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Buffer pH (Henderson-Hasselbalch) Calculator

Enter the pKa and the base and acid concentrations to find the buffer pH.

How it is calculated

The tool divides the base concentration by the acid concentration to get the ratio, then applies pH = pKa + log10([A⁻] ÷ [HA]). [A⁻] is the conjugate base (the salt form) and [HA] the weak acid, entered in the same units, for example mol/L, because only their ratio matters. The pKa is the value of the specific acid at your temperature. Results are shown to two decimals, along with the ratio as x : 1 and the pKa you entered. Both concentrations must be above zero; zero or negative values return no result. The equation is an approximation that works best when both concentrations are much larger than the hydrogen ion concentration, and it ignores activity effects.

Worked example

Take acetic acid with pKa 4.76, with 0.20 of conjugate base (acetate) and 0.10 of acid. The ratio is 0.20 ÷ 0.10 = 2.00 : 1. The log term is log10(2) = 0.301. So pH = 4.76 + 0.301 = 5.061, shown as 5.06. With equal amounts of base and acid the log term would be zero and the pH would equal the pKa, 4.76; doubling the base side shifts it up by about 0.3.

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Common mistakes

  • Swapping base and acid. The base goes on top, so the wrong order flips the sign of the log term and gives a pH that is as far below pKa as it should be above.
  • Entering the amounts of the original acid and strong base added, rather than the concentrations after they have reacted. The formula needs the resulting [A⁻] and [HA].
  • Using a pKa for the wrong proton of a polyprotic acid, such as phosphoric acid, which has several pKa values. Choose the one closest to your target pH.
  • Using very dilute buffers or ratios beyond about 10 : 1 or 1 : 10, where the approximation and the buffering capacity break down.
  • Mixing units between the two concentration fields, such as mmol/L for one and mol/L for the other, which distorts the ratio.

About this calculator

The Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA]), relates the pH of a buffer solution to the acid dissociation constant (pKa) of a weak acid and the ratio of its conjugate base concentration [A⁻] to its own concentration [HA]. When the two concentrations are equal, the log term vanishes and pH equals pKa exactly.

Because it links a measurable ratio of concentrations to pH without needing to solve the full equilibrium expression, this equation is the standard shortcut chemists and biochemists use to design and analyze buffer solutions — mixtures that resist pH change when small amounts of acid or base are added.

Buffers built this way are essential in biochemistry (maintaining physiological pH near 7.4 in blood, using the bicarbonate system), pharmaceutical formulation, and laboratory reagent preparation, where a stable pH is required for reactions or assays to behave predictably.

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Comments (1)

  • Editorial team

    You can run the equation backwards to plan a recipe: the base to acid ratio you need equals 10 raised to the power of (target pH minus pKa). Aiming for pH 5.50 with acetic acid at pKa 4.76 needs a ratio of about 5.5, and entering 0.55 and 0.10 here returns 5.50. Diluting both parts keeps the pH in this formula but lowers the buffering capacity. What pH are you trying to hold, and with which acid?

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