Definition of Bronsted Lowry Acids and Bases

The Bronsted-Lowry definitions for acids and bases are:

• acids are species that donate a proton (H⁺)
• bases are species that accept a proton

Acid example:

In this example, HNO₃ is an acid, donating a proton, and H₂O is acting as a base, accepting a proton.

NO₃^- is called the conjugate base of the acid HNO₃, and H₃O⁺ is the conjugate acid of the base H₂O.

Base example:

In this example, NH₃ accepts a proton, acting as a base, and H₂O is acting as an acid.

NH₄⁺ is the conjugate acid of the base NH₃, and OH^- is the conjugate base of the acid H₂O.

A compound that can act as either an acid or a base, such as the H₂O in the above examples, is called amphoteric. Furthermore, H₂O can be described as amphiprotic, because it can act as both a donor or acceptor of protons, H⁺.

The Bronsted-Lowry definition of acids and bases does not encompass all chemical compounds that exhibit acidic and basic properties. A more general definition is that of Lewis acids and bases:

All Bronsted acids are also Lewis acids, because all Bronsted acids are electron pair acceptors. All Lewis acids are not Bronsted acids, because not all Lewis acids are H+ donors.

A Lewis acid is an electron pair acceptor; a Lewis base is an electron pair donor.

These definitions are broader than the Bronsted-Lowry definition in that they include many compounds that do not have protons, but exhibit acid/base behavior.

The Lewis definition encompasses the Bronsted-Lowry definition: In the 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 we call it a Lewis base.

As an example of a reaction not described by the Bronsted-Lowry definition, Al³⁺ in water is a Lewis acid. It reacts with water to form an aqua complex: the Al³⁺ accepts an electron pair from water molecules with the water acting as a Lewis base.