Identify the following species as Lewis acids or Lewis bases: (a) \(\mathrm{H}^{+}\); (b) \(\mathrm{A}^{3+}\); (c) \(\mathrm{CN}^{-}\); (d) \(\mathrm{NO}_{2}{ }^{-}\).

Lewis acids are commonly identified as species that can accept an electron pair. They often possess a positive charge, or they have a vacant orbital where electrons can be accommodated. Examples of Lewis acids go beyond the simple protons or metal ions; they include molecules with electron-deficient atoms such as boron trifluoride (BF3) and aluminum chloride (AlCl3). In the mentioned problem, H+, as it lacks electrons, is a prime example of a Lewis acid. Similarly, the generic metal ion designated as A3+ signifies electron deficiency due to its positive charge, making it another example of a Lewis acid. This concept is integral to understanding chemical reactions, especially in the context of complex formation, catalysis, and organic synthesis.

Contrary to Lewis acids, Lewis bases are defined by their ability to donate an electron pair. They are typically characterized by their negative charge, or they contain lone pairs of electrons that are not involved in bonding. Common Lewis bases include molecules like ammonia (NH3), with its lone pair of nitrogen, and water (H2O), with lone pairs on oxygen. In the problem set, CN^– and NO2^– are Lewis bases as they both have negative charges denoting available electron pairs for donation. Recognizing Lewis bases is crucial for understanding their role in forming coordinate covalent bonds by donating electron pairs to electron-deficient species.

Electron pair donation is a key concept in the Lewis acid-base theory and involves the provision of an electron pair from one species (the Lewis base) to another (the Lewis acid) to form a coordinate bond. This concept can be visualized when a Lewis base like CN^– approaches a Lewis acid. The base uses its lone electron pair to establish a new bond, effectively donating the pair to the acid. This process is illustrated in various chemical reactions, such as the formation of adducts in coordination chemistry, where ligands (Lewis bases) donate electron pairs to central metal ions (Lewis acids).

The counterpart to electron pair donation is electron pair acceptance, where a species (the Lewis acid) accepts a pair of electrons from another (the Lewis base). Lewis acids such as the positively charged H+ ion and the generic metal ion A3+ exhibit this behavior. For instance, in acid-base reactions, the acid accepts an electron pair from the base to form a new bond, resulting in a Lewis adduct. Understanding electron pair acceptance is fundamental to predicting molecule behavior in numerous chemical reactions, from enzymatic catalysis to industrial processes where catalysts often function as electron pair acceptors, facilitating reaction pathways.