Buffers

Why this matters

Your blood stays near pH 7.4 even after you eat, exercise or hold your breath. Cell cultures, enzyme assays and DNA work all happen in buffered solutions, and many clinical tests depend on buffered reagents. A buffer is the reason a small spill of acid into a reaction mix does not wreck it. Biotech, clinical and chemistry jobs all involve preparing and choosing buffers, so it pays to understand why they work.

The core idea

A buffer is a solution that resists changes in pH when a small amount of acid or base is added. It contains two partners in similar amounts:

• a weak acid, HA, which can neutralise added base, and
• its conjugate base, A⁻, which can neutralise added acid.

The classic example is acetic acid (CH₃COOH) with sodium acetate (CH₃COONa), which supplies the acetate ion CH₃COO⁻.

What happens when you add acid

Added H⁺ is grabbed by the conjugate base:

CH3COO−  +  H+   ->  CH3COOH

The strong acid is turned into a weak acid, which barely ionises, so [H⁺] hardly changes.

What happens when you add base

Added OH⁻ is neutralised by the weak acid:

CH3COOH  +  OH−  ->  CH3COO−  +  H2O

The strong base becomes water and more conjugate base, so again pH changes very little.

Pure water has no such partners. Add 1 mL of 1 M HCl to a litre of pure water and the pH drops from 7 to about 3. Add the same acid to a litre of buffer containing 0.10 M acetic acid and 0.10 M acetate and the pH moves by only about 0.01.

The Henderson–Hasselbalch equation

For a buffer, the pH depends on the acid's strength (its pKa) and the ratio of base to acid:

pH = pKa + log10( [A−] / [HA] )

pKa is −log10 of the acid dissociation constant Ka. A lower pKa means a stronger acid. Acetic acid has pKa ≈ 4.76.

Worked example 1: equal amounts

0.10 M acetic acid and 0.10 M sodium acetate:

pH = 4.76 + log10(0.10 / 0.10) = 4.76 + log10(1) = 4.76 + 0 = 4.76

When acid and conjugate base are equal, pH = pKa. This is where the buffer works best.

Worked example 2: more base than acid

0.10 M acetate and 0.050 M acetic acid:

pH = 4.76 + log10(0.10 / 0.050) = 4.76 + log10(2) ≈ 4.76 + 0.30 = 5.06

Worked example 3: a tenfold ratio

If [A⁻] / [HA] = 10, pH = pKa + 1 = 5.76. If the ratio is 1/10, pH = pKa − 1 = 3.76.

Choosing a buffer

Rule of thumb: a buffer works well within about ±1 pH unit of its pKa. Pick the acid whose pKa is closest to the pH you need.

Buffer capacity

Capacity is how much acid or base a buffer can absorb before its pH changes a lot. It grows with the total concentration of the pair: 0.5 M acetate buffer resists more than 0.05 M at the same ratio. Once nearly all of one partner is used up, the buffer is exhausted and the pH swings quickly.

The blood example

In blood, carbonic acid (from dissolved CO₂) and bicarbonate form a buffer:

CO2 + H2O  <=>  H2CO3  <=>  H+ + HCO3−

The lungs adjust CO₂ and the kidneys adjust bicarbonate, so the body can tune the ratio. Clinical labs measure blood gases and bicarbonate to see whether this system is coping; interpreting those results is part of clinical laboratory training.

Common mistakes

• Thinking a buffer keeps pH constant forever. It resists change only until its capacity runs out.
• Using a strong acid as the buffer acid. HCl and its conjugate base Cl⁻ do not buffer; the acid must be weak.
• Inverting the ratio. It is base over acid: [A⁻]/[HA]. More base means a pH above the pKa.
• Choosing a buffer far from its pKa. A phosphate buffer at pH 4 does very little.

Quick recap

• A buffer is a weak acid plus its conjugate base in similar amounts.
• The base partner absorbs added H⁺; the acid partner absorbs added OH⁻.
• pH = pKa + log10([A⁻]/[HA]); equal amounts give pH = pKa.
• Choose a pKa within 1 unit of the target pH; more concentrated means more capacity.

Try it

Open the Do tab and take the quiz: you will explain how a buffer reacts to added acid and base and estimate buffer pH.