Brønsted–Lowry acids and bases
Section: 7 Equilibria | Syllabus: Cambridge AS Level Physics 9702
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 Pairs Diagram: 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 Pairs Example 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 Fig 7.8 - Degree of Dissociation: Strong vs Weak Acid 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 Link to Enthalpy The standard enthalpy of neutralisation for a strong acid + strong base is approximately −57 kJ mol⁻¹ per mole of water formed.
For weak acid or weak base reactions, the value is less exothermic because energy is required to fully dissociate the weak acid or base. 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 Combinations Diagram: 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 …
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