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Acids and Bases: Arrhenius, Bronsted-Lowry, Lewis Concepts

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Three theories define acids and bases: Arrhenius (1887), Brønsted–Lowry (1923), and Lewis (1923). Each is wider than the last — Arrhenius covers aqueous H+ and OH−, Brønsted–Lowry covers proton transfer in any solvent, and Lewis covers electron-pair transfer with no protons at all.

The wider theories matter because many real reactions — ammonia grabbing a proton, or BF3 accepting an electron pair — cannot be described by the narrower ones. Exams test whether you can pick the right theory and identify conjugate acid–base pairs in any proton-transfer equation.

Key Concepts

  • Arrhenius acid: produces H+ in water (HCl → H+ + Cl−); Arrhenius base: produces OH− in water (NaOH). Works only in aqueous solution.
  • Brønsted–Lowry acid: proton (H+) donor; base: proton acceptor. A conjugate pair differs by exactly one H+, e.g. NH4+/NH3 and H2O/OH−.
  • Amphiprotic substances can act as acid or base: H2O, HCO3−, HSO4−.
  • Lewis acid: electron-pair acceptor (BF3, AlCl3, H+, metal cations); Lewis base: electron-pair donor (NH3, H2O, Cl−). The widest definition.
  • Every Arrhenius acid is a Brønsted–Lowry acid, and every Brønsted–Lowry acid–base reaction is a Lewis one — but not vice versa.

Formulas

  • Ka × Kb = Kw for a conjugate acid–base pair, with Kw = 1.0 × 10^−14 at 25 °C.
  • pH = −log[H+]; pKa + pKb = 14 for a conjugate pair at 25 °C.
  • The stronger the acid (large Ka), the weaker its conjugate base (small Kb).

Worked Example

Consider H2CO3 + H2O ⇌ HCO3− + H3O+ and identify every role:

  1. H2CO3 donates a proton → Brønsted–Lowry acid; H2O accepts it → Brønsted–Lowry base.
  2. Conjugate pairs (differ by one H+): H2CO3/HCO3− and H3O+/H2O.
  3. Under Lewis theory: H2O donates an electron pair to H+ → Lewis base; H+ accepts it → Lewis acid.
  4. Contrast with BF3 + NH3 → F3B–NH3: no proton moves, so Arrhenius and Brønsted–Lowry are silent — only Lewis theory (BF3 acid, NH3 base) explains it.

Common Misconceptions

  • "All acids contain hydrogen." Lewis acids like BF3 and AlCl3 contain no H at all — acidity is about electron pairs, not hydrogen atoms.
  • "Strong acid means concentrated acid." Strength is about extent of ionisation (Ka); concentration is about moles per litre. Dilute HCl is still a strong acid.
  • "Neutralisation always gives pH 7." Only for strong acid + strong base. Salts of weak acids or bases hydrolyse: CH3COONa solution is basic, NH4Cl solution is acidic.

Real-World Applications

  • Buffers: the H2CO3/HCO3− pair holds human blood at pH 7.4 — a direct application of conjugate-pair chemistry.
  • Titrations: choosing indicators and predicting equivalence-point pH needs all three theories plus Ka/Kb.
  • Lewis-acid catalysis: AlCl3 catalyses Friedel–Crafts reactions by accepting electron pairs.

Think of the three theories as nested boxes: Arrhenius inside Brønsted–Lowry inside Lewis. Learn to spot the proton donor and acceptor first, then check whether electron-pair language explains reactions where no proton moves at all.

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