Periodic discussion and general characteristics of alkaline earth metals

Electronic configurations

Group 2 elements are called alkaline Earth metals. These elements have two electrons in the valence shell of their atoms, preceded by the noble gas configuration. Their general configuration is written as [Noble gas] ns² where 'n' represents the valence shell.

Electronic configuration of alkaline Earth metals

General Characteristics of Alkaline Earth Metals

Atomic and ionic radii

The atomic and ionic radii of alkaline Earth metals are smaller than the corresponding members of the alkali metals. This is due to the fact of Group 2 elements having a higher nuclear charge allow electrons to be attracted more towards the nucleus. This reduces the size of atomic and ionic radii.

On moving down the group, the radii increase due to gradual increase in the number of the shells and the screening effect.

Melting and boiling points

The melting and boiling point of alkaline Earth metals are characteristically low and do not show regular trends down the group. However these points are higher than the corresponding alkali metals in the same period as atoms of alkaline Earth metals have smaller size compared to alkali metals. This causes them to be more closely packed in their crystal lattices. They also have two electrons per metal atom in their valence shell (as compared to one for the alkali metals) forming strong metallic bonds for binding the atoms in the crystal lattice of the metal. Hence they have higher melting and boiling points.

Ionization energy

Due to a fairly large size of the atoms, alkaline Earth metals have low ionization energies when compared to 'p' block elements. Down the group the ionization energy decreases as atomic size increases. This is due to new shells being added and increase in the magnitude of the screening effect of inner shell electrons.

Members of group 2 have higher ionization energies values as compared to group 1 because of their smaller size, with electrons being more attracted towards the nucleus of the atoms. Correspondingly they are less electropositive than alkali members.

Although IE₁ values of alkaline Earth metals are higher than that of alkali metals, the IE₂ values of alkaline Earth metals are much smaller than those of alkali metals. This occurs because in alkali metals the second electron is to be removed from a cation, which has already acquired a noble gas configuration. In the case of alkaline Earth metals the second electron is to be removed from a monovalent cation, which still has one electron in the outermost shell. Thus, the second electron can be removed more easily in the case of group 2 elements than in group 1 elements.

Electropositive and metallic character

Because of low ionization energies, alkaline earth metals are reasonably electropositive in nature; among the family members the electropositive character in general increases from Be to Ba.

The alkaline Earth metals are not as strongly electropositive as the alkali metals of group 1 because of comparatively higher ionization energies.

Density

Alkaline Earth metals are denser and harder than the alkali metals of group.

1. The comparatively smaller atomic size of alkaline Earth metals leads to more closely packed crystal lattices and hence stronger metallic bonds. This accounts for their increased hardness and high density.

Characteristic flame coloration

Except for Be and Mg, the alkaline Earth (metal) salts impart characteristic colors to the flame. Since their ionization energies are low, the valence electrons in their atoms can be easily excited to higher but unstable energy state, by the flame of a bunsen burner. When these excited electrons come back to ground state they emit visible radiations, giving colors to the flame.

Beryllium and magnesium atoms being comparatively smaller have ionization energies that are very high. The energy of a Bunsen burner flame is not sufficient to excite their electrons to higher energy levels. These elements, therefore, do not give any color.

Oxidation states

Alkaline Earth metals exhibit a valency of +2 as they can lose two electrons and form bivalent ions, which are stable noble gas configurations. Thus, unlike alkali metals, alkaline Earth metals exhibit +2 oxidation states in their compound.

The second ionization energy (IE₂) of alkaline Earth metals is greater than the first ionization energy (IE₁). We would expect a preference to form +1 ions (M⁺) rather than + 2 ions (M²⁺) by alkaline Earth metals. However, they predominantly show +2 valency, e.g., Mg²⁺, Ca²⁺, Ba²⁺ etc.

This can be explained by the fact that:

• Divalent ions have the stable noble gas configuration.

• In aqueous solution, the +2 ions of alkaline Earth metals are extensively hydrated due to their smaller size and the high hydration energies of M²⁺ ions, making them more stable than M⁺ ions. The amount of energy released when M²⁺ ion is dissolved in water is much more than that for M⁺ion and this large amount of extra energy is more than sufficient to compensate for the high second ionization energy required for the formation of such ions.

• Divalent cations form stronger lattices than monovalent cations in the solid state and a lot of energy of M²⁺ ion, called lattice energy, is released. It is the greater lattice energy that compensates for the high second ionization energy and is responsible for the existence and greater stability of bivalent ions in the solid state, as compared to M⁺ ion.

Reduction properties

Due to their low ionization energies all group 2 elements tend to lose their valence electrons and act as strong reducing agents. The decrease of the reduction potential values down the group, indicate an increase of strength as a reducing agent. However the members of this group are weaker reducing agents than alkali Earth metals because of the latter's higher ionization energy.