Enthalpy

(a) Enthalpy Changes

1. Enthalpy Change of Atomisation (Δ_atH)

The enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state.

Always endothermic (+ve) because energy is required to break bonds.

Example: ½ Cl₂ (g) → Cl (g)

2. Enthalpy Change of Lattice Formation (Δ_latticeH)

The enthalpy change when 1 mole of an ionic solid is formed from its gaseous ions.

Always exothermic (–ve) due to the strong electrostatic attraction between oppositely charged ions.

Example: Na⁺ (g)+ Cl⁻ (g) → NaCl (s)

3. Enthalpy Change of Lattice Breaking (Lattice Dissociation, Δ_latticeH)

The enthalpy change required to break 1 mole of an ionic lattice into its gaseous ions.

Always endothermic (+ve) because energy is needed to overcome ionic bonds.

Example: NaCl(s) → Na⁺ (g) + Cl⁻ (g)

4. Enthalpy Change of Hydration (Δ_hydH)

The enthalpy change when 1 mole of gaseous ions dissolves in water to form an aqueous solution.

Always exothermic (–ve) due to attraction between ions and polar water molecules.

Example: Na⁺ (g) → Na⁺ (aq)

5. Enthalpy Change of Solution (Δ_solH)

The enthalpy change when 1 mole of an ionic solid dissolves in water to form an aqueous solution.

Can be exothermic or endothermic depending on the ionic compound.

Example: NaCl(s) → Na⁺ (aq) + Cl⁻ (aq)

(b) Solubility of Ionic Compounds in Water

The enthalpy change of solution (Δ_solH) determines whether an ionic compound dissolves.

It depends on the balance between:

1. Lattice Breaking (Lattice Dissociation, Δ_latticeH): Energy required to separate ions (endothermic).

2. Hydration Enthalpy (Δ_hydH): Energy released when ions interact with water (exothermic).

If Δ_hydH > Δ_latticeH, Δ_solH is negative (exothermic) → compound dissolves easily.

If Δ_hydH < Δ_latticeH, Δ_solH is positive (endothermic) → compound is less likely to dissolve.

Smaller, highly charged ions have higher hydration enthalpies → more soluble salts.

(c) Born-Haber Cycle for Ionic Compounds

A Born-Haber cycle is a thermochemical cycle used to calculate lattice enthalpy.

Steps involved in forming an ionic compound (e.g., NaCl):

1. Enthalpy of Atomisation (Δ_atH): Formation of gaseous atoms from elements.
   ½ Cl₂ (g) → Cl (g)

2. Ionisation Energy (IE): Energy needed to remove an electron from the metal.
   Na(g) → Na⁺ (g) + e⁻

3. Electron Affinity (EA): Energy change when a non-metal gains an electron.
   Cl(g) + e⁻ → Cl⁻ (g)

4. Lattice Formation Enthalpy (Δ_latticeH): Formation of an ionic lattice from gaseous ions.
   Na⁺ (g) + Cl⁻ (g) → NaCl(s)

The cycle can be used to determine unknown enthalpies using Hess’s Law.

(d) Exothermicity or Endothermicity of Δ_fHϴ and Compound Stability

Enthalpy of Formation (Δ_fHϴ): The enthalpy change when 1 mole of a compound is formed from its elements in their standard states.

Exothermic formation (–ve Δ_fHϴ) → More stable compound (e.g., NaCl).

Endothermic formation (+ve Δ_fHϴ) → Less stable compound (e.g., N₂O).

The more negative Δ_fHϴ, the stronger the bonds and the more thermodynamically stable the compound.