1523
Anal. Chem. 1981, 53, 1523-1528
Adsorption of Sodium Mercaptoacetate on a Mercury Electrode Robert S. Lawton anti Alexander
M. Yacynych”
Department of Chemistty, Rutgers, The State University of New Jersey, New Brunswick, New Jersey 08903
The adsorption of eleclroactlve species at a mercury drop electrode has been investigated for the sodium salt of mercaptoacetic acid. Eledrocapillary curves, linear sweep voltammetry, and cycllc voltammetry were employed in order to study the extent of adsorption and the surface coverage of the electrode. I t has been found that both mercury(I1) and mercury( I) mercaptidr are weakly adsorbed, while the mercaptoacetic anlon I!3 not adsorbed at all.
Mercaptoacetic acid (thioglycollic acid) is an important industrial material with applications in electroplating, plasticizing, and the preparation of thioglycollate of ammonia, a compound used in the permanent waving of hair. A study of electrode adsorption for the sodium salt of mercaptoacetic acid is discussed in this paper. Previous work with mercaptoacetic acid has dealt with its polarography (1-3) and quantitative analysis utilizing cathodic stripping voltammetry (4,5). The cyclic voltammetry of sodium mercaptoacetate exhibits complex behavior due to the adsorption of electrochemical products at the mercury electrode. To date, there have been no studies which deal with this particular case. In the present paper, an investigation of the adsorption of both mercury(1) mercaptoacetate and mercury(I1) mercaptoacetate has been carried out by using cyclic voltammetry, linear sweep voltammetry, and electroaapillary data. EXPERIMENTAL SECTION Apparatus. All voltammetric experiments that were carried out at slow scan rates (less than 180 mV/s) were performed with a three-electrode potentiostat constructed in-house and were recorded on a Houston-Instruments X-Y recorder. All voltammetric experimentsthat were carried out at 4 V/s were performed with a Wenking potentiostat, Model 61RS, and were recorded with a Hewlett-Packard oscilloscope, Model 120B. The Wenking potentiostat was driven with a triangular wave generator constructed in-house. Plots of drop time vs. potential were obtained with a Heath dropping mercury electrodle (DME),Model EWA-19-6. The drop times were measured frono a trace of current vs. time, recorded on a Houston-Instruments strip-chart recorder, Series 5000. The hanging mercury drop electrode (HMDE) employed for all voltammetric experiments was a Metrohm E410. Both the DME and the HMDE were filled with triply distilled Bethlehem Instrument mercury. The size of the mercury drop (one-tenth revolution of micrometer head) which was used for each voltamcm2, metric experiment had an approximate area of 1.9 x The SCE reference electrode and the Ag-wise counterelectrode were isolated from the cell by porous frits. Materials. The sodium salt of mercaptoacetic acid wasobtained from both the Aldrich Chemical Co. and the Fisher Scientific Co. In each case, the reagent was used without further purification. It was discovered that the mercaptoacetate had a finite “shelf life”. This conclusion is based on the observation that a cathodic peak eventually appears at -0.55 V (vs. SCE) in the cyclic voltammograms (pH 5.7). A comparison with a cyclic voltammogram of dithioglycollic acid revealed that the cathodic peak at -0.55 V was due to the reduction of the disulfide, indicating that a small portion of the original mercaptoacetate had been oxidized by the atmosphere. Whenever this occurred, a fresh lot of sodium
mercaptoacetate was purchased and the old lot was discarded. Mercury(1) nitrate, reagent grade, was obtained from the G. Frederick Smith Co., and mercury(I1)nitrate, reagent grade, was obtained from the J.T. Baker Chemical Co. Dithioglycollicacid was obtained from the Aldrich Chemical Co. AU of these reagents were used without further purification. All other chemicals used were reagent grade. Also, all experiments were run in a potassium hydrogen phthalate (KHP) buffer of pH 5.7, and all solutionswere prepared with triply distilled water. Procedures. The cyclic voltammograms were recorded in the following manner: Steady-state experiments were done with switching potentials of +0.2 V. and -0.75 V. The scan rate employed was either 75 mV/s or 4 V/s. The only exception is the data in Figure 8, which was collected from steady-state voltammograms with various scan rates and with switching potentials of 0.0 V and -0.75 V. A 1-min,open-circuit,equilibration period was used before the start of all steady-state and linear sweep voltammograms. The first 40 s were done with stirring, the last 20 s were quiescent. The experiments with linear sweep voltammetry were performed with various scan rates and start potentials. Data for Figures 5 , 7 , 9 , and 10 were obtained in the following manner: Standard solutions of various concentrations were prepared in volumetric flasks with deaerated buffer. These solutions were prepared fresh for each calibration curve. This was done in order to minimize evaporation of electroactive species, which would result in undetermined changes in concentration. The electrochemical cell contained 100 mL of deaerated buffer at the start of each set of experiments. Aliquots from the standard solutions were pipetted into the cell, and the new concentration was calculated based on dilution considerations. This particular procedure was employed because it was less time-consumingthan preparing a 100-mLsolution of known concentrationfor each point on the calibration curve. RESULTS AND DISCUSSION Assignment of t h e Voltammetric Peaks. The voltammograms were characterized by four peaks: a set of anodic and cathodic peaks a t approximately +0.1 V and a second set of peaks a t approximately -0.35 V. The assignment of individual reactions to each of the waves has been done polarographically by several groups (1-3). On the basis of our own results, we agree with the conclusions of Kolthoff and his co-workers. The following sets of reactions have been postulated for the voltammetric peaks observed: At +0.1 V HgSCH,COO-
e Hg/s\
CH,
b-&
+
e-
(11
At -0.35 V
+ HSCH2COO-
+ H+ + e-
(2) In order to confirm the above postulated mechanisms, we performed several sets of experiments. In the fist experiment, cyclic voltammograms were taken with a solution which was 0.12 mM in sodium mercaptoacetate and 0.12 mM in mercury(I1) nitrate. A slight cloudiness was observed upon addition of the mercury(I1) nitrate solution to the solution of sodium mercaptoacetate. This cloudiness is attributed to the formation of slightly soluble mercury(I1) mercaptide. A 1-min equilibration, with open-circuit, was employed prior to the start of the first scan. Hg
HgSCH&OO-
0003-2700/81/0353-1523$01.25/00 1981 American Chemical Society
1524
ANALYTICAL CHEMISTRY, VOL. 53,NO. 9, AUGUST 1981
I
I
I
I
I
- 40
00 VOLTS
vs
I
- 80 SCE
Cyclic voltammograms of (A) a solution which Is 0.12 mM In both mercury(I1) nitrate and sodium mercaptoacetate and (B) a solution which is 0.12 mM in sodium mercaptoacetate alone. Scan rate 75 mV/s. Flgure 1.
The data in Figure 1A demonstrate a significant difference between the first and second scans for the cathodic peak at +0.1 V. The tremendous enhancement of the first scan is attributed to the adsorption of mercury(I1) mercaptide during the equilibration period, and its subsequent reduction at +0.1 V. Referring to Figure lB, this notion is confirmed by the cyclic voltammograms on a solution of sodium mercaptoacetate alone. With only the sodium mercaptoacetate in solution, there is no significant difference between the first and second scans. Cyclic voltammograms taken at different pH values also yielded confirmation of the postulated electrode reactions. Referring to Figure 2, the set of peaks at -0.35 V were observed to shift an average value of 63 mV per change in pH unit, indicating that one proton is reacting per molecule of mercaptide. This is consistent with eq 2. Also, the peaks at +0.1 V were found to be invariant with hydrogen ion concentration, indicating that eq 1 is also correct. The pH dependence of mercaptoacetic acid has been previously described by Liberti (6),who observed that both cathodic peaks are dependent on proton concentration below pH 4 and both peaks are independent of proton concentration above pH 10, while only the peak at -0.35 V has any pH dependence between these two values. An inspection of the dissociation constants obtained by Kolthoff for mercaptoacetic acid (1) reveals that the observed change in pH dependence stems from the fact that the mercaptoacetic acid is largely undissociated at high proton concentration, while it exists as a monovalent anion above pH 4 and as a divalent anion above pH 10. Weak Adsorption vs. Strong Adsorption. There are nine different possible cases of adsorption or nonadsorption which could characterize a voltammetric peak. These different cases are outlined below: R = reactant, P = product, NA = not adsorbed, WA = weakly adsorbed, SA = strongly adsorbed; (1)RNA / PNA, (2) RWA / PNA, (3) RNA / PWA, (4) RSA ' / PNA, ( 5 ) RSA / PWA, (6) RNA / PSA, (7) RWA / PSA, (8) RSA / PSA, (9) RWA / PWA. There exist certain criteria to differentiate between strong and weak adsorption. In the case of strong adsorption of a product, the voltammogram will reveal a prepeak on the cathodic scan and a postpeak on tha anodic scan. The opposite will be true if a reactant is adsorbed. This phenomenon is caused by the difference in free energy between a molecule which is strongly adsorbed and a molecule which is not. A weakly adsorbed molecule undergoes electron transfer at
+
0.1 0
-0.70
-0.30 VOLTS
vs.
SCE
Steady-state cyclic voltammograms of three solutions at different pH values. Concentration of sodlum mercaptoacetate was 87 ppm for each solution. Scan rate = 75 mV/s.
Flgure 2.
nearly the same potential as the diffusing molecule, so that there are no prepeaks or postpeaks. However, the weakly adsorbed molecule will still cause enhancement of the peak current, as compared to strictly diffusion-controlled conditions (7). The exception to the above statements is the case of equally strong adsorption of both product and reactant, where the voltammogram will reveal neither a postpeak nor a prepeak. This is because both species will have the same free energy of adsorption. The prepeaks and postpeaks are eliminated because the stabilization due to adsorption is equivalent for each species, and the product molecules, once formed, can displace the adsorbed reactant molecules from the electrode surface (8). Therefore, the electrode is covered at all times and the ratio of the surface concentrations of reduced and oxidized species is Nernstian (7). The net effect of this is that cases 8 and 9 can become difficult to distinguish. In the particular case of the mercaptoacetate system, at a pH of 5.7, no prepeaks or postpeaks were observed for either of the redox couples. Therefore, either of the redox couples could be described by any of the cases 1-3 or 9. The instance of equally strong adsorption of both product and reactant, case 8, would be possible only if all three electroactive species had the same free energy of adsorption. Otherwise, one of the two redox couples would have had a postpeak and a prepeak. Qualitative Inspection of Cyclic Voltammograms. A qualitative examination of cyclic voltammograms,performed at various concentrations and scan rates, has proved informative with the mercaptoacetate system. Referring to Figure 3A,B, it was observed that the cathodic peaks at +0.1 V and -0.35 V have the symmetrical appearance which is indicative of an adsorption-controlled process (7).It was also observed that these cathodic peaks exhibit peculiar behavior at sufficiently high concentrations of sodium mercaptoacetate. Referring to Figure 3C, in the voltammogram performed at 313 ppm, it was observed that the cathodic peak current at +0.1 V had actually decreased. In addition, a cathodic current spike has appeared at -0.40 V. This current spike does not appear on the first reduction scan so that it can only be studied with steady-state cyclic voltammograms. This spike rapidly achieves a constant peak current which does not change with further increases in concentration. In contrast, the cathodic
ANALYTICAL CHEMISTRY, VOL. 53,NO. 9, AUGUST 1981
1525
t-
z
w LL
a 3
0
- 0.4
0.0 VOLTS
- 0.6
- 0.2
+0.2
V3LTS
vs. h
vs.
SCE
Steady-state cyclic voltammograms of a 220 ppm solution of sodlum mercaptoacetate: (A) 9 mV/s, (B) 75 mV/s. Flgure 4.
*
SCE
n
-0.8
~ p / C o vs
LOG
#
Co
h
C
E, Q
\
m a
I v-
0.250 -
I
0.125 -
.-P
t
Y
3
0
00
1.0
I
1 + 0.2
I
vs.
mV/s. i p vs. v
112
1
SCE
7.0
Steady-state cyclic voltammograms at various concentrations of sodium mercaptaacetiate: (A) 7 ppm, (B) 99 ppm, (C) 313 ppm, (D) 588 ppm. Scan rate = 75 mV/s. Flgure 3.
n
peak current at -0.35 V continues to grow with further increases in concentration. ,4t 588 ppm (Figure 3D),the current spike has been engulfed by the peak at -0.35 V. In addition, the voltammogram at 588 ppm reveals that the cathodic peak at +0.1 V has actually become inverted (anodic current on a cathodic scan). It was also observed that slower scan rates favor the inverted peak at +0.1 V. Referring;to Figure 4,it can be seen that the peak in question is inverted at a scan rate of 9 mV/s, while it becomes positive at a wan rate of 75 mV/s. All of the above data indicate that adsorbed species have a profound effect on the electrochemistry of sodium mercaptoacetate a t a mercury electrode. A more complete discussion of these qualitative observations will be presented later in this paper, after the elucidation of the electrode mechanisms. Dependence of Peak Height on Concentration and Scan Rate. A diffusion-controlled process will demonstrate a linear dependence both on bulk concentration, Co*, and on the square root of the scan rate, VI2.On the other hand, an adsorption-controlled process will demonstrate a nonlinear dependence on the bulk concentration and a linear depen-
3 0
Figure 5. Ratio of peak current to log bulk concentration as a function of the bulk concentration for the cathodic peak at -0.35 V. All points were determined from steady-state experiments. Scan rate was 75
- 0.6
-0 2 VOLTS
-
I
2.0
CO N C E NT R A T IO N
LOG
2
v
Q
._
4.0
1 I I
t
1.0
J
i
J 5
( SCA N
10
15
R A T E )1’2
Peak current as a function of the square root of the scan rate (mV/s) for the cathodic peak at -0.35 V. Concentration of sodium mercaptoacetate was 550 ppb. Flgure 6.
dence on the scan rate, V (7). Plots of the peak current dependence on concentration and scan rate were prepared for the mercaptoacetate system. A plot of i,/Co* (peak current/bulk concentration) vs. log Co* was prepared for the cathodic peak at -0.35 V (Figure 5). The above data yielded a reverse S-shaped curve, corresponding to the transition from Henry’s law behavior to a diffusion-controlled dominated situation. Plots of peak current vs. VI2were prepared for the cathodic peak at -0.35 V, using linear sweep voltammetry with a start
1526
ANALYTICAL CHEMISTRY, VOL. 53, NO. 9, AUGUST 1981 vs. co*
iD
0
-
O.O1
i 150
450
COYCENTRATION
(ppm)
Figure 7. Ratio of peak current to bulk concentration as a function of bulk concentration for the anodic peak at +O. 1 V (A)and the anodic peak at -0.35 V (D). Scan rate = 75 mVls. Data points measured from steady-state cyclic voltammograms.
1.0
I
I
I
I
150
300
450
I
600
CONCENTRATION
(
ppm)
Peak current as a function of bulk concentration for the cathodic peak at 4-0.1 V. Scan rate = 4 V/s. Steady-state experiment. Flgure 9.
I p vs. c o t
c 0.0
-5.0
5 (SCAN
10
Flgure 8. Ratio of anodic peak current to cathodic peak current as a function of the square root of the scan rate (mV/s) for the redox couple at -0.35 V. Concentration of sodium mercaptoacetate was 7 ppm. Switching potentials were 0.00 V and -0.75 V.
potential of 0.00 V. Referring to Figure 6, the nonlinear behavior at 550 ppb is indicative of an adsorption-controlled process. The same type of plot at 220 ppm is linear, demonstrating that the cathodic process at -0.35 V has become diffusion controlled. These data correlate nicely with the graph in Figure 5 , further substantiating the postulated adsorption of mercury(1) mercaptide. In the case of the anodic peaks at +0.1 V and -0.35 V, plots of ip/Co* vs. Co* were prepared. The data for both peaks are graphed on @e same axis in Figure 7. The nonlinear behavior of the peak,current at +0.1 V is expected, since this peak corresponds to the oxidation of adsorbed mercury(I1) mercaptide. The linear behaviok of the anodic peak at -0.35 V indicates that the mercaptoacetic anion is not adsorbed. Scan rate experiments were performed at 7 ppm in order to confiim this conclusion. Figure 8 is a plot of the ratio of the anodic peak current to the cathodic peak current vs. the square root of the scan rate, for the redox couple at -0.35 V. The shape of the graph clearly demonstrates the weak adsorption of mercury(1) mercaptide and the lack of adsorption of the mercaptoacetic anion (7). In the case of the cathodic peak at +0.1 V, a plot of peak current vs. scan rate yielded a linear relationship at 100 ppm, indicating that the mercury(I1) mercaptide is adsorbed. This type of behavior is consistent with the symmetrical appearance of the cathodic peak at +0.1 V. Calculation of Surface Coverage. Once having demonstrated that the mercury(1) and mercury(I1) mercaptides were adsorbed, a surface coverage of the electrode was calculated for the mercury(I1) mercaptide. A calibration curve was constructed at a scan rate of 4 V/s with steady-state cyclic
1
150
I
I
300
CONCENTRATION
15
RATE)^'^
L
(pprn)
Peak current as a function of bulk concentration for the cathodic peak at +0.1 V. Scan rate = 75 mV/s. Steady-state cyclic voltammograms. Figure 10.
voltammograms, using the oscilloscope as a recorder. The data are plotted in Figure 9. The constant peak current attained at high concentrations corresponds to monolayer coverage of the mercury electrode (9). At all concentrations investigated, the cathodic peak at +0.1 V retains its symmetrical appearance, indicating that diffusion contributes only a minor portion of the total current. The negligible contribution of diffusing species is responsible for the plateau exhibited by the Calibration curve at high concentrations. Integration of the peak area for the reduction at +0.1 V, using a 220 ppm solution, yields a surface coverage of approximately 18-22 A2 per molecule. This is in good agreement with an end-on orientation of the mercury(I1) mercaptide in which the plane of the ring is perpendicular to the surface of the electrode. In the case of the cathodic peak at -0.35 V, the calibration curve does not exhibit a flat plateau. Also, unlike the symmetrical peak a t +0.1 V, the cathodic peak at -0.35 V exhibits tailing at a scan rate of 4 V/s. These observations indicate that diffusion is still making a considerable contribution to the reduction current at -0.35 V, even though the scan rate is 4 V/s. Therefore, a surface coverage was not calculated for the mercury(1) mercaptide. In the case of the cathodic peak at +0.1 V, the calibration curve was quite different when constructed at a scan rate of 75 mV/s (Figure 1)as compared to the curve constructed at 4 V/s (Figure 9). Referring to Figure 10, it was observed that the peak current achieves a maximum value at approximately 80 ppm, before falling off with further increases in concentration. Integration of the peak with the greatest height in Figure 10 yields the same amount of charge per unit area as was obtained for monolayer coverage of the electrode. This indicates that the anomalous behavior at +0.1 V (i.e., decreasing peak current with increasing concentration) requires
ANALYTICAL CHEMISTRY, VOL. 53,
-00
-08
-04 VOLTS
v5
-12
SCE
Flgure 11. Drop time as EL function of potential, recorded at several
concentrations of sodium mercaptoacetate: (A) supporting electrolyte, (B) 10 ppm, (C)1050 ppm, (D) 2350 ppm. the prior formation of monolayer coverage of the electrode with mercury(I1) mercaptide. A comparison of Figure 9 with Figure 10 reveals that the inverted peak is favored by high bulk concentrations of sodium mercaptoacetate and slow scan rates. This is entirely consistent with the observations that were made from Figures 3 and 4. At this point, it should be noted that chronocoulometry was considered by the authors as an alternative method for calculating surface coverage of the electrode. In the sodium mercaptoacetate system, however, the adsorbed species exist only within those potential windows that correspond to electrochemical reactions. This violates the boundary conditions of the chronocoulometrictechnique. Accordingly, the authors did not employ this technique. Drop Time vs. Potemtial. Plots of drop time vs. potential were made for the sodium mercaptoacetate system. Referring to Figure 11, the most significant features of these plots are the discontinuities at approximately +0.1 V and -0.35 V, reflecting sudden changes in the interfacial tension of the mercury drop. Similar behavior has also been observed for adsorbed cadmium com,plexes (10). Referring to the discontinuity at +0.1 V, it was observed that this phenomenon is not apparent at 10 ppm, while it has become very prominent a t 1050 ppm. It is interesting to note that this discontinuity appears at the same potential and within the same range of concentration as the inverted peak in the cyclic voltammograms that were performed at 75 mV/s. The significance of this will be discussed in the next section. Discussion o f Qualitative Observations. The results now allow a much more complete explanation of the data presented in Figures 3 and 4. Referring to the cathodic current spike observed at -0.40 ’V in Figure 3, we have attributed this phenomenon to a film rearrangement of adsorbed mercury(1) mercaptide. Bard (11) has observed a similar phenomenon for cysteine, which he attributes to the formation of a “compact” film of adsorbed species, caused by the attractive interactions between the adsorbed molecules. In view of the above explanation, it is not surprising that the current spike rapidly attains a maximum peak height that does not change with further increases in concentration. This would be expected since the “compact” film should eventually achieve maximum electrode coverage. We have also been able to explain the inverted peak observed in Figure 4. To begin with, the inverted peak represents an oxidation current on a cathodic scan. It has been demonstrated that this oxidation current is favored by slow scan rates and high bulk concentrations (Figures 3 and 4). These data indicate that there iis competition between the reduction of one adsorbed species and the oxidation of a second, diffusing species. Therefore, the peak at +0.1 on the cathodic
NO.9,
1527
AUGUST 1981
sweep is actually the sum of two different processes. This interpretation of the data is consistent with the graph in Figure 10. The downward sloping portion of the plot is due to the increasing contribution of the diffusion-controlled oxidation current at higher bulk concentrtions of sodium mercaptoacetate. In order to explain the oxidation of diffusion-supplied mercury(1) mercaptide on a cathodic sweep, it is necessary to refer to the electrocapillary curves. In Figure 11,the data demonstrate that the dissolution of the film of mercury(I1) mercaptide, and the subsequent formation of a layer of mercury(1) mercaptide, is accompanied by a dramatic change in the surface tension of the mercury drop. More importantly, a comparison of the cyclic voltammograms (75 mV/s) and Figure 11demonstrates that the inverted peak occurs at the same potential and concentration as the discontinuity in the electrocapillary curves. In other words, the oxidation of mercury(1) mercaptide on the cathodic sweep does not begin until the dissolution of the layer of mercury(I1) mercaptide. The combination of all of the above data leads to a possible explanation of the inverted peak. The most logical conclusion is that the adsorbed layer of mercury(I1) mercaptide, when it attains monolayer converage, hinders the electron transfer of nearby, mercury(1) mercaptide molecules. This results in a localized excess of mercury(1) mercaptide, as well as a scarcity of mercury(I1) mercaptide, in the proximity of the electrode. The electrode itself is covered with an adsorbed layer of mercury(I1) mercaptide. When the integrity of the mercury(I1) mercaptide layer is destroyed by reduction on the return cathodic sweep, the nearby mercury(1) species begins to transfer electrons as the adsorbed mercury(I1) species is being reduced. The resulting current is then a sum of the two processes. The above postulated mechanism also explains another aspect of the inverted peak. At approximately +0.06 V, immediately following the inverted peak, there is a small, positive peak. With regard to the above postulated mechanism, this small amount of reduction current can be attributed to the reduction of “extra” mercury(I1) mercaptide that was formed by the oxidation of mercury(1) mercaptide on the cathodic sweep. SUMMARY Referring to our earlier discussion of the different possible cases of adsorption which might characterize a voltammetric peak, the data indicate that the redox couple at -0.35 V is described by case 2 (RWA/PNA), while the redox couple at +0.1 V is described by case 9 (RWA/PWA). This work also demonstrates that the mercury(I1) mercaptide is adsorbed onto the electrode surface with an end-on orientation. The various mechanisms which have been elucidated are as follows: E = +0.1 V; Hg’SR (adsorbed) + HgI’SR (adsorbed) leE = -0.35 RSH (solution) + Hg’SR (adsorbed) le- 1H+
+
v;
+
+
LITERATURE CITED (1) Leussing, D. L.; Kolthoff, I. M. J . Nectrochem. SOC. 1953, 700, 334-338. (2) Libertl, A.; Cervone, E. Ann. Chim. (Rome) 1951, 4 1 , 95-104; Chem. Abstr. 1951, 45, 10100e. (3) Gorokhovskaya, V. I.; Klimova, L. 1. Zh. Obshch. Khim. 1965, 35 ( l l ) , 1913-1916; Chem. Abstr. 1966, 6092 g. (4) Florence, T. M. J . Electroanal. Chem. 1970, 97. 219-236. (5) Moore, W. Michael; Gaylor, V. Frances Anal. Chem. 1977, 49 (9), 1386- 1390. (6) Sartori, G.; Liberti, A.; Colazolarl, C. Com. Int. Thermodyn. Chef. Electrochim., C . R . Reun. 2nd301-304; Chem. Abstr. 1952, 46, 10011g. (7) Wopschall, R. H.; Shain, I. Anal. Chem. 1967, 39, 1514-1527. (8) Heyrovsky, J.; Kuta, J. “Principles of Polarography”,Academic Press: New York, 1966; p 293. (9) Schmidt, C. L.; Swofford, H. S. Anal. Chem. 1979, 51, 2028-2033.
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Anal. Chem. 1901, 53, 1528-1530
(IO) Kalvoda, R.;Anstine, W.; Heyrovsky, M. Anal. Chlm. Acta 1970, 50, 93-102. (11) Stankovich, Marian T.; Bard, Allen J. J. Electroanal. Chem. 1977, 75, 487-505.
RECEIVED for review June 9, 1980. Resubmitted January 23,
1981. Accepted April 17, 1981. A M Y . thanks Rutgers Research BiomedicalResearch Grants, and the National Institutes of Health (Grant No. GM 28125-01) for research support. R.S.L. thanks Rutgers University for the award of a University Fellowship.
Stopped-Rotat ion Volta mmet ry Joseph Wang Department of Chemistry, New Mexico State Universiiy, Las Cruces, New Mexico 88003
The technique of stopped-rotation voltammetry at a rotated disk electrode, whlch measures the current difference with the rotation speed switched on and off, Is described. Sensitlvlty, preclslon, mass-transport properties, and linearity of response are reported. Well-defined current-potential curves are obtalnable and compared with those obtalned by dlfferentlal pulse voltammetry. An asymmetric hydrodynamic modulatlon waveform and a relatively rapld stopped-rotation procedure (wlthout the achievement of steady states) are employed to shorten the analytlcal cycle. Due to its inherent sensitivlty and slmpllclty, the technique seems well suited for the measurement of low concentrations of electroactlve materials.
In recent years, the use of solid electrodes for electroanalytical purposes has gained popularity, one of the primary reasons being their applicability to anodic oxidations. However, practical difficulties have hampered the attainment of precise quantitative data employing classical or potential pulse voltammetry at solid electrodes. Hydrodynamic modulation at solid electrodes has been shown to be a feasible technique for obtaining reproducible current-voltage data a t low concentrations of electroactive species. Such modulation involves measurement of the current difference between two rates of convective transport, resulting with efficient discrimination against the nonconvective background current. Miller, Bruckenstein, and co-workers have done a great deal of work in developing the sinusoidal modulation of a disk electrode's rotation speed about a center value (1, 2). The analytical usefulness of this approach has been exploited for obtaining voltammograms a t submicromolar concentrations. The sinusoidal component of the current, which follows directly from the Levich equation i s given by
where n is the number of electrons transferred per molecule, F is the Faraday, A is the disk area, D is diffusion coefficient, v is kinematic viscosity, Cb is the bulk concentration of the electroactive species, Am is the peak to peak amplitude of the change in the rotation speed, and wo is the center value of the rotation speed. Aw is always much smaller than 00 (usually about 1-10% of coo), resulting with exploitation of only a very small portion of the steady-state analytical current of interest. In order to decrease the attenuation of the steady-state analytical response (i.e., increased current difference), Blaedel and Engstrom have suggested the pulsed-rotation voltammetry
(PRV) in which the rotation speed of a rotated disk electrode is switched between two values (3)
The subscripts H and L designate the high and the low rotation speeds, respectively. Since low and high rotation speeds of 500 and 1500 rpm, respectively, are usually selected as the two current states, more than half the steady-state analytical current is not exploited for the quantitation (eq 2). Further increased current difference has been achieved by incorporating the PRV approach with the high surface area of a rotated porous disk electrode (4). Analogous procedures have been developed for electrochemical flow detectors, with the solution flow rate being pulsed to the surface of a stationary electrode (5,6),or with hydrodynamically modulated rotating disk electrode in a flow cell (7). The purpose of the following work is to demonstrate the characteristics of stopped-rotation voltammetry (SRV), in which differential current measurements are made between zero and high rotation speeds. The main advantages of this procedure are its inherent sensitivity (Le,, exploitation of all the rotation-dependent analytical signal of interest for the quantitation (100% modulation)) and the simplicity of operation (turning on and off the rotational speed). Since the lower current (at w = 0) is not zero but a very small value, due to diffusion (linear and radial) and natural convection (8),an expression for the on-off current is not obtained simply by substituting wL = 0 into eq 2; instead, and because of the theoretical limitations to evaluate the zero rotation speed current (8), Ai may be described by the following relationship:
Ai = 0.62nFAD2/3v~1~6Cbco1/2 - L* = O
(3)
Compared to previously reported hydrodynamic modulated rotating disk electrode approaches (discussed above) the SRV does not require any programming circuity for changing the rotation speed, and thus it can be incorporated with every rotating disk assembly. These advantages and the various characteristics of the SRV are elucidated by application to the measurement of micro- and submicromolar concentrations of ascorbic acid and ferrocyanide.
EXPERIMENTAL SECTION Apparatus. The rotating electrode assembly (Model PIR, Pine Instruments Co., Grove City, PA) with a 0.75-cm diameter glassy
carbon disk electrode (ModelDDI 15, Pine Instruments Co.) was used in conjuction with a cell of 200-mL capacity made from Pyrex glass. The salt bridges of the reference electrode (Ag/AgCl,Model RE-1, Bioanalytical Systems Inc., West Lafayette, IN) and of the counterelectrode(a Pt coil immersed in 0.1 M phosphate buffer) join the cell through two holes in its Plexiglas cover. The working
0003-2700/81/0353-1528$01.25/00 1981 American Chemical Society
