Benzene is hydrocarbon made up of 6 carbons and 6 hydrogen atoms. If forms a six sided ring.
It was once believed to be a structure that had alternating double and single bonds.
If this were the case, the bond lengths would be longer for the single bonds and shorter for the double. However this is not the case. All the bonds are the same length. Which means all the bonds of the ring are the same. You can think of the bonds as 1.5 bonds (not a single but not a double). A bond is made up of 2 electrons. Each carbon carbon bond has 3 electrons shared (1.5 bonds). Therefore all the bonds are the same length.
Chemists are very efficient, so here is how we draw benzene.
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History of Benzene (from Purdue)-
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 this compound had the empirical formula: CH. Ten years later, Eilhardt Mitscherlich produced the same material by heating benzoic acid with lime. Mitscherlich named this substance benzin, which became benzene when translated into English. He also determined that the molecular formula of this compound is C6H6.
Benzene is obviously 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 Br2 to the C=C or CC bonds, whereas benzene only reacts with bromine in the presence of a catalyst: FeBr3. Furthermore, when benzene reacts with Br2 in the presence of FeBr3, the product of this reaction is a compound in which a bromine atom has been substituted for a hydrogen atom, not added to the compound the way an alkene adds bromine.
Other compounds were eventually isolated from coal that had similar properties. Their formulas suggested the presence of multiple C=C bonds, but these compounds were not reactive 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 19th century. The first step toward solving this problem was taken by Friedrich August Kekule in 1865. (Kekule'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, Kekule 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 Kekule to propose that benzene consists of a ring of six carbon atoms with alternating CC single bonds and C=C double bonds. Because there are two ways in which these bonds can alternate, Kekule proposed that benzene was a mixture of two compounds in equilibrium.
Kekule's structure explained the formula of benzene, but it did not explain why benzene failed to behave like an alkene. The unusual stability of benzene 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 these structures. Benzene, for example, is a resonance hybrid of the two Kekule 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 mixture of these structures.
One way to probe the difference between Kekule's idea of an equilibrium between two structures and the resonance theory in which benzene is a hybrid mixture of these structures would be to study the lengths of the carbon-carbon bonds in benzene. If Kekule's idea was correct, we would expect to find a molecule in which the bonds alternate between relatively long CC single bonds (0.154 nm) and significantly shorter C=C 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 this molecule are the same length (0.1395 nm). The crystal structure of benzene is therefore more consistent with the resonance model of bonding in benzene than the original Kekule structures.
The resonance theory does more than explain the structure of benzene it also explains why benzene is less reactive than an alkene. The resonance theory assumes that molecules that are hybrids of two or more Lewis structures are more stable than those that aren't. It is this extra stability that makes benzene and other aromatic derivatives less reactive than normal alkenes. To emphasize the difference between benzene and a simple alkene, many chemists replace the Kekule structures for benzene and its derivatives with an aromatic 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.
The significance of the circle in the center of this aromatic ring might best be understood by asking: What is wrong with the Kekule structures for benzene? We start by building a sigma-bond skeleton for the benzene ring in which each of the carbon atoms is sp2 hybridized. This leaves us with one electron in a 2p orbital on each of the six carbon atoms.
If we assume that the interaction between the 2p orbitals is localized between a pair of carbon atoms, we get one of the Kekule structures for benzene. Switching the pairs of atoms that form bonds gives us the other Kekule structure.
But if we allow the six electrons in the six 2p orbitals to interact to form a set of molecular orbitals, we can delocalize the electrons so that they are free to move from one carbon atom to another around the ring.
It is this delocalization of electrons around the aromatic ring that is conveyed by the circle that is often written inside the ring. It is also the delocalization of electrons that makes benzene less reactive than a simple alkene.
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