1. Bonding in the Benzene
Dr. K. Shahzad Baig
Memorial University of Newfoundland (MUN)
Canada
Petrucci, et al. 2011. General Chemistry: Principles and Modern Applications. Pearson Canada Inc., Toronto, Ontario.
Tro, N.J. 2010. Principles of Chemistry. : a molecular approach. Pearson Education, Inc.

2. Delocalized Electrons: Bonding in the Benzene Molecule
Benzene is built from:
hydrogen atoms (1s1) and
carbon atoms (1s22s22px
12py
1).
Each carbon atom is attached with 1 hydrogen and 2 carbons.
Each carbon atom doesn't have enough unpaired electrons to form the required number
of bonds, therefore, it needs to promote one of the 2s2 pair into the empty 2pz orbital.
an electron is
promoted from the
2s to the empty 2p to
give 4 unpaired
electrons
C6H6

3. In benzene each carbon is only joining to 3 other atoms,
when the carbon atoms hybridise their outer orbitals before forming bonds, they only
need to hybridise three of the orbitals rather than all four.
[They use the 2s electron and two of the 2p electrons, but leave the other 2p electron
unchanged]
The new orbitals formed are sp2 hybrids. The three sp2 hybrid orbitals arrange
themselves as far apart as possible - which is at 120° to each other in a plane.
End-to-end overlap of the orbitals produces σ-bonds.
The π- bonds are not localized between specific carbon atoms but are spread out
around the six-membered ring. Hence, called delocalized
The six remaining orbitals construct the π-bonds

4. The π- bonds are not localized between specific carbon atoms but are spread out
around the six-membered ring. Hence, called delocalized
The six remaining orbitals construct the π-bonds

5. The three bonding molecular orbitals,
distribute electron charge in the molecule in
the form of two doughnut-shaped regions:
one above and one below the plane of the C
and H atoms.
Because they are spread out among all six C
atoms instead of being concentrated between
pairs of C atoms, these molecular orbitals are
called delocalized molecular orbitals

6. Structures with Delocalized MOs
1. The VSEPR theory, gives a trigonal-planar electron-group geometry with a bond angle 117o for O3. The
sp2 hybridization is chosen for the central O atom and terminal atoms, Thus, each O atom uses the
orbital set sp2+ p
2. 2. Of the 18 valence electrons in O3, assign 14 to the hybrid orbitals of the σ-bond framework. 4 of
these are bonding electrons and 10 are lone-pair electrons.
3. The three unhybridized 2p orbitals combine to form 3 molecular orbitals of the π type. One of these
orbitals is a bonding molecular orbital, and the second is antibonding. The third is a nonbonding molecular
orbital.
A nonbonding molecular orbital has the same energy as the atomic orbitals from which it is formed, and it
neither adds to nor detracts from bond formation
4. The remaining 4 valence electrons are assigned to the π molecular orbitals is 2 go into the bonding
orbital and 2 into the nonbonding orbital. The antibonding orbital remains empty
5. The bond order associated with the π molecular orbitals is (2-0) /2 = 1. This π bond is distributed
between the two O-O bonds and amounts to ½ of a π bond for each.

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