Benzene resonance description

Twentieth century theories of bonding m benzene gave us a clearer picture of aromatic ity We 11 start with a resonance description of benzene... 

Because the carbons that are singly bonded m one resonance form are doubly bonded m the other the resonance description is consistent with the observed carbon-carbon bond distances m benzene These distances not only are all identical but also are intermediate between typical single bond and double bond lengths... 

Many of the properties of phenols reflect the polarization implied by the resonance description The hydroxyl oxygen is less basic and the hydroxyl proton more acidic in phenols than m alcohols Electrophiles attack the aromatic ring of phenols much faster than they attack benzene indicating that the ring especially at the positions ortho and para to the hydroxyl group is relatively electron rich... 

Having just seen a resonance description of benzene, let s now look at the alternative molecular orbital description. We can construct -tt molecular orbitals for benzene just as we did for 1,3-butadiene in Section 14.1. If six p atomic orbitals combine in a cyclic manner, six benzene molecular orbitals result, as shown in Figure 15.3. The three low-energy molecular orbitals, denoted bonding combinations, and the three high-energy orbitals are antibonding. 

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. 

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)]. 

Benzene is conjugated, so we must use resonance and orbitals to describe its structure. The resonance description of benzene consists of two equivalent Lewis structures, each with three double bonds that alternate with three single bonds. 

The resonance description of benzene matches the Kekule description with one important exception. The two Kekule representations are not in equilibrium witb each other. Instead, the true structure of benzene is a resonance hybrid of the two Lewis structures, with the dashed lines of the hybrid indicating the position of the 7i bonds. 

Having just seen a resonance description of benzene, let s now look at the alternative molecular orbital description. An orbital view of benzene makes clear the cyclic conjugation of the benzene molecule and the equivalence of the. six carbon-carbon bonds. Benzene is a planar molecule with the shape of a regular hexagon. All C-C-C bond angles are 120 , all six carbon atoms are sp -hybridized, and each carbon has a p orbital perpendicular to the plane of the six-membered ring. 

The most common example of the resonance theory is the description of the benzene structure. The experimentally precisely determined and accurately known carbon-carbon bond length is consistent with the model as average of the resonance structures. When Pauling s resonance description of the benzene structure was criticized, the physicist Edward Teller and his colleagues provided spectroscopic evidence to support it [40]. The Nobel laureate physicist Philip Anderson was oblivious of Teller s and his co-workers paper (Private communication from Philip Anderson to the author by e-mail in 2009), and 68 years after Teller s contribution, in 2008, Anderson communicated another supportive paper for Pauling s model [41]. 

Structurally benzene is the simplest stable compound having aromatic character, but a satisfactory graphical representation of its formula proved to be a perplexing problem for chemists. Kekule is usually credited with description of two resonating structures which. 

Benzene has already been mentioned as a prime example of the inadequacy of a connection table description, as it cannot adequately be represented by a single valence bond structure. Consequently, whenever some property of an arbitrary molecule is accessed which is influenced by conjugation, the other possible resonance structures have to be at least generated and weighted. Attempts have already been made to derive adequate representations of r-electron systems [84, 85]. 

Compare atomic charges and electrostatic potential maps for the three cations. For each, is the charge localized or delocalized Is it associated with an empty a-type or Tt-type orbital Examine the lowest-unoccupied molecular orbital (LUMO) of each cation. Draw all of the resonance contributors needed for a complete description of each cation. Assign the hybridization of the C" atom, and describe how each orbital on this atom is utilized (o bond, n bond, empty). How do you explain the benzene ring effects that you observe ... 

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