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:

HNO3(aq) + H2O(aq) ⇌ NO3-(aq) + H3O+(aq)

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:

NH3(aq) + H2O(aq) ⇌ NH4+(aq) + OH-(aq)

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+.

Lewis Acids and Bases

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.

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