Lewis Acids and Bases
G.N. Lewis (1923) expanded the definitions of acids and bases beyond those established by Arrhenius and Brönsted-Lowry. In this framework, an acid is defined as a species that accepts electron pairs, while a base is one that donates electron pairs. This conceptual shift allows for the categorization of many acid-base reactions where there is no H+ involved.
Key Examples
A classic example illustrating this concept is the reaction between the electron-deficient species boron trifluoride (BF3) and ammonia (NH3):
Reaction Example
$$
BF3 + :NH3 \rightarrow BF3:NH3
$$
In this reaction, BF3 acts as a Lewis acid by accepting a lone pair from the ammonia molecule, which acts as a Lewis base. This distinctive perspective encompasses compounds lacking protons that can still exhibit acidic behavior, such as AlCl3, Co3+, and Mg2+, which are all Lewis acids because they can accept electron pairs from bases like water and ammonia.
In contrast, substances such as H2O, NH3, and OH– serve as Lewis bases due to their ability to donate their electron pairs.
This definition is particularly useful in explaining several acid-base reactions that do not conform to traditional Arrhenius or Brönsted-Lowry theories, thereby enriching the understanding of chemical interactions.