Valance Bond theory describes chemical bonding in determining the shapes of molecules. It states that when half-filled orbitals of two atoms overlap, the electrons paired up as a result of this overlap. Thus, a covalent bond is formed. The bond strength depends on the overlapping of orbitals. Larger the overlap, the more powerful the bond is. VBT discusses the two types of overlapping orbitals. Sigma and pi. Sigma bond forms when orbitals overlap in the internuclear axis. Pi bond formation takes place when there is overlap of two unhybridized p – orbitals. Single bond contains one sigma bond, double bonds have one sigma and one pi while triple covalent bond consists of one sigma and two pi bonds.
Valence Bond Theory (VBT)
VSERP theory predicts and discusses the shapes of molecules however does not provide factors for the formation of bonds. VBT is concerned with both the formation of bonds and the shapes of molecules. This approach of describing a covalent bond considers the molecule as a combination of atoms.
According to the quantum mechanical technique, a covalent bond is formed when half-filled orbitals in the outer or valence shells of 2 atoms overlap, so that a pair of electrons, one electron from each atom, inhabits the overlapped orbital. As a result of this overlap, the electrons with opposite whirls ended up being paired to support themselves.
The larger the overlap, the more powerful is the bond. The essential condition for chemical bonding is that the orbitals of atoms taking part in bond formation should overlap and the direction of the bond is figured out by the directions of the two overlapping orbitals.
History of Valence Bond Theory
Valence bond theory draws from Lewis structures. G.N. Lewis proposed these structures in 1916, based upon the idea that two shared bonding electrons formed chemical bonds. Quantum mechanics was applied to explain bonding properties in the Heitler-London theory of 1927.
This theory explained chemical bond formation in between hydrogen atoms in the H2 molecule using Schrödinger’s wave equation to merge the wavefunctions of the two hydrogen atoms. In 1928, Linus Pauling integrated Lewis’s pair-bonding idea with the Heitler-London theory to propose valence bond theory.
Valence bond theory was developed to describe resonance and orbital hybridization. In 1931, Pauling published a paper on valence bond theory entitled, “On the Nature of the Chemical Bond.” The first computer system programs utilized to describe chemical bonding utilized molecular orbital theory, but since the 1980s, principles of valence bond theory have actually become programmable.
Today, the modern versions of these theories are competitive with each other in terms of accurately explaining genuine behavior.
Sigma (σ) and Pi (π) Bonds
There are 2 types of overlapping orbitals: sigma (σ) and pi (π). Both bonds are formed from the overlap of 2 orbitals, one on each atom. σ bonds happen when orbitals overlap in between the nuclei of two atoms, likewise known as the internuclear axis. π bonds take place when 2 (unhybridized) p-orbitals overlap.
The p-orbitals, in one π bond, lie above and below the nuclei of the atoms. By inhabiting the area of space that is above, listed below, and on the sides of an atom’s nuclei, two π bonds can form.
Both types of overlapping orbitals can be related to bond order. Single bonds have one sigma bond. Double bonds consist of one σ and one π bond, while triple bonds consist of one σ and 2 π bonds.
Examples of Sigma (σ) and Pi bonds (π)
An example of Sigma bonds would be the simple C-H bond or C-X bond, while examples of pi bonds would be C =O and C ≡ N, where the first bond is a sigma bond, and the second/third bond are pi bonds.
Postulates of Valence Bond Theory
The essential postulates of the valence bond theory are listed below.
- Covalent bonds are formed when 2 valence orbitals (half-filled) coming from two different atoms overlap each other. The electron density in the location between the two bonding atoms increases as a result of this overlapping, consequently increasing the stability of the resulting molecule.
- The presence of lots of unpaired electrons in the valence shell of an atom enables it to form multiple bonds with other atoms. The paired electrons present in the valence shell do not take part in the formation of chemical bonds according to the valence bond theory.
- Covalent chemical bonds are directional and are also parallel to the area corresponding to the atomic orbitals that are overlapping.
- Sigma bonds and pi bonds differ in the pattern that the atomic orbitals overlap in, i.e. pi bonds are formed from sidewise overlapping whereas the overlapping along the axis containing the nuclei of the two atoms causes the development of sigma bonds.
Uses of VBT
Valence bond theory can often explain how covalent bonds form. The diatomic fluorine molecule, F2, is an example. Fluorine atoms form single covalent bonds with each other. The F-F bond arises from overlapping pz orbitals, which each consist of a single unpaired electron.
A similar situation takes place in hydrogen, H2, but the bond lengths and strength are various between H2 and F2 particles. A covalent bond forms in between hydrogen and fluorine in hydrofluoric acid, HF.
This bond forms from the overlap of the hydrogen orbital and the fluorine 2 p z orbital, which each have an unpaired electron. In HF, both the hydrogen and fluorine atoms share these electrons in a covalent bond.
Limitations of VBT
The imperfections of the valence bond theory consist of:
- Failure to describe the tetravalency exhibited by carbon.
- No insight provided on the energies of the electrons.
- The theory presumes that electrons are localized in specific locations.
- It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination substances.
Number of Orbitals and Types of Hybridization
According to VBT theory, the metal atom or ion under the influence of ligands can utilize its (n-1) d, ns, np, or ns, np, and orbitals for hybridization to yield a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, square planar and so on. These hybrid orbitals are allowed to overlap with ligand orbitals that can contribute electron sets for bonding.