The Arrhenius theory represents the first modern scientific attempt to define and categorize acids and bases. This foundational concept was developed by the Swedish chemist Svante Arrhenius in the late 19th century, starting with his 1884 doctoral thesis on the conductivity of electrolytes. His ideas were initially met with skepticism, but his work laid the groundwork for modern physical chemistry and earned him the Nobel Prize in Chemistry in 1903. The theory provided a simple classification system, setting the stage for more complex models that would follow.
The Requirement of Hydroxide Ions
The Arrhenius definition of a base is specific: it is any substance that, when dissolved in an aqueous solution, increases the concentration of hydroxide ions (\(\text{OH}^-\)). The term “aqueous” is mandatory, meaning the substance must be dissolved in water to be classified under this theory. The increase in hydroxide ion concentration gives Arrhenius bases their characteristic alkaline properties.
This basic behavior is achieved through dissociation, where the ionic compound separates into its constituent ions upon dissolving. For example, a strong base like sodium hydroxide (\(\text{NaOH}\)) is an Arrhenius base because it completely dissociates in water, breaking down into a sodium cation (\(\text{Na}^+\)) and a hydroxide anion (\(\text{OH}^-\)).
Another common example is potassium hydroxide (\(\text{KOH}\)), which similarly releases \(\text{OH}^-\) ions. These metal hydroxides, particularly those of Group 1 and some Group 2 elements, fit cleanly into the Arrhenius model because they directly contain and release the required hydroxide ion. The degree of dissociation determines the strength of the Arrhenius base.
Understanding the Arrhenius Acid
To fully understand the Arrhenius base, one must consider its counterpart, the Arrhenius acid. An acid is defined as any substance that dissociates in an aqueous solution to increase the concentration of hydrogen ions (\(\text{H}^+\)). Since a free hydrogen ion is a bare proton, it immediately combines with a water molecule (\(\text{H}_2\text{O}\)) to form a hydronium ion (\(\text{H}_3\text{O}^+\)).
The relationship between Arrhenius acids and bases is demonstrated through the neutralization reaction, which occurs when they are mixed in an aqueous solution. The \(\text{H}^+\) ions from the acid combine with the \(\text{OH}^-\) ions from the base to form neutral water (\(\text{H}_2\text{O}\)).
The remaining ions, such as the sodium ion from \(\text{NaOH}\) and the chloride ion from hydrochloric acid (\(\text{HCl}\)), combine to form a salt. This reaction, where the defining ions cancel each other out to form water, is the central concept of the Arrhenius theory.
Why the Definition Is Limited
Despite its historical significance, the Arrhenius theory is not universally applicable to all substances that exhibit basic properties. Its primary restriction is the requirement for an aqueous solution as the solvent. The model cannot explain acid-base behavior in non-water solvents, such as liquid ammonia or organic liquids, failing to describe a large number of chemical reactions.
The most significant conceptual failure concerns substances that act as bases but do not contain a hydroxide ion. For instance, ammonia (\(\text{NH}_3\)) is a well-known base that neutralizes acids. Since \(\text{NH}_3\) does not contain the \(\text{OH}^-\) group, the Arrhenius definition cannot classify it as a base. This necessitated the development of broader theories, such as the Brønsted-Lowry and Lewis definitions, which classify acids and bases based on proton transfer and electron movement.