Upper secondary · 15 min

pH-metric titration and the buffer region

Titrate acetic acid with NaOH and watch the curve flatten where the $\ce{CH3COOH/CH3COO^-}$ pair buffers the pH — at half-equivalence pH = pKa.

Goal

Locate the buffer plateau and read pKa of acetic acid at the half-equivalence point (pH ≈ 4.76); explain why the slope is minimal there.

Apparatus and reagents

Beaker of 0.1 mol/L acetic acid, burette of 0.1 mol/L NaOH, calibrated pH meter with glass electrode, magnetic stirrer.

Procedure

Select the weak-acid option (acid = 1) with C = 0.1 mol/L.
Increase Vb step by step and record pH; mark where the curve is flattest — the buffer plateau.
Find the equivalence volume Ve at the steepest jump, then set Vb = Ve/2 and read pH: it equals pKa.
Compare the initial pH and the equivalence pH (>7) with a strong-acid titration (acid = 0).

What to observe

The curve starts higher than for a strong acid and runs through a nearly flat plateau centred on pH ≈ pKa before jumping steeply past equivalence to pH ≈ 12–13.
At Vb = Ve/2 the acid and its conjugate base are present in equal amounts and pH ≈ 4.76, the pKa of acetic acid; at equivalence the solution is basic because $\ce{CH3COO^-}$ is a weak base.

Explanation

As NaOH is added, $\ce{CH3COOH}$ is converted into $\ce{CH3COO^-}$ . While comparable amounts of both coexist, the Henderson–Hasselbalch relation $\mathrm{pH}=\mathrm{p}K_a+\log\frac{[\ce{CH3COO^-}]}{[\ce{CH3COOH}]}$ keeps pH almost constant — the buffer region — and at half-neutralisation pH = pKa. At equivalence only $\ce{CH3COO^-}$ remains, so the solution is basic.

History of the experiment

The pH scale was defined in 1909 by Søren Sørensen at the Carlsberg brewery laboratory; the glass electrode made pH-metric titration routine. The buffer plateau is the same chemistry that keeps blood near pH 7.4 with the

\ce{H2CO3/HCO3^-}

pair.

Chemists behind it

Svante Arrhenius Wilhelm Ostwald