Theories of acids and bases have evolved over time to explain their chemical behavior and interactions in various contexts. The most prominent theories include Arrhenius, Brønsted-Lowry, and Lewis theories. Each theory builds upon its predecessors, providing deeper insight into the nature of acids and bases. Below is an introduction to these major theories:
Arrhenius Theory of Acids and Bases
Arrhenius defined an acid as a substance that liberates hydrogen ions (H+), or protons, in water. For example, HCl (hydrochloric acid) dissociates in water to release H+:
$$ HCl→H^++Cl^- $$
And a base as a substance that supplies hydroxyl ions upon dissociation. For instance, NaOH (sodium hydroxide) dissociates in water as follows:
$$ NaOH→Na^++OH^- $$
The neutralization reaction, where acids and bases react to form water and a salt:
$$ H^++OH^-→H_2O $$
While Arrhenius’s definitions were a major breakthrough at the time, they have a number of limitations. For example, they only apply in aqueous solution, but acid-base reactions can occur in other solvents too. Also, it is difficult to see how the familiar base ammonia (NH3) can function, as it does not contain an OH group.
Bronsted–Lowry Theory of Acids and Bases
According to the Bronsted-Lowry theory, an acid is a substance, whether charged or uncharged, that is capable of donating a proton, while a base is a substance, whether charged or uncharged, that is capable of accepting a proton from an acid.The relative strengths of acids and bases are measured by the tendencies of these substances to give up and take on protons.
Hydrochloric acid is a strong acid in water because it readily gives up its proton, whereas acetic acid is a weak acid because it gives up its proton only to a small extent. The strength of an acid or a base varies with the solvent. For example, hydrochloric acid is a weak acid in glacial acetic acid, while acetic acid is a strong acid in liquid ammonia. Consequently, the strength of an acid depends not only on its ability to donate a proton but also on the ability of the solvent to accept the proton from the acid. This property is referred to as the basic strength of the solvent.
Hydrogen chloride contains covalent HCl molecules. When hydrogen chloride dissolves in water, the following reaction takes place:
$$ \underset{Acid}{HCl}+\underset{Base}{H_2O}→Cl^-+H_3O^+ $$
This is an acid-base reaction in which a proton is transferred from HCl to H2O. In this reaction, H2O behaves as a base.
The reaction between HCl and NH3, is also an acid-base reaction:
$$ \underset{Acid}{HCl}+\underset{Base}{NH_3}→NH_4^+Cl^- $$
In this reaction NH3, is acting as a Bronsted-Lowry base (but not an Arrhenius base) by accepting a proton from HCl. This is an example of an acid-base reaction that does not involve water.
In the Bronsted-Lowry classification, acids and bases can be anions such as HSO4– and CH3COO–, cations such as NH4+ and H3O+, or neutral molecules such as HCl and NH4. Water can act as either an acid or a base, making it amphiprotic. Acid-base reactions occur when an acid reacts with a base to form a new acid and a new base. Since these reactions involve the transfer of a proton, they are referred to as protolytic reactions or protolysis.
Lewis Electronic Theory
Other theories have been suggested for describing acid–base reactions, the most familiar of which is the electronic theory of Lewis.
According to the Lewis theory, an acid is a molecule or anion that accepts an electron pair to form a covalent bond. A base is a substance that provides the pair of unshared electrons by which the base coordinates with an acid. Certain compounds, such as boron trifluoride and aluminum chloride, although not containing hydrogen and consequently not serving as proton donors, are nevertheless acids in this scheme. Many substances that do not contain hydroxyl ions, including amines, ethers, and carboxylic acid anhydrides, are classified as bases according to the Lewis definition.
Reference:
- Sinko, P. (2011). Martin’s Physical Pharmacy and Pharmaceutical Sciences. Baltimore, : Lippincott Williams & Wilkins, a Wolters Kluwer business.
