When moist phosphorus slowly oxidises in the air, it not only forms phosphorous and phosphoric acids, but also hypophosphoric acid, H4P2O6, which when in a dry state easily splits up at 60° into phosphorous and metaphosphoric acids (H4P2O6 = H3PO3 + HPO3), but differs from a mixture of these acids in that it forms well-characterised salts, of which the sodium salt, H2Na2P2O6, is but slightly soluble in water (the sodium salts of phosphoric and phosphorous acids are easily soluble), and that it does not act as a reducing agent, like mixtures containing phosphorous acid.[11]
Judging by the general law of the formation of acids (Chapter [XV.]), the series of phosphorus compounds should include the following ortho-acids and their corresponding anhydrides, answering to phosphuretted hydrogen, H3P:—
| H3PO4, | phosphoric acid, and | P2O5, | anhydride, |
| H3PO3, | phosphorous acid, and | P2O3, | anhydride, |
| H3PO2, | hypophosphorous acid, and | P2O, | anhydride.[12] |
The last of these (the analogue of N2O) is almost unknown. Phosphoric anhydride (P2O5) with a small quantity of water does not at first give orthophosphoric acid, PH3O4, but a compound P2O5,H2O, or PHO3, whose composition corresponds with that of nitric acid; this is metaphosphoric acid. Even with an excess of water, combining with phosphoric anhydride, this metaphosphoric acid, and not the ortho-, passes at first into solution. Metaphosphoric acid in solution only passes into orthophosphoric acid when the solution is heated or after a lapse of time.
Orthophosphoric acid[13] is obtained by oxidising phosphorus with nitric acid until the phosphorus entirely passes into solution and the lower oxides of nitrogen cease to be evolved. The reaction takes place best with dilute nitric acid, and when aided by heat. The resultant solution is evaporated to a syrup. If a weighed quantity of phosphorus (dried in a current of dry carbonic anhydride) be taken, a crystalline mass of the acid can be obtained by evaporating the solution until it consists only of the quantity[14] of phosphoric acid corresponding with the amount of phosphorus taken (from 31 parts of P, 98 parts of solution). The acid fuses at +39°; specific gravity of the liquid 1·88. Phosphorus pentachloride, PCl5, and oxychloride, POCl3 (see further on), give orthophosphoric acid and hydrochloric acid with water. The two other varieties of phosphoric acid, with which we shall presently become acquainted, give the same ortho-acid when under the influence of acids, with particular ease when boiled and more slowly in the cold. By itself orthophosphoric acid (either in solution or when dry) does not pass into the other varieties; it does not oxidise, and therefore presents the limiting and stable form. When heated to 300°, it loses water and passes into pyrophosphoric acid, 2H3PO4 = H2O + H4P2O7, whilst at a red heat it loses twice as much water and is converted into metaphosphoric acid, H3PO4 = H2O + HPO3. In aqueous solution orthophosphoric acid differs clearly from pyro- or metaphosphoric acids, because the solutions of these latter acids give different reactions: thus orthophosphoric acid does not precipitate albumin, does not give a precipitate with barium chloride, and forms a yellow precipitate of silver orthophosphate, Ag3PO4, with silver nitrate (in the presence of alkalis, but not otherwise); whilst a solution of pyrophosphoric acid, H4P2O7, although it does not precipitate albumin or barium chloride, gives a white precipitate of silver pyrophosphate, Ag4P2O7, with silver nitrate; and a solution of metaphosphoric acid, HPO3, precipitates both albumin and barium chloride, and gives a white precipitate of silver metaphosphate, AgPO3, with silver nitrate. These points of distinction were studied by Graham, and are exceedingly instructive. They show that the solution of a substance does not determine the maximum of chemical combination with water, that solutions may contain various degrees of combination with water, and that there is a clear difference between the water serving for solution and that entering into chemical combination. Graham's experiments also showed that the water whose removal or combination determines the conversion of ortho- into meta- and pyrophosphoric acids differs distinctly from water of crystallisation, for he obtained the salts of ortho-, meta-, and pyrophosphoric acids with water of crystallisation, and they differed in their reactions, like the acids themselves. This water of crystallisation was expelled with greater ease than the water of constitution of the hydrates in question.[14 bis]
Orthophosphoric acid has a pleasant acid taste and a distinctly acid reaction; it is used as a medicine, and is not poisonous (phosphorous acid is poisonous). Alkalis, like sodium, potassium, and ammonium hydroxides, saturate the acid properties of phosphoric acid when taken in the ratio 2NaHO : H3PO4—that is, when salts of the composition HNa2PO4 are formed. When taken in the ratio NaHO : H3PO4, a solution having an acid reaction is obtained, and when 3NaHO : H3PO4—that is, when the salt Na3PO4 is formed—an alkaline reaction is obtained. Hence many chemists (Berzelius) even regarded the salts of composition R2HPO4 as normal, and considered phosphoric acid to be bibasic. But the salt Na2HPO4 also shows a feeble alkaline reaction, so that it is impossible to judge the characteristic peculiarities of acids by the reactions on litmus paper, as we already know from many examples. Orthophosphoric acid is tribasic, because it contains three equivalents of hydrogen replaceable by metals, forming salts, such as NaH2PO4, Na2HPO4, and Na3PO4. It is also tribasic, because with silver nitrate its soluble salts always give Ag3PO4,[15] a salt with three equivalents of silver, and because by double decomposition with barium chloride it forms a salt of the composition Ba3(PO4)2, and silver and barium hardly ever give basic salts. With the metals of the alkalis, phosphoric acid forms soluble salts, but the normal salts of the metals of the alkaline earths, R3(PO4)2 and even R2H2(PO4), are insoluble in water, but dissolve in feeble acids, such as phosphoric and acetic, because they then form soluble acid salts, especially RH4(PO4)2.[16]
Phosphoric anhydride, or any of its hydrates, when ignited with an excess of sodium hydroxide, carbonate, &c., forms normal or trisodium orthophosphate, Na3PO4, but when a solution of sodium carbonate is decomposed by orthophosphoric acid, only the salt Na2HPO4 is formed; and when an excess of sodium chloride is ignited with orthophosphoric acid, hydrochloric acid is evolved, and the acid salt H2NaPO4 alone is formed. These facts clearly indicate the small energy of phosphoric acid with respect to the formation of the tri-metallic salt, which is seen further from the fact that the salt Na3PO4 has an alkaline reaction, decomposes in the presence of water and carbonic acid, forming Na2HPO4, corrodes glass vessels in which it is boiled or evaporated, just like solutions of the alkalis, disengages, like them, ammonia from ammonium chloride, and crystallises from solutions, as Na3PO4,12H2O, only in the presence of an excess of alkali. At 15° the crystals of this salt require five parts of water for solution; they fuse at 77°.
Disodium orthophosphate, or common sodium phosphate, Na2HPO4, is more stable both in solution and in the solid state. As it is used in medicine and in dyeing, it is prepared in considerable quantities, most frequently from the impure phosphoric acid obtained by the action of sulphuric acid on bone ash. The solution thus formed—which contains, besides phosphoric and sulphuric acids, salts of sodium, calcium, and magnesium—is heated, and sodium carbonate added so long as carbonic anhydride is disengaged. A precipitate is formed containing the insoluble salts of magnesium and calcium, whilst the solution contains sodium phosphate, Na2HPO4, with a small quantity of other salts, from which it may be easily purified by crystallisation. At the ordinary temperature its solutions, especially in the presence of a small amount of sodium carbonate, give finely-formed inclined prismatic crystals, Na2HPO4,12H2O; when the crystallisation takes place above 30° they only contain 7H2O. The former crystals even lose a portion of their water of crystallisation at the ordinary temperature (the salt effloresces), and form the second salt with 7H2O; whilst under the receiver of an air-pump and over sulphuric acid they also part with this water.[17] When ignited they lose the last molecule of water of constitution, and give sodium pyrophosphate, Na4P2O7.