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The Bronsted-Lowry definitions for acids and bases:
Acids are species that donate a proton (H+).
Bases are species that accept a proton.
HNO3 (aq) + H2O NO3-(aq) + H3O+(aq)
In this example, HNO3 is an acid and H2O is acting as a base.
NO3- is called the conjugate base of the acid HNO3, and H3O+ is the conjugate acid of the base H2O.
NH3 (aq) + H2O NH4+(aq) + OH-(aq)
In this example, NH3 is 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 amphiprotic.
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:
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 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 the electron-pair from water molecules. In this example the water acts as a Lewis base.
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