QCE Chemistry — Unit 3
Acid-Base Equilibria — Flashcards & Quiz
QCE Chemistry Unit 3 examines acid-base equilibria through the Brønsted-Lowry proton-transfer model. You need to identify conjugate acid-base pairs, calculate pH and pOH, and work with Ka for weak acid systems. Buffer behaviour links these ideas together and is a favourite exam scenario because it combines Le Chatelier reasoning with equilibrium calculations.
Key Points
- Brønsted-Lowry: acids are proton donors, bases are proton acceptors. Every reaction has two conjugate acid-base pairs.
- pH = -log₁₀[H⁺]; pOH = -log₁₀[OH⁻]; pH + pOH = 14 at 25°C.
- Strong acids (HCl, HNO₃, H₂SO₄) fully dissociate; [H⁺] equals molar concentration directly.
- Weak acids partially dissociate: Ka = [H⁺][A⁻]/[HA]; for dilute solutions [H⁺] ≈ √(Ka · C).
- Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C; Kw is temperature-dependent.
- Buffers resist pH change — a weak acid plus its conjugate base neutralises added acid or base via Le Chatelier shifts.
Common Mistakes to Avoid
- Using [H⁺] = molar concentration for weak acids — only strong acids fully dissociate.
- Forgetting the minus sign in pH = –log[H⁺].
- Applying pH + pOH = 14 outside 25°C without noting Kw changes.
- Confusing Ka (equilibrium constant for acid dissociation) with pKa (= –log Ka).
- Missing the buffer requirement — you need BOTH weak acid AND conjugate base present.
Exam Strategy
QCAA Unit 3 acid-base questions often give you a weak acid, buffer, or titration scenario. Method: (1) identify whether the acid is strong or weak, (2) for weak, apply Ka and the square-root approximation, (3) for buffers, use pH = pKa + log([A⁻]/[HA]), (4) always state assumptions explicitly. Show working step by step — marks are awarded for method, not just the final number.
Sample Flashcards
Q1: Define acids and bases according to the Brønsted-Lowry theory.
A Brønsted-Lowry acid is a proton (H⁺) donor. A Brønsted-Lowry base is a proton (H⁺) acceptor. In any acid-base reaction, there are two conjugate acid-base pairs. Water is amphoteric — it can act as both an acid and a base depending on the reaction partner.
Q2: Distinguish between strong and weak acids in terms of equilibrium.
Strong acids fully dissociate in water — equilibrium lies completely to the right (Ka is very large). Examples: HCl, H₂SO₄, HNO₃. Weak acids partially dissociate — equilibrium lies to the left (Ka is small). A smaller Ka means a weaker acid. Weak acids have both undissociated acid molecules and ions present at equilibrium.
Q3: Define Ka and explain what it indicates.
Ka is the acid dissociation constant: Ka = [H₃O⁺][A⁻] / [HA]. It measures the strength of a weak acid. Larger Ka = stronger acid (more dissociation). Ka is constant at a given temperature. For strong acids, Ka is so large it is not usually quoted. Water is excluded from the expression.
Sample Quiz Questions
Q1: A Brønsted-Lowry acid is defined as a proton acceptor.
Answer: FALSE
A Brønsted-Lowry acid is a proton (H⁺) DONOR. A proton acceptor is a Brønsted-Lowry BASE.
Q2: Strong acids fully dissociate in aqueous solution.
Answer: TRUE
Strong acids like HCl, H₂SO₄ and HNO₃ completely dissociate into ions in water. There are essentially no undissociated acid molecules remaining.
Q3: Water is amphoteric, meaning it can act as both an acid and a base.
Answer: TRUE
Water can donate a proton (acting as an acid) or accept a proton (acting as a base) depending on the reaction partner.
Revision Tip
pH calculations need procedural fluency — drill a Revizi deck with 10+ problems covering strong acids, weak acids, and buffer systems until the steps are automatic.
Related Concepts
Last updated: March 2026 · 3 flashcards · 3 quiz questions