The usual measure of the strength of an acid is its acid dissociation constant (), which can be determined experimentally by titration methods. Stronger acids have a larger  and a smaller logarithmic constant () than weaker acids. The stronger an acid is, the more easily it loses a proton, . Two key factors that contribute to the ease of deprotonation are the polarity of the  bond and the size of atom A, which determine the strength of the  bond. Acid strengths also depend on the stability of the conjugate base.
While the  value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water or DMSO), the tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weak aniline base) is measured by its Hammett acidity function, the  value.  Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the  and  values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water ( = 3.2) or DMSO ( = 15), has  values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an  value of –15, making it a more strongly protonating medium than 100% sulfuric acid and thus, by definition, a superacid.  (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its  value ( < –1.74).  This usage is consistent with the common parlance of most practicing chemists.)
When the acidic medium in question is a dilute aqueous solution, the  is approximately equal to the pH value, which is a negative logarithm of the concentration of aqueous  in solution. The pH of a simple solution of an acid in water is determined by both  and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH  value is a better measure of acidity than the pH.