The common normal calcium phosphate, Ca3(PO4)2, occurs in minerals, in animals, especially in bones, and also probably in plants, although the ash of many portions of plants, as a rule, contains less lime than the formation of the normal salt requires. Thus 100 parts of the ash (from 5,000 parts of grain) of rye grain contain 47·5 of phosphoric anhydride and only 2·7 of lime, and even the ash of the whole of the rye (including the straw) contains twice as much phosphoric anhydride as lime, and the normal salt contains almost equal weights of these substances. Only the ash of grasses, and especially of clover, and of trees, contains in the majority of cases more lime than is required for the formation of Ca3P2O8. This salt, which is insoluble in water, dissolves even in such feeble acids as acetic and sulphurous, and even in water containing carbonic acid. The latter fact is of immense importance in nature, since by reason of it rain water is able to transfer the calcium phosphates in the soil into solutions which are absorbed by plants. The solubility of the normal salt in acids takes place by virtue of the formation of an acid salt, which is evident from the quantity of acid required for its solution, and more especially from the fact that the acid solutions when evaporated give crystalline scales of the acid calcium phosphate, CaH4(PO4)2, soluble in water. This solubility of the acid salt forms the basis of the treatment by acids of bones, phosphorites, guano, and other natural products containing the normal salt and employed for fertilising the soil. The perfect decomposition requires at least 2H2SO4 to Ca3(PO4)2, but in reality less is taken, so that only a portion of the normal salt is converted into the acid salt. Hydrochloric acid is sometimes used. (In practice such mixtures are known as superphosphates). Certain experiments, however, show that a thorough grinding, the presence of organic, and especially of nitrogenous, substances, and the porous structure of some calcium phosphates (for example, in burnt bones), render the treatment of phosphoric manures by acids superfluous—that is, the crop is not improved by it.
[19] In this sense the ortho-acid itself might be regarded as an anhydro-acid, counting P(HO)5 as the perfect hydrate, if PH5 existed; but as in general the normal hydrates correspond with the existing hydrogen compounds with the addition of up to 4 atoms of oxygen, therefore PH3O4 is the normal acid, just as SH2O4 and ClHO4; while NHO3, CH2O3 are meta-acids, or higher normal acids (NH3O4 and CH4O4) with the loss of a molecule of water.
In order to see the relation between the ortho-, pyro-, and metaphosphoric acids, the first thing to remark in them is that the anhydride P2O5 is combined with 3, 2, and 1 molecules of water. In the absence of data for the molecular weight of ortho- and pyrophosphoric acids it is necessary to mention that all existing data for metaphosphoric acid indicate (Note [21]) that its molecule is much more complex and contains at least H3P3O9 or H6P6O18. The explanation of the problems which here present themselves can, it seems to me, be only looked for after a detailed study of the phenomena of the polymerisations of mineral substances, and of those complex acids, such as phosphomolybdic, which we shall hereafter describe (Chapter [XXI.]) A similar instance is exhibited in the solubility of hydrate of silica (produced by the action of silicon fluoride on water) in fused metaphosphoric acid, with the formation, on cooling, of an octahedral compound (sp. gr., 3·1) containing SiO2,P2O5. A certain indication (but no proof) that ordinary orthophosphoric acid is polymerised is given by Staudenmaier (1893), who obtained a salt, K5H4P3O12, by the action of a solution of KH2PO4 upon K2CO3; and a compound, KH3P2O8, corresponding to the doubled molecule of H3PO4, by the action of KH2PO4 upon H3PO4 itself.
[19 bis] According to Watson (1893) the ortho-acid is partially transformed into the pyro-acid at 230°, whilst at 260° the latter begins to volatilise. At 300° the meta-acid only is formed.
[20] The method of preparation of the acid itself consists in converting the sodium salt, Na4P2O7, by double decomposition with water and a salt of lead, into insoluble lead pyrophosphate, Pb2P2O7, which is then suspended in water and decomposed by sulphuretted hydrogen; lead sulphide is thus precipitated, and pyrophosphoric acid remains in solution. This solution cannot be heated, or the pyro-acid will pass into the ortho-, but must be evaporated under the receiver of an air-pump. It concentrates to a syrup and crystallises, and when ignited in this form loses water, and forms metaphosphoric acid. It resembles orthophosphoric acid in many respects; its salts with the alkalis are also soluble, and the others insoluble in water but soluble in acids. When heated in solution with acid it gives orthophosphoric acid, as well as when fused with an excess of alkali.
Witt heated ammonium chloride with phosphoric acid (hydrochloric acid was evolved), ignited the residue to drive off ammonia, and obtained pyrophosphoric acid in the residue.
[21] As when using phenolphthalein as an indicator in neutralising by an alkali metaphosphoric acid is monobasic, and orthophosphoric acid is bibasic, it is possible by means of this difference to follow the transition of meta- into orthophosphoric acid. Sabatier (1888) carried on an investigation of this nature, and found that the rate of transformation is dependent on the temperature, and is subject to the general laws of the rate of chemical transformations which belongs to physical chemistry.
Metaphosphoric acid has a particular interest in respect to the variations to which its salts are subject. The metaphosphates are formed by the ignition of the acid orthophosphates, MH2PO4, or MNH4HPO4, or of the acid pyrophosphates, M2H2P2O7, or M2(NH4)2P2O7, water and ammonia being given off in the process. The properties of the metaphosphates, which have a similar composition to nitrates—for instance, NaPO3, or Ba(PO3)2—vary according to the duration of the ignition to which the ortho-, or pyrophosphates from which they are prepared have been subjected. When the salts NaH2PO4 or NH4NaHPO4 are strongly ignited, a salt NaPO3 is formed, which deliquesces in the air, and gives a gelatinous precipitate with salts of the alkaline earths. But, as Graham (in 1830–40), and many others, especially Fleitmann and Henneberg (in 1840–50), and Tamman (in the nineties), observed, under other conditions the salts of the same composition acquire other properties. The above chemists recognise five polymeric forms of metaphosphates, (HPO3)n. We will follow the nomenclature and researches of Fleitmann.
Monometaphosphoric acid. The salts are distinguished for their insolubility in water; even the salts NaPO3, KPO3, are insoluble. They are obtained by igniting the monometallic orthophosphates—for example, RH2PO4—up to the temperature at which all water is evolved (316°), but not to fusion. No double salts are known.
Dimetaphosphoric acid, on the contrary, easily forms double salts—for example, KNaP2O6, and also the copper potassium salt, &c. The copper salt is obtained by evaporating a solution of copper oxide in orthophosphoric acid. A blue ortho-salt, CuRHO4, first separates from the solution, then a light-blue pyro-salt, Cu2P2O7; and above 350°, when metaphosphoric acid itself begins to volatilise, the dimetaphosphate, CuP2O6, is formed. The residue is washed with water, and decomposed with a hot solution of sodium sulphide, when the sodium salt, Na2P2O6, is obtained in solution. This salt, when evaporated with alcohol, gives crystals containing 2 mol. H2O, which, however, retain their solubility (in 7 parts of water) after the water is driven off at 100°. When fused, these crystals give a deliquescent salt (hexa-metaphosphate). The solution of the salt has a neutral reaction, which only after prolonged boiling becomes acid, owing to the formation of orthophosphate, NaH2PO4. The soluble salts of dimetaphosphoric acid give the insoluble silver salt, Ag2P2O6, with silver nitrate, and a precipitate of BaP2O62H2O with barium chloride.