pH is the negative decadic logarithm of hydrogen ion (H + ) activity in a solution. Water dissociates into equal amounts of H + and OH − , the reaction is neutral (10 −7 mol H + per kg water, pH 7.0 at 25°C, 6.8 at 37°C). Acids (Ac) dissociate to H + and the negatively charged rest of the acid (Ac − ) or conjugate base, bases to OH − and the positively charged base rest or conjugate acid. Strong acids or bases dissociate nearly completely. The tendency to dissociate is smaller in weak acids and bases because of the molecular structure and distribution of charges; it is described by the dissociation constant K = ([H + ] × [Ac − ])/[Ac], pK being the negative logarithm. If pH = pK, the acid is half-dissociated. This can be artificially obtained by mixing one part of a weak acid (e.g., CH3COOH, acetic acid) with one part of its salt with a strong base (e.g., Na + CH3COO − , sodium acetate), since salts are fully dissociated in.
