The hybridization of all carbon atoms in Benzene molecule is

sp²

Understanding the Structure of Benzene

1. Bonding in Benzene:
   • The benzene molecule consists of six carbon atoms and six hydrogen atoms.
   • Each carbon atom forms three sigma (σ) bonds – two with neighboring carbon atoms and one with a hydrogen atom.
   • There is one additional bond that each carbon forms, but this is not a sigma bond. Instead, it is a pi (π) bond, which is delocalized over the entire ring.
2. The Nature of Carbon–Carbon Bonds in Benzene:
   • In benzene, the carbon atoms are sp² hybridized. This means that each carbon atom in the molecule undergoes sp² hybridization to form three bonding orbitals: one for each sigma bond with another carbon and a hydrogen atom.
   • The remaining unhybridized p orbital on each carbon overlaps with the p orbitals of adjacent carbon atoms to form a pi (π) bond. These π bonds are delocalized, meaning they are spread over all six carbon atoms in the benzene ring, rather than being confined between two carbon atoms.

The sp² Hybridization

1. Orbital Overlap:
   • Each carbon atom in benzene uses three of its four available electrons to form bonds: two electrons form sigma bonds with adjacent carbon atoms, and one electron forms a sigma bond with a hydrogen atom.
   • The remaining electron occupies an unhybridized p orbital, which participates in the formation of the π bond.
2. Geometry of the Carbon Atom:
   • The three sp² hybrid orbitals arrange themselves in a trigonal planar geometry around each carbon atom. This results in a bond angle of approximately 120° between the sigma bonds.
   • The delocalized π electrons are free to move over the ring, providing benzene with its characteristic stability.

Each carbon atom in the benzene molecule undergoes sp² hybridization to form three sigma bonds and one delocalized pi bond. The unique structure and delocalization of electrons in benzene give it its aromatic stability.