Chemistry • Year 12 • Module 6 • Lesson 5
Strong vs Weak Acids & Bases: The Critical Distinction
Build Band 5–6 extended-response technique: synthesise data, evaluate flawed sources, and construct an evidence-based judgement about acid strength and its consequences.
1. Data-based extended response — comparing 0.1 mol/L HCl and 0.1 mol/L CH₃COOH (Band 5–6)
8 marks Band 5–6
Scenario. A NESA-prescribed investigation requires students to compare properties of equal-concentration solutions of a strong acid and a weak acid. The table below summarises measurements obtained for 0.1 mol/L HCl and 0.1 mol/L CH₃COOH at 25°C.
| Property | 0.1 mol/L HCl | 0.1 mol/L CH₃COOH |
|---|---|---|
| pH | 1.0 | 2.9 |
| [H⁺] (mol/L) | 0.10 | ≈0.0013 |
| Conductivity (mS/cm) | 42.1 | 0.53 |
| Degree of ionisation (%) | ~100% | ~1.3% |
| Reaction with Mg ribbon | Vigorous H₂ evolution | Slow H₂ evolution |
| Arrow in ionic equation | → | ⇌ |
Data collected at 25°C using calibrated pH meter, conductivity probe, and visual observation. CH₃COOH sourced from glacial acetic acid (AHA food acid 260), HCl sourced from analytical-grade 32% solution.
Q1. Analyse and evaluate the data to explain the differences between 0.1 mol/L HCl and 0.1 mol/L CH₃COOH. In your response you must:
- Define acid strength and distinguish it clearly from concentration.
- Account for the difference in pH using the concept of degree of ionisation (reference specific data values).
- Explain why the conductivity of HCl is approximately 80 times greater than that of CH₃COOH despite equal concentration.
- Connect the correct arrow notation for each acid to the physical reality of ionisation.
- Reach an evidence-based conclusion about which property — strength or concentration — determines the pH of these two solutions, and why pH alone cannot identify which acid is strong.
2. Source critique — evaluate a student’s exam response (Band 5–6)
7 marks Band 5–6
Source — Year 12 student’s exam response (from a practice paper, 2025):
“In the experiment I compared 0.1 mol/L HCl and 0.1 mol/L CH₃COOH. The HCl had a lower pH of 1.0 because it is more concentrated. The CH₃COOH had a higher pH of 2.9 because it is a weak acid, meaning it is less concentrated. This shows that weak acids are dilute and strong acids are concentrated. The ionic equation for CH₃COOH can be written with either arrow — both → and ⇌ are acceptable depending on how much acid is present. The conductivity of HCl was much higher, which makes sense because HCl is more reactive and dangerous than acetic acid. In real life, acetic acid in vinegar is safe because it is a weak acid.”
Q2. This student’s response contains multiple significant scientific errors. Identify and correct each error, and explain in each case what the correct chemistry is and how the error could be detected experimentally. Your response must:
- Identify all errors (there are at least four distinct scientific errors).
- State the correct chemistry for each error.
- For at least two errors, describe an experimental observation that would contradict the student’s claim.
- Conclude by stating the correct relationship between acid strength, concentration, [H⁺], and pH.
Q1 — Marking criteria (8 marks)
1 mark — Define acid strength / distinguish from concentration: Acid strength (quantified by Ka) is the intrinsic tendency of an acid to donate a proton to water, measured by the fraction of molecules that ionise at equilibrium. Concentration (mol/L) is the total amount of acid dissolved per litre, independent of Ka. These two properties are completely independent of each other.
2 marks — Account for pH difference using degree of ionisation and data: HCl is a strong acid and ionises completely (degree of ionisation ~100%): HCl(aq) → H⁺(aq) + Cl⁻(aq), so [H⁺] = 0.10 mol/L and pH = −log(0.10) = 1.0. CH₃COOH is a weak acid and ionises only ~1.3% at this concentration: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq), so [H⁺] ≈ 0.0013 mol/L and pH ≈ 2.9. The pH difference of 1.9 units means [H⁺] in HCl is approximately 79 times greater — all because of the strength distinction, not a concentration difference. (1 mark for correct HCl calculation; 1 mark for CH₃COOH reasoning with reference to data.)
2 marks — Explain conductivity difference: Conductivity measures ion concentration. HCl ionises completely, so essentially all HCl molecules become H⁺ and Cl⁻ ions — giving 0.10 mol/L of each ion at 0.10 mol/L HCl. CH₃COOH ionises only ~1.3%, so the vast majority of molecules remain as intact CH₃COOH. The ion concentration is ~79× lower, explaining why conductivity (42.1 vs 0.53 mS/cm) is ~80× lower. (1 mark for linking complete vs partial ionisation to ion count; 1 mark for quantitative comparison with data values.)
1 mark — Connect arrow notation to physical reality: HCl uses → (single arrow) because ionisation is complete and irreversible — essentially no Cl⁻ accepts H⁺ back. CH₃COOH uses ⇌ (equilibrium arrow) because the reaction is reversible and partial — forward ionisation and reverse recombination both occur at equilibrium.
2 marks — Evaluative conclusion about pH and strength: At equal concentration, the difference in pH is caused by the difference in strength (Ka), not concentration. However, pH alone cannot identify acid strength: a concentrated weak acid (e.g. 1 mol/L CH₃COOH, pH ~2.4) can have a lower pH than a dilute strong acid (e.g. 0.001 mol/L HCl, pH 3.0). To determine whether an acid is strong or weak, one must compare pH at equal concentration or measure conductivity, or know the Ka. (1 mark for correct evaluative claim; 1 mark for using a counter-example to show pH alone is insufficient.)
Q2 — Marking criteria (7 marks)
Error 1 (1.5 marks): “HCl has lower pH because it is more concentrated.”
Correct chemistry: Both solutions are at the same concentration (0.10 mol/L). HCl has a lower pH because it is a strong acid that ionises completely, producing [H⁺] = 0.10 mol/L. CH₃COOH has a higher pH because it is a weak acid that ionises only ~1.3%, producing [H⁺] ≈ 0.0013 mol/L.
Experimental contradiction: Measuring the concentration of both solutions (e.g. by titration or by weighing mass dissolved) would confirm both are 0.10 mol/L.
Error 2 (1.5 marks): “Weak acids are dilute; strong acids are concentrated.”
Correct chemistry: Strength and concentration are completely independent. You can have a concentrated weak acid (e.g. glacial CH₃COOH at ~17 mol/L — highly concentrated but still a weak acid with low degree of ionisation) or a dilute strong acid (0.001 mol/L HCl — very dilute but 100% ionised). The existence of commercial glacial acetic acid contradicts this claim directly.
Error 3 (1.5 marks): “Both → and ⇌ are acceptable for CH₃COOH depending on how much acid is present.”
Correct chemistry: Arrow notation reflects the intrinsic degree of ionisation (Ka), not the concentration. CH₃COOH always uses ⇌ because it is always a weak acid with Ka = 1.8 × 10⁻⁵ regardless of concentration. Using → for CH₃COOH is always wrong and will lose marks in every ionic equation context.
Experimental contradiction: Measuring conductivity of CH₃COOH at any concentration from 0.001 to 10 mol/L would show conductivity far below the equivalent strong acid at the same concentration, confirming partial ionisation at all concentrations.
Error 4 (1.5 marks): “Acetic acid in vinegar is safe because it is a weak acid.”
Correct chemistry: “Weak” refers strictly to the degree of ionisation — not to concentration, hazard, or biological danger. Glacial acetic acid (~17 mol/L, a concentrated weak acid) is classified as a corrosive dangerous good and causes severe burns. The safety of vinegar (~0.83 mol/L, ~5% by mass) arises from its low concentration, not from it being a weak acid. HF is another example: a weak acid (Ka = 6.8 × 10⁻⁴) that is acutely toxic and can cause fatal systemic effects from skin contact alone.
1 mark — Correct concluding statement: Acid strength (Ka) is an intrinsic property of the acid molecule at a given temperature, independent of concentration. [H⁺] depends on both strength (Ka) and concentration. pH is then determined by [H⁺]. Therefore: same concentration + different Ka = different pH. Same Ka + different concentration = different pH. pH alone cannot tell you which factor caused the difference.