«THE COPPER AND SILVER GROUPS. PRECIPITATION WITH HYDROGEN SULPHIDE»
The sulphides of the metal ions of the zinc group are readily
precipitated by ammonium or sodium sulphide, but hydrogen sulphide,
in the presence of a small excess of a strong acid, such as
hydrochloric acid, does not precipitate any of these sulphides (or
any of the sulphides of the aluminium, the alkaline earth and the
alkali groups). Under the same conditions ‹the sulphides of the metal
ions of the silver group›, Ag^{+}, Hg^{+}, Pb^{2+}, of ‹the copper
group›, Hg^{2+}, Pb^{2+}, Bi^{3+}, Cu^{2+}, Cd^{2+}, and also of ‹the
arsenic group›, As^{3+}, As^{5+}, Sb^{3+}, Sb^{5+}, Sn^{2+}, Sn^{4+},
Pt^{2+}, Pt^{4+}, Au^{+}, and Au^{3+}, are precipitated. Advantage
is taken of these relations in the following way, in systematic
analysis: after the separation of the silver group, by precipitation
of the difficultly soluble chlorides, hydrogen sulphide, in the
presence of an excess of acid, is used to precipitate the sulphides
of the ions of the copper and the arsenic groups, the two groups
being precipitated together. Hydrogen sulphide, under these
conditions, does not precipitate any sulphides of the zinc group or
those of any of the remaining groups. Hydrogen sulphide is used, in
this way, as one of the most valuable reagents in analytical work,
enabling the analyst to separate whole groups of metal ions from
other groups. There is also no other agent, equally important, which
is more likely to be used in a wrong way and to lead to error.
«The Ionization of Hydrogen Sulphide.»—Before taking up the theory
of the separation of these groups by precipitation with hydrogen
sulphide in the presence of a strong acid, a discussion of some
of the characteristics of the reagent will be in place. Hydrogen
sulphide, like carbonic acid and other dibasic acids, ionizes in
two stages; it produces first hydrogen-ion and hydrosulphide-ion,
HS^{−}, and this ion in turn dissociates, producing hydrogen-ion and
sulphide-ion, S^{2−}. [p200]
For the first dissociation, HSH ⇄ H^{+} + HS^{−}, we have
[H^{+}] × [HS^{−}] / [H_{2}S] = K_{1} (I)
The value of the constant[395] for this primary ionization of
hydrogen sulphide is 0.91E−7. It is apparent that hydrogen sulphide,
even in the primary ionization, is a very weak acid and produces
a very small concentration of hydrosulphide-ion. In a solution,
saturated at 25°, the total concentration of hydrogen sulphide is
approximately 0.1 molar, and the concentration of hydrosulphide-ion,
therefore, in the absence of any foreign acid, at most[396] 0.95E−4.
The concentration of the dissolved, nonionized hydrogen sulphide,
[H_{2}S], is practically a constant, if solutions saturated with
hydrogen sulphide under a given pressure, say under atmospheric
pressure, are considered. For such solutions, then, we may put more
simply[396]
[H^{+}] × [HS^{−}] = ‹k› = (0.95E−4)^2 = 0.9E−8. (II)
The concentration of hydrosulphide-ion is, therefore, inversely
proportional to the concentration of hydrogen-ion. It is clear that
the addition of a strong acid, readily yielding concentrations of
hydrogen-ion very much greater than 0.95E−4 (the value of [H^{+}]
in a saturated aqueous solution of hydrogen sulphide) will, as the
result of the greatly increased total hydrogen-ion concentration,
reduce the concentration of hydrosulphide-ion to correspondingly
low values. For instance, the presence of 0.1 molar hydrochloric
acid will increase the concentration of hydrogen-ion close to a
thousandfold and will reduce the concentration of hydrosulphide-ion
to 0.9E−7.
For the secondary ionization (see p. 101) of hydrogen sulphide,
HS^{−} ⇄ H^{+} + S^{2−}, we have
[H^{+}] × [S^{2−}] / [HS^{−}] = K_{2}. (III)
The value of this constant has recently been determined[397] and
found to be 1.2E−15. Recalling the fact that the concentrations
[H^{+}] and [HS^{−}] of hydrogen-ion and hydrosulphide-ion, [p201]
respectively, resulting from the primary ionization, are each[398]
0.95E−4, we have for the concentration of the sulphide-ion, in
aqueous solution saturated with hydrogen sulphide at atmospheric
pressure and 25°, [S^{2−}] = 1.2E−15.
Combining equations (II) and (III), we have, further:[399]
[H^{+}]^2 × [S^{2−}] = ‹k› × K_{2} = ‹k›_{2} = 1.1E−23, (IV)
‹which shows, directly, the relation between the concentration
of the sulphide-ion and that of the hydrogen-ion, the relation
of primary importance in considering the precipitation of
metal sulphides in acid solutions.› The ‹concentration of the
sulphide-ion› is, thus, ‹inversely proportional› to the ‹square›
of the ‹concentration› of ‹the hydrogen-ion›. A thousandfold
increase in the concentration of the latter, which is very nearly
the effect produced by the presence of 0.1 molar hydrochloric acid
([H^{+}] = 0.091), reduces the concentration of sulphide-ion in the
saturated aqueous solution a millionfold: If we call [S^{2−}]_{Ac.}
the concentration of the sulphide-ion in the acid solution,
[S^{2−}]_{Ac.} = (1.1E−23) / (0.091)^2 = 1.3E−21, whereas, in the
absence of acid, as found above, [S^{2−}] = 1.2E−15.
On the other hand, the addition of alkali to hydrogen sulphide,
by neutralizing and suppressing the hydrogen-ion, and by forming
the ‹salts› MeSH and Me_{2}S, will very greatly increase the
concentrations of the hydrosulphide-ion and of the sulphide-ion.
Since the constant for the secondary ionization of hydrogen sulphide
shows that HS^{−} is an ‹exceedingly weak› acid, its salts, Me_{2}S,
are very largely hydrolyzed, the constant for water being somewhat
greater[400] than its own. According to Knox, in a 0.1 molar solution
of Na_{2}S about 99% of the sulphide is hydrolyzed: Na_{2}S +
H_{2}O ⥂ NaSH + NaOH. In spite of this almost complete hydrolysis,
sufficient sodium sulphide remains in a solution of this [p202]
substance, to yield a concentration of the sulphide-ion that is far
greater than that obtained from a solution of hydrogen sulphide. In
a 0.1 molar solution of sodium sulphide the concentration of the
sulphide-ion is, approximately, [S^{2−}]_{alk.} = 0.9E−3, as compared
with [S^{2−}] = 1.2E−15 in a saturated solution of hydrogen sulphide
(25°, 760 mm.), and with 1.3E−21 in the same solution in the presence
of 0.1 molar hydrochloric acid.
Ammonium sulphide (NH_{4})_{2}S is the salt of an extremely weak
acid with a much weaker base than sodium hydroxide, and it is
correspondingly more completely decomposed by water. In a 0.1 molar
solution of the sulphide (NH_{4})_{2}S, we find the approximate
concentration[401] of the sulphide-ion [S^{2−}]_{am.} = 1.8E−6,
as compared with 0.9E−3 in a similar solution of Na_{2}S. But the
concentration of sulphide-ion is still enormously greater than
its concentrations in hydrogen sulphide in the absence and in the
presence of acids (see above).
The following table[402] contains a ‹summary of the concentrations
of sulphide-ion› in the various solutions discussed, as well as its
concentration in the presence of 0.2 molar hydrochloric acid. In
the separation of the copper and arsenic groups from the zinc and
aluminium groups, a concentration of hydrogen-ion corresponding to
the presence of 0.15 to 0.25 molar hydrochloric acid is satisfactory
for an accurate separation for ordinary purposes.
Solution. [S^{2−}]
0.1 molar Na_{2}S: 0.9E−3
0.1 molar NaSH: 0.8E−5
0.1 molar (NH_{4})_{2}S: 1.8E−6
0.1 molar (NH_{4})SH: 1.4E−8
0.1 molar H_{2}S, sat. aq. sol., 25°: 1.2E−15
0.1 molar H_{2}S, 0.1 molar HCl:[403] 1.3E−21
0.1 molar H_{2}S, 0.2 molar HCl: 3.5E−22
[p203]
«Precipitation of Sulphides by Hydrogen Sulphide.»—Although the
hydrosulphide-ion most likely takes part in the precipitation of
sulphides, the main effect appears to be due to the sulphide-ion.[404]
The sulphides seem to be both more stable and less soluble, than the
hydrosulphides. It is very likely that hydrosulphides, mixed with
the sulphides, are precipitated to a certain extent, but they are
unstable, lose hydrogen sulphide and go over into sulphides much
more readily than hydroxides, as a rule, change into oxides.[405]
The final condition of equilibrium, if one waited, in a given
case, until equilibrium was established, would depend, under these
conditions, rather on the concentration of the sulphide-ion than on
that of the hydrosulphide-ion, although the latter is present in much
greater concentrations than the former. For the sake of simplifying
the discussion of the theory of the separation of metal ions by
precipitation with hydrogen sulphide, the discussion will be limited
to the consideration of the sulphide-ion as the active precipitating
agent. As stated, this seems to be in accordance with the known facts.
The absolute values of the solubilities of the various sulphides,
which are involved in the discussion, are known, with any degree of
accuracy, only in a few cases. The aim of the discussion will be,
therefore, to develop, rather, the relations in the values involved,
which may be readily determined. Wherever absolute quantities can be
given, they also will be referred to.
«Theory of the Separation of Sulphides by Precipitation with Hydrogen
Sulphide. I. Precipitation of Ferrous Sulphide.»—If we prepare a
solution of ferrous sulphate, containing 27.8 grams of the salt,
FeSO_{4}, 7 H_{2}O, in one liter (0.1 molar), we may call [Fe^{2+}]
the concentration of the ferrous-ion in the solution. If such a
solution is saturated with hydrogen sulphide, under atmospheric
pressure, (‹exp.›), ferrous sulphide is not precipitated. We would
decide, on the basis of the principle of the solubility-product,
that the reason no precipitate of ferrous sulphide is formed is,
that the product of the ion concentrations is ‹smaller› than the
[p204] solubility-product constant characteristic of the sulphide
(p. 151); [Fe^{2+}] × [S^{2−}] < K_{FeS}, in which [S^{2−}] is the
concentration of sulphide-ion in the solution thus saturated with
hydrogen sulphide. If an alkali, sodium or ammonium hydroxide, is
added to the solution, a heavy precipitate of ferrous sulphide is
immediately formed (‹exp.›). The salts of hydrogen sulphide, like
all common salts, as we have just seen, are very much more highly
ionized than is hydrogen sulphide itself, and the addition of the
alkali has the effect of increasing enormously[406] the concentration
of sulphide-ion, say to ‹x›[S^{2−}] (p. 202). Under these conditions,
the product of the ion concentrations has evidently grown larger than
the constant: [Fe^{2+}] × ‹x›[S^{2−}] > K_{FeS}.
«II. Precipitation of Zinc Sulphide.»—Now, if a solution of
zinc sulphate is prepared so as to contain 28.7 grams of
ZnSO_{4}, 7 H_{2}O, per liter, which would make it of the same
molar concentration as the ferrous sulphate solution, we may, in
view of the fact that analogous salts ionize approximately to the
same extent in solutions of the same concentration, consider the
concentration, [Zn^{2+}], of zinc-ion to be practically the same
as the concentration, [Fe^{2+}], of the ferrous-ion in the ferrous
sulphate solution. We may put [Zn^{2+}] = [Fe^{2+}].
If the solution of zinc sulphate is saturated with hydrogen sulphide,
under the same conditions as were used with the ferrous salt solution
(‹exp.›), or if we add the zinc sulphate solution to the mixture
of ferrous sulphate and hydrogen sulphide (‹exp.›), we immediately
obtain heavy white precipitates of zinc sulphide. We would decide,
therefore, on the basis of the principle of the solubility-product,
that in this case [Zn^{2+}] × [S^{2−}] > K_{ZnS}. Since we have used
hydrogen sulphide under practically the same conditions, we may
consider that [S^{2−}], in this experiment,[407] is the same as in
the [p205] test with ferrous sulphate, and, by the conditions of the
experiment, we have also made [Zn^{2+}] = [Fe^{2+}]. The two factors
of the product are, therefore, the same, for the first moment, and we
may put [Fe^{2+}] × [S^{2−}] = [Zn^{2+}] × [S^{2−}] = P.
Since P is smaller than K_{FeS} and larger than K_{ZnS}, it is
clear that the ‹solubility-product constant for zinc sulphide must
be smaller than that for ferrous sulphide›. The solubility-product
constants, for similar salts, are a measure of their solubilities in
water. We may obtain their values by determining the solubilities
of salts in pure water, whenever the solubility is not affected by
other chemical changes. In the present instance, the quantitative
measurements, that have been made in this way, are open to question,
owing to the considerable hydrolysis which sulphides, as salts of a
very weak acid, undergo in solutions of such extreme dilution.[408]
Until such relations have been taken into account quantitatively, it
is better to limit ourselves for the present to the more accessible
question of ‹relative› solubility.
It is a comparatively easy matter to determine the ‹relative
solubility› of zinc and ferrous sulphides. If equal quantities of the
equivalent solutions are mixed and a precipitant, ammonium sulphide,
‹which would precipitate either sulphide, if its salt were present
alone›, is carefully and gradually added to the mixture, it will
precipitate first the less soluble one (see p. 163); and that one
alone can be present permanently (‹i.e.› in equilibrium) in contact
with the solution containing the two salts. As a matter of fact,[409]
[p206] ‹we find that zinc sulphide is precipitated first, under
conditions permitting the precipitation of ferrous sulphide if no
zinc sulphate were present›, and the precipitate of zinc sulphide
remains unchanged in the presence of a mixture of the composition
indicated (‹exp.›).
It is clear, therefore, that the prediction, based on the
conclusion drawn from the application of the principle of the
solubility-product, is verified by experiment.
Now, closer examination of the solution of zinc sulphate, from
which zinc sulphide has been precipitated by the action of hydrogen
sulphide, shows, after the hydrogen sulphide has precipitated as
much sulphide as it can and the solution has been passed through a
filter, that a very considerable proportion of zinc salt is still
present in the filtrate, and we must ask why hydrogen sulphide
fails to precipitate the zinc completely. The concentration of the
zinc-ion has grown somewhat smaller, but that is not the cause of
the nonprecipitation of zinc sulphide under the new conditions,
since hydrogen sulphide will precipitate the sulphide, if it is
passed into a solution of zinc sulphate which contains even a smaller
concentration of zinc-ion than the filtrate, in which it fails to
give any further precipitate. If the filtrate is examined, it is
found to be ‹strongly acid›, since sulphuric acid has been liberated
by the action of hydrogen sulphide on zinc sulphate: ZnSO_{4} +
H_{2}S → ZnS ↓ + H_{2}SO_{4}. Sulphuric acid is a strong acid, which
is very highly ionized, much more so than the exceedingly weak acid
hydrogen sulphide, and consequently, as the precipitation of zinc
sulphide proceeds, the ‹concentration of hydrogen-ion in the solution
rapidly grows larger and larger›. But, the greater the concentration
of the hydrogen-ion, the smaller is that of the sulphide-ion, since
the product [H^{+}]^2 × [S^{2−}] is a constant [equation (IV), p.
201] for a solution kept saturated with hydrogen sulphide. The
sulphide-ion is reduced in concentration very much more rapidly
than is the zinc-ion.[410] As [S^{2−}] is a factor in the [p207]
solubility-product of zinc sulphide, it is clear that the value of
this product must grow rapidly smaller during the precipitation of
zinc sulphide from the solution, and that it may well, eventually,
grow too small to surpass the value of the solubility-product
constant K_{ZnS}. Precipitation of zinc sulphide will then cease.
Obviously, the suppression of the sulphide-ion may be accomplished
by the addition of hydrochloric, sulphuric or any other strong acid
to the zinc sulphate solution in the first place, and then hydrogen
sulphide should fail to precipitate any zinc sulphide at all. In
fact, if to 50 c.c. of the 0.1 molar zinc sulphate solution 2 c.c. of
hexamolar hydrochloric acid is added,[411] hydrogen sulphide does not
precipitate even a trace of zinc sulphide (‹exp.›).
We have found, then, that 0.1 molar zinc sulphate solution,
acidified with a small excess of hydrochloric acid, fails to
produce a precipitate of zinc sulphide, when it is saturated
with hydrogen sulphide. We must conclude that, under these
circumstances, the product of the ion concentrations is smaller
than the solubility-product constant for zinc sulphide:
([Zn^{2+}] × [S^{2−}] / ‹x›) < K_{ZnS}, the new concentration of the
sulphide-ion being represented by the symbol [S^{2−}] / ‹x›.
It would follow, from these considerations, that the action of
hydrochloric or sulphuric acid, in preventing the precipitation
of zinc sulphide, depends on their producing a sufficiently high
concentration of the hydrogen-ion, to keep the concentration of the
sulphide-ion, in a mixture of zinc sulphate and hydrogen sulphide,
below the point where the solution could become supersaturated with
zinc sulphide. [p208] ‹For exactly similar reasons, none of the
sulphides of the zinc group is precipitated by hydrogen sulphide in
(sufficiently) acid solutions.›
It is evident, further, that, if a solution of zinc acetate (without
the addition of any acid) is substituted for the zinc sulphate
solution and is treated with hydrogen sulphide, an entirely different
result, quantitatively considered, must be obtained. By the action of
hydrogen sulphide on the acetate, acetic acid is liberated, according
to Zn(CH_{3}CO_{2})_{2} + H_{2}S ⥂ ZnS ↓ + 2 CH_{3}COOH. As a ‹weak›
acid, acetic acid produces much less hydrogen-ion than is formed
in equivalent solutions of sulphuric acid. Consequently, ‹a much
slighter suppression of the sulphide-ion› and a much more ‹complete
precipitation of zinc sulphide› from the acetate, than from the
sulphate solution, must result. Such is the case. Zinc sulphide is,
indeed, precipitated ‹quantitatively› by hydrogen sulphide from the
‹acetate› solution.
This behavior of zinc acetate[412]—and zinc salts of other ‹weak
acids› show, of course, the same behavior—represents one of the
pitfalls, into which the unwary analytical chemist is liable to fall,
when he uses hydrogen sulphide. The separation of groups by hydrogen
sulphide depends, as stated, on the fact, that, in the presence
of a certain concentration of hydrogen-ion, hydrogen sulphide
will not precipitate zinc sulphide and the remaining sulphides
of the zinc group. To secure this concentration of hydrogen-ion,
some hydrochloric acid is added to solutions, from which hydrogen
sulphide is expected to precipitate none but sulphides of the copper
and arsenic groups—and, as a rule, the purpose is accomplished, as
desired. It is evident, however, that if a solution contains an
acetate, say sodium acetate, or the salt of any other weak acid,
‹e.g.› a borate or a phosphate, the addition of hydrochloric acid
will result, at least at first, in the ‹liberation› of the ‹weaker
acid› and will not produce the excess of hydrogen-ion, required
for the analysis. Zinc sulphide, and possibly nickel and cobalt
sulphides,[412] may, under such conditions, be precipitated with
the sulphides of the groups mentioned. Unless provision is made,
therefore, ‹to insure a certain excess of hydrogen-ion› (p. 213),
or unless we are on our guard and look for zinc, nickel and cobalt
in [p209] the analysis of the precipitate formed by hydrogen
sulphide,[413] serious errors obviously could result. To add an
inordinately large excess of hydrochloric acid to mixtures, in order
to avoid this pitfall, will, as we shall presently see, only throw us
more certainly into still another error, to which we are exposed in
the use of this important reagent, hydrogen sulphide.
«III. Precipitation of Cadmium Sulphide.»—Now, if a 0.1 molar
solution of cadmium sulphate (22.6 grams of the salt,[414]
CdSO_{4}, H_{2}O, ‹per› liter) is prepared, equivalent, in
concentration, to the solutions of ferrous and zinc sulphates used
previously, we may consider that the concentration of cadmium-ion
is, approximately, the same as the concentrations of the ferrous-ion
and zinc-ion in the solutions of their sulphates. If the solution
of cadmium sulphate is added to the acidified solution of zinc
sulphate, which was saturated with hydrogen sulphide but from which
no zinc sulphide could be precipitated (‹exp.›, p. 207), a heavy
precipitate of cadmium sulphide is at once obtained (‹exp.›). Or,
if 2 c.c. of hexamolar hydrochloric acid is added to 50 c.c. of
the cadmium sulphate solution and the mixture is saturated with
hydrogen sulphide,[415] cadmium sulphide is precipitated readily
(‹exp.›) and, in fact, quantitatively. We would decide, on the basis
of the principle of the solubility-product, that, in this mixture,
precipitation results because ([Cd^{2+}] × [S^{2−}] / ‹x›) > K_{CdS}.
By the conditions of the experiment we started with a concentration
of the cadmium-ion, [Cd^{2+}], equal to the concentration, [Zn^{2+}],
of the zinc-ion in the zinc sulphate solution, and with the same
concentration,[415] [S^{2−}] / ‹x›, of sulphide-ion as was used
when hydrogen sulphide failed to precipitate zinc sulphide. The
corresponding factors of the products of the ion concentrations
are equal, at the beginning of the two experiments, and we may put
([Cd^{2+}] × [S^{2−}] / ‹x›) = ([Zn^{2+}] × [S^{2−}] / ‹x›) = P′.
We recall the fact, that we have already concluded, on the basis
of the principle of the solubility-product, that ([Zn^{2+}] ×
[S^{2−}] / ‹x›), or P′, is ‹smaller› than K_{ZnS} (p. 207), and that
([Cd^{2+}] × [S^{2−}] / ‹x›), or P′, is ‹greater› than K_{CdS}.
P′ being smaller than K_{ZnS} and larger than K_{CdS}, it is clear
that [p210] K_{CdS} is smaller than K_{ZnS} and that cadmium
sulphide must be the ‹less soluble› of the two sulphides. As a
matter of fact, if ammonium sulphide is carefully added to a mixture
of equal quantities of the two salt solutions, cadmium sulphide
is precipitated first, and when practically all of the cadmium
is precipitated, a final precipitate of white zinc sulphide is
obtained (‹exp.›; see note, p. 205). Or, if zinc sulphide is first
precipitated by the addition of a little ammonium sulphide to 25
c.c. of the 0.1 molar zinc sulphate solution, care being taken
to have zinc sulphate in excess, and if 25 c.c. of the 0.1 molar
cadmium sulphate solution is then added to the mixture, the white
zinc sulphide immediately gives way to the less soluble yellow
cadmium sulphide (‹exp.›; see p. 165). Cadmium sulphide is thus
proved to be the ‹less soluble› of the two sulphides, a result which
confirms the prediction made above with the aid of the principle
of the solubility-product, and we may indeed conclude that the
solubility-product constant K_{CdS} of cadmium sulphide must be
smaller than the constant K_{ZnS} of zinc sulphide (see pp. 163–168,
on fractional precipitation). We are, therefore, also justified
in deciding that CdS may well be precipitated from acidulated
solutions by hydrogen sulphide, when ZnS is not thus precipitated,
simply because K_{CdS} is ‹sufficiently small› (CdS is sufficiently
insoluble) ‹to make the product of the ion concentrations›
[Cd^{2+}] × [S^{2−}] / ‹x›, ‹in spite of the extremely small value
of› [S^{2−}] / ‹x›, ‹greater than the constant› K_{CdS}, whereas
the same small value of [S^{2−}] / ‹x› makes it impossible for the
product [Zn^{2+}] × [S^{2−}] / ‹x› to reach the value of the
‹larger constant› K_{ZnS}, required for the precipitation of ZnS.
Since cadmium sulphide may be precipitated quantitatively under the
conditions given, it is also evident that it may be precipitated even
when the concentration of the cadmium-ion also has a rather small
value. The relations, in regard to this point, will be discussed
presently.
«The Separation of the Copper and Arsenic Groups from the
Zinc Group.»—Ferrous-ion and zinc-ion may be taken as typical
representatives of the ions of the ‹zinc group, whose sulphides are
not precipitated by hydrogen sulphide in the presence of a definite
concentration of hydrogen-ion›, cadmium-ion as a representative of
the ‹copper›, ‹silver› and ‹arsenic› groups, ‹whose sulphides are
precipitated under the same conditions›. The separation of the groups
depends, therefore, on the different solubilities of the sulphides of
these [p211] groups, the sulphides of the zinc group being the most
soluble. Since even these sulphides are precipitated quantitatively
by ammonium sulphide and are very difficultly soluble, the separation
is a kind of fractional precipitation of difficultly soluble salts,
in which the fractionation is made possible and convenient by the
use of an agent, hydrogen sulphide, the concentration of whose
active precipitating component, the sulphide-ion, S^{2−}, is easily
regulated and readily made sufficiently small, not to precipitate
even the difficultly soluble sulphides of the iron group.
Solubilities vary from salt to salt, and we have already found that,
in the zinc group, zinc sulphide is less soluble than the sulphide
of a second member of the group, ferrous sulphide, and that the
difference is revealed in a somewhat different behavior of their
salts toward hydrogen sulphide, when the action is studied in some
detail. Similar differences must be expected to exist among the
sulphides of the groups that hydrogen sulphide precipitates even in
the presence of an excess of hydrochloric acid. As these differences
are the sources of some of the most common and most serious errors
which analysts are liable to commit, the detailed study of the action
of hydrogen sulphide must be continued a little further.
«The Effect of a Large Excess of Acid.»—The precipitation of cadmium
sulphide depends on the relation of the product of the concentrations
of the cadmium-ion and the sulphide-ion to the solubility-product
constant for cadmium sulphide (see the equation, p. 209). Now, it
is clear that if a ‹larger› and ‹larger excess of hydrogen-ion› is
introduced by the addition of more concentrated hydrochloric acid to
the cadmium sulphate solution, the concentration of sulphide-ion is
correspondingly ‹reduced›.[416] The point might be reached, where the
sulphide-ion factor becomes so small, that the product of the ion
concentrations remains smaller than the value K_{CdS}, required for
precipitation of cadmium sulphide.
In fact, if a large excess[417] of hydrochloric acid is added to 50
c.c. [p212] of the 0.1 molar solution of cadmium sulphate, hydrogen
sulphide fails to precipitate any of the sulphide (‹exp.›).
But, if a few cubic centimeters of a 0.1 molar solution of cupric
sulphate (25.0 grams of CuSO_{4}, 5 H_{2}O, per liter) are added
to the solution from which hydrogen sulphide fails to precipitate
cadmium sulphide, cupric sulphide is at once precipitated. And, if
15 c.c. of concentrated hydrochloric acid are added to 50 c.c. of
the 0.1 molar cupric sulphate solution, there results a mixture
corresponding to the cadmium sulphate solution from which hydrogen
sulphide fails to precipitate CdS; we find that hydrogen sulphide
will ‹precipitate› the ‹sulphide of copper› very readily, even under
these adverse conditions (‹exp.›). Cupric sulphide must be even
less soluble in water than cadmium sulphide,[418] and there is no
difficulty in showing that such is the case. If ammonium sulphide,
or hydrogen sulphide, is gradually introduced into a mixture of
25 c.c. each of the 0.1 molar sulphate solutions, cupric sulphide
is precipitated first, and, if the precipitate is collected in
fractions, [p213] pure yellow cadmium sulphide is precipitated
last.[419] Or, if 25 c.c. of 0.1 molar cupric sulphate is added to
the mixture in which a precipitate of cadmium sulphide displaced the
more soluble zinc sulphide (p. 210), the yellow sulphide will, in
turn, give way to the less soluble black sulphide of copper (‹exp.›).
We find thus that the precipitation of cadmium sulphide, by hydrogen
sulphide in acid solution, ‹can be prevented by the presence
of an excess of hydrochloric acid›, which does not prevent the
precipitation of the less soluble cupric sulphide.[420] The fact,
then, that, in an analysis of some unknown mixture, hydrogen sulphide
produces a precipitate in acid solution, must not be considered as
evidence that the conditions are such as to insure the precipitation
of all the sulphides of the groups, which we intend to precipitate.
To avoid error, conditions must be such as to insure the complete
precipitation of the more soluble as well as the less soluble
sulphides. The sulphides of ‹cadmium› and ‹lead›, in particular,
and, to a lesser degree, the sulphides of antimony and tin, are
most liable to remain unprecipitated and thus escape detection in
systematic analysis. This is a matter of special importance, also,
in detecting traces of the ions of these metals, especially of lead,
which is a slow cumulative poison, even when absorbed in minute
amounts, and which analysts must therefore be able to detect, even in
traces, with absolute certainty. It is clear, from a consideration of
the product of the ion concentrations, as affecting the precipitation
or nonprecipitation of such a sulphide, that a much smaller excess
of acid will prevent the precipitation of the last traces of lead
sulphide, and, therefore, of all of it, if only traces are present,
than will interfere with the precipitation of the sulphide in bulk.
«The Desirable Concentration of Acid (of Hydrogen-ion) and an
Indicator for Correct Acidification.»—Summarizing the conclusions
reached in regard to the conditions necessary for a successful
separation of the copper and arsenic groups, by means of hydrogen
sulphide, from the zinc and aluminium groups, we find [p214] that
the concentration of the hydrogen-ion, in the solution to be treated
with hydrogen sulphide, is the most important factor. Too great a
concentration, as has just been shown, will prevent the precipitation
of all, or part, of the more soluble sulphides of the former groups,
notably of the sulphides of cadmium and lead, which is a common
error in the laboratory. Too small a concentration, which may result
when a salt of a weak acid, such as an acetate, borate or phosphate
is present, may lead, as was shown above, to the precipitation of
part of the zinc group, notably of zinc and possibly of nickel
and cobalt, with the copper and arsenic groups, and thus lead to
other errors. For the ordinary purposes of analysis, requiring the
precipitation of say one milligram of any ion from 100 c.c. solution
(one-tenth per cent, if one gram of substance is used for analysis),
a concentration of hydrogen-ion of 0.1 to 0.3 gram-ion per liter
forms a good basis for work.[421] The presence of this concentration
of hydrogen-ion, irrespective of the possible presence of weak
organic or inorganic acids, may be readily insured by a simple test
with an appropriate indicator. ‹Methyl-violet›[422] is suitable for
such a purpose, because it is sensitive only to the rather high
concentrations of hydrogen-ion required: 0.1 c.c. or two drops of
0.05 to 0.1 molar hydrochloric acid, added to an equal volume of a
very dilute solution (1 : 12,500) of the indicator, changes its color
to a ‹pure blue›; 0.2 to 0.25 molar acid, used similarly, turns the
indicator to a ‹blue-green› tint, and 0.33 molar acid produces a
‹yellow› or ‹yellow-green› hue. The ‹blue-green tint›, with which one
becomes easily familiar, and which can, indeed, always be prepared
for matching tints, may be used to recognize speedily, and with
sufficient accuracy, a concentration of the hydrogen-ion of the
strength desired, irrespective of its source.[423]
If an analyst aims to find even smaller quantities of a particular
metal ion, ‹e.g.› traces of lead, the ordinary method of analysis
[p215] may be modified, the source of error in the precipitation of
traces of lead sulphide being kept in mind.[424]
Besides the complications mentioned, and provided against in the
way discussed, there is still one more complication in the use
of hydrogen sulphide: this is in the matter of the precipitation
of ‹arsenic› sulphide from solutions containing ‹arsenic in the
pentavalent condition›. Since the interpretation of this complication
and the explanation of the methods for avoiding the errors, which
may arise therefrom, are necessarily intimately connected with the
chemical behavior of arsenic acid, this subject will be considered in
the discussion of the arsenic group (Chapter XIII).
FOOTNOTES:
[395] Auerbach, ‹Z. phys. Chem.›, «49», 220 (1904).
[396] See footnote 1, p. 201.
[397] Knox (in Abegg's laboratory), ‹Trans. Faraday Soc.›, «4», 44
(1908).
[398] The concentration of the hydrogen-ion is really a little
greater than that of the hydrosulphide-ion, as a result of the
ionization of the latter, but the amount of hydrogen-ion formed in
this way (about 1E−15) is so minute, compared with that formed by
the primary ionization, that it is negligible.
[399] We can obtain the relation, directly, from H_{2}S ⇄ 2 H^{+} +
S^{2−} and [H^{+}]^2 × [S^{2−}] / [H_{2}S] = K = 1.1E−22. For
a given pressure of the hydrogen sulphide, [H_{2}S], expressing
its solubility (about 0.1 molar at 25°), is constant, and
therefore [H^{+}]^2 × [S^{2−}] = a constant, as given in equation
(IV). Putting [H_{2}S] = 0.1, we have [H^{+}]^2 × [S^{2−}] =
0.1 × 1.1E−22 = 1.1E−23.
[400] On account of the great mass of water, we compare (see
equation, p. 176) [H^{+}] × [HO^{−}] = 1.2E−14 (at 25°) with
[H^{+}] × [S^{2−}] / [HS^{−}] = 1.2E−15.
[401] The calculation was made by the method used by Knox (‹loc.
cit.›) for a molar solution. The degree of ionization of the
salt was not considered and the correct ionization constant for
ammonium hydroxide was used, 1.8E−5 in place of 2.3E−5. The
latter, evidently, was used by Knox as the result of overlooking a
correction, which Bredig made in his (Bredig's) first calculations
of the constant; ‹cf.› Bredig, ‹Z. phys. Chem.›, «13», 293,
footnote. The same erroneous constant is found in Kohlrausch and
Holborn, ‹loc. cit.›, p. 194.
[402] For further values and for the method of calculation, see
Knox, ‹loc. cit.›
[403] [S_{2−}] = 1.1E−23 / [H^{+}]^2, according to equation (IV),
p. 201.
[404] Knox's work leads to that conclusion.
[405] The precipitation of sulphides, from a solution containing
much more of the hydrosulphide-ion than of the sulphide-ion, is
comparable with the precipitation of mercuric oxide, HgO, and of
silver oxide, Ag_{2}O, by sodium or potassium hydroxide.
[406] On account of the presence of a small, unknown amount of
sulphuric acid in the original solution, resulting from the
hydrolysis of ferrous sulphate, the exact value of [S^{2−}] in the
first solution cannot be calculated without further examination;
but, according to the values given in the table on page 202, the
value of ‹x›, indicating the growth in the concentration of S^{2−},
is ‹at least› 10^{12}, if 2 equivalents of NaOH, and 10^9, if 2
equivalents of NH_{4}OH are used to convert the 0.1 molar hydrogen
sulphide into the corresponding sulphide Me_{2}S, of 0.1 molar
concentration.
[407] [S^{2−}] is exactly the same in the two products, when equal
volumes of the zinc and ferrous sulphate solutions are mixed and
the mixture is saturated with hydrogen sulphide; ‹zinc sulphide› is
precipitated.
[408] The difference in the values obtained, when hydrolysis is
considered or neglected, is very considerable. ‹Vide› Bodländer,
on the solubility of calcium carbonate, ‹Z. phys. Chem.›, «35», 23
(1900), and Stieglitz, ‹Carnegie Institution Publications›, No.
«107», 249 (1909).
[409] In carrying out this ‹fractional precipitation› a very
dilute solution of ammonium sulphide is used, so as to prevent
the mechanical enclosure of black ferrous sulphide, which would
discolor the white sulphide. The ammonium sulphide solution should
be saturated with hydrogen sulphide, to prevent the precipitation
of green ferrous-ferric oxide by an excess of free ammonia. It
is best to prepare a set of the precipitates and to preserve
them in well-stoppered vessels, and not to try to take the
time and care necessary to effect a perfect fractionation as a
lecture experiment. The presence of the ferrous sulphate, in the
supernatant liquid above the first precipitate of zinc sulphide,
may be readily demonstrated by pouring off some of the solution
and adding an excess of ammonium sulphide to it. Of course, it is
also perfectly legitimate, and easier, to precipitate first zinc
sulphide from a pure zinc sulphate solution and then to add ferrous
sulphate solution to the mixture and to preserve the mixture. If
the zinc sulphide were not the less soluble, it would be rapidly
converted into the black ferrous sulphide. (See p. 165, and see
below, pp. 210, 213, where similar transformations are carried out
as lecture experiments.)
[410] When 10% of the zinc in a 0.1 molar solution has been
precipitated, 0.01 molar sulphuric acid has been formed. For
the sake of a rough approximation, the acid may be considered
completely ionized and then [H^{+}] = 0.02, which is 200 times the
value of [H^{+}] in a saturated H_{2}S solution (p. 200); if the
presence of a little sulphuric acid in the original zinc sulphate
solution, resulting from a slight hydrolysis of the salt, is
ignored, the concentration of the sulphide-ion is decreased roughly
(200)^2 or forty thousandfold, while the concentration of zinc-ion
falls 10%. The corrections, that have been indicated, would change
the quantities involved, but they would not modify the character of
the result.
[411] This proportion of acid, making the concentration of the
hydrogen-ion, approximately, [H^{+}] = 0.2, is used, not because
it represents the minimum concentration of the hydrogen-ion, which
will prevent the precipitation of zinc sulphide in 0.1 molar
zinc sulphate solution, but because it represents the practical
conditions under which the precipitation of zinc sulphide is
avoided, when the copper and arsenic groups are precipitated in
qualitative analysis (see p. 213).
[412] Nickel and cobalt sulphides are also precipitated by
hydrogen sulphide in the presence of free acetic acid, if sodium
or potassium acetate is added, to suppress the hydrogen-ion of the
acetic acid (p. 112).
[413] They would be found in the copper group.
[414] The sulphate, of this composition, is obtained by drying the
crystallized sulphate in an air bath at 100–105°.
[415] ‹Cf.› the corresponding experiment with zinc sulphate, p. 207.
[416] [S^{2−}] = ‹k› / [H^{+}]^2. See equation (IV), p. 201.
[417] Any immediate precipitation of cadmium sulphide will be
prevented by the addition of 10 c.c. of concentrated acid (sp.
gr. 1.19) to 50 c.c. of the 0.1 molar solution, and 15 c.c. will
completely prevent any precipitation of the sulphide. Of course, a
smaller excess would prevent the precipitation of small quantities
of the sulphide (‹e.g.› a half milligram of cadmium), which should
easily be found in 50 c.c. (see p. 214).
[418] The value of the solubility-product constant for cupric
sulphide, at 25°, was determined by Knox (‹loc. cit.›): [Cu^{2+}] ×
[S^{2−}] = 1.2E−42, corresponding to a concentration of 1.1E−21
of cupric-ion. Mercuric sulphide was found even less soluble:
[Hg^{2+}] × [S^{2−}] = 2.8E−54, and its behavior agrees with such
a relation (Lab. Manual, p. 50, § 2). The solubility-product
constant for lead sulphide, which resembles cadmium sulphide in
the fact that a large excess of acid prevents its precipitation,
was found to be [Pb^{2+}] × [S^{2−}] = 2.6E−15, the constant
being about 10^{27} times as large as the constant for cupric
sulphide. This value for the solubility-product constant for
lead sulphide must either be considerably larger than the true
value or lead must be easily precipitated as a hydrosulphide,
Pb(SH)_{2}, since solutions in which the product of the ion
concentrations, [Pb^{2+}] × [S^{2−}], is very much smaller
than the constant given, readily precipitate lead sulphide.
Thus Noyes and Bray [‹J. Am. Chem. Soc.›, «29», 137 (1907)]
report it possible to precipitate 1 to 2 milligrams of lead-ion
in 100 c.c. of solution (say [Pb^{2+}] = 1E−4) with hydrogen
sulphide in the presence of 4 c.c. of hydrochloric acid (sp. gr.
1.12), for which, approximately, [H^{+}] = 0.25. Then (equation
(IV), p. 201) [S^{2−}] = (1.1E−23) / (0.25)^2 = 1.8E−22, and
[Pb^{2+}] × [S^{2−}] = 1E−4 × 1.8E−22 = 1.8E−26, which is a much
lower value than that given by Knox, and which still is not claimed
to represent the limit of insolubility. Experiments, made in this
laboratory, confirm this result and show further, that lead-ion in
a concentration of 1E−5 is precipitated in the presence of 0.25
molar hydrochloric acid ([H^{+}] = 0.22). Then [S^{2−}] = 2.3E−22
and [Pb^{2+}] × [S^{2−}] = 2.3E−22 × 10^{−5} = 2E−27, which does
not yet express the limit of insolubility.
[419] The fractions are not prepared in the lecture, but the first
fraction is kept suspended in part of the solution of the two
sulphates and may be kept so for years. The last fraction is kept
in a separate container.
[420] A large excess of acid is liable to interfere with the
precipitation of the ‹last traces› of cupric sulphide and is
avoided in exact work.
[421] Noyes and Bray use, approximately, [H^{+}] = 0.25 [‹J. Am.
Chem. Soc.›, «29», 137 (1907)]. Tests in this laboratory showed
that 1 milligram of cadmium-ion, or of lead-ion, in 100 c.c., is
readily precipitated by hydrogen sulphide in the presence of 0.25
molar hydrochloric acid, ([H^{+}] = 0.22).
[422] Kahlbaum's "Krystallviolett,"
[(CH_{3})_{2}NC_{6}H_{4}]_{2}C : C_{6}H_{4}N(CH_{3})_{2}Cl, is
referred to.
[423] An indelible ink pencil (violet) may, in most cases, be used
in place of the solution. The details for the application of the
indicator are given in the instructions for laboratory practice,
Lab. Manual, pp. 31, 102, 103.
[424] See Blyth, ‹Poisons, etc.›, p. 608 (1895), in regard to the
detection of traces of lead.
[p216]