Covalent bonding theory, Lewis

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 final acid-base theory that we shall consider was proposed by chemist Gilbert Lewis in the early 1920s. The Lewis Theory is the most general, including more substances under its definitions than the other theories of acids and bases. A Lewis acid is a substance that accepts a pair of electrons to form a covalent bond. A Lewis base is a substance that provides a pair of electrons to form a covalent bond. In order for a substance to act as a Lewis base, it must have a pair of unshared electrons in its valence shell. An example of this is seen when a hydrogen ion attaches to the unpaired electrons of oxygen in a water molecule, as shown here ... 

The formulas of the chemical compounds are no accident. There is an NaCl, but no NaCl2 there is a Cap2, but no CaF. On the other hand, certain pairs of elements form two, or even more, different compounds, e.g. C]u20, CuO N2O, NO, NO2. In the case of ionic compounds the relative number of positive and negative ions in a formula is governed simply by the rule of electrical neutrality. In covalent compounds, or within polyatomic ions (like NO ), structures are formed by covalent bonds (i.e., electron sharing). A hierarchy of covalent bonding theories exists, of which the simplest, the drawing of Lewis structures, is emphasized in this and in most elementary texts. 

Valence bond theory pictures bonding in complex ions as arising from coordinate covalent bonding between Lewis bases (ligands) and Lewis acids (metal ions). Ligand lone pairs occupy hybridized metal-ion orbitals to form complex ions with characteristic shapes. 

This chapter walks you through the evolution of covalent bond theory, starting with the Lewis dot structures you likely covered in General Chemistry and then advancing to valence bond and molecular orbital theories that stem from quantum mechanics calculations. As with most chemistry, this chapter is purely about following the bouncing balls (electrons). 

In covalent bonds, atoms share their electrons. We represent covalent bonding in Lewis theory by letting atoms share their dots such that some dots count for the octet of more than one atom. For example, we learned in Chapter 3 that elemental chlorine exists as the diatomic (two atom) molecule CI2. Lewis theory explains why. Consider the Lewis structure of chlorine ... 

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. 

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. 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

A proton (H+) is an electron pair acceptor. It is therefore a Lewis acid because it can attach to ( accept") a lone pair of electrons on a Lewis base. In other words, a Bronsted acid is a supplier of one particular Lewis acid, a proton. The Lewis theory is more general than the Bronsted-Lowry theory. For instance, metal atoms and ions can act as Lewis acids, as in the formation of Ni(CO)4 from nickel atoms (the Lewis acid) and carbon monoxide (the Lewis base), but they are not Bronsted acids. Likewise, a Bronsted base is a special kind of Lewis base, one that can use a lone pair of electrons to form a coordinate covalent bond to a proton. For instance, an oxide ion is a Lewis base. It forms a coordinate covalent bond to a proton, a Lewis acid, by supplying both the electrons for the bond ... 

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. 

In the same year that Bronsted and Lowry proposed their definition of acids and bases, an American chemist named Gilbert Lewis proposed an alternative definition that not only encompassed Bronsted-Lowry theory but also accounted for acid-base reactions in which a hydrogen ion isn t exchanged. Lewis s definition relies on tracking lone pairs of electrons. Under his theory, a base is any substance that donates a pair of electrons to form a coordinate covalent bond with another substance, while an acid is a substance that accepts that electron pair in such a reaction. As we explain in Chapter 5, a coordinate covalent bond is a covalent bond in which both of the bonding electrons are donated by one of the atoms forming the bond. 

What he does not seem to realize is that a perfectly good explanation existed for chemical bonding prior to the advent of the quantum mechanical explanation, namely Lewis s theory whereby pairs of electrons form the bonds between the various atoms in a covalently bonded molecule. Although the quantum mechanical theory provides a more fundamental explanation in terms of exchange energy and so on is undeniable but it also retains the notion of pairs of electrons even if this notion is now augmented by the view that electrons have anti-parallel spins within such pairs. 

At about the same time that Bronsted proposed his acid-base theory, Lewis put forth a broader theory, A base in the Lewis theory is the same as in the Brpnsted one, namely, a compound with an available pair of electrons, either unshared or in a tt orbital. A Lewis acid, however, is any species with a vacant orbital.1115 In a Lewis acid-base reaction the unshared pair of the base forms a covalent bond with the vacant orbital of the acid, as represented by the general equation... 

Before 1927 there was no satisfactory theory of the covalent bond. The chemist had postulated the existence of the valence bond between atoms and had built up a body of empirical information about it, but his inquiries into its structure had been futile. The step taken by Lewis of associating two electrons with a bond can hardly be called the development of a theory, since it left unanswered the fundamental questions as to the nature of the interactions involved and the source of the energy of the bond. Only in 1927 was the development of the theory of the covalent bond initiated by the work o Condon28 and of Hertler and London27 on the hydrogen molecule, described in the following paragraphs. 

Thus, by definition, electrophiles are electron-pair acceptors and nucleophiles are electron-pair donors. These definitions correspond closely to definitions used in the generalized theory of acids and bases proposed by G. N. Lewis (1923). According to Lewis, an acid is any substance that can accept an electron pair, and a base is any substance that can donate an electron pair to form a covalent bond. Therefore acids must be electrophiles and bases must be nucleophiles. For example, the methyl cation may be regarded as a Lewis acid, or an electrophile, because it accepts electrons from reagents such as chloride ion or methanol. In turn, because chloride ion and methanol donate electrons to the methyl cation they are classified as Lewis bases, or nucleophiles ... 

Considerable progress in the development of theoretical and synthetic coordination and organometallic chemistry was made with the use of electron ideas. Lewis elaborated in 1923 the classic electron theory of acids and bases [30], and used it to explain the coordination ideas of Werner [31] (in Ref. 32, this achievement is ascribed to Sidgwick). A Lewis acid (A) is a acceptor of the electron pair and a Lewis base (B) is its donor [33], In other words, A is a species that can form a new covalent bond by accepting a pair of electrons and B is a species that can form a new covalent bond by donating a pair of electrons. The fundamental Lewis acid-base theory is described by a direct equlibrium [Scheme (1.1)], leading to the formation of the adduct (acid-base complex) ... 

If the system contains three electrons, the two occupying 4 will be stabilized, and the other one, localized in XV2, destabilized. Here, the stability of the molecule depends upon the relative energies of 4, Tf, and the AOs thus, HHe dissociates spontaneously, but the three-electron bond in He2+ is moderately robust. Note that, in contradiction with Lewis theory, a covalent bond may be formed with one or three electrons. Electron-deficient bonds (where there are fewer than two electrons per bond) are particularly prevalent amongst boron compounds. 

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