In chemistry, valence bond (VB) theory is one of two basic theories—along with molecular orbital (MO) theory—that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the atomic orbitals of dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule.
According to this theory, a covalent bond is formed by the overlap of the half-filled valence orbitals in atoms containing an unpaired electron. A valence bond structure is similar to a Lewis structure, but where a single Lewis structure cannot be written, several valence bond structures are used. Each of VB structure represents a specific Lewis structure. The combination of valence bond structures is the main point of resonance theory.
Mechanisms of Bonding in VB Theory
Valence bond theory dictates that overlapping atomic orbitals of participating atoms form a chemical bond. Because of the overlap, it is probable that electrons are found in the bond region. VB theory views bonds as weakly coupled orbitals (small overlap). VB theory is typically easier to employ in ground state molecules.
SIGMA AND PI BONDS
There are two types of overlapping orbitals: sigma and pi. Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head. Pi bonds occur when the two overlapping orbitals are parallel. For example, a bond between two s-orbital electrons is a sigma bond, because two spheres are always coaxial. In terms of bond order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond, and triple bonds contain one sigma bond and two pi bonds. However, the atomic orbitals for bonding may be hybrids. Often, bonding atomic orbitals have a character of several possible types of orbitals. The methods to get an atomic orbital with the proper character for bonding is called hybridization.
VB and MO Compared
Valence bond theory complements molecular orbital (MO) theory, which does not adhere to the VB idea that electron pairs are localized between two specific atoms in a molecule. MO theory suggests that electrons are distributed in sets of molecular orbitals that can extend over the entire molecule. MO theory can predict magnetic and ionization properties in a straightforward manner. Valence bond theory gives similar results but is more complicated.
Valence bond theory views aromatic properties of molecules as the result of resonance between Kekulé, Dewar, and possibly ionic structures. Molecular orbital theory views it as delocalization of the π-electrons. The underlying mathematics of VB theory is also more complicated, limiting VB treatment to relatively small molecules. VB theory provides a much more accurate picture of the reorganization of electronic charge that takes place when bonds are broken and formed during the course of a chemical reaction. Specifically, valence bond theory correctly predicts the dissociation of homonuclear diatomic molecules into separate atoms. Molecular orbital theory predicts dissociation into a mixture of atoms and ions.
MODERN VALENCE BOND THEORY
More recently, several groups have developed what is often called modern valence bond theory. This replaces the overlapping atomic orbitals with overlapping valence bond orbitals expanded over a large number of basis functions. These functions are either centered on one atom to give a classical valence bond picture or centered on all atoms in a molecule. The resulting energies are more competitive, with energies from calculations where electron correlation is introduced based on a Hartree–Fock reference wave function. The most recent text discussing modern valence bond theory is written by Shaik and Hiberty.
An important aspect of the VB theory is the condition of maximum overlap which leads to the formation of the strongest possible bonds. This theory is used to explain the covalent bond formation in many molecules. In the F2 molecule the F–F bond is formed by the overlap of pz orbitals of the two F atoms, each containing an unpaired electron. Since the nature of the overlapping orbitals are different in H2 and F2 molecules, the bond strength and bond lengths differ between H2 (Figure 1) and F2 molecules. In an HF molecule, the covalent bond is formed by the overlap of the 1s orbital of H and the 2pz orbital of F, each containing an unpaired electron. Mutual sharing of electrons between H and F results in a covalent bond in HF.