Chemistry • Year 12 • Module 6 • Lesson 6
Strong/Weak Mastery — Consolidation
Apply strong/weak classification rules to real-world solutions, interpret pH and conductivity data, and reason about cause-and-effect relationships in acid/base chemistry.
1. Interpret pH and conductivity data for six solutions
A student prepared six aqueous solutions at identical concentrations (0.10 mol/L) and measured the pH and electrical conductivity of each at 25°C. The conductivity was measured using a calibrated conductivity meter and is reported in millisiemens per centimetre (mS/cm) as a proxy for total dissolved ion concentration. Results are shown in Table 1.
| Solution | Substance | Formula | pH | Conductivity (mS/cm) |
|---|---|---|---|---|
| A | Hydrochloric acid | HCl | 1.0 | 39.1 |
| B | Acetic acid (ethanoic acid) | CH₃COOH | 2.9 | 0.53 |
| C | Nitrous acid | HNO₂ | 2.1 | 3.8 |
| D | Sodium hydroxide | NaOH | 13.0 | 21.8 |
| E | Ammonia solution | NH₃(aq) | 11.1 | 0.46 |
| F | Sodium chloride | NaCl | 7.0 | 11.4 |
1.1 Using only the conductivity data, rank Solutions A–E from highest to lowest degree of ionisation and justify your ranking. Explain why NaCl (Solution F) cannot be included in the same comparison. 4 marks
1.2 Classify each of Solutions A–E as strong acid, weak acid, strong base, or weak base. Provide the correct ionic equation for the ionisation of Solutions A, B, and E. 5 marks
1.3 A student claims: “Solution C has a lower pH than Solution E, so nitrous acid must be a stronger acid than ammonia is a base.” Evaluate this reasoning. 3 marks
1.4 Using the conductivity data, calculate the approximate degree of ionisation of HNO₂ (Solution C) and CH₃COOH (Solution B). For a rough estimate, assume each mole of ionised acid produces approximately one mole of H⁺ and one mole of anion per litre, and that conductivity scales proportionally with the conductivity of Solution A. 3 marks
2. Interpreting a pH vs concentration graph
The graph below shows the calculated pH of aqueous solutions of hydrochloric acid (HCl) and acetic acid (CH₃COOH) at a range of molar concentrations at 25°C. The x-axis uses a logarithmic scale.
2.1 Describe the difference in the shapes of the two lines. Explain what this difference reveals about how the degree of ionisation changes with concentration for each acid. 3 marks
2.2 Estimate from the graph the pH of each acid at a concentration of 0.10 mol/L. Use these values to calculate the approximate degree of ionisation of CH₃COOH at this concentration. Show your working. 3 marks
2.3 At approximately what concentration would a CH₃COOH solution have the same pH as a 0.001 mol/L HCl solution? Use the graph to estimate, then explain why the concentration of CH₃COOH needed is much lower than 0.001 mol/L. 3 marks
3. Compare real-world solutions
Six everyday substances are listed below. Complete the comparison table by classifying each and providing the information required. 12 marks
| Substance | Approx. pH | Strong or weak acid/base, or salt solution? | Ionisation arrow (→ or ⇆) | Main species in solution (1–2 ions/molecules) |
|---|---|---|---|---|
| Blood (bicarbonate buffer) | 7.4 | |||
| Vinegar (5% acetic acid) | 2.4–3.4 | |||
| Household ammonia (NH₃ solution) | ~11 | |||
| Car battery acid (H₂SO₄, conc.) | <1 | |||
| Baking soda solution (NaHCO₃) | ~8.3 | |||
| Seawater | ~8.1 |
3.1 A student says: “Battery acid has a very low pH and is therefore a much stronger acid than acetic acid.” Using specific values from the table and your knowledge of acid strength, evaluate this claim. 3 marks
4. Predict and justify
Read the scenario, then answer the question. 4 marks
Scenario: A chemistry teacher sets up two conductivity demonstrations at the front of the class. Each demonstration uses a simple circuit with a light bulb, two electrodes, and a solution. In Demonstration 1 the bulb glows brightly. In Demonstration 2 the bulb barely glows. The teacher then tells students that both solutions contain exactly 0.10 mol/L of an acid, and that the two acids have the same molar mass.
4.1 Predict which demonstration contains the strong acid and which contains the weak acid. Justify your prediction by explaining the relationship between conductivity, ion concentration, and degree of ionisation. Then predict which solution has the lower pH and explain whether pH alone would be sufficient to identify which acid is strong. 4 marks
Q1.1 — Conductivity ranking
Higher conductivity = higher ion concentration = higher degree of ionisation at the same molar concentration. Ranking (highest → lowest degree of ionisation): A (HCl, 39.1) → D (NaOH, 21.8) → C (HNO₂, 3.8) → B (CH₃COOH, 0.53) → E (NH₃, 0.46). NaCl cannot be included because it is not an acid or base that ionises; it is a salt that dissociates fully into Na⁺ and Cl⁻ regardless of any acid/base strength — comparing it to degree of ionisation of an acid/base would conflate two different processes. Marking notes: 2 marks for correct ranking with justification; 1 mark for explaining the NaCl exclusion.
Q1.2 — Classification and ionic equations
A (HCl): strong acid. B (CH₃COOH): weak acid. C (HNO₂): weak acid. D (NaOH): strong base. E (NH₃): weak base.
A: HCl(aq) → H⁺(aq) + Cl⁻(aq) [single forward arrow].
B: CH₃COOH(aq) ⇆ H⁺(aq) + CH₃COO⁻(aq) [equilibrium arrow].
E: NH₃(aq) + H₂O(l) ⇆ NH₄⁺(aq) + OH⁻(aq) [equilibrium arrow].
Marking notes: 1 mark for all five classifications correct; 1 mark for each correct ionic equation including the correct arrow type (3 marks).
Q1.3 — Evaluating the student’s claim
The reasoning is flawed. Both HNO₂ and NH₃ are weak — they are on different sides of the acid/base divide (one is an acid, one is a base), so comparing their Ka and Kb values against each other using pH is not valid. Furthermore, pH reflects [H⁺] in an acidic solution but [H⁺] in a basic solution is derived from Kw, not from the base’s Kb directly. A lower pH for Solution C relative to the pH above 7 of Solution E tells us only that Solution C is acidic and E is basic — it says nothing about which is intrinsically “stronger” in a meaningful cross-class comparison. Marking notes: 1 mark for identifying the claim as invalid; 1 mark for explaining that pH of an acidic solution cannot be directly compared to pH of a basic solution to rank Ka vs Kb; 1 mark for noting that different type (acid vs base) means different equilibrium constants are being compared.
Q1.4 — Approximate degree of ionisation
Assume conductivity scales proportionally with ion concentration. Solution A (HCl, fully ionised at 0.10 mol/L) gives 39.1 mS/cm. This represents 100% ionisation.
HNO₂ (Solution C): fraction = 3.8 / 39.1 ≈ 9.7% degree of ionisation.
CH₃COOH (Solution B): fraction = 0.53 / 39.1 ≈ 1.4% degree of ionisation.
Marking notes: 1 mark for the proportional method; 1 mark each for the two correct answers (accept ±1%).
Q2.1 — Graph shape
HCl shows a straight line with a constant gradient: each 10-fold increase in concentration decreases pH by exactly 1 unit, consistent with [H⁺] = c (100% ionised; degree of ionisation is constant at 100% regardless of concentration). CH₃COOH shows a curved line that flattens at higher concentrations: the pH drops more slowly per unit increase in concentration, because the degree of ionisation actually increases slightly as concentration decreases (the equilibrium shifts right on dilution, a consequence of Le Chatelier’s principle). Marking notes: 1 mark for correctly describing HCl as linear; 1 mark for describing CH₃COOH curve shape; 1 mark for explaining the link between curve shape and degree of ionisation.
Q2.2 — pH read-off and degree of ionisation calculation
From the graph at 0.10 mol/L: HCl pH ≈ 1.0; CH₃COOH pH ≈ 2.9.
[H⁺] for CH₃COOH = 10⁻2.9 ≈ 1.26 × 10⁻³ mol/L.
Degree of ionisation = (1.26 × 10⁻³ / 0.10) × 100% ≈ 1.3%. Marking notes: 1 mark for correct pH read-off for each (accept ±0.1); 1 mark for correct calculation showing working; total 3 marks.
Q2.3 — Estimating equivalent pH concentration
The 0.001 mol/L HCl solution has pH = 3.0. From the graph, CH₃COOH reaches approximately pH 3.9 at 0.001 mol/L — so to reach pH 3.0, acetic acid would need a concentration well below 0.001 mol/L (read from the graph, approximately <0.0001 mol/L is needed). The reason is that CH₃COOH only produces a small fraction of [H⁺] relative to c; to achieve the same [H⁺] as 0.001 mol/L HCl, CH₃COOH actually needs an even lower starting concentration combined with a relatively higher fractional ionisation at that dilution — this is counterintuitive and illustrates the interaction between Ka, c, and the equilibrium shift on dilution. Marking notes: 1 mark for correct estimation from graph; 1 mark for the direction of the effect (acetic acid needs lower, not higher, concentration for same pH at very dilute range); 1 mark for referencing equilibrium shift on dilution. Accept well-reasoned alternative estimates.
Q3 — Real-world solutions table
Blood: pH 7.4; basic salt solution / buffer (bicarbonate, HCO₃⁻, conjugate of weak acid H₂CO₃); ⇆; main species HCO₃⁻, H₂CO₃, CO₂.
Vinegar: pH 2.4–3.4; weak acid (CH₃COOH); ⇆; mainly CH₃COOH molecules with small amounts of CH₃COO⁻ and H⁺.
Household ammonia: pH ~11; weak base (NH₃); ⇆; mainly NH₃ molecules with small amounts of NH₄⁺ and OH⁻.
Battery acid (H₂SO₄ conc.): pH <1; strong acid (first ionisation complete); →; HSO₄⁻, H⁺, SO₄²– (second ionisation is moderately strong).
Baking soda (NaHCO₃): pH ~8.3; basic salt solution (conjugate of weak acid H₂CO₃); HCO₃⁻ hydrolyses ⇆; Na⁺, HCO₃⁻, small OH⁻.
Seawater: pH ~8.1; slightly basic (dissolved CO₂/bicarbonate equilibrium, Mg²⁺ / Ca²⁺ salts); ⇆ (equilibrium); Na⁺, Cl⁻, HCO₃⁻ dominant.
Q3.1 — Battery acid claim evaluation
The claim is partially correct but the reasoning is incomplete. H₂SO₄ is indeed a strong acid (first ionisation complete: H₂SO₄ → H⁺ + HSO₄⁻) and acetic acid is a weak acid. However, the very low pH of battery acid reflects both its strength (Ka effectively infinite for the first ionisation) and its very high concentration in car batteries (~10 mol/L vs ~0.9 mol/L for 5% vinegar). At the same concentration, H₂SO₄ would produce far more H⁺ than CH₃COOH, confirming it is stronger. But attributing the low pH to strength alone without acknowledging the much higher concentration is an incomplete argument. Marking notes: 1 mark for correctly classifying both as strong and weak; 1 mark for noting that concentration also contributes to the low pH; 1 mark for stating that at the same concentration H₂SO₄ would still be far more acidic.
Q4.1 — Predict and justify
Demonstration 1 (bright bulb) contains the strong acid; Demonstration 2 (dim bulb) contains the weak acid. Conductivity depends on total ion concentration. A strong acid at 0.10 mol/L produces 0.10 mol/L each of H⁺ and its anion (total ≈0.20 mol/L ions). A weak acid at 0.10 mol/L produces far fewer ions (only the small fraction that has ionised), so the conductivity — and hence bulb brightness — is much lower. The strong acid (Demonstration 1) has the lower pH because [H⁺] = c (full ionisation), whereas the weak acid produces [H⁺] << c. pH alone would not identify which is strong if the solutions had different concentrations; however, at the same concentration, the lower pH definitively belongs to the strong acid. Marking notes: 1 mark for correct identification of both demonstrations; 1 mark for linking conductivity to ion concentration and degree of ionisation; 1 mark for identifying which has lower pH with justification; 1 mark for the caveat about same concentration.