Group 2: thermal stability and solubility trends

· Focus elements: magnesium → calcium → strontium → barium.
· Group 2 ions all form M²⁺ ions; down the group, ionic radius increases and charge density decreases.
· Main CIE explanations required: polarisation of large anions, lattice energy, enthalpy change of hydration, and enthalpy change of solution, ΔH⦵sol.
· Key exam skill: explain trends qualitatively, not by memorising numerical data.

Group 2 metals are in the second column of the Periodic Table. In this topic, CIE focuses on the trends from magnesium to barium, especially how increasing ionic radius affects stability and solubility. Source

Thermal decomposition of Group 2 carbonates

· General equation: MCO₃(s) → MO(s) + CO₂(g).
· Carbonates become more thermally stable down Group 2: MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃.
· Therefore, more heat is needed to decompose carbonates as you go down the group.
· Explanation: smaller M²⁺ ions have higher charge density and greater polarising power.
· Mg²⁺ polarises/distorts the large CO₃²⁻ ion most strongly, weakening it and making decomposition easier.
· Down the group, M²⁺ ionic radius increases, so charge density decreases, polarisation decreases, and the carbonate becomes harder to decompose.
· Exam phrase: less polarisation of the large carbonate ion means greater thermal stability.

The diagram shows how a small, highly charged Group 2 cation can distort the electron cloud of the carbonate ion. Greater distortion makes decomposition easier. This explains why carbonates are less thermally stable near the top of Group 2. Source

Thermal decomposition of Group 2 nitrates

· General equation: 2M(NO₃)₂(s) → 2MO(s) + 4NO₂(g) + O₂(g).
· Nitrates become more thermally stable down Group 2: Mg(NO₃)₂ < Ca(NO₃)₂ < Sr(NO₃)₂ < Ba(NO₃)₂.
· Products include metal oxide, brown nitrogen dioxide gas, and oxygen gas.
· Explanation is the same as for carbonates: smaller M²⁺ ions have stronger polarising power.
· Mg²⁺ polarises the large NO₃⁻ ion most strongly, making nitrate decomposition easier.
· Down the group, ionic radius increases, charge density decreases, and polarisation of NO₃⁻ decreases.
· Exam phrase: thermal stability increases down the group because the cation polarises the anion less strongly.

Polarisation: the core explanation

· Polarisation = distortion of an anion’s electron cloud by a nearby cation.
· Polarising power increases when the cation has higher charge density.
· All Group 2 cations have the same charge, +2, so the key factor is ionic radius.
· Smaller ion = higher charge density = stronger polarising power.
· Larger ion = lower charge density = weaker polarising power.
· Large anions such as CO₃²⁻ and NO₃⁻ are more easily polarised than small anions.
· Stronger polarisation weakens the anion, so less heat is needed for decomposition.

The carbonate ion has delocalised electrons spread over the ion. This large electron cloud can be distorted by small, highly charged cations. That distortion is the key reason Group 2 carbonate stability changes down the group. Source

Solubility of Group 2 hydroxides

· Hydroxide solubility increases down Group 2: Mg(OH)₂ least soluble → Ba(OH)₂ most soluble.
· Trend: Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂ in solubility.
· More soluble hydroxide = more OH⁻(aq) released = more alkaline solution.
· Explanation uses ΔH⦵sol = lattice energy effect + hydration enthalpy effect.
· Dissolving requires energy to overcome the ionic lattice but releases energy when ions become hydrated.
· Down the group, lattice energy becomes less exothermic / lattice dissociation becomes easier because M²⁺ ions get larger.
· For hydroxides, the decrease in lattice energy magnitude is more important than the decrease in hydration enthalpy magnitude.
· Result: ΔH⦵sol becomes more favourable, so hydroxides become more soluble.

Solubility of Group 2 sulfates

· Sulfate solubility decreases down Group 2: MgSO₄ most soluble → BaSO₄ least soluble.
· Trend: MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄ in solubility.
· BaSO₄ is very insoluble, which is a common exam example.
· Explanation again uses ΔH⦵sol, lattice energy, and enthalpy change of hydration.
· Down the group, both lattice energy magnitude decreases and hydration enthalpy magnitude decreases.
· For sulfates, the large SO₄²⁻ ion means the lattice energy does not decrease enough to compensate for the large fall in hydration enthalpy magnitude.
· Result: ΔH⦵sol becomes less favourable, so sulfates become less soluble.
· Exam phrase: hydration enthalpy becomes much less exothermic down the group, making sulfate solubility decrease.

The diagram summarises the decreasing solubility of Group 2 sulfates. Magnesium sulfate is relatively soluble, while barium sulfate is very insoluble. This trend is explained by the balance between lattice energy and hydration enthalpy. Source

Lattice energy vs hydration enthalpy

· Lattice energy: energy change linked to forming or separating an ionic lattice.
· Stronger ionic attraction = larger lattice energy magnitude.
· Down Group 2, M²⁺ radius increases, so attraction between ions decreases and lattice energy magnitude decreases.
· Hydration enthalpy: energy released when gaseous ions become surrounded by water molecules.
· Smaller ions attract water molecules more strongly, so they have more exothermic hydration enthalpies.
· Down Group 2, hydration enthalpy becomes less exothermic because M²⁺ ions get larger.
· Solubility depends on the balance: ΔH⦵sol = energy needed to separate lattice + energy released by hydration.
· CIE expects you to compare the relative magnitudes, not just state that both decrease.

Comparing hydroxides and sulfates

· Hydroxides: solubility increases down the group.
· For hydroxides, the fall in lattice energy magnitude dominates, so dissolving becomes more favourable.
· Sulfates: solubility decreases down the group.
· For sulfates, the fall in hydration enthalpy magnitude dominates, so dissolving becomes less favourable.
· The anion matters: OH⁻ is smaller, while SO₄²⁻ is large, changing how lattice energy varies down the group.
· Key contrast: Group 2 hydroxides and sulfates have opposite solubility trends.

High-yield exam wording

· Thermal stability increases down Group 2 because M²⁺ ionic radius increases and polarising power decreases.
· Smaller M²⁺ ions polarise CO₃²⁻ / NO₃⁻ more strongly, so the anion is more easily decomposed.
· Hydroxide solubility increases down Group 2 because ΔH⦵sol becomes more favourable.
· Sulfate solubility decreases down Group 2 because hydration enthalpy becomes much less exothermic relative to lattice energy changes.
· Always compare relative changes in lattice energy and hydration enthalpy.

Common mistakes to avoid

· Do not say thermal stability increases because the bonds “get stronger” without explaining polarisation.
· Do not confuse thermal stability with reactivity.
· Do not state “solubility decreases down Group 2” for all compounds: hydroxides increase, sulfates decrease.
· Do not ignore ΔH⦵sol when explaining solubility; CIE specifically requires it.
· Do not use only “larger atoms” as an explanation; use larger M²⁺ ionic radius, lower charge density, and weaker polarising power.