n ti DONALD 6.DAVIS Department o f Chemistry, Louisiana State University in New Orleans, New Orleans, l a .
B The chronopotentiometry of cerium (IV) has been investigated using a shielded plane platinum electrode designed to prevent disturbance of the diffusion layer by convection. The constancy of i~’/e/C was established over a concentration range of 2 to 20mM and a transition time range of 10 to 100 seconds, Errors caused by the reduction of platinum oxide films on mixtures of cerium(IV) and iron(ll1) were investigated. Indirect methods for the determinations of oxalic acid and arsenic(lll) were developed and the kinetics of the reaction between cerium(lV) and tartaric acid were studied,
has received considerable attention from the theoretical point of view and to gain a qualitative understanding of elec’ trode reactions. However, only a few analytical applications have been actually developed. In many cases a lack of constancy of i+/C has prevented analytical application, although several correction systems have been developed (2, 6). The work of Bard (3) has shown that a shielded plane platinum electrode, such as has been used for diffusion studies, allows accurate chronopotentiometry over a M-ide range of transition times. This study was made to apply chronopotentiometry to the extensive and analytically useful field of cerate oxidimetry (10). Most, if not all, of the indirect determinations of reducing agents accomplished by treatment with excess standard cerium(1V) solution folIowed by titration with a standard reducing agent can also be successfully done chronopotentionietrically. A slight loss in accuracy is compensated by a n increase in speed. The chronopotentiometry of cerium (IV) has been investigated (5, 6). In one case (5) an unshielded gold electrode was used and the other (6) was a brief study which indicated that cerium(1V) was difficult to detelmine using a bright platinum electrode. HRONOPOTENTIOnIETRI’
EXPERIMENTAL
The basic instrumentation and technique of chronopotentiometry have been discussed (7). The constant current source was the same as that previously used (5). Chronopotentiograms were recorded with a Sargent potentiometric
recorder which had a pen speed of 1 second full scale. The cell used in this study has been described by Bard (3). The working electrode, in Bard’s cell, is a platinum disk sealed in soft glass tubing in such a manner that the glass tube provides a shield or mantle around the electrode surface. The electrode is constructed in this way to prevent disturbance of the diffusion layer formed during electrolysis. The electrode can be oriented either up or down to accommodate either an increase or a decrease in density of solution due to electrolysis. It was found that for the reduction of cerium(1V) in both sulfuric and perchloric acids the value of W 2 / C was more constant when the electrode m s oriented up, indicating that the density of the solution of reduction products was greater than the solution being analyzed. -411 further work was done with the electrode oriented for downward diffusion. The projected area of the electrode was 0.958 sq. cm. Contact was made to the platinum disk by pushing a copper wire. coiled in the shape of a spring, against the back side of the disk. The wire was replaced periodically. As in Bard’s cell, a glass tube was attached at the upper end to a rubber bulb to flash the reduction products out the glass mantle of the electrode after each chronopotentiogram was taken. Between 1 and 2 minutes were then allowed for the solution to become quiescent before the next measurement was made. The auxiliary electrode was a coiled heavy platinum wire enclosed in a compartment separated from the main cell by a medium porosity sintered-glass disk. This compartment v-as filled with part of the solution under investigation. The reference electrode was a standard saturated calomel electrode such as is generally used with a pH meter. With sulfuric acid solutions this electrode was placed directly in the main solution close to the norking electrode. When perchloric acid was used, a larger S.C.E. was constructed that could conveniently be placed in contact with the test solution by means of a salt bridge. When desirable, oxygen was removed by passing nitrogen through the solution for 15 minutes. Traces of oxygen in the nitrogen were removed by a vanadous chloride gas scrubber. The removal of oxygen was not necessary when the transition time for the reduction of cerium only was of interest. -411 measurements were made with the cell placed in a constant temperature 0.1’ C. The stirring bath a t 25.0 motor of the bath was not operated
while measurements were being made to reduce vibration. Reagent grade chemicals were used throughout this work. Cerium(1V) solutions were standardized against primary standard sodium oxalate. Periodically the working electrode was cleaned by soaking with a solution of argentic oxide in 51W nitric acid, followed by soaking with nitric acid alone for about 1 hour. The latter was done to remove any adsorbed silver ions. The first chronopotentiogram after each cleaning was disregarded since it included a large tontribution from the reduction of the heavy oxide film on the electrode surface. Actually cerium(1V) also oxidized the electrode surface but to a lesser extent than silver(I1) solutions. Transition times were measured a t a predetermined potential. This pobential mas selected as the potential a t which the chronopotentiogram becomes linear, as indicated by the intersection of the dotted line and the chronopotentiogram (Figure 1). In the case of the reduction of cerium(1V) in 1.M sulfuric acid transition times mere measured a t a potential $0.70 volt us. S.C.E. The potential a t which the chronopotentiogram becomes linear appeared t o be essentially independent of current density and concentration. Thus T was taken as the number of seconds between t = 0 and the time a t which the potentia! becomes equal to the preselected ~ a l u e . RESULT5 AND DISCUSSION
Chronopotentiometry of Cerium (IV). Chronopotentiograms for the reduction of cerium(1V) in sulfuric, perchloric, or nitric acids were well defined even though apparently due to an irreversible reduction. The constancy of i+I2/C for the reduction of cerium(1V) was established over a concentration range of 2 to 20mJI in both 1M sulfuric Rnd 3JI perchloric acids. The results for sulfuric acid are reported in Table I. The data taken in perchloric acid 11ere characterized by relative standard de\ istions falling in the same range (0.4 to 0.7%1, but the average value of i r 1 , 2 / @was 208. The values of ir‘ C for concentrations beheen 6 and 20mJI are very precise, as slioun by a relative standard deviation of O.32yG. However, the value of i+2/C increases a t concentrations below about 5mIll so that the relative standard deviation for ail concentrations becomes 1.5%. -4n increase in i r 1 ~ * / Cwas also VOL. 33, NO. 13, DECEMBER 1961
1839
UJ
50
.’d 0 . 7 5 c!
U J
uj
> w
0.25-
Ii
3 0
-
SEC TIME
Figure 2. Chronopotentiograrn of 1 O.OrnM cerium(IV) and 3.94mM iron(ll1) in oxygen-free 1 M H&04
\ I 30 SEC.-
0.25 TIME
Figure 1 . in
Chronopotentiogram of: 18.4mM ceriurn(lV)
1M HtS04 Current, 0.6733 ma.; oxygen not removed From solution
noticed when r was less than about 10 seconds. Bard (3) and Reinmuth (9) have considered reasons for this increase. A part can be attributed to the double layer capacity of the electrode but most
Table I. Chronopotentiometry of Cerium( IV) in 1 M HzSOl Transition times measured at $0.70 v. vs.
S.C.E.
20 00 18 40 14 00 10.00 6.00
4.60
4.00 2.00 a Each value of ~ ~ 1 represents ’ 2 average of at least 12 chronopotentiograms taken
at at least 6 different current levels. * Relative standard deviation within each concentration is between 0.4 and 0.77~.
Table 11. Chronopotentiometry of Cerium(lV) Corrected for Platinum Oxide Reduction Transition time measured at $0.28 v. us.
S.C.E.
18.4 10.0 4.6
225
219
239 2 50
227 232
Average of at least four measuremente a t different currents.
B
ANALYTICAL CHEMISTRY
of the effect must be due to something else. The possibility of adsorption of ceric ions on the electrode surface exists ( I ) and this might also cause observed transition times to be too large. Roughness of the electrode has been blamed (3) but in this case probably the beginning of the reduction of the platinum oside film on the electrode makes a measurable contribution to the total transition time. Unfortunately the rate of oxidation of the electrode by cerium(1V) is so rapid that this error cannot be eliminated by using a reduced electrode. In any case the accuracy of rhronopotentiometry under the conditions used here is limited a t low concentrations. The use of a calibration irable for strictly analytical work. With unshielded, vertical plane electrodes a variation of about 8% in i+/C was found (5) when r was allowd to vary between 10 and 60 seconds. If, however, the current density is adjusted such that the transition time is approximately constant no matter what the concentration, results almost as precise as those reported here were obtained. The chronopotentiogram in Figure 1 shows a small potential holdup following the main one due to the reduction of cerium(1S’). The break is due to the combined reduction of oxygen in the solution and the reduction of the platinum oxide film formed on the electrode surface by the action of cerium(1V). When oxygen is removed completely, the size of the transition
Current, 0.451 0 ma.
time for the second reduction was reduced to about one half its former value. It was of interest to test the recently proposed equation of Lingane (8) to see if the contribution of the reduction of oside film could be corrected for if transition times were measured a t $0.25 volt us. S.C.E. rather than at 0.70 volt. Actually this procedure has no practical value since the difference in the reduction potentials of the two reactions in question is sufficient to allow easy measurement of either transition time. The equation proposed by Lingane is i corr. = i obs.
- -Q
where i obs. is the current measured in milliamps, i con. is the current used for the reduction of ceric ion, r is the transition time, and Q is the quantity of electricity in millicoulombs necessary to reduce the oside film. The value of Q must be determined for each electrode by a separate experiment. As can be seen in Table I1 the application of the correction does make the values of i+’*/C more nearly constant but is not nearly sufficient. The value estimated for i+Z/’C a t +0.25 volt is 220 which is in good agreement with the corrected value for 18.4mM. The failure of this method to correct for the effect of oxide film probably can be ascribed to the assumption upon which Equation 1 is based. This assumption is that when two electrode reactions of the types involved here occur at a n electrode, the current efficiency of the main reaction can be thought of as being reduced to a calculable value below 100%. The situation is actually more complex but a rigorous mathematical treatment exists (8). Further considerations of oxide film corrections were reserved for a more practical case of mixtures of cerium(IV) and iron(II1).
c 4.6
c
Table 111. Indirect Determination of Arsenic(ll1) and Oxalic Acid
Arsenic(III), Meq. Oxalic Acid, Meq Taken Found Taken Found
I 4.41
3.50 3.00 2.50 2.00 1.50 1.00 0.70
3.48 2.96 2.50 1.97 1.52 0.96 0.71
3.50 3.00 2.50 2.00 1.50 1.00
3.47 2.99 2.51 2.02 1.50 1.02
3.8 0
E LL
0 2.9
zV
z
0 V
CURRENT, MlLLlAMPS
Figure 3. Concentration of iron(IH) vs. current plots for mixtures of cerium(lV) and Iron(ll1) ~
The solution m s then diluted to a known volume and the amount of cerium(1V) remaining was determined chronopotentiometrically. reducing agent determined acid. In this case the medium used was 3M perchloric acid. No was necessary in this case. The results reported in were calculated on the basis of the values of i P / C from Table I for the determination of arsenic in 1M acid and on the basis of i F / " for the determination of oxali 8iM perchloric acid. Examination of Table I11 indicates the accuracy of this
6.OOmM Ce(lV1 3.94mM Fe(lll) 4.00mM Ce(lV)
e 1.97mM Fe(ll1)
1 O.OmM Ce[lV)
1.97mM Fe(lll)
Chronopotentiometry of Cerium(I1V) Although mixtures of cerium(IV) and iron(II1) could be determined by a series of redox titrations or by controlled potential coulometry (4, the analysis of this mixture was attempted chronopotentiometrically to see if the effect of oxide film could be corrected for in this case. Figure 2 shows a typical chronopotentiogrm of cerium(1V) and iron(II1) in oxygen-free I M sulfuric acid. Section A of the curve is due to the reduction of cerium(1V) and section C is due to the reduction of iron(II1). Just before the reduction of iron the beginning of a small wave, B, can be noticed. This is due to the start of the reduction of the oxide film. Although the reduction of the oxide fiIm starts before the reduction of iron, the greater part of these reactions occur together. Analytically useful results for iron cannot be obtained unless a correction can be made for the effect of oxide film. The complexity of the rigorous correction for the case of a single soluble electroactive substance discourages attempts along this line for two substances. The mathematical complexity of the complete treatment may be avoided by the use of a graphical method (8). A plot of the apparent or found iron concentration vs. current is extrapolated
and Iron(III) Mixtures.
to zero current. The zero current intercept is in agreement with the concentration of iron taken, which would be observed in the absence of electrode oxidation. Figure 3 shows a series of such plots for several different combinations of cerium(1V) and iron(III)a With care the concentration of iron can be obtained with an accuracy of between 1 and 2% by this method but ease of calculation is obtained at the expense of experimental simplicity. Each determination of iron requires several chronopotentiograms to establish the line for extrapolation. Even though this method cannot be recommended for analysis, it is interesting that correction for oxide film reduction can be made even in this relatively complex case. Indirect Methods. Many of the indirect methods for the analy~isof reducing agents proposed by Smith (10) may be carried out chronopotentiometrically more rapidly, and almost as accurately, by back titration with a standard ferrous solution. I n addition the need for a standard solution of reducing agent is eliminated. Two examples of the large number possible were tested. Arsenic(II1) was determined by adding an excess of standard cerium(IV) solution in 1M sulfuric acid and a few drops of 0.01M osmium tetroxide a5 a catalyst.
taric acid. The action have been
rate law as: - d Ce+4 7 = k[Ce+']
[tartaric acid]
The effect of sulfate, cerous, or hy-
VOL 33, NO. 13, DECEMBER
ACKNOWLEDGMENT
The author thanks Daniel Orgeron, who did much of the experimental work, and the National Science Foundation for its financial support.
(2) Anson, F. C.,Lingane, J. J., Ibid., 79, 1015 (1967). (3) Bard, A. J., ANAL. CHEM.33, 11 (1961). (4) Davis, D. G., J . Electroanal. Chem. 1, 73 (1959). ( 5 ) Davis, D. G., Ganchoff, J., Zbid., 1,248
(1960).
LITERATURE CITED
(1) Anson, F. C., J. Am. Chem. Xoc. 83, 2387 (1961).
(6) Gierst, L., Mechelynck, P., Anal. Chim. Acta 12, 79 (1955).
( 7 ) Lingane, J. J., “Electroanalytical Chemistry,” Chap. 22, 2nd ed., Interscience, New York, 1958.
(8) Lingane, 9. J., J. Eleclroanal. Chem. 1, 379 (1960). (9) Reinmuth, W. B., AKAL. CHEM.33, 486 (1961). (10) Smith, G. F., “Cerate Qxidimetry,” G. .F. Smith Chemical Co., Columbus, Ohio, 1942.
RECEIVEDfor review July . 1961. Accepted October 2, 1961. Division of Analytical Chemistry, 140th Meeting, ACS, Chicago, Ill., September 1961.
Direct Amperometry of Cyanide at Extreme Dilution JAMES A. McCLOSKEYl U. S. Army Chemical Corps Biological Laboratories, Fort Detrick, Frederick, Md.
b Direct amperometric response of a rapidly rotating silver anode in 0.03M N a O H has been used for determination of cyanide ion in extremely dilute solutions. The lower limit of linear response is about 1 X l o 4 gram of cyanide for 7-ml. samples or 5 X 10-10 gram for 0.5-pI. samples introduced into 1Q-ml. of supporting electrolyte. Interferences were measured from a thousandfold excess of each of 76 unions.
A
of solid electrode amperometry over conventional polarography for measuring very small diffusion currents have been discussed (4). In addition, the use of two solid polarized electrodes offers greater compactness and versatility in many casey. However, few applications of amperometry to the analysis of submicromolar solutions have been reported. Employing a method developed by Potter and White (6),Barkley and Thompson (9) determined iodine in sea water in concentrations down to 5 pg. per liter with a standard deviation of 2.1 pg. per liter. A potential of 0.3 volt was applied across two platinum electrodes and the resulting current flow was measured by determining the potential across a resistor in series with the electrodes. Baker and Morrison (1) determined microgram quantities of cyanide by measuring the “spontaneous electrolysis” current flowing between a silver anode having a large surface area and a platinum cathode a t zero applied potential. The present study demonstrates the direct response of a system composed of a rotating silver anode and a stationary platinum cathode to extremely small amounts of cyanide ion a t an applied potential of 150 mv. A simple versatile circuit has been constructed that allows measurements of DVANTAGES
1 Present address, Department of Chemistry, Massachusetts Institub of Technology, Cambridge 39, Mass.
42
e
ANALYTICAL CHEMISTRY
limiting currents from quantities of cyanide in the l0-1D-gram range. Use of a silver anode was suggested by potentiometric studies (6, 7) involving the detection and estimation of small amounts of cyanide. EXPERIMENTAL
Apparatus. Amperometric measurements were performed using the circuit shown in Figure 1. Two 1.5volt dry cells were connected in parallel for compensating current, which may be used to cancel out unusually large residual or limiting currents, thus setting the background current a t zero if desired. Compensating currents and applied potentials were regulated through Helipot multiturn potentiometers (Beckman Instruments, Inc.). A 1-megohm variable resistor was placed in series with the galvanometer when the potential was first turned on to prevent accidental damage to the galvanometer from unexpectedly large currents. Currents were measured with a Leeds & Northrup ballistic galvanometer with a sensitivity of 0.0090 pa. per cm. and a 25-cm. scale on either side of zero. Potentials applied to the electrodes were monitored with a Keithley Model 200B vacuum tube voltmeter. Difficulty was encountered in obtaining a sufficiently tight seal of silver wire in glass, so silver anodes were prepared by silver plating 20-gage platinum wire
a t a current density of 3 ma. per sq. cm. for 15 to 20 minutes. After plating, the anodes exhibit unusually high and erratic residual currents. Their behavior will usual1 return to normal after polarbation L r 15 to 20 minutes a t 150 mv. in 0.03M NaOH. The anode used for most of the present work consisted of a 13-cm. length of wire of 3.30-sq. cm. apparent surface area coiled into a helix 0.8-cm. in diameter and 0.8-cm. high. After several weeks of continual use the anode became slightly discolored, but maintained original sensitivity for months. A hollow-spindle synchronous motor was used to rotate the anode. The cathode was a 2.30-sq. cm. platinum foil. With time, oxide formation on the platinum cathode may result in a sudden decrease in sensitivity. Original sensitivity may be maintained for several weeks when the equipment is operated in 0.03M base, but in supporting electrolyte of 0.1M base or greater, the drop may occur after several hours of continual use. Original sensitivity may be restored by treating the electrode in 1:1 ”03 for 30 minutes or longer, then stripping the surface clean by shorting against a calomel cell in de-aerated 0.1M HClOl until the current decays to a constant value. Electrical contact with both electrodes was made in the usual way through a mercury column. When not in use the electrodes were stored in distilled water. A 17-ml. weighing bottle was a convenient cell. Considerably larger cells may be used, as long as the rotating anode is able to mix the sample rapidly into the supporting electrolyte. Samples smaller than 0.25 ml. were added from a syringe-type microburet (No. 263 NRL, L. S. Starrett Go., Athol, Mass.). Reagents and Solutions. Stock solutions consisted of 4 X lO-3M potassium cyanide (Merck reagent grade), from which dilutions xere made for working standards. Stock solutions were stable for several months, but tenfold dilutions of the stock maintained original sensitivity for about 1 week. Cyanide solutions more dilute than IO-SM in 0.lM base were unstable within an houi after preparation,
