth upward, remove the glass cover, and quickly apply
a lighted match. Does the ammonia burn?

(e) Moisten the sides of a dry cylinder with concen-
trated hydrochloric acid, and place it over the third cyl-
inder of the gas collected, so that the mouths of the two
vessels are separated only by the glass cover over

26 EXPERIMENTAL CHEMISTRY

the receiver of gas. Remove the glass plate, and observe
carefully what takes place. What is formed?

(f) Examine the water into which the gas has been
passing. How does it affect litmus paper? Boil some of
the solution in a test tube. What gas is given off? How
does the solution compare with the laboratory reagent?
The ammonia gas should be passed through a cylinder of
quicklime or caustic soda to obtain it perfectly dry. Why
could not calcium chloride be used?

NITROUS OXIDE.

19. (a) Place 15 to 20 grams of ammonium nitrate
in a retort or generating flask, and arrange the apparatus
as in the preparation of oxygen. Heat gently, and after
the air has been expelled, collect three receivers of the gas
over warm water. Nitrous oxide is somewhat soluble in
cold wrater. If the flask is heated too fast, some nitric
oxide will be formed which takes oxygen from the air to
form nitrogen peroxide, No.,. Nitrous oxide may contain
some chlorine also. In order to obtain the gas pure, it is
usually passed through two Wolff's bottles, — the first
containing a warm solution of ferrous sulphate which
removes the nitric oxide; the other, containing a warm
solution of caustic potash to remove the chlorine.

(b) Note the odor, color, and taste of the gas. Intro-
duce a lighted splinter into one of the receivers. Does
the gas support combustion? Put a small piece of phos-
phorus in a deflagrating spoon,. ignite, and introduce into
a second receiver of the gas. Introduce burning sulphur
into the third receiver. What gas does nitrogen monox-
ide resemble?

EXPERIMENTAL CHEMISTRY 27

NITRIC OXIDE.

20. (a) Arrange a generating flask, and place in it
a quantity of metallic copper. J*ust cover the copper with
water, and then slowly pour concentrated nitric acid
through the funnel tube until action begins. What is the
color of the gas in the flask at first? What is the color
after action has continued for a short time? Collect four
cylinders of the gas over water, and observe the color.
Allow the nitric oxide from the delivery tube to escape
into the air and notice the change. The gas unites with
the oxygen of the air to form nitrogen peroxide. Explain
the appearance of the colored gas in the flask at the begin-
ning of the experiment.

(b) Introduce a lighted splinter into one of the cyl-
inders of gas. Does the gas support combustion? Does
it burn?

(c) Put a small piece of phosphorus into a deflagrat-
ing spoon, ignite, and lower into a cylinder of gas.
What is the effect?

(d) Repeat (c) using sulphur instead of phosphorus.

NITRIC ACID. ,

21. Place in a retort 30 grams of sodium nitrate or
potassium nitrate, and about 20CC concentrated sulphuric
acid. Pass the end of the retort into the neck of a flask,
and arrange so that the flask may be surrounded with ice,
or so that a stream of water may continually flow over it,
(a large test tube may be used instead of the flask. It
can be kept cool by immersing in cold water. ) Heat the
retort gently until the operation is finished. In what
form is the nitric acid? What is its odor and color?
When pure it is colorless, but it usually contains some of
the oxides of nitrogen which give it a yellowish color.

28 EXPERIMENTAL CHEMISTRY

(b) Add a small amount of nitric acid to 4 or 5CC of
indigo solution in a test tube, and warm.

(c) Put about 20CC of dilute nitric acid in an evapor-
ating dish, and drop into it some quill clippings or small
pieces of silk. Evaporate. What color is imparted to
animal substances?

(d) Test the action of nitric acid, both the concen-
trated and the dilute, (one volume of concentrated acid to
three of water) upon the following substances: copper,
zinc, tin, lead, iron and magnesium. Put about 5CC of acid
into a test tube and then drop in a small piece of the metal
to be tested. It is well to arrange as many test tubes as
there are metals, and test all at the same time. Do all
the metals dissolve ? Which one seems to dissolve most
readily? In which acid, concentrated or dilute, is action
most violent?

(e) Put 2CC of concentrated nitric acid into a test tube
and drop into it a small piece of platinum. Warm in a
flame, and notice carefully for any change.

(f) Repeat (e) using 3CC of concentrated hydrochloric
acid instead of the nitric.

(g) Unite the two acids used in (e) and (f). What
is the effect? A combination of the two acids form Aqua
Regia. The odor produced is that of chlorine. The nas-
cent chlorine unites with the platinum and platinum chlor-
ide is produced. If gold is used, auric chloride is formed.

HYDROCHLORIC ACID.

22 (a) Put 50 grams of common salt in a flask and
attach a delivery tube. Insert a thistle, or funnel tube,
and add about 75CC of dilute sulphuric acid. Connect the
delivery tube to a series of two or three Wolff's bottles
about half full of water. Heat the flask gently over asbes-

EXPERIMENTAL CHEMISTRY 29

tos or in a sand bath. When the air has been displaced,
the gas will be absorbed as soon as it comes in contact
with the water. Write the equation which shows the
reaction that takes place. After the gas has passed for
fifteen to twenty minutes, disconnect the flask from the
Wolff's bottles. What do you notice? Blow your breath
on the escaping gas. What effect has this? Why? Try
to light the gas at the end of the tube. Will it burn?

(b) Fill two wide-mouthed bottles with hydrochloric
acid gas by extending the delivery tube to the bottom of
the bottle which should be loosely covered with a glass
plate or a piece of cardboard. Note the color of the gas.
Is it transparent? Drop a moist piece of blue litmus paper
into one of the bottles. What is the nature of the gas?
Insert a burning splinter or a candle into the bottle.
Does the gas support combustion?

(c) Take the second bottle of gas, hold the glass cover
securely, and invert the bottle bringing the mouth under
water in a beaker or evaporating dish. Remove the glass
plate. What happens? Explain. Test the water with
blue litmus. Taste the water. Why cannot this gas be
collected over water like hydrogen and oxygen? Why
was it not necessary to invert the bottles when collecting
the gas?

(d) Examine the liquid in the first Wolff's bottle.
How does it affect litmus paper? Put about a gram of
granulated zinc into a test tube and add about 10CC of the
liquid. What is the effect? What gas is given off?
Compare the liquid with the laboratory hydrochloric acid.
Are they the same?

Hydrobromic acid and hydriodic acid are very much
like hydrochloric acid, and are made in the same way.
Give the equations representing their preparation.

30 EXPERIMENTAL CHEMISTRY

CHLORINE.

23. This experiment should be performed in the hood
or out of doors. Be careful not to inhale the gas given off.

(a) Arrange a flask like the one used in the last
experiment, and put in it about 30 grams of manganese
dioxide. Pour through the funnel tube about 50CC of
hydrochloric acid — enough to cover the manganese diox-
ide completely. Heat the flask gently as in the last
experiment, and fill six cylinders or wide-mouthed bottles
with the gas by downward displacement of the air. As the
chlorine gas collects, the quantity can be noted by means
of the color. What is the color? Write the equation
representing the action. (Sulphuric acid could be used
instead of hydrochloric acid if an equal weight of common
salt be added to the manganese dioxide. Write the
equation). After the bottles are full, try to light the gas
escaping from the generator. Will it burn? Allow the
escaping gas to run into a receiver of water while you are
testing the gas collected in the wide-wouthed bottles. Is
the gas soluble in water? What is the color of chlorine
water? Save the solution. Pour it into a bottle, cork,
and set in a place protected from bright light.

(b) 1. Into one of the vessels containing chlorine
put some pieces of red and blue litmus paper; a red rose;
a green leaf; a piece of cotton print; a piece of newspa-
per; and a piece of paper with some writing on it. The
substances used must be moist. Use two vessels of chlor-
ine if necessary.

2. Put some of the dry articles into another receiver
of chlorine. In a short time examine the substances in
the receivers, and observe the effect upon each.

3. I/ower a lighted candle or taper into a receiver of
chlorine. Does the gas support combustion?

EXPERIMENTAL CHEMISTRY 31

4. Warm a little oil of turpentine in an evaporating
dish, moisten a piece of filter paper or blotting paper with
it, and quickly introduce into a receiver of chlorine.
Observe what happens.

5. Sprinkle into another receiver of chlorine a little
finely powdered copper, antimony, or arsenic; and observe
the effect.

6. Arrange a hydrogen generator as in experiment
7. (Observe the caution). When the air has all been
expelled, light the hydrogen and bring the burning jet
into the remaining receiver of chlorine. Does it continue
to burn? What is formed?

(c) Put about 5CC of indigo solution into a test tube
and add 5CC of chlorine water. Do you notice any change?
Treat in the same way 5CC of potassium bichromate solu-
tion. Note — Teachers should explain, as far as possible,
the changes which take place in this experiment.

The preparation of bromine and iodine is similar to the
preparation of chlorine. The apparatus used is the same
as that used in the preparation of nitric acid.' The bro-
mine condenses in the receiver, and forms a reddish liquid.
Iodine forms a sublimate of a violet color. Write equa-
tions illustrating the reactions that take place when bro-
mine and iodine are made.

CARBON.

24. (a) Hold a plate or piece of porcelain in the
flame of a candle, or of a Bunsen burner with the air
supply shut off. Observe the color of the deposit. This
is an impure form of carbon. What name is given to
this particular form of the element? Hold the deposit in
the flame of an alcohol lamp, or in a Bunsen flame with
the valve at the base of the burner open. Does the
deposit disappear? Is carbon a combustible element?

32 EXPERIMENTAL CHEMISTRY

(b) Put some pieces of wood into a small Hessian cru-
cible and cover them with sand. Heat the crucible strongly,
and when smoking stops, cool, remove the contents, and
examine. This is impure carbon also. What name is
given to it? What became of the other products of the
wood? Name other forms of the element carbon.

(c) Put about 2 grams of sugar into a porcelain
evaporating dish and heat till the sugar is black. Sugar
is composed of the elements carbon, hydrogen, and
oxygen; when the sugar is heated, water passes off arid
the carbon is left. The carbon can be removed from the
evaporating dish with a strong solution of sodium
hydrate.

(d) Place a few grains of sugar into an evaporating
dish, and add a few drops of concentrated sulphuric acid.
What results? Try starch in the same way.

(e) Test the solubility of carbon (use charcoal) in
water, acids, alkalies and alcohol.

(f) Heat some powdered charcoal on a piece of plat-
inum foil. What occurs?

25. (a) Arrange a bone-black filter by placing a
paper filter into a funnel and adding bone-black. Filter
a dilute solution of indigo and examine the filtrate. Is
there any change in the appearance of the solution? Fil-
ter in the same way a dilute solution of litmus.

(b) The same effect may be produced by putting
bone-black into the solution, boiling for a short time, and
then filtering through a paper filter. Try this with any
solution colored with animal or vegetable coloring matter;
as cochineal, indigo, litmus, iodine. Take about 50CC in
a beaker and place in it about one or two grams of bone-
black; heat for a short time, shake well and filter.

EXPERIMENTAL CHEMISTRY 33

into a beaker, and note the odor. Add about a gram of
bone-black; warm the solution a few moments, shake well
and filter. Note the odor again.

(d) Collect over mercury in a test tube some ammonia
gas, made by heating a strong solution of ammonia and
passing it over quick-lime to dry it. Heat a piece of
charcoal; and, without removing the inverted receiver of
gas, press it through the mercury into the mouth of the
receiver. What is the result? Why was the charcoal
heated before introducing it into the test tube?

(e) Mix together 3 or 4 grams of powdered copper
oxide, CuO, and about one gram of powdered charcoal;
heat strongly for some time in an ignition tube. What
remains in the tube? What gas is given off? Write the
equation which shows the reaction that takes place. Is
carbon an oxidizing or a reducing agent? (The gas given
off may be collected by attaching a delivery tube before
heating the mixture.) After the tube has become cool,
pour in a little strong nitric acid. What takes place?
State some of the uses of carbon which have been illus-
trated by the experiments you have performed.

CARBON DIOXIDK.

26, (a) Put some small pieces of marble (calcium
carbonate), 15 or 20 grams, into a generating flask, and
pour dilute hydrochloric acid on them through the funnel
tube. After action has continued for a time, pass the gas
into a test tube of lime-water. Is there any change in
appearance? Note the odor and color of the gas. Try
to ignite it as it escapes from the delivery tube. Will
carbon dioxide burn?

(b) Collect two or three cylinders (or bottles) full of
the gas by downward displacement. Into one of the cyl-

34 EXPERIMENTAL CHEMISTRY

inders of gas introduce successively a lighted candle or
burning taper, and a small piece of phosphorus in a
deflagrating spoon. Pour a cylinder of the gas, as you
would water, over a burning splinter. Does the gas sup-
port combustion? Is it heavier than air?

(c) Pass the gas from the generator into a test tube
containing distilled water; after a short time, taste the
water. Test it with litmus paper.

(d) Pass the gas for about ten minutes into a test tube
containing a solution of caustic potash; then add hydro-
chloric acid to the solution. What gas is given off?
Write equations representing what has taken place.

(e) Pass the gas into clear lime-water as in the
beginning of the experiment, but allow the action to con-
tinue for a much longer period. Remove the delivery
tube and heat the solution. State the changes that you
have observed; and if you do not understand all of them,
ask for an explanation.

(f) Blow your breath through a clear solution of
lime-water by means of a piece of glass tubing. What is
formed?

CARBON MONOXIDK.

27. (a) Put about ten grams of oxalic acid crystals,
H,C2O4, into a flask; add about 50CC of concentrated sul-
phuric acid, and connect the flask with a Wolff's bottle
containing a solution of caustic soda. Attach a delivery
tube to the other neck of the bottle.

(b) Heat the flask gently and collect some of the gas
in a test tube over water. Avoid inhaling the gas as it is
poisonous. Observe its color. Try to set fire to the gas.
Will it burn, or is it a supporter of combustion? Sul-
phuric acid has great affinity for water, and so it with-
draws hydrogen and oxygen from the oxalic acid in the

EXPERIMENTAL CHEMISTRY 35

right proportion to form water. The carbon and oxygen
combine to form carbon dioxide and carbon monoxide.
What becomes of the carbon dioxide?

(d) Put some copper oxide into a small hard glass
tube open at both ends. Support the tube in a horizontal
position and connect one end with the delivery tube from
the caustic soda solution. Heat the copper oxide and
pass the carbon monoxide over it for a short time.
Explain the changes that take place. Is carbon monox-
ide an oxidizing or a reducing agent?

Problem: — How much carbon monoxide is given off
by the decomposition of 10 grams of oxalic acid crystals?

FLAME.

28. (a) Place a lighted candle in front of some dark
object and examine the flame closely. Notice its shape,
and see how many parts you can distinguish. Make
drawings.

(b) Hold a plain white card, or piece of smooth white
pine, horizontally in the flame of a candle so that it nearly
touches the wick. Remove it before it becomes ignited
and hold the other side in a vertical position against the
wick. Repeat the experiment until you obtain a vertical
section and a transverse section well outlined on the card.
Make drawings of each section. Does combustion take
place in all parts of the flame?

*(c) Light the gas of a Bunsen burner and examine
the flame as you did the candle flame under (a). Has
the flame the same number of parts? Obtain several
small pieces of soft pine, and hold them successively in a

* Omit sections (c), (d) and (e) if Bunsen burners are not used in the la

36 EXPERIMENTAL CHEMISTRY

horizontal position in the flame and try to determine in
what part the combustion is most complete.

(d) Insert a glass tube into the inner cone of the
Bunsen flame; slant the tube upward, and apply a lighted
match to the upper end. Is a flame produced? What is
the effect of placing the glass tube in other parts of the
flame?

(e) Bring a piece of wire gauze down upon the flame
of a Bunsen burner. What is the effect? Is there any
change produced after holding the gauze in the same position
for some time? Turn off the gas, then turn it on again
without lighting; hold the gauze about 5cm above the
burner and see if you can light the gas above it. Explain
all the phenomena observed. This experiment represents
the principle upon which the miner's safety lamp (called
the Davy safety lamp) is constructed.

(f) Notice which portion of the Bunsen (or alcohol)
flame is the most luminous. Sprinkle some drv sawdust
or powdered charcoal in the flame. Is the illuminating
power of the flame greater or less than before? Does it
appear to remain constant? * Close the openings at the base
of the Bunsen burner. How does it change the light of
the flame? What are the openings for? What seems to
be the cause of light in the flame? State any additional
points you can in regard to the combustion, kindling
point, heat, and light of flame.

SULPHUR.

29. (a ) What is the color, taste and odor of sulphur?
Test the solubility of the element in water, alkalies, acids,
alcohol and carbon bisulphide.

(b) Put a few grams of sulphur in a Hessian or por-
celain crucible; heat until the sulphur is melted, and then

EXPERIMENTAL CHEMISTRY 37

allow it to cool slowly. As soon as a crust forms on the
surface of the sulphur, make a hole through it and pour
out the liquid. Examine carefully the crystals attached
to the crust and to the inside of the crucible. What
is the shape of the crystals?

(c) Dissolve one or two grams of powdered roll
sulphur in a little carbon bisulphide. Put the solution
in an evaporating dish or beaker, and allow it to evapo-
rate in the atmosphere. Examine the sulphur crystals.
Have they the same shape as those formed in (b)?

(d) Put in a retort a small amount of sulphur, and
pass the end of the retort into the neck of a flask. Heat
the sulphur to boiling, and after a time examine the
sulphur that has collected in the flask. It is called
flowers of sulphur.

(e) Put some sulphur into a test tube and heat
slowly till it melts. Observe the color of the liquid.
Continue the heating and notice carefully any changes that
take place. When it becomes a thin liquid, pour a little
into an evaporating dish containing water. Allow the
heating to go on until the liquid begins to boil, and then
pour some more into cold water.

Sulphur boils at about 450° and just above this point
it takes fire. Take the two deposits out of the water and
examine them; replace them in the water and after they
have remained for some time, examine again. Do you
notice any change?

What is formed when sulphur burns in oxygen?

SULPHUR DIOXIDE.

30. (a) Note the color and odor of the substance
formed when sulphur burns in air.

Put some pieces of sheet copper or copper filings into

38 EXPERIMENTAL CHEMISTRY

a generating flask, and add enough concentrated sulphuric
acid to cover them. Heat the flask gently and collect
two cylinders of the gas by downward displacement of the
air. What is the color of the gas? Has it an odor?
Avoid inhaling much of the gas.

(b) Introduce a lighted splinter into the mouth of
one of the receivers. Does sulphur dioxide support
combustion? Put some moist pieces of litmus paper in
a receiver of the gas. What is the effect?

(c) Pass some of the gas into a beaker containing
water. Is the gas soluble in water? Kxamine the water
and test it with litmus. What are its properties?

(d) Put a gram or two of sulphur in a porcelain
crucible and place it under a tripod or other support upon
which are placed as many of the following things as may
be had: moist flowers, green leaves, a piece of white woolen
yarn, a ripe apricot (or peach) cut in halves, a piece of
moistened unbleached silk. Ignite the sulphur and place
a large bell-jar over it so as to enclose the articles on the
tripod. lyeave undisturbed until the fumes settle
and then examine. What change takes place in the
substan ces?

(e) A solution of sulphur dioxide in water forms
sulphurous acid, H2SO3. If left in contact with the air, it
combines with oxygen and is converted into sulphuric
acid, H2SO4. It has the power of taking oxygen from
other substances also. Is it an oxidizing or a reducing
agent? Add a little of the sulphurous acid obtained in (c )
to a solution of potassium permanganate or potassium
bichromate containing a little sulphuric acid. What
change is noticeable?

EXPERIMENTAL CHEMISTRY 39

V ••*•

HYDROGEN SULPHIDE.

31. (a) Put a few pieces of iron sulphide, FeS, in a
generating flask, and pour enough dilute sulphuric acid
through the funnel tube to cover them. Write the
equation representing the reaction.

(b) Collect some of the hydrogen sulphide by the
displacement of air. Make the necessary tests to enable
you to answer the following questions: 1. What is its
odor and color? 2. Will it burn? 3. Will it support
combustion? 4. Is it soluble in water? 5. Is it acid or
alkaline?

(c) Place in different test tubes 5CC of each of the
following solutions: Copper sulphate, CuSO4; lead nitrate,
Pb(NO3)2; mercuric chloride, HgCl2; arsenic chloride,
AsCl3; antimony chloride, SbCl3; stannous chloride,
SnCl2. Pass some of the gas from the generator into each
tube. Sulphides of the metals are formed; note the color
of each precipitate and write the equations which repre-
sent the reactions that take place.

(d) Repeat (c) using, instead of the gas, water
through which the gas has been passed. What is the
result?

(e) Hydrogen sulphide is thus used in analytical
operations, and is a group reagent. Pass some of the gas,
or pour some of the solution used in (d) into 5CC of a
solution of barium nitrate. Is a precipitate formed? Mix
5CC of the lead nitrate solution with 5CC of the barium
nitrate solution in a test tube, and pass the gas into the
liquid. What is precipitated? How could you separate
the lead from the barium?

(f) Pour a little lead acetate solution on a piece of
filter paper and allow the gas to come in contact with it,
or pour a little of the water solution over it. Note the

40 EXPERIMENTAL CHEMISTRY

change. What is formed? This is a characteristic test.

SULPHURIC ACID.

32. This experiment should be performed with the
assistance of the teacher as it requires careful manip-
ulation.

(a) Take a large flask and insert a stopper having
five openings. Pass a delivery tube from three of these
openings to three smaller flasks, and leave the others open
to the air. It is advisable, when convenient, to force air
into the large flask through one of the openings by means
of a pair of bellows. Into one of the small flasks put
some copper turnings, or foil, and concentrated sulphuric
acid; into another copper turnings or foil and dilute nitric
acid; into the third water. Heat the flasks containing
water, and sulphuric acid and copper. Notice the fumes
that pass into the large flask. A current of steam is
formed; sulphur dioxide; and oxides of nitrogen — mainly
nitric oxide. These gases react upon one another
in contact with the air and form sulphuric acid.

This experiment may be performed on a still smaller
scale by using test tubes instead of flasks.

(b) 1. Test the liquid formed with red and blue
litmus paper. 2. Pour 15CC of water into an evaporating
dish and add 5°° of concentrated sulphuric acid. What
is the effect upon the temperature of the solution? Save
the dilute acid. 3. Take 5CC of the dilute acid in a test
tube and add barium chloride solution, BaCl.2. What is
the color of the precipitate formed? What is it? Write
the equation. This is the characteristic test for sulphuric
acid and soluble sulphates. Try the solubility of the
precipitate in acids.

4. Put some wood shavings into an evaporating dish
and pour over them some concentrated sulphuric acid.

EXPERIMENTAL CHEMISTRY 41

Note the effect. How does the product compare with the
charcoal formed in (b), Exp. 24? Is there any analogy
between the two methods? Which is the ordinary
method of preparation?

ACIDS, BASKS, SALTS, ETC.

Acids consist of hydrogen with a negative element; i.
e., a non-metal, (Hydracids; as hydrochloric, HC1, and
hydrogen sulphide, H.,S), or with oxygen and a negative
element (Oxyacids, as nitric, HNO3, and sulphuric,
H.2SOJ. Thio-acids contain sulphur in place of oxygen.
Acids generally redden litmus, especially those which are
soluble — all the stronger acids are soluble. Anhydrides
are oxyacids minus water. Mono- di- tri- and tetra-
basic acids contain one, two, three, and four atoms of
basic or replaceable hydrogen to the molecule; as, HC1,
H2S04> H3PO,, H4Si04.

Elements which combine with hydrogen or hydrogen
and oxygen to form compounds with acid properties are
called acid-forming elements or non-metals; such as
sulphur, nitrogen, fluorine, chlorine, bromine, iodine,
phosphorus, arsenic, carbon, and silicon.

Metals are those elements which combine with
oxygen, or with oxygen and hydrogen, to form com-
pounds which have basic properties. They are called
base-forming elements or metals; such as potassium,
sodium, calcium, magnesium, aluminum, zinc, iron, lead,
tin, copper, silver, mercury, platinum, and gold.

Bases (hydroxides or hydrates) contain a metal or
positive element with hydrogen and oxygen. The
strongest bases are the Caustic Alkalis, and the Alkaline
Earths. The Caustic Alkalis, potassium hydroxide,
KOH; sodium hydroxide, NaOH; and ammonium-
hydroxide, NH4OH, are very soluble in water, and their

42 EXPERIMENTAL CHEMISTRY

solutions turn litmus blue. The Alkaline Earths, barium
hydroxide, Ba(OH),; strontium hydroxide, Sr(OH),;
and Calcium Hydroxide, Ca(OH)2, are less soluble but
turn litmus blue. Most other bases are insoluble in
water and do not turn litmus blue.

The substances formed when an acid neutralizes a
base are called salts. Water is formed at the same time.
Besides litmus, other indicators, chiefly organic, as coch-
ineal, methyl orange, etc., may be used to distinguish
acids from alkalies by change of color. The salts formed
by neutralizing sulphuric acid are called sulphates; those
formed from nitric acid are called nitrates; and those
derived from hydrochloric acid are called chlorides.

A normal salt is one which is formed by replacing all
the hydrogen of an acid with a metal; as sodium carbon-
ate, Na2CO3; zinc sulphate, ZnSOj. An acid salt is one
which is formed by replacing only a part of the hydrogen
of an acid with a metal, as KHSO,, and NaHCO3. A
basic salt is one in which all the hydrogen is replaced by
a metal and a further quantity of oxide or hydrate enters
into the molecule, as ZnSO4. ZnO.

33. (a) Evaporate to dry ness in a porcelain evap-
orating dish, lO00 of dilute hydrochloric acid, note the
appearance and amount of residue.

(b) Repeat (a) with 10CC sodium hydrate solution,
NaOH.

(c) Take 10 to locc of dilute hydrochloric acid in an
evaporating dish, add sodium hydrate drop by drop
till the solution no longer affects either red or blue litmus.
Test by inserting small bits of litmus paper, or put one
or two drops of litmus solution into the liquid.
Evaporate to dryness, cool, and examine the residue.
What does it taste like? Dissolve a little in water.

EXPERIMENTAL CHEMISTRY 43

Does it affect litmus? Write the equation which shows
what has taken place.

(d) Repeat (c) using sulphuric acid and sodium
hydrate.

(e) Repeat (c) using nitric acid and potassium
hydrate.

NEUTRALIZATION.

34. (a) Fasten two burettes to a ring-stand. Fill
one with sulphuric acid, the other with prepared sodium
hydrate containing .05 grams of actual sodium hydrate
to the cubic centimeter. (Ask for the liquids to be used.)
In using burettes, if previously used for solutions of
different strength or composition, or if not perfectly dry
and clean, rinse well with water, and then with a small
quantity of the liquid to be used. Always fill with a
clean dry funnel. Read the level carefully by the lower
meniscus, before and after drawing out the liquid.

(b) Run about 10CC of the soda solution, noting the
exact amount, into a clean beaker standing on white
paper; add two or three drops of litmus solution for an
indicator, then run in acid little by little, stirring
constantly, till neutral; the slightest excess of acid is
indicated by a faint permanent pink color. Note the
exact volume of each liquid taken and then determine
the ratio— 1:X.

Form of Noting Burette Readings:

Alkali (NaOH) Acid (H.SOJ

2.6 0.

12.8 15.5

Volume 10.2 : 15.5=1 :X. X=1.5+

44 EXPERIMENTAL CHEMISTRY

(d) Take about 15CC of the acid solution in a clean
beaker, add the indicator, and then run in the soda
solution till neutral (end of reaction is disappearance of
pink color). Compare the value of X found, and take
the mean of the two most concordant as the true value of
X. What does the experiment show regarding
combination in definite proportions? Write reactions
corresponding to what has taken place. As one cubic
centimeter of sodium hydrate is known to contain .05
grams of sodium hydrate, calculate: How many grams
of sulphuric acid will neutralize lcc of the sodium
hydrate. How many grams of sulphuric acid there are
in lcc of the sulphuric acid solution. By this last
calculation you will standardize the acid solution. Use
this standard to determine by a similar procedure:

(e) What weight of potassium hydrate there is in
the total amount of test solution given you by your
instructor.

Acidimetry and Alkalimetry are based on these
methods of neutralization (saturation); a known quantity
of pure acid or pure alkali being taken as a starting-
point.

Problem: — Write the equation and calculate how
many grams of actual sulphuric acid are needed to
neutralize 5 grams of actual sodium hydrate. How much
sodium sulphate wjll be formed?

PART II.

THE METALS AND THEIR COMPOUNDS.

POTASSIUM.

35. (a) (Always handle potassium and sodium
with forceps.)

Cut off a small piece of potassium from one of the sticks
in the bottle, and note the appearance of the cut portion.
Place the piece of potassium on your ring-stand and
leave exposed to the air. Note the changes it undergoes.

(b) Take a small piece of potassium, about one-
fourth the size of a pea, from the bottle; and, after
absorbing the adhering oil with a filter paper, drop it
upon the surface of water in an evaporating dish. Is
potassium heavier or lighter than water? Why is
potassium kept under oil? Drop a second piece of
potassium upon the water and look at the flame through
a blue glass. What is the color of the flame as seen
through the glass?

(c) Examine the water in the evaporating dish; wet
the fingers with it; try its action on red litmus paper.
Write the equation which represents the reaction that has
taken place.

(d) Some volatile substances produce characteristic
colors. It is necessary to moisten some with hydrochloric
acid or concentrated sulphuric acid. Put some fine
potassium chloride in a watch crystal and moisten with
hydrochloric acid. Make a loop in the end of a piece of

46 EXPERIMENTAL CHEMISTRY

platinum wire, and dip it into the potassium chloride.
Hold it in the outer edge of a Bunsen or alcohol flame
and observe the color of the flame.

(e) Repeat the flame test using another potassium
salt. lyook at the flame through a blue glass. What is
its appearance? In most cases the platinum wire may be
dipped into strong solutions of the salt, or the wire
moistened and dipped into a little of the salt.

POTASSIUM COMPOUNDS.

36. (a) The Hydroxide.

Dissolve 10 grams of potassium carbonate in about
125CC of water and boil in an iron (or silver) dish. Why?
Slowly add 5 grams of slaked lime to the boiling liquid,
and stir constantly until all the lime has been added.
Allow the solution to cool and then decant off the clear
liquid into a bottle and use when required. Try its
action upon litmus paper. In what other way has
potassium hydrate been prepared?

(b) How may potassium sulphate, and potassium
nitrate (saltpetre) be made? Review Experiment 33.

Most of the potassium nitrate in use is made from
sodium nitrate (chili saltpetre) by treating with potassium
chloride. Write the equation.

Saltpetre or nitre is used in the preparation of
sulphuric acid (For what purpose?), and in the
preparation of gunpowder.

Mix together intimately in a mortar about 4 grams of
potassium nitrate; and one gram of charcoal and
sulphur, using a little more charcoal than sulphur. Pour
the powder on the iron base of your ring-stand and set
fire to it. Potassium nitrate is easily broken up and

EXPERIMENTAL CHEMISTRY 47

contains enough oxygen for the combustion of the mass.
Gases, principally carbon dioxide and nitrogen, are
liberated; and, at the time of combustion, occupy several
hundred times the volume of the solid powder. Hence
the great explosive force when the powder is confined.

(d) Colored flame.

Mix intimately in a mortar 2 grams of powdered stron-
tium nitrate; 2 grams of powdered potassium chlorate; and
one-half gram of flowers of sulphur. Set fire to the
mixture as before. What is the color of the flame?

(e) Repeat the same experiment using barium
nitrate instead of strontium nitrate. What is the color
of the flame produced?

SODIUM.

37. (a) Bxamine a piece of metallic sodium, and
leave it exposed to the air as in the case of potassium.

(b) Throw a piece of sodium upon the surface of
water; and when it has disappeared, throw a second piece
upon the water and touch it with a lighted match. The
same effect may be produced by throwing the sodium upon
hot water. Try it. How does the action of sodium upon
water differ from the action of potassium upon water?

(c) Repeat (d) and (e) Experiment 35, using
sodium salts instead of potassium salts. What is the
color of the flame? What effect has the blue glass?

(d) Try a mixture of a sodium salt and a potassium
salt. What is the color of the flame? Could you detect
the presence of potassium in the mixture by looking at
the flame?

lyook at the flame through the blue glass. What is the
appearance of the flame? How could you distinguish
between potassium and sodium compounds?

48 EXPERIMENTAL CHEMISTRY

(e) Examine a number of sodium compounds as to
color and solubility in water.

(f) Compare briefly the properties of potassium and
its compounds with the properties of sodium and its
compounds.

SODIUM COMPOUNDS.

38. (a) How may the following sodium compounds
be prepared? Sodium chloride, sodium sulphate, and
sodium nitrate. See Kxp. 33. Sodium chloride is
obtained in large quantities from sea water. How? The
preparation of sodium hydroxide is similar to that of
potassium hydroxide. Write the equation.

(b) Sodium carbonate. (Solvay or ammonia
process. )

Take about 2-r)cc of ammonia solution and pass carbon
dioxide from a generator into it until the carbon dioxide
is no longer absorbed. A solution of acid ammonium
carbonate is formed.

NH4OH+CO2=HNH4CO8.

Slowly add to this a strong solution of sodium
chloride as long as a precipitate is formed. This is
monosodium carbonate, which is comparatively difficultly
soluble in water.

HNH4CO3-fNaCl=HNaCO3-fNH4Cl.

What is left in solution?

Filter the precipitate and dry in an oven. (It may be
partially dried by pressing between filter papers.)

Transfer the precipitate to a crucible or ignition tube,
and heat until carbon dioxide is no longer given off. The
residue is sodium carbonate.

2HNaCOs=NaiCO,-fCOa+H20.

Put some of the residue in a test tube and add a little

EXPERIMENTAL CHEMISTRY 49

dilute hydrochloric acid. What is given off? Determine
the presence of the metal by the flame test. What is
the color of the flame?

AMMONIUM COMPOUNDS.

39. (a) Ammonium is a hypothetical metal, being
too unstable to exist alone.

Review the experiments on ammonia. What does
ammonia gas dissolved in water form? How does it affect
litmus paper? Name a characteristic test for ammonia.
See (a) Experiment 18.

(b) Examine a number of the compounds of
ammonium as to color and solubility in water. How do
they compare with potassium and sodium compounds in
these respects? Do ammonium salts give a characteristic
flame.

Always have your platinum wire clean before trying
the flame test. Clean the wire by dipping it into hydro-
chloric acid and igniting. Sometimes it is necessary to
scrape the wire slightly with a knife.

Take 50CC of a comparatively strong solution of
ammonia, and divide it into two parts of about 25CC each.
Pass hydrogen sulphide from a generator into one part
until the solution is saturated. A characteristic odor will
be noted when the solution is near saturation.

The product is a solution of ammonium hydrosul-
phide. Add to this the other part of the ammonia
solution and ammonium sulphide will be formed. Note
the odor and color of the liquid. Pour it into a bottle and
cork securely.

Ammonium sulphide is a group reagent and is used
to precipitate those sulphides which are soluble in dilute

50 EXPERIMENTAL CHEMISTRY

hydrochloric acid. Take two test tubes containing two
or three cubic centimeters of zinc sulphate solution.
Pour a little hydrochloric acid into each test tube. Is a
precipitate formed?

Add a few drops of hydrogen sulphide solution to one
and a few drops of ammonium sulphide to the other.
What are the results?

Repeat the tests using cobaltous chloride instead of
zinc sulphate.

(d) How are the following compounds formed: —

Ammonium chloride, ammonium nitrate, and ammo-
nium sulphate?

Write the equations.

CALCIUM COMPOUNDS.

40. The element, calcium, is not found uncombined.
It has no practical application, hence it is not made in any
considerable quantity.

(a) Calcium Hydrate.

Put 10 grams of quicklime (calcium oxide) in an
evaporating dish and slowly add about 20CC of water to it.
Observe what takes place. Slaked lime (calcium hydrate)
is formed, and the process is called slaking.

Add 300 or 400°° of water to the slaked lime and pour
the whole into a bottle and cork securely. Examine it
the next laboratory hour. Has any of the lime dissolved?
Test some of the clear solution with litmus paper. Is
there any change? The solution is known as lime-water.
It may be poured from your bottle into the bottle labeled
" Lime- water."

What takes place when you blow your breath through
clear lime-water? What change does it undergo when
exposed to the air? Review (a) and (f) Experiment 26.

EXPERIMENTAL CHEMISTRY 51

Add lime-water to a solution of ferric chloride. Fer-
ric hydroxide is thrown down. What is its color?

Add a little sulphuric acid to lime-water. What is
formed?

(b) Calcium sulphate (gypsum).

Heat some powdered gypsum in an ignition tube.
Does it contain water of crystallization?

Put 4 or 5 grams of gypsum in a crucible and heat
for twenty or thirty minutes in an air bath or over a
flame at about 150° — avoid heating above 200°.

Add enough water to the residue to form a paste and
allow it to stand. Moisten some gypsum that hks not
been heated and allow it to stand for a time. Is there any
difference between the two? Gypsum heated in this way
forms a powder known as plaster of Paris.

(c) Take the carbonate, the chloride, and the sul-
phate of calcium and try their solubility in water, in acids
(hydrochloric and nitric), and in alcohol.

What has calcium chloride been used for in previous
experiments? Water containing calcium carbonate, or
water containing calcium sulphate in solution, is called
a hard water. The former, temporarily hard because
the hardness is easily removed by boiling; the latter,
permanently hard because the hardness is not removed
by boiling.

(d) Take a little calcium chloride and see what color
it gives to the flame. Compare it with sodium.

What tests would you employ to distinguish between
the chlorides of potassium, sodium, ammonium, and
calcium?

ZINC.

41. (a) Examine some metallic zinc. What is its
color? What effect has the common acids (hydrochloric,

52 EXPERIMENTAL CHEMISTRY

sulphuric and nitric) upon ordinary zinc? What gas is
usually evolved?

Try the effect of caustic alkalies upon zinc. See if
you can write the equations expressing the reactions.

For what purpose has zinc been previously used?
Mention any uses of zinc with which you are familiar.
What is the equivalent weight of zinc?

(b) Heat a small piece of zinc on charcoal in the
reducing flame. Does the zinc burn? What is the resi-
due formed?

Note the color of the sublimate while hot and when
cold? When it becomes cool moisten with dilute cobalt
nitrate and heat again. What is the color of the
sublimate?

ZINC COMPOUNDS.

42. (a) Examine several of the compounds of zinc.
What is their usual color?

(b) Take four test tubes and pour into each about
5°° of zinc sulphate solution. Add to one of these a
solution of potassium hydroxide; to another, a solution of
sodium carbonate; to the third, ammonium hydrate
solution; and to the fourth, a solution of ammonium
sulphide. Give the colors of the precipitates. Write the
equations to show what is formed in each case. Under-
line the parts which represent the precipitates.

(c) Take the test tube in which you added potassium
hydroxide, and add an excess of the reagent. What is
the effect? Divide the solution into two parts. Dilute
one part, and boil the other. What are the results?

OXIDATION AND REDUCTION.

Oxidation takes place when oxygen or any
negative element or group is added to an element or

EXPERIMENTAL CHEMISTRY 53

compound, and when hydrogen or any positive element
or group is removed — it consists in the increase of the
proportion of the more negative constituents.

Reduction takes place when oxygen or any negative
element or group is removed from a compound, or when
hydrogen or any positive element or group is added — it
is the reverse of oxidation.

The valence of an element (monad, dyad, triad, etc. )
is the number of bonds or units of chemical force it
possesses. It is measured by the number of atoms of
hydrogen or chlorine which one of its atoms combines
with or replaces. Oxidation or reduction takes place
when the number of acting bonds of an atom of any
element suffers change.

The following are the principal elements which form
two series of salts: — Iron, mercury, tin, antimony,
arsenic, manganese, and copper. By suitable means, the
lower (ous) salts of each can be converted into the higher
(ic) salts and vice versa.

In the following experiments on oxidation and
reduction, use very small quantities of the substances
called for, and have the vessels you use well cleaned.
Notice carefully the nature of each reaction.

43. (a) Burn a little sulphur in air. What is
formed? Has oxidation or reduction taken place?

(b) Put a little finely powdered sulphur in a test
tube and add about 5CC of nitric acid. Warm for about
ten minutes. Part of the sulphur is oxidized to SO3.

S4-2HNO3=H2SO4-f2NO.

Test the presence of sulphuric acid by diluting the
solution to four or five volumes and adding barium
chloride. A white precipitate, insoluble in acids,
indicates a sulphate. Write the equation.

54 EXPERIMENTAL CHEMISTRY

(c) Place a drop or two of mercuric chloride solution
on a clean piece of copper or zinc. It is reduced finally
to metallic mercury. Write the equation. What is the
nature of the change in the nitric acid in (b), and of the
zinc or copper in (c)?

(d) Take 5CC of mercuric chloride, HgCl.,, in a test
tube, and add one or two drops of stannous chloride,
SnCl2; warm and note the result. Now add a considerable
amount of stannous chloride, heat and observe the result
as before. Stannic chloride, SnCl4, is formed in each
case; in the first mercurous chloride, Hg2Cl,, and in the
second mercury is precipitated.

(1) 2HgCl24SnCli=Hg2Cl2-fSnCl4.

(2) HgCl2+SnCl2=Hg-fSnCl4.

(e) Put a small globule of mercury in a test tube,
and add 5CC of dilute (diluted 5 to 1 ) nitric acid. Warm
very gently for a few minutes, if necessary to produce
action. Mercurous nitrate is formed. Test the liquid
with a few drops of hydrochloric acid. What is the color
of the precipitate formed? What does it indicate? Write
the equation.

When hot concentrated nitric acid is used in excess,
mercuric nitrate is formed which gives no precipitate
when treated with hydrochloric acid. Mercuric chloride
is soluble.

Iodine as a Test for Oxidizing and Reducing Agents.

The liberation of iodine from potassium iodide or
hydriodic acid indicates an oxidizing agent, the free
iodine is recognized by the brown color it gives to
water; and the wine color it gives to kerosene, carbon
bisulphide, etc. The test can be made more delicate by
introducing freshly prepared starch paste which is colored

EXPERIMENTAL CHEMISTRY 55

blue by a mere trace of free iodine. (The color of the
iodide of starch disappears temporarily while heated.)

Conversely, the removal of this blue color generally
indicates a reducing agent.

44.. (a) Dissolve a small crystal of potassium iodide
in about 25CC of water, or take 5CC of the prepared
potassium iodide solution and dilute it.

Put 3 or 4CC of the solution into each of three test
tubes, and add to each about 2CC of kerosene or a few
drops of carbon bisulphide. Into one test tube add a
small quantity of ferric chloride solution; into another, a
little copper sulphate solution; and into the third, dilute
nitric acid. Observe any changes in appearance.

(1) Fe8Cl6+2KI=I2H-2KCl-f2FeCl2.

(2) 2CuSO4-t-4KI=I2-|-2K2SO4-f-Cu2I2.

(3) 4HNO3-f 3KI=IS-+ 3KNO3+NO+2H2O.

(b) Prepare some starch paste and add some
potassium iodide solution to it; stir well and keep
covered. Take small quantities in test tubes and add
mere traces of the following:

(1) Ferric chloride, (2) copper sulphate, (3)
mercuric chloride, (4) chlorine water, (add this reagent
drop by drop and finally in considerable quantity).
What changes do you notice?

Heat some of the solutions and then allow them to cool.
What is the effect?

HgCl2+2KI=2KCl-fHgI2.

Cl-f KI=KC1-|-I (with comparatively little chlorine).

(c) Make some of the starch paste, containing
potassium iodide, blue by adding a drop or two of
bromine or chlorine water.

To small quantities of this in three test tubes, add
solutions of the following:

56 EXPERIMENTAL CHEMISTRY

(1) Hydrogen sulphide, (2) stannous chloride, (3)
sodium hydrate. Is there any change in color produced?

(1) H2S+2I=2HI+S.

(2) 2SnCl24-4I=SnCl4-fSnI4.

(3) 6NaOH-f6I=r3H,O-j-5NaI+NaIO3.

IRON.

45. (a) What is the color of a freshly broken piece
of iron? What is ''iron rust"?

Heat a little ferric chloride (or any iron compound)
on charcoal. What is the color of the residue? Allow it
to cool and test with a magnet. Has it magnetic
properties? Try another compound.

Examine several of the ferric and the ferrous salts and
solutions. Can you make any distinction as to color?

IRON COMPOUNDS (OXIDATION AND REDUCTION).

46. (a) Place about one-half gram of fine iron
wire or filings in a small flask and add about 50CC of
dilute sulphuric acid, and then a little sodium carbonate.
(Carbon dioxide forms and expels the air. ) Insert a
delivery tube and warm gently till action ceases.

Carbon, silicon, etc., remain undissolved. What is in
solution?

Write the equation representing the action of
sulphuric acid on iron.

Allow the undissolved portion to settle, and then pour
off the clear solution — one-half into a beaker, the other
half into a bottle. Dilute the solution in the bottle by
adding an equal volume of water; add a pinch of sodium
carbonate to displace the air, and when action ceases,
cork the bottle.

Heat the ferrous sulphate in the beaker to boiling and
add 4 or 5CC of concentrated nitric acid and evaporate to

EXPERIMENTAL CHEMISTRY 57

one-half. If necessary, add more nitric acid and boil
again until you notice a distinct change in color. Ferric
sulphate is formed. Observe the characteristic colors of
ferrous and ferric solutions.

(b) Take about a cubic centimeter of the ferrous
sulphate solution and add at once a solution of sodium
hydroxide; ferrous hydroxide is precipitated.

Treat a similar amount of ferric sulphate with the
sodium hydroxide solution; ferric hydroxide is precipitated.
What is the appearance of each precipitate? Write the
equation expressing the reaction in each case.

Pour a little of the ferrous sulphate solution into a
test tube and leave it, and also the test tube containing the
ferrous hydroxide, open to the air. Shake them from
time to time. Do you notice any change? What forms
in each case?

(c) 1. Treat a small quantity of ferric sulphate with
about 10CC of hydrogen sulphide solution; warm slightly.
Fe2(SOJ3H-H2S==2FeSO4-4-H.,SO4-fS. (white precipitate).

2. Place a piece of zinc in a little of the ferric
sulphate solution and leave for ten minutes. What change
do you notice? Write the equation expressing the reac-
tion.

Mention several oxidizing agents, and several reduc-
ing agents.

COPPER.

47. (a) Examine some copper wire or sheet copper;
also examine some of the metallic copper in the stoppered
bottle. Is there any difference in appearance between
that in the bottle and that exposed to the moist air of the
room?

Mention any uses of copper with which you are
familiar.

58 EXPERIMENTAL CHEMISTRY

Test the effect of the common acids, both the concen-
trated and the dilute, upon metallic copper. Does heat
modify the action? •

(b) Examine several of the compounds of copper.
Which compound have you already examined, and for
what purpose?

What seems to be the general color of the solutions of
the copper compounds? Take a copper compound,
moisten a little with hydrochloric acid, and see what color
it gives to flames.

Test some of the solutions to see if they have either
an acid or an alkaline reaction.

(c) How is copper sulphate made? Obtain some of
the solution and pour about 5CC into each of three test
tubes; then try the effect of each of the following solu-
tions: Potassium hydroxide, hydrogen sulphide, and
potassium ferrocyanide. Give the color of each precip-
itate. Write the equations and underline the parts
representing the precipitates.

Heat the test tube into which potassium hydroxide
was added. Copper oxide is formed. What is its color?
write the equation.

EQUIVALENT WEIGHT OF COPPER.

48. (a) Brighten some narrow strips of pure zinc,
and weigh out exactly one gram. Clean thoroughly a
comparatively large evaporating dish and place the zinc
strips, loosely coiled, in the bottom. Weigh on a rough
balance some copper sulphate crystals, about ten times
the weight of zinc taken, dissolve in about 100CC of hot
water, and then filter the hot solution upon the zinc in the
basin. Place the dish on your ring-stand on a wire gauze
and heat from one to two hours (or, if more convenient,

EXPERIMENTAL CHEMISTRY 59

allow it to remain covered till the next laboratory
period. ) until the deposit ceases to increase.

(b) Allow the deposit to settle, and then decant off
all the liquid possible through a filter paper into a large
beaker. Add water to the basin, heat, allow the copper
to settle, and decant as before. Wash the copper several
times in this manner until a portion of the wash water in
a test tube gives no turbidity when barium chloride is
added. Finally transfer the copper deposit to the filter,
wash again to settle it, and dry in an oven at a temper-
ature not above 100°.

(c) Carefully transfer the metallic copper to a watch
glass and weigh accurately. This amount of copper has
been replaced by one gram of zinc. Hence, knowing the
equivalent of zinc from Experiment 15, calculate the
equivalent weight of copper.

Copper may also be precipitated from its salts by iron
and some other metals. Hence their equivalent weight
may be determined in a similar manner.

LEAD.

49 (a) Take a piece of metallic lead and scrape off
the outer surface with a knife. What is its color?

What is the appearance of lead that has been exposed
to the air for some time? See if you can write on paper
with the metal. Will the common acids dissolve lead?
Try them.

(b) Place on a piece of charcoal a little lead sulphide
(galena) . Cover it with sodium carbonate and heat in the
reducing flame. Note the color of the sublimate while
hot and when cold. Is it volatile? Remove the beads of
metal and test properties. Can you mark on paper with
them?

60 EXPERIMENTAL CHEMISTRY

(c) 1. Place in a test tube apiece of bright sheet
zinc and pour over it 5 or 10CC of lead acetate solution.
Set the tube to one side and leave undisturbed for some
time. Explain the action that has taken place. Write
the equation.

2. Remove the zinc, wash thoroughly, and scrape
the deposit off into a test tube. Pour enough dilute nitric
acid into the test tube to cover the lead deposit. Warm
slightly and as soon as action ceases, dilute to about 10CC.

3. Divide the solution into four parts. To one part,
add hydrogen sulphide solution; to another, sulphuric
acid; to the third, potassium chromate; and to the fourth,
hydrochloric acid. What is the color of the sulphide?
the sulphate? the chromate? and the chloride? Write
the equations expressing the reactions.

4 Divide the chloride into two parts, and test the
solubility in cold water and in hot water. Is it soluble?

Divide the hot solution into two parts. Allow one
part to cool, and add potassium bichromate solution to the
other. What are the results?

SILVER.

50. (a) Dissolve half of a silver dime in a little
dilute nitric acid in a beaker. Warm gently till action
is over, then evaporate nearly to dryness and dilute with
100CC of water.

The solution contains silver nitrate, copper nitrate,
and nitric acid. (Traces of gold are sometimes found in
the form of little black specks. These may be separated
from the solution by filtering.) Which of the compounds
gives the solution its color?

All United States silver coins are alloyed with about
10 per cent, of copper.

EXPERIMENTAL CHEMISTRY 61

Heat the solution to boiling, add hydrochloric acid
slowly, and stir till the precipitate gathers in a lump.
What is the color of the precipitate? What is it? What
remains in solution? Filter out the precipitate and dry in
an air bath heated to 1UO°-110°. When the precipitate is
dry, mix a little sodium carbonate with it, and heat on
charcoal in the reducing flame of the blowpipe. White
beads of metallic silver are formed. Remove these
carefully from the charcoal and place them in an
evaporating dish. Dissolve in a little dilute nitric acid
and evaporate to dryness in a water bath. What is
formed? Dissolve the residue in water and filter into a
bottle of dark colored glass, or a bottle covered with dark
paper.

(b) Pour about one-half cubic centimeter of the
clear silver nitrate solution into each of three test tubes,
and add 5CC of water to each. Now add sodium chloride
solution to one, potassium bromide solution to another,
and potassium iodide solution to the third. Name each
precipitate and give its color?

Expose all to the light and watch closely for any
changes.

Silver salts are used extensively in photography on
account of the change produced by light.

51. (a) Place in a test tube a strip of metallic zinc
and pour over it 3 or 4CC of the silver nitrate solution
prepared in Experiment 50. Let it stand for some time
without disturbing. Remove the strip from the tube and
wash thoroughly with water. What is the appearance of
the deposit? Scrape it off and dissolve in dilute nitric
acid.

(b) Divide the solution into two parts and dilute
each with about 5CC of water.

62 EXPERIMENTAL CHEMISTRY

Add a few drops of hydrochloric acid to one part.
How does the precipitate compare with the precipitate
formed in (a) Kxp. 50? Divide the precipitate into three
parts and test its solubility in hot water, nitric acid and
in ammonium hydroxide. Compare with lead chloride,
(c) 4, Experiment 49.

(c) Take the other part of the silver solution
retained from (b), and add a little hydrogen sulphide
solution. What is the color of the precipitate? What
is it?

Divide the precipitate into three parts. Add cold
nitric acid to one part, hot nitric acid to another, and
potassium cyanide to the third part. Note results.

MERCURY.

52. (a) In what form is metallic mercury? Can
you mention any other element that is usually in the
same form? What is the appearance of mercury?

(b) Take four test tubes: put a piece of zinc into one,
a piece of iron into another, a piece of copper into a third,
and a piece of tin into the fourth. Cover each metal with
a solution of mercurous nitrate. After a few minutes,
remove the metals and examine. Rub each with a piece
of soft cloth. Are all permanently affected?

Drop a globule of metallic mercury upon a piece of
sheet zinc and rub it with a cloth. Do you obtain the
same result as with the piece of zinc previously used?

What are alloys of mercury with other metals called?

For what purpose has mercuric oxide been used
before?

EXPERIMENTAL CHEMISTRY 63

(d) Mix some mercurous chloride (calomel) with a
little sodium carbonate and heat in a closed tube Is the
compound reduced to the metal?

(e) Put a globule of mercury in a test tube and dis-
solve in nitric acid. Dilute and then add a few drops of
hydrochloric acid to the solution. What is the color of
the precipitate? What is it?

Divide the precipitate into two parts: add a little
ammonium hydroxide to one part. What is the effect?
Add water to the other part and boil it. Compare with
the solubility of lead chloride and silver chloride in
water, (c) 4. Exp. 49, and (b) Exp. 51.

How could you separate lead nitrate, silver nitrate,
and mercurous nitrate from one another?

Hydrochloric acid gives no precipitate when added to
mercuric solutions.

ACTION OF ACIDS ON THE METALS.

53. (a) In every case brighten metals used by filing
or with sand paper.

Put into a test tube 2CC of dilute sulphuric acid and
insert a small piece of zinc. Note the action. Warm
(not boil) the acid — is the action modified? Add more
zinc and continue heating; can action be carried on
indefinitely?

(b) Take two test tubes each containing about 5CC of
dilute sulphuric acid. Put a piece of smooth zinc into
each, warm slightly, and note the character of the action.

Drop on the zinc in one tube a piece of platinum or
copper wire, or a silver coin; in the other tube put a
small crystal (or two or three drops of a solution) of
copper sulphate. Is action more, or less, energetic?

64 EXPERIMENTAL CHEMISTRY

by 3mm, and place it in a test tube. Add enough dilute
sulphuric acid to half cover the zinc; warm gently.

At what part of the liquid is the action most
energetic?

(d) Take two pieces of zinc of equal size and
equal thickness (about 3cm by 3mm), and fold one up
closely several times.

Place both in dilute sulphuric, or hydrochloric, acid
and observe which dissolves first. Expain.

(e) Take two small pinches of about equal size of
iron filings; place in test tubes and add to each about 5CC
of dilute sulphuric, or hydrochloric, acid. Shake or stir
continuously the contents of one tube, letting the other
stand in the rack. Which dissolves first? Suggest the
cause.

(f) Review your experiments and make a tabulated
outline of the action of concentrated, moderately strong
(one volume of concentrated acid to one of water), and
dilute sulphuric acid; nitric acid; and hydrochloric acid
upon the following metals: zinc, iron, copper, lead, and
mercury.

Perform tests that have not been previously made, and
repeat tests whenever you are in doubt as to the action.

FORMATION OF THE COMPOUNDS OF THE METALS.

A number of the compounds of the metals have already
been considered; a few of these, by way of review, and
some others will now be studied according to their classi-
fication. These will be sufficient to present the general
methods of preparation.

CHLORIDES.

54. (a) What compound is formed when zinc is
treated with hydrochloric acid? Review (a) Experiment 6.

EXPERIMENTAL CHEMISTRY 65

(b) Review (b) 5, Experiment 23, and state what
was formed when you did the experiment.

(d) What takes place when caustic soda is treated
with hydrochloric acid? Review (c) Experiment 33.

(e) What is the effect of adding hydrochloric acid
to calcium carbonate? See (a) Experiment 26.

As illustrated above, the chlorides are made by treating
the metals with chlorine or hydrochloric acid; and by
treating an oxide, hydroxide, or carbonate (a salt of a
volatile acid) with hydrochloric acid.

(f ) Are the chlorides of lead, silver, and mercury
soluble, difficultly soluble, or insoluble in water? Review
Exps. 49, 51, 52. Test the solubility, in water, of several
of the other chlorides of the metals. Are they soluble
or insoluble?

If a precipitate is formed when hydrochloric acid is
added to a solution, what metal or metals, would you sup-
pose to be present?

OXIDKS.

55. (a) Review Experiments 4 and 8, and state
how the metals, iron and magnesium, were changed.

(b) What takes place when copper hydrate is
heated? Represent the change by an equation. Review
(c) Experiment 47.

(c) Put some lead nitrate in an ignition tube and
heat. Oxygen and nitrogen peroxide are given off
leaving lead oxide behind. What is the appearance of
the residue?

(d) When certain carbonates are heated strongly,

66 EXPERIMENTAL CHEMISTRY

carbon dioxide is given off. What is left? How is lime
made?

From the above, state how the oxides are made.

HYDROXIDES.

56. (a) How is calcium hydroxide formed? Review
(a) Experiment 40.

(b) Put 5CC of magnesium sulphate solution in a test
tube and add sodium hydroxide. The precipitate is
magnesium hydroxide. What is its appearance? Write
the equation representing the reaction. Is the precipitate
soluble in hydrochloric acid?

What chlorides are insoluble or difficultly soluble in
water?

(c) Add caustic soda solution to a little ferric
chloride solution in a test tube. The precipitate is ferric
hydroxide. What is its color? Compare with (b)
Experiment 46. Would you expect it to be soluble in
hydrochloric acid? Why?

Observe that in (b) and (c) a soluble hydroxide has
been added to solutions containing metals whose hydrox-
ides are insoluble; hence, the insoluble hydroxides are
precipitated.

Name the soluble hydroxides. How is potassium
hydroxide made? Do all the hydroxides turn litmus
blue?

SULPHIDES.

57. (a) How is ferrous sulphide made? Review
(d) Experiment 2. What use have you previously
made of ferrous sulphide?

(b) Heat some sulphur to near boiling and introduce
into it a piece of bright copper — copper foil or sheet

EXPERIMENTAL CHEMISTRY

ta
t/N/

copper. What change takes place? What is t
product called?

(c) Put a drop or two of sodium sulphide solution, or
hydrogen sulphide solution (any soluble sulphide) on a
bright silver coin, and leave for a few minutes. Rinse
the coin and examine. How has it changed in appear-
ance? The formation is silver sulphide.

(d) Review (c) Exp. 31, and (c) Exp. 39.
How may the sulphides be made?

What reagent is used to precipitate the sulphides
which are soluble in hydrochloric acid?

NITRATES.

58. (a) What is left in solution when a silver coin
is dissolved in nitric acid.-* See (a) Experiment 50.

(b) What is the effect of adding nitric acid to a
metal? Review (d) Exp. 21; (c) Exp. 49; and (e)
Exp. 52.

Write the equation expressing the reactions which
take place when nitric acid is neutralized with ammonium
hydroxide, calcium hydroxide, and potassium hydroxide.

(d) Add nitric acid to a little sodium carbonate in a
test tube. What is the result?

(e) What is the effect of adding sulphuric acid to
sodium nitrate? Review Experiment 21.

Sulphuric acid decomposes all nitrates liberating
nitric acid. What gas is given off when a nitrate is
heated strongly in a closed tube?

(f) Heat some copper nitrate crystals in a closed
tube. Does the salt contain water of crystallization?

(g) Try the solubility, in water, of a number of the
nitrates.

68 EXPERIMENTAL CHEMISTRY

SULPHATES.

59. (a) What is the action of sulphuric acid on the
metallic hydroxides? What is the color of copper sul-
phate? Does the removal of the water of crystallization
affect the form or color of the crystals? See Experi-
ment 14.

(b) Review the following experiments, and state
what sulphates were formed and how: Experiments 6, 15,
21, 22, and 32.

(c) Obtain dilute solutions of lead nitrate, strontium
nitrate, and barium chloride; and a strong solution of
calcium chloride. The sulphates of lead, strontium, and
barium are insoluble in water; the sulphate of calcium
difficultly soluble. Hence a precipitate is formed when
sulphuric acid is added to solutions of these metals.

Add sulphuric acid to the solutions, and note the
color of each precipitate.

(d) The same insoluble sulphates may be formed by
adding any soluble sulphate to the solutions.

Obtain a number of the sulphates from the shelf and
test their solubility in water. Then take samples of the
solutions named in (c) and form the sulphates by adding
soluble sulphates. What is the characteristic test for
soluble sulphates?

(e) Heat some iron sulphate on charcoal in the
reducing flame. Note the odor of the gas given off.
What is it? Sulphates of all the heavy metals give off
the same gas.

CARBONATES.

60. (a) Review Experiment 26. What carbonates were
formed and how? How is sodium carbonate made?

(b) Test the solubility, in water, of the carbonates

EXPERIMENTAL CHEMISTRY 69

of potassium, sodium, and ammonium. What is the effect
of boiling the solutions?

(d) Obtain solutions of the following compounds:
Lead nitrate, iron sulphate, copper sulphate, barium
chloride, and calcium chloride. Add to each a little of a
solution of a soluble carbonate. Precipitates (carbon-
ates of the metals) will be formed in the solution of all
the salts of the metals whose carbonates are insoluble in
water.

How do the results compare with the tests made in (c)?
Filter off the precipitates in each case, wash, and test
for carbonates with acids. What gas is given off?

(e) What is usually the effect of heating a carbonate
strongly? Review Experiment 55.

CR YST ALUZ ATION .

' 'Crystals are solids bounded by plane faces inclined
at definite angles." While there are a great number of
forms of crystals, it has been found that every form can be
referred to one or other of six systems. But crystals of
the same form may differ widely in appearance and habit;
for example, the faces may be unequally developed, but
he angles remain constant. A given chemical compound
usually crystallizes, under the same conditions, in the
same form. The form, "habit," and optical properties of
a crystal serve as a guide to its composition.

Crystallization can only take place when the particles
or molecules are free to arrange themselves. The most
slowly formed crystals are, as a rule, the largest and most
perfect.

61. (a) Carefully clean and dry some glass plates
and put on them, by means of a glass rod or pipette, a

70 EXPERIMENTAL CHEMISTRY

large drop of each of the solutions named and in the order
given. Use saturated solutions. Sodium chloride,
potassium bromide, sodium nitrate, potassium nitrate,
potassium chlorate, mercuric chloride, ammonium chlor-
ide, copper sulphate, ferrous sulphate, zinc sulphate,
potash alum, ammonium alum, chrome alum, sulphur (in
carbon bisulphide). Leave two or three hours and then
examine with a lens. Note the appearance and resem-
blance of crystals of certain salts.

Crystals of salts of similar molecular structure often
occur in the same or very similar forms. Such crystals
are said to be isomorphous; thus, all the alums crystallize
in octahedra; so also sodium chloride, potassium chloride,
potassium bromide, and potassium iodide all crystallize in
cubes.

The same substance may, under different conditions,
crystallize in two or more distinct forms, — e. g. sulphur
(see Bxp. 29.) and zinc sulphate. Such substances are
said to be dimorphous.

Water of crystallization is essential to the form of some
crystals, and in many cases to the color (see Exp. 59),
while other crystals contain none. (See Experiment 12.)

Crystals may be obtained from the gaseous state in the
case of those substances which do not liquefy before
solidifying.

(b) Put a small piece of iodine in a watch glass and
place over it a glass plate; warm the lower glass very
gently.

Notice the iodine crystals formed on the glass plate.

From a mixed solution certain salts will crystallize out
together as a double salt, which is not a mere mixture,
as the number of molecules bear a definite ratio; thus a
solution of the sulphates of aluminum and potassium yield,

EXPERIMENTAL CHEMISTRY 71

upon evaporation, crystals of alum of the composition
K2SO4.A12(SO4)8.24H2O.

Note — A solution should be prepared, concentrated,
and poured into a shallow crystallizing dish so that the
class may observe the growth of the crystals from day to
day. The temperature must not fluctuate if good crys-
tals are to be had.

Crystals deposited from impure solutions contain less
impurity than the solution, i. e. a less proportion; hence
crystallizable soluble substances can, by repeated solution
and partial recrystallization, be almost absolutely free from
all impurities — excepting isomorphous substances, and
those salts with which they form double salts.

DETERMINATION OF ATOMIC WEIGHTS

The atoms of most elements in the solid state have
approximately the same capacity for heat; as their atomic
weights increase, their specific heats decrease, so that the
product of the atomic weight times the specific heat
approximates to a constant value. The average value
is 6.4.

Hence, atomic weight x specific heat=6.4.

"The equivalent of an element always bears some
numerical relation to its atomic weight. As a rule, this
relation is a simple one. With hydrogen and chlorine,
equivalent and atomic weight are equal; the atomic
weight of oxygen is twice its equivalent; while that of
nitrogen is three times its equivalent."

Elements whose equivalent and atomic weight are
equal are called Monads; those whose atomic weight is
twice the equivalet are called Dyads; etc.

"The word Quantivalence or Valency is applied to
denote generally the state of an element as regards its
function as a monad, dyad, etc."

72 EXPERIMENTAL CHEMISTRY

The valency of an element, then, is expressed by the
number of times which the equivalent is contained in the
atomic weight.

As the product 6.4 varies somewhat with different
metals, the method cannot be used to determine atomic
weights exactly, but it decides whether the equivalent, or
combining weight, is equal to the atomic weight, or is
some fractional proportion of it.

62. (a) Weigh accurately a piece of metallic copper
(copper wire loosely coiled if obtainable), and suspend it
by a fine wire in a beaker of boiling water.

Counterpoise a calorimeter, and then add somewhat
more than enough water to cover the metal, — take a
round number of grams (200, 300 or 500) to simplify
calculations, and weigh accurately.

In weighing liquids it is best to add a little more than
the desired amount and remove the excess by absorbing
it with a piece of blotting paper.

Note the exact temperature of the water in the calori-
meter. Remove the copper from the boiling water, jerk
off as much water as possible, and lower it immediately
into the calorimeter. Stir the water with the metal for
about a minute. Note the temperature from time to time
and take the highest reading.

Calculate from the data obtained the specific heat of
copper.

(b) Determine, in a similar manner, the specific heat
of tin and of lead.

Tabulate the values found. Write the names of the
metals (copper, tin, and lead, in the order given) in the

EXPERIMENTAL CHEMISTRY 73

first column; the corresponding atomic weight in the
second; corresponding specific heat in the third; and the
product of the atomic weight by the specific heat in the
fourth column. Get the average of the three products.

Are the atomic weights of the metals in the order of
magnitude, or not? Are the specific heats in the same, or
the reverse, order?

How does the atomic weight of copper as determined
compare with its equivalent weight found in Exper-
iment 48?

APPENDIX.

TABLE OF COMMON ELEMENTS. .

NAME, SYMBOL, AND ATOMIC WEIGHTS IN EVEN NUMBERS-

Aluminum Al... 27 Lead Pb...205

Antimony Sb...l20 Magnesium Mg.. 24

Arsenic As... 74 Manganese Mn .. 55

Barium Ba...l36 Mercury Hg..l99

Bismuth BL..207 Nickel Ni... 58

Boron.. B ... 11 Nitrogen N ... 14

Bromine Br... 80 Oxygen O ... 16

Calcium Ca... 40 Phosphorus P ... 31

Carbon C 12 Platinum Pt...l94

Chlorine Cl... 35 Potassium K... 39

Chromium Cr... 52 Silicon Si... 28

Cobalt Co... 59 Silver ....Ag..l07

Copper Cu .. 63 Sodium Na.. 23

Fluorine F 19 Strontium Sr... 87

Gold Au ..196 Sulphur S 32

Hydrogen , H... 1 Tin Sn...ll8

Iodine I 126 Zinc Zn .. 65

Iron Fe... 56

COMPLETE LIST OF CHEMICALS REQUIRED
FOR THE COURSE.

ESTIMATE FOR A CLASS OF 12.

Antimony chloride ^ Ib.

Antimony, powdered ^ Ib.

Arsenic chloride i oz.

Alcohol 4^ gal. if used for lamps

Alum, potash... % Ib.

" chrome % "&.

11 ammonium Ib.

76 EXPERIMENTAL CHEMISTRY

Ammonium hydrate, cone 3 Ib.

nitrate 2 Ib.

sulphate # Ib.

sulphide # Ib.

chloride % Ib.

carbonate i Ib.

Barium chloride ^ Ib.

nitrate # Ib.

Bromine i oz.

Boneblack ^ Ib.

Borax % Ib.

Calcium oxide 2 Ib.

carbonate, lumps 2 Ib.

chloride % Ib.

sulphate # Ib.

Carbon bisulphide ^ Ib.

Cobalt chloride i oz.

" nitrate i oz.

Copper oxide i oz.

sulphate, pure ^ Ib.

nitrate % Ib.

foil X Ib.

sheet X Ib.

" wire yz Ib.

Chlorine water (made in Lab.)

Cochineal i oz.

Charcoal 2 doz. sticks

Candles .' ^ dozen

Cotton print few pieces

Cardboard few pieces

Ether ^ Ib.

Ferrous sulphate ^ Ib.

sulphide 2 Ib.

EXPERIMENTAL CHEMISTRY 77

Ferric chloride 3 oz

Filter paper, 5 in. diam 4 packages

Gypsum, powdered 4 oz.

Hydrochloric acid, cone 6 Ib.

Hydrogen sulphide (made in Lab.)

Iron, sheet i Ib.

" filings, fine i Ib.

Iodine i oz.

Indigo solution % Ib.

Kerosene i Ib.

Lead nitrate % ft>-

<l acetate, pure % Ib.

" sulphide # Ib.

11 sheet % Ib.

Litmus i oz .

Litmus paper, red and blue i box each

Labels i box each of large and small

Magnesium ribbon ^ oz.

sulphate # Ib.

Manganese dioxide, coarse 2 Ib.

Mercury i Ib.

Mercurous chloride i oz.

nitrate % Ib.

Mercuric chloride X Ik-
Mercuric oxide X Ik-

Nitric acid, cone 6 Ib.

Oxalic acid crystals i Ib.

Oil of turpentine % Ib.

Paraffine i Ib.

Phosphorus i oz .

Potassium, metallic X oz-

hydrate, sticks i Ib.

chloride Ib.

78 EXPERIMENTAL CHEMISTRY

Potassium, nitrate i Ib.

carbonate X Ib.

iodide X Ib.

bromide i oz.

'£ cyanide ^ Ib.

permanganate % Ib.

bichromate .'. . . ^ Ib.

chlorate 2 Ib.

chromate ^ Ib.

ferrocyanide % Ib.

Platinum foil 2 pieces i in. square

wire i ft.

Silk, unbleached 25 cts

Sodium, metallic i oz.

chloride, pure i Ib.

sulphide i oz.

sulphate ^ Ib.

hydrate, sticks i Ib.

nitrate 2 Ib.

carbonate % Ib.

Sulphur, roll y2 Ib.

flowers y2 Ib.

Sulphuric acid, cone 9 Ib.

Stannous chloride i oz.

Starch y2 Ib.

Strontium nitrate % Ib.

Sugar i Ib.

Sealing Wax 2 sticks

Tin foil, pure # Ib.

Yarn, white woolen 10 cts.

Zinc, granulated 2 Ib.

" sheet i Ib.

" sulphate

EXPERIMENTAL CHEMISTRY 79

APPARATUS.

ESTIMATE FOR A CLASS OF TWELVE WITH AN ALLOW-
ANCE FOR SOME BREAKAGE.

12 Asbestos sheets i in. thick, 4x4 in.

1 Air-bath (drying oven.)

2 Bell jars, $ gal.
1 Balance, rough.

1 Balance, delicate, with metric weights.
12 Blowpipes, plain with moisture bulb.

2 doz. Beakers, 3 to 7 oz.
2 Burettes 100CC each.

2 pieces Blue Glass, 4x4 in.

1 Barometer.

6 Calcium Chloride U-tubes, plain, 6 in.
10 doz. corks, best quality, Nos. 3 to 10.

2 doz. corks, 1 to 2 in. diameter.

2 Condensers, 18 in. long.
1 doz. Clamps, small.

4 Calorimeters, metal.

3 Crystallizing dishes, 8 in. diam.
10 Deflagrating spoons, -J in. bowl.
12 Delivery tubes (made by class).

4 Desiccators, 4 in. diam.

1 Electric battery, 2 cells.

1 Electrolysis apparatus, simple form.

2 doz. Evaporating dishes, 2 and 3 oz.
1-i doz. Flasks, flat bottom, 12 oz.

^ doz. Flasks, flat bottom 8 oz.

1 doz. Funnels 3-| in. diam.

2 doz. Funnel tubes, small, 10 in. long.
4 doz. Files, triangular, 6 in. long.

80 EXPERIMENTAL CHEMISTRY

3 Files, rat tail, 6 in. long.

1 doz. Forceps, common steel, 4 in. long.

2 Graduated cylinders, 50CC and 200CC.

2 Gasoline Lab. lamps, Dangler's. (Not required if gas

is used.)
1-J doz. Glass stirring rods, solid glass.

2 Ib. Glass tubing, hard and soft, i in. diam.

1 Ib. Glass tubing, soft, ^ and -^ in. diam.

4 doz. Glass plates, 3x3 in.

3 nests Hessian crucibles, 4 in nest, small 5s.
-£ doz. Iron dishes, 4 in. diam.

2 doz. Ignition tubes, 4 and 6 in. length.

\ doz. Iron water (or gas) pipes 50cm x 2cm diam.
6 Metric rulers.

2 Mortars and pestles, 1 iron 2 pt., 1 porcelain 4^ in.
diam.

2 Magnets, horse shoe, 4 in long.

1 Magnifying glass.

% doz. Porcelain crucibles, 1-J in. diam.

3 Porous earthen cups, 2 in. diam., 4 in. long.

2 Pipettes, bulb.

Pneumatic troughs (movable shelf for each basin).
1^ doz. Pinch-cocks, Mohr's.

12 Rubber corks, two preforations, for 12 oz. flasks.
40 ft. Rubber tubing, 20 ft. i in.; 10 ft. -J in; 10 ft. — in.

3 Retorts, glass, 4 oz. Bohemian glass.
1 Retort, copper, 1 qt.

12 Supports, iron stands with two rings each.
12 Spirit lamps, glass globe shape; or 12 Bunsen burn-
ers if gas is used.

3 Sand baths, 4 in. diam.

1 Spatula, 5 in. blade.

1 set cork borers, 6 in a set.

EXPERIMENTAL CHEMISTRY 81

1 horn spoon.

10 doz. Test tubes; two doz. each of 4, 5, 6, 7 and 8 in.

2 Thermometers, 200 degrees Centigrade.
\ doz. Triangles, wire pipe stem covered.

1 doz. Test Tube racks, for 13 tubes, with drying pins.

\ doz. Test tube brushes, with sponge on end.

1 doz. Watch glasses, 2^ in diam.

1 doz. Wolff's bottles, two necks, 4 oz. each.

1 doz. Wire gauze 5 in. square.

INDIVIDUAL APPARATUS

Each pupil should be provided with the following
apparatus selected from the preceding list:
1 Iron stand support with two rings and one clamp.
1 Test tube rack.
6 Test tubes.
1 Blowpipe.
1 Stirring rod.
1 piece asbestos
1 Piece wire gauze.
1 Watch crystal

1 ft. 6 in. Rubber tubing.

2 short pieces rubber tubing.

1 Florence flask with rubber cork.
1 Funnel tube.
1 Pinch-cock.
1 Pair forceps.

3 Glass plates.

4 Wide-mouthed bottles or cylinders (furnished by pupil) .
1 Evaporating dish.

1 Beaker.
1 Glass plug.

A few pieces of red and blue litmus paper.
1 Alcohol lamp or Bunsen burner.

82 EXPERIMENTAL CHEMISTRY

PREPARATION OF SOLUTIONS.

(a) The common acids (sulphuric, nitric and hydro-
chloric), and ammonium hydrate purchased from dealers
are concentrated. For ordinary use, unless otherwise
indicated, they should be diluted by mixing with four
times their volume of pure water. Small bottles (say,
4 oz. ) of these reagents should be kept on the shelf over
the student's table.

Great care should be taken in working with the con-
centrated liquids, as they attack the skin and often cause
painful burns or ulcers. For this reason it is well for the
teacher to prepare large bottles of the dilute solutions for
the student to use in filling his reagent bottles. Small
bottles of the concentrated acids should also be filled for
the student's use.

(b) The following solutions are called for in the
course. They should be prepared, filtered into bottles,
and placed on the shelves for use when required. Use
30 to '50 grams in about 500CC water:

Ammonium chloride, NH4C1.

Alum, potash, KA1 (SOJ2.

Alum, ammonium, (NH4)Al(SO4).r

Alum, chrome, KCr(SO4)2.

Arsenic chloride, AsCl3. (1).

Antimony chloride, SbCl3. (2).

Barium nitrate, Ba(No3)2.

Barium chloride, BaCl2.

Copper sulphate, CuSo4.

Cobaltous chloride, CoCl .

Calcium chloride, CaCl2.

Calcium hydrate, (lime-water) Ca(OH)2.

Ferric chloride, FeCl3.

EXPERIMENTAL CHEMISTRY 83

Ferrous sulphate, FeSo^. (3).
Lead nitrate, Pb(NO3)2.
Lead acetate, Pb(C2H3O,)2.
Mercuric chloride, HgCl2.
Mercurous nitrate, HgNO3. (4).
Magnesium sulphate, MgSO4.
Potassium hydrate, KOH.
Potassium bichromate, K2Cr207.
Potassium permanganate, KMnO4.
Potassium Ferrocyanide, KtFe(CN)(i.
Potassium Cyanide, KCN.
Potassium chromate, K2CrO4.
Potassium bromide KBr.
Potassium iodide, KI.
Potassium nitrate, KNO3.
Potassium chlorate, KC1O3.
Sodium chloride, NaCl.
Sodium nitrate, NaNO3.
Sodium hydrate, NaOH.
Sodium carbonate. Na,CO3.
Sodium sulphide, Na,S.
Strontium nitrate, Sr(NO3)2,
Stannous chloride, SnCl,. (5).
Zinc sulphate, ZnSO4.

(1) Treat the solid with a small amount of water;
with much water it is changed into the oxide and hydro-
chloric acid.

(2) Called butter of antimony. Acidify the water
with a considerable amount of hydrochloric acid. It may
also be made by treating the metal antimony with hydro-
chloric acid containing a little nitric acid.

(3) Ferrous sulphate oxidizes when left in solution;
it should be prepared as needed.

84 EXPERIMENTAL CHEMISTRY

(4) Dissolve mercurous nitrate crystals in water
acidified with nitric acid. It may be made by treating
mercury with a small amount of dilute nitric acid.

(5) Acidify water with considerable . hydrochloric
acid. Pieces of metallic tin should be kept in the bottle
to prevent oxidation to some extent. Fresh solutions
should be prepared occasionally.

Hydrogen sulphide, H,S. See method of preparation,
Experiment 31.

Cochineal solution. Pulverize 4 or 5 grams of cochi-
neal, add water, allow to stand for some time; then stir

well and filter.

Iodine solution. Dissolve the solid in alcohol.

Indigo solution. Slowly add 2 or 3 grams of powdered
indigo to about 10CC concentrated sulphuric acid in an
evaporating dish. Cover the dish and allow it to stand
two or three days; and then add 500CC water, stir well,
and filter.

Litmus solution. Pulverize litmus cubes and add
water; allow to stand for some time, stir well and filter.

EXPERIMENTAL CHEMISTRY 85

ADDRESSES OF SUPPLY HOUSES.

Catalogues, from the following houses have been
received by the writer: —

JOHN TAYLOR & Co.,

Corner First and Mission Sts.,
San Francisco, Cal.

Also Agent for ZIEGLER ELECTRIC Co.,
Boston, Mass.

ALFRED L. ROBBINS Co.,

149 and 151 East Huron St ,

Chicago, 111.
BAUCH & LOME OPTICAL Co.,

515-543 North St. Paul St.,

Rochester, N. Y.

THE CHICAGO LABORATORY SUPPLY & SCALE Co.,
31-45 Randolph St.,
Chicago, 111.

HENRY HEIL CHEMICAL COMPANY,
208-212 South Fourth St.,

St. Louis, Mo.

F. A. BECKETT, Agent for HENRY HEIL Co.,
220 Sutter St.,

San Francisco, Cal.

Supplies may also be purchased of
SALE & SON'S,

220 South Spring St.,

Los Angeles, Cal

The author has had occasion to deal with John
Taylor & Co., San Francisco; and with Alfred L.
Robbins Co., Chicago. He has found both houses to
be thoroughly reliable.

OF TMC

UNIVERSITY OF CALIFORNIA LIBRARY

THIS BOOK IS DUE ON THE LAST DATE
STAMPED BELOW

NOV ig 19U

OCT 26 1915

r.».ft 20

,,i.t gg 1916

NOV 10 1918
aw 4 1919

WAV 20 ISJj
AU6 4 1919
SEP 8
3EP 11 1919

ySi

SEP 21 1921

FEt ^ 19Z-
16

30m-6,'14

U.C.BERKELEY LIBRARIES

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