Resonance hybrid electron delocalization, hybridization

Additionally sp hybridization of the hydroxyl oxygen allows one of its unshared electron pairs to be delocalized by orbital overlap with the tt system of the carbonyl group (Figure 19 1) In resonance terms this electron delocalization is represented as... 

Pyrrole has a planar, pentagonal (C2 ) stmcture and is aromatic in that it has a sextet of electrons. It is isoelectronic with the cyclopentadienyl anion. The TT-electrons are delocalized throughout the ring system, thus pyrrole is best characterized as a resonance hybrid, with contributing stmctures (1 5). These stmctures explain its lack of basicity (which is less than that of pyridine), its unexpectedly high acidity, and its pronounced aromatic character. The resonance energy which has been estimated at about 100 kj/mol (23.9 kcal/mol) is intermediate between that of furan and thiophene, or about two-thirds that of benzene (5). 

Resonance (Section 1.9) Method by which electron delocalization may be shown using Lewis structures. The true electron distribution in a molecule is regarded as a hybrid of the various Lewis structures that can be written for a molecule. 

Valence bond theory (Chapter 7) explains the fact that the three N—O bonds are identical by invoking the idea of resonance, with three contributing structures. MO theory, on the other hand, considers that the skeleton of the nitrate ion is established by the three sigma bonds while the electron pair in the pi orbital is delocalized, shared by all of the atoms in the molecule. According to MO theory, a similar interpretation applies with all of the resonance hybrids described in Chapter 7, including SO S03, and C032-. 

Taft and his coworkers14-16 developed a diparametric model which separated the electrical effect into contributions from the inductive (actually the field) and resonance effects. This separation depends on the difference in the extent of electron delocalization when a substituent is bonded to an sp3 -hybridized carbon atom in one reference system and to an sp2-hybridized carbon atom in another. As the first case represents minimal delocalization and the second extensive delocalization, we have referred to the two effects as the localized and delocalized electrical effects. This diparametric electrical effect model can be written in the form ... 

Many molecules that have several double bonds are much less reactive than might be expected. The reason for this is that the double bonds in these structures cannot be localized unequivocally. Their n orbitals are not confined to the space between the double-bonded atoms, but form a shared, extended Tu-molecular orbital. Structures with this property are referred to as resonance hybrids, because it is impossible to describe their actual bonding structure using standard formulas. One can either use what are known as resonance structures—i. e., idealized configurations in which n electrons are assigned to specific atoms (cf pp. 32 and 66, for example)—or one can use dashed lines as in Fig. B to suggest the extent of the delocalized orbitals. (Details are discussed in chemistry textbooks.)... 

Formally, pyrrole can be described as azacyclopentadiene, i.e. cyclopen-tadiene in which a CH2 unit has been replaced by a NH group. However, this translates into a classical structure that does not adequately describe the compound, for pyrrole has aromatic character, even though there are only five atoms in the ring The aromaticity arises because the nitrogen atom contributes two electrons and the four carbons one electron each to form a delocalized sextet of 7t-electrons. In valence bond terms the structure of pyrrole can be represented as a resonance hybrid (Scheme 6.1)... 

Exercise 6-10 Set up an atomic-orbital model of each of the following structures with normal values for the bond angles. Evaluate each model for potential resonance (electron delocalization). If resonance appears to you to be possible, draw a set of reasonable valence-bond structures for each hybrid. ... 

Clearly, it is inconvenient and tedious to write the structures of the contributing forms to show the structure of a resonance hybrid. A shorthand notation is therefore desirable. Frequently, dashed rather than full lines are used where the bonding electrons are expected to be delocalized over several atoms. For benzene, 16a or 16b is quite appropriate ... 

Dithiolenes are best considered to be a resonance hybrid of the limiting structures (1)—(3). In both bis- and tris-dithiolenes the electron delocalization is not limited to the ligand, but includes the metals to give rise to cyclic delocalization ( aromaticity ). To symbolize this electron delocalization in dithiolenes, they can be represented, in a manner similar to that used for benzene, by formulas containing a ring inside the framework given by the metal, sulfur and carbon atoms. We will use this notation, shown in (4), throughout this chapter. 

The bond order is between one (single) and two (double). The resonance hybrid is often pictured with a circle in the ring to indicate the delocalized electron distribution in the molecule. 

A molecule for which resonance forms can be written is more stable than any of the contributing resonance forms. Thus the carboxylate ion (a resonance hybrid) is more stable than either of the contributing resonance forms. The difference in energy between the energy of the molecule and the energy of the most stable resonance form is die resonance energy (RE) of die molecule. The resonance energy represents die stabilization of die molecule due to die delocalization of electrons. 

Although this process can be written to give a single canonical form A, it must be realized that the enolate is a delocalized species and resonance forms A and B can be generated as discussed previously using curved-arrow notation. This is not a mechanistic step since the delocalized product is a resonance hybrid of A and B—that is A is not converted to B, but rather the curved arrows merely indicate the changes in electron distribution that must be used to describe the canonical form B. 

In amides, the nitrogen electron pair is n-r conjugated with the carbonyl group and this electronic delocalization is normally expressed by resonance structures 1, 2, and 3. As a result, the amide function is essentially planar and it is assumed that the three atoms (C, ii, and 0) of this function are sp hybridized. The amide function can be illustrated in three dimensions by structure 4. The electronic distribution can also be viewed as the result of the delocalization of two n electron pairs, one from the oxygen atom and one from the nitrogen atom (cf. 1 and 3 versus 2) and on that basis, it is referred to here as the primary electronic delocalization of the amide function. 

The resonance hybrid B with the double bond at a bridgehead violates Bredt s rule. Therefore, without the delocalization of the nitrogen lone pair of electrons, the carbonyl moiety behaves more like a ketone than an amide. 

In the valence bond or hybridization model for CO2, we have two resonance (or canonical) structures, as shown in Fig. 3.4.7. In both structures, the two a bonds are formed by the sp hybrids on carbon with the 2pz orbitals on the oxygens. In the left resonance structure, the jv bonds are formed by the 2px orbitals on C and Oa and the 2p-y orbitals on C and O. In the other structure, the jv bonds are formed by the 2px orbitals on C and Ob and the 2py orbitals on C and Oa. The real structure is a resonance hybrid of these two extremes. In effect, once again, we get two a bonds, two jv bonds, and four lone pairs on the two oxygens. This description is in total agreement with the molecular orbital picture. The only difference is that electron delocalization in CO2 is... 

These structures are called canonical forms, and the actual resulting blended structure is called a resonance hybrid. They are usually represented as an aromatic sextet of delocalized electrons, represented by a circle within the benzene ring ... 

Benzene is actually a resonance hybrid of the two Kekule structures. This representation implies that the pi electrons are delocalized, with a bond order of 1 between adjacent carbon atoms. The carbon-carbon bond lengths in benzene are shorter than typical single-bond lengths, yet longer than typical double-bond lengths. 

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