Electrochemistry

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EXPERIMENT V

POTENTIOMETRIC TITRATION

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INTRODUCTION

Many Acid-Base titrations are difficult to accomplish using a visual indicator for one of several reasons. Perhaps the analyst is color-blind to a particular indicator color change; there may not be a suitable color change available for a particular type of titration or the solutions themselves may be colored, opaque or turbid. It may be desired to automate a series of replicate determinations. In such situations, potentiometric titration, using a glass hydronium ion selective electrode, a suitable reference electrode and a sensitive potentiometer (a pH meter) may be advantageous.

THEORY

Any acid-base titration may be conducted potentiometrically. Two electrodes, after calibration [to relate potential in millivolts (mV) to a pH value] are immersed in a solution of the analyte. One is an indicator electrode, selective for H3O+ and the other a stable reference electrode. The potential difference, which after calibration is pH, is measured after the successive addition of known increments of acid or base titrant.

When a potentiometric titration is being performed, interest is focused upon changes in the emf of an electrolytic cell as a titrant of known concentration is added to a solution of unknown. The method can be applied to all titrimetric reactions provided that the concentration of at least one of the substances involved can be followed by means of a suitable indicator electrode. The critical problem in a titration is to recognize the point at which the quantities of reacting species are present in equivalent amounts. The titration curve can be followed point by point, plotting as ordinant, successive values of the cell emf (pH) vs the corresponding volume of titrant added. A typical titration curve is presented in Figure 2. Figure 3 represent another method for determining the equivalence point from the titration curve data. Table I, in Appendix I, presents typical data obtained from a potentiometric titration.

The Reference Electrode

Most commonly, the reference electrode is the silver/silver chloride electrode. The potential is based on the following equilibrium:

AgCl(s) + e [pic] Ag(s) + Cl-(aq)
The half cell is:
Ag[(AgCl(Sat'd),KCl(xM)][pic]

NOTE:Electrodes respond to the activity of the electroactive species in solution. However, as a practical matter it is more convenient experimentally to use concentration. For this reason the discussion in this laboratory experiment will be made in terms of concentration rather than the more correct activities.

The practical version of this electrode is a silver wire dipping into a saturated solution of KCl; when fabricated this way its electrode potential is 0.199V (vs. Normal Hydrogen Electrode, NHE) @ 25oC. The potential is a function of temperature and the concentration of KCl in the solution. Such an electrode is comparatively rugged, reliable, and inexpensive.

The Indicator Electrode

The heart of the glass electrode is a thin glass membrane, specially fabricated to preferentially exchange H3O+. The outside of the membrane is in contact with the analyte solution containing the unknown [H3O+]. The inside of the membrane contacts a hydrochloric acid solution of fixed concentration. A silver wire, coated with AgCl dips into this solution; the other end of the wire is connected to the measuring device. A combination glass electrode with silver/silver chloride reference may be represented as shown in Figure 1 below:

FIG. 1. pH Electrode

The Mechanism of the Response

A change in hydronium ion concentration causes a change in...
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