Group IIA: The Alkaline Earth Metals - Preparation, Reactions, Abundance of alkaline earth metals

  Group 2A: The Alkaline Earth Metals (Group IIA)

Electronic Configuration of Alkaline earth Metals (Group IIA metals)

Element	symbol	electron configuration
Beryllium	Be	[He]2s2
Magnesium	Mg	[Ne]3s2
Calcium	Ca	[Ar]4s2
Strontium	Sr	[Kr]5s2
Barium	Ba	[Xe]6s2
Radium	Ra	[Rn]7s2

 

These elements are never found in metallic form in nature.

 

Abundance of Alkaline earth metals:

§  Magnesium is the second most abundant metallic element in the sea.

§  Calcium is found abundantly as CaCO₃ in marble, limestone, and chalk.

§  The most common source of beryllium is the mineral beryl Be₂SO₄.

 

Preparation of alkaline earth metals (Methods of Preparation):

§  By direct reduction of magnesium oxide by carbon

§  By reduction of their oxides by reducing metals.

Properties of Alkaline earth metals

 

1)    Hardness:

  The group IIA metals are all harder than the alkali metals this because the metallic bonding in alkaline earth metals is higher than in alkali metals. The trend of increasing softness with increasing atomic number occurs in group IIA.

 

2)    The melting and boiling points:

  The melting and boiling temperatures and the enthalpies of vaporization of the group IIA metals are much higher than those of the alkali metals, this also because the metallic bonding in alkaline earth metals is higher than in alkali metals.

 

3)    The atomic and the ionic radii

The atomic radii and the ionic radii of the group IIA elements increase as atomic number increases.

The atomic radii and the ionic radii are smaller than the corresponding elements in group IA. The decrease in ionic size in going from group IA to group IIA has a simple explanation. The ion pairs Li⁺ and Be⁺² , Na⁺ and Mg⁺², K⁺ and Ca⁺² and so on, are isoelectronic, each pair has the same electronic structure, but The alkaline earth has a higher nuclear charge than the corresponding alkali metal ion and this factor causes the decreased size of the alkaline earth ion.

 

4)  Enthalpy of Hydration (Heat of Hydration): Reactions of alkaline earth metals with water

Enthalpy of Hydration: It is the heat evolved on hydration of the ions. The hydration enthalpy becomes smaller as the size of the ion increases.
The heats of hydration of Alkaline earth metals are higher than Alkali metals. Due to the increased charge on the alkaline earth ions. For example Na⁺ and Ca⁺² have nearly the same radii, but the enthalpy of hydration of Ca⁺²  is nearly four times that of Na⁺. This shows that the hydration enthalpy is proportional to the square of the charge on the ion.
This increase in hydration enthalpy with ionic charge that is responsible for the existence of the group IIA ions in the +2 oxidation state only in aqueous solutions. Because the gained hydration energy compensates for the extra energy required to remove the second electron.

 

5)    Ionization energy:

 The ionization energies of the elements are higher than group IA elements.

Chemical Properties of Alkaline Earth Metals (Reactions of Alkaline Earth Metals)

The Oxides and Hydroxides of Alkaline Earth Metals:

 The oxides of magnesium and the heavier group IIA elements can be prepared by direct combination of the elements, or by thermal decomposition of the carbonates:

Where M = Mg, Ca, Sr, or Ba.

 The oxides of group IIA elements (except Be) are basic oxides, react with water to from hydroxides of the general formula M(OH )₂. Their hydroxides are strong bases.

 The solubility of alkali metal hydroxides in water are limited but it increases with increasing atomic number of cations i.e. solubility order is:

     Be(OH)₂  <  Mg(OH)₂   <  Ca(OH)₂   <  Sr(OH)₂   <  Ba(OH)₂  

Beryllium oxide:

Beryllium oxide is harder and higher melting than the oxides of the heavier metals.

The acid-base behavior of beryllium oxide is amphoteric, it reacts with concentrated strong acids to give solutions of hydrated ions Be(H₂O)₄⁺², and also reacts with strong bases and forms Be(OH)₄^-2.

The amphoteric behavior of beryllium oxide is due to:

a)     Small size and

b)    Large charge of the beryllium ion.

Consequently the beryllium ion has high polarizing power and polarizes surrounding molecules by drawing electrons from them. Thus the situation in Be(H₂O)₄²⁺ might be as represented in the Figure.

The beryllium ion withdraws electronic charge from the surrounding water molecules (four water molecules) and thereby facilitate removal of their protons and the formation of Be(OH)₄^2-.Thus Be(H₂O)₄²⁺  is an acid while Be(OH)₄^2- is a base.

 

The Halides Alkaline Earth Metals:

All the group IIA metals combine directly with halogens to form metal halides of the general formula MX₂.

The bond in chlorides change from covalent to ionic bonding as the atomic number of the metal increases. So, the bond in BeCl₂ is covalent and has the following properties:

-         lower melting and boiling points

-         low electrical conductivity

-         soluble in some organic solvents.

     The appearance of covalent character in the beryllium halides is due to small, highly charged Be ion which lead to high polarizing power.

     In the gas phase, BeCl₂ is a liner, symmetric molecule and Be atom is sp hybridized. While in the solid phase, BeCl₂ displays the bridged structure shown in the Fig. Each beryllium atom is surrounded by four chlorine atoms, with a Cl-Be-Cl angle of 90. This bridge occurs using the lone pair of chloride. This type of bridge structure also occurs in the aluminum and gallium halides.

The chlorides, bromides, and iodides of the heavier alkaline earth metals arc ionic solids that are soluble in water.

 Other Salts: Chromate, Sulphate, and Hydroxide:

The solubility of sulfates and chromates of the alkaline earth metals decrease as the atomic number of the metal ion increases, and this behavior is opposite to that observed for hydroxides (Explain Why?)

The reason:

For Chromate and sulphate:

As we proceed in down from BeSO₄ to BaSO₄, the hydration enthalpy for the cation tends to be smaller due to increase in ionic radius. This dcrease the solubility of the salts of the heavier metal ions. However, the lattice energies of the metal sulfates and metal chromates, do not change remarkably when moving from beryllium to barium, because of the dependence of lattice energies on radius. Thus, the trend in the solubilities of the group IIA metal salts with large size anions such as SO₄^-2 and CrO₄^-2 is affected by the trend in the enthalpies of hydration.

For Hydroxides:

Since OH^- is a small ion, its lattice energy should change greatly as the size of the cation changes Therefore the trend in the solubilities depend on the lattice and hydration energy.