Which of the following salts would show an appreciable pH-dependent solubility?

Question

a. AgCl
b. Mg(OH)2
c. BaF
d. both b and c
e. all of the above

I put c as my answer (not sure if its right, but can you explain how you would solve this)

Answer 1

The correct answer is d; i.e., both b and c.
You look at the solubility as well as if the H^+ reacts with cation or anion to produce a weak acid. If so, then it is pH dependent.
AgCl, for example, ==> Ag^+ + Cl^-.
Adding H^+ might try to form HCl but that is not a weak acid; therefore, changing the pH is not a problem. (It is generally understood that the solution cannot become too basic or pptn of Ag2O will occur.)
Mg(OH)2 obviously will be affected by the pH because of the OH^- ion and H^+ + OH^- => H2O.

BaF2 (it's BaF2 and not BaF) is affected because of the possible formation of HF, a weak acid.
BaF2(s) ==> Ba^+2 + 2F^-
Note that F^- + H^+ ==> HF
Thus as the solution is made more acidic, the H^+ removes the F^- (actually both remove each other). Remember Le Chatelier's Principle: when we remove the F^- (by adding more H^+), the reaction (the solubility) shifts to the right which means that more of the BaF2(s) dissolves to form more Ba^+ and 2F^-. Therefore, the more acid we add, the more F^- is removed, and the more the reaction shifts to the right, and the more soluble BaF2 is in the solution.