Buffer pH Calculator
The Henderson-Hasselbalch equation predicts the pH of a buffer solution from the pKa of the weak acid and the ratio of conjugate base to weak acid concentrations. Buffer solutions resist pH changes upon addition of small amounts of acid or base, making this calculation fundamental in biochemistry, physiology, and analytical chemistry.
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Formula
pH = pKa + log₁₀([A⁻] / [HA])
pH is the buffer's equilibrium pH. pKa is the negative log of the acid dissociation constant of the weak acid. [A⁻] is the molar concentration of the conjugate base (the deprotonated form). [HA] is the molar concentration of the weak acid (the protonated form). When [A⁻] = [HA], the log term equals zero and pH = pKa — this is the optimal buffering point. The equation is most accurate when both concentrations are between 0.1× and 10× the ratio (i.e., within one pH unit of pKa).
How to use the Buffer pH Calculator
- 1
Enter your pka of weak acid
e.g. acetic acid pKa = 4.74, carbonic acid pKa₁ = 6.35
- 2
Enter your conjugate base concentration [a⁻]
Value should be in M.
- 3
Enter your weak acid concentration [ha]
Value should be in M.
- 4
Read your results instantly
Results update in real time as you type.
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The Henderson-Hasselbalch equation and buffer design
The Henderson-Hasselbalch equation is derived directly from the equilibrium expression for a weak acid dissociation. Because it expresses pH in terms of a ratio of concentrations rather than absolute concentrations, buffers maintain their pH even when diluted — as long as the ratio [A⁻]/[HA] is preserved. To design a buffer at a target pH, choose a weak acid with a pKa within one unit of the target pH, then adjust the [A⁻]/[HA] ratio. For example, to prepare an acetate buffer at pH 5.0 (pKa of acetic acid = 4.74), you need log₁₀([A⁻]/[HA]) = 0.26, so [A⁻]/[HA] ≈ 1.82. Mix sodium acetate and acetic acid in roughly a 1.82:1 molar ratio.
Biological buffers and the carbonate system
The most important biological buffer is the bicarbonate buffer system in blood: CO₂(aq) + H₂O ⇌ HCO₃⁻ + H⁺, with an effective pKa of about 6.1. Blood plasma pH of 7.4 corresponds to a bicarbonate-to-carbonic acid ratio of about 20:1. This buffer is remarkably effective not because of its chemical capacity alone but because the lungs can regulate CO₂ (and thus carbonic acid concentration) by adjusting breathing rate. Common laboratory buffers include phosphate buffered saline (PBS) at pH 7.4, HEPES at pH 7.0–7.5, citrate at pH 3–6, and Tris at pH 7–9. Choosing the wrong buffer can inhibit enzyme activity or precipitate metal cofactors.
Tips & Insights
Buffer works best within ±1 pH unit of pKa
The Henderson-Hasselbalch equation is most accurate and buffer capacity is highest when the [A⁻]/[HA] ratio is between 0.1 and 10, corresponding to pH = pKa ± 1. Outside this range, the system loses buffering capacity rapidly.
Equal concentrations give pH = pKa
When [A⁻] = [HA], log₁₀(1) = 0, so pH = pKa. This is the half-equivalence point in a titration and the pH of maximum buffering capacity.
Temperature shifts pKa and therefore buffer pH
pKa values are temperature-dependent. Tris buffer's pKa shifts by about −0.031 pH units per degree Celsius increase, so a Tris buffer prepared at 4°C will have a different pH at 37°C. Always equilibrate and measure buffer pH at the intended use temperature.
Worked Examples
Acetate buffer at equimolar concentration
Buffer pH = 4.74 — equal concentrations of acetate and acetic acid give pH exactly equal to pKa.
Phosphate buffer shifted toward base
Buffer pH = 7.80 — a 4:1 base-to-acid ratio shifts the pH by log₁₀(4) ≈ 0.60 units above pKa.
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Frequently Asked Questions
What is a buffer solution?
A buffer solution resists changes in pH when small amounts of acid or base are added. It is made from a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations.
What is pKa?
pKa = −log₁₀(Ka) where Ka is the acid dissociation constant. A lower pKa indicates a stronger weak acid. pKa determines the pH range over which an acid-base pair can effectively buffer.
Can I use Henderson-Hasselbalch for very dilute solutions?
No. For very dilute buffers (below about 1 mM), the autoprotolysis of water contributes significantly to [H⁺] and the simple equation breaks down. A full equilibrium calculation is needed.
Why choose a buffer close to the target pH?
Buffer capacity — the ability to resist pH change — is maximum at pH = pKa and drops off sharply beyond ±1 unit. A buffer outside this range provides little protection against pH shifts.
Does dilution change buffer pH?
Ideally, no. Diluting a buffer preserves the [A⁻]/[HA] ratio (both concentrations decrease proportionally), so pH should remain constant. In practice, very high dilution allows water autoprotolysis to become significant and pH can drift.
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