Lewis structure electron-pair delocalization

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 ... 

These considerations imply that we may generate an "increased-valence" structure (V) from the standard Lewis structure (VI) by delocalizing one electron of the lone-pair of atom C into a vacant bonding BC orbital... 

Electron Delocalization Our need for more than one Lewis structure to depict O3 is due to electron-pair delocalization. In a single, double, or triple bond, each electron pair is localized between the bonded atoms. In a resonance hybrid, two of the electron pairs (one bonding and one lone pair) are delocalized their density is spread over a few adjacent atoms. 

In molecular orbital theory, electrons occupy orbitals called molecular orbitals that spread throughout the entire molecule. In other words, whereas in the Lewis and valence-bond models of molecular structure the electrons are localized on atoms or between pairs of atoms, in molecular orbital theory all valence electrons are delocalized over the whole molecule, not confined to individual bonds. 

The boranes are electron-deficient compounds (Section 3.8) we cannot write valid Lewis structures for them, because too few electrons are available. For instance, there are 8 atoms in diborane, so we need at least 7 bonds however, there are only 12 valence electrons, and so we can form at most 6 electron-pair bonds. In molecular orbital theory, these electron pairs are regarded as delocalized over the entire molecule, and their bonding power is shared by several atoms. In diborane, for instance, a single electron pair is delocalized over a B—H—B unit. It binds all three atoms together with bond order of 4 for each of the B—H bridging bonds. The molecule has two such bridging three-center bonds (9). 

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. 

SAMPLE SOLUTION (a) When using curved arrows to represent the delocalization of electrons, begin at a site of high electron density, preferably an atom that is negatively charged. Move electron pairs until a proper Lewis structure results. For nitrate ion, this can be accomplished in two ways ... 

Electrons involved in resonance structures are said to be delocalized. Delocalization means that the sharing of an electron pair is distributed over several pairs of atoms and cannot be identified with just one pair of atoms. Because we cannot draw one Lewis structure that portrays this delocalization, we depict the structure of a molecule that undergoes resonance by drawing all the contributing Lewis structures. 

Although the resonance structures of benzene show it as a cyclo-hexatriene, because of its fully delocalized n system and the closed shell nature of this n system, benzene does not undergo addition reactions like ordinary unsaturated compounds. The destruction of the n electron system during addition reactions would make the products less stable than the starting benzene molecule. However, benzene does undergo substitution reactions in which the fully delocalized closed n electron system remains intact. For example, benzene may be reacted with a halogen in the presence of a Lewis acid (a compound capable of accepting an electron pair) to form a molecule of halobenzene. 

To get an accurate picture for the formate anion, it is necessary to use delocalized MOs that involve all three of the overlapping p orbitals (one from each O and one from C). These three AOs overlap to form three pi MOs.There are four electrons (shown in the Lewis structures in parts 0 and as the pi electrons and one unshared pair on oxygen) in these three MOs.This drawing does not attempt to show the shapes of these three MOs, only the orbitals that overlap to form them. More on their energies and shapes will be presented in later chapters. 

Species requiring multiple Lewis structures often involve electron delocalization with some electron pairs distributed over more than two atoms. 

A satisfactory localized view of ferrocene can be developed by noting from the Lewis structure that there are nine electron pairs around iron three from each cyclopentadi-enide ring and three from the metal atom. From our delocalized MO picture we note that the occupied orbitals in Fig. 21 were generated from a set of GO s composed of an s, all three p s and all of the d s. These GO s were seen to be of the same type as die valence shell AO s actually employed by the iron atom in forming MO s. The Lewis structure suggests that we should localize three bonding electron pairs between each ring and the metal atom. To accomplish this we first hybridize the GO s pO, pi, pi", dO, dl, dl" associated with the lowest six delocalized MO s in Fig. 21 which yields... 

The total bond energy of a substance for which resonance structures are written is greater than would be expected if there were only one formal Lewis structure. This additional stabilization is called resonance energy. It arises from the same principle that is responsible for covalent bond energy, the delocalization of electrons about the atoms forming the bond. As a result of resonance in ozone, for example, the electrons constituting the second pair of the double bond are delocalized around the 3... 

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