Buffer Solution
A buffer solution is an aqueous system that resists significant changes in pH when small amounts of acid or base are added, typically composed of a weak acid and its conjugate base (or a weak base and its conjugate acid) in similar concentrations. Buffers operate by consuming added H⁺ or OH⁻ through equilibrium reactions, maintaining the pH within a narrow range. Buffer systems are critical in biological organisms (blood pH 7.35–7.45 maintained by carbonate/bicarbonate), pharmaceuticals, laboratory experiments, and industrial fermentation.
Key Formula
pH = pKa + log([A⁻]/[HA]) (Henderson-Hasselbalch equation)
LaTeX: \text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}
| Symbol | Meaning | Unit |
|---|---|---|
| pH | pH of the buffer solution | dimensionless |
| pKa | Negative log of the acid dissociation constant | dimensionless |
| [A⁻] | Concentration of conjugate base | mol/L |
| [HA] | Concentration of weak acid | mol/L |
Worked Example
Problem
Prepare a buffer using acetic acid (pKa = 4.74). What ratio of sodium acetate (CH₃COONa) to acetic acid (CH₃COOH) is needed for a buffer with pH = 5.04?
Solution
Step 1: Apply the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA]) Step 2: 5.04 = 4.74 + log([CH₃COO⁻]/[CH₃COOH]) Step 3: log([CH₃COO⁻]/[CH₃COOH]) = 5.04 - 4.74 = 0.30 Step 4: [CH₃COO⁻]/[CH₃COOH] = 10⁰·³⁰ = 2.0 Step 5: Use 2 parts sodium acetate to 1 part acetic acid.
Answer
[CH₃COO⁻] / [CH₃COOH] = 2:1
Common Buffer Systems and Their pH Ranges
| Buffer System | Weak Acid / Conjugate Base | pKa | Effective pH Range |
|---|---|---|---|
| Acetate buffer | Acetic acid / Sodium acetate | 4.74 | 3.8 – 5.8 |
| Phosphate buffer | H₂PO₄⁻ / HPO₄²⁻ | 7.20 | 6.2 – 8.2 |
| Carbonate buffer | H₂CO₃ / HCO₃⁻ | 6.37 | 5.4 – 7.4 |
| Ammonium buffer | NH₄⁺ / NH₃ | 9.25 | 8.3 – 10.3 |
| TRIS buffer | TRIS·H⁺ / TRIS | 8.06 | 7.0 – 9.0 |
| Citrate buffer | Citric acid / Citrate | 3.13–6.40 | 2.1 – 7.4 |
Interactive Tools
Wikimedia Commons, CC BY-SA
Related Terms
Acid Dissociation Constant (Ka)
The acid dissociation constant (Ka) is the equilibrium constant for the ionisation of an acid in water, quantifying the extent to which an acid donates its proton to water at a given temperature. A large Ka (> 1) indicates a strong acid that dissociates almost completely, while a small Ka (< 10⁻³) indicates a weak acid with limited dissociation. Ka values are fundamental in predicting reaction directions, calculating pH of weak acid solutions, and designing buffer systems.
Acid-Base Titration
Acid-base titration is a quantitative analytical technique in which a solution of known concentration (titrant) is gradually added to a solution of unknown concentration (analyte) until the reaction reaches its equivalence point, where the moles of acid exactly equal the moles of base. An indicator or pH meter is used to detect the endpoint of the titration, allowing calculation of the unknown concentration. Titration is widely used in medicine, food testing, environmental science, and quality control to determine the concentration of acids or bases in samples.
pH Scale
The pH scale is a logarithmic measure of the hydrogen ion concentration [H⁺] in a solution, ranging from 0 to 14 at 25 °C, where values below 7 indicate acidic conditions, 7 is neutral, and above 7 is basic or alkaline. Introduced by Danish chemist Søren Peder Lauritz Sørensen in 1909, the scale compresses a trillion-fold range of H⁺ concentrations into a convenient 0–14 range. pH measurement is critical in agriculture, biology, medicine, food science, and environmental monitoring.
The word "buffer" derives from the Old French "buffe" (a blow) and Middle English "buff" (to absorb a blow), reflecting the solution's ability to absorb the "impact" of added acid or base. The Henderson-Hasselbalch equation, which describes buffer behaviour quantitatively, was developed independently by Lawrence Joseph Henderson (1908) and Karl Albert Hasselbalch (1916).