«THE ARSENIC GROUP. SULPHO-ACIDS AND SULPHO-SALTS»
The analytical groups, which we have heretofore discussed, contain
elements, whose oxides are preëminently ‹base-forming›. The methods
of separation of these groups, from each other, involve, primarily,
physical[492] differences between the groups—in the matter of the
relative insolubility of analogous salts. Thus, barium, strontium
and calcium carbonates are precipitated, and separated from the
alkalies, by means of ammonium carbonate, not because the alkalies do
not form carbonates when their salts, in solution, are treated with
ammonium carbonate, but wholly because barium, strontium and calcium
carbonates are very difficultly, the alkali carbonates easily,
soluble in water. The hydroxides of the aluminium group and the
sulphides of the zinc group are less soluble than the hydroxides and
sulphides of the alkaline earths and alkalies. The sulphides of the
copper and the arsenic groups, again, are still less soluble than the
sulphides of the zinc group, and thus the former may be precipitated
by hydrogen sulphide, even when its precipitating power is reduced
by the suppression of its sulphide (and hydrosulphide) ions by the
addition of a strong acid.
On the other hand, the separation of the arsenic group (arsenic,
antimony, tin, gold and platinum) from the copper group, with
which it is precipitated by hydrogen sulphide from acid solutions,
depends, essentially, on a ‹chemical› difference between the
groups. The oxides, especially the ‹higher› oxides, of the arsenic
group, are preëminently ‹acid›-forming; the higher oxides form such
acids as arsenic acid, H_{3}AsO_{4}, antimonic acid, H_{3}SbO_{4},
stannic acid, H_{2}SnO_{3}, platinic acid, H_{2}PtO_{3}, and auric
acid, HAuO_{2}. These [p243] hydroxides are, however, all more
or less ‹weakly basic› in character ‹as well›. The hydroxides of
the lower oxides of the metals are, as one must expect, much more
strongly basic, but most of them—arsenious, antimonous and stannous
hydroxides—still show sufficient acid character to be distinctly
amphoteric in behavior. But, with the exception of arsenious acid,
the basic ionization of the hydroxides of the lower oxides is more
pronounced than their acid ionization.
The basic ionization of the hydroxides of their lower and higher
oxides brings these elements into the plan of analysis for the
metal or positive ions in systematic analysis. In the presence of
hydrochloric acid they form chlorides, which yield positive ions in
sufficient quantity[493] to allow their extremely insoluble sulphides
to be precipitated by hydrogen sulphide in acid solution, together
with the, likewise, very insoluble sulphides of the copper group.
The acid-forming properties of the oxides of the arsenic group are
maintained in their sulphides. Again, this is especially evident
in the ‹higher› sulphides. The element sulphur is substituted for
the closely related element oxygen without any profound change
in the chemical behavior of the compounds. Advantage is taken of
this acid-forming power to separate the sulphides of the arsenic
group from the sulphides of the copper group, which either are not
acid-forming at all, or exhibit this property only to a very slight
degree.[494]
«Sulpho-Salts.»—The similarity in the behavior of oxygen and sulphur
derivatives, in this respect, is general and is not restricted to the
metal sulphides we are discussing. For instance, the acid-forming
power of carbon dioxide is shown also by carbon disulphide, the
corresponding sulphur derivative of carbon. Just as the former
combines with potassium hydroxide to form a carbonate, so carbon
disulphide dissolves in solutions of potassium hydrosulphide and
potassium sulphide to form potassium sulpho-carbonate:
CO_{2} + 2 KOH ⇄ K_{2}CO_{3} + H_{2}O. (1)
CS_{2} + 2 KSH ⇄ K_{2}CS_{3} + H_{2}S. (2)
CS_{2} + K_{2}S ⇄ K_{2}CS_{3}. (3)
[p244]
The higher sulphides of the arsenic group, and arsenious and
antimonious sulphides among the lower ones, combine with the alkali
sulphides to form soluble alkali ‹salts› of sulpho-acids, in the same
way as carbon bisulphide does. In the case of arsenious sulphide, for
instance, we have the action
As_{2}S_{3} + (NH_{4})_{2}S ⇄ 2 NH_{4}AsS_{2}. (4)
The salt, ammonium sulpharsenite, is ionized as follows:
NH_{4}AsS_{2} ⇄ NH_{4}^{+} + AsS_{2}^{−}.
«Sulpho-Acids.»—The free ‹sulpho-acids› are liberated, from their
salts, by any stronger acid, such as hydrochloric acid; the acids are
extremely unstable and revert rapidly to the sulphides, from which
their salts were originally obtained. We have, for instance,
2 HAsS_{2} ⥂ As_{2} S_{3} ↓ + H_{2}S ↑.
This instability of the sulpho-acids is entirely analogous to
the instability of the metal hydrosulphides (p. 203) and to the
instability of certain oxygen acids, notably of carbonic acid.
Sulpho-carbonic acid, H_{2}CS_{3}, is the best-known free acid of
this type. It may be precipitated, undecomposed, as an oil, and its
gradual decomposition into hydrogen sulphide and carbon disulphide
may be observed.
«Sulpho-Bases.»—Since one meets, in this group, ‹sulpho-acids› and
‹sulpho-salts›, corresponding to the oxygen acids and their salts,
one is naturally led to inquire,[495] whether the third great
class of oxygen compounds is not found duplicated among sulphur
compounds, whether ‹sulpho-bases›, as well as sulpho-acids and
salts, are known. One would look for such bases, with the most
pronounced basic character, among the compounds obtained by the
substitution of sulphur for oxygen in the strongest oxygen bases. To
a certain extent potassium and sodium hydrosulphides and sulphides
show, in fact, properties, which are akin to those fundamentally
characteristic of ordinary bases. They combine with acid sulphides to
form sulpho-salts, as the oxygen bases combine with acid oxides. To
a very considerable extent they neutralize all but the very weakest
acids; hydrogen sulphide, itself a very weak acid, is driven out of
its salts by all stronger acids, and the latter are almost completely
neutralized. For instance, we have: [p245] KOH + HCl ⥂ KCl + H_{2}O,
and KSH + HCl ⥂ KCl + H_{2}S and K_{2}S + 2 HCl ⥂ 2 KCl + H_{2}S. The
concentration of the hydrogen-ion is reduced most decidedly in each
of these reactions.
EXP.—A solution of potassium hydrosulphide is saturated with
hydrogen sulphide, in order to prevent hydrolysis and the formation
of potassium hydroxide[496] (p. 180), as far as possible, and the
solution is added to an acid (hydrochloric) solution of methyl
orange;[497] the acid color is changed to orange, as a result of
the almost complete neutralization of the acid. The potassium
hydrosulphide (the hydrosulphide-ion HS^{−}) neutralizes the
hydrogen-ion (of hydrochloric acid), that converts methyl orange
into its pink salt, and hydrogen sulphide is formed, which is too
weak an acid to affect the color of the indicator (p. 79).
The objection that potassium sulphide and hydrosulphide are salts,
the salts of hydrogen sulphide, might be raised against the
conception of their possessing a certain measure of basic functions;
but the common oxygen base, potassium hydroxide, is also a salt,
the salt of a still weaker acid, water. Indeed, the characteristic
properties of ordinary bases are due essentially to the fact, that
they are the more or less readily ionizable salts of an extremely
weak acid, water, and these properties may well be duplicated by
salts of other ‹weak› acids, duplicated in a ‹very much weaker› way,
in proportion as the acids are stronger than water. The difference
is, then, really one of ‹degree› and not of kind.[498]
Owing to the fact that hydrogen sulphide is a much stronger
acid than water, the action of potassium hydrosulphide on an
acid sulphide, like carbon disulphide (equation (2), p. 243),
is reversed to a correspondingly greater degree than the action
of potassium hydroxide on carbon dioxide[499] (equation (1), p.
243). The dissociation constant for the secondary ionization
of hydrogen sulphide (HS^{−} ⇄ H^{+} + S^{2−}) is very much
smaller than the constant for the primary ionization (HS^{−} is
a much weaker acid than HSH), and so we find that a sulphide
like K_{2}S exhibits very much stronger basic functions than
do the hydrosulphides, as, for instance, in forming salts with
acid-forming sulphides [p246] (equation (3), p. 243) and in
neutralizing acids. There can be no question that, if we could have
an aqueous solution of potassium oxide, K_{2}O, it would show,
similarly, the characteristic actions of strong bases ‹even more
powerfully› than the hydroxide, KOH; for instance, in acting on
acid-forming oxides (equation (1), p. 243), in neutralizing acids,
in saponifying esters (p. 81), and so forth. It is, in fact, on
account of this property, that potassium oxide is decomposed by
water. It is a salt involving the ‹secondary ionization› of water,
(HO^{−} ⇄ H^{+} + O^{2−}), which has a much smaller dissociation
constant even than the primary ionization (H_{2}O ⇄ H^{+} + HO^{−}).
The oxide, K_{2}O, is decomposed by ‹neutralizing hydrogen
ions formed by the primary ionization of water›. We have
2 K^{+} + O^{2−} + H^{+} + HO^{−} ⥂ 2 K^{+} + 2 HO^{−},
which is entirely analogous, in principle, to K^{+} + HO^{−} +
H^{+} + Cl^{−} ⥂ K^{+} + Cl^{−} + HOH.
«Sulphoxy-Salts.»—The close relations between the oxygen and the
sulphur series are seen also in the fact that an oxygen base may
be combined with an acid sulphide, and ‹vice versa›; arsenious
sulphide, for instance, dissolves even in the solution of so weak
a base as ammonium hydroxide (‹exp.›). The salts produced by this
"crossing" are usually "hybrid" salts, partly sulpho-, partly
oxygen-salts. There is, for instance, a series of ‹arseniates›,[500]
Me_{3}AsO_{4}, Me_{3}AsSO_{3}, Me_{3}AsS_{2}O_{2}, Me_{3}AsS_{3}O
and Me_{3}AsS_{4}. In analytical work the pure types are ordinarily
utilized, rather than the mixed types.
«Complex Sulphide Ions.»—The ions of the sulpho-acids, like
the ions of oxygen-acids (p. 238), may also be treated as
‹complex ions›—of the positive metal ions and the sulphide-ion,
S^{2−}. Ammonium sulphide combines with stannic sulphide,
forming ammonium sulphostannate: SnS_{2} + (NH_{4})_{2}S ⇄
(NH_{4})_{2}SnS_{3}. If the action is considered to be the
result of interactions of the ions of stannic and ammonium
sulphides, we can resolve the equation into the following one:
Sn^{4+} + 2 S^{2−} + 2 NH_{4}^{+} + S^{2−} ⇄ 2 NH_{4}^{+} +
SnS_{3}^{2−}.
The ammonium-ion, appearing with the same coefficient on both sides
of the last equation, evidently takes no direct part in the action
and we have more simply: Sn^{4+} + 3 S^{2−} ⇄ SnS_{3}^{2−}.
For the condition of equilibrium between the complex and its
components we have:[501][502]
[Sn^{4+}] × [S^{2−}]^3 / [SnS_{3}^{2−}] = K.
[p247]
«Sulphurization of Sulphides.»—Since the solubility of the arsenic
group of sulphides in ammonium sulphide solution—the reagent commonly
used in analysis—depends on the formation of soluble sulpho-salts,
due consideration must be taken of the fact that some of the ‹lower›
sulphides—notably stannous, aurous and platinous sulphides—do not
possess acid-forming properties in any marked degree; even antimonous
sulphide is soluble only in a considerable excess of ammonium
sulphide. In order, then, to insure a more complete separation of the
arsenic from the copper group, precautions are taken to sulphurize
the lower sulphides to higher, stronger acid-forming, sulphides,
in the course of the separation. For this purpose so-called
"yellow" ammonium sulphide, containing persulphides of ammonium,
(NH_{4})_{2}S_{2}, etc., is used in place of a solution of ammonium
sulphide and ammonium hydrosulphide. Stannous sulphide, for instance,
is dissolved by the reagent as ammonium sulphostannate: SnS +
(NH_{4})_{2}S_{2} ⇄ (NH_{4})_{2}SnS_{3}.
«Behavior of Arsenic Acid toward Hydrogen Sulphide.»—In conclusion,
special consideration must still be given to the behavior of arsenic
acid, H_{3}AsO_{4}. As its name indicates, it is an acid, and,
in fact, a rather strong acid, of the order of strength[503] of
phosphoric acid, H_{3}PO_{4}, which it resembles in composition and
in many of its properties. As a strong acid, arsenic acid, when it is
ionized, yields chiefly negative arseniate ions, H_{2}AsO_{4}^{−},
HAsO_{4}^{2−} and AsO_{4}^{3−}. Any basic properties, which it may
and, most likely, does possess, must be extremely weak. It is,
therefore, not surprising to find that unusual difficulties are
experienced in precipitating arsenic sulphide, by hydrogen sulphide,
from arseniate [p248] solutions, in as much as hydrogen sulphide
is an agent for the precipitation of the sulphides of ‹cations›.
Arsenious acid, the hydroxide of the lower oxide, on the other hand,
is a much weaker acid and shows more pronounced basic (amphoteric)
properties, and arsenic trisulphide is precipitated, without
difficulty, from solutions of arsenious acid in hydrochloric acid.
We have: 2 As^{3+} + 3 S^{2−} ⥂ As_{2}S_{3} ↓. When arsenic acid is
reduced to arsenious acid by sulphurous acid, by iodides (see Chap.
XVI) or by hydrogen sulphide (see below), no further difficulty in
precipitating a sulphide (As_{2}S_{3}) is experienced.
When a solution of arsenic acid, containing the usual small amount of
hydrochloric acid (0.3 molar), is treated with hydrogen sulphide at
ordinary temperatures, the following three reactions take place, but
‹exceedingly slowly›:
2 H_{3}AsO_{4} + 5 H_{2}S ⥂ As_{2}S_{5} ↓ + 8 H_{2}O (1)[504]
H_{3}AsO_{4} + H_{2}S ⥂ H_{3}AsO_{3} + S ↓ + H_{2}O (2)[505]
2 H_{3}AsO_{3} + 6 HCl + 3 H_{2}S ⇄ 2 AsCl_{3} + 3 H_{2}O +
3 H_{2}S ⥂ As_{2}S_{3} ↓ + 6 HCl + 3 H_{2}O (3)
Even in the presence of a considerable amount of arsenic acid,
precipitation, either of the trisulphide or of the pentasulphide, may
not occur for some time, and, unless one takes account of that fact,
the dangerous element, arsenic, would ‹easily› be ‹overlooked›. ‹Heat
accelerates› both the precipitation of the pentasulphide and the
reduction of arsenic acid and the subsequent precipitation of arsenic
trisulphide.[506]
The interesting observation has also been made that, in the presence
of an unusually ‹large excess› of hydrochloric acid and of a rapid
stream of hydrogen sulphide, the precipitation of the ‹pentasulphide›
(equation (1)) is favored and accelerated.[507] For instance, if 100
c.c. of concentrated hydrochloric acid (sp. gr. 1.2) are added to
50 c.c. of a 0.1 molar solution of potassium arseniate and a rapid
stream of hydrogen sulphide is passed through the mixture at the
ordinary temperature, a copious precipitate is formed within a minute
(‹exp.›). The precipitate formed under these conditions [p249] is
the ‹pentasulphide›.[508] On the other hand, a mixture of 5 c.c.
of hexanormal hydrochloric acid and 50 c.c. of 0.1 molar potassium
arseniate fails, for a long time, to give a precipitate when treated
in the same way (‹exp.›).
The acceleration of the precipitation of the pentasulphide by the
presence of a large excess of hydrochloric acid forms a problem of
peculiar interest and importance, and no complete explanation of
it has yet been offered.[509] The following considerations lead
to one explanation, that has been suggested. Arsenic acid, by
virtue of its close relations to antimonic, stannic and arsenious
acids, may be assumed to have extremely weak basic, as well as
pronounced acid, properties. For its ionization, we would have
3 H^{+} + AsO_{4}^{3−} ⇄ H_{3}AsO_{4} (+ H_{2}O) ⇄ As(OH)_{5} ⇄
As^{5+} + 5 HO^{−}. Further, the precipitation of As_{2}S_{5}
may be assumed to result, ultimately,[510] from the action
of the sulphide-ion S^{2−} on the positive ion As^{5+}
(2 As^{5+} + 5 S^{2−} ⇄ As_{2}S_{5} ↓). The favorable action of the
hydrochloric acid might, consequently, be thought to result from
the fact, that it ‹facilitates› the ‹ionization of arsenic acid›
as a base and the formation of a salt[511] AsCl_{5}. It could thus
greatly increase the concentration of the ion As^{5+} and facilitate
its combination with the sulphide-ion.
Treatment of a solution of arsenic acid with a concentrated acid,
yielding a large concentration of hydrogen-ion, would carry the
series of actions, represented in the above ionization equation
for arsenic acid, decidedly toward the right—suppressing the
arseniate-ion AsO_{4}^{3−} and increasing the concentration of
the arsenic-ion As^{5+}. Since we cannot apply the equilibrium
laws (or the principle of the solubility-product) to solutions
as concentrated as the one under discussion, a quantitative
theoretical treatment of the subject cannot be given. The following
may be suggested: The action of the acid would be favorable to
the precipitation of As_{2}S_{5} by suppressing the arseniate-ion
AsO_{4}^{3−} and thus increasing the concentration of the hydroxide
As(OH)_{5}, available for ionization as a base and for the
production of the ion As^{5+}. But the further favorable effect
of the hydrochloric acid, in converting the hydroxide into a salt
AsCl_{5} and increasing thereby the concentration of As^{5+}, would
be ‹very largely offset› by the action of the acid in suppressing
the [p250] sulphide-ion ([S^{2−}] = ‹k› / [H^{+}]^2; see p. 201).
For systems to which the equilibrium laws could be applied, the
concentration of As^{5+} (except for the suppression of the ion
AsO_{4}^{3−}) would grow, approximately, with the fifth power[512]
of the concentration of the hydrogen-ion, and the concentration of
the sulphide-ion would decrease, approximately, proportionally to
the square of the concentration of the hydrogen-ion. Further, the
precipitation of As_{2}S_{5}, in a system to which the principle of
the solubility-product were applicable, would depend on the relation
of the product [As^{5+}]^2 × [S^{2−}]^5 to the solubility-product
constant; it is evident that the value for [As^{5+}]^2 would
‹increase› proportionally to the tenth power of [H^{+}] and the
value of [S^{2−}] ‹decrease› proportionally to the tenth power of
the same factor [H^{+}]. The two effects would consequently offset
each other under such conditions. However, the equilibrium laws
cannot legitimately be applied to such concentrated solutions and
the relation has been developed only to indicate opposing factors,
which must be taken into account. An experimental study of the
problem would be extremely interesting.[513] Since it involves the
question of the ‹minute basic ionization of a moderately strong
acid› (H_{3}AsO_{4}), which may be open to ‹measurement› (see Chap.
XVI), the problem is one of particular interest and importance.
The analytical precautions, taken to insure the precipitation, by
hydrogen sulphide, of arsenic sulphide, when arsenic is present
in quinquivalent form, are based on the observations described;
in quantitative analysis, for the sake of securing a precipitate
of ‹uniform composition›, the aim is to precipitate the pure
‹pentasulphide› and a considerable excess of hydrochloric acid
is used. In qualitative analysis, where the composition of the
precipitate is a matter of indifference and a large excess of acid
would seriously interfere with the precipitation of certain sulphides
(‹e.g.› CdS, see p. 211), a smaller excess of acid is used and the
precipitation of arsenic sulphide is insured by prolonged treatment
of a solution with hydrogen sulphide ‹at a high temperature›.
FOOTNOTES:
[492] The ‹weak basic properties› of the hydroxides of the
aluminium group, as compared with those of the zinc group, ‹a
chemical difference›, and the resulting great instability of
the carbonates of the former group, are used in the separation
of the aluminium from the zinc group, by barium carbonate; but
the physical element of extreme insolubility of the trivalent
hydroxides enters also as an important factor (see footnote 3, p.
194).
[493] See p. 247, in regard to the behavior of arsenic acid in this
respect.
[494] See p. 246, footnote 3, in regard to the action of sodium
sulphide on mercuric and bismuth sulphide.
[495] ‹Vide› Nilson, ‹J. prakt. Ch.›, «14», 150 (1876).
[496] Such a solution does not react alkaline to phenolphthaleïn.
[497] Hydrogen sulphide rapidly destroys the indicator and the
experiment is best carried out by preparing 50 c.c. of a saturated
aqueous solution of hydrogen sulphide, containing 1 or 2 c.c. of
normal hydrochloric acid, and by adding a considerable excess of
methyl orange to the solution immediately before the addition
of potassium hydrosulphide solution, which has been prepared as
described in the text.
[498] See the discussion on p. 177. See also the discussion by
Remsen on acidic and basic halides, ‹Am. Chem. J.›, «11», 300
(1889) Stud.
[499] In both cases acid salts, KHCO_{3} and KHCS_{3}, are also
formed.
[500] McCay, ‹Z. anorg. Chem.›, «29», 36 (1901).
[501] On p. 238 the analogous equation for the condition of
equilibrium of the anion of an oxygen acid with its components was
developed. Applying the result to the ion SnO_{3}^{2−} of stannic
acid, H_{2}SnO_{3}, we have:
[Sn^{4+}] × [O^{2−}]^3 / [SnO_{3}^{2−}] = K.
It is evident, from the form of the equation, that for the stronger
oxygen acids, ‹which are most stable as acids and ionize as bases
at most in traces›, the value of the constant must be extremely
small.
[502] Mercuric sulphide is somewhat soluble in potassium and
sodium sulphides, forming the salts Me_{2}HgS_{2}, and the
complex ion HgS_{2}^{2−}. A liter of 0.1 molar Na_{2}S dissolves,
at 25°, 1.9 grams (0.0082 mole) of HgS [Knox, ‹Trans. Faraday
Society›, «4», 36 (1908)]. While the oxide (hydroxide) shows no
perceptible tendency toward acid ionization, mercuric salts, it
will be recalled, show in many cases an abnormally small tendency
to form the mercuric-ion (see p. 115), and the latter also
shows a particularly great tendency towards forming very stable
complex ions of all kinds (‹e.g.› HgI_{4}^{2−}, in K_{2}HgI_{4},
HgCl_{4}^{2−} in K_{2}HgCl_{4}, Hg(CN)_{4}^{2−}, etc.). Knox found
for [Hg^{2+}] × [S^{2−}]^2 / [HgS_{2}^{2−}] = ‹k›, the approximate
value of ‹k› to be 1 / 10^{53}. Bismuth sulphide is also very
sparingly soluble in sodium or potassium sulphide, but not in
ammonium sulphide. Solid salts, KBiS_{2} and NaBiS_{2} are known
[Knox, ‹J. Chem. Soc.› (London), «95», 1760 (1909)].
[503] See the table, p. 104.
[504] See the discussion of the reaction, given below.
[505] See Chap. XVI, for the interpretation of the reduction as an
ionic reaction.
[506] Bunsen, ‹Ann.› (Liebig), «192», 305 (1878). Brauner and
Tomicek, ‹J. Chem. Soc.› (London), «53», 145 (1888). Usher and
Travers, ‹ibid.›, «87», 1370 (1905).
[507] Neher, ‹Z. anal. Chem.›, «32», 45 (1893).
[508] Neher, ‹loc. cit.›
[509] The theory of the relations favoring the precipitation
expressed in equation (1) as against the reduction expressed in
equation (2), forms a second interesting problem.
[510] Intermediate derivatives, such as H_{3}AsSO_{3} (p. 246),
could be the result of the ionization of As(OH)_{5}, or of
AsCl_{5}, in stages (see p. 106). Neher, ‹loc. cit.›, McCay, ‹loc.
cit.›
[511] Neher (‹loc. cit.›) suggested that the favorable action
of the large excess of hydrochloric acid might well be due to
the formation of AsCl_{5}. McCay, ‹J. Am. Chem. Soc.›, «24», 661
(1902), discusses the ionization of arsenic acid as a base, in
connection with the precipitation of As_{2}S_{5}.
[512] As(OH)_{5} is considered to be an ‹extremely weak› base and
AsCl_{5} to be an ionizable salt.
[513] To a certain extent, the effect of the acid may be to
coagulate and precipitate the ‹colloidal› sulphide. Possibly, also,
the concentrated acid renders inactive a considerable portion of
the water present (forming oxonium salts OH_{3}Cl, etc., see p.
238), which tends, by hydrolysis, to reverse the formation of the
chloride As(OH)_{5} + 5 HCl ⇄ AsCl_{5} + 5 H_{2}O. Possibly, the
formation of the pentasulphide is not wholly an ‹ionic reaction›,
its precipitation being always a more or less ‹slow process›,
and there may be intermediate products whose formation could be
‹accelerated› by the presence of acids (see Bredig and Walton, ‹Z.
Elektrochem.›, «9», 114 (1903) for the study of a simple inorganic
action involving such ‹catalytic› effects of acids).
[p251]