Unit 17: Group 2 Elements — Long Questions

Reactivity with water Beryllium does not react with water or steam even on heating, because of its strong protective oxide layer. Magnesium reacts very slowly with cold water to form Mg(OH)₂ and H₂: Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g); with steam it reacts faster, giving MgO and H₂: Mg(s) + H₂O(g) → MgO(s) + H₂(g). Calcium reacts with cold water and steam more vigorously than magnesium, giving a sparingly soluble white precipitate of Ca(OH)₂ and H₂: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). Reactivity with water increases down the group (Sr and Ba react even more vigorously than Ca, e.g. Ba(s) + 2H₂O(l) → Ba(OH)₂(aq) + H₂(g)) because of increasing atomic size, increasing shielding, and decreasing ionization energy.

Reactivity with dilute HCl Group 2 metals react with dilute acids to give a salt and hydrogen gas, generally M(s) + 2HCl(aq) → MCl₂(aq) + H₂(g). Beryllium reacts with dilute acids to form BeCl₂ and H₂, but its protective oxide layer slows this reaction. Magnesium reacts readily: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). Calcium reacts vigorously: Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g); however with H₂SO₄ the reaction of Ca can stop early because the sparingly soluble CaSO₄ formed coats the metal and protects it from further attack: Ca(s) + H₂SO₄(aq) → CaSO₄(s) + H₂(g). Overall, reactivity of Group 2 metals with both water and acids increases down the group (Be < Mg < Ca < Sr < Ba) due to increasing atomic size and decreasing ionization energy, which makes electron loss progressively easier.

Group 2 carbonates decompose on heating to give a metal oxide and carbon dioxide gas: MCO₃(s) → MO(s) + CO₂(g), e.g. MgCO₃(s) → MgO(s) + CO₂(g) and CaCO₃(s) → CaO(s) + CO₂(g). Group 2 nitrates decompose on heating to give a metal oxide, brown nitrogen dioxide gas, and oxygen gas: 2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g), e.g. 2Ca(NO₃)₂(s) → 2CaO(s) + 4NO₂(g) + O₂(g).

Reason the thermal stability of both carbonates and nitrates depends on the polarizing power of the M²⁺ cation. Smaller cations with higher charge density (e.g. Be²⁺) distort the large anion's (CO₃²⁻ or NO₃⁻) electron cloud more strongly, weakening its internal bonds and allowing decomposition at a lower temperature. As the cation gets larger down the group, its polarizing power decreases, the anion is distorted less, and a higher temperature is needed to decompose the compound.

Relative order of ease of decomposition (least stable to most stable) — Carbonates: BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃ (decomposition temperatures 250 °C, 540 °C, 900 °C, 1289 °C, 1360 °C respectively). Nitrates: Be(NO₃)₂ is the least stable and Ba(NO₃)₂ is the most stable, following the same increasing trend down the group.

Solubility of an ionic compound depends on the balance between the exothermic enthalpy of hydration (ΔH_hyd, energy released when gaseous ions are hydrated) and the endothermic lattice enthalpy (ΔH_lattice, energy needed to break the ionic lattice into gaseous ions). A more negative (or less positive) overall enthalpy of solution (ΔH_soln) favours greater solubility.

Sulphates The solubility of Group 2 sulphates decreases down the group, e.g. MgSO₄ is soluble (1.83 × 10⁻¹ mol/100 g water) while BaSO₄ is essentially insoluble (9.43 × 10⁻⁷ mol/100 g water). The sulphate ion is relatively large, so its lattice enthalpy does not fall much as the cation grows down the group, but the hydration enthalpy of the cation falls significantly as the cation gets larger. Since the decrease in hydration enthalpy is more significant than the decrease in lattice enthalpy for this large anion, ΔH_soln becomes less favourable down the group, so solubility decreases.

Hydroxides The solubility of Group 2 hydroxides increases down the group, e.g. Mg(OH)₂ is sparingly soluble (2 × 10⁻⁵ mol/100 g water) while Ba(OH)₂ is quite soluble (1.5 × 10⁻² mol/100 g water). The hydroxide ion is small, comparable in size to the Group 2 cations, so lattice enthalpy depends strongly on cation size and falls off faster down the group than hydration enthalpy. Since the decrease in lattice enthalpy outweighs the decrease in hydration enthalpy, ΔH_soln becomes more favourable down the group, so solubility increases.

The Dow Process extracts magnesium metal from the magnesium ions (Mg²⁺) present in seawater in three main steps:

1. Precipitation: Seawater is treated with calcium hydroxide, Ca(OH)₂ (obtained from limestone or oyster shells), which precipitates the dissolved Mg²⁺ as insoluble magnesium hydroxide: Mg²⁺(aq) + Ca(OH)₂(s) → Mg(OH)₂(s) + Ca²⁺(aq).

2. Conversion to chloride: The precipitated Mg(OH)₂ is filtered off and reacted with hydrochloric acid to form magnesium chloride solution: Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l).

3. Electrolysis: The MgCl₂ solution is concentrated and dried, and the molten MgCl₂ is electrolysed in large electrolytic cells: MgCl₂(l) → Mg(l) + Cl₂(g). This final step produces pure magnesium metal at the cathode and chlorine gas at the anode.

This process is industrially significant because it extracts a valuable metal from an abundant but very low-concentration source (seawater) on a massive scale, and has revolutionized magnesium production.