Acid-base definition according to Lewis
The American chemist Gilbert N. Lewis developed another acid-base concept. He extended the acid concept to all substances that can accept a free electron pair - do my homework . Therefore, Lewis acids are also called electron pair acceptors. Thus, Lewis also counts compounds among the acids that do not contain protons themselves, such as metal cations or compounds like aluminium trichloride (AlCl3). One of the most important Lewis acids is the proton (= H+).
Bases, on the other hand, are substances according to Lewis that have a free electron pair and can donate it or make it available to form a bond - take my class . They are also called electron pair donors. These include water, but also all anions such as the halide anions.
One of the best-known examples of a Lewis base is ammonia. The nitrogen atom in ammonia has three bonds to hydrogen atoms and then has two free outer electrons in the form of a free electron pair.
The Lewis theory is an extension of the Bronsted-Lowry theory, which is very useful in organic chemistry. For example, the course of a nucleophilic substitution on aromatics with aluminium chloride as catalyst can be well illustrated by means of the Lewis theory - chemistry problem solver . According to Lewis , however, the strength of acids and bases cannot be described quantitatively. Moreover, compounds such as water or hydrogen chloride are no longer considered acids, although the latter in particular clearly reacts acidically. The most illustrative model of acids and bases for many scientific processes is therefore still the Bronsted-Lowry theory.