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valance bond theory


Valence bond theory was proposed by Heitler and London in 1927 based on atomic orbitals and their configuration.

According to valence bond theory:

A covalent bond is formed when pure, valence atomic orbital of one atom overlaps with another pure, valence atomic orbital of another atom. During covalent bond formation energy releases that means stability increases.

each of the overlapping orbitals contains the unpaired electron of opposite spin.The electron pair is shared by both the atoms. The strongest bond is formed when the orbitals of the two atoms overlap to the maximum extent. based on the overlapping of orbitals, two types of covalent bonds are formed. These are known as sigma and pi bonds.

Sigma bonds are formed by the end-to-end overlap of atomic orbitals along the inter-nuclear axis known as a head-on or axial overlap.

Types of orbital overlap :-

Depending upon the type of overlapping, the covalent bonds are of two types, known as sigma (σ ) and pi (π) bonds.

(i) Sigma (σ bond):- Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis

The axial overlap involving these orbitals is of three types :-

• s-s overlapping :- In this case, there is overlap of two half-filled s-orbitals along the internuclear axis.


•  s-p overlapping :- This type of overlapping occurs between half-filled s-orbitals of one atom and half filled p-orbitals of another atoms.


. P-P overlapping :- this type of overlapping occurs between half filled p- orbitals of  the two approching atoms .


(2). Pi (π bonds):- π bond is formed by the atomic orbitals when they overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis.The orbital formed is due to lateral overlapping or side wise overlapping.


• strength  of sigma and pf bonds :-

Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be stronger bond in comparison to a π-bond.    Distinction between sigma and n bond.

 Valance bond theory postulates:-

The valence bond theory’s main postulates are mentioned below :-

 (1). When two valence spins half-       filled from two separate atoms overlap on one other, covalent bonds are created. As a result of this overlapping, the electron density in the intersection of two bonding atoms increases, boosting the stability of the resulting molecule

(2).  An atom’s valence shell has several unpaired electrons, allowing it to make many bonds with other atoms. According to the valence bond theory, the paired electrons in the valence shell need not participate in the creation of chemical bonds.

(3).  Chemical bonds that are covalent are directed and parallel to the region corresponding to the overlapping atomic orbitals.

(4).  Sigma bonds and pi bonds are distinguished by the pattern in which the atomic orbitals overlap, i.e., pi bonds are produced by overlapping along the axis having the nuclei of the two atoms, whereas sigma bonds are formed by overlapping along the axis comprising the nuclei of the two atoms

The head-to-head collision of the atomic orbitals involved in the bond forms sigma bonds. Pi bonds, but at the other end, involve the overlapping of atomic orbitals in a parallel manner. 

• strength  of sigma and pf bonds :-

Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be stronger bond in comparison to a π-bond.

•A covalent bond  is directional. Such a bond is also parallel to the region of overlapping atomic orbital.

• In case the atomic orbitals possess more than one unpaired electron, more than one bond can be formed and electrons paired in the valence shell cannot take part in such a bond formation. 

The valance bond theories limitations :- 

 • While it provides a qualitative depiction of the complex, it does not provide a quantitative interpretation of the complex’s stability

 • It does not account for the complexes spectra (colour)

• It predicts no distortion in symmetrical compounds, but predicts distortion in all copper (II) and titanium (III) complexes

• It provides no specific information regarding the complexes’ magnetic characteristics. It cannot, for example, account for the complexes’ temperature-dependent paramagnetism

• 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

• It fails to account satisfactorily for the occurrence of inner and outer orbital complexes

• 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

• Electron spin resonance demonstrate that the electron is not in the 4p level in Cu(II) complexes, indicating that the complex is planar

• 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.

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