Lewis’s Theory

The Lewis structures encountered in Chapter 2 are two-dimensional representations of the links between atoms—their connectivity—and except in the simplest cases do not depict the arrangement of atoms in space. The valence-shell electron-pair repulsion model (VSEPR model) extends Lewis s theory of bonding to account for molecular shapes by adding rules that account for bond angles. The model starts from the idea that because electrons repel one another, the shapes of simple molecules correspond to arrangements in which pairs of bonding electrons lie as far apart as possible. Specifically ... 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

Lewis s theory also fails to account for the compound diborane, B2H6, a colorless gas that bursts into flame on contact with air. The problem is that diborane has only 12 valence electrons (three from each B atom, one from each H atom) but, for a Lewis structure, it needs at least seven bonds, and therefore 14 electrons, to bind the eight atoms together Diborane is an example of an electron-deficient compound, a compound with too few valence electrons to be assigned a valid Lewis structure. Valence-bond theory can account for the structures of electron-deficient compounds in terms of resonance, but the explanation is not straightforward. 

Unlike Lewis s theory, molecular orbital theory can account for the existence of electron-deficient compounds and the paramagnetism of oxygen. 

The application of the quantum mechanics to the interaction of more complicated atoms, and to the non-polar chemical bond in general, is now being made (45). A discussion of this work can not be given here it is, however, worthy of mention that qualitative conclusions have been drawn which are completely equivalent to G. N. Lewis s theory of the shared electron pair. The further results which have so far been obtained are promising and we may look forward with some confidence to the future explanation of chemical valence in general in terms of the Pauli exclusion principle and the Heisenberg-Dirac resonance phenomenon. 

Lewis s theory and Heitler and London s extension permitted the reasonably certain attribution of specific electronic formulas to a great many compounds. In other cases, however, it was possible to set up a number of alternative electronic formulas for a molecule or crystal, and often no sound argument could be advanced supporting any one of them against the others. For example, Lewis gave the perchlorate ion the... 

Lewis s theory of electron valence, 136-137, 154, 186, 190 opposition to permanent polar valences, 137 n.137 inductive effect, 208 static electron atom, 241 ... 

What he does not seem to realize is that a perfectly good explanation existed for chemical bonding prior to the advent of the quantum mechanical explanation, namely Lewis s theory whereby pairs of electrons form the bonds between the various atoms in a covalently bonded molecule. Although the quantum mechanical theory provides a more fundamental explanation in terms of exchange energy and so on is undeniable but it also retains the notion of pairs of electrons even if this notion is now augmented by the view that electrons have anti-parallel spins within such pairs. 

There is an implicit assumption contained in all of the above The two bonding electrons are of opposite spin. If two electrons are of parallel spin, no bonding occurs, but repulsion instead curve /, Fig. 5.1). This is a result of the Pauli exclusion principle. Because of the necessity for pairing in each bond formed, the valence bond theory is often referred to as the electron pair theory, and it forms a logical quantum-mechanical extension of Lewis s theory of electron pair formation. 

Electrophiles (i.e., electron-deficient species) are of fundamental importance to chemistry. The concept of nucleophiles (lit. nucleus seeking ) and electrophiles (lit. electron seeking ) was suggested by Ingold following similar views implied by Lapworth s description of anionoid and cationoid reagents, Robinson s concepts, and Lewis s theory of bases (electron donors) and acids (electron acceptors).1... 

The Limitations of Lewis s Theory Valence Bond (VB) deficiencies... 

In trying to explain why atoms form bonds, G. N. Lewis proposed that an atom is most stable if its outer shell is either filled or contains eight electrons and it has no electrons of higher energy. According to Lewis s theory, an atom will give up, accept, or share electrons in order to achieve a filled outer shell or an outer shell that contains eight electrons. This theory has come to be called the octet rule. 

The simplicity of the VSEPR model is one of its primary strengths. In addition, the model provides a continuity in the development of the qualitative ideas about the nature of the chemical bond and its correlation with molecular structure. Abegg s octet rule (see, e.g.. Ref. [3-71]) and Lewis s theory of the shared electron pair [3-72] may be considered as direct forerunners of the model. 

Lewis s cubical atom [3-72] deserves special mention. It was instrumental in shaping the concept of the shared electron pair. It also permitted a resolution of the apparent contradiction between the two distinctly different bonding types, viz., the shared electron pair and the ionic electron-transfer bond. In terms of Lewis s theory, the two bonding types could be looked at as mere limiting cases. Lewis s cubical atoms are illustrated in Figure 3-51. They are also noteworthy as an example of a certainly useful though not necessarily correct application of a polyhedral model. 

Lewis s theory explains the relations between ordinary valency, electrovalency, and covalency found in Chapter 8 (e.g. V= e = C = 1 for hydrogen). 

G.N. Lewis to R. MiUikan, October 28, 1919, Lewis Archive, University of California, Berkeley, as cited in R.E. Kohler, The Origin of G.N. Lewis s Theory of the Shared Pair... 

G. N. Lewis developed a model for chemical bonding that you have learned in this chapter. His theory was extremely successful and is used today at all levels of chemistry, from the introductory class to the research laboratory. Why was Lewis s theory so successful ... 

Valence bond theory predates molecular orbital theory, being an extension of G. N. Lewis s theory of bonding. It had the early advantage of a conceptually simple, yet accurate depiction of electron distributions in chemical bonds. It has subsequently been eclipsed by MO theory, which shares its conceptual simplicity but is much easier to formulate mathematically. 

In Lewis s theory, component (C) was emphasized relative to components (A) and (B). In Heitler and London s spin theories of valence, component (C) had priority over... 

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