Chemical Equilibrium Focusing on Acid-Base Systems
Core ideas
Acid-base equilibrium applies dynamic equilibrium and Le Chatelier's principle to proton transfer. Write Ka for HA ⇌ H+ + A- and Kb for B + H2O ⇌ BH+ + OH-. Conjugate acid-base pairs differ by one H+; stronger acids have weaker conjugate bases. Kw = [H+][OH-] = 1.0×10^-14 at 25°C, so pH + pOH = 14 only at that temperature. pH = -log[H3O+]; pOH is defined similarly.
For weak acids, do not set [H+] equal to the initial concentration. Use an ICE table: Ka ≈ x^2/([HA]0 - x); if x is less than about 5% of [HA]0, the approximation is acceptable. Salts hydrolyze: a salt of a weak acid and strong base gives a basic solution; a salt of a weak base and strong acid gives an acidic solution. Buffers resist pH change when they contain a weak acid/base pair in comparable amounts; Henderson–Hasselbalch is pH = pKa + log([A-]/[HA]). On a titration curve, the equivalence point is where moles of acid and base match, not necessarily pH 7. Choose an indicator whose pKa is near the equivalence-point pH.
Traps: using pH 7 as the equivalence point for weak acid–strong base; ignoring dilution; and applying Kw at temperatures other than 25°C without adjusting.
Worked example: 0.10 M HA with Ka = 1.8×10^-5: x^2/0.10 = 1.8×10^-5, x = 1.3×10^-3 M, pH = 2.87. Check: x is 1.3% of 0.10 M, so the approximation is valid.