Calcium Orthophosphates

Three calcium orthophosphates exist and have an extensive literature, which is often contradictory owing to the failure to appreciate the consequences of combining a comparatively strong base with a weak acid. The importance of a careful study of their properties arises from a widespread use of phosphates as fertilisers. Especially is it necessary to examine their behaviour with water. There have been considerable discrepancies between the results of solubility determinations, because the extent of the decomposition of the phosphates by water has not been taken fully into consideration. Both tricalcium and monocalcium phosphates give an acid solution when treated with water, the amount of acid depending on temperature and on the ratio of the mass of the water to the mass of the solid phosphate. In the case of the monocalcium phosphate the dicalcium salt is often precipitated. The latter salt is comparatively little affected, but, by treatment with successive quantities of water, the residue is ultimately transformed, with more or less completeness, into tricalcium phosphate.

In more recent times the ternary system, lime: phosphoric acid: water, has been carefully studied with a view to determining the conditions of stability of the different compounds. The problem is complicated by the following factors: the existence of several different hydrates, the precipitation of the phosphates in both gelatinous and crystalline forms, and the slowness with which equilibrium is established. Bassett1 determined three quintuple points and formulated the changes taking place at these points by the following equations, the portions in brackets indicating the composition of the solutions in contact with the solid phases: -

At 21° C., CaHPO₄.2H₂O + 0.138CaH₄(PO₄)₂.H₂O = 1.0747CaHPO₄ + (2.374H₂O + 0.10007P₂O₅ + 0.0627CaO).
At 36° C., CaHPO₄.2H₂O = 0.9985CaHPO₄ + 0.000384Ca₃(PO₄)₂.H₂O + (2.00036H₂O + 0.00033CaO + 0.00036P₂O₅).
At 152° C., CaH₄(PO₄)₂.H₂O = 0.495CaH₄(PO₄)₂ + 0.427CaHPO₄ + (1.80H₂O + 0.292P₂O₅ + 0.078CaO).

The solubility of the monocalcium phosphate is much greater than that of either of the other two. The phosphates are soluble in all strong acids, and the solubility is considerably increased by the presence of organic acids, for example, acetic, citric, and humic, a fact which has an important bearing on the availability of phosphates in the soil. Addition of phosphoric acid increases the solubility at first, but a very high concentration of acid results in the deposition of monocalcium phosphate. The presence of carbonic or sulphurous acids also makes the phosphates more soluble. Sulphurous acid probably forms a compound with tricalcium phosphate, 3CaO.SO₂.P₂O₅.2H₂O, consisting of microscopic hexagonal crystals.

Egg albumen increases the solubility, and also most neutral salts except calcium salts. This is especially the case with ammonium salts, probably owing to the formation of complex ions. Potassium chloride and sodium nitrate appear to increase the amount of lime in solution and to decrease the amount of phosphorus pentoxide.

Tricalcium Orthophosphate, Ca₃(PO₄), occurs naturally, combined with calcium chloride or fluoride, as the mineral apatite. It is also found in layers and nodules in many sedimentary rocks (phosphatic chalk, Florida pebbles, coprolites, etc.). The hills of Christmas Island, and of many other Pacific islands, are capped by layers of pure tricalcium phosphate produced by the action of guano deposits on the coral rocks. There is a large variation in the phosphate content of these sources, which are worked largely for fertiliser. Florida hard rock contains as much as 75-80 per cent, of tricalcium phosphate, the soft sandy deposits of North Africa 57-65 per cent., and the deposits worked in France and Belgium 40-50 per cent. Before the War the production was gradually increasing. The total world production in 1907 was not quite five million long tons, and in 1913 it was over seven million.

Calcium phosphate constitutes the chief part of the mineral matter of teeth and bones, being 50-60 per cent, of the whole, or about 85 per cent, of the ash. A deficiency in phosphates results in diseased or rachitic bones. Milk also contains calcium phosphate.

Calcium phosphate, along with basic phosphates, is a by-product of the basic Bessemer process.

It is obtained as a white gelatinous precipitate by the action of excess of lime on phosphoric acid solution, or by the precipitation of a calcium salt by a tertiary alkali phosphate, or by an acid phosphate in ammonia solution. On drying, an amorphous powder is produced containing varying quantities of water, 1 molecule, 2 molecules, 5 or 5½ molecules.

The melting-point is 1670° C.

If calcium phosphate is heated to 1200° C. with silica and carbon calcium silicate and phosphorus are obtained. This reaction is the basis of a commercial method for the production of phosphorus.

A colloidal solution of calcium phosphate can be prepared by adding slowly, and with continuous shaking, a hot normal solution of calcium chloride containing a definite amount of protective colloid for example, gelatine, gum arabic, blood serum, or starch. It can also be obtained by the action of a solution of orthophosphoric acid on a solution of calcium hydroxide containing gelatine. It forms an opaque solution, bluish-white by reflected light, and might prove of value in therapeutics.

A sample of calcium phosphate, showing the optical phenomena associated with liquid crystals, has been obtained.

Apatite, 3Ca₃ (PO₄)₂. Ca(Cl,F)₂, analogous with the natural mineral form, can be prepared artificially by the action of calcium phosphate on a mixture of calcium fluoride and chloride with a little ammonium chloride at red heat. On treatment with water, crystals in the form of regular hexagonal prisms are isolated. Fluorapatite, chlorapatite, and brom-apatite, containing only one halogen each, have been prepared. Bassett regards the apatites as complex compounds in which the chlorine is united directly with the phosphorus. This view seems to be justified by the existence of a compound of calcium oxide with phosphorus oxy-chloride, CaO.2POCl₃, which may be regarded as a halogen derivative of CaO.2PO(OH)₃, or CaH₄P₂O₈.H₂O.

Other chloro- and bromo-phosphates, of composition Ca₃(PO₄)₂. CaCl₂, and Ca₃(PO₄)₂.CaBr₂, have been obtained.

Dicalcium Orthophosphate, CaHPO₄, is found naturally as the mineral monetite, combined with 2 molecules of water of crystallisation as the mineral brushite, and also as metabrushite. It is also present in teak wood, and forms the chief constituent of stone separating in certain diseases of the kidneys.

It is obtained by precipitating an acid solution of a calcium salt with neutral sodium phosphate, or by the action of water on mono-calcium phosphate, or of an acid on tricalcium phosphate. Artificial brushite of density 2.317 may be obtained by concentration of a solution of the precipitated phosphate in 25 per cent, acetic acid, and monetite by heating together brushite and water in a closed tube at 150° C.

Whether the dicalcium salt is obtained in orthorhombic crystals combined with 2 molecules of water of crystallisation, or whether it is anhydrous, depends on the conditions of preparation. Precipitated at 100° C., or by alcohol, or from concentrated solutions at lower temperatures, it is anhydrous. At 25° C. the stable solid phase in contact with solutions of concentrations up to 15 per cent, phosphorus pentoxide is CaHPO₄.2H₂O. According to Quartaroli, there are two different forms of the dihydrate, one amorphous and the other crystalline, and possessing different chemical properties.

Other hydrates have been mentioned, but they probably do not exist as single chemical individuals.

By the combined action of water on dicalcium phosphate, microscopic crystals are obtained which appear to have the composition 8CaHPO₄.Ca₃(PO₄)₂.2H₂O. A phosphate of composition P₂O₅.2CaO.P₂O₅.3CaO.10H₂O crystallises from solutions containing 0.375 – 0.870 grm. of phosphorus pentoxide per litre in the state of monocalcium phosphate.

H₂O.Monocalcium Orthophosphate, Ca(H₂PO₄)₂.H₂O, has been found in a phosphatic rock in Algeria. It can be obtained by evaporation of a solution of di- or tri-calcium phosphate in phosphoric acid. It crystallises with 1 molecule of water in orthorhombic plates of density 2.02. When pure it is unaffected by exposure to the air, but it is often hygroscopic through contamination with phosphoric acid. In order that this salt may be the stable phase in contact with solutions at 25° C., the latter must contain more than 317 grm. of phosphorus pentoxide per litre.

When heated, the crystallised salt melts at 152° C. If it is allowed to lose water it is transformed into the metaphosphate. By slowly raising the temperature to 170°-180° C., pouring off the clear liquid after a short time, cooling the residue, and finally washing with acetone, triclinic crystals of anhydrous monocalcium phosphate are obtained. It may also be prepared by crystallisation at 160° C. from a solution of calcium carbonate in orthophosphoric acid if the ratio P₂O₅/CaO be sufficiently large. The optimum value is 4.6. If it is less than 3 the monohydrate separates. According to Spring, pressure causes the transformation of mono- into di-calcium phosphate.