Year 12 Chemistry Module 6, Acid/Base Reactions ⏱ ~45 min Lesson 5 of 19 IQ2

Strong vs Weak Acids & Bases, The Critical Distinction

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.

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Today's hook, In 1884, Arrhenius measured two solutions, same concentration, one 70× more conductive than the other. He nearly failed his doctorate over the explanation. What did he see, and why does it matter for every Ka calculation you'll ever do?

0/5QUESTS

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Get oriented

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions

Before You Read

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.

→

📐 Arrow Notation Rules, The Non-Negotiable Framework

Strong acid: HA(aq) → H⁺(aq) + A⁻(aq)

→ single arrow: complete, irreversible ionisation [H⁺] = initial concentration of acid (100% ionised) e.g. HCl(aq) → H⁺(aq) + Cl⁻(aq)

Weak acid: HA(aq) ⇌ H⁺(aq) + A⁻(aq)

⇌ equilibrium arrow: partial, reversible ionisation [H⁺] << initial concentration of acid (only partially ionised) e.g. CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)

Strong base: NaOH(aq) → Na⁺(aq) + OH⁻(aq)

→ single arrow: complete dissociation [OH⁻] = initial concentration of base (100% dissociated) Ca(OH)₂ → Ca²⁺ + 2OH⁻ [OH⁻] = 2 × [Ca(OH)₂]

Weak base: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

⇌ equilibrium arrow: partial proton acceptance [OH⁻] << initial concentration of base Key independence rule: Strength (Ka) is fixed at a given temperature, independent of concentration. [H⁺] depends on BOTH.

Learning Intentions

Know

Key facts

The six strong acids: HCl, H₂SO₄ (1st), HNO₃, HClO₄, HBr, HI
The four strong bases: NaOH, KOH, Ca(OH)₂, Ba(OH)₂
All other acids and bases encountered in HSC are weak
Arrow notation rule: strong → single arrow; weak → equilibrium arrow

Understand

Concepts

Strength = degree of ionisation (Ka), intrinsic, temperature-dependent, not concentration-dependent
Concentration and strength are completely independent properties
Why a concentrated weak acid can have a lower pH than a dilute strong acid
Why HF is a weak acid despite being a hydrogen halide like HCl, HBr, HI

Can do

Skills

Write correct ionic equations with → or ⇌ for any acid or base
Identify acid strength from pH data at equal concentration
Explain why two equal-concentration solutions have different pH values
Correct common errors: "dilute = weak", "weak = safe", "HF is strong"

Scan these before reading

Strong acid

An acid that completely dissociates in water; [H⁺] = initial acid concentration (e.g., HCl, H₂SO₄, HNO₃).

Weak acid

An acid that only partially dissociates in water; establishes an equilibrium with a small Ka.

Degree of dissociation

The fraction of acid or base molecules that have ionised; much higher for strong electrolytes.

Conductivity

A measure of a solution's ability to conduct electricity; strong acids/bases conduct better than weak at the same concentration.

pH vs concentration

Changing concentration changes pH for both strong and weak acids; changing from strong to weak also lowers [H⁺] at the same concentration.

Strength vs concentration distinction

Acid strength (strong/weak) is fixed by Ka; concentration describes how many moles per litre, they are independent properties.

Strength vs concentration, the critical independence

Core Content

Cross-lesson links: The strength vs concentration distinction introduced here is the foundation for Ka and Kb calculations (L09), ICE table pH calculations (L09–L10), buffer design (L12–L13), and titration curve shape interpretation (L14–L17). Any time an exam question distinguishes "strong" from "concentrated," this lesson is the source.

Strength vs Concentration, The Critical Independence

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:

A concentrated strong acid (12 mol/L HCl, many molecules, all ionised)
A dilute strong acid (0.001 mol/L HCl, few molecules, all ionised)
A concentrated weak acid (10 mol/L CH₃COOH, many molecules, few ionised)
A dilute weak acid (0.001 mol/L CH₃COOH, few molecules, few ionised)

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

Fraction of molecules that ionise

Determined by Ka value

Does NOT change with dilution

Acid concentration

Total moles of acid per litre

Determined by amount dissolved

Yes, dilution reduces concentration

[H⁺] in solution

Actual proton concentration

Determined by BOTH strength AND concentration

Yes, depends on both

Critical Rule

In every pH calculation in Module 6, the first question you must ask is: is this acid strong or weak? If strong, [H⁺] = concentration (complete ionisation, use directly). If weak, [H⁺] ≠ concentration (partial ionisation, must use Ka and ICE table, covered in L09). Selecting the wrong method gives a completely wrong answer.

Common Error

"Dilute HCl is a weak acid." This is completely wrong. HCl is a strong acid at any concentration, 12 mol/L, 0.1 mol/L, or 0.000001 mol/L. Strength describes the intrinsic tendency of an acid to donate protons, not how much acid is present. Diluting HCl makes it less concentrated; it does not make it weak. The correct term for a low-concentration HCl solution is "dilute strong acid", never "weak acid."

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?

The six strong acids

Strong Acids, The Complete List and Their Ionic Equations

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

Must Know

Memorise these six strong acids: HCl, H₂SO₄ (1st ionisation), HNO₃, HClO₄, HBr, HI. Every acid not on this list is treated as weak in HSC contexts. This list must be recalled instantly when deciding arrow notation and calculation method in any question involving acids.

Common Error

HF (hydrofluoric acid) is a weak acid despite being a hydrogen halide like HCl, HBr, and HI. Students assume all hydrogen halides are strong, this is wrong. The H–F bond is unusually short and strong (bond enthalpy 570 kJ/mol vs H–Cl 432 kJ/mol), making proton donation much less favourable. HF has Ka = 6.8 × 10⁻⁴, it ionises only partially. The ionic equation for HF must use ⇌, not →.

Safety

"Weak" never means "safe." Hydrofluoric acid (HF) is a weak acid but is extraordinarily dangerous: it penetrates skin and attacks bone and nerves, and concentrated HF can be fatal from skin contact alone. HF is not a school practical reagent, you study it here only through equations, teacher demonstration footage, or secondary data, never by handling it yourself. The same warning applies to glacial (concentrated) acetic acid, a "weak" acid that is still a classified corrosive dangerous good.

Insight

The strength order of hydrogen halides (HF << HCl < HBr < HI in terms of acid strength) is explained by bond strength decreasing down the group. As the halide ion gets larger, the H–X bond lengthens and weakens, making proton donation progressively easier. HF is the exception, the unusually high H–F bond strength makes it significantly harder to donate the proton than for HCl.

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?

The four strong bases

Strong Bases, The Complete List and Their Ionic Equations

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:

NaOH

NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Most common strong base in HSC

KOH

KOH(aq) → K⁺(aq) + OH⁻(aq)

Group 1 hydroxide, fully dissociated

Ca(OH)₂

Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH⁻(aq)

Limited solubility; strong when dissolved; [OH⁻] = 2 × [Ca(OH)₂]

Ba(OH)₂

Ba(OH)₂(aq) → Ba²⁺(aq) + 2OH⁻(aq)

More soluble than Ca(OH)₂; produces 2 mol OH⁻ per mol

NH₃ (weak base, for comparison)

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

Weak base, partial proton acceptance (⇌)

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.

Must Know

For Ca(OH)₂ and Ba(OH)₂, each formula unit produces TWO moles of OH⁻ per mole of base dissolved. When calculating [OH⁻] from a Ca(OH)₂ solution: [OH⁻] = 2 × [Ca(OH)₂]. Forgetting this factor of 2 in pH calculations is a consistent exam error.

Common Error

"Ca(OH)₂ is a weak base because its solution has a low pH." Wrong, Ca(OH)₂ is a strong base. Its solution has a relatively low [OH⁻] because of low solubility limiting the total amount dissolved, not because of partial dissociation. Mg(OH)₂ is weak; Ca(OH)₂ is strong.

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.

Arrow notation, the non-negotiable rule

Arrow Notation, The Non-Negotiable Rule

→ 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.

Strong acid (e.g. HCl)

Correct: HCl → H⁺ + Cl⁻

Wrong: HCl ⇌ H⁺ + Cl⁻

⇌ implies partial ionisation, HCl ionises completely

Weak acid (e.g. CH₃COOH)

Correct: CH₃COOH ⇌ H⁺ + CH₃COO⁻

Wrong: CH₃COOH → H⁺ + CH₃COO⁻

→ implies complete ionisation, only ~1% ionised at 0.1 mol/L

Strong base (e.g. NaOH)

Correct: NaOH → Na⁺ + OH⁻

Wrong: NaOH ⇌ Na⁺ + OH⁻

⇌ implies partial dissociation, NaOH dissociates completely

Weak base (e.g. NH₃)

Correct: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

Wrong: NH₃ + H₂O → NH₄⁺ + OH⁻

→ implies complete proton acceptance, NH₃ only partially accepts protons

Critical Rule

Before writing any ionic equation in Module 6, ask yourself two questions: (1) Is this acid or base strong or weak? (2) Does it appear on the strong acids/bases list? If yes → use →. If no → use ⇌. This two-question check takes three seconds and prevents the most common error in the entire module.

Common Error

Using ⇌ for strong acid dissociation is the single most frequently penalised arrow notation error in HSC Module 6. The equation HCl ⇌ H⁺ + Cl⁻ is chemically wrong, it implies Cl⁻ has a meaningful tendency to accept H⁺ back from solution, which it essentially cannot. A correct formula with the wrong arrow loses the mark.

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?

The practical, and the evidence for strength

The Practical, Measuring pH to Demonstrate Strength vs Concentration

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:

HCl

Concentration: 0.1 mol/L

Expected pH: 1.0

[H⁺] = 0.1 mol/L

Strong acid, 100% ionised

CH₃COOH

Concentration: 0.1 mol/L

Expected pH: ~2.9

[H⁺] ≈ 0.0013 mol/L

Weak acid, ~1.3% ionised

NaOH

Concentration: 0.1 mol/L

Expected pH: 13.0

[OH⁻] = 0.1 mol/L

Strong base, 100% dissociated

NH₃

Concentration: 0.1 mol/L

Expected pH: ~11.1

[OH⁻] ≈ 0.0013 mol/L

Weak base, ~1.3% ionised (Kb = 1.8 × 10⁻⁵)

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.

Practical Report

In a practical report question, describe what observation (pH reading) provides evidence for the strength distinction, and connect it explicitly to the molecular-level explanation: "The pH of 0.1 mol/L CH₃COOH was 2.9 compared to pH 1.0 for 0.1 mol/L HCl, demonstrating that CH₃COOH produces far fewer H⁺ ions per mole dissolved, consistent with partial ionisation." This is the minimum complete response.

Common Error

Students describe the practical result as "HCl is more acidic than CH₃COOH." While technically true for equal concentrations, this phrasing confuses strength with absolute [H⁺] and does not address the concept being demonstrated. The correct description: "At the same concentration, HCl produces a higher [H⁺] than CH₃COOH because HCl ionises completely while CH₃COOH ionises only partially, this is the distinction between strong and weak acids."

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?

VISUAL SUMMARY

At 0.1 mol/L, HCl ionises fully while ethanoic acid barely ionises, so their pH values differ.

Molecular representation at equal concentration (0.1 mol/L), strong acid fully ionised; weak acid mostly intact

Misconceptions to Fix

✗ "Dilute HCl is a weak acid."

✓ HCl is a strong acid at any concentration. 0.0001 mol/L HCl is still a dilute strong acid, every molecule still donates its proton to water completely. Use "dilute strong acid", never "weak acid", for low-concentration HCl.

✗ "HF is a strong acid because it is a hydrogen halide like HCl, HBr, HI."

✓ HF is a weak acid (Ka = 6.8 × 10⁻⁴). The H–F bond is unusually strong due to fluorine's small atomic radius and very high electronegativity, making proton donation energetically unfavourable. HF must use ⇌: HF(aq) ⇌ H⁺(aq) + F⁻(aq).

✗ "Weak acids are safe because they don't ionise much."

✓ Glacial acetic acid (pure, ~17 mol/L) is a classified corrosive dangerous good. Concentrated HF causes fatal systemic toxicity from skin absorption alone. "Weak" describes the fraction of molecules that ionise, never the concentration, hazard level, or biological danger of the substance.

✗ "Ca(OH)₂ is a weak base because limewater has a relatively low pH."

✓ Ca(OH)₂ is a strong base, the dissolved fraction dissociates 100%. The low [OH⁻] in limewater is due to low solubility (~0.02 mol/L at 25°C), not partial dissociation. Solubility and strength are independent properties.

Worked examples and apply

✏️ Worked Examples

Worked Example 1, Classifying Acids and Bases · Band 3

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)

Worked Example 2, Explaining pH Differences · Band 4–5

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.

Worked Example 3, Extended Response · Band 6

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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The Six Strong Acids

HCl, hydrochloric
H₂SO₄, sulfuric (1st ionisation only)
HNO₃, nitric
HClO₄, perchloric
HBr, hydrobromic
HI, hydroiodic
All others = WEAK (esp. HF, CH₃COOH, HNO₂)

The Four Strong Bases

NaOH, sodium hydroxide
KOH, potassium hydroxide
Ca(OH)₂, calcium hydroxide (2 OH⁻ per mol!)
Ba(OH)₂, barium hydroxide (2 OH⁻ per mol!)
All others = WEAK (esp. NH₃, Mg(OH)₂)
Ca(OH)₂ strong but low solubility, not weak!

Arrow Notation Rules

Strong acid: HA → H⁺ + A⁻ (→ always)
Weak acid: HA ⇌ H⁺ + A⁻ (⇌ always)
Strong base: MOH → M⁺ + OH⁻ (→ always)
Weak base: B + H₂O ⇌ BH⁺ + OH⁻ (⇌)
Wrong arrow = wrong mark, always

Strength vs Concentration

Strength = Ka = fraction ionised (intrinsic)
Concentration = mol/L (changes with dilution)
[H⁺] depends on BOTH, pH alone ≠ strength
Dilute strong acid ≠ weak acid
Concentrated weak acid can have lower pH than dilute strong acid

Use the pH Calculator. The pH of a 0.01 mol/L strong acid solution is…

🔀 Sort the Steps +7 XP

Arrange these criteria for distinguishing strong from weak acids in order of reliability for HSC responses:

pH for same concentration (strong → lower pH)

Conductivity (strong → higher for same concentration)

Reaction rate with metals (strong → faster)

Equation arrow used (→ for strong, ⇌ for weak)

Degree of ionisation (strong = complete, weak = partial)

Interactive · Strong vs Weak Explorer

Use the explorer. Compare 0.10 mol/L HCl with 0.10 mol/L CH₃COOH (K_a = 1.8 × 10⁻⁵). Which statement is correct?

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Practice questions

Activities

Strong, Weak, or Wrong Arrow?, Classification Challenge

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

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Multiple Choice

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✍️ 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.

Show All Answers

Activity 1, Sort + Classify

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).

MC Explanations

Reading pH data at equal concentration. 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.

Spotting the wrong arrow. 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 →.

Dilution leaves Ka unchanged. 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.

Ca(OH)₂: strong but low-solubility. 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.

Interpreting the practical result. 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.

Short Answer Model Answers

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].

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Review

What are the six strong acids?

HCl, H₂SO₄ (1st ionisation), HNO₃, HClO₄, HBr, HI. All others are treated as weak in HSC.

What arrow is used for a strong acid and why?

→ (single arrow) because strong acid ionisation is complete and irreversible, the equilibrium lies so far to the right it is treated as going to completion.

Why is HF a weak acid despite being a hydrogen halide?

The H–F bond is unusually strong (bond enthalpy 570 kJ/mol) due to fluorine's small atomic radius and high electronegativity, making proton donation energetically unfavourable. Ka = 6.8 × 10⁻⁴, partial ionisation only.

Can a concentrated weak acid have a lower pH than a dilute strong acid?

Yes. For example, 10 mol/L CH₃COOH (pH ≈ 1.9) has lower pH than 0.001 mol/L HCl (pH = 3.0) because [H⁺] depends on BOTH strength AND concentration, high concentration compensates for low Ka.

What is the correct ionic equation for Ca(OH)₂ dissolving in water, and why must [OH⁻] = 2 × [Ca(OH)₂]?

Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH⁻(aq). Each formula unit of Ca(OH)₂ contains two OH⁻ ions, so complete dissociation of 1 mol Ca(OH)₂ produces 2 mol OH⁻.

Extended Response

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 high-mark response explaining the relationship between acid strength, concentration, and [H⁺]. (6 marks)

How did your thinking change?

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.

Did your explanation use the phrase "degree of ionisation", or did you use vaguer language like "more acidic"?
Can you now write the two ionic equations with correct arrow notation and explain what the arrow communicates about the reaction at the molecular level?
Did you predict that a concentrated weak acid could have a lower pH than a dilute strong acid? Could you now construct a specific numerical example demonstrating this?

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