Delocalized and Localized Lone Pairs

35 Use resonance structures to help you identify all sites of high electron density (8—) in the following compound  

In this section, we will explore some important differences between lone pairs that participate in resonance and lone pairs that do not participate in resonance. 

Recall that one of our five patterns was a lone pair that is allylic to a it bond. Such a lone pair will participate in resonance and is said to be delocalized. When an atom possesses a delocalized lone pair, the geometry of that atom is affected by the presence of the lone pair. As an example, consider the strucmre of an amide  

The rules we learned in Section 1.10 would suggest that the nitrc en atom should be hybridized and trigonal pyramidal, but this is not correct. Instead, the nitrogen atom is actually sp hybridized and trigonal planar. Why The lone pair is participating in resonance and is therefore delocalized  

Whenever a lone pair participates in resonance, it will occupy a p orbital rather than a hybridized orbital, and this must be taken into account when predicting geometry. This will be extremely important in Chapter 25 when we discuss the three-dimensional shape of proteins. 

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Molecular Orbital Theory (Section 1.8) Delocalized and Localized Lone Pairs (Section 2.12)... 

Pyridine, symmetry group C2v, has six electrons in a system delocalized around the ring, and two lone-pair electrons in an orbital localized at the Nitrogen atom. The Is electrons, as well as the electrons in orbitals describing the a bonds, need not be considered explicitly in describing the resonance stabilization and low-lying excited states of pyridine. The simple molecular orbital description has the following characteristic assumptions ... 

This molecular orbital formation moves electrons localized on oxygen into orbitals shared between carbon and oxygen. We can represent this in curly arrow terms as a delocalization of the lone pair electrons. 

Rather similar results are obtained by comparing the bond angles in the silyl and methyl ethers (Fig. 18.5) and isothiocyanates (Fig. 18.6). In dimethyl ether the oxygen is hybridized approximately sp with two lone pairs on the oxygen atom as compared to an approximate sp hybrid in disiloxane with ir bonding. In the same way the methyl isothiocyanate molecule, CHjN=C=S, has a lone pair localized on the nitrogen atom, hence is bent (N sp ), but the delocalization of this lone pair into a ir orbital on the silicon atom of H3SiN=C=S leads to a linear structure for this molecule. 

Resonance attempts to correct a fundamental defect in Lewis formulas. Lewis formulas show electrons as being localized they either are shared between two atoms in a covalent bond or are unshared electrons belonging to a single atom. In reality, electrons distribute themselves in the way that leads to their most stable arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei. What we try to show by the resonance description of ozone is the delocalization of the lone-pair electrons of one oxygen and the electrons in the double bond over the three atoms of the molecule. Organic chemists often use curved arrows to show this electron... 

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 ... 

Because the most common model for bonding considers localized a and tt bonds between adjacent atoms, we need a different approach to modeling the bonding and antibonding orbitals when the bonds are not localized. This approach wiU be true for any system where we can generate contributing structures that delocalize tt bonds and/or lone pairs over three or more atoms. To picture the orbitals involved in delocalized systems, we use only MO theory, dropping the VB theory localization principle. 

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. 

The diethylamino group exhibits a localized lone pair, and as such, it is expected to exist primarily as a charged ammonium ion 10 see pA a table on inside cover of textbook) at physiological pH. In contrast, the other nitrogen atom exhibits a highly delocalized lone pair, and it is not expected to be protonated at physiological pH (see discussion of the Henderson-Hasselbalch equation in Section 21.3). 

Now let s search for a nucleophUic center in the other starting material. There are two amino groups in the other starting material however, one of them exhibits a lone pair that is delocalized (via resonance) into the aromatic ring. Since the lone pair of this amino group is delocalized, it is expected to be less reactive as a nucleophile. The other amino group has a localized lone pair, and is a much better nucleophile. 

In standard quantum-mechanical molecular structure calculations, we normally work with a set of nuclear-centred atomic orbitals Xi< Xi CTOs are a good choice for the if only because of the ease of integral evaluation. Procedures such as HF-LCAO then express the molecular electronic wavefunction in terms of these basis functions and at first sight the resulting HF-LCAO orbitals are delocalized over regions of molecules. It is often thought desirable to have a simple ab initio method that can correlate with chemical concepts such as bonds, lone pairs and inner shells. A theorem due to Fock (1930) enables one to transform the HF-LCAOs into localized orbitals that often have the desired spatial properties. 

Figure 24.3 Arylamines have a larger positive AG for protonation and are therefore less basic than alkylarnines, primarily because of resonance stabilization of the ground state. Electrostatic potential maps show that lone-pair electron density is delocalized in the amine but the charge is localized in the corresponding ammonium ion.

Overall, it can be seen from Figs. 4 and 5 that in predicting the p parameters almost identical results are obtained in five out of the eight cases. Significantly perhaps, these include the more highly localized O- lone pair and C H orbitals. The worst agreement in both molecules is seen for the more delocalized ring 7t orbital. 

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