Chemical elements
    Physical Properties
    Chemical Properties
      Alkali Phosphides
      Alkaline Earth Phosphides
      Copper Silver and Gold Phosphides
      Zinc Group Phosphides
      Aluminium Phosphide
      Titanium Group Phosphides
      Tin Phosphides
      Lead Phosphides
      Arsenic Phosphides
      Antimony Phosphides
      Bismuth Phosphides
      Chromium Phosphides
      Molybdenum and Tungsten Phosphides
      Manganese Phosphides
      Iron Phosphides
      Cobalt Phosphides
      Phosphonium Chloride
      Phosphonium Bromide
      Phosphonium Iodide
      Hydrogen Phosphides
      Phosphorus Trifluoride
      Phosphorus Pentafluoride
      Phosphorus Trifluorodichloride
      Phosphorus Trifluorodibromide
      Fluophosphoric Acid
      Phosphorus Dichloride
      Phosphorus Trichloride
      Phosphorus Pentachloride
      Phosphorus Chlorobromides
      Phosphorus Chloroiodides
      Phosphorus Tribromide
      Phosphorus Pentabromide
      Phosphorus Diiodide
      Phosphorus Triiodide
      Phosphorus Oxytrifluoride
      Phosphorus Oxychloride
      Pyrophosphoryl Chloride
      Metaphosphoryl Chloride
      Phosphoryl Monochloride
      Phosphoryl Dichlorobromide
      Phosphoryl Chlorodibromide
      Phosphoryl Tribromide
      Metaphosphoryl Bromide
      Phosphoryl Oxyiodides
      Phosphorus Thiotrifluoride
      Phosphorus Thiotrichloride
      Phosphorus Thiotribromide
      Mixed Phosphorus Thiotrihalides
      Phosphorus Suboxides
      Phosphorus Trioxide
      Phosphorus Dioxide
      Phosphorus Pentoxide
      Hypophosphorous Acid
      Phosphorous Acid
      Meta- and Pyro-phosphorous Acids
      Hypophosphoric Acid
      Tetraphosphorus Trisulphide
      Diphosphorus Trisulphide
      Tetraphosphorus Heptasulphide
      Phosphorus Pentasulphide
      Phosphorus Oxysulphides
      Phosphorus Thiophosphites
      Phosphorus Thiophosphates
      Phosphorus Selenophosphates
      Phosphorus Sulphoselenides
      Diamidophosphorous Acid
      Phosphorus Triamide
      Monamidophosphoric Acid
      Diamidophosphoric Acid
      Triamidophosphoric Acid
      Dimetaphosphimic Acid ≡P=
      Trimetaphosphimic Acid
      Tetrametaphosphimic Acid
      Penta- and Hexametaphosphimic Acid
      Monamidodiphosphoric Acid
      Diamidodiphosphoric Acid
      Triamidodiphosphoric Acid
      Nitrilotrimetaphosphoric acid
      Monothioamidophosphoric Acids
      Thiophosphoryl Nitride
      Di- Tri-imido- and -amido-thiophosphates
      Imidotrithiophosphoric Acid =
      Phosphorus Chloronitrides
      Triphosphonitrilic Chloride
      Tetraphosphonitrilic Chloride
      Pentaphosphonitrilic Chloride
      Hexaphosphonitrilic Chloride
      Heptaphosphonitrilic Chloride
      Triphosphonitrilic Bromide
      Phosphorus Halonitrides
      Phosphorus Nitride
      Pyrophosphoric Acid
      Phosphoric acids
    Slow Oxidation
    Phosphatic Fertilisers

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

Aqueous Solutions of Hypophosphorous Acid

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


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/litre0.50040.25020.12510.06250.03128

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 Qd Cals. per gram-ion and the undissociated portion of the free acid is 1 - α, then the total heat of neutralisation Qn will be given by the equation

Qn = 13.52 + (1 - α)Qd Cals.

At 21.5° C. a is 0.449 at a certain concentration and (1 - α)Qd has the value +1.769 Cals. Therefore Qn 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

H3PO2H+ + 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.

Oxidation in Solution

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.

The Hypophosphites

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.

Detection and Estimation of Phosphites and Hypophosphites

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.


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.
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