Lewis acid-base definition adduct

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 significance of the Lewis concept is that it is much more general than other definitions. Lewis acid-base reactions include many reactions that do not involve Brpnsted acids. Consider, for example, the reaction between boron trifluoride (BF3) and ammonia to form an adduct compound (Figure 15.11) ... 

The FMO definition also helps explain why Pearson s hard-hard and soft-soft interactions form stable complexes. Hard compounds have a large HOMO-LUMO gap, as shown in Figure 14.9 for F. Therefore, hard Lewis acid-base complexes tend to form strongly ionic compounds, such as LiF, where the interaction is dominated by electrostatic attractions. Soft compounds, on the other hand, have a small HOMO-LUMO gap, as shown in Figure 14.9 for I, so that these types of interactions form covalently bonded acid-base adducts, where the strength of the interaction is controlled primarily by the energies of the FMOs that participate in the bonding. 

Although Lewis and Bronsted bases comprise the same species, the same is not true of their acids. Lewis acids include bare metal cations, while Bronsted-Lowry acids do not. Also, Bell (1973) and Day Selbin (1969) have pointed out that Bronsted or protonic acids fit awkwardly into the Lewis definition. Protonic acids cannot accept an electron pair as is required in the Lewis definition, and a typical Lewis protonic add appears to be an adduct between a base and the add (Luder, 1940 Kolthoff, 1944). Thus, a protonic acid can only be regarded as a Lewis add in the sense that its reaction with a base involves the transient formation of an unstable hydrogen bond adduct. For this reason, advocates of the Lewis theory have sometimes termed protonic adds secondary acids (Bell, 1973). This is an unfortunate term for the traditional adds. 

The Lewis definition thus encompasses all reactions entailing hydrogen ion. oxide ion. or solvent interactions, as well as the formation of acid-base adducts such as R,NBF, and all coordination compounds. Usage of the Lewis concept is extensive in both inorganic and organic chemistry, and so no further examples will be given here, but many will be encountered throughout the remainder of the book.11... 

In a further generalization, Lewis [4] advanced a definition that was no longer restricted to proton-exchange reactions. An acid was defined as any substance that can accept a pair of electrons from a donor substance, and a base as any substance able to donate an electron pair and form a "dative" covalent bond. In these more general terms, the acid-base interaction includes the formation of covalent bonds and applies to any chemical reaction in which an addition compound (adduct) is formed through a coordinative bond ... 

The heats of adduct formation with many Lewis acids have been measured experimentally in the gas phase using the technique of ion cyclotron resonance [10]. The absolute proton affinity (PA) of a gaseous base molecule has a precise thermodynamic definition in relation to the negative of the enthalpy variation for a hypothetieal reaction of attachment of an isolated proton to a molecule M in the gas phase. 

Note that BF3 neither provides hydrogen ions in solution nor is a proton donor as required by the Arrhenius and Bronsted-Lowry definitions, respectively. Similarly, NH3 neither provides hydroxide ions in solution nor acts as a proton acceptor. (There are other instances in which ammonia is a proton acceptor, but it does not play that role in this reaction.) Therefore, this is not an acid-base reaction under these more restricted definitions. Boron trifluoride is, however, an electron-pair acceptor and ammonia an electron-pair donor, so the reaction can be classified as acid-base under the Lewis definitions. (In a sense, the product of a reaction between a Lewis acid and a Lewis base could be called a Lewis salt. However, the more technical term for such a product is a Lewis adduct.)... 

In the original Lewis definition (1923 [29], 1938 [30]), aeids are electron-pair acceptors and bases are electron-pair donors. The fundamental reaction between a Lewis acid A and a Lewis base B is the formation of a complex (or adduct or coordination compound or addition compound) A-B (reaction 1.17) ... 

First and foremost is the existence of isolable coordination compounds. By analogy with ammonium ions and the boron trifluoride adduct of ammonia (discovered by Gay-Lussac in 1809) it has been possible to isolate and characterize oxonium salts of alcohols, ethers, and ketones (49,176,314) and to prepare Lewis acid adducts of the ethers, sulfides, and phosphines (329), as well as many other compounds which fall under our definition of weak bases. 

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