Valence Bond Theory
There have been numerous approaches proposed to explain the nature of bonding in coordination compounds. The Valence Bond Theory is one of them. The Valence Bond Theory was developed to explain chemical bonding using the quantum mechanics method. This theory is primarily concerned with the formation of individual bonds from the atomic orbitals of the atoms involved in the formation of a molecule.
The electrons in a molecule, according to the valence bond theory, occupy atomic orbitals rather than molecular orbitals. On bond formation, the atomic orbitals overlap, and the greater the overlap, the stronger the bond.
Metal bonding is primarily covalent in nature, and metallic structure involves the resonance of electron-pair bonds between each atom and its neighbours.
In the year 1931, scientist Linus Pauling proposed the innovative concept of hybridization. He called the process hybridization and characterized it as the shifting of the energy of particular atoms’ orbitals to produce new orbitals of equivalent energy. New orbitals, known as hybrid orbitals, emerge as a result of this process. Some types of hybridization are-
- sp-hybridization: One s and one p-orbital are combined together to generate two sp–hybrid orbitals with a linear structure and a bond angle of 180 degrees. For example, when BeCl2 is formed, the first atom is in the excited state 2s1 2p1, which is then hybridized to generate two sp–hybrid orbitals. BeCl2 is formed when these hybrid orbitals collide with the two p-orbitals of two chlorine atoms.
- sp2–hybridization: One s and one p-orbital are combined together to generate three sp2– hybrid orbitals with a planar triangular shape and a bond angle of 120 degrees.
- sp3-hybridization: One s and three p-orbitals are merged in this hybridization to generate four sp3– hybrid orbitals with a tetrahedral structure and a bond angle of 109 degrees 28′, or 109.5 degrees.
What is the need for Valence Bond Theory?
Lewis’ theory explained the structure of molecules. It did not, however, explain the formation of chemical bonds. VSEPR theory, on the other hand, explained the shape of simple molecules. However, it had a very limited application. It also failed to explain the geometrical properties of complex molecules. As a result, scientists were forced to develop the theory of valence bonds in order to address and overcome these limitations.
Valence Bond Theory
The Lewis approach to chemical bonding failed to shed light on chemical bond formation. Furthermore, the valence shell electron pair repulsion theory (or VSEPR theory) had only a few applications (and also failed in predicting the geometry corresponding to complex molecules). The valence bond theory was proposed by German physicists Walter Heinrich Heitler and Fritz Wolfgang London to address these issues. The Schrodinger wave equation was also used to explain how a covalent bond between two hydrogen atoms formed.
This theory focuses on the concepts of electronic configuration, atomic orbitals (and their overlapping), and atomic orbital hybridization. Chemical bonds are formed by the overlapping of atomic orbitals, with electrons localised in the bond region. The valence bond theory also explains the electronic structure of molecules formed by the overlapping of atomic orbitals. It also emphasizes the fact that the nucleus of one atom in a molecule is drawn to the electrons of the other atoms.
Postulates of Valence Bond Theory
- When two valence orbitals (half-filled) from different atoms overlap on each other, covalent bonds form. As a result of this overlapping, the electron density in the area between the two bonding atoms increases, increasing the stability of the resulting molecule.
- An atom’s valence shell contains many unpaired electrons, allowing it to form multiple bonds with other atoms. According to the valence bond theory, the paired electrons in the valence shell do not participate in the formation of chemical bonds.
- Covalent chemical bonds are directional and parallel to the region corresponding to the overlapping atomic orbitals.
- Sigma bonds and pi bonds differ in the pattern in which the atomic orbitals overlap, i.e. pi bonds are formed by sidewise overlapping, whereas sigma bonds are formed by overlapping along the axis containing the nuclei of the two atoms.
Applications of Valence Bond Theory
- The valence bond theory’s maximum overlap condition can explain the formation of covalent bonds in several molecules.
- This is one of its most important uses. The difference in the length and strength of chemical bonds in H2 and F2 molecules, for example, can be explained by differences in their overlapping orbitals.
- The covalent bond in an HF molecule is formed by the overlap of the hydrogen atom’s 1s orbital and the fluorine atom’s 2p orbital, as explained by the valence bond theory. The covalent bond in an HF molecule is formed from the overlap of the 1s orbital of the hydrogen atom and a 2p orbital belonging to the fluorine atom, which is explained by the valence bond theory.
Limitations of Valence Bond Theory
- Failure to account for carbon’s tetravalency.
- There is no information provided about the energies of the electrons.
- The theory assumes that electrons are concentrated in specific locations.
- It does not provide a quantitative interpretation of coordination compounds’ thermodynamic or kinetic stabilities.
- There is no differentiation between weak and strong ligands.
- There is no explanation for the colour of coordination compounds.
Modern valence bond theory now complements molecular orbital theory, which rejects the valence bond idea that electron pairs are localised between two specific atoms in a molecule and instead believes that they are distributed in sets of molecular orbitals that can span the entire molecule. Molecular orbital theory predicts magnetic and ionisation properties clearly, whereas valence bond theory produces similar but more complicated results.
Aromaticity, on the other hand, is viewed as the delocalization of the π-electrons in molecular orbital theory. Because of the lack of orthogonality between valence bond orbitals and between valence bond structures, valence bond treatments are limited to relatively small molecules, whereas molecular orbitals are orthogonal. Valence bond theory, on the other hand, provides a much more accurate picture of the reorganization of electronic charge that occurs when bonds are broken and formed during the course of a chemical reaction.
The condition of maximum overlap, which leads to the formation of the strongest possible bonds, is an important aspect of the valence bond theory. This theory is used to explain the formation of covalent bonds in many molecules. In the case of the F2 molecule, for example, the F-F bond is formed by the overlap of the pz orbitals of the two F atoms, each of which contains an unpaired electron. Because the nature of the overlapping orbitals differs between H2 and F2 molecules, the bond strength and bond lengths differ. The covalent bond in an HF molecule is formed by the overlap of the 1s orbital of H and the 2pz orbital of F, both of which contain an unpaired electron. The mutual sharing of electrons between H and F results in the formation of a covalent bond in HF.
Question 1: What is the valence bond theory?
It is a theory that explains chemical bonding. According to VBT, the overlap of partially filled atomic orbitals results in the formation of a chemical bond between two atoms. The unpaired electrons are shared, resulting in the formation of a hybrid orbital.
Question 2: What are the merits of the valence bond theory?
The maximum overlap condition described by the VBT can be used to explain how covalent bonds form in many molecules. The theory can also shed light on the ionic nature of chemical bonds.
Question 3: How are sigma and pi bonds formed?
Sigma bonds are formed when the atomic orbitals involved in the bond overlap head-to-head. Pi bonds, on the other hand, involve the atomic orbitals overlapping in parallel.
Question 4: Based on the overlapping of orbitals, how many types of covalent bonds are formed and what are they?
Two types of covalent bonds are formed as a result of orbital overlapping. These are referred to as sigma and pi bonds.
- The end-to-end overlap of atomic orbitals along the inter-nuclear axis, known as a head-on or axial overlap, forms sigma bonds. End-on overlapping is classified into three types: s-s overlapping, s-p overlapping, and p-p overlapping.
- When atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis, a pi bond is formed.
Question 5: What is the orbital overlap concept?
This concept states that a covalent bond formed between atoms results in the overlap of orbitals belonging to atoms with opposite spins of electrons. The type of overlapping between the atomic orbitals determines the stability of the molecular orbital.
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