Acid-Base Strength and Dissociation

In chemistry, 'strong' doesn't mean 'concentrated.' It refers to dissociation—how easily an acid breaks into ions. Strong acids (like $HCl$ or $H_2SO_4$) are 'all-in' players; they dissociate $100\%$ in water. Every molecule releases its $H^+$ ion.

Weak acids, however, exist in a state of chemical equilibrium. Only a small fraction of molecules break apart, while the rest stay together as whole molecules. This is why a $0.1 M$ solution of $HCl$ is far more corrosive than a $0.1 M$ solution of acetic acid (vinegar). The $HCl$ solution has a much higher concentration of free $H_3O^+$ ions.

To measure exactly how weak an acid is, we use the acid dissociation constant ($K_a$). For a generic acid reaction $HA(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + A^-(aq)$, the equilibrium expression is:

Note that water is omitted because it is a pure liquid. A **larger $K_a$ value indicates a stronger acid (more products/ions), while a smaller $K_a$** indicates a weaker acid. Bases follow the same logic using $K_b$ to measure the production of $OH^-$ ions.

Calculating the pH of a weak acid is more complex than a strong acid because we don't know the final $[H_3O^+]$ immediately. We use an ICE Table (Initial, Change, Equilibrium).

For an initial concentration $C$, we assume $x$ amount dissociates. At equilibrium, we have $[H_3O^+] = x$ and $[HA] = C - x$. If $K_a$ is very small (specifically if $C/K_a > 400$), we can use the approximation rule, assuming $C - x \approx C$ to simplify the math.

The percent ionization tells us what percentage of the original acid molecules actually turned into ions. It is calculated as:

An intriguing property of weak acids is that as you dilute the solution (lower the initial concentration), the percent of molecules that ionize actually increases, even though the total concentration of $H_3O^+$ decreases. This is a direct result of Le Chatelier's Principle!