A buffer solution contains appreciable amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. Its defining behavior is that small additions of strong acid or strong base cause much smaller pH changes than they would in an unbuffered solution.

Consider a buffer containing a weak acid $\mathrm{HA}$ and its conjugate base $\mathrm{A^-}$:

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

If a small amount of strong acid is added, the added hydronium is consumed mainly by the conjugate base:

$$\mathrm{A^-+H_3O^+\rightarrow HA+H_2O}.$$

If a small amount of strong base is added, hydroxide is consumed mainly by the weak acid:

$$\mathrm{HA+OH^-\rightarrow A^-+H_2O}.$$

The buffer therefore converts a strong added acid or base into one member of a weak conjugate pair. The concentrations of $\mathrm{HA}$ and $\mathrm{A^-}$ change, but the hydronium concentration changes relatively little while both components remain present in substantial amount.

For example, an acetic-acid/acetate buffer can absorb added $\mathrm{H_3O^+}$ through acetate and added $\mathrm{OH^-}$ through acetic acid. Neither component is inert: resistance to pH change comes from proton-transfer reactions coupled to the weak-acid equilibrium.

A buffer does not keep pH exactly constant and has finite capacity. If enough strong acid is added to consume most of the conjugate base, or enough strong base is added to consume most of the weak acid, buffering is exhausted and pH changes rapidly.

Buffers are therefore equilibrium systems with a practical stoichiometric reserve of both members of a conjugate acid-base pair. Their pH is controlled mainly by the acid strength and by the ratio of conjugate base to acid.