Ionic bond Lewis model

His second observation concerns the polarity of chemical compounds. Diatomic molecules like F2 or CI2 are completely nonpolar. The highest polarity is found in ionic compounds like NaF or KCl, but in between there appears to be a large number of molecules of intermediate polarity. This observation suggests that the formation of bonds is frequently accompanied by a partial transfer of electrons between the bonded atoms. A good model for the chemical bond should therefore be flexible enough to allow a gradual transition from nonpolar to ionic bonding. Lewis postulated that ... 

Chapter 1 discusses classical models up to and including Lewis s covalent bond model and Kossell s ionic bond model. It reviews ideas that are generally well known and are an important background for understanding later models and theories. Some of these models, particularly the Lewis model, are still in use today, and to appreciate later developments, their limitations need to be clearly and fully understood. 

The familiar Lewis structure is the simplest bonding model in common use in organic chemistry. It is based on the idea that, at the simplest level, the ionic bonding force arises from the electrostatic attraction between ions of opposite charge, and the covalent bonding force arises from sharing of electron pairs between atoms. 

This chapter provides a substantial introduction to molecular structure by coupling experimental observation with interpretation through simple classical models. Today, the tools of classical bonding theory—covalent bonds, ionic bonds, polar covalent bonds, electronegativity, Lewis electron dot diagrams, and VSEPR Theory—have all been explained by quantum mechanics. It is a matter of taste whether to present the classical theory first and then gain deeper insight from the... 

The Bond Valence Model has its roots in the ionic models of Pauling, but it can be equally well derived from the covalent models of Lewis. It thus spans the full range between ionic and covalent bonds without making any distinction between them. In the formal development of the model, a chemical structure is treated as a network of bonds in which each bond is associated with a valence that expresses its strength. The two network equations, viz the Valence Sum Rule (Equation 10.5) and the Equal Valence Rule (Equation 10.6), can be used to predict bond valences, and hence bond lengths, when the bond network is known. The influence of one part of the structure on another is transmitted through the network by application of the network equations. 

A small, positively charged ion (a cation) in an ionic compound can attract the electrons of a neighbouring negatively charged ion (an anion) towards it and distort the anion. When this happens, the anion is said to be poiarized. This distortion can be better represented by using the electron density model of an ionic compound, rather than using Lewis symbols. The outer electron density contours of a purely ionic bond, and an ionic bond which is polarized, are shown in Fig. 4.7. 

Three Types of Bonding Lewis Symbols and the Octet Rule The Ionic Bonding Model... 

Types of Bonding Three Ways Metals and Nonmetals Combine 277 Lewis Symbols and the Octet Rule 278 The Ionic Bonding Model 280 Why Ionic Compounds Form The Importance of Lattice Energy 280 Periodic Trends in Lattice Energy 281 How the Model Explains the Properties of Ionic Corrpounds 283... 

We represent chemical bonding with Lewis theory. In this model, valence electrons are represented as dots surrounding the chemical symbol for the element. Compounds are formed by allowing the electrons to be transferred between atoms (ionic bonding) or shared between atoms (covalent bonding) so that all atoms acquire an octet. 

Lewis theory was a pre-quantum theory and based on the simple idea that stable electron structures were those where the atoms could get a noble gas configuration by sharing or transferring electrons. The ideas of covalent and ionic bonds came from this model. Resonance structures were introduced because no unique bonds could be defined for some molecules, for example, benzene and some oxyanions, such as sulfate(vi) and nitrate(v), and were therefore an extension of Lewis ideas to such cases. 

The success of the HL model and its relation to Lewis model, posed a wonderful opportunity for the young Pauling and Slater to construct a general quantum chemical theory for polyatomic molecules. They both published, in the same year, 1931, several seminal papers in which they each developed the notion of hybridization, the covalent-ionic superposition, and the resonating benzene picture.Especially effective were Pauling s papers that linked the new theory to the chemical theory of Lewis, and that rested on an encyclopedic command of chemical facts. In the first paper, Pauling presented the electron-pair bond as a superposition of the covalent HL form and the two... 

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