Limitations of Valence Bond Theory

Valence bond theory provides an explanation for the structure and magnetic characteristics of a large number of coordination compounds. The valence bond theory was used to explain the structure of coordination molecules and their bond connections.

When a metal atom or ion is impacted by ligands, it can utilise its (n–1)d, ns, np, nd orbitals for hybridization, resulting in a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, square planar, and so on. The orbitals of the ligands that can provide electron pairs for bonding may overlap with the hybridised orbitals.

1. N. Lewis hypothesised in 1916 that chemical bonds are formed by the interaction of two shared bonding electrons, and molecules are represented as Lewis structures. In 1921, chemist Charles Rugeley Bury proposed that electrons in shells with eight and eighteen electrons produce stable structures. Bury postulated that the electron configurations of transition metals were determined by the valence electrons in their outer shells. In 1916, Kossel proposed his idea of the ionic chemical bond (octet rule), which Gilbert N. Lewis independently expanded the following year.  Walther Kossel advanced a hypothesis comparable to Lewis’, except that his model presupposed complete electron transfers between atoms, and hence was a model of ionic bonding. Lewis and Kossel both built their bonding models around Abegg’s rule (1904).

A covalent link between two atoms is produced by the overlap of their half-filled valence atomic orbitals, each of which contains one unpaired electron. While a valence bond structure is similar to a Lewis structure, it is employed in situations where a single Lewis structure cannot be written. Each of these VB structures is a Lewis structure in its own right. The primary focus of resonance theory is on this combination of valence bond configurations. According to the valence bond theory, a chemical bond is formed when the overlapping atomic orbitals of the involved atoms. Due to the overlapping, electrons are most likely to be in the bond region. Bonds are viewed as weakly connected orbitals in the Valence bond theory (small overlap). In general, valence bond theory is easier to apply to ground-state molecules. During bond formation, the core orbitals and electrons stay basically unaltered.

The Valence Bond Theory’s Limitations

1. While it provides a qualitative depiction of the complex, it does not provide a quantitative interpretation of the complex’s stability.
2. It does not account for the complexes spectra (colour).
3. It predicts no distortion in symmetrical compounds, but predicts distortion in all copper (II) and titanium (III) complexes.
4. It provides no specific information regarding the complexes’ magnetic characteristics. It cannot, for example, account for the complexes’ temperature-dependent paramagnetism.
5. It does not explain why, at times, the electrons must be placed in contravention of Hund’s rule, while at other times, the electrical configuration remains unaltered.
6. It fails to account satisfactorily for the occurrence of inner and outer orbital complexes.
7. Occasionally, the theory requires electrons to be transferred from a lower energy level (Example 3d) to a higher energy level (4p), which is highly implausible in the absence of a source of energy.
8. Electron spin resonance demonstrates that the electron is not in the 4p level in Cu(II) complexes, indicating that the complex is planar.
9. It is unable to account for why some complexes are more labile than others. Complexes that are labile are those in which one ligand can be easily displaced by another. Inert complexes, on the other hand, are ones in which ligand displacement is sluggish.

Applications

The requirement of maximum overlap is a critical feature of the valence bond theory, as it results in the production of the strongest possible bonds. This theory is used to explain the development of covalent bonds in a wide variety of compounds.

For instance, in the F2 molecule, the F–F bond is produced by the overlap of the two F atoms’ pz orbitals, each of which contains an unpaired electron. Due to the fact that the nature of the overlapping orbitals in H2 and F2 molecules is different, the bond strength and bond lengths of H2 and F2 molecules are different.

The covalent bond in an HF molecule is produced by the overlap of the H 1s and F 2pz orbitals, each of which contains an unpaired electron. The mutual electron sharing between H and F leads to the formation of a covalent bond in HF.

Conclusion 

The Valence Bond Theory (VBT) examines the interaction of atoms in order to explain chemical bonding. It is one of two important ideas that contribute to our understanding of how atoms combine. The valence bond hypothesis explains how covalent bonds are formed. Additionally, it assists in determining the electrical structure of molecules. Additionally, one can explain the shape of an atom in a molecule using VBT and hybridization. VBT, on the other hand, is unable to account for the existence of inner and outer orbital complexes.