Inorganic Chemistry: Group 2

Inorganic Chemistry: Group 2

1. TRENDS

   • Ionisation Energy
     • Nuclear charge increases → force of attraction for the e- being removed increases → increase in IE down group
     • Each added quantum shell → energy of outermost e- increases
     • Number of filled inner shells increases → force of repulsion on e- being removed increases → decrease in IE down group
     • Combined effect of the last 2 factors outweighs the effect of 1st factor → decrease in IE down group
   • Reactivity
     • Reactivity of G2 metals increases down group
     • Atomic radii increases → more shielding → nuclear attraction decreases → easier to remove outer e-
   • Melting points
     • Melting points decrease down the group
     • Metallic bonding weakens as atomic size increases → distance between cations + delocalised e- increases → electrostatic attractive forces between cations + delocalised e- weaken
   • Atomic Radius
     • Atomic Radius increases down the group
     • More e- → more shells of e- → bigger atom

2. REACTIONS

   • With Oxygen
     • Group 2 metals burn in oxygen
     • Magnesium burns with a bright white flame + formation of white solid
     • GE: 2M(s) + O2(g) → 2MO(s)
   • With Chlorine
     • Group 2 metals combine with chlorine when heated in the gas
     • GE: M(s) + Cl2(g) → MCl2(s)
   • With Water
     • Magnesium reacts slowly with water + doesn't completely react
     • Calcium, Strontium & Barium react w increasing vigour → increase in effervescence
     • GE: M(s) + 2H2O(l) → M(OH)2(aq) + H2(g)
     • Reaction with Calcium (Calcium Hydroxide is only slightly soluble in water, so liquid turns cloudy: Ca(s) + 2H2O(l) → Ca(OH)2(s) + H2(g)
   • Magnesium & Steam
     • When heated in steam, Magnesium Oxide is formed (white solid) and Hydrogen gas in a vigorous reaction
       • Mg(s) + H2O(g) → MgO(s) + H2(g)

3. REACTIONS OF GROUP 2 OXIDES & HYDROXIDES, & TRENDS IN SOLUBILITY

4. THERMAL STABILITY OF GROUP 2 COMPOUNDS & THE COMPARISON WITH GROUP 1

   • Carbonates
     • G2 carbonates decompose on heating → produce G2 oxides + CO2 gas
       • More thermally stable as you go down the group → bigger cations → less of a polarising effect → distort carbonate ion less → C–O bond is less weakened → doesn't break down easily
       • EXP: Heat known mass of carbonate in boiling tube + pass gas produced through lime water
       • MgCO3(s) → MgO(s) + CO2(g)
     • G1 carbonates do not decompose apart from Lithium
       • Only have +1 charges → don't have a big enough charge density to polarise the carbonate ion
       • Li → small enough to have polarising effect
       • Li2CO3(s) → Li2O(s) + CO2(g)
   • Nitrates
     • G2 nitrates decompose on heating to produce G2 oxides, O2 gas and NO2 gas
       • Observe brown gas evolving (NO2) + white nitrate solid melt to colourless solution, then resolidify
       • More thermally stable as you down the group → bigger cations → less of a polarising effect → distort carbonate ion less → N–O bond is less weakened → doesn't break down easily
       • Mg(NO3) decomposes the easiest → Mg2+ ion is smallest + has greater charge density → causes more polarisation of nitrate anion → weakens N–O bond
       • 2Mg(NO3)2 → 2MgO + 4NO2 +O2
     • G1 nitrates do not decompose in the same way as G2 (apart from Lithium nitrate) → decompose to give nitrate (III) salt + O2 gas
       • 2NaNO3 → 2NaNO2 + O2

5. FLAME TESTS & THE TEST FOR AMMONIUM IONS

   • RED: Lithium, Calcium, Rubidium
   • YELLOW: Sodium
   • LILAC: Potassium
   • BLUE: Caesium
   • GREEN: Barium
   • Heat causes an electron to move to higher energy level → e- is unstable at higher e- level → drop back down → energy is emitted in form of visible light energy → with wavelength of observable light