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Alkaline Earth Metals

  1. Reducing agents: The high negative value of standard electrode potential indicates that in aqueous solution these elements are good reducing agents, quite comparable to alkali metals, and this is due to their great hydration energies. The high negative E° values of these elements mean that all react vigorously with water also.
  2. Coloration to the flame, except Be and Mg: The chlorides of these elements produce characteristic flames due to easy excitation of electrons to higher energy levels.
  3. Conductors of heat and electricity: All alkaline earth metals are good conductor of heat and electricity.
  4. Size of atoms and ions: Group II atoms are large, but are smaller than the corresponding group I elements as the extra charge on the nucleus draws the orbital electrons in. Similarly, the ions are large, but are smaller than those of group I, especially because the removal of two orbital electrons increases the effective nuclear charge even further. Thus, these elements have higher densities than group I metals.
  5. Solubility and lattice energy: The solubility of most salts decreases with increased atomic weight, though the usual trend is reversed with the fluorides and hydroxides in this group. Solubility depends on the lattice energy of the solid and the hydration energy of the ions. With most compounds, on descending the group, the hydration energy decreases more rapidly than the lattice energy; hence the compounds become less soluble as the metal gets larger. However, with fluorides and hydroxides the lattice energy decreases more rapidly than the hydration energy, and so their solubility increases on descending the group.
  6. Solutions of the metals in liquid ammonia: These metals dissolve in liquid ammonia as do the group I metals. Dilute solutions are bright blue in color due to the spectrum from the solvated electron. These solutions decompose very slowly, forming amides and evolving hydrogen, but the reaction is accelerated by many transition metals and their compounds.
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    Evaporation of the ammonia from the solution of group I metals yields the metal, but with group II metals, evaporation of ammonia gives hexammoniates of the metals. These slowly decompose to give amides.
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    Concentrated solutions of the metals in ammonia are bronze colored due to the formation of metal clusters.

Oxides and peroxides

  1. Action of oxygen/air: All the metals of this group (except Be and Mg) are easily oxidized by the atmospheric oxygen. Barium readily inflames in air. All alkaline earth metals have affinity towards oxygen.
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    • Nature of oxides: BeO is amphoteric, MgO is weakly basic, CaO is more basic, while SrO and BaO are extremely basic.
    • Solubility: BeO and MgO are insoluble in water, while the other oxides react with water to give corresponding hydroxides of the type M(OH)2. BeO and MgO are insoluble in water due to their large lattice enthalpies.
  2. Sulphates: The sulphates of alkaline earth metals are less soluble than the sulphate of corresponding alkali metal. The solubility of the sulphates of alkaline earth metals decreases in going down the group
    The lattice enthalpies of alkaline earth metal sulphates are higher than those of the alkali metal sulphates. This is why the sulphates of alkaline earth metals are less soluble than those of alkali metals.
  3. Halides: The anhydrous halides are polymeric. Beryllium chloride vapor contains BeCl2 and (BeCl2)2, but the solid is polymerized.
    Both BeCl2 and Al2Cl6 are covalent and have a bridged polymeric structure. Both these chlorides are soluble in organic solvents and act as a strong Lewis acid.
  4. Carbides: The carbides of beryllium (Be) and aluminium (Al) react with water to give
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