CIE A Level Chemistry

Revision Notes

Syllabus Edition

First teaching 2020

Last exams 2024

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6.1.1 Similarities, Trends & Compounds of Magnesium to Barium

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Effect of Ionic Radius on Thermal Stability of Group 2 Nitrates & Carbonates

  • 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

Trends in Solubility & Enthalpy Change of Solution of Group 2 Hydroxides & Sulfates

  • 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

Group 2 - Solubility of Group 2 elements table, downloadable AS & A Level Chemistry revision notes

Enthalpy change of hydration and lattice energy

  • The standard enthalpy of solution (ΔHsol) 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 ΔHsol can be either exothermic or endothermic

  • For example, the ΔHsol 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

ΔHsol = ΔHhyd - ΔHlatt 

  • 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 ΔHlatt is always exothermic
    • For example, the ΔHlatt 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 ΔHhyd is always exothermic
    • For example, the ΔHhyd 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 ΔHlattof 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 ΔHlatt becomes less exothermic

  • Going down the group, the ΔHhyd 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 ΔHhyd , therefore, becomes less exothermic

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

  • For Group 2 sulfates:
    • Sulfate ions are relatively large ions
    • The ΔHlatt falls slower than the ΔHhyd enthalpy
    • The ΔHsol will become more endothermic going down the group

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

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