Resonance Delocalized Electron-Pair Bonding

We can often write more than one Lewis structure, each with the same relative placement of atoms, for a molecule or ion with double bonds next to single bonds. Consider ozone (O3), a serious air pollutant at ground level but a life-sustaining absorber of harmful ultraviolet (UV) radiation in the stratosphere. Two valid Lewis structures (with lettered O atoms for clarity) are 

In structure I, oxygen B has a double bond to oxygen A and a single bond to oxygen C. In structure II, the single and double bonds are reversed. These are not two different O3 molecules, just different Lewis structures for the same molecule. 

Resonance structures are not real bonding depictions O3 does not change back and forth from structure I at one instant to structure II the next. The actual molecule is a resonance hybrid, an average of the resonance forms. 

Consider these analogies. A mule is a genetic mix, a hybrid, of a horse and a donkey it is not a horse one instant and a donkey the next. Similarly, the color 

Our need for more than one Lewis structure to depict the ozone molecule is the result of electron-pair delocalization. In a single, double, or triple bond, each electron pair is attracted by the nuclei of the two bonded atoms, and the electron density is greatest in the region between the nuclei each electron pair is localized. In the resonance hybrid for O3, however, two of the electron pairs (one bonding and one lone pair) are delocalized their density is spread over the entire molecule. In O3, this results in two identical bonds, each consisting of a single bond (the localized electron pair) and a partial bond (the contribution from one of the delocalized electron pairs). We draw the resonance hybrid with a curved dashed line to show the delocalized pairs  

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The organic chemist made an important step in the understanding of chemical reactivity when he realized the importance of electronic stabilization caused by the delocalization of electron pairs (bonded and non-bonded) in organic molecules. Indeed, this concept led to the development of the resonance theory for conjugated molecules and has provided a rational for the understanding of chemical reactivity (1, 2, 3). The use of "curved arrows" developed 50 years ago is still a very convenient way to express either the electronic delocalization in resonance structures or the electronic "displacement" occurring in a particular reaction mechanism. This is shown by the following examples. 

In other words, the unshared electron pair of the base, acetate ion, is delocalized (spread over both oxygens) by resonance. This electron pair is stabilized and less available for bonding to the proton, which localizes this electron pair in the sigma bond and costs resonance energy. The most common effect of resonance on an acid-base reaction is to delocalize and stabilize the unshared electron pair of the conjugate base, resulting in a stronger acid. 

The second and often more powerful influence is the resonance or delocalization effect. For any molecule for which more than one electron-pair bond structure can be written, the true structure will be intermediate between these structures and the molecule will be more stable than expected. Well-known examples include benzene (Figure 3.3a) and charged structures such as the allyl cation (Figure 3.3b). A carbocation can be stabilized either by an adjacent lone pair, which can be donated to form a double bond (Figure 3.3c) or by donation of electrons from an adjacent multiple bond (Figure 3.3b). Anions can be stabilized by donation of the lone pair to an adjacent multiply bonded atom (Figure 3.3d). 

The organic chemist made an important step in the understanding of chemical reactivity when he realized the importance of electronic stabilization caused by the delocalization of electron pairs (bonded and non-bonded) in organic molecules. Indeed, this concept led to the development of the resonance... 

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,... 

If one wishes to use only Heitler-London wave-functions for all electron-pair bonds, as indicated above when considering Coulson-Fischer A-B orbitals, we can still speak of the delocalization of a Y electron of structure (4) into the antibonding A-B orbital. When this is done, we obtain Lewis structures (14) and (15), with configurations (y) (a) (b) and (y) (a) (b) that involve Heitler-London formulations of the wave-functions for the Y-A and long Y-B bonds. Resonance between structures (14) and (15) is equivalent to the utilization of increased-valence structure (1) (with a Heitler-London type wave-function for the fractional Y-A bond). 

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-. 

Resonance structures result from a phenomenon known as electron delocalization. The electron pairs in the three double bonds in a benzene ring are delocalized. These are electrons that belong to no particular atom or bond. As a consequence, no ordinary double bonds exist in a benzene ring. The electrons are in an orbital that extends across adjacent atoms. This smear of electrons is usually represented as a circle within the ring. 

In practice, the NBO program labels an electron pair as a lone pair (LP) on center B whenever cb 2 > 0.95, i.e., when more than 95% of the electron density is concentrated on B, with only a weak (<5%) delocalization tail on A. Although this numerical threshold produces an apparent discontinuity in program output for the best single NBO Lewis structure, the multi-resonance NRT description depicts smooth variations of bond order from uF(lon) = 1 (pure ionic one-center) to bu 10n) = 0 (covalent two-center). This properly reflects the fact that the ionic-covalent transition is physically a smooth, continuous variation of electron-density distribution, rather than abrupt hopping from one distinct bond type to another. 

The t3C resonances of xanthone (21) have been assigned on the basis of shift analysis of the structurally related compounds anthrone and xanthene (77P735) and by correlation with the component molecules diphenyl ether and benzophenone (78T1837). Delocalization of the non-bonding electron pairs at oxygen, leading to an increase in aromaticity of the pyranone ring, accounts for the deviation of the chemical shifts of C-4a, C-8a and the carbonyl carbon atom from the calculated values. 

The anomeric effect in terms of a stabilizing effect can be illustrated by the concept of "double-bond - no-bond resonance" (14, 15) shown by the resonance structures 4 and 2 or by the equivalent modern view (16, 17) that this electronic delocalization is due to the overlap of an electron pair orbital of an oxygen atom with the antibonding orbital of a C —OR sigma bond (12). 

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... 

An electron-donor molecule possesses a high electron density due to the presence of delocalized electrons or non-bonding electrons, such as those associated with nitrogen, oxygen, and sulfur. An electron-acceptor molecule, on the other hand, is able to accept electron density and stabilize the additional charge through resonance or induction. According to the polymerizability of the components, photo-induced reactions of donor/acceptor pairs are divided into three types both are non-polymerizable (Type I), one is polymerizable (Type II), or both are polymerizable (Type III). 

Which is the stronger base, aniline or ammonia The same reasoning applies. Aniline is stabilized by resonance. The basic pair of electrons is delocalized into the ring and is less available to form a covalent bond, so it is a weaker base. (Of course, if the conjugate acid of aniline is stronger than ammonium ion, then aniline must be a weaker... 

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. 

The structural formulas used to represent molecules are based on valence bond theory. Double and triple bonds simply represent additional pairs of shared valence electrons. But structural formulas, while useful, don t tell the whole story about the nature of the bonds between atoms in a molecule. Valence bond theory falls flat when it tries to explain delocalized electrons and resonance structures. To get at what is really going on inside molecules, chemists had to dig deeper. 

Each resonance structure implies that electron pairs are localized in bonds or on atoms. In actuality, resonance allows certain electron pairs to be delocalized over two or more atoms, and this delocalization of electron density adds stability. A molecule with two or more resonance structures is said to be resonance stabilized. We will return to the resonance hybrid in Section 1.5C. 

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