10.1 Trends in Group 2 metals and compounds

Group 2 metals form oxides when heated in oxygen: 2M + O₂ → 2MO. The products contain M²⁺ and O²⁻ ions in an ionic lattice.

Reactivity generally increases down the group because the outer electrons are farther from the nucleus and more shielded, so they are easier to remove.

Magnesium burns in oxygen to form MgO. Barium reacts more readily than magnesium under comparable conditions because its outer electrons are less strongly held.

A more reactive metal does not necessarily produce a different oxidation state here. Group 2 metals commonly form +2 ions; the trend is about ease of reaction.

Group 2 oxides are basic. Oxides react with water to form hydroxides, and both oxides and hydroxides neutralise acids to form salts and water.

Hydroxide solubility generally increases down the group, so alkaline solutions become easier to obtain. Keep solubility separate from the intrinsic basicity of the oxide lattice.

CaO + H₂O → Ca(OH)₂. The hydroxide then reacts with HCl: Ca(OH)₂ + 2HCl → CaCl₂ + 2H₂O.

Do not treat all Group 2 hydroxides as equally soluble. The reaction can occur at the surface even when the hydroxide is only sparingly soluble.

Heating Group 2 carbonates gives a metal oxide and carbon dioxide; heating nitrates gives a metal oxide, nitrogen dioxide and oxygen. Thermal stability generally increases down the group.

Larger M²⁺ ions have lower charge density and polarise the anion less, so the carbonate or nitrate ion is less distorted and decomposes less readily.

MgCO₃ decomposes more readily than BaCO₃. The products follow MgCO₃ → MgO + CO₂ and 2M(NO₃)₂ → 2MO + 4NO₂ + O₂.

Do not explain the trend only by “stronger bonds down the group”. The key comparison is cation polarising power and anion stability.

Down Group 2, ionic radius increases while charge remains +2. Charge density therefore decreases, affecting hydration, lattice interactions, solubility and thermal stability.

Link a prediction to the relevant process: lower polarisation stabilises oxyanions, while hydration and lattice-energy changes determine solubility trends.

Ba²⁺ is larger than Mg²⁺, so BaCO₃ is more thermally stable. Hydroxide solubility rises down the group, whereas sulfate solubility falls.

Not every property changes in the same direction. State which competing energies control the particular trend instead of applying one arrow to all compounds.

Group 2 hydroxide solubility generally increases down the group, while sulfate solubility generally decreases. Each trend reflects a balance between lattice energy and hydration energy.

A dissolving solid must pay lattice energy and gain hydration energy. The relative change down the group is different for OH⁻ and SO₄²⁻, so do not assume one trend applies to both.

Ba(OH)₂ is more soluble than Mg(OH)₂, but BaSO₄ is far less soluble than MgSO₄. This contrast is useful in qualitative analysis.

“Larger ion means always more soluble” is false. Solubility is a free-energy balance, not a single-size rule.