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A sigma bond (σ bond) is formed by the head-on overlap of atomic orbitals, resulting in the sharing of a pair of electrons. This type of bond is the most common and strongest covalent bond found in molecules. It usually occurs between two hybrid orbitals or a hybrid orbital and a pure atomic orbital (e.g., an orbital from the s, p, d, or f subshells).

A pi bond (Ï€ bond) is formed by the side-to-side overlap of atomic orbitals, resulting in the sharing of two electrons. The interaction between the orbitals forms parallel electron-sharing regions above and below the plane of the atoms, creating a weaker bond compared to a sigma bond. A pi bond usually occurs between two unhybridized p orbitals.

1. Formation: Sigma bonds are formed by the head-on overlap of atomic orbitals, whereas pi bonds are formed by side-to-side overlap of atomic orbitals. 2. Strength: Sigma bonds are stronger than pi bonds due to their greater orbital overlap and electron density between the nuclei of the bonding atoms. 3. Rotation: Sigma bonds allow for free rotation around the bond axis, while pi bonds restrict rotation due to the parallel electron-sharing regions that would be disrupted if the atoms rotate about the bond. 4. Occurrence: Sigma bonds are present in all covalent bonds, whereas pi bonds are present only in double and triple covalent bonds. A double bond consists of one sigma bond and one pi bond, while a triple bond consists of one sigma bond and two pi bonds. 5. Electron Density: In a sigma bond, the electron density is concentrated along the bonding axis, directly between the two bonding atoms. Conversely, in a pi bond, the electron density is concentrated above and below the plane of the bonded atoms, creating two parallel regions. By understanding these basic differences, we can better visualize how sigma and pi bonds form and function in molecular structures, ultimately influencing their properties and reactivity.