NB In this chapter we discuss the weak acids and weak bases. The concepts "strong" and "weak" in this chapter are relative and in fact mean "weak species rather strong" and "weak species rather weak"
The reactions $HB$ $+$ $H_2O$ $\rightleftarrows$ $ H_3O^+$ $+$ $B$ (acid dissociation) $B$ $+$ $H_2O$ $\rightleftarrows$ $ OH^-$ $+$ $HB$ (base hydrolysis) have shown us that for every Broenstedt acid $HB$ it is a base $B$
$HB$ weak acid, $B$ weak base: $(HB,B)$ is a couple $(weak \;acid,weak \;base)$
The relation $pK_a$ $+$ $pK_b$ $=$ $14$ shows that
To a strong acid it corresponds a weak base and vice versa.
See the → Table of acid-base couples
Example: Chloric acid ($pK_a$ $=$ $-1$) is strong, whereas its corresponding base, the chlorate ion ($pK_b$ $ = $ $14-(-1)$ $ = $ $15$) is weak.
The formula $K_a$ $=$ $\frac{[H_3O^+][B]}{[HB]}$ evaluates the distribution of $B$ and $HB$ in a given medium $pH$:
$\frac{[B]}{[HB]}$ $ =$ $ \frac{K_a}{[H_3O^+]}$
Example: At $pH=3$, for hypochlorous acid $pK_a=2$ and its corresponding base, hypochlorite ion, we have: $\frac{[ClO^-]}{[HClO]}$ $ =$ $ \frac{K_a}{[H_3O^+]}$ = $\frac{10^{-2}}{10^{-3}}$ $ = $ $10$ which means that there is ten times more hypochlorite than hypochlorous acid!