Aromatic chemistry involves the study of
molecules containing benzene rings, many of these compounds are aromatic hydrocarbons, these compounds are often referred to
as arenes.
When you think of the word aromatic you probably imagine sweet smells and fragrances and indeed many
have pleasant smells and odours but from my own personal experience many aromatic compounds
have very
unpleasant smells and odours and some have no smell at all.
Benzene is perhaps the most well known
molecules in organic
chemistry. It is a pale yellow
liquid at room temperature and was first isolated by the fractional distillation of
whale oil in 1825 by Michael Faraday. At the time whale oil was in common use in oil lamps
used to light
homes. Perhaps the reason that benzene is such a well known
molecule is due to the problems and difficulties that
arose in trying to figure out the structure of this unusua molecule.
Friedrich Kekulé was a prominent German organic chemist who suggested in 1865 that benzene consisted of a ring of six carbon atoms- a cyclohexatriene (cyclo= ring, hexa= 6 carbon atoms, tri=3 double C=C bonds, -ene for alkenes). Kekulé claimed that his idea for the structure of benzene came to him in a dream while riding an open topped horse drawn bus in London!! The two skeletal drawings below are often used to represent the structure of benzene as proposed by Kekulé.
Kekulé's proposed structure of benzene as a cyclic unsaturated triene molecule is not without its problems- namely the fact that the molecule is highly unsaturated yet it is very reluctant to undergo characteristic electrophilic addition reactions that would be expected from an unsaturated molecule.
One of the hydrogen atoms in a benzene ring can be easily replaced by a halogen atom such as chlorine or bromine. This is an electrophilic substitution reaction. If only one hydrogen atom on the benzene ring is replaced by a halogen atom then only one isomer is produced. The example below shows bromobenzene, here one of the hydrogen atoms in the benzene ring has been replaced or substituted by a bromine atom. All the carbon atoms are equivalent so it does not matter which hydrogen is replaced, you can only obtain one isomeric product since all the carbon atoms are in equivalent environments.
If more bromine is added then obviously you can replace more of the hydrogen atoms in a benzene ring with bromine atoms. While only one bromobenzene isomer is produced, in theory four isomers of dibromobenzene should be able to be produced if Kekulé proposed structure for benzene is correct. The structures of these four isomers are shown below:
The problem Kekulé had was simple, his theory predicted that there should be 4 isomers of a disubstituted benzene ring such as the one shown above. However in practice only 3 isomers could be synthesised and isolated. The problem was that only one isomer of 1,2-dibromobenzene could be isolated. You can probably guess what Kekulé's answer to this was! Simple the two possible 1,2-dibromobenzene isomers cannot be isolated because they interconvert rapidly between each isomer. The position of the carbon carbon double bonds is rapidly changing and oscillating between the two possible positions as shown below so it is not possible to isolate any one particular isomer:
You may feel and probably quite correctly that Kekulé's ideas to account for the structure of benzene and his predictions for the number of isomers for the disubstituted benzene rings seem a bit of a fudge. However in his defence they offer an explanation at least for the correct number of isomers for mono and disubstituted benzene rings. The one key point that Kekulé's model fails to address is the fact that benzene's empirical formula clearly indicates that it is unsaturated and yet benzene fails to undergo the expected electrophilic addition reactions you would expect of unsaturated compounds such as alkenes. Indeed in its reactions benzene undergoes electrophilic substitution reactions and not electrophilic addition reactions.
In 1899 the theory of resonance was put forward to try and understand why benzene does not undergo electrophilic addition reactions that we might expect from an unsaturated molecule, this is covered in more detail below. For now consider another piece of unusual behaviour that was observed, that is the enthalpies of hydrogenation of benzene. The hydrogenation of cyclohexene is shown below:
Cyclohexene (C6H10) is an unsaturated molecule containing 1 carbon carbon double bond (C=C). The hydrogenation or addition of hydrogen across the C=C bond is a simple and straightforward reaction. The enthalpy change for this exothermic reaction is slightly less than -120kJ mol- (ΔH =-120kJmol-1).
Now if we imagine the hydrogenation of Kekulé's cyclohexa-1,3,5-triene, as shown below, it would seem reasonable to expect the enthalpy change for this hydrogenation reaction to simply be three times that of the enthalpy of hydrogenation of cyclohexene since the reaction simply involves the hydrogenation of three carbon carbon double (C=C) bonds as opposed to a one C=C bond in cyclohexene. However when the hydrogenation reaction is carried out the enthalpy change (ΔH) is only -208kJmol-1 and not the expected -360kJmol-1, that is -152kJmol-1 less energy is given out than expected. This suggests that benzene is about 150 kJmol-1 more stable than Kekulé's cyclohexa-1,3,5-triene.
If we imagine a simple energy profile diagram similar to the one shown below we can get a better idea of what exactly is happening. If benzene contained alternate carbon carbon double bonds (C=C) and carbon carbon single bonds (C-C) as proposed by Kekulé then when the C=C are hydrogenated in cyclohexene H-H and C=C are broken and C-H bonds are being formed in the product, cyclohexane. The same bonds are being formed and broken in Kekulé's benzene structure and so we would expect an enthalpy of hydrogenation of -360kJ mol-1. However the fact that only 208kJmol-1 are released tells us that the bonds being broken in benzene are NO simple C=C. The energy released by forming the C-H bonds in the product and the H-H bonds in the reactants does not change, the only bonds which must somehow be different are those in the benzene molecule. The fact that there is less energy released means that the benzene molecule contains less energy than expected, that is it is more stable than expected. This is shown in the energy profile diagram where you can see that in the second diagram there is less energy stored in the benzene molecule so it sits lower down the energy y-axis. This means that the enthalpy change for the hydrogenation reaction will be considerable less than expected.
The fact that the bonds in benzene are not C-C or C=C comes from evidence from measuring the bond lengths of the carbon carbon bonds in benzene, the carbon carbon bond lengths in benzene are found to have intermediate length (135pm) between carbon carbon double bonds (135pm) and carbon single bonds (147pm). You may recall that pi(π) bonds can form by the sideways overlap of p-orbitals on carbon atoms in molecules which contain double bonds (C=C), as shown in the image below:
In a benzene molecule each of the carbon atoms has one electron in a 2p-orbital. Within the ring all C-C and C-H bonds are formed by the sigma overlap of sp2 hybrid orbitals (hybrid orbitals are not covered/required in the A-level specification), but essentially this results in the formation of a flat ring of carbon and hydrogen atoms. The 2p-orbitals on each of the carbon atoms have lobes that stick out above and below this flat ring of carbon and hydrogen. Now in a normal carbon double bond (C=C) you would expect the p-orbitals on two carbon atoms to overlap and form a pi bond. However in benzene something unusual happens, the lobes on the p-orbitals merge with all the adjacent lobes on other p-orbitals of adjacent carbon atoms. This results in the formation of new molecular orbitals above and below the plane of the flat carbon atoms, essentially we have areas of high electron density above and below the flat ring of carbon atoms. This is outlined below:
To help explain many of the properties of benzene and to help visualise what is happening to
the electrons within a benzene
ring when it undergoes reactions we will use the ideas put forward in
resonance theory. Resonance is a term you will
come across many times when dealing with organic compounds as it is often used to explain the
reactivity or stability
of certain molecules and ions you will meet in the course
of your studies. Resonance is the simply the idea that
in a molecule the nuclei
stay in fixed positions but that the electrons within the molecule/ion are free to move. This
is the case with benzene where we have 6 delocalised
electrons within the electron clouds or orbitals above and below
the flat ring of 6 carbon atoms.
The diagram below shows two resonance hybrid structures for
benzene. To draw these two
resonance structures we can imagine that
the carbon and hydrogen nuclei stay fixed but that the
6 delocalised electrons simply shift or move. In each of the two
resonance hybrid structures which I have drawn there
are carbon carbon double bonds (C=C) and carbon carbon single
bonds (C-C), now from our discussion above we know that in reality these bonds do NOT exist since
all the carbon
carbon bonds in benzene have the same length, intermediate
between double and single carbon bond lengths, but how would or could
you draw a molecule to show this??
The double headed arrow between the two resonance hydrids
does NOT MEAN that the two molecules
oscillate backwards
and forwards between the two different forms, instead it means that the actual structure of
benzene is a mix or hybrid
of the two drawn molecules. This is very different
from Kekulé's model of the benzene
structure where there are two distinct
forms that oscillate
backwards and forwards between the two different molecules. In Kekulé's
model there are double and single bonds present
within a molecule of benzene.
Resonance also implies stability within a
molecule. The more resonance forms that
can be drawn for a particular molecule the more stable that
molecule is likely to be.
You will often see the structure of
benzene drawn as a ring with a circle in the middle where
the circle represents the six
delocalised pi electrons.
It is a matter of personal choice as to the way you decide to draw the structure of
benzene and other aromatic
compounds which contain benzene rings. Personally I prefer
to use the Kekulé model showing the presence
of carbon carbon double and single bonds, simply because I find it easier to keep track of the
electrons and to follow
where they are going when writing out mechanisms but other people prefer to use the ring with the
circle in the middle. Ultimately the choice is up to you as both models are acceptable despite
their obvious limitations
in describing exactly what is happening, but this limitation is simply due to the method used to
construct
organic mechanisms.