Electron distribution, in covalent

Structural formulas, such as shown in Fig. 9-1. represent the valence electron distributions in covalent molecules and ions. These structures are not meant to indicate actual bond angles in three-dimensional varieties such as CH3CI, NH3, and NH4 they merely show the number of bonds connecting... 

Updated and new Problem-Solving Tips enhance student use of versatile, effective strategies for writing and understanding Lewis formulas and electron distributions in covalent molecules and ions. 

The unequal distribution of electron density in covalent bonds produces a bond dipole, the magnitude of which is expressed by the dipole moment, having the units of charge times distance. Bonds with significant bond dipoles are described as being polar. The bond and group dipole moments of some typical substituents are shown in Table 1.7. 

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the valence bond model, which uses hybrid orbitals to account for the observed shapes of organic molecules. Before going on to a systematic study of organic chemistry, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the electrons in a bond are not shared equally between atoms. 

This description of a covalent bond in terms of overlapping atomic orbitals stems from the desire to retain and use the atomic orbital concept. Atomic orbitals were designed to describe electron distributions in isolated atoms however, one might expect distortions of the isolated atom electron cloud when it is in close proximity to the positive charge of another atom. A more detailed description of the covalent bond would have to include this distortion. 

The bond in sodium chloride, for instance, is largely ionic. Sodium has transferred an electron to chlorine to give Na" and Cl" ions, which are held together in the solid by electrostatic attraction. The C-C bond in ethane, however, is fully covalent. The two bonding electrons are shared equally by the two equivalent carbon atoms, resulting in a symmetrical electron distribution in the bond. Between these two extremes lie the great majority of chemical bonds, in which the electrons are attracted somewhat more strongly by one atom than by the other. We call such bonds, in which the electron distribution is unsymmetrical, polar covalent bonds. 

In chemical usage [23, Section 14.11] an electron is said to be delocalized if its molecular orbital cannot be ascribed to a two-center bond otherwise it is localized. It is, however, always possible, but perhaps rarely convenient, to describe the electron distribution in a molecule with delocalized orbitals only. The situation in a covalent insulator such as diamond is similar to the molecular case. There are four valence electrons per atom, and four neighbors. Therefore, it is possible to describe the structure with four two-center, two-electron bonds, and localized Wannier orbitals. But keep in mind that the only physical reality is the resulting charge distribution. This reality can also be described by freely moving Bloch electrons. 

Soon after the development of the quantum mechanical model of the atom, physicists such as John H. van Vleck (1928) began to investigate a wave-mechanical concept of the chemical bond. The electronic theories of valency, polarity, quantum numbers, and electron distributions in atoms were described, and the valence bond approximation, which depicts covalent bonding in molecules, was built upon these principles. In 1939, Linus Pauling s Nature of the Chemical Bond offered valence bond theory (VBT) as a plausible explanation for bonding in transition metal complexes. His application of VBT to transition metal complexes was supported by Bjerrum s work on stability that suggested electrostatics alone could not account for all bonding characteristics. 

Electron counting (by any method) does not imply anything about the degree of covalent or ionic bonding it is strictly a bookkeeping procedure, as are the metal oxidation numbers that may be used in the counting. Physical measurements are necessary to provide evidence about the actual electron distribution in molecules. Linear and cyclic organic n systems interact with metals in more complicated ways, as discussed in Chapter 5. 

---