Why does the addition of acid increase the solubility of calcium phosphate (Ca3(PO4)2) in water?

• A. Acid causes the precipitation of calcium ions, removing them from solution and increasing solubility.
• B. Acid reacts with phosphate ions to form H3PO4, reducing the concentration of PO4^{3-} and shifting the dissolution equilibrium toward more dissolved Ca3(PO4)2.
• C. Acid increases the temperature of the solution, which directly increases the solubility of Ca3(PO4)2.
• D. Acid neutralizes calcium ions, forming CaCl2 and increasing the solubility of Ca3(PO4)2.

Answer: (B)

Verified step by step guidance

1. Identify the dissolution equilibrium of calcium phosphate in water: \(\mathrm{Ca_3(PO_4)_2 (s) \rightleftharpoons 3 Ca^{2+} (aq) + 2 PO_4^{3-} (aq)}\).
2. Understand that the solubility depends on the concentrations of \(\mathrm{Ca^{2+}}\) and \(\mathrm{PO_4^{3-}}\) ions in solution, governed by the solubility product constant \(K_{sp}\).
3. Recognize that adding acid introduces \(\mathrm{H^+}\) ions, which react with phosphate ions to form various protonated phosphate species such as \(\mathrm{HPO_4^{2-}}\), \(\mathrm{H_2PO_4^-}\), and ultimately \(\mathrm{H_3PO_4}\).
4. This reaction reduces the free \(\mathrm{PO_4^{3-}}\) ion concentration in solution, disturbing the dissolution equilibrium and causing more \(\mathrm{Ca_3(PO_4)_2}\) solid to dissolve to restore equilibrium according to Le Châtelier's principle.
5. Conclude that the increased solubility is due to the acid removing \(\mathrm{PO_4^{3-}}\) ions from solution by protonation, shifting the equilibrium toward dissolution, rather than precipitation or temperature effects.