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Henderson–Hasselbalch Calculator

Find the pH of a buffer with the Henderson–Hasselbalch equation.

How it is calculated

The tool applies pH = pKa + log10([A⁻] / [HA]). pKa is the acid's pKa, [A⁻] the concentration of the conjugate base (the salt form) and [HA] the concentration of the weak acid. Both concentrations must use the same unit, for example mol/L, because only their ratio matters. The tool also reports that ratio as a number followed by ': 1'.

Both concentrations must be greater than zero; a zero gives no result, as the logarithm of zero or a negative number does not exist. When the base and the acid are equal, the log is 0 and the pH equals the pKa. Each tenfold change in the ratio moves the pH by 1 unit. The equation is an approximation that works best for buffers of moderate concentration. It ignores activity effects and the dissociation of water, so use it for weak acid buffers rather than strong acids.

Worked example

An acetic acid and acetate buffer with a pKa of 4.76, containing 0.2 mol/L of acetate and 0.1 mol/L of acetic acid. The ratio is 0.2 / 0.1 = 2.00, so the ratio is shown as 2.00 : 1. The logarithm of 2 is 0.301. Add it to the pKa: 4.76 + 0.301 = 5.061, shown as a pH of 5.06. Making the two concentrations equal would bring the pH back to 4.76.

Try this example →

Common mistakes

  • Putting the acid in the base box. The ratio is base over acid, so swapping them flips the sign of the log and gives a pH on the wrong side of the pKa.
  • Entering the total buffer concentration in one box. The equation needs the base and the acid separately.
  • Mixing units for the two concentrations, such as mM for one and M for the other. That distorts the ratio.
  • Using the pKb for a basic buffer. The tool needs the pKa of the conjugate acid, so for a base buffer use the pKa of its conjugate acid.
  • Trusting the result for very dilute or very concentrated solutions, or when the ratio is far from 1. Buffers work best within about one pH unit of the pKa.

About this calculator

This calculator finds a buffer's pH using the Henderson–Hasselbalch equation: pH = pKa + log₁₀([A⁻]/[HA]), where pKa is the acid dissociation constant of the weak acid, [A⁻] is the concentration of its conjugate base, and [HA] is the concentration of the undissociated acid. It's derived directly from the acid dissociation equilibrium constant Ka.

Chemists use it to prepare buffer solutions at a target pH for lab work and titrations, and it's foundational in biochemistry and physiology for understanding blood buffering systems like the bicarbonate buffer, which keeps arterial blood pH tightly regulated near 7.4. Enter the pKa and the concentrations of the conjugate base and acid to calculate the resulting pH.

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