Brønsted-Lowry Terminology
- Brønsted-Lowry acid. A proton (H⁺) donor
- Brønsted-Lowry base. A proton (H⁺) acceptor
- Conjugate acid-base pair. Two species that differ by exactly one proton (H⁺)
- Amphiprotic. A substance that can both donate and accept a proton (e.g. H₂O, HCO₃⁻)
- Amphoteric. A broader term for any species that can react as both acid and base (includes amphiprotic species plus species like Al₂O₃)
Proton Transfer
According to Brønsted-Lowry theory (1923), acid-base reactions involve the transfer of a proton (H⁺) from an acid to a base. In aqueous solutions, free protons do not exist independently. They form coordination bonds with water to form the hydronium ion (H₃O⁺). The notations H⁺(aq) and H₃O⁺(aq) are used interchangeably.
Example 1. Strong acid in water:
HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
HCl donates a proton (acid); H₂O accepts it (base).
Example 2. Weak base in water:
NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
NH₃ accepts a proton (base); H₂O donates one (acid). Note the reversible arrow. NH₃ is a weak base.
Proton Transfer: HCl + H₂O
Identifying Conjugate Pairs
A conjugate acid-base pair is two species that differ by exactly one proton. In any acid-base reaction, there are always two conjugate pairs on opposite sides of the equation.
Step-by-Step: Identifying Conjugate Pairs
- Identify which reactant loses H⁺ (the acid) and which gains H⁺ (the base).
- Find the conjugate base: remove one H⁺ from the acid and decrease its charge by 1. E.g. \(\text{H}_2\text{SO}_4 \rightarrow \text{HSO}_4^-\)
- Find the conjugate acid: add one H⁺ to the base and increase its charge by 1. E.g. \(\text{NH}_3 \rightarrow \text{NH}_4^+\)
- Pair them up: Acid ↔ Conjugate Base, Base ↔ Conjugate Acid.
Amphiprotic vs Amphoteric
These terms are often confused but have an important distinction:
- Amphoteric: Any species that can react as both an acid and a base (broad term, includes Lewis theory).
- Amphiprotic: A specific subset. Species that can both donate and accept a proton under Brønsted-Lowry theory.
All amphiprotic species are amphoteric, but not all amphoteric species are amphiprotic (e.g. Al₂O₃ is amphoteric but has no H atoms to donate).
Common Amphiprotic Species
- Water (\(\text{H}_2\text{O}\)). Donates H⁺ → OH⁻ or accepts H⁺ → H₃O⁺
- Hydrogencarbonate (\(\text{HCO}_3^-\)). Donates H⁺ → CO₃²⁻ or accepts H⁺ → H₂CO₃
- Hydrogen sulfate (\(\text{HSO}_4^-\)). Donates H⁺ → SO₄²⁻ or accepts H⁺ → H₂SO₄
- Amino acids. The amino group (−NH₂) accepts H⁺; the carboxyl group (−COOH) donates H⁺ → forms a zwitterion
Neutralisation & Acid Reactions
Neutralisation involves the transfer of a proton from an acid to a base, forming a salt and usually water.
| Reaction Type | Products | Example |
|---|---|---|
| Acid + Metal hydroxide | Salt + Water | HCl + NaOH → NaCl + H₂O |
| Acid + Metal oxide | Salt + Water | H₂SO₄ + CuO → CuSO₄ + H₂O |
| Acid + Carbonate | Salt + Water + CO₂ | 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ |
| Acid + Hydrogencarbonate | Salt + Water + CO₂ | HCl + NaHCO₃ → NaCl + H₂O + CO₂ |
| Acid + Reactive metal | Salt + H₂ (redox!) | Mg + 2HCl → MgCl₂ + H₂ |
Note: Acid + reactive metal is technically a redox reaction (electron transfer), not a neutralisation. Always include state symbols in your exam answers!
Common Pitfalls in Acid-Base Theory
- Amphoteric ≠ Amphiprotic: Al₂O₃ is amphoteric but not amphiprotic (it has no H atoms to donate).
- Forgetting charge adjustments: The conjugate base of H₂PO₄⁻ is HPO₄²⁻, not HPO₄⁻.
- Misidentifying the base: In NaOH(aq), only the OH⁻ ion acts as the Brønsted-Lowry base, not the Na⁺ spectator ion.
- "Neutralisation = pH 7": Not always! The equivalence point pH depends on the relative strengths of the parent acid and base.
- State symbols: IB mark schemes are strict - always include (s), (l), (aq), (g) in neutralisation equations.
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