Hypophosphorous Acid, H₃PO₂

The alkali salts of this acid were discovered among the products of the decomposition of phosphides by water. A method of preparing hypophosphites by boiling milk of lime with phosphorus was also discovered early in the nineteenth century. The resulting solution of calcium hypophosphite could then be decomposed by oxalic acid. Hypophosphite was also prepared by heating the phosphorus with a solution of baryta. The barium salt, Ba(H₂PO₂)₂, is easily recrystallised, and from it the free acid may be prepared by double decomposition of a fifth molar solution with the calculated amount of 20 to 25 per cent, sulphuric acid. The filtered solution may be evaporated first to one-tenth of its volume and then until the temperature rises to 105° C. It is filtered hot and then further evaporated to a temperature of 110° C., and this evaporation by stages is continued until the temperature rises to 130° or even 138° C. without decomposition. The liquid should then be poured into a closed flask and cooled to 0° C., when it nearly all solidifies to a mass of crystals. Crystallisation may be induced if necessary by seeding with a crystal of the acid. The commercial acid usually contains calcium salts. These may be removed by the addition of alcohol and much ether to the evaporated solution, when the salts are precipitated. The alcohol and ether are removed by distillation and the acid is further concentrated as above.

The density of the crystallised acid is given as 1.4625. The melting-point, 17.4° C. or 26.5° C., no doubt varies with small variations in the proportion of water present.

The latent heat of fusion (heat absorbed) is 2.4 Cals., the latent heat of solution of the crystals from -0.2 to -0.17 Cals., and of the fused acid from +2.2 to +2.14 Cals. (per mol in each case).

The heats of formation (heat evolved) from the elements

P (solid) + O₂ (gas) + 1½H₂ (gas)

are +137.7 Cals. (liquid acid), +140.0 Cals. (solid acid), +139.8 Cals. (dissolved acid).

The pure acid decomposes rapidly when heated above 130° C. and below 140° C. mainly according to the equation

3H₃PO₂ = PH₃ + 2H₃PO₃

while between 160° and 170° C. the main decomposition (consecutive reaction) is symbolised as

4H₃PO₃ = PH₃ + 3H₃PO₄

The effect of dilution on the molar conductivity of hypophosphorous acid shows that the acid is moderately strong, but however obeys the law of mass or concentration action sufficiently well to give a "constant" (K below) which remains of the same order although it diminishes steadily with increasing dilution.

Molar conductivities of hypophosphorous acid at 25° C

V	2	4	8	16	32	64	128	256	512	1024	2048	∞
μ	140	172	207	245	281	312	335	352	361	367	368.3	389
K	0.1012	0.0876	0.0757	0.0670	0.0587	0.0508	0.0417	0.0336	0.0234	0.0154	0.0082	...
K'	...	...	0.1015	0.1014	0.1017	0.1014	0.1019	0.1024	0.1008	0.1014	0.1028	...

The values of α (=μ/μ_∞) in the expression α²/(1-α)V = K were calculated from μ_∞=389 based on the work of Arrhenius. The ionic conductivity L_∞ of H⁺ = 347, which gives L_∞ of H₂PO₂^- as 42.0, in agreement with the value 41.8 deduced from the sodium salt.

These valued have been recalculated to the newer units and an empirical constant K' (see table) has been found which shows a better constancy, i.e.:



The constant K has been redetermined by Kolthoff as 0.01 at V = 1000 and 0.062 at V = 20. In another series of results the values of α and K were derived from λ_∞ = 392.5, which was based on the limiting conductivity of NaH₂PO₂ (at 25° C.):—

c mols/litre	0.5004	0.2502	0.1251	0.0625	0.03128
μ	137.1	168.5	205.0	242.9	279.5

whence K may be calculated in the usual manner.

The conductivity increases at first with the temperature as is usual; the rate of increase then diminishes and the conductivity reaches a maximum at about 50° C., the exact temperature varying with the concentration and being 57° C. in the case of normal acid. The conductivity then decreases. It is supposed that the effect of the normal increase in ionic mobility with temperature is diminished and finally reversed by the opposite effect of decreasing dissociation. Since the dissociation constant decreases with rise of temperature the dissociation into ions must take place with evolution of heat, i.e. the heat of ionisation is positive. Therefore the neutralisation of the acid with alkali must result in a production of heat greater than the heat of formation of water from its ions, which may be taken as 13.52 Cals. per mol. If the heat of dissociation is Q_d Cals. per gram-ion and the undissociated portion of the free acid is 1 - α, then the total heat of neutralisation Q_n will be given by the equation

Q_n = 13.52 + (1 - α)Q_d Cals.

At 21.5° C. a is 0.449 at a certain concentration and (1 - α)Q_d has the value +1.769 Cals. Therefore Q_n is 15.289 Cals. by calculation, while the experimental value was 15.316 Cals.

The transport number of the anion was found to be 41.8.

The results of conductivity measurements indicate that only one of the hydrogens is dissociable as ion, the dissociation taking place according to the equation

H₃PO₂ ⇔ H⁺ + H₂PO₂^-

In the process of neutralisation also only one hydrion takes part, as is shown by the fact that the heat evolved practically reaches a limit when one equivalent of alkali has been added to one mol of the acid. Thus when 2, 1 and 0.5 mols of the acid were added to one equivalent of NaOH the heats observed were 15.4, 15.2 and 7.6 Cals.

The monobasicity is confirmed by the formulae of all the known salts.

Hypophosphorous acid and its salts are strong reducing agents and are oxidised to phosphorous acid or finally to phosphoric acid and their salts. A similar reaction is brought about by hydrogen iodide:—

3H₃PO₂ + HI = 2H₃PO₃ + PH₄I

Hypophosphorous acid is also oxidised to phosphorous acid by sulphur dioxide with deposition of sulphur, and by phosphorus trichloride with deposition of phosphorus, thus:—

3H₃PO₂ + PCl₃ = 2H₃PO₃ + 2P + 3HCl

It is oxidised by alkali with evolution of hydrogen:—

NaH₂PO₂ + NaOH = Na₂HPO₃ + H₂

The velocity constant of this reaction has been measured at temperatures slightly below 100° C. It is slightly greater for KOH than for NaOH at equivalent concentration, and increases faster than the concentration of the alkali.

Salts of the noble metals are reduced by hypophosphorous acid or hypophosphites, and in many cases phosphorous acid can be isolated among the products. The reduction of silver nitrate was noticed by the early workers. Phosphoric acid was formed with or without the evolution of hydrogen, thus:—

2NaH₂PO₂ + 2AgNO₃ + 4H₂O = 2H₃PO₄ + 2NaNO₃ + 3H₂ + 2Ag

or the nascent hydrogen gave more silver:—

2H + 2AgNO₃ = 2HNO₃ + 2Ag

The formation of phosphorous acid is represented by the equation

2AgNO₃ + H₃PO₂ + H₂O = 2Ag + H₃PO₃ + 2HNO₃

Copper sulphate is reduced to copper with the production of an acid solution:—

4CuSO₄ + Ba(H₂PO₂)₂ + 4H₂O = 2H₃PO₄ + 4Cu + BaSO₄ + 3H₂SO₄

It has also been stated that copper hydride is first produced, which decomposes with evolution of hydrogen. It appears further that, with excess of the hypophosphite, hydrogen also is evolved, while with excess of copper salt copper only is precipitated. The spongy copper appears to act as a catalytic agent in liberating hydrogen from excess of hypophosphorous acid. Cuprous oxide probably is first formed, and the substance first precipitated from an acid solution may contain this with hydride and phosphate.

Mercuric chloride is reduced to mercurous chloride and mercury. In this case also phosphorous acid was produced, and by a reaction which was quicker than that which led to phosphoric acid, thus:—

H₃PO₂ + 2HgCl₂ + H₂O = H₃PO₃ + Hg₂Cl₂ + 2HCl

Palladous salts oxidise the acid to phosphoric acid with deposition of palladium.

The velocities of some oxidations have been determined. Thus in the reaction

H₃PO₂ + I₂ + H₂O = H₃PO₃ + 2HI

the velocity was independent of the concentration of iodine if this was more than 0.004N, and unimolecular with respect to H₃PO₂. The hypothesis has been advanced that the reducing agent is an active form H₅PO₂, which is produced with a measurable velocity (catalysed by hydrogen ions) when the equilibrium amount is diminished.

Hypophosphites of most of the metals have been prepared by a few general reactions:—

1. By heating aqueous solutions of the alkali or alkaline earth hydroxides with white phosphorus.
2. By the double decomposition of barium hypophosphite with the sulphate of the required metal.
3. By dissolving the hydroxide or carbonate of the metal in hypophosphorous acid.

The hypophosphites are all soluble. Those of barium and calcium, which are the least soluble, dissolve in 2.5 to 3.5 and 6 to 7 parts of water respectively. Those of the alkali metals and ammonium dissolve also in alcohol. Solutions of hypophosphites of the alkali metals are fairly stable, especially in the absence of air, and the salts generally may be obtained in well-crystallised forms by evaporation. More concentrated solutions often decompose, especially if alkaline, with evolution of phosphine. The dry salts are also fairly stable in the cold, but when heated decompose giving phosphine and hydrogen and leaving the pyro- or meta-phosphate.

The electrical conductivities of the sodium salt and of the barium salt give the mobility of the H₂PO₂^- ion.

The formulae of some typical hypophosphites are as follows:—

LiH₂PO₂.H₂O, monoclinic prisms; NaH₂PO₂.H₂O, monoclinic prisms; KH₂PO₂, hexagonal plates; NH₄H₂PO₂, rhombic tables; Mg(H₂PO₂)₂.6H₂O, tetragonal; Ca(H₂PO₂)₂, monoclinic leaflets; Ba(H₂PO₂)₂.H₂O, monoclinic needles or prisms; Cu(H₂PO₂)₂, white precipitate; Pb(H₂PO₂)₂, rhombic prisms; Fe(H₂PO₂)₂.6H₂O, green octahedra; Fe(H₂PO₂)₃.xH₂O, white precipitate; Co(H₂PO₂)₂.6H₂O, red tetragonal; Ni(H₂PO₂)₂.6H₂O, green crystals.

The hexahydrated salts of magnesium and the iron group are said to be isomorphous.

Ammonium and hydroxylamine hypophosphites have been prepared by double decomposition between the sulphates and barium hypophosphite. The ammonium salt, NH₄H₂PO₂ was crystallised from water or alcohol. When heated to about 200° C. it melted and decomposed, giving off ammonia, phosphine and hydrogen and leaving a mixture of pyro- and meta-phosphoric acids. Hydroxylamine hypophosphite, NH₂OH.H₃PO₂, has also been prepared by reaction between KH₂PO₂ and NH₂OH.HCl, and was extracted with hot absolute alcohol. Solutions of this salt must be evaporated in an atmosphere of CO₂, etc. in order to avoid oxidation. The crystals obtained were very deliquescent. They began to decompose at about 60° C., melted below 100° C. and exploded at a higher temperature.

In contradistinction to phosphorus itself and the products of its slow oxidation, hypophosphorous acid and the hypophosphites do not appear to be toxic. Calcium hypophosphite appears to be completely eliminable from the system. Hypophosphites of calcium, sodium and iron have been prescribed in medicine, but although in some cases they appear to have done good there is no conclusive evidence of the value of the hypophosphite radical apart from the basic radical or other constituent of the mixture.

Solutions of hypophosphorous and phosphorous acids, as well as their salts, when evaporated to dryness and heated, give spontaneously inflammable phosphine and a residue of phosphoric acid or phosphate respectively. The reduction of salts of copper and silver, and of mercuric chloride, by these acids may also be used as tests. When the alkali salts are boiled with concentrated alkali hydrogen is evolved and a phosphate is found in solution, thus

KH₂PO₂ + 2KOH = K₃PO₄ + 2H₂
K₂HPO₃ + KOH = K₃PO₄ + H₂

Nascent hydrogen, from zinc and sulphuric acid, reduces both acids to phosphine, which may be detected by the yellow stain which it produces on silver nitrate:—

PH₃ + 6AgNO₃ = PAg₃.3AgNO₃ + 3HNO₃

This is similar to the stain given by arsine and, like AsAg₃.3AgNO₃, is hydrolysed and blackened by moisture:—

PAg₃.3AgNO₃ + 3H₂O = H₃PO₃ + 3HNO₃ + 6Ag

Hypophosphites are all soluble in water and therefore the solutions give no precipitates with the ions of the alkaline earths, silver, etc. Silver salts, however, are rapidly reduced to metallic silver, and phosphoric acid is found in the solution:—

2NaH₂PO₂ + 2AgNO₃ + 4H₂O = 2H₃PO₄ + 2NaNO₃ + 2Ag + 3H₂

When hypophosphites are heated with copper sulphate solution to 55° C. a reddish-black precipitate of Cu₂H₂ is produced, which decomposes at 100° C. into hydrogen and copper. Permanganates are immediately reduced by hypophosphites. The effect of other oxidising agents is mentioned under " Estimation."

Phosphites may be distinguished from hypophosphites by several tests. The ions of barium and lead give white precipitates with solutions of phosphites. Silver nitrate gives a white precipitate in the cold, from which black metallic silver is quickly deposited on warming:—

Na₂HPO₃ + 2AgNO₃ = Ag₂HPO₃ + 2NaNO₃
Ag₂HPO₃ + H₂O = 2Ag + H₃PO₄

Copper sulphate is reduced to brown metallic copper with evolution of hydrogen, but without intermediate formation of a hydride. Permanganates are reduced more slowly by phosphites than by hypophosphites under similar conditions.

Estimation

The course of the neutralisation curves of hypophosphorous and phosphorous acids can be deduced from the values of the dissociation constants. The curves have also been determined with the aid of the glass electrode. The curve for hypophosphorous acid is that of a strong acid with only one inflection. That of phosphorous acid has two inflections, at pH = 3.4 to 4.5 and at pH = 8.4 to 9.2. The first "end-point " can be located with dimethylamino-azobenzene or methyl orange, and the second with phenolphthalein or thymol blue. The first gives the sum of H₃PO₂, H₃PO₃ and H₃PO₄ if present together, while the difference between the first and second "end-points" gives the sum of H₃PO₃ and H₃PO₄. From these data the equivalents of H₃PO₂ can be obtained.

Hypophosphite and phosphite may also be determined by oxidation with iodine. In alkaline bicarbonate solutions phosphites are oxidised quickly to phosphates, while hypophosphites are hardly affected, i.e. they do not use any measurable amount of decinormal iodine after standing for two hours at ordinary temperatures. In acid solution hypophosphorous acid is slowly oxidised to phosphorous acid, but no further, according to the equation

H₃PO₂ + I₂ + H₂O = H₃PO₃ + 2HI

On these reactions is based a method for determining the acids separately or in the same solution.

A phosphite solution of about M/10 concentration is placed in a stoppered flask with an excess of sodium bicarbonate saturated with carbonic acid and an excess of decinormal iodine solution and allowed to stand for about an hour. It is then acidified with acetic acid and the excess of iodine is back-titrated with decinormal bicarbonate- arsenite.

1 c.c. of decinormal iodine solution = 0.004103 gram of H₃PO₃

A hypophosphite solution is acidified with dilute sulphuric acid, treated with a known amount, in excess, of decinormal iodine solution and allowed to stand for 10 hours at ordinary temperatures. A cream of NaHCO₃ is then added until CO₂ evolution ceases, then fifth-normal NaHCO₃ solution, saturated with CO₂, whereby oxidation goes through the next stage, to phosphate. After addition of acetic acid the excess of iodine is titrated with standard arsenite as before.

1 c.c. decinormal iodine solution = 0.00165 gram of H₃PO₂

Phosphite, which usually is present in small amounts in hypophosphite, will be included and must be titrated separately in bicarbonate solution as before.

Acid dichromate and alkaline permanganate may also be used as oxidising agents, the excess being determined iodometrically and with ferrous sulphate solution respectively.

Either of these acids may also be determined by weighing the mercurous chloride which is produced by reaction with mercuric chloride in slightly acid solution, according to the equations

H₃PO₂ + 2H₂O + 4HgCl₂ = 2Hg₂Cl₂ + 4HCl + H₃PO₄
H₃PO₃ + H₂O + 2HgCl₂ = Hg₂Cl₂ + 2HCl + H₃PO₄

The mercurous chloride is washed with water and dried at 110° C., or with alcohol and ether and dried at 95° C.