Changes during Neutralisation

In the neutralisation of phosphoric acid with lime there is only a gradual increase in the pH value until more than one equivalent of lime has been added; that is until some 33 per cent, of the acid has been neutralised, corresponding to the formation of CaH₄P₂O₈. The pH value then begins to increase more rapidly with further addition of lime and a precipitation of dicalcium phosphate, Ca₂H₂P₂O₈, may occur if sufficient time is allowed. When 39 to 45 per cent, of the acid has been neutralised there is a sharp drop in the pH value, which is due probably to the precipitation of dicalcium phosphate from supersaturated solution. This not only leaves the solution relatively poorer in hydroxyl ions through the removal of HPO₄⁼, since by hydrolysis

HPO₄⁼ + H⁺ + OH → H₂PO₄^- + OH^-

but also diminishes the ratio CaO/P₂O₅ in the solution, since a solid corresponding to 67 per cent, neutralisation is being removed from a liquid corresponding to 45 per cent. The precipitation may be delayed until about equivalents of lime have been added and the pH value has become 6.7. Further addition of lime then produces a gradual increase in alkalinity, absolute neutrality (pH = 7) being reached after the addition of about 2 equivalents, i.e. at 67 per cent, neutralisation. A slight further addition of lime then produces a sharp drop in the alkalinity, which is not observed however when the neutralisation is carried out in the presence of dissolved sucrose. This " kink " in the curve is probably due to supersaturation. The process of deposition with increase of acidity in the presence of precipitated solids would probably continue much further if sufficient time were allowed, since it has been calculated from the results of Bassett that the solution in equilibrium with dicalcium phosphate or even with tricalcium phosphate has a much greater acidity (pH approximately 5.5) than the solutions in which these precipitates are first produced. A continuation of the titration with lime up to 3 equivalents (100 per cent, neutralisation) gives only a slight increase in alkalinity, which becomes somewhat greater after the addition of 3 equivalents. In contradistinction then to the alkalies and even baryta and strontia, calcium hydroxide apparently does not give high alkalinities when in the presence of precipitated calcium phosphate. This is due not only to the sparing solubility of Ca(OH)₂, but also to its combination with Ca₃P₂O₈ to give hydroxy-apatite, 3Ca₃P₂O₈.Ca(OH)₂.

The practical significance of these observations is that phosphates can only remain freely soluble in soils with a relatively high acidity (pH less than 5.5); the solid present in contact with such solutions is either CaH₄P₂O₈.H₂O or Ca₂H₂P₂O₈ in a finely divided and anhydrous state. Neutral or even faintly acid solutions (pH 5.5 to 7.0) will contain but little dissolved phosphate, being in equilibrium with CaHPO₄.2H₂O, Ca₃(PO₄)₂ or Ca(OH)₂.3Ca₃P₂O₈, or a mixture of these solids. The solubility of these solids is however sufficient for the requirements of plants, since it is found that the amount of phosphorus in the extract even of a rich soil is only of the order of 1 milligram per litre, while soils of average fertility may contain only 0.1 to 0.2 milligram per litre. The fact that this very low concentration appears to be sufficient for the growth of plants makes it probable that these are able to use insoluble phosphate, as has also been pointed out by N. M. Comber and others. It is important that there should be a sufficient reserve of phosphate in the soil which may gradually become available. The removal of phosphoric acid by plants increases the amount of basic phosphate and should be compensated by the addition of an acid-producing fertiliser such as ammonium sulphate. The value of phosphates in a really soluble form has long been recognised.