Acids, Bases and Salts
Quick Notes
- Acids and bases can be defined by Arrhenius, Brønsted–Lowry, or Lewis concepts.
- Arrhenius: Acids release H+; bases release OH− (only in water).
- Brønsted–Lowry: Acid is a proton donor, base is a proton acceptor.
- Lewis: Acid is electron pair acceptor, base is electron pair donor.
- These definitions expand the range of acid–base behaviour beyond aqueous solutions.
Full Notes
Acid–Base Theories
Over time, different models have been developed to explain acid–base behaviour in broader and more useful ways.
Arrhenius Concept of Acids and Bases
This is the earliest and most familiar theory, focusing on how substances behave in water.
- Acid: Produces H+ ions and increases H+ (or H3O+) concentration in water
Example: HCl (aq) → H+ (aq) + Cl− (aq) - Base: Increases OH− concentration in water
Limitations:
- Only applies to aqueous solutions
- Cannot explain the basic behaviour of substances like NH3, which do not contain OH− but still act as bases
A Note on H+ and H3O+
You’ll often see H+(aq) used to represent the hydrogen ion in solution. Technically, free protons (H+) don’t float around in water — they quickly bond with water molecules to form H3O+(aq) (hydronium ions). A water molecule effectively ‘carries’ the H+ ion.
This means we can write the dissociation of an acid in two ways:
- HCl (aq) → H+ (aq) + Cl− (aq)
- HCl (aq) + H2O (l) → H3O+ (aq) + Cl− (aq)
6.10.2 The Brønsted–Lowry Acids and Bases
This theory extends acid–base ideas beyond water-based reactions by focusing on proton transfer.
- Acid: Proton (H+) donor
- Base: Proton (H+) acceptor
- Forms conjugate acid–base pairs
Example: NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH− (aq)
- NH3 is a Brønsted base (accepts H+)
- H2O is acting as an acid (donates H+)
Conjugate Acid–Base Pair: Formed by the loss or gain of a proton.
Examples: NH4+/NH3 and H2O/OH−
6.10.3 Lewis Acids and Bases
This is the most general theory, explaining acid–base reactions in terms of electron pairs, not just protons.
- Lewis Acid: Electron pair acceptor
- Lewis Base: Electron pair donor
Example: BF3 + NH3 → F3B←NH3
- BF3 is the Lewis acid (accepts lone pair)
- NH3 is the Lewis base (donates lone pair)
This theory does not require H+ ions, making it useful in explaining complex formation, catalysis and organic reaction mechanisms.
Summary
- Arrhenius defines acids and bases by H+ and OH− formation in water.
- Brønsted–Lowry defines acids as proton donors and bases as proton acceptors.
- Lewis defines acids as electron pair acceptors and bases as electron pair donors.
- These models broaden acid–base behaviour beyond aqueous systems.