Why pH Matters
The pH of a solution — the negative logarithm of hydrogen ion concentration — profoundly affects every chemical reaction. Enzyme activity, protein structure, membrane permeability, and drug absorption all depend critically on pH. A change of just 0.3 pH units represents a doubling of hydrogen ion concentration. Life requires pH stability, and buffer solutions provide it.
The Henderson-Hasselbalch Equation
Published in 1917, the Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) is one of the most used equations in biochemistry. It tells us that the pH of a buffer equals its pKa when the weak acid and conjugate base are in equal concentrations. Shifting the ratio changes the pH predictably — a tenfold excess of base over acid raises pH by exactly 1 unit.
Buffer Capacity and Its Limits
A buffer doesn't have infinite capacity. As you add acid, it converts conjugate base (A⁻) into weak acid (HA). When the base is depleted, the buffer is overwhelmed and pH drops sharply. This simulation lets you see the titration curve — the characteristic S-shape where pH changes slowly in the buffer zone (±1 of pKa) but plummets once the buffer is exhausted.
Buffers in the Human Body
Blood pH must stay between 7.35 and 7.45 — a remarkably narrow range maintained by three buffer systems working in concert. The bicarbonate buffer (H₂CO₃/HCO₃⁻) is the primary system, supplemented by phosphate buffers and protein buffers (including hemoglobin). When these systems fail — in diabetic ketoacidosis, kidney failure, or respiratory arrest — the consequences are life-threatening.