James Freeman Sellers.

An elementary treatise on qualitative chemical analysis online

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perature, which are shown by sugar in its solutions, are
identical with those manifested by gases, of which it
will be remembered that the concentration or density
of a given mass varies directly as the pressure and
inversely as the absolute temperature. To the form


of tension exercised by the dissolved sugar he gave the
name osmotic pressure.

Law of Osmotic Pressure. Van't Hoff 1 (1887) found
that a large number of solutions behave like that of
sugar, and announced the following law: The osmotic
pressure of a substance in solution is identical with the
pressure which it would exert were it in the form of a
gas occupying the same volume (i.e., the volume of the
solution) at the same temperature. 2

We may conveniently express the simple law which
governs the phenomena of gas and osmotic pressures in
the following form :


or P=-y-'

wherein M represents the number of molecules 3 in a
given body of gas, T and p the temperature and pres-
sure, and V the volume. Certain gases, such as oxygen,
nitrogen, and hydrogen, are obedient to this law within
very wide limits; but there are vapors whose behavior
with regard to it is apparently anomalous. Evidently
V can be made constant, and T and p can be measured
with any desired degree of accuracy. And therefore
unless there can be a change in the value of M, any
change in T ought to be accompanied by an exactly
proportional change in p. Now we find that certain
vapors such as that of ammonium chloride give
greater pressures than can be accounted for by either
the value of T, or the value of M which is based upon the
commonly accepted molecular weight; and, as has been
indicated, we find the explanation of this behavior in


the fact that the molecule NH 4 C1 is split up, or " disso-
ciated," l when we seek to vaporize it, into the smaller
molecules NH 3 and HC1. The analogies between the
behavior of gases and substances in solution seem to
extend to this phenomenon of dissociation, for it has
been observed of many solutes that their osmotic pres-
sures are so large as to be accounted for only on the
supposition that their molecules are split up in solution
and thereby increased in number. Sugar and other
bodies of its neutral character obey the simple law as
stated above; but acids, bases, arid salts in aqueous
solution usually exhibit anomalous pressures.

Freezing Point Depression. Moreover, this is not the
only evidence which bears upon the question of the
dissociation of the molecules of solutes. It is a matter
of common knowledge that the boiling and freezing
points of aqueous solutions are respectively higher and
lower than those of pure water. These relations were
studied carefully by Raoult, 2 who showed that the phe-
nomenon is a general one and that :

(a) When any substances are dissolved in inactive
solvents, the changes in the freezing and boiling points
of the solvents vary with the amounts of substance

(b) When equal weights of different substances are
dissolved in equal amounts of the same solvent, the
changes vary inversely with the molecular weights of
the solutes.

It was found of many bodies such as sugar that
equal depressions of the freezing point were pro-
duced by the solution of equimolecular proportions in


water ; and in such cases the depressions were exactly
in inverse ratio to the molecular weights. In other
cases, however, the solutions of equimolecular weights
of different substances produced unequal depressions;
and the solution of different weights of a given sub-
stance produced depressions which were not in exact
ratio to the weights so dissolved. In the latter anoma-
lous cases the depressions were greater than seemed to
be called for by the amount of matter which had been
dissolved, as naturally would be the case if the mole-
cules of the dissolved substances were dissociated into
more numerous and smaller molecules; and the sub-
stances which exhibited this behavior were those which
show abnormal osmotic pressures, namely, the majority
of acids, bases, and salts.

In these two pieces of independent evidence we
have a strong demonstration of the fact that many sub-
stances exhibit, when dissolved in water, a peculiar
structural condition in which their molecules are split
up into smaller bodies than are indicated by their
accepted formulae; and we have to inquire what
further evidence we have which will throw light upon
the precise nature of these submolecules. We shall
find this evidence in connection with the behavior of
solutions which are subjected to the passage of an
electric current.

Electrolytes. It lias long been known that the con-
ductivity exhibited by liquids is unlike that of metallic
conductors, in that the latter are not affected chemi-
cally by the passage of a current, whereas the former
are decomposed with separation at the electrodes -


the points where the current enters and leaves the
liquid- of products of varying character. In 1834
Faraday l suggested, in explanation of this phenomenon,
that the liquid which conducts electricity has in solu-
tion a compound whose molecules are divided into
freely moving particles, some of which are charged
with positive and the rest with negative electricity.
He named such compounds electrolytes; and to the
hypothetical fragments of their molecules he gave the
name of ions. Those which were assumed to be posi-
tively charged were called cations, and were either
metals, or atom-complexes, like NH 4 , which react analo-
gously to metals. Those bearing a negative charge
were termed anions, and were such bodies as the halo-
gens and acid radicles. The attraction or neutralizing
effect which ions of opposite polarities were supposed
to exercise upon each other, was held to maintain the
identity of the solute until the solution was subjected
to the passage of an electric current; whereupon the
introduction of electrodes of opposite polarities upset
the equilibrium previously existing between the ions
and caused them to migrate, the negative ions going
toward the positive electrode, and the positive ions in
the opposite direction. The appearance of decomposi-
tion products at the electrodes was explained as being
due to the union of the ions, upon arrival at those
points, to the molecular condition or to compounds
with the elements of water.

In 1887 it was demonstrated by Arrhenius 2 that the
solutions which exhibit normal osmotic pressures and
freezing point depressions are nonconductors of elec-


tricity, and that their solutes are not electrolytes.
Conversely, the solutions which give abnormal osmotic
pressures were proved to contain ionized solutes ;
and it was shown, furthermore, by highly accurate
experimental methods, that the degree of their con-
ductivity is proportional to the amount of dissocia-
tion as measured by the osmotic pressure. Between
the extremes presented by bodies like sugar, which
are characterized by little chemical reactivity and the
absence of conductivity and dissociation, and such
substances as salts and strong acids and bases, which
are distinguished by great reactivity and perfect con-
ductivity and dissociation, were arranged the other
varieties of chemical compounds, which possess various
but proportional activities of the three kinds.

With the establishment of these facts the phenom-
enon of electrolytic dissociation received a new signifi-
cance from the standpoint of analytical chemistry. The
behavior of molecules in solution was seen to be chiefly
dependent upon their tendency toward or from disso-
ciation. The solutions of strongly ionized bodies are
characterized rather by the reactions of the ions than
by the properties of the undissociated molecules. In
the case of sodium chloride, for example, the solution
presents certain definite properties which are charac-
teristic of the chlorine and sodium ions, and practically
none which are characteristic of salt itself. In the case
of sugar solutions, on the contrary, such properties as
are manifested are those of the sugar molecule alone ;
and no indication is to be seen in them of the nature
of the constituent elements of sugar.


Analytical Significance of Ions. Borrowing an illustra-
tion from Ostwald, let us assume that we have to deal
with 50 basic and 50 acidic units of some kind, which
may in theory unite to form 2500 distinct compounds
with as many sets of distinctly individual properties.
Were the analyst compelled to recognize these com-
pounds singly, in the solid condition, he obviously
would have to be familiar with the properties of each
individual among the whole number; and were he to
attempt to identify the individuals that might be
present in a mixture, the task would be beyond
accomplishment. Were the compounds not dissoci-
able in solution, his problem would still be scarcely
less difficult of solution ; but, being dissociable, his
task is made comparatively light. Since the proper-
ties of the solution of an ionized compound are merely
the sum of the properties of its ions, and since the
total number of ions with which we have assumed it
necessary to deal is 100, it follows that the knowledge
of 100 sets of properties is sufficient for the identifica-
tion of any of the 2500 compounds. If, as it some-
times happens, the substance under examination is not
soluble or readily dissociated, the analyst has only to
convert it by appropriate means into a body which
is soluble and dissociable, and then to determine its
nature from the character of the latter substance.

Laws of Electrolytic Dissociation. So far we have con-
sidered only the qualitative effects of electrolytic dis-
sociation; let us now examine briefly the quantitative
effects, which are of no less importance to the analytical


As has been said already, different electrolytes have
been found to show great dissimilarity in conductivity
and ionization, even when dissolved in equimolecular
proportions. But it also has been found that all are
obedient to the same law with regard to the degrees
of their dissociation, and that the dissimilarities are
accounted for by constants which depend upon the
nature of the electrolytes. The observed relations
between the amounts of dissociated and undissociated
electrolyte in a solution are expressed most simply for
binary compounds in the equation

a.b = k.c,

wherein a represents the concentration of the positive
ions, b that of the negative ions, c that of the mole-
cules of undissociated material, and k a constant func-
tion of the electrolyte. By assuming a value, such as
unity, for the total amount of electrolyte in solution,
and by representing the amount of dissociated material
by #, and the volume of the solution by v, we may
expand this equation to a somewhat more instructive
form :

c, concentration of undissociated electrolyte = ;

a and 6, concentrations of the two ions, either ion = -


By substitution we obtain the equation in the form

n = *.

1 a

Inspection of these equations, which are merely the
formal expression of observed fact, reveals :


(1) that increase in a (or b) will be accompanied by

an increase in the ratio - i.e., the free ions will


increase and the molecules will decrease ;

(2) that decrease in a (or b) will have the opposite
effect, i.e., the free ions will decrease and the mole-
cules will increase;

(3) that the degree of dissociation may vary in either
direction according as k is increased or decreased by
variation in the nature of the electrolyte;

(4) that dilution of a solution, and corresponding
increase of v, will call for an increase in the propor-
tion of dissociated solute, the degree of dissociation
approaching totality as its limit, as the dilution is
indefinitely increased;

(5) that concentration will have the opposite effect,
and that the ratio of dissociated to undissociated solute
will reach its minimum limit in a saturated solution.

Further inspection of the equation a.b = k.c will reveal
another fact which is of great practical significance
for the analytical chemist. It is evident that in the
solution of any given electrolyte, at a fixed tempera-
ture, the only possible variants will be a, b, and c.
Let us suppose that it is possible in some way to intro-
duce an added quantity of one ion, so that either con-
centration a or b will be increased. This being done,
the increase in the product a.b will demand an increase
in the value c. But the only way in which c may be
increased is through the return from dissociation of a
certain proportion of the ions. Assuming the concen-
tration b to have been increased, the concentration a


must be diminished until, by the decrease in a. b and
the corresponding increase in <?, the original condition
of equilibrium has been restored. In case that we are
dealing with a saturated solution of the electrolyte,
any increase in c will result in supersaturation of the
solution ; and we shall see that a portion of the solute
may separate in solid form. In fact, we have already
seen this in a practical way in Exp. 1, c.

In the saturated sodium chloride solution of that
experiment, a considerable portion of the solute was
present in the form of Na and Cl ions; and the re-
mainder was present in the molecular condition in
quantity sufficient to produce saturation. The addi-
tion of concentrated HC1, whose solution is very
strongly dissociated, introduced a very large excess
of Cl ions in the salt solution; and, in consequence,
the reunion of sodium and chlorine ions to the molec-
ular state was set up and continued until equilibrium
had been restored. But as the solution had already
been saturated with the molecules of salt, these re-
formed molecules were forced to separate in the solid

If we dissolve together two substances which are dis-
sociated more equally, such as KC1 and NaCl, we find
that less action of this sort takes place ; but when, of
our two solutes with a common ion, one is more strongly
dissociated than the other, the weaker is forced back
to the molecular and inactive condition.

The constant Jc has a very uniform value for neutral
salts, but varies considerably for acids and bases, being
high for strong acids and low for weak ones.


Regarding the dissociation values of k, Ostwald has
separated acids, bases, and salts into three classes :

Class 1 : Neutral salts, strong acids, and strong bases.
The strong acids mentioned are hydrochloric, hydro-
bromic, hydriodic, nitric, chloric, and sulphuric ; the
strong bases are hydroxides of the alkali and alkali-
earth metals.

Class 2 : Moderately strong acids and bases. The
acids are phosphoric, sulphurous, and acetic ; the bases
are the hydroxides of ammonium, silver, and magnesium.

Class 3 : Weak acids and bases. The acids are car-
bonic, hydrosulphuric, hydrocyanic, silicic, and boracic ;
the weak bases are the hydroxides of the trivalent
metals and of those divalent metals not mentioned in
Classes 1 and 2.

Applications. - - This discussion of the theories and
laws of electrolytic dissociation enables us to explain
many important operations and reactions in analytical
chemistry, which otherwise could hardly be understood.

A few of the explanations may be conveniently formu-
lated by questions and answers :

1. How does ionization aid chemical activity?

By dissociation of the solute into its ions, making it
possible for them to combine with other ions.

2. How may heat aid chemical activity ? 1

By producing rapid vibrations of the molecules, which
thus increases the speed of the reaction.

3. How may dilution aid chemical activity?

By expanding the volume, thus decreasing the pree-
suro 3 and increasing the degree of dissociation.


4. Why is the activity of an acid or a base usually
decreased by adding some salt of that acid or base ?

Two examples are given :

(1) 1 The addition of sodium acetate to acetic acid decreases
the solvent power of the acid, since the salt is more strongly
dissociated than the acid, and causes a portion of the latter to
reassume the molecular condition by increasing the concentra-
tion of the C 2 H 3 O 2 ions.

(2) The addition of ammonium chloride to ammonia 2 water
decreases the solvent action of the latter by increasing the con-
centration of the NH 4 ions, and decreasing the dissociation and
activity of the NH 4 OH.

5. When an excess of a normal salt of a weak acid is
added to a solution of a strong acid, why is the activity
of the strong acid destroyed, and that of the resulting
weak add greatly weakened?

If an excess of sodium acetate is added to a solution
of calcium phosphate in very dilute hydrochloric acid,
the phosphate will be precipitated in spite of the fact
that it is soluble in both hydrochloric and acetic acids.
The explanation of this behavior is as follows : 3 Hydro-
chloric acid and sodium acetate react to form sodium
chloride and acetic acid. The latter, in the presence
of the excess of sodium acetate, is forced back into the
inactive molecular condition in which it is no longer
able to hold the phosphate in solution.

6. Why does the addition of a solvent having an ion
in common with that of a solute salt tend to precipitate
the solute ?

This question already has been answered in the


explanation of the precipitation of common salt from
its solution by the addition of hydrochloric acid.

7. Why do reagents behave differently towards the
same elements in different compounds?

For example, hydrogen sulphide precipitates black
cupric sulphide from a solution of cupric sulphate, but
not from a solution of potassium cuprous cyanide.
Another example, silver nitrate precipitates white silver
chloride from a solution of potassium chloride, but not
from a solution of potassium chlorate. The general
answer to the question is that the chemical activity of
a compound depends on its dissociated ions, not on
the presence of certain elements. Hydrogen sulphide,
H 2 S, reacts with cupric sulphate, CuSO 4 , because the
latter is ionized into Cu and SO 4 . Hydrogen sul-
phide does not react with potassium cuprous cyanide,
K 3 Cu(CN) 4 , because the latter gives no free Cu ions,
but the molecule is dissociated into the ions, 3K and
Cu(CN) 4 .

In the second example silver nitrate, AgNO 3 , reacts
with potassium chloride, KC1, because the latter is
dissociated into K and Cl; but silver nitrate does not
react with potassium chlorate, KC1O 3 , as the latter salt
is dissociated into K and C1O 3 .

8. Why do reagents behave alike with various salts of
the same metal ?

When we say of any substance that it is a salt of
a certain metal, such as copper, we imply that it
dissociates in solution with the formation of ions of
that metal. These always react alike, no matter


what the negative ions be with which they are in

9. Why are normal salts usually better precipitant*
than their corresponding acids or bases?

For example, calcium chloride readily reacts with
ammonium carbonate, but not with carbonic acid. The
following equations illustrate the comparative reactivi-
ties of normal salts, acid salts, and acids :

(NH 4 ) 2 CO 3 + CaCl 2 yields an immediate precipitate;

H(NH 4 )CO 3 + CaCl 2 yields a tardy precipitate;

H 2 CO 3 + CaCl 2 yields no precipitate.
Normal salts are most completely dissociated, while
weak acids are very slightly dissociated. Acid salts of
weak acids partake of the nature of both normal salts
and weak acids. As ammonium carbonate is a normal
salt, it is moce completely dissociated than either the
acid salt, H(NH 4 )CO 3 , or the acid H 2 CO 3 , and hence
it reacts with calcium chloride more readily.

10. Why does an excess of a strong basic precipitant
redissolve many precipitates from salts of weak bases?

For example, a weak solution of sodium hydroxide
precipitates aluminum hydroxide from a strong solution
of aluminum sulphate, but on adding an excess of the
precipitant, the precipitate disappears. Two reactions
occur here :

(a) Aluminum hydroxide is formed :

6NaOH + Al 2 (SO 4 ) 3 = 2Al(OH) 3 + 3Na 2 SO 4 ; and on
adding more sodium hydroxide the white precipitate
dissolves, forming sodium aluminate, -



Interpreted in terms of the ionic theory, aluminum
being a very weak basic metal, its hydroxide is easily
influenced by a strong base. In aqueous solution
A1(OH) 3 is in equilibrium, being partly dissociated
into the ions Al + and 3 OH~, and, by loss of water,
partly into the ions H + and A1O 2 ~. When a strong
base like NaOH is added, it neutralizes the acid HA1O 2 ,
forming Na 3 AlO 3 and water. This destroys the equi-
librium, and more H + and A1O 2 ~ are developed, only
to be in turn neutralized by more NaOH. And so
the process continues till all of the A1(OH) 3 goes into
solution as Na 3 AlO 3 .



Object of Separation. It is only in rare cases that the
chemist is able to recognize and identify individual
elements or compounds in the mixtures which contain
them, without having first separated them from the
other bodies there present. In some cases, the mix-
tures being purely mechanical, a mechanical treatment
is sufficient to accomplish the separation ; in other
cases, as when the substances are present in 'solution,
- it is necessary in addition to make use of chemical
or physical processes, by which means the material
under examination is converted into such form that
the recognition of its elements is possible. We have
therefore two classes of separations, the members of
the first class being of a mechanical nature, whereas
those of the second are of either physical or chemical
character. The principal separations of the first class
are brought about by the operations of decantation,
filtration, and washing.


Decantation. When we have a mixture of a solid
with a liquid in which it is insoluble, or a mixture
of two liquids which are mutually insoluble, we may



separate them by this process, provided that their specific
gravities are so different that one of the compounds of the
mixture will settle and separate completely from the other.
From a mixture of liquid with solid, for example, water
and silver chloride, we may remove most of the liquid
by careful pouring or by suction with a pipette. From a
mixture of liquids, such as water and ether, we may
remove either layer with the pipette, or we may draw off
the lower layer by means of a separatory funnel.

Though decantation never separates completely, it is
convenient for the removal of the bulk of liquids from
finely divided precipitates which pass through the filter
paper, or from gelatinous precipitates which clog its
pores. Separation can be hastened by centrifugal shak-
ing of the mixture before decantation.

Experiment 5

(a) Dissolve a few crystals of silver nitrate in 10 c.c. of water
in a test-tube, and then add dilute hydrochloric -acid, drop by
drop, until, by shaking, the white silver chloride settles beneath a
clear liquid. Decant the liquid by pouring it off with a glass
rod held against the edge of the test-tube. Add more water
to the solid and decant again by immersing the tip of a pipette
in the clear liquid and sucking it off with the mouth (never allow
the liquid to rise to the mouth). Close the mouth-end of the
pipette with the tongue, lift out the pipette, and when the tongue
is removed the liquid will flow out.

(6) Mix 5 c.c. each of ether and water in a test-tube by shak-
ing vigorously. The lighter ether will rise to the top. Remove
either the ether or water with a pipette.

Filtration. Filtration is the separation of a solid
residue from a liquid filtrate by means of a porous


partition impervious to the residue. The partition
most frequently used in analytical work is unsized
paper, supported in a glass funnel. A circular paper
is folded twice, so as to form the quadrant of a circle,
and is then fitted into a glass funnel and dampened, so
as to adhere closely to the sides of the funnel. For
rapid filtration it is convenient first to fold the paper
once across the middle, and then to "plait" it on

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Online LibraryJames Freeman SellersAn elementary treatise on qualitative chemical analysis → online text (page 2 of 12)