Recall that solutes can be categorized as non-electrolytes, strong electrolytes, or weak electrolytes, based on the degree to which they dissociate, or separate, during the solvation process. As will be discussed in the following paragraphs, acids and bases can also be ranked and, therefore, compared, according to their relative strengths.
Because Brønsted and Lowry defined an acid as a solute that ionizes to produce protons, H+1, in solution, Brønsted-Lowry acids are electrolytes and, therefore, can be classified according to their dissociative behaviors. Since a non-electrolyte does not separate into charged particles when dissolved in a solvent, and a Brønsted-Lowry acid must, by definition, ionize to generate protons, H+1, in solution, acidic solutes cannot be classified as non-electrolytes. Furthermore, because many chemicals are amphoteric and, therefore, act as Brønsted-Lowry acids in some solutions and as Brønsted-Lowry bases in others, the strength of an acid is dependent on the identities of the other substances that are present in a given solution. Consequently, a particular acid can only be classified as "strong" or "weak" compared to other acids. Therefore, chemists have studied the dissociations of hundreds of Brønsted-Lowry acids and, subsequently, ranked these solutes from "strongest," or most-extensively ionized, to "weakest." However, because so many solutes can be categorized as Brønsted-Lowry acids, chemists classified the six strongest acids, hydroiodic acid, HI, hydrobromic acid, HBr, perchloric acid, HClO4, hydrochloric acid, HCl, sulfuric acid, H2SO4, and nitric acid, HNO3, as "strong acids," and categorized all remaining acids as "weak," "by default."
Chemists initially intended to qualitatively assess the strength of Brønsted-Lowry bases by applying the experimental and analytical principles that are described above. However, because Brønsted and Lowry categorized solvent molecules as bases, and, by definition, electrolytes must be solutes, Brønsted-Lowry bases are not categorized according to their dissociative behaviors. Instead, chemists studied the ionizations of Arrhenius bases, which are defined as solutes that generate hydroxide ions, OH–1, when dissolved in water. Because a non-electrolyte does not separate into charged particles when solvated, and an Arrhenius base must ionize to produce hydroxide ions, OH–1, in water, Arrhenius bases cannot be classified as non-electrolytes. Furthermore, the strength of a base is dependent on the identities of the other substances that are present in a given solution, and, as a result, a basic solute can only be classified as "strong" or "weak" compared to other bases. Therefore, upon investigating the dissociative behaviors of the most common Arrhenius bases, chemists ranked these solutes from "strongest," or most-extensively ionized, to "weakest." In order to align with the acid ratings that are stated above, chemists classified the eight strongest bases, cesium hydroxide, CsOH, rubidium hydroxide, ROH, potassium hydroxide, KOH, sodium hydroxide, NaOH, lithium hydroxide, LiOH, barium hydroxide, Ba(OH)2, strontium hydroxide, Sr(OH)2, and calcium hydroxide, Ca(OH)2, as "strong bases," and categorized all remaining bases as "weak," "by default."
Finally, because a "forward," or left-to-right, arrow is incorporated into a solution equation to indicate that a strong electrolyte has completely dissociated, the same type of arrow should be written in Brønsted-Lowry acid/base equations that represent the reactions of strong acids or strong bases, which also fully ionize when dissolved in solution. Furthermore, an equilibrium arrow should be drawn in Brønsted-Lowry acid/base equations that correspond to reactions of weak acids or weak bases, which, like weak electrolytes, only partially ionize and, subsequently, undergo both "forward" dissociation and "reverse" recombination processes in solution. If one reactant in a Brønsted-Lowry acid/base equation is classified as strong, and the other starting material is categorized as weak, the dissociative behavior of the stronger reactant will predominate, and, therefore, a "forward," or left-to-right, arrow should be utilized to indicate the reactivity of the stronger chemical.