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More properly called the Brønsted-Lowry Theory, this logical extention of the original postulate on the nature of acids and bases called the Arrhenius Theory was independently developed by the chemists J. N. Brønsted and T.M. Lowry in the 1920s. The title comes from the shared credit of the two chemists in furthuring the study of acids and bases.

The basic definition supplied by the Brønsted-Lowry Theory is that an acid is a proton donor (of H+) and a base is a proton acceptor. Thus, an acid-base reaction is the transfer of a proton from an acid to a base. This differed from previous definitions in that it wasn't concerned so much with the make-up of a given acid or base as its characteristics; how easily did it accept or give up a proton? Any molecule that's capable of releasing a hydrogen ion is an acid, and any that can accept it is a base. This flexibility was important, because previously only molecules that contained a hydroxyl group OH- could be considered bases, which didn't comply with new discoveries of stronger acids and bases at the time.

So why the original mix-up? Part of the issue is that bases usually do contain a hydroxyl group. Having a free OH hanging around is the easiest way to snap up a proton (but not the only way!). Some absolutely vital byproducts of this realization included descriptions of autoionizaton of water, dissociation in aqueous solutions, amphoterism, acid and base strength, leveling effects, ternary acids, and salt derivations. None could have been predicted with the original Arrhenius Theory.

Another absolute vital result of Brønsted-Lowry Theory, and included intrinsically in the definition, was the concept of conjugate acids and bases. Once an acid has donated its proton, the story is not over. The molecule now has extra space in which it could feasibly receive a proton, making it a base. Likewise for bases, which once they have received a proton, could feasibly give it away. That doesn't necessarily mean it will happen, however. The strength of an acid or base is inversely proportional to the strength of its conjugate base or acid. An extremely strong acid isn't going to easily take back the proton it was so willing to give away (making it a weak conjugate base) and a base hungering madly for a proton won't relinquish it without a fight (making it a weak conjugate acid). For acids and bases which match fairly closely in their conjugate forms, however, constant reverse reactions are common. Under circumstances where the reaction is ocurring in an aqueous environment, conjugate forms help to explain the presence of both neutral pH water and salts.

Whitten, Kenneth W. Davis, Raymond E. Peck, M. Larry. General Chemistry, 6th ed. South Melbourne: Thomson Learning, Inc., 2000.

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