common-ion effect
Suppose a small shop already has a long queue of one kind of customer at the door. New customers of that same kind are far less likely to push in — the crowd is already full of their type. Ions in a solution feel something similar. If a salt is trying to dissolve into a solution that is already swimming with one of its own ions, that crowding shoves the dissolving back, and less of the salt goes in. That pushback is the common-ion effect.
It is really Le Chatelier's principle applied to a solubility (or dissociation) equilibrium. Adding extra of an ion that the equilibrium already contains — a 'common' ion — disturbs the balance, and the system responds by shifting away from making more of that ion. For a solubility equilibrium that means less solid dissolves; for a weak-acid equilibrium it means the acid dissociates even less.
Analysts use this deliberately. Washing a precipitate with a dilute solution of one of its own ions, instead of plain water, keeps the precipitate from re-dissolving and losing analyte. Buffers, too, lean on the common-ion effect to hold pH steady. The caveat to watch: adding a non-common 'inert' salt does the opposite — it raises ionic strength and can actually increase solubility a little, which is the salt effect, not the common-ion effect.
Silver chloride barely dissolves in pure water, but in 0.1 M sodium chloride it dissolves about a thousand times less still: the flood of chloride ions, shared with AgCl, holds nearly all of it as solid.
An added common ion suppresses solubility sharply.
Pile on far too much common ion and, for some salts, solubility can creep back up because the excess ion forms soluble complexes with the solid (for AgCl, soluble AgCl2- ions form in very high chloride). So the effect is reliable in moderation but can reverse at extremes.