Brønsted–Lowry acids and bases
Section: 7 Equilibria | Syllabus: Cambridge AS Level Chemistry (9701)
Common Acids and Alkalis
| Name | Formula | Strong or Weak? |
|---|---|---|
| Hydrochloric acid | HCl | Strong acid |
| Sulfuric acid | H₂SO₄ | Strong acid |
| Nitric acid | HNO₃ | Strong acid |
| Ethanoic acid | CH₃COOH | Weak acid |
| Sodium hydroxide | NaOH | Strong base (alkali) |
| Potassium hydroxide | KOH | Strong base (alkali) |
| Ammonia | NH₃ | Weak base |
The Brønsted–Lowry Theory
Brønsted–Lowry Acid A proton (H⁺) donor - a species that can donate a hydrogen ion to another species. Brønsted–Lowry Base A proton (H⁺) acceptor - a species that can accept a hydrogen ion from another species.When an acid donates a proton it becomes its conjugate base. When a base accepts a proton it becomes its conjugate acid. These form conjugate acid–base pairs.
Fig 7.7 - Conjugate Acid–Base PairsDiagram: The equation HCl + H₂O ⇌ H₃O⁺ + Cl⁻ displayed with curved arrows. HCl is labelled "acid (proton donor)"; Cl⁻ is labelled "conjugate base of HCl". H₂O is labelled "base (proton acceptor)"; H₃O⁺ is labelled "conjugate acid of water". Two coloured brackets highlight conjugate pair 1 (HCl/Cl⁻) and conjugate pair 2 (H₂O/H₃O⁺).
Identifying Conjugate PairsExample 1: HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl⁻(aq)
Pair 1: HCl (acid) / Cl⁻ (conjugate base)
Pair 2: H₂O (base) / H₃O⁺ (conjugate acid)
Example 2: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Pair 1: NH₃ (base) / NH₄⁺ (conjugate acid)
Pair 2: H₂O (acid) / OH⁻ (conjugate base)
Amphoteric Species Water is amphoteric - it can act as both a Brønsted–Lowry acid (donating H⁺ to form OH⁻) and a Brønsted–Lowry base (accepting H⁺ to form H₃O⁺), depending on what it reacts with.Strong and Weak Acids and Bases
Strong Acid Fully (completely) dissociates in aqueous solution - essentially an irreversible process:HCl(aq) → H⁺(aq) + Cl⁻(aq)
A single arrow (→) is used, not a reversible arrow. Weak Acid Only partially dissociates in aqueous solution - an equilibrium is established between the undissociated acid and its ions. A reversible arrow (⇌) must be used:
CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq) Strong Base Fully dissociates in aqueous solution to produce OH⁻ ions:
NaOH(aq) → Na⁺(aq) + OH⁻(aq) Weak Base Only partially dissociates in aqueous solution:
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Comparing Strong and Weak Acids (same concentration)
| Property | Strong Acid (e.g. HCl) | Weak Acid (e.g. CH₃COOH) |
|---|---|---|
| Dissociation | Complete (→) | Partial (⇌) |
| pH at same concentration | Lower (more H⁺ ions) | Higher (fewer H⁺ ions) |
| Rate of reaction with Mg | Faster | Slower |
| Electrical conductivity | Higher (more ions present) | Lower (fewer ions present) |
| Universal indicator colour | Bright red | Lighter red/orange |
Diagram: Two beakers side by side, both labelled 0.1 mol dm⁻³. Left beaker (HCl): all molecules shown fully dissociated as H⁺ and Cl⁻ ions - no HCl molecules remain. Right beaker (CH₃COOH): mostly intact CH₃COOH molecules with only a few H⁺ and CH₃COO⁻ ions present. pH meter readings displayed: ~1 for HCl and ~3 for CH₃COOH. Caption: "Same concentration - very different [H⁺]."
The pH Scale
- The pH scale is a measure of the hydrogen ion concentration [H⁺] in solution
- pH 7 = neutral (pure water at 25°C)
- pH < 7 = acidic
- pH > 7 = alkaline (basic)
- The scale is logarithmic: each unit decrease in pH represents a 10-fold increase in [H⁺]
Neutralisation and Salt Formation
Neutralisation The reaction between an acid and a base to form a salt and water. The net ionic equation for all neutralisations involving H⁺ and OH⁻ is:H⁺(aq) + OH⁻(aq) → H₂O(l)
The spectator ions left in solution form the salt:
| Acid | Base/Alkali | Salt Formed | Equation |
|---|---|---|---|
| HCl | NaOH | Sodium chloride (NaCl) | HCl + NaOH → NaCl + H₂O |
| H₂SO₄ | KOH | Potassium sulfate (K₂SO₄) | H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O |
| HNO₃ | NH₃ | Ammonium nitrate (NH₄NO₃) | HNO₃ + NH₃ → NH₄NO₃ |
| CH₃COOH | NaOH | Sodium ethanoate (CH₃COONa) | CH₃COOH + NaOH → CH₃COONa + H₂O |
pH Titration Curves
A pH titration curve plots pH (y-axis) against volume of base added from the burette (x-axis). The shape depends on the strength of the acid and base used.
Fig 7.9 - pH Curves for Four Titration CombinationsDiagram: Four separate pH vs volume graphs arranged in a 2×2 grid. (1) Strong acid + strong base: classic S-shaped curve with a sharp near-vertical section at the equivalence point at pH 7. (2) Weak acid + strong base: gradual rise initially (buffer region), then sharp section with equivalence point at pH 8–9. (3) Strong acid + weak base: sharp section with equivalence point at pH 5–6, then gradual rise. (4) Weak acid + weak base: no sharp section, equivalence point near pH 7. All graphs label the equivalence point and show the volume of base at the equivalence point.
| Titration Type | Equivalence Point pH | Sharp Vertical Section? |
|---|---|---|
| Strong acid + strong base | pH 7 | Yes - large and clear |
| Weak acid + strong base | pH 8–9 (alkaline) | Yes |
| Strong acid + weak base | pH 5–6 (acidic) | Yes |
| Weak acid + weak base | ~pH 7 | No - no sharp section |
Choosing a Suitable Indicator
An indicator must change colour within the sharp vertical section of the titration curve. For weak acid–weak base titrations (no sharp section), no indicator is suitable - a pH meter must be used.
Fig 7.10 - pH Ranges of Common IndicatorsDiagram: A horizontal pH scale from 0 to 14. Four coloured bars show the pH transition ranges of: methyl orange (red→orange, pH 3.1–4.4), litmus (red→blue, pH 5–8), phenolphthalein (colourless→pink, pH 8.2–10.0), and thymolphthalein (colourless→blue, pH 9.3–10.5). Vertical dashed lines at pH 5–6, 7, and 8–9 mark the equivalence points of the three titration types, showing which indicator's range overlaps with each equivalence point.
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