Lewis acid-base definition with electron-deficient atoms

In the Lewis acid-base definition, an acid is any species that accepts a lone pair to form a new bond in an adduct. Thus, there are many more Lewis acids than other types. Lewis adds include molecules with electron-deficient atoms, molecules with polar multiple bonds, and metal cations. 

The Lewis acid-base definition focuses on the donation or acceptance of an electron pair to form a new covalent bond in an adduct, the product of an acid-base reaction. Lewis bases donate the electron pair, and Lewis acids accept it. Thus, many species that do not contain El are Lewis acids. Molecules with polar double bonds act as Lewis acids, as do those with electron-deficient atoms. Metal ions act as Lewis acids when they dissolve in water, which acts as a Lewis base, to form an adduct, a hydrated cation. Many metal ions function as Lewis acids in biomolecules. 

The Br0nsted-Lowry definition of acids and bases depends on the transfer of a proton from the acid to the base. The base uses a pair of nonbonding electrons to form a bond to the proton. G. N. Lewis reasoned that this kind of reaction does not need a proton. Instead, a base could use its lone pair of electrons to bond to some other electron-deficient atom. In effect, we can look at an acid-base reaction from the viewpoint of the bonds that are formed and broken rather than a proton that is transferred. The following reaction shows the proton transfer, with emphasis on the bonds being broken and formed. Organic chemists routinely use curved arrows to show the movement of the participating electrons. 

If the acid-base reaction is written with the electron pairs and the arrows, as shown for water and HCl, the Lewis base definition is quite useful. The electron-rich molecule is the base, and the electron-rich atom donates two electrons. The molecule bearing the electron-deficient atom (hydrogen) is the acid. For reactions of organic molecules, it is essential to identify electron-rich and electron-poor components of molecules, to understand the electron flow, and to understand how to predict the products. That process begins with making the transition to thinking in terms of Lewis bases/Lewis acids rather than Br0nsted-Lowry acids and bases. 

There is one important distinction. When a base such as "NH2 donates electrons to the electron-deficient proton of HCl, for example, a conjugate acid is formed H-NH2, with the new covalent N-H bond. This reaction breaks the H-Cl bond, generating Cl. There are two products because the acid-base reaction breaks the bond between the proton and the other atom and forms the conjugate acid. When a Lewis base such as ammonia reacts with a Lewis acid such as BFg, however, a dative bond is formed to give the Lewis acid-Lewis base complexes (an ate complex), so there is one product rather than two. This statement is an overgeneralization, but it offers a useful starting point to allow the two definitions to be distinguished from one another. These points will be emphasized in this section. 

It is clear that atoms other than hydrogen can be electron deficient and function as electron pair acceptors. Can a carbon atom function as a Lewis acid The answer is yes, if the definition is modified somewhat. Various reactions generate carbocation intermediates (see 55 and 58) and a Lewis base can certainly donate electrons to that positive carbon. A species that donates electrons to carbon is called a nucleophile (see Section 6.7), so an electron donor that reacts with 55 or with 58 is a nucleophile. In addition to carbocations, which are charged species, the carbon atom in a polarized bond is electron deficient, and a nucleophile could donate electrons to the 6+ carbon. This is the basis of many organic reactions to be discussed, particularly in Chapter 11. The fundamental concept of a species donating electrons to a carbon is introduced in this section, with the goal of relating this chemical reactivity to the Lewis acid-Lewis base definitions used in previous sections. 

The Lewis definition implies the presence of high electron density centres in Lewis bases, and low electron density centres in Lewis acids. In a reaction between a Lewis acid and a Lewis base the electron pair donated by the base is used to form a new sigma bond to the electron-deficient centre in the acid. The identification of Lewis bases follows basically the same guidelines as the identification of Br0nsted-Lowry bases. They frequently contain atoms that have non-bonding electrons, or lone pairs. In contrast Lewis acids frequently contain atoms with an incomplete octet, a full positive charge, or a partial positive charge. 

Gilbert N. Lewis recognized the similarity in behavior of boron tri-fiuoride and a transferred proton toward a base, and in 1923 enunciated a definition of acid-base reaction in terms of sharing of an electron pair—fl base donates an electron pair in covalent bonding and an acid accepts the pair. The acid is called an ELECTROPHILE, and the base is called a nucleophile. In the base, the atom with the unshared pair of electrons is an electron-rich site, and, in the acid, the atom that accepts the pair of electrons to form a covalent bond is an electron-deficient site. The Lewis theory focuses attention on the electron pair rather than on the proton, and in so doing broadens the concept of acidity. The transferred proton of a so-called Brbnsted acid is a special case of a Lewis acid. 

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