Compounds of Alkaline Earth Metals



Oxides and hydroxides of alkaline earth metals


The alkaline earth metals burn in oxygen to form the monoxide, MO which, except for BeO, have rock-salt structure. The BeO is essentially covalent in nature. The enthalpies of formation of these oxides are quite high and consequently they are very stable to heat. BeO is amphoteric while oxides of other elements are ionic in nature. All these oxides except BeO are basic in nature and react with water to form sparingly soluble hydroxides.

MO + H2O → M(OH)2

The solubility, thermal stability and the basic character of these hydroxides increase with increasing atomic number from Mg(OH)2 to Ba(OH)2. The alkaline earth metal hydroxides are, however, less basic and less stable than alkali metal hydroxides. Beryllium hydroxide is amphoteric in nature as it reacts with acid and alkali both.

Be(OH)2 + 2OH → [Be(OH)4]2–

     Beryllate ion

Be(OH)2 + 2HCl + 2H2O → [Be(OH)4]Cl2


Halides of alkaline earth metals


Except for beryllium halides, all other halides of alkaline earth metals are ionic in nature. Beryllium halides are essentially covalent and soluble in organic solvents. Beryllium chloride has a chain structure in the solid state as shown below:

In the vapour phase BeCl2 tends to form a chloro-bridged dimer which dissociates into the linear monomer at high temperatures of the order of 1200 K. The tendency to form halide hydrates gradually decreases (for example, MgCl2·8H2O, CaCl2·6H2O, SrCl2·6H2O and BaCl2·2H2O) down the group. The dehydration of hydrated chlorides, bromides and iodides of Ca, Sr and Ba can be achieved on heating; however, the corresponding hydrated halides of Be and Mg on heating suffer hydrolysis. The fluorides are relatively less soluble than the chlorides owing to their high lattice energies.


Salts of oxo-acids of alkaline earth metals

Carbonates of alkaline earth metals

Carbonates of alkaline earth metals are insoluble in water and can be precipitated by addition of a sodium or ammonium carbonate solution to a solution of a soluble salt of these metals. The solubility of carbonates in water decreases as the atomic number of the metal ion increases. All the carbonates decompose on heating to give carbon dioxide and the oxide. Beryllium carbonate is unstable and can be kept only in the atmosphere of CO2. The thermal stability increases with increasing cationic size.


Sulphates of alkaline earth metals

The sulphates of the alkaline earth metals are all white solids and stable to heat. BeSO4, and MgSO4 are readily soluble in water; the solubility decreases from CaSO4 to BaSO4. The greater hydration enthalpies of Be2+ and Mg2+ ions overcome the lattice enthalpy factor and therefore their sulphates are soluble in water.


Nitrates of alkaline earth metals

 The nitrates are made by dissolution of the carbonates in dilute nitric acid. Magnesium nitrate crystallises with six molecules of water, whereas barium nitrate crystallises as the anhydrous salt. This again shows a decreasing tendency to form hydrates with increasing size and decreasing hydration enthalpy. All of them decompose on heating to give the oxide like lithium nitrate.


(M = Be, Mg, Ca, Sr, Ba)


Anomalous behaviour of beryllium

 Beryllium, the first member of the Group 2 metals, shows anomalous behaviour as compared to magnesium and rest of the members. Further, it shows diagonal relationship to aluminium which is discussed subsequently.

(i) Beryllium has exceptionally small atomic and ionic sizes and thus does not compare well with other members of the group. Because of high ionisation enthalpy and small size it forms compounds which are largely covalent and get easily hydrolysed.

(ii) Beryllium does not exhibit coordination number more than four as in its valence shell there are only four orbitals. The remaining members of the group can have a coordination number of six by making use of d-orbitals.

(iii) The oxide and hydroxide of beryllium, unlike the hydroxides of other elements in the group, are amphoteric in nature.


Diagonal Relationship between beryllium and aluminium

 The ionic radius of Be2+ is estimated to be 31 pm; the charge/radius ratio is nearly the same as that of the Al3+ ion. Hence beryllium resembles aluminium in some ways. Some of the similarities are:

(i) Like aluminium, beryllium is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.

(ii) Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion, [Be(OH)4]2– just as aluminium hydroxide gives aluminate ion, [Al(OH)4].

(iii) The chlorides of both beryllium and aluminium have Cl bridged chloride structure in vapour phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel Craft catalysts.

(iv) Beryllium and aluminium ions have strong tendency to form complexes, BeF42–, AlF63–.


Compounds of Calcium

Calcium Oxide

By the thermal decomposition of calcium carbonate.



1. It is a basic oxide.
2. Its aqueous suspension is known as slaked lime


3. On heating with ammonium salts, it gives ammonia.


4. It reacts with carbon to form calcium carbide


5. It is used as basic flux, for removing hardness of water, for preparing mortar (CaO + sand + water).


Calcium Hydroxide


By dissolving quicklime in water.

CaO + H2O Ca(OH)2;

ΔH = – 63 kj


1. Its suspension in water is known as milk of lime.

2. It gives CaCO3 (milky) and then Ca(HCO3)2 with CO2,


3. It reacts with Cl2 to give bleaching powder, CaOCI2

Ca(0H)2 + Cl2 → CaOCI2 + H2O


Calcium Carbonate


By passing CO2 through lime water.

Ca(0H)2 + CO2 → CaCO3 ↓ + H2O


It is insoluble in H2O but dissolves in the presence of CO2, due to the formation of calcium bicarbonate.


Calcium Sulphate

Gypsum, Calcium Sulphate Dihydrate (CaSO4 * 2H2O)

It is also known as alabaster.

On heating at 390 K, it gives plaster of Paris.

It is added to cement to slow down its rate of setting.

Plaster of Paris or Calcium Sulphate Hemihydrate (CaSO4 * 1 / 2 H2O)

When it is mixed with water, it forms first a plastic mass which sets into a solid mass with slight expansion due to dehydration and its reconversion into gypsum. It is obtained when gypsum is heated at 393 K.

CaSO4 * 2H2O → CaSO4 * 1 / 2 H2O + 3 / 2 H2O

Above 393 K no water of crystallization is left and anhydrous calcium sulphate is obtained. It is known as dead burnt plaster.



Cement is essentially a mixture of complex silicates and aluminates of Ca containing less than 1.0% free lime and some gypsum (CaSO4.2H2O)



An approximate composition is as follows :

















Iron oxide




Sulphur trioxide




Sodium oxide




Potassium oxide




Ratio of Silica and alumina


Ratio of CaO and 6(SiO2 + Al2O3 + Fe2O3)


White Cement : It does not contain ferric oxide


Two processes are employed (i) Wet process (ii) Dry process

Raw material : Lime and Clay



Clay + Lime Cement clinker Cement

Gypsum regulates the setting time



When mixed with water, the cement forms a gelatinous mass sets to hard mass when three dimensional cross links are formed between ... Si-O-Si---and ---Si-O-Al--- chains.

The reactions involved in the setting of cement are :

Hydration: Hydration of 3CaO.Al2O3 and 2CaOSiO2 forming colloidal gel.

Hydrolysis: Hydrolysis of 3CaOAl2O3 and 3CaO.SiO2 forming precipitates of Ca(OH)2 and Al(OH)3

Fly ash: A waste product of steel industry possess properties similar to cement. It is added to cement to reduce its cost.

Rice Husk: It has high silica content and employed to make cement.


Biological importance of magnesium and calcium


 An adult body contains about 25 g of Mg and 1200 g of Ca compared with only 5 g of iron and 0.06 g of copper. The daily requirement in the human body has been estimated to be 200 – 300 mg.


All enzymes that utilise ATP in phosphate transfer require magnesium as the cofactor. The main pigment for the absorption of light in plants is chlorophyll which contains magnesium. About 99 % of body calcium is present in bones and teeth. It also plays important roles in neuromuscular function, intraneuronal transmission, cell membrane integrity and blood coagulation. The calcium concentration in plasma is regulated at about 100 mg/L. It is maintained by two hormones: calcitonin and parathyroid hormone.