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Dithionic Acid, H2S2O6

Dithionic Acid, H2S2O6, is obtained only in aqueous solution. If sulphur dioxide is passed into an aqueous suspension of ferric hydroxide at 0° C., a red solution of ferric sulphite is first produced, which then changes to a pale green solution of ferrous sulphite and ferrous dithionate:

2Fe(OH)3 + 3SO2 = Fe2(SO3)3 + 3H2O,
Fe2(SO3)3 = FeSO3 + FeS2O6.

By the addition of barium hydroxide to the solution, a solution of barium dithionate can be obtained which, on treatment with the necessary quantity of sulphuric acid, yields a dilute solution of dithionic acid; this can be concentrated to some extent by evaporation at the ordinary temperature, but if the concentration is carried too far, decomposition ensues, with formation of sulphur dioxide and sulphuric acid. If so desired, the barium dithionate itself may be isolated by removing the excess of barium hydroxide with a current of carbon dioxide and then allowing the solution to evaporate until crystallisation occurs.

In the foregoing process it will be seen that the ferric hydroxide, in effect, acts as an oxidising agent. Other oxidising agents are also applicable, but the results are less satisfactory than with ferric hydroxide.

The process by which Gay-Lussac and Welter discovered dithionic acid was the oxidation of sulphurous acid with manganese dioxide; the reaction is frequently represented as

MnO2 + 2SO2 = MnS2O6,

but in reality it is more complex and the dithionate is always accompanied by a preponderating quantity of sulphate. Although it is possible that the dithionate and sulphate are produced by two independent concurrent reactions, the dithionate reaction being favoured by a low temperature whilst rise in temperature favours the production of sulphate, the more probable course of the change is the primary formation of manganic sulphite, which subsequently decomposes, giving manganous sulphite and dithionate, in a manner analogous to that described with ferric hydroxide:

2MnO2 + 3H2SO3 = Mn2(SO3)3 + 3H2O + O,
Mn2(SO3)3 = MnSO3 + MnS2O6,
MnSO3 + O = MnSO4.

This view is favoured by the observation that not only does ferric hydroxide give a similar result, but manganese trihydroxide and cobalt trihydroxide also produce a small quantity of dithionate, whereas the peroxides, which generally are derived from metals of unvarying valency, oxidise sulphurous acid only to sulphuric acid or a sulphate. Lead dioxide reacts but very slightly with sulphurous acid, probably on account of the sparing solubility of the products, which coat the particles of dioxide and prevent further action. Chromates and dichromates in acid solution convert sulphurous acid to a small extent into dithionic acid, and again it will be noticed that the metal, in this case chromium, is of variable valency. Permanganates give a similar result with sulphurous acid, and even gradual atmospheric oxidation of the acid or ammonium hydrogen sulphite solution will give rise to some dithionic acid.

Alkali sulphites in boiling aqueous solution dissolve selenium with formation of a selenotrithionate, which subsequently decomposes giving the alkali dithionate:

K2S2SeO6 = K2S2O6 + Se.

Contrary to earlier views, alkali dithionate is not obtainable by the action of iodine on the corresponding alkali hydrogen sulphite, sulphate being the sole product given. Conversion of sulphite into dithionate, however, can be effected by electrolysis, the results not being due to the coupling of discharged sulphite ions but to anodic oxidation of the sulphite, sulphate also being formed:

2Na2SO3 + O + H2O = Na2S2O6 + 2NaOH.

The electrodes should be of smooth platinum, and the yield is increased if the anode is heated for a short time immediately before use. Previous anode polarisation is also advantageous. Addition of ammonium fluoride, 0.1 per cent, maximal, increases the yield when the anode has not been preheated, but otherwise has an adverse effect. Under favourable conditions a yield of 45 per cent, of dithionic acid may be obtained, the maximum possible from theoretical considerations being about 50 per cent.

When silver sulphite or a mixture of sodium sulphite and silver nitrate is heated in boiling aqueous solution, decomposition occurs with formation of silver dithionate.

Thiosulphates, e.g. sodium thiosulphate, are oxidisable to polythion- ates by the addition of the requisite quantity of hydrogen peroxide. If the reaction mixture is allowed to become alkaline, dithionate, tetra-thionate and sulphate are produced:

4Na2S2O3 + 8H2O2 = 2Na2SO4 + Na2S2O6 + Na2S4O6 + 8H2O.

Oxidation of thiosulphate to dithionate can also be effected in dilute acetic acid solution with potassium permanganate; in alkaline solution, sulphate is the only product.

The acid is known only in aqueous solution and in the form of its salts, the dithionates or hypo sulphates, the latter name now being infrequently used.

The aqueous solution is odourless, but possesses an acid taste; its electrical conductivity shows that dithionic acid is to be classed amongst the strong acids. The heat of formation of aqueous dithionic acid relative to its elements and the solvent is represented by the equation:

H2 + 2S + 3O2 + Aq. = H2S2O6,Aq. + 279.4 Calories.

The molecular weight and basicity of the acid were at one time in question, the fact that no acid salts are formed giving the incorrect impression that the acid was monobasic. By Ostwald's method for the determination of the basicity of an acid from the increase in the molecular conductivity of an aqueous solution of the sodium salt on dilution, and by measurement of the molecular weights of the salts in aqueous solution, it was subsequently demonstrated that the acid is dibasic and of the double molecular formula H2S2O6.

In aqueous solution at the ordinary temperature, the concentration of the acid cannot be raised beyond that corresponding with a specific gravity of 1.35, on account of decomposition according to the equation

H2S2O6 + H2O = H2SO4 + H2SO3,

which occurs at any concentration on heating. In dilute solutions (3 to 4 per cent.) the decomposition proceeds to the extent of 3 per cent, in 945 hours at 25° C., and 20 per cent, in 245 hours at 47° C.

Dithionic acid solutions are remarkably resistant to oxidation; the cold solutions withstand the attack of hypochlorite, hypobromite and permanganic acid, although on boiling, these reagents become reduced by the sulphur dioxide liberated in the decomposition of the dithionic acid; sodium peroxide effects a partial oxidation in the cold. Reduction by sodium amalgam or by zinc and an acid produces sulphurous acid:

H2S2O6 + 2H = 2H2SO3.

Dithionates or Hyposulphates

These salts, like the nitrates, are all soluble in water; only normal salts are known. They may be obtained by neutralising dithionic acid solution with the hydroxide of the base, and also by double decomposition between barium dithionate solution and the sulphate of the base, or between manganese dithionate solution and the hydroxide of the base; Sodium ethyl thiosulphate, , decomposes slowly at 100° C. with the formation of ethyl disulphide and sodium dithionate

Solubility of ditionates in water at 20° C.

Salt.Grams of Salt in 100 grams Solution.
Na2S2O6.2H2O .13.39

Of the dithionates, those of the alkali and alkaline earth metals are the most stable and may be heated in aqueous solution up to 100° C. without decomposition. When heated alone, the dithionates decompose readily with formation of sulphate and sulphur dioxide; some of them indeed are so unstable that they are not obtainable in a pure condition.

Mineral acids decompose dithionate solutions only on boiling, sulphate and sulphur dioxide then being produced; for this reason permanganate is decolorised only by hot acid solutions. Iodine solution is slowly decolorised, due to the gradual formation of sulphurous acid, the final product being an acid sulphate.

In the presence of such reagents as bromate, iodate or dichromate, the rate of oxidation is independent of the nature or concentration of the oxidising agent, but is the same as the rate of decomposition to sulphate and sulphite, so that it is evident that hydrolysis is the first stage in the oxidation.

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