Key A-level Energetics, Electron Affinity, Born–Haber & Thermochemistry

Lattice Energy (ΔHlatt^⊖)
• Enthalpy change when 1 mol of ionic solid forms from gaseous ions under standard conditions (298 K, 10^5 Pa).
• Always negative (exothermic).
• Stronger (more –ve) ΔHlatt ⇒ greater lattice stability.

Atomisation (ΔH_at^⊖)
• 1 mol of gaseous atoms formed from element in standard state.
• Always positive (bond-breaking).

Electron Affinity
• 1st EA (ΔH_ea1^⊖): X(g)+e⁻→X⁻(g) (usually –ve).
• 2nd/3rd EA always +ve (adding e⁻ to anion).

Factors Affecting EA
• ↑Effective nuclear charge → EA more –ve.
• ↓Atomic radius → EA more –ve.
• Shielding, orbital type, electron configuration stability.

Born–Haber Cycles
• ΔHf^⊖ = Σ(steps to ions) + ΔHlatt^⊖.
• ΔHlatt^⊖ = ΔHf^⊖ – Σ(ionisation, atomisation, EA).

Lattice-Energy Trends
• Smaller ions / higher charge → more exothermic ΔH_latt.
• Polarisation: small, highly charged cations + large, highly charged anions increase covalent character.

Solution & Hydration Enthalpy
• ΔHsol^⊖ = ΔHhyd(cation)+ΔHhyd(anion)–ΔHlatt.
• Hydration exothermic; magnitude ↑ with charge density.

Entropy (S)
• Measure of disorder (J K⁻¹).
• S(g)≫S(l)>S(s); increases with T, number of particles, dissolution.
• ΔStotal^⊖ = ΔSsys^⊖ + ΔSsurr^⊖; ΔSsurr^⊖ = –ΔH_rxn^⊖/T.

Gibbs Free Energy
• ΔG = ΔH – TΔSsys (kJ mol⁻¹). • Spontaneous when ΔGcell for electrochemical cells.

Electrochemistry
• Standard electrode potential E^⊖ measured vs SHE (0 V).
• Ecell = Ered(cathode) – E_red(anode).
• Greater E^⊖ ⇒ species stronger oxidising agent.
• Nernst eqn: E = E^⊖ + (RT/zF)ln([Ox]/[Red]).

Group-Trends
• EA becomes less exothermic down groups 16 & 17 (except anomalous F, O).
• Carbonates/nitrates: thermal stability ↑ down Group 2 (larger cations polarise less).

Lattice/Hydration Correlations
• ΔHlatt and ΔHhyd both become less exothermic down a group, but ΔHlatt falls faster for sulfates ⇒ solubility ↓; ΔHhyd falls faster for hydroxides ⇒ solubility ↑.

Born–Haber Worked Examples
• LiF: ΔHlatt = –1049 kJ mol⁻¹. • NaCl: ΔHlatt ≈ –788 kJ mol⁻¹.

Sample Hydration Calc
• Mg²⁺ hydration: ΔH_hyd = –1919 kJ mol⁻¹ (using lattice & solution data).

Faraday’s Laws (Electrolysis)
• Q=It ; m = (It M)/(nF).
• Example: 1.80 A, 45 min deposits 5.44 g Ag.

Buffer Solutions
• pH = pK_a + log([A⁻]/[HA]) (Henderson–Hasselbalch).

Solubility Product (Ksp)
• For AxBy: Ksp=[A^{y+}]^x[B^{x−}]^y.
• Precipitation when ionic product > K_sp; common-ion effect reduces solubility.

Transition-Metal Chemistry (brief)
• Variable oxidation states, complex formation, coloured compounds, catalytic activity.
• Ligand field splits d-orbitals (ΔE = h ν -> colour).
• Ligand substitution, redox titrations (MnO₄⁻/Fe²⁺, etc.).