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,


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.


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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