Acids, Bases, Salts and Ionic Equilibria
This domain tests the quantitative treatment of proton transfer: pH calculations for strong and weak acids, buffers, salt hydrolysis, titration curves, indicators and solubility equilibria.
pH, Kw and strong acids
pH = −log10[H+], and [H+] = 10−pH. Kw = [H+][OH−] = 1.0 × 10−14 at 298 K, so neutral pH is 7.00 only at that temperature; Kw rises with temperature. Strong acids and bases dissociate completely, so a 0.010 M HCl solution has pH 2.00.
Weak acids and bases
For HA with Ka, [H+] ≈ √(Ka·c) when the degree of dissociation is small (check: less than about 5%). pKa = −log Ka. For a weak base use Kb and pOH, then pH = 14 − pOH. Diluting a weak acid raises its pH but by less than for a strong acid, because the degree of ionisation increases.
Buffers
A buffer is a weak acid with its conjugate base (or weak base with conjugate acid). Henderson–Hasselbalch: pH = pKa + log([A−]/[HA]). A buffer works best when pH ≈ pKa; it resists added acid or base because the reservoir components neutralise them.
Salts, titration curves and Ksp
Salt hydrolysis: a salt of a weak acid and strong base gives an alkaline solution; a salt of a weak base and strong acid gives an acidic solution; strong acid–strong base salts are neutral. On a titration curve the steep section determines indicator choice: phenolphthalein for weak acid–strong base, methyl orange for strong acid–weak base. For sparingly soluble salts, Ksp = [M+]m[X−]n and a common ion lowers solubility.
Worked example
Find the pH of 0.100 M ethanoic acid, Ka = 1.8 × 10−5. [H+] = √(1.8 × 10−5 × 0.100) = 1.34 × 10−3 M, so pH = 2.87.
Traps that cost marks
- Confusing 'strong' with 'concentrated'.
- Using the strong-acid formula for a weak acid without checking the approximation.
- Choosing an indicator whose range does not lie on the steep part of the curve.
- Treating the conjugate base of a strong acid as a significant base.
- Forgetting that Kw changes with temperature, shifting neutral pH.