Potentiometric titration

In analytical chemistry, potentiometric titration is a technique similar to direct titration of a redox reaction. It is a useful means of characterizing an acid. No indicator is used; instead, the electric potential is measured across the analyte, typically an electrolyte solution. To do this, two electrodes are used, an indicator electrode (the glass electrode and metal ion indicator electrode) and a reference electrode. Reference electrodes generally used are hydrogen electrodes, calomel electrodes, and silver chloride electrodes. This reference electrode forms an electrochemical half-cell, typically serving as the anode. The indicator electrode forms the other half-cell with the ions of interest in the test solution.

The overall electric potential is calculated as

E_{ m {cell}}=E_{ m {ind}}-E_{ m {ref}}+E_{ m {sol}}.

$E sol$ is the potential drop over the test solution between the two electrodes due to differences in ion diffusion rates across the salt bridge, which is typically negligible. $E cell$ is recorded at intervals as the titrant is added. A graph of potential against volume added can be drawn and the end point of the reaction is when the jump in voltage is observed, similar to other types of titrations. $E cell$ depends on the concentration of the ions of interest with which the indicator electrode is in contact. For example, the electrode reaction may be

{\ce {M}}^{n+}+n\ {\ce {e- -> M}}

As the concentration of $M$ changes, the $E cell$ changes correspondingly according to the Nernst equation. Thus the potentiometric titration involve measurement of $E cell$ with the addition of titrant. Types of potentiometric titration include acid–base titration (total alkalinity and total acidity), redox titration, precipitation titration, and complexometric titration (e.g. free EDTA).