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our knowledge of the physical sciences is increased, the more we
realize their interrelation and their interdependence. The study
of no particular one can be successfully pursued if we exclude the
help afforded and the side lights thrown upon it by others. This
is notably so in the case of electricity. For a proper understanding
of the present accepted theory accounting for the phenomena of
voltaic electricity, we must turn to physical chemistry and to
develop our explanation must begin with certain facts which at
first sight appear to have not even a remote connection with our
announced subject.

The following outline will assist the student in following the
thread of connection between the facts which will now be brought

1. Avogadro's law and a derived corollary applicable to gase-
ous pressure are explained (Par. 256).

2. Exceptions to the law of gaseous pressure are shown to be
due to dissociation which is defined (Pars. 257-258).

3. Osmotic pressure is described and its observation and meas-
urement explained (Pars. 259-262).

4. Osmotic pressure is shown to follow the laws of gaseous
pressure (Pars. 263-266).

5. Abnormal osmotic pressures are, like excessive gaseous
pressures, shown to be capable of explanation under the supposi-
tion of dissociation, otherwise called ionization (Pars. 267-268).

6. Ionization is further explained (Pars. 269-274) .

7. Electrolytic properties are shown to depend upon ionization
(Pars. 275-279).

8. Electricity is shown to be atomic in character (Par. 280).

256. Laws of Variation of Gaseous Pressure. Avogadro's
Law, of fundamental importance in Chemistry, is to the effect
that under like conditions of temperature and pressure, equal


volumes of all gases, simple or compound, contain the same num-
ber of molecules. If we should have a series of cylinders of exactly
the same capacity and should fill one with oxygen, one with hydro-
gen, one with carbon dioxide, one with marsh gas, and so on, each
being at the same temperature and exposed to the same pressure,
then each would contain exactly the same number of molecules.

Suppose one of these cylinders of the same diameter as the
others should be twice as tall. If this one be filled with gas it will,
from the above, contain twice as many molecules as the others.
Place a piston in the mouth of this cylinder and press it down until
the volume of the enclosed gas be reduced one-half, that is, until
it becomes the same as that of the other cylinders. The space
beneath the piston now contains twice as many molecules as the
other cylinders contain. From Mariotte's Law, temperature
remaining constant, the volume of a gas varies inversely as the
pressure. The pressure upon the compressed cylinder is therefore
twice that upon the others. Hence we may state, as a corollary
to Avogadro's law, that for a constant temperature and volume, the
pressure of a gas varies directly as the number of molecules enclosed.

From a combination of Charles' and Mariotte's Laws it is shown
that for constant volume, the pressure produced by an enclosed
gas varies as the absolute temperature. (The absolute tempera-
ture is obtained by adding the constant 273 to the temperature
as indicated on the Centigrade scale.) We therefore see that the
pressure of a gas confined f in a given volume varies (a) with the
number of molecules enclosed and, (b) with the absolute tempera-

257. Decomposition and Dissociation. In general, compound
substances if heated to a sufficiently high temperature are resolved
into simpler ones. If when these simpler substances are cooled
to the primary temperature they remain separate, the original
compound body is said to have undergone decomposition. On the
other hand, if when the temperature falls the simpler substances
recombine and reproduce the original substance, this body is said
to have undergone dissociation. Decomposition is therefore
permanent while dissociation is transient and continues only so
long as the agency which brought it about is operative.

258. Example of Dissociation by Heat. Ammonium chloride,
NH 4 C1, like other ammonium salts, is volatilized with compar-


ative ease. Its molecular weight being 53.5, the gas produced by
the volatilization of 53.5 grams should exert the same pressure
as that produced by a molugram of any other gas confined in an
equal volume and at the same temperature. (A molugram is the
molecular weight expressed in grams, as for example 2 grams of
hydrogen, 28 grams of nitrogen, 44 grams of carbon dioxide, and
so on.) By actual experiment however the pressure is found to be
twice as great. From (a) Par. 256 therefore, there must be twice
as many molecules present in the gaseous NH 4 C1 as there are in
the other gases. The explanation is that the NH 4 C1 has been
dissociated by the heat, each molecule becoming two, one of
ammonia, NH 3 , the other of hydrochloric acid, HC1. That this
is so may be proven in several ways. First, if NH 4 C1 became a
gas without dissociation, the specific gravity of this gas referred
to hydrogen should be 26.7 while it is actually only 13.35 which
is the specific gravity of a mixture of equal volumes of NH 3 and
HC1. Second, the specific gravity of HC1 being 18.2 while that
of NH 3 is only 8.5, if the dissociation takes place in a vertical
closed tube, the heavier HC1 will settle at the bottom, the lighter
NH 3 rising to the top. If by means of a stop cock at the middle
of the tube the two halves be now cut apart and after cooling be
tested separately, the contents of the upper half will be found to
be alkaline, that of the lower half acid.

259. Osmosis and Osmotic Pressure. Suppose the space
below the piston of a vertical cylinder to be filled with a gas under
normal pressure. If the piston be raised, thereby increasing the
space beneath it, the gas will be found to have spread through this
new space completely filling it. There is therefore a force or
pressure which compels a volume of gas to diffuse or to swell
out and occupy a greater space when it has the opportunity to
do so.

Again, if in the bottom of a vessel there be placed a concen-
trated solution of a salt and if then there be poured carefully on
top of this solution a layer of pure water, in a short while the dis-
solved salt, in defiance of gravity, will have spread upward and
throughout the liquid until the latter is all of a uniform density.
By using a colored salt the progress of the diffusion can be easily
observed. There is therefore a force, similar to the gaseous pres-
sure described above, which urges the particles of a dissolved
substance to spread equally throughout the solvent.



There are known various membranes, some animal, some
vegetable, and some artificial, which will permit the passage
through them of certain liquids but will prevent the passage of
other substances dissolved in these liquids. On account of this
property these membranes are called semi-permeable. If a bladder
(which is one of these) be filled with an aqueous solution of a salt,
tied tightly, and then submerged in a vessel of pure water, it will
gradually distend and may finally burst. This is explained by
saying that the substance in solution is urged by the pressure
described above to spread out into the surrounding solvent but
being unable to pass through the membrane it pushes against it
and distends it, thus allowing the water on the outside to enter.
Although this explanation is admittedly a poor one, the phenom-
enon does occur and is called osmosis, and the force exerted
upon the membrane by the dissolved molecules is called osmotic

In the above illustration we have assumed an aqueous solu-
tion of a salt but under proper conditions osmosis takes place
whatever the nature of the solvent or of the dissolved sub-

260. Demonstration of Osmotic Pressure. Osmotic pressure
may be conveniently shown as follows. A membrane is stretched
and tied over the mouth of a glass funnel which is then inverted
and filled to the neck with a solution, say
of copper sulphate. The inverted funnel
is then inserted, as shown in Fig. 117, in a
vessel of pure water until the surface of
the water and that of the liquid in the
neck of the funnel are at the same level.
The copper sulphate solution will be ob-
served to rise slowly in the neck of the
funnel and may continue to do so for
several weeks, attaining its maximum
height when the hydrostatic pressure of
the liquid in the tube just prevents the
passage of additional water through the
membrane and the further dilution of the Flg - 117 -

contained solution. The osmotic pressure and the hydrostatic
pressure are now in balance and by measuring the latter we
determine the former.


261. Measurement of Osmotic Pressure. The arrangement
described in the preceding paragraph is not well suited for the
measurement of osmotic pressures. These are relatively great,
The osmotic pressure produced by a dilute solution of sugar has
driven a column of water to a height of nearly 70 feet, and this
pressure is frequently exceeded. The membrane would not stand
these pressures and it is impracticable to use tubes of such length.
Again, the membrane is not absolutely impermeable to the salt
and some escapes into the surrounding solvent. Also, the mem-
brane is distended, thereby increasing the volume of the confined
solution and materially altering the degree of concentration. For
these reasons accurate determinations of osmotic pressure were
not made until within recent years when it was discovered that
certain colloidal or gelatinous precipitates, notably the ferro-
cyanide of copper, act, so far as permeability is concerned, as
ideal membranes. The strength of a film of such a precipitate is
however very small and in order that it may withstand the pres-
sure to which it is to be subjected it must be supported in some
way. This object is now attained by depositing the film within the
substance of a finely porous unglazed porcelain cup. These cups,
about two inches tall and three-quarters of an inch in diameter,
are first filled with a solution of potassium ferrocyanide which
slowly soaks into the walls. They are then immersed in a solution
of copper sulphate, which soaks in from the outside, and when the

.two liquids encounter each other the precipitate is formed. The
actual process requires several days' time and involves a number
of precautions not necessary to mention here. Into the mouth of
the prepared cup are cemented the tube up which the liquid is to
rise and a second tube with a stop cock by which the solution is
introduced. The vertical or pressure tube is sealed at the top and
the osmotic pressure may be determined by the amount of com-
pression of the air above the liquid. In practice, the pressure tube
is a mercurial manometer. By using these closed tubes to measure
the pressure, the amount of the solvent which enters the cup is
reduced to a minimum and the concentration of the solution is
altered but little.

262. Observations of Pfeffer. The botanist, Pfeffer, in his
investigations in plant physiology, made, with the apparatus just
described, a series of observations upon the osmotic pressure pro-
duced under various conditions by dilute solutions of organic


compounds such as sugars, alcohols, etc. His results were pub-
lished in 1877 but at that time attracted no special attention and
it remained for Arrhenius and Van't Hoff to discover some ten
years later the value of his data and its bearing upon the theory
which we shall shortly explain.

263. Osmotic Pressure Varies Directly with Number of Mole-
cules Dissolved in Given Volume of Solution. Pfeffer found
that for these dilute solutions the osmotic pressure increased
directly with the strength of the solution, that is, if the concen-
tration (and hence the number of molecules in solution) be
doubled, the osmotic pressure is likewise doubled, etc. His
results for cane sugar were as follows:

Strength of Osmotic RotiVt ^

Solution Pressure * vauo g

1% 510 mm. 510

2% 1016 mm. 508

4% 2082 mm. 520

6% 3075 mm. 512

In this table, while the pressures do not bear to each other the
exact theoretical ratio, the variations therefrom are not greater
than are to be expected from experimental errors and from the
fact that the observations were not taken under precisely the same
conditions of temperature, although they were made within a
range of less than three degrees Centigrade.

Comparing different substances, he found that while the osmotic
pressure of a one per cent solution of cane sugar at 15.5 C was
520.5 mm., that of a one per cent solution of raffinose at the same
temperature was only 299 mm. The relation between these two
numbers was not discovered until subsequent investigators
worked upon his data. The formula for cane sugar is C^H^On
and its molecular weight is 342 ; that for raffinose is Ci8H 32 Oi65H 2
and its molecular weight is 594. Therefore, equal weights of the
two do not contain the same number of molecules, a one per cent
solution of raffinose containing fewer than a one per cent solution
of sugar. Let us see how the pressures compare if we take the
same number of molecules of each. Each litre of his cane sugar
solution contained ^ f a molugram. The same fraction of a
molugram of raffinose would be ^W of 594 or 17.37 grams. If
10 grams produced a pressure of 299 mm., what pressure would
17.37 produce? 10 : 299 : : 17.37 : x


whence x = 519.4 mm. as compared to the 520.5 mm. of the sugar

This and other examples show that substances in solution con-
form to Avogadro's Law and to its corollaries, that is, equal
volumes of solutions which at the same temperature exhibit the
same osmotic pressure contain the same number of dissolved
molecules, and also, other conditions being constant, the osmotic
pressure varies directly with the number of molecules in solution.

264. Osmotic Pressure Follows Mariotte's Law. An exami-
nation of Pfeffer's data will reveal the fact that osmotic pressure
also follows the corollary to Mariotte's Law for gaseous pressure,
that is, other conditions being constant the osmotic pressure
varies directly with the absolute temperature. For example, the
osmotic pressure of a one per cent solution of sugar at 14.15 C is
510 mm. and at 32 C is 544 mm. Applying the law to the lower


510 : z=273+14.15 : 273+32

whence x = 541.7 mm.,
agreeing closely with the observed pressure 544 mm.

265. Osmotic Pressure of a Molecule in Solution Equals
Pressure of a Gaseous Molecule under Equal Volume and Tem-
perature. We have seen from Par. 263 above that -& of a molu-
gram of sugar or of other organic substance dissolved in a litre of
water exerts at 15.5 C an osmotic pressure of about 520 mm. Let
us see what pressure the same fraction of a molugram (and hence
the same number of molecules) of hydrogen confined in the same
space and at the same temperature would exert. One gram of
hydrogen at C and 760 mm. measures 11.165 litres, therefore, a
molugram of hydrogen (2 grams), would under these conditions
occupy 22.33 litres, and ^W of a molugram would occupy .6529
litre. At a temperature of 15.5 C this would dilate to .6914
litre and if this be allowed to expand into the space of 1.006 litres
(the space occupied by one litre of water to which 10 grams of
sugar is added), the pressure would drop according to the propor-

760 : x = 1.006 : .6914

whence x = 522 .4 mm.

We see then that the osmotic pressure of the sugar in solution
is the same as the pressure exerted by an equal number of mole-


cules of gas confined in the same space and at the same tem-

266. Van't HofTs Generalization. A consideration of the fore-
going facts led to the generalization by Van't Hoff which is to the
effect that "the osmotic pressure of a substance in solution is the
same as the gas pressure which would be observed if the dissolved
substance alone, in gaseous state and at the same temperature,
occupied the volume of the solution." In other words, these sub-
stances in solution behave, comparing osmotic pressure to gaseous
pressure, precisely as if they had been converted into a gas and
filled alone the space occupied by the solution.

Independent theoretical considerations based upon the lowering
of the freezing point and the raising of the boiling point by sub-
stances in solution lead to the same conclusions and entirely
corroborate Van't Hoff.

267. Exceptions to Van't Hoflf's Generalization. Van't HofFs
generalization applies, as we have seen, to dilute solutions of
organic compounds. If the solutions become concentrated, the
laws of osmotic pressure no longer hold strictly. This is thought
to be parallel to the failure of gases, as they approach their point
of condensation, to follow the laws of Charles and Boyle.

If, now, we turn our attention to solutions of the inorganic
compounds we find that the majority of them are exceptions and
give osmotic pressures in excess of those required by theory.
These exceptions embrace solutions of all the acids, all the bases
and all the salts. It might seem therefore that in announcing as
general a law to which the exceptions outnumber the agreements,
Van't Hoff had overstepped the bounds of prudence.

268. Dissociation Theory of Arrhenius. In Par. 263 above
we saw that osmotic pressure varied directly with the number of
molecules in solution. Since in the exceptional cases the pressure
is always greater than what it should be in theory, there must be
a greater number of molecules present in solution than is indicated
by the weights taken. To account for this greater number,
Arrhenius advanced the theory that just as the excessive pressure
produced by iodine, ammonium chloride, etc., when converted
into vapor is explained by the fact that these substances are dis-
sociated by the heat employed, so the excessive osmotic pressures
are to be explained by the fact that the substances in solution


undergo dissociation, or ionization, that is, split up into a greater
number of parts. It is also a part of his theory that these part
molecules or ions, whether they be atoms or compound radicles,
exert the same osmotic pressure as an undissociated molecule.
Some of the consequences following from this theory were so
startling and so contrary to the views generally held by chemists
that it was at first vigorously combated and reluctantly accepted
as one by one the objections advanced against it were explained
away. A full exposition of these consequences and replies to the
objections would require an extended treatise. We can here do
but little more than allude to a few of those most obviously con-
nected with our subject.

269. Why Ionization Takes Place in Solution. Salts, acids
and bases consist of two parts, a metal or hydrogen (or a radicle
playing a similar part) combined with an acid radicle or, in the
case of the base, with hydroxyl. The metal or hydrogen portion
carries a positive charge of electricity; the remaining radicle
carries an equal negative charge. These two parts may therefore
be regarded as held together by the attraction of these opposite
charges. The charges being relatively great (Par. 278) and the
distance separating the parts being infinitely small, the attraction
is very great (Par. 53). In Par. 90 we saw that if two charged
bodies which in air attract or repel each other with a certain force
were placed in some other medium whose dielectric coefficient is
K, then the force exerted between the two bodies would be only
^th of what it was in air. The dielectric coefficient of water is
given in Par. 92 as 80, or with the exceptions of hydrogen peroxide
and hydrocyanic acid, greater than that yet determined for any
other substance. The force which held the ions together is there-
fore reduced to ^th of itself when the substance is brought into
solution, and the ions drift apart. This view is corroborated by
the variation in dissociation produced by using solvents of dif-
ferent dielectric coefficients.

270. How Ionization Takes Place. Ionization takes place
differently from the dissociation by heat. The metallic salts
split into the metal and the acid radicle; the acids split into hydro-
gen and the acid radicle; the bases split into the metal and the
hydroxyl radicle. Now such radicles as NH 4 , OH, S0 4 , etc., which
this requires, are unknown as separate entities. The ionization of


KC1 supposes the presence in the water of atoms of potassium
and of chlorine. If this be so, some of the chlorine should reveal
itself by its color and odor. Further, it is well known that potas-
sium placed upon water decomposes it with such violence as to
produce flame and forms potassium hydroxide. None of these
effects are produced and this was once regarded as a grave objec-
tion to the theory. This objection is answered by the statement
that ions carrying electrical charges differ from those that do not.
A metallic ion can go into solution only when it has a positive
charge, and once in solution it can not be withdrawn until this
.charge is removed or neutralized. This can be shown experi-
mentally thus. A plate of zinc dipped into hydrochloric acid is
attacked vigorously and goes into solution. If, however, this plate
be charged negatively, the action of the acid immediately ceases.
So long as the potassium ion carries a positive charge it remains
in solution, but when this charge is withdrawn by contact with the
negatively-charged cathode the potassium regains its usual proper-
ties and decomposes the water. It is interesting to note that over
one hundred years ago Davy conjectured that "in this state of
transition or electrical progression the chemical elements are
deprived of their wonted properties, their affinities being rendered
dormant or counteracted by the predominating influence of the
electrical attraction."

271. lonization Incomplete. Should NaCl in solution be com-
pletely ionized, the osmotic pressure produced would be twice
that produced by an equal number of molecules of sugar. Barium
chloride, BaCl 2 , since it ionizes into Ba, Cl, Cl, should produce
three times this pressure. Were this the case, doubts about
Arrhenius' theory would disappear, but it is not the case. The
osmotic pressure of NaCl is not twice that of a sugar solution of
the same molecular concentration. The explanation is that these
salts do not completely ionize. At ordinary temperatures moder-
ately dilute solutions of salts, strong acids and strong bases ionize
from 80 to 90 per cent. However, as the dilution increases so does
the dissociation and it approaches the theoretical figure when the
dilution reaches one molugram per 1000 litres.

272. Experimental Demonstration of Free Ions. The presence
of free ions was shown by Ostwald in the following experiment.
A horizontal glass tube (Fig. 118) about one-half inch in diameter



and some 20 inches long is bent up at right angles at the ends,
these terminal portions being expanded to the size of a test tube
and a piece of platinum wire C being fused through the bottom of
the end B. The tube is filled with dilute sulphuric acid. In the
end A is inserted a rubber stopper through which passes an

Fig. 118.

amalgamated rod of pure zinc. In the end B is inserted a stopper
carrying a slender glass manometer M which is filled with water,
colored for ease of observation. The zinc rod is connected to the
positive pole of a battery of five or six cells, D; the platinum wire
C is connected through the key K to the negative pole. The
instant the key is closed, the manometer indicates an increase of
pressure in B due to the hydrogen released at C.

Just before the key was closed this hydrogen must have existed
in the immediate vicinity of C in the form of free ions. From Par.
270 they must have carried positive charges. But the cathode C
was also positively charged and these ions were therefore repelled.
As soon, however, as the key was closed, the charge on C was
withdrawn, the hydrogen ions moved up to C, gave up their
charges and then recovered their status as free hydrogen atoms.

273. Ions Not from Same Molecule. According to the older
theories, when the circuit was closed the zinc and sulphuric acid
in A reacted, producing zinc sulphate and hydrogen and this
hydrogen travelled from A to B and appeared at C.

Online LibraryWirt RobinsonThe elements of electricity → online text (page 17 of 46)