Four students just explained what makes an acid "weak." Only one of them is correct. Before reading on — can you identify who, and precisely what each of the others got wrong?
This lesson introduces no new dot points. Its purpose is to deepen understanding of the strong/weak distinction through analogies, harder worked examples, salt hydrolysis prediction, and explicit misconception resolution. By the end, you should be able to diagnose and fix the four highest-frequency errors in HSC Module 6 without prompting, and write a Band 6 response distinguishing strength from concentration under exam conditions.
Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
Four students were asked: "A solution of 0.1 mol/L hydrochloric acid and a solution of 0.1 mol/L acetic acid are prepared. Both have the same concentration. Explain why they have different pH values, and use this to define the difference between a strong and a weak acid."
Student A: "HCl has a lower pH because it is a stronger acid — it ionises completely, giving [H⁺] = 0.1 mol/L and pH = 1.0. CH₃COOH only partially ionises, giving [H⁺] << 0.1 mol/L and a higher pH. A strong acid is one that ionises completely in water; a weak acid is one that only partially ionises. Strength is about the degree of ionisation, not the concentration."
Student B: "HCl has a lower pH because it is more concentrated than the acetic acid solution. When you make HCl more concentrated it becomes a stronger acid, so it has more H⁺ ions and a lower pH."
Student C: "HCl has a lower pH because it is a strong acid — it fully dissociates. But acetic acid is weak because it is dilute. If you made the acetic acid more concentrated it would become a strong acid too, because there would be more molecules to ionise."
Student D: "HCl has a lower pH than acetic acid at the same concentration. This shows that HCl is a stronger acid. A weak acid like acetic acid barely ionises at all, so it is barely acidic — you could almost drink it safely because the pH is so high."
Before reading on: Which student is correct? Write a precise identification of the specific error each incorrect student has made. You will return to this analysis at the end of the lesson.
📚 Core Content
The four students represent the four most common ways Year 12 students misunderstand the strong/weak distinction — and identifying exactly what each one got wrong, rather than simply knowing who is right, is what builds the precision needed for Band 6 responses.
Student A is correct. The explanation is complete, accurate, and uses the right language at every step. Strong acid = complete ionisation → [H⁺] = concentration. Weak acid = partial ionisation → [H⁺] << concentration. Strength is degree of ionisation — independent of concentration. This is the full, HSC-quality answer.
Student B says HCl has a lower pH "because it is more concentrated" and that increasing concentration makes an acid stronger. Both claims are wrong. The two solutions in the problem are at the same concentration (0.1 mol/L) — concentration is controlled. HCl's lower pH is entirely due to its greater degree of ionisation (100% vs ~1.3%), not any difference in concentration. Acid strength (Ka) is an intrinsic property of the molecule — it does not change when you increase concentration. 12 mol/L HCl and 0.001 mol/L HCl are both strong acids.
✓ Fix: Concentration affects pH; it does not affect Ka. "Dilute" and "weak" are not synonyms. Always use both descriptors separately: "dilute strong acid," "concentrated weak acid."
This follows from the same confusion as Student B but is even more explicit. Acetic acid has Ka = 1.8 × 10⁻⁵ at 25°C regardless of concentration. Ka describes the intrinsic proton-donating tendency of the CH₃COOH molecule — a property of its molecular structure and bond energies, not of how many molecules are present per litre. At 10 mol/L, acetic acid is still a weak acid — a higher fraction of molecules are still intact than ionised.
✓ Fix: Ka is fixed at a given temperature. The only thing that changes Ka is temperature — not dilution, not adding more solute. "Becoming strong" requires a change in Ka, which does not happen by changing concentration.
Student D conflates ionisation fraction (strength) with safety or absolute acidity level. This is dangerous. Glacial acetic acid (pure CH₃COOH, ~17 mol/L) is a classified corrosive dangerous good that causes chemical burns on contact. Hydrofluoric acid (HF), another weak acid (Ka = 6.8 × 10⁻⁴), is one of the most hazardous laboratory acids — F⁻ ions penetrate tissue and cause systemic hypocalcaemia including cardiac arrest from skin contact alone. "Weak acid" means partial ionisation — it says nothing about safety, concentration, or absolute [H⁺].
✓ Fix: "Weak" refers only to the fraction of molecules that ionise. A concentrated weak acid can be highly acidic and highly dangerous. Never equate acid strength classification with safety classification.
Wrong: Acids always have a pH below 7 and bases always above 7.
Right: pH depends on temperature; at high temperatures, neutral pH is less than 7 due to increased Kw.
The most effective way to hold all three variables — strength, concentration, and [H⁺] — in your mind simultaneously is to map them onto a physical scenario where each variable has an obvious, distinct meaning before any chemistry is applied.
Imagine a concert venue with a door leading from the car park (reactant side) into the arena (product side).
A strong acid: every single person rushes through immediately — the car park empties completely. [H⁺] in arena = total crowd. A weak acid: only a small fraction trickle through — most stay in the car park. [H⁺] in arena << total crowd.
The key insight: Scenario 3 (large crowd, 1% rush) can have more arena occupants than Scenario 2 (small crowd, 100% rush) — a concentrated weak acid can be more acidic than a dilute strong acid. This is the critical fact that Student D missed.
The strength/concentration distinction applies equally to bases, but bases introduce an additional complication — solubility — that must be kept separate from both strength and concentration.
Scenario A (NaOH): Dissolves readily (high solubility). Every dissolved formula unit gives Na⁺ + OH⁻ completely. High solubility + strong → high [OH⁻].
Scenario B (Ca(OH)₂): Most does not dissolve — white suspension at the bottom. But every formula unit that DOES dissolve gives Ca²⁺ + 2OH⁻ completely. Low solubility + strong → moderate [OH⁻] (despite being a strong base).
Scenario C (NH₃): Dissolves readily (high solubility). But only a tiny fraction of the dissolved NH₃ accepts a proton from water. Most NH₃ remains intact. High solubility + weak → low [OH⁻].
| Base | Solubility | Strength (dissociation of dissolved fraction) | [OH⁻] result | Correct description |
|---|---|---|---|---|
| NaOH | High | Strong (100%) | High | Soluble strong base |
| Ca(OH)₂ | Low (~0.02 mol/L) | Strong (100%) | Low-moderate | Sparingly soluble strong base |
| NH₃ | High | Weak (~1% at 0.1 mol/L) | Low | Soluble weak base |
| Mg(OH)₂ | Very low | Weak (incomplete dissociation) | Very low | Sparingly soluble weak base |
Every salt dissolves to produce ions — and whether those ions react with water to shift pH depends entirely on whether they are conjugates of strong or weak acids and bases, a fact that directly connects the strong/weak distinction from L05 to a new predictive tool.
The salt hydrolysis rule:
| Salt | Cation source | Anion source | Cation behaviour | Anion behaviour | Solution pH |
|---|---|---|---|---|---|
| NaCl | NaOH (strong) | HCl (strong) | Neutral spectator | Neutral spectator | ~7 (neutral) |
| CH₃COONa | NaOH (strong) | CH₃COOH (weak) | Neutral spectator | Basic — accepts H⁺ from water | > 7 (basic) |
| NH₄Cl | NH₃ (weak) | HCl (strong) | Acidic — donates H⁺ to water | Neutral spectator | < 7 (acidic) |
| Na₂CO₃ | NaOH (strong) | H₂CO₃ (weak) | Neutral spectator | Basic — accepts H⁺ | > 7 (strongly basic) |
| NH₄NO₃ | NH₃ (weak) | HNO₃ (strong) | Acidic — donates H⁺ | Neutral spectator | < 7 (acidic) |
Three-step process: (1) Identify the acid and base that formed the salt. (2) Check whether each is strong or weak. (3) Apply the rule. For example, CH₃COONa: Na⁺ from NaOH (strong base) = neutral spectator. CH₃COO⁻ from CH₃COOH (weak acid) = basic ion → CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻ → pH > 7.
The errors in this lesson are so consistent across HSC cohorts that they have become predictable — and diagnosing them explicitly, rather than just presenting the correct answer, is the fastest way to eliminate them from your own responses.
What the student writes: HCl ⇌ H⁺ + Cl⁻ (or HNO₃ ⇌ H⁺ + NO₃⁻, etc.)
Why it is wrong: ⇌ communicates that a significant reverse reaction occurs — that Cl⁻ meaningfully accepts H⁺ back from H₃O⁺. This is chemically false. Cl⁻ is the conjugate base of a strong acid — it has essentially no tendency to accept H⁺. The equation implies partial ionisation. Every mark for ionic equations in Module 6 includes an arrow check.
✓ Fix: Memorise the six strong acids. Apply → automatically. If uncertain whether an acid is strong, assume weak (⇌) — but for the six listed, → is non-negotiable.
What the student writes: "We used dilute HCl, which is a weak acid."
Why it is wrong: HCl at any concentration is a strong acid. Strength is Ka — an intrinsic property of the molecule. Diluting HCl changes concentration; it does not change Ka or degree of ionisation.
✓ Fix: Always use both descriptors separately. "Dilute" and "weak" describe different axes — never use them interchangeably. Correct phrasing: "dilute strong acid" (low c, complete ionisation) vs "concentrated weak acid" (high c, partial ionisation).
What the student writes: "0.1 mol/L CH₃COOH → [H⁺] = 0.1 mol/L → pH = 1.0"
Why it is wrong: This treats a weak acid as if it were strong. CH₃COOH is only ~1.3% ionised at 0.1 mol/L — [H⁺] ≈ 0.0013 mol/L, pH ≈ 2.9. Using [H⁺] = c overcalculates [H⁺] by a factor of ~77 and gives a pH far too low (1.0 vs 2.9). This error also typically coincides with using → instead of ⇌.
✓ Fix: Always classify the acid as strong or weak FIRST. If strong → [H⁺] = c. If weak → use Ka and ICE table (covered in L09). Never use [H⁺] = c for a weak acid under any circumstances.
What the student writes: "The strong base NaOH was used — weak bases include KOH, Ca(OH)₂, and Ba(OH)₂."
Why it is wrong: KOH, Ca(OH)₂, and Ba(OH)₂ are all strong bases — they dissociate completely. Only NH₃, Mg(OH)₂, and organic amines are weak.
✓ Fix: Memorise the complete strong base list: NaOH, KOH, Ca(OH)₂, Ba(OH)₂. Any base not on this list encountered in HSC is weak.
📝 Worked Examples
Problem: For each substance: (a) classify as strong or weak acid/base; (b) write the correct ionic equation; (c) where a salt is formed, predict whether its aqueous solution is acidic, basic, or neutral with explanation. (i) HNO₃ in water; (ii) HNO₂ in water; (iii) NH₄NO₃ dissolving in water.
HNO₃: On the strong acid list → strong acid. Single arrow. Not a salt — no hydrolysis prediction needed.
HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
HNO₂: NOT on the strong acid list → weak acid. Equilibrium arrow.
HNO₂(aq) ⇌ H⁺(aq) + NO₂⁻(aq)
NH₄NO₃ salt: Dissolution: NH₄NO₃(aq) → NH₄⁺(aq) + NO₃⁻(aq) (ionic salt fully dissociates).
NO₃⁻ is the conjugate of HNO₃ (strong acid) → neutral spectator, no hydrolysis.
NH₄⁺ is the conjugate of NH₃ (weak base) → acidic ion: NH₄⁺(aq) ⇌ H⁺(aq) + NH₃(aq) — donates H⁺ to water, lowers pH.
NH₄NO₃ solution is acidic (pH < 7). This explains why ammonium nitrate fertiliser gradually acidifies soil — consistent with its farming use alongside lime application.
Answers: (i) Strong acid — HNO₃(aq) → H⁺(aq) + NO₃⁻(aq). (ii) Weak acid — HNO₂(aq) ⇌ H⁺(aq) + NO₂⁻(aq). (iii) NH₄NO₃ → NH₄⁺ + NO₃⁻; acidic solution — NH₄⁺ (conjugate of weak base NH₃) donates H⁺ to water; NO₃⁻ (conjugate of strong acid HNO₃) is neutral spectator.
Problem: A student claims that a 5.0 mol/L solution of acetic acid (Ka = 1.8 × 10⁻⁵) is a stronger acid than a 0.001 mol/L solution of HCl because "the acetic acid solution is more acidic — it has a lower pH." (a) Calculate [H⁺] and pH for each solution. Use the approximation [H⁺] ≈ √(Ka × c) for the weak acid. (b) Evaluate the student's claim.
HCl (strong): [H⁺] = concentration = 0.001 mol/L. pH = −log(0.001) = 3.0.
CH₃COOH (weak): [H⁺] ≈ √(Ka × c) = √(1.8 × 10⁻⁵ × 5.0) = √(9.0 × 10⁻⁵) = 9.49 × 10⁻³ mol/L.
Degree of ionisation check: (9.49 × 10⁻³ / 5.0) × 100% = 0.19% << 5% → approximation valid.
pH = −log(9.49 × 10⁻³) = 2.02.
The student is correct that the 5.0 mol/L CH₃COOH solution has a lower pH (2.02 vs 3.0). In terms of absolute [H⁺] in these specific solutions, the acetic acid solution is more acidic.
However, the student's reasoning is incorrect. A lower pH does not indicate a stronger acid — it reflects the effect of a 5000× higher concentration. The intrinsic acid strength (Ka) defines strong vs weak: HCl has Ka effectively → ∞; CH₃COOH has Ka = 1.8 × 10⁻⁵. At the same concentration, HCl always produces far higher [H⁺] — it is the stronger acid.
At 5.0 mol/L, only 0.19% of CH₃COOH has ionised — it is still a weak acid. At 5.0 mol/L, HCl would give [H⁺] = 5.0 mol/L (pH = −0.70) — far more acidic than the acetic acid at the same concentration.
Answer: (a) HCl: [H⁺] = 0.001, pH = 3.0. CH₃COOH: [H⁺] ≈ 9.49 × 10⁻³, pH ≈ 2.02. (b) The lower pH of CH₃COOH is caused by 5000× higher concentration, not greater Ka. At equal concentrations, HCl always gives a lower pH — it is unambiguously the stronger acid. Strength (Ka) and [H⁺] in a specific solution are different quantities and cannot be compared using a single pH value at different concentrations.
Problem: "The terms 'strong', 'weak', 'concentrated', and 'dilute' are often used interchangeably when describing acid solutions, but each describes a fundamentally different property." (a) Define each of the four terms precisely. (b) Explain how two of the four terms can be combined to describe any acid solution. (c) Use quantitative examples to demonstrate that a concentrated weak acid can have a lower pH than a dilute strong acid. (d) Explain what experimental measurement would allow you to distinguish a strong acid from a weak acid without relying on pH alone.
Four precise definitions:
Strong acid: ionises completely in aqueous solution — Ka → ∞ (effectively); fraction ionised ≈ 100%. → in ionic equation.
Weak acid: ionises only partially in aqueous solution — Ka << 1 (typically 10⁻² to 10⁻¹⁰); fraction ionised << 100%. ⇌ in ionic equation.
Concentrated: high total amount of acid dissolved per litre — high c (mol/L), regardless of how much has ionised.
Dilute: low total amount of acid dissolved per litre — low c, regardless of degree of ionisation.
Strength (Ka) and concentration (c) are independent axes — they do not affect each other.
Four combinations: Any acid solution requires one descriptor from each axis. A dilute strong acid (low c, complete ionisation — e.g. 0.001 mol/L HCl). A concentrated strong acid (high c, complete ionisation — e.g. 12 mol/L HCl). A dilute weak acid (low c, partial ionisation — e.g. 0.001 mol/L CH₃COOH). A concentrated weak acid (high c, partial ionisation — e.g. 10 mol/L CH₃COOH). Using only one descriptor without specifying concentration or strength is incomplete.
Quantitative example: Compare 10 mol/L CH₃COOH (concentrated weak acid) and 0.001 mol/L HCl (dilute strong acid).
HCl: [H⁺] = 0.001 mol/L → pH = 3.0.
CH₃COOH: [H⁺] ≈ √(1.8 × 10⁻⁵ × 10) = √(1.8 × 10⁻⁴) = 0.0134 mol/L → pH = −log(0.0134) = 1.87.
The concentrated weak acid (pH 1.87) is more acidic than the dilute strong acid (pH 3.0) — despite CH₃COOH being intrinsically much weaker. Only 0.134% of CH₃COOH molecules have ionised, yet the sheer number of molecules at 10 mol/L means the absolute [H⁺] produced exceeds that from 0.001 mol/L HCl.
Experimental distinction without pH:
Method 1 — Electrical conductivity: At the same concentration, a strong acid (completely ionised) produces far more ions per litre than a weak acid (partially ionised). Conductivity is proportional to total ion concentration. 0.1 mol/L HCl: ~0.2 mol/L ions. 0.1 mol/L CH₃COOH: ~0.003 mol/L ions. HCl conductivity is approximately 65 times higher at the same concentration — a measurable, pH-independent distinction.
Method 2 — Reaction rate with Mg: A strong acid reacts initially faster because it supplies more H⁺ ions immediately. A weak acid reacts more slowly initially (fewer H⁺ available), though the equilibrium shifts right to replenish H⁺ as it is consumed, so the reaction eventually approaches the same extent.
Summary: (a) Four independent descriptors — strong/weak = Ka (degree of ionisation); concentrated/dilute = c (mol/L). (b) Any acid solution = one term from each axis. (c) 10 mol/L CH₃COOH → pH 1.87 vs 0.001 mol/L HCl → pH 3.0 — concentrated weak acid is more acidic despite weaker Ka because high c overwhelms low ionisation fraction. (d) Conductivity at equal concentration (strong gives ~65× more ions); or initial reaction rate with Mg — both distinguish strong from weak without relying on pH.
🧪 Activities
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
🔢 Multiple Choice
1. A student writes the ionic equation for the dissolution of nitric acid as HNO₃(aq) ⇌ H⁺(aq) + NO₃⁻(aq). What is wrong with this equation, and what does the incorrect arrow imply?
2. A solution of sodium acetate (CH₃COONa) is dissolved in water. Which prediction about the pH of the resulting solution is correct?
3. Two solutions are prepared: Solution P is 0.001 mol/L HCl; Solution Q is 2.0 mol/L CH₃COOH (Ka = 1.8 × 10⁻⁵). A student uses a pH probe and finds that Solution Q has a lower pH than Solution P. The student concludes that CH₃COOH must be a stronger acid than HCl because it produces a more acidic solution. Which response correctly evaluates this conclusion?
4. Which of the following salts produces a basic (pH > 7) aqueous solution when dissolved in water?
5. A student wants to distinguish between 0.10 mol/L HNO₃ (strong acid) and 0.10 mol/L HNO₂ (weak acid) without measuring pH. Which experimental approach would provide the clearest evidence?
✍️ Short Answer
6. For each of the following salts, predict whether its aqueous solution is acidic, neutral, or basic. Justify each prediction by identifying the parent acid and base, classifying each as strong or weak, and applying the salt hydrolysis rule: (a) NH₄Br; (b) NaNO₃; (c) Na₂SO₃; (d) CH₃COONH₄. 4 MARKS
7. A student has two 0.10 mol/L solutions: Solution X contains HCl; Solution Y contains HNO₂ (Ka = 4.5 × 10⁻⁴). (a) Calculate the pH of Solution X and Solution Y. Use [H⁺] ≈ √(Ka × c) for the weak acid. (b) Using the calculated pH values, explain what the difference tells us about the degree of ionisation of HNO₂ at 0.10 mol/L. (c) Explain why the electrical conductivity of Solution X will be significantly higher than Solution Y at the same concentration, despite both solutions containing the same total amount of acid per litre. 5 MARKS
8. Extended Response — Consolidation: A student makes the following claim: "Acid strength and acid concentration are essentially the same thing — a concentrated acid is a strong acid, and a dilute acid is a weak acid." (a) Define acid strength and acid concentration using precise chemical language. (b) Use the concert crowd analogy to explain how [H⁺] depends on both properties independently. (c) Provide a specific quantitative example showing that a dilute strong acid can have a higher pH than a concentrated weak acid. (d) Describe how you would experimentally determine whether an unknown acid solution is strong or weak using two different methods that do not require pH measurement. 7 MARKS
Go back to your Think First analysis of the four students at the top of this lesson.
1. Error type 1 — wrong arrow. HBr is on the strong acid list → must use →. Using ⇌ implies partial ionisation and that Br⁻ meaningfully accepts H⁺ back — both false. Correct: HBr(aq) → H⁺(aq) + Br⁻(aq).
2. Error type 2 — dilute ≠ weak. HCl is a strong acid at any concentration. 0.001 mol/L HCl = dilute strong acid. The high pH (3.0) is caused entirely by low concentration — HCl at 0.001 mol/L is still 100% ionised. Correct: "The 0.001 mol/L solution is a dilute strong acid."
3. Error type 3 — [H⁺] = c for weak acid. Ethanoic acid (CH₃COOH) is a weak acid — only ~1.3% ionised at 0.10 mol/L. [H⁺] ≈ 0.0013 mol/L → pH ≈ 2.9. Must use Ka + ICE table. Correct: pH ≈ 2.9 (not 1.0).
4. Error type 4 — incomplete strong base list. KOH and Ca(OH)₂ are strong bases (fully dissociate). Only NH₃ is weak. Correct: Strong bases = NaOH, KOH, Ca(OH)₂, Ba(OH)₂. NH₃ is a weak base.
5. Salt hydrolysis error. HCl is strong; NH₃ is weak → NH₄⁺ is an acidic ion (conjugate of weak base). It donates H⁺ to water: NH₄⁺(aq) ⇌ H⁺(aq) + NH₃(aq). pH < 7. "Neutralisation produces neutral solutions" is only true when BOTH acid and base are strong.
1. Salt predictions: (a) Na₂SO₄: Na⁺ (NaOH, strong) + SO₄²⁻ (H₂SO₄, strong) → both neutral → pH ≈ 7. (b) NH₄Cl: NH₄⁺ (NH₃, weak base) → acidic ion; Cl⁻ (HCl, strong) → neutral → pH < 7 (acidic). (c) Na₂CO₃: Na⁺ neutral; CO₃²⁻ (H₂CO₃, weak acid) → basic ion → pH > 7 (strongly basic). (d) KNO₃: K⁺ (KOH, strong) + NO₃⁻ (HNO₃, strong) → both neutral → pH ≈ 7. (e) CH₃COONH₄: NH₄⁺ (acidic) and CH₃COO⁻ (basic) — Ka(NH₄⁺) ≈ Ka(CH₃COOH) — approximately cancel → pH ≈ 7.
2. Calculations: Solution A (0.50 mol/L HBr, strong): [H⁺] = 0.50 mol/L → pH = −log(0.50) = 0.30. Solution B (2.0 mol/L HF, weak): [H⁺] ≈ √(6.8 × 10⁻⁴ × 2.0) = √(1.36 × 10⁻³) = 0.0369 mol/L → pH = 1.43. Solution A has the lower pH (0.30) despite having a lower concentration (0.50 vs 2.0 mol/L) — because HBr is 100% ionised while HF is only 1.85% ionised at 2.0 mol/L.
3. Experimental methods: (1) Electrical conductivity at equal concentration: HBr (strong, 100% ionised) produces [H⁺] = [Br⁻] = c → ~2c mol/L ions total. HF (weak, ~1.85% ionised at 2.0 mol/L) produces only ~0.037 mol/L ions total. Conductivity of HBr will be far higher — measurable with a conductivity meter. (2) Rate of reaction with Mg ribbon: HBr supplies [H⁺] = 0.50 mol/L immediately; HF supplies [H⁺] ≈ 0.037 mol/L initially (equilibrium shifts right as H⁺ is consumed, but initial rate is lower). HBr will produce visible bubbles of H₂ far faster than HF at any comparable concentration.
1. B — HNO₃ is a strong acid; the correct arrow is →. Using ⇌ implies partial ionisation and a significant reverse reaction (NO₃⁻ accepting H⁺ back from H₃O⁺) — both chemically false. NO₃⁻ is the conjugate base of a strong acid; it has essentially no tendency to accept H⁺. Option A wrong — HNO₃ is definitively strong. Option C wrong — ⇌ implies equilibrium, not amphoteric behaviour. Option D conflates H⁺ and H₃O⁺ notation (both are acceptable in HSC) without addressing the arrow error.
2. C — CH₃COO⁻ is the conjugate base of acetic acid (weak acid); it has a meaningful tendency to accept H⁺ from water: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. This produces OH⁻ → pH > 7. Na⁺ is neutral (from NaOH, strong base). Option B wrong — neutralisation produces neutral solutions only when both acid and base are strong. Option A wrong — Na⁺ has no acidic proton to donate. Option D wrong — Na⁺ is neutral, not acidic.
3. B — The lower pH of Solution Q is caused by 2000× higher concentration, not greater Ka. HCl (Ka → ∞) is definitively stronger than CH₃COOH (Ka = 1.8 × 10⁻⁵). At equal concentrations, HCl always gives lower pH. pH cannot be used to compare strength unless concentrations are equal. At equal concentration (0.1 mol/L): HCl → pH 1.0; CH₃COOH → pH ≈ 2.9. HCl is stronger.
4. C — Na₂CO₃: CO₃²⁻ is the conjugate base of carbonic acid (H₂CO₃, weak acid) → basic ion. CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻ → pH > 7 (strongly basic). Option A (KNO₃): K⁺ (strong base) + NO₃⁻ (strong acid) → both neutral → pH ≈ 7. Option B (NH₄Cl): NH₄⁺ (weak base conjugate) → acidic → pH < 7. Option D (NH₄NO₃): NH₄⁺ acidic, NO₃⁻ neutral → pH < 7.
5. A — At equal concentration, electrical conductivity is proportional to total ion concentration. HNO₃ (strong, 100% ionised) → [H⁺] = [NO₃⁻] = 0.10 mol/L → ~0.20 mol/L ions. HNO₂ (weak, ~6.7% ionised at 0.10 mol/L) → ~0.013 mol/L ions. HNO₃ conductivity is ~15× higher — clearly distinguishable without pH. Option B: universal indicator uses colour related to pH — this is still pH-based. Option C: density is not significantly different between dilute acid solutions at these concentrations. Option D: both acids react with NaOH in the same mole ratio (1:1) — gas is not produced in acid-base neutralisation.
Q6 (4 marks): (a) NH₄Br: NH₄⁺ from NH₃ (weak base) is an acidic ion — donates H⁺ to water: NH₄⁺ ⇌ H⁺ + NH₃. Br⁻ from HBr (strong acid) is neutral. Solution is acidic (pH < 7) [1]. (b) NaNO₃: Na⁺ from NaOH (strong base) neutral. NO₃⁻ from HNO₃ (strong acid) neutral. Neither ion hydrolyses. Solution is neutral (pH ≈ 7) [1]. (c) Na₂SO₃: Na⁺ neutral. SO₃²⁻ from H₂SO₃ (weak acid) is a basic ion — accepts H⁺ from water: SO₃²⁻ + H₂O ⇌ HSO₃⁻ + OH⁻. Solution is basic (pH > 7) [1]. (d) CH₃COONH₄: NH₄⁺ (from NH₃, weak base) is acidic; CH₃COO⁻ (from CH₃COOH, weak acid) is basic. Ka of NH₄⁺ ≈ Ka of CH₃COOH — the acidic and basic tendencies approximately cancel. Solution is approximately neutral (pH ≈ 7) [1].
Q7 (5 marks): (a) HCl: [H⁺] = 0.10 mol/L (strong, 100% ionised) → pH = −log(0.10) = 1.00 [1]. HNO₂: [H⁺] ≈ √(4.5 × 10⁻⁴ × 0.10) = √(4.5 × 10⁻⁵) = 6.7 × 10⁻³ mol/L → pH = 2.17 [1]. (b) Degree of ionisation = (6.7 × 10⁻³ / 0.10) × 100% = 6.7%. Only 6.7% of HNO₂ molecules have donated their proton at equilibrium — consistent with the definition of a weak acid (partial ionisation). The remaining 93.3% exist as intact HNO₂ molecules [1]. (c) Conductivity is proportional to total ion concentration in solution. HCl (strong, 100% ionised): [H⁺] = [Cl⁻] = 0.10 mol/L → ~0.20 mol/L ions total. HNO₂ (weak, 6.7% ionised): [H⁺] = [NO₂⁻] ≈ 6.7 × 10⁻³ mol/L; ~0.093 mol/L intact HNO₂ molecules remain (uncharged, do not conduct). Total conducting ions ≈ 0.013 mol/L — approximately 15 times lower than HCl at the same total acid concentration. Despite equal total acid (0.10 mol/L), most HNO₂ molecules are intact and non-conducting, so conductivity is far lower [2].
Q8 (7 marks): (a) Acid strength = the degree to which an acid ionises in aqueous solution — the fraction of acid molecules that donate a proton to water at equilibrium, quantified by Ka (the acid dissociation constant). Ka is an intrinsic property of the acid molecule at a given temperature — it does not change with concentration [1]. Acid concentration = the total amount of acid dissolved per litre of solution (mol/L). This is independent of Ka — it describes how many acid molecules are present, not what fraction have ionised [1]. (b) Concert analogy: concentration = crowd size in car park; strength = fraction who rush through the door; [H⁺] = number in the arena. A large crowd (concentrated) with a 1% rush (weak) can produce more arena occupants ([H⁺]) than a small crowd (dilute) with a 100% rush (strong) — because the size of the crowd can overwhelm the effect of the rush fraction. [H⁺] depends on both properties acting together; neither alone determines it [2]. (c) Quantitative example: 0.001 mol/L HCl (dilute strong acid) vs 1.0 mol/L CH₃COOH (concentrated weak acid). HCl: [H⁺] = 0.001 → pH = 3.0. CH₃COOH: [H⁺] ≈ √(1.8 × 10⁻⁵ × 1.0) = 4.2 × 10⁻³ → pH = 2.38. The dilute strong acid has the HIGHER pH (3.0 vs 2.38) despite HCl being intrinsically much stronger — because concentration overwhelms the Ka advantage for these specific values [2]. (d) Method 1: Electrical conductivity at equal concentration — strong acid produces ~65× more ions per litre (fully ionised) than weak acid (partially ionised). Conductivity of strong acid solution will be measurably higher. Method 2: Rate of initial reaction with Mg ribbon — strong acid (higher [H⁺] immediately available) produces H₂ gas visibly faster than weak acid at the same concentration. Both methods measure ion concentration or rate directly rather than relying on the [H⁺]-dependent pH scale [1].
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