Lattice Energy & Enthalpy Change of Atomisation
- 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 (ΔHatꝋ) 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 ΔHatꝋ 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 (ΔHatꝋ) 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) ΔHatꝋ = +107 kJ mol -1
Worked example
Write the equations for the standard enthalpy change of atomisation, (ΔHatꝋ) for:
1. Potassium
2. Mercury
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 (ΔHlattꝋ) is the enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions (under standard conditions)
- The ΔHlattꝋ 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 ΔHlattꝋ 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 ΔHlattꝋ of an ionic compound cannot be determined directly by one single experiment
- Multiple experimental values and an energy cycle are used to find the ΔHlattꝋ of ionic compounds
- The lattice energy (ΔHlattꝋ) 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
Write the equations which represent the lattice energy of:
1. Magnesium oxide
2. Lithium chloride
Answer 1:
- Mg2+ (g) + O2– (g) → MgO (s)
Answer 2:
- Li+ (g) + Cl– (g) → LiCl (s)
Exam Tip
Make sure the correct state symbols are stated when writing these equations – it is crucial that you use these correctly throughout this entire topic
