R3.1 Proton transfer reactionsIB Chemistry SL: Revision notes
Section 1
Brønsted–Lowry acids, bases and conjugate pairs
A Brønsted–Lowry acid is a proton (H⁺) donor; a Brønsted–Lowry base is a proton acceptor. In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, water donates a proton (acid) and ammonia accepts it (base).
A conjugate acid–base pair is two species that differ by exactly one proton: NH₄⁺/NH₃ and H₂O/OH⁻. To find the conjugate base of an acid, remove one H⁺ (and lower the charge by 1); to find the conjugate acid of a base, add one H⁺ (and raise the charge by 1).
The conjugate base of H₂SO₄ is HSO₄⁻, not SO₄²⁻ — only one proton is removed.
Section 2
Amphiprotic species
Some species can act as both acids and bases; they are amphiprotic (they can donate and accept a proton). Examples: H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻.
HCO₃⁻ + H₃O⁺ → H₂CO₃ + H₂O (HCO₃⁻ as a base) HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O (HCO₃⁻ as an acid)
Water is amphiprotic: with NH₃ it is an acid; with HCl it is a base (HCl + H₂O → H₃O⁺ + Cl⁻).
Section 3
pH and the ion product of water
pH = −log₁₀[H⁺] and [H⁺] = 10⁻ᵖᴴ. The scale is logarithmic: a change of 1 pH unit is a tenfold change in [H⁺]; diluting a strong acid tenfold raises the pH by 1.
Water ionises slightly: H₂O ⇌ H⁺ + OH⁻. The ion product constant, Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K, so [H⁺] and [OH⁻] are inversely related. Acidic: [H⁺] > [OH⁻]; neutral: [H⁺] = [OH⁻]; basic: [H⁺] < [OH⁻]. For [OH⁻] = 2.0 × 10⁻³, [H⁺] = 5.0 × 10⁻¹² and pH = 11.30.
Neutral means [H⁺] = [OH⁻], not pH 7: Kw changes with temperature, so neutral pH is only 7.00 at 298 K.
Section 4
Strong and weak acids and bases
Strong acids and bases ionise completely in water (HCl, HNO₃, H₂SO₄; group 1 hydroxides). Weak ones ionise only partially and set up an equilibrium (CH₃COOH ⇌ CH₃COO⁻ + H⁺; NH₃ + H₂O ⇌ NH₄⁺ + OH⁻).
Strength (extent of ionisation) is not concentration. At equal concentration a weak acid has a higher pH, lower conductivity and reacts more slowly with metals or carbonates, but it neutralises the same amount of base.
Acid–base equilibria lie in the direction of the weaker conjugate: HCl + H₂O → H₃O⁺ + Cl⁻ lies fully to the right, because H₃O⁺ is a weaker acid than HCl; for ethanoic acid the equilibrium lies to the left.
Section 5
Neutralisation reactions
Acid + base → salt + water. Key equations:
- Metal oxide: CuO + H₂SO₄ → CuSO₄ + H₂O
- Metal hydroxide: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
- Carbonate: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
- Hydrogencarbonate: NaHCO₃ + HCl → NaCl + H₂O + CO₂ The ionic equation for a strong acid with a strong base is H⁺(aq) + OH⁻(aq) → H₂O(l). Carbonates and hydrogencarbonates give effervescence (CO₂).
Section 6
pH curves: strong acid with strong base
Adding 0.100 mol dm⁻³ NaOH to 25.0 cm³ of 0.100 mol dm⁻³ HCl: the pH starts at 1, rises slowly while H⁺ is in excess, then rises very steeply around the equivalence point (25.0 cm³), from about 3 to about 11, and levels off near 13 as OH⁻ is in excess. The pH at equivalence is 7 (at 298 K). Titrating a base with an acid gives the mirror-image curve, starting high and falling.
When sketching, mark: the starting pH, the equivalence volume, the vertical section around pH 7, and the final plateau.
Must know
- Acid = proton donor; base = proton acceptor; conjugates differ by one H⁺.
- Amphiprotic: H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻.
- pH = −log₁₀[H⁺]; Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K.
- Strong = fully ionised; weak = partially ionised; equilibrium lies towards the weaker conjugate.
- Equations for oxides, hydroxides, carbonates and hydrogencarbonates.
- Strong acid–strong base curve: steep jump around pH 3–11, equivalence at pH 7.
That's the notes covered.
Carry on to the next subtopic.