Acids and Bases

Although the Arrhenius definitions of acids and bases are simplest and most useful, they are not the most widely applicable. Some compounds, like ammonia, NH₃, act like bases in aqueous solution even though they are not hydroxide-containing compounds. Also, the Arrhenius definition assumes that the acid-base reactions are occurring in aqueous solution. In many other cases, water is indeed the solvent. In many cases, however, water is not the solvent. What was necessary was to formulate a definition of acid and base that were independent of the solvent and the presence of H⁺ and OH^– ions.

Such a definition was proposed in 1923 by English chemist Thomas Lowry (1874-1936) and Danish chemists J. N. Brønsted (1879-1947) and N. Bjerrum (1879-1958) and is called the Brønsted-Lowry definition of acids and bases. (Bjerrum seems to have been forgotten.) The central chemical species of this definition is H⁺, which consists merely of a proton. By the Brønsted-Lowry definition, an acid is any chemical species that donates a proton to another chemical species. Conversely, a base is any chemical species that accepts a proton from another chemical species. Simply put, a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor.

The Brønsted-Lowry definition includes all Arrhenius acids and bases, since the hydrogen ion is a proton donor (in fact, it is a proton) and a hydroxide ion accepts a proton to form water:

But the Brønsted-Lowry definition also includes chemical species that are not Arrhenius-type acids or bases. The classic example is ammonia, NH₃. Ammonia dissolves in water to make a slightly basic solution even though ammonia does not contain OH^– ions. What is happening is that an ammonia molecule is accepting a proton from a water molecule to make an ammonium ion (NH₄⁺) and a hydroxide ion:

In essence, the water molecule is donating a proton to the ammonia molecule. The water molecule is therefore acting as the Brønsted-Lowry acid and the ammonia molecule is acting as the Brønsted-Lowry base.

In order to better understand the Brønsted-Lowry definition, it needs to be understood what is meant by a proton. The descriptions proton donor and proton acceptor are easy to remember. But are there actually bare protons floating around in solution? Not really. In aqueous solution, the protons are attached to the oxygen atoms of water molecules, giving them a positive charge. This species is called the hydronium ion and has the chemical formula H₃O⁺. It is more accurate to use the hydronium ion instead of the bare hydrogen ion when writing equations for chemical reactions between acids and bases in aqueous solution. For example, the reaction between the hydronium ion and the hydroxide ion, the typical Arrhenius acid-base reaction, would produce two molecules of water.

Chemical reactions can go forward or backward; when the rates of the reverse reactions are equal, it is at chemical equilibrium. It can be shown that each side of the equilibrium has a Brønsted-Lowry acid and base. For example:

On each side of the reaction there is an acid and a base. The NH ⁺ 4 ion is an acid because in the reverse reaction it donates a proton (H⁺) to the OH^-ion to form NH₃ and H₂O. With respect to the reaction above, the H₂O and OH^- species make up an acid-base pair, called a conjugate acid-base pair, while the NH₃ and NH₄+ species make up another conjugate acid-base pair. All Brønsted-Lowry acid-base reactions can be separated into reactions between two conjugate acid-base pairs. The conjugate acid always has one more H⁺ than the conjugate base.