Lewis electron pair bonding

An advantage of VSEPR is its foundation upon Lewis electron-pair bond theory. No mention need be made of orbitals and overlap. If you can write down a Lewis structure for the molecule or polyatomic ion in question, with all valence electrons accounted for in bonding or nonbonding pairs, there should be no difficulty in arriving at the VSEPR prediction of its likely shape. Even when there may be some ambiguity as to the most appropriate Lewis structure, the VSEPR approach leads to the same result. For example, the molecule HIO, could be rendered, in terms of Lewis theory as ... 

The heavier elements of the group, S, Se and Te all form tetralluorides, EF4 and hexafluorides, EF6- If the central atom in these compounds form Lewis electron pair bonds to all the ligating fluorine atoms, it must accommodate five or six electron pairs in the valence shell, and we refer to the atoms as hypervalent. Sulfur forms no further homoleptic hypervalent derivatives. Selenium forms a solid tetrachloride, but as mentioned in the last paragraph, it decomposes on evaporation. Tellurium forms a solid tetrachloride, which may be evaporated without decomposition, as well as tetraphenyl-, tetramethyl- and hexamethyl- derivatives. 

He, somewhat mischievously, made the following comment on the relationship between his molecular orbital analysis and the Lewis electron-pair model "Now I have a favourite argument that Lewis electron pair bonding is better described by a pair of electrons in a molecular orbital than by the Heitler-London bond. If the chemical bond has any polarity, it is necessary to add an ionic term, that is a Heitler-London plus an ionic term, to represent the bond. That is rather a messy description whereas the molecular orbital- this is not the spectroscopic but the chemical molecular orbital, the delocalized molecular orbital fits very nicely to the... 

This picture of the bond in H2 involving two electrons, each in a a orbital but with opposite spins, is analogous to the Lewis electron-pair bond in H2 (Fig. 2-3). It is convenient to carry along the idea that a full bond between any two atoms involves two electrons. Thus we define as a useful theoretical quantity the number of bonds in a molecule as follows ... 

The essential event in a chemical reaction is the breaking or making of a bond between two centers, be they atoms or complex fragments. In the most common situation these two-center bonds are also two-electron bonds, known as Lewis electron-pair bonds [14]. Following Pauling [15], the Lewis bond A-B is described by a blend of the spin-paired Heitler-London (HL) form and the possible zwitterionic situations, as expressed in Eq. (1) ... 

This is a typical Lewis electron-pair bond. By contrast, the formula for Lij projects homopolar ionic bonding discussed before. 

Orbital hybridization descriptions because they too are based on the shared electron pair bond enhance the information content of Lewis formulas by distinguishing... 

The years from 1923 to 1938 were relatively unproductive for G. N. Lewis insofar as his own research was concerned. The applications of the electron-pair bond came largely in the areas of organic and quantum chemistry in neither of these fields did Lewis feel at home. In the early 1930s. he published a series of relatively minor papers dealing with the properties of deuterium. Then in 1939 he began to publish in the field of photochemistry. Of approximately 20 papers in this area, several were of fundamental importance, comparable in quality to the best work of his early years. Retired officially in 1945, Lewis died a year later while carrying out an experiment on fluorescence. 

The boranes are electron-deficient compounds (Section 3.8) we cannot write valid Lewis structures for them, because too few electrons are available. For instance, there are 8 atoms in diborane, so we need at least 7 bonds however, there are only 12 valence electrons, and so we can form at most 6 electron-pair bonds. In molecular orbital theory, these electron pairs are regarded as delocalized over the entire molecule, and their bonding power is shared by several atoms. In diborane, for instance, a single electron pair is delocalized over a B—H—B unit. It binds all three atoms together with bond order of 4 for each of the B—H bridging bonds. The molecule has two such bridging three-center bonds (9). 

The electron-pair bond as postulated by Lewis consists of two electrons held jointly by two atoms. By assuming that atoms tend to surround themselves with an outer shell of either shared or unshared electron pairs, usually four in number, but sometimes more or less, Lewis... 

How does this NBO description of A—F bonding compare with the classical valence-bond (VB) picture 14 Although it is evident that the NBO Lewis-structure description is very VB-like in its emphasis on localized, transferable electron-pair bonds and lone pairs of the chemist s Lewis diagram, there are important differences in mathematical detail. 

For a general closed-shell AX , species, the Lewis-type assumption of a shared A X electron-pair bond for each coordinated monovalent atom X nominally requires m orbitals on A to accommodate the 2m bonding electrons, plus additional orbitals for any nonbonded pairs. Thus, for m bonds and t lone pairs, apparent octet-rule violations occur whenever... 

Soon after the quantum revolution of the mid 1920s, Linus Pauling and John C. Slater expanded Lewis s localized electronic-structural concepts with the introduction of directed covalency in which bond directionality was achieved by the hybridization of atomic orbitals.1 For normal and hypovalent molecules, Pauling and Slater proposed that sp" hybrid orbitals are involved in forming shared-electron-pair bonds. Time has proven this proposal to be remarkably robust, as has been demonstrated by many examples in Chapter 3. 

The study of the reactions of the simple free radicals begun by Bodenstein and Lind in 1906 on the kinetics of gas phase reactions showed that the reactions of H2 with CI2 and Bt2 were complex processes/ and a radical chain mechanism for these reactions (equations 14-18) was proposed in 1919 by Christiansen, Herzfeld, and Polanyi/ The theoretical basis for understanding these reactions in terms of free radicals was presented by G.N. Lewis in 1916, with the theory of the electron pair bond, and free radicals, or odd molecules / Further studies on chain reactions including the extension to explosions in gaseous systems were made by Hinshelwood and by Semenovwho shared the Nobel Prize in 1956. 

The theoretical basis for the understanding of free radicals was first provided by G.N. Lewis in 1916. His clear recognition of the electron pair bond and the possibility of odd electron systems was heavily influenced by the work of pioneers such as Gomberg, Schlenk, and Wieland, who had showed... 

The acidity of a clay can be either of the Brpnsted (H+ donor) or Lewis (electron pair acceptor) type. Even at temperatures below 100 °C, tertiary carbocation intermediates can be generated on clays with high Brpnsted acidity through protonation of the C=C double bond in secondary alkenes, as in the clay-catalyzed formation of MTBE from methanol and isobutene ... 

Ik this chapter we explore how symmetry considerations can be applied to one of the most pervasive concepts in all of chemistry bonding between atoms by the sharing of pairs of electrons. Though the idea of an electron-pair bond was first introduced in 1916 by G. N. Lewis, it was only after the advent of quantum mechanics that it could be given a proper theoretical basis. This came about through the development of two theories valence bond (VB) theory and localized MO theory both of which describe the electron pair in terms of orbitals of the component atoms of the bond. 

The convenient name covalent bond, which we shall often use in this book. in place of the more cumbersome expressions shared-electron-pair bond or electron-pair bond, was introduced by Langmuir (loc. cit. [7], 868). Lewis preferred to include under the name chemical bond a more restricted class of interatomic interactions than that giwri by ov definition f the chemical bond is at all times and in all molecules men ly a pair. i vtror s held jointly by two atoms --Lewis, op. cit. p. 78). 

Lewis s concept of shared electron pair bonds allows for four-electron double bonds and six-electron triple bonds. Carbon dioxide (C02) has two carbon-oxygen double bonds, and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon-nitrogen triple bond. 

Exercise 2-1 Draw the Lewis electron-pair structure of 2-propanone (acetone) clearly showing the bonding and nonbonding electron pairs in the valence shell of each atom. Draw structural formulas for other compounds having the composition C3H60 that possess... 

---