Acids and Bases: Theories Explained

Chemistry defines acids and bases through several theories. The four main concepts are:

1. Traditional Theory / Concept

2. Arrhenius Theory

3. Brønsted–Lowry Theory

4. Lewis Theory

Traditional Theory

Acids:

- Turn blue litmus paper red

- pH < 7

- Sour taste

- React with bases to form salt + water

- Example: Hydrochloric acid (HCl)

Bases:

- Turn red litmus paper blue

- pH > 7

- Bitter taste

- React with acids to form salt + water

- Example: Sodium Hydroxide (NaOH)

Arrhenius Theory (1884, Svante Arrhenius)

Also called the Theory of Ionisation or Electron Dissociation Theory.

Defines acids and bases based on ion formation in water.

- Acid: A substance that releases H⁺ ions in water.

- Base: Substance that releases OH⁻ ions in water.

Neutralisation Reaction:

Acid + Base → Salt + Water

Hydrogen ions combine with hydroxyl ions to form unionised water molecules.

Limitations:

1. Requires water as a solvent

2. Cannot explain acidity/basicity in non‑aqueous solvents

3. Does not explain the basicity of ammonia (NH₃)

4. Cannot explain the acidity of **BF₃, AlCl₃

5. Fails for the acidity of P‑block oxides (CO₂)

6. Fails for basicity of S‑block oxides (Na₂O)

7. Cannot explain neutralisation without a solvent

Brønsted–Lowry Theory (1923)

Also known as the Proton Theory of Acids and Bases, proposed by Johannes N. Brønsted and Thomas M. Lowry.

- Acid: Proton (H⁺) donor

- Base: Proton (H⁺) acceptor

- Amphoteric species: Can act as both an acid and a base

- Conjugate Base: Species left after an acid donates a proton

- Conjugate Acid: Species formed after a base accepts a proton

Lewis Theory (1923, G.N. Lewis)

Defines acids and bases in terms of electron pair transfer.

- Acid: Molecule/ion that accepts a lone pair of electrons

- Examples: H⁺, NH₄⁺, Na⁺, Cu²⁺, Al³⁺

- Base: Molecule/ion that donates a lone pair of electrons

- Examples: OH⁻, Cl⁻, CN⁻