This AP Chemistry lesson on buffer solutions notes gives the key ideas first, then worked examples and diagrams. After that, practise with the worksheet on buffer solutions and check yourself with the quiz.
Quick answer: A buffer is a solution of a weak acid and its conjugate base that resists pH change. Its pH follows pH = pKa + log([A⁻]/[HA]). In a titration curve of a weak acid with a strong base, pH equals pKa at half-equivalence, and the equivalence point lies above 7. Choose an indicator whose range lies on the steep section.
Level: AP Chemistry · Curriculum: AP Chem Unit 8
Buffer solutions notes: the key ideas
What is a buffer and how does it work?
A buffer is a solution that resists a change in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base (such as CH₃COOH and CH₃COO⁻) or a weak base and its conjugate acid (such as NH₃ and NH₄⁺), both in large amounts.
Added acid reacts with the base component: CH₃COO⁻ + H₃O⁺ → CH₃COOH + H₂O. Added base reacts with the acid component: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O. Either way the ratio [A⁻]/[HA] shifts only a little, and pH depends on that ratio, not on the amounts themselves.
How do you calculate the pH of a buffer?
Henderson–Hasselbalch: pH = pKa + log₁₀([A⁻] / [HA]). Here pKa = −log₁₀Ka, [A⁻] is the conjugate base and [HA] the weak acid, both in mol/dm³ (or in mol, because the volume cancels in the ratio).
A buffer holds 0.20 mol/dm³ CH₃COOH and 0.30 mol/dm³ CH₃COONa (pKa = 4.74). pH = 4.74 + log₁₀(0.30/0.20) = 4.74 + 0.18 = 4.92.
To 1.00 dm³ of that buffer Aarav adds 0.010 mol HCl. The acid converts 0.010 mol of A⁻ to HA, so A⁻ = 0.29 mol and HA = 0.21 mol. pH = 4.74 + log₁₀(0.29/0.21) = 4.88, a drop of only 0.04. The same 0.010 mol HCl in 1.00 dm³ of pure water would give pH 2.00.
What is buffer capacity?
Capacity is the amount of acid or base a buffer can absorb before its pH changes a lot. It grows with the concentrations of HA and A⁻, and it is greatest when the two are equal, where pH = pKa. A buffer works well only when 0.1 ≤ [A⁻]/[HA] ≤ 10, that is, pH within one unit of pKa. To choose a buffer for a target pH, pick the acid whose pKa is closest to it.
| Buffer pair | pKa of the acid | Useful pH range |
|---|---|---|
| HCOOH / HCOO⁻ | 3.75 | 2.75 to 4.75 |
| CH₃COOH / CH₃COO⁻ | 4.74 | 3.74 to 5.74 |
| H₂PO₄⁻ / HPO₄²⁻ | 7.20 | 6.20 to 8.20 |
| NH₄⁺ / NH₃ | 9.25 | 8.25 to 10.25 |
What do titration curves show?
A titration curve plots pH against the volume of titrant added. The equivalence point is where the moles of base added equal the moles of acid originally present. For a strong acid with a strong base the curve is almost vertical around pH 7.
For a weak acid with a strong base, the curve begins at a higher pH, passes through a flat buffer region and ends its steep rise above pH 7, because the conjugate base A⁻ is itself weakly basic. Halfway to equivalence, [HA] = [A⁻], so pH = pKa. This is the quickest way to measure pKa.
An indicator is suitable when its colour-change range falls inside the steep part of the curve. Strong acid with strong base: almost any indicator. Weak acid with strong base (equivalence pH about 8 to 9): phenolphthalein. Strong acid with weak base (equivalence below 7): methyl orange.
Students assume every titration ends at pH 7. Only a strong acid with a strong base does. Also do not use the Henderson–Hasselbalch equation at the equivalence point, where no buffer remains, or before any base has been added.
Always do the reaction first, then the equilibrium. With a strong acid or base added, use a mole table to find what is left, and only then decide which formula (Henderson–Hasselbalch, Ka or leftover strong species) applies.
- Henderson–Hasselbalch: pH = pKa + log₁₀([A⁻]/[HA]).
- Buffer capacity is greatest when [A⁻] = [HA], where pH = pKa; useful range is pKa ± 1.
- At half-equivalence in a weak acid titration, pH = pKa.
- Equivalence pH: 7 (strong–strong), above 7 (weak acid–strong base), below 7 (strong acid–weak base).
- Choose an indicator whose colour-change range lies inside the steep part of the curve.
Buffers, titration curves and indicators: frequently asked questions
Why does a buffer resist changes in pH?
It holds a weak acid to neutralise added base and a weak base to neutralise added acid. The ratio of the two changes only slightly, so the pH barely moves.
When is the Henderson–Hasselbalch equation valid?
It applies when both HA and A⁻ are present in comparable amounts, roughly with a ratio between 0.1 and 10. It does not apply to a strong acid solution or at the equivalence point.
Why is the equivalence point of ethanoic acid with NaOH above pH 7?
At equivalence the solution contains only CH₃COO⁻, a weak base that reacts with water to form OH⁻. The result is a basic solution, about pH 8.7 for 0.050 mol/dm³ ethanoate.
How do you find pKa from a titration curve?
Read the pH at half the equivalence volume. At that point [HA] = [A⁻], so pH = pKa.
How do you choose an indicator for a titration?
Choose one whose colour-change range lies within the steep part of the curve. Phenolphthalein (8.3 to 10.0) suits a weak acid with a strong base, while methyl orange (3.1 to 4.4) suits a strong acid with a weak base.
