Introduction

Weak acids, like acetic acid, HAc (Ac^– being an abbreviation for the acetate anion, CH₃COO^–), undergo partial dissociation such that only a fraction of dissolved molecules undergo dissociation (typically ~10 – 15%), and the remaining portion of molecules remain dissolved but intact. This is known as partial dissociation and is represented in a chemical equation with double arrows.

This situation, in which a reaction does not go to completion, is known as equilibrium. Reactions that result in equilibria will proceed until a balance of reactants and products is reached. This balance is different for each reaction, but it can be quantified with the equilibrium constant, K_eq. The equilibrium constant for the partial dissociation of acids is known as the acid dissociation constant (or just “dissociation constant”), K_a. In fact, the K_a values can be used to quantify the strength of acids (the stronger the acid, the higher the K_a value). The acid dissociation constant for acetic acid would be

Note that K_a is calculated from the equilibrium concentrations of H₃O⁺, Ac^–, and undissociated HAc, as denoted by the subscript “eq.” While the derivation goes beyond the scope of this lab, it can be shown that the equilibrium concentration of H₃O⁺ ions can be approximated from the initial concentration of acetic acid (the total amount added to solution before the dissolving process occurs) and its K_a value.

Bases can also be strong or weak depending on whether they completely or partially dissociate in or react with water, respectively. Weak bases, much like weak acids, dissociate until an equilibrium is reached, and this equilibrium can be described using the base dissociation constant, K_b. For example, NH₄OH only undergoes partial dissociation to form NH₄⁺ and OH^–. The equilibrium for NH₄OH can be described with the following equations,