Bronsted-Lowry acids and bases

The Bronsted-Lowry definition is named for Johannes Bronsted and Thomas Lowry, who independently proposed it in 1923. A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a hydrogen ion (proton). Thus, according to the BL definition, acids and bases must come in what is called conjugate pairs. For example, consider acetic acid dissolved in water:

$${\rm CH}_3{\rm COOH}(aq) + {\rm H}_2{\rm O}(l) \rightleftharpoons {\rm H}_3{\rm O}^+(aq) + {\rm CH}_3{\rm COO}^-(aq)$$

Notice that we have written ${\rm H}_2{\rm O}(l)$ explicitly in these reactions. The reason is that acid/base dissociation occurs by a proton transfer reaction from an acid species to a specific water molecule. The transfer occurs through a hydrogen bond between the acid molecule and a solvating water molecule.

Here, CH$_3$COOH is a BL acid because it can donate a proton, and CH$_3$COO$^-$ its conjugate base because it can accept a proton. Note that ${\rm H}_2{\rm O}$ and ${\rm H}_3{\rm O}^+$ also form such a conjugate pair.

Note that the ${\rm H}_3{\rm O}^+$ rather than ${\rm H}^+$ has been used to denote the nature of ${\rm H}^+$ ions in water in the above reaction. This is really only a very crude representation of the true nature of solvated ${\rm H}^+$ ions. Although we will use it in the context of our discussion of acids and bases, the more modern understanding of the true nature of ${\rm H}^+$ in water will be given in a later section in the context of how acidic solutions conduct electricity.

Similarly when ammonia is dissolved in water, one has

$${\rm H}_2{\rm O}(l) + {\rm N}{\rm H}_3(aq)\rightleftharpoons {\rm N}{\rm H}_4^+(aq) + {\rm OH}^-(aq)$$

Here, NH$_3$ is the BL base and its conjugate acid is NH$_4^+$. Similarly, ${\rm H}_2{\rm O}$ acts as a BL acid and ${\rm OH}^-$ acts as a BL base.

Another important advantage of the BL definition is that we are not limited to water as the solvent. Consider the reaction that occurs when HCl is dissolved in ammonia:

$${\rm HCl}({\rm in}\; {\rm N}{\rm H}_3) + {\rm N}{\rm H}_3(l)... ...; {\rm N}{\rm H}_3) + {\rm Cl}^-({\rm in}\; {\rm N}{\rm H}_3)$$

Here, HCl acts as a BL acid with Cl$^-$ as its conjugate base. Also, NH$_3$ acts as a BL base with NH$_4^+$ as its conjugate acid.

An interesting ambiguity comes up within the BL definition, namely, that some species can act either as a BL acid or a BL base. Such beasts are called amphoteric. An example is the hydrogen carbonate ion, HCO$_3^-$. When dissolved in water, two posible reaction can occur:

$${\rm HCO}_3^-(aq) + {\rm H}_2{\rm O}(l)\rightleftharpoons {\rm H}_3{\rm O}^+(aq) + {\rm CO}_3^{2-}(aq)$$

or

$${\rm HCO}_3^-(aq) + {\rm H}_2{\rm O}(l)\rightleftharpoons {\rm H}_2{\rm CO}_3(aq) + {\rm OH}^-(aq)$$

In the first of these, HCO$_3^-$ acts as a BL acid with CO$_3^{2-}$ as its conjugate base, while in the second it acts as a BL base with H$_2$CO$_3$ as its conjugate acid. The treatment of amphoteric reactions is mathematically a little more hideous than you might think, so we will return to such reactions at the end of our acid-base section, when we have developed the necessary machinery.