[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

As2S3 + (NH4)2S ⇄ 2 NH4AsS2.
(4)

The salt, ammonium sulpharsenite, is ionized as follows:

NH4AsS2 ⇄ NH4+ + AsS2−.

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 HAsS2 ⥂ As2 S3 ↓ + H2S ↑.

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, H2CS3, 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 + H2O, and KSH + HCl ⥂ KCl + H2S and K2S + 2 HCl ⥂ 2 KCl + H2S. 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 hydroxide496 (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 dioxide499 (equation (1), p. 243). The dissociation constant for the secondary ionization of hydrogen sulphide (HS− ⇄ H+ + S2−) 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 K2S 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, K2O, 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+ + O2−), which has a much smaller dissociation constant even than the primary ionization (H2O ⇄ H+ + HO−). The oxide, K2O, is decomposed by neutralizing hydrogen ions formed by the primary ionization of water. We have 2 K+ + O2− + 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 Me3AsO4, Me3AsSO3, Me3AsS2O2, Me3AsS3O and Me3AsS4. 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, S2−. Ammonium sulphide combines with stannic sulphide, forming ammonium sulphostannate: SnS2 + (NH4)2S ⇄ (NH4)2SnS3. 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: Sn4+ + 2 S2− + 2 NH4+ + S2− ⇄ 2 NH4+ + SnS32−.

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: Sn4+ + 3 S2− ⇄ SnS32−.

For the condition of equilibrium between the complex and its components we have:501,502

[Sn4+] × [S2−]3 / [SnS32−] = 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, (NH4)2S2, 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 + (NH4)2S2 ⇄ (NH4)2SnS3.

Behavior of Arsenic Acid toward Hydrogen Sulphide.

—In conclusion, special consideration must still be given to the behavior of arsenic acid, H3AsO4. As its name indicates, it is an acid, and, in fact, a rather strong acid, of the order of strength503 of phosphoric acid, H3PO4, 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, H2AsO4−, HAsO42− and AsO43−. 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 As3+ + 3 S2− ⥂ As2S3 ↓. 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 (As2S3) 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 H3AsO4 + 5 H2S ⥂ As2S5 ↓ + 8 H2O
(1)504
H3AsO4 + H2S ⥂ H3AsO3 + S ↓ + H2O
(2)505
2 H3AsO3 + 6 HCl + 3 H2S ⇄
2 AsCl3 + 3 H2O + 3 H2S ⥂
As2S3 ↓ + 6 HCl + 3 H2O
(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+ + AsO43− ⇄ H3AsO4 (+ H2O) ⇄ As(OH)5 ⇄ As5+ + 5 HO−. Further, the precipitation of As2S5 may be assumed to result, ultimately,510 from the action of the sulphide-ion S2− on the positive ion As5+ (2 As5+ + 5 S2− ⇄ As2S5 ↓). 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 salt511 AsCl5. It could thus greatly increase the concentration of the ion As5+ 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 AsO43− and increasing the concentration of the arsenic-ion As5+. 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 As2S5 by suppressing the arseniate-ion AsO43− and thus increasing the concentration of the hydroxide As(OH)5, available for ionization as a base and for the production of the ion As5+. But the further favorable effect of the hydrochloric acid, in converting the hydroxide into a salt AsCl5 and increasing thereby the concentration of As5+, would be very largely offset by the action of the acid in suppressing the [p250] sulphide-ion ([S2−] = k / [H+]2; see p. 201). For systems to which the equilibrium laws could be applied, the concentration of As5+ (except for the suppression of the ion AsO43−) would grow, approximately, with the fifth power512 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 As2S5, in a system to which the principle of the solubility-product were applicable, would depend on the relation of the product [As5+]2 × [S2−]5 to the solubility-product constant; it is evident that the value for [As5+]2 would increase proportionally to the tenth power of [H+] and the value of [S2−] 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 (H3AsO4), 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.

Chapter XIII 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, KHCO3 and KHCS3, 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 SnO32− of stannic acid, H2SnO3, we have:

[Sn4+] × [O2−]3 / [SnO32−] = 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 Me2HgS2, and the complex ion HgS22−. A liter of 0.1 molar Na2S 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. HgI42−, in K2HgI4, HgCl42− in K2HgCl4, Hg(CN)42−, etc.). Knox found for [Hg2+] × [S2−]2 / [HgS22−] = k, the approximate value of k to be 1 / 1053. Bismuth sulphide is also very sparingly soluble in sodium or potassium sulphide, but not in ammonium sulphide. Solid salts, KBiS2 and NaBiS2 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 H3AsSO3 (p. 246), could be the result of the ionization of As(OH)5, or of AsCl5, 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 AsCl5. McCay, J. Am. Chem. Soc., 24, 661 (1902), discusses the ionization of arsenic acid as a base, in connection with the precipitation of As2S5.

[512] As(OH)5 is considered to be an extremely weak base and AsCl5 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 OH3Cl, etc., see p. 238), which tends, by hydrolysis, to reverse the formation of the chloride As(OH)5 + 5 HCl ⇄ AsCl5 + 5 H2O. 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).

CHAPTER XIV OXIDATION AND REDUCTION REACTIONS. I

[p251] TOC

Oxidation and reduction reactions are frequently met with in analysis, and we shall turn now to the consideration of such reactions, from the point of view of the modern theory of solution and the laws of equilibrium.

Leaving until later the discussion of the most important and most common oxidizing agents, such as oxygen, nitric acid, permanganate, etc., we shall, in order to develop the subject most simply, confine ourselves, for the moment, to the qualitative study of some oxidations and reductions met with early in the study of analytical reactions. One such reaction is the reduction of ferric salts by hydrogen sulphide, and the simultaneous oxidation of the latter to sulphur (exp.). The reaction may be expressed by the equation

2 FeCl3 + H2S → 2 FeCl2 + 2 HCl + S ↓.

If the action is considered to be the result of the interaction of the ionized ferric chloride and hydrogen sulphide, it would be represented by the equation

2 Fe3+ + 6 Cl− + 2 H+ + S2− → 2 Fe2+ + 6 Cl− + 2 H+ + S ↓.

It is then clear that the reacting components, according to such a conception, are the ferric and the sulphide ions, whose electrical charges mutually discharge each other. Considering only those components whose charges are changed, we have

2 Fe3+ + S2− → 2 Fe2+ + S ↓.

The reduction of a ferric to a ferrous salt would then be accomplished by the discharge of one of the three positive charges on the ferric ions; the oxidation of hydrogen sulphide to sulphur would be accomplished by the complete discharge of the sulphide ions.

Ferric salts are reduced, much in the same way, by iodides (exp.), iodine being liberated: 2 Fe3+ + 2 I− → 2 Fe2+ + I2.

Reduction then appears to involve a loss of positive charges by ions, oxidation a loss of negative charges. [p252]

Conversely, we frequently have occasion to oxidize ferrous salts to the ferric condition, and among the most convenient reagents for the purpose are chlorine and bromine water (exp.). For instance, we have 2 FeCl2 + Cl2 → 2 FeCl3, or, considering the action from the point of view of the theory of ionization,514 2 Fe2+ + Cl2 → 2 Fe3+ + 2 Cl−. In this case the oxidation of the ferrous to the ferric ion consists in the assumption of an additional positive charge; reduction of chlorine to the chloride-ion consists in the assumption of negative charges by the chlorine atoms.

Definitions of Oxidation and Reduction in Electric Terms.

—The definitions must then be amplified and oxidation be considered to involve ultimately the assumption of positive, or the loss of negative, electrical charges by ions or atoms, reduction to involve ultimately the assumption of negative, or the loss of positive, charges. According to the electron theory of electricity a unit negative charge is an electron, the unit positive charge, probably, the charge left on an atom when it has lost an electron; and, thus, oxidation may simply be defined, according to the electric theory of oxidation and reduction, as consisting, fundamentally, in the loss of electrons by atoms or ions, reduction as consisting in a gain of electrons. For instance, when hydrogen sulphide reduces a ferric salt, 2 Fe3+ + S2− → 2 Fe2+ + S, the sulphide ions transfer their electrons to the ferric ions.

Oxidations and Reductions by Electric Currents.

—All of the oxidation and reduction reactions which have been discussed may, in fact, be effected by the use of an electric current in the place of chemical agents. Ferrous salts, for instance, are oxidized, at the positive pole, by the current to ferric salts:515
Fe2+ + ⊕ → Fe3+, or Fe2+ − ε− → Fe3+.

Ferric salts are reduced, at the negative pole, to ferrous salts:

Fe3+ + ⊖ → Fe2+, or Fe3+ + ε− → Fe2+.
[p253]

Exp.—A solution of ferrous chloride, freshly prepared from iron wire, or a freshly prepared solution of ferrous-ammonium sulphate, is placed in a very small beaker; a solution of ferric chloride, acidulated with hydrochloric acid to prevent subsequent complete reduction of the ferric-ion to iron, is brought into a similar beaker. A small amount (5 c.c.) of each solution is tested, the former with potassium thiocyanate and the latter with ferricyanide solution, to show the absence of perceptible quantities of ferric and ferrous ions in them, respectively. Platinum electrodes, consisting best of cylinders of platinum gauze, are introduced into the solutions, the solutions are connected by means of a "salt bridge" (a U-tube filled with a solution of sodium chloride and closed at both ends by plugs of filter paper), and a current of 0.2 ampere is passed through the system, the positive current entering the solution containing the ferrous salt. After the current has been allowed to pass for a minute or two, 5 c.c. is withdrawn, by a pipette, from the meshes of the positive electrode and tested with thiocyanate, and 5 c.c., withdrawn in the same way from the negative electrode, is tested with potassium ferricyanide.

Production of Electric Currents by Means of Oxidation and Reduction Reactions.

—Not only may all the oxidations and reductions which we have discussed be accomplished with the aid of the electric current, but vice versa, an electric current may readily be produced by a proper arrangement of the components of any one of these reactions. Care need only be taken to have the simultaneous loss and gain of electrons, characteristic of all oxidation and reduction reactions, occur in separate localities, which must be connected, however, in such a way as to make the transfer of electrons, a flow of electricity, possible.

Exp. For instance, some ferric chloride and sodium chloride solution may be put into a small beaker, some sodium chloride solution into a second beaker of the same size, and the two solutions connected, first by means of a "salt-bridge," and then by means of two platinum electrodes dipping into the solutions and connected with the terminals of a sensitive voltmeter.516 If [p254] all of the connections are made, the introduction of the "salt-bridge" being left to the last, a momentary slight motion of the needle is observed, when the bridge is introduced. The needle then falls back to the zero point (see p. 276). If now some hydrogen sulphide water is poured into the beaker containing sodium chloride, a decided, and continuing, deflection of the needle of the voltmeter is immediately observed, showing the passage of an electric current, and it is in the direction anticipated by the consideration of the reaction equation: 2 Fe3+ + S2− → 2 Fe2+ + S ↓. The positive current passes into the voltmeter from the ferric chloride solution, where ferric ions are giving up their charges; the negative current enters the voltmeter from the solution containing the hydrogen sulphide, where sulphide ions are being discharged. The "salt-bridge" is necessary to complete the electrical circuit and prevent any local accumulation of positive or negative electricity (polarization). For instance, as the ferric ions are discharged, an excess of chloride ions would remain in the beaker, rendering the solution negative and preventing the flow of electricity from the electrode, if negative ions did not move off, through the "salt-bridge," into the beaker containing hydrogen sulphide and, simultaneously, positive ions migrate into the beaker containing the ferric salt. Similarly, the accumulation of positive electricity in the hydrogen sulphide solution, on account of the hydrogen ions left free by the discharge of sulphide ions, is prevented by the flow of positive ions (sodium and hydrogen) through the U-tube into the beaker containing the ferric chloride and the flow of negative (chloride) ions into the hydrogen sulphide solution. Thus a current of electricity passes through the whole circuit.

It is thus possible to reduce ferric chloride in one vessel by hydrogen sulphide poured into another vessel,517 and an electric current may be obtained from the simultaneous discharge of the sulphide and ferric ions in the action.

Effects of Ion Concentrations on the Current.

—Hydrogen sulphide, it will be recalled, is an extremely weak acid (p. 199), only a very small proportion is ionized and, consequently, the concentration of the discharging (reducing) sulphide-ion must be minute in this solution. Its salts, however, are highly ionized, and by the addition of an alkali to the solution containing the hydrogen sulphide, the concentration of the discharging ion would be very greatly increased and the current should therefore be intensified most decidedly—provided the hydrogen sulphide really reduces by means of its negative ion and not by the action of the nonionized acid. As a matter of fact, the anticipated decided increase in the intensity of the current is observed, when alkali is added to the mixture containing the hydrogen sulphide (exp.). Similarly, [p255] we have assumed that the oxidizing agent is the highly charged ferric-ion, not the nonionized ferric salt. Now iron forms rather stable complex ions518 with the fluoride-ion, for instance, FeF63−, which yield ferric ions very much less readily than do ferric salts. Hence the addition of a fluoride—potassium or ammonium fluoride—should, according to this view, reduce the oxidizing power of the iron solution by suppressing the ferric-ion and converting it into the complex FeF63−. In fact, the addition of potassium fluoride immediately reduces the intensity of the current (exp.), and, simultaneously, the deep yellow-brown color of the ferric salt solution gives way to the very pale yellow tint of the complex ion and its salt.519

Further Illustrations.

—If ferrous sulphate solution is put into one beaker and sodium chloride solution into another, connections being made similar to those used in the previous experiment, then a vigorous current is instantly produced (exp.), when some bromine or chlorine water is added to the sodium chloride solution, the positive current flowing into the voltmeter from the beaker containing the bromine (chlorine); the bromine atoms (chlorine atoms) combine with electrons lost by the ferrous ions and are reduced to bromide ions (chloride ions) (see p. 252).

It would appear possible, in fact, to obtain an electrical current from any oxidation-reduction reaction, if the oxidizing and reducing agents can be, experimentally, properly arranged for this purpose.

Summary.

—We find thus that there is a most intimate connection between oxidation and reduction phenomena and electrical charges on atoms or ions. In the first place, an electrical current may be used as an oxidizing and reducing agent; indeed, a current cannot be passed through any solution without simultaneous oxidation and reduction at the positive and negative poles, respectively. And, conversely, an electric current may, in turn, be produced by a proper combination of the reagents in oxidation and reduction reactions.