Acids and bases have a few different definitions. Each definition approaches the concepts of what an acid or base is from a unique perspective.

The most common definition is the Bronsted-Lowry definition:
  • A Bronsted acid is defined as a proton donor. The proton is the H+ ion.
  • A Bronsted base is defined as a proton acceptor.

To demonstrate the Bronsted-Lowry definition, look at how hydrochloric acid interacts with water:

HCl (aq) + H2O (l) --> H3O+ (aq) + Cl- (aq)

In this reaction, the acid is HCl and the base is water. Water accepts the H+ ion (a proton) from HCl, so water is acting as a base.




Another similar definition of acids and bases comes from Arrhenius:
  • an Arrhenius acid is defined as any substance that increases the H+ concentration of a solution;
  • an Arrhenius base is defined as any substance that increases the OH- concentration of a solution.




The third acid-base definition comes from Lewis:
  • a Lewis acid is defined as an electron-pair acceptor;
  • a Lewis base is defined as an electron-pair donor.

Lewis definitions are most commonly used in organic chemistry.

For example, the Lewis diagram of ammonia (shown below, NH3) has a lone pair of electrons on nitrogen that are capable of forming a bond with boron in BF3. Because B accepts a pair of electrons, it is a Lewis acid, even though it contains no protons!



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