Percent Ionization

Ka tells you how far a weak acid's equilibrium sits toward products, and it never changes at a given temperature no matter how much you dilute the acid. Percent ionization is different: it is the fraction of the starting acid that actually ionizes, and that fraction changes with concentration. Dilute a weak acid and a bigger percentage of it falls apart, even though Ka stays fixed. The reason traces back to Le Chatelier: diluting lowers the concentration of every species in HA <→ H+ + A-, but it lowers the ion concentrations more sharply relative to what equilibrium demands, so the reaction shifts right to restore the ratio Ka requires. With less HA around to start, that shift represents a larger slice of the original amount.

To calculate it, divide the equilibrium [H+] by the initial concentration of the acid and multiply by 100. Using the small-x approximation, [H+] = √(Ka times c), so percent ionization = √(Ka/c) x 100. That square root is the key: as c shrinks, Ka/c grows, so percent ionization climbs. Graphed against concentration, the curve falls steeply at first then flattens, since halving an already-small c has less effect on the ratio than halving a large one.

This sits right after Ka calculations and right before buffers, and it is the piece that keeps students from confusing a relative measure with an absolute one. Percent ionization describes how complete the reaction is relative to where it started; pH describes how much H+ is actually floating around. Both responses to dilution are real, but they answer different questions.