Henderson–Hasselbalch equation
/ HEN-der-sun HASS-el-balk /
If you have ever followed a recipe that says mix this much of one thing with that much of another to hit a target, you already grasp the spirit of this equation. It is the recipe card for a buffer: tell it how much weak acid and how much conjugate base you have, and it tells you the pH you will get.
The Henderson–Hasselbalch equation states pH = pKa + log([A-]/[HA]) — the pH equals the acid's pKa plus the logarithm of the ratio of conjugate base to acid. When the two are present in equal amounts, the log term is zero and the pH simply equals the pKa. Tilt the ratio toward the base and the pH rises; tilt it toward the acid and the pH falls.
It matters because it turns buffer design into arithmetic: pick a weak acid whose pKa is near your target pH, then adjust the base-to-acid ratio to fine-tune. The honest caveats are that it is a convenient approximation — it uses concentrations rather than true activities, and it breaks down for very dilute buffers or near the extremes of the pH scale.
To make a pH 7.4 blood-like buffer, a chemist picks a weak acid with pKa near 7.2 and sets the conjugate base slightly higher than the acid, so the log term nudges the pH up the small extra amount needed.
pH = pKa + log(base/acid): the recipe that turns a ratio into a pH.
The equation is just the acid-dissociation equilibrium rearranged and put in logarithmic form — it is not a new law. It is most accurate when the base-to-acid ratio lies between about 1:10 and 10:1, the range where a buffer actually works well.