Bonding in benzene

The pattern of orbital energies is different for benzene from the pattern it would have if the six it electrons were confined to three noninteracting double bonds. The delocalization provided by cyclic conjugation in benzene causes its it electrons to be held more strongly than they would be in the absence of cyclic conjugation. Stronger binding of its it electrons is the factor most responsible for the special stability—the aromaticity—of benzene. 

The TT molecular orbitals of benzene arranged In order of increasing energy and showing nodai surfaces. The six -ir eiectrons of benzene occupy the three lowest energy orbitals, all of which are bonding. 

Later in this chapter we ll explore the criteria for aromaticity in more detail to see how they apply to cyclic polyenes of different ring sizes. The next several sections introduce us to the chemistry of compounds that contain a benzene ring as a structural unit. We ll start with how we name them 

Electrostatic potential map of benzene. The red area in the center corresponds to the region above and below the plane of the ring where the ir electrons are concentrated. 

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Benzene was probably the fust compound in chemical history where the valence bond concept proved to be insufficient. Localizing the nr-systems, one comes up with two equivalent but different representations. The true bonding in benzene was described as resulting from a resonance between these two representations (Figure 2-46). 

Twentieth-century theories of bonding in benzene gave us a clearer picture of ar omaticity. We ll start with a resonance description of benzene. 

An important property of aromatic hydrocarbons is that they are much more stable and less reactive than other unsaturated compounds. Benzene, for example, does not react with many of the reagents that react rapidly with alkenes. When reaction does take place, substitution rather than addition is observed. The Kekule formulas for benzene seem inconsistent with its low reactivity and with the fact that all of the C—C bonds in benzene are the sane length (140 pm). 

Examine the geometry of methylbenzyne. Measure carbon-earbon distances. Which 7C bonds are deloealized and whieh are localized Is there really a triple bond (Compare bond distance to triple bond in hexa-l,5-dien-3-yne and to partial double bonds in benzene). Are you able to draw a single Lewis structure whieh adequately represents the geometry of the molecule ... 

Experimentally, all the C-C bonds in benzene are of equal length, and this is mirrored by the C-C bond orders. 

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 ... 

Bond lengths All the carbon-carbon bonds in benzene are the same length. 

The values in Table 2.4 show how resonance affects the strengths of bonds. For example, the strength of a carbon-carbon bond in benzene is intermediate between that of a single and that of a double bond. Resonance spreads multiple bond character over the bonds between atoms as a result, what were single bonds are strengthened and what were double bonds are weakened. The net effect overall is a stabilization of the molecule. 

FIGURE 3.20 The framework of a-bonds in benzene each carbon atom is sp2 hybridized, and the array of hybrid orbitals matches the bond angles (of 120°) in the hexagonal molecule. The bonds around only one carbon atom are labeled all the others are the same. 

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. 

The difficulty in finding a satisfactory representation for the carbon-carbon bonding in benzene brings home to us the fact that our normal way of writing bonds between atoms as single, double or triple, involving... 

Facile C-H bond activation by Pt(II) metal centers seems to require at least one labile ligand in the coordination sphere of platinum. One of the earliest intermolecular examples of this is the activation of C-D bonds in benzene-f/, by 0 an.S -(PAIe .) Pt(neopentyl)(OTf) at 133 °C, where trifluoromethanesulfonate (triflate, OTf) provides the labile group (Scheme 7, A) (26). 

Applied to benzene this conception provides for a structure which is much more saturated than that possible in olefines with several double bonds. In benzene the points of attack, which are always present in the open chain, are lacking ... 

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. 

The photochemistry of borazine delineated in detail in these pages stands in sharp contrast to that of benzene. The present data on borazine photochemistry shows that similarities between the two compounds are minimal. This is due in large part to the polar nature of the BN bond in borazine relative to the non-polar CC bond in benzene. Irradiation of benzene in the gas phase produces valence isomerization to fulvene and l,3-hexadien-5-ynes Fluorescence and phosphorescence have been observed from benzene In contrast, fluorescence or phosphorescence has not been found from borazine, despite numerous attempts to observe it. Product formation results from a borazine intermediate (produced photochemically) which reacts with another borazine molecule to form borazanaphthalene and a polymer. While benzene shows polymer formation, the benzyne intermediate is not known to be formed from photolysis of benzene, but rather from photolysis of substituted derivatives such as l,2-diiodobenzene ... 

Bonding In benzene can be described In terms of sp hybridisation, o-bonds and a 71 molecular orbital containing delocalised electrons. 

For an overview of the structure and bonding In benzene, visit www. brlghtredbooks.net... 

Kekule, a German chemist, was the first to propose a structure for benzene In 1865. It was a cyclic structure of alternating single and double bonds. The Kekule structure, however, does not fit all the evidence that chemists have since collected. For example, when benzene Is added to a bromine solution, rapid decolourlsatlon does not take place. This Implies that benzene resists addition reactions and Is much more stable than a typical unsaturated hydrocarbon. The reason for this becomes clear when we examine more closely the structure and bonding In benzene. 

Each carbon-carbon double bond is constructed from four electrons. In benzene, the electrons that create the apparent double bonds fall into two classes. Two of the electrons are localized between two carbon atoms, just as we have come to expect. The other two electrons that contribute to the apparent double bonds are, in contrast, delocalized over the entire molecule. Since there are three apparent double bonds, we have a total of six electrons that are delocalized over the six carbon atoms. Think of these as free-range electrons. Basically, each of the carbon-carbon bonds in benzene is a 1.5 bond (technically, we say that the bond order in benzene carbon-carbon bonds is 1.5). Hence, the two models for benzene employed above, though universally used in chemistry, leave something to be desired. Benzene is better thought of as a hybrid of the two. Chemists have struggled with ways to depict the reality of benzene better than the stractures A and B. The struggle has not been notably successful. 

The addition of silyl radicals to double bonds in benzene or substituted benzenes (Reaction 5.2) is the key step in the mechanism of homolytic aromatic substitution with silanes [8,9]. The intermediate cyclohexadienyl radical 2 has been detected by both EPR and optical techniques [21,22]. Similar cyclohex-adienyl-type intermediates have also been detected with heteroaromatics like furan and thiophene [23]. 

The direct silylation of arenes through C—H bond activation provides an attractive route for the synthesis of useful aromatic compounds [64]. Vaska s complex was the first of the iridium catalysts to be reported for activation of the C—H bond in benzene by Si—H of pentamethyldisiloxane to yield phenylsubstituted siloxane [65]. However, a very attractive method for the aromatic C—H silylation with disilanes has been recently reported by the groups of Ishiyama and Miyaura [66-68]. 

The bond lengths of quinoline, which are irregular, support the resonance description thus, the 1,2-, 5,6- and 7,8-linkages are shorter than that of the carbon-carbon bond in benzene (more double bond character ). There is also a dipole of 2.19 D directed towards the nitrogen atom. 

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