This lesson delves into Lewis acids and bases in the context of the octet rule for electron-deficient compounds. Here, the concept is discussed, emphasizing the group 13 elements like boron or aluminium. Since group 13 elements possess three valence electrons, they form trivalent compounds with a sextet of electrons and a vacant orbital for the central atom. Consequently, these electron-deficient compounds accept electrons from other species to complete their octet in a chemical reaction. They are referred to as Lewis acids per the 'generalized theory of acids and bases' proposed by Gilbert N. Lewis.

Lewis's theory dealt with compounds that were not under the purview of Brønsted's definition. He proposed that the electron-deficient compounds act as a Lewis acid where their valence-shell octets are completed in a chemical reaction. Hence, a Lewis acid is the species that accepts a pair of electrons to form a new bond.

In contrast, a Lewis base is defined as the species that donates an electron pair. This is elucidated using the specific example of aluminum chloride reacting with ammonia to form a Lewis acid-base adduct. Here, the electron pair is transferred between the oppositely charged species to satisfy the octet. The Lewis acids and bases concept is further reiterated by the reaction between electron-deficient boron trifluorideand an electron-rich ammonia, as shown in Figure 1.

Figure 1. The reaction between Boron trifluoride and ammonia

Here, a significant charge is developed on the species. As the boron center has an empty orbital that could accept an electron, it localizes a positive charge. In contrast, the nitrogen center in ammonia accumulates a negative charge due to the presence of a lone pair of electrons. Hence when they interact, the lone pair of electrons in the valence shell of nitrogen is transferred to the boron atom in BF₃, indicated by the curved arrow above. Thus, the formal charges on boron and nitrogen are balanced, and as a result, the Lewis acid-base adduct possesses no net charge.

The Lewis theory provides an addendum to the Brønsted theory that only uses proton transfer to define acid-base reactions by incorporating the transfer of a lone pair. Therefore, while all Brønsted–Lowry acids are protic acids, Lewis acids can be protic or aprotic. This is delineated using the example of hydrochloric acid. HCl is an acid, according to the Brønsted-Lowry definition, given its ability to donate a proton. It is also a Lewis acid since its hydrogen atom loses the shared electrons to chlorine while simultaneously accepting the pair of electrons from ammonia.