8.1 Introduction to Acids and Bases

Defining Acids and Bases

There are multiple ways of defining acids and bases. The two most important definitions are the Arrhenius definition and the Bronsted-Lowry definition. The former describes how acids and bases behave when dissolved in water. An Arrhenius acid is a substance that, when dissolved in water, increases the concentration of ${H^+}$ ions (note: since ${H^+}$ doesn’t often exist on its own, it often reacts with ${H_2O}$ and forms ${H_3O^+}$, hydronium. As a result, ${H^+}$ and ${H_3O^+}$ are used interchangeably). So, ${HCl}$, ${HF}$, and ${HNO_3}$ are typically considered acids as they dissociate into ${H^+}$ and an anion.

Depending on the degree of dissociation, acids can be considered strong acids or weak acids. Strong acids dissociate completely in aqueous solutions, releasing all of their ${H^+}$ ions. Weak acids, on the other hand, only dissociate partly. For example, if the weak acid ${HF}$ is placed in water, some of the hydrogen will remain bonded to the fluoride, while others dissociate into their respective ions.

The strong acids can be memorized with the acronym CBSPIN, which represents HydroChloric (${HCl}$), HydroBromic (${HBr}$), Sulfuric (${H_2SO_4}$), Perchloric and chloric (${HClO_4/HClO_3}$) , HydroIodic (${HI}$), and Nitric (${HNO_3}$) acid. Everything else will be considered a weak acid. The strong bases in AP Chem are alkali metal-hydroxides and alkaline earth metal-hydroxides (groups 1 and 2, such as ${NaOH}$ or ${Ca(OH)_2}$).