Group 2: Quantitative Trends in Thermal Stability and Solubility

This guide covers subtopic 27.1, Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds, from Topic 27, Group 2, of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is A Level content — the energetics explanation for trends Group 2: The Alkaline Earth Metals states qualitatively at AS Level without the underlying reasoning.

Before studying this

This resource assumes the AS-level trends themselves (thermal decomposition of nitrates/carbonates, hydroxide and sulfate solubility direction) from Group 2: The Alkaline Earth Metals, and lattice energy and hydration enthalpy from Chemical Energetics: Lattice Energy, Entropy and Gibbs Free Energy. Read that energetics resource first if the terms lattice energy and hydration enthalpy are unfamiliar.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — A Level, Topic 27

27.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds — describing and explaining qualitatively the trend in thermal stability of the nitrates and carbonates, including the effect of ionic radius on polarisation of the large anion; describing and explaining qualitatively the variation in solubility and enthalpy change of solution of the hydroxides and sulfates, in terms of the relative magnitudes of enthalpy change of hydration and lattice energy.

Thermal stability of nitrates and carbonates

AS Level established that thermal stability of Group 2 nitrates and carbonates increases down the group (bigger cations need a higher temperature to decompose). The A Level explanation is about polarisation: a Group 2 cation, being small and doubly charged, distorts (polarises) the electron cloud of the large nitrate or carbonate anion sitting next to it, weakening a bond within that anion enough to trigger decomposition on heating.

Smaller cations polarise the anion more strongly (a smaller cation has a higher charge density — the same +2 charge concentrated into a smaller volume, producing a stronger local electric field), so:

• Mg²⁺ (smallest Group 2 cation) polarises the anion most strongly, weakening it most, giving the least thermally stable nitrate/carbonate — decomposing at the lowest temperature.
• Ba²⁺ (largest) polarises least, giving the most thermally stable compounds — decomposing only at the highest temperature.

This is a purely qualitative, comparative argument: the syllabus doesn’t require actual lattice-energy numbers here, only the polarising-power reasoning connecting ionic radius to bond weakening.

Solubility trends: hydroxides vs sulfates

AS Level established that hydroxide solubility increases down the group while sulfate solubility decreases down the group — genuinely opposite trends for the two anions. The energetics explanation uses the solution cycle from Chemical Energetics: Lattice Energy, Entropy and Gibbs Free Energy:

ΔH_sol = −ΔH_latt + ΔH_hyd(cation) + ΔH_hyd(anion)

Solubility depends on how ΔH_sol changes down the group — more exothermic (or less endothermic) favours greater solubility. Both ΔH_latt and ΔH_hyd(cation) become less exothermic going down the group, since the cation’s radius increases and both terms weaken with distance — but the two anions respond differently because of their very different sizes.

Hydroxide, OH⁻ (a small anion). With a small anion, the lattice energy is quite sensitive to the cation’s size (both ions are comparably small, so the cation-radius increase noticeably increases the interionic distance and weakens the lattice). Down the group, ΔH_latt falls faster than ΔH_hyd(cation) does, so the net effect on ΔH_sol becomes more exothermic — hydroxide solubility increases down the group, matching the observed trend (Mg(OH)₂ is barely soluble; Ba(OH)₂ is soluble enough to use as a laboratory alkali).

Sulfate, SO₄²⁻ (a large anion). With a large anion, the lattice energy is dominated by the anion’s size and is much less sensitive to changes in the (comparatively small) cation’s radius — so ΔH_latt changes only a little down the group. Meanwhile ΔH_hyd(cation) still falls substantially, since hydration enthalpy depends strongly on the cation’s own size regardless of what anion it’s paired with. With ΔH_hyd(cation) falling faster than ΔH_latt, the net effect on ΔH_sol becomes less exothermic (more endothermic) — sulfate solubility decreases down the group, matching the observed trend (MgSO₄ is soluble enough to be Epsom salts; BaSO₄ is so insoluble it’s used as a barium meal for X-ray imaging, safe because none dissolves into the body).

The key idea to hold onto: whether lattice energy or hydration enthalpy “wins” as the cation gets bigger depends on the anion’s own size relative to the cation — a small anion (OH⁻) makes the lattice term move comparably to the hydration term, while a large anion (SO₄²⁻) makes the lattice term nearly constant, leaving the hydration term to dominate the trend.

Common mistakes

Trying to explain thermal stability using lattice energy or hydration enthalpy. That quantitative energetics treatment applies to the solubility trend, not the thermal stability trend — thermal stability is explained by cation polarising power acting on the anion, a different argument entirely. Keep the two explanations separate.

Assuming both hydroxide and sulfate solubility move the same direction down the group. They don’t — this is precisely the point of comparing a small anion (OH⁻) against a large one (SO₄²⁻); memorise the directions as a pair, not as a single “solubility increases down the group” rule.

Forgetting that it’s the relative rate of change of ΔH_latt and ΔH_hyd that matters, not their absolute exothermicity. Both quantities get less exothermic down the group for both anions — what differs between hydroxides and sulfates is only which one falls faster.

Quick revision checklist

• Thermal stability of nitrates/carbonates increases down the group: smaller cation → more polarisation → less stable
• Solubility explanation uses ΔH_sol = −ΔH_latt + ΣΔH_hyd
• Small anion (OH⁻): lattice energy falls faster down the group → ΔH_sol more exothermic → solubility increases
• Large anion (SO₄²⁻): lattice energy barely changes, hydration enthalpy still falls → ΔH_sol less exothermic → solubility decreases
• Two separate arguments: polarisation for thermal stability, energetics cycle for solubility — don’t mix them up

Related resources

• Group 2: The Alkaline Earth Metals — the AS-level trends this resource explains quantitatively
• Chemical Energetics: Lattice Energy, Entropy and Gibbs Free Energy — lattice energy and hydration enthalpy, the tools this topic applies
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.