18.2 - Lewis theory

18.2.1 - Define and apply the terms Lewis acid and Lewis base. A Lewis acid - base reaction involves the formation of a new covalent bond in which both electrons are provided by one species. Such bonds are called dative covalent bonds. The formation of complexes (see 13.2.4 and 13.2.5) is usually a Lewis acid - base reaction.

Lewis acids and bases

Lewis extended the theory of acids to cover both non-aquoeus systems and systems that do not involve proton trqansfers. He defin ed a Lewis acid from the point of view of the electrons rather than from the point of view of hydrogen ions (protons)

An electron pair donor becomes a Lewis base and an electron pair receiver is a Lewis acid.

To see how this affects Arrhenius acid - base behavious consider the reaction between a hydrogen ion and a hydroxide ion

H⁺ + OH^- H₂O

In this reaction the H⁺ ion is accepting a lone pair onf electrons from the hydroxide (OH^-) ion. According to Lewis' definition the H+ is and acid (as we already know).

The hydroxide ion is donating a lone pair of electrons and is defined as a Lewis base

Summary

• Lone pair acceptor - Lewis acid
• Lone pair donor - Lewis base

The advantage of Lewis definition is that it can be applied to systems that do not have hydrogen ions involved at all.

Example

The reaction between ammonia and Boron trifluoride takes place as follows:

:NH₃ + BF₃ H₃N_:BF₃

In this reaction the lone pair on the ammonia is coordinating into the empty orbital of the electron deficient boron trifluoride.

The Ammonia molecule is effectively donating a pair of electrons (:) to the boron trifluoride molecule to form the covalent bond and is therefore a base (by Lewis' definition)

The Boron trifluoride is accepting the lone pair of electrons and is therefore a Lewis acid.

Mnemonic (memory jog)
LEwis theory deals with the Electrons

Scope of Lewis' theory

Lewis theory gives an alternative slant on acid base reactions and extends the ideas of Arrhenius and Brønsted Lowry to sytems that do not involve hydrogen ions.

• All Arrhenius acids are also Brønsted Lowry acids and Lewis acids
• All Brønsted Lowry acids are also Lewis acids
• Lewis acids do NOT have to be either Brønsted Lowry or Arrhenius acids

The advantage of Lewis theory is that it extends to all dative coordinate bond systems. These are particularly common in transition metal chemistry and solvation (hydration when water is the solvent)

Transition metals

All transition metals form coordinate bonds with ligands. This means that they accept electron pairs from the ligands and consequently behave as Lewis acids.

The reacting ligands are Lewis bases

Hydration

The dissolution of an ionic substandce in water involves the water molecules forming dative coordinate bonds to the Positi¡ve ions by donation of the lone pairs from the oxygen atoms of the water. The positive ions could therefore be considered to be Lewis acids in this case.

The coordinating water molecules are Lewis bases

Summary

	Lewis' theory of acids covers all the previous definitions and can be extended to other systems that were previously excluded. Brønsted Lowry covers all of the Arrhenius acid behaviour and extends it to include non-aqueous systems but without having the scope of Lewis theory.

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