Hypophosphoric Acid, H₃PO₃

Hypophosphoric Acid, H₃PO₃, together with phosphorous and phosphoric acids, was first produced by the slow oxidation of phosphorus in the air and in the presence of water at a moderately low temperature (5° to 10° C.); it was isolated through the formation of a sparingly soluble sodium salt.

It was also produced by the slow oxidation of phosphorus with silver nitrate in neutral or ammoniacal solution.

Oxidation of yellow phosphorus by copper nitrate in the presence of nitric acid at about 60° C. also yielded hypophosphoric acid along with phosphoric acid and copper phosphide. These reactions have been expressed by the equations

4P + 5Cu(NO₃)₂ + 8H₂O = Cu₃P₂ + 2Cu + 10HNO₃ + 2H₃PO₄
4P + 4Cu(NO₃)₂ + 6H₂O = Cu₃P₂ + Cu + 8HNO₃ + H₄P₂O₆

Better yields were obtained by the electrolysis of 1 to 2 per cent, sulphuric acid using anodes of copper or nickel phosphide.

The conditions for obtaining a good yield by the oxidation of phosphorus have been worked out by a number of investigators. In the method of Cavalier and Cornec glass rods are placed parallel to one another in a photographic dish and sticks of phosphorus laid across them, each pair of sticks being separated by a glass rod, which is also of course laid transversely to the supporting rods. Water is added until the sticks of phosphorus are half-immersed and the whole is covered with a glass plate resting on wadding at its edges. The air is thus filtered as it enters round the edges. A temperature of 5° to 10° C. is the most favourable. Below 5° C. the oxidation is too slow, while above 10° C. undue quantities of meta- and ortho-phosphoric acids are produced. In winter about 12 per cent, of the phosphorus is oxidised to hypophosphoric acid. The solution may then be treated with saturated sodium acetate, when the acid salt NaHPO₃.3H₂O separates on cooling. Or the boiling solution of the acids may be neutralised to methyl orange with sodium carbonate and concentrated by evaporation, when NaHPO₃.3H₂O crystallises and may be re- crystallised from boiling water. The sodium salt dissolved in hot water may be treated with lead acetate and the sparingly soluble lead hypophosphate filtered off. This is then suspended in water and a current of hydrogen sulphide is passed, which liberates the acid. The filtered solution is concentrated somewhat by heat in the open and finally in a vacuum over concentrated sulphuric acid. Alternatively the dilute solution of the free acid may be prepared from barium hypophosphate and dilute sulphuric acid.

The evaporated solution prepared as above deposited crystals of the hydrate H₂PO₃.H₂O which had the form of four-sided rectangular plates and melted at 70° C. They were very deliquescent, but when kept over sulphuric acid effloresced giving the anhydrous acid, which melted at 55° C. A half-hydrate, H₂PO₃.½H₂O, which melted at about 80° C. was also obtained by evaporation over sulphuric acid.

When the anhydrous acid was kept above its melting-point it suddenly decomposed with considerable evolution of heat, thus

2H₂PO₃ = H₃PO₃ + HPO₃

Phosphine was evolved at about 180° C.

The molar heat of fusion of the anhydrous acid was 3.85 Cals. The hydrate dissolved with absorption of heat.

The electrical conductivity of this acid shows that its first hydrion is largely dissociated.

The molar conductivities of hypophosphoric and phosphorous acids

V	16	32	64	128	256	512	1024
μ H2PO3 at 25.6° C.	184	199	222	246	275	304	370
μ H3PO3 at 25.0° C.	222	252	292	318	337	351	358

The molar conductivities of H₃PO₃ at dilutions up to, and including, 512 were greater than those of H₂PO₃, and this difference would have increased if the latter conductivities had been determined at 25° C. The ions HPO₃^- and H₂PO₃^- probably have practically the same mobility, so that the difference is to be attributed rather to the smaller second dissociation of HPO₃^- (into H⁺ and PO₃⁼) as compared with that of H₂PO₃^- into H⁺ and HPO₃⁼. A very low second constant would also be ascribed to H₂PO₃ on account of the high alkalinity of solutions containing two equivalents of alkali.

The change in the conductivity of the salt NaHPO₃ or Na₂H₂P₂O₆ with dilution shows no effect due to dissociation of a second hydrogen until a dilution is reached between 256 and 512 litres, as is seen from the following table:—

V	16	32	64	128	256	512	1024
μ (25° C)	78.8	81.6	88.1	94.5	100.2	105.9	111.9

The heats of neutralisation indicate a dibasic acid (per atom of phosphorus) as will be seen from the following figures:—

Equivalents of NaOH per mol H2PO3	0.5	1	1.5	2.0	3.0
Heat evolved, Cals.	7.57	15.05	21.36	27.11	27.65

During the neutralisation with alkali a sharp end-point was obtained (to methyl orange) when one equivalent of alkali had been added, while phenolphthalein gradually changed between 1.5 and 2.0 equivalents, showing that the second dissociation constant is lower than that of the majority of organic acids.

These facts can be explained equally well on the assumption that the acid is tetrabasic, with two atoms of phosphorus, i.e. H₄P₂O₆, and this view is in accordance with the existence of four salts—MH₃P₂O₆, M₂H₂P₂O₈, M₃HP₂O₆ and M₄P₂O₆.

The molecular weight as deduced from the freezing-points of aqueous solutions corresponded with doubled molecules, H₄P₂O₆, which are highly dissociated, giving one hydrogen ion, or to the acid H₂PO₃ which is very slightly dissociated even in fiftieth-normal solution.

The proof of the individuality of hypo- phosphoric acid rests largely upon the great differences which exist between it and a mixture of phosphorous and phosphoric acids. These acids were formed irreversibly when hypophosphoric acid was allowed to stand in aqueous solution, especially when this was concentrated and the temperature was 30° C. (or above). A 5 per cent, solution after 3 years contained only phosphorous and phosphoric acids. These acids, on the other hand, showed no tendency to condense together when kept under ordinary conditions. The freezing-point curves of mixtures of the anhydrous acids showed no intermediate maximum and only one eutectic at -13.0° C. and 39 mols per cent, of H₃PO₃.

Decomposition of hypophosphoric acid may be represented as an hydrolysis, thus

H₄P₂O₆+H₂O → H₃PO₃+H₃PO₄

This reaction is catalysed by hydrogen ions and was found to be uni-molecular in normal hydrogen-ion concentrations. The values of the constants were 0.000186 at 25° C. and 0.00631 at 60° C. The sodium salt, NaHPO₃, having only a low hydrogen-ion concentration (ca. 1×10^-4), could be kept for long periods without much change, and alkaline solutions containing Na₂PO₃, etc., were still more stable.

Cold solutions of the acid did not precipitate the metals from gold, silver or mercurous solutions. The acid was not oxidised by iodine, hydrogen peroxide or chromic acid in the cold, but was slowly oxidised by potassium permanganate. Hot concentrated solutions were more easily oxidised. When the neutral salts were heated they gave phosphate and phosphine, or pyrophosphate, elemental phosphorus and phosphine. Reducing agents such as hydrogen sulphide or sulphur dioxide had no effect, and even nascent hydrogen, from zinc and acid, gave no phosphine.

Much investigation has been carried out with a view to ascertaining whether the molecule should be formulated as H₂PO₃ or H₄P₂O₆. The evidence of electrical conductivity does not lead to a definite conclusion. Molar weights of some esters are known. Thus dimethyl and diethyl hypophosphates (from the silver salt and alkyl iodide) gave elevations of boiling-point in ethyl iodide, chloroform, etc., which corresponded to the formulae (CH₃)₂PO₃, etc. On the other hand molecules such as (CH₃)₄P₂O₆ were indicated in benzene solution at its freezing-point. If, as seems probable, the free acid in concentrated solution has the formula H₄P₂O₆, the constitution and hydrolysis of this would be represented by

(HO)₂OP-PO(OH)₂ + H₂O = (HO)₂HPO + (HO)₃PO

Hypophosphoric acid was not formed by any ordinary dehydration of a mixture of phosphorous and phosphoric acids, neither was it produced by the hydration of P₂O₄, which yielded only a mixture of equal mols of H₃PO₃ and H₃PO₄.

The salts of the alkali, alkaline earth and some other metals were prepared and studied by Salzer, Rammelsberg, Schuh, Joly, Palm, Bausa and others. The alkali salts prepared from excess of alkali were soluble, whilst those of the alkaline earths, silver and other metals were only sparingly soluble.

The formulae of some typical hypophosphates are as follows [solubilities in grams per 100 grams of water]:—

Li₂H₂P₂O₆.2H₂O, crystalline, sparingly soluble; Li₄P₂O₆.7H₂O, solubility 0.83; Na₂H₂P₂O₆.6H₂O, monoclinic, solubility 2.0; Na₃HP₂O₆.9H₂O, tabular monoclinic; Na₄P₂O₆.10H₂O, six-sided monoclinic, solubility 1.5; K₃H₅(P₂O₆)₂.2H₂O, rhombic, solubility 40; K₂H₂P₂O₆.2H₂O, monoclinic plates; K₄P₂O₆.8H₂O, rhombic pyramids, solubility 25; (NH₄)₂H₂P₂O₆, granular or needles, stable in air, solubility 7.1; (NH₄)₄P₂O₆.H₂O, prismatic, efflorescent, solubility 3.3; Ag₄P₂O₆, insoluble; CaH₂P₂O₆.6H₂O, monoclinic prisms; Ca₂P₂O₆.2H₂O, gelatinous, insoluble; BaH₂P₂O₆, monoclinic prisms; Ba₂P₂O₆, white

precipitate; Mg₂P₂O₆.12H₂O, gelatinous, slowly crystalline, solubility 0.0067; Zn₂P₂O₆.2H₂O, white precipitate; Cu₂P₂O₆.6H₂O, insoluble; Pb₂P₂O₆, insoluble; Ni₂P₂O₆.12H₂O, prismatic crystals.

A number of double hypophosphates of manganese, cobalt and nickel with the alkali metals have been prepared, such as K₂NiP₂O₆.6H₂O and 3K₂H₂P₂O₆.CoH₂P₂O₆.15H₂O.

Hypophosphites of hydroxylamine resemble those of ammonia. Thus (NH₂OH)₂.H₄P₂O₆ is a crystalline salt, very soluble in water and melting with decomposition. Hydrazine dihydro- and trihydro-hypo-phosphates, (N₂H₅)₂H₂P₂O₆ and (N₂H₅)H₃P₂O₆, were both obtained in the crystalline state. The latter is isomeric with ammonium dimetaphosphate, (NH₄PO₃)₂.

One of the most useful is the sparing solubility of sodium hypophosphate. Also, a sparingly soluble guanidine salt (1 per cent, at 28.5° C.) is precipitated when guanidine carbonate is added to a solution of the acid. The acid may be titrated with alkali and phenolphthalein; 3 mols of NaOH correspond to 2 mols of H₂PO₃, i.e. the salt Na₃HP₂O₆ is formed. The acid may be estimated in the presence of phosphoric and phosphorous acids by precipitation of the silver salt at pH = 1 to 2.