15.1.1 Key Enthalpy Terms

Ionisation energy

• The ionisation energy (ΔH_IE^ꝋ) is the standard enthalpy change that occurs on the removal of 1 mole of electrons from 1 mole of gaseous atoms or positively charged ions
• Ionisation energy is always endothermic as energy is need to overcome the attraction between an electron and the nucleus
• The first ionisation energy (ΔH_IE1^ꝋ) is the energy required to remove one mole of electrons from 1 mole of gaseous atoms of an element to form 1 mole of 1+ ions in the gaseous phase

ΔH_IE1^ꝋAl (g) → Al⁺ (g) + e^–          ΔH_IE1^ꝋ = +577 kJ mol^-1

• The second ionisation energy (ΔH_IE2^ꝋ) is the energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to form 1 mole of 2+ ions in the gaseous phase

ΔH_IE2^ꝋAl⁺ (g) → Al²⁺ (g) + e^–      ΔH_IE2^ꝋ = +1820 kJ mol^-1

Enthalpy of Atomisation

• The enthalpy of atomisation (ΔH_at^ꝋ) is the standard enthalpy change that occurs on the formation of 1 mole of separate gaseous atoms an element in its standard state
• The ΔH_at^ꝋ is always endothermic as energy is always required to break any bonds between the atoms in the element or to break the element into its gaseous atoms
  • Since this is always an endothermic process, the enthalpy change will always have a positive value

Na (s) → Na (g)           ΔH_at^ꝋ = +108 kJ mol ^-1

½Cl (g) → Cl (g)          ΔH_at^ꝋ = +122 kJ mol ^-1

Electron Affinity

• The electron affinity (ΔH_EA^ꝋ) of an element is the energy change when 1 mole of electrons is gained by 1 mole of gaseous atoms of an element to form 1 mole of gaseous ions under standard conditions
• For example, the first electron affinity of chlorine is:

Cl (g)+ e^– → Cl^– (g)          ΔH_EA^ꝋ = -364 kJ mol^-1

• The first electron affinity is always exothermic as energy is released when electrons are attracted to the atoms
• However, the second electron affinity of an element can be endothermic as illustrated by oxygen:

O^– (g) + e^– → O^2- (g)          ΔH_EA^ꝋ = +844 kJ mol^-1

• This is because a large force of repulsion must be overcome between the negatively charged ion and the second electron requiring a large input of energy

Lattice Enthalpy

• The lattice enthalpy (ΔH_lat^ꝋ) is defined as the standard enthalpy change that occurs on the formation of 1 mole of gaseous ions from the solid lattice
• The ΔH_lat^ꝋ is always endothermic as energy is always required to break any bonds between the ions in the lattice
  • Since this is always an endothermic process, the enthalpy change will always have a positive value

NaCl (s) → Na⁺ (g) + Cl^- (g)     ΔH_lat^ꝋ = +790 kJ mol ^-1

Enthalpy of Solution 

• The standard enthalpy change of solution (ΔH_sol^ꝋ) is the enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form an infinitely dilute solution
• The symbol (aq) is used to show that the solid is dissolved in sufficient water
• ΔH_sol^ꝋ can be exothermic (negative) or endothermic (positive)

LiBr (s) → LiBr (aq)    ΔH_sol^ꝋ = -48.8 kJ mol ^-1

KCl (s) → KCI (aq)      ΔH_sol^ꝋ = +17.2 kJ mol ^-1

CaCl₂ (s) → CaCl₂ (aq)       ΔH_sol^ꝋ = -82.8 kJ mol ^-1

Enthalpy of Hydration

• The standard enthalpy change of hydration (ΔH_hyd^ꝋ) is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution

Mg²⁺ (g) → Mg²⁺ (aq)   ΔH_hyd^ꝋ = -1963 kJ mol ^-1

Br^- (g) → Br^- (aq)          ΔH_hyd^ꝋ = -328 kJ mol ^-1

• Hydration enthalpies are the measure of the energy that is released when there is an attraction formed between the ions and water molecules
• Hydration enthalpies are exothermic
• The term solvation is used in place of hydration if water has been replaced by another solvent
• When an ionic solid dissolves in water, positive and negative ions are formed
• Water is a polar molecule with a δ- oxygen (O) atom and δ+ hydrogen (H) atoms which will form ion-dipole attractions with the ions present in the solution
• The oxygen atom in water will be attracted to the positive ions and the hydrogen atoms will be attracted to the negative ions

The polar water molecules will form ion-dipole bonds with the ions in solution causing the ions to become hydrated