«THE COPPER AND SILVER GROUPS» (‹Continued›).—«THE THEORY OF COMPLEX
IONS»
We will now turn to the consideration of a series of reactions
involving the behavior of so-called "complex ions," which are very
frequently met with in the various analytical groups and which
offer valuable methods of separation and identification of ions.
The behavior of silver nitrate solution towards ammonia forms a
convenient point of attack in taking up the general subject.
«Action of Ammonia on Silver Nitrate.»—Addition of ammonium hydroxide
solution to silver nitrate (‹exp.›) results in the formation of a
brown precipitate of silver oxide (and silver hydroxide). We may
consider the supernatant liquid to be saturated with silver hydroxide
(this is in equilibrium with silver oxide), and for the saturated
solution we may put [Ag^{+}] × [HO^{−}] = K_{AgOH}.
If more ammonium hydroxide is added to the mixture, the precipitate
dissolves readily. The excess of ammonium hydroxide must increase the
concentration of hydroxide-ion and, if no other action occurred, we
should, according to the principle of the solubility-product, expect
that the precipitate would thereby be slightly increased (p. 145),
rather than that it should be dissolved so readily. Since solution
results even when the value of the one factor, [HO^{−}], of the
product is increased, we must suspect that the value of the other
factor, the concentration [Ag^{+}] of silver-ion, is in some way made
much smaller by the addition of the excess of ammonium hydroxide.
Recalling the fact that aluminium hydroxide is soluble in excess
of sodium hydroxide, as the result of its amphoteric character, a
solution of sodium aluminate NaAlO_{2} being obtained, we might
suspect that silver hydroxide also has amphoteric properties, ‹i.e.›
that it might be capable of ionizing into "argentate ions," AgO^{−},
and hydrogen ions, AgOH ⇄ AgO^{−} + H^{+}. If such be the case, the
nonionized silver hydroxide is in equilibrium, not only with the
solid phase, [p217] but also with two sets of ions,
Ag^{+} + HO^{−} ⇄ AgOH
AgOH ⇄ AgO^{−} + H^{+}
AgOH ⇄ AgOH ↓,
and we must have ‹two› solubility-product constants, one corresponding
to the basic ionization (see above) and the other corresponding to
the acid ionization, and [AgO^{−}] × [H^{+}] = K′_{AgOH}.
If silver hydroxide have acid properties, the addition of an alkali
must suppress the hydrogen-ion and the hydroxide will go into
solution as an argentate, MeAgO. We find, however, that sodium or
potassium hydroxide, which would form an argentate very much more
readily than ammonium hydroxide, has no solvent action on silver
hydroxide (‹exp.›); on the contrary, quantitative experiments show
that the alkali makes the hydroxide still less soluble than in pure
water—as demanded by the solubility-product for the basic ionization.
It is thus evident, that the solvent action of ammonium hydroxide is
not due to its basic functions. We would suspect that we have here
an action, in which ‹ammonia› is the active component, the product
of a form of dissociation of ammonium hydroxide, of which the fixed
alkalies are incapable.
«The Complex Silver-Ammonium[425]-Ion.»—For a solution of
ammonia, in water, we have the reversible reactions: [p218]
HO^{−} + H^{+} + NH_{3} ⇄ NH_{4}^{+} + HO^{−} ⇄ NH_{4}OH, and we
note that a molecule of ammonia appears to combine first with a
hydrogen ion, to form an ammonium ion, and this then forms ammonium
hydroxide with the hydroxide ion. This suggests that ammonia may have
the capacity to combine with positive ions other than hydrogen ion,
and with metal hydroxides other than water. For an analogous reaction
of ammonia with silver ion and with silver hydroxide, we would have:
NH_{3} + Ag^{+} + HO^{−} ⇄ (NH_{3}Ag)^{+} + HO^{−} ⇄ (NH_{3}Ag)OH.
For the condition of equilibrium between ammonia, the silver-ion and
the silver-ammonium-ion, we would have[426]:
[NH_{3}] × [Ag^{+}] / [NH_{3}Ag^{+}] = K.
Experimental investigations of the quantitative relations, obtaining
in ammoniacal solutions containing silver compounds, show that ‹no
constant› value is obtained for the ratio, as just developed. But the
experimental data show equally conclusively,[427] that a constant is
obtained, ‹when the concentration of the ammonia is raised to the
second power, in the mathematical statement›.
The significance of this change in the mathematical relation, it will
be recalled (p. 94), is that two molecules of ammonia must combine
with one silver ion to form an ion [(NH_{3})_{2}Ag]^{+}, whereas
in the formation of the ammonium ion, NH_{4}^{+}, we have a single
molecule of ammonia combining with one hydrogen ion. We have then
2 NH_{3} + Ag^{+} + HO^{−} ⇄ [(NH_{3})_{2}Ag]^{+} + HO^{−} ⇄
[(NH_{3})_{2}Ag]OH.
We would thus have a silver-ammonium ion, [(NH_{3})_{2}Ag]^{+},
and its hydroxide, silver-ammonium hydroxide, corresponding to the
[p219] ammonium ion and its hydroxide, ammonium hydroxide. The
properties of ammoniacal solutions of silver oxide are in entire
agreement with this conception. The hydroxide is a stronger base than
barium hydroxide.[428] It forms salts, [(NH_{3})_{2}Ag]X, in which
silver appears as part of a so-called positive "complex ion." The
hydroxide, like ammonium hydroxide, is unstable and is only known
in solution and in the presence of free ammonia, exactly as is the
case for ammonium hydroxide. The mathematical equation expressing the
equilibrium conditions for the complex ion,
[NH_{3}]^2 × [Ag^{+}] / [(NH_{3})_{2}Ag^{+}] = K_{Instability},
gives a ‹definite measure› of the ‹stability› of this complex ion.
It is clear, that the ‹larger› the constant, the more ‹unstable› the
complex ion would be, and so the constant is called the ‹Instability
Constant›[429] of the complex silver-ammonium-ion. Bodländer found
the value of the constant to be 6.8E−8 at 25°.[430]
According to the composition of the complex ion, two molecules of
ammonia should be required for every molecule of silver nitrate,
to produce a solution containing the nitrate of the complex ion:
Ag^{+} + NO_{3}^{−} + 2 NH_{3} ⇄ [(NH_{3})_{2}Ag]^{+} + NO_{3}^{−}.
As a matter of fact, 20 c.c. of a molar solution of ammonium
hydroxide (= 200 c.c. of a 0.1 molar solution) must be added to
100 c.c. of a 0.1 molar solution of silver nitrate, to convert the
silver nitrate into the salt of the complex silver-ammonium-ion.
If the ammonium hydroxide solution is allowed to flow slowly, from
a pipette, into the silver nitrate solution, we find that the last
trace of the precipitated silver hydroxide redissolves just as the
‹last› drop or two of the 20 c.c. is added to the mixture (‹exp.›).
Working more exactly, Reychler[431] found that the addition of
ammonia to a silver nitrate solution, in the proportion of two
molecules of the former to one of the nitrate, does not change the
freezing-point of the solution, and therefore [p220] does not
increase the total number of molecules in the solution. This result
agrees with the conception that two molecules of ammonia combine with
one silver ion to form a complex ion.
«Application in Analysis.»—Turning now to the consideration of
the bearing of these relations on the detection of silver-ion in
analysis, we may conclude, in the first place, from the value of
the constant as given, that only a ‹small proportion› of the total
silver in such ammoniacal solutions is present in the form of
silver-ion; but, in the second place, there is, at least, a ‹portion›
of the silver present in the form of its ion—it is ‹not entirely›
suppressed; and, in the third place, it is clear, from the form of
the equilibrium equation, that any ‹excess› of ‹ammonia› must very
rapidly reduce the ‹concentration of silver-ion› in such solutions.
The bearing of these relations, which, it will be noted, concern
‹concentrations of silver-ion›, can best be seen by working with
solutions of definite concentrations.
If the solution we have just prepared is diluted with water to 200
c.c., a 0.05 molar solution of [(NH_{3})_{2}Ag]NO_{3} is formed. In
such a solution, the concentration, [Ag^{+}], of the silver-ion is
only 0.0009,[432] whereas in 0.05 molar silver nitrate solution it
is 0.0435. It is clear that the reactions of the silver-ion will
not be observed as readily in such an ammoniacal solution as in a
solution of silver nitrate, which contains the same concentration of
‹total silver›. That such is the case, may be readily demonstrated as
follows: the addition of 1 c.c. of molar sodium bromate to 10 c.c.
of 0.05 molar silver nitrate immediately forms a heavy precipitate
of the moderately difficultly soluble bromate, AgBrO_{3}, while the
same addition to 10 c.c. of the 0.05 molar silver-ammonium nitrate
solution produces no precipitate whatever (‹exp.›). [p221]
A liter of water dissolves 0.025 mole (6 grams) of silver
bromate[433] at 18°. If the same degree of ionization be assumed
for it as for a 0.025 molar solution of the analogous salt, silver
nitrate, AgNO_{3}, 90% of the silver bromate in the saturated
solution is ionized. The solubility-product constant then is
[Ag^{+}] × [BrO_{3}^{−}] = (0.025 × 0.9)^2 = 0.0005.
When 1 c.c. of molar sodium bromate is added to 10 c.c. of 0.05 molar
silver nitrate, each salt is ionized 80% in the mixture, and
[Ag^{+}] = 0.05 × 0.8 × 10 / 11 = 0.037 and [BrO_{3}^{−}] =
(1 × 1 / 11) × 0.8 = 0.072. Then the product of the ion
concentrations, [Ag^{+}] × [BrO_{3}^{−}] = 0.037 × 0.072 = 0.0027, is
considerably larger than the constant 0.0005 and precipitation
follows.
But, when 1 c.c. of molar sodium bromate is added to 10 c.c. of
a 0.05 molar silver-ammonium nitrate solution, the concentration
of silver-ion[434] is 0.00085 and the product of the ion
concentrations, [Ag^{+}] × [BrO_{3}^{−}] = 0.00085 × 0.072 = 6E−5,
is smaller than the constant; the solution will not be saturated
with silver bromate and no precipitate is formed.
On the other hand, if 1 c.c. of a 0.1 molar solution of sodium
chloride is added to 10 c.c. of 0.05 molar silver-ammonium nitrate
solution, a very decided precipitate of silver chloride is formed
(‹exp.›). The difference in the action of the sodium bromate and
the chloride lies in the fact that silver chloride is 2500 times as
insoluble as is silver bromate, and the chloride may be precipitated
from solutions containing a ‹very much smaller concentration of the
silver-ion› than is required for the precipitation of silver bromate.
The quantitative relations for the chloride are as follows: a
liter of water dissolves 0.002 gram, or 1.4E−5 mole, of silver
chloride at 25°,[435] and the solubility-product constant at
25° is [Ag^{+}] × [Cl^{−}] = (1.4E−5)^2 = 2E−10. Now, if 1
c.c. of 0.1 molar sodium chloride is added to 10 c.c. of 0.05
molar silver-ammonium nitrate, we have,[436] for the first
moment, [Ag^{+}] = 8.9E−4 and [p222] [Cl^{−}] = 0.008, and
[Ag^{+}] × [Cl^{−}] = 8.9E−4 × 0.008 = 7E−6, which is much larger
than the solubility-product constant, and precipitation must take
place. The precipitate will be quite a heavy one: as silver-ion is
removed from solution, the complex ion must decompose and furnish
a new supply of silver-ion, and precipitation must continue until
the excess of ammonia, which is liberated by the decomposition of
the complex ion (Ag(NH_{3})_{2}^{+} + Cl^{−} → AgCl ↓ + 2 NH_{3}),
suppresses the silver-ion sufficiently to satisfy, with the
diminished concentration of chloride-ion, the solubility-product
constant of silver chloride.
It is clear that, while the reactions of silver-ion are not obtained
‹as readily› in the ammoniacal solution as in an equivalent solution
of silver nitrate (bromate experiment), nevertheless more sensitive
tests show that a ‹small portion› of the silver still is present as
‹silver-ion› in the ammoniacal solution (chloride experiment).
This brings us to our third point, the influence of an excess of
ammonia on the concentration of silver-ion and on its reactions. It
is evident, from the form of the equilibrium equation (p. 219), that
any excess of ammonia must very rapidly reduce the concentration of
silver-ion. We may ask ‹what excess will be required to prevent the
precipitation of silver chloride› in the experiment just tried.
The question may be answered as follows: The concentration of
chloride-ion, when 1 c.c. of 0.1 molar sodium chloride is added
to 10 c.c. of 0.05 molar [Ag(NH_{3})_{2}]NO_{3}, no precipitate
being formed, will be 0.1 × (1 / 11) × 0.87, the solution being
diluted 1 to 11 and the percentage of ionization of a salt MeX
being approximately 87% in 0.05 to 0.06 molar concentration. For
a solution containing this concentration of chloride-ion, the
concentration [Ag^{+}] of silver-ion, ‹which may just be present›
«without» ‹leading to the precipitation of silver chloride› (‹i.e.›
for the saturated solution) is, according to the principle of the
solubility-product,
[Ag^{+}] = K_{S.P.} / [Cl^{−}] = (2E−10) / (0.1 × 0.87 × 1 / 11).
Further, in the presence of an excess of ammonia, practically
all of the silver is present in the complex form, and,
[Ag(NH_{3})_{2}^{+}] = 0.05 × 0.87 × 10 / 11 the 0.05 molar solution
being diluted 10 parts to 11 and the salt being 87% ionized.
If we call ‹x› the concentration of free ammonia required to reduce
the concentration of silver-ion to the small value indicated, we may
put
[NH_{3}]^2 × [Ag^{+}] / [Ag(NH_{3})_{2}^{+}] =
[‹x›^2 × 2E−10 / (0.1 × 0.87 × 1 / 11)] / (0.05 × 0.87 × 10 / 11) =
6.8E−8.
Solving for ‹x›, we obtain ‹x› = [NH_{3}] = 0.33. The concentration
of free ammonia, necessary to prevent precipitation of silver
chloride in this system, is then [p223] 0.33, instead of 0.0018,
present in the original solution. Now, 11 c.c. of 0.33 molar ammonia
is equal to 11 × 0.33 / 6, or 0.61 c.c. hexamolar ammonia.[437]
Calculations, based on the solubility-product constant of silver
chloride and on the instability constant of silver-ammonium-ion,
lead, thus, to the conclusion that an excess of 0.61 c.c.
of hexamolar ammonia is required, in 10 c.c. of 0.05 molar
[Ag(NH_{3})_{2}]NO_{3}, to prevent the precipitation of silver
chloride by 1 c.c. of 0.1 molar sodium chloride. Conversely, this
excess of ammonia will be required to ‹redissolve the precipitate›
of silver chloride, formed when 1 c.c. of 0.1 molar sodium chloride
solution is added to 10 c.c. of 0.05 molar [Ag(NH_{3})_{2}]NO_{3}.
The following experiment shows that such is the case.
EXP. 1 c.c. of 0.1 molar sodium chloride is added to 10 c.c. of
0.05 molar silver-ammonium nitrate, prepared as described on p.
220; hexamolar ammonia is slowly added to the mixture from a 1
c.c. pipette, graduated in twentieths of a cubic centimeter. The
precipitate will be seen to be just about ‹completely dissolved›
when 0.6 to 0.65 c.c. of the ammonia solution has been used.
We find, in this way, that the equilibrium equation for the
instability constant of the complex silver-ammonium-ion, together
with the principle of the solubility-product, allows a ‹quantitative
interpretation› of the problem of the behavior of ammoniacal silver
solutions, as far as the detection of silver by the precipitation of
its salts is concerned.
If a still larger excess of ammonia is used (‹exp.›), even the
addition of a 10% solution of sodium chloride fails to precipitate
the chloride, and, ‹vice versa›, ammonia in excess will readily
redissolve a heavy precipitate of silver chloride (‹exp.›). Advantage
is taken of this fact in the separation and identification of
silver-ion (Laboratory Manual, ‹q. v.›).
It is interesting to note that the addition of potassium bromide,
iodide or sulphide to the ammoniacal solution, in which sodium
chloride fails to precipitate any silver chloride, will still
precipitate silver bromide, iodide or sulphide readily (‹exp.›).
Judged by the line of argument used above, in contrasting the
behavior of silver bromate and silver chloride, these silver salts
[p224] must be still less soluble than the chloride. Experiment
proves, that such is, indeed, the case.[438]
Silver iodide is so insoluble, that ammonia[439] may be used to
separate it, with a considerable degree of accuracy, from silver
chloride, and this separation forms the basis of a method to detect
chloride-ion in the presence of an iodide. If a solution with a
‹limited concentration of ammonia› is used, the method may be
extended also to the separation of chlorides from bromides (see
Chap. XVI).[440]
«Complex Metal-Ammonium Ions of Copper, Cadmium, etc.»—Quite a number
of metal ions are capable of forming more or less stable complex
ions with ammonia. For analytical purposes, the most interesting of
such complex ions, aside from the silver-ammonium-ion, are those
formed by cupric, cadmium, zinc and nickel ions, the most important
of which represent bivalent complex ions of the composition[441]
Me(NH_{3})_{4}^{2+}. The following instability constants have been
determined:[442] At 21°,
[Cd^{2+}] × [NH_{3}]^4 / [Cd(NH_{3})_{4}^{2+}] = 1E−7
[Zn^{2+}] × [NH_{3}]^4 / [Zn(NH_{3})_{4}^{2+}] = 2.6E−10.
Cupri-ammonium-ion is far more intensely blue than cupric-ion and
its color is used as one of the tests to identify copper in its
salts. Nickel-ammonium-ion is also blue, a much paler blue, and its
color must not be mistaken as indicating the presence of a dilute
solution of cupric-ammonium-ion. The same kind of relations obtain
for these complex ions as for silver-ammonium-ion. For instance, a
salt like cupric phosphate, which is readily precipitated from cupric
sulphate solutions, is not precipitated from the ammoniacal solutions
containing an excess of ammonia (‹exp.›), while the very much less
soluble[443] cupric sulphide is readily precipitated even from the
ammoniacal solutions (‹exp.›). It may [p225] easily be shown, in the
usual way,[444] that the sulphide of copper is very much less soluble
than its phosphate (‹exp.›).
«The Complex Cyanide Ions.»—Metal ions are capable of forming complex
ions, of importance in analytical work, with a number of components
other than ammonia. Among the most important of these are the complex
ions formed with cyanide-ion. The theory of the complex cyanide ions
is entirely analogous to that of the complex metal-ammonium ions,
but there is a difference in stability that makes their special
consideration desirable, both for practical and for theoretical
purposes. The complex ions of silver-ion and cyanide-ion will be
discussed first.
«The Argenticyanide-Ion.»—When potassium cyanide is added to a
solution of silver nitrate, a very insoluble precipitate of silver
cyanide is obtained, but an excess of potassium cyanide readily
redissolves the precipitate (‹exp.›). Since solution is effected in
spite of the presence of an excess of the precipitating cyanide-ion,
one is led to suspect that the other ion, the silver-ion, which is
needed to form the precipitate, is suppressed by entering into some
kind of complex with the excess of cyanide. As a matter of fact,
the solution contains a salt, potassium argenticyanide KAg(CN)_{2},
in which the silver forms a part of a ‹negative argenticyanide-ion›
(Ag(CN)_{2}^{−}).[445] If a current of electricity is passed through
such a solution, the silver (all but traces), together with the
cyanide groups, moves towards the positive electrode.[446] The
complex has been formed, then, by the combination of a positive
silver-ion with two negative cyanide ions,[447] which produce a
univalent negative argenticyanide-ion, Ag(CN)_{2}^{−}. Recalling
the fact that the complex silver-ammonium-ion is not perfectly
stable, one might suspect that the complex cyanide-ion, in
turn, is not absolutely stable, and that the action, by which
it is formed, is balanced, when equilibrium is reached, by a
reverse action of decomposition. We would have, then, [p226]
K^{+} + Ag^{+} + 2 (CN)^{−} ⇄ K^{+} + [Ag(CN)_{2}^{−}] or, more
simply, Ag^{+} + 2(CN)^{−} ⇄ [Ag(CN)_{2}^{−}].
For the condition of equilibrium between the complex and its
components, the relation
[Ag^{+}] × [CN^{−}]^2 / [Ag(CN)_{2}^{−}] = K_{Instability}
would hold. Bodlaender[448] determined the value of this constant by
measuring the concentrations of the three components under varying
conditions. The value found is 1E−21. The value of the instability
constant for [Ag(NH_{3})_{2}^{+}], of analogous composition,
is 6.8E−8, a very much larger value than the constant of the
[Ag(CN)_{2}^{−}] complex. The latter is, therefore, by far the more
stable. It must, consequently, be much more difficult to obtain
reactions, such as precipitations, of silver-ion in cyanide than
in ammoniacal solutions. In fact, it is impossible to precipitate
silver chloride by the addition of sodium chloride to KAg(CN)_{2}
solution (‹exp.›).[449] Silver sulphide was found to be a much less
[p227] soluble salt than the chloride (p. 224), and ammonium or
sodium sulphide solution, when added to the cyanide solution, readily
precipitates silver sulphide (‹exp.›). (The sulphide is ‹capable› of
‹existence› in the solid phase, therefore, under these conditions.)
In view of the extremely small concentrations of silver-ion in the
cyanide solution, we have here a striking illustration of the extreme
insolubility of the sulphide.
According to the equilibrium equation given above, the larger the
excess of cyanide-ion in the solution, the smaller must be the
concentration of silver-ion which is capable of existence in its
presence. In agreement with this conclusion, we find that the
addition of an excess of potassium cyanide readily redissolves
the precipitated silver sulphide (‹exp.›).[450] In other words,
even the minute concentration of silver-ion, that must be present
in the supernatant liquid above a precipitate of silver sulphide
([Ag^{+}]^2 × [S^{2−}] = K_{Sol. Prod.}, for a saturated liquid),
cannot be permanently present with an excess of potassium cyanide.
Consequently, ‹the solid phase is incapable of existence in the
system›.
The equilibrium equations give us, thus, a comprehensive basis for
the interpretation of the behavior of cyanide solutions containing
silver. First, in accordance with the small value of the constant,
we find it very much more difficult to obtain precipitates of silver
salts in cyanide, than, say, in ammoniacal solutions; secondly,
in accordance with the fact that a very small, but definite,
concentration of the silver-ion may still persist in the [p228]
system, we find it possible, in the absence of an excess of cyanide,
to precipitate such an extremely insoluble silver salt as silver
sulphide represents; and finally, in accordance with the form and
constant of our equation, we find it possible, by using an excess of
potassium cyanide, to suppress the silver-ion to the point where even
this extremely insoluble salt can no longer exist.[451]
«Cuprocyanide and Cadmicyanide Ions.»—Very many of the metal ions
are capable of forming complexes with cyanide-ion, of greater or
smaller degrees of stability, and, as is the case for the complex
ions formed by metal ions with ammonia, a metal ion is frequently
able to form more than one complex with cyanide-ion.[452] A number of
these complex cyanide ions are of particular interest in qualitative
analysis. For instance, we make use of the difference in the
stability of the cuprocyanide and the cadmicyanide ions as offering
us the most convenient method of recognizing cadmium in the presence
of copper. Excepting for the sulphide and the oxide, cadmium does not
form salts and compounds of characteristic colors, and, except in
color, the salts resemble the corresponding copper salts very much in
their physical and chemical behavior. Copper and cadmium consequently
show the same group reactions in systematic analysis. The more
intense colors of the copper compounds—the black sulphide, the
intensely blue cupric-ammonium-ion—mask the cadmium reactions. But
cupric-ion may be converted, by potassium cyanide, into a complex ion
of extreme stability, from the solutions of which hydrogen sulphide
and alkali sulphides ‹fail to precipitate any sulphide of copper,
while cadmium sulphide may be precipitated from the solutions of the
much less stable complex cadmicyanide-ion›.
When potassium cyanide is added to the deep blue ammoniacal
solution of cupric-ammonium sulphate [Cu(NH_{3})_{4}]SO_{4}, the
cupric-ion is reduced[453] [p229] to cuprous-ion, and the latter
is converted, by an excess of cyanide, into the extremely stable
complex ion Cu(CN)_{3}^{2−} and its salt, potassium cuprocyanide
K_{2}[Cu(CN)_{3}]. The instability constant of the complex ion
is: [Cu^{+}] × [CN^{−}]^3 / [Cu(CN)_{3}^{2−}] = 0.5E−27, and
the concentration of cuprous-ion, in a 0.1 molar solution, is
approximately[454] 3.7E−8. Without an excess of cyanide, traces of
cuprous sulphide may still be precipitated, but ‹a few drops excess
will prevent the precipitation entirely›.[455]
With an excess of potassium cyanide, cadmium forms the salt
K_{2}[Cd(CN)_{4}], yielding the ion [Cd(CN)_{4}^{2−}]. The
instability constant[456] of the complex ion is [Cd^{2+}] ×
[CN^{−}]^4 / [Cd(CN)_{4}^{2−}] = 1.4E−17. The concentration of
cadmium-ion, in a 0.1 molar solution of the salt, is then
approximately[457] 8E−5.
If potassium cyanide is added to an ammoniacal solution, containing
both cadmium and copper, until the color of the solution is just
discharged, and if two or three drops excess of the cyanide is then
used, the addition of ammonium sulphide will precipitate pure cadmium
sulphide (‹exp.›), while ammonium sulphide, added to a portion of
the original ammoniacal solution, will precipitate a dark mixture
of cupric and cadmium sulphide (‹exp.›), in which the yellow color
of cadmium sulphide is masked by the black precipitate of cupric
sulphide.
«Cobalticyanide and Nickelocyanide Ions.»—In much the same way, in
the identification of nickel in the presence of cobalt, advantage
may be taken of the fact that cobalt forms an extremely stable
cobalticyanide[458] ion, [Co(CN)_{6}^{3−}], which does not permit of
the precipitation of cobaltic hydroxide, whereas nickel does not form
such an ion, but only forms a not very stable nickelocyanide ion,
[Ni(CN)_{4}^{2−}], which is readily decomposed by bromine and alkali,
nickelic hydroxide being precipitated. [p230] The following are the
chief actions involved in the precipitation of the latter:[459]
Ni(CN)_{4}^{2−} ⥂ Ni^{2+} + 4 CN^{−}
2 Ni^{2+} + Br_{2} ⥂ 2 Ni^{3+} + 2 Br^{−}
Ni^{3+} + 3 HO^{−} ⥂ Ni(OH)_{3} ↓
«Applications and Precautions in Analysis.»—The complex cyanide
ions thus give us a convenient means of ‹interfering› with the
precipitation of certain metal ions, and of enabling us, thereby,
to detect other, closely related, ions in their presence. At the
same time, we must be careful ‹to identify the ions›, which we wish
to suppress, before converting them into these extremely stable
complexes. Potassium cuprocyanide and potassium cobalticyanide
solutions would not give any of the ordinary tests for ions of copper
and cobalt, and to find the latter in such solutions, by these tests,
we would have to take the trouble of destroying the complexes. Should
the destruction of such complexes become necessary (‹e.g.› when
complex cyanide ions are present in the original substance under
examination), evaporation with sulphuric acid, with due precautions
against inhaling poisonous hydrocyanic acid fumes, fusion with
alkali carbonates, and perhaps most conveniently, electrolysis with
sufficiently high potentials,[460] are the methods most frequently
employed for the purpose. It will be recalled that we have used the
method of fusion with potassium carbonate to find iron in potassium
ferrocyanide (p. 89).
«Ferrocyanide and Ferricyanide Ions.»—The ferro- and ferricyanide
ions may also be treated as complex ions. For instance, for the
ferricyanide-ion, we would expect a condition of equilibrium to exist
between the complex ion and the simple ions according to:
Fe(CN)_{6}^{3−} ⇄ Fe^{3+} + 6 CN^{−} and
[Fe^{3+}] × [CN^{−}]^6 / [Fe(CN)_{6}^{3−}] = K.
[p231]
If the fact is recalled that the extremely sensitive tests for the
ferric-ion fail to reveal the least trace of it in a potassium
ferricyanide solution, one must conclude that the ferricyanide-ion
must be extremely stable. The conception of the ferricyanide-ion
as a complex ion, subject to the above equilibrium conditions,
suggests that if a considerable excess of hydrogen-ion is added to
its solutions, the concentration of ferric-ion must be increased:
since hydrocyanic acid is an extremely weak acid (p. 104), the
ratio [H^{+}] × [CN^{−}] / [HCN] having the value 7 / 10^{10},
the addition of some concentrated hydrochloric acid must decidedly
suppress the cyanide-ion in a ferricyanide solution and thus lead
to an increase in the concentration of the ferric-ion. Under these
conditions, direct evidence of the presence of the ferric-ion, and
of the fact that the complex ion is a component in a reversible
reaction, may, indeed, be obtained, as well as further evidence of
the extreme stability of the complex. The presence of traces of
ferric-ion may be detected, namely, in the acid solution by the
thiocyanate test, applied in its most sensitive form, in which any
ferrithiocyanate produced is taken up in ether.
EXP. Potassium ferricyanide, treated with a thiocyanate and ether,
does not show the least trace of color; when some concentrated
hydrochloric acid is added to the mixture, a perfectly plain,
although noticeably faint, pink tint is imparted to the ether
solution.
Mercuric cyanide, as we have seen (p. 115), is exceptional in its
exceedingly small capacity for ionization. Sherrill[461] has found
that the dissociation constant for [Hg^{2+}] × [CN^{−}]^4 /
[Hg(CN)_{4}^{2−}] has the extremely small value 0.4E−41.
Consequently, mercuric-ion must be even more effective than
hydrogen-ion (strong acids), in suppressing the cyanide-ion and
liberating the ferrous- or ferric-ion, in solutions of ferrocyanides
or ferricyanides. In fact, when a solution of potassium ferrocyanide
is warmed, for a moment, with some mercuric oxide, the ferrocyanide
complex is to some extent decomposed; the liberated ferrous-ion is
oxidized, by the excess of mercuric oxide,[462] to ferric-ion, and
the presence of the latter, in quantity, is shown by the abundant
precipitation of ferric ferrocyanide or Prussian blue, when the
mixture is acidified with hydrochloric acid (‹exp.›). [p232]
«The Aurocyanide-Ion.»—Gold forms a particularly stable complex ion
with cyanide-ion. The constant[463] for the ratio [Au^{+}] ×
[CN^{−}]^2 / [Au(CN)^{−}_{2}] has the value 10^{−28}. This makes
potassium cyanide an excellent solvent for insoluble gold compounds,
such as gold sulphide, and the cyanide process for the extraction of
gold ores makes use of this property.
«The Reacting Components in Solutions of the Complex Cyanide
Ions.»—The extraordinary values, obtained for the constants
expressing the condition of equilibrium between cyanide ions and
some of the simple metal ions, such as gold and silver ions, and
their complex ions, have led to inquiries concerning a question
of fundamental interest in the theory of complex ions and in
the theory of ionization itself. In a 0.05 molar solution of
KAg(CN)_{2}, containing an excess of 0.1 mole of CN^{−}, per liter,
the concentration of silver-ion is only 5E−21. And yet, the addition
of potassium hydrosulphide (sufficient to make [KSH] = 0.1 molar)
will precipitate silver sulphide practically instantaneously from
the solution.[464] In solutions containing a larger excess of
cyanide, the concentration of the silver-ion is enormously reduced,
and yet, while we can no longer precipitate silver sulphide from
such a solution, metallic silver may be precipitated by zinc or
by the action of an electric potential at the cathode (Chapter
XV). The question may be asked, ‹whether we must consider that in
these actions the minute quantity of free silver ions, present at
any moment, is alone capable of the reactions› indicated, and that
the ‹decomposition of the complex› into its ‹components›—as one of
these, the silver-ion, is removed by precipitation—takes place with
‹sufficient speed› to account for the ‹rapid actions, wholly, as
direct actions of the silver ions›. The alternative to an affirmative
answer to this question is, that silver sulphide or silver ‹may be
precipitated by direct action of the precipitant› on the ‹complex›
ions, rather than on the ‹silver ions›. We may indicate the first
course suggested for the action, as follows:
2 Ag(CN)_{2}^{−} + S^{2−} ⇄ 4 CN^{−} + 2 Ag^{+} + S^{2−} ⇄
Ag_{2}S ↓ + 4 CN^{−}. I
The second, suggested, course of the action would be the following:
2 Ag(CN)_{2}^{−} + S^{2−} ⇄ Ag_{2}S ↓ + 4 CN^{−}. II
This second action would mean that we could obtain reactions of
silver ions, such as the precipitation of silver sulphide, ‹without
the intermediate formation of the free ions themselves›.
The consequences of the first, the ordinary, conception of such
actions ‹as direct actions of silver ions›, have been analyzed by
Haber[465] from the point of view of the velocities of the actions,
by which the complex must be formed from its components and be
decomposed into them, in order to satisfy the facts concerning the
precipitations. [p233]
We may consider, with Haber, a liter of a 0.05 molar solution of
K_{2}Ag(CN)_{3}, containing an excess[466] of 0.95 mole potassium
cyanide. In such a solution, the concentration of silver-ion
is reduced to 8E−24 gram-ion per liter. Now, according to the
best determinations of the ultimate dimensions of molecules,
about 10^{24} molecules are estimated to be contained in a mole
(gram-molecule), and 10^{24} ions, therefore, in a gram-ion (‹e.g.›
in 108 grams of silver-ion there would be 10^{24} individual silver
ions). Then a liter of the solution we are considering would
contain, at any moment, only eight individual silver ions, which are
different ones from moment to moment, since the reversible reactions
Ag^{+} + 3 CN^{−} ⇄ Ag(CN)_{3}^{2−}, are going on continually. Thus
100 c.c. of the solution would not contain even one silver ion all
the time, but the requirements of the equilibrium conditions could
be met[467] by silver ions "flashing up and disappearing" in such
a way, that the required average concentration in unit time is
maintained. There is nothing irrational in such a conception.
One may ask, however, what must be the velocities, with which the
complex is formed from the components, and is resolved into them,
in order to satisfy an instability constant[466] 10^{−22} and still
enable us to obtain a practically instantaneous precipitation, say
of silver sulphide, the action being analyzed on the basis of the
ordinary conception that only the silver-ion itself, and not the
complex ion, is directly active in the formation of the silver
sulphide. A condition of equilibrium, in a reversible action,
implies that the velocities of the two continuous, opposed reactions
are equal (p. 94). For the action Ag^{+} + 3 CN^{−} → Ag(CN)_{3}^{2−}
the ‹velocity› of ‹formation› of the ‹complex› is proportional to
a characteristic constant, K_{Formation}, to the concentration,
[Ag^{+}], of the silver-ion, and to the third power (see p. 94) of
the concentration, [CN^{−}], of the cyanide-ion. The ‹velocity›
of the ‹opposed reaction› of ‹decomposition› of the complex is
proportional to another characteristic constant, K_{Decomposition},
and to the concentration, [Ag(CN)_{3}^{2−}], of the complex ion.
For the condition of equilibrium, the velocities of the opposed
reactions are equal, and we derive the relation:
[Ag^{+}] × [CN^{−}]^3 / [Ag(CN)_{3}^{2−}] =
K_{Decomposition} / K_{Formation} = 1 / 10^{22}.
The equilibrium constant of the complex ion is, then, the ratio of
the velocity constants of its decomposition and formation (see p.
94). Now, the ‹velocity constant›, K_{Formation}, represents the
concentration, in moles, of the complex ion [Ag(CN)_{3}^{2−}],that
is formed in unit time from unit concentrations of its components
Ag^{+} and CN^{−}, and it may be considered as the ‹reciprocal›
of a ‹time constant›, T_{Formation}, the ‹time› required ‹to form
unit concentration› of the complex ion, while the components are
maintained at unit concentration. The analogous reciprocal relation
holds for the velocity constant, K_{Decomposition}, and a time
constant, T_{Decomposition}. The equilibrium equation, therefore,
expresses also the following relations:[468] [p234]
[Ag^{+}] × [CN^{−}]^3 / [Ag(CN)_{3}^{2−}] =
T_{Formation} / T_{Decomposition} = 1 / 10^{22}.
In words, the time required for the spontaneous decomposition of one
mole of the complex is 10^{22} times as long as the time required
to form one mole of the complex, from uniformly unit concentrations
of the components. If the concentration of silver-ion is reduced
to 1 / 10^{22} and the concentrations of the cyanide-ion and the
complex ion are maintained at 1, the formation of the complex takes
place 10^{22} times as slowly as when [Ag^{+}] = 1, and a condition
of equilibrium is produced, the time required to decompose and to
form the same amount of the complex being now equal.
This ‹relation› of ‹time constants› may be used to obtain some idea
of the consequences of assuming certain limiting values for one or
the other, the ratio being maintained at the value 1 / 10^{22}. If
the time constant T_{Formation} for the ‹formation of the complex›
be taken as one ten-thousandth of a second,[469] then, according
to Haber, a molar solution of potassium argenticyanide would not
be able to form in thousands of years sufficient silver ions to be
discovered by any direct test, a result which is not compatible with
the precipitation of silver sulphide and of metallic silver in a few
minutes, since silver ions could not be ‹supplied rapidly enough›.
It is evident, thus, that the ratio 1 / 10^{22} must indicate an
exceedingly small value for T_{Formation}, if only silver ions form
silver sulphide and silver.
If we assume that the complex is decomposed so fast as to supply
new silver ions rapidly enough, to allow us to consider the
precipitation of silver and silver sulphide as direct actions of the
silver ions, then we may, conservatively, consider T_{Decomposition}
to be about 1 / 100 second. Then T_{Formation} would be only
1 / 10^{24} second. Considering the limiting results for the
dimensions of atoms (and ions) and taking account of the fact that
the formation of the complex ‹involves electrical changes›, that
is, in modern terms, changes in position of electrons,[470] Haber
finds, that to satisfy the above value for the time constant, such
changes must involve a motion of electrical charges at a speed
about a million times as great as the velocity of light. Such a
velocity is, unquestionably, incompatible with our knowledge of the
velocities of light and of electrical charges. We must draw the
conclusion that ‹the complex argenticyanide-ion probably cannot
decompose fast enough into its ions›, to enable the latter to be
the ‹only› components which make it possible to precipitate silver
sulphide or metallic silver from its solutions[471] (see above,
p. 232). That would make it necessary to assume ‹direct action›
[p235] (as given in equation II, p. 232) ‹between the complex› and
the ‹precipitating agent›, to some extent, at least, the extent
being dependent on the concentrations involved in a given case. If
further investigations should confirm such a view, we would probably
find that ‹both› the actions under consideration (equations I and
II, p. 232) must proceed ‹simultaneously›. The second one would
have the advantage of enormously greater concentrations of the
reacting components, ‹e.g.› of the complex ion; the first one would,
probably, be found to have the advantage of an enormously greater
velocity constant. The actual velocities of the two reactions have
never been measured[472] and no final explanation of the relations
can be offered. The problem is a very important one, involving
the whole question of the mode of ionic action (‹cf.› Chap. V,
especially p. 83).
Aside from the theoretical value of the problem that has been
raised, the question of immediate moment to us, from the point
of view of analytical chemistry, is the question whether such
conclusions would invalidate, in any way, the use we have made of
the theory of complex ions, in elucidating the question of the
precipitation and nonprecipitation of salts of simple ions from
solutions of their complex ions.
The existence of a precipitate in contact with a solution is a
question of a ‹condition› of ‹equilibrium›; the question raised, as
the result of Haber's calculations, deals simply with the ‹problem
of the path, the mechanism by which equilibrium is reached›, but
the answer to it ‹does not affect the conditions, on which the
maintenance of equilibrium depends›. All the conclusions, drawn in
our discussions of precipitation from solutions of complex ions,
are concerned with ‹final conditions for equilibrium›, ‹i.e.› with
the conditions under which a ‹precipitate can exist›, and not with
the mechanism of its formation. The conclusions reached are valid,
therefore, irrespective of what the decision may ultimately be in
the question, whether the simple ions alone are acted upon, when
their salts are precipitated, or whether the complex ions are also
immediately concerned in the action. The precipitation of silver
chloride from an ammoniacal solution[473] may serve to illustrate
this point.
In the first place, the precipitation of silver chloride from an
ammoniacal solution, say by sodium chloride, may be considered to be
the result of the direct interaction of chloride ions with the small
quantity of silver ions present, ‹the complex serving only to renew
the supply of silver ions›, as the latter are removed from solution,
by the precipitation. The course of the action would be expressed by
the equations
[Ag(NH_{3})_{2}]Cl ⇄ [Ag(NH_{3})_{2}^{+}] + Cl^{−} ⇄
2 NH_{3} + Ag^{+} + Cl^{−} ⇄ AgCl + 2 NH_{3}. (1)
AgCl ⇄ AgCl ↓
When the precipitation is ended and equilibrium established, a
trace of silver chloride is in solution, in contact with the
precipitate, and, according [p236] to the principle of the
solubility-product, we must have [Ag^{+}] × [Cl^{−}] = K_{AgCl}.
Bodländer's experiments,[474] on the solubility of silver chloride
in ammonia, prove that this relation is in perfect agreement with
the facts. For the silver-ammonium-ion, the free ammonia and the
silver-ion present in the solution, we must have the relation
[Ag^{+}] × [NH_{3}]^2 / [Ag(NH_{3})_{2}^{+}] = K_{Instab. Const.}.
This relation, according to the experimental evidence, is also found
to hold.
Now, we might, on the other hand, assume that the primary or main
action, leading to the precipitation of silver chloride, is the
interaction of the chloride ions ‹with the complex ions›, rather
than with silver ions. Silver-ammonium chloride, [Ag(NH_{3})_{2}]Cl,
might first be formed, for instance, and then decompose ‹directly›
into silver chloride and ammonia. This is the simplest assumption we
can make for this kind of action and is sufficiently illustrative of
any kind of direct action between the chloride ions and the complex
ions. The path of the action would then be expressed by the equations
[Ag(NH_{3})_{2}^{+}] + Cl^{−} ⇄ [Ag(NH_{3})_{2}]Cl ⇄
2 NH_{3} + AgCl ⇄ 2 NH_{3} + Ag^{+} + Cl^{−}. (2)
AgCl ⇄ AgCl ↓
It is clear, from a comparison of this series of equations with
equations (1), that they represent the same reversible reactions in
a somewhat different order. ‹Since equilibrium conditions for any
reversible reaction› (‹e.g.› for A + B ⇄ C + D) ‹are independent
of the order in which the components› (‹e.g.› A, B, C, D) ‹are
brought into the system› (p. 94), the difference of order, indicated
by equations (1) and (2), cannot affect the final condition of
equilibrium in the system under discussion. For instance, since we
again have silver chloride in contact with its saturated solution,
we again must have [Ag^{+}] × [Cl^{−}] = K_{AgCl}, and ‹the
experimental confirmation› of this relation (and similarly of the
relation [Ag^{+}] × [NH_{3}]^2 / [Ag(NH_{3})_{2}^{+}] = K) ‹agrees
as well with this second path of the action as with the first›.
Conversely, since the path, by which the equilibrium is reached,
does not affect the condition of equilibrium, it is perfectly
legitimate to draw conclusions from equilibrium constants, ‹without
assuming to know anything at all about the path by which the
condition is reached›.[475]
For analytical work, the vital point is the ‹ultimate condition of
equilibrium, which determines whether a precipitate may exist and be
formed in a given system or not›. The instability constants of the
complex ions and the solubility-product constants of precipitates
are the constant factors involved in the ultimate conditions for
equilibrium, and they do not depend on the path by which equilibrium
is reached. Consequently, we find that the application of [p237]
the theory of complex ions to analytical problems of precipitation
has been in no wise invalidated by the problems presented by Haber.
The theory forms now, as before, in fact, the best quantitative
basis for the expression of the experimental results. The stability
constants, and the concentrations of the components used, determine
the limiting concentrations in which a given metal ion is capable of
continued existence, and determine, therefore, the question whether
an ionogen, of a given solubility, is capable of existence as a
solid phase in a given system.
«The Structure of Complex Ions.»—The ability of ammonia to form
complex ions with simple metal ions is commonly ascribed to the
unsaturated condition of ammonia (see p. 65). It is interesting
to recall the fact, that the marked power of the cyanide-ion to
form complexes of extraordinary stability, is also associated with
a similarly unsaturated condition of the cyanide-ion (p. 66).
According to the results of Nef's researches,[476] we have, in
hydrocyanic acid and the cyanides, unsaturated or bivalent carbon;
potassium cyanide, for instance, has the structure K—N=C<. The two
free valences of the carbon atom may, according to the electrical
theory of valence, be again considered to consist of one negative
and one positive charge. The possibility of the formation of
complexes is then self-evident, and we can readily see, how silver
cyanide, Ag—N=C±, should absorb cyanide-ion,[477] ∓C=N^{−}, and
form a complex (Ag—N=C=C=N^{−}) or Ag(CN)_{2}^{−}, whose potassium
salt would be potassium argenticyanide, K[Ag(CN)_{2}]. For the
potassium salt K_{2}Ag(CN)_{3}, of the second complex ion[478]
[Ag(CN)_{3}^{2−}] of silver and cyanogen, the most likely structure
is
Ag—N=C—C=N—K.
\ /
C=N—K
The other complex cyanide ions, ferrocyanide, ferricyanide,
cobalticyanide, etc., are considered to have structures entirely
similar to those given to the argenticyanide ions.[479]
«Complex Halide, Sulphide, Oxide and Oxonium Ions.»—In conclusion,
there are other elements besides nitrogen (in ammonia and its
derivatives) and carbon (in cyanides), which form complex ions
with simple ions; notably the halogens form such complexes.
Chloroplatinic acid, H_{2}PtCl_{6}, and its salts (Lab. Manual, ‹q.
v.›), fluorosilicic acid, H_{2}SiF_{6}, potassium-mercuric iodide,
[p238] KHgI_{3} and K_{2} HgI_{4},[480] which forms a most sensitive
reagent for the detection of traces of ammonia,[481] are instances
of complexes of the halogens that are of importance in analysis.
It is worthy of note that the most likely structure[482] for these
compounds, ‹e.g.› Cl_{2}=Pt=(Cl=Cl—H)_{2}, bears a very striking
analogy to the structures frequently assigned to the cyanide and
ammonia complex ions. Oxygen and sulphur form complex ions, of great
stability, with many elements, and we shall presently have occasion
to discuss in detail some of the complex ions formed by sulphur. All
of the oxygen acids may be treated as containing complex ions of
oxygen and some other element, and their instability constants[483]
are factors in determining their chemical behavior. For instance,
such is the case, most likely, for their behavior as oxidizing and
reducing agents (see Chapter XVI). The unsaturated condition of
oxygen in water, (H_{2}O±), makes possible, also, the formation of
complex ions, [(H_{2}O)_{‹x›}Me]^{+}, etc., called oxonium ions
and comparable with metal-ammonium ions. They form a most inviting
field for rigorous investigation. It is altogether probable that
the hydrogen-ion is intimately related, in aqueous acid solutions,
to the complex oxonium-ion, [(H_{2}O)_{‹x›}H]^{+}, comparable with
the ammonium-ion, (NH_{3}H)^{+}, and, possibly, the greater part
of the hydrogen-ion, in aqueous solutions, exists in this form of
combination.[484]
«Complex Ions of Organic Oxygen Derivatives.»—A further group of
complex ions, derived from organic derivatives of oxygen, are of
particular importance in analytical work. Many organic compounds,
such as sugars, glycerine, tartrates, citrates, interfere, more or
less, with the precipitation of metal hydroxides and certain of their
salts. For instance, the addition of cane sugar or of rochelle salt
(sodium potassium tartrate) to a solution of [p239] cupric sulphate
prevents the subsequent precipitation of cupric hydroxide by alkali
(‹exp.›). In place of a precipitate of the hydroxide, a clear,
intensely blue, solution is formed. In an analogous way, the same
substances and similar compounds interfere with the precipitation
of the hydroxides of aluminium and chromium, and since aluminium
and chromium should be precipitated as hydroxides in systematic
analysis (Chap. X), the presence of such organic compounds must be
most carefully considered, to avoid error.[485] The precipitation of
moderately difficultly soluble salts, such as phosphates, is also
rendered, appreciably, more difficult. Only extremely insoluble
salts, such as the sulphides of the arsenic, copper and zinc groups,
are precipitated, from solutions of organic substances of the
character indicated, without any appreciable interference.[486]
These relations clearly recall the characteristic behavior of
ammoniacal and cyanide solutions, in which complex ions are formed,
and the interference of the organic compounds with precipitation
is of a similar nature—complex ions are formed by metal ions with
these organic compounds, and the complexes are, in many instances,
sufficiently stable to reduce the concentrations of the metal ions
to the point, where only very difficultly soluble salts can be
precipitated.
The relations may be illustrated by the discussion of the complex
formed by the cupric-ion with tartrates. The structure of sodium
tartrate is expressed by NaO_{2}C—CH(‗OH‗)—CH(‗OH‗)—CO_{2}Na. The
underscored hydroxide groups ‗OH‗ are known in organic chemistry
as ‹alcohol groups›.[487] Now, alcohols resemble water in very
many properties and, among others, in the capacity to form metal
derivatives or ‹alcoholates›, in which the hydrogen (ion) of
the hydroxide group is replaced by metal ions. The alcoholates
correspond, thus, to the metal hydroxides, which are the analogous
[p240] derivatives of water. Exactly as there is a vast difference
in the readiness with which the various metal hydroxides, or bases,
ionize, many of them being only slightly ionizable (the weakest
bases), so certain alcoholates are much less readily ionizable than
others. The alkali alcoholates are most readily ionized.
When sodium tartrate is mixed with an excess of sodium hydroxide,
some of the ‹readily ionizable› sodium salt of the ‹alcohol› groups
of the sodium tartrate is formed and we have:
(CHONa)_{2}(CO_{2}^{−})_{2} + 2 Na^{+} ⇄
(CHO^{−})_{2}(CO_{2}^{−})_{2} + 4 Na^{+}.
When cupric sulphate is added to this mixture, a ‹slightly ionizable
complex cupri-tartrate-ion› is formed by the union of the cupric-ion
with the "alcoholate-tartrate-ion":
(CHO^{−})_{2}(CO_{2}^{−})_{2} + Cu^{2+} ⥂
[(CHO)_{2}Cu](CO_{2}^{−})_{2}.
The complex ion is not perfectly stable and so the action is a
reversible one, as indicated in the equation. The greater portion of
the copper, however, is present ‹as part of the complex negative ion›
of cupric-tartaric acid and its salts. This may be demonstrated by
subjecting the solution to electrolysis in a U-tube (p. 45). It is
readily seen (‹exp.›) that ‹a deep blue ion›, obviously containing
copper, ‹migrates to the positive pole›.[488] The concentration of
cupric-ion is so small that its hydroxide and its phosphate are not
precipitated from the solution by the addition, respectively, of
alkali or of a soluble phosphate (‹exp.›). Cupric sulphide, however,
is so insoluble that it may be precipitated completely from the
solution by the addition of a sulphide (‹exp.›), the concentration
of cupric-ion being much smaller in the saturated solution of the
sulphide than in the solution of sodium cupri-tartrate.
Citrates, sugars, glycerine, contain alcoholic groups of the same
nature as found in the tartrates, and they are capable of forming
similar complexes, or little ionizable ‹alcoholates›, with metal ions.
Certain organic acids, which contain no alcohol groups,[489] are
[p241] also capable of forming fairly stable complexes with
metal ions: thus, acetates form a complex with lead-ion, that is
sufficiently stable to render lead sulphate, which is difficultly
soluble in water, readily soluble in ammonium acetate solution[490]
(‹exp.›). Soluble oxalates readily combine with ferric, ferrous,
cupric and other oxalates and interfere, more or less, with the
detection of the metal ions, as the result of complex formations.
All of these complexes are ‹decomposed› rather readily ‹by the
addition of strong acids›, whose hydrogen-ion breaks up the complex
ions, by suppressing the anions[491] of the much weaker organic
acids and alcohols. Consequently, these organic compounds do not
interfere with tests which may be carried out in strongly acid
solution. For instance, the addition of potassium ferrocyanide to a
solution of ammonium ferrioxalate (NH_{4})_{3}Fe(C_{2}O_{4})_{3},
to which an excess of ammonium oxalate has been added, gives only
a slight indication of the presence of ferric-ion (a greenish blue
solution is obtained); when hydrochloric acid, in excess, is added
to the mixture, Prussian blue, ferric ferrocyanide, is immediately
precipitated in quantity (‹exp.›).
FOOTNOTES:
[425] The naming of the complex ions, which ammonia forms
with metal ions, has not yet been satisfactorily settled.
English writers frequently speak of "ammonio-argentic" ion
and "ammonio-argentic" nitrate. German writers speak of
"Silber-ammoniak" ion (Abegg, ‹Handb. der anorg. Chem.›, II, 728),
which would read "silver-ammon‹ia›" ion in English terms. The
terminology "silver-ammoni‹um›" ion, used in this book, is based on
the idea, that all these complex ions are essentially ‹of the same
nature› as the well-known ‹ammonium-ion›, NH_{4}^{+}, the positive
charge being, almost certainly, carried by ‹nitrogen› in these
complex ions, as it is in ammonium-ion. The latter is a ‹complex
ion of ammonia› with ‹hydrogen›-ion. The name "ammonio-argentic"
ion does not bring out this close relationship and puts the
emphasis on the silver, which is probably little concerned in the
reactions of the complex as such. The names "silver-ammon‹ia›" ion
and "silver-ammon‹ia›" nitrate sound badly and do not emphasize
the relation to ammoni‹um›, potassi‹um›, sodi‹um› and similar
positive ions and their salts. The term "‹ammonium›" is, for the
reasons given, used here in a ‹generic› sense for all complex ions
of ammonia with simple metal ions (such as H^{+}, Ag^{+}, Cu^{2+},
Zn^{2+} etc.), and the number of ammonia molecules, entering into
the composition of a complex ion, is not indicated in the names.
A similar nomenclature has long been in vogue, and has worked
well, for the complex ions of metal ions with the cyanide-ion
(see below). We speak of ferro‹cyanide›, Fe(CN)_{6}^{4−},
argenti‹cyanide›, Ag(CN)_{2}^{−} and Ag_{2}(CN)_{3}^{2−}, etc.,
without indicating the number of cyanide groups, CN, in the
complex, and we use the same generic ‹ending› "cyanide" as is used
to designate the simple cyanide ion, ‹e.g.› to designate the ion
formed from potassium cyanide, KCN ⇄ K^{+} + CN^{−}.
[426] The hydroxide-ion appears with the same coefficient, 1, on
both sides of the equilibrium equation and need not be included in
the mathematical statement; it would appear as a factor in both
terms of the ratio given and would cancel out.
[427] Bodländer and Fittig, ‹Z. phys. Chem.›, «39», 602 (1903).
[428] Bonsdorff, ‹Ber. d. chem. Ges.›, «36», 2324 (1903).
[429] It is also frequently called the ‹dissociation constant› of
the complex ion, indicating the tendency of the complex ion to
dissociate into its components.
[430] Two independent experimental methods were used and gave
concordant results—one having as its basis the solubility of silver
salts (chloride, bromide), the other the electrolytic potentials of
silver against ammoniacal silver solutions (see Chap. XV).
[431] ‹Bull. de la Soc. Chim. de Paris›, (3), «13», 386 (1895).
[432] We may consider the salt to be ionized to about the same
extent as ammonium or potassium nitrate in 0.05 molar solutions,
or, approximately, 87%. If we call ‹x› the concentration of
silver-ion, formed by the decomposition of the silver-ammonium-ion,
then 2 ‹x› is the concentration of the free ammonia, and
(0.05 × 0.87 − ‹x›) is the concentration of the complex ion. Since
‹x› is a small number in comparison with 0.0435, we may write, with
sufficient accuracy for our purposes,
[NH_{3}]^2 × [Ag^{+}] / [(NH_{3})_{2}Ag^{+}] =
(2 ‹x›)^2 × ‹x› / 0.0435 = 6.8E−8.
Then, ‹x› = [Ag^{+}] = 0.0009.
[433] Kohlrausch and Holborn, ‹loc. cit.›, p. 202.
[434] The dilution of the silver-ammonium nitrate (10 c.c.
to 11 c.c.) and the decrease in ionization due to the added
salt reduce the concentration of silver-ion from 0.0009 to
0.00085. [Ag(NH_{3})_{2}^{+}] = (0.05 × 10 / 11) × 0.8 = 0.0364
and 4 ‹x›^3 = 6.8E−8 × 0.0364 (see footnote, p. 220). Then
‹x› = [Ag^{+}] = 0.00085.
[435] Thiel (‹cf.› Bodländer and Fittig, ‹loc. cit.›). The
solubility given in the table at the end of the laboratory manual
refers to 18°. The constant for the complex ion was determined at
25°.
[436] The combined concentration of the salts is 0.055
and their degree of ionization may be taken as 87%, the
same as the degree of ionization of 0.05 to 0.06 molar
KNO_{3}. Then [Cl^{−}] = (0.1 × 1 / 11) × 0.87 = 0.008.
[Ag(NH_{3})_{2}^{+}] = (0.05 × 10 / 11) × 0.87 = 0.04 and
‹x› = [Ag^{+}] = 0.00089 (see the method of calculation in the
footnote, p. 220).
[437] The strong solution of ammonia is used in order to avoid
unnecessary dilution, and in the experiment, described below,
the dilution of the liquids by the added ammonia is considered
negligible.
[438] The following solubilities have been determined at 25°:
[Ag^{+}] × [Cl^{−}] = 2E−10; [Ag^{+}] = 1.4E−5.
[Ag^{+}] × [I^{−}] = 1E−15; [Ag^{+}] = 1E−8.
[Ag^{+}]^2 × [S^{2−}] = 4E−50; [Ag^{+}] = 4.3E−17.
[439] 100 c.c. molar ammonia dissolves at 25° only 0.6 milligram of
silver iodide (Bodländer, ‹loc. cit.›, p. 606).
[440] Fresenius, ‹Qualitative Analysis›, p. 378.
[441] In regard to Cu(NH_{3})_{4}^{2+} see Locke and Forssall,
‹Am. Chem. J.›, «31», 268, 297 (1904), and Dawson, ‹J. Chem. Soc.›
(London), «89», 1674 (1906).
[442] Euler, ‹Ber. d. chem. Ges.›, «36», 3403 (1903).
[443] See footnote, p. 212.
[444] See pp. 165, 210 and 213.
[445] The acid HAg(CN)_{2}, corresponding to the salt, is
crystallizable and is a strong acid. It is largely decomposed, by
water, into silver cyanide and hydrocyanic acid.
[446] See the experiments described on pp. 45 and 89.
[447] In solutions containing an excess of potassium cyanide
greater than 0.05 molar, the salt K_{2}[Ag(CN)_{3}] is formed.
The dissociation or instability constant for the complex ion
Ag(CN)_{3}^{2−} is 1E−22.
[448] ‹Z. anorg. Chem.›, «39», 222 (1904).
[449] The solubility-product constant for silver chloride at
25° is 2E−10. If the concentration of chloride-ion be made
1.0 by the addition of potassium chloride to a 0.05 molar
solution of KAg(CN)_{2}, then the concentration of silver-ion,
necessary for the precipitation of the chloride, would be
K_{S.P.} / [Cl^{−}] = 2E−10 gram-ion. Neglecting the fact
that the complex salt is not completely ionized and putting
[Ag(CN)_{2}^{−}] = 0.05, and calling ‹x› the concentration of the
cyanide-ion just necessary to prevent the precipitation of the
chloride, we have:
[Ag^{+}] × [CN^{−}]^2 / [Ag(CN)_{2}^{−}] =
2E−10 × ‹x›^2 / 0.05 = 10^{−21}.
We find ‹x› = 5E−7 mole, or approximately 0.03 milligram potassium
cyanide (cyanide-ion) per liter. This minute quantity of free
cyanide, if not originally present in the solution used, would be
formed by the liberation of potassium cyanide from the complex
(according to KAg(CN)_{2} + KCl → AgCl + 2 KCN) as soon as 2.5E−7
mole, or 0.036 milligram, of silver chloride per liter have been
formed, a quantity too small to be perceptible.
When potassium cyanide is added to a silver nitrate solution,
the precipitate formed is found to be silver argenticyanide,
Ag[Ag(CN)_{2}], the silver salt of the extremely stable complex,
rather than the simple salt, silver cyanide, AgCN [‹cf.› Bodländer,
‹Z. anorg. Chem.›, «39», 223 (1904)]. Ag[Ag(CN)_{2}] is even less
soluble than silver chloride, the solubility-product constant for
[Ag^{+}] × [Ag(CN)_{2}^{−}] being 2.25E−12. An excess of only 2E−6
mole, or about 0.15 milligram, of potassium cyanide (cyanide-ion)
per liter is sufficient to prevent the precipitation of silver
cyanide (silver argenticyanide) from a 0.1 molar solution of
KAg(CN)_{2}, and, conversely, at least this minute excess of
potassium cyanide is used in the preparation of a clear 0.1 molar
solution of KAg(CN)_{2}, by the addition of potassium cyanide to
silver nitrate, until the silver cyanide, first precipitated, is
just redissolved (Bodländer, ‹loc. cit.›). This excess, as just
explained, is more than sufficient to prevent the precipitation of
silver chloride from the cyanide solution, even by a large excess
of potassium or sodium chloride. Unless one takes into account, in
the manner indicated, this marked influence of a minute excess of
cyanide-ion in decidedly reducing the concentration of silver-ion
in these solutions, one could be led, wrongly, to infer from the
value of the instability constant of the complex ion and that of
the solubility-product constant of silver chloride, that silver
chloride should still be precipitated by the addition of sodium
chloride to a solution of KAg(CN)_{2}.
[450] For this reason potassium cyanide is an excellent cleansing
agent for stained silverware (sulphide stains), and, since it is an
intense poison, cleaning powders should be examined for it.
[451] For the quantitative relations see Lucas, ‹Z. anorg. Chem.›,
«41», 192 (1904).
[452] See Bodlaender, ‹Z. phys. Chem.›, «39», 597 (1902); ‹Ber. d.
chem. Ges.›, «36», 3933 (1903).
[453] Potassium cyanide is a powerful reducing agent (see p.
89) and is readily oxidized to potassium cyanate. The action,
presumably, takes the following course (see Chapters XIV and XV):
2 Cu^{2+} + 4 HO^{−} + KNC^{±} → 2 Cu^{+} + 4 HO^{−} + KNC^{2+} →
2 Cu^{+} + 2 HO^{−} + KNCO + H_{2}O.
[454] Put [Cu^{+}] = ‹x›, and [CN^{−}] = 3 ‹x›, and, neglecting
the degree of ionization, [Cu(CN)_{3}^{2−}] = 0.1, ‹x› being
so small that it need not be subtracted from 0.1. Then
‹x› × (3 ‹x›)^3 = 0.1 × 0.5E−27, and ‹x› = 3.7E−8.
[455] Treadwell and Girsewald, ‹Z. anorg. Chem.›, «38», 92 (1904).
[456] Euler, ‹Ber. d. chem. Ges.›, «36», 3404 (1903).
[457] Putting [Cd^{2+}] = ‹y›, we have ‹y› × (4 ‹y›)^4 = 0.1E−17,
and ‹y› = 8E−5. In view of the values of the constants, a ‹small›
excess of potassium cyanide will have a much smaller suppressing
effect on the cadmium-ion than on the cuprous-ion. For the
excess [CN^{−}] = 0.01, [Cu^{+}] = 5E−23, [Cd^{2+}] = 10^{−10}
as compared with [Cu^{+}] = 4E−8 in a 0.1 molar solution of the
salt K_{3}[Cu(CN)_{3}], and with [Cd^{2+}] = 8E−5 in a 0.1 molar
solution of K_{2}Cd(CN)_{4}.
[458] Bromine water is a convenient agent for oxidizing cobaltous
to cobaltic ions (see Chapter XV).
[459] The heavy arrows «→» [See Transcriber's Note] indicate the
main course the reversible actions take, ‹under the influence of
the reagents used›. Since the oxidation of nickel-ion by bromine
is accomplished only after the bromine has oxidized any excess of
cyanide used—potassium cyanide is a powerful reducing agent (p.
89)—the addition of cyanide, beyond a very small excess, must be
avoided (see laboratory instructions).
[460] ‹E.g.› for the precipitation of silver, copper, nickel,
cobalt and certain other metals from cyanide solutions; ‹cf.› Edgar
F. Smith, ‹Electro-Analysis› (1907).
[461] ‹Z. phys. Chem.›, «43», 705 (1903). ‹Vide› also Haber, ‹Z.
Elektrochem.›, «11», 847 (1905).
[462] 2 Fe^{2+} + Hg^{2+} → 2 Fe^{3+} + Hg ↓. If the treatment
with mercuric oxide is carried to completion the final products of
the reaction are ferric hydroxide, mercuric cyanide, mercury and
potassium hydroxide (Rose, ‹Z. anal. Chem.›, «1», 300 (1862)):
2 K_{4}[Fe(CN)_{6}] + 7 HgO + 7 H_{2}O →
3 Hg[Hg(CN)_{4}] + 8 KOH + 2 Fe(OH)_{3} ↓ + Hg ↓
[463] Bodlaender, ‹loc. cit.›
[464] The solubility-product constant of silver sulphide at 25°
is 0.5E−51; for [S^{2−}] = 0.8E−5 (p. 202), we would have in the
present case [Ag^{+}]^2 × [S^{2−}] = 2E−46, which is greater than
the constant. ‹Vide› quantitative data by Lucas, ‹loc. cit.›
[465] ‹Z. f. Elektrochem.›, «10», 433 and 773 (1904).
[466] See footnote 4, p. 225.
[467] Ostwald, ‹Allgem. Chemie›, Vol. II, part 1, p. 881 (1893).
[468] Haber, ‹loc. cit.›
[469] Then T_{Decomposition} = 10^{−4} × 10^{22} = 10^{18}
seconds, and, since there are 3.15E7 seconds in a year,
T_{Decomposition} = 3E10 years.
[470] See p. 42.
[471] Since there are still smaller "instability constants" than
that of the argenticyanide-ion (‹e.g.› for the gold-cyanide-ion the
constant is 1 / 10^{28}), there is a large margin of safety for the
plausibility of Haber's argument. For full details, his articles
(‹loc. cit.›), and the discussion (by Abegg, Bodlaender, Danneel,
‹ibid.›) aroused by them should be consulted.
[472] See Le Blanc and Schick, ‹Z. phys. Chem.›, «46», 213 (1903),
on measurements of the speed of ionic actions. The values obtained
agree, in general, with Haber's contention.
[473] The concentrations of silver-ion are large, in comparison
with those in cyanide solution, and the action is, most likely,
essentially an ionic one; but the argument applies with equal force
to cyanide systems.
[474] ‹Loc. cit.›
[475] An equilibrium constant, as we have seen, is a ‹ratio› of
velocity constants of balanced reactions (pp. 94, 233) and involves
therefore at least ‹two unknown› velocity constants. By determining
the actual ‹rate› of ‹change› with known concentrations of
reacting components, ‹i.e.› by determining the velocity constants
themselves, rather than their ratio, a definite conclusion as to
the mechanism or path of a given reaction can often be reached (see
p. 80).
[476] ‹Proceedings Amer. Academy›, 1892.
[477] In the absence of any added cyanide, it combines with
itself. Silver cyanide, according to Bodländer's results, is, in
saturated solutions, chiefly (AgCN)_{2} or Ag[Ag(CN)_{2}], ‹i.e.›
Ag—[N=C=C=N—Ag].
[478] See p. 225, footnote 4.
[479] Werner has developed quite a different theory of the
structure of complex ions. (‹Cf.› Nernst, ‹Theoretical Chemistry›,
p. 374 (1904).)
[480] Sherrill, ‹Z. phys. Chem.›, «43», 721 (1903).
[481] In Nessler's reagent, Fresenius' ‹Qualitative Analysis›, p.
141.
[482] ‹Cf.› Remsen, ‹Am. Chem. J.›, «11», 291 (1899); «14», 81
(1892) («Stud.»).
[483] For instance, for arsenious acid we have
3 H^{+} + AsO_{3}^{3−} ⇄ H_{3}AsO_{3} ⇄ As^{3+} + 3 HO^{−}
and, therefore, [As^{3+}] × [HO^{−}]^3 / ([AsO_{3}^{3−}] ×
[HO^{+}]^3) = ‹k›_{1}. Since [H^{+}] = ‹k′›_{HOH} / [HO^{−}] (p.
176), we have further, [As^{3+}] × [HO^{−}]^6 / [AsO_{3}^{3−}] =
‹k›_{2}. And since we may derive the relation [HO^{−}]^2 =
‹k›_{3} × [O^{2−}], by considering the primary and the secondary
ionization of water (see pp. 246, 278), we have, finally,
[As^{3+}] × [O^{2−}]^3 / [AsO_{3}^{3−}] = ‹K›. The constants for
the primary and the secondary ionization of water are included in
the value of ‹K›.
[484] Fitzgerald and Lapworth, ‹J. Chem. Soc.› (London), «93»,
2163 (1908); Lapworth, ‹ibid.›, 2187. ‹Vide› also Franklin on the
characteristics of the NH_{4}^{+} ion in liquid ammonia, ‹Am. Chem.
J.›, «23», 305 (1900).
[485] See the laboratory instructions, in regard to the precautions
used, to avoid errors from this source.
[486] On the other hand, ‹colloidal› organic substances, such as
casein, glue or albumen, interfere with the precipitation of even
the most insoluble sulphides, by producing ‹colloidal suspensions›
of the latter (see Chap. VII; ‹cf.› Müller, ‹Allgemeine Chemie der
Kolloide›, p. 56 (1907)).
[487] In alcohols the hydroxide group is held by a carbon atom,
whose remaining valences are satisfied by hydrogen or carbon atoms,
as in ordinary or ethyl alcohol, H_{3}C—CH_{2}(OH).
[488] Küster, ‹Z. Elektrochem.›, «4», 117 (1897).
[489] The most common organic acids contain the acid group
—CO(‗OH‗), as in acetic acid, CH_{3}CO(‗OH‗). The hydroxide group
‗OH‗ of the alcohols, ‹e.g.› in CH_{3}CH_{2}(‗OH‗), is still found
in these organic acids, but its tendency to form hydrogen-ion is
very much increased by the replacement of two hydrogen atoms of the
alcohols by the oxygen atom, as found in the acids. To a certain
degree, the properties of the alcohol hydroxide are maintained in
the properties shown by the acid hydroxide group. Thus, the organic
acids, on the whole, are still rather weak acids, and their salts,
in many instances, are appreciably less ionizable than the salts of
strong inorganic acids. The organic acids, further, may combine,
to a certain extent, with water and thus form hydrates (‹e.g.›
CH_{3}COOH + H_{2}O ⇄ CH_{3}C(OH)_{3}) containing a number of
hydroxide groups: the second and third hydroxide groups must have
a very much smaller tendency to form hydrogen ions and ionizable
salts, than has the first hydroxide group (p. 102), and the former,
thus show, more nearly, the behavior of alcoholic hydroxide groups.
Finally, organic acids also show a tendency to combine with
themselves, forming complex acids (‹e.g.›, (CH_{3}COOH)_{2} or
CH_{3}C(OH)_{2}OOCCH_{3}), from which complex salts may be derived,
which may be little ionizable. The power of the organic acids to
form complex ions—which they share with many inorganic acids—is
most likely intimately connected with the relations described.
[490] Lead acetate, itself, is less ionized than most salts and
this property contributes to the solubility of lead sulphate in
acetate solutions. (‹Cf.› Noyes and Bray, ‹loc. cit.›)
[491] On p. 231, the same effect is discussed, in detail, in
connection with the ferricyanide-ion.
[p242]