Valence Bond Theory Localized Chemical Bonding

Valence Bond Theory (Localized Chemical Bonding)  - Structure and Bonding

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Valence Bond Theory

The valence bond (VB) theory of bonding is mainly based on the work of Heitler, London, Pauling, and Slater. According to this theory, a covalent bond is formed by the overlapping of two half-filled atomic orbitals containing electrons of opposite spins. For example, let us consider the formation of a hydrogen molecule (H₂) from two hydrogen atoms. When two hydrogen atoms come enough close to each other, a covalent bond is formed between them by overlapping of their half-filled Is orbitals containing electrons of opposite spins. The greater the overlapping, the stronger is the resulting bond. However, total overlapping is prevented by repulsion between the nuclei.

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After this overlapping, each of the two electrons can be found in the atomic orbital of one or the other atom and they can exchange t; their positions. Further, the probability of finding the two electrons of the two overlapped orbitals is much greater in the space between the two atomic nuclei than that in other places. Consequently, both the electrons are attracted by either nucleus, and the internuclear repulsion is also shielded by them. Thus, there is a decrease in the potential energy of the system which causes the overlapping of atomic orbitals to form a covalent bond. Thus, the resulting hydrogen molecule has 104 kcal mole less energy than the constituent hydrogen atoms. This is called the bond strength of the H-H bond.
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If the atoms come too close to each other, there are internuclear and interelectronic repulsions that increase the energy of the system. Thus, there is a critical internuclear distance, called bond length or bond distance, for the overlapping of atomic orbitals to form a covalent bond. This is 0.74 A for the formation of the hydrogen molecules. When this distance is less or more than 0.74 A, the formation of an H-H covalent bond cannot take place.
Similar to that described above for the formation of a covalent bond in the hydrogen molecule, one can apply the VB theory to the covalent bond formation in other molecules also. Non-spherical orbitals (e.g., p and d orbitals) show directional preference and tend to form bonds in the direction of maximum electron density within the orbital, hence a covalent bond is generally directional. The direction of the electron density in the bonds determines the bond angles.

The principles of overlap can be summarized as :

(i) Only those orbitals overlap which participate in the bond formation. Orbitals of the bonded atoms retain their individual identity.
(ii) The greater overlapping of the atomic orbitals leads to a lowering of energy due to the attractive forces between electrons and the nuclei between atomic orbitals. Thus the greater the overlap between atomic orbitals, the stronger is the covalent bond (This is known as the principle of maximum overlap).
(iii) The direction of the electron density in the bonds determines the bond angle.

Limitations of Valence Bond Theory: The following are chief drawbacks of VB theory :

(i) It has to introduce hypothetical concepts of hybridization, resonance, and hyperconjugation to explain the formation and structures of various molecules.
(ii) It does not explain why atomic orbitals of bonded atoms should retain their identity, whereas the nuclei of approaching atoms are bound to affect nearly all atomic orbitals of each other.
(iii) It does not explain the paramagnetic behavior of oxygen molecules.

Types of Overlapping

(i) s-s overlapping: Overlapping between s-s orbitals of two similar or dissimilar atoms is known as s-s overlapping. This forms a single covalent bond as shown below:

(ii) s-p-overlapping: Overlapping between sand p-orbitals is known as s-p overlapping.

(iii) p-p overlapping: p-p overlapping is produced by the end to end overlapping of two p-orbitals.

Multiple Bonds

If two atoms are joined by more than one bond then the bond is called a multiple bond. In multiple bonds there are two types of bonds known as sigma and pi bonds. 

Sigma (σ) Bond : A bond formed by coaxial overlapping of two atomic orbitals is known as cr bond. Since two atomic orbitals overlap along their axes, maximum overlap is possible. Hence, the bond thus formed is a strong bond. cr bonds are formed as a result of maximum overlapping of various atomic orbitals as outlined below* :