Enter the initial total concentration and volume so component moles can be updated before applying the buffer equation.
Choose a goal, enter the shown values, and select Calculate.
Acetate buffer. pK\(_a\)=4.76, target pH=4.50, \(C_t=0.10\ \mathrm{M}\), \(V=1.00\ \mathrm{L}\).
The Henderson–Hasselbalch equation connects the pH of a buffer with its acid dissociation constant and the ratio of conjugate base to acid: \[ \mathrm{pH} = \mathrm{p}K_a + \log_{10}\!\left(\frac{[A^-]}{[HA]}\right). \] It follows from the definition \(K_a=\tfrac{[H^+][A^-]}{[HA]}\) by solving for \([H^+]\) and taking negative logarithms. In practice, it lets you think in ratios: every 1.0 increase in \(\log_{10}([A^-]/[HA])\) raises pH by 1.0.
Capacity is the resistance to pH change upon adding acid or base. It is largest when \([A^-]\approx[HA]\) and when the overall buffer concentration is higher. Two practical levers therefore are: pick a system with pKa close to your target pH, and use a sufficient total concentration (balanced with solubility, ionic strength, and biological compatibility).
Rules of thumb: if pH = pKa ± 0.30, then the ratio is about 2:1 or 1:2; at ±1.0, it’s ~10:1 or 1:10. These mental anchors help during rough planning and sanity checks.
Bottom line: Henderson–Hasselbalch turns buffer planning into a clean ratio problem. Stay near pKa, keep total concentration appropriate, and validate with a pH meter for final adjustments.
An acid buffer uses weak acid HA and conjugate base A⁻:
Plain text: pH = pKa + log10([A−]/[HA])
A base buffer uses weak base B and conjugate acid BH⁺. The calculator finds pOH, then converts it with temperature-dependent pKw:
Plain text: pOH = pKb + log10([BH+]/[B]); pH = pKw − pOH
Preset values are approximate references, mostly at 25 °C. They stay editable because ionic strength, concentration, solvent, and temperature can shift them.
| System | Pair used | Reference value | Approx. range |
|---|---|---|---|
| Acetate | CH₃COOH / CH₃COO⁻ | pKa 4.76 at 25 °C | 3.76–5.76 |
| Phosphate | H₂PO₄⁻ / HPO₄²⁻ | pKa₂ 7.21 at 25 °C | 6.21–8.21 |
| Citrate | H₂Cit⁻ / HCit²⁻ | pKa₂ 4.76 at 25 °C | 3.76–5.76 |
| MES | MES-H / MES⁻ | pKa 6.10 at 25 °C | 5.10–7.10 |
| PIPES | PIPES-H / PIPES⁻ | pKa 6.76 at 25 °C | 5.76–7.76 |
| HEPES | HEPES-H / HEPES⁻ | pKa 7.48 at 25 °C | 6.48–8.48 |
| Tris | Tris-H⁺ / Tris | pKa 8.06 at 25 °C | 7.06–9.06 |
| Borate | B(OH)₃ / B(OH)₄⁻ | pKa 9.24 at 25 °C | 8.24–10.24 |
| Ammonia/ammonium | NH₃ / NH₄⁺ | pKb 4.751 at 25 °C | pH ≈ 8.25–10.25 |
| Bicarbonate | H₂CO₃ / HCO₃⁻ | pKa₁ 6.35 at 25 °C | 5.35–7.35 |
This tool applies the concentration form of Henderson–Hasselbalch. In base mode, pKw is estimated for pure liquid water at ambient pressure from 0–100 °C with a compact temperature correlation; high-accuracy or non-ambient work should use the full IAPWS formulation. Strong-acid/base stress uses complete 1:1 stoichiometric conversion before reapplying the buffer equation.
No reviewed-by attribution is shown because this page does not document an identifiable qualified reviewer and review process.
For an acid buffer, use pH = pKa + log10([A−]/[HA]). For a base buffer, first calculate pOH = pKb + log10([BH+]/[B]), then use pH = pKw − pOH at the selected temperature.
Rearrange the same equation: pKa = pH − log10([A−]/[HA]), or [A−]/[HA] = 10^(pH − pKa). The base-buffer forms use pOH, pKb, and [BH+]/[B].
The base-10 logarithm of 1 is zero, so equal conjugate-base and acid concentrations make the Henderson–Hasselbalch equation reduce to pH = pKa.
An acid buffer contains a weak acid and its conjugate base. A base buffer contains a weak base and its conjugate acid; its equation is commonly written for pOH and converted to pH with pKw.
Ideal dilution leaves the component ratio unchanged, so the calculated pH is unchanged, but buffer capacity falls. Real solutions can shift because activities and ionic strength change.
Temperature can change pKa, pKb, pKw, and activity coefficients. Use dissociation data measured near the working temperature and verify the prepared solution at that temperature.
Choose the conjugate pair whose pKa brackets the target pH and treat that dissociation step as the buffer pair. Other equilibria may matter near overlapping pKa values.
Henderson–Hasselbalch uses concentrations as activity approximations. Temperature, ionic strength, reagent purity, hydrate form, carbon dioxide uptake, calibration, and volume adjustment can all change the measured pH.
At pH = pKa, \([A^-]=[HA]\) and buffer capacity peaks—exactly the halfway point of a weak-acid titration curve.
Every 1.0 pH unit away from pKa is a 10× ratio shift; 0.30 pH is about a 2× shift—handy for mental checks.
Your blood stays near pH 7.4 because bicarbonate buffers at a ~20:1 base:acid ratio (pKa ≈ 6.1) aided by CO₂ breathing.
Keeping the ratio fixed means pH barely moves when you dilute, but the capacity halves if you halve total concentration—less moles to soak up acid/base.
Many pKa values shift by 0.01–0.03 per °C; a 10 °C swing can nudge pH by ~0.1–0.3 at constant ratio.
