Understanding Acids and Bases: Theories, Strength, and Dissociation
Arrhenius Theory of Acids and Bases
An electrolyte is a substance that, in aqueous solution, conducts electricity. Svante Arrhenius established that when an electrolyte dissolves in water, it dissociates into two electrically charged ions with opposite charges (positive: cation, negative: anion). This theory is valid for any substance that conducts electricity in aqueous solution. Since acids and bases conduct electricity, these ideas were applied to them. An acid is any substance that dissociates to form hydrogen ions (H+), and a base is any substance that dissociates in aqueous solution to form hydroxide ions (OH–). This theory suffers from some limitations: it restricts the concept to aqueous solutions and presents difficulties in interpreting the properties of ammonia (NH3).
Brønsted-Lowry Theory of Acids and Bases
Acids and bases do not act in isolation but within what are known as acid-base reactions. According to Brønsted and Lowry, an acid is any substance in solution able to donate protons, and a base is any substance that accepts protons in solution. Brønsted and Lowry introduced the new concept of conjugate acid-base pairs. Acids and bases do not exist in isolation; as Brønsted and Lowry proposed, if an acid yields protons, there must be another substance capable of accepting them.
Relative Strength of Acids and Bases
An acid is considered strong when it has a great tendency to donate protons, and a strong base is one with a greater tendency to accept protons. Quantifying strength is challenging because it varies depending on the substance with which the acid or base reacts. To establish a qualitative scale of acid and base strength, we need to compare the acid or base with respect to a common reference substance, typically water. Quantitatively, the strength of an acid can be defined by its percentage ionization in water.
The strength of an acid can also be measured as its degree of ionization, expressed as a fraction (0 to 1). The ionization of an acid can be considered as the dissociation of a molecule into two parts: an anion and a hydrated proton. Thus, the degree of dissociation is:
An acid is considered strong when it is fully ionized in dilute aqueous solution, i.e., its degree of dissociation equals 1. The prototype of a strong base is the hydroxide ion (OH–). Hydroxides of alkali and alkaline earth metals (e.g., sodium, potassium) also behave as strong bases. These species are strong electrolytes; when dissolved in water, they completely ionize, producing high concentrations of OH– ions that readily accept protons. In general, most acids and bases are not fully ionized when dissolved in water. At equilibrium, the reacting species coexist with the products of the reaction.
Dissociation Constants: Ka and Kb
When a weak acid dissociates into ions in an aqueous solution, it establishes an equilibrium. According to the law of mass action, the equilibrium constant is as follows:
In dilute aqueous solutions, water is present in vast excess, so its concentration can be considered constant. The product of this constant and the water concentration is a new constant, called the dissociation constant of an acid (Ka):
We can apply the same principle to a weak base:
When determining the strength of an acid or base, the values of Ka and Kb are crucial. A higher Ka indicates a stronger acid, and a higher Kb indicates a stronger base. Consequently, stronger acids have weaker conjugate bases, and stronger bases have weaker conjugate acids.
Ion Product of Water (Kw)
It is well established that pure water weakly conducts electricity, indicating it is a very weak electrolyte.
In pure water, the hydronium ion (H3O+) concentration equals the hydroxide ion (OH–) concentration, as each ionized water molecule forms one of each type of ion. This self-ionization reaction of water is represented as:
