Arrhenius Theory (1884)

Svante Arrhenius proposed the first modern definitions of acids and bases based on their behaviour in water.

• Arrhenius acid: a substance that produces H⁺ (or H₃O⁺) ions when dissolved in water.
• Arrhenius base: a substance that produces OH⁻ ions when dissolved in water.

Limitation: The Arrhenius model is only applicable to aqueous (water) solutions. It cannot explain why NH₃ — which contains no OH⁻ — acts as a base, since it doesn't dissolve to release OH⁻ directly.

Brønsted-Lowry Theory (1923)

Johannes Brønsted and Thomas Lowry independently proposed a broader, more general theory:

• Brønsted-Lowry acid: a proton (H⁺) donor — any species that can give away a proton.
• Brønsted-Lowry base: a proton (H⁺) acceptor — any species that can receive a proton.

This model does not require water. It explains NH₃ as a base because NH₃ accepts a proton (H⁺) from water, forming NH₄⁺ and leaving OH⁻:

Conjugate Acid-Base Pairs

When an acid donates H⁺, the species that remains is its conjugate base — it differs from the acid by exactly one proton (H⁺). Every acid-base reaction involves two conjugate pairs simultaneously.

• HA donates H⁺ to H₂O → HA is the acid; A⁻ is its conjugate base (lost one H⁺)
• H₂O accepts H⁺ from HA → H₂O is the base; H₃O⁺ is its conjugate acid (gained one H⁺)

Amphoteric Substances

An amphoteric substance can act as EITHER an acid OR a base depending on what it reacts with.

• H₂O: accepts H⁺ from strong acids (behaves as a base); donates H⁺ to strong bases (behaves as an acid).
• HCO₃⁻: acts as an acid (donates H⁺ to give CO₃²⁻) or as a base (accepts H⁺ to give H₂CO₃).
• HSO₄⁻: similarly amphoteric — can donate or accept a proton.

Strong vs Weak Acids and Bases

The strength of an acid or base refers to the degree of ionisation in water — not the concentration.

• Strong acid: essentially complete ionisation in water (→). [H₃O⁺] equals the initial acid concentration.
  Strong acids: HCl, HNO₃, H₂SO₄ (first ionisation), HBr, HI, HClO₄
• Weak acid: partial, reversible ionisation (⇌); an equilibrium is established.
  Ka = [H₃O⁺][A⁻] / [HA] — the larger Ka, the stronger the weak acid.
  Examples: CH₃COOH (Ka = 1.8 × 10⁻⁵); HF (Ka = 6.8 × 10⁻⁴); H₂CO₃ (Ka₁ = 4.3 × 10⁻⁷)
• Strong base: complete dissociation in water. NaOH, KOH, Ca(OH)₂, Ba(OH)₂
• Weak base: partial reaction with water; equilibrium established.
  NH₃ + H₂O ⇌ NH₄⁺ + OH⁻; described by Kb (base dissociation constant)

The pH Scale

pH is a logarithmic measure of the hydrogen ion concentration in solution at 25°C:

• pH < 7 → acidic solution
• pH = 7 → neutral solution (at 25°C)
• pH > 7 → basic/alkaline solution

The ion product of water: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C. This means pH + pOH = 14 always holds at 25°C.

pH Calculations:

• Strong acid: [H₃O⁺] = c(acid) (fully ionised); pH = −log(c)
  Example: 0.01 mol·L⁻¹ HCl → pH = −log(0.01) = 2.0
• Strong base: [OH⁻] = c(base); pOH = −log(c); pH = 14 − pOH
  Example: 0.01 mol·L⁻¹ NaOH → pOH = 2.0 → pH = 12.0
• Weak acid approximation: [H₃O⁺] ≈ √(Ka × c) when Ka << c
  Example: 0.10 mol·L⁻¹ CH₃COOH, Ka = 1.8 × 10⁻⁵
  [H⁺] = √(1.8 × 10⁻⁵ × 0.10) = √(1.8 × 10⁻⁶) = 1.34 × 10⁻³ mol·L⁻¹ → pH = 2.87

Neutralisation

When an acid and a base react, they form a salt and water in an exothermic reaction:

• HCl + NaOH → NaCl + H₂O
• H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O

The reaction is exothermic — heat is released to the surroundings.

Hydrolysis of Salts

The pH of a salt solution depends on the relative strengths of the acid and base that formed it:

Explanation: CH₃COONa — the acetate ion (CH₃COO⁻) is the conjugate base of a weak acid; it accepts H⁺ from water, producing OH⁻ → basic solution. NH₄Cl — the ammonium ion (NH₄⁺) is the conjugate acid of a weak base; it donates H⁺ to water, producing H₃O⁺ → acidic solution.

Acid-Base Indicators

Indicators are weak acids (HIn) whose ionised form (In⁻) has a different colour from the unionised form:

• In acidic solution: equilibrium shifts left → colour 1 predominates
• In basic solution: equilibrium shifts right → colour 2 predominates
• The colour transition range is approximately pKa ± 1