The Henderson–Hasselbalch equation relates the pH of a buffer solution to the acid’s dissociation constant and the ratio of conjugate base to acid. Core form
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For an acid HA with conjugate base A−, the equation is:
pH = pKa + log10([A−]/[HA]) -
Equivalently, using concentrations: pH = pKa + log10(C(A−)/C(HA))
Key interpretations
- When [A−] equals [HA], pH equals pKa.
- The buffer capacity is highest when the pH is near pKa, and the ratio [A−]/[HA] is near 1.
- Changes in total buffer concentration (the absolute amounts of HA and A−) primarily affect buffering capacity, not the pH, as long as the ratio remains constant.
- For basic buffers (conjugate base B− and acid BH), the equation can be written as pOH = pKb + log10([BH]/[B−]), or pH = 14 − (pKb + log10([BH]/[B−])) when using water at 25°C.
Common applications
- Designing buffer solutions at a desired pH by selecting an acid with an appropriate pKa and adjusting the HA/A− ratio.
- Interpreting titration curves, buffering regions, and the response of pH to small added amounts of strong acid or base.
Limitations
- Assumes the buffer components are the dominant species and that the system behaves ideally (activity approximated by concentrations).
- Less accurate when either the weak acid or its conjugate base is in very low or very high concentration, or at extreme pH values where water autoionization and other equilibria become relevant.
If you have a specific buffer system (e.g., acetic acid/acetate, or a phosphate buffer) and want to calculate the pH, provide the acid’s pKa and the desired or measured concentrations of HA and A−, and the pH can be computed directly using the Henderson–Hasselbalch equation.
