This Demonstration simulates the titration of an acid with a strong base, such as NaOH. As base in the pipette is added to the acid in the beaker, the of the solution is monitored by a meter or by use of an indicator. The strength of the acid is determined by the acid dissociation (or ionization) constant , more conveniently expressed as . For a strong acid, (or negative). A number of weak acids, for example acetic acid, have in the neighborhood of 5. A diprotic acid, such as or , is characterized by two acid constants and that pertain to successive ionizations producing the two hydrogen ions. For simplicity, the initial concentrations of the acid and base are assumed to be equal. Thus a diprotic acid requires twice the volume of base for neutralization.[more]
A neutral solution has [ or . (Most of the hydrogen ions actually exist in aqueous solution as hydronium ions ). Around an equivalence point, where stoichiometric neutralization is approached, the versus volume curves increase steeply. The inflection point of the curve marks the precise location of the equivalence point. The progress of the titration automatically slows down in the Demonstration so that volume can be determined more accurately.
A weak acid exhibits a relatively flat curve to the left of the equivalence point. This provides the basis for buffer solutions, which maintain a nearly constant over a range of concentrations. This behavior can be approximated by the Henderson-Hasselbalch equation . Thus a weak acid is most effective as a buffer around .
Before electronic meters were developed, indicators were commonly used to follow the progress of neutralization. Each indicator undergoes a characteristic color change around a specific value. Five common acid-base indicators are available in the Demonstration.[less]
Snapshot 1: titration curve for strong acid-strong base neutralization
Snapshot 2: titration curve for diprotic acid
Snapshot 3: buffer region for weak acid with
Reference: P. W. Atkins and Julio de Paula, Atkins' Physical Chemistry, 8th ed., New York: Oxford University Press, 2006 pp. 240–245.