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with the oxygen to form water, and the metal is liberated. The
skeleton equation (p. 42) is: Fe 3 O 4 + H > H 2 O + Fe. We then
reason that Fe 3 will give 3Fe. Since all the oxygen is removed
from the compound, O 4 will give 4H 2 O. To produce this, 8H is
required. Hence :

Fe 3 O 4 + 8H -> 4H 2 O + 3Fe.

This interaction is classed as a displacement. In describing it the
chemist would also say that the hydrogen has been oxidized and that
the oxide of the metal has been reduced (pp. 52-53).

Specific Chemical Properties. If the foregoing section on the
chemical properties of hydrogen and the corresponding section under
oxygen (p. 48) are now reexamined, the nature of the facts contained
in, and to be learned from such a section will be seen. Under this
head we describe the chemical behavior of a substance, (1) enumerat-
ing the other substances, simple or compound, with which it unites
or interacts, (2) stating the conditions peculiar to each action, and
(3) estimating the intensity of the tendency to chemical change in
each case. In the case of a simple substance like oxygen we note
particularly with how many of the other elements it can form com-
pounds, how far it unites with them directly, and in how many cases
the compounds have to be made by indirect means. In general,
we call those simple substances active which unite with many other
simple substances and do so by direct union. Oxygen, for example,
is active, and nitrogen (q.v.) is relatively inert.

The Speed of Chemical Actions ; a Means of Measuring Ac-
tivity. One means of measuring the relative chemical activi-
ties of several substances is to observe the speed with which they
undergo the same chemical change (p. 19). Thus we may compare
the activities of the various metals by allowing them separately to
interact with hydrochloric acid and collecting and measuring the
hydrogen liberated per minute by each. It will be seen, even in the
roughest experiment, that magnesium is thus much more active than



76 COLLEGE CHEMISTRY

zinc. The comparison must be made with such precautions, how-
ever, as will make it certain that the conditions under which the
several metals act are all alike. Thus, in spite of the heat evolved
by the action, means must be used, by suitable cooling, to keep the
temperature at some fixed point during the experiment, for all
actions become more rapid when the temperature rises (p. 53).
Again, the pieces of the various metals must be arranged so that
equal surfaces are exposed to the acid in each case. It is found that
the order in which this comparison places the metals is much the
same as that in which they are placed by a study of other similar
actions. A single table suffices, therefore, for all purposes (see
Electromotive series, also footnote to p. 24).

Exercises. 1. What are the valences of the negative radicals
of phosphoric acid (p. 51), and of acetic acid (p. 64)? What must
be the formulae of calcium phosphate, cupric acetate, aluminium
phosphate, ferrous carbonate, ferrous sulphate, cupric chloride?

2. What is the valence of phosphorus in phosphoric anhydride
(p. 51)? What must be the formulae of, (a) the corresponding
chloride and sulphide of phosphorus, and (6) of aluminium oxide?

3. What are the valences of the elements in the following: LiH,
NH 3 , SeH 2 , BN?

4. What are the valences of the metals and radicals in the follow-
ing: Pb(NO 3 ) 2 , Ce(SO 4 ) 2 , KC1, KMnO 4 (potassium permanganate)?
Name all the substances in 3 and 4.

5. Write the formulae of ferrous and ferric oxides, of ferrous and
ferric nitrates, of stannous and stannic sulphides.

6. Make equations to represent, (a) the reduction of lead dioxide
(Pb0 2 ) by hydrogen, (b) the actions of aluminium upon cold water
and, (c) upon steam at a red heat.



CHAPTER VIII
WATER

THE great quantity of water which occurs in nature makes it one
of the most familiar chemical substances. Water is found also in the
bodies of both animals and plants in large amounts, and is indeed
essential to the working of living organisms.

Natural Waters. Sea-water holds about 3.6 per cent of solid
matter in solution. Rain-water is the purest natural water, but
contains nitrogen, oxygen, and carbon dioxide dissolved from the
air. Well-waters often contain calcium sulphate, calcium bicarbo-
nate, and compounds of magnesium in solution, and are then described
as hard. Other waters contain compounds of iron, and still others
are effervescent and give off carbon dioxide. These are called min-
eral waters. All of the dissolved substances are obtained by the
water in its progress over or under the surface of the ground.

Water which is to be used for domestic purposes is examined, par-
ticularly, for organic matter. This usually gains access to the water
by admixture of sewage (p. 52). It is not this organic matter
itself which is deleterious, but the bacteria of putrefaction and disease
and their products (ptomaines) which are likely to accompany it.

Purification of Water. The foreign materials which water may
contain are divisible into two kinds, dissolved matter and sus-
pended matter. No water is free from either of these varieties of
impurity. In chemical laboratories distilled water is always em-
ployed. Yet it is difficult to keep the liquid pure even for a short
time. Ordinary glass dissolves in water to a very noticeable extent.

For ordinary purposes the suspended matter which water contains
is removed by nitration (p. 7). In the laboratory this takes place
through unsized paper. The pores of the paper are sufficiently small
to retain suspended particles, while permitting the passage of the
water with its dissolved matter. On a large scale, beds of gravel are

77






78 COLLEGE CHEMISTRY

employed. In the household, the Pasteur filter is more compact and
efficient. The water is forced by its own pressure through the pores
of a closed tube made of unglazed porcelain. Care must be taken
to clean these tubes at frequent intervals, so that organic and per-
haps putrescent matters may not accumulate upon them.

Matter in solution cannot be removed by ordinary filtration, and
is eliminated by distillation (p. 26). Since the water is converted
into steam and is condensed in platinum or tin pipes, only gases or
volatile liquids dissolved in it can pass into the distillate.

Physical Properties of Water. A deep layer of water has a
blue or greenish-blue color. At a pressure of 760 mm., water ex-
ists as a liquid between and 100. Below it becomes solid,
above 100 a gas. Of all chemical substances it is the one which we
use most, so that its physical properties, discussed below, should
be studied attentively. Then too, what is said of water is in gen-
eral true of all other liquids, from which it differs only in details.

Ice. The raising or lowering of the temperature of a gram of
water through one degree involves the addition or removal of one
calorie of heat. The conversion, however, of a gram of water at
to a gram of ice at requires the removal of 79 calories. The
mere melting of a gram of ice causes an absorption of heat to the
same amount, called the heat of fusion of ice. At a mixture of
ice and water will remain in unchanged proportions indefinitely.
Any cause which tends permanently to lower or raise the temperature
by a fraction of a degree, however, will bring about the disappearance
of the water or of the ice respectively. This temperature is called
the melting or the freezing point. A temperature, like this, at which
a substance passes from one physical state of aggregation to another is
called a transition point.

Steam and Aqueous Tension. At atmospheric pressure, water
passes into steam rapidly at 100, but at lower temperatures, and
even when frozen, it does the same thing more slowly. The quantity
of the vapor present is defined by the gaseous pressure it exercises,
the value being called the vapor pressure of water vapor (or of the
vapor of any other volatile substance) in the location in question.

The most significant fact about vapor pressure is that, when excess
of the liquid is present, the pressure of the vapor quickly reaches a



WATER



79



definite maximum value for each temperature. In the absence of
excess of the water, less than this maximum pressure may exist.
More than the maximum pressure proper to a given temperature, if
produced by compression, cannot be maintained, for a part of the
vapor condenses to the liquid state. The magnitude of this maxi-
mum vapor pressure, at a given temperature, depends on the ability
of the particular liquid to generate vapor. This maximum vapor
pressure is held, therefore, to represent the vapor tension of the liquid,
at the given temperature, and this is a specific property of the
substance.

I/ The vapor tension may be shown by allowing a few drops of water
to ascend into a barometric vacuum (Fig. 25). The tube on the left
shows the mercury when nothing presses on its
surface. The tube on the right shows the result
of admitting the water. The difference in the
height of the two columns gives the value of
the vapor pressure of the water vapor. With
excess of water, the value is that of the vapor
tension, called, in the case of water, the aqueous
tension. The jacket surrounding the tube on the
right enables us, by adding ice or warm water,
to maintain any temperature between and
100. When ice is used outside, and a piece of
it is introduced into the vacuum, the vapor it
gives off quickly reaches a pressure of 4.5 mm.
The vapor pressure of the ice takes the place of
4.5 mm. of mercury in balancing the atmospheric
pressure, and so the mercury column falls by this
amount. Similarly, water at 10 causes a fall of
9.1 mm. and at 20 of 17.4 mm., so that these
represent the mercury-height values of the
aqueous tension at these temperatures. The
quantity of water used makes no difference, so
long as a little more is present than is required
to fill the available space with vapor. With
ether, instead of water, at 10 the fall is 28.7 mm.
With water at higher temperatures the fall of
the mercury column becomes much greater. At 50 it is 92 mm.,
at 70 it is 233.3 mm., at 90 it is 525.5 mm., and at 100 it is 760




FIG. 25.



80 COLLEGE CHEMISTRY

mm., or one atmosphere. At 121 the aqueous tension is two
atmospheres, at 180 it is ten atmospheres.

When water at a certain temperature has given the full amount of
water vapor to the space above it that its aqueous tension permits,
we say that the space is saturated with vapor. That concentration of
vapor which constitutes saturation varies with the temperature of
the water and depends therefore solely on the power of the water to
give off vapor. It has nothing to do with the size of the space, and
is even independent of other gases the space may already contain
(p. 60, also pp. 90-91. See footnote to p. 24).

The space immediately above the surface of the ground, which is
mainly occupied by atmospheric air, is, on an average, less than two-
thirds saturated with water vapor. That is to say, such air, when
inclosed in a vessel containing water, will take up about one-half
more than it already contains. The vapor of water at 100 in
an open vessel displaces the air entirely, and, if the required heat is
furnished, the liquid boils. This temperature, like the freezing-point,
is a transition point.

A gram of water at 100, in turning into a gram of steam at 100,
takes up 537 calories. This is called its heat of vaporization. Steam,
in fact, contains much more internal energy than an equal weight
of water at the same temperature, just as water, in turn, contains
more energy than ice.

All our substances and apparatus have traces of water, derived
from the atmosphere, condensed on their surfaces. This water is,
in a sense, in an abnormal condition, for it does not evaporate even
in dry air. It is observed to pass off in vapor, however, when we
have occasion to heat the substance or apparatus.

Water as a Solvent. One of those physical properties of water
which are most used in chemical work is its tendency to dissolve
many substances. This subject is so important and extensive that
we shall presently devote a complete chapter to some of its simpler
and more familiar aspects.

Chemical Properties of Water. Water is so very frequently
used in chemical experiments in which it is a mere mechanical ad-
junct, that the beginner has difficulty in distinguishing the cases in
which it has itself taken part in the chemical interaction. The four



WATER 81

kinds of chemical activity which it shows should therefore receive
careful notice:

1. Water is a relatively stable substance.

2. It combines with many oxides, forming bases or acids.

3. It combines with many substances, chiefly salts, forming
hydrates.

4. It interacts with some substances in a way described as hydro-
lysis. This property will not be discussed until a characteristic case
is encountered.

Water a Stable Compound: Dissociation. In the case of a
compound, the first chemical property to be given is always, whether
the substance is relatively stable or unstable. Usually the specifica-
tion is in terms of the temperature required to produce noticeable
decomposition. Thus, potassium chlorate gives off oxygen at a low
red heat. Now, water vapor, when heated, is progressively decom-
posed into hydrogen and oxygen, yet at 2000 the decomposition
reaches only 1.8 per cent, and reunion occurs gradually as the tem-
perature is lowered.

A decomposition which thus proceeds at higher temperatures,
while at lower temperatures combination of the constituents can
take place, is called a dissociation. The decomposition of potassium
chlorate (p. 47) is not a dissociation because it is not reversible;
oxygen will not under any circumstances reunite with potassium
chloride.

Union of Water with Oxides. When sodium combines with
oxygen under certain conditions we obtain sodium oxide (Na^O).
The product unites violently with water to form sodium hydroxide:

Na 2 O + H 2 O -+ 2NaOlT.

The slaking of quicklime is a more familiar action of the same kind :
CaO + H 2 - Ca(OH) 2 .

No other products are formed. The clouds of steam produced in the
second instance are due to evaporation of a part of the water by the
heat produced in the formation of calcium hydroxide. The aqueous
solutions of these two products have a soapy feeling, and turn red
litmus blue, and the substances therefore belong to the class of
alkalies (pp. 66, 51) or bases. Very many hydroxides which are of



82



COLLEGE CHEMISTRY



the same nature, for example ferric hydroxide Fe(OH) 3 and tin
hydroxide Sn(OH) 2 , are formed so slowly by direct union of the
oxide and water that they are always prepared in other ways.

Some oxides, although they unite with water, give products of an
entirely different character. Phosphoric anhydride and sulphur
dioxide are of this class and, as we have seen (p. 51), yield acids.

These two classes of final products are so different that we make
the distinction the basis for classification of the elements present in
the original oxides. The elements, like sodium and iron, whose
oxides give bases, are called metallic elements ; those, like phosphorus,
whose oxides give acids, are called non-metallic elements. The dis-
tinguishing words are selected because the division corresponds, in a
general way at least, with the separation into two sets to which
merely physical examination of the elementary substances would lead.

Hydrates. Many substances when dissolved in water and
recovered by spontaneous evaporation of the solvent enter into
combination with the liquid. The products,
which are solids, are called hydrates. That
they are regular chemical compounds is
shown by the following three facts: (1)
These compounds show definite chemical
composition expressible by formula in
terms of chemical unit weights of the
constituents. (2) Often much heat is
given out in their formation. Thus, in the
case of washing soda, the decahydrate of
sodium carbonate (Na 2 CO 3 , 10H 2 0), the heat of the union
(p. 55) is 8800 cal. (3) The hydrates have physical properties
entirely different from those of their components. Thus, cupric
sulphate, often called anhydrous cupric sulphate to distinguish
it from the compound with water, is a white substance crystal-
lizing in shining, colorless, needle-like prisms. The pentahydrate
(blue-stone or blue vitriol), which crystallizes from the aqueous solu-
tion, is blue in color, and forms larger but much less symmetrical
(asymmetric or triclinic) crystals (Fig. 26) :

CuSO 4 + 5H 2 O <= CuS0 4 , 5H 2 0.

The chemical properties show hydrates to be relatively unstable.
When heated, the hydrates, as a rule, lose none of the constituents of




FIG. 26.



WATER 83

the original compound, but only the water, in the form of vapor.
When melted, or when dissolved in water, the hydrates are disso-
ciated (p. 81) into water and the original substance. The aqueous
solutions made from the anhydrous substances and from the hydrates
have identical physical and chemical properties. Hence the cheaper
of the two forms is generally purchased, and many of the chemicals
used in the laboratory are in the form of hydrates. In consequence
of the ease with which hydrates give up water we write their for-
mulae (e.g. CuSO 4 , 5H 2 O) so that the water and original substance
are separate. A formula thus modified, so as to show some favorite
mode of behavior of the substance, is called a reaction (p. 11) for-
mula. The formula H lo CuS0 9 , which would be equally correct, is
never employed, because its use would disguise the relation of the
substance to cupric sulphate.

The less stable hydrates dissociate very readily. Thus the deca-
hydrate of sodium sulphate, Na 2 SO 4 ,10H 2 O (Glauber's salt), loses
all the water it contains (effloresces) when simply kept in an open
vessel. When kept in a dosed bottle, a very little of it loses water,
and then the decomposition ceases. The cause of this we discover
when a crystal of the hydrate is placed above mercury, like the ice
or water in Fig. 25 (p. 79). It shows a definite aqueous tension. At
9 the value of this is 5.5 mm. As its temperature is raised, the
tension increases. When the temperature is lowered, on the other
hand, the tension diminishes, the mercury rises, and a part of the
water enters into combination again. Different hydrates show
different aqueous tensions at the same temperature. For example, at
30, that of water itself is 31.5 mm., strontium chloride (SrCl 2 , 6H 2 O)
11.5 mm., cupric sulphate (CuSO 4 , 5H 2 0) 12.5 mm., barium chloride
(BaCl 2 , 2H 2 O) 4 mm.

In view of these facts, we perceive that loss of water by efflores-
cence is like evaporation. Those hydrates which, like Glauber's
salt and washing soda, have a vapor tension approaching that of
water itself, lose their water at ordinary temperatures at a rapid
pace. Now, atmospheric air is usually less than two-thirds saturated
with water vapor, and the partial pressure (p. 60) of this vapor
opposes the dissociation and tends to prevent the liberation of the
water. Thus at 9, the vapor tension of water being 8.6 mm., the
average vapor pressure of water in the atmosphere will be about
5 mm. Any hydrate with a greater aqueous tension than 5 mm.,



84 COLLEGE CHEMISTRY

at 9, such as Glauber's salt, will therefore decompose spontaneously
in an open vessel. But those with a lower vapor tension, such as the
pentahydrate of cupric sulphate with a tension of 2 mm. at 9, will
not do so (see pp. 90-93).

The water of hydration is known colloquially in chemistry as
water of crystallization. The term was introduced when it was first
observed that a hydrate, in decomposing, crumbles and loses its
original crystalline form. But the term is misleading. All pure
chemical substances, in solid form, when in stable physical condi-
tion, are crystalline. Amorphous (i.e., non-crystalline) substances,
like wax and glass, are supercooled liquids.

Composition of Water. The proportion of hydrogen to oxy-
gen, in water, by weight, is 2 : 15.879 or 2.015 : 16. The proportion
by volume is 2.0027 volumes of hydrogen to 1 volume of oxygen.
That the proportion by volume is very close to 2 : 1 may easily
be shown by mixing hydrogen and oxygen in this proportion, in a
strong tube, and exploding the mixture by means of a spark from
an induction coil. The resulting steam condenses and the whole
gas vanishes. If different proportions are used, the excess of one
of the gases remains uncombined.

Gay-Lussac's Law of Combining Volumes. The almost
mathematical exactness with which small integers express this pro-
portion is not a mere coincidence. Whenever gases unite, or gaseous
products are formed, the proportions by volume (measured at the same
temperature and pressure) of all the gaseous bodies concerned can be
represented very accurately by ratios of small integers. This is called
Gay-Lussac's law of combining volumes (1808). Thus, when the above
experiment is carried out at 100, in order that the product, water,
may be gaseous also, it is found that the three volumes of the con-
stituents give almost exactly two volumes of steam. For example,
15 c.c. of hydrogen and 7.5 c.c. of oxygen give 15 c.c. of steam. Of
course the hydrogen, oxygen, and steam must be measured at the
same pressure, and the temperature must remain constant (100)
during the experiment. Proper manipulation secures the former,
and a jacket filled with steam (Fig. 27) the latter condition. Strips
of paper, 1, 2, and 3, are pasted on the jacket in such a way that
equal lengths of the eudiometer, in this case a straight one, are laid
off. The three divisions being filled with a mixture of hydrogen



WATER



85



and oxygen in the proper proportions, the gas, after the explosion,
shrinks so as to occupy, at the same pressure, only two of them.
From this universal truth in regard to the combination of gases,
we draw the important inference that the chemical unit-weights of
simple substances, and the formula- weights
of compounds, in the gaseous condition,
occupy at the same temperature and pressure
volumes which are equal, or else stand to
one another in the ratio of small integers.

Exercises. 1. Name some familiar
transitions (p. 78) from one physical
state to another.

2. What evidence is there in the com-
mon behavior of ether and chloroform to
show that these liquids have high vapor
tensions?

3. If the pressure of the steam in a
boiler is ten atmospheres, at what tem-
perature is the water boiling (p. 80) ?

/ 4. How many grams of water could be
heated from 20 to 100 by the heat
required to melt 1 kgm. of ice at 0?

5. What do you infer from the fact that
alum and washing soda lose their water
of hydration when left in open vessels,
while gypsum does not?

6. Which fact shows most conclusively
that hydrates are true chemical compounds?

7. Gypsum is a hydrate of calcium sulphate (CaS0 4 ). If 6 g.
of gypsum, when heated, lose 1.256 g. of water, what is the formula
of the hydrate?

8. In what ways does a hydrate differ from, (a) a solution, (6) an
hydroxide?

9. Should you expect to find any difference, in respect to chemical
activity, between the three forms of water? Have we had any
experimental confirmation, or the reverse, of this conclusion (p. 66) ?

10. Name some crystalline substances which are not used, or do
not occur in the form of hydrates.




FIG. 27.



CHAPTER IX
THE KINETIC-MOLECULAR HYPOTHESIS

As soon as we have constructed a law (p. 4) we desire immediately
to find out the basis of the constant mode of behavior it epitomizes.
If no explanation, that is, more detailed description, is forthcoming
as the result of closer observation, we proceed to imagine one (p. 5).
This always takes a mechanical form, often crude at first, and later
undergoing refinement. Thus, at first, the phenomena of light were
explained by the conception of clouds of fine corpuscles emanating
from the luminous body. The chances of hitting upon an objective
reality by guess-work like this is obviously remote. Whether such
particles did really fly about was not the main question, however.
Their value lay in the fact that they could be pictured concretely and
gave a basis for further thought and perhaps suggestions for new
experiments. Such a structure of the imagination is called an
hypothesis.

Tlie Molecular Hypothesis. The only mechanical basis we can
imagine to account for the physical properties of matter is a discon-
tinuous structure of some description. The fact that all kinds of



Online LibraryAlexander SmithGeneral chemistry for colleges → online text (page 8 of 47)