Delocalized Electrons Explain Benzenes Structure

Because early organic chemists did not know about delocalized electrons, they were puzzled by benzene s stmcture. They knew that benzene had a molecular formula of CgHg, that it was an unusually stable compound, and that it did not undergo the addition reactions characteristic of alkenes (Section 6.0). They also knew that when a different atom was substituted for any one of benzene s hydrogens, only one product was obtained and, when the substituted product underwent a second substitution, three products were obtained. 

For every two hydrogens that are missing from the general molecular formula, C H2 +2, a hydrocarbon has either a tt bond or a ring. 

What kind of structure would you predict for benzene if you knew only what the early chemists knew The molecular formula (CsH ) tells us that benzene has eight fewer hydrogens than an acyclic alkane with six carbons (C H2 +2 = C6H14). Benzene, therefore, has a degree of unsaturation of four. In other words, the total number of rings and TT bonds in benzene is four (Section 5.1). 

Because only one product is obtained regardless of which of the six hydrogens of benzene is replaced with another atom, we know that all the hydrogens must be identical. Two sttuc-tures with a degree of unsaturation of four and six identical hydrogens are shown here  

Neither of these stmctures, however, is consistent with the observation that three compounds are obtained if a second hydrogen is replaced with another atom. The acyclic stmcture yields only two disubstituted products. 

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The resonance-delocalized picture explains most of the structural properties of benzene and its derivatives—the benzenoid aromatic compounds. Because the pi bonds are delocalized over the ring, we often inscribe a circle in the hexagon rather than draw three localized double bonds. This representation helps us remember there are no localized single or double bonds, and it prevents us from trying to draw supposedly different isomers that differ only in the placement of double bonds in the ring. We often use Kekule structures in drawing reaction mechanisms, however, to show the movement of individual pairs of electrons. 

CsHi4 CsHs + 4H2 Benzene is the simplest aromatic hydrocarbon. It shows characteristic electrophilic substitution reactions, which are difficult to explain assuming a simple unsaturated structure (such as Kekule s (1865) or Dewar s (1867) formulae). This anomolous behavior can now be explained by assuming that the six pi electrons are delocalized and that benzene is, therefore, a resonance hybrid. It consists of two Kekule structures and three Dewar structures, the former contributing 80% and the latter 20% to the hybrid. Benzene is usually represented by either a Kekule structure or a hexagon containing a circle (which denotes the delocalized electrons). 

In contrast to VBT, "full-blown" MOT considers the electrons in molecules to occupy molecular orbitals that are formed by linear combinations (addition and subtraction) of all the atomic orbitals on all the atoms in the structure. In MOT, electrons are not confined to an individual atom plus the bonding region with another atom. Instead, electrons are contained in MOs that are highly delocalized—spread across the entire molecule. MOT does not create discrete and localized bonds between neighboring atoms. An immediate benefit of MOT over VBT is its treatment of conjugated tt systems. We don t need a "patch" like resonance to explain the structure of a carboxylate anion or of benzene it falls naturally out of the delocalized nature of the MOs. The MO models of simple molecules like ethylene or formaldehyde also lead to bonding concepts that are pervasive in organic chemistry. 

The concept of a sea of electrons not belonging to any particular atom is reminiscent of the resonance structures covered earlier. The valence electrons in a metal are delocalized just as they are in resonance molecules. The mobile electrons in a bar of sodium are not associated with any particular ion core, just as the electrons in the double bonds of benzene are not associated with any particular atom. To explain this phenomenon in metals, one must apply molecular orbital theory. 

Cyclic compounds like benzene are not the only ones with delocalized molecular orbitals. Let s look at bonding in the carbonate ion (CO3 ). VSEPR predicts a trigonal planar geometry for the carbonate ion, like that for BF3. The planar structure of the carbonate ion can be explained by assuming that the carbon atom is ip -hybridized. The C atom forms sigma bonds with three O atoms. Thus the unhybridized 2p orbital of the C atom can simultaneously overlap the 2p orbitals of all three O atoms (Figure 10.29). The result is a delocalized molecular orbital that extends over all four nuclei in such a way that the electron densities (and hence the bond orders) in the carbon-to-... 

The equal sharing of the six electrons of the p orbitals results in a rigid, flat ring structure, in contrast to the relatively flexible, nonaromatic cyclohexane ring. The model also explains the unusual chemical stability of benzene and its resistance to addition reactions. The electrons of the tt cloud are said to be delocalized. That means they have much more space and freedom of movement than they would have if they were restricted to individual double bonds. Because electrons repel one another, the system is more stable when the electrons have more space to occupy. As a result, benzene is unusually stable and resists addition reactions typical of alkenes. 

Delocalization of the leftover p electrons is believed to make benzene more stable than would be expected if the Kekule structures were correct, and explains why benzene does not react as if it had three double bonds. The model predicts that all the carbon-carbon bond lengths in benzene would be of the same length, which they are. The length of each bond is between those normally observed for C=C and C-C. 

Resonance structures extend the utility of Lewis theory in explaining electron delocalization In certain structures. Resonance structures are an integral part of Lewis theory that allows rationalization of chemical and physical properties In a simple and predictive manner. Lewis theory follows the Idea of electron pair bonds, and a typical example where resonance Is used Is for describing the cyclic structure of benzene,... 

An alternative but equivalent model for describing benzene (and other resonance-stabilized structures) is molecular orbital theory. We have already seen how this theory can explain the formation of molecular structures such as methane, ethene and others. In localized molecules like ethene, C2H4, two unhybridized p orbitals overlap to form a Jt molecular orbital in which a pair of electrons is shared between tbe nuclei of two carbon atoms. In molecular orbital theory, resonance-stabilized structures are described in terms of delocalized Jt orbitals where the Jt electron clouds extend over three or more atoms. 

Quantum mechanics, even in the crudest approximation as formulated in Hiickel theory [23], is suitable for explaining the n aromaticity. Application of this theory clarifies the planarity and stability of benzene and leads to the well-known (4n + 2)n electron rule for annulenes. The delocalized n electrons affect not only the structure (planarity and high symmetry) and enhanced stability of aromatic systems they also... 

Mesomery is the delocalization of certain electrons on several atoms of a molecule. This delocalization is generally referred to as resonance or conjugation. It is accompanied by an energetic stabilization of the chemical structure. The mesomeric effect concerns pi (n) electrons, pairs of free electrons, and charges. The cases of benzene and butadiene are often used as examples to explain the phenomenon of conjugation. 

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