Common Ion Effect explanations From my textbook, it is stated that:
e.g. If we add a basic salt $\ce{NaF}$, it will fully dissociate to give $\ce{F-}$. Then if we add $\ce{HF}$, since it is a weak acid, the high concentration of $\ce{F-}$ already present in solution shifts the equilibrium to the left, hence fewer $\ce{H+}$  forms than expected so higher $\mathrm{pH}$.
But can I explain like this: after dissociation of $\ce{NaF}$, the formed $\ce{F-}$ reacts with water to give $\ce{OH-}$ as the following equation:
$$\ce{F- + H2O &lt;=&gt; HF + OH-}$$. Due to the presence of $\ce{OH-}$, upon addition of $\ce{HF}$, the $\mathrm{pH}$ will be higher than expected. Are there any flaws in this explanation?
Both your paragraphs are valid, as all present components are balanced by 3, mutually linked equilibrium reactions, with the respective equilibrium constants:
$\ce{HF + H2O &lt;=&gt; F- + H3O+}$, $K_\mathrm{a}$
$\ce{F- + H2O &lt;=&gt; HF + OH-}$, $K_\mathrm{b}$
$\ce{2 H2O &lt;=&gt; H2O+ + OH-}$, $K_\mathrm{w}$
with the relation $K_\mathrm{a} \cdot K_\mathrm{b} = K_\mathrm{w}$
$\mathrm{pH}= \mathrm{p}K_\mathrm{a} + \log{\frac{\ce{[F-]}}{\ce{[HF]}}}$