Cyclobutadiene and Cyclooctatetraene

The relative stability of benzylic carbocations, radicals, and carbanions makes it possible to manipulate the side chains of aromatic rings. Functionalization at the benzylic position, for example, is readily accomplished by free-radical halogenation and provides access to the usual reactions (substitution, elimination) that we associate with alkyl halides. 

To illustrate, consider the synthesis of (Z)-1 -phenyl-2-butene. A major consideration— controlling the stereochemistry of the double bond—can be achieved by catalytic hydrogenation of the corresponding alkyne. 

The question then becomes one of preparing l-phenyl-2-butyne. A standard method for the preparation of alkynes is the alkylation of acetylene and other alkynes. In the present case, a suitable combination is propyne and a benzylic halide. The benzylic halide can be prepared from toluene. 

Use retrosynthetic analysis to describe the preparation of frans-2-phenylcyclopentanol from cyclopenfylbenzene and write equations showing suitable reagents for the synthesis. 

During our discussion of benzene and its derivatives, it may have occurred to you that cyclobutadiene and cyclooctatetraene might be stabilized by cyclic tt electron delocalization in a maimer analogous to that of benzene. 

The same thought occurred to early chemists. However, the complete absence of naturally occurring compounds based on cyclobutadiene and cyclooctatetraene contrasted starkly with the abundance of compounds containing a benzene unit. Attempts to synthesize cyclobutadiene and cyclooctatetraene met with failure and reinforced the growing conviction that these compounds would prove to be quite unlike benzene if, in fact, they could be isolated at all. 

The first breakthrough came in 1911 when Richard Willstatter prepared cyclooctatetraene by a lengthy degradation of pseudopelletieriney a natural product obtained from the bark of the pomegranate tree. Today, cyclooctatetraene is prepared from acetylene in a reaction catalyzed by nickel cyanide. 

Cyclooctatetraene is relatively stable, but lacks the special stability of benzene. Unlike benzene, which we saw has a heat of hydrogenation that is 152 kJ/mol (36 kcal/mol) less than three times the heat of hydrogenation of cyclohexene, cyclooctatetraene s heat of hydrogenation is 26 kJ/mol (6 kcal/mol) more than four times that of c/5-cyclooctene. 

Both cyclooctatetraene and styrene have the molecular formula CgHg and undergo combustion according to the equation 

Willstatter s most important work, for which he won the 1915 Nobel Prize in chemistry, was directed toward determining the structure of chlorophyll. 

Thermochemical measurements suggest a value of only about 20 kJ/mol (about 5 kcal/mol) for the resonance energy of cyclooctatetraene, far less than the aromatic stabilization of benzene (152 kJ/mol 36 kcal/mol). 

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One of molecular orbital theories early successes came m 1931 when Erich Huckel dis covered an interesting pattern m the tt orbital energy levels of benzene cyclobutadiene and cyclooctatetraene By limiting his analysis to monocyclic conjugated polyenes and restricting the structures to planar geometries Huckel found that whether a hydrocarbon of this type was aromatic depended on its number of tt electrons He set forth what we now call Huckel s rule... 

Benzene cyclobutadiene and cyclooctatetraene provide clear examples of Huckel s rule Benzene with six tt electrons is a An + 2) system and is predicted to be aromatic by the rule Square cyclobutadiene and planar cyclooctatetraene are An systems with four and eight tt electrons respectively and are antiaromatic... 

The data of Figure 21-13 provide a rationale for the instability of cyclobutadiene and cyclooctatetraene. For cyclobutadiene, we can calculate that four 77 electrons in the lowest orbitals will lead to a predicted 77-electron energy of 2(a + 2/3) + 2(a) = 4a + 4/3, which is just the 77-electron energy calculated for two ethene bonds (see Figure 21-3). The delocalization energy of the 77 electrons of cyclobutadiene therefore is predicted to be zero ... 

The simple resonance theory fails to explain the singular lack of effectiveness of delocalization in cyclobutadiene and cyclooctatetraene, but we may turn to molecular orbitals for the solution. 

In contrast to aromatic molecules which have An + 2 n electrons, cyclobutadiene and cyclooctatetraene do not have An + 2 7r electrons and are not aromatic. In fact, diese molecules, which contain An jt electrons (n is an integer), are less stable than die planar model compounds and are termed antiaromatic. Bodi of these molecules adopt shapes that minimize interactions of die n orbitals. 

Figure 18. (a) jt—a-Interplay diagrams for the twin states of /7-electron—/7-center antiaromatic species. (b,c) Calculated geometries for the twin states of cyclobutadiene and cyclooctatetraene (from refs 164a, 177, 188, and 219). 

For the double bonds to be completely conjugated, the annulene must be planar so the p orbitals of the pi bonds can overlap. As long as an annulene is assumed to be planar, we can draw two Kekule-like structures that seem to show a benzene-like resonance. Figure 16-3 shows proposed benzene-like resonance forms for cyclobutadiene and cyclooctatetraene. Although these resonance structures suggest that the [4] and [8]annulenes should be unusually stable (like benzene), experiments have shown that cyclobutadiene and cyclooctatetraene are not unusually stable. These results imply that the simple resonance picture is incorrect. 

Cyclobutadiene and cyclooctatetraene have alternating single and double bonds similar to those of benzene. These compounds were mistakenly expected to be aromatic. 

Valence bond theory, in the terms defined by Pauling, is not able to account for the 4n+2 rule, and the properties of cyclobutadiene and cyclooctatetraene. It has been suggested that the problem with these molecules is the strain associated with the bond angles in the planar structures.10 However, this was shown to be incorrect by the observation that the addition of two electrons to cyclooctatetraene leads to the planar dianion. It is only recently that it has been recognized that cyclic permutations must be included in order to properly treat cyclic systems via valence bond theory.11 One of Pauling s few failures in structural theory is his nonrecognition of the problems associated with the 4n molecules. 

Just as the unusual stability and reactivity of benzene are placed into their proper context by comparison with cyclobutadiene and cyclooctatetraene, the 4 -electron homo-logues of benzene, it is instructive to compare the formally homoantiaromatic bicyclo [3.1.0]hexenyl/cyclohexadienyl cation systems with the homocyclopropenium and homo-tropenylium ions (Scheme 14). Such a comparison not only puts in context the properties of the latter two homoaromatic cations, but also reveals a different mode of cyclopropyl conjugation that occurs in the 4 -electron systems. 

In accord with the HOckel rule of 4/i + 2 electrons, both cyclobutadiene and cyclooctatetraene (cot) are nonaromatic. Cyclooctatetraene contains alternating bond lengths and has a tub-shaped conformation ... 

According to Hiickel s rule, annulenes with 4n tt electrons are not aromatic. Cyclobutadiene and cyclooctatetraene are [4n]-annulenes, and their properties are more in accord with their classification as cyclic polyenes than as aromatic hydrocarbons. Among higher [4n]-annulenes, [16]-annulene has been prepared. [16]-Annulene is not planar and shows a pattern of alternating short (average 134 pm) and long (average 146 pm) bonds typical of a nonaromatic cyclic polyene. 

Q Draw MO diagrams for the 7i-electron systems of cyclobutadiene and cyclooctatetraene. 

Cyclobutadiene has two pairs of rr electrons, and cyclooctatetraene has four pairs of tt electrons. Unlike benzene, these compounds are not aromatic because they have an even number of pairs of rr electrons. There is an additional reason why cyclooctatetraene is not aromatic—it is not planar but, instead, tub-shaped. Earlier, we saw that, for an eight-membered ring to be planar, it must have bond angles of 135° (Chapter 2, Problem 28), and we know that sp carbons have 120° bond angles. Therefore, if cyclooctatetraene were planar, it would have considerable angle strain. Because cyclobutadiene and cyclooctatetraene are not aromatic, they do not have the unusual stability of aromatic compounds. 

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