CHEM10007 · Fundamentals Of Chemistry
Acids, Bases & Aqueous Equilibria
Equilibrium applied to water chemistry. You define acids and bases by the Brønsted–Lowry model, identify conjugate pairs, and use the pH/pOH scale with Kw. For strong acids and bases pH comes straight from concentration; for weak acids you set up an ICE table with Ka (small-x approximation). The chapter closes with buffers and solubility equilibria (Ksp) and the common-ion effect.
What this chapter covers
- 01Brønsted–Lowry acids (proton donors) and bases (proton acceptors); conjugate acid–base pairs; amphiprotic species
- 02Self-ionisation of water: Kw = [H3O+][OH−] = 1.0 × 10−14 at 25 °C
- 03pH = −log10[H3O+], pOH = −log10[OH−], pH + pOH = 14.0
- 04Strong acids/bases: pH directly from concentration (full dissociation)
- 05Weak acids/bases: Ka / Kb, pKa = −log Ka, pKa + pKb = 14.0
- 06ICE-table method for weak-acid pH and the small-x approximation; polyprotic acids
- 07Buffers: a weak acid plus its conjugate base resisting pH change
- 08Solubility equilibria: Ksp, molar solubility, the common-ion effect, and Q vs Ksp for precipitation
pH of a weak acid via an ICE table
- 1 mark — set up ICE tableEquilibrium: CH3COOH + H2O ⇌ CH3COO− + H3O+. ICE: initial 0.200, change −x/+x/+x.
- 1 mark — K<sub>a</sub> expression with approximationKa = x2 / (0.200 − x) ≈ x2 / 0.200 = 1.8 × 10−5 (small-x approximation).
- 1 mark — solve for [H<sub>3</sub>O<sup>+</sup>] and checkx2 = 3.6 × 10−6, so x = [H3O+] = 1.9 × 10−3 M (x is ≈ 0.95 % of 0.200, so the approximation is valid).
- 1 mark — pHpH = −log(1.9 × 10−3) = 2.72.
- 1 mark — conjugate base and pK<sub>b</sub>Conjugate base = CH3COO−; pKa = −log(1.8 × 10−5) = 4.74, so pKb = 14.0 − 4.74 = 9.26.
Key terms
- Brønsted–Lowry acid/base
- An acid is a proton (H+) donor and a base is a proton acceptor; donating or accepting a proton creates a conjugate base or acid.
- Kw (ionic product of water)
- [H3O+][OH−] = 1.0 × 10−14 at 25 °C, linking pH and pOH through pH + pOH = 14.0.
- ICE table
- An Initial–Change–Equilibrium bookkeeping table used to solve weak-acid/base equilibria, typically with the small-x approximation.
- Buffer
- A solution of a weak acid and its conjugate base (or vice versa) that resists pH change when small amounts of acid or base are added.
- Solubility product (Ksp)
- The equilibrium constant for a sparingly soluble salt dissolving; molar solubility derives from it, and a common ion lowers solubility.
Acids, Bases & Aqueous Equilibria FAQ
When can I use the small-x approximation?
When the acid is weak and reasonably concentrated, x (the amount ionised) is tiny compared with the initial concentration, so 0.200 − x ≈ 0.200. Always check afterwards: if x exceeds about 5 % of the initial value, redo the calculation with the full quadratic.
How is strong-acid pH different from weak-acid pH?
A strong acid dissociates completely, so [H3O+] equals the acid concentration and pH follows directly. A weak acid only partially ionises, so you need Ka and an ICE table to find [H3O+].
What is the common-ion effect on solubility?
Adding an ion already present in the solubility equilibrium shifts it left (Le Chatelier), so the salt becomes less soluble. For example, AgCl is far less soluble in NaCl solution than in pure water.
Exam move
The ICE-table weak-acid calculation is one of the highest-yield Section B archetypes — practise it until the set up, approximation and validity check are automatic. Keep pH ↔ pOH ↔ Kw and pKa ↔ pKb conversions fluent, and learn to read Ka, Kb and Ksp from the supplied appendix. For Ksp, drill both the pure-water molar-solubility case and the common-ion case, since they are easy to mix up under time pressure.
