Lewis acid-base definition defined

In these four sections the terms Bronsted and Lewis acid sites have been widely used. While Bronsted sites can be reasonably well defined as proton-donating sites, the properties of Lewis acids are not so clear. A status report on Lewis acid-base definitions is highly relevant for clarifying these problems. 

The final acid-base concept we consider was developed by Gilbert N. Lewis, whose contribution to understanding the importance of valence electron pairs in molecular bonding we discussed in Chapter 9. Whereas the Brpnsted-Lowry concept focuses on the proton in defining a species as an acid or a base, the Lewis concept highlights the role of the electron pair. The Lewis acid-base definition holds that... 

Lewis acid-base definition A model of acid-base behavior in which acids and bases are defined, respectively, as species that accept and donate an electron pair. (606)... 

In modem usage, a Lewis acid would be defined as any substance capable of accepting electron density and a Lewis base as any substance capable of donating electron density. Many substances are capable of being either one or the other and some materials e.g., H2O) are capable of being both. A Lewis acid-base interaction requires coordination of the two so that the bonding electron density is shared by both the acid (acceptor) and the base (donor). With this definition, it becomes difficult... 

The Lewis bonding model with its electron pairs can be used to define a more general kind of acid-base behavior of which the Arrhenius and Bronsted-Lowry definitions are special cases. A Lewis base is any species that donates lone-pair electrons, and a Lewis acid is any species that accepts such electron pairs. The Arrhenius acids and bases considered so far fit this description (with the Lewis acid, H, acting as an acceptor toward various Lewis bases such as NH3 and OH , the electron pair donors). Other reactions that do not involve hydrogen ions can still be considered Lewis acid-base reactions. An example is the reaction between electron-deficient BF3 and electron-rich NH3 ... 

The American chemist G. N. Lewis formulated such a definition. He defined what we now call a Lewis base as a substance that can donate a pair of electrons. A. Lewis acid is a substance that can accept a pair of electrons. For example, in the protonation of ammonia, NH3 acts as a Lewis base because it donates a pair of electrons to the proton H, which acts as a Lewis acid by accepting the pair of electrons. A Lewis acid-base reaction, therefore, is one that involves the donation of a pair of electrons from one species to another. Such a reaction does not produce a salt and water. 

Concepts of acidity and basicity are, in practice, defined and evaluated by their utility. Since overly formd definitions can be restrictive the concepts of acidity evolve towaids more comprehensive definitions. For example the Lewis definition includes the Broensted definition simply regarding the proton as an electron acceptor. Because the interaction of Broensted acids and bases in solutions involves a common process, protic transfer, scales of acidity can be established, for example the Hammett [1] acidity function. For Lewis acid-base interaction there is no common process to provide a unique basis for comparisons of acid strength. Experimentally, the strength of a Lewis acid depends upon the particular Lewis base. The classification of acids and bases as hard or soft in the principle of hard and soft acids and bases (HSAB principle) clarifies the interactions of Lewis acids and bases [2a]. Strong interactions occur between hard acid and hard base, or between soft acid and soft base, hi the hard-hard interaction there is a considerable electrostatic contribution to bonding and in the soft-soft interaction there is a major covalent contribution to bonding. The use of density functional analysis has clarified the concepts of hardness and softness and an empirical ranking of Lewis acids, based on local hardness is, proposed [2c]. 

Classically acids are defined as substances that give off protons and bases are those that generate hydroxide ions when dissolved in water. These restrictive definitions are no longer adequate and have been modified or supplemented by others. Currently the most widely used acid-base definitions are due to Bronsted and Lowry, and Lewis. 

In 1923, Gilbert Newton Lewis defined an acid as an electron pair acceptor and a base as an electron pair donor. This definition is even more inclusive than the previous one because it includes all Bronsted-Lowry acids and bases as a subset and provides the foundation for the field of coordination chemistry. A coordination compound is the product of a Lewis acid-base reaction, such as the one shown in Equation (14.11) and Figure 14.5, in which the metal ion (Lewis acid) and ligand (Lewis base) are held together by a coordinate covalent bond. 

Lewis acids have low-lying empty valence orbitals that readily interact with lone pair(s) on other atoms (Lewis bases) to move towards a closed shell electronic configuration. These definitions do not specify that a lone pair of electrons must be transferred from one atom to another—only that an electron pair, residing originally on one atom (the Lewis base) must be shared with the Lewis acid. Neutralization is defined as coordinate covalent (or dative) bond formation. This results in a bond in which both electrons are furnished by one atom or ion (the Lewis base). 

We have seen that a base can be defined as combining with a proton and, therefore, requires at least one lone pair of electrons. A more general definition of acids and bases, due to G. N. Lewis, describes a base as any species (atom, ion or molecule) which can donate an electron pair, and an acid as any species which can accept an electron pair— more simply, a base is an electron-pair donor, an acid an electron-pair acceptor. Some examples of Lewis acids and bases are ... 

This electrophile/nucleophile dichotomy can be looked upon as a special case of the acid/base idea. The classical definition of acids and bases is that the former are proton donors, and the latter proton acceptors. This was made more general by Lewis, who defined acids as compounds prepared to accept electron pairs, and bases as substances that could provide such pairs. This would include a number of compounds not previously thought of as acids and bases, e.g. boron trifluoride (39),... 

Lewis defined a base as an electron pair donor and an acid as an electron pair acceptor. Lewis electron pair donor was the same as Bronsted-Lowry s proton acceptor, and therefore, was an equivalent way of defining a base. Lewis acids were defined as a substance with an empty valence shell that could accommodate a pair of electrons. This definition broadened the Bronsted-Lowry definition of an acid. The three definitions of acids and bases are summarized in Table 13.3. 

A Lewis acid is defined as a species that can accept a pair of electrons from a base. This is a very general definition of an acid proposed by G. N. Lewis in 1923 (LI). In the case of structure (I), the Al atom is not completely coordinated, i.e., it is bonded to three oxygen atoms instead of four. The aluminum atom thus has a total of six valence electrons instead... 

Before we review the methods used to determine surface acidity, we wish to define the type of acidity that should be measured. An acid is an electron-pair acceptor. In our opinion, the term acid should be limited to this definition rather than broadening the term to include oxidizing agents as well. We agree with Flockhart and Pink (10) who suggest a clear distinction be made between Lewis acid-Lewis base reactions (which involve coordinate bond formation) and oxidation-reduction reactions (which involve complete transfer of one or more electrons). 

In 1923 G. N. Lewis" proposed a definition of acid-base behavior in terms of electron-pair donation and acceptance. The Lewis definition is perhaps the most widely used of all because of its simplicity and wide applicability, espeaally in the field of organic reactions. Lewis defined a base as cn electron-pair donor and an acid as an electron-pair acceptor. In addition to all of llie reactions discussed above, the Lewis definition includes reactions in which no ions are formed and no hydrogen ions or other ions are transferred 2... 

Acids and bases were defined and described by early chemists, including Boyle, Lavoisier, Davy, Berzelius, Liebig, and Arrhenius. At the present time, depending on objectives, one of two definitions of acids and bases is likely to be accepted. These two definitions, by Bronsted and Lowry and by Lewis, were proposed about the same time. According to the Bronsted definition acids are substances having a tendency to lose a proton, and bases are those having a tendency to accept a proton. Thus, for an acid HA the acid-base half-reaction is... 

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