Acid-Base Titration

The Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) is one of the most used equations in biochemistry. Blood plasma is buffered at pH 7.4 by the carbonic acid/bicarbonate system (pKa 6.1): even a 1 mmol/L change in CO₂ would be lethal without it. Pharmaceutical formulations, enzyme assays, and electrophoresis all depend on buffers tuned to the right pKa. The equation is only valid in the buffer region (roughly pKa ± 1.5 pH units) — outside that range the approximation breaks down and exact equilibrium expressions must be used.

A titration answers a simple question: how much acid is in this solution? You add base drop by drop until it exactly neutralises the acid — the equivalence point — and because you know the base's concentration and volume, you can work backwards to the acid's. The trick is the pH: right around that point one extra drop swings the pH by several units, so an indicator's colour change pinpoints the endpoint. Drag the V_b slider to watch the swing.

Raise the pKa slider and watch the jump at the equivalence point flatten out. As a rule, an acid gives a sharp, usable endpoint in water only if its Ka is above about 10⁻⁸ (pKa below ~8); weaker than that, the conjugate base grabs so many protons back from water that the curve has no clear step for any indicator to catch. Chemists get around this by titrating such acids in non-aqueous solvents that don't compete with the reaction.