20.1    Introduction to Lattice Energy

What is lattice energy?

 

• In a solid ionic crystal lattice, the ions are bonded by strong ionic bonds between them. These forces are only completely broken when the ions are in gaseous state

 

• Lattice energy(or lattice enthalpy) is the enthalpy change when one mole of solid ionic lattice is formed from its scattered gaseous ions

 

• Lattice energy is always negative. This is because energy is always released when bonds are formed

 

• Use sodium chloride, NaCl as an example. It can be described as the enthalpy change when one mole of sodium chloride is formed from its scattered gaseous sodium and chloride ions. The value is -787 kJ mol⁻¹.

 

• Lattice energy is a good indication of the strength ionic bonds. The higher the magnitude of lattice energy, the stronger the ionic bond formed. This is because more energy is released when strong bonds are formed

 

 

Factors affecting lattice energy

 

• There are two factors which govern the magnitude of lattice energy:
  i) Charge on the ions.
  ii) Radius of the ions.
• • The higher the charge on the ion, the higher the lattice energy
  • This is because ions of higher charge have stronger attraction towards each Hence more energy is released when bonds are formed between them.
  • For example, magnesium oxide, MgO has a higher lattice energy than sodium chloride. This is because Mg²⁺ and O²⁻ have higher charge than Na⁺ and Cl⁻ ions
• • The larger the ionic radius, the lower the lattice energy
  • This is because the attraction between ions gets weaker as the distance between them increases. Hence less energy is released when bonds are formed between them
  • For example, the lattice energy of sodium halides gets lower from sodium fluoride to sodium iodide. This is because as the size of the halide ion increases, the distance between sodium and halide ions increases

 

Electron affinity

• The first electron affinity, ΔHea1 is the energy released when one mole of electrons is added to each atom in one mole of the atoms of  the element to form one mole of gaseous 1- ions
  For example:
  Cl(g) + e⁻ → Cl⁻(g)             ; ΔHea1 = -349 kJ mol⁻¹

• The use of electron affinities are almost always confined to Group 6 and 7 elements.

 

• Generally, first electron affinities are exothermic, this is because there is an attraction between the positive nucleus and the incoming electron. Energy is released when this attraction is set up

 

• Electron affinity is a measure of the attraction between the incoming electron and the nucleus – the stronger the attraction, the more energy is released.

 

• The factors that govern the magnitude of electron affinity is the same as the ones for ionisation energy – nuclear charge, distance between nucleus and incoming electron and shielding effect of  electrons.

 

• Going down a Group, the magnitude of first electron affinity decreases. This is because the distance between incoming electron and the nucleus increases, therefore the attractive force between them is less. The increasing nuclear charge is offset by the increasing shielding effect by inner electrons

 

• Take Group 7 as an example, the elements follow the pattern described, except for fluorine. This is because fluorine is a very small atom with high electron A repulsion between electrons is created hence the attraction between nucleus  and incoming electrons is less. In Group 6, oxygen and the rest follow the exact same trend.

 

• Group 6 elements have a lower electron affinity compared to their corresponding Group 7 elements. This is because the nuclear charge of Group 6 elements is lower, less attraction is set up between the nucleus and outer electrons, therefore less energy is released

• The second electron affinity, ΔHea2 is the energy released when one mole of electrons is added to each gaseous 1- ions of to form one mole of gaseous 2- ions.
  For example:
  O⁻(g) + e⁻ → O²⁻(g)           ;  ΔHea1 = +844 kJ mol⁻¹

 

• Second electron affinities are usually endothermic, this is because energy is required to add electrons into an already negative region due to the repulsion between electrons

 

Recall some enthalpy changes

• Standard enthalpy change of formation, ΔH°f is the enthalpy change when one mole of a compound is formed from its elements under standard  condition
  The reactants and products must be in their standard states.
  2Fe(s) + 1½O2(g) → Fe2O3(s)    ;  ΔH°f [ Fe2O3(s) ] = -824.5 kJ mol⁻¹

 

• Standard enthalpy change of atomisation, ΔH°at is the enthalpy change when one mole of gases atoms is formed from its element under standard conditions
  ½H2(g) → H(g)   ;  ΔH°at [ ½H2 ] = +218 kJ mol⁻¹