Alkaline Earth Metals (Grp 2)

Both groups show increasing reactivity, atomic radius, and ionic radius down the group, and decreasing ionisation energy and electronegativity. However, Group 2 metals have two valence electrons and a 2+ ionic charge, giving them higher charge density than Group 1. This means Group 2 metals are harder (stronger metallic bonding from two delocalised electrons per atom), have higher melting points (Mg: 650 °C vs Na: 98 °C), and are less reactive with water — Mg reacts only slowly with cold water, whereas Na reacts vigorously. Group 2 first ionisation energies are lower than Group 1 of the same period (due to larger atomic radius from additional shielding), but second ionisation energies remain accessible, allowing stable M²⁺ ions. The higher charge density of M²⁺ compared to M⁺ has profound consequences for lattice enthalpies and hydration enthalpies of salts.

The hydroxides of Group 2 show a clear trend: solubility increases down the group. Be(OH)₂ is almost insoluble, Mg(OH)₂ is sparingly soluble (used as milk of magnesia), Ca(OH)₂ is slightly soluble (limewater), Sr(OH)₂ is moderately soluble, and Ba(OH)₂ is the most soluble. This trend is explained using enthalpy of solution (ΔH_sol = lattice enthalpy + hydration enthalpy). As ionic radius increases down the group, lattice enthalpy decreases in magnitude (ions are further apart, so the lattice is less stable and requires less energy to break). Hydration enthalpy also decreases, but it decreases more slowly than lattice enthalpy for hydroxides because the OH⁻ ion is relatively small and its hydration enthalpy does not change much. The net result is that the overall enthalpy of solution becomes less endothermic (more energetically favourable) going down the group, so solubility increases. Entropy effects also favour dissolution.

The sulphates show the opposite trend: solubility decreases down the group. MgSO₄ (Epsom salt) is very soluble, CaSO₄ is sparingly soluble (gypsum), SrSO₄ is poorly soluble, and BaSO₄ is virtually insoluble — a fact exploited in the barium meal X-ray procedure and in the gravimetric test for sulphate ions. The explanation again lies in the relative rates of decrease of lattice enthalpy and hydration enthalpy. The SO₄²⁻ ion is large; therefore its hydration enthalpy is relatively small and does not change much as the cation changes. However, as the M²⁺ cation grows larger down the group, its hydration enthalpy decreases significantly. The lattice enthalpy also decreases but not as fast as the hydration enthalpy, so the enthalpy of solution becomes increasingly endothermic (unfavourable), and solubility falls. The key insight: for sulphates, the dominant change is in the cation hydration enthalpy; for hydroxides, the dominant change is in the lattice enthalpy.

All Group 2 carbonates decompose on heating to give the metal oxide and carbon dioxide: MCO₃(s) → MO(s) + CO₂(g). The temperature required for decomposition increases markedly down the group: MgCO₃ decomposes at ~300 °C, CaCO₃ at ~840 °C, SrCO₃ at ~1100 °C, and BaCO₃ at ~1450 °C. This trend is explained by the polarising power of the M²⁺ cation. A small, highly charged cation (like Mg²⁺, charge density = charge/radius²) has high polarising power: it distorts the electron cloud of the carbonate ion, weakening the C–O bonds and making decomposition easier at lower temperatures. As the ionic radius increases down the group, polarising power decreases, the carbonate ion is less distorted, C–O bonds are harder to break, and a higher temperature is needed. The same argument explains why Group 1 carbonates (larger, lower-charge-density cations) are even more thermally stable — Li₂CO₃ decomposes at ~720 °C but Na₂CO₃, K₂CO₃, Rb₂CO₃, and Cs₂CO₃ do not decompose below 800 °C under normal conditions, with stability increasing further down Group 1.

• \ce{BaCl2} — Barium chloride (aq): Soluble Group 2 salt used as the test reagent for sulphate ions; always acidified with HCl to prevent false positives from BaCO₃ or BaSO₃.
• \ce{BaSO4} — Barium sulphate (s): White precipitate; virtually insoluble due to the large, low-hydration-enthalpy Ba²⁺ cation paired with the large SO₄²⁻ anion — a classic demonstration of the decreasing solubility of Group 2 sulphates.
• \ce{Na2SO4} — Sodium sulphate (aq): Source of sulphate ions; the sodium ions are spectator ions and do not participate in the net ionic equation \ce{Ba^{2+}(aq) + SO4^{2-}(aq) -> BaSO4(s)}.

Problem. Explain why barium sulphate is virtually insoluble in water, whereas magnesium sulphate is very soluble, using the concepts of lattice enthalpy and hydration enthalpy. Then predict whether barium hydroxide or magnesium hydroxide is more soluble, giving a reason.

Answer. BaSO₄ is virtually insoluble because the low hydration enthalpy of the large Ba²⁺ ion cannot compensate for the energy needed to break the lattice. MgSO₄ is soluble because the high hydration enthalpy of Mg²⁺ compensates adequately. Ba(OH)₂ is more soluble than Mg(OH)₂ because the lattice enthalpy of Ba(OH)₂ is much smaller (weaker lattice), while the hydration enthalpy of OH⁻ is nearly constant, so dissolution is more energetically favourable.

• Alkali Metals (Group 1) — Group 1 metals share the s-block configuration and many periodic trends with Group 2 but differ in charge density and reactivity — understanding both groups together clarifies how charge and ionic size drive chemical behaviour.
• Lattice Enthalpy and Born-Haber Cycles — The solubility trends of Group 2 sulphates and hydroxides are explained quantitatively using lattice enthalpy and hydration enthalpy values, which are best understood through Born-Haber thermodynamic cycles.
• Polarisation and Fajans' Rules — The thermal stability of Group 2 carbonates depends on the polarising power of the cation — a concept formalised in Fajans' Rules, which predict when ionic bonding departs from ideal behaviour.
• Periodic Trends: Ionisation Energy — Group 2 first and second ionisation energies explain why these metals form 2+ ions rather than 1+ or 3+ ions, and why their reactivity differs from Group 1.