Upper secondary · 15 min

Building an acetate buffer

Mix acetic acid with sodium acetate and watch the pH settle near pKa = 4.76, as Henderson–Hasselbalch predicts.

Goal

Prepare a buffer of a target pH by choosing the [salt]/[acid] ratio, and verify it resists small additions of acid or base.

Apparatus and reagents

0.1 M acetic acid, 0.1 M sodium acetate, burette, flask, pH meter or universal indicator, deionised water.

Procedure

Place 25 mL of 0.1 M acetic acid in the flask and read the starting pH.
Add the acetate solution slowly; note pH every 2 mL and look for the flat region near the half-equivalence point.
At [salt] = [acid] (12.5 mL added here), confirm pH ≈ pKa = 4.76.
Continue to double the salt amount and see how slowly pH climbs through the buffer region.

What to observe

The pH barely moves while acid and conjugate base coexist — the flat middle of the curve is the buffer region.
The indicator changes colour around pH ≈ 4.8–5.5, bracketing the acetate pKa.

Explanation

A buffer is a conjugate pair in one pot: added H⁺ is consumed by CH₃COO⁻ (→ CH₃COOH), added OH⁻ is consumed by CH₃COOH (→ CH₃COO⁻ + H₂O). Taking logarithms of Ka gives Henderson–Hasselbalch: pH = pKa + lg([A⁻]/[HA]). Buffer capacity peaks at [A⁻] = [HA] and stays useful over roughly pKa ± 1. The simulation’s glassware shows the colour and volume behaviour; the lab’s pH curve is the one students should sketch.

History of the experiment

Arrhenius framed acids and bases as donors and acceptors of ions in 1884, and Henderson wrote the pH relation in 1908 before Hasselbalch put it in logarithmic form in 1917 to study blood pH. Buffer control of blood pH is still a life-critical example of the same chemistry.

Chemists behind it

Svante Arrhenius Wilhelm Ostwald