Lewis theory half bond

Most chemists still tend to think about the structure and reactivity of atomic and molecular species in qualitative terms that are related to electron pairs and to unpaired electrons. Concepts utilizing these terms such as, for example, the Lewis theory of valence, have had and still have a considerable impact on many areas of chemistry. They are particularly useful when it is necessary to highlight the qualitative similarities between the structure and reactivity of molecules containing identical functional groups, or within a homologous series. Many organic chemistry textbooks continue to use full and half-arrows to indicate the supposed movement of electron pairs or single electrons in the description of reaction mechanisms. Such concepts are closely related to classical valence-bond (VB) theory which, however, is unable to compete with advanced molecular orbital (MO) approaches in the accurate calculation of the quantitative features of the potential surface associated with a chemical reaction. 

The valence bond (VB) theory is an extension of Lewis s concept of electron sharing that originated with the work of Walter Heitler" and Fritz London in 1927. To begin our discussion of VB theory, let s consider the formation of an H2 molecule from two H atoms, as discussed on p. 174. The Lewis theory desaibes the H—H bond in terms of the pairing of the two electrons on the H atoms. In VB theory, the covalent H—H bond is formed by the overlap of the two half-filled F orbitals in the H atoms. By overlap, we mean that the two orbitals share a common region in space (Figure 3.4). 

A stable molecular species has more electrons in bonding orbitals than in antibonding orbitals. For example, if the excess of bonding over antibonding electrons is tzvo, this corresponds to a single covalent bond in Lewis theory. In molecular orbital theory, we say that the bond order is 1. Bond order is one-half the difference between the number (no.) of bonding and antibonding electrons (e ), that is. 

H2 This species has a single electron. It enters the o-js orbital, a bonding molecular orbital. Using equation (11.1), we see that the bond order is (1 — 0)/2 = This is equivalent to a one-electron, or half, bond, a bond type that is not easily described by the Lewis theory. 

Molecular Properties. Corresponding to its nuclear charge number, the nitrogen atom possesses seven shell electrons. One electron pair in in the ground state 1 s(K shell), and five electrons are distributed over the four orbitals with the principal quantum number 2 (L shell). Of these, one electron pair occupies the 2 s level and three unpaired electrons, respectively, a half of the remaining three levels, 2 px, 2 pv, 2 pz. The unpaired electrons can enter into electron-pair bonds with the 1 s electron of three hydrogen atoms. Thus, the three half occupied orbitals of the L shell become about fully accupied (formation of an octet of the neon type in accordance with the octet theory of Lewis-Langmuir). 

Lewis published these ideas in his 1923 book Valence and the Structure of Atoms and Molecules, and they were widely taken up and developed in the U.S.A. and Europe, for example, by N. V. Sidgwick at Oxford, whose Electronic Theory of Valency appeared in 1927. The Nobel Prize in Chemistry was left unfilled in 1919, 1924 and 1933 for lack of candidates of suitable stature, and Lewis would have been an appropriate candidate for any of these years. In fact, he was nominated for a Nobel Prize by the inorganic chemist and historian of chemistry, J. R. Partington (1886-1965) at the University of London. For the first half-century after the award of the first Nobel Prize in Chemistry to van t Hoff in 1901, the chemistry prize went to those who had discovered or characterised new chemical elements, new physico-chemical principles, new chemical reactions, or had elucidated the structure and accomplished the synthesis of natural products. The first award for research into the nature of the chemical bond and its application to the elucidation of the structure of complex substances went in 1954 to Linus Pauling at Caltech. 

The Lewis definitions of acid-base interactions are now over a half a century old. Nevertheless they are always useful and have broadened their meaning and applications, covering concepts such as bond-formation, central atom-ligand interactions, electrophilic-nucleophilic reagents, cationic-anionic reagents, charge transfer complex formation, donor-acceptor reactions, etc. In 1923 Lewis reviewed and extensively elaborated the theory of the electron-pair bond, which he had first proposed in 1916. In this small volume which had since become a classic, Lewis independently proposed both the proton and generalized solvent-system definitions of acids and bases. He wrote ... 

Bond order in a Lewis formula, the number of pairs of electrons in a bond. (9.10) In molecular orbital theory, one-half the difference between the number of bonding electrons and the number of antibonding electrons. (10.5)... 

We have learned that the Lewis model portrays a chemical bond as the transfer or sharing of electrons represented as dots. Valence bond theory portrays a chemical bond as the overlap of two half-filled atomic orbitals. What is a chemical bond according to molecular orbital theory ... 

In contrast to the Lewis model, in which a covalent chemical bond is the sharing of electrons represented by dots, in valence bond theory a chemical bond is the overlap of half-filled atomic orbitals (or in some cases the overlap between a completely filled orbital and an empty one). 

The concept of a chemical bond as a localized interaction between two neighboring atoms has been a central part of chemistry for the past century and a half, yet our current description of chemical bonds is still empirical it is a collage of ill-defined and largely incompatible models that are based on assumptions that do not always correspond to physical reality. The ionic and covalent models are mutually incompatible, and both the Lewis and orbital models have serious flaws [3, 4]. They do not conform to modem views of atomic stmcture, and consequently their predictions sometimes fail. While the bond valence theory belongs to this tradition of localized bond models, it is derived from a realistic, though simplified picture of the atom, one that is compatible with more sophisticated atomic descriptions. It can be used to derive powerful and quantitative theorems about chemical stracture. The mles of both the traditional ionic and covalent models can be derived as two special cases of this model (Sects. 5 and 7.2). 

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