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i'^
National Library
of Canada
Bibliotheque nationale
du Canada
HIGH SCHOOL CHEMISTRY.
BY
A. P. KNIGHT, M.A, M.D.,
Queen's University, Kingston,
AND
W. S. ELLIS, B.A., B.Sc,
Collegiate Institute, Kingston.
^ttthorUeb bg the ^e^mrtmcnt of (Ebucatwii for ©nhmo.
TOROI^TO:
THE COPP, CLARK COMPANY, LIMITED.
,/ /
Entered according to Act of the Parliament of Canada, in the year one thousand
eight hundred and ninety-five, by The Copp, CiiARK Compant, Limited, Toronto,
Ontario, in the Office of the Minister of Agriculture.
■ — ■'^ \ \ O *"
PREFACE.
This book is intended to aid the teacher, not to
supersede him. For this reason particular care has been
tal<en to arrange the work in such a way that the student
can learn only by doing the experiments under intelli-
gent guidance. Except in rare cases the wording of an
experiment or question gives no hint as to the result
required.
Questions obviously arising out of the experiments, or
such as a teacher would probably ask, for the purpose of
developing his lesson, have been designedly omitted.
Under the heading " Additional Exercises " a con-
siderable number of experiments have been inserted for
the purpose of furnishing material (i) for a variety in the
work from year to year, (2) to enable teachers to adapt
their courses somewhat to the equipment of their schools,
(3) to provide work for students who get along faster
than the class, or who are going over their chemistry a
second time.
It is assumed that students, before beginning chem-
istry, have a knowledge of elementary physics.
In an appendix will be found the syllabus of the Pass
Matriculation course in chemistry for the next three
years ; also a list of books of reference suitable for this
portion of high school work. The contents of this book
are necessarily confined to the subjects there laid down,
and it is intended that the work be done in one year.
Kingston, 1895.
CONTENTS.
CHAPTER I.
Physical and Chemical Change
CHAPTER II.
Conditions that Promote Chemical Change
CHAPTER III.
Theory of Chemical Action
CHAPTER IV.
Elements
CHAPTER VII.
Chemical Notation.
CHAPTER VIII.
Oxygen
CHAPTER IX.
Hydrogen Dioxide,
Nascent Condition.
CHAPTER X.
rvi
Paob.
8
13
17
CHAPTER V.
To Find out if Water is an Element 22
CHAPTER VI.
Preparation and Properties of Hydrogen 23
37
41
50
53
^ CONTENTS.
CHAPTER XT.
Paoii.
Acids, Bases and Salts . . 55
CHAPTER XTI.
Chemical Nomenclature 81
CHAPTER XITI.
Valency or Atomicity 53
CHAPTER XIV.
Synthesis of Water gg
CHAPTER XV.
Definite Proportions 71
CHAPTER XVI.
Some Chemical Calculations 74
CHAPTER XVII.
Combustion "jy
CHAPTER XVIII.
The Atmosphere gQ
CHAPTER XIX.
Nitrogen, its Properties and Preparation 85
CHAPTER XX.
Compounds of TJitrogen and Oxygen 87
CHAPTER XXI.
Acids of Nitrogen 96
CHAPTER XXII.
Avogadro's Law 103
CONTENTS. vii
CHAPTEH XXIII.
Paob.
Nitrogen and Hydrogen 1Q5
CHAPTER XXIV.
Percentage Composition and Formulae 115
CHAPTER XXV.
Carbon , 122
CHAPTER XXVI.
Carbon and Oxygen j25
CHAPTER XXVII.
Relation between Volumes of Constituents and Volume of Com-
pound formed, when all are Gases 136
CHAPTER XXVIII.
Carbon and Hydrogen 1 jo
CHAPTER XXIX.
Coal Gas and Flame I^w
CHAPTER XXX.
Chlorine I gg
CHAPTER XXXL
Chlorine and Hydrogen igo
CHAPTER XXXII.
Some Compounds of Chlorine Igo
CHAPTER XXXIII.
Sulphur J--
CHAPTER XXXIV.
Oxides of Sulphur I»g
Vlil CONTENTS.
CHAPTER XXX '.
Paoi.
Acids of Sulphur . . , 182
CHAPTER XXXVI.
Calculation of Formula) 188
CHAPTER XXXVII.
Impurities in Air and Water . . 189
•
CHAPTER XXXVIII.
Molecules of Elements usually consist of more than one Atom .... 193
CHAPTER XXXIX.
Selected Questions 198
APPENDIX.
EXPERIMENTAL CHEMISTRY
CHAPTER I.
The object of this chapter is to illustrate the various
kinds of changes and apparent changes which matter
may undergo.
1.— Physical Change
Experiments.
1. Heat a piece of platinum
wire in the flame of a spirit
lamp or gas burner.
2. Fill a test-tube full of
water and invert it over a
tumbler or beaker half-full of
water, as in Fig. i. Heat the
upper end with a spirit lamp
until steam forms, and then
allow the test-tube to cool.
Fia. 1.
3. Hold a rod of hard glass in a gas flame until it
becomes red hot, then pull the ends of the rod apart.
After cooling can you observe any change in the pro-
perties of the glass at the place where it was heated ?
[1]
2
MIXTURE.
2.— Chemical Change.
Experiments.
1. Treat a piece of magnesium wire as you did the
platinum of ex. i, sect. i.
2. Lay a bit of phosphorus about half as large as a
pea on a piece of sheet zinc, set this on a plate, touch it
with a hot wire, and then cover it with a bell jar.
3. Hold a splinter of wood or a piece of sealing wax
in the gas flame until it takes fire.
3.— Questions and Exercises.
1. In sections i and 2, compare similady numbered experiments
and state what you observed in each case.
2. By comparison of experiments, as indicated in the last ques-
tion, try to determine in what cases, only some of the properties of
the substances acted on were changed, and in what cases entirely
new substances were formed.
3. Mix a little sulphuric acid with twice its volume of water
and divide the mixture into two parts, pouring each into a test-tube.
Into one drop a piece of platinum wire, into the other a piece of
magnesium wire, then hold a lighted match to the mouth of each
tube. Is there a chemical change in either case? Upon what
evidence is your answer based ?
4. In sect. I, the changes were physical, and in sect. 2,
chemical ; from this express in your own words what these con-
sist of and what is the difference between them.
4.— Mixture and Combination.
Experiments.
I Prepare some fine iron filings (those formed when
saws are sharpened answer well) ; to these add about
MIXTURE AND COMBINATION.
double their weight of powdered sulphur, and shake on a
piece of paper until the two are thoroughly intermingled,
then examine with a magnifying glass. Draw a magnet
or a magnetized knife blade a number of times through
the mixture. Divide the mixture into two parts, on one
drop some hydrochloric acid and carefully smell the gas
that comes off. Heat the other part in a test-tube or
small crucible until it glows; after cooling, again examine
it with glass and magnet, and then drop on it some
hydrochloric acid and notice the odor of the gas formed.
2. Repeat this experiment but use copper filings in-
stead of iron.
3. Mix some sand and common table salt (sodium
chloride), both dry, after stirring them together examine
with a magnifying glass. Pour water on the mixture,
stirring it meantime, and when it has stood for a few
minutes filter it; evaporate the filtrate (that which passes
through the paper), also collect what remains on the
paper and dry it.
4. Make a mixture of powdered sal ammoniac and
charcoal, then heat over a water or sand bath for some
time.
5.— Questions.
I. In ex. I, sect. 4, were the iron and the sulphur distinguishable
at any time after they were put together ? Would your answer be
appUcable during the whole treatment which they received ? What
evidence is there that a new substance was formed ?
A mixture (or mechanical mixture as it is some-
times called) is an intermingling of masses, even though
very minute, of two or more substances, but each retains
its own properties and identity.
SOLUTION.
6.— Solution.
Experiments.
1. In a small beaker or test-tube place a little salt,
then pour water on it until the vessel is nearly full ; after
standing awhile, taste the liquid by lifting a drop on the
end of a piece of glass rod and placing it on the tongue.
2. Repeat this experiment using sugar instead of salt.
Have the salt and sugar vanished ?
3. Mix, in a test-tube, i part of sulphuric acid with 20
parts of water. Taste the mixture, and say whether the
acid has dissolved in the water. Repeat this experiment
using equal parts of alcohol and water.
4. Fit a test-tube with a good cork and delivery
tube, as in Fig 2. Put in this tube about 2 cc. of
spirits of hartshorn and invert a large dry test-tube
over the mouth of the delivery tube; heat the harts-
horn, and when the smell of ammonia is plainly
discernible about the inverted tube, carefully place
this tube, still inverted, with its mouth under
Fio. 2. water and let it stand there for a few minutes.
Explanation. — If there is a doubt as to whether a
solid is soluble in a liquid, shake the two together, and
after the solid particles settle to the bottom, lift out two
or three drops of the clear liquid and gently evaporate
this on a strip of clean glass, a bit of mica sheet, or a
piece of platinum foil. If no trace of sediment is left the
solid is insoluble, but if any sediment appears on the slip,
it must have come from the clear fluid ; hence some of
the solid was dissolved in it.
When a substance disappears in a liquid, as in the case
of salt or sugar in water, it is said to dissolve, the liquid
SOLUTION. 6
in which it disappears is called a solvent, and the result-
ing fluid a solution. The solution is saturated when
the liquid does not take up any more of the substance ;
that is, when the liquid does not undergo any further
change in presence of the substance. A solution is
concentrated when it contains a relatively large por-
tion of the dissolved substance, and it is dilute when
the liquid is present in considerable excess.
The solution of soluble substances is hastened by
(i) powdering the solid, (2) stirring it and the solvent
together, (3) suspending it in the liquid, (4) heating the
liquid. From the study of these methods it will readily
be seen that solution is entirely a surface operation.
7.— Questions and Exercises.
1. Fill a small test-tube to the depth of 4 c. with water, and then
add chloroform, or sulphuric ether, to the depth of i c. Shake well,
and allow the mixture to stand for a few minutes, after which,
examine it carefully.
2. Make a solution of iodine in iodide of potash solution ; dilute
until the colour is a seal brown, then add a few drops of chloroform
and shake the two together. What does this show about the
relative solubility of iodine in water and in chloroform ? Carefully
pour off the water, then turn the chloroform into an evaporating
dish and let it stand for half an hour. Does this experiment give you
any hint as to ho»v iodine may be separated from water? Would
chloroform act in the same way on an alcoholic solution of iodine ?
3. Pour about 2 cc. each of water, alcohol, sulphuric ether, chlo-
roform, and carbon bisulphide into separate small test-tubes, then
into each let fall a couple of drops of oil and shake well. Can you
form any opinion as to the solubiUty of oil in these substances?
4. Where does a solid go to when dissolved ?
6
SOLUTION.
■ !
i
!i^
5. How may a solid be obtained from a liquid in which it is
dissolved ?
6. Devise means of separating salt from sand, sugar from char-
coal, iodine from iron filings ; in each case saving both substances.
7. Perform the following operations and decide whether the.
change is a physical or chemical one in each case : —
(i) Heat, on a piece of mica, some white sugar until it blackens.
(2) Put some sugar in an evaporating dish and pour a little
strong sulphuric acid on it, after standing for a few
minutes, wash it well by turning a very light stream of
water on it. Taste it.
(3) Hold the end of a strip of zinc in the gas flame until it melts.
(4) Take another piece of this zinc, put it in a test-tube and
pour over it a little sulphuric acid dilpted with twice its
volume of water; after all bubbles cease to rise, evaporate
the liquid.
8. How do you account for the bubbles of gas rising out of a
bottle of soda water or of ginger ale when the stopper is removed?
Pour a little soda water or ginger ale out of a newly-opened bottle
into a test-tube and warm it.
9. Scatter some iron filings over a sheet of paper, then bring a
magnet near the under side of the paper, and move it about beneath
the filings. Is the change produced in the filings a physical or a
chemical one ?
8.— Additional Exercises.
1. Prepare some powdered nitrate of potash and some powdered
charcoal, shake them together on a piece of paper, then divide into
two parts, and heat one of these on a piece of mica. Compare the
result with the other part of the substance that was on the paper.
Have you a mixture, a compound, or both ?
2. Take an ounce each of powdered alum, washing soda, and
copper sulphate, and dissolve them separately in a fluid ounce of
water each. Stir with a glass rod, and observe the extent to
SOLUTION. 7
which each one dissolves. Now heat the solutions until they boil,
stir again, and notice whether there is any variation in solubility.
3. Carefully counterbalance on a scale a small beaker or dish
containing a couple of grams of bicarbonate of soda, then add water
little by little until the white solid disappears. Evaporate over a
sand bath, or in a water bath, until the liquid all disappears. After
thoroughly drying and cooling, compare the weight of the vessel
and its contents with what it was at first.
4. Try, judging by the colour, if iodine is soluble in water, alcohol,
solution of iodide of potash, ether, chloroform or bisulphide of
carbon.
5. Try if nitrate of potash, red lead, bichloride of mercury
(corrosive sublimate), sulphur and charcoal are soluble in water.
6. A blacksmith burns coal in his forge that he may make the
iron hot and more easily worked ; how would you classify the
changes in the coal and iron ? Can you give any other examples
of these changes from mechanical or domestic operations ?
7. Gunpowder consists of charcoal, sulphur and nitre. Separate
these substances in a given sample of gunpowder.
8. On some copper filings pour a little nitric acid, and gently warm;
after the filings have disappeared, evaporate the substance left. Is
the solid obtained soluble in water .■* How would you classify
the action between the copper and the acid .-'
9. Tell how you would find out how much sand there is in 100
grams of the augar of which your teacher will supply you with a
sample. If the mixture had contained some copper filings also, how
would you proceed ?
10. How does the weight of a solution compare with the weights
of the solid and liquid which form it ?
11. Compare the density, color, taste and smell of a solution
with the substances that enter into it.
8
CONDITIONS THAT PROMOTE CHEMICAL CHANOB.
CHAPTER II.
l! '
:.i
l.—Oonditions that Promote Ohemical Change.
In some cases substances will combine chemically if
simply mixed, but generally it is necessary to resort to
some special treatment in order to bring about chemical
combination; similarly,compounds sometimes decompose
spontaneously, but in most cases, the breaking up of a
compound into its constituents results from methods
adopted for that purpose.
The conditions that tend to promote chemical action
are generally (i) either simply mixing, rubbing together,
or dissolving the constituents, (2) exposing to higher tem-
perature, (3) using electrical energy, (4) exposing to light,
(5) using vital energy. Of these, solution and change
of temperature are the ones more commonly employed.
'i
2.— Intimate Mixtm*e.
The production of chemical combination is often de-
pendent on the bringing of the minute parts of the con-
stituents into very near contact with each other. The
methods adopted for accomplishing this result are gener-
ally (i) stirring the substances together, (2) rubbing or
pounding them together, (3) mixing solutions of them.
Experiments.
I. Wet the inside of a slightly warmed glass beaker
with strong aqua ammoniac, and the inside of another
beaker with a strong solution of hydrochloric acid ^ cover
i
INTIMATE MIXTURE.
9
the first beaker with a glass plate and invert over the
second. Then draw out the plate.
2. Cut a thin slice from the end of a stick of phos-
phorus.* Dry it well and place on a plate, then sprinkle
over it a little powdered iodine. Cover with a wide-
mouthed bottle.
3. In a small dish or test-tube put a piece of freshly
cut phosphorus and cover it with a strong solution of
silver nitrate ; let it stand for 24 hours.
4. Powder a little chlorate of potash and dry it well
on a warm glass or on mica, then mix with it half its
own bulk of sulphur, by shaking them together on a
piece of paper (they must not be stirred or rubbed).
Place a little of the mixture on some hard object,
such as a smooth stone or an iron plate ; either strike
this mixture with a hammer or rub it with a large pestle.
This experiment is dangerous unless the directions are
followed.
5. Mix a teaspoonful of baking soda (sodium bi-
carbonate, NaHCOg) and half as much oxalic acid,
H2C2O4, in a large test-tube ; shake them well together,
then pour in a little water. Vary this experiment by
mixing the two substances when dry; also dissolve por-
tions of them separately and mix the solutions.
6. Dissolve a few crystals of iodide of potassium, KI,
and of lead acetate, Pb(C2H302)2, in separate test-tubes,
then mix the solutions.
♦Always cut phosphorus under water, and always hold it witk a pair of forceps
— never in the finger^.
10
HEAT. — LIGHT.
> 'i
II '
I '■■)
3.— Heat.
Substances when heated will often undergo chemical
change, both of combination and decomposition, while at
ordinary temperatures they are chemically inert.
Experiments.
I Place about half an inch in depth of chlorate of
potash, KClOg, in a test-tube and heat it strongly until it
melts, and bubbles of gas begin to come off, then hold in
the mouth of the tube a glowing splinter. After the
heating has been continued for about five minutes, allow
the tube and its contents to cool, then dissolve the solid
residue. At the same time make a solution of some of
the original chlorate ; into each, drop a little silver nitrate
solution.
2. Heat some red oxide of mercury, HgO, in a test-
tube and hold a glowing splinter in the mouth of the
tube.
3. Make a mixture of some powdered chlorate of pot-
ash and white sugar, put a little of this on a piece of mica
and heat it. Vary the experiment by letting fall a drop
or two of sulphuric acid from the end of a glass rod on a
portion of the mixture.
4.— Light.
Experiments.
1 . Moisten a piece of paper with nitrate of silver solu-
tion, then lay on this paper a leaf of a plant, cover the
whole with a piece of glass and expose to sunlight.
2. Repeat experiment i, but use bichromate of potash
solution instead of silver nitrate.
ill.
ELECTRICITY.
11
3. To some dilute solution of silver nitrate add some
solution of common salt, allow the precipitate which will
be formed to stand in sunlight for a time.
/, A precipitate (p'p'te.) is a solid substance formed in a
liquid or a mixture of liquids, and is consequently in-
soluble in the fluid in which it is formed.
5.— Electricity.
I. Pass a current of electricity through water in a
decomposition -of- water apparatus, as in Fig. 3. This will
require a current from about 4 bichromate cells " in
series." (See High School Physical Science).
ZINC
FiQ, 3.
6.— Vital Force.
Experiment.
I. Make a weak solution of sugar in water and add to
it a little yeast powder. Let this stand for a few days
in a warm place. Taste the liquid.
7.— Questions and Exercises
I. Make a mixture of powdered sulphate of iron (copperas),
FeS04, and ferrocyanide of potash (yellow prussiate of potash),
12
ADDITIONAL BXERCISES.
K4FeCye. Drop some water on the mixture. What does this
show?
2. A match may be lit by rubbing it against a rough surface, or
by holding it against a hot object. Give reasons why.
3. liy what agents are the chemical actions promoted in the fol-
lowing cases : — printing a photograph, electroplating a spoon, com-
bustion of coal in a furnace.
4. Stir the parts of a Seidlitz powder together when dry, then
throw the mixture into a large beaker half-full of water. What
does the result prove ?
5. Write your name on a sheet of white paper with a solution of
sulphate of iron (copperas), when this is dry dip the paper in solu-
tion of ferrocyanide of potash. Repeat the experiment, but use for
the writing fluid iodide of potash solution, and for the bath, bichlo-
ride of mercury solution. How do you explain the results ?
it
8.— Additional Exercises.
1. Powder some iodide of potash and some bichloride of mercury,
then stir the two together.
2. Rub together in a mortar a drop of mercury and a little iodine.
If a drop or two of alcohol be added to dissolve the iodine, the
combination will be more readily obtained. If the resultant com-
pound is green in colour, add a little more iodine ; if red, add a little
mercury.
3. Dissolve a httle common salt (sodium chloride), NaCl, and in
another test-tube a little silver nitrate, AgNOg ; mix the solutions.
4. Heat to redness a piece of bright copper, also a piece of bright
iron.
5. Make a mixture of equal parts of powdered nitrate of potash,
white sugar and sulphur ; heat a /tWe of this on a piece of mica.
6. Heat some nitrate of potash, KNO3, in a test-tube, until it
melts, then drop into it a piece of charcoal about as large as a pea.
7. Repeat experiment 5, sec. 7, but use silver nitrate solu-
tion for the writing fluid, then dip the paper into some weak
i' !
I I
THEORY or CHEMICAL ACTION.
18
hydrochloric acid, quickly remove it and expose to bright light for
a few minutes. Try the experiment again, and put the paper in a
dark drawer after dipping it in the acid. What agent produces the
result in this case ?
8. Attach a bright piece of iron, three or four inches long, to the
terminal wire connected with the zinc of a three or four celled elec-
tric battery, and then immerse both the iron and the other terminal
wire in a solution of sulphate of copper contained in a glass beaker
or clean wooden trough. After a time attach the iron to the other
terminal and repeat the experiment.
9. Place on a glass plate a few drops of strong solution of silver
nitrate, then place the terminals of the battery in this solution at a
distance of about half an inch from each other and hold them still.
As the dark solid forms, keep the other terminal moved away from
it.
In electrolytic decompositions those substances which
are attracted to the positive electrode (that is the ter-
minal attached to the copper, carbon or platinum plate),
are called electro-negative, and those which appear at
the negative electrode (the terminal connected with the
zinc plate), are electro-positive.
CHAPTER III.
1.— Theory of Chemical Action.
From the experiments of the last chapter, it is quite
evident that two or more substances may have their parts
intimately mixed up with one another, yet each retain
its own identity and have all its properties unchanged.
In other cases, however, it is quite impossible to observe
u
THEORY OP CHEMICAL ACTION.
i
'!i
any traceof either constituent in the resultant substance,
and the distinguishing properties of tlie kinds of matter
that were acted on have been altered so that an entirely
new material has been formed. It becomes necessary
now to give a very brief outline of the theory which
offers an explanation of this phenomenon. At the same
time, it is well to warn the student that it is only a theory
which at present cannot be proved; but it affords a
reasonable and consistent explanation of a marvellous
number of observed facts, and accounts very generally
for the phenomena of chemistry, so that in all proba-
bility it is the true theory.
From the observation of both physical and chemical
action, it is reasonably certain that all matter, of every
state and condition, is made up of separated particles,
very minute, indivisible by physical means, yet exist-
ing as individual portions. Such particles are called
/- molecules. There are also good reasons for believing
that these molecules are in rapid vibration, sometimes
moving freely among one another, sometimes so con-
fined that their vibrations may not carry them outside of
a limited space. (See High School Physical Science.)
2. In chemistry we have two kinds of matter to deal
with — elementary and compound. When the parts that
go to make up the molecules are all of one kind, that
matter is said to be elementary, because when divided or
broken up as much as possible it yields only the one
substance. If, however, the individual parts of the
molecules are dissimilar, the substance is said to be a
compound, because when properly divided up it yields
matter of different kinds.
1^
ATOMS.
15
The portions of matter that go to form a molecule are
called atoms.
Chemical theory further supposes that when combina-
tion takes place the atoms of one element join with the
atoms of one or more other elements to form groups of
atoms all exactly alike in composition ; thus when the
magnesium (Chap. I., ex. i,sec. 2) burned, an atom of the
metal united with an atom of oxygen, one of the gases of
the atmosphere, to form a group of two atoms, i.e., a mole-
cule, of oxide of magnesium, which is the chemical name
of the white ash produced.
Sulphuric acid consists of two atoms of hydrogen, one
of sulphur and four of oxygen (these three substances
are elements). Now in Chap. I., ex. 3, sec. 3, the chemi-
cal action consists in an atom of magnesium crowd-
ing out the two atoms of hydrogen that are in every
molecule of the acid ; this would manifestly give rise
to a molecule different from that of the acid. The
hydrogen is a gas, and it was the crowded-out atoms of
this element that had congregated into masses, and
formed the bubbles of gas that rose to the surface. Had
the remaining water been evaporated, a white salt would
have been found ; this would have been the matter made
up of the new molecules each composed of an atom of
magnesium, one of sulphur and three of oxygen.
3. From what has been said, these statements follow : —
(i) An atom is the smallest part of an element
that can enter into the composition of a mole-
cule ; hence, that can take part in chemical
action.
16
CHEHldlt.
*
I) VI I
IM
!|-
(2) A molecule is the smallest part of a substance,
whether elementary or compound, that can
have a separate existence.
(3) Molecules are made up of atoms, and, for the
same kind of matter, their composition is con-
stant, that is, in all molecules of the same
chemical substance there are equal numbers of
the same kinds of atoms.
(4) If chemical combination is the union of atoms
to form molecules, then decomposition must
consist not in the separation of molecules from
one another, but in the breaking of them up
into atoms or into groups of atoms.
4. Ghemism. — When masses of the same kind join
together to form a single mass it is said that they cohere,
or that they are held together by cohesion. When the
suostances that join togethor are of different kinds, they
are said to adhere. When, however, atoms join to-
gether to form molecules, a new force comes into play,
which is known as chemism, or chemical aflBnity.
This differs from both cohesion and adhesion, because it
is capable of acting through only infinitesimally short
spaces, such as those which separate the molecules of a
substance.
We have learnt in the preceding chapter that in very
many instances substances will not act chemically among
one another, no matter how finely they may be powdered
or how intimately mixed, until means are employed to
bring the molecules into still closer contact.
Chemical affinity is not equally strong among all
substances. Hydrogen and chlorine can scarcely be pre-
]t I
EL£MENT8.
17
vented from combining if mixed, while hydrogen and
nitrogen can be made to unite only with the greatest
difficulty. Compounds of chlorine and nitrogen, obtained
by decomposition of other substances, are held so loosely
in union that they are liable to break up with violent
explosions, while compounds of chlorine and iron (as
well as most other metals), are very stable, that is, are
not easily decomposed.
References to chemical affinity may be found: Remsen, 12; Lodge's
Modern Views of Electricity, 2nd Ed,, 83-4 ; Muir & Slater, 175 : Remsen
Theoretical Chem., 14; Tilden, 203; Wurtz, 224, 311.
Atoms and molecules are discussed at length in Ramsay's Chem. Theory
for Beginners, p. 61 ; R. & S., p. 69-80; Wurtz, 33-47, 305-332; Tilden,
3-4,85; Remsen Th. Chem., 17-18,36; Muir & Slater, 203-16 ; Rem-
sen, 68-80.
CHAPTER IV.
1.— Elements.
I. It has been found that by far the greater number
of substances with which chemists have to deal are capa-
ble of being decomposed into simpler ones. There are,
however, about sixty-eight or seventy substances that
have never been so divided ; these are called elemeikts ;
and from them, all kinds of matter have been formed, so
far as we know at present. It is not likely, though, that
these are all the elements, because within late years a
number of new ones have been discovered, and it is
likely that more will be found out in the future.* On the
* Since these lines were written, two Enelish scientists have discovered and
isolated a substance that seems to be quite plentiful in the atmosphere, though its pres-
ence was not suspected until a few months ago. This substance has been named Argon.
As far as investigation has gone at present (Feb'v. i8qs) it may be either a new element
or a mixture of two or more elements hitherto unknown.
2
«'i'
iil
!'M I
ViW-
in;
m : i
li n
18
ELEMENTS.
other hand, it is quite possible that some of those now
treated as elements may be found to be compounds*
when methods of research improve and a more exact
knowledge of the laws of matter is gained.
2. The following list contains the names, symbols and
atomic weights of the sixty-eight elements at present
known ; two doubtful ones are omitted : —
Nahh op Element.
SYMBOIj.
Atomic
Weioht.
Aluminium
Antimony
Arsenic
Barium
Beryllium
Bismuth
Boron
Bromine
Cadmium
Caesium
Calcium
Carbon
Cerium
Chlorine
Chromium
Cobalt
Copper
Diaymium
Erbium
Al
Sb (Stibium) .....
As
Ba . .
Be
Bi
B
Br
Cd
Cs
Ca
C
Ce
CI
Cr
Co
Cu (Cuprum)
D
E
27 3
122
75
137
9
210
11
80
112
133
40
12
141
35-5
52
68-7
63
147
166
Fluorine
Gallium
Germanium
Gold
Hydrogen
Indium
Iodine
Iridium
Iron
Lanthanum
Lead
Lithium
Magnesium
Manganese
Mercury
F
Ga
Ge
Au (Aurum)
H
In
I
Ir
Fe (Ferrum)
La
Pb (Plumbum) ....
Li
Mg
Mn
Hg (Hydrargyrum) . .
19
70
72 32
197
1
113-4
127
193
56
139
207
7
24
55
200
M£TALs — Non-metals.
Id
Nahr of Elsmbnt.
Molybdenum
Nickel . .
Niobium . .
Nitrogen ■
Osmium . .
Oxygen . .
Palladium
Phosphorus
Platinum . .
Potassium
Rhodium .
Rubidium
Ruthenium .
Scandium .
Selenium . .
Silver . .
Silicon . ■
Sodium • .
Strontium
Sulphur
Tantalum . .
Tellurium .
Thallium . .
Thorium . .
Tin ... .
Titanium
Tungsten . .
Uranium
Vanadium
Ytterbium .
Yttrium . .
Zinc . . .
Zirconium
Symbol.
Mo . . . .
Ni . . . .
Nb . . . .
N . . . .
Os
V^ • • • • •
Pd
p
Pt ....
K (Kalium) .
Ro ....
Rb . . . .
Ru . . . .
Sc
Se . . . .
Ag (Argentum)
Si . . . .
Na (Natrium) .
Sr
S
Ta
Te . . . .
Tl
Th . . . .
Sn (Stannum) .
Ti ....
W (Wolfram) .
U . . . .
V
Yt . . . .
Y
Zn . . . .
Zr
Atomic
Wbioiit.
5
1
5
4
5
96
58-7
94
14
199
16
106-5
31
197
39
104
85
103
44
79-5
108
28
23
87-5
32
182
128
203-5
234
118
50
184
240
51
173
89-6
65
89-6
Some of the more important ones are printed in full-faced type.
3. Metals and Non-Metals.— These elements are
classified into metals and non-metals. The former are
characterized by their peculiar appearance (metallic
lustre) and by being good conductors of heat and of
electricity. It must be remembered, however, that there
is no sharp distinction between the two classes. The
elements classed as non-metals are hydrogen, bromine
20
STMB0L8.
)!■ i
ii:ili:
■■t:
chlorine, iodine, fluorine, nitrogen, phosphorus, arsenic,
boron, carbon, silicon, oxygen, sulphur, selenium, tellu-
rium. Arsenic serves as the connecting link between
the metals and non-metals. In appearance, and most of
its physical properties, it is metallic, but chemically it is
a non-metal because it does not form certain compounds
which are characteristic of all true metals.
4. Ssrmbols. — In the column headed " Symbols," cer-
tain letters are placed opposite the names of the ele-
ments. These letters serve as a convenient, short way
of indicating the substance; thus, in chemistry, H. stands
for hydrogen, Ca. for calcium, Hg. for mercury, and so
on through the list. It will be noticed that in most cases
the symbols are formed of the first letters of the names
of the elements ; or, where two or more elements begin
with the same letter, the symbol is formed of the initial
letter joined with one of the other letters that is promi-
nently sounded in the word ; thus, C. stands for carbon,
Ca. fo'* calcium, Cd. for cadmium and Cs. for Caesium.
Generally, the names of the elements are formed after
the manner of Latin nouns in " urn," but in some few
cases a popular name has, in ordinary use, supplanted the
Latin form ; though the symbol is that derived from the
Latin word. Iron, copper, silver, mercury, potassium,
serve as examples of this.
^ 5. Atomic Weights. — By atomic weight of an element
is meant the number of times that an atom of the sub-
stance is heavier than an atom of hydrogen. Of course
it would be absurd to think of weighing out an atom of
any substance and comparing its weight with that of an
atom of hydrogen. The pupil must understand that
these numbers have been derived from the results of a
ATOMIC WEIGHTS.
21
long series of difficult experiments, which he can not
comprehend at the present stage.
Any other element might be adopted as the unit for
atomic weight instead of hydrogen ; and if this were
done the atomic weights of all the other elements would
be relatively changed.
6. —Questions and Exercises.
I. Take a piece of roll sulphur and of iron or copper wire ; hold
them, one in each hand, and dip them into a vessel containing
boiling water. Note the one along which the heat travels quickest
to the hand.
Fio. 4.
2. Insert these same substances, in turn, into the circuit of a
galvanic battery. Attach a galvanometer to the circuit, as in Fig.
4, and by its aid note which substance acts as a conductor of
electricity. A toy compass will do for the galvanometer.
3. Compare the surface appearance of copper, silver, and other
metals with that of sulphur and phosphorus.
4. Connect a galvanometer into a battery circuit, then cut the
circuit and dip the cut ends of the wire into a vessel of mercury.
5. What have you noticed which would aid you in classifying
mercury as a metal or non-metal ?
For basis of theory of elements, see H. S. Physical Science.
For articles on metals and non-metals, see Muir & Slater's Elementary
Chemistry, p. 99; Tilden's Chem. Phil, pp. 244-256; Roscoe & Schor-
lemmer, vol. i., pp. 53-4 ; Richter, p. 253 ; Remsen, Inorganic Chem. , pp.
452-55; D. <fc W., 12.
i
'%'
m
22
COMPOSITION OP WATER.
CHAPTER V.
1.— To Find out if Water is an Element.
Experiments.
Pig. 5.
I. Take a test-tube about
2 centimetres in diameter,
and lo or I2 centimetres in
length. Fill it with acidu-
lated water (i to 6o), and
invert it over a beaker con-
taining water. Under the
mouth of the test-tube, place
the terminal wires of a gal-
vanic battery, as in Fig. 5.
The ends of these wires
should consist of platinum,
and should not touch each
other when placed under the
mouth of the test-tube.
2. When all the water has been expelled by the accu-
mulated gas in the preceding experiment, raise the tube,
keeping it mouth downward, and apply a lighted match
to it.
3. Repeat experiment i, using two test-tubes full of
acidulated water, inverted over a soup plate. Place the
end of a wire under each tube. Each wire must be insu-
lated where it touches the water, except about i centi-
metre at the end. When gas has filled one of the test-
m'^
HYDROGEN.
23
tubes, stop the current, and examme the gases. Put a
glowing splinter of wood into the one with least gas in
it, and apply a lighted taper to the full one.
4. Place the battery terminals in acidulated water
both in the same vessel. When the gas begins to rise
freely, touch the terminals together.
5. Repeat ex. 3, but before applying the splinter and
taper turn each tube mouth upwards for a few seconds.
This process of decomposing a compound by making
it part of an electric circuit, is called the Electrolysis
of it, or the electrolytic decomposition of it.
t\
2.— Questions and Exercises.
1. Could the gas in the tube in ex. i, sect, i, have been produced
by decomposition of the battery terminals .'*
2. Was the gas ordinary air ? What reason have you for your
answer ?
3. How does the mixture of gases differ from each one separately,
when tested with a blazing splinter "i
"A
CHAPTER VI.
1.-— The Preparation and Properties of Hydrogen.
In the preceding chapter it was found that water is a
compound made up of at least two substances, which
physically resemble each other somewhat, but chemically
are quite different. As water is one of the commonest
-m
24
HYDROGEN.
substances known, it has been chosen as a starting point
for the study of the chemical properties of matter. Of
the two gases of which water is composed, the one that
came off in greater quantity, and which burned when
brought into the presence of a flame, is known as
hydrogen. The methods of preparing this gas and the
stud}' of its chief properties will occupy the remainder
of this chapter.
2.
14
.,■•:!
l|,
Experiments.
I. Throw a bit of freshly-cut potassium on some water
on a plate or in a wide dish. Repeat the experiment,
but tinge the water red with litmus solution.
(a) Again repeat the experiment in both ways, but
use sodium instead of potassium.
2. Confine a bit of sodium
or potassium not larger than
a pea in a cage of wire gauze,
and hold it under a tube that
has been inverted full of water
over a dish of water, as in
Fig. 6. When the sodium
has disappeared, another piece
may be put in the cage. (If
large pieces are used a violent
explosion may occur.) When
the tube is filled with gas it
may be lifted and a lighted
taper applied to its mouth.
FlQ. 6.
HYDROGEN.
26
3. Make a mixture of water and strong sulphuric acid
in the proportion of 6 to i. Fill a test-tube with this
mixture, and invert it over a plate containing some of
the same mixture. Below the test-tube, which must
always be kept with its mouth below the level of the acid
and water, place some small pieces of zinc.
3 (a). Vary this experiment by using a piece of mag-
nesium instead of zinc.
Explanation. — Such chemical actions as those be-
tween sodium or potassium and water, zinc and sulphuric
acid, magnesium and sulphuric acid, come under the
class of substitutions in which one or more of the
atoms of a molecule, generally the hydrogen, are dis-
placed by atoms of other elements.
4. Take a wide-mouthed flask and fit it with a good
cork perforated by two glass tubes, one of which passes
nearly to the bottom of the bottle, and has on its upper
end a funnel-like expansion ; the other tube merely
passes through the
cork, is bent at right
angles, and has a rub-
ber tube attached to it
for conveying the gas to
a " pneumatic trough,"
as in Fig. 7. Place
some clippings of zinc
(or better, some granu-
lated zinc, prepared by melting common sheet zinc in
an iron ladle, then pouring it from a height of 3 or 4
feet into a pail of cold water), in the bottle, fill it about
one-third full of water, and then pour down the funnel
Fro. 7.
%
I t J
i.
'"^
r
jil
26
HYDROGEN.
tube about one-tenth as much sulphuric acid. The gas
begins to form quickly, and is collected in bottles pre-
viously filled with water and kept mouth downwards in
the water in the pneumatic trough. Collect two or three
bottles or large test-tubes full of gas and preserve them
for future experiments. Preserve the liquid which re-
mains in the bottle after the gas has ceased to come off.
Filter this liquid and either evaporate it over a lamp
flame, or allow it to stand in an open vessel for a day or
two. Then examine carefully.
5. Keep the mouth of one of the bottles downward,
and plunge a lighted taper upward into it. Then with-
draw the taper slov/ly, allowing the burnt end to remain a
moment or two at the mouth. Note exactly what phe-
nomena occur (i) just as the taper enters the tube,
(2) when the taper is inside the tube, and (3) just as it is
withdrawn.
6. Take the second jar, and quickly turn its mouth
upwards under the mouth of a similar jar filled with air.
Let it remain thus for a few seconds, then apply a burn-
ing taper to the mouth of the upper vessel.
7. Pass the gas from the generating apparatus into
soap-suds, and set free some bubbles in air.
8. Pass the gas also into a collodion balloon until it is
full, then let it free in the air.
Note. — Before collecting this ^as, and before bringing a flame near it, be certain
that all air has been driven out of the generating flask and tubes. To do this allow the
gas to escape for a few minutes, then collect a test-tube full over water, as in Fig. 7 ;
bring the test-tube rapidly, mouth downwards, to a lighted lamp. If the gas burns
quietly it may be collected, but if there is either a sharp explosion or a whistling sound
the air is not all driven out. Neglect of this caution to test the gas in this way will almost
certainly lead to violent and very dangerous explosions.
HTDROOEN.
27
Fio. 8.
9. Prepare hydrogen gas, using the apparatus Fig. 8
and putting into it granulated
zinc and hydrochloric acid.
Pass the gas through a tube
containing fragments of calcic
chloride, for the purpose of
drying the gas ; and, through
the cork which should tightly
fit the end of this tube, pass
a glass tube drawn to a fine
point. After all air has escaped
apply a lighted match to the gas-jet.
10. Introduce a piece of fine
iron or steel wire into the flame.
Try the effect of the flame
on platinum wire and copper
wire.
11. The Chemical Harmoni-
cum. — Bring down over the jet a
tube about 4 centimetres wide
and 40 or 50 centimetres long,
as in Fig. 9. Use tubes of dif-
ferent diameters and different
lengths, and move them slowly
up and down.
12. Invert a long, dry, wide-
mouthed bottle or bell-jar over
the jet
1 3. Have a vessel made out of tin or sheet copper in
the form of a double cone, as in Fig. 10. At one end,
B, have a neck for a cork, and at the other end, A, a
Fro. 9.
J'.
St;
IS
i
28
HYDROGEN.
i
small opening about one-eighth of an inch in diameter.
The vessel should be about five inches long
and two and a half inches wide in the
middle. Pass a hydrogen delivery tube
in through B until the air and hydrogen
become well mixed (or fill the vessel
with oxygen and hydrogen mixed in a
jar), then close B tightly with a cork
and hold the end A to a flame. Hold
the apparatus so that when the cork blows
out no one will be struck.
Fio. IOl
ri
III
3.— Questions and Exercises.
1. Make a list of the properties of hydrogen which you have
observed in the preceding experiments.
2. Compare the phenomena you observed in the cases when
sodium and potassium were thrown on water.
3. Float a piece of filtering paper on some water, then drop on
this a piece of sodium ; compare the result with that noticed in the
first experiment, and also with that of i (a) in sec. 2.
4. Heat some water in a metal dish nearly to boiling, then drop
into it a little piece of sodium.
Explanation.— Potassium and sodium are two metals
that decompose water at ordinary temperatures. Each
molecule of water is formed of two atoms of hydrogen
and one of oxygen, and one atom of the sodium or of
the potassium displaces one of the atoms of the hydro-
gen, so that instead of the original molecule there is now
a new one consisting of one atom of sodium or potas-
sium, one atom of hj'drogen and one of oxygen. The
displaced atom of hydrogen escapes, and masses of these
m
tCXRROISEB.
29
form the bubbles of gas that rise to the surface when
either of these metals is sunk in water. The combustion
which goes on in some of the cases when these metals
float on the surface of the water, is due to the hydrogen
becoming ignited on account of the heat generated by the
rapidity of the chemical action. The different colours of
the flames are owing to small portions of the metals be-
coming vaporized and burning along with the hydrogen.
5. Try whether the colour of the flame would change if a platinum
or copper nozzle were used instead of a glass one, in ex. 9, sec. 2.
6. Is this flame hotter than that of a spirit lamp? Devise some
experiment to show that your answer is conect.
7. What phenomena occurred in ex. 11, sec. 2? How was the
pitch of the note made to vary ? Did the shape of the fl une
change? How?
8. Where did the moisture on the inside of the bottle, ex. 12 "C
2, come from? Give a reason for concluding that it could not have
come from the generating flask.
9. What became of the zinc in ex. 4, sec. 2 ? How do you know
that the gas was not air?
10. Try if zinc and strong sulphuric acid will yield hydrogen.
After the gas has ceased to come off, pour the zinc and acid out on
a plate, then pick out two or three pieces of the zinc, dry them
without rubbing and examine their surfaces ; put them in water for
a little while, then drop them back into the acid on a part of the
plate by themselves. Next try the effect of diluting the acid.
Does the hydrogen come from the strong acid or from the water?
11. Point out the resemblances between the gas obtained by the
action of zinc and sulphuric acid, sodium and water, and one of
those that resulted from the decomposition of wafer by electricity.
12. If there were two jars, one full of hydrogen and the other full
of air, how could you f.nd, out which jar contained each gas ?
13. What reasons are there for believing that when hydrogen
burns, chemical combination is going on ?
\
-4--
30
EXERCISES.
1 f:
r
14. Take two rubber bags, and fill one with hydrogen, the
other with oxygen. Subject both bags to an equal amount of
pressure between two boards, or otherwise, then connect them with
the apparatus known as the oxy-hydrogen blow-pipe, and having
turned on the hydrogen, ignite it, then carefully and very
gradually turn on oxygen gas from the other bag. A form of blow-
pipe apparatus is represented in Fig. 11. Gas holders may be sub-
Fio. 11.
stituted for the rubber bags, if more convenient ; but the gas
should be driven out with considerable force.
15. Introduce into the oxy-hydrogen flame, separately, a piece of
platinum wire, of steel, of zinc and of quick-lime.
16. Use an ordinary mouth blow-pipe and connect it with a
supply of oxygen by means of a rubber tube ; when the gas is
escaping under pressure hold the nozzle horizontally in a lamp or
gas flame. Try this for heating effect.
17. If sodium forms a compound with water when thrown upon
it, where is that compound ? Devise an experiment to test the
correctness of your answer.
4.— Additional Exercises.
I. To avoid any risk of explosion in the preparation of hydrogen,
sodium amalgam may be used instead of sodium. The amalgam is
prepared as follows : — Drop into a four or five inch test-tube about
a half of a cubic centimetre of mercury, heat this to boiling, then
throw into it bits of sodium cut small. Be careful to keep the
mouth of the test-tube turned in such a direction that no one will
be burned by the hot metal which may spurt out. When sodium,
about equal to the mercury in bulk, has been added, pour the liquid
quickly out on a cold plate. When cooled there should be a
STEAM.
31
it
(i
brittle, silvery white solid, — sodium amalgam. Some of this may
be put under the mouth of a test-tr.be that has been filled with
water and inverted as in Fig. 7. Test the gas with a lighted
taper, keeping the tube mouth downwards.
2. Try if results similar to those of ex. i, and i (a\ sec. 2, can be
obtained by using ice instead of water.
3. Try if hydrogen is given off in the following cases : —
(i) Iron is treated with nitric acid.
(2) Copper is treated with hydrochloric acid.
(3) Copper ** " sulphuric
(4) Iron
(5) Iron " " hydrochloric "
(6) Zinc " " nitric "
4. A gas jar fitted with a stop-cock is to be pressed mouth down-
wards into water until about f of the air is driven out, the stop-
cock is then closed and hydrogen equal in volume to about § of
the remaining air pressed into the jar. A delivery tube is then to
be fitted to the stop-cock and soap bubbles inflated on a metal dish
(an earthenware or glass one will likely be broken) with the mixed
gases. By means of a long splinter, ignite the gas in these bubbles.
Why should there occur an explosion here, when hydrogen will
bum quietly in a jar or at the mouth of a tube ?
5. If some dilute hydrochloric acid be poured on a little carbonate
of sodium (washing soda or sal soda) a gas will be generated, is it
hydrogen ?
6. Connect a piece of charcoal in a battery circuit in which is also
a galvanometer. Compare the conducting power of the charcoal
and of a piece of copper wire for electricity, by this means. Compare
their conducting powers for heat.
7. Did the gas that arose from the sodium amalgam come from
the sodium or the mercury ?
5. — Hydrogen and Steam.
As steam is one of the forms into which water may be
changed without altering its chemical composition, any
d2
HYDROGEN.
result obtained from steam by action between it and
another substance is really a result of water acting on
that substance.
Experiments.
I. Take an iron tube about ^ of a meter in length
and 2 centimeters in diameter (an old gun barrel will do
well) ; fill it nearly full of clean iron filings ; fit each end
with a tightly-fitting cork and tube, the one leading to a
pneumatic trough, the other connected with a flask con-
i!|f"
Fig. 12.
taining boiling water, as shewn in Fig. 12. Place the
iron tube in a charcoal fire, built upon bricks, or heat it
by gas flames. When it is red-hot, boil the water in the
flask and force steam through the tube. After the tube
has cooled, turn out the iron filings and examine them
carefully.
2. Collect a few jars of the gas, prepared in the
foregoing manner, and test for hydrogeri, as in former
experiments.
3. Take a hard glass tube drawn out to a point (such
a one as is user' for organic analysis with the point
REDUCTION.
33
nipped off, is just suitable), and about its middle, place
a imall quantity of black oxide of copper. Connect it
with a tightly-fitting cork and tube to a hydrogen
t' n
!l.
^^m—mz::
Fig. 18.
generating apparatus. The gas must be dried by pass-
ing it through calcic chloride or sulphuric acid, before it
reaches the tube with the oxide of copper in it. (See Fig.
13). After allowing the hydrogen to escape for a few
minutes, so as to drive out all the air from the apparatus,
heat the oxide of copper. Condensed vapor of water
should pass out of the point of the tube.
What compound is formed in passing steam over red-
hot iron in a tube? Compare the action in this case
with that of sodium or potassium on water.
DefinitiOD. — When the electro- positive element of a
compound is freed from some or all of the atoms of the
electro-negative element in combination with it, it is
said to have been reduced.
li .
A-.,
1
\>
'UK'
StM
1:'
1: 1
Note. — The oxide of copper may be introduced at the proper place without smearing
the tube by cutting -x long narrow strip of paper, folding it up the middle, placii.^ the
oxide in the crease thus formed, then gently shoving the paper into the tube and tipping
it over so as to spill the powder at the proper place.
■^ rw
34
CHEMICAL ACTIONS.
\m>
6— Additional Exercises.
1. Ascertain by experiment whether red-hot copper has the same
action on steam that red-hot iron has.
2. Repeat ex. 3, sect. 5, but instead of copper oxide use pow-
dered iron rust, which is an oxide of iron, Fe.jOjj. After the rust
has turned black, heat it strongly for some time, then quickly dis-
connect the combustion tuVje and spill the black powder in the air.
If the experiment is successful the particles should glow brightly as
they fall.
3. Try if red oxide of lead can be reduced in the same way that
iron oxide is.
4. Compare the phenomena observed when the copper oxide
and the iron oxide were heated in the current of hydrogen. How
do you account for the change of color in each case .'*
5. Heat to redness a piece of folded magnesium wire in a current
of steam.
6. How do you account for the glowing of the particles of iron
when scattered out of the tube after having been reduced? Has
the smallness of the particles anything to do with this phenomenon.?
7— Classification of Chemical Actions.
For convenience of reference, chemical actions are
classified as follows : —
Simple Combination. —When two or more substan-
ces unite to form a compound, but without causing the
decomposition of any existing compound. Illustrations
of this kind of change are found when magnesium burns
in air, thus uniting with the oxygen to form magnesic
oxide ; when oxygen unites with hydrogen to form
water, and when sulphur unites with iron to produce
sulphide of iron.
Simple Decomposition— When a compound breaks
up into simpler compounds, or into its constituent ele-
NOTES ON HYDROGEN.
35
meats. Examples of this are seen when mercuric oxide
decomposes into mercury and oxygen ; when chlorate of
potassium breaks up into oxygen and potassic chloride.
Decomposition by Displacement —When one or
more of the constituents of a compound are displaced by
other substances. This occurs when sodium displaces
part of the hydrogen of water, and when zinc displaces
the hydrogen of sulphuric acid.
Double Displacement.— When two compounds so
act on one another that they interchange elements or
groups of elements to form two new compounds. This
is sometimes known as metathesis, and takes place in
the case of potassic chloride acting on silver nitrate,
KCH-AgN03=KN03 + AgCl. Sodium carbonate and
hydrochloric acid yield sodium chloride and carbonic
acid, Na2C03 + 2HCl = 2NaCl + H2C03.
Questions and Exercises.
1. Classify the chemical actions of the experiments of sect. 3.
2. Under which class would you place the chemical action when
water is decomposed (a) by electricity, (d) by being passed over
red-hot iron, (c) by sodium amalgam .''
3. What kind of decomposition goes on when hydrogen is formed
from magnesium and dilute sulphuric acid, when hydrogen is formed
from iron and hydrochloric acid, when hydrogen is passed over
hot iron oxide ?
1"*
1
till
»<i
u.
•..3
■;,)
8.— Notes on Hydrogen.
I. A gram of hydrogen at 760 mm. pressure and 0° C.
occupies I r 1636 (ir2 nearly) litres; hence a litre of
hydrogen weighs tt-tVij¥ = '089578 ('0896 nearly) grams.
36
NOTES OS IIYDROGKN.
m
r \
2. The atomic weight of hydrogen is i, its molecular
weight is 2 ; this means that hydrogen, when freed from
combination with other substances, does not continue to
exist in the atomic state, but that its atoms unite in
groups, and these groups consist each of at least two
atoms. Its molecular volume is also 2, which means
that each molecule occupies the space of the two atoms
of which it is composed, hence there is no condensation
in the change into the molecular condition.
3. Hydrogen occurs chiefly in combinations such as
water, the acids, the hydrides of many elements, and as
a constituent of all organic bodies.
4. Hydrogen when pure is odourless. The disagree-
able smell which it has when prepared from zinc is due
to impurities of the metal or acid. The chief of these are
arsenic, which causes the smell, and lead and carbon
which give rise to the black flakes that float on the sur-
face of the acid and water
5. It has already been shown that potassium and sodium
decompose water at ordinary temperatures. The rarer
metals barium, strontium and calcium act similarly.
Vapour of water led over red-hot iron was decomposed.
Zinc, nickel, tin, antimony, lead, bismuth, copper, and
some other of the rarer metals will, at a red heat
(1000° F.), also decompose steam.
Mercury, silver, gold and platinum do not decompose
water at all.
6. Hydrogen is used to aid combustion, for reducing
metallic compounds, and for filling balloons. It is one
of the mixture of gases which we burn, and which we
CHEMICAL NOTATION.
37
call coal gas ; it is freed sparingly when coal is burned
in a furnace, and when water is raised to a very high
temperature, especially in the presence of red-hot carbon.
CHAPTER VII.
1.— Chemical Notation.
Symbols. — It was stated in a previous chapter, (iv.),
that, very generally in chemistry, the symbols of the
elements are used instead of the full names. This gives
us a kind of chemical shorthand which is brief, ex-
pressive, and easily intelligible.
In our present system of chemical notation each
symbol stands not only for the name of an element, but
also for a definite weight of that element, called its
atomic weight ; e.g.^ H always stands for i part by weight
of hydrogen, and CI for 355 parts by weight of chlorine,
hence the symbol of an element stands for three* distinct
things : —
(i) The name of the element ;
(2) One atom of the element ;
(3) The atomic weight of the element.
Thus, the symbol O stands for: (i) the name oxygen ;
(2) one atom of oxygen ; (3) 16 parts by weight of
oxygen.
III
M.
"1
-11 .
.,u
II in
1:1::
.lie
* A symbol stands for these three things taken collectively — not separately. The same
emark applies to formulae.
38
FORMULAE.
Hi 9
A small numeral written at the lower right hand cor-
ner of a symbol denotes that the atom is doubled, tripled,
etc,, e.g., O2, Ng, P5.
If we take i centigram as our unit of weight, then H
signifies i centigram of hydrogen, and O, 16 centigrams
of c^ygen, and so on. O^ would then represent 32 cen-
tijj.ams of oxygen ; N3, 42 centigrams of nitrogen : CI4,
142 centigrams of chlorine, etc.
Pormulee. — A chemical formula consists of two or
more symbols written side by side, and denotes that the
elements for which the symbols stand have united to
form a chemical compound. The symbol of the most
electro -positive constituent of a compound stands first in
its formula.
When water was decomposed by electricity, the hydro-
gen was given off from the electrode that was connected
with the zinc^ or negative pole of the battery, and since
oppositely electrified bodies attract each other, the
hydrogen is said to be more electro-positive than the
oxygen is.
The formula of a compound substance stands for :
(i) The name of the compound ;
(2) One molecule of the compound ;
(3) The molecular weight of the compound.
The molecular weight of a compound is found by
taking the sum of the atomic weights of its constituent
elements.
A numeral placed before a formula multiplies every
atom and atomic weight in it, as far as the first comma.
EQUATIONS.
39
plus sign, or period. For example, in 4H0O, the 4 multi-
plies both the atoms and the atomic weights, and means
4 molecules, each consisting of two atoms of hydrogen
and one atom of oxygen.
The formula H^O expresses the following facts :
(i) That water consists of hydrogen and oxygen ;
(2) That its molecule consists of 3 atoms : 2 of
hydrogen and i of oxygen ;
(3) That its molecular weight is 18.
The formula C^2^22^iv signifies :
(i) That cane sugar consists of carbon, hydrogen,
and oxygen ;
(2) That its molecule consists of 45 atoms: 12 of
carbon, 22 of hydrogen, and 1 1 of oxygen ;
(3) That its molecular weight is 342.
The molecule, O2, consists of 2 atoms.
TJ (t « «
P(l , <(
4» 4
CgHgO, (common alcohol) consists of 9
atoms.
Equations. — A chemical equation consists of signs
and formulae, and expresses the fact that certain sub-
stances do, of themselves, or by means of some force
applied to them, decompose, and re-arrange their atoms
so as to form other substances.
For example the chemical equation —
H2 + 0=H,0
2+16 i8
«
«
it
a
u
u
i<r.
•'3
11/
"'5 1
O ,
n '
-11
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J.
..,4
""5
.,1111
I'll'
hi'"
d"
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itf
i f'-J
IT
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ill
40
EQUATIONS.
expresses the fact that 2 centigrams of hydrogen unite
with i6centigrams of oxygen and form 18 centigrams of
water, when a centigram is the unit of weight chosen.
The equation is equally true for any other unit of weight,
for example, that of one atom of hydrogen.
In the same way, the equation
CaC03 + 2HCl = CaCl2+H20 + C02
100 4-73
iii-f 18 +44
may be thus translated : mix 100 grams (or ounces) of
marble with a solution of 73 grams of hydrochloric acid
and they will yield ill grams of calcic chloride, 18
grams of water, and 44 of carbonic anhydride.
When two symbols or groups of symbols are connected
by the sign " + ", it means that the substances are mixed,
but not in chemical union ; when symbols are written
side by side without the connecting sign, the meaning is
that the substances represented by them are in combina-
tion. The sign " = " is not to be understood as used in
its algebraic sense of equality ; it may be read " gives,"
or " produces," or " forms." Thus : —
KHO+HCl = KCl-}-H,0
means that the compound consisting of one part of
potassium, one part of hydrogen and one part of oxygen
when mixed with another compound consisting of one
part of hydrogen and one of chlorine, may, by proper
treatment be made to yield two new compounds which
will also be mixed ; one of these is made up of one part
of potassium and one part of chlorine, the other of two
parts of hydrogen and one of oxygen.
OXYOKN.
41
The sum of all the atoms of any element on one side
of the equation must equal the sum of the atoms of that
same element on the other side ; hence, the total number
of atoms on one side equals the total number on the
other, and the sum of the atomic weights on one side is
identical with the sum on the other side.
CHAPTER VIII.
1— Oxygen.
When water was decomposed by electricity two gases
were obtained ; one of these was oxygen ; and this chapter
treats of the preparation and the more important pro-
perties of this substance.
Experiments.
I. Put a couple of grams of chlorate of potash in a
test-tube fitted with a cork and delivery tube, as in Fig.
F». 14.
14, then heat the tube ; when bubbles of gas come off
freely, hold a glowing splinter (one that is on fire but not
"J
'"3
'"!>
42
OXYUEN.
HI!
H
(!■: >
blazing) close to the end of the delivery tube, or better,
collect a tube full of the gas, as shown in the figure, and
put the splinter in it. When the gas has ceased to come
off lift the delivery tube out of the water, then remove
the flame from under the tube, and when the latter has
cooled, dissolve the white residue ; also dissolve a little
chlorate of potash separately, test ea ' 'vith a drop of
silver nitrate solution.
2. Weigh out I gram of red oxide of mercury and
place in a test-tube fitted with a cork and delivery tube,
as in Fig. 14. Heat uniformly at first so as not to break
the glass. Oxygen is driven off, and collects in the up-
right glass vessel standing in the ** pneumatic trough."
The gas that collects at first should be thrown away.
Raise the tube containing the oxygen out of the water
and plunge a glowing splinter into it.
Examine with a magnifying glass e grey deposit
formed in the test-tube. After gently . ilmg out what
oxide is still unchanged, strike the mouth of the test-
tube sharply on the palm of the hand to drive out
some of this grey substance.
3. The best method of preparing oxygen in moder-
ately large quantities in a small laboratory, is the follow-
ing : — Into a mortar put 10 grams of potassic chlorate
and 2)^ grams of manganese dioxide. Powder them,
and place in a dry Florence flask. Use chemically pure
material only, as serious explosions have occurred from
Note. — When a delivery tube from a heated vessel dips under water never remove
the source of heat until the tube has been lifted out of water, else there will be great
danger of the water running into the heated vessel and breaking it if made of glass.
Or an explosion might result from the steam suddenly formed if the vessel be made of
metal.
%
<)XY(}KN.
43
orp^anic matter, or carbon beinpj mixed with the man-
ganese dioxide.*
Meat strongly but carefully, and collect four or five
small jars of the gas ; or better still, collect in a gas-
holder. Manganese dioxide is mixed with the potassic
chlorate to cause the latter to part with its oxygen at a
lower temperature than it otherwise would. The man-
ganese dioxide remains unchanged at the end of the
experiment.
4. Into a jar of oxygen plunge a piece of glowing
charcoal. After combustion has ceased, pour some water
into the jar, shake, then test with blue litmus.
5. Put a little sulphur in a deflagrating spoon, ignite,
and place in a jar of oxygen. Pour some water into
the jar after the combustion has ceased. Test with
blue litmus paper ; also taste the solution.
6. Draw the temper from a piece of fine watch spring
by p. sing it slowly through a lamp flame ; file the end
of the ^)ring very thin, then tip with sulphur, or better,
wrap the thin end of the spring tightly round a piece of
charcoal, and ignite. Place in a jar of oxygen, and
it will burn with beautiful scintillations. Add water
to the solid that forms, then taste and test, first with
•"3
"J
'")
*To test the manganese dioxide heat a little of it on a deflagrating spoon ; if any
glowing or combustion is observed there are dangerous impurities in it.
Note. — To make red and blue solutions, "steep some solid litmus in water or weak,
alcohol. Divide the liquid which you pour off from the sediment, into two parts ; to one
add a few drops of weak sulphuric acid ; to the other add a little solution of caustic soda.
You will then have red and blue litmus solutions, and if you add them to the products
of your experiments with oxygen, you will be able to test whether a new compound has
been formed in case other evidence of a chemical change is wanting."
Litmus paper is prepared by dipping strips of white filtering paper in these solutions,
and allowing them to dry.
44
OXYGEN.
* ;
ill
blue litmus, then with red litmus, being careful to see
that the results are not masked by the combustion
products of the carbon or sulphur. Observe closely the
deposit on the sides of the jar, and examine the pellets
that fell to the bottom.
7. Clean the spoon used in the last experiment, and
place a piece of phosphorus on it about the size of a pea.
Ignite, and place in a jar of oxygen. After combus-
tion has ceased, remove the spoon, and pour some
water into the jar. Shake up the water with the pro-
duct of the combustion, taste, and then add some blue
litmus solution.
8. Place a small piece of sodium in a deflagrating
spoon, hold it in a lamp flame until it begins to burn,
and then plunge into a jar of oxygen. Add water,
taste, and test with reddened litmus.
9. Pour a little potassium hydroxide into a jar of
oxygen. Shake, and if no change in volume takes place,
add a small quantity of pyrogallic acid. Shake again,
and note any change. Modify this experiment by
selecting two test-tubes such that one will ju.st pass
mouth first into the other, fill the smaller one with
oxygen gas and half fill the larger with solution of
caustic potash ; drop into this as much pyrogallic acid
as will rest on a 10 cent piece, and quickly pass the ether
tube mouth first into it ; let the whole stand for some
hours.
10. Make a large bubble on a plate with oxygen gas,
then hold a magnet near it Does the magnet attract
the bubble film, or the oxygen in the bubble ? Devise
an experiment to determine this point.
TEStS FOR OXYGEN.
45
2.— Tests.
1. If free oxygen be present to any considerable ex-
tent a glowing splinter will rekindle when put into the
gas.
2. If copper clippings be heated with weak nitric acid
and the resultant gas allowed to stand over waU-r it will
become colourless. Oxygen passed into this gas causes
it to turn brown and be dissolved in water.
3. A mixture of pyrogallic acid and caustic potash
solutions (pyrogallate of potash) will turn dark brown in
presence of oxygen.
3.— Questions and Experiments.
1. How would you separate manganese dioxide from the other
ingredients as it occurs in the residue from the preparation of
oxygen ? How would you show that there is a substance in this
residue different from either of those taken at first ?
2. Point out the resemblances and differences you have observed
between oxygen and hydrogen.
3. Oxygen is said to support combustion much more energetically
than air does. What reasons are there for such a statement as
this ?
4. When charcoal or sulphur is burn«^d in oxygen, the gases pro-
duced differ from air. How may this difference be shown ?
)
i1
.t-i
■'4
")
4. — Notes.
Oxygen : atomic weight, 16 ; mol. weight , '^2 ; mol.
vol. 2 ; specific weight (air=i) I'lo^dj.
Oxygen occurs free in the atmosphere, of which it
forms about 21% by volume; it also exists largely in
46
NOTES ON OXYGEN.
(I
i
combination, as in water, in many minerals, and in all
organic bodies.
Oxygen is one of the substances entering into most
chemical actions that are known as combustion, and is
the gas necessary for the support of life.
It is possible that in the preparation of oxygen from
chlorate of potash and manganese dioxide the latter
substance may act by absorbing oxygen from the
chlorate to form a higher oxide, Mn207, which on account
of its instability is immediately decomposed. There are
other substances which may be substituted for the
Mn02 in this operation, viz.: — CuO, Fe.^O.j and PbO.
Some of these, at least, are capable of being further
oxidized to CUO2, Pb02 ; (R. & S. I., 174). The object of
using the Mn02 is to cause the oxygen to be given off
at a temperature much lower than if the chlorate were
heated alone, hence it exerts some effect which causes
the molecules of the chlorate to be decomposed more
readily.
When iron is burned in oxygen the black brittle
globules that fall to the bottom are magnetic oxide of
iron ; they have the composition Fe304 and are generally
considered to be a union of Fe02 and Fe^O^, ferrous and
ferric oxides respectively, hence sometimes called ferro-
ferric oxide. The red powder which settles on the side
of the jar is ferric oxide, Fe^,03. Oxygen may be
separated from other gases with which it is mixed by
using pyrogallate of potash.
5.— Additional Exercises.
I. Heat some red oxide of lead in the same manner as in experi-
ment 2, sec. I. Test as before.
EXERCISES.
47
)
2. Cut zinc foil into fine strips, tip thorn with sulphur, place on a
small piece of sheet zinc on a cork ; ignite and place in a jar of
oxygen. After combustion has ceased, add water to the compound
formed, tlien taste and test with litmus paper.
3. Repeat experiment 8, sec. i, using, first, potassium, and then
magnesium.
4. Burn a piece of charcoal in a jar of oxygen, then shake some
lime water up with the gas that is in the jar. Shake some lime
water up with oxygen. What does this demonstrate ? Does
a similar result follow from the burning of sulphur in oxygen ?
5. Stir some red lead ':.ith dilute nitric acid, after a brown
powder, PbOj, peroxide of lead, has formed, filter, dry the powder ;
and try if, when heated, it will yield oxygen.
6. Bend a piece of wire into the form shown in Fig. 15,
place a piece of phosphorus about as big as a pea on top
of it, then set it on a plate of water and place a bell jar
filled with oxygen over it. Let it stand for two or three
qJ •) days, being careful that there is plenty of water on the
Fio. 15. plate. Then test the water with litmus paper.
7. Try if nitrate of potash (saltpetre) when heated gives off
oxygen. When the nitrate begins to boil, drop into it a bit of char-
coal or the end of a match.
8. Mix some manganese dioxide with sulphuric acid and heat
gently in a test-tube. Determine if oxygen is given off. Pass the
gas that comes off through cold water, and test again for oxygen.
9. When copper oxide was heated in a current of hydrogen
what happened ? Try if the operation can be reversed by heating
some fine copper filings in a current of oxygen.
10. Place some zinc filings in a hard glass tube, heat this while a
current of oxygen is passing through it. Repeat the experiment,
using lead filings.
v."
' .1
Note. — A bell jar full of gas may easily be transferred from one place to another
by covering the mouth with a piece of ground glass, v
' /
48
OZONE.
6.— Ozone.
1:
Experiments.
1. Suspend a stick of clean phosphorus in a closed
bottle that has a little water in the bottom of it ; let the
phosphorus remain for a couple of days, then remove
the stopper and smell the gas in the bottle. (If the
phosphorus is covered with a brown coating scrape this
off under water.)
2. Repeat the experiment and test the gas formed in
the bottle, with starch paper, prepared as follows : —
Boil some starch to a paste, drop into it some
potassic iodide, dip in this strips of white unsized
paper (strips from a leaf of a scribbler will answer). Hold
one of these test papers in ozone.
3. In a wide-mouthed bottle or beaker put two
grams of crystals of permanganate of potash, KMnO^,
and on these pour some sulphuric acid, but do not warm
the mixture. Test the gas with the starch paper. Smell
the gas.
2KMn04+HaSO, = KoS04 + Mn207+H,0-K2S04+
2Mn62 + H20 + 03.
4. When the gas is coming off freely from die mixture
of permanganate and sulphuric acid, hang in it, but so as
not to touch the liquid, a piece of paper saturated with
turpentine. This should burst into flame after a short
time. i
Note. — Care must be taken
heptoxide, explodes violently, wh
be set in a larger vessel o^water.
toW
■ep this mixture well cooled as Mn207, m«asciii..:.c
*eated. The beaker containing the mixture should
OZONE.
49
Fio. 16.
5. Fit up a bottle and three test-tubes, as shown in
Fig. 16; into the bottle put a
mixture of permanganate of
potash and sulphuric acid, and
into one test-tube a faintlyblue
solution of indigo, into a
second a weak solution of log-
wood, and into the third a
slightly purple solution of per-
manganate of potash ; allow
the gas that comes off to bubble through the three
solutions for a few hours.
6. Pass some of the gas prepared, as in the last experi-
ment, through a solution of iodide of potassium, KI.
Explanation. — The gas formed in these experi-
ments is ozone. The symbol for it is usually written
O3, but better OOo. Chemically it is oxygen, but it has
properties which ordinary oxygen has not. This may
be demonstrated by using oxygen instead of ozone in
Ex. 4, 5 and 6. The theory is, that ozone exists in a
molecular state different from that of oxygen ; for theo-
retically, while the molecule of oxygen consists of two
atoms, that of ozone consists of three. This three-atom
molecule is very unstable, so is easily decomposed into
ordinary oxygen. Thu.s, 203=303 ; but as each mole-
cule of O3 breaks up, one atom is set free as an atom, not
as a molecule. The subsequent union of two of these
free atoms forms the third molecule.
An Oxide is a compound formed by the union of
oxygen with some other element.
It
•'3
•■J
'•"»
i
ih
60
HYDROGEN DIOXIDE.
7.— Additional Exercises.
r. Pass the gas prepared, as described in Fig. i6, into some sul-
phuric ether, after a time pour a little of this ether into weak solutions
of indigo, logwood and litmus. Try if the ether before its treatment
with the gas will produce the same effect on these substances.
2. Pass sparks between the terminals of an electric machine in
dry air, after a little while the odour of ozone should be detected.
3. Try if ozone will support combustion as oxygen does by
causing a glowing splinter to burst into flame.
References for ozone, R. & S., vol. I., 194-201 ; Muir & Slater, 224;
R., 85-90.
CHAPTER IX.
Hydrogen Dioxide.
Hydrogen and oxyi^en form two compounds ; one of
these has already been considered in Chapter v., the
other is hydrogen dioxide, or hydrogen peroxide, and
it is intended *:o discuss this substance in the present
chapter. Its formula is H^Og.
1.— Hydrogen Peroxide.
Barium oxide, B.iO, when heated to dull redness in a
current of oxygen, or in a free supply of air, changes into
barium dioxide, Ba02.
Experiments.
I. Treat some barium peroxide, BaO^, with hydro-
chloric acid diluted with two or three times its own bulk
nVDROOEN DIOXIDE.
61
of water ; after the white oxide has all dissolved, drop
in some sulphuric acid, and when the precipitate ceases
to form, filter. The filtrate is a solution of hydrogen
dioxide. As this is a very unstable compound it cannot
be evaporated in the ordinary way by heating, because it
then decomposes into water and oxygen.
BaO, + 2HCI = BaCl^ + Hp.,.
The sulphuric acid was added to form an insoluble com-
pound with the barium chloride, BaCl^, which then
separated from the liquid as a precipitate.
BaCl2 + H,S04 = BaSO^ + 2HCI.
The solution of hydrogen dioxide may be evaporated
under an air pump, if required.
2. Divide the solution of hydrogen dioxide into a
number of parts, into one drop some logwood solution,
into another some indigo, into a third some litmus.
3. Prepare some test papers as directed for ozone, and
dip one of them into hydrogen dioxide solution. If it
does not turn blue at once add a few drops of clear
ferrous sulphate (copperas), FeS04, solution. Dip
another piece of the paper into another part of the
hydrogen dioxide solution, and let it stand for several
hours.
4. Make an acid solution of potassium permanganate,
KMnO^, but only slightly purple in colour, add hydrogen
dioxide solution and shake the two together, then let the
mixture stand for some time. Try if the permanganate
solution loses colour when no hydrogen dioxide is added
to it.
•■•
.i
iW\
•I
I:
•k ■
on I AKiu oULLLut OF EUUUA I lUN
02
EXERCISES.
5. Heat a portion of this hydrogen dioxide solution
in a corked test-tube ; place a strip of the test paper
above the solution to find if this compound conies off as
gas. After heating, test the liquid with the test paper
and ferrous sulphate solution.
Explanation. — The chemical action of hydrogen per-
oxide is due to the same cause as that of ozone, viz.: the
weak chemical attraction existing between one of the
oxygen atoms and the other parts of the molecule,
that is, HgO. Whatever may be the molecular struc-
ture of the group of atoms, it is clear from the result of its
decomposition that one atom of oxygen is held very
loosely, hence easily breaks away from the others as an
atom. These loose atoms join in pairs to form oxygen
molecules, thus: — HjjO^ = H^O-f-0 or better, 2H202 =
2H2O + O2.
i.
3.— Additional Experiments.
1. Will barium dioxide, when treated with dilute sulphuric acid,
yield hydrogen dioxide ?
2. Try if barium dioxide, when treated with dilute nitric acid, and
the result precipitated with a few drops of sulphuric acid, will give
a solution of hydrogen dioxide ?
3. Does the manner of the decomposition of the hydrogen diox-
ide molecule indicate in any way that the formula should be written
H2O2, rather than as two hydroxy! molecules, 2 HO .-*
NASCENT STATE.
53
IH'
CHAPTER X.
Nascent State.
Many substances, particularly elements, at the instant
at which they are freed from combination, possess a
chemical activity in the way of forming molecules which
requires a special explanation. This will be briefly
given in the present chapter.
Ordinary oxygen does not bleach indigo, logwood,
litmus or permanganate of potash solutions, yet ozone,
which is only oxygen in a somewhat altered molecular
combination, does destroy the colours of these substances.
Hydrogen gas may be led for days, through silver
chloride held in suspension in water, yet the chloride
will not be decomposed. If, however, some silver
chloride (prepared by dropping hydrochloric acid, or a
solution of a chloride, into a solution of silver nitrate) be
spread on a piece of zinc and the whole immersed in
dilute sulphuric acid, the silver chloride will, in a few
hours, be reduced to metallic silver, (i). Zn-f H2S04=
ZnS04 + 2H. (2). H + AgCl=HCl + Ag. Similarly, a
current of hydrogen passed through potassic chlorate
solution has no effect on it, but hydrogen generated in
the solution from some pieces of zinc and dilute sulphuric
acid will reduce the chlorate to the chloride of potassium,
Zn+H,SO, = ZnSO, + 2H, and 6H + KC10,-KC1+
3H2O Free hydrogen has no effect on nitric acid, but
hydrogen freed in presence of the acid from a compound,
at once reduces the acid, (i). Zn + 2HN03=Zn (NO.,)^
+ 2H. (2). 2H+2HN03=2H20 + 2N02. This is the
■M'
M it*
3 .1
•1' I
"J
: ,1"
54
NASCKNT STATE.
reason that zinc and nitric acid do not yield free hydrogen
but an oxide of nitrogen generally. Many substances in
solution are oxidized by passing chlorine gas through
the liquid. Now chlorine contains no oxygen, so we are
obliged to look elsewhere for a reason for this change. It
is well known that chlorine has a.grcat affinity for hydro-
gen; so strong indeed is this attraction, that it breaks up
the water molecules, appropriates the hydrogen for the
formation of hydrochloric acid, and sets the oxygen free,
and it is this latter which oxidizes the substances,
2Cl-|-H^O = 2HClH-0, though a stream of oxygen gas
produces no such effect. Numerous instances might be
given of similar chemical action brought about by ele-
ments at the instant at which they are freed from com-
bination, though they do not retain the power for any
appreciable length of time. When the molecule of a com-
pound is decomposed, the constituents pass off as atoms,
and these may either unite with other elements to form
new combinations, or may remain uncombined with any
other substance ; but in the latter case they combine with
each other, and, since the combining powers of atoms are
limited in amount, though for different elements these
amounts are different, it follows that if two atoms of the
same kind combine with each other their affinity for
other atoms is lessened by the amount of attraction by
which they are held together. Their chemical activity
in the way of forming new combinations will therefore
be reduced ; hence, at the instant at which atoms are
freed from molecules and exist as individual atoms, their
chemical attraction for other atoms is stronger than it is
after they have joined in groups. At the time at which
a portion of an element exists as atoms, and before these
ACIDS.
55
have combined to form molecules, it is said to be in the
nascent state. When the molecule of ozone breaks up
into a molecule of oxygen and an atom of oxygen, the
latter is in the nascent condition, that is, uncombined
with any other atom, so that its powers of combination
are not impaired in any way. On this account it oxidizes,
and thus destroys the colouring matters spoken of; unites
so vigorously with turpentine that combustion is set up,
CioHi,i+280.j== ioCO.,+8H.,0+280^, ; and decomposes
potassic iodide by oxidizing the potassium, 0.5+2KI =
Oo+K^O+I.j. Similarly the hydrogen atoms whon first
liberated decompose silver chloride, reduce nitric acid,
reduce potassic chlorate and decompose sulphuric acid
under proper conditions. The nascent oxygen resulting
from the spontaneous decomposition of the hydrogen
peroxide molecule, or from the action of chlorine on the
water molecule, acts in a way precisely similar to that in
which the oxygen did, when the ozone molecule was
broken up.
For more detailed treatment, see Tilden, p. 125-6; Wiirtz, p. 207-8;
Muir and Slater, p. 233 ; Renisen, Th. Ch. 50 ; R., go ; D. and W., 69.
f
•1'
■;;3
' 1"*
the
for
by
CHAPTER XI.
ACIDS, BASES AND SALTS.
It has been necessary several times to mention sub-
stances which have been called acids. These form one
of three classes that include a great many chemical com-
pounds. The other two classes are bases and salts.
56
ACIDS.
Salts and acids are always compounds ; bases also are
compounds, and are either oxides or hydroxides of
metals. Oxides of the non-metals are generally acid-
forming substances, — never bases ; and the hydroxides
of the non-metals are acids.
Hydroxides or hydrates are compounds formed by
the union of an oxide with water.
1.— Acids.
In order to learn some of the general properties of
acids, perform the following experiments, using any three
or four substances labelled acids, which you can find upon
your working table : —
1. Pour some of the acid, or drop a crystal about as
big as a pea if the acid is a solid, into twenty or thirty
times its own volume of water in a test-tube. Taste the
solution.
2. Half-fill a small test-tube with blue litmus solution
and add to it some of the diluted acid. Add some of
the acid to some red litmus solution also.
3. Place some " bread soda," bicarbonate of sodium,
NaHCOy, in a test-tube and pour some of the dilute acid
upon it.
Tabulate your results as follows : —
Name of
Acid.
Taste.
Action on
Rfd Litmus.
Acti"" on
Blue
So>
Remarks.
BASKS.
57
All acids contain replaceable hydrogen ; that is, hydro-
gen which may be driven out of the molecule by one or
more atoms of some other substance. A familiar ex-
ample is in the preparation of hydroi^cn gas from
sulphuric acid, in which two hydrogen atoms are replaced
by one of zinc.
2.— Bases.
Any metallic oxide or hydroxide, the metal of which
is capable of replacing the hydrogen of an acid is a
base. The hydrogen of acids may also be replaced by
metals, the term base, however, is usually applied only to
the oxides and hydrates. The latter require a little
attention. A molecule of an hydroxide is formed by
the union of a metallic atom with one or more hydroxyl
groups (HO). Examples of these have already been
met v/ith when potassium and sodium were thrown
on water. These hydrates are formed either by direct
action of the metal on water (this occurs only with some
of the alkalies) or by dissolving the oxide in water.
(There is a third method which need not be discussed
here, as it belongs essentially to the chemistry of the
metals).
Experiments.
1. Take a piece of the metal potassium, about the size
of a pea, place it in an iron spoon, and heat it over a
spirit lamp until it has ceased to burn. Then add a little
water, and test the solution with red, and with blue
litmus, as before. Taste the solution.
2. Repeat this experiment, using the metals, magne-
sium and sodium,
•i
y
li
58
SALTS.
3. Obtain a piece of quick lime, place a bit of it on a
piece of dry litmus paper ; then put a piece of the lime
as big as a bean in a large test-tube full of water, shake
it up, and let the whole stand until the water becomes
clear. Taste it, and test with litmus and with turmeric
paper.
4. Repeat, using barium oxide.
Tabulate results as in the case of the acids.
3.— Salts.
When the hydrogen of an acid is replaced by a metal,
or the metal of a base, the resulting compound is a
salt.
Experiments.
1. Take a piece of "caustic soda" (sodic hydrate)
NaHO, about the size of a pea, and dissolve it com-
pletely in a test-tube of water, then add to it hydrochloric
acid, drop by drop, until a piece of blue litmus paper
placed in the solution slowly begins to turn red, pour
half of this solution into an evaporating dish, place on a
sand bath and heat until all the water is driven off.
Carefully examine the residue. Taste it.
Pour the rest of the solution into a flat dish of any
kind, and allow it to remain for a day or two in a warm
room.
2. Perform similar experiments using potassium hy-
droxide, KHO (caustic potash), and nitric acid ; also
sodium hydroxide r,nd sulphuric acid. .
3. Use the acids as in the last experiment, but instead
of the base take copper or zinc.
SALTS.
59
Tabulate your results, especially with regard to their
effc Jt on litmus.
Acids in which there is one replaceable hydrogen atom
are monobasic Nitric, hydrochloric and acetic acids
are examples. These have the formuLne HNO3, HCl,
HC2H3O2 respectively. Those acids in which there are
two atoms of replaceable hydrogen are dibasic. Sul-
phuric, H2SO4, and carbonic, H.^CO.^, acids are examples.
Either one or both atoms of the hydrogen may be
replaced, thus forming either acid or neutral salts; thus
and
or
H.,S04+K = KHS04+H,
KHS04-t-K = K,S04+H,
H,SO,+2K-K2S04+2H.
Of course such salts are possible only when the valency
of the base is less than the basicity of the acid ;— the
basicity being determined by the number of atoms of
replaceable hydrogen in the molecule.
Acids with three atoms of replaceable hydrogen are
tribasic Phosphoric acid H3PO4, is a good example of
this class. Either one, two or three of its hydrogen
atoms may be displaced, and these not necessarily by the
same base, thus: —
NaH.PO^, Na,HPO„ NagPO^, Ca(H,POJ.,. NaNH.HPO^
are some of the salts formed from it.
Experiments.
I. When nitric acid is prepared, a difficultly soluble,
white solid was left in the retort ; procure a lump of this,
wash it for a mir ute or two in a stream of water to free
it from any sulphuric acid that might be adhering to it.
ll»i
m ■
j'li'i 1
■■..}
m
\\m
^ss
■i.i'
.1?
60
SALTS.
then dissolve it. Test a part of the solution (i) with
litmus, (2) with a solution of barium nitrate, — the former
shows the acid nature of the solution, the latter that it is
a sulphate. To the part of the solution left add,
cautiously, caustic potash until it is neutralized, then
evaporate to dryness. A white salt should be obtained
which answers to the test for a sulphate but is neutral to
litmus.
2. Weigh out three and a half grams of strong sul-
phuric acid, and separately 2 grams of solid potassic hy-
drate, dissolve the latter in a measuring glass. Add half
the pc tash solution to the whole of the acid, evaporate to
dryness, dissolve the salt and test a drop of the solution
with litmus and with barium chloride. Add the rest of
the potash solution, again evaporate to dryness and test
as before. This time the salt should be neutral to litmus.
In naming them, acid salts are frequently distinguished
from neutral ones by the prefix bi-, thus, NaHSO^ and
NaHCOg are known as acid sulphate and acid carbonate
of sodium, bisulphate and bicarbonate of sodium, or
sodium hydrogen sulphate and sodium hydrogen car-
bonate.
These experiments show that it is possible to replace
the hydrogen of sulphuric acid in two separate stages,
and that two distinct substances are obtained as a result
of these displacements. These illustrate the double
basicity of the acid. A similar attempt with potassic
hydrate and nitric acid will result in only one salt being
formed, hence nitric acid is monobasic.
There is still another class of salts known as basic,
but, as there will be no occasion to refer to them in this
book, all discussion of them will be omitted.
I
Nomenclature.
61
4.— Questions and Exercises.
1. Given the acids HNO,, HgCO;,, and copper, potassium,
sodium hydrate and silver, write the formuht for all the salts
that theoretically could be formed. Write the names of these salts.
2. Which one out of each of the following pairs is correct and
why.? (i) PbNOg or Pb(N03)2 ; (2) ZnllSO^ or ZnSO^ ; (3)
NaSO^ or NaHSO^ ; (4) AgSO^ or Ag,S04.
3. From what acid is each of the following prepared, and what is
its basicity, as shown by the salt.? KNOg, Pb(N0.2)o, NagSOa,
CaSO^, Ba(N03)2.
I
CHAPTER XII.
Chemical Nomenclature.
The object of this chapter is to explain the principles
upon which the names of the compounds in inorganic
chemistry are based. To the beginner these names
doubtless appear bewildering in their variety, but a
single lesson, or at most two lessons, should put him in
a position to readily name the inorganic compounds,
once he knows a few formulae, (chiefly of acids), and the
valency of the elements and radicals.
'J
■lO
1.
Binary compounds, that is, those of two ek ments, have
names that end in -ide. The most electro-positive ele-
ment stands first (which one this is, will be learned by
practice), and its name may be in either the noun or
J . = —
62
nomknclature.
ii
nil
IS'
I'
11
r:
adjectival form in the complete designation. Thus KCl
is potassic chloride, potassium chloride or chloride of
potassium ; H.^S is hydric sulphide, hydrogen sulphide
or sulphide of hydrogen.
When the electro-negative element unites in mire
than one proportion with the other, the number of parts
of it in any particular combination are indicated by the
prefixes mono or mon-, di-, tri-, tetr-, or tetra-, and
pent-. Thus H2O (water) is chemically hydric mon-
oxide ; H^O^ is h^ drogen dioxide ; CO is monoxide of
carbon ; CO^ is carbon dioxide ; PCI3 is phosphorus tri-
chloride ; CCI4 is carbon tetrachloride ; P2O5 is phos-
phorus pcntoxide. An old ending, -uret, is sometimes
used instead of -ide ; and prot- is an old prefix used
instead of mono-.
In the names of acids the endings -OUS and -ic very
generally occur, and they indicate that the acid whose
name ends in -iC has in its composition a greater
quantity of oxygen than the one whose name ends
in -OUS. This does not mean that either of these
endings points to a fixed quantity of oxygen, but that
relatively the -ous acid always has in it less oxygen
than the -ic acid. Thua HNO3 is nitric acid, and HNO2
nitrous acid ; HCIO^ chlorous acid, HCIO3 chloric acid ;
H2SO3 sulphurous acid, and H0SO4 sulphuric acid. The
prefixes hypo (beneath) and per (above) are used also
with regard to the quantity of oxygen in the molecule of
the acid. T;<us HCIO is hypochlorous acid ; HClOjj
chlorous acid ; HCIO3 chloric acid, and HCIO4 per-
chloric acid.
VALENCY.
63
In naming salts the prefixes belonging to the names
of the acids a'*e preserved, but in the salt the ending
-OUS of the acid is changed into -ite, and the ending ic-
of the acid into -ate. Thus KCIO is hypochlorite of
potassium; NaClO.j chlorite of sodium; AgClOg chlorate
of silver, and KCIO^ perchlorate of potassium; CaS04 is
sulphate of calcium, and Cu(NOj2 i^ nitrite of copper.
In salts the adjectival form of the name of the base may
be used. Thus AgSOg is argentic sulphite, or silver
sulphite, or sulphite of silver ; KNO3 is potassic nitrate,
potassium nitrate or nitrate of potassium.
CHAPTER XIII.
1.— Valency or Atomicity.
The student who has followed this book thus far will
probably have had some difficulty in deciding in what
proportions elements unite with one another. The object
of this chapter is to explain this difficulty, as far as can be
done at this elementary stage.
The following are the formulae of some well known
compounds: —
HCl, HA NH3, CH^; KCl, FeCU, PCI3, CCl^.
In the first four, hydrogen unites in proportions of
one, two, three and four aroms with one atom re'spec-
I.I"
4
;'
••■
:.>
'^
64
VALENCY.
it!
tfi!
Pi >"
.1 >"'
IW*'
$i^
lively of chlorine, oxygen, nitrogen and carbon. The
second group shows similar compounds of chlorine. It
seems, therefore, that one atom of hydrogen unites with
one atom of chlorine to form a definite stable compound,
while it takes two atoms of hydrogen to form such a com-
pound with one of oxygen, three atoms of hydrogen with
one of nitrogen, and four of hydrogen with one of carbon.
Just as hydrogen is taken as the unit of atomic weight,
so it is taken here as the unit of combining power, and
the other elements are valued with reference to this one.
Thus an element, one atom of which unites with one
atom of hydrogen, or replaces an atom of hydrogen in
combination is said to be a univalent or a monad ele-
ment, that is, it is worth one. Similarly an element, one
of whose atoms is capable of uniting with two atoms of
hydrogen, or of replacing two atoms of hydrogen in a
compound, is called a bivalent or diad element, — worth
two. According to this classification elements are divided
into monad or univalent, diad or bivalent, triad or
trivalent, tetrad or tetravalent, pentad or quinqui-
valent, and hexad or hexevalent elements. The
combining forces of an atom of an element, or its replac-
ing power, in terms of the number of atoms of hydrogen
with which it unites or which it displaces, is known as its
valency or atomicity. Even for the same elem.ent this
valency is frequently a variable quantity. The reasons
for this cannot be considered until a later stage, but a few
examples will make clear its importance. H,,0 and H2O2
are both compounds of hydrogen and oxygen ; sulphur
and oxygen form SO, and SO3 ; carbon and oxygen
unite to produce the oxides CO and COg ; FeO, Fe203
1(1
VALENCY.
65
and Fe304 are three oxides of iron ; and of lead and
oxygen we have the compounds PbO, Vb.fi.^, ^h.fi^
and PbOg.
Monads unite with one another in the proportion of
one to one ; diad elements also unite with one another in
the proportion of one to one, but with monads in the pro-
portion of one to two.
If there are six elements whose valencies are indicated
by the Roman numerals, (the common way of marking
it), and whose names are represented by letters, thus,
A*, B", C", D^ E\ F^', they may form combinations as
follows: AgB, A3C, A4D ; C3.,, B2D, EgBg, B3F, D3C4, etc.
It must not be inferred from what is said above that
every element is capable of uniting with every other one.
There are many cases in which tvvo elements have never
been known to unite directly with each other, or to form
a group from the breaking down of higher compounds ;
examples are, oxygen and fluorine, potassium and
sodium, hydrogen and bismuth.
Artiads. — Atoms of elements of even
termed artiads.
.omicity are
Perissads. — Atoms of elements of uneven or odd
atomicity are called perissads.
The sum of the atomicities (f the elements in a com-
pound is always an even number ; and when an element
has more than one valency it changes two degrees at a
time, so that an element can never be both an artiad and
perissad.
it*
■''.]
;.i
i
"I
*"
66
VALENCV.
||:
I' III
III':
%
2.
The following tabic <^nves the valency of the principal
elements. It will be found useful to commit it to
memory : —
(75
-J)
H
o
H
Monads.
DVADS.
Triads.
Trtradh.
PRNTADH.
Hkxads.
Bromine.
Oxygen.
Boron.
Carhon.
Nitrogen.
Sulplmr.
Chloriiii".
Sulphur.
Nitrogen.
Silicon.
Piiosphorus.
Iluoriiie.
Phosphorus.
Sulphur.
Arsenic.
Ilydroyeii.
Arsenic.
Iodine.
Calcium.
Aluminium
Potassium.
Antimoiiv.
Chromium.
yo<Uum.
Copper.
Bismuili.
Cobalt.
Antimony.
Manganese.
Silver.
Magnesium
Gold.
Iron.
Bisnmth.
Iron.
Mercury.
Aluminium.
Lead.
«
Manganese.
Manganese.
Strontium.
Nickel.
Zinc.
Platinum.
Iron.
Tm.
Barimn.
Lead.
Many of the elements do not unite directly with hydro-
gen ; when this is the case their atomicities are calculated
from their union with other elements whose combina-
tions with hydrogen are known.
3. —Radicals.
In many cases a group of atoms acts in combination
and replacement in the same way that single atoms do.
ii!i
CIIKMICAL EQUIVALKXT.
67
An example of this we have already met in the hydroxyl
group HO, which unites with potassium and sodium to
form the hydrates of these metals.
K+H,0+ = KHO+H.
These compound radicals, as they are called, do not
exist in the free state, but when separated in the
decomposition of compounds, they unite either with each
other or with some other substance present.
4.— Equivalent.
Chemical equivalent is a term used to express the
proportions, by weight, in which elements combine with
one another, or displace one another in a compound, one
part, by weight, of hydrogen being taken as the unit.
Thus, the chemical equivalents of oxygen, sulphur,
chlorine and nitrogen are respectively 8, i6, 35 '5, 4"66.
Chemical equivalents must not be confounded with
combining proportions. The former is taken with refer-
ence to hydrogen only ; the latter may be taken with
reference to any other element. Thus from CH4 we get
the equivalent of carbon, but from CO, and CO2, are
obtained the two proportions in which carbon and
oxygen combine.
References for this chapter. — D. & W., 192 ; Tilden, 139 ; Richter, 169 ;
R., 81 and 427 ; R. & S., 95, vol. I.; Wurtz, 226, 233 ; Rem. Th. Chem.,
79 ; Ramsay's Chem. Theory, 80.
'ii-
r "
'i;!
,<'
68
SYNTHESIS OF WATIiR.
CHAPTER XIV.
1.— Synthesis of Water.
When a compound is separated into its constituents
and these determined, the compound is said to be
analysed. Analysis may be of two kinds, — qualitative,
when the operator simply determines what the consti-
tuents are ; quantitative, when he goes further and cal-
culates the proportions by weight or volume in which
these constituents enter into the compound. The oppo-
site of analysis is synthesis. This consists in bringing
together the constituents and treating them in such a
way that they unite to form the compound required.
When water was decomposed by electricity, and it was
shown that hydrogen and oxygen composed it, we had
a qualitative analysis of water ; on the other hand, when
oxygen and hydrogen are caused to unite, and when it
is shown that water is the result of the union, we have
synthesis of water.
2.
Experiments.
I. Take a graduated tube called a eudiometer, fill it
with mercury, and invert it over mercury in a soup-plate
or saucer. Then pass into it a known volume of oxygen
and twice the volume of hydrogen, measuring both at
the same temperature and pressure. Pass a spark from
a Leyden jar, or from a Ruhmkorff's coil, through the
mixed gases, Figure 17. Before igniting the gases, press
the eudiometer firmly on a rubber pad placed at the
COMPOSITION OF STEAM.
69
bottom of the plate or saucer. After I'miitioii examine
the top of the eudiometer with a good lens.
Fid. 17.
2. Repeat this experiment, using equal volumes of the
two gases. Test any gas that remains.
3. Try the experiment again, using twice as much
oxygen as hydrogen. Test as before any gas that
remains.
4. Would it be possible to burn 2 litres of hydrogen in
2 litres of oxygen ? Could the process be reversed and
2 litres of oxygen burned in 2 litres of hydrogen ? What
gas, and what volume of it, would be left in each case?
3.— Ooraposition, by Volume, of Steam.
We have now to find out how many volumes of steam
will be produced by the union of two volumes of
hydrogen and one of oxygen.
Experiment.
Fill a eudiometer one-third full of a mixture of hydro-
gen and oxygen gases — using two volumes of the former
to one of the latter. Cover the eudiometer with a large
tube, into the top and bottom of which pass tiglitly-fitting
corks perforated with tubes, admitting steam at the top
and giving exit to it at the bottom, Fig. 18. The wires
"*(>
■}
70
COMPOS IT ION OP STEAM.
M*
n.
I'
VI
I If
III
from the battery to the eudiometer should pass into the
jacket through its
upper cork. After
the steam has been
admitted, mark the
height of the mer-
cury above that in
the trough, and also
the volume of the
contained gases,then
explode them. Af-
ter explosion de-
press the eudio-
meter, until the mer-
cury in the tube stands the same height above that in
the trough as before. Then measure the volume of the
water-gas (steam) in the eudiometer, and compare this
volume with that of the original mixture.
If we represent equal volumes of oxygen and of
hydrogen by equal squares, and then place in these
squares the first letter of the name of these elements, we
can represent to the eye, by another figure, the volume
of water-gas or steam formed, and the diminution in
volume which occurs after union. Thus :
Fio. 18.
y +
I vol.
o
2 vols.
Steam
How would you account for the change of volume ?
DEFINITE PUOrOUTlONS.
71
CHAPTER XV.
Definite Proportions.
The object of this chapter is to show that chemical
action, whether of coinbination or decomposition, takes
place only between definite weights of the constituents.
EXPKRIMENTS.
1. Into a hard glass tube, A, Fig. 19, introduce a weighed
quantity of copper oxide. A bout one gram is a convenient
portion to work with. Pass a jet of dry hydrogen through
this tube and after all air is expelled heat the tube and
contained oxide to redness. Find the weight of the
remaining copper. From the result of your work
calculate the weight of oxygen that unites with 63*5
parts, by weight of copper.
2. Alter the last experiment by heating the copper
that was left in the tube, in a current of air, and find
the weight of the black substance (copper oxide) that is
formed. Calculate how many parts of oxygen unite with
63 5 parts of copper.
Note. — The student must not expect an absolutely correct result in q\ianlitative work
of this kind. The following notes of an actual experiment indic::ite such a degree of
correctness as may be looked for. The true result is to be foiuul in the average of many
experiments. For this reason the teacher shouhl keep a record of the best results from
year to year.
Weight of empty tube, i9'oo45 grams.
Weight of tube with copper oxide in it, zo'oos grams.
Weight of copper oxide, i"ooo5 grams.
After heating in current of hydrogen, weight of tube and contents, ig'8o2, grams.
Loss of weight, "203 grams ; weig t of copper, '7975 grams.
From this it follows that the weight of oxygen that combines with 63'5 parts of copper
is 1617.
i '■
72
DEFINITE PROPORTIONS.
3. Vary Ex. i by passini'- the hydrogen through a
drying tube filled with lumps of caustic potash before it
enters the combustion tube, then passing the escaping gas
with the resr.lts of the combustion through another
drying tube, also filled with pieces of caustic potash.
This latter tube must be carefully weighed both before
the experiment begins and after it is completed. This
will give the weight of the water formed, and from it the
weight of hydrogen that unites witli 16 parts of oxygen
may be found.
A
''^^'"•^''^ •-
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Fia. 19.
4. Heat in a hard glass tube, closed at one end, a
weighed quantity of silver nitrate crystals. After all
brown fumes cease to come off, find the weight of the
resultant solid — pure silver in this case. How many
parts of nitrogen trioxide are in union with 108 of silver?
5. Weigh a beaker and a sheet of mica (better, plati-
num), place on the mica a weighed piece of magnesium
wire, invert ^he beaker over Ihe wire, then heat the mica
until all combustion ceases. Weigh again.
6. Close the end of a piece of hard glass tubing
about 4 cm. long, and yi cm.
2
^
v_.
in diameter, A, Fig. 20. Place
in this a weighed quantity
of red oxide of mercury,
then pass the end of this
tube into a larger tube B,
and afte: weighing the whole, heai: until the red powder
Fio. -M.
DEFINITE PHOPOllTIONS.
73
has all disappeared, weigh again and calculate from
your result the proportions in which mercury and
oxygen unite in mercuric oxide, HgO.
7. In one beaker prepare some dilute nitric acid, and
in another, a solution of potassic hydrate. Pour about 10
cc. of the dilute acid into a graduated measuring glass or
tube, add enough of the potash solution to just neutralize
this. Observe how much of the latter has been required,
then pour the whole into a weighed evaporating dish,
evaporate to dryness, but without boiling. Find the
weight of the solid substance in the dish. Repeat, at
least twice, using different volumes of the acid, and
determine if a fixed quantity, by volume, of the acid
requires a fixed volume of the alkaline solution to
neutralize it, and if it produces relatively an unvarying
weight of the salt.
8. Repeat the last experiment, but use hydrochloric
acid, and carbonate of sodium.
The results of such experiments as these should show,
beyond question, that substances do not take part in
chemical actions in random proportions.
These results lead to one of the fundamental laws of
chemistry, which is, that each element (or compound
radical) unites with other .substances in certain fixed and
invariable proportions by weight. These proportions are
either multiples or sub-multipies of the atomic weight in
the case of an element. In the case of a compound,
these proportions are either the molecular weight or
some simple multiple of it.
One of the chief differences between a solution and a
chemical compound is that, in the former, varying
p'Wii
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74
CHKMICAL CALCULATIONS.
quantities of the substances may take part in the action.
When a centigram of salt is dissolved in a litre of water
the result is just as truly a solution as when lo grams of
ihe salt are used. Another difference is that solutions,
unless agitated artificially, are seldom homogeneous ;
while chemical combination throughout a mass must be
absolutely the same in every part.
See Tilden, 80 ; D. & W., 39 j R., 14; Wmiz, 3 ; R. & S., 63.
CHAPTER XVI.
Some Chemical Calculations.
It has been said in the chapter on symbols and
equations (vil), that in any chemical equation the sum of
the atomic weights of the elements on one side must
equal the sum of those on the other. Some applications
of this principle will now be made.
1.
When hydrogen is prepared from zinc and dilute
sulphuric acid, the following equation expresses the
reaction that takes place : —
Zn + H,SO,-f H.p-ZnSO^ + H.+ HgO.
From this we see that 65 parts, by weight, of zinc, 98
of sulphuric acid and 18 of water, yield 161 parts, by
weight, of zinc sulphate, 2 of hydrogen and 18 of water.
The water evidently takes no part in the chemical action,
so far as the evolution of hydrogen is concerned, and
CHKMICAl. CALCULATIONS.
<•)
SO, in our calculation, it may be neglected. We are now
ready to work out some numerical problems of which
the following are examples : —
1. Suppose we required i r2 grams of hydrogen, how
much zinc and how much sulphuric acid would be used
up in obtaining it ?
65 " parts " by weight of zinc and 98 of sulphuric acid
yield 2 of hydrogen, then I r2 "parts" by weight of
hydrogen come from 65 x -¥^ of zinc and 98 x ^^'^ of
sulphuric acid, but a part may be any unit of weight
whatever, since it i. a general term, and is used as the
unit throughout the problem, hence 1 1'2 grams of hydro-
gen come from --f j^-^-^ of zinc and --.j^--^ of sulphuric
acid.
2. Iron filings treated with hydrochloric acid yield
hydrogen according to the equation Fe + 2HCl = FeCl.,H-
If 40 grams of iron were used in the experiment, how
much pure hydrochloric acid should be taken, and how
much hydrogen would be obtained ?
Solution :—
56 parts by weight of iron, and 73 of hydrochloric acid
yield 127 of chloride of iron and 2 of hydrogen; then
40 of iron would require *[| of 73 parts of acid and
would yield 5|fXi27 of the chloride, and |J} X 2 of
hydrogen.
3. When oxygen is prepared from chlorate of potassium
and manganese dioxide the equation is :
KClO,+ MnO,-KCl-f MnO,+ 30, hence 122-6 parts
by weight of potassic chlorate, heated with Sy parts of
t
.A
II
76
CHEMICAL CALCULATIONS.
manganese dioxide, yield 87 parts of dioxide, 48 of oxy-
gen and 74'6 parts of chloride of potassium. Here also
the dioxide is unchanged and may be neglected. Then
X grams of the chlorate will yield ttIvo ^4^ grams of
oxygen, and x grams of oxygen may be obtained
from -jy-x I22'6 grams of chlorate.
!:■
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2.— Questions and Exercises.
1. Five grams of sodium are placed on water, and the hydrogen
resulting from the chemical action is collected, afterwards the
water is evaporated and the whitv-i salt that is obtained is weighed.
Theoretically, how much hydrogen and how much of this salt
would there be ?
2. If 5 grams of potassium had been used in the last question
what would then have been the answers ?
3. If 5 grams of copper oxide, CuO, were reduced in a current of
hydrogen, what products would be obtained and how much of each
by weight ? How much hydrogen by weight would be required to
complete this chemical action ?
4. If 10 grams of lead peroxide, PbOg, are reduced to lead oxide,
PbO, how much oxygen would be given off in the operation, and if
this oxygen immediately united with hydrogen, how many grams
of the compound would be formed .-*
5. An excess of iron filings is treated with 50 grams of a solution
of hydrochloric acid, containing 25% by weight of pure acid,
how many grams of hydrogen will be produced, and how
many grams of the compound of iron and acid will be formed ?
6. If 50 grams of chlorate of potr.sh w^re entirely decomposed
by heat into potassium chloride and oxygen, and the latter collected
over water, then if a jet of burning hydrogen were passed into the
jar and kept there until all the oxygen was used up, what would be
the weight of the resultant compound ?
Mv
COMBUSTION.
77
7. 10 grams of water arc decomposed by electricity, 10 grams are
decomposed by the action of sodium, 10 grams are decomposed by
the action of potassium, and 10 grams are converted into steam and
passed over red hot iron filings. How much hydrogen would be
obtained in each case ?
CHAPTER XVII.
Combustion.
It has been customary to classify bodies as combus-
tible or as supporters of combustion. The object of this
chapter is to show that there is no such division of
substances; that combustion is a chemical action in which
at least two substances are equally concerned, and that
the phenomena of combustion are produced by the
energy with which the chemical union goes on.
1.
When two substances enter into chemical union, o*- act
upon each other chemically in any way, the action is
usually accompanied by change of temperature, change
of volume, or by both these phenomena. When the
combination of two substances is accompanied by light
and heat (the lii;ht as a consequence of the heat) there
is said to be either gloiving or combustion — glowing, if a
mass of solid matter simply becomes red hot — combus-
tion, either if a flame is produced, as in the case of any
burning gas, or if the solid, while glowing, gradually
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78
BURNING.
changes into an incandescent gaseous compound, as in
the case of charcoal.
Burning or comhustion is generally caused by some
substance uniting with oxygen, hence the popular
assertion that oxygen is a supporter of combustion.
Examples of this may be found in the experiments
under oxygen. As some of the following experiments
will show, however, oxygen is not necessary for com-
bustion.
2.
Experiments.
I. In Fig. 21, A is a glass tube about 3 or 4 centi-
metres in diameter, drawn to about half that diameter at
Fig. 21.
the upper end and closed with a perforated stopper at
the lower end ; C leads either to an ordinary gas cock
or to a gas holder. After the air is expelled from A, fire
the escaping gas at the top. B is a gas bag, or gas holder,
filled with oxygen, and the delivery tube D is gradually
lowered until the nozzle "passes into the interior of A, at
the same time the oxygen is being driven out of B by
nil
COMBUSTION.
79
pressure. The gas that escapes throuG^h D o. ght to
take fire at the top of A just as the tube D is lowered
through the mouth of A, and continue to burn in the
interior. Try the experiment using hotli coal gas and
hydrogen passing througli C, and both oxygen and air
passing through D. Next, reverse the operation.
2. Heat some sulphur in a test-tube, until the tube be-
comes filled with the vapour, then lower into it some copper
foil or a braid of very fine copper wire.
3. Fit up an apparatus, as in Fig. 40; put into the flask
equal portions of common salt (sodium chloride) and
manganese dioxide, then pour on this some sulphuric
acid and gently heat the whole. Soon a greenish yellow
gas will come off which will collect in the jar at the
right. This gas is chlorine. As soon as the jar is full
of the gas, as may be seen by the colour, drop into it
some powdered antimony or arsenic.
4. Prepare another jar of chlorine and lower into it a
jet of burning hydrogen.
5. Use the apparatus. Fig. 21, but cause a jet of chlorine
to pass through D, and hydrogen through C.
These experiments are sufficient to show that "com-
bustible " and " supporter of combustion " hardly express
the relations between the pairs of substances which here
enter into combination with each other. A better way
to state the case would be, that two kinds of matter which
act on each other as do oxygen and hydrogen, or hydro-
gen and chlorine are mtitiiaUy combustible ; or that if
there are two gases, A and B, of which A burns in pre-
sence of Bj it is equally true that B burns in presence
of A.
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^
80
THE ATMOSPHERE.
3.— Questions.
1. When coal is heated to a certain degree it unites with part of
the air. In this case should we say that the coal burns, the air
burns, or tiiey both burn ?
2. Air burns in coal gas, coal gas burns in air, which is the
supjjorter of combustion ?
CHAPTER XVIII.
l.~Is Air an Element, a Compound, or a Mixtures
of Gases ?
It has been learned that water, one of the substances
of most common occurrence, is a compound of two gases.
We are now in a position to study another very common
substance — the air. The first step will be to determine,
if possible, whether it is an element, a compound, or a
mixture of two or more gases.
tiiM
Experiments.
I. Cover a cork, about two inches in diameter, with a
piece of tin and float on a soup-plate full of water.
Take a piece of phosphorus about the
size of a pea, and place it on the cork.
Now set fire to the phosphorus, and
then cover it quickly with a beaker or
small bell jar, placed mouth down-
wards, as in Fig. 23. Allow it to stand
Fig. 23. ^^^^ ^^'' 1 5 or 20 minutcs.
Composition of air.
81
After the white fumes liave entirely disappeared, lower
the plate and jar into water until the water stands at the
same height on the inside and the outside of the jar ;
then test the gas by (a) passing a lighted splinter into it,(/;)
passing a little of it through lime water, (c) driving
some of it into a test-tube inverted over mercury, then
passing pyrogallate of potash into this tube. In the last
case, observe if a material darkening of the liquid takes
place after it enters the tube, or if there is any consider-
able decrease in the volume of the gas.
2. Invert a test-tube of air over mercury and pass into
it some pyrogallate of potash to the depth of a couple of
centimetres.
3. Wet the inside of a bell jar with water, and drop
into it some fine iron filings. Then shake the jar so that
the inside may become closely sprinkled with the filings.
Place the jar, mouth downward, over a soup-plate filled
with water, and allow the whole to stand for a day or two.
Carefully invert the jar by slipping the hand under its
mouth, and as it is turned mouth upward, allow the water
in it to run to the bottom ; thus an influx of air is pre-
vented. Test the gas in the jar with a burning splinter.
2. Questions.
r. Pvrogallate of potash turns brown in the atmosphere and takes
part of it up. What is the inference ?
2. Is air a mixture of at least two substances, a chemical com-
pound of two substances; or a single element ? What reasons have
you for your answer ?
Explanation. — The gas that was left in the jar after
the phosphorus burned, and the fumes were absorbed,
was almost entirely nitrogen.
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82
COMPOSITION OF AIR.
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3. -Volumetric Composition of the Atmosphere.
The quantities, by volume, in which oxygen and nitro-
gen are mixed in tlie atmosphere may be determined in
two ways, — (i) by causing the oxygen to unite with some
substance to form an oxide which may be got rid of by
solution ; (2) by causing the oxygen to unite with hydro-
gen to form steam, whose volume when condensed to
water may be neglected in the calculation.
EXI'KRIMENT.
Burn a piece of phosphorus as in experiment i of
section i, but do not allow bubbles of gas to escape,
and after all white fumes have disappeared measure
the condensation which the gas in the jar has under-
gone. This may be done by marking the level of the
water in the jar just when the combustion of the phos-
phorus is completed ; then after all oxide is dissolved
depress the jar in a tank until the water stands at the
same level inside and outside of it ; again mark the
height of the water. Now invert the jar, fill it with water
and measure the contents in a graduated glass. Measure
also the volume between the marks and compare this
volume with the volume of the whole jar.
if ^
\
I
4.— Additional Exercises.
1. Try if a jet of burning hydrogen will continue to burn when
plunged into nitrogen. Reverse the process and try if a jet of
nitrogen will burn in hydrogen.
2. Will nitrogen burn in oxygen, or oxygen burn in nitrogen ?
3. Is there any free phosphorus left on the tin after combustion
has ceased in experiment i , sec. i .'' Would the use of a larger
COMPOSITION OF AlU.
83
quantity of phos|)hoius cause any variation in the volume of gas
remaining after chemical action has ceased?
4. Take a graduated tul)e aliout 30 centimetres long and 3 centi-
metres wide, closed at one end by a stop-
cock, or by a piece of rubber tube and
a pinchcock. Invert it over a vessel
9 or 10 centimetres in de|)th, filled with
water. Fix the tul)e in a support, taking
care that the water stands at the same
level on the inside as it docs on the out-
side of the tube. Then pass uj) to the
top of the tube a piece of phosphorus
attached to a copper wire, as in l''ig. 24,
To attach the wire to the phospliorus, fuse
it under water in a test-tube, introduce
the end of the wire into it, and then let it
cool. Leave the whole twenty-four hours,
then withdraw the phosphorus and adjust
the level of the water inside and outside
the tube ; read off the volume of the gas
remaining in the graduated tul)e. The
volume of gas at the beginning and at the end of the experiment
must be reduced to that at standard temperature and pressure.
5. Pass into a eudiometer over mercury a measmed volume of air,
dried by passing it over lumps of calcium chloride in a tube, then
pass in half as much dried hydrogen, and exi)lode the mixture.
Note reduction in volume of the gases.
Explanation.— It has been learned before that hydro-
gen and oxygen unite in the proportion by volume of 2
to I, hence one-third of the reduction of the volume of
the mixed gases is due to the oxygen of the air uniting
with the hydrogen. From careful determinations made
in this way, oxygen has been found to form 2i/^ by
volume of the air, and nitrogen 79%. By weight, the
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AIR A MIXTURE.
percentage is somewhat different. In terms of volumes
of hydrogen, we have
21 vols, of oxygen become by weight 21 x i6= 336
79 " nitrogen " " 79x14=1106
1442
That is -Yi^^ of the whole is oxygen, = 23-3%.
In addition to the experimental reasons given in this
chapter for considering air to be a mixture, the following
are further proofs of the same fact : —
(i) A mixture of oxygen and nitrogen may be
made which cannot be distmguished from air,
either chemically or physically,
(2) The gases are not present in the proportions
corresponding to their atomic weights, or in
simple multiples of their atomic weights, yet
they could not be present in other proportions
if air were a compound.
(3) Air is soluble in water, and if the air dissolved
in water be expelled by heating or by the air
pump, and the gases collected, the oxygen
and nitrogen are found not in the same pro-
portion in which they were before solution ;
more of the latter gas in proportion is given
off from the solution ; this, of course, would
be impossible in the case of a compound.
(4) Air that has been collected at different places
gives, on analysis, slightly different propor-
tions by weight.
NITROGEN.
85
CHAPTER XIX.
r;
1— Nitrogen- Its Properties and Preparation.
Experiments.
I. The simplest method of obtaining nitroj-en is by
burning phosphorus, or some other combustible which
will form a readily soluble oxide, in a- bell jar over
water. The student should repeat the first experiment
of the previous chapter. After the white fumes have
been absorbed, depress the jar in a trough of water until
the level of the water inside is the same as that outside.
Then remove the cork and test the gas that remains with
htmus paper, with a glowing splinter, and with a jet of
burning hydrogen.
2. Prepare a hydrogen apparatus with a delivery tube
as shown in Fig. 25. After the hydrogen is burning
freely at the mouth of
the tube invert a gas
cylinder over it so that
the mouth of the cylin-
der will be under water
in the dish. Watch
closely what occurs.
The instant the flame
goies out disconnect the
delivery tube from the
flask. After the water
has ceased to rise in the cylinder slip a glass cover under
it and turn it mouth upwards, but do not let the water
Fio. 25.
I*'
23
,..r
86
NITROGEN.
i::
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run out, else air will become mixed with the gas in the
cylinder. Test this gas as in the previous experiment.
Tabulate the properties and appearance of nitrogen,
oxygen, hydrogen, and air.
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Fig. 2&
2.— Additional Exercises.
1. The apparatus represented in Figure 26 may be used for pre-
paring nitrogen by passing air
over red-hot copper. A is a U
tube filled with calcic chloride,
B is a straight tube filled with
fine, bright copper filings, and
C is a large -mouthed bottle
used as a gas holder, one of
its tubes passing through the
cork, the otiier passing to the
bottom of the bottle ; D is a
piece of rubber tubing attach-
ed to the latter of these tubes
in order to convert it into a
syphori, and draw ofif the water from the bottle. E is a spring clip
which may be kept slightly open by putting a small wedge in it.
On starting the experiment, the bottle must be full of water. When
the copper has been made red-hot, the syphon must be made to act
very slowly by regulating the clip, and as the water flows out of the
bottle, air is drawn through the U tube and passes over the red-hot
copper. The nitrogen is collected in the bottle. At what stage in
this experiment would air begin to pass out of the tube ?
2. Heat some potassic nitrate in a tube or crucible until it gives
readily an alkaline reaction. It has then changed into the nitrite.*
Mix this salt with ammonium chloride, heat them in a test-tube
and collect over water the gas given off.
KN02 + NH4C1 = KC1-H2H20-1-2N.
Explain the chemical actions mvolved and classify them.
*The alkalinity is due lo the nitrite beginning to inidergoaiurther decomposition, by
which the oxide of pota&bium is lorined, and tt is a strong alkali.
r M
COMPOUNDS OF NITROGEN,
87
3.— Notes on Nitrogen.
Nitrogen: symbol N.; atomic weight, 14.; molecular
weighty 28 ; molecular volume , 2.
On account of its chemical inertness, n'trogen forms
many unstable compounds, and of these a large number
are liable to sudden and often very violent decomi-
position : for example, nitro-glyceri.ie, and chloride of
nitrogen.
Nitrogen occurs free in the air, of which it forms by
volume 79%, and by weight yf/^ ; it is also a constituent
of almost all organic substances, and enters into the
composition of a large number of inorganic compounds.
It serves to dilute the oxygen in the atmosphere, but
takes no direct part in the support of animal life. It is
doubtful if it acts directly in support of vegetable life,
neither does it enter into direct chemical combinations
except very rarely and by special treatment.
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CHAPTER XX.
Compounds of Nitrogen and Oxygen.
Nitrogen and oxygen do not unite directly to form
compounds as do hydrogen and oxygen ; there are.
however, five oxides of nitrogen known, all of them
being obtained by the decomposition of other com-
pounds.
1
i 1^
88 NITROUS OXIDE.
These oxides are :
Nitrous oxide or nitrogen monoxide NgO,
Nitric " " dioxidt NO,
Nitrogen trioxide NgOg.
Nitrogen peroxide or nitrogen tetroxide NOo,
Nitrogen pentoxide
§ 1.— Nitrous Oxide
Experiments.
I. Put 25 grams of commercial ammonic nitrate,
NH4NO3, into an oxygen generating apparatus, con-
nected with three bottles, as in Fig. 27. The first bottle
should contain a solution of ferrous sulphate, the second
Fig. 27.
a solution of caustic potash, and the third, water. Heat
the nitrate gently, and nitrofren monoxide will be given off.
Thus prepared, the gas will be found mixed with nitrogen
^
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NITROUS OXIDK.
89
dioxide and chlorine gas. The first will be removed by
passing through the ferrous sulphate solution, and the
second by passing through the caustic potash solution.
If the nitrate be chemically pure, the wash bottles
may be omitted. The reaction may be thus represented:
NH,N03 = 2H,0+N,0.
2. Collect several jars of the gas over warm water and
perform the following experiments : —
(a). Plunge a lighted taper into the first jar ; also test
it with a glowing splinter, as in the case of oxygen.
(d). Burn a piece of phosphorus, carbon or sulphur in
another of the jars.
(c). Explode a mixture of the gas with hydrogen.
(d). Place another jar, mouth downward, over co/d
water, and then shake. Let it stand for 24 hours.
Vary this experiment by filling a 4 or 5 inch test-tube
with the gas, put a little water in it, close the mouth
tightly by putting the thumb over it, then shake the
water up with the gas, invert the tube, dip the mouth
under water and remove the thumb. Test the water
that rose in the tube with litmus.
3. Try the effect of dry nitrous oxide upon litmus.
4. How could you distinguish nitrous oxide from
oxygen ?
5. Using the apparatus of Fig. 25, burn a jet of
hydrogen in nitrous oxide. What are the substances
formed ? Write equation. What change in volume, if
any, takes place during the combustion? Explain.
6. Try if a mixture of hydrogen and nitrous oxide
will explode when an electric spark is passed through it.
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90
NITRIC OXIDE.
■•': f
2. —Notes on Nitrous Oxide.
Nitrous Oxide : formula, N,fi ; molecular weighty /j.4. ;
vapour density, 22.
" Laujrhing gas " is an old name for this compound.
It derives this name from the fact that many persons
after inhaling a mixture of the gas and air are compelled
to laugh — nolens volens. On inhaling more of the gas
temporary unconsciousness is produced ; it is therefore
frequently used as an anaesthetic for minor operations in
surgery.
Nitrous oxide is soluble in cold water to the extent of
130 per cent of its own volume ; it may be condensed to
a liquid by a cold of o" C, and a pressure of 30 atmos-
pheres. Liquid nitrous oxide when mixed with carbon
disulphide, CS.^, forms a freezing mixture capable of
producing a cold of — 140° C.
3.— Nitric Oxide or Nitrogen Dioxide.
Experiments.
I, Place some copper filings in a hydrogen generating
apparatus similar to that in Fig. 7, add some warm
water, and then pour down the funnel tube some strong
nitric acid. The gas that first forms should be allowed
to escape. The reaction may be thus represented : —
3Cu + 8HN03=3Cu(N03)2+4H,0 + 2NO.
The reaction here represented is really the result of
two separate and successive ones, thus : —
(i). Cu + 2HN03=Cu(N03)2-l-2H.
(2). 3H-hHN03=NO+2H20.
The explanation of this will be found in chap, xxi, sect. 5.
Nitric oxide.
91
Collect, over water, two bell jars and two large test-
tubes full of the gas.
2. Place one of the tubes that is full of gas, mouth
downwards, over a small quantity of water in a dish,
then pass air, a little at a time, into the tube.* Test the
water with litmus both before the air is passed in, then
again after the brown fumes have disappeared.
Unless the operator is careful in this experiment a
wrong result will be obtained. When the jar is filled
with the gas it should be removed to a clean plate with
a little wp.ter on it, so that the gas will be tested and not
a solution of the brown fumes formed by bubbles of it
coming in contact with air.
3. Test the gas in one of the bell jars with a glowing
splinter, a blazing splinter, a burning taper, a piece of
slightly-ignited phosphorus, a piece of brightly-burning
phosphorus.
4. Lift one of the test-tubes full of the gas and place
it, mouth downwards, in a vessel containing a cold solu-
tion of copperas (ferrous sulphate), FeSO^.
Vary this experiment as follows : —
Pour some well-cooled solution of ferrous sulphate,
FeS04, into a beaker full of the gas ; then hold the hand
over the beaker's mouth and shake vigorously. Note
the two phenomena that occur.
* The air may readily be driven into the tube by using an empty flask fitted up
like the one in figure 25. When water is poured down the funnel, air is forced out
through the delivery tube.
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92
MTKIO OXIDE.
4.— Notes on Nitric Oxide.
Nitric oxide: formula^ NO ; molecular weighty jo ;
vapour density, 75.
Nitric oxide condenses to a liquid at — ii^C. and a
pressure of 104 atmospheres. It does not unite with
water to form an acid. One test for this gas is its
reaction with air or free oxygen ; another is that with
a solution of ferrous sulpiiate a dark ring or layer is
formed on the liquid, as seen in ex. 4, in the preceding
section.
iiif
5.— Composition by Volume of Nitrous Oxide
and Nitric Oxide.
Prepare a hard glass tube, bent as
A in the Fig. 28. Fill this with
washed nitrous oxide g?.?i, having
previously dropped into the tube a
piece of sodium, or of potassium,
about as large as a pea. Dip the
mouth of the tube, when filled with
gas, under mercury, and by jarring
it, get the sodium into a position just below A. Then
heat it strongly. The hot sodium decomposes the
nitrous oxide to form oxide of sodium and the nitrogen
is left. The volume of the nitrogen should be the same
as that of the original gas.
Nitric oxide may be decomposed in the same way,
but the volume of nitrogen in this case is only one-half
that of the oxide taken,
Fia.28.
NlTROOEN TRIOXIDB.
93
G.— Questions and Exercises.
1. Pass nitrogen into nitric oxide.
2. Is it the air as a whole, or one of the constituents of it, that
causes the brown coloured gas to appear with nitric oxide ?
3. What reasons have you for believing that nitric oxide does not
burn ?
4. Will dry nitric oxide change litmus ?
5. Pass oxygen into a jar of nitric oxide over water very slowly
so that the brown furies may disappear as rapidly as formed.
Account for the result obtained.
6. When air was passed into a jar of nitric oxide over water until
brown fumes ceased to appea.'- and be dissolved, what remained.^
Apply tests to find out if your conclusion is a correct one. Was the
result different when oxygen was passed in ?
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7.— Nitrogen Trioxide.
Experiment.
Fit a Florence flask with
a cork and delivery tube,
and place on a retort stand,
as in Fig. 29. To the de-
livery tube attach a U tube,
immersed in a freezing
mixture of salt and snow.
Connect the other end of
the U tube with a glass F"*- 29-
tube leading to a vessel containing ice-Water. Place
10 grams of starch in the flask and cover with nitric
acid. Gently heat the generating flask and nitrogen tri-
oxide will be plentifully produced, part of it being con-
densed in the U tube, and the remainder passing on into
the ice-water.
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94
NITROOKN TETROXIDK.
!
I
Instead of starch, white arsenic, A ^.Oy, may be used.
The reaction in this case may be thus represented :
2HNOa + Asa03 + 2H20=:N203 + 2M>,As04 (Arsenic acid).
Notice the colour of the gas. It is condensed to a
h'quid by a temperature of — i8°C. Try to collect some
of the gas over water. Has it any smell ?
The gas, as condensed in the U tube, is green in
colour. This is owing to nitrogen peroxide being mixed
with it. If the generating flask be disconnected and a
current of nitric oxide passed through the U tube, the
brown gas that passes off, if again condensed, will be
indigo blue in colour ; this will be pure nitrogen tri-
oxide.
By using for a condenser a piece of thick glass tubing
drawn out, as shown in Fig. 30, the liquid may be pre-
served ; for by using a blowpipe
flame the tube may readily be sealed
Ifi^ at A and B ; and internal pressure
Fig. 30. will then prevent the fluid from
evaporating. The tube must be strong enough, however,
to withstand the pressure.
7.— Nitrogen Tetroxide or Nitrogen Peroxide.
This gas is prepared by heating lead nitrate and con-
densing the gas, as in the case of the trioxide. It is the
substance most largely formed when nitric oxide comes
in contact with air. In the preparation, trioxide and
tetroxide are mixed, but the former may be changed
into the latter by a current of oxygen. The liquid per-
MULTIPm PIIOPOKTIONS.
95
oxide is yellowish or brownish in colour. This substance
has the formula NCX and is introduced here because
of theoretical considerations.
There is still another oxide of nitrojren, viz., the pent-
oxide, N.Pj, but as it is difficult of preparation and of
no practical value its study may be omitted.
Fill out the followinj^ schedule: —
API'RARANCK.
Solubility in
Watkr.
ACIDITV OF Hy-
DKATKH.
PhVJICAL STATR
AT 0°.
N,0
N
N,0,
»
N 0.,
•
8.— Law of Multiple Proportions.
These oxides of nitrogen illustrate the Law of Multiple
Proportions in chemistry. Beginning with the lowest
oxide and going to the highest, there are successively
one, two, three, four and five volumes of oxygen united
with two volumes of nitrogen. Expressed in another way,
the quantity of oxygen, which is the variable element
here, is an integral multiple in every case, both of unit
volume and of atomic weight. The relative quantities
of oxygen are in the ratio of the numbers i, 2, 3, 4 and 5,
and these are the only proportions in which the elements
can be made to unite.
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96
NITKOUS ACID.
■Jll
CHAPTER XXI.
1.— Acids of Nitrogen.
Nitrogen, in union with hydrogen and oxygen, forms
two well defined acids that have the formula HNOg and
HNO3, and are named nitrous and nitric acids respec-
tively. Some salts corresponding to a third acid, hypo-
nitrous, are known, but the acid itself has not been
isolated ; and, as its salts are somewhat rare and of little
value in elementary work, it will be passed over with
the remark that, if separated, its formula would be HNO,
and that its salts have the composition MNO when M is
a monad base.
2.— Nitrous Acid.
Experiments.
1. Pass nitrojjen trioxide into cold water and test for
acid properties.
Definition. — An oxide which unites chemically with
the water, and thus forms an acid, is called an anhy-
dride.
2. Pass nitrogen trioxide into a solution of potassic
hydrate until it is neutral, then evaporate ; the salt
obtained is potassic nitrite.
3. Add a {^v^ drops of nitrous acid to a solution of
potassium permanganate.
3.— Questions and Exercises.
I. Nitrous acid is an unstable compound decomposing, upon
standing, into nitric acid, nitric oxide, and water, thus : —
3HN02 = HNO;5+2NO + H20.
NITRIC ACID.
97
This may be shown by filling a bottle with nitrous oxide within an
inch of the top, tightly corking it and letting it stand for a couple of
days. On removing the cork, brown fumes are formed. After
standing for some time longer, it will answer to the test for nitric
acid.
2. Potassium permanganate is very readily broken up by sub-
stances that absorb oxygen. How can you account for the result
observed in experiment 3, sec. 2?
2KMn04 + 5 H NO. = 2MnO + sHNOg + KjO.
3. Heat some nitrous acid solution, then test the residue for nitric
acid.
4. Water tainted with sewage always contains nitrites and nitrates
in solution. The common test for this is to pour some of the water
into a weak solution of permanganate of potash and watch for
decoloration. How do you explain this chemically.'*
5. When nitrites are acidulated with acetic acid they give a white
precipitate with nitrate of silver. Nitrites are soluble in water.
Try nitrite of potash, as prepared in ex. 2, sec. 2, for these reactions.
4.— Nitric Acid.
Experiments.
I. Put into a tubulated glass retort 30 grams of pow-
dered nitrate of potash, KNO3, ^"^ ^" equal weight of
strong sulphuric acid, H2SO4. Place the end of the
retort in a flask which
is made to float on a
basin of water as in
Fig. 31. Apply heat
to the retort. Soon a
yellowish colored liquid
distils over and is col-
lected in the cool flask.
The reaction may be Pio. 31.
represented as taking place in two successive stages, the
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NITRIC ACID.
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first requiring a comparatively low, the second a high,
temperature.
(«.)2KN03+H2S04=HKS04+HN03 + KN03.
On increasing the heat more acid comes off, the second
reaction being represented as follows : —
{b) HKS04 + KN03 = K2S04 + HN03.
Sodic nitrate, NaNO,^, may be used instead of potassic
nitrate in the preparation of nitric acid ; in fact sodic
nitrate is generally used when this acid is to be manu-
factured on a large scale.
2. Heat a few drops of the acid until nearly boiling,
then hold close to its surface a piece of glowing charcoal.
Vary this by heating strongly some fine charcoal dust,
then dropping on it some strong nitric acid.
3. Warm a few drops of the acid in a small evaporat-
ing dish, then drop into it a bit of phosphorus.
4. Immerse some undyed wool, silk or other organic
substance in a little of the acid.
5. Add a few drops of the acid to a solution of indigo.
6. Place some copper filings in the bottom of a test-
tube, and then pour in some of the acid. When all
action has ceased, evaporate to dryness the solution
which has been formed.
5.— Notes on Nitric Acid.
The anhydride of nitric acid is nitrogen pentoxide,
N2O5.
Nitric acid is said to be a powerful oxidizing agents
that is, it readily yields its oxygen to substances which
TESTS POR NITKIC ACID.
90
have an affinity for that element. This oxygen, at the
moment of its liberation from the acid, is said to be in
its nascent state.
2HN03=H20 + 2N02 + 0.
Nitric acid is a strong monobasic acid. When metals
act on it, nitrates are formed by the replacement of
the hydrogen by the metal, but the nascent hydrogen at
once acts on part of the nitric acid present, forming with
it, according to accompanying circumstances, nitrous acid
or one of the oxides of nitrogen, or sometimes even
reducing it to ammonia.
The brown fumes that arise when nitric acid is being
prepared are caused by its partial decomposition ; this
occurs at the boiling point of the acid, about 68" C.
Similar fumes are always present over very strong nitric
acid contained in a glass-stoppered bottle.
In preparing the acid, hydro-potassic sulphate, KHSO^,
the acid salt, is first formed ; but at a high temperature
the further reaction resulting in the formation of the
neutral potassic sulphate, K^SO^, takes place. The heat
necessary for this causes the decomposition of the acid
formed, however.
:i'
6.— Tests.
1. Nitric acid heated with copper filings gives off
brown fumes of NgOg and NO.,.
2. Dissolve a few crystals of ferrous sulphate, FeS04,
in water in a test-tube. Add a few drops of sulphuric
acid and allow the whole to cool. Then turn the test-
tube sideways and gently pour nitric acid, or a nitrate in
100
QUESTIONS AND EXERCISES.
i-l":;
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11
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solution, down its side. The phenomenon which results
will always enable us to recognize nitric acid or a nitrate.*
3. Nitric acid bleaches indigo solution.
7.— Questions and Exercises.
1. Zinc treated with dilute sulphuric acid gives hydrogen. What
is produced when nitric acid is substituted for sulphuric ?
2. C + 2HN03 =
4PH-ioHN03 =
Complete these equations.
3. Nitric acid acts on copper and forms the salt cupric nitrate
Cu(N03)2 ; find out whether it acts similarly on other common
metals such as lead, zinc, iron and mercury.
4. The principle of atomicity is employed in writing the formulas
of salts from nitric acid by replacing one atom of the hydrogen
of the acid with one atom of a monad metal ; two atoms of the
hydrogen of the acid in two molecules, with one atom of a diad
metal, and so on. For example :
Acid.
Salt.
Nam^ of Salt.
HNO3
AgNOs
Silver Nitrate.
2HNO3
Cu(N03)3
Copper Nitrate.
3HNO3
Bi(N03)3
Bismuth Nitrate.
In the same way symbolize the salts which nitric acid may form
with the following metals : Potassium, calcium, lead.
5. Explain the action of nitric acid on a solution of sulphate of
indigo.
6. Twenty grams of sodic nitrate are heated with an excess of
sulphuric acid, assuming that the acid sulphate of sodium alone is
•Another method of performing this experiment is to drop a little nitric acid or a
solution of a nitrate into a porcelain evaporating dish, add a little sulphuric acid, then set
the dish on ice or ii cold water until cool. After it has cooled, drop into it some solution
of ferrous sulphate. If a precipitate is thrown down by the sulphuric acid, filter and test
the filtrate.
NITRATES.
101
formed, and that the acid produced is led into a solution of potassic
hydroxide, what salt would be formed, and how much of it ? Write
equations for the reactions.
8— Nitrates.
Experiments.
1. Test as many nitrates as you can find for solubility.
2. Heat different nitrates, and test them by the ordi-
nary method to ascertain whether oxygen is given off;
then sprinkle on them some powdered charcoal, also
some sulphur. Do all the nitrates give off brown fumes
when heated ?
9.~Notes on Nitrates.
All nitrates are soluble, and yield oxygen when heated,
thus,— KN03=K NO2 + O.
To this ammonium nitrate, NH4NO3, seems to be an
exception, probably on account of being composed of
two unstable radicals. When gently heated it yields, as
we have seen, N2O + 2H2O; but when strongly heated
breaks up into water, oxygen and free nitrogen, the
oxygen being first given off, thus :
NH,N03 = NH,N02 + = 2N + 2H20 + 0.
The nitrates of the metallic bases, when heated, form
the oxide of the metal, free oxygen and nitrogen per-
oxide, thus :
Pb(N03)2=PbO + 2N02 + 0.
Those of the alkaline bases lose oxygen on heating,
become first reduced to nitrites and finally to the oxides
of the metals.
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102
ADDITIONAL EXERCISES.
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A substance supposed to be a nitrate may be tested by
treating it with sulphuric acid, which frees the nitric
acid, the latter may then be tested for, as in sec. 6.
10.— Additional Exercises.
1. In the test for nitric acid by the use of ferrous sulphate, the
brown layer is due to the formation of nitric oxide with ferric
sulphate. Compare Ex. 4, Sec. 3, Chap, xx :
2HNOy + 3HoS04+6FeS04 = 2NO + 3Fe2(S04)3+4H20.
If nitric oxide be prepared and passed (i) into warm solution of
ferrous sulphate, (2) into cold solution, the brown ring will be
found in the latter case only, thus showing that the oxide is soluble
in warm sulphate solution.
2. Investigate some of the properties of potassic nitrite by
{a) throwing some of it upon red-hot charcoal, (<^)by placing a drop
of any strong acid upon it.
3. Boil some starch in water so as to form a paste. Then add
some iodide of potassium solution, and allow the whole to cool.
The reaction, which occurs on adding free nitrous acid to the
mixture, forms, when taken in connection with the nitrate of silver
test, a sure indication of the presence of nitrous acid.
4. Put into a test-tube a little strong nitric acid, then plug the
mouth of the tube with horsehair or fine wood shavings, and heat
the acid to boiling.
5. Dissolve some common salt (sodium chloride), also some silver
nitrate, mix the solutions, filter and evaporate the filtrate, test the
result and see if it is a nitrate. The white precipitate was silver
chloride. Why did it appear as a precipitate, and the other salt not
appear .?
6. Heat a little strong nitric acid in a test-tube, then cautiously
drop into it a little turpentme, CjoHig.
7. Repeat the experiment but use benzine, CqH^
AVOGADROS LAW.
103
8. Place a little carbolic acid in a test-tube, then pour in a
few drops of fuming nitric acid. The carbolic acid must be largely
diluted, and the tube turned so that the fluid will do no harm if it
spurts out.
Explanation. — The staining of organic substances yellow by
nitric acid is probably due to the formation of picric acid, a deep
yellow dye. When nitric acid acted on carbolic acid the result was
picric acid ; the replacement is interesting :
C6HeO + 3HN03 = CeH3(N02), + 3H204-0.
9. Assuming that air is by volume 21% oxygen and 79% nitrogen,
calculate the weight of a litre of air, and find a multiplier which
would enable one to change the specific gravity of a gas from air
= I to H = I.
10. Will a metallic nitrate, such as that of copper or zinc, when
distilled with sulphuric acid yield nitric acid.-*
11. Pour some dilute nitric acid on some powdered marble,
CaCOg. After the powder disappears, evaporate, test the solid to
find if it is a nitrate, then try if it will yield nitric acid when distilled
with sulphuric acid.
For nitric acid and nitrates, consult Bloxam, p. 167 ; Muir & Slater, 159;
Richter, 203; R. & S., 399; R., 277; Fownes, 152.
I
\
1
CHAPTER XXII.
1.— Avogadro's Law.
Gaseous bodies have their volumes altered by changes
in temperature and pressure. Thus the volume occupied
by a quantity of gas varies inversely as the pressure to
which it is subjected (Boyle's Law) ; and the volume of
a gas varies directly as the absolute temperature (Law of
104
AVOGADRO'S LAW.
H
Charles). Since all gases, whether light or heavy, ele-
mentary or compound, vary according to these two laws,
it follows that the change must be dependent, not on the
chemical, but on the physical properties of gaseous sub-
stances. Decrease of volume by compression and increase
of volume caused by increase of temperature are both due
to the overcoming of molecular forces of attraction and
repulsion. It follows, then, that if exactly similar forces
cause equal changes in equal volumes of different gases,
that equal forces are being overcome ; but equal forces
exist only among equal numbers of molecules ; hence
the law enunciated by Avogadro, an Italian chemist,
which is as follows : Equal volumes of gases, ivhether of
the same or of different kinds, contain equal numbers of
molecules ^ under like conditions of temperature and
pressure.
Of course this law has not been demonstrated true,
but it seems to follow directly not only from the physical
reasons already given, but also from the atomic theory.
We know from experiment that two volumes of hydrogen
unite with one volume of oxygen to form two volumes
of steam. If, then, elementary matter is composed of
atoms or groups of atoms, there must be two of these
individuals or groups of hydrogen in steam for every
one that there is of oxygen ; but the volume of the hydro-
gen is double that of the oxygen ; hence, in equal
volumes there must be equal numbers of atoms, or of
groups of atoms. (Some reasons for believing that these
are groups, that is, molecules, and not individual atoms
will be found in Chap, xxxvill).
NITROGEN AND HYDROGEN.
105
CHAPTER XXIII.
1.— Nitrogen and Hydrogen.
There is one well known compound of nitrogen
and hydrogen — ammonia, which is of very common
occurrence in compounds, and wliich is of consider-
able economic importance on account of its use in tiie
arts. This substance has been prepared synthetically in
small quantities by passing an electric discharge through
a mixture of the two gases of which it is composed.
This method of preparation is only of theoretical value,
and, as it is not suited for class work, no further reference
will be made to it. The ordinary preparation of ammonia
and its chief properties will be illustrated by the follow-
ing experiments.
2.
Experiments.
1. Take about 20 grams of dry ammonic chloride and
an equal quantity of dry quick-
lime ; powder them finely in a
mortar. Smell the mixture, and
then transfer it to a flask with
tightly-fitting cork and long tube
bent upwards. Heat gently.
Hold a large test-tube over the
delivery tube, and fill it with gas
by downward displacement of
air, as in Fig. 32.
2NH4Cl-hCaO-CaCl2 + 2NH3+H20.
2. Pass a lighted taper up into the test-tube full of gas.
Fio. 32.
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106
NITROGEN AND HYDROGEN.
3. Pass some of the gas into reddened litmus. Upon
the result of this, devise a means of knowing when a
bottle is full of this gas.
4. i'our 4 or 5 drops of hydrochloric acid into a large
beaker and, by shaking, spread the acid over the bottom
and sides of the vessel, then hold it mouth downwards
over the delivery tube.
A more convenient method of obtaining ammonia gas
is by liL-ating the spirits of hartshorn, the liquor am-
monice of the drug shops. Hartshorn is only a solution
of ammonia gas in water.
5. Fit a flask with a rubber stopper and tube as in Fig.
33. Place in the flask a little hartshorn and
heat it to boiling. When ammonia gas begins
to escape from the tube, invert it, and place the
open end in some water coloured pink with
litmus.
6. Fill a graduated tube, such as a eudio-
meter, with dried ammonia gas over mercury,
then lift the tube and place it mouth downward
in cold water.
7. Fill a large test-tube with ammonia gas
over mercury. Take a piece of porous char-
coal (that from pine wood, or pine bark is best),
hold it in a flame until it begins to burn over most of its
surface, then pass it into the tube without raising the
latter out of the mercury ; let the whole stand for a
couple of hours. Is the result due to any action of the
mercury ?
8. Fit up a piece of apparatus as is shown in Fig. 34.
Fig. 33,
NITROGEN AND HYDUOOEN.
107
The tube, A, is a large test-tube drawn out to a nozzle
about ^ of an inch in /^ ^
diameter, the tube, B,
must slide somewhat
freely through the per- *"'■<*• ^•*-
foration in thcstop[)cr and just project beyond the opening
in A ; C is connected with an oxygen supply, and B is
joined to the flask in which ammonia is generated. When
the current of ammonia is passing through B try to
light it, then turn on the oxygen and after it is escaping
from the nozzle try again to ignite the ammonia.
Vary the experiment by passing ammonia through C
and oxygen through B, but pull B backwards until its
end is just inside the nozzle of A.
Vary the experiment again by pulling B backwards
until it projects through the cork only as far as C does ;
as the mixed gases escape from the nozzle try to ignite
them.
Explanation. — When ammonia burns with oxygen it
is decomposed, thus : —
2NH3+30 = 2N + 3H,0.
When slowly oxidized, it forms ammonium nitrite and
water, thus : —
2NH3+30 = NH,NO,+H,0.
This is the result of two separate chemical actions
though, which are : —
NH3+30 = HN02+H,0, and
HNO^-fNHg^NH^NOo.
.t
108
NITROGEN AND HYDROQEN
In the former, ammonia unites with oxygen to form
nitrous acid and water, then in the latter this nitrous
acid combines with part of the free ammonia to form
ammonium nitrite.
9. Form about a foot of platinum wire into a spiral by
winding it around a lead pencil or bit of glass rod. Heat
some hartshorn in a flask until ammonia is coming off,
then heat the wire red hot and suspend it in the escaping
gas. If kept in the neck of the flask just where the
ammonia and air are mixed the spiral should glow for
several minutes, while in the air it almost instantly cools.
The explanation is that the red hot platinum promotes
the union of the ammonia and the oxygen of the air, the
resultant products being those of the combustion of am-
monia, and the heat developed in the chemical action
is sufficient to keep the platinum at the glowing point.
10. Put about 2 c.c. of dilute hydrochloric acid into a
test-tube, pass into this ammonia gas until it is neutral
to litmus, then evaporate, but without heating more than
is necessary.
11. Vary ex. 10 by dropping hartshorn very slowly
into hydrochloric acid until it is neutral to litmus, then
evaporating. Compare the results in this and in the
preceding experiment with that obtained in experi-
ment 4.
3.— Questions and Exercises.
I. How can you tell when a jar is full of this gas, if you collect
by displacement of air ?
3. Powder some soft coal coarsely in a mortar. Then place in a
■i
AMMONIUM AND AMMONIUM UYDROXIDE.
109
hard glass tube and heat. Smell the gas that comes off. Is the
liquid that forms acid or alkaline ?
3. What became of the ammonia in ex. 7 ? How could it be
shown that mercury did not cause the change .-*
4. Pass a current of oxygen from a gas holder into an open flask
containing hartshorn, warm the latter and apply a light to the
mixed gases that are escaping from its mouth.
4.— Ammonium and Ammonimn Hydroxide.
The radical NH4 in ammonium salts is very similar to
the metal potassium in its chemical characteristics, and
it is generally believed that this unisolated base has the
properties of an alkaline metal. The following experi-
ment illustrates one of the reasons for believing in the
existence of such a substance. Mercury forms peculiar
combinations with many metals ; these combinations
are not definite in quantity, and seem to be a sort of
connecting link between mixtures and chemical com-
pounds. They are exactly analogous to solutions and
are called amalgams. Ammonium amalgam can be
easily prepared ; and as amalgams are formed only with
metals, it would seem that ammonium belongs to that
class of substances.
Experiment.
Make a strong solution of ammonium chloride in a
beaker and put some sodium amalgam into it. The mass
should rapidly swell up to many times its former bulk,
have a soapy feeling when rubbed, and give off bubbles
of gas that smell strongly of ammonia. After a few
hours the original mercury will be found in the beaker
and the fluid will have a salty taste.
i
( I
110
DETECTION OF AMMONIA.
r,.f
i
The reaction may be represented as follows, but the
sign ' X ' must be understood as indicating that the
elements whose symbols it joins have formed an amal-
gam, and this not necessarily in the proportion of one
to one : —
Hg + Na = Hg X Na,
Hg X Na + NH4CI = Hg X (NH4) + NaCl,
HgxNH4 = NH3 + Hg + H.
Ammonium Hydrate.
When potassium is thrown on water, the water mole-
cules are broken up into hydroxyl m.olecules and
hydrogen, the former uniting with the potassium to
form potassic hydrate, and the latter escaping as free
gas. It is generally considered that ammonium hydrate
is the basic substance that unites with acids to form salts.
It has the composition NH^OH, that is, ammonium in
union with hydroxyl. When ammonia gas, NH3 dis-
solves in water this hydrate is produced, thus : —
NHg-f-Hp-NH.HO.
It will be necessary to keep very clearly in mind that
when compounds of this substance are formed they are
compounds of ammomum, not of the gas ammonia.
5.— Detection of Ammonia
It is of importance that ammonia should be capable
of easy and accurate detection, as its presence in water,
to any appreciable extent, is usually an indication of the
1 ';l
NOTES ON AMMONIA.
Ill
unfitness of that water for drinking. The following
form the tests generally applied : —
1. Pungent smell, if present in quantity.
2. Dissolve any ammonic salt in water, pour into this
solution, in a test-tube, some solution of potassic, sodic
or calcic hydrate, then heat, and the odour of ammonia
should be perceived.
3. When present in minute quantities, as it frequently
is in drinking water, ammonia is best detected by what
is known as Nessler's test : " To a solution of potassic
iodide add solution of mercuric chloride until the preci-
pitate formed just ceases to be re-dissolved, then add an
equal volume of strong solution of caustic potash, and
allow the whole to stand until clear. A few drops of
this solution will give a yellowish red precipitate,
with even the slightest trace of ammonia."
!■!
6.— Notes on Ammonia.
Ammonia: symbol^ ^H.^; viol, vol., 2 ; vapour density^ 8' 5.
Ammonium is a strongly alkaline^ monad base.
Ammonia is soluble to upwards of 700 times its bulk
in water at I5°C; it becomes a liquid at — 40°C,and may
even be frozen at — 75°C.
It occurs in the urine and in some other products of
animals ; also in air as the result of the decay or decom-
position of nitrogenous animal matter, hence it exists
in rain water. Its compounds, ammonium chloride and
ammonium carbonate, are found sparingly in nature.
'fr
1 1
112
QUESTIONS AND EXERCISES.
li '" •■ii
It is prepared for commercial purposes from the waste
products of gas works. The illuminating gas, cfter com-
ing from the retort, is washed in cold water to free it
from soluble impurities ; the ammonia is thus dissolved
and afterwards separated as the chloride.
Dilute liquor ammoniae is used in medical practice
and in manufactures to neutralize acids. Its com-
pounds are used in medicine as stimulants in cases
of fainting or of syncope from overdoses of chloroform,
ether or laughing gas. It is also used in dyeing. The
latent heat resulting from its volatility, renders it valu-
able for cooling purposes in some manufactures.
7.— Questions and Exercises.
1. Complete the following equations : —
NH4HO + HNO3 =
2NH4HO+H2S04 =
NH4N0a + KH0 =
2NH40H + H2C03 =
NH^Cl + NaOH =
2. Potassic chloride and potassic nitrate are written KCl and
KNO3 respectively, while ammonia has the formula NH3. but the
chloride and nitrate are written NH4CI and NH4NO3. How is
the extra atom of hydrogen accounted for ?
3. Ammonium nitrate has its formula written NH4NO3, what
objection is there to writing it N2H4O3 ?
4. What weight of ammonia gas at 60° F. can be obtained from
214 grams of ammonic chloride ?
5. What weight of quick-lime is required to decompose 107
grams of ammonic chloride, and what will be the weight of the
calcic chloride and water produced ? What weight of ammonia
gas will ';e evolved ?
ADDITIONAL EXERCISES.
113
if
6. Ammonium chloride is heated with caustic soda, the resultant
gas led into water, this solution neutralized with nitric acid, then
evaporated to dryness, and heated in a test-tube until decomposed.
Write the equations for the various steps of this process.
IS
ma
8.— Additional Exercises.
The combining power of substances in the nascent state is well
shown by the formation of ammonia, as illustrated by the following
experiments : —
1. Make a mixture of hydrogen and nitrogen. Try by testing
with litmus and by smelling the gas if any trace of ammonia can
be detected.
2. Mix in a mortar 30 centigrams of fine iron filings with sixty
centigrams of solid caustic potash. Then transfer the mixture to a
test-tube fitted with cork and delivery tube, and heat until gas
escapes. Collect some of this gas and test for hydrogen.
5Fe-l-(oKHO = 5FeO-f 5K2O + 5H2.
Iron and caustic potash yield ferrous oxide, potassic oxide and
hydrogen.
3. Repeat the experiment, substituting 20 centigrams of nitre for
the 60 of potash. In this case test for nitrogen.
5Fe-f2KN03-5FeO-f-K„0 + 2N.
Iron and potassic nitrate give ferrous oxide, potassic oxide and
nitrogen.
4. Now mix 30 centigrams of iron filings, 30 of caustic potash and
6 of nitre, place in a test-tube and heat as before. Smell the gas
that is evolved.
loFe-f ioKHO-f2KN03=ioFeO-l-6K20-l-2NHjj-f-2H2.
Iron, potassic hydrate and potassic nitrate yield ferrous oxide,
potassic oxide, ammonia and water — the ammonia instead of free
hydrogen and free nitrogen ; though free hydrogen and nitrogen,
when formed separately and then mixed, will not form ammonia.
5. Pass ammonia for some time into a long test-tube of ice cold
water. Note any changes in temperature and volume of the water.
8
■ %
, 5
lU
COMPOSITION OP AMMONIA.
I:.
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ill
■? ■».
9.— Determination of the Composition of
Ammonia.
I. Take a eudiometer, fill it with mercury and invert
over a small trou^^h or saucer, also containing mercury.
Heat some ammonium hydrate and pass 20 c. c. of the
gas into the eudiometer. Then pass a series of electric
sparks from an induction coil through the gas, taking
care to insert a Leyden jar in the circuit, so as to in-
crease the heating effect. When the gas no longer
expands, pass 30 cc. of pure oxygen into the eudio-
meter and explode. Depress the eudiometer so as to
bring the mercury to the same level inside and outside ;
then note the volume of the gas remaining in it.
10.— Questions and Exercises.
1. What must be the composition of the gases remaining after
the explosion ? What voknne of each ? What thei: are the consti-
tuent^^s of ammonia by volume ?
2. How many atoms of each element must there be in the mole-
cule of ammonia ? What therefore must be its formula .''
3. Assuming the truth of Avogadro's Law, what space does the
molecule of ammonia occupy as compared with that of an atom of
hydrogen ?
4. In performing Ex. 5, sec. 2, the bottom of the flask broke into
small pieces, but these all flew inwards. How may such a result
be accounted for.?
5. In preparing ammonia from sal ammoniac and lime, try if it
will answer to make the substances into a paste with water.
6. Into some dilute sulphuric acid in a test-tube put some pieces
of zinc, after hydrogen begins to come ofif freely, add drop by drop
some nitric acid until the gas ceases to pass out of the liquid.
Then add potassic hydrate and heat. Ammonia should escape.
How may its formation be accounted for }
, |i !^.
:fi
PERCENTAGE COMPOSITION AND PORMULiE.
115
7. Place a small piece of ammonium chloride upon a strip of
platinum foil and heat for some time. Devise an experiment to
determine whether the change is a physical one or a chemical one.
8. How many volumes of ammonia will be produced by uniling
one volume of nitrogen and three volumes of hydrogen ? Upon
what experimental evidence is the answer based?
9. It is said that any ammoniacal salt heated with any alkaline
hydrate or oxide will yield ammonia. Try with as many of the
following as you can get, and write the equation for the reaction, in
each case in which the gas is obtained : —
{a) Ammonium nitrate, NH^NO,, with potassic hydrate KHO.
(d) Ammonium carbonate (NH4)2C03, with barium oxide BaO.
(c) Ammonium sulphate (NH4)2S04, with calcium hydrate Ca(OH)a.
{d) Ammonium chloride NH^Cl, with sodium hydrate NaOH.
10. Pour some hartshorn into a beaker and place a thermometer
in it to get its temperature. Tie a piece of cloth or filtering paper
loosely about the lower end of the thermometer, then let a few
drops of this same hartshorn run down the stem and drip through
the cloth or paper. Compare the fall in temperature with that when
alchohol, water, ether, or chloroform is used.
CHAPTER XXIV.
PERCENTAGE COMPOSITION AND FORMULAE.
sees
rop
lid.
ipe.
1.— To Calculate the Formula of a Compound
when its Percentage Composition is Known.
As soon as we have determined what elements are
present in an unknown compound, and what their
' relative weights are, we can use this knowledge to
construct a formula which will show the molecular
II
?s
116
PERCENTAGE COMPOSITION AND FORMULiC.
composition of the substance. For instance, a gaseous
compound may yield oxygen and nitrogen on analysis,
but this does not settle what oxide of nitrogen it is.
Some examples worked out will make clear the
methods of solution and the principles on which they
are based.
Examples.
I. A certain compound when analyzed yielded sodium
27%, nitrogen i6"5^, oxygen 56"5%. What is its formula?
The percentages give the relative weights of the con-
stituents in a unit weight of the compound ; therefore
the relative weights of the different kinds of atoms in a
molecule of the compound. In a molecule of this com-
pound, if the weight of the sodium atoms be represented
by 27, that of the nitrogen will be 16 j4, and that of the
oxygen S6}4- We must next find the relative num-
bers of these atoms in the molecule. To do this the
percentage weight must be divided by the atomic weight
of the element.
Sodium, 27 ^23 = 1.171.
Nitrogen, 16 ^I4=ii7±.
Oxygen, 56.5-^16=3.5 ±.
We now know that for every v\y atoms of sodium there
are ri7 atoms of nitrogen and 3*5 atoms of oxygen.
These figures are not absolutely correct, but in practical
work, errors of experiment render m ithematical accuracy
an impossibility. The next step will be to determine
what integral numbers will- express these ratios. To do
this, we divide the smallest of the quotients obtained in
the last operation into each of the others.
PERCENTAOB COMPOSITION AND FORMULJG.
117
Thus, for sodium,
" nitrogen,
oxygen,
((
11 7
117
117
11 7
3-5
1-1 7
= 1.
= I.
= 3.
i'l
This tells us that the numbers of atoms of sodium,
nitrogen and oxygen are as i, i and 3. Hence the
formula for the compound is NaNOg, or some multiple
of this, n(NaNO.,)
2. A compound gave on analysis 78*3% silver, 4'3%
carbon, 174% oxygen. Find a formula for it.
Solution.
78.3 -r 108 =. 725 ; 4.3-ri2 = .358; 17.4-^16=1.09.
.725^.358 = 2±
.358^.358=1
1.09 -f- .358 = 3±.
There are, therefore, two atoms of silver to one of carbon
to three of oxygen. Hence a formula for the substance
is AggCOg. This is not necessarily the exact formula,
because Ag^C^Og or AggnCnOgn would answer just as
well, so far as the data of the question applies. There is
an element left out in stating the problem which is
necessary for an exact solution. This is the vapour
density, which will be dealt with in another chapter.
For present practice, the lowest number of atoms per-
missible may be taken as the proper formula, which is
then said to be empirical.
Definition.
An empirical formula expresses the proportions by
weight in which the constituents of a substance unite to
form it.
f
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lif
118
EXERCISES.
li::
1 '•■
ii .. '
II I^V
Ml
'111
!
The proper empirical formula for each compound sub-
stance is fixed by accurate chemical analysis.
To solve all similar problems observe the following
rule:
7. Divide the percentage amount of each constituent
element by its own atomic weight.
2. Divide each of the quotients thus obtained by the
lowest of them and the numbers obtained will express the
proportional number of atoms of each element in the com-
pound.
Exercise.
The following are percentage compositions of various sub-
stances ; determine a formula for each.
1. Carbon, 42*86 ; oxygen, 57*14.
2. Hydrogen, 273 ; chlorine, 97"27.
3. Hydrogen, '83 ; sodium, I9'i7 ; sulphur, 26"66 ; oxygen, 53*33
4. Sodium, 39*31 ; chlorine, 6069.
5. Nitrogen, 82*35 ; hydrogen, 17 "65.
6. Phosphorus, 91*17 ; hydrogen, 8*83.
7. Carbon, 26*67 > hydrogen, 2*22 ; oxygen, 71*11.
8. Carbon, 75 ; hydrogen, 25.
9. Carbon, 1 2 ; calcium, 40 ; oxygen, 48.
10. '9 gram of a substance containing carbon, hydrogen and
oxygen is found on analysis to yield "24 gram of carbon and "02 of
hydrogen. Calculate its simplest formula.
11. *9 gram of a substance containing carbon, oxygen and hydro-
gen is found on analysis to yield '06 of hydrogen and 48 of oxygen.
Calculate its simplest formula.
12. A portion of a substance is found on analysis to yield *36
gram of carbon, *o55 gran) of hydrogen, and '44 of oxygen. Calcu-
ate its formula,
PERCENTAGE COMPOSITION PROM THE FORMULA. 119
2.— To Calculate Percentage Composition from
the Formula.
Sometimes we are given the formula of a substance
and are asked to calculate the percentage composition.
This is easily done. Proceed as follows : —
Find the molecular weight of the compound by taking
the sum of the atomic weights of its constituents, then
divide this separately into the weights of the atoms of
the different elements in the molecule.
Example.
Calculate the percentage composition of sulphate of
copper, CUSO4.
Copper 63 . 5
Sulphur 32
Oxygen (16x4) 64
1595
If, by weight, in 1 59^ parts of sulphate of copper there
are 63^ parts of copper, how many parts by weight of
the metal will there be in 100 of sulphate.
159^ of sulphate yield 633^ of copper.
.-. I will yield ,^33^
i59>^
and .*. 100 will yield . . — 5_2x^^=^Q.8i per cent.
159.5 I
The percentage of sulphur and oxygen in this com-
pound may be found in the same way.
m
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I I
120
GRAPHIC AND UATIONAL PORMULiR.
Exercise.
What is the percentage composition of each of the following
substances : —
1. Arsenious oxide, AsjOg.
2. Chloride of gold, All CI.,.
3. Arseniuretted hydrogen, AsH.,.
4. Potassium ferrocyanide, K4 P^eCgNo.
5. Magnesium sulphate, MgSO^.
6. Copper nitrate, Cu(N03)a.
k
3.— Graphic ani Rational FormulsB.
Rational Formulse.— A rational formula besides ex-
pressing the proportions by weight in which the elements
are united, expresses also the way in which the elements
are supposed to be grouped within the molecule of a
compound. For example,
HOI
HOI ^^^ ^^ ^^^^ rational formula for sulphuric acid, — that
is, two hydroxyl molecules joined with sulphur dioxide.
Graphic Formulae. — Graphic formulae express more
fully than rational formulae the manner in which we sup-
pose atoms to be associated in forming compounds. For
example, the graphic formula for water is H — O — H.
For nitric acid the empirical formula is HNO3; the
rational formula is NO2 (OH), and the graphic formula,
O
II
N— O— H.
II
O
In graphic formulae, the lines indicate the manner in
which the atomicities of each element are disposed of;
GRAPHIC AND RATIONAL FORMUI.ifJ.
121
nitrogen being joined to the other elements by five links,
oxygen by two, and hydrogen by one.
In the formula for water, oxygen is shown to be a diad,
because it has two combining powers joining it to the
two atoms of hydrogen. In a similar way, nitrogen is a
pentad, each atom of oxygen a diad and hydrogen a
monad.
The graphic formula for sulphuric acid may be written
O-H
o=s=o
O— H
The graphic formulae for ammonic chloride, ammonic
nitrate, copper sulphate, copper nitrate, are respectively :
H H
\ /
N CI,
/ \
H H
OO
11/ I
S Cu,
II \l
OO
H H O
\ / II
N O— N,
/ \
H H
O
O
O
= N— Cu— N = 0.
I I
O o
Exercises.
Construct graphic formulae for hydrogen peroxide, ozone, potassic
nitrate, sodium hydrate, calcic hydrate, Ca(0H)2, acid sulphate
of potassium, zinc sulphate, potassic chlorate.
T^"
122
EXPERIMENTS WITH CHARCOAL.
CHAPTER XXV.
1.— Oarbon.
This element enters into the composition of every
organic substance, and is a constituent of a very large
number of mineral substances. It exists in three forms
that, physically, are entirely different from one another,
but chemically are identical.
The word alJotropism is used to express the fact
that some elements exist in very unlike states physically,
or with very different properties, but all the while pre-
serve their fundamental chemical identity.
The allotropic forms of carbon are (i) charcoal, an
impure form derived by roasting organic matters out of
contact with air ; (2) graphite, plumbago or black
lead, a mineral found chiefly in metamorphic rocks ; and
(3) diamond. The last is the purest form.
ii 'i
2.— Experiments with Charcoal.
I. Partly fill a narrow test-tube with white paper, dry
sawdust, or pieces of dry wood ; stop the mouth loosely
with a piece of chalk, or pour sand to a depth of half an
inch over the substance in the tube. Hold the tube in a
nearly horizontal position with its lower end in a lamp-
flame. As the heating goes on, notice whether any odour
is evolved. Place separate pieces of blue and red litmus
paper within the mouth of the tube. When nearly all
action has ceased, turn out and examine what remains in
the test-tube,
■
QUESTIONS AND EXEKCISES.
123
2. Place a little of the residue on platinum foil, or on
a sheet of mica, and heat over a lamp flame for some
time. Observe what is left.
3. Clean as well as you can the test-tube used in the
preccdin<j experiment, and then repeat it, using a piece
of woollen cloth, silk cloth, or dry lean meat.
4. Heat on a sheet of mica the residue obtained in
experiment 3.
5. Heat some suf:^ar upon a sheet of mica.
6. Place a wet filter paper inside a funnel, and then
cover the inside of the filter paper with a thick coating
of animal charcoal or bone black. Now filter through
the paper a wine-glassful of ale or porter, or a - Intion of
dark brown sugar.
Definition.
Roasting a substance out of contact with air is known
as destructive distillation of it.
3. -Questions and Exercises.
1. How is wood charcoal prepared? Animal charcoal? Is
wood an element ? Give reasons for your answer.
2. Mention one difference between the liquid produced in the
destructive distillation of wood and that obtained in the destructive
distillation of lean meat.
3. Charcoal is said to be an impure form of carbon. What
impurities have you found ? Is the statement regarding impurity
true of animal, as well as of wood charcoal ?
4. Put into an alcohol lamp a mixture of fluids, two-thirds alcohol
and one-third turpentine. Hold a cold plate in its flame until a
thick coating of soot is formed on it, scrape this coating off and
heat it on mica. This black powder is lamp-black, Is it carbon?
^
w^
124
ADDITIONAL EXERCISES.
5. Set fire to one corner of a sheet of paper and let it slowly burn
away, then try if the charred sheet will again burn. What is finally
left?
6. Let a piece of meat stand until it begins to smell putrid, then
cover it with powdered charcoal. How is the foul smell affected.?
How long does this continue ?
7. Ascertain what effect animal charcoal has upon solutions of
(a) litmus, (d) indigo, (c) potassium permanganate, (ci) writing ink,
when these fluids are filtered through it.
8. Fill a test-tube with air over mercury, pass into the tube a
piece of freshly-burned charcoal and let the whole stand for an*
hour.
9. Water filters are frequently made of layers of sand and char-
coal laid alternately. From your experiments, have you reason to
believe that this would prove a serviceable arrangement? Would
it be likely to \i^ permanently effective ?
4.— Additional Exercises.
1. Place a piece of charcoal in a test-tube and then pour upon
it a little strong sulphuric acid. Observe whether the charcoal
changes in any way. Try whether an alkali will produce any
change in the charcoal.
2. Wet the inside of a large test-tube with liquor ammonias.
Now drop into the tube some wood charcoal previously heated in a
covered crucible. Cork the tube, and after a few minutes remove
the cork and ascertain by smelling whether there is any ammonia
left in the test-tube.
3. Secure a piece of electric-light carbon ; try if it will burn (i) in
an ordinary flame, (2) in a blowpipe flame.
4. Heat, in a test-tube, a piece of soft coal covered with a layer
of sand. Does it undergo destructive distillation ? Try with
hard coal.
5. Cut two little blocks of wood abort as large as peas. Heat
one of these over a flame on a piece of mica j notice what is left
after all action ceases.
CARBOI^ AND OXYGKN.
125
Heat the other similarly on a piece of mica, but take the precau-
tion to cover it to a depth of naif an inch with some incombustible
substance like sand, clay, chalk-dust or lime.
6. Repeat the last experiment, but substitute lean meat for wood.
7. From a sample of coarse brown sugar prepare some that will
be pure and white.
8. Devise an experiment to ascertain how long charcoal will
retain its decolorizing and deodorizing power when in use as a
deodorant.
9. If, after animal black has lost its power as a decolorizer, it
were again roasted, would it regain that power ? Test your conjec-
ture by an experiment.
Definition.^ — ^A sub.stance which has the power of
destroying offensive smells is called a deodorant.
CHAPTER XXVI.
1.— Carbon and Oxygen.
There are two well-known oxides of carbon, viz., car-
bon monoxide (carbonic oxide), and carbon dioxide
(carbonic anhydride, carbonic acid gas, or choke damp).
The latter is of general distribution in the atmosphere ;
and, dissolved in water, has played a very important part
in the formation of the rocky crust of the earth. It is
also necessary for the support of vegetable life.
2.— Oarbon Dioxide.
The test for the presence of this gas is lime water.
This is prepared by pouring clean water on lime, letting
v.-
I
P- H
A,.
At
3
3
J!
M
126
CARBON DIOXIDE.
it stand for several hours, with frequent stirring, then
allowing it to settle and decanting the clear liquid. This
should have a distinct alkaline taste, and should be free
from turbidity. When carbon dioxide is passed through
lime water a white precipitate is thrown down.
Experiments.
1. Twist one end of a piece of fine wire round a bit of
charcoal, hold the latter in a lamp flame until it is glow-
ing brightly, then lower it into a bottle and insert the
cork beside the wire ; when the charcoal ceases to burn
withdraw it and shake up some lime water with the gas
in the bottle. Contrive a means of driving the gas in the
bottle through lime water in a test-tube.
2. Take the apparatus used in preparing hydrogen
and in it place some powdered limestone or white
marble. Then cover the marble with water, and pour
down the thistle-tube some hydrochloric acid. Collect,
over the pneumatic trough, two or three bell jars and
two or three large test-tubes full of the gas. The equa-
tion for the re-action is
CaC03+2HCl = CaCl2+H20-l-C02.
3. Test the gas by lowering into it a lighted candle.
Substitute a piece of burning phosphorus for the candle.
4. Invert one of the test-tubes full of the gas over
water and let it stand for some hours.
5. Test, with litmus, the water that rises in the tube, in
the previous experiment. This may be done by slipping
the hand under the mouth of the tube, then raising the
whole and inverting it. If the litmus is not immediately
changed, let it .stand for a time.
QUESTIONS AND EXERCISES.
127
6. Invert a tube full of carbon dioxide over mercury,
then with a curved pipette pass a solution of potassic
hydrate up into the tube. Vary the experiment by try-
ing to collect the gas over a solution of caustic potash.
in
3.— Questions and Exercises.
1. Devise experiments to show that carbon dioxide is formed in
(a) a burning lamp, (b) burning coal gas.
2. Show, by means of the lime water test, that this gas is given
off during respiration.
3. It is said that carbon dioxide is one of the common impurities
of ill-ventilated school rooms. Devise a means of testing the accu-
racy of this statement.
4. What is the gas that comes off from ginger ale, soda water or
ale when the bottle is unstopped ? To find out, prepare a cork with
a delivery tube that will fit the neck of the bottle, tiien unstop the
bottle and quickly insert the prepared cork ; allow the escaping gas
to pass through lime water. Why does the gas escape with a rush
when the cork is drawn ?
5. Twenty-five grams of sodium carbonate is heated with just
enough dilute hydrochloric acid to complete the chemical action ;
the gas that comes ofT is passed into solution of ammonia which is
afterwards evaporated to dryness.
(a) Write the equation for the reaction.
(d) If the hydrochloric acid were 25% pure, what weight of it
would be required .'* (Ordinary hydrochloric acid may
be assumed to be a 20% solution by weight of the
gaseous acid in water.)
(c) What weight of each of the salts would be formed .-'
6. Mention substances that when burned do not yield carbonic
acid gas as one of the products of combustion.
7. Lime water is a solution of calcic hydrate, Ca(H0)2. This in
contact with carbon dioxide forms calcic carbonate ; thus :
Ca(OH)a 4- C02=CaC0;j + HgO.
128
ADDITIONAL fiXRRCISfiS.
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Calcic carbonate may take the form of limestone, marble or chalk.
It is called by the last name when in fine white powder.
Is chalk dust soluble in water ? Is sodium carbonate soluble in
water? From your observations, tell why a solution of calcic
hydrate rather than a solution of sodic hydrate is used as the test
for carbon dioxide.
4.— Additional Exercises.
1. Hang, by a piece of bent wire, a lighted taper with a small
flame in a gas cylinder, then pour carbon dioxide into this cylinder,
out of another vessel. Try to pour carbon dioxide upwards into a
jar, then test with lime water to see if the gas really passed into the
upper vessel.
2. Try if this gas can be prepared from any of the following salts
when treated with any one of the acids given : Ammonium carbonate
(NH4)oCO;3, sodium bicarbonate, NaHCOy, potassium carbonate,
K0CO3, ammonium bicarbonate (NH^)HC03, lead carbonate,
PbCOa (white lead, PbCOaPbOH will do), barium carbonate,
BaCO;} ; for acid, either dilute nitric, dilute hydrochloric or dilute
sulphuric may be taken. In each case write the equation.
3. Take a large bottle, to hold a gallon if possible, fill it with
water ; and, in the worst ventilated room in the school, empty this
water into another vessel. Have ready a large test-tube full of
lime water, pour tiiis into the bottle, cork the latter and shake the
lime water thoroughly with the air, then pour it back into the test-
tube. Repeat the experiment, but collect the air outside the
building and use an equal quantity of fresh lime water. Compare
the quantity of precipitate in the two tubes.
4. Counterpoise a large beaker on a scale, then pass a stream of
carbon dioxide into this beaker. How is the scale affected .'' How
can you tell when the beaker is full of the gas ?
5. Try whether you can syphon carbon dioxide from one jar to
another ?
CALCIC CARBONATE IN CARBON DIOXIDE.
129
5.
-Solution of Calcic Carbonate in Solution of
Carbon Dioxide.
A fact of immense importance in the history of the
world is that a solution of carbon dioxide in water is capa-
ble of dissolving limestone which is insoluble in pure
water. The following experiments illustrate this.
I. Pass a current of carbon dioxide through lime water
until the precipitate formed at first nearly all dissolves.
Filter and divide the filtrate into several parts. Boil one
of them, it should turn milky in appearance. Drop into
another a little ammonia ; into another, some potassic
hydrate solution ; into another, some more lime water.
A white precipitate should be formed in each case.
Explanation.
When the carbon dioxide is first passed into the lime
water a precipitate of calcic carbonate, CaCOa, is thrown
down ; when the water becomes saturated with the
dioxide this precipitate is dissoh'cd ; perhaps because of
the formation of an acid carbjnate, CaM.,(C0.5)2 ; but
any treatment that takes the carbon dioxide out of the
solution causes the precipitate to be again formed, be-
cause water alone will not dissolve the neutral carbonate.
Boiling expels the carbonic acid gas, and the alkaline
hydrates unite with it to form other carbonates.
The quantity of carbon dioxide that water will dissolve
varies as the pressure to which the water is subjected,
hence, springs that come from far below the surface some-
times contain much of this gas dissolved, and if the solu-
tion has trickled over limestone the latter will be also
dissolved and carried to the surface, where it will be
9
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n» ' .,1
11,
130
UEDUCINO POWER OF CARBON.
deposited, because, the pressure being removed from the
water, the dioxide escapes.
From what has been shown in the preceding experi-
ments can you explain the cause of the deposition of
" lime " (furring) of steam boilers, tea kettles and other
vessels in which much water has been boiled ?
6. -ileducing Power of Carbon.
Carbon is very largely used in metallurgy as a reducing
agent, the monoxide, and dioxide being formed in the
process. The following experiments illustrate the prin-
ciple on which it acts.
1. Mix in a hard glass tube, closed at one end, some
copper oxide and powdered charcoal ; on top of this place
a thick layer of charcoal, then heat strongly. Test the
gas coming off for carbon dioxide. Examine what re-
mains in the tube.
2. Repeat the preceding experiment but use arsenic
trioxide for copper oxide. What appears on the cold
part of the tube ?
3. Try if red lead and iron rust are changed, when sub-
jected to treatment similar to that applied to the copper
oxide.
7.— Carbonic Acid.
It has been found in previous experiments that when
carbon dioxide, prepared either from burning carbon, or
from the decomposition of a carbonate, is dissolved in
water, it forms a distinctly acid solution. The solution
::;:I
CARBONIC ACID.
131
is dilute carbonic acid, formed thus :— Ho04-CO^ =
H.^COa. The acid cannot be obtained in a concentrated
form, because it readily breaks up again into water and
carbon dioxide. Though the acid is thus unstable, its
salts are of very general occurrence, and form a consider-
able part of the earth's mass. It forms carbonates with
all the common metals and alkalies, the most common
of these being calcium carbonate in its various forms
of chalk, limestone atid marble.
All carbonates are very readily decomposed by the
acHon of any strong acid upon them, with the evolution
of carbon dioxide gas. Thus :
MCO3 -f 2HCI = MCI2+ HXO,
= H20 + C02.
The carbonic acid at once breaks up into water and car-
bon dioxide. The decomposition of the acid may be
represented graphically as follows :
O
II
H— O— C— O— H
When the two hydrogen atoms are freed, the two oxygen
atoms to which they are attached are thrown into an
unstable condition and the result is a rearrangement ;
thus :
= C = 04-H— O— H.
The alkaline carbonates alone are soluble in water. Most
carbonates when strongly heated break up into carbon
dioxide and the oxide of the metal. The well-known
burning cf limestone to form lime serves as an example.
CaCOg^CaO-fCOg.
ONTARIO COLLEGE OF EUUUA! ION
ITff'
U'2
CARBONIC ACID.
Experiments.
1. Take a hydrogen generating apparatus and place in
it some powdered chalk, or oyster shells. Cover with
water. Insert the cork tightly and then pour down the
thistle-tube some hydrochloric acid. Collect the gas
that escapes and test it for carbon dioxide.
2. The following is a simple method, when you want
to prepare quicklime on a small scale ; but you need a
Bunscn burner with which to do it. Take a piece of
calc-spar and round it, wind some platinum wire. Hold
the calc-spar in the hottest part of the Bunsen flame for
a short time.
3. Moisten a piece of red litmus paper and press it
against a piece of limestone or marble. Now press a
similar piece of litmus against a piece of recently burnt
quicklime.
4. To a lump of quicklime add, little by little, water
until it becomes pasty. What became of the first water
that was added ? What did the change of temperature
indicate ?
\'h '
8.— Additional Exercises.
Carbon Dioxide Contains Carbon and Oxygen.
Experiments.
I. Lead a current of carbon dioxide through a hard glass tube
which contains some thin sHces of sodium. After the gas has been
passing long enough to exclude all air from the tube, heat the
sodium until it ignites. After a few minutes remove the tube and
scrape out the black flakes. Prove that these are carbon by burning
them in a current of air and leading the product of combustion
through lime water. For this, use the apparatus of Fig. 26, but
NOTES ON CARBON DIOXIDE.
133
insert a bottle of lime water between the aspirator and combustion
tube. The white solid left in the tube is sodium oxide, and when
dissolved in water it forms sodium hydrate. Test it for sodium.
2. A piece of burning magnesium, if put into carbon dioxide will
continue to burn. Try this, and after the combustion has ceased,
a little weak nitric acid will dissolve the magnes'c oxide and leave
the carbon particles.
3. Take thin pieces of marble, chalk or oyster shell, and weigh
them. Place them upon an old tin plate, or in a crucible, and heat
in a fire to redness, for an hour or two. Remove them and try
whether they will now yield carbon dioxide on the addition of
hydrochloric acid. The residue after the roasting is called
quicklime, or simply lime. Note differences in colour and
weight between lime and limestone.
9.— Notes on Carbon Dioxide.
Carbon dioxide: symbol, CO^; mol. iveight, ^^; vapour
density y 22 ; sp. gr.^ i'$2 (air=i); soluble to the extent
of about 3 vols, of gas in 2 of water ; exists in air to the
extent of about 4 vols, in 10,000. This supply is main-
tained by the breath of animals, the combustion of car-
bonaceous matters, the decomposition of carbonates, and
by fermentation.
This gas is the choke damp of miners; so named
because of the formation of large quantities of it by the
explosion of another gas (fire damp) in the mines, and its
deadly effect when breathed continuously. As it contains
no free oxygen it is incapable of supporting life, so as-
phyxiation results if it is inhaled for any length of time.
The carbon dioxide in the atmosphere is largely absorbed
and decomposed by plants and part of it used by them
as a food.
f
134
CARBON MONOXIDE.
li
10.— Carbon Monoxide.
Experiments.
1. Into a Florence flask put 8 or lo grams of oxalic
acid, and about 50 c. c. of sulphuric acid. Fit with a
tight cork and tube and attach, as in Fig. 35, to a wash
bottle containing a
strong solution of
caustic potash. From
the wash bottle, a
delivery tube should
pass tothe pneumatic
trough. Apply heat
cautiously to theflask,
regulating it so that
the gas may come
off in a slow, steady
stream. After the
air has been expelled from the apparatus, collect three
bottles of the gas, and allow these bottles to stand over
water for some time before using them. Meanwhile, sub-
stitute for the delivery tube one whose end has been
drawn to a fine point. Apply a lighted match to the jet.*
2. Raise one of the bottles of gas from the water,
and apply a lighted taper to its mouth.
3. Try to pour the gas from one bottle to another,
then test the result with a lighted taper.
* Unless the experimenter is careful he will get enough carbon dioxide mixed with
the monoxide to spoil his result. To guard against this he should use a large volume
of the potash solution in a tall vessel, then if necessary he should cause the gas to bubble
slmvly through this a second time, or better still, let it stand for several hours over caustic
potash solution.
Carbon monoxide is an exceedingly poisonous gas, therefore should not be inhaled.
Fig. 35.
QUESTIONS AND EXERCISES.
135
4. Purify thoroughly the gas in a third bottle, by
shaking it up well with caustic potash or caustic soda
solution, then test the gas with clear lime water.
11— Questions and Exercises.
1. If carbon monoxide burns in air, try to find out whether vapour
of water is formed during the process of combustion? Whether
carbon dioxide is formed ?
2. On what data can you conclude that carbon dioxide contains
more oxygen than carbon monoxide ?
3. Hydrogen and carbon monoxide are both combustible. Can
you conclude from this resemblance that carbon monoxide and
oxygen will form an explosive mixture ? Test the accuracy of your
conclusion.
4. Oxalic acid has the formula of C2H2O4, and sulphuric acid
takes from this the elements that, when united, form water. Write
the equation for the reaction, and explain the necessity for the wash
bottle filled with potassic hydrate solution in experiment i, sec. 10.
5. Try if a burning splinter, or burning phosphorus will continue
to burn if thrust into a jar of carbon monoxide.
6. Thirty grams of oxalic acid is treated with excess of sulphuric
acid, the resultant gas led through a solution of sodium hydrate in
which all the carbon dioxide is dissolved, the remaining gas is
burned, and the product of combustion passed into lime water where
it unites with the dissolved hydrate. Write equations for the
several reactions, and determine what weight of each compound is
formed.
7. How do you account for the blue blaze that spreads over coal,
when newly thrown on a hot fire ?
8. Trace the chemical changes that occur when air enters a
furnace by the lower damper, passes through a fire pot filled with
white hot coals, into the air space above the coals, thence out
through the smoke pipe at the top of the furnace.
9. In what particulars (3 at least) does hydrogen resemble carbon
monoxide? How would you distinguish them, if similar jars were
filled, one with each gas, and given to you ?
If
136
RELATION BISTWKKM VOLUMES OF OASES.
' <* >
i
12.— Additional Exercises.
1. If we can form carbon dioxide from carbon monoxide by
supplying it with oxygen, it ought to be possible to prepare carbon
monoxide from carbon dioxide by supplying it with carbon. Test
the accuracy of this induction by using apparatus similar to that in
Fig. 12, and passing dry carbon dioxide over red-hot charcoal.
2. Pass carbonic oxide and oxyg' ' ito a eudiometer over
mercury. Do they unite when an elec . spark is passed through
the mixture? How can you tell whether or not a union has taken
place ?
3. Try if carbonic oxide will burn in nitrous oxide gas.
4. By using the apparatus of Fig. 45 (appendix), or some similar
device, drive a current of carbon monoxide through a delivery tube,
ignite the escaping gas and allow it to burn under a stoppered bell
jar or inverted bottle, whose mouth dips under water. After the
combustion ceases, let the jar or bottle stand over a solution of
caustic potash for some hours. What gas remains ?
CHAPTER XXVII.
Ill
III
]S ■
1. — Relation Between Volumes of Constituents
and the Volume of the Compound Formed,
when all are Gases.
We have already seen that 2 volumes of hydrogen
unite with i of oxygen to form two volumes of steam.
The compound formed occupies only ^ of the volume
which we would expect it to occupy. The process of
union has brought about a shrinkage in volume which
we must try to understand.
RELATION nETWEEV VOLUMES OP CASES.
137
On the other hand, when nitrous oxide is decomposed,
we find that 2 vohimcs of the gas procliice 2 vokimes of
nitrogen and i volume of oxygen. In other words, the
process of decomposition has given rise to an increase of
volume. Both these results appear to contradict the well
known fact that 2 and i make 3, but in reality they are
easily understood if we use Avogadro's law as the basis of
explanation. According to this law, one molecule of any
substance in the gaseous state occupies the same space
as one molecule of any other substance under like condi-
tions. Now, the molecules of the mixed gases in the first
example, in the process of union, re-arrangc their atoms
and combine them so that the number of new molecules
of steam formed are reduced in the proportion of 3 to 2,
and consequently the total number of molecules in the
steam is only ^ of what they were in the mixture, hence,
as all molecules occupy the same space, the total volume
is diminished to ^ of what it was at first.
The very opposite takes place in the decomposition of
nr^rous oxide. Here the new molecules are increased in
number, in the proportion of 3 to 2, and hence the total
volume is increased to one half more than in the nitrous
oxide.
When we come to study hydrochloric acid we shall
find that i volume of hydrogen unites with i volume of
chlorine to form 2 volumes of the gas. Similarly, 2 vol-
umes of nitric oxide may be decomposed into i volume of
nitrogen and i of oxygen. In both these cases, there is
neither an increase nor a diminution in volume, because
in the re-arrangement of the atoms to form new mole-
cules, the number of molecules in the products are exactly
thQ same as at first.
138
VAPOR DENSITY.
t
3
J
Stating these results generally : there will be in the
final product of any chemical reaction an increase, a
diminution, or equality in volume to that which entered
into the reaction, according as the atoms re-arrange
themselves so as to produce more molecules, fewer mole-
cules, or an equal number of molecules to those which
first entered into the reo...tion. If the number of mole-
cules remain the same at the end of a reaction as at the
beginning, the volume will be the same : if the number
of molecules has increased, the volume must have in-
creased : and if the number has diminished, the volume
must have diminished. In all cases the volume is
entirely independent of the number of atoms in each
molecule, but depends, as already stated, upon the num-
ber of molecules.
1. Apply this principle and explain what change of
volume will take place when ammonia and nitrogen tri-
oxide are decomposed.
2. What change in volume, if any, will take place
when (a) sulphur is burned in oxygen, (d) carbon dioxide
is passed over red-hot charcoal, (c) marsli gas is burned
in air ? In the first case, devise an experiment to ascer-
tain whether your conclusion is correct or not.
2.— Vapor Density.
The molecule of a gaseous body occupies the space of
two atoms of hydrogen (this will be clear if it is remem-
bered that steam when decomposed yields its own volume
of hydrogen, and nitrous oxide its own volume of nitro-
gen). It follows, therefore, that the density of a gas is
VAPOR DENSITY.
139
expressed by a number equal to one-half its molecular
weight (H= i).
It also follows that 22 4 litres of any gas weigh the
number of grams expressed by tlie molecular weight of
that gas ; for i r2 litres of hydrogen weigh i gram, then
22'4 litres of hydrogen weigh 2 grams ; also 22*4 litres of
another gas, whose molecular weight is x will weigh
2 XXX y2=x.
This result is constantly used in cliemical calculations in
translating weights into volumes and vice versa.
Example.
How many litres of nitrous oxide may be obtained by
the decomposition of 30 grams of ammonium nitrate?
Solution,—
NH4N03=N,0 + 2H,0.
80 parts give 44 parts by weight,
•*• 30 grams give |^ x 30 — "'^- grams.
And since 22"4 litres weigh 44 grams, '\f- grams occupy
-_- 4 V -'—V ':.\J>
1_
4 4
8.4 litres.
3.
It may be well to remind the student here of two physi-
cal laws that affect the volume of a fixed weight of a gas,
and must, therefore, be taken into account in all calcula-
tions, except when the standard conditions of tempera-
ture and pressure exist. These laws are Boyle's and
Charles'. The former relates to pressure, and tells us
that, when temperature is constant, the volume of a quan-
tity of gas confined in a closed vessel vanes mversely as
140
QUESTIONS AND EXERCISES
i:!5
w
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the pressure to which it is subjected. The latter relates to
temperature, and according to it the volume of a quantity
of gas increases -^^^ part of its volume at o°C. for every
degree centigrade through which its temperature is raised.
If a quantity of gas occupied ' a^ vols, at pressure jr, and
if the pressure changed to '^,' the volume would then be
ax- ; also if the temperature at first was f °C., then if
this changed to/°C., the correction would be obtained
by usinc: the factors — -^—, and -^ — -, the former to
change to volume at o", the latter to change from volume
at o° to that at p°. The entire corrections for temperature
and pressure would then be ^ x - x
^ y 273+/
4.— Questions and Exercises.
Assume standard temperature and pressure, unless others are
given.
1. How many grams of hydrogen will occupy 224 litres at the
standard temperature and pressure ?
2. Steam is passed through a tube containing red-hot iron filings,
18 litres of hydrogen pass out at the other end. Wliat volume of
steam was decomposed, and how much are the iron filings increased
in weighty
3. How much sulphuric acid and zinc must be taken to form 1 12
litres of hvdrogen at 7° C. .'*
4. In 285 grams of caustic potash how many grams of potassium.'*
of hydrogen .-*
5. What weight of sodium must be taken to obtain 20 grams of
hydrogen from a litre of water? If temperature were raised to 70°
C. and pressure to 800 mm., what would the answer then be?
QUESTIONS AND EXERCISES.
141
6. A reservoir of hydrogen gas holds 89'6 litres. What weight
of water will be formed in burning the gas in air? What volume
of air will be required for the combustion, assuming that oxygen
forms 21% of the volume of air?
7. I want 220 grams of oxygen. If I obtain it from potassic
chlorate, how much of it must I use? If from water, how much ?
If from mercuric oxide, how much? What volume at 15° C. and
200 mm. pressure would this gas occupy ?
8. A gas bag is capable of containing 56 litres, how much potassic
chlorate must be taken to procure enough oxygen to fill it at 35° C.
and 750 mm.?
9. 25 litres of oxygen are exploded with 36 of hydrogen. What
volume of gas (if any) remains? What volume of steam is pro-
duced? And what is its weight?
10. How much oxygen can be obtained from 435 grams of man-
ganese dioxide by heating it to a red heat ? (The reaction is
3MnOo=:Mn304 4-02). What volume will it occupy at 30° C.
and 780 mm. ?
1 1. What volume will 80 grams of oxygen occupy at the standard
temperature and pressure ? At 62° and 790 mm. ?
12. How much potassium will be required to decompose no
grams of carbon dioxide ? What weight of each of the products
will be formed ?
13. If 10 litres of carbon dioxide be passed over red-hot charcoal,
what gas, and how many litres of it, will be formed at 30° C. ?
What weight of it ?
14. 20 litres of carbonic oxide are burned in oxygen gas. What
gas is produced, what volume at 40° C. and what weight of it ?
15. How much carbon can be obtained from 264 grams of carbon
dioxide? Would change of pressure vary the answer?
16. What volume of oxygen at 10° C. is required to burn 66 grams
of carljon ? How would the volume of the gas formed compare
with that of the oxygen used ?
17. In question 5, what volume of air at 740 mm. would be needed
for the combustion of the hydrogen at 0° C?
U2
OAKBON AND HYDROGEN.
i8. What volume do no grams of carbon dioxide occupy at 760
mm. pressure and 0° C?
19. What volume do 140 grams of carbonic oxide occupy at
standard temperature and pressure ?
20. What weight of carbon dioxide can be obtained from 250
grams of pure limestone by treating with hydric chloride ?
21. What weight of carbonate of lime and hydric chloride must
be decomposed to produce 35 2 grams of carbon dioxide .-*
22. What volume will 98 grams of carbonic oxide occupy at 72,
mm. pressure and 40° C?
23. If 270 grams of oxalic acid be decomposed by sulphuric acid,
find the volume of the gases produced at 750 m. pressure and 20° C.
CHAPTER XXVIII.
a..:.^
CARBON AND HYDROGEN.
1— Hydrocarbons.
A large number of compounds of carbon and hydrogen
are known under the general name of hydrocarbons. So
numerous are these compounds, and those which carbon
forms with oxygen, nitrogen, sulphur and phosphorus,
that their mere names would fill a small x^olume. The
study of the carbon compounds forms a distinct branch
of chemistry under the name of organic chemistry. For-
merly, this name included the study of those compounds
which, it was supposed, were formed only by the agency
of life ; but it was soon found that there was no essential
difference between chemical substances whethci of animal,
vegetable or mineral origin. The division of chemistry
MKTHANE.
I4;i
at
therefore, into organic and inorganic, is a pure matter of
convenience. The last-named division of the science
treats of the composition and properties of air, earth and
water, whereas organic chemistry may be raid to be
the chemistry of the carbon compounds.
Marsh gas is the first of a series of hydrocarbons known
as the marsh gas series. Each member of it differs from
the following one by CHo. There is a difference of
30° between the boiling points of successive members.
All are inflammable. There is also a regular increase or
decrease of other physical properties. Such series are
called homologous series. The general algebraic formula
for the series is CJi-^a+i.
2.— Methane.
Methane (Marsh Gas^ Light Carburetted Hydrogen,
*' Fire-damp^' ), CH^; molecular weighty 16; vapour
density, 8.
Experiments.
I. Take a hard glass test-tube or Florence flask and
fit it with a cork and fine delivery tube (a copper retort
used for preparation of oxygen is preferable as there is
no danger of breaking it). Place in the test-tube 4 grams
of acetate of sodium, NaC^H.^Og, 8 grams of sodium
hydroxide, and 4 grams of finely-powdered quicklime,
CaO. Heat. After collecting a beaker or two of the
gas, light the jet and observe the colour of the f.^vme.
Before lighting the gas, test it in the same way as
hydrogen to see that it is not mixed with air.
1'
■i
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bj
Ri
1
m
144
NOTES ON METHANE.
Tlie formula for this reaction is
NaC,H,p,-f-CaO + NaHO=:CaCO,4-Na,0 + CH,.
If a little water were added, so as to change the oxide of
calcium into the hydrate, a second decomposition would
have occurred, thus :
(1) NaQH,0,+Ca(H0)2= CaCO,4-NaHO+CH,.
(2) NaC,H302+NaHO=Na,C03+CH4.
hence :
(3) 2NaC,H302+Ca(HOX, = Na,C03+CaC03-|-2CH4.
2. Fill a small soda water bottle with a mixture of one
part of me gas and two parts of oxygen. Ignite the
mixture. Express the reaction by an equation.
3. Take a stoppered bottle and fill it with a mixture
of equal volumes of marsh gas and chlorine. Expose to
sunlight for a day, then test the contents with blue
litmus. Note any change in colour.
3.— Questions and Exercises.
1. Devise an experiment to ascertain whether marsh gas is acid
or basic in reaction.
2. Devise simple experiments to prove that the gas contains
carbon and hydrogen as constituent elements.
3. Prove that the gas is lighter than air. How would you distin-
guish the gas from air .''
4.— Notes.
Methane is generated in marshes by the decomposition
of vegetable matter containing carbon and hydrogen. It
is formed in coal mines also, and, on being mixed with
OLEPIANT GAS.
146
air and ignited, causes fearful explosions, hence in
miners' lan.cruage it is " fire damp." To prevent these
accidents Sir H. Davy invented his celebrated Safety
Lamp. ^
5— Olefiant Gas.
This is the old name of another gas that may be taken
as a type of a second series of the hydrocarbons. The
general formula of the homologues of this series is C,.H.,.
Ethylene (Ethene. Olefiant gas. Heavy carburettZ
hydrogen X C.,H,; molecular weight, 28; vapour density, i^.
Experiments.
I. Into a Florence flask pour 50 or 60 c.c. of strong
sulphuric acid and half that volume of alcohol. Insert a
tightly.fitting cork and delivery tube. Place the flask
on a retort stand and heat gently. After the air has all
been expelled, collect two or three jars of the gas.
Alcohol has the formula C,H,0. The sulphuric acid
acts here in the same way that it does on oxalic acid in
the preparation of carbonic oxide, viz , by extracting the
elements of water.
C,Hp + H,SO, = C,H, + H,SO, + H,0.
2. Ascertain whether the gas will burn or not. Has it
any taste ?
3. Devise simple experiments to prove that the gas
contains hydrogen and carbon.
4- Find out whether the gas is heavier or lighter than
air
10
I~
la;
II
146
ACKTVLENE.
:«
i
5. Remove a jar of the gas, let it drain well, then turn
it mouth upward and place over it another jar of the
same size filled with chlorine. After standini^ for some
time, note whether the colour of the chlorine changes.
Observe closely what forms at the bottom of the lower
jar ("Dutch liquid.")
6. Ascertain whether the gas will explode when mixed
with air or oxygen.
7. Is it soluble in water?
6.— Acetylene.
Acetylene, another hydrocarbon that has the formula
C2H2, may become of great economic importance for
illuminating purposes, because of a recently-discovered
method by which it can be produced at small expense
in large quantities. When calcium carbonate and
powdered coal are heated in an electric furnace, calcium
carbide, CaC.^, is formed. This, when immersed in water,
readily yields acetylene, and the gas burns with great
luminosity in the air, when ignited. The products of
the reaction by which the acetylene is formed are calcic
hydrate and acetylene, thus : —
CaC2 4-2HoO = C2H2-hCa(HO)2.
COAL GAS AND FLAME.
147
CHAPTER XXIX.
COAL GAS AND FLAME.
1.— Coal Gas.
Coal gas is formed by the distillation of coal in large
iron retorts. The process of manufacturing it may be
illustrated by strongly heating some powdered soft coal
in a common clay pipe. The mouth of the pipe should
be closed with kneaded clay. The average composition
of coal gas (for it is really a mixture of many gases) is
about as follows : —
Hydrogen ^^.
Marsh gas --.
Carbonic oxide -.
Olefiant gas
Butylene ^
Hydric sulphide q.-
Nitrogen ^
Carbon dioxide ^.g
T^^'^^ loo- vol^.
When the gas comes from the retort it contains a
much larger quantity of hydric sulphide carbon dioxide
and ammonia. It undergoes certain purifying processes
notably, a washing in a stream of cold water which
causes the precipitation of tarry ingredients and the
solution of a large portion of the ammonia. Any ingre-
dient of coal gas which either does not burn with a lumi-
nous flame, or does not help to support the combustion of
the other substances is objectionable, and should be got
148
LUMINOSITY OP PLAME.
J
It
! •''
I't ' ■
I. 1
^>
m '
Bi
rid of, if possible. Hydric sulphide will burn readily,
but one of the products of its combustion is a very un-
desirable substance to have mixed with the air of dwell-
ings. The water in which the gas is washed yields, on
proper treatment, our supply of two important substances,
ammonia and coal tar. From the latter, are manufac-
tured the beautiful aniline dyes so extensively used in
recent years.
2.— Luminosity of Flame.
Now that we have learned something of carbon, hydro-
gen, oxygen, and a (cw of the compounds which they
form, we are in a position to study somewhat more fully
than we have hitherto done, the subject of combustion,
with reference to the luminosity, or light giving proper-
ties of flames.
Experiments.
1. Sprinkle into the flame of an alcohol lamp or Bun-
sen burner, some fine iron filings.
2. Rub together over the top of any non-luminous
flame, two pieces of charcoal. Repeat, using two pieces
of chalk.
3. Hold a piece of platinum wire or a piece of lime in
the flame of hydrogen gas.
4. Observe closely the flame of an alcohol lamp, and
if possible take it into a dark room to test its illuminat-
ing powers. Hold a cold plate horizontally in the flame
for a minute ; then blow out the flame and put into the
lamp about one-half as much spirits of turpentine as there
is of alcohol, quickly light the lamp and observe the
I 5 ;
STRUCTURE OF FLAME.
149
charjf^e that comes over the character of the flame ; again
hold a cold plate in it. Repeat, but use a teaspoonful of
benzine instead of the turpentine. »
3.— Questions and Exercises.
1. From what source did the light emanate in all three experi-
ments ?
2. Mention one way of changing a non-luminous flame into a
luminous one.
3. Explain the source of the black mark formed on a white plate
by holding it horizontally across the flame of a candle, or of a coal
oil lamp. What is the black substance.''
4. How do you explain the facts observed in your experiments
with oxygen, viz., that sulphur and phosphorus give more light
when burned in oxygen than in air .''
5. In experiment 4, sec. 2, turpentine is a hydro-
carbon having the formula Cjo Hio, and benzine
has the formula C,,Hf,, can you now account for
the change in the flame and for the dark spot on
the plate ?
6. Close the holes at the base of a Eunsen
burner, turn on the gas and light it. What kind of
flame is there .'' Turn off" the gas and invert a com-
mon funnel over the burner, as in Fig. 36, still
leaving the holes at the base closed ; turn on the
gas and after a few seconds light it at the end of the
funnel stem. What sort of a flame is there now ?
Fio. 36.
Why
4. — Structure of Flame.
I. Spread out the wick of a candle or alcohol lamp,
light it, and then thrust into the middle of the flame the
phosphorus end of a friction match.
150
STRUCTURE OF FLAME.
m]
2. Float a small cork upon a little common alcohol or
methylated spirits placed in the bottom of a small saucer.
Place a few ^i^rains of gunpowder upon the cork, and
then ignite the alcohol.
3. Bring a piece of wire gauze
down horizontally upon the flame
of a candle, of a coal oil lamp, or
alcohol lamp.
4. Light a candle and observe its
flame carefully. Note how many
parts there are in it. Take a nar-
row bent glass tube, about four or
five inches long, and thrust one end
of it into the dark cone in the
middle of the flame, as in Fig. ^y.
Try to light the vapours which rise
through the tube. Repeat the ex-
periment but use a Bunsen burner
Fig. 37. instead of the candle Try, both
when the holes at the base are open and when they are
closed. Why can you not get this result with a common
coal oil lamp? Try a lamp that has a circular wick.
Explanation.— It is customary to speak of a flame as
being made up of three parts ; these are — ist, the central
cone, consisting of gas that is not i^i^nited ; 2nd, the
luminous mantle in which co' ci s going on; this is
the chief light giving } I i|y| iter mantle which
is usually but slightly li linou Jid m v/hich combustion
of the gaseous substance is c )mpleted. In the central
cone the gas is still unburned. In the candk flame it is
formed by the fatty constituents of the wa or tallow
..'it!
'\4
IGNITION.
ir,i
being drawn up through tlie wick from the h'ttle reservoir
of melted matter surrounding it, and these are freed at the
top of the wick in the form of vapour by the heat of the
surrounding flame. As no oxygen is in contact with this
gas it will not burn. A current of air is, however, being
drawn into the vapour at tiie base of the flame around the
wick, and the oxygen of this air causes the outer layer of
the central cone to undergo constant combustion, with
the result that the carbon particles in this mantle become
incandescent. In the outer layer or mantle of the flame,
air has become freely mixed with the vapour, and this
vapour has been largely deprived of scjlid carbon particles
while it was in the middle (luminous) mantle, so that it
now burns with a nearly non-luminous flame. The
middle and outside mantles correspond respectively to a
Bunsen's flame with the holes at the base of the burner
closed and then open. The central cone is called that
of non-combustion, the next one, that of incomplete
combustion, and the outer one that of completed
combustion.
5.— Ignition.
The temperature at which a substance begins to burn
is called its temperature of ignition, or its kindling point.
Do all substances ignite at the same temperature .<*
Experiments.
I. Pour a little carbon bisulphide into a large test-tube.
Close with the thumb, and shake well, so that the vapour
will fill the tube. Then warm a glass rod and place it
in the vapour. K there be no result, heat the rod morQ
ptrongly and again place it in the vapour,
152
QUESTIONS AND EXERCISES.
2. Try to light coal gas with a rod similarly warmed ;
then with an iron rod nearly red hot ; and lastly with an
iron rod at a white heat.
■ i.
)
II *"
.1
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fi;
6.— Questions and Exercises.
1. Why is it harder to light a coal fire than a wood one?
2. If you cool a burning substance, will it cease to burn? In-
vestigate tliis point by making a small cone-shaped helix of cc/pper
wire and covering a candle flame with it. Heat the helix to red-
ness and again place it over the flame.
3. Investigate the principle of the Davy lamp, used to prevent
explosions in coal mines. The " fire damp" burns on the inside of
the wire gauze which surrounds the flame. Why does not the gas
outside take fire? Hold a piece of fine wire gauze two or three
inches above a Bunsen's burner, turn on the gas, and after a few
seconds Ignite it on the upper side of the gauze.
4. Why does blowing on a flame "//^/ // onf'^ ?
5. 5 grams of methane are burned in air. What would be
the weight of the products of combustion ? How could these
products be collected for weighing? What weight of air would be
required in order to complete the combustion ? Change all weights
into volumes, giving the answers in litres.
6. Try by experiment what changes a blowpipe makes in me
flame of a candle, (i) as regards its size, and (2) as regards its heat?
7. In using a blowpipe, which is preferable to use for the blast,
air exhaled from the lungs, or air that has been simply drawn into
the mouth ?
8. If a splinter of wood be placed across the flame of an alcohol
lamp where does the wood begin to burn first ? Test the accuracy
of your conjecture by actual experiment.
9. Whit shaped mark will a candle flame make upon a piece of
white letter-paper, when pressed for an instant horizontally upca the
flame ? Press the paper down almost to the wick, and remove
quickly.
f!
ADDITIONAL EXEIICISES.
153
JO. In the case of a candle flame, whence comes the gas that
forms the flame ?
1 1. Why sliould the outer mantle (of complete combustion) be
more prominent in a candle flame than in that of a liuiisen burner?
12. Which is hotter, the luminous or the non-luminous flame of
a Bunsen burner? Test by noting the time required to heat a
piece of wire red hot, when held in about the same position in both
flames.
13. What effect has a lamp chimney on the luminosity and tem-
perature of a coal-oil lamp flame ?
14. " Flame is incandescent gas." " Only gases burn with a
flame." Examine these statements in the light of the experiments
you have just made, or of observations which you have made on
flames in coal or wood stoves.
15. If flame is incandescent gas, whence comes the gas that pro-
duces the flame when phosphorus burns in oxygen, sulphur in air
or sodium on water ?
16. A piece of paper dipped in turpentine, CiqH^p,, when ignited
in the air burns with a sooty smoke coming off from it. A piece of
paper dipped in alcohol, CgHgO, when ignited, burns with an
almost non-luminous flame and without smoke. Why the differ-
ence ?
7.— Additional Exercises.
1. Examine the flame of a Bunsen burner, first, when the air
holes are closed, then, when open. Note consequent changes in
the temperature and luminosity of the flame. Pass in nitrogen or
carbon dioxide through the air holes instead of air, and compare
changes thus produced with these that took place in admitting air.
2. It in sec. 6, question :;, the gas burned had been olefiant gas
instead of methane, what would then have been the answers to the
questions ?
3. Ten volumes of hydrogen, 10 volumes of carbonic oxide,
10 volumes 01' marsh gas, and 10 volumes of olefiint gas are each
mixed with 25 volumes of oxygen and the mixtures burned. If the
m
t-^Sf
154
BLOWPIPE FLAME.
si ■
lUi I
II
Ijf'ir
products of combustion remain as gases, state what would be their
volume at ioo°C. and 750 mm. pressure.
4. A mixture of hydrogen and carbonic oxide, obtained by blow-
ing jets of steam into white hot coals, is used as an illuminant, and
is known technically as water gas (to distmguish it from coal gas
obtained by the distillation of coal). Explain the chemical reac-
tions, with equations, that go on in the preparation of this gas ;
also the chemical changes, and the value of each constituent for
illuminating purposes.
5. Why is the burner of a gas jet made with an opening in the
form of a slit instead of a round hole .'' and why is an argand
burner made to allow air to pass up the middle of it ? -,
6. Use a two-necked WoulflPs jar as a hydrogen
generator, arrange it as shown in Fig. 3^^, putting a
plug of cotton wool in the wide tube ; after all the air
is driven out set fire to the hydrogen escaping
from both tubes. Remove the stopper from the large
tube, pour two or three drops of benzine (CgH^) on
the cotton, replace the stopper and again ignite the
gas. How do you account for the change .'*
7. " .Soft coal " or bituminous coal burns with a
bright luminous flame, " hara coal " or anthracite glows, and
burns away almost entirely without flame. How do you account
for the difference .''
8. A lamp flame turned too high will smoke ; of what does this
smoke consist? Whence does it come } Why is it formed only when
the flame is turned too high i* How do you account for the clouds
of black smoke that come out of factory chimneys? If you are told
that this smoke escapes most freely just after fresh coal has been
put on the tire, would the statement confirm your explanation ?
Explain the " burning of smoke " in factories.
Fig, 38.
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■if
8.— Blowpipe Flame.
Three zones are observed when flame has a jet of air
jplown into it from the nozzle of a blowpipe. The inner
I!
CHLORINE.
155
mantle or zone of incomplete combustion, R, Fig. 39, is
technically known as the reducing flame, because here
the supply of oxygen is limited, hence the carbon has
been oxidized only to carbonic oxide, and there is, there-
fore, a great tendency to take oxygen away from any
Fig. 39.
substance that will part with it. The outer mantle, O,
is the oxidizi7ig flame, because the supply of oxygen
is plentiful, and the heating of a substance to a high
degree in contact with oxygen of course promot^'es
chemical union between the two, if that is possible.
References— Tilden, 64; R., 399; R. and S., 187.
CHAPTER XXX.
CHLORINE.
1.— The Halogens.
There are four elements — chlorine, bromine, iodine
and fluorine— that are closely related to one another, and
are knowi) in chemistry as halogens (salt producers).
is:
m
166
EXPERIMENTS WITH CHLORINE.
Chlorine is the most important of these. They all form
acids that do not contain oxygen ; these are sometimes
called haloid acids, and the salts which they form, haloid
salts ; they are thus distinguished from salts and acids
which contain oxygen.
'An:
2.— Experiments with Chlorine.
1. Into a test-tube put one part of manganese dioxide,
two parts of salt, and three of sulphuric acid. Fit the
test-tube with a cork and delivery tube. Heat gently.
Cautiously smell the gas that comes off. Note its colour.
2. To prepare the gas on a larger scale, take a 4 oz.
Florence flask and place in
it about 20 grams of man-
ganese dioxide and 100 c.c.
of strong hydrochloric acid.
Use fittings similar to those
in Fig. 40. Apply a very
gentle heat. The delivery
tube should pass almost to
the bottom of the jar. Fill
several jars, taking care
that little or none of the
gas escapes into the room.
Puce a wet glass cover over each jar. Afterwards pass
the gas into a flask perfectly ///// of water ; in about
ten minutes place this flask aside for future use. Smell
the water. The decompositions, which result in free
Fia. 40.
EXPERIMENTS WITH CIILOHINE.
157
chlorine beiwg formed in the first of these experiments,
may be re^jresentcd by the following equations : —
MnO,-f2NaCl + 3H,SO,= MnSO, + 2NaHSO,+ 2Cl.
This occurs in two steps thus :
NaCl+H,SO,-^NaHSO,+ HCl
2HCl+MnO,+ H,SO,=MnSO,+2Cl+2H,0.
If these reactions go on together, there are being pro-
duced at the same time nascent oxygen, hydrogen and
chlorine ; of these the oxygen and hydrogen unite and
the chlorine remains free. The sum of these two reac-
tions is indicated in the first equation.
The reaction that occurs in the second experiment
may be represented as follows : —
Mn02+4HC1 = MnCL-f- 2H,0+2C1.
3. Take the flask full of chlorine water, prepared in
the last experiment, and fit it with
a cork and tube. The outer end of
the tube must be drawn to a fine
point. Insert the cork so that there
is not a bubble of air left in the
flask. Invert the flask as in Fig. 41,
and expose to direct sunlight for a
day. Then place the flask on the
table, remove the cork, and quickly
bring a glowing splinter to the
mouth of the flask. Test the water
in the flask with blue litmus solu-
tion. Taste it.
4. Lower very slowly a lighted taper into a jar of
chlorine. At the same time suspend a piece of blue
Fio. 41.
m
158
EXPKKIMKNTS WITH CHLORINE.
:
1 1^
m
litmus paper at the mouth of the jar. Smell the gas
that is formed during the combustion.
5. Take a few pieces of the metal antimony and powder
them ; then place on a sheet of paper and shake the
powder into a jar of chlorine. Try if arsenic acts
similarly.
6. Fill a small jar with hydrogen, then bring its
mouth below the mouth of another jar of chlorine. Keep
the jars mouth to mouth, and invert them several times,
so as to mix the gases thorouL,hly. Then separate the
jars, carefully corking one, and applying a lighted match
to the other. Wrap a towel around the jar which you
have corked, so as to exclude the light, and carry
it to where the sun is shining, either in a room or out-
side. Place it on the floor or on the ground and quickly
unroll the towel so as to send the jar a short distance
from you. Chlorine and hydrogen should never be
mixed excepting in dim or diffused ligJit.
7. Shake a tube full of the gas up with cold water to
test its solubility.
8. Pass a current of chlorine, or pour some chlorine
water, into solutions of logwood and of indigo.
9. Prepare chlorine as before, and cause it to bubble
slowly through some strong sulphuric acid in order to
dry it. Collect a jar of the dry gas and place in it pieces
of dry calico of various colours. Close the jar tightly.
Allow the calico to remain in the jar for about fifteen
minutes. Then open the jar ; quickly remove the calico,
wet it in water, and return to the jar for fifteen minutes
more.
BLEACHING BY CULOllINE.
159
gas
3.— Bleaching by Chlorine.
Experiment 3, sec. 2, demonstrated the formation of
oxygen by the action of chlorine on water ; at the same
time an acid was formed, and, as water and chlorine were
the only substances present it is reasonable to assume
that the chlorine decomposed the water molecules, joined
with the hydrogen to form hydrochloric acid, and set the
oxygen free. The latter gas, in the nascent state, is
as we have already learned, a powerful oxidizing agent,
and consequently unites with the colouring matters
present and destroys them. Chlorine is, on this account,
said to bleach by oxidation.
It is possible that in a few cases chlorine may combine
directly with the hydrogen of the coloring matter, thus
breaking up the compound. These cases are rare, for
chlorine nearly always requires the presence of water for
bleaching.
4.— Questions and Exercises.
1. Half fill a test-tube with hydrogeri ov r water, then fill it up
with chlorine, and let it stand over water fo; a few hours in diffused
light. Test, with litmus paper, and with nitrate of .silver solution,
the water that passes into the tube.
2. Treat the refuse from preparing oxygen from manganese
dioxide and chlorate of potash with sulphuric acid, and observe
what gas is evolved. How do you account for this.?
3. Take a piece of printed paper, write some words on it with
lead pencil, and some with ink, then moisten the paper and drop
It into a jar of chlorine. After half an hour examine it.
4. If chlorine bleaches by oxidation, and you wish to remove ink-
stains from a handkerchief, why not plunge it into a jar of oxygen
rather than one of chlorine 1
M
'
100
ADDltlONAL EXKRCISliS.
H !i
in.
(It.r »
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1 i!
5. Lower a piece of glowing charcoal gradually into a jar of
chlorine ; from the negative result of this experiment, explain the
formation of the black smoke which escapes from the candle when
burning in chlorine.
6. If the waste ])ipe of a kitchen sink were foul smelling, devise
a method of deodorizing it.
,
5. —Additional Exercises.
1. Try if other chlorides may be substituted for common salt and
hydric chloride in the preparation of chlorine, as in experiments
I and 2, sec. 2.
2. Drop chlorine water into sodic hydrate solution until the
mixture smells of chlorine after stirring, then evaporate to dryness.
What remains.? Why did a precipitate not appear .''
3. Lower a piece of freshly cut phosphorus on a deflagrating
spoon into ajar of chlorine.
4. The oxidizing power of chlorine may be shown in the following
way : —
There are three oxides of lead, the protoxide, PbO, a buff or
yellow powder, the red oxide, l'b304, a scarlet powder and the
peroxide, PbOg, a dark brown powder. They are all insoluble in
water. If a little of the protoxide and of the red oxide be shaken up
separately with water in test-tubes, and chlorine be then passed
through the mixtures until the water is saturated, and the whole
allowed to stand for some hours, the yellow and red powders will
both be changed to brown, thus showing the change to the peroxide.
This change maybe hastened by using solution of potassic or sodic
hydrate instead of water as the liquid with which the powder is
mixed. Why ?
5. Wet a piece of blotting paper with oil of turpentine, CioHje,
and then place it in another jar of the gas. Use fresh and perfectly
fluid turpentine.
6. Make a saturated solution of chlorine in water in a bottle, then
immerse the bottle in a freezing mixture of ice and salt. Yellow
• 1
MII-OKISK AM) TIIK AF-KALINK IIYDKATKS.
161
scale-like crystals should separate. These are supposed to be
chlorine hydrate, CI-H5H2O. When warmed they decompose
into chlorine and water.
6. -- Notes.
CJdorine : symbol CI ; mol. vol. 2 ; sp.gr. 2\fy, (air~i).
This element exists commonly as a i^as, but may be
obtained in the li([uid form by enclosinc^ crystals of
chlorine hydrate in a strong tube and seah'ng off the
part in which the yellow liquid chlorine condenses.
Chlorine is not valuable alone as a bleachini^ agent.
In sanitary operations it is largely used as a disinfectant
and deodorant. It here acts in the same way that it
does in bleaching, viz., by indirect oxidation, or, possibly,
at times by the direct union with the hydrogen in the
noxious compound.
A disinfectant is a substance which arrests the spread
of specific disease, by destroying the special agent that
enters the body from without and causes the disease.
7.— Chlorine and the Alkaline Hydrates.
When chlorine is passed into a cold solution of potassic
hydrate, a chemical action according to the following
equation occurs : —
2K0H+CU-KC1 + KC10+H,0.
Potassic chloride, potassic hypochlorite and water are
produced.
If the solution of the hydrate were hot a different
combination would occur, thus ; — •
6KHOf3CL-5KCl-f KClOy+sH.O.
11
m
102
OlILOKIKK AND HYDROGEN.
Ik.
[
,.,,1
1 =
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ini.
liiii
The reason for this change from the former result is that
an alkaline hypochlorite in solution of an alkaline
hydrate is easily decomposed by heat into the chloride
and chlorate. The reaction expressed above really occurs
in two steps, thus : —
6KH0+3C1.,-3KC1+3KC10+3H,0
and 3KC10 = 2KC1 + KC103.
This shows the effect of altered temperature in changing
the products of a chemical experiment.
Chlorine, though very common in combination, does
not occur free in nature. This is readily accounted for
by its energetic chemical action with many elements,
such as phosphorus, hydrogen, sodium, antimony and
others.
The student should be able now to devise tests for
chlorine and to give examples of its uses.
CHAPTER XXXI.
^ii;
fi^
1.— Chlorine and Hydrogen.
Only one compound of chlorine and hydrogen is
known, that is the haloid compound HCl, called
hydric chloride or more commonly hydrochloric acid.
A popular name for this substance is muriatic acid, from
murta, sea salty because the acid was prepared from the
salts got by evaporating sea water.
■\\
IIYDHOCIILOUIC ACID.
163
2.— Hydrochloric Acid.
Experiments.
It has been found that chlorine readily unites with
hydrogen either when heated or exposed to light ; also
that chlorine decomposes water, forming a combination
with the hydrogen and setting the oxygen free. In each
case the combination was an acid.
1. Place some ammonic chloride in a medium-sized
test-tube, and add a few drops of sulphuric acid. Bring
a lighted match to the mouth of the test-tube ; also a
piece of blue litmus paper ; and lastly, a glass rod dipped
in ammonium hydroxide. Smell very cautiously.
2. Repeat this experiment, using a large test-tube or
flask fitted with a cork or delivery tube, and substituting
sodic chloride, NaCl, for the ammonic chloride. Use
twice the weight of sulphuric acid that you do of salt,
and apply heat very carefully. Collect some of the
gas by passing the delivery tube to the bottom of an
"empty" jar. Cover its mouth with a glass, plate.
Having filled the jar, remove the plate cover and turn
the jar mouth downward over some water coloured
blue with litmus. Then shake slightly. Devise a means
of finding out when the jar is full of the gas.
3. Fill a second jar with this gas, as before. Place
two or three globules of sodium the size of a pea, in a
deflagrating spoon, the handle of which passes through
a cork that exactly fits the jar. Heat the sodium to
ignition, and lower the spoon into the second jar of gas.
After all action has ceased, withdraw the cork and quickly
bring a lighted taper to the mouth of the jar. Dissolve
the solid on the spoon and taste the solution.
164
Aqua KKoiA.
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4. Pass the ^as into a solution of sodium hydrate
until it is neutral to litniiis, thcMi cvai)oratc. What com-
pound have you ? Write the equation.
5. Place a few pieces of zinc in a test-tube, and then
pour upon them about 2 c.c. of hydrochloric acid. After
all effervescence has ceased, remove the surplus zinc and
evaporate the solution to dryness.
Commercial Acid.
A solution of this <^as in water is what is usually sold
by dru<]^gists under the name of hydrochloric or muriatic
acid. How can the gas be obtained from such a solu-
tion? The commercial acid is prepared as a by-product
in the manufacture of common "soda" by Leblanc's
process. The solution of the acid is of varying strength,
but about 33% by weight is the strongest solution that
is permanent. Commercial acid contains about 20% of
the gas.
3.— Aqua Regia.
A mixture of three volumes of hydrochloric acid and
one volume of nitric acid is called aqua regia.
Experiments.
1. Place a little piece of gold-leaf in chlorine water
and let it stand for some time. Try what effect dry
chlorine gas has on the gold-leaf.*
2. Place a few fragments of gold-leaf in a test-tube
and pour upon them about i c.c. of hydrochloric acid.
* That which is sold as gold-leaf is often only a base alloy, so that these experi-
ments may fail through impurity of the metal used. The gold-leaf should be tried first m
strong nitric acid ; if it dissolves it is no good ; if, however, it is unaffected in the acid, it
will be pure enough for the work here indicated.
COMPOSITION OF HYDKOCIILOUIC ACID.
1G5
Warm slightly. After a minute or two add a few drops
of nitric acid.
3. Repeat the prccedini^ experiment, using small
scraps of platinum instead of gold.
Explanation. -The solvent action of aqua rcs^ia is
due to chlorine which is freed by the action c»f the two
acids on each other, thus :
HN0,+3HC1-2H,0+N0C1+2C1.
(NOCl is chloronitrous gas, or nitrosyl chloride). The
chlorine readily attacks the gold or platinum to form the
chloride, even when not in the nascent state.
4.— Composition of Hydrochloric Acid.
Experiments.
I. Take a bent tube like that in Fig. 42. Partly fill the
tube, as indicated, with hydrochloric
acid, and insert in the ends the terminal
wires of a battery. These terminals
should be carbon. Bring a lighted
match to that end of the tube connected
with the zinc of the battery. Moisten
a piece of coloured calico and place it
over the other end of the tube. Colour
the acid with litmus solution.
^
Fig. 42.
2. Pass about 25 c.c. of hydrochloric acid gas into a
eudiometer over mercury, then introduce sodium amal-
gam until the eas ceases to contract in volume ; test the
gas that remains.
h
166
COMPOSITION OF HYDROCHLORIC ACID.
I?
5.— Questions and Exercises.
1. Will hydrochloric acid burn in air? Will it support the com-
bustion of a candle ?
2. An analysis of a piece of gold coin is required, how would you
dissolve the metal ?.
3. In the experiments relating to the composition of hydrochloric
acid, write the equations for all the reactions.
4. Suggest any reason why the terminal wires should be tipped
with carbon, in the electrolysis of y ydrochloric acid.
5. How is the " solvent " power of the acid increased ?
6. Explain the effect of the acid gas on quicklime. To do this
fill a tube with the gas over mercury, and then pass a piece of
quicklime up under the mouth of the tube. Is this gas absorbed
by charcoal ? Find out by experiment.
7. How could you distinguish between this gas and chlorine ?
The following simple method of determining the com-
position of hydrochloric acid gas is described in Rey-
nold's Chemistry, Part II., page 69 : — -
Open the stopcock, Fig. 43, and pass a current of
hydrochloric acid gas through the U tube for some time,
then quickly close the stopcock and pour mercury
enough into the U tube to close the bend and half fill
II the open arm. Open the stopcock slightly and
^ allow gas to escape until the mercury stands at
nearly the same height in both arms. Mark the
height of the mercury in the closed arm. Next
drop into the open arm some sodium amalgam
and fill to the top with mercury ; close the open
end with tiie thumb, and pass the gas backward
Fio. 43. and forward a number of times through the mer-
cury by tilting the tube. Finally hold the tube erect,
ADDITIONAL EXERCISES.
167
raise the thumb and allow air to enter the open arm.
Pour in o«' remove mercury until it is at the same height
in both arms. The gas in the closed arm should now
occupy half the volume that it did at first. Tilt the
tube so that the mercury will press on the gas in the
closed arm ; cautiously open the stopcock and hold the
nozzle to a flame.
6.— Additional Exercises.
1. Try it any other chloride besides that of sodium can be used
for the preparation of hydrochloric acid.
2. Would nitric acid answer instead of sulphuric .'' Base your
answer on experiment.
3. Take 3 test-tubes ; half fill the first with a solution of silver
nitrate, Ag NO3 ; the second with a solution of mercuric nitrate,
Hg(N03)2 ; and the third with a solution of acetate of lead,
Pb(C2H302)2. Into each tube pour a few drops of hydrochloric
acid. Try to write the equations.
4. Place a little cupric oxide, CuO, in a test-tube and pour some
hydrochloric acid upon it. When the oxide ceases to dissolve,
filter, and evaporate the solution to dr) ness.
5. Try to form other chlorides by warming- h\drochloric acid
with metals, oxides or hydrates which you can find in the laboratory,
and which you have not yet used.
6. Decompose some hydrochloric acid solution in a dark room
by electricity. After this has gone on for some time so as to allow
the liquid to become saturated with the escaping chlorine,
collect some of the mixed gases and expose a measured volume to
the action of solution of iodide of potash. When shrinkage has
ceased, test the remaining gas.
7. What weight of common salt and sulphuric acid must be taken
if it be required to liberate 146 grams of hydric chloride ?
8. Calculate the amount of hydro-sodic sulphate that will be
produced in generating 219 grnms of hydric chloride from salt and
sulphuric acid, at a moderate temperature.
I^B^
108
NOTES OS HYDROCHLORIC ACID.
l!'' 3
!ii
*i
9. Explain why we believe that hydrogen and chlorine are united
in the proportions by weight of i to 35 "5.
10. What volume will 73 grams of hydric chloride occupy at the
standard temperature and pressure?
11. 10 grams of zinc are treated with dilute sulphuric acid in
excess, the gas that comes off is burned in an atmosphere of chlorine,
what volume would the product of the combustion occupy at 30^
C and 750 mm. pressure ?
7.— Notes on Hydrochloric Acid.
Hydrochloric acid gas ox hydric cJiloride: formula^ HCl;
inol. 7>oL, 2 ; TdoL weigJit, 73; sp. gr.^ i^j (air=i) ; soluble
in water to the extent of 480 times its own volume.
Commercial hydrochloric acid is one of the by-products
in the preparation of soda, where sodium chloride is
treated with sulphuric acid, and the escaping gas is in-
tercepted and dissolved.
The white fumes observed when hydric chloride escapes
into the air are due to the affinity 01' the acid gas for
water, which causes a condensation of the aqueous vapour
in the atmosphere.
The test for chlorides, including hydric chloride, is the
white curdy precipitate they form with solution of silver
nitrate.
I. Dissolve some common salt in half a test-tube full
of pure water, and then add a few drops of r"*;rate of
silver solution. Shake. Now pour half of the solution
into k second test-tube ; add a little nitric acid to the
one test-tube and ammonium hydrate to the other. Boil
the one to which you added nitric acid.
3. Repeat this experiment, using any soluble chloridtj
in place of common salt.
SOME COMPOUNDS OF CHLORINE.
169
CHAPTER XXXII.
1. — Some Compounds of Chlorine.
The compounds of chlorine arc numerous and import-
ant, but only a few of them can be referred to here.
The chlorides of the alkalies can be prepared by either
treating the metal with hydric chloride, by passing
chlorine into the hydrate, or by bringing the metal into
contact with chlorine. Most other metals form chlorides
with hydrochloric acid.
The followmg paragraphs refer to the preparation and
properties of a few of the important chlorine compounds.
11
2.— Bleaching Powder
Bleaching powder, or chloride of lime, is an important
article of commerce which is extensively used in bleach-
ing the coarser kinds of cotton and linen goods. Its
manufacture is illustrated in the following experiments.
Experiments.
I. Cover the inside of a bell jar with slaked lime, and
then pass chlorine into it for some time. The chemical
changes which take place may be thus represented :
2C].2-}-2CafKO)2 = 2H,0-f-CaCl, + CarC10)o.
SlakcU liihe.
Cak-ic,
chloride.
Calcic
hypochlorite.
It is this Hiixture of calcic chloride and calcic hypo-
chlorite which forms the most imjiortant ingredients of
what is popularly known as " bleaching powder."
170
POTASSIC CHLORATE.
■>►...,
I::
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kl
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2. Remove from the jar the product obtained in the
preceding experiment. Place it in a soup-plate and add
about ICX) c. c. of water, stirring the mixture for five or
ten minutes. Immerse in the solution thus prepared, a
piece of printed calico. After a few minutes remove the
calico and immerse it in a very dilute solution of sul-
phuric acid.
Bleaching powder, when acted on by sulphuric acid,
yields chlorine slowly.
Ca(C10X,-l-CaCl,-f-2H.^S04=--2CaS04-h2H20-f-2Cl2.
This is the result of three separate actioiiS.
(i) Ca(C10),-}-H,S04-CaSO, + 2HC10
(hypcohlorous acid.)
(2) Caa,-FH.,S04-CaSO,+2HCl.
(3) " HCl-hHC10-H20 + 2Cl.
3. -Potassic Chlorate.
Experiments.
I. Boil a strong solution of caustic potash in a test-
tube and pass into it a current of chlorine for half an
hour. Evaporate the solution to a small quantity and
then allow it to cool slowly. Both potassic chloride
KCI, and potassic chlorate KCIO3, will be formed in the
solution. The latter being the least soluble crystallizes
out JirsL The liquid that remains contains the potassic
chloride in solution. Pour off this liquid. To purify the
crystals re-dissolve them in a little hot water and allow
them to reform. (Compare Chap. XX x, sec. 7.)
All the oxygen compounds of chlorine are unstable,
and most of them are explosive, breaking up into chlorine
and oxygen.
11
POTASSIC IHLORATE.
171
2. Put a crystal of chlorate of potash about the size of
a pea into a test-tube, then drop in a little sulphuric acid
and heat gently. While pouring in the acid, and while
heating, keep the tube carefully pointed away from your-
self and other persons near you.
The rather violent decomposition that goes on is
expressed by the following equation :
3KC103+2H,SO,-2C102H-KC104+2KHSC4+H20.
This is the sum of the following reactions :'
2KC103 + 2H,S04 = 2KHSO, + 2HC103 (chloric acid).
2HC103+KC103=2C102+H,0-!-KC104.
The sulphuric acid and chlorate give rise to chloric acid,
which immediately breaks up into chloric peroxide,
water and oxygen, the latter uniting with a molecule of
the chlorate to oxidize it to perchlorate.
3. In a conical vessel, such a as graduate, place a few
crystals of chlorate of potash, on these lay two or three
bits of freshly cut phosphorus, cover the whole with
water to a depth of a couple of inches ; then, by means
of a pipette, introduce a few drops of strong sulphuric
acid among the lumps of chlorate. Compare this experi-
ment with the preceding one and explain the result.
All chlorates yield oxygen readily. The following
experiments go to prove the truth of this statement.
4. Powder a few of the crystals of chlorate of potash
with a little charcoal and heat the mixture on a piece of
mica.
5. Powder some more of the salt with dry sugar. Place
172
OXIDES.
the mixture on a tin plate or piece of cardboard, and add
a drop or two of sulphuric acid with a pipette.
6. Powder separately, and dry on a warm plate, some
sulphur and chlorate of potash. Rub a little of tiie mix-
ture on an iron plate with a pestle or hammer. This is
dangerous, so only small quantities of the mixture should
be made and used.
7. Dissolve a crystal of a chlorate in water ; add a
little indigo solution, and then a few drops of sulphuric
acid. Explain the cause of the change of colour.
Tests for a chlorate. — Experiments 2 and 4 give tests
sufficient to distinguish chlorates from other compounds.
ill
4.— Oxides.
Oxygen forms with chlorine three well known oxides
and two hypothetical ones.
FOKMIM.A.
N A M K.
COUKKSPONDING AciD.
CI2O.
Hypochlorous anhydride.
HCIO Flypochlorousacid
Cl.,03.
Chlorous anliydridc.
HCiO^ Chlorous acid.
C1,0„(C10.,'
Chloric peroxide.
No corresponding acid.
C1005.
Not eliminated.
HCIO;. Chloric acid.
ci^o,.
Not eliminated.
HCIO 4 Perchloric acid.
Just as sodium nitrate NaNO.j yields nitric acid when
treated with sulphuric acid ; and soclic chloride NaCl
yields hydrochloric acid ; so potassmm chlorate KCIO^
QUKSTIONS AND KXKUCISIiS.
173
yields chloric acid HCIO., ; and potassium hypochlorite
yields hypochlorous acid liClO. Thus :
2NaNO,-:-H,SO,=Na,SO, + 2lTN03,
2NaCl + H,SO,=NaoSO, + 2H(:i,
2KC10,4-H,SO, = K,SO, + 2HC10,,
2KC10+H,SO, = K,SO, + 2HC10.
As chloric and hypochlorous acids, however, break up
with dangerous explosions as soon as formed, the student
is warned not to attempt to prepare them in this way.
O.
5.— Questions and Exercises.
1. How can the chlorate of potash he converted into the chloride?
2. What physical state do the compounds formed in an explosion
usually assume ?
3. How much chlorine by weight and volume can be obtained
from 1460 grams of hydric chloride ?
4. How much chlorine can be libeiated from 585 grams of
common salt? What volume will it occu])y at 60' F. ?
5. What volume will 284 grams of chlorine occupy at 80° F. ?
6. What quantities of manganic sulphate, hydro-sodic sulphate,
water, and chlorine, will be tormed by the decompositions of 351
grams of common salt, with manganese dioxide and sulphuric acid?
7. If 142 grams of chlorine gas be passed into steam at a red
heat, what substances will be formed, and what weight of each ?
8. What weight of hydric chloride will 261 grams of manganese
dioxide require for its decomposition ?
9. V/hy should chlorate of potash decompose quietly when
heated alone, but violently when heated with sulphuric acid ?
10. Write the formul.ie for sodium chlorite, magnesium chlorate,
barium perchlorate, calcium chlorite, potassium hyposulphite.
'wr
HH
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174
ADDITIONAL EXERCISES.
r ^
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7;',i<>1ii
1 1. Mention any advantage that results from the use of bleaching
powder as a disinfectant in a room, rather than chlorine prepared
as for experimental work.
12. Powder some chlorate of potash. Heat a little of it on a
piece of mica having previously placed two or three small lumps of
charcoal on the powder. Compare this with experiment 4, sec. 3.
Account for the different results.
6.— Additional Exercises.
1. Make a solution of asafnetida oy dissolving some of the
substance in a little alcohol in a small beaker. Now add a few
drops of sulphuric acid to a solution of bleaching powder, and then
try what effect this solution will have on the solution of asafcetida.
Does the oduor change ?
2. Would any other acid added to bleaching powder jjroduce the
same effect that sulphuric acid does when added to it?
3. Had nitric acid been used instead of sulphuric in sec. 3,
experiment 2, what would have been the result ?
4. Try the effect of chlorine water on any bad smelling solution
which you can secure, ammonium sulphide, for example.
5. Determine by experiment if other alkaline hydrates such as
those of sodium and potassium will yield the analogues of bleaching
powder. Write the equations in full for the reactions. Will the
substances obtained bleach dyed goods? In these experiments
keep the solutions cool.
6. What would be the result of passing a current of chlorine into
Ume water?
SULPHUR.
175
CHAPTER XXXIII.
1.— Sulphur.
Experiments.
1. Place some sulphur on a metal spoon and hold it
in a flame until ignited, then withdraw it. Notice the
colour of the flame and the smell of the oxide that comes
off. Do this by wafting the fumes toward the face with
the open hand.
2. Place 15 or 20 grams of sulphur in a large test-tube
and heat slowly over a lamp flame until the sulphur
boils. As the heating goes on, note changes in the
appearance of the substance. When it begins to boil,
pour it into a vessel of cold water. When it has cooled,
remove and examine. Burn a little of it as in the last
experiment. Compare the colour of the flame and the
odour of the vapour with those noticed before. Keep in a
dry place for a few days, and then examine again.
3. Try if yellow sulphur is soluble in water, alcohol,
chloroform, ether, spirits of turpentine, or carbon bi-
sulphide.
4. Powder some iron pyrites, FeSo, and place in a hard
glass test-tube. Hold the test-tube nearly horizontally in
the lamp flame and heat the lower end strongly for some
time. Test the residue with a magnet. Observe what
collects in the cool part of the test-tube.
2.— Questions and Exercises.
I. What physical property of sulphur is illustrated in the fourth
experiment 1
17G
AI)r)ITlO>fAL KXKIUMSKS.
2. The residue from the iron pyrites has the composition FcaS.j.
Write the equation e\pressin>^ the reaction that took place.
3. IVeparc some sulphur crystals by melting 40 or 50 grams of
fltnvers of sulpluu" in a porcelain or earthenware dish, and then
allowing it to cool slowly. When a thin crust has formed on the
surface punch a couple of holes in it with a glass rod, pour out the
liquid sulphur from the interior, and after the dish has cooled, care-
fully remove the sinface crust.
4. Try to prepare some more crystals by dissolving" sulphur in
carbon disulphide and allowing the solution to evaporate.
5. Try whether plastic sulphur is soluble, as in experiment 3, sec. i.
6. Describe in detail the changes which you noticed when sulphur
was heated in a tube to boiling.
7. Fuse a little sulphur with some sodium carbonate ; lay the
fused mass on a piece of clean silver, and let a drop of water fall
on it.
This is a test for free siilpiiur.
2Na2C03 + 3S=2Na.S 4- SO2 + 2CO.J,
and NagS+HaO^NaaO'+HaS,
and HoS + 2Ag=AgoS + 2H. "
3.— Additional Exercises.
1. Bend a piece of hard glass tubing as shown in Fig. 44. Place
in it at A some iron pyrites ;rnd heat
strongly in a flame, holding the
longer arm of the tube at about
45° with the vertical. Smell the gas
that comes off. How do you account
for the result being different from
that obtained in experiment 4, sec. i ?
Repeat the experiment, using ga-
lena, or lead sulphide, instead of pyrites.
2. Boil some sulphur in a tube, and when the vapour is coming
off freely lower into it separately a bit of sodium, a httle coil of
fine copper wire, and some fine iron filings. Then drop some
hydrochloric acid on some fresh copper and iron filings, smell the
gas that comes off. Treat the coil and filings that were in the
sulphur in a similar way.
ft.;
AUNOHMAL VAPoUn DHNSITY OP SULPHUR.
177
Explanation — When iron is roasted with sulphur,
the sulphide FcS is obtained. This, when acted on by
hydrochloric acid, HCl, is decomposed, ferrous chloride,
F'jCl2, and sulphuretted hydroiren, 1I._,S, beiii^ the result.
The latter was the foul smelling gas that escaped.
ga-
4.— Abnormal Vapour Density of Sulphur.
Read Chapter XXVII.
The following is a tabulated statement of the molecular
weights and densities of a few of the gases we have
already become ac(juaintcd with : —
Moii.
FORMl'bA.
Moii.
Wkioiit.
MOL.
Voiil'MK.
2
2
2
o
2
o
2
2
2
2
Vapour
Dkssity.
Hydrogen
Oxvcen
H,
0.
N,
H^O
x,o
NO
NH,
CO,
CO
CH^
2
32
28
18
44
30
17
44
28
16
1
16
Nitrogen
Ac^ueous Vapour
Nitrous oxide .
Nitric oxide
Ammonia
14
9
22
15
8-5
Carbon dioxide
Carbon monoxide
Marsh cas
22
14
8
From the examples given above it will be seen thai in
the column headed " Vapour Density" the numbers ex-
12
IMAGE EVALUATION
TEST TARGET (MT-3)
10
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1.8
1.25
1.4
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WEBSTER, N.Y. 1458C
(716) 872-4503
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178
NOTES ON SULPHUR.
press the ratio between the molecular weight and the
molecular volume of any gas. From Avogadro's Law we
know that equal volumes of gases under like conditions
contain equal numbers of molecules, hence any unusual
vapour density must be due to an unusual number of
atoms in the molecule rather than to any departure from
the number of molecules in a unit volume of the gas
showing the irregularity.
In the case of sulphur, it is found that near its boiling
point (485°) the vapour density is 96, while at 850° this
density is 32, in round numbers. When the vapour
density is 96 the molecular weight must be 192, but at
the higher temperature mentioned the molecular weight
will be 64. Chemical analysis leads to the belief that
sulphur has an atomic weight of 32. From this it is
clear that, at about 500°, the molecule of sulphur consists
of 6 atoms ; but at 850° and above that to 1200°, the
molecule has in it only 2 atoms. A change of tempera-
ture, therefore, modifies the molecular structure of the
substance, as is indicated by the change in the vapour
density.
The vapour density of sulphur is treated in R. & S., p. 292, Vol. I.;
Rem., pp. 89, 192 ; Wurtz, pp. 69-71 ; Tilden, p. 127-128; Rem. Th.
Ch., 44.
5.— Notes on Sulphur.
Symbol, S ; atomic iveight, 32 ; specific weight in the
form of crystals^ 20^ {waters i).
Sulphur, known also as brimstone, is found native in
many volcanic regions ; it occurs also in the ores of
some of the common metals. Iron pyrites FeSg, galena
OXIDES OF SULPHUR.
179
PbS, cinnabar HgS, gypsum CaS04+2H20, and heavy
spar BaSO^, all contain sulphur. It is also found, as
sulphuretted hydrogen, HgS, dissolved in the waters of
some springs, and in many substances of organic origin,
such as albumen and coal.
6.— Allotropic Modifications.
Sulphur is known in three different forms ; of these,
two are modifications of the shape in which sulphur
crystalizes ; the third one, known as plastic sulphur, is
prepared as described in experiment 2, sec. i, and is
used for making moulds of coins, etc.
Common yellow sulphur, known as flowers of ^:ulphur,
or when run into cylindrical moulds, roll sulphur, may be
taken as the first allotropic form. The crystalline struc-
ture can be studied with a magnifying glass. The second
crystalline form is prepared as in experiment 3, sec. 2.
CHAPTER XXXIV.
Oxides of Sulphur.
There are two oxides of sulphur known — the dioxide
SO2 and the trioxide SO^. Of these only the former will
be treated of here, because the latter is difficult of pre-
paration and comparatively of little importance.
¥
' i
180
SULPHUR DIOXIDE.
1.— Sulphur Dioxide.
Experiments.
1. In a flask, fitted with a cork and delivery tube, heat
some copper clippings and sulphuric acid.
Cu+H2S04=:Cu£04 + 2H.
and 2H+H,S04=2H,0-f-S02.
at
These equations are usually written thus : —
Cu + 2H,SO,-CuS04 + 2H20+S02.
2. Collect a couple of bell jars full of the gas, test for
solubility, acidity and inflammability.
3. Hang a red rose or other high-coloured flower in the
second jar. If any change takes place in the flower, re-
move it and place in pure air.
4. An easy method for the preparation of sulphur
dioxide is to treat some sodmm hyposulpJdte Na^SgOg
(which is chemically sodium thiosulphate) with sulphuric
acid.
Na2S203 + H2S04= Na2S04 + H.S.Og
and H2S203=H20 + S02 + S.
Try this method.
5. Pass a current of the gas into solutions of logwood,
indigo and permanganate of potash.
2— Questions and Exercises.
I. Put some sulphur on a plate, set fire to it and turn an empty
bell jar over it. After combustion has ceased compare the gas as
regards solubility, odour and acidity with that obtained from copper
and sulphuric acid.
I
NOTES ON SULPHUR DIOXIDE.
181
2. Zinc and cold dilute sulphuric acid yield free hydrogen. Try
if the same result follows if the acid be strong and hot. Smell the
escaping gas, collect a test-tube full of it and try if it is all, or in
part, hydrogen. Account for the results you obtain.
3.-~ Additional Exercises.
1. Pass some sulphur dioxide gas through a wash bottle, then
into a solution of barium nitrate, add hydrochloric acid.
2. Pass some of the gas into water until the solution is distinctly
acid, let this stand in an open vessel for a day, then test with
barium nitrate and hydrochloric acid.
3. Heat together sulphur and sulphuric acid.
S + H3S04 = 2S02 + 2H,
2H + H2S04 = 2H20 + S02,
hence S + 2H2S04.=:3SO.^ + 2H20.
4. Heat together cliarcoal and sulphuric acid.
C + 2H2S04=C02+ (complete this).
5. Heat together manganese dioxide and sulphur.
Mn02 4-2S=MnS + S02.
6. Prepare some thin starch paste, add to it a little solution of
potassic iodide, then drop into it a little chlorine water. After
it has become blue make a thin solution of it in water in a test-tube,
pass a current of sulphur dioxide through this.
7- Try if any sulphite which you can obtain will yield sulphur
dioxide when treated with any strong acid.
4.— Notes on Sulphur Dioxide.
Symbol, SO^; molecular weight, 64; vapour density
(H=i)32.
Sulphur dioxide, on account of its strong affinity for
oxygen, is a pov/erful reducing agent. It is used exten-
sively both as a bleaching agent and as a disinfectant
182
ACIDS OF SULPHUR.
The power of sulphur diox'de to act in this way is
supposed to be due to its affinity for oxygen. According
to this theory the dioxide unites with the oxygen of the
water which is used to moisten the article to be bleached;
the hydrogen of the water then combines with the color-
ing matter and forms colourless compounds.
Tests (i) Odour.
{2) See ex. 1,2 and 6, sec. 3.
This gas is the anhydride of sulphurous acid, H.^SO^.
CHAPTER XXXV.
Acids of Sulphur.
Culphur, in union with hydrogen and oxygen, forms
two well known acids, sulphurous and sulphuric
The following exercises will illustrate their preparation
and some of their chief properties.
1.— Sulphurous Acid.
Experiments.
1. Pass some sulphur dioxide gas slowly into water in
a bottle. After some time test the water with litmus.
2. Divide the solution made in the previous experi-
ment into two parts, with one of these fill a small bottle,
cork it tightly and set it away for a couple of days.
Pour a few drops of the other part into a test-tube, add
SULPHURIC ACID.
183
a little solution of barium nitrate (there should be no
precipitate), then a drop or two of silver nitrate solution;
set the remainder of this part away in an open beaker
beside the bottle just mentioned. At the end of two or
three days test each part of the liquid again. (Barium
nitrate alone gives a white precipitate with sulphuric
acid, but in the case of sulphurous acid this precipitate
does not appear until silver nitrate is added.)
2.— Sulphuric Acid.
Experiments.
1. Invert a jar of sulphur dioxide gas over a small
beaker containing a little strong nitric acid.
2. Pour a little water into a large flask, then pass into
it simultaneously brown fumes from a heated metallic
nitrate and sulphur dioxide gas, also allow air free access
into the flask. After about ten minutes, remove the
connections, shake the water well with the gas in the
flask, then test the liquid with barium nitrate.
Explanation. — Sulphur dioxide in presence of water
or steam will readily reduce either nitrogen trioxide or
peroxide to nitric oxide, and thus become itself oxidized
to sulphuric acid.
SO,-f-N02+H20=:H2S04 + NO.
or, SO2 + N2O3+H2O-H2SO4+2NO.
In either case the nitric oxide in presence of air at once
changes back to the trioxide and is ready to yield up
oxygen again ; thus, this gas serves simply as a medium
for transferring oxygen from the air to the sulphurous
acid.
184
I.
i 1 :
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HYDROGKN AND SULPHUR.
3.— Questions and Exercises.
1. Is barium sulphate soluble in nitric acid, sulphuric acid or
ammonia? Is lead sulphate soluble in these liquids?
2. Sulphuric acid is prepared in quantity by leading into a lead-
lined chamber the gases formed : — •
(a) By roasting in air iron pyrites,
2FeS2 + iiO=Fe203 + 4S02.
(6) By heating together sodic nitrate and sulphuric acid,
NaN03 + H»S04=
Steam is also passed into the chamber. The reaction may be
represented as follows : —
SO.+ +H20=H2S04 +
Complete the equations.
Why use lead as a lining for the chamber in preference to iron or
zinc ? Would you suspect that lead sulphate would be an impurity
of the commercial acid ?
3. Measure out accurately about 20 c.c. of strong sulphuric acid
and an equal portion of water, then pour the acid in a thin stream
into the water. What volume is there of the mixture? If change
of volume and change of temperature indicate chemical action,
how bhould this result be classed ?
4. Given some sulphur, nitric acid and other necessaries, how
would you proceed to prepare sulphuric acid ?
5. A bottle containing sulphurous acid should be kept quite full.
Why ?
4.— Hydrogen and Sulphur.
Sulphur and hydrogen form one compound that is of
very general application in chemical operations, and is an
agent extensively applied in the separation of some of
the more commonly occurring elements. This is hydro-
gen sulphide^ sulphuretted hydrogen or hydrosulphurie acid,
HoS.
HYDROGEN AND SULPHUR.
185
Experiments.
1. On some iron sulphide (prepared by roasting a mix-
ture of iron filings and about two-thirds as much sulphur,
by weight, in a closed crucible) pour some dilute sul^
phuric acid. FeS+ H^SO^-^what ?
2. Try if the gas given off is soluble in water ; if you
have reason to believe that it is, find whether the sup-
posed solution is acid, alkaline or neutral.
3. Try whether the gas given off will burn. In case
that it does, hold a cold, dry beaker inverted over the
flame. Fasten on a piece of glass rod a strip of moist
litmus paper, blue at one end, red at the other, and raise
it into the beaker. What is being formed in the beaker?
4. Dip some pieces of soft white paper in a solution of
acetate of lead, hold one of these in a current of hydro-
gen sulphide, dip another into a solution of the gas.
5. Let a drop of solution of the gas fall on a bright
silver coin.
6. Make a bottle half full of strong solution of the
gas, cork it tightly and set it away for a month. How
do you account for the gray powder? Smell the liquid.
Test it with the acetate of lead paper.
7. Pass a current of chlorine into a solution of hydroo-en
sulphide. 2Cl-{-H2S = what ?
5.— Questions and Exercises.
I. What is your opinion regarding the stability of hydrogen
sulphide as a chemical compound ? Quote as many proofs as you
have seen in support of that opinion.
186
ADDITIONAL EXERCISES.
|i
2. Hydrogen sulphide is one of the constituents of illuminating
gas before it is purified. It is exceedingly objectionable, although
readily combustible. Why .''
3. Hydrogen will reduce heated copper oxide to metallic copper ;
sulphur will burn freely to form the dioxide. What should be the
effect of hydrogen sulphide on such a substance (oxide of copper)
when heated ?
6.— Additional Exercises.
1. Drop a lump of white sugar into some strong sulphuric acid
in an evaporating dish ; let it stand for 24 hours, then dilute largely
with water, filter, wash with water, dry and examine carefully.
Try if a little of the black substance will burn on mica. Heat some
of it in a combustion tube and lead the gas into lime water.
2. Repeat the preceding experiment, but instead of sugar use
sawdust or wood shavings.
3. Try if any other sulphide such as that of copper (prepared in
a way similar to that of iron) or lead, (galena) will yield sulphuretted
hydrogen. Is it necessary that sulphuric acid should be used.''
4. Repeat experiment 3, sec. 4, but use a bell jar fitted with
a stopper instead of the beaker. When the jar is filled with the
acid gas, set it, mouth downward, on a plate that has a little water
on it, then extinguish the burning gas and turn the current of
sulphuretted hydrogen into the bell jar.
SSOa + sH-jS^sS + HaSgOg (pentathionic acid).
5. Make solutions of jny salts of lead, iron, zinc, copper, mercury,
barium, antimony, pota =iiu':>. sodium, magnesium. Pass sulphur-
etted hydrogen into a part of each (if more convenient, a strong
solution of the gas may be prepared and some of this poured into a
solution of each salt). If no precipitate is obtained add hydrochloric
acid to a part of the solution until it is distinctly acid, then use the
sulphuretted hydrogen again. If no precipitate still, add ammonia
to each solution until it is alkaline, then try again with the gas.
6. Ammonium sulphide (N H4)2S is prepared by passing hydrogen
sulphide through a solution of ammonia ; do it. Will this give
NOTES ON HYDHOGEN SULPHIDE.
187
precipitates with any of the solutions used in the preceding experi-
ment?
7. Pass a current of steam through a combustion tube in which
some sulphur is heated. What comes off? Substitute hydrogen
for steam. What now is formed?
8. Given a constant current of sulphide of hydrogen, some nitric
acid and some pieces of iron wire, how could you illustrate the
preparation of sulphuric acid .•*
9. Compare oxide of hydrogen and sulphide of hydrogen as to
(i ) the compounds from which they are prepared, (2) their syntheses,
(3) their acidity, (4) their stability, (5) their chemical action with
metallic salts, (6) their molecular constitution.
7.— Notes on Hydrogen Sulphide.
Symbol, H,S ; mol. wt., j^ ; mol. vol., 2 ; sp.gr. rig,
(air—i).
It is a poisonous gas, soluble in water in proportion
of 3 to I by volume, occurs frequently in natural water,
particularly of springs, and is formed largely in the
decay of organic matters.
In analysis of chemical compounds sulphuretted
hydrogen is a valuable group test for the metals, that is,
by its aid they are divided into three classes :
(i) Those whose sulphides are precipitated from
acid solutions (copper group or hydrogen
sulphide group). The common metals of this
group are copper, lead, mercury, silver, bis-
muth, antimony, arsenic, tin.
(2) Those whose sulphides are precipitated from
alkaline solutions only, such as iron, nickel,
cobalt, zinc.
188
CALCULATION OP FORMULA.
I
(3) Metals whose sulphides are not precipitated from
any solution ; the alkalies are examples.
Tests for the gas are (i) its odour, (2) its effect on
acetate of lead solution, or paper dipped in it, and (3) its
effect on silver.
CHAPTER XXXVI.
Calculation of PormulsB.
•
In chapter XXV the calculation of empirical formulae of
compounds was discussed when the percentage composi-
tion was known. At that time one important fact for
the accurate determination ot the molecular formulae had
not been learned, viz.: that the vapour density of a sub-
stance is one-half its molecula" weight. Vapour density
is always taken with hydrogen as the unit. A couple of
examples will be solved to show the application of this
principle : —
I. A compound, on analysis yielded
hydrogen, 2-25%
carbon, 26-65 %
oxygen, 71-2 %
Its vapour density is 45, find its formula.
2*25 -^ I = 225
26-65 -r 12 =^ 2-25
71-2 ^ 16 = 4-45
Neglecting what are probably errors of experirnent, the
IMPURITIES IN AIR AND WATEll.
189
elements are present in proportion of i, i and 2. The
formula may, therefore, be either HCO,, HaC^O,, H.C.Oo,
etc. The vapour density is 45, therefore the molecular
weight is 90. Now, starting with the lowest empirical
formula, we find that it gives a molecular weight of 45,
just half that required. Wc must, therefore, double the
number of atoms, and write the formula Hfi.fi^ (oxalic
acid).
2. A hydrocarbon, when analyzed, gave hydrogen,
77 % ; carbon, 92*2 % ; its vapour density is 39, determine
its formula.
77 T- I = 77
922 -^- 12 = 77
Therefore the proportions of hydrogen and carbon are as
I to I. Hence the formula is H,.C„, where n is any
integer. The molecular weight of the substance is
39x2-^78.
The molecular weight of HC is 13.
78-7-13=6, hence formula is CqUq (benzine).
CHAPTER XXXVn.
Impurities in Air and Water.
It is desirable that every one should be able to deter-
mine, approximately at least, the degree of purity of the
two substances which are most necessary for our exist-
ence, in order that hurtful impurities may be removed
190
IMPURITIES OP AIR.
or rejected. These substances are the air we breathe
and the water we use for drinking and for domestic
purposes.
i::
; I
1.— Air.
The atmosphere is a mixture of a number of gaseous
substances some of which are quite variable in quantity ;
but it is generally considered that a mixture of oxygen
and nitrogen in the proportion of 21% by volume of the
former to 79% of the latter shall be taken as pure air.
The chief gases mixed with these are aqueous vapour,
carbon dioxide, and traces of ammonia It would seem
from recent investigations that are not yet finished (Feb'y,
1895) that there is a third constant constituent of the
atmosphere, Argon, and possibly a fourth ; but at present
it may be passed over with a mere mention.
Two experiments maybe repeated for determining the
carbon dioxide and vapour of water in the atmosphere.
The first is 3, Chap. XXVII, sec. 4. This experiment may
be varied by driving a measured volume of air backward
and forward a number of times through a weighed
quantity of strong solution of caustic potash, then weigh-
ing the solution.
When it is necessary to test the purity of air for breath-
ing in such places as school rooms, dwellings and lecture
halls, the quantity of carbonic acid gas per thousand
volumes is generally taken as the test of purity^ This is
not an absolute test, for tliere may be and indeed gener-
ally are other objectionable and deleterious products of
respiration present, but as they always accompany the
■
IMPURITIES OF WATER.
191
Mcrh-
carbon dioxide the latter is used as the basis of the mea-
surement. There are about 4 parts of carbon dioxide
to 10,000 of air in the atmosphere. When the proportion
rises above 10 in 10,000, on account of impurities due to
respiration, the air becomes very objectionable for
breathing.
i|i!
2.— Water.
Pure water is both scarce and difficult to prepare.
Probably the purest natural water is that which has
recently fallen as rain, away from the neighbourhood of
towns and factories. It is then much in the condition of
the water prepared by distillation. Water which has lain
in contact with the earth for some time is sure to become
impregnated with mineral salts and decaying matters of
various kinds.
Natural waters are classified as hard and soft.
Water that contains magnesium, and calcium salts,
and that curdles soap, is said to be Jiard; water that does
not contain these salts is soft. Hardness is usually con-
sidered as being of two kinds, viz., temporary and per-
manent. The former is due to the presence of calcic and
magnesic carbonate, the latter to the presence of salts of
calcium and magnesium other than the carbonates, such
as the sulphates and nitrates.
Water that is temporarily hard may be softened by
boiling, because the carbonates are held in solution by
the carbonic acid dissolved in the water or the bicar-
bonate is itself soluble. Boiling expels the carbon
dioxide or decomposes the bicarbonate and the carbonate
is precipitated.
192
IMPURITIES OP WATEft.
Water that is permanently hard may be frequently
softened by the use of washing soda — neutral sodium
carbonate, NaaCOg. (Compare Chap. XXVI.)
Complete the two following equations. When com-
plete they represent the reaction of washing soda on two
kinds of hard water.
H2Ca(C03)2 + Na2C03=
CaS04 + Na2C03=
Water suitable for drinking is described as potable.
That which comes from springs generally, contains
mineral salts, such as the carbonate of calcium or other
substances through which the water has trickled, in solu-
tion. These salts are not necessarily objectionable —
indeed the flat taste of rain water and of distilled water
is due to the absence of them, and to the lack of aeration.
Organic matters, howe-ver, when held in solution,
frequently render water dangerous to use. One test for
such impurities depends on the decolorization of per-
manganate of potash by them.
1. Place the water to be tested in a flask, and add
to it, first, a few drops of sulphuric acid, and then
enough of a solution of permanganate of potash to give
to the whole a purple tint. Set to one side for three or
four hours in a warm place, and if the solution loses its
colour, orL:^anic impurities are present. Water that will
thus decolorize permanganate of potash is in all probabil-
ity unfit for drinking. If it is necessary to use such water,
it should first be boiled for at least half an hour.
2. Rub a chalk crayon on the hands until they become
covered with the dust, then try to make a lather with a
little water and soap.
'i
MOLECULES OF ELEMENTS.
193
CHAPTER XXXVIII.
1.— Molecules of Elements Usually Consist of
More Than One Atom.
The only perfectly reliable means which we possess
for ascertaining the molecular weight of a compound is
the determination of its vapour dettsity.
It follows from Avogadro's Law that the weights of
individual molecules of different gases is proportional to
the weights of equal volumes of these gases. All we
have to do then, in order to find the relative weights of
molecules of different gases, is to weigh equal volumes of
them under like conditions of temperature and pressure,
and the numbers thus obtained will represent the
relative weights of a single molecule of each gas.
Manifestly, any gas might be taken as a standard with
which to compare the weights of all other gaseous
substances ; but, for many reasons, it has been found
preferable to take hydrogen as the unit of comparison.
The following facts have been established by actual
weii^hing : —
1 litre of oxygen weighs .... 1 .429 gram.
I " nitrogen " .... 1.2553 "
I " chlorine " .... 3.167
I •' hydrochloric acid gas weighs 1.6283
I *' hydrogen weighs . . . .0896
(The weight of hydrogen is obtained by calculation from the two
preceding data, because it is exceedingly difficult to weigh a litre of
hydrogen accurately, on account of its lightness.)
Now, using hydrogen as the standard of comparison,
it follows from the above data that oxygen is nearly
13
((
«
(t
194
OXYGEN MOLECULES.
sixteen times heavier than hydrogen ; nitrogen, nearly
fourteen times heavier; and chlorine, 35.34 times heavier.
Hence, these figures represent the number of times that
a molecule of each of these elements is heavier than a
molecule of hydrogen. It follows, further, that if we
know the actual number of atoms composing each of these
molecules, we should be able to calculate their atomic
weights. If there are the same number of atoms (say
two) in each molecule of these elements, the above
figures will also represent their atomic weights, one atom
of hydrogen being taken as the standard. Of course, no
one knows how many atoms there are in the molecule
of any of the elements, but the following considerations
will help the pupil to understand the conventional ideas
on this subject.
2.— The Molecule of Oxygen Consists of at
Least Two Atoms.
Two volumes of hydrogen and one volume of oxygen
unite to form two volumes of steam.
From this it follows that two molecules of hydrogen
and one molecule of oxygen unite to form two molecules
of water. In one molecule of water there must be one
molecule of hydrogen and half a molecule of oxygen,
therefore this half molecule must consist of at least one
atom.
3.— The Hydrogen Molecule Consists of Two
Atoms at Least.
We y seen, in Chap, xxxi, that two volumes of
hydrocL.oi ic acid gas may be broken up into one volume
HYDROGEN MOr-ECULES.
195
of hydrogen and one volume of chlorine. One volume
of the acid may, therefore, be divided into half a volume
of hydrogen and half a volume of chlorine. Then one
molecule of the acid consists of half a molecule of each
constituent, and this half molecule must be at least one
atom ; hence, the molecule of hydrogen has in it two
atoms at least.
Since this substance is the standard for vapour density
comparison, we have to rely on other considerations for
the proof that the molecule consists of only two atoms.
Some of these are : — In compounds of hydrogen with
monad elements the combinations and decompositions
take place each at one stage; never is part of the hydro-
gen freed from the other element, and then by changed
or intensified treatment, the other part liberated. On
the other hand, when hydrogen unites with a diad
element, half of it may frequently be displaced at once,
and the other half at another time. Thus : —
H20+Na= NaHO+H
and NaHO+Zn = ZnNaO+H.
With monads such displacements are manifestly
impossible.
When decomposition of such compounds (hydrogen
with monad elements) occurs, the hydrogen always
occupies one half the space of the original gas ; hence,
from two molecules of the compound, one molecule of
hydrogen is set free. The same conclusion is arrived
at from the consideration that in equal volumes oi
hydrogen and hydrochloric acid gas, the weight of
hydrogen in the latter, when freed, is just half that of
the former ; hence, in equal volumes of hydrogen and
196
NITKOOEN MOLECULKS.
hydrochloric acid, the number of molecules being equal,
the number of hvdrofjen molecules formed from the
latter gas equals half that existing in the former. Since
in chemical decompositions the quantity of hydrogen
freed from combination with monad elements is the unit
of volume, of which that liberated from other combina-
tions is always an integral multiple, it is reasonable to con-
clude that we have here the smallest subdivisions of the
hydrogen molecule which have existence, ?>., half
molecules or atoms.
4.— Nitrogen Molecules.
When nitrous oxide was decomposed by burning
potassium (see Chap. XX), a volume of nitrogen equal
to that of the original gas remained. When nitric oxide
was similarly treated the nitrogen remaining was half
that of the gas taken. Now it will be evident that equal
volumes of the two oxides contain equal numbers of
molecules, and that every molecule of the nitrous oxide
contains nitrogen sufficient to form one molecule of that
gas, while in the case of the nitric oxide each molecule
contains only half a molecule of nitrogen, hence the mole-
cule of nitrogen is divisible into two equal parts, hence,
contains at least two atoms.
5. — Chlorine Molecules.
We have learned that the hydrogen molecule has in it
two atoms ; also one volume of hydrogen unites
with one volume of chlorine to form two volumes
of hydrochloric acid. The analysis of the latter
OTHER ELfilMENTS.
197
shows that it is composed of equal parts, by volume, of
hydrogen and chlorine; then, since one volume of
the gas is made up of half a volume of hydrogen and
half a volume of chlorine, it follows that one molecule of
it is composed of half a molecule of hydrogen and half a
molecule of chlorine ; hence, the chlorine molecule is
divisible into two equal parts, or at least into two atoms.
6.— Other Elements.
Starting with the compounds marsh gas and sulphur
dioxide, the conclusion follows that sulphur and carbon
molecules are also divisible into at least two equal parts
or atoms.
The student must not understand, however, that this
is a proof that there are only two atoms in the molecules of
these substances. While that is probably the case with
most elements, it has already been shown that for sul-
phur there are six atoms in the molecule at certain tem-
peratures. Phosphorus and arsenic have each a four-
atom molecule, while ozone has three, and mercury one.
In the case ot compounds, it follows directly from the
atomic theory that the molecule must consist of a group
of atoms,— one at least from each constituent.
19S
SELECTED QUESTIONS.
CHAPTER XXXIX.
The following questions have been selected, partly
from recent Pass Matriculation and Junior Leaving
examination papers, partly from other sources. They
are here simply as an indication of the standard of effi-
ciency which scholars have been required to reach in the
past few years.
I. SERIES.
1. On what grounds do you consider Hydrogen and Oxygen
to be chemical elements, and water to be a compound of these two
elements ?
2. Describe as fully as you can, the phenomena of a solution
of a salt in water.
3. A test-tube is known to contain distilled water, or a solu-
tion of one of the following: Ammonia Gas, Potassium Hydrate,
Potassium Chloride, Nitric Acid. How would you determine most
simply which the test-tube contains ?
4. Explain by means of equations, how each of the following
substances bleaches :
(a) Chlorine in the air.
(d) Chlorine in a solution of water.
(c) Sulphur Dioxide Gas.
5. Sulphur 1 >ioxide : how prepared ? how converted into
Sulphur Trioxide .'' How would you prove that Sulphur Dioxide
contains its own volume of Oxygen .''
6. Calculate the weight of the product or products in each of
the following cases :
{a) One gram of Carbon Monoxide burned in Oxygen.
{d) One gram of Ammonia Gas burned in Oxygen.
{c) One gram of Sodium burned in Chlorine.
aELEOTED QUESTIONS.
109
7. How is Ammonia Gas prepared from Ammonium Chloride?
Calculate how much heavier it is than Hydrogen and how much
lighter than Nitrogen ? How could you show that it contains both
Nitrogen and Hydrogen?
8. Describe experiments showing how you would distinguish
(a) Carbon Monoxide from Hydrogen,
(d) Carbon Dioxide from Nitrogen.
(c) Marsh Gas from Hydrogen. '
II. SERIES.
1. (a) Describe experiments to show that one cc' of Hydrogen
Gas and one cc. of Chlorine Gas are found in two cc. of Hydro-
chloric Acid Gas, and one cc. of Oxygen Gas and two cc of
Hydrogen Gas in two cc. of Water Gas.
(d) Draw the inference from the above experiments that the
ratio of the weight of two cc. of each of these compound gases
to the weight of one cc of Hydrogen is twice the Specific Gravity
of the Compound Gases compared to Hydrogen.
2. Discuss the question as to the distinction between a com-
bustible substance and a supporter of combustion. Illustrate by
equations the chemical reactions which occur in the combustion of
(a) Hydrogen in Chlorine.
{&) Oxygen in Marsh Gas.
(c) Carbon Monoxide in Oxygen.
(d) Sodium in Hydrochloric Acid Gas.
(<?) Hydrogen Sulphide in Oxygen.
3. Explain the meaning assigned by Chemists to the following
terms : (a) Oxidizing Agents, (d) Reducing Agents ; write equa-
tions showing instances of oxidation, (c) by Oxygen Gas, (d) by
Chlorine Water, {e) by Nitric Acid ; of reduction (/) by heat, (g)
by Charcoal, (A) by Nascent Hydrogen.
4. Describe the physical changes and illustrate by equations the
chemical changes which occur when each of the following sub-
stances is heated in a test-tube, (a) Ammonium Nitrate, (d) Potas-
sium Nitrate, (c) Lead Nitrate, (d) Calcium Carbonate, (^) Ammonium
Chloride.
B
200
SELECT RD QUESTIONS.
5. Name and give the formul.ne of the substances formed by the
action of hot Concentrated Sulphuric Acid upon each of the fol-
lowing bodies : (a) Lopper, (d) Charcoal, (c) Potassium Chlorate,
{(/) Ammonium Nitrate, {c) Ammonium Chloride, (/) Calcium
Carbonate.
6. Explain the chemical and physical reactions which occur in
the following experiments — give equations in each case :
(a) A small piece of Sodium is thrown upon Water.
(d) A small piece of Potassium is thrown upon Water.
(c) Chlorine Gas is mixed with Hydrogen Sulphide.
(d) Charcoal is heated with Sulphur Vapour.
(e) Nitrogen Trioxide is mixed with Sulphur Dioxide.
7. Describe experiments showing how you would distinguish
(a) Oxygen from Nitrous Oxide.
(d) Nitrous Oxide from Nitric Oxide.
8. Write equations explaining the reactions when Chlorine is
passed into
(i ) Dry Ammonia,
(ii.) Solution of Potassium Iodide,
(iii.) Hot solution of Potassium Hydrate.
9. In what respects do the properties of Hydric Nitrate differ
from Potassium Nitrate and Potassium Hydrate ?
III. SERIES.
1. (a) State the effect of Carbon Bisulphide upon each of
the forms of Sulphur.
(d) What is the action of Hydrogen Sulphide upon Sulphur
Dioxide ?
2. How would you prove the presence of
(^i!) Hydrogen and Sulphur in Hydrogen Sulphide,
{d) Carbon in Carbon Dioxide,
(c) Nitrogen in Ammonia ?
3. Give one illustration in each case showing the relations of
Electricity, Heat, and Light, as a cause, and an effect of chemical
action.
SELECTED QUKSTIONS.
201
4. Describe, giving equations, what occurs in each of the follow-
ing experiments :
(a) Copper wire and strong Sulphuric Acid are heated together
in a flask and the gaseous product passed into a solution of caustic
potash.
{d) Hydrochloric Acid is added to pulverized Barium Dioxide
and the resulting mixture boiled.
5. Each of five bottles contains one of the following gases :
Hydrochloric Acid Gas, Sulphur Dioxide, Carbon Monoxide,
Nitric Oxide, Carbon Dioxide.
Describe how you would most easily determine the gas in each
bottle.
6. (rt) State briefly one of the theories usually held regarding
solution.
(d) Describe two methods of determining the percentage com-
position of sand and ammonium carbonate present in a mixture of
100 grams of these substances.
7. When houses are heated with hot water passing through
iron pipes from an iron furnace, a substance which plumbers call
" air " collects in the uppermost parts of the pipes. This "air"
burns with a pale-blue flame and forms a mist on any cold solid
held over it. Name the gas and explain its formation by means of
an equation. Give also the product of its combustion.
8. When a coal fire gets low, then throwing much coal on it,
or greatly increasing the draft will frequently put the fire entirely
out. Why.-* Describe an experiment which illustrates the correct-
ness of your explanation.
9. Heat in a flask fitted with cork and delivery tube a mixture of
dry powdered quicklime and amnionic chloride. Pass the gas that
comes off into pure water until no more will dissolve. Neutralize
this water with pure nitric acid and then evaporate to dryness.
Heat on a piece of mica the solid that remains. Name the final
products and explain the whole series of changes.
10. Explain the meaning of the following equations: — H;jP04
— HaO^^HPOs- How is the operation carried out in the labora-
tory ? What is the Usl for the last substance or its salts ?
202
SELECTKD QUKHTIONS.
1 1. (fi) If you are in doubt as to whether A solid is soUible in water
or not, describe an experiment wliich you would perform to decide
the point.
(d) How would you separate a finely powdered mixture of sand,
sugar and iron, so as to preserve the first two substances?
12. (a) Describe fully how you would [irepare some metallic copper
from copper sulphate, and metallic silver from silver nitrate.
{d) How would you prove the presence of the acids in these
salts ?
13. Explain, using equations, the reactions that occur when
(a) Carbon dioxide is passed over red hot charcoal,
(d) Dry hydrogen is passed over red hot copper oxide.
14. Compare the action of hot sulphuric acid on copper with that
of strong nitric acid on copper. Give equations.
IV. SERIES.
1. Nitric acid may be prepared by heating sodic nitrate with
sulphuric acid. When the other substance formed is acid sulphate
of sodium, find in what proportions the substances must be taken
that none of either may be left. What is acid sulphate of sodium ?
When is such a salt possible ? Write the equation for the other
reaction possible between the substances.
2. A current of ammonia gas is passed into water for some time,
then nitric acid is added until the solution is neutral ; afterwards
the liquid is evaporated and the residue heated in a test-tube, what
will be the final products ? Give equations.
3. Four volumes of methane are mixed with six volumes of
oxygen and the mixture exploded. Find the volume of the gas
in the vessel and state its composition (i) at a temperature of 120°,
(2) after the products of combustion have stood in a room at 20° c.
for some time.
4. Make a strong solution of ammonic chloride in a beaker, test
the solution with litmus. Then hang a piece of litmus paper just
above the liquid, but not touching it, place another piece in the
solution, and boil the contents of the beaker for half an hour. State
what will occur, and explain the chemical action.
SELECTED QUESTIONS.
203
5. Some chloiine, carbon dioxide and sulplnir dioxide are mixed
in a jar, this jar is tlicn placed mouth downwards over a dish con-
taining-^ a sf Union of sodium hydrate in excess. Explain, with
equations, the chemical changes that will have taken place at
the end of a couple of days.
6. In a flask, A, is heated a mixture of sal ammoniac and potassic
hydrate,' both in solution. The resultant gas is led into another
flask, B, which contains a little water. At the same time some
manganese dioxide and sulphuric acid are added to a third flask, C,
in which has been placed some of the liquid residue from the
preparation of carbon dioxide, by the action of hydrochloric acid
on an excess of calcic carbonate. 1 he flask, C, is also heated, and
the gas formed led into B. Explain the series of chemical changes.
7. 10 grams of sulphuric acid is diluted and added to an excess
of sulphide of iron ; the gas that comes otT is burned in air. How
many grams of air will be required to complete the combustion ?
Find the volume of each product of the combustion at 120° C. and
750 mm. pressure.
8. You are given copper clippings, iron wire, sulphur, sulphuric
acid, saltpetre, water, all necessary apparatus. Mention at least
a dozen compounds that could be formed from them. State how
you would proceed in each case.
9. .Some sodium is thrown upon water, then hydrochloric acid is
added until the solution is neutral to litmus. The liquid is evapor-
ated to dryness and divided into two parts ; both are heated with
sulphuric acid, but to one manganese dioxide has been added as
well. The resultant gases are led into separate solutions of
sodium carbonate. After a few hours, argentic nitrate is added to
both solutions. Explain the series of chemical actions, with
equations.
V. SERIES.
I . You are given (a) 5 white powders and are told that they are
sodium carbonate, sodium nitrate, potassium chloride, zinc sulphate,
potassium oxide.
(d) 5 bottles of liquid and are told that they contain solutions
of chlorine, carbon dioxide, ammonia, hydrochloric acid and
hydrogen sulphide.
204
SELECTED QUESTIONS.
(c) 8 jars of transparent gas containing separately, hydrogen,
nitrous oxide, nitric oxide, carbon monoxide, carbon dioxide,
oxygen, nitrogen and air. I )escribe how you would test each
group in order to determine the substances.
2. A piece of sodium was completely converted into chloride by
uniting with 200 c. c. of CI, at the standard temperature and pres-
sure. What was the weight of the sodium ?
3. Name the chief oxidizing agents with which you have experi-
mented, and explain the theory of the action of each.
4. What experiments have you made that illustrate the direct
replacement of hydrogen by steam ?
5. How many grams of nitric acid containing 67*2 % of pure
HNO3, will neutralize 54*4 grams of ammonia containing 36 % of
NH3?
6. Chlorine was formerly regarded as a compound of hydro-
chloric acid gas with oxygen. Describe experiments, proving that
this was an incorrect view.
7. State in each case the simplest mode of determining when a
receiver is full, in the preparation of ammonia, chlorine, carbon
dioxide, and sulphur dioxide. H«)w would you transfer each of
these gases from one receiver to another ?
8. With some water containing CO.^ in solution, is shaken up a
mi::ture of pure sand and NaCl.
(i) How would you separate these four substances?
(2) How would you prove that you had separated them ?
9. Matter is said to be composed of the elements. Give several
illustrations cf the facts which lead to this theory, and explain why
you consider a mixture of equal volumes of hydrogen and chlorine
gases to be a " mixture " before explosion, and to be replaced by a
" compound " after explosion.
VI. SERIES.
I. A solid substance contains both a carbonate and an easily
dissolved sulphide. How would you prove the presence of these
two bodies ?
SELECTED) QUESTIONS.
206
2. Carbonate of ammonia and nitric acid are given you. How
would you prepare nitrous oxide from these.'* Give equation.
Draw apparatus used.
3. You have given to you some sulphur, water, and nitric acid.
Describe how you would make sulphuric acid from these materials.
4. How could carbon monoxide be shown to contain half its
volume of oxygen .-'
5. Each of two flasks contains some hydrochloric acid, into one is
dropped some iron filings, into the other some manganese dioxide,
then both are heated. The gases that come off are led simul-
taneously through a hot tube, thence into a solution of potassic car-
bonate. Trace the chemical changes throughout.
6. Nitrogen may be prepared by using copper clippings, nitric
acid and air. It may also be got from ammonium nitrate.
Explain the process in each case.
7. If carbon were the element of reference for atomic weights,
find what numbers would express the atomic weights of magnesium,
chlorine, sulphur, arsenic, iron, mercury. Find also the molecular
weights of hydroxyl, carbonic anhydride, nitric acid and ozone.
8. The formula for alcohol is CgHgO. If a lamp burns 10 grams
of alcohol; find the weight of water produced by the combustion,
also how many litres of air would be necessary to furnish the oxygen
required. '
9. Some copper wire is gently warmed with nitric acid, and the
gas that comes off is collected over water. A jar full of this gas is
placed, mouth downward, over calcic hydrate solution, then oxygen
is very slowly forced into it. What chemical changes go on ?
10. Show, from experiment, that (a) change of temperature may
affect the chemical results when the substances act on each other,
(d) that the masses of the two substances may affect the result, {c)
that the degree of concentration of one or both constituents may
alter the substance formed.
11. The following mixtures of gases are given you, and you are
required to separate them, but to preserve the first of each group as a
gas, (a) nitrogen and oxygen, (d) nitrogen and hydrogen, (c) carbon
monoxide and carbon dioxide, (d) hydrogen and sulphur dioxide,
{e) oxygen and hydrochloric acid, ( /) nitrous oxide ana nitric oxide.
How would you proceed in each case .''
a
c
c
t]
1
APPENDIX.
The following books should form a portion of the
reference library of every high school. The publishers'
names are appended, but the latest editions should be
asked for by those who will purchase them. The letters
after the names are the contractions used in the refer-
ences in this book : —
Roscoe and Schorlemmer's Treatise on Chemistry (R. & S.),
Vols. I. and il. (2 parts). Macmillan & Co.
Remsen's Inorganic Chemistry (R.). Advanced series. H. Holt
& Co.
Reynold's Experimental Chemistry. Parts i. to IV. Longmans,
Green & Co.
Muir and Slater's Elementary Chemistry. Cambridge University
Press.
Remsen's Theoretical Chemistiy (Rem. Th. Ch.). Henry C. Lea.
Wurtz Atomic Theory (Wiirtz). International Scientific Series,
Vol. XXX. D. Appleton & Co.
Richter's Inorganic Chemistry. Blakiston, Son & Co.
Chemical Theory for Beginners, by Dobbin and Walker (D. & W.).
Macmillan & Co.
Ramsay's Proofs of Chemical Theory. Macmillan & Co.
Tilden's Introduction to Chemical Philosophy (Tilden). Long-
mans, Green & Co.
Bloxam's Chemistry. Henry C. Lea.
The following is the work in chemistry prescribed for
the years 1896-97-98. An experimental course defined
as follows : — Properties of Hydrogen, Chlorine, Oxygen,
Sulphur, Nitrogen, Carbon, and their more important
compounds. Nomenclature. Laws of Combination of
the Elements. The Atomic Theory and Molecular
Theory,
[2(>7]
208
APPENDIX.
Teachers who are not supplied with gas holders will
find a contrivance similar to that of the accompanying
figure a great convenience.
It consists of a large bottle (a glazed earthenware jar
will answer well) fitted with a perforated rubber stopper,
and a pail with a stopcock at the
bottom. The bottle may be filled
with gas just as any other vessel
would be, then the stopper may
be inserted and the clips closed.
In this way the gas may be stored
until wanted. When the gas is
required it is only necessary to
connect the apparatus as in the
figure, open the clips and allow
the water to flow — the rate may
be regulated by the stopcock.
When it is necessary to draw a
current of gas through any
apparatus, the tube B may be
disconnected from the stopcock,
the positions of the bottle and
pail interchanged, the bottle
filled with water, and a piece of
tubing connected with B long
enough to reach belowthe bottom
of the bottle. As the water is
Fig. 45. siphoucd out at B, air or other
gas will flow in at A. A more convenient arrangement
still is to have the tube B about 4 or 5 feet long, then
to cause a flow one way or the other it is simply neces-
sary to interchange the positions of pail and bottle.
INDEX
14
[209}
^^WPm^^H"
ttl
INDEX.
Abnormal vapour density of sul
phur, 177.
Acetylene, 146.
Acid of carbon, 125.
Acids of chlorine, 163, 172.
" nitrogen, 96.
" sulphur, 182.
Acid, definition of, 57
Acids, bases and salts, 53.
Acid salts, 59.
Air, a mixture, 84.
Air, impurities of, 190.
Alkaline hydrates and chlorine, 161.
Allotropism, 122.
Allotropic modifications of carbon
122.
Allotropic modifications of sulphur
179. *^ '
Ammonia, 105.
Ammonia, tests for, 110.
Ammonia, notes on. 111.
Ammonia, composition of, 114.
Ammonium, 109.
Ammonium hydrate, 109.
Analysis, 68,
Anhydride, 96.
Aqua regia, 164.
Artiad, 65.
Atmosphere, 80.
Atmosphere, composition of, 82.
Atomicity, 63.
Atoms, 15, 193.
Atomic weights, 20.
Atomic theory, 14.
Avogadro's Law, loa
[2
I]
B.
Base, defined, 57.
Bases, acids and salts, 53.
Basicity of acids, 59.
Bivalent, 64.
Black lead, 122.
Bleaching by chlorine, 159.
Bleaching powder, 169.
Blowpipe flame, 154.
Books of reference, 206.
Boyle's Law, 139.
Calcic carbonate, solution of, 129.
Calculation of formulae, 115, 188.
Carbon, 122.
Carbon, reducing power of, 135.
Carbon dioxide, 125.
test for, 125.
composition of, 132.
" notes on, 133.
Carbonic acid gas, 125.
Carbonic acid, ISO.
Carbon monoxide, 134.
Carburetted hydrogen, 143, 145.
Change, physical, 1.
Change, chemical, 2.
Charcoal, 122.
Charles' Law, 139.
Chemical action, theory of, 13.
affinity, 16.
change, 2.
harmonicum, 27.
calculations, 74.
equivalents, 67.
nomenclature, 61.
(I
((
212
INDEX.
<<
(<
Chemical notation, 37.
Chemism, 16.
Chlorate of potash, 170.
Chloric acid, 172.
Chlorine, 156.
bleaching by, 159.
and the alkaline hydrates,
notes on, 161. [161.
" and hydrogen, 162.
" oxides of, 172.
Choke damp, 125.
Classification of chemical actions, 34.
Coal gas, 147.
Combination and mixture, 2.
Combination, simple, 34.
Combustion, 77.
Commercial hydrochloric acid, 164.
Composition of the atmosphere, 82.
" ammonia, 114.
" nitrous and nitric
oxides, 92.
** hydrochloric acid,
" steam, 69. [165.
" water, 22.
" percentage, 119
Compounds of nitrogen and oxygen,
88.
Constituents, volumes of, compared
with volume of compound, 136.
D.
Decomposition, simple, 34.
" and displacement,
35.
" of steam, 32.
Definite proportions, 71.
Density, vapour, 138.
Deodorant, 125.
Diad, 64.
Dibasic acids, 59.
Dioxide of hydrogen, 50.
" nitrogen, 90.
•' carbon, 125.
Disinfectant, 161.
Distillation, destnictive, 123.
Double displacement, 35.
E.
Electricity and chemical change, 17.
Electrode, 13.
Electrolysis, 23.
Electrolytic decomposition, 23.
Electro-positive, 1,% 39.
Electro-negative, 13, 39.
Elements, 17.
list of, 18.
Empirical formulse, 117.
Equations, 39.
Equivalent, 67.
Ethene, 145.
Ethylene, 145.
F.
Filtrate, 3.
Fire damp, 143, 145.
Flame, structure of, 149.
" luminosity of, 148.
" blowpipe, 134.
Force, vital, and chemical change, 11
Formulae, 37.
empirical, 117.
graphic, 120.
rational, 120.
calculation of, 115, 188.
«<
G.
Granulated zinc, 25.
Graphite, 122.
Graphic formulae, 120.
H.
Halogens, 155.
Haloid, 156.
Hardness of water, 191.
Harmonicum, 27.
Heat and chen)ical change, 10.
INDEX.
21.3
7.
Heavy carburetted hydrogen, 145.
Hexad, 64.
Hexevalent, 64.
Hydrate, 56, 57.
Hydrocarbons, 142.
Hydrochloric acid, 163.
" composition of,
165.
" " notes on, 168.
Hydrogen, 23.
" dioxide, 50.
" notes on, 35.
" and chlorine, 162.
" and oxygen, 50.
sulphide, 184.
" notes on, 187.
Hydrosulphuric acid, 184.
Hydroxides, 56, 57.
Hydroxide of ammonium, 109.
Hypochlorous acid, 172.
I.
Ignition, 151.
Impurities in air and water, 189.
Intimate mixture and chemical
change, 8.
L.
Lampblack, 123.
Laughing gas, 90.
Law of Avogadro, 103.
Boyle, 139.
Charles, 139.
Lead, black, 122.
Light carburetted hydrogen, 143.
" and chemical change, 10.
Lime water, 125.
List of atomic weights, 18.
" elements, 18.
Luminosity of flame, 148.
M.
Marsh gas, 143.
Metals and non-metals, 19,
Metathesis, 35.
Methane, 143.
Methane, notes on, 144,
Mixture, defined, 3.
Mixture and combination, 2.
Mixture of gases in air, 84.
Molecular volume, 138.
Molecules, 14, 16, 193.
of elements n.ade up (.f
groups of atoms, 193.
Monad, 64.
Monobasic acids, 59.
Multiple proportions, law of, 95.
N.
Nascent state, 53, 90.
Neutral salts, 59.
Nitrates, notes on, 101.
Nitric acid, 97.
" notes on, 98.
" tests for, 99.
oxide, 90.
Nitrogen, 85.
and oxygen, 87.
" notes on, 87.
" trioxide, 93.
" tetroxide, 94.
" peroxide, 94.
acids, 96.
Nitrous oxide, 88.
Nitrous and nitric oxides, composi-
tion of, 92.
Nomenclature, 61.
Non-metals, 19.
Notation, 57.
Notes on ammonia, 101,
carbon dioxide, 133.
chlorine, 161.
hydrochloric acid, 168.
hydrogen, 35.
methane, 144.
oxygen, 45.
sulphur, 178,
«
214
INDEX.
<<
<t
Notes on sulphuretted hydrogen,
187.
o.
Olefiant gas, 145.
OxideH of carbon, 125.
chlorine, 172.
hydrogen, 50.
nitrogen, 87.
" sulphur, 179.
Oxygen, 41.
" notes on, 45.
" tests for, 45.
Ozone, 48.
P.
Pentad, 64.
Percentage composition, 119.
Perissad, 65.
Peroxide of hydrogen, 50.
" " nitrogen, 94.
Physical change, 1.
Plumbago, 122.
Potable water, 192.
Potassium t)n water, 28.
Potasaic chlorate, 170.
Precipitate, 11.
Proportions, definite, law of, 71.
" multiple, law of, 95.
Q.
Qualitative analysis, 68.
Quantitative analysis, 68.
Quinquivalent, 64.
Questions, selected, 198.
R.
Radicals, 66.
Rational formulae, 120.
Reduction, 33.
Reducing power of carbon, 130.
Reference books, 206.
s.
Safety lamp, 145.
Salts, bases and acids, 55.
" acid and neutral, 59.
Saturated solution, 5.
Selected questions, 198.
Simple decomi)08ition, 34.
" combination, 34.
Sodium on water, 28.
Solution, 4.
Solution of calcic carbonate, 129.
Solvent, 5.
Steam, composition of by voliune, 69.
" decomiMJsition of, 32.
Structure of flame, 149.
Substitutions, 25.
Sulphide of hydrogen, 184.
Sulphur, 175.
" allotropic forms of, 179.
" vapour density of, 177.
" notes on, 178.
" oxides of, 179.
" dioxide, 180.
Sulphurous acid, 182.
Sulphuric acid, 183.
Suljihuretted hydrogen, 184.
Symbols, chemical, 25.
Synthesis, 68.
T.
Table of valency, 66.
Tests for ammonia, 111.
" oxygen, 45.
** nitric acid, 99.
** nitrous acid, 97.
** nitrates, 101.
" carbon dioxide, 125.
" sulphuric acid, 182.
" sulphurous acid, 182.
Tetrad, 64.
Tetravalent, 64.
Tetroxide of nitrogen, 94.
Theory of chemical action, 13.
INDEX.
215
Triad, 64.
Triliasic acids, 59.
Trioxide of nitrogen, 98,
Trivalent, 64.
Univalent, 64.
u.
V.
Valency, 13.
tal)le of, 66.
Vapour density, i;^.
" of sulphur, 177.
<(
Vital force, 11.
Volunifs of con8titnentH and com-
IK)unds, l,%.
Volume of a molecule, 13H.
w.
Water, composition of, 22.
" impuritieH of, 1}>1.
" hardness of, 1!)1.
Weights, atomic, 20.
" list of, 18.
Z.
Zinc, granulated, 25.