As will be exemplified in the following sections of this chapter, many molecules contain hydrogen atoms that can dissociate as protons during the hydration process and, therefore, can be classified as Arrhenius acids. However, the number of compounds that contain a hydroxide ion, which must be generated in order for the corresponding electrolyte to be categorized as an Arrhenius base, is much more limited. Additionally, Arrhenius investigated the dissociative behaviors and, subsequently, defined the acid/base reactivity, of strong and weak electrolytes in water. However, while water is known as the "universal solvent," due to its ability to dissolve more substances than any other chemical, certain solutes are immiscible in water. When mixed, these insoluble chemicals do not dissolve in water and, instead, form a heterogenous bilayer, as described in Section 7.3.
In 1923, Danish chemist Johannes Brønsted and English chemist Thomas Lowry independently recognized the limitations that are inherent in the Arrhenius system, and both scientists designed and executed experimental protocols that investigated the reactivity of solutes that ionized to produce H+1 ions in solution. As stated above, a large number of solutes contain hydrogen atoms that can dissociate as protons, and fewer molecules ionize to form hydroxide ions. Therefore, Brønsted and Lowry initially focused their studies on molecules that contain ionizable protons, which they classified as acids, in alignment with the terminology that was originally established by Arrhenius. However, in contrast to the aqueous solutions that Arrhenius studied, Brønsted and Lowry did not explicitly identify the solvent in which their acids must be dissolved and, consequently, were able to investigate the dissociative behaviors of solutes that are not water-soluble. As a result, a larger number of molecules can be classified as Brønsted-Lowry acids, which are defined as proton, H+1, donors in solution, relative to the number of chemicals that can be classified as Arrhenius acids. Because water is a specific solvent, the Arrhenius system for classifying acids and bases is a subcategory of the Brønsted-Lowry definition. As a result, all Arrhenius acids can also be categorized as Brønsted-Lowry acids. However, if an alternative solvent, such as ethanol, is utilized to dissolve a Brønsted-Lowry acid, that solute cannot be classified as an Arrhenius acid, which, by definition, requires that the chemical of interest be dissolved in water.
Furthermore, during their investigations of solutes that dissociate to form H+1 ions in solution, Brønsted and Lowry discovered that the resultant protons did not exist independently in those solutions. As stated in Section 3.6, while ionization improves the stability of a particle's electron configuration, relative to that of an atom, the charged state of the resultant ion is inherently destabilizing. As a result, a single ion cannot exist alone in nature and, instead, typically combines with an ion of the opposite charge to form an ionic compound. As an H+1 ion consists, by definition, of only a single proton, the hydrogen ion is extremely unstable, as it does not contain any electrons that would typically off-set some of the charge imbalance that results from the ionization process. Therefore, when a Brønsted-Lowry acid dissociates, the hydrogen ion that is produced is immediately absorbed by a nearby solvent molecule. Brønsted and Lowry regarded this phenomenon as chemically-significant and, as a result, classified these solvent molecules as bases. Because a Brønsted-Lowry base is defined as a proton, H+1, acceptor in solution, a Brønsted-Lowry acid/base reaction involves the transfer of a proton from a solute particle, which is classified as a Brønsted-Lowry acid, to a solvent particle, which is categorized as a Brønsted-Lowry base. Due to the synergistic relationship that exists between Brønsted-Lowry acids and bases, this classification system is inherently more sophisticated than the Arrhenius system, in which two ions are generated and analyzed as independent chemical entities.