Buffers: Solutions That Resist pH Change

Core Concept

A buffer is a solution that resists changes in pH when small amounts of strong acid or base are added. It consists of a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable concentrations. Because both members are present, the weak acid can neutralize added OH⁻ and the conjugate base can neutralize added H₃O⁺, converting a strong acid/base assault into a tiny shift of a weak-acid equilibrium. This is why blood, seawater, and many biological fluids hold their pH nearly constant.

Key Ideas

• Composition: a weak acid + its conjugate base (e.g., CH₃COOH/CH₃COO⁻) or a weak base + its conjugate acid (NH₃/NH₄⁺).
• How it absorbs acid: the conjugate base reacts: A⁻ + H₃O⁺ → HA + H₂O.
• How it absorbs base: the weak acid reacts: HA + OH⁻ → A⁻ + H₂O.
• Key requirement: both members must be present in significant, comparable amounts; a buffer is NOT a strong acid/strong base mix.
• Buffer range: works best within about ±1 pH unit of pKa.
• Buffer capacity: the amount of strong acid/base a buffer can absorb; highest when [A⁻] ≈ [HA] and when total concentrations are high.

Equations and Variables

• Buffer equilibrium: HA ⇌ H₃O⁺ + A⁻
• pH = pKa + log([A⁻]/[HA]) (Henderson–Hasselbalch)
• Acid absorbed: A⁻ + H₃O⁺ → HA + H₂O
• Base absorbed: HA + OH⁻ → A⁻ + H₂O
• Buffer range: pKa ± 1

How It Works

1. A buffer holds both halves of a conjugate pair. The weak acid (HA) sits ready to donate a proton; the conjugate base (A⁻) sits ready to accept one.
2. Add strong acid (H₃O⁺): the added protons are consumed by A⁻ → HA. The pH barely moves because the strong acid is replaced by the weak acid HA.
3. Add strong base (OH⁻): the hydroxide is consumed by HA → A⁻ + H₂O. Again, a strong species is swapped for a weak one.
4. The ratio [A⁻]/[HA] changes slightly, and the pH shifts by the log of that ratio — a tiny change for modest additions.
5. The buffer fails (capacity exceeded) once one member of the pair is essentially used up.

Worked Example

A buffer is 0.10 M CH₃COOH and 0.10 M CH₃COONa (pKa = 4.74). (a) Find its pH. (b) Find the new pH after adding 0.010 mol of NaOH to 1.0 L of the buffer.

(a) pH = pKa + log([A⁻]/[HA]) = 4.74 + log(0.10/0.10) = 4.74 + 0 = 4.74.

(b) Stoichiometry first: NaOH (0.010 mol) reacts with HA:

HA + OH⁻ → A⁻ + H₂O

New amounts: HA = 0.10 − 0.010 = 0.090 M; A⁻ = 0.10 + 0.010 = 0.110 M.

pH = 4.74 + log(0.110/0.090) = 4.74 + log(1.222) = 4.74 + 0.087 = 4.83.

Adding 0.010 mol of strong base moved the pH by only +0.09 units. The same NaOH added to pure water would jump the pH from 7 to ~12.

How it works

1. A buffer holds both halves of a conjugate pair. The weak acid (HA) sits ready to donate a proton; the conjugate base (A⁻) sits ready to accept one.
2. Add strong acid (H₃O⁺): the added protons are consumed by A⁻ → HA. The pH barely moves because the strong acid is replaced by the weak acid HA.
3. Add strong base (OH⁻): the hydroxide is consumed by HA → A⁻ + H₂O. Again, a strong species is swapped for a weak one.
4. The ratio [A⁻]/[HA] changes slightly, and the pH shifts by the log of that ratio — a tiny change for modest additions.
5. The buffer fails (capacity exceeded) once one member of the pair is essentially used up.

Common confusions

• "A buffer is any acid + base mixture." — It must be a weak acid with its conjugate base; HCl + NaCl is not a buffer (Cl⁻ is too weak a base).
• "A buffer keeps pH exactly constant." — It resists change; pH still moves a little, and it fails entirely once capacity is exceeded.
• "The conjugate base can be any anion." — It must be the conjugate base of the buffer's weak acid (acetate with acetic acid, not chloride).
• "Strong acid + strong base makes a good buffer." — No; they neutralize to a neutral salt with no buffering ability.
• "Buffers work at any pH." — Each buffer only works near its pKa (±1 unit).

Quick review

• Buffer = weak acid + conjugate base (or weak base + conjugate acid).
• Acid absorbed by A⁻, base absorbed by HA.
• pH = pKa + log([A⁻]/[HA]).
• Max capacity when [A⁻] = [HA]; range = pKa ± 1.
• Biological example: blood's H₂CO₃/HCO₃⁻ system.