In chemistry, sigma (σ) and pi (π) bonds are types of covalent bonds formed between atoms. The distinction between sigma and pi bonds arises from the types of orbitals involved in their formation.
Sigma bonds are formed by the overlap of atomic orbitals along the internuclear axis (the line connecting the nuclei of the bonded atoms). The orbitals involved in sigma bonding are typically hybridized orbitals, which are created by mixing different types of atomic orbitals. Hybridization allows the orbitals to have the appropriate spatial orientation and shape to maximize the overlap and form strong sigma bonds.
For example, in the case of sp³ hybridization, one s orbital and three p orbitals combine to form four sp³ hybrid orbitals. These hybrid orbitals have a characteristic shape and orientation, which allows them to overlap effectively with other hybrid orbitals or atomic orbitals to form sigma bonds. This type of hybridization is commonly observed in molecules with tetrahedral geometry, such as methane (CH₄).
On the other hand, pi bonds are formed by the sideways overlap of unhybridized p orbitals, which are perpendicular to the internuclear axis. These p orbitals are not involved in hybridization because they retain their original shape and orientation. As a result, pi bonds do not require hybridized orbitals for their formation.
In a double bond, such as in ethene (C₂H₄), one sigma bond is formed by the overlap of sp² hybrid orbitals, while the pi bond is formed by the sideways overlap of two unhybridized p orbitals. The pi bond is weaker than the sigma bond due to the lower extent of overlap and is more susceptible to disruption.
In summary, sigma bonds involve hybridized orbitals because hybridization provides the necessary orbital geometry and orientation for effective overlap along the internuclear axis. Pi bonds, on the other hand, do not require hybridization since they involve the overlap of unhybridized p orbitals that are perpendicular to the internuclear axis.