According to Bronsted theory, an acid is a "proton donor" and a base is a "proton acceptor." 

Consider the equation,

HCl + NH₃ → NH⁴⁺ + Cl^-

Here, hydrochloric acid (HCl) "donates" a proton (H⁺) to ammonia (NH₃) which "accepts" it , forming a positively charged ammonium ion (NH₄⁺) and a negatively charged chloride ion (Cl^-). Therefore, HCl is a Brønsted-Lowry acid (donates a proton) while the ammonia is a Bronsted-Lowry base (accepts a proton). Also, Cl^- is called the conjugate base of the acid HCl and NH₄⁺ is called the conjugate acid of the base NH₃.

In the Lewis theory of acid-base reactions, bases donate pairs of electrons and acids accept pairs of electrons. A Lewis acid is therefore any substance, such as the H⁺ ion, that can accept a pair of nonbonding electrons. In other words, a Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such as the OH^- ion, that can donate a pair of nonbonding electrons. A Lewis base is therefore an electron-pair donor.

Consider the chemical reaction,

Al³⁺(aq) + 6 H₂O(l) Al(H₂O)₆³⁺(aq)

This is an example of a Lewis acid-base reaction. The Lewis structure of water suggests that this molecule has nonbonding pairs of valence electrons and can therefore act as a Lewis base. The electron configuration of the Al³⁺ ion suggests that this ion has empty 3s, 3p, and 3d orbitals that can be used to hold pairs of nonbonding electrons donated by neighboring water molecules.