Buffer pH Calculator
This calculator works out the pH of a buffer solution using the Henderson-Hasselbalch equation, pH = pKa + log([A-]/[HA]), the standard method for predicting how a mixture of a weak acid and its conjugate base will behave. Enter the pKa of your weak acid manually, or pick from a list of common buffer systems including acetic acid/acetate, phosphate, Tris, HEPES, MES, ammonia/ammonium and the three citric acid pKa values, then enter the molar concentrations of the weak acid [HA] and the conjugate base [A-]. The calculator instantly returns the buffer pH, the pOH, the [A-]/[HA] ratio, and the effective buffering range (pKa plus or minus one pH unit), alongside a calculation breakdown showing the ratio, log value and final pH, plus buffer properties including hydrogen and hydroxide ion concentration and whether your buffer sits inside its effective range. A pH scale bar and a plain-English verdict show how strong your buffer is and whether the acid or base side holds more capacity to neutralise additions. Use it to check a buffer you have already made, or to work backwards and choose a suitable pKa and ratio for a target pH. Because the equation assumes ideal behaviour and becomes less accurate at very low concentrations or extreme pH, treat the result as a theoretical estimate and confirm with a calibrated pH meter for precision laboratory work.
1. Buffer System
2. Concentrations
Calculation Breakdown
Buffer Properties
Worked Example (matches defaults)
Acetate buffer: acetic acid / sodium acetate, pKa = 4.76, [HA] = 0.1 mol/L, [A-] = 0.1 mol/L
Ratio = [A-] / [HA] = 0.1 / 0.1 = 1.000
log(1.000) = 0.000
pH = 4.76 + 0.000 = 4.76
When acid and conjugate base concentrations are equal, the buffer pH equals the pKa exactly.
The Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation expresses the pH of a buffer solution in terms of the pKa of the weak acid and the ratio of the concentrations of conjugate base to weak acid:
pH = pKa + log([A-] / [HA])
Where:
- pKa is the negative base-10 logarithm of the acid dissociation constant (Ka)
- [A-] is the molar concentration of the conjugate base (the deprotonated form)
- [HA] is the molar concentration of the weak acid (the protonated form)
The equation is derived from the equilibrium expression for the dissociation of the weak acid (HA) into H+ and A-. It is an excellent approximation when the concentrations of acid and base are both substantially greater than the H+ concentration (typically above about 1 mmol/L) and the ratio [A-]/[HA] is between 0.1 and 10.
How Buffer Solutions Work
A buffer solution resists changes in pH when small amounts of acid or base are added. It does this through two reactions:
- Added acid: H+ is consumed by the conjugate base: A- + H+ → HA
- Added base: OH- is consumed by the weak acid: HA + OH- → A- + H2O
The buffer is most effective at resisting pH change when both [A-] and [HA] are present in substantial amounts. A buffer is essentially exhausted when one component is fully consumed.
Effective Buffering Range
The Henderson-Hasselbalch equation shows that the buffer pH equals pKa when [A-] = [HA] (since log(1) = 0). As the ratio moves away from 1, the pH moves away from pKa. At a ratio of 10:1 (base:acid), pH = pKa + 1. At a ratio of 1:10 (base:acid), pH = pKa - 1. Outside the range pKa ± 1, the buffer loses much of its capacity to resist pH change, and a different buffer with a more appropriate pKa should be selected.
Common Buffer Systems
| Buffer System | pKa | Useful pH Range | Common Use |
|---|---|---|---|
| Acetic acid / Acetate | 4.76 | 3.8 to 5.8 | Food chemistry, biochemistry |
| Carbonic acid / Bicarbonate | 6.35 | 5.4 to 7.4 | Blood physiology, geology |
| Phosphate H2PO4- / HPO42- | 7.20 | 6.2 to 8.2 | Biochemistry, cell biology |
| Tris | 8.06 | 7.1 to 9.1 | Molecular biology, gel electrophoresis |
| Ammonia / Ammonium | 9.25 | 8.3 to 10.3 | Analytical chemistry |
| HEPES | 7.48 | 6.8 to 8.2 | Cell culture, protein biochemistry |
| MES | 6.10 | 5.5 to 6.7 | Plant biology, protein chromatography |
Preparing a Buffer at a Target pH
To prepare a buffer at a specific target pH, rearrange the Henderson-Hasselbalch equation to find the required ratio of conjugate base to weak acid:
[A-] / [HA] = 10^(pH - pKa)
For example, to prepare a phosphate buffer at pH 7.40 using the H2PO4-/HPO42- pair (pKa 7.20):
- Ratio = 10^(7.40 - 7.20) = 10^0.20 = 1.585
- So you need approximately 1.585 parts of HPO42- (dibasic) to every 1 part of H2PO4- (monobasic)
Limitations of the Henderson-Hasselbalch Equation
The equation assumes ideal solution behaviour and that the autoionisation of water is negligible compared to the concentrations of the buffer components. It becomes less accurate at very low concentrations (below about 1 mmol/L), at extreme pH values (below 3 or above 11), or when the ratio [A-]/[HA] is outside the range 0.01 to 100. For high-precision work, activity coefficients and the contribution of water autoionisation should be accounted for.
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Sources and method: Henderson-Hasselbalch equation as derived from the weak acid equilibrium expression (Lawrence J. Henderson, 1908; Karl A. Hasselbalch, 1917). Atkins' Physical Chemistry (Oxford University Press). Chang, R., Chemistry (McGraw-Hill). pKa values from the NIST Chemistry WebBook and standard biochemistry references.
This calculator provides theoretical pH values based on the Henderson-Hasselbalch approximation. It assumes ideal behaviour and that concentrations are well above the ionisation of water. For high-precision laboratory work, verify against measured pH using a calibrated pH meter.