Buffer pH Calculator (Henderson-Hasselbalch)
Written by Thierno Sadou Diallo, formula verified per our methodology • Last checked on 9/8/2026
The Henderson-Hasselbalch equation gives pH = pKa + log₁₀([conjugate base] ÷ [acid]). When the concentrations of the weak acid and conjugate base are equal, the buffer solution's pH is exactly equal to the acid's pKa.
Explanation
A buffer solution resists pH changes when a small amount of acid or base is added to it, thanks to the simultaneous presence of a weak acid and its conjugate base in solution. The Henderson-Hasselbalch equation relates this solution's pH to the weak acid's pKa (a constant characterizing its strength) and the ratio of concentrations between the conjugate base and the acid. This equation reveals a remarkable and widely used property in the lab: when the acid and conjugate base concentrations are exactly equal, the ratio equals 1, whose logarithm is zero — the solution's pH is then exactly equal to the acid's pKa, a practical reference point for preparing a buffer at a precise target pH. Departing from this 1:1 ratio shifts the pH on either side of the pKa: more conjugate base raises the pH (more basic solution), more acid lowers it (more acidic solution). This equation is distinct from our pH calculator, which calculates the pH of a strong acid alone from its concentration, with no conjugate base or buffering effect: the two calculations answer different chemistry questions, even though they share the same pH concept. Buffer solutions are essential in biology and medicine (human blood itself is kept at a stable pH by natural buffer systems) and in the lab, where they keep the pH constant during an experiment despite the gradual addition of a reagent. The exact concentration of each buffer component is generally prepared from a stock solution, like the one calculated by our stock solution molarity calculator.
Example: acetate buffer with twice as much conjugate base as acid
Inputs
pKa (acetic acid): 4.76. Conjugate base concentration: 0.2 mol/L. Acid concentration: 0.1 mol/L.
Calculation
pH = 4.76 + log₁₀(0.2 ÷ 0.1) = 4.76 + log₁₀(2) ≈ 4.76 + 0.301 = 5.061.
Result
This buffer solution has a pH of about 5.061, more basic than acetic acid's pKa since it contains more conjugate base than acid.
Frequently asked questions
Why is pH exactly equal to pKa when concentrations are equal?
Because the logarithm of a ratio of 1 (two identical concentrations) is exactly 0, which entirely cancels the second term of the equation. Only pH = pKa remains, a very useful property in practice: to prepare a buffer at a precise target pH, an acid whose pKa is close to the desired pH is generally chosen, then the concentration ratio is finely adjusted around 1:1.
Over what pH range does a buffer solution remain effective?
A buffer generally remains effective within about ±1 pH unit around the pKa of the acid used (the concentration ratio then staying between 1:10 and 10:1). Beyond this range, one of the two components (acid or conjugate base) becomes too much of a minority to keep effectively absorbing an addition of acid or base without significantly shifting the pH.
Why does human blood need buffer systems?
Blood pH must stay within a very narrow range (about 7.35 to 7.45) for proteins and biological processes to function normally; a large deviation can have serious consequences for the body. The body uses several natural buffer systems, notably the carbonic acid/bicarbonate pair, to absorb the acidity variations constantly produced by metabolism (cellular respiration, digestion) and keep this pH stable.