Buffer pH (Henderson–Hasselbalch) Calculator
The Henderson–Hasselbalch equation describes how the pH of a buffer made of a weak acid HA and its conjugate base A⁻ depends on the pKa and the [A⁻] / [HA] ratio. Enter pKa (or pick a common laboratory buffer preset) and the two concentrations to get the buffer pH, the [A⁻] : [HA] ratio, the effective buffering window (pKa ± 1) and the fraction of each species — or switch to reverse mode to find the ratio needed to hit a target pH.
Enter valid values: a finite pKa, positive acid/base concentrations, or a finite target pH.
Buffer pH
4.76
Inside effective buffering range
—
[A⁻] : [HA] ratio
1 : 1
—
Species fractions
—
Formula
pH = pKa + log₁₀([A⁻] ÷ [HA])
Assumes an ideal dilute solution where activities ≈ concentrations and ionic strength is moderate. At high ionic strength, apply a Davies / Debye–Hückel correction.
Formula
pH = pKa + log₁₀([A⁻] ÷ [HA]) Reverse: [A⁻] ÷ [HA] = 10^(pH − pKa) At [A⁻] = [HA], pH = pKa — the point of maximum buffering capacity.
- · Only the [A⁻] : [HA] ratio enters the equation — units cancel, so mol/L, mmol/L or any consistent measure of "amount" works as long as both sides match.
- · The textbook "useful" buffering window is pKa ± 1, equivalent to a ratio between 1:10 and 10:1 (A⁻ fraction between ~9 % and ~91 %; Skoog, "Fundamentals of Analytical Chemistry" 9e §9-5).
- · Preset pKa values are taken from the CRC Handbook of Chemistry & Physics, 97th ed. §5 (aqueous, 25 °C).
- · Assumes an ideal dilute solution where activities ≈ concentrations and the weak acid is the dominant proton donor. At high ionic strength (> 0.1 M) or very dilute regimes, apply a Davies or Debye–Hückel correction.
- · Carbonic acid has two pKa values (≈6.35 and ≈10.33) and phosphoric acid three (≈2.15 / 7.20 / 12.35); picking the wrong one shifts the predicted pH by several units.
- · Calculation happens entirely in your browser — values are never uploaded, and the URL stays in sync so you can share a link that reproduces the same result.
Frequently asked
Does it matter if I use mol/L versus mmol/L?
No. Only the [A⁻] / [HA] ratio enters the equation, and units cancel between numerator and denominator. So 0.1 M and 0.1 M, 100 mM and 100 mM, or 5 g and 5 g of the same species all give the same pH — as long as both sides share the unit. In practice you still need to pick a sensible total concentration: too dilute (< 1 mM) and the buffer has little capacity to resist added acid/base; too concentrated and it may interfere with other reactions or precipitate.
When does the Henderson–Hasselbalch equation break down?
In three main regimes: (1) strong acids/bases — the equation assumes a partially-dissociated weak acid/base; (2) very dilute buffers (< 1 mM, especially near pKa 7), where water's autoionisation contributes to pH and you need the full proton-balance equation; (3) high ionic strength systems (physiological saline, plasma, mixed solvents), where you must swap concentrations for activities using a Davies or Debye–Hückel correction. Errors grow once |pH − pKa| > 2, but for routine lab buffer design it remains the fastest, most intuitive approximation.
Why is the "effective buffering range" pKa ± 1?
Buffer capacity β peaks at [A⁻] = [HA] (i.e. pH = pKa) and falls off geometrically as |pH − pKa| grows. By |pH − pKa| = 1 the ratio is already 10 : 1 or 1 : 10, meaning one species is below 10 % — adding a little more acid or base depletes the minor partner and pH swings sharply. Textbooks (and analytical-chemistry references like Skoog 9e) therefore quote pKa ± 1 (A⁻ between ~9 % and ~91 %) as the useful buffering window; outside it, switch to a different buffer whose pKa is closer to your target pH.
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