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>> Compounds of Alkaline Earth Metals
The alkaline earth metals combine directly with halogen at appropriate temperature forming halides MX2.
These halides can also be prepared by the action of halogen acids (HX) on metals, metals oxides, hydroxides and carbonates.
M + 2HX —→ MX2 + H2
MO + 2HX —→ MX2 + H2O
M (OH)2 + 2HX —→ MX2 + 2H2O
MCO3 + 2HX —→ MX2 + H2O + CO2
(1) All beryllium halides are essentially covalent and are soluble in organic solvents. They are hydroscopic and fume in air due to hydrolysis. On hydrolysis, they produce acidic solution.
BeCI2 + 2H2O —→ Be (OH)2 + 2HCI
(2) The halides of all other alkaline earth metals are ionic. Their ionic character, however increases as the size of the metal ion increase.
(3) Except BeCl2 all other chlorides of group 2 form hydrates but their tendency to form hydrates decreases for eg – MgCl2.6H2O, CaCl2.6H2O.
(4) The hydrated chloride, bromides and iodides of Ca, Sr and Ba can be dehydrated on heating but those of Be and Mg undergo hydrolysis.
(5) BeF2 is very soluble in water due to the high hydration energy of the small Be+2ion. The other fluorides (MgF2, CaF2, SrF2 and BaF2) are almost insoluble in water. Since on descending the group lattice energy decreases more rapidly than the hydration energy. Therefore whatever little solubility these fluorides have that increase down the group.
The chlorides, bromides and iodides of all other elements i.e. Mg, Ca, Sr, Ba are ionic have much lower melting points than the fluorides and are readily soluble in water. The solubility decreases some what with increasing atomic number.
(6) Except of BeCl2 and MgCl2, the other chlorides of alkaline earth metals impart characteristics colour to flame.
CaCl2 = Brick red colour
BaCl2 = Grassy green colour
(i) Calcium fluoride or fluorospar (CaF2) is by far the most important of all the fluorides of the alkaline earth metals since it is the only large scale source of fluorine.
(ii) CaCl2 is widely used for melting ice on roads, particularly in very cold countries because 30% eutectic mixture of CaCl2/ice freezes at 218 K as compared to NaCl /ice at 255K.
(iii) CaCl2 is also used as a desiccant (drying agent) in the laboratory.
(iv) Anhydrous MgCl2 is used in the electrolytic extraction of magnesium.
Solubility and Thermal stability of oxo salts
The salts containing one or more atoms of oxygen such as oxides, hydroxides, carbonates, bicarbonates, nitrites, nitrates, sulphates, oxalates and phosphates are called oxo salts.
The sulphates of alkaline earth metals (MSO4) are prepared by the action of sulphuric acid on metals, metals oxides, hydroxides and carbonates.
M + H2SO4 —→ MSO4 + H2
MO + H2SO4 —→ MSO4 + H2O
M (OH)2 + H2SO4 —→ MSO4 + 2H2O
MCO3 + H2SO4 —→ MSO4 + CO2 + H2O
Properties of sulphates
The sulphates of alkaline earth metals are all white solids.
The solubility of the sulphates in water decreases down the groups i.e. Be > Mg > Ca > Sr > Ba.
Thus BeSO4 and MgSO4 are highly soluble, CaSO4 is sparingly soluble but the sulphates of Sr, Ba and Ra are virtually insoluble.
The magnitude of the lattice energy remains almost constant as the sulphate is so big that small increase in the size of the cation from Be to Ba does not make any difference. However the hydration energy decreases from Be+2 to Ba+2 appreciably as the size of the cation increase down the group. Hence, the solubilities of sulphates of alkaline earth metals decrease down the group mainly due to the decreasing hydration energies from Be+2 to Ba+2. The high solubility of BeSo4 and MgSO4 is due to high hydration energies due to smaller Be+2 and Mg+2 ions.
The sulphates of alkaline earth metal decompose on heating giving the oxides and SO3.
MSO4 —→ MO + SO3
The temperature of decomposition of these sulpahtes increases as the basicity of the hydroxide of the corresponding metal increase down the group