Similarities, Trends & Compounds of Magnesium to Barium (CIE A Level Chemistry)

• The Group 2 nitrates and carbonates become more thermally stable going down the group
• The charge density of the cation (Group 2 metal ion) and the polarisation  of the anion (the nitrate and carbonate ion) attribute towards this increased stability

Trends in thermal stability going down the group

• All Group 2 metals form 2+ ions as they lose two electrons from their valence shells
• The metal cations at the top of the group are smaller in size than those at the bottom
  • For example, the atomic radius of beryllium (the first element in Group 2) is 112 pm whereas the atomic radius of calcium (further down the group) is 197 pm

• The metal cations at the top of Group 2, therefore, have the greatest charge density as the same charge (2+) is packed into a smaller volume
• As a result, smaller Group 2 ions have a greater polarising effect on neighbouring negative ions
• When a carbonate or nitrate ion approaches the cation, it becomes polarised
  • This is because the metal cation draws the electrons in the carbonate or nitrate ion towards itself

• The more polarised the anion is, the less heat is required to thermally decompose them
• Therefore, the thermal stability increases down the group
  • As down the group, the cation becomes larger
  • Thus has a smaller charge density
  • And a smaller polarising effect on the carbonate or nitrate anion
  • So the anion is less polarised
  • Therefore, more heat is required to thermally decompose them

• The solubility of Group 2 hydroxides increases down the group
• In contrast, the Group 2 sulfates show a decrease in solubility going down the group
• Compounds that have very low solubility are said to be sparingly soluble
  • For example, calcium sulfate (CaSO₄) has low solubility as only 0.21 g of CaSO₄ dissolves in 100 g of water

• Most of the sulfates are soluble in warm water with the exception of barium sulfate which is insoluble

Solubility of Group 2 elements table

Enthalpy change of hydration and lattice energy

• The standard enthalpy of solution (ΔH_sol^ꝋ) is the energy absorbed or released when 1 mole of ionic solid dissolves in enough water to form a dilute solution (under standard conditions)
  • The ΔH_sol^ꝋ can be either exothermic or endothermic

• For example, the ΔH_sol^ꝋ of sodium chloride (NaCl) is +3.9 kJ mol^-1

NaCl (s) + aq → NaCl (aq)

OR

NaCl (s) + aq → Na⁺ (aq) + Cl^- (aq)

• This means, that 3.9 kJ mol^-1 of energy is absorbed when one mole of NaCl is dissolved in enough water to form a dilute solution

ΔH_sol^ꝋ = ΔH_hyd^ꝋ - ΔH_latt^ꝋ 

• The lattice (formation) energy is the energy released when gaseous ions combine to form one mole of an ionic compound under (standard conditions)
  • Since energy is released when an ionic compound is formed, the ΔH_latt^ꝋ is always exothermic
  • For example, the ΔH_latt^ꝋ of NaCl is -787 kJ mol^-1

Na⁺ (g) + Cl^- (g) → NaCl (s)   

• This means, that 787 kJ mol^-1 of energy is released when NaCl is formed from its gaseous ions
• The standard enthalpy of hydration is the energy released when gaseous ions dissolve in enough water to form a dilute solution (under standard conditions)
  • Since energy is released when gaseous ions become hydrated, the ΔH_hyd^ꝋ is always exothermic
  • For example, the ΔH_hyd^ꝋ of the sodium (Na⁺) ion is -406 kJ mol^-1

Na⁺ (g) → Na⁺ (aq)

• This means, that 406 kJ mol^-1 of energy is released when Na⁺ ions become hydrated

Trends of enthalpy change of solution

• Going down the group, the ΔH_latt^ꝋ of the ionic compounds decreases
  • Going down the group, the positively charged cations become larger
  • There is more space between the negatively and positively charged ions in the ionic compound so there are weaker attractive forces between the ions
  • As there are weaker electrostatic forces between the ions, there is less energy released upon formation of the ionic compound from its gaseous ions
  • Therefore, the ΔH_latt^ꝋ becomes less exothermic

• Going down the group, the ΔH_hyd^ꝋ also decreases
  • Again, the positively charged ions become larger going down the group
  • As a result, the ion-dipole bonds between the cations and water molecules get weaker
  • This means that less energy is released when the gaseous Group 2 ions become hydrated
  • The ΔH_hyd^ꝋ , therefore, becomes less exothermic

• For Group 2 hydroxides:
  • Hydroxide ions are relatively small ions
  • The ΔH_latt^ꝋ falls faster than the ΔH_hyd^ꝋ
  • The enthalpy change of solution is, therefore, more exothermic going down the group

• For Group 2 sulfates:
  • Sulfate ions are relatively large ions
  • The ΔH_latt^ꝋ falls slower than the ΔH_hyd^ꝋ enthalpy
  • The ΔH_sol^ꝋ will become more endothermic going down the group

• The more exothermic the ΔH_sol^ꝋ the more soluble the compound
  • This is why the sulfates become less soluble going down the group and the hydroxides more soluble