Bonding cyclohexatriene

Our experience has been that some 125 kJ/mol (30 kcal/mol) is given off when ever a double bond is hydrogenated When benzene combines with three molecules of hydrogen the reaction is far less exothermic than we would expect it to be on the basis of a 1 3 5 cyclohexatriene structure for benzene... 

Does this mean that benzene is in fact an equilibrium mixture of two 1,3,5-cyclohexatriene molecules Or, does this mean that benzene is a single molecule with some intermediate type of bonding ... 

Cyclohexatriene to benzene displays a sequence of structures from 1,3,5-cyclohexatriene (withCC single and double bonds initially set to 1.5 and 1.3 A, respectively) to benzene (witb all CC bonds set to 1.4 A) and back to cyclohexatriene. Plot energy (vertical axis) vs. CC bond length (horizontal axis). How many energy minima are there Do the minima look more like 1,3,5-cyclohexatriene or benzene What is the correct interpretation of the resonance picture ... 

It is important to note that benzene does not behave like a typical cyclic olefin in that the benzene ring undergoes ionic substitution rather than addition reactions the ring also resists hydrogenation and is chemically more inert. Despite this, it is still a common practice to represent benzene with three double bonds as if it were 2,4,6-cyclohexatriene,... 

Canonical forms of benzene that are calculated to contribute about 22% to the resonance stabilization of benzene. Such resonance structures have no separate physical reality or independent existence. For the case of benzene, the two Kekule structures with alternating double bonds i.e., cyclohexatriene structures) contribute equally and predominantly to the resonance hybrid structure. A dotted circle is often used to indicate the resonance-stabilized bonding of benzene. Nonetheless, the most frequently appearing structures of benzene are the two Kekule structures. See Kekule Structures... 

Canonical forms of benzene with alternating double bonds Le., cyclohexatriene ), structures which contribute equally and predominantly to benzene s resonance hybrid structure. 

If tetrahydro[60]fullerene (C5QH4) is formed by additions to two [6,6] double bonds, which are two 1,2-additions with respect to the cyclohexatriene units in C q, then eight regioisomers are possible [37, 38] (see Chapter 10). This very plausible assumption is corroborated by the theoretical investigations of multiple additions to C50 in a 1,2- and a 1,4-mode. These investigations predict 1,2-additions to be favorable over 1,4-additions up to the formation of CjqHj2 (see also Table 5.5 below)... 

Following these ideas, Table 15.10 shows results of 6-3IG calculations of the TV system of normal benzene and benzene distorted to have alternating bond lengths matching standard double and single bonds, which we will call cyclohexatriene. 

When we look at the energies from Table 15. 10, perhaps the most striking fact is that the correlation energy in the n system makes so little difference in the Ai values. As we indicated above, the experimental value for the resonance energy from heats of hydrogenation is —1.54 eV, in quite satisfactory agreement with the result in Table 15.10. The fact that our value is a little lower than the experimental one may be attributed to the small amoimt of residual resonance remaining in the cyclohexatriene, whereas the isolated double bonds in the experiment are truly isolated in separate molecules. ... 

Pyridine and benzene conform to Hiickers rule, which predicts that planar cyclic polyenes containing (4n + 2) -electrons ( = 0, or an integer) should show added stability over that anticipated for theoretical polyenes composed of formal alternate single and double bonds. This difference is sometimes called the empirical resonance energy. For example, benzene, where n = 1, is estimated to be 150 kJ moT more stable than the hypothetical molecule cyclohexatriene (Box 1.8) for pyridine, the empirical resonance energy is 107 kJ mol . ... 

The heat of hydrogenation of the cyclic diene (8) is very nearly twice that of cyclohexene (7), and the heat of hydrogenation of the three double bonds in a Kekule structure might thus be expected to be of the order of 3 x -120 kJ ( — 28-6 kcal)mol = -360 kJ ( — 85-8 kcal) mol but when real benzene is hydrogenated only —208 kJ (-49-8 kcal) mol are evolved. Real benzene is thus thermodynamically more stable than the hypothetical cyclohexatriene by 151 kJ (36 kcal) mol this compares with only 17 kJ (4 kcal)mol by which a conjugated diene i stabilised, with respect to its analogue in which there is no interaction between the electrons of the double bonds. 

Problem 10.1 Benzene is a planar molecule with bond angles of 120°. All six C-to-C bonds have the identical length, O.I39nm. Is benzene the same as 1,3,5-cyclohexatriene <... 

No. The bond lengths in 1,3,5-cyclohexatriene would alternate between 0,153nm for the single bond and 0.132nm for the double bond. The C-to-C bonds in benzene are intermediate between single and double bonds. 

In like manner the cycloparaffins form the structural basis of the cyclic group. Cycloolefins, -diolefins, and -triolefins may be regarded as derivatives of the cycloparaffins having the same carbon skeleton. The aromatic series forms a particular case among the cyclohexatrienes, in which the three double bonds are spaced symmetrically around a six-atom carbon ring. 

We have already shown that the energy of the system of six n electrons in benzene is equal to 8p. We must now calculate, in units of / , the energy of a Kekule structure, that is, the energy of the hypothetical molecule cyclo-hexatriene. In cyclohexatriene there are three localized n bonds. When two atomic n orbitals, say < , and (j)2, interact to form a two-center bond, two MOs, j/1 and yr2, are formed. In order that these be real, normalized, and... 

Experimentally, the delocalization energy of benzene is estimated in the following way. The actual enthalpy of formation of benzene can be determined by thermochemical measurements. The energy of the hypothetical molecule cyclohexatriene can be estimated by using the bond energies for C—C, C=C, and C—H found in other molecules such as ethane and ethylene. The difference between these energies is the experimental value of the delocalization energy. We then evaluate / , since... 

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