[13] Phosphoric acid, being a soluble and almost non-volatile substance, cannot be prepared like hydrochloric and nitric acids by the action of sulphuric acid on the alkali phosphates, although it is partially liberated in the process. For this purpose the salts of barium or lead may be taken, because they give insoluble salts, thus Ba3(PO4)2 + 3H2SO4 = 3BaSO4 + 2H3PO4. Bone ash contains, besides calcium phosphate, sodium and magnesium phosphates, and fluorides and other salts, so that it cannot give directly a pure phosphoric acid.
[14] If this is not done the orthophosphoric acid, PH3O4, loses a portion of its water, and then, as with an excess of water, it does not crystallise.
[14 bis] The difference between the reactions of ortho-, meta- and pyrophosphoric acids, established by Graham (see p. [163]), is of such importance for the theory of hydrates and for explaining the nature of solutions, that in my opinion its influence upon chemical thought has been far from exhausted. At the present time many such instances are known both in organic (for instance, the difference between the reactions of the solutions of certain anhydrides and hydrates of acids), and inorganic chemistry (for example, the difference between the rose and purple cobalt compounds, Chapter [XXII.] &c.) They essentially recall the long known and generalised difference between C2H4 (ethylene), C2H6O (ethyl alcohol = ethylene + water), and C4H10O (ethyl ether = 2 ethylene + water = 2 alcohol - water); but to the present day the numerous analogous phenomena existing among inorganic substances are only considered as a simple difference in degrees of affinity, distinguishing the water of constitution (hydration), crystallisation, and solution without penetrating into the difference of the structure or distribution of the elements, which exists here and gives rise to a distinct isomerism of solutions. In my opinion the progress of chemistry, especially with regard to solutions, should make rapid strides when the cause of the isomerism of solutions, for instance, of ortho- and pyrophosphoric acids, has become as clear to us as the cause of many well-studied instances of the isomerism, polymerism, and metamerism of organic compounds. Here it forms one of those many important problems which remain for the chemistry of the future in a state of only indistinct presentiments and in the form of facts empirically known but insufficiently comprehended.
[15] Silver orthophosphate, Ag3PO4, is yellow, sp. gr. 7·32, and insoluble in water. When heated it fuses like silver chloride, and if kept fused for some length of time it gives a white pyrophosphate (the decomposition which causes this is not known). It is soluble in aqueous solutions of phosphoric, nitric, and even acetic acids, of ammonia, and many of its salts. If silver nitrate acts on a dimetallic orthophosphate—for instance, Na2HPO4—it still gives Ag3PO4, nitric acid being disengaged: Na2HPO4 + 3AgNO3 = Ag3PO4 + 2NaNO3 + HNO3. When alcohol is added to silver orthophosphate, Ag3PO4, dissolved in syrupy phosphoric acid, it precipitates a white salt (the alcohol takes up the free phosphoric acid) having the composition Ag2HPO4, which is immediately decomposed by water into the normal salt and phosphoric acid.
[16] The researches of Thomsen showed that in very dilute aqueous solutions the majority of monobasic acids—nitric, acetic, hydrochloric, &c. (but hydrofluoric acid more and hydrocyanic less)—HX evolve the following amounts of heat (in thousands of calories) with caustic soda: NaHO + 2HX = 14; NaHO + HX = 14; 2NaHO + HX = 14; that is, if n be a whole number nNaHO + HX = 14 and NaHO + nHX = 14. Hence reaction here only takes place between one molecule of NaHO and one molecule of acid, and the remaining quantity of acid or alkali does not enter into the reaction. In the case of bibasic acids, H2R″ (sulphuric, dithionic, oxalic, sulphuretted hydrogen, &c.), NaHO + 2H2R″ = 14; NaHO + H2R″ = 14; 2NaHO + H2R″ = 28; nNaHO + H2R″ = 28; that is, with an excess of acid (NaHO + 2H′2R″) 14 thousand units of heat are developed, and with an excess of alkali 28. When phosphoric acid is taken (but not all tribasic acids—for instance, not citric) the general character of the phenomenon is similar to the preceding, namely, NaHO + 2H3PO4 = 14·7; NaHO + H3PO4 = 14·8; 2NaHO + H3PO4 = 27·1; 3NaHO + H3PO4 = 34·0; 6NaHO + H3PO4 = 35·3; or, in general terms, NaHO + nH3PO4 = 14 (approximately) and nNaHO + H3PO4 = 35 and not 42, which shows a peculiarity of phosphoric acid. In the case of energetic acids, when one equivalent (23 grams) of sodium (in the form of hydroxide) replaces one equivalent (1 gram) of hydrogen (with the formation of water and in dilute solutions), 14,000 heat units are evolved; and this is true for phosphoric acid when in H3PO4, Na or Na2 replaces H or H2, but when Na3 replaces H3 less heat is developed. This will be seen from the following scheme based on the preceding figures: H3PO4 + NaHO = 14·8; NaH2PO4 + NaHO = 12·3; Na2HPO4 + NaHO = 5·9; with Na3PO4 + NaHO, a very small amount of heat is evolved, as may be judged from the fact that Na3PO4 + 3NaHO = 1·3, but still heat is evolved. It must be supposed that in acting on phosphoric acid in the presence of a large quantity of water, a certain portion of the sodium hydroxide remains as alkali uncombined with the acid. Thus, on increasing the mass of the alkali, heat is still evolved, and a fresh interchange between Na and H takes place. Hence water shows a decomposing action on the alkali phosphates. The same decomposing action of water is seen, but to a less extent, with Na2HPO4, as may be judged both from the reactions of this salt and from the amount of heat developed by NaH2PO4 with NaHO. Such an explanation is in accordance with many facts concerning the decomposition of salts by water already known to us. Recent researches made by Berthelot and Louguinine have confirmed the above deductions made by me in the first edition (1871) of this work. At the present time views of this nature are somewhat generally accepted, although they are not sufficiently strictly applied in other cases. As regards PH3O4 it may be said that: on the substitution of the first hydrogen this acid acts as a powerful acid (like HCl, HNO3, H2SO4); on the substitution of the second hydrogen as a weaker acid (like an organic acid); and on the substitution of the third, as an alcohol, for instance phenol, having the properties of a feeble acid.
[17] Na2HPO4,12H2O has a sp. gr. 1·53. Poggiale determined the solubility in 100 parts of water (1) of the anhydrous ortho-salt Na2HPO4, and (2) of the corresponding pyro-salt Na4P2O7:—
| 0° | 20° | 40° | 80° | 100° | |
| I. | 1·5 | 11·1 | 30·9 | 81 | 108 |
| II. | 3·2 | 6·2 | 13·5 | 30 | 40 |
At temperatures of 20° to 100° the ortho-salt is so very much less soluble that this difference alone already indicates the deeply-seated alteration in constitution which takes place in the passage from the ortho- to the pyro-salts.
[18] The ammonium orthophosphates resemble the sodium salts in many respects, but the instability of the di- and tri-metallic salts is seen in them still more clearly than in the sodium salts; thus (NH4)3PO4, and even (NH4)2HPO4, lose ammonia in the air (especially when heated, even in solutions); NH4H2PO4 alone does not disengage ammonia and has an acid reaction. The crystals of the first salt contain 3H2O, and are only formed in the presence of an excess of ammonia; both the others are anhydrous, and may be obtained like the sodium salts. When ignited these salts leave metaphosphoric acid behind; for example, (NH4)2HPO4 = 2NH3 + H2O + HPO3. Ammonia also enters into the composition of many double phosphates. Ammonium sodium orthophosphate, or simply phosphate, NH4NaHPO4,4H2O, crystallises in large transparent crystals from a mixture of the solutions of disodium phosphate and ammonium chloride (in which case sodium chloride is obtained in the mother liquid), or, better still, from a solution of monosodium phosphate saturated with ammonia. It is also formed from the phosphates in urine when it ferments. This salt is frequently used in testing metallic compounds by the blow-pipe, because when ignited it leaves a vitreous metaphosphate, NaPO3, which, like borax, dissolves metallic oxides, forming characteristic tinted glasses.
When a solution of trisodium phosphate is added to a solution of a magnesium salt it gives a white precipitate of the normal orthophosphate Mg2(PO4)2,7H2O. If the trisodium salt be replaced by the ordinary salt, Na2HPO4, a precipitate is also formed, and MgHPO4,7H2O is obtained. It might be thought that the normal salt Mg3(PO4)2 would be precipitated if disodium phosphate was added to ammonia and a salt of magnesium, but in reality ammonium magnesium orthophosphate, MgNH4PO4,6H2O, is precipitated as a crystalline powder, which loses ammonia and water when ignited, and gives a pyrophosphate, Mg2P2O7. This salt occurs in nature as the mineral struvite, and in various products of the changes of animal matter. If we consider that the above salt parts with ammonia with difficulty, and that the corresponding salt of sodium is not formed under the same conditions (MgNaPO4,9H2O is obtained by the action of magnesia on disodium phosphate), if we turn our attention to the fact that the salts of calcium and barium do not form double salts as easily as magnesium, and remember that the salts of magnesium in general easily form double ammonium salts, we are led to think that this salt is not really a normal, but an acid salt, corresponding with Na2HPO4, in which Na2 is replaced by the equivalent group NH3Mg.