Proton Transfer Reactions

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Notes de leçon

Brønsted–Lowry Acids & Bases

  • A Brønsted–Lowry acid is a species that donates a proton (H⁺).
  • A Brønsted–Lowry base is a species that accepts a proton using a lone pair of electrons.
  • The theory applies to proton transfer and is not limited to aqueous solutions; it also applies to gas-phase reactions.
  • In HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq), HCl acts as the acid and water acts as the base.
  • The formation of ions in this way is called dissociation or ionisation.
  • A hydrogen ion (H⁺) is simply a proton because a hydrogen atom that loses its electron leaves only a proton.

Proton transfer between an acid and a base

Proton transfer between an acid and a base

Conjugate Acids & Bases

  • A conjugate acid–base pair consists of two species that differ by one H⁺ ion.
  • In CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq), CH₃COOH and CH₃COO⁻ form one conjugate pair, and H₂O and H₃O⁺ form the other.
  • The acid and its conjugate base are related by the loss of a proton; the base and its conjugate acid are related by the gain of a proton.
  • Strong acids produce weak conjugate bases; weak acids produce strong conjugate bases.
  • The chloride ion, Cl⁻, is the conjugate base of HCl and is a very weak base because the reverse reaction is virtually non-existent.
  • The hydroxide ion, OH⁻, is a strong base; its conjugate acid is water, which is a weak conjugate acid.

Amphiprotic Species

  • An amphiprotic species can act as both a proton donor and a proton acceptor.
  • Water is amphiprotic: it donates a proton to ammonia (acting as an acid) and accepts a proton from hydrochloric acid (acting as a base).
  • An amphoteric substance can react as both an acid and a base, but it does not necessarily transfer protons.
  • All amphiprotic substances are amphoteric, but not all amphoteric substances are amphiprotic.
  • Aluminium oxide, Al₂O₃, is amphoteric because it reacts with both HCl and NaOH, but it is not amphiprotic.
  • The anion C₃H₃O₄⁻ can be classified as amphiprotic because it can both donate and accept a proton.

The pH Scale

  • pH is defined as pH = –log[H⁺], where [H⁺] is the concentration of hydrogen ions in mol dm⁻³.
  • The hydrogen ion concentration can be found from pH using [H⁺] = 10^(–pH).
  • The pH scale is logarithmic with base 10, so a change of one pH unit corresponds to a tenfold change in [H⁺].
  • Acidic solutions have pH below 7; basic solutions have pH above 7; pure water at 298 K has pH 7.
  • The higher the [H⁺], the lower the pH; the higher the [OH⁻], the higher the pH.
  • A pH meter gives the most accurate pH reading; universal indicator paper is less accurate and is compared to a colour chart.

The pH scale

The pH scale

The Ion Product of Water

  • Water dissociates slightly: H₂O(l) ⇌ H⁺(aq) + OH⁻(aq).
  • The ion product of water is Kw = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K.
  • The product [H⁺][OH⁻] is always 1.00 × 10⁻¹⁴ at 298 K, so if one concentration increases, the other must decrease.
  • In neutral water at 298 K, [H⁺] = [OH⁻] = 1 × 10⁻⁷ mol dm⁻³, giving pH 7.
  • The ionisation of water is endothermic; increasing temperature shifts the equilibrium to the right, increasing Kw and decreasing the pH of pure water.
  • Even when the pH of pure water is below 7 at higher temperatures, it is still neutral because [H⁺] = [OH⁻].

Strong & Weak Acids

  • A strong acid dissociates almost completely in aqueous solution; the equilibrium lies far to the right and the reaction can be shown as irreversible.
  • Examples of strong acids include HCl, HNO₃ and H₂SO₄ (first ionisation).
  • A weak acid dissociates only partially in aqueous solution; an equilibrium is established and the equilibrium lies to the left.
  • Examples of weak acids include organic acids such as ethanoic acid, HCN, H₂S and H₂CO₃.
  • For a strong acid, pH can be calculated directly from the acid concentration; for a weak acid, Ka is needed to find [H⁺].
  • Acid strength down Group 17 increases as HF < HCl < HBr < HI because the H–X bond becomes longer and weaker, making H⁺ easier to release.
  • Strong and weak acids can be distinguished by pH, electrical conductivity and reactivity with metals.

Neutralisation Reactions

  • A neutralisation reaction occurs when an acid and a base react to form water and a salt: acid + base → salt + water.
  • The essential reaction is H⁺(aq) + OH⁻(aq) → H₂O(l); spectator ions form the salt.
  • The salt produced depends on the acid used: hydrochloric acid forms chlorides, nitric acid forms nitrates, sulfuric acid forms sulfates, ethanoic acid forms ethanoates.
  • Acid + metal → salt + hydrogen; acid + metal oxide → salt + water; acid + metal hydroxide → salt + water.
  • Acid + metal carbonate → salt + water + carbon dioxide; acid + metal hydrogencarbonate → salt + water + carbon dioxide.
  • The enthalpy of neutralisation for a strong acid and strong base is approximately –57 kJ mol⁻¹ because the reaction is essentially the same in each case.

Universal indicator across the pH scale

Universal indicator across the pH scale

pH Curves

  • A pH curve shows how the pH of a solution changes as acid or base is gradually added during a titration.
  • All pH curves have an S-shape; the midpoint of the vertical section is the equivalence point.
  • From a pH curve you can identify the initial pH, the pH at the equivalence point, the volume of base added at the equivalence point, and the pH range of the sharp change.
  • For a strong acid–strong base titration, the pH at the equivalence point is 7.0.
  • During a strong acid–strong base titration, pH changes slowly at first, then sharply near the equivalence point, then slowly again.
  • Calculating pH during a titration involves determining the moles of excess H⁺ or OH⁻ and dividing by the total volume.

Diapos

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  1. 1.Which statement defines a Brønsted–Lowry acid?

    Easy
    • AA species that donates a proton (H⁺)
    • BA species that accepts a proton (H⁺)
    • CA species that releases hydroxide ions (OH⁻) in water
    • DA species that accepts a lone pair of electrons

  2. 2.In the reaction HSO₄⁻(aq) + H₂O(l) ⇌ H₂SO₄(aq) + OH⁻(aq), which species acts as a Brønsted–Lowry base?

    Easy
    • AHSO₄⁻
    • BH₂O
    • CH₂SO₄
    • DOH⁻

  3. 3.In the reaction CO₃²⁻(aq) + H⁺(aq) ⇌ HCO₃⁻(aq), which species acts as a Brønsted–Lowry acid?

    Easy
    • ACO₃²⁻
    • BH⁺
    • CHCO₃⁻
    • DNone of these

  4. 4.Which statement correctly describes the difference between an amphiprotic and an amphoteric species?

    Medium
    • AAll amphoteric species are amphiprotic
    • BAmphiprotic species can donate and accept protons, while amphoteric species can act as both an acid and a base but not necessarily by proton transfer
    • CAmphoteric species can donate and accept protons, while amphiprotic species can act as both an acid and a base but not necessarily by proton transfer
    • DThere is no difference; the terms are interchangeable

  5. 5.Which of the following is the conjugate base of the hydroxide ion, OH⁻?

    Medium
    • AO²⁻
    • BH₂O
    • CH₃O⁺
    • DOH⁻

  6. 6.In the equilibrium CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq), which pair is a conjugate acid–base pair?

    Medium
    • ACH₃COOH and H₂O
    • BH₂O and H₃O⁺
    • CH₂O and CH₃COO⁻
    • DCH₃COO⁻ and H₃O⁺

  7. 7.Which species in the following reaction is acting as a Brønsted–Lowry base? H₂PO₄⁻(aq) + H₂O(l) → HPO₄²⁻(aq) + H₃O⁺(aq)

    Medium
    • AH₂PO₄⁻
    • BH₂O
    • CHPO₄²⁻
    • DH₃O⁺

  8. 8.Which of the following is a strong acid?

    Medium
    • ACH₃COOH
    • BHCN
    • CHNO₃
    • DH₂CO₃

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