Acid-Base Titration: Reading the Shape of a pH Curve

Adding base one drop at a time

A titration is a slow, controlled reaction: a strong base of known concentration, delivered from a burette, is added drop by drop to a weak acid of unknown concentration while a pH meter (or an indicator dye) tracks the solution. The resulting plot of pH against volume of base added is not a straight line — it has a distinctive S-shape with a flat middle stretch and a steep rise near the end, and every feature of that shape is set by the underlying equilibrium chemistry of the weak acid.

The equilibrium underneath everything

A weak acid HA does not dissociate fully in water; it settles into an equilibrium governed by its acid dissociation constant Ka:

HA ⇌ H⁺ + A⁻         Ka = [H⁺][A⁻] / [HA]

Henderson-Hasselbalch equation:
  pH = pKa + log( [A⁻] / [HA] )

Every drop of strong base added converts a little HA into its conjugate base A⁻, shifting the ratio [A⁻]/[HA] and, through the logarithm, the pH. Because the relationship is logarithmic, large changes in that ratio produce only modest changes in pH over most of the titration — which is exactly why the curve is flat in the middle.

Four landmarks on the curve

The titration curve of a weak acid with a strong base has four recognisable regions. The initial pH, before any base is added, is set purely by the weak acid's own equilibrium in water. The buffer region is the long, flat stretch where roughly comparable amounts of HA and A⁻ are both present; the solution resists pH change here because added base is consumed converting HA to A⁻ rather than accumulating as free hydroxide. The half-equivalence point, exactly halfway through the buffer region, is where [A⁻] = [HA] and Henderson-Hasselbalch collapses to pH = pKa — the single most direct way to read a weak acid's pKa off an experimental curve. The equivalence point, where moles of base added exactly equal the original moles of acid, is where the steep, near-vertical jump in the curve is centred.

Why the equivalence point isn't pH 7

At equivalence, essentially all of the original HA has been converted to A⁻ dissolved in water — but A⁻ is itself a weak base, and it reacts with water: A⁻ + H₂O ⇌ HA + OH⁻. That hydrolysis generates a small excess of hydroxide, so the equivalence-point pH sits above 7, not at it. The weaker the original acid (the smaller its Ka), the stronger its conjugate base and the further above 7 the equivalence point lands — a useful diagnostic in its own right when identifying an unknown acid from its titration curve.

Choosing an indicator

An indicator is itself a weak acid-base pair whose two forms have different colours, and it changes colour over a narrow pH window centred near its own pKa. Choosing the right one means matching that window to the pH at the equivalence point, not to pH 7 — for a weak acid titrated with a strong base, that means an indicator like phenolphthalein (colourless below about pH 8.2, pink above about pH 10) rather than one centred near neutral pH, because the equivalence point genuinely sits in the basic range.

Frequently asked questions

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