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# Lewis Acid

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 Sub Topics Lewis acid refers to a definition of acid published by Gilbert N.Lewis in 1923, at approximately the same time as Bronstead and Lowry, Gilbert Lewis also proposed definitions of acids and bases. Lewis defined an acid as an electron-pair acceptor and a base as an electron pair donor. The Lewis definition encompasses some species not included within the Bronstead-Lowry definition. Lewis noted that a common feature of all Bronstead-Lowery bases is the presence of an unshared pair of electrons and he defined a base as a substance that can donate a pair of electron, now called a Lewis base. Rather than limiting the definition of an acid to a proton donor, he defined an acid as any substance that can accept a pair of electrons-now called a Lewis acid.

## Definition

Acid base theory was broadened considerably by G.N.Lewis in 1923. Striking at what he called "the cult of the proton", Lewis proposed that acids be defined as "electron pair acceptors and bases be defined as electron pair donors. In the Lewis acid base theory, proton donors are not the only acids many other species are acids as well. "

A Lewis acid is defined as an electron pair acceptor. The Lewis definition relies on a behavior that is not vastly different from the Bronstead-Lowery interactions. In the Lewis definition the focus of the reaction is no longer on the proton, but instead the electrons forming the coordinate covalent bond.

## Examples

According to Lewis acid is an electron pair acceptor. For example,

Cd2+(aq) + 4Cl-(aq) $\rightleftharpoons$ CdCl42-(aq)

Any species that is electron deficient and capable of accepting an electron pair is also a Lewis acid. Common examples of Lewis acids include BF3 and AlCl3. These compounds contain elements in group 3A of the periodic table that can accept an electron pair because they do not have filled valence shells of electrons.

The Lewis acid-base definition is much broader than the Bronstead definition. In other words it includes acids and bases that neither the Arrhenius nor Bronsted definition covers. For example, when an electron poor molecule, such as BF3 reacts with ammonia, the BF3 acts as an acid. In the Lewis model electron pairs in an acid-base reaction are not given away from base to acid, but rather are shared.

## List

Some of the hard and soft Lewis acids in solution are listed below.

Hard acids

H+, Li+, Na+, K+, (Rb+, Cs+), Mg2+, Ca2+, Sr2+, (Ba2+), Ti4+, Zr4+, Cr3+, Cr6+, MoO3+, Mn2+, Mn3+, Fe3+, Co3+, Al3+, Si4+, CO2

Boderline acids

Fe2+, Co2+, Ni2+, Cu2+, Zn2+, (Pb2+)

Soft acids

Cu+, Ag+, Au+, Cd2+, Hg+, Hg2+, Ch3Hg+, pi-acceptors such as quinones, bulk metals.

## Catalyst

Catalysts can also be described as electrophilic or nucleophilic depending on the catalysts electronic nature. Catalysis by Lewis acids and Lewis bases can be classified as electrophilic and nucleophilic respectively. In free-radicals reactions the initiator often plays a key roles. An initiator is a substance that can easily generate radical intermediates. Radical reactions often occur by chain mechanisms, and the role of the initiator is to provide the free radicals that start the chain reaction.

Lewis acid catalysts has been well appreciated during the past three decades. Although we treat transition metal catalysis separately from Lewis acid catalysis, it should be noted that as long as electron pair donors/acceptors are involved, the interactions between transition metals and corresponding substrates are always Lewis acid/Lewis base interactions and thus any electron pair acceptor catalyst initiated asymmetric reaction could be regarded as the chiral Lewis acid catalyzed reaction in its broadest sense.

## Strength

Example: AF3, AlCl3, AlBr3, AlI3 $\Rightarrow$ Acidic nature