Lewis theory of bonding

Although electron-dot structures are often useful, they have some limitations. The Lewis theory of bonding itself has some limitations, especially in explaining the three-dimensional geometries of molecules. For this purpose in particular, we will discuss how another theory of bonding, involving orbitals, is more useful. 

Many chemical models are known to be incorrect, for example, the Lewis theory of bonding, or an incomplete description of chemical phenomena, for example, the octet rule (Chapter 4). For the purposes of teaching chemistry it is often preferable to use simple approximate bonding models that give correct predictions for the majority of cases, than to use a more accurate but complicated model, such as the quantum mechanical model (Chapter 12), which is based on the electron s wave properties. 

According to the Lewis theory of bonding, covalent bonding results when atoms share valence electrons. The atoms in molecules and those in polyatomic ions are held together by covalent bonds. 

Lewis dot symbol. 278 Lewis structure. 284 Lewis theory of bonding. 284 Lone pair. 284 Multiple bond, 285 Nonpolar. 288... 

Strength The Lewis theory of bonding enables us to make qualitative predictions about bond... 

Weakness Because the VSEPR model is based on the Lewis theory of bonding, it also fails to explain why bonds form. 

Strength The Lewis theory of bonding enables us to make qualitative predictions about bond strengths and bond lengths. Lewis structures are easy to draw and are widely used by chemists [M Section 8.9],... 

The Lewis theory of bonding describes a covalent bond as the sharing of a pair of electrons, but this does not necessarily mean that each atom contributes an electron to the bond. A covalent bond in which a single atom contributes both of the electrons to a shared pair is called a coordinate covalent bond. 

Initially, the noble gases were thought to be chemically inert. This apparent inertness helped provide a theoretical framework for the Lewis theory of bonding. Subsequently it was found that compounds of xenon can be made... 

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 ... 

As we have seen, the Lewis theory of acid-base interactions based on electron pair donation and acceptance applies to many types of species. As a result, the electronic theory of acids and bases pervades the whole of chemistry. Because the formation of metal complexes represents one type of Lewis acid-base interaction, it was in that area that evidence of the principle that species of similar electronic character interact best was first noted. As early as the 1950s, Ahrland, Chatt, and Davies had classified metals as belonging to class A if they formed more stable complexes with the first element in the periodic group or to class B if they formed more stable complexes with the heavier elements in that group. This means that metals are classified as A or B based on the electronic character of the donor atom they prefer to bond to. The donor strength of the ligands is determined by the stability of the complexes they form with metals. This behavior is summarized in the following table. 

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. 

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. 

Electron donation-acceptance reactions, which are considered to be Lewis acid-base interactions, also include the formation of coordination compounds, complex formation through hydrogen bonding, charge transfer complex formation, and so on. It should be apparent that the Lewis theory of acids and bases encompasses a great deal of both inorganic and organic chemistry. 

Hypervalent molecules and ions are fascinating species, since they appear to violate the traditional Lewis-Langmuir theory of bonding by expanding the valence shell of a main gronp element. Althongh species snch as PCI5,... 

G. N. Lewis (1875-1946) not only came up with important theories of bonding but also gave a new definition to acids and bases. 

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 way to deal with the specific geometry of a molecule is to return to Coulomb s law, that is, to look at electron-electron repulsion. VSEPR is such an electrostatic theory of bonding. As with Lewis dot structures, it ignores specific orbitals. The observed geometry reflects the attempt to minimize electron-electron repulsion by maximizing the distance between electrons. Bond angles are determined solely by the number of valence electrons around a central atom. It is instructive to use examples ... 

Obviously the Lewis theory of valence is unable to provide any satisfactory explanation of the stability and the physical origin of the chemical bond. Moreover, Lewis did not really understand the mechanism of the pairing of electrons. However, as shown below, quantum mechanics confirms very nicely most of the intuitive ideas of Lewis and consequently the essential features of its chemical formulas. 

The Lewis theory of chemical bonding, although useful AND easy to apply, DOES NOT TELL US HOW AND WHY BONDS FORM. A proper understanding of bonding comes from quantum mechanics. Therefore, in the second part of this chapter we will APPLY quantum mechanics TO THE STUDY OF THE GEOMETRY AND STABILITY OF MOLECULES. 

In the preceding Section we discussed the valence bond (VB) or electron-pair theory of bonding. The basic qualitative idea in this theory is essentially Lewis idea that each bonded atom pair in a molecule is held together by an electron pair or perhaps several electron pairs. These electron pairs are localized between particular pairs of nuclei. Moreover, it is assumed that the wave functions for these electrons are just the products of atomic wave functions. The MO theory starts with a qualitatively different assumption. 

In Lewis theory, covalent bonding occurs when atoms share electrons because the sharing concentrates electron density between the nuclei. In valence-bond theory, we visualize the buildup of electron density between two nuclei as occurring when a valence atomic orbital of one atom shares space, or overlaps, with a valence atomic orbital of another atom. The overlap of orbitals allows two electrons of opposite spin to share the space between the nuclei, forming a covalent bond. 

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. 

These reactions are Lewis acid-base reactions. The Lewis theory of acids and bases defines an acid as a substance capable of accepting a pair of electrons and a base as a substance that donates a pair of electrons. The terms acceptor and donor are sometimes used for acid and base, respectively. A Lewis acid-base reaction results in the formation of a coordinate bond, equation (9). 

---