• 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^?) 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^??can be either?exothermic?or?endothermic
• For example, the?ΔH_sol^??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
• The?ΔH_sol^??is the sum of the?lattice energy?(ΔH_latt^?) and the standard enthalpy change of hydration?(ΔH_hyd^?)

ΔH_sol^??=?ΔH_latt^??+?ΔH_hyd^?

• 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^??is always?exothermic
  • For example, the?ΔH_latt^??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^??is always?exothermic
  • For example, the?ΔH_hyd^??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^??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^??becomes?less exothermic
• Going down the group, the?ΔH_hyd^??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^??, therefore, becomes?less exothermic
• For?Group 2 hydroxides:
  • Hydroxide ions are relatively small ions
  • The?ΔH_latt^??falls faster than the?ΔH_hyd^?
  • 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^??falls slower than the?ΔH_hyd^??enthalpy
  • The?ΔH_sol^??will become more?endothermic?going down the group
• The more?exothermic?the?ΔH_sol^??the?more soluble?the compound
  • This is why the?sulfates?become?less?soluble going down the group and the?hydroxides?more?soluble