Covalent bonds structures Localized electron

Since covalent bonding is localized, and forms open crystal structures (diamond, zincblende, wurtzite, and the like) dislocation mobility is very different than in pure metals. In these crystals, discrete electron-pair bonds must be disrupted in order for dislocations to move. 

Aromatic compound (Section 11.3) An electron-delocalized species that is much more stable than any structure written for it in which all the electrons are localized either in covalent bonds or as unshared electron pairs. 

For each molecule, ion, or free radical that has only localized electrons, it is possible to draw an electronic formula, called a Lewis structure, that shows the location of these electrons. Only the valence electrons are shown. Valence electrons may be found in covalent bonds connecting two atoms or they may be unshared. The student must be able to draw these structures correctly, since the position of electrons changes in the course of a reaction, and it is necessary to know where the electrons are initially before one can follow where they are going. To this end, the following rules operate ... 

In this contribution it is shown that local density functional (LDF) theory accurately predicts structural and electronic properties of metallic systems (such as W and its (001) surface) and covalently bonded systems (such as graphite and the ethylene and fluorine molecules). Furthermore, electron density related quantities such as the spin density compare excellently with experiment as illustrated for the di-phenyl-picryl-hydrazyl (DPPH) radical. Finally, the capabilities of this approach are demonstrated for the bonding of Cu and Ag on a Si(lll) surface as related to their catalytic activities. Thus, LDF theory provides a unified approach to the electronic structures of metals, covalendy bonded molecules, as well as semiconductor surfaces. 

It is difficult to give a localized orbital description of the bonding in a period 3 hypervalent molecule that is based only on the central atom 3s and 3p orbitals and the ligand orbitals, that is, a description that is consistent with the octet rule. One attempt to do this postulated a new type of bond called a three-center, four-electron (3c,4e) bond. We discuss this type of bond in Box 9.2, where we show that it is not a particularly useful concept. Pauling introduced another way to describe the bonding in these molecules, namely, in terms of resonance structures such as 3 and 4 in which there are only four covalent bonds. The implication of this description is that since there are only four cova-... 

Soon after the quantum revolution of the mid 1920s, Linus Pauling and John C. Slater expanded Lewis s localized electronic-structural concepts with the introduction of directed covalency in which bond directionality was achieved by the hybridization of atomic orbitals.1 For normal and hypovalent molecules, Pauling and Slater proposed that sp" hybrid orbitals are involved in forming shared-electron-pair bonds. Time has proven this proposal to be remarkably robust, as has been demonstrated by many examples in Chapter 3. 

It seems to be realistic to relate catalytic activity to the most stable [111] plane of fee metals. Bond (135) describes the electron structure of the this plane. So-called 2g electron orbitals point toward those interstices where metal atoms in the subsequent overlayer would be accommodated. These orbitals have metallic character. So-called orbitals point toward the next nearest neighbor. These are localized and able to form real covalent bonds. The degree of hybridization of these orbitals is imknown. Knor (136) assumes that only orbitals would stick out of the plane, but they are almost completely hybridized. He assumes that the /2g electrons are parts of the electron gas of the metal. The and sites are by no means equivalent. 

Two later sections (1.6.5 and 1.6.6) look at the crystalline structures of covalently bonded species. First, extended covalent arrays are investigated, such as the structure of diamond—one of the forms of elemental carbon—where each atom forms strong covalent bonds to the surrounding atoms, forming an infinite three-dimensional network of localized bonds throughout the crystal. Second, we look at molecular crystals, which are formed from small, individual, covalently-bonded molecules. These molecules are held together in the crystal by weak forces known collectively as van der Waals forces. These forces arise due to interactions between dipole moments in the molecules. Molecules that possess a permanent dipole can interact with one another (dipole-dipole interaction) and with ions (charge-dipole interaction). Molecules that do not possess a dipole also interact with each other because transient dipoles arise due to the movement of electrons, and these in turn induce dipoles in adjacent molecules. The net result is a weak attractive force known as the London dispersion force, which falls off very quickly with distance. 

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