From a chemist's viewpoint, the three acid-base definitions are different ways of defining and categorizing chemical reactions that involve the transfer of a proton. Each definition provides a different level of insight into how these reactions occur and can be applied to different chemical systems.
The Arrhenius definition defines an acid as a substance that produces H+ ions in solution and a base as a substance that produces OH- ions in solution. This definition is useful for understanding the behavior of acids and bases in aqueous solutions but is limited because it only applies to solutions and does not explain why certain substances behave as acids or bases.
The Brønsted-Lowry definition defines an acid as a proton donor and a base as a proton acceptor. This definition provides a more general framework for understanding acid-base reactions because it can be applied to non-aqueous systems and explains why certain substances act as acids or bases. For example, ammonia can act as a base by accepting a proton from a molecule with a weaker bond to hydrogen.
The Lewis definition defines an acid as an electron pair acceptor and a base as an electron pair donor. This definition provides the most general framework for understanding acid-base reactions because it can be applied to all chemical systems, including those that do not involve proton transfer. For example, a metal ion can act as a Lewis acid by accepting a pair of electrons from a molecule with a lone pair of electrons.
In summary, the Arrhenius definition is limited to aqueous solutions, the Brønsted-Lowry definition is more general and can be applied to non-aqueous systems, and the Lewis definition is the most general and can be applied to all chemical systems. Therefore, the Lewis definition provides deeper insight than the Brønsted-Lowry definition, while the Brønsted-Lowry definition is of more theoretical value than the Arrhenius definition.