Valence Bond Theory

Many coordination compounds exhibit structures and magnetic properties that are explained by valence bond theory. Valence bond theory is used to explain the bond linkages of coordination compounds.

By using its orbitals for hybridisation, a metal atom or ion can create a set of equivalent orbitals having the same geometry, such as octahedral, tetrahedral, square planar, etc., when it is influenced by ligands. Hybridised orbitals and ligand orbitals can overlap when electron pairs may be required for bonding.

Valence Bond Theory in Coordination Compounds

A major contribution of Pauling was the valence bond theory of bonding in complexes. Among the three theories, it is the simplest one and provides satisfactory explanations for the structure and properties of a large number of coordination compounds.

Valence Bond Theory Applications

1. The valence bond theory can explain how covalent bonds form in several molecules by describing the maximum overlap condition.
2. The theory can help explain how covalent bonds form in several molecules. As an example, the difference in the overlapping orbitals in H2 and F2 molecules can be explained by the differences in the length and strength of their chemical bonds.
3. The valence bond theory states that the covalent bond in an HF molecule is formed by the overlap of the hydrogen 1s orbital and the fluorine 2p orbital, which is the reason for the covalent bond.

History of the Valence Bond Theory

A Lewis approach to chemical bonding was not successful in shedding light on chemical bonding. Further, the valence shell electron pair repulsion theory (or VSEPR theory) is not applicable to complex molecules (and fails to predict their geometry).

Fritz Wolfgang London and Walter Heinrich Heitler proposed the valence bond theory to address these issues. As well as explaining covalent bonds, the Schrodinger wave equation was also used to explain the formation of hydrogen bonds between two atoms of hydrogen.

The concept of electronic configurations, atomic orbitals and their overlaps, and hybridisation of these orbitals is central to the theory. In chemical bonds, electrons are distributed in the corresponding bond region due to the overlapping of atomic orbitals.

In addition to explaining the electronic structure of molecules formed by the overlapping of atomic orbitals, furthermore, this makes it clear that the nucleus of one atom in a molecule attracts the electrons of all the other atoms in that molecule.

Postulates of Valence Bond Theory

 Listed below are some of the important postulates of valence bond theory.

1. Bonds between covalently bound atoms occur when the half-filled valence orbitals of two different atoms overlap. Due to the overlapping between the two bonding atoms, the electron density in the area between increases, resulting in a more stable molecule.
2. An atom’s valence shell contains a large number of unpaired electrons, which enable it to form multiple bonds with other atoms. As defined by the valence bond theory, paired electrons present in the valence shell are not involved in chemical bond formation.
3. Chemical bonds that are covalent are directional and parallel to the atom’s overlapping orbitals.
4. Sigma bonds and pi bonds differ in how their atomic orbitals overlap, i.e., pi bonds are formed when the orbitals overlap side by side, whereas sigma bonds form when the orbitals overlap along the axis containing the nuclei of the two atoms.

Limitations of Valence Bond Theory

1. In spite of the fact that it provides a qualitative picture of the complex, it does not provide a quantitative understanding of the stability of the complex.
2. The spectra (colour) of the complexes cannot be explained by it.
3. In symmetrical complexes, it does not predict any distortions, whereas all copper (II) and titanium (III) complexes have distorted structures.
4. The temperature-dependent paramagnetism of the complexes cannot be explained by this model.
5. It does not explain why at one time the electrons must be arranged against Hund’s rule, while at other times, the electronic configuration is not disturbed.
6. The existence of inner and outer orbital complexes is not well explained by this theory.

7. In some cases, the theory requires electrons to be transferred from a lower energy level (Example. 3d) to the higher energy level (4p).

8. Nuclear spin resonance demonstrates that (4p) is very small in Cu(II) complexes. 

9. There is no explanation for why some complexes are more labile than others. Complexes classified as labile are those in which a ligand can easily be displaced by another. Complexes that are inert, however, are those whose ligand displacement is slow.

Conclusion

The Valence Bond Theory (VBT) looks at the interaction between atoms to explain chemical bonding. Atoms can be explained by this theory in part because it is one of the two popular theories. Bonds between covalent compounds can be explained by the theory of valence bonds. Valence bond theory can also help determine a molecule’s electronic structure. Additionally, hybridization and VBT can provide an explanation of the geometry of individual atoms within a molecule. Yet, VBT is unable to explain the existence of orbital complexes in the inner and outer orbits of the planet.