In the early twentieth century, S. Arrhenius defined an acid as a compound that liberates hydrogen ions and a base as a compound that liberates hydroxide ions. In his acid-base theory, a neutralization is the reaction of a hydrogen ion with a hydroxide ion to form water.
The weakness of Arrhenius's theory is that it is limited to aqueous systems. A more general acid-base theory was devised by Brønsted and Lowry a couple decades later. In their theory, an acid is any compound that can donate a proton (hydrogen ion). A base is similarly defined as any substance that can accept a proton. This definition broadened the category of bases. In a Brønsted-Lowry neutralization, an acid donates a proton to a base. In the process, the original acidic molecule becomes a conjugate base; that is, it can accept a proton. Likewise, the base that accepted the proton becomes a conjugate acid, and it can donate a proton. Thus, in a Brønsted-Lowry neutralization reaction, conjugate acid-base pairs are generated.
The ability of a compound to liberate protons is a measure of its strength as an acid. For a compound to easily liberate a proton, its conjugate base must be weak. Similarly, a substance that liberates protons poorly must have a conjugate base that is strong. Thus, the conjugate bases of strong mineral acids are weak, while the conjugate bases of weak inorganic and organic acids are strong.
Review of General Chemistry
Structure of Organic Molecules
