Buffer pH Calculator
Enter a weak acid's pKa, conjugate base concentration [A⁻] and weak acid concentration [HA]; using the Henderson–Hasselbalch equation the tool instantly computes the buffer pH and the [A⁻]/[HA] ratio — ideal for preparing buffers.
Input Data
Results
At a glance:The Henderson–Hasselbalch equation describes the relationship between a buffer's pH and its weak-acid/conjugate-base concentration ratio: pH = pKa + log₁₀([A⁻]/[HA]). Here pKa is the negative logarithm of the weak acid dissociation constant, [A⁻] is the conjugate base concentration and [HA] the weak acid concentration. When [A⁻] = [HA], the log term is 0 and pH equals pKa, where the buffer capacity is strongest. This equation is the core tool for preparing biochemistry and molecular-biology buffers (phosphate, acetate, Tris, etc.).
Formula
Buffer pH: pH = pKa + log₁₀([A⁻] / [HA]).
[A⁻] = [HA] makes log term 0, pH = pKa (strongest buffer capacity).
[A⁻] > [HA] gives pH > pKa; [A⁻] < [HA] gives pH < pKa.
$$pH = pK_a + \log_{10}\!\left(\dfrac{[A^-]}{[HA]}\right)$$$$[A^-] = [HA] \Rightarrow pH = pK_a$$How to Use
- Enter the weak acid pKa (e.g. acetic acid 4.76, dihydrogen phosphate 7.21, Tris 8.06).
- Enter the conjugate base concentration [A⁻] and weak acid concentration [HA] (same unit).
- The right panel instantly shows the buffer pH and the [A⁻]/[HA] ratio.
pKa of common buffer systems (25°C)
| Buffer System | Effective pH Range | pKa |
|---|---|---|
| Acetate / sodium acetate | 3.8–5.8 | 4.76 |
| Phosphate (H₂PO₄⁻/HPO₄²⁻) | 6.2–8.2 | 7.21 |
| Tris | 7.0–9.0 | 8.06 |
| Bicarbonate (blood) | 6.1 (physiological) | 6.1 |
A buffer works well roughly within pKa ± 1. Choose a system whose pKa is closest to the target pH for best capacity.
Case Studies
Preparing pH 4.76 acetate buffer
Use acetic acid (pKa 4.76); sodium acetate [A⁻] = 0.1 M, acetic acid [HA] = 0.1 M.
pH = 4.76 + log₁₀(0.1 / 0.1) = 4.76 + 0 = 4.76.
Here [A⁻]/[HA] = 1 and buffer capacity is strongest, resisting pH change from added strong acid or base.
Raising to pH 5.06
Same acetic system, double the conjugate base: [A⁻] = 0.2 M, [HA] = 0.1 M.
pH = 4.76 + log₁₀(0.2 / 0.1) = 4.76 + 0.301 ≈ 5.06.
Raising the conjugate-base ratio makes pH above pKa; to lower pH, add more weak acid.
FAQ
Why is buffer capacity strongest when [A⁻] = [HA]?
When conjugate base and weak acid concentrations are equal, the ratio is 1 and the log term is 0, so pH = pKa. The system then has plenty of both A⁻ (to neutralise added acid) and HA (to neutralise added base), so it most resists pH change from small amounts of strong acid or base.
How to choose a buffer system?
Pick a weak acid whose pKa is closest to the target pH. The effective range is about pKa ± 1; outside it capacity drops sharply. For example, to make a pH 7.4 physiological buffer, phosphate (pKa 7.21) is far better than acetate (pKa 4.76).
Must the concentration unit be M?
Not necessarily. The equation only uses the ratio [A⁻]/[HA], so as long as both share a unit (M, mM, or moles), the ratio and hence pH are correct. But total concentration affects buffer capacity — higher concentration resists pH change better.
What are the limits of this equation?
Henderson–Hasselbalch assumes the acid/base dissociation does not much alter the total concentration; it is most accurate at moderate concentrations and pH within pKa ± 1. For very dilute solutions, strong acids/bases, or work needing activity coefficients, use full equilibrium calculation.
Why is blood pH maintained at 7.4?
Blood mainly relies on the bicarbonate buffer (H₂CO₃/HCO₃⁻, pKa ≈ 6.1). Physiologically [HCO₃⁻]/[H₂CO₃] ≈ 20:1, giving pH = 6.1 + log₁₀(20) ≈ 7.4. Breathing (adjusting CO₂) and kidneys (adjusting HCO₃⁻) together maintain this ratio, keeping blood pH at 7.35–7.45.
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References
Content review: Calculatorism Science Team. Results are for reference only; please refer to the relevant authorities for the official figures.