James Richard Fromm
From a consideration of the intrinsic strengths of acids and bases themselves we now proceed to consideration of their effective strength in various aqueous solutions. This effective strength in solution, or acidity, will vary with both the intrinsic strength of the acid or base itself and the amount of the acid or base which is present in the aqueous solution.
The qualitative ranking of solutions as "more acidic" or "more basic" is possible on an empirical basis of taste and qualitative tests, but such an approach is not useful for quantitative discussion. Therefore the concept of acidity has been defined quantitatively as follows: the acidity of a solution is measured by and is equal to the activity of hydrogen ion in that solution.
The acidity of a solution was originally viewed as the molar concentration of H3O+ present in it. This idea was used by the Swedish chemist S. P. L. Sorenson when in 1909 he defined the acidity of a solution expressed in logarithmic notation as pH = -log[H3O+].
The value of the solution pH is a quantitative measurement of the acidity of a solution. Although the analogous concept of basicity is much less commonly used, the pOH notation is sometimes convenient. Only one scale need be used, since they are related through the ion product of water:
[H3O+][OH-] = Kw = 1 x 10-14
in molar concentration units. Taking the logarithm of both sides,
log[H3O+] + log[OH-] = logKw = -14.
-log[H3O+] + (-log[OH-]) = -logKw = +14, pH + pOH = pKw = 14
Sorenson later modified his concept of pH after the idea of chemical activity was introduced, and now pH is defined as
pH = -log a(H3O+).
In dilute aqueous solution, the original and modified definitions are effectively the same.
Since the concentration of protons can change over several orders of magnitude, it is convenient to use a logarithmic acidity scale. In water we find an actual range of pH from about 0 to 14, although in solutions containing high concentrations of strong acids or bases this range can be exceeded somewhat. It is convenient to classify aqueous solutions according to their pH. If the pH of a solution is less than 7, the solution is called acidic; if the pH is about 7, the solution is neutral; if the pH is greater than 7, the solution is is called basic. In an acidic solution, then, the concentration of hydrogen ions is greater than the concentration of hydroxide ions. In a neutral solution, the concentrations of hydrogen ions and hydroxide ions are roughly equal. In a basic solution, the concentration of hydroxide ions is greater than the concentration of hydrogen ions.
The solvent water itself can act as a base, and the autoionization constant reflects this. Thus the product of acidity and basicity, if basicity is logically defined as the activity of hydroxide ion, is a constant for any aqueous species, so that as the acidity of a solution increases (solution becomes more acidic) the basicity of that solution decreases (solution becomes more basic). This is what would be expected from the qualitative understanding of acidity or basicity. Since acidity and basicity are related through a constant in water (and they are in all other protonic solvents as well; only the numeric value of the constant changes with solvent) we need be concerned with only one of them -- acidity.
The ionization constant of an acid or base is a measure of the intrinsic strength of the acid or base itself in a particular solvent. On the other hand, acidity is a measure of the activity of protons in a particular solution. Equal acidity can be obtained from a higher concentration of a weaker acid or a lower concentration of a stronger acid.