Lewis Acids and Bases

Recall the Arrhenius description of acids and bases:

• An Arrhenius acid reacts in water to release a proton
• An Arrhenius base reacts in water to release a hydroxide ion

In the Bronstead-Lowry description of acids and bases:

• A B-L acid reacts to donate a proton
• A B-L base accepts a proton

A B-L base, therefore, is a compound with an unshared pair of electrons that can from a bond with a proton:

G.N.Lewis thought about acids and bases in terms of donation and acceptance of unshared pairs of electrons:

• A Lewis acid is defined as an electron-pair acceptor
• A Lewis base is defined as an electron-pair donor
• In the above example with ammonia, the ammonia is acting as a Lewis base (donates a pair of electrons), and the proton is a Lewis acid (accepts a pair of electrons)

The description of an acid and a base by Lewis is consistent with the description by Arrhenius, and with the definition by Bronstead-Lowry. However, the Lewis description, a base is not restricted in donating its electrons to a proton, it can donate them to any molecule that can accept them.

Since we are so used to thinking about aqueous solutions and protons as the electron pair acceptor (i.e. acid), any molecule like BF₃ that can act as an "acid" according to Lewis' definition is explicitly referred to as a "Lewis acid" (and not just as an "acid").

• Lewis acids include molecules that have less than an octet of electrons in the valence shell
• Many simple cations can function as Lewis acids. For example, Fe³⁺ can interact strongly with cyanide ions to form Ferricyanide ion, Fe(CN)₆^3-. The Fe³⁺ ion has vacant valence shell orbitals that can accept electrons from the cyanide ions (thus forming a covalent bond). The Fe³⁺ ion is acting as a Lewis acid. Many metal ions can thus behave as acids (in the Lewis definition).
• Compounds with multiple bonds can also behave as Lewis acids (i.e. they can accept pairs of electrons and form new bonds:

• The CO₂ is acting as a Lewis acid - it can accept a pair of electrons because the multiple bonds means there is a deficiency of valence electrons around the carbon. The product shown is unstable. The O^- group will act as a Lewis base and the neighboring H will act as a Lewis acid. This will result in an O-H bond (and the O on the left will have two unshared pair of electrons).

The solutions of many metal ions exhibit acidic properties

• Metal ions are cations. The positive charge attracts an unshared pair of electrons on a water molecule.

• The water molecules surrounding a metal ion serve to hydrate (solvate) the ion. However, the charge on the ion helps to increase the polar nature of the O-H bond of the water molecule - the metal draws electrons away from the oxygen, the oxygen in turn draws the shared electrons in the O-H bond away from the H atom.
• This increase in the polar nature of the O-H bond shifts the K_w of these particular waters, and they are more acidic than bulk solvent water molecules.
• The waters involved in hydration of a metal ion give up protons (i.e. are ionized) more readily than bulk solvent waters

Fe(H₂O)₆³⁺(aq) ó Fe(H₂O)₅(OH)²⁺(aq) + H⁺(aq)

K_a = 2 x 10^-3