The Bonding in Benzene

Kekule warned against publishing dreams when he said, Let us learn to dream, and perhaps then we shall learn the truth. But let us also beware not to publish our dreams until they have been examined by the wakened mind. In 1895, he was made a nobleman by Emperor William II of Germany. This allowed him to add von Stradonitz to his name. Kekule s students received three of the first five Nobel Prizes in Chemistry van t Hoff in 1901, Fischer in 1902, and Baeyer in 1905. 

Between 1865 and 1890, other possible structures were proposed for benzene such as those shown here. Considering what nineteenth-century chemists knew about benzene, which is a better proposal for benzene s structure, Dewar benzene or Ladenburg benzene Why  

Each of the six tt electrons of benzene, therefore, is localized neither on a single carbon nor in a bond between two carbons (as in an alkene). Instead, each tt electron is shared by all six carbons. In other words, the six tt electrons are delocalized—they roam freely 

Delocalized Electrons and Their Effect on Stability, p/Q, and the Products of a Reaction 

Electron delocalization is shown by double-headed arrows ( — ), whereas equilibrium is shown by two arrows pointing in opposite directions 

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A difficulty arises when we attempt to represent the bonding in benzene and its derivatives. We might write... 

Both of these structures satisfy the formal valence rules for carbon, but each has a serious fault. Each structure shows three of the carbon-carbon bonds as double bonds, and three are shown as single bonds. There is a wealth of experimental evidence to indicate that this is not true. Any one of the six carbon-carbon bonds in benzene is. the same as any other. Apparently the fourth bond of each carbon atom is shared equally with each adjacent carbon. This makes it difficult to represent the bonding in benzene by our usual line drawings. Benzene seems to be best represented as the superposition or average of the two structures. For simplicity, chemists use either one of the structures shown in (30) usually expressed in a shorthand form (SI) omitting the hydrogen atoms ... 

A possibly more accurate value for the double bond character of the bonds in benzene (0.46) id obtained by considering all five canonical structures with weights equal to the squares of their coefficients in the wave function. There is some uncertainty aS to the significance of thfa, however, because of- the noii -orthogOnality of the wave functions for the canonical structures, and foF chemical purposes it fa sufficiently accurate to follow the simple procedure adopted above. 

We know today that benzene is a cyclic compound with the equivalent of three double bonds and three single bonds, as shown in Figure 1.9(A). However, the electrons that form the double bonds in benzene are spread out and shared over the whole molecule. Thus, benzene actually has six identical bonds, each one half-way between a single and a double bond. These bonds are much more stable than ordinary double bonds and do not react in the same way. Figure 1.9(B) shows a more accurate way to represent the bonding in benzene. Molecules with this type of special electron sharing are called aromatic compounds. As mentioned earlier, benzene is the simplest aromatic compound. 

With use of Equation 7-5 we calculate 1.420 A for the length of a carbon-carbon bond with bond order 1.5. This may be taken a8 the length that the bonds in benzene would have if there were no stabilizar tion (and consequent bond shortening) by resonance. The actual bond length in benzene is 1.397 A, and we conclude that resonance between the two Kekul6 structures decreases the bond length by about.0.023 A. 

This, in fact, does not happen. The benzene ring is a very stable structure that can be involved in many reactions in which the reaction products retain the benzene ring. Thus the bonds in benzene rings must be much more stable than double bonds. 

Modern instrumental studies confirm earlier experimental data that all the bonds in benzene are of equal length, approximately 1.40 pm. (A picometer equals 1 x 10 meter.) This bond length falls exactly halfway between the length of a carbon-carbon single bond (1.46 pm) and a carbon-carbon double bond (1.34 pm). In addition, these studies confirm that all bond angles are equal (120°) and that the benzene molecule has a planar (flat) structure. 

Benzene has two major resonance structures that contribute equally to the resonance hybrid. These are sometimes called Kekule structures because they were originally postulated by Kekule in 1866. You may also encounter benzene written with a circle inside the six-membered ring rather than the three double bonds. This representation is meant to show that the bonds in benzene are neither double nor single. However, the circle structure makes it difficult to count electrons. This text uses a single Kekule structure to represent benzene or its derivatives. You must recognize that this does not represent the true structure and picture the other resonance structure or call upon the MO model presented in Section 16.3 when needed. 

Thus the central circle in the resonance hybrid representation actually captures the physical reality of the bonding in benzene. 

We will illustrate this procedure by considering the bonding in benzene, an important industrial chemical that must be handled carefully because it is a known carcinogen. The benzene molecule (C6H6) consists of a planar hexagon of carbon atoms with one hydrogen atom bound to each carbon atom, as shown in Fig. 14.48(a). In the molecule all six C—C bonds are known to be equivalent. To explain this fact, the LE model must invoke resonance [see Fig. 14.48(b)]. 

A better description of the bonding in benzene results when we use a combination of the models, as described above. In this description it is assumed that the cr bonds to carbon involve sp2 orbitals, as shown in Fig. 14.49. These cr bonds are all centered in the plane of the molecule. 

A special class of cyclic unsaturated hydrocarbons is known as the aromatic hydrocarbons. The simplest of these is benzene (C6H6), which has a planar ring structure, as shown in Fig. 22.11(a). In the localized electron model of the bonding in benzene, resonance structures of the type shown in Fig. 22.11(b) are used to account for the known equivalence of all the carbon-carbon bonds. But as we discussed in Section 14.5, the best description of the benzene molecule assumes that sp2 hybrid orbitals on each carbon are used to form the C—C and C—H a bonds, while the remaining 2p orbital on each carbon is used to form 77 molecular orbitals. The delocalization of these 1r electrons is usually indicated by a circle inside the ring [Fig. 22.11(c)]. 

This definition takes us naturally to the following question are all attractive interactions bonds which can be rewritten as where do we put the bonds This issue is faced in complex molecules and intermolecular aggregates, and originates in the fact that Pauling s definition (except for the trivial case of a diatomic molecule) only allows to identify the existence of a bond, but not the atoms involved in it. We can illustrate the problem in two examples, selected as prototypes of intramolecular and intermolecular bonds where do we plot the bonds in benzene and in water dimer ... 

If the carbon-carbon bonds all have the same length, they must also have the same number of electrons between the carbon atoms. This can be so, however, only if the tt electrons of benzene are delocalized around the ring, rather than each pair of tt electrons being localized between two carbon atoms. To better understand the concept of delocalized electrons, we ll now take a close look at the bonding in benzene. 

We will illustrate the general method by considering the bonding in benzene, an important industrial chemical that must be handled carefully because it is a known carcino-... 

What are aromatic hydrocarbons Benzene exhibits resonance. Explain. What are the bond angles in benzene Give a detailed description of the bonding in benzene. The TT electrons in benzene are delocalized, while the tt electrons in simple alkenes and alkynes are localized. Explain the difference. 

Benzene is stabilized by delocalized molecular orbitals. The C-C bonds are equivalent, rather than alternating single and double bonds. The additional stabilization makes the bonds in benzene much less reactive chemically than isolated double bonds such as those in ethylene. 

This structure represents the fact that electrons from the carbon-carbon double bonds are spread out or delocalized over the whole molecule, which makes the molecule more stable. A Lewis structure is not sophisticated enough to describe the bonding in benzene. In order to obtain a more satisfactory description of the bonding, the electrons must be treated as clouds (rather than crosses or dots ). For a brief description of the bonding model, see Box 17.8. 

The observed properties of benzene can be explained using a more sophisticated model, than that proposed by Kekule, for the bonding in benzene. Each carbon atom in the ring is bonded to a hydrogen and to one carbon on either side of it by ct bonds. Every carbon has one unused p orbital, containing a single electron, perpendicular to the plane. This gives a p/anar skeleton as shown in Eig. 17.5. 

Because we caimot describe the bonds in benzene as individual bonds between... 

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