This assignment attained a distinction, with full explanation of structure of benzene, evidence for structure, chemical properties, addition and substitution mechanisms for benzene and their comparison, and analysis of different monosubstituents on benzene ring.
Benzene
Benzene belongs to a family of aromatics. It is a six-carbon aromatic ring made up of
hydrogen and carbon atoms. Today, it is used as a carcinogen. Benzene has an electron
configuration of 1s2, 2s2, 2p2 (with 6 electrons in total), and a molecular formula of C6H6.
There were three problems of Benzene that was suggested by Kekule. Figure 1 - Benzene
In 1865, Kekule proposed benzene with three double bonds to explain the molecular
formula. It was a cyclic ring. The carbons are arranged in a hexagon, with
alternating double and single bonds between them. Each carbon atom
had a hydrogen atom attached to it. However, this structure did not
justify the chemical behaviour of benzene.
Problem 1: Lack of reactivity with Benzene
All normal alkenes show the characteristic behaviours that’s expected from the double
bonds. Alkenes with their double bonds react very quickly with bromine (brown) to produce
a colourless, bromoalkane solution, in the dark and at room temperature. However, benzene
did not undergo this reaction. Instead, it only reacted with bromine in a very bright light or
using a catalyst. It was concluded that benzene was an unusual alkene that doesn’t undergo
addition reactions.
Problem 2: Thermodynamic stability of Benzene
Every thermochemistry calculation using Kekules structure was wrong by roughly 150kJmol-
1. This is shown with hydrogenation. When hydrogen is added to cyclohexene (one double
bond) It forms cyclohexane. The double bond is replaced by a single bond and the ‘CH’
groups become ’CH2’ The enthalpy of hydrogenation of to cyclohexene was found as -
119kJmol-1. Figure 2 - Hydrogenation of Cyclohexene
Since cyclohexene contained one double bond. Benzene contains three double bonds. The
expected/theoretical enthalpy of hydrogenation of Benzene should be -360kJmol-1 (three
times the energy of cyclohexene). However, the Figure 3 - Thermochemical analysis
experimentally determined value for the
hydrogenation of benzene was found to be -
208kJmol-1. (Wrong of approximately
150kJmol-1). This suggests that real benzene is -
152kJmol-1 more stable (delocalisation energy)
than Kekules structure.
Figure 4 - Hydrogenation of Benzene
1
, Osman Yousuf – Unit 14 – Assignment 2
Problem 3: Bond lengths of Benzene Figure 5 - X-ray crystallography of Benzene
The bond lengths in benzene are all the same
length, experimental data shows all carbons in
benzene have a bond length of 0.139nm. This
creates an equal and symmetrical hexagon, with
the same bond angles. However, single carbon-
carbon bonds have a length of 0.154nm and
double carbon-carbon bonds have a length of
0.134nm (shorter). If there were alternating single
and double carbon bonds, the shape of benzene
would not be equal, it would cause distortion and
become an irregular shape, with different bond
angles. Kekules structure suggests that there are three long bonds and three short bonds.
X-ray crystallography demonstrates that all bond lengths and bond angles are the same,
creating a perfect hexagon. The positions of carbon atoms are shown as equally separated
spikes.
Accepted Structure of Benzene
Figure 6 - Promotion of an electron
Benzene is a hydrocarbon, made up of only hydrogen and carbon atoms. Carbons have 6
electrons (1s2, 2s2, 2p2) There are
insufficient number of unpaired
electrons in the 2p orbital to form
bonds. Therefore, one electron from
the 2s orbital is promoted to the 2pz
orbital. This gives four unpaired electrons in total, with a configuration of 1s 2, 2s1, 2p3. This is
a favourable process as it provides less repulsion and more stability.
Three orbitals (one s and two p orbitals) hybridise (combine) Figure 7 - Single Carbon, Sp2 Hybridised
to form three new orbitals. All three orbitals are equivalent.
Resulting in three sp2 orbitals and a 2p orbital that remains
unhybridised with a bond angle of 90° to planar orbitals. The
three sp2 orbitals repel into a planar arrangement. They move
as far away as possible; creating a bond angle of 120° to each
other in a plane. This results in a single carbon that is Sp2
hybridised.
A Sp2 orbital forms covalent bonds with other carbon sp2
orbitals due to an overlap. This forms a single carbon- Figure 8 - Cyclical 3D structure of Benzene
carbon bond. The bond is now called a Sigma bond. Each
carbon has three sigma bonds with an electron left over in
the unhybridised 2p orbital.
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