Henderson–Hasselbalch Calculator
Buffer pH from the pKa and the base-to-acid ratio, the ratio needed for a target pH, or how much weak acid and conjugate base make a buffer of a given volume and concentration.
Result…
How it's calculated
pH = pKa + log₁₀([A⁻] ÷ [HA]) Ratio for a target pH: [A⁻] ÷ [HA] = 10^(pH − pKa) Share as base = ratio ÷ (1 + ratio) [A⁻] = total × share; [HA] = total − [A⁻]; moles = concentration × volume
The Henderson–Hasselbalch equation rearranges the acid dissociation constant of a weak acid, Ka = [H⁺][A⁻] ÷ [HA], into a form for buffers: the pH equals the pKa plus the base-10 log of the base-to-acid ratio. When the two are equal the pH is the pKa; each tenfold change in the ratio moves the pH by one unit. It works best when the pH is within about one unit of the pKa and the concentrations are much larger than Ka (at least 100 times, per the LibreTexts general chemistry text), because it assumes the acid barely dissociates on its own. It uses concentrations rather than activities, so it is a starting point for making a buffer, not the final pH.
Your numbers, step by step
About the Henderson–Hasselbalch calculator
This Henderson–Hasselbalch calculator does the three buffer jobs that come up in the lab and in exams. It finds the pH of a buffer from the pKa and the concentrations of the weak acid and its conjugate base. It finds the base-to-acid ratio you need for a target pH. And it plans a buffer: give the volume, the total concentration and the target pH, and it splits that into moles of acid and conjugate base, and grams too if you add both molar masses. It also warns when the pH is more than one unit from the pKa or when a concentration is too low for the equation to hold. Each step, from the ratio to the logarithm to the final amounts, is shown with your numbers, so you can check it against your own working.
How to use it
- Under “Find”, choose “pH of a buffer”, “Ratio for a pH” or “Make a buffer”.
- Enter the “pKa of the weak acid” from a data table. For an ammonia-type buffer use the pKa of the conjugate acid.
- For a pH, enter the “Concentration of the conjugate base [A⁻]” and the “Concentration of the weak acid [HA]”. Otherwise enter the “Target pH”.
- To make a buffer, enter the “Volume of buffer” with its unit and the “Total buffer concentration”, and choose “From molar masses” under “Masses too” to get grams.
- Read the result and the note, which flags pH values far from the pKa.
Worked example
To make 500 mL of a 0.1 mol/L buffer at pH 5.00 from an acid with pKa 4.76: the ratio [A⁻] ÷ [HA] = 10^(5.00 − 4.76) = 10^0.24 = 1.738, so 1.738 ÷ 2.738 = 63.47% is in the base form. That is 0.06347 mol/L of base and 0.03653 mol/L of acid; in 0.5 L, 0.03174 mol of base and 0.01826 mol of acid. With molar masses of 82.034 g/mol for the salt and 60.052 g/mol for the acid, weigh out 2.604 g of the salt and 1.097 g of the acid.
Updated
Questions people ask
What is the Henderson–Hasselbalch equation?
pH = pKa + log₁₀([A⁻] ÷ [HA]), where [A⁻] is the conjugate base and [HA] the weak acid. When the two are equal, the pH equals the pKa.
How do I find the ratio of base to acid for a pH?
Raise 10 to the power of (pH − pKa). One unit above the pKa needs ten times as much base as acid; one unit below needs ten times as much acid.
When is the Henderson–Hasselbalch equation not accurate?
When the pH is more than about one unit from the pKa, and when concentrations are not much larger than Ka; the LibreTexts general chemistry text puts the limit at concentrations at least 100 times Ka. It also ignores activity, so concentrated or salty buffers differ.
Can I use moles instead of concentrations?
Yes, for the pH. Both species are in the same volume, so the ratio of moles equals the ratio of concentrations.
Why is my measured buffer pH different from the calculated one?
The equation uses concentrations, not activities, and pKa shifts with temperature and ionic strength. Make the buffer from the calculated amounts, then adjust with a calibrated pH meter. For a strong acid or base on its own, use the pH calculator instead.