A Level Chemistry: Group 2 — Thermal Stability and Solubility Trends Practice Questions

These are original questions written for Marlbridge, in the style and at the standard of the examination. They are not reproduced past-paper questions — examination boards hold copyright in their own papers. Use these alongside the official past papers available free from your board.

Related: Group 2: Thermal Stability and Solubility Trends study guide

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Section A

1. State how ionic radius affects the polarising power of a Group 2 cation. [2]

2. State the equation linking enthalpy change of solution, lattice energy and enthalpy changes of hydration. [1]

3. State whether hydroxide solubility and sulfate solubility follow the same trend down Group 2. [1]

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Section B

4. Explain why magnesium carbonate decomposes at a lower temperature than barium carbonate on heating. [4]

5. Explain, in terms of ionic radius and polarising power, why calcium nitrate is thermally more stable than magnesium nitrate but less stable than barium nitrate. [3]

6. Magnesium hydroxide is only sparingly soluble, but barium hydroxide is soluble enough to be used as a laboratory alkali.

(a) State how lattice energy and hydration enthalpy of the cation change going from Mg²⁺ to Ba²⁺. [2]

(b) Explain why this makes ΔH(solution) become more exothermic down the group for the hydroxides, and link this to the observed solubility trend. [3]

7. Explain why magnesium sulfate is soluble but barium sulfate is so insoluble that it is used safely inside the body for medical imaging. [5]

8. A student claims: “Lattice energy always decreases in size (becomes less exothermic) as the cation gets bigger, for any Group 2 salt.” Using the sulfate and hydroxide comparison, explain why this statement, while true, is not by itself enough to predict the solubility trend of a Group 2 salt. [3]

9. A student explains that barium carbonate is more thermally stable than magnesium carbonate because “barium’s lattice energy is less exothermic than magnesium’s.” Explain why this argument is not appropriate, and give the correct explanation. [3]

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Answers

1. A smaller cation has the same +2 charge concentrated into a smaller volume, giving it a higher charge density [1]. This produces a stronger local electric field, so it distorts (polarises) a neighbouring anion’s electron cloud more strongly [1].

2. ΔH(solution) = −ΔH(lattice) + ΔH(hydration, cation) + ΔH(hydration, anion) [1].

3. No — hydroxide solubility increases down the group while sulfate solubility decreases down the group [1].

4. Mg²⁺ is a smaller cation than Ba²⁺ (higher charge density) [1], so it polarises the large carbonate anion more strongly [1], weakening a bond within the carbonate ion more [1], meaning less thermal energy is needed to trigger decomposition — so MgCO₃ decomposes at a lower temperature than BaCO₃ [1].

5. Ca²⁺ is smaller than Ba²⁺ but larger than Mg²⁺ [1]. A smaller cation polarises the nitrate anion more strongly, weakening it and lowering its decomposition temperature [1], so thermal stability increases in the order Mg < Ca < Ba, exactly as observed [1].

6. (a) Both lattice energy and hydration enthalpy of the cation become less exothermic (fall in magnitude) going from Mg²⁺ to Ba²⁺, since both terms weaken as the interionic/ion-water distance increases with cation size [2].

(b) OH⁻ is a small anion, so lattice energy is quite sensitive to the cation’s size and falls faster than the cation’s hydration enthalpy does going down the group [1]. This means the net effect on ΔH(solution) becomes more exothermic down the group [1], which favours greater solubility — matching the observed increase in hydroxide solubility from Mg(OH)₂ to Ba(OH)₂ [1].

7. SO₄²⁻ is a large anion, so lattice energy is dominated by the anion’s own size and changes only a little as the cation gets bigger down the group [1]. Meanwhile, hydration enthalpy of the cation still falls substantially going from Mg²⁺ to Ba²⁺, since it depends strongly on the cation’s own size regardless of the anion present [1]. With hydration enthalpy falling faster than lattice energy stays roughly constant, the net effect on ΔH(solution) becomes less exothermic (more endothermic) down the group [1], so solubility decreases from MgSO₄ (soluble, Epsom salts) to BaSO₄ (essentially insoluble) [1]. BaSO₄’s extreme insolubility means it passes through the body without dissolving into the bloodstream, which is why it is safe to use as a barium meal for X-ray imaging despite barium being toxic in soluble form [1].

8. The statement about lattice energy alone is true, but solubility depends on the relative rates of change of lattice energy and cation hydration enthalpy, not on lattice energy in isolation [1]. For a small anion like OH⁻, lattice energy falls faster than hydration enthalpy, giving more exothermic ΔH(solution) and increasing solubility [1]; for a large anion like SO₄²⁻, lattice energy barely changes while hydration enthalpy still falls, giving less exothermic ΔH(solution) and decreasing solubility — so the same lattice-energy trend produces opposite solubility trends depending on the anion [1].

9. Lattice energy and hydration enthalpy are the tools used to explain the solubility trend of hydroxides and sulfates, not the thermal stability of nitrates and carbonates [1]; using them here mixes up two separate explanations. The correct explanation is about polarising power: Ba²⁺ is a larger cation than Mg²⁺, with a lower charge density, so it polarises the carbonate anion less strongly, weakening it less [1]; this means barium carbonate needs more thermal energy to trigger decomposition, making it more thermally stable than magnesium carbonate, whose smaller, more polarising cation weakens the anion more [1].

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Where marks are usually lost

• Using the lattice-energy/hydration-enthalpy argument to explain thermal stability, or the polarising-power argument to explain solubility — these are two separate explanations for two separate trends.
• Stating that hydroxide and sulfate solubility both increase (or both decrease) down the group, rather than recognising they move in opposite directions.
• Explaining a solubility trend by quoting only one of lattice energy or hydration enthalpy, rather than comparing which one changes faster.
• Forgetting that it is the anion’s size (small for OH⁻, large for SO₄²⁻) that determines which energy term dominates the trend.