Traditionally, analytical chemists have relied on the classical method of titrimetry to determine the composition of substances.
These methods require no specialized equipment beyond a burette, beakers, heating, and filtration apparatus, and yet they can
outperform many spectrophotometric methods in terms of precision. A titration can often yield a result which is both precise
and accurate to 4 significant figures.
The idea behind a titration is that a reagent of precisely known concentration
(the titrant) is slowly added to a known amount of
an analyte until some event occurs which signals the end of the reaction. A species which is deliberately added to produce such
a signal is called an indicator. When the signal occurs, the volume of added titrant is recorded. Knowing the
the reaction, and both the volume and concentration of titrant, the composition of the analyte can be found.
Classical methods of analysis rely on the chemist's ability to manipulate equilibria. There are four basic
types of equilibria which are used in titrimetry, namely acid-base, solubility, complexation and redox reactions. The common example of a titration is the
strong-acid-strong base titration using phenolphthalein as an indicator. The signal which accompanies the end of the reaction in
this case is a change from a colorless to a pink solution. Ideally this signal, called the endpoint, occurs at exactly the same time
that the stoichiometry of the reaction is satisfied , i.e. at the equivalence point. A huge number of indicators are known and
tabulated, many of which rely on very clever applications of one of the four types of equilibria.
The type of titration we will focus on in this session is one that uses an electrode as an
indicator. Electrodes can be made to be sensitive to one species
only. For example, the pH electrode responds to changes in [H+], and the fluoride electrode responds
to changes in [F-]. The fluoride electrode is basically a metal electrode which is separated from the analyte solution by a crystal
of a substance in which F- ions are mobile. A high concentration of
F- in the solution induces more of the ions to move closer to
the electrode, with a corresponding change in potential.
Such electrodes are available for nitrogen oxides, oxygen, phosphate, halides, glucose and a host of other species. These
electrodes can be made to directly measure the concentration of a species in solution via the Nernst equation (below). These
other electrodes function in basically the same way as a pH electrode
A + e- = B ; E° = 0.00 V
A simple chloride ion-selective electrode can be constructed by coating a silver wire with a layer of silver chloride. The two
relevant equilibria at this electrode are:
Ag+(aq) + e- = Ag(s); E° = 0.800 V
AgCl(s) = Ag+(aq) + Cl-(aq); Ksp = 1.78 x
These two equations can be combined to yield a third half-reaction :
AgCl(s) + e- = Ag(s) + Cl-(aq) E° = 0.222 V
The Nernst equation for this half-reaction is as follows:
E = 0.222V - 0.059 log[Cl-]
Thus, this electrode's potential will be determined by the amount of chloride in the solution. Of course, the potential will be
influenced by other species in solution, especially if their silver salts have
Ksp values smaller than silver chloride.
To use this electrode to measure chloride concentrations directly would require a series of standards and a calibration curve.
However, this electrode can be used as an indicator electrode in a titration without any special calibration procedure. This
advantage is due to the fact that we will be looking for the largest change in the potential, rather than the absolute value of the
potential. Another advantage is the fact that we will be able to determine the
equivalence point itself from the data, rather than
In this experiment, we will attempt to determine the identity of an alkali metal-chloride salt by potentiometric titration. A known
mass of unknown will be titrated against a solution of Ag+(aq). The electrode potential will be plotted vs. the volume of titrant
added, and the endpoint will be determined. The data will look similar to a normal pH titration curve.
Fig. 1 Typical potentiometric titration curve.