Lowry – Bronsted Theory (Proton Theory)

In 1923, Lowry and Bronsted proposed a more general definition: 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 -- in short, an acid is a proton donor and a base is a proton acceptor. When $HCl$ dissolves in water it donates a proton to water, so $HCl$ behaves as the acid and $H_2O$ as the base; when $NH_3$ dissolves in water it instead accepts a proton from water, so $NH_3$ behaves as the base and $H_2O$ as the acid -- showing that water can act as either, depending on the partner it reacts with.

Conjugate acid-base pairs in the HCl + H2O equilibrium. Writing the reverse reaction of $HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$ as a general Bronsted-Lowry equilibrium $Acid_1 + Base_2 \rightleftharpoons Acid_2 + Base_1$: $H_3O^+$ donates a proton back to $Cl^-$ to re-form $HCl$, i.e. the products behave as an acid and a base too. The species left behind after a proton is donated is itself a base (the conjugate base of that Bronsted acid); chemical species differing only by a proton are called a conjugate acid-base pair. Here $HCl$/$Cl^-$ and $H_3O^+$/$H_2O$ are the two conjugate acid-base pairs: $Cl^-$ is the conjugate base of $HCl$ (equivalently $HCl$ is the conjugate acid of $Cl^-$), and $H_3O^+$ is the conjugate acid of $H_2O$.

Limitation of the Lowry–Bronsted theory. Substances such as $BF_3$ and $AlCl_3$ are well known to behave as acids (they readily accept an electron pair) even though they possess no proton to donate at all -- a case the strictly proton-based Lowry-Bronsted definition cannot classify as acidic, motivating the still more general Lewis concept. …