The Characteristic Properties of Acids and Bases

7.1 The Characteristic Properties of Acids and Bases

Definition and Ion Formation in Aqueous Solutions

• Aqueous solutions of acids contain hydrogen ions (H⁺)
  • In aqueous (water-based) solutions, acids release hydrogen ions (protons) into the solution.
  • Example:
    • HCl(aq) → H⁺(aq) + Cl⁻(aq)
    • HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
  • The presence of H⁺ ions is responsible for acidic properties such as sour taste, ability to turn blue litmus red, and reactivity with metals and carbonates.
• Aqueous solutions of alkalis contain hydroxide ions (OH⁻)
  • Alkalis are soluble bases that release hydroxide ions when dissolved in water.
  • Example:
    • NaOH(aq) → Na⁺(aq) + OH⁻(aq)
    • KOH(aq) → K⁺(aq) + OH⁻(aq)
  • OH⁻ ions are responsible for alkaline properties such as slippery feel, turning red litmus blue, and neutralising acids.

Definitions of Acids and Bases (Brønsted–Lowry Theory)

• Acid: A proton donor (substance that donates H⁺ ions in a reaction).
  • Example: HCl donates H⁺ to OH⁻ in neutralisation.
    • HCl + OH⁻ → Cl⁻ + H₂O
• Base: A proton acceptor (substance that accepts H⁺ ions in a reaction).
  • Example: NH₃ accepts a proton from water to form NH₄⁺:
    • NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
• This definition applies to reactions in both aqueous and non-aqueous solutions.

Bases, Oxides, Hydroxides, and Alkalis

• Bases
  • Metal oxides (e.g., CuO, MgO) or metal hydroxides (e.g., NaOH, Ca(OH)₂) that neutralise acids to form salt and water.
  • May be soluble or insoluble in water.
• Alkalis
  • Soluble bases (hydroxides of Group 1 metals, some Group 2 metals, and ammonium hydroxide).
  • Release OH⁻ ions in aqueous solutions.
  • Examples: NaOH, KOH, Ba(OH)₂, NH₄OH.
• Difference:
  • All alkalis are bases, but not all bases are alkalis.

Characteristic Reactions of Acids

Acid + Metal → Salt + Hydrogen

• General equation:
  • Acid + Metal → Salt + H₂(g)
• Example:
  • 2HCl(aq) + Mg(s) → MgCl₂(aq) + H₂(g)
• Observable changes: effervescence (bubbles of H₂), metal dissolves, exothermic reaction.

Acid + Base → Salt + Water

• Neutralisation reaction.
• Example:
  • H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)
• Used in preparation of soluble salts by reacting insoluble bases with acids.

Acid + Carbonate → Salt + Water + Carbon Dioxide

• General equation:
  • Acid + Carbonate → Salt + H₂O + CO₂(g)
• Example:
  • 2HCl(aq) + Na₂CO₃(s) → 2NaCl(aq) + H₂O(l) + CO₂(g)
• Observable changes: effervescence due to CO₂ gas, which turns limewater milky.

Characteristic Reactions of Bases

Base + Acid → Salt + Water

• Typical neutralisation.
• Example:
  • NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l)

Base + Ammonium Salt → Salt + Ammonia + Water

• Example:
  • NH₄Cl(s) + NaOH(aq) → NaCl(aq) + NH₃(g) + H₂O(l)
• Observable changes: ammonia gas released with pungent smell, turns damp red litmus blue.

Neutralisation Reactions

• Definition: Reaction between an acid and a base to form salt and water.
• Ionic equation for acid + alkali:
  • H⁺(aq) + OH⁻(aq) → H₂O(l)
• Exothermic process – releases heat.

Indicators and Colour Changes

• Litmus:
  • Acidic: turns blue litmus red
  • Alkaline: turns red litmus blue
• Thymolphthalein:
  • Acidic: colourless
  • Alkaline: blue
• Methyl Orange:
  • Acidic: red
  • Alkaline: yellow
  • Neutral: orange

Strong and Weak Acids

• Strong Acid: Completely dissociates in aqueous solution to release maximum H⁺ ions.
  • Example:
    • HCl(aq) → H⁺(aq) + Cl⁻(aq)
    • HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
    • H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)
• Weak Acid: Partially dissociates in aqueous solution, establishing equilibrium between undissociated and dissociated forms.
  • Example:
    • CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)

Examples of Strong Acids and Their Equations

• Hydrochloric acid: HCl(aq) → H⁺(aq) + Cl⁻(aq)
• Nitric acid: HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
• Sulfuric acid: H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)

Examples of Weak Acids and Their Equations

• Ethanoic acid: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
• Carbonic acid: H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)

Comparison of Acidity, Alkalinity, and pH Using Universal Indicator

• pH scale:
  • Acidic: pH < 7
  • Neutral: pH = 7
  • Alkaline: pH > 7
• Colour changes with universal indicator paper:
  • pH 1–3: red (strong acid)
  • pH 4–6: orange/yellow (weak acid)
  • pH 7: green (neutral)
  • pH 8–11: blue (weak alkali)
  • pH 12–14: violet/purple (strong alkali)
• Relative acidity and alkalinity:
  • Higher [H⁺] → lower pH → stronger acid
  • Higher [OH⁻] → higher pH → stronger alkali