Chemistry QUANTUM MECHANICAL APPROACH TO COVALENT BONDING : VALENCE BOND THEORY

Valence Bond Theory :

According to this theory proposed by Heitler and London and developed by Linus Pauling and Slater.

`ast` For the formation of a covalent bond, a half-filled atomic orbital of one atom overlaps with the half-filled atomic orbitals of another atom. These atomic orbitals belong to the outermost shell of the atoms.

`ast` A covalent bond is formed by the overlapping of atomic orbitals having electrons with opposite spins.

`ast` The atomic orbitals containing paired electrons do not participate in the processes of overlapping. These electrons are called non-bonding electrons.

`ast` Due to the directional nature of most of the orbitals, overlapping is possible only when orbitals are properly oriented .

Sigma and Pi Bonding :

When two hydrogen atoms form a bond, their atomic orbitals overlap to produce a greater density of electron cloud along the line connecting the two nuclei. In the simplified representations of the formation of `H_2O` and `NH_3` molecules, the `O - H` and `N- H` bonds are also formed in a similar manner, the bonding electron cloud having its maximum density on the lines connecting the two nuclei. Such bonds are called sigma bonds (`sigma` -bond).

Limitation of Valance Bond Theory :

Llmitation 1 :

When the bonding scheme of certain polyatomic molecules is discussed, our simple valence bond theory fails to accurately describe what is going on. Methane is one of these molecules where we must expand upon our valence bond theory. The 2s orbital is filled with two electrons and the `2P_y` orbital has no electrons for use in bonding. This leaves only the `2P_x` and the `2p_z` orbitals as the orbitals that can participate in bonding .We know however, that in `CH_4`, four hydrogen atoms must make four bonds with the central carbon atom.

What happens is that the carbon atom promotes one of its 2s electrons to the `2p_y` oribital. This leaves one electron in the 2s orbital and one electron in the `2P_y` orbitai. AII together, the central carbon now has four unpaired valence electrons to bond to each of the
four hydrogen atoms. Thus out of four `C-H` bonds one should be different from remaining three `C-H ` bonds. But experimentally it is found that all four `C -H` in methane are identical all aspect.

Limitation 2 :

Another one of the shortcomings of the Valence Bond Theory is the inability to predict molecular geometries. If we take a look at the bonding orbitals of `H_2O`, only two of the `p` orbitals of sulphur bond with a ls orbital from `H`.The molecular geometry predicted by these two bonding orbitals is `90^(circ)`. This is obviously incorrect because the valence bond theory does not take into account the lone pair electrons on the central `O` atom. The two lone pair electrons create a ang le of `104.5^(circ)` between the two hydrogen atoms.