Limitations:
(i) Relative stability of different molecules cannot be explained.
(ii) Paramagnetic nature of complexes or molecules cannot be explained.
(iii) Different shape of molecules cannot be explained.
Hybridization:
Hybridization is the mixing of atomic orbitals prior to overlap that change the partial distribution of orbitals.
Due to this hybridization mainly two different kind of bonds forms, (i) sigma (σ ) and (ii) pi (π) bond.
Figure 2.18. Sigma (σ) and pi (π) bonding and anti - bonding by the combination of s and p orbitals.
Sigma (σ) bond: A covalent bond established between two atoms having the maximum density of the electron cloud along the axis connecting the centers of the two participating atoms is called sigma (σ) bond.
Pi (π) bond: A bond is formed by the lateral overlap between two atomic orbitals possessing maximum electron density on both sides of the overlapping axis is known as pi (π) bond.
Therefore, from the definition it is clear that sigma (σ) bonds are more strong than pi (π) bonds.
Sigma (σ) bonds may arises from the overlap between s, p, and d orbitals, like (i) s - s orbitals, (ii) s - p prbitals, (iii) p - p orbitals, (iv) s - d orbitals, so on.
s orbitals are spherical, while p , d orbitals are dumbbell shaped. Hence, s - s over lap is weaker compared to s - p , p - p overlaps.
**Herein, it is important to note that due to hybridization of two atomic orbitals two molecular orbitals formed. One is called bonding orbital and the other is called anti - bonding orbital. Hence, n number of atomic orbitals forms n /2 bonding and n /2 anti - bonding orbitals.
Examples:
In the formation of CH4 one s orbital and three p orbital of the central carbon atom participate.
Similarly, in the CO2 molecule one s orbital and three p orbital of the central carbon atom participate but here two p orbitals of carbon atom forms two π - bonds. Therefore, CH4 is sp 3 hybridized and CO2 is sp 2 hybridized. NOTE: Orbitals participate in π - bond formation is not counted to the hybridization.