Chemistry Laboratory #8: Potentiometric Titration of a Halide Mixture*

*Adapted from "Exploring Chemical Analysis", D. Harris, Freeman, 1997.

February 27, 2001
 

Purpose: The goal of this lab is to determine the composition of a mixture of potassium chloride and potassium iodide by means of a potentiometric titration.
 

Introduction: In many titrations, a visual change in the solution signals the endpoint. But it is possible to determine the endpoint without a visible change. Other types of physical changes can also be used to determine endpoints, such as the physical effects arising from relatively large changes in the concentrations that occur near equivalent points. For example, a pH-meter can be used to determine the endpoint in an acid-base titration by showing where a rapid change in pH occurs. A pH measurement is an example of a potentiometric measurement, that is, the measurement of an electric potential across two electrodes. If the potential at one electrode is a function of concentration of a given species, the voltage across the electrodes will be function of the concentration of that species provided that the potential of other electrode is stable.
 

In this experiment, you will measure the potential across two electrodes. The half-reactions are:

If the solutions have a constant composition, the measured potential should be constant. This is not the case in a titration however. In this experiment you will be titrating a mixture of chloride and iodide ions with a solution of silver nitrate. As the titration progresses, the silver ion concentration will change by many orders of magnitude. This will change the potential of the Ag/Ag⁺ half-reaction. There will be no change in the Cu/Cu²⁺ potential, so the change in the measured potential difference will be entirely due to the changing silver ion concentration.
 

Neither silver iodide nor silver chloride is soluble so there are two precipitation reactions that will take place:

Since the silver iodide is significantly less soluble, it will precipitate from solution first. As the titration progresses, the silver iodide will come out of solution until the iodide is nearly entirely consumed. During the precipitation of the silver iodide, there will be a large excess of iodide in solution (the silver will be virtually entirely removed upon its addition). The actual concentration of silver can be calculated using the K_sp relationship:

Once the iodide is consumed, the silver chloride will begin to precipitate and the silver concentration will be limited by the K_sp of silver chloride:

Since the K_sp of silver chloride is so much higher than that of silver iodide, the concentration of silver ion will be several orders of magnitude greater after the silver iodide endpoint than it was prior to it. This increase in the silver ion concentration will be reflected by a very sharp change in the measured potential between the electrodes. When potential is plotted as a function of volume of silver added, this sharp change in potential signals the first endpoint; from the volume of silver added, the amount of iodide in the solution initially can be determined.
 

The second endpoint is another point at which the silver concentration increases by several orders of magnitude. This occurs when the chloride is completely consumed. After this point, any silver added to the solution stays in solution since there is no longer a suitable ion with which it can form a precipitate. The volume of silver added between the first and second endpoints can then be used to determine the amount of chloride that was present in the unknown.

Procedure:

1) Pour all of your unknown into a 50 or 100 mL beaker. Dissolve in ~20mL of deionized water and pour into a 100 mL volumetric flask. Thoroughly rinse the beaker with several small portions of deionized water and transfer the rinse water into the volumetric flask. Dilute to the mark and mix well.

2) Set-up the electrodes in your titration apparatus as described in lab. Both electrodes must be in contact with the solution. Use a magnetic stir plate to keep the solution well mixed as the titration progresses.

3) Pipet 25.00 mL of the unknown solution into a beaker and begin stirring. Begin the titration by recording the initial volume of AgNO₃ (Warning!!!) in your buret. Titrate with ~1 mL aliquots (record to the nearest 0.01 mL); record the potential between the electrodes after each addition. You do not need to wait more than 15 to 30 seconds for each point. Identify the approximate endpoints (± 1 mL).

4) After both endpoints are identified, empty the beaker into the silver waste receptacle and clean the beaker and electrodes thoroughly.

5) Repeat the titration with another 25.00 mL aliquot, this time using approximately 0.1 mL aliquots within ± 1 mL of the endpoints. Use 1 mL aliquots at other regions of the titration. Again, you do not need to allow more than 30 seconds between data points.
 

Prepare two graphs of your data:

a) potential vs. volume Ag⁺ added, and

b) a first derivative plot of the raw data (dE/dV vs V_av)

Prepare a formal written lab report that includes a discussion of the theory of the experiment (e.g., why potential varies with concentration, what chemical reactions are taking place, etc.) as well as your graphs and results of the mixture analysis.