The solubility of $\ce{PbSO4(s)}$ increase with the addition of $\ce{H2SO4}.$ Why?
I don't quite understand this. When dissolving $\ce{PbSO4(s)}$ we get the equilibrium equation:
$$\ce{PbSO4(s) <=> Pb^2+ + SO4^2-}$$
I know $\ce{H2SO4}$ is a strong acid, that will dissociate completely (or almost completely), like this:
$$ \begin{align} \ce{H2SO4 &-> H+ + HSO4-}\\ \ce{HSO4- &-> H+ + SO4^2-} \end{align} $$
So as far as I can see, the concentration of $\ce{SO4^2-}$ increases, which should give a lower solubility of $\ce{PbSO4(s)}$ (common ion effect). What am I missing?