pH & Buffers: Why Some Solutions Resist Changing Acidity

pH is just -log of a concentration

pH measures how acidic or basic a solution is by tracking the concentration of hydrogen ions it contains, on a logarithmic scale so that the enormous range found in real chemistry — from battery acid to drain cleaner — fits into a number line roughly 0 to 14.

pH = -log10( [H+] )        pure water at 25 C: [H+] = 1e-7 M, pH = 7

Weak acids and the equilibrium that resists change

A strong acid like HCl dissociates completely, so its pH is set directly by how much you added. A weak acid HA, like acetic acid, only partially dissociates, settling into an equilibrium with its conjugate base A⁻:

HA  <-->  H+  +  A-        Ka = [H+][A-] / [HA]        pKa = -log10(Ka)

Ka (and its more convenient log form, pKa) describes how far that equilibrium sits toward dissociation: a small pKa means a stronger acid that dissociates more readily, a large pKa a weaker one. Every weak acid has its own fixed pKa, a physical constant of the molecule, independent of concentration.

The Henderson-Hasselbalch equation

Rearranging the equilibrium expression and taking logarithms gives the single equation that governs every buffer:

pH = pKa + log10( [A-] / [HA] )

When the acid and its conjugate base are present in equal amounts, [A⁻] = [HA], the log term is log10(1) = 0, and pH = pKa exactly — the half-equivalence point of a titration. Moving away from that 1:1 ratio shifts the pH, but only logarithmically: a tenfold change in the ratio shifts pH by exactly one unit. That slow, logarithmic response to changing ratio is precisely what makes a mixture of a weak acid and its conjugate base resistant to pH swings — it is a buffer.

Why buffers resist added acid or base

Add a small amount of strong acid to an HA/A⁻ buffer and the extra H+ reacts with the reservoir of A⁻ to make more HA, consuming most of the perturbation instead of letting free [H+] spike. Add a strong base and it is neutralised by HA turning into A⁻. Either way the ratio [A⁻]/[HA] moves only a little, so by Henderson-Hasselbalch the pH moves only a little too. This buffering effect is strongest when [A⁻] and [HA] are comparable — buffer capacity is maximal exactly at pH = pKa, and is usually considered practically useful within about one pH unit either side of pKa, after which one of the two reservoirs runs out and the pH starts changing rapidly again.

The titration curve: buffer region and equivalence point

Plotting pH against the volume of strong base added to a weak acid produces a characteristic S-shaped curve. Near the start, pH rises quickly as the first drops of base react. It then flattens into the buffer region centred on pH = pKa, where large additions of base change the pH only slowly. As the acid is nearly all consumed, the curve turns upward again and rises steeply through the equivalence point — where moles of base added equal the original moles of acid — before flattening out at high pH as the solution becomes dominated by excess strong base. An indicator is chosen so that its own colour-change pH range brackets the equivalence point as tightly as possible, which is what lets a titration be read visually rather than only with a pH meter.

Frequently asked questions

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