In the Laboratory
A Variety of Electrochemical Methods in a Coulometric Titration Experiment A. Lötz Institut für Physikalische Chemie der Universität München, Munich, Germany
Introduction and General Description of the Experiment Coulometric titration represents a direct application of Faraday’s laws of electrolysis and is one of the methods that appear most frequently in instrumental analysis courses (1). Whereas most earlier articles on coulometric analysis in this Journal (2–20) concentrated on single methods applied to single substances, this note presents an experiment that utilizes a variety of methods of coulometric titration and electrochemical endpoint indication applied to a mixture of KI and HCl. The experiment, which was introduced in our practical course in physical chemistry some years ago, uses common laboratory equipment of moderate cost. The coulometric current source can easily be constructed from low-cost, readily available electronic components. The iodide content of the mixture of KI and HCl is determined by titration with bromine generated at the anode from bromide: 4 Br { → 2 Br 2 + 4 e{ 2 I { + Br 2 → I 2 + 2 Br { I 2 + Br 2 + 2 Br { → 2 I Br 2{ There is also a direct oxidation of I { leading to the same turnover of electrical charges for the same amount of I{ consumed and IBr2{ produced: 2 Br { → Br 2 + 2 e{ 2 I { → I 2 + 2 e{ I 2 + Br 2 + 2 Br { → 2 I Br 2{ The titration is followed with a twin polarizable platinum electrode whose indication can be understood qualitatively from the amount of the easily reducible I 2 and (after the endpoint) Br2 present during the titration (Fig. 1, top) (21). The electrolysis current of the twin polarizable platinum electrode is smaller than that of the coulometric electrolysis by approximately a factor of 1000, so the change in concentrations induced by the indicator electrode pair is negligible. Analysis of the chloride content is performed indirectly by titration of the acid at the cathode: and
2 H2O + 2 e{ → H2↑ + 2 HO{ 2 H+ + 2 e{ → H2↑
The indicator is a glass electrode combined with a Ag/AgCl reference electrode (Fig. 1, center). Finally I { and Cl { are determined simultaneously by anodic dissolution of a silver wire: Ag + X { → Ag X ↓ + e{ ; X = I, Cl
The course of the titration is recorded with a silver electrode (Fig. 1, bottom). A source of confusion with coulometric titration using electrochemical detection is the presence of two pairs of electrodes. The first pair is used for the generation of the titrating species—for example, Ag+, in the method presented last. Actually, the above chemical reactions take place at only one electrode of the pair, the working electrode, whereas the second electrode, the counter electrode, is necessary for closing the current circuit. The counter electrode is placed in a separate compartment with electrolytic connection to the working electrode in order to avoid interference with the reactions at the working electrode. Depending on whether the counter electrode is the cathode or anode, hydrogen from an acidic solution or oxygen from an alkaline solution is generated at the counter electrode in the present experiments. Whereas the working and counter electrode carry a current of approximately 10 mA, much smaller currents flow through the pair of electrodes that dip into the solution surrounding the working electrode and monitor the course of the titration. In the case of potentiometric measurements (determination of Ag+ and H+ concentrations, in the present experiments) the current through the monitoring electrodes is virtually zero, and is very low (in the µA range) in amperometric measurements (e.g. monitoring the concentration of I 2). Some Details of the Experiment The students get a mixture of HCl and KI whose absolute and relative content is unknown to them (20 ± 5 mL of each of two 0.01 M solutions), dilute it to 100 mL, and take 20-mL aliquots. These are filled to approximately 50 mL with the chemicals listed in the third column of Table 1. In each of the three titrations, the working electrode and the counter electrode dip into solutions in different beakers connected with a salt bridge filled with 1 M KNO3 solution. The ends of the bridge are plugged with agar-agar gel up to a length of 1 to 2 cm (3% solution in hot water, solidified in the bridge). The use of a salt bridge avoids the problem of diffusion between the anodic and cathodic compartments. The bridges are stored in KNO3 solution in separate containers for the three titrations (1-L plastic containers that solid chemicals are delivered in) and can be reused until they lose their plugs by accident. In the titration of I { and Cl { with Ag+, a straight silver wire of 0.5 mm diameter and 5 cm length is used as the working electrode, of which approximately 2 cm dips into the solution. It can be reused once or twice after being cleaned by drawing across sandpaper. The potential of the silver electrode for the endpoint detection (not to be confused with the silver wire) is measured against a glass electrode because the KNO3
JChemEd.chem.wisc.edu • Vol. 75 No. 6 June 1998 • Journal of Chemical Education
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In the Laboratory
I = 10.6 mA
Current / µA
50
40
30
20
10
100
200
300
400
500
600
700
800
Time / s
Constant-Current Source and Other Electronic Equipment
I = 10.6 mA
10 9
pH
8 7 6 5 4 0
100
200
300
400
500
600
700
Time / s
I = 10.6 mA
Potential / mV
0
-100
-200
-300
-400 0
100
200
solution of the reference electrode of commercially available combined silver electrodes was often found to contain halides under the conditions of the practical course. The connection of the glass electrode to the pH/mV-meter must be made with a shielded coaxial cable whose shield is at the potential of the silver electrode in our setup. The twin polarizable platinum electrode used for the bromometric titration is stored dry. The other platinum electrodes used in the experiment are cylindrical, approximately 0.5 cm in diameter and 0.3 cm in length. All are commercially available. Further details of the experiment are collected in Table 1.
300
400
500
600
700
800
900
Time / s
Figure 1. Coulometric titration of a KI/HCl mixture. Thin lines: first derivatives of the thick titration curves. Dashed lines: end points. Top: bromometric titration of I { with a twin polarizable platinum electrode as indicator. Center: acidimetric titration of HCl with a monitoring glass electrode. Bottom: titration of I { and Cl { with a silver electrode. Performance: 20 mL 0.01 N K I and HC l analyzed; found: (top) 19.9 ± 0.1 mL K I, (center) 19.9 ± 0.2 mL HC l, (bottom) 19.7 ± 0.3 mL K I and 20.5 ± 0.5 mL HC l (40.2 ± 0.5 mL halide). Errors calculated from estimated errors of the end points.
The electrolysis current in our experiment is fixed to 10.6 mA, which results in an electrolysis time of about 6 min for the acidimetric titration and twice as much for the other two titrations. The constant-current source (Fig. 2, upper part) runs between +15 and {10 V for a wide regulation range. The current is boosted by a transistor, since in our experience a common 741 operational amplifier is too much stressed by delivery of 10 mA during longtime use. The voltage of 140 mV for the twin polarizable electrode is obtained from the circuit of Figure 2, lower part, whose feedback resistor of 10 k Ω transforms 1 µA of indicator current to 10 mV at the output. This circuit has its own power supply, which means that the twin polarizable electrode is floating with respect to the electrodes of the constant-current circuit. Data collection during the titration was first done by reading from the pH/mV-meter in regular intervals. We later connected a recorder to the pH/mV-meter, but found that the students were so much engaged with handling the recorder that the titration itself threatened to become an appendage of the recorder. The situation improved greatly when the recorder was equipped with a switch by which preset bias and span for each of the three titrations could easily be established. We now run the experiment with a Toshiba T2100 laptop computer appropriately programmed for the experiment in Visual Basic. This is how the three parts of Figure 1 were produced. Since the pH/mV-meter that we use did not have a digital data output, the computer was equipped with a ADconverter interface (PCM-DAS16D/12, ComputerBoards Inc., Mansfield, MA). Severe hum on the analog signal line due to the different floating potentials was suppressed by a 330-k Ω /1-µ F filter incorporated in a connector of the signal line.
Table 1. Electrodes and Solutions for Coulometric Titration of KI/HCl Mixtures Beaker 1 Ion I{ {
+
Working Electrode
Electrolytea
Pt (anode)
HNO3b + KBr
Pt/Pt
Counter Electrode/ Electrolyte Pt/HNO3b
H (Cl )
Pt (cathode)
H2O + KCl
glass/Ag, AgCl
Pt/1 M KOH
I {, Cl {
Ag wire (anode)
HNO3b + Ba(NO3)2
Ag/glass
Pt/HNO3b
a20
776
Beaker 2 Monitoring/ Reference Electrode
mL KI/HCl mixture + 30 mL liquid listed + 2 g salt listed. bConcn 0.1 M.
Journal of Chemical Education • Vol. 75 No. 6 June 1998 • JChemEd.chem.wisc.edu
In the Laboratory
Concluding Remarks As Figure 1 shows, the performance is good, the only difficulty being the uncertainty in the location of the first endpoint in the simultaneous titration of I { and Cl {, also well known from the corresponding volumetric titration. We believe that this experiment with its variety of electrochemical methods and its internal check of the consistency of the results is a valuable addition to the range of experiments performed in the course. Acknowledgments I thank Mr. Horbach and Mr. Bachmeier for assembling the electronic equipment. The very useful comments of the three referees are gratefully acknowledged. Literature Cited 1. Harris, H. H.; O’Brien, J. J. J. Chem. Educ. 1992, 69, A266–A269. 2. Reilley, C. N. J. Chem. Educ. 1954, 31, 543–545. 3. Van Lente, K. A.; Van Atta, R. E.; Willard, H. H. J. Chem. Educ. 1959, 36, 576–578. 4. Head, W. F.; Marsh, M. M. J. Chem. Educ. 1961, 38, 361–362. 5. Van Lente, K. A. J. Chem. Educ. 1966, 43, 306–307. 6. Evans, D. H. J. Chem. Educ. 1968, 45, 88–90. 7. Stock, J. T. J. Chem. Educ. 1968, 45, 736–738. 8. Vincent, C. A.; Ward, J. G. J. Chem. Educ. 1969, 46, 613–614. 9. Stock, J. T. J. Chem. Educ. 1969, 46, 858–860. 10. Beilley, A. L.; Landowski, C. A. J. Chem. Educ. 1970, 47, 238–239. 11. Tackett, S. L. J. Chem. Educ. 1972, 49, 52–54. 12. Stock, J. T. J. Chem. Educ. 1973, 50, 268–269. 13. Marsh, D. G.; Jacobs, D. L.; Veening, H. J. Chem. Educ. 1973, 50, 626–628. 14. Lieu, V. T.; Kalbus, G. E. J. Chem. Educ. 1975, 52, 335–335. 15. Muha, G. M. J. Chem. Educ. 1976, 53, 465–466. 16. Bell, D. A. J. Chem. Educ. 1978, 55, 815–815. 17. Grimsrud, E.; Amend, J. J. Chem. Educ. 1979, 56, 131–133. 18. Greenspan, P. D.; Burchfield, D.E.; Veening, H. J. Chem. Educ. 1985, 62, 688–690. 19. Bertotti, M.; Vaz, J. M.; Telles, R. J. Chem. Educ. 1995, 72, 445–447. 20. Swim, J.; Earps, E.; Reed, L. M.; Paul, D. J. Chem. Educ. 1996, 73, 679–683. 21. Wooster, W. S.; Farrington, P. S.; Swift, E. H. Anal. Chem. 1949, 21, 1457–1460. Figure 2. Top: constant current (~10 mA) source for the coulometric titration. Bottom: 140 mV supply for the twin polarizable electrode and conversion of the indicator current to voltage.
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