In 1884, Svante Arrhenius measured the electrical conductivity of 0.1 mol/L HCl and 0.1 mol/L acetic acid — same concentration, but HCl conducted over 70 times more current. He published this data in his Uppsala doctoral thesis, showing that equal concentrations of different acids produce vastly different numbers of ions. His examiners nearly failed him. Forty years later he received the Nobel Prize.
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
A student is handed two beakers. Beaker A contains 0.1 mol/L hydrochloric acid. Beaker B contains 0.1 mol/L acetic acid (vinegar). Both solutions have exactly the same concentration — the same number of acid molecules per litre.
The student dips a pH probe into each beaker. Beaker A reads pH 1.0. Beaker B reads pH 2.9. Same concentration, different pH — nearly 100 times more H⁺ in the HCl solution than in the acetic acid solution.
Before you read on: Write down your explanation for why two solutions with identical concentrations produce such different pH readings. What is fundamentally different about the two acids at the molecular level? You will return to this at the end of the lesson.
Core Content
These two properties are completely independent · Confusing them is the most common Module 6 error
Before a single equation is written, the conceptual distinction between strength and concentration must be clear — because these two properties are completely independent of each other, and confusing them is the single most common error in the entire module.
Acid strength describes the degree to which an acid ionises in water — what fraction of the original acid molecules donate their protons to water at equilibrium. A strong acid ionises completely — essentially every molecule donates its proton. A weak acid ionises only partially — most molecules remain intact.
Concentration describes the total amount of acid dissolved per litre of solution — regardless of how much of it has ionised. These two properties are completely independent. You can have:
The critical consequence: a concentrated weak acid can have a lower pH than a dilute strong acid. For example, 10 mol/L acetic acid (Ka = 1.8 × 10⁻⁵) produces approximately 0.013 mol/L H⁺ (pH ≈ 1.9). This is a lower pH than 0.001 mol/L HCl (pH = 3.0) — despite HCl being a strong acid, its very low concentration means fewer H⁺ ions are present.
Acid strength (degree of ionisation, measured by Ka) and acid concentration (mol/L) are completely independent — [H⁺] depends on both; diluting an acid reduces concentration but never changes its strength; a concentrated weak acid can have a lower pH than a dilute strong acid (e.g. 10 mol/L CH₃COOH pH ≈ 1.9 vs 0.001 mol/L HCl pH = 3.0).
Pause — copy the highlighted definition into your book before moving on.
Which of the following correctly describes the relationship between acid strength and concentration?
Six strong acids only · Everything else is weak · → for all strong acid equations
We just saw that acid strength and concentration are independent, and that [H⁺] depends on both. That raises a question: Which acids are actually strong — completely ionised? This card answers it → there are exactly six: HCl, HNO₃, H₂SO₄ (1st), HClO₄, HBr, HI — everything else is weak.
There are only six common strong acids — and because there are so few of them, the fastest way to identify a weak acid is to check whether it appears on this list; if it does not, assume it is weak until evidence suggests otherwise.
Strong acids ionise essentially completely in dilute aqueous solution. At the molecular level, the forward reaction (proton donation to water) is so strongly favoured that the reverse reaction is negligible — the equilibrium lies so far to the right that we treat it as irreversible. This is why the ionic equation for a strong acid uses →, not ⇌.
| Strong acid | Ionic equation | Arrow type | Conjugate base | Conjugate base character |
|---|---|---|---|---|
| HCl (hydrochloric) | HCl(aq) → H⁺(aq) + Cl⁻(aq) | → (single) | Cl⁻ | Extremely weak base — spectator ion |
| HNO₃ (nitric) | HNO₃(aq) → H⁺(aq) + NO₃⁻(aq) | → (single) | NO₃⁻ | Extremely weak base — spectator ion |
| H₂SO₄ (sulfuric, 1st ionisation) | H₂SO₄(aq) → H⁺(aq) + HSO₄⁻(aq) | → (single) | HSO₄⁻ | Weak acid (2nd ionisation partial ⇌) |
| HClO₄ (perchloric) | HClO₄(aq) → H⁺(aq) + ClO₄⁻(aq) | → (single) | ClO₄⁻ | Extremely weak base — spectator ion |
| HBr (hydrobromic) | HBr(aq) → H⁺(aq) + Br⁻(aq) | → (single) | Br⁻ | Extremely weak base — spectator ion |
| HI (hydroiodic) | HI(aq) → H⁺(aq) + I⁻(aq) | → (single) | I⁻ | Extremely weak base — spectator ion |
The six strong acids are HCl, HNO₃, H₂SO₄ (1st ionisation only — 2nd is weak: HSO₄⁻ ⇌ H⁺ + SO₄²⁻), HClO₄, HBr, HI — all use → in ionic equations; all acids not on this list are weak in HSC; HF is weak (Ka = 6.8 × 10⁻⁴) despite being a hydrogen halide.
Add the highlighted list to your notes before the check below.
Which of the following is a weak acid?
Four strong bases · Ca(OH)₂ and Ba(OH)₂ give 2 mol OH⁻ · Solubility ≠ strength
We just saw the six strong acids and their ionic equations — everything not on that list is weak. That raises a question: What about bases — which bases are strong and give [OH⁻] = c? This card answers it → four strong bases (NaOH, KOH, Ca(OH)₂, Ba(OH)₂); Ca(OH)₂ and Ba(OH)₂ give [OH⁻] = 2 × c; solubility ≠ strength.
Just as with strong acids, there are relatively few strong bases — and identifying them correctly determines whether you write a single arrow or an equilibrium arrow, and whether you can assume complete dissociation in every calculation that follows.
Strong bases are those that dissociate completely in aqueous solution to give OH⁻ ions. They are all ionic hydroxides of Group 1 and the heavier Group 2 metals:
Ca(OH)₂ and Ba(OH)₂ are strong bases — they dissociate completely — but they have limited solubility. A saturated Ca(OH)₂ solution (limewater) has a concentration of only about 0.02 mol/L at 25°C. The low [OH⁻] is due to low solubility, not weakness. Solubility and strength are different properties.
Mg(OH)₂ is a special case: it has very low solubility AND its dissolved fraction does not fully dissociate — making it weak by the dissociation criterion. All nitrogen-containing bases (NH₃ and organic amines) are weak bases.
The four strong bases are NaOH, KOH, Ca(OH)₂, and Ba(OH)₂ — all dissociate completely using →. For Ca(OH)₂ and Ba(OH)₂: [OH⁻] = 2 × c(base) because each formula unit releases two OH⁻ ions. Low [OH⁻] in limewater reflects low solubility, not weakness — solubility and strength are independent properties.
Pause — write the highlighted definition into your book.
Ca(OH)₂ is a weak base because limewater has a relatively low pH.
→ for all strong acids and bases · ⇌ for all weak acids and bases · No exceptions
We just saw the four strong bases and why [OH⁻] = 2 × c for Ca(OH)₂ and Ba(OH)₂. That raises a question: How does knowing strong vs weak actually change what you write on paper? This card answers it → → for strong (complete), ⇌ for weak (partial) — no exceptions; wrong arrow loses the mark.
The choice between → and ⇌ in an ionic equation is not a stylistic preference — it communicates a physical reality about whether a reaction goes to completion or reaches dynamic equilibrium, and using the wrong arrow changes the meaning of the equation entirely.
Strong acids and strong bases use → because ionisation is complete — the reaction goes essentially to completion, and the reverse reaction is negligible. Weak acids and weak bases use ⇌ because ionisation is partial — forward and reverse reactions both occur at significant rates, establishing a dynamic equilibrium.
Arrow rule: strong acids and bases use → (complete reaction, no reverse); weak acids and bases use ⇌ (partial ionisation, dynamic equilibrium). The decision is binary — check the strong list. If the species is on it, use →; if not, use ⇌. A correct equation with the wrong arrow loses the HSC mark.
Add the highlighted point to your notes before the check below.
Which ionic equation uses the correct arrow notation?
NESA-prescribed investigation · Same concentration, different pH = evidence of different ionisation extent
We just saw the arrow notation rule and how it flows directly from the strong/weak classification. That raises a question: How do we actually observe the strength difference in a lab — and what is the NESA-prescribed way to report it? This card answers it → pH probe shows HCl at 1.0 vs CH₃COOH at 2.9 (same 0.1 mol/L); the gap proves partial ionisation experimentally.
Place a pH probe into 0.1 mol/L HCl: it reads pH 1.0. Place it into 0.1 mol/L acetic acid: it reads pH 2.9. Same concentration — yet acetic acid has only about 1.3% of the H⁺ concentration of HCl. That gap between the two readings is the experimental signature of the difference between strong and weak acids, and it is exactly what the NESA-prescribed investigation asks you to measure and explain.
In the NESA-prescribed practical investigation for IQ2, students measure pH of a range of acid and base solutions using a calibrated digital pH probe. The key comparison that demonstrates strength vs concentration uses equal-concentration solutions of a strong acid and a weak acid:
The pH difference between HCl and CH₃COOH (both 0.1 mol/L) is approximately 1.9 pH units — meaning [H⁺] in HCl is about 79 times higher than in CH₃COOH at the same concentration. This numerical difference is the measurable, experimental proof of the strength distinction.
NESA investigation: at 0.1 mol/L, HCl reads pH 1.0 and CH₃COOH reads pH ~2.9 — the 1.9-unit gap means [H⁺] in HCl is ~79 times higher at identical concentration. This is experimental proof that HCl ionises completely while CH₃COOH ionises only ~1.3%; report it as: "higher [H⁺] at equal concentration because of complete vs partial ionisation."
Pause — copy the highlighted definition into your book before moving on.
In a practical investigation, 0.1 mol/L HCl gives pH 1.0 and 0.1 mol/L CH₃COOH gives pH 2.9. What does this show?
Molecular representation at equal concentration (0.1 mol/L) — strong acid fully ionised; weak acid mostly intact
✏️ Worked Examples
For each of the following, state whether it is a strong or weak acid/base, and write the correct ionic equation using appropriate arrow notation:
(a) HBr dissolving in water; (b) HF dissolving in water; (c) Ba(OH)₂ dissolving in water; (d) NH₃ dissolving in water.
HBr: HBr is on the strong acid list (HCl, H₂SO₄, HNO₃, HClO₄, HBr, HI). Strong acid → single arrow → complete ionisation.
HBr(aq) → H⁺(aq) + Br⁻(aq)
HF: HF is NOT on the strong acid list. HF is a weak acid (Ka = 6.8 × 10⁻⁴). Weak acid → equilibrium arrow → partial ionisation.
HF(aq) ⇌ H⁺(aq) + F⁻(aq)
Ba(OH)₂: Ba(OH)₂ is on the strong base list (NaOH, KOH, Ca(OH)₂, Ba(OH)₂). Strong base → single arrow → complete dissociation. Each formula unit gives 2 OH⁻.
Ba(OH)₂(aq) → Ba²⁺(aq) + 2OH⁻(aq)
NH₃: NH₃ is NOT on the strong base list. NH₃ is a weak base — it partially accepts protons from water via equilibrium.
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Answers: (a) Strong acid — HBr(aq) → H⁺(aq) + Br⁻(aq) (b) Weak acid — HF(aq) ⇌ H⁺(aq) + F⁻(aq) (c) Strong base — Ba(OH)₂(aq) → Ba²⁺(aq) + 2OH⁻(aq) (d) Weak base — NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
A student measures the pH of four solutions, each at 0.10 mol/L, and obtains: HNO₃ = pH 1.0; HNO₂ = pH 2.1; HCl = pH 1.0; CH₃COOH = pH 2.9.
(a) Identify which acids are strong and which are weak. (b) Explain why HNO₃ and HCl give the same pH despite being different compounds. (c) Explain why HNO₂ gives a different pH from HNO₃ despite having a similar formula. (d) A student argues that since CH₃COOH has a higher pH than HNO₂, acetic acid must be more dilute. Evaluate this claim.
HNO₃ — on the strong acid list → strong. HCl — on the strong acid list → strong. HNO₂ — NOT on the strong acid list → weak (Ka = 4.5 × 10⁻⁴). CH₃COOH — NOT on the strong acid list → weak (Ka = 1.8 × 10⁻⁵).
HNO₃ and HCl are both strong acids at 0.10 mol/L. Both ionise completely: HNO₃ → H⁺ + NO₃⁻ and HCl → H⁺ + Cl⁻. For both, [H⁺] = 0.10 mol/L (100% ionisation). pH = −log(0.10) = 1.0 for both. The identity of the spectator anion (NO₃⁻ vs Cl⁻) has no effect on [H⁺].
HNO₂ is a weak acid despite its similar formula to HNO₃. HNO₂ only partially ionises: HNO₂ ⇌ H⁺ + NO₂⁻. At 0.10 mol/L, only a fraction of HNO₂ molecules donate their protons — [H⁺] << 0.10 mol/L. At pH 2.1, [H⁺] = 10⁻²·¹ ≈ 7.9 × 10⁻³ mol/L — only about 7.9% ionised. The structural difference (HNO₃ has one more oxygen, making the conjugate base NO₃⁻ more stable via resonance) makes HNO₃ donate its proton far more readily.
The student's claim is incorrect. Both solutions are at 0.10 mol/L — the same concentration. The difference in pH (2.9 vs 2.1) is not due to a difference in concentration but to a difference in acid strength. CH₃COOH has Ka = 1.8 × 10⁻⁵ (smaller than HNO₂'s Ka = 4.5 × 10⁻⁴) — it ionises to a lesser extent, producing fewer H⁺ ions per mole. This is a direct example of the strength vs concentration distinction: same concentration, different pH, because of different Ka values.
Answers: (a) Strong: HNO₃, HCl. Weak: HNO₂, CH₃COOH. (b) Both 100% ionised at same concentration → [H⁺] = 0.10 mol/L → pH = 1.0 for both. (c) HNO₂ weak — partial ionisation ~7.9%; structural difference explains strength difference. (d) Claim wrong — both are 0.10 mol/L; pH difference reflects weaker Ka of CH₃COOH, not lower concentration. Strength and concentration are independent.
A student claims: "A 0.001 mol/L solution of hydrochloric acid must be a weak acid because its pH is 3.0, which is much less acidic than a 1.0 mol/L solution of acetic acid, which has a pH of 2.4." Identify the errors in the student's reasoning and write a complete, accurate explanation of the relationship between acid strength, concentration, and pH.
Error 1 — Using pH to determine acid strength: The student is using pH as the criterion for determining whether an acid is strong or weak. This is incorrect. Acid strength is determined by Ka (the degree of ionisation) — not by the pH of a particular solution. pH depends on BOTH strength AND concentration. A strong acid at very low concentration can have a higher pH than a weak acid at high concentration.
Error 2 — Misclassifying HCl: HCl is a strong acid at any concentration. 0.001 mol/L HCl ionises completely: HCl → H⁺ + Cl⁻. [H⁺] = 0.001 mol/L. pH = −log(0.001) = 3.0. The pH of 3.0 is caused entirely by the low concentration — not by partial ionisation. 0.001 mol/L HCl is correctly described as a dilute strong acid.
The acetic acid comparison: 1.0 mol/L CH₃COOH has pH 2.4. [H⁺] = 10⁻²·⁴ = 4.0 × 10⁻³ mol/L. Despite being at 1.0 mol/L total concentration, only 0.40% of the acetic acid molecules have ionised. The lower pH (2.4) compared to the dilute HCl (pH 3.0) is a result of high concentration partially compensating for low Ka — not evidence that acetic acid is strong.
Correct relationship: Acid strength (Ka) is an intrinsic property of the acid molecule at a given temperature — it does not change with concentration. HCl is always strong; CH₃COOH is always weak. [H⁺] depends on both Ka and concentration — which is why pH alone cannot distinguish strong from weak. At 0.001 mol/L, HCl is 100% ionised and CH₃COOH is still only ~1.3% ionised — conductivity of 0.001 mol/L HCl would be far higher, which is how strength is properly measured.
Answer: Error 1 — pH cannot determine acid strength; strength is defined by Ka, not pH of a particular solution. Error 2 — HCl is always strong; pH 3.0 reflects low concentration, not partial ionisation. The lower pH of 1.0 mol/L CH₃COOH vs 0.001 mol/L HCl reflects the effect of high concentration producing more total H⁺ despite low Ka. Concentration and strength are independent. 0.001 mol/L HCl = dilute strong acid; 1.0 mol/L CH₃COOH = concentrated weak acid.
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Activities
For each substance or equation below: (i) classify as strong acid, weak acid, strong base, or weak base; (ii) write the correct ionic equation with appropriate arrow notation; (iii) if the equation given is wrong, identify the error and write the correction.
| # | Species / Equation given | Classification | Correct ionic equation | Error (if any) |
|---|---|---|---|---|
| 1 | HClO₄ in water | Write here | Write here | — |
| 2 | HF in water | Write here | Write here | — |
| 3 | HNO₃ ⇌ H⁺ + NO₃⁻ | Write here | Write here | Identify error |
| 4 | Ca(OH)₂ in water | Write here | Write here | — |
| 5 | CH₃COOH in water | Write here | Write here | — |
| 6 | NH₃ + H₂O → NH₄⁺ + OH⁻ | Write here | Write here | Identify error |
| 7 | H₂SO₄ (1st ionisation) | Write here | Write here | — |
| 8 | HSO₄⁻ (2nd ionisation) | Write here | Write here | — |
Check Your Understanding
1. A student measures the pH of 0.10 mol/L solutions of four acids and obtains: Acid W pH 1.0; Acid X pH 3.2; Acid Y pH 1.0; Acid Z pH 2.6. Which conclusion is best supported by this data?
2. Which of the following ionic equations contains an error in arrow notation?
3. A 0.10 mol/L solution of a weak acid HA has pH 3.5 at 25°C. A student dilutes this solution to 0.010 mol/L. Which statement correctly predicts the effect of dilution on the acid's strength and pH?
4. Which of the following correctly describes Ca(OH)₂ in aqueous solution?
5. In a NESA-prescribed practical, a student measures pH of 0.10 mol/L HCl (pH 1.0) and 0.10 mol/L CH₃COOH (pH 2.9). Which of the following correctly interprets this observation as evidence for acid strength?
✍️ Short Answer
ApplyBand 3(4 marks) 6. Write the correct ionic equation with appropriate arrow notation for each of the following dissolving in water: (a) HI; (b) HNO₂; (c) KOH; (d) CH₃NH₂ (methylamine, a nitrogen-containing organic base). For each, state whether the substance is strong or weak and justify your arrow choice.
AnalyseBand 4(4 marks) 7. A chemist prepares two solutions: Solution X: 5.0 mol/L ethanoic acid (CH₃COOH, Ka = 1.8 × 10⁻⁵) and Solution Y: 0.001 mol/L hydrochloric acid. Without doing a full Ka calculation, explain: (a) which solution has the lower pH, and (b) why it is misleading to say that "the HCl is a stronger acid than the ethanoic acid, therefore it must produce a more acidic solution." Your answer must address the distinction between acid strength, concentration, and [H⁺].
EvaluateBand 6(6 marks) 8. Four students were asked to explain why 0.1 mol/L HCl has a lower pH than 0.1 mol/L CH₃COOH.
Student A: "HCl ionises completely, giving [H⁺] = 0.1 mol/L and pH = 1.0. CH₃COOH partially ionises, so [H⁺] << 0.1 mol/L and pH is higher. Strength is about the degree of ionisation — independent of concentration."
Student B: "HCl has a lower pH because it is more concentrated than the CH₃COOH solution. Stronger acids have higher concentrations."
Student C: "HCl is strong so it ionises more. If you made the CH₃COOH more concentrated, it would also become strong because there would be enough molecules to ionise."
Student D: "Weak acids like CH₃COOH barely ionise so they are barely acidic. You could almost drink concentrated acetic acid safely."
Identify which student is correct. For each incorrect student, identify the specific error in their reasoning.
1. HClO₄: Strong acid (on strong acid list). HClO₄(aq) → H⁺(aq) + ClO₄⁻(aq). → because complete ionisation.
2. HF: Weak acid (NOT on strong acid list; Ka = 6.8 × 10⁻⁴). HF(aq) ⇌ H⁺(aq) + F⁻(aq). ⇌ because partial ionisation.
3. HNO₃ ⇌ H⁺ + NO₃⁻: Strong acid. INCORRECT ARROW — HNO₃ is strong → must use →. Correct: HNO₃(aq) → H⁺(aq) + NO₃⁻(aq).
4. Ca(OH)₂: Strong base. Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH⁻(aq). → because complete dissociation.
5. CH₃COOH: Weak acid. CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq). ⇌ because partial ionisation.
6. NH₃ + H₂O → NH₄⁺ + OH⁻: Weak base. INCORRECT ARROW — NH₃ is weak → must use ⇌. Correct: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq).
7. H₂SO₄ (1st ionisation): Strong acid. H₂SO₄(aq) → H⁺(aq) + HSO₄⁻(aq).
8. HSO₄⁻ (2nd ionisation): Weak acid. HSO₄⁻(aq) ⇌ H⁺(aq) + SO₄²⁻(aq).
1. B — At 0.10 mol/L, a strong acid produces [H⁺] = 0.10 mol/L → pH = 1.0. W and Y both give pH 1.0 — consistent with complete ionisation → both strong. X (pH 3.2) and Z (pH 2.6) give higher pH than expected for complete ionisation — both weak. X has higher pH than Z at the same concentration → X has smaller Ka → X is weaker than Z. Option A wrong — same pH does not mean same acid. Option C wrong — all solutions are at the same concentration.
2. C — HF is a weak acid (Ka = 6.8 × 10⁻⁴) — ionises only partially. The correct equation uses ⇌: HF(aq) ⇌ H⁺(aq) + F⁻(aq). Using → implies complete ionisation, which is incorrect for HF. Options A and B correct — HClO₄ and H₂SO₄ first ionisation are strong, using →.
3. B — Ka is an intrinsic property of the acid at constant temperature — it does not change with dilution. Diluting from 0.10 to 0.010 mol/L reduces total molecules per litre. The absolute [H⁺] still decreases. Net effect: [H⁺] decreases → pH increases.
4. C — Ca(OH)₂ is a strong base — the dissolved fraction dissociates 100%: Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH⁻(aq). Each mole dissolved gives 2 mol OH⁻. The low [OH⁻] in limewater is because only ~0.02 mol/L dissolves — limited solubility, not partial dissociation.
5. B — At equal concentration (0.1 mol/L), HCl produces [H⁺] = 0.1 mol/L (pH 1.0) while CH₃COOH produces [H⁺] ≈ 0.0013 mol/L (pH 2.9). This is direct experimental evidence that HCl ionises to a much greater extent than CH₃COOH at the same concentration.
Q6 (4 marks): (a) HI: strong acid (on strong acid list). HI(aq) → H⁺(aq) + I⁻(aq). Single arrow — complete ionisation [1]. (b) HNO₂: weak acid (NOT on strong acid list; Ka = 4.5 × 10⁻⁴). HNO₂(aq) ⇌ H⁺(aq) + NO₂⁻(aq). Equilibrium arrow — partial ionisation [1]. (c) KOH: strong base (on strong base list). KOH(aq) → K⁺(aq) + OH⁻(aq). Single arrow — complete dissociation [1]. (d) CH₃NH₂: weak base — nitrogen-containing organic amine, NOT on strong base list. CH₃NH₂(aq) + H₂O(l) ⇌ CH₃NH₃⁺(aq) + OH⁻(aq). Equilibrium arrow — partial ionisation [1].
Q7 (4 marks): (a) Solution X (5.0 mol/L CH₃COOH) has the lower pH. Even though CH₃COOH is weak, 5.0 mol/L is an extremely high concentration. [H⁺] ≈ √(Ka × c) ≈ √(1.8 × 10⁻⁵ × 5.0) ≈ 0.0095 mol/L → pH ≈ 2.0. Solution Y: [H⁺] = 0.001 mol/L → pH = 3.0 [1]. (b) The statement is misleading because it conflates acid strength (Ka) with [H⁺] in solution [1]. HCl is indeed stronger (Ka >> CH₃COOH) but [H⁺] depends on BOTH strength AND concentration [1]. At 0.001 mol/L, HCl has very few molecules even though 100% ionised — total [H⁺] = 0.001 mol/L. At 5.0 mol/L, CH₃COOH ionises partially but has so many molecules the absolute [H⁺] exceeds the dilute HCl [1].
Q8 (6 marks): Student A is correct [1]. Student B error: confuses strength with concentration. The two solutions are at the same concentration (0.1 mol/L). HCl's lower pH is entirely due to its complete ionisation (100%) vs CH₃COOH's partial ionisation (~1.3%). Acid strength (Ka) is an intrinsic molecular property — it does not change when concentration changes [2]. Student C error: states weak acid becomes strong at high concentration. Ka is determined by molecular structure and bond energies — it is fixed at a given temperature regardless of concentration. At 10 mol/L, CH₃COOH is still a weak acid [2]. Student D error: conflates ionisation fraction with corrosiveness or safety. Glacial acetic acid (~17 mol/L) is a classified corrosive substance. "Weak" refers strictly to ionisation fraction — never to safety or absolute acidity [1].
Look back at what you wrote about Beaker A (HCl) and Beaker B (CH₃COOH). Recall Arrhenius's 1884 conductivity data: 0.1 mol/L HCl conducted 70× more current than 0.1 mol/L acetic acid at the same concentration — because HCl fully ionises and acetic acid only partially ionises (Ka = 1.8 × 10⁻⁵). You explained why two solutions with identical concentrations produce such different pH readings.
Review
A student examines two solutions: Solution A is 0.001 mol/L HCl (pH 3.0) and Solution B is 1.0 mol/L CH₃COOH (Ka = 1.8 × 10⁻⁵, pH ≈ 2.4). The student concludes "CH₃COOH must be a strong acid because it has a lower pH than HCl." Identify all errors in this reasoning and write a complete Band 6 response explaining the relationship between acid strength, concentration, and [H⁺]. (6 marks)
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