Acids, Bases and Equilibrium
What the exam tests
Acids, Bases and Equilibrium covers Brønsted–Lowry definitions, strong and weak acids, pH and pOH, Kw, Ka, titration curves, indicators, buffers and Le Chatelier's principle.
Core rules
- Brønsted–Lowry: an acid is a proton donor, a base a proton acceptor; they differ by H⁺ to form a conjugate pair.
- Strong acids ionise completely; weak acids ionise only partially and have a Ka.
- pH = −log₁₀[H⁺]; pOH = −log₁₀[OH⁻]; at 25 °C pH + pOH = 14.
- Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 25 °C.
- Ka = [H⁺][A⁻]/[HA]; a larger Ka means a stronger weak acid.
- Le Chatelier: a system at equilibrium shifts to oppose any change in concentration, pressure or temperature.
- A buffer resists pH change; it needs a weak acid and its salt.
Worked example
Find the pH of 0.010 M HCl and of 0.010 M NaOH at 25 °C. HCl is strong, so [H⁺] = 0.010 M and pH = −log(0.010) = 2.0. NaOH is strong, so [OH⁻] = 0.010 M, pOH = 2.0, pH = 14 − 2.0 = 12.0.
Traps
- Using 14 − pH for a solution not at 25 °C.
- Assuming a weak acid of the same concentration has the same pH as a strong one.
- Forgetting that dilution moves pH towards 7, not past it.
- Shifting the equilibrium in the wrong direction when temperature changes.
- Choosing an indicator whose range does not span the equivalence pH.