Q.Discuss the Lowry – Bronsted concept of acids and bases.

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.