5.2:
Lewis Acids and Bases
In the periodic table, all elements in group 13 have three valence electrons. Consequently, they form trivalent compounds in which the central atom has a sextet of electrons and a vacant orbital.
These electron-deficient compounds achieve their octets by receiving electrons from other species in a chemical reaction.
For example, aluminum chloride reacts with ammonia by accepting the lone pair from the nitrogen atom.
This makes aluminum chloride — the electron-pair acceptor — a Lewis acid, and ammonia — the electron-pair donor — a Lewis base.
As the two oppositely charged species attract each other, the electron pair is transferred from the base to the acid. This completes the octet of the Lewis acid, resulting in a Lewis acid—base adduct.
Now, consider another Lewis acid—base reaction between the electron-deficient boron trifluoride and ammonia.
The electrostatic potential map of boron trifluoride indicates a significant positive charge centered on the boron atom.
In the case of ammonia, the nitrogen atom carries a negative charge on account of its nonbonding electron pair.
The opposite charges cause boron trifluoride and ammonia to react. During the reaction, the lone pair of ammonia fills the valence shell of boron to give it an octet of electrons.
In the adduct, boron carries a formal negative charge and nitrogen carries a formal positive charge — thus conserving the net charge of the compound.
Now, examine the reaction between a protic acid, like hydrochloric acid, and a base — ammonia. By the Brønsted–Lowry definition, HCl is an acid owing to its ability to donate a proton.
Although HCl is not an electron-deficient compound, in its reaction with ammonia, the hydrogen atom of HCl loses its shared electrons to chlorine while simultaneously accepting a pair of electrons from ammonia. HCl is, therefore, also a Lewis acid.
5.2:
Lewis Acids and Bases
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 trifluoride and 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.
Suggested Reading
- Brown, W.H., & Iverson, B.L., & Anslyn, V.E., & Foote S.C. (2014). Organic Chemistry. Mason, Ohio: Cengage Learning, 118-123.
- Solomons, G., & Fryhle, C. & Snyder, S. (2015). Organic Chemistry. New Jersey, NJ: Wiley, 192-194.
- Loudon, M., & Parise, J. (2016). Organic Chemistry. New York, NY: Macmillan Publishers, 230-234.