15.14: Lewis Acids and Bases

In 1923, G. N. Lewis proposed a generalized definition of acid-base behavior in which acids and bases are identified by their ability to accept or to donate a pair of electrons and form a coordinate covalent bond.

A coordinate covalent bond (or dative bond) occurs when one of the atoms in the bond provides both bonding electrons. For example, a coordinate covalent bond occurs when a water molecule combines with a hydrogen ion to form a hydronium ion. A coordinate covalent bond also results when an ammonia molecule combines with a hydrogen ion to form an ammonium ion. Both of these equations are shown here.

Reactions involving the formation of coordinate covalent bonds are classified as Lewis acid-base chemistry. The species donating the electron pair that compose the bond is a Lewis base, the species accepting the electron pair is a Lewis acid, and the product of the reaction is a Lewis acid-base adduct. As the two examples above illustrate, Brønsted-Lowry acid-base reactions represent a subcategory of Lewis acid reactions, specifically, those in which the acid species is H⁺. A few examples involving other Lewis acids and bases are described below.

The boron atom in boron trifluoride, BF₃, has only six electrons in its valence shell. Being short of the preferred octet, BF₃ is a very good Lewis acid and reacts with many Lewis bases; a fluoride ion is the Lewis base in this reaction, donating one of its lone pairs:

In the following reaction, each of two ammonia molecules, Lewis bases, donates a pair of electrons to a silver ion, the Lewis acid:

Nonmetal oxides act as Lewis acids and react with oxide ions, Lewis bases, to form oxyanions:

Many Lewis acid-base reactions are displacement reactions in which one Lewis base displaces another Lewis base from an acid-base adduct, or in which one Lewis acid displaces another Lewis acid:

Another type of Lewis acid-base chemistry involves the formation of a complex ion (or a coordination complex) comprising a central atom, typically a transition metal cation, surrounded by ions or molecules called ligands. These ligands can be neutral molecules, like H₂O or NH₃, or ions, such as CN^– or OH^–. Often, the ligands act as Lewis bases, donating a pair of electrons to the central atom.

This text is adapted from Openstax, Chemistry 2e, Section 15.2: Lewis Acids and Bases.

The Brønsted-Lowry model defines acids and bases in terms of protons, where acids are proton donors, and bases are proton acceptors. In contrast, the Lewis model defines acids and bases in terms of electron pairs, where Lewis acids are electron-pair acceptors, and Lewis bases are electron-pair donors.   

In a Brønsted acid, like acetic acid, the hydrogen can also act as a Lewis acid because it has an empty orbital to accept electrons donated from a base, like water, acting as a Lewis base.

The advantage of the Lewis model is that it allows scientists to classify a greater number of compounds as acids—including the ones that do not have ionizable protons. For example, boron trifluoride cannot be classified as an acid by the Brønsted-Lowry model because it does not contain hydrogen.

However, boron trifluoride possesses an incomplete octet with an empty orbital that can accept an electron pair from a Lewis base, such as ammonia, and, therefore, can act as a Lewis acid.

The resultant product formed by such Lewis acid-base reactions is called a Lewis acid-base adduct.

Some molecules, like carbon dioxide, can rearrange their electrons to act as a Lewis acid.

For example, in the reaction between water and carbon dioxide, an electron-pair moves from the carbon-oxygen pi bond to the terminal oxygen of the carbon dioxide.

The resultant empty orbital on the carbon atom allows it to accept the electron pair from a water molecule and act as a Lewis acid. As the water molecule donates the electron pair, it acts as a Lewis base.

In further rearrangement, a proton is transferred from the water oxygen to the terminal oxygen of the carbon dioxide, resulting in the formation of the carbonic acid adduct.

Small metal cations, like Al(III), can reacquire electron pairs and act as Lewis acids. For example, Al(III) accepts lone pairs of electrons from water and forms hexaaquaaluminum ions. Here, water molecules donate electron pairs and act as a Lewis base.