At the turn of the nineteenth century, one of the signs of living the good life was having gas lines connected to your house, so that you could use gas lanterns to light the house af-ter dark. The gas burned in the lanaf-terns was called coal gas because it was produced by heating coal in the absence of air. The principal component of coal gas was methane, CH4. In 1825, Michael Faraday was asked to analyze an oily liquid with a distinct odor that collected in tanks used to store coal gas at high pressures. Faraday found that the com-pound had the empirical formula CH. Ten years later, Eilhardt Mitscherlich produced the same material by heating benzoic acid with lime. Mitscherlich named the substance ben-zin, which became benzene when translated into English. He also determined that the mol-ecular formula of the compound was C6H6.
Benzene must be an unsaturated hydrocarbon because it has far less hydrogen than the equivalent saturated hydrocarbon: C6H14. But benzene is too stable to be an alkene or alkyne. Alkenes and alkynes rapidly add Br2to their CPC and CqC bonds, whereas benzene reacts with bromine only in the presence of a catalyst, FeBr3. Furthermore, when benzene reacts with Br2 in the presence of FeBr3, the product of the reaction is a com-pound in which a bromine atom has been substituted for a hydrogen atom, not added to the compound the way an alkene adds bromine.
FeBr3
C6H6 Br2 8888n C6H5Br HBr
Other compounds were eventually isolated from coal that had similar properties. Their for-mulas suggested the presence of multiple CPC bonds, but the compounds were not reac-tive enough to be alkenes. Because they often had a distinct odor, or aroma, they became known as aromatic compounds.
The structure of benzene was a recurring problem throughout most of the nineteenth century. The first step toward solving the problem was taken by Friedrich August Kekulé in 1865. (Kekulé’s interest in the structure of organic compounds may have resulted from the fact that he first enrolled at the University of Giessen as a student of architecture.) One day, while dozing before a fire, Kekulé dreamed of long rows of atoms twisting in a snakelike motion until one of the snakes seized hold of its own tail. This dream led Kekulé to propose that benzene consists of a ring of six carbon atoms with alternating COC sin-gle bonds and CPC double bonds. Because there are two ways in which the bonds can al-ternate, Kekulé proposed that benzene was a mixture of two compounds in equilibrium.
H H
H H
H H
H
H H H
H H
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Kekulé’s structure explained the formula of benzene, but it did not explain why ben-zene failed to behave like an alkene. The unusual stability of benben-zene wasn’t understood until the development of the theory of resonance. This theory states that molecules for which two or more satisfactory Lewis structures can be drawn are an average, or hybrid, of the structures. Benzene, for example, is a resonance hybrid of the two Kekulé structures.
The difference between the equilibrium and resonance descriptions of benzene is subtle, but important. In the equilibrium approach, a pair of arrows is used to describe a reversible reaction, in which the molecule on the left is converted into the one on the right, and vice versa. In the resonance approach, a double-headed arrow is used to suggest that a benzene molecule does not shift back and forth between two different structures; it is a hybrid mix-ture of the strucmix-tures.
One way to probe the difference between Kekulé’s idea of an equilibrium between two structures and the resonance theory in which benzene is a hybrid mixture of the structures would be to study the lengths of the carbon–carbon bonds in benzene. If Kekulé’s idea was correct, we would expect to find a molecule in which the bonds alternate between relatively long COC single bonds (0.154 nm) and significantly shorter CPC double bonds (0.133 nm). When benzene is cooled until it crystallizes, and the structure of the molecule is studied by X-ray diffraction, we find that the six carbon–carbon bonds in the molecule are the same length (0.1395 nm). The crystal structure of benzene is there-fore more consistent with the resonance model of bonding in benzene than the original Kekulé structures.
The resonance theory does more than explain the structure of benzene—it also ex-plains why benzene is less reactive than an alkene. The resonance theory assumes that mol-ecules that are hybrids of two or more Lewis structures are more stable than those that aren’t. It is the extra stability that makes benzene and other aromatic derivatives less re-active than normal alkenes. To emphasize the difference between benzene and a simple alkene, many chemists replace the Kekulé structures for benzene and its derivatives with an aro-matic ring in which the circle in the center of the ring indicates that the electrons in the ring are delocalized; they are free to move around the ring.
It is the delocalization of electrons around the aromatic ring that makes benzene less re-active than a simple alkene.
Aromatic compounds were being extracted from coal tar as early as the 1830s. As a result, many of the compounds were given common names that are still in use today. A
H
H H
H H
H H
H
H H
H H
H
H H H
H H
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few of these compounds are shown below.
There are three ways in which a pair of substituents can be placed on an aromatic ring.
In the ortho (o) isomer, the substituents are in adjacent positions on the ring. In the meta (m) isomer, they are separated by one carbon atom. In the para (p) isomer, they are on opposite ends of the ring. The three isomers of dimethylbenzene, or xylene, are shown below.
Exercise O1.8
Predict the structure of para-dichlorobenzene, one of the active ingredients in mothballs.
Solution
Para isomers of benzene contain two substituents at opposite positions in the six-membered ring. para-Dichlorobenzene therefore has the following structure.
Aromatic compounds can contain more than one six-membered ring. Naphthalene, an-thracene, and phenanthrene are examples of aromatic compounds that contain two or more fused benzene rings.
A ball-and-stick model of anthracene is shown in Figure O1.18.
H
H H
H H
H H
H H
H H
H H H
H H
H H
H H
H H
H H H H H H
Naphthalene (C10H8) Anthracene (C14H10)
Phenanthrene (C14H10)
Cl Cl
p-Dichlorobenzene
CH3
CH3
CH3
CH3
CH3
CH3
Para Meta
Ortho
CH3 OH OCH3 NH2 CO2H
Toluene Phenol Anisole Aniline Benzoic acid
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