A Level Chemistry: Group 2 Trends — Revision Notes

Condensed for revision. For the full explanation, use the Group 2: Thermal Stability and Solubility Trends study guide, then test yourself with the practice questions. For the AS facts these explanations build on (reactions of the elements, and the trends stated without the energetics), see the AS Group 2 revision notes. For the definitions of lattice energy and enthalpy of hydration, see the Lattice Energy, Entropy and Gibbs Free Energy revision notes.

Syllabus: Cambridge International AS & A Level Chemistry 9701, 2025–2027, A Level content: subtopic 27.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds.

The one idea behind both trends (27.1)

Every Group 2 cation has the same charge, 2+. Down the group, from Mg²⁺ to Ba²⁺, the ionic radius increases, so the charge density of the cation decreases.

Cation	Relative ionic radius	Charge density	Polarising power
Mg²⁺	smallest	highest	greatest
Ca²⁺	↓	↓	↓
Sr²⁺	↓	↓	↓
Ba²⁺	largest	lowest	least

The similarities come from the common 2+ charge: every nitrate decomposes in the same way, every carbonate decomposes in the same way, and every hydroxide and sulfate has the formula M(OH)₂ or MSO₄. The trends come from the increasing ionic radius.

Thermal stability of nitrates and carbonates (27.1, outcome 1)

Decomposition equations

All the Group 2 nitrates and carbonates (magnesium to barium) decompose on heating to the oxide (barium carbonate only at very high temperatures, beyond a Bunsen burner).

2Mg(NO₃)₂(s) → 2MgO(s) + 4NO₂(g) + O₂(g)
2Ba(NO₃)₂(s) → 2BaO(s) + 4NO₂(g) + O₂(g)

MgCO₃(s) → MgO(s) + CO₂(g)
BaCO₃(s) → BaO(s) + CO₂(g)

Nitrates: a brown gas (NO₂) is seen and the oxygen relights a glowing splint. Carbonates: the CO₂ turns limewater milky.

The trend

Thermal stability increases down the group for both nitrates and carbonates. MgCO₃ and Mg(NO₃)₂ decompose at the lowest temperature; BaCO₃ and Ba(NO₃)₂ need the highest temperature.

The explanation: polarisation of the large anion

1. The nitrate ion (NO₃⁻) and the carbonate ion (CO₃²⁻) are large anions with a large, easily distorted electron cloud.
2. A small cation with a high charge density attracts the electron cloud of the neighbouring anion towards itself: it polarises (distorts) the anion.
3. Polarisation weakens the N–O bonds in NO₃⁻ (or the C–O bonds in CO₃²⁻), so less energy is needed to break the anion apart into NO₂ and O₂ (or CO₂), leaving the oxide.
4. Down the group the ionic radius increases, so the charge density and polarising power fall. The anion is distorted less, its bonds are weakened less, and a higher temperature is needed to decompose the compound.

Chain to learn: larger cation → lower charge density → less polarisation of the anion → anion bonds weakened less → more thermally stable.

Solubility and enthalpy of solution of hydroxides and sulfates (27.1, outcome 2)

The trends

Compound	Trend down the group (Mg → Ba)	Ends of the series
Hydroxides, M(OH)₂	solubility increases	Mg(OH)₂ only sparingly soluble; Ba(OH)₂ much more soluble
Sulfates, MSO₄	solubility decreases	MgSO₄ soluble; BaSO₄ insoluble

The energy cycle

Using the 9701 convention that lattice energy is exothermic (the enthalpy change when one mole of the solid lattice forms from its gaseous ions):

ΔHsol = −ΔHlatt + ΔHhyd(cation) + ΔHhyd(anions)

For a hydroxide: ΔHsol = −ΔHlatt + ΔHhyd(M²⁺) + 2ΔHhyd(OH⁻). For a sulfate: ΔHsol = −ΔHlatt + ΔHhyd(M²⁺) + ΔHhyd(SO₄²⁻).

• Lattice energy term: energy put in to separate the ions. Its size depends on the charges and on the distance between ion centres, r(cation) + r(anion).
• Hydration terms: energy released as water molecules surround each gaseous ion. Because every Group 2 cation is 2+, the size of ΔHhyd(cation) depends only on the cation’s own radius.
• In simple terms, the more exothermic (or less endothermic) ΔHsol is, the more soluble the compound tends to be. The syllabus asks for the explanation in terms of the relative magnitudes of the enthalpy change of hydration and the lattice energy (entropy is not required).

Down the group both the lattice energy and ΔHhyd(cation) become less exothermic, because the cation gets larger. The trend in ΔHsol depends on which one falls faster, and that depends on the size of the anion.

Hydroxides: a small anion

• OH⁻ is small, so the cation radius is a large part of r(cation) + r(anion).
• As the cation gets larger, the inter-ionic distance increases by a large fraction, so the lattice energy falls a lot.
• The lattice energy falls faster than the hydration enthalpy of the cation.
• The energy needed to break the lattice decreases by more than the energy released on hydration decreases, so ΔHsol becomes more exothermic (less endothermic) down the group: solubility increases.

Sulfates: a large anion

• SO₄²⁻ is large, so r(anion) dominates r(cation) + r(anion).
• As the cation gets larger, the inter-ionic distance increases by only a small fraction, so the lattice energy changes only a little.
• The hydration enthalpy of the cation still falls significantly, because it depends only on the cation’s own radius.
• The hydration enthalpy falls faster than the lattice energy, so ΔHsol becomes less exothermic (more endothermic) down the group: solubility decreases.

	Hydroxides (small OH⁻)	Sulfates (large SO₄²⁻)
Lattice energy down the group	becomes much less exothermic	becomes only slightly less exothermic
ΔHhyd(cation) down the group	becomes less exothermic	becomes less exothermic
Which changes more?	lattice energy	hydration enthalpy
ΔHsol down the group	more exothermic	less exothermic
Solubility down the group	increases	decreases

Exam traps

• Hydroxides and sulfates go in opposite directions: hydroxides more soluble down the group, sulfates less soluble. Do not write one “solubility increases” rule.
• Thermal stability increases down the group. Mg compounds decompose most easily, not Ba compounds.
• Use polarisation for thermal stability and the energy cycle for solubility. Do not mix the two explanations.
• It is the small cation that polarises the large anion, not the other way round.
• Say ionic radius and charge density, not atomic radius, when explaining these trends.
• Both lattice energy and hydration enthalpy become less exothermic down the group for both types of compound. The mark is for saying which changes more, and why (the size of the anion).
• Nitrate equations need 2 in front of M(NO₃)₂: 2M(NO₃)₂ → 2MO + 4NO₂ + O₂. Check the oxygen atoms: 12 on each side.

Self-test

1. Write a balanced equation, with state symbols, for the thermal decomposition of strontium nitrate.
2. Write a balanced equation for the thermal decomposition of calcium carbonate.
3. State the trend in thermal stability of the Group 2 carbonates from magnesium to barium.
4. Explain why magnesium carbonate decomposes at a lower temperature than barium carbonate.
5. Which bonds are weakened when a Group 2 cation polarises a nitrate ion?
6. State the trends in solubility of the Group 2 hydroxides and of the Group 2 sulfates down the group.
7. Write the expression linking ΔHsol to the lattice energy and the enthalpies of hydration for MSO₄.
8. Explain why the lattice energy of the Group 2 sulfates changes only a little down the group.
9. Explain, in terms of lattice energy and enthalpy of hydration, why the solubility of the Group 2 hydroxides increases down the group.
10. Which is more soluble in water, magnesium sulfate or barium sulfate? Which is more soluble, magnesium hydroxide or barium hydroxide?

Answers:

1. 2Sr(NO₃)₂(s) → 2SrO(s) + 4NO₂(g) + O₂(g)
2. CaCO₃ → CaO + CO₂
3. Thermal stability increases down the group.
4. Mg²⁺ is smaller than Ba²⁺, so it has a higher charge density. It polarises (distorts) the electron cloud of the large carbonate ion more, weakening the C–O bonds more, so less energy is needed to decompose MgCO₃ to MgO and CO₂.
5. The N–O bonds within the nitrate ion.
6. Hydroxides: solubility increases down the group. Sulfates: solubility decreases down the group.
7. ΔHsol = −ΔHlatt + ΔHhyd(M²⁺) + ΔHhyd(SO₄²⁻)
8. SO₄²⁻ is large, so it dominates the inter-ionic distance r(M²⁺) + r(SO₄²⁻). A larger cation increases this distance by only a small fraction, so the lattice energy changes only slightly.
9. OH⁻ is small, so an increase in cation radius greatly increases the inter-ionic distance and the lattice energy becomes much less exothermic. The hydration enthalpy of the cation also becomes less exothermic, but by less. So ΔHsol becomes more exothermic (less endothermic) down the group and solubility increases.
10. Magnesium sulfate; barium hydroxide.

These are original notes written for revision. The trends are given qualitatively, as the syllabus requires, and no data values are quoted. Check the full syllabus wording in the official 9701 syllabus.