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.
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.
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 :
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.
The potential energy curve showing variation of energy with internuclear distance between two hydrogen atoms
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.
There are two types of covalent bonds depending upon the nature of the overlap :
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 :
When there is an overlap of two half filled s-orbitals along the internuclear axis, the overlapping is called an s-s 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.
This type of overlapping occurs between two half filled p-orbitals.
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.
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 :
Some important points of difference are given in the table below :
|σ Bond||π Bond|
|Sigma bond is formed by end-to-end overlapping of orbitals (axial overlapping).||Pi bond is formed by sideway overlapping of orbitals (lateral overlapping).|
|Sigma bond can be formed by overlapping between s-s, s-p or p-p orbitals.||Pi bond involves the overlap of p-p orbitals only.|
|Sigma bond is a strong bond||Pi bond is a weaker bond.|
|Sigma bond consists of only one electron cloud, symmetrical about the internuclear axis.||Pi bond consists of two electron clouds, one above the plane of atomic nuclei and the other below it.|
|Free rotation of atoms around sigma bond is possible.||Rotation of atoms around pi bond is not possible without breaking the pi bond.|