Calcium phosphate is a mineral compound that forms the hard structures of bones and teeth. It is known for its extremely low solubility in water, meaning it does not readily dissolve under normal conditions. However, adding an acid dramatically increases the amount of calcium phosphate that can dissolve. This increased dissolution is a chemical reaction involving one of the mineral’s component ions.
Understanding Calcium Phosphate’s Solubility
Calcium phosphate is an ionic solid, held together by strong electrostatic attractions between positive calcium ions (\(Ca^{2+}\)) and negative phosphate ions (\(PO_4^{3-}\)). When placed in water, a small amount of the solid dissociates into these individual ions. This process continues until the solution reaches saturation, where no more solid can dissolve.
This state of saturation is known as a chemical equilibrium. In this balanced state, the rate of the solid dissolving into ions is exactly matched by the rate of the ions rejoining to form the solid. The concentration of dissolved calcium and phosphate ions remains constant. The low natural solubility of calcium phosphate is a result of this equilibrium strongly favoring the solid, undissolved form.
Phosphate: A Chemical Base
The key to understanding the acid effect lies with the phosphate ion (\(PO_4^{3-}\)), which is a chemical base. A chemical base has a strong tendency to accept a proton, which is a positively charged hydrogen ion (\(H^+\)). The phosphate ion is highly reactive toward these protons.
When an acid is added to water, it increases the concentration of free hydrogen ions (\(H^+\)) in the solution. The phosphate ion immediately reacts with available \(H^+\) ions, effectively removing them from the solution. This reaction is a multi-step process where the phosphate ion can accept up to three protons.
The initial reaction converts the phosphate ion (\(PO_4^{3-}\)) into the hydrogen phosphate ion (\(HPO_4^{2-}\)). If more acid is present, this new ion can accept another proton to become the dihydrogen phosphate ion (\(H_2PO_4^{-}\)), and finally, it can fully convert into phosphoric acid (\(H_3PO_4\)). These newly formed phosphate species are highly soluble in water and do not readily combine with calcium to re-form the solid.
How Acid Addition Shifts the Equilibrium
The addition of acid leverages the basic nature of the phosphate ion to disrupt the solubility equilibrium established by the solid calcium phosphate. When the acid is introduced, the hydrogen ions scavenge the dissolved phosphate ions (\(PO_4^{3-}\)) from the solution to form the highly soluble protonated species. This reaction lowers the concentration of the dissolved phosphate ions.
According to Le Châtelier’s principle, a system at equilibrium will shift to counteract any stress applied to it. The stress here is the removal of the dissolved phosphate ion from the saturated solution. To relieve this stress and restore the concentration of the phosphate ion, the system must produce more of it.
The only way for the system to produce more dissolved phosphate ions is for the solid calcium phosphate to dissolve further. The equilibrium is thus forced to “shift right,” meaning the reaction favors the dissolving of the solid and the creation of more dissolved calcium and phosphate ions. This continuous removal of the phosphate ion by the acid drives more of the calcium phosphate solid into the solution, increasing solubility.