Lattice Energy & Enthalpy Change of Atomisation (CIE A Level Chemistry)

• Enthalpy change (ΔH) refers to the amount of heat energy transferred during a chemical reaction, at a constant pressure

Enthalpy change of atomisation

• The standard enthalpy change of atomisation (ΔH_at^ꝋ) is the enthalpy change when 1 mole of gaseous atoms is formed from its element under standard conditions
  • Standard conditions in this syllabus are a temperature of 298 K and a pressure of 101 kPa

• The ΔH_at^ꝋ is always endothermic as energy is always required to break any bonds between the atoms in the element, to break the element into its gaseous atoms
  • Since this is always an endothermic process, the enthalpy change will always have a positive value

• Equations can be written to show the standard enthalpy change of atomisation (ΔH_at^ꝋ) for elements
• For example, sodium in its elemental form is a solid
• The standard enthalpy change of atomisation for sodium is the energy required to form 1 mole of gaseous sodium atoms:

Na(s) → Na(g) ΔH_at^ꝋ = +107 kJ mol ^-1

Worked example: Writing equations for the standard enthalpy change of atomisation

Answer

Answer 1: Potassium in its elemental form is a solid, therefore the standard enthalpy change of atomisation is the energy required to form 1 mole of K(g) from K(s)

K(s) → K(g) 

Answer 2: Mercury in its elemental form is a liquid, so the standard enthalpy change of atomisation of mercury is the energy required to form 1 mole of Hg(g) from Hg(l)

Hg(l) → Hg(g)

Lattice energy

• The lattice energy (ΔH_latt^ꝋ) is the enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions (under standard conditions)
• The ΔH_latt^ꝋ is always exothermic, as when ions are combined to form an ionic solid lattice there is an extremely large release of energy
  • Since this is always an exothermic process, the enthalpy change will always have a negative value
  • Because of the huge release in energy when the gaseous ions combine, the value will be a very large negative value

• The large negative value of ΔH_latt^ꝋ suggests that the ionic compound is much more stable than its gaseous ions
  • This is due to the strong electrostatic forces of attraction between the oppositely charged ions in the solid lattice
  • Since there are no electrostatic forces of attraction between the ions in the gas phase, the gaseous ions are less stable than the ions in the ionic lattice
  • The more exothermic the value is, the stronger the ionic bonds within the lattice are

• The ΔH_latt^ꝋ of an ionic compound cannot be determined directly by one single experiment
• Multiple experimental values and an energy cycle are used to find the ΔH_latt^ꝋ of ionic compounds
• The lattice energy (ΔH_latt^ꝋ) of an ionic compound can be written as an equation
  • For example, magnesium chloride is an ionic compound formed from magnesium (Mg²⁺) and chloride (Cl^-) ions
  • Since the lattice energy is the enthalpy change when 1 mole of magnesium chloride is formed from gaseous magnesium and chloride ions, the equation for this process is:

Mg²⁺(g) + 2Cl^-(g) → MgCl₂(s)

Worked Example: Writing equations for lattice energy

Answer

Answer 1: Mg²⁺(g) + O^2-(g) → MgO(s)

Answer 2: Li⁺(g) + Cl^-(g) → LiCl(s)