Energetics

• Lattice energy (Δ_lattH) — the energy released when gaseous ions come together to form 1 mole of ionic solid. Always exothermic, always huge and negative.
• Born–Haber cycles — a Hess's Law diagram that breaks the formation of an ionic compound into small steps (atomisation, ionisation, electron affinity, lattice formation) so we can calculate the one step we can't measure directly.
• Key enthalpy definitions — atomisation, first/second electron affinity, ionisation energy — each with its own energy "sign" (endo or exo) that you must know cold.
• Polarisation — real ionic bonds aren't 100% ionic. Small, highly-charged cations distort ("polarise") large anions, adding covalent character — which is why experimental and theoretical lattice energies disagree.
• Enthalpy of solution & hydration — what happens energetically when an ionic solid dissolves in water: breaking the lattice apart (endothermic) and hydrating the free ions (exothermic).
• Trends — how ionic radius and ionic charge affect the size of lattice energy and hydration enthalpy (charge density is the master variable).
• Predicting solubility — using entropy (ΔS_total) alongside enthalpy of solution to explain why some ionic solids dissolve and others don't.

Think of an ionic solid like NaCl as a crowd of positive and negative ions that have "fallen into place" next to each other, held tight by electrostatic attraction. Lattice energy is the enthalpy change when that crowd forms from scratch — starting from ions floating freely as a gas, with no attraction between them at all.

Because oppositely-charged ions want to be near each other, snapping them together into a solid lattice releases a huge amount of energy. That's why lattice energy is always exothermic (a big negative number, often in the hundreds or thousands of kJ mol⁻¹).

Why so negative? The bigger (more negative) the lattice energy, the stronger the ionic bonding — because more energy had to be released to pull the ions together, which means the electrostatic attraction between them was very strong.

Before you can turn an element into ions, you first have to get it into single, separate gas atoms. That's atomisation — and it always costs energy, because you're breaking whatever bonds or forces hold the element together (metallic bonds in a solid metal, or covalent bonds in a diatomic gas like Cl₂).