Key ideas
The pH scale measures acidity on a logarithmic scale. Strong acids and bases dissociate completely, weak ones only partly, and a buffer resists changes in pH.
- pH = −log₁₀[H⁺], and at 25 °C pH + pOH = 14. Each pH unit is a factor of ten in [H⁺].
- A strong acid gives [H⁺] equal to its concentration. A weak acid gives [H⁺] ≈ √(K_a·C) when little of it dissociates.
- A buffer is a weak acid with its conjugate base. The Henderson–Hasselbalch equation is pH = pKa + log₁₀([A⁻] ÷ [HA]).
- At the equivalence point of a titration, moles of acid equal moles of base: M_aV_a = M_bV_b for a one-to-one reaction.
Key equations
pH = −log₁₀[H⁺]pH + pOH = 14 (25 °C)K_a = [H⁺][A⁻] ÷ [HA][H⁺] ≈ √(K_a·C)pH = pKa + log₁₀([A⁻] ÷ [HA])M_aV_a = M_bV_b
Worked example
For 0.10 M acetic acid with K_a = 1.8 × 10⁻⁵, [H⁺] = √(1.8 × 10⁻⁵ × 0.10) = 1.3 × 10⁻³ M, so pH = 2.87. A buffer of 0.20 M acetic acid and 0.10 M acetate (pKa = 4.74) has pH = 4.74 + log(0.10 ÷ 0.20) = 4.74 − 0.30 = 4.44.
Common mistakes
- Treating a weak acid as if it dissociated completely.
- Putting the acid over the base in the Henderson–Hasselbalch ratio instead of base over acid.
- Forgetting that pH is logarithmic: pH 3 is ten times more acidic than pH 4, not 25 % more.
Practice
Test yourself with practice questions on this topic: Acids, bases and buffers