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Full text of "Experimental chemistry for high school students"

273 filD 



XPERIMENTAL 
CHEMISTRY 



BY WALTER MONROE 




J&N 61906 



LIBRARY 

OF THE 

UNIVERSITY OF CALIFORNIA. 

GIFT OF" 

^ ^ r 

/^Vw^.. HA ,...iL, 

C/^S5 



EXPERIMENTAL CHEMISTRY 

FOR 

HIGH SCHOOL STUDENTS 



BY 



WALTER MONROE, 

Instructor in Chemistry in the Monrovia High School 




MONROVIA, CALIFORNIA, 

1899 



COPYRIGHTED 1899 

BY 

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. 
5051 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 10 CC 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 5 CC of water. 
Taste the solution. What change has taken place? 
Always give reasons for your answers. Add slowly 5 CC 
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 10 CC 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 100 CC 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 15 cm 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 50 cm 
long and 20 mm 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 30 cm 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 2 cm 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 10 cm , 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 100 CC 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 5 CC 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. 

(c) Treat in a similar manner potassium bichromate 
crystals and compare with (a) and (b). 

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

(c) Put 5 CC 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 5 CC 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 5 CC 
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 fo r 
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 760 mm 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 
C. and 760 mm pressure. 

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

Its volume at 2 C. and760 mm pressure=V-j-~V. 

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

273+tx 
V l 273 )' 

Its volume at t C. andp mm pressure=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 

4.6 mm 21 t , 18.49 mm 

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 C. and 
760 mm pressure; what would be the volume if measured 
at 17 C. and 950 mm pressure? 

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

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

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



22 EXPERIMENTAL CHEMISTRY 

and 760 mm pressure; what volume would the gas occupy 
measured dry at C. and 760 mm 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 10 CC 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 C. and 760 mm pressure. 
Hence calculate the volume of gas liberated by one gram 



EXPERIMENTAL CHEMISTRY 23 

of zinc. One liter (1000 CC ) 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 4 cm 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 20 cm 
long, closed at one end, and from 10 to 15 mm in diameter. 
Bend the closed end to a right angle making the bend 
about 5 cm 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 w r ater rises about 3 cm 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 w r ater. 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 20 CC 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 5 CC of 
indigo solution in a test tube, and warm. 

(c) Put about 20 CC 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 5 CC 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 2 CC 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 3 CC 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 75 CC 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 10 CC 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 50 CC 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 5 CC of indigo solution into a test tube 
and add 5 CC of chlorine water. Do you notice any change? 
Treat in the same way 5 CC 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 50 CC in 
a beaker and place in it about one or two grams of bone- 
black; heat for a short time, shake well and filter. 

(c) Pour 10 or 15 CC of hydrogen sulphide solution 



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,C 2 O 4 , into a flask; add about 50 CC 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? 

(c) Repeat (c) in Experiment 26, using carbon mon- 
oxide instead of 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 5 cm 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, H 2 SO 3 . If left in contact with the air, it 
combines with oxygen and is converted into sulphuric 
acid, H 2 SO 4 . 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 5 CC of each of the 
following solutions: Copper sulphate, CuSO 4 ; lead nitrate, 
Pb(NO 3 ) 2 ; mercuric chloride, HgCl 2 ; arsenic chloride, 
AsCl 3 ; antimony chloride, SbCl 3 ; stannous chloride, 
SnCl 2 . 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 5 CC of a 
solution of barium nitrate. Is a precipitate formed? Mix 
5 CC of the lead nitrate solution with 5 CC 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 15 CC 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 5 CC 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, HNO 3 , and sulphuric, 
H. 2 SOJ. 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, 
H 2 S0 4> H 3 PO,, H 4 Si0 4 . 

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, NH 4 OH, 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, Na 2 CO 3 ; 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 NaHCO 3 . 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 ZnSO 4 . ZnO. 

33. (a) Evaporate to dry ness in a porcelain evap- 
orating dish, lO 00 of dilute hydrochloric acid, note the 
appearance and amount of residue. 

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

(c) Take 10 to lo cc 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 10 CC 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+ 

(c) Rinse the beaker and redetermine with a some- 
what larger quantity of sodium hydrate. 



44 EXPERIMENTAL CHEMISTRY 

(d) Take about 15 CC 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 l cc of the sodium 
hydrate. How many grams of sulphuric acid there are 
in l cc 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 
125 CC 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. 

(c) 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. 

NH 4 OH+CO 2 =HNH 4 CO 8 . 

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. 

HNH 4 CO 3 -fNaCl=HNaCO 3 -fNH 4 Cl. 

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. 

2HNaCO s =Na i CO,-fCO a +H 2 0. 

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. 

(c) Ammonium sulphide. 

Take 50 CC of a comparatively strong solution of 
ammonia, and divide it into two parts of about 25 CC 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 20 CC 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 5 CC of nitric acid. Warm for about 
ten minutes. Part of the sulphur is oxidized to SO 3 . 

S4-2HNO 3 =H 2 SO 4 -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 5 CC of mercuric chloride, HgCl.,, in a test 
tube, and add one or two drops of stannous chloride, 
SnCl 2 ; warm and note the result. Now add a considerable 
amount of stannous chloride, heat and observe the result 
as before. Stannic chloride, SnCl 4 , is formed in each 
case; in the first mercurous chloride, Hg 2 Cl,, and in the 
second mercury is precipitated. 

(1) 2HgCl 2 4SnCl i =Hg 2 Cl 2 -fSnCl 4 . 

(2) HgCl 2 +SnCl 2 =Hg-fSnCl 4 . 

(e) Put a small globule of mercury in a test tube, 
and add 5 CC 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 25 CC of water, or take 5 CC of the prepared 
potassium iodide solution and dilute it. 

Put 3 or 4 CC of the solution into each of three test 
tubes, and add to each about 2 CC 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) Fe 8 Cl 6 +2KI=I 2 H-2KCl-f2FeCl 2 . 

(2) 2CuSO 4 -t-4KI=I 2 -|-2K 2 SO 4 -f-Cu 2 I 2 . 

(3) 4HNO 3 -f 3KI=I S -+ 3KNO 3 +NO+2H 2 O. 

(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? 

HgCl 2 +2KI=2KCl-fHgI 2 . 

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) H 2 S+2I=2HI+S. 

(2) 2SnCl 2 4-4I=SnCl 4 -fSnI 4 . 

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

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 50 CC 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 5 CC 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 10 CC of hydrogen sulphide solution; warm slightly. 
Fe 2 (SOJ 3 H-H 2 S==2FeSO 4 -4-H.,SO 4 -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 5 CC 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 100 CC 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 10 CC 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 10 CC . 

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 
100 CC 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 5 CC 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 4 CC 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 5 CC 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? 

(c) Heat a little mercuric oxide in a closed tube. 
What change in appearance do you notice? Explain the 
change. 

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 2 CC 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 5 CC 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? 

(c) Cut a long, narrow strip of zinc, about 10 cm 



64 EXPERIMENTAL CHEMISTRY 

by 3 mm , 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 3 cm by 3 mm ), 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 5 CC 
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. 

(c) Add hydrochloric acid to a little quicklime. 
What takes place? Write the equation. 

(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 5 CC 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. 

(c) What is formed when nitric acid acts upon 
copper? Review Experiment 20. 

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? 

(c) Test the solubility, in water, of other carbonates. 

(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 
K 2 SO 4 .A1 2 (SO 4 ) 8 .24H 2 O. 

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. 

(c) Having found the specific heats of the metals, 
calulate from the formula given their respective atomic 
weights. 

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 y z 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 y 2 Ib. 

flowers y 2 Ib. 

Sulphuric acid, cone 9 Ib. 

Stannous chloride i oz. 

Starch y 2 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 100 CC 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, 50 CC and 200 CC . 

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 50 cm x 2 cm 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 500 CC water: 

Ammonium chloride, NH 4 C1. 

Alum, potash, KA1 (SOJ 2 . 

Alum, ammonium, (NH 4 )Al(SO 4 ). r 

Alum, chrome, KCr(SO 4 ) 2 . 

Arsenic chloride, AsCl 3 . (1). 

Antimony chloride, SbCl 3 . (2). 

Barium nitrate, Ba(No 3 ) 2 . 

Barium chloride, BaCl 2 . 

Copper sulphate, CuSo 4 . 

Cobaltous chloride, CoCl . 

Calcium chloride, CaCl 2 . 

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

Ferric chloride, FeCl 3 . 



EXPERIMENTAL CHEMISTRY 83 

Ferrous sulphate, FeSo^. (3). 
Lead nitrate, Pb(NO 3 ) 2 . 
Lead acetate, Pb(C 2 H 3 O,) 2 . 
Mercuric chloride, HgCl 2 . 
Mercurous nitrate, HgNO 3 . (4). 
Magnesium sulphate, MgSO 4 . 
Potassium hydrate, KOH. 
Potassium bichromate, K 2 Cr 2 7 . 
Potassium permanganate, KMnO 4 . 
Potassium Ferrocyanide, K t Fe(CN) (i . 
Potassium Cyanide, KCN. 
Potassium chromate, K 2 CrO 4 . 
Potassium bromide KBr. 
Potassium iodide, KI. 
Potassium nitrate, KNO 3 . 
Potassium chlorate, KC1O 3 . 
Sodium chloride, NaCl. 
Sodium nitrate, NaNO 3 . 
Sodium hydrate, NaOH. 
Sodium carbonate. Na,CO 3 . 
Sodium sulphide, Na,S. 
Strontium nitrate, Sr(NO 3 ) 2 , 
Stannous chloride, SnCl,. (5). 
Zinc sulphate, ZnSO 4 . 

(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. 

(c) Ammonium sulphide, (NH 4 ) 2 S. See Experiment 
39 (c) for method of preparation. 

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 10 CC concentrated sulphuric acid in an 
evaporating dish. Cover the dish and allow it to stand 
two or three days; and then add 500 CC 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 



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