Sample Notes: Chemical Energetics

A2 Level Chemistry – Detailed Notes

Chapter 23: Chemical Energetics

23.1 Lattice Energy and Born–Haber Cycles

Key Definitions

• ΔHᵃₜ (Enthalpy change of atomisation):
  Energy required to produce 1 mole of gaseous atoms from the element in its standard state.
  Example:
  ½Cl₂(g) → Cl(g)
• ΔHˡᵃₜᵗ (Lattice energy):
  Energy released when 1 mole of an ionic solid is formed from its gaseous ions.
  Example:
  Na⁺(g) + Cl⁻(g) → NaCl(s)

Electron Affinity

• First Electron Affinity (EA):
  Enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous negative ions.
  Example:
  Cl(g) + e⁻ → Cl⁻(g)
• Factors Affecting EA:
  • Atomic radius: Larger atoms = less attraction for added electron = less negative EA.
  • Nuclear charge: Higher nuclear charge = stronger attraction = more negative EA.
  • Electron repulsion: Second EA is always endothermic due to repulsion.
• Trends:
  • Group 17: EA becomes less negative down the group due to increased atomic size and shielding.
  • Group 16: Similar trend but less negative than Group 17 due to added electron entering a doubly occupied orbital.

Born–Haber Cycle Construction

For ionic solids (e.g., NaCl, MgCl₂)

Steps:

1. Enthalpy of formation (ΔHf)
2. Atomisation of metal
3. Ionisation energy (1st or 2nd IE)
4. Atomisation of non-metal
5. Electron affinity (1st and 2nd EA)
6. Lattice energy (ΔHlatt)

General energy cycle:
ΔHf = ΔHat (metal) + ΔHat (non-metal) + IE + EA + ΔHlatt
→ Rearranged to calculate ΔHlatt if unknown.

Factors Affecting Lattice Energy (ΔHlatt)

• Ionic charge: Greater charge = stronger electrostatic attraction = more negative ΔHlatt.
• Ionic radius: Smaller ions = stronger attraction = more negative ΔHlatt.
• Charge density ∝ charge / radius²

23.2 Enthalpies of Solution and Hydration

Definitions

• ΔHsol (Enthalpy change of solution):
  Enthalpy change when 1 mole of an ionic compound dissolves in water to form an infinitely dilute solution.
  Example:
  NaCl(s) → Na⁺(aq) + Cl⁻(aq)
• ΔHhyd (Enthalpy change of hydration):
  Enthalpy change when 1 mole of gaseous ions becomes hydrated (dissolved in water).
  Example:
  Na⁺(g) → Na⁺(aq)

Energy Cycle for ΔHsol

ΔHsol = –ΔHlatt + ΣΔHhyd
(Where ΔHlatt is endothermic and ΔHhyd values are exothermic)

Cycle:

• Solid ionic lattice → gaseous ions (ΔHlatt)
• Gaseous ions → aqueous ions (ΔHhyd)

Factors Affecting ΔHhyd

• Ionic charge: Higher charge = more exothermic ΔHhyd.
• Ionic radius: Smaller ions = higher charge density = more exothermic ΔHhyd.

23.3 Entropy Change, ΔS

Definition of Entropy

• A measure of the disorder or randomness of a system.
• Symbol: S
  Units: J mol⁻¹ K⁻¹

Predicting the Sign of ΔS

• Changes of State:
  • Solid → Liquid → Gas → ΔS increases
  • Gas → Liquid → Solid → ΔS decreases
• Temperature Increase:
  • Increases particle energy → ΔS increases
• Change in Number of Gaseous Molecules:
  • More moles of gas on product side → ΔS > 0
  • More moles on reactant side → ΔS < 0

Entropy Change of Reaction

ΔS°reaction = ΣS°(products) – ΣS°(reactants)

• Use tabulated standard entropy values.

23.4 Gibbs Free Energy Change, ΔG

Definition

• Gibbs Free Energy indicates whether a reaction is thermodynamically feasible.

Gibbs Equation

ΔG° = ΔH° – TΔS°

• ΔH°: Enthalpy change (kJ mol⁻¹)
• ΔS°: Entropy change (J mol⁻¹ K⁻¹)
• Convert ΔS° to kJ by dividing by 1000

Feasibility

• ΔG < 0 → reaction is feasible/spontaneous
• ΔG > 0 → reaction is not feasible
• ΔG = 0 → reaction is at equilibrium

Effect of Temperature on Feasibility