Q.Identify the conjugate acid–base pair for the following reactions in aqueous solution:

Proposed independently by Lowry and Bronsted in 1923, this theory defines acids and bases by proton transfer rather than by what they release in water: an acid is a substance with a tendency to donate a proton to another substance, and a base is a substance with a tendency to accept a proton from another substance -- an acid is a proton donor, a base is a proton acceptor.

When $HCl$ dissolves in water, $HCl+H_2O \rightarrow H_3O^++Cl^-$: $HCl$ donates a proton (acts as the acid) and $H_2O$ accepts it (acts as the base). When $NH_3$ dissolves in water, $H_2O+NH_3 \rightleftharpoons NH_4^++OH^-$: $NH_3$ accepts a proton (acts as the base) and $H_2O$ donates it (acts as the acid) -- so water can play either role depending on its partner.

Conjugate acid-base pairs. In the general reaction $Acid_1+Base_2 \rightleftharpoons Acid_2+Base_1$, the species left behind after an acid donates its proton is itself a base -- the conjugate base of that acid. Species differing only by a proton are called a conjugate acid-base pair. For $HCl+H_2O \rightleftharpoons H_3O^++Cl^-$, the two conjugate pairs are $HCl/Cl^-$ and $H_3O^+/H_2O$: $Cl^-$ is the conjugate base of $HCl$, and $H_3O^+$ is the conjugate acid of $H_2O$. A strong acid's conjugate base is always weak, and vice versa, because a strong acid's dissociation lies almost entirely to the right, leaving its conjugate base with only a negligible tendency to re-accept a proton.

Amphiprotic species such as $HCO_3^-$ or $HPO_4^{2-}$ can act as either a Bronsted acid (donating a proton) or a Bronsted base (accepting one), depending on the reaction partner -- they sit in the middle of a chain of successive conjugate pairs.

Limitation. Substances like $BF_3$ and $AlCl_3$ behave as acids (accepting an electron pair) despite having no proton to donate at all -- a case the Lowry-Bronsted definition cannot classify, which the Lewis concept resolves.