Covers Lessons 1–6: Acid-Base Models, Nomenclature & Reactions, Enthalpy of Neutralisation, Everyday Applications, and Strong vs Weak Acids & Bases.
Lesson Summaries — Quick Review
The Arrhenius model defines acids as H⁺ producers and bases as OH⁻ producers in water. Brønsted-Lowry extends this: acids are proton donors, bases are proton acceptors. Free H⁺ does not exist in water — it immediately bonds to H₂O to form H₃O⁺ (the hydronium ion). Every acid–base reaction involves conjugate pairs.
Acids are named systematically: binary acids (HCl → hydrochloric acid), oxyacids (H₂SO₄ → sulfuric acid). Acids react predictably with metals, metal oxides/hydroxides, carbonates, and hydrogen carbonates. Indicators change colour across pH ranges — choose an indicator whose endpoint matches the equivalence point of the reaction.
Neutralisation is exothermic. For strong acid + strong base, ΔH ≈ −57 kJ mol⁻¹ because the same net ionic reaction occurs: H⁺(aq) + OH⁻(aq) → H₂O(l). Calculate using q = mcΔT, then n = c × V, then ΔH = q/n. Weak acids/bases give less exothermic values because energy is consumed breaking incomplete dissociation.
Neutralisation is used in antacids (Mg(OH)₂, CaCO₃), agriculture (lime to raise soil pH), water treatment, and industrial processes. Excess stomach acid is neutralised — but not completely (pH would overshoot). The choice of neutralising agent depends on cost, availability, reaction speed, and safety.
Strong acids/bases dissociate completely (→); weak acids/bases partially dissociate (⇌). Strength ≠ concentration. A 0.1 M HCl solution has pH ≈ 1 (strong); 0.1 M CH₃COOH has pH ≈ 2.9 (weak). Strong electrolytes conduct electricity better. Indicators of strength: conductivity, rate of reaction with metals, pH for equivalent concentrations.
Consolidation of IQ1 so far: identifying strong/weak acids and bases, writing correct arrow notation, predicting comparative pH and conductivity, interpreting experimental evidence. Band 6 responses connect particle-level explanations (extent of dissociation) to observable properties (pH, conductivity, reaction rate).
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Answers are checked automatically. Short answer marking guide is in the Answers accordion below.
Multiple Choice — 20 Questions (1 mark each)
According to the Brønsted-Lowry model, which of the following is the conjugate base of H₂SO₄?
Why is H⁺ represented as H₃O⁺ in aqueous solutions?
Which equation correctly represents the reaction of sodium carbonate with hydrochloric acid?
A student mixes 50.0 mL of 1.0 mol L⁻¹ HCl with 50.0 mL of 1.0 mol L⁻¹ NaOH and records a temperature rise of 6.8°C. Using q = mcΔT (assume ρ = 1.0 g mL⁻¹, c = 4.18 J g⁻¹ K⁻¹), what is the enthalpy of neutralisation?
The standard enthalpy of neutralisation for any strong acid with any strong base is approximately −57 kJ mol⁻¹ because:
Short Answer — 3 Questions
4 marksA student dissolves equal concentrations of HCl and CH₃COOH in water. Compare and explain the pH and electrical conductivity of the two solutions at the particle level.
3 marksA student mixes 100 mL of 0.5 mol L⁻¹ H₂SO₄ with excess CaCO₃. Write the balanced equation for this reaction, identify the type of reaction, and predict one observable change.
5 marksExplain why antacids containing Mg(OH)₂ are preferred over NaOH for treating excess stomach acid. In your response, refer to the nature of the neutralisation reaction and patient safety considerations.
The conjugate base is formed when H₂SO₄ donates one proton (H⁺), yielding HSO₄⁻ (hydrogen sulfate ion). SO₄²⁻ would be the conjugate base if both protons were donated simultaneously, but Brønsted-Lowry defines a conjugate base as the species remaining after ONE proton is removed.
A proton (H⁺) is a bare nucleus with extremely high charge density. It is immediately attracted to the lone pair electrons on a water molecule's oxygen, forming a coordinate covalent bond and producing H₃O⁺. This is why "H⁺(aq)" is more accurately written as "H₃O⁺(aq)".
Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂. The carbonate reacts with two moles of acid (balancing the 2 Na⁺ ions and the CO₃²⁻). Products are a salt, water, and carbon dioxide gas.
q = mcΔT = 100 g × 4.18 J g⁻¹ K⁻¹ × 6.8 K = 2842.4 J = 2.842 kJ. n(HCl) = 0.050 L × 1.0 mol L⁻¹ = 0.050 mol. ΔH = −2.842 / 0.050 = −56.8 kJ mol⁻¹. Negative because the reaction is exothermic (temperature rose).
Strong acids and bases are fully dissociated, so the reaction is simply H⁺(aq) + OH⁻(aq) → H₂O(l) regardless of the counter-ions. This constant net ionic equation means the energy change is always the same (~−57 kJ mol⁻¹).
• HCl (strong acid) dissociates completely: HCl → H⁺ + Cl⁻, giving a higher [H⁺] and therefore a lower pH (~1 for 0.1 M). (1 mark)
• CH₃COOH (weak acid) only partially dissociates: CH₃COOH ⇌ CH₃COO⁻ + H⁺, giving fewer H⁺ ions and a higher pH (~2.9 for 0.1 M). (1 mark)
• Conductivity is proportional to ion concentration. HCl produces more ions per mole dissolved → higher conductivity. (1 mark)
• CH₃COOH solution has fewer free ions → lower conductivity despite equal concentration. (1 mark)
• Balanced equation: H₂SO₄ + CaCO₃ → CaSO₄ + H₂O + CO₂ (1 mark — must be balanced).
• Acid–carbonate reaction (type: neutralisation / acid–base / acid–carbonate). (1 mark)
• Observable change: effervescence/bubbling (CO₂ gas produced); solid CaCO₃ dissolves; temperature increases slightly. (1 mark — any one valid observation)
• NaOH is a strong base that dissociates completely → rapid, potentially over-neutralisation (pH spikes above 7), causing alkaline burns to stomach lining. (1 mark)
• Mg(OH)₂ is a weak base / sparingly soluble → releases OH⁻ slowly, giving a gentler pH increase. (1 mark)
• The neutralisation reaction: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O. (1 mark)
• Mg(OH)₂ acts as a buffer of pH — its low solubility self-limits the OH⁻ release, preventing the stomach from becoming dangerously alkaline. (1 mark)
• Mg²⁺ ions are non-toxic at therapeutic doses, making Mg(OH)₂ safe for oral consumption. (1 mark)
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