Solutions, Electrolytic Dissociation and Ionic Equilibria
Solutions, Electrolytic Dissociation and Ionic Equilibria covers concentration units, the behaviour of strong and weak electrolytes, pH, buffers, hydrolysis of salts and solubility products.
Core rules
- Molarity c = n/V (mol/L) depends on volume and therefore on temperature; molality is temperature-independent. Dilution: c₁V₁ = c₂V₂.
- Strong acids, strong bases and soluble salts dissociate completely; weak acids and bases obey Ka, Kb and Ostwald's dilution law, α ≈ √(Ka/c).
- Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C, so pH + pOH = 14.
- A buffer of weak acid and its salt: pH = pKa + log([A⁻]/[HA]).
- Salt hydrolysis: strong acid + weak base gives an acidic solution; weak acid + strong base gives a basic one; strong + strong is neutral.
Worked example
Find the pH of 0.10 M NH₃ (Kb = 1.8 × 10⁻⁵). [OH⁻] = √(Kb × c) = √(1.8 × 10⁻⁵ × 0.10) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ mol/L. pOH = 2.87, so pH = 14 − 2.87 = 11.13. The approximation is valid because α ≈ 1.3 %, well below the 5 % threshold.
Traps
- Applying pH = −log c to a weak acid without checking the degree of dissociation.
- Forgetting that tenfold dilution of a weak acid raises pH by only about 0.5, not a full unit.
- Swapping Ka and Kb, or using pKw = 14 at a temperature other than 25 °C.
- Ignoring the common-ion effect: adding AgNO₃ to saturated AgCl lowers chloride solubility.