[jakfak]

I'm trying to grow some lettuce hydroponically in beach sand.
I've gotten fed up with adjusting the pH constantly from the shells (calcium carbonate) in this medium, so want to dissolve it all out. I have a good supply of phosphoric acid and tried this on some of the sand. Much effervescence, of course, but I'm wondering what is in the filtered fluid when this reaction has ceased, with still some calcium carbonate left in excess. I wonder what the equilibrium pH would be here. And what the rough proportions of the anionic species would be present.
I understand the common ion effect, and le Chatelier's principle, but obviously tri-calcium phosphate is more soluble in phosphoric acid than distilled water. What principle is in effect here? Thanks for any light, regards,    Jak

[Borek]

In low pH PO₄^3- gets protonated.

[jakfak]

In low pH PO₄^3- gets protonated.

Thanks Borek
So my resultant solution is likely to be mainly calcium, and dihydrogen phosphate ions to saturation with an acid pH?
Roughly what sort of pH would not liberate carbon dioxide from available solid carbonate?
Regards,    Jak
ps Just as a general principle, if the anion could not be protonated, would the solubility product be the only factor influencing precipitation? Other than the common ion effect, has pH no effect on solubility?

[jakfak]

In low pH PO₄^3- gets protonated.

A further question occurred to me, and that is just why does the phosphate ion get protonated?
I can think of several reasons, but I would like the opinion of an expert, rather than relying on my elderly memory.
Perhaps the added hydrogen ions force the reaction to the direction where hydrogen ions are absorbed into a new ionic species? Does the increased solubility product of calcium ions and the protonated phosphate ions have any bearing on this reaction? Regards, Jak

[Borek]

Every acidic solution will liberate carbon dioxide, even at pH 5.

You have a rather complicated equilibrium involving multistep dissociation of two acids and precipitation of several weakly soluble salts (it is not just carbonate and phosphate, CaHPO₄ solubility is not that low either). Most likely some Ca/phosphates complexes are also present.

ps Just as a general principle, if the anion could not be protonated, would the solubility product be the only factor influencing precipitation? Other than the common ion effect, has pH no effect on solubility?

pH alone - if H⁺ doesn't take part in acid or ligand protonation - doesn't influence equilibrium directly. It's presence may change ionis strength of the solution, but as far as I understand you are working with solutions concentrated enough for that effect to be not too important.

why does the phosphate ion get protonated?

Phosphoric acid is a weak one (especially second and third protons) so it prefers to be protonated. The higher the pH, the lower the concentration of H⁺, the more likely the phosphoric acid to lost its protons, but at pH around 11 it will be mostly in HPO₄^3- form. These things can be easily calculated from acid dissociation constants.