Acids are defined as compounds which can reversibly lose protons to the solution. The pH at which this occurs (the halfway point) is the pK_a. Strong acids such as HCl will give up protons even at very low pH (hence low pK_a) and weak acids will only give up protons if the pH is very high (i.e. the free proton concentration is very low). The pH at which this occurs can easily be above 7.

Most biological acids, however, are weaker acids than HCl. The major class of biological acids is carboxylic acids. Because the difference in electronegativity between oxygen and hydrogen in a carboxylic acid is not as dramatic as it is with HCl, the tendency of a carboxylic acid to give up its proton is much less than that of HCl. However, carboxylic acids dissociate more readily than water due to the presence of two electronegative oxygens.

CH₃COO (acetic acid) CH₃COO^- (acetate ion) + H⁺ ; pK_a = 4.8;

Despite the differences between acids and bases, pK_a can also be used to quantify the relative strength of bases. Notice that while the pK_a values for acids are generally less than 7, pK_a values for bases are usually greater than 7. For instance, ethanolamine has a pK_a of 9.5 and is thus half protonated (charge = +1/2) when the pH is 9.5 (Figure 4). An important exception to this rule of thumb is the amino acid histidine, which is a base but has a pK_a around 6.0.