Definition of Lewis Acids and Bases

What are Lewis Acids and Bases?

A Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor. They can react with each another such that a covalent bond forms, with both electrons provided by the Lewis base.

Lewis acids have an unoccupied low-energy atomic or molecular orbital. Lewis bases have occupied relatively high energy atomic or molecular orbitals.

This is a more flexible definition of acids and bases than the more traditional Bronsted-Lowry definitions, which say that acids are species that donate a proton (H+), and bases are species that accept a proton.

The Lewis definition encompasses compounds that the Bronsted-Lowry definition says are acids or bases, and also encompasses compounds that do not donate protons, but still exhibit acid/base behavior.

Lewis acids and bases can be described as hard or soft.

Examples of Lewis Acids: H+, K+, Mg2+, Fe3+, BF3, CO2, SO3, RMgX, AlCl3, Br2.

Examples of Lewis Bases: OH-, F-, H2O, ROH, NH3, SO42-, H-, CO, PR3, C6H6.

Lewis Acid Example

An example of an acid/base reaction that can't be described by the Bronsted-Lowry definition is Al3+ in water.

Aluminum Hexa-aqua Complex

Al3+ is a hard Lewis acid.

It reacts with water to produce an aqua complex Al(H2O)63+.

In this complex, the Al3+ accepts electron-pairs from six water molecules.

The water donates electron-pairs, so it is acting as a Lewis base.

Traditional Acids as Lewis Acids

In the traditional acid/base neutralization reaction of H+ and OH-, H+ is a Lewis acid because it accepts an electron pair from the OH-. Since the OH- donates an electron pair it is a Lewis base.

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