Orbital Overlap and Chemical Bonding

To explain how two waves—or two orbitals—might interact with one another, we will need to invoke the concept of interference, which provides a clear contrast between waves and particles. When two particles occupy the same space, they collide, but when two waves occupy the same space, they interfere. As you may have learned in a high school physics or physical science class, this 

The name valence bond model arises from the fact that only valence electrons are considered. 

In the quantum mechanical model, electrons are simply one type of wave, and we can picture the formation of chemical bonds as an example of constructive interference between electron waves. But for electron waves from different atoms to interact, they must first occupy the same region of space. Another way of saying this is that the atoms must be positioned so that the valence orbitals from one atom must overlap those of the other atom if a bond is to form. This idea forms the basis for the valence bond model of chemical bonding, in which all bonds are seen as the result of overlap between atomic orbitals. Let s examine this idea of orbital overlap by looking at the simplest possible molecule, H2. 

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Chapters 4 and 5 emphasize the band gap in materials. This is a critical parameter of a material determining its utility in transistors, lasers, and detectors. Until recently, altering a material s band gap involved chemical modifications to affect bond lengths, atomic orbital overlaps, and electronegativity. A nice example of this is the series GaP cAsi f where the band gap is systematically varied from 2.27 to 1.40 eV as x varies from 1 to 0 (Fig. 12.15). This represents a traditional chemical approach to achieving a desired property. A more recent approach to preparing materials with specific band gaps involves what have become known as photonic materials. 

Silicon and germanium crystallize in the diamond structure. However, they have somewhat weaker covalent bonds than carbon as a consequence of less efficient orbital overlap. These weaker bonds result in lower melting points for silicon (1420°C for Si and 945°C for Ge, compared with 4100°C for diamond) and greater chemical reactivity. Both sihcon and germanium are semiconductors, described in Chapter 7. 

In the previous paper (13), we have discussed the chemical bonding nature of uranyl nitrate dihydrate and found that the bonding interaction is mainly due to the U 5f, 6d - O 2p components. In the present work, we carry out orbital overlap population analysis to understand contribution of each atomic orbital to the chemical bonding. The orbital overlap populations indicate strength of covalent bonds (19,20). 

Indirect evidence for small redox-dependent differences has been presented based on titration studies of the S -phosphate for NMN and NMNH IT). Such determinations, however, are highly suspect because the effects of titration of the 5 -phosphate are not limited to through-space interactions. Chemical shifts also depend on orbital overlap ofdie intervening bonds and localized bond polarization, neither of which is a direct function of the interatomic distances 4). 

The most important orbitals are those in the outer shells, which are involved in the formation of chemical bonds. Covalent bonds are formed when atomic orbitals overlap and merge to form molecular orbitals (Chapter 14). 

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