91Ó°ÊÓ

Atoms form bonds by sharing electrons in their atomic orbitals. The type of bond (single, double, or triple) depends on the number and type of orbitals involved in the bond formation. Single bonds involve the sharing of one pair of electrons (one from each atom) in a sigma (σ) bonding orbital, while double and triple bonds involve the sharing of additional electron pairs through pi (π) bonding orbitals.

In a single bond, atoms share a pair of electrons between their nuclei in a sigma (σ) bonding orbital. This σ bond is formed by the end-to-end overlap of the atomic orbitals along the internuclear axis, creating an area of electron density between the two nuclei. Since the electron density is distributed symmetrically along the internuclear axis, the atoms can rotate without breaking the bond.

In double and triple bonds, there are additional pi (π) bonding orbitals formed by the sideways overlap of p orbitals. The electron density in these π bonds is concentrated above and below the internuclear axis, forming electron-rich regions above and below the bond. In a double bond, one σ and one π bond are formed, while in a triple bond, one σ and two π bonds are formed.

The presence of π bonds in double and triple bonds restricts rotation about the internuclear axis. This is because rotating one atom about the internuclear axis would cause the p orbitals to lose their alignment, breaking the π bond’s overlap and disrupting electron density. Thus, breaking the π bond is necessary for rotation to occur in double and triple bonds. In conclusion, atoms in a single bond can freely rotate about the internuclear axis because the electron density in the σ bond is symmetrically distributed along the axis. In contrast, atoms in double and triple bonds cannot rotate without breaking the bond due to the presence of π bonds, which restrict rotation by requiring proper alignment of the p orbitals involved in the overlap.