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Bronsted-Lowry Definition

Bronsted and Lowry, in 1923, defined acids as proton donors and bases as proton nacceptors and aced, base reactions are regarded as proton-transfer reactions. When applied to water as a solvent,

2H2O  <==>  H3O+  + OH-
                       acid        base

this definition does not differ form the already known definition of Arrhenius [who defined acids which ionize to give H+ and bases as those which ionize to give hydroxide ions (OH-)]. The usefulness of this definition is in its capacity to categorizing species as ‘acids’ or ‘bases’ in proton containing solvents other than water, such as liquid NH3, sulphuric acid or hydrofluoric acid.

     2NH3  <==>  NH4+  +  NH2-
                           (acid)       (base)
2H2SO4  <==>  H3SO4+  HSO4-
                            (acid)            (base)
      3HF  <==>  H2F+  +  HF2-
                           (acid)   (base)

According to this definition any solute which when dissolved in ammonia ionizes to given NH4+ (ammonium ion) may be termed as an acid and those solutes which ionize in liquid ammonia to give NH2- (amide ion) will be defined as bases in this solvent. Species that differ form each other only in terms of a transferred proton are termed as conjugates. For example:

HB1    +      B2           <==>    HB2     +      B1
(Acid-1)       (Base-2)           (Acid-2)     (Base-1)

The base formed by the loss of proton form an acid is termed as the conjugate base, that is, B1 is the conjugate base of the acid HB1 and B2 is the conjugate base of the acid HB2.

HCI        +      NH3         →        NH4+       +       CI-
Acid               Base                  Conjugate          Conjugate
                                                        acid                      base

Here CI- is the conjugate base of the acid HCI and NH3 is the conjugate base of the acid NH4+. Stronger acid and stronger base of each conjugate pair react to form the weaker acid and the weaker base.

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