Overlapping of
Atomic Orbitals

Atoms combine together to lower down the energy of the system to attain stability — in layman's terms, the rule simply says that the less energy you need, the easier for you to survive. For example, hydrogen prefers to exist as H₂ molecule instead of isolated hydrogen atom because H₂ requires lesser energy than isolated H atom.

Orbital Overlapping

When two atoms combine together to form a covalent bond, their energy is minimum when they are so close to each other that their orbitals are partially merged. This partial merging of atomic orbitals is known as orbital overlapping or overlapping of atomic orbitals.

Orbital Overlap Concept

Molecules like H₂ and F₂ contain a single bond. Even so, the bond dissociation enthalpies of H₂ (435.8 KJ/mol) and F₂ (155 KJ/mol) are different. The concept of orbital overlap helps us explain the reason. Let us take the example of H₂ molecule:

Overlap concept in H₂ molecule : Let us assume that two hydrogen atoms A and B having nuclei nA and nB; and electrons eA and eB are coming close to each other. When the two atoms are far apart, their is no interaction between them. As these two atoms come close to each other, new attractive and repulsive forces begin to operate. These forces are written below :

1. Force of attraction between nucleus of A and electron of B (attraction between nA and eB)
2. Force of attraction between nucleus of B and electron of A (attraction between nB and eA)
3. Force of repulsion between nuclei A and B (repulsion between nA and nB)
4. Force of repulsion between electrons A and B (repulsion between eA and eB)

Attractive forces tend to bring two atoms close to each other whereas repulsive forces tend to move them away. In hydrogen, the magnitude of the new attractive forces is greater than that of new repulsive forces. As a result, two atoms come close to each other and potential energy decreases. The atoms approach each other until the equilibrium stage is reached where the net force of attraction balances the force of repulsion and system acquires minimum energy. At this stage, the bond length — distance between two nuclei — of hydrogen is 74 pm.

Since energy is released during the formation of H₂ molecule, H₂ molecule is more stable than isolated hydrogen atoms.

Explanation for unequal bond dissociation enthalpies in molecules like H₂ and F₂: In general, stronger bonds are the result of greater overlap. The overlap in H₂ is greater than that of F₂. As a result, more energy is required to dissociate an H₂ molecule into isolated hydrogen atoms.

Types of Covalent Bond Based on nature of overlap

There are two types of covalent bonds depending upon the nature of the overlap :

1. Sigma (σ) bond
2. Pi (π) bond

Sigma (σ) Bond

Sigma bond is formed by the end to end overlap of bonding orbitals along the internuclear axis. This type of overlap is called a head on overlap or axial overlap.

The head on overlap can take place in any one of the following ways :

s-s overlapping

When there is an overlap of two half filled s-orbitals along the internuclear axis, the overlapping is called an s-s overlapping.

s-p overlapping

When there is an overlap between a half filled s-orbital and a half filled p-orbital, the overlapping is called an s-p overlapping.

p-p overlapping

This type of overlapping occurs between two half filled p-orbitals.

Pi (π) Bond

A pi (π) bond is formed when the atoms overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. This type of overlap is called a sideway overlap.

Strength of Sigma & Pi Bonds

In σ bond, the overlapping occurs to a larger extent. Hence, σ bond is stronger than π bond, where the overlapping occurs to a smaller extent. Further, a π bond between two atoms is formed in addition to a sigma bond. The presence of σ and π bonds in single, double and triple bonds is as follows :

• A single bond contains one σ bond.
• A double bond contains one σ and one π bonds.
• A triple bond contains one σ and two π bonds.

Difference between Sigma and Pi Bonds

Some important points of difference are given in the table below :