pH and Solubility

Every ionic solid has a dissolution equilibrium: solid <→ cation + anion, with some small Ksp. Adding acid only speeds up dissolving if the H+ you add has somewhere to go — specifically, if the anion in solution is willing to grab that H+. An anion reacts with H+ only if it's the conjugate base of a weak acid, because conjugate bases of weak acids are decent proton acceptors. When that anion gets consumed by H+, its concentration in solution drops, and by Le Chatelier's principle the dissolution equilibrium shifts forward to replace it, pulling more solid into solution.

Compare CaF2 and AgCl. F- is the conjugate base of HF, a weak acid, so F- reacts with added H+ to form HF, dropping [F-] and dragging more CaF2 into solution. Cl- is the conjugate base of HCl, a strong acid — strong acids dissociate completely, which means their conjugate base has essentially zero pull toward grabbing a proton back. Added H+ has nothing to react with, so AgCl's solubility barely budges in acid. Same logic explains limestone: CO3 2- is the conjugate base of the weak acid HCO3-, so it reacts with H+, which is exactly why acid rain eats away at marble and limestone statues while ordinary rain, with its tiny H+ concentration, does not.

This is really the common-ion effect run in reverse. Adding a common ion suppresses dissolving by pushing equilibrium backward; removing an ion by chemical reaction — here, converting the anion into its conjugate acid — pulls equilibrium forward. Same Le Chatelier logic, opposite direction, same equilibrium expression underneath.