John H. Kennedy
and Frank Adamo University of California Santa Barbara, 93106
I
~ontroliedPotentid Ilertrolysis An experiment for elementary quantitative analysis
O n e of the most fundamental electrochemical techniques is that of controlled potential electrolysis. Students are taught that if an electrode is maintained at a constant potential versus a reference potential, a specific electrochemical reaction may take place, and all of the current which flows is due to this reaction. When the solution has been exhaustively electrolyzed a t this potential, the total electrical charge is given by and the charge, Q, is related to the amount of electroactive substance present by Faraday's Law. The technique can be used for quantitative analysis if "n," the number of electrons transferred, for the reaction is known, or can he used for determining n if the number of moles present is known. However, in the laboratory the technique is rarely used especially in elementary courses. The reason for this situation is basically cost and availability of instrumentation. A commercial instrument capable of controlling working electrode potential versus a reference and electronic integration of the current costs several thousand dollars. Also, instrument time runs high since. each electrolysis may require 30-60 min. A class of 12-15 students would need at least 3-5 instruments and preferably one instrument for each pair of students. A breakdown of the requirements consists of three units, (1) variable power supply, (2) potential control, (3) integration of current. Inexpensive power supplies are readily available, and from a pedagogical point of view, manual potential control is ideal. This was accomplished in our laboratory by using pH meters, available for acid-base titration, for monitoring the working electrode potential versus a calomel reference. The student would adjust the power supply periodically to maintain the correct potential. By using fairly low concentrations to limit initial currents to approximately one mA this manual control was not difficult, and all students were able to control the potential easily to =t20 mV. Alternatively, a simple and relatively inexpensive potentiostat, which has been recently reported,' may be used to control the working electrode potential automatically. The third unit, an integrator, was obJACKETP, S. L., AND KNOWLEB, J. A,, J. CAEM.EDUC.,43, 428 (1966). a The E-cell Data-Stor Units (approximately $14 each) and the EDR-300 (approximately $375) are available from the Bissett-Berman Carp., 2941 Nebraska Ave., Santa Monica, California.
tained with the use of a new commercial product, an E-cell Data-Stor Unit combined with an EDR-300 Unit for reading the c~ulometer.~ The E-cell Data-Stor Unit is a small silver coulometer. Coulombs of electric charge are tabulated by plating silver on a gold electrode during the controlled potential electrolysis. The silver is then stripped from the gold using a pulse discharge method, and the charge is read directly on a counter in millicoulombs (mA-sec). The pulse discharge method is more accurate for integration than a constant current readout. Although an EDR-300 Unit is somewhat costly, only one is required for the lab since each student would have his own Data-Stor Unit and would read it when convenient after the electrolysis was completed. The experiments described below require 30-60 min of electrolysis time, but reading out the E-cell requires only 5-15 min depending on the amount of material in the sample. Some of the early Data-Stor Units tested were very sluggish requiring longer readout times, but the newer units were far superior. The small sample size, convenient for the E-cell and for easy potential control, also demonstrates to the students that an accurate analysis can be performed on less than a milligram of material. The other experiments in the course usually required 10&1000 mg samples. Controlled potential electrolysis can be used for either anodic or cathodic reactions, hut the requirements are different. Anodic oxidations were accomplished using a platinum working electrode, hut had the advantage that air could be present. Slightly higher results (an increase of about 1-2% in some cases) were obtained when air was excluded during Fe(I1) oxidation since Fe(I1) was slowly air oxidized. The disadvantage was cost of platinum since the speed of the electrolysis depends on electrode surface/solution volume ratio. About 1 in.? of platinum was used as the working electrode., Cathodic reductions were accomplished using a mercury pool working electrode with a nitrogen purge. Since nitrogen was not readily available in the lab, student results are for anodic oxidations only. However, in courses where polarography is discussed, controlled potential electrolysis at a mercury pool complements this technique, and successive electrolyses of mixtures might be possible, particularly if the working electrode potential is automatically controlled. Certainly there are many other possible examples, but the ones chosen to illustrate the approach are oxidation of iron(II), iodide, and arsenic(III), plus reduction of lead(I1) and cadmium(I1). Volume 47, Number 6, June 1970
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POWER SUPPLY
pH METER
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Diagram of oloctricd circuit and cell used in the axperimsnts
Experimental Apparofus
The figure illustrate the electrical circuit and the cell used for controlled potential electrolysis. The electrolysis cell was oonstmcted from a 30-ml beaker and a 10-mm sealing tube with a medium porosity fritted disc. Approximately 25 mm from the fritted disc, the tube was attached to the side and near the bottom of the beaker, which served as the working electrode compartment. The side tube was bent upright about 25 mm on the other side of the fritted disc to form the auxiliary electrode compartment. Before using, the side tube was partially filled with a 4% agar-saturated potassium chloride gel. For cathodic electrolysis, a polyethylene cap was used to cover the working electrode compartment to prevent oxygen interference. The solution in the working electrode compartment was stirred with a '/s-in. X in. magnetic stirring bar. A l-in. length of 16-gauge platium wire was used rts the auxiliary electrode, while a saturated calomel electrode was used as a reference. The working electrode for anodic electrolysis was constructed from a 0.002 in. X 0.5 in. X 2.0 in. platinum foil spot-welded to a 2-in. length of 16-gauge platinum wire to form a cylindrical-shaped electrode. For cathodic electrolysis, approximately 4 ml of triple distilled mercury connected to the circuit by a 2-in. length of platinum wire sealed in glass served as the working electrode. A pH meter was used to measure the working electrode potential versus the S.C.E. The working electrode potential was manually controlled by adjusting the output of a regulated power supply (input 60-400 cps, output 6-32 V, 0-2 A) with a 10-turn potentiometer. To measure the amount of charge collected during the electrolysis, an E-cell integrator, Model 302-0002, was incorporated into the circuit. An EDR-300 was used to discharge the E-cell coulometer, thereby displaying the number of millicoulombs collected. Procedures The Anodic Electrolysis of Iodide. In this experiment, each student received an accurately weighed sample containing 200500 mg of reagent grade potassium iodide in a clean 50-ml bes?ner. The student quantitatively transferred the sample to a
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500-ml volumetric flask and diluted to volume. Alternatively, a known aliquot of a stock solution of potassium iodide could have been given to each student. After rinsing the electrolysis cell and the electrodes, each sturlenr plnred the elcrrn~defI n the sppropriare ir,nlpartments. Appmxinmte.y Id ml of (1.5JI HAO, ,l~pportingrlertrol\.re was added 10 thc workiw elerrrode rompsrtmenr. To minimize the time required for ~ o & ~ l eelectrolyks, te it was important for the working electrode to be placed as close to the bottom of the compartment as possible, with just enough supporting electrolyte added to cover the electrode. The auxiliary electrode compartment was filled to near the top with the same electrolyte solution. After discharging an Ecell and inserting it in the circuit, each student proceeded to electrolyze the supporting electrolyte st a working electrode potential of +0.70 + 0.02 V versus S.C.E. for 5 min. The E-cell was then disconnected from the circuit and discharged on the EDR-300 instrument. If the reading was above 1.0 millicoulomb, which indicated impurities, the blank was electrolyzed for another 5 min. or until the coulometer cleared with less than 1.0 mC. A 1-ml aliquot of the unknown iodide solution was then added to the working electrode compartment which was electrolyzed at a working electrode potential of +0.70 i 0.02 V versus S.C.E. for 30 mi". The electrolysis was discontinued and the number of millicoulombs collected was determined by discharging the E-cell coulometer. The electrolysis was continued for successive Frmin periods until the amount of charge collected in a 5-min period was less than 1.0 mC. Since students could not standardize the rate of stirring, a definite electrolysis time could not be established. Thirty minutes was usually sufficient, but the extra 5-min period showed the student that the solution was indeed free of iodide and gave him confidence that the electrolysis was completed. The amount of charge to complete the electrolysis was determined by adding up the total number of millicoulombs collected. The electrolysis was generally repeated two times by electrolyzing an additional 1-ml aliquot of the unknown, which was added to the existing electrolyzed solution. Students reported total milligrams of potassium iodide in the sample. The Anodie Electrolysis of Arsenie(1II). A stock solution of primary standard AslOs was prepared by dissolving shout 2 g of the reagent, accurstely weighed out, in a 250-ml volumetric flask with the addition of 100 ml of distilled water and 20 drops of saturated NaOH solution. The solution was then acidified dropwise with concentrated HzS04, cooled to room temperature, and diluted to volume. Each student received a 2-5-1111 aliquot of the stock solution ins. clean 100-ml volumetric flask, which was then diluted to volume. Following procedure described for the electrolysis of iodide, each student electrolyzed 1.0 M HzSO, supporting electrolyte at a working electrode potential of 4-1.05 & 0.02 V versus S.C.E. for 10 min. If the E-cell was discharged with more than 2 or 3 mC, indicating impurities, the student continued to electrolyze the blank for successive 10-min periods until the Ecell cleared with less than 3 mC. A l-ml aliquot of the unknown arsenic solution was then electrolyzed s t a working electrode potential of +1.05 i00.2 V versus S.C.E. far 45-50 min. The coulometer was discharged and the electrolysis continued for successive 10min periods until the amount of charge collected in a 10-min period was less than 3 mC. The number of millicoulombs was totaled to determine the smount of charge required for complete electrolysis. After the electrolysis of a l-ml aliquot sample of the unknown arsenic solution was repeated two times, the total milligrams of srsenic trioxide was determined. The Anodic Elec1l-olysi.s of I?on(II). A ferrous ammonium sulfate ~nknownsample,~ contsining5-15% Fe, was given to each student. About 0.7 g of the sample was accurately weighed out and transferred to a 250-ml volumetric flask. After adding about 150 ml of distilled water and 5 ml of concentrated HB04, the solution was cooled to room temperature and diluted to volume. Each student then proceeded to electrolyze the 1.0 M HISO, supporting electrolyte and the l-ml aliquot samples of the unknown iron solution at a working electrode potential of +0.90 =t 0.02 V versus S.C.E. following the procedure given for the electrolysis of arsenic (111). The percent of iron in the sample was reported. V e r m u s ammonium sulfate standard samples me available from Thorn Smith, Royal Oak, Mich.
Table
1.
Material analvzed
Procedure Checks for Controlled Potential Electrolysis Experiments
Amt. taken fmC)
Amt. found (mC) . .. Currenttime error^ % Ecell area
Std. dev. No./ (%'n) trials
Beckman Eledroscan Pohtia2 Control
iodide
fi
Manual Potential Control +0.72 264.5 266.4 +0.89 439.0 442.9 643.8 651.7 f1.23 643.8 648.9" f0.79 fflR.fi* +0.75
0.24 0.25 0.18
4 3 3
Four Ecells connected in series to check variability. Solutions deaerrtted with nitrogen. Error based on E-cell results.
The Calhodie Electrolysis of Lead(N) and Cadmim(I1). A solution of reagent grade lead nitrate was prepared by dissolving about 0.2 g of the reagent, accurately weighed out, in a 250-ml volumetric flask and diluting to volume. A reagent grade cadmium bromide solution was similarly prepared by dissolving 0.35 g of the reagent in 500 ml. The electrolysis cell and the electrodes were rinsed with distilled water and wiped dry. After adding about 4 ml of tripledistilled mercury to the working electrode compartment, the platinum wire lead and the polyethylene caplug were inserted. The platinum wire was placed in the mercury pool before the introduction of the supporting electrolyte to prevent the platinum surface from wetting. Ten milliliters of 0.5 M KC1 was added to the working electrode compartment, and the auxiliary electrode compartment was nearly filled with the same electrolyte. After placing the reference and auxiliary electrodes in the appropriate compartments, the solution was purged with nitrogen for at least 10 min before starting the electrolysis and purging was continued during electrolysis. A discharged E-cell was inserted into the circuit, and the supporting electrolyte was electrolyzed at a working electrode potential of -0.60 =t0.02 V venus S.C.E. for precisely 5 min. The electrolysis was repeated for 5-min periods until a constant residual current was ohtained. For a cadmium analysis, the working electrode potential was controlled at -0.80 + 0.02 V versus S.C.E. A 1-ml aliquot of the lead(I1) solution was added to the working electrode compartment and electrolyzed at a working electrode 0.02V versus S.C.E. for precisely 20 min. potential of -0.60 Alternatively a 1-ml sliquot of cadmium(I1) was electrolyzed at -0.80 10.02 Vversns S.C.E. To correct the originalelectrolysis for residual current, the electrolysis was continued for 5-minute periods until a constant residual current was obtained. Since the rate of stirring and the working electrode potential significantly affect the residual current,' it was mential that both of these variables remained as constant as possible throughout the electrolysis. The amount of charge to complete the electrolysis was determined by subtracting the amount of residual current which flowed during the time period from the original values obtained from the E-cell readout. The electrolysis was repeated three or four times by electrolyzing an additional 1-ml aliquot of the metal ion solution.
Results and Discussion
Using a Beckman Electroscan for controlling the potential, the procedures described in the experimental section were checked and also the E-cell integrator was compared with the recorded current-time curve. Results for the iron and arsenic experiments are given in Table 1. There was no significant difference between the E-cell integration and the area computed from the current-time recorder trace. Higher results were obtained for iron when the solution was purged with nitrogen, and this anticipated 1-2y0 air oxidation was taken into account when assigning grades to student results. The iodide procedure was checked using manual potential control, and the results are given in Table 1. Variability in E-cells was checked by using four in series for one electrolysis. They were all within 1% of the true value and the range was only 0.2%. One quantitative analysis class performed the iodide analysis, and the results are given in Table 2. Another class performed the iron(I1) analysis, and the results are given in Table 3. A few students in this class were given the option of carrying out the arsenic(II1) oxidation, and the results are given in Table 4. Students found the iodide oxidation easiest to perform. The arsenic oxidation at +1.05 V was close to Table 2. Amt. K I taken
Student Results for Iodide Oxidation No./
Table 3.
Range of
Error
Student Results for lran(ll) Oxidation
% Iron true
Av. amt. found
Trial values
Av. % Fe Found
Error
(%)
*
'MACNEVINW. M., 1994 (1955).
AND
M C I V ER. ~ D., A d . Chem., 27,
a EDR readouts for trials in millicoulombs showing electrolysis reachine comoletion 60 min 672.6 659.7 673.2 677.2 10 min 15.7 18.5 9.2 7.5 10 min 4.5 6.4 692.8 684.6 682.4 684.7
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Table 4. Asto8 taken (mg)
Student Results far Arsenic(lll) Oxidation
Trial values
Av. AslOs found (mg)
Error
(%)
--
supporting electrolyte &composition and some students had high blanks. The 50-60 min required for ironfI1) oxidation allowed small amounts of iron to be -air oxidized which accounts for the low results. The results of the electrochemical reduction of lead (11) and cadmium(II), given in Table 1, show that the procedure given is suitable for student participation. The variables in the experiment must, however, he more rigidly controlled. The rate of stirring, in particular, greatly affects the residual current. ~
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Conclusion
Oxidation of iodide was the most convenient experiment because of the short time required for complete electrolysis. However, larger working electrodes would enable students to carry out iron(I1) and arsenic(II1) oxidations as easily. The accuracy of an E-cell integrator is to about 1% at the time of writing this paper, but may be improved with time since it is a relatively new commercial product having no inherent errors. The E-cell is in principle a silver coulometer which has been known for high accuracy ever since the experiments of Faraday. Acknowledgment
The authors thank the Bissett-Berman Corp. for a loan of an EDR-300 Pulse Discharge Readout device and the Data-Stor Units.
