Periodic table, bond energies

Consequently, it must be emphasized that precautions have to be taken with the conventional rough description of molecules based on the chemical bond pattern. In a molecule that contains at least two atoms which do not belong to the first row of the periodic table, the energy and all the monoelectronic properties are literally spread out over the whole molecule. Obviously, the concept of chemical bond, based as it is on the principle of topological proximity, is inadequate on its own for a correct description of the chemical and physical behavior of such a molecule. 

Because bond length increases down a column of the periodic table, bond dissociation energies are a quantitative measure of the general phenomenon noted in Chapter 1—shorter bonds are stronger bonds. 

Semiconductors can be divided into two classes, elemental semiconductors, which contain only one type of atom, and compound semiconductors, which contain two or more elements. The elemental semiconductors all come from group 4A. As we move down the periodic table, bond distances increase, which decreases orbital overlap. This decrease in overlap reduces the energy difference between the top of the valence band and the bottom of the conduction band. As a result, the band gap decreases on going from diamond (5.5 eV, an insulator) to silicon (1.11 eV) to germanium (0.67 eV) to gray tin (0.08 eV). In the heaviest group 4A element, lead, the band gap collapses altogether. As a result, lead has the structure and properties of a metal. 

The one usually chosen is one of the five 3d orbitals (Figure 8.7). These normally become occupied only when we reach the next row of the Periodic Table. The energy needed to promote the electron is then recouped by forming five bonds instead of three. 

These data also show that the bond energies decrease as the atomic number of the metal increases in the same Group of the Periodic Table. 

Consider the fluorides of the second-row elements. There is a continuous change in ionic character of the bonds fluorine forms with the elements F, O, N, C, B, Be, and Li. The ionic character increases as the difference in ionization energies increases (see Table 16-11). This ionic character results in an electric dipole in each bond. The molecular dipole will be determined by the sum of the dipoles of all of the bonds, taking into account the geometry of the molecule. Since the properties of the molecule are strongly influenced by the molecular dipole, we shall investigate how it is determined by the molecular architecture and the ionic character of the individual bonds. For this study we shall begin at the left side of the periodic table. 

Elements beyond the second row of the periodic table can form bonds to more than four ligands and can be associated with more than an octet of electrons. These features are possible for two reasons. First, elements with > 2 have atomic radii that are large enough to bond to 5, 6, or even more ligands. Second, elements with > 2 have d orbitals whose energies are close to the energies of the valence p orbitals. An orbital overlap description of the bonding in these species relies on the participation of d orbitals of the inner atom. 

The data show that bond energies for these three diatomic molecules increase moving across the second row of the periodic table. We must construct molecular orbital diagrams for the three molecules and use the results to interpret the trend. 

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