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:
Which of the following best explains why agricultural lime (CaCO₃) is added to acidic soil?
A factory neutralises acidic wastewater using crushed limestone rather than NaOH. The main advantage of limestone is:
Which equation correctly represents the behaviour of ammonia (NH₃) in water?
At 25°C, which 0.10 mol L⁻¹ solution has the highest pH?
When measuring electrical conductivity of 0.10 mol L⁻¹ solutions, which produces the highest reading?
A student adds equal masses of magnesium ribbon to separate flasks containing 0.5 mol L⁻¹ HCl and 0.5 mol L⁻¹ CH₃COOH. Which observation correctly distinguishes the acids?
A 0.001 mol L⁻¹ solution of HCl and a 0.10 mol L⁻¹ solution of CH₃COOH are compared. Which statement is correct?
At the particle level, why does a 0.10 mol L⁻¹ HCl solution have a lower pH than a 0.10 mol L⁻¹ HF solution, even though both are acids?
A student records the following data for two acids at 0.10 mol L⁻¹:
Acid X: pH = 1.0, conductivity = 42 mS cm⁻¹
Acid Y: pH = 2.9, conductivity = 1.1 mS cm⁻¹
What is the most reasonable conclusion?
Which reaction CANNOT be classified as an acid-base reaction using the Arrhenius model but CAN be classified using the Brønsted-Lowry model?
When 25.0 mL of 2.0 mol L⁻¹ HCl is mixed with 25.0 mL of 2.0 mol L⁻¹ NaOH, the temperature rises by 13.6°C. What is the enthalpy of neutralisation?
A student titrating a strong acid against a strong base should choose an indicator with:
Which of the following species acts as a weak base in aqueous solution?
A farmer notices that crop yields have decreased. Soil testing reveals a pH of 5.2. Which treatment is most appropriate and why?
A student proposes that "all acids are dangerous and should be neutralised completely." Which evaluation best addresses this statement?
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).
CaCO₃ is a sparingly soluble salt, not a strong base. It reacts slowly with soil acidity: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. This gradual neutralisation prevents pH from overshooting above 7, which would make the soil alkaline and lock up essential micronutrients (e.g. iron, manganese).
NaOH is a strong base that dissociates completely → rapid neutralisation → risk of over-neutralisation (pH > 7) and it is corrosive/expensive. Crushed limestone (CaCO₃) is cheap, abundant, safe to handle, and its low solubility means it self-limits the neutralisation rate.
NH₃ is a weak base, it only partially accepts protons from water. The equilibrium arrow (⇌) must be used. Option A uses → (complete reaction, wrong), B uses ← (reverse reaction, wrong), C uses ⇄ (not standard notation).
NH₃ is a weak base → produces OH⁻ → pH > 7 (~11 for 0.10 M). HCl and HNO₃ are strong acids (pH ~1). CH₃COOH is a weak acid (pH ~2.9). Therefore NH₃ has the highest pH.
H₂SO₄ is a strong diprotic acid: H₂SO₄ → 2H⁺ + SO₄²⁻. One mole produces three moles of ions → highest conductivity. CH₃COOH and NH₃ are weak → few ions. HF is weak (despite F being electronegative, the H–F bond is very strong).
HCl dissociates completely → [H⁺] = 0.5 mol L⁻¹. CH₃COOH partially dissociates → [H⁺] << 0.5 mol L⁻¹. Reaction rate with Mg depends on [H⁺] (collision theory). More H⁺ ions → more frequent effective collisions → faster H₂ production.
0.001 M HCl: [H⁺] = 0.001 → pH = 3. 0.10 M CH₃COOH: [H⁺] ≈ 10⁻³ to 10⁻² (partial dissociation) → pH ≈ 2–3. Because CH₃COOH is ~100× more concentrated, even its partial dissociation produces more H⁺ than the dilute HCl → lower pH for CH₃COOH.
HF is actually a weak acid (not strong), but the key reasoning stands: HCl dissociates completely (→) producing maximum [H⁺], while HF partially dissociates (⇌). The H–F bond is unusually strong due to high bond polarity and small atomic size, making dissociation difficult. More free H⁺ → lower pH.
The data shows Acid X has pH = 1.0 (high [H⁺]) and high conductivity (many ions) → strong acid. Acid Y has pH = 2.9 (lower [H⁺]) and low conductivity (few ions) → weak acid. Both are at the same stated concentration (0.10 M), so concentration is controlled.
The Arrhenius model requires water as solvent and defines bases as OH⁻ producers. In the gas-phase reaction NH₃ + HCl → NH₄Cl, there is no water and no OH⁻ produced, so Arrhenius cannot classify it. Brønsted-Lowry can: NH₃ accepts a proton (base), HCl donates a proton (acid).
q = mcΔT = 50 g × 4.18 × 13.6 = 2842.4 J. n(HCl) = 0.025 × 2.0 = 0.050 mol. ΔH = −2.842/0.050 = −56.8 ≈ −57 kJ mol⁻¹. For any strong acid + strong base, the net ionic equation is identical (H⁺ + OH⁻ → H₂O), so the value is always ~−57 kJ mol⁻¹.
Strong acid + strong base → equivalence at pH 7. The indicator must change colour across this pH. Bromothymol blue (6.0–7.6) brackets pH 7. Methyl orange (3.1–4.4) would change too early, before the true equivalence point.
NaOH, KOH, and Ca(OH)₂ are strong bases (group 1 and 2 hydroxides of low atomic weight) → dissociate completely. NH₃ is a weak base → only partially reacts with water to produce OH⁻.
pH 5.2 is acidic (optimal crop pH is 6.0–7.5). CaCO₃ neutralises excess H⁺ gradually. NaOH would over-neutralise and damage soil biology. Ammonium sulfate is acidic (NH₄⁺ hydrolyses to produce H⁺). Distilled water dilutes but does not remove H⁺, buffering by soil minerals prevents pH change.
The statement conflates acid strength with danger. Weak acids (citric, carbonic, acetic) are safe at typical concentrations. Stomach acid (HCl, pH 1.5–3.5) is essential for digestion and pathogen defence. Complete neutralisation would cause achlorhydria → impaired digestion and infection risk.
• 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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