“Thermochemistry and Energetics”

🟢 Lite — Quick Review (1h–1d)

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“Thermochemistry and Energetics” — Key Facts for NECO SSCE

• Exothermic: releases heat (ΔH < 0) — products more stable than reactants
• Endothermic: absorbs heat (ΔH > 0) — products less stable than reactants
• Enthalpy (ΔH): heat change at constant pressure
• Hess’s Law: ΔH is independent of path — use intermediate steps
• Bond dissociation energy: energy required to break a bond
• ⚡ Exam tip: If ΔH is negative → exothermic → heat released → temperature rises

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🟡 Standard — Regular Study (2d–2mo)

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“Thermochemistry and Energetics” — NECO SSCE Chemistry Study Guide

Key Concepts:

• System vs surroundings: the reacting mixture is the system; everything else is surroundings
• Enthalpy change (ΔH): heat absorbed or released during a reaction at constant pressure
• Standard enthalpy (ΔH°): measured at standard conditions (298 K, 1 atm)
• Exothermic reactions: ΔH < 0 — heat flows out of system to surroundings (e.g., combustion)
• Endothermic reactions: ΔH > 0 — heat flows into system from surroundings (e.g., decomposition)

Enthalpy of Reaction: ΔH = Σ ΔHf (products) − Σ ΔHf (reactants)

Hess’s Law of Constant Heat Summation:

• Total enthalpy change is the same regardless of the route taken
• Useful when direct measurement is difficult
• Can be calculated from: bond energies, formation enthalpies, or combustion enthalpies

Bond Energy Calculations: ΔH = Σ (energy of bonds broken) − Σ (energy of bonds formed)

• Breaking bonds = +energy; forming bonds = −energy

Lattice Energy (for ionic compounds): Energy released when one mole of gaseous ions forms one mole of an ionic solid

• Stronger attraction → higher lattice energy → more negative ΔH

Born-Haber Cycle: A Hess cycle that calculates lattice energy using formation enthalpy, atomisation, ionisation, electron affinity, and sublimation energies

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🔴 Extended — Deep Study (3mo+)

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“Thermochemistry and Energetics” — Deep Dive for NECO SSCE

Detailed Enthalpy Types:

• Enthalpy of formation (ΔHf): enthalpy change when 1 mole of compound forms from its elements
• Enthalpy of combustion (ΔHc): enthalpy change when 1 mole of substance burns in oxygen
• Enthalpy of neutralisation (ΔHn): enthalpy change when 1 mole of H⁺ reacts with 1 mole of OH⁻ (≈ −57.3 kJ/mol for strong acid/base)
• Enthalpy of atomisation: energy to convert elements in standard states to gaseous atoms
• Enthalpy of solution: overall enthalpy change when solute dissolves in solvent
• Enthalpy of hydration: energy released when gaseous ions become aqueous ions
• Enthalpy of fusion (ΔHfus): solid → liquid at melting point
• Enthalpy of vaporisation (ΔHvap): liquid → gas at boiling point

Born-Haber Cycle — Worked Example (NaCl):

1. Na(s) → Na(g): sublimation energy (+107 kJ/mol)
2. Na(g) → Na⁺(g) + e⁻: ionisation energy (+496 kJ/mol)
3. ½ Cl₂(g) → Cl(g): bond dissociation energy (+122 kJ/mol)
4. Cl(g) + e⁻ → Cl⁻(g): electron affinity (−349 kJ/mol)
5. Na⁺(g) + Cl⁻(g) → NaCl(s): lattice energy (unknown, to be calculated)
6. Na(s) + ½ Cl₂(g) → NaCl(s): ΔHf = −411 kJ/mol

Using Hess’s Law: ΔHf = ΔHsub + ΔHI + ½ΔHdiss + ΔHEA + ΔH lattice −411 = 107 + 496 + 122 + (−349) + ΔHlattice ΔHlattice = −787 kJ/mol

Kirchhoff’s Law: How enthalpy changes with temperature: ΔH(T₂) = ΔH(T₁) + ∫ΔCp dT Where ΔCp = Cp(products) − Cp(reactants)

Application in Industry:

• Haber Process (NH₃ synthesis): exothermic (ΔH = −92 kJ/mol) — low temperature favours yield but slows rate; optimal compromise at 400–500°C with catalyst
• Contact Process (H₂SO₄ production): exothermic — conditions optimised for equilibrium position
• Cement production: highly endothermic reactions in kiln require large fuel input

Energy Profile Diagrams:

• Activation energy (Ea): energy barrier — must be overcome for reaction
• Catalysts lower Ea by providing alternative pathway
• ΔH shown as difference between reactant and product energy levels
• Exothermic: products lower than reactants (ΔH < 0)
• Endothermic: products higher than reactants (ΔH > 0)