Atomistry » Sulphur » Compounds » Hydrosulphites
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Hydrosulphites are much more stable than the free acid; the best known salts are those of the alkali and the alkaline earth metals, for example Na2S2O4.2H2O, K2S2O4.3H2O, CaS2O4.1½H2O. As will be seen from these formulae, the acid is dibasic, but only normal salts are known. The zinc salt is also fairly stable, especially in the form of double salts with the alkali hydrosulphites, for example Na2Zn(S2O4)2. Crystalline hydrosulphites of aromatic amines have also been obtained by interaction of the sodium salt with the hydrochloride of the amine.

The hydrated salts obtained by ordinary crystallisation from water are relatively unstable, and tend in the solid condition or in aqueous solution with exclusion of air to decompose into thiosulphate and sulphite (the latter as pyrosulphite or acid sulphite, according to the conditions); the further interaction of these primary products complicates the final result. In the case of the sodium salt in aqueous solution the reaction is bimolecular and follows the scheme

2Na2S2O4 = Na2S2O3 + Na2S2O5,
Na2S2O5 + H2O = 2NaHSO3.

In the presence of alkali, hydrosulphites react with alkali polysulphides with the formation of sulphite and sulphide:

Na2S2O4 + Na2S2 + 4NaOH = 2Na2S + 2H2O + 2Na2SO3,

a similar result being obtained with thiosulphates, although more slowly:

Na2S2O4 + Na2S2O3 + 4NaOH = Na2S + 2H2O + 3Na2SO3.

As reducing agents the hydrosulphites are exceptionally active and generally become oxidised to sulphites, although with relatively strong oxidising agents such as hydrogen peroxide or iodine, sulphate may be formed.

Ferric salts are reduced to the ferrous condition and -chromates are reduced to chromic salts. Salts of gold, silver, copper, antimony, bismuth and mercury are reduced to the free metals, which are frequently obtained as colloidal solutions if the original solutions are weak; with the exception of the first-named there is a tendency for the liberated metal to be accompanied by sulphide, especially if excess of hydro- sulphite is used. Chloroplatinic acid is reduced to red chloroplatinous acid solution. Tellurous and telluric acids, selenious acid and arsenic compounds, are reduced to the free elements.

Indigo and many other colouring matters are bleached by sodium hydrosulphite. Organic nitro-compounds are reduced to amino-compounds, the group -NO2 being converted into -NH2.

Commercially the anhydrous sodium salt is manufactured in large quantities on account of its greater stability than the hydrated Na2S2O4.2H2O. The latter can be dehydrated by extraction with warm alcohol or acetone at 60° to 70° C. It is also possible to produce the anhydrous salt directly by precipitation of the aqueous solution with alcohol at 60° to 70° C., or even by "salting out" the solution at this temperature by the addition of sodium hydroxide, chloride, sulphate, carbonate, nitrate or acetate; with sodium hydroxide solution of 50 per cent, concentration the anhydrous salt can be separated even at 20° C.

The hydrosulphites are very liable to atmospheric oxidation, especially in the hydrated or moist condition; in their preparation and preservation it is therefore desirable to exclude air. When exposed to the atmosphere a solution of a hydrosulphite first becomes yellow, probably due to the transient presence of free hydrosulphurous acid, but soon becomes colourless, the final product being sulphite. Much heat is evolved during this oxidation, as may readily be observed with the exposed moistened salt:

2Na2S2O4 + O2 + 2H2O = 4NaHSO3.

With hydrogen sulphide the aqueous solution reacts with formation of thiosulphate and sulphur, the equation being

Na2S2O4 + H2S = Na2S2O3 + H2O + S.

In feebly acid solution the salts of such metals as nickel, cobalt, lead, zinc and cadmium react with hydrosulphites yielding the corresponding metal sulphide and no free metal.

With aldehydes and ketones sodium hydrosulphite readily forms additive compounds, the most important being that derived from formaldehyde. This product appears to have the composition 2CH2O.Na2S2O4.4H2O, but is separable by recrystallisation from water into the sodium hydrogen sulphite derivative of formaldehyde, viz. CH2O.NaHSO3.H2O, and an analogous compound CH2O.NaHSO2.2H2O, a crystalline solid of 63° to 64° C. The latter is known as " Rongalite " and is of especial commercial importance on account of its stability at the ordinary temperature, although at steam heat it exerts all the reducing power of the hydrosulphites; on this account " Rongalite," or sodium formaldehydesulphoxylate, is a very convenient form of reducing agent where storage for prolonged periods may be necessary before use. The aqueous solution may be stabilised by addition of a soluble zinc salt. Various other methods have been recommended for the preparation of " Rongalite," for example, the interaction of hydrosulphite and formaldehyde in the presence of an

Na2S2O4 + CH2O + NaOH = CH2O.NaHSO2 + Na2SO3,

and the reduction of sodium hydrogen sulphite solution with zinc dust and zinc oxide in the presence of formaldehyde, recrystallising from water at a temperature not exceeding 70° C. the crystals first obtained.


Sodium hydrosulphite and sodium formaldehydesulphoxylate ("Rongalite") are largely applied industrially, generally for bleaching purposes. Sugar, soap, straw, etc., can be improved in colour by treatment with these reagents. In the dyeing industry also these substances are used, for example to convert insoluble dyes like indigo into a soluble form, by which means a vat suitable for dyeing can be obtained; the reducing agent may also be applied locally to a dyed cloth and a pattern thereby produced.


Although the formation of sodium formaldehyde- sulphoxylate from sodium hydrosulphite (see before) might appear to indicate that the latter is a mixture of sodium hydrogen sulphite with sodium hydrogen sulphoxylate, NaHSO2, it is generally accepted that the hydrosulphites are definite compounds; a simple experimental observation in favour of this view is that a slightly alkaline solution of sodium hydrosulphite becomes acid on atmospheric oxidation, whereas an alkaline solution of sodium sulphite and sodium sulphoxylate should produce only normal sulphite and sulphate. The sulphoxylates are little known except for their formaldehyde derivatives, only the sodium and zinc salts having been prepared. Except under exceptional conditions, therefore, the hydrosulphites are much more stable than the sulphoxylates.

Schutzenberger ascribed to sodium hydrosulphite the formula NaHSO2 which, curiously enough, was also one of the earliest formulae suggested for sodium thiosulphate. Oxidation processes are of assistance in indicating the nature of hydrosulphurous acid; thus when oxidised to the stage of sulphite, for example by ammoniacal copper sulphate solution, one atomic weight of oxygen is required per molecule of acid, whilst on oxidation to the stage of sulphate, which can be effected by iodine or by ammoniacal copper sulphate with the addition of ammonium chloride, a total of three atomic weights of oxygen is consumed; these results accord with the following equations:

Na2S2O4 + H2O + O = 2NaHSO3,
Na2S2O4 + H2O + 3O = 2NaHSO4,
S2O3 + O = 2SO2,
2SO2 + O2 = 2SO3,

according to which hydrosulphurous acid corresponds with an oxide S2O3, whereas the oxide required for a salt NaHSO2 would be SO. In this way a decision between the formulae NaHSO2 and Na2S2O4 is easily made, whereas the percentage of sodium and sulphur in the two salts is almost the same and is of relatively little value in enabling a selection between the two to be made.

Confirmatory evidence of the probable correctness of the formula Na2S2O4 is supplied by electrical conductivity measurements on aqueous solutions and comparison of the results with those given by the normal salts of the other sulphur oxy-acids, namely Na2SO4, Na2SO3, Na2S2O3 and Na2SxO6. All these salts dissociate in two stages. Cryoscopic measurements with the sodium salt also indicate the formation of three ions per molecule.

The foregoing considerations were of especial importance in the earlier history of the hydrosulphites, before the salts had been obtained in a pure condition and before their syntheses by the direct interaction of sulphur dioxide and the metal or metallic hydride were available as evidence.

The direct sulphur to sulphur linking in the first formula is in harmony with the formation of the salt by reduction of the sulphite and with the stability of the salt in the presence of alkalis, but it is discounted by the absence of dithionate from the oxidation products, and by the easy fission of the substance into sulphite and sulphoxylate on treatment with an aldehyde.

In support of this contention, the hydrolysis of the sodium salt in alkaline solution in the presence of sodium plumbite may be cited; at room temperature lead sulphide and lead are slowly precipitated, but if the solution is boiled before the addition of the plumbite only the sulphide is precipitated. The reaction in the cold is similar to that of formaldehydesulphoxylate under similar conditions, so that sulphoxylate is evidently a product of the hydrolysis of the hydrosulphite. When the solution is boiled the sulphoxylate is decomposed, being converted into sulphide and sulphite as fast as it is formed. The presence of the S-O-S chain is also in agreement with the formation of benzyl sulphoxide, benzylsulphonic acid and benzylsulphinic acid on treatment with benzyl chloride and potassium hydroxide, and the only important fact possibly at variance with this constitutional representation is the stability of the substance towards alkalis, since it might be expected that a substance of such structure would resemble an acid anhydride in properties.

The third formula perhaps expresses the general behaviour of hydrosulphites more closely than the second, particularly with respect to the action with various oxidising agents. With mild oxidising agents sulphite is the main product, but gaseous oxygen is exceptional in producing approximately equivalent quantities of sulphate and sulphite.

Aqueous solutions of hydrosulphites become orange-yellow on acidification and the presence of sulphur dioxide is soon evident. The yellow colour is not due to colloidal sulphur, nor, since its salts are colourless, would it be expected that the free acid should be yellow, unless there is some change in constitution. It has therefore been suggested that the decomposition on acidification involves the formation of a coloured isomeride of the type (HO)2S.SO2 by co-ordination of a molecule of sulphur dioxide with the sulphur atom of sulphoxylic acid. The decomposition may proceed thus:

(i) + H2OS(OH)2 + H2O + SO2,

(ii) S(OH)2 + SO2 ⇔ (HO)2S.SO2; (Coloured isomeride)


(iii) ⇔ (HO)2S.SO2.

Further action results in the formation of pyrosulphite and thiosulphate, which may be explained thus:

(iv) + (HO)2S.SO2H2S2O5 + H2S2O3.

In addition, the further breaking down of these decomposition products results in the formation of polythionates.


As the chief value of hydrosulphites lies in their reducing power, the method of estimation frequently consists of a measurement of the amount of a standard indigo solution which can be reduced by the sample. Another method is to titrate with standard ammoniacal copper sulphate solution until the hydrosulphite is almost completely oxidised and then to complete the process using indigo-carmine as indicator. Ammoniacal silver nitrate solution has also been suggested as oxidising agent, in which case it is convenient to weigh the silver produced.

Instead of making a direct volumetric determination of the hydrosulphite it is possible to modify the process by estimating volumetrically the product of a primary reaction. For example the hydrosulphite solution may be submitted to atmospheric oxidation and the resulting acidity determined with standard alkali,2 or a mercuric salt may be reduced, the mercury produced being estimated subsequently by the addition of standard iodine solution and titration of the excess of iodine; one molecule of hydrosulphite is equivalent to an atom of mercury and therefore to two atoms of iodine. Similarly, instead of the gravimetric estimation of silver as described above, the latter may be redissolved in nitric acid and determined volumetrically.

Hydrosulphites may also be determined accurately by titration with standard ferricyanide solution using ferrous ammonium sulphate as indicator; the reaction is expressed by the equation:

2K3Fe(CN)6 + Na2S2O4 + 2H2O = 2K3NaFe(CN)6 + 2H2SO3.

Alternatively, the titration may be conducted in the presence of alkali, the end-point being determined electrometrically.

An iodometric method consists in estimating the iodine liberated from a mixture of potassium iodide and iodate, excess of which is added to the hydrosulphite solution, the reaction being according to the equation:

3Na2S2O4 + 4KIO3 + 2KI = 3I2 + 3Na2SO4 + 3K2SO4.

Excess of standard thiosulphate is added after the action and back- titrated with standard iodine.

Another method consists in titrating the hydrosulphite with iodine;

in the presence of excess of formaldehyde the formaldehydesulphoxylate alone reacts, thus:

Na2S2O4 + CH2O + 2I2 + 4H2O = NaHSO4 + 4HI + NaHSO3.CH2O + H2O.

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