Cambridge International AS and A Level Chemistry 27: Group 2

Group 2 is Cambridge International Chemistry 9701 Topic 27. The A Level boundary compares magnesium to barium through nitrate and carbonate thermal stability, then explains hydroxide and sulfate solubility and enthalpy of solution through the competing magnitudes of hydration enthalpy and lattice energy.

1. The comparison set

Cambridge asks for the Group 2 sequence magnesium, calcium, strontium and barium. Each forms a two-plus ion, so ionic charge remains constant while ionic radius increases down the group.

That controlled change in radius drives the required explanations. A larger ion has lower charge density and attracts nearby electron density less strongly.

Beryllium is not part of this stated comparison, so do not use its atypical chemistry to define the requested trend.

2. Thermal decomposition of carbonates

Group 2 carbonates decompose on heating to the metal oxide and carbon dioxide. The general pattern is metal carbonate giving metal oxide plus carbon dioxide.

Thermal stability increases from magnesium carbonate to barium carbonate. Therefore, progressively stronger heating is needed down the group.

Stability concerns resistance to thermal decomposition, not whether the carbonate is soluble in water.

3. Thermal decomposition of nitrates

Group 2 nitrates decompose on heating to the metal oxide, nitrogen dioxide and oxygen. A balanced general equation uses two nitrate formula units to form two oxide units, four nitrogen dioxide molecules and one oxygen molecule.

The nitrates also become more thermally stable down the group. Magnesium nitrate decomposes more readily than barium nitrate under comparable conditions.

Do not import the different nitrate products used for some Group 1 compounds.

4. Polarisation of the large anion

A small two-plus cation has high charge density and a strong electric field. It attracts electron density in a nearby carbonate or nitrate ion and distorts, or polarises, that large anion.

Polarisation weakens covalent bonding within the anion and makes decomposition easier. The magnesium ion therefore destabilises carbonate or nitrate more than the larger barium ion does.

Down Group 2, cation radius increases, polarising power decreases and anion distortion decreases. Thermal stability consequently increases.

5. Building the thermal-stability explanation

A complete trend explanation has four linked statements: ionic radius increases; charge density decreases because charge stays two-plus; polarising power falls; the large anion is less distorted and decomposes less easily.

Naming only stronger or weaker bonds is insufficient unless the cation-to-anion polarisation link is given.

The relevant radius is the metal cation radius, while carbonate and nitrate are the large polarisable anions.

6. Enthalpy change of solution cycle

For an ionic solid to dissolve, its lattice must be separated into gaseous ions and those gaseous ions must be hydrated by water.

Using a lattice-dissociation convention, enthalpy of solution equals positive lattice-dissociation enthalpy plus negative hydration enthalpies of the ions. Equivalently, a lattice-formation convention requires reversing its sign before adding hydration terms.

State the convention and use signs consistently. The syllabus asks for relative magnitudes, not a slogan that one term always wins.

7. Lattice energy down Group 2

As the metal ion grows, the distance between oppositely charged ions in the lattice increases. Electrostatic attraction weakens, so the magnitude of lattice energy generally decreases down the group.

The exact extent depends on the anion and crystal structure. Hydroxide and sulfate do not show identical numerical changes.

When using lattice-dissociation enthalpy, a smaller magnitude means less positive energy is needed to separate the lattice.

8. Hydration enthalpy down Group 2

Hydration is exothermic because ion-dipole attractions form between ions and water. A smaller, more charge-dense metal ion attracts water more strongly.

As Group 2 cation radius increases, metal-ion hydration enthalpy becomes less negative. The anion is unchanged within a hydroxide series or sulfate series, so its hydration contribution is broadly constant for the qualitative comparison.

Both lattice and hydration magnitudes decrease down the group, but not by the same amount.

9. Hydroxide solubility trend

Group 2 hydroxides become more soluble down the group. Magnesium hydroxide is sparingly soluble, while barium hydroxide is substantially more soluble.

For hydroxides, the decrease in lattice-energy magnitude down the group has the greater favourable effect relative to the simultaneous decrease in hydration-energy magnitude. The enthalpy of solution becomes more favourable overall.

Explain the competition. Saying only that larger ions have weaker lattices omits the opposing loss of hydration attraction.

10. Sulfate solubility trend

Group 2 sulfates become less soluble down the group. Magnesium sulfate is soluble, whereas barium sulfate is very sparingly soluble.

For sulfates, the hydration contribution becomes less exothermic down the group by an amount that is more important than the favourable weakening of the lattice. The enthalpy of solution becomes less favourable overall.

The same increasing cation radius therefore produces the opposite observed solubility trend because the relative changes differ for sulfate and hydroxide lattices.

11. Solubility and enthalpy of solution

Enthalpy of solution helps explain energetic favourability, but solubility is an equilibrium property and also depends on entropy. Cambridge's stated learning outcome asks for variation in both solubility and enthalpy of solution in terms of relative hydration and lattice magnitudes.

Do not claim that a negative enthalpy of solution automatically means unlimited solubility or that a positive value means no dissolution.

Use the known solubility trend, then explain how the changing energy competition supports it.

12. Comparing the two series

For both hydroxides and sulfates, larger metal ions weaken lattice attraction and metal-ion hydration. The decisive question is which magnitude changes more.

Hydroxides: lattice weakening dominates, so dissolution becomes more favourable and solubility rises. Sulfates: loss of hydration favourability dominates relative to lattice weakening, so dissolution becomes less favourable and solubility falls.

This comparison is stronger than memorising two arrows because it uses one common energetic framework.

Worked application: explaining four linked observations

A student is told that barium carbonate needs stronger heating than magnesium carbonate, barium hydroxide is more soluble than magnesium hydroxide and barium sulfate is less soluble than magnesium sulfate. The carbonate result follows because the larger barium ion has lower charge density, polarises carbonate less and therefore destabilises it less. For both solubility series, increasing cation radius reduces lattice and hydration magnitudes. In the hydroxides, favourable lattice weakening outweighs less favourable hydration, so solubility rises. In the sulfates, the hydration change is relatively more important than lattice weakening, so enthalpy of solution becomes less favourable and solubility falls. One radius trend explains all observations only after the competing terms are separated.

Common misconceptions and corrections

• Including beryllium in the required sequence. The stated comparison is magnesium to barium.
• Saying ionic charge rises down the group. It remains two-plus.
• Reversing the radius trend. Cation radius increases down Group 2.
• Calling thermal stability the same as solubility. They are different properties.
• Producing metal plus carbon dioxide from a carbonate. The solid product is the oxide.
• Omitting oxygen from nitrate decomposition. Balance all three products.
• Using a Group 1 nitrate rule. Group 2 nitrates form oxide, nitrogen dioxide and oxygen.
• Saying large cations polarise more. Smaller, charge-dense cations polarise more strongly.
• Claiming the cation is the large polarisable anion. Carbonate or nitrate is distorted.
• Stopping the explanation at charge density. Link it to anion distortion and decomposition.
• Saying stability decreases down Group 2. It increases for both nitrates and carbonates.
• Mixing lattice-formation and lattice-dissociation signs. Declare one convention.
• Calling hydration endothermic. Formation of ion-dipole attractions is exothermic.
• Saying hydration becomes more negative down the group. It becomes less negative.
• Assuming lattice magnitude increases with ion size. Greater separation weakens attraction.
• Explaining solubility with lattice energy alone. Hydration opposes or supports the trend.
• Saying both energy terms change equally. Their relative changes determine the outcome.
• Reversing hydroxide solubility. It increases down the group.
• Reversing sulfate solubility. It decreases down the group.
• Using the hydroxide explanation unchanged for sulfate. The dominant relative change differs.
• Equating exothermic dissolution with complete solubility. Entropy and equilibrium also matter.
• Claiming a positive solution enthalpy prevents dissolution. It does not decide equilibrium alone.
• Listing trends without the magnesium-to-barium direction. State both property and direction.

Assessment guidance

Thermal-stability answers should state the magnesium-to-barium direction and link increasing cation radius to lower charge density, weaker polarisation of nitrate or carbonate, less anion distortion and harder decomposition. Solubility explanations must discuss both lattice and hydration magnitudes under a consistent sign convention. State that both weaken down the group, then identify the relatively larger change: lattice weakening supports increasing hydroxide solubility, whereas reduced hydration favourability supports decreasing sulfate solubility. Keep enthalpy of solution distinct from solubility because equilibrium also contains an entropy contribution.

Retrieval practice

Balance the general carbonate and nitrate decomposition equations and explain their stability trend from cation radius to anion polarisation. Draw a solution-enthalpy cycle using one lattice convention, then predict how the metal-ion lattice and hydration terms change from magnesium to barium. Reconstruct the opposing hydroxide and sulfate solubility trends without memorised arrows, and explain why enthalpy of solution alone does not completely determine equilibrium solubility.

Return to the Cambridge International AS and A Level Chemistry notes hub.

Official source

Cambridge International, Chemistry 9701 syllabus for examinations in 2025, 2026 and 2027, section 27.1 for Similarities and Trends in the Properties of the Group 2 Metals, Magnesium to Barium, and Their Compounds.