Lewis theory development

W. M. Latimer and W. H. Rodebush, J. Am. Chem. Soc. 42 (1920), 1419. This paper is often cited as the discovery of H-bonding (but see Jeffrey, note 5). Its title, Polarity and ionization from the standpoint of the Lewis theory of valence, reflects the strong influence of G. N. Lewis on all aspects of the early development of H-bond theory. 

This model originated with the theory developed by G.N. Lewis in 1916, and it remains the most widely-used model of chemical bonding. The essential elements of this model can best be understood by examining the simplest possible molecule. This is the hydro-... 

Let us consider the diffusion of radicals. According to the two-film theory developed by Lewis and Whitman (1924) for mass tranter across the... 

In the LCAO MO description, the H2 molecnle in its ground state has a pair of electrons in a bonding MO, and thus a single bond (that is, its bond order is 1). Later in this chapter, as we describe more complex diatomic molecules in the LCAO approximation, bond orders greater than 1 are discussed. This quantum mechanical definition of bond order generalizes the concept first developed in the Lewis theory of chemical bonding—a shared pair of electrons corresponds to a single bond, two shared pairs to a double bond, and so forth. 

One example of the development of a sedimentary structure has been worked out in Table 1 for the commonly present subject of acids and bases (De Vos Pilot, 2001). Several contexts followed each other over the years from Lavoisier, the ionic theories, the equilibrium theory, the Brpnsted-Lowry and Lewis theory, to contexts of our every day life and biochemical contexts. All these contexts are still being present in the contemporary curriculum. This confronts us with an incoherent acid-base theory, which is very difficult to learn and to teach, and even contains apparent inconsistencies between the layers of the sediment (see third column of the table). 

The first quantitative theory of chemical bonding was developed for the hydrogen molecule by Heitler and London in 1927, and was based on the Lewis theory of valence in which two atoms shared electrons in such a way that each achieved a noble gas structure. The theory was later extended to other, more complex molecules, and became known as valence bond theory. In this approach, the overlap of atomic orbitals on neighbouring atoms is considered to lead to the formation of localized bonds, each of which can accommodate two electrons with paired spins. The theory has been responsible for introducing such important concepts as hybridization and resonance into the theory of the chemical bond, but applications of the theory have been limited by difficulties in generating computer programs that can deal efficiently with anything other than the simplest of molecules. 

Modified Lewis theory . Chemistry Education Research and Practice, 2001, Vol. 2, pp. 67-72 (Part 1), 179-182 (Part 2) [http /www.uoi.gr/cerp] Recent Research Developments in Inorganic Chemistry, 2004, Vol. 4, pp. 1-11. 

This theory, developed by Lewis and Whitman, supposes that motion in the two phases dies out near the interface and the entire resistance to transfer is considered as being contained in two fictitious films on either side of the interface, in which transfer occurs by purely molecular diffusion. It is postulated that local equilibrium prevails at the interface and that the concentration gradients are established so rapidly in the films compared to the total time of contact that steady-state diffusion may be assumed. 

Since the time of G. N. Lewis, other, more powerful theories have been developed to explain chemical bonding, yet Lewis theory is still incredibly useful. Use what you know about the scientific method and the nature of scientific theories to explain how this can be so. 

Lewis, Q. N. (1875-1946) American chemist at the University of California at Berkeley who developed a simple theory for chemical bonding, now called Lewis theory. 

The movement of an electron pair during an acid-base reaction is the basis of the Lewis theory of acidity developed by Gilbert Lewis (see Chapter 4). A Lewis acid is defined as a substance that can accept a pair of electrons from another atom to form a coordinate covalent bond. A Lewis base is defined as a substance that can donate a pair of electrons to another atom to form a dative covalent (coordinate) bond. In the simple examples above, the proton (H" ) is the Lewis acid and the ammonia molecule and water molecule are the Lewis bases. 

He was the first to observe the very stable adsorbed monatomic films on tungsten and platinum filaments, and was able, after experiments with oil films on water, to formulate a general theory of adsorbed films. He also studied the catalytic properties of such films. Langmuir s work on space charge effects and related phenomena led to many important technical developments which have had a profound effect on later technology. In chemistry, his interest in reaction mechanism caused him to study structure and valence, and he contributed to the development of the Lewis theory of shared electrons. 

Furthermore, natural bond orbital (NBO) analysis of the first-order density has also been used to quantify aromaticity [73,74]. More recently Boldyrev and Zubarev [75] developed the adaptive natural density partitioning (AdNDP) algorithm attempting to combine the ideas of Lewis theory and aromaticity. The results obtained by the application of the AdNDP algorithm to the systems with non-classical bonding can be readily interpreted from the point of view of aromatic-ity/antiaromaticity concepts. 

To this aim, the ligand field theory comes to complete the presence of the attractive forces in chemical bonding, by combining the crystal field theory developed by Bethe and van Vleck (about 1930) with the contemporary theory of the valence bond of Pauling, which assumes the molecular complex formation as a reaction between the Lewis bases (ligands) and the Lewis acids (the metals or the metallic ions). 

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