The hydrogen requirement of Arrhenius and Brønsted–Lowry was removed by the Lewis definition of acid–base reactions, devised by Gilbert N. Lewis in 1923, in the same year as Brønsted–Lowry, but it was not elaborated by him until 1938. Instead of defining acid–base reactions in terms of protons or other bonded substances, the Lewis definition defines a base (referred to as a Lewis base) to be a compound that can donate an electron pair, and an acid (a Lewis acid) to be a compound that can receive this electron pair.
For example, boron trifluoride,  is a typical Lewis acid. It can accept a pair of electrons as it has a vacancy in its octet. The fluoride ion has a full octet and can donate a pair of electrons. Thus
is a typical Lewis acid, Lewis base reaction. All compounds of group 13 elements with a formula  can behave as Lewis acids. Similarly, compounds of group 15 elements with a formula , such as amines, , and phosphines, , can behave as Lewis bases. Adducts between them have the formula  with a dative covalent bond, shown symbolically as ←, between the atoms A (acceptor) and D (donor). Compounds of group 16 with a formula  may also act as Lewis bases; in this way, a compound like an ether, , or a thioether, , can act as a Lewis base. The Lewis definition is not limited to these examples. For instance, carbon monoxide acts as a Lewis base when it forms an adduct with boron trifluoride, of formula .
Adducts involving metal ions are referred to as co-ordination compounds; each ligand donates a pair of electrons to the metal ion. The reaction
can be seen as an acid–base reaction in which a stronger base (ammonia) replaces a weaker one (water).
The Lewis and Brønsted–Lowry definitions are consistent with each other since the reaction
is an acid–base reaction in both theories.