General Characteristics of Compounds of the Alkaline Earth Metals

Group 2 elements show essentially only the +2 oxidation state ($M^{2+}$) in their compounds. Those compounds are still predominantly ionic, but noticeably less so than the analogous compounds of the alkali metals — the higher nuclear charge and smaller size of the $M^{2+}$ ion give it greater polarising power. This covalent contribution is most visible in beryllium and magnesium, whose oxides and other compounds are distinctly more covalent than those of the larger, heavier members calcium, strontium and barium.

1. Oxides and Hydroxides Burnt in oxygen, the Group 2 metals give the monoxide $MO$; every one of these except $BeO$ adopts the rock-salt (NaCl-type) crystal structure, while $BeO$ itself is essentially covalent. All the oxides have high (strongly exothermic) enthalpies of formation and are correspondingly very stable to heat. $BeO$ is amphoteric, but every other oxide in the group is basic and hydrolyses in water to the (sparingly soluble) hydroxide:

   $MO + H_2O \rightarrow M(OH)_2$

   Down the group from Mg(OH)₂ to Ba(OH)₂, the hydroxides become more soluble, more thermally stable and more strongly basic — though, taken as a whole, Group 2 hydroxides remain both less basic and less thermally robust than the corresponding Group 1 hydroxides. Beryllium hydroxide stands apart as genuinely amphoteric, dissolving in both acid and alkali:

   $Be(OH)_2 + 2OH^- \rightarrow [Be(OH)_4]^{2-} \quad (\text{beryllate ion})$

   $Be(OH)_2 + 2HCl + 2H_2O \rightarrow [Be(OH)_4]Cl_2$

2. Halides Apart from beryllium's, all the Group 2 halides are ionic. Beryllium halides, by contrast, are essentially covalent and dissolve in organic solvents; solid beryllium chloride adopts a chain structure (see cover) in which each Be centre bridges to its neighbours through chlorine atoms, giving every beryllium a four-coordinate (tetrahedral-style) environment. In the vapour phase $BeCl_2$ instead forms a chlorine-bridged dimer, which only breaks down into the simple linear monomer at high temperature (around 1200 K). The tendency of these halides to crystallise as hydrates falls off steadily down the group — for example $MgCl_2\cdot 8H_2O$, $CaCl_2\cdot 6H_2O$, $SrCl_2\cdot 6H_2O$ and finally $BaCl_2\cdot 2H_2O$. Heating drives off this water cleanly for the calcium, strontium and barium halides, but the corresponding hydrated halides of beryllium and magnesium instead undergo hydrolysis when heated. Across the group, the fluorides are consistently less soluble than the other halides, owing to their higher lattice energies.
3. Salts of Oxoacids Carbonates: All Group 2 carbonates are insoluble in water and can be precipitated straightforwardly by adding sodium or ammonium carbonate to a solution of the metal's soluble salt. Solubility falls as the metal ion gets larger, and every carbonate decomposes on heating to the oxide plus $CO_2$ — with thermal stability rising as cationic size increases. Beryllium carbonate is so unstable it can only be kept under an atmosphere of $CO_2$. …