A polyprotic acid can donate more than one proton per molecule. The protons are normally transferred in separate equilibrium steps, each with its own acid-ionization constant.

For a diprotic acid $\mathrm{H_2A}$,

$$\mathrm{H_2A+H_2O\rightleftharpoons H_3O^+ + HA^-}$$

has constant

$$K_{a1}=\frac{[\mathrm{H_3O^+}][\mathrm{HA^-}]}{[\mathrm{H_2A}]},$$

while

$$\mathrm{HA^-+H_2O\rightleftharpoons H_3O^+ + A^{2-}}$$

has

$$K_{a2}=\frac{[\mathrm{H_3O^+}][\mathrm{A^{2-}}]}{[\mathrm{HA^-}]}.$$

Successive dissociations usually become less favorable:

$$K_{a1}>K_{a2}>K_{a3}>\cdots.$$

After one proton has been removed, the remaining species is more negatively charged, so removing another positive proton is generally less favorable.

For example, carbonic acid undergoes

$$\mathrm{H_2CO_3+H_2O\rightleftharpoons H_3O^+ + HCO_3^-}$$

followed by

$$\mathrm{HCO_3^-+H_2O\rightleftharpoons H_3O^+ + CO_3^{2-}}.$$

The intermediate $\mathrm{HCO_3^-}$ is amphiprotic: it can accept a proton to reform carbonic acid or donate one to form carbonate.

When the successive $K_a$ values differ by several orders of magnitude, the first dissociation often dominates the hydronium concentration of an ordinary acid solution. Later dissociations can then be treated as smaller corrections. This is why many introductory pH calculations for polyprotic acids begin with the first ionization rather than solving every equilibrium simultaneously.

However, later proton-transfer steps become essential when the pH is changed over a wide range, as in buffers, titrations, carbonate chemistry, phosphate chemistry and many biochemical systems.

Polyprotic behavior is therefore not a special replacement for ordinary acid equilibrium. It is a sequence of coupled Brønsted-Lowry equilibria involving successive conjugate acid-base pairs.