Sodium hyposulphite, Na₂S₂O₄

The Sodium hyposulphite was first prepared in solution by Schutzenberger by reducing a solution of sodium hydrogen sulphite with zinc, but the yield is unsatisfactory. Bernthsen and Bazlen obtained the pure substance by reducing the primary sulphite with zinc-dust in presence of sulphurous acid,

2NaHSO₃ + SO₂ + Zn = Na₂S₂O₄ + ZnSO₃ + H₂O,

precipitating the zinc and removing the excess of acid by addition of milk of lime, the hyposulphite being salted out in the form of dihydrate by addition of sodium chloride. They recommend increasing the stability of the dihydrate by washing with acetone, at first dilute and finally pure, and drying in a vacuum. Jellinek advises heating the dihydrate at 60° C. in vacuum, the process transforming it into the stable anhydrous salt. The hyposulphite can also be produced under special experimental conditions by the electrolytic reduction of sodium hydrogen sulphite, Jellinek having obtained a 7 to 8 per cent, solution by this means. Prepared by this process, it finds application in the reduction of indigo to indigo-white.

The anhydrous hyposulphite is also formed by the interaction of sulphur dioxide and sodium hydride:

2NaH+2SO₂=Na₂S₂O₄+H₂.

The anhydrous salt is converted into the dihydrate by crystallization from water.

A laboratory method for the preparation of sodium hyposulphite depends on the interaction of sodium formaldehydesulphoxylate, NaCH₃O₃S,2H₂O, and sodium hydrogen sulphite, the product having a purity of 80 to 85 per cent., and the yield being 55 to 60 per cent, of that theoretically possible. The sodium formaldehydesulphoxylate can be obtained by the reduction of commercial "hydrosulphite " by zinc-dust and zinc oxide in presence of formaldehyde solution, the crude product being recrystallized from water at 70° C.

Sodium hyposulphite dihydrate forms colourless vitreous prisms, readily soluble in water. On heating it loses its water of crystallization, melts at red heat, and burns with a blue flame, evolving sulphur dioxide. On addition of acid in small proportion the solution develops a red colour, due to separation of sulphur. Molecular-weight determinations by the cryoscopic method indicate it to be the salt of a dibasic acid, and to be represented by the formula Na₂S₂O₄.

At temperatures between 0° and 30° C. the salt decomposes in aqueous solution in accordance with the scheme

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

Hydrogen sulphide has no action on dry sodium hyposulphite, but in presence of water it reacts as indicated by the equation

Na₂S₂O₄+H₂S=Na₂S₂O₃+H₂O+S.

The sulphur is obtained in a weighable form, and Sinnatt advocates the estimation of sodium hyposulphite by this method.

The method devised by Seyewetz and Bloch for the valuation of this substance depends on the oxidizing action of an ammoniacal solution of silver chloride:

Na₂S₂O₄+2AgCl+4NH₄OH =2(NH₄)₂SO₃+2NaCl+2H₂O + 2Ag.

Hollingsworth Smith has improved the speed of the operation, and has also eliminated the defect due to solid impurities in the hyposulphite being weighed with the silver.

Sodium hyposulphite is oxidized by iodine to sodium hydrogen sulphate, a reaction applied by Bazlen and Bernthsen to the quantitative estimation of the salt:

Na₂S₂O₄+6I+4H₂O =2NaHSO₄+6HI.

The action of selenium and tellurium on sodium hyposulphite has been investigated by Tschugaev and Chlopin.

The hyposulphite was employed by Julius Meyer in the preparation of metallic colloidal solutions. It finds technical application as a reducer in the dye-industry. When administered intravenously, it exerts a toxic effect.