Electron-pair splitting

Free-radical reactions were discussed in the context of combustion in Chapter 9, and the same principles apply here. Radical reactions start with the formation of the radicals by initiators such as peroxides. As shown in Figure 13.25, radicals react with monomers to create more radicals that combine with monomers. The process continues imtil the chain reaction is stopped by one of three mechanisms, all of which involve the combination of two radicals. Monomers capable of free-radical polymerization include styrene, vinyl chloride, and methyl methacrylate. Chemical initiators include hydrogen and other peroxides and potassium persulfate. Most initiators have weak O—O bonds that cleave homolytically (i.e., an electron pair splits in half) to form radicals, as in peroxides. 

The monomers are electron pair acceptors, and donor molecules are often able to split the dimeric halide molecules to form adducts thus, whilst the dimeric halides persist in solvents such as benzene, donor solvents such as pyridine and ether appear to contain monomers since adduct formation occurs. Aluminium halides, with the one exception of the fluoride, resemble the corresponding boron halides in that they are readily hydrolysed by water. 

Pure anhydrous aluminium chloride is a white solid at room temperature. It is composed of double molecules in which a chlorine atom attached to one aluminium atom donates a pair of electrons to the neighbouring aluminium atom thus giving each aluminium the electronic configuration of a noble gas. By doing so each aluminium takes up an approximately tetrahedral arrangement (p. 41). It is not surprising that electron pair donors are able to split the dimer to form adducts, and ether, for example, forms the adduct. 

To see why this splitting occurs, consider what happens when six ligands (e.g., HzO, CN-, NH3) approach a central metal cation along the x-, y-, and z-axes (Figure 15.9). The unshared electron pairs on these ligands repel the electrons in the d orbitals of the cation. 

In a covalent compound of known structure, the oxidation number of each atom is the charge remaining on the atom when each shared electron pair is assigned completely to the more electronegative of the two atoms sharing it. A pair shared by two atoms of the same element is split between them. 

Crystal field splitting energy compared to the electron pairing energy. 

Figure S6.3 Electron configurations possible for d" cations in an octahedral crystal field. For the ions d4 to d7, two configurations are possible. When the crystal field splitting is small, electrons avoid each other and produce a high-spin (HS) configuration. When the crystal field splitting is large, the electrons pair and produce a low-spin (LS) state.

In 1925, Wolfgang Pauli gave chemists what they wanted from the physicists a physical principle underlying electron-pair valency. Pauli built on the fact that in addition to the continuous, line, and band spectra, there is a fine structure of doublets, triplets, and multiple lines, some of which are split in a magnetic field (Zeeman effect). 

To transport the two electrons from NADPH to the acceptor molecule (A), the one-electron transfer reactions must proceed in two consecutive steps. These two enzymes demonstrate how nature is making use of one and the same redox system to split the incoming electron-pair into single electrons of equipotential energy to reduce a particular acceptor system. 

Anisotropy due to nuclear-electron hyperfine splitting, to spin-orbit coupling ( -factor), and particularly to strong electron-electron dipolar splitting in molecular triplets and radical pairs, provides a great deal of orientational information. It also makes it possible to shift peak positions in the spectrum, creating new windows through which to observe minor components that would be completely obscured in powder spectra. 

The electron pair that formed the carbon-carbon bond has split into two separate electrons, as indicated by the dot on the C atom in each -CH3. Because the radical has only seven electrons, it is not possible for them all to pair. 

These level splittings are the essential features that determine the type of chemical interaction, i.e., 2c-le, 2c-2e or 2c-3e bond. However, as discussed in the preceding section, there are not many examples in which these bonds occur in such a pure state. Archetypal representatives of these bonding modes (e.g.,, H2, HeH) are in fact rather atypical for chemical bonds in general. In nearly all other cases, we have to deal with mono- or polyatomic fragments, carrying additional electrons in other orbitals. These electrons will interfere with and affect the nature of the primary frontier orbital interactions. Here we focus on the nature of the electron pair bond and how this bond can be influenced by Pauli repulsion effects due to other fragment orbitals such as lone pairs.54 72... 

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