Brønsted–Lowry Acid-Base Theory

Introduction to the Brønsted–Lowry Theory

• Brønsted–Lowry Acid: A substance that donates protons (hydrogen ions/H⁺).
• Brønsted–Lowry Base: A substance that accepts protons.

This theory provides insight into reactions beyond aqueous environments, extending its relevance to non-aqueous reactions.

Historical Context

• Arrhenius Theory: Concentrates solely on aqueous solutions with acids releasing H⁺ and bases releasing OH⁻.
• Lewis Theory: Involves reactions with electron pairs.
• Brønsted–Lowry Theory: Broadens the scope to include reactions in solvents like ammonia and methanol.

Base Dissociation and Dissociation Constants

Base Dissociation in Water

• Base Dissociation: The process by which bases dissociate in water to yield hydroxide ions (OH⁻).
• Example: Ammonia (NH₃) reacting with water: $\mathrm{NH}_3 + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{OH}^-$

Define Kb and pKb

• Kb (Base Dissociation Constant): Represents the strength of a base in an aqueous solution.
  • A higher Kb value indicates a stronger base.

• pKb: The logarithmic form of Kb, where a lower pKb signifies stronger bases.

$\mathrm{pK}_b = -\log_{10}(K_b)$

Mathematical Representation

• Formula: $K_b = \frac{[\mathrm{OH}^-][\mathrm{HB}^+]}{[\mathrm{B}]}$
• Calculation Example for NH₃:
  • Initial concentrations: [NH₄⁺] = 0.01 M, [OH⁻] = 0.01 M, [NH₃] = 0.1 M.
  • Calculation: $K_b = \frac{0.01 \times 0.01}{0.1} = 1.0 \times 10^{-3}$

Comparing Kb and Ka

• Understand their relationship in chemical equilibria.
• Ammonia (base): $\mathrm{NH}_3 + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{OH}^-$
• Acetic acid (acid): $\mathrm{CH}_3\mathrm{COOH} + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{CH}_3\mathrm{COO}^- + \mathrm{H}^+$

Exploring Neutralisation Reactions

Neutralisation Basics

• Neutralisation reactions: Processes where acids and bases produce water and salt.
• Equivalence Point: The stage at which the moles of H⁺ equal the moles of OH⁻, achieving complete neutralisation.
• Example Process:
  • HCl dissociates into H⁺ and Cl⁻.
  • NaOH dissociates into Na⁺ and OH⁻.
  • H⁺ and OH⁻ form water, while Na⁺ and Cl⁻ form salt (NaCl).

Visualising and Analysing Neutralisation

• Titration Curves: Display pH variations during neutralisation.

  

• pH Indicators: Phenolphthalein changes from pink to clear upon neutralisation.

  

Knowledge Check

• Exercise 1: Calculate Kb for ammonia, [NH₄⁺] = 0.01 M, [OH⁻] = 0.01 M, [NH₃] = 0.1 M. Solution: Using the formula $K_b = \frac{[\mathrm{OH}^-][\mathrm{NH}_4^+]}{[\mathrm{NH}_3]}$, we substitute the known values: $K_b = \frac{0.01 \times 0.01}{0.1} = \frac{0.0001}{0.1} = 1.0 \times 10^{-3}$

• Exercise 2: Given $K_b = 1.8 \times 10^{-5}$, find the $\mathrm{pK}_b$. Solution: Using the formula $\mathrm{pK}_b = -\log_{10}(K_b)$: $\mathrm{pK}_b = -\log_{10}(1.8 \times 10^{-5}) = -(-4.7447) = 4.74$

Glossary

• Kb (Base Dissociation Constant): A measure of a base's capability to dissociate in water.
• pKb: A logarithmic representation to compare base strength.
• Dissociation: The breakdown of a compound into ions in solution.
• Resonance: The distribution of electron density that stabilises a molecule.
• Substituents: Atoms or groups attached to a molecule affecting its chemical characteristics.