Find a buffer solution’s pH from its acid dissociation constant and the ratio of conjugate base to weak acid concentration.
How it works
The Henderson-Hasselbalch equation is pH = pKa + log([A⁻] ÷ [HA]). When the base and acid concentrations are equal, pH equals pKa exactly — a buffer’s most effective point.
What this does not include
This calculates pH for a buffered solution within its effective range. It doesn’t model a solution far outside that range, where the buffer’s capacity to resist pH change breaks down.
How to use this calculator
- Enter the pKa of the weak acid.
- Enter the conjugate base and weak acid concentrations.
A worked example
A buffer with pKa 4.76, equal base and acid concentrations (0.1 each): pH = pKa + log(base/acid) = 4.76 + log(1) = 4.76 — when concentrations are equal, pH exactly equals pKa.
Same pKa, base concentration 1, acid concentration 0.1: pH = 5.76 — more base relative to acid raises the pH.
What the variables mean
| Variable | Meaning |
|---|---|
| pKa | The acid’s dissociation constant, in pKa form |
| Base concentration | Concentration of the conjugate base |
| Acid concentration | Concentration of the weak acid |
Edge cases worth knowing
When base and acid concentrations are equal, pH always equals pKa exactly — the first example above shows this directly, since log(1) = 0.
Zero acid concentration makes the ratio undefined — dividing by zero has no meaning, so the calculator declines to show a result for that case.
Why does pH equal pKa when the concentrations are equal?
The logarithm of a 1:1 ratio is zero, so the equation reduces to pH = pKa exactly — this is also the point where a buffer resists pH change most effectively.
What is a buffer solution used for?
A buffer resists changes in pH when small amounts of acid or base are added, which is essential in biological systems, laboratory work, and industrial processes that need a stable pH.
Why is the acid called “weak” specifically?
A strong acid dissociates completely, leaving no equilibrium between HA and A⁻ to buffer against — the Henderson-Hasselbalch equation specifically models a weak acid’s partial, reversible dissociation.
