Buffer Solution Calculators
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Henderson-Hasselbalch Equation
pH = pKa + log([A⁻]/[HA])
When [A⁻] = [HA]: pH = pKa. Buffer range: pKa ± 1 pH unit. Example: acetate buffer (pKa 4.75): ratio [CH₃COO⁻]/[CH₃COOH] needed for pH 5.0: log(ratio) = 5.0 − 4.75 = 0.25; ratio = 10^0.25 = 1.78.
How Buffers Work
Adding H⁺: A⁻ + H⁺ → HA (conjugate base absorbs acid — pH barely changes). Adding OH⁻: HA + OH⁻ → A⁻ + H₂O (weak acid neutralizes base — pH barely changes). After many additions, buffering capacity is exhausted when either A⁻ or HA is nearly depleted.
Blood Bicarbonate Buffer
CO₂ + H₂O ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺. pH = 6.1 + log([HCO₃⁻]/[H₂CO₃]). Normal blood: [HCO₃⁻] ≈ 24 mM; [H₂CO₃] ≈ 1.2 mM; pH = 6.1 + log(24/1.2) = 6.1 + log(20) = 6.1 + 1.301 = 7.40.
Common Lab Buffers
- Acetate: pH 3.8–5.8 (pKa 4.75)
- Phosphate (PBS): pH 6.2–8.2 (pKa₂ 7.20)
- HEPES: pH 6.8–8.2 (pKa 7.48)
- Tris-HCl: pH 7.0–9.0 (pKa 8.06); temperature-sensitive
Glossary
Frequently Asked Questions
A buffer resists pH changes by containing both a weak acid (HA) and its conjugate base (A⁻) in solution: When acid (H⁺) is added: A⁻ + H⁺ → HA — the conjugate base absorbs the protons, preventing a large pH drop. When base (OH⁻) is added: HA + OH⁻ → A⁻ + H₂O — the weak acid neutralizes the base, preventing a large pH rise. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Buffer works best when [A⁻] ≈ [HA] (i.e., pH ≈ pKa) — maximum buffering capacity. Range: effective within ±1 pH unit of pKa. When one component is nearly consumed, buffering capacity is exhausted.
pH = pKa + log([A⁻]/[HA]). Example 1: prepare 0.1 M acetate buffer (pKa = 4.75) using 0.06 M CH₃COO⁻ and 0.04 M CH₃COOH: pH = 4.75 + log(0.06/0.04) = 4.75 + log(1.5) = 4.75 + 0.176 = 4.93. Example 2: needed pH 7.20 using phosphate (pKa₂ = 7.20): log([HPO₄²⁻]/[H₂PO₄⁻]) = 7.20 − 7.20 = 0; [HPO₄²⁻]/[H₂PO₄⁻] = 1.0 → equal concentrations of both forms. At pH = pKa: exactly 50% of the buffer is in each form; this is also the point of maximum buffering capacity.
Blood pH is maintained at 7.35–7.45 by the bicarbonate buffer: CO₂ (dissolved) + H₂O ⇌ H₂CO₃ ⇌ HCO₃⁻ + H⁺. pH = 6.1 + log([HCO₃⁻]/[CO₂]). Normal values: [HCO₃⁻] = 24 mEq/L; [CO₂] = 1.2 mEq/L (pCO₂ = 40 mmHg × 0.0301). pH = 6.1 + log(24/1.2) = 6.1 + 1.301 = 7.40. Respiratory regulation: lungs control pCO₂ by adjusting ventilation rate — hyperventilation blows off CO₂ → pH rises; hypoventilation retains CO₂ → pH falls. Metabolic regulation: kidneys control [HCO₃⁻] by excreting or reabsorbing bicarbonate. Clinical: metabolic acidosis (low HCO₃⁻); metabolic alkalosis (high HCO₃⁻); respiratory acidosis (high CO₂); respiratory alkalosis (low CO₂).
Buffer capacity (β) = ability to resist pH change per unit of acid or base added. Maximum β at pH = pKa (where [A⁻] = [HA]). β is proportional to total buffer concentration (C): β ≈ 2.303 × C × Ka × [H⁺] / (Ka + [H⁺])². At pH = pKa: β_max ≈ 0.576 × C. Doubling buffer concentration doubles β — the buffer can neutralize twice as much acid or base before the pH changes significantly. Practical: for experiments where large pH perturbations are expected (acid-producing cells, enzyme reactions generating acid), use higher buffer concentrations (50–200 mM) rather than the minimum (10–25 mM). Limitation: high buffer concentration increases osmolarity, which may affect cellular assays.