Buffer

When a weak acid is dissolved in water, only a few of its molecules dissociate to form only a few hydrogen ions; the rest of the acid molecules remain as undissociated, neutral molecules that do not affect the pH. For example, whenever acetic acid is added to water, the following three species will be in the solution:

To make a buffer solution out of this system, we can add many more acetate ions to the solution in the form of sodium acetate, which is a strong electrolyte and dissociates completely. Ignoring the sodium ions that come along with the sodium acetate because they do not affect the acidity at all, we then have:

This solution will resist having its hydrogen ion concentration changed. To see how that works, first consider what would happen if we were to add some acid-some extra hydrogen ions-to this solution. According to LeChâtelier's principle, the equilibrium will be shifted to the left. That is, the added hydrogen ions will react with some of the acetate ions to form more acetic acid molecules. The result is that almost all of the added hydrogen ions are used up to form "harmless" neutral molecules; they therefore are not available to increase the acidity of the solution. The solution has resisted having its pH lowered more than a little bit.

What if we were to add some base-hydroxide ions to the buffer solution? Hydroxide ions to react with hydrogen ions, because the resulting molecule, H₂O, is so stable.

Therefore, the added hydroxide ions will quickly remove hydrogen ions from the buffer solution, which according to LeChâtelier's principle will then shift its equilibrium to the right, making more acetate ions out of acetic acid molecules. Thus, the added hydroxide ions will have been used up, and only "harmless" acetate ions will have been formed. (Acetate ions are slightly basic, however, so the pH of the buffer solution does increase slightly.)