The Effect of Ionic Charge & Radius on Enthalpy (OCR A Level Chemistry A): Revision Note

The Effect of Ionic Charge & Radius on Enthalpy

Factors affecting lattice enthalpy

• The two key factors which affect lattice energy, ΔH_latt^ꝋ, are the charge and radius of the ions that make up the crystalline lattice

Ionic radius

• The lattice energy becomes less exothermic as the ionic radius of the ions increases

• This is because the charge on the ions is more spread out over the ion when the ions are larger

• The ions are also further apart from each other in the lattice

  • The attraction between ions is between the centres of the ions involved, so the bigger the ions the bigger the distance between the centre of the ions

• Therefore, the electrostatic forces of attraction between the oppositely charged ions in the lattice are weaker

• For example, the lattice energy of caesium fluoride (CsF) is less exothermic than the lattice energy of potassium fluoride (KF)

  • Since both compounds contain a fluoride (F^-) ion, the difference in lattice energy must be due to the caesium (Cs⁺) ion in CsF and potassium (K⁺) ion in KF

  • Potassium is a Group 1 and Period 4 element

  • Caesium is a Group 1 and Period 6 element

  • This means that the Cs⁺ ion is larger than the K⁺ ion

  • There are weaker electrostatic forces of attraction between the Cs⁺ and F^- ions compared to K⁺ and F^- ions

  • As a result, the lattice energy of CsF is less exothermic than that of KF

The lattice energies get less exothermic as the ionic radius of the ions increases

Ionic charge

• The lattice energy gets more exothermic as the ionic charge of the ions increases

• The greater the ionic charge, the higher the charge density

• This results in stronger electrostatic attraction between the oppositely charged ions in the lattice

• As a result, the lattice energy is more exothermic

• For example, the lattice energy of calcium oxide (CaO) is more exothermic than the lattice energy of potassium chloride (KCl)

  • Calcium oxide is an ionic compound which consists of calcium (Ca²⁺) and oxide (O^2-) ions

  • Potassium chloride is formed from potassium (K⁺) and chloride (Cl^-) ions

  • The ions in calcium oxide have a greater ionic charge than the ions in potassium chloride

  • This means that the electrostatic forces of attraction are stronger between the Ca²⁺ and O^2-compared to the forces between K⁺ and Cl^-

  • Therefore, the lattice energy of calcium oxide is more exothermic, as more energy is released upon its formation from its gaseous ions

  • Ca²⁺ and O^2- are also smaller ions than K⁺ and Cl^-, so this also adds to the value for the lattice energy being more exothermic

Factors affecting enthalpy of hydration

• The standard enthalpy change of hydration (ΔH_hyd^ꝋ) is affected by the amount that the ions are attracted to the water molecules

• The factors which affect this attraction are the ionic charge and radius

Ionic radius

• ΔH_hyd^ꝋ becomes more exothermic with decreasing ionic radii

  • Smaller ions have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution

  • Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋ becomes more exothermic

• For example, the ΔH_hyd^ꝋ of magnesium sulfate (MgSO₄) is more exothermic than the ΔH_hyd^ꝋ of barium sulfate (BaSO₄)

  • Since both compounds contain a sulfate (SO₄^2-) ion, the difference in ΔH_hyd^ꝋ must be due to the magnesium (Mg²⁺) ion in MgSO₄ and barium (Ba²⁺) ion in BaSO₄

  • Magnesium is a Group 2 and Period 3 element

  • Barium is a Group 2 and Period 6 element

  • This means that the Mg²⁺ ion is smaller than the Ba²⁺ ion

  • The attraction is therefore much stronger for the Mg²⁺ ion

  • As a result, the standard enthalpy of hydration of MgSO₄ is more exothermic than that of BaSO₄

Ionic charge

• ΔH_hyd^ꝋ is more exothermic for ions with larger ionic charges

  • Ions with large ionic charges have a greater charge density resulting in stronger ion-dipole attractions between the water molecules and the ions in the solution

  • Therefore, more energy is released when they become hydrated and ΔH_hyd^ꝋbecomes more exothermic

• For example, the ΔH_hyd^ꝋ of calcium oxide (CaO) is more exothermic than the ΔH_hyd^ꝋof potassium chloride (KCl)

  • Calcium oxide is an ionic compound that consists of calcium (Ca²⁺) and oxide (O^2-) ions

  • Potassium chloride is formed from potassium (K⁺) and chloride (Cl^-) ions

  • Both of the ions in calcium oxide have a greater ionic charge than the ions in potassium chloride

  • This means that the attractions are stronger between the water molecules and Ca²⁺ and O^2-ions upon hydration of CaO

  • The attractions are weaker between the water molecules and K⁺ and Cl^- ions upon hydration of KCl

  • Therefore, the ΔH_hyd^ꝋ of calcium oxide is more exothermic as more energy is released upon its hydration

The enthalpy of hydration is more exothermic for smaller ions and for ions with a greater ionic charge