In 1923, Johannes Bronsted and Thomas Lowrie introduced a significant definition for acids and bases, which is foundational in understanding acid-base chemistry. According to the Bronsted-Lowry theory, an acid is defined as a proton donor, while a base is a proton acceptor. This framework allows us to analyze chemical reactions involving acids and bases in a more dynamic way.
For instance, when hydrobromic acid (HBr) reacts with water (H2O), HBr donates a proton (H+) to water, resulting in the formation of the hydronium ion (H3O+). In this reaction, HBr acts as the Bronsted-Lowry acid, giving away its H+ ion, while water serves as the base by accepting the proton. The remaining species after this reaction is bromide ion (Br-), illustrating the transformation of HBr into its conjugate base.
Conversely, when ammonia (NH3) interacts with water, ammonia functions as a base by accepting a proton from water, which acts as the acid in this scenario. The acceptance of the proton converts ammonia into ammonium ion (NH4+), while water, after donating the proton, is left as hydroxide ion (OH-). This exchange highlights the reciprocal nature of acid-base reactions, where acids donate protons and bases accept them.
Overall, the Bronsted-Lowry definitions enhance our comprehension of acid-base behavior, emphasizing that acids are characterized by their ability to donate H+ ions, while bases are defined by their capacity to accept H+ ions.